Electrochemistry

Electrochemistry studies the interconversion of chemical and electrical energy through redox reactions. The PMDC MDCAT 2026 syllabus expects you to give the characteristics of a redox reaction, define oxidation and reduction in terms of oxidation-number change, balance redox equations, and use the standard hydrogen electrode (SHE) as the reference for standard electrode potentials. Expect 2-3 MCQs per paper.

PMC Table of Specifications. This chapter covers four PMDC subtopics — balancing chemical reactions, oxidation/reduction definitions, redox reactions, and the standard hydrogen electrode. Most MCQs target either the SHE convention or oxidation-state assignment.

Balancing Chemical Reactions

Every chemical equation must obey the law of conservation of mass (atoms balanced) and, in ionic form, conservation of charge. For redox reactions the simplest, fool-proof method is the half-reaction (ion-electron) method.

Half-reaction method (acidic medium)

Write separate oxidation and reduction half-reactions.
Balance atoms other than O and H first.
Balance O by adding H₂O.
Balance H by adding H⁺.
Balance charge by adding electrons (e⁻).
Multiply each half-reaction so electrons cancel; add the half-reactions.

For basic medium, add the same number of OH⁻ to both sides as there are H⁺ and combine with the H⁺ to form H₂O.

Common trap. Students often forget to verify charge balance after balancing atoms. A correctly balanced redox equation must balance both atoms and net charge on each side — check both before moving on.

Oxidation and Reduction

Modern definitions are based on electron transfer:

Oxidation: Loss of electrons; increase in oxidation number; addition of oxygen or removal of hydrogen.
Reduction: Gain of electrons; decrease in oxidation number; removal of oxygen or addition of hydrogen.
Oxidising agent: Species that causes oxidation (gets reduced itself). Examples: KMnO₄, K₂Cr₂O₇, Cl₂, O₃.
Reducing agent: Species that causes reduction (gets oxidised itself). Examples: H₂, C, Na, Zn, SnCl₂.

Rules for assigning oxidation numbers

Free elements have oxidation number 0 (e.g., Na, O₂, P₄).
Monoatomic ion = its charge (Na⁺ → +1, S²⁻ → −2).
Group I metals = +1, Group II metals = +2, Al = +3.
Hydrogen = +1 (with non-metals), −1 (in metal hydrides like NaH).
Oxygen = −2 (general), −1 in peroxides (H₂O₂), +2 in OF₂, −1/2 in superoxides (KO₂).
Sum of oxidation numbers = net charge of the species.

Mnemonic: OIL RIG. Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons). Or: LEO the lion says GER — Loss of Electrons = Oxidation, Gain of Electrons = Reduction.

Redox Reactions

A redox reaction is one in which electrons are transferred from a reducing agent to an oxidising agent. Oxidation and reduction always occur simultaneously: the electrons lost by one species must be gained by another.

Galvanic (voltaic) cell

Spontaneous redox reaction generates electrical energy. The Daniell cell is the prototype: Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s). Zinc is oxidised at the anode (negative electrode), copper ions are reduced at the cathode (positive electrode). A salt bridge maintains electrical neutrality. E°_cell = E°_cathode − E°_anode = +0.34 − (−0.76) = +1.10 V.

Anode and cathode

Anode: The electrode at which oxidation occurs. In a galvanic cell it is the negative electrode.
Cathode: The electrode at which reduction occurs. In a galvanic cell it is the positive electrode.
Electrode potential: The tendency of an electrode to lose or gain electrons when in contact with a solution of its own ions. It cannot be measured on its own, only against a reference.

Oxidation is always at the anode and reduction is always at the cathode — remember the vowel/consonant pair (an-ode/ox-idation, cath-ode/red-uction).

Standard Hydrogen Electrode (SHE)

An absolute electrode potential cannot be measured — only differences are accessible. The standard hydrogen electrode is therefore defined as the reference, with E° = 0.00 V.

