To learn the laws of thermodynamics in understanding homogenous and heterogenous equilibria and the theories and concept of electrochemistry in redox system and cells.
Gibbs and Helmholtz functions, Gibbs function (G) and Helmholtz function (A) as thermodynamic quantities, A & G as criteria for thermodynamic equilibrium and spontaneity, their advantages over entropy change, variation of G & A with P, V & T.
Equilibrium constant and free energy, thermodynamic derivation of law of mass action. Le Chatelier’s principle,reaction isotherm and reaction isochore- Clapeyron equation and Clausius -Clapeyron equation, applications
Introduction to phase, component and degree of freedom, derivation of Gibbs phase rule; phase equilibria of one component system-water, CO2 and sulphur system. Phase equilibria of two component system-solid-liquid equilibria, simple eutectic – Bi-Cd, Pb-Ag systems, desilverisation of lead.
Solid solutions: compound formation with congruent melting point (Mg-Zn) and incongruent melting point (NaCl-H2O), (FeCl3 – H2O) and (CuSO4 – H2O) system; freezing mixtures (acetone – dry ice).
Types of reversible electrodes – gas-metal ion, metal-metal ion, metal-insoluble salt-anion and redox electrodes; electrode reactions, Nernst equation, EMF of a cell and its measurements, computation of cell EMF, calculation of thermodynamic quantities of cell reactions (G, H & K),derivation of cell E.M.F. and single electrode potential; standard hydrogen electrode- reference electrodes, standard electrode potential, sign conventions, electrochemical series and its significance.
Electrolytic and Galvanic cells: reversible and irreversible cells, conventional representation of electrochemical cells. Concentration cell with and without transport, liquid Junction potential, applications of concentration cell - valency of ions, solubility product, activity coefficient, potentiometric titrations.
Definition of pH and pKa, determination of pH using hydrogen, quinhydrone and glass electrodes and by potentiometric method.