An electrochemical cell is a device capable of either generating electrical energy from chemical reactions or using electrical energy to cause chemical reactions. They are fundamentally divided into two major classes: Galvanic (Voltaic) Cells and Electrolytic Cells.

Definition: A Galvanic cell (or Voltaic cell) is an electrochemical cell that converts chemical energy into electrical energy through a spontaneous redox (reduction-oxidation) reaction.

A classic example of a Galvanic cell is the Daniell Cell. It consists of two separate half-cells connected externally by a wire and internally by a salt bridge:

1. Anode (Negative Electrode, -):
Zinc rod immersed in 1.0 M ZnSO₄ solution. Site of Oxidation (loss of electrons).
2. Cathode (Positive Electrode, +):
Copper rod immersed in 1.0 M CuSO₄ solution. Site of Reduction (gain of electrons).
3. External Circuit:
Metallic wire with a voltmeter connecting both electrodes to allow electrons to flow from Anode to Cathode.
4. Salt Bridge:
A U-tube filled with an inert electrolyte gel (e.g., KCl or KNO₃ in agar-agar). Function: Maintains electrical neutrality in both half-cells and completes the circuit.

Working Principle & Reactions:

Zinc has a higher tendency to lose electrons than copper. At the anode, zinc oxidizes into $\text{Zn}^{2+}$ ions, releasing electrons into the external circuit. These electrons travel through the wire to the copper cathode, where $\text{Cu}^{2+}$ ions gain them and deposit as copper metal.

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)

Cell Notation:            Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)

Primary Applications:
• Primary Batteries: Non-rechargeable dry cells (zinc-carbon) and alkaline batteries used in remotes and clocks.
• Secondary Batteries: Lead-acid car batteries and Lithium-ion batteries in smartphones and electric vehicles (during discharge).
• Fuel Cells: Hydrogen-oxygen fuel cells used in space vehicles to produce electricity and water.

Definition: An Electrolytic cell is an electrochemical cell that converts electrical energy into chemical energy to drive a non-spontaneous redox reaction through a process called electrolysis.

Construction & Components:

Unlike Galvanic cells, an electrolytic cell consists of a single container (cell vessel) holding the electrolyte and two electrodes connected to an external Direct Current (DC) power supply:

1. Direct Current (DC) Power Source:
Forces current into the system (Battery/Power Supply).
2. Anode (Positive Electrode, +):
Connected to the positive terminal of the DC battery. Attracts anions (negative ions); site of Oxidation.
3. Cathode (Negative Electrode, -):
Connected to the negative terminal of the DC battery. Attracts cations (positive ions); site of Reduction.
4. Electrolyte:
Molten ionic salt or aqueous solution containing free mobile ions.

Working Principle & Reactions (Example: Electrolysis of Molten NaCl):

When direct current is applied, positive $\text{Na}^+$ cations move to the cathode to gain electrons, while negative $\text{Cl}^-$ anions move to the anode to lose electrons.

At Cathode (reduction):   2Na⁺(l) + 2e⁻ → 2Na(l)

At Anode (oxidation):     2Cl⁻(l) → Cl₂(g) + 2e⁻

Overall Net Reaction:    2NaCl(l) → 2Na(l) + Cl₂(g)

Primary Applications:
• Electroplating: Coating metals with gold, silver, nickel, or chromium for protection or decoration.
• Electro-refining: Refining crude metals (e.g., purification of copper up to 99.99% purity).
• Extraction of Reactive Metals: Industrial extraction of sodium, aluminium (Hall-Héroult process), magnesium, and calcium from fused salts.
• Chlor-Alkali Industry: Industrial production of chlorine gas, hydrogen gas, and caustic soda ($\text{NaOH}$) from brine.