Electrochemistry and Redox Reactions
This domain tests oxidation numbers, half-equation balancing in acidic and alkaline media, galvanic cell predictions from standard electrode potentials, and the quantitative link between charge, moles of electrons and mass deposited in electrolysis.
Oxidation numbers
- A free element has oxidation number zero; a monatomic ion equals its charge; oxygen is usually −2 but −1 in peroxides; hydrogen is +1 except in hydrides; the sum equals the overall charge.
- Oxidation is loss of electrons and an increase in oxidation number; reduction is gain of electrons and a decrease.
- Disproportionation is simultaneous oxidation and reduction of the same element, as in Cl2 with cold dilute NaOH.
Balancing redox equations
- Split into half-equations, balance atoms other than O and H, balance O with H2O, balance H with H+, and balance charge with electrons.
- In alkaline solution, add OH− to both sides to neutralise H+, then cancel water.
- Multiply each half-equation so the electron counts match before adding.
Cells and potentials
- E°cell = E°cathode − E°anode; a positive value means the reaction is spontaneous as written.
- E° is intensive: do not multiply it when you double a half-equation. ΔG° = −nFE° with F = 96 485 C mol−1.
- Oxidation occurs at the anode. In a galvanic cell the anode is the negative electrode; in an electrolytic cell it is the positive one.
Worked example
Silver is deposited from AgNO3 solution using a current of 2.00 A for 30.0 min. Q = It = 2.00 × 1800 = 3600 C. n(e−) = 3600/96 485 = 0.0373 mol. Since Ag+ + e− → Ag, n(Ag) = 0.0373 mol and m = 0.0373 × 107.9 = 4.03 g.
Traps that cost marks
- Reversing the sign conventions for galvanic and electrolytic cells.
- Multiplying E° values when balancing a cell reaction.
- Forgetting to balance charge with electrons before adding half-equations.
- Assigning oxygen −2 in peroxides or hydrogen +1 in metal hydrides.
- Mixing seconds and minutes in Q = It, or using the wrong number of electrons per ion.