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J. Rohonczy: Inorganic Chemistry I

J. Rohonczy: Inorganic Chemistry I

Dr. János Rohonczy

Lecture Notes

Eötvös Loránd University, Budapest Faculty of Sciences

Dr. János Rohonczy

INORGANIC CHEMISTRY I.

Lecture Notes

Eötvös Loránd University

Faculty of Sciences

BUDAPEST 2017. János Rohonczy: I.

Lecture Notes. Copyright © 2017 Dr. János Rohonczy, Eötvös Loránd University, Budapest, Faculty of Sciences

All Right are Reserved. No part of this publication may be reproduced, stored in a retrieval system or transmitted in any form or by any means: electronic, electrostatic, magnetic tape, mechanical, photographical, photocopying, recording or otherwise, without permission in writing form the publisher.

This book is written utilized the lecture notes of the Inorganic Chemistry lectures of the author at the Department of Inorganic Chemistry of Eötvös Loránd University, Budapest.

Revised, and the fullerene and cluster topics remarked by Dr. Béla Csákvári professor emeritus.

First edition 2017

Edited and cover page made by Dr. János Rohonczy

Publisher: Eötvös Loránd University, Faculty of Sciences ISBN: 978-963-284-853-2 DOI: 10.21862/

3

Table of Contents

Introduction 7

1. 8 1.1. Hydrogen compounds 9

2. Halogens: F, Cl, Br, I, At 10 2.1. Hydrogen halides 13 2.2. Interhalogens 14 2.3. Polyhalogen and interhalogen , organic derivatives 16

3. (16th column) O, S, Se, Te, Po 17 3.1. (O) 17 3.1.1. 19 3.1.2. Halogen and oxygen halides 21 3.1.3. Halogen oxoacids and their salts 24 3.1.4. Halogen oxofluorides and fluorinated oxoacids 28 3.2. (S) 29 3.2.1. Sulfur containing compounds 31 3.2.2. Sulfur – Sulfanes 32 3.2.3. Sulfur halides 32 3.2.4. Sulfur oxohalides 33 3.2.5. Sulfur //iodides 33 3.2.6. Sulfur oxohalides 34 3.2.7. Sulfur oxides 34 3.2.8. Sulfur oxoacids 35 3.3. (Se), (Te), (Po) 39 3.3.1. Preparation, produce and application 39 3.3.2. Se, Te, Po Polycations 41 3.3.3. Se, Te, Po Hydrides 41 3.3.4. Se, Te, Po Halides 42 3.3.5. Oxohalides 42 3.3.6. Oxides 42 3.3.7. – Oxoacids 43

4. (15-th column) N, P, As, Sb, Bi 44 4.1. (N) 44 4.1.1. , azides and nitrido compounds 45 4.1.2. Nitrogen hydrides 45 4.1.3. Nitrogen halides 47 4.1.4. Nitrogen oxides and oxoacids 48 4.1.5. Nitrogen oxoacids 49 4.1.6. Sulfur nitrides 50 4.2. (P) 51 4.2.1. Phosphorus hydrides 52 4.2.2. Phosphorus halides 52 4.2.3. Phosphorus oxides / / oxosulfides, oxoacids 53 4.2.4. Phosphorus nitrides and phosphorus organic compounds 55 János Rohonczy: Inorganic Chemistry I.

4.3. (As), (Sb), (Bi) 56 4.3.1. Arsenic, antimony, bismuth hydrides 58 4.3.2. Arsenic, antimony, bismuth halides 58 4.3.3. Arsenic, antimony, bismuth oxides and sulfides 59 4.3.4. Elemento organic compounds 60

5. (14-th column) C, Si, Ge, Sn, Pb 61 5.1. (C) 61 5.1.1. Carbides 64 5.1.2. and carbon halides 65 5.1.3. Carbon oxides, carbonic 66 5.1.4. Carbon sulfides 67 5.1.5. Carbon nitrides 68 5.2. 69 5.2.1. Binary compounds 70 5.2.2. – Hydrosilicons 71 5.2.3. Silicon halides 71 5.2.4. Silicones - silicon organic compounds 72 5.2.5. Silicon oxides, silicic , silicates 73 5.2.6. Silicon–sulfur compounds 75 5.2.7. Silicon – nitrogen compounds 76 5.3. (Ge), (Sn), (Pb) 77 5.3.1. Ge/Sn/Pb hydrides, and hydrido halides 79 5.3.2. Ge/Sn/Pb halides 79 5.3.3. Ge/Sn/Pb oxides, and hydroxides 80 5.3.4. Ge/Sn/Pb oxoacid salts 81 5.3.5. Ge/Sn/Pb 81 5.3.6. Anions with cluster skeleton 81 5.3.7. Ge/Sn/Pb organic compounds 81

6. (13th column) B, Al, Ga, In, Tl 83 6.1. Boron (B) 83 6.1.1. Borides 84 Metal rich borides 84 6.1.2. Boron halides 85 6.1.3. Boron , boric acids, borates 86 6.1.4. Boron nitrogen and other boron containing compounds 87 6.1.5. Other boron compounds 87 6.1.6. 88

6.2 (Al), (Ga), (In), Tallium(Tl) 91 6.2.1. Al/Ga/In/Tl-hydrides 93 6.2.2. Al/Ga/In/Tl halides 93 6.2.3. Al/Ga/In/Tl oxides, hydroxides and complicated oxides 94 6.2.4. Ternary and more complicated oxides 95 6.2.5. Al/Ga/In/Tl chalcogenides 95 6.2.6. Binary III-V compounds of Al/Ga/In/Tl 95 6.2.7. Al/Ga/In organic compounds 96 6.2.8. Al / Ga / In / Tl and nitrogen bounded compounds 96 6.2.9. Aluminum salts 97

7. Be, Mg, Ca, Sr, Ba, Ra – Alkaline Earth Metals 98 7.1. Elements 98 7.1.1. Preparation and application 99 7.1.2. Chemical properties 100

5 János Rohonczy: Inorganic Chemistry I.

7.2. Compounds 101 7.2.1. Be, Mg and hydrides 101 7.2.2. Be, Mg and alkaline earth metal halides 101 7.2.3. Be, Mg and alkaline earth metal oxides and hydroxides 101 7.2.4. Be, Mg and alkaline earth metal carbonates 102 7.2.5. Be, Mg and alkaline earth metal , 102 7.2.6. Be, Mg and alkaline earth metal salts 102 7.2.7. Be, Mg and alkaline earth metal complexes 103 7.2.8. Be, Mg and alkaline earth metal organometallic compounds 103

8. (1st column) Li, Na, K, Rb, Cs, Fr – Alkali Metals 104 8.1. Elements 104 8.1.1. Discovery 104 8.1.2. Natural appearance 104 8.1.3. Preparation 104 8.1.4. Application 105 8.2. compounds 106 8.2.1. Alkali metal hydrides 106 8.2.2. Alkali metal halides, pseudo halides, oxohalides 106 8.2.3. Alkali metal oxides 106 8.2.4. Alkali metal hydroxides 107 8.2.5. Other alkali metal salts and complexes 107 8.2.6. Alkali metal organic compounds 109

9. (18th column) He, Ne, Ar, Kr, Xe, Rn – Noble Gases 110 9.1. Elementary properties 110 9.2. compounds 111 9.2.1. Xe fluorides, reactions 111 9.2.2. Xe organic compounds, Xe-Si/N bonded compounds 113 9.2.4. Chemistry of and 113

References 114

6 János Rohonczy: Inorganic Chemistry I. Introduction

Introduction This book was written for supporting the lectures of B.Sc. students for chemistry and material science in inorganic chemistry at Eötvös Loránd University Faculty of Natural Sciences Department for Inorganic Chemistry. Inorganic chemistry contains the introduction of all elements forming our nature and the wide characterization the compounds thereof. The elements are discussed in groups in order of their position in the . Following the traditions at Eötvös Loránd University the discussion of hydrogen is followed by the elements of p-block starting with halogens and then going backwards column by column. The reason of this treatment is that the elements with high electonegativity often form molecules with covalent bonds thus we are moving from the simpler to the more complicated structures of materials. When speaking about an element the history, the mineral appearance, the production and the main application will be discussed. As for the physical characteristics their state, thermal and electric conductivity, main isotopes and allotropes will be shown. The introduction of molecule structure is based on the Lewis structure and VSEPR theory. The chemical nature of elements is characterized by the oxidation number as well as the reactions with , acids and bases. This is followed by the systemic discussion of binary and ternary compounds, showing their electron configuration as well as the production, typical and the important reactions in practice. The relations between the chemical structure and the chemical reactivity are often discussed as well. Molecules with multiple carbon are just slightly mentions in inorganic chemistry. On the other hand the element–carbon bonds are often discussed as being the bases of element-organic compounds getting more and more important on the field of organic syntheses and chemical industry. Due to the fast development of semiconductor based electronic industry in the 1970s the search for inorganic compounds having special electric, optical, thermal, mechanic or magnetic properties has increased. The improvement of the high tech industry gave a huge jump-start to create new materials, thus initializing the material science separated from inorganic chemistry. When discussing the elements we will mention the roll of their and their compounds’ in the high-tech technologies, their applications. At the end of this book I call your attention on some textbooks in Hungarian and in English as well as some relevant articles in Wikipedia. Finally, I wish to express my acknowledgement to my old professor Dr. Béla Csákvári for vetting the manuscript and his notes to cluster chemistry.

Budapest, 2017 The Author

János Rohonczy: Inorganic Chemistry I. 1. Hydrogen

1. Hydrogen shell: 1s1 Physical properties: Colorless gas, low , low viscosity Occurrence in Universe: Stars, Sun, interstellar gas clouds, giant planets (Jupiter etc.) Occurrence on Earth: Water, hydrocarbons, water Isotopes: Protium, deuterium, tritium Allotropes: Ortho-H2 and para-H2 (with parallel and anti parallel nuclear spins) Industrial preparation: 1) 3 Fe + 4 H2O(vapor) → Fe3O4 + 4 H2 (900 °C) Lavoisier, 1783-1900. 2) C + H2O(vapor) → CO + H2 (1000 °C) till ca. 1945. CO + H2O → CO2 + H2 3) CH4 + H2O(vapor) → CO + 3 H2 (400 °C) today 4) Electrolysis of water (KOH ), very pure. Red cylinder, 150-200 Bar. Laboratory preparation: 1) Zn + 2 HCl → ZnCl2 + H2 (Kipp apparatus) 2) CaH2 + 2 H2O → Ca(OH)2 + 2 H2 3) Electrolysis of Na2SO4 solution. Pure H2 gas on cathode (negative pole), O2 on anode. Cleaning: Diffusion through (Pd) foil. Reactivity: H2 + F2 → 2 HF cold gas, explosion in darkness, initial step: F2 →2 F⋅ H2 + Cl2 → 2 HCl explosion in blue light H2 + Br2 → 2 HBr heating is needed H2 + I2 → 2 HI equilibrium at elevated temperature. Synthesis of HI: see reactions of PI3.

2 H2 + O2 → 2 H2O (radical reaction) Initial step: H2 → 2 H⋅ ∆H = 255 kJ/mol, high energy required, e.g. flame, spark.

Complicated chain reaction: ∙H + O2 → ∙OH + :O ∙OH + H2 → H2O + ∙H :O + H2 → ∙OH + ∙H Recombination steps: ∙H + ∙OH → H2O ∙H + ∙H → H2 ∙OH + ∙OH → H2O2

Detection of H2O2 : H2 flame is cooled on ice rock: molten water contains H2O2. 2+ 2+ TiO + H2O2 → H2O + TiO2 (intensive yellow color) H2 + x S → H2Sx (ca. 600 °C) (Thermal dissociation of S–S bonds at higher temperature) 3 H2 + N2 → 2 NH3 (Haber-Bosch synthesis) equilibrium: Le Chatelier's proinciple. János Rohonczy: Inorganic Chemistry I.

Reactions with metals: If EN () of the metal is low (alkali metals, etc.), then ionic hydrides are formed: – Ca + H2 → CaH2 (- with H ). This case H2 gas is oxidizer. If the EN of metal is high (and electrode potential is positive), H2 gas is reducing agent: Cu2O + H2 → 2 Cu + H2O WO3 + 3 H2 → W + 3 H2O

1.1. Hydrogen compounds a) Covalent hydrides: (p-group, clumns 14-17). Hydrides of 14th column elements are neutral, hydrides of 15th column elements are Lewis-bases, hydrides of 16th groups elements are weak acids, hydrides of 17th group elements are acids. Bond strength, thermal stability and basicity is decreasing, is increasing in top to down direction.

Table 1. Name and composition of covalent hydrides. 14th column – s2p2 15th column – s2p3 16th column – s2p4 17th column – s2p5 – CH4 – NH3 water – H2O – HF – SiH4 – PH3 hydrogen – H2S hydrogen – HCl – GeH4 – AsH3 – H2Se hydrogen – HBr stannane – SnH4 stibine – SbH3 hydrogen – H2Te – HI Autoprotolysis + – -14 2 H2O → H3O + OH K = 10 , neutral pH = 7 + – -30 2 NH3 → NH4 + NH2 K = 10 , neutral pH = 15 b) Polymer hydrides: BmHn (m∼25), CmHn (m>40), SimH2m+2(m≤8), GemH2m+2(m≤5), BeH2 (polymer as well) c) Ionic hydrides: low EN of metals is required, e.g. LiH (NaCl lattice), NaH, KH, RbH, CsH, CaH2, SrH2, BaH2, UH3. Electrolysis of molten LiH results H2 gas on anode(+). d) Metallic (interstitial) (with elements of d- and f-fields). Hydrogen atoms in the interatomic holes of the metal atoms. Metallic hydrides are not fully stoichiometric compounds Stoichiometric compunds: 3-rd column: ScH2 / YH2 / LaH2 4-th column: TiH2 / ZrH2 / HfH2 5-th column: VH2 / VH / NbH / NbH2 / TaH 6-th column: CrH 10-th column: (NiH) / PdHx (x<1) 11-th column: CuH 12-th column: ZnH2 Lanthanoids: CeH2 / PrH2 Actinoids: ThH2 / UH2 / NpH2 e) Complex hydrides: LiBH4 / LiAlH4 / NaBH4 / Al(BH4)3 2– 2– 2– 2– 2– [PtH2] / [PtH4] / [RhH4] / [RuH6] / [ReH9] (last one is the ion with highes known coordination number).

9 János Rohonczy: Inorganic Chemistry I. 2. Halogens

2. Halogens: F, Cl, Br, I, At Valence shell: ns2np5 Table 2. Atomic properties of the halogens. Halogen 1. 2. 3. 4. 5. 6. radius (Å) Covalent radius (Å) EN F 2 7 0.57 0.72 3.98 Cl 2 8 7 0.97 0.99 3.16 Br 2 8 18 7 1.12 1.14 2.96 I 2 8 18 18 7 1.32 1.33 2.66 At 2 8 18 32 18 7 1.43 1.45 2.2

Table 3. Physical, chemical properties of covalent X2 molecules.

X2 MP(°C) BP (°C) State 1st ionization energy (eV) Oxidation number F2 -220 -188 Gas 17.4 -1 Cl2 -101 -35 Gas 13.0 ±1, +3, +5,+7 Br2 -7 59 Liquid 11.8 ±1, +5 I2 114 184 10.5 ±1, +5, +7 At 302 337 Solid/metal 9.5 ±1, +3, +5, +7

Table 4. Discovery data of the halogens.

X2 Year Discover Naming F2 1886 Moissan Fluoros Cl2 1774 Scheele Khloros Br2 1826 Balard Bromos I2 1804 Courtois Iodes At 1940 Corson, MacKenzie,Segre Astatos 19 35/37 79/81 127 206 Important stable nuclides: F, Cl, Br, I, At (t1/2 = 30 min)

IUPAC names + – Na Cl – chloride, HCl – , NaClO3 – sodium[trioxo chlorate] KClO4 – [tetraoxo-chlorate], [Al(H2O)6]Cl3 – [hexaaqua-aluminium(III)] trichloride

General characterisation

Fluorine Occurrence: In oceans 1.2 ppm, fluorite CaF2, cryolite Na3AlF6, fluor-apatite Ca5(PO4)3F, topaz Al2SiO4(OH,F)2 Biological impact: Toxic: 2-3 ppm F2, or 150 mg NaF, but if < 1 ppm, non-toxic (in toothpaste) Physical properties: light (greeny) yellow gas. Occurrence: in oceans 3.5%, -mines. History: NaCl – romans, HCl / HNO3 aqua regia, cc. HCl – Glauber 1648, Cl2 – Scheele. Physical properties: light yellow gas. Occurrence: in oceans Cl:Br = 300:1, AgBr bromirite (Mexico), mineral water. János Rohonczy: Inorganic Chemistry I. 2. Halogens

History: crimson(snail): 6,6-dibrom-indigo – bible, mineral water in Montpellier: MgBr2 + Cl2 → MgCl2 + Br2 (1826, Balard). Physical properties: braun liquid.

Iodine Occurrence: salinous mineral USA, Japan (100 ppm!), lantarite Ca(IO3)2 (Chile) Physical properties: black , sublimates: violet vapor History: ash of alga + cc. H2SO4: violet vapor (1811, Courtois)

Astacium Physical properties: Radioactive, short life time, very low amount in Earth's crust (max 44 mg). 209 4 211 1 Synthesis: Bi + He → At + 2 n, t1/2 =7.2 hours

Preparation, application

Fluorine Electrolysis: KF:HF = 2:3 (steel cathode, graphite anode,72 °C, 10-6000 A, 8-12 V) Theoretically interesting: K2MnF6 + 2 SbF5 → 2 KSbF6 + MnF3 + ½ F2. Application: most important (70-80%) preparation of UF6. Other products: SF6, ClF3, BrF3, IF5, WF6, ReF6. Inorganic and organic fluorides are made by reactive fluorides (e.g. SbF5, CoF3 etc.) instead of F2 gas. Chlorine Industrial preparation: electrolysis of NaCl solution. Diaphragm: asbestos, Nafion- membrane, or Hg cathode technology. Laboratory preparation: 4 HCl(cc) + MnO2 → MnCl2 + 2 H2O + Cl2 16 HCl(cc) +2 KMnO4 → 2 KCl + 2 MnCl2 + 8 H2O + 5 Cl2

Application: 70%: chlorinated organic moleclues (e.g. CH2=CH2 + Cl2 → CH2Cl–CH2Cl (etilene-dichloride); 20%: whitening, sterilization (e.g. paper, textil, swimming pool, drinking water); 10%: inorganic compounds (e.g. HCl, Cl2O, HOCl, NaClO3, AlCl3, SiCl4, SnCl4, PCl3, PCl5, POCl3, AsCl3, SbCl3, SbCl5, S2Cl2, SCl2, SOCl2, ClF3, ICl, ICl3, TiCl4, MoCl5, FeCl3, ZnCl2, Hg2Cl2, HgCl2) Bromine Industrial preparation: see/mineral water with high Br– concentration: – – 2 Br + Cl2 + air → Br2(vapor) + 2 Cl (at pH = 3.5)

Application: pharmaceutical industry: CH3Br fungicide; CH2BrCH2Br (earlier in leaded petrol), C3H5Br2Cl etc. fire-roof materia; AgBr photography; paint industry: HBr, KBr, KBrO3. Industrial preparation methods: 1) Iodide containing mineral water (Japan): – – 2 I + Cl2 + air → I2(vapor) + 2 Cl – – 2) I + AgNO3 → AgI + NO3 2 AgI + Fe → 2 Ag + FeI2

11 János Rohonczy: Inorganic Chemistry I. 2. Halogens

FeI2 + Cl2 → FeCl2 + I2 Ag recycling: Ag + 2 HNO3 → AgNO3 + NO2 + H2O

– – – 3) 3 I + Cl2 → 2 Cl + I3 , – I3 ions are trapped on ion exchanging resins. Elution with NaOH solution. Regeneration of resin with NaCl solution.

4) From Chile salpetre (NaIO3 contamination): – – – 2– + IO3 + 3 HSO3 → I + 3 SO4 + 3 H – – + IO3 + 5 I + 6 H → 3 I2 + 3 H2O

Application: 50%: organic iodine containing compounds, 35%: catalist in rubber industry, pigments, inks, photo industry,15%: I2, KI. Special: K2HgI4 (Nessler reagent – detection of NH3), Mayer-reagens (detection of alkaloids), Cu2HgI4 thermocolor pigment, Ag2HgI4 (most ionic solid at 20 °C).

General reactivity

Fluorine Reactivity Most reactive element: reacts wit all other elements except: He, Ne, Ar. Oxidation number in all compounds: -1 F2(f) + O2(f) → O2F2 (-196 °C, 3 MeV γ-iradiation) F2(g) + O2(g) → O2F2 (720 Hgmm, silent discharge) Passivating of Al, Fe, Ni, Cu, Mg (solid protecting fluoride layer on the surface) Fine powered metals: Ag + F2 → AgF2 Reaction with noble metals: F2 + Xe → XeF2, further more XeF4, XeF6. Extremely strong oxidizer: highest oxidation numbers can be reached: IF7, PtF6, PuF6, BiF5, TbF4, CmF7, KAgF4, AgF2 – F-bridge containing structures: [As2F11] (coordination number = 2) MgF2, MnF2 (coordination number = 3), CaF2, SrF2, PbF2 (coord number = 4), NaF, CsF (coord number = 6).

Chlorine, bromine, iodine Decreasing reactivity: Cl2 > Br2 > I2 Exampes: Cl2 + CO → COCl2 (phosgene) Cl2 + 2 NO → 2 NOCl (nitrosyl-chloride) Cl2 + SO2 → SO2Cl2 (sulphuryl-chloride) But Br2 / I2 + CO / NO / SO2 → NO REACTION.

Increasing ligandum size and decreasing ionization potential → decreasing number of : Re + Cl2 → ReCl6, Re + Br2 → ReBr5, Re + I2 → ReI4

Physical solution of X2: Good solvents: EtOH, Et2O, CS2 , EtBr, CHCl3, hexane, etc.

+ – : F2 + H2O → 1/2 O2 + 2 H + 2 F , other products: O3, H2O2, HOF. – E0(F2/F ) = +2.866 V and E0(1/2 O2/H2O) = + 1.229 V – + Cl2 + H2O → HOCl + Cl + H , at acidic/neutral pH.

12 János Rohonczy: Inorganic Chemistry I. 2. Halogens

Br2 and I2: similar reactions. – – – Cl2 + 2 OH → OCl + Cl + H2O, in solution. Br2 and I2: similar reactions. – – – Over room temperature: 3 OCl → ClO3 + 2 Cl (disproportion) Gas hydrates – clathrates: in icy water Cl2⋅8 H2O / Br2⋅10 H2O. X2 molecules in the holes of the crystal lattice (yellow crystal). Solvatation: I2 is violet in non-polar organic solvents (CCl4) , is red-brown in benzene, I2 is dark brown in or . – I2 is low in water, but the solubility is much higher in the presence of I ions: – – I2+ I ↔ I3 , tri-iodide ion is formed.

2.1. Hydrogen halides

HX form molecules in gas phase. Complete dissociation appears in water, they are strong acids. Preparation, application

H2F2 Preparation: a) from fluorite: CaF2 + cc. H2SO4 → CaSO4 + H2F2 (200 °C) SiO2 contamination reacts with H2F2 as well SiO2 + 2 H2F2 → SiF4 + 2 H2O SiF4 + H2F2 → H2SiF6

b) from fluor-apatite: Ca5(PO4)3F + cc. H2SO4 → CaSO4 + CaHPO4 + H2F2 (byproduct during the preaparation of phosphate containing artifical fertilizer).

Application: Preparation of freons (e.g. CCl2F2, CCl3F) and Teflon. Synthesis of Na3AlF6 (kriolite). Other compounds: UF4, UF6, NaF, SnF2, HBF4, H2SiF6. 2+ + Biological influence: HF: dehidration, broken equilibrium of Ca /K ions (insoluble CaF2). In the case of skin bite: 15 minutes washing with water, MgSO4 pulp, Ca-gluconate injection into the skin. HCl Preparation: Leblanc method: NaCl + H2SO4 → NaHSO4 + HCl (150 °C) NaCl + NaHSO4 → Na2SO4 + HCl (500 °C)

Hasgreaves method: 4 NaCl + 2 SO2 + O2 + 2 H2O → 2 Na2SO4 + 4 HCl (450 °C)

Direct synthesis, high purity: H2 + Cl2 → 2 HCl

Byproduct of organic : CH2Cl–CH2Cl → CH2=CHCl + HCl (500 °C) Application: HCl + SiC → SiCl4 HCl + NH3 → NH4Cl MxNy + HCl → MClx MO + 2 HCl → MCl2 (M = Ti, Zr, Hf, Nb, Ta, Cr, Mo, W) Al + 3 HCl → AlCl3 + 3/2 H2 2 HCl + NaClO3 → ClO2 + 1/2 Cl2 + NaCl + H2O (catalyst: Ti/Mn) Use of aqeous HCl solution: as cheap strong acid, eching of rust, synthesis of PVC, etc.

13 János Rohonczy: Inorganic Chemistry I. 2. Halogens

HBr / HI Preparation: NaBr + H3PO4 → NaH2PO4 + HBr H2 + Br2 → 2 HBr 2 I2 + N2H4 → 4 HI + N2 I2 + H2S → 2 HI + S

Laboratory synthesis: P(red) + H2O + I2 → HI + H3PO3 Special syntheses: Tetrahydro-naphthalene + Br2 → tetrabromo-naphthalene + 4 HBr HBr(aq.) + P4O10 → HBr(sicc) 3 D2O + PBr3 → 3 DBr + D3PO3 Use: HBr – inorganic bromides, alkyl-bromides, HBr small/big cylinders. HI: small cylinders for laboratory use only. HX physical / chemical behaviours Physical properties: (HF)x colorless liquid with low viscosity. BP = 19.5 °C, two dimensional H-bridge structure. HCl (BP = -84 °C), HBr (BP = -67 °C), HI (BP = -35 °C). Colorless gases. Azeotropic mixtures with water: Diluted solution: mostly H2O evaporates. Concentrated solution: mostly HX evaporates. Table 5. Physical properties of azeotropic HX acids.

HX/H2O azeotropes HF HCl HBr HI BP (°C) 112 109 124 127 Conc. (g HX/100 g solution) 38 20 48 57 Density (g/cm3) 1.14 1.1 1.5 1.7 Protolytic dissotiation. Strength of acidity: HF << HCl < HBr < HI.

+ – (HF)x autoprotolysis: 3 HF � H2F + HF2

2.2. Interhalogens

Binary compounds of halogen atoms. Types: neutral molecules, cations, anions, covalent organic derivatives.

Neutral molecules

Stable products, composition: XY, XY3, XY5, XY7 (X – central atom has the higher atomic number).

XY Structure: linear AXE3

Table 6. Physical properties of interhalogens with XY constitution. Interhalogen ClF BrF IF BrCl ICl IBr Color Colorless Yellow Instable Red-brown Red Black crystal gas gas gas crystal MP (°C) – < -23 – < -66 ca. 20 41

14 János Rohonczy: Inorganic Chemistry I. 2. Halogens

Preparation: Cl2 + F2 → 2 ClF (225 °C), Br2 + BrF3 → 3 BrF (high temperature) I2 + AgF → IF + AgI Reaction: W + 6 ClF → WF6 + 3 Cl2 SO2 + ClF → Cl–SO2–F SF4 + ClF → SF5Cl H2O + 2 ClF → 2 HF + Cl2O + – BF3 + 2 ClF → [Cl2F] [BF4] + – Electric conductivity: 3 ICl � I2Cl + ICl2

XY3 Structure: AX3E2, T-shape molecules.

Table 7. Physical properties of interhalogens with XY3 constitution.

Interhalogene ClF3 BrF3 IF3 I2Cl6 Color Colorless liquid Light yellow Yellow crystal Intensive yellow crystal liquid MP (°C) -76 9 101(16 bar) BP (°C) 12 126 -28 (decomposes) –

Preparation. Cl2 + 3 F2 → 2 ClF3 (direct synthesis) I2 + 3 XeF2 → 2 IF3 + 3 Xe Reaction: U(solid) + ClF3(liquid) → UF6(gas) + 3 ClF(gas) (70 °C) + – – AsF5 + ClF3 → [ClF2] [AsF6] (F donation) AgCl + ClF3 → AgF2 + 1/2 Cl2 + ClF (cross fluorination)

XYZ2 IFCl2, IF2Cl.

XY5 19 Structure: IF5: tetragonal pyramid, AX5E at 20 °C ( F NMR signals in 4:1 ratio) At 115 °C: fast molecular re-arrangement, Berry-type pseudo-rotation: only one 19F NMR signal.

Table 8. Physical properties of interhalogens with XY5 constitution.

Halogen pentafluorides ClF5 BrF5 IF5 MP (°C) -103 -60 +9 BP (°C) -13 +4 +105

Preparation: KBr + 3 F2 → KF(solid) + BrF5 (at 25 °C) Reactions: ClF5 + 2 H2O → FClO2 + 4 HF + – ClF5 + AsF5 → [ClF4] [AsF6] BrF5 + 3 H2O → HBrO3 + 5 HF (!) + – : 2 IF5 � IF4 + IF6 (little dissociation of IF5)

XY7 Structure: close to penthagonal bipyramid, AX7 (see VSEPR) IF7 synthesis: I2(g) + 7 F2 → 2 IF7 (300 °C) PdI2 + 8 F2 → PdF2 + 2 IF7 (PdI2 is not air sensitive) Properties: IF7 strong fluorinating agent ( sublimation at 4.8 °C): 2 IF7 + SiO2 → 2 IOF5 + SiF4

15 János Rohonczy: Inorganic Chemistry I. 2. Halogens

2.3. Polyhalogen and interhalogen ions, organic derivatives

Polyhalogen cations

+ + + + XY2 Structure: V-shape, AX2E2. Homonuclear: e.g. I3 , Br3 . Heteronuclear: e.g. ICl2 . + + Paramagnetic, e.g. I2 , or diamagnetic e.g. ICl2 . + – Preparation: Br2 + BrF + AsF5 → Br3 ⋅AsF6

Polyhalogen anions

– – – – ClF2 , Cl3 , BrF2 , IF2 – – – – – Structure: linear, AX2E3. Several I3 analogons: Cl3 , BrF2 , IF2 , ClF2 . – – – Known ternary ions: IBrCl , IBrCl3 , six coordinated: IF6 . + – Synthesis example: IF5 + CsF → Cs IF6 .

Covalent organic interhalogen derivatives Known R(aryl)–XFn organic interhalogen compounds, where X = I or Br, n = 2, or 4. There are R(alkyl, aryl)XF2, and R(aryl)XF derivatives as well. Typical synthesis: low temperature flourination of alkyl-X or aryl-X compounds (where X = Br and I).

16 János Rohonczy: Inorganic Chemistry I. 3.1. Oxygen

3. (16th column) O, S, Se, Te, Po Valence shell: ns2np4 elements. They react with metals and non-metals.

Table 9. Physical properties of chalogen elements. Name MP BP E.N. 1st ioniz. pot. Therm. conduct. El. conduct. Metal conduct. (°C) (°C) (kJmol-1) (W/cm K) (1/cm Ω) ? O – oxygen -219 -183 3.44 1.3140 2⋅10-4 – No S – sulfur 119 445 2.58 0.9996 2⋅10-3 5⋅10-18 No Se – selenium 217 685 2.55 0.9409 2⋅10-2 1⋅10-6 semiconductor Te – tellurium 449 990 2.10 0.8693 2⋅10-2 2⋅100 semiconductor Po – polonium 254 962 2.0 0.812 2⋅10-1⋅ 2⋅104 Yes

3.1. Oxygen (O)

Valence shell: 1s22s2np4 History: 15th century. Leonardo da Vinci: air contains a component which feeds the burning. Discovery: 1774 Pristley (England) heating of HgO → O2, 1774 Scheele (Sweden) heating of KNO3 / HgO / Mg(NO3)2 → O2. Name: 1777 Lavoisier (France): oxus gennan – "forms acids". Allotropy: O2 (dioxygen) and O3()

Isotopes, abundance, application

Table 10. Data of oxygen isotopes. 15O 16O 17O 18O -2 -1 t1/2 = 2 min. 99.8% 3.10 % 2.10 % I = 2.5 NMR Stable, tracing Tracing of photosynthesis: * * H2O + CO2 → O2 + [CH2O]n (sugar). Occurrence: 21% in air, 86% in oceans, 45.5% in lithosphere (silicates, carbonates, phosphates, etc.).

Atomic O Non stable, synthesis: N2O/O2/NO2 → O:⋅ (UV light) * Detection: O: + NO2 → O2 + NO, O: + NO → NO2 → NO2 + hν (yellow-green light) Reaction: O: + O2 → O3, or 3 O: + CH4 → CO2 + H2O

O2 Properties: colorless, odorless, tasteless gas. Two electronic states: 3 – 3 Common: triplet ( Σg ) O2 contains even number of electrons, but the two electrons of the highest occupied level are with parallel spins on two different – degenerated - π* orbitals → paramagnetic gas. 1 1 Excited singlet ( ∆g) O2 (colroless gas, blue liquid at -183 °C), diamagnetic, bigger bond length, higher energy, higher reactivity.

János Rohonczy: Inorganic Chemistry I. 3.1. Oxygen

3 Laboratory synthesis of common O2: a) 2 KMnO4 + 5 H2O2 + 3 H2SO4 → 5 O2 + 2 MnSO4 + K2SO4 + 8 H2O b) K2Cr2O7 + 3 H2O2 + 4 H2SO4 → 3 O2 + Cr2(SO4)3 + K2SO4 + 7 H2O – – c) KOH solution → 4 OH - 4e → O2 + 2 H2O (on anode) d) 2 H2O2 → 2 H2O + O2 (Ni / Pt / MnO2 catalysts) e) Oxoacid salts: 2 KClO3 → 2 KCl + 3 O2 (400 °C or 200 °C with MnO2) 2 KMnO4 → K2MnO4 + MnO2 + O2 (200 °C) 3 Industrial synthesis of O2: fractional distillation of liquid air (-183 °C), Note: N2 is more volatile than O2.

Reactivity: General properties: very reactive, reacts with many elements, decomposition of O2 molecule is endothermic, but after that exothermic reactions: burning, explosion. Reacts with C / H2 / metals / inorganic / organic compounds. There is no direct reaction, but stable oxides: W, Pt, Au. There are no oxides of He, Ne, Ar.

Application: 100Mt/year: Bessemer steel, glass melting furnace, oxide, TiO2, rocket fuel, medical/biological application. Blue cylinder, 150-200 Bar (oil contamination can cause explosion). 1 Synthesis of O2: 3 1 a) 2 O2 + hν → 2 O2 (630 nm, both molecules are excited) 3 1 1 3 b) O2 + Sens + hν → O2 + Sens (Sens = sensitizer, e.g. fluorescein) – – 1 c) H2O2 + OCl → Cl + H2O + O2 (in solution) 1 Reaction: H2C =CH2 + O2 → 2 H2C=O (synthesis, polymer oxidizer, air chemistry)

O3 (ozone) Properties: blue, diamagnetic, toxic gas, characteristic ozone odor (copy machines produce). Name: Greek ozein = smell. Dark blue liquid at -112 °C, violet-black crystals at -193 °C. Thermodynamically instable gas, but the decomposition into O2 is slow at 200 °C. (UV light or heavy metal catalysts speed up the decomposition). Liquid ozone is brisant: 2 O3 → 3 O2. Ozone layer absorbs the UV light. Solubility: hydrocarbons, freon, CO, F2, etc. V-shape molecule. Synthesis: with ozonizer – metal plated glass tube, 10-20 kV, silent discharge, 25 °C, 1 Bar. Other methods: + – * a) O2 + O2 → O3 + O or O2 + O2 → O3 + O (ca. 10%) b) O2 → O3 (UV light, low concentration) c) H2SO4 → H2S2O8 + O2 + O3 (-10 °C, electrolysis) d) F2 + H2O → H2F2 + O3 (with other by products)

Concentration measurement of O3 – – O3 + 2 I + H2O → O2 + I2 + 2 OH (buffer, I2 measurement: titration with Na2S2O3)

Reaction: Strong oxidizer: F2 > F2O > O3 > – – O3 + CN → OCN + O2 O3 + 2 NO2 → Ο2 + N2O5 + – – O3 + 2 H + 3 I → I3 + H2O + O2 + + 3+ O3 + CO2 + 2 H → 2 Co + O2 + H2O

– O3 (ozonides) Structure: V-shape. Properties: KO3 red-brown, paramagnetic powder. Most stable salts: CsO3 / Ba(O3)2.

18 János Rohonczy: Inorganic Chemistry I. 3.1. Oxygen

Synthesis: 5 O3 + 2 KOH → 2 KO3 + 5 O2 + H2O (-10 °C, powder) Reaction: Forms potassium by heating: KO3 → KO2 + 1/2 O2 Hydrolysis: 4 KO3 + 2 H2O → 4 KOH + 5 O2 Application: organic syntheses, sterilization.

EtCH CH2 Et–CH=CH2 + O3 → → Et–CHO + HCHO. O3

Table 11. Oxidation states of oxygen. oxidation number -2 -1 -1/2 -1/3 0 +1/2 2– 2– – – + ion O O2 O2 O3 O3 / O2 O2 name oxide superoxide allotropes oxygen cation + – example Li2O Na2O2 KO2 KO3 O3 / O2 [O2] [PtF6] Table 12. Coordination numbers of oxygen. 1 2 3 3 4 6 8 + + CO H2O [O(HgCl)3] H3O AgO CaO Li2O linear V-shape triangle planar pyramid tetrahedral octahedral anti-fluorite

3.1.1. Oxygen compounds

O2 as ligand. Loosely bonded O2

a) O2 + � O2·hemoglobin

b) [Ir(CO)Cl(PPH3)2] + O2 � [Ir(CO)Cl(O2)(PPH3)2] (Vaska-complex, 1963) 8 Structure: complexes with coordination number 4: square planar, 16 electrons, AX4, ks-d , D4h symmetry. Complexes with coordination number 6: octahedral, 18 electron, bidental O2: Oh symmetry, π-bonds.

