NEET Electrochemistry Notes by Physics Wallah

Electrochemistry links redox chemistry with electrical energy, and NEET questions often combine the two. One question may ask you to identify the anode, the next to find a cell potential under non-standard conditions, and another to calculate the mass deposited after a given current has passed. Most errors come from unbalanced equations, a wrong value of n or a formula applied without checking its conditions.

Physics Wallah’s notes on electrochemistry organise the chapter into manageable topics, from cell construction and electrode potentials to conductance, Kohlrausch’s law and electrolysis. Check the units, balance the reaction before finding n and keep the key formulas handy so you can apply them quickly and accurately in the exam.

Electrochemical Cells

Electrochemical cells convert chemical energy into electrical energy, or use electrical energy to drive chemical reactions.

• Galvanic (voltaic) cell: a spontaneous redox reaction produces electrical energy.

• Electrolytic cell: an external electrical supply drives a non-spontaneous reaction.

In both types, oxidation occurs at the anode and reduction occurs at the cathode. The electrode signs differ: the anode is negative in a galvanic cell and positive in an electrolytic cell.

Daniell Cell

A Daniell cell has a zinc electrode and a copper electrode. Zinc is oxidised at the anode, and copper ions are reduced at the cathode.

• Anode: Zn(s) → Zn²⁺(aq) + 2e⁻

• Cathode: Cu²⁺(aq) + 2e⁻ → Cu(s)

• Overall: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)

Electrons flow from zinc to copper through the external circuit. The zinc electrode loses mass, while copper is deposited on the copper electrode.

Salt Bridge

A salt bridge contains an electrolyte that does not react with the substances in the cell. It:

• maintains electrical neutrality in both half-cells

• completes the internal ionic circuit

• minimises liquid junction potential

Anions migrate towards the anode compartment, where positive ions are produced. Cations migrate towards the cathode compartment, where positive ions are consumed.

The electrolyte must be chosen carefully. KCl is unsuitable when silver ions are present because AgCl precipitates:

Ag⁺(aq) + Cl⁻(aq) → AgCl(s)

Cell Representation and Inert Electrodes

A galvanic cell is written with the anode on the left and the cathode on the right:

Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)

A single vertical line (|) represents a phase boundary, and a double vertical line (||) represents the salt bridge.

An inert electrode, such as platinum, provides a conducting surface without taking part in the overall reaction. It is used when the reacting species cannot act as a solid electrode themselves.

Electrode Potential and Electrochemical Series

Electrode potential indicates the tendency of an electrode reaction to occur as oxidation or reduction. Standard electrode potentials are measured relative to the standard hydrogen electrode (SHE), whose standard reduction potential is defined as zero.

For the same electrode under the same conditions:

E°(ox) = −E°(red)

• A higher standard reduction potential means a stronger tendency to be reduced, so the oxidised species is a stronger oxidising agent.

• A more negative standard reduction potential means the reduced species is a stronger reducing agent.

The standard cell potential is calculated from reduction potentials:

E°(cell) = E°(cathode) − E°(anode)

A positive E°(cell) means the reaction is spontaneous under standard conditions.

Nernst Equation and Gibbs Energy

The Nernst equation gives the cell potential under non-standard conditions. At 298 K:

E(cell) = E°(cell) − (0.0591 / n) log Q

For other temperatures, use the general form:

E(cell) = E°(cell) − (RT / nF) ln Q

Here, n is the number of electrons transferred in the balanced overall reaction, and Q is the reaction quotient.

At equilibrium, E(cell) = 0 and Q = K, so at 298 K:

E°(cell) = (0.0591 / n) log K

The relationship between Gibbs energy and cell potential is:

• ΔG = −nF E(cell)

• ΔG° = −nF E°(cell)

A positive cell potential corresponds to a negative ΔG, indicating a spontaneous reaction. At equilibrium, ΔG = 0 and E(cell) = 0.

