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Class 12 Chemistry Notes

Electrochemistry Class 12 Notes

Complete, exam-ready notes on electrochemistry: electrochemical cells, standard electrode potential, the Nernst equation, electrical conductance, Faraday's laws of electrolysis and commercial cells — written for CBSE boards, JEE and NEET revision.

Class12SubjectChemistryCoversCBSE · JEE · NEET

Written byDeep Narayan· Science & Mathematics EducatorReviewed byPushpanjali

What is an electrochemical cell in one line?

An electrochemical cell converts chemical energy into electrical energy (galvanic cell) or the reverse (electrolytic cell) using redox reactions — oxidation at the anode and reduction at the cathode.

Types of Electrochemical Cells

Galvanic (voltaic) cell

A device that converts chemical energy into electrical energy through a spontaneous redox reaction. The two electrodes are connected by a salt bridge and an external circuit.

Electrolytic cell

A device that uses electrical energy to drive a non-spontaneous redox reaction (electrolysis).

  • Anode:\text{Anode}:
  • where oxidation occurs (loses electrons); in a galvanic cell it is negative, in an electrolytic cell positive.
  • Cathode:\text{Cathode}:
  • where reduction occurs (gains electrons); in a galvanic cell positive, in an electrolytic cell negative.
  • The salt bridge completes the circuit and balances charge, maintaining electrical neutrality.

Electrode Potential and Standard Electrode Potential

Standard electrode potential

E∘E^{\circ}

The potential of an electrode measured at 1 M concentration, 1 bar pressure and 298 K, relative to the standard hydrogen electrode (SHE) taken as zero. It measures the tendency of an electrode to gain or lose electrons.

Predicting spontaneity

A larger (more positive)E∘E^{\circ} means a stronger oxidising agent and greater tendency to be reduced. A redox reaction is spontaneous if the cell potentialEcell∘=Ecathode∘−Eanode∘>0E^{\circ}_{\text{cell}} = E^{\circ}_{\text{cathode}} - E^{\circ}_{\text{anode}} > 0.

The Nernst Equation

Ecell=Ecell∘−0.059nlog⁡QE_{\text{cell}} = E^{\circ}_{\text{cell}} - \frac{0.059}{n}\log Q
Nernst equation (at 298 K)

Cell potential

EcellE_{\text{cell}}

The potential of the cell under non-standard conditions. At 298 K the general Nernst equation becomes the compact form above, where nn is the number of moles of electrons transferred and QQ is the reaction quotient.

  • At equilibrium:\text{At equilibrium}:
  • Q=KQ = K
  • and
  • Ecell=0E_{\text{cell}} = 0
  • , giving
  • log⁡K=nEcell∘0.059\log K = \frac{n E^{\circ}_{\text{cell}}}{0.059}
  • .
  • EMF and Gibbs energy:\text{EMF and Gibbs energy}:
  • ΔG∘=−nFEcell∘\Delta G^{\circ} = -nFE^{\circ}_{\text{cell}}
  • .

Electrical Conductance

Conductance measures how easily charge flows through a solution. It depends on the nature and concentration of the electrolyte. Key terms: conductivity, molar conductivity and its concentration dependence.

  • Conductivity:\text{Conductivity}:
  • κ=1ρ\kappa = \frac{1}{\rho}
  • (S m⁻¹), follows Kohlrausch's law.
  • Molar conductivity:\text{Molar conductivity}:
  • Λm=κ×1000c\Lambda_m = \frac{\kappa \times 1000}{c}
  • (S cm² mol⁻¹).
  • Strong electrolytes: Λₘ decreases slowly with √c (linear, Debye-Hückel).
  • Weak electrolytes: Λₘ rises sharply at low concentration and can't be extrapolated to Λ⁰ₘ directly — use Kohlrausch's law.

Kohlrausch's law

The limiting molar conductivity of an electrolyte equals the sum of the limiting molar conductivities of its constituent ions:Λm∘=λ+∘+λ−∘\Lambda^{\circ}_m = \lambda^{\circ}_+ + \lambda^{\circ}_-. Used to find Λ⁰ₘ of weak electrolytes.

Faraday's Laws of Electrolysis

First law

The mass of a substance deposited at an electrode is directly proportional to the quantity of electricity passed: W=Z Q=Z I tW = Z\,Q = Z\,I\,t.