Binding of O2 to central metal atoms: superoxo / monodental: M–O=O, superoxo / bidental: M–O=O–M, (M = Rh, Co), peroxo / bidental: M=O2 or M=O2=M, (M = Ir, La). – – 2– Reactive oxygen: MLnO2 + NO → NO2 / + NO2 → NO3 / + SO2 → SO4

Strongly bonded O2, as ion. 2– Peroxo anion, O2 M–O–O–M, e.g. Na2O2; M=O2, (M = Fe, Co, Rh); – Superoxo anion, O2 K–O=O, + + – cation, O2 O2 + PtF6 → [O2] [PtF6] , ionic lattice, O2 is in oxidized state!

Binary compounds with hydrogen

H2O (dihydrogen oxide, water) Occurrence: oceans, lakes, rivers, air, minerals. Water purification: physical / chemical / biological. Sedimentation / coagulation: Fe(OH)3, Al(OH)3. 2+ 2+ Water softening: remove of Mg -, Ca -ions. Disinfection of water with Cl2, O3.

Physical properties: H2O, D2O, T2O (see hydrogen), ice: at least 8 crystal forms. Common structural properties : four O atoms surround a central one, tetrahedral arrangement, O– H→O bond. + 2+ Crystal water: with cations H2O→M , e.g.: [Ni(H2O)6] , with oxoanions: H-bond in solid 2– phase: CuSO4⋅5 H2O (water is bonded to SO4 ).

19 János Rohonczy: Inorganic Chemistry I. 3.1. Oxygen

Water in zeolite: water molecules are in the cavities. Clathrates: 12-46 water molecules around the guests: H2S / Ar / Kr / CH4 / Cl2 / Br2 etc. Chemical properties: Polar solvent with big dielectric constant (ε = 80), Good solvent for: salts/anhydrides and polar covalently bonded molecules: 3+ – AlCl3 + H2O → [Al(H2O)6] + 3 Cl (solution of ionic compound), Further chemical reactions during the solution of covalent molecules: dissociation/hydrolysis: P4O10 + H2O → 4 H3PO4 + – Na2S + H2O → H2S + 2 Na + 2 OH (similar covalent hydrides: PH3, SiH4) + – Acid/base reaction: 2 H2O � H3O + OH (autoprotolysis, pH = 7 at 20 °C) + + + Oxonium ions: H3O , in solid phase other ones as well: H5O2 , [H(H2O)n] , n = 1, 4, 6. – – Hydrogen-bond with OH : only in HCl⋅2 H2O: [H3O2] Quantitative measurement of water by Karl Fischer method (1935). 2 H2O + I2 + SO2 + Py / MeOH (anhydrous) → 2 HI⋅Py + H2SO4 (Py = pyridine)

H2O2 (dihydrogen-dioxide, hydrogen-peroxide, 1818) Synthesis: a) Ba + O2 → BaO2 then BaO2 + H2SO4 → BaSO4 + H2O2 b) by electrochemical oxidation: – – 2 HSO4 - 2 e → HO3SOOSO3H (peroxodisulfuric acid) – + H2S2O8 + 4 H2O → 2 HSO4 + 2 H3O + H2O2

Laboratory synthesis: K2S2O8 + 2 D2O → 2 KDSO4 + D2O2.

Industrial synthesis: H2 + ethyl anthraquinone + O2 → H2O2 +ethyl anthraquinone (0.5 Mt/year):

O OH

alkohol/CH H + 2 Raney Ni

O OH

+ O2 / viz Etil-antrakinon H2O2 / viz

Structure: Dihedral angle: 111° in gas phase, 90° in solid phase, 129° in H2O2⋅H2O, 180° in NH4F⋅H2O2. Physical properties: viscous, colorless liquid, MP = 0 °C, BP = 150 °C, dielectric constant = 70, electric conductivity: like water. Chemical properties: Oxidation number = +1. Stronger acid, than the water.

Decomposition: 2 H2O2 → 2 H2O + O2, inhibitor: . Heterogeneous catalyst: Pt, Ag, MnO2. 4– + 3– Oxidation agent: 2 [Fe(CN)6] + H2O2 + 2 H → 2 [Fe(CN)6] + 2 H2O (yellow → red) 2+ 3+ 2– 2– Fe → Fe , SO3 → SO4 NH2OH → HNO3 (hydroxylamine)

Reduction agent: (both O atoms of O2 come from peroxide {indicated by isotope tracing}): – + 2+ MnO4 + H2O2 + H → Mn + H2O + O2 4+ 3+ + Ce + H2O2 → Ce + 2 H + O2

20 János Rohonczy: Inorganic Chemistry I. 3.1. Oxygen

+ – 1 HOCl + H2O2 → H3O + Cl + O2 (singlet O2, light emission) – – 1 Cl2 + H2O2 + OH → Cl + H2O + 1/2 O2 (light emission) + – 2– Acid/base reaction: (H2OOH) , (OOH) , O2 ions: + + (H2OOH) + H2O � H2O2 + H3O + – H2O2 + Na → Na + OOH + 1/2 H2 + – H2O2 + NH3(f) → NH4 + OOH (white solid, MP = 25 °C) + – + + H2O2 + HF + AsF5 → [H3O2] [AsF6] (and [H3O2] → ½ O2 + H3O ) 2– Combined anion: S2O8 peroxodisulfate with S–O–O–S bonds.

Other oxides Most elements form oxides except the light noble gases (He, Ne, Ar).

MP / BP: Wide thermal range: CO (BP = -192 °C) → ZrO2 (MP = 3265 °C).

Electric conductivity: MgO (insulator) → NiO (semiconductor) → ReO3 (el. conductor).

Thermal stability: Al2O3 (stable, exothermic) → SiO2 → H2O → Cl2O (instable, endothermic). Acid/base character: CO2/SO3 (acid) → BeO/Al2O3/Bi2O3/ZnO (amphoteric) → CO/NO (inert) → Li2O/CaO/La2O3 (base). Elements in periodic system: from left to right: bases → acids. From up to down: increasing acidity.

Structure: CO/OsO4/Sb2O3/P4O10 (molecules) → HgO/SeO2/CrO3 (chain polymers) → SnO/AS2O3/Re2O7 (layers) → SiO2/MgO (3D lattice).

Stoichiometry: CO/H2O (stoichiometric) → UO2+x (0

3.1.2. Halogen oxides and oxygen halides

Fluorine has bigger EN than the oxygen: only the O can be central atom.

Oxygen fluorides:

OF2 () Structure: V-shape. Properties: light yellow (toxic) gas over 145 °C. Less reactive in liquid/solid phase: mixtures are stable with H2, CH4 or CO at RT, but sparks cause explosion.

Synthesis: F2 + NaOH (2% solution)→ OF2 + NaF + H2O (20 °C), or electrolysis of HF/KF . – – Reaction: hydrolysis in bases: OF2 + OH → O2 + F + H2O fluorination agent: OF2 + P → PF5 + POF3 OF2 + S → SO2 + SF4

21 János Rohonczy: Inorganic Chemistry I. 3.1. Oxygen

O2F2 (dioxygen difluoride) Structure: like H2O2, long F–O, short O–O bonds.

Synthesis: F2(gas) + O2(gas) → O2F2 (silent discharge, 10 mBar) F2(gas) + O2(gas) → O2F2 (-200 °C / γ-ray) Properties: yellow liquid below -57 °C / solid below -150 °C. Radical dissociation: O2F2 → F⋅ + ⋅O–O–F Strong oxidation and fluorination agent: H2S + 4 H2O2 → SF6 + 2 HF + 4 O2

Further, exotic known compounds: O2F (dioxygen fluoride). Synthesis: O2 + F → O2F O4F2 (tetraoxygen difluoride). (stable below -183 °C only).

Oxygen chlorides

Cl2O () Physical properties: brawn-yellow gas / red-brawn liquid. Structure: Cl–O–Cl, V shape. Synthesis: endothermic compound, Cl2 + O2 do not react. Both in industry and in labor (since 1834): 2 HgO(yellow) + 2 Cl2 → HgCl2·HgO + Cl2O(gas) Chemical properties: water soluble. Heating or spark: explosion. Major reaction: Cl2O + H2O ↔ 2 HOCl (HOCl is not stable) Cl2O + NH3 → N2 + NH4Cl + H2O

Chlorine oxides: Cl2O3 / ClO2 / Cl2O4 / Cl2O6 / Cl2O7

ClO2 () (Since 1811.) Structure: O=Cl=O, V-shape. Physical properties: yellow gas, dark red liquid/solid, paramagnetic.

Preparation: 3 KClO3 + 3 H2SO4(cc) → 2 ClO2 + HClO4 + H2O + 3 KHSO4 – 2– + 2 ClO3 + C2O4 + 4 H → 2 ClO2 + 2 CO2 + 2 H2O – – + Industrial synthesis: ClO3 + Cl + 2 H → ClO2 + 1/2 Cl2 + H2O (cheap, but Cl2 by-product) – 2– 2 ClO3 + SO2 → 2 ClO2 + SO4 (better than the previous) Chemical properties: Explosion over -40 °C and 60 mBar. Strong oxidizer agent, reacts with metals, e.g.: 2 Cl2O + Mg → Mg(ClO2)2 Dissociation in bases by disproportion: – – – 2 Cl2O + 2 OH → ClO2 + ClO3 + H2O

Hydrolysis in darkness: ClO2⋅(6-10) H2O (dark green solution). At lighting: ClO2 → ClO + 1/2 O2 ClO + H2O → H2ClO2 (+ClO) → HCl + HClO3. Application: whitening of paper, cellulose and textile. Produce: 0.1 Mt/year in USA.

Cl2O7 () Discovery: 1900. Structure: O3Cl–O–ClO3 Physical properties: colorless oil, it is possible to distillate in vacuum.

Synthesis: 2 HClO4 → Cl2O7 + H2O (cc. H3PO4, -10 °C, -H2O).

22 János Rohonczy: Inorganic Chemistry I. 3.1. Oxygen

Chemical properties: hydrolysis: Cl2O7 + H2O ↔ 2 HClO4 (reversible) Thermal decomposition: Cl2O7 → ClO3 + ClO4 (explosion)

Other exotic compounds Cl2O3 (dichlorine trioxide) (1967) Synthesis/structure: 2 ClO2(solid) → Cl–O–ClO2 + ½ O2 (hν, -78 °C) Physical properties: dark brawn liquid, explosion over 0 °C.

Cl2O4 (dichlorine tetroxide) (1970). Synthesis: CsClO4 + ClOSO2F → CsSO3F + Cl–O–ClO3 Structure: Cl–O–ClO3 V-shape. Physical properties: light yellow liquid Chemical properties: decomposition at 20 °C: Cl–OClO3 → Cl2 + O2 + ClO2 + Cl2O6

Cl2O6 (dichlorine hexoxide) (1843). Synthesis: 2 ClO2 + 2 O3 → Cl2O6 + 2 O2 () + – Structure: O3Cl–ClO3, or O2Cl:O2:ClO2, or [ClO2] [ClO4] () Physical properties: dark red liquid / yellow solid Chemical properties: 2 Cl2O6 ↔ 2 ClO3(decomposition) → 2 ClO2 + O2 Cl2O6 + H2O → HO–ClO2 + HClO4 ( + )

Bromine oxides

Br2O (dibromine monoxide) Structure: Br–O–Br, V-shape. Physical properties: dark brawn liquid. MP = -17.5 °C

Synthesis: 2 HgO(yellow) + 2 Br2(vapor) → HgBr2·HgO + Br2O, (similar to Cl2O) 2 BrO2 → Br2O + 3/2 O2 (low pressure, heating)

Chemical properties: 5 Br2O + 6 I2 → I2O5 + 10 IBr – – – 6 Br2O + 6 OH → 5 BrO3 + Br + 3 H2O (in base)

BrO2 (bromine dioxide)

Synthesis: Br2 + 4 O3 → 2 BrO2 + 4 O2 (in freon at -78 °C, ozonolysis) Structure: O–Br–O, V-shape Physical properties: light yellow crystal, explosion over -40 °C. – – – Chemical properties: hydrolysis, 6 BrO2 + 6 OH → 5 BrO3 + Br + 3 H2O.

Iodine oxides

Stable ones: I4O9 / I2O4 / I2O5. Unknown: I2O.

I2O5 (iodine pentoxide) Discovery: 1813. Structure: O2I–O–IO2 Physical properties: most stable white crystal, water soluble, hygroscopic. Synthesis: 2 HIO3 → I2O5 + H2O (by heating) Chemical properties: I2O5 → I2 + 5/2 O2 (300 °C) I2O5 + CO → CO2 + I2 (fast reaction at RT) + 2– I2O5 + SO3 → [IO2]2 [SO4] (iodyl sulphate) Other reactions: I2O5 + H2O → HI3O8 (I2O5·HIO3) HI3O8 + H2O → 3 HIO3

23 János Rohonczy: Inorganic Chemistry I. 3.1. Oxygen

III V HIO3 → I4O9 + H2O, (dehydration with cc. H3PO4) white, hygroscopic product: I (I O3)3 + – HIO3 → I2O4 + H2O (dehydration with cc. H2SO4) lemon-yellow crystal [IO] [IO3]

3.1.3. Halogen oxoacids and their salts

Fluoro oxoacids: EN of F2 ≈ 4, E0 = +2.866 V, so only one fluoro oxoacid is known.

HOF () Discovery: 1968. Physical properties: white solid below MP = -117 °C, light yellow liquid.

Laboratory synthesis: F2 gas on ice surface at -40 °C, by-products: HF, H2O Synthesis, industrial: F2 + H2O → HOF + HF (frozen into solid N2)

Reaction: HOF + H2O → HF, O2, H2O2 (slow decomposition) HOF + Ag+ → Ag2+, – – HOF + BrO3 → BrO4 (like elementary F2 in water) HOF + F2 → OF2 + HF Salts: Unknown. Covalent derivatives: F–O–NO2, (BP = -50 °C), KNO3 + F2 → F–O–NO2 + KF

Chloro, bromo and iodo oxoacids

Table 13. Selected data of halogen oxoacids. Oxidation number chlorine acids bromine acids iodine acids anions +1 HOCl* HOBr* HOI* hypohalogenite +3 HOClO* ? – halogenite * * +5 HOClO2 HOBrO2 HOIO2 halogenate * +7 HOClO3 HOBrO3 HOIO3 perhalogenate (HO)5IO H4I2O7 *stable only in water solution Usually potential strongly depends on pH: – + – BrO3 + 6 H3O + 5e ↔ 1/2 Br2 + 9 H2O (E0 = +1.495 V) – – – BrO3 + 3 H2O + 5e ↔ 1/2 Br2 + 6 OH (E0 = +0.519 V)

Increasing the pH: E0 decreases, oxidation ability decreases.

– – – Disproportion properties: a) 3 XO ↔ 2 X + XO3 (fast at T > 70 °C-on, X = Cl, Br, I) – – – b) 4 ClO3 ↔ Cl + 3 ClO4 (still slow at T ≈ 100 °C)

HOX (hypohalic acids, hypohalites) + – Properties: weak acids, thermally instable: 2 HOX → 2 H + 2 X + O2 Both the HOCl and OCl– anions are strong oxidizer agents + – Synthesis of acid: X2 + H2O → HOX + H + X , (add HgO/Ag2O to push the equilibrium to AgX) – Industrial synthesis of HOCl: Cl2O + H2O → 2 HOCl, (0 °C, Cl free method)

– – – – Synthesis of OX : a) X2 + 2 OH → X + OX + 2 H2O (cold solution) b) halogen exchange: X– + OCl– → OX– + Cl– (X = Br– , I–)

24 János Rohonczy: Inorganic Chemistry I. 3.1. Oxygen

Table 14. Differences in oxidation strength of HOCl acid and OCl– ion.

Reaction partner with HOCl at pH<7 with OCl– in base NH3 NCl3 NH2Cl – – – Br Br2 OBr / BrO3 – – – I I2 OI / IO3 H2O2 O2 − – – NO2 NO3 − 2– S SO4 − – – ClO2 − ClO3 CN– − OCN– 2– 2– SO3 − SO4 2+ – Mn − MnO4

HOBr + R–NH2 → N2 (quantitative reaction) – – 2– – 3 OBr + (H2N)2CO(urea) + 2 OH → N2 + CO3 + 3 Br + 3 H2O Application: 1) Bleach = NaOCl, LiOCl, Ca(OCl)2⋅CaCl2⋅2H2O (oxidation agents, bleaching, disinfection compounds) 2) halogenation strength: OI– > OBr– > OCl– – – – R–COCH3 + 3 OBr → RCO2 + 2 OH + CHBr3 (bromoform test), 3) industrial synthesis: NH3 + NaOCl → N2H4 + NaCl + H2O (water/gelatin inhibitor) 4) synthesis of α-glycols: HOCl + H2C=CH2 → H2C(OH)–CH2Cl {+ H2O} → H2C(OH)–CH2OH + HCl

Produce: NaOCl 0.2 Mt/year, Ca(OCl)2 0.1 Mt/year, LiOCl 5000 t/year

HXO2 (halogenous acids / halogenites) − Known: HClO2 and ClO2 salts. Structure: H–O–Cl=O: Properties: very instable, only in water solution.

HClO2 (). Oxidation number of Cl = +3. Synthesis: 1) 2 ClO2 + H2O → HClO2 + HClO3 2) Ba(OH)2 + ClO2 + H2O2 → Ba(ClO2)2 + 2 H2O + O2 (H2O2 reduces) Ba(ClO2)2 + H2SO4(diluted) → BaSO4 + 2 HClO2 (medium strong acid) Decomposition (depends on the conditions): – + a) 5 HClO2 → 4 ClO2 + Cl + H + 2 H2O – – + b) 3 HClO2 → 2 ClO3 + Cl + 3 H – + c) HClO2 → Cl + H + O2

+ 2+ + 2+ 2+ + M(ClO2)x salts: M: Ag , Pb , Hg , Ba , Sr , Na Properties: crystalline salt, explosion by heating. Most stable is: NaClO2.

Synthesis: a) 2 NaOH + ClO2 + H2O2 (or Na2O2) → 2 NaClO2 + 2 H2O + O2 2+ b) Ba + ClO2 + H2O → Ba(ClO2)2 + Ba(ClO3)2

Application: NaClO2: 20,000 t/year. Bleaching, production of ClO2, oxidation of H2S, HCN, RSH, R2S, RCHO in industrial smoke.

HXO3 (halogenic acids, halogenates) Oxidation number of halogen atom: +5

25 János Rohonczy: Inorganic Chemistry I. 3.1. Oxygen

HClO3 (chloric acid) Synthesis from salts: Ba(ClO3)2 + H2SO4 → BaSO4 + 2 HClO3. Properties: water soluble, stable till 30%, decomposition by heating, e.g.: 3 HClO3 → HClO4 + H2O + 2 ClO2 (byproducts: Cl2, O2)

HBrO3 (), similar to HClO3, decomposition by heating: products Br2 + O2.

HIO3 () Synthesis: I2 + cc. HNO3 → HIO3 + NO2 + H2O. Properties: colorless solid (MP = 200 °C), very stable even in anhydrous state, by heating: – – –1 2 HIO3 → I2O5 + H2O. In water: IO3 + HIO3 � [H(IO3)2] , stable dimer (K = 41 mol ).

− XO3 (halogenate salts) Structure: AX3E, pyramid shape. Synthesis: − − − a) 3 X2 + 6 OH → XO3 + 5 X + 3 H2O (use hot base, disproportion of X = Cl, Br) b) electrolysis: − − 2 Cl → Cl2 + 2 e (anode(+) reaction) − − 2 H2O + 2 e → H2 + 2 OH (cathode(–) reaction) Mixing: − − − Cl2 + 2 OH → Cl + OCl + 2 H2O − − − 3 OCl → ClO3 + 2 Cl (disproportion) oxidation on anode(+): − − + − OCl + 2 H2O → ClO3 + 4 H + 4 e

– BrO3 (bromate) – – – + Synthesis: Br + Cl2 + H2O → BrO3 + Cl + H3O

– IO3 (iodate) – Synthesis: I2 + NaClO3 → NaIO3 + Cl2 (at high temperature IO3 is more stable) I I Salts: M H(IO3)2, more over M H2(IO3)3 etc.: hydrogen diiodate, dihydrogen triiodate. Properties: thermal decomposition: 4 NaClO3 → NaCl + 3 NaClO4 (fast reaction at 200 °C) 2 NaClO3 → 2 NaCl + 3 O2 (with MnO2 catalyst) 4 NaBr/IO3 → 2 Na2O + 2 Br2/I2 + 5 O2 NH4XO3 → HXO3 + NH3 → N2 + H2O (explosion danger)

– – – Redox reaction: (oxidizer strength: BrO3 ≥ CO3 > IO3 , strong pH dependence), e.g.: – – + 2 BrO3 + 2 Cl + 12 H → Br2 + Cl2 + 6 H2O Chemical clock reaction: autocatalytic (1885 Landolt) – 2– – 2– – – + IO3 + 3 SO3 → I + 3 SO4 (initial step), 5 I + IO3 + 6 H → 3 I2 + 3 H2O (production of I2), starch is blue. 2– – + 2– 3 I2 + 3 SO3 + 3 H2O → 6 I + 6 H + 3 SO4 , I2 disappears, starch is colorless. Oscillating reaction, chemical wave: complicated kinetics HIO3 + 5 H2O2 → 5 O2 + I2 + 6 H2O (production of I2), I2 + 5 H2O2 → 2 HIO3 + 4 H2O (decrease the I2).

26 János Rohonczy: Inorganic Chemistry I. 3.1. Oxygen

– 2− Application: synthesis of ClO2: 2 ClO3 + SO2 → 2 ClO2 + SO4 (in acid), bleaching of cellulose (without destruction), production of other: MClO3 and MClO4 salts. KClO3 strong oxidizer agent in pyrotechnics (non-hygroscopic).

HXO4 (perhalogenic acids, perhalogenate salts) Oxidation number = +7, thermally very stable, kinetic stability, lower oxidation ability.

HClO4 (perchloric acid) Physical properties: colorless liquid, HOClO3 molecules in gas phase, AX4 structure. Synthesis: NaClO4 + cc. HCl → HClO4 + NaCl (distillation from salt with cc. HCl) Ba(ClO4)2 + 2 HCl → 2 HClO4 + BaCl2 (azeotropic at 203 °C, distillation on cc. H2SO4, production of anhydrous acid) Chemical properties: Strong, colorless acid: Mg + 2 HClO4 → Mg(ClO4)2 + H2 Ag2O + 2 HClO4 → 2 AgClO4 + H2O In diluted solution weak oxidizer agent (kinetic hindrance): + – H + ClO4 + H2S/SO2/HNO2/HI→ "no reaction". Concentrated acid can dissolve the gold metal. Anhydrous acid is explosive in the presence of organic compounds or by mechanical strike. Thermal decomposition results in HCl, Cl2, Cl2O, O2. There are lots of hydrates. + – Weak electric conductor: 3 HClO4 � Cl2O7 + H3O + ClO4

– ClO4 (perchorates) Industrial / laboratory synthesis: electrolysis of water solution: – – + – ClO3 + H2O → ClO4 + 2 H + 2 e (on Pt-anode(+)) (Na2Cr2O7 additive to protect the perchlorate reduction on the cathode) Only in laboratory: a) KClO3 + cc. H2SO4 → ClO2 + KClO4 (dangerous) – 2– – b) ClO3 + O3/S2O8 /PbO2 → ClO4 c) NaClO4 + NH4Cl → NH4ClO4 + NaCl.

Application: Mg(ClO4)2 electrolyte in batteries. KClO4 oxidizer in pyrotechnics. NaClO4 (30,000 t/year) + NH4Cl → NH4ClO4, NH4ClO4 in solid rocket fuel: NH4ClO4 : Al(powder) = 7 : 3. – HBrO4 / BrO4 ( / perbromates) – – – – Synthesis (1968): BrO3 + F2 + 2 OH → BrO4 + 2 F + H2O (ca. 20%) – – + BrO4 + F + 2 Ag → AgBrO4 + AgF (precipitate), dissolution from cation exchanging resin → HBrO4

Properties: white crystalline powder. Stable salts: e.g. KBrO4, NH4BrO4. Thermal decomposition: KBrO4 → – – KBrO3 + 1/2 O2 (at 270 °C). Oxidation properties: stronger than HClO4: reacts with I , Br .

Periodic acids / periodates

Structure: several alterations: H5IO6 (orthoperiodic acid) HIO4 (metaperiodic acid)

27 János Rohonczy: Inorganic Chemistry I. 3.1. Oxygen

"H3IO5" → H6I2O10 (meso- or diperiodic acids) H7I3O14 (triperiodic acid) Properties: increasing the pH in water solution: deprotonation, dimerization, several anions are known, which are white, water soluble powders. – – 5– Synthesis of salts: oxidation of I / I2 / IO3 → IO6 : a) by Cl2: NaIO3 + 6 NaOH + Cl2 → Na5IO6 + 3 H2O + 2 NaCl – – – 5– b) by electrolysis on anode(+): IO3 + 6 OH - 2e → IO6 + 3 H2O c) oxidation of NaI in base: NaI + 2 Na2O + O2 → Na5IO6 (at 350 °C)

Synthesis of other salts: Na5IO6 + 2 H2O → Na3H2IO6 + 2 NaOH (hydrolysis) Ba(IO3)2 → Ba5(IO6)2 + 4 I2 + 9 O2 + + – Na3H2IO6 + K → KIO4 + 3 Na + 2 OH (in H2O/HNO3)

+ – Synthesis of acids from salts: Ba5(IO6)2 + 10 H + 10 NO3 → 5 Ba(NO3)2 + 2 H5IO6 Heating: very stable compounds: H5IO6 → H7I3O14 (120 °C) H5IO6 → HIO4 + 2 H2O → O2+ H2O + I2O7⋅I2O5 → I2O5 + O2 (at 100 °C in vacuum) + – + – Reaction with strong acids: H5IO6 + H + ClO4 → [I(OH)6] ⋅ClO4 (white salt) 5– 2+ + – – Oxidation agent in acids: 5 IO6 + 2 Mn + 14 H → 5 IO3 + 2 MnO4 + 7 H2O

III Synthesis of complexes: Na3K[H3Cu (IO6)2]⋅14 H2O. Stable complexes with exotic cations: Ni4+ / Cu3+ / Ag3+.

3.1.4. Halogen oxofluorides and fluorinated oxoacids

Halogen oxofluorides

Oxides of interhalogens. Anions and cations are formed from neutral molecules by F– addition and elimination.

Table 15. Halogen oxofluoride molecules, cations and anions. (Central Atom: Cl, Br, I).

Formula FClO FClO2 FClO3 F3ClO F3ClO2 F5IO Oxidation No. +3 +5 +7 +5 +7 +7 Structure AX2E2 AX3E AX4 AX4E AX5 AX6 Ion + F– + F– - F– - F– – – + + Ion formula F2ClO2 F4ClO F2ClO F2ClO2 Structure AX4E AX5E AX3E AX4

Synthesis: Cl2O + 2 F2 → F3ClO + ClF (NaF, -78 °C) – + – Reaction: as Lewis base: F donor: F3ClO + AsF5 → [F2ClO] [AsF6] – + – as Lewis acid: F acceptor: F3ClO + CsF → [Cs] [F4ClO] Halogenated oxoacids Structure: F–O–ClO3 / Cl–O–ClO3 / Br–O–ClO3. Synthesis: CsClO4 + Cl–O–SO2F → Cl–O–ClO3 + CsSO3F.

28 János Rohonczy: Inorganic Chemistry I. 3.2. Sulfur

3.2. Sulfur(S)

Valence shell: [Ne]3s23p4 General properties: Yellow non-metal crystals, 6 electrons on the valence shell. Oxidation number: -2, (+2) ,+4 ,+6. Occurrence: Elementary form (S8) in volcanic gas, over salt-bed, in the atmosphere of Venus. 2– 2– Minerals: sulfides (S ), earth gas, mineral oil, oil shale, mineral coal; (S2 ): pyrite FeS2. Sulfide containing ores: Mo, Fe, Ni, Cu, Ag, Zn, Cd, Hg, Ga, In, Tl, Pb, As, 2– Sb, Bi, Se, Te. SO2 (): flue gas – acid rain, SO4 (): CaSO4⋅2 H2O (gypsum), CaSO4 (anhydrite), MgSO4⋅7 H2O (epsomite), in K, Al, Fe, Cu, Pb containing minerals. Stable isotopes: 32S 95% / 33S 0.75%, I = 3/2 NMR active (hard to measure) / 34S 4 % / 36S 0.02%. Mw = 32.06±0.01 depends on the origin: mineral oil / asteroid / ocean / sulfur-mine.

35 Radioactive isotopes: 6 ones. Most stable: S t1/2 = 88 days. * * * * Commercial products: S 8, H2S , S OCl2, KS CN

Synthesis: a) Frasch-method: from 100-200 meter below the surface of the Earth. Essence: 1. Sulfur is melted by high pressure hot water (165 °C). 2. The mixture of liquid sulfur, water and air is coming up to the surface in the middle pipe of a three member concentric pipe system. (Outer line: hot water inlet, internal pipe: air inlet). Purity: 99.5%. b) From earth gas: 1. absorbing of H2S in ethanol (H2N–CH2–CH2OH) 2. partial oxidation to elementary sulfur (S8): H2S + 1/2 O2 → S + H2O H2S + 3/2 O2 → SO2 + H2O 2 H2S + SO2 → 3 S + 2 H2O (Fe2O3 catalyst / 300 °C) c) From mineral oil: Partial reduction: S → H2S, after that by method b) → S8. II d) Roasting of sulfides: M S + O2 → MO + SO2 { + O2 + H2O} → H2SO4. e) CaSO4 + 2 NH3 + CO2 + H2O → CaCO3 + (NH4)2SO4 (fertilizer).

Application: a) SO2/SO3/H2SO4: 88%: fertilizer (superphosphate), (NH4)2SO4, galvanic bath, batteries, petrol chemical processes, synthetic detergents (sulfonic acids, sulfates), (nitration), etc. b) CS2: cellophane, viscose rayon, CCl4, rubber industry. c) S8: curing of rubber, fungicides, match, gunpowder. d) SO2: bleaching, heat exchanger in refrigerators (liquid SO2), sugar industry. Allotropes: Several allotropes: –S–S– chains (stable, flexible), several crystal lattices. dS–S = 180-260 pm / αS–S–S = 90-180° / ΦS–S–S–S = 0-180°.

Natural allotropes: (all are cyclo-S8)

α-S8 (orthorhombic) Properties: ρ = 2,07 g/cm3, MP = +119.6 °C, good thermal conductor, good electric insulator János Rohonczy: Inorganic Chemistry I. 3.2. Sulfur

Solubility: soluble in CS2, S2Cl2, Me2CO, benzene, CCl4, EtOH. Structure: symmetrical, all α = 107,9°.

β-S8 (monoclinic) Structure: distorted angles, loosely packed crystal lattice, lower density: ρ = 1.94 g/cm3. Properties: stable over 95.3°. MP = +115.1°, reversible metamorphosis of α-S8 � β-S8 Enantiotropic pair. β-S8 MP = +119° > 95.3° = Tmetamorphic (see: β-quartz � β-tridymite � β- cristobalite reversible enantiotropic polymorphism, opposite: diamond � graphite, irreversible monotropic metamorphosis). Synthesis: melting of α-S8 at 120 °C, slow cooling, light yellow pin crystals for several days: β-S8, after that yellow powder, which is α-S8 – stable at RT.

γ-S8 (orthorhombic) Properties: MP = 106.8°, ρ = 2.19 g/cm3. Synthesis: slow cooling from 150 °C or crystallization from hot, concentrated .

Synthetic cyclo allotropes cyclo-S6 (ε) (rhombohedral) Structure: strained but symmetric 6 member ring, angle = 100°. Properties: (1891) ρ = 2.21 g/cm3 – highest density! Synthesis: H2S4 + S2Cl2 → cyclo-S6 + 2 HCl (in diluted ethanol). cyclo-Sn (n = even number) Synthesis: S12-nCl2 + H2Sn → cyc-S12 + 2 HCl. Similar: S6, S10, S12, S18, S20. cyclo-Sn (n = odd number) 5 Synthesis: SxCl2 + [Ti(η -C5H5)2S5] → cyc-Sx+5 + [Ti(Cp)2Cl2]. Similar: S7, S9, S10, S11. or SO2Cl2 + [Ti(Cp)2S5] → cyc-S10 + 2 [Ti(Cp)2Cl2] + 2 SO2. Properties: cyc-S7 four allotropes. Bond angle: 105.5-107.5° cyc-S10 light sensitive, rare crystal lattice. Bond angle: 78.5-110° cyc-S12 MP = +148 °C (high). Synthesis: H2S8 + S4Cl2 → S12 + 2 HCl (in diethyl ether). cyc-S18 two allotropes, built from two helices. MP = 128 °C, lemon yellow crystal. Synthesis: H2S8 + S10Cl2 → S18 + 2 HCl. cyc-S20 lemon yellow, dissociation at 124 °C. Synthesis: H2S10 + S10Cl2

Synthetic poly catena allotropes Structure: they contain chains, helices elastic-S (stretching: 15x), χ plastic-S / φ, ψ fibrous-S (bunches of well defined helices, hexagonal crystal lattice) / µ polymer-S / ω insoluble-S / white-S. Difficult structure elucidation. Synthesis: from ca. 400 °C molten sulfur. Meta stable mixture: cyclo-S8 and catena-S∞

Liquid sulfur Properties: At T ≈ 160 °C cyclo-S8 and catena-S8. Increasing the temperature the viscosity, color, electric conductivity is changing: catena-S8 + catena-S8 → catena-S16 (light sensitive, radical mechanism) catena-S16 → catena-S200.000 (180 °C, high viscosity), brown → yellow

30 János Rohonczy: Inorganic Chemistry I. 3.2. Sulfur

catena-S1000 (T = 400 °C) → catena-S100 (600 °C), decreasing viscosity, thermal degradation. Sulfur vapor 3 Properties: Sn, 2≤ n ≤10, S3 dark red, T > 720 °C: S2 violet gas, molecular states: triplet S2, 3 1 like O2, and singlet S2 (excited). 3 3 Synthesis of gas S2: CS2 → S2 (by UV light), 1 1 Synthesis of gas S2: 2 COS → 2 CO + S2 (by UV light).

Sulfur (solid) Chemical properties: Reactive element (at higher temperature), break of S–S bonds. Binary compounds: by direct synthesis (except noble gases – there are no binary sulfides): N2 / Te / I2 / Ir / Pt / Au elements can bond to sulfur. Reaction with sulfur, sulfides H2 + S → H2S F2 + S → SF Cl2 + S → S2Cl2 + SCl2 Br2 + S � S2Br2 I2 + S → physical solubility only. O3 + S → SO2 + SO3 with hot O2 + S → SO2 + SO3 . Reaction with B, C, Si, Ge, P, As, Sb, Se at high temperature. Reaction with Cu, Ag, Hg, Pb, Sn, Bi at RT. Violent reaction with Fe, Co, Ni, Mn, W, La, Ac at high temperature. Reactivity: cyclo-S8 < catena-S8 < S2

3.2.1. Sulfur containing compounds

Oxidation number: -2 , +6. Bond type: covalent. Sulfides:, coordination number: 2-10 ligands. Ionic or metallic character.

2+ Poly atomic Sn ions

History: (1804) S8 + oleum → dark yellow / red / blue liquid. 2+ – Synthesis with strong oxidizing agents: S8 + 3 AsF5 → [S8 ][AsF6 ]2 + AsF3, dark blue liquid. Structure: S–S bridge inside the S8 ring. 2+ – 2+ – Similar: [S4 ][SbF6 ]2, light yellow crystal [S19 ][SbF6 ]2

Sulfur ligands: Terminal, μ1 or μ2-6 positions, soft ligand, easy to deform.

–S– bridge

µ2 µ3 µ4 µ6-8 2 el. donor 4 el. donor 6 el. donor ionic µ µ 2-S(Au)2 3-S(Au(PPH3))3

S5 ligand. Synthesis (1903): [Ti(Cp)2Cl2] + Na2S5 → [Ti(Cp)2S5] + 2 NaCl. Structure (1903): TiS5, 6 member ring. 2– Analogous: H2PtCl6 + (NH4)2Sx → (NH4)2[Pt(S5)3] , chirality

31 János Rohonczy: Inorganic Chemistry I. 3.2. Sulfur

3.2.2. Sulfur hydrides – Sulfanes

H2S (dihydrogen sulfide) Properties: Stable, abundant in nature, smell: rotten egg, very toxic. More volatile than the water, there is no H-bridge, low dielectric constant (9), weakly polar. Chemical solubility: soluble in water, acids, bases + – H2S + H2O � H3O + HS pK = 7 – + 2– HS + H2O � H3O + S pK = 14 (weak acid) Hydrate: H2S⋅5,75 H2O. Synthesis: a) FeS + HCl → FeCl2 + H2S (diluted acid, Kipp-apparatus) b) CaS/Al2S3 + HCl → H2S + CaCl2/AlCl3 (higher purity) c) H2 + S8 → H2S (at 600 °C, highest purity) Reaction: H2S + air → H2O + SO2 / S (burning) + – H2S + SbF5 + HF → [H3S] [SbF6 ] (H2S protonable with very strong acids)

H2Sn(2-8) () Properties: Non stable, catena-S Structure, increasing n → increasing viscosity. H2S2 colorless, H2Sn yellow, smelly. Synthesis: Na2S + S8 → Na2Sx (boiling of water solution). Na2Sx + diluted HCl → H2Sx + 2 NaCl (at -10 °C), yellow oil, distillation in vacuum. Reaction: H2Sx → H2S + S8 (thermal decomposition, catalysts: glass, alkaline metal salts) – – H2S2 + CN → H2S + SCN 2– 2– H2S2 + SO3 → H2S + S2O3

S2– (sulfides) Occurrence: important minerals, ores. Application: MS + air → MO + SO2 / MSO4 / M + SO2 Synthesis: a) Fe + S → FeS, b) Na2SO4 + 4 C → Na2S + 4 CO 2+ + c) Cd + H2S → CdS + 2 H d) NaOH + H2S → NaHS + H2O NaHS + NaOH → Na2S + H2O. Properties: soluble ones: alkaline and alkaline earth metal salts; hydrolysis, ph>7, production of H2S. Insoluble salts: with heavy metals. Different colors (qualitative analysis of ions).

