Superhalogens: Properties and Applications (SpringerBriefs in Molecular Science) 303137570X, 9783031375705

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SpringerBriefs in Molecular Science Ambrish Kumar Srivastava

Superhalogens Properties and Applications

SpringerBriefs in Molecular Science

SpringerBriefs in Molecular Science present concise summaries of cutting-edge research and practical applications across a wide spectrum of fields centered around chemistry. Featuring compact volumes of 50 to 125 pages, the series covers a range of content from professional to academic. Typical topics might include: • A timely report of state-of-the-art analytical techniques • A bridge between new research results, as published in journal articles, and a contextual literature review • A snapshot of a hot or emerging topic • An in-depth case study • A presentation of core concepts that students must understand in order to make independent contributions Briefs allow authors to present their ideas and readers to absorb them with minimal time investment. Briefs will be published as part of Springer’s eBook collection, with millions of users worldwide. In addition, Briefs will be available for individual print and electronic purchase. Briefs are characterized by fast, global electronic dissemination, standard publishing contracts, easy-to-use manuscript preparation and formatting guidelines, and expedited production schedules. Both solicited and unsolicited manuscripts are considered for publication in this series.

Ambrish Kumar Srivastava

Superhalogens Properties and Applications

Ambrish Kumar Srivastava Department of Physics Deen Dayal Upadhyaya Gorakhpur University Gorakhpur, Uttar Pradesh, India

ISSN 2191-5407 ISSN 2191-5415 (electronic) SpringerBriefs in Molecular Science ISBN 978-3-031-37570-5 ISBN 978-3-031-37571-2 (eBook) https://doi.org/10.1007/978-3-031-37571-2 © The Editor(s) (if applicable) and The Author(s), under exclusive license to Springer Nature Switzerland AG 2023 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 Switzerland AG The registered company address is: Gewerbestrasse 11, 6330 Cham, Switzerland

Let noble thoughts come to me from all directions… —Rig Veda

Preface

Halogens possess a very rich chemistry due to their high electronegativity, high electron affinity, and consequently, high reactivity. It was always interesting to explore whether any system can have more electron affinity than halogen. This led to the conceptualization of “superhalogen” in 1981 by G. L. Gutsev and A. I. Boldyrev, two scientists from Russia (then, USSR). After the experimental verification of this concept in 1999, the field became popular and also populated by works of other researchers from the USA, Poland, and China. This topic appeared very fascinating to me when I was admitted to the Ph.D. program in 2012. I immediately started working in this field, which led to my first publication in late 2013. There was a cluster of opportunities in the field. In 2016, I started communicating with Dr. Gutsev to work together on some problems. However, it could not be possible, partially due to a problem in the computational cluster and partially due to the submission of my Ph.D. thesis in the same year. Till 2020, I have published more than half a century of papers on various aspects of superhalogens. In 2021, I was invited to contribute an article for a special issue dedicated to Prof. Boldyrev. In 2022, I got the opportunity to edit an article collection on this topic having Dr. Gutsev as a co-editor to which Prof. Boldyrev’s group contributed as well. This field gave me much more than I expected. It was my turn to pay back to this field. I missed the lack of concise literature with the latest trends in the research in the field of superhalogen in the last two decades. This prompted me to plan this book in the form of the SpringerBriefs. This concise book is divided into seven chapters. Chapter 1 offers a general introduction to the subject including the basic idea and recent trends in the design of superhalogens. Chapter 2 exclusively discusses the design of transition metal core-based superhalogens. The subsequent chapters deal with the applications of superhalogens as strong oxidizers

vii

viii

Preface

and in superacids and electrolytic salts. These are followed by a chapter on some miscellaneous applications explored only recently. The last chapter concludes the book with some future perspectives. I believe that the contents of the book will not only provide readers with an overview of this relatively new field but also motivate other researchers to explore this field. A warm welcome to the world of superhalogens! Gorakhpur, India May 2023

Ambrish Kumar Srivastava

Acknowledgements I am thankful to the Council of Scientific and Industrial Research (CSIR), India, for Junior and Senior Research Fellowships (2012–2016), Science and Engineering Research Board (SERB), India, for the National Post-doctoral Fellowship (2017), and the University Grants Commission (UGC), India, for Startup Grant (2019–2021). These funds enabled me to carry out my research works tirelessly, which form a major part of this book.

Contents

1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.1 Background . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.2 Theoretical Prediction of Superhalogens . . . . . . . . . . . . . . . . . . . . . . . 1.3 Adiabatic Electron Affinity Versus Vertical Detachment Energy . . . 1.4 Experimental Confirmation of Superhalogens . . . . . . . . . . . . . . . . . . . 1.5 Quest for New Ligands . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.6 Towards Maximization of VDEs . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.7 Electron Counting Rules . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.8 Non-Coordinating Anions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 1.9 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

1 1 2 3 4 5 8 10 10 11 12

2 Transition Metal Fluorides and Oxides as Superhalogens . . . . . . . . . . . 2.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 2.2 Transition Metal with Fluorine Ligands . . . . . . . . . . . . . . . . . . . . . . . . 2.3 Transition Metal with Oxygen Ligands . . . . . . . . . . . . . . . . . . . . . . . . 2.4 Transition Metal with Other Ligands . . . . . . . . . . . . . . . . . . . . . . . . . . 2.5 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

13 13 14 16 19 21 21

3 Superhalogens as Strong Oxidizers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.2 Oxidation of Water and Its Clusters . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.3 Oxidation of Benzene . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.4 Oxidation of Fullerene . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.5 Oxidation of Carbon Dioxide . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.6 Oxidation of Transition Metal Oxides . . . . . . . . . . . . . . . . . . . . . . . . . 3.7 Oxidation of Heterocycles . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.8 Oxidation of Nitric Oxide . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 3.9 Formation of Noble Gas Compounds . . . . . . . . . . . . . . . . . . . . . . . . . . 3.10 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

23 23 24 25 25 26 27 27 28 29 29 30 ix

x

Contents

4 Superhalogens in the Design of Superacids . . . . . . . . . . . . . . . . . . . . . . . . . 4.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4.2 Acidity of HAln F3n+1 Species . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4.3 Protonated Superhalogens as Superacids . . . . . . . . . . . . . . . . . . . . . . . 4.4 Origin of (Super)acidity . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4.5 Path for Stronger Superacids . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 4.6 Other Superhalogen-Based Superacids . . . . . . . . . . . . . . . . . . . . . . . . 4.7 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

33 33 33 34 36 37 40 42 43

5 Superhalogens in the Design of Electrolytic Salts . . . . . . . . . . . . . . . . . . . 5.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5.2 Anions in Typical Electrolytic Salts . . . . . . . . . . . . . . . . . . . . . . . . . . . 5.3 New Electrolytic Salts Based on Halogen-Free Superhalogens . . . . 5.4 New Electrolytic Salts Based on Superhalogen Dianions . . . . . . . . . 5.5 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

45 45 46 46 51 52 54

6 Miscellaneous Applications of Superhalogens . . . . . . . . . . . . . . . . . . . . . . 6.1 Introduction . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 6.2 Organic Superconductors . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 6.3 Ionic Liquids . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 6.4 Liquid Crystals . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 6.5 Nonlinear Optical Materials . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 6.6 Summary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

55 55 55 56 58 61 62 63

7 Conclusion and Future Perspectives . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 65 References . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 68 Index . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 69

Chapter 1

Introduction

Abstract Superhalogens are species explored more than four decades earlier in 1981. Since their confirmation in 1999, such species were paid much attention. In this Chapter, an introduction to the field of superhalogens has been offered. Starting from their background, theoretical prediction, and experimental verification, the recent progress in the design of superhalogens has been discussed. Various ligands used in the design of superhalogens and the strategy to design polynuclear superhalogens have been revealed. Some related topics such as electron counting rules and noncoordinating anions have been also mentioned. Keywords Superhalogens · Electron affinity · Vertical detachment energy · Photoelectron spectroscopy · Electron-counting rules · Non-coordinating anions

1.1 Background The energy released upon the addition of an electron to any neutral atom or molecular system defines its electron affinity (EA). The upper limit of EA of elemental atom is 3.63 eV [1] for chlorine (Cl) which belongs to a special group in the periodic table, known as halogen. Due to the virtue of high EA, halogens are strong oxidizers. Stable molecules have fairly low (even negative) EA and very high ionization energy (IE). For instance, the EA and IE of O2 are 0.45 and 12 eV, respectively [2]. In 1962, Bartlett and Lohmann [3] synthesized an unusual ionic compound O2 + PtF6 − which showed that O2 can be oxidized by PtF6 . The EA of PtF6 was estimated to be 8.6–9.3 eV by Compton and coworkers in 1978 [4]. Evidently, the oxidation of O2 is a consequence of the high EA of PtF6 .

© The Author(s), under exclusive license to Springer Nature Switzerland AG 2023 A. K. Srivastava, Superhalogens, SpringerBriefs in Molecular Science, https://doi.org/10.1007/978-3-031-37571-2_1

1

2

1 Introduction

1.2 Theoretical Prediction of Superhalogens In 1981, Gutsev and Boldyrev [5] reported some complex anions belonging to the formula of MXk+1 − , where M is a central atom with valence k and X is some electronegative ligand (F, O, etc.) as shown in Fig. 1.1. They used the Hartree–Fock (HF) method to obtain (first) ionization energy, commonly known as the vertical detachment energy (VDE), of these anions, also given in Fig. 1.1. Assuming this VDE to be approximately equal to the EA of corresponding neutral (radical), they coined the term ‘superhalogen’ for MXk+1 species. Thus, superhalogens are species whose EA exceeds that of halogen or whose anions have larger VDE than that of halide. According to the above formula, MX2 (M = Li, Na, K; X = F, Cl, Br) should be superhalogens. This is indeed the case as the VDEs of LiF2 − , LiCl2 − , NaF2 − , and NaCl2 − are calculated to be 6.51, 5.88, 6.18, and 5.77 eV, respectively [6]. Now the question is, why is the EA of these neutral systems or the VDE of their anions so high? Let us take LiF2 as an example. LiF is a stable molecule in which the F atom takes an electron from the Li atom to complete its octet. Now another F atom attached to the LiF molecule requires an extra electron to attain its stability according to the octet rule [7]. This results in the high EA of LiF2 , making it a superhalogen. Equivalently, LiF2 − anion (with an extra electron) possesses high VDE due to its stability as per the octet rule.

Fig. 1.1 Typical superhalogens introduced by Gutsev and Boldyrev along with their EA values from Ref. [5]

1.3 Adiabatic Electron Affinity Versus Vertical Detachment Energy

3

1.3 Adiabatic Electron Affinity Versus Vertical Detachment Energy Since EA and VDE are deciding factors for the superhalogen behavior of systems, it would be obvious to discuss briefly, how these are calculated theoretically and measured experimentally. The EA or more precisely adiabatic EA (AEA) are obtained by the difference of total electronic (including zero-point) energies of neutral structures and anions, both at their optimal geometries. These optimal geometries and energies are obtained by appropriate electronic structure methods such as ab initio methods or density functional theory (DFT). However, the VDE is obtained by the difference in total energies of neutral and anions, both at the optimal geometries of anions. Thus, the VDE does not take the structural relaxation into account due to the removal of extra electron from anions as shown schematically in Fig. 1.2. As a matter of fact, VDE is always larger than AEA. The importance of VDE is due to the fact that the anions have stable and closedshell structures, which can be realized experimentally. The experimental VDEs are generally obtained by photoelectron spectroscopy (PES). In this technique, a photon of appropriate energy is used to detach an electron from anions based on the photoelectric effect. The binding energy of an electron is measured by taking the difference of the kinetic energy of the electron (EKE) from that of photon energy (hv), which is equal to the VDE. This is also shown schematically in Fig. 1.2. Although the AEA can not be measured directly from the experiments, it can be estimated using the measured VDE and reorganization energy. Fig. 1.2 Schematic diagram of potential energy surface for the estimation of VDE and AEA

4

1 Introduction

1.4 Experimental Confirmation of Superhalogens One of the apparent challenges in the PES of superhalogens was the large VDE of these anions. In 1999, Wang et al. [8] developed an experimental set-up by combining a PES device with an electrospray ionization (ESI) source and performed photodetachment at 193 nm (6.42 eV) using an ArF excimer laser. Subsequently, they recorded photoelectron spectra of MX2 − anions [9], which were calibrated and converted to kinetic energy spectra. The VDEs were obtained through binding energy spectra by subtracting kinetic energy from photon energy (6.42 eV) as shown in Fig. 1.3. The spectra were interpreted and assigned theoretically. Their measured VDEs of 5.94 and 5.86 eV of LiCl2 − and NaCl2 − , respectively were in good agreement with their previously reported values. This not only confirmed the existence of superhalogens but also validated the reliability of theoretical results, which are to be discussed in the next chapter onwards. However, they could not succeed in measuring the VDE of MF2 − (M = Li, Na, K)

Fig. 1.3 Photoelectron spectra of MXk+1 − along with their VDEs from Ref. [9] with permission of American Institute of Physics, copyright 1999

1.5 Quest for New Ligands

5

either due to their too-high VDEs or due to weak mass signals. Note that the VDE of LiF2 − (6.51 eV) is larger than that of photon energy (6.42 eV).

1.5 Quest for New Ligands In MXk+1 − superhalogen anions, ligand X is usually an electronegative atom such as F, Cl, Br, I, O, H, etc. Further studies on superhalogens were focused on searching for alternative electronegative ligands. In 2009, Anusiewicz [10, 11] used electronwithdrawing groups such as NO2 , CF3 , CCl3 , SHO3 , COOH, COOCH3 , CHO, CONH2 , etc. and acidic functional groups such as ClO4 , ClO3 , ClO2 , ClO, NO3 , PO3 , H2 PO4 , HSO4 , HCO3 , SH, etc. as ligands (X) to design new superhalogen anions. The underlying idea was that these functional groups work as strong electronegative reagents in organic reactions. She used Na and Mg atoms as central core (M) and calculated the VDE of MXk+1 − anions using post-HF methods as listed in Table 1.1. All these ligands were found to be capable of increasing the VDEs and designing new superhalogens, except CHO and CONH2 . Further, these superhalogen anions were found to be stable, except Na(COOCH3 )2 − , which was unstable. Moreover, the largest VDE values were obtained for the anions containing SHO3 (6.0– 8.2 eV) and CF3 (5.2–6.6 eV) ligands. In the same year, Smuczynksa and Skurski [12] employed halogenoids (pseudohalogens) such as CN, NC, OCN, NCO, SCN, and NCS as ligands (X) to design MXk+1 − anions for M = Li, Na, Be, Mg, Ca, B, Al and calculated their VDEs also listed in Table 1.1. These all species were also found to be superhalogens. The largest VDEs correspond to the anions having CN or NC ligands, being as large as 8.94 eV for Al(CN)4 − and 9.20 eV for Al(NC)4 − . These works suggested further possibilities of using novel strategies such as new ligands, new cores, and new formulae to explore superhalogens. Six years later, Sikorska [13] utilized OF ligand and designed M(OF)k+1 − anions (her study was inspired by the work of Srivastava and Misra [14] on Au(OF)n reported a month earlier in which they showed that Au can successively bind with six OF ligands such that its EA increases from 4.34 eV for n = 2 to 6.43 eV for n = 6 and consequently, turning them into superhalogens). From Table 1.1, one can note that the VDE of M(OF)k+1 − anions lies in the range 4.99–7.56 eV, making all of them superhalogens. More recently, Srivastava [15] employed BO ligands to design the M(BO)k+1 − series of anions. Their VDE values, ranging between 4.44 and 7.64 eV, advocated their superhalogen properties (see Table 1.1). Among all these ligands, well-known CN ligands have been very popular particularly due to their higher EA, 3.9 eV, than all ligands and even higher than that of Cl. This was the reason why CN ligands have been further explored with different core atoms [16]. The main disadvantage with the CN ligands is that they have a very strong tendency of forming a dimer, i.e., its dimerization energy (or the binding energy of the dimer) is very high. This prohibits the binding of all CN ligands directly to the core atom. For instance, in the most stable structure of Au(CN)3 , central Au bind with the CN group and (CN)2 dimer [16]. This affects the binding freedom of the

