Electrochemistry - SS2 Chemistry Past Questions and Answers - page 2

Electrochemical cells are devices that convert chemical energy into electrical energy through redox reactions. They consist of two half-cells: an oxidation half-cell and a reduction half-cell, connected by a conductive pathway. Each half-cell contains an electrode immersed in an electrolyte solution.

In the oxidation half-cell, oxidation occurs as a species loses electrons and forms cations. The electrode in this half-cell is called the anode. The anode releases electrons into the external circuit, allowing the oxidation reaction to proceed.

In the reduction half-cell, reduction occurs as a species gains electrons and forms anions. The electrode in this half-cell is called the cathode. The cathode attracts electrons from the external circuit to facilitate the reduction reaction.

The oxidation and reduction reactions in the respective half-cells are connected by the flow of electrons through the external circuit. Electrons released from the anode travel through the external circuit to the cathode. This electron flow creates an electrical current that can be utilised for various applications.

Meanwhile, in the electrolyte solutions, ions flow to balance the charge as the oxidation and reduction reactions occur. The anions migrate towards the anode, and the cations migrate towards the cathode. This movement of ions within the electrolyte maintains charge neutrality.

Cell potential, also known as electromotive force (EMF), is a measure of the driving force behind a redox reaction in an electrochemical cell. It indicates the voltage or electrical potential difference between the electrodes of the cell. Cell potential is denoted by the symbol Ecell and is measured in volts (V).

The cell potential determines the spontaneity of a redox reaction. If the cell potential is positive (Ecell > 0), the reaction is spontaneous and proceeds in the forward direction. Conversely, if the cell potential is negative (Ecell < 0), the reaction is nonspontaneous, and an external energy source is required for the reaction to proceed.

The cell potential is calculated using the Nernst equation:

Ecell = E°cell - (RT/nF) x ln(Q)

where E°cell is the standard cell potential, R is the gas constant, T is the temperature in Kelvin, n is the number of moles of electrons transferred, F is the Faraday constant, and Q is the reaction quotient.

The standard electrode potential (E°) is the cell potential when all reactants and products are in their standard states (usually at 298 K, 1 atm pressure, and 1 M concentration). It is a measure of the tendency of a half-cell reaction to occur. The standard electrode potentials are tabulated and provide a reference point for comparing different redox reactions.

To calculate the cell potential using the standard electrode potentials, the following equation is used:

E°cell = E°cathode - E°anode

The standard electrode potentials allow for the determination of the cell potential without the need for actual cell measurements. They provide valuable information about the relative strengths of oxidising and reducing agents and allow predictions about the feasibility and direction of redox reactions.