Dry Cell Electrolytic Cells and Galvanic Cells

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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.

$Z n (s) \to Z n^{2 +} (a q) + 2 e^{-}$

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

$C u^{2 +} (a q) + 2 e^{-} \to C u (s)$

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. ( $E_{cell} < 0$ )

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.

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

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

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

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

Feature	Galvanic Cell	Electrolytic Cell
Spontaneity	Spontaneous (ΔG<0)	Non-spontaneous (ΔG>0)
Energy Conversion	Chemical → Electrical	Electrical → Chemical
Anode	Negative	Positive
Cathode	Positive	Negative
Salt Bridge	Required (to maintain neutrality)	Not required
Reaction	Redox reactions occur to generate current	Redox reactions occur due to external current

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 (NH4Cl) 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):

$Z n (s) \to Z n^{2 +} (a q) + 2 e^{-}$

Cathode (Reduction):

$2 M n O_{2} (s) + 2 N H_{4}^{+} (a q) + 2 e^{-} \to M n_{2} O_{3} (s) + 2 N H_{3} (a q) + H_{2} O (l)$

Overall Reaction:

$Z n (s) + 2 M n O_{2} (s) + 2 N H_{4}^{+} (a q) \to Z n^{2 +} (a q) + M n_{2} O_{3} (s) + 2 N H_{3} (a q) + H_{2} O (l)$

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.

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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.

$Z n (s) \to Z n^{2 +} (a q) + 2 e^{-}$

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

$C u^{2 +} (a q) + 2 e^{-} \to C u (s)$

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. ( $E_{cell} < 0$ )

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.

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

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

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

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

Feature	Galvanic Cell	Electrolytic Cell
Spontaneity	Spontaneous (ΔG<0)	Non-spontaneous (ΔG>0)
Energy Conversion	Chemical → Electrical	Electrical → Chemical
Anode	Negative	Positive
Cathode	Positive	Negative
Salt Bridge	Required (to maintain neutrality)	Not required
Reaction	Redox reactions occur to generate current	Redox reactions occur due to external current

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 (NH4Cl) 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):

$Z n (s) \to Z n^{2 +} (a q) + 2 e^{-}$

Cathode (Reduction):

$2 M n O_{2} (s) + 2 N H_{4}^{+} (a q) + 2 e^{-} \to M n_{2} O_{3} (s) + 2 N H_{3} (a q) + H_{2} O (l)$

Overall Reaction:

$Z n (s) + 2 M n O_{2} (s) + 2 N H_{4}^{+} (a q) \to Z n^{2 +} (a q) + M n_{2} O_{3} (s) + 2 N H_{3} (a q) + H_{2} O (l)$

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.

Frequently Asked Questions

Why is the anode negative in a galvanic cell but positive in an electrolytic cell?

What is the role of the salt bridge in a galvanic cell?

Why are dry cells not rechargeable?

What is the purpose of the depolarizer in a dry cell?

Join ALLEN!

Galvanic cell, Electrolytic cell, and Dry cell

1.0Galvanic Cell (Voltaic Cell)

Principle and Working

Characteristics

2.0Electrolytic Cell

Principle and Working

Characteristics

3.0Galvanic vs. Electrolytic Cells

4.0Dry Cell: A Practical Galvanic Cell

Construction and Chemistry

Reactions in a Dry Cell

Uses and Limitations

Table of Contents

Frequently Asked Questions

Why is the anode negative in a galvanic cell but positive in an electrolytic cell?

What is the role of the salt bridge in a galvanic cell?

Why are dry cells not rechargeable?

What is the purpose of the depolarizer in a dry cell?

Join ALLEN!

Galvanic cell, Electrolytic cell, and Dry cell

1.0Galvanic Cell (Voltaic Cell)

Principle and Working

Characteristics

2.0Electrolytic Cell

Principle and Working

Characteristics

3.0Galvanic vs. Electrolytic Cells

4.0Dry Cell: A Practical Galvanic Cell

Construction and Chemistry

Reactions in a Dry Cell

Uses and Limitations

Table of Contents