17.1 Redox Chemistry Revisited

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Chapter Outline 17.1 Redox Chemistry Revisited 17.2 Electrochemical Cells 17.3 Standard Potentials 17.4 Chemical Energy and Electrical Work 17.

8 Electrolytic Cells and Rechargeable Batteries 17.9 Fuel Cells 19-1 1 Redox Chemistry Revisited (Ch 8, section 8.6, p.

1 Chapter Outline 17.1 Redox Chemistry Revisited 17.2 Electrochemical Cells 17.3 Standard Potentials 17.4 Chemical Energy and Electrical Work 17.5 A Reference Point: The Standard Hydrogen Electrode 17.6 The Effect of Concentration on E cell 17.7 Relating Battery Capacity to Quantities of Reactants 17.8 Electrolytic Cells and Rechargeable Batteries 17.9 Fuel Cells Redox Chemistry Revisited (Ch 8, section 8.6, p. 332 Oxidation-Reduction Reactions (Redox): Characterized by gain or loss of electrons by atoms involved in the reaction. Oxidation: Historical definition = Modern definition = Reduction: Historical definition = Modern definition = 1

Table 8.4 - Oxidation Number Rules The charge the atom would have in a molecule (or an ionic compound) if electrons were completely transferred, i.e. applies to both ionic and covalent compouinds. 1.

2 Table Oxidation Number Rules The charge the atom would have in a molecule (or an ionic compound) if electrons were completely transferred, i.e. applies to both ionic and covalent compouinds. 1. O.N. = 0 for atoms in pure elements. 2. O.N = the charge on monovalent ions 3. O.N. of fluorine = -1 for all of its compounds Table Oxidation Number Rules 4. O.N. of oxygen = -2 in nearly all of its compounds 5. O.N. of hydrogen = +1 in nearly all of its compounds 6. O.N. values of the atoms in a neutral molecule sum up to zero 7. O.N. values of the atoms in a polyatomic ion sum up to the charge on the ion 2

Half Reactions Zn(s) + Cu 2+ (aq) Cu(s) + Zn 2+ (aq) Chapter Outline 17.1 Redox Chemistry Revisited 17.2 Electrochemical Cells 17.3 Standard Potentials 17.4 Chemical Energy and Electrical Work 17.

3 Half Reactions Zn(s) + Cu 2+ (aq) Cu(s) + Zn 2+ (aq) Chapter Outline 17.1 Redox Chemistry Revisited 17.2 Electrochemical Cells 17.3 Standard Potentials 17.4 Chemical Energy and Electrical Work 17.5 A Reference Point: The Standard Hydrogen Electrode 17.6 The Effect of Concentration on E cell 17.7 Relating Battery Capacity to Quantities of Reactants 17.8 Electrolytic Cells and Rechargeable Batteries 17.9 Fuel Cells

Electrochemical (Galvanic or Voltaic) Cells The difference in electrical potential between the anode and cathode is called: e- e- cell voltage or potential (Volts) electromotive force (E) Cell

4 Electrochemical (Galvanic or Voltaic) Cells The difference in electrical potential between the anode and cathode is called: e- e- cell voltage or potential (Volts) electromotive force (E) Cell Diagram Zn (s) + Cu 2+ (aq) Cu (s) + Zn 2+ (aq) [Cu 2+ ] = 1 M & [Zn 2+ ] = 1 M < Zn (s) Zn 2+ (1 M) Cu 2+ (1 M) Cu (s) > Anode (-) Cathode (+) The circuit is completed via the Salt Bridge e- e- Common salt bridge = Na 2 SO 4 SO 4 2- Na + gaining (+) charge losing (+) charge 4

5 Chapter Outline 17.1 Redox Chemistry Revisited 17.2 Electrochemical Cells 17.3 Standard Potentials 17.4 Chemical Energy and Electrical Work 17.5 A Reference Point: The Standard Hydrogen Electrode 17.6 The Effect of Concentration on E cell 17.7 Relating Battery Capacity to Quantities of Reactants 17.8 Electrolytic Cells and Rechargeable Batteries 17.9 Fuel Cells 19-9 The cell voltage is the difference in potential between the cathode and the anode: E 0 = E 0 cell cathode - E0 anode cathode: Cu 2+ (aq) + 2e - Cu(s) E o cathode anode: Zn 2+ (aq) + 2e - Zn(s) E o anode E o cathode and E o anode are called Standard Reduction Potentials; measured and tabulated (Table A6.1) Measured under Standard Conditions = 1 atm, 1.0 M, 298 K 5

