Galvanic cell, Electrolytic cell, and Dry cell 

Electrochemical cells are devices that either produce electrical energy from a chemical reaction or use electrical energy to cause a chemical reaction. They are categorized into two main types: Galvanic cells and Electrolytic cells. The dry cell is a specific, common example of a galvanic cell. Understanding these cells is crucial for JEE preparation, as they form the basis of electrochemistry.

1.0Galvanic Cell (Voltaic Cell)

A galvanic cell, also known as a voltaic cell, is an electrochemical cell that converts chemical energy into electrical energy through a spontaneous redox reaction. The process has a negative change in Gibbs free energy (ΔG<0) and a positive cell potential (Ecell​>0).

Principle and Working

A classic example is the Daniel cell, which uses zinc and copper electrodes.

• Anode (Negative Electrode): Zinc metal is oxidized, losing electrons.

$Zn(s)\to Z{n}^{2+}(aq)+2e^{-}$

• Cathode (Positive Electrode): Copper ions in the solution are reduced, gaining electrons.

$C{u}^{2+}(aq)+2e^{-}\to Cu(s)$

• Electron Flow: Electrons spontaneously flow from the zinc anode to the copper cathode through an external wire. This flow of electrons constitutes an electric current.
• Salt Bridge: A salt bridge connects the two half-cells, allowing ions to flow and maintain electrical neutrality, completing the circuit.

Characteristics

• Spontaneous Reaction: The redox reaction occurs on its own.
• Energy Conversion: Chemical energy → Electrical energy.
• Electrode Polarity: Anode is negative, Cathode is positive.
• Function: Used as a power source, like in batteries.

2.0Electrolytic Cell

An electrolytic cell is a device that uses electrical energy to drive a non-spontaneous redox reaction. This process requires an external power source and has a positive change in Gibbs free energy (ΔG>0) and a negative cell potential. ($E_{\text{cell}}<0$)

Principle and Working

An example is the electrolysis of molten sodium chloride (NaCl).

• External Power Source: A battery is used to supply energy. The positive terminal of the battery is connected to the anode, and the negative terminal to the cathode.
• Anode (Positive Electrode): Chloride ions are attracted to the anode and are oxidised.

$2C{l}^{-}(l)\to C{l}_2(g)+2e^{-}$

• Cathode (Negative Electrode): Sodium ions are attracted to the cathode and are reduced.

$N{a}^{+}(l)+e^{-}\to Na(s)$

• Overall Reaction: The non-spontaneous decomposition of molten NaCl into sodium metal and chlorine gas.

Characteristics

• Non-spontaneous Reaction: The reaction requires an external energy input.
• Energy Conversion: Electrical energy → Chemical energy.
• Electrode Polarity: Anode is positive, Cathode is negative.
• Function: Used for industrial processes like electroplating, metal refining, and the production of elements.

3.0Galvanic vs. Electrolytic Cells

4.0Dry Cell: A Practical Galvanic Cell

A dry cell is a type of galvanic cell that uses a paste-like electrolyte instead of a liquid solution. It's a common, single-use battery found in flashlights and remote controls.

Construction and Chemistry

• Anode: Zinc container, which also serves as a negative electrode.
• Cathode: A central carbon rod, which acts as the positive electrode.
• Electrolyte: A moist paste of ammonium chloride (NH4​Cl) and zinc chloride (ZnCl2​).
• Depolarizer: Manganese dioxide (MnO2​) is mixed with the carbon to prevent the buildup of hydrogen gas, which would otherwise stop the reaction.

Reactions in a Dry Cell

• Anode (Oxidation):

$Zn(s)\to Z{n}^{2+}(aq)+2e^{-}$

• Cathode (Reduction):

$2Mn{O}_2(s)+2N{H}_4^{+}(aq)+2e^{-}\to M{n}_2{O}_3(s)+2N{H}_3(aq)+H_{2}O(l)$

• Overall Reaction:

$Zn(s)+2Mn{O}_2(s)+2N{H}_4^{+}(aq)\to Z{n}^{2+}(aq)+M{n}_2{O}_3(s)+2N{H}_3(aq)+H_{2}O(l)$

Uses and Limitations

• Uses: Widely used in low-drain devices.
• Limitations: They are non-rechargeable. The zinc casing can corrode over time, and the paste can dry out, limiting its shelf life.