Electrochemistry — NEET Notes

Electrochemical cells — two directions of the same arrow

An electrochemical cell is any device that physically separates the oxidation and reduction halves of a redox reaction so that the electrons must travel through an external wire. Two families exist, and they are mirror images of each other. A galvanic cell (also called a voltaic cell) lets a spontaneous redox reaction do electrical work — chemical energy out, current in the circuit. An electrolytic cell reverses the arrow: it uses an external power source to force a non-spontaneous reaction. NCERT introduces them through the same hardware — a Daniell cell. Below its natural emf (1.1 V) it behaves as a galvanic cell; raise the external opposing voltage above 1.1 V and the same hardware becomes an electrolytic cell, with zinc plating onto the zinc rod and copper dissolving off the copper rod.

Four distinct cell types appear in the NEET syllabus. They share the same building blocks — two electrodes, an electrolyte, an electron path — but differ in what drives the reaction and what the cell is for.

Galvanic cells & electrode potential

The Daniell cell is NEET's archetypal galvanic cell. A zinc rod sits in 1 M ZnSO₄, a copper rod in 1 M CuSO₄, the two solutions are connected by a salt bridge (typically KCl or KNO₃ in agar), and the two rods are linked through an external voltmeter. The spontaneous reaction is

Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)    E°_cell = 1.10 V

The Daniell-cell reaction — NCERT's working example

The reaction is the sum of two half-reactions. At the zinc rod, zinc atoms surrender electrons and dissolve as Zn²⁺ — this is the oxidation half, and by IUPAC convention the zinc electrode is the anode. At the copper rod, Cu²⁺ ions pick up electrons from the external wire and deposit as metallic copper — this is the reduction half, and the copper electrode is the cathode. Electrons flow externally from zinc to copper; conventional current is in the opposite direction.

At the interface between any metal electrode and its electrolyte, two opposing processes compete: metal atoms try to escape into solution as ions, and metal ions in solution try to plate out onto the rod. At equilibrium, the two rates are equal and a potential difference exists between the rod and the surrounding solution. This is the electrode potential. When all species are at unit concentration (1 M for ions, 1 bar for gases) and 298 K, it is called the standard electrode potential (E°).

The convention for representing a galvanic cell is rigid: anode on the left, cathode on the right. A single vertical line shows the metal–electrolyte interface; a double vertical line shows the salt bridge. The Daniell cell is written

Zn(s) | Zn²⁺(aq, 1 M) || Cu²⁺(aq, 1 M) | Cu(s)

Cell notation — anode (left) | salt bridge (||) | cathode (right)

The cell emf is then calculated as E°_cell = E°_right − E°_left = E°_cathode − E°_anode. Both potentials must be looked up as reduction potentials; you do not flip the sign of the anode value yourself — the subtraction handles it. This is the convention NEET 2022 tested directly: among Zn → Fe, Fe → Cu, Cu → Ag and Ag → Cu, only the last (2 Ag + Cu²⁺) gives a negative E°_cell and therefore cannot occur.

Standard hydrogen electrode & electrochemical series

An individual electrode potential cannot be measured in isolation — only differences can. Chemistry needs a reference, and the IUPAC reference is the standard hydrogen electrode (SHE). A platinum strip coated with platinum black is dipped in a 1 M H⁺ solution and pure hydrogen gas at 1 bar is bubbled over it. By convention, the SHE is assigned E° = 0.00 V at all temperatures, for the half-reaction

H⁺(aq, 1 M) + e⁻ ⇌ ½ H₂(g, 1 bar)    E° = 0.00 V

The reference half-cell — every other E° is measured against this

To measure the standard electrode potential of any half-cell, you build a galvanic cell with the SHE as one half and the unknown half-cell as the other, all at standard conditions, and read the emf. Because E°_SHE = 0, the measured cell emf is the standard electrode potential of the unknown — sign included. The Cu²⁺|Cu half-cell gives +0.34 V; the Zn²⁺|Zn half-cell gives −0.76 V. The Daniell cell emf is therefore 0.34 − (−0.76) = 1.10 V, exactly what you measure.

