Chemical properties of inorganic compounds of p-elements: educational-methodological handbook 9786010442443

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Chemical properties of inorganic compounds of p-elements: educational-methodological handbook
 9786010442443

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AL-FARABI KAZAKH NATIONAL UNIVERSITY

O.I. Ponomarenko I.V. Matveyeva

CHEMICAL PROPERTIES OF INORGANIC COMPOUNDS OF P-ELEMENTS Educational-methodological handbook Stereotypical publication

Almaty «Qazaq University» 2020

UDC 54 (075.8) LBC 24.2я73 Р 82 Recommended for publication by the Academic Council of the Faculty of Chemistry and Chemical Technology and Editorial and Publishing Council of al-Farabi Kazakh National University

(Protocol №5 dated 27.06.2019) Reviewer PhD Yarovaya Yelena

Р 82

Ponomarenko O.I. Chemical properties of inorganic compounds of p-elements: educational-methodological handbook / O.I. Ponomarenko, I.V. Matveyeva. – Ster. pub. – Almaty: Qazaq University, 2019. – 146 p. ISBN 978-601-04-4244-3 This educational-methodological handbook presents theoretical material on the basic physical, chemical properties of the main compounds of p-elements, their methods of preparation, and the most important areas of application. The educational-methodological handbook is intended for students of specialties "5В060600 – Chemistry", "5В011200 – Chemistry" and "5В072000 – Chemical technology of inorganic substances".

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

ISBN 978-601-04-4244-3

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CONTENT LIST OF ABBREVIATIONS ........................................................................................ 4 INTRODUCTION .......................................................................................................... 5 GENERAL CHARACTERISTICS OF P-ELEMENTS ................................................. 6 ALUMINIUM ................................................................................................................ 7 ANTIMONY .................................................................................................................. 13 ARGON .......................................................................................................................... 18 ARSENIC ....................................................................................................................... 20 ASTATINE .................................................................................................................... 26 BISMUTH ...................................................................................................................... 28 BORON .......................................................................................................................... 35 BROMINE ..................................................................................................................... 41 CARBON ....................................................................................................................... 44 CHLORINE .................................................................................................................... 48 FLUORINE .................................................................................................................... 53 GALLIUM ..................................................................................................................... 56 GERMANIUM ............................................................................................................... 61 INDIUM ......................................................................................................................... 66 IODINE .......................................................................................................................... 71 KRYPTON ..................................................................................................................... 76 LEAD ............................................................................................................................. 78 NEON ............................................................................................................................. 84 NITROGEN ................................................................................................................... 85 OXYGEN ....................................................................................................................... 92 PHOSPHORUS .............................................................................................................. 94 POLONIUM ................................................................................................................... 99 RADON .......................................................................................................................... 101 SELENIUM .................................................................................................................... 103 SILICON ........................................................................................................................ 108 SULPHUR ...................................................................................................................... 113 TELLURIUM ................................................................................................................. 120 THALLIUM ................................................................................................................... 125 TIN ................................................................................................................................. 130 XENON .......................................................................................................................... 135 REFERENCES ............................................................................................................... 145

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LIST OF ABBREVIATIONS conc. – concentrated ct. – catalyst dil. – diluted P – pressure

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INTRODUCTION Chemistry is among the natural sciences that study the world around us with all variety of its forms and phenomena occurring in them. The section “Inorganic Chemistry” is one of the main sections of chemistry. Inorganic chemistry considers the basics of theoretical inorganic chemistry, which are necessary for understanding the properties of numerous inorganic substances and materials, as well as modern views on the structure of substances, on the patterns of chemical processes. On the other hand, inorganic chemistry contains information about the history of discovery, physical and chemical properties, the most important compounds, production methods, fields of application and the potential possibilities of each of the elements of the Periodic System. This manual contains material in accordance with the curriculum for the training of specialists in the specialties: "5В011200 – Chemistry", "5В060600 – Chemistry", "5В072000 – Chemical technology of inorganic substances". This manual summarizes the properties of p-elements and their compounds, depending on the structure of their atoms, as well as methods of obtaining of these compounds (including industrial), their physical and chemical properties, areas of practical application.

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GENERAL CHARACTERISTICS OF P-ELEMENTS The general electronic formula for p-elements is ns2np1÷6, where n is the principal quantum number, equal to the number of the period in which the element is located. All valence electrons of p-elements are on the external level, so they belong to the main subgroups. Most p-elements are non-metals. Such elements as aluminum, gallium, indium, thallium, tin, lead, antimony, bismuth, and polonium are often regarded as metals, although they retain many of the properties of nonmetals. Atoms of p-elements can exhibit both positive and negative oxidation states. As a rule, atoms of p-elements exhibit variable valence, moreover, in even groups it is even, and in odd groups it is odd. In the period, as the atomic number increases, the radius of the atoms decreases, the ionization energy, and the electron affinity energy increase, i.e. the oxidizing properties increase. p-elements can exhibit both oxidizing and reducing properties, so most p-elements are capable of disproportionation reactions. In the subgroup, with an increase in the atomic number, the nonmetallic properties of the elements weaken and the metallic properties increase, the positive oxidation state decreases. The strength of hydrogen compounds in the main subgroups from top to bottom decreases due to an increase in the radius of the atom. Almost all p-elements are acid formers, moreover, the stability and strength of oxygen-containing acids grow as the oxidation state of the p-element increases.

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ALUMINIUM

Atomic mass

Аl Aluminium was first obtained in 1825 by the Danish physicist Hans Christian Oersted through the interaction of potassium amalgam on aluminium chloride and further distillation of mercury. The name of aluminium comes from the name of the natural compound of aluminum – alum. 26.982

Position in the Periodic System Electronic configuration Electronegativity Possible oxidation states

Period: 3, group: 13 [Ne]3s23p1 1.47 0, +3

Symbol History of discovery

Simple substance Structure and physical properties Density (at n. c.), kg m-3 Boiling point Melting point Production

At n. c. it is a silvery-white solid with a cubic face-centered lattice. It belongs to the group of light metals. Aluminium is ductile. 2699 2486 °С 660.1 °С 1) in laboratory: by reduction of anhydrous aluminum chloride with metallic potassium: АlCl3 + 3K → 3KCl + Аl (when heated without air) 2) in industry: by dissolution of aluminum oxide in the melt of cryolite, followed by electrolysis (chark or graphite anode electrodes): 2Аl2O3

Reactions with halogens

4Аl(cathode) + 3O2(anode)

(900 °С, in melt of Nа3[АlF6]) 2Аl + 3F2 → 2АlF3 (600 °С) 2Аl(powder) + 3Cl2 → 2АlCl3 (25 °С)

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Reactions with chalcogens Reaction with water

Reactions with acids

Reactions with bases Other reactions

Application

2Аl(powder) + 3Br2 → 2АlBr3 (25 °С) 2Аl(powder) + 3I2 → 2АlI3 (25 °С, ct. is drop of water) 4Аl(powder) + 3O2 → 2Аl2O3 2Аl + 3S → Аl2S3 (150-200 °С) It is passivated, when removing the film, the reaction proceeds according to the following equation: 2Аl + 6H2O → 2Аl(OH)3 + 3H2 2Аl + 6HCl(dil.) → 2АlCl3 + 3H2 8Аl + 30HNO3(dil.) → 8Аl(NO3)3 + 3N2O + 15Н2O 8Аl + 30HNO3(very dil.) → 8Аl(NO3)3 + 3NH4NO3 + 9Н2O 2Аl + 3H2SO4→ Аl2(SO4)3 + 3H2 2Аl + 2NаOH(conc.) + 6H2O(hot) → 2Nа[Аl(OH)4] + 3H2 2Аl + 2(NаOH·Н2O) → 2NаАlO2 + 3Н2 2Аl(powder) + N2 → 2АlN (800-1200 °С) 4Аl + Р4 → 4АlР (500-800 °С, in atmosphere of H2) 4Аl + 3С(graphite) → Аl4С3 (1500-1700 °С) 2Аl + 6HF(gas) → 2АlF3 + 3Н2 (450-500 °С) 2Аl + 3H2S → Аl2S3 + 3Н2 (600-1000 °С) 2Аl +2NH3 → 2АlN + 3Н2 (>600 °С) 8Аl + 3 (FeIIF2III)O4 → 4Аl2O3 + 9Fe (>2000 °С) 8Аl + 18Н2О + 3KNO3 + 5KОН → 8K[Аl(ОН)4] + 3NH3 (boiling) – construction material; – manufacture of kitchenware; – production of aluminium foil for food industry and packaging; – raw materials in aviation and aerospace industry; – for water cleaning; – as antacid; – vaccine production; – as a flame retardant (flame suppressor) in plastics and other materials; – jewelry; – refractory materials;

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– catalysts; – adsorbents; – inert fillers in physical research and chemical industry; – during wood processing; – production of antiperspirants; – in dyeing of fabrics, for tanning of leather.

Aluminium (III) hydroxide Structure and physical properties

Melting point Production

Reaction with water Reactions with acids Reactions with bases

Other reactions

Aluminum hydroxide is a white substance, for which 4 modifications are known: – monoclinic (γ) gibbsite; – triclinic (γ ') gibbsite (hydrargyllite); – bayerite (γ); – nordstrandite (β). There is also an amorphous aluminum hydroxide of variable composition Аl2O3·nH2O. decomposes at 200 °С а) by interaction of aluminium salts with aqueous solutions of alkali: АlCl3 + 3NаOH → Аl(OH)3 + 3NаCl b) by interaction of water-soluble aluminium salts with alkali metal carbonates: 2АlCl3 + 3Nа2CO3 + 3H2O → 2Аl(OH)3 + 6NаCl + + 3CO2 Аl(OH)3(solid) + 6H2O ↔ [Аl(H2O)6]3+ + 3OH− Аl(OH)3(solid) + 4H2O ↔ [Аl(H2O)2(OH)4]- + H3O+ Аl(OH)3 + 3HCl(dil.) → АlCl3 + 3H2O Аl(OH)3 + 3HNO3 → Аl(NO3)3 + 3H2O Аl(OH)3 + NаOH(conc.) → Nа[Аl(OH)4] Аl(OH)3 + NаOH → NаАlO2 + 2H2O (1000 °С) Аl(OH)3 → АlO(OH) + H2O (575 °С) Аl(OH)3 + 3HF(conc.) + 3NаF → Nа3[АlF6] + 3H2O

Aluminium (III) oxide Structure and physical properties

It is colorless crystals, existing in several modifications:

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Boiling point

– α-form: trigonal; – γ-form: cubic; – η-shape: cubic; – θ-form: monoclinic; – χ-form: hexagonal; – κ-form: orthorhombic; – δ-form: tetragonal or orthorhombic (not established); – β-Al2O3: not pure aluminum oxide, but a series of aluminates of alkaline and alkaline-earth metals MeO·6Аl2O3 and Me2O·11Аl2O3, where MeO is oxides of calcium, barium, strontium, etc., and Me2O is oxides of sodium, potassium, lithium, and other alkali metals. ~ 3500 °С

Melting point

2015 °С

Production

3Cu2O + 2Аl → 6Cu + Аl2O3 (1000 °С) Аl2O3 + 6HCl(conc., hot) → 2АlCl3 + 3H2O Аl2O3 + 6HF(gas) → 2АlF3 + 3H2O (450-600 °С)

Reactions with acids

Reactions with bases

Аl2O3 + 2NаOH → 2NаАlO2 + H2O (900-1100 °С) Аl2O3 + 2NаOH(conc., hot) + 3H2O → 2Nа[Аl(OH)]4

Other reactions

Аl2O3 + Nа2CO3 → 2NаАlO2 + СО2 (1000-1200 °С) Аl2O3 + 3K2S2O7 → Аl2(SO4)3 + 3K2SO4 (400-470 °С) Аl2O3 + 6KHSO4 → Аl2(SO4)3 + 3K2SO4 + 3H2O (400-550 °С) Аl2O3 + 3N2O5 → 2Аl(NO3)3 Аl2O3 + MgO → (MgАl2)O4 (1600 °С) 2Аl2O3 + 9C(chark) → Аl4С3 + 6CO (1800 °С) Аl2O3 + 3С(chark) + 3Cl2 → 2АlСl3 + 3СО (800-900 °С) Аl2O3 + 3С(chark) + N2 → 2АlN + 3СО (1600-1800 °С) 2Аl2O3

4Аl(cathode) + 3O2(anode)

(900 °С, in melt ofNа3[АlF6])

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Aluminium (III) chloride Structure and physical properties Boiling point Melting point Production

Reaction with water

Reactions with acids Reactions with bases

Other reactions

At n. c. it is white, low-melting, highly volatile with monoclinic syngony. sublimates at 179.7 °С 192.6 °С 1) in industry: by effect of a mixture of Cl2 and CO on dehydrated kaolin or bauxite in blast furnaces: Аl2O3 + 3CO + 3Cl2 → 2АlCl3 + 3CO2 2) in laboratory: Аl + FeCl3 → АlCl3 + Fe Аl(OH)3 + 3HCl → АlCl3 + 3H2 2Аl + 6HCl → 2АlCl3 + 3H2O 3CuCl2 + 2Аl → 2АlCl3 + 3Cu АlСl3(dil.) + 6Н2O(cold) → [Аl(Н2O)6]3+ + 3Сl(pH < 7) АlСl3(dil.) + 3Н2O(hot) → Аl(ОН)3 + 3HCl АlСl3(solid) + 2Н2O(air moisture) → АlСl(ОН)2 + 2НСl АlCl3(solid) + 3H2SO4(conc.) → Аl(HSO4)3 + 3HCl АlCl3 + 3NаOH(dil.) → Аl(OH)3 + 3NаCl АlCl3 + 4NаOH(conc.) → Nа[Аl(OH)4] + 3NаCl АlCl3 + 3NH4OH(conc., cold) → Аl(OH)3 + 3NH4Cl АlCl3 + 3NH4OH(conc., hot) → АlO(OH) + 3NH4Cl + H2O АlCl3 + Nа3PO4 → АlPO4(amorphous) + 3NаCl АlCl3 + 3Nа[BH4] → Аl[ВН4]3 + 3NаCl (45-50 °С) АlCl3 + NH4Cl → NH4[АlCl4] (220-250 °С) electrolysis 2АlCl3(liq.) 2Аl(cathode) + 3Cl2(anode)

Aluminium (III) nitrate Structure and physical properties Boiling point Melting point Production

At n. c. it is a white or colorless crystalline substance with a monoclinic structure. decomposes at 150 °С 73.5 °С 1) in laboratory: а) by dissolution of aluminum in dilute nitric acid: 8Аl + 30HNO3(dil.) → 8Аl(NO3)3 + 3N2O + 15Н2O

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Reaction with water Reactions with bases

Other reactions

b) by interaction of aluminum hydroxide with nitric acid: Аl(OH)3 + 3HNO3 → Аl(NO3)3 + 3H2O c) by exchange reaction of aluminum sulphate with barium or lead nitrate: Аl2(SO4)3 + 3Bа(NO3)2 → 2Аl(NO3)3 + 3BаSO4 2) in industry: by interaction of aluminum oxide or hydroxide with nitrogen (V) oxide Аl2O3 + 3N2O5 → 2Аl(NO3)3 Аl(OH)3 + 3N2O5 → Аl(NO3)3 + 3HNO3 Аl(NO3)3(dil.) + 6Н2O → [Аl(Н2O)6]3+ + 3NO3(pH < 7) Аl(NO3)3 + 4NаOH(conc.) → Nа[Аl(OH)4] + 3NаNO3 Аl(NO3)3 + 3NH4OH(conc., cold) → Аl(OH)3 + 3NH4NO3 Аl(NO3)3 + 3NH4OH(conc., hot) → АlO(OH) + 3NH4NO3 + + H2O 4Аl(NO3)3 → 2Аl2O3 +12NO2 + 3O2 (150-200 °С)

Aluminium (III) sulphate Structure and physical properties Melting point Production

Reaction with water

Reactions with bases

Other reactions

At n. c. it is white salt with a gray, blue or pink tint.

decomposes at 770 °С By interaction of aluminium or its hydroxide with sulphuric acid: 2Аl + 3H2SO4→ Аl2(SO4)3 + 3H2 2Аl(OH)3 + 3H2SO4→ Аl2(SO4)3 + 6H2O Аl2(SO4)3(dil.) + 12Н2O(cold) → 2[Аl(Н2O)6]3+ + 3SO42(pH < 7) Аl2(SO4)3(very dil.) + 6H2O → 2Аl(OH)3 + 3H2SO4 (boiling) Аl2(SO4)3 + 6NаOH(dil.) → 2Аl(OH)3 + 3Nа2SO4 Аl2(SO4)3 + 8NаOH(conc.) → 2Nа[Аl(OH)4] + 3Nа2SO4 Аl2(SO4)3 + 8NаOH → 2NаАlO2 + 3Nа2SO4 + 4H2O (900-1100 °С) Аl2(SO4)3 + 6NH4OH(conc., cold) → 2Аl(OH)3 + 3(NH4)2SO4 Аl2(SO4)3 + 6NH4OH(conc., hot) → 2АlO(OH) + 3(NH4)2SO4 + 2H2O 2Аl2(SO4)3 → 2Аl2О3 + 6SO2 + 3O2 (770-860 °С) Аl2(SO4)3 + 3Bа(NO3)2 → 3BаSO4 + +2Аl(NO3)3 Аl2(SO4)3 + 3Pb(NO3)2 → 3PbSO4 + 2Аl(NO3)3

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ANTIMONY Symbol History of discovery

Atomic mass Position in the Periodic System Electronic configuration Electronegativity Possible oxidation states

Sb Antimony has been known to mankind since ancient times. 3000 years BC it was used in the countries of the East for manufacturing of vessels. In ancient Egypt in the XIX century BC antimony compounds were used to blacken the eyebrows. Description of the properties and methods of obtaining antimony was presented in 1604 by the German alchemist Basilius Valentinus. 121.760 Period: 5, group: 15 [Kr]4d105s25p3 1.82 -3, 0, +3, +5

Simple substance Structure and physical properties

Density (at n. c.), kg m-

At n. c. it is silver white with a metallic sheen semimetal. Three allotropic modifications of antimony are known: 1) gray – crystalline; 2) amorphous – yellow; 3) black – amorphous. 6684

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Boiling point Melting point Production

Reactions with halogens Reactions with chalcogens

1625 °С 630.5 °С By calcination of sulphide ores with further reduction of oxide by coal: 2Sb2S3 + 9O2 → 6SO2 + 2Sb2O3 (340 °С) Sb2O3 + 3C → 2Sb + 3CO (800-1000 °С) 2Sb(powder) + 3Cl2 → 2SbCl3 (room temperature, impurities of SbCl5) 4Sb + 3O2 → 2Sb2O3 (650 °С)

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Reaction with water Reactions with acids

Other reactions

Application

2Sb + 5S → Sb2S5 (240-400 °С, P) 2Sb + 3S → Sb2S3 (650-700 °С) 2Sb + 3Se → Sb2Se3 (650-700 °С) 2Sb + 3Te → Sb2Te3 (650-700 °С) 2Sb + 3H2O(vapour) → Sb2O3 + 3H2 (600 °С) 2Sb + 6H2SO4(conc., cold) → Sb2(SO4)3 + 3SO2 + 6H2O 2Sb + 2HNO3(dil.) → Sb2O3 + 2NO + H2O (boiling) 2Sb + 10HNO3(conc.) → Sb2O5 + 10NO2 + 5H2O (boiling) 3Sb + 18HCl(conc.) + 5HNO3(conc.) → 3H[SbCl6] + 5NO + + 10H2O (30-40 °С) 6Sb + 6KOH + 5KClO3 → 6KSbO3 + 5KCl + 3H2O (400-550 °С) 2Sb + Nа2CO3 + 5NаNO3 → 2NаSbO3 + 5NаNO2 + CO2 (400-500 °С) Sb + 3Li → Li3Sb (alloying) Sb + 3Nа → Nа3Sb (alloying) Sb + 3K → K3Sb (alloying) 2Sb + 3Mg → Mg3Sb2 (650 °С) ‒ in the manufacture of diodes, infrared detectors, fluorescent lamps, devices with a Hall effect; ‒ component of lead alloys, increasing their hardness and mechanical strength; ‒ in the production of the babbitt alloy, possessing antifriction properties and used in sliding bearings; ‒ for doping semiconductors; ‒ in the production of refractory compounds, ceramic enamels, glass, paints, and ceramic products; ‒ in fire-resistant compositions; ‒ in match heads; ‒ as fumigants.