3.2.3. Sulfur halides

Sulfur fluorides: Oxidation number: +1…+6, coordination number: 1...6.

S2F2 (disulfur difluoride) Physical properties: BP = 15 °C, instable. Structure: F–S–S–F.

Synthesis: 3 S8 + 16 AgF → 8 S2F2 + 8 Ag2S (at 125 °C).

SSF2 (thio-thionyl fluoride) Physical properties: BP = -11 °C. Structure: S=SF2 Synthesis: S2F2 → SSF2 (KF catalyses the isomerisation)

32 János Rohonczy: Inorganic Chemistry I. 3.2. Sulfur

Reaction: SSF2 + O2 → SO2 + SOF2 +SO2F2 SSF2 + H2O → S8 + HF + H2S4O6 SSF2 → S8 + SF4 (thermal decomposition)

SF2 (sulfur difluoride). Structure: F–S–F, instable.

Synthesis: SCl2 + HgF2 → SF2 + HgS + HgCl2 (at 150 °C), byproducts: S=SF2, S2F2, Cl–S–S–F, Cl–S–SF3, F3S–SF.

SF4 (sulfur tetrafluoride) Physical properties: BP = -38 °C. Synthesis: 3 SCl2 + 4 NaF → S2Cl2 + SF4 + 4 NaCl (at 75 °C in MeCN solvent). Properties: strong, selective fluorination agent:

Reactions: SF4 + H2O → HF + SO2, SF4 + R2C=O → R2CF2/R–COOH/R–CF3 SF4 + Cl2 + CsF → SClF5 + CsCl, SF4 + I2O5 → 2 IF5 + 5 O=SF2

S2F10 Physical properties: BP = 30 °C. Structure: F5S–SF5. Properties: stronger fluorination agent than SF4.

Synthesis: 2 SClF5 + H2 → S2F10 + 2 HCl (UV light).

SF6 (sulfur hexafluoride) Physical properties: BP = -64 °C. Synthesis: S + 3 F2 → SF6, or SF4 + F2 → SF6 Properties: colorless, tasteless, non-toxic, non-flammable, insoluble, very good electric INSULATOR (5 bar, 5 cm, 1 million V), non-reactive (noble) gas. Reaction: SF6 + Na(molten) → reduction of S! SF6 + H2O → no reaction (kinetic barrier) SF6 + metals, As, P → no reaction (even at 500 °C).

3.2.4. Sulfur oxohalides

Sulfur oxofluorides

Table 16. Molecular structures and names of sulfur oxofluorides. Thionyl fluoride Peroxo-sulfuryl fluoride

O=SF2 O2SF2 F–SO2–OO–SO2–F F–SO2–OH

Synthesis: SO3+ HF → FSO2(OH) FSO2(OH) + SO3 + SbF5 ("super acid" or "": names of strongest acids).

3.2.5. Sulfur chlorides/bromides/iodides

S2Cl2 (disulfur dichloride) Physical properties: BP = 138 °C, toxic yellow liquid. Synthesis: S8 + Cl2 → S2Cl2. Structure: like H2O2. Application: vulcanisation of rubber, exploration of metals, extraction.

33 János Rohonczy: Inorganic Chemistry I. 3.2. Sulfur

SCl2 (sulfur-dichloride) Physical properties: BP = 59 °C, cherry-red liquid. Synthesis: S2Cl2 + Cl2 → SCl2 (FeCl3 catalyst). Reaction: 2 SCl2 � S2Cl2 + Cl2 (decomposition, inhibitor: PCl5) SCl2 + H2O → H2S, SO2, H2SO3, H2SO4 (hydrolysis) SCl2 + O2 → SOCl2, SO2Cl2 (oxidation) 3 SCl2 + 4 NaF → SF4 + S2Cl2 + 4 NaF (disproportion) Application: SCl2 + 2 H2C=CH2 → S(CH2–CH2–Cl)2 ()

SxCl2 (polysulfur dichloride) Properties: orange liquid.

SCl4 (sulfur tetrachloride) Physical properties: instable. + – Structure: [SCl3 ][Cl ] + – Reaction: SCl4 + AlCl3 → [SCl3 ][AlCl4 ] (stable salt).

SBr2 (sulfur dibromide) Physical properties: stable only at very low temperature.

S2Br2 () Synthesis: S8 + Br2 ↔ S2Br2 (at 100 °C).

SI2 / S2I2 (sulfur diiodide and disulfur diiodide) Physical properties: Endothermic compounds, still not prepared.

+ – Derivatives: I2 + S8 + SbF5 → [S7I ][SbF6 ] + SbF3 (dark orange crystal)

2+ – 2 I2 + 1/4 S8 + 3 AsF5 → [S2I4 ][AsF6 ]2 + AsF3

3.2.6. Sulfur oxohalides

SOCl2 () Physical properties: liquid, SOBr2 is known as well. Laboratory synthesis: SO2 + PCl5 → SOCl2 + O=PCl3 Industrial synthesis: SO3 + SCl2 → SOCl2 + SO2 Application: AlCl3⋅6 H2O + SOCl2 → AlCl3 + 6 SO2 +12 HCl (removal of crystal water).

SO2Cl2 () Physical properties: colorless liquid, instable, SO2Br2 is known as well. Synthesis: SO2 + Cl2 → SO2Cl2 (activated carbon or FeCl3 catalyst, cooling). Application: chlorination of organic compounds, chlorosulfonation, build-up SO2Cl group.

3.2.7. Sulfur oxides

SO2 (sulfur dioxide) Synthesis: burning of S, H2S, roasting of FeS2. 2 H2S + 2 O2 → SO2 + 2 H2O Properties: colorless, pungent, toxic gas. Good water solubility, chemical dissociation. Weak acid in water, SO2⋅6 H2O, low dielectric constant, ε = 15.4. Solvents: SOX2, PCl3, CS2, alcohols, . 2+ 2– Autodissociation: 2 SO2 � SO + SO3 . Solvolysis:2 SO2 + UCl6 � UO2Cl2 + 2 SOCl2

34 János Rohonczy: Inorganic Chemistry I. 3.2. Sulfur

Application: synthesis of SO3: 2 SO2 + O2 → 2 SO3 (catalyst: V2O5), bleaching, disinfection, conservation, refrigerant, aprotic solvent, synthesis of salts: , , sulfo chlorination of organic compounds. Removal: a) SO2 + Ca(OH)2 → CaSO3/CaSO4. Ca(HSO3)2 is used in cellulose industry, b) SO2 + H2 → H2S + H2O { + O2 } → 2 H2S + SO2 → 3 S + H2O. Detection: a) SO2 + H2O2 → H2SO4, b) flame photometry, 9 c) pulsed fluoroscopy ( SO2 is fluorescent in UV light, sensibility: 1:10 ). 1 1 2 Structure: V-shape (angle: 119°), SO2 as ligand: η -SO2, η -bridge-SO2, η -SO2, etc.

SO3 () Structure: in gas phase: planar triangle, in liquid: S3O9 trimer (γ-SO3), in solid: chain polymer: –S–O–S–O... Synthesis: a) 2 SO2 + O2 → 2 SO3, polymerization in liquid, inhibitor: B2O3, SOCl2 b) heating of sulfates, c) distillation of oleum: cc. H2SO4 + 25-65% SO3 → SO3

Reaction: a) SO3 + H2O → H2SO4 b) SO3 + HCl → HO–SO2–Cl () c) SO3 + carbohydrates → H2SO4 + "C" d) SO3 + NH3 → H2N–SO3H () Removal: SO3 + Fe3O4 → FeSO4 + Fe2(SO4)3

"Lower oxides" – SnO2 / SnO

Synthesis: a) cyclo-Sn + CF3–CO–OOH → SnO2, SnO, S6O2, S7O2 ... b) SO2 + S → S12O2 c) SO2 → S2O / S2O2 / SO.

"Higher oxides"

SO4 Synthesis: a) SO2 / SO3 + O2 → [–O–SO2–O–O–SO2–O–]n b) SO3 + O3 → SO4. Structure: O2–SO2(?). Hydrolysis: SO4 + H2O → H2SO4 + H2SO5 + H2O2 + O2 + ...

3.2.8. Sulfur oxoacids

Numerous stable/known oxoacids. Most of them are important. Several acids are instable, but their salts are stable and useful.

Table 17. Sulfur oxoacids, anions and their names. Formula Name Oxidation No. Anion Name of anion 2– H2SO4 +6 SO4 – HOSO3 hydrogen-sulfate 2– H2S2O7 +6 S2O7 disulfate 2– H2S2O3 +6/-2 or +4/0 SSO3

35 János Rohonczy: Inorganic Chemistry I. 3.2. Sulfur

2– H2SO5 peroxomonosulfuric acid +6 OOSO3 peroxomonosulfate 2– H2S2O8 peroxodisulfuric acid +6 S2O8 peroxodisulfate * 2– H2S2O6 +5 S2O6 2– H2Sn+2O +5/0/+5 S n+2O6 polithionate 6 * 2– H2SO3 +4 SO3 – HOSO2 hydrogen-sulfite * 2– H2S2O5 disulfurous acid +5/+3 S2O5 disulfite * 2– H2S2O4 dithionous acid +3 S2O4 *instable acid

Logical relations: H2O + SO3 → H2SO4 + SΟ3 → H2S2O7 ↓ + SO2 H2S2O6 H2O + SO2 → H2SO3 + SO2 → H2S2O5 H2O2 + SO3 → H2SO5 + SO3 → H2S2O8 H2S + SO3 → H2S2O3 + SO3 → H2S3O7 H2Sn + SO3 → (H2Sn+1O3) + SO3 → H2Sn+2O6 Redox properties: several kind of reactions. Extremes: 2– + – – a) S2O8 + 2 H + 2 e → 2 HSO4 E0 = +2.123 V (strong oxidizer) + – – b) 2 H2SO3 + H + 2 e ← HS2O4 + 2 H2O E0 = –0.082 V (strong reducing agent)

H2SO4 / H2S2O7 (sulfuric acid / disulfuric acid) Properties: Anhydrous sulfuric acid is colorless, viscous liquid, decomposes at 300 °C. Well soluble in water, very exothermal reaction. Strong acid. Dissolves SO3 gas. Dielectric constant = 100. Synthesis: a) FeSO4⋅4 H2O → H2SO4 + Fe2O3 b) SO2 + HNO3/NO2 → SO3 + NO + H2O → [H2SO4⋅NO] (blue acid) [H2SO4⋅NO] + H2O + air → H2SO4 + NO2 c) S / FeS2 / Cu2S / NiS + O2 → SO2 + 1/2 O2 → SO3, catalysis: SO2 + V2O5 → SO3 + V2O4, V2O4 + 1/2 O2 → V2O5 SO3 + H2SO4 → H2S2O7 (and few H2S4O13) H2S2O7 + H2O → H2SO4 (dilution with water) Produce: > 100 Mt/year. + – Properties: Autodissociation: 2 H2SO4 � H3SO4 + HSO4 (electric conductivity), + – hydrolysis: 2 H2SO4 + H2O � H3O + HS2O7 + – H2SO4 + H2O → H3O + HSO4 non volatile: H2SO4 + KCl → KHSO4 + HCl + – + used in nitration: HNO3 + 2 H2SO4 → NO2 + 2 HSO4 + H3O

2– – Salts (sulfates, hydrogen-sulfates) Structure: SO4 , HOSO3 , tetrahedral, AX4 Synthesis: a) Fe + H2SO4 → FeSO4 + H2 b) Ca(OH)2 + H2SO4 → CaSO4 + 2 H2O c) Na2CO3 + H2SO4 → Na2SO4 / NaHSO4 + H2O + CO2 d) CaCl2 + Na2SO4 → CaSO4 + 2 NaCl e) 3 K2SO3 + 2 KMnO4 + H2O → 3 K2SO4 + 2 MnO2 + 2 KOH

36 János Rohonczy: Inorganic Chemistry I. 3.2. Sulfur

H2S2O3 (thiosulfuric acid). Properties: instable in water as well, its salts are stable. Synthesis: Na2S2O3 + 2 HCl → H2S2O3 + 2 NaCl – – 2– Synthesis of salts: a) 2 HS + 4 HSO3 → 3 S2O3 + 3 H2O – 2– + b) 3/8 S8 + 3 HSO3 + 3 H2O → 3 S2O3 + 3 H3O c) Na2SO3 + S → Na2S2O3

Industrial synthesis: a) Na2S5 + O2 → Na2S2O3 + S.

b) CaS2 + O2 → CaS2O3

Application: AgBr + 3 Na2S2O3(aq) → Na5[Ag(S2O3)3](aq) + NaBr(aq) 2– 2– – 2 S2O3 � S4O6 + 2 e , E0 = +0.169 V 2– – 2 S2O3 + I2 → S4O6 + 2 I (tetrathionate) 2– – + – S2O3 + 4 Cl2 + 13 H2O → 2 HSO4 + 8 H3O + 8 Cl 2– S2O3 + Br2 → depends on concentrations, temperature, etc.

H2SO5 (, Caro acid). Properties: white solid, MP = +45 °C, explosive. Synthesis: cc. H2O2 + ClSO2(OH) → HOOSO2OH + HCl.

H2S2O8 () Properties: white solid, Decomposition at 65 °C. Synthesis: by electrolysis of (K/NH4)2SO4 solution. On anode: → (K/NH4)2S2O8, weakly soluble, crystalline. With H2SO4 form the free acid. Strong oxidizer agent: E0 > +2V).

H2S2O6 (dithionic acid) Properties: instable, but salts are stable. Synthesis of salts: 2 MnO2 + 3 SO2 → MnSO4 + MnS2O6 , {+BaCl2} → BaS2O6 BaS2O6 + H2SO4 → H2S2O6 (in solution)+ BaSO4 2– 2– Properties: S2O6 + X2 / KMnO4 / K2Cr2O7 → SO4 2– 2– S2O6 + Na(Hg) → S2O4 (dithionite)

H2SnO6 (polythionic acids, 1808) Structure: HO3S–Sn-2–SO3H (polysulfan sulfonic acids) 2– 2– – Synthesis: a) S2O3 + I2 → [S4O6] + 2 I ,

Other synthesis:

– 2– a) S2Cl2 + 2 HSO3 → [S4O6] + 2 HCl

b) H2S + SO2 → H2S5O6 + H2O,

c) KNO2 + K2S2O3 + cc. HCl → K2S6O6

H2SO3 (sulfurous acid). Properties: exists in aqueous solution, stable salts and acidic salts. Synthesis: SO2 + H2O � H2SO3. Salts: NaOH + SO2 + H2O → NaHSO3 NaHSO3 + NaOH → Na2SO3 + H2O. Industrial synthesis: SO2 + H2O + Na2CO3 + H2O → NaHSO3 + NaHCO3 NaHSO3 + Na2CO3 → Na2SO3 + NaHCO3

37 János Rohonczy: Inorganic Chemistry I. 3.2. Sulfur

NaHSO3 + SO2 + Na2CO3 → Na2S2O5 (disulfite) + NaHCO3. Application: Na2SO3/K2SO3 1 Mt/year : paper industry / photo industry. Structure: AX3E – – – + Reaction: a) H2O + HSO3 + I2 → HSO4 + 2 I + 2 H 2– 2– – + b) 2 H2O + 2 SO3 + 2 Na/Hg → S2O4 + 4 OH + 2 Na

2– – 2– 2– – c) 2 SO3 + 4 HCO2 → S2O3 + 2 C2O4 + 2 OH + H2O

H2S2O5 (disulfurous acid). Properties: only the salts are stable. – 2– Synthesis: 2 HSO3 � S2O5 + H2O. 2– + – Reaction: S2O5 + H → HSO3 + SO2

H2S2O4 (dithionous acid). Properties: only the salts are stable. Synthesis: a) NaHSO3 + SO2 + 2 Zn + H2O → ZnSO3 + Na2S2O4 + 2 H2O b) NaHSO3 + NaBH4 → Na2S2O4 – 2– 3+ 2+ 3+ + Properties: reduction agent: I2 → I / CrO4 → Cr / TiO → Ti / Cu → Cu Application: reduction agent, paper-, cellulose industry.

38 János Rohonczy: Inorganic Chemistry I. 3.3. Selenium, Tellurium, Polonium

3.3. Selenium (Se), Tellurium (Te), Polonium (Po) Valence shell: ns2np4 Occurrence: In earth's crust Se: 66-th (0.05 ppm), Te 73-rd (0.002 ppm), Po: (3⋅10-10 ppm), mainly selenides, tellurides and oxides of the following elements: Ag, Hg, Cu, As, Zn, Cd, Bi, Pb, Fe, Ni. Other minerals: NiSeO3, PbTeO3, SeO2, TeO2.

Table 18. Discovery of selenium, tellurium and polonium. Year Discoverer Origin of the name Notes Se 1817 Berzelius Selene – Moon burning of CuS2: dark red cinders Te 1782 Franz Müller von Tellus – Earth faked antimony ore (Transylvania) Reichenstein Po 1898 Marie Curie Polska – Poland Fractional crystallization. From pitchblende 0.1 mg Po in 1 ton ore

3.3.1. Preparation, produce and application

Synthesis a) electrochemical refining of Cu → in anode slime: Pt, Ag, Au, Se, Te b) Ag2Se + Na2CO3 + O2 → 2 Ag + Na2SeO3 + CO2 Cu2Te + Na2CO3 + O2 → 2 CuO + Na2TeO3 + CO2 c) Cu2Se + 3/2 O2 → CuO + CuSeO3 → 2 CuO + SeO2 Ag2SeO3 → 2 Ag + SeO2 + 1/2 O2 2– 2– d) SeO3 + TeO3 + H2SO4 → H2SeO3 + TeO(OH)2 H2SeO3 + 2 SO2 + H2O → Se + H2SO4 e) Na2TeO3 + H2O → Te + 2 NaOH + O2

Synthesis of Po: 209Bi (n, γ) → 210Bi (β–) → 210Po(α) Refine: fractional distillation of bismuth. 210 206 Properties: strong α radiation: Po (α) → Pb, t1/2 = 140 days Produce: Se: 1500 t/year, Te 150 t/year.

Application: Se: red color of glass, pigments, ink, thermo-color paint [Cd(S,Se)] in xerography (photo-semiconductor layer on cylindrical drum), Fe/Se stainless steel, catalyst in vulcanization of rubber: Et2N–CS2–Se–CS2–NEt2, laboratory : SeO2, Na2SeO3, Na2SeO4, SeOCl2 Te: in stainless steel, glass paint, catalyst in rubber industry, laboratory precursors: Te, Fe/Te, TeO2, Na2TeO4, Te(Et2NCSS2)2 Po: a) in neutron generator: 9Be (α,n) → 12C, 1 neutron / 104 α particles b) thermocouple

Allotropes

Se allotropes. 6 natural allotropes. Most stable is gray-Se (MP = 221 °C), others: red α,β,γ−cyclo-Se8, amorphous red and glassy black-Se. János Rohonczy: Inorganic Chemistry I. 3.3. Selenium, Tellurium, Polonium

Gray-Se Synthesis: slow cooling of molten Se or condensation of Se vapor. th st Structure: catena – helices (similar to fibrous-sulfur): [–Se–]n chains: 4 Se over 1 one. Properties: Hexagonal gray crystals, insoluble in CS2. Photo-semiconductor.

Black glassy Se Synthesis: Quick cooling of molten Se. Structure: random, ca. 1000 atoms in chains and rings

Red ciklo-Se8 Synthesis: a) boiling of CS2 solution of black-Se, and cooling down to form crystals. α,β-cyclo-Se8, b) thermal decomposition of Se containing compounds: γ-cyclo-Se8 Properties: well soluble in CS2.

Red amorphous-Se Synthesis: H2SeO3 + SO2 → H2SO4 + Se(red) Properties: bad solubility in CS2. Chain polymer structure, electric insulator.

Te allotropes: Only one allotrope, Structure like crystalline gray-selenium.

Po allotropes: Two metal allotropes (cubic and rhombohedral crystals). Properties: silvery white, estimated phase transition temperature (self heating by radioactivity).

Nuclear properties

Se 6 stable isotopes: 80Se (50%) stable, 77Se (7,5%) (I = 1/2 NMR active), but 82 20 Se (9,5%) t1/2 = 10 year

Te 8 stable isotopes: 130Te (34%) / 128Te (32%) / 123Te (1%) (I = 1/2 NMR active)

125Te (7%) (I = 1/2 NMR and Mössbauer active)

208 Po 27 known, but all radioactive isotopes, longest half-life: Po t1/2 = 2,9 year / 209 210 Po t1/2 = 100 year / Po t1/2 = 139 days. Chemical reactivity. They are less reactive than sulfur, increasing metallic behavior: O,S: insulators; Se, Te: semiconductors, Po: metal. Increasing cationic, base character from top to down: Te, Po are bases.

Binary compounds: Se2– selenides, Te2– tellurides, Po2– polonides. Most stable: alkaline, alkaline earth metal and compounds, and O/F/Cl/Br compounds. Decreasing thermal stability from top to down, e.g. hydrides: H2O > H2S > H2Se > H2Te > H2Po. Weaker chain forming and π-bond forming character than sulfur: CO2, CS2 π-bonded monomers, CSe2 polymer, CTe2 instable, CPo2 is unknown.

Increasing coordination number from top to down: SO2 gas, SeO2 solid chain (coordination number = 3), TeO2 solid chain (coordination number = 4), PoO2 solid (coordination number = 8).

Typical oxidation number = +4. Most abundant than in the sulfur chemistry.

40 János Rohonczy: Inorganic Chemistry I. 3.3. Selenium, Tellurium, Polonium

Table 19. Reactivity differences between the sulfur and other . Sulfur Se, Te, Po S + HNO3 → NO + H2O + H2SO4 Se + HNO3 → NO + H2O+ H2SeO3 H2SO4 + P4O10 → H3PO4 + SO3 H2SeO4 + P4O10 → H3PO4 + SeO2 + 1/2 O2 R2SO2, stable compounds R2SeO2, instable, hard to synthesize 4+ + 4+ S ion is unknown PO4 /Te stable ions Biological activity: All Se, Te, Po compounds are strong poisons, elemento-organic compounds are extremely strong ones. Inorganic selenium in trace quantity is medicine (natural in glutathione peroxydase enzyme). 150µg Se/day input is required.

BiologicalCoordination activity compounds Se, Te as coordination centers: similarity to sulfur, coordination number = 2-8. Se as ligand: e.g. µ2-Se2 compounds: µ2-Se2[(Co)3Fe]2 Te as ligand: non typical.

3.3.2. Se, Te, Po Polycations

2+ 2+ 2+ 2+ 2+ 2+ Similarity to sulfur: Se4 , Se8 , Se10 , Te4 , Te6 , or mixed [Se4Te2] Synthesis: Se8 + 5 SbF5 → [Se8][Sb2F11]2 + SbF3

3.3.3. Se, Te, Po Hydrides

H2Se (hydrogen selenide) Properties: colorless, smelly, toxic gas. Synthesis: a) H2 + Se → H2Se b) Al2Se3 + 6 H2O → 3 H2Se + 2 Al(OH)3 c) Fe/Se + 2 HCl → FeCl2 + H2Se.

H2Te () Properties: colorless, toxic gas. Synthesis: a) electrolysis of H2SO4 with Te cathode: Te + H2 → TeH2 b) Al2Te3 + 6 HCl → 3 TeH2 + 2 AlCl3 Te + H2 → NO REACTION

H2Po (hydrogen polonide) -10 Synthesis: Just in traces: 10 g: Po/Mg + HCl → PoH2 + H2 Common properties: In top-to-down order: acidity increases, MP/BP slowly increases, thermal stability decreases. 2+ Reaction: a) M + H2Se → MSe, b) H2Se/H2Te + O2 → SeO2/TeO2 c) H2Se + SO2 → S8, Se8, H2SO4, .... Metal Selenides, tellurides Occurrence: minerals, together with sulfides. Types: with IA/IIA elements in direct synthesis: colorless, water soluble salts – Na2Se, MgPo Polyselenides are less stable: Na + Se8 → Na2Se3 + ... With transition metals: non-stoichiometric compounds, low EN difference, alloy type

41 János Rohonczy: Inorganic Chemistry I. 3.3. Selenium, Tellurium, Polonium compounds: Ti0,9Se, Ti3Se4, Ti5Se8. Properties: special optical, electric, thermo-electric etc. properties. New high-tech materials. 3.3.4. Se, Te, Po Halides Various compositions similar to the sulfur. Even with iodine in higher . Table 20. Physical properties of Se, Te and Po halides. Oxidation No. F Cl Br I +1/+2 Se2F2/SeF2 Se2Cl2 Se2Br2 – in matrix yellow liquid red liquid +4 SeF Se Cl Se Br – Se 4 4 16 4 16 colorless liquid white solid red solid +6 SeF6 – – – colorless gas <1 – Te2Cl/Te3Cl Te2Br Te2I gray solid gray solid gray solid +1/+2 – (TeCl2) (TeBr2) Te4I4 ? ? black Te +4 TeF4 Te4Cl16 Te4Br16 Te4I16 colorless liquid yellow solid yellow solid black solid +6 TeF6 – – – colorless gas +2 – PoCl2 PoBr2 (PoI2) red solid brown solid decomposes Po +4 – PoCl4 PoBr4 PoI4 yellow solid red solid black solid

Mixed halides: TeBr2Cl2, PoBr2Cl2, etc.

Synthesis: 2 Se + Cl2 → Se2Cl2 3 Te + Cl2 → Te3Cl2 SeO2 + SF4 → SeF4 + SO2 Te + 2 Cl2 → TeCl4 Se + 3 F2 → SeF6 Te + 3 F2 → TeF6 TeCl4 + F2 → TeClF5 + ClF Complexes: TlF + SeF4 → [Tl][SeF5], 2 NaCl + TeCl4 → Na2[TeCl6]

3.3.5. Oxohalides

SeOF2 / SeOCl2 / SeOBr2 (selenyl fluoride, etc.) Structure: similar to SOCl2 Synthesis: SeO2 + SeCl4 → 2 SeOCl2

+ – Properties: 2 SeOCl2 → SeOCl + SeOCl3 . Dielectric constant = 46 (polar) SeO2F2 ( difluoride). Synthesis: SeO3 + SeF4 → SeO2F2 F5Se–O–SeF5, F5Te–O–TeF5, F5Se–O–O–SeF5 Synthesis: SeO2 + F2/N2 → F5Se–O–SeF5 Pseudohalides Properties: stable compounds: Se(CN)2 / Se2(CN)2 / Se(SCN)2 / Te(CN)2

3.3.6. Oxides

SeO / TeO Properties: just in flames

PoO Properties: black solid

42 János Rohonczy: Inorganic Chemistry I. 3.3. Selenium, Tellurium, Polonium

SeO2 Properties: white crystal, MP = +340 °C, water soluble, easy to reduce Synthesis: Se + O2 → SeO2. Reaction: 3 SeO2 + 4 NH3 → 3/8 Se8 + 2 N2 + 6 H2O

TeO2 Properties: 2 variants, yellow and white, at 730 °C red liquid. Synthesis: a) Te + O2 → TeO2 b) H2TeO3 → TeO2 + H2O (by desiccators)

PoO2 Properties: yellow solid → brown → Po + O2 (500 °C).

SeO3 Properties: white, hygroscopic solid. Structure: Se4O12, cyclic tetramer, Se–O–Se ring. Synthesis: 2 SeO2 + O2 → 2 SeO3

TeO3 Properties: 2 variants, insoluble in water, strong oxidizer, solution in base: – 6– TeO3 + 6 OH → TeO6 + 3 H2O.

3.3.7. Hydroxides – Oxoacids

H2SeO3 / H2TeO3 ( ) Properties: white . Synthesis: a) SeO2 + H2O → H2SeO3 b) 3 Se + 4 HNO3 + H2O → 3 H2SeO3 + 4 NO c) TeCl4 + 3H2O → H2TeO3 + 4 HCl – Chemical properties: Existing acidic salts (HSeO3 ) H2SeO3 + H2O2 → H2SeO4

Exotic acids: H2Se2O / H4Se3O11 / H2SeO5.

PoO(OH)2 (polonium oxide dihydroxide) Properties: PoO(OH)2 + 2 KOH � K2PoO3 + 2 H2O (amphoteric).

H2SeO4 () Properties: colorless viscous liquid. Synthesis: Se + 3 Cl2 + 4 H2O → H2SeO4 + 6 HCl. Reaction: 2 Au + 6 H2SeO4 → Au2(SeO4)3 + 3 H2SeO3 + 3 H2O.

Te(OH)6 (H6TeO6) (orthotelluric acid) 2– 2– Structure: isoelectric with [Sn(OH)6] and [Sb(OH)6] ions. Properties: white crystalline powder, dissociation in several steps. Salt: Na6TeO6⋅4 H2O. Synthesis: 5 Te + 6 HClO3 + 12 H2O → 5 H6TeO6 + 3Cl2 Properties: H6TeO6 + 3 SO2 → Te + 3 H2SO4 (oxidizer) H6TeO6 → Te–O–Te–O (polymerization)

HO–Se(O)–OOH (peroxoselenic acid)

43 János Rohonczy: Inorganic Chemistry I. 4.1. Nitrogen

4. (15-th column) N, P, As, Sb, Bi Valence shell: ns2np3 These elements exhibit big difference in the thermal coefficient of the electric conductivity. Nitrogen, phosphorus and arsenic are typical non-metals, but in the latter case metallic α-As is known as well. α-arsenic, antimony and bismuth are metals, but there are big differences in their metallic behaviors. The thermal coefficients of the electric conductance of rhombohedral, stable (25oC, 1 bar) α-As is gently negative. Stable allotrope of Sb is typical metal at RT. (But it is easy to prepare the non-metallic Sb4 allotrope as well.). Bismuth is metal. Non-metal allotrope is unknown in this case.

Table 21. Some physical parameters of the elements of 15th column. MP (°C) BP (°C) State Electric conductivity E.N. 1st ioniz. potential (V) (1/cmΩ) N2 -210 -196 Gas – 3.04 14.5 -11 P4 +44 +280 Molecular lattice 1⋅10 2.19 10.5 α-As +817 613 Metal 3.5⋅104 2.18 9.8 (28 bar) (sublimates) α-Sb 630 1750 Metal 2.9⋅104 2.05 8.6 α-Bi 271 1560 Metal 8.7⋅104 2.02 7.3

4.1. Nitrogen (N)

Valence shell: 1s22s22p3 Properties: N2 colorless, odorless, tasteless, inert gas. Discovery: Rutherford, Scheele, Cavendish – non-burning part of the air (1772). Naming: "nitron gennan" – forms niter / "azotos" – no life (still used in names: aso-, diaso-, - azides). Known: NH4Cl (Herodotos, 5th century B.C.), niter, . Occurrence: 78% in air, in Earth's crust low (19 ppm, just 33.), Minerals: KNO3 / NaNO3. – – N2-cycle: N2 fixing bacteria, lots of : NO3 , NO2 . Laboratory synthesis: a) NH4NO2 → N2 + 2 H2O b) 2 NH3 + CuO → N2 + 3 Cu + 3 H2O, c) (NH4)2Cr2O7 → N2 + Cr2O3 + 4 H2O.

Other: 2 NaN3 → 2 Na + 3 N2 8 NH3 + 3 Br2 → N2 + 6 NH4Br(aq)

Industrial synthesis: fractional distillation of liquid air (see O2), Quantity > 30 Mt/year. Application a) inert gas: steel industry, oil refining, laboratories, industrial packing technology...

b) liquid N2: cooling liquid, processes on rubber, refrigerate (blood, etc.). Nuclear properties: Stable isotopes: 14N 99,6% (I = 1) and 15N 0,4% ( I = 1/2). Radioactive ones: 13N, 16N

Enrichment of 15N up-to 99,5%. Synthesis: fractional distillation of liquid NO, or by using the 15 14 – 14 15 – partition equilibrium: NO(gas) + NO3 (aq) � NO(gas) + NO3 (aq) Structure: :N≡N: diamagnetic, diatomic molecule: inert, dissociation over 4000K, Hdissociation = +945 kJ/mol (high) János Rohonczy: Inorganic Chemistry I. 4.1. Nitrogen

Reaction at R.T: N2 + 6 Li → 2 Li3N.

Reaction at high temperature: H2 → NH3 / C → (CN)2 / IA, IIA, B, Al, Si, Ge, Ln, An → nitrides. * Dissociation: electric discharge at low pressure: atomic N. Recombination: 2 N → N2 2+ 2+ N2 absorption at R.T.:: [Ru(NH3)5(H2O)] + N2 � [Ru(NH3)5N2] + H2O. Reversible.

1 2 N2-ligand in complexes: η -complexes: M←:N≡N: η -complex: :N≡N: bonded with π-bond

4.1.1. Nitrides, azides and nitrido compounds

N3– () Structure: a) ionic: Li3N, Be3N2 (IA/IIA), b) covalent: (CN)2, P3N5, S4N4, (BN)x c) interstitial: TiN, TaN, UN (glittering opaque, very hard, inert, high MP, electric conductors.)

Synthesis: 3 Ca + N2 → Ca3N2, 3 Ca + 2 NH3 → Ca3N2 + 3 H2 3 CaO + 3 C + N2 → Ca3N2 + 3 CO, 3 CaCl2 + 3 H2 + N2 → Ca3N2 + 6 HCl.

– + Complexes: nitrido ligands: [O3Os≡N] or [N(HgCH3)4]

– N3 (azide): NaN3 linear ion. AgN3: increasing character → explosive.

4.1.2. Nitrogen hydrides

NH3 (ammonia) Properties: colorless, pungent, toxic gas. Well soluble in water, easy to condense (on 3-4 bar). MP = -78 °C, BP = -34 °C. Liquid in cylinders at RT. – + Laboratory synthesis: a) NH4Cl + NaOH → NH3 + Cl + Na + H2O, b) 3 Ca + N2 → Ca3N2, and Ca3N2 + 6 H2O → 3 Ca(OH)2 + 2 NH3 Industrial synthesis: Haber-Bosch (1913 - BASF). N2 + 3 H2 � 2 NH3 Exothermal, moderate high temperature, decreasing volume, equilibrium is pushed by high pressure. catalyst / 400 °C / 200 Bar / efficiency: 15% (condensation/recirculation). Produce: 100 Mt/year. Produce of H2: CH4 + H2O → CO/CO2 + H2, Production of N2: air + H2 → H2O + N2 Removal of CO: CO + H2O → CO2 + H2 Removal of CO2: CO2 + H2O + K2CO3 → 2 KHCO3 CO traces must be removed (reacts with Fe catalyst). Byproducts: noble gases.

Structure: AX3E, trigonal pyramid. Ammonia inversion frequency: 23.79 GHz. (Atomic clock) In liquid phase few H-bond, low viscosity, not too high . Diver solution: NH4NO3/NH4I/NH4SCN + NH3: very high boiling point over RT, H-bonds. + – -30 Autoprotolysis: 2 NH3 � NH4 + NH2 Dielectric constant = 22, K=10 , neutral pH = 15. Brönstedt acidity in liquid ammonia: + – – 3– acid: NH4 (NH4Cl), bases: NH2 (NaNH2), NH2 (Na2NH), N neutralization: NH4NO3 + KNH2 → KNO3 + 2 NH3

45 János Rohonczy: Inorganic Chemistry I. 4.1. Nitrogen

Reaction in liquid ammonia: Ba(NO3)2 + 2 AgBr → BaBr2 + 2 AgNO3 Special reactions: synthesis of dry AgNO3. Solubility of salts in NH3(liq.) can differ from their solubility in water.

Reaction with water: + – NH3 + H2O � NH4 + OH , pKbase = 4,74 , weak base. NH4OH molecule is unknown, 3 dimension H-bonded H2O⋅NH3 network. Other reactions: 4 NH3 + 5 O2 → 4 NO + 6 H2O 2 NH3 + 6 F2 → 2 NF3 + 3 (HF)2 NH3 + Cl2 → NH4Cl, NH2Cl, NHCl2, NCl3 NH3 + S → H2S + N4S4, NH3 + C(hot) → NH4CN + H2 + – NH3(liq.) + Na/Li/Ca/K → Na + e solvated (blue liquid) NH3(liq.) + Na/Li/Ca/K (+Fe catalyst) → H2 + NaNH2 (colorless solution). 2+ NH3 + O2 + Cu → [Cu(NH3)4] Application a) fertilizers: NH3, NH4NO3, carbamide (urea): (NH2)2CO, (NH4)3PO4, (NH4)2SO4 b) explosives: NH4NO3, nitro-glycerin, nitro-cellulose, TNT c) synthetic fibers: hexamethylenediamine, poly-amides, poly-urethanes d) refrigerant, buffer, other N-containing compounds.

