Seawater Batteries: Principles, Materials and Technology (Green Energy and Technology) 981190796X, 9789811907968

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Green Energy and Technology

Youngsik Kim Wang-geun Lee

Seawater Batteries Principles, Materials and Technology

Green Energy and Technology

Climate change, environmental impact and the limited natural resources urge scientific research and novel technical solutions. The monograph series Green Energy and Technology serves as a publishing platform for scientific and technological approaches to “green”—i.e. environmentally friendly and sustainable—technologies. While a focus lies on energy and power supply, it also covers “green” solutions in industrial engineering and engineering design. Green Energy and Technology addresses researchers, advanced students, technical consultants as well as decision makers in industries and politics. Hence, the level of presentation spans from instructional to highly technical. **Indexed in Scopus**. **Indexed in Ei Compendex**.

More information about this series at https://link.springer.com/bookseries/8059

Youngsik Kim · Wang-geun Lee

Seawater Batteries Principles, Materials and Technology

Youngsik Kim Seawater Resources Technology (SRT) Center, School of Energy and Chemical Engineering Ulsan National Institute of Science and Technology (UNIST) Ulsan, Korea (Republic of)

Wang-geun Lee Seawater Resources Technology (SRT) Center, School of Energy and Chemical Engineering Ulsan National Institute of Science and Technology (UNIST) Ulsan, Korea (Republic of)

4TOONE Corporation Ulsan, Korea (Republic of)

ISSN 1865-3529 ISSN 1865-3537 (electronic) Green Energy and Technology ISBN 978-981-19-0796-8 ISBN 978-981-19-0797-5 (eBook) https://doi.org/10.1007/978-981-19-0797-5 © The Editor(s) (if applicable) and The Author(s), under exclusive license to Springer Nature Singapore Pte Ltd. 2022 This work is subject to copyright. All rights are solely and exclusively licensed by the Publisher, whether the whole or part of the material is concerned, specifically the rights of translation, reprinting, reuse of illustrations, recitation, broadcasting, reproduction on microfilms or in any other physical way, and transmission or information storage and retrieval, electronic adaptation, computer software, or by similar or dissimilar methodology now known or hereafter developed. The use of general descriptive names, registered names, trademarks, service marks, etc. in this publication does not imply, even in the absence of a specific statement, that such names are exempt from the relevant protective laws and regulations and therefore free for general use. The publisher, the authors and the editors are safe to assume that the advice and information in this book are believed to be true and accurate at the date of publication. Neither the publisher nor the authors or the editors give a warranty, expressed or implied, with respect to the material contained herein or for any errors or omissions that may have been made. The publisher remains neutral with regard to jurisdictional claims in published maps and institutional affiliations. This Springer imprint is published by the registered company Springer Nature Singapore Pte Ltd. The registered company address is: 152 Beach Road, #21-01/04 Gateway East, Singapore 189721, Singapore

Preface

Batteries are a significant source of energy in modern life, with applications ranging from small devices to electronic equipment, electric vehicles, auxiliary power devices, and energy storage systems. This scalability was enabled in particular by the development of secondary batteries, the performance of which has been continuously enhanced over decades to make them capable of storing and supplying higher power. However, such a rapid growth in battery demand inevitably raises worries about the sustainability of energy resources, particularly lithium and metals. This is why, despite the enormous success of lithium-ion batteries, research is being conducted on next-generation batteries that use less expensive and more ecofriendly post-lithium materials. Sodium is perceived as a sustainable energy storage medium due to its accessibility. In this perspective, a seawater battery that stores energy utilizing sodium ions in seawater is one of the most promising technologies as it directly exploits sodium ions included inside the semi-infinite seawater, the world’s most plentiful resource. Additionally, it is acknowledged as one of the most appropriate batteries for use in a variety of marine devices, machines, and ships that operate in a marine environment. Unlike other batteries, however, the rigorous and systematic research and development necessary to implement the seawater battery was initiated until lately. This is attributed to the irony that abundant and environmentally friendly seawater is difficult to handle mechanically and electrically. This necessarily implies research into seawater-resistant materials, robust energy storage/production architecture, and detailed charging/discharging controls. Due to these difficulties, seawater battery development has concentrated on primary batteries for specialized applications such as military or communication during the last few decades. The rechargeable secondary seawater battery was just recently demonstrated and produced. Secondary seawater batteries, as an environmentally friendly and sustainable sodium-based power supply device, are expected to be used in a wide range of applications, not just in maritime equipment and devices, but also in coastal and inland areas. The purpose of this book is to provide an overview of these seawater battery technologies. Accordingly, this book introduces the characteristics of seawater as an energy resource, and then deals with the materials, principles and technology v

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of the first developed primary seawater battery, and the recently researched and developed secondary seawater battery. In Chapter 1, we tried to describe the value of seawater as the next-generation resource and energy materials, as long as a circulation medium of the earth’s energy and resources system. In Chapter 2, we tried to help the understanding of the primary seawater battery for special purposes by examining the development process of the primary seawater battery and examining various materials, operating principles, and cell design. In Chapter 3, the characteristics and operating principles of each part of the recently implemented secondary seawater battery, as well as the latest research on cell development, were to be explained in relatively detail. Through this book, it is expected that readers would get a general understanding of seawater batteries, but also develop an integrated understanding of materials exploration, cell design, and electrochemical in research and development of batteries. Furthermore, we hope that this book will not only provide clues to the energy resource problems facing mankind, but also contribute to suggesting technological directions for a more prosperous sustainable society. Ulsan, Korea (Republic of)

Youngsik Kim Wang-geun Lee

Acknowledgments

We’d like to express our gratitude to YK group members, Wooseok Go, Jongwoo Kim, Jihun Cho, Youngjae Jung, Seyoung Lee, Dowan Kim, Jeongwoo Park, Haeun Kim, Kwangho Shin and Dongyeop Kim, who assisted us greatly with data collection, content organization, and figure correction. We thank you also to the rest of the group members for their dedicated research and collaboration. This work was funded by the 2021 Research Fund (1.210051.01) of the Ulsan National Institute of Science and Technology (UNIST) and the Ministry of Trade, Industry & Energy, Republic of Korea (20215610100040).

vii

Contents

1 Seawater and Its Resources . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.1 Sea and the Earth . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.1.1 Sea and the Earth Ecosystem . . . . . . . . . . . . . . . . . . . . . . . . . . 1.1.2 Circulation of Seawater . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.1.3 Seawater and Mankind . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.2 Characteristics of Seawater . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.2.1 Components of Seawater . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.2.2 Physical Properties of Seawater . . . . . . . . . . . . . . . . . . . . . . . . 1.2.3 Chemical Properties of Seawater . . . . . . . . . . . . . . . . . . . . . . . 1.3 Seawater Resource Technologies . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.3.1 Fresh Water (Desalination) . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.3.2 Mineral Mining . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.3.3 Hydrogen (H2 ) Production . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.3.4 Energy Production . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.3.5 Seawater Electrochemical Energy Systems . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

1 1 1 3 8 10 10 12 17 24 25 26 28 28 29 31

2 Primary Seawater Batteries . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.1.1 Brief History of Batteries . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.1.2 Types of Batteries . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.1.3 The Birth of Primary Seawater Batteries . . . . . . . . . . . . . . . . . 2.2 Operating Principles . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.2.1 General Principles . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.2.2 Reactions in a Primary Seawater Battery . . . . . . . . . . . . . . . . 2.2.3 Features of Primary Seawater Batteries . . . . . . . . . . . . . . . . . . 2.3 Materials . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.3.1 Magnesium Battery . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.3.2 Zinc Battery . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.3.3 Aluminum Battery . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.3.4 Li Battery . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

37 37 37 39 43 50 50 52 53 56 60 67 70 72

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2.4 Cell Engineering . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.4.1 General Construction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.4.2 Types of Primary Seawater Batteries . . . . . . . . . . . . . . . . . . . . 2.4.3 Applications of Primary Seawater Batteries . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

73 73 74 77 83

3 Secondary Seawater Batteries . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.1 Introduction and Operating Principle . . . . . . . . . . . . . . . . . . . . . . . . . . 3.1.1 Birth of the Secondary Seawater Battery . . . . . . . . . . . . . . . . . 3.1.2 Operating Principle and Features . . . . . . . . . . . . . . . . . . . . . . . 3.1.3 Structure and Components . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.2 Cathode . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.2.1 Reactions Occurring at SWB Cathodes . . . . . . . . . . . . . . . . . . 3.2.2 OER/ORR . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.2.3 Cl− Capture/Release . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.2.4 Na+ Intercalation/Deintercalation . . . . . . . . . . . . . . . . . . . . . . . 3.2.5 Redox Reactions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.3 Anodes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.3.1 Anode Parts of SWBs: Roles and Requirements . . . . . . . . . . 3.3.2 Sodium Metal . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.3.3 Na Intercalation in Hard Carbon . . . . . . . . . . . . . . . . . . . . . . . . 3.3.4 Alloy Materials . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.3.5 Liquid Electrolytes and Anolytes . . . . . . . . . . . . . . . . . . . . . . . 3.4 Solid Electrolytes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.4.1 Role and Requirement of Solid Electrolyte for SWBs . . . . . 3.4.2 Solid Electrolyte Materials . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.4.3 Sodium Beta Alumina . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.4.4 NASICON . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.5 Cell Engineering . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.5.1 From Proof-of-Concept to Cell . . . . . . . . . . . . . . . . . . . . . . . . . 3.5.2 Coin Cell . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.5.3 Prismatic Cells . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.5.4 Module Fabrication . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.5.5 Pilot System for Square Cell Fabrication . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

91 91 91 94 98 100 100 102 141 145 160 171 171 174 185 196 209 221 221 223 227 233 250 250 254 258 267 273 275

Chapter 1

Seawater and Its Resources

Abstract Water (H2 O) is the most abundant and renewable substance on Earth. Approximately 71% of the Earth’s surface is covered with water. Seawater accounts for 97.5% of the total available water, with the remaining 2.5% being fresh water. Seawater has an impact on the global environment and humanity through its constant circulation, consequently leading to the circulation of the Earth’s water, matter, and energy. Seawater contains approximately 3.5% salts, which confer several physical and chemical properties that differ from those of freshwater. As such, seawater comprises various minerals and presents with ionic conductivity. Thus, it can be used in various seawater resource technologies, such as freshwater production, mineral extraction, hydrogen production, and energy storage/conversion. Throughout this chapter, the characteristics and roles of seawater and its systems, the main characteristics of seawater, and particularly, the value of seawater as an abundant substance and energy resource are examined. In addition, electrochemical energy storage and production technology will be considered. This chapter aims to inspire collective creative imagination to help resolve the global energy and resource problems plaguing humanity.

1.1 Sea and the Earth 1.1.1 Sea and the Earth Ecosystem The pale blue dot. —Carl Edward Sagan. To the best of our knowledge, the Earth is a unique planet among the many planets in the universe. This tiny blue dot-like planet, Earth, which seems insignificant in the vast universe, supports a variety of lifeforms. Coincidentally, this blue planet contains water at an appropriate temperature and pressure to result in suitable atmospheric activity. One of the most important features that makes the Earth unique is the existence of the oceans. The Earth is a water-rich planet and is also called the ‘planet of water’ (Fig. 1.1). The Earth’s surface has a larger ocean surface © The Author(s), under exclusive license to Springer Nature Singapore Pte Ltd. 2022 Y. Kim and W. Lee, Seawater Batteries, Green Energy and Technology, https://doi.org/10.1007/978-981-19-0797-5_1

1

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1 Seawater and Its Resources

Fig. 1.1 Image of Earth from space

area (approximately 71% of the surface area) than the land portion [1, 2]. The total amount of water on Earth is estimated at 1.386 billion/km3 [3, 4]. Most of this water exists in the form of seawater. Of all the water on Earth, oceans contain approximately 97% of the water, which excludes glaciers, groundwater, lakes, soil moisture and atmospheric water [3, 5]. The water-rich ocean is an important storage area for energy and matter in the Earth’s energy and matter cycle, contributing to the absorption and release of solar heat and the circulation of water and salts. The water cycle passing through the ocean is called the hydrosphere, which interacts closely with the Earth’s other circulatory systems: the atmosphere, the geosphere, and the biosphere. In other words, the presence of the ocean is a key feature and element of the Earth system that sets it apart from other planets. Humanity has built an astonishing civilization over the history by exploiting the numerous resources and energy sources, particularly, since the Industrial Revolution. It is only now, in the age of the telecommunications and energy revolutions, that humanity has begun to consider the consumption of energy and resources and the associated environmental issues. From this perspective, the oceans and seawater could become major sources of energy and resources in the future. Thus, the answer to the current energy and resource crises may lie in the ocean.

1.1.1.1

Origin of Seawater

There are few clues as to when the Earth became the watery planet it is today. However, water has existed for a long time in the early Earth environment, and it is presumed that it played an important role before the formation of the present oceans and atmosphere. Geologically, fossils of primitive algae and invertebrates

1.1 Sea and the Earth

3

estimated to have been formed 500–600 million years ago and those of carbonate sedimentary rocks formed 1 billion years ago are believed to have formed in an aquatic environment [1, 6]. Precambrian fossils show that bacteria and blue-green algae existed approximately 3.3 billion years ago, indicating the prior existence of water on the planet [7]. There are two main hypotheses regarding the origin of the water on earth. One is that it has existed since the formation of the Earth, and the other is that it originated from outer space [6, 8–10]. The former theory claims that during the formation of the earth, gas was emitted from the high temperature environment within the Earth’s core. As a result, hydrogen and oxygen gases were converted at high temperatures into a large quantity of water vapor [9]. In contrast, the latter theory asserts that the continuous comet and asteroid collisions were responsible for introducing a large amount of water to the Earth. To verify the credibility of this scenario, various studies are being conducted, such as a comparison of the deuterium ratio in Earth’s water and that in the ice contained in an asteroid [10]. However, no clear evidence has been found to date.

1.1.1.2

Marine Earth Ecosystem

Oceans play an important role on Earth in the absorption and release of external energy. On the surface of the ocean, the water absorbs solar radiation and vaporizes. Vaporized water molecules are lighter than nitrogen and oxygen molecules in the atmosphere and hence rise upward, with a portion condensing to form clouds. The clouds formed produce rain that falls back to the earth’s surface in the form of freshwater. Through this process, a large amount of solar energy is transferred to the land surface. On a global scale, this has the effect of dispersing solar energy concentrated in lower latitudes to higher latitudes [11]. In addition to this energy cycle, as fresh water from the ocean flows over land, the ocean drives the Earth’s material cycles (Fig. 1.2). Approximately 413,000 km3 of water evaporates from the ocean annually, of which approximately 40,000 km3 of water falls back on land in the form of rain [12]. The rain water then flows through the soil and rocks on the surface, dissolving and carrying various chemical substances. This water continues to flow and eventually returns to the sea. This cycle is called the hydrological cycle and plays an important role in the energy and material cycle related to the Earth’s water cycle and related ecosystems [11, 13].

1.1.2 Circulation of Seawater Seawater is constantly moving. Generally, the global flow of seawater can be described as three types—waves, surface water circulation, and thermohaline circulation—according to the main force and action characteristics of the flow.

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Fig. 1.2 Global annual water cycle in the 1990s. The arrows indicate the schematic flow of water substances in various forms (Units: Thousand km3 for storage, and thousand km3/yr for exchanges) Reproduced with permission of Ref. [12]

1.1.2.1

Waves

Tides are periodic sea-level fluctuations caused by the motion of the Moon, Sun, and Earth [14]. As the moon orbits the Earth, it attracts the Earth and its water based on the law of gravitation. Tides are the most extended wave of all sea waves, with a wavelength equivalent to half the circumference of the Earth. Unlike the other waves, this gigantic shallow sea wave is classified as a forced wave because it is still under the influence of wave force. Moreover, the tidal difference is constantly changing. During a full moon, the Sun, the Earth, and the Moon are aligned such that the tidal force of the Sun is combined with that of the Moon, as shown in Fig. 1.3a. As the Earth and the Moon are closely located at its perigee, a greater force is applicable. Contrary to this high tide, during a half-moon, the Sun and Moon are positioned at 90° in relation to the Earth, and the tidal forces nullify each other, as shown in Fig. 1.3b. Generally, ocean tides can be explained by equilibrium tidal theory, which considers the equilibrium of forces as the planets orbit the Sun and the Moon in turn orbits the Earth [15, 16]. The equilibrium tidal theory is based on the assumption that the sea level immediately responds to the force that affects the deformation, that is, the sea level is always in equilibrium with a force exerted on it. According to this, a factor of a maximum of 55 cm by the Moon and a maximum of 24 cm

1.1 Sea and the Earth

5

Fig. 1.3 Illustration on the equilibrium tidal theory. The a portrays tides during the full moon where the tidal force of the Sun is combined with that of the Moon. The b portrays tides during the half-moon where these tidal forces nullify each other. Reproduced with permission of Ref. [14]

by the Sun is generated, which is very small compared to the average value of the Earth, which is approximately 2 m. This difference is attributed to the fact that the motion of the Sun and Moon is much faster than the movement of the sea, and the surface of the sea is not in perfect. Laplace proposed the dynamic theory of tides in 1775 by combining the concept of fluid motion with Newton’s discoveries in celestial mechanics. This theory explains the difference between the actual tides and that estimated by the equilibrium tide theory [16].

1.1.2.2

Surface Circulation

The circulation of surface seawater in the world’s oceans is closely related to the rotation of the Earth. The rotation of the Earth changes the acting centrifugal force that depends on the latitude, which causes differences in seawater flow in the northern and

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1 Seawater and Its Resources

southern hemispheres. The vertical surface component of the centrifugal force nullifies the effect of the Earth’s gravity. Although it has no effect, the horizontal surface component drives ocean currents toward lower latitudes. Fluids in the northern hemisphere always curve to the right, and vice versa in the southern hemisphere. The deflection effect of these fluids is called the Coriolis force or the Coriolis effect, after Coriolis [17]. This effect is more prominent at higher latitudes and does not exist at the equator. Seawater circulation is closely related to the air, another mixture of gases that makes up the Earth’s surface. Wind stress causes the water to move in the direction of the wind, and this force acts up to 50 m beneath the surface. The seawater layer affected by this force is called the Ekman layer [15, 18]. The balance between wind stress, Coriolis force, and frictional force causes the surface water to flow at an angle of 45° to the right of the wind direction. As the depth of the water increases, the flow velocity decreases, and the direction of the current is directed further to the right (Fig. 1.4a). The flow of surface seawater leads to the formation of an area where surface seawater gathers and one where it separates. In the convergence zone where seawater gathers, subsidence occurs, wherein the surface seawater sinks, whereas in the divergence zone, where the surface seawater separates, upwelling occurs wherein water from the lower part of the Ekman Formation rises to the surface, which then causes vertical movement of the seawater (Fig. 1.4b).

Fig. 1.4 Illustration of Ekman transport in the northern hemisphere (a) and upwelling of seawater by Ekman transport in the equatorial region (b) Reproduced with permission of Ref. [18]

1.1 Sea and the Earth

1.1.2.3

7

Thermohaline Circulation

Deep-ocean currents are propelled by density variations in water, which are controlled by temperature (thermo) and salinity (haline). Thermohaline circulation is the technical term for this phenomenon [19]. Ocean water freezes in the Earth’s Polar Regions, generating sea ice. As a result, the surrounding ocean becomes saltier, as salt remains in the water when sea ice forms. Seawater becomes denser as it becomes saltier, and it begins to sink. Surface water is pulled in to replenish sinking water, which gets sufficiently cold and salty to sink (Fig. 1.5). This is the mechanism that drives the worldwide seawater conveyer belt of the deep-ocean currents [15, 20]. The North Atlantic and Antarctic Oceans are particularly known for dense water masses descending into deep basins. Winds and low ambient temperatures significantly reduce the surface water temperature of the North Atlantic Ocean. Winds blowing over water increase evaporation, which results in a fall in temperature associated with latent heat (referred to as evaporative cooling). Evaporation raises the salinity of residual saltwater by removing only water molecules, thus, increasing the density of the seawater mass and lowering the temperature. Evaporative cooling dominates in the Norwegian Sea, and a sinking water mass, the North Atlantic Deep Water (NADW), fills the basin and flows south through a gap in the seabed sill linking Greenland, Iceland, and the United Kingdom [21, 22]. However, unlike deep water formation, its upwelling is less localized and more difficult to observe than convection [23]. These thermohaline cycles occur over extended periods of time and contribute to the movement of energy and materials.

Fig. 1.5 Thermohaline circulation in Earth. Deep-ocean currents are driven by temperature and salinity; saltier and colder seawater sinks, and surface water is drawn inward to replace the sinking water, which eventually forms the global conveyer belt of seawater. Reproduced with permission of Ref. [14]

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Fig. 1.6 Distribution percentage (%) of total water, fresh water, and surface water (fresh) on Earth

1.1.3 Seawater and Mankind 1.1.3.1

Seawater and Fresh Water

One of the most abundant substances on the Earth is water (H2 O, approximately 1.3 × 1018 tons)[24], which covers approximately 71% of the earth’s surface [2]. Based on the concentration of salt, water is categorized into seawater (also called saltwater) and fresh water. Seawater makes up approximately 97.5% of total available water, while the remaining 2.5% is constituted by fresh water (Fig. 1.6) [5, 25]. Of the total amount of available fresh water, approximately 69% is in the form of glaciers, 30% is underground, and < 1% lies in lakes, rivers, and swamps. Of that, only 1% of the fresh water on the Earth’s surface is usable by humans, with the other 99% of the utilizable amount being underground. Additional details about the salt concentration of seawater and fresh water are described in Sect. 1.2.1. As it is believed that life originated from the sea, fresh water is an essential resource for most living beings, including humans. However, humans comprising 0.9% salt cannot consume seawater, which contains 3.5% salt. This means that most of our drinking water comes from fresh water and not the sea. Even though the absolute amount of fresh water on Earth is sufficient compared to the amount of water needed by the total human population, many parts of the world are experiencing water shortages.

1.1.3.2

Seawater and Climate Change

With the advent of the industrial revolution, mankind began to utilize an increasing amount of resources. Among them, fossil fuels have been widely used for personal

1.1 Sea and the Earth

9

transport, freight transportation, and power generation. As a result, the amount of carbon dioxide, a chemical by-product of fossil fuels, in the atmosphere has increased dramatically. Carbon dioxide is one of the representative greenhouse gases and is accelerating global warming. The ocean is sink for this carbon dioxide and is estimated to absorb more than 25% of the carbon dioxide emitted by humans in the last 20 years [26–28]. One of the key mechanisms by which the ocean absorbs carbon dioxide from the atmosphere is the process of calcium carbonate formation by microorganisms on the sea surface [29]. When the microbes that produce calcium carbonate die, the calcium carbonate particles sink to the bottom and dissolve in the deep sea or on the seabed. As such, the ocean can absorb carbon dioxide; however excess carbon dioxide absorption can acidify the ocean [30]. Acidified oceans are known to slow down microbial decay and reduce the ocean’s ability to absorb carbon. Although seawater can tolerate a certain amount of pH change similar to a buffer solution and no noticeable global pH change has been observed yet, it is essential to continuously monitor the pH, especially in certain areas with high carbon dioxide emissions. Seawater not only absorbs carbon dioxide, a representative greenhouse gas, but also directly absorbs the atmospheric or solar heat. Approximately 93% of the solar heat that reaches the Earth is absorbed by the ocean [15]. Unlike land, the sea, which has a higher specific heat than land, is less sensitive to diurnal temperature changes and seasonal changes, which allows the global environment to achieve a more stable thermal balance. The ocean is an indicator of climate change. As mentioned earlier, the ocean plays a role in maintaining the Earth’s temperature and climate, either through the absorption of gases such as carbon dioxide or the direct absorption of solar heat. However, due to changes in the amount of carbon dioxide in the atmosphere or changes in the Earth’s radiant heat, the ocean is expected to gradually achieve a new thermal equilibrium. Some observations and interpretations of changes in the temperature of seawater, such as huge heat reservoirs, show that the average ocean temperature has increased since industrialization (Fig. 1.7). The elevated seawater temperature can cause a change in the existing seawater circulation structure, along Fig. 1.7 Annual temperature change in ocean and land surfaces after the pre-industrial era. Reproduced with permission of Ref. [32]

10

1 Seawater and Its Resources

with a decrease in the amount of dissolved carbon dioxide and consequent changes in marine ecology. Similarly, the ocean is directly affected by climate change, such that it can be used as a representative indicator for climate change observations [31, 32].

1.2 Characteristics of Seawater 1.2.1 Components of Seawater Seawater is an aqueous solution containing a wide range of water-soluble salts that mainly comprise sodium chloride. It consists of approximately 96.5% water and 3.5% salt [5, 33, 34]. Water with a salinity of 0.05% or less is considered to be freshwater; thus, seawater contains 70 times more salt than typical freshwater (such as in rivers, lakes, groundwater, and glaciers). Seawater has a salt content of approximately 3.3–3.7% and its proportion of salt varies depending on evaporation, rainfall, and freshwater inflow. In the Baltic Sea, for example, where freshwater flows from nearly 200 rivers, the dissolved salt concentration is low. In contrast, the Red Sea, surrounded by deserts, has a low freshwater inflow from rainfall or rivers and greater frequency of evaporation, presenting a high quantity of salt [1]. The salt content of seawater is referred to as salinity and is denoted by the symbol ‘S’. S represents the mass unit, gram) of the dried salt residue contained in 1 kg of seawater. For instance, seawater with a salinity of 35 (i.e., 35 practical salinity units, 35 PSU) refers to 35 g of salt is dissolved in 965 g of water, and is also expressed in 35 parts per thousand, 35 ppt (or 35,000 parts per million, 35,000 ppm) [15, 34]. This indicates that the density, temperature, and electrical properties of seawater determine its physical and chemical properties. Conversely, the amount of salinity of seawater can be estimated by its temperature and electrical conductivity. Although the salinity of seawater may have a distribution range, the proportion of its main inorganic components shows a remarkably constant ratio. In other words, for any sample of seawater from various areas, its ratio of major inorganic components is explicitly uniform. This is possible because the speed of solution in seawater for maintaining the inorganic salt composition is faster than the speed of changing salinity in seawater. Consequently, the ratios among inorganic salt constituents in different seas remain constant. Note that the salt in seawater exists in the form of dissolved ions. The main ions (that is, cations and anions), comprising of 99.9% of salt, are as follows [1]. Cations: Na+ > Mg2+ > Ca2+ > K+ > Sr2+ Anions: Cl− > SO4 2− > HCO3 − > Br− > BO3 2− > F− These inequality signs represent the order of reduction of ionic bonds. The major concentrations of dissolved ions in the salinity of standard seawater are listed in Table 1.1 [4, 15]. Here, Na+ and Cl− ions are the dominant dissolved ions in seawater; Na+ and Cl− ions together comprise approximately 85% of all salt

1.2 Characteristics of Seawater

11

Table 1.1 Major composition of dissolved ions in the salinity of standard seawater. Reproduced with permission of Refs. [1, 4, 15] Chemical ions

Percentage of total salt content (%)

Concentration (ppm) g/kg

Cl−

55.0

19.345

30.6

10.752

SO4 2− (Sulfate)

7.6

2.701

Mg2+

3.7

1.195

(Chloride)

Na+ (Sodium) (Magnesium)

Ca2+ (Calcium)

1.2

0.416

K+ (Potassium)

1.1

0.390

HCO3 − (Bicarbonate)

0.4

0.145

Br− (Bromide)

0.2

0.066

BO3 2− (Borate)

0.08

0.027

Sr2+

0.04

0.013

0.003

0.001

(Strontium)

F− (Fluoride)

content (10.7 g and 19.3 g, respectively). Table 1.2 shows the comparison between the composition of dissolved ions in the salinity of seawater and river water. The sum of the six main components (i.e., chloride, sodium, sulfate, magnesium, calcium, and potassium) of seawater constitutes more than 99% of seawater. In particular, the portion of Na+ and Cl− is deceived in composing salinity and this enables seawater to achieve higher salinity than river water [15]. Table 1.2 Comparison of salt concentration in dissolved ions in seawater and river water. Reproduced with permission of Ref. [15] Chemical ions

Percentage of total salt content in seawater (%)

Cl− (Chloride)

55.0

Na+ (Sodium)

30.6

6.9

SO4 2− (Sulfate)

7.6

12.4

Mg2+ (Magnesium)

3.7

4.6

Ca2+ (Calcium)

1.2

16.6

K+ (Potassium)

1.1

2.6

HCO3 −

(Bicarbonate)

Percentage of total salt content in river water (%) 8.6

0.4

31.9

Br− (Bromide)

0.2

−

BO3 2− (Borate)

0.08

−

Sr2+

0.04

−

(Strontium)

F− (Fluoride)

0.003

−

SiO2 (Silica)

−

14.6

Fe2+ (Iron)

−

0.7

NO3 − (Nitrate)

−

1.1

12

1 Seawater and Its Resources

1.2.2 Physical Properties of Seawater Since seawater is an aqueous solution, it shows similar characteristics as the physical properties of water. However, there are slight differences due to salinity and distribution, which mutually affect various physical properties of seawater. Some of the main features are briefly described as follows.

1.2.2.1

Temperature Distribution

The temperature distribution of surface seawater is basically caused by the difference in solar heat received depending on latitude (Fig. 1.8). The equatorial region is at its highest at approximately 30 °C, and the Polar Regions drop as low as −2 °C [1, 15]. The altitude of the Sun changes according to the season, and the temperature distribution of the seawater changes accordingly, which results in the shift of the thermal equator (the connected line where the seawater temperature is at its maximum). In the northern hemisphere, the thermal equator moves to 20° north latitude in summer and descends south in winter. Ocean currents carry water with the characteristics of one latitude region to another. The Gulf Stream in the Atlantic Ocean and the Kuroshio Current in the North Pacific move warm water northward to high latitudes. In contrast, the Labrador Current and the Oyashio Current flow south to high latitudes and carry cold water. At the point where the warm and cold currents

Fig. 1.8 Surface seawater temperature distribution. Reproduced with permission of Ref. [36]

1.2 Characteristics of Seawater

13

meet, a sharp temperature change line is created [35]. It is characterized by a large amount of nutrients and plankton, forming a habitat for a variety of organisms. The temperature of the oceans drops as depth increases. At deeper depths, there are no seasonal fluctuations. Whereas, at the surface water layer the temperature ranges from 30 to 1 °C [1, 15]. As with salinity, the temperature in the deep water layer is determined by the circumstances at the surface. At low latitudes, the temperature difference between the surface water layer and deep water layer is significant. In the upper temperate and arctic areas, the formation of a thick water surface layer descends to the deep water layer and results in near-isothermal conditions with depth. In areas of the seas where the surface temperature changes annually, a shallow wind-mixed layer of higher temperature forms in the summer. The temperature quickly declines with depth under this isothermal layer, which is roughly 10– 20 m thick, forming a shallow thermocline (i.e., a layer of steep vertical temperature change) approximately between 200 and 1000 m, as seen in Fig. 1.9 [1, 15]. Winter cooling and enhanced wind mixing at the ocean surface result in the erasure and deepening of the thermocline by convective overturning and mixing. When summer returns, the seasonal thermocline re-forms. A weaker nonseasonal thermocline separates water from temperate and subpolar sources at greater depths. The temperature gradually lowers below this permanent thermocline. Temperatures may slightly increase with depth in the very deep ocean basins. This occurs when the ocean’s deepest regions are flooded with water of the same temperature from a single source. As this water sinks, it undergoes an adiabatic temperature rise. As the higher temperature is caused by compression, which increases the density of the water, such a temperature rise creates a stable water column. For example,

Fig. 1.9 Vertical seawater temperature distribution. Reproduced with permission of Ref. [15]

14

1 Seawater and Its Resources

lowering 2 °C surface seawater to a depth of 10,000 m (about 33,000 ft) results in a temperature increase of around 1.3 °C [1]. When monitoring deep-sea temperature, the adiabatic temperature rise may be computed from the observed temperature to produce the potential temperature; this can be used to identify a parcel of water and track its origin.

1.2.2.2

Density

At 20 °C, the normalized average density of seawater is between 1,024 and 1,030 kg/m3 , whereas for fresh water it is approximately 1000 kg/m3 [15]. In general, salt ions are heavier than water molecules, making seawater denser than fresh water. Temperature and salinity directly affect the density of surface seawater. Pressure also exerts an influence on seawater, but it is not as strong as the effects of other parameters. Regarding the density of surface seawater regionally, the equatorial region with higher temperatures shows the lowest density, and the coldest polar region shows the highest density. One interesting fact is that, while pure water has the highest density at 4 °C, seawater has the highest density at −2 °C, i.e., the freezing point (Fig. 1.10). Accordingly, the density distribution of seawater shows a mirrored temperature distribution—the lower the temperature, the higher the density [37]. The vertical density distribution of seawater also shows that the density is directly affected by the temperature distribution (Fig. 1.11). The density of seawater is lowest on the warmest surface of the water and increases with increasing depth. For surface seawater, the density of seawater is 1.07 g/cm3 at a depth of 100 m. Earlier, we referred to a section where the temperature rapidly decreases with water depth, the thermocline. In this section, the density increases rapidly under the influence of temperature. This is called the pycnocline and it coincides with the thermocline in most cases. Under the pycnocline, the value of the density is almost as constant as the temperature [15]. Fig. 1.10 Surface seawater density change with seawater temperature. Reproduced with permission of Ref. [38]

1.2 Characteristics of Seawater

15

Fig. 1.11 Seawater density change with seawater temperature. Reproduced with permission of Ref. [15]

1.2.2.3

Acoustics

Sound travels faster in liquids than in gases. This means that seawater can transmit sound much better than air. The average speed of sound in seawater reaches approximately 1560 m/s, 4–5 times the speed of sound in air, which is approximately 343.2 m/s, while the impedance of sound waves in seawater is 1.54 × 106 [1]. Sound waves can be more useful than electromagnetic waves in the marine environment. Electromagnetic waves are strongly attenuated in seawater, while sound waves only suffer a small transmission loss. Sound can determine the position and distance of an object from the reflected sound waves when it encounters obstacles [39]. Marine creatures such as dolphins also use sound waves to identify obstacles. The speed of sound in seawater can change under the influence of several variables such as temperature, salinity, and pressure; it increases by 3–4 m/s for every 1 °C increase in seawater temperature. For every one unit of increase in salinity, the speed of sound increases by 1.1 m/s and for every 10 atmosphere pressure units (approximately 100 m of depth), it increases by 1.8 m/s [1]. Similarly, changes in these factors can enhance the attenuation of sound waves and the speed of sound [40]. Considering the temperature, salinity, and pressure distribution of seawater, the main factor that directly affects the speed of sound is the temperature at the surface of the ocean. If there is a temperature change when the sound wave moves along the surface layer of seawater, we can inversely estimate the temperature and current of a large area of the sea from the change in the speed of the sound wave. This thermal analysis using sound waves is called acoustic thermometry and is a representative tomography technique [41]. This technology estimates temperature by accurately

16

1 Seawater and Its Resources

Fig. 1.12 Vertical acoustic velocity distribution and the SOFAR channel in seawater. Reproduced with permission of Ref. [42]

measuring the time taken for sound to travel between a known pair of acoustic sources and receivers. Several measurements spread out over a 100–5000 km range can be used to estimate the temperature distribution. In deep seawater of depth approximately 1 km or less, the pressure change has a greater effect on the speed of sound than the temperature and salinity change. As the depth of the water increases, the pressure rises, which causes the speed of sound to increase. From the vertical characteristics of seawater temperature and pressure, sound waves can have a minimum speed of sound at a certain depth. This region is called the SOund Fixing And Ranging (SOFAR) channel [42]. Sound waves generated in the SOFAR channel can be trapped by refraction (Fig. 1.12). The sound wave transmitted to the upper or lower layer of the SOFAR channel, where the speed of sound is faster, is refracted toward the SOFAR channel with a smaller reflection angle compared to the angle of incidence. As a result, sound waves in the SOFAR channels can propagate over very long distances with little attenuation (loss of energy).

1.2.2.4

Optics

Light is a form of electromagnetic radiation, and its travel speed is 3 × 108 m/s in a vacuum. However, when light passes through seawater, its speed is reduced to 2.2 × 108 m/s [43]. Once sunlight enters seawater, the intensity of light decreases exponentially with increasing depth from the surface, and this decrease is called attenuation [1]. This exponential intensity loss of light is caused by two main factors: (i) absorption and (ii) scattering. Absorption occurs in the presence of algae, inorganic and organic particulate matter in suspension, dissolved organic molecules, and water itself (the electromagnetic energy of light is converted to heat in the molecules of

1.2 Characteristics of Seawater

17

Fig. 1.13 Transmission of light in seawater. Reproduced with permission of Ref. [45, 46]

water). Scattering occurs due to the bouncing of light at the air–water interface and also dust and other objects. Here, 60% of light is absorbed over a depth of 1 m depth, while 80% is absorbed over 10 m, such that there is no light penetration below 1000 m [44]. Generally, sunlight for photosynthesis is observed within a depth of 200 m, a zone for primary photosynthetic production. Note that blue light can penetrate seawater to a greater depth than other colors in the visible light spectrum (Fig. 1.13) because all other visible light is absorbed by water, causing the water to appear blue [45, 46].

1.2.3 Chemical Properties of Seawater Seawater, one of the most renewable and abundant resources, is an aqueous solution characterized by different chemical properties. In this section, the major chemical properties of seawater are discussed in detail.

18

1.2.3.1

1 Seawater and Its Resources

Thermal Properties

Water is a substance with high heat capacity due to hydrogen bonding between the water molecules. The specific heat of pure water is 1 cal/(g°C), by definition. In seawater, salinity weakens the hydrogen bonds between water molecules, which reduces the specific heat. Seawater with an average salinity (35 PSU) has a specific heat of 0.932 [47]. Salinity also affects freezing and boiling points (Fig. 1.14). Pure water freezes at 0 °C and boils at 100 °C under normal pressure conditions, but the addition of salt lowers the freezing point and raises the boiling point compared to a dilute solution [48, 49]. In other words, the crystallization of the water molecules in seawater is delayed by impurities that decrease its freezing point to −1.9 °C. The increase in the boiling point can also be explained by the vapor pressure drop due to presence of impurities. As mentioned in the previous section, the temperature of seawater ranges from approximately −2–30 °C. In other words, seawater is partially frozen in the Polar Regions, and this phase change releases a large amount of latent heat, which acts as a buffer against temperature changes. Pure water has the highest density at 4 °C. Addition of salt lowers the temperature of maximum density, which decreases faster than the freezing point due to the salt. At 24.70 PSU, the freezing point and maximum density temperature coincide at −1.332 °C. In other words, seawater (generally 35 PSU) always has the greatest density at its freezing point [50, 51].

Fig. 1.14 Phase diagram of saline water. Reproduced with permission of Ref. [50, 51]

1.2 Characteristics of Seawater

1.2.3.2

19

pH

Potential of hydrogen (pH) is a scale used to state the acidity and basicity of an aqueous solution. It is indicated as a logarithm and is inverse to the value of the concentration of hydrogen ions. Generally, the range of seawater pH varies from 7.5 to 8.5, but it normally reflects a pH value of 8.1, which is slightly basic relative to freshwater (pH = approximately 7). The presence of ions, especially carbonate ions, turns seawater into a weakly basic solution. However, due to a lack of consensus on measurements and knowledge of the pH in seawater, [52, 53] the hydrogen ion concentration is difficult to measure clearly unlike other ion concentrations such as K+ and Cl− . This is because H+ ions are involved in the autoionization of water and forming bonds with other components like CO3− . Therefore, various methods of measuring the pH in seawater are used for differing reasons. One method for measuring pH is by measuring active products such as HCl instead of H+ or OH− [54]. The definitions of pH are various, including free pH (pHf ) and seawater scale pH (pHSWS ) The pH from the International Union of Pure and Applied Chemistry (IUPAC) [55] is shown in Eq. 1.1 as follow: p H = −log(a H + ) = −log

m

H +γH + m0

 (1.1)

where aH+ , aHCl , and aCl− are activities of H+ , HCl, and Cl− ; mH+ mHCl , mCl− are molar concentrations of H+ , HCl, and Cl− ; rH+ and rCl− are activity coefficients of H+ and Cl− ; and m0 is the standard molarity (1 mol/kg H2O ). The activity formula of HCl expresses hydrogen ion activity in Eqs. 1.2 and 1.3 as follows: 1

a H C I = [(a H + )(γC I − )(m C I − )] 2

(1.2)

(a H + ) = (a H C I )2 /[(γC I − )(m C I − )]

(1.3)

As above, the activity of HCl, the concentration of Cl− , and the activity coefficient can be measured except for the activity of hydrogen ions. This is one method to determine the level of seawater pH [55]. Another way to estimate pH is a geochemical method, the FREZCHEM model [56], using the correlation of four parameters: pH, carbon dioxide partial pressure, total basicity, and dissolved carbon minerals. The FREZCHEM model calculates the level of pH based on the pCO2 − TA relationship. Equation 1.4 shows the charge balance of seawater as follows:

20

1 Seawater and Its Resources



             H + + N a + + K + + Mg O H + + 2 Mg 2+ + 2 Ca + + 2 Sr 2+       2−   −   2− + H S O− + CI = H C O− 3 + 2 C O3 4 + 2 S O4    −   − − − + O H + Br + B(O H )4 + [F ] (1.4)

The rearranging equations for HCO3 – , CO3 2− , and OH− is then,               i = H + + N a + + K + + Mg O H + + 2 Mg 2+ + 2 Ca + + 2 Sr 2+   2−   −   −     − − C I − Br − B(O H )− (1.5) − H S O− 4 − 2 S O4 4 − [F ] If the concentrations of HCO3 – , CO3 2– , and OH− are expressed using the equilibrium constant equation, Eq. 1.6 is as follows: i =

K H K 1 f C O2 aW K W aW 2K H K 1 K 2 f C O2 aW   + + + 2 (H + )(γ H C O3 ) (γ + O H − )(H ) (H ) γC O 2− 3

(1.6)

where KH is Henry’s constant for CO2 , K1 and K2 are the first and second dissociation constants for carbonic acid, Kw is the dissociation constant for water, (H+ ) is the activity of the hydrogen ion, fCO2 is the fugacity of CO2 , aw is the activity of water, γ values are activity coefficients, and i is the difference in non-carbonate species that are either measured directly (e.g., Na+ or Cl− ) or estimated mathematically using algorithms for minor species (e.g., MgOH+ , HSO4 − , F− , and B(OH)4 − ). Using the mentioned models, the level of pH can be measured by substituting all other parameters except hydrogen ion concentration—equilibrium constant value, concentrations, carbon dioxide fugacity, activity coefficient, etc. The level of pH can be determined by selecting one model from various models using their specific parameters. Figure 1.15 illustrates the relationship Fig. 1.15 Model-calculated dependence of pH on salinity. Superscript M and FC represent Miami and FREZCHEM models, respectively, and subscripts MacInnes, F, and SWS represent differences according to MacInnes assumption, free pH definition, and seawater scale pH definition, respectively. Reproduced with permission of Ref. [53]

1.2 Characteristics of Seawater

21

between seawater pH and salinity. Differences according to inference models (Miami model and FREZCHEM model), pH definition (pH, pHF , pHSWS ), and assumptions (MacInnes convention) can be confirmed.

1.2.3.3

Ionic Conductivity

Seawater has a high ionic conductivity of approximately 50,000 μS/cm (at 20 °C) due to dissolved and dissociated salt ions [57–60]. The conductive degree is much higher than other types of water such as deionized water (approximately 5.5 μS/cm at 25 °C) or typical drinking water (5–50 mS/cm), as shown in Table 1.3 [61]. The size and conductivity of bare and hydrated ions are compared in Table 1.4. Seawater, which has a high salinity compared to freshwater, also exhibits high Table 1.3 Comparison of conductivity value between types of water at 20 °C. Reproduced with permission of Ref. [61, 66]

Table 1.4 Comparison of the sizes of bare and hydrated ions and ionic conductivity values. Reproduced with permission of Ref. [65−69]

Water source

Conductivity

Distilled water

5.5 μS/m

Drinking water

5–50 mS/m

Seawater

50,000 μS/cm

Industrial wastewater

10,000 μS/cm

Melted snow

2–42 μS/cm

Ion

Bare ion size (Å)

Hydrated ion size (Å)

Ionic conductivity (S cm2 mol−1 )

H+

1.15

2.80

350.1

Li+

0.60

3.82

38.69

Na+

0.95

3.58

50.11

K+

1.33

3.31

73.5

NH4 +

1.48

3.31

73.7

Mg2+

0.72

4.28

106.12

Ca2+

1.00

4.12

119

Ba2+

1.35

4.04

127.8

Cl−

1.81

3.32

76.31

NO3 −

2.64

3.35

71.42

SO4 2−

2.90

3.79

160.0

OH−

1.76

3.00

198

ClO4 −

2.92

3.38

67.3

PO4 3−

2.23

3.39

207

2−

2.66

3.94

138.6

CO3

22

1 Seawater and Its Resources 70

Seawater Conductivity (mS/cm)

S=40 60

S=35 S=30

50

S=25 40

S=20

30

S=15

20

S=10 S=5

10 0 0

5

10

15

20

25

30

o

Seawater Temperature ( C) Fig. 1.16 Seawater conductivity with respect to temperature (from 0 to 30 °C) and salinity ((S =) from 5 to 40). Reproduced with permission of Ref. [38, 47]

conductivity due to ions in seawater, such as Na+ and Cl− . This indicates that seawater, as an electrolyte, has a lower equivalent series resistance (ESR) than freshwater. In practice, the differences of ionic conductivity of seawater occur depending on salinity and temperature as well as the electromagnetic attenuation characteristics and layers of seawater (Fig. 1.16). The ionic conductivity values, for instance, vary with depth due to the differences in the vertical distribution of seawater and its temperature and salinity. Surface water with a relatively high temperature has higher ionic conductivity than deep seawater. However, a deep seawater layer, at a depth of 1.8–2.6 km, possesses a salinity and pressure which increases the ionic conductivity rather than changing the temperature [59]. Therefore, seawater acts as an electrolyte and can be used in water electrolysis and electrochemical cells [62–64]. Note that the conductivity of seawater is determined with regard to temperature and salinity [15, 57, 60].

1.2.3.4

Dissolved Gas in Seawater

Seawater contains a variety of gases, including nitrogen (N2 ), oxygen (O2 ), and carbon dioxide (CO2 ), all of which are vital to life on Earth [15, 34]. These gases are exchanged with the atmosphere to maintain the ocean–atmosphere equilibrium.

1.2 Characteristics of Seawater

23

Table 1.5 Saturation coefficient of gases (O2 , N2 , and CO2 ) in water depending on chlorinity and temperature in equilibrium with 1 atm pressure. Reproduced with permission of Refs. [70, 71, 73] Chlorinity (%)

Gas

Temperature

0

O2

49.24

N2

23.00

CO2

1715

O2

40.1

N2

15.02

CO2

1489

O2

38.0

N2

14.21

CO2

1438

0 °C (mL/L)

16

20

12 °C (mg/L)

(mL/L)

4.40

36.75

2.06

17.80

77.0

1118

3.60

30.6

1.73

11.56

66.8

980

3.40

29.1

1.64

10.99

64.5

947

24 °C (mg/L)

(mL/L)

3.28

29.38

1.59

14.63

50.2

782

2.75

24.8

1.33

9.36

44.0

695

2.61

23.6

1.26

8.96

42.5

677

(mg/L) 2.62 1.31 35.1 2.22 1.08 31.2 2.12 1.03 30.4

The mixing of the surface water layer caused by wind and waves facilitates this interaction. On the surface of seawater, the dissolved gas is composed of 48% nitrogen, 36% oxygen, and 15% carbon dioxide. However, beneath the surface water layer, the dissolve gas is composed of 83% carbon dioxide, 11% nitrogen, and 6% oxygen, demonstrating seawater’s ability to store carbon dioxide [15]. Marine life is profoundly affected by these dissolved gases. Most of marine plants, algae, phytoplankton, and marine animals require oxygen to respire. Marine plants require carbon dioxide for photosynthesis. In a nitric acid fixation reaction, bacteria convert dissolved nitrogen to usable nitrogenous chemicals. The amount of dissolved gas is typically determined by temperature, salinity, and pressure (Table 1.5). Cold water contains a greater concentration of gas than warm water. Due to the exothermic nature of the solvation process by which gas is dissolved, oxygen, nitrogen, and carbon dioxide all dissolve more readily at lower temperatures. Salinity has an effect on solubility as well. Low-salinity seawater contains more gas than high-salinity seawater [70]. This is explained by the salting-out effect, a phenomenon in which the solubility of an insoluble gas reduces with increasing salt concentration. In the deep sea, pressure-induced changes in gas solubility can be seen. It is thought that there may be more gas dissolved in the deep sea due to the increasing water pressure. Along with these physical considerations, changes in the amount of dissolved gas induced by biological activity must be considered. Due to the critical role of dissolved oxygen in biological activities, the dissolved oxygen distribution in the seas has been extensively studied. Other types of trace gases can also be dissolved in seawater. Ammonia, argon, helium and neon have all been detected in seawater, and hydrogen is present in trace amounts. Hydrogen sulfide can exist in an environment free of dissolved oxygen, furthermore, other decomposition products, such as methane, can occur in

24

1 Seawater and Its Resources

stagnant water [71]. Finally, the presence of helium and neon in brine has also been confirmed [72].