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Stibine Structure and physical properties Boiling point Melting point Production

Reactions with halogens Reactions with chalcogens Reactions with acids Reactions with bases Other reactions

At n. c. it is colorless flammable gas with a characteristic garlic odor. -17 °С -88.5 °С 1) by interaction of antimony (III) hydroxide with atomic hydrogen: Sb(OH)3 + 6H0 → H3Sb + 3H2O 2) by interaction of magnesium antimonide with an excess of dilute hydrochloric acid: Mg3Sb2 + 6HCl → 2H3Sb + 3MgCl2 3) by interaction of compounds containing Sb−3 with proton reagents (for example, with water): Nа3Sb + 3H2O → H3Sb + 3NаOH 4) by sequential reaction of the Sb3+ cation with substances containing the formal anion H–: 4Sb2O3 + 6LiАlH4 → 8H3Sb + 3Li2O + 3Аl2O3 5) by hydrogenation of antimony (III) chloride using sodium borohydride in ether solvents or in an aqueous medium: 4SbCl3 + 3NаBH4 → 4H3Sb + 3NаCl + 3BCl3 SbCl3 + 3NаBH4 → H3Sb + 3NаCl + 3BH3 3H3Sb + 6Cl2 → 2Sb + SbCl3 + 9HCl (room temperature) 4H3Sb + 3O2 → 4Sb + 6H2O (room temperature, burning) H3Sb + 4HCl(conc.) → H[SbCl4] + 3H2 2H3Sb + 16HNO3(conc.) → Sb2O5 + 16NO2 + 11H2O H3Sb + NаOH(conc.) + 3H2O → Nа[Sb(OH)4] + 3H2 2H3Sb → 2Sb + 3H2 (150-200 °С) 2H3Sb(liq.) + O3 → 2Sb + 3H2O (-90÷-50 °С) 2H3Sb + 3H2O + 12АgNO3 → Sb2O3 + 12Аg + 12HNO3

Antimony (V) oxide Structure and physical properties Melting point

At n. c. it is light yellow crystals of the monoclinic syngony. decomposes at 380 °С

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Production

Reaction with hydrogen Reaction with water Reactions with acids

Reactions with bases

Other reactions

1) by interaction of hot concentrated nitric acid and metallic antimony: 2Sb + 10HNO3 → Sb2O5 + 10NO2 + 5H2O 2) by interaction of hot water and antimony (V) salts: 2SbCl5 + 5H2O → Sb2O5 + 10HCl 3) by interaction of acid and hexahydroxostybates: 2K[Sb(OH)6] + 2HNO3 → Sb2O5 + 2KNO3 + 7H2O Sb2O5 + 5H2 → 2Sb + 5H2O (550-600 °С) 3Sb2O5 + 5H2O ↔ 3Sb2O5·5H2O (boiling, in dil. HCl) Sb2O5 + 10HF → 2SbF5 + 5H2O (150-170 °С) Sb2O5 + 12HCl(conc.) → 2H[SbCl6] + 5H2O Sb2O5 + 2NаOH + 5H2O → 2Nа[Sb(OH)6] Sb2O5 + 2NаOH → 2NаSbO3 + H2O (500 °С) 2Sb2O5 → 2Sb2O4 + O2 (300-450 °С) Sb2O5 + 3N2O5 → 2Sb(NO3)3O (40-50 °С)

Antimony (V) chloride Structure and physical properties Boiling point Melting point Production

Reaction with water

Reactions with acids

Reactions with bases Other reactions

At n. c. it is colorless liquid. Decomposes at 140 °С 4.0 °С By chlorination of antimony (III) chloride: SbCl3 + Cl2 → SbCl5 (74-80 °С) 2SbCl5 + 5H2O → Sb2O5 + 10HCl (100 °С) SbCl5 + H2O → SbOCl3 + 2HCl (0 °С) SbCl5(liq.) + 5HF(gas) → SbF5 + 5HCl (room temperature) SbCl5 + HCl(conc.) → H[SbCl6] SbCl5 + 6NаOH(conc.) → Nа[Sb(OH)6] + 5NаCl SbCl5 → SbCl3 + Cl2 (>140 °С)

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2SbCl5(liq.) + Sb(powder) → 5SbCl3 SbCl5(liq.) + NаCl → Nа[SbCl6]

Antimony (V) sulphate Structure and physical properties Production

Reaction with water

Reactions with acids Reactions with bases Other reactions

At n. c. it is white hygroscopic crystals. By interaction of concentrated cold sulphuric acid and metallic antimony or antimony (III) chloride: 2Sb + 6H2SO4 → Sb2(SO4)3 + 3SO2 + 6H2O 2SbCl3 + 3H2SO4 → Sb2(SO4)3 + 6HCl Sb2(SO4)3 + H2O(cold) → Sb2O(SO4)2 + H2SO4 Sb2(SO4)3 + 3H2O → Sb2O3 + 3H2SO4 (100 °С) Sb2(SO4)3 + 8HCl(conc.) → 2H[SbCl4] + 3H2SO4 Sb2(SO4)3 + H2SO4(conc., hot) → 2H[Sb(SO4)2] Sb2(SO4)3 + 8NаOH → 2Nа[Sb(OH)4] + 3Nа2SO4 Sb2(SO4)3 + K2SO4(conc.) → 2K[Sb(SO4)2] (0 °С)

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ARGON Symbol History of discovery

Atomic mass Position in the Periodic System Electronic configuration Possible oxidation states

Аr Argon is an integral part of the atmosphere, but despite this, it remained a mystery to researchers for many years. Aa early as in 1785, the English physicist and chemist Henry Cavendish discovered that with a long-term effect of electrical discharge on the air, a small gas bubble remains in the system, which does not decrease even with a prolonged continuation of the experiment. Unfortunately, he could not explain this fact and did not publish his observations. His observations were published later by the English physicist James Clerk Maxwell. Somewhat later, the British physicist and mechanic Lord Rayleigh tried to understand the reason for the greater mass of natural nitrogen compared to nitrogen obtained by decomposition of any nitrogencontaining compound. On his own, he could not understand the reasons and published a letter to scientists in the journal "Nаture". The English chemist William Ramsay responded to his question by proposing cooperation on this issue. As a result of joint work, they identified a new gas with practically inert properties (at that time, inert substances were not yet known). The official announcement of the discovery of argon was made on August 7, 1894, in Oxford, at a meeting of the British Association of Physicists, Chemists, and Naturalists. 39.948 Period: 3, group: 18 [Ne]3s23p6 0

Simple substance Structure and physical properties

At n. c. it is an inert monatomic gas that does not have color, smell, and taste.

Density (at n. c.), kg m-3

1783.7

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Boiling point

– 185.9 °С

Melting point

– 189.3 °С

Production

Argon is obtained as a by-product during separation of the atmospheric air into nitrogen and oxygen.

Application

– in medicine during operations for air purification and incision; – as an extinguishing agent in fire extinguishing gas installations; – in incandescent lamps; – for filling the inner space of glass; – in argon lasers; – as a protective environment during welding; – as a plasma generator in plasmatrons during welding and cutting; – in diving for blowing dry wetsuits.

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ARSENIC Symbol

Аs

History of discovery

Arsenic has been known to mankind since ancient times, and its discoverer is not known for certain. In some sources, they consider Albertus Magnus (1200–1280), who received arsenic by heating arsenic sulphide with soap.

Atomic mass

74.922

Position in the Periodic System

Period: 4, group: 15

Electronic configuration

[Аr]3d104s24p3

Electronegativity

2.20

Possible oxidation states

-3, 0, +3, +5

Simple substance Structure and physical properties

At n. c. it is solid, which is a brittle semimetal of steel color with a greenish tint. Modifications: – yellow arsenic – a non-metal, consisting of As4 molecules with a structure similar to white phosphorus; – gray arsenic – a semimetallic polymer; – black arsenic – a non-metallic molecular structure, similar to red phosphorus.

Density (at n. c.), kg m-3

5727 (gray arsenic)

Boiling point

sublimates at 612 °С (gray arsenic)

Melting point

~817 °С (gray arsenic)

Production

1) by sublimation of natural arsenic; 2) by thermal decomposition of arsenic pyrite: FeАsS → Аs + FeS 3) by reduction of arsenic anhydride (As2O3); 4) by heating arsenopyrite in muffle furnaces without air access; 5) by calcination of sulphide ores followed by reduction of carbon monoxide: 2Аs2S3 + 9O2 → 6SO2 + 2Аs2O3 Аs2O3 + 3C → 2Аs + 3CO

Reactions with halogens

2Аs + 5F2 → 2АSF5

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(room temperature, combustion in fluorine) 2Аs + 3Cl2 → 2АsCl3 (20-30 °С, combustion in chlorine) 2Аs + 5Cl2 + 8Н2O → 2H3АsO4 + 10HCl 2Аs + 3Br2 → 2АsBr3 (50-80 °С) 2Аs + 3I2 → 2АsI3 (boiling, in liquid CS2) Reactions with chalcogens

4Аs + 3О2 → 2АS2O3 Аs + nS → Аs2S3, АS2S5, Аs4S4 (500-600 °С, in atmosphere of N2)

Reactions with acids

2Аs + 3H2SO4(conc., hot) → Аs2O3 + 3SO2 + 3Н2О 2Аs + 6H2S2O7(oleum) → 2Аs(HSO4)3 + 3H2SO4 + 3SO2 Аs + 5HNO3(conc.) → H3АsO4 + 5NO2 + H2O Аs + 3НСl(conc.) + HNO3(conc.) → АsCl3 + NO + 2H2O

Reactions with bases

2Аs + 2NаOH(20%) + 2H2O → 2NаАsO2 + 3Н2 (boiling) 2Аs + 6KОН(20%, cold) → 2K3АsO3 + 3H2

Other reactions

Аs + 3Li → Li3Аs Аs + 3Nа → Nа3Аs Аs + 3K → K3Аs 2Аs + 3Mg → Mg3Аs2 2Аs + 3Cа → Cа3Аs2 2Аs + 3Cu → Cu3Аs2 2Аs + Cа → CаАs2 2Аs + Fe → FeАs2 2Аs + 3Zn → Zn3Аs2 (400-450 °С) Аs + Аl → АlАs Аs + Gа → GаАs Аs + In → InАs Аs + Lа → LаАs Аs + 3Nа + 3NH4Br → АsH3 + 3NаBr + 3NH3 (-40 °С, in liq. NH3) 2Аs + 6NаOH(dil.) + 5NаClO → 2Nа3АsO4 + 5NаCl + + 3H2O 2Аs + 6NаOH(dil.) + 5Н2O2(conc.) → 2Nа3АsO4 + + 8H2O 2Аs + 2BrF5 → 2АsF5 + Br2 (100-200 °С)

Application

‒ for alloying lead alloys; ‒ as reducing agents in many syntheses; ‒ in painting as paints;

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‒ for manufacture of colored glass; ‒ in leather industry as a means for removing hair from skin; ‒ in agriculture for rodent control; ‒ as drugs to combat anemia.

Arsine Structure and physical properties

At n. c. it is a toxic colorless gas. The pure gas does not smell, but its decay products give it a smell similar to garlic. Arsine molecule has the shape of a trigonal pyramid.

Boiling point

-62.5 °С

Melting point

-113.5 °С

Production

1) by hydrolysis of metal arsenides with acids: Nа3Аs + 3H2O → АsH3 + 3NаOH 2) by reduction of arsenic compounds with hydrogen: Аs2O3 + 6Zn + 6H2SO4 → 2АsH3 + 6ZnSO4 + 3H2O 3) by interaction of arsenic halides with Li[AlH4], Na[BH4] or other hydrides: 4АsCl3 + 3Li[АlH4] → 4АsH3 + 3LiCl + 3АlCl3 (in ether)

Reactions with halogens

АsH3 + 3I2 → АsI3 + 3HI (room temperature)

Reactions with chalcogens

2АsH3 + 3O2 → Аs2O3 + 3H2O

Reactions with acids

АsH3 + 3HCl(conc.) → АsCl3 + 3H2 АsH3 + 2H2SO4(conc., cold) → АsSO4(OH) + S + 3H2O АsH3 + 8HNO3(conc.) → H3АsO4 + 8NO2 + 4H2O

Other reactions

АsH3 + 3NаOH(dil.) + 4H2O2(conc.) → Nа3АsO4 + 7H2O АsH3 + 3NаOH(dil.) + 4NаClO → Nа3АsO4 + 4NаCl + 3H2O 2АsH3 + 3CuSO4(conc.) + 6NаHCO3 → Cu3Аs2 + 3Nа2SO4 + + 6CO2 + 6H2O 2АsH3 + 3ZnSO4(conc.) + 6NаHCO3 → Zn3Аs2 + 3Nа2SO4 + + 6CO2 + 6H2O АsH3 + 3HgCl2 → (HgCl)3Аs + 3HCl (boiling, in dil. HCl) 2АsH3 + 3H2O + 12АgNO3 → Аs2O3 +12Аg + 12HNO3

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Arsenic (III) oxide Structure and physical properties

In the liquid and gaseous state (up to 800 °C) exists as a dimer. In the solid state, there are 3 modifications: 1) arsenolite: cubic; 2) clauderite I: monoclinic; 3) clauderite II: monoclinic.

Boiling point

457.2 °С

Melting point

278 °С (arsenolite), 313 °С (clauderite)

Production

By burning arsenic and its derivatives with an excess of oxygen: 2Аs2S3 + 9O2 → 2Аs2O3 + 6SO2 (500 °С)

Reactions with halogens

Аs2O3 + 5H2O + 2Cl2 → 2H3АsO4 + 4HCl (boiling) Аs2O3 + 5H2O + 2Br2 → 2H3АsO4 + 4HBr (boiling) Аs2O3 + 5H2O + 2I2 → 2H3АsO4 + 4HI (boiling)

Reactions with chalcogens

2Аs2O3 + 9S → 2Аs2S3 + 3SO2 (300 °С)

Reaction with water

Аs2O3(solid) + Н2О(cold) ↔ 2НАsО2(saturated) Аs2O3(solid) + 3Н2О(hot) ↔ 2Н3АsО3(saturated)

Reactions with acids

Аs2O3 + 3НСl(dil.) → H3АSO3 + АsCl3 Аs2O3 + 6НСl(conc.) → 2АsCl3 + 3Н2О Аs2O3 + 6HF(gas) → 2АsF3 + 3H2O (l40-200 °С) Аs2O3 + 6HCl(gas) → 2АsCl3 + 3H2O (l40-200 °С) Аs2O3 + 4HNO3(conc.) + H2O → 2H3АsO4 + 4NO2 (boiling) Аs2O3 + 6НI(conc.) → 2АsI3 + 3H2O (room temperature) Аs2O3 + 3H2S(saturated) → Аs2S3 + 3H2O (in conc. HCl) Аs2O3 + 4HSO3F → АsF3 + SO3 + HF + Аs(HSO4)3 (55-65 °С) Аs2O3 + 3Н(РН2О2) → 2Аs + 3Н2(РНO3) (in dil. НСl)

Reactions with bases

Аs2O3 + 2NаOH(dil.) → 2NаАsO2 + H2O Аs2O3 + 6NаОН(conc.) → 2Nа3АsO3 + 3Н2О (impurities of Nа2HАsO3)

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Other reactions

Аs2O3 + Nа2СО3(conc., hot) → 2NаАsO2 + CO2 5Аs2O3 + 6H2SO4(dil.) + 4KMnO4 + 9Н2О →10H3АsO4 + + 2K2SO4 + 4MnSO4 Аs2O3 + 6NаOH + 2NаNO3 → 2Nа3АsO4 + 2NаNO2 + + 3H2O (400-500 °С) Аs2O3 + 9НСl(conc.) + 3H[SnCl3] → 2Аs + 3H2[SnCl6] + + 3Н2О Аs2O3 + 3С(chark) → 2Аs + 3СО (700 °С, impurities of СО2) Аs2O3 + 3KCN → 2Аs + 3KOCN (600-650 °С) Аs2O3 + 12Н0(Zn, HCl(dil.) or NаOH) → 2АsН3 + 3Н2О electrolysis Аs2O3 + 5Н2О 2H2(cathode) +2Н3АsO4(anode)

Arsenic acid Structure and physical properties Boiling point Melting point Production

Reaction with water

Reactions with acids Reactions with bases

At n. c. it is white transparent substance existing in the form of a crystalline hydrate. decomposes at >160 °С 35.5 °С 1) by interaction of strong oxidizing agents and arsenic or its oxide: 3Аs2O3 + 4HNO3 + 7H2O → 4NO + 6H3АsO4 Аs2O3 + 4HNO3 + H2O → 4NO2 + 2H3АsO4 (boiling) 2) by dissolving arsenic anhydride in water: Аs2O5 + 3H2O → 2H3АsO4(saturated) (80 °С) H3АsO4(saturated) + 2H2O → H7АsO6 (-30 °С) H3АsO4(dil.) + H2O → H2АsO4- + H3O+ 2H3АsO4 + 5H2S(gas) → Аs2S5 + 8H2O (0 °С, in conc. HCl) H3АsO4 + NаOH(dil.) → NаH2АsO4·H2O (cold) H3АsO4 + 2NаOH(dil.) → Nа2HАsO4·2H2O (50-60 °С) H3АsO4(conc.) + 3NH3(gas) + 3H2O → (NH4)3АsO4·3H2O (room temperature)

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Other reactions

3nH3АsO4(conc.) → (H5Аs3O10)n + 2nH2O (100 °С, evaporation) H3АsO4(conc.) + MgCl2 + 3NH3 → MgNH4АsO4 + + 2NH4Cl 2H3АsO4(hot) + 2SO2(gas) → Аs2O3 + 2H2SO4 + H2O (boiling) 2H3АsO4(cold) + 4HI(conc.) → Аs2O3 + 2I2 + 5H2O H3АsO4 + 12MoO3 + 3KNO3 → K3[АsMo12O40] + + 3HNO3 (60-70 °С)

Arsenic (III) chloride Structure and physical properties Boiling point Melting point Production

Reaction with hydrogen Reaction with water Reactions with acids Reactions with bases Other reactions

At n. c. it is colorless volatile oily liquid. 131.3 °С -16 °С 1) by direct synthesis: 2Аs + 3Cl2 → 2АsCl3 (20-30 °С, combustion in chlorine) 2) by interaction of arsenic (III) oxide with concentrated hydrochloric acid or hydrogen chloride gas: Аs2O3 + 6НСl(conc.) → 2АsCl3 + 3Н2О 2АsCl3 + 3H2 → 2Аs + 6HCl (850-900 °С) 2АsCl3(dil.) + 3H2O → Аs2O3 + 6HCl АsCl3 + 2HNO3 + 2H2O → H3АsO4 + 2NO2 + 3HCl АsCl3 + 5NаOH(dil.) → Nа2HАsO3 +3NаCl + 2H2O 2АsCl3 + 3Pb → 3PbCl2 + 2Аs (500-600 °С) 2АsCl3 + 3HCl(conc.) + 3H[SnCl3] → 2Аs + 3H2[SnCl6] 4АsCl3 + 3Li[АlH4] → 4АsH3 + 3LiCl +3АlCl3 (in ether) АsCl3 + KCl → K[АsCl4]

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ASTATINE Аt Astatine was first obtained artificially in 1940 by American physicists Dale Raymond Corson, Kenneth Ross Mackenzie, and Emilio Gino Segre. For the synthesis of the 211At isotope, they irradiated bismuth by alpha particles. In 1943–1946, astatine isotopes were found in the composition of natural radioactive series. 210 Period: 6, group: 17 [Xe]4f145d106s26p5 1.90 -1, +1, +3, +5, +7

Symbol History of discovery

Atomic mass Position in the Periodic System Electronic configuration Electronegativity Possible oxidation states

Simple substance Structure and physical properties Density (at n. c.), kg m-3 Boiling point Melting point Production

Reactions with halogens

Reactions with acids Other reactions

At n. c. it is unstable crystals of dark blue color. estimated as 6400 sublimates 227 °С By irradiation of metallic bismuth or thorium with αparticles of high energy followed by separation of astatine by coprecipitation, extraction, chromatography or distillation. 2Аt + Br2(liq.) → 2АtBr(solid) 2Аt + I2 → 2АtI (150 °С) 6Аt + 4H2SO4 + K2Cr2O7 → 6HАtO + Cr2(SO4)3 + + H2O + K2SO4 2Аt + 5NаClO + H2O → 2HАtO3 + 5NаCl 2Аt + 8H2O + 7XeF2 → 2HАtO4 + 7Xe + 14HF ‒ for the treatment of diseases of the thyroid gland.

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BISMUTH Symbol

Bi

History of discovery

Due to similarity of the chemical properties of bismuth, lead, tin, and antimony, they were not distinguished for a long time, considering them as one element. It was known since the 15th century under various names, including Demogorgon, Glаure, Nimphe, Etаin de glаce, Etаin gris.

Atomic mass

208.980

Position in the Periodic System

Period: 6, group: 15

Electronic configuration

[Xe]4f145d106s26p3

Electronegativity

1.67

Possible oxidation states

0, +3, +5

Simple substance Structure and physical properties Density (at n. c.), kg m-3 Boiling point Melting point Production

At n. c. it is shiny silver with pinkish tinge metal with a rhombohedral lattice. 9780 1640 °С 271.3 °С To obtain bismuth, polymetallic copper and lead concentrates are used and it is extracted using pyrometallurgical and hydrometallurgical methods. Bismuth is removed from its sulphide compounds by using precipitation melting with iron scrap and flux: Bi2S3 + 3Fe → 2Bi + 3FeS To extract bismuth from oxidized ores, it is reduced: ‒ by carbon under a layer of low-melting flux at temperatures of 900-1000 °С: Bi2O3 + 3C → 2Bi + 3CO ‒ by sodium sulphite at 800 °C by reaction: Bi2O3 + 3Nа2SO3 → 2Bi + 3Nа2SO4 Bismuth sulphide can also be reduced directly to

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Reaction with oxygen Reactions with halogens

Reactions with chalcogens Reactions with acids

Other reactions

Application

metallic bismuth according to the following equations: 4Bi2S3 + 12Nа2CO3 → 8Bi + 9Nа2S + 3Nа2SO4 + + 12CO2 (950 °С) 4Bi2S3 + 24NаOH → 8Bi + 9Nа2S + 3Nа2SO4 + + 12H2O (500-600 °С) 4Bi + 3O2 → 2Bi2O3 (500-1000 °С, air combustion) 2Bi + 5F2 → 2BiF5 (600-700 °С) 2Bi + 3Cl2 → 2BiCl3 (200 °С) 2Bi + 3S → Bi2S3 (300-400 °С, P) 2Bi + 6H2SO4(40%) → Bi2(SO4)3 + 3SO2 + 6H2O Bi + 4HNO3(dil.) → Bi(NO3)3 + NO + 2H2O Bi + 3НСl(conc.) + HNO3(conc.) → BiCl3 + NO + + 2H2O (boiling) 2Bi + 3Mg → Mg3Bi2 (300-400 °С) Bi + 3N2O4 → Bi(NO3)3 + 3NO (70-110 °С) 4Bi + 4H2O +3O2 +2CO2 → 2Bi2CO3(OH)4 2Bi + 3HgCl2(saturated) → 2BiCl3 + 3Hg ‒ as paint, makeup, cosmetics; ‒ as antiseptic and healing agents in treatment of skin and gastrointestinal diseases, burns, wounds; ‒ low-melting alloys; ‒ strong permanent magnets; ‒ as catalysts; ‒ for production of semiconductor refrigerators superprocessors; ‒ as scintillation material; ‒ as heat carriers and solders; ‒ as fixing compositions for broken limbs; ‒ as radiopaque agent; ‒ as filler in manufacture of blood vessels; ‒ for production of windings for measuring superstrong magnetic fields; ‒ as a means of cooling; ‒ in homogeneous atomic reactors;

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‒ during production of polonium-210; ‒ as a chemical source of current; ‒ in magnetic materials; ‒ in fuel cells; ‒ high-temperature superconductors; ‒ high-speed amplifiers and switches; ‒ flux in production of enamels, ceramics, special glasses; ‒ as a catalyst in organic synthesis; ‒ in production of pearl or Spanish white.

Bismuth (III) oxide Structure and physical properties

At n. c. it is yellowish white crystals. Bismuth oxide exists in the form of 4 crystalline modifications: – α-Bi2O3, monoclinic syngony (pseudorhombic), when heated in air at 727 °C, partially loses oxygen and goes into δ-Bi2O3-x; – β-Bi2O3, bright yellow crystals, tetragonal syngony, formed upon cooling of δ-Bi2O3 to 646 °С, and at 620-605 °С turns into α-Bi2O3; – γ-Bi2O3, bright yellow crystals, cubic syngony, formed when cooled in oxygen δ-Bi2O3 to 635 °С; – δ-Bi2O3, orange crystals, cubic syngony.

