HYDROGEN Occurence: 89 % Universe 0,88 % Earth (tj. 15.4 at. %), Earth's crust 0.15 % Isotope H D T 99.844 % 0.0156 % Relative atomic mass 1.007825 2.014102 3.016049 Nuclear stability stable stable T1/2=12.35 let Melting point / °C -259.193 -254.65 -252.53 Boiling point / °C -252.76 -249.48 -248.11 Dissotiation heat / kJ/mol 435.88 443.35 446.9 Comment: 2H = D ; 3H = T Deuterium (D) and tritium (T) Deuterium can be obtained by electrolysis of water Nuclear reactions leading to the formation of tritium H (2H , p) H 14N (n , 3H) 12C 6Li (n , a) H Used for tritium production Storing of gaseous tritium: in form UT, The thermal decompositin at 400 °C yields gaseous tritium 2 UT3 -> 2 U + 3 T2 Isotopic effect • observed at compounds where some atom is replaced by another isotopic atom having other atomic mass • the change af atomic mass has influence on physical properties of the compound. This effect is relatively strongiest just for hydrogen isotopes - replacing H-atom by D-atom with double mass. Mean kinetic energy of gas molecules heavier molecules move slowly Speed of chemical reactions reaction with heavier isotopes are going on with other speed Vibration of chemical bond wavelengths of vibrations are changed Melting point „light water 0 °C, „heavy water" 3.82 °C Diffusion speed Graham law , different separation speed of 235UF6 , 238UF6 Compound marked with deuterium or tritium „ISOtope marking" (specific or non-specific) of compounds with heavier hydrogen isotopes leads to the formation of compounds that enables to follow behavior of this marked atom in reactions or other processes with the aim to investigate the mechanism of the process. Marking can be often carried out by simple contact of the intended compound with the compound that can yield (e.g. due to dissociation) free heavier hydrogen particle - isotopic exchange Non-specific marking CH3OH + D2O -> CH3OD + HDO Specific marking (hydrogen isotope is placed on the required position in the molecule => it needs a sophisticated synthetic approach) Nuclear isomers of hydrogen isotopes Comment: conversion ortho -— para is slightly exotermic => problems with storing of liquid hydrogen Hydrogen - laboratory preparation MS 1 211(1 M^'li 1 "in statu nascendi" Zn + 2NaOH + 2H20---► Na2[Zn(OH)4] + H2 Na + H20---► NaOH + Y2H^ 2II30+ + 2e"---► 2H20 + H 40H'---1 211,0 + O, + 4c cathode anode CaH2 + 2H20---► Ca(OH)2 + 2H- 300 °C 2UIIj----► 2U + 3H „Water gas" ^ 1 000 °c Preparatory stage ( [ I ,<> ■(.'() I 1 ^O",,!!,) Conversion CO + ii2o 500 °C , CO, i 11, l;0,O, (r:(). Cll4 t n2o 900 °C Ni/AljO, CO + 311 1200°C CI I. . C I 211, 2Na(IIg)x f 2II20-- (decomposition of sodium amalgam in the NaOH production • 2NaOH +• H, l 2x1 Ig Hydrogen - production Chemical (non-electrolytic) decomposition of water 2H20 -2H2 + 02 f 1- CaBr2 + H20 750'C CaO + 2 HBr 2. Hg + 2 HBr 100C t HgBr, + H2 1st sequence I 25'C -» 3. HgBr2 + CaO HgO + CaBr, 4. HgO 500" C Hg + 1/2 02 2nd sequence 1 3 FeCI2 + 4 H20 500'C ♦ Fe304 + 6 HCI + H 2. Fe304 + 3/2CI+6 HCI 1 ■ 3 FeCI3 + 3 H20 + 1/2 02 3. 3 FeCU 300 C ■ 3 FeCI + 3/2 CI Hydrogen - utilization in: Food industry Metallurgy | Food industry Fertilizers, plastics Missile fuel elements Hydrogen - bonding Relations among bonds Ionic bond • e Hydrogen - bonding Bonging possibilities of hydrogen a) Formation of molecular particles: H2 ; ; H2 + b) Formation of atomic particles: - very small: 1,5.10-3 pm, (for comparison common radius of atoms are 50 - 220 pm) H+ - dissociation of acids, very reactive, searching for a stabilizing partner HA — H+ + A- H+ + HO — H 0+ 2 3 H arises only in the process of ionic hydrides melting, e.g. NaH Hydrogen c) Formation of hydrogen bonds: bonding energy 10 - 60 kJ mol-1 molecule 1 molecule 2 molecule 42 fl n <2L Hydrogen bond intermolecular Hydrogen bond intramolecular X-H- Y Hydrogen - bonding 100 II 100 Boiling points of some binary hydrogen compounds as a result of the existence of hydrogen bonding ■200 n Hydrogen - reactivity a) Reduction properties (typical) PdCl2 (aq) + H2 (g)---► Pd (s) + 2HC1 (aq) ' I H2 I X2---► 2HX I 2H2 I 02----► 2H20 I 2H2 + CO---> CH3OH I b) Oxidation properties (only in the case of ionic hydrides formation) 2Na + H —> 2 NaH Hydrogen - hydrides Groups of pse Ionic hydrides 1, 2, 3 + LnH2 (in LnmH2 e-) Transition hydrides 4, 5 Metal hydrides Cr, Ni, Pd, ... (solid solutions) Covalent molecular hydrides H2O, NH3, ... Covalent polymeric hydrides Be, Mg, 12th and 13th groups Ionic hydrides (Na, Ca, etc.) have strong reduction properties, e.g. in reaction with water free hydrogen is released. H- + H2O -- OH- + H2 Alkali metals, ns1 Lithium, sodium, potassium, rubidium, caesium, francium • Alkali metals are silvery-grey metals, fresh cut are of glossy appearance, only caesium is coloured to gold-yellow tone. • It is necessary to keep these metals under paraffinic liquids or in inert atmosphere (Rb, Cs). • All elements of this group are very electropositive, caesium is most positive common element at all (except radioactice and very rare Fr). Alkali metals - properties Element Li Na K Rb Cs Fr Atomic number 3 19 37 55 87 Density / g cm-3 0,534 0,968 0,856 1,532 1,90 ■ Melting point / °C 180,5 97,8 63,2 39,0 28,5 27 Boiling point / °C 1347 881,4 765,5 688 705 667 Metallic radium / pm 152 186 227 248 265 ? ■ Ionic radius (for number = 6) / pm 76 102 138 152 167 180 1st ionization energy / eV 5,390 5,138 4,339 4,176 3,893 4,0 2nd ionization energy / eV 75,62 47,29 31,81 27,36 23,4 ? ■ Electronegativity (Allred-Rochow) 0,97 1,01 0,91 0,89 0,86 0,86 Alkali metals - Minerals I j lepidolit K2Li3Al4Si7021(OH,F)3 Lspod u men LiAlSi20Ŕ halit NaCl Mil trona Na2C03.NaHC03.2H20 \l/ kryolit Na,AlFť Chile nitrate N:iNG3 ?^ \ karnalit KCLMgCl2.6H20 kainit KCl.MgS04.3H20 sylvite KCl R b lepidolit CS pollucit Cs4Al4Si90,fi.H,0 (lake Bernic, Manitoba) Fr HALITE NaCI SYLVITE KCl Alkali metals - reactivity Chemistry is relatively simple and is related to the easy formation of oxidation state + I, exceptionally also -I > formation of ionic compounds > Li - more covalent character of bonds > Li chemistry is similar to Mg chemistry (diagonal similarity) Li+76pm Mg2+ 72 pm Na+ 102 pm > formation of complexes is not typical, most known are complexes with macrocyclic ligands (e.g. crowns) > existence is possible only in complexes of a macrobicyclic type - cryptates, e.g. [Na(crypt)]+Na- Alkali metals - typical reaction • in most cases direct reactions • on air - metals react with oxygen - formation of a layer containing oxidation products - oxides, peroxides, hyperoxides, hydroxides, carbonates Reduction action of alkali metals Reaction with water 2M + 2H20-----► 2MOH + H2 Reaction of non-metallic halogenides SiF4 + 4K-----► Si + 4KF Alkali metals - Lithium Lithium: electrolysis in melt LiCl (55 %) a KCl (45 %) at 450 °C LiCl production 1. annealed spodumene is extracted by conc. sulfuric acid => lithium sulfate is converted to lithium chloride Li2SO4 + Na2CO3 — Li2CO3l + Na2SO4 Li2CO3 + HCl — 2 LiCl + H2O + CO2 2. spodumene is annealed together with limestone, products is leached using water => 2 LiOH + 2 HCl — 2LiCl + H2O Li - metal lowest density Alloys Li + Mg + Al (e.g. so called. LA141) of the composition 14 % Li, 1 % Al a 85 % Mg) is utilized as a construction material in missile technology Alkali metals - Sodium Production: Electrolysis of melt NaCl / CaCl2 (4:6) at 580 °C, (NaCl melts at 808 °C) e Alkali metals - production Potassium: > electrolysis KCl melt > reduction of KCl melt using sodium Caesium: reduction caesium dichromate melt by Zr 500 C. vacuum Cs2Cr207 + 2Zr ------------> 2Cs + 2Zr02 + Cr203 Alkali metals - compounds with oxygen Alkali metals form: oxides (O2-), peroxides (O2 2-), hyperoxides (O2-), suboxides, respectively. Na2O production Na2O2 and NaO2 2 Na + O2 Na2O2 Ozonides Ozonide reactions Na2O2 + O2 2 NaO2 6 KOH + 4 O3 -> 4 KO3 + KOH.H2O + O2 Alkali metals - compounds with oxygen Alkali metals - compounds with oxygen Practical applications of oxygen-containing sodium compounds Reaction of sodium peroxide with water yields hydrogen peroxide |Na202 + 2H20 ---- t 2NaOH + H202 Reaction of sodium peroxide with CO2 yields alkali metal carbonates Na202 + C02-----v Na2CQ3 + ^02 Na202 + 2K02 + 2C02-----► Na2C03 + K2CQ3 + 202 Reaction of sodium peroxide with CO and CO2 are exploited in breathing devices (firemen, submarines, space ships): 2 Na2O2 + 2 CO2 2 Na2CO3 + O2 Na2O2 + CO Na2CO3 Alkali metals - compounds with sulfur Li Na K, Rb, Cs Na2S KHS M2SX x= 2 5 6 Alkali metals sulfides are : Soluble in water ♦> Crystallize with many water molecules ♦> Hydrolysis is observed in aqueous solutions Alkali metals - metal hydrides Direct synthesis (LiH is most stable) Reduction effects LiH + H2O -- H2 + LiOH Reaction of LiH is very vigorous, NaH and other hydrides react even explosively. 4 NaH + TiCl 400 C > Ti + 4 NaCl + H2 Sodium formiate production NaH + CO2 — Na(HCOO) Complex hydrides (used in organic synthesis) 4 LiH + BF3 — Li[BH4] + 3 LiF 4 NaH + AlBr3 — Na[AlH4] + 3 NaBr Alkali metals - carbides and organometallic compounds Acetylides M + C2H2-----> for Li also LiHC2 Alkyl- and aryl lithium (used for alkylation or arylation reactions) 2 Li + Cl-CH2CH2CH2CH3 -> Li-CH2CH2CH2CH3 + LiCl Li- CH2CH2CH2CH3 + Ar-I -> Ar-Li + I- CH2CH2CH2CH3 • reaction solvents petrolether, cyclohexane, benzene, diethyl ether • alkyl- and aryl lithium derivatives are very sensitive towards water, air-moisture, oxygen, and CO2 => manipulation only in inert atmoshere Alkali metals - reaction with nitrogen a N-compounds Other reactions of alkali metals: > Lithium + N2 -> Li3N a Li2NH (direct reaction) > Alkali metals in liquid ammonia => intense blue solutions with reduction properties K2[Ni(CN)4] + 2 K NH3("33°C) > K4[Ni(CN)4] => presence of solvated electrons Q Na+ e- (NH3)2-3 ) => solutions are not too stable and amides are formed 2 M + 2 NH3 -> MNH2 + H2 Remark: Similar solutions are formed also in the process of dissolving alkali metals in amines, polyethers, etc. Alkali metals - salts Common properties of alkali metal compounds: • cations are colourless • chemical properties of salts are given by the character of the central atom in anionic part of the compound (e.g. colour) • salts are formed most frequently by neutralization • salts are mostly well-soluble in water (strong electrolytes) • salt of weak acids are partially hydrolyzed • analytically significant are little soluble sodium salts: Na[Sb(OH)6] and NaZn(UO2)3(CH3COO)96H2O • K+, Rb+, Cs+ - can be precipitated as perchlorates, hexanitrocobaltates (III), tetraphenylborates, and hexachloroplatinates (IV) Alkali metals - salts Nitrates ?"» < 2NaN03______, 2NaN02 + 02 2NaN03 -----^ Na20 + N2 + 5/202 NaN03 + Pb-----► NaN02 + PbO Na2C03 + NO + N02----> 2NaN02 + CO Halogenides NaCl, KCl, CsCl; NaBr, KBr, CsBr; NaI, KI, CsI Alkali metals - Technically important alkali metal compounds Sodium hydroxide a) NaOH production - soda caustification (not used) Na2C03 + Ca(OH)2-----> CaC03l + 2NaOH b) NaOH production by electrolysis of brine (70% NaCl in water) Diaphragma method Amalgam method Alkali metals - Technically important alkali metal compounds Na2CO3 (Le Blanc way) Na2S04 + 2C-----> Na2S + 2C02 Na2S + CaC03-----> Na2C03 + CaS Na2CO3 (Solvay way from brine) NH3 + H20 + C02_____> NH4HC03 NaCl + NH4HC03_____► NaHC03 I + NH4CI NaHCO3 thermically decomposes (calcination) to Na2CO3 Remark: ammonium chloride reacts with Ca(OH)2 and relased NH3 is used in soda production. The only real wastecominh from this process product is CaC^ . Alkali metals - Technically important alkali metal compounds K2CO3 (Engel way) MgC03.3H20 + KCl + H20 + C02 2MgC03.KHC03.4H20-----► K2C03 + 2MgC03.3H,0 + CO, + H20 Alkali metals - complexes Crown-ethers Alkali metals - complexes B) LtMU natride 2Na + N{ (NH4),BcF4 + H2 (g) Be alloys Beryllium bronze Be/Cu Properties of Be > m.p. 1300 °C > Be chemistry is similar to Al Chemistry- diagonal similarity > reaction with water is slow, the surface of a metal is covered with a layer of bad soluble hydroxide > Be can be dissolved in acids (H2 is formed) > Be2+ does not exist in aqueous solutions, only in the form hydrated ions > Be in conc. HNO3 is passivated > Be is amphoteric - dissolves also in alkali hydroxides > Be form tetrahedral complexes - SP3 ( Td) > soluble Be compounds are toxic!! Be compounds Simple Be compounds: BeC2 Be.C 2 m.p. 2750 °C, Mohs hardness scale 9 Be halogenides Be(OH)2 + 2HF BeF2 + 2F- BeF2. 4H2O [BeF4]2- Other halogenides are synthetized by direct reaction or with dry HHal Be compounds Beryllium hydride (cannot be highly polymeric pared by direct synthesis from elements) Preparation: BeCl2 + 2 LiH BeH2 + 2 LiCI Hydrolysis: BeH2 + H2O Be(OH)2 + H2 Solvolysis: BeH2 + CH3OH —► Be(OCH3)2 + H2 Complex compounds of Be Be compounds slowly undergo to hydrolysis in aqueous solution [Be(OH)]n 11+ Complex fluorides: Other complexes: Bc40(N03), H3C 3 \ CH3 HC Q. O-CH Be. HC-Q Q CH H3C CH 3 Beryllium acetylacetonate Organometallic compounds of Be Bc(CH3)2 Bc(Bu)2 (direct bond Be - C) [Fe"(115-C5H5)2] (^F^ Ferrocene [Be(n l-C5H^Xi15-C5H5)] „Sandwich" complex type Use of Be and its compounds > Enter windows of X-ray and Geiger-Müller tubes (low absorption of radiation) > Beryllium bronze > Tritium production 94Be + 2iH -----------> 2 42He + 3iH > Neutron 241Am / Be source Magnesium Mg occurance: 2.76% In sea water v. dolomit CaC03.MgC03 magnezit MgC03 brucit Mg(OH)2 kainil KCl.MgS04.3H,0 peri k las VI gO epsomil MgS04.7H20 epsomit karnaiit oliviine I talc asbestos MgSO4.7H2O KCl.MgCl2 (MgFe)2SiO4 Mg3Si4O10(OH)2 Mg3Si2O5(OH)4 spinei semi-presious stone) MgAl2O4 Mg production: 300 000 tons / year 2 MgO.CaO + FeSi ferrosilicon 2 Mg + CaSiO4 + Fe Electrolysis of molten MgCl2 Properties of Mg > reaction with water is slow, the surface of a metal is covered with a layer of bad soluble hydroxide > Mg can be dissolved in acids (H2 is formed) > Mg2+ exists only in the form hydrated ions in aqueous solutions > non- solubility in alkali metal hydroxides - Mg is not amfoteric > Mg burns even in water vapour => cannot be used for fire extinguishing Magnesium hydride Mg + H2 MgH2 Direct synthesis at 20 Mpa, catalysis using MgI2 Reaction with water and alcohols similar as Be: MgH2 + H2O Mg(OH)2 + H2 MgH2 + CH3OH Mg(OCH3)2 + H2 Other binary magnesium compounds Carbides: MgC Nitride: Mg3N. Halogenides: anhydrous are less stable as Be analogues bad soluble is F- 2 MgCl2 + H2O Mg2OCl2 + 2 HCl (termal decomposition) ^ principle of Sorell cement hardening: Mg(OH)2 + conc. MgCl2 solution Hydroxide: Non-amphoter Sulfide: Mg + S MgS Hydrolysis in water 2 MgS + 2 H2O Mg(OH)2 + Mg(HS)2 Mg(HS)2 + H2O Mg(OH)2 + H2S Important Mg salts Carbonates: not soluble Significant analytical reaction for gravimetric P determination: NH4MgP04.6H,0 annealing ] Mg,P,0- Magnesium perchlorate - one of best siccatives: Mg(C104)2 .2H20 .4H20 .6H20 anhydron Organometallic Mg compounds Grignard reagents RX = alkyl- or arylhalogenide Used for alkylation or arylation: R—X + R'— MgX H3C^ ^CH3 O dietherate H3C7 i Mg O \ H3C CH3 R — R' + MgX2 2R — MgX + CdBr2 R2Cd + MgBr2 4 Ph — MgBr + K[BF4] K[B(Ph)4] + 4 MgBrF |[,0 IIjCMgl i II,C 0---C,II50-Mg-I---> C,IIsOII i Mg(OII)I RMgl +■ CO,-----► RCOOMgl------► RCOOII \ Mg(OII)I Significant coordination Mg compounds 9 „sandwich" complex ) CHH with cyklopentadiene Phytyl ;. ■■' ■'.; |. ■'.; Chlorophyll Use of Mg and some compounds Mg - technically very important metal > Light alloys > Construction material in aeronautics, car industry, space ship building MgO Grignard reagents in organic synthesis Calcium, strontium, barium Sources of Ca: limestone (calcite) I Island limestone CaCO: [Gypsum C;iSO,.2il20 ^Anhydrite [ľílSO, CaF? Fluorite Apatite C';i.:(ľO );\ Marmor, chalk, travertine Sources of Sr: Sources of Ba: Production: electrolysis of molten chlorides Remark: Soluble Ba compounds are toxic Hydrides MH2: direct synthesis, reaction with water- prompt H2 source Carbide and calcium cyanamide fertilizer Nitrides: production of deuterited ammonia Alkali earth s metals compounds Sulfides: Ca + S CaS BaSO4 + 2 C BaS + CO2 Alkali earth's metals compounds Oxides: calcination of carbonates at approx. cca 900 °C CaCO3 CaO + CO2 burnt lime Hydroxides: CaO + H2O Ca(OH)2 lime hydrate in mortars Ca(OH)2 + MgCl2 Mg(OH)2 + CaCl2 Mg extraction from sea water Peroxides: Ca(OH)2 + H2O2 CaO2.8 H2O + 2 H2O 2 BaO + O2 2 BaO2 Calcination at 500 °C BaO2 + H2SO4 H2O2 + BaSO4 This reaction was used for hydrogen peroxide production Alkali earth's metals compounds Fluorides: commonly little soluble in water CaF2 used for fluorine production (electrolysis of molten salt) Chlorides: CaCl2. 