Construction

A platinum electrode coated with platinum black is dipped into a 1 M H⁺ solution at 25°C while H₂ gas is bubbled at 1 atm pressure over it. The half-cell reaction is 2H⁺(aq, 1 M) + 2e⁻ ↔ H₂(g, 1 atm).

Standard conditions

Temperature = 25°C (298 K).
[H⁺] = 1 M.
P(H₂) = 1 atm.
E°(SHE) = 0.00 V (by convention).

Using SHE to find standard electrode potentials

Couple any half-cell with SHE. The voltmeter reading is the E° of that half-cell. Sign convention: electrodes that reduce H⁺ better than H₂ have E° > 0 (e.g. Cu²⁺/Cu = +0.34 V); electrodes that are oxidised more easily than H₂ have E° < 0 (e.g. Zn²⁺/Zn = −0.76 V).

High-yield exam point. Memorise the three SHE conditions: 1 M H⁺, 1 atm H₂, 25°C. Examiners frequently change one parameter and ask whether the cell potential is still 0.00 V (it is not).

Worked MCQs

Five MCQs that capture the high-yield testing patterns for electrochemistry. Read every explanation — the deeper concept lives there.

Q1. The oxidation number of manganese in KMnO₄ is:

K = +1, each O = −2 (four oxygens = −8). Total must equal 0: +1 + Mn + (−8) = 0 → Mn = +7. This is also why KMnO₄ is such a strong oxidising agent.

Q2. In the reaction Zn + Cu²⁺ → Zn²⁺ + Cu, which species is the reducing agent?

Zn
Cu²⁺
Zn²⁺
Cu

Zinc gives up two electrons (it is oxidised) and forces Cu²⁺ to gain electrons. The species that gets oxidised is, by definition, the reducing agent.

Q3. The standard hydrogen electrode (SHE) operates at:

1 M H⁺, 0°C, 1 atm H₂
1 M H⁺, 25°C, 1 atm H₂
0.1 M H⁺, 25°C, 1 atm H₂
1 M H⁺, 25°C, 760 mm H₂O

Standard conditions for SHE are 1 mol/dm³ H⁺, 1 atm H₂(g) and 298 K (25°C). E° is defined as exactly 0.00 V under these conditions.

Q4. For the Daniell cell with E°(Cu²⁺/Cu) = +0.34 V and E°(Zn²⁺/Zn) = −0.76 V, the cell EMF is:

0.42 V
−1.10 V
+1.10 V
+0.42 V

E°_cell = E°_cathode − E°_anode = +0.34 − (−0.76) = +1.10 V. A positive value confirms the reaction is spontaneous — consistent with our experience that zinc displaces copper from copper sulphate.

Q5. In the reaction MnO₂ + 4HCl → MnCl₂ + Cl₂ + 2H₂O, the element that is oxidised is:

Track the oxidation-number changes. Mn goes from +4 in MnO₂ to +2 in MnCl₂ — a decrease, so Mn is reduced. Chlorine goes from −1 in HCl to 0 in Cl₂ — an increase, so Cl is oxidised. H stays +1 and O stays −2 throughout.

Quick Recap

Oxidation = loss of electrons / increase in oxidation number; reduction = gain of electrons / decrease in oxidation number (OIL RIG).
Oxidising agent gets reduced; reducing agent gets oxidised.
Half-reaction method: balance atoms → balance O with H₂O → balance H with H⁺ → balance charge with e⁻.
Anode = oxidation (negative in a galvanic cell); cathode = reduction (positive in a galvanic cell).
SHE: Pt + H₂(1 atm) + H⁺(1 M) at 25°C; E° = 0.00 V by definition.
E°_cell = E°_cathode − E°_anode.
Assign oxidation numbers first: the atom whose number rises is oxidised, the one whose number falls is reduced.

Test yourself. Take a timed practice test or browse the topic-wise MCQs to lock these concepts in.