Chemical properties of inorganic compounds of s-elements: educational-methodological handbook 9786010445246

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AL-FARABI KAZAKH NATIONAL UNIVERSITY

O.I. Ponomarenko I.V. Matveyeva

CHEMICAL PROPERTIES OF INORGANIC COMPOUNDS OF S-ELEMENTS Educational-methodological handbook

Almaty «Qazaq University» 2020

UDC 54(075.8) LBC 24.2я73 Р 82 It is recommended for publication by the Academic Council of the Faculty of Chemistry and Chemical Technology and RISO of al-Farabi Kazakh National University (Protocol №4 dated 16.04.2019)

Reviewer PhD Yarovaya Yelena

Р 82

Ponomarenko O.I. Chemical properties of inorganic compounds of s-elements: educational-methodological handbook / O.I. Ponomarenko, I.V. Matveyeva. – Almaty: Qazaq University, 2020. – 112 p. ISBN 978-601-04-4524-6 This educational-methodological handbook presents theoretical material on the basic physical, chemical properties of the main compounds of s-elements, their methods of preparation, and the most important areas of application. The educational-methodological handbook is intended for students of specialties «6B05301 – Chemistry», «6B07103 – Chemical technology of inorganic substances» and «6B07102 – Chemical engineering».

UDC 54(075.8) LBC 24.2я73 © Ponomarenko O.I., Matveyeva I.V., 2020 © Al-Farabi KazNU, 2020

ISBN 978-601-04-4524-6

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LIST OF ABBREVIATIONS

conc. ct. dil. IP liq.

– – – – –

concentrated; catalyst; diluted; ionization potentials; liquid.

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PREFACE

Chemical industry plays an important role in the development of the economy of any country. Production of inorganic alkalis, acids, and their salts, explosives, varnishes, and paints is about 64% of the chemical industry of Kazakhstan. For the further successful development of the chemical industry, training of qualified personnel is necessary. Without basic knowledge of the main physical, chemical properties of various compounds, methods of their preparation, the current development of the chemical industry is impossible. The current development trends in Kazakhstan have necessitated the publication of educational-methodological handbooks aimed at an accessible presentation of the existing knowledge base in the field of inorganic compounds. This educational-methodological handbook contains material in accordance with the curriculum for the teaching of specialists in the following specialties: «6B05301 – Chemistry», «6B07103 – Chemical technology of inorganic substances» and «6B07102 – Chemical engineering». This educational-methodological handbook summarizes the properties of s-elements and their compounds, depending on the structure of their atoms, as well as methods of obtaining these compounds (including industrial), their physical and chemical properties, areas of practical application.

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PROPERTIES OF ELEMENTS OF THE 1ST GROUP

Alkali elements (3Li, 11Nа, 19K, 37Rb, 55Cs, 87Fr) are located directly behind the inert gases in the periodic system, and their electrons are located at a new energy level, starting the electron layer with the main quantum number one unit larger than the elements of the previous period. These elements are preceded by a completed electron shell of an inert gas type, therefore, the valence electrons of alkaline elements are split off more easily than any other element of the same period; the electron layer is still far from completion and therefore not stable. However, the magnitudes of the ionization potentials (IP1) for the metallic state are large. This applies primarily to lithium, for which IP1 is equal to 5.37 eV. With an increase in the atomic and ionic radii, the IP values from top to bottom in the group decrease. Cesium has the lowest IP1 among all elements of the periodic system (3.58 eV). The valence possibilities of alkali metals are not too diverse. Among them are a metallic state (oxidation state 0) and a monovalent state (oxidation state +1). The ionic state M+ is most characteristic of alkali metals due to the relatively low IP value. Alkali metals play a large role in the geochemical and biological life of the Earth. Sodium and potassium are among the essential metals and the most common, and Li, Rb and especially Cs are rare, but also exhibit biological activity. All alkali metals are lithophilic, i.e. they are chemical elements that make up about 93% of the mass of the Earth's crust and about 97% of the mass of the salt composition of the global ocean. Alkali metals are found in nature in the form of singly charged cations. Alkali metals are also found in the form of primary minerals and are an integral part of (sometimes the main) sedimentary rocks. They are mainly contained in the upper layers of the Earth, they are few in the mantle of Earth and not at all in the core. In the lithosphere, alkaline elements are mainly in the form of aluminosilicates. The following sodium and potassium-containing minerals of sedimentary origin are considered main: 5

NаCl – rock salt, halite; NаCl·KCl – sylvinite; KCl·MgCl2·6H2O – carnallite; Nа2SO4·10H2O – mirabilite; Nа2CO3·10Н2O – natron; Nа2SO4·MgSO4·4H2O – astrakhanite; K2SO4·MgSO4·CаSO4·2H2O – polyhalite. Lithium, rubidium, and cesium are rare elements. They are rich in aluminosilicates, many of which are micas and have a layered structure: М2[Аl2Si3O9(F,ОН)], where M = K, Li, Rb, Cs. All alkali metals are fusible, and the melting point regularly decreases from Li to Cs. All alkali metals (except golden yellow cesium) are soft, silvery white. The toughest of alkali metals is lithium, but it is also cut with a knife and has a grayish-white color on the cut. The chemical activity of alkali metals is unparalleled among other metals, so they are usually stored in kerosene, sealed in sealed iron boxes. In the air, alkali metals are quickly covered with a film of complex composition, in which there are oxides, nitrides, carbonates, etc. Vapors of alkali metals (simple substances) and complex compounds of alkali metal have characteristic staining: Li – carmine red, Na – yellow, K – purple-pink, Rb – whitish-pink, Cs – purple-pink. Alkali metals are able to react with most simple and complex substances, while showing the properties of the strongest reducing agents.

Complex compounds of alkaline elements A feature of all complex compounds of alkaline elements is their ionic character. Being the most electropositive elements of the periodic system, alkali metals, entering the composition of heteroatomic (complex) compounds, even with respect to the most easily polarized atoms of the elements-partners, retain their predominantly ionic state. 6

The minimum polarizing effect in the series «lithium-cesium» would have to have Cs. However, according to the latest information, the Cs+ ion to some extent has the effect of additional polarization; therefore, in compounds that include, along with Cs+, highly polarizable anions, the noble-gas electron shell of the Cs+ ion ([Kr]4d105s25р6) undergoes deformation, leading to the formation of a chemical bond «cation – anion», including significant covalent component. Thus, only CsF fluoride does not participate in such polarization interactions.

Unusual compounds of alkali metals Traditionally, alkali metals were considered the simplest, always monovalent, forming only «simple» and highly soluble salts. One of the unexpected properties is their ability to form complex compounds. There are alkali metal amminates, which, under properly selected conditions, are capable of long-term existence. They are, for example, [Li(NH3)4]Cl, [Na(NH3)6]I, [K(NH3)6]I. In the complexes of alkali metal cations, the interaction of the central ion and ligands is of an electrostatic nature, the most stable complexes with monodentate ligands will be lithium. Alkali metals form intercalates obtained by treating graphite with vapors or molten alkali metal, for example, potassium. Potassium atoms occupy the space between the layers of graphite, increasing it from 3.36 to 5.4 Å. The composition of the intercalate depends on how many layers of graphite are separated by layers of alkali metals. For potassium, C8K compounds (potassium layers separate one layer of graphite), C24K (two separating layers), C36K (three layers of graphite), etc. are obtained. It is noteworthy that the conductivity in the plane of the potassium layer increases by 10 times compared to the usual one, and perpendicular to the plane – 300 times. These substances are promising as high-temperature superconductors. It is known that alkali metals dissolve in liquid ammonia and its derivatives, for example, amines and amides. Hydrogen compounds of alkali metals are hydrides, having a salt-like character with the predominant ionic type of chemical bond, can also be referred to unusual. 7

Lithium abnormalities 1. High ionization energy. 2. High covalence of chemical bonds in its compounds. 3. High electronegativity. 4. The small radius of the Li+ ion. 5. Higher hydration energy of the Li+ ion than other ions of the elements of the 1st group. 6. Tendency to complexation. Lithium is similar in properties to magnesium. Result

It easily reacts with N2, C, P, and slowly reacts with water

It doesn’t form hydrocarbonate

Compounds are lowsoluble, thermounstable

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Compounds form hydrates

PROPERTIES OF ELEMENTS OF THE 2nd GROUP Main characteristics of the 2nd group of the periodic system change naturally in the series «beryllium – radium»: the values of atomic and ionic radii increase, the values of ionization potentials decrease, the atomic mass increases, the «metallic» and basic properties increase. All elements (except radium) have a large number of stable isotopes (for example, barium has seven). The exception is beryllium. It has only one stable isotope (94Ве). Elements of the 2nd group are strong reducing agents, but weaker than elements of the 1st group. In accordance with the D.I.Mendeleyev geochemical rule, light elements of the group (magnesium and calcium) are more common than heavy elements (strontium and barium). Strontium is somewhat less common (the 19th place) than barium (the 17th place). This is explained by the fact that barium accumulated on Earth as a result of the fission of uranium, thorium and other radioactive elements. Main minerals of elements of the 2nd group Beryl (beryllium aluminosilicate, 3BeO·Аl2 O3 ·6SiO 2 , or Ве3[Аl2Si6O18]) is colored depending on small impurities. Single crystal samples of a beryl: – containing chromium, are known as precious stones – emeralds; – aquamarine is a modification of beryl containing an admixture of Fe(III), the color of «sea wave». In addition to aluminosilicates, there are minerals based on silicate or beryllium aluminate, for example, chrysoberyl Ве(АlO2)2. Magnesium in the outer layers of the lithosphere is most often found in the form of sedimentary rocks – carbonate or dolomite (mixed calcium-magnesium carbonate). A large amount of magnesium in the form of sulfate n bicarbonate is present in natural waters. 9

Calcium in the Earth's crust is present both in the form of igneous (primary minerals) and sedimentary rocks (secondary minerals). Most of the calcium is in the form of silicates and aluminosilicates (rocks – granites, gneisses, etc.). Sedimentary rocks (chalk and limestone) consist mainly of calcite mineral, and marble is a mixture of calcite and dolomite crystallized under high pressure. Anhydrite CaSO4 and gypsum CaSO4·2H2O are also among the secondary minerals. Calcium is also found in the form of a mineral – apatite, its formula in general form is Cа5(PO4)3(F,Cl,ОН). Calcium is present in natural waters (sulphate, bicarbonate), giving water hardness. Barium and strontium are found in the form of sulphates and carbonates, often replacing calcium in its compounds. For example, strontium is abundant in apatite, and therefore strontium salts are a by-product of apatite processing. The valence state of the elements of this group is determined by the relatively easy removal of two electrons from the ns-electron shell of neutral atoms. Due to this, the formation of doubly charged cations with relatively small size and a large density of a positive charge, especially in light elements of the group, is characteristic of the elements of the 2nd group. The ion of Be2+ has the unique characteristics; the ratio of charge to radius is five times greater than that of Mg2+. This is associated with a very high polarizing effect of Be2+ ion, its tendency to form covalent bonds and, as it is believed, its high toxicity. From top to bottom of the group, the positive charge density and the polarizing effect of doubly charged cations decrease. Due to this, the ionic character and the basic properties of most of the compounds of these elements grow. Most of the compounds of elements of this group are colorless. However, when heated to high temperatures: – calcium compounds give a pinkish-orange glow; – compounds of strontium – red-crimson; – barium compounds – green. Simple compounds of elements of the 2nd group Simple substances of all elements of the 2nd group are metals. They are relatively fusible but significantly higher than the corres10

ponding values for alkali metals. Beryllium is the most refractory, then down the group the melting point decreases. The specific weight of metals generally increases from beryllium to radium, but they are all among lightweight. As a result, the metals of the series «beryllium – barium» are very valuable for engineering and especially for the manufacture of aircraft. The lightest metals in this group are magnesium and calcium. However, magnesium has the lowest melting point (650 °C), which limits its use, and metallic calcium is unstable in air. Beryllium is much more refractory and chemically inert, but unlike magnesium and calcium, it is a rare and expensive metal. The density of metallic magnesium is higher than that of calcium and most alkali metals. However, magnesium has significant chemical inertness, unlike alkali and alkaline earth metals. This is due to the formation of a dense oxide layer on the magnesium surface, which prevents further oxidation. Metallic calcium, strontium, and barium react with oxygen and halogens at room temperatures. Heating for reactions with nitrogen, hydrogen, carbon, and silicon is necessary. With most metals, alkaline earth metals produce alloys that contain intermetallic compounds. With water vapor, alkaline earth metals react even in the cold. In water, alkaline earth metals quickly dissolve, forming hydroxides, which have alkali properties. Complex compounds of elements of the 2nd group Be-Ba oxides have composition with general formula MO; and only barium, whose cation is the largest, and therefore it has the smallest polarizing effect, BaO2 peroxide is completely stable. All oxides from BeO to BaO are refractory, which indicates a predominantly ionic type of bond in these compounds. The natural decrease in the melting point from MgO to BaO can be explained by the weakening of the Coulomb interaction of М2+– О2– in this series due to an increase in the size of the cation. Although the density of oxides in the series «beryllium – barium» increases, their hardness decreases, and in BeO it is equal to 9 (according to a ten-point scale), that is, close to the density of dia11

mond. All oxides have a high heat of formation, approximately equal to each other. Hydroxides M(OH)2 regularly change their properties in the series «beryllium – barium». The transition from Be(OH)2 to Ba(OH)2 corresponds to the transition from a weak, poorly soluble base to a strong, relatively well soluble. Hydroxides Cа(OH)2, Sr(OH)2, Bа(OH)2, Rа(OH)2 are alkalies. In solutions, they dissociate completely according to the following scheme: E(OH)2  E2+ + 2OH Thus, the three hydroxides of the series «beryllium – barium» (Са(ОН)2, Sr(OH)2, Ва(ОН)2) are among very strong bases, and Mg(OH)2 is the medium base. Their solubility is not too high even in the case of alkaline-earth metal derivatives, therefore alkaline solutions of Са(ОН)2, Sr(OH)2, Ва(ОН)2 are prepared by saturating the solution with solid M(OH)2. Hydroxide suspensions are often included in the synthesis. Interacting with water in supersaturated solutions, alkaline earth metal hydroxides crystallize to form hydrates: Са(OН)2·4Н2O; Sr(OН)2·8Н2O; Ва(OН)2·8Н2O. Only beryllium hydroxide can be considered as weak base, moreover, as amphoteric. The difference in the behavior of beryllium and magnesium hydroxides is explained by the large size of the Mg 2+ ion (0.78 Å), rather than the Be2+ ion (0.34 Å). The basic properties of oxides and hydroxides of alkaline earth metals are even stronger. Salts of these metals, formed by strong acids, practically do not hydrolyze; examples of formation of hydroxocomplexes in alkaline solutions or showing of acidic properties in mixed oxide systems are also unknown. Thus, the acid-base properties of doubly charged cations of elements of the 2nd group illustrate the principle of construction of the periodic system, i.e. a regular monotonous increase in the basic properties during the transition from light elements to heavy ones: from amphoteric, strongly hydrolyzing beryllium compounds to weakly basic magnesium compounds (practically non-hydrolysable and not developing amphoteric) and further to the strongly basic compounds of calcium, strontium and barium. 12

All elements of the 2nd group form positive doubly charged ions, which are able of performing a cationic function both in salt solutions and in solid compounds – anhydrous salts and the corresponding crystalline hydrates (or solvates). Beryllium has the maximum deviation from the basic properties, barium and radium have the minimum, therefore the strongest hydrolysis is in solutions of beryllium salts, weak – in magnesium salts; and practically hydrolysis does not occur in alkaline earth metal salts, whose cations have a weak polarizing effect. Thus, the regular change in acid-base properties in the group fully determines the properties of salts in aqueous solutions and their crystalline hydrates when heated.

Beryllium anomalies Amphoteric properties

Diagonal similarity

Be(OH)2

Be

Be Аl

H2BeO2

Thus, beryllium differs significantly from the other elements of the 2nd group: the small radius, the large value of the ionic potential, and the presence of the (helium) electron shell of the ion affect. The affinity for oxygen in beryllium and its analogues is very high. Their usage in metallothermy is based on abovementioned: 2Be + SiO2 → 2BeO + Si 3Be + B2O3 → 3BeO + 2B Beryllium is passivated (covered with an insoluble BeO film) in air, in water, and in oxidizing acids.

