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Oxygen on Earth H2O (oceans) O2, CO2 (atmosphere) CO3− (rocks, coral, seashells) SiO2, silicates (sand, clay, rocks) Oxygen Content Earth Crust Air Made commercially by fractional distillation of air (b.p. = 90K) ALLOTROPES OF OXYGEN O2 Paramagnetic (why?) O3 Higher energy form - important UV absorber in the stratosphere Light or electrical discharge 3O2 2 O3 decomposition Ozone (O3) is a strong oxidizing agent, highly toxic Kills bacteria (replacement for Cl2 in municipal water treatment) Irritating component of photochemical smog OXYGEN IONS Oxide Ion ⇒ O2− (most compounds) e.g. Li2O = 2Li+ O2− Peroxide Ion ⇒ O22− = −O – O− e.g. Na2O2 = 2 Na+ −O – O − Also, H2O2 (hydrogen peroxide) Superoxide Ion ⇒ O2− e.g. KO2 = K+ O2− Can have positive oxidation states in combination with fluorine + 2 in OF2 HYDROGEN PEROXIDE • Strong oxidizing agent (30-85% solutions) e.g. bleaching wood pulp to produce white paper • Hair bleach (~6% solution) • Antiseptic (3% solution) H2O2 decomposition can be explosive: 2 H2O2 → 2 H2O + O2 ΔH = −200 kJ/mol (disproportionation reaction) HYDROGEN PEROXIDE Reduction to H2O: H2O2 + 2H+ + 2I− → I2 + 2H2O Oxidation to O2: 2MnO4− + 5 H2O2 + 6H+ → 2Mn2+ + 5O2 + 8H2O SULFUR Sources: Sulfide Minerals (S2-): FeS2 (Pyrite) Cu3FeS3 (Bornite) PbS (Galena) ZnS (Zinc Blende) - Iron Ore (Fool’s Gold) Source of Cu Source of Pb Source of Zn Sulfate Minerals (SO42-) e.g. Na2SO4 , MgSO4 Also, CaSO4 · (H2O)2 (Gypsum) Used for wallboard, plaster of Paris. COMMERCIAL SOURCES OF SULFUR 1) Sulfur Mines – Along Gulf of Mexico, deposits of S8 ⇒ Frasch Process. 2) Byproduct from other manufacturing processes. a) Production of Zn, Pb, and Cu from their sulfide ores. b) Petroleum – 3% S. c) Coal – 5%. SO2 forms when coal is burned. SO2 + H2O → H2SO3 SO2 +[O] →SO3 +H2O → H2SO4 CaCO3 + H2SO4 → CaSO4 + H2O + CO2 (Marble) Acid Rain Sulfur Two allotropes S8 yellow, cyclic Polymer Sx⇒ red-brown polymer: zigzag chains of sulfur atoms S8(s) → S8(l) → Melts at 113°C SSSSSSS T> 150°C Sulfur Common Oxidation States +6 SO3; H2SO4 (sulfuric acid) can’t be oxidized can only be reduced +4 SO2; H2SO3 (sulfurous acid) can be both oxidized AND reduced -2 H2S; S2can’t be reduced can only be oxidized Compounds of S • H2SO4 most important industrial chemical • H2S (rotten egg smell) (S2- ) source: metal sulfides + strong acid e.g. ZnS + HCl → ZnCl2 + H2S(g) – – poisonous tarnishes Ag in presence of O2 4Ag + 2H2S + O2 → 2Ag2S + 2H2O – Organic sulfides, e.g. C4H9SH Strong odor – added to natural gas • S2O32- (thiosulfate) used in photography: forms water soluble complexes with Ag Uses of H2SO4 • Making phosphate fertilizer Ca3(PO4)2 + 3H2SO4 → 3CaSO4 + 2H3PO4 ~65% of H2SO4 • Manufacture of chemicals • Metal refining • Petroleum refining (as catalyst) • Strong oxidizing agent • Drying agent Selenium, Tellurium Source: metal sulfides byproducts of Cu, Pb refining Uses: semiconductors e.g. Se: Has low electrical conductivity in the dark which increases in light - photoconductor Used in photocopiers, light meters in cameras Compounds: form covalent bonds Oxides and hydroxides are acidic (typical of nonmetals) Se Te Po non metal semi-metal metal Nitrogen Nitrogen (N2) is very unreactive triple bond energy = 941kJ/mol Source fractional distillation of air (78% of air is N2) KNO3 water soluble salts NaNO3 found in deserts Nitrogen fixation: formation of N containing compounds from N2 N is an essenial element in proteins, nucleic