1.3 Seawater Resource Technologies As discussed in the previous sections, seawater is part of the Earth’s energy system and a source of water and various minerals. This massive body of fluid transports a large amount of matter and energy across the Earth. Before modern times, humans mainly used the sea for biological resources and salt, however, with technological development we began to acquire various seawater materials and utilize the sea as an energy source. Seawater resource technology utilizes material and energy resources from seawater. Considering seawater as a source of matter, water and minerals can be separated from seawater and used as resources. Technological developments have enabled the utilization of tidal motion, temperature difference, and salinity difference as energy resources. A brief outline is given below for each seawater resource technology in Fig. 1.17. Seawater resource technology is important for reducing and diffusing human´s dependence on terrestrial resources. Historically, industries have depended on terrestrial sources of carbon and minerals, and in the near-future are likely to face resource scarcity leading to access imbalances, price volatilities, and environmental problems. However, seawater exists in abundance on Earth and, theoretically, can be used virtually infinitely through circulation. Seawater resource technology is expected to solve these problems based on abundant material and energy resources derived from Fig. 1.17 Seawater resource technologies

1.3 Seawater Resource Technologies

25

seawater. This section examines the technologies for obtaining hydrogen, the nextgeneration energy resource, freshwater, and minerals from seawater, and introduces technologies for producing or storing energy using seawater.

1.3.1 Fresh Water (Desalination) Freshwater availability is decreasing worldwide owing to population growth, enhanced lifestyles, and rapid growth of industrial and agricultural activities. By 2025, it is expected that humans will have consumed 90% of available fresh water [74, 75]. Hence, it is extremely urgent to find an alternative to fresh water in the coming decades. The seawater desalination process is a potential method to overcome the freshwater scarcity problem due to the abundance of seawater on Earth. Therefore, in recent decades, considerable research has been conducted and technologies have been developed. Reverse osmosis (RO) and thermal desalination technologies are well established and commercialized, but RO technology is the most commonly utilized method (Fig. 1.18). In addition, capacitive deionization (CDI), electrodialysis (ED), and desalination battery technologies [74, 75] have also received considerable attention but are still in the testing phase. The use of fresh water produced by these technologies depends on the level of salinity. For instance, water with a salinity between 3000–12,000 mg/L is permissible for irrigation use [76]. Whereas, water with salinity below 500 mg/L is required for drinking purposes (Department of Water, Government of Western Australia). The salinity is calculated to be less than 180 mg/L by feeding saline water at 55 Pa through an RO membrane [77].

Fig. 1.18 Illustration for water resources and Reverse Osmosis technique

26

1 Seawater and Its Resources

1.3.2 Mineral Mining Seawater is a source of numerous chemical elements; nearly all elements in the periodic table can be found in seawater. Mining minerals from seawater is becoming more attractive as land-based mining industries are currently facing several difficulties including decreased availability of high-grade ores, environmental issues associated with overproduction and health hazards posed to miners, and demand for large quantities of land, water, and energy [24, 78]. Mining minerals from seawater is not novel and is as ancient as common salt, NaCl, that has been extracted from seawater since before 2200 BC in China [33] and was reportedly practiced in the old testament period [24]. Currently, there are several metals that are commercially extracted from seawater and seawater brine specifically, such as Na, Mg, Ca, K, and Li in the form of Cl− , SO4 2− , and CO3 2− (Table 1.6). Magnesium is also mined in the form of Table 1.6 Major metallic elements dissolved in the sea, their concentrations, estimated deposits and land deposits of those elements. Reproduced with permission of Ref. [4] Element

Concentration in seawater (ppm)

Total oceanic abundance (tons)

Mineral reserves on Land (tons)

Na

10,800

1.40 × 1016

−

Mg

1,290

1.68 × 1015

2.20 × 109

Ca

411

5.34 ×

−

K

392

5.10 × 1014

8.30 × 109

Li

0.178000

2.31 × 1011

4.10 × 106

0.021000

2.73 ×

1010

1.90 × 108

Mo

0.010000

1.30 ×

1010

8.60 × 106

Ni

0.006600

8.58 × 109

6.70 × 107

0.005000

6.50 ×

109

1.80 × 108

Fe

0.034000

4.42 ×

109

1.50 × 1011

U

0.003300

4.29 × 109

2.6–5.4 × 106

0.001900

2.47 ×

109

1.30 × 107

Ti

0.00100

1.30 ×

109

7.30 × 108

Al

0.001000

1.30 × 109

2.50 × 1010

0.000900

1.17 ×

109

4.90 × 108

Mn

0.000400

5.20 ×

108

4.60 × 108

Co

0.000390

5.07 × 108

7.00 × 108

0.000280

3.64 ×

108

6.10 × 106

Cr

0.000200

2.60 ×

108

4.75 × 108

Cd

0.000110

1.43 × 108

4.90 × 105

0.000030

3.90 ×

107

7.90 × 107

Au

0.000011

1.43 ×

107

4.20 × 104

Th

0.0000004

5.20 × 105

1.30 × 106

Ba

Zn

V

Cu

Sn

Pb

1014

1.3 Seawater Resource Technologies

27

Fig. 1.19 Concentrations of elements dissolved in seawater and their cost effectiveness. Reproduced with permission of Refs. [24, 80, 81]

MgO [24, 79]. However, most of the minor minerals present in seawater have not been commercially extracted as their market price is much lower than the capital and operation costs of extraction [24]. Seawater is a rich source of precious minerals, many of which are uncommon and expensive on land. However, only a few minerals, those that are abundant in large quantities, are currently extracted from the sea (Fig. 1.19). Due to the current difficulties encountered by land-based mining businesses as a result of mineral depletion, sustainable water and energy demand, and environmental concerns, saltwater mining is becoming an appealing option. Solar evaporation, electrodialysis (ED), membrane distillation crystallization (MDC), and adsorp-tion/desorption are all described as extraction processes. The criteria for producing salt through the evaporation of seawater are as follows: (1) a hot, dry climate with dry winds, (2) available land and proximity to the sea, (3) nearly impermeable soil, (4) large areas of flat ground at or below sea level, (5) little rainfall during the evaporation months, (6) no possibility of dilution by freshwater streams, and (7) inexpensive transportation or nearby markets. The extraction of salts from brine released as effluent following the distillation of fresh water from saltwater has received increased attention. The use of brine in mineral extraction has a number of significant advantages. Firstly, the distillation plant pays for the pumping of the brine; second, the brine temperature is relatively high, which assists in evaporation; and third, the salt concentrations in these effluents are up to four times that in primary seawater.

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1 Seawater and Its Resources

1.3.3 Hydrogen (H2 ) Production Hydrogen is considered as the most viable candidate for future energy supplies, as an alternative to fossil fuels. Hydrogen is one of the most abundant chemical elements in the universe. It is a colorless, odorless, nonmetallic, tasteless, and highly flammable diatomic gas. In addition, hydrogen has the highest gravimetric energy density (approximately 142 MJ/kg (higher heating value) and approximately 120 MJ/kg (lower heating value) at 25 °C)) [82] than any other type of fuel. Generally, four technologies have been used to produce hydrogen: (i) thermochemical, (ii) electrochemical, (iii) photobiological, and (iv) photoelectrochemical [83]. Among these, electrochemical and photoelectrochemical technologies have attracted considerable attention for hydrogen production, owing to the use of water as a proton source. However, these two technologies use extremely pure distilled water-based synthetic electrolytes, which are usually highly alkaline or acidic, or a buffer system. Therefore, it is believed that the direct use of seawater could be a better alternative to the use of pure distilled water-based synthetic electrolytes. Seawater is a natural electrolyte with a pH of approximately 8.1 (slightly alkaline) and the most abundant water resource on Earth, which makes it suitable for low-cost, mass hydrogen production [84]. However, experts are examining the feasibility of using seawater in electrochemical [85, 86] and photoelectrochemical [87, 88] technologies for mass hydrogen production via a suitable OER/HER electrocatalyst and photoelectrode, respectively. Consequently, these technologies using seawater are still in the early experimental stages and require further research to achieve practical applicability.

1.3.4 Energy Production The movement of seawater in seas/oceans creates waves and tides. The waves are formed by the action of winds on the surface of seas, while tides are caused by the gravitational forces of the Earth and their interaction with the gravitational forces exerted by the Moon and Sun. Several decades ago, engineers and scientists found that sea/ocean waves and tides could be utilized to produce power, leading to the development and commercialization of wave and tidal energy technologies [89]. These technologies were considered to be renewable, green energy technologies because they do not produce any greenhouse gases (e.g., carbon dioxide) similar to wind and solar power technologies. Moreover, there are several renewable energy technologies [90–92] such as marine current energy, ocean thermal energy, and salinity gradient energy, which have been investigated and exploited using seawater in many coastal countries and regions. The rise and fall of the tide (range) offers the opportunity to harness the high tide, delaying its descent using a barrage or fence, and then utilizing its potential energy before the next tidal cycle. The worldwide theoretical power of tides (including tidal currents) has been estimated at around 7,800 TWh/yr [93].

1.3 Seawater Resource Technologies

29

Since Georges Claude pioneered Ocean Thermal Energy Conversion (OTEC) approximately 80 years ago [94], generations of engineers have wished to exploit the ocean thermal gradient as a renewable resource. Considerable work began during the 1970s oil price shocks, but these efforts diminished over the next two decades due to less favorable political and economic situations. Meanwhile, proponents of OTEC and academics recognized that the ocean thermal gradient could be utilized for more than just electricity generation; it could also be used for desalination, cooling, and aquaculture. Though these technologies were created independently, these additional deep ocean water applications (DOWA) have been frequently envisioned as co-products capable of assisting OTEC in overcoming initial economic non-viabilities. The resource’s total sustainable size is restricted by the rate of formation of deep cold saltwater, but significantly large estimations based on solar fluxes are frequently proposed. Orders of magnitude between 3 and 10 TW are plausible, i.e., roughly double the current total power consumption to roughly half the current primary energy demands [95, 96]. The lower threshold indicates the possibility of local temperature gradient deterioration under extremely intense OTEC conditions. OTEC locations with favorable conditions are located far offshore, implying that significant development of OTEC would require floating structures rather than land-based plants. Nonetheless, tropical locations with steep bathymetries continue to be the best choices, and therefore OTEC presents a significant opportunity for energy production for tropical nations, including small island developing states, and large island nation states such as Indonesia, the Philippines, Papua New Guinea, and Taiwan. Moreover, Brazil has a lengthy coastline with strong ocean thermal gradients, while the Gulf of Mexico may present significant prospects for the United States. Considerable OTEC development is unlikely to occur in areas with significant logistical obstacles (e.g., a lack of infrastructure) or in governments that cannot readily afford the risk and burden associated with innovative capital-intensive technologies. In the longer term, the development of remote OTEC regions will almost certainly involve the production of energy vectors such as liquid fuels rather than direct power transmission to shore.

1.3.5 Seawater Electrochemical Energy Systems 1.3.5.1

Seawater Supercapacitors

Supercapacitors (SCs) have received significant attention as energy storage systems owing to their rapid charging/discharging rate, high power density, excellent life cycle, and high reliability compared to lithium-ion batteries [97–99]. The most common electrolyte solutions including KOH and H2 SO4 in SCs are considered as corrosive and dangerous, whose leakage during use can be extremely harmful to humans in the vicinity. Therefore, as an alternative to KOH and H2 SO4 electrolytes, seawater could be used for devising green and economical SCs. Seawater

30

1 Seawater and Its Resources

is a readily available mildly alkaline natural electrolyte that is low-cost and safe [97, 98]. Nonetheless, the development of SCs with seawater remains in the early stages of development.

1.3.5.2

Seawater Batteries

Seawater batteries have received considerable interest for marine applications (e.g., signal lights, life jackets, GPS, sensors for monitoring water temperatures and flow rates especially for climate change monitoring, and deep-sea applications). Earlier seawater batteries were developed using metal anodes of Mg, Al and their alloys, wherein seawater was mostly used to activate the battery as the electrolyte. The operating voltage of such batteries mainly depends on the anodes used, for example, for Mg, it ranges from approximately 1.8 to 1.4 V and for Al and Zn, it is below 1 V [100]. However, this operating voltage is very low compared to a state-of-the-art lithium-ion battery. In addition, these classical types of seawater batteries are nonrechargeable primary batteries and suffer from high self-discharging and electrode clogging due to the formation of a solid-electrolyte interphase (metal hydroxides) [101, 102]. It is well known that seawater contains a significant amount of Na+ ions, approximately 0.46 M; hence, it is possible to use seawater as a source of Na-ion cathodes, such as the cathodes that are used in Na-ion batteries. Considering this prospect, a novel rechargeable seawater battery has recently been developed (http://www.ykimre search.com/: Ulsan National Institute of Science and Technology, South Korea) by utilizing seawater as a cathode and electrochemically-consumed Na metal (−2.71 V vs standard hydrogen electrode, SHE) obtained from seawater as the anode [103]. This novel and advanced seawater rechargeable battery can deliver a high-voltage range of approximately 2.7–3 V [104–106]. A selective Na-ion conductive solid electrolyte (Na3 Zr2 Si2 PO12 ) is employed in this seawater battery to separate the anode compartment from the seawater cathode [103]. Moreover, the development of rechargeable seawater batteries has recently begun. Currently, this seawater battery technology is in the R&D stage for its commercialization and additional attempts have been made to modify this technology by considering desalination and hydrogen production features. In this chapter, we looked at several physicochemical properties of seawater, as well as technologies that utilize seawater as an energy and resource. Later chapters of this book aim to summarize the scientific and technical knowledge regarding primary and rechargeable seawater batteries by covering their background, fundamentals, overall development, and applications.

References

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73. Fox CJJ (1909) On the coefflcients of absorption of the atmospheric gases in distilled water and sea water: part ii: carbonic acid. ICES J Mar Sci 1:3–31. https://doi.org/10.1093/icesjms/ s1.44.3 74. Pasta M, Wessells CD, Cui Y, La Mantia F (2012) A desalination battery. Nano Lett 12:839– 843. https://doi.org/10.1021/nl203889e 75. Zhang Y, Senthilkumar ST, Park J, Park J, Kim Y (2018) A new rechargeable seawater desalination battery system. Batter Supercaps 1:6–10. https://doi.org/10.1002/batt.201800012 76. Xu P, Cath TY, Robertson AP, Reinhard M, Leckie JO, Drewes JE (2013) Critical review of desalination concentrate management, treatment and beneficial use. Environ Eng Sci 30:502– 514. https://doi.org/10.1089/ees.2012.0348 77. Yu YH, Jenne D (2018) Numerical modeling and dynamic analysis of a wave-powered reverseosmosis system. J Mar Sci Eng 6. https://doi.org/10.3390/jmse6040132 78. Diallo MS, Kotte MR, Cho M (2015) Mining critical metals and elements from seawater: opportunities and challenges. Environ Sci Technol 49:9390–9399. https://doi.org/10.1021/ acs.est.5b00463 79. Romero VCE, Tagliazucchi M, Flexer V, Calvo EJ (2018) Sustainable electrochemical extraction of lithium from natural brine for renewable energy storage. J Electrochem Soc 165:A2294–A2302. https://doi.org/10.1149/2.0741810jes 80. Mineral Commodity Summaries. In: U.S. Dep. Inter. U.S. Geol Surv. http://minerals.usgs. gov/minerals/pubs/mcs/index.html 81. Turekian KK (1968) Oceans. Prentice-Hall 82. Amikam G, Nativ P, Gendel Y (2018) Chlorine-free alkaline seawater electrolysis for hydrogen production. Int J Hydr Energy 43:6504–6514. https://doi.org/10.1016/j.ijhydene. 2018.02.082 83. Haryanto A, Fernando S, Murali N, Adhikari S (2005) Current status of hydrogen production techniques by steam reforming of ethanol: a review. Energy Fuels 19:2098–2106. https://doi. org/10.1021/ef0500538 84. Jin H, Liu X, Vasileff A, Jiao Y, Zhao Y, Zheng Y, Qiao SZ (2018) Single-crystal nitrogenrich two-dimensional Mo5N6 nanosheets for efficient and stable seawater splitting. ACS Nano 12:12761–12769. https://doi.org/10.1021/acsnano.8b07841 85. Li A, Sun Y, Yao T, Han H (2018) Earth-abundant transition-metal-based electrocatalysts for water electrolysis to produce renewable hydrogen. Chem - A Eur J 24:18334–18355. https:// doi.org/10.1002/chem.201803749 86. Lu X, Pan J, Lovell E, Tan TH, Ng YH, Amal R (2018) A sea-change: Manganese doped nickel/nickel oxide electrocatalysts for hydrogen generation from seawater. Energy Environ Sci 11:1898–1910. https://doi.org/10.1039/c8ee00976g 87. Hsu SH, Miao J, Zhang L, Gao J, Wang H, Tao H, Hung SF, Vasileff A, Qiao SZ, Liu B (2018) An earth-abundant catalyst-based seawater photoelectrolysis system with 17.9% solar-to-hydrogen efficiency. Adv Mater. https://doi.org/10.1002/adma.201707261 88. Kumari S, Turner White R, Kumar B, Spurgeon JM (2016) Solar hydrogen production from seawater vapor electrolysis. Energy Environ Sci 9:1725–1733. https://doi.org/10.1039/c5ee03 568f 89. Scruggs J, Jacob P (2009) Harvesting ocean wave energy. Science 323(80):1176–1178. https:// doi.org/10.1126/science.1168245 90. Von Jouanne A, Brekken TKA (2017) Ocean and geothermal energy systems. Proc IEEE 105:2147–2165. https://doi.org/10.1109/JPROC.2017.2699558 91. Esteban M, Leary D (2012) Current developments and future prospects of offshore wind and ocean energy. Appl Energy 90:128–136. https://doi.org/10.1016/j.apenergy.2011.06.011 92. Khan N, Kalair A, Abas N, Haider A (2017) Review of ocean tidal, wave and thermal energy technologies. Renew Sustain Energy Rev 72:590–604. https://doi.org/10.1016/j.rser.2017. 01.079 93. Tidal & Currents. In: Ocean Energy Syst. https://www.ocean-energy-systems.org/about-oes/ what-is-ocean-energy/tidal-currents/ 94. Claude G (1930) Power from the tropical seas. Mech Eng J Am Soc Mech Eng 52:1039–1044

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Chapter 2

Primary Seawater Batteries

Abstract A primary seawater battery is a purpose-built electricity storage system that are usually designed as a reserve-type battery, used in the seawater. The primary seawater battery specially developed for military purposes for use in the marine environment has the characteristics of long shelf life, fast activation when used and higher power then other primary batteries. After the development of Mg-AgCl batteries using magnesium and silver chloride as electrodes and seawater as electrolyte, research on cheaper materials and higher output and energy have been conducted. The purpose of this chapter is to provide an overview of primary (un-rechargeable) seawater batteries focusing on battery science, materials, and their technologies. The chapter starts with a brief history of batteries in general, followed by an introduction to the history, principles, composition, and classification of primary seawater batteries. Understanding the primary seawater battery introduced in this chapter is necessary for understanding the secondary (rechargeable) seawater battery introduced in Chap. 3.

2.1 Introduction 2.1.1 Brief History of Batteries Electricity is the foundation of modern civilization. Electrochemical cells were the only means of harnessing electrical power for many years after the discovery of electricity and continue to play an essential role in electrical technology today. A battery is a device that stores chemical energy and converts it into electricity, producing direct current upon the connection of its terminals through a closed electrical circuit. The term “battery” was coined by Benjamin Franklin in the 1740s, when he was experimenting with electricity using a set of linked Leyden jars (i.e., capacitors) [1, 2]. Interestingly, the first battery structures are possibly much older (dating to nearly 2000 years ago). Archaeologist Wilhelm König suggested that a jar found near Baghdad, Iraq, containing an iron rod wrapped in a copper sheet, could have generated electricity if filled with an acidic solution such as vinegar or wine [3].

© The Author(s), under exclusive license to Springer Nature Singapore Pte Ltd. 2022 Y. Kim and W. Lee, Seawater Batteries, Green Energy and Technology, https://doi.org/10.1007/978-981-19-0797-5_2

37

38

2 Primary Seawater Batteries

Whether the so-called “Baghdad battery” was in fact used for this purpose remains a matter of debate (Fig. 2.1). In 1800, Alessandro Volta, created the first true electrochemical (or electric) battery [4, 5]. A voltaic pile (Fig. 2.2) [4] can be structurally viewed as a prototype of a modern cell. It consisted of a pair of zinc and copper (or silver) discs, separated by a piece of cloth soaked in brine. Zn served as the anode, Cu or Ag as the cathode. Brine, which can conduct ions, acted as the electrolyte, with the cloth separating the anode and cathode to avoid a short circuit. In addition, Volta discovered

Fig. 2.1 Composition and structure of the Baghdad Battery

Fig. 2.2 Schematic drawing of an early voltaic pile. Reproduced with permission from Ref. [4]

2.1 Introduction

39

that the voltage could be increased if the cells were serially connected, as illustrated in Fig. 2.2. Following the invention of the voltaic pile, other types of batteries were developed to satisfy various chemical and performance requirements [6, 7]. The Daniell battery, invented in 1836, consisted of two vessels connected by a potassium chloride bridge; one vessel contained a zinc plate immersed in a zinc sulfate solution, the other a copper plate immersed in a copper sulfate solution [8]. The gas voltaic cell, invented in 1839, was a cell in which the electrolysis of water was reversed; this principle is the basis of the modern fuel cell [9]. The lead–acid battery, the first secondary battery, was invented in 1859, [10] and the Leclanché battery, a precursor of the manganese battery used today, was invented in 1866 [1, 11]. The nickel–cadmium battery was invented in 1899, and the nickel–iron battery in 1901. The alkaline battery, invented in the 1950s, uses potassium hydroxide as an electrolyte and has the advantages of minimal self-discharge and long life in adverse environmental conditions [12]. A primary battery having a lithium anode was created in 1970. Lithium has the advantages of very low potential energy and mass density; the goal of subsequent research has been to increase the electromotive force and energy density. The Matsushita Battery Industrial Co. introduced Li/(CF)n batteries in 1973, after which Sanyo Electric Co. and others developed Li/MnO2 and Li/SOCl2 batteries [13]. Ni/Cd and NiH2 batteries were developed in the 1960s and have been the subject of extensive research; [1, 14] their use as secondary batteries dates to 1989. Lithium-ion batteries have been produced since the 1970s and have been the most studied and exploited batteries to date. The rare earth elements La, Ce, Pr, and Nd were combined with the transition metals Ni, Co, and Al. The lithium-ion battery, developed by M. Stanley Whittingham, J. O. Besenhard, and John B. Goodenough, among others, was initially marketed in 1991 by Sony and Asahi Kasei, based on a concept launched by Yoshino in 1985 [13, 15]. Originally employed in small mobile devices, they have recently been widely utilized in electric vehicles [16–18]. Table 2.1 lists representative batteries that have been developed over time.

2.1.2 Types of Batteries Early batteries were named after the scientist who developed them or for their main characteristics. Today, batteries may be classified according to the main active material or metal electrode, the mechanism of the reaction occurring at the negative (or positive) electrode, or the intended use, shape, or size. However, in general, batteries are classified first and foremost according to whether they can be recharged, that is, whether they are primary or secondary batteries. Primary batteries are non-rechargeable: the electrochemical reaction that occurs inside the cell is irreversible [1, 25]. Therefore, primary batteries cannot be reused following complete discharge. They are generally cheaper, simpler to use, and more compact than rechargeable batteries, and offer significantly higher energy densities and lower self-discharge rates. Historically, primary batteries have dominated in

Name (type)

Baghdad Battery (Primary)

Voltaic Battery (Primary)

Daniell Cell (Primary)

Gas Voltaic Cell (Fuel cell)

Year

B.C. 250

1800

1836

1839

A: 2H2 → 4H+ + 4e− C: O2 + H2 O + 4e− → 4OH− E: Sulfuric acid EF: 1.23 V

A: Zn (s) → Zn2+ (aq) + 2e− C: Cu2+ (aq) + 2e− → Cu(s) E: Copper sulfate solution, Zinc sulfate solution EF: 1.1 V

A: Zn(s) → Zn2+ (aq) + 2e− C: 2H+ (aq) + 2e− → H2 (g) E: Dilute sulfuric acid EF: 0.76 V

A: Fe C: Cu E: Vinegar or wine (hypothesized) EF: 0.8–2.0 V

Material and Potentials

[20, 21]

[19]

The first historically confirmed battery

The earliest known battery [3]

Note

Table 2.1 Representative batteries by year of development. (A: Anode, C: Cathode, E: Electrolyte, EF: Electromotive force)

(continued)

40 2 Primary Seawater Batteries

Name (type)

Lead-acid Battery (Secondary)

Leclanché Cell (Primary)

Ni/Cd Battery (Secondary)

Ni/Fe Battery (Secondary)

Year

1859

1866

1899

1901

Table 2.1 (continued)

[11]

[21]

Note

A: Fe + OH− → Fe(OH)2 + 2e− Fe(OH)2 + OH− → Fe(OH)3 + e− 2Fe(OH)3 + Fe(OH)2 → Fe3 O4 + 4H2 O C: NiO(OH) + H2 O + e− → Ni(OH)2 + OH− E: Potassium hydroxide aqueous solution EF: 1.4 V

[21]

A: Cd + 2OH− → Cd(OH)2 + 2e− [22] C: NiOOH + H2 O + e− → Ni(OH)2 + OH− E: Potassium chloride aqueous solution EF: 1.32 V

A: Zn (s) → Zn2+ (aq) + 2e− C: 2MnO2 + 2NH4 + + 2e− → Mn2 O3 + 2NH3 + H2 O E: Ammonium chloride EF: 1.4 V

A: Pb(s) + H2 SO4 (l) → PbSO4 (s) + 2e− + 2H+ (aq) C: PbO2 (s) + 2e− + H2 SO4 (l) → PbSO4 (s) + 2OH− (aq) E: Concentrated sulfuric acid EF: 2.1 V

Material and Potentials

(continued)

2.1 Introduction 41

Name (type)

Alkaline Battery (Primary)

Lithium Metal Battery (Primary)

Nickel Metal Hydride Battery (Secondary)

Lithium-Ion Battery (Secondary)

Year

1957

1970s

1960–1990s

1970–1990s

Table 2.1 (continued) Material and Potentials

[24]

[23]

Note

A: Lix C6 → xLi+ + 6C + xe− C: Li1-x CoO2 + xLi+ + xe− → LiCoO2 E: EF: 4.1 V

[15]

A: MH + OH− → M + H2O + e− [22] C: NiOOH + H2 O + e− → Ni(OH)2 + OH− E: E: 1.35 V

A: Li → Li+ + e− C: SOCl2 , (CF)x, MnO2 etc MnO2 + Li+ + e− → MnO2 − (Li+ ) (CF)n + ne− → nC + nF− 2SOCl2 + 4Li+ + 4e− → 4LiCl + S + SO2 E: LiClO4 , LiAlCl4 etc EF: Max 4 V

A: Zn → Zn2+ + 2e− Zn2+ + 2OH− → ZnO + H2 O C: MnO2 + H2 O → MnO + 2OH− E: KOH (Alkaline solution) EF: 1.5 V

42 2 Primary Seawater Batteries

2.1 Introduction

43

two areas: consumer products for which minimizing the initial cost of the battery is essential, and electronic products (such as watches, hearing aids, and pacemakers) where drains are low or recharging is infeasible. The most common primary-battery chemistries are carbon–zinc (Leclanché), alkaline–manganese dioxide, lithium– manganese dioxide, lithium–sulfur dioxide, lithium–iron disulfide, lithium–thionyl chloride (LiSOCl2 ), silver–oxide, and zinc–air. Secondary batteries can be recharged electrically following every discharge throughout their lifetime. Their rechargeability stems from the reversibility of the electrochemical reactions that drive them. The lead–acid battery, developed by Gaston Planté in France in 1859, was the first secondary battery; [10] it remains one of the most widely used, owing to its stability, reliability, and economy. The bestknown secondary batteries are the lead–acid battery, the Li-ion battery, the Li/Na-S battery, and the Ni/MH battery [13]. These various rechargeable batteries are used in laptop computers, mobile phones, and all types of vehicles. Other types of batteries have been developed with novel structures and performances tailored to satisfy the requirements of specific fields (military, meteorological, medical, etc.) despite high development costs. A typical example of a specialpurpose battery is the reserve battery [21, 26]. This is a primary battery from which certain components have been intentionally excluded; it is generally used when a long shelf life is required. A step is necessary to activate such a battery before its use. For example, a reserve battery may exclude the electrolyte, which must be injected immediately before use. Reserve-type batteries include seawater batteries, silver–zinc batteries, lead–lead oxide batteries, and lithium reserve batteries. Table 2.2 lists selected battery systems and their theoretical and actual voltages, capacities, and powers.

2.1.3 The Birth of Primary Seawater Batteries The most frequently used type of seawater battery is the disposable, non-rechargeable “seawater-activated battery,” developed in the 1940s for military purposes [26]. It is a typical example of the reserve batteries designed in the mid-twentieth century to ensure reliable operation of military equipment and weaponry. Unlike the conventional batteries of the time, it had an extended shelf life and high energy density; it was therefore used as a special-purpose marine battery despite the high cost of its components and the inconvenience associated with battery activation. The reason for its unlimited shelf life was that the electrode components were stored in a dry and inert environment, away from seawater or electrolytes; they did not begin producing electricity until submerged in seawater. There are several different types of primary seawater batteries. All use metal as their anode material: throughout the cell’s functional life, the metal anode dissolves continuously in seawater. The numerous possible choices of cathode material, including silver chloride (AgCl), cuprous chloride (CuCl), lead chloride (PbCl2 ), and dissolved oxygen (DO) in seawater, have led to a variety of primary seawater

Zn

Mg

Zn

Zn

Cd

Zn

Zn

Zn

Li

Li

Li

Li

Magnesium

Alkaline MnO2

Mercury

Mercad

Silver oxide

Zinc/O2

Zinc/air

Li/SOCl2

Li/SO2

LiMnO2

Li/FeS2

Anode

Leclanché

Primary batteries

Battery type

FeS2

MnO2

SO2

SOCl2

Ambient air

O2

Ag2 O

HgO

HgO

MnO2

MnO2

MnO2

Cathode

3.25 2.64 3.50 1.38

1.34 0.91 1.6 1.65 1.65

Cd + HgO + H2 O → Cd(OH)2 + 2Ag Zn + Ag2 O + H2 O → Zn(OH)2 + 2Ag Zn + 1/2O2 → ZnO Zn + 1/2O2 → ZnO 4Li + 2SOCl2 → 4LiCl + S + 3.65 SO2 3.1

Zn + HgO → ZnO + Hg

2Li + 2SO2 → Li2 S2 O4 Li + MnIV O2 → MnIV O2 (LI+ ) 3.5 4Li + FeS2 → 2Li2 S + Fe 1.8

1.22

1.5

Zn + 2MnO2 → ZnO + Mn2 O3

1.52

5.55

6.15

5.27

4.46

3.69

2.8

Mg + 2MnO2 + H2 O → Mn2 O3 + Mg(OH)2

4.46

1.6

726

286

379

403

820

658

180

163

190

224

271

224

1307

1001

1175

1471

1353

1085

288

148

255

358

759

358

Wh/kg

1.5

3.0

3.0

3.6

1.5

−

1.6

0.9

1.35

1.5

1.7

1.5

Wh/kg

Practical values Ah/kg

V

g/Ah

Theoretical values V

Zn + 2MnO2 → ZnO Mn2 O3

Reaction mechanism

Table 2.2 Selected battery systems, with theoretical and practical values of key parameters. Reproduced with permission from Ref. [6]

−

500(5) (continued)

535(5) 260(5)

415(5)

1100(4)

1300(6)

230(5)

260(5)

590(4)

370(6)

−

525(6)

230(6)

55(6) 135(6)

470(6)

100(6)

400(4)

195(4)

100(4) 145(4)

165(4)

85(4)

Wh/L

44 2 Primary Seawater Batteries

Li

Li

Li/(CF)n

Li/I2 (3)

Zn

Li

Zinc/silver oxide

Thermal

Ni oxide

Ni oxide

Fe

Cd

Zn

H2

MH(1)

Zn

Edison

Nickel–cadmium

Nickel-zinc

Nickel-hydrogen

Nickel-metal hydride

Silver-zinc

AgO

Ni oxide

Ni oxide

Ni oxide

Pb

PbO2

FeS2

AgO

CuCl

I2 (P2VP)

(CF)n

Cathode

Lead-acid

Secondary batteries

Mg

Cuprous chloride

Reserve batteries

Anode

Battery type

Table 2.2 (continued) Wh/kg

1.4 1.35 1.73 1.5 1.35 1.85

Fe + 2NiOOH + 2H2 O → 2Ni(OH)2 + Fe(OH)2 Cd + 2NiOOH + 2H2 O → 2Ni(OH)2 + Cd(OH)2 Zn + 2NiOOH + 2H2 O → 2Ni(OH)2 + Zn(OH)2 H2 + 2NiOOH → 2Ni(OH)2 MH + NiOOH → M + Ni(OH)2 Zn + AgO + H2 O → Zn(OH)2 + Ag

1.8

2Li + FeS2 → Li2 FeS2 2.1

1.81

Zn + AgO + H2 O → Zn(OH)2 + Hg

Pb + PbO2 + 2H2 SO4 → 2PbSO4 + 2H2 O

1.6

2.8

Li + 1/2I2 → LiI Mg + Cu2 Cl2 → MgCl2 + 2Cu

3.1

3.53

5.63

3.46

4.64

5.52

4.46

8.32

1.38

3.53

4.14

4.99

1.42

283

178

289

215

181

224

120

726

283

241

200

706

524

240

434

372

244

314

252

1307

512

386

560

2189

1.5

1.2

1.2

1.6

1.2

1.2

2.0

1.8

1.5

1.3

2.8

3.0

Wh/kg

Practical values Ah/kg

V

g/Ah

Theoretical values V

nLi + (CF)n → nLiF + nC

Reaction mechanism

105

75

55

60

35

30

35

40(9)

30(8)

60(7)

245

250(5)

Wh/L

(continued)

180(10)

240(5)

60

120

100(5)

55(10)

70(10)

100(9)

75(8)

80(7)

900

635(5)

2.1 Introduction 45

Cd

Zn

Zn

Lix C6

Li

Li(Al)

Li(Al)

Na

Na

Silver-cadmium

Zinc/chlorine

Zinc/bromine

Lithium-ion

Lithium/manganese dioxide

Lithium/iron disulfide(2)

Lithium/iron monosulfide(2)

Sodium/sulfur(2)

Sodium/nickel chloride(2)

H2

H2

H2 /O2

H2 /air

Fuel cells

Anode

Battery type

Table 2.2 (continued)

Ambient air

O2

NiCl2

S

FeS

FeS2

MnO2

Li(1-x) CoO2

Br2

Cl2

AgO

Cathode Wh/kg

2.65

2Li(Al) + FeS → Li2 S + Fe + 1.33 2Al 2.1 2.58

2Na + 3S → Na2 S3 2Na + NiCl2 → 2NaCl + Ni

1.23

2.90

1.73

2Li(Al) + FeS2 → Li2 FeS2 + 2Al

1.23

3.50

Li + MnIV O2 → MnIV O2 (Li+ ) 3.5

H2 + 1/2O2 → H2 O

3.50

4.1

LixC6 + Li(1-x) CoO2 → LiCoO2 + C6

H2 + 1/2O2 → H2 O

9.98

1.85

Zn + Br2 → ZnBr2

0.037

0.336

3.28

4.17

2.54

2.12

Zn + Cl2 → ZnCl2

4.41

1.4

26,587

2975

305

377

345

285

286

100

309

394

227

32,702

3660

787

792

459

493

1001

410

572

835

318

−

−

2.6

2.0

1.3

1.7

3.0

4.1

1.6

−

1.1

Wh/kg

Practical values Ah/kg

V

g/Ah

Theoretical values V

Cd + AgO + H2 O → Cd(OH)2 + Ag

Reaction mechanism Wh/L

−

(continued)

−

−

190(11)

−

345(11) 115(11)

220(11)

350(11)

265

400(5)

60

-

120(10)

170(11)

130(11)

180(11)

120

150

70

−

70

46 2 Primary Seawater Batteries

CH3 OH

CH3 OH

Methanol/O2

Methanol/air

Ambient air

O2

Cathode Wh/kg

1.24 1.24

CH3 OH + 3/2O2 → CO2 + 2H2 O 0.20

0.50 5020

2000 6225

2480 −

−

Wh/kg

Practical values Ah/kg

V

g/Ah

Theoretical values V

CH3 OH + 3/2O2 → CO2 + 2H2 O

Reaction mechanism

(1) MH = metal hydride, data based on 1.7% hydrogen storage (by weight) (2) High-temperature batteries (3) Solid-electrolyte battery (Li/I2 (P2VP)) (4) Cylindrical bobbin-type batteries (5) Cylindrical spiral-wound batteries (6) Button-type batteries (7) Water activated (8) Automatically activated, 2–10-min rate (9) With lithium anodes (10) Prismatic batteries (11) Value based on cell performance

Anode

Battery type

Table 2.2 (continued)

−

−

Wh/L

−

−

2.1 Introduction 47

48

2 Primary Seawater Batteries

batteries [27–30]; Mg/AgCl, Mg/CuCl, Mg/DO, and Mg/water batteries are the most widely known [31]. The features of and reactions occurring at the various anodes and cathodes that can be employed in seawater batteries are discussed in Sect. 2.2. The primary seawater battery utilized in the early 1940s was a direct seawater electrolyte battery [1]. Such batteries, which are activated by immersion in seawater, were termed seawater-activated batteries and can be considered to be the predecessors of the primary seawater battery. Various anode metals such as zinc and lead were evaluated in the developmental days of primary seawater batteries. Magnesium electrodes were used initially, with silver chloride as the cathode. The Mg/AgCl battery was created in 1943; initially utilized for military purposes, it found civilian applications during the years 1945–1952. The AgCl electrode of the Mg/AgCl battery is generally resistant to common environmental conditions such as low temperature or high humidity. Such a battery should therefore have consistent performance and a long shelf life. However, the high cost of silver creates a significant barrier to the use of AgCl in electrodes meant to be utilized a single time before disposal. Consequently, numerous metal halides have been considered as replacements for silver chloride. A case in point is the cuprous chloride battery, invented in 1949 [26]. However, CuCl is strongly hygroscopic, which violates the requirement that a primary seawater battery have a long shelf life in the actual usage environment. By contrast, the AgCl electrode has a near-infinite shelf life. Nonetheless, the development of CuCl batteries stimulated the further development of primary seawater cell systems using inexpensive and stable anode materials such as PbCl2 , PbO2 , and Hg2 Cl2 instead of AgCl [27, 28]. Although these materials could replace the expensive silver chloride in some applications, their performance was less stable than that of silver chloride. Since the 1990s, Mg/CuI and Mg/Cu2 O primary seawater batteries have also been developed [30]. Since the commercialization of magnesium–silver chloride batteries, the majority of primary seawater batteries have been developed using magnesium as the negative electrode; magnesium’s lightness, high energy density, and ease of storage and activation have incentivized its adoption. Metals such as aluminum, zinc and lithium have also been considered [1, 24, 32, 33], but they have high reactivities toward oxygen or moisture; consequently, their reliability is reduced upon prolonged storage, and their voltage efficiency and capacity are reduced because of various side reactions [24, 32, 33]. Batteries developed for deep-sea applications and long-term underwater operation since the late 1980s use seawater not only as an electrolyte, but also as an active material in the anode reaction [31]. This decreases the cathode material’s weight and volume, permitting higher energy densities. Magnesium–seawater and aluminum–seawater batteries of this type, developed in the 1990s, have demonstrated their operability in the deep sea for extended periods of time. Such metal–seawater batteries utilize the DO or the water decomposition reactions occurring in natural seawater, such as DO reduction or hydrogen evolution. Several anode materials have been investigated for them, including magnesium and its alloys, zinc and its alloys, and aluminum. However, issues such as cell blockage and excessive self-discharge induced by side reactions (e.g., metal oxide production) have been encountered.

2.1 Introduction

49

Fig. 2.3 History of primary seawater batteries

Zn/AgO torpedo batteries, developed after 1960 [34], start to operate when seawater and electrolyte are injected into the battery at the time of launching. Even with a high load, these batteries can generate a stable current, but owing to the presence of the zinc electrode, they initially had limited shelf lives compared to other reserve batteries. Aluminum–silver oxide cells have been created to attain higher energy densities. The Al/AgO battery stores the soluble solid alkaline material in an electrolyte tank that dissolves after seawater flows in. It has the potential to achieve power and energy densities of up to 1200 and 250 Wh/kg, respectively [35, 36], but in practice, additional equipment such as the electrolyte tank, a gas separator, and a temperature-control valve are required. Recently, a lithium–seawater battery has been developed by the PolyPlus Battery Co.; it is expected to increase output and energy density by utilizing the high oxidation potential of lithium [37]. Primary seawater batteries were developed for very specific reasons; despite their long history, they remain confined to a narrow range of applications. This is in contrast to the utility of lithium-ion and other types of batteries. Given the variety of devices and applications in the marine field to which the seawater battery was initially applied, as well as the future demand for eco-friendly ships and marine-exploration robots, high-power, high-energy seawater batteries that outperform existing ones will be required. This may entail the construction of a high-utilization secondary seawater battery. Figure 2.3 shows a schematic of the developmental history of primary seawater batteries and their applications.

50

2 Primary Seawater Batteries

2.2 Operating Principles Batteries are devices for storing electrical energy. The redox reactions at the anode and cathode (or rather, the two half-reactions in combination), result in the final cell reaction and are thus the most important factors determining the operation of a battery [18, 38, 39]. The energy delivered by a battery can be quantified by the battery’s voltage and capacity (current and usage time). The operating principles of batteries have been covered in detail in numerous places [40–43]; in this section, we will briefly explain the functioning of a general battery before looking at the operating principle of a primary seawater one.

2.2.1 General Principles The electrical pressure or potential difference between two points is defined as the work done in moving a unit charge from one of the points to the other. In a battery, this is the difference in electrical potential between the two electrodes; it causes charged particles to flow, establishing an electrical current. The cell potential is the difference between the cathode and anode potentials: ◦ ◦ = E Cathode − E ◦Anode E Cell

(2.1)

Numerous metals and compounds exhibit useful redox properties; thus, there various combinations of materials can be used in batteries to produce the desired voltage. Once the cathode and anode materials have been chosen, the theoretical cell voltage can be obtained from their redox potentials. For example, in a Zn/Cu pair, the overall cell potential can be calculated from the redox potentials of Zn and Cu. Zinc oxidizes due to its lower reduction potential, whereas copper reduces due to its higher reduction potential: Anode: Z n 2+ + 2e− → Z n, Er◦ed = −0.76V

(2.2)

Cathode: Cu → Cu 2+ + 2e− , Er◦ed = 0.34V

(2.3)

Therefore, in a galvanic cell containing these metals, zinc is the anode and copper the cathode. The cell has a theoretical voltage of 1.1 V: ◦ = 1.1V Overall: Z n 2+ + Cu → Z n + Cu 2+ , E cell

(2.4)

The voltage of an actual battery may differ from the theoretical voltage owing to the presence of resistive elements. Furthermore, the cell voltage is influenced by temperature and humidity. In a typical electrochemical cell arrangement, the anode is constructed from a metal that oxidizes at a voltage that is between 0.5 and 4 V higher than that of the cathode.

2.2 Operating Principles

51

Another factor that determines the amount of energy stored in a cell is the charge. The Coulomb (C), the SI unit of charge, is the charge of approximately 6.25 × 1018 electrons or protons. Conversely, the charge of an electron or proton is 1.60 × 10–19 C. The faraday, the charge of one mole of electrons, is 96,485 coulombs. The amount of charge flowing per unit time is called the current; its unit is the ampere (A). In battery applications, charge is traditionally expressed not in C but in ampere-hours (Ah). The battery capacity is the total charge (in Ah) that flows from the battery under given conditions. Theoretically, the battery capacity depends only on the cathode and anode materials: a cell constructed using zinc and copper electrodes requires 32.7 g of zinc to release one mole of electrons and 31.8 g of copper to accept them. (Actual electrode reactions may include reactions involving hydrogen and oxygen and electrolytes as well as zinc and copper.) The energy stored in a battery is the product of the voltage and capacity of the battery, i.e., voltage × current × time. This is equivalent to power × time, so the conventional unit of battery energy is the watt-hour (Wh). As previously stated, a battery’s energy and power are largely governed by the materials that comprise its electrodes. Thus, the development of batteries capable of storing greater amounts of energy and delivering higher power outputs is closely related to the development of electrode materials; it necessitates the careful design of reactions and side reactions based on the selection of the electrode materials. This is why research on each component of the battery has been ongoing for decades. The cathode is generally composed of a metal oxide, sulfide, or chloride, which receives electrons along with ions and is therefore reduced. To convert the energy of the anode–cathode reaction into electrical energy, a closed electrical circuit connecting the anode and the cathode and an electrolyte in which the two electrodes are immersed are required. The circuit provides a pathway for the movement of electrons; energy can be extracted from their movement. The charge imbalance generated by electron migration must be balanced by ions migrating through the electrolyte, a liquid or solid that contains one or more dissociating chemicals. The structure and components of a Zn/Cu cell are illustrated in Fig. 2.4. Fig. 2.4 Illustration of a typical battery (the Daniel cell) and its key components

52

2 Primary Seawater Batteries

2.2.2 Reactions in a Primary Seawater Battery The primary seawater battery was developed in earnest along with the development and military application of the magnesium–silver chloride battery. The reactions occurring at the two electrodes of a seawater-activated magnesium battery are as follows: 2+ + 2e− Anode: Mg(s) → Mg(aq)

(2.5)

− Cathode: AgCl(s) + e− → Ag(s) + Cl(aq)

(2.6)

2+ Overall: Mg(s) + 2 AgCl(s) → 2 Ag(s) + Mg(aq) + 2Cl −

(2.7)

The silver that is used in the anode is expensive; moreover, its storage capacity is limited. To overcome this shortcoming, various substitutes have been developed. The following are the reactions undergone by some representative cathode materials: PbCl2 cathode: PbCl2 + 2e− → Pb + 2Cl −

(2.8)

Cu2 l2 cathode: Cu2 l2 cathode: Cu 2 I2 + 2e− → 2Cu + 2I −

(2.9)

CuSCN cathode: 2Cu SC N + 2e− → 2Cu + 2SC N −

(2.10)

MnO2 cathode:2Mn O2 + H2 O + 2e− → Mn 2 O3 + 2O H −

(2.11)

Additionally, a battery using the oxygen or hydrogen reaction of seawater has been developed for deep-sea applications or long-term underwater operation. Such a seawater battery uses oxygen as a reactant at the anode. This type of seawater battery is different from those discussed previously that use seawater as the electrolyte and seawater or DO as the active material for the positive electrode. To exploit the oxygen reaction, which is relatively slow compared to the reaction undergone by Ag, more reactive anode materials such as aluminum, zinc, and magnesium have been considered. However, increases of the alkalinity due to cathodic reaction products may result in the precipitation of insoluble magnesium and calcium salts, degrading the battery efficiency. The oxygen-reduction reaction (ORR) and the hydrogen-evolution reaction (HER) generally occur as follows in seawater of pH 8: O2 reduction at cathode : O2 + 2H2 O + 4e− → 4O H −

(2.12)

H2 reduction at cathode : 2H2 O + 2e− → H2 + 2O H −

(2.13)

2.2 Operating Principles

53

Fig. 2.5 Illustration of Mg/H2 O battery structure. Reproduced with permission from Ref. [144]

Seawater is an excellent electrolyte for batteries, assuming it is sufficiently saline. However, when in contact with the electrode, it accelerates self-discharge. This is also why, in contrast to existing primary batteries, seawater batteries are designed to be reserve batteries. The majority of metals used as anodes in seawater batteries, including magnesium, exhibit side reactions, sometimes including corrosion, when exposed to an aqueous environment containing chlorine or other ions. For example, when a cathode such as magnesium is exposed to seawater, the following hydrogengeneration reaction may occur between the anode substrate and the seawater, resulting in severe self-discharge: Self−discharge reaction at anode: Mg + 2H2 O → Mg(O H )2 + H2

(2.14)

In general, these side reactions cause energy loss and unstable voltage reformation in the cell. However, immersion-type seawater batteries remove insoluble magnesium hydroxide using the generated hydrogen. This improves their performance in the event of battery failure by mitigating the obstruction of ion flow due to magnesium hydroxide deposits (Fig. 2.5).

2.2.3 Features of Primary Seawater Batteries The primary seawater battery is purpose-built. As previously stated, it is equipped with a reserve battery that has a longer life and a higher output than standard batteries.