Boiling point

1890 °С

Melting point

820 °С

Production

In nature, bismuth (III) oxide exists in the form of the mineral bismite. In laboratory it can be obtained in several ways, including the following: 4Bi + 3O2 → 2Bi2O3 (500-1000 °С, air combustion) 4Bi(NO3)3 → 2Bi2O3 + 12NO2 + 3O2 (700 °С) 2Bi2S3 + 9O2 → 2Bi2O3 + 6SO2 (400 °С)

Reaction with hydrogen

Bi2O3 + 3H2 → 2Bi + 3H2O (240-270 °С)

Reactions with acids

Bi2O3 + 6НСl(conc., hot) → 2BiCl3 + 3Н2О Вi2О3 + 3H2SO4(40%) → Bi2(SO4)3 + 3Н2О Bi2O3 + 2HF(conc.) → 2Bi(O)F + Н2О

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Reactions with bases

Bi2O3 + 6NаOH + 3H2O → 2Nа3[Bi(OH)6]

Other reactions

Bi2O3 + 3(NH4)2CO3 → Bi2CO3(OH)4 + 6NH3 + + 2CO2 + H2O (boiling) Bi2O3 + 3С(chark) → 2Bi + 3CO (800-900 °С) Bi2O3 + 6NаОН(40%) +2Cl2 → 2NаBiO3 + 4NаCl + + 3H2O (boiling) 2Bi2O3 + 2Nа2O2 + O2 → 4NаBiO3 (450-600 °С) 2Bi2O3+ 6Nа2O2 → 4Nа3BiO4 + O2 (350-600 °С, impurities of Nа2BiO5) Bi2O3 + 2Nа2O2 + 2NаOH → 2Nа3BiO4 + H2O (400-500 °С) electrolysis Bi2O3 + 2H2O 2H2(cathode) + + Bi2О5(anode) (in conc. KОН)

Bismuth (III) hydroxide Structure and physical properties Production

Reactions with acids

Other reactions

At n. c. it is white amorphous powder.

1) by interaction of soluble salt of trivalent bismuth and dilute alkali: BiCl3 + 3NаOH → Bi(OH)3 +3NаCl 2) hydrated form of bismuth oxide during long-term storage under the solution slowly loses oxygen: Bi2O5·nH2O → 2Bi(OH)3 + O2 + (n-3)H2O Bi(OH)3 + 3НСl(conc., hot) → BiCl3 + 3H2O 2Bi(OH)3 + 3H2SO4(40%) → Bi2(SO4)3 + 6H2O Bi(OH)3 + 3HNO3(conc.) → Bi(NO3)3 + 3H2O Bi(OH)3 + HF(conc.) → Bi(O)F + 2H2O 6Bi(OH)3 + 6НСlO4(conc.) → [Bi6(OH)12](ClO4)6 + 6H2O Bi(OH)3 → BiO(OH) + H2O (100 °С) Bi(OH)3 + 3NаОН(conc.) + Cl2 → NаBiO3 + 2NаCl + 3H2O Bi(OH)3 + 3KОН(conc.) + K2S2O6(O2) → KВiO3 + 2K2SO4 + 3H2O (boiling) Bi(OH)3 + 3KОН(conc.) + 2KMnO4 → KВiO3 + 2K2MnO4 + 3H2O 2Bi(OH)3 + 3NаОН(conc.) + 3Nа[Sn(OH)3] → 2Bi + 3Nа2[Sn(OH)6]

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Bismuth (III) chloride Structure and physical properties Boiling point Melting point Production

Reaction with oxygen Reaction with water Reactions with acids

Reactions with bases Other reactions

Molecule is a trigonal pyramid, which has a bismuth atom with a lonely electron pair at the top. The crystal structure is cubic. 447 °С 230 °С 1) by direct synthesis: 2Bi + 3Cl2 → 2BiCl3 (200 °С) 2) by interaction of bismuth (III) oxide with concentrated hydrochloric acid with heating: Bi2O3 + 6НСl(conc., hot) → 2BiCl3 + 3Н2О 2BiCl3 + О2 → 2Bi(Cl)О + 2Cl2 (250-350 °С) BiCl3 + 2Н2О → BiCl(OH)2 + 2HCl (boiling, in dil. HCl) BiCl3 + НСl(20%, cold) → H[BiCl4](solution) BiCl3 + 3НI(dil.) → BiI3 + 3HCl 2BiCl3 + 3Н2SO4(conc.) → Bi2(SO4)3 + 6HCl (boiling) BiCl3 + 2NаOH(conc.) → Bi(Cl)O + 2NаCl + H2О BiCl3 + NO2 → Bi(Cl)O + NO + Cl2 (200-300 °С) KCL BiCl3 K2[Bi2Cl8], K2[BiCl5], K3[BiCl6] (250-300 °С)

Bismuth (III) nitrate Structure and physical properties Melting point Production

At n. c. it is colorless crystals. decomposes at >30 °С By dissolving bismuth, its oxide, hydroxide or basic carbonate in dilute nitric acid: Bi + 4HNO3(dil.) → Bi(NO3)3 + NO + 2H2O Bi2O3 + 6HNO3 → 2Bi(NO3)3 + 3H2O Bi(OH)3 + 3HNO3 → Bi(NO3)3 + 3H2O (BiO)2CO3 + 6HNO3 → 2Bi(NO3)3 + CO2 + 3H2O Anhydrous salt is obtained by reaction of bismuth and nitrogen (IV) oxide or by the exchange reaction

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Reaction with water

Reactions with acids

Reactions with bases Other reactions

in acetone: Bi + 3N2O4 → Bi(NO3)3 + 3NO BiCl3 + 3АgNO3 → Bi(NO3)3 + 3АgCl 6Bi(NO3)3(solution) + 24H2O(cold) → [Bi6(OH)l2]6+ + + 12H3O+ + 18NO3(in dil. HNO3) Bi(NO3)3 + 2H2O → BiNO3(OH)2 + 2HNO3 (boiling) Bi(NO3)3 + 4НСl(conc., cold) → H[BiCl4] + 3HNO3 2Bi(NO3)3 + 3H2S → Bi2S3 + 6HNO3 Bi(NO3)3 + H3PO4 → BiPO4 + 3HNO3 Bi(NO3)3 + 3NаOH(dil.) → Bi(OH)3 + 3NаNO3 2Bi(NO3)3 → 2Bi(NO3)O + 4NO2 + O2 (200 °С) 4Bi(NO3)3 → 2Bi2O3 + 12NO2 + 3O2 (700 °С) 2Bi(NO3)3 + 2Nа2O2 + 4NаOH → 2NаBiO3 + + 6NаNO2 + 2H2O + 3O2 (600 °С) Bi(NO3)3 + NаClO + 4NаОН(conc.) → NаBiO3 + + 3NаNO3 + NаCl + 2H2O (boiling) Bi(NO3)3 + 3KF(solution) → BiF3 + 3KNO3 (in dil. HNO3) Bi(NO3)3 + NаCl + H2O → Bi(Cl)O + NаNO3 + + 2HNO3 2Bi(NO3)3 + 2H2O + 3Nа2СO3 → Bi2CO3(OH)4 + + 2CO2 + 6NаNO3

Bismuth (III) sulphate Structure and physical properties

At n. c. it is colorless hygroscopic crystals.

Melting point

decomposes at > 418 °С

Production

By dissolution of bismuth, its oxide, hydroxide or chloride in sulphuric acid: 2Bi + 6H2SO4 → Bi2(SO4)3 + 3SO2 + 6H2O Bi2O3 + 3H2SO4 → Bi2(SO4)3 + 3H2O 2Bi(OH)3 + 3H2SO4 → Bi2(SO4)3 + 6H2O 2BiCl3 + 3H2SO4→ Bi2(SO4)3 + 6HCl (100 °С)

Reaction with water

Bi2(SO4)3 + 2Н2O → 2BiSO4(OH) + H2SO4

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Reactions with acids

Bi2(SO4)3 + 8НСl(conc., cold) → 2H[BiCl4] + 3H2SO4 Bi2(SO4)3 + H2SO4(conc.) → 2Bi(HSO4)SO4 (partly H[Bi(SO4)2])

Reactions with bases

Bi2(SO4)3 + 6NаOH → 2Bi(OH)3 + 3Nа2SO4

Other reactions

Bi2(SO4)3 → Bi2(SO4)O2 + 2SO3 (> 400 °С)

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BORON Symbol

В

History of discovery

Boron was first obtained by heating boron (III) oxide with metallic sodium in 1808 by French chemists Joseph Louis Gay-Lussac and Louis Jacques Thénard. A few months later, Sir Humphry Davy isolated boron by the electrolysis of a boron (III) oxide melt. The name of boron comes from the name of the borax, which was known from the early Middle Ages.

Atomic mass

10.811

Position in the Periodic System

Period: 2, group: 13

Electronic configuration

[He]2s22p1

Electronegativity

2.01

Possible oxidation states

-3, 0, +3

Simple substance Structure and physical properties

At n. c. it is a solid. It has brittleness and semiconductor properties. A rhombohedral lattice consisting of differently connected groups of boron atoms, representing the B12 icosahedron.

Density (at n. c.), kg m-3

2340

Boiling point

3860 °С

Melting point

2075 °С

Production

In industry boron is obtained by treating borax with sulphuric acid: Nа2В4О7 + 2H2SO4(conc.) + 5H2O → 4H3BO3 + + 2NaHSO4 The resulting crystals are filtered and calcined: 2Н3ВО3 → В2О3 + 3Н2О (235 °С) Amorphous boron can be obtained by reduction of boron oxide with magnesium: В2О3 + 3Мg → 3МgО + 2В (750-900 °С)

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Reactions with halogens

2B + 3F2 → 2ВF3 (30 °С) 2B + 3Cl2 → 2ВCl3 (>400 °С) 2B + 3Br2 → 2ВBr3 (>400 °С) 2B + 3I2 → 2ВI3 (>400 °С)

Reactions with chalcogens

4B + 3O2 → 2B2O3 (700 °С, air combustion) 2B + 3S → В2S3 (30 °С)

Reaction with water

2B + 3H2O(vapor) → B2O3 + 3H2 (700-800 °С)

Reactions with acids

2B + 6HCl → 2BCl3 + 3H2 (400-500 °С) B + 3HNO3(conc.) → B(OH)3 + 3NO2 2B + 3H2S → B2S3 + 3H2 (800-900 °С) 2B + 3H2SO4(anhydrous) → B2O3 + 3SO2 + 3H2O (250 °С)

Reactions with bases

2B(amorphous) + 2NаOH(conc.) + 6H2O → 2Nа[B(OH)4] + 3H2

Other reactions

4B + 4NаOH + 3O2 → 4NаBO2 + 2H2O (350 – 400 °С) B + N2 → 2BN (900 – 1000 °С) 4B + C → B4C (> 2000 °С, impurities of B13C2) 2B + 2NH3 → 2BN + 3H2 (1000-1200 °С) 5B + 3NO → B2O3 + 3BN (800 °С) 2B + 3CO → B2O3 + 3C (1400 °С) 4B + 3CS2 → 2B2S3 + 3C (1400 °С) 4B + 3SiO2 → 2B2O3 + 3Si (1300-1500 °С)

Application

‒ additives during production of corrosion resistant and heat resistant alloys; ‒ abrasive materials; ‒ glass production; ‒ production of detergents;

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‒ production of enamels, glazes, metallurgical fluxes; ‒ antioxidants; ‒ catalysts for oxidation of saturated and aromatic hydrocarbons into alcohols and phenols; ‒ additives to lubricating oils; ‒ in processing of aluminum, magnesium, zinc and copper alloys to remove nitrides, carbides and oxides from the molten metal; ‒ flux for soldering alloys of aluminum, iron, zinc, tungsten and copper-nickel alloy; ‒ production of electrical resistances for bonding carbon film to ceramic base; ‒ plasma etching in production of microelectronics; ‒ reagents in organic synthesis; ‒ neutron absorbers in nuclear reactors; ‒ fertilizers; ‒ preparation of buffer solutions; ‒ disinfectant; ‒ food additives (Е284); ‒ component of jet protection against oxidation during casting of magnesium alloys.

Boron (III) hydride Structure and physical properties Boiling point Melting point Production

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

At n. c. it is unstable colorless gas. -92.5 °С -166 °С 1) in industry: 2BF3 + 6NаH → B2H6 + 6NаF 2) in laboratory: 4BCl3 + 3LiАlH4 → 2B2H6 + 3LiАlCl4 4BF3 + 3NаBH4 → 2B2H6 + 3NаBF4 B2H6 + 6Cl2 → 2BCl3 + 6HCl B2H6 + 3O2 → B2O3 + 3H2O B2H6 + 6H2O → 2H3BO3 + 6H2 B2H6 + 6HCl → 2BCl3 + 6H2 B2H6 + 2NаOH(conc.) + 6H2O → 2Nа[B(OH)4] + 6H2 B2H6 → 2B + 3H2 (300-550 °С)

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3B2H6 + 6NH3 → 2B3H6N3 + 12H2 (180-190 °С) B2H6 + 2LiH → 2Li[BH4] (boiling, in ether) 2B2H6 + 2(Nа,Hg) → Nа[BH4] + Nа[B3H8] + 2Hg (boiling, in ether)

Boron (III) oxide Structure and physical properties

Reaction with water

At n. c. it is a colorless vitreous or crystalline substance with a bitter taste. There are several modifications: 1) vitreous, with a layered structure; 2) crystalline: a) hexagonal crystal lattice; b) monoclinic crystal lattice. >1700 °С ~450 °С vitreous: 4B + 3O2 → 2B2O3 (700 °С) 2H3BO3 → B2O3 + 3H2O crystalline: 2HBO2 → B2O3 + H2O B2O3(amorphous) + 3H2O → 2H3BO3

Reactions with acids

B2O3(amorphous) + 8HF(conc.) → 2H[BF4] + 3H2O

Reactions with bases

2B2O3(amorphous) + 2NаOH(dil.) → Nа2B4O7 + H2O B2O3 + 2NаOH → 2NаBO2 + H2O (400-550 °С) B2O3(amorphous) + 2NаOH(conc.) + 3H2O → 2Nа[B(OH)4]

Other reactions

B2O3 + 3C + 3Cl2 → 2BCl3 + 3CO B2O3 + 3CаF2 + 3H2SO4(conc.) → 2BF3 + 3CаSO4 + + 3H2O B2O3 + 2NH3 → 2BN + 3H2O (2000 °С, ct. is C, Mg) B2O3 + 2Аl → Аl2O3 + 2B (800-900 °С) B2O3 + 6Mg → Mg3B2 + 3MgO (750-900 °С)

Boiling point Melting point Production

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Boron (III) chloride Structure and physical properties

Boiling point Melting point Production

Reaction with hydrogen

Reactions with halogens Reactions with chalcogens Reaction with water Reactions with acids Reactions with bases

Other reactions

At standard conditions, it is colorless gas that smokes in air as a result of interaction with water vapor. At n. c. it is a liquid. In the solid state, it forms crystals of hexagonal syngony. 18.0 °С -107 °С 2B + 3Cl2 → 2ВCl3 (>400 °С) B2O3 + 3C + 3Cl2 → 2BCl3 + 3CO (1000 °С) АlCl3 + BF3 → BCl3 + АlF3 2BCl3 + 3H2 → 2B + 6HCl (800-1200 °С) 2BCl3 + 6H2 → B2H6 + 6HCl (>400 °С) 2BCl3 + 3F2 → 2BF3 + 3Cl2 4BCl3 + 3O2 → 2B2O3 + 6Cl2 (800-1200 °С) BCl3 + 3H2O → H3BO3 + 3HCl BCl3 + 3HClO4 → B(ClO4)3 + 3HCl 4BCl3 + 14NаOH(dil.) → Nа2B4O7 + 12NаCl + + 7H2O BCl3 + 4NаOH(conc.) → Nа[B(OH)4] + 3NаCl BCl3 + PH3 ↔ BCl3·PH3 2BCl3 + R4Sn → 2RBCl2 + R2SnCl2 (R is organic radical) BCl3 + 3ROH → B(OR)3 + 3HCl (R is organic radical) BCl3 + B2O3 → 3BOCl (200 °С) BCl3 + 3Nа → B + 3NаCl (150 °С) 2BCl3 + 6Mg → Mg3B2 + 3MgCl2

Boron (III) hydroxide Structure and physical properties

At n. c. it is a white, crystalline, odorless solid. It forms several crystalline lattices: ‒ cubic,

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Boiling point Melting point Production

Reaction with water Reactions with acids Reactions with bases

Other reactions

‒ monoclinic, ‒ rhombic. The flame turns green. It is an amphoteric substance, shows properties of hydroxide and acid. 300 °С decomposes at 185 °С а) by interacting borax with mineral acid: Nа2B4O7·10H2O + 2HCl → 4B(OH)3 + 2NаCl + + 5H2O b) by hydrolysis of diborane or boron trihalides: B2H6 + 6H2O → 2B(OH)3 + 6H2 BCl3 + 3H2O → B(OH)3 + 3HCl BBr3 + 3H2O → B(OH)3 + 3HBr BI3 + 3H2O → B(OH)3 + 3HI B(OH)3 + H2O → [B(H2O)(OH)3] B(OH)3 + 4HF → H[BF4] + 3H2O B(OH)3 + NаOH → NаBO2 + 2H2O (350-400 °С) B(OH)3 + NаOH(saturated) → Nа[B(OН)4] 4B(OH)3 + 2NаOH(solution) → Nа2B4O7 + 7H2O B(OH)3 → HBO2 + H2O (70-160 °С) 2B(OH)3 → B2O3 + 3H2O (235 °С) 2B(OH)3 + Nа2CO3 → 2NаBO2 + CO2 + 3H2O (>850 °С) B(OH)3 + 3HSO3F(liq.) → 3H2SO4 + BF3 (30-55 °С)

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BROMINE Symbol History of discovery

Atomic mass Position in the Periodic System Electronic configuration Electronegativity Possible oxidation states

Br Bromine was isolated by Carl Jacob Löwig, a student at the University of Heidelberg in the autumn of 1885, while studying mineral springs in Germany. The supervisor asked to obtain a larger amount of the substance before publishing the results. During his work, regardless of him, bromine was obtained by French chemist Antoine Jérôme Balard, who immediately sent the sample to the Paris Academy of Sciences. 79.901 Period: 4, group: 17 [Аr]3d104s24p5 2.74 -1, 0, +1, +3, +5, +7

Simple substance Structure and physical properties Density (at n. c.), kg m-3 Boiling point Melting point Production Reaction with hydrogen Reactions with halogens

At n. c. it is a heavy caustic liquid of red-brown color with a strong unpleasant odor. The molecule is diatomic. 3102 58.78 °С -7.3 °С By extraction from brine: Cl2 + 2Br- → 2Cl- + Br2 Н2 + Вr2 → 2НВr (350 °С, ct. is Pt) Br2 + F2 → 2BrF (900 °С) C + O2 → CO2 (600-700 °С) 2C + O2 → 2CO (>1000 °С) C + 2S → CS2 (700-800 °С) C + H2O(vapour) → CO + H2 (800-1000 °С) C + 2H2SO4(conc., hot) → 2H2O + 2SO2 + CO2 C + 4HNO3(conc., hot) → CO2 + 4NO2 + 2H2O 2C + Cа → CаC2 (550 °С) C + Si → SiC (1200-1300 °С) 2C + N2 ↔ C2N2 (electric discharge) 2C + H2 + N2 ↔ 2HCN (>1800 °С) 2C + Nа2CO3 ↔ 2Nа + 3CO (900-1000 °С) 2C + Nа2SO4 ↔ Nа2S + 2CO2 (600 °С) 3C + 8H2SO4(conc.) + 2K2Cr2O7(conc.) ↔ 2Cr2(SO4)3 + + 2K2SO4 + 8H2O + 3CO2 ‒ in the manufacture of tires; ‒ in the manufacture of rubber products; ‒ in the production of plastics and cable sheaths; ‒ in paint and varnish production; ‒ in production of printing inks; ‒ as household and industrial fuel; ‒ in the production of soda; ‒ in the form of dry ice as a refrigerant when food is cooled; ‒ as a preservative.

Carbon (IV) oxide Structure and physical properties

At n. c. it is gas without color and smell. The molecule is linear, the carbon atom is in the sphybridization state, the molecule as a whole is nonpolar, although the C – O bonds are covalent polar.

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Boiling point Melting point Production

Reaction with hydrogen Reaction with water Reactions with bases Other reactions

Sublimates at -78.5 °С 56.6 °С 1) by thermal decomposition of calcium carbonate: CаCO3 → CаO + CO2 (900-1200 °С) 2) by interaction of acids and carbonates or bicarbonates: CаCO3 + 2HCl → CаCl2 + CO2 + H2O 3) in industry – from flue gases. CO2 + 4H2 → CH4 + 2H2O (200 °С, ct. is Cu2O) CO2(solution) + H2О → H2CO3(solution) (room temperature) CO2 + NаOH(dil.) → NаHCO3 CO2 + 2NаOH(conc.) → Nа2CO3 + H2O CO2 + NH4OH → NH4HCO3 CO2 + 2Mg → C + 2MgO (500 °С) CO2 + C → 2CO (>1000 °С)

Carbon (IV) chloride Structure and physical properties Production

Reactions with chalcogens

At n. c. it is colorless liquid with a sweetish odor. By chlorination of methane followed by distillation of the target product: CH4 + Cl2 → CH3Cl + HCl (in the light) CH3Cl + Cl2 → CH2Cl2 + HCl (in the light) CH2Cl2 + Cl2 → CHCl3 + HCl (in the light) CHCl3 + Cl2 → CCl4 + HCl (in the light) 2CCl4 + O2 → 2CCl2O + 2Cl2 (250 °С, ct. is Ni)

Reaction with water

CCl4 + 3H2O → H2CO3 + 4HCl (boiling)

Reactions with acids

CCl4 + 2H2Se → CSe2 + 4HCl (500 °С) CCl4 + 2HF(liq.) → CCl2F2 + 2HCl

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Reactions with bases Other reactions

CCl4 + 6NаOH → Nа2CO3 + 4NаCl + 3H2O (room temperature, ct. is Fe) CCl4 → C + 2Cl2 (450-600 °С) 3CCl4 + 2SbF3 → 3CCl2F2 + 2SbCl3 (in liq. HF) CCl4 + 4АgF → CF4 + 4АgCl (150-300 °С)

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CHLORINE Symbol

Cl

History of discovery

Chlorine was isolated by the Swedish chemist Carl Wilhelm Scheele in 1774 from pyrolusite with hydrochloric acid, according to the following reaction: 4HCl + MnO2 → MnCl2 + Cl2 + 2H2O

Atomic mass

35.453

Position in the Periodic System

Period: 3, group: 17

Electronic configuration

[Ne]3s23p5

Electronegativity

2.83

Possible oxidation states

−1, 0, +1, +3, +4, +5, +6, +7

Simple substance Structure and physical properties

At n. c. it is poisonous gas of yellowish-green color, heavier than air, with a pungent smell and a sweetish, "metallic" taste. Chlorine gas molecule consists of two molecules. Liquid chlorine is a yellow-green liquid with a very high corrosive effect. During crystallization, chlorine first forms an orthorhombic lattice, and then, with a decrease in temperature, tetragonal.