4.69, 5.55 6.04, 6.77

Na(CCl3 )2 − , Mg(CCl3 )3 −

−,

Na(SHO3 )2

Na(COOH)2 − , Mg(COOH)3 −

CCl3

SHO3

COOH

PO3

Na(PO3 )2

−,

Mg(PO3 )3

−

Na(NO3 )2 − , Mg(NO3 )3 −

NO3

Mg(ClO)3

Na(ClO)2

−

ClO

−,

Na(ClO2 )2 − , Mg(ClO2 )3 −

ClO2

Mg(ClO3 )3

Na(ClO3 )2

ClO3

−

Na(ClO4 )2 − , Mg(ClO4 )3 −

ClO4

−,

Mg(COOCH3 )3

COOCH3

−

7.61, 8.67

6.68, 8.29

4.98, 6.14

4.47, 5.29

6.65, 7.55

7.84, 8.91

5.08

4.06, 5.20

5.16, 6.59

Na(CF3 )2 − , Mg(CF3 )3 −

CF3

Mg(SHO3 )3

5.14, 6.50

Na(NO2 )2 − , Mg(NO2 )3 −

NO2

−

VDE (eV)

Superhalogen anions

Ligand

NCS

SCN

NCO

OCN

NC

CN

Ligand

Table 1.1 Superhalogen anions and their VDEs using different ligands, Ref. [10-15]

Li(NCS)2 − , Na(NCS)2 − Be(NCS)3 − , Mg(NCS)3 − , Ca(NCS)3 − B(NCS)4 − , Al(NCS)4 −

Li(SCN)2 − , Na(SCN)2 − Be(SCN)3 − , Mg(SCN)3 − , Ca(SCN)3 − B(SCN)4 − , Al(SCN)4 −

Li(NCO)2 − , Na(NCO)2 − Be(NCO)3 − , Mg(NCO)3 − Ca(NCO)3 − B(NCO)4 − , Al(NCO)4 −

Li(OCN)2 − , Na(OCN)2 − Be(OCN)3 − , Mg(OCN)3 − Ca(OCN)3 − B(OCN)4 − , Al(OCN)4 −

Li(NC)2 − , Na(NC)2 − Be(NC)3 − , Mg(NC)3 − , Ca(NC)3 − B(NC)4 − , Al(NC)4 −

Li(CN)2 − , Na(CN)2 − Be(CN)3 − , Mg(CN)3 − , Ca(CN)3 − B(CN)4 − , Al(CN)4 −

Superhalogen anions

(continued)

5.11, 5.01 5.70, 5.83, 5.78 6.00, 6.26

5.21, 5.17 6.19, 6.31, 5.84 6.00, 6.51

5.74, 5.61 6.34, 6.57 6.51 6.56, 7.10

6.07, 5.84 7.42, 7.33 6.96 8.08, 8.23

7.43, 5.87 8.24, 8.41, 6.95 8.82, 9.20

7.23, 7.09 8.13, 8.24, 8.11 8.50, 8.94

VDE (eV)

6 1 Introduction

Na(SH)2

Mg(SH)3

4.28, 5.11

6.78, 7.37

Na(HCO3 )2 − , Mg(HCO3 )3 −

HCO3

SH

7.31, 8.60

Na(HSO4 )2 − , Mg(HSO4 )3 −

HSO4

−

6.26, 7.23

Na(H2 PO4 )2 − , Mg(H2 PO4 )3 −

H2 PO4

−,

VDE (eV)

Superhalogen anions

Ligand

Table 1.1 (continued)

BO

OF

Ligand

Li(BO)2 − , Na(BO)2 − , K(BO)2 − Be(BO)3 − , Mg(BO)3 − , Ca(BO)3 − B(BO)4 − , Al(BO)4 −

Li(OF)2 − , Na(OF)2 − , K(OF)2 − Be(OF)3 − , Mg(OF)3 − , Ca(OF)3 − B(OF)4 − , Al(OF)4 −

Superhalogen anions

5.54, 4.85, 4.44 6.61, 6.38, 6.18 7.81, 7.64

5.43, 5.15, 4.99 6.63, 6.66, 6.45 7.07, 7.56

VDE (eV)

1.5 Quest for New Ligands 7

8

1 Introduction

Fig. 1.4 Model structures of the dimer of CN, OF, and BO

Au atom and tends to reduce the EA of Au(CN)3 . The structures of (CN)2 , (OF)2 , and (BO)2 dimers are displayed in Fig. 1.4. Although the dimerization energy of OF (1 eV) is lower than that of CN (4 eV), (OF)2 is bent unlike linear (CN)2 . To compare, the dimerization energy of F is only 0.7 eV. However, BO is isoelectronic to CN and its dimer is linear like (CN)2 (see Fig. 1.4). Further, its dimerization energy (2.5 eV) is lower than that of CN. This makes BO a suitable alternative and inorganic analog of CN.

1.6 Towards Maximization of VDEs In early 1984, Gutsev and Boldyrev [17] proposed some binuclear superhalogens including Au2 F11 , whose EA reaches as high as 10 eV. Other binuclear superhalogen anions reported later were Mg2 F5 − (9.38 eV) by Anusiewicz and Skurski [18] and Mg2 (CN)5 − (8.63 eV) by Yin et al. [19]. Note that MgF3 − is already reported to be a superhalogen with a VDE of 8.79 eV [20]. We could expect higher VDEs if we were capable of incorporating more F ligands with the core atom, Mg. This happens due to the fact that the extra electron is delocalized over several ligands. However, there are a few concerns while increasing the number of ligands. The fixed valency of the core atom puts a constraint on the bonding and the ligand repulsion, the steric hindrance, causes to destabilize the anions. The former can be overcome by using the core atom with variable oxidation states such as transition metal atoms as discussed in Chap. 2. Later can be managed by using a strategy as shown in Fig. 1.5 for binuclear superhalogens. For instance, one can replace one of the F atoms in MgF3 − by MgF3 superhalogen moiety to obtain an Mg2 F5 − anion, whose VDE is 9.38 eV, i.e., larger than that of MgF3 − (8.79 eV). The VDE can be further increased by substituting MgF3 moieties

1.6 Towards Maximization of VDEs

9

Fig. 1.5 Strategy to design binuclear superhalogens taken from Ref. [19] with permission of American Institute of Physics, copyright 2014

for other F atoms. This process can be replicated to maximize the VDE value subject to the condition that the ligand repulsion can be avoided (see Fig. 1.5). This enables to design the polynuclear superhalogens (anions) having even higher EA (VDE) values. This was already seen in 2004 in a joint experimental and theoretical study on Nax Clx+1 − anions by Alexandrova et al. [21]. They measured the VDE value of 5.6–7.0 eV for x = 1–4 using photoelectron spectroscopy. In 2007, Smuczynska and Skurski [22] showed that the H-atom can be used as a central core to design superhalogen anions. They reported mononuclear HF2 − and binuclear H2 F3 − anions having VDEs of 6.45 and 7.96 eV, respectively. In 2010, Freza and Skurski [23] extended this work and designed polynuclear Hn Fn+1 − (n = 2–5, 7, 9, 12) superhalogen anions with the VDE as high as 13.87 eV for n = 12 as shown in Fig. 1.6. It can be noticed that the VDE of Hn Fn+1 − have logarithmic relation with n with the correlation coefficient of 0.9961. Thus, the VDE increases sharply from n = 1–5 and begins to increase slowly with the further increasing n. They predicted that it would not be possible to reach the VDE value of 16.22 eV (which is actually the ionization energy of the HF molecule) and beyond. To test their prediction, the authors estimated, by extrapolation, that VDE for n = 25 can not exceed 14 eV. The maximal VDE limit reported, thus far, is limited to 13.87 eV for any superhalogen anion. Several other polynuclear superhalogens have been reported subsequently. In 2010, Willis et al. [24] performed a combined theoretical and experimental study on some complex superhalogens in which they employed superhalogen as ligands, just as in polynuclear superhalogens. For instance, they studied Au as a core atom with two BO2 ligands. Note that BO2 itself is a superhalogen (see Fig. 1.1). The calculated VDE of Au(BO2 )2 − was 5.66 eV, which was in good agreement with the experimentally measured VDE value of 5.9 ± 0.1 eV by photoelectron spectroscopy. The authors used the term ‘hyperhalogen’ to describe such complex superhalogens, which are formed by employing the superhalogen moieties as ligands.

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1 Introduction

Fig. 1.6 Correlation between the number of core atoms (n) and VDE of Hn Fn+1 − anions from Ref. [23] with permission of Elsevier, copyright 2010

1.7 Electron Counting Rules As mentioned earlier, the stability and hence, high VDE of MXk+1 − anions are associated with the octet rule. There are some other electron-counting rules, which can be exploited to design superhalogens such as the 18-electron rule [25], Hückel rule of aromaticity [26], and Wade–Mingos rule [27]. The 18-electron rule applies to the stability and high VDE of metal clusters and organometallic compounds. According to this rule, the central atom requires 18 electrons (2 s-electrons, 6 pelectrons, 10 d-electrons) to complete the electronic shell, e.g., M@Au12 − (M = V, Nb, Ta) superhalogen anions [28]. Hückel rule of aromaticity applies to planar and cyclic molecular systems with conjugated rings having (4n + 2)π electrons. For example, XC5 (CN)6 (X = B, Al, and Ga) [29] has 5π electrons and requires 1 electron, and consequently, XC5 (CN)6 − is an aromatic superhalogen anion. The Wade-Mingo rule applies to carboranes, based on polyhedral skeleton electron pair theory. As per this rule, a closed deltahedral system with n vertices requires (n + 1) electrons to complete the shell. For instance, CB11 H12 (n = 12) needs 13 electrons for stability and hence, the VDE of CB11 H12 − is quite high making it a superhalogen [30].

1.8 Non-Coordinating Anions The anions which interact only weakly with the metal center are referred to as noncoordinating anions (NCAs) or more precisely, weakly coordinating anions. Superhalogen anions introduced by Gutsev and Boldyrev (see Fig. 1.1) are actually classic

1.9 Summary

11

Fig. 1.7 Schematic structures of BArF20 − and BArF24 −

examples of NCAs. For instance, ClO4 − , BF4 − , PF6 − , etc. form complexes with Li such as LiClO4 , LiBF4 , LiPF6 , etc. which are used as electrolytic salts in lithium-ion batteries as discussed further in Chap. 5. The definition of NCA would not suffice without mentioning coordinated anion. According to Rosenthal [31], a “coordinated anion is one which more closely approaches the metal ion than a non-coordinated anion. This will affect the metal ion in such a way that its magnetic moment and spectral properties will be significantly different from what they would be if the anion were not coordinated to it”. Obviously, the vibrational spectrum does not help for NCAs and X-ray diffraction is the only reliable probe. After 1990, several studies have been performed to design weaker coordinating anions [32, 33] than classic NCAs described above using various electron counting rules. The BARF is a popular example of such NCAs, which represent a class of NCAs including [(C6 F5 )4 B]− and [{3,5-(CF3 )2 C6 H3 }4 B]− , abbreviated as BArF20 − and BArF24 − , respectively due to the presence of aryl groups as shown in Fig. 1.7. This establishes a clear and direct relation of superhalogen anions with NCAs. However, little or nothing has been said about this relationship in the literature so far.

1.9 Summary In this Chapter, we provided an introduction to superhalogens. Starting from their background and theoretical prediction in the 1980s, the research on superhalogens was greatly influenced after their experimental confirmation in 1999. Attempts have been made to utilize ligands other than halogen and oxygen atoms. With the introduction of polynuclear superhalogens, the efforts to maximize the VDE have been made. Various electron-counting rules have been exploited to design superhalogen anions. The close relation between superhalogen and non-coordinating anions has also been revealed. It should be noticed that these superhalogen anions can be often found in several complex compounds. For instance, potassium permanganate (KMnO4 ) is well known oxidizing agent, which is easily available in chemical laboratories and

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1 Introduction

used for titration. KMnO4 can be thought of as an ionic salt between K+ and MnO4 − , where MnO4 is a superhalogen with an EA of 5 eV [34]. Thus, superhalogens have the capability to rewrite chemistry in their own ways.

References 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. 31. 32. 33. 34.

Hotop H, Lineberger WC (1985) J Phys Chem Ref Data 14:731–750 https://webbook.nist.gov/cgi/cbook.cgi?ID=C7782447&Mask=20 Bartlett N, Lohmann DH (1962) Proc Chem Soc 115 Compton RN, Reinhardt PW, Cooper CD (2023) J Chem Phys 1978:68 Gutsev GL, Boldyrev AI (1981) Chem Phys 56:277–283 Gutsev GL, Bartlett RJ, Boldyrev AI, Simons J (1997) J Chem Phys 107:3867–3875 Langmuir I (1919) J Am Chem Soc 41:868–934 Wang L-S, Ding C-F, Wang X-B, Barlow SE (1999) Rev Sci Instrum 70(4):1957–1966 Wang XB, Ding CF, Wang LS, Boldyrev AI, Simons J (1999) J Chem Phys 110:4763–4771 Anusiewicz I (2009) J Phys Chem A 113:6511–6516 Anusiewicz I (2009) J Phys Chem A 113:11429–11434 Smuczynska S, Skurski P (2009) Inorg Chem 48:10231–10238 Sikorska C (2015) Chem Phys Lett 638:179–186 Srivastava AK, Misra N (2015) New J Chem 39:9543–9549 Srivastava AK (2022) J Phys Chem A 126:513–520 Samanta D, Wu MM, Jena P (2011) Inorg Chem 50:8918–8925 Gutsev GL, Boldyrev AI (1984) Chem Phys 108:250–254 Anusiewicz I, Skurski P (2007) Chem Phys Lett 440:41–44 Yin B, Li T, Li J-F, Yu Y, Li J-L, Wen Z-Y, Jiang Z-Y (2014) J Chem Phys 140:094301 Anusiewicza I, Sobczyka M, Dazbkowska I, Skurski P (2003) Chem Phys 291(2):171–180 Alexandrova AN, Boldyrev AI, Fu Y-J, Yang X, Wang X-B, Wang L-S (2004) J Chem Phys 121:5709–5719 Smuczynska S, Skurski P (2007) Chem Phys Lett 443:190–193 Freza S, Skurski P (2010) Chem Phys Lett 487:19–23 Willis M, Gotz M, Kandalam AK, Gantefor GF, Jena P (2010) Angew Chem Int Ed 49:8966– 8970 Langmuir I (1921) Science 54:59–67 Hückel E (1930) Eur Phys J A 60:423–456 Wade K (1971) J Chem Soc D 15:792–793 Zhai H-J, Li J, Wang L-S (2004) J Chem Phys 121:8369–8374 Giri S, Child BZ, Jena P (2014) ChemPhysChem 15:2903–2908 Pathak B, Samanta D, Ahuja R, Jena P (2011) ChemPhysChem 12:2423–2428 Rosenthal MR (1973) J Chem Educ 50:331–335 Strauss SH (1993) Chem Rev 93:927–942 Krossing I, Raabe I (2004) Angew Chem Int Ed 43:2066–2090 Gutsev GL, Rao BK, Jena P, Wang X-B, Wang L-S (1999) Chem Phys Lett 312:598–605

Chapter 2

Transition Metal Fluorides and Oxides as Superhalogens

Abstract Transition metal atoms possess variable oxidation states, which permit them to bind with several ligands. In this Chapter, the superhalogen behavior of the species with transition metal atom (M) as a central core with increasing fluorine (F) and oxygen (O) ligands. It was noticed that the EA of MFn increases successively with the increase in n such that the species tend to become superhalogens. The gold (Au) based species have been exclusively focused and AuXn (X = F, Cl, CN, and OF) are found to be superhalogens for n ≥ 2, irrespective of the nature of the ligands. The formation of dimer and complex was demonstrated by using AuO3 superhalogen. Keywords Superhalogens · Electron affinity · Transition metal fluorides · Transition metal oxides · Oxidation states · Gold compounds

2.1 Introduction As mentioned in Chap. 1, the superhalogen properties of MXk+1 species were based on the s and p block elements as central core (M) with F or O as ligands (X). However, these core atoms have fixed valence, allowing them to bind with a fixed number of ligands. To overcome this, transition metal atoms were employed as a central core due to the fact that their oxidation states may vary and consequently, they can bind with several ligands. The increase in ligands not only increases the oxidation state of the transition metal atom but also enhances the EA of the resulting species. It should be noted that the superhalogen behavior of various transition metal hexafluorides and tetroxides have been studied by Gutsev and Bolydrev [1–3] in early 1984, but systematic reports appeared only after 2009.