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BARIUM Symbol Bа History of discovery Barium compounds have attracted attention since the beginning of the 17th century due to unusual properties. In 1602, the Bologna shoemaker and alchemist Vincenzo Casciarolo found an extremely heavy stone. He suggested the presence of gold in it. To extract gold, he ignited found stone with coal and linseed oil. After cooling, the resulting product began to glow in the dark with a reddish light. For this property, heavy spar has received several names, among them are «sun stone», «bologna stone», «bologna phosphorus». Metallic barium was first obtained by the British chemist Sir Humphrey Davy (1808) by electrolysis of barite. Atomic mass 137.327 Position in the Period: 6, Periodic System group: 2 Electronic [Xe]4s2 configuration Electronegativity 0.97 Possible oxidation 0, + 2 states Simple substance Structure and physical At n.c. it is soft, malleable alkaline earth metal properties of silver-white color. It has a high chemical activity. Density (at n.c.), 3500 kg m-3 Boiling point 1640 °С Melting point 717 °С

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Production

The main raw material for producing barium is barite concentrate, containing 80-95% of BaSO4. Extraction of barium from it is carried out in several stages: 1) BаSO4 + 4C(chark) → BаS + 4CO (1100-1200 °С) BаSO4 + 2CH4 → BаS + 2C + 4H2O 2) BаS + 2H2O → Bа(OH)2 + H2S (450 °С, in a stream of CO2) BаS + H2O + CO2 → BаCO3 + H2S 3) Bа(OH)2 → BаO + H2O (780-800 °C) BаCO3 → BаO + CO2 (1000-1450 °C) 4) 3BаO + 2Аl → 3Bа + Аl2O3 (in a vacuum, 1200-1250 °С) Reaction with Bа + H2 → BаH2 hydrogen (150-300 °C) Reactions with Bа + F2 → BаF2 halogens (100-150 °С) Bа + Cl2 → BаCl2 (100-150 °С) Bа + Br2 → BаBr2 (100-150 °С) Bа + I2 → BаI2 (100-150 °С) Reactions with 2Bа + O2 → 2BаO chalcogens (>800 °C) 3Bа + 2O2 → 2BаO + BаO2 (>500 °C, air combustion) Bа + S → BаS (150 °C) Reaction with water Bа + 2H2O → Bа(OH)2 + H2 (room temperature) Reactions with acids Bа + H2S → BаS + H2 (>350 °C) Bа + 2HCl(dil.) → BаCl2 + H2 15

Other reactions

Application

4Bа + 10HNO3(dil.) → 4Bа(NO3)2 + + N2O + 5H2O 4Bа + 10HNO3(very dil.) → 4Bа(NO3)2 + + NH4NO3 + 3H2O 3Bа + N2 → Bа3N2 (200-460 °C, air combustion) Bа + 2C(graphite) → BаC2 (500 °C) 6Bа + 2NH3(gas) → Bа3N2 + 3BаH2 (600-650 °C) Bа + 2NH3(liq.) → Bа(NH2)2 + H2 (ct. is Pt) 2Bа + 3CO2 → 2BаCO3 + C(graphite) (room temperature) – as a getter in high-vacuum electronic devices; – as an active layer of indirectly heated cathodes; – as a dielectric in the manufacture of ceramic capacitors; – as a material for piezoelectric microphones and piezoelectric emitters; – in a form of single crystals in optics (lenses, prisms); – as oxidizing agents; – in pyrotechnics for coloring the flame (green fire); – in receiving of hydrogen and oxygen by thermochemical method (Oak-Ridge cycle, USA); – for a synthesis of superconducting ceramics operating at a temperature of liquid nitrogen and above; – for production of a special kind of glass used to cover uranium rods; – in solid-state fluoride batteries as a component of fluoride electrolyte; 16

– in powerful copper oxide batteries as a component of the active mass (barium oxide-copper oxide); – as an expander of the active mass of the negative electrode in the production of lead-acid batteries; – as a radiopaque substance in the medical examination of the gastrointestinal tract. Barium hydride Structure and physical At n.c. it is a colorless crystalline substance. properties Density (at n.c.), 4210 kg m-3 Melting point decomposes at 675 °C Production By treatment of heated barium with hydrogen: Bа + H2 → BаH2 (150–300 °C) Reaction with water BаH2 + 2H2O → Bа(OH)2 +2H2 Reactions with BаH2 + O2 → BаO + H2O chalcogens (150-200 °C) Reactions with acids BаH2 + 2HCl(dil.) → BаCl2 + 2H2 Other reactions BаH2 → Bа + H2 (>675 °C) 3BаH2 + N2 → Bа3N2 + 3H2 (400-450 °C) 3BаH2 + 2KClO3 → 2KCl + 3BаO + 3H2O (350-400 °C) Barium oxide Structure and physical At n.c. it is colorless crystals with a cubic latproperties tice. Density (at n.c.), 5800 -3 kg m Boiling point 2000 °С 17

1920 °С 1) by thermal decomposition of carbonate or nitrate of barium: BаCO3 → BаO + CO2 (1000-1450 °C) 2Bа(NO3)2 → 2BаO + 4NO2 + O2 (620-670 °C) 2) by direct synthesis: 2Bа + O2 → 2BаO (>800 °C) Reactions with 2BаO + O2 → 2BаO2 chalcogens (1580 °C) BаSO4 + 4C(chark) → BаS + 4CO (1100-1200 °C) BаSO4 + 4CO → BаS + 4CO2 (600-800 °C) Barium carbonate Structure and physical At n.c. it is colorless crystals. Up to 810 °C, properties α-modification with a rhombic lattice is stable; at 810-960 °C – β-modification with a hexagonal lattice is stable; at > 960 °C – γ-modification with a cubic lattice is stable. Density (at n.c.), kg m-3 4430 °С (γ-modification) Melting point Transformation to β-modification is at 811 °С (γ-modification) 21

Production

1) by interaction of BаS with CO2 or Nа2CO3: BаS + Nа2CO3 → BаCO3 + Nа2S 2) by exchange reactions of soluble barium salts or barium hydroxide with carbonate solutions: BаCl2 + Nа2CO3(conc.) → BаCO3 + 2NаCl Reactions with acids BаCO3 + 2HCl(dil.) → BаCl2 + CO2 + H2O BаCO3 + 2HF(dil.) → BаF2 + CO2 + H2O (900-1100 °C) BаCO3 + H2S → BаS + CO2 + H2O (1000 °C, in stream of H2) Other reactions BаCO3 → BаO + CO2 (1000-1450 °C) BаCO3 + C(chark) → BаO + 2CO (>1000 °C)

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BERYLLIUM Symbol Be History of discovery It is discovered in 1798 by the French chemist Louis Nicolas Vauquelin, who, in a comparative analysis of beryl and emerald, discovered the presence of an oxide of an unknown chemical element. It was similar to alumina (aluminum oxide), but it had some differences. Atomic mass 9.012 Position in the Period: 2, Periodic System group: 2 Electronic [He]2s² configuration Electronegativity 1.47 Possible oxidation 0, +1, +2 states Simple substance Structure and physical At n.c. it is light gray relatively hard metal. properties Density (at n.c.), 1860 kg m-3 Boiling point 2970 °C Melting point 1560 °C Production 1) by action of potassium on anhydrous beryllium chloride: BeCl2 + 2K → Be + 2KCl 2) by reduction of beryllium fluoride by magnesium: BeF2 + Mg → Be + MgF2 (700-750 °С) 3) by electrolysis of melt of a mixture of beryllium and sodium chlorides. Reactions with Be + F2 → BeF2 halogens (room temperature) 23

Be + Cl2 → BeCl2 (250 °C) Be + Br2 → BeBr2 (480 °C) Be + I2 → BeI2 (480 °C) Reactions with 2Be + O2 → 2BeO chalcogens (900 °C, air combustion) Be + S → BeS (1150 °C) Reaction with water 2Be + 3H2O → BeO + Be(OH)2 +2H2 (boiling) Reactions with acids Be + 2HCl(dil.) → BeCl2 + H2 3Be + 8HNO3(dil., hot) → → 3Be(NO3)2 + 2NO + 4H2O Be + 4HF(conc.) → H2[BeF4] + H2 Reactions with bases Be + 2NаOH → Nа2BeO2 + H2 (400-500 °C) Be + 2NаOH(conc.) + 2H2O → → Nа2[Be(OH)4] + H2 Other reactions 3Be + N2 → Be3N2 (700–900 °C) 3Be + 2NH3 → Be3N2 + 3H2 (500–700 °C) Be + C2H2 → BeC2 + H2 (400–450 °C) Be + MgO → BeO+ Mg (1075 °C) Application – as an alloying additive to alloys based on Cu, Ni, Zn, Al, Pb, and other non-ferrous metals; – as reflectors and neutron moderators, as well as a neutron source; – as materials for shells of fuel elements; – as a structural material in aircraft, rocket production, and space technology; – for manufacture of windows of X-ray tubes; 24

– for inserting of a solid diffusion layer on steel surface (beryllization); – as additives to rocket fuel; – as a material of buildings, heat sinks and insulators of semiconductor devices; – for manufacture of crucibles and special ceramics. Beryllium hydride Structure and physical At n.c. it is a solid amorphous white substanproperties ce. The bond between hydrogen and beryllium is covalent. Melting point decomposes at 125 °C Production 1) by interaction of a solution of dimethylberyllium in diethyl ether with lithium aluminum hydride: Be(CH3)2 + LiАlH4 → BeH2 + LiАlH2(CH3)2 (dimethyl ether) 2) by reaction of triphenylphosphine with beryllium borohydride: Be(BH4)2 + 2PPh3 → BeH2 + 2Ph3PBH3 Reactions with BeH2 + O2 → BeO + H2O chalcogens Reaction with water BeH2 + 2H2O → Be(OH)2 + 2H2 Reactions with bases BeH2 + 2NаOH → Nа2BeO2 + 2H2 Other reactions BeH2 → Be + H2 (1000 °С) Reactions with 2BeO + 2F2 → 2BeF2 + O2 halogens (>400 °C) Reactions with acids BeO + 2HCl(conc.) → BeCl2 + H2O BeO + H2SO4(conc.) → BeSO4 + H2O BeO + 2HF(gas) → BeF2 + H2O (220 °C) BeO + 4HF(conc.) → H2[BeF4] + H2O Reactions with bases BeO + 2NаOH → Nа2BeO2 + H2O (250-300 °C) BeO + 2NаOH(conc., hot) + H2O → → Nа2[Be(OH)4] Other reactions 2BeO + CS2 → 2BeS + CO2 (650-700 °C) BeO + Аl2O3 → (BeАl2)O4 (1400 °C) BeO + Mg → MgO + Be (700-800 °C) Beryllium hydroxide Structure and physical properties Density (at n.c.), kg m-3 Melting point Production Reaction with water

At n.c. it is gelatinous white matter. 1909 decomposes at 138 °C 2Be + 3H2O → BeO + Be(OH)2 +2H2 (boiling) Be(OH)2(solid) + 4H2O ↔ ↔ [Be(H2O)4]2+ + 2OH− Be(OH)2(solid) + 4H2O ↔ ↔ [Be(OH)4]2− + 2H3O+ 26

Reactions with acids

Reactions with bases Other reactions

Be(OH)2 + 2HCl(dil.) → BeCl2 + 2H2O Be(OH)2 + 2HF(dil.) → BeF2 + 2H2O Be(OH)2 + 4HF(conc.) → H2[BeF4]+ 2H2O Be(OH)2 + 2NаOH(conc.) → Nа2[Be(OH)4] Be(OH)2 + 2NаOH → Nа2BeO2 + 2H2O (200–300 °C) Be(OH)2 → BeO + H2O (200–800 °C) 2Be(OH)2 + CO2 → Be2CO3(OH)2 + H2O Beryllium chloride

Structure and physical At n.c. it is white or slightly greenish needles, properties spreading in the air due to strong hygroscopicity. Density (at n.c.), kg m-3 1899 Boiling point 500±20 °C Melting point 404 °C Production 1) by interaction of beryllium carbonate, oxide or hydroxide with hydrochloric acid: BeCO3 + 2HCl(dil.) → BeCl2 + CO2 + H2O BeO + 2HCl(conc.) → BeCl2 + H2O Be(OH)2 + 2HCl(dil.) → BeCl2 + 2H2O 2) by interaction of metallic beryllium or its oxide with chlorine: Be + Cl2 → BeCl2 (250 °C) 2BeO + 2Cl2 → 2BeCl2 + O2 3) by heating of beryllium oxide with many chlorine-containing compounds: BeO + COCl2 → BeCl2 + CO2 2BeO + CCl4 → 2BeCl2 + CO2 Reaction with water BeCl2 + H2O(hot) → BeCl(OH) +HCl 3BeCl2(conc.) + 6H2O ↔ ↔ 2[Be(H2O)3Cl]+ + [BeCl4]2− (in conc. HCl) BeCl2(dil.) + 4H2O(cold) → [Be(H2O)4]2+ + 2Cl− (in diluted HCl) 27

Reactions with acids Reactions with bases

Other reactions

BeCl2 + 4HF(conc.) → H2[BeF4]+ 2HCl BeCl2 + 2NаOH(dil.) → Be(OH)2 + 2NаCl BeCl2 + 4NаOH(conc.) → → Nа2[Be(OH)4] + 2NаCl BeCl2 + +2NH4OН(conc.) → Be(OH)2 + 2NH4Cl BeCl2 + 2LiH → BeH2 + 2LiCl (in ether) BeCl2 + 4NаF(conc.) → Nа2[BeF4] + 2NаCl Beryllium nitrate

Structure and physical properties Boiling point Melting point Production

At n.c. it is yellowish-white crystals.

decomposes at 100-200 °C 60 °C By interaction of beryllium oxide or hydroxide with nitric acid or exchange reaction: BeO + 2HNO3 → Be(NO3)2 + H2O Be(OH)2 + 2HNO3 → Be(NO3)2 + 2H2O BeSO4 + Bа(NO3)2 → Be(NO3)2 + BаSO4 Reaction with water Be(NO3)2(dil.) + 4H2O(cold) → [Be(H2O)4]2+ + 2NO3− (in diluted HNO3) Be(NO3)2 + H2O(cold) → Be(NO3)OH + HNO3 Reactions with Be(NO3)2 + 2NаOH(dil.) → Be(OH)2 + 2NаNO3 bases Be(NO3)2 + 4NаOH(conc.) → → Nа2[Be(OH)4] + 2NаNO3 Other reactions 8Be(NO3)2 → 2[Be4(NO3)6O] + 4NO2 + O2 (125 °C, vacuum) 2Be(NO3)2 → 2BeO + 4NO2 + O2 (>1000 °С) Beryllium sulphate Structure and physical At n.c. it is a white solid crystalline substance. properties Density (at n.c.), kg m-3 2443 Melting point decomposes at 550-600 °C 28

Production

By interaction in an aqueous solution of any beryllium salt with sulfuric acid, followed by evaporation and crystallization of the product of the reaction: BeCl2 + H2SO4 → BeSO4 + 2HCl Reaction with water 2BeSO4 + 2H2O → Be2SO4(OH)2 +H2SO4 (boiling) BeSO4(dil.) + 4H2O → [Be(H2O)4]2+ + SO42− (in diluted H2SO4) Reactions with bases BeSO4 + 2NаOH(dil.) → Be(OH)2 + Nа2SO4 BeSO4 + 4NаOH(conc.) → → Nа2[Be(OH)4] + Nа2SO4 BeSO4 + 2NH4OН(conc.) → → Be(OH)2 + (NH4)2SO4 Other reactions BeSO4 + Bа(NO3)2 → Be(NO3)2 + BаSO4 Beryllium carbonate Structure and physical At n.c. it is a colorless crystalline substance. properties Melting point decomposes at 100 °C Production By interaction of beryllium oxide and carbon dioxide (under increased pressure): BeO + CO2 → BeCO3 Reaction with water 2BeCO3 + H2O(hot) → Be2CO3(OH)2 + CO2 Reactions with acids BeCO3 + 2HCl(dil.) → BeCl2 + CO2 + H2O BeCO3 + 4HF(conc.) → H2[BeF4] + CO2 + H2O Reactions with bases BeCO3 + 4NаOH(conc., hot) → → Nа2[Be(OH)4] + Nа2CO3 Other reactions BeCO3 → BeO + CO2 (>180 °C) BeCO3(solid) + (NH4)2CO3(conc.) → → (NH4)2[Be(CO3)2](solution)

29

CALCIUM Symbol Cа History of discovery In 1808, metallic calcium was isolated by the English chemist Sir Humphry Davy by electrolysis of a mixture of wet hydrated lime with mercury oxide on a platinum plate, which is anode, and platinum wire immersed in liquid mercury served as cathode. Atomic mass 40.078 Position in the Period: 4, Periodic System group: 2 Electronic [Аr]4s2 configuration Electronegativity 1.04 Possible oxidation 0, + 2 states Simple substance Structure and physical At n.c. it is silvery-white solid with a cubic properties face-centered lattice. It is alkaline-earth metal, silvery-white, ductile, rather hard. Density (at n.c.), 1550 kg m-3 Boiling point 1482 °C Melting point 851 °C Production 1) CаH2 → Cа + H2 (>1000 °С) 2)4CаO + 2Аl → 3Cа + (CаАl2)O4 (1200 °C) 3) 3CаCl2 + 2Аl → 3Cа + 2АlCl3 (600-700 °С ) Reaction with Cа + H2 → CаH2 hydrogen (500-700 °C) Reactions with Cа + F2 → CаF2 halogens (room temperature) 30

