B.Sc. 2nd Year Chemistry — Electrochemistry
Electrochemistry is the branch of chemistry that deals with the relationship between chemical reactions and electrical energy. Many chemical reactions are associated with the transfer of electrons. Such electron-transfer reactions are called oxidation-reduction or redox reactions.
An electrochemical cell is a device in which chemical energy is converted into electrical energy or electrical energy is used to produce a chemical change. Electrochemical cells have important applications in batteries, fuel cells, corrosion studies, electroplating, electrolysis and industrial processes.
An electrochemical cell consists of two electrodes in contact with an electrolyte or electrolytes and provides a medium for electron transfer. Depending on the direction of energy conversion, electrochemical cells are mainly classified into galvanic cells and electrolytic cells.
| Galvanic Cell | Electrolytic Cell |
|---|---|
| Converts chemical energy into electrical energy. | Converts electrical energy into chemical energy. |
| Reaction is spontaneous. | Reaction is generally non-spontaneous. |
| Produces electrical energy. | Consumes electrical energy. |
| Example: Daniell cell. | Example: electrolysis of molten NaCl. |
A galvanic cell is an electrochemical cell in which a spontaneous redox reaction produces electrical energy.
Oxidation occurs at the anode and reduction occurs at the cathode. Electrons flow through the external circuit from the anode to the cathode.
An electrolytic cell uses external electrical energy to drive a non-spontaneous chemical reaction. Electrolysis is therefore the process in which electrical energy produces chemical change.
Examples include electrolysis of water, molten sodium chloride and aqueous solutions of various electrolytes.
When a metal electrode is dipped into a solution containing its ions, a potential difference develops between the metal and the solution. This potential difference is called electrode potential.
It arises because of the tendency of the metal to lose electrons and the tendency of its ions to gain electrons.
The tendency of an electrode to undergo oxidation and release electrons is called oxidation potential.
The tendency of an electrode to accept electrons and undergo reduction is called reduction potential.
The electrode potential measured under standard conditions is called standard electrode potential, represented by E°.
Standard conditions generally include:
Standard reduction potentials are commonly used to compare the oxidizing and reducing tendencies of different electrode systems.
The standard hydrogen electrode (SHE) is used as the primary reference electrode for measuring electrode potentials.
It consists of a platinum electrode immersed in a solution containing hydrogen ions of unit activity while hydrogen gas is passed over the platinum surface at standard pressure.
The standard electrode potential of the standard hydrogen electrode is defined as zero volts.
Reference electrodes provide a stable and reproducible potential against which the potential of another electrode can be measured.
The calomel electrode consists of mercury, mercurous chloride and potassium chloride solution.
It is widely used as a reference electrode in electrochemical measurements.
Another commonly used reference electrode is the silver-silver chloride electrode.
The Daniell cell is a common example of a galvanic cell. It consists of a zinc electrode immersed in zinc sulfate solution and a copper electrode immersed in copper sulfate solution. The two half-cells are connected by a salt bridge.
The zinc electrode acts as the anode and the copper electrode acts as the cathode.
A galvanic cell is conveniently represented using cell notation.
A single vertical line represents a phase boundary, while a double vertical line represents the salt bridge or liquid junction.
The electromotive force (EMF) of a cell is the maximum potential difference between the two electrodes when no current is being drawn from the cell.
Under standard conditions:
The Nernst equation gives the electrode or cell potential under non-standard conditions.
At 298 K, the equation becomes:
Where E is the electrode or cell potential, E° is the standard potential, R is the gas constant, T is absolute temperature, n is the number of electrons transferred, F is the Faraday constant and Q is the reaction quotient.
The electrical work produced by a galvanic cell is related to the Gibbs free energy change of the reaction.
Under standard conditions:
For a spontaneous cell reaction, Ecell is positive and ΔG is negative.
For a reaction at equilibrium, the relationship between standard Gibbs free energy and equilibrium constant is:
Since:
Therefore:
At 298 K:
A concentration cell is an electrochemical cell in which the two electrodes are chemically identical but the concentrations of the electrolyte are different.
The EMF is produced because of the difference in concentration.
Concentration cells are useful for determining ionic activities, concentration differences and transport properties.
The concentration difference occurs in the electrodes or gases associated with the electrodes.
The electrodes are identical but the electrolyte concentrations differ.
When two solutions of different concentrations or different electrolytes are in contact, ions diffuse from one solution to another at different rates. This produces a potential difference called liquid junction potential.
The salt bridge helps reduce liquid junction potential by allowing ions of similar mobility to migrate between the two half-cells.
A salt bridge is a tube containing an inert electrolyte, usually a solution of KCl or KNO₃, connecting the two half-cells of an electrochemical cell.
The arrangement of elements or ions according to their standard reduction potentials is called the electrochemical series.
A species with a more positive reduction potential has a greater tendency to undergo reduction, whereas a species with a more negative reduction potential generally has a greater tendency to undergo oxidation in its reduced form.
Corrosion is the gradual deterioration of a metal due to chemical or electrochemical reaction with its environment.
Rusting of iron is a common example of electrochemical corrosion.
In electrochemical corrosion, different areas of the metal surface behave as anode and cathode.
Ferrous ions react with hydroxide ions and subsequently undergo further oxidation and hydration to produce hydrated iron oxides, commonly called rust.
A battery is a combination of one or more electrochemical cells used as a source of electrical energy.
Primary batteries are generally not designed for repeated recharging. Examples include dry cells and alkaline batteries.
Secondary batteries can be recharged by passing current in the reverse direction. Examples include lead-acid and lithium-ion batteries.
A fuel cell converts the chemical energy of continuously supplied fuel and oxidant directly into electrical energy.
The hydrogen-oxygen fuel cell is an important example.
Fuel cells are efficient and can produce water as the major product when hydrogen and oxygen are used.
Electrochemistry deals with the relationship between chemical reactions and electrical energy. Electrochemical cells are mainly classified into galvanic and electrolytic cells. Galvanic cells produce electrical energy from spontaneous redox reactions, whereas electrolytic cells use electrical energy to drive non-spontaneous reactions. Electrode potential, standard electrode potential, EMF, the Nernst equation and electrochemical series are important concepts for understanding cell behavior. Salt bridges maintain electrical neutrality and complete the circuit. Electrochemical principles are widely applied in batteries, fuel cells, corrosion control, electroplating and industrial processes.