Chemical cells are devices that convert chemical energy directly into electrical energy through spontaneous electrochemical reactions. From the simple voltaic cell to the rechargeable lead-acid battery in a military vehicle, cells are the backbone of portable power in defence operations — powering radios, night-vision devices, missile guidance systems, and submarine propulsion. Understanding how cells work requires grasping oxidation, reduction, electrodes, and electrolytes.
📌 What to expect in NDA (based on 2022–2025 pattern): (1) Oxidation and reduction — definitions, electron transfer, OIL RIG mnemonic; (2) Voltaic cell — Zn-Cu setup, anode (+/−), EMF ≈ 1.1 V; (3) Dry cell (Leclanché) — composition, EMF ≈ 1.5 V, local action, polarisation; (4) Lead-acid accumulator — composition, EMF per cell ≈ 2 V, charging/discharging reactions; (5) Nickel-cadmium cell — composition, EMF ≈ 1.2 V, advantages for defence; (6) Primary vs secondary cell distinction — rechargeability; (7) Faraday's laws of electrolysis — mass deposited proportional to charge.
The language of all electrochemical cells — losing and gaining electrons
All chemical cells operate through redox reactions — one substance loses electrons (oxidation) while another gains electrons (reduction). These happen simultaneously — you cannot have one without the other. The driving force for this electron flow is the difference in electrode potentials, which appears as the cell's EMF.
⚡ Oxidation, Reduction & Key Definitions
OIL RIG (Memory device):
OIL = Oxidation Is Loss (of electrons)
RIG = Reduction Is Gain (of electrons)
Oxidation: A → A⁺ + e⁻ (loses electrons; oxidation state increases)
Reduction: B⁺ + e⁻ → B (gains electrons; oxidation state decreases)
In a cell:
ANODE = negative electrode = where OXIDATION occurs
CATHODE = positive electrode = where REDUCTION occurs
Memory: AN OX = ANnode OXidation
RED CAT = REDuction at CAThode
Electrolyte: ionic conductor (solution or molten salt) — carries ions between electrodes
Electrode potential: tendency of electrode to gain/lose electrons
Standard Hydrogen Electrode (SHE): reference = 0 V
More positive electrode potential → better oxidising agent (cathode)
More negative electrode potential → better reducing agent (anode)
EMF of cell = E_cathode − E_anode (= reduction potential difference)
Electrochemical series (standard reduction potentials, selective):
Li⁺/Li: −3.04 V (strongest reducing agent — anode)
Zn²⁺/Zn: −0.76 V
Fe²⁺/Fe: −0.44 V
H⁺/H₂: 0.00 V (reference)
Cu²⁺/Cu: +0.34 V
Ag⁺/Ag: +0.80 V
Au³⁺/Au: +1.50 V (strongest oxidising agent — cathode)
Key rule: A metal higher in the electrochemical series (more negative reduction potential) displaces a metal lower in the series from its salt solution. Zinc displaces copper from CuSO₄ because Zn is higher (more reactive) than Cu.
Non-electrolyte: no ionisation — sugar, alcohol, urea
Conductivity increases with concentration (up to a point)
Metals: increase R with T; electrolytes: decrease R with T
📝 TOPIC-WISE PYQ
Electrochemistry Basics — NDA Pattern Questions
Q1. In a galvanic cell, oxidation occurs at the:
(a) Cathode (b) Anode (c) Electrolyte (d) Both electrodes equally
Answer: (b) Anode
In all electrochemical cells (galvanic and electrolytic), oxidation always occurs at the anode and reduction at the cathode. Memory: AN OX, RED CAT. In a galvanic cell, the anode is the negative terminal (source of electrons).
Q2. Which of the following is a strong electrolyte?
Answer: (c) Sulphuric acid (H₂SO₄)
Strong electrolytes are fully ionised in solution. H₂SO₄, HCl, NaOH, NaCl are strong electrolytes. Acetic acid (CH₃COOH) and ammonia (NH₃) are weak electrolytes (partially ionised). Glucose is a non-electrolyte (does not ionise).
