CHEM 163 Chapter 21

1. CHEM 163Chapter 21 Spring 2009

2. 3-minute review • What is a redox reaction?

3. Half-Reactions Split overall reaction into two reactions Oxidation Reduction Step 1. Divide reaction into half reactions. Step 2. Balance atoms in each half reaction. Do O and H last! Need O? Add H2O Need H? Add H+ Step 3. Balance charges in each half reaction. Add e- Step 4. Make # e- gained equal # e- lost. Multiply by integer! Step 5. Add reactions together. Step 6. Check that atoms and charges are balanced.

4. Electrochemical Cells Voltaic (Galvanic) Cell Electrolytic Cell ∆G < 0 ∆G > 0 Erxt > Eprod Elost electricity Erxt < Eprod Electricity  rxn Sys does work on surr Surr do work on sys • Electrodes: • Conduct electricity between cell and surroundings • Anode (oxidation) • Cathode (reduction) • Electrolyte: contains ions

5. Fig. 21.3

6. Voltaic Cells Half-cells: to complete the circuit, electrons must flow externally Zn (s) + Cu2+ (aq)  Zn2+ (aq) + Cu (s) • Oxidation half-cell: • Anode (Zn) • “reactant” • Electrolyte • Reduction half-cell: • Cathode (Cu) • “product” • Electrolyte Fig. 21.5

7. Zn (s) + Cu2+ (aq)  Zn2+ (aq) + Cu (s) Voltaic Cells • Electrode charges: • e- flow left to right • e- created at anode, used up at cathode • Anode has excess e- Anode Cathode - + • Salt bridge: • Completes circuit • Keeps each cell neutral • Direction of ions Anode Cathode

8. Electrodes • Conduct electricity between cell and surroundings Active Electrodes: electrodes are components of half-reactions Inactive Electrodes: conduct electrons but are not reactants or products Ex. Graphite, Pt 2I- (aq)  I2 (s) + 2e- MnO4- (aq) + 8H+ (aq) + 5e- Mn2+ (aq) + 4H2O (l) Anode Cathode

9. How much electricity? • Zn gives up electrons more easily • Zn is a stronger reducing agent • Potential difference between two electrodes • Cell potential (Ecell) • Cell voltage • Electromotive force (emf) Zn (s) + Cu2+ (aq)  Zn2+ (aq) + Cu (s) Zn (s) Zn2+ (aq) + 2 e- Cu (s) Cu2+ (aq) + 2 e- Ecell > 0 (spontaneous process)

10. Standard Cell Potentials • Ecell at standard conditions • Specific T (usually 298 K) • All components in standard states • 1 M (aq) • 1 atm (g) • Pure solid • Standard Electrode Potential • Half-cell potential • Always shown as a reduction reduction oxidation

11. How can you measure a half-cell? • Half-cell potentials are relative to a standard Standard Hydrogen Electrode (SHE) 2H+ (aq; 1 M) + 2e- H2 (g; 1 atm) • Stronger oxidizing agents… • are easily reduced themselves • Reduction reaction occurs more easily • have more positive Eo • are weaker reduction agents M+(aq) + e- M (s)

12. Writing spontaneous redox reactions 2 Ag+ (aq) + Sn (s) 2 Ag (s) + Sn2+ (aq) ? • Which is the oxidizing agent? • Write reduction rxn for oxidizing agent (incl. Eo) • Flip oxidation rxn for reducing agent (incl. -Eo) • Multiply to make # e- lost = # e- gained • Add together Ag+ (aq) + e- Ag (s) Sn2+ (aq) + 2e- Sn (s) Ag Eo value does not change!

13. Activity Series of Metals • Metals that can displace H2 from acid • Ecell is positive for reaction with H+ • Any negative Ehalf-cell (reduction potential) • Metals that cannot displace H2 from acid • Ecell is negative for reaction with H+ • Any positive Ehalf-cell (reduction potential) • Metals that can displace H2 from water • Ecell is positive for reaction with water

14. How much Work? Volt Coulomb Joule electrical potential electrical charge energy Max work: How much charge flows? Faraday constant # mols of e- transferred Charge of 1 mol of e- = 96,485 C / mol e- = 96,485 J/V mol e- (standard state)

15. Spontaneous At equilibrium Nonspontaneous

16. Effect of Concentration on Ecell Nernst Equation (at 298 K)

17. Concentration Cells Cells with different concentrations of same half-reaction ? Not standard conditions!

18. Primary Batteries Nonrechargable • Alkaline Zn (s) + MnO2 (s) + H2O (l) ZnO (s) + Mn(OH)2 (s) • Mercury and Silver • Zn anode; Hg or Ag cathode • Steady output • Primary Lithium Batteries • High energy/mass ratio • Lithium metal anode • Implanted medical devices, watches E = 1.5 V

19. Secondary Batteries Rechargeable Reverse reaction using electricity • Lead-Acid PbO2 (s) + Pb (s) + 2H2SO4 (aq)  2 PbSO4 (s) + 2 H2O (l) Ecell = 2.1 V • Nickel-Metal Hydride (Ni-MH) • Lithium-Ion • Anode: Li atoms between graphite sheets • Cathode: Lithium metal oxide

20. Corrosion Natural redox: metal  metal oxides and metal sulfides Anodic regions: • Dents, ridges • Iron loss Cathodic regions: • Surface • Forms water Fe2+ reacts with O2: • Rust deposits

21. Electrolytic Cells electrical energy  nonspontaneous reaction Ecell < 0 • oxidation at anode • reduction at cathode • anode is positive • cathode is negative

22. Electrolysis • Splitting a substance using electrical energy • Way to harvest elements (for industrial use) from substances What types of substances? • Pure molten salts • Isolate metal or nonmetal • Mixed molten salts • Isolate more easily reduced metal (based on EA)

23. Electrolysis of Water Anode Cathode Net Not at standard state: [H+] = [OH-] = 10-7 M Ecell determined using Nernst equation:

24. Electrolysis of Aqueous Salts Which is going to react: water or salt? • Reduction with less negative Eelectrode occurs • Oxidation with less positive Eelectrode occurs Example: KI (aq) Reduction: Oxidation: H2 forms at cathode I2 forms at anode

25. Electrolysis of Aqueous Salts (con’t) • Overvoltage: Additional voltage used to produce gases (including H2 and O2) at electrodes • Usually 0.4 – 0.6 V So what forms? • Cations of less active metals are reduced • Cations of more active metals are not reduced; Water is reduced instead • Anions that are oxidized are typically halides • F-, common oxoanions are not oxidized; water is oxidized instead

26. How much product forms? The amount of product is directly proportional to quantity of charge that flows How long does it take to produce 0.0423 mol of Cl2 (g) by electrolysis of NaCl (aq) with power supply current of 12 A?

27. Homework due TUESDAY, May 19th Chap 21: #16, 21, 30, 33, 38, 42, 56, 60, 70, 89, 94, 105