acids & necessary to maintain soil fertility Compounds of Nitrogen Oxidation states of −3 to +5 Compounds with H 1. NH3 (−3 oxidation state) 2. N2H4 (−2 oxidation state) strong reducing agent: N2H4 → N2(g) + 2 H2(g) forms N2 readily: ΔH = − ΔS = + 3. Dimethyl hydrazine (rocket fuel) N compounds with oxygen N2O colorless, odorless gas used as anesthetic (laughing gas) propellant in whipping cream NO formed in car engines: N2 +O2 ↔ 2NO N2O3 blue solid, decomposes: N2O3 → NO + NO2 NO2 brown gas; component of smog N2O4 2NO2 ↔ N2O4 N2O5 unstable, decomposes to NO2 HNO3 Produced from NH3 by Ostwald process (catalytic oxidation). Uses: fertilizer NH3 + HNO3 → NH4NO3(s) strong acid strong oxidizing agent. cleaning agent to make explosives (e.g. nitroglycerine, TNT) Hydrolysis of oxides Hydrolysis: reaction with water N is a non metal: oxides are acidic. Oxide + H2O = hydroxide N2O3 + H2O → 2HNO2 (nitrous acid) 3NO2 + H2O → 2HNO3 + NO N2O5 + H2O → 2HNO3 (nitric acid) PHOSPHORUS Source: Phosphate Minerals Ca3(PO4)2 contains PO43- (tetrahedral P) P is made by heating Ca3(PO4)2 and coke in an electric furnace. 2Ca3(PO4)2(s) + 10C(s) + 6SiO2 → 6CaSiO3(s) + 10CO(g) + P4(g) Two allotropes: P White P: P4, tetrahedral P P P Red P: Polymeric P P P P P P PHOSPHORUS ALLOTROPES White phosphorus (P4) burns spontaneously in air. P4(s) + 5O2(g) → P4 O10 ΔH = −3000 kJ/mole Red phosphorus (polymeric) is more stable. Not volatile. Does not react with air at 25°C. 600oC Red P White P Let Stand OXIDES OF PHOSPHORUS PHOSPHORUS OXYACIDS P4O10 + 6H2O → 4H3PO4 phosphoric acid P4O6 + 6H2O → 4H3PO3 phosphorous acid Also H3PO2 hypophosphorous acid USES OF PHOSPHORUS Fertilizer P is essential for plant growth Ca3(PO4)2 + 3H2 SO4 → 2H3PO4 + 3CaSO4 H3PO4 + 3NH3 → (NH4)3PO4 Detergent Complexes metal ions Biological molecules (DNA, RNA) Biochemical energy source (ATP) COMPARISONS IN GROUP V Nitrogen N2(g) N≡N NH3 is stable. Non-metal ⇒ oxides dissolve to give acidic solutions N2O3 + H2O →2HNO2 N2O5 + H2O → 2HNO3 3NO2 + H2O → 2HNO3 + NO PHOSPHORUS Allotropes: White P ⇒ P4, tetrahedral. Red P ⇒ polymer. PH3 burns in air. Non-metal ⇒ oxides dissolve to give acidic solutions: P4O10 + 6H2O → 4H3PO4 P4O6 + 6H2O → 4H3PO3 ARSENIC Allotropes: Yellow As ⇒ As4 Gray As ⇒ brittle solid. AsH3 ignites spontaneously in air. As4O10 – acidic oxide: As4O10 + 6H2O → 4H3AsO4 As4O6 is amphoteric, but is more soluble in base. ANTIMONY – Sb Brittle gray metalloid. Sb4O6 is amphoteric. There is no Sb4O10. BISMUTH - Bi Bismuth is a metal. Bi4O6 is basic and dissolves only in acids. Bi(OH)3 is basic. Bi5+ is rare. OXIDATION STATES P5+ dominates. As3+, As5+ are equally common. Sb3+ dominates. Bi3+ dominates. Inert Pair Effect HYDRIDE STABILITY NH3 is stable. PH3 is stable but burns in air. AsH3 decomposes easily. SbH3, BiH3 are very unstable. GROUP V TRENDS Going down the periodic table: 1) 2) 3) 4) 5) Electronegativity decreases. Switch from non-metallic to metallic. Hydroxides and oxides become more basic. Hydrides become less stable. “Inert pair effect” becomes more pronounced: +3 becomes more stable as compared to +5. CARBON and Group IV Carbon Sources: 1) Elemental form – coal. 2) Carbonate rocks (CO3 2-) Limestone, marble, chalk = CaCO3 Dolomite = MgCO3 ALLOTROPIC FORMS OF CARBON 1) Diamond - used as abrasive, in drill bits and cutting tools, and as a gem. 2) Graphite - used in batteries, pencils, and lubricants. 