2. Primary Cells
2.1
Voltaic Cell (Galvanic Cell)
The first chemical cell — Zinc and Copper in sulphuric acid
The Voltaic cell (invented by Alessandro Volta, 1800) is the simplest galvanic cell. It consists of a zinc electrode (anode, −) and a copper electrode (cathode, +) dipped in dilute sulphuric acid (electrolyte). The zinc dissolves (oxidation) and hydrogen is released at copper (reduction), creating a potential difference of approximately 1.1 V.
⚡ Voltaic Cell — Reactions & Key Points
Electrode reactions:
Anode (Zn, negative): Zn → Zn²⁺ + 2e⁻ (oxidation)
Cathode (Cu, positive): 2H⁺ + 2e⁻ → H₂↑ (reduction)
Overall: Zn + H₂SO₄ → ZnSO₄ + H₂
EMF = E_cathode − E_anode = 0.00 − (−0.76) = 0.76 V
(Theoretical; actual ≈ 1.1 V in practice with copper cathode)
Defects of Voltaic Cell:
1. Local action: impurities in zinc react with H₂SO₄ → wasteful
Cure: amalgamation of zinc (coating with mercury)
2. Polarisation: H₂ bubbles form on copper cathode → insulates it
→ reduces EMF and increases internal resistance
Cure: use a depolariser (KMnO₄, MnO₂) to oxidise H₂ to H₂O
Amalgamation of zinc (rubbing with mercury) removes impurities from the zinc surface, preventing local action and making the zinc electrode more uniform. This significantly extends cell life.
Fig. 1 — Voltaic Cell. Zinc anode (green) dissolves, releasing electrons (oxidation). Electrons flow through external circuit to copper cathode where H⁺ ions are reduced to H₂ gas (bubbles). EMF ≈ 1.1 V. Defects: local action (zinc impurities) and polarisation (H₂ bubbles).
2.2
Dry Cell (Leclanché Cell)
The everyday battery — portable, sealed, and still used in billions of devices
The dry cell (carbon-zinc or Leclanché cell) is the most widely used primary cell. It is "dry" because the electrolyte is a paste, not a liquid — making it portable and spill-proof. Used in torches, clocks, remote controls, and field radios.
🔌 Dry Cell Composition
Anode (−): Zinc cylinder (outer casing)
Cathode (+): Carbon rod (central)
Electrolyte: Paste of NH₄Cl + ZnCl₂ + MnO₂
Depolariser: MnO₂ — oxidises H₂ to prevent polarisation
EMF: ≈ 1.5 V (fresh cell)
Sealed — no spillage; portable
🚫 Limitations of Dry Cell
Not rechargeable — once discharged, must be discarded
EMF drops with use as Zn is consumed
Cannot deliver high currents for long (high internal resistance)
Heavy metals (Zn, Mn) — environmental disposal concern
Q1. In a Leclanché (dry) cell, the role of MnO₂ is:
(a) To act as the anode (b) To act as the electrolyte (c) To prevent polarisation (depolariser) (d) To increase EMF
Answer: (c) To prevent polarisation (depolariser)
Polarisation occurs when hydrogen gas accumulates on the cathode, creating a back-EMF and increasing internal resistance. MnO₂ acts as a depolariser — it oxidises H₂ to H₂O, keeping the cathode clear. Without MnO₂, the dry cell would rapidly lose its EMF.
Q2. The EMF of a fresh dry cell (Leclanché cell) is approximately:
(a) 2.0 V (b) 1.5 V (c) 1.2 V (d) 12 V
Answer: (b) 1.5 V
A fresh dry cell (carbon-zinc or Leclanché type) gives an EMF of approximately 1.5 V. Lead-acid cell: ≈ 2 V per cell. Nickel-cadmium: ≈ 1.2 V. Standard car battery (12 V) = 6 lead-acid cells in series.
Q3. Local action in a voltaic cell is prevented by:
(a) Adding depolariser (b) Amalgamating the zinc electrode (c) Replacing H₂SO₄ with HCl (d) Using copper anode
Answer: (b) Amalgamating the zinc electrode
Local action is caused by impurities in zinc reacting with the acid even when no external current is drawn. Amalgamation (coating zinc with mercury) removes surface impurities and makes the zinc surface uniform — preventing this wasteful local reaction. Depolarisers prevent polarisation, which is a separate problem.
🤔 TRICKY QUESTIONS
Primary Cells — Conceptual Traps
T1. In a voltaic cell, which electrode is the anode — is it positive or negative? Students often get confused. Explain clearly.