© The Author(s), under exclusive license to Springer Nature Switzerland AG 2023 A. K. Srivastava, Superhalogens, SpringerBriefs in Molecular Science, https://doi.org/10.1007/978-3-031-37571-2_2

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2.2 Transition Metal with Fluorine Ligands Jena and coworkers [4] presented the first comprehensive report on the investigation of superhalogen behavior in transition metal fluorides by studying the CuFn species and their anions. The authors chose a Cu core bonded with one to six F atoms consecutively. The ground-state structures of CuFn clusters and their anions were obtained by considering different spin multiplicities employing density functional theory using B3LYP functional. The objectives of this study were as follows: How does a copper atom bind with successive fluorine ligands? What may be the maximum possible oxidation states of the Cu metal atom? How does the electron affinity vary with the increase in the number of ligands? How does the species tend to become superhalogens? Do they really mimic the properties of halogens?

To answer these questions, the authors explored CuFn (n = 1–6) in their neutral and anionic forms. The ground-state structures of CuFn are displayed in Fig. 2.1. Except CuF3 − and CuF5 − , all species prefer the lowest spin-multiplicities. Generally, the oxidation states of Cu are +1 and +2 as exemplified in Cu2 O and CuO. Hence, it is worthwhile to investigate how more than two F atoms are bonded to the Cu atom. One can see from Fig. 2.1, Cu can bind with up to six F atoms to form CuFn species. To verify their stability, authors calculated fragmentation energies of CuFn against the F atom and F2 molecule as plotted in Fig. 2.2a and b. One can see that CuF5 , CuF6, and CuF6 − are unstable against dissociation to the F2 molecule due to negative fragmentation energy. Considering the stability of CuF4 − and CuF5 − , it could be concluded that the oxidation state of Cu might be +3 and +4 as well. Such a high oxidation state of Cu was explained due to the involvement of 3d electrons in bonding. The authors also plotted the contribution of d electrons in bonding in Fig. 2.2c. They noticed that this contribution increases with the increase in F ligands in CuFn . Finally, the authors calculated adiabatic EA (see Chap. 1) of CuFn and plotted it as a function

Fig. 2.1 Equilibrium structures of CuFn species and their anions (n = 1–6) from Ref. [4] with permission of American Institute of Physics, copyright 2009

2.2 Transition Metal with Fluorine Ligands

15

Fig. 2.2 The plots of fragmentation energy against F atom (a), F2 molecule (b), number of delectrons in bonding (c), and EA of CuFn (d) as a function of n from Ref. [4] with permission of American Institute of Physics, copyright 2009

of n, see Fig. 2.2d. One can see that the EA of CuFn species increases with the increase in n up to 5 and becomes as high as 7.2 eV for CuF5 . They concluded that the EA of CuFn sufficiently exceeds that of Cl for n ≥ 3 and consequently, all these species may behave as superhalogens. The authors also showed that CuF4 superhalogen forms ionic complex K-CuF4 similar to KF salt. A comparative study on 11th group transition metal fluorides, namely, CuFn , AgFn , and AuFn was performed later by Jena and coworkers [5] for n = 1–7. In this study, the authors considered all possible structures. They noticed that CuFn species have a tendency to make complexes such as (CuF3 )F2 , and (CuF4 )F2 with F2 molecule bonded to core atom (Cu) atom or ligand atom (F) for n = 5 and 6. Therefore, it is unlikely to obtain Cu in the oxidation state exceeding +3. This finding was consistent with some earlier results as mentioned previously [6]. It should be noted that these coinage metals have different oxidation states, despite being in the same group. Like CuFn , AgFn species (n ≤ 4) adopt geometries having core (Ag) bonded directly to all ligands (F). These species tend to form complexes like (AgFn−1 )F or (AgFn−2 )F2 for n = 5–7, independent of their neutral or anionic nature. The relativistic effects, however, dominate in the case of gold (Au), making it unique in the group. For instance, it is capable of binding atomically with as large as six F atoms, irrespective of the charge of resulting species AuFn (n = 7), on the contrary, exists in the form of complexes such as (AuF5 )F2 for neutral species and (AuF6 )F for the anion. Thus, the maximum oxidation state of Au seems to be limited to +5.

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Fig. 2.3 The EA of transition metal fluorides, MFn at the B3LYP level taken from Refs. [5], [7–13]

These findings of Jena and co-workers invited several studies on other transition metal fluorides such as ScFn , TiFn , VFn, and MnFn by Pradhan et al. [7, 8], FeFn by Ding et al. [9], PdFn and RhFn by Siddiqui et al. [10, 11], CoFn , NiFn and OsFn by our group [12, 13], etc. The EA of all these clusters is plotted in Fig. 2.3. One can see that the increase in the ligand F atoms leads to an increase in the EA of MFn species. This happens due to the delocalization of negative charge over all ligand atoms. However, one should keep in mind that such an increase in the EA with ligand atoms is governed by the nature of the core atom itself. To exemplify, contrary to AuFn and OsFn , the EA of WFn is relatively very small, e.g., 1.86–3.72 eV [14] for n = 1–5 at the same level of theory.

2.3 Transition Metal with Oxygen Ligands Unlike fluorine, the use of oxygen as a ligand to design transition metal oxides-based superhalogens was not paid much attention due to its limited bonding nature. This is presumably because of the fact that its valency is 2. In 2013, our group carried out a comprehensive study [15] for the first time to explore the superhalogen properties in palladium oxides, PdOn for n = 1–5 using the B3LYP functional. The ground state structures of these species along with their anions are displayed in Fig. 2.4. It is evident that PdOn clusters have O atoms bonded to Pd for n ≤ 4. For n = 5, it adopts the structure of (PdO3 )O2 complex, irrespective of its charge. The energetic stability of these clusters was confirmed by positive dissociation energy values for fragmentation to the O2 molecule and O atom as plotted in Fig. 2.5 a and b.

2.3 Transition Metal with Oxygen Ligands

17

Fig. 2.4 Equilibrium structures of PdOn (n = 1–5) and its anion with spin multiplicities from Ref. [15] with permission of Wiley, copyright 2013

Apparently, the formal oxidation state of Pd is limited to +2 due to the formation of PdO and palladium acetate. Therefore, its bonding with more than two O atoms can be expected by considering the involvement of inner shell d-electrons as displayed in Fig. 2.5c. One can see that as the number of O atoms increases, more d-electrons contribute to the bonding of PdOn clusters, which explains the bonding of the Pd with more than two O atoms. From Fig. 2.5d, it can be noticed that the EA of PdOn species rises up with an increase in n up to 4. More interestingly, PdO4 tends to become a superhalogen due to its EA of 4.57 eV. The increase in the EA happens as a consequence of the delocalization of the electronic charge over a number of ligands as mentioned above. In order to verify this, we analyzed the charge localized on the Pd atom in PdOn species and their anions, not included in Fig. 2.5 (see, Ref. [15]). We noticed that for PdO, there is a localization of electronic charge on the Pd atom such that the positive charge on Pd is reduced. With the increase in O atoms, however, the positive charge on the Pd atom is not reduced, i.e., the extra electronic charge starts delocalizing over ligands consequently, increasing the EA of PdOn . Just to mention, almost 90% of extra charge is delocalized over all O atoms in PdO4 , resulting in its high EA, i.e., superhalogen property (Fig. 2.5d). Subsequently, Srivastava and Misra [16] studied AuOn (n = 1–4) in their neutral and various anionic forms. We noticed that although Au can bind up to four O atoms dissociatively. The EA of AuOn increases up to 4.25 eV for n = 3 and then is reduced to 4.10 eV for n = 4. We concluded that AuO3 can indeed behave as a superhalogen, whose anion is displayed in Fig. 2.6. We also explored the tendency of AuO3 to form a dimer, (AuO)3 and complex K-AuO3 , just analogous to typical halogen atoms (see Fig. 2.6). We have also performed [17] a comparative study on MOn for M = Ni, Pd, and Pt; n = 1–5. It was observed that Pt was capable of binding with as large as five O ligands, contrary to Ni or Pd. Therefore, one can expect the higher oxidation states of Pt to exist. Our group further investigated other transition metal oxides [18–21] and found that several transition metal atoms, e.g., Re, Ru, and Co were capable of binding with

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Fig. 2.5 The plots of dissociation energy against O atom (a), O2 molecule (b), number of d-electrons in bonding (c), and EA of PdOn species (d) as a function of n from Ref. [15] with permission of Wiley, copyright 2013

Fig. 2.6 AuO3 − superhalogen anion, its dimer and complex from Ref. [16] with permission of Wiley, copyright 2013

2.4 Transition Metal with Other Ligands

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Fig. 2.7 The EA (in eV) of transition metal oxides, MOn at the B3LYP level taken from Refs. [15], [17– 21]

five O atoms. According to theoretical calculations, these transition metal pentoxides (MO5 ) were obtained to be stable. However, the difficulties in the gas phase reactions may put some constraints on their experimental viability. On the other hand, Au and V were not capable of binding with five O atoms, and consequently, their oxidation states were found to be limited. The EA of transition metal oxides (MOn ) is plotted in Fig. 2.7, which reveals the superhalogen properties of MOn species for n ≥ 3.

2.4 Transition Metal with Other Ligands Note that transition metal atoms with Cl and Br ligands have been also explored for their superhalogen behavior [7, 8]. Our group attempted to employ both O and F ligands simultaneously to investigate the superhalogen properties of MnOx Fy [22] and AuOx Fy [23] clusters. However, the EA of the clusters with mixed O and F ligands are marginally decreased as compared to respective oxides and fluorides. To mention, the EA of MnO2 F2 (4.64 eV) is smaller than those of MnO4 (5.06 eV) and MnF4 (5.60 eV) at the same level. Therefore, mixing O and F atoms is also capable of inducing superhalogen behavior, but with relatively lower EA values. Gold (Au) has been extensively studied as a core atom with several ligands. After the study on AuFn in 2010 [5], Samanta et al. [24] studied the superhalogen properties of gold cyanides, Au(CN)n (n = 1–6) in 2011. The authors found that the Au is capable of binding with six CN groups in Au(CN)n and their anions. They noticed that the EA of Au(CN)n is raised from 6.0 eV for n = 2 to 8.4 eV for n = 6.

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In 2013, Srivastava and Misra [25] reported AuCln (n = 1–6) species to investigate their superhalogen nature. They suggested that the Au is capable of binding with five Cl atoms in neutral and six Cl atoms in anions. The EA of AuCln is found to be increased from 4.7 to 6.3 eV for n = 2 to 5, respectively, and reduced to 5.2 eV for n = 6. Subsequently, Srivastava and Misra [26] carried out a systematic investigation on Au(OF)n in 2015 and found that Au is capable of binding with six OF ligands. The EA of Au(OF)n becomes in the range 4.3–6.4 eV for n = 2–6. All these studies were performed at the same level of theory. This provided the opportunity to compare their EA as plotted in Fig. 2.8. One can note that the EA of AuFn and Au(CN)n species shows a quite similar pattern. The peaks in the EA of AuFn for even values of n have already been explained on the basis of ionic resonance [5]. It is also clear that the EA values of Au(CN)n are higher than all other ligands. This is due to the fact that the EA of CN is larger than all other ligands. Interestingly, a similar pattern has also been observed in the EA values of AuCln and Au(OF)n species, i.e., they increase successively only up to n = 5. Although the EA of Cl is larger than that of F, the EA of AuCln are smaller than those of AuFn species but comparable to those of Au(OF)n . As known, the electron taken by molecular species goes to the lowest unoccupied molecular orbital (LUMO). It was noticed that usually heavier ligands possess a larger number of electrons and therefore, the energy of LUMO is lowered. Although the EA of OF moiety, calculated to be only 2.38 eV [26], is much smaller than that of Cl, its

Fig. 2.8 A comparative plot of the EAs of AuFn , AuCln , Au(CN)n, and Au(OF)n species from Ref. [26] with permission of Royal Society of Chemistry, copyright 2015

References

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LUMO energy level is comparable to that of Cl. This can explain why the EAs of Au(OF)n are comparable to those of the corresponding AuCln species. Finally, it can be concluded that the EAs of metal complexes are closely associated with the electronic structures of the ligands. Further, it is evident that all the Au-complexes (see Fig. 2.8) behave as superhalogens for n ≥ 2, irrespective of the nature of the ligands.

2.5 Summary In this Chapter, we discussed transition metal atoms as a central core to design superhalogens. Due to their variable oxidation states, they can bind with several ligands. We have chosen various transition metal atoms (M) with F, O, and other ligands. The binding with successive F atoms not only increases the EA of MFn species but also pushes up the oxidation state of M. In general, the EA of MFn and MOn increases with the increase in n, which is primarily due to the delocalization of extra electron over several ligand atoms. However, the EA of MOn is smaller than corresponding MFn , which is expected due to low EA and double valency of O atom as compared to F. Taking Au core as an example, we found that the EAs of AuXn (X = F, Cl, CN and OF) are closely related to the electronic structures of the ligands (X) such that they all behave as superhalogens for n ≥ 2, irrespective of the nature of the ligands. These superhalogens can form dimer and complexes like halogens do. For instance, the complex salt of AuF6 superhalogen, Li-AuF6 has already been synthesized [27].

References 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18.

Gutsev GL, Boldyrev AI (1983) Chem Phys Lett 101:441–445 Gutsev GL, Boldyrev AI (1984) Mol Phys 53:23–31 Gutsev GL, Boldyrev AI (1984) Chem Phys Lett 108:255–258 Wang Q, Sun Q, Jena P (2009) J Chem Phys 131:124301 Koirala P, Willis M, Kiran B, Kandalam AK, Jena P (2010) J Phys Chem C 114:16018–16024 Riedel S, Kaupp M (2009) Coord Chem Rev 253:606 Pradhan K, Gutsev GL, Jena P (2010) J Chem Phys 133:14430 Pradhan K, Gutsev GL, Weatherford CA, Jena P (2011) J Chem Phys 134:234311 Ding L-P, Kuang X-Y, Shao P, Zhong M-M, Zhao Y-R (2013) J Chem Phys 139:104304 Siddiqui SA, Pandey AK, Rasheed T, Mishra M (2012) J Fluor Chem 135:285–291 Siddiqui SA, Bouarissa N (2013) Solid State Sci 15:60–65 Srivastava AK, Misra N (2015) Mol Phys 113:36–44 Srivastava AK, Misra N (2014) J Fluor Chem 158:65–68 Srivastava AK, Panday AK, Misra N (2015) J Chem Sci 127:1853–1858 Srivastava AK, Misra N (2014) Int J Quantum Chem 114:328–332 Srivastava AK, Misra N (2014) Int J Quantum Chem 114:521–524 Srivastava AK, Misra N (2014) Mol Phys 112:1639–1644 Srivastava AK, Misra N (2014) Mol Phys 112:1963–1968

22 19. 20. 21. 22. 23. 24. 25. 26. 27.