Arranging every half-cell in order of its standard reduction potential gives the electrochemical series. Fluorine sits at the top with E° = +2.87 V — the strongest oxidising agent, the species most eager to be reduced. Lithium sits at the bottom at −3.05 V — the weakest oxidising agent (as Li⁺) and therefore the strongest reducing agent (as Li metal). Five rules drop straight out of the table:

• A more positive E° means a stronger oxidising agent on the left side of the half-reaction.
• A more negative E° means a stronger reducing agent on the right side.
• Any metal below hydrogen in the series (negative E°) can displace H⁺ from dilute acid — Zn dissolves in HCl, Cu does not.
• Any metal can displace from solution the metal of higher reduction potential — Zn displaces Cu²⁺, Cu displaces Ag⁺.
• For any candidate reaction, calculate E°_cell; if positive, the reaction proceeds spontaneously.

The Nernst equation — when concentrations are not 1 M

Standard electrode potentials assume unit concentration. Real cells almost never operate at unit concentration. The Nernst equation (Walther Nernst, 1889) gives the cell potential at any concentration. For the general electrode reaction Mⁿ⁺(aq) + ne⁻ → M(s),

E_cell = E°_cell − (RT/nF) ln Q

General form (any temperature)

At 298 K, substituting R = 8.314 J K⁻¹ mol⁻¹, F = 96,485 C mol⁻¹, and converting natural log to base-10, the equation collapses to the NEET-ready form

E_cell = E°_cell − (0.0591 / n) log Q

Nernst equation at 298 K — memorise this version

Here n is the number of electrons transferred in the balanced cell reaction, and Q is the reaction quotient — written for the cell reaction (products over reactants), with activities of solids and pure liquids set to 1, gas pressures in bar, and ion concentrations in mol L⁻¹. The Nernst expression has a simple consequence: increasing the concentration of products lowers the emf; increasing the concentration of reactants raises it.

For the Daniell cell at arbitrary concentrations, Q = [Zn²⁺]/[Cu²⁺], n = 2, and

E_cell = 1.10 − (0.0591 / 2) log ( [Zn²⁺] / [Cu²⁺] )

Daniell cell at 298 K

NEET 2017 exploited this: starting from a Daniell cell with [Zn²⁺] = 0.01 M and [Cu²⁺] = 1 M (giving E₁), the question flipped the concentrations (E₂). Because Q increases from 0.01 to 100, the log term flips sign, and E₂ < E₁. The correct answer was E₁ > E₂.

Equilibrium constant from the Nernst equation

At electrochemical equilibrium, the cell is "dead" — no net current flows, E_cell = 0, and the reaction quotient Q has reached the equilibrium constant K. Substituting into the Nernst equation gives a powerful bridge between electrochemistry and equilibrium:

0 = E°_cell − (0.0591 / n) log K   ⇒   log K = (n × E°_cell) / 0.0591

From Nernst to K at 298 K

The same E°_cell also fixes the Gibbs energy change. From the maximum non-PV work theorem in thermodynamics,

ΔG° = −nFE°_cell   and   ΔG° = −RT ln K

Three relations, one master equation

Combining: nFE°_cell = RT ln K. A single experimental measurement — the standard cell emf — therefore gives the Gibbs energy of the reaction and its equilibrium constant. For the Daniell cell, E°_cell = 1.10 V and n = 2, so ΔG° = −2 × 96,485 × 1.10 ≈ −212.3 kJ/mol, and log K ≈ 37.2 — K is astronomically large, confirming the reaction goes essentially to completion.

Conductance of electrolytic solutions

An electrolytic solution conducts electricity because dissolved ions migrate under an applied field. The quantitative description borrows three quantities from physics and adds one of its own. Resistance R (in ohms) of a conductor is R = ρ × l/A, where ρ is the resistivity, l the length and A the cross-sectional area. Conductance G = 1/R (in siemens, S, or ohm⁻¹). Conductivity κ (kappa, in S cm⁻¹) is the reciprocal of resistivity: κ = (1/R) × (l/A). The geometric factor l/A is called the cell constant (G\*), with units cm⁻¹.