N2H4 (hydrazine). Properties: colorless, water soluble, toxic, carcinogenic. Synthesis: 2 NH3 + NaOCl + glue → N2H4 + NaCl + H2O (1907, Raschig). 2+ 2– Purification: N2H4 + H2SO4 → [N2H6] [SO4] , recrystallization of the salt or distillation of base. Produce: 10,000 t/year. Application: a) rocket fuel (+ N2O4, Apollo mission: Lunar Module), b) water solution: reducing agent (mirror coating of Ag, Cu, etc.), c) remove of O2 from stream: N2H4 + O2 → N2 + 2 H2O, ion free process. d) organic chemical industry. Structure: H2N–NH2, twisted molecule: dipolar moment = 1.85 D. + – -6 Reaction: in water double base: N2H4 + H2O � N2H5 + OH , Kb = 10 + 2+ – -15 N2H5 + H2O � N2H6 + OH , Kb = 10 reducing agent: N2H4 + KIO3 + 2 HCl → N2 + KCl + ICl + 3 H2O

NH2OH (hydroxylamine). Properties: colorless, unstable, water soluble. + – Laboratory synthesis: HNO3 + 6 H + 6 e → 2 H2O + NH2OH, (on Hg/Pb anode) + – NH2OH +HCl → [NH3(OH)] Cl

Industrial synthesis: NH4NO2 + SO2 → [NH3(OH)]2SO4 + (NH4)2SO4

+ – Weak base: NH2OH + H2O � [NH3(OH)] + OH , similar to N2H4 and H2O2. Structure: H2N–OH, two conformers. Application: reducing agent, antioxidant.

HN3 (hydrogen azide). Properties: only in water solution, medium strong acid, salts: azides + – HN3 + H2O � H3O + N3 . – Structure: H–N=N=N, N3 pseudohalide ion. Synthesis: a) N2O + NaNH2 → NaN3 + H2O

b) NaNO3 + 3 NaNH2 → NaN3 + 3 NaOH + NH3

46 János Rohonczy: Inorganic Chemistry I. 4.1. Nitrogen

c) 3 N2O + 4 Na + NH3 → NaN3 + 3 NaOH + 2 N2

d) Pb(NO3)2 + 2 NaN3 → Pb(N3)2 + 2 NaNO3 Application: Ag-, Pb-, Cu-azides are explosives – used in detonators.

N2H2 (diimine) Structure: HN=NH, nonlinear, unstable.

N4H4 () Structure: H2N–NH–N=NH (only below 0 °C) → N2 + N2H4

4.1.3. Nitrogen halides

Structure: AX3E geometry: pyramid. Stability: NF3 (stable) > NCl3 (unstable) > NBr3 (explosive) > NI3 (explosive).

NF3 () Properties: Most stable (discovery: 1928), colorless, odorless gas. Insoluble in water, diluted acids and diluted bases. Non-reactive. Synthesis: a) 4 NH3 + 3 F2 → NF3 + 3 NH4F (N2F4, N3F is), b) NH4F/HF → NF3 (Electrochemical oxidation on Pt anode). Application: fluorinating agent at high temperature: 2 NF3 + Cu → N2F4 + 2 CuF

N2F4 (dinitrogen tetrafluoride) Properties: Colorless, reactive, strong fluorinating agent + – Reaction: S + N2F4 → SF4 + SF5NF2 + …, AsF5 + N2F4 → [N2F3] [AsF6]

NCl3 () Properties: (1811) viscous, unstable, explosive liquid. Synthesis: NH4Cl → NCl3 (pH < 4, electrochemical oxidation on Pt anode) Reaction – hydrolysis: NCl3 + 3 H2O → NH3 + 3 HOCl (bleaching agent) oxidation: 2 NCl3 + 6 NaClO2 → 6 ClO2 + 6 NaCl + N2

NBr3 () Properties: explosive over -100 °C.

NI3⋅NH3 () adduct Properties: brown precipitate, dry powder is explosive to the touch . Structure: NI4 tetrahedrons, N–I–N–I bonds.

Synthesis: NH3 + I2 → NI3⋅NH3.

Nitrogen oxohalides

XNO ( halides) Structure: V-shape. Properties: reactive gases, FNO colorless, ClNO yellow, BrNO red. Synthesis: a) X2 + 2 NO → 2 XNO b) N2O4 + KCl → ClNO + KNO3 Reaction: XNO + H2O → HNO2 + HX XNO + R–OH → HX + R–O–N=O (alkyl nitrite)

XNO2 ( halides) Structure: planar triangle, reactive gases: X = F, Cl.

47 János Rohonczy: Inorganic Chemistry I. 4.1. Nitrogen

Synthesis: ClSO3H + HNO3 → ClNO2 + H2SO4 Reaction: ClNO2 + H2O → HNO3 + HCl

4.1.4. Nitrogen oxides and oxoacids

N2O (dinitrogen monoxide, nitrous oxide). Properties: BP = -88 °C, colorless, odorless, sweet, dormitive gas: "laughing gas". Structure: N=N=O, isoelectric with CO2, linear, non-reactive. Synthesis: NH4NO3 → N2O + 2 H2O (slow heating below 250 °C). Reaction: inert at 20 °C. No reaction with X2 / O3 / Na, but it is reactive at high temperature: N2O + H2 → N2 + H2O N2O + C2H2 → N2 + CO2 + H2O N2O → N2 + "O" N2O + NaNH2 → NaN3 + H2O N2O + H2O → no reaction, not real acid anhydride.

NO (nitrogen monoxide). Properties: BP = -152 °C, water insoluble, colorless gas, paramagnetic – free radical molecule. Synthesis: a) Cu + HNO3 → NO + Cu(NO3)2, the by-product NO2 is soluble in NaOH solution. b) KNO2 + KI + H2SO4 → NO + K2SO4 + H2O + 1/2 I2 Reaction: 2 NO + O2 → 2 NO2 (fast) 2 NO + X2 → 2 NOX 6 NO + 4 NaOH → 4 NaNO2 + N2 + 2 H2O (slow). 2+ Known ligand in complexes: [Fe(H2O)5NO] (nitroso ferrous sulphate, brown ring)

N2O3 () Properties: blue liquid, MP = -100 °C, at higher temperature decomposes: N2O3 � NO + NO2 Synthesis: a) 2 NO + N2O4 → 2 N2O3, b) 4 NO + O2 → 2 N2O3 Reaction: N2O3 + H2O → 2 HNO2 (nitrous acid), – – N2O3 + 2 OH → 2 NO2 (nitrite ion) + H2O.

NO2 / N2O4 ( / ) Properties of N2O4: MP = -11 °C, colorless solid, below R.T. colorless gas. At +100 °C: 2 NO2 � N2O4. Brown gas mixture of 90% NO2 : 10% N2O4 Structure: NO2 paramagnetic, odd number of electrons. One electron on a high energy antibonding orbital. Brown color. N2O4 diamagnetic, colorless. Long, weak N–N bond. In solid phase: planar O2N–NO2, in gas phase twisted geometry and isomerisation: O=N–O– NO2 Synthesis: a) 2 Pb(NO3)2 → 4 NO2 + 2 PbO + O2. Separation of gases by cooling. b) 2 NO + O2 � 2 NO2 Reaction: a) N2O4 + H2O → HNO2 + HNO3 (common anhydrate of two acids)

+ – b) N2O4 � NO + NO3 (heterolytic dissociation)

Acid–base reaction in liquid N2O4: NOCl + AgNO3 → AgCl + N2O4 Redox reaction: TiI4 + N2O4 → Ti(NO3)4 + 4 NO + 2 I2

48 János Rohonczy: Inorganic Chemistry I. 4.1. Nitrogen

N2O5 (). Properties: white crystal, light sensitive, real anhydrate of the HNO3 Synthesis: 2 HNO3 → N2O5 (dehydration by anhydrous P4O10) Reaction: N2O5 + H2O → 2 HNO3, N2O5 + H2O2 → HNO3 + HOONO2 Oxidation agent: N2O5 + Na → NaNO3 + NO2 N2O5 + I2 → I2O5 + N2 Thermal decomposition: N2O5 → NO2 + NO + O2 + – Heterolytic dissociation N2O5 � NO2 + NO3

4.1.5. Nitrogen oxoacids

Exotic acids HNO (), Structure: H–N=O, V-shape. Properties: reactive intermediate, ligand in complexes. H2N2O2 (), N2O is only its formal anhydride. Structure: HO−N=N−OH � ΗΟ−ΝΗ−Ν=O, known stable salts.

Synthesis: 2 NaNO3 + 8 Na/Hg → Na2N2O2 + 8 NaOH + 8 H2 Ag2N2O2 + 2 HCl → H2N2O2 + 2 AgCl. Reaction: decomposition: H2N2O2 → N2O + H2O.

H2N2O3 (hyponitric acid), unstable. Stable salt: Na2N2O3. Structure: HO−NH−NO2 H4N2O4 (nitroxyl acid), explosive, stable salt: Na4[O2N=NO2]. HOONO (peroxonitrous acid). Properties: strong oxidizer, unstable, no solid salts known. Synthesis: H2O2 + N2O3 HOONO2 (peroxonitric acid. Properties: explosive crystals, no solid salts known.

HNO2 (nitrous acid), its anhydrate is: N2O3. Synthesis: a) Ba(NO2)2 + H2SO4 → BaSO4 + 2 HNO2 b) NaNO2 + HCl → NaCl + HNO2 (not clean) Industrial synthesis: NO + NO2 + 2 NaOH → 2 NaNO2 + H2O, after that NaNO2 + diluted HCl / H2SO4 → HNO2 + Na-salt Structure: HO–N=O. + – Properties: weak acid: HNO2 + H2O � H3O + NO2 + – decomposition: 3 HNO2 � H3O + NO3 + 2 NO (disproportion) + – oxidation: HNO2 + H + e → NO + H2O – + – reduction: HNO2 + H2O → NO3 + 3 H + 2 e – Salts: NO2 (nitrite ion). Structure: V-shape. NaNO2 gently toxic, white crystal. Application: organic chemistry / industry – diazotization. Complexes: Metal-NO2 (nitro ligand), or Metal-O-NO (nitrito ligand).

HNO3 (nitric acid), its anhydrate is: N2O5. Structure: HO–NO2. Properties: strong acid, stable, anhydrous acid exists – strong oxidizing agent. Laboratory synthesis: KNO3 + cc. H2SO4 → KHSO4 + HNO3 Industrial synthesis: catalytic oxidation of NH3 with O2 (Challenge: N2 and N2O byproducts are thermodynamically more stable compounds than the NO). Technology: high O2 concentration, Pt/Rh catalyst / 1 ms contact time. Conversion: 96%. Produce: 1.5 Mt/year. Formal description of production of HNO3: NH3 + O2 → NO {+O2} → NO2 → {+H2O} → HNO3 Application: in 80%: NH4NO3 (mixture with CaCO3: fertilizer), pure: industrial explosive. in 5%: synthesis of cyclohexanone, synthesis of fibers, in 5% nitration: nitro-glycerin, TNT,

49 János Rohonczy: Inorganic Chemistry I. 4.1. Nitrogen nitro-cellulose, other: etching, rocket fuel, pyrotechnics. + – Reaction: HNO3 + H2O � H3O + NO3 (strong acid), + + – cc. HNO3 + cc. H2SO4 � NO2 + H3O + 2 HSO4 (nitration acid) cc. HNO3 + cc. HCl � Cl2 , ClNO (aqua regia: Au → HAuCl4) cc. HNO3 + Ag/Cu → AgNO3/ Cu(NO3)2 (aqua fortis)

– NO3 (nitrates) Structure: symmetrical planar anion. Properties: most of them are water soluble salts: KNO3, NaNO3, AgNO3, Pb(NO3)2 Thermal decomposition: various products: metal nitrites, metal oxides, or metals: 2 NaNO3 → 2 NaNO2 + O2 , 2 KNO3 → K2O + N2 + 5/2 O2 Pb(NO3)2 → PbO + 2 NO2 + 1/2 O2 , 2 AgNO3 → 2 Ag + 2 NO2 + O2 NH4NO3 → N2O + 2 H2O or NH4NO3 → 2 N2 + O2 + 4 H2O (explosion over 250 °C) – NO3 as ligand: several coordination numbers, M–O bonds.

H3NO4 (orthonitric acid). Properties: Free acid is unknown. Salts: orthonitrates: M3NO4 – 3– Synthesis: NaNO3 + Na2O → Na3NO4. Structure: tetrahedral NO4 , like PO4 Reaction: Na3NO4 + H2O → NaNO3 + 2 NaOH.

4.1.6. Sulfur nitrides

S4N4 (tetrasulfur tetranitride) Properties: crystalline. Light yellow at -30 °C, orange at 20 °C, dark red at 60 °C (thermo color). Most stable sulfur-nitrogen compound, but endothermic. Structure: ring. e– donor: Lewis base. Synthesis: 10 S + 4 NH3(liquid) � S4N4 + 6 H2S

Other syntheses: 6 S2Cl2 + 16 NH3 → S4N4 + S8 + 12 NH4Cl 6 S2Cl2 + 4 NH4Cl → S4N4 + S8 + 16 HCl – 2– 2– Reaction: S4N4 + 6 OH + 3 H2O → S2O3 + 2 SO3 + 4 NH3

S2N2 (disulfur dinitrogen) Structure: 4 atom ring, unstable Synthesis: S4N4 + 8 Ag → 4 Ag2S + N2 (at 250 °C, 1 mBar) S4N4 → 2 S2N2 (Ag2S catalyst)

S8-n(NH)n (sulfur imides) Structure: S8-type ring, S → NH isoelectric replacement: same Structure Synthesis: S4N4 + SnCl2 → Sn(NH)4 + SnCl4 (in benzene/ethanol solution), or S2Cl2 + NH3 → S8 + S7(NH) + ...(yellow, MP = +114 °C).

Known: S7NH, isomers of S6(NH)2, where the NH position is 1,3-, 1,4-, 1,5-… Cyclic, condensed, heterocyclic compounds: O / C / Si / P / As, SxNyOz, SxNyHlgz.

(SN)x Properties: fibrous, bronze shiny color, superconductor at very low temperature. (SNBr0,4)x Properties: electron conductor at RT.

50 János Rohonczy: Inorganic Chemistry I. 4.2. Phosphorus

4.2. Phosphorus (P) Valence shell: [Ne]3s2 3p3 Discovery: Brandt (1669) Name: Φωσφόρος – phosphoros, meaning "light-bearer" (Latin: Lucifer). Properties: several allotropes, P–P chains, stable oxidation number = +5, coordination number: 4 (and 5, 6). Occurrence: in Earth minerals: +5 oxidation state in Ca3(PO4)2 (phosphorite) and Ca5(PO4)3X (apatite), where X = F, Cl, OH. -3 oxidation state in iron meteorites (Fe,Ni)3P. In organism: DNS, RNS, ATP, phospholipids, tooth, bone, etc. Phosphate cycle: minerals → plants → animals → surface water → minerals. Phosphorus in surface water: fertilizers, detergents, pesticides. Synthesis: 2 Ca3(PO4)2 + 6 SiO2 + 10 C → 6 CaSiO3 + 10 CO + P4. Two reaction steps, but only both together result in elementary phosphorus: 1) 2 Ca3(PO4)2 + 6 SiO2 → 6 CaSiO3 + P4O10 2) P4O10 + 10 C → P4 + 10 CO Supplementary reaction using fluorapatite: CaF2 → SiF4 + {Na2CO3 + H2O} → Na2SiF6 And from Fe2O3 → Fe2P (ferrophosphorus)

Application: in 90%: P4O10 and H3PO4, others: P sulfides, chlorides and organic P compounds. Produce: > 1 Mt/year. Allotropy

α-P4 (white phosphorus). Structure: tetrahedral geometry, molecule lattice. Properties: MP = 44 °C, BP = 280 °C, white, soft, heavier than the water, soluble in: CS2 > PCl3 > POCl3 > SO2(liq.) > NH3(liq.) > benzene > fats. Reaction: very reactive, instable on air, flammable over 34 °C, toxic.

β-P4 Synthesis: from P4 vapor at -196 °C.

Pn polymers: several types and colors, all have atomic lattices built from P4 tetrahedral units. Insoluble, non-toxic, non-flammable allotropes. Most important ones:

Amorphous red phosphorus (discovery: 1848) Synthesis: heating of white phosphorus in inert atmosphere.

Crystalline red phosphorus allotropes are made from the amorphous one. Synthesis: P(red) → P4(gas) → P4(liq.) → α-P4(white), by cracking and condensation.

Violet phosphorus (Hittorf, 1865) Synthesis: P(red) is dissolved in molten lead. During the crystallization phase separation: violet, monoclinic elementary phosphorus. Black phosphorus Structure: 3D atomic lattice. Properties: thermally most stable, high polymerization degree, 3D atomic lattice, high density, semiconductivity. Several polymorphs: orthorhombic, rhombohedral, cubic crystal lattice. o First synthesis: from P(white) (P4) at 12000 Bar and 200 C (Bridgman, 1916). János Rohonczy: Inorganic Chemistry I. 4.2. Phosphorus

Isotopes: 31P is the only one stable isotope: precise mass, I = 1/2, good NMR nucleus. 32 P t1/2 = 14,5 day, β-radiation, trace analysis. Synthesis: 32S(n, p) → 32P (5 kg/year).

Chemical properties: Typical oxidation number: ±3, +5. Coordination number 3-6 (more over: 7-9). Binary compounds: almost with all elements (except: Sb, Bi, noble gases). Violent reactions: P4 + O2 / S8 / alkaline metals P4 + 5 O2 → P4O10 / P4O6 + hν at RT, chemiluminescence, exothermal reaction. P4 + X2 → PX3 / PX5 – – P4 + OH → PH3 + [H2PO2] (hypophosphite), P4 + S8 → P4S6 /.../ P4S10 P4 + Na → Na3P

Phosphides

Table 22. Some metal-rich, stoichiometric and phosphorus-rich metal phosphides. Metal rich Stoichiometric Phosphorus rich M:P > 1 M : P = 1 M : P < 1 Hard, metal shiny, high MP ZnS lattice P2 units, lower stability, lower MP atomic: TiP, ZrP, FeP, MnP Ni3P, T2P, Ru2P, Ni5P2 GaP, InP (LED), FeP2, OsP2, PtP2 WP, NbP, TaP (durable) ionic: Li3, Th3P4, Ca3P2

4.2.1. Phosphorus hydrides

PH3 (phosphine) PnHn+2, homologues: n = 1-6. Structure: pyramid, like NH3. Properties: less stable than NH3, colorless, toxic, garlic smelling gas. Laboratory synthesis: Ca3P2 + 6 H2O → 2 PH3 + 3 Ca(OH)2 (and P2H4) Industrial synthesis: P4(white) + 3 H2O + 3 KOH → PH3 + 3 K[H2PO2] Chemical properties: slightly soluble in water, pH ≈ 7, + – weak base: PH3 + HCl → PH4 + Cl (phosphonium cation) + – weak acid: PH3 + NH3(liq.) → NH4 + PH2 (phosphide anion) Burning: PH3 + 2 O2 → H3PO4 Good electron donor ligands: :PH3, and :PPH3 (triphenyl phosphine)

P2H4 (diphosphine) Properties: self burning gas. Synthesis: like PH3, and 2 PH3 → P2H4 + H2

4.2.2. Phosphorus halides

Types: PX3 / P2X4 / PX5, known mixed ligands and pseudo halides as well.

PX3 (phosphorus trihalides) Structure: AX3E.

Table 23. Physical properties of phosphorus trihalides.

PF3 PCl3 PBr3 PI3 Colorless gas Colorless liquid, 250,000 t/year colorless liquid red crystal

Synthesis: P4 + 6 Cl2 → 4 PCl3, 2 PCl3 + 3 CaF2 → 2 PF3 + 3 CaCl2

52 János Rohonczy: Inorganic Chemistry I. 4.2. Phosphorus

Reaction: PCl3 + Cl2 → PCl5 PCl3 + 1/2 O2 → POCl3 PCl3 + AsF3 → PF3 + AsCl3 PCl3 + 3 H2O → H3PO3 + 3 HCl (and H4P2O5) PCl3 + 6 NH3 → P(NH2)3 + 3 NH4Cl PCl3 + 3 Li–R → PR3 + 3 LiCl (R = alkyl-, aryl-) PCl3 + 3 R–Mg–Cl → PR3 + 3 MgCl2 (Grignard reaction in absolute Et2O) PCl3 + Ni(CO)4 → Ni(PCl3)4 + 4 CO.

P2X4 (diphosphorus tetrahalides) All halides are known. Most stable is P2I4.

PX5 (phosphorus pentahalides) Structure: AX5, trigonal bipyramid. 19 X-ray diffraction: Two kind of F atoms are in PF5; F NMR: only one kind of fluorine. Solution: fast Berry-rearrangement, pseudo-rotation. Synthesis: PCl3 + Cl2 → PCl5 (20,000 t/year). + – Properties: 2 PCl5 ↔ [PCl4] [PCl6] , mostly ions are in MeNO2. Reaction: PCl5 + H2O → 2 HCl + POCl3 {+3 H2O} → H3PO4 + 3 HCl PCl5 + 6 KF → K[PF6] + 5 KCl.

Phosphorus oxohalides

POX3 E.g.: POCl3 (phosphorus oxide trichloride, phosphoryl chloride) Physical properties: colorless liquid Laboratory synthesis: PI3 + S → PSI3, PCl5 + SO2 → POCl3 + SOCl2 Industrial synthesis: 2 PCl3 + O2 → 2 POCl3 Application: POCl3 + 3 ROH → (RO)3PO + 3 HCl ( esters), surfactants, non- inflammables. Produce: 10,000 t/year.

4.2.3. Phosphorus oxides / sulfides / oxosulfides, oxoacids

P4O6 () Synthesis: P4 + 3 O2 → P4O6 (white powder). Structure: cluster molecule.

P4O10 () Synthesis: P4 + 5 O2 → P4O10 Structure: cluster molecule. Properties: white powder, strong desiccative, frequently used in laboratories, real anhydrate of H3PO4. Application: P4O10 + 6 Et2O → 4 PO(OEt)3 (synthesis of phosphate esters).

P4O7-9 Synthesis: x P4O6 + y P4O10 → P4Ox (by melting).

P4S6-10 Synthesis: P4 + S8 → P4Sx (by melting).

P4S10 Synthesis: 4 Fe2P + 18 FeS2 → P4S10 + 26 FeS. Application: P4S10 + R–OH → (RO)2P(S)SH (dialkyldithiophosphoric acid esters) oil additive, fungicide, flotation of Cu, Zn ores). Produce: 250,000 t/year.

53 János Rohonczy: Inorganic Chemistry I. 4.2. Phosphorus

Matches: P4S10+ KClO3 + glass powder + ZnO + Fe2O3 + glue + water. On match box: Sb2S3 + Pred + Fe2O3 + gum arabic.

P4S5O5, P4S4O6 Structure: like upper one compounds, but S is in terminal position: S=P(O–)3

Phosphorus oxoacids Structure: P coordination number = 4 always. At least one O=P≡ and at least one HO–P(O)< bond, optionally H–P≡ bond and P–P / P–O–P / P–O–O–P / P–O–OH groups. Table 24. Phosphorus oxoacids and the oxidation numbers of phosphorus.

H4P2O2n-1 H3POn HPOn-1 Oxidation No. di- / pyro- ortho- meta- H3PO2 +1 H(H2PO2) H4P2O5 H3PO3 "HPO2" +3 H2(HPO3) not discovered diphosphorous acid orthophosphorous acid metaphosphorous acid H4P2O7 H3PO4 (HPO3)n +5 diphosphoric acid orthophosphoric acid metaphosphoric acid

H3PO2 (hypophosphorous acid) or phosphinic acid (salts: ). (+1) Structure: H[H2P O2], or O=P(H)2OH. Properties: Monovalent, unstable, known acid. – – Synthesis: P4 + 4 OH + 4 H2O → 4 H2PO2 + 2 H2, byproduct: P4 + 2 Ca(OH)2 + 2 H2O → Ca(HPO3)2 + 2 PH3. Salts: NaH2PO2 (stable, industrial reduction agent). – 2+ – 2– Application: Ni-coating on plastic surfaces: H2PO2 + Ni + 3 OH → HPO3 + Ni + 2 H2O.

H3PO3 (orthophosphorous acid) Structure: O=P(H)(OH)2, bivalent acid. Synthesis: PCl3 + 3 H2O → H2[HPO3] + 3 HCl (white crystals, MP = 70 °C) Derivatives: neutral salts, acidic salts, R–P(O)(OR')2, alkyl-compounds.

H4P2O5 (diphosphorous acid). Structure: H2[H2P2O5], only bivalent.

H3PO4 (orthophosphoric acid) Structure: O=P(OH)3, trivalent acid. Salts: primary phosphates – NaH2PO4, galvanic bath. Secondary phosphates – Na2HPO4 buffers (in milk powder). Tertiary phosphates – Na3PO4, TSP, sodium phosphate, water softening. Synthesis: a) P4 + 5 O2 → P4O10 {+ 6 H2O} → 4 H3PO4 b) Ca3(PO4)2 + 2 H2SO4 → Ca(H2PO4)2 + 2 CaSO4, (superphosphate). Application: Fertilizer (10 Mt/year), galvanic bath, chemical polishing of Al, detergents, pharmaceutical industry, food industry. Non-toxic, weak mineral acid: Coke, toothpaste.

H4P2O7 (diphosphoric acid) Structure: (HO)2P(O)–O–P(O)(OH)2 Properties: tetravalent, stronger than H3PO4.

– – – H5P3O10 () Structure: [Ad–O–P(O)(O )–O–P(O)(O )–O–P(O)(O )2]. Properties: pentavalent: Neutral salt: Na5P3O10, industrial detergent.

54 János Rohonczy: Inorganic Chemistry I. 4.2. Phosphorus

Natural occurrence: ATP – . High energy phosphates, exothermal 4– 3– 2– + hydrolysis: ATP + 2 H2O → ADP + HPO4 + H3O , ∆G = -41 kJ/mol.

[HPO3]n (metaphosphoric acid) – Structure: chain polymers or rings, glassy, [–O–P(O)(O )–]n

II and IV oxidation state acids (P–P bonded acids)

H2[H2P2O4] Structure: HO–P(O)(H)–P(O)(H)–OH, only one .

H4[P2O6] Structure: (HO)2P(O)–P(O)(OH)2, known structural isomers.

H3[HP2O6] Structure: (HO)2P(O)–O–P(O)(H)–OH, isomers.

Peroxo acids. Phosphorus is exclusively in +5 oxidation state.

H4[P2O8] Known salts. Structure: (HO)2P(O)–O–O–P(O)(OH)2

Polyphosphates –P–O–P–O–P– / –P–P–O–P– / and their combinations

4.2.4. Phosphorus nitrides and phosphorus organic compounds

Phosphazenes Synthesis: PCl5 + NH4Cl → (NPCl2)n + HCl, n = 3, 4, 5. Structure: alternating [–N=PX2–]n groups, 6 member ring, π-system is broken at P atom, ψ- aromatic. Reaction: bases, N is protonable. Ligand in complexes: [TiCl4(N3P3(Me6)], known : Cl / Ph / F / Me / OR / OAr. Properties: thermally very stable oligomers, no hydrolysis, no degradation, insoluble, non- swelling, slightly hydrophilic, non-irritant, non-flammable, has temperature independent elasticity. Application: Ph-, F-containing plastics: O-rings.

P(NH2)3 (phosphorus triamide), unstable, Structure: P–N σ-bond. Synthesis: PCl3 + 6 NH3 → P(NH2)3 + 3 NH4Cl.

P(NH)(NH)2 (phosphorus-amide-imide), Structure: HN=P–NH2 Synthesis: PCl3 + 5 NH3 → PN2H3 + 3 NH4Cl.

(PN)x (phosphorus nitride), Properties: amorphous polymer, very stable. Synthesis: x PCl3 + 4x NH3 → (PN)x + 3x NH4Cl.

Phosphorus organic compounds. P–Corganic bond.

P(CH3)3 / PPH3 (trimethylphosphine / ) Synthesis: PCl3 + 3 Li–Ph → PPH3 + 3 LiCl. Properties: very toxic, hydrolysis, they do not appear in the nature.

55 János Rohonczy: Inorganic Chemistry I. 4.3. Arsenic, Antimony, Bismuth

4.3. Arsenic (As), Antimony (Sb), Bismuth (Bi)

Valence shell: ns2np3

Table 25. History of As, Sb and Bi. Element As Sb Bi Compound and ancient times, ancient times, application As2S3 – orpigment Sb2S3 - black antimonite yellow pigment, epilator eyebrow pigment Discoverer Albertus Magnus Baril Valentine Agricola preparation As2S3 + soap → As roman / greek / persian greek: "anti monos" = against German: Name "az-zarnikh" = yellow monophobia(?) "wissmuth" = pigment Latin: stibium white mass Table 26. Occurrence of As, Sb and Bi compared to .

U As Sb Bi Ppm 2.2 1.8 0.2 0.008

Rare but well known compounds: enrich in sulfide minerals, low solubility.

Table 27. As, Sb and Bi minerals. As Sb Bi

As4S4 realgar Sb2S3 stibnite, antimonite α-Bi2O3 bismite As2S3 orpigment Fe/Co/Ni/Pb/Cu/Ag/Hg- Bi2S3 bismuthinite As2O3 arsenolite S/Sb ores, oxides (BiO)2CO3 bismutite FeAs2 / CoAs / NiAs elementary Sb elementary Bi (in Ag, Pb ores) FeAsS arsenopyrite CoAsS cobaltite elementary As Synthesis: As a) FeAsS → FeS + As(vapor) → As(gray). b) FeAsS + O2 → As2O3 / Fe2O3 / SO2 {+ carbon powder} → As(vapor) + CO/CO2

Sb a) Sb2S3 + 3 Fe → 3 FeS + 2 Sb(vapor), sublimates b) Sb2O3 + C → Sb(vapor) + CO/CO2 c) electrolysis, depends on the types of ores / contaminations / application requirements.

Bi a) 2 Bi2S3 + 9 O2 → 2 Bi2O3 + 6 SO2 { + C } → Bi + CO/CO2 b) Bi2S3 + 3 Fe → 2 Bi + 3 FeS

Application: As hardening of Pb-alloys (accumulator), As+Al/Ga/In semiconductors: LED, Laser-diodes, Hall-sensor, etc., glass industry, preserve of woods, pesticides, pharmaceutics (less important). Produce: As2O3 – 50,000 t/year. Sb Pb-alloys: accumulator, type-metal, bullet, tin-lead solder, semiconductors. 80,000 t/year.

Bi alloys: type-metal, melting fuse, X-ray contrast, Hall-sensor János Rohonczy: Inorganic Chemistry I. 4.3. Arsenic, Antimony, Bismuth

Physical properties: gray, metal shine, decreasing electric conductivity with increasing temperature – real, but weak electric conductor, rigid. Bi is strong diamagnetic, strong Hall- effect, increasing volume during freezing: +3,3% Atomic properties: Odd atomic number, few stable isotopes (several radioactive ones).

Table 28. As, Sb, Bi isotopes. Isotope Abundance Spin (I) 75As 100% 3/2 121Sb 57% 5/2 123Sb 43% 7/2 209Bi 100% 9/2 Allotropy: As As4 tetrahedral molecules in vapor, yellow sublimate, unstable, α-As solid, gray, metallic. Structure: hexagonal layers, ε-As solid, isomorphic with the rhombic, black phosphorus. Sb α-Sb isomorphic with α-As. Further 5 solid forms: yellow, black, "explosive", etc., Bi α-Bi isomorphic with α-As. Several allotropes, unknown structures. Chemical properties: amphoteric oxides As stable oxide layer at R.T., reactive at high temperature: As + air → As4O6/As4O10 As + X2 → AsF5 / AsCl5 / AsCl3 / AsBr3 / AsI3 (burning) As + 3 NaOH → Na3AsO3 + 3/2 H2 As + dil. HNO3 → H3AsO3 As + cc. HNO3 → H3AsO4 As + cc. H2SO4 → As4O6 Oxidation number = -3: Na3As: non ionic, intermetallic compound (higher than of Na and As). 3– 3– Composite ions: As7 , As11 , e.g.: K3As11 , cluster structures.

Oxidation number = +3: AsX3, covalent molecules. Oxidation number = +5: strong oxidation agent (in contrast to phosphorus).

Sb Similar to As chemistry, less reactive element. Reaction with strong oxidizing acids: Sb + cc. HNO3 → Sb2O5⋅x H2O Sb + aqua regia → SbCl5 Sb + cc. H2SO4 → Sb2(SO4)3

Bi Bi2O3 base anhydrate. Stable salts: Bi2(SO4)3, Bi(NO3)3. (Comparison to Sb2O3, As2O3 – amphoteric compound / P2O5, N2O5 acid anhydrates) Bi + (O2 / S8 / X2) → Bi2O3 / Bi2S3 / BiX3 – burning. Solution in strong oxidizing acids: see As, Sb.

3+ 2+ 5+ Cations: Bi5 (trigonal bipyramid), Bi8 (antiprism) clusters. Bi ion is unknown.

Binary arsenides, antimonides, bismuthides Intermetallic compounds with most metals: many stoichiometric and non-stoichiometric types. Mostly, simple to prepare them by melting of the components in inert atmosphere. Non ionic, conductive compounds, e.g. Na3As / Ca3Sb2

57 János Rohonczy: Inorganic Chemistry I. 4.3. Arsenic, Antimony, Bismuth

CoAs3 skutterudite diamagnetic, semiconductor mineral. As / Sb / Bi + Al / Ga / In → important semiconductors. Pb alloys: important intermetallic systems

4.3.1. Arsenic, antimony, bismuth hydrides

Thermal stability: decreasing in order: AsH3 > SbH3 > BiH3 → Bi + 3/2 H2. + Properties: toxic, colorless, smelly gases. Neutral compounds: AsH4 is unknown. Synthesis: a) As2O3 + 6 Zn + 12 HCl → 6 ZnCl2 + 3 H2O + 2 AsH3 and AsH3 → As + 3/2 H2 (Marsh-test, is running with Sb as well), b) 4 AsCl3 + 3 LiAlH4 → 4 AsH3 + 3 LiCl + 3 AlCl3 (or with NaBH4). Reaction: 2 AsH3 + 3 O2 → As2O3 + 3 H2O, similar reaction with S, Se. SbH3 + Si → Si(Sb) + H2, n-type semiconductor layer. Known other hydrides and alkyles: BiH3, As2H4, Bi(CH3)3.

4.3.2. Arsenic, antimony, bismuth halides

AX3 (trihalides) A = As, Sb, Bi, X = F, Cl, Br, I (all known). Properties: AsF3 (colorless gas), AsCl3 (colorless liquid), AsBr3 (colorless crystal), AsI3(red crystal), SbF3/BiF3 (liquid),.... SbI3/BiI3, (dark crystal). Structure: AX3E, AsI3 covalent molecule, BiI3 ionic salt. Synthesis: M2O3 + 6 HF → 2 MF3 + 3 H2O (cc. H2SO4 + CaF2) M2O3 + 6 HCl → 2 MCl3 + 3 H2O (cc. HCl / H2SO4 + NaCl) Reaction: MX3 + 3 H2O → H3MO3 + 3 HX (hydrolysis) AsX3 + 3 ROH + 3 NH3 → (RO)3As + 3 NH4Cl (alcoholysis) SiCl4 + SbF3 → SiCl3F / SiCl2F2 / SiClF3 + SbCl3 (fluorination) + – 2 SbF3 � [SbF2] [SbF4] (autodissociation)

AX5 (pentahalides) AsF5 (gas) / SbF5 (liquid) / BiF5 (solid) and SbCl5 (liquid). Synthesis: 2 M + 5 F2 → 2 MF5, SbCl3 + Cl2 → SbCl5 (BiCl5 is unknown), +5 oxidation number is not preferred, f-field contraction after . Exception: BiF5, where oxidation number = +5. F has highest EN, Bi–F bond with Bi s2 electrons. Structure: MX5 molecule, trigonal bipyramid. SbF5 is very viscous, [SbF]n F-bridged polymer. Properties: Very strong fluorinating agents, strongest one is BiF5 BiF5 + H2O → O3 / OF2 (explosion) BiF5 + UF4 → UF6 + BiF3 BiF5 + Br2 → BrF / BrF3 / BrF5 + BiF3 BiF5 + Cl2 → 2 ClF + BiF3

Mixed halides

Synthesis: MX3 + M'Y3 � MX4Y + M'XY4 � etc. (slow ligand exchange)

MX5 + M'Y5 � MX4Y + M'XY4 � etc. (fast ligand exchange through ion pair) + – MX5 + M'Y5 → [MX4] [MY5X] , (ion pair). + – Application: SbF5 + 2 HF → [H2F] [SbF6] (super acid). Lower oxidation state halides

As2I4 Properties: stable. Reaction: by heating: 3 As2I4 → 4 AsI3 + 2 As.

58 János Rohonczy: Inorganic Chemistry I. 4.3. Arsenic, Antimony, Bismuth

Complexes + – + – [ICl2] [SbCl6] / [AsF2] [SbF6] (ion pairs, salts),

F2SbF→SbF5 (adduct) often F, Cl, O bridges.

Oxo halides NOX analogues: Increasing stability: AsOF < AsOCl < SbOCl < BiOCl < BiOI. Synthesis: BiCl3 + H2O → BiOCl + 2 HCl.

POX3 analogues: Stable AsOF3, AsOCl3. 2– Complexes: [Sb2OCl6] (F, Cl, O bridges).

