Electrochemistry Notes for NEET and JEE Main 2027: Key Points and PDF

Electrochemistry is Chapter 2 of NCERT Class 12 Chemistry Part I and part of Unit 7, Redox Reactions and Electrochemistry, in both the NEET UG and JEE Main syllabus. These notes cover galvanic and electrolytic cells, electrode potential, the Nernst equation, Gibbs energy, conductance, Kohlrausch’s law, Faraday’s laws, batteries, fuel cells and corrosion, with a free PDF.

Electrochemistry Notes for NEET and JEE Main

Section-wise key points from NCERT Chapter 2, formula and comparison tables, high-yield facts and common mistakes for 2027 revision.

Electrochemistry at a Glance

The topic list follows the NEET UG 2026 syllabus (NMC public notice dated 22 December 2025) and the JEE Main 2026 syllabus published by NTA, which stay the reference until the 2027 versions are out. The full lists are on the NEET syllabus and JEE Main syllabus pages. Faraday’s laws, the mercury and nickel-cadmium cells and corrosion are in the NCERT chapter but are not named separately in the syllabus text.

What are electrochemical cells?

An electrochemical cell links a redox reaction to electricity: a galvanic cell turns the Gibbs energy of a spontaneous reaction into electrical work, and an electrolytic cell uses electrical energy to drive a non-spontaneous reaction.

• Daniell cell: Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s), with a potential of 1.1 V when Zn2+ and Cu2+ are at unit concentration.
• If an opposing external voltage below 1.1 V is applied, the cell still works as a galvanic cell: zinc dissolves at the anode and copper deposits at the cathode.
• At exactly 1.1 V no current flows and the reaction stops.
• Above 1.1 V the reaction reverses: zinc deposits, copper dissolves and the device works as an electrolytic cell.

Galvanic Cells and Electrode Potential

Electrode potential is the potential difference that develops between an electrode and its electrolyte; when every species in the half-cell is at unit concentration it is the standard electrode potential.

• Each half-cell (redox couple) is a metal electrode in an electrolyte. The two solutions are joined by a salt bridge; it is not needed if both electrodes dip in the same electrolyte.
• IUPAC convention: standard reduction potentials are now called standard electrode potentials.
• Cell notation puts the anode on the left and the cathode on the right, with a single vertical line between metal and solution and a double line for the salt bridge, e.g. Cu(s)|Cu2+(aq)||Ag+(aq)|Ag(s).
• E°cell = E°right − E°left = E°cathode − E°anode. The emf is the cell potential when no current is drawn.
• The potential of a single half-cell cannot be measured, only the difference between two.
• Standard hydrogen electrode (SHE), Pt(s)|H2(g)|H+(aq): platinum coated with platinum black, H2 at 1 bar, H+ at 1 M. It is assigned zero potential at all temperatures.
• Inert electrodes such as platinum or gold only provide a surface for the reaction and for electron conduction.

Selected Standard Electrode Potentials at 298 K (NCERT Table 2.1)

• Going down the table, E° falls, the oxidising power of the species on the left decreases and the reducing power of the species on the right increases.
• A positive E° means the reduced form is more stable than hydrogen gas; Cu2+ is reduced more easily than H+, so copper does not dissolve in HCl. In nitric acid copper is oxidised by the nitrate ion, not by H+.
• A negative E° (as for zinc) means H+ can oxidise the metal.
• For the Daniell cell: E°cell = 0.34 − (−0.76) = 1.10 V.

Nernst Equation, Equilibrium Constant and Gibbs Energy

The Nernst equation gives the electrode or cell potential at any concentration; at equilibrium the cell potential becomes zero, which links E°cell to the equilibrium constant.

• R = 8.314 J K–1 mol–1 and F = 96487 C mol–1 (about 96500 C mol–1 for approximate work).
• The Daniell cell potential rises when [Cu2+] increases and falls when [Zn2+] increases.
• Use the same number of electrons n for both electrodes, e.g. n = 2 for Ni(s) + 2Ag+ → Ni2+ + 2Ag(s).
• Ecell is an intensive property, but ΔrG is extensive and depends on n: doubling the equation doubles ΔrG (−2FE becomes −4FE) while E stays the same.