The cell constant is measured by filling the conductance cell with a solution of known κ — typically standard KCl — measuring R, and rearranging κ = G × G*. NEET 2023 (Q.52) gave exactly this experiment: 0.01 M KCl with κ = 0.0210 S cm⁻¹ and cell resistance 60 Ω. Solving G* = κ × R = 0.0210 × 60 = 1.26 cm⁻¹.

For comparing electrolytes regardless of how much of them is dissolved, chemists use molar conductivity Λ_m — the conductance of all the ions produced by one mole of electrolyte:

Λ_m = κ × 1000 / c    (S cm² mol⁻¹, with c in mol L⁻¹)

Molar conductivity — concentration-corrected conductivity

Two patterns emerge as you dilute a solution. Conductivity κ falls with dilution — fewer ions per unit volume to carry charge. Molar conductivity Λ_m rises with dilution — each mole's worth of ions has more room, less interionic friction. For a strong electrolyte (KCl, HCl), Λ_m rises gradually and extrapolates linearly to a limit at infinite dilution called limiting molar conductivity Λ°_m (Debye-Hückel-Onsager treatment: Λ_m = Λ°_m − A√c). For a weak electrolyte (CH₃COOH, NH₄OH), Λ_m rises steeply at very low concentration because dilution drives further dissociation; Λ°_m cannot be obtained by extrapolation.

Kohlrausch's law of independent migration of ions

Friedrich Kohlrausch observed in 1875 that at infinite dilution, ions move independently of one another — the cation contributes one fixed amount to the molar conductivity and the anion another, regardless of who the counter-ion is. Stated formally:

Λ°_m = ν₊ λ°₊ + ν₋ λ°₋

Kohlrausch's law — ions migrate independently at infinite dilution

where ν₊ and ν₋ are the stoichiometric numbers of cation and anion in the formula unit, and λ°₊ and λ°₋ are their ionic limiting molar conductivities. Two NEET-ready applications follow.

Application 1 — limiting conductivity of a weak electrolyte. Λ°_m of CH₃COOH cannot be found by extrapolation. But it can be assembled from three strong-electrolyte Λ°_m values:

Λ°_m(CH₃COOH) = Λ°_m(CH₃COONa) + Λ°_m(HCl) − Λ°_m(NaCl)

The NEET 2021 PYQ identity — verify by cancelling Na⁺ and Cl⁻

Substituting 91.0 + 426.16 − 126.45 = 390.71 S cm² mol⁻¹ — exactly the answer asked in NEET 2021 (Q.62).

Application 2 — degree of dissociation of a weak electrolyte. The ratio α = Λ_m / Λ°_m gives the fraction of solute molecules that have actually dissociated. Once α is known, the dissociation constant follows from K_a = Cα² / (1 − α). NEET 2021 (Q.93) used this for 0.007 M acetic acid: Λ°_m = 400, Λ_m = 20, so α = 1/20, and K_a ≈ 1.75 × 10⁻⁵ — the textbook value for acetic acid.

Electrolytic cells & electrolysis

When a substance is decomposed by passing electricity through it, the process is called electrolysis, and the apparatus is an electrolytic cell. The chemistry has three rules. (1) Cations migrate to the cathode and are reduced. (2) Anions migrate to the anode and are oxidised. (3) When multiple species can react at an electrode, the one with the most favourable potential wins, modified by overpotential. NCERT's standard examples illustrate the third rule:

• Molten NaCl: only Na⁺ at the cathode (gives Na metal), only Cl⁻ at the anode (gives Cl₂). Used in the Down's process for industrial sodium.
• Aqueous NaCl (brine): at the cathode, H₂O is reduced in preference to Na⁺ (more positive reduction potential) — gives H₂. At the anode, Cl⁻ is oxidised in preference to H₂O because of high O₂ overpotential — gives Cl₂. Sodium hydroxide accumulates. This is the chlor-alkali process.
• Dilute H₂SO₄ on platinum: at the cathode, H⁺ is reduced — H₂. At the anode, water is oxidised — O₂. The net reaction is electrolysis of water. NEET 2020 (Q.137) tested this directly.
• Aqueous CuSO₄ with copper electrodes: Cu dissolves at the anode, Cu plates out at the cathode. This is the basis of electrorefining of copper.

