Why does Leclanche voltage drop but lead storage holds voltage well?

In Leclanche cells Zn2+ builds up in solution and the Nernst equation predicts steady voltage drop. In lead storage batteries all electrodes and products are solids at unit activity; only H2SO4 concentration changes, giving a flatter discharge.

Why is H2-O2 fuel cell pollution-free?

The only product is water, unlike fossil-fuel combustion which releases CO2, NOx, and SOx. However H2 production is truly clean only if made from renewable sources.

Common Cell

Q: Why is a Leclanche cell called a dry cell if it contains a paste?

Dry refers to the absence of free-flowing liquid electrolyte. The NH4Cl-ZnCl2 paste is moist enough to allow ion flow but not liquid, so the cell works in any orientation and won't spill.

Q: Why is lithium used in modern batteries?

Lithium is the lightest metal, has the most negative E-standard giving highest cell EMF, and its small ion size allows fast intercalation into graphite and layered oxides for rapid charge and discharge.

Q: Why does iron rust faster than aluminium even though Al is more reactive?

Al forms a dense adherent Al2O3 layer that seals the surface (passivation). Iron's oxide (rust) is porous and flaky, letting oxygen and moisture reach fresh metal underneath.

Q: Difference between galvanisation and cathodic protection?

Galvanisation coats iron with a thin Zn layer acting as both a barrier and a sacrificial anode. Cathodic protection connects a large Mg or Zn block by wire to a buried structure so the block corrodes and protects the structure.

Q: Why are fuel cells more efficient than heat engines?

Heat engines are limited by Carnot efficiency (35-40%). Fuel cells convert chemical energy directly to electricity with no thermal step, so Carnot does not apply. The ceiling is delta-G over delta-H, about 83%, and practical cells reach roughly 70%.

Q: Can iron corrode in the complete absence of oxygen?

Yes, but very slowly and by a different mechanism, often via anaerobic sulphate-reducing bacteria producing H2S that reacts with iron. In pure oxygen-free water without such organisms, corrosion is negligible.

ChemistryElectrochemistryFor JEE aspirants

Common cells are practical galvanic cells that power everyday devices, from remote controls to cars to satellites. They fall into three families: primary cells (single-use, non-rechargeable: Leclanche dry cell, alkaline, mercury), secondary cells (rechargeable: lead storage, Ni-Cd, lithium-ion), and fuel cells (continuous supply of reactants: $H_{2}$ - $O_{2}$ fuel cell). A parallel topic is corrosion, the electrochemical destruction of metals like the rusting of iron, and its prevention by galvanisation, cathodic protection, and coatings.

Key Formulas - Quick Reference

Lead storage overall (discharge): $Pb + PbO_{2} + 2 H_{2} SO_{4} \to 2 PbSO_{4} + 2 H_{2} O$
$H_{2}$ - $O_{2}$ fuel cell overall: $2 H_{2} + O_{2} \to 2 H_{2} O$
Fuel cell efficiency ceiling: $η_{m a x} = \frac{Δ G ^{\circ}}{Δ H ^{\circ}} = \frac{237}{286} \approx 83%$
Fuel cell EMF (standard): $E_{cell}^{\circ} = 1.23$ V (at 25°C, from $Δ G^{\circ} = - 237$ kJ/mol water)
Iron corrosion (rusting): $2 Fe + O_{2} + 2 H_{2} O \to 2 Fe(OH)_{2} \to Fe_{2} O_{3} \cdot x H_{2} O$ (rust)
Nernst for corrosion cell: $E = 1.23 - 0.0591 pH$ (oxygen cathode reaction at 1 atm)

1. Classification of Practical Cells

Family	Rechargeable?	Examples
Primary cells	No; cell reaction is irreversible once reactants are consumed	Leclanche dry cell, alkaline, mercury
Secondary cells	Yes; reaction is reversed by passing current in the opposite direction	Lead storage, Ni-Cd, Ni-MH, lithium-ion
Fuel cells	Continuous operation as long as fuel + oxidiser are supplied	$H_{2}$ - $O_{2}$ , methanol, phosphoric acid

2. Leclanche (Dry) Cell

2.1 Construction

Anode: zinc container (also serves as the case).
Cathode: a graphite (carbon) rod surrounded by a paste of $MnO_{2}$ and carbon powder.
Electrolyte: moist paste of $NH_{4} Cl$ and $ZnCl_{2}$ (not liquid, hence "dry").

