Case study: Batteries and fuel cells are very useful forms of galvanic cell. Any battery or cell that we use a

Question

Q. Case study: Batteries and fuel cells are very useful forms of galvanic cell. Any battery or cell that we use as a source of electrical energy is basically a galvanic cell. For a battery to be of practical use it should be reasonably light, compact and its voltage should not vary appreciably during its use. There are mainly two types of batteries — primary batteries and secondary batteries. In primary batteries the reaction occurs only once and after use the battery becomes dead and cannot be reused, whereas secondary batteries are rechargeable. Production of electricity by thermal plants is inefficient and a major source of pollution; to solve this, galvanic cells are designed so that energy of combustion of fuels is directly converted into electrical energy — these are known as fuel cells. One such fuel cell was used in the Apollo space programme. Answer the following questions:

(a) How do primary batteries differ from secondary batteries?

(1)

(b) The cell potential of Mercury cell is 1.35 V, and remains constant during its life. Give reason.

(1)

(OR)

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Part (a)Concept understanding — Lead Acid Battery Recharging

Lead Acid Battery Recharging – From Intuition to Precision

Think of a lead acid battery like a rechargeable water tank. When you use the battery (discharge), you're letting water flow out to do work. Recharging is the process of pumping that water back in, restoring the tank to its full capacity. But instead of water, we're moving charged particles (ions) and electrons.

The Intuition: Reversing the Chemical Spill

During discharge, the battery's internal chemicals react to produce electricity. Lead dioxide ( $PbO_2$ ) and sponge lead ( $Pb$ ) react with sulfuric acid ( $H_2SO_4$ ) to form lead sulfate ( $PbSO_4$ ) and water. This reaction releases electrons that flow through your circuit.

Recharging is the exact opposite. You force electrons back into the battery by applying an external voltage (from a charger). This reverses the chemical reaction, converting lead sulfate back into lead dioxide and sponge lead, and regenerating sulfuric acid. The battery is "refilled" chemically.

The Precise Statement

Recharging a lead acid battery is the process of applying an external electrical potential (greater than the battery's open-circuit voltage) to force a non-spontaneous reverse of the discharge reaction, restoring the active materials and electrolyte concentration.

The key chemical equation for the overall cell reaction during discharge is:

Pb + PbO_2 + 2H_2SO_4 \rightarrow 2PbSO_4 + 2H_2O

During recharging, the external voltage drives the reverse:

2PbSO_4 + 2H_2O \rightarrow Pb + PbO_2 + 2H_2SO_4

What Happens Step-by-Step

Charger applies voltage – The charger must supply a voltage slightly higher than the battery's resting voltage (typically 13.8–14.4 V for a 12 V battery). This overcomes the battery's internal resistance and the chemical back-EMF.
Electrons flow backwards – Electrons are forced into the negative terminal and pulled out of the positive terminal. This reverses the direction of current compared to discharge.
Lead sulfate dissolves – At the negative plate, $PbSO_4$ gains electrons and converts back to sponge lead ( $Pb$ ). At the positive plate, $PbSO_4$ loses electrons and reforms $PbO_2$ .
Sulfuric acid regenerates – The sulfate ions ( $SO_4^{2-}$ ) released from both plates recombine with hydrogen ions ( $H^+$ ) from water to form $H_2SO_4$ . The electrolyte's specific gravity rises back to its charged state.
Gassing (if overcharged) – Once all lead sulfate is converted, further charging splits water into hydrogen and oxygen gas. This is why you see bubbling in a flooded battery – it's a sign of full charge (or overcharging).

Watch out

Never recharge a lead acid battery with a voltage too high. Excessive voltage causes rapid gassing, water loss, and can damage the plates. The correct charging voltage depends on temperature and battery type (flooded vs. sealed).

Why It's Not Perfect …

Why this formula?

Lead Acid Battery Recharging: Why the Key Formulas Hold

Let's start with the core chemical reaction during recharging — understanding why the formulas work requires knowing what happens inside the battery.