Remember: Balance the redox reaction before finding n. Cell potential is an intensive property, so it does not change when the balanced equation is multiplied by a coefficient.

Conductance of Electrolytic Solutions

Metallic conduction occurs through electrons, whereas electrolytic conduction occurs through the movement of ions. As temperature rises, metallic conductivity generally decreases, while electrolytic conductivity generally increases because ion mobility improves.

• Conductance: G = 1/R

• Resistance: R = ρ l / A

• Conductivity: κ = 1/ρ

• Cell constant: l / A

• Conductivity from conductance: κ = G × cell constant

Molar and Equivalent Conductivity

Molar conductivity represents the conducting ability of all the ions produced by one mole of an electrolyte. When κ is in S cm⁻¹ and concentration is in mol L⁻¹:

• Λm = κ × 1000 / C

• Λeq = κ × 1000 / N (N is normality)

On dilution:

• Conductivity decreases, because the number of ions per unit volume falls.

• Molar and equivalent conductivity generally increase, because interionic interactions decrease. For weak electrolytes, greater ionisation also contributes.

Kohlrausch’s Law

Kohlrausch’s law states that at infinite dilution, each ion contributes independently to the limiting molar conductivity of an electrolyte.

Λm° = ν₊ λ₊° + ν₋ λ₋°

Here, ν₊ and ν₋ are the numbers of cations and anions in the formula unit, and λ° is the limiting ionic conductivity of each ion.

For calcium chloride:

Λm°(CaCl₂) = λ°(Ca²⁺) + 2 λ°(Cl⁻)

The law is used to find the limiting molar conductivity, the degree of dissociation and the ionisation constant of weak electrolytes.

Degree of dissociation: α = Λm / Λm°

Faraday’s Laws of Electrolysis

First law: The mass of a substance deposited or liberated during electrolysis is directly proportional to the charge passed.

• Charge: Q = It

• Mass deposited: m = M I t / (n F)

Here, m is the mass deposited, M the molar mass, I the current, t the time, n the number of electrons required per ion and F the Faraday constant (about 96500 C mol⁻¹).

Second law: When the same charge passes through different electrolytes, the masses deposited are proportional to their equivalent masses.

m₁ / m₂ = (M₁/n₁) / (M₂/n₂)

For numericals, convert time into seconds and find n from the electrode reaction.

Products of Electrolysis

The products depend on the ions present, their discharge tendencies, the electrode material and the electrolyte concentration.

• Molten NaCl: sodium at the cathode, chlorine at the anode.

• Aqueous NaCl (concentrated brine): hydrogen at the cathode, chlorine at the anode.

• Dilute H₂SO₄ (inert electrodes): hydrogen and oxygen.

• Aqueous CuSO₄ (inert electrodes): copper at the cathode, oxygen generally at the anode.

• Aqueous CuSO₄ (copper electrodes): copper dissolves from the anode and is deposited at the cathode.

Batteries, Fuel Cells and Corrosion

Primary batteries are used without recharging, while secondary batteries can be recharged and reused. Examples include the Leclanché cell, mercury cell, lead-storage battery and nickel-cadmium cell.

The mercury cell is a primary cell with a nearly constant voltage of about 1.35 V. A hydrogen-oxygen fuel cell converts chemical energy into electrical energy, with water as the overall product.

Corrosion is the gradual deterioration of a metal through chemical or electrochemical reactions with its surroundings. Galvanisation protects iron by coating it with zinc, which oxidises preferentially and shields the iron from corrosion.

Electrochemistry becomes easier when you treat each formula as a tool with its own conditions: Nernst and ΔG for cell potential, κ and Λm for conductance, Kohlrausch’s law for dilution and Faraday’s laws for electrolysis. Balance every reaction first, track your units and practise a mix of conceptual and numerical questions. PW’s Chemistry notes can support your NEET revision with focused, topic-wise practice.

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