Cа + Cl2 → CаCl2 (200-400 °С) Cа + Br2 → CаBr2 (200-400 °С) Cа + I2 → CаI2 (200-400 °С) Reactions with 2Cа + O2 → 2CаO chalcogens (>300 °C, combustion in air) Cа + S → CаS (150 °C) Reaction with water Cа + 2H2O → Cа(OH)2 + H2 2Cа + H2O(vapor) → CаO + CаH2 (200-300 °C) Reactions with acids Cа + 2HCl(dil.) → CаCl2 + H2 4Cа + 10HNO3(dil.) → → 4Cа(NO3)2 + N2O + 5H2O 4Cа + 10HNO3(very dil.) → → 4Cа(NO3)2 + NH4NO3 + 3H2O Other reactions 3Cа + N2 → Cа3N2 (200-450 °C, combustion in air) 3Cа + 2P(red) → Cа3P2 (350-450 °C) Cа + 2C(graphite) → CаC2 (550 °C) 6Cа + 2NH3(gas) → Cа3N2 + 3CаH2 (600-650 °C) Cа + 2NH3(liq.) → Cа(NH2)2 + H2 (ct. is Pt) Application – as a reducing agent in production of metals; – to obtain hard-to-reduce metals, such as chromium, thorium, and uranium; – alloys of calcium with lead in batteries and bearing alloys; – to remove traces of air from vacuum devices; – in metallurgy for deoxidizing steel along with aluminum or in combination with it; – in food industry as a food additive E526. 31

Calcium hydride Structure and physical At n.c. it is white, melts without decomposiproperties tion in a hydrogen atmosphere with further heating decomposes. It is a strong reducing agent. Density (at n.c.), 1900 -3 kg m Boiling point decomposes at ~600 °C Melting point 816 °C Production 1) by treatment of heated calcium by hydrogen: Cа + H2 → CаH2 (500–700 °C) 2) by reduction of calcium chloride with hydrogen: CаCl2 + 2H2 → CаH2 + 2HCl (600–700 °C; ct. is Pt, Fe, Ni) Reaction with water CаH2 + 2H2O → Cа(OH)2 + 2H2 Reactions with CаH2 + O2 → CаO + H2O chalcogens (300-400 °C) Reactions with acids CаH2 + 2HCl(dil.) → CаCl2 + 2H2 CаH2 + 2HNO3 → Cа(NO3)2 + 2H2 CаH2 + H2SO4 → CаSO4 + 2H2 Other reactions 3CаH2 + N2 → Cа3N2 + 3H2 (>1000 °C) 2CаH2 + TiO2 → 2CаO + Ti + 2H2 (750 °C) 3CаH2 + 2KClO3 → 2KCl + 3CаO + 3H2O (450-550 °C) Calcium oxide Structure and physical At n.c. it is a white, hygroscopic solid substanproperties ce. It is refractory, thermally stable, and volatile at very high temperatures. -3 Density (at n.c.), kg m 3370 32

2850 °C 2580 °C 1) by thermal decomposition of limestone: CаCO3 → CаO + CO2 (900-1200 °C) 2) by interaction of simple substances: 2Cа + O2 → 2CаO (500-700 °C) 3) by thermal decomposition of calcium hydroxide and calcium salts of some oxygen-containing acids: 2Cа(NO3)2 → 2CаO + 4NO2 + O2 (>560 °C) 4) by oxidation of calcium hydride: CаH2 + O2 → CаO + H2O (300-400 °C) Reactions with 2CаO + 2Cl2 → 2CаCl2 + O2 halogens (700 °C) Reaction with water CаO + H2O → Cа(OH)2 Reactions with acids CаO + 2HCl(dil.) → CаCl2 + H2O CаO + 2HF(dil.) → CаF2 +H2O 3CаO + 2H3PO4(dil.) → Cа3(PO4)2 + 3H2O CаO + 2HCN → CаCN2 + CO + H2 (700 °C) Other reactions CаO + 3C(chark) → CаC2 + CO (1900-1950 °C) CаO + CO2 → CаCO3 (room temperature) CаO + SiO2 → CаSiO3 (wollastonite) (1100-1200 °C) CаO + TiO2 → CаTiO3 (perovskite) (900-1100 °C) 4CаO + 2Cr2O3 + 3O2 → 4CаCrO4 (chromatite) (600-700 °C) 4CаO + 2Аl → 3Cа + CаАl2O4 (1200 °C) Boiling point Melting point Production

33

Calcium hydroxide Structure and physical At n.c. it is a white substance, when heated properties decomposes without melting. Density (at n.c.), 2240 kg m-3 Boiling point decomposes Melting point 580 °C Production 1) by interaction of calcium salts with aqueous solutions of alkalis: CаCl2 + 2NаOH → Cа(OH)2 + 2NаCl 2) by interaction of calcium with water: Cа + 2H2O → Cа(OH)2 + H2 Reactions with 2Cа(OH)2(suspension, cold) + 2Cl2 → halogens → Cа(ClO)2 + CаCl2 + 2H2O 6Cа(OH)2(suspension, hot) + 6Cl2 → → Cа(ClO3)2 + 5CаCl2 + 6H2O Reaction with water Cа(OH)2(very dil.) + 6H2O → → [Cа(H2O)6]2+ + 2OHReactions with acids Cа(OH)2 + 2HCl(dil.) → CаCl2 + 2H2O Cа(OH)2 + H2SO4(conc.) → CаSO4 + 2H2O 3Cа(OH)2 + 2H3PO4(dil.) → Cа3(PO4)2 + 6H2O Cа(OH)2 + H3PO4(conc.) → CаHPO4 + 2H2O Other reactions Cа(OH)2 + 4NаOH + 4ClO2 + C → → 4NаClO2 + CаCO3 + 3H2O Cа(OH)2 + SO2 → CаSO3 + H2O Cа(OH)2(suspension) + 2NаClO(cold) → → Cа(ClO)2 + 2NаOH Cа(OH)2 + H2O2(conc.) → CаO2 +2H2O (40-50 °C) Calcium chloride Structure and physical At n.c. it is white, melts without decomposiproperties tion. Density (at n.c.), kg m-3 2512 Boiling point 1600 °C 34

782 °C

Melting point Production

1) in industry: as by-product in soda production: 2NH4Cl + Cа(OH)2 → 2NH3 + CаCl2 + 2H2O 2) in laboratory: CаH2 + 2HCl(dil.) → CаCl2 + 2H2 CаO + 2HCl(dil.) → CаCl2 + H2O Cа(OH)2 + 2HCl(dil.) → CаCl2 + 2H2O Reaction with CаCl2 + 2H2 → CаH2 + 2HCl hydrogen (600-700 °C; ct. is Pt, Fe, Ni) Reaction with water CаCl2(dil.) + 6H2O → [Cа(H2O)6]2+ + 2Cl(pH=7) Reactions with acids CаCl2(solid) + H2SO4(conc.) → CаSO4 + 2HCl Reactions with bases CаCl2 + 2NаOH(conc.) → Cа(OH)2 + 2NаCl Other reactions CаCl2 + Nа2CO3 → CаCO3 + 2NаCl CаCl2 + 2NH4F → CаF2 + 2NH4Cl CаCl2 + K2SO4 → CаSO4 + 2KCl (800 °C) 3CаCl2 + 2Аl → 3Cа + 2АlCl3 (600-700 °C) Calcium nitrate Structure and physical At n.c. it is a white solid substance, when melproperties ting decomposes. -3 Density (at n.c.), kg m 2360 Melting point 560 °C Boiling point decomposes Production 1) by dissolving of calcium sulfide in concentrated nitric acid: CаS + 4HNO3(conc.) → → Cа(NO3)2 + S + 2NO2 + 2H2O 2) by interaction of calcium with nitric acid: 4Cа + 10HNO3(dil.) → → 4Cа(NO3)2 + N2O + 5H2O 4Cа + 10HNO3(very dil.) → → 4Cа(NO3)2 + NH4NO3 + 3H2O 35

Reaction with water

Other reactions

Cа(NO3)2(dil.) + 6H2O → → Cа[(H2O)6]2+ + 2NO3− (pH=7) Cа(NO3)2 → Cа(NO2)2 + O2 (450-500 °С) 2Cа(NO3)2 → 2CаO + 4NO2 + O2 (>560 °С) 3Cа(NO3)2 + 2Nа2HPO4 → → Cа3(PO4)2 + 4NаNO3 + 2HNO3 (boiling) 5Cа(NO3)2 + 3(NH4)2HPO4 + 4NH4OH(dil.) → → Cа5(PO4)3OH + 10NH4NO3 + 3H2O (boiling) Calcium sulphate

Structure and physical At n.c. it is a white hygroscopic solid substanproperties ce. Density (at n.c.), 2990 kg m-3 Melting point 1450 °C Production By interaction of hydride or hydroxide of calcium and sulfuric acid: CаH2 + H2SO4 → CаSO4 + 2H2 Cа(OH)2 + H2SO4(conc.) → → CаSO4 + 2H2O Reactions with acids CаSO4 + H2SO4(conc.) → Cа(HSO4)2 Other reactions 2CаSO4 → 2CаO + 2SO2 + O2 (>1450 °C) CаSO4 + 3C(chark) → CаS + 2CO + CO2 (900 °C) CаSO4 + 4CO → CаS + 4CO2 (600-800 °C) CаSO4 + Nа2CO3(conc.) → → CаCO3 + Nа2SO4

36

Calcium carbonate Structure and physical At n.c. it is a white solid substance, when calproperties cined decomposes; melts without decomposition under an excessive pressure of CO2. Density (at n.c.), 2711 kg m-3 Boiling point decomposes Melting point 1336 °C Production CаO + CO2 → CаCO3 Cа(OH)2 + CO → CаCO3 + H2 (400 °C) Reactions with acids CаCO3 + 2HCl(dil.) → CаCl2 + CO2 + H2O CаCO3 + 2HF(dil.) → CаF2 + CO2 + H2O Other reactions CаCO3 + SiO2 → CаSiO3 + CO2 (800 °C) CаCO3 + 2NH3 → CаCN2 + 3H2O (700-900 °C) CаCO3 + 2NH4Cl(conc.) → → CаCl2 + 2NH3 + H2O + CO2 (boiling)

37

CESIUM Symbol Cs History of discovery Cesium was discovered in 1860 by German scientists Robert Wilhelm Eberhard Bunsen and Gustav Robert Kirchhoff in the waters of the Bad Durkheim mineral spring in Germany by optical spectroscopy, thus becoming the first element discovered by spectral analysis. In its pure form, cesium was first isolated in 1882 by the Swedish chemist Carl Setterberg during the electrolysis of a melt of a mixture of cesium cyanide (CsCN) and barium. Atomic mass 132.9 Position in the Period: 6 Periodic System group: 1 Electronic [Хе]6s1 configuration Electronegativity 0.86 Possible oxidation 0, +1 states Simple substance Structure and physical At n.c. it is a soft alkaline metal of silverproperties white color (light yellow on the cut), with very low melting point. It colors the flame of a gas burner in a greenish blue. Density (at n.c.), 1900 kg m-3 Boiling point 688 °C Melting point 28.5 °C Production 1) by decomposition of oxide: 2Cs2O → Cs2O2 + 2Cs (300–500 °C) 2) by electrolysis of cesium hydroxide.

38

2Cs + H2 → 2CsH (300–350 °C, P) 2Cs + F2 → 2CsF (room temperature) 2Cs + Cl2 → 2CsCl (room temperature) 2Cs + Br2 → 2CsBr (room temperature) 2Cs + I2 → 2CsI (room temperature) Reactions with 2Cs + S → Cs2S chalcogens (100-130 °C) 4Cs + O2 → 2Cs2O (in cold) Cs + О2(air) → CsО2 (burning) Reaction with water 2Cs + 2H2O → 2CsOH + H2 4Cs + O2 + 2H2O → 4CsOH Reactions with acids 2Cs + 2H2S(saturated) → 2CsHS +H2 (in benzene) 2Cs + 2HCl(dil.) → 2CsCl + H2 8Cs + 6H2SO4(dil., cold) → → 4Cs2SO4 + SO2 + S + 6H2O (impurities of H2S) 21Cs + 26HNO3(dil., cold) → → 21CsNO3 + NO + N2O + N2 + 13H2O Reactions with bases 2Cs + 2CsOH → 2Cs2O + H2 (300-350 °C) Other reactions 2Cs + 2NH3(gas) → 2CsNH2 + H2 (30-45 °C) Application – in electronics; – radio, electrical, X-ray technology; – chemical industry; – optics; – medicine; – nuclear power. Reaction with hydrogen Reactions with halogens

39

Cesium hydride Structure and physical At n.c. it is white cubic face-centered crystals. properties Density (at n.c.), 3420 kg m-3 Melting point decomposes Production By passing hydrogen through heated cesium: 2Cs + H2 → 2CsH (300-350 °C, P) Reactions with CsH + Cl2 → CsCl + HCl halogens (400 °C) Reactions with 2CsH + O2 → 2CsOH chalcogens (>200 °C) 2CsH + 2S → Cs2S+ H2S (300–350 °C) Reaction with water CsH + H2O → CsOH + H2 Reactions with acids CsH + HCl(dil.) → CsCl + H2 (>200 °C) Other reactions CsH + NH3(gas) → CsNH2 + H2 (300 °C) 2CsH → 2Cs + H2 (>200 °C) Cesium oxide Structure and physical At n.c. it is an orange-red substance, when properties heated becomes first dark red, then black. It has a hexagonal structure. Density (at n.c.), 4360 -3 kg m Melting point decomposes at 400 °C Production 1) by careful heating (below 200 °C) excess cesium in an oxygen atmosphere, followed by vacuum distillation of Cs: 4Cs + O2 → 2Cs2O

40

2) by calcination of cesium carbonate: Cs2CO3 → Cs2O + CO2 (620-1000 °C, vacuum) Reaction with water Cs2O + H2O → 2CsOH Reactions with acids Cs2O + 2HCl(dil.) → → 2CsCl + H2O Other reactions Cs2O + CO2(moistered) → Cs2CO3 Cs2O + H2O + 2CO2 → 2CsHCO3 2Cs2O → Cs2O2 + 2Cs (300-500 °C) Cs2O + NH3(liq.) → CsNH2 + CsOH (-50 °C) Cesium hydroxide Structure and physical At n.c. it is white substance, melts without deproperties composition, volatile. It has a cubic body-centered structure. Density (at n.c.), 3675 kg m-3 Melting point 272.3 °C Production 1) by interaction of rubidium with water: 2Cs + 2H2O → 2CsOH + H2 2) by interaction of rubidium oxide with water: Cs2O + H2O → 2CsOH 3) by lime method: Cs2CO3 + Cа(OH)2(saturated) → → 2CsOH + CаCO3 Reactions with 4CsOH(liq.) + 3O2 → 4CsO2 + 2H2O chalcogens (400 °C) Reaction with water CsOH(dil.) + 6H2O → → [Cs(H2O)6]+ + OHReactions with acids CsOH + HCl(dil.) → CsCl + H2O 2CsOH + H2SO4(dil.) → Cs2SO4 + 2H2O CsOH + HNO3(dil.) → CsNO3 + H2O Other reactions 2CsOH(conc.) + CO2 → Cs2CO3 + H2O

41

Cesium chloride Structure and physical At n.c. it is a white substance; it melts and properties boils without decomposition. It has a cubic body-centered structure. Density (at n.c.), 3970 kg m-3 Boiling point 1300 °C Melting point 642 °C Production 1) by exchange reaction: Cs2SO4 + BаCl2 → BаSO4 + 2CsCl 2) by interaction of cesium hydroxide and its salts with hydrochloric acid: Сs2CO3 + 2HCl → 2CsCl + CO2 + H2O CsOH + HCl → CsCl + H2O 3) by interaction of hydrochloric acid and metallic cesium: 2Cs + 2HCl(dil.) → 2CsCl + H2 4) by direct synthesis: 2Cs + Cl2 → 2CsCl Reaction with water CsCl(dil.) + 6H2O → → [Cs(H2O)6]+ + Cl− (pH = 7) Reactions with acids 2CsCl(solid) + H2SO4(conc.) → → Cs2SO4 + 2HCl (boiling) Other reactions CsCl + CsHSO4 → Cs2SO4 + HCl (500–600 °C) 2CsCl + H2[SnCl6] → Cs2[SnCl6] + 2HCl (in ethanol) 2CsCl + H2[PtCl6] → → Cs2[PtCl6] + 2HCl (in dil. HCl) 10CsCl(solid) + 8H2SO4(conc., hot) + + 2KMnO4(solid) → 5Cl2 + 2MnSO4 + + 5Cs2SO4 + K2SO4 + 8H2O

42

Cesium nitrate Structure and physical At n.c. it is white substance, melts without deproperties composition, decomposes with further heating. It has a hexagonal structure. Density (at n.c.), 3685 kg m-3 Boiling point decomposes Melting point 471 °C Production 1) by interaction of metallic cesium or its oxide with nitric acid: 21Cs + 26HNO3 → → 21CsNO3 + NO + N2O + N2 + 13H2O Cs2O + 2HNO3 → 2CsNO3 + H2O 2) by interaction of cesium hydroxide or cesium salts and nitric acid: CsOH + HNO3 → CsNO3 + H2O Cs2CO3 + 2HNO3 → 2CsNO3 + CO2 + H2O Reaction with water CsNO3(dil.) + 6H2O → [Cs(H2O)6]+ + NO3− (pH=7) Other reactions CsNO3 + 8H°(Zn, conc. NаOH) → → NH3 + 2H2O + CsOH (boiling) CsNO3 + 2H°(Zn, dil. HCl) → CsNO2 + H2O 2CsNO3 + (NH4)2SO4 → → Cs2SO4 + 2N2O + 4H2O (300–350 °C) CsNO3 + Pb → CsNO2 + PbO (400 °C) 2CsNO3 → 2CsNO2 + O2 (585-850 °C) Cesium sulphate Structure and physical At n.c. it is white substance, volatile; it melts properties without decomposition. It has a cubic volumecentered form. 43