3) Fullerenes - More recently discovered molecules such as C60 which has the shape of a soccer ball. Carbon Black – Soot Amorphous form of carbon used in tires, inks, pigments, and carbon paper. CARBIDES 1) Ionic Carbides Contain C4- or C22- (-C≡C-) C4-: Be2C, Al4C3 react with water to give CH4. C2 2-(-C≡C-): CaC2 reacts with water to give HC≡CH. 2) Covalent Carbides Carbon is bound covalently to a metal or metalloid. SiC - almost as hard as diamond, does not react w/water 3) Interstitial Carbides Metals with carbon atoms found in between the metal atoms in the structure. Steel – often harder than the pure metal. SILICON Second most abundant element. Found in combination with O. Silicate Minerals: [Si2O52-]n, SiO44Sand: SiO2 (this is also quartz). With aluminum in aluminosilicates (clay, feldspars). Prepared by: SiO2(s) + 2C(s) → Si(l) + 2CO(g) sand coke 98% (3000°C) Very pure silicon (<1 ppb impurity) is required for electronics applications. GROUP IV TRENDS Going down the periodic table: 1) The +2 oxidation state becomes more stable than +4 due to the “inert pair” effect. +2 is rare for C, Si, Ge. +2 in some compounds, +4 most common for Sn. +4 is unstable for Pb ⇒ strong oxidizing agent (PbO2) 2) Basicity of oxides and hydroxides increases. CO2, SiO2, GeO2 are weakly acidic. SnO, SnO2, PbO are amphoteric. 3) Hydrides become less stable. Enormous number of stable hydrocarbons. SiH4 is stable but is spontaneously flammable. Ge, Sn, Pb hydrides are very unstable. Orbital Hybrids and Valence 2s Li Be 2p B C N O F 3s Na Mg 3p Al Si P S Cl The differences between the 2nd and 3rd periods: 2nd period: Only s and p orbitals are possible with n = 2 Therefore, the maximum number of bonds is 4 (single and/or double bonds) Examples: CH4, NF4+, BH43rd (and higher periods): can use d-orbitals to make bonds E.g. PF5 P atom is sp3d SF6 S atom is sp3d2 Let’s look at valences: N can gain 3 electrons or lose 5 to make an octet But, N can only make 4 bonds (maxiumum for n=2) Therefore N usually has a valence of 3 (NH3, NCl3, CH3NH2 - all have 3 bonds and one lone pair on the N atom) N with oxidation state 5 never has more than four bonds: O e.g., NO3 N=O (4 bonds to N) O NO2+ O=N=O (4 bonds to N, like CO2) Likewise, O usually makes 2 bonds: H2O, OF2, H2C=O Likewise, C can gain 4 or lose 4 electrons to make an octet (valence = 4) So carbon always makes 4 bonds CH4 (4 single bonds) O=C=O (2 double bonds) H-C≡C-H (1 single + 1 triple bond) H2N H2N C=O (2 single + 1 double bond) (urea) What about 3rd (and higher) periods - Si, P, S…? For these elements, double bonds are very uncommon (usually only single bonds) Si, P, S, ... Ge, As, Se, ... Sn, Sb, Te, ... Compounds have only single bonds (double and triple are very rare) Reason: atoms past second row are too big C Si good sideways overlap of p orbitals (double and triple bonds OK) C Si poor overlap of p orbitals ---- no multiple bonds (can still make single bonds) So CO2 is molecular (O=C=O, has double bonds) But SiO2 (quartz, sand, glass…) is a 3-dimensional solid network: O2 is molecular (O=O, has a double bond) But S forms rings (e.g., S8) Nitrogen (N2) has a triple bond N≡N (very stable molecule) But phosphorus is found in several forms (white, red, black), all of which have only single bonds. The chemistry of carbon is unique because: • It has a valence of 4 (highest in 2nd period) • It can make stable bonds with itself • It can make multiple bonds to C, N, O • The C-H bond is nonpolar, but bonds to other elements (N, O, halogens) are polar This is why life is based on the chemistry of carbon (organic chemistry)