Redox Reactions and Electrochemistry
This domain tests electron bookkeeping and the conversion of chemical energy into electrical energy and back: oxidation numbers, half-equations, standard electrode potentials, galvanic cells and electrolysis.
Oxidation numbers and balancing
Oxidation is loss of electrons (increase in oxidation number); reduction is gain. Rules: elements 0; monatomic ion = charge; O usually −2 (but −1 in peroxides); H usually +1 (but −1 in hydrides); sum equals overall charge. Balance half-equations by atoms, then O with H2O, H with H+, then charge with electrons, and equalise electron numbers before adding.
Standard electrode potentials
E° values are measured against the standard hydrogen electrode (H+/H2, 1 M, 100 kPa, 298 K), defined as 0.00 V. E°cell = E°(cathode) − E°(anode), and a positive E°cell means the reaction is spontaneous under standard conditions. E° is intensive: doubling the equation does not double E°.
Cells and electrolysis
In a galvanic cell, oxidation occurs at the anode (negative electrode) and reduction at the cathode (positive). In electrolysis the anode is still where oxidation happens, but it is the positive electrode. Products depend on relative electrode potentials and concentration. Faraday's laws: Q = It, and n(e−) = Q/F with F = 96 500 C mol−1.
Worked example
For Zn|Zn2+‖Cu2+|Cu with E°(Cu2+/Cu) = +0.34 V and E°(Zn2+/Zn) = −0.76 V: E°cell = 0.34 − (−0.76) = +1.10 V. The positive value confirms zinc is oxidised and copper deposited spontaneously.
Traps that cost marks
- Reversing the anode sign convention: in a galvanic cell the anode is negative.
- Multiplying E° when multiplying a half-equation.
- Assuming standard E° predictions hold at non-standard concentrations.
- Forgetting that in electrolysis of aqueous solutions, water may be oxidised or reduced in preference to the solute.
- Mixing up Q = It units; current must be in amperes and time in seconds.