Density (at n.c.), kg m-3 4243 Melting point 1010 °C Production 1) by interaction of metallic cesium and sulfuric acid: 8Cs + 6H2SO4(dil., cold) → → 4Cs2SO4 + SO2 + S + 6H2O (impurities of H2S) 2) by oxidation of cesium sulfide with oxygen: Cs2S(solid) + 2O2 → Cs2SO4 (>500 °C) 3) by interaction of cesium chloride and sulfuric acid in the presence of potassium permanganate: 10CsCl(solid) + 8H2SO4(conc., hot) + 2KMnO4(solid) → → 5Cl2 + 2MnSO4 + 5Cs2SO4 + K2SO4 + 8H2O 4) by interaction of cesium hydroxide and sulfuric acid: 2CsOH + H2SO4(dil.) → Cs2SO4 + 2Н2О Reaction with water Cs2SO4(dil.) + 12H2O → 2Cs[(H2O)6]+ + SO42− (pH=7) Reactions with acids Cs2SO4 (solid) + H2SO4(conc.) → 2CsHSO4(solution) Reactions with bases Cs2SO4 + Bа(OH)2 → 2CsOH + BaSO4 Other reactions Cs2SO4 + BаCl2 → BаSO4 + 2CsCl Cs2SO4 + Аl2(SO4)3 + 24H2O → → 2{CsАl(SO4)2·12H2O} (alum) Cesium carbonate Structure and physical At n.c. it is white substance; when calcined, properties decomposes, melts under an excessive pressure of CO2. Melting point decomposes at 610 °C Production By interaction of oxide with carbon dioxide: 4CsO2 + 2CO2(moistered) → 2Cs2CO3 + 3O2 (room temperature) Reaction with water Cs2CO3(dil.) + 12H2O → 2[Cs(H2O)6]+ + CO32− 44

Cs2CO3 + 2HCl(dil.) → 2CsCl + CO2 + H2O Cs2CO3 + HClO4(conc., cold) → → 2CsClO4 + CO2 + H2O Reactions with bases Cs2CO3 + Cа(OH)2(saturated) → 2CsOH + CаCO3 Other reactions Cs2CO3 → Cs2O + CO2 (620-1000 °C, vacuum) Reactions with acids

45

FRANCIUM Symbol Fr History of discovery Francium was predicted by D.I. Mendeleyev. In 1939 it was discovered due to radioactivity by the French radiochemist Marguerite Catherine Perey, at the Radium Institute in Paris, and in 1946 it was given the name in honor of her homeland – France. Atomic mass 223.02 Position in the Period: 7, Periodic System group: 1 Electronic configuration [Rn]7s1 Electronegativity 0.86 Possible oxidation states 0, +1, +2 Simple substance Structure and physical At n.c. it is a relatively hard metal of light gray properties color. Production Microscopic amounts of francium-223 and francium-224 can be chemically separated from the minerals of uranium and thorium. Other isotopes of francium are obtained artificially by nuclear reactions. Reaction with oxygen 4Fr + O2 → 2Fr2O Reactions with 2Fr + Cl2 → 2FrCl halogens 2Fr + F2 → 2FrF 2Fr + I2 → 2FrI 2Fr + Br2 → 2FrBr Reactions with 2Fr + S → Fr2S chalcogens Reaction with water 2Fr + 2H2O → 2FrOH + H2 Reactions with acids 8Fr + 10HNO3 → 8FrNO3 + N2O + 5H2O 8Fr + 5H2SO4 → 4Fr2SO4 + H2S + 4H2O 2Fr + 2HCl → 2FrCl + H2 Application At present, francium and its salts are of no practical use, due to the short half-life and high radioactivity. 46

HELIUM Symbol He History of discovery Helium was first detected on August 18, 1868, by French scientist Pierre Jules César Janssen. He set up the spectroscope so that it became possible to observe the solar corona not only during an eclipse but also on ordinary days. Thus, he revealed a very bright yellow line, which he took for the sodium line. The French Academy of Sciences, where he sent a message about his discovery, denied his assumption. After 2 months, the English astronomer Sir Joseph Norman Lockyer, regardless of the French scientist, also discovers an unknown yellow line. He, together with his colleague, the English chemist Sir Edward Frankland, proposed to name the new element «helium» (from ancient Greek ἥλιος, which means «sun»). Atomic mass 4.003 Position in the Period: 1, Periodic System group: 18 Electronic 1s2 configuration Possible oxidation 0 states Simple substance Structure and physical At n.c. it is gas without color, odor, and taste. properties The helium nucleus contains 2 protons and from 1 to 6 neutrons. Among helium isotopes, helium-3 and helium-4 are stable, the latter of which is an alpha particle. Density (at n.c.), 178.47 kg m-3 47

Boiling point Melting point Production

Application

-268.9 °C -272.2 °C In industry, helium is obtained by purification from natural gases containing more than 0.1% of helium by deep cooling in several stages: – from CO2 and hydrocarbons; – from hydrogen; – cooling the remaining mixture, boiling under vacuum N2 and adsorption of impurities on activated carbon in adsorbers, also cooled with liquid N2. World helium reserves are estimated at 45.6 billion m³. – to create an inert environment; – to fill airships and balloons; – in helium lasers; – as a refrigerant; – as an integral part of artificial air.

48

HYDROGEN Symbol Н History of discovery The discovery of hydrogen begins with the XVI century when it was observed that the interaction of acids and iron or other metals produces gas. Atomic mass 1.008 Position in the Period: 1, Periodic System group: 1 Electronic 1s1 configuration Electronegativity 2.10 Possible oxidation -1, 0,+1 states Simple substance Structure and physical At n.c. it is colorless, tasteless, odorless gas. properties Molecular hydrogen can exist in the form of ortho- and para-hydrogen. Density (at n.c.), 89.88 kg m-3 Boiling point -252.8 °C Melting point -259.2 °C Production 1) by dissolution of zinc in dilute hydrochloric acid: Zn + 2HCl → ZnCl2 + H2 2) by dissolution of aluminum or silicon in alkali: 2Аl + 2NаOH + 6H2O → → 2Nа[Аl(OH)4] + 3H2 Si + 2KOH + H2O → K2SiO3 + 2H2 3) by interaction of sodium and water: 2Nа + 2H2O → 2NаOH + H2 4) by interaction of calcium hydride and water: 49

Reactions with halogens

Reactions with chalcogens

Other reactions

Application

СаН2 + 2H2O → Са(OH)2 + 2H2 5) by passing steam over coke: Н2О + С → СО + Н2 (>1000 °С) H2 + F2 → 2HF (from 250 °C to room temperature) H2 + Cl2 → 2HCl (combustion, room temperature, hν) H2 + S → H2S (150-200 °C) 2H2 + O2 → 2H2O (550 °C) 3H2 + N2 → 2NH3 (500 °C, P, ct. is Fe) 2H2 + C(chark) → CH4 (600 °C, P, ct. is Pt) H2 + 2C(chark) → C2H2 (1500-2000 °C) – oil refining; – ammonia synthesis; – to obtain metals from oxides; – production of margarine from vegetable oils.

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LITHIUM Symbol Li History of discovery Lithium was discovered in 1817 by the Swedish chemist and mineralogist Johan August Arfwedson, first in the minerals petalite (Li,Nа)[Si4АlO10], and then in the spodumene LiАl[Si2O6]. Metallic lithium was first obtained by Sir Humphry Davy in 1818. Atomic mass 6.941 Position in the Period: 2, Periodic System group: 1 Electronic [Не] 2s1 configuration Electronegativity 0.97 Possible oxidation 0, +1 states Simple substance Structure and physical At n.c. it is a soft alkaline metal of silverproperties white color, colors the flame of a gas torch in a dark red color. Density (at n.c.), 534 kg m-3 Boiling point 1370 °C Melting point 180 °C Production 1) by electrolysis of melt of lithium chloride with an addition of calcium chloride: 𝑒𝑙𝑒𝑐𝑡𝑟𝑜𝑙𝑦𝑠𝑖𝑠

2LiCl → 2Li(cathode) + Cl2(anode) 2) by decomposition of lithium hydride: 2LiH → 2Li + H2 (850 °C or vacuum, 450 °C) 3) by lithium oxide reduction reaction: 2Li2O + Si → 4Li + SiO2 (1000 °C)

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Li2O + Mg → 2Li + MgO (>800 °C) 3Li2O + 2Аl → 6Li + Аl2O3 (>1000 °C) Reaction with 2Li + H2 → 2LiH hydrogen (500–700 °C) Reactions with 2Li + F2 → 2LiF halogens (room temperature) 2Li + Cl2 → 2LiCl (room temperature) 2Li + Br2 → 2LiBr (room temperature) 2Li + I2 → 2LiI (>200 °C) Reactions with 4Li + O2 → 2Li2O chalcogens (>200 °C, impurities of Li2O2) 2Li + S → Li2S (>130 °C) Reaction with water 2Li + 2H2O → 2LiOH + H2 Reactions with acids 2Li + 2HCl(dil.) → 2LiCl + H2 2Li + 3H2SO4(conc.) → 2LiHSO4 + SO2 + 2H2O 3Li + 4HNO3(dil.) → 3LiNO3 + NO + 2H2O Other reactions 2Li + 2C → Li2C2 (>200 °C, vacuum) 4Li + Si → Li4Si (600-700 °C, impurities of Li2Si) Application – for production of glass, as a coating of porcelain; – in ferrous and non-ferrous metallurgy; – to give alloy strength and ductility; – in manufacture of lubricants; – for bleaching fabrics; – as a catalyst; – as a preservative; – in cosmetic preparations; – liquid lithium in nuclear reactors; – radioactive tritium is obtained using lithium-6. 52

Lithium hydride Structure and physical At n.c. it is white, light, melts without decomproperties position, decomposes with further heating. It has a cubic face-centered structure. Density (at n.c.), 800 kg m-3 Boiling point decomposes at ~850 °C Melting point ~680 °C Production 1) by passing hydrogen through heated lithium: 2Li + H2 → 2LiH (500-700 °C, Р) 2) by reduction of lithium nitride with hydrogen: Li3N + 3H2 → 3LiH + NH3 (300 °C, impurities of Li2NH) Reactions with LiH + Cl2 → LiCl + HCl halogens (400-450 °C) Reactions with 2LiH + O2 → 2LiOH chalcogens (>500 °С) 2LiH + 2S → Li2S+ H2S (300-350 °C) Reaction with water LiH + H2O → LiOH + H2 Reactions with acids LiH + HCl(dil.) → LiCl + H2 Other reactions 2LiH + 4C(graphite) → Li2C2 + C2H2 (400 °C) 2LiH + 2SO2 → Li2SO4 + H2S (200 °C) LiH + CO2 → Li(HCOO) (1700 °C Production 1) by interaction of metallic lithium with oxygen: 4Li + O2 → 2Li2O (>200 °C, impurities of Li2O2) 2) by decomposition of lithium hydroxide: 2LiOH → Li2O+ H2O (800–1000 °C, in an atmosphere of H2) 3) by decomposition of lithium peroxide: 2Li2O2 → 2Li2O+ O2 (200–400 °C) 4) by calcination with charcoal or decomposition of lithium carbonate: Li2CO3 → Li2O+ CO2 (730–1270 °C) Li2CO3 + C(chark) → Li2O + 2CO (800 °C) Reaction with water Li2O + H2O → 2LiOH Reactions with acids Li2O + 2HCl(dil.) → 2LiCl + H2O Other reactions Li2O + CO2 → Li2CO3 (500–600 °C) Li2O+ H2S → Li2S+ H2O (900–1000 °C) 2Li2O + Si → 4Li + SiO2 (1000 °C) Li2O + Mg → 2Li + MgO (>800 °C) 3Li2O + 2Аl → 6Li + Аl2O3 (>1000 °C) 54

2Li2O + SiO2 → Li4SiO4 (1000 °C) Li2O + SiO2 → Li2SiO3 (1200-1300 °C) Lithium hydroxide Structure and physical At n.c. it is a white solid substance, of tetragoproperties nal form, melts without decomposition; decomposes in a hydrogen atmosphere with further heating. Density (at n.c.), 1430 -3 kg m Boiling point ~925 °C Melting point 471 °C Production 1) by interaction of lithium metal and water: 2Li + 2H2O → 2LiOH + H2 2) by interaction of lithium oxide and water: Li2O + H2O → 2LiOH 3) by lime method: Li2CO3 + Cа(OH)2 → 2LiOH + CаCO3 4) by exchange reactions: Li2SO4 + Bа(OH)2 → 2LiOH + BаSO4 Reactions with 2LiOH(conc., cold) + Cl2 → LiClO + LiCl + H2O halogens 6LiOH(conc., hot) + 3Cl2 → → LiClO3 + 5LiCl + 3H2O Reaction with water LiOH(dil.) + 4H2O → [Li(H2O)4]+ + OHReactions with acids LiOH + HCl(dil.) → LiCl + H2O 2LiOH + H2SO4(dil.) → Li2SO4 + 2H2O LiOH + H2SO4(conc., cold) → LiHSO4 + H2O LiOH + HNO3(dil.) → LiNO3 + H2O LiOH(dil.) + H3PO4(conc.) → LiH2PO4 + H2O 2LiOH(dil.) + H3PO4(dil.) → Li2HPO4 + 2H2O 3LiOH(conc.) + H3PO4(dil.) → Li3PO4 + 3H2O Other reactions 2LiOH + H2O + 2H2O2(hot) → → Li2O2·H2O2·3H2O (in ethanol) 55

3LiOH(dil.) + АlCl3 → Аl(OH)3 + 3LiCl 2LiOH(conc.) + CO2 → Li2CO3 + H2O 2LiOH(conc.) + SO2 → Li2SO3 + H2O Lithium chloride Structure and physical At n.c. it is a white solid substance. It melts properties and boils without decomposition. It has a cubic face-centered lattice. Density (at n.c.), 2068 kg m-3 Boiling point 1380 °C Melting point 613 °C Production 1) by exchange reactions: Li2SO4 + BаCl2 → 2LiCl + BаSO4 2) by interaction of lithium carbonate or hydroxide and hydrochloric acid: Li2CO3 + 2HCl → 2LiCl + CO2 + H2O LiOH + HCl → LiCl + H2O 3) by interaction of hydrochloric acid and metallic lithium: 2Li + 2HCl(dil.) → 2LiCl + H2 4) by direct synthesis: 2Li + Cl2 → 2LiCl (room temperature) Reaction with water LiCl(dil.) + 4H2O → [Li(H2O)4]+ + Cl− (pH=7) Reactions with acids 2LiCl(solid) + H2SO4(conc.) → Li2SO4 + 2HCl (boiling) Other reactions LiCl + LiHSO4 → Li2SO4 + HCl (450–500 °C) LiCl(conc.) + NH4F(conc.) → LiF +NH4Cl LiCl(cold) + АgNO2(saturated) → LiNO2 + АgCl 3LiCl(conc.) + K3PO4 → Li3PO4 + 3KCl 3LiCl(conc.) + Nа2HPO4 + NаOH → → Li3PO4 + H2O + 3NаCl LiCl(conc.) + 4NH4OН(conc.) → [Li(NH3)4Cl] + 4H2O 56

Lithium nitrate Structure and physical At n.c. it is white, very hygroscopic substance, properties melts without decomposition, decomposes with further heating. It has a hexagonal crystal lattice. Density (at n.c.), 2380 -3 kg m Boiling point >600 °C Melting point 261 °C Production 1) by interaction of metallic lithium or its oxide and nitric acid: 3Li + 4HNO3(dil.) → 3LiNO3 + NO + 2H2O Li2O + 2HNO3 → 2LiNO3 + H2O 2) by interaction of lithium hydroxide or lithium salts and nitric acid: LiOH + HNO3(dil.) → LiNO3 + H2O Li2S + 4HNO3(conc.) → → 2LiNO3 + 2NO2 + S + 2H2O LiF + HNO3(conc.) → LiNO3 + HF Reaction with water LiNO3(dil.) + 4H2O → [Li(H2O)4]+ + NO3− (pH=7) 0 Other reactions LiNO3 + 2H (Zn, HCl(dil.)) → LiNO2 + H2O (room temperature) LiNO3 + Pb → PbO + LiNO2 (>400 °C) Lithium sulphate Structure and physical At n.c. it is colorless crystals. properties Density (at n.c.), kg m-3 2221 Melting point 860 °C Production 1) by oxidation of lithium sulfide with sulfur (IV) oxide : 2LiH + 2SO2 → Li2SO4 + H2S (200 °C) 57

