Chapter 18 Notes1 Chapter 18 Electrochemistry 1. review of terms; balancing redox equations 2. galvanic cell notation, relationships 3. standard reduction potentials 4. Nernst equation: example problems 5. pH electrodes; batteries 1. review of terms; balancing redox equations terms: the oxidation number of the oxidized species increases the oxidation number of the reduced species decreases an oxidizing agent is reduced a reducing agent is oxidized

Chapter 18 Notes2 algorithm for balancing redox reactions: 1. Separate the reaction into half-reactions, one for oxidation, one for reduction. 2. Balance each half-reaction for atoms other than H or O. 3. Balance O atoms by adding H 2 O to the side of the reaction deficient in O. 4. Balance H atoms by adding H 1+ to the side deficient in H. 5. If the reaction occurs in basic solution, convert each H 1+ to H 2 O by adding 1 OH 1- for each H 1+ to each side of the equation. 6. Balance charge by adding e 1- to the side that is more positive. 7. Combine half-reactions so that the e 1- cancel out (find the lowest common denominator).

Chapter 18 Notes3 2. galvanic cell notation, relationships

semi-permeable membrane: salt bridge Zn Cu ZnSO 4 (aq)CuSO 4 (aq) Daniell Cell electrode compartment: oxidation electrode compartment: reduction Pt wire electrode: anode electrode: cathode R e 1- Zn Zn 2+ e 1- Cu 2+ Cu e 1- Zn 2+ SO 4 2-

Chapter 18 Notes5 when R is large (open circuit) the cell is in a metastable equilibrium state at equilibrium the change in free energy as the reaction progress continues is 0 (recall the graph of G vs. reaction progress) the free energy includes contributions from the chemical species, mixing, and the electrons (the electric potential)

Chapter 18 Notes6 Nernst Equation