54

2 Primary Seawater Batteries

Consequently, the battery’s electrode material is selected to have a high discharge voltage and a high energy density. Additionally, it needs to remain stable in air to maintain a long shelf life. Because simplicity of use may conflict with other criteria, the choice of electrode materials for primary seawater batteries is limited [6, 21]. Several metals can be used as anodes in seawater batteries, as listed in Table 2.3. The density, redox potential, and theoretical specific capacity of each metal is provided alongside its listing. The most common anode material for primary seawater batteries is magnesium; magnesium anodes generate a fairly stable output when seawater or saltwater is employed as the electrolyte [25]. Other materials exhibit various shortcomings. Although aluminum has a lower reduction potential and a large theoretical ability to enhance the energy density, it is difficult to achieve a consistent output at high temperatures using an aluminum anode. Because lead anodes require an acidic environment to function, lead reserve batteries require injecting a stored electrolyte rather than immediately injecting saltwater. Alkali metals such as lithium, sodium, and potassium react rapidly with water and oxygen. A non-aqueous electrolyte has been developed for use in lieu of an aqueous electrolyte, resulting in the development of a new type of battery. A reduction reaction occurs at the cathode of a primary seawater battery. Metal-ion precipitation, chloride or oxide reactions, or electrolytic water- or DO-reduction reactions may occur. The cathode reaction of the most widely known primary seawater battery is a silver chloride reaction that results in the reduction of the silver chloride electrode and the release of chlorine ions. Silver chloride has excellent physical Table 2.3 Properties of different metal anodes used in seawater batteries. Reproduced with permission from Ref. [45] Metal anode

Density (g/cc)

Redox potential (Eν · (V)) @ 25 °C

Theoretical capacity (mAh/g)

Environment compatibility

Mg

1.738

−2.363

2200

Neutral, non-aqueous solvents, fused salts

Zn

7.13

−0.76

820

Mildly acidic, neutral, alkaline, and non-aqueous solvents

Al

2.702

−1.66

2980

Neutral, alkaline, non-aqueous solvents, and fused salts

−0.13

260

Acidic

Li

0.534

−3.04

3860

Non-aqueous solvents and fused salts

Na

0.971

−2.71

1166

Non-aqueous solvents and fused salts

K

0.862

−2.92

685

Non-aqueous solvents and fused salts

Pb

11.35

2.2 Operating Principles

55

stability and can be safely stored for an extended period of time prior to use. Its high power and energy densities make it ideal for reserve-type batteries; it is instantly activated when used, even after prolonged storage, and exhibits stable voltage characteristics. However, the battery is expensive, because it contains silver, and the battery, once activated, discharges rapidly, making it suitable only for a single use. Various replacements for the expensive silver in the anode have been investigated [27–30]. Other metal chlorides, such as CuCl and PbCl2 , as well as halogenated compounds such as CuI and oxides such as MnO2 , have been used [27, 29, 46, 47]. As stated previously, oxygen-based processes have also been utilized [31, 48]. However, the physical stability and output power of silver chloride remain unmatched by these non-silver anodes. Table 2.4 shows the advantages and disadvantages of AgCl electrodes compared to others. Primary seawater batteries are designed for irregular use: they are meant to produce high output for a short period of time and to have a long inactive lifespan. Seawater, the electrolyte, is not supplied until the battery is used. This design is suitable for many military and aerospace applications. Product specifications for seawater reserve batteries include voltage, capacity, reserve capacity, energy density, operating temperature, and terminal connections. As shown in Table 2.5, the shelf life of Mg-AgCl primary seawater battery is superior to that of other types of batteries. The battery’s missing components can be added in a variety of ways prior to use. Water or electrolyte can be manually introduced to the battery, and the battery can be triggered when the system is submerged in water (e.g., in a sonobuoy). A variety of activation triggers can be used, depending on the cell’s nature and purpose. Primary seawater batteries can also be categorized as immersion-, dunk-, and flowtype batteries, according to the design of the seawater inlet for activation and terminal connection. Table 2.4 Comparison between batteries using silver and non-silver cathodes. Reproduced with permission from Ref. [21]

Advantages

Disadvantages

Silver chloride cathodes Reliable, safe High raw-material costs High power and energy density High rate of self-discharge Good response to pulse after activation loading Instantaneous activation Long inactivate shelf life No maintenance Non-silver cathodes Abundant domestic supply Low raw-material cost Instantaneous activation Reliable, safe Long inactivated shelf life No maintenance

Require supporting conductive grid Operate at low current densities Low energy density compared to silver High rate of self-discharge after activation

56 Table 2.5 Representative primary and secondary batteries with their temperatures and shelf lives. Reproduced with permission from Ref. [48, 49]

2 Primary Seawater Batteries Battery type

Chemistry

Temperature

Shelf life in years

Primary

Mg-AgCl (reserve)

−60–65 °C

Semi-infinite (5 year warranty)

Alkaline

−18–55 °C

5–10

Zn-carbon

−18–55 °C

3–5

Li (Primary)

−40–60 °C

10–12

Lead-acid

−20–50 °C

0.5

Ni–Cd

−20–65 °C

1.5–3

Ni-MH

−20–65 °C

3–5

Li (Secondary)

−20–60 °C

2–4

Secondary

2.3 Materials As indicated in Table 2.3, the anode of a primary seawater battery can be fabricated from a variety of metals. Seawater batteries based on magnesium, zinc, aluminum, and lithium have been studied, with magnesium-based batteries being the most widely commercialized. Various metal chlorides, notably AgCl, have been investigated as cathodes [50]. Silver chloride can be melted and cast into ingots, as well as rolled into sheet stock having a thickness of 0.08 mm or greater. Due to the malleability and ductility of this material, it can be employed in practically any configuration. Silver chloride is nonconductive until it is immersed; then, the surface is reduced to silver. There is no reason to utilize a default grid for silver chloride. Whereas non-silver cathodes are typically prismatic and flat in shape, silver chloride cathodes are fabricated in a variety of forms, both flat and wavy. Despite the advantages of AgCl, a variety of other anodes are being developed to circumvent the high price of silver [27, 29, 46, 47]. Copper iodide, copper thiocyanide, lead chloride, and copper chloride anodes have been investigated. Seawater-activated batteries are designed to work in an infinite electrolyte: the world’s seas. However, using ocean water for design, development, and quality control purposes is impractical. Instead, the industry standard is to use simulated ocean water, a commercial product that contains all of the required ingredients. When the temperature is above freezing, dunk-type batteries can operate on water or seawater. They are activated by pouring the electrolyte into the battery, where it is absorbed by the separator. Special electrolytes can be utilized at lower temperatures. The use of a conductive aqueous electrolyte results in a more rapid accumulation of voltage. However, adding salts to the electrolyte accelerates the rate of self-discharge. Table 2.6 summarizes the various cell electrodes studied in seawater or saltwater, their reactions, voltages, and energies.

Zn

ZnS O4 (aq)

alkaline

ZnS O4 /MnS O4 (aq)

Mn O2

AgO

α − MnO2

1.65

Mg + 2H2 O → Mg(OH)2 + H2 Zn + 2OH− → Zn(OH)2 + 2e−

SW

SW(H2 )

0.9–1.7

1.0–1.8

6MnO2 + 3Zn + H2 O + ZnSO4 → 1.0–1.8 6MnOOH + ZnSOS4 [Zn(OH2 )]3 H2 O

Zn + AgO + H2 O → Zn(OH)2 + Ag 1.86

2MnO2 + H2 O + 2e− → Mn2 O3 + 2OH−

2MnO2 + H2 O + 2e− → Mn2 O3 + 2OH−

−

Mg + O2 + 2H2 O → 2Mg(OH)2

SW

SW(O2 )

Alkaline

1.34

Mg + 2OH− → Mg(OH)2 + 2e−

Alkaline

Air

Mn O2

3.09

Mg +H2 O2 +2H+ → Mg2+ +2H2 O

SW − NaCl, H2 SO4 , H2 O2

H2 O2

Alkaline

4.14

Mg + 2MnO2 + H2 O → Mn2 O3 + Mg(OH)2

Mg(ClO4 )2

MnO2

Air

1.2 2.8

2.4

Mg + K2 S2 O8 → K2 SO4 + MgSO4 Mg + PbCl2 + 2e− → Pb + MgCl2

1.5

Mg + 2CuCl → MgCl2 + 2Cu

Salt solution

1.6

Mg + 2AgCl → MgCl2 + 2Ag

PbCl2

Salt solution

CuCl

Theoretical voltage (V)

Overall reaction

K2 S2 O8

Salt solution

AgCl

Mg

Electrolyte

Cathode

Anode

1.2–1.35

1.3–1.5

1.3

1.37

1.0–1.2

−

1.0–1.3

1.2–1.4

1.7–1.8

1.9–2.0

0.9–1.05

1.6–2.0

1.2–1.4

1.1–1.5

Operating voltage (V)

Table 2.6 Anode–cathode combinations for primary seawater batteries, and their electrochemical properties

−

447

336

336

1300

−

−

6800

3814

759

−

−

386

Mg

Theoretical specific energy (Wh/Kg)

(continued)

[101–108]

−

[102]

[97–101]

[95, 96]

[31, 93, 94]

[89–92]

−

[83–88]

[81, 82]

[27, 78–80]

[76, 77]

[62–75]

[51–61]

References

2.3 Materials 57

Al

Anode

Liquid

Alkaline,K2 S

KOH

H2 O2

S

Fe(CN)3− 6

NiOOH

Alkaline

MnO2

AgO

Zn(CH3 COO)2 / Zn(CH3 F3 SO3 )2

Na3 V2 (PO4 )3

−

−

Al + 3NiOOH + Al + 3NiOOH + OH− + 3H2 O →Al(OH)− 4 + 3Ni(OH)2

−

Al + 3OH− + 3 2.2 4− → 3Fe(CN) + Al(OH) Fe(CN)3− 3 6 6 −

1.3

1.8

− 2Al + S2− 4 + 4H2 O + 2OH → 4HS− + Al(OH)3

2Al + 3H2 O2 + 2OH− → 2Al(OH)− − 4

2.0

1.2

−

2.7

0.8–1.7

0.75–0.9

1.3

Operating voltage (V)

−

Al + 3MnO2 + 3H2 O → 3MnO2 H + Al(OH)3

2Al + 3AgO + 2OH− + 3H2 O → 2Al(OH)− 4 + 3Ag

NaV2 (PO4 )2 + xZn → Zn6 NaV1 (PO4 )3

V2 O5 + Zn2+ + 2e− + nH2 O → ZnV2 O5 · nH2 O

Zn (CF3 SO3 )2 (aq)

Vx Oy

0.35–1.3

1.0–1.8

Mn2 O3 + Zn → ZnMn2 O3 (x = 2, y = 3)

ZnS O4 /MnS O4 (aq)

Mnx Oy

Theoretical voltage (V)

Overall reaction

Electrolyte

Cathode

Table 2.6 (continued)

−

−

647

−

−

1020

−

322

−

Theoretical specific energy (Wh/Kg)

[133]

[132]

(continued)

[130, 131]

[127–129]

[125, 126]

[122–124]

[44, 119–121]

[113–118]

[109–112]

References

58 2 Primary Seawater Batteries

Li

Anode

SW

SW

SW (H2 )

SW

SW (H2 )

sw(O2 )

SW

sw(O2 )

Alkaline

Alkaline

Air

Air

Electrolyte

Cathode

Table 2.6 (continued)

−

2.5

→ H2

+ 2OH−

2H2

3.8

O2 + 2H2 O + 4e− → 4OH− O + 2e−

3.45

Li + OH− → Li(OH) + e−

→ H2

+ 2OH−

2H2

−

O + 2e−

2.7

O2 + 2H2 O + 4e− → 4OH−

Theoretical voltage (V)

Al + 3OH− → Al(OH)3 + 3e−

Overall reaction

−

−

2.4

−

−

1.2–1.6

Operating voltage (V)

−

−

13,300

−

−

8100

Theoretical specific energy (Wh/Kg)

−

[138, 139]

[137]

[136]

[134, 135]

−

References

2.3 Materials 59

60

2 Primary Seawater Batteries

Separators, cell-to-cell connections, and termination configurations are required to manufacture a battery. Separators made of non-conductive fabrics, non-woven fabrics, and polymers allow free entry of electrolytes and escape of corrosion products. When multiple single cells are stacked, cell-to-cell connection (which depends on the voltage and capacity) is required. With silver electrodes, silver foil is used; nonsilver electrodes are connected through an insulator via copper or other conducting staples. Terminations for connection to wires are joined to the positive poles by direct soldering or welding. Seawater enters through an electrolyte inlet in the final casing; gases and reactants exit through another hole on the opposite side. Cell performance can be maximized when the other parts are sufficiently insulated when exposed to seawater [21].

2.3.1 Magnesium Battery Magnesium is an alkaline-earth metal possessing exceptional specific strength and stiffness. Its density is approximately two-thirds that of aluminum, making it the lightest of all commonly used metals [140]. Magnesium typically exists in its oxidized state in nature as divalent ions. It reacts with air, but rapidly forms a thin superficial layer of oxide to protect the bulk from further oxidation [141]. Because of the these properties of magnesium alloys, magnesium can be used as a battery electrode [142]. Primary seawater batteries have been developed to alleviate the dissolution of magnesium, and the discharge performance of magnesium anodes has been improved through alloy research and advancements in manufacturing procedures. The half-reaction at the magnesium anode is Anode: Mg → Mg 2+ + 2e− , E ◦ = −2.363V

(2.15)

Magnesium and its cation serve respectively as the anode and active charge transport agent in a magnesium battery. The chemistries of non-rechargeable (primary) and rechargeable (secondary) cells have been investigated. Magnesium primary cell batteries are employed as reserve and general-purpose batteries [143, 144]. Silver chloride, copper chloride, copper iodide, copper thiocyanate, manganese dioxide, and air (oxygen) have been explored as cathodes for reserve battery chemistries [27, 29, 46, 47]. By 1943, a commercially usable water-activated magnesium reserve battery was available. The magnesium/silver chloride (Mg/AgCl) system was used in the first commercially produced seawater-activated battery, and it was utilized in batterypowered torpedoes. Today, small devices such as sonobuoys, radio buoys, life-saving equipment, nautical markers, and emergency lights are frequently powered by these batteries. However, because silver compounds are used in the cathode, these batteries are quite expensive. To lower costs, water-activated batteries using low-cost cathode materials have been utilized despite their inferior performance. The electrochemical couplings Mg/PbCl2 , Mg/CuCl, and Mg/K2 S2 O8 are currently used in wateractivated batteries. Magnesium–cuprous chloride batteries, are dunk-type batteries,

2.3 Materials

61

in which the separator placed between the electrodes absorbs and retains the electrolyte [21]. Seawater-activated batteries using copper iodide–sulfur, copper sulfate, and manganese dioxide cathodes are being developed, although none of these can rival the magnesium–silver chloride battery in terms of performance. The dissolved-oxygen–magnesium battery was developed for extended use in buoys and autonomous underwater vehicles [31]. Theoretically, a magnesium–air fuel cell operates at 3.1 V and has an energy density of 6.8 kWh/kg. Oxygen dissolved in saltwater serves as the cathode reactant, while magnesium serves as the anode. Magnesium dry-batteries of type BA-4386 have been fully commercialized, with costs per unit approaching those of zinc batteries; they have higher volumetric capacities and longer shelf lives than identical zinc-carbon cells. From 1968 to approximately 1984, the BA-4386 was widely employed by the US military, although it was later superseded by the lithium–thionyl chloride battery [145]. As early as the 1960s, General Electric developed a magnesium–air fuel cell that operated in a neutral NaCl solution. Although the magnesium–air battery is a primary cell, it can be refueled by replacing the anode and the electrolyte. These batteries have been commercialized and are being used as land-based backup systems as well as undersea power sources [146]. The Mark 44 torpedo is powered by a magnesium battery that is ignited by water [147]. The development of aqueous magnesium-ion batteries (MIBs) has been hampered by the limited number of cathode materials available for this system that do not suffer from the sluggish diffusion of Mg2+ into the host lattices. The passivation of the magnesium anode significantly also inhibits further transit of Mg2+ ions, necessitating the use of a passivation-resistant anode material [148]. Consequently, magnesium alloys are widely employed as anodes in seawater-activated battery systems [149, 150]. Research on magnesium alloys suitable for fabricating commercial electrodes from has been conducted aggressively. A higher discharge voltage than that of the pristine magnesium electrode can be developed by modulating the response of the magnesium electrode by the introduction of other elements. For example, the alloy AZ61A creates less sludge and has a higher discharge voltage than alloy AZ31B [151]. AP65, developed later, has a slightly higher discharge voltage (~0.1–0.3 V), but suffers from excessive sludge generation [152]. However, employing AP65 in force-flow-type batteries can partially resolve the sludge-generation problem. Table 2.7 details the chemical compositions and ratios of the constituent elements in these alloys [21].

2.3.1.1

Mg/AgCl Battery

Mg metal, AgCl, and seawater serve as the anode, cathode, and electrolyte, respectively, in this battery system. A magnesium alloy of aluminum, zinc, and manganese is the most frequently used anode material, because of its resistance to corrosion in seawater. Although silver chloride is an insulator, it becomes a conductor when a silver coating is formed on the surface of the electrode [153].

62

2 Primary Seawater Batteries

Table 2.7 Composition of Mg alloys used in electrodes (in weight %) Reproduced with permission from Ref. [21] Element

AZ31

AZ61

AP65

MELMAG75

Al

2.5–3.5

5.8–7.2

6.0–6.7

4.6–5.6

Zn

0.6–1.4

0.4–1.5

0.5–1.5

0–0.3

Pb

−

−

4.4–5.0

−

Ti

−

−

−

6.6–7.6

Mn

0.15–0.7

0.15–0.25

0.15–0.30

0–0.25

Si

0–0.1

0–0.05

0–0.3

0–0.3

Ca

0–0.04

0–0.3

0–0.3

0.3

Cu

0–0.05

0.05

0.005

−

Ni

0–0.005

0–0.005

0–0.005

0–0.005

Fe

0–0.006

0–0.006

0–0.01

0–0.006

The following electrochemical reaction occurs at the cathode during discharge: Cathode: 2 AgCl + 2e− → 2 Ag + 2Cl − , E ◦ = 0.222V

(2.16)

Combining Eqs. 2.15 and 2.16, the overall reaction undergone by the Mg-AgCl system during discharge is: Mg + 2 AgCl → MgCl2 + 2 Ag, E cell = 2.585V

(2.17)

The theoretical cell voltage of the Mg/AgCl battery is 2.585 V. However, because of wasteful corrosion of the Mg anode, the actual cell-discharge voltage is lower (~1.6 V). Although these batteries are designed to maintain high current densities, they discharge rapidly upon activation, in addition to having high raw material costs. They are stored in a dry condition until immediately before use, and generally start working within seconds to tens of seconds when activated by seawater (Fig. 2.6). Fig. 2.6 The structure of an Mg/AgCl seawater battery. Reproduced with permission from Ref. [154]

2.3 Materials

2.3.1.2

63

Mg/CuCl Battery

The widely used Mg/CuCl battery became commercially available in 1949 [28]. It is significantly less expensive and more reliable than the Mg/AgCl battery, activates instantly, and requires no maintenance, although it has a lower energy density, a lower rate capability, and a high rate of self-discharge following activation. This system must be fully sealed during storage, because CuCl is strongly hygroscopic. Consequently, the seal that prevents seawater from entering the cell needs to be broken prior to use. Generally, the Mg/CuCl system is interchangeable with the Mg/AgCl system. During discharge, the electrochemical reaction occurring at the CuCl cathode is Cathode: 2CuCl + 2e− → 2Cu + 2Cl − , E ◦ = 0.124V

(2.18)

The overall reaction is Mg + 2CuCl → MgCl2 + 2Cu, E cell = 2.487V

(2.19)

Although the theoretical cell voltage of the Mg/CuCl battery is 2.487 V, the practical open circuit voltage ranges between 1.7 and 1.8 V.

2.3.1.3

Mg/PbCl2 Battery

Mg/PbCl2 batteries are also intended as replacements for Mg/AgCl batteries [47]. While copper chloride is strongly hygroscopic, lead chloride is moderately stable in air and soluble in seawater. The reduction potential of the PbCl2 electrode is lower than that of the CuCl electrode. The discharge reaction is Cathode: PbCl2 + 2e− → Pb + 2Cl − , E ◦ = −0.268V

(2.20)

The overall reaction follows from Eqs. 2.15 and 2.20: Mg + PbCl2 → MgCl2 + Pb, E cell = 2.095V

(2.21)

This battery has an open circuit voltage of 1.2 V but a load voltage of 0.9–1.06 V, which is lower than that of Mg/AgCl and Mg/CuCl batteries. Consequently, multiple cells need to be connected in series to form a high voltage battery system. PbCl2 can be dissolved in water, releasing chlorine ions; as this establishes ionic conductivity, the battery may be used in both freshwater and seawater. Caution should be exercised, however, because lead is more harmful to the environment than silver or copper.

64

2.3.1.4

2 Primary Seawater Batteries

Mg/Seawater Battery

In the magnesium–seawater cell, the anode reaction involves either water or oxygen, and seawater serves as the source of DO or water for the anode. Mg/seawater cells were created in the 1990s as high-energy–density cells for long-term applications such as deep-sea drilling. Mg/seawater batteries have two modes of operation. One employs the ORR for water molecules; the cathode acts as the current collector in seawater utilizing DO as the cathode oxidizer. The other employs the HER through the consumption of seawater-derived water molecules as a cathodic oxidizer (Figs. 2.7 and 2.8) [21, 31]. Fig. 2.7 The structure of a magnesium–seawater battery. Reproduced with permission from Ref. [21]

Fig. 2.8 The structure of a magnesium–air battery. Reproduced with permission from Ref. [21]

2.3 Materials

65

Mg/seawater Battery Based on ORR DO is one of the most electrically active molecules in seawater. Consequently, when a suitable positive electrode current collector is employed to discharge an Mg/seawater battery, the DO supplied by seawater can be reduced on the surface of the cathode current collector [31]. Since the product of the ORR is the hydroxide ion, the reduction potential and rate are also influenced by the pH. At the cathode, the electrochemical reaction occurring in common seawater (pH = 8) is Cathode: O2 + 2H2 O + 4e− → 4O H − , E ◦ = 0.77V

(2.22)

The overall reaction follows from Eqs. 2.15 and 2.22: 2Mg + O2 + 2H2 O → 2Mg(O H )2 , E ◦ = 3.133V

(2.23)

Despite the high theoretical voltage, this battery operates at 1–1.3 V. This suggests that the ORR is kinetically much more limited than the metal chloride reduction reaction: i.e., because the diffusion rate of DO into the cathode current collector is limited and the ORR kinetics are sluggish, the actual cell voltage is substantially lower than the theoretical value. In general, oxygen has a concentration of approximately 0.5 mol/m3 and low solubility, and the ORR proceeds through a sluggish four-electron mechanism. These constraints predict a low current density during operation or a reduction in the power output of the battery system. To overcome these shortcomings, multiple techniques are used to enhance the current or power output while using largearea cathode current collectors instead of magnesium anodes. However, ORR-based Mg/seawater batteries are frequently used for low-current operation. Mg/Seawater Battery Based on HER In some Mg/seawater batteries, H2 O is directly reduced with high current mode [31]. This theoretically permits operating at a lower voltage than the oxygen reaction, but without the challenges associated with slow oxygen diffusion and DO concentration constraints. When a magnesium–seawater battery utilizes water instead of oxygen as the oxidizing agent, the water accepts electrons to generate hydroxide ions and hydrogen gas. The principle is identical to that of the HER, and this reaction is also influenced by the pH. The electrochemical reaction occurring in seawater (pH = 8) is Cathode: 2H2 O + 2e− → H2 + 2O H − , E ◦ = 0.47V

(2.24)

The overall reaction follows from Eqs. 2.15 and 2.24: Mg + 2H2 O → Mg(O H )2 + H2 , E ◦ = 1.893V

(2.25)

The HER mode of operation has faster kinetic characteristics than the ORR mode and can improve the current density. The gain on the output side, however, may fall when the reaction voltage decreases. The HER-mode Mg/seawater battery’s working

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2 Primary Seawater Batteries

Fig. 2.9 Structure of Mg/H2 O2 seawater battery. Reproduced with permission from Ref. [156]

cell voltage is extremely low (~0.4–0.5 V), resulting in a reduced energy density. In addition, the hydrogen bubbles generated during the HER may impede the movement of ions between electrodes in the cell or may be a limiting factor for military use.

2.3.1.5

Mg/H2 O2 Battery

A cell having a magnesium anode and a hydrogen peroxide solution-phase cathode (Mg/H2 O2 ) in a seawater electrolyte exhibits semi-fuel-cell characteristics. Such cells typically have voltages in the range of 1.0–1.8 V and are intended to function at low current densities (10–50 mA/cm2 ). They were developed for long-lasting underwater maritime applications [31, 155]. Catalytic electrodes are occasionally used to enhance the reduction of hydrogen peroxide [156]. Figure 2.9 shows the structure of the Mg/H2 O2 battery. The cathodic reaction is Cathode: H O2− + H2 O + 2e− → 3O H − , E ◦ = 0.88V

(2.26)

The overall reaction follows from Eqs. 2.15 and 2.26: Mg + H O2− + H2 O → Mg 2+ + 3O H − , E ◦ = 3.58V

2.3.1.6

(2.27)

Additives for Mg Batteries

By supplementing a magnesium–saltwater battery with AgCl and Ag2 S utilizing CuI as the anode, Renuka [29] was able to enhance the performance of the copper iodide electrode to a certain extent, demonstrating that battery performance can be improved by a trace amount of silver. The objective was to maximize the output of silver chloride or silver sulfide while maintaining the competitiveness of copper

2.3 Materials

67

Table 2.8 Possible cathode compositions with additives for magnesium–seawater batteries (in weight %). Reproduced with permission from Ref. [21] Silver chloride

Cuprous iodide

Cuprous thiocyanide

Lead chloride

Cuprous chloride

Depolarizer

100

73

75–78

80.7–82.5

95–100

Sulfur

−

20

10–12

−

−

Additive

−

−

0–4

2.3–4.4

−

Carbon

−

7

7–10

9.6–9.8

−

Binder

−

−

0–2

1.5–1.6

0–5

Wax

−

−

−

3.8

−

iodide. These additives may help reduce the internal resistance and enhance cell conductivity to a certain extent. Sulfur has been proposed as an additive to the Cu2 O electrode (using Ni foam as the substrate) in an Mg/Cu2 O seawater battery [157]. Sulfur can participate in the electrochemical reaction at the cathode, and adding it enhances the cell voltage by 200–350 mV, depending on the ratio of Cu2 O to S. This is the result of confirming the electrochemical performance of sulfur as an addition in electrodes that have been appropriately built. The Ni foam support electrode possesses electrochemical characteristics compatible with seawater batteries with sulfur. A depolarizer (normally cathodic active material, used to remove gas collected at the electrodes) and a current collector comprise the negative electrode. [21] Depolarizers are powder-based and do not conduct electricity; to make the depolarizer function, a special type of carbon must be added to increase its conductivity. A binder is also added for cohesion and the negative electrode is placed on a metal grid to permit cell-to-cell contact and battery termination. Additional experiments have been conducted to improve the stability of the anode output by adding a trace amount of silver, sulfur, carbon, binder, or wax. Table 2.8 lists possible cathode compositions.

2.3.2 Zinc Battery Owing to its large theoretical capacity, low redox potential, low cost, and excellent safety, zinc has been regarded as an ideal anode in a variety of primary and secondary batteries [158–160]. Zinc has been widely utilized as a general-purpose primarybattery anode since the invention of zinc–manganese dioxide batteries. Because of its longer lifetime and lower cost in intermittent usage compared to alkaline batteries, it is widely employed in a variety of primary batteries. Zinc is oxidized in a humid environment; however, the zinc carbonate (Zn5 (OH)6 (CO3 )2 ) coating formed when CO2 in the air reacts with zinc prevents further internal corrosion [161]. The anode reaction is

68

2 Primary Seawater Batteries

Anode: Z n → Z n 2+ + 2e− , E ◦ = −0.76V

(2.28)

Zinc is an easily obtainable and processable metal that has been employed as an electrode material since the infancy of battery development because of its strong tendency to ionize. The Volta pile and the Leclanché cell used zinc electrodes. Since then, dry batteries, alkaline batteries, zinc-mercury batteries, zinc-silver oxide batteries, etc. have utilized Zn as the negative electrode [162]. Recently, as energy density expectations and application requirements have increased, research on secondary-battery technologies such as zinc–oxygen batteries has accelerated [33]. Despite the development of a zinc-based aqueous oxygen battery, the reactivity and reaction by-products of zinc have limited the metal’s use in seawater batteries [163]. Zinc batteries with basic solutions have been introduced instead.

2.3.2.1

Zn/MnO2

The zinc–MnO2 or zinc–carbon primary battery generates direct current electrochemically through the interaction of zinc and manganese dioxide (MnO2 ). At the MnO2 cathode, the half-reaction is Cathode: 2Mn O2 + H2 O + 2e− → Mn 2 O3 + 2O H −

(2.29)

The overall reaction follows from Eqs. 2.28 and 2.29: Z n + 2Mn O2 + H2 O → Z n 2+ + Mn 2 O3 + 2O H −

(2.30)

Typically, a Zn/MnO2 cell produces approximately 1.5 V. Manganese batteries are highly stable, simple to make, and inexpensive. The current zinc–manganese oxide battery dates to 1868, when the Leclanché cell was invented [12]. Originally, it was a wet battery with an aqueous (ammonium chloride) electrolyte, but later, the electrolyte was hardened with gypsum and other materials. Since then, potassium hydroxide solution has been utilized as the electrolyte; the resulting alkali batteries have longer lives than manganese batteries, and are increasingly replacing them. They are useful for low-drain or intermittent applications such as remote controls, flashlights, watches, transistor radios, and other portable devices. Alkaline secondary batteries are also the subject of much recent research and development (Fig. 2.10).

2.3.2.2

Zn/Ag2 O

Zinc–silver oxide batteries have high energy and power densities [164]. Furthermore, when the Ag2 O constituting the anode is reduced to metallic Ag, it does not generate an intermediate product and delivers a stable voltage. Because of their high production cost, these batteries have mainly been used in special fields such as aerospace and weapons engineering. A button-type or ultra-thin small primary battery with

2.3 Materials

69

Fig. 2.10 Structure of a Zn/MnO2 battery

Ag2 O has been put into more general practical use in devices such as hearing aids [153]. The cathode half-reaction is Cathode: Ag2 O + 2H + + 2e− → 2 Ag + H2 O, E ◦ = 1.17V

(2.31)

The overall reaction follows from Eqs. 2.29 and 2.31: Z n + Ag2 O + H2 O → Z n 2+ + 2 Ag + 2O H − , E ◦ = 1.93V

(2.32)

A battery based on silver peroxide (AgO) was developed to enhance the energy density of silver oxide [165]. Silver peroxide batteries may have higher energy densities; however, they are unstable in alkaline electrolytes. Consequently, they are difficult to handle, and high energy densities cannot be realized in small batteries, limiting their practical application. Even at high loads, the Zn/Ag2 O battery maintains consistent voltage characteristics. It is ideally suited for devices that require a large power output over a short period of time, such as torpedoes, and plays a critical role in military applications [1]. The high-power silver oxide battery is a reserve seawater battery that is stored separately and injected into a dry battery when activated by an electrical stimulus. In addition to the silver oxide electrode’s exceptional high-rate discharge properties, it has a long shelf life due to the reserve-type design. Typically, zinc electrodes in zinc–silver oxide batteries are designed to have a high porosity and reaction area, allowing for simultaneous or rapid applications of loads. However, during a highpower discharge, a significant volume of hydrogen gas may be generated; to minimize this, a high-overvoltage substance such as mercury may be added (Fig. 2.11).

70

2 Primary Seawater Batteries

Fig. 2.11 Structure of a Zn/Ag2 O battery, the LR700(DS) cell. Reproduced with permission from Ref. [166, 167]

2.3.3 Aluminum Battery Aluminum is remarkably light, inexpensive, and nontoxic. The oxidation–reduction couple of aluminum can release three electrons; therefore, it has very high volumetric (8040 mAh/cm3 ) and gravimetric (2980 mAh/g) capacities. Furthermore, Al is a highly recyclable metal. Because of these advantages, research and development on aqueous aluminum-ion batteries (AIBs) using aluminum anodes has intensified. The anode reaction and reduction potential of aluminum are [168] Anode: Al → Al 3+ + 3e− , E ◦ = −1.66V

(2.33)

Aluminum batteries, despite their advantages, are only available commercially for a limited range of uses. This is because an aqueous electrolyte (4 ≤ pH ≤ 8) forms a protective oxide (Al2 O3 ) coating on the surface of aluminum, preventing a stable reaction [169, 170]. Oxide production results in decreased cell efficiency,

2.3 Materials

71

decreased electrode potential, and non-uniform aluminum corrosion. In basic solutions, aluminum corrodes into aluminum hydroxide (Al(OH)3 or Al(OH)4 – ). This results in a non-uniform electrochemical reaction and a decrease in power generation. Due to these issues, water-based aluminum batteries remain experimental, with only a few developed for key high-energy systems. Aluminum–water (Al/H2 O) energy technology has been developed by the company L3Harris for equipment operating on the seabed [171]. Because it is aluminum-based, this battery system is safe and scalable and has very high energy density and low toxicity; thus, it is expected to prove suitable for a variety of marine equipment, including unmanned underwater vehicles (UUVs), sensors, and navigation systems. Thanks to their increased energy density, aluminum-water cells could be especially useful for challenging missions such as submarine communications, mine hunting, and pipeline surveying. Aluminum–air batteries consist of an aluminum-metal cathode, an anode that enables oxygen transport and reduction, and an alkaline solution consisting of a suitable electrolyte, usually sodium hydroxide (NaOH), potassium hydroxide (KOH), or sodium. Research on aluminum–air batteries has been actively conducted since 2010, and many countries, including China, Korea, and countries in North America, have developed and secured patents for the technology [169]. In 2017, Frost Sullivan’s TechVision Division declared that aluminum-ion batteries were one of the potential replacements for lithium batteries [172]. Recently, a solid fiber aluminum–air battery was reported [173]. Using carbon nanotubes with a cross-stacked structure, this battery demonstrated a specific energy of 1,168 Wh/kg. Nonetheless, these batteries suffer from numerous issues, including aluminum corrosion, high self-discharge rates, sluggish discharge kinetics, a fundamental lack of rechargeability, and limited shelf life (Figs. 2.12 and 2.13).

Fig. 2.12 Conceptual illustration of an aluminum–water battery. Reproduced with permission from Ref. [171]

72

2 Primary Seawater Batteries

Fig. 2.13 Conceptual illustration of an aluminum–air battery. Reproduced with permission from Ref. [32]

2.3.4 Li Battery The metal anodes of Mg, Mg alloy, Zn, and Al primary seawater batteries have low redox potentials, so the output voltages remain low (1–1.8 V). By contrast, lithium has a high oxidation potential and a high capacity (3860 mAh/g); therefore, a lithiumbased battery system is expected to have a high energy density. The lithium–seawater battery, like the magnesium–seawater battery, can operate via two cathodic reactions: ORR and HER. The theoretical voltages of ORR- and HER-based lithium-seawater batteries are 3.8 and 2.5 V (vs. Li/Li+ ), respectively, at a pH of 8; they operate at 3 and 2.3 V (vs. Li/Li+ ), respectively. The anode half-reaction is Anode: Li → Li + + e− , E ◦ = −3.05V

(2.34)

2.3 Materials

73

Fig. 2.14 Proposed lithium–seawater battery with protected lithium electrode (PLE). Reproduced with permission from Ref. [37]

However, using an alkali metal as an electrode material is not straightforward. Having a redox potential capable of producing a high voltage also suggests a high reactivity. Because metals such as lithium or sodium are easily oxidized when exposed to oxygen or moisture in the air, they need to be well-protected when used as electrodes. Despite these difficulties, the PolyPlus Battery Co. may soon build the first seawater-activated lithium battery with a protected lithium electrode (PLE); see Fig. 2.14 [174]. Such batteries can use a variety of cathodes to suit the application and marine environment. Utilizing an anode that is powered by DO in saltwater achieves the highest energy density and durability possible, while a solid cathode enables high-power applications. The energy densities of the anode is 300 Wh/kg, while that of the cathode is 400 Wh/kg (600 Wh/l). These batteries are safe, compact, and rarely self-discharge, owing to adequate protection of the lithium.

2.4 Cell Engineering 2.4.1 General Construction As discussed previously, seawater batteries are reserve batteries in which seawater (or DO) is isolated from the electrolyte or anode active material. Each cell is comprised of a positive electrode, a negative electrode, a separator, a termination, and a casing. A primary battery is formed by connecting multiple cells in series or parallel to maximize the output or obtain sufficient energy density to balance the conflicting properties of high output and long shelf life. To enable rapid activation of this system, a series–parallel design is required. When connecting batteries in series to increase the voltage in a marine environment, the separator must be carefully constructed.

74

2 Primary Seawater Batteries

Fig. 2.15 General cell structure of a primary seawater battery: cross-sectional view. Reproduced with permission from Ref. [21]

Additionally, battery capacity (i.e., the amount of active material between the negative and positive electrodes) and other factors must be addressed in order to achieve a sufficient running time for the intended application. The reaction area of the electrode must be increased in proportion to the current required for operation. In actual use, the power output of the battery depends on the temperature and salinity of the electrolyte, which, in the case of seawater, often decrease at increased depth. Optimization of the seawater intake design is also necessary. Primary seawater batteries can be divided into immersion type, forced flow type, and dunk-type batteries according to the design of the seawater inflow system. The general unit cell structure of primary seawater batteries considering cell-stacking is illustrated in Figs. 2.15 and 2.16.

2.4.2 Types of Primary Seawater Batteries There are three broad types of primary saltwater reserve battery, depending on the seawater inflow design: immersion type, forced flow type, and dunk type. Immersion batteries are designed to be activated by a natural influx of seawater when immersed in the electrolyte. As illustrated in Fig. 2.17, they are constructed in such a way that seawater enters through the bottom vent and internal air departs through the top vent. These batteries are designed to create target voltage by cell-stacking in devices such as sonobuoys.

2.4 Cell Engineering

75

Fig. 2.16 General cell structure of primary seawater batteries: a silver; b non-silver. Reproduced with permission from Ref. [21]

A forced-flow battery (Fig. 2.18) is one in which the electrolyte is introduced by hydraulic pressure or a similar external stimulus. This is effective when a device apable of causing seawater flow is used, such as an electric torpedo. This type of battery mitigates to some extent the issues associated with by-products and gas formation at the electrodes by circulating the electrolyte. Consequently, it is suitable for generating a high power output. However, additional challenges arise, such as designing the casing and recirculation pumps to optimize seawater flow. Therefore, this type of battery is most suitable for applications requiring high power outputs in systems larger than a threshold size, such as torpedoes. Considering the heat created

76

2 Primary Seawater Batteries

Fig. 2.17 Structure of immersion-type seawater battery: a side view; b front view. Reproduced with permission from Ref. [21]

Fig. 2.18 Schematic representation of torpedo-battery construction: a cell construction; b battery configuration. Reproduced with permission from Ref. [21]

2.4 Cell Engineering

77

Fig. 2.19 Schematic representation of a dunk-type battery. Reproduced with permission from Ref. [21]

during circulation, this type of battery may function at a cathode surface current density of up to 500 mA/cm2 when the torpedo is fired and may have an operational time of several tens of minutes. Dunk-type batteries are charged by pouring electrolyte into the battery. An absorbent separator between the electrodes absorbs and activates the electrolyte. This type of battery is suitable for high-altitude meteorological observation devices such as radiosondes; Mg/CuCl batteries have been used to replace Mg/AgCl batteries that operate at temperatures as low as –50 °C. Their mechanism of operation is identical to that of other reserve batteries, commencing operation within 1–10 min of activation. At room temperature, they have a discharge time of up to 3 h. Figure 2.19 illustrates a dunk-type Mg/CuCl battery.

2.4.3 Applications of Primary Seawater Batteries Although conventional batteries have many maritime applications, it is frequently preferable to use a battery activated by seawater. With sufficient salinity and a high enough temperature, seawater can be an excellent water-based electrolyte for batteries. Special-purpose seawater reserve batteries are suitable for military and communication devices that require long shelf lives, undergo intermittent use, and need to become operational within seconds of submersion. The sonobuoy, introduced by the US Navy in 1942 to track German submarines, is a representative piece of military equipment in which seawater batteries are employed. A sonobuoy is a system for locating enemy vessels by measuring sound waves; it is deployed by being thrown into the sea (generally using a rotochute instead of a parachute) from a ship or aircraft. Upon immersion, it is activated by seawater,

78

2 Primary Seawater Batteries

Fig. 2.20 Primary seawater batteries used for military purposes in the AN/CRT-1 sonobuoy. Reproduced with permission from Ref. [175]

and operates for several hours. Mg/AgCl seawater-activated batteries are typically used in sonobuoys. Early buoys had a battery life of only four hours. However, the design has been modified to include a separate container enclosing an additional battery to extend the use of the sonobuoy. Active sonobuoys employ sonar; passive ones do not. Usually, an active sonobuoy requires more energy, which necessitates a larger-capacity battery (Fig. 2.20). Seawater batteries can be applied to systems requiring high power such as torpedoes. Initially, torpedoes used thermal-energy-based propulsion systems, but the noise and combustion products facilitated tracking. To avoid this, electrical propulsion systems were developed, and the battery has become one of the most important elements of the torpedo. At first, reserve-type lead-acid batteries were used; however, they suffered from low energy density and high energy loss. Zinc–silver oxide and aluminum–silver oxide batteries have been developed and applied to overcome this shortcoming. Magnesium–silver chloride and aluminum–silver oxide seawater batteries are used as light torpedo batteries. In the magnesium batteries used initially, the voltage dropped under load; this problem was partially solved by recirculating the electrolyte. Aluminum–silver oxide batteries have been developed to meet the demand for higher power outputs. Compared to magnesium batteries, aluminum batteries have twice the energy density, as well as longer shelf life. Recently, lithiumbased batteries having even higher energy densities have been attracting attention (Figs. 2.21, 2.22 and 2.23). The majority of seawater batteries contain metal anodes and use seawater as an electrolyte and a metal salt as the cathode. The SWB600 and SWB1200 are batteries developed by Kongsberg Simrad AS, Norway [177, 178], and used in shallow water. These batteries generate energy by the electrolytic dissolution of their magnesiumrod anodes; the cathodes are composed of carbon fibers. These batteries have the

2.4 Cell Engineering

79

Fig. 2.21 Mark 32 torpedo powered by a primary seawater battery. Reproduced with permission from Ref. [34]

80

2 Primary Seawater Batteries

Fig. 2.22 Schematic of Mg/AgCl seawater battery used in torpedoes: (a) cell structure; (b) battery configuration. Reproduced with permission from Ref. [34]

Fig. 2.23 A primary seawater battery used in a torpedo. Reproduced with permission from Ref. [176]

advantages of high energy density, low cost, good safety characteristics, and infinite storage capability. Wilcock and Kaufmann conducted a deep-water test of magnesium–graphite and magnesium–copper batteries at a depth of approximately 2200 m in the NE Pacific [90]. They reported that the output voltage from seawater cells is lower in deep water than in shallow. The SWB600 was improved for applications at depths exceeding

2.4 Cell Engineering

81

6000 m; a single cell of the SWB600 generated an electric power of 2 W at the depth of approximately 700 m. Seawater batteries have also been developed for sea floor observation. In 2000, the WP-2 sea-floor-borehole broadband seismic observatory, which used seawater batteries, was installed in the northwestern Pacific basin during Ocean Drilling Program (ODP) Leg 191 to collect data for a study of the Earth’s interior (Fig. 2.24) [179]. A water-activated lead chloride battery for life-saving equipment was introduced by the GS Yuasa Corp. (Fig. 2.25), although it has been replaced by other batteries and is not currently being produced [180]. L3 OpenWater Power has developed and modularized an aluminum–seawater battery for deep sea applications (Fig. 2.25). These developments can be seen as efforts to produce greater energy and output than in the marine environment with a special environment and purpose (Fig. 2.26) In fact, primary seawater batteries are suitable for many underwater or deep-sea applications including:

Fig. 2.24 Schematic depiction of the WP-2 borehole seismological observatory, powered by seawater-battery cells. Reproduced with permission from Ref. [178]

82

2 Primary Seawater Batteries

Fig. 2.25 Primary seawater battery (Mg/PbCl2 ) that powers life-saving equipment. Reproduced with permission from Ref. [180]

Fig. 2.26 Aluminum–seawater energy module for undersea power generation. Reproduced with permission from Ref. [171]

• • • • • •

Meteorological electronic equipment Airborne electronic equipment Signal lights or flares Floating buoys or remote sea buoys Sonobuoys Air-sea rescue equipment

2.4 Cell Engineering

• • • •

83

Rescue lights on lifeboats Life jackets with emergency signals Underwater sensors and defense systems Torpedo propulsion In the next chapter, secondary seawater batteries will be discussed.

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158. Lee J-S, Kim ST, Cao R, Choi N-S, Liu M, Lee KT, Cho J (2011) Metal-air batteries with high energy density: Li-air versus Zn-air. Adv Energy Mater 1:34–50. https://doi.org/10.1002/ aenm.201000010 159. Pan J, Xu YY, Yang H, Dong Z, Liu H, Xia BY (2018) Advanced architectures and relatives of air electrodes in Zn–air batteries. Adv Sci 5: https://doi.org/10.1002/advs.201700691 160. Song M, Tan H, Chao D, Fan HJ (2018) Recent advances in Zn-ion batteries. Adv Func Mater 28:1–27. https://doi.org/10.1002/adfm.201802564 161. Kaleva A, Tassaing T, Saarimaa V, le Bourdon G, Väisänen P, Markkula A, Levänen E (2020) Formation of corrosion products on zinc in wet supercritical and subcritical CO2: In-situ spectroscopic study. Corros Sci 174:108850 . https://doi.org/10.1016/j.corsci.2020.108850 162. Song J, Xu K, Liu N, Reed D, Li X (2021) Crossroads in the renaissance of rechargeable aqueous zinc batteries. Mater Today 45:191–212. https://doi.org/10.1016/j.mattod.2020. 12.003 163. Susanto A, Baskoro MS, Wisudo SH, Riyanto M, Purwangka F (2017) Performance of Zn-Cu and Al-Cu electrodes in seawater battery at different distance and surface area. Int J Renew Energy Res 7:298–303 164. Torabi F, Ahmadi P (2020) Chapter 7 - Zinc–silver oxide batteries. In: Torabi F, Ahmadi P (eds) Simulation of battery systems. Academic, pp 217–262 165. Schmidbaur H, Cihonski JL (2003) Noble metals (chemistry). In: Meyers RA (ed) Encyclopedia of physical science and technology, 3rd edn. Academic, New York, Third Edit, pp 463–492 166. Hasvold (2009) Applications - transportation | submersibles: batteries. Encyclop Electrochem Power Sourc i:367–380. https://doi.org/10.1016/B978-044452745-5.00371-3 167. Yardney LR700(DS) cell. https://yardneyfilters.com/ 168. Shen PK, Tseung ACC, Kuo C (1994) Development of an aluminium/sea water battery for sub-sea applications. J Power Sourc 47:119–127. https://doi.org/10.1016/0378-7753(94)800 55-3 169. Leisegang T, Meutzner F, Zschornak M, Münchgesang W, Schmid R, Nestler T, Eremin RA, Kabanov AA, Blatov VA, Meyer DC (2019) The aluminum-ion battery: a sustainable and seminal concept? Front Chem 7:1–21. https://doi.org/10.3389/fchem.2019.00268 170. Wang C, Yu Y, Niu J, Liu Y, Bridges D, Liu X, Pooran J, Zhang Y, Hu A (2019) Recent progress of metal-air batteries-a mini review. Appl Sci (Switzerland) 9:1–22 171. L3Harris AL-H2O Aluminum-water energy modules. In: L3Harris Technologies. https:// www.l3harris.com/all-capabilities/al-h2o-aluminum-water-energy-modules 172. Frost & Sullivan (2017) Post li-ion battery R&D trends. Mountain view 173. Xu Y, Zhao Y, Ren J, Zhang Y, Peng H (2016) An all-solid-state fiber-shaped aluminum-air battery with flexibility, stretchability, and high electrochemical performance. Angewandte Chemie - Int Ed 55:7979–7982. https://doi.org/10.1002/anie.201601804 174. U.S. Department of Energy (2016) Manufacturing of protected lithium electrodes for advanced batteries manufacturing of protected lithium electrodes for advanced lithium- batteries. Advanced manufacturing office 175. Holler R (2014) Evolution of the sonobuoy from world war ii to the cold war. US Navy J Underwater Acoust 25:322–346 176. Maritime Park Association (1943) U.S. Navy Torpedo Mark 18 (Electric), OP 946 177. Shinohara M, Araki E, Kanazawa T, Suyehiro K, Mochizuki M, Yamada T, Nakahigashi K, Kaiho Y, Fukao Y (2006) Deep-sea borehole seismological observatories in the western Pacific: temporal variation of seismic noise level and event detection. Ann Geophys 49. https:// doi.org/10.4401/ag-3131 178. Shinohara M, Araki E, Kanazawa T, Suyehiro K, Mochizuki M, Yamada T, Mochizuki K, Nakahigashi K, Kaiho Y (2007) Seafloor borehole broadband seismic observatories in the western pacific and performance of recovered seismic data. In: International symposium on underwater technology, UT 2007 - international workshop on scientific use of submarine cables and related technologies 2007, pp 682–690 . https://doi.org/10.1109/UT.2007.370821

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Chapter 3

Secondary Seawater Batteries

Abstract A rechargeable seawater battery is a novel cost-efficient and high-density electricity storage system that uses sodium ions and dissolved oxygen from the seawater to convert chemical energy into electricity. In particular, the fact that the active material (Na+ ) is obtained from a practically unending supply of seawater and that it functions in natural seawater elevates the expectation that this secondary seawater battery will potentially be used in many more applications. However, we must consider that this secondary seawater battery is still in the early stages of its development, and to be used as a commercial battery, it must go through many stages of verification and validation, including but not limited to material research, cell development, and manufacturing process optimization. In this regard, it is noteworthy that recent research on secondary seawater batteries has gone beyond laboratory-level development to commercial-grade production techniques, and that various types of cells and pilot manufacturing processes are being reported. This chapter introduces the background, operation principle, characteristics, research on each component, and the cell development process of secondary seawater batteries. In addition, seawater utilization technologies using electrochemical methods related to seawater batteries are also briefly introduced.