2 H2O CaCl2 anhydrous - siccative that can be regenerated by heating All anhydrous halogenides are soluble in many organic solvents (alcohols, ethers, etc.) - formation of solvates. Alkali earth's metals compounds - salts Calcium nitrate : fertilizer Calcium carbonate: mountain range are formed from substance calcite aragonite Karst effect: CaCO3 + H2O + CO2 Ca(HCO3)2 Ca(HCO3)2 in water - leads to temporary „hardness" of water Alkali earth's metals compounds - oxygen-containing salts Phosphates: Ca3(P04)2 CaHP04 Ca(H2P04)2 non-soluble soluble Sulfates: commonly little soluble compounds CaSO4 - its presence in water leads to permant water hardeness CaS04.2H20 _. ______ „_ 100"C gypsum ^CaS04.1/2H,0 plaster BaSO4 ■ very insoluble, used for gravinetric determination of sulfates ■ pigment ■ contrast agent used in X-ray examinations of digestive tract Tendency of solubility Little soluble are: hydroxides, sulfates, oxalates, carbonates Chromates, phosphates, fluorids Hydroxides Sulfates Oxalates Be Mg Ca Ca t Ca 1 Sr Sr Sr Ba t Ba Ba ▼ little solubility great solubility Alkali earth's metals compounds - coordination compounds ■ Formation of complexes is not typical ■ Well-known are compleses with polydentate ligands (EDTA, macrocycic ligands) 3rd group PSE, ns2np1 Boron, aluminium, gallium, indium, thallium ♦> B is non-metal, Al, Ga, In, and Tl are typical metals ♦> formation of boranes is typical for B ♦> B forms compouds containing covalent bond, compounds of other elements are mostly ionic ♦> electropositivity in group increases Selected properties of 3rd group elements B Ga In Tl El. configuration (He) (Ne) (Ar)3d10 (Kr)4d10 (Xe)5d10 2s22p1 3s23p1 4s24p1 5s25p1 6s26p1 Electronegativity 2,0 1,5 1,8 1,5 1,4 Radius / pm) atomic 98 143 141 166 171 ionic M(III) - 54 62 80 89 covalent 82 125 126 142 144 Most stable III III I, III I, III I, III oxidation states Melting point / °C 3180 660 30 157 304 Boiling point / °C 3650 2476 2400 2080 1457 Density / g.cm-3 2,35 2,70 5,90 7,31 11,85 Boron Occurance: Production: H3BO3 2 B2O3 + 6 Mg B + 6 MgO 2 BCl3 + 3 Zn 2 B + 3 ZnCl2 BI3 2 B + 3/2I2 Decomposition on heated W-fibre t Boron - common properties ❖ B chemistry is similar to Si - diagonally similarity ❖ B atom has 4 bonding orbitals (s +3p), but only 3 valence electrons ❖ B chemistry is given by small boron atom, high ionization energy, and electronegativity => a lot of interesting compouds are formed ❖ B is typically 3-bonded, another electron pair is accepted => B is then 4-bonded ❖ Formation of polycentric electron deficit bonds is typical (boranes) ❖ Existence of these type of bonding leads to semiconductivity of B Elemental boron t -\ 1 - B12 in upper layer 2 - B12 in lower layer a -trigonal boron Boron - reactivity V Crystalline boron is very little reactive, while amorphous boron is more reactive V At high temperatures, direct reaction with oxygen, nitrogen, halogens, and sulfur => formation of B2O3, BN, BX3 and B2S3 v Boiling HNO3 and molten NaOH lead to B oxidation B + 3 HNO3 2 B + 6 NaOH H3BO3 + 3 NO2 Na3BO3 + 3 H2 Boron - use > Additive to alloys (moderating material in nuclear technology) > Fibrous boron containing W-core is used in cosmic technology > Boron nitrides are very hard - used for cutting edges of instruments and for metal surface treatment Boron - compounds OrideS - binary, often non-stoichiometric componds of boron and metal |lf|§ M5B - MB66 200 compounds, very hard materials Sc2O3 + 7 B — 2 ScB2 + 3 BO BCl3 + W + 1/2 H2 -► WB + Cl2 + HCl 2 TiCl4 + 4 BCl3 + 10 H2 -► 2 TiB2 + 20 HCl Eu2O3 + 3 B4C -► 2 EuB6 + 3 CO Boride production - in eletric oven Application of borides: ❖ abrasive materials ❖ Used as extremely exerted material for turbine paddles, rocket jets, etc. Boron - structure of borides Boron - boranes Boranes - very large group of covalent boron compounds - Nobel price 1976) Preparation and production 4NaH + B(OCH3)3 - NaBH4 + 3CH3ONa 2NaBH4(s) + 2H3P04 l, anhydrous B2H6 (g) + 3NaBH4 + 4BF3.(C2H5)20 - 2BJ-L + .... Diborane reactions B2H6 + 3 O2 -► B2O3 + 3 H2O B2H6 + 6 H2O -► 2 H3BO3 + 6 H2 Boron - boranes Diborane B2H6 Diborane reactions B2H6 + 3 O2-► B2O3 + 3 H2O Boron - boranes Other boranes (besides B - H - B bonds B - B - B bonds are also present) CIOSO- boranes B„hL+2 ^H,/" (n = G al 12) nido-\ boranes 3„hL+4 n/(n+1) arachno- boranes B„hL+6 n/(n+2) hypho- boranes B„H„+8 n/(n+3) conjuncto- boranes Boron - boranes Boron - boranes Boron - boranes Isomers of conjuncto - B20H184- 6 - 6 10 - 10 6 - 10 Sloučeniny boru Carboranes - boron atoms are partially replaced by C-atoms => anions Boron - borane derivatives Boron - borane derivatives Boron - borane derivatives Chloroderivive of cobalt dicarbollide H{DKCoCl7] is a strong acid => used for Cs extraction from nuclear fuel waste solutions Boron - oxides Boron trioxide B + O2 —► B2O3 B2O3 - polymer, that can be prepared by careful dehydratation of H3BO3 (the reaction is reversible) ■ amorphous non-easily crystallizing compound ■ polymeric character ■ contains planar not regularly arranged BO3 groups connected through O- atom ■ in crystalline form, basic units are BO4 tetrahedrons forming chains Boron - acids Trihydrogenboric acid (orthoboric) - H3BO3 II3BO3 2H20 : > II30+ + B(OII) Weak acid, volumetric determination is done in the presence of mannite Preparation Na2B4O7 + H2SO4 + 5 H2O-► 4 H3BO3 + Na2SO4 layer structure ■ layers are formed by trigonal BO3 units connected by hydrogen bridges ■ Reaction with H3BO3 + 3 CH3OH -► B(OCH3)3 + 3 H2O a|coho|s Boric acid trimethyl ester Boron - acids (HBO2)n can be prepared by very careful dehydratation of H3BO3 at 180°C Polymeric compound containing trimer units B3O3(OH)3 Similar to H3BO3 Hydrogenbori (metabonc) - (HBO2)n Boron - borates Borates - their structures has a lot of common features as compared with silicates _, 1"" ' Only few compounds Basic building units: - planar BO3 and tetrahedral BO4 groups ■ joined through O-atoms to chains or cycles - both unit types are often present in one molecule Boron - borates Borates Na6H4B2O8 Peroxoborates ■ derived from borates (e.g. NaBO34H2O2) ■ containig peroxidic -O - O - group bonded to boron atom ■ having significant oxidation ability => used in washing agents Boron - sulfides Boric sulfide B2S3 , white crystalline compound decomposing in water B2S3 + 6 H2O -► 2 H3BO3 + 3 H2S Boron - halogenides BX3 (X = F Cl Br I) BF3 gaseous, BCl3 and BBr3 liquids, BI3 solid Preparation and B2O3 + 6 HF —* 2 BF3 + 3 H2O production B2O3 + 3 C + 3 Cl2 -► 2 BCl3 + 3 CO Reactions Formation and hydrolysis of tetrafluoroborates BCI3 + 3H20 ■ H3BO3 + 3HCI BCl3 hydrolysis 2 H3BO3 + 8 HF + 2 HBF4 + 6 H2O Other possibility of HBF4 preparation Boron - adducts BF3 + NH3-► BF3.NH3 Aduct formation BF + Et O _► BF 2 Et O BF3. 2 Et2O - liquid, enabling easy storing 3 aduct BCl3 with acetonitrile BCl3.CH3CN Reactions of boron halogenides with Grignard agents in anhydrous media lead to the formation of organometallic boron compounds (R = alkyl, aryl) BX3 + 3 RMgX * BR3 + 3 MgX2 Boron - nitrides Boron nitride BN ■ very stable white substance ■ extremely hard ■ BN is formed in burning of boron in nitrogem atmosphere or by annealing of many nitrogen-containing boron compounds in nitrogen (e.g. borazol) Boron carbide B4C Boron - borazole Borazole B3N3H6 Pseudoaromatic compound, isoelectronic with benzene Reactivity of borazole nad benzen are similar Total hydrogenation yields to B3N3H12 Preparation and production Borazol reactions Boron - borazole analoques BN analogues of naphtalene and biphenyl Occurance Aluminium 3rd most spread element in the nature, present mostly in aluminosilicates (feldspars, micas, zeolites, clay) bauxite (böhmit y-AIO(OH)) corundum a-AI203 Sapphire, ruby, emerald, cryolite Aluminium production Electrolysis of the molten and chemically treated bauxite and cryolite with addition of CaF2 and AlF3 (these additives lead to the decrease of melting point). Metal is reduced on cathode formed by steel vessel and covered with carbon Oxygen is formed on graphite anode, reacts with carbon to CO2. Aluminium - chemical treatment of bauxite before electrolysis Bauxite contains impurities: Fe2O3 SiO2, aluminosilicates, etc. 1st step: Removing of impurities is based on solubility of bauxite in alkali media whereas impurities are insoluble. AlO(OH) + NaOH -> Na[Al(OH)4] 2nd step: Filtration and acidifying of the solution by CO2 Na[Al(OH)4] + CO2 -> Al(OH)3 3rd step: Calcination Al(OH)3 Al2O3 4th step: Electrolysis of Al2O3 melt Aluminium - common properties ❖ formation of covalent compounds ❖ covalent bonds, due to low electronegativity of Al (Al is considered to be a metal, compared with B), are strongly polar ❖ ionic character is observed at compounds with most electronegative partners, e.g. AlF3 ❖ common coordination number in compounds is 4 (sp3 hybridization, tetrahedral) or 6 (sp3d2 octahedral) ❖ most obvious oxidation number is III+ ❖ compounds with oxidation number I+ are also known (AlCl) ❖ aluminium is silver-similar, soft, light, and very malleable metal ❖ relatively good electrical conductor ❖ resistant towards air corrosion - compact Al2O3 layer on the surface is formed ❖ no reaction with water; only after removing protective corrosion oxide or hydroxide layer, e.g. by amalgamation using mercury Aluminium - chemical behavior Aluminium is amphoter - soluble in acids and hydroxides Remark: concentrated and oxidizing acids lead to the passivation of the aluminium surface Aluminium salts hydrolyze Example: Aluminium - chemical rections Direct reactions ❖ high affinity to oxygen, Al2O3 formation ❖ Al reacts with sulfur to Al2S3 ❖ with halogens corresponding halogenides of the type AlX3 (anhydrous are known in the form of dimers Al2X6 ; AlX also exist ❖ in the presence of phosphorus, AlP can be obtained ❖ reaction with C yields carbide Al4C3 (methanide) Reaction with oxygen at higher temperatures is strong exothermic (aluminothermic reaction) Cr2O3 + 2 Al -► 2 Cr + Al2O3 Fe2O3 + 2 Al — -► 2 Fe + Al2O3 "Termite" - used for welding Aluminium - compounds with hydrogen Aluminium hydride (alane): Preparation, production Al2Cl6 + 6 LiH 2 AlH3 + 6 LiCl Et20 3LiAIH4 + AICI3 ____> 4[AIH3.(Et20)n] + 3LiCI ❖ Polymeric character ❖ (AlH3)x - bonds Al-H-Al ❖ Al is octahedral coordinated ❖ Decomposes in the presence of humidity AlH3 + 3 H2O -► Al(OH)3 + 3 H2 Aluminium - compounds with hydrogen Tetrahydridoalumates v reaction of alane or aluminium halogenides with alkali metal hydrides in ether AlH3 + LiH Li[AlH4] high-pressure synthesis from elements in industry v Na + Al + 2 H2 hydrolysis in moist air and water Na[AlH4] Used in preparative chemistry as reduction agents. Aluminium - compounds with oxygen Aluminium oxide Al2O3 - white, hard, and very inert substance that can be obtained by combustion of Al in oxygen or by calcination of Al(OH)3 Occurance in several modifications Corundum oc-Al2O3 with anions O2- - most tight hexagonal arrangement with octaedral cavities, occupied from 2/3 Al3+ ions (p=4 g.cm-3) If rest cavities are occupied by other ions => coloured precious stones (red ruby - Cr3+, blue sapphire - Fe3+, green emerald V3+) Cubic y-Al2O3 ("activated" aluminium oxide), higher reactivity, strong sorption capability; (p=3.4 g cm-3), at high temperatures yields a-modification - ALUMINA Al2O3 - fiber form, similar to ZrO2 - 0 3 um, length up to several cm, thermal stability to 1400 °C, used instead of harmful asbestos as insulating and filtrating material, carrier for catalysts, etc. Spinels - MeAl2O4 - double oxides formed together with metals (Me = Ca, Mg ) Use: abrasive pastes, standard for thermal analysis, sorption material Aluminium - compounds with oxygen a-Al2O3 Y-M2O3 Aluminium - compounds with oxygen Spinel MgAl2O4 Aluminium - compounds with oxygen Aluminium oxide- hydroxide - AlO(OH) •:• known in two forms (a-diaspore a y-bohmite); in bauxite. ❖ preparation by slow precipitation from aluminium salts at higher pH Aluminium hydroxide Al(OH)3 ❖ two modifications: bayerite a-Al(OH)3 y-Al(OH)3 (gibbsite - hydrargillite) ❖ white voluminous precipitate of amphoteric character AI(OH)3 + OH- [Al(OH)J- AI(OH)3 + 3 H3O+ * [AI(H2O)6]3+ * Alumates - e.g. Ca3Al2O6, a component of Portland cement Aluminium - compounds with oxygen Structure of cyclic alumate Ca3Al2O6 Aluminium — salts Aluminium salts V aluminium sulfate Al2(SO4)3 . X H2O (x = -> 18), soluble in water, hydrolysis => acidic solution [Al(H2O)6]3+ + H2O [Al(H2O)5(OH)]2+ + H3O+ In waterworks used for water cleaning => adsortion of impurities on surface of soluble , voluminous, and little soluble hydroxocomplexes V Aluminium nitrate Al(NO3)39H2O good solubility V Aluminium acetate Al(CH3COO)3, used in medicine for treatment of swellings as a compress Alums M'MeIII(SO4)2-12H2O (Me = Al, Fe, Cr, V aj.) V alums are white (K- Al) but also coloured (violet K-Cr) substances V isomorfous mutually V cubes, in corners [M(H2O)6]2+ and [Me(H2O)6]3+ alternate Aluminium - halogenides Aluminium halogenides AlX3, Al2X6> AlX a AlX2 are also known) Preparation: reaction Al with anhydrous HX or direct reaction of elements (except AlF3) 700 °C Al2O3 + 6 HF -► 2 AlF3 + 3 H2O ❖ AlF3 is most stabte typical ionic compound, high m.p. (over 1200 °C ), exists in two modifications (a a p) ❖ other halogenides form dimer molecules Al2X6 easily; two tetrahedrons merged with edge ❖ AlX36 H2O ❖ anhydrous halogenides cannot be prepared form hydrates by heating => hydrolysis Aluminium - halogenides AICI, (s) CI—Al ^Cl c.n.= 3 Al2Cl6 catalyst for many organic reactions (Friedel-Craft reactions) Aluminium - halogenides Friedel-Crafts acylation Aluminium - coordination compounds [Al(H2O)g] 3+ in aqueous solutions [AlH4] -, [AlX4] Al in anionic form, coordination number 4 [AlFg] 3-, [Al(C2O4)3] 3- Al in anionic form, coordination number 6 Aluminium - coordination compounds Fluoroalumates cryolite Al2()3 + 3Na2CX)3 + 12HF---► 2Na3(AlF6) + 3C02 + 6H2() 2 Al(OH)3 + 3 Na2CO3 + 12 HF -> 2 Na3[AlF6] + 3 CO2 + 9 H2O Cryolite - used in Al production (a component of electrolyzed mixture) Aluminium - organometallic compounds Mg3AI2 + 6C2H5CI--> AI2(C2H5)6 + 3MgCI2 6Mg(CH3)CI + 2AICI3--> AI2(CH3)6 + 6MgCI2 4Mg(CH3)CI + 2AICI3 -► AI2(CH3)4CI2 + 4MgCI2 Very reactive, on air self-igniting substances, rapid reaction with water AlR3 + 3 H2O -> Al(OH)3 + 3 RH Used in synthesis as alkylation or arylation agents Together with TiCl3 Ziegler-Natta catalysts (alkene production) Aluminium - utilization ❖ light alloys in aircraft, spacecraft and car industry dural (Al + Cu + Mg + Mn) magnalium electron silumin (+ Si) ❖ Al as conductor in electrotechnics ❖ Al as reduction agent (aliminothermy) ❖ Al for production of thin foils ❖ Al in catalytically affecting mixtures IV th group PSE, ns2np2 Carbon, silicon, germanium, tin, lead > C and Si are metals, Ge is semimetal, Sn and Pb are typical metals > C is present in many organic substances (because of carbon atoms chain-formation and possibility to form various bond types) exist (=> organic chemistry) > Inorganic chemistry of C is poor, as compared with organic chemistry IV th group PSE - properties of elements C Si Ge Sn Pb Atomic number 6 14 32 50 82 Density / g cm-3 3,51 2,32 5,38 7,31 11,48 M.p. / °C 4070 1686 1232 505 600,7 B.p. / °C 4620 2570 2970 2543 2010 Covalent radius / pm 77 117 122 140 154 Ionization energy / eV I1 11,25 8,15 7,89 7,34 7,41 24,37 16,34 15,93 14,63 15,03 47,87 33,46 34,21 30,49 31,93 64,5 45,1 45,7 40,7 42,3 Oxidation degrees -IV II, IV -IV (II), IV II, IV II, IV II, IV Electronegativity (Allred-Rochow) 2,50 1,7 2,0 1,7 1,6 tendency to the formation of "inert electron pair" ns2 IV th group PSE - properties of elements ❖ acid oxides at higher oxidation states (CO2, SiO2 i PbO2) ❖ Sn(OH)2 and Pb(OH)2 are amphoters ❖ compounds with hydrogen, MH4, are volatile. ❖ their stability decreases from C to Pb with decreasing energy of bonds C Si Ge Sn H O F Cl C 347 322 297 222 414 351 485 330 Si 322 176 293 465 540 360 Ge 297 159 310 360 356 Sn 222 142 259 343 [kJ mol-1] ❖ tendency to form chains of atoms is decreasing in series C—C, Si—Si ❖ C-C chains are very stable, Si—Si, Ge—Ge chains undergo to oxidation easily ❖ the same tendency is also observed for many covalent compounds containing bonds (X-F, X-Cl, X-Br, etc.) IV th group PSE - bonding ❖ Only C forms np bonds - C=C, C=C or C=O, C=N, C=N sp, sp2sp3 ❖ Si, Ge, Sn, and Pb do not form this type of bonds ❖ Si analogues of CO2, CaCO3, R2CO have quite different structure and properties ❖ Si is able to create npd bonds (non occupied d-orbitals) sp3 and sp3d2 ❖ Sn, Pb - oxid. state II - tendency to „inert electron pair"s2 => non-equivalent sp2 hybridization => in SnCl2 (g) bond angle Cl—Sn—Cl is 95° (expected value is 120°) Carbon C in nature 98.89 % 12C 1.11 % 13C radioactive 14 c 14C (p-irradiator, half-time 5570 years) > formed in high atmosphere layers > ratio 12C/14C is constant > radiocarbon method of dating Inorganic carbon: diamond, graphite, fullerenes Carbonates: Limestone CaCO3; Dolomite CaCO3.MgCO3; Magnesite MgCO3 Organic carbon : coal, crude oil, asphalt, natural gas Carbon - allotropic modifications Cubic diamond, tetrahedral lattice C in sp3, in the tetrahedron centre C—C 154 pm Bond angles C-C , 109,5 o => extraordinary hardness (10) high m.p. => low chemical reactivity => non- conductor Diamond is metastable carbon modification Conversion to