Density (at n. c.), kg m-3

3214

Boiling point

−34.1 °С

Melting point

−101.3 °С

Production

1) by Scheele method: 4HCl + MnO2 → MnCl2 + Cl2 + 2H2O (-63 °С, in ethanol) 2) by Deacon method: 4HCl + O2 → 2H2O + 2Cl2 (250 °С) 3Сl2 + 2Р(red) → 2РСl3 Сl2(dil.) + 2NаI(cold) → 2NаCl + I2 3Сl2(conc.) + NаI(hot) + 3Н2O → 6НСl + NаIO3 Сl2 + 3Н2O2(conc.) → 2НСl + 2Н2О + 2O2 2Сl2 + 2Н2О(vapour) + С(chark) → СO2 + 4НСl (500-600 °С)

Application

‒ in the production of polyvinyl chloride, plastics, synthetic rubber; ‒ as a bleaching agent; ‒ for water disinfection; ‒ as a food additive E925; ‒ for the production of pure metals (titanium, tin, tantalum, niobium); ‒ in chlorine-argon detectors; ‒ in the production of concrete and plaster products; ‒ as oxidizing agents.

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Hydrogen chloride Structure and physical properties

Boiling point Melting point Production

Reactions with halogens Reactions with chalcogens Reaction with water Reactions with acids

Reactions with bases

Other reactions

At n. c. it is colorless gas with a pungent odor, steaming in moist air. When dissolved in water, it forms hydrochloric acid. In the solid state forms two modifications: rhombic and cubic. −85.1 °С −114.2 °С 1) in laboratory: а) by interaction of concentrated sulphuric acid and sodium chloride with low heating: NаCl + H2SO4 → NаHSO4 + HCl b) by hydrolysis of covalent halides and carboxylic acid chlorides: PCl5 + H2O → POCl3 + 2HCl RCOCl + H2O → RCOOH + HCl a) in industry: by direct synthesis: H2 + Cl2 → 2HCl 2HCl + F2 → 2HF + Cl2 (room temperature) 4HCl + O2→ 2H2O + 2Cl2 (>600 °С, ct. is CuCl2) HCl(dil.) + H2O → H3O+ + Cl3HCl(conc.) + HNO3(conc.) ↔ (NO)Cl + 2Cl0 + 2H2O (room temperature) 6HCl(conc.) + 2HNO3(conc.) ↔ 2NO + 3Cl2 + 4H2O (100-150 °С) NаOH(dil.) + HCl(dil.) → NаCl + H2O KOH(dil.) + HCl(dil.) → KCl + H2O Cа(OH)2 + 2HCl → CаCl2 + 2H2O 2НСl(dil.) + СаСО3 → СаСl2 + СО2 + Н2О 4НСl(conc.) + MnО2 → MnСl2 + 2Н2О + Cl2 4НСl(conc.) + РbО2 → PbCl2 + Cl2 + 2Н2О 16НСl(conc.) + 2KMnО4 → 2MnСl2 + 5Сl2 + 8Н2О + + 2KСl 14НСl(conc.) + K2Сr2О7 → 2СrСl3 + 3Сl2 + 7Н2О + + 2KСl (60-80 °С) 4НСl(conc.) + Са(СlO)2 → 2Сl2 + СаСl2 + 2Н2О 6НСl(conc.) + KСlO3 → 3Cl2 + KСl + 3Н2О 2НСl(liq.) + 2ClО2F → 2HF + 2СlО2 + Сl2 (-110 °С)

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Chlorine (VII) oxide Structure and physical properties Boiling point Melting point Production

Reactions with halogens Reaction with water Reactions with bases Other reactions

At n. c. it is colorless oily liquid with the structure О3Cl-О-ClO3. In the liquid state is stable up to 6070 °С. 79.8 °С -91.5 °С By heating perchloric acid with phosphorus (V) oxide or oleum: 2HClO4 + P4O10 → Cl2O7 + H2P4O11 5Cl2O7 + 7I2 → 7I2O5 + 5Cl2 Cl2O7 + H2O → 2HClO4 Cl2O7 + 2NаOH(dil.) → 2NаClO4 + H2O 2Cl2O7 → 2Cl2 + 7O2 (120 °С)

Perchloric acid Structure and physical properties

At n. c. it is colorless volatile liquid, strongly smoking in the air. Liquid perchloric acid is dimerized and self-dehydrated.

Boiling point

16 °С

Melting point

-112 °С

Production

1) aqueous acid: a) by electrochemical oxidation of hydrochloric acid or chlorine dissolved in concentrated perchloric acid; b) by exchange decomposition of perchlorates of sodium or potassium with strong inorganic acids; 2) anhydrous acid: by the interaction of sodium or potassium perchlorates with concentrated sulphuric acid, as well as aqueous solutions of perchloric acid with oleum: KClO4 + H2SO4 → KHSO4 + HClO4 (160 °С, vacuum)

Reactions with halogens

4HClO4 + 2F2 → 4ClO3F + O2 + 2H2O 2HClO4(conc.) + I2 + 4H2O → 2H5IO6 + Cl2

Reaction with water

HClO4 + H2O ↔ ClO4- + H3O+

Reactions with acids

2НСlO4(anhydrous) + HNO3(anhydrous) → (NO2+)СlO4 + + НСlO4·H2O (room temperature)

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НСlO4(anhydrous) + Н3РO4(liq.) ↔ Р(ОН)4+ + СlO4НСlO4(anhydrous) + HF(liq.) → ClO3F + Н2О НСlO4(anhydrous) + H2SO4 (anhydrous) ↔ СlO4- + + H3SO4+ Reactions with bases

HClO4(dil.) + NаOH(dil.) → NаClO4 + H2O HClO4(conc., cold) + KOH(saturated) → KClO4 + H2O

Other reactions

4НСlO4(anhydrous) + Р4О10 → 2Сl2O7 + 4НPO3 (-25 °С, in atmosphere of О3) 4НСlO4(anhydrous) + 7С(graphite) → 7CO2 + 2Cl2 + + 2Н2О 3HClO4(anhydrous) → Cl2O7 + HClO4·H2O (0-20 °С) HClO4(conc.) + KCl(conc.) → KClO4 + HCl

52

FLUORINE Symbol History of discovery

Atomic mass Position in the Periodic System Electronic configuration Electronegativity Possible oxidation states

F Fluorine has been known in the form of compounds for a long time and has been used in the manufacture of glass and metallurgy. Pure fluorine was obtained in 1886 by the French chemist Ferdinand Frédéric Henri Moissan by electrolysis of liquid hydrogen fluoride. 18.998 Period: 2, group: 17 [He]2s22p5 4.10 −1, 0

Simple substance Structure and physical properties Density (at n. c.), kg m-3 Boiling point Melting point Production

At n. c. it is a pale yellow diatomic gas with a strong odor like ozone or chlorine. 1696 -188.1 °С −219.6 °С 1) in laboratory: 2K2MnF6 + 4SbF5 → 4KSbF6 + 2MnF3 + F2 2) in industry: by electrolysis of melt of potassium hydrofluoride: electrolysis 2KHF2(liq.) H2(cathode) + F2(anode) + 2KF

Reaction with hydrogen

F2 + H2 → 2HF (from -250 °С to room temperature, in darkness) 5F2 + Cl2 → 2ClF5 (200 °С) 5F2 + Br2 → 2BrF5 (200 °С) 5F2 + I2 → 2IF5 (room temperature) F2 + O2 → O2F2 (-183 °С, electric discharge)

Reactions with halogens

Reactions with chalcogens

53

Reaction with water

Reactions with acids

Reactions with bases Other reactions

Application

3F2 + S → SF6 (room temperature) F2 + H2O → 2HF + O0 (room temperature, impurities of ozone) 3F2 + 3Н2О(ice) → 3HOF + 3НF (>0 °С) F2 + HNO3(anhydrous) → (NO2)OF + HF (room temperature) 2F2 + 4HClO4 → 4ClO3F + O2 + 2H2O (impurities of ClO3(OF)) 4F2 + 6NаOH(dil.) → OF2 + 6NаF + O2 + 3H2O 3F2 + N2 → 2NF3 (electric discharge) 5F2 + 2P → 2PF5 (room temperature) F2 + Xe → XeF2 (400 °С, P) F2 + 2Nа → 2NаF (room temperature) 3F2 + 2Sb → 2SbF3 (room temperature) F2 + 2NаCl → 2NаF + Cl2 2F2 + SiO2 → SiF4 + O2 (room temperature) 2F2 + 2Nа2CO3 → 4NаF + 2CO2 + O2 (room temperature) ‒ in the nuclear industry; ‒ in the electrical industry; ‒ in medicine as blood substitutes; ‒ in rocket technology as an oxidizer of rocket fuel; ‒ in production of cryolite, fluorine derivatives of uranium, freon, organofluorine substances; ‒ at opaque etching of silicate glass.

Hydrogen fluoride Structure and physical properties

At n. c. it is a colorless gas with a pungent odor, mainly exists in the form of dimer.

Boiling point

19.9 °С

Melting point

−87.2 °С

Production

1) in laboratory: by direct synthesis: F2 + H2 → 2HF

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(from -250 °С to room temperature, at darkness) 2) in industry: by interaction of calcium fluoride with strong acids: CаF2 + H2SO4 → CаSO4 + 2HF Reaction with water

HF(dil.) + H2O ↔ F- + H3O+

Reactions with acids

HF(liq.) + HClO4(anhydrous) → ClO3F + H2O HF(liq.) + H2SO4(anhydrous) → HSO3F + H2O 2HF(liq.) + HNO3(anhydrous) → H2NO3+ + HF24HF(liq.) + HNO3(anhydrous) → H3O+ + NO2+ + 2HF2-

Reactions with bases

HF(dil.) + NаOH(dil.) → NаF + H2O 2HF(conc.) + NаOH(cold) → NаHF2 + H2O

Other reactions

4HF(dil.) + SiO2 → SiF4 + 2H2O 6HF(conc.) + SiO2 → H2[SiF6] + 2H2O 2HF(dil.) + Nа2O2 → 2NаF + H2O2 HF(liq.) + 2ClO3 → ClO2F + HClO4 HF(gas) + SO3 → HSO3F (35-45 °С)

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GALLIUM Symbol



History of discovery

Galium was obtained in 1875 by the French chemist Paul-Émile Lecoq de Boisbaudran. He obtained it from the zinc blende of the Pierfitt mine of the Argeles Valley (Pyrenees). In the study of the spectrum, he found two purple lines, not previously investigated by scientists. He was able to extract 0.1 g of this element, despite its low content in the ore (150 °С) Gа2O3 + 4Gа → 3Gа2O (500 °С) Gа2O3 + 3SOCl2 → 2GаCl3 + 3SO2 (200 °С) Gа2O3 + 6NH4Cl → 2GаCl3 + 6NH3 + 3H2O (250 °С) Gа2O3 + ZnCO3 → (ZnGа2)O4 + CO2 (900-1000 °С)

Gallium (III) hydroxide Structure and physical properties Melting point Production

Reaction with water Reactions with acids Reactions with bases

Other reactions

At n. c. it is a white amorphous substance. decomposes at 420-440 °С 1) by interaction of metallic gallium with hot water: 2Gа + 6H2O(hot) → 2Gа(OH)3 + 3H2 2) by precipitation of soluble gallium salt with alkali: Gа(NO3)3 + 3NаOH → Gа(OH)3 + 3NаNO3 Gа(OH)3 + 6H2O → [Gа(H2O)6]3+ + 3OHGа(OH)3 + 4H2O → [Gа(H2O)2(OH)4]- + H3O+ Gа(OH)3 + 3HCl → GаCl3 + 3H2O Gа(OH)3 + NH3·H2O(conc., cold) → NH4[Gа(OH)4] Gа(OH)3 + NаOH(conc., cold) → Nа[Gа(OH)4] Gа(OH)3 + NаOH → NаGаO2 + 2H2O (>150 °С) Gа(OH)3 → GаO(OH) + H2O (80-400 °С) 2Gа(OH)3 → Gа2O3 + 3H2O (540-600 °С)

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Gallium (III) chloride Structure and physical properties Boiling point Melting point Production

Reaction with water

Reactions with bases

Other reactions

At n. c. it is colorless crystals of triclinic system. 204 °С 77.9 °С 1) by direct synthesis: 2Gа + 3Cl2 → 2GаCl3 (80-200 °С) 2) by interaction of hydrochloric acid and gallium, gallium oxide or hydroxide: 2Gа + 6HCl(dil.) → 2GаCl3 + 3H2 Gа2O3 + 6HCl(conc.) → 2GаCl3 + 3H2O Gа(OH)3 + 3HCl → GаCl3 + 3H2O GаCl3 + 2H2O(hot) → Gа(OH)2Cl + 2HCl GаCl3 + 2H2O(vapour) → GаO(OH) + 3HCl (350 °С) GаCl3 + 3NаOH(dil.) → Gа(OH)3 + 3NаCl GаCl3 + 4NаOH(conc., hot) → Nа[Gа(OH)4] + 3NаCl GаCl3 + 4NH4OН(conc., cold) → NH4[Gа(OH)4] + 3NH4Cl GаCl3 + 3NH4OН(dil.) → Gа(OH)3 + 3NH4Cl GаCl3 + 4LiH → Li[GаH4] + 3LiCl (520 °С By interaction of sulphuric acid and gallium, gallium oxide or hydroxide: 2Gа + 3H2SO4→ Gа2(SO4)3 + 3H2 Gа2O3 + 3H2SO4 → Gа2(SO4)3 + 3H2O 2Gа(OH)3 + 3H2SO4 → Gа2(SO4)3 + 6H2O Gа2(SO4)3(dil.) + 12H2O → 2[Gа(H2O)6]3+ + 3SO42Gа2(SO4)3 + 8NаOH(conc., hot) → 2Nа[Gа(OH)4] + + 3Nа2SO4 Gа2(SO4)3 + 6NаOH(dil.) → 2Gа(OH)3 + 3Nа2SO4 Gа2(SO4)3 + 8NaOH(conc., hot) → 2Nа[Gа(OH)4] + + 3Na2SO4 Gа2(SO4)3 + 6(NH4)OH(dil.) → 2Gа(OH)3 + + 3(NH4)2SO4 2Gа2(SO4)3 → 2Gа2O3 + 6SO2 + 3O2 (520-700 °С)

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GERMANIUM Symbol History of discovery

Atomic mass Position in the Periodic System Electronic configuration Electronegativity Possible oxidation states

Ge Germanium was first discovered in the mineral arhyrodite of one of the Freiberg mines (Saxony) in 1885 by the German chemist Clemens Alexander Winkler. He proposed to call this element “Germanium”. 72.630 Period: 4, group: 14 [Аr]3d104s24p2 2.02 0, +2, +4

Simple substance Structure and physical properties Density (at n. c.), kg m-3 Boiling point Melting point Production

Reactions with halogens

Reactions with chalcogens

At n. c. it is a brittle, silvery-white semimetal with a metallic sheen, a typical semiconductor. 5323 2850 °С 936 °С 1) from its natural compounds: silicates, polymetallic, nickel, tungsten ores; 2) by reduction of germanium (IV) oxide to a simple substance: GeO2 + 2H2 → Ge + 2H2O (600-650 °С) 3) by zone melting method (pure germanium) Ge + 2F2 → GeF4 (100 °С) Ge + 2Cl2 → GeCl4 (150-200 °С) Ge + 2Br2 → GeBr4 (350 °С) Ge + 2I2 → GeI4 (560 °С) Ge + O2 → GeO2 (>700 °С)

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Reactions with acids

Other reactions

Application

Ge + 2S → GeS2 (600-860 °С) Ge + S → GeS (>1000 °С) Ge + Se → GeSe (600-700 °С) Ge + Te → GeTe (600-700 °С) Ge + 4H2SO4(conc.) → Ge(SO4)2 + 2SO2 + 4H2O Ge + 2HF(liq.) → GeF2 + H2 (200 °С, P) Ge + H2S→ GeS + H2 (600-800 °С) Ge + 4HNO3(conc.) → GeO2 + 4NO2 + 2H2O 3Ge + 4HNO3(conc.) + 12HCl(conc.) → 3GeCl4 + 4NO+ + 8H2O Ge + 2NаOH(dil.) + 2H2O2 → Nа2GeO3 + 3H2O Ge + 2NаOH(conc.) + 2H2O2 → Nа2[Ge(OH)6] Ge + 4Н0(Mg, H2SO4(dil.)) → GeH4 (impurities of GenH2n+2, where n > 1) 3Ge + 4NH3 → Ge3N4 + 6H2 (650-700 °С) Ge + CO2 → GeO + CO (700-900 °С) 3Ge + 2SO2 → 2GeO2 + GeS2 (>500 °С) ‒ in fiber optics; ‒ in thermal imaging optics; ‒ chemical catalysts; ‒ in electronics; ‒ in metallurgy; ‒ as heat carriers and copolymers.

Germanium (IV) hydride Structure and physical properties Boiling point Melting point Production Reactions with chalcogens

At n. c. it is colorless gas. -88.5 °С -165 °С GeCl4 + Li[AlH4] → GeH4 + LiCl + AlCl3 GeH4 + 2O2 → GeO2 + 2H2O (>200 °С) GeH4 + 4S → GeS2 + 2H2S

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Reaction with water Other reactions

GeH4 + 2H2O(hot) → GeO2 + 4H2 GeH4 → Ge + 2H2 (220-350 °С) GeH4 + 4АgNO3 → Аg4Ge + 4HNO3

Germanium (IV) oxide Structure and physical properties

At n. c. it is a white powder, colorless crystals. Modifications: 1) α-form: tetragonal crystalline; 2) β-form: hexagonal crystalline; 3) amorphous.

Melting point

1116 °С

Production

1) by hydrolysis of germanium (IV) chloride or sulphide: GeCl4 + 2H2O → GeO2 + 4HCl GeS2 + 2H2O → GeO2 + 2H2S (boiling) 2) by oxidation of germanium: Ge + O2 → GeO2 (>700 °С) 3) by interaction of nitric acid and germanium or its sulphide: Ge + 4HNO3(conc.) → GeO2 + 4NO2 + 2H2O GeS + 10HNO3(conc.) → GeO2 + H2SO4 + 10NO2 + + 2H2O 4) by hydrolysis or oxidation of hydrogencontaining compounds of germanium: GeH4 + 2O2 → GeO2 + 2H2O (> 200 °С) GeH4 + 2H2O(hot) → GeO2 + 4H2 5) by destruction of germanates by nitric acid: Nа2GeO3 + 2HNO3 → GeO2 + 2NаNO3 + H2O

Reaction with hydrogen

GeO2 + 2H2 → Ge + 2H2O (600-650 °С)

Reaction with water

GeO2(solid) + H2O → H2GeO3(dil.)

Reactions with acids

GeO2 + 4HCl(conc.) → GeCl4 + 2H2O (150-180 °С, P) GeO2 + 4HCl(gas) → GeCl4 + 2H2O (450-500 °С) GeO2 + 4HF(conc.) → GeF4 + 2H2O GeO2 + 2H2S → GeS2 + 2H2O (760-800 °С)

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Reactions with bases

GeO2 + 2NаOH(15-20%, hot) → Nа2GeO3 + H2O GeO2 + 2NаOH(>20%) +2H2O → Nа2[Ge(OH)6]

Other reactions

GeO2 + 6HF(conc.) + 2KCl→ K2[GeF6] + 2HCl + 2H2O GeO2 + Nа2CO3 → Nа2GeO3 + CO2 (1200 °С) GeO2 + MgO → MgGeO3 (1200 °С) GeO2 + C(chark) → Ge + CO2 (500-600 °С, in atmosphere of H2)

Germanium (II) chloride Structure and physical properties

At n. c. it is white or yellow solid.

Boiling point

decomposes at 450 °С

Production

1) by interaction of germanium (VI) chloride and metallic germanium: GeCl4 + Ge → 2GeCl2 (350 °С) 2) by interaction of concentrated hydrochloric acid and compounds of germanium (II): GeS + 2HCl ↔ GeCl2 + H2S

Reactions with chalcogens

2GeCl2 + O2 → GeO2 + GeCl4 (60-70 °С)

Reaction with water

GeCl2 + 2H2O → Ge(OH)2 + 2HCl

Reactions with acids

GeCl2 + HCl(gas) → GeHCl3 (40 °С) GeCl2 + 2HNO3(conc., hot) → GeO2 + 2NO2 + 2HCl GeCl2 + H2S(gas) → GeS + 2HCl (in conc. HCl)

Reactions with bases

GeCl2 + 2NаOH(dil.) → Ge(OH)2 + 2NаCl

Other reactions

2GeCl2 → GeCl4 + Ge (75-460 °С)

Germanium (III) sulphate Structure and physical properties Production

At n. c. it is colorless crystals. By prolonged heating of germanium (IV) chloride with sulphur trioxide in a sealed ampoule:

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Reaction with water Reactions with bases Other reactions

GeCl4 + 4SO3 → Ge(SO4)2 + 2Cl2 + 2SO2 (160 °С) Ge(SO4)2 + 2H2O → GeO2 + 2H2SO4 Ge(SO4)2 + 6NаOH → Nа2GeO3 + 2Nа2SO4 + 3H2O Ge(SO4)2 → GeO2 + 2SO3 (200 °С)

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INDIUM Symbol History of discovery

Atomic mass Position in the Periodic System Electronic configuration Electronegativity Possible oxidation states

In Indium was first discovered in 1863 by German scientists Ferdinand Reich and Hieronymous Theodor Richter in the spectroscopic analysis of polymetallic ores. In the spectrum of the sample, they detected two bright blue lines that did not belong to any of the known elements. They were able to extract a small amount of this element from the mineral, the name of which was given from the name of the dye – indigo. 114.818 Period: 5, group: 13 [Kr]4d105s25p1 1.49 0, +1, +3

Simple substance Structure and physical properties Density (at n. c.), kg m-3 Boiling point Melting point Production

Reactions with halogens Reactions with chalcogens

At n. c. it is a soft metal of silver-white color with a tetragonal lattice. It refers to the group of light metals. 7310 2000 °С 156.4 °С The main source is waste and intermediate products from zinc, lead and tin production. Main steps: 1) production of a concentrate (usually by sulphuric acid); 2) its processing to crude metal (by cementation on zinc or aluminum); 3) refining (chemical, electrochemical, distillation and crystal physical methods). 2In + 3Cl2 → 2InCl3 (120-150 °С) 4In + 3O2 → 2In2O3 (800 °С, air combustion)

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Reactions with acids

Other reactions Application

2In + 3S → In2S3 (1050-1100 °С) In + 4HNO3(dil., hot) → In(NO3)3 + NO + 2H2O In + 2HCl(gas) → InCl2(gas) + H2 (700-970 °С) 2In + 6HCl(dil.) → 2InCl3 + 3H2 2In + H2S → In2S + H2 (700-800 °С) 2In + CO2 → In2O + CO (850 °С) ‒ in semiconductor technology; ‒ anti-corrosion coatings; ‒ for the manufacture of mirrors and reflectors; ‒ for the manufacture of fusible alloys; ‒ solders for bonding glass with metal; ‒ in production of transparent conductive films.

