Galvanic Cell (original) (raw)

Last Updated : 23 Jul, 2025

**Galvanic Cell also called Voltaic Cell is an electrochemical device that converts spontaneous chemical energy generated in a redox reaction into electrical energy.

Table of Content

• What is Galvanic Cell?
• Cell Definition
• Electrolytic Cell Definition
• Electrochemical Cell Definition
• Primary Cell & Secondary Cell
• What is a Galvanic Cell?
• Parts of Galvanic Cell
• Constructions of Galvanic Cell
• Principle and Working of Galvanic Cell
• Electrode Potential
• Standard Electrode Potential
• Cell Potential
• Example of Galvanic Cell
• What is Salt Bridge?
• Difference between Galvanic Cell and Electrolytic Cell
• Solved Examples on Galvanic Cells

What is Galvanic Cell?

We define a Galvanic cell as a device that converts the chemical energy of the redox reaction to electrical energy, this is a type of electrochemical cell that uses electrolytes to produce the electrical energy.

To understand the concept of a Galvanic Cell, let's first understand what is cell and what are its types.

Cell Definition

A Cell is an electrical device that when connected to a circuit generates a potential difference which results in the flow of charge or ions from higher potential to lower potential. A Cell is a unit source of power. When cells are combined together to create potential differences then it is called Battery.

Depending upon the conversion of energy from chemical to electrical or electrical to chemical there are two types of cells:

• Electrolytic Cell
• Electrochemical Cell

Electrolytic Cell Definition

An Electrolytic Cell is a device that converts electrical energy into chemical energy. It means it already has the power supply which is used in 'lysis' means the breaking of electrolytes into ions which then moves towards electrodes to constitute current and produce electrical energy. In an electrolytic cell, Anode is +ve while Cathode is -ve. In this type of cell, the flow of electrons is from Anode to Cathode.

**Learn more about, **Electrolytic Cells and Electrolysis

Electrochemical Cell Definition

An Electrochemical Cell is a device that converts chemical energy into electrical energy. It means the chemical energy stored in the cell undergoes a reaction to produce electrical energy. In this type of cell, Anode is -ve while Cathode is +ve. The flow of electrons is from Cathode to Anode. This cell is the reverse of an electrochemical cell.

Let's learn about Galvanic Cells in detail in this article.

Primary Cell & Secondary Cell

**Primary Cells are non-rechargeable cells. Once, the chemical energy stored in the cell consumes the cell becomes useless. It is disposable in nature.

**Secondary Cells are rechargeable cells. It is based on the principle of reverse chemical reactions i.e. electrical energy is used to produce electrons inside the cell which is used to run the device. It is economically and environmentally advantageous than Primary Cells.

What is a Galvanic Cell?

The devices in which chemical reaction is used to produce electrical energy are called Galvanic Cells or Voltaic Cells. In these devices, the Gibbs Energy of the spontaneous Redox Reaction is converted into electrical work that can be used to drive a motor or to power electrical equipment such as heaters, fans, geysers, etc.

These cells are greatly important because of their many practical applications. An early example of the Galvanic Cell is Daniel's Cell invented by the British chemist John Daniels in 1836.

Daniel's cell was constructed based on the redox reaction:

**Zn(s) + Cu² + (aq)+ → Zn² + (aq)+ Cu(s)

Half-Cell Reactions

In these cells, Oxidation and Reduction reactions occur in autonomous containers called half cells or Redox Couples. The half cell of the reaction is represented as follows

**Anode(Oxidation): Zn **→ Zn2+ + 2e-

Cathode(Reduction): Cu2+ + 2e-**→**Cu

Oxidation occurring at Anode is referred to as Oxidation Half Cell while the Reduction occurring at Cathode is called Reduction Half Cell. Although these half-cell reactions are occurring in separate containers they are connected with each other internally and externally. Internally they are connected via Salt Bridge while externally they are connected via wire, voltmeter, and switch.

The above-shown redox reaction is spontaneous. General Representation of a Galvanic Cell:

**M 1 (s)/M 1 n+ (aq) || M₂ n+ (aq)/M₂(s)

Parts of Galvanic Cell

Various parts of the Galvanic Cell include,

• **Anode: The electrode at which oxidation occurs is the anode.
• **Cathode: The electrode at which reduction occurs is the cathode.
• **Salt bridge: It is a tube filled with electrolytes that maintain the neutrality of the Galvanic Cell.
• **Half-cells: Two separate beakers where oxidation and reduction occur are called Half-Cell.
• **External circuit: It helps to conduct the flow of electrons between electrodes.

Constructions of Galvanic Cell

• A Galvanic Cell is made by combining two electrodes an oxidation electrode and a reduction electrode. Both electrodes individually are called the half cell.
• The two half-cells are individually filled with different electrolytic solutions which helps in their particular reaction. Both the half cell are connected to each other using a Salt Bridge internally and externally via wire, switch, and voltmeter.
• Oxidation occurs at the oxidation electrode which releases the free electrons that accumulate on the electrode and provide a negative potential. The electrode at which oxidation occurs is called the Anode.
• Reduction occurs at the reduction electrode that generates the positive charge and provides the positive potential. The electrode at which reduction occurs is called the Cathode.
• Connecting these electrodes via wire, switch, and voltmeter initiates the flow of electrons from one electrode to another resulting in a flow of electric current. For a galvanic cell, the anode is negatively charged and the cathode is positively charged.

A Galvanic cell image is added below in the article.