stable graphite is possible by heating at 1750 °C in oxygenless atmosphere. Carbon - allotropic modificatioms Diamond properties v found diamonds form well-developed octahedrons v pure and cut diamond - brilliant, refraction is very high (2.42) V mass of diamonds is given in carats (1 carat = 0.2 g) v non-transparent or black diamonds and artifitial diamonds are used as abrasive material V synthetic diamonds are made from graphite dissolved in molten metal (Ni, Co) at 2000 - 3000 °C, 10 GPa V chemical reactivity of diamonds is low V heating of diamond up to 930 °C - combustion V strong oxidation mixtures (melting with KNO3, conc. H2SO4 + K2Cr2O7) lead to diamond oxidation to CO2 Carbon - allotropic modifications - diamond Carbon - allotropic modifications - graphite Carbon - allotropic modifications - graphite ❖ C atoms in 6-membered cycles, sp2 hybridization ❖ C-C distance 141,5 pm, i. e. shorter as compared with single C-C bond => bond order 1,33 => n - electrons are delocalized ❖ good thermal and electric conductivity of graphite Microcrystalline graphite forms: :■;■[■[ j_■: -e-e-e-e-e-e-e-e^-e-e-e-e-e-e-e-e-i-e-e-e-e-e-e-e-e-e-e-e ❖ Black carbon Black carbon ❖ „Shiny" carbon i-e-e-e-e-e-e-e-e-e-e-t-e-e-e-e-e-e-e-e-e-e-e-i-e-e-e-e-e-e-e-e-e-e-e i-e-e-e-e-e-e-e-e-e-e-t-e-e-e-e-e-e-e-e-e-e-e-i-e-e-e-e-e-e-e-e-e-e-e Soot Carbon - allotropic modifications Graphite is more reactive than diamond, mainly at higher temperatures graphite + O2 graphite CO2 F2 CF4, C2F2, C5F12, CFx metals carbides H2 + 1 catalyst 1 -- Si C2H2 SiC S CS2 H2O Carbon - allotropic modifications - graphite Intercalates CxOy (x:y = 2:1,H) C24X (X = HS031 NO3) chloridess MC ,{x = 2 - 6) Dxides ( : S02: NO sulfides (V2S3, PdS, Sb2S5) Carbon - allotropic modifications - graphite Carbon - allotropic modifications - fullerenes (known from 1985) ❖ structure od fullerenes resembles football ❖ alternating 5- and 6-memberes cycles ❖ ratio of cycles leads to the formation of clusters: C60, C70, C76, C78, C90 , etc. C60 C70 Carbon - aIIotropic modifications - fuIIerenes BuckminsterfuIIerene C6Q FuIIerene C54Q Carbon - allotropic modifications - fullerenes Architect: J. Buckminster Fuller Carbon - allotropic modifications - fullerenes Fullerene compounds C60OsO4(NC5H4But)2 Nanotubes with outstanding properties - perspective of utilitization, e.g. in catalysis Carbon - allotropic modifications - fullerenes C60 [C60 (ferrocene)2] Carbon - other allotropic modifications Glossy carbon (amorphous) - production by controlled thermal decomposition of some polymers (polyacrylates). Used in electrochemical processes. Coal - anthracite, black and brown coal, lignite -(„Impurities" in graphite) Carbon - carbides Ionic carbides (mostly as acetylides) MzQa (M - Cu, Ag, Au) MC2 (M = Zn, Cd, Hg) MHC2 (M = Li, Interstitial carbides TiC: ZrC: VC: V,C: MoC: Mo,C: WC, W,C r > 130 pm ❖ Metal atoms in structures can be replaced stepwise by C-atoms => existence of many non-stoichiometric compounds ❖ Electric conductivity is usually preserved ❖ Increasing C-content => m.p., hardness is higher, as well as other physical constants are also changed ❖ TiC, ZrC, Mo2C, WC „hard metals" - made by sintering - "sintered carbides" - for cutting instruments B Carbon - carbides Transition carbides r < 130 pm Between interstitial and ionic carbides Covalent carbides Extra hard, diamond structure, chemically very stable, high m.p., used as cutting material SiO2 + 2 C — SiC + CO2 (carborundum) 2 B2O3 + 8 C — B4C + 6 CO Carbon - carbides Sorting according to the reaction with water a) methanides Be2C a Al4C3 Be2C + 4 H2O 2 Be(OH)2 + CH4 b) acetylides yield ethin (acetylene), e.g. CaC2 CaC2 + 2 H2O —> Ca(OH)2 + C2H2 (acetylene production) c) Mg2C3 hydrolyses, releasing of allylen Mg2C3 + 4 H2O 2 Mg(OH)2 + HC = C-CH3 Carbon - carbides Carbide prepations: 1) direct synthesis from elements at high temperatures: 2 V + C V2C 2) reaction of C with metal oxides at high temperatures SiO2 + 2 C SiC + CO2 3) reaction of „acid" hydrocarbons with metals or their compounds C2H2 + 2 [Ag(NH3)2]+ Ag2C2 + 2 NH4+ C2H2 + R2Zn ZnC2 + 2 RH Carbon - oxocompounds ^Suboxides" C3O2, C5O2 or C12O9 (no practical significance) Malonic acid anhydride Mellitic acid anhydride c,2o9 Carbon - oxocompounds - CO Carbon monoxide CO Preparation HCOOH HlS°J (COOH)2 ^ » CO + -* CO + HaO C02 Production 2 C + O2 - - 2 CO CO2 + C —" 2 CO C + H2O CO + H2 + H,0 Generator gas - 25 % CO, 70 % N?1 4 % CO Boudoard equilibrium (g) AH0 = 172.6 kJ mol"1 Water gas - 40% CO, 50 % H,, 5 % COPI 5 % N. Carbon - oxocompounds - CO Reaction of CO CO + atfivatedjpe h2 Ni, Co h2 ZnO o2 NaOH CH3ONa Cl2 metals petrol CH4 CH3OH CO2 HCOONa CH3COONa COCl2 carbonyls Carbon - oxocompounds - carbonyls Metal carbonyles CO is a ligand of the type donor a- acceptor n Direct synthesis, e.g. Ni(CO)4, Fe(CO)5, Cr(CO)6 Donor a-bonding is relatively weak. Stability of these compounds can be ascribed to back-donation of metal d-electrons into antibonding n* orbitals Weak a-donor properties were observed towards some weak Lewis acids B2H6 + 2 CO —> 2 BH3CO CO is very poisonous - a carbonyl complex with Fe atom in haemoglobin is formed - this complex is more stable as compared with similar oxygen containing complex Carbon - oxocompounds - C02 Carbon dioxide C (s) + 02 (g) - CO, (g) 0=C=0 C02 CaC03 —i____CaO + C02 t * acid gas, easily can be liquidified * soluble in water and some less polar solvents * CO2 is formed in the combustion process of organic compounds at sufficient access of oxygen * CO2 is formed in reaction between carbonates and acids CaCO3 + 2 HCl — CaCl2 + H2O + CO2 MgCO3 + H2SO4 - MgSO4 + H2O + CO2 * when released from a bomb => a solid „dry ice" is formed Carbon - oxocompounds - H2CO3 CO2 is unstable carbonic acid anhydride CO .(x-l)H-O CO/" + H20 i HCO; + H20 . * HCCV + OH" ' H2C03 + OH Hydrolytic reaction Very weak acid, forming two salt series: • hydrogencarbonates • carbonates Carbon - oxocompounds - carbonates Carbonates (M1 , M11 , Mm) Preparation: 2 NaOH + CO2 Na2CO3 ❖ Only ammonium and alkali metal carbonates are soluble (except Li-carbonate). ♦♦♦ As a result of hydrolysis - their aqueous solutions are strong alkaline. ❖ Thermally stable are only alkali metal carbonates (except Li2CO3 and ammonium carbonate => different decomposition mechanism). ❖ Other carbonates yield CO2 and metal oxide at higher temperatures. ZnCO3 CO2 + ZnO Carbon - oxocompounds - carbonates Trivial names of some carbonates soda Na2CO3. 10 H2O calcinated soda Na2CO3 soda bicarbonate NaHCO3 potash K2CO3 ammonium ("confectionery yeast") NH4HCO3 Structure of NaHCO3 Carbonate utilization: • glass production (soda, potash) • soap production (soda, potash) • in building industry (limestone CaCO3, magnesite MgCO3> etc.) Carbon - oxocompounds - peroxocarbonates Peroxodicarbonates Anodic oxidation od concentrated solutions of alkali metal carbonates: 2 CO32- -- C2O62- + 2 e- Peroxocarbonates M2CO4x H2O are also known, their composition is estimated as peroxohydrates M2CO3H2O2y H2O. Peroxocarbonates are used instead of more expensive peroxo borates in washing detergents. Carbon - oxocompounds - C + S compounds Carbonyl sulfide COS - colourless smelling gas, COS is formed in reaction of CO with sulfur vapour Carbon disulfide CS2 = carbon dioxide thioderivative Production: 2S(g) + C(s)____> CS2 (g) AIT = 104.2 kJ mol" (sulfur vapour is drifted over glowing coal) CH4 + 4 S -- CS2 + 2 H2S ~ 600 °C Al^ ' . 4 2 2 Catalysis Al2O3 or gel. SiO2 (natural gas or methane) Properties: excellent solvent for white phosphorus, extremely flammable Reactions: CS2 + 2 H2O — CO2 + 2 H2S CS2 + 3CI2 CCI4 + S2CI2 CS2 + K2S K2CS3 Tetrachlormethane production thiocarbonates (dichlodisulfane is useful by-product) Carbon - halogenides CS2 + 3CI2-----> CCI4 + S2CI2 SbF freons ■4HCI -2HCI 2CHCI3 + 4HF--> 2CF2CIH----► C2F4---> 1/n(C2F4)n Teflon Freons - earlier used as cooling media for refrigerators Demaging effect on ozon layer - their utilization is forbidden. Carbon - function derivatives of carbonic acid Carbonyl halogenides COX2 (X = f, ci, Br) Preparation: CO + X2 COX2 Phosgene is very toxic. Very reactive, used in organic synthesis Carbonyl dichloride COCl2 (Phosgene) COCl2 + 2 H-Y Y2CO + 2 HCl Y = OH, OR, NH2, NHR, NR2. CSCl2 (thiophosgen) exists, too. Carbon - function derivatives of ca rbon acid Esters (RO)CO(NR2) (carbamates) - pesticide production Carbon - CN compounds Hydrogen cyanide HCOONH4 + P2O5 -> HCN + 2 HPO3 HCN Preparation: 2 KCN + H2SO4 -> K2SO4 + 2 HCN Production: * HCN is very toxic, * easily-liquefiable gas (b. p. 25,6 °C), formation of H-bonds * very stable polar triple bond -C=N == dissociation H-C=N in aqueous solution. HCN is weak acid (Ka = 2,11Q9) 60 % produced HCN is used in production of acetonitrile, iion: acrylonitrile, and methyl methacrylate. Cyanides Production: Properties: Carbon - CN compounds CaC2 + N2 — CaCN2 + C CaCN2 + C + Na2CO3 — CaCO3 + 2 NaCN 2 NaNH2 + C — Na2CN2 + 2 H2 Na2CN2 + C — 2 NaCN ❖ CN- is isostructural and isoelectronic with CO ❖ excelent ligand ❖ hydrolysis in aqueous solutions (=> high pH ) ❖ heavy metal cyanides are explosive Utilization: Gold leaching from ores - cyanide procedure: Carbon - CN compounds Cyanogen (CN)2 - gas of bitter- almond smell, very toxic Production: Structure: Disproportionation: Similar reaction course in halogene group is observed == (CN)2 is called pseudohalogene Cyanogen is stable also at higher temperatures, presence of impurities leads to dicyane polymerization at 280-380 - == paracyanogen (CN)x ^ > /C>^ y"C^ /C\ / N VN N Carbon - CN compounds Cyanogen is very reactive: ❖ Reaction with water gives oxalic acid diamide (CN)2 + 2 H2O (CONH2)2 ❖ Reduction with hydrogen yields 1,2 - diaminoethane (ethylendiamine - known ligand) (CN)2 + 2 H2 H2NCH2CH2NH2 Carbon - CNO compounds Hydrogen cyanate and its isomers H-N=C=0 Hydrogen isocyanate IN=C-0-H Hydrogen cyanate IN=C-0 h—c=n—o Fulminic acid ic=n-oi Isomers can be distinguished at esters Preparation: CO(NH2)2 HNCO + NH3 Hydrolysis: HNCO + H2O CO2 + NH3 Carbon - CNO and CNS compounds Isocyanates CNO- can be easily prepared by cyanide oxidation KCN + PbO V KNCO + Pb Hydrogen thiocyanate H- N=C=S is strong acid Thiocyanates (rhodanides) Oxidation of cyanides with sulfur KCN + S t KNCS SCN- anion is frequent ligand in complexes . The bond to metal ion can be curried out through N-atom (cyanate complexes) or S-atom (isothiocyanate complexes). Carbon - oth er CN compounds Halogene cyanides, chlorine cyanide CICN NaCN + Cl2 — CICN + NaOl Calcium cyanamide CaCN2 - fertilizer Cyanuric acid H I N I I HCV*N OH I 1 N n c I OH C3H3N3O3 —> 3 HNOO t Silicon Occurance 27, 2 %, quartz SiO2 and silicates Si production: In electrical furnace Production of pure Si:: SiO2 + C (+ Fe) -> (Si,Fe) + 2 CO ferrosilicon ("technical silicon") 1. pure SiCl4 - distillation, reduction by hydrogen in glow to Si 2. thermal decomposition of SiH4 3. reduction of SiCl4 by Mg 4. exotermic reaction Na2SiF6 + 4 Na Si + 6 NaF Production of extremelly pure Si: 99,99 % purity O rozebíratelný spoj \^ _JH posuv vzhůru rotace Monocrystal Graphite vessel í * » ■ * * ■ *" i.V.' 71 Inert gas Germinal crystal Si melt Zone melting Silicon - properties ❖ electron configuration 3s2px1py1 + free d-orbitals ♦> bonding and chemical propertie of C and Si differ significantly ❖ covalent compounds formation ❖ bond energy Si—Si i Si—H is substantially lower than energy of C—C or C—H bonds => silicon analogues of organic compounds are not stable ❖ Si—O bond is more stable than C—O == compounds containing Si—O nebo Si—O—Si bonds are typical for Si ❖ absence of np bonds == absence of alkene and alkyne analoques, graphite, aromatic compounds, etc. ❖ Si atom contains non-occupied 3d orbitals == formation both o- bonds, and TTpd interaction == consequencies in structure and reactivity of many Si compounds. Silicon - bonding Hybridization Bonds Example sp3 4a SÍH4, (CH3)4Si 4a + 277d deloc. SiO44-, SiF4, SiCl4 sp3d2 6a SiF62- Silicon - reactivity ❖ pure Si is grey crystalline substance (cubic, diamond structure similar (distance Si—Si is 235 pm). ❖ very hard, but fragile ❖ chemically not too reactive, more reactive at higher temperatures. Silicon - reactions Direct reactions Si + °2 S Si°2 SiS2 2- g low I H2° Si°3 + H2 Si°2 C halogene SiC SiX4 metals silicides Insoluble in all acids (HF is an exception), in alkali hydroxides silicates are formed Silicon - silanes Silanes - binary compounds with hydrogen m.p. (oC) b.p. (oC) density 103 kg m-3 / oC SiH4 -185 -112 0,68 / -186 -132 -14 0,686 / -25 -117 53 0,725 / 0 Si4H10 -90 108 0,82 / 0 Production of silanes and halogenderivatives: SiCl4 + LiAlH4____► SiH4 + LiCl + A1C13 SiH4 \ HCl-----► SiH3Cl \ H2 Silanes are very reactive / difference from alkanes / - low energy of Si—Si a Si—H bonds) => self-igniting, sensitive to air-moisture SiH4 + 2 O2 SiO2 + 2 H2O 4110 2SiO: 711: Silicon - silicides Silicides (partially similar to carbides) to v only some have stoichiometric composition, e.g. Mg2Si V most of them have character of intermetallic compounds V contain chains or space structures, Si—Si lenghts are similar to (Mo3Si, U3Si2, USi2, CaSi2, BaSi3). chemically very stable V preparation is based on direct synthesis, or on reduction of SiO2 in excess of a metal. Silicon - carbide and nitride Silicon carbide, SiC ("carborundum") Si02(s) + 3C(s) -1 SiC(s) \ 2CO(g) AH°=539 kJraol In electrical furnace Very hard material having diamond structure - used as abrasive agent Silicium nitride, Si3N4 Powder Si3N4 is formed at 1200-1400 0C Si(s) + 2 N2 (g) = Si3N4 Silicon - sulfide Silicon sulfide, SiS2 Different structure, as compared with oxide, as a result of possible greater deformation of bond angles (Si hybridization sp3 is preserved SiS4 tetrahedrons in chains have common edge Si /Si /Si /Si S S S \ Production: Si + 2 S SiS2 Properties: SiS2 is sensitive to air-moisture == decomposition: SiS2 + 2 H2O SiO2 + 2 H2S Silicon - halogenides (formally, they can be considered as silane halogenderivatives ) SiX4, SinX2n+2 (n je pro F = 14, Cl = 6, Br, I = 2) SiF4 Colourless gas b.p. - 95 °C SiCl4 Colourless liquid b.p. 57 °C SiBr4 Colourless liquid b.p . 153 °C SiI4 Colourless crystals m.p. 120 °C Silicon - halogenides Preparation and production Si + 2 X2 SiX4 SiO2 + 2 C + 2 Cl2 —SiCl4 + 2 CO t ^ SiO2 + 4 HF SiF4 + 2 H2O Principle of glass etching) Reaction of silicon halogenides SiCl4 + 2 H2O -- SiO2 + 4 HCl Hydrolysis is possible as a consequence of the presence d- orbitals == therefore CCl4 does not hydrolyse Silicon - halogenides Hexafluorosilicic acid SiF4 + 2 HF H2[SiF6] - very strong acid ❖ stable to 13 % concentration, salts are very stable Sir4 + 21 IF + 21I20----► SiF62" H- 21130+ 3Sil-4 + 6H20----^Sily" + 41130+ + Si02 ❖ anion [SiF6]2- has octaedral structure, Si atom in sp3d2 hybridization Silicon - alkyl and aryl- compounds Alkyl- and arylsilanes V Formally, derived from silanes by substitution of hydrogen atom by alkyl or aryl V Covalent compounds, soluble in non-polar solvents V More stable compared with silanes, are not self-igniting Production SiCl4 + 3 CH3MgCl (CH3)3SiCl + 3 MgCl2 Reactions (CH3)3SiCl + H2O HCl + (CH3)3SiOH (trimethyl silanol) 2 (CH3)3SiOH H2O + (CH3)3Si-O-Si(CH3)3 ("siloxane") Hexamethyl disiloxane (HMDSO) Silicon - alkyl and aryl- compounds - siloxanes Siloxanes (silicons) alkysilane product of hydrolysis product of condenzation RSiCl3 RSi(OH)3 R R R 1 1 1 —O-Si-O-Si-O-Si-O— 1 1 1 000 1 1 1 —O-Si-O-Si-O-Si-O— 1 1 1 RRR R2SiCl2 R2Si(OH)2 R R 1 1 —O-Si-O-Si-O— 1 1 RR R3SiCl R3Si(OH) R3Si - O - SiR3 Silicon - alkyl and aryl- compounds - silazanes Silazanes ♦> Analogous to siloxanes. ♦> Production is similar, only amino-compounds were used for solvolysis (instad of hydrolysis ♦> Bonding: Si-N-Si ■ i Silicon - alkyl and aryl- compounds - properties and utilization Properties of silicons and silazanes Silicones are formed in combination of mono-, di- a trihalogen-alkylsilanes, and used solvolytical conditions => molecular mass and physical properties of formed technical silicone or silazane can be significantly influenced. Silicones and silazanes - liquids, oils, resins Very thermally stable, hydrofobic, electrically and thermal non-conductive. Utilization ❖ silicone lubricants, silicone oils ❖ insulators ❖ rubbers (silicone rubber) ❖ hydrofobizing liquids for reconstruction od buildings (Lukofob) ❖ hydrofobizing liquids for conservation Silicon - oxygen-containing compounds SiO Can be formed at high temperatures, non-stable, easy oxidation to SiO2, (burning on air) Silicon dioxide, SiO2 ❖ diametrically differs from CO2 ❖ Si atom - sp3 hybridization, Si in tetrahedron centre, O-atoms can be found in tetrahedron tops ❖ SiO2 structure is macromolecular, SiO4tetrahedrons are merged through their tops ❖ Arrangement of tetrahedrons SiO4 => existence of three modifications: quartz, tridymite and cristoballite (3-quartz 867oC 1470oC 1731oC ■ » p-tridymit ^-» p-cristobalit - " 573 oC a-quartz 120-160 oC a- tridymit 200-280 oC a-cristobalit melt