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Air

Oxygen on Earth. H 2 O (oceans) O 2 , CO 2 (atmosphere) CO 3  (rocks, coral, seashells) SiO 2 , silicates (sand, clay, rocks). Oxygen Content. Earth Crust. Air. Made commercially by fractional distillation of air (b.p. = 90K). ALLOTROPES OF OXYGEN. O 2 Paramagnetic (why?)

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Air

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  1. 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)

  2. 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

  3. 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 withfluorine • + 2 in OF2

  4. 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 H = 200 kJ/mol (disproportionation reaction)

  5. HYDROGEN PEROXIDE Reduction to H2O: H2O2 + 2H+ + 2I I2 + 2H2O Oxidation to O2: 2MnO4 + 5 H2O2 + 6H+ 2Mn2+ + 5O2 + 8H2O

  6. SULFUR Sources: Sulfide Minerals (S2-): FeS2 (Pyrite) - Iron Ore (Fool’s Gold) Cu3FeS3 (Bornite) - Source of Cu PbS (Galena) - Source of Pb ZnS (Zinc Blende) - Source of Zn Sulfate Minerals (SO42-) e.g. Na2SO4, MgSO4 Also, CaSO4 · (H2O)2 (Gypsum) Used for wallboard, plaster of Paris.

  7. 1) Sulfur Mines – Along Gulf of Mexico, deposits of S8 Frasch Process. • 2) Byproduct from other manufacturing processes. • 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 • Acid Rain • SO2 +[O] SO3 +H2O  H2SO4 • CaCO3 + H2SO4 CaSO4 + H2O + CO2 • (Marble) COMMERCIAL SOURCES OF SULFUR

  8. Sulfur Two allotropes S8 yellow, cyclic Polymer Sx red-brown polymer: zigzag chains of sulfur atoms S8(s)  S8(l)  Melts at 113C T> 150C

  9. 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; S2- can’t be reduced can only be oxidized

  10. Compounds of S • H2SO4most 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

  11. 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

  12. 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 non metal Te semi-metal Po metal

  13. Nitrogen Nitrogen (N2) is very unreactive triple bond energy = 941kJ/mol Source fractional distillation of air (78% of air is N2) KNO3water soluble salts NaNO3found 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

  14. Compounds of Nitrogen Oxidation states of 3 to +5 Compounds with H • NH3 (3 oxidation state) • N2H4 (2 oxidation state) strong reducing agent: N2H4 N2(g) + 2 H2(g) H =  forms N2 readily: S = + 3. Dimethyl hydrazine (rocket fuel)

  15. 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

  16. 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)

  17. 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)

  18. 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)

  19. PHOSPHORUS ALLOTROPES White phosphorus (P4) burns spontaneously in air. P4(s) + 5O2(g)  P4O10 H = 3000 kJ/mole Red phosphorus (polymeric) is more stable. Not volatile. Does not react with air at 25°C.

  20. OXIDES OF PHOSPHORUS

  21. PHOSPHORUS OXYACIDS P4O10 + 6H2O  4H3PO4phosphoric acid P4O6 + 6H2O  4H3PO3 phosphorous acid Also H3PO2 hypophosphorous acid

  22. USES OF PHOSPHORUS Fertilizer P is essential for plant growth Ca3(PO4)2 + 3H2SO4 2H3PO4 + 3CaSO4 H3PO4 + 3NH3 (NH4)3PO4 Detergent Complexes metal ions Biological molecules (DNA, RNA) Biochemical energy source (ATP)

  23. COMPARISONS IN GROUP V NitrogenN2(g) NN NH3 is stable. Non-metal  oxides dissolve to give acidic solutions N2O3 + H2O 2HNO2 N2O5 + H2O  2HNO3 3NO2 + H2O  2HNO3 + NO

  24. 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

  25. ARSENIC Allotropes: Yellow As  As4 Gray As  brittle solid. AsH3ignites spontaneously in air. As4O10– acidic oxide: As4O10 + 6H2O  4H3AsO4 As4O6 is amphoteric, but is more soluble in base.

  26. 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.

  27. 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.

  28. GROUP V TRENDS Going down the periodic table: 1) Electronegativity decreases. 2) Switch from non-metallic to metallic. 3) Hydroxides and oxides become more basic. 4) Hydrides become less stable. 5)“Inert pair effect”becomes more pronounced: +3 becomes more stable as compared to +5.

  29. CARBON and Group IV • Carbon Sources: • Elemental form – coal. • Carbonate rocks (CO32-) • Limestone, marble, chalk = CaCO3 • Dolomite = MgCO3

  30. 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.

  31. 1) Ionic Carbides • Contain C4- or C22- (-CC-) • C4-: Be2C, Al4C3 react with water to give CH4. • C22-(-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. CARBIDES

  32. SILICON Second most abundant element. Found in combination with O. Silicate Minerals: [Si2O52-]n, SiO44- Sand: SiO2 (this is also quartz). With aluminum in aluminosilicates (clay, feldspars). Prepared by: SiO2(s) + 2C(s)  Si(l) + 2CO(g) (3000C) sand coke 98% Very pure silicon (<1 ppb impurity) is required for electronics applications.

  33. 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. GROUP IV TRENDS

  34. 2s Li Be 2p B C N O F 3s 3p Na Mg Al Si P S Cl Orbital Hybrids and Valence 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+, BH4- 3rd (and higher periods): can use d-orbitals to make bonds E.g. PF5 P atom is sp3d SF6 S atom is sp3d2

  35. 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

  36. 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 C=O (2 single + 1 double bond) H2N (urea) What about 3rd (and higher) periods - Si, P, S…? For these elements, double bonds are very uncommon (usually only single bonds)

  37. 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)

  38. 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)

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