Principle and Working of Galvanic Cell

The working of Galvanic Cell is discussed below:

• Electrodes are exposed to the electrolyte at the electrode-electrolyte interface, which generates ions in the electrolyte solution making one metal electrode negatively charged.
• For the other electrode, the metal ions in the electrolyte solution deposit on the other metal electrode making the electrode positively charged.
• Due to this charge separation, a potential difference is developed between electrodes and electrolytes. This potential difference is called electrode potential.
• The electrode, where oxidation occurs, is the anode while the electrode where reduction occurs is the cathode.
• Anode is at a negative potential with respect to the solution while the cathode is at a positive potential with respect to the solution.
• The potential difference between these two electrodes is called the potential of a Galvanic cell and is responsible for the flow of electrons in the circuit.

Electrode Potential

In each half-cell, there is a movement of electrons between the electrodes and the electrolyte. Since there is a flow of charge between the electrode and electrolyte there develops a potential called Electrode Potential. There are two types of Electrode Potential, Oxidation Potential, and Reduction Potential. Their representation is given below:

**Oxidation Potential: M **→ Mn+ + ne-

**Reduction Potential: Mn+ + ne- **→ M

The Electrode Potential is affected by the nature of metal and ion, its concentration, and temperature. The Electrode Potential of a half-Cell is written in terms of its Reduction Potential. Hence, the Electrode Potential of the Oxidation half-cell is -ve and that of the Reduction half-cell is +ve.

Standard Electrode Potential

The Electrode Potential calculated above is relative in nature. In order to find the individual potential of an electrode we use a Standard Hydrogen Electrode whose potential is zero to calculate Standard Electrode Potential.

A Standard Hydrogen Electrode consists of a Platinum Wire covered with Platinum foil in a test tube which is immersed in a 1M concentration of HCl which liberated H+ ion and hydrogen gas is bubbled in it at 1atm at 298K of temperature.

Standard Electrode Potential is the potential of an Electrode dipped in 1M concentration of its salt in a half cell and this half cell is connected to a Standard Hydrogen Electrode via a salt bridge.

It is represented as

**M | M n+ (1M) || H + (1M) | H 2 (1 atm), Pt

Cell Potential

Cell Potential refers to the potential difference between the cathode and anode of the Galvanic Cell. When no current is drawn from it i.e. the two electrodes are not connected with each other then it is called Cell Electromotive Force or EMF of Galvanic Cell.

The convention to represent cell potential follows that the anode potential is written on the left side and the cathode potential is written on the right side and both are separated by two vertical lines (||). For example, **M 1 (s)/M 1 n+ (aq) || M₂ n+ (aq)/M₂(s).

Hence, the left potential is Anode Potential while the right potential is Cathode Potential. Hence, Cell Potential, Ecell is given as

**E cell = E right - E left

Example of Galvanic Cell

Daniel’s cell is the most common example of a galvanic cell. The galvanic cell converts chemical energy into electrical energy. For a Galvanic Cell Copper Ions are reduced at the cathode and Zinc Ions are oxidized at the anode.

Reactions taking place at the cathode and anode of a Galvanic cell are:

**At Anode: Zn → Zn2+ + 2e–

**At Cathode: Cu2+ + 2e– → Cu

What is Salt Bridge?

Salt Bridge is a U- shaped tube that contains a concentrated solution of inert electrolytes. Some examples of electrolytes used in the salt bridge are KCl, KNO3, K2SO4, etc. These inert electrolytes do not participate in the cell reaction.

Salt Bridge allows the movement of ions from one solution to the other without mixing two solutions. The salt bridge also helps to maintain the electrical neutrality of the solution in the two half-cells.

Difference between Galvanic Cell and Electrolytic Cell

Galvanic Cells and Electrolytic Cells are both electrochemical cells and the major difference between them is as follows:

**Read More:

• Electrochemical Cells
• Redox Reactions and Electrode Processes
• Chemical Effects of Electric Current

Solved Examples on Galvanic Cells

**Example 1: Calculate Δ r G φ for the reaction:

**Mg(s)+Cu 2 + (aq) → Mg 2 (aq)+Cu(s)

**Given E 0 cell =2.71 V, 1F = 96500 C mol -1

**Solution:

ΔrGφ = -nF
Eocell = 2.71 V,
1 F = 96500 C mol-1,
n = 2

ΔrGφ = -2×96500 C mol-1 ×2.71 V

= -523030 J mol-1 (1CV = 1J)

= -523.080 kJ mol-1

**Example 2: The ΔG φ for the Daniell cell has been found to be -212.3 kJ at 25°C. Calculate the equilibrium constant for the cell reaction.

**Solution:

ΔGφ =-RT ln Kc

ΔGφ = -212.3 kJ = -212300 J,
T = 298 K

Here
R=8.314.K-1 mol-1

ln(Kc) = 212300 / (8.314 × 298)
= 85.69

Kc = 1.64 × 1037

**Example 3: What does the negative sign in the expression E o Zn2+/Zn =-0.76 V mean?

**Solution:

It means that zinc is more reactive than hydrogen. When zinc electrode is connected to SHE, zinc will get oxidized and H+ will get reduced.

**Example 4: A galvanic cell has an electrical potential of 1.1 V. If an opposing potential of 1.1 V is applied to this cell. What will happen to the reactance of the cell and the current flowing in the cell?

**Solution:

When the opposing potential becomes equal to the electric potential, the reaction of the cell stops and no current flows through the cell. Thus, no chemical reaction takes place.