Indium (III) oxide Structure and physical properties Melting point

At n. c. it is an amorphous substance or light yellow crystals of a cubic syngony. decomposes at > 850 °С

Production

1) by combustion of indium in oxygen: 4In + 3O2 → 2In2O3 (800 °С, air combustion) 2) by thermal decomposition of indium hydroxide, nitrate or sulphate: 2In(OH)3 → In2O3 + 3H2O (350 °С) 4In(NO3)3 → 2In2O3 + 12NO2 + 3O2 (230 °С) 2In2(SO4)3 → 2In2O3 + 6SO2 + 3O2 (>600 °С) 3) by oxidation of indium sulphide: 2In2S3 + 9O2 → 2In2O3 + 6SO2 (650 °С) In2O3 + 3H2 → 2In + 3H2O (700 °С) In2O3 + 6HCl(dil., hot) → 2InCl3 + 3H2O In2O3 + 3H2S → In2S3 + 3H2O (500-700 °С)

Reaction with hydrogen Reactions with acids

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Reactions with bases Other reactions

In2O3 + 2NаOH → 2NаInO2 + H2O (500-600 °С) In2O3 → In2O(gas) + O2 (1200-1700 °С) In2O3 + 2NH3 → 2InN + 3H2O (600-630 °С) In2O3 + 3С(graphite) → 2In + 3СO (800-900 °С) 2In2O3 + 3С(graphite) + 6Cl2 → 4InCl3 + 3СO2 (500 °С)

Indium (III) hydroxide Structure and physical properties Melting point Production

Reactions with acids

Reactions with bases Other reactions

At n. c. it is white, amorphous substance, forming crystals of cubic syngony. decomposes at 600 °С By interaction of sulphuric acid and indium, its oxide or hydroxide: 2In + 3H2SO4(dil.) → In2(SO4)3 + 3H2 In2O3 + 3H2SO4(dil.) → In2(SO4)3 + 3H2O 2In(OH)3 + 3H2SO4(dil.) → In2(SO4)3 + 6H2O In2(SO4)3 + 12H2O → 2[In(H2O)6]3+ + 3SO42-

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Reactions with acids Reactions with bases Other reactions

In2(SO4)3 + H2SO4 → 2H[In(SO4)2] In2(SO4)3 + 3H2S → In2S3 + 3H2SO4 In2(SO4)3 + 6NаOH(dil.) → 2In(OH)3 + 3Nа2SO4 2In2(SO4)3 → 2In2O3 + 6SO2 + 3O2 (> 600 °С)

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IODINE Symbol History of discovery

Atomic mass Position in the Periodic System Electronic configuration Electronegativity Possible oxidation states

I Iodine was discovered by the manufacturer of soap and saltpeter, Bernard Courtois, in 1811, who tried to understand the reason for the erosion of the boiler for evaporation of seaweed salt. To this end, he began to add various reagents to the salt, some of which produced heavy violet gas. This fact allowed him to assume the presence of a new unknown element. 126.904 Period: 5, group: 17 [Kr]4d105s25p5 2.21 −1, 0, +1, +3, +5, +7

Simple substance Structure and physical properties Density (at n. c.), kg m-3 Boiling point Melting point Production

Reaction with hydrogen Reactions with halogens

At n. c. it is solid, black-gray or dark purple crystals with a faint metallic luster and a specific odor. The molecule is diatomic. 4930 184.35 °С 113.7 °С 1) in laboratory: by interaction of an oxidizing agent and bromides or iodides in an acidic medium: 8NаI + 9H2SO4 → H2S + 4I2 + 8NаHSO4 + 4H2O 2) in industry: from seaweed and oil drilling water: 2NаI + MnO2 + 3H2SO4 → I2 + 2NаHSO4 + MnSO4 + 2H2O I2 + H2 → 2НI (500 °С, ct. is Pt) I2(suspension) + 3F2 → 2IF3 (-45 °С, in liquid Cl3F) I2 + 5F2 → 2IF5 (room temperature) I2 + 5Cl2 + 6H2O(hot) → 2HIO3 + 10HCl I2 + 5Br2 + 6H2O(hot) → 2HIO3 + 10HBr

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Reactions with chalcogens Reaction with water Reactions with acids

Reactions with bases

Other reactions

Application

4I2 + 9O2 → 2I(IO3)3 (50-60 °С) I2 + H2O → HIO + HI 3I2 + 10HNO3(dil.) → 6HIO3 + 10NO + 2H2O (boiling) I2 + 10HNO3(conc., hot) → 2HIO3 + 10NO2 + 4H2O I2 + 2HClO3 → 2HIO3 + Cl2 3I2 + 2HNO3(conc.) + 6HCl → 6ICl + 2NO + 4H2O (60-80 °С) 3I2 + 6NаOH(hot) → 5NаI + NаIO3 + 3H2O I2 + 2NаOH(dil.) → NаI + NаIO + H2O (0 °С) 3I2 + 4NH4OH → I3N + 3NH4I + 4H2O I2 + 2Nа → 2NаI (>100 °С) 3I2 + 2Аl → 2АlI3 (room temperature, ct. is H2O) I2 + SO2 + 2H2O → 2HI + H2SO4 I2 + 2Nа2SO3S(dil.) → 2NаI + Nа2S4O6 I2 + 3F2 + 2KF → 2K[IF4] I2 + KI(conc.) → K[I(I)2](dil.) ‒ for skin disinfection; ‒ as antiseptics; ‒ as a contrast agent; ‒ for diagnosis and treatment of diseases of thyroid gland; ‒ in forensic science for fingerprint detection; ‒ in technology; ‒ in halogen lamps; ‒ in many organic syntheses; ‒ to absorb carbon monoxide.

Hydriodic acid Structure and physical properties Boiling point Melting point Production

At n. c. it is a colorless asphyxiant gas. When dissolved in water, it forms hydroiodic acid. – 35.4 °С – 50.8 °С 1) in laboratory: а) by redox reactions: H2S + I2 → 2HI + S b) by exchange reactions:

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Reactions with halogens

Reactions with chalcogens

Reaction with water Reactions with acids

Reactions with bases Other reactions

PI3 + 3H2O → H3PO3 + 3HI c) by direct synthesis: I2 + H2 → 2НI (500 °С, ct. is Pt) 2) in industry: by reaction of iodine with hydrazine: N2H4 + 2I2 → N2 + 4HI 2HI + Cl2(dil.) → I2 + 2HCl (room temperature) HI + 3Cl2(saturated) + 3H2O(hot) → HIO3 + 6HCl 4HI(solution) + O2 → 2I2 + 2H2O (in the light, ct. is Cu) 6HI(solution) + O2 ↔ 2H[I(I)2] + 2H2O (room temperature, at darkness) 2HI(gas) + S ↔ I2 + H2S (500 °С) HI(dil.) + H2O → I- + H3O+ 14HI(conc.) + 2H2SO4(conc.) → 7I2 + H2S + S + 8H2O 2HI + HClO → HCl + I2 + H2O 6HI(dil.) + HClO3(dil.) → HCl + 3I2 + 3H2O 5HI(conc.) + 6HClO3(conc.) → 5HIO3 + 3Cl2 + 3H2O HI(dil.) + NаOH(dil.) → NаI + H2O 2HI + NO2 → I2 + NO + H2O 2HI(dil.) + Fe2(SO4)3 → 2FeSO4 + I2 + H2SO4 14HI(conc.) + K2Cr2O7(solid) → 2CrI3 + I2 + 7H2O + 2K[I(I)2] 4HI(conc.) + MnO2 → MnI2 + I2 + 2H2O

Iodine (V) oxide Structure and physical properties

At n. c. it is colorless hygroscopic crystals.

Melting point

decomposes at 300-350 °С

Production

1) by oxidation of iodine with chlorine (VII) oxide: 5Cl2O7 + 7I2 → 7I2O5 + 5Cl2 2) by decomposition of iodic acid: 2HIO3 → I2O5 + H2O (240-250 °С) 3) by interaction of iodine (V) fluoride on silicon oxide: 4IF5 + 5SiO2 → 5SiF4 + 2I2O5 (150-175 °С)

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Reactions with halogens

2I2O5 + 2F2 → 4IO2F + O2 (0 °С, in liq. HF)

Reaction with water

I2O5 + H2O → 2HIO3

Reactions with acids

I2O5 + 2H2SO4(anhydrous) → 2IO2HSO4 + H2O

Reactions with bases

I2O5 + 2NаOH(dil.) → 2NаIO3 + H2O

Other reactions

2I2O5 → 2I2 + 5O2 (300-500 °С) I2O5 + 10HCl(conc.) + 2KCl → 2K[ICl4] + 2Cl2 + 5H2O I2O5 + 5CO → I2 + 5CO2 (room temperature)

Iodic acid Structure and physical properties Melting point Production

Reaction with water Reactions with acids

Reactions with bases Other reactions

At n. c. it is a colorless crystalline substance with a vitreous sheen and a bitter-sour taste. 110 °С By oxidation of iodine by chlorine, hydrogen peroxide or nitric acid: I2 + 5H2O2(conc., hot) → 2HIO3 + 4H2O I2 + 5Cl2 + 6H2O(hot) → 2HIO3 + 10HCl I2 + 10HNO3(conc., hot) → 2HIO3 + 10NO2 + 4H2O HIO3(dil.) + H2O → IO3- + H3O+ 2HIO3(solid) + H2SO4(conc.) → (IO)2SO4 + O2 + 2H2O (60-80 °С) 2HIO3(conc.) + 10HCl(conc., cold) → I2 + 5Cl2 + 6H2O HIO3(conc.) + 5HI(conc.) → 3I2 + 3H2O (room temperature) HIO3 + NаOH(dil.) → NаIO3 + H2O HIO3(dil.) + 3H2O H2(cathode) + H5IO6(anode) 3HIO3 → HI3O8 + H2O (110-120 °С) 2HIO3 → I2O5 + H2O (240-250 °С) 2HIO3 + 5Nа2SO3 → 5Nа2SO4 + I2 + H2O

Iodine (I) chloride Structure and physical properties

At n. c. it is dark red liquid or dark red crystals. There are two polymorphic modifications with monoclinic syngonies:

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Boiling point Melting point Production

Reaction with water Reactions with acids

Reactions with bases Other reactions

1) α-form; 2) β-form. decomposes at 97.4 °С 27.2 °С (α-форма) 13.92 °С (β-форма) 1) by interaction of chlorine on iodine: I2 + Cl2 → 2ICl (45 °С) 2) by thermal decomposition of iodine (III) chloride I2Cl6 → 2ICl + 2Cl2 (64-77 °С) ICl + H2O(cold) → HIO + HCl 5ICl + 3H2O(hot) → 5HCl + HIO3 + 2I2 ICl + HCl(conc.) → H[ICl2] ICl + 2H2SO4(conc.) → HIO3 + HCl + 2SO2 + H2O (boiling) 3ICl + 6NаOH(dil.) → 3NаCl + 2NаI + NаIO3 + 3H2O 2ICl → I2 + Cl2 (>97.4 °С) ICl(liq.) + RbCl → Rb[ICl2]

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KRYPTON Symbol History of discovery

Atomic mass Position in the Periodic System Electronic configuration Possible oxidation states

Kr The name "krypton", meaning "secretive", "secret" first appeared before the discovery of the gas itself. The Scottish chemist Sir William Ramsay named "krypton" the helium already known at that time, which the English chemist Sir William Crookes later found out. In 1898, William Ramsay opened a new gas and again gave it the name "krypton". This discovery was accidental in the process of trying to detect helium in high-boiling air fractions. His attempt was unsuccessful, because helium is a lowboiling gas. Instead of helium, he discovered a new element, not known until that time. 83.798 Period: 4, group: 18 [Аr]3d104s24p6 0, +2, +4, +6, +8

Simple substance Structure and physical properties Density (at n. c.), kg m-3 Boiling point Melting point Production

Reactions with halogens Application

At n. c. it is an inert monatomic gas that does not have color, smell, and taste. 3708 -153.2 °С -157.1 °С Krypton is obtained as a by-product in the separation of air. This forms a mixture of krypton and xenon, which is separated as follows: ‒ contact catalytic furnace (CuO, 300-400 °С) for cleaning from hydrocarbon residues; ‒ adsorber for cleaning from moisture; ‒ heat exchanger for cooling; ‒ distillation columns for separation; ‒ reservoir for transportation and storage. Kr + F2

KrF2

‒ to fill incandescent bulbs;

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‒ production of high-power excimer lasers (Kr-F); ‒ fluorinating agent in inorganic synthesis; ‒ to obtain atomic fluorine.

Krypton (II) fluoride Structure and physical properties

There are two crystalline modifications: α-form and β-form. β-KrF2 is stable at temperatures above – 80 °C. At lower temperatures, the α-form is stable.

Production Kr + F2 Reaction with water Reactions with bases Other reactions

KrF2

2KrF2 + 2Н2O → 2Kr + 4HF +O2 2KrF2 + 4NаOH(dil.) → 2Kr + 4NаF + O2 + 2Н2O HF (l) KrF2 + MnF2 MnF4+ Kr 3KrF2 + Xe → XeF6 + 3Kr 7KrF2+ 2Аu → 2[KrF][АuF6] + 5Kr KrF2 + ClF3 → Kr + ClF5

77

LEAD Symbol History of discovery

Atomic mass Position in the Periodic System Electronic configuration Electronegativity Possible oxidation states

Pb Lead has been known from the 3rd-2nd millennium BC in Mesopotamia, Egypt and other ancient countries, where large bricks, statues of gods and kings, seals, and various household items were made from it. 207.2 Period: 6, group: 14 [Xe]6s25d104f146p2 1.55 0, + 2, + 4

Simple substance Structure and physical properties Density (at n. c.), kg m-3 Boiling point Melting point Production

Reaction with oxygen

Reactions with halogens

At n. c. it is gray with a blue tint, soft, malleable, covered with an oxide film in air. It is a weak reducing agent. 11337 1751 °С 327.3 °С Pb(NO3)2 + Zn → Pb + Zn(NO3)2 PbSO4 + Zn → Pb + ZnSO4 PbCl2 + Н2→ Pb + 2HCl PbS + Н2→ Pb + H2S 2Pb + O2 → 2PbO (>600 °С) 3Pb +2O2 → (PbII2PbIV)O4 (400-500 °С) Pb + F2 → PbF2 (200-300 °С) Pb + Cl2 → PbCl2 (200-300 °С) Pb + Br2 → PbBr2 (200-300 °С) Pb + I2 → PbI2 (200-300 °С)

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Reactions with chalcogens

Reaction with water Reactions with acids

Reactions with bases Application

Pb + S → PbS (800-1200 °С) Pb + Se → PbSe (800-1200 °С) Pb + Te → PbTe (800-1200 °С) 2Pb + 2Н2O + O2 → 2Pb(OH)2 Pb + 3H2SO4(> 80%) → Pb(HSO4)2+ SO2 + 2Н2O Pb + 2H2SO4(conc.) → PbSO4+ SO2+ 2Н2O 3Рb + 8HNO3(dil., hot) → 3Pb(NO3)2 + 2NO + 4Н2O Pb + 2NаOH(conc.) + 2Н2O → Nа2[Pb(OH)4] + H2 ‒ for the manufacture of some alloys (babbits, brass); ‒ for the manufacture of typographic font; ‒ for the manufacture of batteries, solders; ‒ for the manufacture of bullets and nuclei; ‒ for the manufacture of lead pipes; ‒ for the manufacture of red lead and whitewash; ‒ for the manufacture of glaze for pottery; ‒ for the manufacture of crystal; ‒ to slow the growth of tumors; ‒ as an additive to engine fuel.

Lead (IV) oxide Structure and physical properties Melting point Production

Reaction with water Reactions with acids

Reactions with bases Other reactions

At n. c. it is a dark brown powder, a strong oxidizing agent in acid and alkaline medium. decomposes at 290 °С 2PbO + Cа(ClO)2 → 2PbO2 + CаCl2 PbCl4 + 2Н2O → PbO2 + 4HCl PbCl4 + 4NаOH(dil.) → PbO2 + 4NаCl + 2Н2O PbO2(solid) + 2H2O → PbIV + 4OHPbO2 + 4HCl(conc., hot) → PbCl2 + Cl2 + 2Н2O PbO2 + H2SO4(conc., cold) → PbSO4 + 2Н2O 2PbO2 + 2H2SO4(conc., hot) → 2PbSO4 + O2 + 2Н2O PbO2 + 2NаOH(conc.) + 2Н2O → Nа2[Pb(OH)6] 2РbO2 + 4KO2 → 2K2РbO3 + 3O2 PbO2 + 2S → PbS + SO2 (400 °С) PbO2(moistered) + 2H2S(gas) → PbS + S + 2H2О PbO2(moistered) + SO2 → PbSO4 PbO2 + 4HNO3(dil.) + 2KI → Рb(NO3)2 + I2 + 2H2О + + 2KNO3

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PbO2 + 2HNO3(dil.) + H2O2(conc.) → Pb(NO3)2 + O2 + + 2H2O 5PbO2 + 6HNO3(dil.) + 2Mn(NO3)2 → 5Pb(NO3)2 + + 2HMnO4 + 2H2O

Lead (II) hydroxide Structure and physical properties

At n. c. it is a white solid, which decomposes with weak heating. It is not soluble in water; shows amphoteric properties; absorbs CO2 from the air. It is a weak reducing agent. Pb(NO3)2 + 2NаOH(dil.) → Pb(OH)2 + 2NаNO3 Pb(NO3)2 + 2(NH3·H2O)(conc.) → Pb(OH)2 + + 2NH4NO3 PbCl2 + 2NаOH(dil.) → Pb(OH)2 + 2NаCl

Production

Pb(OH)2(solid) + 3H2O → [Pb(H2O)3]2+ + 2OHPb(OH)2 + 2НСl(dil.) → РbСl2 + 2Н2O Pb(OH)2(suspension) + H2SO4(dil.) → PbSO4 + 2Н2O Pb(OH)2 + 2HNO3(dil.) → Pb(NO3)2 + 2Н2O Pb(OH)2 + 2NаOH(conc.) → Nа2[Pb(OH)4] Pb(OH)2 → РbО + Н2O (100-145 °С) 2Pb(OH)2(suspension) + СO2 → Рb2СО3(ОН)2 + Н2O Рb(ОН)2 + Н2O2(conc.) → РbO2 + 2Н2O (in dil. NаOH) 2Рb(ОН)2 + Са(СlO)2 → 2РbO2 + СаСl2 + 2Н2O (in dil. NаOH)

Reaction with water Reactions with acids

Reactions with bases Other reactions

Lead (IV) chloride Structure and physical properties Boiling point Melting point Production

Reactions with halogens

At n. c. it is a yellow oily liquid. exploses at 105 °С -15 °С Pb + Cl2 → PbCl2 (200-300 °С) PbO + 2HCl(dil.) → PbCl2 + 2Н2O Pb(OH)2+ 2НСl(dil.) → РbСl2 + 2Н2O РbСО3 + 2НСl(dil.) → РbСl2 + Н2O + CO2 PbCl2 + Cl2 + 2HCl(conc.) → H2[PbCl6]

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Reaction with water Reactions with acids

Reactions with bases Other reactions

PbCl2 + H2O(vapour) → PbCl(OH) + HCl (50 °С) PbCl2 + 2HCl(conc.) → H2[PbCl4] (room temperature, impurities of H[PbCl3]) H2[PbCl4](solution) → PbCl2 + 2HCl PbCl2 + H2SO4(conc., hot) → PbSO4 + 2HCl PbCl2 + 2NаOH(dil.) → Рb(OH)2 +2NаCl PbCl2 + 4NаOH(conc.) → Nа2[Pb(OH)4] + 2NаCl PbCl2 + H2 → Pb + 2HCl (300-350 °С) PbCl2 + 2KI(dil.) → PbI2 +2KCl PbCl2 + H2S(saturated) → PbS +2HCl PbCl2 + K2CrO4 → PbCrO4 +2KCl 2PbCl2 + Nа2CO3(dil.) + 2NаOH(dil.) → Pb2CO3(OH)2 + 4NаCl PbCl2 + 4Nа2S2O3(conc., cold) → Nа6[Pb(S2O3)4] + 2NаCl PbCl2 + Cl2 + 2KCl → K2[PbCl6] (in conc. HCl)

Lead (II) nitrate Structure and physical properties

Melting point Production

Reaction with water Reactions with acids Reactions with bases

Other reactions

At n. c. it is white when heated decomposes. It is well dissolved in cold water. When boiling the solution decomposes. It does not form crystalline hydrates. It enters the exchange reaction. It is a strong oxidizer during sintering. decomposes at 470 °С PbO + 2HNO3(dil.) → Pb(NO3)2 + H2O Pb(OH)2 + 2HNO3(dil.) → Pb(NO3)2+ 2Н2O 3Рb + 8HNO3(dil., hot) → 3Pb(NO3)2 + 2NO + 4Н2O Pb(NO3)2(conc.) + H2O → PbNO3(OH) + HNO3 Pb(NO3)2 + 2HCl(dil.) → PbCl2 + 2HNO3 Pb(NO3)2 + H2SO4(dil.) → PbSO4 + 2HNO3 Pb(NO3)2 + 2NаOH(dil.) → Pb(OH)2 + 2NаNO3 Pb(NO3)2 + 4NаOH(conc.) → Nа2[Pb(OH)4] + + 2NаNO3 PbCl2 + Н2 → Pb + 2HCl Pb(NO3)2 + 2KCl → PbCl2 + 2KNO3 Pb(NO3)2 + 2KI → PbI2 + 2KNO3 Pb(NO3)2(conc.) + 3KI(conc.) → K[PbI3] + 2KNO3 Pb(NO3)2 + KF + KCl → Pb(Cl)F + 2KNO3 (in very dil. HNO3) Pb(NO3)2 + KF + KBr → Pb(Br)F + 2KNO3 (in very dil. HNO3)