Faraday's laws of electrolysis

Michael Faraday quantified electrolysis in 1834. His two laws — one of NEET's most reliable scoring grounds — state:

First law: the amount of any substance liberated at an electrode is proportional to the quantity of electricity passed through the electrolyte.

m = Z × Q    and    Q = I × t    so    m = Z × I × t

Faraday's first law — Z is the electrochemical equivalent

Here m is the mass deposited or liberated (g), Q the charge passed (coulombs), I the current (amperes), t the time (seconds), and Z the electrochemical equivalent — mass deposited by 1 coulomb (g/C).

Second law: when the same quantity of electricity is passed through different electrolytes, the masses of the substances liberated are proportional to their equivalent weights. The crucial constant is the Faraday F = 96,485 C/mol — the charge of one mole of electrons. One Faraday deposits one mole of a singly-charged ion (Na, Ag), half a mole of a doubly-charged ion (Cu, Zn), or one-third of a mole of a triply-charged ion (Al).

Batteries — primary and secondary

A battery is a galvanic cell or stack of cells packaged for portable use. NCERT splits them into two families. Primary batteries are one-shot — once the reactants are consumed, the cell is dead. Secondary batteries are rechargeable — the discharge reaction is reversed by passing current in the opposite direction. Two of each are NEET-relevant.

Fuel cells — open-ended galvanic cells

A fuel cell is a galvanic cell that is continuously fed with reactants instead of being sealed with a finite charge. The cell never "discharges" in the battery sense — it keeps producing electricity as long as fuel and oxidant flow in. The flagship example is the hydrogen-oxygen fuel cell, used in the Apollo space programme and now in fuel-cell vehicles.

Two porous carbon electrodes infused with platinum or palladium catalyst dip into a hot aqueous KOH electrolyte. H₂ flows past the anode, O₂ past the cathode. The half-reactions are

Anode:   2 H₂(g) + 4 OH⁻(aq) → 4 H₂O(l) + 4 e⁻
Cathode:   O₂(g) + 2 H₂O(l) + 4 e⁻ → 4 OH⁻(aq)
Overall:   2 H₂(g) + O₂(g) → 2 H₂O(l)

H₂–O₂ fuel cell in alkaline electrolyte

The cell has three remarkable virtues NEET likes to flag. Efficiency exceeds 70% — far above thermal power plants (~40%). Pollution-free — the only exhaust is water. Quiet and continuous — no moving parts, no scheduled discharge. The catch is the cost of platinum catalysts and the difficulty of safely storing H₂.

Corrosion — galvanic cells you didn't ask for

Corrosion is the slow electrochemical destruction of a metal by environmental redox. The textbook case is the rusting of iron. Two tiny half-cells set up spontaneously on the metal surface wherever a film of moisture is present. At an anodic site, iron oxidises:

Anode:   2 Fe(s) → 2 Fe²⁺(aq) + 4 e⁻    E° = +0.44 V (oxidation)

Iron dissolving as Fe²⁺

Electrons travel through the metal itself to a cathodic site, where dissolved oxygen — concentrated in carbonic-acid-laden surface water — is reduced:

Cathode:   O₂(g) + 4 H⁺(aq) + 4 e⁻ → 2 H₂O(l)    E° = +1.23 V

Oxygen reduction in acidic surface water

The overall E°_cell is +1.67 V — strongly positive — which is why iron rusts so readily. The Fe²⁺ ions diffuse out, react with O₂ to form Fe³⁺, and precipitate as hydrated ferric oxide Fe₂O₃·xH₂O — the orange flaky rust. Because rust occupies more volume than the parent iron, the protective layer cracks and exposes fresh metal — corrosion is self-accelerating.

NCERT lists four prevention strategies, every one of which targets the electrochemical circuit:

• Barrier coatings — paint, oil, lacquer — keep water and oxygen away.
• Galvanisation — coating with zinc. Zinc has a more negative E° than iron and corrodes preferentially, sacrificially.
• Cathodic protection — connecting iron to a more electropositive "sacrificial anode" (Mg or Zn block), used on pipelines and ship hulls.
• Alloying — stainless steel (Fe + Cr + Ni) forms a self-healing chromium oxide film that blocks the anodic reaction.