1. The Recharging Reaction (The "Why" Behind the Voltage)

During discharge, lead sulfate ( $PbSO_4$ ) forms on both electrodes. Recharging reverses this:

At the negative terminal (cathode during recharge):

PbSO_4(s) + 2e^- \rightarrow Pb(s) + SO_4^{2-}(aq)

At the positive terminal (anode during recharge):

PbSO_4(s) + 2H_2O(l) \rightarrow PbO_2(s) + SO_4^{2-}(aq) + 4H^+(aq) + 2e^-

Overall recharge reaction:

2PbSO_4(s) + 2H_2O(l) \rightarrow Pb(s) + PbO_2(s) + 2H_2SO_4(aq)

Key insight: Recharging regenerates sulfuric acid ( $H_2SO_4$ ) — this is why the electrolyte's specific gravity increases during charging.

2. The Voltage Formula: $V_{\text{cell}} = 2.1 + \frac{0.059}{2} \log \frac{[H^+]^4}{[SO_4^{2-}]^2}$

Why this formula holds — step by step:

Step 1: Nernst equation for each half-cell

For the negative half-cell (Pb/PbSO₄):

E_- = E_-^\circ - \frac{0.059}{2} \log \frac{1}{[SO_4^{2-}]}

For the positive half-cell (PbO₂/PbSO₄):

E_+ = E_+^\circ - \frac{0.059}{2} \log \frac{[H^+]^4}{[SO_4^{2-}]}

Step 2: Cell voltage = $E_+ - E_-$

V = (E_+^\circ - E_-^\circ) - \frac{0.059}{2} \log \frac{[H^+]^4}{[SO_4^{2-}]} + \frac{0.059}{2} \log \frac{1}{[SO_4^{2-}]}

Step 3: Simplify using standard potentials

$E_+^\circ - E_-^\circ = 2.1\ \text{V}$ (standard cell voltage). Combining logs:

V = 2.1 + \frac{0.059}{2} \log \frac{[H^+]^4}{[SO_4^{2-}]^2}

Why this matters for recharging:

As $H_2SO_4$ concentration increases during charging, $[H^+]$ rises and $[SO_4^{2-}]$ rises → the log term becomes more positive → voltage rises.
This is why a fully charged battery shows ~2.6–2.7 V per cell during charging (higher than the 2.1 V open-circuit voltage).

3. The Charging Current Formula: $I_{\text{charge}} = \frac{C}{10}$ (for constant current)

Why "C/10" is the standard rule:

Reasoning:

C = battery capacity in Ah (e.g., 100 Ah)
C/10 = 10 A for a 100 Ah battery

Why this rate?

At higher rates (e.g., C/5), the internal resistance causes excessive heat and gassing (water electrolysis into $H_2$ and $O_2$ ).
At lower rates (e.g., C/20), charging takes too long and may not fully reverse sulfation.
C/10 balances:
- Reaction kinetics: The reverse reaction rate matches the diffusion of ions.
- Thermal management: Heat generated = $I^2R$ — kept low enough to avoid damage.
- Gas evolution: At C/10, most current goes to the desired reaction, not water splitting.

Exam tip: For lead-acid, the charging voltage is typically 2.4–2.45 V per cell (not 2.1 V) — this overpotential is needed to overcome the activation energy of the reverse reaction.

--- …

Part (b)Concept understanding — Fuel Cell Fuels

Fuel Cell Fuels — From Intuition to Precision

Imagine you have a battery that never runs out as long as you keep feeding it a specific fuel. That's the core idea of a fuel cell. But unlike a car engine that burns petrol in a violent explosion, a fuel cell quietly combines a fuel with oxygen to produce electricity, water, and heat — no flames, no moving parts.

The question is: what can you feed it? That's where fuel cell fuels come in.

The Intuition: What Makes a Good Fuel?

A fuel cell works by stripping electrons from a fuel molecule at one electrode (the anode) and passing those electrons through an external circuit to do work. At the other electrode (the cathode), oxygen grabs those electrons and combines with the fuel's leftover ions.