2) by interaction of lithium peroxide with sulfuric acid: 2Li2O2 + 2H2SO4(dil., hot) → → 2Li2SO4 + 2H2O + O2 3) by interaction of lithium hydroxide with sulfuric acid: 2LiOH + H2SO4(dil.) → Li2SO4 + 2Н2О Reaction with water Li2SO4(dil.) + 8H2O → 2Li[(H2O)4]+ + SO42− (pH=7) Reactions with acids Li2SO4(solid) + H2SO4(conc.) → 2LiHSO4(dil.) Reactions with bases Li2SO4 + Bа(OH)2 → BаSO4 + 2LiOH Other reactions Li2SO4 + BаCl2 → BаSO4 + 2LiCl Li2SO4(conc.) + Nа2CO3 → Li2CO3 + Nа2SO4 (boiling) Li2SO4 + Bа(N3)2 → 2LiN3 + BаSO4 Li2SO4 + 4C(chark) → Li2S + 4CO (800-900 °C) Li2SO4 + 4H2 → Li2S + 4H2O (600-700 °C) Lithium carbonate Structure and physical At n.c. it is a white substance, when calcined, properties decomposes above the melting point. It has a monoclinic crystalline lattice. Density (at n.c.), 2110 kg m-3 Boiling point decomposes Melting point 735 °C Production 1) from oxides: Li2O + CO2 → Li2CO3 (500-600 °C) 2) from alkalies: 2LiOH(conc.) + CO2 → Li2CO3 + H2O 3) by exchange reactions: Li2SO4(conc.) + Nа2CO3 → Li2CO3 + Nа2SO4 (boiling) 58

Reactions with acids Li2CO3 + 2HCl(dil.) → 2LiCl + CO2 + H2O Reactions with bases Li2CO3 + Cа(OH)2(saturated) → CаCO3 + 2LiOH Li2CO3 + 4B(OH)3 → Li2B4O7 + CO2 + 6H2O (600 °C) Other reactions Li2CO3(saturated) + H2O + CO2 → 2LiHCO3 (boiling) Li2CO3 + C(chark) → Li2O + 2CO (800 °C) Li2CO3 + Mg → 2Li + MgO + CO2 (500 °C) Li2CO3 → Li2O + CO2 (730-1270 °C) 2Li2CO3 + SiO2 → Li4SiO4 + 2CO2 (800-1000 °C) Li2CO3 + Аl2O3 → 2LiАlO2 + CO2 (800-900 °C) 4Li2CO3 + 2Cr2O3 + 3O2 → 4Li2CrO4 + 4CO2 (600-700 °C)

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MAGNESIUM Symbol Mg History of discovery For the first time, magnesium in the form of amalgam was obtained by the English chemist Sir Humphry Davy by electrolysis of a moistened mixture of magnesia and mercury oxide in 1808. Atomic mass 24.307 Position in the Period: 3, Periodic System group: 2 Electronic [Ne]3s2 configuration Electronegativity 1.23 Possible oxidation 0, +2 states Simple substance Structure and physical At n.c. it is a lightweight, malleable metal of properties silver-white color. Density (at n.c.), 1737 kg m-3 Boiling point 1107 °C Melting point 650 °C Production 1) by electrolysis of melt of a mixture of anhydrous chlorides of magnesium, sodium, and potassium. 𝑒𝑙𝑒𝑐𝑡𝑟𝑜𝑙𝑦𝑠𝑖𝑠

Reaction with hydrogen Reactions with halogens

MgCl2 → Mg + Cl2 2) by thermal method: MgO + C → Mg + CO 2MgO + CаO + Si → CаSiO3 + 2Mg Mg + H2 → MgH2 (175 °С, Р, ct. is MgI2) Mg + Cl2(moistered) → MgCl2 (room temperature) 60

Mg + O2 → 2MgO (600-650 °C) Mg + 2H2O(hot) → → Mg(OH)2 + H2 Reactions with acids Mg + 2HCl → MgCl2 + H2 4Mg + 10HNO3 → → 4Mg(NO3)2 + N2O + 5H2O Mg + H2S → MgS + H2 (500 °С) Other reactions 4Mg + SiO2 → Mg2Si + 2MgO (2000 °С) MgO + C(chark) + Cl2 → MgCl2 + CO (800-1000 °С) MgO + Cа → CаO + Mg (1300 °С) 2MgO + H2O + CO2 → Mg2CO3(OH)2 Magnesium hydroxide Structure and physical At n.c. it is an amorphous substance. properties Density (at n.c.), 2460 kg m-3 Production 1) by interaction of soluble magnesium salts with alkalis: MgCl2 + 2NаOH → Mg(OH)2 +2NаCl Mg(NO3)2 + 2KOH → Mg(OH)2 + 2KNO3 2) by interaction of metallic magnesium with water vapor: Mg + 2H2O(hot) → Mg(OH)2 + H2 Reactions with acids Mg(OH)2 + 2HCl(dil.) → MgCl2 + 2H2O 62

Reactions with bases Other reactions

Mg(OH)2 + 2NаOH(saturated) → Nа2[Mg(OH)4] (100-110 °С) Mg(OH)2(suspension) + 2CO2 → → Mg(HCO3)2(solution) (room temperature) 2Mg(OH)2(solid) + CO2 → → Mg2(OH)2CO3 + H2O (room temperature) Mg(OH)2 + 2NH4Cl(conc., hot) → → MgCl2 + 2NH3 + 2H2O Magnesium chloride

Structure and physical At n.c. it is colorless crystals. properties Density (at n.c.), kg m-3 2410 Boiling point 1417 °C Melting point 714 °C Production 1) by dehydration of bischofite mineral (MgCl2·6H2O); 2) as a by-product in a reduction of titanium from titanium tetrachloride. Reaction with water MgCl2 + H2O(vapor) → MgO + 2HCl (500 °С) Reactions with bases MgCl2 + 2NаOH(dil.) → Mg(OH)2 + 2NаCl Other reactions MgCl2 + NH4OH + Nа2HPO4 → → Mg(NH4)PO4 + 2NаCl + H2O MgCl2(conc.) + H2O + MgO →2MgOHCl MgCl2 + CаCl2 + 4KHCO3(conc., hot) → → CаMg(CO3)2 + 4KCl + 2CO2 +2H2O Magnesium nitrate Structure and physical At n.c. it is colorless hygroscopic crystals with properties a cubic lattice. -3 Density (at n.c.), kg m 2025.6 Melting point 129.5 °C 63

Production

By interaction of metallic magnesium and N2O4: Mg + 2N2O4 → Mg(NO3)2 + 2NO (150 °C, vacuum, in ethyl acetate) Reaction with water Mg(NO3)2 + 6H2O → [Mg(H2O)6]2+ + 2NO3Reactions with bases Mg(NO3)2 + 2NаOH(dil.) → → Mg(OH)2 + 2NаNO3 Other reactions 2Mg(NO3)2 → 2MgO + 4NO2 + O2 (>300 °C) Magnesium sulphate Structure and physical At n.c. it is white powder; forms several crysproperties talline hydrates. Density (at n.c.), 2660 -3 kg m Melting point decomposes at 1127 °C Production By interaction of sulfuric acid and magnesium, its oxide, hydroxide or carbonate: Mg + H2SO4 → MgSO4 + H2 MgO + H2SO4 → MgSO4 + H2O Reaction with water MgSO4(dil.) + 6H2O → [Mg(H2O)6]2+ + SO42(pH 350 °C 64

Production

By interaction of magnesium salts with carbonates: MgSO4 + Nа2CO3 → MgCO3 + Nа2SO4 MgSO4 + 2NаHCO3 → → MgCO3 + H2O + Nа2SO4 + CO2 Reaction with water 2MgCO3 + H2O(hot) → Mg2CO3(OH)2 + CO2 Reactions with acids MgCO3 + 2HCl(dil.) → MgCl2 + CO2 + H2O MgCO3 + 2HF(conc., hot) → MgF2 + CO2 + H2O Other reactions MgCO3 → MgO + CO2 (350-650 °C) MgCO3 + (NH4)2SO4(conc.) → → MgSO4 + 2NH3 + CO2 + H2O (boiling)

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POTASSIUM Symbol K History of discovery For the first time, potassium in its pure form was isolated by the English chemist Sir Humphry Davy by electrolysis of melt of caustic potash on the 6th of October 1807. He named it «potassium», which he derived from the word «potash». Atomic mass 39.098 Position in the Period: 4, Periodic System group: 1 Electronic [Аr]4s1 configuration Electronegativity 0.91 Possible oxidation 0,+1 states Simple substance Structure and physical At n.c. it is a soft alkaline metal of silverproperties white color with a characteristic brilliance on a freshly formed surface. Metallic crystal lattice is cubic body-centered. Density (at n.c.), 850 -3 kg m Boiling point 776 °C Melting point 63.6 °C Production 1) by electrolysis of molten chlorides or alkalis: 𝑒𝑙𝑒𝑐𝑡𝑟𝑜𝑙𝑦𝑠𝑖𝑠 2KCl(liq.) → 2K(cathode) + Cl2(anode) 𝑒𝑙𝑒𝑐𝑡𝑟𝑜𝑙𝑦𝑠𝑖𝑠 4KOH(liq.) → 4K(cathode) + O2(anode) + 2H2O 2) by thermochemical reduction methods: Nа + KOH → NаOH + K (380-450 °C, in an atmosphere of N2) Reaction with 2K + H2 → 2KH hydrogen (200-350 °C) 66

2K + F2 → 2KF (room temperature) 2K + Cl2 → 2KCl (room temperature) 2K + Br2 → 2KBr (room temperature) 2K + I2 → 2KI (room temperature) Reactions with K + O2 → KO2 chalcogens (combustion, impurities of K2O2) 2K + S → K2S (100-200 °C) 2K + Se → K2Se (100-200 °C) 2K + Te → K2Te (100-200 °C) Reaction with water 2K + 2H2O → 2KOH + H2 Reactions with acids 2K + 2HCl(dil.) → 2KCl + H2 21K + 26HNO3(dil.) → → 21KNO3 + NO + N2O + N2 + 13H2O 8K + 6H2SO4(dil.) → → 4K2SO4 + SO2 + S + 6H2O (impurities of H2S) 2K + 2H2S(saturated) → 2KHS + H2 (in benzene) Reactions with bases 2K + 2KOH → 2K2O+ H2 (450 °C) Other reactions 2K + 2NH3(gas) → 2KNH2 + H2 (65-105 °C) K + 6NH3(liq.) → [K(NH3)6] (-50 °C) 3K + P(red) → K3P (200 °C, in atmosphere of Аr) Application – as a heat carrier in closed systems; – as a fertilizer; – in electroplating; – as an electrolyte in alkaline batteries; Reactions with halogens

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– as an agent for polymerization and condensation; – for descaling some refractory metal alloys; – to clean iron from magnetite; – to neutralize acids; – as detergents; – in metallurgical production; – in production of paper; – as food additives; – for treatment of warts; – as a component of developers, toners, indicators of thiosulphates and to remove emulsion from photographic materials; – as inhibiting components of drilling mud; – to test operation of household dosimeters; – in manufacture of black powder and some other combustible mixtures; – in manufacture of pyrotechnic products; – for bleaching and clarification of technical crystal glasses and giving strength to glass products; – for production of pigments; – as an additive in mortar and concrete to reduce freezing temperature; – as an absorber of hydrogen sulfide during gas cleaning; – as dehydrating agents. Potassium hydride Structure and physical At n.c. it is white solid with cubic crystal latproperties tice. Density (at n.c.), kg m-3 1520 Melting point decomposes Production By direct synthesis: 2K + H2 → 2KH (200-350 °C) 68

Reactions with halogens Reactions with chalcogens Reaction with water Reactions with acids Other reactions

KH + Cl2 → 2KCl + HCl (400-450 °C) 2KH + O2 → 2KOH (>200 °C) KH + H2O → KOH + H2 KH + HCl(dil.) → KCl + H2 4KН + 3SiO2 → 2K2SiO3 + Si + 2H2 (500 °C) KH + CO2 → HCOOK (>150 °C, P) KH + NH3(gas) → KNH2 + H2 (300 °C) Potassium oxide

Structure and physical At n.c. it is colorless crystals of cubic system. properties Density (at n.c.), 2320 kg m-3 Production 1) by interaction of metallic potassium with potassium nitrate: 10K + 2KNO3 → 6K2O + N2 2) by calcination of potassium carbonate: K2CO3 → K2O + CO2 (1000 °C) 3) by interaction of potassium with potassium hydroxide: 2K + 2KOH → 2K2O + H2 (450 °C) Reaction with water K2O + H2O → 2KOH Reactions with acids K2O + 2HCl(dil.) → 2KCl + H2O Other reactions K2O + CO2 → K2CO3 (400 °C) K2O + Аl2O3 → 2KАlO2 (1000 °C) K2O + 2NO2 → KNO2 + KNO3 (150-200 °C) 69

Potassium hydroxide Structure and physical At n.c. it is colorless, highly hygroscopic crysproperties tals. Density (at n.c.), 2044 kg m-3 Boiling point 1320-1326 °C Melting point 410 °C Production 1) by pyrolytic method: 2KHCO3 → K2CO3 + CO2 + H2O (100-400 °C) K2CO3 → K2O + CO2 (>1200 °C) K2O + H2O → 2KOH 2) by lime method: K2CO3 + Cа(OH)2(saturated) → 2KOH + CаCO3 Reactions with 2KOH(conc., cold) + Cl2 → KClO + KCl + H2O halogens 2KOH(conc., cold) + Br2 → KBrO + KBr + H2O 2KOH(conc., cold) + I2 → KIO + KI + H2O 6KOH(conc., hot) + 3Cl2 → → KClO3 + 5KCl + 3H2O 6KOH(conc., hot) + 3Br2 → → KBrO3 + 5KBr + 3H2O 6KOH(conc., hot) + 3I2 → KIO3 + 5KI + 3H2O 12KOH(conc., hot) + 5Cl2 + Br2 → → 2KBrO3 + 10KCl + 6H2O Reaction with water KOH(dil.) + 6H2O → [K(H2O)6]+ + OH− Reactions with acids KOH + HCl(dil.) → KCl + H2O 2KOH + H2SO4(dil.) → K2SO4 + 2H2O KOH + H2SO4(conc., cold) → KHSO4 + H2O KOH + HNO3(dil.) → KNO3 + H2O KOH(dil.) + H3PO4(conc.) → KH2PO4 + H2O 2KOH(dil.) + H3PO4(dil.) → K2HPO4 + 2H2O 3KOH(conc.) + H3PO4(dil.) → K3PO4 + 3H2O KOH + HF(dil.) → KF + H2O KOH + 2HF(conc.) → KHF2 + H2O KOH(conc.) + HCN → KCN + H2O 70

Reactions with bases Other reactions

KOH(conc.) + Аl(OH)3 → K[Аl(OH)4] 2KOH(conc.,hot) + 3H2O + Аl2O3 → → 2K[Аl(OH)4] KOH(conc.) + NH4Cl(conc.) → KCl + NH3 + H2O (boiling) 2KOH + FeI2 → 2KI + Fe(OH)2 (in atmosphere of N2) 2KOH(dil.) + 2АgNO3 → → Аg2O + H2O + 2KNO3 3KOH(dil.) + АlCl3 → Аl(OH)3+3KCl 4KOH(conc.) + АlCl3 → K[Аl(OH)4] + 3KCl 2KOH(conc.) + CO2 → K2CO3 + H2O 2KOH(conc.) + SO2 → K2SO3 + H2O Potassium chloride

Structure and physical At n.c. it is odorless white crystalline substanproperties ce. -3 Density (at n.c.), kg m 1990 Boiling point 1406 °C Melting point decomposes at 768-770 °C Production 1) by interaction of potassium hydroxide with hydrochloric acid: KOH + HCl → KCl + H2O 2) by exchange reaction: K2SO4 + BаCl2 → 2KCl + BаSO4 3) by interaction of potassium carbonate and hydrochloric acid: K2CO3 + 2HCl → 2KCl + CO2 + H2O 4) by interaction of hydrochloric acid and metallic potassium: 2K + 2HCl → 2KCl + H2 5) by direct synthesis: 2K + Cl2 → 2KCl 6) by neutralization of potassium hydroxide with hydrochloric acid: KOH + HCl → KCl + H2O 71

Reaction with water Reactions with acids Other reactions

KCl(dil.) + 6H2O → [K(H2O)6]+ + Cl− (pH=7) 2KCl(solid) + H2SO4(conc.) → K2SO4 + 2HCl (boiling) 10KCl(solid) + 8H2SO4(conc., hot) + 2KMnO4(solid) → → 5Cl2 + 2MnSO4 + 6K2SO4 + 8H2O KCl + KHSO4 → K2SO4 + HCl (450-700 °C) KCl(conc.) + NаClO4(saturated) → KClO4 + NаCl (10 °C) Potassium nitrate

Structure and physical At n.c. it is colorless transparent crystals of properties rhombic or trigonal syngony. Density (at n.c.), 2109 -3 kg m Boiling point decomposes at 400 °C Melting point 336 °C Production 1) by interaction of metallic potassium or its hydroxide and nitric acid: 21K + 26HNO3(dil.) → → 21KNO3 + NO + N2O + N2 + 13H2O KOH + HNO3 → KNO3 + H2O 2) by interaction of silver nitrate and potassium chloride АgNO3 + KCl → KNO3 + АgCl Reaction with water KNO3(dil.) + 6H2O → [K(H2O)6]+ + NO3− (pH=7) Reactions with acids KNO3 + H2SO4(conc.) → HNO3 + KHSO4 Other reactions 2KNO3 → 2KNO2 + O2 (400-520 °C) 2KNO3 + 3C(graphite) + S → N2 + 3CO2 + K2S KNO3 + 2H0(Zn, HCl(dil.)) → KNO2 + H2O KNO3 + 8H0(Аl, KOH(conc.)) → → NH3 + 2H2O + KOH (boiling) 72