Figure 1: A Leclanche dry cell in cut-away. Two details matter for exam questions: the graphite rod only collects current (the real oxidant is

MnO_{2}

), and the rod must never reach the zinc, or the cell shorts itself out.

2.2 Cell Reactions

Anode

Zn \to Zn^{2 +} + 2 e^{-}

Cathode

2 MnO_{2} + 2 NH_{4}^{+} + 2 e^{-} \to Mn_{2} O_{3} + 2 NH_{3} + H_{2} O

Overall

Zn + 2 MnO_{2} + 2 NH_{4}^{+} \to Zn^{2 +} + Mn_{2} O_{3} + 2 NH_{3} + H_{2} O

EMF: about 1.5 V. The $NH_{3}$ produced combines with $Zn^{2 +}$ to form $[Zn(NH_{3})_{4}]^{2 +}$ , preventing pressure buildup.

2.3 Limitations

Voltage drops steadily as $Zn^{2 +}$ concentration rises (Nernst effect).
Non-rechargeable: cell reactions are not readily reversible.
Short shelf life: internal reactions continue slowly even when not in use.
Zinc container corrodes and leaks after prolonged use or discharge.

3. Alkaline (Dry) Cell

An upgrade over the Leclanche cell where $NH_{4} Cl$ paste is replaced by KOH.

Anode

Zn (s) + 2 OH^{-} (a q) \to ZnO (s) + H_{2} O (l) + 2 e^{-}

Cathode

2 MnO_{2} (s) + H_{2} O (l) + 2 e^{-} \to Mn_{2} O_{3} (s) + 2 OH^{-} (a q)

Overall

Zn (s) + 2 MnO_{2} (s) \to ZnO (s) + Mn_{2} O_{3} (s)

Advantages over Leclanche: no gas produced, more stable voltage (1.5 V), longer shelf life, better performance at high current drain. Widely used in remote controls, torches, cameras.

4. Mercury (Button) Cell

Compact "button" cell used in watches, hearing aids, calculators. Very steady 1.35 V output over its lifetime.

Anode

Zn(Hg) + 2 OH^{-} \to ZnO + H_{2} O + 2 e^{-}

(zinc amalgam)

Cathode

HgO (s) + H_{2} O (l) + 2 e^{-} \to Hg (l) + 2 OH^{-}

Overall

Zn(Hg) + HgO (s) \to ZnO (s) + Hg (l)

Because the reaction products stay solid or liquid (not gaseous), voltage is exceptionally stable. Environmental concern: mercury toxicity has led to phase-out in favour of silver oxide and lithium button cells.

5. Lead Storage Battery

The workhorse secondary cell used in cars, inverters, and UPS units. Consists of 6 cells (2 V each) in series to give 12 V.

5.1 Construction

Anode: spongy lead ( $Pb$ ) plates.
Cathode: lead grid packed with $PbO_{2}$ .
Electrolyte: 38% $H_{2} SO_{4}$ solution (density about 1.28 g/mL when fully charged).

Figure 2: One cell of a lead storage battery; six of these in series give the familiar 12 V. Because both plates end up as the same solid,

PbSO_{4}

, a fully discharged cell has nothing left to drive a current, which is why deep discharge ruins it.

5.2 Reactions on Discharge

Anode

Pb (s) + SO_{4}^{2 -} (a q) \to PbSO_{4} (s) + 2 e^{-}

Cathode

PbO_{2} (s) + SO_{4}^{2 -} (a q) + 4 H^{+} (a q) + 2 e^{-} \to PbSO_{4} (s) + 2 H_{2} O (l)

Overall

Pb (s) + PbO_{2} (s) + 2 H_{2} SO_{4} (a q) \to 2 PbSO_{4} (s) + 2 H_{2} O (l)

EMF per cell: 2.05 V. As the cell discharges, $H_{2} SO_{4}$ is consumed and water is produced, so the electrolyte density drops. This is how a hydrometer is used to check battery state of charge.