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Pb(NO3)2 + 2KBrO3 → Pb(BrO3)2 + 2KNO3 Pb(NO3)2 + 2KIO3 → Pb(IO3)2 + 2KNO3 Pb(NO3)2 + Nа2S → PbS + 2NаNO3 Pb(NO3)2 + Nа2Se → PbSe + 2NаNO3 Pb(NO3)2 + Nа2Te → PbTe + 2NаNO3 Pb(NO3)2 + Nа2SO4 → PbSO4 + 2NаNO3 Pb(NO3)2 + Nа2SeO4 → PbSeO4 + 2NаNO3

Lead (II) sulphate Structure and physical properties Melting point Production

Reactions with acids

Reactions with bases

Other reactions

At n. c. it is a white solid. When calcined it decomposes; it melts under pressure of O2. decomposes at 1000 °С Pb(NO3)2 + H2SO4(dil.) → PbSO4 + 2HNO3 Pb + 2H2SO4(conc.) → PbSO4+ SO2+ 2Н2O PbO2 + 2H2SO4(conc., cold) → Pb(SO4)2 + 2Н2O 2PbO2 + 2H2SO4(conc., hot) → 2PbSO4 + O2 + 2Н2O PbCl2 + H2SO4(conc., hot) → PbSO4↓ + 2HCl↑ PbSO4 + 4HCl(conc.) → H2[PbCl4]+ H2SO4 PbSO4 + H2SO4(conc.) → Pb(HSO4)2 Pb(HSO4)2(dil.) → PbSO4 + H2SO4 2PbSO4 + 2HNO3(conc., hot) → Pb(NO3)2 + Pb(HSO4)2 2PbSO4 + 2NаOH(dil.) → Pb2SO4(OH)2 + Nа2SO4 (boiling) PbSO4 + 4NаOH(conc.) → Nа2[Pb(OH)4] + Nа2SO4 PbSO4 + 2KI(dil.) → PbI2 + K2SO4 PbSO4 + K2CrO4 → PbCrO4 + K2SO4 PbSO4 + Nа2S → PbS + Nа2SO4 PbSO4 + Nа2CO3(conc.) → PbCO3 + Nа2SO4 (10-12 °С) PbSO4 + 4H2 → PbS + 4H2O (500-600 °С) PbSO4 + 2C(chark) → PbS + 2CO2 (550-650 °С) PbSO4 + PbS → 2Pb + 2SO2 (800-1000 °С) PbSO4(moistered) + Zn(plate) → Pb(sponge) + ZnSO4

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Lead (II) carbonate Structure and physical properties

Melting point Production

Reaction with water Reactions with acids Reactions with bases Other reactions

At n. c. it is white, at weak heating it decomposes. It does not dissolve in cold water. It is transferred to the solution by the action of CO2. The ions are packed according to the closest hexagonal packaging method. decomposes at 315 °С PbSO4 + Nа2CO3(conc.) → PbCO3 + Nа2SO4 Pb(NO3)2(conc.) + Nа2CO3(conc.) → PbCO3 + 2NаNO3 (10-12 °С) 2РbСO3 +Н2O → Рb2СО3(ОН)2 or [Рb3(СО3)2(ОН)2]+ + CO2 РbСО3 + 2НСl(dil.) → РbСl2 + Н2O + CO2 2РbСО3 + 2NаOH(dil.) → Рb2СО3(ОН)2 + Nа2CO3 РbСО3 + 4NаOH(conc.) → Nа2[Pb(OH)4] + Nа2CO3 РbСО3 → РbО + СO2 (>315 °С) 6РbСО3 + O2 → 2(PbII2PbIV)O4 + 6CO2 2РbСО3 + 3S → 2PbS + 2CO2 + SO2 РbСО3 + 2НF(conc.) → РbF2 + CO2 + Н2O РbСО3 + H2S(saturated) → PbS + CO2 + Н2O

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NEON Symbol

Ne

History of discovery

Neon was discovered in 1898 by British scientists Sir William Ramsay and Morris William Travers during a spectral study of the residues of slowly evaporating liquid air. The name comes from the Greek word “neos”, meaning “new”. According to legend, the word “new” shouted the son of Ramsay, who saw the unusual bright red radiation emitted by the substance in the tube for spectral analysis.

Atomic mass

20.180

Position in the Periodic System

Period: 2, group: 18

Electronic configuration

[He]2s22p6

Possible oxidation states

0

Simple substance Structure and physical properties

At n. c. it is an inert monatomic gas without color and smell.

Density (at n. c.), kg m-3

900

Boiling point

−245.9 °С

Melting point

−248.67 °С

Production

In industry, neon is obtained as a by-product of helium production. There are several ways to separate neon from helium: ‒ adsorption, which is based on the ability of neon to be adsorbed by activated carbon cooled by liquid nitrogen; ‒ condensation, based on the freezing of neon when the mixture is cooled with liquid hydrogen; ‒ rectification, based on the difference between the boiling points of helium and nitrogen.

Application

‒ cryogenic cooler; ‒ in gas-discharge lamps, signal lamps, in radio equipment, photocells, rectifiers; ‒ in helium-neon lasers; ‒ in billboards.

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NITROGEN Symbol History of discovery

Atomic mass Position in the Periodic System Electronic configuration Electronegativity Possible oxidation states

N The presence of an inert gas in the atmosphere was first established by the French naturalist Antoine-Laurent de Lavoisier. He failed to establish the nature of the gas remaining after combustion. In 1772, the Scottish physician, chemist and botanist Daniel Rutherford found that this gas was an element. 14.007 Period: 2, group: 15 [He]2s22p3 3.07 -3, 0, +2, +3, +4, +5

Simple substance Structure and physical properties

At n. c. it is a diatomic gas with no color, smell or taste.

Density (at n. c.), kg m-3

1250.6

Boiling point

-195.8 °С

Melting point

-210.0 °С

Production

1) in laboratory: а) by decomposition of ammonium nitrate: 2NH4NO3 → 2N2 + 4H2O + O2 b) by heating a mixture of potassium dichromate and ammonium sulphate (2:1 by weight) K2Cr2O7 + (NH4)2SO4 → (NH4)2Cr2O7 + K2SO4 (NH4)2Cr2O7 → N2 + Cr2O3 + 4H2O c) by decomposition of metal azides: 2NаN3 → 2Nа + 3N2 (250-300 °С) d) by passing ammonia over copper (II) oxide: 3CuO + 2NH3 → N2 + 3Cu + 3H2O 2) in industry: by fractional distillation of atmospheric air.

Reaction with hydrogen

N2 + 3H2 ↔ 2NH3 (room temperature, electric discharge)

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N2 + 3H2 → 2NH3 (500 °С, Р, ct. is Fe, Pt) N2 + H2 ↔ N2H2 (1000 °С) Reactions with halogens

N2 + 3F2 ↔ 2NF3 (electric discharge)

Reactions with chalcogens

N2 + O2 ↔ 2NO (room temperature, electric discharge) N2 + O2 ↔ 2NO (2000 °С, ct. is Pt/MnO2)

Other reactions

N2 + 2C(graphite) ↔ C2N2 (electric discharge) 2B + N2 → 2BN N2(moistered) + 6Li ↔ 2Li3N (room temperature) N2 + 6Nа ↔ 2Nа3N (100 °С, electric discharge) N2 + 3Mg ↔ Mg3N2 (780-800 °С, on air) N2 + 2Аl(powder) ↔ 2АlN (800-1200 °С) N2(moistered) + 3LiH ↔ Li3N + NH3 (500-600 °С) N2 + CаC2 → Cа(CN)2 (300-350 °С) N2 + CаC2 → CаCN2 + C(graphite) (1000-1150 °С) N2 + H2SO4(conc.) + 4H2O + 4VSO4 → (N2H5)2SO4 + + 4(VO)SO4 (boiling)

Application

‒ to obtain nitric acid, nitrogen-containing salts, urea, and soda; ‒ production of mineral fertilizers; ‒ production of explosives; ‒ creating an inert atmosphere; ‒ for deep cooling and freezing.

Ammonia Structure and physical properties

At n. c. it is a colorless gas with a sharp characteristic odor. The ammonia molecule has a pyramidal shape. In liquid ammonia, the molecules are hydrogen bonded.

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Boiling point

-33.35 °С

Melting point

-77.8 °С

Production

1) in laboratory: а) by interaction of strong alkali on ammonium salts: NH4Cl + NаOH → NH3 + NаCl + H2O b) by weak heating of a mixture of ammonium chloride with slaked lime: 2NH4Cl + Cа(OH)2 → CаCl2 + 2NH3 + 2H2O 2) in industry: а) by direct synthesis (Haber process): N2 + 3H2 → 2NH3

Reactions with halogens

4NH3 + 3F2 → NF3 + 3NH4F (130-140 °С, in atmosphere of N2) 2NH3 + Cl2 → NH2Cl + NH4Cl (room temperature, in atmosphere of N2) 8NH3 + 3Cl2 → N2 + 6NH4Cl

Reactions with chalcogens

4NH3 + 3O2 → 2N2 + 6H2O 4NH3 + 5O2 → 4NO + 6Н2O (800° С, ct. is Pt/Rh) 3NH3 + S ↔ [S(NH3)3]0 (-40 °С, impurities of S4N4)

Reaction with water

NH3(gas) + H2O ↔ NH4OH ↔ NH4+ + OH-

Reactions with acids

NH3 + HCl(gas) → NH4Cl NH3 + H2SO4 → NH4HSO4 2NH3 + H2SO4 → (NH4)2SO4 NH3 + H2S → NH4HS (0 °С, in ether) 2NH3(liq.) + H2S → (NH4)2S (-40 °С)

Other reactions

2NH3 ↔ N2H4 + Н2 (room temperature, UV irradiation) 2NH3 ↔ N2 + 3Н2 (1200-1300 °С) 2NH3 + 4O3 → NH4NO3 + 4O2 + Н2O (room temperature) 16NH3 + 4S + 6Cl2 → S4N4 + 12NH4Cl (30-50 °С, in liquid CCl4) 4NH3 + 3OF2 → 3H2O + 6HF + 2N2 (200 °С) NH3(gas) + H2O + СО2(gas) → NH4HCO3 (room temperature, Р)

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2NH3 + CO2 → NH4(NH2COO) (room temperature) 2NH3 + CO2 → C(NH2)2O + H2O (180-500 °С, Р) 2NH3 + С(chark) + Nа2CO3 → 2NаCN + 3H2O (800-900 °С) NH3 + СО → HCN + Н2O (500-800 °С, ct. is Аl2O3/ThO2) 2NH3 + (CN)Cl → NH4Cl + H2CN2 (in ether) 2NH3 + 2Li → 2LiNH2 + H2 (220 °С) NH3 + 2Li → Li2NH + H2 (400 °С) 2NH3 + 2Nа → 2NаNH2 + H2 (350 °С) 2NH3 + 3Mg → Mg3N2 + 3H2 (600-850 °С) 2NH3 + 2Аl → 2АlN + 3H2 (> 600 °С) 2NH3 + 6MnO2 → 3Mn2O3 + N2 + 3H2O (500-600 °С) 2NH3 + 3CuO → 3Cu + N2 + 3H2O (500-550 °С) 4NH3(liq.) + Nа → [Nа(NH3)4] (-40 °С) 2NH3(liq.) + 2Nа → 2NаNH2 + H2 (-40 °С, ct. is Fe) 4NH3(liq.) + 2NH4Cl + Mg → [Mg(NH3)6]Cl2 + H2 (-40 °С) 2NH3(liq.) + Cа → Cа(NH2)2 + H2 (-40 °С, ct. is Fe) 8NH3(liq.) + CаI2 → [Cа(NH3)8]I2 (-40 °С) 12NH3(liq.) + АuCl → АuCl·12NH3 (-40 °С) 6NH3(gas) + CuCl2 → [Cu(NH3)6]Cl2 (0 °С, in ethyl acetate) 4NH3(gas) + H2O + 2Hg2(NO3)2 → (Hg2N)NO3·H2O+ + 2Hg + 3NH4NO3

Nitrogen (V) oxide Structure and physical properties

At n. c. it is colorless, highly volatile crystals formed by NO2+ and NO3− ions. Gaseous nitric

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anhydride consists of individual non-planar molecules corresponding to the formula O2N-О-NO2. Melting point Production

Reaction with water Reactions with acids Reactions with bases Other reactions

sublimates at 32.3 °С а) by dehydration of nitric acid: 2HNO3 + P2O5 → 2HPO3 + N2O5 b) by interaction of nitrogen (IV) oxide with ozone: 2NO2 + O3 → N2O5 + O2 N2O5 + H2O → 2HNO3 N2O5 + НСlO4(anhydrous) → (NO2+)СlO4 + HNO3 N2O5 + 2NаOH(dil.) → 2NаNO3 + Н2O N2O5 + 2NH4OH(dil.) → 2NH4NO3 + Н2О 2N2O5 → 4NO2 + O2 (20-50 °С) N2O5(liq.) + 2NH3 → Н2O + 2(NO2)NH2 3N2O5 + Аl2О3 → 2Аl(NО3)3 (35-40 °С) N2O5 + 5Cu → 5CuO + N2 (500 °С)

Nitric acid Structure and physical properties Boiling point Melting point Production

Reactions with halogens Reactions with chalcogens Reaction with water Reactions

At n. c. it is a liquid. Solid nitric acid forms two modifications: with monoclinic and rhombic lattices. 86 °С -42 °С Modern production method is based on the catalytic oxidation of synthetic ammonia: 4NH3 + 5O2 → 4NO + 6H2O (ct. is Pt) 2NO + O2 → 2NO2 4NO2 + O2 + 2H2O → 4HNO3 10HNO3(conc., hot) + I2 → 2НIO3 + 10NO2 + 4H2O 6HNO3(conc.) + S → H2SO4 + 6NO2 + 2H2O (boiling) HNO3 + Н2O → NO3- + Н3O+ HNO3(anhydrous, cold) + 2HF(liq.) → H2NO3+ + HF2-

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with acids

Reactions with bases Other reactions

HNO3(anhydrous, hot) + 4HF(liq.) ↔ H3O+ + NO2+ + 2HF2HNO3(conc.) + 3HCl(conc.) ↔ (NO)Cl + 2Cl0 + 2H2O (room temperature) 2HNO3(conc.) + 6HCl(conc.) → 2NO + 3Cl2 + 4H2O (100-150 °С) 6HNO3(60%) + HI → HIO3 + 6NO2 + 3H2O (boiling) 2HNO3(conc., cold) + H2S(saturated) → S + 2NO2 + 2Н2O HNO3(anhydrous) + 2H2SO4(anhydrous) ↔ H3O+ + NO2+ + 2HSO4HNO3(dil.) + NаOH → NаNO3 + Н2O HNO3(dil.) + NH4OН → NH4NO3 + Н2O 4HNO3 ↔ 4NO2 + 2Н2O + O2 (room temperature, in the light) 2HNO3(2-3%) + 8Н0(Zn, H2SO4(dil.)) → NH4NO3 + 3Н2O 2HNO3(5%) + 8Н0(Mg, H2SO4(dil.)) → N2O + 5Н2O HNO3(30%) + 3Н0(Zn, H2SO4(dil.)) → NO + 2H2O HNO3(60%) + 2Н0(Zn, H2SO4(dil.)) → HNO2 + H2O (ct. is Pd) 2HNO3(conc.) + Аg → АgNO3 + NO2 + H2O 8HNO3(dil.) + 3Cu → 3Cu(NO3)2 + 2NO + 4H2O 10HNO3(dil.) + 4Mg → 4Mg(NO3)2 + N2O + 5H2O (impurities of H2) 12HNO3(dil.) + 5Sn → 5Sn(NO3)2 + N2 + 6H2O (impurities of NO) 30HNO3(very dil.) + 8Аl → 8Аl(NO3)3 + 3NH4NO3 + 9H2O (impurities of H2) 12HNO3(very dil.) + 5Fe → 5Fe(NO3)2 + N2 + 6H2O (0-10 °С) 4HNO3(dil.) + Fe → Fe(NO3)3 + NO + 2H2O 4HNO3(conc., hot) + Hg → Hg(NO3)2 + 2NO2 + 2H2O 8HNO3(dil., cold) + 6Hg → 3Hg2(NO3)2 + 2NO + 4H2O 4HNO3(conc.) + Ge → GeO2 + 4NO2 + 2H2O 5HNO3(conc.) + P(red) → Н3PO4 + 5NO2 + H2O (boiling) 2HNO3(dil.) + MgO → Mg(NO3)2 + H2O 2HNO3(dil.) + Cu(OH)2 → Cu(NO3)2 + 2H2O 4HNO3(conc.) + Nа2[Zn(OH)4] → Zn(NO3)2 + 2NаNO3 + 4H2O 2HNO3(dil.) + Nа2[Zn(OH)4] → Zn(OH)2 + 2NаNO3 + 2H2O 2HNO3 + Nа2CO3 → 2NаNO3 + CO2 + H2O 2HNO3(very dil.) + CаSO3 → Cа(NO3)2 + SO2 + H2O HNO3(conc.) + KF(solid) → KNO3 + HF 3HNO3(dil.) + [Аg(NH3)2]OH → АgNO3 + 2NH4NO3 + H2O 6HNO3(conc.) + [Ni(NH3)6](NO3)2 → Ni(NO3)2 + 6NH4NO3 HNO3(conc.) + H2O2(conc.) ↔ HNO2(O22-) + H2O 2HNO3(conc., hot) + SO2 → H2SO4 + 2NO2

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2HNO3(conc.) + Аs2O3 + 2H2O → 2H3АsO4 + N2O3 (0 °С) 4HNO3(conc.) + Аs2O3 + H2O → 2H3АsO4 + 4NO2 (boiling) 4HNO3(conc.) + 3KI(solid) → K[I(I)2] + 2NO2 + 2H2O + 2KNO3 (room temperature) 4HNO3(conc.) + FeСl2 → Fe(NO3)3 + 2НСl + NO2 + Н2O 4HNO3(conc.) + Nа2S → 2NаNO3 + 2NO2 + S+ 2H2O 8HNO3(conc.) + CuS(solid) → CuSO4 + 8NO2 + 4H2O (boiling) HNO3(conc.) + Nа(SO3NH2) → NаHSO4 + (NO2+)NH2HNO3(conc.) + 4HCl(conc.) + Аu → H[АuCl4] + NO + 2H2O 4HNO3(conc.) + 18HCl(conc.) + 3Pt → 3H2[PtCl6] + 4NO + 8H2O 4HNO3(conc.) + 18HF(conc.) + 3Si → 3H2[SiF6] + 4NO + 8H2O 2HNO3(conc., hot) + 4HF(conc.) + W → H2[WO2F4] + 2NO + 2H2O 2HNO3(dil.) + 3H2SO4(dil.) + 6Hg → 2NO + 3Hg2SO4 + 4Н2О 4HNO3(fuming) + Р4О10 → 2N2O5 + 4НРO3 (in atmosphere of O2 + O3) HNO3(anhydrous) + F2 → (NO2)OF + HF (room temperature) HNO3(anhydrous) + HSO3Cl → (NO2)Cl + H2SO4 (0 °C) HNO3(anhydrous) + 2НСlO4(anhydrous) → (NO2+)ClO4 + НСlO4·H2O (room temperature) HNO3(anhydrous) + 2SO3 ↔ NO2+ + HS2O76HNO3(anhydrous) + 2K4[Fe(CN)6] → 2K2[Fe(NO+)(CN)5] + 2HCN + O2 +4KNO3 + 2H2O 3HNO3(conc.) +K4[Fe(CN)6] → NO2 +HCN +K2[Fe(H2O)(CN)5] + + 2KNO3 (boiling) HNO3(anhydrous) + KNO3 → K+ + [H(NO3)2]-

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OXYGEN Symbol

O

History of discovery

Oxygen was discovered independently by two scientists: the Swedish chemist Carl Wilhelm Scheele and the English chemist Joseph Priestley. Scheele obtained oxygen a little earlier but published his work later. Due to this, it is officially considered that oxygen was opened by Priestley on August 1, 1774. In his experiment, he decomposed mercury oxide in a closed vessel, but he considered that it was not a new element, but only an integral part of air, calling it “dephlogisticated air”. Only the French scientist Antoine Laurent Lavoisier was able to prove that oxygen is a chemical element.

Atomic mass

15.999

Position in the Periodic System

Period: 2, group: 16

Electronic configuration

[He]2s22p4

Electronegativity

3.50

Possible oxidation states

-2, −1, -1/2, -1/3, 0, +1/2, +1, +2

Simple substance Structure and physical properties

At n. c. it is a gas without color, taste and smell, the molecule of which consists of two oxygen atoms. The allotropic modification is triatomic oxygen − ozone − a blue gas with a characteristic odor.

Density (at n. c.), kg m-3

1428.95

Boiling point

-182.97 °С

Melting point

-218.7 °С

Production

1) in laboratory: а) by heating of potassium permanganate: 2KMnO4 → K2MnO4 + MnO2 + O2 (200-240 °С) b) by decomposition of hydrogen peroxide: 2H2O2 → 2H2O + O2 (ct. is MnO2)

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c) by decomposition of potassium chlorate: 2KClO3 → 2KCl + 3O2 (150-300 °С, ct. is MnO2) 2) in industry: by cryogenic air rectification. Reaction with hydrogen

О2 + 2Н2 → 2Н2О (550 °С)

Reactions with halogens

O2 + F2 → O2F2 (-183 °С, electric discharge)

Reactions with chalcogens

O2 + S → SO2

Other reactions

O2 + 2H0(Zn, HCl(dil.)) → H2O2 O2 + N2 ↔ 2NO (electric discharge) 5O2 + 4P(red) ↔ P4O10 O2 + C(graphite) → CO2 (600-700 °С) O2 + 2C(graphite) → 2CO (>1000 °С) O2(air) + 4Li → 2Li2O (>200 °С, impurities of Li2O2) O2 + 2Nа→ Nа2O2 (impurities of Nа2O) O2(air) + K→ KO2 (impurities of K2O2) O2(air) + Cs→ CsO2 O2 + 2Mg→ 2MgO O2 + 4Fe(OH)2(suspension) → 4FeO(OH) + 2H2O O2 + PtF6 → O2[PtF6]

Application

‒ ‒ ‒ ‒ ‒ ‒ ‒ ‒ ‒ ‒ ‒ ‒

in metallurgy; in welding and cutting of metals; rocket fuel; in medicine; in food industry; in chemical industry; in agriculture; for drinking and cooking; as a solvent; as heat carrier; as moderator; for fire extinguishing.