3.1 Introduction and Operating Principle 3.1.1 Birth of the Secondary Seawater Battery Primary seawater batteries, developed for a specific purpose as described previously, utilize the current primary battery’s cathode and anode materials and exploit the separated seawater as an electrolyte (with high ionic conductivity of 50 mS/cm at 20 °C), which is an incomplete battery with high power output as well as endurance [1]. Several anodes, including magnesium, were considered, and cathodes with reliable performance such as metal chloride, especially silver chloride, were also potential candidates [2]. Later developments saw primary seawater batteries utilizing seawater or dissolved oxygen as a cathode reaction, with research focused towards boosting

© The Author(s), under exclusive license to Springer Nature Singapore Pte Ltd. 2022 Y. Kim and W. Lee, Seawater Batteries, Green Energy and Technology, https://doi.org/10.1007/978-981-19-0797-5_3

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Fig. 3.1 Concept of Secondary seawater battery

the energy density. However, the immense labor and expenses involved in the production, activation methods, and the one-time use nature of the batteries, have limited their scope to primarily industries receiving national investment, such as military and communications. In comparison to commercialized primary and secondary batteries that have been developed and applied to a variety of prototypes and fields, the reservetype primary seawater battery remains relatively unchanged from its initial development form and therefore posits a difficult challenge in expanding its application to other fields [2–5]. A secondary seawater battery on the other hand overcomes the constraints of a primary seawater battery and also improves its utility [6, 7]. The most distinguishing aspects of such a battery include the fact that it can be charged and discharged in seawater, and also that both the cathode and anode active elements are derived from seawater. In other words, the secondary seawater battery bears the advantage of actively employing the salt in seawater as an energy storage medium rather than simply using it as an electrolyte, as depicted in Fig. 3.1. A basic type of secondary seawater battery utilizes sodium as the active material for the anode and dissolved oxygen in the saltwater as the active material for the cathode. The secondary seawater battery offers an affordable and effective energy storage medium by utilizing sodium and oxygen, two of the most abundant elements in the marine environment. The operational properties of seawater batteries are highly dependent on the ocean’s characteristics due to the direct employment of saltwater and its salts. We briefly reviewed the huge energy-matter circulation system of the ocean in the Chap. 1. In particular, we examined that the seawater salts, i.e., the various ions, are one of the most important characteristics of the seawater. On an average, seawater has 35 PSU of salt, i.e., 35 g of salt in 1 L of seawater. Sodium chloride accounts for about 90% of the total ionic mass (about 10.7 g of sodium and 19.3 g of chloride ions). When all 10.7 g of sodium participates in the oxidation/reduction process, this corresponds to an energy output capacity of approximately 12.5 Ah. However, in a realistic scenario, an almost infinite amount of sodium ions can be utilized when the seawater battery is immersed. Dissolved oxygen participating in the cathodic reaction also has the merit of indefinite supply from the atmosphere. Seawater can be used as an active material source as well as a good electrolyte. The developmental history of the seawater battery (SWB) is depicted in Fig. 3.2, from the primary system to the secondary system and its applications. Mg-AgCl batteries, utilizing seawater as the electrolyte for the first time, were created in the

Fig. 3.2 Timeline of development of Seawater Batteries (SWBs) from a primary to a secondary system and their categories depending on materials, role of seawater, and types. Reproduced with permission from Ref. [6]

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1940s and were employed for military (1943–1945) and later civilian (1945) applications. The following years witnessed an active effort in the development of cheaper materials (CuCl, PbCl2 , Hg2 Cl2 , CuI, and Cu2 O) as a substitute to the relatively costly Ag, but these are essentially reserve-type batteries that operate on saltwater as the electrolyte. This period was followed by the invention of a primary SWB using seawater as the cathode. These batteries utilize oxygen and hydrogen reactions to maximize the energy density on the cells. However, it is still a reserve type primary battery that has been created for specialized applications such as submarine research. It wasn’t until 2014 that the rechargeable SWB technology was introduced [8, 9]. This secondary SWB utilizes seawater as an unlimited source of Na+ ions for storage of electrical energy. The idea that dissolved oxygen in the seawater can be used as a cathode has the potential to further increase the energy density of these secondary SWBs. Secondary SWBs are being developed using practical use after proof of concept and material research in coin cell units can be conducted [10–15]. The study of each material will be dealt with in detail in Sects. 3.2.4, and the development of the cell structure in Sect. 3.5.

3.1.2 Operating Principle and Features The anode of a secondary SWB uses sodium as the active material, and the cathode uses dissolved oxygen in the saltwater as the active material. At the anode, sodium stores energy in a reduced metal form, and when it is oxidized, releases and discharges electrons and positive sodium ions. The oxidation–reduction potential of sodium is −2.73 V [16]. In the cathode, hydroxide ions are oxidized to oxygen and water during charging, and during discharge, dissolved oxygen and water are reduced and converted into hydroxide ions. The redox potential could vary according to the participation of hydroxide ion in the reaction, which effectively determines the pH of the solution. In a seawater environment of pH = 8, a redox potential of about 0.75 V is obtained. The half-cell reaction and overall reaction during charging and discharging are as follows: Anode : Na+ e− ↔ Na, Cathode : O2 + 2H2 O + 4e− ↔ 4OH− ,

E ◦ = −2.73 V

(3.1)

E ◦ = 0.75 V(at p H = 8)

(3.2)

Overall : 4Na + O2 + 2H2 O ↔ 4Na+ + 4OH- ,

◦ E cell = 3.48 V

(3.3)

This novel and advanced seawater rechargeable battery could effectively store sodium in a real seawater environment and provide a stable source of high voltage (~2.7~ 3 V range) during the discharging process [13, 14, 17–19]. Although this operating voltage is slightly lower than the theoretically calculated value, its efficiency is quite remarkable compared to that of primary SWBs.

3.1 Introduction and Operating Principle

95

By employing seawater as the cathode, the secondary SWB can significantly lower the volume and mass of the cathode material, contributing to its high energy density. However, this renders the cathode reaction of the secondary SWB much more complicated than the anode [17, 20–23] due to the competition between the oxygen and the chlorine reaction at the cathode [8, 9, 24–26]. Oxygen reactions, that is, OER (Oxygen Evolution Reaction) and ORR (Oxygen Reduction Reaction), are known to be slow kinetic processes [25, 27–29], which affects the charging and discharging efficiency in addition to competing with the chlorine reaction that occurs at a larger potential with a high overvoltage. While in seawater the cathode is a characteristic and advantage of SWBs, the anode can give rise to its own set of technical difficulties [30–32]. An appropriately designed complex of electrolyte, separator, and cell structure is required to store sodium in a stable manner [14, 33]. A selective Na-ion conductive solid electrolyte (Na3 Zr2 Si2 PO12 ) is employed in this SWB to separate the anode compartment from the seawater cathode. Due to this solid electrolyte, the sodium of the anode is shielded from moisture and air and exists in the form of a metal, an alloy, or an intercalation. An organic electrolyte is frequently introduced in the anode part, which is surrounded by a solid electrolyte. Figure 3.3 shows the movement of electrons and ions in the anode, cathode, and solid electrolyte during the charging and discharging process of a SWB. Unlike primary SWBs, a secondary SWB can be used over multiple iterations as a charge/discharge system. This unique capability allows us to provide a competitive battery option for a wide range of devices used at sea and for future demands [6]. In

Fig. 3.3 Schematic representation of anode and cathode electrochemical reactions in rechargeable SWBs

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Table 3.1 Comparison of primary and secondary SWB Materials

Characteristics

Energy density

Primary SWB

Secondary SWB

Seawater mostly as electrolyte

Seawater as cathode

Electrode materials (i.e., Mg, AgCl) are consumed

Na can be possibly harvested from seawater

Non-rechargeable

Rechargeable

Reaction product can inhibit internal space

Highly stable even after first start to use

1500 Wh/L (practical)

3051Wh/L (theoretical)

9600 Wh/L (theoretical)

other words, the secondary SWB is expected to expand the application of SWBs and to supply the optimal marine power source system to various marine equipment and offshore structures. The primary and secondary SWBs’ materials, properties, and energy densities are listed in Table 3.1. Demand for secondary batteries has been increasing explosively as it can be used in various portable devices and large-capacity ESSs. Lead-acid batteries, nickel– cadmium batteries, nickel-hydrogen batteries, lithium polymer batteries, and lithiumion batteries are some of the most well-known secondary batteries that have been developed [16] since the invention of the technology. Lithium-ion secondary batteries (LIBs) have been continuously improved in performance and capacity due to the small-scaled nature of their applications. The high output of 250 kW h−1 is expected to drive the demand for LIBs in electric vehicle market and electric grids explosively in the future [34]. However, the high price of LIBs (USD 250 per kWh) due to the supply and demand instability of lithium and cobalt has been slowing the expansion of these LIB applications, apart from the safety issues that have also been emerging in the field of LIB-based grid energy storage [35, 36]. The high-cost issue stems from the fact that LIB-based systems require expensive lithium and cobalt-based raw materials for optimal functioning. Not only is cobalt expensive, but it is also poisonous and has a limited supply, like lithium. Safety issues with LIBs arise from their operation in flammable organic electrolytes with a low flash point. Batteries, apart from possessing a high energy density, should also be affordable, durable, and safe for widespread use; this situation prompted researchers to seek an alternative (or complementary) rechargeable battery system to LIBs. Battery scientists realize that using earth plentiful metals such as aluminum, magnesium, sodium, potassium, and calcium in batteries would drastically reduce the expenditure involved in production due to their high abundance, and consequently, low costs [37–41]. From this point of view, secondary SWBs are an excellent candidate for nextgeneration secondary batteries [7]. However, in comparison to commercially used and manufactured secondary batteries, secondary SWBs are still in the early stages of their development. Yet, SWBs hold a lot of promise since they are extremely economical due to their usage of sodium ions, the 6th most abundant element on Earth

3.1 Introduction and Operating Principle

97

accounting for about 2.8% of the Earth’s crust [42]. Furthermore, because Na+ ions are abundant in seawater, they can be gathered in operation through charging in the ocean environment. The concentration of Na+ ions in seawater is about 0.47 M, which is sufficient to supply sodium to the SWB anodes. The secondary SWB generates a theoretical voltage of approximately 3.47 V from a Na+ reduction potential of − 2.71 V and an oxygen redox potential of 0.76 V (at seawater, pH = 8), which is the second highest voltage produced by secondary batteries, only behind lithium [38]. In comparison to the development and widespread use of lithium-ion batteries, secondary SWBs have not yet reached commercialization on a large scale. However, it is worth noting that seawater secondary battery development is a relatively much newer field on which research began only a few years ago. As proven by the development history of several batteries, including lithium-ion batteries, long-term research, idea formation, and numerous field tests are required for commercialization. Currently, secondary SWB technology is being evaluated for application in a variety of fields following the concept’s validation. Numerous nextgeneration secondary batteries are undergoing commercialization studies concurrently with the development of materials and cell unit technologies. This rechargeable SWB has a bright future in a variety of applications, including on/offshore tidal power and wind power, as well as in military energy supplies and submarine special batteries, replacing the utilization of conventional primary batteries (Fig. 3.4). Some striking qualities of SWBs include their ability to function in marine environments due to their usage of seawater as an active substance i.e., the cathode is the open sea. These qualities can provide numerous benefits during the construction and operation of these batteries, which can be broadly summarized into the following categories: (1) cost savings, (2) improved thermal control, and (3) increased cycle life. To begin with, material and manufacturing costs can be significantly lowered in the production of SWBs as we mentioned earlier, especially due to the simplicity

Fig. 3.4 Comparison of LIB and Secondary SWB

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of designing the cathode. Cathode material expenses account for roughly half of all material costs in LIBs [43]. With the spike in demand for batteries for applications such as electric vehicles, renewable energy storage, and smart grids, supply shortage (material reserves and supply reliability) of expensive LIB raw materials, particularly lithium and cobalt, is a serious cause for concern. SWB battery systems that utilize natural seawater in the ocean as well as in well-ventilated areas will offer significant cost savings, depending on the use and application [44]. Secondly, owing to its capability for operation in open seawater environments, SWBs provides solid advantage for heat control. This results in increased stability, extended lifespan, and explosion safety in case of an emergency. When a battery fails as a result of excessive charging and discharging, or due to suboptimal operating conditions (high temperature or corrosion), it results in the occurrence of unintended, automatic reactions such as a sudden spike in temperature. In severe circumstances, the temperature rises past a threshold, resulting in an irreversible exothermic reaction known as “thermal runaway”. Large-scale battery systems (ESS) in particular can result in massive fires, making thermal control a critical aspect of consideration in the design of these batteries. Usually, cooling packs or heat exchangers are suggested to manage the thermal load on big battery systems. However, SWB being an open cathode design allows for it to utilize seawater circulation for effective heat management, considerably reducing the risk of explosion in the event of an emergency. In other words, SWB-based ESSs offer cost efficiency in terms of installation and maintenance per volume (or weight) of the system, as the internal active cooling system can be eliminated or minimized. Finally, SWB can continuously deliver Na+ ions, an active substance, when charging in a driving environment, i.e., the marine environment [6]. During the early charging and discharging cycles of a conventional battery, the active material generates a solid-electrolyte interphase (SEI) layer, reducing the battery capacity by around half. Repeated utilization of such a restricted active material steadily degrades the capacity due to the side reactions that occur during the charging-discharging process. However, SWBs are projected to overcome this issue and maintain a steady capacity for a much longer period due to the unlimited supply of active material (Na+ ) from saltwater.

3.1.3 Structure and Components The basic structure of the unit cell is schematically illustrated in Fig. 3.5 In general, the rechargeable SWB is comprised of three components: a cathode part, an anode part, and a solid membrane. A secondary SWB’s cathode is composed of seawater and a current collector (and catalyst). Since the cell does not fully form until seawater, the cathode material, is applied, it can also be stored and utilized as an immersiontype reserve battery. Water, OH− , and dissolved oxygen are the active components in the electrochemical processes in seawater. However, chlorine species (Cl− , ClO− ) 2− or carbonate species (CO2 , HCO− 3 , CO3 ) may also participate in a competitive

Fig. 3.5 Component of secondary SWB and their candidate materials

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reaction, apart from interference from a variety of organic compounds. As discussed in Chap. 1, seawater is also an electrolyte with a pH of 8 and an ionic conductivity of 50 mS/cm due to the presence of dissolved salts. The current collector facilitates electron mobility and acts as a catalyst for the anode reaction. It should be electrically conductive and have a large surface area to facilitate a sluggish oxygen reaction. Stability in the water (seawater environment) and low material cost should also be considered. Generally, carbon current collectors are suitable for secondary SWB, but titanium materials are occasionally utilized in conjunction with carbon for higher conductivity and physical rigidity [45]. The anode part is where sodium is reduced and stored, and it is typically made up of a non-aqueous electrolyte, a current collector, and an anode material (Na metal, alloy, intercalating material). When sodium is extracted from seawater, the initial anode component may be devoid of sodium active material (Na-free). The non-aqueous electrolyte works as a buffer, allowing sodium ions to transmit electrons from the anode current collector to the separator (solid electrolyte) in a stable manner. The current collector ensures that electrons are transported uniformly to the anode active material. While the cathode has a limitless supply of saltwater and dissolved oxygen, the anode has a finite supply of sodium. This is because of the limited volume of the anode and the surrounding solid electrolyte. As a result, the anode part decides the capacity of the secondary SWB. The solid electrolyte separates the anode and cathode parts in SWBs, allowing for steady sodium storage. Because sodium metal oxidizes rapidly when exposed to water or air, the anode part of the battery must be structurally isolated from seawater. The solid electrolyte acts as a separator between the aqueous anode and the non-aqueous cathode, allowing sodium to be reduced steadily at the anode, yielding a high working voltage. Simultaneously, the separator also equalizes the charge via ion permeation. If this process is accompanied by the transport of other ions, side reactions may occur, necessitating the use of selective ion permeation. NASICON (Na3 Zr2 Si2 PO12 ) performed well when compared to other solid electrolytes, transmitting exclusively positive Na-ions during the charging-discharging process. The solid membrane must possess extreme electrochemical stability and resistance to both seawater and nonaqueous electrolytes.

3.2 Cathode 3.2.1 Reactions Occurring at SWB Cathodes Although rechargeable SWBs primarily utilize the redox reactions of O2 , the ions present in seawater (e.g., OH− and Cl− ) can also donate electrons to the cathode during charging. As a result, depending on the participating ion, several reactions can take place at the SWB cathode, namely (1) the OER/ORR, (2) Cl− capture/release, (3) Na+ intercalation/deintercalation, and (4) the redox reactions of redox couples

3.2 Cathode

101

in the catholyte (Fig. 3.6). As these reactions are affected by the properties of the current collector and the electrode material, they can be used to design SWBs for a broad range of specific applications. In the case of the OER/ORR, the OH− ions present in seawater donate electrons to the cathode during charge and are oxidized to form O2 and water. This process is known as the oxygen evolution reaction (OER) and is the reverse of the oxygen reduction reaction (ORR) occurring during discharging (Fig. 3.6, Type I; Eq. 3.4). Although the OER/ORR combination offers the advantage of limitless resources (seawater), the slow kinetics of these reactions in pH-neutral seawater results in high overpotentials during operation and thus necessitates the use of bifunctional electrocatalysts. O2(aq) + 2H2 O + 4e− ↔ 4OH− ,

E ◦ = 0.75 V vs. SHE( p H = 8)

(3.4)

In the case of Cl− capture/release, the Cl− ions present in seawater donate electrons to the cathode during charge and are trapped in the form of chlorides produced by the oxidation of the metal cathode with the in situ generated elemental chorine (Fig. 3.6, Type II). When a Cl− capture/release cathode is employed, the adsorption/desorption of chloride ions results in oxidation– reduction cycling and, hence, in a substantially smaller charge/discharge gap compared to that observed for ORR/OER cathodes. However, such cathodes suffer from anode-restricted storage capacity and poor cycling stability. Ag is commonly used as a cathode material for Cl− capture/release (Eq. 3.5). − Ag(s) + Cl− (aq) ↔ AgCl(s) + e ,

E ◦ = 0.22 V vs. SHE

(3.5)

The intercalation/deintercalation of Na+ in seawater can be used as a cathode reaction when suitable electrode materials are used instead of seawater (Fig. 3.6,

Fig. 3.6 Schematics of cathode compartment. Reproduced with permission of Ref. [46]

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Type III). These reactions are usually fast and therefore show relatively high voltage efficiencies. Cell capacity, however, depends on the specific capacity of the cathode active material. Hence, cathode capacity can be increased via the hybridization of Na+ intercalation/deintercalation and OER/ORR cathodes. Nickel hexacyanoferrate (NiHCF) and layered oxides (Na0.5 Co0.5 Mn0.5 O2 ) are exemplary Na+ intercalation/deintercalation materials, with the involved reactions presented in Eqs. 3.6 and 3.7. Na0.61 Ni1.78 FeII (CN)6 ↔ Na0.61−x Ni1.78 FeIII (CN)6 + xNa+ + xe−

(3.6)

Na0.5 Co0.5 Mn0.5 O2 ↔ Na0.5−x Co0.5 Mn0.5 O2 + xNa+ + xe−

(3.7)

Rechargeable SWBs can be operated with a solution of redox couples as the catholyte instead of seawater (Fig. 3.6, Type IV), in which case the cathode reaction is the redox reaction of the active material. Such reactions are usually rapid and therefore lead to relatively high voltage efficiencies. However, as the area provided by the catholyte tank is limited, the battery capacity depends on catholyte volume and active material solubility. Hence, cathode capacity can be increased through the hybridization of redox and intercalation/deintercalation cathodes. Sodium hexacyanoferrate (NaHCF) is an exemplary redox catholyte, engaging in the following redox reaction (Eq. 3.8). 4− 3−   II Fe (CN)6 (aq) ↔ FeIII (CN)6 (aq) + e− ,

E ◦ = 0.36 − 0.51 V vs. SHE

(3.8)

3.2.2 OER/ORR SWB charge/discharge is based on the redox reactions of Na+ and seawater at the anode and cathode, respectively. Unlike conventional batteries, which operate in a closed system, SWBs use flowing seawater and an open-structure cathode exposed to ambient air. Although seawater has a low specific capacity (12.6 mAh g−1 , 12.9 Ah L−1 ), the above configuration enables an infinite supply of reactants and the removal of reaction products, thus allowing one to maintain a constant concentration of reactants at the cathode during operation. As charging relies on the extraction of Na+ from seawater and the transport of these ions to the anode side through a membrane, SWBs may employ seawater as the only active material, which eliminates the requirement for an initial Na metal anode (anode-free) or a Na+ storage anode. Both the OER and the ORR use ambient O2. As seen in Fig. 3.7, O2 is introduced into the seawater catholyte by aeration, and the thus produced dissolved oxygen (DO) participates in the catholyte reduction reaction (ORR) during discharge. Figure 3.8a depicts a simulated Pourbaix diagram of seawater, with potentials defined relative to the SHE (left) and the Na+ /Na couple (right). When seawater with pH 8 is oxidized

3.2 Cathode

103

Fig. 3.7 Schematic SWB structure and operation mechanism upon charge/discharge. Reproduced with permission of Ref. [46]

Fig. 3.8 a Simulated Pourbaix diagram for seawater at 25 °C, b typical SWB charge–discharge curve. Reproduced with permission of Ref. [46]

during charging, the OER (Eq. 3.9) is thermodynamically preferable to hypochlorite production (Eq. 3.10). However, both reactions compete for energy at high charging currents because of the slow kinetics of the OER [47]. The corresponding half-cell and overall reactions are given below (Eq. 3.11), and the corresponding potentials are given for 100% saturation of seawater with O2 (Po = 0.206 atm), pH ≈ 8, and [Na+ ] ≈ 0.47 M. DO(aq) + 2H2 O + 4e− ↔ 4OH− ,

E ◦ = 0.75 V vs. SHE(pH = 8)

− − − OCl− (aq) + H2 O + 2e ↔ Cl(aq) + 2OH ,

4Na(s) + DO(aq) + 2H2 O ↔ 4NaOH(aq) ,

(3.9)

E ◦ = 1.24 V vs. SHE(pH = 8) (3.10) E ◦ = 3.48 V vs. Na+ /Na

(3.11)

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The theoretical maximum cell voltage (E cell ) of SWBs is 3.48 V. However, when the ORR involves two-electron transfer (Eq. 3.12) or in the case of anoxic (oxygendeficient) conditions at the cathode during discharge (Eq. 3.13), the following cathode half-reactions may occur, resulting in a decrease of E cell to 2.9 or 2.2 V, respectively. − DO(aq) + H2 O + 2e− ↔ HO− 2 + OH ,

2H2 O + 2e− ↔ H2(g) + 2OH− ,

E ◦ = 0.21 V vs. SHE (pH = 8) (3.12)

E ◦ = −0.47 V vs. SHE (pH = 8)

(3.13)

Figure 3.8b shows a typical SWB charge–discharge curve. Typically, SWBs have an open-circuit voltage of 2.9–3.4 V versus Na+ /Na, depending on the cathode current collector or electrodes. The utilized materials should be developed with the marine environment in mind. The cathode compartment of rechargeable SWBs contains the current collector, the catalyst, and seawater. The catalyst-coated cathode current collector is directly exposed to seawater for electron exchange, which has an effect on the OER/ORR overpotential. As a result, the cathode material should not only be physically and chemically stable in seawater but should also assure the overall voltage efficiency via OER/ORR catalytic activity. While collectors and catalysts should occasionally be examined together, they are treated separately in this section based on their roles and compositions.

3.2.2.1

Carbon-Based Current Collectors

The cathode current collector supplies reaction sites and electron transport channels for the cathode reaction and should therefore exhibit high electrical conductivity, excellent (electro)chemical stability, a large surface area (i.e., numerous active sites), and low cost. In this regard, further studies on cathode current collectors and the influence of catholyte hydrophilicity and flow properties are required. Commercially available fibrous carbon materials such as carbon felts (CFs), carbon papers, and carbon cloths have been used as effective current collectors in a variety of energy storage devices. However, in view of the presence of sizing agents for mechanical reinforcement, CFs generally have hydrophobic surfaces and are therefore not well wetted by aqueous electrolytes such as seawater. Consequently, CFs were heated to 500 °C in air [48], and their surface wettability was assessed by contact angle measurements. Figure 3.9a illustrates the contact angles of seawater on pristine (raw) carbon felt (PCF) and heated carbon felt (HCF), showing that even after 5 min, PCF was not completely wetted and featured a contact angle of 135°. In contrast, HCF absorbed the water droplet within 1 s (contact angle could not be determined), i.e., heating in air rendered the CF surface strongly hydrophilic. Figure. 3.9b presents the initial charge–discharge curves of SWBs with PCF and HCF air electrodes (OFF state). The charge and discharge voltages of the PCF cell were 4.1 and 2.5 V, respectively, i.e., the charge–discharge voltage difference was

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105

Fig. 3.9 a Contact angles of water on PCF and HCF and the related photographs (inset), b galvanostatic charge–discharge curves of SWBs with PCF and HCF air electrodes. Reproduced with permission of Ref. [15]

as high as 1.6 V. On the other hand, the HCF-based cell featured a lower voltage gap of 1.4 V. Thus, the reaction kinetics during charge–discharge cycling could be improved by increasing the wettability of the cathode current collector (in this case, CF) and thus decreasing the voltage gap. The effects of heat treatment on CF surface chemistry and surface area were probed using XPS and N2 sorption isotherm measurements. Heating in air increased the C–O to C=O bond ratio [49–51] and the specific surface area of CFs (Fig. 3.10a and b), thus increasing surface wettability and decreasing the voltage gap to improve the energy efficiency of the corresponding SWBs. SWBs make use of naturally flowing seawater to steadily supply reactants (e.g., water and O2 ) to the reaction sites of the cathode current collector and remove reaction products (e.g., HO− ) during cell operation. This circulation promotes the conversion of ambient O2 (g) to DO in seawater (aeration), decreases the thickness of the mass-transfer boundary layer at the electrode, and therefore increases the limiting

Fig. 3.10 a, b O 1 s spectra of PCF and HCF. Insets show the related S BET values. Reproduced with permission of Ref. [15]

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Fig. 3.11 a Charge–discharge profiles recorded in flow-ON and -OFF states; b discharge voltage and DO content profiles recorded in flow-ON and -OFF states. Reproduced with permission of Ref. [15]

current density, particularly for the ORR during discharge. A comparison between the flow-ON and -OFF states of the seawater catholyte revealed the influence of flow on cell performance, demonstrating that this flow results in very small voltage gaps between the charge and discharge profiles during flow-ON operation. Figure 3.11a presents charge–discharge profiles recorded in OFF and ON modes. When the flow was turned off, the charge and discharge terminal voltages were 3.9 and 2.8 V, respectively, resulting in a large voltage gap of 1.1 V. On the other hand, seawater flow (ON state) significantly reduced the voltage gap (0.7 V), with the charge and discharge voltages equaling 3.7 and 3.0 V, respectively. It should be noted that seawater flow was more efficient in decreasing the overpotential generated during discharging than in decreasing that generated during charging. The DO content of seawater (catholyte) was measured during discharge in ON and OFF states (Fig. 3.11b). The initial discharge voltage in the ON state equaled 2.8 V and remained nearly constant during discharge, while DO concentration declined from 9.2 to 8.2 mg L−1 . Thus, discharge consumed DO, as indicated previously. Immediately after the flow was switched off, the voltage and DO content dropped to 1.9 V and 5.9 mg L−1 , respectively, but were almost fully recovered when the flow was subsequently turned on. These results demonstrate that the circulation of seawater (catholyte) enables the continuous supply of reactants to the active areas of the cathode current collector. Specifically, this circulation promotes the continuous dissolution of ambient O2 gas in seawater and thus ensures an adequate supply of DO to the submerged cathode current collector during discharge, mitigates cathode-side concentration polarization, and lowers the overpotential generated by the OER/ORR during operation. In summary, the wettability of CF cathode current collectors by seawater was investigated, as well as the influence of this wettability on charge–discharge cycling performance. Heating in air rendered the CF surface hydrophilic, thus decreasing the voltage gap (1.4 V; cf. 1.6 V for the cell with pristine CF). Additionally, the circulation of seawater (catholyte) decreased the voltage gap (0.7 V; cf. 1.1 V in

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107

the flow-OFF state), most likely by facilitating reactant (water and DO) diffusion to the cathode current collector and thus decreasing concentration polarization during charge and discharge [52, 53]. Despite the performance enhancement offered by the increased wettability of carbon current collectors and catholyte circulation, SWB cathodes still suffer from low voltage efficiency due to the slow kinetics of the OER and the ORR. Thus, commercially available fibrous carbonaceous materials such as CFs, carbon sheets, and carbon cloths were studied as current collector materials for SWB cathodes (Table 3.2). The hybridization of the electrical double layer (EDL) method with electrocatalytic activity (OER/ORR) was validated using carbon cloth (ACC 5092– 20, Table 3.2) as the current collector of the SWB cathode. Charge storage and transport via the EDL are significantly faster than OER/ORR and other faradic processes, as is the case for supercapacitor-type electrodes in hybrid energy storage (HES) devices [54–59]. These devices contain two electrodes, one of which is a supercapacitor capable of storing ions via EDL formation, i.e., via the adsorption of electrolyte ions on the material surface, while the other is a battery capable of storing charges via a redox or Faradaic reaction, i.e., via intercalation/deintercalation or conversion [60, 61]. Thus, similar to supercapacitor electrodes in HES devices, the SWB cathode current collector is submerged in an electrolyte (seawater), and electrochemical reactions take place at the electrode/electrolyte interface. This similarity allows the EDL process to be implemented on the current collector (Fig. 3.12) [62]. Commercially available activated carbon cloth (ACC) with a large surface area (2038 m2 g−1 ) and CF with a low surface area (2 m2 g1 ) were compared as current collector materials. Figure 3.13a and c show low-magnification SEM images of CF and ACC, respectively, whereas Fig. 3.13b and d present the respective highmagnification images. In CF, the carbon fibers were not uniformly organized, with the bundle thickness roughly equaling 3 mm (Fig. 3.13a). Additionally, the carbon fiber surface in CF exhibited a slit-shaped pore structure (Fig. 3.13b). ACC contained carbon fibers that were arranged in bundles (Fig. 3.13c) to afford a cloth-like structure and featured abundant surface nanopores (Fig. 3.13d). Figure 3.13e and f show the N2 sorption isotherms and pore size distributions of CF and ACC, respectively. CF demonstrated a type IV isotherm, thus containing mesopores, whereas ACC demonstrated a type I isotherm, thus containing micropores [63, 64]. The average pore diameter/specific surface area of CF and ACC were determined as 14.55 nm/2.04 m2 g−1 and 1.76 nm/2038 m2 g−1 , respectively, i.e., ACC had a much larger surface area than CF. Thus, when used in SWBs, ACC may provide a significantly greater capacity for charge storage via EDL formation and more active sites for OER/ORR than CF. The Raman spectra (Fig. 3.14a) of CF and ACC featured two broad peaks at 1350 cm−1 (D band) and 1591 cm−1 (G band), which were assigned to disordered carbon and graphitic carbon, respectively [65]. The corresponding peak intensity ratio (I D /I G ) was larger for CF (1.34) than for ACC (1.01), suggesting the existence of more defect carbon structures in the former case and the presence of a wellbalanced graphitic and disordered carbon structure in the latter case. Additionally,

Manufacturer

Kynol, Wizmac

Kynol, Wizmac 400 °C, 4 h

Kynol, Wizmac

Kynol, Wizmac

CNF

CNF 500 °C, 4 h

Sigmatech, FRP shop

Product name

ACC-5092–20

Heated ACC-5092–20

ACC-5092–10

ACC-5092–15

CF

Heated CF

Carbon cloth (T300)

Image

SEM

Sample1

CNC-1500

Activated CF

1071 HCB

Px30

Ox

UT-050

Product name

Samjung

Trustchem

Namtong

AVcarb

Zoltek

Zoltek

SamjungCNG

Manufacturer

Table 3.2 Commercially available fibrous carbonaceous materials investigated as materials for SWB current collectors Image

(continued)

SEM

108 3 Secondary Seawater Batteries

Manufacturer

Sigmatech 500 °C, 4 h

Sutong

Sutong

Sutong

Product name

Heated carbon cloth

STF-1800

STF-1600

STF-1300

Table 3.2 (continued) Image

SEM

Carbon paper

BET2000 ACF-CNT

BET1300

Sample2

Product name

Dissol

China Beihai Fiberglass Co., Ltd Dissol

China Beihai Fiberglass Co., Ltd

Samjung

Manufacturer

Image

SEM

3.2 Cathode 109

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3 Secondary Seawater Batteries

Fig. 3.12 Schematic of metal-free SWB with a hard carbon anode and electrochemical processes at the cathode part. Reproduced with permission of Ref. [17]

the surface functional groups of CF and ACC were probed by XPS. Figure 3.14b shows the survey spectra of CF and ACC, revealing the presence of C 1 s and O 1 s peaks. Figures 3.14c and d present the core-level C 1 s spectra of CF and ACC, respectively, showing that these spectra could be deconvoluted into three and four peaks, respectively. These peaks were assigned to C=C (284.5 ± 0.1 eV), C–C, (285.3 ± 0.1 eV), hydroxyl or ether C–O groups (286.2 ± 0.1 eV), and C=O groups (287.1 ± 0.1 eV) [50, 51, 66]. The C–O and C=O groups are hydrophilic and may form hydrogen bonds with water molecules to promote their absorption [66, 67]. Compared to CF, ACC had more C=O groups (Fig. 3.14c and d) and a higher surface area, and was therefore expected to exhibit better hydrophilic and electrocatalytic properties. CV and galvanostatic charge–discharge cycling were used to investigate the electrochemical properties of SWBs with CF and ACC as cathode materials. Figure 3.15a presents the obtained CV curves (2.5–4 V vs. Na+ /Na), revealing that large currents were observed between 3 and 3.5 V for the ACC-SWB because of the large surface area of ACC, which facilitated EDL formation. Conversely, only minimal current was observed for the CF-SWB, i.e., the EDL was not formed in this case because of the low surface area of the electrode material. Thus, the substantial surface area of ACC enabled a hybrid electrochemical process combining EDL formation and OER/ORR activity. Figure 3.15b presents the charge–discharge profiles of ACC- and CF-SWBs, showing marked differences between the two cases. For the CF-SWB, flat voltage plateaus attributable to the OER and the ORR were observed at 3.94 and 2.61 V, respectively, whereas two distinct regions (slope and plateau) were observed for the ACC-SWB. The slope zone was ascribed to ion adsorption during EDL formation, whereas the plateau reflected the occurrence of the OER. It is worth noting that galvanostatic charge–discharge tests corroborated the results of CV measurements.

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Fig. 3.13 (a–d) SEM images of (a, b) CF and (c, d) ACC at (a, c) low and (b, d) high magnification. (e, f) N2 sorption isotherms and pore size distributions (insets) of (e) CF and (f) ACC. Reproduced with permission of Ref. [17]

According to Fig. 3.15b, the EDL process in the ACC-SWB accounted for 50% of the overall capacity, whereas this contribution was as low as 2% for the CFSWB. Moreover, the overpotential was attained within only 30 min of charging and discharging for the CF-SWB, whereas a longer time of 3.75 h was required for the ACC-SWB because of EDL formation in the latter case. The overpotentials at the end of charge–discharge cycling were higher for the ACC-SWB (3.64 and 2.81 V) than for the CF-SWB (3.94 and 2.61 V), i.e., the OER/ORR electrocatalytic activity was higher in the former case. The charge–discharge cycling of the above SWBs was performed at currents of 0.2–2 mA to shed more light on galvanodynamic electrochemical processes. The expansion of the voltage gap with increasing current (Fig. 3.16a) was ascribed to the sluggishness of OER/ORR kinetics. On the other hand, even at a high current

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Fig. 3.14 a Raman spectra and b X-ray photoelectron survey spectra of CF and ACC. C 1 s core-level spectra of c ACC and d CF. Reproduced with permission of Ref. [17]

Fig. 3.15 a CV curves and b initial charge–discharge profiles of ACC- and CF-SWBs. Reproduced with permission of Ref. [17]

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Fig. 3.16 a Galvanostatic charge–discharge profiles of the ACC-SWB recorded at various currents and b polarization curves and power densities of ACC- and CF-SWBs. Reproduced with permission of Ref. [17]

of 2 mA, the charge–discharge profiles of the ACC-SWB featured both slope and plateau regions, whereas only the latter region was observed for the CF-SWB. Moreover, at this current, the ACC-SWB featured a lower voltage gap and a greater energy efficiency (1.16 V and 69%, respectively) than the CF-SWB (2.94 V and 47%, respectively). Additionally, a different discharge current was employed to study the power performances of CF- and ACC-SWBs, with the obtained polarization curves and power characteristics shown in Fig. 3.16b. The ACC-SWB exhibited higher cell voltages and maximum power densities than the CF-SWB. Specifically, a power density of 16.2 mW cm−2 at 1.84 V (@ 8.75 mA) was observed in the former case, whereas a power density of 5.5 mW cm−2 at 1.64 V (@ 3.25 mA) was observed in the latter case. Notably, polarization curve slopes (Fig. 3.16b) indicated that the ACC-SWB had a lower total internal resistance (168 ) than the CF-SWB (447 ). In summary, the operation of the ACC-SWB involved a hybrid electrochemical process including both EDL formation and OER/ORR electrolytic activity, whereas the operation of the CF-SWB involved only the latter activity. In view of the advantages of the hybrid electrochemical process, the ACC-SWB demonstrated a smaller voltage gap, a higher voltage efficiency, and a higher power performance than the CF-SWB. Electrochemical analysis indicated that the increased surface area of ACC allowed the hybridization of EDL formation with electrocatalytic activity (OER/ORR) at the cathode current collector of the SWB. Thus, the above findings describe the mechanism of the hybrid electrochemical process occurring at the SWB cathode and demonstrate its benefits for enhancing voltage/energy efficiency and power performance when large-surface-area current collectors are used. A previous study demonstrated that the use of high-surface-area current collectors such as ACC enhances voltage/energy efficiency and power performance relative to those obtained for CF. However, the use of ACC as a cathode current collector resulted in an increased voltage gap and a decreased energy efficiency after 80 cycles (Fig. 3.17b). Additionally, the steady increase in overpotential even during

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Fig. 3.17 a Schematic illustration of cell components, b galvanostatic charge–discharge profiles recorded during the first and 80th cycles for an ACC current collector, c cycling performance showing recovery after cathode replacement. Reproduced with permission of Ref. [23]

early cycles and the restoration of voltage efficiency following cathode replacement (Fig. 3.17c) indicated that the cathode is the primary site of cell performance degradation, which is mainly due to (1) gradual impedance build-up caused by surface carbon oxidation and (2) structural damage after a critical operation time. Further, the failure causes of ACC-based cathode current collectors in SWBs were explored and linked to the occurrence of cathode side reactions alongside the anticipated target reactions (OER/ORR) during charge and discharge. Carbon corrosion was initially confirmed as a side reaction by monitoring gas development using differential electrochemical mass spectrometry (DEMS) during galvanostatic SWB charging (Fig. 3.18a). This corrosion ultimately resulted in CO2 evolution (Eqs. 3.14–3.16) and was caused by the thermodynamic instability of carbon in seawater environments at high electrochemical potentials. C + 2H2 O ↔ CO2 + 4H+ + 4e− ,

E◦ = 0.207 V vs. SHE

C + H2 O ↔ CO + 2H+ + 2e− , CO + H2 O ↔ CO2 + 2H+ + 2e− ,

(3.14)

E◦ = 0.518 V vs. NHE

(3.15)

E◦ = −0.103 V vs. NHE

(3.16)

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115

Fig. 3.18 a CO2 evolution monitored by DEMS during the galvanostatic charging of the ACC current collector, and b, c SEM images of the ACC pre-charged at 20 mAh. Insets show images of pristine ACC. Reproduced with permission of Ref. [23]

The above process begins at the edges and corners of carbon fiber basal planes, which are more sensitive to electrochemical corrosion, and the oxygenated surface species are degraded to CO2 while functioning as corrosion intermediaries [68, 69]. The rate of carbon corrosion is influenced by the shape and surface functional groups of the ACC cathode. According to Fig. 3.18a, the fast degradation of more vulnerable edge sites and functional groups causes an early CO2 evolution peak, which is followed by a rise in gas evolution owing to corrosion. The shape and microstructure changes of the ACC cathode due to carbon corrosion were examined by SEM. Figure 3.18b and c show SEM images of a pre-charged ACC cathode, revealing that compared to pristine fibers (inset), degraded fibers had a more distorted structure. Figure 3.18b shows the occurrence of surface delamination, while Fig. 3.18c reveals the occurrence of main body splitting, which reduced the number of connections. Carbon corrosion may cause structural changes that negatively impact cell performance, e.g., peeling and fracturing, by reducing specific surface area and electrical conductivity. Both pristine and pre-charged ACC featured Type I N2 sorption isotherms (Fig. 3.19a), which was indicative of a microporous structure similar to that of the pristine ACC surface seen in the insets of Fig. 3.18b and c. In view of the availability of surface reaction sites, ACC had a relatively high specific surface area of 2077 m2 g−1 and an average pore diameter of 1.71 nm. The adsorbed N2 volume significantly decreased after charging, and so did the specific surface area (1285 m2 g−1 ) and mean pore diameter (1.65 nm). This behavior suggested that charging-induced carbon corrosion resulted in the deterioration of the SWB cathode and, hence, in the reduction of cell performance. The effects of carbon corrosion on the surface functional groups of ACC were probed by XPS. Figure 3.19b shows the survey X-ray photoelectron spectra of ACC

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Fig. 3.19 a N2 sorption isotherms and pore size distributions (inset) of pristine and charged ACC. The inset on the right shows a SEM image of pristine ACC. b Survey X-ray photoelectron spectra of charged and pristine ACC. Reproduced with permission of Ref. [23]

cathodes before and after charging, revealing the dominance of C 1 s and O 1 s peaks, with other peaks corresponding to S 2p, N 1 s, Na 1 s, and Mg 1 s. As charging increased the intensity of the O 1 s peak and decreased that of the C 1 s peak, it was concluded to induce the electrochemical oxidation of the ACC surface. Thus, XPS data confirmed that carbon corrosion involved oxidation reactions. The electrochemical performances of pristine and pre-charged ACC cathodes were studied using the galvanostatic charging of the corresponding SWBs. Figure 3.20a shows the voltage curves recorded during charge tests, revealing that the pre-charged ACC cathode featured a greater voltage plateau and therefore indicating that the polarization degree increased during pre-charging. A three-electrode method was used to record the CV curves of ACC cathodes before and after pre-charging and thus determine their electrochemically active surface area (ECSA). As shown in Fig. 3.20b, the pristine cathode showed a certain initial capacitive activity, and an

Fig. 3.20 a Galvanostatic charge profiles and b CV curves of pristine and charged ACC cathodes. Reproduced with permission of Ref. [23]

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117

EDL was subsequently formed. Pre-charging reduced the area between the cathodic and anodic curves, resulting in the loss of ECSA and EDL. The current density at the peak of the CV curve decreased for the OER but increased for the ORR, which indicated a reduction in electrochemical activity. Unexpected carbon corrosion occurred during charging, reducing surface area and electrochemical activity. As discussed previously, carbon corrosion resulted in carbon fiber cracking and reduced specific surface area. Thus, the above data suggested that carbon corrosion affected the performance of the ACC cathode. The XRD pattern of the precipitate formed on the discharged ACC cathode (Fig. 3.21a) matched that of aragonite, i.e., CaCO3 precipitation occurred as a side reaction during SWB discharge. This precipitation was ascribed to the combination of CO2 circulation in seawater and the absorption of CO2 due to the ORR occurring at the cathode during discharge (Eqs. 3.17–3.19). The increase in the pH of saltwater near the ACC cathode facilitated the formation of carbonate ions, which reacted with the naturally occurring cations in seawater to generate solid carbonates, among which CaCO3 was the most stable and had the lowest Gibbs free energy of formation [42]. CO2 + H2 O ↔ H2 CO3 ↔ H+ + HCO− 3

(3.17)

2− − HCO− 3 + OH → CO3 + H2 O

(3.18)

Ca2+ + CO2− 3 → CaCO3(s)

(3.19)

Further insights into the physical impacts of CaCO3 precipitation were provided by the SEM imaging of ACC cathode fibers after discharge (Fig. 3.21b and c). These

Fig. 3.21 a XRD pattern of the precipitate formed on discharged ACC, with the CaCO3 reference pattern presented for comparison. b, c SEM images of pre-discharged ACC demonstrating CaCO3 precipitation on the fiber surface, with insets showing images of pristine ACC. Reproduced with permission of Ref. [23]

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images showed that the insulating crystalline CaCO3 precipitates covering the ACC micropores prevented the access of reactants to the reaction sites and hindered electron transport, which agreed with the previously acquired XRD data and suggested that this precipitation should adversely affect cell performance. The N2 sorption isotherms of ACC before and after discharge were recorded to determine the effect of CaCO3 precipitation on the surface area and pore size distribution (Fig. 3.22a). The ACC cathode was subjected to a similar procedure to artificially precipitate more CaCO3 crystals and was compared to pristine ACC. As a result, CaCO3 precipitation decreased the specific surface area (from 2077 to 1441 m2 g−1 ) and the mean pore diameter (from 1.71 to 1.67 nm) of ACC. The reduction in surface area was most likely due to CaCO3 precipitates surrounding the ACC surface. The XPS analysis of a comparable discharged sample revealed the effects of CaCO3 precipitation on the surface functional groups of ACC. As shown in Fig. 3.22b, the spectrum of the discharged sample featured minor S 2p, N 1 s, Na 1 s, Mg 1 s, Ca 2 s, and Ca 2p peaks as well as major C 1 s and O 1 s peaks. This finding, together with the fact that the contents of O and Ca increased after charging, confirmed the formation of CaCO3 precipitates during SWB discharge. The electrochemical performance of the ACC cathode in SWBs was tested before and after pre-discharge, with the recorded discharge voltage profiles shown in Fig. 3.23a. The pre-discharged cathode performed worse than the pristine ACC cathode, featuring a discharge voltage plateau between 2.92 and 2.75 V. The deposition of CaCO3 on the ACC cathode reduced cell performance, as reflected by the rise in polarization intensity during discharge. The effects of CaCO3 precipitation on the ECSA of the discharged ACC cathode were probed by CV (Fig. 3.23b). The ECSA of the pre-discharged cathode was lower than that of the pristine cathode but higher than that of the cathode subjected to carbon corrosion. The apex current density decreased for the OER and increased for the ORR, which indicated a loss of electrochemical activity at the ACC cathode. The deposition of CaCO3 on the carbon

Fig. 3.22 a N2 sorption isotherms and pore size distributions (inset) of charged and discharged ACC. Inset shows a SEM image of pristine ACC (inset). b Survey X-ray photoelectron spectra of pristine and discharged ACC. Reproduced with permission of Ref. [23]

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119

Fig. 3.23 a Galvanostatic discharge profiles and b CV curves of pristine and discharged ACC cathodes. Reproduced with permission of Ref. [23]

fiber surface resulted in surface area loss, as demonstrated by the above SEM images and N2 sorption isotherms. CV data confirmed that CaCO3 precipitation precluded the reactants from accessing the active sites of the ACC cathode. The comparison of CV data for carbon corrosion and CaCO3 precipitation reveals that carbon corrosion degraded electrochemical performance to a greater extent than CaCO3 precipitation. The side reactions of carbon corrosion and CaCO3 precipitation occurring at the ACC cathode during SWB charging and discharging were investigated further. Furthermore, it was shown that the detrimental effects on cell functioning were exerted when the artificial aggravation of the related side reactions was initiated. Nevertheless, the impacts of these cathode side reactions have not yet been determined under the conditions of real-life operational cycling. An experimental setup consisting of two cathodes and a single anode was employed to better understand the effects of each distinct side reaction during cycling. Notably, the above setup allows one to physically decouple the OER from the ORR so that they occur independently at alternate time intervals at two distinct ACC cathodes while sharing the same catholyte (seawater). That is, under cycling conditions, each cathode side reaction can be confined to the respective SWB cathode. Figure 3.24a depicts the results of this experiment. For 5 h, each ACC cathode was either charged or discharged, which resulted in significant variances in the observed voltage plateaus. A constantly climbing voltage plateau was observed after repeated charging, suggesting that the polarization extent increased with time. In the case of discharge, a consistent voltage plateau was maintained even after >50 cycles. These findings demonstrate that although carbon corrosion has an irreversible negative impact on cell performance, CaCO3 precipitation during charge/discharge was reversible and could be hindered, thus having no marked influence on battery operation. Figure 3.24b presents the postulated reasons for the irreversibility of carbon corrosion and the reversibility of CaCO3 precipitation, along with the related consequences. The pH shift of the seawater catholyte during the OER/ORR as well as SWB charge and discharge are some of the causes of the problem. When pCO2 increases,

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Fig. 3.24 a Galvanostatic cycling profiles obtained using a three-electrode setup with physically detached cathodes, b schematic depiction of carbon corrosion and CaCO3 precipitation/dissolution at the carbon fiber surface during cycling, and c in situ pH estimation of the seawater catholyte during galvanostatic charge–discharge. Reproduced with permission of Ref. [23]

CaCO3 dissolves, in line with the equilibria in the carbonic acid–water system [70]. In terms of electrochemical reactions in the SWB, seawater pH decreases during charging because of the consumption of hydroxide ions by the OER, and the pH increase during discharge results from the generation of hydroxide anions by the ORR. Figure 3.24c presents the results of in situ measurements of seawater pH during SWB charge and discharge.

3.2.2.2

Electrocatalysts

In view of the slow kinetics of the OER and the ORR, rechargeable SWBs with the aforementioned carbon (e.g., CF) current collectors still exhibited poor voltage efficiencies and large voltage gaps between charge and discharge curves. These gaps could be reduced through the deposition of electrocatalysts on the cathode current collector to enhance OER and ORR activity. Electrocatalysts are defined as substances that accelerate electrochemical reactions without being consumed in the process and can be divided into those based on noble and non-noble metals. This section discusses research conducted on electrocatalysts to improve the performance of rechargeable SWBs.

