Batteries — Primary vs Secondary, Lead Acid, Lithium Ion Comparison

Question

What is the difference between primary and secondary batteries, and how do lead-acid and lithium-ion batteries work?

Solution — Step by Step

graph TD
    A[Batteries] --> B[Primary - Non-rechargeable]
    A --> C[Secondary - Rechargeable]
    B --> B1[Dry cell / Leclanche cell]
    B --> B2[Mercury cell]
    B --> B3[Alkaline cell]
    C --> C1[Lead-acid battery]
    C --> C2[Nickel-cadmium NiCd]
    C --> C3[Lithium-ion Li-ion]

Primary batteries undergo irreversible reactions — once the reactants are consumed, the battery is dead. Secondary batteries use reversible reactions — passing current in the opposite direction regenerates the reactants.

Used in cars, inverters, UPS systems.

Anode (Pb):

\text{Pb} + \text{SO}_4^{2-} \rightarrow \text{PbSO}_4 + 2e^-

Cathode (PbO2):

\text{PbO}_2 + 4\text{H}^+ + \text{SO}_4^{2-} + 2e^- \rightarrow \text{PbSO}_4 + 2\text{H}_2\text{O}

Overall:

\text{Pb} + \text{PbO}_2 + 2\text{H}_2\text{SO}_4 \rightarrow 2\text{PbSO}_4 + 2\text{H}_2\text{O}

Cell EMF = 2.05 V per cell. A car battery has 6 cells in series = 12.3 V.

During charging, the reaction reverses — PbSO4 converts back to Pb and PbO2.

Used in phones, laptops, EVs.

Anode: Graphite intercalated with lithium

\text{LiC}_6 \rightarrow \text{C}_6 + \text{Li}^+ + e^-

Cathode: Metal oxide (LiCoO2 or LiFePO4)

\text{Li}^+ + e^- + \text{CoO}_2 \rightarrow \text{LiCoO}_2

Li-ion advantages over lead-acid:

3-4x higher energy density (lighter, smaller)
No memory effect
Thousands of charge cycles
No toxic lead

Feature	Lead-Acid	Lithium-Ion	Dry Cell (Primary)
Type	Secondary	Secondary	Primary
Voltage	2.05 V/cell	3.6-3.7 V/cell	1.5 V
Energy density	Low (30-50 Wh/kg)	High (150-250 Wh/kg)	Low
Cycle life	500-1000 cycles	2000-5000 cycles	Single use
Cost	Low	High (dropping)	Very low
Common use	Cars, inverters	Phones, laptops, EVs	Torches, remotes
Electrolyte	Dilute H2SO4	Organic solvent with Li salt	NH4Cl paste

Why This Works

All batteries convert chemical energy to electrical energy through redox reactions. The key difference between primary and secondary is whether the electrode reactions are reversible. In lead-acid, both electrodes form PbSO4 during discharge — a solid product that stays on the electrode surface, making it easy to reverse. In dry cells, the zinc casing physically dissolves, making reversal impossible.

CBSE boards ask for the reactions in lead-acid batteries (2-3 marks). Write the anode, cathode, and overall reactions. JEE may ask you to calculate the EMF using standard electrode potentials.

Alternative Method

For understanding why lead-acid gives 2 V per cell, use the standard reduction potentials:

$E^o(\text{PbO}_2/\text{PbSO}_4) = +1.69$ V
$E^o(\text{PbSO}_4/\text{Pb}) = -0.36$ V

E^o_{\text{cell}} = E^o_{\text{cathode}} - E^o_{\text{anode}} = 1.69 - (-0.36) = 2.05 \text{ V}

Common Mistake

Students write that “during charging of a lead-acid battery, H2SO4 is consumed.” This is wrong — it is the opposite. During DISCHARGING, H2SO4 is consumed (converted to water). During CHARGING, H2SO4 is regenerated. That is why the specific gravity of the electrolyte drops when the battery is low and rises when fully charged. Getting this backwards is a common exam error.

Question

Solution — Step by Step

Why This Works

Alternative Method

Common Mistake

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