5.3 Recharging

Applying an external voltage greater than 2.05 V per cell reverses all the reactions:

2 PbSO_{4} (s) + 2 H_{2} O (l) charging Pb (s) + PbO_{2} (s) + 2 H_{2} SO_{4} (a q)

Both electrodes regenerate their original forms and the electrolyte density rises again. Practical cycle life is around 1000-2000 cycles depending on depth of discharge.

6. Nickel-Cadmium (Ni-Cd) Cell

Sealed rechargeable cell with EMF around 1.35 V. Robust, tolerates deep discharge, but Cd is toxic (largely superseded by Ni-MH and Li-ion).

Anode

Cd (s) + 2 OH^{-} (a q) \to Cd(OH)_{2} (s) + 2 e^{-}

(discharge)

Cathode

NiO_{2} (s) + 2 H_{2} O (l) + 2 e^{-} \to Ni(OH)_{2} (s) + 2 OH^{-} (a q)

(discharge)

Overall

Cd (s) + NiO_{2} (s) + 2 H_{2} O (l) \to Cd(OH)_{2} (s) + Ni(OH)_{2} (s)

(discharge)

7. Lithium-Ion Battery

The dominant modern rechargeable battery, used in phones, laptops, electric vehicles. High energy density, long cycle life, no memory effect.

7.1 Construction

Anode: graphite intercalated with lithium (Li $_{x}$ C $_{6}$ ).
Cathode: a lithium-metal oxide, typically $LiCoO_{2}$ , $LiMn_{2} O_{4}$ , or $LiFePO_{4}$ .
Electrolyte: lithium salt (e.g., $LiPF_{6}$ ) in an organic solvent (no water; Li reacts violently with water).

7.2 Reactions on Discharge

Anode

Li_{x} C_{6} \to x Li^{+} + x e^{-} + C_{6}

Cathode

Li_{1 - x} CoO_{2} + x Li^{+} + x e^{-} \to LiCoO_{2}

OverallLi

^{+}

ions shuttle from anode (graphite) through the electrolyte to the cathode (

LiCoO_{2}

) during discharge, and back during charging.

EMF: about 3.7 V per cell (much higher than aqueous cells, since no water constraint). Energy density around 250 Wh/kg, several times that of Pb-acid.

8. Fuel Cells

A fuel cell is a galvanic cell where reactants are supplied continuously from outside, and products are removed continuously. It's not a "storage" cell (like a battery) but an energy-conversion device: chemical energy directly to electrical energy without going through heat.

8.1 $H_{2}$ - $O_{2}$ Fuel Cell (used in Apollo missions)

Anode: porous graphite impregnated with Pt catalyst; $H_{2}$ gas is bubbled in.
Cathode: similar electrode with $O_{2}$ gas.
Electrolyte: aqueous KOH (or an ion-exchange membrane in modern PEM fuel cells).

Anode

2 H_{2} (g) + 4 OH^{-} (a q) \to 4 H_{2} O (l) + 4 e^{-}

Cathode

O_{2} (g) + 2 H_{2} O (l) + 4 e^{-} \to 4 OH^{-} (a q)

Overall

2 H_{2} (g) + O_{2} (g) \to 2 H_{2} O (l)

;

E^{\circ} = 1.23

Figure 3: An alkaline

H_{2}

O_{2}

fuel cell. Nothing is stored: reactants flow in and water flows out, so the cell runs as long as it is fed. Because there is no heat step, it is not capped by the Carnot limit.

Advantages of fuel cells:

Efficiency around 70% in practice, against 35 to 40% for a thermal power plant, because there is no heat step and therefore no Carnot limit. The thermodynamic ceiling is $Δ G^{\circ} /Δ H^{\circ} = 237/286 \approx 83%$ .
No pollution: only product is water.
No moving parts, silent, reliable.

Disadvantages: expensive Pt catalyst; difficult storage of

H_{2}

(high-pressure or cryogenic); infrastructure limited.

9. Corrosion

Corrosion is the electrochemical oxidation of a metal in contact with air, water, or a corrosive environment. Iron rusting is the most familiar and costly example.