Second law

When the same quantity of electricity is passed through different electrolytes, the masses deposited are proportional to their chemical equivalents (equivalent weights).

W=E I t96500W = \frac{E\,I\,t}{96500}
Mass of product (1 F = 96500 C = 1 equivalent)

Commercial Cells and Batteries

  • Primary cells:\text{Primary cells}:
  • cannot be recharged — e.g. dry cell (Leclanché), mercury cell.
  • Secondary cells:\text{Secondary cells}:
  • can be recharged — e.g. lead-acid battery, nickel-cadmium, lithium-ion.
  • Lead-acid battery discharge:
  • Pb+PbO2+2H2SO4→2PbSO4+2H2O\text{Pb} + \text{PbO}_2 + 2\text{H}_2\text{SO}_4 \to 2\text{PbSO}_4 + 2\text{H}_2\text{O}
  • .

Fuel cells

A hydrogen-oxygen fuel cell uses the combustion of hydrogen to produce electricity with water as the only product — clean and efficient, used in spacecraft.

Solved Examples

Example: For the cell Zn|Zn2+(0.1 M)∣∣Cu2+(0.5 M)∣Cu\text{Zn|Zn}^{2+}(0.1\,M)||\text{Cu}^{2+}(0.5\,M)|\text{Cu} with Ecell∘=1.10 VE^{\circ}_{\text{cell}} = 1.10\,V and n=2n = 2, find the cell EMF.

Solution: Ecell=1.10−0.0592log⁡[Zn2+][Cu2+]=1.10−0.0295log⁡0.2≈1.12 VE_{\text{cell}} = 1.10 - \frac{0.059}{2}\log \frac{[\text{Zn}^{2+}]}{[\text{Cu}^{2+}]} = 1.10 - 0.0295\log 0.2 \approx 1.12\,V.

Revision

Key formulas at a glance

Memorise these before attempting numericals — most exam questions hinge on one of them.

Cell potential

Ecell=Ecell∘−0.059nlog⁡QE_{\text{cell}} = E^{\circ}_{\text{cell}} - \frac{0.059}{n}\log Q

Gibbs energy

ΔG∘=−nFEcell∘\Delta G^{\circ} = -nFE^{\circ}_{\text{cell}}

Equilibrium constant

log⁡K=nEcell∘0.059\log K = \frac{nE^{\circ}_{\text{cell}}}{0.059}

Molar conductivity

Λm=κ×1000c\Lambda_m = \frac{\kappa \times 1000}{c}

Kohlrausch's law

Λm∘=λ+∘+λ−∘\Lambda^{\circ}_m = \lambda^{\circ}_+ + \lambda^{\circ}_-

Faraday's first law

W=ZIt=EIt96500W = ZIt = \frac{EIt}{96500}

Exam tips

How this chapter is asked

Where this topic appears in CBSE, JEE Main and NEET papers.

  • In a galvanic cell, anode is negative; in an electrolytic cell, anode is positive.
  • Salt bridge maintains electrical neutrality; it does not take part in the reaction.
  • At 298 K, Nernst equation uses 0.059/n in place of (2.303 RT)/nF.
  • 1 Faraday = 96500 C deposits 1 gram-equivalent.
  • Strong electrolytes follow Debye-Hückel; weak ones need Kohlrausch's law for Λ⁰ₘ.

FAQ

Common questions

What is the role of a salt bridge?

A salt bridge completes the circuit and maintains electrical neutrality by allowing movement of inert ions, preventing charge build-up in the half-cells.

How do I know if a cell reaction is spontaneous?

The cell is spontaneous if the standard cell potential E°cell = E°cathode − E°anode is positive (so ΔG° = −nFE°cell is negative).

What is the difference between molar conductivity and conductivity?

Conductivity (κ) is the ability of 1 m³ of solution to conduct current; molar conductivity (Λₘ) is the conductance of all ions produced from one mole of electrolyte dissolved in a known volume.

Why does electrolysis deposit mass proportional to time?

Faraday's first law gives W = ZIt: the mass deposited grows linearly with the current and the time for which it flows, with Z the electrochemical equivalent of the substance.

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Chemical Kinetics NotesSolutions NotesSolid State NotesClass 12 Chemistry NotesChemistry Tutor in Mohali

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