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Air Earth Crust Oxygen on Earth H 2 O (oceans) O 2, CO 2 (atmosphere) CO 3  (rocks, coral, seashells) SiO 2, silicates (sand, clay, rocks) Made commercially.

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Presentation on theme: "Air Earth Crust Oxygen on Earth H 2 O (oceans) O 2, CO 2 (atmosphere) CO 3  (rocks, coral, seashells) SiO 2, silicates (sand, clay, rocks) Made commercially."— Presentation transcript:

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

2 ALLOTROPES OF OXYGEN Ozone (O 3 ) is a strong oxidizing agent, highly toxic Kills bacteria (replacement for Cl 2 in municipal water treatment) Irritating component of photochemical smog O 2 Paramagnetic (why?) O 3 Higher energy form - important UV absorber in the stratosphere Light or electrical discharge 3O 2 2 O 3 decomposition

3 OXYGEN IONS Oxide Ion  O 2  (most compounds) e.g. Li 2 O = 2Li + O 2  Peroxide Ion  O 2 2  =  O – O  e.g. Na 2 O 2 = 2 Na +  O – O  Also, H 2 O 2 (hydrogen peroxide) Superoxide Ion  O 2  e.g. KO 2 = K + O 2  Can have positive oxidation states in combination with fluorine + 2 in OF 2

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) H 2 O 2 decomposition can be explosive: 2 H 2 O 2  2 H 2 O + O 2  H =  200 kJ/mol (disproportionation reaction)

5 HYDROGEN PEROXIDE Reduction to H 2 O: H 2 O 2 + 2H + + 2I   I 2 + 2H 2 O Oxidation to O 2 : 2MnO 4  + 5 H 2 O 2 + 6H +  2Mn O 2 + 8H 2 O

6 SULFUR Sources: Sulfide Minerals (S 2- ): FeS 2 (Pyrite) -Iron Ore (Fool’s Gold) Cu 3 FeS 3 (Bornite) -Source of Cu PbS (Galena) - Source of Pb ZnS (Zinc Blende) -Source of Zn Sulfate Minerals (SO 4 2- ) e.g. Na 2 SO 4, MgSO 4 Also, CaSO 4 · (H 2 O) 2 (Gypsum) Used for wallboard, plaster of Paris.

7 COMMERCIAL SOURCES OF SULFUR 1) Sulfur Mines – Along Gulf of Mexico, deposits of S 8  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%. SO 2 forms when coal is burned. SO 2 + H 2 O  H 2 SO 3 Acid Rain SO 2 +[O]  SO 3 +H 2 O  H 2 SO 4 CaCO 3 + H 2 SO 4  CaSO 4 + H 2 O + CO 2 (Marble)

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

10 Sulfur Common Oxidation States +6SO 3 ; H 2 SO 4 (sulfuric acid) can’t be oxidized can only be reduced +4SO 2 ; H 2 SO 3 (sulfurous acid) can be both oxidized AND reduced -2H 2 S; S 2- can’t be reduced can only be oxidized

11 Compounds of S H 2 SO 4 most important industrial chemical H 2 S(rotten egg smell) (S 2- ) source: metal sulfides + strong acid e.g. ZnS + HCl  ZnCl 2 + H 2 S(g) –poisonous –tarnishes Ag in presence of O 2 4Ag + 2H 2 S + O 2  2Ag 2 S + 2H 2 O – Organic sulfides, e.g. C 4 H 9 SH Strong odor – added to natural gas S 2 O 3 2- (thiosulfate) used in photography: forms water soluble complexes with Ag

12 Uses of H 2 SO 4 Making phosphate fertilizer Ca 3 (PO 4 ) 2 + 3H 2 SO 4  3CaSO 4 + 2H 3 PO 4 ~65% of H 2 SO 4 Manufacture of chemicals Metal refining Petroleum refining (as catalyst) Strong oxidizing agent Drying agent

13 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) Senon metal Tesemi-metal Pometal

14 Nitrogen Nitrogen (N 2 ) is very unreactive triple bond energy = 941kJ/mol Source fractional distillation of air (78% of air is N 2 ) KNO 3 water soluble salts NaNO 3 found in deserts Nitrogen fixation: formation of N containing compounds from N 2 N is an essenial element in proteins, nucleic acids & necessary to maintain soil fertility

15 Compounds of Nitrogen Oxidation states of  3 to +5 Compounds with H 1.NH 3 (  3 oxidation state) 2.N 2 H 4 (  2 oxidation state) strong reducing agent: N 2 H 4  N 2 (g) + 2 H 2 (g)  H =  forms N 2 readily:  S = + 3.Dimethyl hydrazine (rocket fuel)

16 N compounds with oxygen N 2 O colorless, odorless gas used as anesthetic (laughing gas) propellant in whipping cream NO formed in car engines: N 2 +O 2  2NO N 2 O 3 blue solid, decomposes: N 2 O 3  NO + NO 2 NO 2 brown gas; component of smog N 2 O 4 2NO 2  N 2 O 4 N 2 O 5 unstable, decomposes to NO 2

17 HNO 3 Produced from NH 3 by Ostwald process (catalytic oxidation). Uses: fertilizer NH 3 + HNO 3  NH 4 NO 3 (s) strong acid strong oxidizing agent. cleaning agent to make explosives (e.g. nitroglycerine, TNT)

18 Hydrolysis of oxides Hydrolysis: reaction with water N is a non metal: oxides are acidic. Oxide + H 2 O = hydroxide N 2 O 3 + H 2 O  2HNO 2 (nitrous acid) 3NO 2 + H 2 O  2HNO 3 + NO N 2 O 5 + H 2 O  2HNO 3 (nitric acid)