9.1 Mechanism of Rusting

A tiny water droplet on iron acts as a mini galvanic cell:

Figure 4: Rusting is a corrosion cell, not a simple oxidation. The oxygen gradient across a single droplet is enough to split the surface into anodic and cathodic patches, and the electrons never leave the metal.

Anode

Fe (s) \to Fe^{2 +} (a q) + 2 e^{-}

;

E_{ox}^{\circ} = 0.44

V (centre of drop, oxygen-poor)

Cathode

O_{2} (g) + 4 H^{+} (a q) + 4 e^{-} \to 2 H_{2} O (l)

;

E_{red}^{\circ} = 1.23

V (edge, oxygen-rich)

Overall

2 Fe (s) + O_{2} (g) + 4 H^{+} (a q) \to 2 Fe^{2 +} (a q) + 2 H_{2} O (l)

;

E_{cell}^{\circ} = 1.67

Then

Fe^{2 +}

is oxidised further to

Fe^{3 +}

, which precipitates as hydrated ferric oxide (rust):

Fe_{2} O_{3} \cdot x H_{2} O

9.2 Factors Accelerating Corrosion

Presence of moisture: water provides the electrolyte medium.
Acidity (low pH): $H^{+}$ is a reactant at the cathode; acid rain accelerates rusting.
Sea water (dissolved salts): $NaCl$ increases conductivity and speeds up ion migration. This is why ships and coastal metals corrode faster.
Presence of impurities in the metal: impure iron with carbon or other metals sets up more micro-galvanic cells.
Temperature: higher temperature speeds up all electrochemical reactions.

9.3 Prevention of Corrosion

Method	Principle
Barrier coating (paint, oil, grease, enamel)	Physically blocks moisture and $O_{2}$ from reaching the metal surface.
Galvanisation (coating with Zn)	Zn oxidises preferentially (Zn is more reactive); it's a sacrificial anode. Even if the coating cracks, Zn continues to protect Fe.
Tinning (coating with Sn)	Only works as long as the coating is intact. If scratched, Fe corrodes faster (Sn is less reactive than Fe).
Cathodic protection	Connect the metal to a more reactive metal (Mg, Zn) or a DC source, making the protected metal the cathode. Used for underground pipelines and ship hulls.
Alloying (e.g., stainless steel = Fe + Cr + Ni)	Cr forms a passive oxide layer that prevents further oxidation.
Electroplating with Ni, Cr, Au, Ag	Decorative and protective barrier.

Figure 5: The same scratch, opposite outcomes. A coating protects iron only if it is more reactive than iron; otherwise breaking the coat turns the whole coated sheet into a corrosion cell with the iron as the anode.

10. Solved Examples

Solved Example 1

Which of the following metals can be used for galvanisation of iron: Cu, Al, Zn, Sn?

Solution:

Galvanisation requires a metal more reactive than Fe (more negative reduction potential), so it can act as a sacrificial anode. $E^{\circ}$ values: Zn ( $- 0.76$ ), Al ( $- 1.66$ ), Fe ( $- 0.44$ ), Sn ( $- 0.14$ ), Cu ( $+ 0.34$ ).

Both Zn and Al are more reactive than Fe and can protect it. In practice Zn is used (Al forms a passivating oxide layer too quickly). Cu and Sn are less reactive than Fe and would accelerate corrosion if the coating scratched.

Solved Example 2

Why does sea water promote corrosion of iron?

Solution:

Sea water contains dissolved $NaCl$ (about 3.5%), which greatly increases the electrical conductivity of the water film on iron. Higher conductivity means faster ion transport in the mini galvanic cells on the metal surface, and higher current $=$ faster rusting. Additionally, $Cl^{-}$ ions penetrate protective oxide films and destabilise them.

Solved Example 3

Calculate the standard EMF of the

H_{2}

O_{2}

fuel cell given

Δ G^{\circ} = - 237

kJ/mol of water formed,

n = 2

Solution:

$Δ G^{\circ} = - n F E^{\circ} ⟹ E^{\circ} = - \frac{Δ G ^{\circ}}{n F} = - \frac{- 237000}{2 \times 96500} = 1.228$ V $\approx 1.23$ V.