Conductance of Electrolytic Solutions

Conductivity (κ) is the conductance of a unit cube of solution; molar conductivity (Λm = κ/c) is the conductance of the volume that contains one mole of electrolyte.

• Electronic (metallic) conductance is due to electrons and decreases with rise in temperature. Ionic (electrolytic) conductance increases with rise in temperature.
• Ionic conductivity depends on the nature of the electrolyte, the size of the ions and their solvation, the nature and viscosity of the solvent, concentration and temperature.
• Pure water has about 10–7 M H+ and OH– ions and a very low conductivity of 3.5 × 10–5 S m–1.
• Measurement uses alternating current (direct current would change the solution’s composition) and a conductivity cell with platinised platinum electrodes in a Wheatstone bridge.
• The cell constant is found with KCl solutions of known conductivity, e.g. 0.1 M KCl has κ = 1.29 S m–1 at 298.15 K.
• MacDiarmid, Heeger and Shirakawa discovered conducting polymers (polyacetylene); they got the 2000 Nobel Prize in Chemistry.

How do conductivity and molar conductivity change with concentration?

On dilution conductivity always decreases, because there are fewer ions per unit volume, while molar conductivity increases, because the volume holding one mole of electrolyte grows.

• The constant A depends on the electrolyte type: NaCl is a 1-1, CaCl2 a 2-1 and MgSO4 a 2-2 electrolyte; all electrolytes of one type have the same A.
• Kohlrausch’s law of independent migration of ions: Λ°m of an electrolyte is the sum of the limiting molar conductivities of its ions, Λ°m = ν+λ°+ + ν–λ°–.
• Kohlrausch’s observation: Λ°m(KX) − Λ°m(NaX) is nearly constant (about 23.4 S cm2 mol–1) for any anion X.
• Degree of dissociation of a weak electrolyte α = Λm/Λ°m, and Ka = cα2/(1 − α).
• Λ°m(HAc) = Λ°m(HCl) + Λ°m(NaAc) − Λ°m(NaCl); with NCERT’s data this gives 390.5 S cm2 mol–1.
• Limiting ionic conductivities in water at 298 K (S cm2 mol–1): H+ 349.6, OH– 199.1, Na+ 50.1, K+ 73.5, Cl– 76.3, Ca2+ 119.0, Mg2+ 106.0, SO4 2– 160.0, CH3COO– 40.9.

Electrolysis and Faraday’s Laws

In an electrolytic cell an external voltage drives the reaction; the amount of product depends on the charge passed, Q = It.

• First law: the amount of chemical reaction at an electrode is proportional to the quantity of electricity passed.
• Second law: the amounts of different substances liberated by the same quantity of electricity are proportional to their chemical equivalent weights (atomic mass ÷ electrons needed to reduce the cation).
• One faraday (F) is the charge on one mole of electrons: 6.02 × 10^23 × 1.6021 × 10–19 C = 96487 C mol–1.
• Ag+ needs 1F per mole, Mg2+ needs 2F and Al3+ needs 3F.
• Impure copper is refined by making it the anode; pure copper deposits at the cathode.
• Na and Mg are obtained by electrolysis of their fused chlorides; Al by electrolysis of alumina in the presence of cryolite.
• Inert electrodes (Pt, Au) only supply or remove electrons; reactive electrodes take part in the reaction.

Products of Electrolysis

Aqueous NaCl gives NaOH, Cl2 and H2 overall.

Batteries and Fuel Cells

Batteries are galvanic cells: primary cells work only once, secondary cells can be recharged, and fuel cells run as long as fuel is supplied.

Corrosion as an Electrochemical Process

Corrosion is the slow coating of a metal surface with its oxide or salts; rusting of iron behaves like a small electrochemical cell set up on the metal.

