General Inorganic Chemistry PDF

Group 15, 5A Valence-shell configuration: ns2 np3 Exhibits varied chemical properties. 1. N and P are nonmetals; 2. As and Sb are metalloids; 3. Bi is a metal (the heaviest non-radioactive element) Some Physical Properties, Sources, and Methods of Prep...

Group 15, 5A Valence-shell configuration: ns2 np3 Exhibits varied chemical properties. 1. N and P are nonmetals; 2. As and Sb are metalloids; 3. Bi is a metal (the heaviest non-radioactive element) Some Physical Properties, Sources, and Methods of Preparation  Oxides of Group 5A Elements Nitrogen: N2O, NO, N2O3, NO2, N2O4, N2O5; Phosphorus: P4O6 & P4O10; Arsenic: As2O3 (As4O6) & As2O5; Antimony: Sb2O3 & Sb2O5 Bismuth: Bi2O3 & Bi2O5 phosphorus in orange, oxygen in red  Chlorides of Group 5A Elements Nitrogen: only NCl3; Phosphorus: PCl3 and PCl5; Arsenic: AsCl3 and AsCl5; Antimony: SbCl3 and SbCl5; Bismuth: BiCl3 All are molecular compounds.  Reactions of Oxides and Chlorides 3NO2(g) + H2O(l) → 2HNO3(aq) + NO(g); N2O5(g) + H2O(l) → 2HNO3(aq) P4O10(s) + 6H2O(l) → 4H3PO4(aq); As2O5(s) + 3H2O(l) → 2H3AsO4(aq); PCl5(s) + 4H2O(l) → H3PO4(aq) + 5HCl(aq); AsCl5(s) + 4H2O(l) → H3AsO4(aq) + 5HCl(aq); 2SbCl5(s) + 5H2O(l) → Sb2O5(s) + 10HCl(aq);  The Chemistry of Nitrogen The triple bonds (N≡N) in N2 provide high stability to the molecule; Oxidation states from -3 to 5 NH2OH (-1), N2H4 (-2), and NH3 (-3) Many reactions involving nitrogen gas are endothermic and compounds containing nitrogen decompose exothermically to the elements. N2(g) + O2(g) → 2NO(g) ΔH° = 180 kJ 2NO2(g) → N2(g) + O2(g); ΔH° = -68 kJ N2H4(g) → N2(g) + 2H2(g); ΔH° = -95 kJ  Nitrogen Fixation The process of transforming N2 to other nitrogen–containing compounds. Atmospheric fixation (occurs naturally during thunderstorm): N2(g) + O2(g) → 2NO(g); ΔHo = 180 kJ 2NO(g) + O2(g) → 2NO2(g); ΔHo = -112 kJ 3NO2(g) + H2O(l) → 2HNO3(aq) + NO(g); ΔHo = -140 kJ  Biological Nitrogen Fixation Fixation of atmospheric N2 by bacteria living in soils and water; some live in root nodules; Plants such as legumes and alfalfa have root nodules that contain nitrogen-fixing bacteria – they benefit directly from these bacteria; Other plants benefit when the bacteria die and release absorbable forms of nitrogen (NH3, NH4+, and NO3-) to the soils;  Biological Nitrogen Fixation In nitrogen-fixing bacteria 1. Atmospheric N2 is first reduced to NH3; 2. In bacterial cells, NH3 becomes NH4+, oxidized to NO2- and then to NO3-; 3. NH3, NH4+, and NO3- can be released into the surroundings (water or soils) and become available to plants; Denitrifying bacteria (in soils) change NO3- back to NO2-, NH3, and finally to N2 to complete the biological nitrogen cycle. Biological Fixation and The Nitrogen Cycle  Industrial Nitrogen Fixation Industrial Fixation (the Haber Process): N2(g) + 3H2(g) → 2NH3(g) ΔH° = -92 kJ Most NH3 are converted to: 1. Fertilizers (~70%) 2. Nitric acid, HNO3 (~20%) 3. Hydrazine, N2H4, and monomers for various plastics and nylons. The Haber Process  Important Hydrides of Nitrogen Ammonia, NH3 (most important hydride)  Production of fertilizers (NH4NO3, (NH4)2SO4, (NH4)3PO4, and CO(NH2)2), HNO3, and N2H4 Hydrazine, N2H4  Rocket propellant, manufacture of plastics, agricultural pesticides; Monomethylhydrazine, CH3N2H3  Rocket fuels Production of HNO3 Oswald Process: 1. NH3(g) + O2(g) → NO(g) + H2O(g) 2. 2NO(g) + O2(g) → NO2(g) 3. 3NO2(g) + H2O(l) → 2HNO3(aq) +NO(g) HNO3 - a strong acid and an oxidizing agent; Reactions with metals does not produce H2 Cu(s) + 4HNO3(16 M) → Cu(NO3)2(aq) + 2NO2(g) + 2H2O(l) 3Cu(s) + 8HNO3(aq, 6 M) → 3Cu(NO3)2(aq) + 2NO(g) + 4H2O(l) 4Zn(s) + 10HNO3(aq, 3 M) → 4Zn(NO3)2(aq) + N2O(g) + 2H2O(l)  Allotropes of Phosphorus White Phosphorus: P4 (tetrahedral) - very reactive Black Phosphorus: crystalline structure - much less reactive Red Phosphorus: amorphous with P4 chains  Allotropes of Phosphorus White Black Red very reactive P(white) (heat, 1 atm, no air) P(black) P(white) or P(red) (high pressure) P(black)  Oxides of Phosphorus Reaction of white phosphorus with oxygen: 1. P4(s) + 3O2(g) → P4O6(l); (o.s. of P = +3) 2. P4(s) + 5O2(g) → P4O10(s); (o.s. of P = +5) Reactions of phosphorus oxides with water: 