For this to work efficiently, the fuel must:

Be easy to ionise (give up electrons readily)
React quickly at the anode (good kinetics)
Not poison the catalyst (usually platinum or nickel)
Be storable and transportable in practice

The simplest fuel that satisfies all these? Hydrogen gas ( $H_2$ ). It's the gold standard.

The Precise Statement

Fuel cell fuels are substances that can be oxidised at the anode of a fuel cell to release electrons, producing an electric current. The most common fuel is hydrogen, but other fuels like methanol, ethanol, natural gas (methane), and even ammonia or formic acid are used depending on the application.

The general anode reaction for any fuel is:

\text{Fuel} \rightarrow \text{oxidised products} + \text{electrons}

And the overall cell reaction combines this with oxygen reduction at the cathode:

\text{Fuel} + O_2 \rightarrow \text{oxidation products} + \text{water} + \text{electricity}

The Main Fuel Types (Exam-Ready)

Fuel	Anode Reaction (simplified)	Key Advantage	Key Challenge
Hydrogen ( $H_2$ )	$H_2 \rightarrow 2H^+ + 2e^-$	Highest energy density by mass; zero carbon emissions	Storage (low density); production cost
Methanol ( $CH_3OH$ )	$CH_3OH + H_2O \rightarrow CO_2 + 6H^+ + 6e^-$	Liquid at room temp; easy to transport	Lower efficiency; $CO_2$ produced; catalyst poisoning
Ethanol ( $C_2H_5OH$ )	$C_2H_5OH + 3H_2O \rightarrow 2CO_2 + 12H^+ + 12e^-$	Renewable (bioethanol); less toxic than methanol	Slower reaction kinetics; C–C bond breaking is hard
Methane ( $CH_4$ )	$CH_4 + 2H_2O \rightarrow CO_2 + 8H^+ + 8e^-$	Abundant (natural gas); existing infrastructure	Requires high temperature (solid oxide fuel cells)
Ammonia ( $NH_3$ )	$2NH_3 \rightarrow N_2 + 6H^+ + 6e^-$	Carbon-free; easier to store than $H_2$	Toxic; requires cracking to $H_2$ first

Watch out

A common mistake: thinking any fuel can be used in any fuel cell. Each fuel cell type is designed for a specific fuel. For example, a Proton Exchange Membrane Fuel Cell (PEMFC) runs on pure hydrogen — putting methanol in it would destroy the membrane. Direct methanol fuel cells (DMFCs) are built differently.

Why Not Just Burn the Fuel?

This is the key conceptual leap. In a fuel cell, the fuel's chemical energy is converted directly to electricity, bypassing the heat → mechanical work → generator steps of a thermal power plant. This means:

Higher efficiency (40–60% vs 30–35% for combustion engines)
No moving parts (silent, low maintenance)
No pollutants (if using hydrogen, the only byproduct is water)

Note

For hydrogen fuel cells: $2H_2 + O_2 \rightarrow 2H_2O$ — the only exhaust is pure water. This is why hydrogen fuel cells are considered zero-emission.

The Real-World Catch …

Why this formula?

Fuel Cell Fuels: Why the Key Formulas Hold

Fuel cells convert chemical energy directly into electrical energy. The core idea is electrochemical combustion — instead of burning a fuel to make heat, then turning a turbine, you let the fuel react electrochemically with oxygen to produce electricity directly.

The Fundamental Reaction

For a hydrogen–oxygen fuel cell (the simplest and most common):

2\text{H}_2 + \text{O}_2 \rightarrow 2\text{H}_2\text{O}

But this happens in two half-reactions:

Anode (oxidation):

2\text{H}_2 \rightarrow 4\text{H}^+ + 4e^-

Cathode (reduction):

\text{O}_2 + 4\text{H}^+ + 4e^- \rightarrow 2\text{H}_2\text{O}

The electrons flow through an external circuit — that’s your electricity.