Solved Example 4

During the discharge of a lead storage battery, what is the change in the concentration of

H_{2} SO_{4}

Solution:

Overall discharge: $Pb + PbO_{2} + 2 H_{2} SO_{4} \to 2 PbSO_{4} + 2 H_{2} O$ . Two moles of $H_{2} SO_{4}$ are consumed and two moles of water are produced per cycle. So $H_{2} SO_{4}$ concentration (and hence electrolyte density) decreases as discharge proceeds. A hydrometer reading below 1.20 g/mL indicates the battery needs recharging.

Solved Example 5

Why is the EMF of a lithium-ion battery (3.7 V) much higher than that of a lead storage cell (2.05 V) or a Ni-Cd cell (1.35 V)?

Solution:

Lithium has the most negative standard reduction potential of any metal ( $E^{\circ} = - 3.03$ V), giving it a very high oxidation potential. Combined with an intercalation cathode chemistry, the effective cell EMF is around 3.7 V. Also, Li-ion batteries use a non-aqueous electrolyte (organic solvent), which avoids the water-electrolysis limit of about 1.23 V that caps aqueous cells at low voltages.

Solved Example 6

Why doesn't a lead storage battery lose voltage significantly with discharge, unlike a Leclanche cell?

Solution:

In a lead storage cell, the reactants (Pb, $PbO_{2}$ ) and products ( $PbSO_{4}$ ) are all solids with unit activity. Only $H_{2} SO_{4}$ concentration changes, and even that changes slowly, so the Nernst correction is small and voltage stays near 2 V until the acid is depleted. In a Leclanche cell, $Zn^{2 +}$ builds up in solution as an aqueous ion, and its concentration keeps changing, causing the voltage to drop steadily.

Common Mistakes to Avoid

Watch out

Confusing primary and secondary cells: primary $=$ single-use (Leclanche, alkaline, mercury); secondary $=$ rechargeable (lead, Ni-Cd, Li-ion).
Thinking a fuel cell is a battery: fuel cells need continuous fuel supply, they are not storage devices.
Forgetting that in the lead storage battery, both electrodes convert to $PbSO_{4}$ during discharge, making a "self-shorting" cell (which is why deep discharge damages it).
Assuming Zn coating on Fe works only as a physical barrier: it works even better as a sacrificial anode when the coating is scratched. Sn coating does not (Sn is less reactive than Fe).
Confusing galvanisation (Zn on Fe) with tinning (Sn on Fe): they behave oppositely when scratched.
Ignoring the role of $O_{2}$ in rusting: without $O_{2}$ , iron does not rust even in water. Both water and oxygen are needed.
Thinking sea water corrodes because of $Cl^{-}$ chemistry alone: the main effect is enhanced conductivity from dissolved ions.

Frequently Asked Questions

Q1. Why is a Leclanche cell called a "dry" cell if it contains a paste?

"Dry" refers to the absence of free-flowing liquid electrolyte. The $NH_{4} Cl$ - $ZnCl_{2}$ paste is moist enough to allow ion flow but not liquid, so the cell can be used in any orientation and won't spill. This makes it portable, unlike older wet cells that had to stay upright.

Q2. Why does the voltage of a Leclanche cell drop with use but a lead storage battery holds voltage well?

In a Leclanche cell, $Zn^{2 +}$ builds up in solution as its concentration rises; the Nernst equation gives steady voltage drop. In a lead storage battery, both electrodes and products are solids with unit activity; only $H_{2} SO_{4}$ concentration changes, giving a much flatter discharge curve.

Q3. Why is lithium used in modern batteries?

Three reasons: (a) lightest metal (density 0.53 g/cm $^{3}$ ), giving high energy per unit mass; (b) most negative $E^{\circ}$ ( $- 3.03$ V), giving highest cell EMF; (c) small ion size allows fast intercalation into graphite and layered oxides, enabling rapid charge/discharge without structural damage.

Q4. Why does iron rust faster than aluminium even though Al is more reactive?