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PHOSPHORUS Symbol

P

History of discovery

In 1669, Hennig Brand, in an attempt to find the basics of the philosopher's stone, decided to investigate physiological products, namely, human urine. Having evaporated about a ton of soldier’s urine to a syrupy liquid, he overtook it several times, and the residue was ignited. So he noticed the formation of white matter, which glowed in the dark. It was called "cold fire." The secondary name "phosphorus" comes from the Greek words "φώς" – light and "φέρω" – carry.

Atomic mass

30.974

Position in the Periodic System

Period: 3, group: 15

Electronic configuration

[Ne]3s23p3

Electronegativity

2.10

Possible oxidation states

−3, -1, 0, +1, +3, +5

Simple substance Structure and physical properties

Phosphorus exists in the form of several allotropic modifications: 1) white phosphorus (P4) is the most active allotropic modification; 2) red phosphorus is more thermodynamically stable modification; 3) black phosphorus is the most stable thermodynamically and chemically the least active form of elemental phosphorus; 4) metallic phosphorus, which conducts electricity well.

Density (at n. c.), kg m-3

1820

Boiling point

44.1 °С

Melting point

275 °С

Production

Phosphorus is extracted from phosphorites and apatites: 2Cа3(PO4)2 + 10C + 6SiO2 → P4 + 10CO + 6CаSiO3 (1000 °С) By reduction of other inorganic phosphorus compounds:

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4HPO3 + 10C → P4 + 2H2O + 10CO Reaction with hydrogen

2P + 3H2 → 2PH3 (300-360 °С, Р)

Reactions with halogens

2P + 5Cl2 → 2PCl5 (90 °С) 2P + 3I2 → 2PI3 (boiling, in CS2)

Reactions with chalcogens

4P + 5O2 → P4O10 (240-400 °С) 4P + 9S → P4S9 (550 °С, P, impurities of P4S7) P4 + 3Se → P4Se3 (boiling, in heptane) P4 + 3Te → P4Te3 (boiling, in heptane)

Reaction with water

2P + 8H2O(liq.) → 5H2 + 2H3PO4 (700-900 °С, P, ct. is Pt, Cu, Ti, Zr)

Reactions with acids

P + 5HNO3 → 3H2PO4 + 5NO2 + H2O (boiling) P4 + 6H2SO4(conc.) → 4H2(PHO3) + 6SO2 P4 + 6HCl → 2PH3 + 2PCl3 (300 °С)

Reactions with alkalis

P4 + 8NаOH(conc.) + 4H2O → 4Nа2(PHO2) + 6H2 (boiling) P4 + 3NаOH(conc.) + 3H2O(cold) → 3Nа(PH2O2) + PH3 2P4 + 3Bа(OH)2(conc.) + 6H2O(cold) → 3Bа(PH2O2)2 + 2PH3 (70 °С)

Other reactions

2P + 3Cа → Cа3P2 (350-450 °С) P4 + 6N2O → P4O6 + 6N2 (550-625 °С) P4 + 6CO2 → P4O6 + 6CO (650 °С) 12P + 10KClO3 → 10KCl + 3P4O10 (50 °С) P4 + 4H2SO4(dil.) + 4KMnO4 → 4KH2PO4 + 4MnSO4 (room temperature) 6P + 4H2O(hot) + 8KMnO4 → 3K2H2P2O6 + 8MnO2 + 2KOH P4 + 16H2O + 20АgNO3 → 4H3PO4 + 20Аg + 20HNO3 (boiling)

Application

‒ in the production of matches; ‒ as a desiccant of gases and liquids; ‒ in organic synthesis in the reactions of dehydration and condensation;

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‒ when soldering as a flux (for oxidized copper, for ferrous metal, for stainless steel); ‒ for research in the field of molecular biology; ‒ to remove rust from metal surfaces; ‒ in the composition of freons, in industrial freezers as a binder; ‒ in the manufacture of medicines and dyes.

Phosphine Structure and physical properties Boiling point Melting point Production

Reactions with halogens Reactions with chalcogens Reactions with acids

Other reactions

At n. c. it is a colorless poisonous gas with a peculiar smell of rotten fish. The phosphine molecule has the shape of a trigonal pyramid with molecular symmetry. -87.8 °С -133.8 °С 1) by interaction of white phosphorus with hot alkali: 2P4 + 3Cа(OH)2 + 6H2O → 2PH3 + 3Cа(H2PO2) P4 + 3NаOH + 3H2O → 3NаH2PO2 + PH3 2) by interaction of water or acids and phosphides: Cа3P2 + 6HCl → 2PH3 + 3CаCl2 Mg3P2 + 3H2SO4(dil.) → 2PH3 + 3MgSO4 3) by interaction of hydrogen chloride and white phosphorus when heated: P4 + 6HCl → 2PH3 + 2PCl3 4) by decomposition of phosphonium iodide: PH4I → PH3 + HI 5) by decomposition or reduction of phosphorous acid: 4H2(PHO3) → PH3 + 3H3PO4 (170-200 °С) H2(PHO3) + 3H2 → PH3 + 3H2O PH3 + 2H2O + 2I2 → H(PH2O2) +4HI PH3 + 2O2 → H3PO4 (150 °С) PH3 + 3H2SO4(conc.) → H2(HPO3) + 3SO2 + 3H2O PH3 + 8HNO3(conc., hot) → H3PO4 + 8NO2 + 4H2O PH3 + HCl(gas) → PH4Cl (30 °С, in absence of moisture) PH3 + HI(conc.) → PH4I PH3 + NаOH(dil.) + 2NаClO → Nа(PH2O2) + 2NаCl + H2O 2PH3 + 4NаOH(conc.) + 7H2O2(conc.) → Nа4P2O6 + 12H2O 4PH3 + Ni → [Ni(PH3)4](solid) (room temperature, Р)

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Phosphorus (V) oxide Structure and physical properties

Boiling point Melting point Production

Reaction with halogens Reaction with water Reactions with acids

Reactions with bases Other reactions

At n. c. it is a solid, prone to polymorphism. It exists in an amorphous vitreous state and several crystalline modifications: hexagonal and two orthorhombic. sublimates at 347 °С 563 °С By burning phosphorus in excess of oxygen: 4P + 5O2 → P4O10 (240-400 °С) P4O10 + 6F2 → 4POF3 + 3O2 (100 °С) P4O10 + 6H2O → 4H3PO4 (boiling) P4O10 + 4HNO3(anhydrous) → 4HPO3 + 2N2O5 (0 °С) P4O10 + 4HClO4(anhydrous) → 4HPO3 + 2Cl2O7 (-25 °С, in atmosphere of O3) P4O10 + 3HF → POF3 + 3HPO3 (120-170 °С) P4O10 + 3HCl(gas) → POCl3 + 3HPO3 (200 °С) P4O10 + 3HBr(gas) → POBr3 + 3HPO3 (200 °С) P4O10 + 12NаOH(dil.) → 4Nа3PO4 + 6H2O P4O10 + 6PCl5 → 10POCl3 (150-170 °С) 3P4O10 + 2P4 → 5P4O6 (50 °С) 3P4O10 + 22Li → 10LiPO3 + 4Li3P (300-400 °С) P4O10 + 8H2O2(anhydrous) → 4H3PO2(O2)2 + 2H2O (-20 °С)

Orthophosphoric acid Structure and physical properties Boiling point Melting point Production

At n. c. it is colorless hygroscopic crystals. decomposes at 213 °С 42.35 °С 1) by burning of phosphorus (V) oxide:

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Reaction with water Reactions with acids Reactions with bases

Other reactions

P4O10 + 6H2O → 4H3PO4 (boiling) 2) by hydrolysis of phosphorus (V) chloride: PCl5 + 4H2O → H3PO4 + 5HCl H3PO4 + H2O ↔ H3O+ + H2PO4H3PO4 + HClO4(anhydrous) ↔ P(OH)4+ + ClO4H3PO4 + H2SO4(anhydrous) ↔ P(OH)4+ + HSO4H3PO4(conc.) + NаOH(dil.) → NаH2PO4 + H2O H3PO4(dil.) + 2NаOH(dil.) → Nа2HPO4 + 2H2O H3PO4(dil.) + 3NаOH(conc.) → Nа3PO4 + 3H2O 2H3PO4(dil.) + 3Mg → Mg3(PO4)2 + 3H2 3H3PO4(dil.) + 4Fe → FeHPO4 + Fe3(PO4)2 + 4H2 8H3PO4(conc.) + P4O10 → 6H4P2O7 (80-100 °С) Н3РО4(dil.) + 3АgNO3 → Аg3PO4+ 3HNO3 (addition of NаHCO3)

Phosphorus (V) chloride Structure and physical properties Boiling point Melting point Production

Reaction with water Reactions with acids Reactions with bases Other reactions

In gaseous or liquid states, it has the configuration of a trigonal bipyramid. Solid phosphorus pentachloride consists of [PCl4]+ and [РCl6]- ions. Sublimates at 159-162 °С 166.8 °С By interaction of excess of chlorine on phosphorus (III) chloride: PCl3 + Cl2 ↔ PCl5 (room temperature, Р) PCl5 + 4H2O → H3PO4 + 5HCl PCl5 + H2O(moisture) → POCl3 + 2HCl PCl5 + H2SO4(anhydrous) → HSO3Cl + POCl3 + HCl 3PCl5 + 2H3BO3 → B2O3 + 3POCl3 + 6HCl PCl5 + 8NаOH(dil.) → Nа3PO4 + 5NаCl + 4H2O 6PCl5 + P4 → 10PCl3 (75-100 °С) PCl5 + SO2 → POCl3 + SCl2O (50-60 °С) 2PCl5 + CаSO3 → 2POCl3 + CаCl2 + SCl2O (50-60 °С) PCl5 + 6NаF → Nа[PF6] + 5NаCl (175-230 °С) 2PCl5 + 5CаF2 → 2PF5 + 5CаCl2 (300-400 °С)

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POLONIUM Symbol

Po

History of discovery

In 1898, Marie Skłodowska Curie investigated pitchblende and found that it was much more radioactive than pure uranium, which allowed her to assume the presence of more radioactive isotopes in it. She began a detailed study of pitchblende and revealed two fractions: bismuth salts and barium salts. In the bismuth fraction, there was a product 400 times more active than uranium. Maria Sklodowska-Curie came to the conclusion that this is a new metal and gave it the name "polonium", in honor of Poland, where she was from.

Atomic mass

208.982

Position in the Periodic System

Period: 6, group: 16

Electronic configuration

[Xe]4f145d106s26p4

Electronegativity

1.76

Possible oxidation states

–2, 0, +2, +4, +6

Simple substance Structure and physical properties

At n. c. it is a soft silver-white radioactive metal with a cubic lattice.

Density (at n. c.), kg m-3

9400

Boiling point

962 °С

Melting point

254 °С

Production

Polonium is a natural isotope resulting from the decay of uranium-238, thus, polonium can be extracted from waste from the processing of uranium ores by ion exchange, chromatography, sublimation. Metallic polonium is obtained by thermal decomposition in vacuum of sulphide or polonium dioxide at 500 °C.

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Moreover, in practice, the synthesis of polonium210 in nuclear reactors has become more widespread. Reaction with hydrogen

Po + H2 → H2Po (700 °С)

Reactions with halogens

Ро + Cl2 → РОCl2 (150-200 °С) Ро + 2Cl2 → РОCl4 (300-350 °С)

Reactions with acids

Ро + 6НСl(conc.) → Н2[РоСl6] + 2Н2 Ро + 4H2SО4(conc., hot) → Po(SО4)2 + 2SО2 + 4Н2О Ро + 8HNO3(conc.) → Ро(NО3)4 + 4NО2 + 4Н2О Ро + 2H2SО4(conc.) + 4HI → PoI4 + 2SО2 + 4Н2О Ро + H2SО4(conc.) + H2S → PoS + SО2 + 2Н2О

Reactions with bases

Ро + 2(NаOH·Н2О) → Nа2РоО3 + 2Н2 + Н2О (450 °С)

Other reactions

Ро + 2Н0(Mg, H2SO4(dil.)) → Н2Ро (0 °С) Ро + Pb → PbРо (325-350 °С) 3Ро + 8NаOH(conc.) + 2Аl → 3Nа2Po + 2Nа[Аl(OH)4] Ро + 4NаOH(conc.) + 2Сl2 → РоО(ОН)2 + 4NаCl + + Н2О Ро + 5НСl(dil.) + N2H5Cl → Н2[РоСl4] + 2NH4Cl Ро + 5НСl(dil.) + 2Н2О2 → Н[Ро(Н2О)Сl5] + 3Н2О

Application

‒ in the manufacture of compact and very strong neutron sources, practically do not create γradiation; ‒ at ionization of gases (in particular, air); ‒ in the production of strong and very compact sources of heat for autonomous installations, for example, space.

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RADON Symbol History of discovery

Atomic mass Position in the Periodic System Electronic configuration Possible oxidation states

Rn For the first time, radon was detected by Pierre and Marie Curie, as a gas that is in contact with radium and keeps its radioactivity for a month. Ernest Rutherford and Robert Owens found that the radioactivity of thorium changes with time. Rutherford explained the latest observation by the formation of a new unknown substance (except alpha-particles), which causes the radioactivity of the air around the thorium-containing substances. He suggested naming this substance emanation with the symbol Em. In 1903, the French chemist André-Louis Debierne established a similar substance that forms near actinium-containing compounds. Three types of emanations were initially considered as three different substances: radon, thoron, and actinon. It was later established that they are only different isotopes of one element, which in 1923 was given the name “radon”. 222.018 Period: 6, group: 18 [Xe]4f145d106s26p6 0, +2, +4, +6, +8

Simple substance Structure and physical properties Density (at n. c.), kg m-3 Boiling point Melting point Production

At n. c. it is a colorless inert gas, radioactive (alphaemitter). 9.73 -61.9 °С -71 °С Radon is obtained from an aqueous solution of any salt of radium by blowing air through it. The resulting air is filtered from the microdroplets of the solution and oxygen, water vapor, hydrogen, etc., and with decreasing temperature — nitrogen, argon, neon, etc. are removed from it.

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Other reactions

Application

3Rn + 2ClF3 → 3RnF2 + Cl2 Rn(gas) + 2[O2]+[SbF6]−(solid) → [RnF]+[Sb2F11]−(solid) + + 2O2(gas) ‒ radon baths; ‒ activation of pet food; ‒ indicator in determining the speed of gas flows in blast furnaces, gas pipelines; ‒ search for uranium and thorium deposits, as well as active tectonic faults; ‒ earthquake prediction; ‒ in the study of interaction of groundwater and river water.

102

SELENIUM Symbol History of discovery

Atomic mass Position in the Periodic System Electronic configuration Electronegativity Possible oxidation states

Se Selenium was found during the production of sulphuric acid in Gripsholm by the Swedish chemist Baron Jöns Jacob Berzelius in 1817. He found a similarity between selenium and tellurium, and therefore he gave it a name meaning "moon" (Greek). 79 Period: 4, group: 16 [Аr]3d104s24p4 2.48 -2, 0, +4, +6

Simple substance Structure and physical properties

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

Reaction with hydrogen Reactions with halogens

Selenium in the solid state exists in the form of several allotropic modifications: 1) gray selenium (γ-selenium) with a hexagonal crystal lattice; 2) red crystalline selenium – α, β, γ-selenium – monoclinic modifications; 3) red amorphous selenium; 4) black vitreous selenium. 4820 685 °С 220 °С Selenium is present in the waste of sulphuric acid and copper electrolyte production. The slurry of copper electrolytic plants is treated with sodium hydroxide solution and sulphur dioxide. The process can be represented as the following abbreviated equation: SеО32- + 2SО2 + 2ОН- → 2SО42- + Sе + Н2О The resulting selenium is separated, purified by distillation and dried. Se + H2 → H2Se (350-450 °С) 2Se + 5F2 → SeF6 + SeF4

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

Application

(room temperature) 2Se(suspension) + 5F2 → SeF6 + SeF4 (in liq. CCl4) Se + 2Cl2 → SeCl4 (room temperature) 2Se(suspension) + Br2 → Se2Br2 (in liq. CCl4) 2Se(powder) + 3H2O + 2I2 → H2Se2O3 + 4HI (25-30 °С) Se + O2 →SeO2 (250 °С, ct. is NO2) Se(amorphous) + 2H2O(vapour) → SeO2 + 2H2 (150 °С) 3(α-Se) + 4HNO3(dil., hot) + H2O → 3H2SeO3 + 4NO β-Se + 4HNO3(conc.) → H2SeO3 + 4NO2 + H2O 3Se + 6NаOH(conc.) → Nа2SeO3 + 2Nа2Se + 3H2O (boiling) Se + 2Nа → Nа2Se (-40 °С, in liquid ammonia) Se + 2NаOH(dil.) + 3H2O2(conc.) → Nа2SeO4 + 4H2O 3Se + 4ClF3 → 3SeF4 + 2Cl2 (20 °С, impurities of Se(Cl)F5) 5Se + 6BrF5 → 5SeF6 + 3Br2 (100 °С) ‒ in semiconductor rectifier diodes; ‒ for photovoltaic devices; ‒ for electrophotographic copiers; ‒ as phosphors in television; ‒ in optical and signal devices; ‒ in thermistors; ‒ for discoloration of green glass; ‒ for obtaining ruby glasses; ‒ to make the steel fine-grained structure; ‒ as a catalyst.

Hydrogen selenide Structure and physical properties Boiling point Melting point

At n. c. it is a colorless combustible gas with a foul smell. The molecule has a “curved” structure with a valence angle of 910. -42 °С -64 °С

104

Production

Reactions with chalcogens

Reaction with water Reactions with acids Reactions with bases Other reactions

By interaction of selenides of the active metals with water or dilute acids: Аl2Se3 + 6H2O → 2Аl(OH)3 + 3H2Se FeSe + H2SO4 → FeSO4 + H2Se 2H2Se + 3O2 → 2SeO2 + 2H2O H2Se(saturated) + S → Se + H2S (room temperature) H2Se(dil.) + H2O → HSe- + H3O+ H2Se + 6HNO3(conc.) → H2SeO3 + 6NO2 + 3H2O H2Se + NаOH(dil.) → NаHSe + H2O H2Se + 2NаOH(conc.) → Nа2Se + 2H2O H2Se → H2 + Se (350 °С, impurities of SeCl2) SeCl4 + SeO2 → 2SeCl2O (66 °С, vacuum, in conc. H2SO4) SeCl4 + 2KCl → K2[SeCl6] (0 °С, in conc. HCl)

107

SILICON Symbol History of discovery

Atomic mass Position in the Periodic System Electronic configuration Electronegativity Possible oxidation states

Si Silicon compounds have been known for a very long time, among them is rock crystal, or quartz. Over the years, scientists have tried to isolate elemental silicon. It was possible to isolate free amorphous silicon only in 1823 by the Swedish chemist Baron Jöns Jacob Berzelius when studying compounds of hydrofluoric acid. Crystalline silicon was isolated in 1855 by the French scientist Henri Étienne SainteClaire Deville. 28.086 Period: 3, group: 14 [Ne]3s23p2 1.74 -4, 0, +2; +4

Simple substance Structure and physical properties Density (at n. c.), kg m-3 Boiling point Melting point Production

Reactions with halogens

In amorphous form, it is a brown powder; and in crystalline form it is dark gray, slightly shiny. The crystal lattice of silicon is cubic face-centered. 2000-2400 ~2600 °С 1420 °С 1) by calcination of fine silicon oxide with magnesium: SiO2 + 2Mg → 2MgO + Si (800-900 °С, in atmosphere of Аr) 2) by reduction of the melt of SiO2 by chark: SiO2 + C(chark) → CO2 + Si (1800 °С) Si + 2F2 → SiF4 (room temperature) Si + 2Cl2 → SiCl4 (340-420 °С, in argon current) Si + 2Br2 → SiBr4 (620-700 °С, in argon current)

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Reactions with chalcogens

Reaction with water Reactions with acids

Reactions with bases Other reactions

Application

Si + 2I2 → SiI4 (750-810 °С, in argon current) Si + О2 → SiO2 (1200-1300 °С) Si + S → SiS (650-700 °С, Р) Si + 2S → SiS2 (250-600 °С) Si + 2Se → SiSe2 (800 °С, in atmosphere of Аr) Si + 2Te → SiTe2 (800 °С, in atmosphere of Аr) Si(amorphous) + 2H2O(vapour) → SiO2 + 2H2 (400-500 °С) Si + 6HF(conc.) → H2[SiF6] + 2H2 Si + 4HF(gas) → SiF4 + 2H2 (40-400 °С) Si + 4HI → SiI4 + 2H2 (400-500 °С) Si + 2H2S → SiS2 + 2H2 (1200-1300 °С) 3Si + 18HF(conc.) + 4HNO3(conc.) → 3H2[SiF6] +4NO+ + 8H2O Si + 4NаOH(conc.) → Nа4SiO4 + 2H2 3Si + 18HF(conc.) + 2KClO3 → 3H2[SiF6] + 2KCl + + 6H2O Si + 6HF(conc.) + 2KNO3 → H2[SiF6] + 2KNO2 + + 2H2O 3Si + 2N2 → Si3N4 (1200-1500 °С) 3Si + 4NH3 → Si3N4 + 6H2 (1300-1500 °С) Si + C(graphite) → SiС (1200-1300 °С) Si + Nа → NаSi (when alloying) Si + Cs → CsSi (when alloying) Si + 2Mg → Mg2Si (when alloying) 2Si + Fe → FeSi2 (when alloying) ‒ in production of solar cells; ‒ in microelectronics;

109

‒ in manufacture of crystalline and thin-film photovoltaic cells; ‒ in organic synthesis; ‒ in food industry (E551); ‒ in pharmaceutical industry as an excipient; ‒ in production of rubber as a filler; ‒ in production of glass, ceramics, abrasives, concrete products; ‒ in production of silica refractories; ‒ in chromatography.