Key Formula 1: Cell Potential from Gibbs Free Energy

The maximum electrical work a fuel cell can do equals the change in Gibbs free energy ( $\Delta G$ ) of the overall reaction:

W_{\text{max}} = -\Delta G

For an electrochemical cell, electrical work is:

W_{\text{electrical}} = n F E

where:

$n$ = number of electrons transferred per mole of fuel (here $n=4$ for $\text{H}_2$ )
$F$ = Faraday constant ( $96\,485\ \text{C/mol}$ )
$E$ = cell potential (volts)

Why this holds:

The Gibbs free energy is the “useful” energy available at constant temperature and pressure. In a fuel cell, that energy is converted into the work of moving electrons through a potential difference. Equating the two gives:

\Delta G = -n F E

So the reversible cell voltage is:

E^\circ = -\frac{\Delta G^\circ}{nF}

For the $\text{H}_2/\text{O}_2$ fuel cell at standard conditions:

$\Delta G^\circ = -237\ \text{kJ/mol}$ (for liquid water product)
$n = 4$
$F = 96\,485\ \text{C/mol}$

E^\circ = -\frac{(-237\,000)}{4 \times 96\,485} \approx 1.23\ \text{V}

This is the theoretical maximum voltage — no fuel cell can exceed this because it’s set by thermodynamics.

Key Formula 2: Efficiency Limit

The thermodynamic efficiency of a fuel cell is:

\eta = \frac{\Delta G}{\Delta H}

where $\Delta H$ is the enthalpy change (the total chemical energy released).

Why this holds:

$\Delta H$ is the total energy released when the fuel burns (the “heat of combustion”).
$\Delta G$ is the portion of that energy that can be converted to electrical work.
The rest ( $T\Delta S$ ) is lost as heat due to entropy change.

For hydrogen:

$\Delta H^\circ = -286\ \text{kJ/mol}$ (higher heating value)
$\Delta G^\circ = -237\ \text{kJ/mol}$

\eta_{\text{max}} = \frac{237}{286} \approx 0.83\ (83\%)

Why this is higher than a heat engine:

A Carnot engine is limited by temperature difference. A fuel cell is not a heat engine — it converts chemical energy directly, so its efficiency limit is set by $\Delta G/\Delta H$ , not by Carnot. This is why fuel cells can theoretically exceed 50% efficiency while internal combustion engines are stuck below ~40%.

Key Formula 3: Nernst Equation for Real Conditions

In practice, the cell voltage depends on concentrations (or partial pressures) of reactants and products:

E = E^\circ - \frac{RT}{nF} \ln Q

where $Q$ is the reaction quotient.

For the hydrogen fuel cell:

Q = \frac{P_{\text{H}_2\text{O}}}{P_{\text{H}_2}^2 \cdot P_{\text{O}_2}}

So: …

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Part (a): primary batteries are single-use (irreversible reaction), secondary ones rechargeable (reversible); the mercury cell holds 1.35 V because its solid/liquid reaction and paste electrolyte keep ion concentrations fixed; recharging the lead cell reverses discharge to regenerate Pb, PbO₂ and H₂SO₄. Part (b): fuel cells are more efficient (direct energy conversion) and cleaner/continuous (only water produced).

Part (a)

Primary vs secondary batteries. A primary battery runs on an irreversible reaction; when the reactants are used up it is dead and cannot be recharged (dry Leclanché cell, mercury cell). A secondary battery uses a reversible reaction — passing current backwards regenerates the reactants, so it is rechargeable and reusable (lead storage cell, Ni–Cd cell).
Constant 1.35 V of the mercury cell. The overall cell reaction is $\ce{Zn(s) + HgO(s) -> ZnO(s) + Hg(l)}$ Every species is a pure solid or liquid and the electrolyte is a moist paste of KOH/ZnO, so the ionic concentrations (and hence the reaction quotient $Q$ in the Nernst equation) do not change as the cell discharges. With $Q$ constant, $E_{cell}$ stays constant at 1.35 V throughout its life.
Recharging the lead storage battery. Recharging drives the discharge reactions in reverse, converting $\ce{PbSO4}$ back to Pb and PbO₂: $\text{Cathode: } \ce{PbSO4(s) + 2e^- -> Pb(s) + SO4^2-(aq)}$ $\text{Anode: } \ce{PbSO4(s) + 2H2O(l) -> PbO2(s) + SO4^2-(aq) + 4H+(aq) + 2e^-}$ …

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Same / Similar Concept — real previous-year questions on the same or a closely similar concept, not this exact question.

