Electrochemistry, Corrosion & Its Prevention Notes

Last reviewed 16 Sept 2026 · 11 min read

Electrochemistry basics

Electrolytes conduct electricity by movement of ions (in solution or molten state) — strong (NaCl, HCl) and weak (CH₃COOH, NH₄OH).
Conductance $G = 1/ R$ ; specific conductance (conductivity) κ; molar conductivity increases with dilution.
Electrochemical cell (galvanic/voltaic) — converts chemical energy → electrical energy spontaneously (batteries).
Electrolytic cell — uses electrical energy to drive non-spontaneous reactions (electroplating, electrolysis).

Electrodes

Anode — oxidation (loss of electrons).
Cathode — reduction (gain of electrons).
In a galvanic cell, anode is negative, cathode positive; in an electrolytic cell, anode is positive.
Mnemonic: "An Ox, Red Cat" (anode oxidation, reduction cathode).

Daniell cell: Zn | Zn²⁺ || Cu²⁺ | Cu — zinc anode (Zn → Zn²⁺ + 2e⁻), copper cathode (Cu²⁺ + 2e⁻ → Cu); salt bridge completes the circuit and maintains neutrality.

Electrode potential and EMF

Standard electrode potential E° — measured against the standard hydrogen electrode (SHE), assigned 0.00 V, at 25 °C, 1 M ion concentration, 1 atm.
Standard reduction potentials (examples):

Half reaction	E° (V)
Li⁺ + e⁻ → Li	−3.04
Mg²⁺ + 2e⁻ → Mg	−2.37
Al³⁺ + 3e⁻ → Al	−1.66
Zn²⁺ + 2e⁻ → Zn	−0.76
Fe²⁺ + 2e⁻ → Fe	−0.44
2H⁺ + 2e⁻ → H₂	0.00
Cu²⁺ + 2e⁻ → Cu	+0.34
O₂ + 2H₂O + 4e⁻ → 4OH⁻	+0.40
Ag⁺ + e⁻ → Ag	+0.80
Au³⁺ + 3e⁻ → Au	≈ +1.5

Electrochemical series — metals with more negative E° are more active (stronger reducing agents, more easily oxidised/corroded) and displace metals below them from solution.

FormulaCell EMF and Nernst equation

E_{ce l l}^{\circ} = E_{c a t h o d e}^{\circ} - E_{an o d e}^{\circ}

Nernst equation (25 °C):

E = E^{\circ} - \frac{0.0591}{n} lo g_{10} Q

$n$ = electrons transferred; $Q$ = reaction quotient. For an electrode $M^{n +} / M$ : $E = E^{\circ} + \frac{0.0591}{n} lo g [M^{n +}]$ .

Spontaneity: $Δ G = - n F E$ — reaction is spontaneous when $E_{ce l l} > 0$ .

Reference electrodes and pH

Calomel electrode (Hg/Hg₂Cl₂/KCl), silver–silver chloride electrode — stable secondary references.
Copper–copper sulphate electrode (CSE) — used in half-cell potential surveys of reinforced concrete.
Glass electrode — pH measurement; pH $= - lo g [H^{+}]$ .

Batteries and fuel cells

Type	Examples	Features
Primary (non-rechargeable)	Dry cell (Leclanché, zinc–carbon), alkaline cell	Single use
Secondary (rechargeable)	Lead–acid (Pb anode, PbO₂ cathode, H₂SO₄; ~2 V per cell), nickel–cadmium, nickel–metal hydride, lithium-ion	Rechargeable; Li-ion — high energy density (vehicles, electronics, grid storage for solar)
Fuel cells	Hydrogen–oxygen fuel cell (H₂ + ½O₂ → H₂O)	Continuous supply of fuel; high efficiency; only water as product (hydrogen fuel cells)

Electrolysis — Faraday's laws

FormulaFaraday's laws

First law: mass deposited ∝ charge passed

m = Z I t = \frac{M I t}{n F}

$Z$ = electrochemical equivalent; $M$ = molar mass; $n$ = valency (electrons per ion); $F$ = Faraday constant ≈ 96 485 C/mol.

Second law: for the same charge, masses deposited are proportional to chemical equivalent weights ( $M / n$ ).

Applications: electroplating (chromium, nickel, zinc), electrorefining of copper, electrometallurgy of aluminium, electrochemical chloride extraction and re-alkalisation of concrete (rehabilitation techniques).

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