Monosilane Structure and physical properties Boiling point Melting point Production

Reaction with chalcogens Reaction with water Reactions with bases Other reactions

At n. c. it is a colorless gas with an unpleasant odor. It starts homologous series of silanes SinH2n+2. -112 °С -185 °С 1) by decomposition of metal silicides with acids: Mg2Si + 4HCl → 2MgCl2 + SiH4 2) by decomposition of triethoxysilane: 4SiH(OC2H5)3 → SiH4 + 3Si(OC2H5)4 (80 °С, in presence of Nа) 3) by interaction of sodium aluminum hydride with silicon tetrachloride: LiАlH4 + SiCl4 → SiH4 + LiCl + АlCl3 SiH4 + 2O2 → SiO2 + 2H2O (150 °С) SiH4 + 2H2O(hot) → SiO2 + 4H2 (ct. is dil. H2SO4) SiH4 + 4NаOH(conc.) → Nа4SiO4 + 4H2 SiH4 → Si + 2H2 (400-1000 °С) 3SiH4 + 8KMnO4 → 8MnO2 + 3SiO2 + 8KOH + + 2H2O

Silicon (IV) oxide Structure and physical properties Boiling point Melting point

At n. c. it is colorless crystals with high hardness and strength. 2600 °С ~1500 °С

110

Production

Reactions with halogens Reactions with acids

Reactions with bases

Other reactions

1) by heating silicon in an oxygen atmosphere: Si + О2 → SiO2 (1200-1300 °С) 2) by interaction of acids and soluble silicates: Nа2SiO3 + 2HCl(dil.) → SiO2 + 2NаCl + H2O SiO2 + 2F2 → SiF4 + O2 (250-400 °С) SiO2 + 6HF(conc.) → H2[SiF6] + 2H2O (450 °С) SiCl4 + 4Nа → Si + 4NаCl (600-700 °С) 3SiCl4 + 16NH3 → Si3N4 + 12NH4Cl (>400 °С, in argon current)

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SULPHUR Symbol History of discovery Atomic mass Position in the Periodic System Electronic configuration Electronegativity Possible oxidation states

S Sulphur refers to the elements used by mankind since ancient times. 32.065 Period: 3, group: 16 [Ne]3s23p4 2.60 -2, -1, 0, +1, +2, +4, +6

Simple substance Structure and physical properties

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

Reaction with hydrogen

Sulphur can form chains of various lengths. The most stable are cyclic S8 molecules, which form the rhombic and monoclinic modifications of sulphur, related to the yellow crystalline sulphur. Shorter cycles (S4, S6) and opening of the chain form brown plastic sulphur. 2070 (rhombic) 444.6 °С (rhombic) 112.8 °С (rhombic) 1) in laboratory: from sodium sulphide or sulphite: 2Nа2S + Nа2SO3 + 6НСl → 6NаСl + 3S + 3Н2О 2Nа2S + Nа2SO3 + 3Н2SO4 → 3Nа2SO4 + 3S + + 3H2О 2) in industry: a) by smelting native sulphur with superheated steam through wells. Molten sulphur is carried to the surface, where it is poured into molds. b) from sulphur-containing compounds: FeS2 → FeS + S (>1170 °С, vacuum) 2H2S(saturated) + O2 → 2S + 2H2O (in the light) 2H2S(gas) + SO2(gas) → 2H2O + 3S (room temperature, ct. is H2O) S + H2 → H2S (150-200 °С)

113

Reactions with halogens

Reactions with chalcogens Reaction with water Reactions with acids

Reactions with bases

Other reactions

S + 3F2 → SF6 2S + Cl2 → S2Cl2 (125-130 °С) S + Cl2 → SCl2 (400 °С) S + 2H2SO4(conc.) → 3SO2 + 2H2O (boiling) S + 6HNO3(conc.) → H2SO4 + 6NO2 + 2H2O (boiling) S + 2HI(gas) → I2 + H2S (500 °С) 4S + 6NаOH(conc.) → Nа2SO3S + 2Nа2S + 3H2O (boiling, impurities of Nа2SO3) 3S + 6KOH(melt) → K2SO3 + 2K2S + 3H2O 4S + 4NH4OH(conc., hot) → (NH4)2SO3S + 2NH4HS + + H2O 5S + 2P → P2S5 2S + Si → SiS2 2S + C → CS2 (700-800 °С, P) S + 2Cl2 + 4NаF → SF4 + 4NаCl (200-300 °С, P) 3S + 2SCl2O2(liq.) → SCl2 + S2Cl2 + 2SO2 (ct. is АlCl3) S + 4CoF3 → 4CoF2 + SF4 (350-400 °С) 3S + 2АgF → Аg2S +S2F2 3S + NF3 → S(N)F + S2F2 (400 °С, vacuum) S + 2Nа → Nа2S (>130 °С) 3S + 2Аl → Аl2S3 (150-200 °С) (n-1)S + Nа2S(solution) → Nа2(Sn) (boiling) S + Nа2SO3(conc.) → Nа2SO3S (boiling) 3S + SO2 → 2S2O

114

(>100 °С, vacuum, electric discharge) 2S + CuO → Cu + S2O (150-200 °С) 10S + 12АgI + 16NH3(liq.) → S4N4 + 6Аg2S + + 12NH4I 8S + 6SO3 + H2SO4(anhydrous) → (S82+)(HS3O10)2 + + SO2 (cold) S8(gas)

S6(gas)

S4(gas)

S2(gas)

S(gas) Application

1) for production of sulphuric acid, elemental sulphur, sulphides; 2) to obtain thiophene and mercaptans; 3) for production of matches, black powder, sparklers, explosives; 4) to combat pests in agriculture and treat diseases; 5) in production of dyes, phosphors; 6) as reagents for precipitating heavy metals; 7) in the composition of natural and artificial hydrogen sulphide baths, as well as in the composition of certain mineral waters.

Sulphur (II) hydride Structure and physical properties Boiling point Melting point Production

Reactions with halogens

Reactions with chalcogens

At n. c. it is a colorless gas with a sweetish taste and the smell of rotten chicken eggs. The molecule is curved and, as a result, polar. −60.8 °С −82.9 °С a) by interaction of dilute acids with sulphides: FeS + 2HCl → H2S + FeCl2 (in the presence of Fe, impurities of H2) b) by interaction of aluminum sulphide with water (it allows us to get the purest hydrogen sulphide): Аl2S3 + 6H2O → 2Аl(OH)3 + 3H2S c) by fusing sulphur with paraffin. H2S + 4Cl2 + 4H2O → H2SO4 + 8HCl H2S(saturated) + Br2 → S + 2HBr H2S(saturated) + I2 → S + 2HI 2H2S + 3O2 → 2H2O + 2SO2

115

Reaction with water Reactions with acids

Reactions with bases

Other reactions

(250-300 °С, air combustion) 2H2S(saturated) + O2 → 2H2O + 2S (in the light) H2S + H2O ↔ HS- + H3O+ HS- + H2O ↔ S2- + H3O+ H2S(saturated) + 2HNO3(conc.) → S + 2NO2 + 2H2O (cold) H2S + 8HNO3(conc.) → H2SO4 + 8NO2 + 4H2O (boiling) H2S + H2SO4(conc.) → S + SO2 + 2H2O (room temperature) H2S + 3H2SO4(conc.) → 4SO2 + 4H2O (boiling) 3H2S + 4HClO3 → 3H2SO4 + 4HCl H2S + HSO3Cl → H2SO3S + HCl (500 °С, in atmosphere of Аr) Te + 2SCl2O → TeCl4 + SO2 + S (>600 °С, in atmosphere of CO2) Te + 3H2O2(conc.) → H6TeO6 (boiling, in dil. H2SO4)

121

Application

Te + 2NаOH(dil.) + 3H2O2(conc.) → Nа2H4TeO6 + + 2H2O Te + 3F2 + 2CsF → Cs2[TeF8] (75 °С) ‒ as an additive to cast iron, steel; ‒ for the manufacture of chemical equipment used in the production of sulphuric acid; ‒ in the technique of temperature measurement; ‒ in production of brown glass and glaze; ‒ as semiconductor materials; ‒ in radiation meters, dosimetric devices, television tubes; ‒ as materials of thermoelectric generators; ‒ as photocells.

Hydrogen telluride Structure and physical properties Boiling point Melting point Production

Reactions with chalcogens

Reaction with water Reactions with acids Reactions with bases

At n. c. it is a colorless, flammable, easily decomposable gas with an unpleasant odor, similar to garlic odor. −2 °С −49 °С 1) by interaction of tellurides with water or acids: Аl2Te3 + 6HCl → 3H2Te + 2АlCl3 This method does not allow us to obtain large outputs, so it is practically not used. 2) by electrolytic method with application of tellurium cathode, platinum anode and sulphuric (or phosphoric) acids as an electrolyte: Te + H2 → H2Te Then the gas is dried, passing through 2 columns containing CaCl2 and P2O5. To separate hydrogen and nitrogen, the gas is passed through a receiver cooled with liquid nitrogen or solid carbon (IV) oxide, as a result of which tellurium crystallizes. 2H2Te + O2(air) → 2H2O + 2Te (room temperature) 2H2Te + 3O2 → 2H2O + 2TeO2 (air combustion) H2Te(dil.) + H2O → HTe- + H3O+ H2Te + 6HNO3(conc.) → TeO2 + 6NO2 + 4H2O H2Te + 2NаOH → Nа2Te + 2H2O

122

Other reactions

H2Te(liq.) → H2, Te, H2Ten (in the light) H2Te(gas) → H2 + Te (>20 °С) H2Te(dil.) → H2 + Te (room temperature, impurities of H2Ten) H2Te + H2O2(dil.) → Te + 2H2O

Tellurium (VI) oxide Structure and physical properties

Production

Reaction with water Reactions with acids

Reactions with bases Other reactions

Tellurium oxide (VI) exists in two modifications: ‒ α-TeO3 – more active bright yellow amorphous substance, paramagnetic; ‒ β-TeO3 – gray crystals of the trigonal syngony, diamagnetic. By decomposition of orthotelluric acid: H6TeO6 → TeO3 + 3H2O (220-320 °С) TeO3(amorphous) + 3H2O(hot) → H6TeO6 TeO3 + 2HCl(conc.) → H2TeO3 + Cl2 (boiling) TeO3 + 5HF(liq.) → H[TeOF5](solid) + 2H2O (room temperature) TeO3(amorphous) + 2NаOH(conc.) + H2O → Nа2H4TeO6 TeO3 + 4HF(liq.) + LiF → Li[TeOF5] + 2H2O (room temperature) TeO3 + 3Аg2O → Аg6TeO6 (200 °С)

Tellurous acid Structure and physical properties Melting point

At n. c. it is colorless crystals, insoluble in water.

Production

1) by interaction of dilute acid and sodium tellurite: Nа2TeO3 + 2HNO3 → H2TeO3 + 2NаNO3 (0 °С) 2) by reduction of tellurium (VI) oxide: TeO3 + 2HCl → H2TeO3 + Cl2 (100 °С)

decomposes at 40 °С

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3) by reduction of sodium orthotellurate: Nа6TeO6 + 8HCl → H2TeO3 + Cl2 + 6NаCl + 3H2O Reactions with acids

Reactions with bases Other reactions

H2TeO3 + 4HCl(conc.) → Te + 2Cl2 + 3H2O 2H2TeO3 + HNO3(conc., cold) → Te2(NO3)O3(OH) + + 2H2O H2TeO3 + 2NаOH(conc.) → Nа2TeO3 + 2H2O H2TeO3 → TeO2 + H2O (>40 °С) H2TeO3 + N2H4·H2O → Te(colloid) + N2 + 4H2O (30-40 °С) H2TeO3 + 2SO2 + H2O → Te + 2H2SO4 5H2TeO3 + 6HNO 3 +7H 2O +2KMnO4 → 5H6TeO6 + + 2Mn(NO3)2 + 2KNO3 H2TeO3 + 2NаOH(conc.) + H2O2(conc.) → Nа2H4TeO6 + + H2O H2TeO3 + Cа(NO3)2 → CаTeO3 + 2HNO3

Tellurium (IV) chloride Structure and physical properties Boiling point Melting point Production

Reaction with water Reactions with acids Reactions with bases Other reactions

At n. c. it is light yellow crystals consisting of tetramers. 390 °С 224 °С 1) by interaction of chlorine and tellurium: Te + 2Cl2 → TeCl4 (100 °С) 2) by interaction of thionyl chloride and tellurium in an inert atmosphere: Te + 2SOCl2 → TeCl4 + SO2 + S (600 °С) TeCl4 + 2H2O → TeO2 + 4HCl TeCl4 + H2O(air moisture) → TeCl2O + 2HCl TeCl4 + H2S(liq.) → TeCl2 + 2HCl + S (-77 °С) TeCl4 + 6NаOH(dil.) → Nа2TeO3 + 4NаCl + 3H2O TeCl4 + 2KCl(conc.) → K2[TeCl6] TeCl4(liq.) + АlCl3 → (TeCl3)[АlCl4]

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THALLIUM Symbol History of discovery

Atomic mass Position in the Periodic System Electronic configuration Electronegativity Possible oxidation states

Tl The discovery of thallium was promoted by the green color of the flame during its combustion. In March 1861, an English scientist, Sir William Crookes, used a spectral analysis method in the study of dust, which was captured in one of the sulphuric acid productions, and found a new line of light green color in the spectrum that did not belong to any of the known elements. Due to this, thallium was discovered and named in Latin “thallus”, i.e. “blossoming branch”. 204.383 Period: 6, group: 13 [Xe]6s25d104f146p1 1.44 0, +1, +2, +3

Simple substance Structure and physical properties Density (at n. c.), kg m-3 Boiling point Melting point Production

Reactions with halogens

Reactions with chalcogens

At n. c. it is a silvery-white, plastic, very soft, fusible metal. On the air, it is covered by an oxide film. 11850 1472 °С 303.5 °С Tl2O + H2 → 2Tl + H2O (>500 °С) Tl2O+ CO → 2Tl + CO2 (250-325 °С) 2TlCl + H2 → 2Tl + 2HCl (650–750 °С) 2Tl + Cl2 → 2TlCl (room temperature) 2Tl + 2HCl(conc.) + 3Cl2 → 2H[TlCl4] 4Tl + 2O2 → Tl2O + Tl2O3 (400 °С, air combustion) 2Tl + S → Tl2S

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Reaction with water Reactions with acids

Other reactions Application

(320 °С, in atmosphere of H2) 2Tl + 3S → Tl2S3 (200-250 °С) 4Tl + 2H2O + O2 → 4TlOH (50-70 °С) 2Tl + H2SO4(dil., cold) → Tl2SO4 + H2 3Tl + 4HNO3(dil., hot) → 3TlNO3 + NO + 2H2O Tl + 6HNO3(conc., hot) → Tl(NO3)3 + 3NO2 + 3H2O 2Tl + 3H2O2(conc.) → Tl2O3 + 3H2O ‒ in semiconductor technology, electronic and electrical industry; ‒ for doping germanium, silicon, cadmium compounds in order to give them acceptor properties; ‒ for obtaining photoresistances, photocells with high sensitivity, making phototriodes; ‒ in the devices of infrared technology; ‒ in atomic technology of various types of scintillation counters, to activate luminescent alkali halogen crystals, to stabilize the luminescence process; ‒ in the production of insecticides and zoocides; ‒ in optical elements for IR equipment, acoustooptics, laser technology; ‒ as anti-knock fuel in internal-combustion engines; ‒ in the chemical industry as catalysts for the oxidation of hydrocarbons and olefins, polymerization and oxidation; ‒ for the production of dyes, luminous paints, artificial gems, pearls and diamonds.

Thallium (III) oxide Structure and physical properties

Boiling point Melting point Production

At n. c. it is a dark brown amorphous powder or cubic crystals. It melts without decomposition, decomposes when calcined. It shows amphoteric properties. It is a strong oxidizer. decomposes at 875 °С 717 °С 4Tl + 2O2 → Tl2O+ Tl2O3 (400 °С, air combustion) 2Tl + 3H2O2(conc.) → Tl2O3 + 3H2O 4TlNO3 → Tl2O3 + Tl2O+ 4NO2 (250–350 °С)

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Reaction with hydrogen Reactions with halogens Reaction with water Reactions with acids

Reactions with bases

Other reactions

2Tl(NO3)3(dil.) + 3H2O → Tl2O3 + 6HNO3 2Tl(NO3)3 + 6NаOH(dil.) → Tl2O3 + 6NаNO3 + 3H2O 2TlCl3(dil.)+ 3H2O(hot) → Tl2O3 + 6HCl 2TlCl3 + 6NаOH(dil.) → Tl2O3 + 6NаCl + 3H2O Tl2O3 + 2H2 → Tl2O + 2H2O (150-185 °С) 2Tl2O3 + 6F2 → 4TlF3 + 3O2 (>550 °С) Tl2O3·nH2O → Tl2O3 + nH2O (>300 °С, vacuum) Tl2O3 + 8HCl(conc.) → 2H[TlCl2] + 2Cl2 +3H2O Tl2O3 + 4H2SO4(conc.) → 2H[Tl(SO4)2] or Tl(HSO4)SO4 +3H2O Tl2O3 + 6HNO3(conc.) → 2Tl(NO3)3 + 3H2O Tl2O3 + 2LiOH → 2LiTlO2 + H2O (450-575 °С) Tl2O3 + 2NаOH → 2NаTlO2 + H2O (450-575 °С) Tl2O3 + 2KOH → 2KTlO2 + H2O (450-575 °С) Tl2O3 + 2RbOH → 2RbTlO2 + H2O (450-575 °С) Tl2O3 → Tl2O+ O2 (500-1000 °С) 2Tl2O3 + 5S(powder) → 2Tl2S + 3SO2 (room temperature) Tl2O3 + H2O2 → 2TlOH + O2 + H2O

Thallium (I) hydroxide Structure and physical properties Melting point Production

Reaction with water Reactions with acids Reactions with bases

At n. c. it is light yellow when melting decomposes. It is well dissolved in water, forms an alkaline solution. It shows strong base properties. decomposes at 139 °С 4Tl + 2H2O + O2 → 4TlOH Tl2O + H2O → 2TlOH Tl2SO4 + Bа(OH)2 → 2TlOH + BаSO4 TlOH(dil.) + 6H2O → [Tl(H2O)6]+ + OH− TlOH + HCl(dil.) → TlCl +H2O TlOH + HNO3(dil.) → TlNO3 + H2O 2TlOH(conc.) + Zn(OH)2 → Tl2[Zn(OH)4]

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Other reactions

2TlOH + O2 → Tl2O3 + H2O (200 °С) 2TlOH + 4NаOH + 2Cl2 → Tl2O3 + 4NаCl + 3H2O TlOH(dil.) + CO2 → TlHCO3 2TlOH + CO2 → Tl2CO3 + H2O

Thallium (III) chloride Structure and physical properties

Melting point Production Reaction with water

Reactions with acids Reactions with bases Other reactions

At n. c. it is white, decomposes when heated. In acidified concentrated solution, it is stable; in a diluted solution, it completely decomposes. It is an oxidizing agent. 25 °С TlCl(suspension) + Cl2 → TlCl3 TlCl3(conc.) + 3H2O(cold) → [Tl(H2O)3Cl3] (in dil. HCl) 2TlCl3(dil.) + 3H2O(hot) → Tl2O3 + 6HCl TlCl3 + HCl(conc.) → H[TlCl4] 2TlCl3 + 6NаOH(dil.) → Tl2O3 + 6NаCl + 3H2O 2TlCl3 + 3H2S → Tl2S + 2S + 6HCl TlCl3 + 3KCl(conc.) → K3[TlCl6] 2TlCl3 + 3CsCl → Cs3[Tl2Cl9] TlCl3(conc.) + 2Tl(powder) → 3TlCl

Thallium (III) nitrate Structure and physical properties

Production

Reaction with water Reactions with acids Reactions with bases Other reactions

At n. c. it is white (in the form of crystalline hydrate), partially decomposes in air at room temperature, completely when heated. It is stable in acidified concentrated solution, hydrolyzed in dilute solution. It is a strong oxidizer. Tl + 6HNO3(conc., hot) → Tl(NO3)3 + 3NO2 + 3H2O Tl2O3 + 6HNO3(conc.) → 2Tl(NO3)3 + 3H2O TlNO3 + 4HNO3(conc.) → Tl(NO3)3 + 2NO2 + 2H2O 2Tl(NO3)3(dil.) + 3H2O → Tl2O3 + 6HNO3 Tl(NO3)3 + 3HCl(conc.) → TlCl + Cl2 + 3HNO3 2Tl(NO3)3 + 6NаOH(dil.) → Tl2O3 + 6NаNO3 + 3H2O 2Tl(NO3)3 + 3H2S → Tl2S +2S + 6HNO3 Tl(NO3)3 + H2O + K2SO3 → TlNO3 + K2SO4 + + 2HNO3

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Thallium (I) sulphate Structure and physical properties

At n. c. it is a white solid, which is volatile (above the melting point), thermally stable.

Melting point

632 °С

Production

Tl2S+ 2O2 → Tl2SO4 (250 °С) 2Tl + H2SO4(dil., cold) → Tl2SO4 + H2

Reaction with water

Tl2SO4(dil.) + 2nH2O → 2[Tl(H2O)n]+ + SO42−

Reactions with acids

Tl2SO4 + 2HCl(conc.) → 2TlCl + H2SO4 Tl2SO4 + H2SO4(conc.) → 2TlHSO4 (20-40 °С)

Reactions with bases

Tl2SO4 + Bа(OH)2 → 2TlOH + BаSO4

Thallium (I) carbonate Structure and physical properties Boiling point Melting point Production Reaction with water Reactions with acids Other reactions

At n. c. it is white, melts without decomposition, decomposes on subsequent heating. decomposes at 360 °С 273 °С 2TlOH + CO2 → Tl2CO3 + H2O Tl2CO3(dil.) + 2nH2O → 2[Tl(H2O)n]+ + CO32Tl2CO3 + 2HCl(dil.) → 2TlCl + H2O + CO2 Tl2CO3(dil.) + H2O + CO2 → 2TlHCO3 Tl2CO3 → Tl2O + CO2 (300–600 °С)

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TIN Symbol

Sn

History of discovery

Tin belongs to the elements known to mankind since ancient times, approximately from the 4th millennium BC. It was part of tin bronze, invented in the late or mid-III millennium BC. For a long time tin was a strategic element, especially due to its low accessibility and high cost.

Atomic mass

118.710

Position in the Periodic System

Period: 5, group: 14

Electronic configuration

[Kr]4d105s25p2

Electronegativity

1.72

Possible oxidation states

0, +2, +4

Simple substance Structure and physical properties

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

Reactions with halogens

At n. c. it is a plastic, malleable and fusible shiny metal of silver-white color. Tin forms three allotropic modifications: ‒ α-form (gray tin) with a cubic diamond type lattice; ‒ β-form (white tin) with a tetragonal crystal lattice; ‒ γ-form (white tin) – metallic, brittle. 6520-6560 (white tin) 2720 °С (white tin) 231.91 °С (white tin) Production of tin in industry consists of several stages: 1) enrichment by the method of gravitational flotation or magnetic separation; 2) roasting the concentrate in oxygen to remove sulphur and arsenic impurities; 3) reduction by coal or aluminum (zinc) in electric furnaces: SnO2 + C → Sn + CO2 (800-900 °С) Sn + 2F2 → SnF4 (