6 E 0 = E 0 cell cathode - E 0 anode e- e- Cu 2+ (aq) + 2e - Cu(s) Zn 2+ (aq) + 2e - Zn(s) Sign conventions: E > 0 spontaneous E = 0 equilibrium E < 0 nonspontaneous Standard Reduction Potentials at 298 K F 2 (g) + 2 e - 2 F - (aq) V 2 H 3 O + (aq) + 2 e - H 2 (g) + 2 H 2 O(l) 0.00 V Li + (s) + e - Li(aq) V 6

reactions are written as a reduction: E 0 red the

the substance to be reduced strong oxidizing agents

half-cell reactions are reversible the sign of E 0

7 reactions are written as a reduction: E 0 red the more positive E 0 is, the greater the tendency for the substance to be reduced strong oxidizing agents at the top strong reducing agents at the bottom the half-cell reactions are reversible the sign of E 0 changes when the reaction is reversed = oxidizing potential 7

The Zinc-Air Battery Cell potentials when the number of electrons transferred is different for each half reaction - Changing the stoichiometric coefficients

8 The Zinc-Air Battery Cell potentials when the number of electrons transferred is different for each half reaction - Changing the stoichiometric coefficients of a half-cell reaction does not change the value of E 0 Anode: Zn(s) + 2 OH - (aq) ZnO(s) + H 2 O(l) + 2e - Cathode: O 2 (g) + 2H 2 O(l) + 4e - 4OH - (aq) 8

9 Chapter Outline 17.1 Redox Chemistry Revisited 17.2 Electrochemical Cells 17.3 Standard Potentials 17.4 Chemical Energy and Electrical Work 17.5 A Reference Point: The Standard Hydrogen Electrode 17.6 The Effect of Concentration on E cell 17.7 Relating Battery Capacity to Quantities of Reactants 17.8 Electrolytic Cells and Rechargeable Batteries 17.9 Fuel Cells Chemical Energy and Electrical Work G = -nfe cell Faraday constant (F) is C/(mol e - ) n = number of moles of electrons 9

10 Button Batteries 10

11 Chapter Outline 17.1 Redox Chemistry Revisited 17.2 Electrochemical Cells 17.3 Standard Potentials 17.4 Chemical Energy and Electrical Work 17.5 A Reference Point: The Standard Hydrogen Electrode 17.6 The Effect of Concentration on E cell 17.7 Relating Battery Capacity to Quantities of Reactants 17.8 Electrolytic Cells and Rechargeable Batteries 17.9 Fuel Cells A Reference Point: The Standard Hydrogen Electrode 2 H 3 O + (aq) + 2 e - H 2 (g) + 2 H 2 O(l) 0.00 V H+ (1.00 M) H2(g, 1.00 atm) Pt 11

12 12

13 Chapter Outline 17.1 Redox Chemistry Revisited 17.2 Electrochemical Cells 17.3 Standard Potentials 17.4 Chemical Energy and Electrical Work 17.5 A Reference Point: The Standard Hydrogen Electrode 17.6 The Effect of Concentration on E cell 17.7 Relating Battery Capacity to Quantities of Reactants 17.8 Electrolytic Cells and Rechargeable Batteries 17.9 Fuel Cells The Effect of Concentration on E cell The Nernst Equation for aa + bb = cc + dd, from Thermodynamics we know - G = G o + RT lnq and G = nfe 13

14 The Lead-Acid Battery The Lead-Acid Battery Cathode: Anode: 14

15 The Lead-Acid Battery Both cells kept at 4.5 M H 2 SO 4 15

16 E o and K aa + bb = cc + dd E E o log Q n at 25 o C at equilibrium - 16

17 Chapter Outline 17.1 Redox Chemistry Revisited 17.2 Electrochemical Cells 17.3 Standard Potentials 17.4 Chemical Energy and Electrical Work 17.5 A Reference Point: The Standard Hydrogen Electrode 17.6 The Effect of Concentration on E cell 17.7 Relating Battery Capacity to Quantities of Reactants 17.8 Electrolytic Cells and Rechargeable Batteries 17.9 Fuel Cells

18 Electrolytic Cells and Rechargeable Batteries Electrolysis is the process in which electrical energy is used to cause a nonspontaneous chemical reaction to occur. Electrolysis of NaCl to Produce Na Metal Anode: Cl 2 (g) + 2 e 2 Cl - (aq) E o = Cathode: Na + (aq) + e Na(s) E o = 18

19 Electrolysis and Mass Calculations charge (C) = current (A) x time (s) 1 mole e - = 96,500 C 19

Assigning Oxidation Numbers:

Assigning Oxidation Numbers: 1. Oxidation number of a free element or diatomic molecule is zero. Ex: Na(s), Cu(s), H 2 (g), F 2 (g) 2. In most cases the oxidation number of hydrogen is +1, oxygen is -2,