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Exhaustive Guide: Batteries & Corrosion | Class 12 Chemistry

Exhaustive Guide: Batteries & Corrosion | Class 12 Chemistry | ChemCA

Batteries, Fuel Cells & Corrosion

Module 5 | CBSE Class 12 Chemistry | Electrochemistry Chapter

1. Introduction to Batteries

Any battery (actually it may have one or more than one cell connected in series) or cell that we use as a source of electrical energy is basically a Galvanic cell where the chemical energy of a spontaneous redox reaction is converted into electrical energy.

For a battery to be of practical use, it should be reasonably light, compact, and its voltage should not vary appreciably during its use. Batteries are broadly classified into two types: Primary and Secondary.

2. Primary Batteries

Primary Batteries: In primary batteries, the reaction occurs only once. After use over a period of time, the battery becomes dead and cannot be reused again (non-rechargeable).

2.1 Dry Cell (Leclanche Cell)

The most familiar example of a primary cell is the dry cell, used commonly in our transistors and clocks. It was invented by Georges Leclanche.

  • Anode: A zinc cylinder (which also acts as the container).
  • Cathode: A carbon (graphite) rod surrounded by powdered manganese dioxide (MnO2) and carbon.
  • Electrolyte: A moist paste of ammonium chloride (NH4Cl) and zinc chloride (ZnCl2) placed between the electrodes.

Cell Reactions:

Anode: Zn(s) → Zn2+ + 2e-

Cathode: MnO2 + NH4+ + e- → MnO(OH) + NH3
Important NCERT Detail: In the cathode reaction, manganese is reduced from the +4 oxidation state to the +3 state. The ammonia (NH3) produced in the reaction is not released as gas; instead, it combines with Zn2+ to form a complex ion [Zn(NH3)4]2+, maintaining the cell's integrity. The cell has a potential of nearly 1.5 V.

2.2 Mercury Cell

Suitable for low current devices like hearing aids, watches, etc.

  • Anode: Zinc-mercury amalgam.
  • Cathode: A paste of HgO and carbon.
  • Electrolyte: A paste of KOH and ZnO.

Cell Reactions:

Anode: Zn(Hg) + 2OH- → ZnO(s) + H2O + 2e-

Cathode: HgO + H2O + 2e- → Hg(l) + 2OH-

Overall Reaction: Zn(Hg) + HgO(s) → ZnO(s) + Hg(l)
Why does a mercury cell provide a constant voltage (1.35 V)?
As seen in the overall reaction, there are no ions in solution whose concentration can change during its life time. Thus, the cell potential remains constant throughout its useful life.

3. Secondary Batteries

Secondary Batteries: A secondary cell after use can be recharged by passing current through it in the opposite direction so that it can be used again. A good secondary cell can undergo a large number of discharging and charging cycles.

3.1 Lead Storage Battery

This is the most important secondary cell, commonly used in automobiles and inverters.

  • Anode: Spongy lead (Pb).
  • Cathode: A grid of lead packed with lead dioxide (PbO2).
  • Electrolyte: A 38% solution of sulphuric acid (H2SO4) by mass (Density = 1.30 g/mL).

Discharging Reactions (When battery is in use):

Anode: Pb(s) + SO42-(aq) → PbSO4(s) + 2e-

Cathode: PbO2(s) + SO42-(aq) + 4H+(aq) + 2e- → PbSO4(s) + 2H2O(l)

Overall: Pb(s) + PbO2(s) + 2H2SO4(aq) → 2PbSO4(s) + 2H2O(l)
Charging the Battery: During charging, the battery acts as an electrolytic cell. The reactions are exactly reversed. PbSO4(s) on anode and cathode is converted back into Pb and PbO2, respectively, and H2SO4 is regenerated.

3.2 Nickel-Cadmium (Ni-Cd) Cell

Another important secondary cell which has a longer life than the lead storage cell but is more expensive to manufacture.

Overall Reaction during discharge:
Cd(s) + 2Ni(OH)3(s) → CdO(s) + 2Ni(OH)2(s) + H2O(l)

4. Fuel Cells

Production of electricity by thermal plants is not very efficient (~40%) and causes massive pollution. Galvanic cells that are designed to convert the energy of combustion of fuels like hydrogen, methane, methanol, etc., directly into electrical energy are called fuel cells.

Hydrogen-Oxygen Fuel Cell

One of the most successful fuel cells uses the reaction of hydrogen with oxygen to form water. This cell was famously used for providing electrical power in the Apollo space programme. The water vapors produced during the reaction were condensed and added to the drinking water supply for the astronauts.