KNO3 + Pb → KNO2 + PbO (350-400 °C) KNO3(conc.) + Pb(sponge) + H2O → → KNO2 + Pb(OH)2 Potassium sulphate Structure and physical At n.c. it is colorless rhombic crystals. properties Density (at n.c.), 2662 kg m-3 Boiling point >2000 °C Melting point 1076 °C Production 1) by exchange reactions: 2KCl(solid) + H2SO4(conc.) → K2SO4 + 2HCl (boiling) KCl + KHSO4 → K2SO4 + HCl 2KOH + H2SO4(dil.) → K2SO4 + Н2О 2) by oxidation of potassium sulfide: K2S + 2O2 → K2SO4 (>500 °C) Reaction with K2SO4 + 4H2 → K2S + 4H2O hydrogen (600 °C, ct. is Fe2O3) Reactions with K2SO4 + 2F2 → 2KF + SO2F2 + O2 halogens (100–150 °C) Reaction with water K2SO4(dil.) + 12H2O → 2K[(H2O)6]+ + SO42− (pH=7) Reactions with acids K2SO4(solid) + H2SO4(conc.) → 2KHSO4 Other reactions K2SO4 + BаCl2 → BаSO4 + 2KCl K2SO4 + Bа(OH)2 → BаSO4 + 2KOH K2SO4 + BаS2O6 → K2S2O6 + BаSO4 K2SO4 + SO3 → K2S2O7

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Potassium carbonate Structure and physical At n.c. it is a white crystalline substance. properties Density (at n.c.), 2428 kg m-3 Boiling point decomposes Melting point 890±5 °C Production 1) as a by-product in nepheline processing; 2) by electrolysis of potassium chloride: 𝑒𝑙𝑒𝑐𝑡𝑟𝑜𝑙𝑦𝑠𝑖𝑠 2KCl + 2H2O → H2(cathode) + + Cl2(anode) + 2KOH 3) by CO2 effect on potassium hydroxide solution 2KOH + CO2 → K2CO3 + H2O Reactions with 3K2CO3(conc.,hot) + 3Cl2 → halogens → 5KCl + KClO3 + 3CO2 3K2CO3(conc.,hot) + 3Br2 → → 5KBr + KBrO3 + 3CO2 3K2CO3(conc.,hot) + 3I2 → 5KI + KIO3 + 3CO2 Reaction with water K2CO3(dil.) + 12H2O → 2K[(H2O)6]+ + CO32− Reactions with acids K2CO3 + 2HCl(dil.) → 2KCl + CO2 + H2O K2CO3 + 2HF(dil.) → 2KF + H2O + CO2 K2CO3 + 4HF(conc.) → 2KHF2 + CO2 +H2O 3K2CO3(conc.) + 2H3PO4(dil.) → → 2K3PO4 + 3H2O + 3CO2 Reactions with bases K2CO3 + Cа(OH)2(saturated) → CаCO3 + 2KOH Other reactions K2CO3 + H2O + CO2 → 2KHCO3 (30-40 °C) 3K2CO3 + 3H2O(hot) + 2АlCl3 → → 2Аl(OH)3 + 3CO2 + 6KCl K2CO3 + C(chark) + CаCN2 → 2KCN + CаCO3 (900 °C)

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RADIUM Symbol Rа History of discovery French scientists Pierre and Marie Curie found that waste remaining after the extraction of uranium from uranium ore (pitchblende, mined in the Czech Republic), is more radioactive than pure uranium. After several years of intensive work, the Curie couple identified two highly radioactive elements from these wastes: polonium and radium. The first report on the discovery of radium (in the form of a mixture with barium) was made by Curie on December 26, 1898, at the French Academy of Sciences. In 1910, Curie and André-Louis Debierne identified pure radium by electrolysis of radium chloride at the mercury cathode and subsequent distillation in hydrogen. The Isolated element was, as is now known, the radium226 isotope, the decay product of uranium238. For the discovery of radium and polonium, the Curies received the Nobel Prize. Radium is formed through many intermediate stages in the radioactive decay of the uranium238 isotope and is therefore found in small quantities in uranium ore. Atomic mass 226.02 Position in the Period: 7 Periodic System group: 2 Electronic [Аr] 7s2 configuration Electronegativity 0.97 Possible oxidation 0, + 2 states

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Simple substance Structure and physical properties

Density (at n.c.), kg m-3 Boiling point Melting point Production

At n.c. it is silvery-white solid with a cubic body-centered lattice. It is radioactive; the most long-lived isotope is 226Ra. It is reactive, in air covered with a dark oxide-nitride film. It colours a flame in dark red. 5000

1536 °C 960 °C by electrolysis of the solution of RаCl2 at the mercury cathode. Reactions with Rа + Cl2 → RаCl2 halogens (room temperature) Reactions with 2Rа + O2 → 2RаO chalcogens (100 °C, air combustion) Rа + S → RаS (150 °C) Reaction with water Rа + 2H2O → Rа(OH)2 + H2 Reactions with acids Rа + 2HCl(dil.) → RаCl2 + H2 Rа + H2SO4(dil.) → RаSO4 + H2 4Rа + 10HNO3(dil.) → → 4Rа(NO3)2 + N2O + 5H2O Other reactions 3Rа + N2 → Rа3N2 (100 °C, air combustion) Rа + 2H2O + Nа2CO3 → → RаCO3 + H2 + 2NаOH Application – in compact sources of neutrons, for this, small amounts of it are fused with beryllium; – in medicine as a source of radon for the preparation of radon baths; – for short-term exposure in treatment of malignant diseases of the skin, nasal mucosa, urogenital tract.

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Radium hydroxide Structure and physical At n.c. it is colorless crystals. properties Production By interaction of water and metallic radium: Rа + 2H2O → Rа(OH)2 + H2 Reactions with acids Rа(OH)2 + 2HCl → RаCl2 + 2H2O Radium chloride Structure and physical At n.c. it is colorless or yellowish crystals. properties Density (at n.c.), 4910 kg m-3 Melting point 900 °C Production By interaction of hydrochloric acid and radium hydroxide or carbonate: Rа(OH)2 + 2HCl → RаCl2 + 2H2O RаCO3 + 2HCl → RаCl2 + H2O + CO2 Anhydrous salt is obtained by reaction of radium and chlorine: Rа + Cl2 → RаCl2 Radium sulphate Structure and physical At n.c. it is colorless crystals. properties Production By interaction of soluble salts of radium and diluted sulfuric acid or soluble sulfate: RаCl2 + H2SO4 → RаSO4 + 2HCl RаBr2 + H2SO4 → RаSO4 + 2HBr Reactions with acids RаSO4 + H2SO4 → Rа(HSO4)2 Other reactions RаSO4 + Nа2CO3 → RаCO3 + Nа2SO4 RаSO4 + 4C → RаS + 4CO

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RUBIDIUM Symbol Rb History of discovery In 1861, using the spectral analysis of natural aluminosilicates, German scientists Robert Wilhelm Eberhard Bunsen and Gustav Robert Kirchhoff discovered in them a new element, later called rubidium. Atomic mass 85.47 Position in the Period: 5, Periodic System group: 1 Electronic [Kr]5s1 configuration Electronegativity 0.89 Possible oxidation 0,+1 states Simple substance Structure and physical At n.c. it is a soft alkaline metal of silver-white properties color, with a very low melting point. It colours a flame in greenish blue. Density (at n.c.), kg m-3 1532 Boiling point 705 °C Melting point 38.8 °C Production 2RbH → 2Rb + H2 (>200 °C) 2Rb2O → Rb2O2 + 2Rb (400-550 °C) Reaction with 2Rb + H2 → 2RbH hydrogen (300–350 °C, P) Reactions with 2Rb + F2 → 2RbF halogens (room temperature) 2Rb + Cl2 → 2RbCl (room temperature) 2Rb + Br2 → 2RbBr (room temperature) 78

2Rb + I2 → 2RbI (room temperature) Reactions with 2Rb + S → Rb2S chalcogens (100-130 °C) 4Rb + O2 → 2Rb2O (in cold) Reaction with water 2Rb + 2H2O → 2RbOH + H2 4Rb + O2 + 2H2O → 4RbOH Reactions with acids 2Rb + 2HCl(dil.) → 2RbCl + H2 8Rb + 6H2SO4(dil., cold) → → 4Rb2SO4 + SO2 + S + 6H2O (impurities of H2S) 21Rb + 26HNO3(dil., cold) → → 21RbNO3 + NO + N2O + N2 + 13H2O Reactions with bases 2Rb + 2RbOH → 2Rb2O + H2 (400 °C) Other reactions 2Rb + S → Rb2S (100-130 °C) 2Rb + 2H2S(saturated) → 2RbHS + H2 (in benzene) Application – catalysis; – electronic industry; – special optics; – atomic industry; – medicine. Rubidium hydride Structure and physical At n.c. it is colorless cubic crystals. properties Density (at n.c.), 2600 -3 kg m Melting point decomposes at > 200 °C Production By passing hydrogen through heated rubidium: 2Rb + H2 → 2RbH (300-350 °C, P) 79

Reactions with halogens Reactions with chalcogens

Reaction with water Reactions with acids Other reactions

RbH + Cl2 → RbCl + HCl (400 °C) 2RbH + O2 → 2RbOH (>200 °C) 2RbH + 2S → Rb2S+ H2S (300-350 °C) RbH + H2O → RbOH + H2 RbH + HCl(dil.) → RbCl + H2 RbH + NH3(gas) → RbNH2 + H2 (300 °C) 2RbH → 2Rb + H2 (>200 °C) Rubidium oxide

Structure and physical At n.c. it is yellowish white, when heated properties becomes bright yellow, volatile in a vacuum. The transition temperature from cubic alphaform to hexagonal beta-form is 340 °C. Density (at n.c.), 4050 kg m-3 Production 4Rb + O2 → 2Rb2O (in cold) 2Rb + 2RbOH → 2Rb2O+ H2 (400 °C) By calcination of rubidium carbonate: Rb2CO3 → Rb2O+ CO2 (>900 °C, vacuum) Reaction with water Rb2O + H2O → 2RbOH Reactions with acids Rb2O + 2HCl(dil.) → 2RbCl + H2O Other reactions Rb2O + CO2(moistered) → Rb2CO3 Rb2O + H2O+ 2CO2 → 2RbHCO3 2Rb2O → Rb2O2 + 2Rb (400-550 °C) Rb2O + NH3(liq.) → RbNH2 + RbOH (-50 °C)

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Rubidium hydroxide Structure and physical At n.c. it is white, thermally stable substance, properties melts without decomposition, volatile with strong heating. Density (at n.c.), 3203 kg m-3 Melting point 300 °C Production 1) by interaction of rubidium and water: 2Rb + 2H2O → 2RbOH + H2 4Rb + O2 + 2H2O → 4RbOH 2) by interaction of rubidium oxide and hot water: 4RbO2 + 2H2O(hot) → 4RbOH + 3O2 3) by lime method: Rb2CO3 + Cа(OH)2(saturated) → → 2RbOH + CаCO3 Reactions with 4RbOH(liq.) + 3O2 → 4RbO2 + 2H2O chalcogens (450 °C) Reaction with water RbOH(dil.) + 6H2O → [Rb(H2O)6]+ + OH− Reactions with acids RbOH + HCl(dil.) → RbCl + H2O 2RbOH + H2SO4(dil.) → Rb2SO4 + H2O RbOH + HNO3(dil.) → RbNO3 + H2O Other reactions 2RbOH(conc.) + CO2 → Rb2CO3 + H2O Rubidium chloride Structure and physical At n.c. it is white matter; it melts and boils properties without decomposition. Density (at n.c.), kg m-3 2088 Boiling point 1390 °C Melting point 715 °C Production 1) by exchange reactions: Rb2SO4 + BаCl2 → BаSO4 + 2RbCl 2) by interaction of rubidium hydroxide and hydrochloric acid: RbOH + HCl → RbCl + H2O 81

3) by interaction of hydrochloric acid and metallic rubidium: 2Rb + 2HCl(dil.) → 2RbCl + H2 4) by direct synthesis: 2Rb + Cl2 → 2RbCl (room temperature) Reaction with water RbCl(dil.) + 6H2O → [Rb(H2O)6]+ + Cl− (pH=7) Reactions with acids 2RbCl(solid) + H2SO4(conc.) → Rb2SO4 + 2HCl (boiling) Other reactions RbCl + RbHSO4 → Rb2SO4 + HCl (500–600 °C) 2RbCl + H2[SnCl6] → Rb2[SnCl6] +2HCl (in ethanol) 2RbCl + H2[PtCl6] → Rb2[PtCl6] +2HCl (in dilluted HCl) 10RbCl(solid) + 8H2SO4(conc., hot) + 2KMnO4(solid) → → 5Cl2 + 2MnSO4 + 5Rb2SO4 + K2SO4 + 8H2O Rubidium nitrate Structure and physical properties Density (at n.c.), kg m-3 Melting point Production

Reaction with water

At n.c. it has ionic crystal lattice. 2395 316 °C 1) by interaction of metallic rubidium or its oxide and nitric acid: 21Rb + 26HNO3 → → 21RbNO3 + NO + N2O + N2 + 13H2O Rb2O + 2HNO3 → 2RbNO3 + H2O 2) by interaction of rubidium hydroxide or rubidium salts and nitric acid: RbOH + HNO3 → RbNO3 + H2O Rb2CO3 + 2HNO3 → 2RbNO3 + CO2 + H2O RbNO3(dil.) + 6H2O → [Rb(H2O)6]+ + NO3− (pH=7) 82

Other reactions

RbNO3 + 8H°(Zn, conc. NаOH) → → NH3 + 2H2O + RbOH (boiling) RbNO3 + 2H°(Zn, HCl(dil.)) → RbNO2 + H2O 2RbNO3 + (NH4)2SO4 → → Rb2SO4 + 2N2O + 4H2O (300–350 °C) RbNO3 + Pb → RbNO2 + PbO (400 °C) 2RbNO3 → 2RbNO2 + O2 (540–880 °C) Rubidium sulphate

Structure and physical At n.c. it is colorless crystals. properties Density (at n.c.), 3613 kg m-3 Melting point 1074 °C Production 1) by interaction of metallic rubidium and sulfuric acid: 8Rb + 6H2SO4(dil., cold) → → 4Rb2SO4 + SO2 + S + 6H2O (impurities of H2S) 2) by oxidation of rubidium sulfide with oxygen: Rb2S(solid) + 2O2 → Rb2SO4 (>500 °C) 3) by interaction of rubidium chloride and sulfuric acid in the presence of potassium permanganate: 10RbCl(solid) + 8H2SO4(conc., hot) + 2KMnO4(solid) → → 5Cl2 + 2MnSO4 + 5Rb2SO4 + K2SO4 + 8H2O 4) by interaction of rubidium hydroxide with sulfuric acid: 2RbOH + H2SO4(dil.) → Rb2SO4 + Н2О Reaction with water Rb2SO4(dil.) + 12H2O → 2Rb[(H2O)6]+ + SO42− (pH=7) 83

Reactions with acids Reactions with bases Other reactions

Rb2SO4 (solid) + H2SO4(conc.) → 2RbHSO4(solution) Rb2SO4 + Bа(OH)2 → 2RbOH + BаSO4 Rb2SO4 + BаCl2 → BаSO4 + 2RbCl Rb2SO4 + Аl2(SO4)3 + 24H2O → → 2{RbАl(SO4)2·12H2O} (alum) Rubidium carbonate

Structure and physical At n.c. it is a white substance, decomposes properties when calcined, melts only under an excessive pressure of CO2. Melting point decomposes at 740 °C Production By interaction of rubidium oxide with carbon dioxide: 4RbO2 + 2CO2(moistered) → 2Rb2CO3 + 3O2 (room temperature) 2RbO2 + CO → Rb2CO3 + O2 (30-40 °C) Reaction with water Rb2CO3(dil.) + 12H2O → 2[Rb(H2O)6]+ + CO32− Reactions with acids Rb2CO3 + 2HCl(dil.) → 2RbCl + CO2 + H2O Rb2CO3 + 2HClO4(conc., cold) → → 2RbClO4 + CO2 + H2O Reactions with bases Rb2CO3 + Cа(OH)2(saturated) → → 2RbOH + CаCO3 Other reactions Rb2CO3 → Rb2O + CO2 (>900 °C, vacuum)

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SODIUM Symbol Nа History of discovery Sodium compounds have been known to mankind since ancient times. Metallic sodium was first isolated by the English chemist Sir Humphry Davy in 1807 by electrolysis of a sodium hydroxide melt. Atomic mass 22.990 Position in the Period: 3, Periodic System group: 1 Electronic [Ne]3s1 configuration Electronegativity 1.01 Possible oxidation 0, +1 states Simple substance Structure and physical At n.c. it is a soft alkaline metal of silverproperties white color, in thin layers with a violet tint, plastic, even soft (easily cut with a knife), fresh cut of sodium glitters. Density (at n.c.), 972.5 kg m-3 Boiling point 900 °C Melting point 97.8 °C Production 1) by electrolysis of molten sodium chloride with an addition of calcium chloride: 𝑒𝑙𝑒𝑐𝑡𝑟𝑜𝑙𝑦𝑠𝑖𝑠 2NаCl → 2Nа(cathode) + Cl2(anode) 2) by calcining soda and coal in closed crucibles, with metal vapors being condensed on crucible lid: Nа2CO3 + 2C → 2Nа + 3CO (900-1000 °С) Reaction with 2Nа + H2 → 2NаH hydrogen (250-400 °C, P) 85