Silicon - SiO2 Properties of SiO2 ❖ all SiO2 forms are chemically very resistant (see Si-O bond energy). ❖ reduction with C or Mg, Al, resp., at high temperatures ❖ SiO2 reacts only with HF and alkali hydroxides or carbonates, Si—O—Si bonds are split SiO2 + 2 NaOH — Na2SiO3 + H2O SiO2 + Na2CO3 — Na2SiO3 + CO2 Utilization: technical SiO2 (sand form) - glass production and in building industry quartz-glass for chemistry Silicon - SiO2 Quartz glass ❖ Melting and quick cooling lead to the collapse of crystal structure in SiO2 ❖ SiO4 tetrahedrons are merged accidentally => quartz glass ♦> Quartz glass - amorphous glassy substance, having some good practical properties (low expansivity coefficient, high m.p., transparency for UV). ❖ Prolonged heating near m.p. leads to the formation of small crystals - glass disintegrates. Utilization of quartz glass: > makinf parts of quartz apparatus > cuvettes for UV spectroscopy > quartz bulbs for UV quelles /lamps Silicon - Si02 A lot of crystalline and amorphous, anhydrous and partially hydrated minerals can be found in nature as semi-presious stones - used in jewellery Smoky quartz Chalcedony Rosy quartz finr Agate | MP^^ 1 Amethyst Opal Silicon - SiO2 meta acid ortho acid very weak acid pK « 12 H4SiO4 can be relaased from silicates by acidification non-stable, immediately condensates in acid media => formation of Si—O—Si bonds and amorphous gels of polymeric Si IV acids silica gel Silicon - acids Si acids o o 0—H H—O h o 0—H H—0 h Silicon - silica gel ♦> large surface - absorption of water, gases, impurities, etc. • absorption of water and gases is reversible Utilization: • adsorption material for chromatography • Silufol for TLC (thin-layer chromatography) • drying agnet used in exsiccators Silicon - silicates Alkali silicates - soluble in water Si02 + 2NaOH (Na.co,)-----* Na^SiO^ + H20 (co2) "water glass" 6SiOj + NatCO} + CaCO,--> Na2O.Ca0.6SiOj + 2COj "insoluble " glass Silicon - silicates Other Silicates: (insoluble in water) ❖ various arrangement of SiO4 tetrahedrons => variability of structures ❖ bonding with neighbouring tetrahedrons through 1, 2, 3, 4, resp., bridges => chains (1, 2 bridges), planar (3 bridges) or three-dimensional (4 bridges) structures. ❖ some of Si atoms can be replaced by divalent (Be, Mg), trivalent (B, Al), tetravalent (Ti), but also fivevalent (P) elements Silicon - silicates > olivine (Mg,Fe)2SiO4, > garnet Me3IIMe2III(SiO4)3, kde Me" = Ca, Mg, Fe and Mem = Al, Cr, Fe > hemimorphite Zn4(OH)2Si2O7H2O > benitoite BaTiSi3O9 > wollastonite a-Ca3Si3O9 > beryl Be3Al2Si6O18 Silicon - silicates ZIRCON Garnet Silicon - silicates Siliocates with chainy or band structure pyroxene apmfibole Spodumene LiAl(SiO?)2 Tremolite Ca>lg5(Si/)n)2(OH)2 Silicon - silicates AMFIBOLE A' WE Silicon - silicates Silicates with planar structure Silicon - silicates TURMALINE Silicon - silicates kaolinite Al2(OH)4Si2O5 talc Mg3(OH)2(Si2O5)2 muscovite Kal(OH)2(Si3AlO10) Silicon - silicates BIOTITE Silicon - alumosilicates Alumosilicates with three-dimensional structure Silicon - alumosilicates PLAGIOKLAS ORTHOKLAS Silicon - alumosilicates Zeolites - properties ❖ Zeolites differs from feldspars - contain water that can be reversibly removed (as at silica gel). ❖ Crystal net is formed by units containing cavities of given size. Water molecues or other substances are bonded by van der Waals forces. ❖ Similarly, metal cations can be reversibly exchanged (inorganic ionex) Synthetic zeolitic materials - Synthetically, molecular sieves with determined size of cavitis (400 - 1200 pm) can be prepared Molecular sieves serve for selective adsorption - separation of liquids, gases, siccative etc. Silicon - alumosilicates Zeolites Vth group PSE, ns2np3 Nitrogen, phosphorus, arsenic, antimony, bismut ❖ N and P are non-metals, creating compounds with covalent bonds, As, Sb are semi-metals, Bi - typical metal ❖ N - formation of np bonds, ❖ P - contais vacant 3d-orbitals, npd interaction with strong electronegative elements (F, O) ❖ P - 3d-orbitals are often used in the process of formation higher coordinated compounds (5 a 6) - trigonal bipyramide and octahedral arrangement ❖ As and Sb - amphoters ❖ As and Sb - oxidation effect in oxidation degree V ❖ As prefers coordination number 4, Sb in antimonates - coordination number 6 ❖ Towards H - formal oxidation degree -III , its stability decreases with increacing atomic number (=> wit decreasing energy of M-H bond) Vth group elements - common properties N P As Sb Bi Atomic number 7 15 33 51 83 Density / g cm-3 1,027 1,828 (white) 5,73 6,68 9,80 M.p. / oC - 210 44,1 816 (4 MPa) 630,7 271,4 B.p. / oC - 195,8 280,5 615 (subl.) 1587 1564 Covalent radius / pm 70 110 121 141 146 Ionization energy/ eV I1 14,53 10,488 9,81 8,639 7,287 29,60 19,72 18,63 16,5 16,68 I3 47,43 30,16 28,34 25,3 25,56 I4 77,5 51,4 50,1 44,1 45,3 I5 97,9 65,0 62,3 56 56 Oxidation degrees -III to +V -III to +V -III, +V -III +III +V -III +III +V Electronegativity 3,07 2,06 2,20 1,82 1,9 Nitrogen - element • Earth's core and atmosphere (cca 78 %) • NaNO3 (Chile nitrate) • Ammonium salts • Biogene element - in peptides ❖ Molecule N is isostructural and isoelectronic with : |N=N| |C=O| |N =O|+ |C=N|- ❖ Symmetrical distribution of electron density in N2 molecule + high energy of |N=N| bond => low reactivity of elemental nitrogen. Occurance in nature |N = N| 946,2 kJ.mol-1 |N - N| 159,1 kJ.mol-1 Nitrogen - bonding Hybridization! Bonding types Examples sp3 NH4+, (CH^NO, [Ag(NH3)2J+ 3<7 + 1 fp NH3, NF3, NH2-NH2 2<7 + 2 fp Na+NH2- 1g + 3 vp Li22+NH2- sp2 3g + 1;r deloc. HNO3, NO2Cl, NO3- 2g + 1 fp + 1 deloc. NOF, NO2- 2g + 2;r deloc. NO2+ 1<7 + 1 fp + 2;r NNO (ending N atom) fp - free electron pair Nitrogen - reactivity ❖ Elemental N2 reacts only with few N2 + 2 O2 2 NO2 elements or substances, mostly at high temperatures and presence of catalysts n2 + 3 H2 £== 2 NH3 (Fe, Al2O3) ❖ Reaction of N2 with metals at high temperatures => nitrides: Mg, Ca, Sr, Ba, B, Al, Si a Ti. ❖ Industrial significance of the reaction: CaC2 + N2 75Q°C > CaCN2 + C cya^l^cte ❖ N2 was identified as ligand in complexes in recent tim e - co par iso wi th isoelectron ic CO, NO+ a CN- ( easy complex formation) [C oH(N2) (PPh3)2], [Ru(NH3)5N2]Cl Nitrogen - preparation, production, utilization Preparation (NH4)2Cr2O7 NH4NO2 N2 + Cr2O3 + 4H2O — N2 + 2 H2O 4 NH3 + 3 O2 2 N2 + 6 H2O Preparation of very pure nitrogen Ba(N3)2 Barium azide Ba + 3 N2 Production B.p. (°C) N2 - 196 O2 - 183 Fraction distillation of liquedified air Utilization • ammonia production, calcium cyanamide CaCN2 • inert (protective) atmosphere Nitrogen - N-H compounds Bases: Ammonia ^............'j NH3 Hydrazine 2 4 Acids: Azoimide - Salts: Ammonium azide HN Hydrazinium azide (+1) N2H5N3 Unstable: Diazene (diimide) t-[M.) V) <-30 V) Non-existence: Ammonium hydride N11411 ■Tetraazene: N4H4 Colourless gas, characteristic smell, m.p -77,7 °C, b.p. - 33,4 °C Mg3N2 + 6 H2O — 3 Mg(OH)2 + 2 NH3 Preparation CaCN2 + 3 H2O — CaCO3 + 2 NH3 2 NH4Cl + CaO — CaCl2 + 2 NH3 + H2O NH4Cl + NaOH — NH3 + H2O + NaCl Preparation of deuterated ammonia Production > „ammonia" waters (waste product from gasworks and coke production - reaction with Ca(OH)2 > Haber - Bosch direct synthesis N2 + 3 H2 ff°E 2 NH3 AH = - 92 kJ mol-1, 20-100 MPa Nitrogen - N-H compounds - ammonia Reactions of ammonia Ammonia is well soluble in water 3CuO i 2NII, t 3Cu i 3ll:0 i N: 4nh3 + 30: + 6h.o + 2n. 4 NH3 + 5 O2 75<°Q Pt , 4 NO + 6 H2O Reduction effect of ammonia t Nitrogen - N-H compounds - ammonia Reactions in liquid ammonia 2NII, Nil, i Nil, K « 10-3° 2H20-----H30+ + OH" Ammonolysis leads to amides, imides or nitrides PC13 + 6NH,----► P(NH2)3 + 3NH4C1 Gel4 + 6NH3----► Ge(NII)2 + 4NH SbCl3 + 4NH3---->SbN + 3NH4C14I Nitrogen - N-H compounds - ammonia Utilization ❖ ammonium salts as fertilizers ❖ starting compound for nitric acid production ❖ soda production (Solvay) ❖ NH3(l) - in cooling devices ❖ 25 % aqueous solution ❖ NH3(l) in bombs Nitrogen - N-H compounds - amides, imides Natrium amide - industrial production, used in organic syntheses Na + 2 NH3 (g) —x—> 2 NaNH2 + H2 Hydrolysis in water and hydroxide solutions NH2 + H20 NH3 + OH" ^^^H NH2 + OH NH3 + O2" —\ Metal imides - only few are known, e.g. Li2NH, CaNH. / Imides A Formation by spontaneous thermal decomposition of llllllllll^K amides. • |||§§||W|2LiNH2----v LiAJI + NH3 Pbl2 + 2KNH2----► PbNHl + 2KI + NH3 Nitrogen - nitrides Preparation Ionic Interstitial XN (X = Ti,Zr, V.Nb.U), X?N (X = Mo, W), (Mn3N25 U2N3) Covalent (A1N, BN, SxNy) Hydrolysis AlN + 3 H2O - Al(OH)3 + NH3 Nitrogen - N-H compounds - hydrazine Colourless, on air fuming liquid, m.p. 1,4 °C, b.p. 113,5 °C Hydrazine J Heating can cause decomposition (often with explosion) N2H4 intermediate Production: [ Mi j NH3 + NaOCI — NH2CI + NaOH I NH2CI + NaOH + NH3 — N2H4 + NaCI + H2O CO(NH2)2 + NaOCI + 2 NaOH — N2H4 + NaCI + Na2CO3 + H2O Hydrazine: > is relatively unstable => isolation in form of soluble N2H5+(HSO4)\ > is miscible with water at any ratio. > N2H4- H2O is formed in water (sale form). Nitrogen - N-H compounds - hydrazine Reactions of hydrazine N2H4 — N2 + 4 NH3 disproportionation Reducing porperties of hydrazine - used in precious metal production Hydrogen in hydrazine is „acidic" => hydrazides 2 N2H4 + 2 Na — 2 Na+2H3- + H2 Reaction with acid chlorides => acid hydrazides HSO3Cl + N2H4 + py — HSO3NHNH2 + pyHCl Nitrogen - N-H compounds- hydrazines Utilization of hydrazine: ❖ N2H4 and its methylderivatives are used in as fuel for spaceships ❖ Hydrazine as reduction excellent agent - advantage: reaction products (nitrogen and water) are volatile Preparation of tetrafluorohydrazine NF3+Cu -+N2F4 Nitrogen - N-H compounds - azoimide = Hydrogen azide AZoimide \ Colourless liquid, m.p. - 80 °C, b.p. 35,7 °C, explosive HN / Aqueous solutions have acidic reaction (Ka = 1,8-105). 3 Production N2H5+ + HNO2 — HN3 + H3O+ + H2O Reactions with iodine Nitrogen N-H compounds - azides 3 NaNH2 + NaNO3 2 NaNH2 + N2O 750oC NaN3 + NaOH + NH3 190oC NaN3 + NaOH + NH3 Azides of alkali meta il and alkali ear th elements are more s itable, soluble i n water, the rmal deco mposition witho )ut explosion a nd form ation of elemen tal nitrogen and metal - used in airbags. Heavy metal azides are explosive: AgN3, Pb(N3)2 a fuse Hg(N3)2 used in detonating Hydroxylamine NH2OH > White crystalline substance (m.p. 32 °C) > Well soluble in water. > When anhydrous => disproportionation > Heating leads to explosion > Weaker base than ammonia, Kb = 6,6-109. > Both reduction and oxidation agent. Production Industry CH3NO2 + conc. H2SO4 120oC [NH3OH]HSO4 + CO Nitrogen - N-H compounds - hydroxylamine Hydroxylamine salts: e.g.. [NH3OH]Cl, [NH3OH]2SO4J etc. Ligand: [Zn(NH2OH)2Cl2]. Reaction with aldehydes and ketones => oximes >C = O + NH2OH -> >C = NOH + H2O Diacetyl dioxime (R= CH3) (Tschugaev agent) -used for gravimetric determination of Ni2+, Pd2+ Cyklohexanon oxime - precursor in polymer chemistry £ - caprolactam (polyamides Silon, Nylon, Perlon) Nitrogen - oxides Oxides in oxidation states I - V Azoxide - colourless gas (m.p. -90.8 °C, b.p. -88,5 °C) NO2- Production Chemically reactive, thermal decomposition at higher temperatures (N2 + O2) oxidation properties. Utilization: anaesthetic gas, in small bombs - for making whipped cream Nitrogen oxides H2N2O2 - weak acid, little stable, decomposes under explosion Production N-OH N| cis- trans- HONH, + ONOH--- H2N2O2 -► N2O + H2O HON=NOH + H,0 Salts are more stable : 2 NaNO2 + 8 Na/Hg + 4 H2O Na2N2O2 + 8 NaOH + 8 Hg Reduction by sodium amalgam NO Production ❖ colourless paramagnetic gas ❖ neutral oxide, no reaction with water. ❖ (m.p. -163,6 °C, b.p. - 151,8°C). ❖ non-pair electron is delocalizated along whole molecule, => formation of dimer N2O2 is not possible. ❖ very toxic Preparation INO Reactions with oxygen is very easy: 2 NO + O2 -- 2 NO2 Powerful oxidation agent => HNO3 Reduction ^> NH3, NH2OH, N2O Nitrogen - oxides Nitrosyl compounds (X = f, ci, Br) 2 NO + Cl2 2 NOCI covalent compound NOCl + ROH — HCl + RONO Preparation of nitrous acid esters NO+ compounds NO+HSO4-, NO+CIO4-, NO+BF4- > Cation NO+ je isoelectronic with CO, CN- N2. > Ligand in complexes, e.g. natrium nitroprusside Na2[Fe(CN)5NO] -used for qualitative evidence of Fe Nitrogen - oxides Exists only at low temperatures as a pale-blue solid (m.p. -102 °C). 0 Planar arrrangent of atoms with non-ususal long N-N distance Disproportionation in liquid state, above 3°C (b.p.) is 100% _ <+// N-N O N2O3 NO + NO2 => Equimolar mixture NO + NO2, resp. NO + O2 in correct ratio have same properties as N2O3 . Nitrogen - oxides N2O3 can be considered nitrous acid anhydride N2O3 + H2O — 2 HNO2 — decomposition In alkali media N2O3 yields nitrite N2O3 + 2 OH- — 2 NO2- + H2O Reaction with strong acids => cations NO+ N2O3 + 3 H2SO4 — 2 NO+ + H3O+ + 3 HSO4- Nitrogen - nitrous acid Nitrous acid HNO2 unstable acid, Ka = 6-10"4) easy disproportionation at room temperature 3 HNO2 HNO3 + 2 NO + H2O Preparation Ba(NO2)2 + dil. H2SO4 -> BaSO4 + 2 HNO2 Redox properties of HNO2 2 HNO2 + 2 HI 2 NO + I2 + 2 H2O HNO2 + 3 H2S NH3 + 3 S + 2 H2O HNO2 + NH2OH [H2N2O2] N2O + 2 H2O Nitrogen - nitrites Nitrites as solids are stable NaNO2 industrialy production 2 NaOH + NO + NO2 — 2 NaNO2 + H2O Oxidation by powerful oxidation agents 5 NO2- + 2 MnO4- + 6 H+ — 5 NO3- + 2 Mn2+ + 3 H2O Complex formation (nitro- and nitrito- complexes) e.g. [Co(NO2)6]3- /r R nitroalkanes N — ®\ \ R-O N P alkylnitrites Nitrogen - nitrites Important: condensation reactions with ammonia yield diazonium salts, Aniline Benzendiazonium chloride HCI, aq CfiH5NH2 + HN02-----► (Q,H5NN)C1+2H20 Diazonium salts give by following reaction with alifatic or aromatic amines azodyes. Nitrites are toxic! Nitrogen oxides NO2, resp. N2O 4 very toxic Colourless diamagnetic form N—N Brown paramagnetic form -N O/ Q -N Unusually long N-N distance Preparation 2Pb(N03) Production NO + O2 2 NO2 Nitrogen - oxides Reactivity of NO2 Disproportionate 2 no2 + h2o hno2 + hno3 follows: 3 HNO2 - HNO3 + 2 NO + H2O 2 no + o2 2 no2 Principle of nitric acid production In alkali media 2 no2 + 2 oh- no3- + no2- + h2o Decomposition at >150 °C 2 no2 2 no + h2o Decomposition is complete at 650 °C. n2o4 no+ + no3- Nitrogen - oxides Reactions with concentrated acids N2O4 + H2SO4 NO+ + HSO4- + HNO3 follows HNO3 + 2 H2SO4 NO2+ + 2 HSO4- + H3O+ Nitryl cation Application Ag + N2O4 — AgNO3 + NO KCl + N2O4 (l) — KNO3 + NOCl NaClO3 + N2O4 (l) — NaNO3 + NO2 + ClO2 Nitrogen - oxides +V N2O5 Production by careful dehydratation of HNO3 2 HNO3 + 1/2 P4O10 2 HPO3 + N2O5 NCVNCV m 0(2) (m.p. 30 °C) N2O5 is considered as nitric acid anhydride N2O5 + H2O — 2 HNO3 Powerful oxidation effect Na + N2O5 NaNO3 + NO2 I2 + N2O5 I2O5 + N2 N2O5 ionic disociation nitryl -salts N2O5 + 3 H2SO4 2 NO2+ + H3O+ + 3 HSO4-N2O5 + 2 SO3 2NO2+ + S2O72- Nitrogen - nitric acid Nitric acid, HNO3 O Production scheme > Colourless in crystalline state > As a liquid is yellow coloured > (m.p. - 41,6 °C, b.p. 84 °C) HNO3 form hydrates HNO3 • H2O and HNO3 • 3 H2O HNO3 is distributed as 68 % azeotrope (b.p. 121,9 °C) Nitrogen - nitric acid Anhydrous HNO3 NaNO3 + H2SO4 — NaHSO4 + HNO3 NaNO3 + NaHSO4 — Na2SO4 + HNO3 or by vaccum distillation of conc. HNO3 with conc. H2SO4, + Mg(ClO4)2 Dissolving NO2 in anhydrous HNO3 => fuming nitric acid used for nitrations in organic syntheses. Autoprotolysis as base Nitrogen - nitric acid Difference in behavior of diluted (< 5 %) and conc.HNO3 Zn + HNO3 (dil.) — Zn(NO3)2 + H2 Cu + 4 HNO3(conc.) — Cu(NO3)2 + 2 NO2 + 2 H2O 3 Cu + 8 HNO3(conc.) — 3Cu(NO3)2 + 2 NO + 4 H2O => reaction course depends on HNO3 concentration. Au, Pt, Rh and Ir are not soluble in HNO3 , Ag and Hg are dissolved „Aqua regia" Mixture of conc. HCl and conc. HNO3 (3 : 1, vol.) Powerful oxidation properties - reaction with Au, Pt 3 HCl + HNO3 — Cl2 + NOCI + 2 H2O Similar properties has also mixture H2SeO4 + HCl. Nitrogen - nitric acid Oxidation properties of conc. HNO3 towards non-metals: Phosphorus is oxidated to H3PO4, sulfur to H2SO4, iodine to HIO3 . Nitrates Allmetga nitrates are soluble Preparation HNO3 + KOH — KNO3 + H2O , ,. ■■■^^3 ,^,^^3 neutralization possibilities HNO3 + NH3 - NH4NO3 CaCO3 + 2 HNO3 — Ca(NO3)2 + CO2 + H2O conversion Nitrogen - nitric acid > Aqueous nitrate solutions miss oxidation properties. > Nitrates in melts are powerful oxidation agents. > All nitrates are thermally unstable, decomposition course depends on cation. 2 KNO3 2 KNO2 + O2 (only alkali metals) 2 Pb(NO3)2 —- 2 PbO + 4 NO2 + O2 (metals) 2 AgNO3 —l- 2 Ag + 2 NO2 + O2 (precious metals) NH4NO3 —N2O + 2 H2O Nitrogen - nity I haIoge nides /q 2 NO2 + F2 — 2 NG2F Ä-£ NOCI + O3 — NG2CI + O2 HNO3 + HSO3CI — NG2CI + H2SO4 Nitrogen - nitric acid Utilization of HNO3 Utilization of HNO3 Production of fertilizers, nitrites, nitro-compounds, Organic dyes, pharmaceuticals, etc. Nitrogen - orthonitric acid NaNO3 + Na2O -> Na3NO4 Nitrogen - N-Hal compounds Starting compound F Cl Br I NH2F NH2Cl NH2Br expl. NH2I expl. NHF2 NHCl2 NHBr2 expl. Very stable NCl3 NBr36NH3 expl. NI3NH3 expl. N2H4 N2F4 HN3 expl. ClN3 expl. BrN3 expl. IN3 expl. Only chloramine NH2Cl (also important intermediate in hydrazine production) has practical use as disinfectant. Nitrogen - N-Hal compounds fluorides Nil F, (11 = (i -2) | Electrolysis in melt [M11,1 IFJalso [n.i-. ) 2NF3 + 3H2 N, + 6HF -4 L\j NHnCl3.