Galvanic cells can be broadly classified based on whether their underlying chemical reactions can be reversed by passing an external electrical current through them. They fall into two main categories: Primary Cells and Secondary Cells.

Definition: Primary cells are non-rechargeable electrochemical cells in which the chemical reactions occur in only one direction. Once the active chemical reactants are completely exhausted, the cell stops producing electricity and becomes dead or discharged.

Key Characteristics:

• Irreversible Reactions: The redox reactions cannot be reversed by applying an external electric current.
• Internal Resistance: They possess higher internal resistance, which increases as reaction products build up over time.
• Portability: Convenient for low-drain devices, highly lightweight, inexpensive, and ready for immediate use.
• Examples: Daniell Cell, Leclanché (Dry) Cell, Mercury Button Cell, Alkaline Batteries.

Definition: Secondary cells (also known as storage cells or accumulators) are rechargeable electrochemical cells where the chemical reactions are reversible. They store electrical energy as chemical energy during charging and release it on demand during discharge.

Key Characteristics:

• Dual Mode Operation:
  1. Discharging: Operates as a Galvanic cell, converting internal chemical energy into electrical energy.
  2. Charging: Operates as an Electrolytic cell, using an external DC supply to drive the reverse reaction and restore the active materials.
• Low Internal Resistance: Capable of delivering high surge currents when required.
• Examples: Lead-Acid Accumulator, Nickel-Cadmium (Ni-Cd) Cell, Lithium-ion (Li-ion) Battery.

Definition & Background: Invented by British chemist John Frederic Daniell in 1836, the Daniell cell is a classic primary cell designed to eliminate polarization (hydrogen gas bubble accumulation on the electrode) that affected earlier cell designs.

It consists of two half-cells connected externally by a conducting wire and internally by a porous pot or salt bridge:

1. Anode (Negative Electrode, -): Zinc ($\text{Zn}$) rod immersed in 1.0 M Zinc Sulfate ($\text{ZnSO}_4$) solution. Site of oxidation.
2. Cathode (Positive Electrode, +): Copper ($\text{Cu}$) container or rod immersed in 1.0 M Copper(II) Sulfate ($\text{CuSO}_4$) solution. Site of reduction.
3. Porous Barrier / Salt Bridge: Keeps the two electrolyte solutions separate to prevent direct chemical mixing while allowing ions to migrate to maintain electrical neutrality.

Working Principle & Reactions:

Zinc undergoes oxidation more readily than copper, releasing electrons that flow through the external circuit to the copper electrode, where $\text{Cu}^{2+}$ ions are reduced and deposited as metallic copper.

At Anode (oxidation):     Zn(s) → Zn²⁺(aq) + 2e⁻

At Cathode (reduction):   Cu²⁺(aq) + 2e⁻ → Cu(s)

Overall Cell Reaction:    Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)

Electromotive Force:      E°cell = +1.10 Volts

Definition & Background: Invented by Georges Leclanché in 1866 and later developed into the portable "dry cell" format, this primary cell is widely used in low-drain household appliances.

To make the cell portable and leak-resistant, liquid solutions are replaced with a moist paste:

1. Anode Container (Negative Electrode, -): A cylindrical Zinc ($\text{Zn}$) outer vessel that serves as both the container and the anode.
2. Cathode Rod (Positive Electrode, +): A central Carbon/Graphite rod surrounded by a packed mixture of Manganese(IV) Oxide ($\text{MnO}_2$) and powdered carbon.
3. Moist Electrolyte Paste: A moist paste containing Ammonium Chloride ($\text{NH}_4\text{Cl}$) and Zinc Chloride ($\text{ZnCl}_2$).
4. Depolarizer Role: $\text{MnO}_2$ acts as a depolarizer by oxidizing hydrogen gas produced during operation into water, maintaining steady current flow.

Working Principle & Reactions:

At Anode (oxidation):     Zn(s) → Zn²⁺(aq) + 2e⁻

At Cathode (reduction):   2NH₄⁺(aq) + 2MnO₂(s) + 2e⁻ → Mn₂O₃(s) + 2NH₃(aq) + H₂O(l)

Overall Cell Reaction:    Zn(s) + 2NH₄⁺(aq) + 2MnO₂(s) → Zn²⁺(aq) + Mn₂O₃(s) + 2NH₃(aq) + H₂O(l)

Electromotive Force:      E°cell ≈ 1.50 Volts

Definition & Background: Invented by Gaston Planté in 1859, the Lead-Acid battery is the primary example of a rechargeable secondary cell. It is heavy, reliable, and capable of delivering extremely large surge currents for automobile ignition and backup power (UPS).

A standard 12V lead-acid battery comprises six individual 2V cells connected in series inside a heavy-duty casing:

1. Anode Grid (Negative Electrode, -): A metallic grid stuffed with porous spongy Lead ($\text{Pb}$).
2. Cathode Grid (Positive Electrode, +): A metallic grid packed with Lead(IV) Oxide ($\text{PbO}_2$).
3. Electrolyte Solution: Aqueous solution of Sulfuric Acid ($\text{H}_2\text{SO}_4$) with a specific gravity of around 1.28 when fully charged.

Working Principle & Reactions (Discharging & Charging):

During discharge, both electrodes convert into Lead(II) Sulfate ($\text{PbSO}_4$), consuming acid. During charging, an external power supply forces current in reverse to regenerate lead and lead oxide.

DISCHARGING (Galvanic Mode):
At Anode (oxidation):     Pb(s) + SO₄²⁻(aq) → PbSO₄(s) + 2e⁻
At Cathode (reduction):   PbO₂(s) + 4H⁺(aq) + SO₄²⁻(aq) + 2e⁻ → PbSO₄(s) + 2H₂O(l)

Overall Discharge:        Pb(s) + PbO₂(s) + 2H₂SO₄(aq) → 2PbSO₄(s) + 2H₂O(l)

CHARGING (Electrolytic Mode):
Overall Charging:         2PbSO₄(s) + 2H₂O(l) → Pb(s) + PbO₂(s) + 2H₂SO₄(aq)

Electromotive Force:      ~2.0 Volts per cell (12 Volts for a standard 6-cell battery)