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121

Noble metal–based electrocatalysts (Pt/Co) Generally, commercial electrocatalysts mostly contain noble metals such as Pt, Ru, and Ir. As IrO2 -based electrocatalysts are preferred for the OER, while Pt/C is preferred for the ORR, neither of these catalysts can efficiently promote both the OER and the ORR in rechargeable SWBs [71–80]. However, a bifunctional Pt–Co alloy OER/ORR electrocatalyst for application in high-performance SWBs has been recently described. Carbothermal shock (CTS) has been used to synthesize multi-component metallic nanoparticles [81–83]. Pt and Co were chosen as ORR and OER catalysts, respectively. While precious metal oxides such as IrO2 and RuO2 have been extensively employed as OER catalysts, the first-row transition metal Co has recently drawn attention because of its potential OER activity and durability [84–86]. Pt–Co alloys were deposited on a HCF current collector using aqueous solutions of H2 PtCl6 ·6H2 O and CoCl2 in molar ratios of 1:0, 2:1, 1:1, 1:2, and 0:1 as precursors (Fig. 3.25a). Notably, the color of the above solutions changed from intense yellow to pale yellow when the Pt:Co ratio changed from 2:1 to 1:2. UV–vis absorbance spectroscopy (Fig. 3.25b) indicated that the introduction of CoCl2 broadened the strong absorption band at 265 nm associated with the PtCl6 2− anion and shifted it to longer wavelengths [87]. Figure 3.25c illustrates the technique used to form Pt–Co alloy catalysts on the HCF current collector. The HCF current collectors were impregnated with the precursor solution, dried in a vacuum oven at 60 °C, and the subjected to CTS to rapidly produce metal catalysts. Finally, the carbon current collectors containing nanoparticles were used to create a SWB with a Na metal anode, a NASICON separator, and seawater as the catholyte (Fig. 3.25d) [15]. Prior to examining the electrochemical performance of Pt–Co alloys in SWBs, we prepared single-component metal catalysts and characterized their galvanostatic properties. For comparison, pristine HCF was also employed as a cathode current collector (Fig. 3.26a). Single-component nanoparticulate Pt and Co catalysts were produced on HCF current collectors by applying a current of 10 A for 1–10 s. Figures 3.26b and c present SEM images of Pt and Co nanoparticle–containing HCF current collectors, respectively, revealing the presence of well-developed nanoparticles. The CTS-induced production of metal nanoparticles on HCF was confirmed by XPS. The C 1 s spectrum of pristine HCF featured a peak at 285 eV, which indicated significant sp2 bonding [89]. The Pt 4f spectrum featured distinctive peaks at 74.7 and 71.3 eV corresponding to the Pt 4f5/2 and Pt 4f7/2 transitions of Pt nanoparticles on HCF, respectively. Additionally, the Co 2p spectrum indicated that CTS induced the formation of CoO nanoparticles on HCF. Thus, CTS enabled the effective reduction of metal precursors on HCF current collectors and was concluded to be well suited for the large-scale synthesis of metal nanoparticles. Subsequently, the electrochemical performance of pristine and Pt–Co nanoparticle–decorated HCF current collectors for usage in SWBs was evaluated. Figure 3.26g presents the representative galvanostatic charge/discharge curves of a pristine HCF-containing SWB, showing that the charge and discharge potentials

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Fig. 3.25 a Digital photograph of aqueous H2 PtCl6 :CoCl2 solutions with different Pt:Co ratios. b UV–vis spectra of precursor solutions with Pt:Co ratios of 2:1 (green), 1:1 (yellow), and 1:2 (blue). c Schematic CTS-based synthesis of Pt–Co alloys. d Schematic of SWB containing Pt–Co alloy nanoparticles on a carbon current collector. Reproduced with permission of Ref. [88]

equaled 4.10 and 2.64 V versus Na/Na+ , respectively (cf. HCF decomposition potential of ~4.10 V vs. Na/Na+ ). This large potential gap may be caused by the slow kinetics of the pristine HCF [68, 69]. Compared to pristine HCF, that decorated with Pt nanoparticles exhibited a higher ORR activity, featuring a discharge potential of 2.85 V. However, the charge potential of 4.10 V (Fig. 3.26h) was still close to the breakdown potential of carbon, resulting in an inevitable parasitic reaction [82]. A

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Fig. 3.26 SEM images and X-ray photoelectron spectra of a, d pristine, b, e Pt nanoparticle– decorated, and c, f Co nanoparticle–decorated HCF. Galvanostatic charge–discharge plots of SWBs containing g pristine, h Pt nanoparticle–decorated, and i Co nanoparticle–decorated HCF. Reproduced with permission of Ref. [88]

possible explanation for the small influence of metal nanoparticles on ORR activity is that Pt often enhances ORR activity during discharge, which ultimately necessitates the use of an OER catalyst. To increase OER activity, a carbon cathode containing Co nanoparticles was used (Fig. 3.26i). Although the resulting increase in ORR activity was very small, the Co nanoparticles effectively lowered the charge overpotential to 3.92 V. CTS was used to synthesize a bifunctional OER/ORR electrocatalyst by forming Pt–Co alloys with different blending ratios, and the generation of these allows was probed by grazing incidence X-ray measurements. As illustrated in Fig. 3.27a, the high flux of the synchrotron X-ray source allowed us to observe the shift of the peak at q = 2.77 A−1 to a higher value of 2.81 A−1 , which indicated lattice contraction due to Pt–Co alloy formation. These results well agreed with those of conventional XRD measurements and the fast Fourier transform (FFT) of a high-resolution transmission electron microscopy image. Additionally, the Pt 4f spectra of matching samples revealed that the Pt 4f5/2 and Pt 4f7/2 peaks migrated to higher binding energies owing to the loss of the 5d electron in Pt-based alloys, implying the formation of

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Fig. 3.27 a Grazing incidence X-ray patterns of pristine HCF (gray), HCF with Pt (red), and HCF with Pt–Co alloys (Pt:Co = 1:2 (orange), 1:1 (blue), and 2:1 (green)). Magnified grazing incidence X-ray patterns of the region highlighted in yellow (right). b Pt 4f spectra of HCF containing Pt and Pt–Co alloys with Pt:Co ratios of 1:2, 1:1, and 2:1. c–e SEM images and corresponding EDX elemental mappings of Pt and Co in Pt–Co alloys with Pt:Co ratios of c 2:1, d 1:1, and e 1:2. Reproduced with permission of Ref. [88]

a Pt–Co alloy (Fig. 3.27b) [90–93]. The generation of the above alloy was further probed by SEM coupled with EDX mapping, with the SEM images of CTS-treated HCF current collectors following infusion with precursor solutions (Pt:Co ratio = 1:2, 1:1, and 2:1) shown in Fig. 3.27c–e. As seen in plane-view SEM images, the HCF current collectors were uniformly covered by particles with diameters of several tens of nanometers. Pt and Co mapping images further revealed the distribution of Pt and Co atoms in these nanoparticles. The galvanostatic charge–discharge plots of SWBs containing HCF decorated with Pt–Co alloy nanoparticles (Pt:Co = 1:2, 1:1, and 2:1) are shown in Fig. 3.28a. Unlike single-component Pt and Co catalysts, those containing Pt–Co alloys promoted both the OER and the ORR in SWBs. The lowest discharge potential and charge overpotential gap were obtained at a Pt:Co ratio of 2:1. When ORR performance was probed using a revolving disk electrode, the best activity was observed for a Pt:Co ratio of 2:1 (Fig. 3.28c). Thus, the catalyst with the 2:1 Pt:Co ratio achieved a high current density and low overpotential, thus showing good ORR activity, with similar behavior observed for OER activity. The durability and stability of the Pt–Co alloy-SWB were characterized at several current densities (Fig. 3.29). Pristine HCF has an inconsistent cycling capability and could only sustain a current density of 0.2 mA cm−2 . The high charging potential

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Fig. 3.28 a Galvanostatic charge–discharge plots of SWBs containing HCF decorated with Pt–Co alloy nanoparticles having Pt:Co ratios of 2:1 (green), 1:1 (blue), and 1:2 (orange). Plots of b OER and c ORR activities of Pt-, Co-, and Pt–Co alloy (2:1)-containing HCF measured in Ar and O2 . Reproduced with permission of Ref. [88]

Fig. 3.29 Galvanostatic charge–discharge profiles recorded at various current densities for SWBs containing pristine HCF (black) and HCF with Pt–Co alloy nanoparticles (green). Reproduced with permission of Ref. [88]

at high current rates eventually resulted in the decomposition of the carbon current collector, causing severe performance deterioration. Unlike that with a pristine HCF current collector, the SWB with HCF decorated by Pt–Co alloy nanoparticles (Pt:Co = 2:1) could be cycled at 0.3 mA cm−2 for >500 h without substantial performance deterioration. Bifunctional Pt–Co alloy electrocatalysts appear to be important for reducing the overpotential gap. Figure 3.30 a illustrates the CTS-based synthesis of bifunctional Pt–Co alloy electrocatalysts for large-scale prismatic SWB cells. HCF was totally submerged in the Pt:Co = 2:1 aqueous solution to create Pt–Co nanoparticles on a large carbon current collector using CTS and was subsequently supplied with a current of 60 A for 10 s to examine the uniform black body radiation of the current collector (Fig. 3.30b). The formation of Pt–Co alloy nanoparticles on the carbon current collector was confirmed by SEM imaging of four randomly chosen points (Fig. 3.30c–f). The

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Fig. 3.30 a Schematic illustration of a potential SWB application based on a stacked SWB cell as an energy storage system. b Digital photographs of a large-area SWB cell containing an anode compartment (left) and a HCF current collector subjected to CTS treatment (right). c–f Plane-view SEM images of Pt–Co alloy nanoparticles on HCF acquired at four random spots in (b) (right). g Galvanostatic charge–discharge profiles of large-area SWB cells with and without Pt–Co alloy nanoparticles on HCF. Reproduced with permission of Ref. [88]

observed morphology resembled that in Fig. 3.27, which indicated that the CTS approach can be easily upscaled for large-area use. The galvanostatic cycling performance of prismatic SWB cells with and without Pt–Co alloy current collectors is shown in Fig. 3.30g. Notably, the latter cell featured an enhanced electrochemical reaction and a lower overpotential. Thus, the CTS-based decoration of the HCF current collector with Pt–Co alloy nanoparticles not only improved the electrochemical performance of the related SWBs but was also suitable for the fabrication of large-scale prismatic SWB cells. (RuPOM) To ensure the effective operation of rechargeable SWBs, one should develop appropriate electrocatalysts promoting the extremely sluggish four-electron transfer reactions of the OER and the ORR [94]. In view of their molecular structure, structural variety, functional tunability, and high chemical stability, polyoxometalates

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(POMs) are attractive candidates for the development of such catalysts [95–98]. As POMs feature covalently bonded metal oxide clusters, they are intrinsically stable under highly oxidative conditions where most metallic and organic catalysts fail. The ability of POMs to embed a variety of metal ions and their great structural tunability allow them to perform complex redox chemistry and have a wide range of catalytic applications [97]. POMs have the potential to be very atom-efficient electrocatalysts, as all catalytic elements (e.g., Co or Ru) are accessible to external media and can therefore participate in electrochemical processes. Tetraruthenium polyoxometalate (RuPOM) has been investigated as a bifunctional OER/ORR catalyst for SWBs (Fig. 3.31). The electrochemical activity of RuPOM as a bifunctional OER/ORR catalyst in seawater was assessed using a molecular formula of [Ru4 (μ-O)4 (OH)2 (H2 O)4 (γSiW10 O36 )2 ]10– . Until recently, RuPOM has been used only as an OER catalyst [95, 100, 101], while its use as an ORR catalyst has remained undocumented. Historically, most POMs, including RuPOM, were evaluated primarily as homogeneous OER catalysts. For OER/ORR activity evaluation, RuPOM was coated on a glassy carbon electrode with PVDF as a binder in the presence and absence of Ketjen black (KB). Transmission electron microscopy (TEM) and energy-dispersive X-ray spectroscopy (EDS) investigations convincingly demonstrated the molecular nature of RuPOM on the scale of 1–2 nm (Fig. 3.32a, b) and its uniform distribution in the RuPOM/KB mixture (Fig. 3.32c–f). As a benchmark, well-known catalysts such as Pt/C and RuO2 were used. Pt was present in the commercially available Pt/C at a content of 20 wt%. For RuPOM/KB and RuO2 /KB, the active catalyst (i.e., RuPOM or RuO2 ) loading was held constant at 80 wt%. Linear sweep voltammetry (LSV) measurements were used to probe OER activity in N2 -saturated seawater (Fig. 3.33a), showing that RuPOM/KB outperformed other

Fig. 3.31 Schematic illustration of the structure and OER/ORR catalytic activity of RuPOM. Reproduced with permission of Ref. [99]

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Fig. 3.32 Analysis of a, b pure RuPOM and c–f RuPOM/KB by a–e high-resolution TEM and f EDS elemental mapping. The lattice fringes of RuPOM were observed by high-magnification TEM imaging of the locations marked in (b). For EDS elemental mapping of RuPOM/KB, C K, Ru L, W L, and W M edges were analyzed. Reproduced with permission of Ref. [99]

Fig. 3.33 Electrocatalytic activity of RuPOM/KB. Polarization curves of Pt/C, RuPOM/KB, RuPOM, and RuO2 /KB for a the OER and b the ORR. For the ORR test, catalyst-loaded electrodes were characterized using a RRDE. c Numbers of electrons transferred per oxygen molecule in the ORR promoted by different catalysts. d Schematic explanation of the role of KB in ORR. e Comparison between the specific activities of RuPOM/KB, RuO2 /KB, and Pt/C for the OER and the ORR. Reproduced with permission of Ref. [99]

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samples (including a well-known OER catalyst, RuO2 ) in terms of onset potential and current density. The fact that RuPOM devoid of KB and Pt/C exhibited negligible OER activity strongly implies that KB is crucial for enabling charge transfer between the underlying electrode and the RuPOM catalyst. RuPOM/KB, RuPOM, RuO2 /KB, and Pt/C achieved onset potentials of 1.64, 1.67, 1.73, and 1.79 V versus RHE, respectively. It is worth noting that Pt/C can resist corrosion during long-term OER testing [102, 103]. Subsequently, the ORR activities of the above catalysts were determined using a rotating ring-disk electrode (RRDE) in O2 -saturated seawater (Fig. 3.33b). As anticipated, Pt/C exhibited exceptional ORR activity. Surprisingly, the performance of RuPOM/KB was equivalent to that of Pt/C, whereas pure RuPOM showed low activity. Although RuPOM/KB had a somewhat lower (by 124 mV) onset potential than Pt/C, its saturation current density was equivalent to that of Pt/C. These findings reinforce the importance of KB for achieving high electrocatalytic activity. To verify the ORR reaction pathways (i.e., the two- and four-electron pathways leading to the formation of peroxides and water, respectively), the number of electrons transferred per reaction (n) was calculated for different electrocatalysts using RRDE voltammetry data: n=

4I D I D + I R /N

(3.20)

where I D is the disk current, I R is the ring current, and N is the collection efficiency of the ring current. The number of electrons involved in the ORR was strongly dependent on the presence of conductive KB. In the presence of conductive KB, RuPOM, like Pt/C, could enhance four-electron reductive routes in preference to two-electron pathways (Fig. 3.33c, d). On the other hand, RuO2 /KB exhibited low ORR activity, which was most likely related to the low activity of RuO2 rather than to low conductivity. Notably, KB alone had a negligible ORR catalytic activity. Following that, the specific catalytic activities of RuPOM/KB, RuO2 /KB, and Pt/C for the OER and the ORR were examined. RuPOM/KB exhibited a much more specific activity than RuO2 /KB and Pt/C (Fig. 3.33e). For OER, the activity of RuPOM/KB was approximately 24 and 98 times that of RuO2 /KB and Pt/C, respectively, at 1.7 V versus RHE. RuPOM/KB had a greater activity for the ORR per unit mass (g) of noble metal than the best-performing ORR catalysts such as Pt/C; featuring an activity that was around 15 and 6 times that of RuO2 /KB and Pt/C, respectively, at 0.0 V versus RHE. These findings imply that RuPOM may be a feasible substitute for currently available Pt/C catalysts. PVDF binder, which is frequently used in SWB air cathodes, allowed the easy and uniform deposition of RuPOM and KB on HCF (Fig. 3.34a). SEM imaging (Fig. 3.34b, c) verified the uniformity of this deposition and indicated that it resulted in a discernible morphology change. XPS analysis (Fig. 3.34b and c, insets) revealed the presence of a F 1 s peak at 680 eV (PVDF) as well as O 1 s, Ru 3d, Si 2p, and W 4f peaks at 530, 250, 110, and 53 eV, respectively (RuPOM).

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Fig. 3.34 Application of RuPOM/KB in SWBs. a Schematic fabrication of SWB air cathode by the slurry coating method. Inset shows the employed slurry coating solution. b, c SEM images and (insets) X-ray photoelectron spectra of b bare and c RuPOM/KB-coated HCFs. d Charge−discharge profiles of SWB cells containing HCF coated with RuPOM/KB at different RuPOM:KB mass ratios. e Charge−discharge profiles recorded at current densities of 0.1−0.5 mA cm−2 for SWBs prepared using a RuPOM:KB mass ratio of 8:2. Reproduced with permission of Ref. [99]

Encouraged by these findings, we used RuPOM/KB-coated HCF as an air cathode and examined the half-cell performances of the corresponding SWBs during charge– discharge cycling. As expected, the RuPOM/KB coating considerably enhanced the charge–discharge performance of the HCF electrode (Fig. 3.34d), with the optimal RuPOM:KB mass ratio determined as 8:2. Specifically, the deposition of this coating increased the HCF discharge potential from 2.73 to 2.84 V and decreased the charge potential from 4.03 to 3.60 V. Additionally, the optimal RuPOM/KB coating performed well at higher charge–discharge rates (Fig. 3.34e).

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Fig. 3.35 Performance of SWB with RuPOM/KB compared to that of SWB with commercially available Pt/C. Bare and KB-coated HCF were also tested. a Charge−discharge profiles. b Results of stability testing through charge–discharge cycling. c Current-dependent output power densities. Reproduced with permission of Ref. [99]

SWBs with RuPOM/KB-coated HCF as cathodes were compared to those with Pt/C, and a HCF cathode coated solely by KB was evaluated as a control. Notably, KB had a minor influence on charge and discharge potentials in the absence of RuPOM. On the other hand, the modification of HCF with RuPOM/KB or Pt/C greatly reduced the gap between charge and discharge voltages. RuPOM/KB was more successful than Pt/C at charging and somewhat less effective at discharging, thus featuring an overall performance close to that of Pt/C (Fig. 3.35a). It is worth noting that despite the better specific catalytic activity of RuPOM/KB, it performed similarly to Pt/C, as both catalysts were loaded using the same mass of catalytic slurry. For instance, although Pt accounted for 20 wt% of Pt/C, the catalytically active Ru constituted only 4.85 wt% of RuPOM/KB (8:2). Charge–discharge cycling was used to determine the stability of HCF cathodes. As seen in Fig. 3.35b, both RuPOM/KB- and Pt/C-modified HCF cathodes exhibited reliable performances after 50 cycles, featuring only small performance deterioration. These findings suggested that the observed performance decline was due to the degradation of elemental carbon in HCF rather than to RuPOM/KB deterioration. The charge–discharge voltage gap increased from 0.76 to 0.81 V after 50 cycles. Moreover, modification with RuPOM/KB significantly increased the output power density of the HCF cathode (Fig. 3.35c), and the maximal power density of 12.3 mW cm−2 was comparable to that obtained for modification with Pt/C (12.5 mW cm–2 ). These data clearly indicate the superior performance of RuPOM/KB-modified HCF. Non-precious metal–based electrocatalysts (GNS/MC) The high cost of precious metals such as Pt, Ir, and Ru has inspired the development of novel low-cost bifunctional electrocatalysts such as nanostructured carbons doped with heteroatoms and transition metals [104–115]. In this section, we discuss the structure of highly integrated, bifunctional oxygen electrocatalysts based on graphitic nanoshell/mesoporous carbon (GNS/MC) nanohybrids. Figure 3.36a presents the fabrication of bifunctional GNS/MC electrocatalysts from Ni and Fe phthalocyanines (NiPc and FePc) by nanocasting. Typically, the metal Pc (1 g) was pyrolyzed with SBA-15 silica at 1000 °C in N2 , and the silica

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Fig. 3.36 a Schematic synthesis, b TEM images, and c HR-TEM image of GNS/MC. Reproduced with permission of Ref. [116]

template was then etched to afford a carbon-based product. Specifically, Ni-MC, FeMC, and Nix Fey -MC nanostructures were obtained from NiPc, FePc, or their mixture. For comparison, OMC was prepared from metal-free Pc as a precursor. The sample prepared using a NiPc:FePc mass ratio of 3:7 exhibited the best OER/ORR activity. The related SEM (Fig. 3.36b, inset) and TEM (Fig. 3.36b, c) images revealed the presence of highly graphitic nanoshells with a 50-nm hollow core. The corresponding HR-TEM image (Fig. 3.36c) showed the presence of 8-nm-thick graphitic nanoshells embedded into mesoporous carbon structures. The bifunctional electrocatalytic activities of GNS/MC, Ni-MC, and Fe-MC catalysts were examined using RRDE measurements in 0.1 M KOH saturated with O2 . The electrocatalytic activities of OMC, Pt/C (20 wt% Pt), and Ir/C (20 wt% Ir) were determined as references. LSV polarization curves for the OER and the ORR (Fig. 3.37a) as well as data pertaining to oxygen electrode activity (Fig. 3.37b) demonstrated the superiority of GNS/MC among the evaluated samples. In the case of the OER, GNS/MC exhibited an overpotential of 340 mV at a current density of 10 mA cm−2 , which was much less than the overpotentials of Ni-MC (480 mV) and Fe-MC (500 mV). This behavior demonstrates the synergetic effects of Ni and Fe metals in enhancing OER activity. Bifunctional oxygen electrode activity was defined as the difference between the OER potential at 10 mA cm−2 and the ORR potential at −3 mA cm−2 . According to

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Fig. 3.37 a OER and ORR polarization curves of GNS/MC, Ni-MC, Fe-MC, OMC, Ir/C, and Pt/C catalysts. OER curves were obtained via CV cathodic sweeping. ORR LSV curves are presented after background current correction. b Oxygen electrode activities (E OER − E ORR ) of GNS/MC, Ni-MC, Fe-MC, OMC, Ir/C, and Pt/C catalysts. Reproduced with permission of Ref. [116]

this definition, the activity of GNS/MC (0.72 V) was the highest, greatly surpassing those of Ni-MC (0.96 V), Fe-MC (0.88 V), Ir/C (0.91 V), and Pt/C (1.01 V). The Tafel slopes of GNS/MC (OER: 80 mV per decade, ORR: 50 mV per decade) were comparable to those of Ir/C (90 mV per decade for the OER) and Pt/C (60 mV per decade for the ORR). Comparison of the oxygen electrode activities obtained herein with those of previously reported high-performance catalysts suggests that GNS/MC is one of the most effective bifunctional ORR/OER catalysts (Fig. 3.38). Specifically, compared to doped carbon catalysts, GNS/MC (0.72 V) exhibited an unprecedentedly high oxygen electrode activity. Furthermore, the activity of GNS/MC exceeded that of the most active transition metal–based Co3 O4 /N-rmGO catalyst (0.71 V) despite the fact that the former contains less metal (0.39 wt%) than the latter (70 wt%). Additionally, GNS/MC exhibited remarkable durability for the OER and the ORR, as revealed by chronopotentiometric measurements (Fig. 3.39a). Notably, the OER activity of this catalyst was sustained for up to 20 h after the initial voltage decrease, which was indicative of high durability in oxidizing environments. In the ORR durability test, GNS/MC was subjected to 30,000 potential cycles between 0.6 and 1.0 V (Fig. 3.39b) and exhibited exceptional durability, as evidenced by a nearly unchanged CV curve, a negligible negative shift of E 1/2 by 15 mV, and retained morphology. The origins of the exceptional catalytic activity and endurance of GNS/MC were further examined using X-ray absorption spectroscopy (XAS). Figure 3.40 presents the X-ray absorption near-edge structure (XANES) spectra of GNS/MC and Ni reference samples measured in situ in 0.1 M KOH. At an open-circuit voltage (OCV), GNS/MC and Ni-MC featured lower peak intensity at 8337 eV than the NiPc precursor but exhibited a new peak at 8332 eV (Fig. 3.40a). These concurrent alterations implied the coordination of one or more axial ligands to the Ni centers in

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Fig. 3.38 Comparison of GNS/MC with previously reported catalysts (including carbon-, metal oxide-, and perovskite-based catalysts) in terms of oxygen electrode activity. Reproduced with permission of Ref. [116]

Fig. 3.39 a OER chonoamperometric response of GNS/MC. The noise and drifts in the recorded potential during the OER durability test are attributed to the accumulation of O2 bubbles at the electrode. b ORR polarization curves of GNS/MC before and after 30,000 potential cycles. Reproduced with permission of Ref. [116]

GNS/MC and Ni-MC [117]. When a potential of 1.8 V was supplied (OER condition), the white line in the XANES spectrum of GNS/MC exhibited an intensity rise, which suggested the additional oxidation of Ni centers, which was ascribed to the adsorption of reactive oxygenated species. The in situ Ni K-edge extended X-ray absorption fine structure (EXAFS) spectrum of GNS/MC (Fig. 3.40b) recorded at OCV exhibited a prominent peak at ~1.5 Å,

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Fig. 3.40 a Ni K-edge XANES and b EXAFS spectra of GNS/MC, NiPc, and Ni references. Reproduced with permission of Ref. [116]

which was ascribed to metal–N bonding [118]. Interestingly, this peak gained intensity when an oxidative potential of 1.8 V was applied, which indicated the generation of oxygenated adsorbates (O–Ni–Nx ). Notably, the intensity of this peak was higher for GNS/MC than for Ni-MC, possibly because of the higher OER activity of the former catalyst due to the presence of both Ni and Fe. Figure 3.41a and b present the results of in situ Fe K-edge XAS measurements, revealing that compared to that of FePc, the EXAFS spectrum of GNS/MC featured a considerably less intense Fe–N peak at 1.4 Å (Fig. 3.41b) and exhibited a new peak at 2.2 Å. These findings indicate that during high-temperature pyrolysis, the Fe–N moieties were partly converted into Fe metal, oxide, or carbide phases undetectable by HR-TEM and XPS. According to recent works on high-performance Fe–N/C ORR catalysts, the conserved Fe–N moieties and Fe carbide phases may contribute

Fig. 3.41 a Fe K-edge XANES and b EXAFS spectra of GNS/MC, FePc, and Fe references. The EXAFS and XANES spectra of GNS/MC were collected under in situ electrochemical conditions. Reproduced with permission of Ref. [116]

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Fig. 3.42 a SWV curves of GNS/MC, Ni-MC, Fe-MC, and OMC recorded in N2 -saturated 0.1 M KOH and b corresponding expansions. Reproduced with permission of Ref. [116]

to the high ORR activity of GNS/MC and Fe-MC catalysts. According to XAS data, the OER was largely catalyzed by Ni species with Fe serving as a cofactor, whereas the ORR was promoted by Fe species. Metal–N coordination was further explored using square wave voltammetry (SWV) measurements (Fig. 3.42). The SWV curves of GNS/MC and Ni-MC featured a minor peak above 1.4 V due to the oxidation of metal–N ligands [119]. In contrast, neither Fe-MC nor OMC exhibited an oxidation peak. Interestingly, the ligand oxidation peak shifted to higher potentials upon going from Ni-MC to GNS/MC, which was attributed to the concomitantly increasing strength of axial OH group binding. Thus, the Ni–N coordination in GNS/MC was more conducive to OH group adsorption than that in the Ni–MC catalyst, which provided more evidence for the superior OER activity of GNS/MC. The anodic shift of Ni redox peaks was also observed for an Fe-doped nickel oxyhydroxide (Ni1−x Fex OOH) catalyst [120]. The graphite contents of catalysts were determined by Raman spectroscopy (Fig. 3.43). The spectra of all catalysts exhibited two prominent peaks corresponding to disordered (D-band, 1350 cm−1 ) and graphitic (G-band, 1580 cm−1 ) carbon. The corresponding intensity ratio (I D /I G ) increased in the order of GNS/MC (0.51) < FeMC (0.86) < Ni-MC (0.98) < OMC (0.99), i.e., the shell of GNS/MC had the highest graphitization degree. Moreover, GNS/MC exhibited a greater number of clearly defined Raman peaks above 2250 cm−1 , which indicated that the graphitic shell of this catalyst contained numerous defects [121]. The shift of the G-band to lower wavenumbers in the case of GNS/MC (Fig. 3.43b) was ascribed to the presence of stretched or bent graphitic surfaces. Notably, earlier research has demonstrated that the presence of N atoms in graphitic carbon defects can endow carbon-based catalysts with increased OER activity [115]. Additionally, the highly graphitic structure of GNS/MC was held responsible for its exceptional longevity.

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Fig. 3.43 a Raman spectra of GNS/MC, Ni-MC, Fe-MC, and OMC and b expansions around the G-band. Reproduced with permission of Ref. [116]

According to the above explanation, the synergistic effect of Ni and Fe as well as the graphitic nanoshell structure may contribute to the high catalytic activity and durability of GNS/MC. It should be emphasized that no direct correlation could be established between metal content and oxygen electrode activity. For instance, in a series of Ni3 Fe7 -MC catalysts prepared by pyrolysis at 800–1200 °C, the amount of both Fe and Ni steadily decreased with increasing pyrolysis temperature, yet the greatest oxygen electrode activity was observed for Ni3 Fe7 -MC prepared at 1000 °C. It is thought that various factors, including metal type, conductivity, and porosity collectively contribute to OER/ORR activity. GNS/MC was subsequently used as an air cathode of an aqueous Na-air battery. Figure 3.44a illustrates the charge–discharge curves of aqueous Na–air batteries with various cathodes. The voltage gap, i.e., the difference between charge and discharge voltages, increased in the order of GNS/MC (115 mV) < Pt/C (179 mV) < Ir/C (364 mV) < catalyst-free carbon paper (698 mV). This behavior demonstrated the superiority of GNS/MC in the full-cell arrangement and was also observed in halfcell tests performed using the RRDE method, although the voltage difference order observed in the latter case was different (GNS/MC > Ir/C > Pt/C). Figure 3.44b presents the charge–discharge curves of the aqueous Na–air battery with GNS/MC-coated carbon paper, demonstrating that this battery could be reversibly charged and discharged. The discharge potential was almost fully retained after 10 cycles, demonstrating the high stability of the ORR activity of GNS/MC. To the best of our knowledge, this is the first report of a rechargeable aqueous Na–air battery. Thus, GNS/MC was concluded to be a promising electrocatalyst for aqueous Na–air batteries, offering the benefits of low voltage difference, high round-trip efficiency, high energy density, and outstanding charge–discharge stability. In summary, GNS/MC nanohybrids were highly active and durable bifunctional OER/ORR electrocatalysts, and their high performance was ascribed to the presence of several structural motifs. Specifically, the optimal GNS/MC catalyst with highly

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Fig. 3.44 a Voltage differences between the charge and discharge plateaus of aqueous Na–air batteries containing carbon paper coated with various electrocatalysts. b Charge–discharge curves of aqueous Na–air battery with GNS/MC-coated carbon paper recorded for 10 cycles. Reproduced with permission of Ref. [116]

Fig. 3.45 Schematic synthesis of Co3 V2 O8 nanoparticles and their use in a SWB. Reproduced with permission of Ref. [20]

graphitized defect-rich nanoshells and low quantities of Ni and Fe entities exhibited one of the greatest performances (high OER and ORR activities as well as a high oxygen electrode activity of 0.72 V) among carbon-based catalysts, outperforming Ir/C and Pt/C catalysts. (Co3 V 2 O8 ) Catalysts based on mixed transition metal oxides are promising alternatives to noble metal-based catalysts, offering the advantages of lower cost, adjustable composition (which allows them to outperform single-metal oxides in terms of OER activity), and the synergetic effects of nanostructures [86, 122, 123]. Among these catalysts, cobalt-based oxides exhibit superior activity and stability [109, 124–129]. Recently, metal oxide–supported vanadium species with variable oxidation states and redox characteristics have been actively explored as oxidation electrocatalysts. Thus, the synergy between catalytically active cobalt oxides and redox-active vanadium oxides is expected to afford electrocatalysts with enhanced performance.

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Figure 3.45 illustrates the synthesis of Co3 V2 O8 nanoparticles via precipitation/annealing and the construction of a SWB cell with a Co3 V2 O8 -coated carbon current collector. An aqueous solution of NH4 VO3 was treated with a solution of LiOH, and the mixture was supplemented with a solution of CoCl2 . In view of the increasing surface charge of metal oxide precipitates, they were generated evenly without notable aggregation [130, 131]. The high surface energy of atoms on the particle surface during the first-step stirring at 80 °C made them less stable than those in the interior. As a result, the particles interacted with the surrounding molecules and lowered the surface energy via Ostwald ripening, thus growing in size [132]. In the second stage, the mixed-phase precipitates were annealed at 350 °C in air to generate crystalline Co3 V2 O8 . Figure 3.46a presents a field-emission scanning electron microscopy (FESEM) image of as-prepared Co3 V2 O8 nanoparticles, revealing that annealing induced the formation of clusters (100–200 nm) without strongly affecting the morphology of individual particles. The voids of the produced network structure possibly provided numerous active sites for the OER and the ORR. The HRTEM imaging of Co3 V2 O8 nanoparticles revealed lattice fringes with an interplanar spacing of 0.254 nm, which corresponded to the (122) plane of orthorhombic Co3 V2 O8 (Fig. 3.46b). The spots observed in the related FFT pattern were assigned to the (002), (122), (042), and (442) planes of orthorhombic Co3 V2 O8 (inset in Fig. 3.46b). The XRD pattern of as-synthesized nanoparticles (Fig. 3.46c) confirmed the presence of orthorhombic Co3 V2 O8 . Scanning transmission electron microscopy (STEM)-EDS elemental mappings (Fig. 3.47a) revealed that Co, V, and O were uniformly distributed throughout nanoparticles. The chemical structure of Co3 V2 O8 nanoparticles was further probed by XPS, and the related wide-scan spectrum revealed no detectable contaminants. Figures 3.47b and c show the narrow-scan Co 2p and V 2p spectra of the above nanoparticles deconvoluted using a Gaussian–Lorentzian function. The primary signals of Co 2p3/2 and Co 2p1/2 were detected at 782.2 and 798.1 eV, respectively, each with a satellite peak. The deconvoluted Co 2p3/2 peak confirmed the coexistence of Co2+ (781.8 eV) and Co3+ (783.9 eV) on the nanoparticle surface, indicating that Co2+ is the dominant oxidation state. The main peaks of V 2p3/2 and V 2p1/2 were

Fig. 3.46 a FESEM image, b HRTEM images and FFT pattern, and c XRD pattern of Co3 V2 O8 nanoparticles. Reproduced with permission of Ref. [20]

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Fig. 3.47 a STEM-EDS mapping images, b Co 2p spectrum, and c V 2p spectrum of Co3 V2 O8 nanoparticles. Reproduced with permission of Ref. [20]

observed at 518.6 and 526.2 eV, respectively. The deconvoluted V 2p3/2 peak showed the co-presence of V4+ (517.1 eV) and V5+ (518.6 eV), with V5+ being the dominant state [133]. Subsequently, we examined the performance of the SWB cell with Co3 V2 O8 nanoparticle–coated HCF as a cathode current collector. Figure 3.48a presents the initial discharge–charge voltage profiles of SWBs with bare and Co3 V2 O8 -coated HCF, revealing that the deposition of Co3 V2 O8 reduced polarization in both charged

Fig. 3.48 Comparison of SWB cells with bare HCF (black) and Co3 V2 O8 -coated HCF (red). a Initial discharge–charge profiles, b discharge–charge profiles recorded at different current densities, c polarization curves (symbols) and power characteristics (lines), and d cycling performance (20 cycles, 400 h in total). Reproduced with permission of Ref. [20]

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and discharged states, i.e., the charge voltage dropped from 3.88 to 3.76 V, while the discharge voltage rose from 2.80 to 2.87 V. Thus, the voltage efficiency of the cell with Co3 V2 O8 (76%) exceeded that observed without this catalyst (72%). Rate capacity testing was performed at current densities of 0.01–0.5 mA cm−2 . As demonstrated in Fig. 3.48b, when the current density increased by a factor of 50, the voltage gaps between the charge and discharge curves also increased. The Co3 V2 O8 cell had a lower voltage gap than the pure-HCF cell, thus featuring electrocatalyst-enhanced kinetics. Figure 3.48c compares the polarizations and power densities of these cells, showing that the maximal power densities of pure-HCF and Co3 V2 O8 cells were 5.1 and 5.9 mW cm−2 , respectively. Thus, the latter cell outperformed the former in terms of discharge ability, as Co3 V2 O8 enhanced the kinetics of the OER and the ORR in seawater. Figure 3.48d presents the results of charge–discharge cycling (20 cycles, 400 h in total), revealing that the Co3 V2 O8 cell demonstrated an increase in charge voltage with increasing cycle number while featuring a lower voltage gap than the bare HCF cell. Both bare and Co3 V2 O8 cells had an average voltage gap of 1.17 V, which was indicative of high OER/ORR activity and stability in seawater.

3.2.3 Cl− Capture/Release 3.2.3.1

Ag Electrodes

The generation of gases on SWB electrodes results in significant internal polarization and, hence, in a considerable difference between charge and discharge voltages (1.0 V) [9, 134, 135], which hinders the further development of SWBs and metal– air batteries. In this regard, a chloride ion–capturing electrode (CICE) containing Ag foil is used as a cathode. During charge–discharge, Ag may reversibly react (capture/release) with Cl− ions in the seawater catholyte at 2.93 V versus Na+ /Na, which allows one to bypass the OER/ORR pathway and thus considerably reduce the voltage gap and increase voltage efficiency. The operation of CICE-containing SWBs can be summarized as follows. During charging, Na+ ions diffuse from seawater to the anode via the solid electrolyte, whereas Cl− ions diffuse and interact with the Ag cathode to produce AgCl (top right in Fig. 3.49). During discharge, the opposite reaction happens (bottom right in Fig. 3.49). The overall reaction occurring during charge–discharge cycling is given by Eq. 3.21: Ag(s) + NaCl(aq) ↔ Na(s) + AgCl(s) ,

E ◦ = 2.932 V vs. Na+ /Na

(3.21)

This mechanism was validated by comparing the galvanostatic charge–discharge behaviors of SWBs with a CICE and a carbonaceous cathode. Figure 3.50a illustrates the initial galvanostatic charge–discharge curves of the CICE-SWB, revealing that it featured a voltage gap of 291 mV and charge and discharge plateaus (3.09 and 2.80 V, respectively) located near the predicted cell voltage of 2.932 V. The charge and discharge overpotentials (158 and 133 mV, respectively) indicated that the redox

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Fig. 3.49 Schematic cell structure and operation mechanism of a CICE-containing SWB. Reproduced with permission of Ref. [13]

Fig. 3.50 a Galvanostatic first charge–discharge profiles of SWBs with a Ag electrode (red) and a carbonaceous electrode (black), b XRD patterns of pristine Ag foil and Ag electrodes after the first charge and discharge. Reproduced with permission of Ref. [13]

reaction of Ag/AgCl in the SWB was reversible. The SWB with a carbonaceous cathode had a considerable voltage gap (1.58 V) and featured charge and discharge voltage plateaus located at 4.07 and 2.49 V, respectively, which was ascribed to substantial OER/ORR overpotentials. The specific charge and discharge capacities of the CICE-SWB equaled 119.9 and 118.3 mAh g−1 , respectively. The phase, morphological, and compositional changes of the CICE during charge and discharge were probed using XRD, SEM, and EDS. Figure 3.50b shows the

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Fig. 3.51 Plan-view SEM images and EDS data of a pristine Ag foil and b–e Ag electrodes after first charge (b, c) and discharge (d, e). f Schematic surface morphology change of the Ag electrode during charge and discharge. c and e show the enlarged image and the EDS results of the regions indicated by symbols in (b) and (d). Reproduced with permission of Ref. [13]

XRD patterns of the Ag electrode before and after the first charge–discharge cycle, revealing that after charging, the sample comprised primarily cubic AgCl with a trace of cubic Ag from the foil itself. After discharge, the AgCl phase vanished, and only Ag was left. The formation/removal of the AgCl phase upon charge–discharge suggested the occurrence of a reversible redox interconversion of Ag and AgCl on the CICE of the SWB. The surface morphology and elemental composition of the CICE before and after the first cycle are shown in Fig. 3.51a–e. The surface of pristine Ag foil (Fig. 3.51a) showed many parallel microstretches, which were presumably produced during manufacture and/or transportation as a roll. Following the initial charge (Fig. 3.51b), the surface became coated with spherical particles (1–7 μm). Elemental mapping revealed the precipitates to be AgCl (red box in Fig. 3.51b), which agreed with XRD data. Following the initial discharge, AgCl disappeared, and the electrode surface became covered with irregular and rectangular Ag particles (Fig. 3.51d and e). This roughened surface was created during the reduction of AgCl to Ag and Cl as well as during the elimination of Cl− during discharge. The above morphological change might be due to the oxidative etching of the Ag surface caused by the Cl− ions and O2 dissolved in seawater during discharge [136]. Notably, the Ag cathode repeatedly experienced the formation and removal of AgCl precipitates throughout subsequent cycles, which afforded a substantially rough Ag surface featuring multiple nanoparticles. Figure 3.51f schematically illustrates these findings. Charging resulted in the formation of spherical AgCl precipitates on the CICE surface, which were then reduced to Ag and Cl− during discharge to afford a roughened surface with rectangular Ag nanoparticles.

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Figure 3.52a illustrates the galvanostatic charge–discharge performance of the CICE-SWB during 20 cycles at 4.76 mA g−1 (500 h in total). Throughout these cycles, the cell demonstrated steady performance, i.e., the initial voltage gap of 289 mV decreased to 263 mV by the 20th cycle. Figure 3.52b illustrates the evolution of discharge capacity retention, voltage efficiency, and Coulombic efficiency upon the cycling of the CICE-SWB, showing that discharge capacity retention was maintained at 99.8%, while voltage efficiency was maintained at 90.3–93.5% over 20 cycles. Thus, the CICE-SWB could be reversibly charged and discharged, outperforming SWBs with carbonaceous cathodes and conventional metal–air batteries. Previously reported SWBs utilizing OER/ORR pathways showed a low first-cycle voltage efficiency of 73% because of the significant OER/ORR overpotentials at the cathode [9, 137]. During cycling, the Coulombic efficiency of the CICE-SWB was maintained at 98.6–98.7%. As previously stated, these outstanding voltage and Coulombic efficiencies were attributed to the reversible redox reaction of Ag/AgCl at the cathode. After 20 cycles, the surface of the Ag electrode became much rougher because of the

Fig. 3.52 Characterization of the CICE-SWB. a Cycling performance (20 cycles), b evolution of capacity retention (black triangles), voltage efficiency (red circles), and Coulombic efficiency (violet triangles) during cycling. c Plan-view SEM images and d EDS spectrum of the Ag electrode after 20 cycles. The EDS spectrum was acquired for the area marked by a yellow cross in (c). Insets reveal that the surface of the Ag electrode after 20 cycles became considerably rougher, while the particle size was reduced to a submicron scale, which induced an increase in the electrode surface area. Reproduced with permission of Ref. [13]

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deposition of submicron-sized particles (Fig. 3.52c and d). The voltage efficiency of the CICE-SWB (90.3%) exceeded those of previously reported metal–air batteries and OER/ORR-based SWBs [8, 9, 134, 138–141].

3.2.4 Na+ Intercalation/Deintercalation 3.2.4.1

Nickel Hexacyanoferrate (NiHCF)

Aqueous rechargeable SWBs have drawn considerable interest due to their low cost, environmental friendliness, and safety [142]. One of the key drawbacks of SWBs is their relatively low energy density, which is caused by the limit imposed on the operating voltage (~2 V) by water decomposition. Thus, much effort has been directed at increasing cell voltage. Figure 3.53 displays the potential intercalation/deintercalation cathodes for SWBs. NiHCF was used as a suitable cathode material in seawater as the electrolyte in this proof of concept investigation. NiHCF has a Prussian blue (PB) hexacyanometalate structure (face-centered cubic lattice) in which the Ni2+ ions are six-fold coordinated by nitrogen atoms while the Fe2+ ions are octahedrally coordinated by the carbon atoms of CN ligands. This arrangement affords a broad ionic channel parallel to the axis, allowing for the easy insertion and extraction of Na+ ions [143]. Most PB-based compounds contain coordinated water and cannot therefore be used in conventional non-aqueous electrolytes [144]. However, this is not the case for SWBs, which use seawater as the catholyte. Additionally, PB-based materials are cheap, easy to manufacture, and display excellent cycling stability in both aqueous and non-aqueous environments [143, 145–147].

Fig. 3.53 Schematics of possible intercalation/deintercalation cathodes for SWBs. Reproduced with permission of Ref. [18]

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NiHCF nanoparticles were synthesized using a straightforward co-precipitation technique. The XRD pattern of the prepared sample (Fig. 3.54a) revealed a high degree of crystallinity, and all peaks were indexed to the face-centered cubic PB phase. As shown in Fig. 3.54b and c, the NiHCF nanoparticles exhibited a consistent size of 30–50 nm. This very small size may lower the diffusion length of Na+ in the solid phase, hence improving the SWB rate performance. According to ICP-OES measurements, the composition of NiHCF was identified as Na0.61 Ni1.78 Fe(CN)6 ·nH2 O. Figure 3.54d presents a SEM image of NiHCF-coated carbon paper, demonstrating the uniform dispersion of well-attached NiHCF particles (see inset). Equation 3.22 summarizes the electrochemical process occurring in SWBs: Na0.61 Ni1.78 FeII (CN)6 ↔ Na0.61−x Ni1.78 FeIII (CN)6 + xNa+ + xe−

(3.22)

During charge, Na+ ions are deintercalated from NiHCF, Fe2+ is oxidized to Fe3+ , and the electrons are released via an external circuit. The deintercalated Na+ ions move through seawater, NASICON, and a non-aqueous electrolyte before being reduced by electrons from the external circuit and deposited at the anode. The opposite process occurs during discharge. Notably, the solid structure of the NASICON

Fig. 3.54 a XRD pattern of NiHCF, b SEM image of NiHCF, c TEM image of NiHCF, and d SEM image of NiHCF-coated carbon paper. Inset shows a magnified view of a single carbon fiber. Reproduced with permission of Ref. [18]

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Fig. 3.55 CV curves of a SWB with a NiHCF cathode recorded a during the first 10 cycles and b at scan rates of 2.8–3.8 V. Inset shows the dependence of peak current on the square root of scan rate. Reproduced with permission of Ref. [18]

separator does not allow the passage of protons, water, and hydrated/solvated ions [8, 148]. CV and galvanostatic charge–discharge measurements were used to characterize the electrochemical parameters of a SWB with a NiHCF cathode. Figure 3.55a presents the CV profiles for the first 10 cycles in the potential range of 2.8–3.8 V versus Na/Na+ , revealing a pair of well-defined symmetric peaks at 3.25 and 3.46 V corresponding to the intercalation and deintercalation of Na+ , respectively. According to this figure, the deintercalation of Na+ ions starts at 3.1 V, and the position of the intercalation peak remains almost constant over 10 cycles, which suggests a highly reversible redox process. During Na+ insertion, Fe2+ is oxidized to Fe3+ , whereas Fe3+ is reduced to Fe2+ during Na+ extraction (Eq. 3.22). However, in view of the narrow voltage range of 2.8–3.8 V employed for CV measurements, no redox reactions of Ni2+ ions were detected [146]. Figure 3.55b presents the CV curves obtained at scan rates of 0.1–0.5 mV s−1 . The absolute values of anodic and cathodic peak currents linearly increased with the square root of scan rate, i.e., the battery’s redox reaction was diffusion-controlled (inset in Fig. 3.55b), which is typical of intercalation/deintercalation electrodes [149]. The galvanostatic charge–discharge voltage profiles of the SWB using NiHCF are shown in Fig. 3.56a. The sloping areas at 3.2–3.64 V on the charge curve and at 3.6–3.15 V on the discharge curve corresponded to the Fe2+ /Fe3+ redox pair [146], in line with the results of CV measurements. The battery had a maximum capacity of 56.58 mAh g−1 at 20 mA g−1 during the first discharge. This capacity was computed using the mass of NiHCF [146] and was comparable to that obtained in earlier investigations for NiHCF and zinc hexacyanoferrate cathodes [150]. Additionally, the current battery system was evaluated at various current densities (2, 30, 40, 50, 80, and 100 mA g−1 ; Fig. 3.56b and c). When the current density increased from 20 to 100 mA g−1 , the capacity decreased from 56.58 to 19.8 mAh g−1 (Fig. 3.56b). The former capacity value was comparable to those determined for other (non-)aqueous

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Fig. 3.56 Galvanostatic charge–discharge profiles recorded a over the first 10 cycles and b at different current densities (20–100 mA g−1 ); c rate performance and d cycling performance of the SWB with a NiHCF cathode. Reproduced with permission of Ref. [18]

battery cathode materials. After 80 cycles, a capacity of ~46 mAh g−1 was recovered when the current density was lowered to the original value of 20 mA g−1 . These findings revealed the superior capacity retention performance of NiHCF (Fig. 3.56c). Figure 3.56d illustrates the cycling performance and Coulombic efficiency of the SWB system at a current density of 20 mA g−1 . After 200 cycles, the capacity reached 34.53 mAh g−1 , which accounted for 62% of the initial capacity (56.58 mAh g−1 ). After several cycles, the Coulombic efficiency increased from 95 to 98%, mostly because of the concomitant disintegration of interstitial water molecules in the structure [151, 152]. Charging induced the decomposition of interstitial water molecules within NiHCF, resulting in a low Coulombic efficiency during the first few cycles. However, after this decomposition was completed, a high Coulombic efficiency of 98% was sustained in subsequent cycles. After 200 cycles, the spent NiHCF cathode was replaced with a fresh one. Surprisingly, the cell with the fresh cathode showed a capacity of 56.5 mAh g−1 during the first cycle, which is close to 99.8% of the initial discharge capacity of the spent cathode. Following 200 cycles, the cell retained 61.4% of its capacity (60.49% of the initial discharge capacity of the spent cathode). These results indicate the outstanding replaceability of the investigated cathode and the reusability of the related battery. A SWB was fabricated with NiHCF as the cathode and sodium as the anode. The well-defined open framework created by the covalent bonding of CN ligands to the

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Fig. 3.57 Charge–discharge profiles of SWBs with the Na-ion intercalating NiHCF cathode and an OER/ORR current collector made of HCF. Reproduced with permission of Ref. [46]

octahedral Fe and tetrahedral Ni sites endowed the NiHCF cathode with excellent electrochemical activity and a reversible capacity of 56.58 mAh g−1 . Figure 3.57 illustrates the charge–discharge profiles of cells with a Na-ion intercalating NiHCF cathode and an OER/ORR current collector made of HCF. In view of the extremely reversible and rapid Na-ion intercalation/deintercalation capabilities of NiHCF, the fabricated SWB demonstrated a comparatively small voltage gap between charge and discharge curves and a typical discharge voltage of 3.4 V. Along with that of the NiHCF cathode, the use of naturally available seawater as the catholyte makes the developed SWB environmentally benign and cheap, allowing for deployment in large-scale EESs. The presented SWB exhibited an average discharge voltage of 3.4 V, excellent cycling stability, and a Coulombic efficiency of 98% after 200 cycles. Additionally, in view of its open nature, the cathode could be simply replaced at the end of its cycle lifetime.

