Redox and Electrochemistry
Redox and electrochemistry tests electron transfer: oxidation numbers, half-equations, electrochemical cells, standard electrode potentials and electrolysis.
Core rules
- Oxidation is loss of electrons and an increase in oxidation number; reduction is gain of electrons and a decrease in oxidation number.
- Oxidation-number rules: O is usually −2, H usually +1, Group 1 +1, Group 2 +2; the sum equals the overall charge.
- Balance half-equations by atoms first, then O with H2O, H with H+, and charge with electrons. In alkaline solution, add OH− to both sides to remove H+.
- In any cell, oxidation occurs at the anode and reduction at the cathode. For standard cells, Ecell = E(right) − E(left), using reduction potentials. A positive Ecell indicates a feasible reaction.
- Standard conditions: 298 K, 1 mol dm−3 solutions and 100 kPa for gases.
- Electrolysis uses electrical energy to drive a non-spontaneous change; Q = It and moles of electrons = Q/96 500.
Worked example
Find the oxidation number of Mn in MnO4−. Let Mn = x. Four O atoms contribute 4 × (−2) = −8. The ion charge is −1, so x − 8 = −1, giving x = +7. In a cell with Zn2+/Zn (E = −0.76 V) and Cu2+/Cu (E = +0.34 V), Ecell = 0.34 − (−0.76) = +1.10 V, so the reaction is spontaneous under standard conditions.
Traps
Reversing the sign of Ecell; flipping electrode potentials instead of subtracting; placing oxidation at the cathode; forgetting electrons in a half-equation; using the wrong oxidation-number rule for peroxides or hydrides; ignoring standard-state conditions when judging feasibility.