• Anode spot: 2Fe(s) → 2Fe2+ + 4e–, E°(Fe2+/Fe) = −0.44 V.
• Cathode spot: O2(g) + 4H+(aq) + 4e– → 2H2O(l), E° = 1.23 V. The H+ is believed to come from H2CO3 formed when CO2 dissolves in water.
• Overall: 2Fe + O2 + 4H+ → 2Fe2+ + 2H2O, E°cell = 1.67 V. Fe2+ is further oxidised to rust, Fe2O3.xH2O.
• Other examples: tarnishing of silver and the green coating on copper and bronze.
• Prevention: paint or chemicals such as bisphenol; coating with metals such as Sn or Zn; a sacrificial electrode of Mg or Zn that corrodes instead of the object.

Most-asked Facts for NEET and JEE Main: Electrochemistry

Revise these one-line facts from the NCERT chapter just before the exam.

• Daniell cell E°cell = 1.10 V; its Kc is about 2 × 10^37 at 298 K.
• SHE: Pt black, H2 at 1 bar, 1 M H+, E° = 0 at all temperatures.
• F2 is the strongest oxidising agent (+2.87 V); Li is the strongest reducing agent in water (−3.05 V).
• E°cell = E°cathode − E°anode; anode on the left in cell notation.
• Nernst at 298 K: Ecell = E°cell − (0.059/n) log Q.
• ΔrG° = −nFE°cell and ΔrG° = −RT ln K.
• 1 F = 96487 C mol–1 (about 96500).
• Conductivity falls and molar conductivity rises on dilution.
• Λm = Λ°m − A c^1/2 holds for strong electrolytes only.
• Λ°m of a weak electrolyte comes from Kohlrausch’s law; α = Λm/Λ°m.
• H+ (349.6) and OH– (199.1) have the highest limiting ionic conductivities in NCERT’s table.
• Cell constant G* = l/A = R × κ, unit m–1.
• Aqueous NaCl electrolysis gives H2 at the cathode and Cl2 at the anode (oxygen overpotential).
• Lead storage battery uses 38% H2SO4; both electrodes become PbSO4 on discharge.
• Mercury cell keeps a constant 1.35 V; dry cell gives about 1.5 V.
• H2-O2 fuel cell powered Apollo; about 70% efficient.
• Rusting: E°cell = 1.67 V; Mg or Zn act as sacrificial electrodes.

Common Mistakes to Avoid

Watch for these mix-ups when you answer questions on this chapter.

• Giving the anode the same sign in both cells. In a galvanic cell the anode is negative; in an electrolytic cell it is connected to the positive terminal. Oxidation always happens at the anode.
• Multiplying E° by the stoichiometric coefficient when an equation is doubled. Electrode potential is intensive; only ΔrG doubles.
• Writing E°cell = E°anode − E°cathode. It is cathode minus anode (right minus left).
• Saying conductivity increases on dilution. Conductivity decreases; molar conductivity increases.
• Extrapolating Λm to zero concentration for acetic acid. That works only for strong electrolytes.
• Forgetting the unit conversion: 1 S m2 mol–1 = 10^4 S cm2 mol–1, and 1 mol L–1 = 1000 mol m–3.
• Assuming O2 forms at the anode in aqueous NaCl electrolysis. Cl2 forms because of the overpotential of oxygen.
• Using direct current to measure the conductivity of a solution. AC is used so the composition does not change.
• Calling the mercury cell rechargeable. It is a primary cell; the lead storage and nickel-cadmium cells are secondary.

Electrochemistry Notes PDF and Related Pages

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• Source text: NCERT Chapter 2 (NCERT, Class 12 Chemistry Part I)

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• Some Basic Concepts of Chemistry
• Structure of Atom
• Chemical Kinetics
• Equilibrium
• Organic Chemistry: Some Basic Principles and Techniques
• Electrochemistry
• Amines
• Solutions
• Thermodynamics
• Coordination Compounds
• Haloalkanes and Haloarenes
• p-Block Elements
• Alcohols, Phenols and Ethers
• Hydrocarbons
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