Showing the 12 most recent of 16 on this concept.

CBSE 2026Set V11 markMCQ
$Q. Among the following cells, the cell used in the apollo space program for providing electric power is$
(a) SHE
(b) $\text{H}_2\text{-O}_2$ fuel cell
(c) Daniel cell
(d) Mercury cell
›Reveal solutionSolution
The Apollo space programme was powered by the H $_2$ –O $_2$ fuel cell.

A fuel cell converts the energy of combustion of a fuel directly into electrical energy. In the hydrogen–oxygen fuel cell, H $_2$ and O $_2$ are bubbled through porous carbon electrodes in concentrated aqueous NaOH/KOH:

Anode: $\;2\text{H}_2(g) + 4\text{OH}^-(aq) \rightarrow 4\text{H}_2\text{O}(l) + 4e^-$

Cathode: $\;\text{O}_2(g) + 2\text{H}_2\text{O}(l) + 4e^- \rightarrow 4\text{OH}^-(aq)$

Overall: $\;2\text{H}_2(g) + \text{O}_2(g) \rightarrow 2\text{H}_2\text{O}(l)$ …
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CBSE 2026Set ANNUAL1 markMCQ
$Q. The cell which is used in automobiles and inverters are$
(a) Lead storage battery
(b) Nickel cadmium battery
(c) Dry cell
(d) Fuel cell
›Reveal solutionSolution
The lead storage (lead-acid) battery is a rechargeable secondary cell widely used in vehicles and inverters because it can deliver high current and be recharged repeatedly.

The lead storage battery consists of a lead anode and a grid of lead packed with lead dioxide (PbO2) as cathode, with dilute H2SO4 as electrolyte. During discharge, both electrodes get converted to PbSO4, and during charging (by an external current, as in a running car engine or an inverter's mains supply) the reaction is reversed.

Nickel-cadmium battery: used in small rechargeable electronic devices, more expensive, longer-lasting but less common in vehicles. …
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CBSE 2026Set ANNUAL1 mark
$Q. Fill in the blank: Fuel cells produce electricity with an efficiency of about ______ percentage compared to thermal plants whose efficiency is about 40%.$
›Reveal solutionSolution

Fuel cells work at about 70% efficiency, versus about 40% for thermal plants.

A fuel cell (e.g. the H2-O2 fuel cell) converts the chemical energy of a fuel directly into electrical energy through controlled electrochemical reactions, avoiding the wasteful heat-to-work steps of a thermal plant. Its efficiency is therefore high, a …

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CBSE 2025Set 56/6/11 markMCQ
$Q. Which of the following cell converts the energy of combustion of fuel into electrical energy ? (A) Mercury cell (B) Fuel cell (C) Dry cell (D) Lead storage cell$
›Reveal solutionSolution
A fuel cell directly converts the chemical energy from the combustion of a fuel (like hydrogen) into electrical energy, without burning the fuel in a flame. The correct answer is (B).

The key idea here is the direct conversion of chemical energy to electrical energy. In a fuel cell, the fuel (e.g., hydrogen) and an oxidant (e.g., oxygen) are supplied continuously. The reaction that happens inside the cell is essentially the same as combustion — hydrogen combining with oxygen to form water — but it occurs electrochemically, not as a flame. This means electrons are forced to travel through an external circuit, producing electricity.

Let’s look at each option to see why only one fits.