2 Transition Metal Fluorides and Oxides as Superhalogens Srivastava AK, Misra N (2014) Eur Phys J D 68:309 Srivastava AK, Misra N (2015) Theor Chem Acc 134:93 Shukla DV, Srivastava AK, Misra N (2017) Main Group Chem 16:141–150 Srivastava AK, Misra N (2014) Mol Phys 112:2820–2826 Srivastava AK, Misra N (2015) Chem Phys Lett 630:106–110 Samanta D, Wu MM, Jena P (2011) Inorg Chem 50:8918–8925 Srivastava AK, Misra N (2014) Int J Quantum Chem 114:1513–1517 Srivastava AK, Misra N (2015) New J Chem 39:9543–9549 Graudejus O, Elder SH, Lucier GM, Shen C, Bartlett N (1999) Inorg Chem 38:2503

Chapter 3

Superhalogens as Strong Oxidizers

Abstract Stable molecules and inert gas atoms usually possess high IE, which makes their ionization difficult. Due to high EA, however, superhalogens possess high oxidizing power. In this Chapter, the strong oxidizing power of superhalogens has been discussed. It has been shown that superhalogens are not only capable of ionizing a variety of stable inorganic molecules such as H2 O, NO, metal oxides, etc., and organic molecules such as CO2 , C6 H6 , heterocycles, etc. but also of forming stable noble gas compounds. For successful ionization of systems, the EA of superhalogens should be comparable to the IE of the systems. Keywords Superhalogens · Electron affinity · Oxidation · Ionization · Noble gas compounds · Ionic compounds

3.1 Introduction In early 1962, Bartlett and group [1, 2] synthesized some complex compounds such as O2 PtF6 and XePtF6 . The stability of these compounds was due to the capability of PtF6 to oxidize the Xe atom and O2 molecule, both having ionization energy (IE) exceeding 12 eV. This was possible due to the high EA of PtF6 , 7.00 ± 0.35 eV [3], making it a superhalogen as discussed in Chap. 2. The high IE of stable molecules such as O2 , CO2 , C6 H6 , etc., and inert atoms such as Ar, Xe, Rn, etc. make their oxidation very difficult. After several decades, it was realized that the superhalogens are capable of oxidizing stable molecules and inert atoms, and forming stable ionic complexes with them.

© The Author(s), under exclusive license to Springer Nature Switzerland AG 2023 A. K. Srivastava, Superhalogens, SpringerBriefs in Molecular Science, https://doi.org/10.1007/978-3-031-37571-2_3

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3.2 Oxidation of Water and Its Clusters The IE of H2 O is as large as 12.6 eV whereas that of its small clusters, (H2 O)n (n = 2–4) is slightly lower (12.3–11.8 eV) [4, 5]. In 2013, Marchaj et al. [6] employed BF4 and AlF4 superhalogens and study their interaction with water clusters, (H2 O)n for n = 1–4. They noticed that the H2 O and its clusters form ionic complexes with BF4 / AlF4 superhalogen as shown in Fig. 3.1 for BF4 (the structures are almost similar for AlF4 ). This was confirmed by calculating the charge transfer of 0.53–0.83e and 0.63–0.83e for BF4 and AlF4 , respectively. The binding energies of [H2 O]n + BF4 − complexes are large enough to confirm their stability, which is, even more, higher for [H2 O]n + AlF4 − . This was due to a higher EA of AlF4 (7.88 eV) than that of BF4 (6.66 eV). To check whether it is possible to ionize (H2 O)n cluster with superhalogens with lower EAs, they chose BCl4 and AlCl4 superhalogens, which are similar in structures with BF4 and AlF4 but with the EA of 5.32 and 5.92 eV, respectively. The structure of (H2 O)3 BCl4 and (H2 O)AlCl4 are also displayed in Fig. 3.1, which are non-ionized complexes. It can be seen that unlike (H2 O)3 + BF4 − , there is no H3 O moiety in (H2 O)3 BCl4 due to a very small charge transfer of 0.29e. Therefore, they concluded that BF4 and AlF4 , but not BCl4 and AlCl4 , superhalogens are capable of ionizing or oxidizing water clusters.

Fig. 3.1 Equilibrium structures of (H2 O)n BF4 complexes for n = 1–4, (H2 O)3 BCl4 and (H2 O)AlCl4 from Ref. [6] with permission of Elsevier, copyright 2013

3.4 Oxidation of Fullerene

25

Fig. 3.2 Equilibrium structures of (C6 H6 )X complexes for X = LiF2 , MgF3 , AlF4, and LiCl2 from Ref. [8] with permission of Elsevier, copyright 2015

3.3 Oxidation of Benzene Benzene (C6 H6 ) is the most stable aromatic molecule with an IE of 9.24 eV [7]. In 2015, Czapla et al. [8] studied the ionization of C6 H6 by using LiF2 , MgF3 , and AlF4 as well as LiCl2 superhalogens. Although they found two stable isomers for each complex, only the lowest energy structures are shown in Fig. 3.2. Due to the significant amount of charge transfer, these complexes are characterized as ionic. The binding energy of these ionic complexes, (C6 H6 )+ X− (~35–60 kcal/mol), increases from X = LiF2 to MgF3 to AlF4 due to their increased VDEs from 6.51 to 8.79 to 9.79 eV and charge transfer from 0.48 to 0.58 to 0.71e, respectively. On the contrary, the complex of benzene with LiCl2 is weakly bound due to low binding energy (~5 kcal/mol) and low charge transfer (0.39e) as a consequence of its lower VDE (5.90 eV). Therefore, AlF4 , MgF3, and LiF2 can oxidize benzene molecule and form ionic products, while LiCl2 does not.

3.4 Oxidation of Fullerene Buckminsterfullerene (C60 ) is a stable nanoparticle with an IE of 7.58 eV [9]. In 2016, Sikorska [10] reported the ionization of C60 by LiX2 , MgX3, and AlX4 superhalogens (X = F, Cl). One can note that LiF2 , MgF3 and AlF4 superhalogens, whose anions have VDEs in the range 6.51–9.79 eV, are capable of ionizing fullerene and form strongly bound ionic complexes C60 + LiF2 − , C60 + MgF3 − and C60 + AlF4 − as reflected by their binding energies and charge transfer displayed in Fig. 3.3. Although LiCl2 , MgCl3 and AlCl4 can also ionize C60 , the binding energy (~17–29 kcal/mol) and charge transfer (0.63–0.80e) are relatively smaller (not shown in Fig. 3.3). This can be expected as LiCl2 − , MgCl3 − and AlCl4 − are relatively weak electron acceptors due to their lower VDEs, 5.90–6.68 eV. The author also noticed that LiI2 with VDE of 4.57 eV can not ionize C60 due to low charge transfer and its complex is not stable due to negative binding energy.

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Fig. 3.3 Equilibrium structures of (C60 )X complexes for X = LiF2 , MgF3, and AlF4 from Ref. [10] with permission of Royal Society of Chemistry, copyright 2016

3.5 Oxidation of Carbon Dioxide Carbon dioxide (CO2 ) is an extremely stable molecule with a high IE of 13.8 eV [11]. Czapla and Skurski [12] attempted the ionization of CO2 using the Sbn F5n+1 series of superhalogens for n = 1–3. The VDE of these polynuclear superhalogen anions ranges 10.2–13.3 eV [13]. The lowest energy structures of these complexes are displayed in Fig. 3.4. They observed that SbF6 and Sb2 F11 are incapable of ionizing CO2 due to the charge transfer of 0.07–0.11e and their complexes are weakly bound with the binding energy less than 15 kcal/mol. On the contrary, Sb3 F16 can indeed oxidize CO2 molecule with the charge transfer of 0.88e such that strongly bound ionic complex (CO2 )+ (Sb3 F16 )− possesses the binding energy exceeding 26 kcal/ mol. The authors suggested that this can be expected due to the fact that the VDE of Sb3 F16 − (13.3 eV) is comparable to the IE of the CO2 molecule. According to them, Sb3 F16 was probably the first system to ionize carbon dioxide successfully.

Fig. 3.4 Equilibrium structures of CO2 (Sbn F5n+1 ) complexes for n = 1, 2, and 3 from Ref. [12] with permission of Royal Society of Chemistry, copyright 2015

3.7 Oxidation of Heterocycles

27

Fig. 3.5 Equilibrium structures of Mg3 F7 − anion, NiO(Mg3 F7 ), and TiO2 (Mg3 F7 ) complexes from Ref. [14] with permission of American Chemical Society, copyright 2018

3.6 Oxidation of Transition Metal Oxides Sikorska [14] worked on the oxidation of some transition metal monoxides, MO (M = Co, Cu, Ni, Zn) and dioxides, MO2 (Mn, Si, Ti) using polynuclear superhalogen, Mg3 F7 , whose anion has VDE of 10.47 eV [15]. The IE of these metal oxides ranges 9.38–12.58 eV. She noticed that Mg3 F7 can ionize these metal oxides and form strongly bound complexes, (CoO)+ (Mg3 F7 )− , (CuO)+ (Mg3 F7 )− , (MnO2 )+ (Mg3 F7 )− , (NiO)+ (Mg3 F7 )− , (TiO2 )+ (Mg3 F7 )− , (ZnO)+ (Mg3 F7 )− having binding energies 4.89–8.26 eV by transferring a charge of 0.47–0.71e. The lowest energy structures of Mg3 F7 − superhalogen anion, (NiO)+ (Mg3 F7 )− and (TiO2 )+ (Mg3 F7 )− complexes are shown in Fig. 3.5. On the contrary, Mg3 F7 was not capable of ionizing SiO2 , whose IE exceeds 12 eV. Instead, it tends to form (Mg3 F5 )+ (F2 SiO2 )− complex.

3.7 Oxidation of Heterocycles Xue et al. [16] employed LiF2 , LiCl2 , MgF3 , and AlF4 superhalogens to oxidize some heterocyclic compounds such as pyrrole (C4 H5 N), furan (C4 H4 O), thiophene (C4 H4 S) and pyridine (C5 H5 N). The IEs of these aromatic heterocyclic compounds lie in the range of 8.10–9.51 eV [17]. The lowest energy structures of complexes of heterocycles with superhalogens are displayed in Fig. 3.6. The authors noticed that AlF4 , MgF3, and LiF2 superhalogens were capable of oxidizing C4 H5 N, C4 H4 O, and C4 H4 S due to charge transfer of 0.51–0.82e and form strong ionic compounds with binding energies 33–89 kcal/mol. However, these superhalogens were probably not appropriate for ionizing C5 H5 N having relatively high IE. They found that the interaction between C5 H5 N and superhalogens results in merely charge redistribution with a little charge transfer such that C5 H5 N retains its original geometry (see Fig. 3.6). Similarly, LiCl2 (due to its smaller EA) form weakly bound complexes with C4 H5 N, C4 H4 O, and C4 H4 S with binding energies 0.5–10 kcal/mol. Further, the complex of LiCl2 with C5 H5 N becomes unstable due to negative binding energy.

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Fig. 3.6 Equilibrium structures of (Het)X complexes for Het = C4 H5 N, C4 H4 O, C4 H4 S, and C5 H5 N; X = AlF4 , MgF3 , LiF2, and LiCl2 from Ref. [16] with permission of Elsevier, copyright 2021

3.8 Oxidation of Nitric Oxide Srivastava [18] reported the interaction of nitric oxide (NO) with LiF2 , BeF3 , BF4 superhalogens. The ionization of NO is difficult to its IE of 9.26 eV [19]. The author noticed that the interaction resulted in (NO)X complexes having two isomers. The lowest energy structures correspond to ones in which superhalogen interacts to N atom of NO as displayed in Fig. 3.7. The binding energy (2.43–2.85 eV) and charge transfer (0.49–0.66e) of these complexes increase with the increase in the size of

3.10 Summary

29

Fig. 3.7 Equilibrium structures of (NO)X complexes for X = LiF2 , BeF3, and BF4 from Ref. [18]

superhalogens (LiF2 → BeF3 → BF4 ), which is evidently due to increase in their EAs (5.10–7.05 eV). This charge transfer is even more (0.65–0.79e) in the case of its isomer in which the superhalogen interacts with the O atom of NO. Nevertheless, this causes to ionize NO to NO+ . He noticed that the size of superhalogens and their EA play a crucial role in their oxidizing tendency.

3.9 Formation of Noble Gas Compounds As mentioned earlier, superhalogens can form stable compounds with noble gases (Ng’s) due to their high EA. Note that the existence of HNgX (X = halogen) compounds is limited to the fact that these are metastable despite substantial ionic bonding between Ng and X atoms. In 2014, Samanta [20] studied HNgF (Ng = Ar and Kr) and replaced X with superhalogens, BO2 and BF4 . The structures of these compounds are displayed in Fig. 3.8. The author found that the stability of HNgBO2 and HNgBF4 is increased as compared to HNgF compounds. This was due to increased charge transfer to superhalogen, which resulted in stronger ionic bonds between Ng and superhalogens. This makes it possible to design Ng compounds by employing superhalogens. In 2015, Chakraborty and Chattaraj [21] reported a Xe compound with some transition metal trifluorides XeMF3 (M = Ru, Os, Rh, Ir, Pd, Pt, Ag, Au). Note that all these transition metal trifluorides (MF3 ) behave as superhalogens (see Chap. 2). They concluded that the interaction between Xe and M in these compounds is partially covalent, and orbital interaction plays a dominant role in XeMF3 compounds.

3.10 Summary In this Chapter, we have shown that superhalogens can act as strong oxidizers due to their high EA. They can not only ionize molecules with high IE but also form stable noble gas compounds. By choice of appropriate superhalogens, one can ionize stable inorganic compounds such as water, metal oxides, and nitric oxide as well as organic compounds such as benzene and aromatic heterocycles. It is even possible to ionize

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Fig. 3.8 Equilibrium structures of (NgH)X complexes for Ng = Ar and Kr; X = F, BO2, and BF4 from Ref. [20] with permission of American Chemical Society, copyright 2014

carbon dioxide having IE exceeding 13 eV. The ionizing or oxidizing ability of superhalogen is directly proportional to its EA, i.e., superhalogens with higher EA are more capable of ionizing molecules and forming ionic complexes. Consequently, the stability of ionic complexes increases with the ionizing ability of the superhalogen involved. In general, it has been found that superhalogens are capable of oxidizing molecules when their EA is comparable to the IE of the molecules to be ionized. It is possible, in principle, to oxidize any molecule with a cleverly chosen superhalogen.

References 1. N. Bartlett, Proc. Chem. Soc., 1962, 218. 2. N. Bartlett and D. H. Lohmann, J. Chem. Soc., 1962, 5253–5261. 3. Korobov MV, Kuznetsov SV, Sidorov LN, Shipachev VA, Mitkin VN (1989) Int J Mass Spectrom Ion Processes 87:13–27 4. Hayashi H, Watanabe N, Udagawa Y, Kao C-C (2000) Proc Natl Acad Sci USA 97:6264 5. Segarra-Martí J, Merchán M, Roca-Sanjuán D (2012) J Chem Phys 136:244306 6. Marchaj M, Freza S, Rybacka O, Skurski P (2013) Chem Phys Lett 574:13–17 7. Nemeth GI, Selzle HL, Schlag EW (1993) Chem Phys Lett 215:151 8. Czapla M, Freza S, Skurski P (2015) Chem Phys Lett 619:32–35 9. Devries J, Steger H, Kamke B, Menzel C, Weisser B, Kamke W, Hertel IV (1992) Chem Phys Lett 188:159–162 10. Sikorska C (2016) Phys Chem Chem Phys 18:18739–18749

References 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21.