In a galvanic cell: Anode = NEGATIVE (−).
This contradicts electrolytic cell convention (where anode is +), confusing students. In a galvanic cell: Zinc (anode) releases electrons → electrons flow out of zinc through external circuit → zinc is the electron-rich, negative terminal. Copper (cathode) receives electrons → positive terminal. The rule is always: Oxidation at Anode. But in galvanic cells, the anode is negative (source of current); in electrolytic cells, the anode is positive (connected to +ve terminal of power supply).
T2. Why does the EMF of a dry cell drop when a large current is drawn, but recover slightly when the load is removed?
Internal resistance causes voltage drop; electrochemical equilibrium partially restores it.
Under load: V = EMF − Ir (terminal voltage drops due to internal resistance r). When a large current I flows, the voltage drop Ir is large → V decreases. When load is removed: no current flows → no Ir drop → terminal voltage = EMF again. Also, during rest, the electrolyte concentration redistributes slightly and depolariser action partially reverses — so the cell "recovers" some capacity. This is why switching off a torch briefly and back on sometimes appears to restore brightness momentarily.
3. Secondary Cells (Rechargeable Batteries)
3.1
Lead-Acid Accumulator
The most common rechargeable battery — in every car, truck, and submarine
The lead-acid accumulator (invented by Gaston Planté, 1859) is the most widely used secondary cell. Each cell has an EMF of approximately 2 V, and a 12 V car battery consists of 6 cells in series. It is heavily used in military vehicles, submarines, and UPS systems.
⚡ Lead-Acid Accumulator — Composition & Reactions
Composition:
Anode (−): Lead (Pb) plates
Cathode (+): Lead dioxide (PbO₂) plates
Electrolyte: Dilute sulphuric acid (H₂SO₄, density ≈ 1.28 g/mL when charged)
DISCHARGE reactions (cell produces current):
Anode (Pb): Pb + SO₄²⁻ → PbSO₄ + 2e⁻ (oxidation)
Cathode (PbO₂): PbO₂ + 4H⁺ + SO₄²⁻ + 2e⁻ → PbSO₄ + 2H₂O (reduction)
Overall: Pb + PbO₂ + 2H₂SO₄ → 2PbSO₄ + 2H₂O
On discharge: H₂SO₄ consumed → density of electrolyte decreases
PbSO₄ formed on both electrodes
CHARGE reactions (pass current to restore):
Reverse of above: PbSO₄ converted back to Pb and PbO₂
H₂SO₄ regenerated → density of electrolyte increases back to 1.28
EMF: ≈ 2.0 V per cell
12 V car battery = 6 cells in series (6 × 2 = 12 V)
State of charge check: density of H₂SO₄ (hydrometer test)
Fully charged: density ≈ 1.28 g/mL
Fully discharged: density ≈ 1.10 g/mL
Memory for discharge: Both electrodes become PbSO₄ (white) during discharge — "both go white when discharged." The H₂SO₄ density falls (check with a hydrometer). Charging reverses this completely.
Fig. 2 — Lead-acid accumulator. Pb plates (anode, grey) and PbO₂ plates (cathode, red-brown) in dilute H₂SO₄. On discharge: both form PbSO₄; H₂SO₄ consumed (density falls). On charging: original electrodes restored; H₂SO₄ regenerated (density rises to 1.28 g/mL).
🔸 Advantages of Lead-Acid Battery
Rechargeable — hundreds of cycles possible
High current output — can start car engines (high surge current)
Stable voltage during discharge (2 V/cell)
Reliable and well-understood technology
State of charge easily checked (hydrometer — density test)
🚫 Disadvantages
Very heavy (lead is dense) — poor energy/weight ratio
Contains toxic lead and corrosive H₂SO₄ — disposal hazard
Cannot be over-discharged (sulphation damages plates)
Self-discharge during storage
Produces H₂ gas on charging — explosion risk if sealed
3.2
Nickel-Cadmium (NiCd) Cell
Compact, rechargeable, and robust — the choice for field electronics
The nickel-cadmium (NiCd) cell is a secondary cell preferred where weight, cycle life, and reliability matter over cost. It powers portable field radios, night-vision devices, and emergency systems. It can be deeply discharged and recharged with minimal damage.
