Oxidation & Reduction Electrochemistry

1. Oxidation & ReductionElectrochemistry BLB 11th Chapters 4, 20

2. Chapter Summary • Oxidation and Reduction (redox) – introduced in chapter 4 • Oxidation Numbers • Electron-transfer • Balancing redox reaction • Electrochemical cells • Corrosion • Electrolysis

3. 20.1, 4.4 Oxidation-Reduction Reactions • Oxidation • Loss of electrons • Increase in oxidation number • Gain of oxygen or loss of hydrogen • Reduction • Gain of electrons • Decrease in oxidation number • Loss of oxygen or gain of hydrogen Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)

4. Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s) • Oxidizing agent or oxidant – reactant that contains the element being reduced; is itself reduced • Reducing agent or reductant – reactant that contains the element being oxidized; is itself oxidized

5. Oxidation Numbers (p. 132) Assign according to the following order: • Atoms zero (since neutral) • Ions equal to charge of the ion • Nonmetals • O −2 • H +1 (when bonded to other nonmetals) −1 (when bonded to metals) • F −1 • X −1 except when combined with oxygen Sum of the oxidation numbers equals zero or the charge of the polyatomic ion.

6. O2 CH4 NO3¯ CH3OH Cr2O72- CH2O Cu2+ OCl¯ Oxidation numbers practice

7. Redox Reactions • Combustion, corrosion, metal production, bleaching, digestion, electrolysis • Metal oxidation • Activity Series (Table 4.5, p. 136) • Some metals are more easily oxidized and form compounds than other metals. • Displacement reaction – metal or metal ion is replaced through oxidation A + BX → AX + B

9. 20.2 Balancing Redox Reactions • Goal: Balance both the atoms and the electrons • Examples: Al(s) + Zn2+(aq) → Al3+(aq) + Zn(s) MnO4¯(aq) + Cl¯(aq) → Mn2+(aq) + Cl2(g)

10. The Rules (p. 830-1) In acidic solution: • Divide equation into two half-reactions (ox and red). • Balance all elements but H and O. • Balance O by adding H2O. • Balance H by addingH+. • Balance charge by adding electrons (e-). • Cancel out electrons by integer multiplication. • Add half reactions & cancel out. • Check balance of elements and charge.

11. MnO4¯(aq) + Cl¯(aq) → Mn2+(aq) + Cl2(g)

12. CH3OH(aq) + Cr2O72-(aq) → CH2O(aq) + Cr3+(g)

13. The Rules (p. 833) In basic solution: Proceed as for acidic solution through step 7. • Add OH¯ to neutralize the H+. (H+ + OH¯ → H2O) • Cancel out H2O. • Check balance of elements and charge.

14. Cr(s) + CrO4¯(aq) → Cr(OH)3(aq)

15. 20.3 Voltaic Cells • A spontaneous redox reaction can perform electrical work. • The half-reactions must be placed in separate containers, but connected externally. • This creates a potential for electrons to flow. • Reactant metal is the most reactive; product metal the least. Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s) Line notation: Zn(s)|Zn2+(aq)||Cu2+(aq)|Cu(s)

16. 20.3 Voltaic Cell Net reaction: Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)

17. Movement of Electrons Zn(s) → Zn2+(aq) + 2 e¯ Cu2+(aq) + 2 e¯ → Cu(s) e¯ Net reaction: Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)

18. 20.4 Cell Potentials Under Standard Conditions • EMF – electromotive force – the potential energy difference between the two electrodes of a voltaic cell; Ecell; measured in volts • E°cell – standard cell potential (or standard emf) • For the Zn/Cu cell, E°cell = 1.10 V • electrical work = Coulombs x volts J = C x V

19. Standard Reduction (Half-cell) Potentials • E° - potential of each half-cell • E°cell = E°cell(cathode) -E°cell(anode) • For a product-favored reaction: • ΔG° < 0 • E°cell > 0 • Measured against standard hydrogen electrode (SHE); assigned E° = 0 V.

22. App. E, p. 1064 More E° values

24. Problem Voltaic cell with: Al(s) in Al(NO3)3(aq) on one side and a SHE on the other. Sketch the cell, determine the balance equation, and calculate the cell potential.

25. Problem Voltaic cell with: Pb(s) in Pb(NO3)2(aq) on one side and a Pt(s) electrode in NaCl(aq) with Cl2 bubbled around the electrode on the other. Sketch the cell, determine the balance equation, and calculate the cell potential.

26. 20.5 Free Energy and Redox Reactions • ΔG° < 0 • E°cell > 0 ΔG° for previous problems

28. 20.6 Cell Potentials Under Nonstandard Conditions • Concentrations change as a cell runs. • When E = 0, the cell is dead and reaches equilibrium. • Nernst equation allows us to calculate E under nonstandard conditions:

29. Concentration Cells • A cell potential can be created by using same half-cell materials, but in different concentrations. Problem 69

30. Problem 69

31. Cell EMF and Equilibrium • When E = 0, no net change in flow of electrons and cell reaches equilibrium. K of previous problems

32. 20.7 Batteries and Fuel Cells Batteries • self-contained electrochemical power source • More cells produce higher potentials • Primary – non-rechargeable (anode/cathode) • Alkaline: Zn in KOH/MnO2 • Secondary – rechargeable (anode/cathode) • Lead-acid: Pb/PbO2 in H2SO4 • nicad: Cd/[NiO(OH)] • NiMH: ZrNi2/[NiO(OH)] • Li-ion: C(s,graphite)/LiCoO2

34. Hydrogen Fuel Cells • Convert chemical energy directly into electricity • Fuel and oxidant supplied externally continuously • Products are only electricity and water cathode: O2(g) + 4 H+(aq) + 4 e¯ → 2 H2O(l) anode: 2 H2(g) → 4 H+(aq) + 4 e¯ overall: 2 H2(g) + O2(g) → 2 H2O(l)

35. PEM Fuel Cell ►

36. 20.8 Corrosion • RUST! • Anode: M(s) → Mn+(aq) + n e¯ • Cathode: O2(g) + 4 H+(aq) + 4 e¯ → 2 H2O(l) or: O2(g) + 2 H2O(l) + 4 e¯ → 4 OH¯ (aq)

37. Preventing Corrosion • Anionic inhibition • painting • oxide formation • coating • Cathodic inhibition • sacrificial anode – attach a metal (like Mg) more easily oxidized • galvanizing steel – coating with zinc

38. 20.9 Electrolysis • Electrical energy  chemical change

39. Hall-Héroult Process for Al Production C(s) + 2 O2-(l) → CO2(g) + 4 e¯ 3 e¯ + Al3+(l) → Al(l)