3.2.4.2

Dual-Mode Electrodes (PB/CF)

The use of Na+ intercalating/deintercalating materials such as NiHCF results in a smaller voltage gap compared to that observed for the OER/ORR process. However, the related energy output is restricted by the existing cathode chemistry. In light of the above, a dual-mode seawater energy extraction system has been developed, offering high power density due to Na-ion intercalation and high energy due to the reduction of DO [153]. The key to realizing this dual-mode system is the use of a cyanometallate coordination polymer (PB, Mx Fe[Fe(CN)6 ]y ·nH2 O, where M = alkali metal cation, 0 < x < 2, 0 < y < 1) as an electrode to accept Na+ ions in their interstitial sites via a redox reaction between the Fe2+ /Fe3+ nodes in a narrow potential window or via electron transfer to DO in seawater [151]. As the coordination polymer electrode directly contacts seawater, it should be sufficiently flexible to respond to the dynamic and unpredictable flux of seawater. CF

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Fig. 3.58 Demonstration of the flexibility of the PB film a adhesively bound to CF and b grown on CF. Reproduced with permission of Ref. [154]

may be used as a carrier for PB particles [153]. However, unlike CF, the coordination polymer particles and adhesives are not flexible, and the numerous interfaces between the adhesives, PB, and CF make the overall electrode brittle (Fig. 3.58a). Typically, electrodes fail when they are bent. Thus, the manufacture of a flexible cyanometallate coordination polymer felt and its usage as an electrode of a dual-mode sea-water energy extraction system have been investigated. The deposited layer grew on the electrochemically modified CF. The modified surface featured reactive hydroxyl and carboxylic acid groups capable of forming chemical bonds with the PB layer (Fig. 3.58b). Chemical bonding enabled the close adherence of the PB layer to CF, allowing the composite electrode to withstand deformation and the effects of seawater flow during power generation. CF was used as a substrate because of its high electron conductivity and resilience in seawater [153, 155, 156]. The original CF contained interwoven fibers with a smooth surface (Fig. 3.59a), was impervious to wetting by the reactant solution, and featured few functional groups available for chemical linking with PB. The interaction of PB with CF can be enhanced through the modification of the latter surface with hydroxyl or carboxyl groups, which can engage in chemical interactions with the exposed iron ions on the surface of PB crystals. As the superior electron conductivity of CF ensures that electrons may reach each fiber, surface oxidation was predicted to be uniform. SEM imaging indicated that the smooth surface of unmodified CF became rough after electrochemical alteration (Fig. 3.59b), indicating the occurrence of a certain chemical transformation. The nature of functional groups on the CF surface was probed by XPS. The C 1 s spectrum of modified CF featured a larger shoulder peak with a small shift (0.5 eV) to higher energies and two extra broad peaks at 286.18 and 288.28 eV (Fig. 3.60a). These signals could be deconvoluted into four peaks centered at 284.2, 285.6, 286.4, and

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Fig. 3.59 SEM images of (a-1, a-2) unmodified CF, (b-1, b-2) electrochemically treated CF, and (c-1, c-2) the PB-CF composite. Reproduced with permission of Ref. [154]

Fig. 3.60 a C 1 s core-level spectra of modified and raw CF. b Deconvoluted C 1 s spectrum of modified CF. c O 1 s core-level spectra of modified and raw CF. d Deconvoluted O 1 s spectrum of modified CF. Reproduced with permission of Ref. [154]

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Fig. 3.61 a XRD patterns and b FTIR spectra of CF before and after PB deposition. Reproduced with permission of Ref. [154]

288.4 eV and corresponding to C–C, C–OH, C=O, and COOH groups, respectively [157, 158]. The C 1 s spectrum of raw CF did not contain the peaks of C–OH, C=O, or COOH groups. Figure 3.60c shows the O 1 s spectra of pristine and treated CF. The spectrum of the latter featured a wide peak that was deconvoluted into three signals corresponding to C=O (530.9 eV), C–OH (531.7 eV), and COOH (532.8 eV) groups [159, 160]. This result was consistent with C 1 s data, implying that CF treatment resulted in surface oxidation and afforded a substrate suitable for PB film deposition. After wet chemical deposition, the CF surface was completely coated by cubic crystals with a side length of ~300 nm (Fig. 3.59c). The XRD pattern of the deposited film displayed a strong peak at 26.01°, which was ascribed to the (002) facet of CF (Fig. 3.61a). The remaining peaks were indexed to a face-centered cubic structure with a lattice parameter of a = 10.166 Å. This crystal structure was consistent with that of usual PB analogs, which indicated that PB was successfully deposited on CF. The results of XRD analysis agreed with those provided by FTIR spectroscopy (Fig. 3.61b). The FTIR spectrum of coated CF featured a strong absorption peak around 2080 cm−1 and was ascribed to the C≡N stretching vibration of cyanide groups in the deposited PB. For flexibility evaluation, the composite electrode was bent by ~120° and twisted by ~90°. Throughout the deformation process, the PB layer remained firmly attached to CF (Fig. 3.62). After the external pressure was removed, the distorted electrode reverted to its original shape. The absence of PB powder removal indicated that the coating and substrate deformed in lockstep within the respective elastic deformation ranges. The mass change during repeated bending was negligible for the electrode with chemically linked PB but was substantial for the electrode with adhesively bonded PB (Fig. 3.62d). The effect of water flow was investigated by submerging the electrodes in seawater upon continuous stirring at 300 rpm. After 30 min, the composite electrode lost only 3% of its mass, whereas a mass loss of >60% was observed for the electrode with adhesively bonded PB (Fig. 3.62e). The high flexibility of the composite electrode inspired us to investigate its energy generation performance. Figure 3.63 schematically illustrates the prototype device

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Fig. 3.62 a–c Images of the composite electrode featuring PB grown on CF: a in the as-produced state, b during and immediately after bending, c during and immediately after twisting. Mass changes of electrodes during d the bending test and e continuous stirring in seawater. Reproduced with permission of Ref. [154]

Fig. 3.63 Schematic mechanism of a dual-mode electricity generation system. Reproduced with permission of Ref. [154]

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and the operation mechanism. According to this mechanism, electrons are automatically transmitted from the metal anode to various acceptors, depending on the operation mode. As O2 has a higher reduction potential than PB, electrons are often first received by O2 . Inexhaustible DO equates to limitless energy as long as a sufficient amount of electrons can be supplied. As a result, O2 reduction is a high-energy phase. However, the low concentration of DO in seawater is insufficient for the high uptake of electrons per unit time, i.e., when the output power is significantly increased, the electrons must be directed to PB. Using the redox reaction of the Fe2+ /Fe3+ nodes, which is driven by electron transfer in a restricted potential window, Na+ ions are accommodated in the interstitial sites of the PB lattice throughout this process. The sufficient concentration of alkali metal ions in seawater and their fast insertion can result in high output power. When simulated seawater was supplied to the system, an electric current was created between the magnesium alloy anode and the composite cathode with an average OCV of 1.7 V. This current was sustained for >90,000 s at a current density of 0.025 mA cm−2 (Fig. 3.64a), which suggested that the electrode could be operated in a high-energy mode. Through current density control, the voltage difference between the anode and cathode was varied between 1.7 and 1.5 V. At a current density of 0.025 mA cm−2 , the voltage difference equaled 1.7 V, corresponding to an output power density of 0.0425 mW cm−2 . At a current density of 5 mA cm−2 , the voltage difference equaled 1.5 V, corresponding to an output power density of 7.5 mW cm−2 (Fig. 3.64a). This shift demonstrated that the flexible electrode–based system could flip between two modes depending on the output power density, similarly to the system with a rigid electrode. The electrochemical performance of the above system was evaluated when the electrode was distorted or exposed to water flow. When the electrode was bent by almost 180°, its performance barely changed (Fig. 3.64b). A current with an average OCV of 1.7 V was produced between the magnesium anode and the composite cathode over 90,000 s at a current density of 0.025 mA cm−2 . Similar to planar electrodes, the studied system could switch between two operating modes. Figure 3.64b (right) illustrates how the output voltage could be automatically altered in response to changes in output power density. Magnetic stirring was used to replicate the directional underwater flow found in the ocean. Figure 3.64c demonstrates that the system could be continuously operated in the high-energy mode even at a stirring rate of 250 rpm. At a current density of 0.025 mA cm−2 , the average circuit voltage equaled 1.7 V, corresponding to an output power density of 0.0425 mW cm−2 . Dual-mode switching could also be accomplished under continuous stirring (Fig. 3.64c, right). In summary, PB was grown on modified CF to create a flexible coordination polymer electrode. The modified CF featured hydroxyl and carboxyl surface groups that could engage in strong chemical bonding with the deposited PB film. In view of the close contact between the PB coating and CF, this electrode was very flexible, which allowed for efficient electrochemical power generation even after deformation or under the influence of a water flow. Depending on the output current density, high-energy and high-power modes could be utilized interchangeably.

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Fig. 3.64 a Performance of the system with a planar electrode. The left figure depicts the results for high-energy-mode operation. The right figure illustrates the activity occurring upon mode switching. b Performance of the system with a bent electrode. The left figure depicts the results for high-energymode operation. The right figure illustrates the activity occurring upon mode switching. c System performance upon continuous solution stirring. The left figure depicts the results for high-energymode operation. The right figure illustrates the activity occurring upon mode switching. Reproduced with permission of Ref. [154]

3.2.4.3

Layered Oxides (Na0.5 Co0.5 Mn0.5 O2 )

The catalytic performance of earth-abundant 3d transition metals and their oxides suggests that they can be employed as bifunctional ORR/OER electrocatalysts [129, 161–165]. The different oxidation states of 3d metals in their oxides may enhance OER/ORR activity during electrochemical processes at the electrode surface. Consequently, we analyzed the Na+ intercalation/deintercalation process in conjunction with OER/ORR electrocatalytic processes to improve SWB performance via a dual process utilizing Na-ion electrode materials spanning the integrated voltage profiles.

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Fig. 3.65 a Powder XRD pattern of P2-type layered Na0.5 Co0.5 Mn0.5 O2 (900 °C for 12 h) with Rietveld refinement; b annular bright-field SEM image of the local structure, with faint contrasting dark dots referring to Na atoms; c magnified view and schematic layered structure with Co/MnO6 octahedra and Na trigonal prisms. Reproduced with permission of Ref. [167]

Na+ intercalation/deintercalation on a Co-Mn electrocatalyst would be the primary mechanism responsible for the reduction of the initial voltage polarization and prominent OER/ORR features. The suggested dual process was demonstrated using a P2type layered Na0.5 Co0.5 Mn0.5 O2 oxide that exhibited a propensity for Na+ intercalation/deintercalation and OER/ORR catalytic activity as the primary electrochemical reaction. A mixed hydroxycarbonate method was used to construct the P2-type layered Na0.5 Co0.5 Mn0.5 O2 electrocatalyst according to a reported procedure, and the composition of this catalyst was confirmed by ICP-OES. All signals in the powder XRD pattern of the above catalyst were successfully indexed to the P2-type phase with a hexagonal layered structure. The lattice parameters and unit cell volume were obtained using Rietveld refinement analysis as a = 2.8363 Å, c = 11.3152 Å, and V = 78.83 Å3 (Fig. 3.65a). The local structure of the layered Na0.5 Co0.5 Mn0.5 O2 electrocatalyst was validated through ABF-STEM imaging (Fig. 3.65b), which revealed that Co/MnO6 octahedral units corresponded to the atomic arrangement of Co/Mn with O atoms and the Na trigonal prisms corresponded to Na with O atoms, demonstrating the layered stacking (Fig. 3.65c) [166]. Taking into consideration the suggested dual method, a SWB was designed with P2-type layered Na0.5 Co0.5 Mn0.5 O2 as the cathode electrocatalyst and compared to SWBs with CF and a 20 wt% Pt/C catalyst (Fig. 3.66a). The voltage versus time curves for the CF substrate (red) and the 20 wt% Pt/C catalyst (blue) demonstrated similar OER and ORR characteristics and flat voltage profiles. However, two distinct voltage profiles were observed for the P2-type layered Na0.5 Co0.5 Mn0.5 O2 electrocatalyst, with the slope profile corresponding to Na+ intercalation/deintercalation and the flat profile corresponding to the OER/ORR characteristics of the layered Na0.5 Co0.5 Mn0.5 O2 electrocatalyst (Fig. 3.66a (green)). In half-cell SWBs, the P2type layered electrocatalyst outperformed 20 wt% Pt/C (blue) and CF (red), which was attributed to the increased exposure of active sites in the layered material. The Na+ intercalation/deintercalation reactions and OER/ORR characteristics were explored

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Fig. 3.66 a Intercalation–deintercalation and OER/ORR properties of layered Na0.5 Co0.5 Mn0.5 O2 compared to those of CF and 20 wt% Pt/C. b OER and ORR characteristics with intercalation– deintercalation characteristics proportional to the supplied charge–discharge capacity. Reproduced with permission of Ref. [167]

further (Fig. 3.66b). The supplied discharge capacity was calculated as 30 mAh g−1 , which is consistent with the Na+ intercalation/deintercalation process (slope voltage profile) of Na0.5 Co0.5 Mn0.5 O2 . The slope voltage profile of the Na-ion half-cell in Fig. 3.67a (voltage vs. capacity) was ascribed to the Na+ intercalation/deintercalation process of Na0.5 Co0.5 Mn0.5 O2 . Therefore, this electrocatalyst was explored as a possible SWB cathode capable of operating at high voltages while maintaining a high-rate capability. The corresponding half-cell was discharged at 52 mAh g−1 for 50 cycles between 3.0 and 4.4 V at 0.1 C. Additionally, CV analysis between 2.5 and 4.0 V confirmed the presence of distinct Co3+ /Co4+ redox peaks at 3.31/3.04 V, which were attributed to the Na+ intercalation/deintercalation process of Na0.5 Co0.5 Mn0.5 O2 and were not

Fig. 3.67 a Charge–discharge curves of a Na-ion half-cell with the Na0.5 Co0.5 Mn0.5 O2 electrocatalyst and b CV curves of the half-cell in (a) compared to those of the half-cell with CF. Reproduced with permission of Ref. [167]

158

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observed for CF (Fig. 3.67b). Thus, Na0.5 Co0.5 Mn0.5 O2 improved the redox process involving the Co3+ /Co4+ redox species in SWBs. Subsequently, we probed the efficiency of Na0.5 Co0.5 Mn0.5 O2 as a bifunctional ORR/OER electrocatalyst for use in SWBs. The catalyst ink generated from Na0.5 Co0.5 Mn0.5 O2 was deposited on an RDE, which was then subjected to LSV analysis, and the obtained ORR activity was compared to those of IrO2 and 20 wt% Pt/C as reference catalysts. The above analysis was performed in O2 -saturated seawater (Fig. 3.68b) and O2 -saturated 0.1 M KOH (Fig. 3.68c) at rotation speeds of 400–2225 rpm. At 2225 rpm, the ORR catalytic potentials of Na0.5 Co0.5 Mn0.5 O2 equaled −0.8 V at 1.6 mA cm−2 (seawater) and −0.8 V at 2.9 mA cm−2 (0.1 M KOH). CV analysis (0.4 to −0.7 V) convincingly demonstrated the ORR activity of Na0.5 Co0.5 Mn0.5 O2 in both seawater and 0.1 M KOH (Fig. 3.68d), revealing that a reductive peak due to ORR catalytic activity was observed when these media were saturated with O2 but not when these media were saturated with N2 . The anodic OER catalytic activities of Na0.5 Co0.5 Mn0.5 O2 and IrO2 were determined using LSV

Fig. 3.68 OER and ORR performances of Na0.5 Co0.5 Mn0.5 O2 . a OER anodic polarization curves of Na0.5 Co0.5 Mn0.5 O2 and IrO2 recorded in seawater and 0.1 M KOH. b, c ORR cathodic polarization curves of Na0.5 Co0.5 Mn0.5 O2 recorded in O2 -saturated b seawater and c 0.1 M KOH. d CV curves recorded to probe the ORR activity of Na0.5 Co0.5 Mn0.5 O2. . Reproduced with permission of Ref. [167]

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experiments in seawater and 0.1 M KOH (Fig. 3.68a). Na0.5 Co0.5 Mn0.5 O2 was OERactive at 1.37 V in seawater and 0.74 V in methanol (0.1 M KOH), thus exhibiting dual OER/ORR activity suitable for the fabrication of SWB cathodes. After identifying the Na+ intercalation/deintercalation and OER/ORR properties of Na0.5 Co0.5 Mn0.5 O2 , we examined its performance in a half-cell seawater charge–discharge cycling study. Compared to CF and 20 wt% Pt/C, the above catalyst featured a significantly better voltage versus time performance for the initial five cycles (Fig. 3.69a). The voltage difference of 0.78 V obtained for Na+ intercalation/deintercalation and the OER/ORR activity of Na0.5 Co0.5 Mn0.5 O2 were less than those obtained for CF and were almost identical to those obtained for 20 wt% Pt/C (Fig. 3.69a). The investigation of half-cell SWBs revealed that the one with Na0.5 Co0.5 Mn0.5 O2 exhibited superior cycling stability (50 cycles) compared to those with CF and 20 wt% Pt/C (Fig. 3.69b). The voltage versus time cycling graph (Fig. 3.69c) demonstrated the catalytic properties of Na+ intercalation/deintercalation and OER/ORR in relation to the integrated voltage profile corresponding to the slope and flat voltage regions. Additionally, the half-cell SWB displayed a consistent voltage difference of 0.78 V and an efficiency of 80% (Fig. 3.69d).

Fig. 3.69 Performance of half-cell SWB with Na0.5 Co0.5 Mn0.5 O2 as an electrocatalyst. a Voltage profile and comparison with SWBs containing CF and 20 wt% Pt/C. b Cycling stability and comparison with SWBs containing CF and 20 wt% Pt/C. c Cycling profile of Na0.5 Co0.5 Mn0.5 O2 . d Cycling reliability performance. Reproduced with permission of Ref. [167]

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In a rechargeable SWB, Na0.5 Co0.5 Mn0.5 O2 performed exceptionally well, and the structure of this electrocatalyst was stable within a voltage range of 3.0–4.4 V. The high performance of the tested SWB was validated by the successful intercalation– deintercalation of Na+ ions during electrochemical OER/ORR cycling, as revealed by the related integrated voltage profiles. Compared to those of SWBs with CF and 20 wt% Pt/C, the voltage profile of the half-cell SWB with Na0.5 Co0.5 Mn0.5 O2 featured a steep-slope area corresponding to Na+ ion intercalation/deintercalation and a flat area corresponding to catalytic OER/ORR. CV analysis revealed a discharge capacity of 30 mAh g−1 for Na+ ion intercalation/deintercalation and Co3+ /Co4+ redox peaks at 3.31 V (oxidation) and 3.04 V (reduction). ORR catalytic activity was demonstrated using a rotating disk electrode method. The half-cell SWB with Na0.5 Co0.5 Mn0.5 O2 maintained a voltage difference of 0.78 V and an efficiency of 80% over 50 cycles at 0.1 mA.

3.2.5 Redox Reactions 3.2.5.1

Sodium Hexacyanoferrate (NaHCF)

Solutions of redox couples can be used as SWB catholytes instead of seawater, in which case, however, the battery system must be closed. Figure 3.70 illustrates the redox couples suited for aquatic environments. Among such aqueous catholytes, those containing NaHCF (Na4 Fe(CN)6 ·10H2 O) are most suitable. NaHCF offers the

Fig. 3.70 Redox couples suitable for use in aqueous catholytes for SWBs and the related standard redox potentials. Reproduced with permission of Ref. [169]

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Fig. 3.71 a Schematic illustration of a redox catholyte–containing SWB and the electrochemical reactions occurring at the anode and cathode (1. anode, 2. organic electrolyte, 3. NASICON ceramic electrolyte, 4. aqueous catholyte (NaHCF in water), 5. CF current collector, 6. catholyte tank, and 7. pump). b Schematic illustration of Na+ transport through the ceramic separator (NASICON) from the cathode to the anode and vice versa. c Chemical structures of [Fe(CN)6 ]4− and [Fe(CN)6 ]3− . Reproduced with permission of Ref. [169]

benefits of low toxicity (due to the strong bonding between the CN ligand and the Fe center; the formation constant for Fe(CN)6 4− is 1036 ) and low volatility [168], while NaHCF-containing catholytes are less corrosive (i.e., more user-friendly) than those containing bromine and iodine. Figure 3.71a illustrates the operation principle of SWBs with redox catholytes. During charging, Na+ ions from the catholyte are transported through the ceramic electrolyte (NASICON) and deposited at the Na-metal anode, while [FeII (CN)6 ]4− ions are oxidized to [FeIII (CN)6 ]3− ions, whereas the reverse processes occur during discharge (Fig. 3.71b). The overall charge/discharge reaction is given by Eq. 3.23. 4− 3− Na+ (aq) + [Fe(CN)6 ](aq) ↔ [Fe(CN)6 ](aq) + Na(s)

(3.23)

Both [FeIII (CN)6 ]3− and [FeII (CN)6 ]4− ions exhibit an octahedral structure (Oh geometry with a central Fe atom) (Fig. 3.71c), featuring Fe ions in 3d5 and 3d6 configurations, respectively. Thus, the 3d orbitals of FeII /FeIII in [FeII (CN)6 ]4− /[FeIII (CN)6 ]3− are responsible for the sequence of oxidation/reduction (i.e., electron release/gain) events occurring during charge/discharge. In view of the lack of information regarding the solubility of NaHCF, we tested the solubility of commercially available NaHCF in water at room temperature (25 °C), showing that dissolution was fast and determining the concentration of the saturated solution as 0.68 M (Fig. 3.72a). All aqueous NaHCF solutions had a pH of 8. At this pH, water has an electrochemical stability window of 0.47–0.76 V versus SHE.

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Fig. 3.72 a Photograph of aqueous NaHCF solutions with different concentrations, b CV curves recorded using 0.4 M aqueous NaHCF as the catholyte for the first 10 cycles, c galvanostatic charge–discharge profiles recorded for various NaHCF concentrations (0.1–0.68 M), and d effect of NaHCF concentration in the catholyte on SWB capacity. Reproduced with permission of Ref. [169]

However, in practice, water electrolysis (i.e., the hydrogen and oxygen evolution reactions (OER and HER, respectively) typically does not occur at the theoretical values of 0.47–0.76 V, as the electrical energy input must provide the entire amount of energy required for the OER and the HER. As a result, the theoretical potential window is expanded to 0.87–1.13 V versus SHE, and flow cells employing the NaHCF aqueous catholyte may function over a broad potential range of ~1.84–3.84 V versus Na+ /Na. Although the theoretical redox potential of the [Fe(CN)6 ]4− /[FeIII (CN)6 ]3− couple is 0.36 V versus SHE, the formal redox potential fluctuates with the solution’s concentration/ionic strength and reaches a maximum of ~0.40–0.51 V versus SHE [170– 172]. Thus, a maximum cell voltage of ~3.11–3.22 V versus Na+ /Na may be reached in SWBs with NaHCF-based aqueous catholytes. The electrochemical characteristics of the above SWBs were initially examined in static mode using CV and galvanostatic charge–discharge measurements. CV measurements were performed in the potential range of 2.5–3.8 V using 0.4 M aqueous NaHCF as the catholyte. The resulting curves (Fig. 3.72b) featured welldefined redox peaks due to the [Fe(CN)6 ]4− /[FeIII (CN)6 ]3− redox pair but no OER or HER peaks, which indicated that the catholyte was stable within the potential range of 2.5–3.8 V. Additionally, the CV curves for the first 10 cycles nearly overlapped, and repeated scanning did not significantly affect the peak current or intensity (Fig. 3.72b), which confirmed the reversibility and stability of the [Fe(CN)6 ]4−/ [FeIII (CN)6 ]3− redox pair. Figure 3.72c illustrates first-cycle galvanostatic charge–discharge profiles for aqueous catholytes with several NaHCF concentrations, revealing that the volumetric capacity linearly increased with increasing

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163

Fig. 3.73 Electrochemical performance of SWB with 0.68 M aqueous NaHCF as redox catholyte determined in static mode and characterization of the NASICON and catholyte before and after cycling. a Galvanostatic charge–discharge profiles, b effects of cycling on capacity and Coulombic efficiency, c electrochemical impedance spectra recorded before and after 100 cycles. Reproduced with permission of Ref. [169]

NaHCF concentration and equaled 2.6, 10.4, 15.3, and 17.3 Ah L−1 for 0.1, 0.4, 0.6, and 0.68 M NaHCF, respectively. Thus, capacity could be increased through the use of a more concentrated catholyte. According to Fig. 3.72d, 97% of the theoretical capacity was realized with 0.1 M and 0.4 M NaHCF, while a value of 95% was obtained for 0.6 M and 0.68 M NaHCF. Figure 3.73a illustrates the first 10 galvanostatic charge–discharge curves recorded for 0.68 M aqueous NaHCF as the catholyte, demonstrating no substantial potential deterioration. Additionally, no discernible variations in discharge capacity were observed over 100 cycles, and a 100% Coulombic efficiency was maintained after the first cycle (Fig. 3.73b). The excellent charge–discharge and cycling performance substantiated the viability of utilizing redox catholytes in SWBs. Electrochemical impedance spectra were recorded before and after cycling to determine the cell’s stability and compatibility with the catholyte (0.68 M aqueous NaHCF). Figure 3.73c presents the Nyquist plots recorded before and after 100 cycles. After 100 cycles, the interfacial resistance decreased because of the grain boundary resistance of the NASICON ceramic electrolyte and the interfacial resistance between NASICON and both electrolytes (anolyte/aqueous catholyte). This behavior might be due to the etching of NASICON grain boundaries during cell testing [173]. Figure 3.74a illustrates the effects of cycling on the morphology and appearance of the NASICON ceramic electrolyte. After cycling, the NASICON surface on the anode side remained virtually unchanged in color, whereas the surface on the cathode side turned yellow, possibly because of the adsorption of aqueous NaHCF. SEM images demonstrated that cycling affected the surface roughness of NASICON, possibly by promoting its etching. However, as this surface corrosion happened on a microscale, it was assumed to have no effect on cell performance. Subsequently, both sides of cycled NASICON were investigated using XRD measurements, and the results were compared to those obtained for pristine NASICON (Fig. 3.74b). Notably, cycling did not induce any changes in phase composition or peak position, i.e., NASICON was extremely stable in both aqueous and organic electrolytes. The Na-ion content of the aqueous catholyte after cycling was determined by ion chromatography (Fig. 3.74c) and slightly decreased (by 0.056 M) after 100 cycles, possibly because of the initial

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Fig. 3.74 a Photographs and SEM images of NASICON before and after cycling (1. pristine, 2. aqueous cathode side after cycling, and 3. non-aqueous anolyte side after cycling), b XRD patterns of NASICON before/after cycling and XRD pattern of the NASICON reference, c effects of cycling on Na-ion concentration in the catholyte. Reproduced with permission of Ref. [169]

irreversibility of cycling. However, the aqueous catholyte was stable over subsequent cycles. Finally, as a proof of concept, we evaluated the cycling performance of a SWB with concentrated aqueous NaHCF (0.68 M) as the catholyte (Fig. 3.75) at currents of 1.5, 3, 4, and 5 mA (0.75, 1.5, 2, and 2.5 mA cm−2 , respectively), as seen in Fig. 3.75a. The investigated flow cell had a volumetric capacity of 17.7 Ah L−1 , which accounted for 97.14% of the theoretical capacity (18.22 Ah L−1 ), and exhibited a maximum

Fig. 3.75 Electrochemical performance of a SWB containing 0.68 M aqueous NaHCF as the redox catholyte and operated in a flow-through mode. a Galvanostatic charge–discharge profiles at different currents, b polarization curves, c representative charge–discharge curves, d plots of capacity, CE, VE, and EE versus cycle number, e photograph of assembled flow-cell setup, and f, g charged flow-cell setup powering multicolored glowing LEDs. Reproduced with permission of Ref. [169]

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165

average discharge potential of 3.06 V when operated at 1.5 mA. However, when the current was increased to 5 mA, the discharge potential decreased to 2.74 V. At 1.5 and 5 mA (0.75 and 2.5 mA cm−2 , respectively), the cell exhibited energy densities of 54.16 and 33.92 Wh L−1 , respectively. The former energy density accounted for 96.71% of the theoretical energy density (56 Wh L−1 ). The related polarization curves are shown in Fig. 3.75b. The maximum power density was determined as 15.3 mW (7.65 mW cm−2 ), corresponding to a discharge potential of 2.05 V. At 7 mA (3.5 mA cm−2 ), the flow cell rapidly approached its maximum power output. The potential drop observed at high currents of >4 mA (2 mA cm−2 ) was ascribed to the sluggish kinetics of redox reactions and the limited Na-ion mobility in NASICON. Thus, battery performance may be enhanced by increasing the ionic conductivity of NASICON. Figure 3.75c shows representative galvanostatic charge–discharge curves of the flow cell for the first seven cycles, revealing that the cell could be charged and discharged within the indicated window. Notably, the yellow NaHCF solution became virtually colorless during charging and then turned yellow again during discharging (insets in Fig. 3.75c). Figure 3.75d presents the effects of cycling on the flow cell’s capacity, Coulombic efficiency, VE, and EE. After 50 cycles, CE, VE, EE, and capacity retention equaled 98%, 79%, 79%, and 99%, respectively. Figure 3.75e depicts the fully completed flow-cell configuration, which included a pump, a catholyte reservoir, a flow controller, and a cell. The flow cell could power 10 multicolored luminous light emitting diodes (LEDs) when charged at 3 mA (Fig. 3.75f and g). In summary, a proof-of-concept SWB was developed and manufactured using a ceramic electrolyte (NASICON) separator and aqueous NaHCF as the redox catholyte in the absence of a supporting electrolyte. The catholyte supplied Na+ ions and a redox couple for anodic and cathodic processes, respectively. When 0.68 M NaHCF was used, the SWB featured a potential of 3.08 V and an impressive capacity and energy density of 17.7 Ah L−1 and 54.16 Wh L−1 , respectively (calculated based on catholyte volume). The energy density can presumably be increased further by (i) operating the battery system at 100 °C to increase the concentration of the saturated NaHCF solution to 1.5 M (which corresponds to a theoretical energy density of 123 Wh L−1 ), (ii) using an aqueous catholyte with multiple redox couples, and (iii) using a new high-concentration catholyte. However, at the current level of manufacturing, it is most important to ensure scalability. Future works should focus on optimizing flow cell engineering, the design of the cathode current collector, and flow rate as well as on increasing the conductivity of the NASICON ceramic separator and substituting the sodium metal anode for a safer alternative (e.g., hard carbon dispersion or semi-liquid sodium biphenyl) to achieve significant performance improvement and increased safety. As a result, it is anticipated that this technology will set the foundation for the next-generation high-energy–density energy storage systems.

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3 Secondary Seawater Batteries

Dual Electrolytes (NiHCF/NaHCF)

The charge storage capacity of SWBs can be increased by combining electrolyte redox chemistry with a Na-ion intercalation/deintercalation cathode through the use of dual electrolytes. Whereas conventional electrolytes do not contribute to the total charge storage capacity of batteries, redox-active electrolytes do so via charge transfer processes at the electrode–electrolyte interface (Fig. 3.76). Herein, aqueous NaHCF was used as the redox-active electrolyte, and NiHCF was used as the Naion intercalation/deintercalation cathode. The NaHCF electrolyte is less corrosive than iodine- and bromine-based redox electrolytes and provides an electrochemically active [FeII (CN)6 ]4− /[FeIII (CN)6 ]3− couple for the cathode reaction as well as an adequate amount of Na+ ions for the anode reaction [169]. Additionally, NiHCF is a potential Na-ion intercalation/deintercalation cathode material because of its high rate capability and cycle stability in aqueous electrolytes [174, 175]. The electrochemical properties of NaHCF (redox-active electrolyte) and Na2 SO4 (traditional electrolyte) were first determined using a galvanostatic charge–discharge technique. The cut-off voltages were chosen to lie between 2.8 and 3.8 V versus Na/Na+ to account for the cathode’s water oxidation and oxygen reduction restrictions. The charge–discharge curves obtained for Na2 SO4 and NaHCF during the first cycle are shown in Fig. 3.77a. The average discharge voltage and voltage difference for Na2 SO4 were determined as 3.27 and 0.09 V, respectively, whereas the respective values for NaHCF equaled 3.24 and 0.07 V. The initial discharge capabilities of Na2 SO4 and NaHCF were 47.8 and 107.11 mAh g−1 , respectively (the calculated capacity was based on the mass of the cathode active material unless otherwise specified). Thus, the average discharge voltage and the voltage difference were not significantly different for the two electrolytes, whereas the capacity obtained for NaHCF was almost two times that obtained for Na2 SO4 . This finding implies that

Fig. 3.76 Electrochemical processes occurring in dual-electrolyte SWBs with aqueous redoxactive electrolytes. Reproduced with permission of Ref. [176]

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167

Fig. 3.77 a Charge–discharge profiles of NaHCF and Na2 SO4 electrolytes and b illustration of the cathodic reaction at the cathode (NiHCF@CF) with the NaHCF electrolyte. Reproduced with permission of Ref. [176]

the improved charge storage performance of the dual-electrolyte SWB was due to the redox reactions of the [FeII (CN)6 ]4− /[FeIII (CN)6 ]3− couple provided by NaHCF and the Na-ion intercalation/deintercalation reactions at NiHCF. The Na2 SO4 electrolyte acted exclusively as a channel for the transfer of Na+ ions to and from NiHCF. Thus, when this electrolyte was used, charge storage was restricted by the cathode. However, the NaHCF electrolyte acted as a channel for the transport of Na+ ions during intercalation/deintercalation at NiHCF and provided electrochemically active species ([FeII (CN)6 ]4− /[FeIII (CN)6 ]3− ) for an extra charge transfer reaction at the cathode (Fig. 3.77b). Thus, the combination of the NaHCF electrolyte with NiHCF allowed for a considerable increase in charge storage capacity. As seen in Fig. 3.78a, both Na2 SO4 and NaHCF electrolytes displayed a steady capacity with a Coulombic efficiency of around 99% after 100 cycles. The high reversible discharge capacity and outstanding cycling performance of the NaHCF electrolyte indicated that the [FeII (CN)6 ]4− /[FeIII (CN)6 ]3− redox reactions and the rapid Na+ ion intercalation/deintercalation reactions at NiHCF did not interact with each other during charge–discharge cycling. In real-time applications, extended cycling performance is critically needed. Consequently, the rate performance of the NaHCF electrolyte was evaluated by cycling at various current densities, and capacities of 107.11, 104.49, 101.01, 96.41, 90.65, and 26.61 mAh g−1 were attained at current densities of 0.05, 0.1, 0.2, 0.4, 0.6, and 1 mA cm−2 , respectively (Fig. 3.78b).

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Fig. 3.78 a Effects of cycling on the capacity retention and Coulombic efficiency of NaHCF and Na2 SO4 electrolytes, b rate performance capability of the NaHCF electrolyte. Reproduced with permission of Ref. [176]

Additionally, the capacities achieved at 0.1, 0.4, and 1 mA cm−2 were 97.55, 90.01, and 24.84% of that attained at 0.05 mA cm−2 , respectively, and more than 99% of the capacity was regained when the current density was lowered from 1 to 0.05 mA cm−2 . In Fig. 3.79a, traces (i), (ii), and (iii) represent the CV curves recorded for Na2 SO4 , Na-HCF, and NaHCF without NiHCF, respectively, and well-defined redox peaks within the 2.8–3.8 V region are observed in all cases. Peaks (i), (ii), and (iii) correspond to Na-ion intercalation/deintercalation at NiHCF, [FeII (CN)6 ]4− /[FeIII (CN)6 ]3− redox reactions in combination with Na-ion intercalation/deintercalation reactions at NiHCF, and [FeII (CN)6 ]4− /[FeIII (CN)6 ]3− redox reactions [169]. It should be emphasized that in order to preserve the charge neutrality of NiHCF, Na-ion intercalation/deintercalation occurred via a redox mechanism involving the FeII /FeIII pair [143, 145]. However, a significant area was observed under the current–voltage curve for the trace of NaHCF electrolyte + NiHCF (curve (ii)), in line with the increased charge storage capacity of this system. Additionally, the positive-scan oxidation onset potential of curve (ii) was comparable to that of curve (iii), while the negative-scan reduction onset potential of curve (ii) was similar to that of curve (i). These findings indicated that in the NaHCF-containing system, oxidation started with the oxidation of [FeII (CN)6 ]4− during the positive scan (red

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169

Fig. 3.79 a CV curves of systems containing Na2 SO4 (i), NaHCF (ii), and NaHCF without NiHCF (iii), b charge–discharge profiles of NaHCF including the step-by-step charge and discharge routes, and c kinetic model of the cathode in charged and discharged states. Reproduced with permission of Ref. [176]

arrow in Fig. 3.79a), while reduction started with the Na+ intercalation process during the negative scan (green arrow in Fig. 3.79a). Figure 3.79b and c explain the electrochemical behavior and processes occurring at the cathode of the cell with the NaHCF electrolyte. When the electrode–electrolyte interface is charged, [FeII (CN)6 ]4− ions migrate toward the cathode current collector and are rapidly oxidized to [FeIII (CN)6 ]3− . With increasing charge voltage, NiHCF becomes positively charged, and Na+ ions are ejected from the crystal lattice, while FeII is oxidized to FeIII in NiHCF. In view of their increased mobility in the liquid phase, free Na+ ions rapidly pass through the electrolyte (Fig. 3.79b and c(i)). Upon discharge, NiHCF becomes negatively charged, which allows the reinsertion of Na+ ions into the crystal structure, while FeIII is reduced to FeII in NiHCF. [FeIII (CN)6 ]3− is converted to [FeII (CN)6 ]4− at the electrode–electrolyte interface at a more negative potential (Fig. 3.79c(ii)). To address the practical applications of NaHCF, we investigated the electrochemical performance of a system where a semi-liquid NaC12 H10 (sodium biphenyl

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Fig. 3.80 a Image of a NaHCF-based battery with a semi-liquid NaC12 H10 anode, b reaction between Na metal and water and that between semi-liquid NaC12 H10 and water, c charge–discharge profiles and d cycling capacity retention and efficiency of a battery containing NaHCF with semiliquid NaC12 H10 , e current density–capacity and f polarization curves of batteries with Na metal and semi-liquid NaC12 H10 anodes. Reproduced with permission of Ref. [176]

complex) anode was used in place of the Na metal anode (Fig. 3.80a). The electrochemical reaction of semi-liquid NaC12 H10 is given by NaC12 H10 ↔ [C12 H10 ]• + Na+ + e− . NaC12 H10 was produced by dissolving Na metal in a solution of C12 H10 in TEGDME. Unlike classical Na-ion insertion/deinsertion Na anodes, the semi-liquid NaC12 H10 anode can operate as an electrolyte, is safer than Na metal, and can survive direct contact with aqueous media. Figure 3.80b demonstrates that when water was brought in contact with semi-liquid NaC12 H10 , no fire occurred, in contrast to the case of Na metal. The reaction of NaC12 H10 with water is given by 2NaC12 H10 + 2H2 O → 2C12 H10 + 2NaOH + H2 (Fig. 3.80b) [177]. Na-metal-free NaHCF featured an average discharge capacity of 104 mAh g−1 and an average discharge voltage of 3.11 V (Fig. 3.80c; the calculated capacity was based on the mass of the cathode active material unless otherwise specified). This capacity was very close to that of sodium metal (107.11 mAh g−1 ; Fig. 3.80c). However, the voltage of 3.11 V was somewhat lower than that of the Na-metal NaHCF (3.24 V), possibly because of the lower redox potential of NaC12 H10 . Figure 3.80d shows that even at a high current density of 0.4 mA cm−2 , Na metal-free NaHCF retained about 92.17% of its capacity after 500 cycles. Additionally, capacities of 102.28, 99.18, 96.22, 93.37, 88.11, 81.25, and 58.29 mAh g−1 were obtained at current densities of 0.1, 0.2, 0.4, 0.6, 1, 1.25, and 1.5 mA cm−2 , respectively. The rate performance of NaHCF was enhanced by the semi-liquid NaC12 H10 anode. Moreover, the use of this anode increased power from ~12.5 to ~13.8 mW (which translates into a power density increase from ~6.25 to ~6.9 mW cm−2 ) (Fig. 3.80f). This power output exceeded that of previously reported dual-electrolyte Li batteries (4 mW cm−2 (Li/S battery) [178] and 5.9 mW cm−2 (Li/[Fe(CN)6 ]3−

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171

battery [179]). The increased rate capability and power performance were attributed to a reduction in resistance due to the usage of the semi-liquid NaC12 H10 anode.

3.3 Anodes 3.3.1 Anode Parts of SWBs: Roles and Requirements The anode is a critical compartment for determining the energy density of rechargeable SWBs, because it stores the reduced sodium ions during the charging process. Because the seawater is used in an open-structured cathode, an infinite supply of Na+ ions is available to the anode during charging. This vast Na+ supply enables the use of high-capacity and nanostructured materials such as alloying and conversion reaction-based materials, which typically lose Na+ ions irreversibly during repeated charge–discharge cycles in a conventional closed battery system. In addition, the liquid electrolyte of the anode part plays a critical role in the cyclability of SWBs: the irreversible loss of Na+ at the anode is caused by the formation of an SEI layer on the anode’s surface through side reactions with the electrolyte. Furthermore, because the NASICON solid electrolyte separates the anode and cathode electrically and physically, seawater cannot affect the anode material or liquid electrolyte of the anode components [180] (Fig. 3.81). The anode (Na-metal) half-cell reaction in a conventional liquid electrolyte is : 4Na+ + 4e− ↔ 4 Na,

E = −2.71V (vs. SHE).

Fig. 3.81 The anode part of a seawater battery. Reproduced with permission from Ref. [181]

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3 Secondary Seawater Batteries

Anode-Material Requirements

In a rechargeable SWB, an unlimited number of sodium cations can be obtained from seawater. Thus, the anode determines the battery’s performance in terms of energy density. A variety of anode materials have been investigated in sodium-based batteries; most are also suitable for seawater batteries. The anode of an SWB, like that of conventional sodium-ion battery, must have a redox potential near ENa+ /Na , high reversible capacity, stable structural/electrochemical characteristics, and high electrical conductivity. Figure 3.82 depicts the redox potentials of representative anode candidates against their specific capacities. Note that anode materials (excluding Na metal) may be divided into three categories according to their reaction mechanism during the sodiation/desodiation process: insertion reaction materials, alloying reaction materials, and conversion reaction materials [182]. Because Na metal is more reactive than Li metal, its reversible deposition/stripping performance is more important for operating safety [180]. This necessitates a method for generating a stable SEI layer on the surface, as well as assuring reversible behavior via anode components such as the anode current collector and non-aqueous electrolyte, and improving charge/discharge procedures during cell operation [6].

Fig. 3.82 Candidate active anode materials for seawater batteries. Reproduced with permission from Ref. [6]

3.3 Anodes

3.3.1.2

173

Requirements for Liquid Electrolytes (Anolytes)

Non-aqueous liquid electrolytes are among the most critical factors impacting cycle performance. They function as the ionic transport channels between solid electrolytes and anodes throughout the charging and discharging processes and guarantee the stability of the electrochemical performance of the anode. Several comparative studies of how organic electrolytes enhance the anode performance of full-cell SWBs have been conducted. To avoid internal short circuits, the electrical conductivity of liquid electrolytes in classic sodium-ion batteries should typically be minimized. The sodium superionic conductor (NASICON) ceramic membrane used in SWBs prohibits electron transport between the anode and cathode compartments. As a result, SWB electrolytes have different requirements from non-aqueous electrolytes, including higher ionic conductivity, adequate electrochemical stability at low voltage ranges, low cost, and low toxicity[6]. Also, liquid electrolytes can have high electronic conductivity (Tables 3.3, 3.4 and 3.5). Table 3.3 Non-aqueous solvents. Reproduced with permission from Ref. [183] Solvent

Structure

MP (°C)

BP (°C)

Viscosity (cP) 25 °C

Dielectric constant 25 °C

Acceptor (donor) number

Ethylene carbonate (EC)

36.4

248

1.9 (40 °C)

89.78

(16.4)

Propylene carbonate (PC)

-48.8

242

2.53

64.92

18.3 (15.1)

Dimethyl carbonate (DMC)

4.6

91

0.59

3.107

Diethyl carbonate (DEC)

−74.3

126

0.75

2.805

Ethylmethyl carbonate (EMC)

−53

110

0.65

2.958

Dimethoxyethane (DME)

−58

84

0.46

7.18

10.9 (18.6)

DEGDME

−64

162

1.06

7.4

9.9 (19.2)

TEGDME

−46

216

3.39

7.53

10.5 (14)

(16.0)

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Table 3.4 Na salts for non-aqueous liquid electrolytes. Reproduced with permission from Ref. [183] MW (g mol−1 )

MP (°C)

Ionic conductivity (mS cm−1 )

NaClO4

122.4

468

6.4

NaBF4

109.8

384

NaPF6

167.9

300

NaCF3SO3

172.1

248

NaTFSI

303.1

257

NaFSI

203.3

117

Salt

Anion structure

7.98

6.2

Table 3.5 Polyaromatic hydrocarbons for anolytes. Reproduced with permission from Ref. [184] Poly-aromatic hydrocarbons

Potential (V vs. Na/Na+ )

LUMO (eV)

Biphenyl

0.095 (Bp/Bp·− )

−0.68

(Nap/Nap·− )

−0.97

Naphthalene

0.154

Phenanthrene

0.224 (Phe/Phe·− )

Anthracene

0.725 (Ant/Ant·− ) 0.143 (Ant·− /Ant·2− )

−1.00

(Tet/Tet·− ),

Tetracene

1.118 0.508 0.123 (Tet·2− / Tet·3− )

Pyrene

0.621 (Pyr/Pyr·− )

Perylene

1.034

(Per/Per·− ),

(Tet·− /

Tet·2− )

(Per·− /

Per·2− )

−1.62 −2.09 −1.45

0.418

−1.84

3.3.2 Sodium Metal Sodium-metal anodes are promising candidates for the next-generation rechargeable batteries that will replace LIBs. Sodium compounds are much more abundant than

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lithium in the earth: sodium is the sixth most abundant element in the earth’s crust (2.6~3.0%) and is commonly found in sodium chloride in the oceans. Sodium metal is extremely reactive, so pure elemental sodium is rare in nature; it is produced artificially by electrolytic methods such as the Deville and Down processes. Na-metal anodes, like Li-metal ones, have several advantages. Compared with other anode-material candidates, Na metal possesses a high theoretical capacity (1165 mAh g−1 ) and low electrochemical potential (−2.71 V vs. SHE). These electrochemical properties are responsible for the high theoretical energy densities of many sodium-metal battery systems such as Na–S (~1672 mAh g−1 , 1230 Wh kg−1 ) and Na–O2 (1105 Wh kg−1 ) [185]. However, Na-metal anodes also face challenges during the charge/discharge process. The first of these is Na-dendrite formation and short circuits. In the charge process, Na+ ions come into contact with electrons on the anode surface and electrodeposition of Na metal occurs; this electrodeposition is affected by the uneven distribution of current and irregular structure of anode surfaces. Another critical factor for dendrites is the formation of solid electrolyte interphase (SEI) between the liquid electrolyte and Na metal anode. This unstable SEI layer leads to non-uniform Na+ ion flux, resulting in Na dendrite growth. A second challenge is posed by low coulombic efficiency (CE) and poor cycling performance. During SEI layer formation, the liquid electrolytes decompose; the consumption of electrolyte and electrons leads to irreversible capacity and low CE. In addition, the combination of dendrite formation and unstable SEI formation leads to the accumulation of “dead” Na that is electronically insulated so that cannot participate in the charge/discharge reaction [185]. The SWB anode compartment, like that of the sodium-metal battery, consists of anode material and liquid electrolyte. Therefore, the sodium-metal anode, which has been studied extensively because of its outstanding theoretical capacity and high energy density, provides insight into the fundamental mechanisms of SWBs. However, sodium-metal battery systems also face challenges such as dendrite formation and poor electrochemical cycle performance, so additional study is needed to improve anode performance. An especially important problem is dendrite formation at the anode, which can eventually lead to cell failure despite the existence of a solid electrolyte that suppresses the short circuit. Furthermore, the high reactivity of Na metal with ambient air and moisture makes it difficult to manufacture sodium foil, whereas lithium foil is available commercially.

