An electrochemical cell is a device that either generates electrical energy from chemical reactions or uses electrical energy to drive chemical reactions.
It consists of two electrodes (anode and cathode) in contact with an electrolyte, a medium containing ion that conduct electricity.
The electrodes, typically made of conductive materials like metals or carbon, may also have catalysts to facilitate the reactions.
Types of Electrochemical Cells
1) Galvanic cells:
In a galvanic cell, a spontaneous redox reaction occurs between the anode (site of oxidation) and cathode (site of reduction).
Electrons flow from the anode to the cathode through an external circuit, generating electric current, while the electrolyte maintains charge neutrality by allowing ion movement.
Example: The Daniell cell, with a zinc anode in zinc sulfate and a copper cathode in copper sulfate, connected by a salt bridge.
2) Electrolytic cells:
In electrolytic cells, an external voltage source drives a non-spontaneous redox reaction.
The anode (connected to the positive terminal) is the site of oxidation, and the cathode (negative terminal) is the site of reduction.
Applications include electroplating, electrorefining, and the electrolysis of water to produce hydrogen and oxygen.
Electrode potential:
The voltage developed by an electrode in an electrochemical cell when it is in contact with a solution containing its own ions.
It represents the tendency of the electrode to lose or gain electrons in a redox reaction.
Standard electrode potential:
The electrode potential of an electrochemical half-cell when all reactants and products are in their standard states (1M concentration, 1 atm pressure) and at a constant temperature (usually 25°C).
It indicates the tendency of a redox couple to undergo reduction or oxidation relative to the Standard Hydrogen Electrode (SHE).
Salt bridge:
A device used in electrochemical cells to connect two half-cells, allowing the flow of ions between them to maintain charge neutrality and complete the electrical circuit while minimizing the mixing of the two solutions.
It prevents the direct mixing of the solutions in each half-cell while allowing ions to flow and balance the charges.
Nernst equation:
The Nernst equation calculates the electrode potential under non-standard conditions and is expressed as:
E = E° - (RT/nF) * ln(Q)
Where:
E = electrode potential under non-standard conditions
E° = standard electrode potential of the half-cell
R = gas constant (8.314 J/mol·K)
T = temperature in Kelvin (K)
n = number of electrons involved in the redox reaction
F = Faraday's constant (96,485 C/mol)
Q = reaction quotient (ratio of product to reactant concentrations, raised to the power of their stoichiometric coefficients)
At 25°C (298 K), the Nernst equation simplifies to:
In this form, the natural logarithm (ln) is replaced by the common logarithm ), and the constant is approximated to 0.05916 V.