Electrochemical Cells and Cell Potentials - SS2 Chemistry Lesson Note

Electrochemical cells are devices that convert chemical energy into electrical energy through redox reactions. Electrochemical cells consist of two half-cells connected by a conductive material, such as a wire or salt bridge. Each half-cell contains an electrode immersed in an electrolyte solution.

Half-cell components:

a.    Anode: The electrode where oxidation occurs. It loses electrons and becomes oxidised.

b.    Cathode: The electrode where reduction occurs. It gains electrons and becomes reduced.

c.     Electrolyte: A solution containing ions that allow the flow of charge between the electrodes.

Cell notation: Electrochemical cells are often represented using a shorthand notation, such as:

Anode | Anode solution || Cathode solution | Cathode

For example, in a zinc-copper cell, the cell notation is:

Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)

Cell potential (Ecell): The driving force of an electrochemical cell. It is the measure of the potential difference between the two electrodes.

Standard cell potential (E°cell): The cell potential under standard conditions, which include 1 M concentration of electrolytes, 1 atm pressure, 298 K temperature, and electrodes in their standard states. The standard cell potential (E°cell) is determined by the difference in the standard reduction potentials (E°red) of the two half-reactions involved in the cell.

E°cell = E°red (cathode) - E°red (anode)

Reduction potential: The tendency of a substance to gain electrons and be reduced. It is measured in volts (V) or millivolts (mV). The more positive the reduction potential, the stronger the oxidising agent is. The more negative the reduction potential, the stronger the reducing agent is. The Nernst equation allows the calculation of the cell potential (Ecell) under non-standard conditions:

Ecell = E°cell - (0.0592/n) x log(Q)

where E°cell is the standard cell potential, n is the number of electrons transferred in the balanced equation, and Q is the reaction quotient.

Factors affecting cell potential:

a.    Concentration: Changes in reactant concentrations affect the reaction quotient (Q) and, thus, the cell potential.

b.    Temperature: Changes in temperature influence the reaction rates and, therefore, the cell potential.

c.     Pressure: Changes in gas partial pressures (for gaseous species) can affect cell potential, especially in fuel cells.

Types of electrochemical cells:

a.    Galvanic or Voltaic cells: Spontaneous redox reactions produce electrical energy. Electrons flow from the anode to the cathode through an external circuit.

b.    Electrolytic cells: Non-spontaneous redox reactions require an external power source to drive the reaction in the opposite direction.

Applications of electrochemical cells include batteries, fuel cells, corrosion protection, electroplating, and electrolysis processes.

Understanding electrochemical cells and cell potentials provides insights into the conversion of chemical energy into electrical energy. It allows for the prediction of cell behaviour, optimization of cell design, and practical applications in energy storage, electronics, and industrial processes.