Cell Reactions:

Anode (Oxidation): 2H2(g) + 4OH-(aq) → 4H2O(l) + 4e-

Cathode (Reduction): O2(g) + 2H2O(l) + 4e- → 4OH-(aq)

Overall Reaction: 2H2(g) + O2(g) → 2H2O(l)
Advantages of Fuel Cells:
1. Highly Efficient: They operate at an efficiency of around 70% compared to thermal plants (40%).
2. Pollution Free: The only by-product in the H2-O2 cell is water.
3. Continuous Supply: The cell runs continuously as long as the reactants (fuel and oxidant) are supplied.

5. Corrosion (Rusting of Iron)

Corrosion slowly coats the surfaces of metallic objects with oxides or other salts. Examples: Rusting of iron, tarnishing of silver, development of green coating on copper.

Corrosion is essentially an electrochemical phenomenon. At a particular spot on an object made of iron, oxidation takes place, and that spot behaves as an anode.

The Electrochemical Theory of Rusting:

At Anode: Iron is oxidized to Fe2+.
2Fe(s) → 2Fe2+(aq) + 4e-     [E°(Fe2+/Fe) = -0.44 V]

Electrons released at anodic spot move through the metal to another spot where they reduce oxygen in the presence of H+ ions (which come from H2CO3 formed by dissolution of atmospheric CO2 in water moisture). This spot behaves as a cathode.

At Cathode:
O2(g) + 4H+(aq) + 4e- → 2H2O(l)     [E°(H+/O2/H2O) = 1.23 V]

Overall Cell Reaction for Rusting:
2Fe(s) + O2(g) + 4H+(aq) → 2Fe2+(aq) + 2H2O(l)
E°cell = 1.23 V - (-0.44 V) = 1.67 V

The ferrous ions are further oxidized by atmospheric oxygen to form ferric ions, which come out as rust in the form of hydrated ferric oxide (Fe2O3.xH2O) and with further production of hydrogen ions.

Prevention of Corrosion:

  • Barrier Protection: Coating the surface with paint, oil, grease, or chemicals like bisphenol.
  • Sacrificial Protection (Galvanization): Covering the iron surface with a more reactive metal (one having a more negative E°), such as Zinc. Even if the zinc coating is broken, zinc acts as the anode and is oxidized first, protecting the iron (cathode).

6. Previous Year Questions (PYQs) & Question Bank

Part A: Conceptual & Give Reasons (1-2 Marks)

[CBSE 2017, 2019]

Q1. Why does the voltage of a mercury cell remain constant throughout its useful life?

Answer: The overall cell reaction of a mercury cell is: Zn(Hg) + HgO(s) → ZnO(s) + Hg(l). As evident, the reaction does not involve any ions in the solution whose concentration can change over time. Therefore, the cell potential remains constant (1.35 V) throughout its life.
[CBSE 2018, 2020]

Q2. What happens to the pH of the electrolyte in a lead storage battery during discharging?

Answer: During discharging, sulphuric acid (H2SO4) is consumed to form PbSO4 and water. Because the concentration of H+ ions decreases and water is produced (diluting the acid), the pH of the electrolyte increases.
[CBSE 2015, 2021]

Q3. Mention two advantages of fuel cells over ordinary cells.

Answer: 1. Higher efficiency: Fuel cells are about 70% efficient, whereas thermal plants are only about 40% efficient.
2. Eco-friendly: They are pollution-free (e.g., the H2-O2 fuel cell only produces water).
3. Continuous operation: They do not go "dead" as long as the fuel is continuously supplied.

Part B: Application Based (3 Marks)

[CBSE 2016, 2019]

Q4. Write the cell reactions which occur in a lead storage battery (i) when the battery is in use (discharging) and (ii) when the battery is being charged.

Answer:
(i) During Discharging:
Anode: Pb(s) + SO42-(aq) → PbSO4(s) + 2e-
Cathode: PbO2(s) + SO42-(aq) + 4H+(aq) + 2e- → PbSO4(s) + 2H2O(l)

(ii) During Charging (Reactions are reversed):
Cathode (Reduction): PbSO4(s) + 2e- → Pb(s) + SO42-(aq)
Anode (Oxidation): PbSO4(s) + 2H2O(l) → PbO2(s) + SO42-(aq) + 4H+(aq) + 2e-
[CBSE 2023 Sample Paper]

Q5. Galvanization prevents rusting of iron even if the galvanized coating is scratched. Explain why.

Answer: Galvanization is the coating of iron with Zinc. Zinc is more electropositive (more reactive, lower standard reduction potential, E° = -0.76 V) than Iron (E° = -0.44 V). Even if the scratch exposes the iron, zinc will preferentially undergo oxidation (act as the anode) and supply electrons to the iron, making the iron act as a cathode. Thus, iron is protected from rusting. This is called sacrificial protection.

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This concludes the exhaustive series on the Electrochemistry Chapter for CBSE Class 12.

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