ELECTROCHEMISTRY
9.1 Electrolytes
Electrolyte may be defined as that substance which in the form of its solution or in its fused state conducts electricity and simultaneously undergoes chemical decomposition.
Strong electrolytes
Those electrolyte which possess a high value of equivalent conductance even at high concentrations (i.e., when the dilution is not very large) and whose equivalent conductance increases gradually with dilution and then becomes constant are called strong electrolytes. Examples are mineral acids (HCl, H2SO4, HNO3), alkalies (NaOH, KOH), alkaline earth hydroxides [e.g., Ca(OH)2, Ba (OH2)], and salts (e.g, NaCl, KCl, etc.).
or
Thus it follows that the Specific resistance of a conductor is the resistance in ohms which one centimeter cube of it offers to the passage of electricity.
9.4 Specific Conductivity
9.5 Equivalent Conductance
Relationship between equivalent and specific conductivities in terms of concentration Suppose the solution has a concentration of c gram equivalent per litre. Then the volume V containing 1 gram equivalent of the electrolyte will be 1/c litre, i.e.
Molar conductance of a solution at a dilution V is defined as the conductance of all the ions produced from one mole of the electrolyte dissolved in V. c.c. of the solution when the electrodes are 1 cm apart and the area of the electrodes is so large that whole of the solution is contained between them. It is denoted by m. Like equivalent conductance, it is related to electrolytic conductivity (specific conductance) as below
9.7 Relation between equivalent conductivity and molar conductivity
9.8 Electrolysis
(b) The anions move towards anode and on reaching the anode they lose electrons and converted to neutral atoms.
(At anode) B– —® B + e– (Oxidation)
where ERP = Reduction potential of
cation and 
Standard reduction potential of
cation.
Thus, it is possible that a cation (A+) with lower standard reduction potential getting discharged in preference to cation (B+) having higher standard reduction potential because their concentrations might be such that the reduction potential of A+ is higher than that of B+.
W µ Q
where, W = Mass of ions liberated in gm, Q = Quantity of electricity passed in coulombs.
Thus electrochemical equivalent may be defined as the mass of the ion deposited by passing a current of one ampere for one second (i.e. by passing coulomb of electricity). Its unit is gram per coulomb.
Faraday’s Second Law
Where F is again a proportionality constant and has been found to be 96540 coulombs. It is called faraday.
Thus E = 96540 × Z
An electrochemical cell is a single arrangement of two electrodes and an electrolyte for producing an electric current due to chemical action within the cell, or for producing chemical action due to passage of electricity. Thus electrochemical cells may be used for two purposes namely.
(i) to convert chemical energy into electrical energy.
(i) It prevents voltage drops, i.e. prevents junction potential.
(ii) It allows flow of current by completing the circuit, i.e. migration of anion from anode to cathode half cell.
9.12 Types of Electrodes
In an electrochemical cell, there are two electrodes, positive and negative. Each electrode constitutes a half cell or a single electrode. Although a number of electrodes are possible but the more important of these electrodes are grouped into the following types:
(v) Oxidation-reduction or redox electrodes.
We shall denote these electrodes as right hand ones, i.e., the electrode reactions would correspond to reductions.
and depends on the activity of the metal cation in the solution.
and the electrode potential is given by
The electrode potential is given by
(v) Oxidation-reduction or redox electrodes: These are electrodes in which the emf arises from the presence of ions of a substance in two different oxidation states. These electrodes are set up by dipping an inert metal like gold or platinum into a solution containing ions in two different oxidation states of the substance. For example, a platinum wire immersed in a solution of ferrous and ferric ions or stannous and stannic ions constitutes a redox electrode. These electrodes are represented as
9.13 Electrode Potential
The potential of any electrode is the potential difference between it and the electrolyte surrounding the electrode. The electrode potential depends upon the nature of the metal, concentration of the metallic ions in solution and the temperature of the solution. When the ions are at unit activity and the temperature is 25°C (298K), the potential difference is called a standard electrode potential (SEO). The potential of a single electrode cannot be determined but the potential difference between the two electrodes can be accurately measured. Normal hydrogen electrode (NHE), also called standard hydrogen electrode (SHE), is the standard reference electrode 1. Thus the standard electrode potential of a metal may be defined as the potential difference in volts developed in a cell consisting of two electrodes, the pure metal in contact with a molar solution of one of its ions and the normal hydrogen electrode.
is
(ii) Temperature T
If T increases at constant [Mn+], Eox decreases.
9.14 Cell Potentials (EMF) from Electrode Potentials
As mentioned at start, every galvanic cell is made up of two half cells: the oxidation half-cell (anode or negative electrode) and the reduction half-cell (cathode or positive electrode). The electrode potentials of the two half-cells of a cell differ. Due to this difference in potential, an electric current flows from the electrode at higher potential to the electrode at a lower potential. The difference in potential of the two half-cells of a cell is known as electromotive force (emf) of the cell or cell potential. It is measured directly with a voltmeter and is given in volts. Its value can also be determined by knowing the standard half-cell potentials of the two half-cells. Cell potential (or EMF of the cell) can be determined from electrode potentials in the following three ways.
(iii) Ecell = Oxidation potential of anode + Reduction potential of cathode
If cell potential (Ecell) is positive, the reaction is spontaneous. Further, the more positive the value of Ecell is, the faster the reaction is. In case the value of cell potential is negative, the reaction will not take place spontaneously as written. However, the reaction will take place.
WMax = nFEcell
Where n = Number of moles of electrons transferred through the wire
Therefore, from above equations we get
DG = – n F E0cell
From Gibbs-Helmholtz equation we have
Also DG = DH – TDS
Hence from equations, we obtain
From equation, the enthalpy change for cell reaction can be determined from the measurement of cell emf and the temperature coefficient of the emf.