19 PHOSPHORUS Source: Phosphate Minerals Ca 3 (PO 4 ) 2 contains PO 4 3- (tetrahedral P) P is made by heating Ca 3 (PO 4 ) 2 and coke in an electric furnace. 2Ca 3 (PO 4 ) 2 (s) + 10C(s) + 6SiO 2  6CaSiO 3 (s) + 10CO(g) + P 4 (g)

20 White phosphorus (P 4 ) burns spontaneously in air. P 4 (s) + 5O 2 (g)  P 4 O 10  H =  3000 kJ/mole Red phosphorus (polymeric) is more stable. Not volatile. Does not react with air at 25°C. PHOSPHORUS ALLOTROPES

21 OXIDES OF PHOSPHORUS

22 PHOSPHORUS OXYACIDS P 4 O H 2 O  4H 3 PO 4 phosphoric acid P 4 O 6 + 6H 2 O  4H 3 PO 3 phosphorous acid Also H 3 PO 2 hypophosphorous acid

23 USES OF PHOSPHORUS Fertilizer P is essential for plant growth Ca 3 (PO 4 ) 2 + 3H 2 SO 4  2H 3 PO 4 + 3CaSO 4 H 3 PO 4 + 3NH 3  (NH 4 ) 3 PO 4 Detergent Complexes metal ions Biological molecules (DNA, RNA) Biochemical energy source (ATP)

24 COMPARISONS IN GROUP V NitrogenN 2 (g)N  N NH 3 is stable. Non-metal  oxides dissolve to give acidic solutions N 2 O 3 + H 2 O  2HNO 2 N 2 O 5 + H 2 O  2HNO 3 3NO 2 + H 2 O  2HNO 3 + NO

25 PHOSPHORUS Allotropes: White P  P 4, tetrahedral. Red P  polymer. PH 3 burns in air. Non-metal  oxides dissolve to give acidic solutions: P 4 O H 2 O  4H 3 PO 4 P 4 O 6 + 6H 2 O  4H 3 PO 3

26 ARSENIC Allotropes: Yellow As  As 4 Gray As  brittle solid. AsH 3 ignites spontaneously in air. As 4 O 10 – acidic oxide: As 4 O H 2 O  4H 3 AsO 4 As 4 O 6 is amphoteric, but is more soluble in base.

27 ANTIMONY – Sb Brittle gray metalloid. Sb 4 O 6 is amphoteric. There is no Sb 4 O 10. Bismuth is a metal. Bi 4 O 6 is basic and dissolves only in acids. Bi(OH) 3 is basic. Bi 5+ is rare. BISMUTH - Bi

28 OXIDATION STATES P 5+ dominates. As 3+, As 5+ are equally common. Sb 3+ dominates. Bi 3+ dominates. Inert Pair Effect HYDRIDE STABILITY NH 3 is stable. PH 3 is stable but burns in air. AsH 3 decomposes easily. SbH 3, BiH 3 are very unstable.

29 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.

30 CARBON and Group IV Carbon Sources: 1)Elemental form – coal. 2)Carbonate rocks (CO 3 2- ) Limestone, marble, chalk = CaCO 3 Dolomite = MgCO 3

31 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 C 60 which has the shape of a soccer ball. Carbon Black – Soot Amorphous form of carbon used in tires, inks, pigments, and carbon paper.

32 CARBIDES 1)Ionic Carbides Contain C 4- or C 2 2- (-C  C-) C 4- : Be 2 C, Al 4 C 3 react with water to give CH 4. C 2 2- (-C  C-): CaC 2 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.

33 SILICON Second most abundant element. Found in combination with O. Silicate Minerals: [Si 2 O 5 2- ] n, SiO 4 4- Sand: SiO 2 (this is also quartz). With aluminum in aluminosilicates (clay, feldspars). Prepared by: SiO 2 (s) + 2C(s)  Si(l) + 2CO(g) (3000  C) sand coke 98% Very pure silicon (<1 ppb impurity) is required for electronics applications.

34 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 (PbO 2 ) 2)Basicity of oxides and hydroxides increases. CO 2, SiO 2, GeO 2 are weakly acidic. SnO, SnO 2, PbO are amphoteric. 3)Hydrides become less stable. Enormous number of stable hydrocarbons. SiH 4 is stable but is spontaneously flammable. Ge, Sn, Pb hydrides are very unstable.

35 Orbital Hybrids and Valence Li Na Be Mg B Al C Si N P O S F Cl 2p 3p 2s 3s 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: CH 4, NF 4 +, BH 4 - 3rd (and higher periods): can use d-orbitals to make bonds E.g. PF 5 P atom is sp 3 d SF 6 S atom is sp 3 d 2

36 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 (NH 3, NCl 3, CH 3 NH 2 - 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., NO 3 - N=O(4 bonds to N) O NO 2 + O=N=O (4 bonds to N, like CO 2 ) Likewise, O usually makes 2 bonds: H 2 O, OF 2, H 2 C=O

37 Likewise, C can gain 4 or lose 4 electrons to make an octet (valence = 4) So carbon always makes 4 bonds CH 4 (4 single bonds) O=C=O(2 double bonds) H-C  C-H(1 single + 1 triple bond) H 2 N C=O (2 single + 1 double bond) H 2 N (urea) What about 3rd (and higher) periods - Si, P, S…? For these elements, double bonds are very uncommon (usually only single bonds)

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39 So CO 2 is molecular (O=C=O, has double bonds) But SiO 2 (quartz, sand, glass…) is a 3-dimensional solid network: O 2 is molecular (O=O, has a double bond) But S forms rings (e.g., S 8 )

40 Nitrogen (N 2 ) 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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