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Electrochemistry: concepts, formulas & solved PYQ-style examples

Electrochemistry is calculation-heavy but very predictable: cell EMF, the Nernst equation and Gibbs energy questions come up regularly.

Write the cell reaction first and count the electrons (n) — every formula in this chapter needs n.

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Key concepts

Electrode potentials

E°cell = E°cathode − E°anode using reduction potentials.

Nernst equation

EMF at non-standard concentrations.

Gibbs energy and EMF

ΔG° = −nFE°; a positive EMF means a spontaneous cell.

Conductance

Molar conductivity and Kohlrausch’s law.

Electrolysis

Faraday’s laws: mass deposited ∝ charge passed.

Formula sheet

Cell EMF

Ecell∘=Ecathode∘−Eanode∘E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}

Nernst equation (298 K)

E=E∘−0.0591nlog⁡QE = E^\circ - \dfrac{0.0591}{n}\log Q

Gibbs energy

ΔG∘=−nFE∘\Delta G^\circ = -nFE^\circ

Equilibrium constant

log⁡K=nE∘0.0591\log K = \dfrac{nE^\circ}{0.0591}

Faraday’s first law

m=M I tnFm = \dfrac{M\,I\,t}{nF}

What PYQs keep asking

Patterns that repeat across previous-year JEE and NEET papers for this chapter.

  • EMF of a cell at given ion concentrations using the Nernst equation.
  • ΔG° or equilibrium constant from E°.
  • Mass deposited in electrolysis for a given current and time.
  • Trends in molar conductivity with dilution.

Solved examples

Example 1

For the Daniell cell Zn | Zn²⁺(0.1 M) || Cu²⁺(1 M) | Cu, find E°cell and Ecell at 298 K (E°Cu²⁺/Cu = +0.34 V, E°Zn²⁺/Zn = −0.76 V).

  1. 1Ecell∘=0.34−(−0.76)=1.10E^\circ_{cell} = 0.34 - (-0.76) = 1.10 V.
  2. 2Reaction Zn+Cu2+→Zn2++Cu\text{Zn} + \text{Cu}^{2+} \to \text{Zn}^{2+} + \text{Cu}, so n=2n = 2 and Q=[Zn2+][Cu2+]=0.1Q = \dfrac{[\text{Zn}^{2+}]}{[\text{Cu}^{2+}]} = 0.1.
  3. 3E=1.10−0.05912log⁡0.1=1.10+0.0296E = 1.10 - \dfrac{0.0591}{2}\log 0.1 = 1.10 + 0.0296.
Answer: E° = 1.10 V, E ≈ 1.13 V
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Example 2

Find ΔG° for the Daniell cell (E° = 1.10 V, F = 96500 C mol⁻¹).

  1. 1ΔG∘=−nFE∘\Delta G^\circ = -nFE^\circ with n=2n = 2.
  2. 2ΔG∘=−2×96500×1.10=−212300\Delta G^\circ = -2 \times 96500 \times 1.10 = -212300 J mol⁻¹.
  3. 3Negative ΔG° confirms the cell reaction is spontaneous.
Answer: ΔG° ≈ −212.3 kJ mol⁻¹
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Common traps

  • Multiplying E° by the stoichiometric coefficient — electrode potentials are intensive.
  • Writing Q upside down (products over reactants is correct).
  • Using the wrong n — count electrons transferred in the balanced reaction.

Practise with

Molarity CalculatorChemical ThermodynamicsChemical BondingGeneral Organic Chemistry (GOC)4 Electrochemistry practice questionsVideo solutions

FAQ

Does E° change if I multiply the cell reaction by 2?

No. E° is intensive; ΔG° doubles, but E° stays the same.

What is 0.0591 in the Nernst equation?

It is 2.303RT/F at 298 K, in volts.

Is electrochemistry important for NEET?

Yes — EMF, Nernst equation and conductance questions are regular and fully scoring.

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