⚡ Nickel-Cadmium Cell — Composition & Reactions
Composition:
Anode (−): Cadmium (Cd)
Cathode (+): Nickel hydroxide (NiO(OH) / Ni(OH)₂)
Electrolyte: Potassium hydroxide (KOH) — alkaline
DISCHARGE:
Anode: Cd + 2OH⁻ → Cd(OH)₂ + 2e⁻ (oxidation)
Cathode: 2NiO(OH) + 2H₂O + 2e⁻ → 2Ni(OH)₂ + 2OH⁻ (reduction)
EMF: ≈ 1.2 V per cell
Advantages of NiCd:
✓ Rechargeable (500–1000 cycles)
✓ Good performance at low temperatures
✓ Can be discharged completely without permanent damage
✓ Low internal resistance — good for high current loads
✓ Long shelf life and robust construction
Disadvantage:
✗ Memory effect — capacity reduces if not fully discharged before charging
✗ Contains cadmium (toxic heavy metal — serious environmental hazard)
✗ Replaced in many applications by Li-ion (no memory effect, higher energy density)
The memory effect of NiCd cells: if repeatedly recharged without full discharge, the battery "remembers" the shorter cycle and delivers only that partial capacity. Solution: fully discharge before recharging (periodic deep discharge).
⚓ Defence Use of NiCd Cells: Field radios (Motorola type), night-vision goggles, portable GPS units, and emergency lighting systems in Indian military use NiCd batteries because they maintain performance in extreme cold (Himalayas, Arctic ops) and can be fully discharged without damage — critical when deep field operations may leave no charging opportunity for extended periods. Submarine emergency systems also rely on NiCd due to their sealed construction and H₂ safety.
📝 TOPIC-WISE PYQ
Secondary Cells — NDA Pattern Questions
Q1. The electrolyte in a lead-acid accumulator is:
(a) H₂SO₄ (b) HCl (c) NH₄Cl (d) KOH
Answer: (a) H₂SO₄ (dilute sulphuric acid)
The lead-acid accumulator uses dilute H₂SO₄ as electrolyte (density ≈ 1.28 g/mL when fully charged, drops to ≈ 1.10 when discharged). NiCd cells use KOH. Dry cells use NH₄Cl paste.
Q2. A 12 V car battery consists of how many lead-acid cells in series?
(a) 4 (b) 6 (c) 8 (d) 12
Answer: (b) 6
Each lead-acid cell gives EMF ≈ 2 V. Total battery voltage = 12 V. Number of cells = 12/2 = 6 cells in series. (Truck batteries are 24 V = 12 lead-acid cells in series.)
Q3. During the discharge of a lead-acid battery, what happens to the density of the electrolyte?
(a) Increases (b) Remains constant (c) Decreases (d) First increases then decreases
Answer: (c) Decreases
During discharge, H₂SO₄ is consumed (reacts with Pb and PbO₂ to form PbSO₄ + H₂O). As H₂SO₄ concentration drops, the density of the electrolyte decreases. This is how a hydrometer measures the state of charge — lower density = more discharged.
Q4. Which secondary cell is known for the memory effect?
Answer: (c) Nickel-cadmium
NiCd cells suffer from the memory effect — if repeatedly recharged without full discharge, the battery delivers only the partial capacity it was "trained" to use. Lead-acid and lithium-ion cells do not exhibit this effect significantly.
🤔 TRICKY QUESTIONS
Secondary Cells — Deep Reasoning
T1. During charging of a lead-acid battery, which electrode becomes the anode (where oxidation occurs)?
During charging: the PbSO₄ on the positive plate (PbO₂ plate) is oxidised back to PbO₂.
During charging, an external power supply reverses the cell reactions. The plate that was the cathode during discharge (PbO₂ plate, + terminal) now becomes the anode of the electrolytic process — PbSO₄ is oxidised back to PbO₂. The negative plate (Pb) receives electrons and PbSO₄ is reduced back to Pb. In short: during charging, the conventional electrodes switch roles (as electrolytic cell), but we still identify them by their physical position (+/−) in the circuit.
T2. A hydrometer is placed in a car battery and reads a density of 1.10 g/mL. What does this indicate, and what should be done?
The battery is nearly fully discharged (density should be ≈ 1.28 g/mL when charged).