4.3.3. Arsenic, antimony, bismuth oxides and sulfides

As2O3 (), several alterations. Laboratory synthesis: a) 4 As + 3 O2 → As2O3 (only till +3 oxidation number), b) 2 AsCl3 + 3 H2O → As2O3 + 6 HCl, Industrial synthesis: 2 FeAsS + 3/2 O2 → FeS + As2O3 Structure: in vapor: As4O6 molecules, in solid phase: higher coordination number, atomic lattice. Reaction: As2O3 + 3 H2O � H3AsO3 (amphoteric): – 3– in bases: H3AsO3 + 3 OH � AsO3 + 3 H2O + 3+ in acids H3AsO3 + 3 H3O � As + 6 H2O Application: synthesis of acids, esters, AsH3, As halides, As sulfides, As organic compounds.

Sb2O3 () Synthesis: like As2O3 Reaction: Sb2O3 + 3 H2O → 2 "H3SbO3" → HSbO2 + H2O (only meta acid is stable) – – HSbO2 + OH � SbO2 + H2O (amphoteric: in base: meta-antimonate anion) + + HSbO2 + H3O � SbO + 2 H2O (in acid: antimonyl cation)

Bi2O3 / Bi(OH)3 (bismuth trioxide / bismuth ), 3+ – Synthesis: Bi + 3 OH → Bi(OH)3 → Bi(O)OH → Bi2O3 + – Bi(NO3)3 + H2O � Bi(O)NO3 + 2 H + 2 NO3 (very low solubility) BiO+ (bismuthyl cation)

As2O5 (). Synthesis: As2O3 + cc. HNO3 → 2 H3AsO4⋅H2O + NO. Reaction: 2 H3AsO4 → As2O5 + 3 H2O (with P4O10 desiccative) 2 H3AsO4 → As2O3 + H2O + O2 (at 300 °C)

Sb2O5 (). Synthesis: SbCl5 + 6 H2O → H3SbO4⋅2 H2O + 5 HCl, + – Structure: H [Sb(OH)6] (crystalline, monovalent acid)

Bi2O5 (bismuth-pentoxide). Synthesis: 3 Bi2O3 + 2 KClO3 → 3 Bi2O5 + 2 KCl. III V Mixed oxides. Synthesis: melting: Sb2O3 + Sb2O5 → Sb Sb O4 (like with phosphorus)

59 János Rohonczy: Inorganic Chemistry I. 4.3. Arsenic, Antimony, Bismuth

Arsenic antimony, bismuth sulfides

As4S4 (realgar, tetrahedral cluster with two As–As bonds). Synthesis: 2 As2O3 + 9/8 S8 → 2 As2S3 + 3 SO2 (yellow orpigment)

4.3.4. Elemento organic compounds

Arsenic organic compounds: usually As replaces the nitrogen, strong poisons. C6H5As (arsenobenzene) Structure: like pyridine

Atoxyl(R), Thomas (1905) against sleeping sickness (encephalitis lethargica) p-H2N–C6H4–As(O)(OH)(ONa)

Salvarsan, Ehrlich (1909) against syphilis.

HO As As OH

H2N NH 2 Sb, Bi organic compound: similar compounds, less studied topic.

60 János Rohonczy: Inorganic Chemistry I. 5.1. Carbon

5. (14-th column) C, Si, Ge, Sn, Pb Valence shell: ns2np2 Carbon, silicon and germanium are non-metals, tin and lead are metals, characterized by the temperature dependence of their electric conductivity, and their crystal lattice. Note: the stable graphite is good electron conductor, silicon and germanium are native semiconductors.

Table 29. Physical properties of stable allotropes. Element MP (°C) BP (°C) Crystal lattice Electric conductivity (1/cmΩ) C 3500 4827 graphite lattice 6⋅102 Si 1410 2355 diamond lattice 2.5⋅10-6 Ge 937 2830 diamond lattice 1.5⋅10-2 Sn 232 2270 metal 9.2⋅104 Pb 328 1740 metal 4.8⋅104

Chain forming character: (C–C)n >> (Si–Si)n ≈ (Ge–Ge)n > (Sn–Sn) ≈ (Pb–Pb). Inertness of s2 electrons is increasing in top-down direction: C (Oxidation number = +4, in 2+ exotic :CR2 and Ox. No.= +2) → at lead the most stable is Pb . Reason: in the case of 2nd row rs ≈ rp, from 3rd period rs < rp (rs "belongs to the core").

5.1. Carbon (C)

Valence shell: 1s2 2s22p2 Occurrence: Elementary forms: graphite, diamond, graphene, fullerene (in few minerals, in soot, asteroids). Coal – not elementary carbon, but poly-condensed aromatic compounds. Natural compounds: CO2 (CO2-cycle, greenhouse effect), carbonates (limestone: CaCO3, dolomite, Ca/Mg carbonate). History: Egypt: hieroglyph, India: soot containing ink. 1564 – pencil. 1779. Sheele: graphite is carbon, 1796. Tennant diamond is also carbon. Name: carbon(Latin) – wood coal, γράφω (Greek) – to write, diamond (Greek) – transparent and hard. Allotropy: α-graphite, β-graphite, diamond, graphene, lonsdaleite, chaoite, carbon-VI and other LAC's (Linear acetylenic carbon), fullerenes, nanotubes, glassy carbon, carbon nanofoam.

Graphite Structure: lattice of hexagonal planar layers. α-graphite: ABAB layers, hexagonal, dC–C = 142 pm. dlayer = 335 pm. Bond order = 1.33. β-graphite: ABCABC layers, rhombohedral lattice. Properties: density 2.3 g/cm3, black, metal shine, anisotropic properties: splits to layers, orientation dependent electric conductivity: in layer 103-4 cm-1Ω-1, in perpendicular direction ca. 1 cm-1Ω-1. János Rohonczy: Inorganic Chemistry I. 5.1. Carbon

Occurrence: in metamorphic deposits purity: 25-60%. Purification: flotation, washing with HCl, heating in vacuum. Synthesis: SiO2 + C → SiC → Si(vapor) + C(graphite) Application: electrode, collector brush, steel industry - casting mould, lubricant, break pad, pencil, carbon fiber, carbon filaments, neutron moderator. Produce: 500,000 t/year, obtain 300,000 t/year. Diamond Structure: diamond lattice with tetrahedral units, dC–C = 154 pm (only σ bonds). Properties: hardest natural material, colorless, density: 3.5 g/cm3 (high), high MP = 4000 °C, high refraction, good thermal conductivity (6x more than ), bad electric conductivity (semiconductor). Metastable at 1 Bar, phase transfers at high temperature into graphite. It burns on air over 600 °C. Occurrence: volcanic crater. Mass unit: 1 metric carat = 0.2 g. Synthesis: High pressure, high temperature, slow crystallization: small, black crystals for industry. Production: crushing, washing, separation on fatty endless belt, 5 t/year jewel, 12 t/year industrial quality. Application: industrial: cutting, drilling, polishing; and jewel. Graphene Structure: Monolayer graphite. Lonsdaleite Structure: Tetrahedral geometry, hexagonal lattice. Occurrence: asteroid impacts, artificial synthesis.

Chaoite White carbon, natural mineral and artificial product.

Carbon-VI, and other LAC (linear acetylenic carbon, 1972) Structure: under study.

Fullerenes (isolated ones: C2n, n = 30-48). Discovery: 1985. Electric arc between graphite electrodes in He atmosphere, Mass Spectrometric (MS) analysis.

Name: C60 geometry: carbon atoms on a ball-like polyhedron surface (Buckminster Fuller architect's "Montreal Biosphère"). Structure: C60 has =C< units form five- and six-member rings, only 6-member rings are around the five member ones. Synthesis: 1) Arc between graphite electrodes in He inert atmosphere. 2) Vacuum evaporation of the black carbon deposit. This is the fullerene soot. 3) Extraction of the mixture of fullerene molecules. 4) HPLC separation of pure C60 and C70 fragments as major components.

Structure: Fullerenes form molecule lattices with 3 center bonds. They are Wade-type electron deficient clusters. Symmetry and color of pure fullerenes: C60 – Ih (dark red or black), C70 – D5h (black), C76 – D2 (yellowy green), C78 – C2v, D3 (brown), C82 – C2, C2v, C3v, D2, D2d (greenish yellow), a C84 – D2, D2d (yellowy green). Only C76 with D2 symmetry has optical isomers.

62 János Rohonczy: Inorganic Chemistry I. 5.1. Carbon

3 Physical properties: density 1.7 g/cm , insulators, soluble in aromatic solvents. C60 has only one 13C NMR signal. Artificial amorphous and theirs application Coke – metallurgy Soot – rubber industry Activated carbon (bone-, wood-, walnut-, almond-coal): sugar industry, air- and water- purification, catalysts. Glassy carbon – electrode.

Mineral coal is not elementary carbon, it is mixture of organic compounds with low water content.

Atomic properties of the carbon. Isotopes: 12C 98,89% (atomic mass unit = mass of 12C isotope/12). 13C 1.2%, nuclear spin, I = 1/2 (NMR active). 14 -10 C 10 % radioactive, β−radiation, t1/2 = 5730 year. 14 14 14 C is used in radiocarbon dating: N + n → C + p → CO2 → living body. 14C containing compounds are trade products for isotope tracing.

Chemical properties of carbon

Graphite chemistry

C6(COOH)6 (mellitic acid or graphitic acid) Synthesis: Cgr + cc. HNO3 → C6(COOH)6

C6(O)x(OH)y (graphite oxides) Properties: lemon-yellow, solid. Synthesis: Cgr + KClO4 → C6(O)x(OH)y

C3.6-4F (graphite fluorides) Synthesis: Cgr + F2/HF → C3.6-4F (at R.T.) Structure: like graphite, F-insertion.

(CFx)n (graphite monofluoride) Synthesis: Cgr + F2 → (CFx)n (at 400 °C) Structure: cyclohexane like, perfluorinated, fluorine ligands in axial position, no C=C bond. Properties: x = 0,7 ...1 color is lightening, electric conductivity is decreasing.

CnF2n+2 (perfluoro-) CF4, C2F6, C5F12 etc. Synthesis: Cgr + F2 → CnF2n+2 Reactivity of the graphite: Cgr + X2 → no reaction even at high temperature (X = Cl, Br, I). Cgr + H2 → CnH2n+2 (Ni catalyst) Cgr + O2 → CO/CO2 Cgr + S8 → CS2 Cgr + Si → SiC Cgr + B → B4C (boron carbide) Cgr + MO → M + CO (reduction of metals)

63 János Rohonczy: Inorganic Chemistry I. 5.1. Carbon

Intercalation compounds of graphite Guest atoms are hosted between two layers of the graphite: Compounds with alkaline metals: Cgr + K → (C8K)x, bronze color.

Table 30. Interlayer distances of alkaline metal – graphite compounds.

Cgr C8K C8Rb C8Cs dlayer (pm) 335 540 561 598 Properties: high electric conductivity, which decreases from left to right. Different magnetic properties. Reaction: violent reactions with air and water. Derivatives: other metals, halogen atoms, metal halides: C8K + MX4 → C8M (MX4 halide) + – + – Guest + graphite atoms build common electron system: Cn + M + e ⇔ M Cn

Fullerene chemistry – – 6– Reactions: Reversible electrochemical reduction: C60 + e → C60 → ... C60 buckid-ions. Strong oxidizer: C60 oxidizes the O2, which is bonded by epoxy bond, increasing the number of members in the cluster: C60O (at high temperature C60 is burning in O2 forming CO). Halogens partly form endohedral compounds (C60Br2 / C60Br4). Fluorine oxidizes the fullerene and forms exohedral C60F50. Transition metals are bonded in oxidized form giving + intercalation compounds with exohedral connections ( C60M , M = Fe, Co, Ni, Cu, Rh), and + bis-fullerene-complexes: [(C60)2Ni] . Endohedral compounds Simple synthesis: C60 + Fe(CO)5 → Fe@C60 + 5 CO More complicated methods: LaCl3 + graphite → La@C60 Known endohedral compounds: La@C82, U@C82, Sc3@C82. Notes: 1. endohedral compounds are related to centered cluster structures, where the guest atoms increase the stability – Wade's PSEPT theory. + 2. The neutral endohedral [He@C60] is more stable than the [He@C60] . Increasing the member number of the cluster: C60 + PH2C=N=N → C61PH2. Occurrence: in and benzene soot (0,003-9 %), in exhaust of Diesel motors, but never in sugar + cc. H2SO4 reaction. Natural occurrence: shungite ( C containing mineral: both C60 and C70) References: Fullerene Science and Technology. An International and Interdisciplinary Journal. (Editor: T. Braun, Eötvös University Budapest) Marcel Dekker, Inc., New York.

5.1.1. Carbides

δ+ δ– Carbides are binary carbon containing compounds: An Cm Synthesis: M + C → (2000 °C) | MO + C | M + CnH2n+2 | M + C2H2 + NH3

Carbides with atomic crystal lattice SiC (silicon carbide = carborundum) Synthesis: SiO2 + 3 C → SiC + 2 CO → Si + C(graphite) (synthetic graphite). Application: in furnaces (till 1400 °C), cutting tools, structural material, LED, nuclear material, etc.

64 János Rohonczy: Inorganic Chemistry I. 5.1. Carbon

B4C (boron carbide) Synthesis: B2O3 + 7 C → B4C + 6 CO Properties: Neutron absorber, harder than diamond.

Ionic carbides – 2– C4 (carbide ion) in Al4C3 and C2 (acetylic ion) in CaC2. Carbon reacts with s- and f-block metals (Ce, La, U, Ln, An):

CaC2 (calcium carbide) Synthesis: a) CaO + 3 C → CaC2 + CO (traditional, 2000 °C), b) Ca + 2 C2H2 → CaC2⋅C2H2 + H2 (-40° C, in liquid NH3) CaC2⋅C2H2 → CaC2 + C2H2 (+300 °C) Application: CaC2 + 2 H2O → Ca(OH)2 + C2H2 (acetylene) CaC2 + N2 → CaN–C≡N: (Ca-cyanamide, and H2N–C≡N: is cyanamide)

Interstitial carbides. Carbon reacts with transition metals. Structure: non stoichiometric compounds, carbon atoms in the holes of the metal lattice, metal-carbon bonds. Properties: metal like electric conductivity, hardness, high MP. E.g. TiC (armouring steel), V2C, Mn3C, Mn15C4, Fe3C (cementite – steel component).

5.1.2. Hydrocarbons and carbon halides

Homologous series: CnHm, (several series, see organic chemistry)

CH4 (methane) There is no C–C bond so it is inorganic compound. Laboratory synthesis: Al4C3 + 12 H2O → 4 Al(OH)3 + 3 CH4 Reactivity: less reactive, but its mixture with air explodes after ignition (6-12 vol. % CH4, fire- damp explosion): CH4 + O2 → CO / CO2 + H2O, exothermal reaction Synthesis gas: CH4 + H2O � CO + 3 H2

Carbon halides

CF4 (carbon tetrafluoride) Properties: very stable, inert gas. Synthesis: SiC + 4 F2 → SiF4 + CF4 or CO2 + SF4 → CF4 + SO2

C2F4 (tetrafluoroethylene) Discovery: 1933. Properties: hydrophobic, organophobic, stable till 600 °C, resistive against acid, bases, oxidizing agents. Disadvantages: soft, gas permeable, big thermal expansion. Synthesis: CHCl3 + 2 HF (SbFCl4 catalyst, CF2ClH → C2F4 → (C2F4)n {(C2F4)n – teflon}.

CCl4 (carbon tetrachloride). Properties: does not exhibit hydrolysis, high density, colorless liquid, non-flammable, good non-polar solvent, stable till 400 °C, toxic, carcinogenic. Synthesis: CS2 + 3 Cl2 → CCl4 + S2Cl2, CH4 + 4 Cl2 → CCl4 + 4 HCl.

CFC's, Freons. CFCl3 / CF2Cl2 / CF3Cl. Properties: easy to condensate, inert, odorless, low viscosity, volatile liquids or gases. Catalyses break down of O3 molecules, ozone hole. Application: refrigerators, sprays. Produce: 100,000 t/year, decreasing. C–H bond containing

65 János Rohonczy: Inorganic Chemistry I. 5.1. Carbon

CFC derivatives are used today. Synthesis: CCl4 + HF → CFCl3 + HCl, CFCl3 + HF → CF2Cl2 + HCl.

CBr4 (carbon tetrabromide) Properties: light yellow liquid, less stable than CCl4, hydrolyses. Synthesis: 6 CCl4 + 4 Al2Br → 6 CBr4 + 4 Al2Cl6

CI4 (carbon tetraiodide) Properties: red crystals, hydrolyses. Synthesis: CCl4 + 4 EtI → CI4 +4 EtCl

Carbon oxohalides Structure: X2C=O, X = F, Cl, Br, I, and mixed halogen ligands.

COCl2 (phosgene) Properties: colorless, very toxic gas (chemical weapon). Synthesis: CO + Cl2 → COCl2, Produce: >1000 t/year. Application: H2N–(CH2)6–NH2 + 2 COCl2 → O=C=N–(CH2)6–N=C=O, HO–R–OH + O=C=N–R'–N=C=O → [–R–O–CO–NH–R'–NH–COO–]n, synthesis of polyurethanes. MO + COCl2 → MCl2 + CO2, synthesis of anhydrous metal halides (reductive halogenation).

COF2 (carbonyl difluoride), Application: laboratory synthesis of fluoro organic compounds. Synthesis: COCl2 + SbF3 → COF2 + SbCl3. Reaction: COF2 + H2O → CO2 + 2 HF (fast hydrolysis).

5.1.3. Carbon oxides,

Several known carbon oxides: CO / CO2 / C3O2 / C5O2 / C12O9

CO (). Structure: ←:C≡O: BP = 78 K (similar to :N≡N: with BP = 77,5 K). Properties: stable transition metal carbonyl complexes, e.g. Fe + 5 CO → [Fe(CO)5] straw- colored liquid. Labor synthesis: HCOOH + cc. H2SO4 → CO + H2O (from formic acid or Na formate) Industrial synthesis: C(hot) + H2O(vapor) → CO + H2 (synthesis gas, syngas) Properties: colorless, odorless, flammable, very toxic gas: CO + Haemoglobin → CO⋅Haemoglobin Chemical detection: CO + PdCl2 + H2O → Pd + CO2 + 2 HCl, or 5 CO + I2O5 = 5 CO2 + I2 Reaction: 2 CO +O2 � 2 CO2 (exothermal, CO is stable over 500 °C), CO + Cl2 → COCl2 CO + S(molten) → COS, CO + NaOH → HCO2Na (Na-formate) CO + H2O → no reaction – formal anhydrate. But: CO + H2Ovapor � CO2 + H2

CO2 () Structure: O=C=O linear. Properties: less toxic than CO, colorless, odorless, heavier than air. Laboratory synthesis: CaCO3 + 2 HCl → CaCl2 + H2O + CO2 Industrial synthesis: CH4 + 2 H2O → CO2 + 4 H2 (first step in NH3 production).

66 János Rohonczy: Inorganic Chemistry I. 5.1. Carbon

Earlier: CaCO3 → CaO + CO2 Occurrence: Natural gas, e.g. in Répcelak, Hungary (contamination: H2S, removal with catalytic oxidation → S8 precipitate). Produce: 10 Mt/year CO2. Application: cooling – dry ice (sublimates, liquid only under pressure), fire-extinguisher, power-gas, carbonated water. Synthesis of carbamide: CO2 + 2 NH3 → NH4[NH2COO] → H2N–CO–NH2 + H2O. Synthesis of silylated carbamic acid derivatives, which are good silylating reagents: 2 (CH3)2NH + (CH3)3SiCl + CO2 → (CH3)2N–COO–Si(CH3)3 + (CH3)2NH2Cl Silylation: R–OH + Me2N–COO–SiMe3 → Me3Si–OR + CO2 + Me2NH. + – – 2– Properties: water soluble: CO2(aq) � "H2CO3" � H3O + HCO3 {+ OH } → CO3 + H2O.

Carbon

C3O2 (tricarbon dioxide) Structure, properties: O=C=C=C=O, linear, cumulated bonds, yellow crystals

Preparation: 3 HOOC-CH2-COOH + 2 P2O5 → 3 C3O2 + 4 H3PO4

O O O O O O C3O2 →

O O O O O O red-violet crystal.

C5O2 (pentacarbon dioxide) Structure: O=C=C=C=C=C=O linear

C12O9 (mellitic acid anhydride) Properties: white solid

O Structure: O O O O

O O O O Table 31. Carbonic acid and derivatives. Carbonic acid Carbamic acid Thiocarbamic acid Carbamide Thiocarbamide HO–CO–OH H2N–CO–OH H2N–CS–OH H2N–CO–NH2 H2N–CS–NH2 decomposes decomposes decomposes white crystal stable

Derivatives are most stable than carbonic acid, which does not exists in anhydrous form.

CO((NH2)2) (carbamide, urea) Synthesis: CO2 + 2 NH3 → NH4[NH2COO] → H2N–CO–NH2 + H2O.

5.1.4. Carbon sulfides

CS2 (carbon ) Structure: S=C=S, linear molecule. Properties: BP = 46 °C, toxic, very flammable, ignition point -30 °C, good non polar solvent. CS2 + 3 O2 → CO2 + 2 SO2 Synthesis: CH4 + 4 S → CS2 + 2 H2S. Produce: 1 Mt/year

67 János Rohonczy: Inorganic Chemistry I. 5.1. Carbon

Application: synthesis of rayon, cellophane, CCl4. Reaction: CS2 + Na2S → Na2CS3 (Na-trithiocarbonate) CS2 + NaOH → Na2CO3 + Na2CS3

– + Important reaction: CS2 + NaOH + EtOH → EtO–CS–S + Na (Na-ethyl dithiocarbonate) In acid: 2 EtO–CS–SNa + H2SO4 → 2 EtOH + Na2SO4 + 2 CS2 (with cellulose: cellophane) 2– 2– – – Ions in solutions: CS2 / CS3 / COS2 / CS2OR / CS2NR2

COS (carbonyl sulfide) Structure: O=C=S, linear. Properties: BP = -50 °C. Synthesis: 2 CO + S2 → 2 COS.

5.1.5. Carbon nitrides

(CN)2 (cyanogen) Structure: :N≡C–C≡N: Properties: toxic, pseudo halogen: MP = -28 °C, BP = -21 °C, – – – Hydrolysis in bases: (CN)2 + 2 OH → CN + OCN + H2O Synthesis: Hg(CN)2 → Hg + (CN)2 2+ – Cu + 4 CN → 2 CuCN + (CN)2 (oxidation)

HCN ( or prussic acid) Structure: H–C≡N: � H–N=C: (99% : 1% ) Properties: almond-like odor, toxic gas, MP = -125 °C, BP = -13,4 °C. KCN, NaCN are water soluble, toxic salt. Industrial synthesis: NaNH2 + C(hot) → NaCN + H2 – 2– Reaction: 2 CN + CO2 + H2O → CO3 + 2 HCN

ClCN (cyanogen chloride) Structure: Cl–C≡N: Properties: very toxic, colorless gas. Synthesis: NaCN + Cl2 → Cl–CN + NaCl (trimer molecule contains [–N=CH–]3 ring).

H2N–CN (cyanamide) Structure: :N≡C–NH2 Trimer: [–N=C(NH2)–]3 (melamine). Synthesis: Cl–CN + 2 NH3 → H2N–CN + NH4Cl, or CaC2 + N2 → CaN–CN + C, Application: Ca salt is fertilizer, Produce: 1 Mt/year. Melamine + gives melamine resin.

HO–CN (cyanic acid). Isomerisation: HO–C≡N � H–N=C=O () The isocyanic acid is the stable compound.

HS–CN ( or rodanic acid) – Salts: SCN ( ion), Fe(SCN)3 complex (water soluble molecule, blood red solution). Isomerisation: HS–C≡N � H–N=C=S, (isothiocyanic acid, this is more stable).

68 János Rohonczy: Inorganic Chemistry I. 5.2. Silicon

5.2. Silicon

Valence shell: [Ne]3s23p2 History: minerals: quartz, silicates, obsidian. Element: Berzelius (1823, synthesis: K2SiF6 + 4 K → Si + 6 KF), 1860 – known compounds: SiF4, SiCl4, SiHCl3, SiH4, SiEt4, siloxanes. Naming: silex (Latin) Rochow (1941): silicon oil, (1960) elementary Si as semiconductor. Occurrence: in earth's crust: 27% lighter silicates ( Al, Na, K, Mg, Ca). In earth's mantle: heavier (Mg,Fe)2SiO4 Physical properties: solid, dark blue/grey, rigid. Structure: diamond-like lattice, but weaker hardness, lower MP = +1410 °C. Very low electric conduct at R.T., increasing temperature: increasing electric conductivity: thermal semiconductor. with (15-th column: P, As, Sb, Bi) - n-type (e–-excess), with (13-th column) B, Al, Ga, In – p-type (e–-lack). Band-structure: conductivity depends on the distance of valence and conductive bands. Overlapping bands: metals. Small gap: p-, n-type or thermal semiconductors. Big gap: electric insulators.

Synthesis: SiO2 + 2 C → Si + 2 CO (96-99% purity) 2 SiC + SiO2 → 3 Si + 2 CO (in SiO2 excess) With iron: SiO2 + C + Fe → CO + Fe/Si (ferrosilicon, Ppoduce: Mt / year) Purification: washing with water, sedimentation, further processes: Si + 2 Cl2 → SiCl4 (silicon tetrachloride) Si + 3 HCl → SiHCl3 + 2 H2 (trichlorosilane, at 300 ºC) SiCl4 + Zn/Mg → Si(powder) + ZnCl2/MgCl2 Sipowder (melting) → Si-ingot Cystallization by Czochralski method → single crystal (purity: 10-9 – 10-12%)

Van Arkel - de Boer method: SiI4 → Si + 2 I2 (thermal decomposition on W-wire) Thermal decomposition, SiH4 → Si + 2 H2 (epitaxial growing) Solar-cell production: Na2SiF6 + 4 Na(molten) → Si + 6 NaF Nuclear behavior: 28Si 92%, I = 0, 29Si 5%, I = 1/2 ( NMR active isotope), 30Si 3% I = 0. 31 30 – 31 Synthesis of Si: Si (n,β ) Si, t1/2 = 2,5 h. Application: n-activation analysis: 1.48 MeV.

IC fabrication: 1) crystal growing – ingot, 2) slicing to wafers, polishing, cleaning, 3) oxidation of the surface: Si + O2/H2O → SiO2 4) photoresist, UV irradiation through a mask, 5) solution of exposed resist by diluted acid, 6) etching: solution of SiO2 by HF, 7) solution of non-exposed resist by organic solvent, 8) deposition: by vapors of III/V elements (13/15. column), 9) etching: solution of SiO2 by HF, 10) repeat step 3) ... 9), 11) interconnect: Al-metal wires by evaporation and etching of Al, 12) packaging: Au, Al wires by thermocompression.

János Rohonczy: Inorganic Chemistry I. 5.2. Silicon

Chemical properties of Si: relatively inert at R.T. (SiO2 surface coating), but at 20 °C: Si + 2 F2 → SiF4 Si + cc. HNO3 + HF → SiF4 + H2O + NO, 100 °C: Si + NaOH → Na2SiO3 + H2 300 °C: Si + 2 Cl2 → SiCl4 Si + R–X → SiR2X2 / SiRX3 / SiR3X/ SiR4 / SiX4. 600 °C: Si + 2 S → SiS2 higher: Si + O2 → SiO2 Si + N2 → SiN / Si3N4, Si + P → Si3P4 Reactivity is similar to the carbons one at high temperature. Other reactions: Si(molten) is reactive: reduces the metal oxides to silicides (alloys). With the elements of 14-th column: SiC is carbide, others: silicides: Ge2Si, Sn2Si, Pb2Si. Molten Si reactions are performed in ZrO2 or Ti/V/Cr-boride vessels. 2– Coordination number: mostly 4 in molecules, but can be 5 in silatranes, or 6 in [SiF6] . Coordination number in minerals: 6, 8, 10.

5.2.1. Binary compounds

SiC (silicon carbide). Name: (1891) carborundum. Hardness: diamond > SiC > corundum (Al2O3). Polymorphism: 3D lattices with tetrahedrally coordinated silicon and carbon atoms. Hexagonal α-SiC Wurtzit lattice, cubic β-SiC diamond type lattice. Physical properties: colorless or yellow crystals. Industrial grade: black, purple, green – Fe contamination. Synthesis: SiO2 + 2 C → Si + 2 CO, Si + C → SiC (carbon excess). Physical properties: hard (polishing material), thermally stable till 2700 °C, resist again diluted HF, stable till 1000 °C in air (SiO2 cover), semiconductor (heater in furnaces, blue LED). Reaction: SiC + 4 Cl2 → SiCl4 + CCl4 (1000 °C) SiC + 2 Cl2 → SiCl4 + C (100 °C) SiC + 4 NaOH + 2 O2 → Na2SiO3 + 2 H2O + Na2CO3 (100 °C)

Silicides δ– δ– Stoichiometry: M6Si ... MSi6 . Properties: similar to borides and not to carbides, but lower MP.

Table 32. Melting point of binary compounds of 14-th column elements.

TaC TaB2 TaSi2 3800 °C 3100 °C 1560 °C Table 33. Bond types and properties of silicides. Covalent Ionic Metallic p-group s-group d-group Ge2Si, Pb2Si Na2Si, Mg2Si Cu5Si, Fe3Si, Mn3Si insulator reactive metallic, eutectic (non reactive) inert with HF/F2/Cl2 they react only with molten NaOH

70 János Rohonczy: Inorganic Chemistry I. 5.2. Silicon

Atomic radius: rSi > rB – borides and silicides are not isostructural compounds. Bond types and physical chemical properties of silicides depend on electronic configuration of partner elements.

Preparation: Si + M → MSi, SiO2 + MO + C/Al → MSi + CO/Al2O3 – 2– Reaction: Ca2Si + H2O/OH → SiO3 + H2 (isolated Si), – CaSi + H2O/OH → SiH4 / SinH2n+2 (polysilane, Si–Si chain) – CaSi2 + H2O/OH → H2 + (SiH2)2 (, Si layer) Mg2Si + H2SO4(aq) → 2 MgSO4 + SiH4 (silane, gas)

5.2.2. Silanes – Hydrosilicons

Structure: SinH2n+2 (silanes, n = 1..4 stable, 1≤n≤8 colorless gas/liquids), known cyclosilanes as well. Reactive, flammable compounds. Thermal stability: a) increasing chain length → decreasing stability, b) C–C > C–Si > Si–Si c) Si–X > C–X (halo-silanes are more stable). Laboratory synthesis: Mg2Si + H2SO4(aq) → MgSO4(aq) + SiH4 Mg2Si + 4 NH4Br(NH3) → 2 MgBr2 + SiH4 + 4 NH3 Industrial synthesis: SiCl4 + LiAlH4 → LiCl + AlCl3 + SiH4, Si or Fe/Si + HCl → SiHCl3 + H2 Reactivity: much more reactive compared to hydrocarbons. Reason: bigger Si radius, and Siδ+–Hδ– bond polarity is fair to nucleophilic attack. Reactions: pyrolysis: SiH4 → Si + 2 H2 (highest purity Si) RSiH2SiH3 → SiH2 + RSiH3 RSiH2SiH3 → :SiRH + SiH4 () hydrolysis: SiH4 + H2O → (insoluble in water) – 2– SiH4 + H2O + 2 OH → SiO3 + 4 H2 (soluble in bases) other: SiH4 + 2 Cl2 → SiH2Cl2 + 2 HCl (explosion) SiH4 + HCl → SiH3Cl + H2 (substitution) SiH4 + 2 AgI → 2 Ag + HI + SiH3I (colorless liquid) SiH4 + K → KSiH3 + 1/2 H2 (colorless crystal) reaction of SiH3I: SiH3I + Ag2S → S(SiH3)2 SiH3I + Li2Te → Te(SiH3)2} reaction of KSiH3: KSiH3 + MeI → SiH3Me, KSiH3 + SiH3Br → Si2H6 (pure)

SiHnX4-n (halo-silanes) Industrial synthesis: Si + 3 HCl → SiHCl3 + H2 (trichlorosilane or silico-). Application: primary materials for high temperature vulcanised silicones.

5.2.3. Silicon halides

SiF4 / SiCl4 (silicon tetrafluoride / silicon tetrachloride) Industrial synthesis: SiO2 + 4 Ca5(PO4)3F + H2SO4 → SiF4, (further: SiF4 + 2 HF → H2SiF6 Si + 2 Cl2 → SiCl4. Properties: reactive Si–X bonds. Results: equilibration and hydrolysis. Equilibration: SiCl4 + SiBr4 � 2 SiCl2Br2 + SiBrCl3 + SiClBr3 (cross-halogenation with

71 János Rohonczy: Inorganic Chemistry I. 5.2. Silicon

SbF3). Hydrolysis: Double attack of H2O on Si–Cl bond, fast reaction, SNi mechanism.

SinX2n+2 Properties: viscous liquid, or solids. More stable than hydrosilanes. Synthesis: Si + SiF4 → 2 SiF2(g) → (SiF2)x → SinH2n+2 (mixture), Si + 3 SiCl4 → 2 Si2Cl6 + .... 5 Si2Cl6 → Si6Cl14 + 4 SiCl4

Reaction: partial hydrolysis: SiF4 < SiCl4 < SiBr4 < SiI4 (explosion), increasing hydrolysis sensitivity. Application: high purity Si, SiO2, Si esters. Produce: > 100,000 t/year.

5.2.4. Silicones - silicon organic compounds

History: Charles Friedel (1863), SiCl4 + 2 Zn(Et)2 → 2 ZnCl2 + Si(Et)4 Ladenburg (1872) synthesis: Si(Et)3Cl; Si(Et)2Cl2; Si(Et)3Cl. Hydrolysis products are sticky organo-siloxanes. Rochow (1941) Silicones as industrial products, 2 CH3Cl + Si → SiMe2Cl2 (70%) + SiMeCl3 (12%) + SiMe3Cl(5%)+ + SiCl4 + SiMe4 + dimers. Isolation by fractional distillation. Reaction: controlled hydrolysis (1), and condensation (2): 1) Me3SiCl + H2O → M3SiOH + HCl (trimethylsilanol) 2) Me3Si–OH + HO–SiMe3 → Me3Si–O–SiMe3 (hexamethyldisiloxane, notation: M2) M2 is stable, colorless liquid, BP = 100.8 °C. Me2SiCl2 + H2O → [–O–SiMe2–]n + 2 HCl, n = 4-6 cyclic (D, Dn). Silicone oil. Synthesis by hydrolysis of mixtures: Industrial synthesis: 2 M+ x D4 + 4x H2O→ M–D4x–M + 2x HCl, x ≈ 10,000-100,000. Properties: increasing viscosity with increasing chain length – planned behavior according to reaction above. Slightly changing viscosity between -100 °C and +300 °C. Low surface tension – anti-foaming agents (fermentation, distillation, cooking oil), hydrofobic, non- toxic. Application: hydraulic fluid (airplanes), oil bath, lipstick, car polish.

Silicone grease. Me, Ph-silicone + SiO2 + (Li-stearate). Silicone rubber Industrial synthesis: MeSiCl3 ("T") + H2O → [–O–SiMe(O–)2] (by hydrolysis of T trifunctional → 3D silicone network, white powder): M + x D4 + y T + H2O → M–Dn–T–Dn - loose 3D network, rubber Properties: elasticity designed by T concentration, temperature independent, very good electric insulator, antistatic, non rigid at low temperature, low degradation at high temperature (350 °C), UV resistant, non degrading, biocompatible.

Application: tubes, insulator on cables, sticky, implants (breast implant, etc.).

Silicone resin: SiCl4 – Q. Hydrolysis of lots of T and/or Q results in 3D network, rigid polymer. Industrial synthesis: Ph–SiCl3 → [–Si(Ph)(OH)–O–] chain + HCl → 3D network + H2O.

72 János Rohonczy: Inorganic Chemistry I. 5.2. Silicon

Application: printed circuit boards, high temperature pigments, thermostable coatings, insulators.

High temperature vulcanization Catalyst: Ph–CO–O–O–CO–Ph → 2 Ph–CO–O⋅ (benzoyl peroxide)

O CH3 Si Si O CH3 [–Me2SiO–]n + RO⋅ → ROH + (cross-linked) Other method with Pt catalyst: [–O–Si(Me)H–] + [–O–Si(Me)(CH=CH2)–] → –Si–CH2–CH2–Si– cross-link

Low temperature vulcanization

a) [–O–Si(Me)(O–C(O)CH3)–] + H2O → [–Si–O–Si–O–] + CH3COOH acetoxymethylsiloxane, catalyst: (n-Bu)2Sn(Ac)2 b) textile–OH + HO–SiR3 → Textile–O–SiR3 hydrofobic Produce: 300,000 t/year.

5.2.5. Silicon oxides, silicic acids, silicates

SiO (). Synthesis: 2 Si + O2 → 2 SiO, or SiO2 + C → SiO + CO. Properties: Brown, black solid, dissociation below 1180 °C: 2 SiO → SiO2 + Si.

SiO2 (). Occurrence: quartz: sand, berg crystal, amethyst, citrine, onyx, jasper. Rare: tridymite, cristobalite, coesite. Amorphous: obsidian, diatomite,

Structure: SiO4 tetrahedrons, O–Si–O angle: close to 109,5°, flexible Si–O–Si angle ≈ 153° ± 20°, results in glass forming ability, p-dπ conjugation. Physical properties: high MP, UV transparent, hard, piezoelectric, glass forming oxide.