3.3.2.1

Sodium Metal on a Graphene-Covered-Copper Current Collector

Several approaches to Li-metal passivation have been developed to prevent adverse reactions from occurring at the Na-metal anode. Such methods include controlling the current density of the current collector by using 3D conducting structures and crumpled graphene balls to extend the Sand’s time; increasing the surface tension for mechanical suppression; using nitrogen atoms for Li formation; and adjusting the

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Fig. 3.83 Comparison of Na-plating processes between pristine and graphene-coated Cu current collectors: a, b growth of dendrites; SEM images of c pure copper foil, d copper foil coated with e graphene; f diminished dendrite growth on graphene-coated copper. Reproduced with permission from Ref. [186]

concentration of the liquid electrolyte. At present, however, there are few methods to control the collector’s interface [186–188]. Time-lapse images of sodium dendritic formation on a pure-copper current collector are shown in Fig. 3.83a. The cell design comprised a bottom Cu current collector and a top Na-metal electrode in a transparent quartz tube filled with a liquid electrolyte; the galvanostatic current density was 3.25 mA cm−2 . In the first 50 min of Na plating on the Cu foil, the plated Na metal remained relatively unaffected, and reduced dendrite growth was observed. Sharp dendrites started to form at 55 min; they selectively grew toward the counter-electrode on the Na surface (Fig. 3.83a, b). Na-dendrite growth severely impacts cycle life; therefore, its alleviation is crucial for the development of efficient Na-metal-based batteries. Numerous factors, including ionic conductivity, current density, and transport number of cations, must be considered in cell-system design in order to properly suppress alkali-metal dendrite formation during the plating process [189–194]. Applying a low current density is vital for delaying the onset of dendritic formation (i.e., Sand’s time) and raising the critical nucleation radius. Applying a high overpotential during metal plating at a high current density effectively decreases the critical radius of nucleation. When many nucleation sites form on the current collector surface with a high nucleation rate, the plated Na metal grain size shrinks, resulting in a rough and inhomogeneous surface.

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Because the structural and chemical composition of the current collector surface have a significant effect on the nucleation rate, it was decided to establish a chemically and physically homogeneous surface for the working electrode in order to avoid dendrite development. On the other hand, conventional current collectors often possess varying heights and surface characteristics that may influence the nucleation locations on the surface throughout Na-ion deposition. SEM images of a pure Cu foil (Fig. 3.83c) indicate a corrugated and heterogeneous surface, which may lower the energy barrier (G* = 16πγ 3 V m 2 /3F 2 |η|2 ) and facilitate the development of Nanucleation sites. Although the critical radius in heterogeneous nucleation is shapefactor independent, the lowered energy barriers significantly improve the nucleation rate, leading to a lower grain size in the deposited Na metal. Using a graphene monolayer offers an efficient strategy for achieving atomic-scale surface homogeneity on the current collector during the Na metal plating process. The uniform interface that is formed between the graphene-coated current collector and the electrolyte controls Na-metal formation on the surfaces, resulting in a far smoother Na surface on the metal than on a pure-Cu current collector. Figure 3.83d shows Cu foil coated with mono-layer graphene following a normal chemical vapor deposition (CVD) process that passivated the surface of the chemically heterogeneous Cu foil. This simple yet effective procedure not only prevented chemical degradation, but also preserved the form of the copper-foil current collector’s surface. By lessening the instability of the corrugated copper surface, the chemically homogeneous graphene layer successfully minimized the number of nuclei formed on the copper foil. As demonstrated by Fig. 3.83f, the atomic-scale homogeneous graphene surface on the Cu foil effectively retarded the commencement of Na-dendrite formation on the graphene-coated current collector. It is worth noting that using a Cu current collector not only improves CE, but also helps growth of graphene directly on the surface. The first stage of Na plating was investigated to further elucidate the graphene coating’s effect on the Cu foil. Figure 3.84a–h are SEM images of the island-like

Fig. 3.84 SEM images of island-like sodium-metal plating on (top row) pure and (bottom row) graphene-coated Cu current collectors. Reproduced with permission from Ref. [186]

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sodium-metal plating morphologies on both current collectors under various current densities. The mean island size fell from 150 to 20 μm as the current density increased from 0.032 to 0.65 mA cm−2 , owing to the proportional correlation between the nucleation overpotential and the current density. Owing to its independence from the current density, the graphene-coated Cu current collector contained fewer Na islands and a smoother surface than the pristine Cu current collector. This impressive result was attributed to the homogeneity of the graphene coating, which mitigated nucleus formation on the collector due to the decreased nucleation potential and greater width of the Na metal islands. Additionally, the surface of the graphene coating was covered with Na metals because the lower number of defects on the graphene coating increased the strength of its adhesion to the Cu current collector [195–197]. Galvanostatic plating and stripping experiments can further characterize the graphene layer’s influence on the Cu foil (Fig. 3.85a, b). Compared to pure copper foil, the graphene-coated current collector shows a 40-mV-lower nucleation overpotential. The 18-mV plateau overpotential in the graphene-coated Cu current collector illustrates the important function of the homogeneous graphene coating in lowering not only the nucleation overpotential at the beginning of the plating process, but also the cell potential during Na plating/stripping (Fig. 3.85b). Moreover, even after 200 cycles (150 h), the graphene-coated Cu current collector retained its CE, whereas that of the pristine Cu current collector declined after only a few cycles (Fig. 3.85c). The impedance was measured to assess the graphene-coated current collector’s increased performance (Fig. 3.85d). The radius of the semicircles corresponding to the graphene-coated Cu collector was nearly 50% smaller than that for the pristine Cu foil at the 5th, 10th, and 15th cycles of the plot, showing that the

Fig. 3.85 Electrochemical characterization of Na-metal plating on pristine and graphene-covered Cu current collectors. Reproduced with permission from Ref. [186]

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179

Fig. 3.86 a Schematic illustration of the SWB architecture and SEM images of plated Na metal on b graphene covered Cu and c pristine Cu current collectors. d galvanostatic charge/discharge plot and e discharge plot of SWB cell. Reproduced with permission from Ref. [186]

homogeneous interface of the graphene layer significantly lowered resistance during the electro-crystallization process. The top-view SEM image (Fig. 3.86b) of deposited sodium metals on a graphenecoated copper foil reveals the grain size for the plated sodium metal species: approximately 150–200 μm. As mentioned earlier, owing to the more homogeneous graphene-coated copper foil surface, the density of nuclei and the surface roughness were both less than those of the pure Cu collector. Prior to measuring the SWB cyclability, Na metal was deposited on the anode current collector by charging with 2 mAh cm−2 , and discharging/charging tests were performed for 5 h at a fixed current of 0.075 mA cm−2 . Figure 3.86d illustrates the degrading galvanostatic plots of the graphene-coated and uncoated Cu current collectors. With a cycling profile that was relatively twice as long, the graphenecoated current collector demonstrated a more stable voltage plateau. The improvement in SWB performance due to controlling the current collector surface is shown by the light intensity difference and sustained lighting time for 49 LED lights connected in an Arduino breadboard (Fig. 3.86e). The luminance of the LED lights for the graphene-coated Cu current collector was recognizable for over 5 h, while the pristine Cu current collector exhibited shorter lifetime (approximately 2.5 h). Thus, employing a surface-controlled current collector with a graphene coating

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3 Secondary Seawater Batteries

offers a realistic solution to the problem of Na metal dendrite formation in SWBs. The homogeneous surface of the graphene-coated Cu foil not only decreased nucleation, but also increased the overall electrochemical performance. Performance was consistent, and the cell lifetime was increased. Deactivating the current-collecting surface in Na metal-based batteries is a practical and trust-worthy technology for SWBs.

3.3.2.2

Sodium-Metal–Carbon-Fiber Composites

Although many attempts have been made to improve the performance of Na metal anodes, the processes used are difficult to sustain and cannot be practically employed in manufacturing. Carbon materials are attractive for battery applications because of their low density, abundance in nature, diverse functions, and wide range of production. Through the modification of their physical shapes and chemical interactions, several carbon species, including hollow carbon spheres, carbon nanofiber films, layered reduced graphene oxide (rGO), and carbonized wood, have shown promising electrochemical performance [198–202]. Commercial carbon fiber is lightweight, possesses excellent mechanical properties and electrical conductivity, and demonstrates thermal stability for a variety of applications. It has a rough, checkered three-dimensional structure with a large surface area, which is conducive to alkali-metal deposition and aids in the control of the local current density. Due to the low affinity of liquid Li and Na metals for carbon, pure fiber is ineffective as a scaffold, demonstrating poor permeability. The scaffold’s affinity for liquid metals is important in metal-infusion processes, because it impacts the adhesion between the alkali-metal species and the fiber surface. The carbon fiber must therefore be modified to boost its attraction toward liquid Li or Na species, in order to ensure their capture [203, 204] (Fig. 3.87). The incorporation of nanocrevasses into the carbon fiber improves its permeability, resulting in enhanced alkali-metal-infusion performance. Heat-treated carbon fiber has superior mechanical and chemical properties, as well as a larger active surface area than Na foil, which conjointly enhance its effectiveness as an alkali metal host scaffold. A simple and rapid synthesis technique for the large-scale production of alkalimetal–carbon composites by soaking carbon fiber cloth in alkali metal liquids has been reported. A carbon scaffold and a simple melt infusion method were used to fabricate the nanocomposite electrodes. Most commercial carbon cloths include a thin polymer coating on the surface that helps maintain the fiber’s shape; because the coating agents used in the experiment exhibited minimal adhesive strength with liquid alkali metals, and thus inhibited the direct interaction between the carbon fiber and the liquid alkali metals during the infusion process, the carbon fibers were heat-treated for 4 h at 500 °C in an ambient atmosphere to eliminate the interfacial polymers. After heat treatment, the coating chemicals were completely evaporated, and the carbon cloth exhibited a higher affinity to liquid metals. As a result, in a

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181

Fig. 3.87 Schematic diagram showing the fabrication of Na–C composites. Reproduced with permission from Ref. [198]

matter of seconds, liquid alkali metals could permeate the carbon fiber across the entire surface. The technique described above is compatible with a variety of polyacrylonitrile (PAN)-based scaffold materials, including carbon filters, carbon felt, and carbon paper. The final composite form and size can be readily modified by selecting or modifying the carbon scaffold. To facilitate large-scale manufacturing, an automated roller machine for alkali metal-infiltrated carbon fiber was designed. This system allows for mass composite fabrication and could pave the way for commercialization in the future. The pure carbon cloth displayed a checkered pattern in SEM observations; its structure comprised many overlapping carbon fibers (Fig. 3.88a). Images of the carbon-fiber cloth before and after heat treatment are shown in Fig. 3.88b, c. The BET data for the pristine and heat-treated carbon fibers are shown in Fig. 3.88d, e. The surface area of the pure carbon cloth was approximately 3.8 m2 g−1 (Fig. 3.88d), which corresponded well with that of the flat surface shown in Fig. 3.88b. By contrast, the surface area of the heat-treated carbon fiber was 454.5 m2 g−1 (Fig. 3.88e). This indicates that the introduction of nanocrevasses significantly increased the surface area of the carbon cloth, which exhibited the roughness observed in Fig. 3.88c. This phenomenon is due to the pore-size dispersion after heat treatment: multiple micropores (2 nm) and mesopores (3.53 nm) with pore volumes two orders of magnitude larger than those of the pure carbon fiber were produced. These crevasses enhanced the surface area of the carbon fibers, enabling them to function as hosts for alkali

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Fig. 3.88 Differences in physical/chemical properties of carbon cloth a, b, d, f ,h before and c, e, g, i after heat treatment. j Infusion of metal droplets into nanocrevasses. Reproduced with permission from Ref. [198]

metals such as rGO. The nanocrevasses were formed when the carbon fibers reacted with O2 during heat treatment, resulting in the formation of CO and CO2 . This also accounts for the difference in permeability between the pristine and heat-treated carbon cloths [205]. The surface chemical compositions of the pristine and heat-treated carbon fiber samples were evaluated using X-ray photoelectron spectroscopy (XPS) (Fig. 3.88f, g). The C–O–C epoxy peak disappeared in the spectra of the heat-treated car-bon cloth sample, whereas in that of the pristine carbon cloth sample (Fig. 3.88f), C–O–C peaks arose as expected (Fig. 3.88g), with a minor shift due to the charging effect. Although the C–OH peak appeared to correspond to the leftovers of the stabilization process, it could also have been newly produced during heat treatment by reaction with air. The XPS data show that heat treatment created a new area with oxygencontaining functional groups on the samples, implying that the composites strongly interacted with rGO; this is similar to the behavior of alkali metals. In conclusion, heat treatment increased the affinity of the carbon fibers to alkali metals by causing physical deformations and chemical alterations [206–209].

3.3 Anodes

183

Molten alkali metals were placed dropwise on the untreated and heat-treated textile samples to assess the effect of heat treatment on carbon wettability (Fig. 3.88h, i). Metal droplets tended to condense on the untreated carbon cloth. On the other hand, the metal droplets exhibited a low contact angle with the heat-treated cloth, indicating a strong affinity. This difference is attributable to the increased capillary forces stemming from the nanoscale crevasses on the heat-treated fibers (Fig. 3.88j), which caused infusion of the molten metal droplets into the nanocrevasses. To demonstrate the universality of this approach across alkali metal anodes, the electrochemical performances of a Na–carbon composite and a bare Na foil symmetric cell were compared. In an ether-based electrolyte solution, the Na–carbon composite cell was more stable than the bare Na foil cell, as illustrated in Fig. 3.89a. It is well known that ether-based electrolytes create flexible SEI layers able to survive volume fluctuations caused by non-uniform plating. Here, despite using an etherbased electrolyte, the bare Na foil cell exhibited poor electrochemical performance, with significant variations indicating an internal soft short circuit. The composite

Fig. 3.89 Electrochemical characterization of sodium-metal–carbon composites: a, b potential measurements at different current densities; c Nyquist plots; d SEM image of Na-metal electrode; e SEM image of composite electrode. Reproduced with permission from Ref. [198]

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3 Secondary Seawater Batteries

cell, by contrast, maintained a low overpotential during over 200 h of continuous testing, with only a slight increase. A larger current density (3.0 mA cm−2 ) was then applied for further experimental verification (Fig. 3.89b). Even after 10 h, the bare Na foil cell exhibited unstable and fluctuating hysteresis, whereas the composite cell maintained a steady cycle performance for over 80 h. The stable performance of the composite cell was attributed to the carbon scaffold, which served as a framework for Na storage [201]. The Nyquist plots in Fig. 3.89c highlight the different characteristics of the composite and pure metal-based cells. At high frequencies, the semicircle represents resistance, attributable to the SEI layer and charge transfer. The composite cell demonstrated a more stable interface and more uniform deposits than did the bare Na cell (Fig. 3.89c). The ex situ SEM profiles highlight the differences between the bare Na metal and Na–carbon composite electrodes (Fig. 3.89d, e). The surface of the bare Na metal electrode was porous and irregular, whereas that of the Na– carbon composite was comparatively smooth. The formation of mossy dendrites and the uneven surface of the bare Na electrode increased its internal resistance, resulting in short-circuiting, whereas the composite electrode’s flat and smooth surface morphology resulted in stable long-term cycling. According to electrochemical analysis, alkali metals infiltrated into the nanocrevasses under the influence of capillary forces and adhered strongly to the carbon scaffold, consistently forming sites for Na metal deposition during electrochemical cycling [210]. The Na–carbon composite was then tested in an SWB (Fig. 3.90a). Coin-type SWBs were fabricated with two distinct anodes: Na–carbon composite and Nametal foil. Figure 3.90b depicts the battery performance of the two cells. As the current increased from 0.5 to 3 mA during galvanostatic cycling, the voltage profile of the Na–carbon composite cell presented smaller gaps than that of the Na-foil cell. The increased surface area of the Na–carbon composite reduced its local current density, significantly improving its performance, while producing smaller voltage gaps, at higher currents. The SWB was fully charged and discharged to confirm the reversibility of the Na–carbon composite (Fig. 3.90c). At full discharge (0 V), the specific capacity was 1106 mAh g−1 , which was 94.8% of the theoretical capacity of Na metal (1166 mAh g−1 ). The entirely stripped Na–carbon composite is depicted in the inset of Fig. 3.90c. By contrast, the specific capacity of the Na-foil cell, which included residual Na (inset of Fig. 3.90c), was 942.8 mAh g−1 , or 80.9% the theoretical capacity of Na. These findings confirmed that liquid electrolyte penetrated the crevasses of the Na–carbon composite via capillary forces during the stripping process, enabling the sufficient infiltration of Na metals. The liquid electrolyte that infiltrated in the crevasses was also thought to maintain an active channel for electric deposition, resulting in a 98% CE and a high stripping/deposition reversibility. Overall, introducing crevasses with liquid-electrolyte soaking behavior on carbon cloths enhances the utilization of Na metal. The shape and size of the Na–carbon composite could be easily modified by cutting the large composite sheet. The rectangular composite material was used as the anode in a prismatic SWB, which was demonstrated in the operation of a TV monitor

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185

Fig. 3.90 Application of the Na–carbon composite in a seawater battery: a general structure; b battery performances of Na foil and composite over time; c potential against capacity; d use of composite to power a TV monitor. Reproduced with permission from Ref. [198]

(Fig. 3.90d). These results demonstrate the scientific and practical advantages of alkali metal/carbon composites with nanocrevasses.

3.3.3 Na Intercalation in Hard Carbon Carbon materials are classified as soft or hard. Graphite, a soft carbon, is widely used as a negative-electrode material for lithium-ion batteries. In sodium-ion battery systems, however, it only delivers 31 mAh g−1 , reflecting the size disparity between sodium ions (0.102 nm) and lithium ions (0.076 nm). Lithium ions can be inserted between graphite interlayers, but the insertion of sodium ions is not thermodynamically advantageous [211]. As a result, sodium storage between graphite interlayers is uneven. Therefore, hard carbon is used in sodium-ion batteries as an alternative to graphite. Hard carbon is non-graphitizable; it typically consists of “turbostratic” nanodomains (few-layer stacks of graphite) with nanovoids interspersed between them [212]. Due to its low average operating potential and relatively high reversible

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specific capacity, hard carbon is a leading carbonaceous anode material for sodium storage in sodium-ion batteries. Although several mechanisms by which sodium ions can be stored in hard carbon have been suggested, the generally accepted one is the “house of cards” model proposed by Stevens et al. in 2000. In this model, sodium ions intercalate between several interlayers of hard carbon inserted into micropores and can also be adsorbed to special sites such as defects, edges, and functional groups. On the surface of hard carbon, solid electrolyte interphase is formed, or sodium metal is plated; additional sodium ions are sometimes stored in this way, but this is generally thought of as an irreversible capacity (Fig. 3.91). When hard-carbon sodiation begins, sodium ions begin to adsorb to defects and edge sites. Then, they fill in the micropores, up to a potential range of around 0.1 V compared to sodium metal. The intercalation reaction of sodium ions between Fig. 3.91 Structure of a hard-carbon particle: a in general; b with stored sodium ions, according to the house-of-cards model. Reproduced with permission from Ref. [213]

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187

Fig. 3.92 The Na+ -ion-storage mechanism of lignin-derived hard carbons. Reproduced with permission from Ref. [214]

graphite interlayers occurs from below 0.1 V to about 0.03 V. Finally, sodium clustering occurs near the cutoff potential, and sodium ions are additionally filled in the micropores of hard carbon [214]. Although hard carbon is a popular anode material in sodium-ion battery systems, some practical problems remain to be solved. The first is low theoretical capacity: hard carbon has a charge/discharge capacity of about 300 mAh g−1 (Fig. 3.92), which is very low compared to that of sodium metal (1166 mAh g−1 ). The initial irreversible capacity that occurs during the first charge/discharge process of hard carbon exacerbates this problem. This initial irreversible capacity in general anode materials is known to be the main factor in the formation of a solid electrolyte interface during the charging process, but in the case of hard carbon, another factor is added: some sodium ions are irreversibly trapped in graphene edges and defects, lowering the CE of the first cycle. The second problem is the conductivity of hard carbon, which is low for a carbon material. In fact, the conductivity of a hard-carbon electrode (0.48 S cm−1 ) is only about 1% that of a graphite anode (42.57 S cm−1 ). This results in a low rate capability and lowers the stability of the hard-carbon electrode at high current. Therefore, research on special treatment or remodeling of raw materials to overcome these characteristics of hard carbon is ongoing.

3.3.3.1

Hard Carbon Applied to Seawater Batteries

As mentioned earlier, hard carbon is the most widely used and promising anode material in sodium-ion battery systems. Since the anode part of an SWB is similar to

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Fig. 3.93 Electrochemical performance of a seawater battery with a hard-carbon negative electrode: a charge and discharge voltage curves of the half-cells and the full-cell; b observed seventh cycle of the full-cell (purple) and prediction from the two half-cell results (orange). Reproduced with permission from Ref. [215]

that of the existing sodium-ion battery, there have been many attempts to apply the hard-carbon anode to the SWB. In structure, sodium ions of seawater pass through the NASICON solid electrolyte and the sodium-conducting liquid electrolyte before being stored in the hard-carbon cathode. There is no longer a risk of explosions in the manufacturing process because sodium metal is not used, and the cycle/capacity retention problem due to metal dendrites (linked to explosions) is also avoided. When hard carbon was first used as a negative electrode in SWB systems, researchers began to look at the charging/discharging profile [215]. Figure 3.93 is a graph showing the charging/discharging profile of the fifth cycle by applying hard carbon to the metal counter-electrode system (half-cell) and the SWB full-cell. In both methods, organic electrolyte and NASICON solid electrolyte were used. Charging and discharging in the SWB system was performed by calculating a capacity of 300 mAh g−1 , known as the theoretical capacity of hard carbon. There was no significant difference between the actual charging/discharging profile of the cell and that predicted theoretically by considering the potential of hard carbon and OER/ORR (which is the cathode reaction of the SWB). Thus, energy can be stored by using a hard carbon anode in a SWB system.

3.3.3.2

Hybrid Hard-Carbon Composites for Higher Energy Density

By applying hard carbon to the SWB, the problems of explosion risk and metaldendrite formation encountered with previous metal anodes can to some extent be solved. However, because of the intrinsic problems of hard carbon mentioned in Sect. 3.3.3, further research is needed. The problem of low theoretical capacity could be overcome by using composites of hard carbon and a sodium-ion-storing polymer [216]. In one study, an oxidation– reduction reaction was performed on a hybrid anode made from such a composite. The

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189

chosen polymer, poly(4-styrenesulfonate) (PSS; C8 H7 O3 S), has sufonyloxy anions in its repeating unit, allowing it to engage swiftly with sodium cations in seawater and transfer them to the hard-carbon lattice [217]. The discharge potential was nearly identical to that of an anode made entirely of carbon. Although PSS alone has a low theoretical coulombic capacity (146 mAh g−1 ), when paired with hard carbon (372 mAh g−1 ), which is also redox-active at the anode, the theoretical coulombic capacity sharply increases to above 500 mAh g−1 . Because both hard carbon and PSS store sodium ions by mutual charge transfer, the functional redox-active anode material produced from them (HC-PSS) has a high discharge capacity. Furthermore, hard carbon enhances the electrical conductivity of organic materials as a bifunctional substance. Fast ionic interactions between the sulfonyloxy radicals and sodium cations increased the electrochemical reaction rate of the complete functionalized hybrid anode. The ion-transfer mechanism in functionalized HC-PSS is shown in Fig. 3.94a. Heterogeneous ion transfer from hard carbon to the coated PSS of the pendant sulfonyloxy group (the stable organic anion constituting the repeating unit in PSS) initiated during discharge. The ion was then homogeneously transferred to the surrounding pendant sulfonyloxy group by a rapid self-exchange reaction. The broadening of the electrolytic electron-spin-resonance (ESR) spectrum of HC-PSS shows intrachain dipole–dipole interactions due to the short distances between the polymerbound sites. The ion self-exchange reaction between adjacent pendant sulfonyloxy groups (Fig. 3.94b) involved a one-ion mechanism. The occurrence of both heterogeneous and homogeneous charge transfer allowed for a quick electrochemical response. Although an electronic conductor and able to store charges, hard carbon has a high reversible reactivity due to the dense SEI that forms on the electrode surface. In theory, each 6C unit of hard carbon can store one positive charge, and each monomer of PSS can also store one positive charge. A reversible sodium-ion transfer between the two components was detected in the HC-PSS composite. Because of its chemical interaction with PSS’s aromatic ring, hard carbon can be employed as a molecular current collector to insulate PSS (Fig. 3.95). The anode had a discharge capacity of over 500 mAh g−1 when the battery was discharged at 0 V. The theoretical capacity of a composite electrode made up of functional redox materials is the sum of those of its components [218]; the discharge capacity measured experimentally was 97% (500.6 mAh g−1 ) of the total theoretical capacity (518 mAh g−1 ). When the hard-carbon electrode made in the above method was applied in an SWB, the same effect of increasing capacity was obtained. Figure 3.96a shows charge/discharge data for 0.5C and 1C operation of SWBs with a HC-PSS negative electrode. The HC-PSS anode exhibited specific capacities of 383.2 mAh g−1 and 295.5 mAh g−1 at each current, respectively. Even when the current was increased, the electrode maintained a high discharge capacity because of the quick electron transport due to self-exchange processes through ionic contacts with pendant sulfonyloxy groups. The cycling performance of the SWB at 0.5C for 100 cycles is shown in Fig. 3.96b. HC-PSS had a capacity retention of 99.7% after 100 cycles, confirming the HC-PSS electrode’s stable cycling performance. The capacity at 0.5C exceeded

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Fig. 3.94 Hard-carbon–PSS hybrid material: a structure of hybrid, b electrochemical mechanism. Reproduced with permission from Ref. [216]

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Fig. 3.95 Schematic of charge transfer reactions in the HC-PSS anode. Reproduced with permission from Ref. [216]

the low theoretical capacity of unmodified hard carbon (300 mAh g−1 ), and the cycle characteristics of the SWB were also excellent.

3.3.3.3

Biomass-Derived Hard Carbon for Special Characteristic

The practical use of hard carbon in sodium-ion battery systems is obstructed by the low theoretical capacity, the high cost of the precursor, the low yield of synthesis, and the low initial CE due to formation of SEI. Carbons obtained from biomass have been considered to address these concerns; they are ecologically friendly, renewable, abundant, and inexpensive [219–222]. Hard carbons can be obtained from a variety of natural sources, many of which are unsuitable as food. An SWB system that uses pine-pollen carbon (PPC) for both anode and cathode active materials has been demonstrated in this paper. At the anode, the PPC electrode can be sodiated/desodiated, and at the cathode, it can catalyze ORR/OER processes. Symmetric batteries can considerably lower battery manufacturing costs due to their

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Fig. 3.96 Seawater cell with HC-PSS anode: a initial voltage profiles of cycling; b electrochemical performance. Reproduced with permission from Ref. [216]

easy fabrication technique and economies of scale. However, lithium or sodium ions are held in symmetric cells of traditional batteries utilizing differing reaction potentials of the same material. The energy density is restricted, despite the fact that it is cost effective. Carbon compounds generated by biomass pine pollen can lower the cost of making active components. Simple heat treatment of pine pollen yielded microporous self-N and -P dual-doped carbon. The sodium storage abilities of PPC were then investigated for SWBs. Heteroatom (N and P) doping was used to boost the electrical conductivity of carbon, and nano-structuring was used to construct many active sites with a large surface area [223–226]. The PPC was created from pine pollen using a simple and quick synthesis procedure that did not require any pre-treatment. Carbon was created by carbonization at 600 °C in an Ar gas environment, as shown in Fig. 3.97. After carbonization, activation was carried out at 400 °C in an air environment to create oxygen functional groups and more surface pores that boost metal-ion storage capacity [227].

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Fig. 3.97 Preparation of pine-pollen carbon. Reproduced with permission from Ref. [219]

The PPC was used in a SWB bifunctional electrode not only as an electrocatalyst in the cathode but also as an active material in the anode for sodiation/desodiation. Current–voltage measurements between 0.1 and 3 V (Fig. 3.98a) revealed two large peaks in the first discharge curve. The reactions between functional groups and sodium ions were responsible for the peak near 0.75 V, and an irreversible process leading to the creation of a SEI layer caused the uneven curve between 0.75 and 0.01 V. These characteristics were absent from the second and third cycles. The insertion or extraction of sodium ions into the graphitic structure was indicated by the peak near 0.01 V. The significant slope between 1 and 0.01 V was due to sodium-ion adsorption on the PPC electrode’s surface [228]. Figure 3.98b illustrates galvanostatic

Fig. 3.98 Sodium-storage performance of pine-pollen carbon in a coin-type half-cell: a current– voltage curves for first three cycles; b galvanostatic charge/discharge curves; c voltage vs capacity (comparison with hard carbon); d electrochemical performance. Reproduced with permission from Ref. [219]

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charge/discharge curves. Because of the creation of an SEI layer at the surface, the first discharge and charge capacity were 371 and 244 mAh g−1 , respectively, resulting in an initial CE of 65.7%. The significant capacitive contribution induced by sodium adsorption creates these sloping voltage profiles in the sodium-ion battery because PPC has a high doping level [229–231]. The reversible capacity was then compared to hard carbon at 1 A g−1 for the fifth cycle (Fig. 3.98c). PPC and hard carbon have discharge capacities of 156 and 103 mAh g−1 , respectively; their cycling abilities and CEs for 100 cycles at 1 A g−1 are shown in Fig. 3.98d. For 100 cycles, PPC’s discharge capacity remained higher than that of hard carbon. PPC thus demonstrated outstanding sodium storage capabilities in coin-type half-cells, making it better than normal hard carbon anode choice for SWBs. Figure 3.99 presents data from an SWB system using PPC for the anode active material; 191 mAh g anode−1 was achieved at the tenth discharge and 178 mAh g anode−1 at the 100th discharge (capacity retention rate: 93.4%). The CE was also sustained at over 90% for 100 cycles, showing PPC | PPC’s steady cycling performance at 50 mA g anode−1 . SWB was successfully produced by applying PPC to the anode, which has good electrochemical performance, cost-efficiency, and safety.

Fig. 3.99 Sodium-storage performance of pine-pollen carbon in a coin-type half-cell. Reproduced with permission from Ref. [219]

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Fig. 3.100 SEM images for a commercial hard carbon and b starch hard carbon, and c XRD spectra of both materials. Reproduced with permission from Ref. [31]

Like pine-pollen carbon, starch-derived hard carbon (S-HC) has been studied as an alternative to commercial hard carbon (C-HC). One study considered an S-HC with an aqueous binder (styrene-butadiene rubber and carboxymethyl cellulose) for the anode [31]. In comparison to a commercial product, the S-HC had enhanced structural and electrochemical properties. Rice starch was chosen as a precursor because of its widespread availability as a standardized chemical. The SEM micrographs shown in Fig. 3.100a and b reveal that the S-HC had a larger particle size with a smoother surface than the C-HC. X-ray diffraction was used to characterize the microstructural variation (Fig. 3.100c). The two broad diffraction peaks at 23° and 44° two degrees suggest that the hard carbons were amorphous and had a highly disordered structure, with C-HC showing more prominent peaks. The interlayer spacing of the (002) planes in S-HC (4.1) was larger than in C-HC (3.8), which favors sodiation/desodiation due to the larger interlayer gap [232]. S-HC was tested in seawater cells cycling at 0.2C and 0.5C rates for 100 and 400 cycles, respectively (Fig. 3.101). The cell performances at the two rates were quite similar. The capacity retention after cycling at 0.2C (Fig. 3.101b) and at 0.5C (Fig. 3.101d) shows that the seawater cell was quite stable. For example, at 0.2C, the 10th cycle discharge yielded 299 mAh g−1 while the 100th cycle yielded 273 mAh g−1 (discharge capacity retention: 91%). The results were 296 mAh g−1 in the 10th cycle and 291 mAh g−1 in the 400th cycle at 0.5C. (98%). In terms of cyclability, the SWB with the hard carbon electrode had better qualities than other types of electrodes. However, the low theoretical capacity (~300 mAh g−1 ) [233–236]—a

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Fig. 3.101 Cycling performance of seawater cells employing starch-derived hard carbon at cycling rates a, b 0.2C and c, d 0.5C. Reproduced with permission from Ref. [31]

difficulty typical of hard carbon materials in general—hampers the directionality of the SWB’s high capacity. The fact that the folding utilized in other battery systems cannot be employed in the SWB adds to the complexity of the situation. As a result, future study should focus on increasing either the capacity of hard carbon or the amount of hard-carbon active material inside a single SWB cell.

3.3.4 Alloy Materials In a sodium-ion-battery anode system, Group IVA and VA metals of the periodic table and their compounds (e.g., oxides, sulfides, and selenides) can achieve high capacities. They form Na-rich intermetallic compounds through alloy reactions [237]. For example, in the case of Sn, when a complete sodium ionization product (Na15 Sn4 ) is formed, a theoretical capacity of 847 mAh g−1 can be attained; this is more than twice the capacity of the carbon-based negative electrode used in the conventional lithiumion and sodium-ion batteries [238]. Red phosphorus has an extremely high theoretical capacity (2596 mAh g−1 ), making it potentially the highest-capacity material for sodium-ion battery cathodes [239].

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Despite this high theoretical capacity, there are many challenges in applying alloybased materials to actual sodium-ion-battery anode systems, in particular the serious capacity degradation that occurs during the charging and discharging process [240]. This is caused by a change in volume, leading to aggregation and pulverization of the active material, and thus to low rate capability caused by loss of diffusion distance and to loss of electrical conductivity [241, 242]. On the other hand, the negative electrode for the sodium-ion battery has a larger volume expansion than that of the lithium-ion battery because of the larger radius of Na+ , which makes the negative electrode more effective in suppressing the stress generated within the active material and buffering the capacity decay of the cell during sodiation and desodiation. In addition, a layer of SEI is continuously formed with the newly exposed surface of the active material during cycling [243–245]. The resulting thick film blocks charge transfer significantly and results in rapid capacitive fading. Moreover, the low initial CE due to the unstable SEI layer also represents an important problem for the sodiumion battery. The poor electrical conductivity of some alloy-based anodes (e.g., P and Si) dramatically limits the specific delivered capacity and rate capacity. Nevertheless, significant progress has been made with alloy-based anode materials. Strategies to improve the performance of alloy-type anodes mainly focus on designing efficient nanostructures and introducing conductive carbon hosts/substrates (e.g., carbon nanofibers or graphene), both of which effectively accelerate the reaction kinetics and mitigate capacity reduction. In addition, engineering bimetal alloys and utilizing appropriate binders and electrolytes (additives) have been found to be effective ways to improve sodium-storage properties (Fig. 3.102).

Fig. 3.102 a Schematic of the structural development of Sn during the sodiation process. Reproduced with permission from Ref. [246]. b Illustration of the pulverization process in crystalline alloys. Reproduced with permission from Ref. [247]

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Tin–Carbon

The high theoretical capacity of tin (847 mAh g−1 ) has sparked interest in employing a Sn-based anode material, resulting in the production of the intermetallic compound Na15 Sn4 . The sodium-storage capacity is usually undermined by the poor cyclability resulting from the large specific-volume changes in charging and discharging. However, the encapsulation of Sn nanoparticles in an amorphous yet conductive carbon matrix (Sn–C nanocomposite) enables the improved the performance of Na-storage [8, 248]. In one study, X-ray diffraction (XRD) spectra of the nanocomposite revealed the existence of metallic tin, while no traces of tin oxide species were observed (Fig. 3.103a). The presence of the Na15 Sn4 phase (JCPDS #31–1327) was clearly demonstrated by ex situ study of the Sn–C-based electrode. The high-resolution XPS core-level profiles of Na 1 s are shown in Fig. 3.103c. The intensity of the Na 1 s peak increased significantly after the first charge, indicating Na+ uptake during charging. Furthermore, the findings of an ex situ SEM–EDX examination of the Sn–C-based electrodes before and after electrochemical sodiation were consistent, which demonstrates a very homogeneous sodium distribution in the nanocomposite (Fig. 3.103d). These findings show that suitable high-capacity alloying materials can be used to replace extremely reactive metallic Na in SWB [249, 250]. For additional confirmation of sodiation in Sn–C composite, high-resolution transmission electron microscopy (HR-TEM) including FFT (fast Fourier transform) analysis of images was conducted. The average size of pristine Sn nanoparticles

Fig. 3.103 a X-ray diffraction patterns, b X-ray photoelectron spectra, and c SEM–EDX analysis of Sn–C anode stored in electrolyte and sodiated Sn–C anode. Reproduced with permission from Ref. [8]

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Fig. 3.104 Ex situ HR-TEM bright-field images of the Sn–C anode: a pristine state; b charged state. Reproduced with permission from Ref. [8]

was ~50 nm. FFT analysis of a selected area clearly showed that the observed dspacing corresponded to the (110) and (107) lattice planes of metallic Sn, confirming the results obtained by XRD (Fig. 3.104a). For the sodiated sample, slightly larger nanocrystallites were observed with an average size of ~70 nm. The larger size was due to the tin-nanoparticle volume expansion in the formation of Sn–Na alloy. Furthermore, FFT analysis of the fully sodiated sample revealed that the observed d-spacing was in very good agreement with the (332) and (310) lattice planes of Na15 Sn4 . This result confirms that, upon sodiation, the composition reaches the end member of the electrochemical alloying process of Sn and Na, Na15 Sn4 . When the Sn–C nanocomposite was employed as the anode in a full-cell design, a first-cycle irreversible capacity of 200 mAh g−1 and a reversible capacity of 300 mAh g−1 were demonstrated. The high irreversible capacity was attributed to electrolyte decomposition at the particle surface, which resulted in the formation of an SEI, as well as the structural rearrangements that occurred within the micronsized composite particles during the first sodiation stage. Following this process, the reversible capacity grew to >300 mAh g−1 by the fifth cycle, with the irreversible capacity steadily declining (90 mAh g−1 by the fifth cycle) [251]. Generally, the cyclability of both the hard carbon and Sn–C nanocomposite anode materials in SWB is very stable upon continuous cycling, showing a remarkably low

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Fig. 3.105 Seawater cell with Sn–C anodes: a voltage vs capacity in various cycles; b cyclability comparison with hard carbon. Reproduced with permission from Ref. [8]

capacity fading after 30 cycles (Fig. 3.105b). These results again highlight the great advantage of accessing an almost infinite supply of sodium ions by employing the open-system seawater cathode to compensate the irreversible capacity.

3.3.4.2

Antimony Sulfides (Sb2 S3 )

An a-Sb2 S3 nanoelectrode exhibits superior Na ion storage capability and electrochemical properties, such as a stronger cycle stability and a higher rate capability, compared to a commercial crystalline Sb2 S3 electrode. Because of these advantages, the a-Sb2 S3 nanoparticle has been applied as an anode in SWBs. A new synthesis method for a spherical, amorphous (a-)Sb2 S3 nano-aggregate anode using a simple polyol route at room temperature, in contrast to other complicated synthetic pathways (mostly, hydrothermal methods) under high-temperature conditions to form crystalline Sb2 S3 , was reported [252–254]. The a-Sb2 S3 nanoparticles were synthesized by reacting an Sb precursor solution (SbCl3 ) with a sulfur precursor solution (thioacetamide) at room temperature in ambient air. As the reaction proceeded, the clear solution mixtures became opaque orange and ultimately crimson. Red precipitates were obtained using washing and drying processes, and the chemical structure was confirmed using electron microscopy and XRD. SEM images of the resulting Sb2 S3 particles are shown in Fig. 3.106a. The particles were composed of spherical aggregates of sub-constituent nanoparticles with sizes ranging between 150 and 300 nm, as verified by TEM (Fig. 3.106b). Figure 3.106c illustrates the XRD pattern of the as-synthesized nanoparticles; it exhibits typical amorphous-species peaks. Initially, a Sb-glycolate complex was formed by dissolving the SbCl3 precursor in EG medium, which acted as a chelating agent, preventing the quick hydrolysis and oxidation of SbCl3 (SbCl3 + H2 O → SbOCl + 2HCl); EG also acted as a dispersant, facilitating the creation of uniform Sb2 S3 spheres. The interaction with thioacetamide, which generated S2– ions in the EG medium, progressively converted the Sb-glycolates to SbSx

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Fig. 3.106 a SEM and b TEM images, c XRD patterns, and d Raman spectrum of the synthesized (red) Sb2 S3 nanoparticles. Reproduced with permission from Ref. [11]

nuclei, the development (Ostwald ripening) of which caused the spherical nanoparticles to aggregate into an amorphous shape. The as-prepared nanoparticles were then heated to 250 °C for crystallization in the presence of an Ar flow. As seen in the center of Fig. 3.106c, the sample exhibited crystalline XRD peaks corresponding to the orthorhombic phase of Sb2 S3 following heat treatment. Raman spectroscopy was also employed to investigate the structural characteristics (Fig. 3.106d). The Raman spectra of the as-prepared nanoparticles revealed very wide bands between 139 and 298 cm−1 , indicating poor crystallinity. By contrast, the spectra of the heat-treated nanoparticles displayed several sharp peaks and matched the spectrum of crystalline Sb2 S3 . For instance, the bands at 281 and 309 cm−1 were attributed to the symmetric vibrations of SbS3 pyramidal units with C 3v symmetry, whereas the bands at 192 and 239 cm−1 revealed the crystalline nature [252]. Galvanostatic cycling with a capacity cut-off (550 mAh ganode −1 ) during charging and a voltage cut-off (0.4 V) during discharging at a current density of 0.05 mA cm−2 was used to test the full-cells. The measured charge–discharge curves for the first and subsequent cycles were very similar to the calculated curves, as shown in Fig. 3.107b

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Fig. 3.107 Voltage profiles of a the antimony sulfide anode half-cell and b the full-cell under capacity cut-off. c Cycling performance of the seawater full-cell. d XRD pattern of the anode, which was disassembled after the 15th full-cell charge process. Reproduced with permission from Ref. [11]

(the bold curves in Fig. 3.107a). The initial discharge capacity was 233 mAh g−1 , resulting in a 42.3% CE. The formation of an SEI layer on the a-Sb2 S3 nanoparticle anode was the main cause of the low initial CE. The lower initial CE compared to the anode half-cell (Na|a-Sb2 S3 , 65%) could be due to capacity-limited charging (550 mAh ganode −1 ), which prevented the a-Sb2 S3 nanoparticle anode from reaching its full charge state. As shown in Fig. 3.107c, discharge capacity and CE rose with increasing cycle number and stabilized around the 50th cycle; particularly, the discharge capacity was 470–485 mAh g−1 with an average discharge voltage of 1.9 V and CE of 83%– 88% during the 50–70th cycles. Subsequently, the charged anode was exposed to ambient air in order to analyze sodiation of the a-Sb2 S3 nanoparticle anode by natural seawater. The XRD pattern (Fig. 3.107d) revealed peaks corresponding to the crystalline NaSb(OH)6 phase, indicating that the a-Sb2 S3 nanoparticle anode stored the Na-ions from natural seawater during the charge process.

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Phosphorus–Carbon

Amorphous-red-phosphorus–carbon (PC) composites have recently been studied as promising anode materials for Na ion batteries because they exhibit a high reversible capacity of 1890 mAh gphosphorus −1 (=1323 mAh gcomposite −1 ), along with a strong cyclability. However, owing to their high irreversibility in the initial cycles, redphosphorus-based electrodes in full-cell systems are rarely studied. The inherent difficulty of the huge irreversible capacity of alloying reaction-based electrode materials, such as phosphorus, could be overcome by SWBs with open-structured cathodes. As a result, PC composite could be a potential anode for SWBs [255–258]. One study used PC as a high-capacity anode material in a rechargeable SWB, attaining a reversible capacity exceeding 920 mAh gcomposite −1 over 80 cycles, as well as an excellent rate capability, with a discharge capacity of more than 850 mAh g−1 even at the high current density of 2.0 A gcomposite −1 [30]. As shown in Fig. 3.108, the SWB achieved a high specific capacity (maximum: 950 mAh g−1 ) and a high specific energy (maximum: 2.4 W h g−1 ). These findings show that PC composite could be a promising anode material for SWBs. A simple high-energy ball-milling procedure was used to make the PC composite. The XRD patterns of the raw red-phosphorus powder, carbon black, and as-prepared

Fig. 3.108 Comparison of the specific capacity and voltage of sodium-ion battery systems (SIBs) and lithium-ion battery systems (LIBs). (PC: amorphous-red-phosphorus–carbon). Reproduced with permission from Ref. [30]

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Fig. 3.109 Study of amorphous-red-phosphorus–carbon (PC) composite: a XRD and b Raman spectra of PC composite, red phosphorus, and carbon black (Super P(C)); c, d SEM images of the PC composite cross-section; e EDS and f TEM and EDS images of PC composite. Reproduced with permission from Ref. [30]

PC-composite powder are shown in Fig. 3.109a. The XRD patterns of the first two substances featured broad peaks, which are characteristic of amorphous species. The XRD pattern of the PC composite, by contrast, had larger peaks, indicating that the high-energy ball-milling process had induced amorphization. As shown in Fig. 3.109b, the Raman spectrum of red phosphorus was normal, whereas the Raman spectrum of carbon contained large bands at 1350–1584 cm−1 ; these were assigned to the D and G bands, respectively. The Raman spectrum of the PC composite revealed no visible bands corresponding to red phosphorus, but it did reveal D and G bands that were significantly stronger than those of pristine carbon black. Figure 3.109c shows an SEM image of the PC-composite particles; as a result of the ball-milling process, they comprised several constituent particles with diameters of several tens of micrometers. The cross-section of a PC-composite particle was observed using dualbeam focused ion beam (FIB) spectroscopy. The porous nature of the cross-section (Fig. 3.109d) is advantageous for reducing the substantial volume changes in the PC composite and facilitating ion movement during repeated charging and discharging. In Fig. 3.109e, the EDS line profile of the cross-sectional surface is also shown. The distribution of phosphorus, carbon, and oxygen within the PC composite particle was found to be uniform, which was confirmed by the TEM-EDS result (Fig. 3.109f). The PC composite appears to be made up of amorphized red-phosphorus particles mixed uniformly with carbon nanoparticles, resulting in a reduction in red-phosphorus peak intensity. The structural characteristics of the composite help to maintain physical contact between the phosphorus and carbon particles and improve the electrical properties of the electrochemical reactions during cell operation [255]. The voltage profiles of the SWB full-cells are shown in Fig. 3.110a. The cells were cycled with a capacity cut-off of 1000 mAh g composite −1 while charging and a voltage

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Fig. 3.110 a Voltage profiles and b coulombic efficiency of a seawater full-cell including an amorphous-red-phosphorus–carbon composite anode. Reproduced with permission from Ref. [30]

cut-off of 0 V while discharging. The measured voltage profiles were nearly identical to the calculated ones. The voltage difference between charging and discharging widened as the cycle progressed. As shown in Fig. 3.110b, the discharge capacity of the first cycle was around 700 mAh gcomposite −1 with a CE of 70%, while discharge capacities of subsequent cycles rapidly increased to over 950 mAh gcomposite −1 with CEs of 90–96%. The continuous development of an SEI layer by capacity-limited charging during the initial cycles and the gradual activation of the PC composite over the subsequent cycles explain this rise in discharge capacity with cycle number. The formation of SEI layers and structural rearrangement of the PC composite during the

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charge process of the initial cycles were linked to the decomposed organic electrolyte on the solid surfaces (PC composite and NASICON), resulting in the relatively high irreversible capacity [8, 11]. In other battery systems, due to the irreversible capacity in the first cycle, all subsequent cycles are affected, but in the SWB system, infinite capacity can be continuously supplied because of the open-system cathode. As a result, even a red-phosphorus material with a poor initial CE can be used with fine cycle retention.

3.3.4.4

Phosphorus–Carbon Semi-liquid Anodes

A system in which the cathode can provide infinite capacity to the anode based on infinite seawater, and the capacity of the anode is determined by the capacity of the SWB, opens new horizons for anode research. The current electrode system uses an electrolyte with no electronic conductivity as well as a thin electrode. However, the use of solid electrolytes did not force the non-electrolyte solution in SWBs, and enables the use of electronically conductive materials. Moreover, with an electroconductive electrolyte, the electrode needs no longer be flat, and charging and discharging are possible even if only the active material exists without additional materials for electrode manufacturing such as binders or conductive materials. This is a form in which an active material of powder exists in an electroconductive electrolyte, and it was named “Semi-liquid anode”. Such a semi-liquid anode based on phosphorus and featuring redox mediators has been developed [259]. The P/C composite powder was mixed with a redoxmediator-based electrolyte solution to make a red-phosphorus semi-liquid anode. Two polycyclic aromatic hydrocarbons (PAHs) were chosen as redox mediators (couples) for the anode. PAHs react with alkali metals to produce complexes of alkali-metal cations and radical anions; when in contact with Na metal, sodium– polycyclic aromatic hydrocarbons (Na-PAHs) are formed. Sodium-biphenyl (NaBP) has been investigated as a liquid anode material for sodium-based batteries; it has good reversibility, low redox potential (vs Na/Na+ ), and appropriate electronic and ionic conductivities. For the P/C composite, Na-BP and sodium-pyrene (NaPYR) were utilized as reducing and oxidizing redox mediators, respectively. The results showed that the red-phosphorus semi-liquid electrode with redox mediators could provide a high-areal-capacity anode for sodium-based batteries, particularly Na-SWB [177, 184]. Na-BP and Na-PYR, as indicated in Fig. 3.111a, act as reducing and oxidizing mediators (couples), respectively. Figure 3.111b shows the working mechanism of these redox mediators in greater detail. Because of the volumetric changes of the P/C composite and the small contact area between the P/C particles in the dispersion, a large polarization is normally observed when sodiated red phosphorus is desodiated. Instead, in the presence of Na-PYR, Na-PYR is desodiated. The P/C composite is then chemically oxidized by the desodiated PYR, a charge transfer reaction. When BP is sodiated, the resulting Na-BP reduces the P/C composite chemically (a charge transfer reaction). In the fifth cycle, differential capacity plots (Fig. 3.111c)

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Fig. 3.111 Red-phosphorus semiliquid anode: a reaction mechanism; b working mechanism of redox mediators in terms of redox potential; c comparison of dQ/dV plots of half-cells (fifth cycle) with phosphorus–carbon composite foil electrode. Reproduced with permission from Ref. [259]

of half-cells with either a conventional P/C composite electrode or Na-PAHs anolytes including supporting electrolyte, Na-BPs and Na-PYRs. In conclusion, the sodiation voltage of Na-BP is lower than that of P/C composite, while the desodiation voltage of Na-PYR is higher. As a result, these two Na-PAHs could be used as redox mediators in P/C-based anodes. Figure 3.112a shows electrochemical impedance spectra of semi-liquid half-cells. In the case of anolytes containing Na-PYRs, the charge-transfer resistance (especially between the working electrode and the NASICON solid electrolyte) was smaller. In comparison to Na-PYRs, the charge transfer resistance was highest when only Na-BPs was present; P/C @ Na-BP-PYRs, Na-PYRs, and Na-BP-PYRs. Densityfunctional theory (DFT) calculations (Fig. 3.112b, c) back up these findings, showing that the sodium-cation charge is higher in Na-PYR (0.87, Fig. 3.112b) than in NaBP (0.76, Fig. 3.112c). Furthermore, the reduction reaction in the PYR-containing

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Fig. 3.112 a Electrochemical impedance spectra of half-cells before the cycling test in opencircuit voltage (OCV) state. Density-functional-theory calculation of b Na-PYRs and c Na-BPs. Reproduced with permission from Ref. [259]

system with G = 222 kJ mol–1 is clearly preferable to the BP-containing system with G = 193 kJ mol–1 . Because the PYR system is more stretched than the BP system, it can accommodate the electron more easily. A small Na-SWB coin cell setup was used to demonstrate the possibility of achieving higher reversible capacity. Although the overpotential varied with current density, the cell had high capacities of 15 mAh (= 7.5 mAh cm–2 ) and stable voltage profiles over 10 cycles, as shown in Fig. 3.113a and b. Furthermore, as shown in Fig. 3.113c, the cell capacity could be increased up to 30 mAh (= 15 mAh cm–2 ) for five cycles by increasing the amount of P/C @ Na-BP-PYRs.

