Redox Reactions and Electrochemistry
Oxidation numbers and balancing
Oxidation is loss of electrons (OIL RIG); reduction is gain. Assign oxidation numbers systematically: elements are 0, monatomic ions equal their charge, oxygen is usually −2, hydrogen is +1, and the sum equals the overall charge. Balance redox half-equations by adding electrons, then combine so electrons cancel.
Galvanic and electrolytic cells
In a galvanic cell, the more negative standard electrode potential is the anode, where oxidation occurs; the cathode is the positive electrode where reduction occurs. Electrons flow through the external circuit from anode to cathode; ions migrate through the salt bridge to maintain charge balance. Cell potential E°cell = E°cathode − E°anode, and a positive value indicates a spontaneous reaction.
In electrolysis, an external supply drives the non-spontaneous reaction. Predict products using relative potentials and concentration: dilute solutions may favour water oxidation or reduction.
Traps
- Reversing the anode/cathode labels between cell types.
- Forgetting to multiply E° values when doubling half-equations; potentials are intensive and are not multiplied.
- Using ΔG without checking the sign convention.
- Forgetting that the salt bridge completes the circuit without allowing electron flow.
- Ignoring concentration effects on electrode potential.
Rule of thumb
For galvanic cells: E°cell = reduction potential of cathode − reduction potential of anode. Positive means spontaneous.
Worked example
Zn²⁺/Zn is −0.76 V and Cu²⁺/Cu is +0.34 V. Zinc is the anode. E°cell = 0.34 − (−0.76) = +1.10 V, so the reaction is spontaneous and zinc dissolves.
Practical link
Clean electrodes before use and note the direction of electron flow from the voltmeter reading.