„ (n = 0 - 2) Chlorine is introduced in salmiac solution NH2C1 pH>8.5 NHC12 pH - 5 NC13 pH<4.5 Reactions N(V) halogenides NI3.NH3 „iodonitrogen" Very explosive ! Its formation have to be always considered in reactions between iodine compounds and ammonia Nitrogen - N-S compounds Single bond N-S is very stable. Sulfonamides R-SO2-NR2 Cyclamates (in sweeteners ) C6H11NH-SO3Na Binary S-N compounds Tetrasulfur tetranitrideS4N4 orange crystals (m.p. 178 °C) very explosive 6 S2Cl2 + 16 NH3 S4N4 + 12 NH4Cl + S8 Thiazenes (S=N) and thiazyl halogenides S = N - halogene Phosphorus N - P: differences in chemical reactivity ❖ P - no ttp bonds => formation of other structures of analogous compounds => other properties ❖ Covalent bonds are possible both with non-metals and metals (phosphorus electronegativity is 2.1) ❖ P — H compounds - no H- bonds ❖ Typical c.n. = 4. ❖ Unoccupied d-orbitals in P-atom => compounds with c.n.= 5 and 6. ❖ npd interaction with electronegative elements (F, O, Cl). ❖ Existence of all oxidation degrees from -III to V. ❖ Even oxidation degrees (e.g. in P2Cl4, H4P2O6) are only formal. Phosphorus - bonding Hybridization Bonding Example sp3 (with octet configuration ) PH4 +, (CH3)4P+ 3a + 1 fp PH3 , PCl3 , P4O6 sp3 (withover- octet configuration ) 4a + 277d deloc. 4a + 177d H3PO4 , P4O™ sp3d 5a PF5 , PCl5 (g) sp3d2 6a PF6- , PCl6- (s) fp - free electron pair Phosphorus - in nature > fluoroapatite Ca5F(PO4)3 > phosphorite Ca3 (PO4)2 > carbonatoapatite > hydroxyapatite Ca5(OH)(PO4)3 > Bones and tees of vertebrate > Biogene element - nucleotides Production: Phosphorite or apatite reduction using C in the presence of SiO2 (electrical furnaces at 1300 °C) ^ vapour of white phosphorus. 2 Ca3(PO4)2 + 6 SiO2 + 10 C -» 6 CaSiO3 + P4 + 10 CO white phosphorus Phosphorus - allotropy White phosphorus P4 yellowish, soft solid (m.p. 44,1 °C, b.p. 280,5 °C) ❖ Bond in P4 molecule have p-character ❖ Bonds are „bent", bond agle 0o (for pure p-bonds 90o is expected) => great tension extreme high reactivity. ❖ Bonds among P-atoms - van der Waals. Properties of white phosphorus > White phosphorus - insolubility in water (=> storing under water). > Well soluble in carbon disulfide and some polar solvents (e.g. PCl3). > Phosphorus pentaoxide is formed under burning on air. > Phosphoresence. > Toxic (lethal dose approx. 50 mg). Phosphorus - allotropy Red phosphorus Pn red solid (t. conversion at 400 °C), c.n. of P-atoms = 3 amorphous (tetragononal triclinic cubic) > Amorphous red phosphorus - can be obtained by heating from white phosphorus in inert atmosphere to 270 - 300 °C. > Splitting of P— P bonds by heating => formation of high polymers > Lower tension in bonds => lower reactivity, lower solubility > Non toxic. Black phosphorus Pn t. conversion 400 °C) amorphous (orthorhombic triclinic cubic) > Black phosphorus - can be obtained by heating together oh P with Hg-vapour at 370 °C, or by heating P4 at high pressures (200 °C, 1,2 GPa) > Layer structure, good thermal and electric conductivity. > Thermodynamically most stable modification, less reactive. Phosphorus - chemical reactions Reactivity depends on P modification. Direct reactions Phosphorus foms binary compounds with all elements, except Sb, Bi, and inert gases. P + O2 S X2 metals HNO3 OH- P4O6 P4S3-10 P4O10 PX3 PX5 phosphides H3PO4 H2, PH3, HPO2-, HPO. Phosphorus - P - H compounds Phosphane ^ Colourless, very toxic, smelling gas PH3 > (m.p.- 133,8 °C, b.p.- 87,8 °C). > Insoluble in water, absence of H-bonds > Better solubility in less polar organic solvents. Ca3P2 + 3H20-->3Ca(OH)2 + 2PH3 2AIP + 3H2S04---► A12{S04)3 + 2PH3 PH4I + KOH-----> 2PH3 + KI + H20 P4 + 3 KOH + 3 H2O 3 KH2PO2 + PH3 Phosphorus - P - H compounds Phosphane reactions PH3 + 2 O2 -- H3PO4 burning PH3 + 2 I2 + 2 H2O NaHCOq H3PO2 + 4 HI hydroformylation PH3 + 4 HCHO + HCl — [(CH2OH)4P]Cl fire retarder for textile fibers Phosphonium salts PH4+ > Thermally little stable, only PH4I is more stable > Easy hydrolysis. Phosphorus - P - H compounds Phosphane complexes |PH3 as ligand => formation of very stable complexes (back donation of metal electron density to vacant d-orbitals). =^> Hydrogen can be substituted with alkyl, aryl - significance of complexes with PR3 derivatives. Diphosphane P2H4 is analogous to hydrazine M.p. -99 °C, b.p. 63.5 °C Little stable, self-igniting, explosive Phosphorus - phosphides Phosphides - binary compounds phosphorus - metal ❖ Direct synthesis from elements. ❖ Reduction of phosphates by C in glow. ❖ Phosphides with elctropositive elements undergo to hydrolysis ❖ Transition metal phosphides have character of intermetallic inert compounds. ❖ Utilization in electrotechnics (GaP). Phosphorus - oxides Stereochemistry of phosphorus and its oxides O adamantanoid structures Phosphorus - oxides P(III) Phosphorous oxide P4O6 adamantanoid structure Reactions of P4O6 > Preparation by controlled oxidation (O:P = 3:1) of white phosphorus at 50 °C. > Fraction sublimation is used for purification (separation from P4O10). > P4O6 cannot be prepared by phosphorous acid dehydratation, but it is considered as its anhydride. P4O6 + 6 H2O — 4 H3PO3 under cooling 4 H3PO3 3 H3PO4 + PH3 disproportionaton Phosphorus - oxides P(V) White (powder) substance in few modifications Production in dry air P4 + 5 O2 — P4O10 > Extremely sensitive to moisture. > Hydrolysis in aqueous media, end-product is phosphoric acid => P4O10 is phosphoric acid anhydride. P4O10 + 2 H2O — H4P4O12 HO > H4P2O7 H2° > H3PO4 Afinity towards water is great: P4O10 removes also water, that is constituonally bonded in hydroxyl compounds => e.g. dehydratation of aqueous solutions of strong acids gives anhydrous forms or oxides C2H5OH yields ethene, RCONH2 - acid nitriles. Utilization: thermic H3PO4 production drying agent for exsiccators and various columns Phosphoric oxide P4O10 adamantanoid structure Phosphorus - oxides P(„IV") Phosphorus (IV) oxide (PO2)n Contains Pm a PV in various ratios Production n P4O6 3 (PO2) + n P in sealed tube Composition depends on a way of production ^> P4O7, P4O8 a P4O9 adamantanoid structures Hydrolysis of „mixed" oxides gives a mixture of H3PO3 and H3PO4 P4O7 + 6 H2O 3 H3PO3 + H3PO4 Phosphorus - P-sulfides and sulfide-oxides adamantanoid structures Phosphorus - P-sulfides adamantanoid structures Phosphorus - halogenides px3 px5 p2x4 PCI 3 m.p . , -93,6 °C b.p. 76,1 °C PC13 + AsF5 -> PF31 + AsCl3 PCI3 + Cl2 —» PC1; PC15 + AsF3 - - PF5- + PF6 PX5 (g) pci4- PClft PBr4' Br" 6PC15 + P4O10--- —* IOPOCI3 =^ POF3 , POI3 PBr5 + (COOH)2 - ----> POBr3 + CO + C02 + 2HBr Phosphorus - halogenides PCI/BCI B -CI3 184.0 pm B -CI4 182.8 pm < Cl - P - Cl 107.5-111.2 < Cl-B-Cl 108.4-110.7 Phosphorus - halogenides Tetraiododiphosphane P2I4 Phosphorus - P(III) halogenides P(III) halogenides are more stable than analogous N(III) halogenides. Characteristic reactions, used also in practice: solvolysis, substitution, and redox-reactions. Solvolytic reactions - examples: PX3 + 3 H2O — H3PO3 + 3 HX PX3 + 3 ROH — (RO)2PH(O) + 2 HX + RX PX3 + 3 ROH + 3 py — P(OR)3 + 3 py.HX PX3 + 3 RCOOH — H3PO3 + 3 RCOCl Phosphorus - P(III) halogenides Substitution reactions - examples: PX3 + 3 RMgX — R3P + 3 MgX2 PCl3 + 3 AgCN — P(CN)3 + 3 AgCl Redox reactions - examples: PCl3 + SO3 — POCl3 + SO2 PCl3 + S ^ PSCl3 PCl3 + Cl2 — PCl5 Phosphorus - P(V) halogenides Additive reaction (for Cl, Br, I): PCl3 + Cl2 — PCl5 Fluorides are prepared by using fluorinating agents: lllllilllB^ + SbF5 — SbCl5 + PF5 2 PCl5 + 5 CaF2 — 5 CaCl2 + 2 PF5 PCl5 in solutions and molten state forms ionic substance [PClJ+[PCl6]-, PBr5 (and probably also PI5) yield [PBrJ+Br-. PF5 reacts with ionic fluorides to complex compounds - sp3d2 PF5 + MF — M[PF6] Phosphorus - P(V) halogenides Stepwise hydrolysis i!!iHSipPCl5 + h2o — poci3 + 2 HCl POCl3 + 3 H2O — H3PO4 + 3 HCl Chlorinating agent in reactions with OH group: PCl5 + H2SO4 — POCl3 + HSO3Cl + HCl lplPCl5 + RCOOH — POCl3 + RCOCl + HCl PCl5 + ROH — POCl3 + RCl + HCl Similar reaction with some oxides: 6 PCl5 + P4O10 — 10 POCl3 |||||||||lPCl5 + SO2 — POCl3 + SOCl2 Phosphorus - P(V) oxide-halogenides POX3 > Derived from OP(OH)3 - replacing all OH groups by halogen atom 1 > Partially substitution at fluorides - stable H2PO3F or HPO2F2. > Tetradedral structure, P=O bond is evidently shorter due npd interaction. > Physical properties correspond to molecular mass: POF3 - gas, POCl3 - liquid POBr3 - crystalline substance > Mixed halogenides are also known: e.g. POX2Y. Phosphorus - P(V) oxide-halogenides POCl is of great practical significance Preparation: PCl5 + (COOH)2 — POCl3 + CO + CO2 + 2 HCl 2 PCl3 + O2 — 2 POCl3 6 PCl5 + P4O10 — 10 POCl3 > P- Cl bond is very reactive (hydrolysis, solvolysis, substitution) => POCl3 can replace PCl5 in chlorinating reactions, > POCl3 is starting compound for syntheses of many orgaphosphorus compounds that are used as insecticides, softeners, detergents, extraction agents, etc. > POCl3 can be used as polar solvent ! Phosphorus - oxoacids Some rules are valid for structures of phosphorus oxoacids: > All P- atoms are tetrahedral coordinated => c.n. = 4 > Minimally one -OH group, able to dissociate, is bonded to P- atom > More OH groups => disociation constants differ up to 5 orders of magnitude. > Some acids contain besides P-O, P-OH groups even non-dissociative P-H bonds. > Tautomerism is possible : H - P = O <-» HO - P| => reduction properties > Formation of isopolyacids is realized by folloving bonds: P-O-P, P-O-O-P or P-P > These bonds hydrolyse in adicic and alkali media. > Exception is p- p bond,, that is very stable in alkali media. Phosphorus - oxoacids - H3P02 Ka = 8,5-io-2 Hypophosphorus acid, H3P02 ph3 + 2I2 + 2H20-->H,P02 + 4HI oh i>. k )ii in () i :n P4 i 40IP + 2H20-----2IIPO32 l 2PH3 > white crystalline substance > anion [H2P02]- has tetragonal structure > Na and Ni (II) salts are used in currentless nickeling. |o|0 oh h-p9— oh =3=5: |p- h h Phosphorus - oxoacids - H3PO3 (k1 = s-iq2, k2 = 210-7) > Alkali metal phosphites are well soluble well soluble while M(II) salts are mostly insoluble. Phosphorus - oxoacids - H3P04 Orthophosphoric acid > Colourless crystaline compound (m.p. 42.3 oC). > Excelently soluble in water. > Formation of crystalline hemihydrate H3PO4. 1/2H2O. > Distributed as 85 % solution. HO O OH OH Preparation > Oxidation of P4 > Oxidation of othet phosphorus containing oxoacids by conc. HN03 > Hydrolysis of phosphorus halogenides or oxide-halogenides. Phosphorus - oxoacids - H3P04 Production apatite Ca3p04: C, SiOz CaSiOj P4 (,n> °2 » p4o10 -^H3P04 combustion "Thermic" phosphoric acid Extraction phosphoric acid 3 Ca3(P04)2.CaF2 + 10 H2S04 10 CaSO4 + 6 H3P04 + 2 HF Phosphorus - properties of H3PO4 H3PO4 (AC, = 7,1-10-3, K2 = 6,3108, K3 = 4,2-10-13). > Three salt series. > Dihydrogen phospates [H2PO4]- are slightly acidic, soluble in water > Hydrogen phosphates [HPO4]2- are slightly alkaline, only alkali metal salts are soluble in water. ^Phosphates [PO4]3- solutions are alkaline as result of hydrolysis. > Other salts are insoluble. > Phosphates are resistant towards reduction, no reaction with hydrogen in statu nascendi was observed - difference from P(I) and P(III) salts. > Reduction is possible only in glow using carbon. Phosphorus - oxoacids - condensation Three types of phosphate groups can participate in condensation reactions: 0 1 End groups -o—p—o- o 0 1 Middle groups 1 -o—p—o o 0 1 Net groups -o—p—o- 0 1 Phosphorus - isopoiyacids H3PO4 H4P2O7 0 1 O 0 1 H5P3O10 HO-P-O-1 1 1 -p-o-L 1 J -P~OH n 1 H^O^ etc. OH OH OH Phosphorus - isopolyacids H3PO4 condensation 220 CC I H4P207 m.p. 61 °C diphosphates M1 [>i): dihydrogen diphosphates M'.H.p.O.. _ f J _Li -£j -£j t 2Na2IIP04 2NaII2P04 Phosphorus - isopolyanions f-^ catena - triphosphate 06- P1,P2 158.6, 162.9 pm 09- P1,P3 160.0, 162.6 pm P = 0 148.8 - 153.2 pm Phosphorus - isopolyanions Metaphosphoric acid (HPO3)x cyclic and linear metaphosphates (MPO3)x xKH2P04---► (KP03)x + xH20 x/2K2H2P207 ---► (KP03)x + x/2H20 (NaP03)x (x = 20a4 500) 3Na4P20 7 + 6NH4C1----> 2(NaP03)3 + 6NaCl + 6NH ^ + 6H20 Cyc/o-hexaphosphate P = 0 148.0 - 149.2 pm P-O-P 159.3 - 160.1 pm P1 -01, 02 147.1, 148.2 pm P1 -04 161.3 pm P2 -03 148.3 pm P2 -04 160.0 pm H3PO4 is used for surface treatment of metals (phosphatization) ^Phosphate as fertitizers (superphosphate, ammonium phosphate) >Phosphoric acid esters (r0)3p0 - extraction agents in actinoide chemistry >Diphospates, triphospates (and commonly also all polyphosphates) are used as detergents (formation of soluble calcium and magnesium complexes). H4P2O6 p4 + o: H4pa + H20 P - 01 155.5 pm P - 02,03 150.3, 150.9 pm P - P 217.8 pm 04 - H 93.5 - 102.8 pm 03...H4 151.1 pm H4P206.2H20 (HO)2OP-PO(OH)2 (HO)(H)OP-0-PO(OH)2 Isodiphosphoric acid H,P03 + H3pq4 P - P bond is very stable, even in 60% KOH under boiling Phosphorus P - N compounds f-\ Phosphazenes V_) = P=N- /-\ Linear: (C6H5),PCI2 + (QH5)NH2----> (QHS)3P=N(QHS) + 2HC1 3PC15 + NH4C1---1 [C13P=N-PC13]+PC16- I 4HC1 f-N Cyclic: nPCl5 + nNH4Cl----> (PNCl2)n + 4nHCl Phosphorus - P - N Compounds Hexafluoro-cyc/o-triphosphazene Cyclo- phosphazenes HexaChloro-cyc/o-triphosphazene 0(1) VIth group PSE, ns2np4 Oxygen, sulfur, selenium, tellurium, polonium ❖ O and S are non-metals (forming covalent bonds), Se, Te semi-metals, Po is typical metal ❖ O - 2nd most electropositive element => creation of octet configuration ❖ formation of np bonds at oxygen ❖ S and other elements - vacant d-orbitals, npd interaction with electronegative elements (O, N) => formation of compounds with c.n. = 5 (trigonal bipyramide) and c.n. = 6 (octahedron) ❖ Oxidation degree VI+ is most stable at S, its stability decreases with atomic number increasing => increase of oxidation properties ❖ stability of IV+ increases with increasing atomic number ❖ Po - typical is II+ Properties of VIth group elements O S Se Te Po Atomic number 8 16 34 52 84 Ar 15,9994 32,06 78,96 127,60 (209) Density / g.cm-3 1,30 2,06 4,82 6,25 9,19 m.p. / oC -218,8 119 217 452 246-54 b.p. / oC -182,97 444,6 685 990 962 Coval. radius / pm 73 104 117 137 164 loniz. energy / eV I1 13,6 10,4 9,75 9,00 8,43 I2 35,1 23,4 21,3 138,1 88,0 Electronegativity 3,50 2,44 2,48 2,01 1,76 Oxygen In nature > most abundant biogene element (45.5 % in hydro-, litho- a atmosphere) > in Earth's core (cca 21 % vol.) - two allotropic forms O2, O3, > in compounds (water, oxides, and other oxygen-containing substances) > on the Moon (44.6 %) > natural oxygen is mixture of 16O, 17O (0,04 %) a 18O (0,2 %) isotopes (17O used in NMR spectroscopy, 18O in IR spectroscopy) Formation by photosynthesis chlorophyll H2O + CO2 + hv "enzymes^ O2 + CH2O (sacharides) Dioxygen - molecule m.p. -218.8 °C b.p. -182.97 °C hu 3(X + 'B 02 + JB B = organic dye > in (s) three crystal modifications > in (l) and (s) - pale-blue > restrictly soluble in water > paramagnetic molecule (two non-pair electrons - triplet oxygen) > singlet oxygen (all electrons in pairs) Singlet oxygen can be formed in higher atmosphere layers. Properties: > very reactive - powerful oxidation agent > many direct reactions with elements. Exception are halogenes, noble gases and some precious metals > reactions are usually exothermic (burning) Oxygen - bonding > Covalent bonds with both metals and non-metals. > As a rule O is electronegative part of a molecule. > In O2F2 is oxygen electropositive element. > O2+ - dioxygenyl cation exists only with anions BF4- , PtF6- , PF6- > Oxygen can create 4 bonds. Sometimes with additive 71-bonding. > C.n. in complexes is up to 8 (in oxides M2O with antifluorite structure). > Dioxygen can serve as ligand: Oxygen - bonding Hybridization Bonding Examples ionic K2O, BaO sp3 4a ZnO, Al2O3, Be4O(CH3COO)6 3(7 + 1 fp H3O+, [Cu(H2O)4J2+ 2a + 2 fp H2O, Cl2O, R2O 1a + 3 fp, event. 