Polymorphism of SiO2

867 ºC 1470 ºC 1713 ºC β-tridimite β-cristobalite molte

573 ºC 120-260 ºC 200-280 ºC

α-quartz α-trdmite α-cristobalite

The most stable is α-quartz.

73 János Rohonczy: Inorganic Chemistry I. 5.2. Silicon

Chemically resistive, but: SiO2 + H2O → decreased MP, Si–O–Si → Si–OH (hydrolysis) SiO2 + 2 F2 → SiF4 + O2 SiO2 + 4 HF → SiF4 + 2 H2O (glass etching) SiO2 + 3 C → SiC + 2 CO (carbide synthesis) SiO2 + FeO → FeSiO3 (iron orthosilicate) SiO2 + Na2CO3 + CaO/CaCO3/PbCO3 → glass (unordered silicates) SiO2 + 2 NaOH(aq) → Na2SiO3 + H2O (Na orthosilicate, water-glass) Water-glass application: water solubility: fire-proof coating; alkaline hydrolysis: detergent, glue, acid-proof cement, silica gel. Properties of silicate glasses: flexible Si–O–Si bond, terminal O with metal cations – Na: softening, Ca: water insoluble, Co/Fe: color, Al2O3/B2O3: hardening, PbO: skeleton forming: non toxic, high optical refraction.

Silicic acids

H4SiO4 (ortho-silicic acid). Synthesis: SiCl4 / Si(OEt)4 + H2O → H4SiO4 or Si(OH)4 (Only in solution) Properties: Si(OH)4 - H2O → (SiO2)x(H2O)y (3D network, polysilicic acid)

H6Si2O7 (disilicic acid) Only in solution.

(H2SiO3)n (metasilicic acid) Polymer, formed from salts. + 2– Synthesis: Na2SiO3 + H2SO4(aq) → H2SiO3 +2 Na + SO4

(H2Si2O5)n (di-metasilicic acid, some salts as minerals). Synthesis: 2 H2SiO3 → H2Si2O5 + H2O.

Silicates. Salts of silicic acids, rock-forming minerals, crystal lattices containing SiO4 tetrahedrons.

4– (a) SiO4 (ortho-, nesosilicates) ZrSiO4, Be2SiO4. Structure: missing Si–O–Si bonds. 6– (b) Si2O7 (di-ortho-, sorosilicates) Sc2Si2O7, [Zn4(OH)2Si2O7]. One common O atom. 2– 12– (c) SiO3 )n (cyclosilicates) Two common O atoms, n = 3, 4, 6, 8. E.g. Be3Al2[Si6O18 ], (d) Inosilicates 2– (SiO3 )∞ , pyroxene group: MgSiO3, CaSiO3. 4– (Si2O6 )∞ spodumene: LiAlSi2O6 6– (Si8O22(OH)2 )∞ amphiboles, tremolite: [Ca2Mg5(Si4O11)2(OH)2] 12– (Si2O8 )∞ (double chain inosilicates) Ca2Al2Si2O8 6– (e) (Si2O5(OH)4 )∞ (2D-layer phyllosilicates) serpentine, kaolinite, montmorillonite, muscovite, mica, talk, biotite, bentonite. (f) Tectosilicates: quartz, feldspars, zeolites

Feldspars: orthoclase, albite, anorthite. 60% of the rocks.

Zeolites: open Al-silicates, holes, tubes. Structure: 24 SiO4 (a)-polyhedron, (c) 1 big hole (B) among 8 (A) smaller ones. Synthesis: alkaline silicate + aluminate gel → zeolite. Application: molecular sieves.

74 János Rohonczy: Inorganic Chemistry I. 5.2. Silicon

Zeolite structures

– – – Ultramarine: blue pigment. Si–O–Al–O–Si. 3D network, guest ions: S2 , S3 , S4 . Blue color. Application: blue painting (Meissen porcelain).

Application of SiO2: α-quartz Properties: piezoelectric crystal (oscillators / frequency filters, quartz clocks) Synthesis: NaOH(aq) + SiO2(powder) → SiO2(α-quartz), hydrothermal crystallization.

Quartz glass. Properties: small dilatation, high thermal stability, UV transparency: sample holder in UV photometers, photochemical reactors. High optical refraction, inert: high purity chemical equipments. High melting point, viscous, with narrow thermal range: expensive. 2 Kieselguhr. Amorphous SiO2, porous (800 m /g). Synthesis: Na2SiO3 + H2SO4 → gel. Washing, drying. Application: desiccator (CoCl2, blue/pink color marker), adsorber, chromatographic phase. Non-toxic: additive in food industry (stick inhibitor).

2 Smoke silica. SiO2, 500 m /g. Synthesis: SiCl4 + H2O + O2 → SiO2 + HCl. Application: additive in synthetic resins, silicon grease, silicon rubber.

Kieselguhr, diatomite. From natural sources: 2 Mt/year: Application: filtration, settling.

Aerogel. solid, "frozen smoke", very good thermal insulators.

5.2.6. Silicon–sulfur compounds

SiS2 (silicon disulfide) Structure: differs from SiO2. There is no delocalization in Si–S–Si bond: non glass-forming. Properties: white, pin crystals, sublimation, MP = 1090 °C. Synthesis: Si + S → SiS2 (similar: SiSe2) Reaction: SiS2 + 2 H2O → SiO2 + 2 H2S SiS2 + 4 Et–OH → Si(OEt)4 + 2 H2S SiCl4 + 4 EtOH → Si(OEt)4 + 4 HCl Application: conserve of buildings, tube coating: hydrolyzed SiO2. Produce: 2,000 T/year.

75 János Rohonczy: Inorganic Chemistry I. 5.2. Silicon

5.2.7. Silicon – nitrogen compounds

Si3N4 (silicon nitride. Synthesis: SiCl4 + NH3 → Si3N4 + NH4Cl, SiO2 + C + N2/H2 → Si3N4 + CO + H2O. Properties: white powder. Inert in air till 1000 °C. Hardness = 9, density: 3.2 g/cm3 (high), stable till 1900 °C, good insulator.

Si(NH2)4 (silicon tetramide) Synthesis: SiCl4 + NH3 → Si(NH2)4 + NH4Cl, non stable.

Si(NH)2 (silicon ) Synthesis: Si(NH2)4 → Si(NH)2 + 2 NH3, SiS2 + 4 NH3 → Si(NH)2 + 2 NH4SH.

N(SiH3)3 (tris(silyl)amine) Synthesis: SiH3I + NH3 → N(SiH3)3 (p-dπ bond: planar structure).

Silazanes. Synthesis: 2 Me3SiCl + NH3 → NH(SiMe3)2 + NH4Cl (hexamethyl- disilazane).

Silatranes. Voronkov, Hencsei Pál (Technical Univ., Budapest). Structure: Si in fivefold-coordinated state. Small Si–N distance.

76 János Rohonczy: Inorganic Chemistry I. 5.3. Germanium, Tin, Lead

5.3. Germanium(Ge), Tin(Sn), Lead(Pb)

Valence shell: ns2 np2 History: Ge Mendeleev (1869) missing element below silicon: ekasilicon. Discovery from Ag8GeS6, argyrodite, C. Winkler (1886). Name: Germany. Sn Known in Old Testament. Name: Latin stannum, bronze: 15% Sn + 85% Cu (prehistoric ages), Sn + Pb (Roman Empire). Pb Know in prehistoric ages. Name: Latin plumbum. Application: painted pottery, Pb-floor, water pipes. Biological influence: Inorganic Ge, Sn are non-toxic, Pb is very toxic, complex formation with oxo and cysteine SH-groups of enzymes. Well soluble, nervine.

Occurrence Ge 1,5 ppm /scattered, few minerals/ expensive. Enrichment: coal cinder, roasting of Zn-minerals: in smoke. Produce: 100 t/year. Sn 2 ppm /abundant/ always in +4 oxidation state. SnO2 (cassiterite). Produce: 200 et/year. Pb 13 ppm /abundant, (final element of radioactive decays) / always in +2 oxidation state. Ores: PbS (galenite), PbSO4 (anglesite), PbCO3 (cerussite). Produce: 4 Mt/year.

Make of Ge: In ZnS roasting smoke: 10% GeS2 GeO2 + 2 H2SO4 → Ge(SO4)2 + 2 H2O (solution) Ge(SO4)2 + 4 NaOH → Ge(OH)4 + 2 Na2SO4 (precipitation) Ge(OH)4 + HC / Cl2 → GeCl4 (BP = 83 °C), and ZnCl2 (BP = 756 °C) (separation) GeCl4 + 2 H2O → GeO2 + 4 HCl (hydrolysis) GeO2 + 2 H2 → Ge + 2 H2O (reduction)

Purification: zone melting. Produce: 100 t/year. Application: semiconductors, IR-transparent (window, prisms, lenses), in superconductor alloys, Mg2Ge: in phosphoresce powders, Synthesis of Sn: SnO2 + 2 C → Sn + 2 CO and Fe2O3 + 3 C → 2 Fe + 3 CO Worry: Fe contamination – brittle, rigid alloys. Solution: O2 excess at 1200 °C. Thermodynamically stable: molten Sn and insoluble FeO. Application: 40% tin-plated iron; 24% solder: Sn(33%) / Pb + Ga / In / Bi; 15% bronze: Sn(10%) / Cu + P / Zn (500,000 t/year); 5% bearing alloy – Babbit metal: Sn(80%) / Cu soft grain in Pb(75%) / Sn(12%)/Sb solid matrix; 3% soft tin (knick-knacks, organ pipes); 0,5% type-metal: Pb / Sn / Sb (Sn / Sb: hard). Nb3Sn alloy: superconductors, electromagnets. Metal Sn or alloys with low melting point: float glass production. Synthesis of Pb: from galenite (PbS) Reduction with carbon: PbS + 3/2 O2 → PbO + SO2 PbO + C → Pb(molten) + CO PbO + CO → Pb(molten) + CO2 János Rohonczy: Inorganic Chemistry I. 5.3. Germanium, Tin, Lead

Reduction without carbon: PbS + 3/2 O2 → PbO + SO2 2 PbO + PbS → 3 Pb(molten) + SO2

Removing contaminants – refinement: by separation: Cu; by oxidation of Sn, As, Sb by NaOH / NaNO3; by electrolysis of Ag, Au; Zn +Cl2 and vacuum distillation of ZnCl2; by electrolysis of Bi in PbSiF6 electrolyte.

Application: 50% accumulators (91% Pb / Sb), solder, fuse, Babbit metal, type-metal, water- pipes, cable coating, plates, etc., lead organic compound: Pb(Et)4, Pb(Me)4, pigments, paintings: Pb3O4 / PbCrO4 / PbMoO4 / PbO, glass industry: Pb silicates: lead-glass, flint- glass.

Nuclear/physical properties Atomic numbers are even numbers – several stable isotopes.

Ge: 5 stable isotopes, gray/white crystal, diamond structure, similar to silicon, but lower hardness, MP, BP and ionization potential (more metallic element). Sn: 10 stable isotopes (most among the elements). 117Sn and 119Sn have I = 1/2, they are NMR active, 119Sn is Mössbauer-active as well. Allotropy: α-Sn below +13 °C. Diamond lattice, gray powder: "tin pest" β-Sn, white tin, tetragonal lattice, metal, stable. Pb 4 stable isotopes. 204Pb is closing element in the decay chains, 207Pb: I = 1/2. Blue/gray color, soft, heavy (ρ = 11.34 g/cm3), low MP, electron conductive metal. Chemical properties Reaction: electro negativity is decreasing in top-down order, MIV → MII tendency is increasing, M–M and M–O–M stability is decreasing, coordination number = 5-6, Sn / Pb can form cluster anions.

Ge is more reactive than Si: slowly soluble in cc. H2SO4 / cc. HNO3, soluble in diluted acids / bases + H2O2 / NaOCl Ge + O2 → GeO2 Ge + H2S / S → GeS2 Ge + Cl2 / Br2 → GeX4 Ge + HCl → GeCl4, GeHCl3 Ge + R–X → GeR2X2

Sn is more reactive than Ge, amphoteric: Sn + H2O(vapor) → SnO2 + H2, 2+ – Sn + 2 HCl → Sn + 2 Cl + H2 with cc. HCl → SnCl2 2– Sn + 2 KOH + 4 H2O → [Sn(OH)4] + 2 H2 Sn + 2 X2 → SnX4 Sn + SnX4 → 2 SnX (where X = F / Cl / Br / I) Sn + S/Se → SnS / SnS2 and SnSe / SnSe2 Sn + Te → SnIITe Pb most reactive, pyrophoric, surface is protected by oxide, carbonate, sulphate, chloride coating. Solubility: 2+ – Pb + HNO3(aq) → Pb + 2 NO3 + H2

78 János Rohonczy: Inorganic Chemistry I. 5.3. Germanium, Tin, Lead

Pb + cc. H2SO4 → PbSO4 + H2 (weakly soluble) PbCl2 is better soluble. Well soluble: Pb(NO3)2 and PbAc2. (lead sugar) Pb + F2/Cl2 → PbX2 (at high pressure: PbX4) Pb + S/Se/Te → PbIIS / PbIISe / PbIITe

5.3.1. Ge/Sn/Pb hydrides, and hydrido halides

GenH2n+2 (hydrogermanes) Properties: 1 ≤ n ≤ 5 colorless gases or liquids Synthesis: Mg2Ge + 4 HCl → GeH4 + 2 MgCl2 GeCl4 + 4 LiAlH4 → GeH4 + LiCl + AlCl3 GeO2 + NaBH4 → GeH4 + NaBO2 SiH4 + GeH4 → H3Si–GeH3 ...

GeH4 (germane) Properties: Not self burning; there is no hydrolysis in bases and acids. + – GeH4 + NH3 → NH4 + GeH3 , reacts as acid. KGeH3 stable salt.

GeHnX4-n Properties: colorless, reactive liquids. X = Cl / Br / I, n = 1, 2, 3 are known. Synthesis: Ge / GeH4 / GeX2 + HX → GeHnX4-n (e.g. GeH3Cl) Reaction: 2 GeH3Cl + H2O → O(GeH3)2

SnH4 (stannane) Properties: BP = -52,5 °C, slowly dissociates at 20 °C: Sn + 2 H2 Synthesis: SnCl4 + LiAlH4 → SnH4 + LiCl + AlCl3 Reactivity: soluble in strong acids / bases. Reduction agent.

Sn2H6 (distannane) Properties: Less stable, known aryl derivatives: RnSnH4-n Synthesis: RnSnCl4-n + LiAlH4 → RnSnH4-n + AlCl3 + LiCl

PbH4 () Properties: Only traces in the reaction written above. Known derivatives: R2PbH2 and R3PbH, all dissociate over -20 °C.

5.3.2. Ge/Sn/Pb halides

Two known series: MX2 / MX4. More stable GeX4 / SnX2 / PbX2

GeX4 (germanium tetrahalides) Properties: colorless liquid, or orange crystal. Synthesis: Ge + 2 X2 → GeX4, or GeO2 + 4 HX → GeX4 + 2 H2O. Reaction: GeX4 + 2 H2O → GeO2 + 4 HX (hydrolysis) GeX4 + n-Li–R → RnGeX4-n + n LiX – 2– GeX4 + 2 X → [GeX6 ], octahedral, X = F/Cl (complex formation)

GeX2 (germanium dihalides) Synthesis: GeF4 + Ge → 2 GeF2 (white, solid, MP = 110 °C, Structure: [GeF3]n chain), GeHCl3 → GeCl2 + HCl, Reaction: GeCl2 + H2O → Ge(OH)2 (yellow) → GeO (brown).

79 János Rohonczy: Inorganic Chemistry I. 5.3. Germanium, Tin, Lead

SnF2 (tin difluoride) Synthesis: SnO + 2 HF → SnF2 + H2O (Sn4F8: tetramer with F-bridges)

SnCl2⋅2 H2O Properties: reduction agent, e.g. creating metal coatings. Synthesis: Sn + 2 HCl(aq) → SnCl2(aq) + H2

SnCl2 (tin dickloride) Synthesis: Sn + 2 HCl(gas) → SnCl2 + H2 (SnCl2, SnBr2 white, SnI2 red crystal)

SnF4 (tin tetrafluoride) Synthesis: SnCl4 + 4 HF → SnF4 + 4 HCl.

SnX4 Synthesis: Sn + 2 X2 → SnX4, colorless liquid or crystal. Application: synthesis of SnO2, Friedel-Crafts catalyst.

PbX4 (lead tetrahalides) Properties: PbF4, MP = 600 °C, very stable. PbCl4 MP = -15 °C. yellow oil, decomposition at 50 °C: PbCl4 → PbCl2 + Cl2. More stable complexes: PbCl2 + Cl2 + 2 KCl → K2[PbCl6], stable yellow salt.

PbX2 (lead dihalides) Properties: More stable than PbX4 compounds. 2+ + Synthesis: Pb + 2 HX → PbX2 + 2 H , known mixed derivatives. 4– Structure: Isolated [PbX6] octahedral units in complexes.

5.3.3. Ge/Sn/Pb oxides, and hydroxides

GeO2 () Properties: white, several polymorphs with 4 and 6 coordination number. Synthesis: Ge + O2 → GeO2. Reaction: GeO2 + Ge → 2 GeO (brown) ← Ge(OH)2 · H2O. SnO (tin oxide) – – Synthesis: SnCl2 + 2 OH → Sn(OH)2 + 2 Cl , white gel, (in crystal: Sn6-cluster), Sn(OH)2 → SnO + H2O, blue-black, metastable. – – 2– Reaction: SnO + H2O + OH → [Sn(OH)3] , (AX3E structure, [Sn(OH)4] is unknown)

SnO2 (tin dioxide – cassiterite) Reaction: Sn(OH)2 + 1/2 O2 → SnO2 + H2O, stable, insoluble in water, diluted acids and – 2– bases. Soluble in concentrated bases: SnO2 + 2 OH + 2 H2O → [Sn(OH)6] , hexahydroxo- stannates 2– + Other reactions: [Sn(OH)6] + 2 H → Sn(OH)4 + 2 H2O → SnO2 K2[Sn(OH)6] → K2SnO3 ← K2O + SnO2 (melting) Application: milk glass, pigments, reflects IR radiation. With In2O3 and Sb/F doping colorless, transparent, electron conductive coating: ITO (Indium Tin Oxide). With Sb2O3 – catalyst. PbO (lead oxide). Polymorphs: red and yellow crystals. Synthesis: 2 Pb(molten) + O2 → 2 PbO Pb(NO3)2 → PbO + 2 NO2 + 1/2 O2

80 János Rohonczy: Inorganic Chemistry I. 5.3. Germanium, Tin, Lead

PbO2 () Properties: Brown solid. Synthesis: oxidation on anode (lead battery). – – Pb3O4 + OH + Cl2 → PbO2 + Cl , or with other strong oxidizer agents.

Properties: thermally unstable: PbO2 → Pb12O19 → Pb12O17 → Pb3O4 → PbO + O2 Application: PbO red/yellow – lead glass, pottery painting, pigments, battery (100,000 t/year) Pb3O4 (minium) – against corrosion. Toxic, prohibited. PbO2 – oxidation agent, electrode in lead/acid battery PbTiO3/PbNb2O6 – ferroelectric crystals.

5.3.4. Ge/Sn/Pb oxoacid salts

Ge salts Known but unstable compounds. Sn salts II IV Properties: Sn /Sn basic salts are stable: Sn2O(PO4)2, Sn(NO3)4, SnSO4, crystalline Sn(SO4)2 ⋅ 2 H2O. Synthesis: by oxidation of SnSO4 solution, quickly hydrolyses in water. Pb salts Properties: Pb(NO3)2, Pb(ClO4)2⋅3 H2O are water soluble. PbSO4, PbCrO4 are weakly soluble.

5.3.5. Ge/Sn/Pb chalcogenides

MX, where M = Ge, Sn, Pb and X = S, Se, Te. All nine combinations are known. Structure: GeS/SnS layered lattice, PbS (galenite) semiconductor, PbSe, PbTe – IR detectors. 2+ 2– Synthesis: Sn + S → SnS, Pb + S → PbS, GeCl2 + H2Se → GeSe + 2 HCl.

5.3.6. Anions with cluster skeleton

2– 2– 4– 4– Sn5 , Pb5 , Sn9 , Pb9 etc., typical Zintl-phases, Wade-cluster structures. Stabilized by multi-toothed ligands.

5.3.7. Ge/Sn/Pb organic compounds

Ge Organogermanium compounds are similar to ones, but they are more reactive and thermally less stable. Synthesis: 2 Me3GeBr + 2 K → 2 KBr + Ge2Me6 (MP = -40 °C, BP = +140 °C) Properties: [–R2GeO–]n partial hydrolysis. Sn Properties: Soluble organostannanes are toxic. Synthesis: SnCl4 + 4 R–Cl + 8 Na → SnR4 + 8 NaCl SnCl4 + SnR4 � equilibration {+LiAlH4} → RnSnH4-n cyclostannane; e.g.: [SnEt2]6, t Application: R.T. vulcanization of silicones: Bu2Sn(OAc)2 is non-toxic. Pb Properties: More than 2000 known organo lead compounds are known. PbEt4, volatile liquid, very toxic. Produce: 1 Mt/year, leaded fuel.

81 János Rohonczy: Inorganic Chemistry I. 5.3. Germanium, Tin, Lead

Laboratory synthesis: Li–R / R–Mg–X / AlR3 + PbCl2 / K2PbCl6 → PbR4 Industrial synthesis: 4 R–X + Pb/Na(alloy) → PbR4 + 4 NaX

Properties: PbEt4 → PbEt2 + 2 Et⋅ , radical trap, terminates chain reactions. In exhaust gases: Pb / PbO vapor (other radical traps are in lead-free fuels). Catalytic converter – vehicle emission control: (Al, Si, Ti, Zr ceramic carrier + Pt/Pd/Rh), 2 CO → + O2 → 2 CO2. It works at high temperature. Byproduct: nitrogen monoxide.

82 János Rohonczy: Inorganic Chemistry I. 6.1. Boron

6. (13th column) B, Al, Ga, In, Tl

Valence shell: ns2np1 Boron is nonmetallic element, similar to carbon in elementary form – e.g. forms clusters like fullerenes. Boron clusters contain tree center bonds. Other elements are metals in this group, they do not have covalently bonded allotropes.

Table 34. Physical properties. MP (°C) BP (°C) Crystal lattice Hardness Oxidation state Electric conductivity (cm-1Ω-1) B 2300 2550 atomic 11 3 1⋅10–6 Al 660,4 2467 metal 2,75 3 3,4⋅105 Ga 29,8 2403 metal 1,5 3 6,8⋅104 In 156,6 2000 metal 1,2 3 1,2⋅105 Tl 303,5 1457 metal 1,2 1 / 3 6,2⋅104

6.1. Boron (B)

Valence shell: 2s22p1

History: Na2B4O7⋅10 H2O (borax) was used in overglaze enamelling in prehistoric ages. Discovery: Davy, Gay-Lussac (1808) – impure product, Moissan (1892) – high-purity. Name: borax-carbon Occurrence: 9 ppm in Earth crush (rare), concentrated at volcanic areas. Minerals: primary: NaCa[B5O6 (OH) 6]⋅5 H2O, ulexite, Na2 [B4O5 (OH) 4]⋅8 H2O – borax, secondary: Ca2 [B3O4 (OH) 3] 2⋅2 H2O – colemanite, Na2 [B4O5 (OH) 4]⋅2 H2O – kernite. Physical properties: Very hard element, high MP, low density, low electric conductivity. Fine powder: black, crystals: dark red. Allotropes: α-rhombohedral. Unit cell: 12 atoms. Structure: B12-icosahedrons. β-rhombohedral. Unit cell: 105 atoms (most stable), complicated structure. α-tetragonal. Unit cell: 50 B atoms. 4 icosahedrons + 2 B atoms. Table 35. Nuclear properties. Isotope Abundance Spin (n,α) neutron absorption cross section (barn–1) 10B 20% +3 3800 11B 80% -3/2 0.005

Preparation: B2O3 + 3 Mg → 2 B + 3 MgO Electrolysis of KBF4 in KF/KCl → Bpowder 2 BBr3 + 3 H2 → 2 B + 6 HBr (decomposition over 2000 °C) (BCl3 is suitable as well, BI3 is expensive, BF3 is too stable) Produce: 4 Mt/year. Chemical properties: Very stable covalent bonds: B-B and B-O (in boric acids and borates). Small atomic radius (rB): borides are alloy type compounds. Electron-pair acceptor, Lewis- acid. Multicenter bonds, cluster structures. C-H fragments and B-H– ions are isoelectronic, bond polarity is opposite: Cδ–-H, Bδ+-H. János Rohonczy: Inorganic Chemistry I. 6.1. Boron

Reactivity: 2 B + 3 F2 → 2 BF3 4 B + 3 O2 → 2 B2O3 (passive oxide coating on surface) High temperature reactions: B + non-metal → BX3 (except: H2, Ge, Te, noble-gases) B + metal → metal borides (except: Ag, Au, Cd, Hg, Ga, In, Tl, Sn, Pb, Sb, Bi) Relatively inert element: B + cc. NaOH /→ B + NaOH(molten) /→ Chemical solubility: B + Na2CO3/NaNO3 → Na-borate or B + cc. H2SO4/cc. HNO3 → Application: Common boron containing compounds: oxide, boric acid, esters, borids, halides, boranes, organo-boranes. 35%: heat-proof glass (Pyrex), glass-wood, fiberglass. 20%: detergent (Na-perborate), soap, cosmetics. 15%: enamelling on porcelain. 10%: herbicides, fertilizers. Other: neutron absorber, metallurgy (borids), tanning, retardation of burning.

6.1.1. Borides δ– MnBm Stoichiometry: M5B ... MB15, and non-stoichiometric compounds as well. Table 36. Structures of borides.

Isolated B atom M / B big Mn4B • Isolated B2 pairs M / B big V3B2 •  • Bn chain M / B ≈ 1 VB, TiB Bn chain + sidechain M / B ≈ 1 Ru11B8

Double chain M3B4 Cr3B4

Layer MB2 TiB2, ZrB2

3D B-B network or MB4, MB6, LaB12, YB66 network of icosahedrons clusters MB10,MB12,MB66

Metal rich borides

Properties: very hard, chemically inert, high MP, metal shiny, good el. conductivity, e.g.: Zr, Hf, Nb, Ta, Ti borides: MP >3000 °C. Preparation: Cr + n B → CrBn BCl3 + W + 3/2 H2 → WB + 3 HCl TiCl4 + 4 BCl3 + 10 H2 → 2 TiB2 + 20 HCl 7 Ti + B2O3 + 3 B4C → 7 TiB2 + 3 CO

Application: TiB2, ZrB2, CrB2 very hard: turbine blade, rocket nozzle, burner. Chemically inert coatings: reactors, thermocouples, high temperature electrodes.

B4C (boron-carbide) Application: nuclear technology, n-absorber: 10B(n) → 7Li+4He, shielding, regulation. Other: polishing material, coupling, brake block, light corselet.

B4C fiber. Structure: B12C3 ... B13C2 (B12 icosahedrons are connected by C atoms).

84 János Rohonczy: Inorganic Chemistry I. 6.1. Boron

Preparation: 4 BCl3 + 6 H2 + C(fiber) → B4C(fiber) + 12 HCl Boron film/fiber: 2 BCl3 + 3 H2 → 2 B + 6 HCl (Produce: 50 t/year) Application: composite materials: aircraft body and wing, tennis and golf racket, bicycle.

6.1.2. Boron halides

Properties: volatile, monomer molecules (even BF3, in spite of ∆EN>2). They are coordinatively unsaturated and reactive compounds.

Structure: trigonal planar, short B-X distance: p-pπ bond,. dB-F is 130 pm in BF3, while dB-F is – 145 pm in BF4 anion. Lewis acid strength: BF3 < BCl3 < BBr3 < BI3. Table 37. Physical properties of boron halides. MP (°C) BP (°C) State BF3 -127,1 -99,9 gas BCl3 -107,0 +12,5 gas / liquid BBr3 -46,0 +91,3 liquid BI3 +49,9 +210,0 solid

BF3 ()

Preparation: 6 CaF2 + Na2B4O7 + 8 H2SO4 → 4 BF3 + 2 NaHSO4 + 6 CaSO4 + 7 H2O 6 KBF4 + B2O3 + 6 H2SO4 → 8 BF3 +6 KHSO4 + 3 H2O Easy way with phenyl diazonium tetrafluoroborate: Ph-N2BF4 → Ph-F + N2 + BF3

Reactions: BF3 + KF � K[BF4] (soluble salt without hydrolysis) o BF3 + H2O → BF3⋅H2O (hydrate at 20 C) + – B2O3 + 8 HF → 2 H + 2 BF4 + 3 H2O {+2 KOH} → 2 KBF4 BF3 + AsF3 → F3As→BF3 (adduct) Application: alkylation of aromatic hydrocarbons in industry: Friedel-Crafts reaction. – + Mechanism: RX + BF3 � {BF3X } + {R } {R+} + PhH � PhR + {H+} + – {H } + {BF3X } � ΒF3 + HX

BCl3 (boron trichloride) Laboratory preparation: 2 BF3 + Al2Cl6 → AlF3 + 2 BCl3 ("equilibration") Industrial preparation: B2O3 + 3 C + 3 Cl2 → 2 BCl3 + 6 CO Reaction: BCl3 + 4 H2O → H[B(OH) 4] + 3 HCl (hydrolysis)

BI3 (boron triiodide) Preparation: LiBH4 + 4 I2 → BI3 + LiI + 4 HI Reaction: BI3 + 3 H2O → B(OH) 3 + 3 HCl (violent hydrolysis)

B2F4 (diboron tetrafluoride). Gas. AX3 structure, but not coplanar arrangement.

B2Cl4 (diboron tetrachloride). Liquid (BP = 65 °C). Structure: Not planar AX3 in gas phase, but planar AX3 in solid phase. Preparation: 2 BCl3 + 2 Hg → B2Cl4 + Hg2Cl2 Reactions: B2Cl4 + Cl2 → 2 BCl3 B2Cl4 → B4Cl4 / B8Cl8, etc. (clusters: closo halogeno boranes) B2Cl4 + 2 NMe3 → [B2Cl4][NMe3]2 (adduct)

85 János Rohonczy: Inorganic Chemistry I. 6.1. Boron

B2Cl4 + LiBH4 → B2H6 / B4H10 etc. (by reduction) B2Cl4 + H2O → B2(OH)4 → (BO)n (glassy oxide, by hydrolysis) B2Cl4 + 4 EtOH → B2(OEt)4 (tetraethoxy diborane)

6.1.3. Boron oxide, boric acids, borates

B2O3 (diboron trioxide). Properties: white powder, glassy, MP = 450 °C, hard to crystallize. Structure: BO3 planar network, BO4 tetrahedral cross links. Preparation: 2 B(OH) 3 → B2O3 + 3 H2O

B(OH)3 (orthoboric acid) Structure: Hydrogen-bonds in solid phase: H O O H H B O O B H H O O H Solubility in water: very weak, monoprotic acid, pKa = 9.25. + – B(OH) 3 + 2 H2O � H3O + [B(OH) 4] Chelate forming increases the acidity: (104 times) O - O B(OH)3 + 2 B + H3O+ + 2 H2O O O HO OH

Reaction: B(OH) 3 + 3 EtOH (+cc. H2SO4) → 3 H2O + B(OEt)3 flammable, volatile ester, typical green flame.

HBO2 (metaboric acid) Structure: Polymorphs: orthorhombic – only BO3 (planar), monoclinic – BO3 / BO4, units, cubic – only BO4 units. Very weak, monoprotic acid, pKa =9,25. Preparation: B(OH) 3 → HBO2 + H2O Properties: 2 HBO2 → B2O3 + H2O

Anhydro borates: Only BO3 units, without crystal water: 3– 4– 3– – BO3 , B2O5 , B3O6 , [BO2 ]n, e.g.: CaB2O4

Hydrated borates: BO3 and BO4 units, with chemically bonded water. 5– – 2– Examples: BO4 , B(OH) 4 , [B2O(OH) 6] . Structure: 4 BO3 + 1 BO4 units. Salts: spiroborate: KB5O8⋅4 H2O ≡ K[B5O6 (OH) 4]⋅2 H2O

Na2B4O7⋅10 H2O (borax) Structure: Na2B4O7⋅10 H2O ≡ Na2 [B4O5 (OH) 4] ⋅ 8 H2, 2 BO3 + 2 BO4 units. Thermogravimetric analysis: 8 mol water is released at 100 °C, 2 moles are released at 200 °C. Partial delocalization of electrons: B–O–B (see Si–O–Si), bond distance shortening (147,5 pm in BO4, 136,6 pm in BO3, 120 pm in terminal -B=O). Glass forming oxide, used in chemical analysis: "borax-glass perl". Produce / application: 2 Mt/year. Glass fiber, borosilicate glass, porcelain, fire retardant.

NaBO3 (sodium perborate)

86 János Rohonczy: Inorganic Chemistry I. 6.1. Boron

– + Reaction: NaBO3 + 3 H2O → B(OH) 4 + H2O2 + Na (hydrolysis at 90 °C, disinfection)

6.1.4. Boron nitrogen and other boron containing compounds Structure: >B=X< is isoelectronic with >C=C< , where X = N, P, As, Sb.

(BN)x () Structure: hexagonal layers, B-N-B atoms of layers are in covering position (AAA layer). Properties: colorless insulator, inert. Exception: 2 BN + 3 F2 → 2 BF3 + N2 BN + 4 HF → NH4BF4

Laboratory preparation: Na2B4O7 + 4 NH4Cl → 4 BN + 2 NaCl + 7 H2O + 2 HCl Industrial preparation: (NH2) 2CO + 2 B(OH)3 → 2 BN + CO2 + 5 H2O BCl3 + NH3 → BN + 3 HCl

R3N·BX3 (amine borane adducts, R = alkyl, H, ... X = alkyl, H) Properties: color crystals, low MP. Amino boranes Structure: dimer: H H N Me Me Me B B N Me HH B3N3H6 (borazine, "inorganic benzene") Structure: Isoelectronic with benzene. Similar physical properties, colorless liquid. N=B bonds form delocalized aromatic π-bond system. Less aromatic.

N B

Preparation: 3 BCl3 + 3 NH4Cl → (BClNH)3 (BClNH)3+ 3 NaBH4 → B3N3H6 + 3 NaCl + 3/2 B2H6

Reaction: hydrolysis, solvolysis: B3N3H6 + 3 H2O → [BH(OH)NH2] 3 → [B(OH)NH] 3 + 3 H2 6 Complexes: analog coordination in η -B3N3Me6⋅Cr(CO)3 and dibenzene chrom (Cr(C6H6)2).

6.1.5. Other boron compounds

B-P / B-As / B-Sb, amino borane analog compound, e.g. trimer (R2P-BH2)3. B-S / B-Se bonded thio and seleno boron compounds, e.g. B(SR)3.

B2S3 (boron sulfide), Light yellow layered crystal structure: 4 and 6 member rings.

B8S16 Polymerization product of B2S3. Colorless, water sensitive crystal. Structure: similar to porphine:

B B B B

B B B B

87 János Rohonczy: Inorganic Chemistry I. 6.1. Boron

6.1.6. Boranes Derivatives: neutral, anionic, carboranes, metalloboranes.

"BH3" (mono borane): Structure: Instable monomer, H can not donate electrons. Arises from thermal decomposition of polyboranes. Immediately forms dimers and polymers. In contrast to short lifetime the structure of the monomer is known, trigonal planar. BH3 monomer can be stabilized in derivatives, e.g. BR2H (R = big alkyl ligand) which are kinetically stable compounds. The attack of the central atom is inhibited. BH3 is soluble and stable in Lewis base solvents, – + + + + e.g. Et2O, NH3 (liq.), polyethers. BH4 ion stable in MBH4 salts (M = Li , Na , K , Rb , Cs+, Be2+, Al3+).

Preparation: BCl3 + 4 NaH → NaBH4 + 3 NaCl Et2O⋅BF3 + 4 LiH → LiBH4 + 3 LiF + Et2O B2H6 + 2 LiH → 2 LiBH4 Al2Cl6 + 6 NaBH4 → 2 Al(BH4)3 + 6 NaCl Application: industrial and laboratory reducing agent in organic chemistry, preparation of shiny metal coatings, e.g. Ni mirror (instead of hypophosphite salt).

B2H6 (diborane) Structure: 2 B-H-B tree-center (2e–-3 center) bonds + – (similar to H3 molecule ion) and 4 terminal B-H (2e -2 center) bonds. H H H B B H H H Properties: self burning on air. Extremely exothermic reaction.

Laboratory preparation: 3 NaBH4 + 4 Et2O⋅BF3 → 2 B2H6 + 3 NaBF4 + 4 Et2O Industrial preparation: 2 BF3 + 6 NaH → B2H6 + 6 NaF

Reaction: B2H6 → B3H9 → B4H10 → B5H11 ...

B2H6 + 3 Cl2 → 2 BCl3 + 6 HCl, (similarly with F2) B2H6 + 6 H2O → 2 B(OH) 3 + 6 H2 (similarly with R-OH) 2 B2H6 + 2 Na → NaBH4 + NaB3H8 (oligo borane) B2H6 + 6 NH3 → 2 (NH-BH) 3 + 12 H2 (borazole) B2H6 + PbEt4 → B2H5Et + PbEt3H (substitution) B2H6 + 6 R-CH=CH2 → 2 B(CH2-CH2-R)3 (addition)

Oligo boranes – – B3H9 (triborane) Structure: 6 x (2e – 2 center) BHB-bonds, and 3 x (2e – 3 center) BH bond. (3 B atoms: 3 x 3 e–) + (9 H = 9 x 1e–) → Σ = 18 e–). There are anionic derivatives. – – – Preparation: 4 BF3 + 5 BH4 → 3 BF4 + 2 B3H8 + 2 H2 Structure: 6 x (2e–– 2 center) BH bonds + 2 x (2e–– 3 center) BHB bonds. – – Preparation: B3H9 + OH → B3H8 + H2O, or – – 2– B3H8 + OH → B3H7 + H2O Structure: 6 x (2e–– 2 center) BH bond + 1 x (2e-–3 center) BHB bond.