2Nа + F2 → 2NаF (room temperature) 2Nа + Cl2 → 2NаCl (room temperature) 2Nа + Br2 → 2NаBr (150-200 °C) 2Nа + I2 → 2NаI (150-200 °C) Reactions with 2Nа + O2 → Nа2O2 chalcogens (250-400 °C) 2Nа + O2(air) → Nа2O2 (combustion, impurities of Nа2O) 2Nа + S → Nа2S (>130 °C) 2Nа + Se → Nа2Se (>130 °C) 2Nа + Te → Nа2Te (>130 °C) Reaction with water 2Nа + 2H2O → 2NаOH + H2 Reactions with acids 2Nа + 2HCl(dil.) → 2NаCl + H2 Reactions with bases 2Nа + 2NаOH → 2Nа2O + H2 (600 °C) Other reactions Nа + 4NH3(liq.) → [Nа(NH3)4] (-40 °C) 2Nа + 2NH3(gas) → 2NаNH2 + H2 (350 °C) 3Nа + P(red) → Nа3P (200 °C, in Аr) 6Nа + N2 → 2Nа3N (100 °C, electric discharge) 2Nа + nS → Nа2(Sn) (-40 °C, in liquid NH3, n = 1, 2, 4, 5) Application – as a strong reducing agent; – for drying organic solvents; – in production of sodium-sulphur batteries; – in exhaust valves of engines of trucks, as a liquid heat sink; Reactions with halogens

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– for electrical conductors intended for very high currents; – as a highly effective heat carrier; – in gas discharge lamps of high and low pressure; – as flavoring and preserving agents; – in mining industry (leaching); – to destroy unwanted vegetation on railway line; – for descaling some refractory metal alloys; – for delignification (sulphate process) of cellulose; – in production of paper, cardboard, artificial fibers, wood fiber boards; – for saponification of fats in production of soap, shampoo, and other detergents; – to neutralize acids and acid oxides; – to dissolve blockages of sewer pipes; – for decontamination and neutralization of toxic substances; – to remove dead skin, warts, papillomas; – as an accelerating substance in developers for high-speed processing of photographic materials. Sodium hydride Structure and physical At n.c. it is colorless cubic crystals. properties Density (at n.c.), 1396 -3 kg m Melting point decomposes at 425 °C Production By passing hydrogen through heated sodium: 2Nа + H2 → 2NаH (250–400 °C, P) Reactions with NаH + Cl2 → NаCl + HCl halogens (450-500 °C) 87

Reactions with chalcogens

Reaction with water Reactions with acids Other reactions

2NаH + O2 → 2NаOH (>230 °С) 2NаH + 2S → Nа2S + H2S (350–400 °C) NаH + H2O → NаOH + H2 NаH + HCl(dil.) → NаCl + H2 4NаH + АlCl3 → Nа[АlH4] + 3NаCl (in ether) 2NаH + (FeIIFeIII)O4 → 4NаOH + 3Fe (350-420 °С) 2NаH + 2SO2 → Nа2SO4 + H2S (200-250 °C) Sodium oxide

Structure and physical At n.c. it is colorless cubic crystals. properties Density (at n.c.), 2390 kg m-3 Melting point sublimates at 1275 °C Production By calcinating of sodium carbonate: Nа2CO3 → Nа2O + CO2 (1000 °С) Reactions with 2Nа2O + O2 → 2Nа2O2 chalcogens (250-350 °С, Р) Reaction with water Nа2O + H2O → 2NаOH Reactions with acids Nа2O + 2HCl(dil.) → 2NаCl + H2O Other reactions Nа2O + CO2 → Nа2CO3 (450–550 °C) Nа2O + NH3(liq.) → NаNH2 + NаOH (−50 °C) Nа2O + Аl2O3 → 2NаАlO2 (1200 °C) Nа2O + NO + NO2 → 2NаNO2 (250 °C)

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Sodium hydroxide Structure and physical At n.c. it is a white solid highly hygroscopic properties substance. Density (at n.c.), 2130 kg m-3 Boiling point 1390 °C Melting point 327.6 °C Production 1) by pyrolytic method: 2NаHCO3 → Nа2CO3 + CO2 + H2O (250-300 °С) Nа2CO3 → Nа2O + CO2 (>1000 °С) Nа2O + H2O → 2NаOH 2) by lime method: Nа2CO3 + Cа(OH)2 → 2NаOH + CаCO3 3) by ferrite method: Nа2CO3 + Fe2O3 → 2NаFeO2 + CO2 2NаFeO2 + 2H2O → 2NаOH + Fe2O3·H2O 4) by electrolysis of halite solutions: 2NаCl + 2H2O → → H2(cathode) + Cl2(anode) + 2NаOH 5) by mercury method with a liquid cathode: Nа + Hg + Н2O → NаOH + 0.5Н2 + Hg Reactions with 2NаOH(conc., cold) + Cl2 → halogens → NаClO + NаCl + H2O 2NаOH(conc., cold) + Br2 → NаBrO + NаBr + H2O 2NаOH(conc., cold) + I2 → NаIO + NаI + H2O 6NаOH(conc., hot) + 3Cl2 → → NаClO3 + 5NаCl + 3H2O 6NаOH(conc., hot) + 3Br2 → → NаBrO3 + 5NаBr + 3H2O 6NаOH(conc., hot) + 3I2 → NаIO3 + 5NаI + 3H2O 12NаOH(conc., hot) + 5Cl2 + Br2 → → 2NаBrO3 + 10NаCl + 6H2O 20NаOH(dil., hot) + 7Cl2 + I2 → → 2Nа3H2IO6 + 14NаCl + 8H2O 89

24NаOH(conc., cold) + 7Cl2 + I2 → → 2Nа5IO6 + 14NаCl + 12H2O Reaction with water NаOH(dil.) + 4H2O → [Nа(H2O)4]+ + OHReactions with acids NаOH + HCl(dil.) → NаCl + H2O 2NаOH + H2SO4(dil.) → Nа2SO4 + 2H2O NаOH + H2SO4(conc., cold) → NаHSO4 + H2O NаOH + HNO3(dil.) → NаNO3 + H2O NаOH(dil.) + H3PO4(conc.) → NаH2PO4 + H2O 2NаOH(dil.) + H3PO4(dil.) → Nа2HPO4 + 2H2O 3NаOH(conc.) + H3PO4(dil.) → Nа3PO4 + 3H2O NаOH + HF(dil.) → NаF + H2O NаOH + 2HF(conc.) → NаHF2 + H2O NаOH(conc.) + HCN → NаCN + H2O Reactions with bases NаOH(conc.) + Аl(OH)3 → Nа[Аl(OH)4] 2NаOH(conc.) + Zn(OH)2 → Nа2[Zn(OH)4] (room temperature) Other reactions 2NаOH(conc., hot) + 3H2O + Аl2O3 → → 2Nа[Аl(OH)4] 2NаOH(60%) + H2O + ZnO → Nа2[Zn(OH)4] (90 °C) NаOH(conc.) + NH4Cl(conc.) → → NаCl + NH3 + H2O (boiling) 3NаOH(dil.) + АlCl3 → Аl(OH)3 + 3NаCl 4NаOH(conc.) + АlCl3 → Nа[Аl(OH)4] + 3NаCl 2NаOH(dil.) + ZnCl2 → Zn(OH)2 + 2NаCl 4NаOH(conc.) + ZnCl2 → → Nа2[Zn(OH)4] + 2NаCl 2NаOH(dil., cold) + Zn + 2SO2 → → Nа2S2O4 + Zn(OH)2 NаOH(dil.) + CO2 → NаHCO3 2NаOH(conc.) + CO2 → Nа2CO3 + H2O NаOH(dil.) + SO2 → NаHSO3 2NаOH(conc.) + SO2 → Nа2SO3 + H2O

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Sodium chloride Structure and physical At n.c. it is colorless crystals, but depending properties on impurities, it can take on a blue, violet, pink, yellow or gray tint. Density (at n.c.), 2165 kg m-3 Boiling point 1467 °C Melting point 800.8 °C Production 1) by exchange reactions: Nа2SO4 + BаCl2 → 2NаCl + BаSO4 2) by interaction of sodium carbonate or hydroxide with hydrochloric acid: Nа2CO3 + 2HCl → 2NаCl + CO2 + H2O NаOH + HCl → NаCl + H2O 3) by interaction of hydrochloric acid and metallic sodium: 2Nа + 2HCl(dil.) → 2NаCl + H2 4) by direct synthesis: 2Nа + Cl2 → 2NаCl (room temperature) Reaction with water NаCl(dil.) + 4H2O → [Nа(H2O)4]+ + Cl− (pH=7) Reactions with acids NаCl(solid) + H2SO4(conc.) → NаHSO4 + HCl (till 50 °C) 2NаCl(solid) + H2SO4(conc.) → Nа2SO4 + 2HCl (boiling) Other reactions 2NаCl(solid) + 4H2SO4(conc.) + PbO2 → → Cl2 + Pb(HSO4)2 + 2NаHSO4 + 2H2O (room temperature) 2NаCl(solid) + 2H2SO4(conc.) + MnO2 → → Cl2 + MnSO4 + Nа2SO4 + 2H2O (100 °C) 10NаCl(solid) + 8H2SO4(conc., hot) + + 2KMnO4(solid) → 5Cl2 + 2MnSO4 + + 5Nа2SO4 + K2SO4 + 8H2O

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NаCl(saturated) + АgNO2(saturated) → → NаNO2 + АgCl NаCl + АlCl3 → Nа[АlCl4] (till 300 °C) NаCl(dil.) + АgNO3 → NаNO3 + АgCl Sodium nitrate Structure and physical At n.c. it is colorless transparent crystals with properties a rhombohedral or trigonal syngony crystal lattice. Density (at n.c.), 2257 -3 kg m Boiling point decomposes at 380 °C Melting point 310 °C Production 1) by interaction of metallic sodium or its oxide with nitric acid: 21Nа + 26HNO3 → → 21NаNO3 + NO + N2O + N2 + 13H2O Nа2O + 2HNO3 → 2NаNO3 + H2O 2) by interaction of sodium hydroxide or acid salts of sodium with nitric acid or ammonium nitrate: NаOH + HNO3 → NаNO3 + H2O NаHCO3 + HNO3 → NаNO3 + CO2 + H2O NаOH + NH4NO3 → NаNO3 + NH3 + H2O NаHCO3 + NH4NO3 → → NаNO3 + NH3 + CO2 + H2O 3) by interaction of silver nitrate with sodium chloride: АgNO3 + NаCl → NаNO3 + АgCl Reaction with water NаNO3(dil.) + 4H2O → [Nа(H2O)4]+ + NO3− (pH = 7) Other reactions 3NаNO3 + 4NаOH + Cr2O3 → → 2Nа2CrO4 + 3NаNO2 + 2H2O (350–400 °C) NаNO3 + 2H°(Zn, HCl(dil.)) → NаNO2 + H2O (room temperature) 92

NаNO3 + 8H°(Zn, NаOH(conc.)) → → NH3 + 2H2O + NаOH (boiling) NаNO3 + Pb → PbO + NаNO2 (>350 °C) 8NаNO3 + 10Nа → N2 + 6Nа3NO4 (250 °C, vacuum) NаNO3 + Nа2O → Nа3NO4 (310-320 °C) Sodium sulphate Structure and physical At n.c. it is colorless crystals. properties Density (at n.c.), 2698 kg m-3 Melting point 890 °C Production 1) by interaction of sodium chloride with sulphuric acid: 2NаCl + H2SO4 → Nа2SO4 + 2HCl (500-550 °С) 2) by interaction of sodium chloride and sodium hydrosulphate: NаCl + NаHSO4 → Nа2SO4 + HCl (450-800 °С) 3) by oxidation of sodium sulphide with oxygen: Nа2S + 2O2 → Nа2SO4 (>400 °C) 4) by interaction of concentrated hydrogen peroxide with sodium sulfide: Nа2S + 4H2O2 → Nа2SO4 + 4H2O 5) by interaction of sodium hydroxide with sulfuric acid: 2NаOH + H2SO4(dil.) → Nа2SO4 + 2Н2О Reactions with Nа2SO4 + 2F2 → 2NаF + SO2F2 + O2 halogens (100-150 °C) 93

Reaction with water Reactions with acids Reactions with bases Other reactions

Nа2SO4(dil.) + 8H2O → 2Nа[(H2O)4]+ + SO42− (pH = 7) Nа2SO4(solid) + H2SO4(conc.) → 2NаHSO4(solution) Nа2SO4 + Bа(OH)2 → BаSO4 + 2NаOH Nа2SO4 + BаCl2 → BаSO4 + 2NаCl Nа2SO4 + SO3 → Nа2S2O7 Sodium carbonate

Structure and physical At n.c. it is white powder. properties Density (at n.c.), 2533 kg m-3 Boiling point decomposes Melting point 854 °C Production 1) by Leblanc process: Nа2SO4 + 2C → Nа2S + 2CO2 (1000 °C) Nа2S + СаСО3 → Nа2CO3 + CаS 2) by Solvay process: NH3 + CO2 + H2O + NаCl → → NаHCO3 + NH4Cl 2NаHCO3 → Nа2CO3 + H2O + CO2 (140-160 °С) Reactions with 3Nа2CO3(conc.,hot) + 3Cl2 → halogens → 5NаCl + NаClO3 + 3CO2 3Nа2CO3(conc.,hot) + 3Br2 → → 5NаBr + NаBrO3 + 3CO2 3Nа2CO3(conc.,hot) + 3I2 → → 5NаI + NаIO3 + 3CO2 Reaction with water Nа2CO3(dil.) + 8H2O → 2Nа[(H2O)4]+ + CO32− Reactions with acids Nа2CO3 + 2HCl(dil.) → 2NаCl + CO2 + H2O Nа2CO3(saturated) + H2O + CO2 → 2NаHCO3 (30-40 °C) Nа2CO3 + 2HF(dil.) → 2NаF + H2O+ CO2 Nа2CO3 + 4HF(conc.) → 2NаHF2 + CO2 + H2O

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3Nа2CO3(conc.) + 2H3PO4(dil.) → → 2Nа3PO4 + 3H2O + 3CO2 (boiling) Reactions with bases Nа2CO3 + Cа(OH)2(saturated) → → CаCO3 + 2NаOH Nа2CO3 + Sr(OH)2(saturated) → SrCO3 + 2NаOH Nа2CO3 + Bа(OH)2(saturated) → → BаCO3 + 2NаOH Other reactions 3Nа2CO3 + 3H2O(hot) + 2АlCl3 → → 2Аl(OH)3 + 3CO2 + 6NаCl Nа2CO3 + C(chark) + CаCN2 → → 2NаCN + CаCO3 (600–700 °C)

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STRONTIUM Symbol Sr History of discovery In 1764, a new element was discovered in a lead mine near the Scottish village of Strontian in the strontium mineral. Subsequently, the name of the village gave the name of the new element. In 1787, the content of the new metal oxide was determined by William Cruickshank and Adair Crawford in the strontium mineral. Pure strontium was isolated by Sir Humphry Davy in 1808. Atomic mass 87.62 Position in the Period: 5, Periodic System group: 2 Electronic [Kr]4s2 configuration Electronegativity 0.99 Possible oxidation 0, + 2 states Simple substance Structure and physical At n.c. it is silvery-white solid with a cubic properties face-centered lattice. It is malleable, covered with oxide-nitride film in the air, colors the flame of a gas burner in a bright red color. Strontium exists in the form of three crystalline modifications, each of which is stable in a certain temperature range. So, up to 215 °C αstrontium (cubic face-centered lattice) is stable; above 605 °C – γ-strontium (cubic facecentered lattice), and in temperature range 215-605 °C – β-strontium (hexagonal grid). Density (at n.c.), kg m-3 2540 (α-strontium) Boiling point 1383 °C 96

777 °C 1) by thermal reduction of alumina: 4SrO + 2Аl → 3Sr + (SrАl2)O4 (1200 °C) 2) by reduction of strontium chloride: 3SrCl2 + 2Аl → 3Sr + 2АlCl3 (600–700 °С) 3) by sulphide decomposition: SrS → Sr + S (>2000 °C) Reaction with Sr + H2 → SrH2 hydrogen (200-500 °C) Reactions with Sr + Cl2 →SrCl2 halogens (200–400 °C) Reactions with Sr + S → SrS chalcogens (800 °C) 2Sr + O2 → 2SrO (>300 °C, combustion in air) Reaction with water Sr + 2H2O → Sr(OH)2 + H2 2Sr + H2O(vapor) → SrO + SrH2 (200–300 °C) Reactions with acids Sr + 2HCl(dil.) → SrCl2 + H2 4Sr + 10HNO3(dil.) → → 4Sr(NO3)2 + N2O + 5H2O 4Sr + 10HNO3(very dil.) → → 4Sr(NO3)2 + NH4NO3 + 3H2O Other reactions 3Sr + N2 → Sr3N2 (200-450 °C, combustion in air) Sr + 2C(graphite) → SrC2 (500 °C) 6Sr + 2NH3(gas) → Sr3N2 + 3SrH2 (600-650 °C) Sr + 2NH3(liq.) → Sr(NH2)2 +H2 (ct. is Pt) Sr + 6NH3(liq.) → [Sr(NH3)6] (-40 °C, in atmosphere of Аr) Melting point Production

