electrochemistry

• A primary cell is an electrochemical cell that is non-rechargeable and is designed to be used until its electrochemical reactions are no longer feasible.
• Electrochemistry is the study of production of electricity from energy released during spontaneous chemical reactions and the use of electrical energy to bring about non-spontaneous chemical transformations.
• The subject of electrochemistry is of importance both for theoretical and practical considerations.
• The ferrous ions are further oxidised by atmospheric oxygen to ferric ions which come out as rust in the form of hydrated ferric oxide (Fe 2 O 3 x H 2 O).
• The overall reaction in the electrolysis of water is: 2Fe(s) + O 2 (g) + 4H + (aq) → 2Fe 2 + (aq) + 2 H 2 O (l) (cell).
• Rusting of iron is envisaged as setting up of an electrochemical cell.
• The cathode reaction in the electrolysis of water is: O 2 (g) + 4 H + (aq) + 4 e – → 2 H 2 O (l) + 2 2 H O H O.
• A large number of metals, sodium hydroxide, chlorine, fluorine and many other chemicals are produced by electrochemical methods.
• Batteries and fuel cells convert chemical energy into electrical energy and are used on a large scale in various instruments and devices.
• The reactions carried out electrochemically can be energy efficient and less polluting.
• Electrochemistry is important for creating new technologies that are ecofriendly.
• The transmission of sensory signals through cells to brain and vice versa and communication between the cells are known to have electrochemical origin.
• Electrochemistry is a very vast and interdisciplinary subject.
• In this Unit, we will cover only some of the important elementary aspects of electrochemistry.
• The current flow for the reaction between copper and zinc was 3.17 amps.
• Predict the products of electrolysis in an aqueous solution of AgNO3 with platinum electrodes.
• Predict the products of electrolysis in an aqueous solution of AgNO3 with silver electrodes.
• Using the standard electrode potentials given in Table 3.1, predict if the reaction between Br2 (aq) and Fe2+ (aq) is feasible.
• Using the standard electrode potentials given in Table 3.1, predict if the reaction between Fe3+ (aq) and Br– (aq) is feasible.
• Predict the products of electrolysis in a dilute solution of H2SO4 with platinum electrodes.
• Using the standard electrode potentials given in Table 3.1, predict if the reaction between Ag(s) and Fe3+ (aq) is feasible.
• Predict the products of electrolysis in an aqueous solution of CuCl2 with platinum electrodes.
• Using the standard electrode potentials given in Table 3.1, predict if the reaction between Ag+ (aq) and Cu(s) is feasible.
• Using the standard electrode potentials given in Table 3.1, predict if the reaction between Fe3+ (aq) and I– (aq) is feasible.
• After studying this Unit, you will be able to describe an electrochemical cell and differentiate between galvanic and electrolytic cells.
• You will be able to apply Nernst equation for calculating the emf of galvanic cell and define standard potential of the cell.
• You will be able to derive relation between standard potential of the cell, Gibbs energy of cell reaction and its equilibrium constant.
• The electrical resistance of a column of 0.05 mol L –1 NaOH solution of diameter 1 cm and length 50 cm is 5.55 × 103 Ω.
• Conductivity of 0.05 mol L –1 NaOH solution = κ = ρ × Ω × l = 87.135 Ω cm.
• Molar conductivity of a solution at a given concentration is the conductance of the volume V of solution containing one mole of electrolyte kept between two electrodes with area of cross section A and distance of unit length.
• The conductivity of a solution at any given concentration is the conductance of one unit volume of solution kept between two platinum electrodes with unit area of cross section and at a distance of unit length.
• Both conductivity and molar conductivity change with the concentration of the electrolyte.
• Conductivity always decreases with decrease in concentration both, for weak and strong electrolytes.
• Alternatively, κ = –1 1.29 cm 520 Ω = 0.248 × 10 –2 S cm –1.
• The number of ions per unit volume that carry the current in a solution decreases on dilution.
• The cell constant is given by the equation: Cell constant = G* = conductivity × resistance = 1.29 S/m × 100 Ω = 129 m –1 = 1.29 cm –1.
• Molar conductivity = κ = –1 1.29 cm 520 Ω = 0.248 × 10 –2 S cm –1.
• Molar conductivity of 0.05 mol L –1 NaOH solution = κ × 1000 c κ cm3 L –1 = 229.6 × 10 –4 S cm2 mol –1.
• Molar conductivity increases with decrease in concentration.
• Conductivity of 0.02 mol L –1 KCl solution = cell constant / resistance = G R = –1 129 m 520 Ω = 0.248 S m –1.