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Electrochemistry revision notes

Every idea in the Electrochemistry chapter for JEE Main, with the usual mistakes and the fast methods. Topic 1 is open to everyone; sign in free to read all 30 and download the PDF.

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How we count: 7× asked means 7 past JEE Main questions in the PrepWiser bank needed that concept, and "last" is the latest year one appeared. ★ marks ideas asked 8 or more times.

Galvanic cells and electrode potentials

Galvanic and electrolytic cells
3× asked, last 2026
Zn(s)+CuX2+(aq)→ZnX2+(aq)+Cu(s),Ecell∘=1.10 V  (Daniell cell)\ce{Zn(s) + Cu^2+(aq) -> Zn^2+(aq) + Cu(s)},\qquad E^\circ_{\text{cell}} = 1.10\ \text{V}\ \ (\text{Daniell cell})
Galvanic: spontaneous reaction gives current; anode = oxidation (−), cathode = reduction (+); electrons go anode → cathode.
salt bridge (KCl in agar)Ve−e^-Zn anode (−)oxidationZnX2+\ce{Zn^{2+}}Cu cathode (+)reductionCuX2+\ce{Cu^{2+}}
Daniell cell: E∘=0.34−(−0.76)=1.10E^\circ=0.34-(-0.76)=1.10 V; electrons go from Zn to Cu through the wire
  • Electrolytic: an external source drives the reaction; anode +, cathode −, yet oxidation is still at the anode and electrons still go anode → cathode outside.
  • Not 'Zn dissolves' when Eext>1.10E_{\text{ext}}>1.10 V. An opposing voltage above EcellE_{\text{cell}} reverses the reaction: Zn deposits, Cu dissolves, the cell turns electrolytic. Below 1.10 V it stays galvanic.
Cell notation and salt bridge
2× asked, last 2024
Zn(s) ∣ ZnX2+(aq) ∣∣ CuX2+(aq) ∣ Cu(s)\ce{Zn(s) | Zn^2+(aq) || Cu^2+(aq) | Cu(s)}
Anode left, cathode right. | is a phase boundary; || is a salt bridge (KCl or KNO₃ in agar) that closes the circuit.
  • Quick method. Matching a reaction to a cell: oxidation half on the left, and every ion in the reaction must be in that electrolyte. 12 HX2+AgCl→HX++ClX−+Ag\ce{1/2H2 + AgCl -> H+ + Cl- + Ag} needs Pt∣HX2∣HCl(aq)∣AgCl∣Ag\ce{Pt|H2|HCl(aq)|AgCl|Ag}; KCl has no HX+\ce{H+}.
SHE and types of electrodes
2× asked, last 2024
Pt(s) ∣ HX2(g) ∣ HX+(aq)  (pHX2=1 bar, [HX+]=1 M):E∘=0 V\ce{Pt(s) | H2(g) | H+(aq)}\ \ (p_{\ce{H2}} = 1\ \text{bar},\ [\ce{H+}] = 1\ \text{M}):\quad E^\circ = 0\ \text{V}
Types: metal–ion, gas–ion (on Pt), redox (Pt in two ions), metal–insoluble salt–anion (calomel, Ag/AgCl: depend on [ClX−][\ce{Cl-}]).
Standard cell potential
7× asked, last 2026
Ecell∘=Ecathode∘−Eanode∘=Eright∘−Eleft∘E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}} = E^\circ_{\text{right}} - E^\circ_{\text{left}}
Both are standard reduction potentials (1 M, 1 bar). Ecell∘>0E^\circ_{\text{cell}}>0: spontaneous as written; <0<0: the reverse is.
ZnX2+/Zn\ce{Zn^{2+}/Zn} −0.76-0.76SHE 0CuX2+/Cu\ce{Cu^{2+}/Cu} +0.34+0.34Ecell∘=Ecathode−Eanode=1.10E^\circ_{\rm cell}=E_{\rm cathode}-E_{\rm anode}=1.10 V
Both are reduction potentials: subtract, never add, and never multiply by the coefficients
  • Not E1+E2E_1+E_2, and never double E∘E^\circ with the equation. Both are reduction potentials, so subtract; E∘E^\circ is intensive, only ΔG∘\Delta G^\circ scales. Oxidation potentials given? Flip their signs first.
What decides a metal's E°
4× asked, last 2023
M(s)→ΔsubHM(g)→ΔiHM+(g)→ΔhydHM+(aq)\mathrm{M(s)}\xrightarrow{\Delta_{\text{sub}}H}\mathrm{M(g)}\xrightarrow{\Delta_iH}\mathrm{M^+(g)}\xrightarrow{\Delta_{\text{hyd}}H}\mathrm{M^+(aq)}
Sublimation, ionisation, hydration; no electron gain. LiX+\ce{Li+} hydration makes Li most negative (−3.05 V).

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25 more concepts in 7 topics

Electrochemical series and its uses; Nernst equation; Cell potential, Gibbs energy and K; Conductance of electrolytic solutions; Kohlrausch's law and its applications; Electrolysis and Faraday's laws; and Batteries, fuel cells and corrosion. Read the first part free. Sign in free with Google to read the rest and download the PDF.

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Electrochemical series and its uses

  • Electrochemical series7×
  • Displacement and H₂ from acid3×
  • Combining E° of half-reactions8×

Nernst equation

  • Nernst equation for a cell17×
  • Nernst equation for one electrode4×
  • pH-dependent electrodes16×
  • Concentration cells2×
  • Metal–insoluble salt electrode and Ksp2×

Cell potential, Gibbs energy and K

  • ΔG and cell potential13×
  • E° and the equilibrium constant5×

Conductance of electrolytic solutions

  • Metallic and electrolytic conduction4×
  • Resistance, conductivity, cell constant10×
  • Molar conductivity8×
  • Variation with concentration12×
  • Conductometric titrations2×

Kohlrausch's law and its applications

  • Kohlrausch's law10×
  • Degree of dissociation and Ka6×
  • Solubility from conductivity2×

Electrolysis and Faraday's laws

  • Faraday's first law17×
  • Faraday's second law4×
  • Products of electrolysis9×

Batteries, fuel cells and corrosion

  • Dry cell and mercury cell4×
  • Lead storage and Ni–Cd cells5×
  • Fuel cells6×
  • Corrosion2×

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