24 Electrochemistry

24.1 Electrolysis

  1. Predicting substances liberated during electrolysis:
    1. In molten electrolytes: The substance liberated at the cathode will be the metal, while at the anode, it will be the non-metal or oxygen.
    2. In aqueous electrolytes: The substance liberated at the cathode depends on the position of the metal in the reactivity series. The substance liberated at the anode depends on the concentration of ions present and their standard electrode potentials.
  2. Relationship between Faraday constant and charge on the electron:

F = Le

The Faraday constant (F) represents the charge of one mole of electrons, and it is related to the charge on a single electron (e) and the Avogadro constant (L).

  1. Calculations during electrolysis:
    1. Quantity of charge passed:

Q = It

where Q is the quantity of charge (in coulombs), I is the current (in amperes), and t is the time (in seconds).

  1. Mass and/or volume of substance liberated: By using stoichiometry and the molar mass or

volume of the substance, you can calculate the mass or volume of the substance liberated during electrolysis.

  1. Determination of Avogadro constant by an electrolytic method: The electrolytic method involves depositing a known mass of metal during electrolysis and calculating the number of moles. By knowing the Avogadro constant (L), you can determine its value.

24.2 Standard Electrode Potentials E⦵, Standard Cell Potentials E⦵cell, and the Nernst Equation

  1. Definitions:
    1. Standard electrode potential:

The potential difference between an electrode and a standard hydrogen electrode (SHE) under standard conditions.

  1. Standard cell potential:

The potential difference between the two half-cells of an electrochemical cell under standard conditions.

  1. Standard hydrogen electrode (SHE):

The SHE is used as a reference electrode with an assigned electrode potential of 0V. It consists of a platinum electrode immersed in an acidic solution with hydrogen gas at 1 atm.

  1. Methods for measuring standard electrode potentials:

(a) For metals or non-metals in contact with their ions in aqueous solution, the potential is measured using a standard hydrogen electrode and a voltmeter.

(b) For ions of the same element in different oxidation states, the potential is measured by constructing half-cells for each oxidation state and comparing them with the SHE.

  1. Calculating standard cell potential: The standard cell potential (E⦵cell) is obtained by subtracting the standard electrode potential of the anode from that of the cathode when the two half-cells are combined.
  2. Use of standard cell potentials:

(a) Polarity and direction of electron flow: The electrode with a more positive standard electrode potential will be the cathode, attracting electrons and undergoing reduction. The electrode with a less positive (or more negative) potential will be the anode, undergoing oxidation. Electrons flow from anode to cathode in the external circuit.

(b) Feasibility of a reaction: If the standard cell potential is positive, the reaction is thermodynamically feasible. If it is negative, the reaction is not feasible under standard conditions.

  1. Relative reactivity of elements, compounds, and ions: By comparing their standard electrode potentials, you can determine the relative reactivity of species as oxidizing agents or reducing agents. Species with higher electrode potentials are better oxidizing agents, while species with lower electrode potentials are better reducing agents.
  2. Constructing redox equations using half-equations: Half-equations represent the oxidation and reduction processes occurring at each electrode in an electrochemical cell. By combining appropriate half-equations, you can construct the overall redox equation.
  3. Variation of electrode potential with ion concentrations: Increasing the concentration of oxidized species or reducing species will affect the electrode potential. Higher concentrations of oxidized species or lower concentrations of reduced species will result in more positive electrode potentials.
  4. Nernst equation: The Nernst equation relates the electrode potential (E) to the concentrations of the oxidized and reduced species in an electrochemical cell. It is given by

E = E⦵ + (0.059/z) * log ([oxidized species]/ [reduced species])

where z is the stoichiometric coefficient of electrons in the balanced half-equation.

  1. Gibbs free energy and cell potential: The relationship between the standard cell potential (E⦵cell) and the Gibbs free energy change (ΔG⦵) is given by

ΔG⦵ = -nE⦵cellF

where n is the number of moles of electrons transferred and F is the Faraday constant.