1a + 2 fp + 1n deloc. R3PO, R2SO sp2 2a + 1fp + 1n deloc. O3 (central atom) 1a + 2fp + 1n ketones sp 1a + 1fp + 2n CO, NO+ fp = free electron pair Oxygen - preparation and production Preparation: > electrolysis of water > thermal decomposition of oxides, peroxides and some salts ^ 2HgO----► 2Hg + 02 2H202 -----► 2H20 + 02 Mn02 2NaC10,------2NaCl + 302 220 °C KMn04--------► K2Mn04 + Mn02 + 02 > reaction of higher oxides with conc. H2SO4 2Mn02 + 2H2S04-----► 2MnS04 + 2H20 + 02 4Cr03 + 6H2S04-----► 2Cr2(S04)3 + 6H20 + 30 Production: fraction distillation of liquidified air b.p. (N2) = -196 °C, b.p. (O2) = -183 °C) Oxygen - ozone > Very reactive, in higher concentrations toxic. > Powerful oxidation effect in gaseous form and in solutions. > More stable in acidic solutions. (-< Determination 03 + 2KI + H20----► 02 + I2 + 2KOH 2Na2S203 + I2----► Na2S406 + 2NaI Preparation : ❖ by electrical discharge in oxygen atmosphere ❖ thermal decomposition of peroxoiodic acid Ozone - properties CN- + O3 — OCN- + O2 H PbS + 4 O3 — PbSO4 + 4 O2 3 I- + O3 + 2 H+ — [I3]- + O2 + H2O Reaction with dry powder-forming hydroxides == ozonides, e.g. KO3 5 O3 + 2 KOH 10 C > 2 KO3 + 5 O2 + H2O Utilization: > sterilization of drinking water > air cleaning > oil and starch whitening Oxygen compounds - oxides Sorting: according to bond type ionic f double s_ - Pcrowskit AB03 ( A b) Ilmcnit ABO3 (ab) Spinel AB.O, covalent polymeric chains layers molecular HgO, (Se02)n, . SnO, M0O3,.... CO,, OsCX acidic according to character B203,S03,.... » B203 + 3H20 2H3BO3 basic Na20, CaO: >- CaO + H,0 Ca(OH)2 > CuO + 2H,SO, ■+ CuS04 + H20 amphoters ZnO + 2H30+ > ZnO + 20H + H20 - ZnO, A1203, .... > Zn2+ + 3H20 — [Zn(OH)4]: indifferent CO, N20, NO, Oxygen compounds- oxides Common preparatiom methods: a) Direct synthesis from elements b) Decomposition of hydroxides Cu(OH)2 CuO + H2O c) Decomposition of salts : CaCO3 CaO + CO2 2 Pb(NO3)2 - 2 PbO + 4 NO2 + O2 d) Reaction of elements with water vapour: C + H2O CO + H2 3 Fe + 4 H2O Fe3O4 + 4 H2 e) Oxidation of elements by oxidation agents f) Thermal decomposition or reduction of higher oxides. Oxygen compounds - water H-bonds => high m.p. and b.p. high value of phase energies • thermally stable • universal solvent • H—O bond, very polar and stable (464 kJ mol1) Structure of ice Ih 100 pm 176 pm • 9 cryst. ice modifications • hexagonal ice has „empty" structure => formation of clathrates Ar-5,75 H2O, Cl27,25 H2O, CHCl317H2O • anomaly of water Oxygen compounds - water > Properties of water are infuenced by quantity of dissolved substances. > Water have to be processed, according to purpuse of use. > Drinking water contains mostly chlorides, sulfates, and Ca and Mg hydrogencarbonates =^>water hardeness and dissolved gases (CO2, NH3, H2S, SO2). Drink water is produced in waterworks: > precipitation od coloids on Fe(OH)3, event. Al(OH)3 surface > filtration > softening using ion exchangers > disinfection using chlorine or ozone Water for laboratory use: distillation or deionization Oxygen compounds - water > water in complex cations e.g. [Be(H2O)4]2+, [Co(H2O)6]2+, [Cr(H2O)6]3+. > crystal water forms salt hydrates LiCl-H2O, KF-2H2O > in [Cu(H2O)4]SO4-H2O one water molecule is through 2 H-bridges bonded to two [SO4]2- anions. > Permitivity 8 = 78 => solvation of both cations and anions. > Excellent solvent for many ionic compounds > Water is miscible with a lot of organic solvents (alcohols, acetone, carboxyl acids, dioxane, tetrahydrofurane, dimethyl formamide, dimethyl sulfoxide, hexamethylphosphor triamide > Many compounds are potential electrolytes => dissotiation in aqueous solutions, e.g. HCl, H2SO4, BF3 . Oxygen compounds - water Hydrolytic reactions: P4O10 + 6 H2O — 4 H3PO4 SiCl4 + n H2O — SiO2.aq + 4 HCl AlCl3 + n H2O — [Al(H2O)6]3+.aq + 3 Cl- (aq) Autoprotolysis of water: , 2 * H3O+ + OH- Dissociation constant of water K = [H+][OH]/[H2O]2 = 1,8-10-16. this low value means that salts of even weaker acids can be hydrolysed => alcoholates, amides, ionic hydrides, nitrides, phosphides, silicides, and borides yield hydroxides and corresponding hydrides (alkohols, NH3, PH3, silanes, boranes). Oxygen compounds - water Deuterium oxide D2O (heavy water) > D2O can be found in common water in low concentration. > production of D2O is based on longterm electrolysis of water => light water is electrolysed more rapidly and D2O remains in electrolyte. >D2O is used: in nuclear technology (cooling medium, neutron moderator), solvent for NMR spectroscopy. > chemically, there in no difference between H2O and D2O (only reactios in D2O media are slower - isotopic effect). > lower permitivity of D2O leads to lower solubility of salts. > autoprotolytic constant is also lower in D2O > pronounced difference is observed in physical constants. Oxygen compounds - water Properties of H2O, D2O a T2O H2O D2O T2O 18,015 20,028 22,032 M.p. / °C 0,00 3,81 4,48 B.p. / °C 100,00 101,42 101,51 Density at 25 oC, / g cm-3) 0.99701 1.1044 1.2138 Max. density / g cm-3 1,000 1,1059 1,2150 Permitivity 78,39 78,06 - Disociation constant 1,82110-16 3,5410-17 1,110-17 Ionic product 1,008 10-14 1,9510-15 610-16 Oxygen compounds - hydrogen peroxide First way of H2O2 production: BaO2 + H2SO4 BaSO4 + H2O2 Today s production ^> anodic oxidation of sulfuric acid 2 HS04" -> H2S208 + 2 e- and following hydrolysis. H2S2Os + 2H20____> H202 + 2H2SO, H2O2 is then distilled in vacuum. Oxygen compounds - hydrogen peroxide > Bond angle H—O—O approx. 96,9 o, > Diedric angle of both -OH planes is 93,6 o, > These value differ in crystalline and gaseous state. Oxygen compounds - hydrogen peroxide > Hydrogen peroxide is stronger acid as compared with water (Ka = 1,78-10-12) => formation of peroxides and hydrogenperoxides. > Peroxides and hydrogenperoxides hydrolyse and elemental oxygen is released. > Boiling leads to complete decomposition, also in alkali media. > Decomposition is catalyzed by metal ions. > H2O2 as oxidation agent Mir IIXX 2 MntX 211,0 > H2O2 as reduction agent 2Mn04 + 5H202 + 6H30 2Mn2+ + 507+ 14H,0 Permanganatometric determination of H2O2 Oxygen compounds - hydrogen peroxide Hydrogenperoxides are known only at alkali metals. NaHO21/2H2O. Peroxides are well studied at alkali metals and alkali earth elements. Na2O2 a BaO2: 2 Na + O2 Na2O2 2 BaO + O2 2 BaO2 All peroxides contain —O—O— bond. Hyperoxides, containg paramagnetic anion O2- , are known with heavier alkali metal ions. Orange to brownish coloured. Preparation: direct synthesis. Hydrolysis in water: 2 O2- + 2 H2O 2 OH- + H2O2 + O2 Utilization: 4 KO2 + 2 CO2 2 K2CO3 + 3 O2 (oxygen recovery in breathing apparatus) Oxygen compounds - hydrogen peroxide H202 production 2-cthylantrachinol OH 2-cthylantrachinon 0 o2 + H202 0 1 Substrate recovery Et H2, Pd OH OH H2 + 02 + H202 Oxygen compounds - hydrogen peroxide Production and storing conditions: > H2O2 is distributed in 30% concentration > 30% H2O2 is can be get by vacuum distillation of aqueous phase > higher H2O2 concentrations are very dangerous and can be get by water evaporation > decomposition of H2O2 occurs in the presence of some metal ions, MnO2, dust, etc. => often explosion :::::::::::::::::::::::::::::::::::::::::::::::::::::::::: > storing is possible in PE bottles > H2O2 is stabilized by addition of H3PO4, H2SO4, urea, acetanilide, etc. Oxygen compounds - hydrogen peroxide Utilization: > whitening of textils, paper, straw, leather, > production of other whitening agents (peroxoborates, peroxocarbonates) > disinfectant > epoxide production > Na2O2 for analytical purposes, alkali oxidation melting Sulfur Occurrence • elemental sulfur S8 occurs in nature • in minerals: sulfates (Gypsum CaSO42H2O, Baryte BaSO4 etc.) sulfides (Sphalerite ZnS, Galena PbS, Pyrite FeS2 etc.) in the atmosphere H2S, SO2 a part of essential amino acids (Cytidine, Cysteine and Methionine) • natural sulfur is a mixture of isotopes 32S, 33S, 34S, 36S sulfur Formula: S Hardness: 1.5 to 2 Streak: white, sometimes pale Color: Yellow, yellow as honey, yellow-brown, yllow-green Transparency: transparent Gloss: the diamond crystal surfaces, the fracture surfaces matt The ability to split: weak Refraction: unequal, marl Crystal system: orthorhombic Occurrence: Vígľašská Huťa, Dubník, Smolník Associated minerals: calcite, aragonite, Celestine Similar minerals: yellow sphalerite Tests: Sulfur melts at low temperature and excludes SO2. Usage: production of H2SO4, chemicals, explosives, usage in paper industry, in rubber industry, the manufacture of matches, using against pests Interests: In the past it was used for producing gunpowder. PYRITE Formula: F©S2 Hardness: 6-6,5 Streak: green-black Colour: yellow, Transparency: opaque Štiepateľnosť: imperfect Refraction: Marl, uneven Crystal system: cubic Occurrency: the most abundant sulfide mineral - Hnúšťa, Banská Štiavnica, Smolník, Zlatá Baňa, ... Accompanying mineral: sphalerite, galena, silica, kalcoit Similar minerals: markazite (different crystal shape, streak - more green) Tests: to impact with hard metal objects spark, to melt relatively easy Usage: production of H2SO4 and polishing powders, sometimes as a source of Co, Cu, Au, Se, ... , bound in the pyrite ore , rarely separate and less perfect crystals are processed as a precious stone Interests: for its color and similarity to chalkopyrite is also called „Caťs gold" - "Fool's Gold" • GALENA Formula: PbS Hardness: 2,5 Streak: grea-black, bright blue Color: light or dark lead-lead, the open fracture with bluish tint Transparency: opaque Gloss: metal Ability to split: very good Refraction: half Marl Crystal system: cubic Occurrency: Banská Štiavnica, Zlatá Baňa, pri Ochtinej Accompanying mineral: sphalerite, chalkopyrite, pyrite, baryte, Silver sulfide Similar minerals: Vzhľadom na farbu, lesk, dokonalú štiepateľnosť je galenit nezameniteľný. Tests: dissolve in HCl and produce H2S smelly gas. Use : the main lead ore - lead^^^^^Hf • CHALKOPYRITE Formula: CuFeS2 Hardness: 3,5 - 4 Streak: green-black Coloc: gold-yellow (sometimes geenish) Transparency: opaque Gloss: metalic Cleavage: not good Refraction: lasturnatý, nerovný Crystal system: tetragonal Occurrence: Smolník, Gelnica, Slovinky, Rožňava, Zlatá Baňa, Banská Štiavnica, Hodruša Accompanying mineral: pyrite, sfalerit, kalcit, fluorit, tetraedrit Similar materials: baryte, dolomit, silica Tests: dissolve in HNO3, firecolor- green Use: the most important cupper ore, use in the electrical industry, the chemical industry and as a precious stone Sulfur - production Elemental sulfur is often obtained from sulfur deposits (the main sources are in U.S.A,former USSR, Canada, Poland, Japan) Frash process Sulfur is obtained from bedrock using overheated water which melts and extrudes the molten sulfur to the surface. This sulfur is very pure, purity is over 99.5 %. Other ways of sulfur production > Oxidation of hydrogen sulfide from natural gas, > From sulfur compounds present in petroleum. Sulfur - bonding > The formation of S2- anion is difficult =^>only sulfides of electropositive metals with low ionization energy (alkali metals) are known. > The reason is low electronegativity of sulfur (only 2.4) and negative electronaffinity (for the transition S — S2- , 3.4 eV). Therefore, sulfur readily forms covalent bonds. > Sulfur has got free 3d- orbitals. Because of presence a- bonds sulfur can form six bonds. a-bonds can be created by using sulfur p-orbitals,or more frequently hybrid sp2, sp3, and sp3d sp3d2 orbitals. > Sulfur is able to create also npd interactions with highly electronegative elements (such as F, O, Cl) due to the presence of 3d- free orbitals. These n-bonds are delocalized and usually are shorter than sum of covalent radii. > Energy of S-S bond is quite high (264 kJ mol-1), therefore there are a number of compounds with this type of bonding. > Due to the low electronegativity of sulfur, hydrogen bonds are not too typical. Sulfur - bonding Type of hybridization Type of bonds Examples sp3 ionic K2S, CaS (cryst.) 4a ZnS (cryst.) 3a + 1 fp R3S+ 2a + 2 fp S8 1a + 3 fp S22- 4a + 2nd deloc. SO42_, H2SO4, (SO3)3 3a + 2nd deloc. + 1 fp SO32- 3a + 1 nd + 1 fp SOCl2 p3 3a + 1 fp H3S+ 2a + 2 fp H2S 1a + 3 fp SH- sp2 3a + 3nd deloc SO3 gas 2a + 2nd deloc + 1 fp SO2 sp3d 4a + 1 fp SF4, SCl4 sp3d2 6a fp - free electron pair SF6 Sulfur - molecule Sulfur forms few allotropic modifications. > The only stable sulfur modification is orthorhombic sulfur Sa, stable at normal pressure to a temperature of 95.3 °C. ^ Sp is monoclinic, it is stable in the range 95,3 - 119 oC. > Both modification are created by cyclo the difference is observed in the arrangement of molecules in the crystal lattice. > Angles S—S—S are 107.8o, hybridization sp3 (sulfur atoms are regullary above and under the plane of cycle S8 Sulfur - properties m. p. = 119 °C b. p. = 444.6 °C Behavior of S in the process of heating > 119-161 oC - sulfur = yellow liquid > > 161 oC - reactivity and viscosity increase - splitting of cycles and creating biradicals • S-(S)6-S • ( Sn) > These radicals are bonded together, they create long chains with higher viscosity. Rapid cooling of the melt (pouring water) lead to the formation of plastic sulfur SM, this sulfur type consists of these very long chains. > SM is not stable, it spontaneously transfers to the Sa . > Other sulfur modifications exist, too, e.g. Sp, S6, S7, S10, S12, S18 _ > 900 oC paramagnetic sulfur S2 is formed. Sulfur- properties and reactivity Sa is soluble in CS2 (very well), in CCl4 (worse), in benzene (very badly), in alcohols and water (not soluble) Reactivity: > good, especially at high temperatures (radicals) > reactions with almost all elements (except noble gases, nitrogen, tellurium, iodine, platinum, iridium and gold) > sulfur creates sulfides with metals; ZnS and HgS are formed at room temperature => removal of spilt mercury Sulfur - compounds - sulfane Sulfane is colorless gas (melting point -85,6 oC, boiling point -60,3 oC) > diluted sulfane smells of bad eggs > concentrated sulfane smells nice a hen is very toxic (more than HCN !). Preparation: H2 + S j- H2S AH = -20 kJ.mol-1 > decomposition of some sulfides by non-oxidizing strong acids FeS + 2 HCl — FeCl2 + H2S Sulfur - compounds - sulfane Sulfane - reducing properties (free electron pairs), is mostly oxidized into elemental sulfur, burns in oxygen to SO2 h2s + Cl2 — 2 HCl + S H2S + H2O2 — 2 H2O + S H2S + 2 FeCl3 — 2 FeCl2 + 2 HCl + S H2S + conc. H2SO4 — S + SO2 + 2 H2O 2HN03 + 3H2S----► 2NO + 3S + 4H20 Sulfur - compounds - sulfane Sulfane is soluble in water - 0,1 m solution > creates two series of salts (hydrogensulfides and sulfides). > hydrogensulfides are generally soluble in water, known only from alkali metals and earths. sulfides are known almost for all metals, only sulfides of alkali metals, alkaline earth'elements are soluble. ammonium sulfide is known only in solution. As a result of hydrolysis -salts react alkaline. S2- + H2O SH- + OH- Sulfur - compounds - sulfides > heavy metals create mostly insoluble sulfides, > sulfides prepared by precipitation are colored- sulfides are mostly dark > some trivalent metal sulfides are easily hydrolyzed (Al2S3, Cr2S3, Ln2S3) > some sulfides ca be precipitated also in acidic media, e.g. PbS, Ag2S, HgS, CdS, CuS, As2S3, SnS2, > some sulfides precipitate in alkaline media, e.g. FeS, MnS, CoS, NiS CuSO4 + H2S CuS + H2SO4 MnSO4 + (NH4)2S MnS + (NH4)2SO4 Some sulfides react with alkaline sulfides and create thio-salts. SnS2 + (NH4)2S (NH4)2SnS3 Sulfur - compounds - polysulfides Melting of alkali metal sulfides or salts with sulfur lead to the formation of polysulfides (yellow coloured) Example: Na2Sn (n = 2 - 6). • ionic character, used in tanneries Reaction with solutions of strong acids leads to desulfuration Na2S4 + 2HC1----> 2NaCl + H2S + 3/8SK Sulfur - compounds - polysulfanes Polysulfanes H2Sn oily yellow liquids > Formation by acidification of aqueous solutions of alkali polysulfides using non-oxidizing acids at low temperatures >Decomposition at higher temperature is accompanied by desulfuration H2Sn — H2S + (n-1) S ^Other method for polysulfane preparation SnCl2 + 2 H2S — 2 HCl + H2Sn+2 > Acidification leads to the elimination of sulfur p*1 pk2 H2S H2S2 H2S3 H2S4 H2S5 6,83 5,0 4,2 3,8 3,5 9,7 7,5 6,3 5,7 14 Sulfur compounds - oxides Unstable SnO, SnO2 (n = 5-10), S2O, SO , peroxide SO4. Without practical application. s Q0 S .0 S .0 sulfur - compounds - sulfur dioxide b.t. -72.5 °C b.v. -10 °C S (s) + 02 (g)----► S02 (g) AH° = -296.8 kJ moľ1 pyrite 4l-'oS, ]](>: -^:0;. SSO: Preparation: a) Reduction of H2SO4 S + 2 H2SO4 3 SO2 + 2 H2O Cu + 2 H2SO4 CuSO4 + SO2 + H2O b) Reaction of sulfites with strong acids Na2SO3 + H2SO4 Na2SO4 + SO2 + H2O Sulfur compounds - sulfur dioxide >SO2 form complexes with transition metals in low oxidation states. > SO2 as reducing agent SO2 + Cl2 SO2Cl2 SO2 + Cl2 + 2 H2O H2SO4 + 2 HCl SO2 + NaOCl + H2O H2SO4 + NaCl 2 SO2 + O2 V2°5 > 2 SO3 HNO, + 3S02 + 2H20----> 2NO + 3H2SO Sulfur compounds - sulfur dioxide SO, as reducing agent SO2 + H2 —^ S + 2 H2O SO2 + 4 HI(g) S + 2 I2 + 2 H2O SO2 + H2S 3 S + 2 H2O 2 SO2 + 2 Na —^ Na2S2O4 Sulfur - compounds - sulfur dioxide as a solvent > liquid SO2 - aprotic solvent for PCl3, CS2, SOX2, Br2, amines, R-OH, iodides.... Some reactions are possible: WCl6 + SO2 — WOCl4 + SOCl2 Solubility in water > 3900 cm3 SO2 in n 100 cm3 at 20 oC > formation of SO2xH2O......SO26H2O Sulfur compounds - sulfur dioxide utilization ❖ Production H2SO4, SO3 2- , bleaching agents, fruit preservation, etc. Environmental hazard acid rains" Sulfur compounds - sulfur trioxide 137 pm -o 142 pm monomer 143 pm cyclic trimer Y-form V 66'J: S = 0 135.9- 144.3 pm S-O 159.3 - 163.8 pm polymeric SO3 a- and ß forms 02 J* S-01 163.2 pm S -02,3 140.6, 141.4 pm Polysulfuric acids Sulfur compounds - sulfur trioxide Preparation: Fe2(SO4)3 Fe2O3 + 3 SO3 2 H2SO4 + P4O10 — 2 HPO3 + 2 SO3 K2S2O7 —— K2SO4 + SO3 H2S2O7 —— H2SO4 + SO3 Catalytic oxidation of SO2 2 SO2 + O2 —— 2 SO3 AH = -195,8 kJ mol-1 Sulfur compounds - hydrogensulfite and their properties Preparation: NaOH + SO2 — NaHSO3 CaCO3 + 2 SO2 + H2O — Ca(HSO3)2 + CO2 > Hydrogensulfites are not thermally stable: 2 NaHSO3 Na2S2O5 + H2O CaCO3 + 2 SO2 + H2O —Ca(HSO3)2 + CO2 sulfur - compounds - sulfite and their properties Preparation - neutralization of hydrogensulfite using hydroxide > Alkali salts are soluble > Salts of Me" metals are not soluble > Oxidating agent: Na2SO3 + Br2 + H2O — Na2SO4 + 2 HBr > Sulfites are not toot stable at higher temperatures: 4 K2SO3 3 K2SO4 + K2S CaSO3 CaO + SO2 Sulfur compounds - hydrogensulfites Tautomerism Sulfur compounds - sulfuric acid > H2SO4 colorless oily liquid, miscible with water in all ratios. > Form two salt series: hydrogensulfates and sulfates Preparation: SO3 + H2O — H2SO4 ah = -130 kJ mol-1 Sulfur compounds - sulfuric acid Oxidation and dehydration S + 2 H2SO4 1 -> 3 SO2 + 2 H2O C + 2 H2SO4 1 -> CO2 + 2 SO2 + 2 H2O 2 Ag + 2 H2SO4 Ag2SO4 + SO2 + 2 H2O 2 HBr(g) + H2SO4 — Br2 + SO2 + 2 H2O 8 HI(g) + H2SO4 — H2S + 4 I2 + 4 H2O Sulfur compounds - sulfuric acid Sulfuric acid as a solvent Autoprotolysis: 2 H2SO4 - — H3SO4+ + HSO4- sulfatacidium Other reactions: 2 H2SO4 = H3O+ + HS2O7- H2S2O7- + H2SO4 =^ H3SO4+ + HS2O7" sulfatacidium sulfur - compounds - hydrogensulfafes MIHSO4 > mostly soluble > crystallize from alkali salts > not thermally stable : 2 KHSO4 K2S2O7 + H2O Sulfur compounds - sulfates > SO42- - are formed with almost all metals > well soluble are alkali metal sulfates > bad soluble SO42- of alkali metals, PbSO4, partially soluble Ag2SO4. Preparation H2SO4 + 2 KOH — K2SO4 + 2 H2O Zn + dil.H2SO4 — ZnSO4 + H2 Hg + conc.H2SO4 — HgSO4 + SO2 + 2 H2O BaCl2 + H2SO4 — BaSO4 I + 2 HCl ZnCO3 + H2SO4 — ZnSO4 + CO2 + H2O Na2SO3 + H2O2 — Na2SO4 + H2O Sulfur compounds - sulfates Utilization: (NH4)2SO4 - fertilizer Na2SO410H2O (Glauber salt) Production of Na2CO3 Vitriols: MIISO4nH2O (M = Zn, Fe, Co, Mn n = 7), M = Cu, Mn, Cr; n = 5), Alums: MWm(SO4)2-12H2O (MI = Na, K, NH4, Rb, Cs aj.; Mm = Al, Cr, Fe, Mn, Ti, V aj.) Plaster CaSO41/2H2O Baryte BaS04 (X-ray examinations of digestive tract) Sulfur compounds - polysulfuric acids and their salts S-O 159.2, 150.7, 166.7, 182.3 pm S = () 139.8 -142.9 pm Use: > Halogenation reactions > Preparation of sulfonamide Sulfur compounds - halogenosulfuric acid and their salts Lithium fluorosulfate S-F 155.6 pm S - Ol 145.6 pin S - 02 140.1 pm Li ... Ol 204.6 pm Li ... 