88 János Rohonczy: Inorganic Chemistry I. 6.1. Boron

Poly-boranes

x– Structure: BnHm , where n ≥ 4, m ≥ n, and x ≥ 0 Poly-borane homolog series with Wade-clusters: closo, nido, arachno, clado, hypo. Structure is determined by the relation between the number of skeleton members and electron population. Several substituted compounds are known, but they are mostly synthesized NOT from diborane. Borane derivatives are important model systems in the cluster chemistry. 2– BnHn (closo-boranes) Name: closo – closed. Only in anionic form. 2– 2– Structure: polyhedrons, e.g. B6H6 – octahedral, B7H7 – pentagonal bipyramidal, 2– B12H12 – icosahedral.

Preparation example: 5 B2H6 + 2 NaBH4 → Na2B12H12 + 13 H2 (Na2B12H12 is stable till 600 °C, known substituted derivatives)

Bn-2C2Hn (dicarba closo-carboranes) – 2– Structure: BH and CH groups are isoelectronic: B10C2H12 similar to B12H12 , but neutral. Preparation: nido-B10H14 + 2 Et2S → B10H12(Et2S)2 + H2 B10H12(Et2S) 2 + C2H2 → closo-1,2-C2B10H12 + 2 Et2S + H2

Isomerisation: 1,2-ortho-closo-C2B10H12 MP = 320 °C 1,7-meta-closo-C2B10H12 MP = 265 °C 1,12-para-closo-C2B10H12 MP = 261 °C

Substitution: 1,12-CH are weakly acidic, they can react with Cl-[SiMe2-O-]n- SiMe2Cl polymer: "DEXSIL" is stable till 500 °C, oxygen resistant. Used as stationer phase in GC.

BnHn+4 (nido-boranes) Structure: Not completely closed neutral molecules. Lipscomb model: nido structure is derived from the parent's closo-cluster deltahedron geometry by elimination of one cluster member (BH fragment). Note: hydrogens in B-H-B tri-center bonds do not have influence on the cluster structure. Example: B5H9 structure is based on the six member closo 2– octahedron (B6H6 ) by the elimination of one fragment. Bond types in B5H9 are 5 x B-H, and 4 x B-H-B. 11 B10H14. Lipscomb's structure is confirmed by B NMR. Industrial preparation: 5 B2H6 → B10H14 + 8 H2 (Produce ≈ 5 t/year, solid rocket fuel) + 2– Reactions: B10H14 + 2 Na → 2 Na + B10H14 (reduction, Wade-cluster) B10H14 + 2 R-Br → B10H12R2 + 2 HBr (substitution) – 2– B10H14 + 2 OH → B10H12 + 2 H2O (deprotonation)

C, S, Se, Te and metal containing borane clusters 2– B9C2H11 Structure: parent closo cluster icosahedron contains 12 members. Electron donor 2– - – behavior of B9C2H101 is similar to the cyclopentadienate, C5H5 (6 e donor). – – Preparation: closo-B10C2H12 + EtO + 2 EtOH → [7,8-B9C2H12] + B(OEt)3 + H2 + 2– Na[B9C2H12] + NaH → 2 Na + [B9C2H11] + H2 Note: closo, nido, arachno, hypo, clado boranes can contain all elements of the periodic system – except noble gases. Cluster chemistry is a dynamically growing area.

89 János Rohonczy: Inorganic Chemistry I. 6.1. Boron

BnHn+6 (arachno-boranes) Structure: by elimination of two vertices of the parent deltahedral closo structure. They are more open than nido-boranes. 2– B4H10 (arachno-tetraborane) Elimination of two vertices from the parent closo B6H6 structure. Bond types: 6 x (2e––2 center) BH bonds + 4 x (2 e–– 3 center) BHB bonds. – 2– Anions: [BnHn+5] , [BnHn+4] – – Reaction: B4H10 + OH → B4H9 + H2O – Electronic structure of B4H9 equals to B4H10, but bond types are different: 5 x (2 e–– 2 center) BH bonds + 4 x (2 e–– 3 center) BBB bonds.

B10H16 (arachno-decaborane) and its isostructural deprotonated anion: Preparation: nido-B10H14 + 2 Na → Na2[B10H14]

BnHn+8 (hypo-boranes) Only derivatives are known, e.g. B5H9 (PMe3)2

BnHn+10 (clado-boranes) Only derivatives are known.

Conjuncto-boranes

Clusters with complicated, versatile structures: common boron atom (spiro position), B-B σ-bond, common B-B edge, common B3 face, common B4 tetrahedron. Huge number or known compounds.

90 János Rohonczy: Inorganic Chemistry I. 6.2. Aluminium, Gallium, Indium, Tallium

6.2 Aluminium (Al), Gallium(Ga), Indium (In), Tallium(Tl) Valence shell: ns2np1 Discovery: Al Oersted, Wöhler (1825): AlCl3 + K/Hg → Al + KCl (not pure), Bunsen (1854): electrolysis of NaAlCl4 – pure but very expensive, Héroult and Hall (1886): electrolysis of Al2O3 dissolved in molten Na3AlF6. Name: alumen (Latin) KAl(SO4) 2⋅12 H2O – known medicine. Ga (1875) de Boisbaudran: by detection of two violet spectrum lines. Name: Gallia (France). Mendeleev's prediction: eka-aluminium. In (1863) Name: indigo blue color spectrum line. Tl (1861) Name: intensive green spectrum line (thallos (Greek) = green twig). Occurrence: Al – 3rd abundant after O and Si in the Earth's crust. Minerals: feldspar, mica, talc, montmorillonite, kaolinite [Al2(OH)4Si2O5], cryolite (Na3AlF6), spinel (MgAl2O4), beryl (Be3Al2Si6O18), turquoise (CuAl6(PO4)4(OH)8·4H2O), corundum (Al2O3) (ruby, sapphire), bauxite (AlOx(OH)y). Produce: 80 Mt/year. Ga – ZnS/bauxite/cinder. There is no own mineral. Only 19 ppm (like Pb), Produce: 10 t/year. In – ZnS/PbS volatile roast smoke, rIn ≈ rZn. Only 0,21 ppm (like Sb). Produce: 50 t/year. Tl – PbS/magmatic K-minerals, feldspars: rTl ≈ rPb and rTl ≈ rRb. Produce: < 5 t/year.

Preparation / Application Al preparation: A) Bayer method based on bauxite with low SiO2 content: 45-60 % Al2O3⋅3 H2O (hydrargillite) and Al2O3⋅H2O (diaspore). 5-10 % Fe2O3, max. 1-5 % SiO2, 1-2 % TiO2. Preparation of alumina (Al2O3): + – Al(O)OH + NaOH → Na + [Al(OH) 4] + red mud (Red mud: Fe2O3, TiO2, V2O5, Al2(SiO3)3) – – [Al(OH) 4] � Al(OH) 3 + OH (by dilution, cooling and crystallization) 2 Al(OH) 3 → Al2O3 + 3 H2O (by calcination) B) Electrolysis: 4,5 V, 105 A, 950 °C: 15 kWh/kg Al. electrolyte: Na3AlF6 + 2-8% Al2O3 + few CaF2/AlF3 Al(OH) 3 + 3 NaOH + 6 HF → Na3AlF6 + 6 H2O (synthetic cryolite)

C) Dry process based on bauxite with 5-12% SiO2 Al2O3 + Na2CO3 + CaCO3 → Ca(AlO2) 2 / NaAlO2, aluminates.

D) Process based on kaolinite/clay + dry HCl → AlCl3, distillation, sublimation. Expensive. Al application: Alloys (Al + Cu / Mn / Si / Mg / Zn etc) – better mechanical properties, better corrosion resistance, higher MP, etc. Constructional metal: buildings, vehicles, airplanes, electric wires, packing materials, pigments, etc. Mirrors. Ga – It is concentrated in the base solution of Bayer process till 300:1 ratio. Electrolysis on Hg cathode, dissolution by NaOH → Na[Ga(OH)4] (Na-gallate). Electrolysis again on steel cathode: 99.9% pure Ga. Ga → GaMe3, distillation / pyrolysis / zone melting → Ga (ultra pure). János Rohonczy: Inorganic Chemistry I. 6.2. Aluminium, Gallium, Indium, Tallium

Properties: shiny blue, wets the glass and porcelain, MP =30°C. 2+ Application: GaAs semiconductor. MgGa2O4 (Mn ) exhibits fluorescence on UV light (Xerox). Manometers, metal baths, low MP solders, thermometers. In – impurity in PbS, ZnS. Preparation by electrolysis. Properties: soft, silver shiny metal. Application: low MP alloys, semiconductors, solder for high vacuum glass equipments, neutron absorber. ITO – indium-tin-oxide: transparent, electric conductor in LCD panels. Tl – Preparation from PbS, ZnS, CdS and selenide minerals: + 2– 1) Tl2S / PbS + H2SO4 → 2 Tl + PbSO4 + SO4 + H2S 2) Tl+ + HCl → TlCl + H+ 3) Tl2SO4 solution: electrolysis on Pt cathode: metal Tl. Properties: very toxic both in elementary form and in compounds. Application: Tl2SO4 – former rat poison. TlBr/TlI single crystals – less soluble in water: transparent far IR detectors. Saturated 1:1 solution of Tl-formiate / Tl-malonate: high density water solution (Clerici solution): ρ = 4.324 g/cm3 – used to separate of minerals. Table 38. Atomic properties: odd number of protons: free stable isotopes 27Al 69Ga 71Ga 113In 115In 203Tl 205Tl I=5/2 I=3/2 I=3/2 I=9/2 I=9/2 I=1/2 I=1/2

Physical properties Low MP, soft, silver shiny white metals with good electric conductivity. Crystal structure: close packed with coordination number = 12. Exception: Ga unit cell contains Ga2 units. Complicated metal ⇔ molecule equilibrium even in liquid phase. Low MP (like of Hg), dilatation during freezing (like Ge, Sb, Bi, H2O).

Chemical Properties They differ very much from boron. Reactive elements at moderate high temperature, and they form cations in water. There are no volatile hydrides, nor homofragment clusters.

Al – Binary compounds with non-metal elements: AlN, Al2S3, Al2X6, Al2O3 – latter forms protective coating on the Al metal surface. Reactions: amphoteric metal in water solution: + – Al + NaOH + 5 H2O → Na + [Al(OH)4(H2O)2] + 3/2 H2 + 3+ Al(OH)3(H2O)3 + 3 H3O → [Al(H2O)6] + 3 H2O (6 coordinated in acids) – – Al(OH)3(H2O)3 + OH → [Al(OH)4(H2O)2] + H2O (in base solution) Amphoteric at anhydrous conditions: Al2O3 + CaO → Ca(AlO2)2 (Ca-aluminate, acts as acid) Al2O3 + 3 SiO2 → Al2(SiO3)3 (Al-silicate, acts as base) Hydrolysis of binary compounds: Al2S3 + 6 H2O → 2 Al(OH)3 + 3 H2S AlN + 3 H2O → Al(OH)3 + NH3 Al4C3 + 12 H2O → 4 Al(OH)3 + 3 CH4 – 2– – Hydrolysis of Cl , SO4 and NO3 salts are not complete: 2 [Al(H2O)6]Cl3 → Al2O3 + 9 H2O + 6 HCl (not only crystal water departs)

92 János Rohonczy: Inorganic Chemistry I. 6.2. Aluminium, Gallium, Indium, Tallium

Ga – Similar to Al: Ga2O3 is amphoteric (but gallates are more stable – more acidic) + – Ga + NaOH + 5 H2O → Na + [Ga(OH)4(H2O)2] + 3/2 H2 (water soluble)

In – In + NaOH(aq) → insoluble. In2O3 is very weakly amphoteric, it is base anhydride. Tl – forms bases. Tl+ is stable ion in +1 oxidation state. TlOH is strong base. TlN3 ( azide) is non explosive. Tl2CrO4, Tl2S and TlCl are insoluble salts in water (like Ag salts). TlCl3 is molecular compound.

6.2.1. Al/Ga/In/Tl-hydrides

AlH3 (aluminum-hydride) Properties: colorless, non volatile crystalline polymer. Structure: Al-H-Al bond. α-Al-hydride: 1 Al is octahedrally surrounded by six H atoms.

Preparation: 3 LiAlH4 + AlCl3 (+Et2O) → 4 [AlH3⋅Et2O]+ 3 LiCl (instable in benzene) Reaction: decomposition at 150 °C. Strong reducing agent, violent reaction with water, forming H2 gas.

LiAlH4 + [NMe3⋅HCl] + NMe3 → [AlH3⋅2 NMe3] +LiCl + H2 (adduct)

GaH3 (gallane) Properties: viscous liquid, MP= -15 °C, decomposition to elements at 20 °C. Forms stable, crystalline adducts.

InH3 / TlH3 Instable compounds even at low temperature.

Table 39. MP data of LiMH4, where M = B, Al, Ga, In, Tl

LiBH4 LiAlH4 LiGaH4 LiInH4 LiTlH4 380 °C 100 °C 50 °C 0 °C 0 °C

Laboratory preparation: 4 LiH + AlCl3 → LiAlH4 + 3 LiCl Industrial preparation: Na + Al + H2 → NaAlH4 (in THF, 140 °C, 350 bar, 99%) Produce: several t/year.

Application: MAlH4 (M = Li, Na) strong reducing agent, agent. Application in organic chemistry: → alkanes, aldehides / → alcohols, amides / cyanides → amines, etc. i New, similar reagent: Bu2AlH (cheaper, more safe).

6.2.2. Al/Ga/In/Tl halides

Al/Ga/In/Tl monohalides

Stable compounds only at high temperature, except fluorides. Disproportion at room temperature: 3 AlX � AlX3 + 2 Al. Strong F-bridge in fluorides, high MP solids. GaX, InX are stable only at high temperature, but Tl(+1)X compounds are stable at room temperature.

Al trihalides

AlF3 (aluminum trifluoride) Properties: stable. Sublimation at 1200 °C, strong covalent Al-F bonds, low tension, insoluble, [AlF6] octahedrons in crystal lattice bonded by F-bridges.

Preparation: AlCl3 + BF3 � AlF3 + BCl3 Gas phase equilibrium pushed in forward direction by the low tension of AlF3. Complex fluorides are very stable compounds. Preparation of synthetic cryolite: Al(OH) 3 + 6 HF + 3 NaOH → Na3AlF6 + 6 H2O

93 János Rohonczy: Inorganic Chemistry I. 6.2. Aluminium, Gallium, Indium, Tallium

AlCl3 (aluminum trichloride) Properties: Lewis acid, important catalyst (Friedel-Crafts catalyst of electrophilic substitution of aromatic hydrocarbons). Preparation: Al + 3/2 Cl2 → AlCl3 (expensive process) Al2O3 + 3 C + 3 Cl2 → 2 AlCl3 + 3 CO (by reductive chlorination) Alternative industrial preparation using phosgene at 700 °C: Al2O3 + 3 COCl2 → 2 AlCl3 + 3 CO2 Properties: treefold-coordinated monomers in vapor over 200 °C. Fourfold-coordinated dimer Al2Cl6 molecules below 192.4 °C. Sixfold-coordinated layered lattice in crystal phase at 20 °C. Application: catalyst, intermediate, methylation agent: AlCl3 + 3 LiMe → AlMe3 + 3 LiCl (self burning)

Al2Br6 / Al2I6 Structure: dimers even in solid phase.

Ga / In / Tl trihalides

– Volatile chlorides, forming adducts, e.g. GaCl3⋅L, stable complexes with anions: (GaCl4 ). Known solid mixed halides. Known TlX3, but they are thermally less stable. NaTlF4 is stable.

6.2.3. Al/Ga/In/Tl oxides, hydroxides and complicated oxides

α-Al2O3 (corundum) Properties: very hard (Mohs hardness = 9), MP = 2045 °C, non volatile, inert, insulator. Structure: Al is octahedrally coordinated by six O atoms, 2/3 of holes are occupied by Al atoms. Other polymorphs: several Al2O3 and Al(O)OH: α-diaspore (in bauxite), γ-diaspore, etc. 3+ – Preparation: Al + 3 OH → Al(OH)3 (gel) 2 Al(OH)3 → Al2O3 + 3 H2O (dehydration, calcination) Application: polishing material, ceramics, metal Al product., Al cement, desiccator, catalysts. Jewels: ruby (CrIII, red), sapphire (FeII/FeIII/TiIV: blue), topaz (FeIII, yellow), emerald (B/ III III Cr /V , green). Synthetic (from Al(OH)3 → at 1200 °C): catalysts. Al2O3 + ZrO2 "Saffil" (1974) contains 3 µm x 2-5 cm long threads mixed to molten metals resulting in very high mechanical stability against stretching.

α,β,γ-Ga2O3 Similar to Al2O3, but β has special structure.

In2O3 / In(O)OH. They are known compounds.

Tl2O Black solid. Preparation: By heating of Tl2CO3.

Tl2O3 Properties: Brown/black solid. Strong oxidizer agent. Cement Composition: 70% CaO + 20% SiO2 + 5% Al2O3 + 3 % Fe2O3 (low Na2O, K2O, MgO, P2O5 content is expected).

Fastening: 2 Ca2SiO4 + 4 H2O → 3 CaO ⋅ 2 SiO2 ⋅ 3 H2O + Ca(OH)2 2 Ca3SiO5 + 6 H2O → 3 CaO ⋅ 2 SiO2 ⋅ 3 H2O + 3 Ca(OH)2

Preparation: CaCO3 (limestone) + alumino silicates (clay) + SiO2 (sand). Milling, sifting, firing (at 1500 °C), milling, doping by CaSO4. Produce: 700 Mt/year.

94 János Rohonczy: Inorganic Chemistry I. 6.2. Aluminium, Gallium, Indium, Tallium

6.2.4. Ternary and more complicated oxides II III MgAl2O4 (spinel) Structure: A B2 O4, contains AO4 and B4O4 units. Properties: structure dependent electric and magnetic properties – magnetite, ferrites.

Na2O⋅11 Al2O3 (Na β-aluminate) Discovered: 1967. Structure: corundum – Na+-ion – corundum layer structure. Application: solid state electrolyte (solid ionic conductor): membrane in Na/S batteries. 18+ Ca3Al2O6 (Ca aluminate) Structure: cyclic [Al6O18] cation. Important component (11%) in Portland cement. Other components: Ca2SiO4 (26%), Ca3SiO5 (51%), Ca4Al2Fe2O10 (1%).

Al2(OH)2Si2O5 (kaoline) Used to make porcelain. Structure: Si–O–Al–O–Si network. Zeolites. Al silicates containing holes. Application: catalysts, filters, adsorbers, etc.

6.2.5. Al/Ga/In/Tl chalcogenides

Properties: solids Table 40. Aluminum chalcogenides.

Al2S3 Al2Se3 Al2Te3 White crystal gray dark gray

Preparation: 2 Al + 3 Se → Al2Se3 (direct synthesis at 1000 °C) Reaction: Al2S3 + 6 H2O → 2 Al(OH) 3 + 3 H2S (hydrolysis)

Ga sulfides / selenides / tellurides

GaS Properties: yellow crystals. Structure: Ga-Ga bond, layered structure.

Ga2S3 Structure: α, β, γ polymorph crystals, incomplete wurtzite lattice. In/Tl chalcogenides. Known compounds, many known structures. Important properties: conductors or semiconductors, optical conductors, superconductors, etc. Under study.

6.2.6. Binary III-V compounds of Al/Ga/In/Tl

Table 41. Simple structures, isoelectronic with Si and Ge, III-V semiconductors B Al Ga In Tl N B-N/s w w w Tl3N/TlN3 2200 °C P s s s s Tl3P/TlP5 2000 °C 1465 °C 1070 °C As s s s s Alloy 1740 °C 1238 °C 942 °C Sb - s s s Alloy 1060 °C 712 °C 525 °C Notes: s sphalerite lattice (like diamond), w wurtzite lattice (hexagonal) and MP data. MP is decreasing in top-down order. Ternary systems: tunable energy gaps. Color LEDs. Preparation: As + Ga → GaAs Ga + NH3 → GaN + 3/2 H2 In2O3 + 2 NH3 → 2 InN + 3 H2O Reaction: AlN is inert, GaN is soluble in bases, InN is soluble in acids/bases. AlP + 3 H2O → Al(OH) 3 + PH3 (hydrolysis)

95 János Rohonczy: Inorganic Chemistry I. 6.2. Aluminium, Gallium, Indium, Tallium

6.2.7. Al/Ga/In organic compounds

Typical electron deficient three center bonds, like in B2H6.

AlR3 R = alkyl or aryl group. Structure: Al2R6 is dimer. NMR data: isolated signals at -75 °C, but one signal and fast exchange at 20 °C. Fast rearrangement processes. Properties: colorless, volatile liquids, low MP, high reactivity, self-burning properties. Several known derivatives with mixed ligands. Precursors of special syntheses.

Table 42. Melting points of AlR3 type aliphatic compounds n Al2Me6 Al2Et6 Al2 Pr6 MP = 15 °C MP = -53 °C MP = -107 °C BP = 126 °C Industrial methylation agent Laboratory preparation: 2 Al + 3 HgMe2 → Al2Me6 + 3 Hg 2 Al + 3 HgPh2 → Al2Ph6 + 3 Hg Industrial preparation: 3 Me-Cl + 2 Al → Me3Al2Cl3 � 1/2 Me4Al2Cl2 + 1/2 Me2Al2Cl4 3 Me4Al2Cl2 + 6 Na → 2 Al2Me6 + 2 Al + 6 NaCl 2 Al + 3 H2 + 2 Al2Et6 → 6 Et2AlH (mixed types, ligand exchange) 2 Et2AlH + 2 CH2=CH2 → Al2Et6 (by addition)

Another addition but without H2: AlEt3 + CH2=CH2 → Et2AlCH2CH2Et → Al[(CH2)nEt]3 (n ≈ 15) + 3+ Al[(CH2) nMe] 3 + O2 + H3O → Al (aq) + CH3-(CH2)n-OH Ziegler-Natta catalysis: Polymerization of alkenes (1963, Nobel-price)

Al2Et6 + CH2=CH2 + TiCl4 → [-CH2-]n (85-95% efficiency) Low temperature, low pressures. Produce: million t/year.

Ga/In/Tl organic compounds. Less studied, less important compounds in industry. MR3 type compounds are already not dimers. MPh3 is solid. Reactivity of M-C bond is decreasing in top-down direction: Al > Ga ≈ In > Tl. For example: TlR2Cl is stable on air, there is no hydrolysis, [TlMe2+] is linear and water soluble ion.

Preparation: 2 Ga + 3 HgR2 → 2 GaR3 + 3 Hg Al2R6 + 2 GaCl3 → 2 GaR3 + 2 AlCl3

Properties: low MP, highly fluid, flammable liquid. InR3 and TlR3: higher MP and BP.

6.2.8. Al / Ga / In / Tl and nitrogen bounded compounds Many Al-N containing compound families are known. Mostly complicated, big polyhedrons, closed structures: 21 AlMe3 + 21 NH2Me → 7 [Me2AlNHMe]3 Simple structures in gas phase, e.g. Me2AlNHMe molecule has C3v symmetry. Ga/In/Tl-N bonds are known just between cluster fragments as tree center bonds.

96 János Rohonczy: Inorganic Chemistry I. 6.2. Aluminium, Gallium, Indium, Tallium

6.2.9. Aluminum salts

Al2(SO4)3⋅18 H2O Properties: white, water soluble crystals, hydrolyses in water. Application: impregnation, water cleaning, tanning.

KAl(SO4)2⋅12 H2O (potassium alum). Typical double sulphate salt. Its solution contains I III hydrated cations. Other isomorphic structures: double salts are M M (SO4)2⋅2 H2O, where I + + + + III 3+ 3+ 3+ M = K , Na , Tl , NH4 , etc. and M = Al , Cr , Fe , etc. Application: tanning. Similar alum compounds can contain MIII= Ga3+, In3+, Tl+3 ions (in the latter case TlIII ions are stabilized in crystalline form).

97 János Rohonczy: Inorganic Chemistry I. 7. Be, Mg and Alkaline Earth Metals

7. Be, Mg, Ca, Sr, Ba, Ra – Alkaline Earth Metals

7.1. Elements

Valence shell: ns2. All of them are metals. Be differs in several properties: it forms mostly 2+ covalent compounds and it is amphoteric. On the other side its aqua complex [Be(H2O)4] is known as well. Mg is between the and the ionic alkaline earth metal elements (Ca, Sr, Ba, Ra). Table 43. Physical properties of the 2nd column elements Symbol Name MP(°C) BP(°C) EN Lattice Color of flame Be Beryllium 1278 2970 1.67 metallic - Mg 639 1090 1.31 metallic - Ca Calcium 839 1484 1.00 metallic brick red Sr 769 1384 0.96 metallic carmine Ba 726 1140 0.89 metallic fallow green Ra Radium 700 1737 0.90 metallic colorful red

Name and discovery: Be – Beryl mineral (berillus - Latin). Discovery: Haüy (1789). Wöhler (1828): BeCl2 + 2 K → Be(metal)+ 2 KCl. Bussy (1890): pure metal by electrolysis

Mg – Magnesia (Greek, MgCO3), other known compound: talc. Davy (1808): pure metal by electrolysis.

Ca – Calx, calcis (Latin): CaCO3 (limestone), Ca(OH)2 (slacked lime). Plaster: in Egypt: CaSO4⋅2 H2O (gypsum), in Rome: lime + sand. Davy (1808): pure metal by electrolysis. Sr – Strontianite mineral (Strontian mine in Scotland). Discovery: Crawford (1787-90). Davy (1808): pure metal by electrolysis.

Ba – Barite – heavy (Greek). Mineral: barite (BaSO4), known from XVI-th century. Davy (1808): pure metal by electrolysis. Ra – Radiation (Latin). Discovery: Pierre and Marie Curie (1898) from uraninite - pitchblende. M. Curie (1910): pure metal by electrolysis.

Occurrence: Be – Rare, similar to Li and B: only 2 ppm. Minerals: beryl (enriches well) variants: morganite (pink beryl), emerald (green beryl), aquamarine (turquoise), heliodor (golden beryl), red beryl: Be3Al2Si6O18. Mg – Very common (6th before Na). Insoluble compounds: carbonate, sulphate, silicates. Minerals: MgCa(CO3)2 (dolomite), MgCO3 (magnezite), MgSO4⋅7 H2O (epsomite), (Mg,Fe)SiO4 (olivine), Mg3Si4O10(OH)2 (talc), Mg3Si2O5 (OH)4 (asbestos), MgAl2O2 (spinel), Mg(OH)2 (brucite). Plants: chlorophyll.

Ca – Very common element (5th). Minerals: microcrystalline CaCO3 (limestone), rhombohedral calcite, orthorhombic aragonite. CaSO4⋅7 H2O (gypsum), CaSO4 (anhydrite), CaF2 (fluorite), Ca5(PO4)3(OH) (apatite). János Rohonczy: Inorganic Chemistry I. 7. Be, Mg and Alkaline Earth Metals

Sr/Ba 14-th and 15-th (between S and F). Minerals: SrSO4 (celestite), SrCO3 (strontianite), BaSO4 (barite). Ra It is contaminant of uranium. Very rare: 10 tons U mineral contains 1 mg Ra. General physical properties: silver white, shiny, relatively soft metals.

7.1.1. Preparation and application Be Preparation: from beryl, at 750 °C. Be3Al2Si6O18 + 3 Na2SiF6 → 3 BeF2 + 2 AlF3 + 6 NaF + 9 SiO2 BeF2 + 2 NaOH → Be(OH) 2 + 2 NaF (precipitation) BeF2 + Mg → Be + MgF2 Electrolysis of BeCl2 / KCl (KCl decreases the MP of the electrolyte) Properties: light, hard metal with high MP. Stable on air: protective oxide coating is similar to the Al surface. Application: Alloys: Ni + 2% Be is flexible even at high temperature: springs, clamps, connectors. Cu + 2% Be is six times harden than copper, good electric conductor, flexible, non magnetic, does not sparking. Metal Be: n-moderator, n-reflector in nuclear reactors. X-ray window: transparent. Ra + Be → n (discovery of neutron, Chadwick: 9Be(α,n) 12C). Mg Preparation: silicothermic Pidgeon process: 2 CaO ⋅ MgO + FeSi → 2 Mg + Ca2SiO4 + Fe (and distillation of Mg) Electrolysis of molten MgCl2⋅x H2O (products: Mg + Cl2) Produce: 615,000 t/year (2005). Properties: light, stable on air, soft metal. Application: engine block, aircraft body and wing: typical alloy: 90% Mg, 7% Al, 2% Zn, 1% Mn + Pr/Nd/Th. Other: 95% Al, 5% Mg: hard, corrosion resistant, weldable. Cathode corrosion protector, oxygen getter, reducing agent: preparation of metal Be/Ti/Zr/Hf/U.

Ca Preparation: electrolysis of molten CaCl2 salt. Properties: weakly protective oxide/nitride coating, machinable. Application: Hardening of Al alloys, graphitization of cast iron, removing Bi from lead. Getter: remove N2/O2 traces from Ar gas. Reduction agent: preparation of Cr/U/Zr/Th. Preparation of CaH2. Removing O, S and P by CaO during the preparation of steel. Produce: ca. 1000 t/year.

Sr / Ba Preparation: 3 SrO / BaO + 2 Al → 3 Sr/Ba + Al2O3 or electrolysis of molten BaCl2/SrCl2. Application: getter, Ni-Ba alloy: ignition plug in cars. Ra Application: α-radiation source, neutron source. Table 44. Stable isotopes of 2nd column elements. 236Ra – common natural isotope. Symbol Stable isotopes Density (g/cm3) Ionization energy (kJ/mol) Be 1 1,848 899 Mg 3 1,738 737 Ca 6 1,55 590 Sr 4 2,63 549 Ba 7 3,62 503 Ra radioactives: (4) 5,5 (236Ra) 509 (236Ra)

99 János Rohonczy: Inorganic Chemistry I. 7. Be, Mg and Alkaline Earth Metals

7.1.2. Chemical properties Beryllium is less reactive, others reactivity and electro positive characters are increasing in Mg < Ca < Sr < Ba order.

Be + H2O → no reaction Mg + 2 H2O → H2 + Mg(OH)2 (at ca. 100 °C) Ca + 2 H2O → Ca(OH)2 + H2 (vehement reaction) Be: 2 Be + O2 → BeO Be(powder) + O2/N2 → BeO / Be3N2 (on air) Be + X2 → BeX2 Be + S → BeS 2 Be + C → Be2C 3 Be + 2 NH3 → Be3N2 + 3 H2 Be + H2 –/→ (no reaction, but BeH2 is known compound) BeMe2 → BeH2 + C2H4 (preparation of BeH2) Be + cc. HNO3 –/→ (no reaction, passivated) Be + 2 HCl(aq) → BeCl2 + H2 (soluble in diluted acids) Be + 2 NaOH(aq) + 2 H2O → Na2 [Be(OH) 4] + H2 Mg: Mg + X2 → MgX2 (burning) Mg + O2/N2 → MgO + Mg3N2 (burning) Mg + H2 → MgH2 (direct reaction) Mg + H2O(vapor) → Mg(OH)2 + H2 Mg + 2 MeOH → Mg(OMe)2 + H2 Mg + RX → R-Mg-X (Grignard reaction) Ca / Sr / Ba: More vehement reaction than with Mg. 3 Ca + N2 → Ca3N2 Ca + H2 → CaH2 (stable, used as desiccator) Ca + 2 C → CaC2 (less stable) 2 Ca + O2 → 2 CaO (simple, stoichiometric oxide) Ba + O2 → BaO2 (Ba forms peroxide) Ca + NH3 (liq.) → physical, reversible solution, blue color liquid (instable) Ca + 2 NH3 → Ca(NH2) 2 + H2 (heavy metal catalyzed decomposition) Thermal stability of oxoacid salts depend on the polarizability behavior of the cations. Small atomic core radius and big charge ⇒ strong polarizer ⇒ easy decomposition (instable compound). See decomposition temperature of carbonates. Increasing atomic number increases the core size. Table 45. Decomposition temperature of 2nd column elements

BeCO3 MgCO3 CaCO3 SrCO3 BaCO3 250 °C 540 °C 900 °C 1289 °C 1360 °C – 2– 2– Solubility of salts: less soluble F , CO3 , SO4 salts compared to alkaline metal salts. Consequence: CaF2, CaSO4, CaCO3 are precipitates, minerals. Exception: BeF2 is well soluble 2+ in water, while [Be(H2O)4] is stable aqua complex.

100 János Rohonczy: Inorganic Chemistry I. 7. Be, Mg and Alkaline Earth Metals

7.2. Compounds

7.2.1. Be, Mg and alkaline earth metal hydrides

Structure: BeH2 is covalent polymer: amorphous, white, stable on air, forms H2 in acids. MgH2 has rutile lattice, like TiO2. Covalent bonds. – CaH2 ... BaH2 are rhombic solids. Ionic bonds: H ions. Thermal stability: covalent: BeH2 (250 °C) > MgH2 (85 °C) ionic: CaH2 (885 °C) > SrH2 (585 °C) > BaH2 (230 °C)

Preparation: M + H2 → MH2 (except: M ≠ Be) (white crystals) BeCl2 + 2 LiH → BeH2 + 2 LiCl Reaction: CaH2 + 2 H2O → Ca(OH)2 + 2 H2 (drying of CHCl3)

7.2.2. Be, Mg and alkaline earth metal halides

BeX2 Preparation: Be + X2 → BeX2 (where X = Cl, Br, I) BeO + C + Cl2 ⇌ BeCl2 + CO (with F2 as well)

BeF2 Structure: BeF2 is glassy, {BeF4} units → crystalline BeF2 (like SiO2 quartz) Preparation: (NH4) 2BeF4 → BeF2 + 2 NH4F (water soluble)

BeCl2 Structure: infinite chains like BeH2. Decomposition with Lewis bases. Stable adducts or complexes: [BeCl2⋅2 Et2O], [Be(H2O)4]Cl2 Application: preparation of Be organic compounds.

MgF2 Structure: rutile lattice with sixfold-coordinated {MgF6}.

CaF2 ... BaF2 Structure: fluorite lattice with eightfold-coordinated {CaF8} (big core), high MP, low solubility. CaF2 is F2 source (obtain: 5 Mt/year) CaF2 properties: white crystal, high MP = 1418 °C, insoluble in water.

MX2 M=Mg/Ca/Sr/Ba, X=Cl/Br/I. Properties: white, water soluble, hygroscopic salts: form hydrates and aqua complexes. Relatively low MP: 700 °C-900 °C (lowest Br/I salts). Eutectic temperature: CaCl2 / H2O = -55 °C. CaCl2⋅6 H2O ⇌ CaCl2 + 6 H2O, desiccator, 2+ – but: CaCl2 + 6 NH3 → [Ca(NH3) 6] + 2 Cl , it is not suitable to dry NH3. BaCl2: non-hygroscopic, water soluble, very toxic. Structure: lower coordination number.

Preparation: M + X2 → MX2 (burning) CaO + 2 HCl → CaCl2 + H2O CaCO3 + 2 HCl → CaCl2 + H2O + CO2 Application: CaCl2, road salting.

7.2.3. Be, Mg and alkaline earth metal oxides and hydroxides

Preparation: MCO3 → MO + CO2 M(OH)2 → MO + H2O BeO Properties: white crystals (wurtzite lattice), high MP (2530 °C), fire-proof, low tension, good thermal conductivity (like metals), water insoluble. There is no Be-peroxide.

101 János Rohonczy: Inorganic Chemistry I. 7. Be, Mg and Alkaline Earth Metals

Be(OH) 2 BeX2 + 2 H2O → Be(OH)2 + 2 HX, very low solubility. Amphoteric. + 2+ Reaction: Be(OH)2 + 2 H3O � [Be(H2O)4] – 2– – 2– 2 Be(OH)2 + 2 OH → [Be2(OH)6] + {2 OH } → 2 [Be(OH)4] MgO White, high MP (2826 °C), fire-proof, good thermal but bad electric conductor. Application: fire-proof melting pots. MgO + C + Cl2 → MgCl2 + CO.

Mg(OH)2 Mineral: brucite, low solubility. Application: MgO + MgCl2⋅6 H2O + sawing + water → "magnesia cement". CaO (quicklime, caustic lime) White solid, high MP (2613 °C). Preparation: CaCO3 → CaO + CO2 (Produce: thousand Mt/year) Application: CaO + H2O → Ca(OH)2. Portland cement. Steel industry: removal of P/S/Si (as dross), produce of silicothermic Mg. Preparation of glass and CaC2.

Ca(OH)2 (slacked lime) Weakly soluble white solid. Preparation: CaO + H2O → Ca(OH)2 Application: Cheap base: neutralize of acids, water softening. Paper industry: Ca(OH)2, precipitated CaCO3 powder (Ca(OH)2 + CO2→). CaCO3 used in toothpaste as well.

SrO / BaO Better soluble bases: Sr(OH)2, Ba(OH)2, the latter is strong base.

Sr(OH)2 Ba(OH)2 + CO2 → BaCO3 + H2O. Carbonates are less or insoluble.