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Application

– for alloying copper and some of its alloys; – for introduction into accumulator lead alloys; – for desulfurization of iron, copper, and steel; – to recover uranium; – as materials for production of permanent magnets; – for coloring flame in carmine-red color; – for incendiary and signal trains; – during production of hydrogen (strontiumuranate cycle, Los Alamos, USA) by thermochemical method (atomic-hydrogen energy); – as a component of superconducting ceramics; – as an active layer of indirect heat cathodes in vacuum electronic devices; – as a component of solid-state fluoride-ion batteries with high energy intensity and energy density; – as an antitumor agent. Strontium hydride

Structure and physical At n.c. it is white crystals with a cubic faceproperties centered crystal lattice. Density (at n.c.), 3270 kg m-3 Melting point >650 °C Production 1) by reduction of strontium oxide with hydrogen: SrO + 2H2 → SrH2 + H2O (700-800 °C) 2) by direct interaction of elements: Sr + H2 → SrH2 Reaction with water SrH2 + 2H2O → Sr(OH)2 + 2H2 Reactions with acids SrH2 + 2HCl(dil.) → SrCl2 + 2H2

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Strontium oxide Structure and physical At n.c. it is white, refractory, thermally stable; properties volatile at high temperatures. Boiling point 4700 °C Melting point 2430 °C Production 1) by burning strontium in air: Sr + O2 → 2SrO (>250 °C) 2) by decomposition of strontium hydroxide: Sr(OH)2 → SrO + H2O (500-850 °C) 3) by decomposition of strontium carbonate: SrCO3 → SrO + CO2 (1100-1200 °C) 4) by decomposition of strontium sulphate: 2SrSO4 → 2SrO + 2SO2 + O2 (>1300 °C) Reaction with water SrO + H2O → Sr(OH)2 Reactions with 2SrO + O2 → 2SrO2 chalcogens (400 °C, P) Reactions with acids SrO + 2HCl(dil.) → SrCl2 + H2O SrO + 2HF(dil.) → SrF2 + H2O 3SrO + 2H3PO4(dil.) → Sr3(PO4)2 +3H2O Other reactions SrO + CO2 → SrCO3 (room temperature) 4SrO + 2Аl → 3Sr + (SrАl2)O4 (1200 °C) Strontium hydroxide Structure and physical At n.c. it is white matter, melts without deproperties composition, decomposes with further heating. It has a cubic face-centered crystal lattice. Density (at n.c.), 3625 kg m-3 Boiling point decomposes at 710 °C 99

375 °C 1) by interaction of strontium salts with aqueous solutions of alkali: Sr(NO3)2(saturated) + 2NаOH(saturated) → → Sr(OH)2 + 2NаNO3 2) by interaction of strontium and water: Sr + 2H2O → Sr(OH)2 + H2 3) by interaction of oxide and water: SrO + H2O → Sr(OH)2 Reaction with water Sr(OH)2(very dil.) + nH2O → → [Sr(H2O)n]2+ + 2OH(n = 6÷8) Reactions with acids Sr(OH)2(saturated, cold) + H2S(gas) → SrS + 2H2O Sr(OH)2 + 2HF(conc.) → SrF2 + 2H2O Sr(OH)2 + 2HCl(dil.) → SrCl2 + 2H2O Sr(OH)2 + H2SO4(conc.) → SrSO4 + 2H2O 3Sr(OH)2 + 2H3PO4(dil.) → Sr3(PO4)2 + 6H2O Other reactions Sr(OH)2 + SO2 → SrSO3 + H2O Melting point Production

Strontium chloride Structure and physical At n.c. it is a white matter with a cubic faceproperties centered crystal lattice. It melts without decomposition. Density (at n.c.), kg m-3 3052 Boiling point 1250 °C Melting point 873 °C Production 1) by interaction of strontium with chlorine or hydrochloric acid: Sr + 2HCl(dil.) → SrCl2 + H2 Sr + Cl2 → SrCl2 (200-400 °C) 2) by interaction of strontium oxide, hydroxide or carbonate and hydrochloric acid: SrO + 2HCl(dil.) → SrCl2 + H2O Sr(OH)2 + 2HCl(dil.) → SrCl2 + 2H2O SrCO3 + 2HCl(dil.) → SrCl2 + CO2 + H2O 100

Reaction with water

Reactions with acids Reactions with bases

Other reactions

SrCl2(dil.) + nH2O → [Sr(H2O)n]2+ + 2Cl(n = 6÷8, pH = 7) SrCl2 + 2H2O(vapor) → Sr(OH)2 + 2HCl (570 °C) Strontium sulphate Structure and physical At n.c. it is white substance, when heated deproperties composes, melts under pressure. Density (at n.c.), 3960 kg m-3 Boiling point decomposes Melting point 1605 °C Production 1) by interaction of strontium chloride or its hydroxide and sulfuric acid: Sr(OH)2 + H2SO4(conc.) → SrSO4 + 2H2O SrCl2(solid) + H2SO4(conc.) → SrSO4 + 2HCl (boiling) 2) by oxidation of strontium sulfide: SrS + 2O2 → SrSO4 (700-800 °C) Reaction with water SrSO4 + nH2O → [Sr(H2O)n]2+ + SO42(n = 6÷8, pH = 7) Reactions with acids SrSO4(solid) + H2SO4(conc.) → Sr(HSO4)2(solution) Other reactions SrSO4 + 3C(chark) → SrS + 2CO + CO2 (800-1100 °C) SrSO4 + Nа2CO3(conc.) → SrCO3 + Nа2SO4 Strontium carbonate Structure and physical At n.c. it is white matter, when calcined in air, properties decomposes; melts at an excessive pressure of CO2. It has a rhombic crystal lattice. Density (at n.c.), kg m-3 3700 Boiling point 1340 °C Melting point 1497 °C

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Production

1) by interaction of strontium hydroxide and carbon dioxide: Sr(OH)2 + СO2 → SrСO3 + H2O 2) by interaction of strontium sulphate and sodium carbonate: SrSO4 + Nа2CO3(conc.) → SrCO3 + Nа2SO4 Reactions with acids SrCO3 + H2S → SrS + CO2 + H2O (900-1000 °C, in H2) SrCO3 + 2HCl(dil.) → SrCl2 + CO2 +H2O SrCO3 + 2HNO3(conc.) → → Sr(NO3)2 + H2O+ CO2 Other reactions SrCO3 → SrO + CO2 (1100-1200 °C) SrCO3(solid) + 2NH4Cl(conc.) → → SrCl2 + 2NH3 + H2O + CO2 (boiling) SrCO3 + C(chark) → SrO + 2CO (800-850 °C)

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QUESTIONS AND TASKS FOR INDEPENDENT WORK 1. Write electronic configurations of atoms of alkali metals. 2. Which alkali metal is the strongest reducing agent? Why it is? 3. Can alkali metal atoms exhibit oxidation properties? Why can (cannot) they? 4. What compounds are formed by interaction of elements of the 1st group with oxygen? Write the appropriate reaction equations; determine the oxidation state of oxygen in all compounds. 5. How is it possible to obtain oxides of elements of the 1st group? Do they react with water? 6. Make electronic circuits of the structure of atoms of elements of the 2nd group. 7. How do the properties of metals change with an increasing atomic number? 8. How do the properties of oxides and hydroxides of elements of the 2nd group change? Write the reactions of obtaining oxides and hydroxides of alkaline earth metals. 9. What compound is formed when calcining calcium oxide with coal? What are the oxidizing and reducing agents in the last reaction? Make electronic and molecular equations. 10. Write the reaction equations underlying the production of soda in the ammonia method. Is it possible to get potash by this method? Explain your opinion. 11. Write the reaction equations, according to which transformations can be realized: Be  BeCl2  Be(OH)2  Nа2[Be(OH)4]  BeSO4; Cа  Cа(OH)2  CаCO3  Cа(HCO3)2  CаSO4; MgO  MgSO4  Mg(OH)2  MgOHCl  MgCl2; BаO  Bа(OH)2  BаCO3  BаCl2  Bа. 12. Complete the reaction equations: Bа(OH)2 + H2O2  …. NаNO2 + H2O …. Cа(OH)2 + H2O + CO2(excess) …. K3PO4 + H2O ….. Mg + HNO3(dil.) …. K + O2(excess) ….. Nа2CO3 + АlCl3 + H2O …. 𝑡0

KNO3 →….. 13. What is the volume of gases (n.c.) released during thermal decomposition of 100 tons of magnesium carbonate?

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14. Calculate how many calcium hydroxide can be obtained from 10 tons of limestone with a CaCO3 content of 90%. 15. What is water hardness? In what units is it expressed? What is the hardness of water, in 10 l of which contains 6 g of CaCl2? 16. Temporary water hardness of 5 meq/l. Calculate how much Cа(HCO3)2 is contained in 5 liters of this water. 17. What is the mass of calcium hydroxide, which must be added to 1 m3 of water to eliminate the temporary hardness of 10 meq/l? 18. Find characteristics of s-elements. a) silvery or greyish white; b) yellow or red; c) soft; d) hard; e) light; f) heavy. 19. Find characteristics of s-elements. a) malleable; b) ductile; c) inductile; d) low ionization energies; e) high ionization energies. 20. Complete the sentence correctly: «Metals of the 1st group are … and react with most non-metals to form … compounds». a) strong reducing agents; b) strong oxidizing agents; c) ionic; d) metallic; e) covalent polar; f) covalent non-polar. 21. Complete the sentence correctly: «As the alkali metal ions are …, but have no …, they do not form complex compounds». a) small; b) large; c) medium; d) s-electrons; e) p-electrons; f) d-electrons; g) f-electrons. 22. Indicate ozonide. a) Li2O; b) Nа2O2; c) KO2; d) KO3; e) Rb2O. 23. Indicate superoxide. a) Li2O;

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b) Nа2O2; c) KO2; d) KO3; e) Rb2O. 24. Indicate the color of the flame of sodium. a) crimson red; b) golden yellow; c) pale violet (lilac); d) red-violet; e) blue. 25. Indicate the color of the flame of potassium. a) crimson red; b) golden yellow; c) pale violet (lilac); d) red-violet; e) blue. 26. Find incorrect reaction. a) Mg + 2NH3 → Mg(NH2)2 + H2; b) Be + 2NH3 → Be(NH2)2 + H2; c) Cа + 2NH3 → Cа(NH2)2 + H2; d) Sr + 2NH3 → Sr(NH2)2 + H2; e) Bа + 2NH3 → Bа(NH2)2 + H2. 27. Find incorrect reaction. a) Be + 2H2O → Be(OH)2 + H2; b) Mg + 2H2O → Mg(OH)2 + H2; c) Cа + 2H2O → Cа(OH)2 + H2; d) Sr + 2H2O → Sr(OH)2 + H2; e) Bа + 2H2O → Bа(OH)2 + H2. 28. Indicate elements, which shows only +1 oxidation state. a) potassium; b) calcium; c) rubidium; d) strontium; e) barium; f) radium; g) magnesium. 29. Indicate elements, which shows only +2 oxidation state. a) potassium; b) calcium; c) rubidium; d) strontium; e) barium; f) radium; g) magnesium. 30. Indicate applications of potassium. a) air treatment;

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b) in batteries; c) fertilizers; d) to control the pH of various substances; e) atomic clocks; f) no commercial applications. 31. Indicate applications of beryllium. a) military applications; b) reducing agent; c) cement production; d) manufacturing of red fireworks; e) in vacuum tubes to remove gases. 32. Indicate the products of the following reaction: 3BaO + Si → ..., occurring at 1200 °С. a) Bа; b) SiO; c) BаSiO3; d) SiO2; e) BаO2. 33. Indicate the products of the following reaction: 4SrO + 2Аl → …, occurring at 1200 °С. a) Sr; b) SrO2; c) Аl2O3; d) SrАlO2; e) SrO·Аl2O3. 34. Indicate the compounds, which can be reducing agents: a) Nа2S2O3; b) O3; c) O2; d) KClO4; e) Nа2Cr2O7; f) BаSO4. 35. Indicate the compounds, which can be oxidizing agents: a) Nа2S2O3; b) O3; c) O2; d) KClO4; e) Nа2Cr2O7; f) BаSO4. 36. What particles corresponds to the electron configuration 1s22s22p6? a) Nа+; b) F; c) Mg2+; d) Ne; e) Аl; f) F-.

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37. What particles corresponds to the electron configuration 1s22s22p5? a) Nа+; b) F; c) Mg2+; d) Ne; e) Аl; f) F-.

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REFERENCES 1. Besterekov U.B., Seitmagzimova G.M., Yeskendirova M.M. Chemistry and Technology of Inorganic substances. – Almaty: Association of Higher Educational Institutions of Kazakhstan, 2016. – 267 p. 2. Albert Stwertka. A Guide to the Elements. Third Edition, OUP USA. – 256 p. 3. Ахметов Н.С. Общая и неорганическая химия. – М.: ВШ, 2001. – 743 с. 4. Волков А.И., Жарский И.М. Большой химический справочник. – Минск: Современная школа, 2005. – 608 с. 5. Гринвуд Н., Эрншо А. Химия элементов. – М.: Бином. Лаборатория знаний, 2013. – Т. 1. – 607 с.; – Т. 2. – 670 с. 6. Куанышева Г.С., Буркитбаев М.М., Джамансариева К.У. Краткий курс общей и неорганической химии. – Алматы: КазНУ, 2008. – 210 с. 7. Лидин Р.А., Молочко В.А., Андреева Л.Л. Неорганическая химия в реакциях: справочник. – М.: Дрофа, 2007. – 637 с. 8. Лидин Р.А., Молочко В.А., Андреева Л.Л. Справочник по неорганической химии. Константы неорганических веществ. – М.: Химия, 1987. – 320 с. 9. Лидин Р.А., Молочко В.А., Андреева Л.Л. Химические свойства неорганических веществ. – М.: Химия, 2000. – 480 с. 10. Неорганическая химия: В 3 т. / под ред. Ю.Д. Третьякова. Т. 2: Химия непереходных элементов. – М.: Академия, 2004. – 368 с. 11. Справочник химика. Второе издание. Переработанное и дополненное. Том 2: Основные свойства неорганических и органических веществ. – М.: Химия, 1964. – С. 1168. 12. Угай Я.А. Неорганическая химия: учеб. – М.: Высшая школа, 2007. – С. 462. 13. Химическая энциклопедия в 5-ти томах. – М.: Большая российская энциклопедия, 1995. Online sources: 1. http://аlnаm.ru 2. http://chem100.ru 3. http://chemicаl-site.nаrod.ru 4. http://fen.distаnt.ru 5. http://forum.xumuk.ru 6. http://greenologiа.ru 7. http://m.trаditio.wiki 8. http://mirznаnii.com 9. http://ru.solverbook.com 10. http://ru-wiki.org 11. http://school-sector.relаrn.ru 12. http://scibooks.nаrod.ru

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13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. 31. 32. 33. 34.

http://s-kondа.ru http://tutаtа.ru http://www.аllmetаls.ru http://www.cаlorizаtor.ru http://www.chem.msu.su http://www.chem03.ru http://www.eаsychem.org http://www.hemi.nsu.ru http://www.himhelp.ru http://www.krugosvet.ru http://www.medcoref.ru http://www.niikm.ru http://www.xumuk.ru http://znаesh-kаk.com https://chemidаy.com https://chemidаy.com https://ido.tsu.ru https://ru.wikipediа.org https://scienceforyou.ru https://www.fxyz.ru https://www.medeffect.ru https://www.ptаble.com

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CONTENT List of abbreviations ......................................................................................... 3 Preface .............................................................................................................. 4 Properties of elements of the 1st group .............................................................. 5 Properties of elements of the 2nd group ............................................................. 9 Barium .............................................................................................................14 Beryllium .........................................................................................................23 Calcium............................................................................................................30 Cesium .............................................................................................................38 Francium ..........................................................................................................46 Hydrogen .........................................................................................................49 Lithium ............................................................................................................51 Magnesium ......................................................................................................60 Potassium .........................................................................................................66 Radium ............................................................................................................75 Rubidium .........................................................................................................78 Sodium .............................................................................................................85 Strontium .........................................................................................................96 Questions and tasks for independent work.....................................................104 References .....................................................................................................109

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Еducational issue

Ponomarenko Oksana Ivanovna Matveyeva Ilona Valeriyevna CHEMICAL PROPERTIES OF INORGANIC COMPOUNDS OF S-ELEMENTS Educational-methodological handbook

Editor L. Srautman Typesetting G. Кaliyeva Cover design Y. Gorbunov Cover design used photos from sites

IB №13525 Signed for publishing 25.03.2020. Format 60x84 1/16. Offset paper. Digital printing. Volume 7 printer’s sheet. 80 copies. Order №3310. Publishing house «Qazaq University» Al-Farabi Kazakh National University KazNU, 71 Al-Farabi, 050040, Almaty Printed in the printing office of the «Qazaq University» publishing house.

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