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Fig. 3.113 Voltage profiles of Na–seawater full-cells including P/C @ Na-BP-PYRs. Reproduced with permission from Ref. [259]

Figure 3.113d shows a test where the P/C @ Na-BP-PYRs was totally discharged; it delivered a capacity of roughly 26 mAh (= 13 mAh cm–2 ). Based on total anode mass and volume, the resulting energy densities were roughly 2200 Wh kg–1 and 237 Wh L–1 , respectively. The cell was then fully charged and cycled at 15 mAh (= 7.5 mAh cm–2 ), demonstrating good reversibility. This is a key aspect of the Na–SWB system, as it allows for infinite Na-ion harvesting from seawater.

3.3.5 Liquid Electrolytes and Anolytes Liquid electrolytes of SWBs mainly affect the stability of the anode parts. Any existing sodium-based liquid electrolyte could be used in an SWB. As described above (Sect. 3.3.1), a liquid electrolyte requires high ionic conductivity, a wide electrochemical stability window (ESW), and thermal stability. Ionic conductivity is crucial for a battery that operates by Na+ transport, and a sufficiently rapid ion mobility is important for ensuring that the cell has high power. The anion type of the Na salt and the solvent according to the viscosity has a great influence on the ion

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conductivity of a non-aqueous electrolyte. The electrochemical stability window of an electrolyte is the energy gap between its LUMO and HOMO. If the LUMO level of the electrolyte is lower than the energy level of the anode, the electrolyte accepts electrons and decomposes on the anode surface; if the HOMO level is higher than the energy level of the cathode, the electrolyte donates electrons and decomposes on the cathode surface. Through the decomposition reaction, an SEI layer is generated on the surface of an electrode. There are no studies reported in the literature describing how to make a stable SEI layer or a stable electrolyte in which an SEI layer is not generated. The safety of liquid electrolytes has also recently attracted attention: it is important to prevent fires from battery heat generation. The thermal stability of liquid electrolytes can be confirmed by analytical methods such as differential scanning calorimetry (DSC) and thermogravimetric analysis (TGA). There are three broad categories of liquid electrolytes: organic liquid electrolytes, ionic liquids, and anolytes [183].

3.3.5.1

Organic Liquid Electrolytes

A metal-free SWB consisting of a hard carbon anode and the TEGDME-based electrolyte was studied in 2014 [135]. The electrolytes used in this work were 1 M NaClO4 in EC/PC and 1 M NaCF3 SO3 in TEGDME. It was the first study to look at the effect of the electrolyte on the performance of a SWB and showed that TEGDME improves the battery cycle life. In the SWB, the EC/PC and TEGDME-based electrolytes were evaluated in terms of charge/discharge performance at room temperature; typical flat curves were obtained for the charge and discharge reactions with a voltage separation of 1.2 and 1.08 V for the EC/PC and TEGDME-based electrolytes, respectively (Fig. 3.114a). These results suggested that the latter had a lower cell resistance, and

Fig. 3.114 Cycling performance of seawater batteries with different electrolytes. Reproduced with permission from Ref. [135]

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thus outperformed the former in terms of higher-charge delivery and discharge capabilities when evaluated at the same current densities. During discharge, TEGDME exhibited a greater active-material utilization rate, but its irreversible capacity was lower than that of the EC/PC. In the first 30 cycles, both electrolytes showed stable cycle performance (Fig. 3.114b). After 30 cycles, however, the cycle performance of EC/PC rapidly declined until reaching 80 cycles; by contrast, cycle performance of TEGDME remained constant after 100 cycles. As a result, when the TEGDMEbased electrolyte was applied, the cell’s general electrochemical performance was better. For sodium-ion batteries with alkyl carbonate-based electrolytes, discharge chemicals such as Na2 CO3 , sodium alkyl carbonate (NaOCO2 R), and polymers such as polycarbonate and polyethylenoxide have been reported to produce an SEI. Also, TEGDME generates a SEI layer with the decomposition products of AOCH2 R and A2 CO3 (A = Li or Na) [260–262]. The surface layers generated by the electrochemical reactions of the two electrolytes were investigated using XPS measurements (Fig. 3.115). The strong peak in the pristine hard-carbon anode spectra was assigned to the C–C bond of graphene. The other peaks came from the polyvinylidene difluoride binder (PVdF; CF2 ) and the hard-carbon electrodes’ functional groups [263]. Na2 CO3 and NaOCO2 R were allocated to the peaks at 290.0 eV, while CH2 was assigned to the peaks at 284.9 eV. The peak intensities associated with CH2 and CO3 in EC/PC were significantly higher than those in the TEGDME electrolyte. Because the CH2 was derived from the alkyl/alkylene groups and polymer species, the surface film formed by the EC/PCbased electrolyte contained a higher concentration of hydrocarbon compounds. Furthermore, the CO3 peak intensity stemming from the alkali and alkyl carbonates was higher in EC/PC, implying a thicker SEI layer. These findings revealed that the intensity of the CF2 peak in the EC/PC spectra decreased, in contrast to the TEGDME case. Furthermore, the F 1 s spectra confirmed the decomposition of the binder (Fig. 3.115). The C-F bond of the PVdF binder was assigned to a single peak in the pristine hard-carbon anode; there were two peaks in the cycled hard carbon (C–F and NaF). The peak with lower binding energy indicated the decomposition of the binder during the electrochemical reaction; this implies that the fluorine partially reacted with sodium when the EC/PC electrolyte was used, because there was no fluorine source in the EC/PC electrolyte itself. Furthermore, the degradation of the binder could result in the insulation of electron transfer on the hard-carbon anode. In the TEGDME-based electrolyte, however, the binder did not degrade; instead, the electrolyte partially dissolved to generate NaF, causing the C-F peak to move to a lower binding energy [260, 262–264]. Thus, the SWB with 1 M NaCF3 SO3 in TEGDME had a high cycle performance, due to an efficient SEI layer and good binder stability.

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Fig. 3.115 X-ray photoelectron C 1 s and F 1 s spectra of the surfaces of hard-carbon anodes cycled in various liquid electrolytes. Reproduced with permission from Ref. [135]

3.3.5.2

Ionic Liquids

Because of their safety and low environmental impact, ionic-liquid-based electrolytes (ILEs) are very appealing for use in electrochemical energy-storage systems. An ILE allows for stable anode cycling and prevents insoluble decomposition products from forming at the solid/liquid electrolyte interface. SWBs with ILEs could be used as large-scale green energy-storage systems because of their improved performance. Ionic liquids have a low vapor pressure, are thermally stable, have good ionic conductivity, and have a large electrochemical window.

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213

An ILE composed of 1:9 mol/mol NaTFSI/Pyr14 TFSI [Pyr14 = N-butyl-Nmethyl-pyrrolidinium, TFSI = bis (trifluoromethanesulfonyl)imide] was studied in 2015 [12]. Sn–C was utilized as an anode in this investigation, and the performance of the ILE was compared to that of a standard liquid electrolyte (LE). At room temperature, the ionic conductivity of the ionic liquid was lower than that of the LE. However, because of its superior interfacial characteristics with the Sn–C anode, the ILE considerably improved the cycle performance of the SWB. Figure 3.116a shows the calculated voltage efficiency in the first cycle: ~69%. The Sn–C anode, by contrast, stored sodium ions reversibly inside its structure at low voltage (1.0 V vs. Na/Na+ ). Figure 3.116b depicts the voltage curves of the Sn–C/ILE/seawater cell, which displayed a gradually shifting voltage profile with an average value of 2.7 V and a discharge capacity of 251 mAh g−1 , as expected from the half-cell results. In the first cycle, the cell with the LE (Fig. 3.116c) yielded a discharge capacity of 328.9 mAh g−1 , albeit with a substantial irreversible capacity (180 mAh g−1 ). At 0.03 mA cm−2 , the second charge and discharge capacities were 452 and 319.7 mAh g−1 , respectively, showing a low CE in the second cycle. Overall, the charge and discharge capacities of the Sn–C/ILE/seawater cell (Fig. 3.116b) were lower; however, a significant reduction in the irreversible capacity resulted in higher coulombic and energy efficiencies. Figure 3.116d compares the long-term cycle

Fig. 3.116 The voltage profiles and cyclability of seawater full-cells and half-cells containing conventional and ionic liquid electrolytes. Reproduced with permission from Ref. [12]

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Fig. 3.117 Liquid electrolyte (LE) and ionic-liquid electrolyte (ILE): a thermogravimetric analysis; b flammability test; c linear-sweep voltammetry. Reproduced with permission from Ref. [12]

performance of the ILE- and LE-based Sn–C/seawater cells. The increased penetration of the ILE into the electrode was a major factor in the improved performance observed during cycling. After 10 cycles, the ILE-based cell had reached the same level of performance as the LE-based cell. More importantly, the ILE-based cell’s performance remained consistent in the following cycles, whereas the LE-based cell’s performance rapidly declined after 30 cycles, falling significantly below that of the ILE-based cell. Thermal stability is another important property that electrolytes should exhibit. To evaluate the thermal performance of the LE and ILE, thermogravimetric analysis (TGA) was performed in air (Fig. 3.117a). Because of its extreme volatility, the LE demonstrated severe weight loss at 110 °C. Furthermore, the high reactivity of the NaClO4 salt resulted in total breakdown of the LE at 250 °C, with a 100% weight loss. By contrast, the ILE exhibited signs of thermal degradation only above 380 °C, with a 95% weight loss; this indicates its good thermal stability in air up to a relatively high temperature. Finally, linear sweep voltammetry (LSV) measurements were conducted to determine the ESW of the electrolytes (Fig. 3.117c). In the potential range of 2.0–4.0 V, no electrochemical reactions were identified for either electrolyte. Nevertheless, the LE demonstrated a lower anodic stability than the ILE, as evidenced by their oxidation levels at ~4.6 and 4.8 V (vs. Na/Na+ ). The LE showed a minor peak at 0.4 V after reduction, attributable to the production of ClO2 and NaCl due to ClO4 anion decomposition. The ILE, by contrast, had excellent cathodic stability down to 0 V

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Fig. 3.118 Chemical structure of the ions in ILE-EC, an ionic liquid electrolyte containing ethylene carbonate. Reproduced with permission from Ref. [265]

(vs. Na/Na+ ), which was theoretically attributed to the formation of an SEI on the stainless-steel electrode surface. A study of the ILE NaFSI-Pyr13 FSI-Pyr13 TFSI-EC (0.1/0.6/0.3/0.5wt %) was published in 2019 [265]. The fabrication process for this ILE, which is comprised of sodium bis(fluorosulfonyl)imide (NaFSI, 0.1 mol fraction), N-methyl-Npropylpyrrolidinium bis(fluorosulfonyl)imide (Pyr13FSI, 0.6 mol fraction), and Nmethyl-N-propylpyrrolidinium bis(trifluoro-methanesulfonyl) (Pyr13 TFSI, 0.3 mol fraction) was described. EC (5 wt%) was also added. In combination with the FSI anion, this allowed the production of a stable and protective SEI layer on reactive anode surfaces [265–272] (Fig. 3.118). Exothermic peaks were present in ILE-EC at –84.3, –94.7, –115, and –123.8 °C, and endothermic peaks at –96.7, –83.2, and –32.1 °C in the first cycle. There were two exothermic peaks at –94.7 and –126 °C in the second cycle, but only one endothermic peak at –96.7 °C. ILE and ILE-EC did not display any significant features at temperatures over –30 and –40 °C, respectively, despite the behavior of the pure components, namely the two ILs and EC. The crystallization process that manifested in ILE below –30 °C was not detected in ILE-EC, indicating that EC had a role in this behavior [270, 273]. Figure 3.119b shows the ionic conductivities of Pyr13 FSI, ILE, and ILE-EC as a function of temperature. Although the ionic liquid, Pyr13 FSI, was still solid, its ionic conductivity increased sharply along with the temperature, indicating the high mobility of the ions. As the temperature increased to –10 °C, the ionic conductivity climbed over two orders of magnitude (from 5.8 × 10–3 mS cm−1 to 5.5 × 10–2 mS cm−1 ). The conductivity of this IL approached 7.6 mS cm−1 at the ambient temperature of 20 °C and grew further to 20 mS cm−1 at 60 °C (Fig. 3.119b). The onset of conductivity in the ternary ILE occurred around –35 °C, after which the ionic conductivity increased along with the temperature. The conductivity rise occurred over a wide temperature range (from –35 to –25 °C), showing that the ions gained mobility before full melting. The ILE conductivity approached 4 mS cm−1 at the ambient temperature of 20 °C, and climbed to 13.3 mS cm−1 at 60 °C. Finally, in the temperature range studied, ILE-EC electrolyte did not demonstrate any significant variations in terms of conductivity: its ionic conductivity was 0.1 mS

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Fig. 3.119 Characterization of two ionic liquid electrolytes (ILE and ILE-EC) and their components. Reproduced with permission from Ref. [265]

cm−1 at –40 °C, 5 mS cm−1 at 20 °C, and 15.8 mS cm−1 at 60 °C. The quaternary electrolyte consisting of Pyr13 FSI, Pyr13 TFSI, NaFSI, and EC demonstrated the largest liquidus temperature range: its ionic conductivity at and above ambient temperature (20 °C ) was comparable to that of Pyr13 FSI, a highly conductive ionic liquid. The modest amount of organic solvent (EC) in the electrolyte, as well as the mixture of distinct anions (FSI and TFSI), which interfered with ion packing, were responsible for these features. The TGA curves of ILE and ILE-EC, as well as those of their components (Pyr13 FSI, Pyr13 TFSI, and EC), are shown in Fig. 3.119c. The lowest thermal signature observed was the evaporation temperature of EC, which was above 120 °C, as expected. However, no weight loss was detected in ILE-EC at temperatures below 200 °C, showing that the molecular solvent interacted significantly with the IL and Na-salt ions. The decomposition of the FSI anion in Pyr13 FSI occurred at ~250 °C, which was also the case for the ILE and ILE-EC electrolytes. The TFSI anion, in contrast, disintegrated in a single abrupt step at a significantly higher temperature, concurrently with the Pyr13 + cation. Overall, the TGA results suggested an exceptional thermal stability for the ionic liquid-based electrolytes, far exceeding the battery working conditions. To further examine the combination of the hard carbon anode and the ILE-EC anolyte in Na-SWBs, the performances of the ILE-EC and a standard electrolyte (1.0 M NaCF3 SO3 in TEGDME, referred to hereafter as LE) were experimentally

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217

compared. The outcomes are depicted in Fig. 3.120. When their cycling performances were compared at 1.0C (Fig. 3.120a, b), it was clear that ILE-EC outperformed LE. The anolyte-containing cell exhibited a high reversible capacity and consistent cycling performance over 300 cycles, whereas the capacity of the LE-containing cell began to diminish after 150 cycles. During the first 100 cycles, the difference of the dQ/dV peaks during charging and discharging was within 0.1 V (3.2 V and 3.1 V, respectively). Figure 3.120c shows a direct comparison between the cell voltage and discharge capacity at the 100th cycle for the two cells. While the areas under the two discharge curves were identical, the area under the LE-containing cell’s charging curve was significantly higher, indicating a lower energy efficiency. The ILE-EC and LE cells

Fig. 3.120 Comparison of the electrochemical performance of Na–seawater cells employing ILEEC or LE: a, b cycling performance; c cell voltage against specific capacity in the 100th cycle; d differential capacity against voltage; e energy efficiency. Reproduced with permission from Ref. [265]

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had a specific energy efficiency of 83.4% and 76.3%, respectively, indicating the higher energy efficiency of the electrolyte in the former. When the dQ/dV curves of the two electrolytes at the 100th cycle were compared (Fig. 3.120d), it was obvious that the LE-containing cell exhibited a substantially larger overvoltage when charged. As seen in Fig. 3.120e, this had a significant impact in terms of energy efficiency during cycling, with LE’s energy efficiency being unstable and significantly lower than that of ILE-EC. Overall, the results confirmed that the ionic-liquid-based quaternary anolyte (ILEEC) significantly improved the reversible capacity, cycle life, and energy efficiency of the Na–seawater cell in which it was tested.

3.3.5.3

Anolytes

The electrolyte requirements of SWBs differ from those of other sodium-based batteries because a solid-state electrolyte physically separates the cathode and the anode. The most notable difference is that liquid electrolytes can have electronic conductivity. For this reason, various types of anolytes that stably perform redox reactions but have both ionic and electronic conductivity have been applied in seawater cells. For example, a saturated sodium/3 M biphenyl (BP)/dimethoxyethane (DME) solution (Na-BP-DME) was proposed as a redox-active functional anolyte capable of allowing high SWB cyclability [181]. Redox-active organic compounds have been frequently used as electrode materials for rechargeable battery systems, but not for SWBs. Unlike traditional SWB electrolytes, the Na-BP-DME electrolyte solution has shown exceptional chemical and electrochemical stability during electrochemical cycling, making it compatible even with high-capacity Na-metal anodes. The sodiation of BP occurs before the Na-metal plating reaction during the charging process, because of the low reduction potential of Na-BP-DME (0.09 V vs. Na/Na+ ) [177, 274–278]. Also, BP is desodiated after the stripping of Na metal while discharging (Fig. 3.121). The conductivities of Na-BP-DME and NaCF3 SO3 -TEGDME, when used as liquid electrolytes in SWBs, are directly compared in Table 3.6; their respective overall conductivities, as determined by a conductivity meter, were 9.04 and 0.83 mS cm−1 . The isothermal transient ionic current (ITIC) approach was used to separate the components of ionic and electronic conductivities in Na-BP-DME. The ionic conductivity of Na-BP-DME anolyte was measured at 1.81 mS cm−1 when 0.1 V was applied to the samples; this is more than twice that of NaCF3 SO3 -TEGDME [279, 280]. Figure 3.122a shows the voltage profiles of a seawater cell using Na-BP-DME as the electrolyte; for more than 100 cycles, the cell maintained a stable cycling performance. On the other hand, the seawater cell utilizing NaCF3 SO3 -TEGDME as the electrolyte performed poorly over 100 cycles, presenting an erratic voltage profile (Fig. 3.122b). Na-BP-DME was totally sodiated during synthesis, because of the chemical saturation of the Na metal. Owing to the ease of electron transfer between

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219

Fig. 3.121 Working mechanism of redox-active material Na-BP-DME. Reproduced with permission from Ref. [181] Table 3.6 Electrical conductivity of two electrolytes. Reproduced with permission from Ref. [181]

Electrolyte

Conductivity [mS cm−1 ] Total

Ionic

Electronic

Na-BP-DME

9.04

1.81

7.23

NaCF3 SO3 -TEGDME

–

0.83

–

Fig. 3.122 Seawater cells using Na-BP-DME and NaCF3 SO3 -TEGDME as the electrolytes: a, b cycle performance of coin cells; c, d differential electrochemical mass spectra; e 1H NMR (400 MHz, CDCl3 ) spectra of BP-DME. Reproduced with permission from Ref. [181]

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the Na-metal anode and Na-BP-DME electrolyte during the Na plating/stripping process, no visible degradation of the electrolyte was detected [281]. An in situ DEMS seawater cell was built to test the stability of the Na-BP-DME anolyte toward sodium metal. Figure 3.122c shows that there was no evident evolution of organic-derived gases during the metal plating and cycling process; the SWB’s cycle performance with the Na-BP-DME electrolyte improved as a result of this. During the charging phase of NaCF3 SO3 -TEGDME (Fig. 3.122d), however, considerable H2 gas evolution was observed. The 1H NMR spectra of the desodiated samples were recorded as well (Fig. 3.122e). After cycling, Na-BP-DME showed no signs of decomposition or polymerization, indicating that it has outstanding electrochemical reversibility and stability [282, 283]. A saturated Na-1 M BP in DEGDME was also studied as an anolyte for SWBs [284]. It was demonstrated that using biphenyl as an anolyte solves the problem of sodium metal anodes by producing Na-BP naturally [177, 285, 286]. Na-BP can operate as a redox mediator, suppressing electrolyte decomposition at the sodium-metal anode (Fig. 3.123b). As a result, the Na-BP anolyte contributes to reversible sodium-cation storage while also allowing for uniform, low-overpotential sodium-metal deposition. The introduction of low-cost biphenyl allows for the suppression of continuous electrolyte degradation and the mitigation of sodium-dendrite formation (Fig. 3.123c). The Na-BP anolyte and the NaCF3 SO3 electrolyte (1 M NaCF3 SO3 in TEGDME) were analyzed using differential electrochemical mass spectroscopy (DEMS). The cells were repeatedly charged and discharged (Fig. 3.124a, b) while gas production in the negative compartment was monitored (Fig. 3.124c, d). Figure 3.124c shows that the cell with the Na-BP anolyte exhibited no significant gas evolution, whereas the cell with the NaCF3 SO3 electrolyte exhibited significant hydrogen evolution— especially during the second charge (Fig. 3.124d), which also explains the fairly noisy voltage profile (Fig. 3.124b). The reaction of trace water with the sodium metal anode and the possible electrolyte decomposition were due to the creation of hydrogen. The lack of hydrogen evolution in the Na-BP-containing cell showed that this electrolyte was more stable, and that the Na-BP molecules may have exhibited an H2 O-scavenging effect. Na-metal deposition was examined on a Cu foil using ex situ SEM. The NaBP anolyte was found to have an additional beneficial effect: substantially larger plated Na-metal grains (Fig. 3.125a) compared to those obtained in NaCF3 SO3 electrolyte (Fig. 3.125b). These imply that the nucleation barrier was lower in the NaOTF-based electrolyte, favoring the deposition of relatively small Na metal grains. Plating/stripping tests in symmetric Na-BP anolyte cells (Fig. 3.125c) revealed that the overpotential for Na metal deposition was higher in the case of the NaCF3 SO3 electrolyte. For the first 10 stripping/plating cycles, the NaCF3 SO3 electrolyte cell was around 270 mV, which was larger than the 220 mV obtained for the Na-BP-based cell (Fig. 3.125d). While the latter’s voltage response remained relatively constant for hundreds of hours, the cell with the NaCF3 SO3 electrolyte quickly faded (Fig. 3.125c, e), and the voltage response became very unstable and noisy, indicating the formation of an unstable SEI and severe Na dendrite formation.

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221

Fig. 3.123 Using sodium biphenyl (Na-BP) as an anolyte in seawater batteries. Reproduced with permission from Ref. [284]

3.4 Solid Electrolytes 3.4.1 Role and Requirement of Solid Electrolyte for SWBs 3.4.1.1

Role of Solid Electrolyte in SWBs

A solid ceramic electrolyte is a key component of a SWB system as the anode storing Na+ can be insulated from direct exposure to seawater. The viability of a rechargeable battery system employing seawater as an electrode material has been established. The solid electrolyte solely transports Na ions between the anode and cathode of the SWB during the charge and discharge processes. The solid electrolyte separates the anode and cathode compartments, preventing the diffusion of ions other than Na+ between them. In this regard, the solid electrolyte in a SWB works as an electrolyte as well as a separator.

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Fig. 3.124 The voltage profiles of the differential electrochemical mass spectroscopy (DEMS) cell with a Na-BP and b NaCF3 SO3 . Gas evolved during the cycle in cells containing c Na-BP and d NaCF3 SO3 . Reproduced with permission from Ref. [284]

3.4.1.2

Required Property

Working electrodes of rechargeable SWB are separated by Na ion conducting solid electrolyte. The solid electrolyte used for the fabrication of SWBs must exhibit high ionic (Na+ ) conductivity under low temperature conditions. The system should also exhibit high ionic (Na+ ) selectivity (it should not allow the passage of any ions other than Na+ ), electrical insulation properties (to separate the anode and cathode), high chemical stability in the presence of seawater and nonaqueous electrolytes, electrochemical stability in a wide electrochemical window, and high mechanical/physical strength (to secure the separation of nonaqueous electrolytes). These properties have been illustrated in Fig. 3.126. Na ion conducting solid electrolytes are classified as inorganic, polymer, and polymer–ceramic composites. Inorganic electrolytes have higher density and mechanical strength than polymer-based electrolytes and can be used to physically separate nonaqueous electrolytes from seawater. However, not all types of inorganic solid electrolytes can be used to fabricate rechargeable SWBs.

3.4 Solid Electrolytes

223

Fig. 3.125 a, b SEM images of plated Na metal with different electrolytes. c Voltage profile of symmetric cells with Na-BP (black) and NaCF3 SO3 (red). Magnification of the d initial and e following stripping/plating of (e). Reproduced with permission from Ref. [284]

3.4.2 Solid Electrolyte Materials 3.4.2.1

Categories of Solid Electrolyte Materials

Various solid materials exhibiting excellent Na+ conducting properties have been reported. The ionic conductivities of a few such materials are summarized in Fig. 3.127. These materials, in general, exhibit high ionic conductivities at high temperatures (>80 °C). Some of the materials conduct ions (ionic conductivity: 1 × 10–5 –1 × 10–3 S/cm) at room temperature (25 °C) and could be used as solid electrolytes for the fabrication of SWBs.

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Fig. 3.126 Required properties of solid electrolytes for rechargeable SWB. Reproduced with permission of Ref. [287]

Fig. 3.127 Ionic conductivities of representative sodium based solid state electrolytes. Reproduced with permission of Ref. [288]

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225

Fig. 3.128 Representative polymer, sulfide, and oxide solid electrolytes and their general properties. Reproduced with permission of Ref. [287]

Solid electrolyte materials can be categorized as organic (polymer) and inorganic (sulfide and oxide), shown in Fig. 3.128. Polymer-based solid electrolytes consist of a polymer matrix and sodium salt. The processes of interchain diffusion and shift are responsible for the transport of Na+ in a polymer-based solid electrolyte. Often, low ionic conductivities are exhibited by polymer-based solid electrolytes at room temperature (range: 10–6 –10−8 S/cm). The materials are also characterized by low density and low Na+ selectivity. This results in the crossover of organic liquid electrolytes and the saltwater in the SWB system. On the other hand, inorganic solid electrolytes (oxide and sulfide electrolytes) exhibit high density, Na+ selectivity, and mechanical strength; adequate to separate non aqueous electrolytes against seawater.

3.4.2.2

Sulfide-Based Solid Electrolytes

Sulfide-based solid electrolytes are regarded as viable solid electrolytes for sodium batteries as they exhibit mechanical ductility and high ionic conductivity (>1 ×

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Fig. 3.129 Example of sulfide solid electrolyte pellet (Li10 SnP2 S12 ). Reproduced with permission of Ref. [289]

10–3 S/cm) at ambient temperature. The sulfide solid electrolyte consists of a sulfide framework with Na+ ions interspersed throughout the matrix. The Na+ ions move by ion hopping from one Na+ site to the next by overcoming the energy barrier, which is controlled by the conduction channel (sulfide framework). The properties of high ionic conductivity and mechanical ductility of sulfide solid electrolytes can be attributed to the fact that sulfur is characterized by a large ionic radius and forms weak electrostatic interactions with Na+ ions. Na3 PS4 and Na3 SbS4 are two representative sulfide electrolytes. At ambient temperature, both materials exhibit high ionic conductivities that are >1 × 10–3 S/cm. These sulfide solid electrolytes, in contrast, exhibit low chemical stability when exposed to humid conditions and air. Na3 PS4 chemically reacts with H2 O and O2 . The reaction results in the production of H2 S, a poisonous gas. When Na3 SbS4 combines with water, it forms the Na3 SbS4 ·H2 O phase, which exhibits low ionic conductivity. Another example of a sulfide solid electrolyte is Li10 SnP2 S12 , shown in Fig. 3.129. The low chemical stability of the sulfide materials (in water) limits their practical applications as solid electrolytes during the fabrication of SWB systems.

3.4.2.3

Oxide-Based Solid Electrolytes

Oxide solid electrolytes were first reported in the 1960s. Since then, they have been extensively studied and used in batteries. The oxide solid electrolytes are characterized by desirable qualities such as high ionic conductivity (1 × 10–3 S/cm) at ambient temperature, broad electrochemical windows, and high mechanical strength. The oxide solid electrolytes are chemically more stable than the sulfide solid electrolytes. An oxide solid electrolyte consists of an oxide framework with Na+ ions interspersed throughout the framework. The Na+ conduction pathway, in this case, involves the hopping mechanism like that observed for the sulfide solid electrolytes. The bonding between Na+ and oxygen is strong enough to result in high mechanical

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227

strength and brittle characteristics. This can be attributed to the fact that oxygen is more electronegative than sulfur. Sodium beta alumina and NASICON (Na super ion conductors) are two examples of oxide solid electrolytes. At room temperature, both materials exhibit high ionic conductivity (1 × 10–3 S/cm) and a wide electrochemical window (0–5 V vs. Na/Na+ ). More information on these materials is presented in the subsequent sections.

3.4.3 Sodium Beta Alumina 3.4.3.1

Introduction to Sodium Beta Alumina

Researchers working with the Ford Motor Company first reported sodium beta alumina in 1967. The discovery sparked a flurry of research on sodium beta alumina because of its high ionic conductivity and electrical resistance. Moreover, researchers investigated the applicability of sodium beta alumina as a solid electrolyte for energy storage. Sodium beta alumina presents a layered structure (Fig. 3.130). A spinel block composed of four layers of O, with Al ions in the octahedral and tetrahedral interstices forming one layer. The oxygen and sodium ions are loosely packed in the other layer, which is often referred to as the conduction plane. The Na+ ions are transported via a 2D channel, which forms the conduction plane. The crystal structures of sodium beta alumina have been identified as Na-β-Al2 O3 and Na-β -Al2 O3 . The chemical stoichiometry and structure of the materials are distinct. Na-β-Al2 O3 possesses a hexagonal structure and is composed of Na2 O·(8– 11)Al2 O3 . Na-β -Al2 O3 presents a rhombohedral structure and is composed of Na2 O·(5– 7)Al2 O3 . Na-β -Al2 O3 exhibits a higher Na+ ion conductivity (2.0 × 10–3 S/cm) at room temperature than Na-β-Al2 O3 . This can be attributed to the presence of a large number of Na+ ions. The relative ratio of the two phases (the Na-β -Al2 O3 to Na-β-Al2 O3 ratio) in the sodium-beta-alumina-based ceramic materials influences their various properties including the ionic conductivity and mechanical properties. The sodium beta alumina ceramic is commercially available. The SEM (Scanning Electron Microscope) images representing the microstructures of the sodium beta alumina obtained from Ionotec. Co., Ltd. are shown in Fig. 3.131a [14]. This morphology is common for typical ceramics and sodium beta alumina, and is attributed to grain formation that occurs during the process of solid–liquid reaction sintering. Complete spinel blocks appear to spread as ledges along the basal planes during the development of the material. The observed lattice parameters for the β -Al2 O3 (reference PDF # 00–035-0438) are a = 5.61 Å and b = 33.484 Å which are close to the values reported in the literature [14]. The X-ray diffraction (XRD) patterns of β -Al2 O3 shown in Fig. 3.131b correspond to a rhombohedral structure (Fig. 3.131b).

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Fig. 3.130 a Digital image of a sintered sodium beta alumina ceramic and b crystal structure of sodium beta alumina. Reproduced with permission of Ref. [290, 291]

3.4.3.2

Ionic Conductivity of Sodium Beta Alumina

Sodium beta alumina can be used as a solid electrolyte as it exhibits high ionic conductivity at room temperature. The Na+ ion conductivity typically recorded for sodium beta alumina is 1.0 × 10–3 S/cm. The ion-conducting ability is influenced by the microstructure, composition, and β/β -phase ratio. Ionic conductivity of the commercial Na beta alumina has been evaluated and shown in Fig. 3.132. Its ionic conductivity is found to be 1.84 × 10–4 S/cm at room temperature.

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229

Fig. 3.131 a SEM image of Na-β -Al2 O3 . b X-ray diffraction (XRD) patterns of sintered Na-β Al2 O3 . Reproduced with permission of Ref. [14]

Fig. 3.132 Electrochemical impedance spectroscopy results of pristine Na-β -Al2 O3 . Reproduced with permission of Ref. [14]

The ionic conductivities recorded for sodium beta alumina are presented in Table 3.7. Single crystal samples are observed to be more conductive than polycrystalline samples. This could be attributed to the fact that the β- and β -Al2 O3 crystals are not characterized by the grain boundary effect and anisotropic sodium ion conduction properties.

3.4.3.3

Mechanical Property of Sodium Beta Alumina

Table 3.8 summarizes the mechanical parameters defining sodium beta alumina. The fracture strength of normal sodium beta alumina is 120–210 MN/m2 . The microstructure (and properties such as porosity, grain size, and phase ratio) exerts a significant

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Table 3.7 Ionic conductivity (σ) and activation energy (Ea ) of sodium beta alumina [292] σ (S/cm)

Type

Ea (eV)

25 °C

300 °C

0.036

–

0.024

–

0.16

0.035

0.21

0.13

0.03

–

0.16

0.025

–

0.15

0.03

0.27

0.14

Polycrystalline β-Al2 O3

0.0012

–

0.27 (25–200 °C)

Single crystal β -Al2 O3

0.04

–

0.22 (25–250 °C)

0.1

–

0.20 (−80–150 °C)

0.22–0.35

0.15–0.26

0.21

0.24 (285–330 °C)

0.36

0.18 (285–330 °C)

Single crystal β-Al2 O3

–

0.014 Polycrystalline β -Al2 O3

0.31 (−80–150 °C)

Table 3.8 Mechanical properties of sodium beta alumina reported in literature Property

Fabrication conditions Isostatically pressed at 60,000 psi and sintered at 1585 °C for 30 min

Isostatically pressed at 60,000 psi and sintered at 1585 °C for 20 min

Isostatically pressed and sintered at 1585 °C for 10 min

Isostatically pressed and sintered at 1605 °C for 20 min

Fracture strength (MN/m2 )

210

120

140–170

120

Critical stress intensity factor (MN/m3/2 )

2.3

2.8

Young’s 1.4 × 1011 2 modulus (N m )

1.2 × 1011

References

[293]

[293]

2.6–3.8

[294]

[295]

impact on the mechanical strength of sodium beta alumina. The addition of ZrO2 to the microstructure of sodium beta alumina enhances the fracture strength, and the maximum value recorded is 350 MPa. The strength is comparable to those of common ceramics such as ZrO2 and Al2 O3 .

3.4 Solid Electrolytes

3.4.3.4

231

Chemical Stability of Sodium Beta Alumina

Sodium beta alumina is moisture sensitive and occurrence of rapid exothermic occlusion of water has been reported on the surface micropores of β -Al2 O3 . The process results in pore saturation in less than 1 h. Following the process of water occlusion, H3 O+ ions slowly diffuse into the β -Al2 O3 lattice. Under these conditions, ion exchange with sodium occurs. Flor et al. used thermogravimetric and X-ray techniques to confirm that β -Al2 O3 is unstable in the presence of moisture. According to Flor et al., rapid penetration occurs along the conduction plane for the first few microns. Changes in the a and c lattice parameters of β -Al2 O3 could be attributed to water absorption.

3.4.3.5

Application of Sodium Beta Alumina

Sodium beta alumina has been successfully used for the fabrication of energy-storage systems such as Na-S and ZEBRA batteries. Its chemical formula is NaAl(5.33–11) O17 . Figure 3.133 shows a schematic of the structure of the Na-S battery. Tubular sodium beta alumina (Fig. 3.133b) is located between sulfur and the molten Na metal anode (Fig. 3.133a). Single Na-S cells (by SICCAS) are presented in Fig. 3.133c. Due to Na+ conducting characteristics, it is believed that Na beta alumina can be a promising option for solid electrolytes in SWBs. Sodium beta alumina was used as a solid electrolyte during the development of SWBs. SWBs fabricated using Na-β Al2 O3 as the solid electrolyte were produced, and their electrochemical performances were studied. The charge/discharge voltage curves generated for a SWB system with hard carbon as the anode and Na-β -Al2 O3 as the solid electrolyte are shown in Fig. 3.134. The voltage plateau can be observed at 3.0 V. This region is followed by a smooth slope, which extends till 3.5 V. The first charge capacity was 220 mAh/g, and the subsequent discharge capacity was 3.00 g/cm3 , and the ionic conductivity is 3.0 × 10–4 S/cm. The size and

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Fig. 3.139 Crystal structures of NASICON and corresponding sodium transfer pathways: rhombohedral (left) and monoclinic (right). Reproduced with permission of Ref. [288]

shape of the NASICON ceramic is relatively limited compared to the sodium beta alumina. Commercially accessible sodium beta alumina has various shapes and sizes [296]. These materials can be tubular, circular, or square in shape, with dimensions ranging from 1 to 300 cm2 . Most of the fabricated NASICON materials are circular in shape, and their diameter is less than 20 pi. Physical properties of the NASICON ceramic purchased from 4TOONE Co. Ltd. have been characterized. The SEM images reflecting the microstructures of the NASICON-ceramic-based electrolyte are shown in Fig. 3.140a. The analysis of the microstructures exhibited a high density. The XRD patterns recorded for the prepared NASICON (reference PDF # 01–084-1190) are presented in Fig. 3.140b. The peaks agreed well with those published in the literature.

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237

Fig. 3.140 a SEM image and b XRD patterns of NASICON ceramic. Reproduced with permission of Ref. [14]

3.4.4.2

Ionic Conductivity of NASICON

The ion conducting property of NASICON makes it suitable for use as a solid electrolyte. High ionic conductivity of 1 × 10–5 –1 × 10–3 S/cm at room temperature has been recorded since its discovery in 1976 [302–308]. Electrochemical impedance spectroscopy (EIS) is used to determine the ionic conductivity of the NASICON ceramics. The ionic conductivity measured with EIS can be classified into two types: bulk and total ionic conductivity. The bulk conductivity of the NASICON ceramic is determined by analyzing the NASICON phase (grain), while the overall ionic conductivity is determined by analyzing the contributions of the amorphous, ZrO2 and NASICON phases [298]. Among the typical NASICON compositions (Na1+x Zr2 Six P3-x O12 ), Na3 Zr2 Si2 PO12 (x = 2) is reported with high ionic conductivity of 1 × 10–4 –1 × 10–3 S/cm. The ionic conductivity of the NASICON ceramic obtained from 4TOONE Co. Ltd. is shown in Fig. 3.141, and the ionic conductivity is found to be 1.03 × 10–3 S/cm at room temperature. Table 3.9 summarizes the ionic conductivities of the NASICON ceramics reported in literature. At room temperature, the ionic conductivity of this material is comparable to that recorded for the sodium beta alumina ceramic (1 × 10–3 S/cm). Although efforts have been devoted towards increasing the ionic conductivity of the NASICON materials by optimizing the synthesis method and composition, in most cases, the materials exhibits a conductivity of 1.0 × 10–3 S/cm.

3.4.4.3

Mechanical Property of NASICON

The NASICON ceramic is fragile as it is an oxide. The bending strength testing method (also known as the fracture strength testing method) is commonly used to assess the mechanical strength of the material. The bending strength of the material is influenced by parameters such as the microstructure, flaws, and chemical composition [309]. The bending strength of common ceramic materials such as

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Fig. 3.141 Impedance spectra of pristine NASICON. Reproduced with permission of Ref. [14]

Table 3.9 Compositions of common NASICON ceramics and their ionic conductivity (σ) Type

Composition Method

Hong-type x = 1.6 NASICON x = 1.8 (Na1+x Zr2 Six x = 2.0 P3-x O12 ) (Na3 Zr2 Si2 PO12 )

σbulk [S/cm] σtotal [S/cm] Year References

Solid-state reaction

7.1 × 10–4

2.0 × 10–4

3.6 ×

1.6 ×

Solid-state reaction

6.8 × 10–4

Sol–gel

~7 ×

Spark-plasma sintering

-

Pechini

10–4

1978 [303]

10–4

2.5 × 10–4

3.3 × 10–4a

1992 [304] ~ 1.2 × 10–4

1999 [305]

1.8 × 10–3

2004 [306]

~1.9 × 10–3 8.8 × 10–4

2017 [302]

10–4

Solid ~1.5 × 10–3 1.0 × 10–3 assisted-solid-state reaction

a

x = 2.0 + 4.8 wt.% Na3 BO3

Liquid-phase sintering

–

~ 1 × 10–3

2018 [307]

x = 2.0 + 5 wt.% Na2 SiO3

Solid-state reaction

~1.6 × 10–3 1.5 × 10–3

2019 [308]

bulk and total conductivity is not separated ~ data read from published graph

Al2 O3 and ZrO2 has been extensively studied. The reports on bending strength of the NASICON material, in contrast, are relatively uncommon. The bending strength of the NASICON ceramic is lower than that of other conventional ceramics. As shown in Fig. 3.142, Al2 O3 is characterized by a bending strength of >300 MPa

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Fig. 3.142 Variation in the bending strength of ZrO2 , Al2 O3, and NASICON ceramics with density

[310–313], and ZrO2 is characterized by a bending strength of >400 MPa [314]. The bending strength recorded for the NASICON material is 50–200 MPa [315–317]. The origin of low bending strength has rarely been discussed. The elements influencing the strength of the material and the ways to improve mechanical strength must be studied as NASICON exhibits a unique microstructure and is characterized by a unique composition.

3.4.4.4

Chemical Stability of NASICON in Seawater

For the fabrication of SWBs, a solid electrolyte that is stable in a seawater environment should be used. The electrolyte must transfer sodium ions from the seawater to the anode and simultaneously act as a separator between the seawater and the anode compartment. In SWB systems, it is critical to keep the solid electrolyte intact and retain its features. The electrochemical properties should be retained, and structural stability should be maintained. The facile migration of sodium ions and a long energy storage system (ESS) lifetime are ensured by the robust electrochemical and structural features of the SWBs. This helps prevent short circuits. The chemical stability of the NASICON in the seawater has been investigated by immersing it in the seawater [14]. The XRD patterns recorded for the NASICONbased solid electrolytes recorded. The 2θ degree values were plotted as a function of immersion time in seawater (Fig. 3.143a). The analysis of the XRD peaks recorded for NASICON immersed in seawater at room temperature for 60 days reveals the

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Fig. 3.143 a X-ray diffraction (XRD) patterns and b impedance spectra of NASICON immersed in seawater for 60 days. Reproduced with permission of Ref. [14]

absence of impurity phases. The EIS spectra shown in Fig. 3.143b confirms the chemical stability of NASICON ceramic in seawater. The long-term chemical stability of NASICON ceramic in the seawater and distilled water (DI water) have also been compared (Fig. 3.144) [301]. The Nyquist plots of pristine NASICON ceramic and after 90 days of immersion in seawater are shown in the Fig. 3.144a. For the case of the seawater immersed NASICON ceramic, the Nyquist plot exhibited little change. The calculated total ionic conductivity is 2.7 × 10–4 S/cm which is close to that of the pristine sample (2.8 × 10–4 S/cm). On the other hand, the Nyquist plot of the DI water immersed sample exhibited noticeable change. Especially the diameter of the semi-circle was significantly increased. The calculated total ionic conductivity of the DI water immersed sample decreased to 1 ×

Fig. 3.144 a Impedance spectra and b ionic conductivity of NASICON immersed in DI water and seawater. Reproduced with permission of Ref. [301]

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241

10–5 S/cm (Fig. 3.144a) [301]. The results of the longer-term (~150 days) immersion test are also shown in Fig. 3.144b. Following confirmation that H3 O+ (hydronium) NASICON can be formed in DI water at room temperature, researchers conducted the 3D time-of-flight secondaryion mass spectrometry (TOF–SIMS) to investigate the extent of distribution of H3 O+ in NASICON immersed for 6 months in DI water and seawater (Fig. 3.145). H3 O+ ions were chosen as the focus of this investigation because they can occupy the sites of the sodium ions present in the NASICON and result in the production of hydronium NASICON. The exchanged H3 O+ ions generally remained on the surface. Relatively, substantial H3 O+ ions were present in the DI-water-immersed NASICON materials than seawater immersed one, and these penetrated the bulk to reach a depth of approximately 500 nm. The origin of the intense hydronium exchange processes that occur in DI water has been investigated using density functional theory (DFT). The DFT has been also used to understand why hydronium NASICON is generated primarily near the surface and penetration into the bulk material is not realized. As dissolved ions are present in the systems, various dielectric constants were set for DI water and seawater to predict how NASICON responds in aqueous solutions. Though the energy of insertion of

Fig. 3.145 Schematic of surface reaction of NASICON in a marine environment and b DI water. c, d Corresponding change in surface concentration of H3 O+ ions investigated through TOF–SIMS 3D mapping. Reproduced with permission of Ref. [301]

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Fig. 3.146 a Energy barrier of Na extraction and H3 O+ insertion of the NASICON structure in DI water and Seawater. b Calculation results of Na extraction followed with composition change and the specific crystal plane of NASICON in the two solutions. c Energy barrier comparison of H3 O+ ion insertion from surface to bulk structure of NASICON. Reproduced with permission of Ref. [301]

hydronium ions was comparable for both the solutions, the fact that low energy was required for the extraction of sodium ions from the NASICON materials immersed in DI water suggested that the development of sodium vacancy in the NASICON structure was critical to produce hydronium NASICON (Fig. 3.146a). Figure 3.146b presents the additional calculations carried out for the generation of sodium vacancies in the NASICON structure. The results help to predict the various structural states and directions realized. When the sodium sites were entirely occupied by the sodium ions (the bad condition), sodium vacancies were easily formed. The energy recorded in DI water was significantly lower than that in a marine environment along the (010) plane. These findings suggest that as the number of sodium ions escaping the structure increases (near the rich condition), the production of sodium vacancies becomes more difficult. The formed vacancies exhibited preferred features in particular planes. The generation of hydronium NASICON thus resulted in planepreferential features. These features were not observed for the sodium vacancies generated under other conditions (Fig. 3.146c). The minimum formation energy was recorded along the (010) plane, and this indicated that the hydronium NASICON formed preferentially in a specific direction and on the surface layer. This result

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243

demonstrates that the H3 O+ occupied the Na sites present near the surface layer. It was difficult to penetrate the inner layers of the NASICON materials. A NASICON pellet immersion test has also been carried out to study the behaviors of the materials under conditions of minor phase dissolution reactions in DI water and seawater at room temperature. Inductively coupled plasma-optical emission spectroscopy (ICP-OES) has been used to identify the types of minor phases dissolved in aqueous solutions (Fig. 3.147a, b). The results reveal that the Zr ions were not effectively dissolved in the two solutions. Nevertheless, the P and Si units could easily dissolve as the minor phases exhibited higher solubility than ZrO2 . After 15 d, the amount of P dissolved in DI water was found to be approximately 5.5 mg/L. This value is 100 times higher than that recorded in a marine environment. The concentrations of dissolved Si ions in DI water are approximately 10 times higher than those recorded in a seawater environment. Minor phases containing phosphorus

Fig. 3.147 ICP-OES results for minor phase dissolution of NASICON pellet in a DI water and b seawater. c pH measurements of DI water and seawater after NASICON pellet immersion. d Comparison of solubilities of minor phases of NASICON powder. Reproduced with permission of Ref. [301]

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were formed; these phases dissolved more easily in the NASICON pellets than the minor phases containing Si. The pH (acidity) of the solution increased to approximately 10 following the process of pellet immersion in DI water (Fig. 3.147c). Similar observations were made in the case of NASICON powder immersion (Fig. 3.147d). However, following pellet immersion, the pH of the seawater was approximately 8. The pH differences in both solutions could be attributed to the minor phase dissolution process in the presence of the NASICON pellets because the minor phases (i.e., Na3 PO4 and Na2 SiO3 ) could increase the pH of the solutions (Fig. 3.147c). When small components were dissolved in water, NaOH was produced, which increased the pH of the solutions. A consistent change in pH was realized in both solutions by dissolving Na3 PO4 and Na3 SiO3 in the solutions. The pH increased in the range of 12–13 in DI water following the dissolution of the pure minor phase chemicals (until a saturation concentration was reached). The pH was