The hydrometer measures the density of H₂SO₄. Fully charged: ≈ 1.28 g/mL. Fully discharged: ≈ 1.10 g/mL. A reading of 1.10 means most H₂SO₄ has been consumed (converted to PbSO₄ + H₂O). Action: the battery must be recharged immediately. Leaving it in this deeply discharged state causes sulphation — the PbSO₄ crystals grow large and become hard, permanently reducing capacity.
4. Faraday's Laws of Electrolysis
4.1
Faraday's Laws & Electroplating
Quantitative rules for electrolysis — mass deposited is proportional to charge
⚡ Faraday's Laws of Electrolysis
First Law:
Mass of substance deposited (m) ∝ Charge passed (Q)
m = Z × Q = Z × I × t
Z = electrochemical equivalent (g/C)
Second Law:
For same charge passed through different electrolytes:
m₁/m₂ = E₁/E₂
E = equivalent weight = Atomic mass / Valency
Combined form:
m = (E/F) × Q = (E × I × t) / F
F = Faraday constant = 96,485 C/mol ≈ 96,500 C/mol ≈ 10⁵ C/mol
One Faraday = charge of 1 mole of electrons
E = equivalent weight (g/equivalent)
Electroplating:
Object (cathode) to be plated
Metal to be deposited (anode) — dissolves to replenish ions
Electrolyte: salt of plating metal
Applications: silver plating (jewellery), nickel plating (tools), chromium plating (automotive)
F = 96,500 C is the charge of one mole of electrons (6.022 × 10²³ electrons × 1.6 × 10⁻¹⁹ C/electron ≈ 96,500 C). Faraday's laws connect chemistry and electricity quantitatively.
🔧 Electroplating Applications
Chrome plating: car bumpers, taps — corrosion resistance
Silver plating: cutlery, jewellery — aesthetics
Zinc plating (galvanising): iron pipes — rust prevention
Q1. A current of 2 A is passed through a CuSO₄ solution for 30 minutes. The mass of copper deposited (E = 32, F = 96,500) is approximately:
(a) 1.19 g (b) 2.38 g (c) 0.60 g (d) 4.76 g
Answer: (a) 1.19 g
m = EIt/F = 32 × 2 × 1800 / 96,500 = 115,200 / 96,500 ≈ 1.19 g. (t = 30 × 60 = 1800 s. Equivalent weight of Cu = 64/2 = 32.)
Q2. In electrolytic refining of copper, the impure copper is made the:
(a) Cathode (b) Anode (c) Electrolyte (d) Both anode and cathode
Answer: (b) Anode
In electrolytic refining: impure metal = anode (it dissolves); pure metal = cathode (it deposits). Impurities either dissolve into solution or fall as anode sludge. Pure copper deposits on the cathode with 99.99% purity. This is the standard industrial method.
5. Fuel Cells & Cell Comparison
5.1
Hydrogen Fuel Cell & Complete Cell Comparison
The future of clean power — and a snapshot of all cell types
A fuel cell continuously converts chemical energy to electrical energy as long as fuel (hydrogen) and oxidant (oxygen/air) are supplied. Unlike batteries, fuel cells do not "run out" — they are not discharged. They are used in spacecraft, submarines (AIP — air-independent propulsion), and emerging hydrogen-powered vehicles.
⚡ Hydrogen-Oxygen Fuel Cell
Anode: H₂ → 2H⁺ + 2e⁻ (oxidation; H₂ supplied)
Cathode: ½O₂ + 2H⁺ + 2e⁻ → H₂O (reduction; O₂ supplied)
Overall: H₂ + ½O₂ → H₂O (only by-product: water!)
EMF: ≈ 1.23 V per cell (theoretical)
Electrolyte: KOH solution or polymer membrane (PEM fuel cell)
Efficiency: ≈ 60–80% (much higher than combustion engines: ≈ 25–35%)
Advantages:
✓ Zero pollution (only water produced)
✓ High efficiency — no heat losses of combustion
✓ Continuous operation (fuel supplied externally)
✓ Silent operation — ideal for stealth submarines (AIP)
✓ Scalable — from milliwatt (electronics) to megawatt (power plants)
Defence use: German Type 212 submarines use H₂-O₂ fuel cells for silent
underwater propulsion (AIP system)
Answer: (c) Lead-acid accumulator
Lead-acid, NiCd, and lithium-ion are secondary (rechargeable) cells. Dry cell and voltaic cell are primary — once exhausted, they are discarded. Secondary cells can be restored to original state by passing current in reverse direction.