Wang L, Reutt JE, Lee YT, Shirley DA (1988) J Electron Spectrosc Relat Phenom 47:167 Czapla M, Skurski P (2017) Phys Chem Chem Phys 19:5435–5544 Czapla M, Skurski P (2015) J Phys Chem A 119:12868 Sikorska C (2018) J Phys Chem A 122:7328–7338 Sikorska C (2018) Int J Quantum Chem 118:1–15 Xue D, Chen Z, Liu J, Liu J, Wu D, Li Y, Li Z (2021) Polyhedron 201:115160 http://webbook.nist.gov/chemistry. Srivastava AK (2021) Main Group Chem 20:33–40 Reiser G, Habenicht W, Muller-Dethlefs K, Schlag EW (1988) Chem Phys Lett 152:119 Samanta D (2014) J Phys Chem Lett 5:3151–3156 Chakraborty D, Chattaraj PK (2015) J Phys Chem A 119:3064–3074

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Chapter 4

Superhalogens in the Design of Superacids

Abstract H2 SO4 is the strongest mineral acid having deprotonation energy as low as approximately 300 kcal/mol. Superacids possess even lower deprotonation energy than 300 kcal/mol, e.g., HSbF6 . Note that SbF6 is a superhalogen. In this Chapter, the possibility of designing new superacids by protonation of superhalogen anions has been discussed. Using a variety of superhalogens, it has been found that their protonated complexes may indeed behave as superacids. Further, their acidity increases with the increase in the VDE of constituting anions. It has been emphasized that the protonation of superhalogen anions is a rational route to design new superacids. Keywords Superhalogens · Vertical detachment energy · Acidity · Superacid · Deprotonation energy · Protonation

4.1 Introduction Strong acids can easily release proton (H+ ) and possess low deprotonation energy. H2 SO4 is the strongest mineral acid, followed by HNO3 and HCl. The Gibbs’ free energy of deprotonation for HCl, HNO3 , and H2 SO4 is reported to be 328.1, 317.8, and 302.2 kcal/mol, respectively. In 1927, Hall and Conant [1] introduced the term ‘superacids’ for the acids, which are more acidic than pure H2 SO4 . HSbF6 was considered to be the strongest superacid, which is 109 times stronger than pure H2 SO4 [2, 3] having Gibbs’ free energy of deprotonation of 255.5 kcal/mol. Note that SbF6 − is luckily a superhalogen anion with a VDE of 10.2 eV. In this Chapter, we will see how superhalogens can be used to design new superacids.

4.2 Acidity of HAln F3n+1 Species In 2015, Czapla and Skurski [4] designed HAln F3n+1 species by the protonation of Aln F3n+1 − anions whose lowest energy structures are displayed in Fig. 4.1.

© The Author(s), under exclusive license to Springer Nature Switzerland AG 2023 A. K. Srivastava, Superhalogens, SpringerBriefs in Molecular Science, https://doi.org/10.1007/978-3-031-37571-2_4

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4 Superhalogens in the Design of Superacids

Fig. 4.1 Equilibrium structures of HAln F3n+1 ; n = 1–4 from Ref. [4] with permission of Elsevier, copyright 2015

Note that Aln F3n+1 − are polynuclear superhalogen anions, whose VDEs lie in the range 9.79–12.40 eV [5]. They calculated Gibb’s free energy of deprotonation (Gdep ) in the range 269.2–246.3 kcal/mol for n = 1–4. Note that these values are lower than that of H2 SO4 (302.2 kcal/mol). Therefore, the authors proposed that the gas phase acidity of HAln F3n+1 is higher than that of the strongest mineral acids, thus, making them superacids. The increasing acidity is a result of decreasing Gdep values. It was also noticed that the Gdep of HAl4 F13 (246.3 kcal/mol) is even lower than that of HSbF6 (255.5 kcal/mol). They also mentioned that although these calculations were performed in the gas phase, the conclusions should be valid in the liquids as well.

4.3 Protonated Superhalogens as Superacids Srivastava and Misra [6] investigated whether the protonation of superhalogens leads to superacids. They employed typical superhalogen anions such as LiF2 − , LiCl2 − , BeF3 − , BeCl3 − , BF4 − , and PF6 − and studied their protonated complexes as displayed in Fig. 4.2. These complexes turned out to be two-component systems having HF/ HCl moieties. Therefore, the authors also checked their stability against dissociation to HF/HCl and found them indeed stable.

4.3 Protonated Superhalogens as Superacids

35

Fig. 4.2 Equilibrium structures of HX for X = LiF2 , LiCl2 , BeF3 , BeCl3 , BF4 , and PF6 from Ref. [6] with permission of Elsevier, copyright 2015

They calculated that the Gdep values of these species, were in the range 290.6– 272.4 kcal/mol, except HLiF2 whose Gdep value was found to be 319.1 kcal/mol. Based on their Gdep values, the gas phase acidities are in the order of HBeCl3 > HPF6 > HLiCl2 > HBeF3 > HBF4 > H2 SO4 > HLiF2 > HCl. This clearly suggests superhalogens as building blocks of superacids. Subsequently, Czapla et al. [7] further explored this idea and reinvestigated whether the protonation of superhalogen anions always leads to superacids. They considered protonated MXk+1 − superhalogen anions for M = Li, Na, K, Be, Mg, Ca, B, Al, Ga, C, Si, Ge, and X = F, Cl. They found that HLiCl2 , HNaCl2 , HBeF3 , HBeCl3 , HMgF3 , HMgCl3 , HCaCl3 , HBF4 , HBCl4 , HAlF4 , HAlCl4 , HGaF4 , HGaCl4 , and HGeF5 have HF/HCl moieties and lower Gdep values than H2 SO4 . Therefore, they are indeed superacids, which is consistent with the previous findings of Srivastava and Misra [6]. On the contrary, HLiF2 , HNaF2 , HKF2 , HKCl2 , HCaF3 , HCF5 , HCCl5 , HSiF5 , HSiCl5 , and HGeCl5 are ionic or weakly bound complexes with higher Gdep values than H2 SO4 and therefore, they can not be considered as superacids. The possible explanation for these results was later given by Srivastava et al. [8].

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4.4 Origin of (Super)acidity Srivastava et al. [8] tried to explore the origin of the acidity of these compounds from a superhalogen perspective. They considered the strongest mineral acid H2 SO4 as a consequence of the protonation of HSO4 − anion, whose VDE is calculated to be 5.12 eV. Evidently, HSO4 − is itself a superhalogen anion. Subsequently, they chose BO2 − and ClO4 − superhalogen anions [9] having VDEs of 4.89 and 5.44 eV, respectively. The equilibrium structures of these anions and their protonated species are shown in Fig. 4.3 along with their VDE and Gdep values, respectively. They noticed that the origin of (super)acidity of the protonated complex is related to the superhalogenity of anions. For instance, the VDE of BO2 − is lower than that of HSO4 − and consequently, the Gdep of HBO2 is higher than that of H2 SO4 . The converse is true in the case of ClO4 − and HClO4 , i.e., higher gas-phase acidity of protonated species appears as a consequence of higher VDE of the anions involved. To further validate this, they designed a Bn H3n+1 − series of polynuclear superhalogen anions for n = 1–5 as displayed in Fig. 4.4a. The VDE of these anions

Fig. 4.3 Equilibrium structures of X− and HX for X = HSO4 , BO2 , and ClO4 from Ref. [8] with permission of Royal Society of Chemistry, copyright 2016

4.5 Path for Stronger Superacids

37

increases successively with an increase in n, lying in the range 4.62–7.61 eV. Note that the VDE of BH4 − (4.62 eV) is lower than that of HSO4 − (5.12 eV) and consequently, the Gdep of its protonated complex should be higher than that of H2 SO4 . The authors designed the protonated complex of these anions, Bn H3n+2 for n = 1–5. The Gdep of BH5 was calculated to be 323.8 kcal/mol, which was indeed higher than that of H2 SO4 . On the contrary, the Gdep of B2 H8 (292.7 kcal/mol) was lower than that of H2 SO4 due to the higher VDE of B2 H7 − (5.71 eV) than that of HSO4 − . The Gdep of Bn H3n+2 species goes on decreasing from 276.0 kcal/mol (n = 3) to 258.8 kcal/mol (n = 5). This led the authors to propose a linear correlation between the VDEs of Bn H3n+1 − anions and the acidity (i.e., reciprocal of Gdep ) of Bn H3n+2 series plotted in Fig. 4.4b. One can see that the acidity of Bn H3n+2 species is directly proportional to the superhalogenity of Bn H3n+1 − anions. This relation indirectly indicates the dependence of Gdep on n, i.e., it is possible to design even stronger superacids by increasing the value of n further. The statistical relation can also be used to predict the acidity of the Bn H3n+2 series for any value of n, in principle.

4.5 Path for Stronger Superacids From the above discussion, it is clear that the concept of polynuclear superhalogens is an effective strategy to design even stronger superacids. In 2015, Czapla and Skurski [10] already reported HInx F3x+1 , HSnx F4x+1 , and HSbx F5x+1 series of compounds by protonation of Inx F3x+1 − , Snx F4x+1 − and SbF5n+1 − anions, respectively for x = 1–3. The structures of all these species are similar in the sense that there is an HF moiety for x = 1 but H interacts with two F atoms for x = 2 and 3. For a visual indication, the equilibrium structures of HSbx F5x+1 species are displayed in Fig. 4.5. The VDEs of anions are 9.93–11.98 eV for Inx F3x+1 − , 10.25–12.56 eV for Snx F4x+1 − and 10.15–13.27 eV for Sbx F5x+1 − , making them polynuclear superhalogens. They calculated Gdep values of 273.6–268.9 kcal/mol for HInx F3x+1 and 275.6–243.5 kcal/mol for HSnx F4x+1 . Evidently, these all are superacids. More importantly, they obtained Gdep of 260.5–230.3 kcal/mol for the HSbx F5x+1 series, making them one of the strongest superacids reported so far. This can be expected due to enormously high VDEs of anions involved as mentioned earlier. Srivastava et al. [11] used the concept of hyperhalogen (see Chap. 1) to design stronger superacids. They chose M(M’F4 )4 − (M, M’ = B, Al) having VDE of exceeding 10 eV. The equilibrium structures of Al(BF4 )4 − and Al(AlF4 )4 − anions along with their protonated species are displayed in Fig. 4.6. All protonated species HM(M’F4 )4 are found to be superacids having Gdep lower than 300 kcal/mol. In particular, the Gdep of HAl(BF4 )4 superacid was found to be 236.4 kcal/mol with Al(BF4 )4 − − polynuclear superhalogen anion of VDE 10.52 eV. This is stronger in acidity than the HAlF3n+1 series of superacids, reported earlier (see Fig. 4.1).

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4 Superhalogens in the Design of Superacids

Fig. 4.4 Equilibrium structures of Bn H3n+2 (a) and the correlation between the acidity of Bn H3n+2 and VDE of Bn H3n+1 − anions (b) from Ref. [8] with permission of Royal Society of Chemistry, copyright 2016

4.5 Path for Stronger Superacids

39

Fig. 4.5 Equilibrium structures of HSbx F5x+1 for x = 1–3 from Ref. [10] with permission of American Chemical Society, copyright 2015

Fig. 4.6 Equilibrium structures of Al(BF4 )4 − and Al(AlF4 )4 − anions along with their protonated species from Ref. [11] with permission of Elsevier, copyright 2017

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4 Superhalogens in the Design of Superacids

4.6 Other Superhalogen-Based Superacids The exploration of superacids was made using the protonation of other superhalogen anions. Zhao et al. [12] utilized carborane-based anions, CBn HXn − (X = H, F, Cl, CN; n = 5, 7, 9, 11, 13). These anions behave as superhalogens due to their enhanced stability according to the Wade–Mingos rule as discussed in Chap. 1. The authors designed and studied a total of 63 complexes by protonation of these anions, some of which are displayed in Fig. 4.7. They calculated Gdep of these complexes, which are plotted as Gacid in Fig. 4.8. The authors noticed that 60 out of 63 complexes become superacids having Gdep (Gacid ) lower than 300 kcal/mol. They found that H/[CB11 HF11 ], H/[CB13 HF13 ], H/[CB11 F12 ], and H/[CB13 F14 ] had Gacid in the range 215–220 kcal/mol. Note that the H/[CB11 HF11 ] is the strongest organic superacid synthesized experimentally. Subsequently, Luo et al. [13] performed a similar study and used silaboranebased superhalogen anions, SiBn Hm Xn-m+1 − (X = H, F, Cl, CN). They studied their protonated complexes in the gas phase as well as in solvents. All these complexes,

Fig. 4.7 Equilibrium structures of protonated carborane-based anions from Ref. [12] with permission of Royal society of chemistry, copyright 2018

4.6 Other Superhalogen-Based Superacids

41

Fig. 4.8 The Gacid of protonated carborane from Ref. [12] and silaborane-based anions from Ref. [13] with permission of royal society of chemistry, copyright 2018 and 2019

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4 Superhalogens in the Design of Superacids

except two, were found to be superacids. Their Gacid values are also plotted for comparison in Fig. 4.8. As per these calculations, the silaborane-based complexes are relatively more acidic than those of carboranes. For instance, the Gacid of H/ [SiB12 HF12 ] is found to be 205 kcal/mol, which is lower than that of H/[CB11 F12 ] by ∼10 kcal/mol. The authors also noticed that the acidities of silaborane-based complexes in solution are also larger than those of carborane-based species, similar to the gas phase. They recommended that for a complete description of acidities, both the values of gas phase Gacid and pK a values in solution should be reported. The authors also found that the species with CN derivatives gains extra stabilization in the solvent and hence, possesses higher Gacid as compared to F/Cl derivatives. According to them, although (bond) dissociation enthalpies play a significant role in some species, the VDE becomes the dominant and crucial parameter for the acidities of protonated superhalogens, at least in the gas phase. In 2019, Zhao et al. [14] also used some conjugated anions such as CB2 H3 − , C5 H5 − , BC5 H6 − , etc., and substituted F, CN, and NO2 . Subsequently, they designed their protonated complexes as shown in Fig. 4.9. Note that these all anions are aromatic following the Hückel rule (see Chap. 1). The substitution of CN and NO2 leads to an increase in the VDE of anions beyond 5 eV. Therefore, 10 out of 20 anions considered were found to be superhalogens. Consequently, their protonated complexes have lower Gdep than 300 kcal/mol. The authors highlighted H/[BC5 (NO2 )6 ], H/a-[BC9 (NO2 )8 ], H/a-[BC9 (CN)8 ], H/b-[BC9 (NO2 )8 ] and H/[BC5 (CN)6 ] as five strongest superacids in their study having Gdep values in the range 243.7–248.1 kcal/mol. The authors highlighted that the design of organic superacids by combining an organic superhalogen and the proton is a rational route as proposed by Srivastava and Misra [6] for inorganic species, not a coincidence.

4.7 Summary In this Chapter, we have discussed the role of superhalogens in the design of superacids. Motivated by the fact that the strongest superacid, namely, HSbF6 contains SbF6 − superhalogen anion, a hypothesis has been made that the superacids can be considered as protonated complexes of superhalogen anions. This hypothesis has been verified using various typical superhalogens, polynuclear superhalogens, carborane, and silaborane-based superhalogen anions. In the case of polynuclear superhalogen anions, the acidity was found to be increasing monotonically with the increase in the VDEs of anions. This was demonstrated by using Bn H3n+1 − superhalogens and the acidity of their protonated complexes, Bn H3n+2 . Although the results discussed in the Chapter were mainly based on the calculations in the gas phase, these can be easily rationalized in solutions. These results advocated that the protonation of superhalogen anions is an effective strategy to design superacids and higher VDEs of anions lead to stronger superacids.

References

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Fig. 4.9 Equilibrium structures of protonated conjugated anions from Ref. [14] with permission of Royal Society of Chemistry, copyright 2019

References 1. 2. 3. 4. 5. 6. 7. 8.

