Redox Reactions and Electrochemistry formulas for JEE Main & Advanced

12 formula entries for Redox Reactions and Electrochemistry, in the order the chapter builds them. Each shows the formula, when to use it and the first trap to avoid.

Standard Electrode Potential (E°)

E_{cell}^{\circ} = E_{cathode}^{\circ} - E_{anode}^{\circ}

When to use. Calculate cell potential and predict spontaneity.

Trap. Using reduction potentials directly without subtracting; always subtract anode from cathode.

Nernst Equation

E = E^{\circ} - \frac{R T}{n F} \ln Q

When to use. Calculate cell potential under non-standard conditions or equilibrium constant.

Trap. Using log base 10 vs ln; remember 0.0591 is for log10 at 298 K.

Gibbs Free Energy and Cell Potential

Δ G = - n F E_{cell}

When to use. Relate electrochemical data to thermodynamics.

Trap. Forgetting that n is number of moles of electrons transferred, not stoichiometric coefficient.

Equilibrium Constant from E°

E^{\circ} = \frac{0.0591}{n} \log K

When to use. Calculate equilibrium constant from standard cell potential.

Trap. Forgetting that K is dimensionless and uses activities, but concentrations are often used.

Faraday's Laws of Electrolysis

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

When to use. Calculate mass, time, or current in electrolysis.

Trap. Forgetting to convert time to seconds if given in minutes.

Conductance in Electrolytes

Λ_{m} = \frac{κ}{C}

When to use. Study electrolyte behavior and determine degree of dissociation.

Trap. Confusing conductivity (κ) with conductance (G); $κ = G \times G^{*}$ , where $G^{*} = l / A$ is the cell constant, so $G = κ / G^{*}$ .

Kohlrausch's Law

Λ_{m}^{\infty} = \sum ν_{i} λ_{i}^{\infty}

When to use. Calculate Λm° for weak electrolytes or find degree of dissociation.

Trap. Forgetting to multiply by stoichiometric coefficients (ν).

Degree of Dissociation (α) from Conductivity

α = \frac{Λ_{m}}{Λ_{m}^{\infty}}

When to use. Determine dissociation constant of weak acids/bases from conductivity data.

Trap. Using this for strong electrolytes; α ≈ 1 for them.

Relation Between ΔG°, E°, and K

Δ G^{\circ} = - n F E^{\circ} = - R T \ln K

When to use. Convert between thermodynamic and electrochemical quantities.

Trap. Forgetting that R = 8.314 J/mol·K, so ΔG° in J, not kJ.

Disproportionation Reactions

E_{cell}^{\circ} = E_{reduction}^{\circ} - E_{oxidation}^{\circ}

When to use. Predict if a species (e.g., Cu+, MnO4²⁻) undergoes disproportionation.

Trap. Thinking all disproportionations are spontaneous; check E°cell.

Standard Hydrogen Electrode (SHE)

2 H^{+} + 2 e^{-} \to H_{2}

When to use. Measure electrode potentials relative to SHE.

Trap. Thinking SHE is used as anode always; it can be cathode if other half has lower E°.

Concentration Cells

E_{cell} = \frac{0.0591}{n} \log \frac{C_{cathode}}{C_{anode}}

When to use. Calculate potential due to concentration difference.

Trap. Thinking E°cell is non-zero; it's zero because same half-reaction.

Values to remember

$E_{SHE}^{\circ} = 0 V$ Standard Electrode Potential (E°)
$\frac{R T}{F} \ln 10 \approx 0.0591 V$ at 298 K Nernst Equation
$F = 96485 {C mol}^{- 1}$ Nernst Equation
$1 V = 1 {J C}^{- 1}$ Gibbs Free Energy and Cell Potential
$E_{{Li}^{+} / Li}^{\circ} = - 3.04 V$ Electrochemical Series
$E_{F_{2} / F^{-}}^{\circ} = + 2.87 V$ Electrochemical Series
$1 Faraday = 96485 C$ Faraday's Laws of Electrolysis
$λ_{H^{+}}^{\infty} \approx 350 {S cm}^{2} {mol}^{- 1}$ Kohlrausch's Law
$λ_{{OH}^{-}}^{\infty} \approx 200 {S cm}^{2} {mol}^{- 1}$ Kohlrausch's Law
$E_{O_{2} / H_{2} O}^{\circ} = + 1.23 V$ Product Prediction in Electrolysis
$E_{{Cl}_{2} / {Cl}^{-}}^{\circ} = + 1.36 V$ Product Prediction in Electrolysis
$E_{lead-acid}^{\circ} \approx 2.0 V$ Batteries and Fuel Cells
$R = 8.314 {J mol}^{- 1} K^{- 1}$ Relation Between ΔG°, E°, and K
$T = 298 K$ (standard) Relation Between ΔG°, E°, and K

Read the full chapter: Redox Reactions and Electrochemistry for JEE, with what JEE asks and the common traps.