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Electrochemistry and Society

Electrochemistry and Society. Corrosion = oxidation of pure metals to their oxides Corrosion Basics Metals (M o ) are easily oxidized to cationic forms (M n+ ) [Table 18.1] e o 1/2 of O 2 gas reduction > oxidation of most metals

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Electrochemistry and Society

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  1. Electrochemistry and Society • Corrosion = oxidation of pure metals to their oxides • Corrosion Basics • Metals (Mo) are easily oxidized to cationic forms (Mn+) [Table 18.1] • eo1/2 of O2 gas reduction > oxidation of most metals • O2 + 4H+ + 4e- 2H2O eo1/2 = +1.23 V • This leads to an eocell that is positive for this process = spontaneous • Mo + O2 MxOyeocell = + • Most metals don’t completely decompose because MxOy protects the vulnerable Mo underneath from further corrosion • Aluminum Example • Al3+ + 3e- Aloeo = -1.70 V • O2 + 4H+ + 4e- 2H2O eo = +1.23 V • Al2O3 produced has eo1/2 = -0.60 V, resulting in a much less favorable corrosion process, once the aluminum underneath is covered. eocell = +2.93 V

  2. The “Noble Metals” (Ag, Au, Cu, Pt) do not react with oxygen as easily • Auoeo1/2 = +1.50 V no corrosion • Agoeo1/2 = +0.80 V Ag2S tarnish formed instead of the oxide • Cuoeo1/2 = +0.16 V Cu2CO3 forms green “patina” • The Corrosion of Iron • This is the most economically important corrosion process due to structural steel • Steel has a non-uniform surface due to physical stress • Anodic Region: Fe Fe2+ + 2e- • Cathodic Region: O2 + 2H2O + 4e- 4OH- • Fe2+ then acts as the salt bridge electrolyte if wet (added salt speeds up corrosion) Cathode: 4Fe2+ + O2 (4 + 2n)H2O 2Fe2O3• nH2O + 8H+

  3. Preventing Corrosion • Paint covers the surface to prevent the contact of oxygen and the metal • Plating steel with Cr or Sn to produce very stable oxides • Galvanizing = coating with Zinc • Fe Fe2+ + 2e- -eo1/2 = +0.44 V • Zn Zn2+ + 2e- -eo1/2 = +0.76 V • Corrosion occurs on Zn rather than Fe (sacrificial coating) • Stainless Steel = Fe + Cr + Ni -eo1/2 ~ Noble metal • Cathodic Protection = protects buried steel or ships with a sacrificial reactant • Active metal (Mg) connected to pipe by a wire -eo1/2 = +2.37 V • Bars of Ti attached to ship -eo1/2 = +1.60 V

  4. Electrolysis = using electric energy to produce chemical change (opposite of cell) • Example • Consider the Cu/Zn Galvanic Cell • Anode: Zn Zn2+ + 2e- • Cathode: Cu2+ + 2e- Cu eocell = +1.10 V • If we attach a power source of eo > +1.10 V, we can force e- to go the other way • Anode: Cu Cu2+ + 2e- • Cathode: Zn2+ + 2e- Zn • Called an Electrolytic Cell

  5. B. Calculations with Electrolytic Cells • How much Chemical Change? Is usually the question. • Find mass of Cuo plated out by passing 10 amps (10 C/s) through Cu2+ solution. • Cu2+ + 2e- Cuo(s) • Steps: current/time, charge (C), moles e-, moles Cu, grams Cu • Example: How long must a current of 5.00 amps be applied to a Ag+ solution to produce 10.5 g of silver metal? • Electrolysis of Water • Galvanic: 2H2 + O2 2H2O (Fuel Cell) • Electrolytic Cell: • Anode: 2H2O O2 + 4H+ + 4e- -eo = -1.23 V • Cathode: 4H2O + 4e- 2H2 + 4OH-eo = -0.83 V • Overall: 6H2O 2H2 + O2 + 4(H+ + OH-) 2H2O 2H2 + O2eocell = -2.06 V • We must add a salt to increase the conductance of pure water [H+] = [OH-] = 10-7

  6. Electrolysis of Mixtures • Mixture of Cu2+, Ag+, Zn2+; What is the order of plating out? • Ag+ + e- Ag eo1/2 = +0.80 V • Cu2+ + 2e- Cu eo1/2 = +0.34 V • Zn2+ + 2e- Zn eo1/2 = -0.76 V • Reduction of Ag+ is easiest (eo = most positive) followed by Cu, then Zn • Example: Ce4+ (eo1/2 = +1.70 V), VO2+ (eo1/2 = +1.00 V), Fe3+ (eo1/2 = +0.77 V) • Commercial Electrolytic Processes • Production of Aluminum • Most metals are found naturally as their oxides, MxOy • Only the noble metals are typically found as the pure metal • Bauxite = aluminum ore; Al2O3 • Aluminum is the third most abundant element on crust (oxygen and silicon) • No commercial process for pure Al until 1854 (eo1/2 = 1.66 V) • Al was more valuable than gold or silver

  7. Hall—Heroult Process • Al3+ + 3 e- Al eo1/2 = -1.66 V • 2H2O + 2e- H2 + 2OH-eo1/2 = -0.83 V • Can’t make Al in water because water gets reduced before Al3+ • Use molten Al2O3/Na3AlF6 mixture at 1000 oC • Aluminum alloys with Zn, Mn are most useful because they are stronger • Aluminum production uses 5% of the electricity consumed in the U.S.

  8. Electrorefining = purification of metals • Impure Cuo anode: Cuo Cu2+ + 2e- • Pure Cuo cathode: Cu2+ + 2e- Cuo (99.95% pure) • Also useful for purification of Zn, Fe • Gold, Silver, Platinum fall to the bottom of the tank as sludge (won’t plate out) • Metal Plating • Coat easily corrodable metal object with a noble metal • Ag+ + e- Ago on a spoon • Electrolysis of NaCl • Production of Na metal from NaCl in a “Downs Cell” • Anode: 2Cl- Cl2 + 2e- • Cathode: Na+ + e- Nao • Cell is designed to to keep products apart so they can’t reform NaCl • Production of Cl2, OH- in a Mercury Cell • Water is reduced to OH- (eo1/2 = -0.83 V) before Na+ (eo1/2 = -2.71 V) • Anode: 2Cl- Cl2 + 2e- • Cathode: 2H2O + 2e- H2 + 2OH- • Chlor-Alkali Process = second largest electricity user in U.S. (after Al)

  9. Downs Cell

  10. Mercury Cell for Chor-Alkali Process

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