Hall NF, Conant JB (1927) J Am Chem Soc 49:3062–3070 Olah GA, Prakash GK, Sommer J (1979) Science 206:13 Olah GA, Prakash GK, Sommer J (1985) Superacids. Wiley, New York Czapla M, Skurski P (2015) Chem Phys Lett 630:1–5 Sikorska C, Skurski P (2012) Chem Phys Lett 536:34–38 Srivastava AK, Misra N (2015) Polyhedron 102:711–714 Czapla M, Anusiewicz I, Skurski P (2016) Chem Phys 465–466:46–51 Srivastava AK, Kumar A, Misra N (2017) New J Chem 41:5445–5449

44 9. 10. 11. 12.

4 Superhalogens in the Design of Superacids

Gutsev GL, Boldyrev AI (1981) Chem Phys 56:277–283 Czapla M, Skurski P (2015) J Phys Chem A 119:12868 Srivastava AK, Kumar A, Misra N (2017) J Fluor Chem 197:59–62 Zhao R-F, Zhou F-Q, Xu W-H, Li J-F, Li C-C, Li J-L, Yin B (2018) Inorg. Chem Front 5:2934–2947 13. Luo L, Zhou F-Q, Zhao R-F, Li J-F, Wu L-Y, Li J-L, Yin B (2019) Dalton Trans 48:16184–16198 14. Zhou F-Q, Zhao R-F, Li J-F, Xu W-H, Li C-C, Luo L, Li J-L, Yin B (2019) Phys Chem Chem Phys 21:2804–2815

Chapter 5

Superhalogens in the Design of Electrolytic Salts

Abstract Lithium-ion battery (LIB) powers modern electronic devices. A typical LIB consists of a cathode, an anode, and electrolytic salts dissolved in organic solvents. In this Chapter, the role of superhalogen in the design of electrolytic salts for LIBs has been discussed. Noting the fact that all commercial electrolytic salts are made up of superhalogen anions, new superhalogen anions can enrich the family of these salts. These new salts might increase the Li+ ion conduction with enhanced environmental safety and aqueous stability. Some of these claims have been verified experimentally. Keywords Superhalogens · Vertical detachment energy · Lithium-ion battery · Electrolytic salt · Li dissociation · Water affinity · Halogen free salts

5.1 Introduction Lithium-ion batteries (LIBs) have become the main energy source for portable electronics such as smartphones, tablets, laptops, etc. The simple diagram of a typical LIB is displayed in Fig. 5.1. It consists of two electrodes- an anode of graphite and a cathode of lithium transition metal oxide such as LiCoO2 , LiMn2 O2, etc. [1]. These electrodes are separated by an electrolyte that acts as an ionic medium. The electrolytes are made up of electrolytic salts dissolved in a polar aprotic solvent. These salts are nothing but Li complexes of some non-coordinating anions (see Chap. 1). The commercially available and frequently used electrolytic salts include LiBF4 , LiPF6 , LiClO4 , etc. In this Chapter, we will see how the concept of superhalogen can be exploited to design new electrolytic salts.

© The Author(s), under exclusive license to Springer Nature Switzerland AG 2023 A. K. Srivastava, Superhalogens, SpringerBriefs in Molecular Science, https://doi.org/10.1007/978-3-031-37571-2_5

45

46

5 Superhalogens in the Design of Electrolytic Salts

Fig. 5.1 A simple diagram of a typical lithium-ion battery (LIB)

5.2 Anions in Typical Electrolytic Salts As mentioned above, most common electrolytic salts such as LiBF4 and LiPF6 contain halogen, which is toxic, and hence, environmental safety remains a grave concern. This demands a search for alternative salts, which are safe and efficient to work as electrolytic salts in LIBs. In the second half of 2014, Giri et al. [2] reported that superhalogens might be useful in designing new electrolytic salts in LIBs. The idea behind this work was that in typical electrolytic salts, e.g., Li+ BF4 – and Li+ PF6 – , BF4 – and PF6 – are superhalogen anions satisfying MXk+1 – formula (see Chap. 1). In fact, the currently employed electrolytic salts are Li-complexes of superhalogen anions. These superhalogens are displayed in Fig. 5.2. The VDE of these superhalogen anions lies in the range 4.32–8.91 eV. Thus, the concept of superhalogen can be exploited to design new Li-complexes for prospective electrolytic salts.

5.3 New Electrolytic Salts Based on Halogen-Free Superhalogens Giri et al. [2] utilized other superhalogens such as BeF3 [3] and AuF6 [4]. Considering the fact that they do contain halogen, they also considered halogen-free superhalogens, e.g., NO3 [3], BH4 [5], B3 H8 , and CB11 H12 [6]. The equilibrium structures of these superhalogens are shown in Fig. 5.3. The VDE of their anions ranges from 4.22 to 8.86 eV. The electrolytic salts contribute a Li+ ion increasing the conductivity of the electrolytic medium. Further, these salts need to possess aqueous stability. In order to assess this, they calculated the Li+ binding energy (E Li+ ) and water affinity (WA) of these salts as listed in Table 5.1. One can see that the E Li+ and WA of typical electrolytic salts are found to lie in the range 5.65–7.38 eV and 0.99–1.41 eV, respectively. Similarly, the E Li+ and WA of the proposed salts were calculated to be 5.08–6.62 eV and 0.92–1.08 eV, respectively. The authors particularly highlighted

5.3 New Electrolytic Salts Based on Halogen-Free Superhalogens

47

Fig. 5.2 Equilibrium structures of commonly used anions from Ref. [2] with permission of Wiley, copyright 2014

LiCB11 H12 as a potential candidate due to its lower E Li+ value, which ensures its easy ionic dissociation. However, it is carbonaceous, unlike other electrolytic salts which are mostly inorganic. Srivastava and Misra [7] proved that the properties of LiBeF3 are similar to those of LiBF4 not only in the gas phase but also in the organic solvent such as diethyl ether. In fact, they were first with the idea of utilizing superhalogens in LIBs. In the first half of 2014, they communicated a paper entitled “Theoretical investigations on lithium-superhalogen (Li-X) complexes (X = LiF2 , BeF3 , BF4 , PF6 ): Searching new electrolytic salts for lithium-ion batteries” to the New Journal of Chemistry. This work was rejected based on reviewers’ suggestions. Subsequently, they communicated the same paper to Molecular Physics in the second half of 2014. Surprisingly, one of the reviewers emphasized that “this work really does not offer new ideas for ions to use in Li-ion battery systems” and suggested, removing “all the inappropriate rationale about ion batteries”. The completely modified version of the same paper was, then, published as mentioned above [7]. Later in 2016, Srivastava and Misra [8] reported a comprehensive study on the potential candidates for electrolytic salts. They employed all halogen-free superhalogen anions such as BO2 − [3], AlH4 − , TiH5 − , VH6 − [5] having VDE values in

48

5 Superhalogens in the Design of Electrolytic Salts

Fig. 5.3 Equilibrium structures of proposed anions for electrolytic salts by Giri et al. from Ref. [2] with permission of Wiley, copyright 2014

the range 4.23–4.72 eV as given in Table 5.1. The structures of these anions and their Li-salts are displayed in Fig. 5.4. The E Li+ and WA of these Li-salts were calculated to be 6.00–6.36 eV and 0.90–0.99 eV, respectively. These proposed electrolytic salts may be more suitable, being halogen-free inorganic salts. Considering the fact that the E Li+ values of proposed salts (see Fig. 5.4) may be a bit higher, they utilized polynuclear Bn H3n+1 − superhalogen anions (see Chaps. 1 and 4) for n = 2–5 having VDEs in the range 5.52–7.28 eV as listed in Table 5.1. The Licomplexes of these anions are shown in Fig. 5.5. Note that LiBH4 was already studied by Giri et al. [2] with E Li+ and WA values of 6.62 eV and 0.92 eV, respectively (see Table 5.1). The E Li+ and WA of LiBn H3n+1 are found to be 6.34–5.45 eV and 0.93– 0.91 eV for n = 2–5. These parameters make them more favorable for electrolytic salts. The authors also noticed that the E Li+ of LiBn H3n+1 salts decreases with the increase in n as plotted in Fig. 5.6, leaving the WA values almost unaltered (see Table 5.1). Therefore, it might be possible to obtain an even lower value of E Li+ by substituting more BH4 superhalogen moieties in place of H atoms.

5.3 New Electrolytic Salts Based on Halogen-Free Superhalogens Table 5.1 Vertical detachment energy (VDE) of X− , Li+ -dissociation energy (E Li+ ), and water affinity (WA) of LiX salts at the wB97xD level from Ref. [2, 8, 9]. All values are in eV

LiX salts

VDE of X−

49 E Li+

WA

LiFePO4

4.32

7.38

1.04

LiClO4

5.83

5.96

1.02

LiN(SO2 F)2

6.89

5.82

1.02

LiN(SO2 CF3 )2

7.01

6.01

0.99

LiBF4

7.66

6.08

1.41

LiPF6

8.55

5.73

1.07

LiAsF6

8.91

5.65

1.09

LiNO3

4.22

6.53

0.96

LiBH4

4.50

6.62

0.92

LiB3 H8

4.72

6.25

0.93

LiCB11 H12

5.99

5.08

1.08

LiBeF3

6.99

6.50

0.98

LiAuF6

8.86

5.50

1.06

LiVH 6

4.23

6.29

0.91

LiBO2

4.44

6.36

0.99

LiAlH4

4.69

6.00

0.97

LiTiH5

4.72

6.11

0.90

LiB2 H7

5.52

6.34

0.92

LiB3 H10

6.25

6.15

0.91

LiB4 H13

6.64

5.80

0.91

LiB5 H16

7.28

5.45

0.93

LiB(OH)4

4.63

7.31

1.09

LiPO3

5.03

6.25

0.97

LiCB4 H5

3.40

5.94

0.99

LiCB5 H6

3.89

5.75

1.04

LiCB6 H7

4.19

5.51

1.17

LiCB7 H8

3.86

5.81

0.99

LiCB8 H9

4.30

5.30

1.23

LiCB9 H10

5.62

5.19

1.06

LiCB10 H11

4.66

5.28

1.12

LiCB12 H13

4.31

4.98

1.07

LiCB13 H14

5.06

4.98

1.10

LiSiB4 H5

3.20

5.91

0.97

LiSiB5 H6

4.18

5.84

1.01

LiSiB6 H7

4.23

5.62

1.08

LiSiB7 H8

3.80

5.57

1.06

LiSiB8 H9

4.77

5.39

1.04 (continued)

50 Table 5.1 (continued)

5 Superhalogens in the Design of Electrolytic Salts

LiX salts

VDE of X−

E Li+

WA

LiSiB9 H10

4.69

5.39

1.09

LiSiB10 H11

4.50

5.23

1.08

LiSiB11 H12

6.34

5.16

1.06

LiSiB12 H13

4.26

5.08

1.06

LiSiB13 H14

5.31

5.02

1.07

Fig. 5.4 Equilibrium structures of proposed anions corresponding Li-salts by Srivastava and Misra from Ref. [8] with permission of Elsevier, copyright 2016

Subsequently, Sun et al. [9] used BO2 − , B(OH)4 − , PO3 − along with carborane (CBn Hn+1 − ) and silaborane (SiBn Hn+1 − ) anions for n = 4–13 using various methods. Note that CB11 H12 − (Giri et al.) and BO2 − (Srivastava and Misra) anions have already been considered. The VDE of B(OH)4 − [10] and PO3 − [3] are 4.63 eV and 5.03 eV, respectively, making them superhalogens. CBn Hn+1 − and SiBn Hn+1 − anions are designed by borane Bn+1 Hn+1 2− dianions by replacing one B atom with C and Si atom, respectively for n = 4–13. The VDE of these anions is also listed in Table 5.1. One can see that most of them behave as superhalogens, which can be understood on the basis of the Wade–Mingos rule (see Chap. 1). The E Li+ value of LiB(OH)4 and LiPO3 are found to be 7.31 eV and 6.25 eV, respectively, which are a bit higher than the commercial electrolytic salts. However, the authors suggested that the Li-salts of carborane and silaborane for n ≥ 9 can be more suitable candidates for electrolytic salts. It should, however, be noticed that their E Li+ values do not follow the trend of the VDE values of their anions, like Bn H3n+1 − (Fig. 5.6).

5.4 New Electrolytic Salts Based on Superhalogen Dianions

51

Fig. 5.5 Equilibrium structures of LiBn H3n+1 complexes (n = 2–5) from Ref. [8] with permission of Elsevier, copyright 2016

5.4 New Electrolytic Salts Based on Superhalogen Dianions The stability of borane cages follows the Wade–Mingos rule as mentioned in Chap. 1. A typical example includes stable B12 H12 2− dianion. The stability of B12 H12 2− can be further increased by the substitution of halogen (F, Cl, Br, etc.) in the place of H atoms [11]. In 2016, Zhao et al. [12] substituted CN moieties in the place of H and reported B12 (CN)12 2− superhalogen anion. The first and second EAs of B12 (CN)12 were found to be 8.56 and 5.28 eV, respectively. They also replaced one B atom with a C atom to form CB11 (CN)12 having first and second EAs of 8.72 and 1.07 eV, respectively. Subsequently, they studied Li-salts of these dianions, namely Li2 B12 (CN)12 and Li2 CB11 (CN)12 as displayed in Fig. 5.8. The binding energies of the first and second Li+ ions in Li2 B12 (CN)12 are found to be 4.66 and 6.83 eV, respectively. these values are comparable to those of other electrolytic salts (see Table 5.1). Note that LiCB11 H12 has already been reported to be better electrolytic salt due to its lower E Li+ value (5.08 eV, see Table 5.1). The E Li+ value of LiCB11 (CN)12 is even lower, i.e., 4.09 eV. Further, CB11 (CN)12 2− is a stable dianion, and therefore, it can form stable Li2 CB11 (CN)12 salt whose E Li+ values are 4.67 and 6.89 eV for the first and second Li+ ion, respectively. The authors also showed its applicability for

52

5 Superhalogens in the Design of Electrolytic Salts

Fig. 5.6 The VDE of Bn H3n+1 − superhalogen anions and E Li+ of LiBn H3n+1 salts for n = 1–5 from Ref. [8] with permission of Elsevier, copyright 2016

magnesium ion batteries (MIBs) as well. Subsequently, Fang and Jena [13] reported a similar study using B12 (SCN)12 2− dianion and its derivative, namely CB11 (SCN)12 − superhalogen anion. They noticed that the Li-salts of these anions possess promising properties to be electrolytic salts, particularly for MIBs.

5.5 Summary In this Chapter, we have described how superhalogens can be used to design alternative electrolytic salts for LIBs. These salts contribute Li+ ions by dissociating into ionic fragments. It was noticed that the commercial electrolytic salts used in LIBs are made up of superhalogen anions. Therefore, new superhalogens can be introduced to modify and enrich these salts. Several theoretical studies were discussed which emphasized that better electrolytic salts can be designed with the choice of appropriate superhalogens. These salts may be halogen-free and hence, less toxic as well as with lower Li+ binding energies and hence, easy dissociation. In addition, they all exhibit remarkable stability in aqueous environment due to their low water affinity. It was also noticed that the use of polynuclear superhalogens with higher VDE values leads to lower Li+ binding energies without affecting its aqueous stability. A

5.5 Summary

53

Fig. 5.7 Equilibrium structures of Li-complexes of carborane and silaborane anions from Ref. [9] with permission of authors, copyright 2016

54

5 Superhalogens in the Design of Electrolytic Salts

Fig. 5.8 Equilibrium structures of Li2 B12 (CN)12 (a) and Li2 CB11 (CN)12 (b) salts from Ref. [12] with permission of Wiley, copyright 2016

few superhalogen dianions have also been recommended to design prospective electrolytic salts. Some of these predictions have already been confirmed experimentally for lithium and magnesium ion batteries [14, 15].

