Electrochemistry and Chemical Kinetics
Domain 7: Electrochemistry and Chemical Kinetics
This domain tests cell potentials, the Nernst equation, conductance, and the rate laws that describe how fast reactions proceed.
Core rules
- Cell notation: anode on left, cathode on right; oxidation at anode, reduction at cathode. E°cell=E°cathode−E°anode. A positive E°cell means a spontaneous galvanic cell.
- Nernst equation at 298 K: E=E°−(0.0592/n)log Q. For a concentration cell, E°=0 and E depends on the ion concentration ratio.
- Conductance increases with ion concentration and mobility; Kohlrausch's law gives limiting molar conductivity as a sum of independent ion contributions. Faraday's laws link charge Q=It to moles via nF.
- Rate law is determined experimentally: rate=k[A]^m[B]^n; order is m+n, not the stoichiometric coefficient. First-order integrated law: ln[A]t=−kt+ln[A]0; t1/2=0.693/k.
- Arrhenius: k=Ae^(−Ea/RT); a plot of ln k versus 1/T has slope −Ea/R. Catalysts lower Ea without changing ΔH.
Worked example
For Zn(s)+Cu2+(aq) -> Zn2+(aq)+Cu(s), E°=1.10 V, n=2. With [Zn2+]=0.10 M and [Cu2+]=0.010 M, Q=0.10/0.010=10. E=1.10−(0.0592/2)log10=1.10−0.0296=1.07 V. Increasing [Cu2+] or decreasing [Zn2+] raises the cell potential.
Traps
- Reversing anode and cathode in E°cell.
- Using the stoichiometric coefficient as reaction order.
- Forgetting that E° is intensive and unchanged by multiplying the equation.
- Mixing half-life formulas for zero-, first-, and second-order reactions.
- Ignoring units of k; first-order k has s^−1.