Q2. Which type of cell produces only water as a by-product and is used in space applications?
Answer: (c) Hydrogen-oxygen fuel cell
The H₂-O₂ fuel cell reaction produces only water. Used in NASA spacecraft (Apollo, Space Shuttle) to provide both electricity and drinking water. Also used in stealth submarines (AIP) where zero pollution and silent operation are critical.
Q3. The difference between a primary cell and a secondary cell is:
(a) Primary cells are larger (b) Primary cells can be recharged; secondary cannot (c) Secondary cells can be recharged; primary cells cannot (d) Primary cells have higher EMF
Answer: (c) Secondary cells can be recharged; primary cells cannot Primary cell: irreversible chemical reaction — once used, discarded (dry cell, voltaic cell). Secondary cell: reversible reaction — can be restored by charging (lead-acid, NiCd, Li-ion). The key difference is reversibility of the electrochemical reaction.
🤔 TRICKY QUESTIONS
Chemical Cells — Exam Surprises
T1. A student connects cells in series. Why does connecting more cells in series increase voltage but not necessarily increase current?
Series cells: voltages add; current is limited by total internal resistance.
In series: V_total = V₁ + V₂ + ... + Vₙ (voltages add → higher EMF). But also: r_total = r₁ + r₂ + ... + rₙ (internal resistances add → higher total resistance). Current I = V_total/(R_ext + r_total). If r_total becomes large, the increase in current is limited. For maximum current (high-drain applications), batteries are connected in parallel — voltage stays the same but total internal resistance decreases (1/r_eff = Σ1/rᵢ) → more current capability.
T2. Why must lead-acid batteries NOT be completely discharged (fully run down)?
Deep discharge causes sulphation — large PbSO₄ crystals permanently damage the plates.
During normal discharge, PbSO₄ crystals form on both electrodes — small and reversible. During deep or prolonged discharge, these crystals grow large and hard (sulphation). Large PbSO₄ crystals are difficult to reconvert during charging — the plate active area is permanently reduced. Result: the battery capacity is permanently diminished after a deep discharge. Military vehicle batteries are carefully monitored to prevent over-discharge — this is why vehicles have battery protection systems that disconnect the load before complete discharge.
⚡ High-Yield Formula Sheet — PN08 Chemical Cells
⚡ Redox Basics
OIL RIG: Oxidation = loss; Reduction = gain (e⁻)
Anode = oxidation; Cathode = reduction (always)
Galvanic cell: anode (−), cathode (+)
EMF = E_cathode − E_anode
🔋 Primary Cells
Voltaic: Zn(−) | H₂SO₄ | Cu(+) — EMF ≈ 1.1 V
Dry cell: Zn(−) | NH₄Cl | C(+) — EMF ≈ 1.5 V
Local action: cure by amalgamation (Hg on Zn)
Polarisation: cure by depolariser (MnO₂)
🔌 Lead-Acid Battery
Pb(−) | dil. H₂SO₄ | PbO₂(+)
EMF ≈ 2 V per cell; 12 V = 6 cells
Discharge: H₂SO₄ consumed; density ↓
Charged density ≈ 1.28 g/mL
Both electrodes → PbSO₄ on discharge
🔋 NiCd & Fuel Cell
NiCd: Cd(−) | KOH | NiO(OH)(+) — EMF ≈ 1.2 V
NiCd: memory effect — fully discharge before recharge
H₂-O₂ fuel cell: EMF ≈ 1.23 V; product = H₂O only
Fuel cell: no discharge — continuous fuel supply
📈 Faraday's Laws
m = ZIt = ZQ (Z = electrochemical equivalent)
m = EIt/F (E = equivalent weight, F = 96,500 C)
F = 96,500 C/mol (charge of 1 mole e⁻)
Second law: m₁/m₂ = E₁/E₂ (same charge)
📌 Quick Comparison
Voltaic: 1.1 V; Dry cell: 1.5 V; NiCd: 1.2 V
Lead-acid: 2.0 V/cell; Li-ion: 3.6 V
Primary: not rechargeable; Secondary: rechargeable
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