References 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15.

Dominey LA (1994) Lithium Batteries. Elsevier Science B. V, Netherlands Giri S, Behera S, Jena P (2014) Angew Chem Int Ed 53:13916–13919 Gutsev GL, Boldyrev AI (1981) Chem Phys 56:277–283 Gutsev GL, Boldyrev AI (1983) Chem Phys Lett 101:441–445 Boldyrev AI, von Nissen W (1991) Chem Phys 155:71–78 Pathak B, Samanta D, Ahuja R, Jena P (2011) ChemPhysChem 12:2423–2428 Srivastava AK, Misra N (2015) Mol Phys 113:866–870 Srivastava AK, Misra N (2016) Polyhedron 117:422–426 Sun Y-Y, Li J-F, Zhou F-Q, Li J-L, Yin B (2016) Phys Chem Chem Phys 18:28576–28584 Swierszcz I, Anusiewicz I (2011) Chem Phys 383:93 Boere RT, Derendorf J, Jenne C, Kacprzak S, Kebler M, Riebau R, Riedel S, Roemmele TL, Rεhle M, Scherer H, Vent-Schmidt T, Warneke J, Weber S (2014) Chem Eur J 20:4447–4459 Zhao H, Zhou J, Jena P (2016) Angew Chem 128:3768–3772 Fang H, Jena P (2017) J Phys Chem C 121:7697–7702 Tutusaus Q, Mohtadi R, Arthur TS, Mizuno F, Nelson EG, Sevryugina YV (2015) Angew Chem Int Ed 54:7900–7904 Tang WS, Unemoto A, Zhou W, Stavila V, Matsuo M, Wu H, Orimo S, Udovic TJ (2015) Energy Environ Sci 8:3637–3645

Chapter 6

Miscellaneous Applications of Superhalogens

Abstract The applications of superhalogens as strong oxidizers and in the design of superacids and electrolytic salts have been widely studied. However, the applications of superhalogens could be endless. In this Chapter, some miscellaneous applications of superhalogens have been discussed, which are less explored to date. It has been shown that typical organic superconductors (OSCs) and ionic liquids (ILs) comprise superhalogen anions. Therefore, novel OSCs and ILs can be designed using the concept of superhalogens. Superhalogens can also be used in the design of new liquid crystals and nonlinear optical materials as discussed herein. Keywords Superhalogens · Applications · Superconductors · Ionic liquids · Liquid crystals · Nonlinear optical materials

6.1 Introduction The applications of superhalogens as strong oxidizers and in the design of superacids and electrolytic salts have been discussed in preceding Chapters. These applications attracted various groups of researchers worldwide. However, the scope of superhalogens is very broad, making its applications in diverse areas. It was noticed by our research group in the last few years. In this Chapter, we discuss some diverse applications of superhalogens, namely, in organic superconductors, ionic liquids, liquid crystals, and nonlinear optical materials.

6.2 Organic Superconductors In early 1980, Bechgaard [1, 2] discovered the first organic superconductor (OSC) consisting of tetramethyl tetraselenafulvalene (TMTSF) and PF6 molecules, (TMTSF)2 PF6 . This resulted in the synthesis of Bechgaard salts, (TMTSF)2 X, a new family of salts named after him [3, 4]. In (TMTSF)2 X, complex anion X is an electron accepter whereas TMTSF molecules act as an electron donor. An isolated © The Author(s), under exclusive license to Springer Nature Switzerland AG 2023 A. K. Srivastava, Superhalogens, SpringerBriefs in Molecular Science, https://doi.org/10.1007/978-3-031-37571-2_6

55

56

6 Miscellaneous Applications of Superhalogens

Fig. 6.1 Equilibrium structure of TMTSF molecule from Ref. [5] with permission of Royal Society of Chemistry, copyright 2017

TMTSF molecule possesses a boat-like conformation having a dihedral angle of 1050 as displayed in Fig. 6.1. In 2017, Srivastava et al. [5] noticed that the typical anions (X− ) in Bechgaard salts such as NO3 − , BF4 − , ClO4 − , PF6 − , etc. are actually superhalogens [6] (see Chap. 1), having VDE in the range, 4.10–8.22 eV. This makes the possibility of using other superhalogen anions to design new Bechgaard salts. Considering other mononuclear such as BO2 − [6], BH4 − [7], AuF6 − [8] and binuclear B2 F7 − superhalogen anions, whose VDE lies in the range 4.35–8.53 eV, the authors designed and studied (TMTSF)2 X complexes as displayed in Fig. 6.2. In (TMTSF)2 X complexes, X is sandwiched between two TMTSF molecules. They noticed that the binding energy of (TMTSF)2 dimer with X in typical (TMTSF)2 X based OSCs ranges from 1.34 to 5.11 eV for X = NO3 to X = PF6 , respectively. Further, the dihedral angle (δ) is found to be increased from 1050 to 158–1730 . For X = BH4 , BO2 , AuF6 , and B2 F7 , the binding energies and δ of (TMTSF)2 X take the values from 0.52 to 6.31 eV and 1600 to 1800 , respectively. Therefore, it might be possible to enrich the family of Bechgaard salts using new superhalogens. It was already reported [9, 10] that in (TMTSF)2 X salts, the TMTSF molecules are deformed from boat-like conformation to become quasi-planar such that better OSCs possess larger stabilization of boat conformation of neutral TMTSF donors and hence, lower energy of deformation. Srivastava et al. also calculated this deformation energy in the range 0.12–0.18 eV for typical OSCs. The authors noticed that all proposed salts, except (TMTSF)2 AuF6 , have deformation energy comparable to or lower than typical OSCs. They, therefore, concluded that the proposed salts may be potential candidates for new OSCs.

6.3 Ionic Liquids Ionic liquids (ILs) are salts of large asymmetric organic cations and complex anions. 1-butyl-3-methylimidazolium (BMIM) is one of the commonly used imidazolium cations as displayed in Fig. 6.3. The ILs with simple anions such as halides are less stable [11, 12], and therefore, ILs consists of complex anions such as NO3 − , BF4 − , PF6 − and so forth. Evidently, all these complex anions are superhalogens [6] as noticed by Srivastava et al. [13] not earlier than 2021. Consequently, they used other

6.3 Ionic Liquids

57

Fig. 6.2 Equilibrium structures of (TMTSF)2 X complexes from Ref. [5] with permission of Royal Society of Chemistry, copyright 2017

58

6 Miscellaneous Applications of Superhalogens

Fig. 6.3 Molecular structure of 1-butyl-3-methylimidazolium (BMIM) from Ref. [13] with permission of American Chemical Society, copyright 2021

superhalogen anions such as BO2 , BeF3 [6], and LiF2 [14] and designed BMIM-X salts. Their objective was to investigate whether the properties of these new salts are similar to those of typical ILs. The equilibrium structures of BMIM-X salts for X = Cl, NO3 , BO2 , LiF2 , BeF3 , BF4 , and PF6 are shown in Fig. 6.4. All these salts are found to be energetically stable such that their binding energy increases with the increase in the EA of X, i.e., from NO3 to PF6 . Stable ILs are charge-transfer salts, which generally prefer to dissociate into ionic fragments. For instance, the dissociation of BMIM-X into (BMIM)+ + X− is energetically favored as compared to BMIM + X. The authors noticed that, except for X = Cl, all other BMIM-X salts preferably dissociate into ionic fragments, irrespective of superhalogen X. Due to charge transfer from BMIM to X, BMIM-X salts are highly polar and polarizable. This resulted in their high dipole moment of 11.2–13.3 a.u. and mean polarizability 118.2–127.3 a.u. These findings suggest the design of more stable ILs by exploiting the concept of superhalogens.

6.4 Liquid Crystals Liquid crystals (LCs) are soft matter with the order like crystals and fluidity like liquids. One of the popular LC series includes n-alkyl-n’-cyanobiphenyl (nCB) in which the biphenyl ring system is connected to the alkyl (Cn H2n+1 ) chain and cyano (CN) group at terminals as displayed in Fig. 6.5. Note that the EA of CN is larger than that of Cl. This is the reason that CN has been employed as a ligand to design several superhalogens as discussed in Chaps. 1 and 2. In 2017, Srivastava et al. [15] studied the functionalization of benzene by superhalogens and found that the properties of BO2 superhalogen [6] substituted phenyl ring are similar to those of CN substituted benzene. A few years later, Srivastava [16] considered one of the compounds of the nCB series, 2CB, and substituted BO2 in the place of CN, thus forming a new compound, 2BB. The kinetic stability of this novel compound was confirmed by molecular dynamics simulations.

6.4 Liquid Crystals

59

Fig. 6.4 Equilibrium structures of anions in BMIM-X ionic liquids from Ref. [13] with permission of American Chemical Society, copyright 2021 Fig. 6.5 Schematic structure of n-alkyl-n’-cyanobiphenyl (nCB) series

60

6 Miscellaneous Applications of Superhalogens

The molecular structures of 2CB and 2BB are displayed in Fig. 6.6. One can see that the structures of both 2CB and 2BB are quite similar. The length-to-breadth ratio (L/B) and ovality (O) are parameters, which determine the shape of LC compounds. For instance, the value of O greater than unity suggests the rod-like structures of LC compounds. The author calculated the L/B ratio and O of 2BB to be 2.71 and 1.21, respectively as compared to 2.62 and 1.20 in 2CB, respectively. This suggested that the 2BB possesses a rod-like structure just like 2CB. Several other parameters such as dipole moment, quadrupole moment trace, and anisotropy of 2BB are also found to be very close to those of 2CB. Therefore, a new LC series can be developed with the substitution of a suitable superhalogen. In a recent study, Kumar et al. [17] proposed nBB series of compounds substituting BO2 at one terminal and increasing alkyl chain length for n = 1–15. They calculated various geometric and electronic parameters of the nBB series for n = 1 to 15. It is well established that the nCB series of compounds exhibit the odd–even effect [18, 19], i.e., their parameters vary according to the odd and even values of alkyl chain length (n). The authors analyzed the variation in various parameters such as dipole moment, quadrupole trace, molecular anisotropy, etc. with the increase in the alkyl chain length (n) and found that the variation in dipole moment values shows a remarkable dependence on n. Therefore, they plotted the dipole moments of the nBB series as a function of n and compared them with those of the nCB series for n = 1–12 as shown in Fig. 6.7. They noticed that the variation in dipole moments of both series follows a similar trend and shows, the so-called, odd–even effect.

Fig. 6.6 Molecular structures of 2BB and 2CB liquid crystal compounds from Ref. [16] with permission of Elsevier, copyright 2021

6.5 Nonlinear Optical Materials

61

Fig. 6.7 Variation of dipole moments of nBB and nCB series for n = 1–12 from Ref. [17] with permission of Springer Nature, copyright 2023

6.5 Nonlinear Optical Materials Cyclic conjugated polymers including cyclic thiophene, pyrrole, and furan are well known for their applications in electronic devices [20]. In 2020, Sajid et al. [21] reported the use of superhalogens to enhance the nonlinear optical (NLO) response of cyclic oligofurans (nCF). They doped BeX3 (X = F, Cl) superhalogens in 5CF and 6CF molecules as displayed in Fig. 6.8 for BeF3 . One can see that the planarity of 5CF and 6CF molecules is disturbed by the interaction with superhalogen. These complexes are found to be stable due to high interaction energies (62–111 kcal/ mol). The authors calculated and analyzed various parameters such as first hyperpolarizability (β o ), hyper-Rayleigh scattering coefficient, electro-optical Pockels effect, second harmonic generation, and nonlinear refractive index to assess the NLO response of these complexes. The authors noticed that β o of BeX3 @nCF increases to 3 × 104 –6 × 104 a.u., which becomes as high as 3 × 105 a.u. for the BeF3 @6CF complex. Likewise, the second-order NLO response coefficient of BeCl3 @6CF is increased to 2 × 109 a.u. The remarkable NLO response of superhalogen-doped cyclic oligofuran complexes is due to the electron affinity of superhalogens which are capable of withdrawing electrons from oxygen atoms of nCF. This suggested that the doping of superhalogens can effectively enhance the NLO response. This leads to the use of superhalogens to increase the NLO properties of various other planar systems. In 2021, Ishaq et al. [22] studied superhalogen-doped borophene (B36 ). B36 is an all-boron atom inorganic analog of graphene as displayed in Fig. 6.9. Subsequently,

62

6 Miscellaneous Applications of Superhalogens

Fig. 6.8 Equilibrium structures of 5CF (a) and 6CF (b) molecules interacting with BeF3 from Ref. [21] with permission of Royal Society of Chemistry, copyright 2020

they doped one and two BeF3 and BF4 superhalogens in B36 also shown in Fig. 6.9 for BF4 . The doping results in bowl shape structure of B36 , consistent with the previous results of Sajid et al. [21]. These all structures were found to be stable. The authors found that the doping of single BF4 in B36 leads to an enormous increase in the NLO response, as the β o of BF4 @B36 becomes 2 × 105 a.u. They proposed that this may be due to its strong interaction energy of ~15 eV as compared to other complexes (1– 2 eV). Further studies in this direction were performed by Khan et al. [23] and Asif et al. [24]. These authors doped AlF4 superhalogen to increase the NLO response of boron nitride (B12 N12 ) [23] and aromatic heterocycles (C4 H4 NH, C4 H4 O, C4 H4 S, and C5 H5 N) [24].

6.6 Summary In this Chapter, we have discussed some miscellaneous applications of superhalogens. Using typical OSCs, which are Bechgaard salts, we noticed that the complex anions in these salts belong to the class of superhalogens. Therefore, new series of Bechgaard salts can be designed by using superhalogens. Similarly, we discussed typical imidazolium-based ILs, which are salts containing superhalogen anions, and found that new ILs can be designed with enhanced stability by incorporating other superhalogens. We have also explored the possibility of the design of new LC

References

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Fig. 6.9 Equilibrium structures of borophene (a) interacting with BeF3 (b) and BF4 (c) from Ref. [22] with permission of Springer Nature, copyright 2021

compounds by considering typical nCB series and substituting BO2 superhalogen in the place of the CN group. We noticed that nBB possess similar structures and exhibits similar odd–even effect as those of the nCB series. These less-known findings require further exploration from other researchers working in this field. The doping of superhalogens is found to enhance the NLO response of various systems.