02 190.4 pm Li ... Li 296.7 pm Sulfur compounds - peroxo acids H2S208 peroxo disulfuric acid o ,0 o 10 — 00 00 _ H-O-S—O-O-S-O-H Olo Production: I Olo (m.p. 65 oC). S-Ol 164.6 pm S = Q 142.1 -142.9 pm <01a-01-S 106.2° 0(3> K2S208 K hygroscopic crystalline substance preparation: by oxidation of sulfuric acid salts well soluble important salts are K2S2O8 and (NH4)2S2O8, (oxidation agents) Sulfur compounds - peroxo acids Peroxo sulfuric acid - intermediate in the preparation process of hydrogen peroxid by hydrolysis. H2S2O8 h2o > h2SO5 + H2SO4 h2o > 2 H2SO4 + H2O2 Peroxosulfuric acid Caroo acid O b.t45*C HSO3CI + H202 ----► H2s05 + HC1 S03(OOII) H — + + -O— S—o- 0|q O H Sulfur compounds - oxo acids with S-S bonds H2S2O6 Do not exist without water O 0 O|G — 00 H-O-S-S-O-H O0 |O0 H2S2O6 — H2SO4 + SO2 Salts of H2S2O6 are dithionanes sulfur - compounds - oxo acides with S-S bonds Dithionane M2S2O6 Not too good oxidation agents in aqueous media MnO2 + 2 SO2 + 2 H2O — MnS2O6 + 2 H2O Fe2O3 + 3 SO2 + 3 H2O — [Fe2(SO3)3] — FeSO3 + FeS2O6 + 3 H2O Potassium dithionane SI - 0 144.6 pm S1 - SI a 215.3 pm 01 K2 Without use Sulfur compounds - oxo acids with S-S bonds Thiosulfuric acid H2S2O3 — <+<+ H-O-S-S-H O 0 Free acids is not stable. H2S2O3 — H2S + SO3 Preparation: H2S + SO3 — H2S2O3 HSO3Cl + H2S — H2S2O3 + HCl Na2S2O3 + 2 HCl — H2S2O3 + 2 NaCl Sulfur compounds - oxo acids with S-S bonds Thiosulfuric acide - reductive properties. „anti-chlorine" Na,S203 + 4C\2 + 5H20 2NaHS04 + 8HC1 iodometry fixer in photography 2Na2S203 + I2----► Na2S406 + 2NaI 2Na,S,03 + AgBr----► Na3[Ag(S203)2] + NaBr HO- HO- O S- O S S H I H I S O O O HO-S-O S S HO-S-O O 2 HI Ssulfur compounds - oxo acids with S-S bonds Polythionic acids H2SnO6 ; n = 3 - 12 > They are produced by reaction between SO2 and H2S (aq) -Wackenroder solution. > The system of parallel and consecutive reactions is very complex, the reaction mixture contains sulfates, sulfites, thiosulfites a mixture of polythionic acids up to n = 6 The only significant acid is thiosulfuric acid H2S2O3. Preparation 2 Na2S2O3 + 4 H2O2 — Na2S3O6 + Na2SO4 + 4 H2O SCl2 + 2 [HSO3]- — [O3S-S-SO3]2- + 2 HCl S2Cl2 + 2 [HSO3]- — [O3S-S-S-SO3]2- + 2 HCl SCl2 + 2 [HSO3]- — [O3S-(S)3-SO3]2- + 2 HCl Sulfur compound s - sulfur halogenides Non-existence of Iodide) Fluoride Chloride Bromide SSF2 (b.p. -10,6 oC) SnCl2 orange (l) SnBr2 red (l) S2F2 (b.p. 15 oC) S2Cl2 yellow (b.p. 138 oC) S2Br2 red (b.p. 54 oC) SF4 (b.p. -38 oC) SCl2 red (b.p. 59 oC) SF6 (subl. -64 oC) SCl4 cryst. decomp. at 31 oC S2F10 (b.p. 30 oC) Sulfur halides are covalent compounds, S-X bonds are polar => great reactivity (except SF6). Sulfur compounds - sulfur halogenides SF4 - very reactive gas 3 SCl2 + 4 NaF CH CN 75 oC S2Cl2 + SF4 + 4NaCl -3—:-> Easily hydrolyses into HF and SO2, Selective fluorination agent: >C=O >CF2, —COOH —CF3 =P(O)OH, P=O, PF2 =PF3 I2O5 -» IF7 Sulfur - compounds - halides S2^2 - dichlor disulfane, yellow, smelly liquid (b. p. 138 oC). > It hydrolyses in water yielding HCl, H2S, S, SO2, H2SO4, and polythionic acids... > Used for preparation of CS2 and as a solvent of sulfur, used in rubber vulcanization process of rubber. SCl2 dichlorsulfane, red coloured liquid > produced by chlorination at room temperature > not stable, undergo to hydrolysis > addition on ethylene CH2=CH2 + SCl2 S(CH2CH2Cl)2 yperite (mustard gas) (blistering warfare agent). Sulfur compounds - acid halogenides Sulfurous acid halides - thionyl halides SOF2 gas, boiling point -44 oC SOClF gas, boiling point 12 oC SOCl2 liquid, boiling point 76 oC SOBr2 red-yellow liquid, boiling point 140 oC The most important substance is thionyl chloride SOCl2> highly reactive liquid with pungent odor. 502 + PCl5 — SOCl2 + POCl3 503 + SCl2 — SOCl2 + SO2 The practical application - reactions with hydroxyl compounds: H2O + SOCl2 — SO2 + 2 HCl ROH + SOCl2 — SO2 + RCl + HCl RCOOH + SOCl2 — SO2 + RCOCl + HCl SOCl2 is used as dehydrating agent in inorganic chemistry. It is used also as a non-aqueous ionizing solvent (the same as liquid SO2). Sulfur compounds - acid halogenides Sulfuric acid halides - sulfuryl halogenides SO2F2 (gas, boiling point -55 oC) SO2Cl2 (liquid, boiling point 69 oC). SO2FCl, SO2FBr and SQ2ClBr exists stoo. camphor SO2 + Cl2 ~^ SO2Cl2 2 HSO3Cl —— H2SO4 + SO2Cl2 Only SO2Cl2 is practically used in organic chemistry. Substitution -OH group by Cl or -SO2Cl. Hydrolysis leads to H2SO4 , HCl, ammonolysis yields sulfuryl diamide SO2(NH2)2. Disulfuryl fluoride S - 01 161.1 pm SO 139.5,140.2 pm S-F 152.1pm < 02 - S - 03 125.4° 4 S - Ol - S' 123.4° Sulfur compounds with S -N bonds (SN).. i« Sulfur-compounds with S- Nbonds SN - compounds Sulfur imides S„(NHV0 \ 5^ 5 Thiazyl halogenides VI W Thiazyl trifluoride N-SF Sulfanur halogenides I X N Sulfu r-compounds wizh S-Nbonds Amidosulfuric acid HSCXNH CO(NH,)2 + 2H2SO,,---► CO, + HS03NH2 + NII.HSO, 0(2) 0(3) S-O 143.8-144.4 pm S-N 177.3 pm N - II 103.4- 103.6 pm Imido-bis(sulfuric) acid HN(S03H): 4CO(NH2)2 + 5H2S04----> 4CO: + 2HN(S03NH4)2 + (NH4)2SO 0(51 S-N 166.4, 166.5 pm S 0 143.8 - 145.1 pm K<2> KM K(3} 013) 0<2J 51 - N 160.8 pm 52 - N 159.9 pm S -0 145.6 -147.6 pm )&li>- N - S2 121.0° Nitrido-tris(sulfuric) acid KN()2 + 3KHSO,----v N(S03K)3 + KOH + H2Q VII th group PSE, ns2np5 Fluorine, chlorine, bromine, iodine, astatine ❖ Group characterics: ❖ Name „halogene" is derived from Greek and denotes „salt-forming" ❖ Oxidation degree VII+ can be expected in compounds with high electronegativity (O, F) , e.g. in HClO4, IF7 ❖ Oxidation degree - I will be realized with electropositive elements. ❖ Fluorine is the most electronegative element at al.l ❖ F is the most powerful oxidation chemical agent at all. ❖ F in compounds is known only in ox. degree -I ❖ Fluorine is extremely reactive => reactions with most elements All halogenes form two-atomic molecules X2 Halogenes - common properties F Cl Br I At Atomic number 9 17 35 53 85 Ar 18,998403 35,453 79,904 126,9045 209,99 Density of liquid /at °C) 1,513 (-188) 1,655 (-70) 3,187 (0) 3,960 (120) - M.p. / °C -218,6 -101,0 -7,25 113,6 302 B.p. / °C -188,1 -34,0 59,50 185,2 330 Ionic radius of X- / pm 133 184 196 220 - 1st ionization energy / kJ/mol 1680,6 1255,7 1142,7 1008,7 (926) Electronegativity (Allred-Rochow) 4,10 2,83 2,74 2,21 1,96 Halogenes - in nature I f 544 ppm fluoroapatite \ \ pi ) í cryolite i fluorite CaF Br 2.8 ppm salt marshland to 0.5 % sea water - 65 ppm CI 126 ppm halite NaCI I 0.46 ppm NaIO3 in Chile nitrate NaNO3 salt marshland to 100 ppm Fluorine - production Fluorine is produced only electrolytically from melt KF : HF = 1:1 to 1:3 at 72 - 240 °C Electrolytical high-temperature device is used for KF:HF 1:1, low-temperature device for KF:HF 1:3 It consists of: > vessel made from pure Ni or Monel metal (Ni+ Cu alloy) > steel cathode > carbon anode > cathode and anode spaces are separated Fluorine is supplied in bombs. It can be also prepared in small size electrolytic devices (for laboratory purposes). Chlorine - production Chlorine is a by-product coming from NaOH production by electrolysis of NaCl solution (brine) Electrolytical device : carbon (anode) steell (cathode) anode and cathode spaces are separated common temperature > hypochlorite and chlorate can be produced at higher temperatures - anode and cathode spaces must notbe separated and electolyte is stirred. > Chlorine is supplied in bombs. Chlorine - preparation 2 KMnO4 + 16 HCl — 5 Cl2 + 2 MnCl2 + 2 KCl + 8 H2O MnO2 + 4 HCl — Cl2 + MnCl2 + 2 H2O K2Cr2O7 + 14 HCl — 3 Cl2 + 2 CrCl3 + 2 KCl + 7 H2O HClO + HCl — Cl2 + H2O Bromine - preparation and production Bromine is industrially produced only by bromide oxidation using chlorine (bromide source - brine from Dead Sea or from salt Michigan slougs) Preparation in labors: K2Cr2O7 + 6 KBr + 7 H2SO4 — 3 Br2 + Cr2(SO4)3 + 4 K2SO4 + 7 H2O Bromine is supplied in sealed ampules made from dark glass in a package containing inert material. Iodine - preparation and production a) From brine: I- + Ag+ — AgI 2 AgI + Fe — FeI2 + 2 Ag FeI2 + Cl2 — FeCl2 + I2 b) From NaIO3 (in Chile nitrate): 2 IO3- + 6 HSO3- — 2 I- + 6 SO42- + 6 H+ 3I- + IO3- + 6 H+ — 3 I2 + 3 H2O Halogenes - solubility Solubility in water ^Fluorine reacts with water 2H20 + 2F2____> 4HF + 02 | > Solubility of chlorine and bromine enables after cooling formation of clathrates in crystalline form. > Iodine is very little soluble in water. > Its solubility can be enhanced in the presence of iodide => triiodide is formed I2 + I- — I3- Solubility in organic solvents Halogenes are usually well soluble in polar and non-polar solvents (carbon disulfide, diethylether, chloroform, carbon tetrachloride, ethanol) Halogenes - fluorine reactivity Fluorine reacts very rapidly with most elements > reaction with hydrogen at -252 °C is explosive. > only Cu and Ni do not react, similarly also nitrogen > Fluorine is in all reactions in the role of oxidizing agent: Si02 + 2F2----> SiF4 + 02 > Fluorine reacts with oxygen in smouldering electrical discharge at low temperatures => unstable O2F2 is formed Halogenes - chlorine reactivity Chlorine also reacts directly with most elements. Reactions are very vigorous. Chlorine + hydrogen: under burning HCl is obtained. No direct reaction was observed in the case of oxygen and nitrogen. Chlorine reacts with water: + H2O HClO + HCl Halogenes - reactivity of bromine and iodine Bromine and iodine have are very similar to chlorine, their reactivity is lower, oxidative effect included.. Halogenes - utilization of free halogenes All halogenes as elements are powerful oxidation agents => many applications ( water chlorinating, tincture of iodine) Halogenes - utilization of halogene derivatives Organic halogene derivatives are very important group of organic substances - solvents, reaction agents, polymer industry, etc. Halogene compounds - hydrogen halogenide and their salts Oxidation number of halogene is - I. HF HCl HBr HI M.p. / °C -83,4 -114,7 -88,6 -51,0 B.p. / °C 19,5 -84,2 -67,1 -35,1 AH° (at 298,15 K) / kJ/mol -271,12 -92,31 -36,4 26,48 Dipol moment |i .1030 / C.m 5,79 3,56 2,62 1,27 Halogene compounds - hydrogen halogenide and their salts -200 Influence of H-bonding on b.p. CaF2 + H2S04 NaCI + H2S04 PX3 + 3H2Q PCl3, PBr3 ■ 2HFT + CaS04 2HCIT + NaHS04 3HX + H3PO3 HF(g) and HF (acid) react with SiO2 (and with glass) SiO2 + 4 HF — SiF4 (g) + 2 H2O Halogene compounds - hydrogen halogenide and their salts Hydrogen bromide and hydrogen iodide on air yield free halogene. Sulfuric acid cannot be used for releasing HBr or HI from salts because of oxidation effect of the acid in higher concentrations, => HBr preparation: P4 + 12 H2O + 6 Br2 12 HBr + 4 H3PO3 H3PO3 + H2O + Br2 2 HBr + H3PO4 ^ HI preparation: I2 + H2S 2 HI + S 2 I2 + N2H4 - 4 HI + N2 (in water) Halogene compounds - hydrogen halogenide and their salts Free hydrogen halogenides form defined hydrates with water. HF.H20 2HF.H20 HCI.2H20 Aqueous solution are denoted as hydrohalogenic acids. Their strength increase form F to I. Hydrofluoric acid is a weak acid s pKa = 3,14 pri 25 °C when diluted. Its strength is increasing with increasin concentration as a consequence of I H- bonding. Other hydrohalogenic acid are strong acid => complete disssociation in water. All hydro halogenic acid form azeotropic mixtures with water: HF 35 % HCl 20,2 % HBr 48 % HI 57 % Sale concentrations of HF is 40 %, for HCl 35%. Halogene compounds - hydrogen halogenide and their salts Ionic [ Polymeric J Molecular S + 3F2 - - SF6 2Fe + 3CI2 ^ 2FeCI3 Hg + l2 - Hgl2 Zn + 2HCI —> ZnCI2 + H2 Ag20 + 2HF —> 2AgF + H20 KOH + HCl - —> KCl + H20 CaC03 + 2HBr CaBr2 + C02 + H20 Pb(N03)2 + 2KI -» PblJ- + 2KN03 CrCl3 + 3HF -» CrF3 + 3HC1 Halogene compounds - ionic halogenides Halogenides of alkali metals, alkali earth's elements and some transition metals (rare earth's and thorium) > Halogenides of alkali metals, alkali earth's elements are basic structural typs for many ionic (NaCl, CsCl, CaF2 ) > High m.p. are typical for ionic halogenides > In moltes state, electric conductivity. > In aqueous solution - strong electrolytes. Formation of crystalline hydrates: LiClH2O BaCl22H2O AlCl36H2O. Halogene compounds - ionic halogenides Some hydrates can be heated under formation of anhydrous salts NiCl26H2O Other salts hydrolyse FeCl36H2O AlCl36H2O Then, is is possible to get anhydrous salts by heating hydrates in stream of dry hydrogen halogenide or by using suitable drying agent. CrCl3.6H2O + 6 SOCl2 CrCl3 + 12 HCl + 6 SO2 Halogene compounds - ionic halogenides > Halogenides CuI AgI TI1 HgI PbII are not soluble in water. > Alkali and rare earth's fluorides are bad soluble in water. > Mercury chloride and bromide are not dissociated in water => HgX2 are present in aqueous solutions. Halogene compounds - covalent halogenides A) Simple molecules retaining is molecular structure also in solid state TiCl4, SnCl4, WF6, NbCl5 aj. B) Non-metal halogenides (S, N, P, Si), and some semi-metals (Te, Se, As, Sb) are: > liquids of low m.p. and b.p. > bad conductors > they are known as gases in most cases > some of them are easily sublimable Halogene compounds - covalent halogenides C) Highly condensed systems Halogenides of elements with electronegativity 1,5 - 2,2 in oxidation states II and III (exceptionallly I) with infinite atom structures (chains, layers, space systems). Examples: Layer structural typ CdCl2 - anhydrous CrCl3, FeCl2, MnCl2, CoCl2, NiCl2 (lower m.p. and b.p. as compared with ionic halogenides, easy sublimation, resp. Some of them form dimer molesules (also in gasoues state). Halogene compounds - covalent halogenides Some halogenides hydrolyse - this reaction is often used for their production SiCl4 + 2 H2O — SiO2 + 4 HCl BCl3 + 3 H2O — 3 HCl + H3BO3 Partially hydrolysis leads to the formation of oxo-halogenides: SbCl3 + H2O SbOCl + 2 HCl SF6, CF4, NF3 CCl4 are inert, no reaction and no miscibility with water. Halogene compounds - polyhalogenides The ability to form polyhalogenides F < Cl < Br < I For iodine up to I - Possible reactions: I - + ICl I2Cl- Interhalogenes > The chemical similarity of halogenes leads to the formation of „mixed" halogenes - interhalogenes (volatile low-molecular compounds of yellow, red or red-brown colour). > Interhalogenes can be prepared by direct synthesis. > Interhalogenes are very reactive. > Reaction with water: ClF + H2O HF + HClO > In anhydrous media, formation of complex anions was observed ICl3 + Cl- -> ICl-4- > Some interhalogenes can serve as halogenating agents: Mo + 2 BrF3 MoF6 + Br2 typ formula properties structure XY ClF colourless gas m.p. -156 °C b.p. -101 °C sp linear BrF light brown gas m.p. -33 °C b.p. 20 °C IF unstable at 20 °C colourless gas m.p. -83 °C b.p. 12 °C sp3d „T" shape BrF3 yellow-green liquid m.p. 8,8 °C b.p. 126 °C XY5 gas sp3d2 tetragonal pyramide IF5 colourless liquid m.p. 9,6 °C b.p. 97 °C XY7 IF7 colourless gas m.p. 5,5 °C sublimation at 4,5 °C fsp3d2 pentagonal bipyramide Halogene compounds - oxo compounds Oxides Cl2O ClO2 Cl2O6 Cl2O7 Br2O BrO2 I2O5 Oxoacids of halogenes ■o Ü < HCI04 HBr04 HI04, H5IOe HCIO3 HBr03 HIO3 HCI02 HBr02 HIO, HCIO HBrO HIO Oxidation effect Halogene compounds - oxo compounds Oxygen fluorides (-\ 2% solution 2F2 + 2NaOH----> OF2 + 2NaF + H20 Dioxygen difluoride O2F2 Dioxygen tetrafluoride O2F4 V_) OF2 is relatively stable, non-explosive (difference from explosive chlorine oxides). Very powerful oxidation agent. Fluorination of ice at low temperatures very unstable hypofluorous acid is formed. _^ F2 + H2O HFO + HF Halogene compounds - oxo compounds Chlorine oxides - relatively unstable compounds decomposing under explosion and yielding oxygen and chlorine. Cl2O, b.p. -2 °C, yellow-brown gas Preparation: HgO + 2 Cl2 -> Cl2O + HgCl2 Production: Cl2 + 2 Na2CO3 + H2O -> 2 NaHCO3 + 2 NaCl + Cl2O Cl2O is considered as anhydride of hypochlorous acid: Cl2O + H2O -> 2 HClO Halogene compounds - oxo compounds Hypochlorous acid HClO Reaction of chlorine with water: Cl2 + H2O — HClO + HCl > Unstable weak acid (pKa = 7,47 at 25 °C), stepwise decomposes to oxygen, chlorine and chloric acid. > Powerful oxidation effect: HClO + H+ + e — V2 Cl2 + H2O E° = +1,63 V => Cr (III) salts are oxidized to chromates =^> lead hydroxide to lead dioxide Halogene compounds - oxo compounds MIS WX;-^ Hypochlorites |q|_q| > extensive hydrolysis in aqueous solutions. Preparation and production: - Reaction of chlorine and cooled alkali hydroxide solutions - Electrolysis of brine, electrode spaces not separated, cooling is neccessary. > Hypochlorite solutions have whitening and antiseptic effect. (SAVO) > Chlorinated lime Ca(ClO)Cl - disinfectant > Hypochlorites disproportionate at higher temperatures: 3ClO- -> ClO3- + 2Cl- Halogene compounds - oxo compounds Chlorite acid and chlorites > Free chlorite acid is very unstable, only diluted solution can be prepared by following reaction: Ba(ClO2)2 + H2SO4 — BaSO4 + 2 HClO2 > medium strong acid, pK« 2. > powerful oxidation effect Chlorites - preparation: 2 ClO2 + 2 OH- — ClO2- + ClO3- + H2O > Chlorite solutions have oxidizing, whitening and antiseptic effect (SAVO) Chlorites - production: 2 ClO2 + O22- — 2 ClO2- + O2 2 ClO2 + Zn — 2 ClO2- + Zn2+ 3 ClO2- — 2 ClO3- + Cl- Halogene compounds - oxo compounds Chlorine dioxide, b.p. 11 °C > Yellow-brown gas, that can be easily liquidified. > Very unstable in liquid state and at higher concentrations, explosive > Soluble in water, dark-green solution, forms hydrates. ClO2 preparation 3 KClO3 + 3 H2SO4 2 ClO2 + HClO4 + 3KHSO4 + H2O (possible explosion,especially in the presence of organic substances) or better (more safe) with oxalic acid: 2KClO3 + 2 (COOH)2 2 ClO2 + 2 CO2 + K2C2O4 + 2 H2O Halogene compounds - oxo compounds Very pure ClO2 can be prepared: 2 AgClO3 + Cl2 2 ClO2 + 2 AgCl + O2 Laboratory and industrial production: 2 NaClO2 + Cl2 2 ClO2 + 2 NaCl ClO2 utilization: • whitening agent in cellulose production • disinfectant in the protection of cultural heritage objects. Halogene compounds - oxo compounds Chloric acid Preparation: Ba(ClO3)2 + H2SO4 BaSO4 + 2 HClO3 > HClO3 cannot be prepared in pure form > spontaneously decomposes at > 30 % concentrations, yielding chlorine, perchloric acid and oxygen, evtl. chlorine dioxide 8 HClO3 4 HClO4 + 2 Cl2 + 3 O2 + 2 H2O 3 HClO3 HClO4 + 2 ClO2 + H2O > HClO3 is very powerful oxidating agent in aqueous soluitons. Halogene compounds - oxo compounds Chlorates Disproportionate of chlorine in hot solutions of alkali hydroxides: 3 Cl2 + 6 OH- — ClO3- + 5 Cl- + 3 H2O Liebig way: 6 Ca(OH)2 + 6 Cl2 — Ca(ClO3)2 + 5 CaCl2 + 6 H2O Ca(ClO3)2 + 2 KCl — 2 KClO3 + CaCl2 Both salts are separated by crystallization - great difference in solubilities. Chlorate production: Electrolysis of hot brine solutions, electrode spaces are not separated. Halogene compounds - oxo compounds Properties of chlorates Pure alkali metal chlorates can be thermally disintegrated: 4 KClO3 — 3 KClO4 + KCl In the presence of impurities of a catalyst (MnO2) the decomposition can be carried out at lower temperatures 2 KClO3 — 2 KCl + 3 O2 Potassium chlorate is a part of pyrotechnic products Sodium chlorate is a powerful herbicide - TRAVEX (nowadays forbidden) Attention! Mixtures of chlorates and organic substances very easyly explosive! Halogene compounds - oxo compounds Chlorine trioxide, ClO3 , b.p. 4 °C, dark-red liquid Preparation: 2 ClO2 + 2 O3 — Cl2O6 + 2 O2 Hydrolysis: Cl2O6 + H2O — HClO3 + HClO4 Cl2O6 reacts with hydroxides to a mixture of chorate and perchlorate Reaction with nitrosyl chloride: 2 Cl2O6 + 2 NOCl — 2 NO+ + 2 ClO2 + Cl2 Halogene compounds - oxo compounds Chlorine heptaoxide b.p. 83 °C Colourless oily liquid > Explosive in contact with organic substances or by heating. > Preparation - dehydratation of anhydrous perchloric acid using P4O10 at -10 °C 4 HClO4 + P4O10 — 2 Cl2O7 + 4 (HPO3)x >Cl2O7 can be distilled from the reation mixture > Cl2O7 is anhydride of perchloric acid. Halogene compounds - oxo compounds Perchloric acid KClO4 + H2SO4 HClO4 + KHSO4 > Isolation of anhydrous acid by distillation. > Sale concentration is 72 % (azeotrope, b.p. 203 °C) > The most strong inorganic acid, miscible with water, solutions are very stable. > H3O+ ClO4- is formed in water, it can be considered as hydroxonium perchlorate HClO4 is also powerful oxidating agent, but kinetically inert, its reactions are slow. Attention! Possible explosion of concetrated solutions about 70 % after contact with organic substances. Halogene compounds - oxo compounds Perchlorates > perchlorates are produced by electrolytical oxidation of chlorates (steel cathode, Pt- anode, or made from PbO2) ^Ammonium perchlorate as a source for solid fuel in rocket fuel systems ^Potassium perchlorate is used in pyrotechnic mixtures > Magnesium perchlorate is very effective drying agent (anhydron). > Magnesium perchlorate can be used as solid electrolyte in so called „dry elements". electrolysis MCI MOH/CI decomposition *> MCIO MCIO MCIO CI cio; ci" CIO Halogene compounds - oxo compounds Br2O Br2 + HgO — Br2O + HgBr2 Reaction with water Br2O + H2O — 2 HBrO Hypobromous acid HBrO Disproportionate (similar to chlorine) Br2 + H2O — HBrO + HBr Hypobromites Br2 + 2 OH" — BrO" + Br- + H2O (coolling) Halogene compounds - oxo compounds Br02 This oxide can be prepared by reaction of bromine with ozone at - 78 °C in CF3Cl as solvent: Br2 + 03 2 + 4 02 Disproportionation in alkali hydroxides 6 Br02 + 6 OH- 5 Br03- + Br- + 3 H20 Halogene compounds - oxo compounds Bromic acid, HBrO3 5HCIO + Br2 + H20---* 2HBr03 + 5HCI |Sj Ba(Br03)2 + H2S04--, 2HBr03 + BaS04 Br2 + 5 Cl2 + 6 H2O — 2 HBrO3 + 10 HCl HBrO3 is very similar to HClO3 in aqueous solutions, powerful oxidating agent Free bromine reacts with hot alkali hydroxides similarly to chlorine: 3 Br2 + 6 OH- — BrO3- + 5 Br- + 3 H2O Bromates Thermally unstable, decompose by heating 2 KBrO3 — 2 KBr + 3 O2 Halogene compounds - oxo compounds Perbromic acid, HBrO4 Prepared recently: BrO3- + XeF2 + H2O BrO4- + Xe + 2HF HBrO4 can be prepared in 55% concentration without any risk HBrO4 in concentated state is powerful oxidating agent, it dissolves easily also corrosion-proof steels. Perbromates Production: BrO3- + F2 + 2 OH- BrO4- + 2 F- + H2O Diluted perbromate solutions have mild oxidating effect. Iodine oxide The only genuine is I2O5 + I4O12 I (V) and I (VII) I2O5 > preparation by thermal dehydratation of HIO3 at 240 °C 2 HIO3 — I2O5 + H2O > further heating over 300 °C causes decomposition to elements > I2O5 is anhydride of iodic acid. Utilization - detection of CO I,05 + 5CO Other binary compounds of iodine and oxygen UO 2W4 Iodosyl iodate [mO [() 1 °' Iodine (III) iodate Im IVO3)3 Halogene compounds - oxo compounds Hypoiodous acid, HIO - unstable, very weak acid Preparation: I2 + H2O HIO + HI or 2 I2 + 2 HgO + H2O — HgO.HgI2 + 2 HIO Possible ionization as a base in aqueaous: HIO + H2O I(H2O)+ + OH- > oxidation effect > dispropotionation to iodate and iodide. Hypoiodites I2 + 2 OH- — IO- + I- + H2O (cooling) Halogene compounds - oxo compounds Iodic acid 3I2 + IOHNO3---, 6HIO3 + 10NO + 2H20 l2 + 5CI2 + 6H20--► 2HI03 + 10HCI I Iodates l2 + 2X03"---» X2 + 2IO3" (X = CI,Br) I2 + NaClO3 — 2 NaIO3 + Cl2 Both have powerful oxidating effect. Halogene compounds - oxo compounds Tetraoxoperiodates can be prepared by iodate oxidation using hypochlorite IO3- + ClO- IO4- + Cl- Halogene compounds - oxo compounds Orthoiodic acid (pentahydrogeniodic) Preparation by thermal decomposition of some iodates 5 Ba(IO3)2 Ba5(IO6)2 + 4 I2 + 9 O2 Ba5(IO6)2 + 5 H2SO4 2 H5IO6 + 5 BaSO4 Thermally decomposes to iodine pentaoxide, oxygen and water. Orthoiodates can be prepared by iodate oxidation using chlorine in alkali media: IO3- + 6 OH- + Cl2 IO65- + 2 Cl- + 3 H2O Halogene compounds - oxo compounds Properties of orthoiodic acid > Strong oxidation effect, Mn(II) is oxidized to permanganate in aqueous media > Ligand in complexes (difference from perchlorates). Octahedral anions can be bidentally bonded - formation of chelate cycle. > Orthoiodates stabilize central atoms od complexes in high oxidation degrees. ASTATINE > Chemistry of At is little investigated because of short half-lifes > 211At T (1/2) = 7,21 h. > Oxidation states of At: 0, -I, V nd posiible I, III, VII. > The only non-disproportionating halogene in oxidation degree 0 - V. > Astatine gives interhalogene compounds of the type AtX that can be extracted to carbon tetrachloride. > The only possible utilization of astatine should be in medicine. It is supposed that its application for thyroid gland treatment is more suitable than in the case of iodine agents. VIIIth group PSE - Noble gases Helium, neon, argon, crypton, xenon, radon He, Ne and Ar no compounds are known Compounds of Kr and Xe are known. Rn is able to create stable fluoride and other types of compounds. But no stable isotopes are known => these compounds have no significance. Noble gases - properties He Ne Ar Kr Xe Rn Atomic number 2 10 18 36 54 86 Atomic mass 4,00260 20,179 39,948 83,80 131,29 222 M.p. / °C -248,61 -189,37 -157,2 -111,8 -71 B.p. / °C -268,93 -246,06 -185,86 -153,35 -108,13 -62 1st Ionization potential (eV) 24,58 21,56 15,76 14,00 12,13 10,75 Heat of evaporation / kJ mol1 0,08 1,74 6,52 9,05 12,65 18,1 Water solubility /cm3 kg-1 8,61 10,5 33,6 59,4 108,1 230 Noble gas - occurrence Noble gases create about 1 % Earth's atmosphere Helium: > The second most common element in the universe (23%). > He is produced in the cores of stars as a product of nuclear fusion of hydrogen atoms. > 4He is a product of cc-decay of heavy metals. > He can be extracted from natural gas after liquidifying other components. Its content varies considerably according to the gas reservoir in the range of 0,4-7 %. Noble gas - occurrence Neon, argon (1 %), crypton, xenon- in the air 40Ar is a product of electron capture (EC) of 40K. Radon is a product of radioactive decay of Ra. Rn use is not relevant because of its short half-life (3.824 day). Radon is a problem for environment ^radioactive emanation released from the rocks (e.g. granite) can accumulate in harmful concentrations . Its decay products are solids and deposit low parts of residential places. Noble gases - use > Use of He and Ar, especially, in metallurgy, in chemical synthesis -creating an inert atmosphere, e.g. in welding processes. > Helium - as a carrier gas in gas chromatography, NMR spectroscopy > He (l) is a superconductor, because it has the lowest b.p.of all elements > Used as a coolant in cryotechnics > In the gaseous state, He has high thermal conductivity => cooling medium for nuclear reactors. Noble gas - clathrates The name comes from Latin "clathratus" (= enclosed in a cage). The concept of host and guest Clathrate consists of Ar, Kr, and Xe, also other molecular gases (SO2, O2, N2, CO), or other molecules (guest) in structures of solid substances (host). Clathrates - arrangement of "host- molecules" in the crystal - the cavities of subsystem are created. Atoms or molecules of different substances in general can be closed into this cavity, they are then bound only by weak van der Waals forces. These substances have a nonstoichiometric composition. They are not typical chemical compounds with typical chemical bonds. Noble gas - clathrates > Clathrates are relatively stable, gas is released out in melting |;|process. > They are formed by crystallization e.g. from water saturated with gas at ^pressure of 1-4 MPa. 8Kr. 46 H2O > Their practical use is associated with the need to prevent releasing radioactive isotopes of noble gases that occur in nuclear reactors. > Clathrate containing 20 % argon under high pressure can be formed The other substance which creates clathrate with Ar, Kr, and Xe is e.g. hydroquinone. Their composition is close to the limit value of the ratio of gas: hydroquinone 1:3 . Noble gases - compounds Stable compounds are created only in the case of Kr and Xe. It is difficult to isolate radon compounds because of their hight activity and rapid radiolysis. N. Bartlett and D.H. Lochman in 1962 carried out reaction between Xe and PtF6 . 25 °C 60 °C Xe + 2 PtF6 [ XeF]+ [PtF6]- + PtF5 [ XeF]+ [Pt2 F11]- Other similar compounds such as XeF2 a XeF4 were prepared later. These reactions were first ones that led to noble gas compounds. Noble gas - xenon compounds Oxidation state Formula M.p. / °C Stereochemistry II XeF2 129 Door, linear IV XeF4 117,1 D4h square VI XeF6 49,5 distorted octahedron XeOF4 -46 ^4v square pyramid XeO2F2 30,8 C2v CsXeOF5 distorted octahedron KXeO3F square pyramid XeO3 explode pyramid VIII XeO4 -35,9 tetrahedron XeO3F2 -54,1 trigonal bipyramid Ba2XeO6 >300 decay Oh octahedron Noble gases - fluorides > They are synthesized directly from Xe and F2, the molar ratio, temperature and pressure are important. > Reaction is done in closed nickel container. > Products are white crystalline compounds. XeF, commercially available, these substances are not yet practicaly use. For research, these compounds are very interesting objects, especially for studies of their chemical bonds. KrF2 is best investigated compound of Kr. It is stable around -153 ° C. Noble gases - xenon fluorides XeF2 linear shape of a molecule, it easily dissolves in water and its solutions are around 0° C in a neutral media stable. In the presence of alkaline media, the most common reaction is hydrolysis: 2 XeF2 + 2 H2O — 2 Xe + 4 HF + O2 Aqueous solution of XeF2 is a weak fluorinating and strong oxidizing agent: 2 Ag+ + XeF2 — 2 Ag2+ + Xe + 2 F" XeF2 + 2 Cl" — Xe + Cl2 + 2 F" Utilization: XeF2 + BrO3" + 2 OH" — Xe + BrO4" + 2 F" + H2O Oxidation reaction of chromium salts to chromates is also realizable.. Noble gases - xenon fluorides XeF4 Square molecule, symetry D4h > easy sublimation > reactions are similar to the reactions XeF2, but XeF4 is stronger fluorinating agent than XeF2 2 Hg + XeF4 Xe + 2 HgF2 Pt + XeF4 Xe + PtF4 2 SF4 + XeF4 Xe + 2 SF6 Hydrolysis: (complicated mechanism) 6 XeF4 + 12 H2O 2 XeO3 + 4 Xe + 3 O2 + 24 HF ■ Noble gases - xenon fluorides XeF6 > More volatile than XeF4 > Water causes rapid decomposition => mixture of products, contains also explosive XeO3. XeF6 is very powerful fluorinating agent - also glass is attacked 2 XeF6 + SiO2 — 2 XeOF4 + SiF4 2 XeOF4 + SiO2 2 XeO2F2 + SiF4 2 XeO2F2 + SiO2 — 2 XeO3 + SiF4 The shape of XeF6 molecule is not sufficiently investigated till now. Noble gases - xenon fluoride-oxide Intermediates arising in the course of hydrolysis, e.g. XeF6 + H2O — XeOF4 + 2 HF They can be obtained in reactions between xenon fluorides and oxides. XeO4 + XeF6 — XeO3F2 + XeOF4 XeO3 + XeOF4 — 2 XeO2F2 Volatile, colorless liquids or low m.p., easily undergoing to hydrolysis in aqueous solutions. Compounds with the Xe-N and Xe-C bonds are also known, but they are not too stable. Noble gases - oxygen compounds XeO3 > dangerous, highly explosive > arises in hydrolytical reaction of xenon fluoride > strong oxidizing agent (in water) > its reactions are slow (it is kinetically inert). > Its solutions are called xenonic acid and they are stable, unless they contain oxidizable substances or alkalis Xenonates are formed in alkaline media: XeO3 + OH- -- HXeO4- Their disproportionation leads to salts of xenon (VIII) and elemental xenon: 2 HXeO4- + 2 OH- — XeO64- + Xe + O2 + 2 H2O Noble gas - oxygen compounds Xenon salts They can be only obtained by precipitation XeO3 with NaOH in solution and in the presence of ozone. Na4XeO62,5H2O Ba2XeO6 They decompose by action of concentrated sulfuric acid under formation of explosive gaseous XeO4 (similar as XeO3).