Chapter 20: Electrochemistry

1. Chapter 20: Electrochemistry Chemistry 1062: Principles of Chemistry II Andy Aspaas, Instructor

2. Oxidation-Reduction reactions • Oxidation-reduction (redox) reaction: transfer of electrons from one species to another • H3O+ becomes simply H+ when dealing with redox reactions to simplify balancing • (still the same species, just different notation) • Skeleton oxidation-reduction equation: involves only the species being oxidized and reduced. • Write oxidation numbers above each species. • No spectator ions, no balancing • Half reaction: shows only one oxidation OR one reduction • Most redox reactions are split into an oxidation half-reaction and a reduction half-reaction • LEO, GER

3. Balancing redox equations in acidic solutions • For each half reaction… • Balance everything except H or O • Balance O by adding H2O to one side • Balance H by adding H+ to one side • Balance charge by adding e- to one side • Multiply each half reaction by a factor so that the electrons cancel when the two half reactions are added together (e- cannot appear in the final equation) • Add the reactions, cancel anything that appears on the left and right, and simplify the coefficients to the smallest integers

4. Practice balancing acidic redox reactions • Balance I2(s) + NO3-(aq)  IO3-(aq) + NO2(g) in acidic solution • Half reactions • Cancel electrons • Add half-reactions • Simplify

5. Voltaic cells • A voltaic cell consists of two half-cells • Each half-cell contains a metal rod dipped in a solution containing that metal ion • Anode: a species is being oxidized • Cathode: a species is being reduced • Cell reaction: redox reaction for entire voltaic cell • Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s) Zn(s)  Zn2+(aq): oxidation half-reaction, anode Cu2+(aq)  Cu(s): reduction half-reaction, cathode

6. Cell notation • Zn(s) + Cu2+(aq)  Zn2+(aq) + Cu(s) • Cell notation: Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s) • Anode || Cathode • Write half reactions and cell reactions for the following cell: Tl(s) | Tl+(s) || Sn2+(aq) | Sn(s)

7. emf, Standard Cell emf, Standard electrode potential • Electromotive force, emf, Ecell = electrical pressure across the conductors of an electrochemical cell • Unit: Volt, V • Measure of the driving force of a cell reaction • Standard cell emf, Eocell = solutes are 1 M, gases are 1 atm, temperature is 25 oC • Standard electrode potenital • By convention, the standard hydrogen electrode has an emf of 0 V • All reactions shown as reductions • Ecell = Ecathode – Eanode • Ecellis positive for spontaneous reactions as written

8. Practice calculating Ecell • Using standard potentials, calculate Ecellfor Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s) • Standard cell potentials are an intensive property • Do not depend on quantity! • If you have to multiply a half-reaction to cancel electrons, do not multiply the Eo for that half-reaction

9. Free energy and K from Ecell • ΔGo = -nFEcell n = moles of electrons transferred F = Faraday’s constant, 96,500 C/mol e- • This gives an answer in J, since 1 J = 1 C·V • Convert to kJ since that’s what ΔGo is usually expressed in • Ecell = (0.0592 / n) log K (Nernst equation)

10. Practice with ΔGoand K • Calculate ΔGoand K for the following cell: Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s) • ΔGo = -nFEcell • Ecell = (0.0592 / n) log K