References 1. Jerome D, Mazaud A, Ribault M, Bechgaard K (1980) J Phys Lett 41:95–98 2. N. Thorup, G. Rindorf, H. Soling and K. Bechgaard, Acta Crystallogr., Sect. B: Struct. Crystallogr. Cryst. Chem., 1981, 37, 1236–1240. 3. Bechgaard K, Jacobsen CS, Mortensen K, Pedersen HJ, Thorup N (1980) Solid State Commun 33:1119–1125 4. Parkin SSP, Ribault M, Jerome D, Bechgaard K (1981) J Phys C: Solid State Phys 14:5305–5326 5. Srivastava AK, Kumar A, Tiwari SN, Misra N (2017) New J Chem 41:14847–14850 6. Gutsev GL, Boldyrev AI (1981) Chem Phys 56:277–283 7. Boldyrev AI, von Nissen W (1991) Chem Phys 155:71–78 8. Gutsev GL, Boldyrev AI (1983) Chem Phys Lett 101:441–445 9. Demiralp E, Goddard WA III (1995) Synth Met 72:297–299 10. Demiralp E, Goddard WA III (1997) J Phys Chem A 101:8128–8131 11. Hallett P, Welton T (2011) Chem Rev 111:3508–3576 12. Poole CF, Poole SK (2010) J Chromatogr A 1217:2268–2286 13. Srivastava AK, Kumar A, Misra N (2021) J Phys Chem A 125:2146–2153 14. Gutsev GL, Bartlett RJ, Boldyrev AI, Simons J (1997) J Chem Phys 107:3867–3875 15. Srivastava AK, Kumar A, Misra N (2017) Chem Phys Lett 671:44–48 16. Srivastava AK (2021) J Mol Liq 344:117968 17. Kumar A, Srivastava AK, Sharma D, Tiwari SN, Misra N (2023) Theor Chem Acc 142:17 18. Dwivedi MK, Tiwari SN (2011) J Mol Liq 158:208–211 19. Kumar A, Srivastava AK, Tiwari SN, Misra N, Sharma D (2019) Mol Cryst Liq Cryst 681:23–31 20. S. Reineke, F. Lindner, G. Schwartz, N. Seidler, K. Walzer, B. Lüssem and K. Leo, 21. Nature, 2009, 459, 234–238. 22. Sajid H, Ullah F, Yar M, Ayub K, Mahmood T (2020) New J Chem 44:16358–16369 23. Ishaq M, Shehzad RA, Yaseen M, Iqbal S, Ayub K, Iqbal J (2021) J Mol Model 27:188 24. Khan AU, Muhammad S, Khera RA, Shehzad RA, Ayub K, Iqbal J (2021) Optik 231:16646 25. Asif M, Sajid H, Gilani MA, Ayub K, Mahmood T (2022) Vacuum 203:111301

Chapter 7

Conclusion and Future Perspectives

Abstract The various aspects of the design of superhalogens and their diverse applications have been systematically presented in the preceding chapters of the book. This Chapter concludes the book and offers some future directions. Keyword Superhalogens; Design; Applications; Progress; Scope; Future directions

In this book, I presented a brief account of historical survey and recent progress in the field of superhalogens. As mentioned in Chap. 1, the superhalogens were introduced in 1981 as the species (radicals) having higher EA than that of halogen. In 1999, the anions of these species were studied experimentally for the first time. Subsequently, the focus shifted from radicals to anions whose VDE exceeds that of halides. Evidently, EA and VDE do not differ for halogen or any other atom and its anion. But this is not the case for molecular species in which EA is smaller than VDE due to the structural relaxation upon the addition of an electron in the anion. In the cases where the EA of species is smaller than that of halogen but the VDE of its anion is larger than halide, making a dilemma to characterize it as superhalogen as highlighted recently [1]. There is no such rule whether it should be characterized as superhalogen or not. However, experience tells that the preference should go to the VDE values, as they can be directly measured from experiments. Apart from this, there is no universal theoretical approach to calculating the VDEs in the absence of experimental values. There is a need for general benchmarking for the calculations of reliable and universally acceptable VDE values. One of the most striking results discussed includes the ultimate limit of VDE value for any anion. In polynuclear anions, Mn Xkn+1 it has been noticed that the VDE can be increased successively by increasing the number of core atoms (n). To which extent, is a question, which is not yet fully answered. So far, the maximal value of VDE is limited to 14 eV [2]. The studies are invited in this direction. Another aspect (not discussed in the book) is that there may be a few molecules (e.g., FeO4 ) whose EA exceeds that of halogen. These have a closed-shell structure, but not their anion. The fate of these superhalogens has not been understood as they are unlikely to behave as halogens, i.e., to form

© The Author(s), under exclusive license to Springer Nature Switzerland AG 2023 A. K. Srivastava, Superhalogens, SpringerBriefs in Molecular Science, https://doi.org/10.1007/978-3-031-37571-2_7

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anion, salt, and dimer. These challenges in the research of superhalogens can be easily turned into opportunities for future researchers [3]. A typical superhalogen consists of a core with few ligands. The beginning phase of research started with the use of a suitable core, which is mostly, but not always, an electropositive atom. There are several examples (e.g., ClO4 ) in which this condition seems to be violated. This makes it difficult to define a core in superhalogen. The core should always be identified with ligands. To design a superhalogen, for instance, H-atom can be used as ligands with metal atoms such as Be, Mg, Ca, etc. and it can be employed as a core with halogen ligands such as F, Cl, Br, etc. As far as ligand is considered, this needs not to be an atom but a combination of two or more atoms such as CN, SCN, OF, BO, etc. The idea behind choosing an appropriate ligand is guided by the fact that ligands should be hungry enough for an electron to complete its shell. Therefore, ligand engineering for superhalogens contains further opportunities. The historical application of superhalogens as strong oxidizers in forming compounds with stable molecules and noble gases, which appeared two decades before the birth of superhalogens, spanned five decades. It has been found that superhalogens are capable of oxidizing molecules when their EAs are comparable to the IE of the molecules to be ionized as discussed in Chap. 3. This makes it possible to ionize, in principle, any molecular system with cleverly chosen superhalogen. This property may also be exploited to design superhalogen-stabilized compounds. The superhalogens with extremely high EAs should be capable to stabilize the compounds of noble gas atoms such as Ar, Kr, etc. The reports are awaited. As far as other applications of superhalogens are concerned, the interesting thing is that all of them have been explored only in the last decade as discussed in Chaps. 4 onwards and shown in Fig. 7.1. Most of these applications have been explored by two or three groups. The most recent applications on ionic liquids and liquid crystals were reported by a single group and no follow-up study has been carried out so far by other researchers. Needless to mention that there is a knowledge gap and consequently, a lot of scope in the applications aspects of superhalogens. Below I highlight a few points that should be taken as future perspectives. As mentioned in Chap. 4, the protonation of superhalogen anions is a rational route for the design of superacids. The acidity of protonated species increases monotonically with the increase in the VDEs of anions. The anions with higher VDEs lead to stronger superacids upon protonation. In this sense, the strongest superacid is yet to be explored and this is likely to be achieved with the maximal VDE value for any anion. Meanwhile, the experimentalists should attempt to synthesize some superacids already predicted using other superhalogen anions. I am very hopeful to see some progress in the near future. The use of superhalogens to design new electrolytic salts for LIBs seems a revolutionary idea (Chap. 5). These new salts may have tailored properties, i.e., low Li+ dissociation energy to facilitate easy ion conduction, low water affinity to provide high aqueous stability as well as less toxicity to increase environmental safety. According to theoretical calculations, the new electrolytic salts might possess all these properties with the choice of appropriate superhalogens in these salts. Further, these conclusions are equally valid for sodium and magnesium ion batteries. These claims were supported by some experimental studies as well.

7 Conclusion and Future Perspectives

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Fig. 7.1 Diagram showing applications of superhalogens explored in the last decade

The route is open to design and suggests new electrolytic salts using this strategy. The idea is yet to be commercialized. Chapter 6 discussed the applications of superhalogens in the design of organic superconductors, ionic liquids, liquid crystals, and NLO materials. These are the less explored to date and this is the reason they could be clubbed together in a single place. Evidently, there are numerous opportunities to work further on this. To mention, the application of superhalogens in organic superconductors has been demonstrated using Bechgaard salts but the idea can be extended to other families such as Fabre salts. Likewise, the conclusions based on imidazolium-based ionic liquids should apply to other ionic liquids. This makes an open invitation for researchers to work on these aspects. Thus, the field of superhalogen contains a bunch of opportunities for future researchers. One of the most important features of this field is its multidisciplinary nature. The research, so far, discussed in the preceding chapters creates an interesting picture as drawn in Fig. 7.2. The presence of superhalogen anions and their structural integrity in a variety of systems for entirely different applications might lead to some interconnections. To be specific, the existence of BF4 in HBF4 (for superacid, see Chap. 4), LiBF4 (for electrolytic salts, see Chap. 5), (TMTSF)2 BF4 (for organic superconductors, see Chap. 6), and BMIM-BF4 (for ionic liquids, see

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Fig. 7.2 Diagram showing the interdisciplinary nature of superhalogens

Chap. 6) establishes a common relation among apparently different fields. Therefore, the fields are all set for future researchers to come and create the bridge across the disciplines. I would like to finish it here with a quote, “If I have seen further than others, it is by standing upon the shoulders of giants.” by Sir Issac Newton.

References 1. Srivastava AK (2023) Chem Commun 59:5943–5960 2. Skurski P (2022) Superhalogens—enormously strong electron acceptors. In: Superatoms: principles, synthesis and applications, Jena P, Sun Q (eds). Wiley, USA. 3. Srivastava AK, Anusiewicz I, Velickovic S, Sun W-M, Gutsev GL (2022) Superhalogens and Superalkalis: exploration of structure, properties and applications. Frontiers Media SA, Laussane, Switzerland

Index

A Acceptors, 25 Acidity, 33–38, 42, 66 Adiabatic, 3, 14 Affinity, 1, 3, 14, 46, 49, 52, 61, 66 Anions, 1–11, 14–18, 20, 25–27, 33–43, 45–48, 50–53, 55, 56, 58, 59, 62, 65–67 Anisotropy, 60 Anode, 45 Applications, 55, 61, 62, 65–67 Aqueous, 45, 46, 52, 66 Aromatic, 10, 25, 27, 29, 42, 62 Au, 5, 9, 13, 15, 17, 19–21, 29

B BARF, 11 Batteries, 11, 45–47, 52, 54, 66 Bechgaard, 55, 56, 62, 67 Benzene, 25, 29, 58 Binding, 3–5, 15, 17, 19–21, 24–28, 46, 51, 52, 56, 58 Bonding, 8, 14–18, 29 Borane, 50, 51 Borophene, 61, 63 Br, 2, 5, 19, 51, 66 Buckminsterfullerene, 25

C Ca, 5, 35, 66 Carbon, 26, 29 Carborane, 40–42, 50, 53 Cathode, 45 Cations, 56

Charge, 15–17, 24–29, 58 Cl, 1, 2, 5, 13, 15, 19–21, 25, 35, 40, 42, 51, 58, 61, 66 Closed-shell, 3, 65 CN, 5, 6, 8, 13, 20, 21, 40, 42, 51, 58, 63, 66 Complexes, 2, 9, 11, 13, 15–18, 21, 23–30, 33–37, 40, 42, 45, 47, 51, 55–57, 61, 62 Compounds, 1, 10, 11, 23, 27, 29, 36, 37, 58, 60, 63, 66 Conduction, 45, 66 Conjugated, 10, 42, 43, 61 Coordinating, 10, 11 Core, 5, 8–10, 13–16, 19, 21, 65, 66 Cu, 14, 15, 27 Cyanobiphenyl, 58, 59

D Deformation, 56 D-electrons, 10, 15, 17, 18 Delocalization, 16, 17, 21 Deprotonation, 33, 34 Design, 1, 5, 9–11, 16, 21, 29, 33, 37, 42, 45, 46, 52, 54–56, 58, 62, 65–67 Detachment, 2, 3, 49 Devices, 4, 45, 61 Dianions, 50–52, 54 Dimerization, 5, 8 Dioxide, 26, 27, 29 Doping, 61–63

E Electrolytic, 11, 45–48, 50–52, 54, 55, 66, 67

© The Editor(s) (if applicable) and The Author(s), under exclusive license to Springer Nature Switzerland AG 2023 A. K. Srivastava, Superhalogens, SpringerBriefs in Molecular Science, https://doi.org/10.1007/978-3-031-37571-2

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70 Electron-counting, 10, 11 Electronegative, 2, 5 Excimer, 4 Experimental, 1, 3, 4, 9, 11, 19, 65, 66

F Fabre, 67 Fluidity, 58 Fluorides, 13–16, 19 Fragmentation, 14, 15, 17 Functionalization, 58

G Geometries, 3, 15, 27 Graphene, 61 Ground-state, 14

H Halogen, 1, 2, 11, 14, 17, 21, 29, 46, 51, 52, 65 Halogen-free, 46, 48, 52 Heterocycles, 23, 27, 29, 62 Hückel, 10, 42 Hyperhalogen, 9, 37 Hyperpolarizability, 61

I Imidazolium-based, 62, 67 Inert, 23 Inorganic, 8, 23, 29, 47, 48, 61 Ionic, 1, 12, 15, 20, 23–27, 29, 35, 45, 47, 52, 55, 56, 58, 59, 66, 67 Ionization, 1, 2, 4, 9, 23, 25, 26, 28 Isomers, 25, 28, 29

K Kinetic, 3, 4, 58

L LIB, 45, 46 Li-complexes, 46, 48, 53 Ligands, 1, 2, 5–9, 11, 13–17, 19–21, 58, 66 Liquids, 34, 55, 56, 58–60, 66, 67 Li-salts, 48, 50–52 Lithium-ion, 11, 45–47 Localization, 17

Index M Magnesium, 52, 54, 66 Materials, 55, 61, 67 Metal, 8, 10, 11, 13–16, 19, 21, 23, 27, 29, 45, 66 Mineral, 33, 34, 36 Molecule, 1, 2, 9, 14, 15, 17, 18, 23, 25, 26, 29, 30, 55, 56, 61, 62, 65, 66 Multiplicities, 14, 17

N Negative, 1, 14, 16, 25, 27 Nitric, 28, 29 Noble, 23, 29, 66 Non-coordinating, 1, 10, 11, 45 Nonlinear, 55, 61 Novel, 5, 55, 58

O Octet, 2, 10 Odd-even, 60, 63 Oligofuran, 61 Optical, 55, 61 Orbital, 20, 29 Organic, 5, 23, 29, 40, 42, 45, 47, 55, 56, 67 Organic Superconductors (OSCs), 55, 56, 62 Ovality, 60 Oxidation, 1, 8, 13–15, 17, 19, 21, 23–28 Oxides, 13, 16, 19, 23, 26–29, 45 Oxidizers, 1, 23, 29, 55, 66

P Pd, 16, 17, 29 Permanganate, 11 Phase, 19, 34–36, 40, 42, 47, 66 Photodetachment, 4 Photoelectron, 3, 4, 9 Photon, 3–5 Polarizability, 58 Polynuclear, 1, 9, 11, 26, 27, 34, 36, 37, 42, 48, 52, 65 Protonation, 33–37, 40, 42, 66 Pseudohalogens, 5

Q Quadrupole, 60 Quasi-planar, 56

Index R Radicals, 2, 65 Rational, 33, 42, 66 Reactions, 5, 19 Relaxation, 3, 65 Response, 61–63 Rule, 1, 2, 10, 11, 40, 42, 50, 51, 65 S Salts, 11, 12, 15, 21, 45–52, 54–56, 58, 62, 66, 67 Series, 5, 26, 36, 37, 58–63 Silaborane, 40–42, 50, 53 Species, 1, 2, 5, 13–21, 33, 35–37, 39, 42, 65, 66 Spectroscopy, 3, 9 Stability, 2, 10, 14, 17, 23, 24, 29, 34, 40, 45, 46, 51, 52, 58, 62, 66 Stable, 2, 3, 5, 19, 23, 25, 26, 29, 34, 51, 56, 58, 61, 62, 66 Strongest, 33, 34, 36, 37, 40, 42, 66 Structures, 3, 5, 8, 11, 14–17, 21, 24–30, 33–40, 43, 46–48, 50, 51, 53, 54, 56–60, 62, 63, 65 Substituted, 42, 51, 58 Superacids, 33–35, 37, 40, 42, 55, 66, 67 Superconductors, 55, 67 Superhalogenity, 36, 37 Superhalogens, 1–17, 19–21, 23–30, 33–37, 40, 42, 45–48, 50–52, 54–56, 58, 60–63, 65–68

71 T Tendency, 5, 15, 17, 29 Theoretical, 1, 2, 4, 9, 11, 19, 47, 52, 65, 66 Thiophene, 27, 61 Trifluorides, 29 Typical, 2, 17, 34, 42, 45, 46, 51, 55, 56, 58, 62, 63, 66

U Unstable, 5, 14, 27

V Valency, 8, 16, 21 Variable, 8, 13, 21 Vertical, 2, 3, 49

W Wade-Mingos, 10, 40, 50, 51 Weakly, 10, 25–27, 35 Withdrawing, 5, 61

X Xe, 23, 29