Ba(OH)2 Basic strength order: Be(OH)2 Mg(OH)2 < Ca(OH)2 ≈ Sr(OH)2 < Ba(OH)2

7.2.4. Be, Mg and alkaline earth metal carbonates

MgCO3 magnezite, (Ca,Mg)CO3 dolomite, CaCO3 limestone, chalk, marble, calcite, aragonite. Low solubility in water, but: CaCO3 + H2O + CO2 ⇌ Ca(HCO3)2 Water hardness: dissolved Ca2+, Mg2+ ions in water. German degrees: 1 °dH ≅ 10 mg CaO. Sweet water: 12-17 °dH. – - 2– Temporary hardness: HCO3 , permanent hardness: Cl , SO4 etc. Water softening: temporary hardness: Ca(HCO3)2 + Ca(OH)2 → 2 CaCO3 + 2 H2O MgCO3 + Ca(OH)2 → Mg(OH)2 + CaCO3

Permanent hardness: CaCl2 + Na2CO3 → CaCO3 + 2 NaCl (other method: using Na3PO4) Ion free water: by ion exchange resins. In laboratory: double distilled water in quartz equipment.

7.2.5. Be, Mg and alkaline earth metal peroxides, ozonides

CaO2 / SrO2 / BaO2 (Peroxides) Be/Mg peroxides are not known yet.

Ca(O2)2 / Sr(O2)2 / Ba(O2)2 () yellow solids.

CaO3 / BaO3 (Ozonides) Not pure compounds. Preparation: Ba + O2 → BaO2 Application: BaO2 + H2SO4 → BaSO4 + H2O2 (cc. )

7.2.6. Be, Mg and alkaline earth metal salts

CaC2, CaNCN Calcium carbide, calcium cyanamide.

102 János Rohonczy: Inorganic Chemistry I. 7. Be, Mg and Alkaline Earth Metals

CaCl(OCl)⋅Ca(OH)2 Calcium , bleach-powder, whitening, oxidizer agent. CaCl(OCl) + CO2 → CaCO3 + Cl2, produces Cl2 without water. CaSO4⋅2 H2O Gypsum, gypsum alabaster. Obtain: 100 Mt/year. CaSO4⋅1/2 H2O Hemihydrate: from gypsum at 150 °C (reversible process). CaSO4 Anhydrite: from hemihydrate at 200 °C (irreversible change) At very high temperature: CaSO4 → CaO + SO3 MgSO4⋅7 H2O (Epsom salt) Magnesium sulphate, osmotic purgative. CaS, SrS, BaS Fluorescent pigments (with 1% transition metal contaminations) Ca(NO3)2, Ba(NO3)2 Pyrotechnical compounds. Ca3(PO4)2 Tertiary phosphate: insoluble in acids/water. CaHPO4 Secondary phosphate: already soluble in acids. Ca(H2PO4)2 Primary phosphate: best soluble in water.

7.2.7. Be, Mg and alkaline earth metal complexes

Be Unique, stable, volatile compounds: carboxilates [OBe4(RCO2)6], where R = H, Me, Et, Pr, Ph, etc. Structure of R = Me compound: central oxygen atom is surrounded by four tetrahedrally oriented Be atoms, each of them are tetrahedrally coordinated by the oxygen atoms of the acetate groups. MP = 285 °C, BP 330 °C. Mg Stable natural complex: chlorophyll. Ca/Sr/Ba Weak complex forming agents. Only crown ether complexes are stable.

7.2.8. Be, Mg and alkaline earth metal organometallic compounds

Ca/Sr/Ba Very reactive metals. Covalent elemento organic derivatives are not known. M-C bond is very polar, practically ionic. Stable alkylation/arylation agents. Preparation: HgR2 + Ca → CaR2 + Hg, at low temperature.

BeR2 Preparation: BeCl2 + 2 LiR → BeR2 + 2 LiCl BeCl2 + 2 R-Mg-Cl → BeR2 + 2 MgCl2 Known: Be(η5-Cp)Y, where Y = Me, acetylene, or η1-Cp.

MgR2 Preparation: RMgCl + LiR → MgR2 + LiCl (not important family) RMgX Grignard reagent (1900). Very important reagent in the synthetic organic chemistry. Preparation: Mg + R-X (abs. Et2O) → R-Mg-X(etherate) Application: Structure: Et O OEt RMgX + 1/2 O2 → RO-MgX {+HX} → ROH + MgX2 (alcohols) 2 2 X RMgX + R'CHO → RR'CHOH (→ alcohol) R Mg Mg R → → RMgX + R'CN RR'CO (cyanide ) X → → Et O OEt RMgX + 1/8S8 RSMgX { + HX} RSH 2 2 { + RMgX} → R2S { + R'I} → RR'S

103 János Rohonczy: Inorganic Chemistry I. 8. Alkali Metals

8. (1st column) Li, Na, K, Rb, Cs, Fr – Alkali Metals

8.1. Elements

Valence shell: ns1. Unique behavior of alkali metals is their strong similarity in physical and chemical properties. Typical metals with strong electropositive behavior, EN<1. Very reactive elements. Low ionization potential, oxidation number is +1 in compounds. Low density. Two atomic molecules in vapor phase, specific flame colors. Increasing ion radius in top-down direction. Deviating element: Li – small ion radius, stronger polarizing behavior, some similarities to LiI and MgII chemistry. Table 46. Physical properties of alkaline metals Symbo MP BP Ion radius EN Ionization potential Density Flame color l (°C) (°C) (pm) (V) (g/cm3) Li 180.5 1347 76 0.98 5.39 0.534 Fire red Na 87.8 881.4 102 0.93 5.14 0.97 Yellow K 63.2 765.5 138 0.82 4.34 0.86 Violet Rb 39.0 688 152 0.82 4.18 1.53 Red violet Cs 28.5 705 167 0.79 3.89 1.90 Blue Fr 27 677 180 0.7 ? ? ?

8.1.1. Discovery Li Arfvedson(1817): new element. Name: lithos = stone (Greek). Preparation: Davy(1818), by electrolysis of molten Li2O. Na Davy (1807) by electrolysis of molten NaOH. Name: Na2CO3 = soda ash. K Davy (1807) by electrolysis of molten KOH. Name: pot ash = al-kalja (Arabic). Rb Bunsen/Kirchoff (1861) by spectroscopy. Name: rubidus = dark red (Latin). Cs Bunsen/Kirchoff (1861) by spectroscopy. Name: sky blue = caesius (Latin). Fr Perey (1939) Name: France. Product of Uranium decay.

8.1.2. Natural appearance

Li rare: 18 ppm, in ferromagnetic silicate minerals instead of Mg: LiAlSi2O6 spodumene Na/K 23000 ppm/18000 ppm. 7.-8. among the most abundant elements. Sea water, NaCl/KCl (salt/silvine), Na2CO3, NaNO3 (Chile saltpeter), Na2SO4, Na2B4O7, KCl⋅MgCl2⋅6 H2O. Rb 78 ppm, (similar to Ni, Cu, Zn). There is no own mineral. Byproduct of Li preparation. Produce: 5 t/year.

Cs 2,6 ppm (similar to Br, Hf, U). Mineral: Cs4Al4SiO26⋅H2O (pollucite). Prepared as byproduct of Li. Produce: 5 t/year. Fr radioactive, very rare: 235U → 227Ac → 223Fr → 223Ra.

8.1.3. Preparation

Li Origin: LiAlSi2O6 (spodumene) – 1-3% Li containing ore.

Technologic steps: 1) milling, flotation, calcination (at 1100 °C) – change of polymorph 2) + H2SO4 (250 °C) cooking, dilution, crystallization → Li2SO4⋅H2O János Rohonczy: Inorganic Chemistry I. 8. Alkali Metals

3) + Na2CO3 → Li2CO3 { + 2 HCl} → 2 LiCl + H2O + CO2 4) desiccation 5) electrolysis. MP = 900 °C, but MP of 55% LiCl / 45% KCl = 450 °C. Electrode potential depends on the temperature: over 500 °C K(Li), but not used to prepare potassium.. Below 500 °C Li(with K contamination). Produce: 1000 t/year. 6) Purification: Li(K) + H2 → LiH + KH → LiH + K + H2 (thermal diss. of KH).

Na MP of NaCl = 1100 °C, but MP of 40% NaCl / 60 % CaCl2 = 600 °C. Electrolysis: Na(Ca) on iron cathode, Cl2 on graphite anode. Purification: distillation, but industry grade Na contains Ca without any problem. Produce: 200,000 t/year.

K Electrolysis of molten KCl + K2CO3: Pure K, but potassium is well soluble in the salt. Volatile, oxygen sensitive forming superoxide. Special preparation only over T > 850 °C: KCl(molten) + Na(metal) → NaCl + K

Rb/Cs 2 Rb/CsCl + Ca → 2 Rb/Cs + CaCl2 (at T=750 °C)

8.1.4. Application

Li Al/Mg/Li alloy: aircrafts, cars, etc., Li-battery, Li compounds. Na Strong reducing agent. Former major application: PbEt4 (Na/Pb alloy). Ti/Zr (from halides), NaH and Na2O2. Na catalyst in rubber industry. Heat exchanger fluid in nuclear reactors: low MP, low viscosity, low n-absorbance, good thermal conductivity, high specific heat.

K Additive of Na alloy as heat exchanger fluid. KO2 (potassium superoxide) Preparation: K + O2 → KO2 (Application: 4 KO2 + 2 CO2 → 2 K2CO3 + 3 O2)

Physical properties

Table 47. Nuclear properties. Few stable isotopes. 6Li 7Li 23Na 39K (40K) 41K 85Rb 87Rb 133Cs 7% 93% 100% 93% 109 year 7% 72% 1010 year 100% I=1 I=3/2 I=3/2 I=3/2 I=3/2 I=5/2 I=7/2 There are few natural radioactive isotopes with long halt time. See over. Soft metals: K < Na < Li < Pb, low MP, silver white. Big atom size, low ionization energy, small EN values, easy excitation.

Solubility in liquid NH3: always blue (bronze shiny) color, metastable liquid. Strong reduction + – activity. Nonlinear electric conductivity. Contains M , M ions, solvated electrons, M2 molecules: complicated equilibrium. Chemical properties

Strong reduction activity in liquid NH3: – – M + NH3 → MNH2 + 1/2 H2 (NH3 + e → NH2 + 1/2 H2) – – EtOH + e → EtO + 1/2H2 (Na-ethylate) – – – – O2 + e → O2 Ge2H6 + 2 e → 2 GeH3 – 2– – – S8 + 2 e → S8 R–C≡CH + e → R–C≡C + 1/2 H2

105 János Rohonczy: Inorganic Chemistry I. 8. Alkali Metals

8.2. Alkali metal compounds

8.2.1. Alkali metal hydrides NaH / LiH Properties: ionic crystals, reducing agents, hydrogenation agents. Preparation: 2 Na + H2 → 2 NaH Application: 4 NaH + BF3 → NaBH4 + 3 NaF (preparation of complex metal hydrides) 4 NaH + AlBr3 → NaAlH4 + 3 NaBr special: NaH + CO2 → HCOONa (Na-formate) 2 NaH + 2 SO2 → Na2S2O4 + H2 (Na-dithionite)

8.2.2. Alkali metal halides, pseudo halides, oxohalides

Alkali metal halides Properties: ionic, colorless crystals, high MP, but MP/BP is decreasing in MF > MCl > MBr > MI order. LiCl is hygroscopic, NaCl is not (natural salt is hygroscopic because of MgCl2 contamination).

Preparation: MOH + HX → MX + H2O M2CO3 + 2 HX → 2 MX + CO2 + H2O

Occurrence, application: NaCl/KCl minerals. Better soluble LiCl/KCl. Note: NH4Cl is similar to KCl. Special application in soldering: NH4Cl → NH3 + HCl (complex formation and dissolution of CuO). NaBr/KBr in medicine: sedatives. KI: non hygroscopic reagent in iodometry. NaI is hygroscopic, not used for this purpose.

Alkali metal pseudo halides

NaCN/KCN Properties: strong toxicity, 2 KCN + CO2 + H2O → K2CO3 + 2 HCN. Laboratory preparation: K2CO3 + C + 2 NH3 → 2 KCN + 3 H2O Industrial preparation: K2CO3 + 4 C + N2 → 2 KCN + 3 CO NaNH2 + C(coke) → NaCN + H2 – Application: dissolving gold as [Au(CN)4] in water, after that easy reduction of gold.

Alkali metal oxohalides NaOCl / KOCl Only in water solution: strong oxidizers, disinfectants: liquid bleach. Preparation: electrolysis of NaCl (without diaphragm): 2 NaOH + Cl2 → NaOCl + NaCl + H2O

KClO3 (potassium chlorate) Non hygroscopic, pyrotechnic application. KBrO3 / KIO3 (potassium bromate/iodate) Analytical reagents. KClO4 (potassium perchlorate) Preparation: 2 KClO3 → KClO4 + KCl + O2

8.2.3. Alkali metal oxides Preparation: burning of metals: 2– 4 Li + O2 → 2 Li2O ( oxide, O ) 2– 2 Na + O2 → Na2O2 (sodium peroxide, O2 ) – K + O2 → KO2 (potassium superoxide, O2 , similar: Rb/Cs)

106 János Rohonczy: Inorganic Chemistry I. 8. Alkali Metals

By controlled oxidation: all other combination: NaNO2 + 3 Na → 2 Na2O + 1/2 N2 LiOH⋅H2O + H2O2 → LiOOH⋅H2O + H2O 2 LiOOH⋅H2O → Li2O2 + H2O2 + 2 H2O

M2O (alkali metal oxides)

Direct synthesis: 4 Li + O2 → 2 Li2O (only lithium) Table 48. The colors of the alkaline metal oxides.

Li2O Na2O K2O Rb2O Cs2O White White Pale yellow Yellow Orange

M2O2 (alkali metal peroxides) Na2O2 is pale yellow solid. Explodes with Al powder. Hydrolysis: Na2O2 + 2 H2O → 2 NaOH + H2O2 (preferred at acidic pH). Application: Na2O2 for bleaching (cellulose, textile). Na2O2 + CO → Na2CO3 (absorption of CO), Li2O2 + CO2 → Li2CO3 + 1/2 O2 (regeneration of O2 in air).

MO2 (alkali metal superoxides) e.g.: KO2. Properties: paramagnetic, darker colors in top-to-down order. MP ≈ 400 °C in all cases.

M2O3 (alkali metal sesquioxides) Properties: paramagnetic, brown. Preparation: 2 MO2 → M2O3 + 1/2 O2

MO3 (alkali metal ozonides). Other known: [Li(NH3)4]O3 Preparation: 3 MOH(cold solid) + 2 O3 (g) → 2 MO3 + MOH⋅H2O + 1/2 O2

MnOm (n>m) (alkali metal suboxides) Preparation: 12 Rb + O2 → 2 Rb6O → Rb9O2 + 3 Rb Dissociation: Rb9O2 → 2 Rb2O + 5 Rb. Other known: Cs7O ... Cs21O3, bronze color.

8.2.4. Alkali metal hydroxides

LiOH (lithium hydroxide). Preparation: LiOH(aq) → LiOH⋅H2O → LiOH + H2O. NaOH / KOH / RbOH Properties: hygroscopic white, solid pastilles, well soluble in water, form hydrates. Preparation: electrolysis of NaCl, KCl solutions (with diaphragms). Reaction: 2 NaOH + 2 HBr → 2 NaBr + H2O (with acid → salts) NaOH + EtOH ⇌ EtONa + H2O (Na ethylate, with alcohols → alkoxides) NaOH + CO2 → NaHCO3 (with acid anhydrides) NaOH + H2S → NaHS + H2O (with weak acids) Dissolves amphoteric oxides: BeO, ZnO, Al2O3, Sc2O3, SiO2, SnO2 Application: LiOH: CO2 absorption. NaOH and KOH are universal bases.

8.2.5. Other alkali metal salts and complexes

Na2S⋅9 H2O (sodium sulfide) Very hygroscopic (used in synthesis of rayon and cellophane).

(NH4)2S ( sulfide) Analytical reagent.

Na2SO3⋅7 H2O / NaHSO3 (sodium sulfite and ) Industrial reduction agents. Preparation: Na2CO3 + 2 SO2 + H2O → 2 NaHSO3 + CO2 (crystallization)

107 János Rohonczy: Inorganic Chemistry I. 8. Alkali Metals

2 NaHSO3 + Na2CO3 → 2 Na2SO3 + H2O + CO2 (crystallization)

Na2SO4⋅10 H2O (Glauber's salt) MP = 32,38 °C, temperature reference compound.

KHSO4 () Non volatile, Acidic disintegration agent: KHSO4 + oxide → soluble sulphates. Basic disintegration agents: Na2CO3 + K2CO3 / NaOH / KOH Basic oxidative disintegration agents: Na2O2/Na2CO3 or NaNO3/Na2CO3 or NaNO2

Na2S2O3⋅5 H2O (photographic fixer) Preparation: Na2SO3 + S → Na2S2O3 (in boiling solution) Reaction: 2 Na2S2O3 + I2 → 2 NaI + Na2S4O6 (iodometry) Na2S2O3 + 4 Cl2 + 5 H2O → Na2SO4 + 8 HCl + H2SO4 3– + – 2 Na2S2O3 + AgBr → [Ag(S2O3)2] + 4 Na + Br

NaNO2 (sodium nitrite) Hygroscopic. Application: diazotizing (in organic chemistry).

KNO3 (potassium nitrate) Non hygroscopic. Used in gunpowder.

NaNO3 (Chile saltpeter) Hygroscopic. Conversion process: NaNO3 + KCl → KNO3 + NaCl.

Na3PO4 (trisodium phosphate) Tertiary salt, TSP (water softening), Na2HPO4, secondary salt, NaH2PO4, primary salt.

NH4NaHPO4 → NaPO3 (sodium metaphosphate) "Phosphate glass" Transition metal oxides form color phosphate glasses: e.g. NaCoPO4 orthophosphate blue, Application: qualitative Chemical analysis.

Li2CO3 (lithium carbonate) Less soluble in water (similar to Li3PO4).

Na2CO3⋅10 H2O (soda ash) Important industrial compound. Application: other Na salts, glass industry, soap boiling. Leblanc process (1794): used minerals: NaCl + S + C (First chemical industrial product) 1) NaCl + cc. H2SO4 → Na2SO4 + 2 HCl (at high temperature) 2) Na2SO4 + C(coke) + CaCO3 → Na2CO3 + CaS + CO Solvay process (1865), till today 1) NaCl + NH3 + CO2 + H2O → NaHCO3 (brutto-equation) H2O + NH3 + CO2 → NH4HCO3 (gas absorption) + – NH4HCO3 + NaCl � NaHCO3 + NH4 + Cl (less soluble salt) 2) 2 NaHCO3 → Na2CO3 + H2O + CO2 (calcination)

K2CO3 ( potash) Preparation: 2 KCl + 3 MgCO3 + CO2 + H2O → 2 MgCO3⋅KHCO3 + MgCl2 2 MgCO3⋅KHCO3 → K2CO3 + 2 MgCO3 + H2O + CO2 Application: soap, glass industry, ceramics, strong base, K salts ...

KHCO3 (potassium hydrogen carbonate) Not hygroscopic analytical standard.

Na2SiO3 (water-glass) Application: flameless coatings, paper industry.

Na2B4O7⋅10 H2O (borax) Application: cosmetics, dissolution of metal oxides in welding, borax perl in qualitative analysis of metals.

108 János Rohonczy: Inorganic Chemistry I. 8. Alkali Metals

Alkali metal complexes

Li adducts are relatively stable. Stable complexes: crow ethers are soluble in hydrophobic solvents: e.g. benzo-15-crown-5 sodium: 170-220 pm holes can capture the Na+ ions:

O O

O O O

8.2.6. Alkali metal organic compounds

R-Li (alkyl lithium) Most important Li organic family (similar to R–MgX). Structure: less ionic alkaline metal organic compounds, covalent bond, increasing ionic character in Li < Na < K < Rb < Cs order. Properties: thermally instable: LiCnH2n+1 → LiH + CnH2n (β-elimination, products). Most stable: Li–Me (methyl lithium) white, solid, covalent. nBu-Li, tBu-Li – tetramer or hexamer units. Preparation: R–X + 2 Li → R–Li + LiX nBu-Li + Ar-I → Ar–Li + nBu-I Sn(CH=CH2)4 + 4 Ph–Li → 4 Li–CH=CH2 + SnPh4 HgAr2 + 2 Li → 2 R–Li + Hg More ionic compounds: 2 Na + 2 C5H6 → 2 NaC5H5 + H2 2 Cs + CH2=CH2 → 2 Cs–CH2–CH2–Cs 2 Li + 2 HC≡CH → 2 Li–C≡C–H + H2 (ionic product: Li-acetylide) Mechanisms: through R– carbanion or radicals. Reaction: LiR + X2 → LiX + RX LiR + H+ → Li+ + RH LiR + R'-I → LiI + R-R' LiAr + CO2 → ArCOOLi {+H2O} → ArCOOH + LiOH LiR + R'CONMe2 → LiNMe2 + R-CR'O 3 LiR + BCl3 → BR3 + 3 LiCl Application: Polymerization catalysts, alkylation reagents instead of Grignard's reagent. Suitable for syntheses of Si/Ge/Sn/P/As/Sb/Bi organic compounds. Special application: intermediate in synthesis of vitamin A. 2 Li + HC≡CH → Li2C2 (Li-dicarbide)

109 János Rohonczy: Inorganic Chemistry I. 9. Noble Gases

9. (18th column) He, Ne, Ar, Kr, Xe, Rn – Noble Gases

9.1. Elementary properties Valence shell: He – 1s2, others ns2np6 Table 49. Selected physical data of the noble gases. Symbol Name MP (K) BP (K) 1st Ioniz. Energy (kJ/mol) He At high pressure 4,21 2372 Ne Neon 24,5 27,1 2080 Ar 83,8 87,3 1520 Kr Krypton 115,9 119,7 1351 Xe Xenon 161,3 165,0 1170 Rn Radon 202,0 211 1037

History: 1785 Cavendish: 1/120 part of the air is not O2, nor N2, but it is inert.

He Lockyer, Frankland (1868): unknown line close to the NaD in the spectrum of the Sun. (Name: helios = Sun (Greek)}. Palmieri (1881): same spectrum line in the gas of Vezuv. Ramsey, Rayleigh (1895): He in the air, in the gas of heated unanimity (pitchblende).

Ar MN2(air) > MN2(from NH3): N2(air) + 3 Mg → Mg3N2 + ?: heavy, monatomic gas. There was no place for it in the periodic system. Kr (1898) detected in liq. air. Name = hidden (Greek). Ne (1898) detected in liq. air. Name = new (Greek). Xe (1898) detected in liq. air. Name = strange (Greek). Rn (1902) Rutherford, Soddy: radium emanation/niton. Isotopes: radon/toron/actinon. Defined place in the periodic system: between alkaline metals an halogens. Common name: Inert gas / . Inertness: Lewis , Kossel (1916): octet of 8 electrons, like atoms in stable molecules. Occurrence: He is 2nd abundant in the Universe. On the Earth: 4 ppm in the air, escapes into space. Origin: 4He: α-radiation + 2 e–

Ar 1% of air. Pollutant of N2/H2 in industrial produce of NH3. Ne/Ar/Kr/Xe He: 5 ppm, Ne: 18 ppm, Ar: 94 ppm, Kr: 1 ppm, Xe: 0.09 ppm, Rn: variable in air. Produce: Condensation and fractional distillation of liquid air. Application: Ar Inert atmosphere in high temperature metallurgy. Filling gas of electric bulb. Produce: 700,000 t/year. Ne, Kr Discharge tubes: neon tube (color depends on the gas mixture). He Ca. 0,4% in earth gas (max. 7%), ca. 5,000 t/year. Application: balloons, inert atmosphere. Diluting gas instead of N2 (scuba diver). Carrier gas in GC, coolant gas in nuclear reactors. Liquid He: coolant liquid for superconductors. Rn Former radioactive isotope in cancer therapy, inspection of welded joints. Today other cheaper isotopes are used for this purpose. Preparation: 1g 226Ra → 0,64 cm3 222Rn, during 30 days. János Rohonczy: Inorganic Chemistry I. 9. Noble Gases

Table 50. Nuclear properties, number of isotopes of the noble gases. He Ne Ar Kr Xe Rn 2 3 3 6 9 (1)

Physical properties

Colorless, odorless, tasteless, monatomic, inert gases. Increasing atom size ⇒ increasing polarizability. Big, but decreasing ionization energy. Crystal lattice: ideal molecular lattice with weak intermolecular van der Waals forces, that is why the MPs and BPs are very low. Fast diffusion through plastics, He: through glass as well. He Unusual property: He-I → He-II (λ point at 2,2K): He-I is ordinary liquid, but He-II is super fluid: there is no boiling, no turbulences. Few hundred atom thick liquid film layer on the surface. Heat capacity 10 times, thermal conductivity 106 times bigger, viscosity = 0. (Reason: quantum liquid).

Chemical Properties

Chemically inert gases, but there are several known compounds. Pauling's hypothesis (1933): KrF6 and XeF6 may exist. 1962: first Xe compound: XePtF6. Till that only the noble gas chlatrates were known: 1) hydroquinone crystals (HO–C6H4–OH) at 10-40 bar: Ar, Kr, Xe (or O2, N2, CO, CO2). 2) hydrates with [M8(H2O)46] ideal composition: M = Ar, Kr, Xe, but M ≠ He, Ne – too small. 3) gas inclusion in zeolite crystals. Application: capture of synthetic radioactive isotopes.

New chapter of noble gas chemistry started in 1962. Bartlett: first ionization energy of O2 and ο + – Xe practically are same, and PtF6 can oxidize the oxygen: PtF6 + O2 → (25 C) O2 [PtF6] (the color of PtF6 changed on air at RT). Xe can react wit PtF6 as well.

Table 51. Ionization energy of O2 and noble gases. + + + + + 1. ioniz. O2 → O2 Xe → Xe Rn → Rn Kr → Kr Ar → Ar kJ/mol 1175 1170 1037 1351 1520

Xe forms similar compound: PtF6 + Xe → Xe[PtF6], dark red → orange. First Xe compound. Xe chemistry is very rich. Kr chemistry describes less stable compounds. Radon has also compounds, but these are very dangerous. Auto-radiolysis. He, Ne, Ar compounds are instable intermediates.

9.2. Xenon compounds

9.2.1. Xe fluorides, reactions

XeF2 (xenon difluoride) Structure: linear F–Xe–F molecule. Properties: white crystals, sublimates, water soluble. Preparation: Xe + F2 → XeF2 (400 °C, Ni reactor) or (hν radiation = sunlight) Reactivity: Slow hydrolysis: 2 XeF2 + 2 H2O → 2 Xe + 4 HF + O2 Water solution is strong oxidizer agent: – – 3+ 2– + 2+ – – XeF2 + 2 Cl → Xe + 2 F + Cl2, Cr → CrO4 , Ag → Ag , BrO3 → BrO4

111 János Rohonczy: Inorganic Chemistry I. 9. Noble Gases

XeF4 (xenon tetrafluoride) Properties: white solid powder, sublimates. Structure: AX4E2 tetragonal planar shape, VSEPR. Preparation: Xe : F2 = 1 : 5 mixture → XeF4 (400 °C, Ni reactor, 6 bar) Reaction: like XeF2, but stronger fluorination agent: XeF4 + 2 Hg → 2 HgF2 + Xe XeF4 + Pt → PtF4 + Xe Hydrolysis: 3 XeF4 + 6 H2O → XeO3 + 2 Xe + 3/2 O2 + 12 HF (explosive XeO3)

XeF6 () Properties: colorless crystal, yellow liquid, yellow gas. + Structure: AX6E distorted octahedron in gas phase. XeF5 is square pyramidal: AX5E. + – + + – In liquid phase: XeF5 ···F ···XeF5 (2, 4, 6 XeF5 units around F anions). Preparation: Xe : F2 = 1 : 20 mixture → XeF6 (300 °C, 50-60 bar) Reactivity: Most reactive, strong oxidizer and fluorination agent. Violent hydrolysis, like XeF4. Violent reaction with SiO2: 2 XeF6 + SiO2 → 2 XeOF4 + SiF4 2 XeOF4 + SiO2 → 2 XeO2F2 + SiF4 2 XeO2F2 + SiO2 → 2 XeO3 + SiF4

XeF Known, instable radical: formed from XeF4 by γ radiation.

Other halides. XeF8 is not confirmed compound. XeCl2, XeCl4, XeBr2 are instable compounds, they are observed in trace quantity only.

Reactions of Xe fluorides With F– acceptors: + – XeF2 a) XeF2 + MF5 → [XeF] [MF6] + – 2 XeF2 + MF5 → [Xe2F3] [MF6] + – XeF2 + 2 MF5 → [XeF] [M2F11] (M= P, As, Sb, I) + – b) XeF2 + MOF4 → [XeF] [MOF5] (M= Mo, W) + – c) Xe + PtF6 → [XeF] [Pt2F11] (Bartlett) + – XeF4 d) XeF4 + 2 SbF5 → [XeF3] [Sb2F11] (M = Sb, Bi) + – XeF6 e) XeF6 + MF5 → [XeF5] [MF6] (M = Sb, Pt) With F– donors: + – XeF6 + CsF → Cs [XeF7] (M = Cs, Rb) 2– 2 Rb[XeF7] → XeF6 + Rb2[XeF8] (M = Na, K, Rb, Cs) + + Na2 [XeF8] → 2 NaF + XeF6 (Na : Tdiss. =100 °C, Cs : T diss. = 400 °C) (XeF6 reacts only. Used to separate from XeF4 and XeF2) Hydrolysis:

XeOF4 Structure: AX5E. Properties: colorless, volatile liquid. Preparation: XeF6 + H2O → XeOF4 + 2 HF (controlled, otherwise till XeO3) 2 XeF6 + SiO2 → 2 XeOF4 + SiF4

XeO2F2 Structure: AX4E. Properties: low MP, colorless solid. Preparation: XeO3 + XeOF4 → 2 XeO2F2

XeO3 Structure: AX3E. Properties: colorless, explosive crystals. Preparation: XeF6(vapor) + 3 H2O → XeO3 + 6 HF Purification: 2 HF + MgO → MgF2 + H2O (vacuum distillation)

112 János Rohonczy: Inorganic Chemistry I. 9. Noble Gases

– HXeO4 (hydrogen tetraoxo xenate) Properties: instable, disproportion. – – 3 Preparation: XeO3 + OH � HXeO4 (oxoacid anion, Kd = 1,5⋅10 ) 4– 2– XeO6 / H2XeO6 ( / dihydrogen perxenates) Colorless solids till 200 °C. 4– + + + 2+ 3+ Structure: XeO6 octahedral anion, AX6 geometry. Stable salts: Na , K , Li , Ba , Am . – – 4– Preparation: 2 HXeO4 + 2 OH → XeO6 + Xe + O2 + 2 H2O – 4– – 2 XeF6 + 16 OH → XeO6 + Xe + O2 + 12 F + 8 H2O Reaction: strong oxidizer agent: 2– 2+ – – + 5 H2XeO6 + 2 Mn (aq) → 5 HXeO4 + 2 MnO4 +H2O + H3O

XeO4 () Properties: explosive gas. Condensates at liquid N2 temperature: still explosive. MP = -35,9 °C. Preparation: Ba2H2XeO6 + 2 H2SO4 → 2 BaSO4 + XeO4 + 2 H2O

Other compounds: by ligand exchange between XeF2 and concentrated acids: XeF2 + HL → F–Xe–L + HF (HL = FSO2OH, HClO4, HOTeF5, CF3COOH, HN(SO2F)2)

9.2.2. Xe organic compounds, Xe-Si/N bonded compounds

Xe(CF3)2 (bis trifluormethyl xenon) Preparation: 6 XeF2 + C2H6 → Xe(CF3) 2 + 6 HF + 5 Xe + [C6F5Xe][AsF6], [XeC6F9][AsF6], [XeC6F11][AsF6], [XeCCR] + + + [F3SiXe] Preparation: [F3SiFH] + Xe → [F3SiXe] (in Mass Spectrometer)

[FXeN(SO2F)2] Preparation: HN(SO2F)2 + XeO2 → [FXeN(SO2F)2]

9.2.4. Chemistry of Krypton and Radon Less studied compounds, but they are similar to the Xe compounds.

Krypton compounds

KrF2 (krypton difluoride) Properties: volatile colorless crystal. Preparation: crystalline product of reaction of Kr + F2 mixture at –196 °C. KrF2 is less soluble in liquid F2. Reactivity: Slow thermal dissociation at room temperature. Extremely strong oxidation and fluorination agent: e.g. RuO4 + F2 → RuOF4 + O2 + + + Known ions: [KrF] , [Kr2F3] , [CF3CNKrF] .

Rn compounds All Radon isotopes have short lifetime. Very special/difficult synthesis methods were used, + - + like isotope tracing to identify Rn compounds: e.g. RnF2, RnO3, [RnF] [TaF6] . Rn can replace Na+ or K+ ions in less soluble alkali salts.

113 János Rohonczy: Inorganic Chemistry I. References

References

1 A. Earnshaw, N.N. Greenwood: Chemistry of the Elements (ChE) Elsevier Ltd., (Second Edition) 1997. ISBN: 978-0-7506-3365-9 http://www.sciencedirect.com/science/book/9780750633659 1H ChE, Chapter 3, Hydrogen 1X ChE, Chapter 17, The Halogens 1O ChE, Chapter 14, Oxygen 1S ChE, Chapter 15, Sulfur 1Se ChE, Chapter 16, Se, Te, Po 1N ChE, Chapter 11, Nitrogen 1P ChE, Chapter 12, Phosphorus 1As ChE, Chapter 13, As, Sb, Bi 1C ChE, Chapter 8, Carbon 1Si ChE, Chapter 9, Silicon 1Ge ChE, Chapter 10, Ge, Sn, Pb 1B ChE, Chapter 6, Boron 1Al ChE, Chapter 7, Al, Ga, In, Tl 1Be ChE, Chapter 5, Be, Mg, Ca, Sr, Ba, Ra 1Li ChE, Chapter 4, Li, Na, K, Rb, Cs, Fr 1He ChE, Chapter 18, Noble Gases

2 Wikipedia. https://www.wikipedia.org/wiki/ 2H H https://en.wikipedia.org/wiki/Hydrogen

2X Halogens https://en.wikipedia.org/wiki/Halogen 2F F https://en.wikipedia.org/wiki/Fluorine 2Cl Cl https://en.wikipedia.org/wiki/Chlorine 2Br Br https://en.wikipedia.org/wiki/Bromine 2I I https://en.wikipedia.org/wiki/Iodine 2At At https://en.wikipedia.org/wiki/Astatine

2KA Oxygen family https://en.wikipedia.org/wiki/Chalcogen 2O O https://en.wikipedia.org/wiki/Oxygen 2S S https://en.wikipedia.org/wiki/Sulfur 2Se Se https://en.wikipedia.org/wiki/Selenium 2Te Te https://en.wikipedia.org/wiki/Tellurium 2Po Po https://en.wikipedia.org/wiki/Polonium

2PN Nitrogen family https://en.wikipedia.org/wiki/Pnictogen 2N N https://en.wikipedia.org/wiki/Nitrogen 2P P https://en.wikipedia.org/wiki/Phosphorus 2As As https://en.wikipedia.org/wiki/Arsenic 2Sb Sb https://en.wikipedia.org/wiki/Antimony 2Bi Bi https://en.wikipedia.org/wiki/Bismuth

2CG Carbon family https://en.wikipedia.org/wiki/Carbon_group 2C C https://en.wikipedia.org/wiki/Carbon 2Si Si https://en.wikipedia.org/wiki/Silicon 2Ge Ge https://en.wikipedia.org/wiki/Germanium 2Sn Sn https://en.wikipedia.org/wiki/Tin 2Pb Pb https://en.wikipedia.org/wiki/Lead

2BG Boron family https://en.wikipedia.org/wiki/Boron_group 2B Boron https://en.wikipedia.org/wiki/Boron 2Al Aluminium https://en.wikipedia.org/wiki/Aluminium 2Ga Gallium https://en.wikipedia.org/wiki/Gallium János Rohonczy: Inorganic Chemistry I. References

2In Indium https://en.wikipedia.org/wiki/Indium 2Tl Thallium https://en.wikipedia.org/wiki/Thallium

2AE Alkaine earth metal https://en.wikipedia.org/wiki/Alkaline_earth_metal 2Be Beryllium https://en.wikipedia.org/wiki/Beryllium 2Mg Magnesium https://en.wikipedia.org/wiki/Magnesium 2Ca Calcium https://en.wikipedia.org/wiki/Calcium 2Sr Strontium https://en.wikipedia.org/wiki/Strontium 2Ba Barium https://en.wikipedia.org/wiki/Barium 2Ra Radium https://en.wikipedia.org/wiki/Radium

2AM Alkaine metal https://en.wikipedia.org/wiki/Alkaline_metal 2Li Lithium https://en.wikipedia.org/wiki/Lithium 2Na Sodium https://en.wikipedia.org/wiki/Sodium 2K Potassium https://en.wikipedia.org/wiki/Potassium 2Rb https://en.wikipedia.org/wiki/Rubidium 2Cs https://en.wikipedia.org/wiki/Caesium 2Fr Francium https://en.wikipedia.org/wiki/Francium

2NG Noble gases https://en.wikipedia.org/wiki/Noble_gas 2He Helium https://en.wikipedia.org/wiki/Helium 2Ne Neon https://en.wikipedia.org/wiki/Neon 2Ar Argon https://en.wikipedia.org/wiki/Argon 2Kr Krypton https://en.wikipedia.org/wiki/Krypton 2Xe Xenon https://en.wikipedia.org/wiki/Xenon 2Rn Radon https://en.wikipedia.org/wiki/Radon

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