Fuel cell (Option B) — This is the only cell that continuously consumes a fuel and an oxidant from outside. The “combustion of fuel” happens at the electrodes: at the anode, fuel is oxidised (loses electrons), and at the cathode, oxygen is reduced (gains electrons). The overall reaction is the same as burning the fuel, but the energy is released as electrical work, not heat. This is exactly what the question describes.

Mercury cell (Option A) — This is a primary (non-rechargeable) battery. It uses a zinc anode and a mercuric oxide cathode. The reaction is not a combustion; it’s a redox reaction between solid chemicals inside the sealed cell. No fuel is burned, and no external fuel is supplied.

Dry cell (Option C) — The common Leclanché cell (zinc-carbon) or alkaline cell. Again, it’s a sealed primary battery. The reactants are inside the cell (zinc, manganese dioxide, etc.). There is no combustion of a fuel — the reaction is a one-time conversion of stored chemicals. …
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CBSE 2025Set ANNUAL1 mark
$Q. Define the following — Fuel Cells$
›Reveal solutionSolution

A fuel cell is a galvanic cell fed continuously with fuel and oxidant to give a steady electric current.

A fuel cell is a galvanic (voltaic) cell in which the chemical energy of combustion of a fuel (such as $H_2$ , $CH_4$ , $CH_3OH$ ) is converted directly into electrical energy. Fuel (at the anode) and oxygen/air (at the cathode) are fed continuously, and the products are removed continuously, so the cell keeps working as long as the reactants are supplied. Example: the $H_2$ – $O_2$ fuel cell, where

Anode: $2H_2 + 4OH^- \to 4H_2O + 4e^-$

Cathode: $O_2 + 2H_2O + 4e^- \to 4OH^-$

…

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CBSE 2025Set ANNUAL1 mark
$Q. Suggest one material other than hydrogen that can be used as fuel in fuel cell.$
›Reveal solutionSolution

Fuel cells can run on other combustible fuels besides hydrogen, with methanol being a commonly cited practical alternative.

A fuel cell converts the chemical energy of a continuously supplied fuel directly into electrical energy through controlled electrochemical oxidation at the anode (paired with reduction of an oxidant, usually oxygen, at the cathode). While hydrogen is the most common fuel used (as in the $H_2$ – $O_2$ fuel cell), other fuels can also be oxidised electrochemically at a fuel cell's anode, such as:

…

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CBSE 2024Set A11 markMCQ
$Q. During discharging of lead storage battery the correct half-cell reaction is ;$
(a) At anode, Pb is converted into $\text{PbO}_2$
(b) At anode, Pb is converted into $\text{PbSO}_4$
(c) At anode, $\text{PbO}_2$ is converted into $\text{PbSO}_4$
(d) At cathode, Pb is converted into $\text{PbSO}_4$
›Reveal solutionSolution

On discharge, at the anode Pb is oxidised to $\text{PbSO}_4$ — option (b).

In a lead storage (lead–acid) battery on discharge:

Anode (oxidation): $\text{Pb(s)} + \text{SO}_4^{2-}\text{(aq)} \rightarrow \text{PbSO}_4\text{(s)} + 2e^-$

Cathode (reduction): $\text{PbO}_2\text{(s)} + 4\text{H}^+ + \text{SO}_4^{2-} + 2e^- \rightarrow \text{PbSO}_4\text{(s)} + 2\text{H}_2\text{O}$ …
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CBSE 2024Set ANNUAL1 markMCQ
$Q. A device that converts energy of combustion of fuels like hydrogen and methane directly into electrical energy is known as$
(a) fuel cell
(b) electrolytic cell
(c) dynamo
(d) Nickel-Cadmium cell
›Reveal solutionSolution

A fuel cell oxidises a fuel like $H_2$ or $CH_4$ at one electrode and reduces $O_2$ at the other, converting combustion chemical energy directly into electrical energy.