Al forms a thin, dense, adherent $Al_{2} O_{3}$ layer that seals the surface and prevents further reaction (called passivation). Iron's oxide (rust) is porous, flaky, and non-adherent, so it lets $O_{2}$ and moisture keep reaching fresh metal underneath. This is why Al windows last decades while iron ones rust quickly.

Q5. What is the difference between galvanisation and cathodic protection?

Both use a sacrificial metal (usually Zn or Mg) more reactive than iron. Galvanisation coats the iron with a thin Zn layer that acts as both a physical barrier and a sacrificial anode. Cathodic protection connects a block of Mg or Zn (buried nearby) to the iron by a wire, so the block corrodes and protects a large structure (pipeline, ship hull) without direct coating.

Q6. Why are fuel cells more efficient than heat engines?

Heat engines are limited by the Carnot efficiency, $η = 1 - T_{c} / T_{h}$ , typically 35-40% for modern power plants. Fuel cells convert chemical energy directly to electricity with no thermal step, so Carnot does not apply. Their ceiling is instead $Δ G^{\circ} /Δ H^{\circ} = 237/286 \approx 83%$ , and practical cells reach about 70%.

Q7. Why is $H_{2}$ - $O_{2}$ fuel cell called "pollution-free"?

The only product is water, unlike fossil-fuel combustion which releases $CO_{2}$ , $NO_{x}$ , and $SO_{x}$ . However, hydrogen production is only truly clean if the $H_{2}$ itself is made from renewable energy (electrolysis of water using solar/wind). If $H_{2}$ comes from methane reforming, there's still a $CO_{2}$ footprint upstream.

Q8. Can iron corrode in the complete absence of oxygen?

Yes, but very slowly, and by a different mechanism. In oxygen-free water, iron can be attacked by anaerobic sulphate-reducing bacteria that produce $H_{2} S$ , which reacts with Fe. In pure oxygen-free water without such organisms, corrosion is negligible. So modern boiler design deaerates water to protect the metal.

Q9. What are the environmental problems with Ni-Cd and mercury batteries?

Both Cd and Hg are toxic heavy metals that accumulate in the food chain. Cd causes kidney damage and cancer; Hg causes neurological damage (Minamata disease). Most countries have banned or heavily restricted mercury cells, and Ni-Cd cells are being replaced by Ni-MH (nickel-metal hydride, non-toxic) and Li-ion cells. Recycling programs are mandatory in many regions.

Previous year questions on Common Cell

2 questions from past papers, each with a step-by-step solution.

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1. Classification of Practical Cells

2. Leclanche (Dry) Cell

2.1 Construction

2.2 Cell Reactions

2.3 Limitations

3. Alkaline (Dry) Cell

4. Mercury (Button) Cell

5. Lead Storage Battery

5.1 Construction

5.2 Reactions on Discharge

5.3 Recharging

6. Nickel-Cadmium (Ni-Cd) Cell

7. Lithium-Ion Battery

7.1 Construction

7.2 Reactions on Discharge

8. Fuel Cells

8.1 H2​-O2​ Fuel Cell (used in Apollo missions)

9. Corrosion

9.1 Mechanism of Rusting

9.2 Factors Accelerating Corrosion

9.3 Prevention of Corrosion

10. Solved Examples

Common Mistakes to Avoid

Frequently Asked Questions

Q1. Why is a Leclanche cell called a "dry" cell if it contains a paste?

Q2. Why does the voltage of a Leclanche cell drop with use but a lead storage battery holds voltage well?

Q3. Why is lithium used in modern batteries?

Q4. Why does iron rust faster than aluminium even though Al is more reactive?

Q5. What is the difference between galvanisation and cathodic protection?

Q6. Why are fuel cells more efficient than heat engines?

Q7. Why is H2​-O2​ fuel cell called "pollution-free"?

Q8. Can iron corrode in the complete absence of oxygen?

Q9. What are the environmental problems with Ni-Cd and mercury batteries?

Previous year questions on Common Cell

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8.1 $H_{2}$ - $O_{2}$ Fuel Cell (used in Apollo missions)

Q7. Why is $H_{2}$ - $O_{2}$ fuel cell called "pollution-free"?