1. P4O6(l) + 6H2O(l) → 4H3PO3(aq); 2. P4O10(s) + 6H2O(l) → 4H3PO4(aq); phosphorus in orange, oxygen in red  Oxyacids of Phosphorus Phosphoric acid, H3PO4 - triprotic Phosphorous acid, H3PO3 - diprotic Hypophosphorous acid, H3PO2 - monoprotic  ATP Adenosine triphosphate ADP Adenosine diphosphate  Phosphorus Halides Reactions of white phosphorus with halogens: 1. P4(s) + 6X2 → 4PX3(l); 2. P4(s) + 10X2 → 4PX5(s); Examples of reactions: PCl3(l) + Cl2(g) ⇄ PCl5(s) PCl3(l) + 3H2O(l) → H3PO3(aq) + 3HCl(aq) PCl5(s) + 4H2O(l) → H3PO4(aq) + 5HCl(aq)  Important Compounds of Phosphorus Ca3(PO4)2 & Ca5(PO4)3F : source of phosphorus Ca5(PO4)3(OH) : forms bones and teeth P4O10 : formation of H3PO4 H3PO4 : production of fertilizers & phosphates H2PO4- & HPO42- : phosphate buffers Na3PO4 : scouring powder and paint remover Na5P3O10 : fabric softeners ADP & ATP : storage of metabolic energy PCl5 : precursor for lithium hexafluorophosphate (LiPF6), an electrolyte in lithium ion batteries;  Group 16, 6A Valence-shell configuration: ns2 np4 O, S, Se, Te, Po None of the Group 6A elements behaves as a typical metal. Elements form covalent bonds with other nonmetals. Some Physical Properties For S, Se, Te - similar gases with characteristic odor, poisonous Hydrogenchalkogenides (HS)–I and chalkogenides S–II Soluble in water alikali metals, alkaline earth and NH4+ Heavy metal sulfides – non-soluble – precipitation from aqueous solutions - Sulfur used for removing small amounts of elemental mercury - Melting of sulfides with sulfur  polysulfides; e.g. Iron(II) disulfide FeS2 (pyrite)  Oxygen O2 makes up 21% of the Earth’s atmosphere. O3 (ozone) exists naturally in the upper atmosphere (the stratosphere) of the Earth. Ozone – forming at electric discharges - pale blue gas  Ozone layer absorbs UV light and acts as a screen to block most uv- radiation from reaching the Earth’s surface.  We now know that Freons (CFCs = chloro-fluorocarbons) are promoting destruction of ozone layer. Strong oxidizing agent  Various Forms of Oxides Metal oxides (ionic) 1. Nonconductor – example: MgO 2. Semiconductor – example: NiO 3. Conductor – example: ReO3 4. Superconductor – example: YBa2Cu3O7 Nonmetal oxides (covalent): Molecular oxides – examples: CO2, NO, NO2, N2O, SO2, P4O10, etc. Covalent network oxide – SiO2  Characteristics of Oxides Metallic oxides – basic or amphoteric Examples: Na2O (basic); Al2O3 (amphoteric) Semi-metallic oxides – mild to weakly acidic Example: B2O3 Nonmetallic oxides – weak to strong acids Examples: 1. SO2(g) + H2O(l) → H2SO3(aq) (weak acid); 2. SO3(g) + H2O(l) → H2SO4(aq) (strong acid);  Redox properties of water in solid Hydrogen peroxide weak acid disproportionation ΔHo of −98.2 kJ/mol; ΔS of 70.5 J/(mol·K) Redox properties in acidic env. in basic env.  Sulfur Sulfur is found in nature both in large deposits of the free element and in ores such as: Galena = PbS, Cinnabar = HgS, Pyrite = FeS2, Gypsum = CaSO4⋅2H2O, Epsomite = MgSO4.7H2O, and Glauberite = Na2Ca(SO4)2 Sulfur Mining: Frasch Process Aggregates of Sulfur  Sulfur Oxides and Oxyacids S(s) + O2(g) → SO2(g) 2SO2(g) + O2(g) → 2SO3(g) SO2(g) + H2O(l) → H2SO3(aq) SO3(g) + H2O(l) → H2SO4(aq) H2SO3 – diprotic; weak acid H2SO4 – diprotic; strong acid  Sulfuric Acid Productions: 1. S8(s) + 8 O2(g) → 8SO2(g); 2. 2H2S(g) + 3 O2(g) → 2SO2(g) + 2H2O(l); 3. FeS2(s) + 11 O2(g) → Fe2O3(s) + 8SO2(g); 4. 2SO2(g) + O2(g) → 2SO3(g); (V2O5/K2O catalyst) 5. 2SO3(g) + H2SO4(l) → H2S2O7(l); 6. H2S2O7(l) + H2O(l) → 2H2SO4(l); Oxyacids of sulfur Sulfurous acid Sulfuric acid Disulfurous acid Polysulfuric acids H2SO4.nSO3 di- Thiosulfurous acid Peroxymonosulfuric acid Dithionous acid Peroxydisulfuric acid Dithionic acid Polythionic acids Thiosulfuric acid  Important Compounds of Sulfur H2SO4 – most important compound, for manufacture of fertilizer, soap, detergents, metal and textile processing, sugar refinery, and organic syntheses; SF4 – for fluoridation SF6 – as insulating and inert blanket Na2S2O3 – as reducing agent and complexing agent for Ag+ in photography (called “hypo”); P4S3 – in “strike-anywhere” match heads  Exercise #8 Draw Lewis structures for the following molecules: 1. SO2 2. SF2 3. SF4 4. SF6 5. H2SO4 6. H2SO3 7. H2S2O7