In a fuel cell (e.g. the $H_2$ – $O_2$ fuel cell used in spacecraft), hydrogen is oxidised at the anode and oxygen is reduced at the cathode, with an electrolyte (often aqueous KOH) between them:

Anode: $2H_2(g) + 4OH^-(aq) \rightarrow 4H_2O(l) + 4e^-$

Cathode: $O_2(g) + 2H_2O(l) + 4e^- \rightarrow 4OH^-(aq)$

Overall: $2H_2(g) + O_2(g) \rightarrow 2H_2O(l)$

This is exactly the combustion reaction of hydrogen, but instead of releasing the energy as heat, the electron transfer is forced through an external circuit, generating electricity directly and with high efficiency (no Carnot-limited heat-engine step) and no pollution (only water is produced).

…

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CBSE 2023Set ANNUAL1 markMCQ
$Q. The reaction taking place at cathode of fuel cells is -$
(a) $O_2(g) + 4H^+(aq) + 4e^- \rightarrow 2H_2O(g)$
(b) $2H_2(g) + 4OH^-(aq) \rightarrow 4H_2O(l) + 4e^-$
(c) $H_2(g) \rightarrow 2H^+(aq) + 2e^-$
(d) $2H_2(g) + O_2(g) \rightarrow 2H_2O(l)$
›Reveal solutionSolution

The cathode is where reduction happens; in a hydrogen–oxygen fuel cell, oxygen is reduced there, matching option (a).

In a fuel cell, the cathode is always the electrode where reduction occurs (the electrode that accepts electrons from the external circuit). In an $H_2$ – $O_2$ fuel cell operating in an acidic medium, oxygen gas is reduced at the cathode:

$O_2(g) + 4H^+(aq) + 4e^- \rightarrow 2H_2O(g) \quad \text{(cathode, reduction)}$

Checking the other options: …

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CBSE 2022Set E1 markMCQ
$Q. What happens when a lead storage battery is charged ?$
(a) Lead dioxide dissolves
(b) Sulphuric acid is regenerated
(c) Lead electrode becomes coated with lead sulphate
(d) The concentration of sulphuric acid decreases
›Reveal solutionSolution

Charging a lead storage battery reverses the discharge reaction, regenerating Pb, PbO2 and sulphuric acid.

During discharge both electrodes turn into PbSO4 and H2SO4 is consumed:

Pb + PbO2 + 2H2SO4 -> 2PbSO4 + 2H2O.

Charging drives this reaction backwards (electrolysis):

2PbSO4 + 2H2O -> Pb + PbO2 + 2H2SO4.

So on charging: PbSO4 is converted back to Pb (cathode) and PbO2 (anode), and sulphuric acid is regenerated (its concentration and density rise again).

…

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CBSE 2022Set E1 markMCQ
$Q. Hydrogen-oxygen cell is which of the following types of cell ?$
(a) Primary cell
(b) Secondary cell
(c) Fuel cell
(d) Lead storage cell
›Reveal solutionSolution
The hydrogen-oxygen cell is a fuel cell: it produces electricity from a continuously supplied fuel (H2) and oxidant (O2).

A fuel cell is a galvanic cell in which the reactants (a fuel and an oxidiser) are fed continuously to the electrodes, converting the energy of combustion directly into electrical energy.

Electrode reactions (in KOH):

Anode: H2 + 2OH- -> 2H2O + 2e-

Cathode: O2 + 2H2O + 4e- -> 4OH- …
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CBSE 2022Set ANNUAL1 mark
$Q. The galvanic cells which are used to convert the energy of combustion of fuels like hydrogen, methane etc into electrical energy are generally called as ______.$
›Reveal solutionSolution

Galvanic cells that convert the energy of combustion of fuels like H2, CH4 etc directly into electrical energy are called fuel cells.

In a fuel cell, a fuel (e.g. H2) is oxidised continuously at one electrode and an oxidant (usually O2) is reduced at the other, with the reactants fed in continuously from outside (unlike a normal galvanic cell where reactants are built into the electrodes). This lets the chemical energy of combustion be converted directly and efficiently into electrical energy, without the losses of first burning the fuel and then running a heat engine.

…

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Part (a)

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