Electrochemistry Class 12 Notes: Formulas, Cells and Numericals

These electrochemistry class 12 notes cover the whole chapter in one place. Electrochemistry links chemical reactions with electricity, so it mixes theory with many numericals. The chapter carries about 9 marks in the CBSE theory paper, which makes it one of the heaviest chapters in Class 12 Chemistry. First, learn the cell types and the Nernst equation. Then practise conductance and Faraday’s laws, since most numericals come from them.

Electrochemistry Class 12 Notes: Galvanic and Electrolytic Cells

A galvanic cell converts the energy of a spontaneous redox reaction into electrical energy. The Daniell cell is the standard example. It has a zinc electrode in zinc sulphate and a copper electrode in copper sulphate, and a salt bridge joins the two half-cells. Zinc oxidises at the anode, so electrons flow through the wire to the copper cathode. Meanwhile, the salt bridge keeps both solutions electrically neutral.

An electrolytic cell works in the opposite way. Here, an external source forces a non-spontaneous reaction, so the anode is positive and the cathode is negative. Therefore, remember that the sign of the electrodes changes between the two cells, although oxidation is still at the anode.

Feature	Galvanic cell	Electrolytic cell
Energy change	Chemical to electrical	Electrical to chemical
Reaction	Spontaneous	Non-spontaneous
Anode sign	Negative	Positive
Cathode sign	Positive	Negative

Electrode Potential and Cell Potential

Every half-cell has an electrode potential. The standard hydrogen electrode is the reference, and its potential is zero by definition. A positive E° means the species is reduced more easily than hydrogen. As a result, a larger E° marks a stronger oxidising agent.

The standard cell potential is E°cell = E°cathode − E°anode. If E°cell is positive, the reaction is spontaneous. Also, the cell notation writes the anode on the left and the cathode on the right, for example Zn | Zn²⁺ || Cu²⁺ | Cu.

The Nernst Equation

For a reaction aA + bB → cC + dD, the Nernst equation is E = E° − (RT/nF) ln Q. At 298 K, it becomes E = E° − (0.0591/n) log Q. Here, n is the number of electrons transferred, and Q is the reaction quotient. At equilibrium, E is zero, so E° = (0.0591/n) log K.

The equation also connects with thermodynamics. Because ΔG° = −nFE°, a positive E° means a negative ΔG°. Likewise, ΔG° = −RT ln K links the cell to the equilibrium constant. Use F = 96,500 C mol⁻¹ in numericals.

Conductance of Electrolytic Solutions

Resistance R and conductance G are inverses, so G = 1/R. Conductivity κ is the conductance of a cube of solution with unit dimensions, and κ = G × (l/A). The ratio l/A is called the cell constant.

Molar conductivity is Λm = κ × 1000 / M, where M is the molarity in mol L⁻¹. For strong electrolytes, Λm rises slowly as dilution increases. For weak electrolytes, however, Λm rises steeply at high dilution because ionisation increases.

Debye–Hückel–Onsager gave Λm = Λ°m − A√c for strong electrolytes. Hence, a plot of Λm against √c is a straight line, and its intercept gives Λ°m.

Kohlrausch’s Law

Kohlrausch’s law states that the limiting molar conductivity of an electrolyte equals the sum of the limiting ionic conductivities of its cation and anion. Consequently, you can find Λ°m of a weak electrolyte from strong ones. For acetic acid, Λ°m(CH₃COOH) = Λ°m(CH₃COONa) + Λ°m(HCl) − Λ°m(NaCl).

The law also gives the degree of dissociation, α = Λm / Λ°m. Then Ka = cα² / (1 − α). Therefore, one table of ionic conductivities solves many problems.

Electrolysis and Faraday’s Laws

Faraday’s first law says that the mass deposited is proportional to the charge passed, so m = Z × I × t. Faraday’s second law says that equal charge deposits masses in the ratio of their equivalent weights. In practice, use this chain: charge Q = I × t, then moles of electrons = Q / 96,500, then moles of substance from the half-reaction.

The product of electrolysis depends on the electrode and the electrolyte. Molten sodium chloride gives sodium and chlorine. Aqueous sodium chloride, however, gives hydrogen at the cathode and chlorine at the anode. In aqueous solutions, water can compete, so compare the electrode potentials to decide.

Batteries and Fuel Cells

A primary battery, such as the dry cell, cannot be recharged. A secondary battery, such as the lead storage battery, can be recharged. The lead storage battery has a lead anode, a lead dioxide cathode and sulphuric acid, and its discharge reaction forms lead sulphate on both plates.

A fuel cell turns the energy of a fuel directly into electricity. The hydrogen–oxygen fuel cell has an efficiency of about 70 percent, and its only product is water. Because it never runs out while fuel is supplied, it works well in space missions.

Corrosion

Corrosion is the slow damage of metals by air and moisture, and rusting of iron is the usual example. It is an electrochemical process: iron acts as the anode and loses electrons, while oxygen is reduced at another spot. Finally, hydrated iron(III) oxide forms as rust.

Prevention methods include painting, galvanising, which coats iron with zinc, and cathodic protection with a more reactive metal. Sacrificial anodes are also used on ships and pipelines.

How to Solve Electrochemistry Numericals

• Write the half-reactions first, and find n from them.
• Use E°cell = E°cathode − E°anode, and check the sign.
• Write Q carefully, because pure solids and liquids do not appear in it.
• Convert units before you substitute: use litres for molarity and seconds for time.
• Finally, check that the answer has the right unit and a sensible size.

Common Mistakes in Electrochemistry

• Many students swap the anode and cathode in a cell diagram, so the sign of E°cell flips.
• Some forget to multiply E° by the stoichiometric coefficient, although E° is an intensive property and must not be multiplied.
• Others use ΔG° = −nFE° without checking the value of n.
• In Kohlrausch problems, students often mix up Λm and Λ°m.

Revision Plan for Electrochemistry

Spend day one on cells and the Nernst equation, and use day two for conductance and Kohlrausch’s law. Next, give day three to Faraday’s laws and batteries. After that, solve ten numericals from the NCERT exercise. For physics-side revision of the same week, read class 12 physics notes next. You can also browse the Chemistry category, the Class 11–12 notes hub and how to stop procrastinating.

Wrap-Up

Electrochemistry class 12 notes work best when you pair each formula with one solved numerical. Keep the Nernst equation and Faraday’s chain on a single page, and revise them daily. Then explore the Chemistry category, or return to the The Note homepage.