Electrochemistry

Electrochemistry studies the interconversion of chemical and electrical energy through redox reactions. The PMDC MDCAT 2026 syllabus expects you to assign oxidation states, balance redox equations, distinguish galvanic from electrolytic cells, and use the standard hydrogen electrode (SHE) to compute cell 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.

Electrolytic cell

Non-spontaneous redox driven by an external power source. Used in electroplating, electrorefining of copper, and the chlor-alkali process. Anode is positive (oxidation still occurs there); cathode is negative.

Galvanic (voltaic) vs Electrolytic cell
Property	Galvanic / Voltaic	Electrolytic
Energy conversion	Chemical → Electrical	Electrical → Chemical
Reaction	Spontaneous (ΔG < 0; E°_cell > 0)	Non-spontaneous (ΔG > 0; needs input)
Anode (oxidation)	Negative electrode	Positive electrode
Cathode (reduction)	Positive electrode	Negative electrode
External power source	None — cell is the source	Required (battery or DC supply)
Two electrolyte compartments?	Yes — connected by salt bridge	Single compartment with electrolyte
Examples	Daniell cell, dry cell, lead-acid battery, fuel cell	Electroplating, electrorefining of Cu, electrolysis of H₂O / NaCl, chlor-alkali process
Common rule	Oxidation always at anode, reduction always at cathode — only the polarity flips

Faraday's laws of electrolysis

First law: mass deposited / liberated ∝ charge passed; m = ZIt where Z is the electrochemical equivalent.
Second law: for the same charge, masses deposited are proportional to their equivalent weights.
1 Faraday (F) = 96 500 C = charge of 1 mole of electrons. ZF = molar mass / n (n = electrons exchanged).

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. One Faraday of charge is approximately equal to:

9 650 C
96 500 C
965 000 C
6.022 × 10²³ C

1 F = N_A × e = (6.022×10²³)(1.602×10⁻¹⁹) ≈ 96 485 C, rounded to 96 500 C. It is the charge carried by one mole of electrons.

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⁻.
Galvanic cell: spontaneous, E°_cell > 0; electrolytic cell: non-spontaneous, requires external EMF.
SHE: Pt + H₂(1 atm) + H⁺(1 M) at 25°C; E° = 0.00 V by definition.
E°_cell = E°_cathode − E°_anode.
Faraday's law: m = ZIt; 1 F = 96 500 C = 1 mol e⁻.

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