Redox reactions, galvanic & electrolytic cells, the electrochemical series and Faraday's laws — exam-focused for the BIEK / Sindh Board paper. Read it, or open the interactive lecture and build a cell to measure its EMF.
1 — Oxidation & reduction (redox)
| Oxidation | Reduction |
|---|
| Electrons | loss of electrons | gain of electrons |
| Oxygen / Hydrogen | gain O / lose H | lose O / gain H |
| Oxidation number | increases | decreases |
OIL RIG — Oxidation Is Loss, Reduction Is Gain (of electrons). They always happen together → a redox reaction.
2 — Oxidation numbers
The oxidation number is the charge an atom would have if all bonds were ionic.
- Free element = 0 (e.g. Zn, O₂).
- Simple ion = its charge (Na⁺ = +1, Cl⁻ = −1).
- O = −2 (except peroxides −1); H = +1 (except metal hydrides −1).
- Sum of oxidation numbers = the overall charge.
find the oxidation number
Oxidation number of S in H₂SO₄?
2(+1) + S + 4(−2) = 0 → S = +6
3 — Oxidising & reducing agents
- Oxidising agent — accepts electrons (and is itself reduced). e.g. KMnO₄, O₂, Cl₂.
- Reducing agent — donates electrons (and is itself oxidised). e.g. metals, H₂, C.
4 — Balancing redox equations
Split into two half-reactions (one oxidation, one reduction), balance atoms then electrons, and add them so the electrons cancel.
Example (Zn + Cu²⁺)Zn → Zn²⁺ + 2e⁻ (oxidation)
Cu²⁺ + 2e⁻ → Cu (reduction)
Zn + Cu²⁺ → Zn²⁺ + Cu (overall)
5 — Galvanic (voltaic) cells
- Galvanic cell — converts chemical energy → electrical energy via a spontaneous redox reaction (a battery).
The Daniell cell: a Zn rod in ZnSO₄ and a Cu rod in CuSO₄, joined by a wire and a salt bridge. Zn is oxidised (the anode, −), Cu²⁺ is reduced (the cathode, +); electrons flow Zn → Cu through the wire.
Anode = oxidation (electrons leave); cathode = reduction (electrons arrive). In a galvanic cell the anode is negative.
6 — Salt bridge & cell notation
- Salt bridge — a tube of inert electrolyte (e.g. KNO₃) that completes the circuit and keeps both half-cells electrically neutral.
Cell notationZn(s) | Zn²⁺(aq) ‖ Cu²⁺(aq) | Cu(s)
(anode on the left, ‖ = salt bridge, cathode on the right)
7 — Standard electrode potential & the electrochemical series
- Standard electrode potential (E°) — the voltage of a half-cell vs the standard hydrogen electrode (E° = 0.00 V) at 298 K, 1 M, 1 atm.
| Electrode | E° (V) |
|---|
| Zn²⁺/Zn | −0.76 |
| Fe²⁺/Fe | −0.44 |
| 2H⁺/H₂ | 0.00 |
| Cu²⁺/Cu | +0.34 |
| Ag⁺/Ag | +0.80 |
A more negative E° = more easily oxidised (a better reducing agent / more reactive metal).
8 — EMF of a cell
Cell EMFE°cell = E°cathode − E°anode (= E°(reduction) − E°(oxidation))
A positive E°cell means the reaction is spontaneous (the cell delivers a voltage).
EMF of the Daniell cell
Zn (−0.76) and Cu (+0.34). Find E°cell.
E°cell = E°(Cu, cathode) − E°(Zn, anode) = (+0.34) − (−0.76) = +1.10 V
9 — Electrolytic cells & electrolysis
- Electrolytic cell — uses electrical energy → chemical change (a non-spontaneous reaction is driven by an external supply).
In electrolysis, cations move to the cathode (reduced) and anions move to the anode (oxidised). Here the anode is positive and the cathode is negative (opposite to a galvanic cell).
Example: electrolysis of molten NaCl gives Na at the cathode and Cl₂ at the anode.
10 — Faraday's laws of electrolysis
- First law: the mass deposited ∝ the quantity of charge (Q = I × t).
- Second law: for the same charge, the mass deposited ∝ the equivalent mass.
Charge & the faradayQ = I t · 1 mole of electrons = 1 faraday = 96 500 C
11 — Applications
- Batteries & cells — dry cell, lead-acid, lithium-ion (galvanic).
- Electroplating — coating a metal with another (silver, chromium).
- Extraction & refining of metals (Al, Cu) by electrolysis.
- Corrosion (rusting) is electrochemical; prevented by galvanising or sacrificial anodes.
12 — Exam recap
- Redox: oxidation = loss, reduction = gain (OIL RIG); oxidation numbers.
- Oxidising vs reducing agents; balancing by half-reactions.
- Galvanic cell (Daniell), salt bridge, cell notation.
- Standard electrode potential & the electrochemical series.
- E°cell = E°cathode − E°anode; positive → spontaneous.
- Electrolysis; Faraday's laws (Q = It, 1 F = 96 500 C).