ELECTROCHEMISTRY
Syllabus:-
Ø Electrolytic
cells & Galvanic cells
Ø Redox
reaction, EMF of cell, standard electrode potential
Ø Nernst
equation and its application to Chemical cell
Ø Conductance
in electrolytic solutions, Conductance, Conductivity, Cell Constant
Ø Specific
& Molar Conductivity, Variation of Conductivity and molar Conductivity with Concentration, Limiting Molar Conductivity
Ø Kohlrausch’s
Law, Applications of Kohlrausch Law
Ø Electrolysis
& Laws of Electrolysis
Ø Chemistry
of Cells and Batteries, dry cell, Lead Storage Battery, Mercury Cell, fuel cell
Chemistry of Corrosion.
Electrochemistry: It is a branch of
chemistry that deals with the relationship between chemical energy and
electrical energy and their inter conversions.
Redox Reactions: Oxidation
is the process which involves loss of electrons & reduction is a process in
which it involves gain of electrons. The reactions which involve both that
reaction simultaneously are called as redox reaction.
Electrochemical
Cells: These are devices that convert
chemical energy of some redox reactions to electrical energy. They are also
called Galvanic cells or Voltaic cells. An example for Galvanic cell is
Daniel cell.
It is constructed by dipping a Zn rod in
ZnSO4 solution and a Cu rod in CuSO4 solution. The two
solutions are connected externally by a metallic wire through
a voltmeter and a switch and internally by a salt bridge.
A salt bridge is a U-tube containing
an inert electrolyte like NaNO3 or KNO3 in
a gelly like substance.
The functions of a salt
bridge are:
1. To complete the electrical circuit
2. To
maintain the electrical neutrality in the two half cells.
(i) Cu2+ + 2 e- → Cu(s) (reduction half
reaction)
(ii) Zn(s) → Zn2+ + 2 e- (oxidation half
reaction)
These reactions occur
in two different portions of the Daniel cell. The reduction half reaction
occurs on the copper electrode while the oxidation half reaction occurs on the
zinc electrode. These two portions of the cell are also called half-cells
or redox couples. The copper electrode may be called the reduction half-cell and
the zinc electrode, the oxidation half-cell.
Electrode Potential: This
tendency of a metal to lose or gain electron when it is in contact with its own
solution is called electrode
potential.
Standard electrode potential (E0): The electrode potential measured at standard conditions.ie at
298K,1 atm pressure and at 1 molar concentration.
Standard hydrogen electrode (SHE): The reference electrode used to measure single
electrode potential. Its potential is assumed to be zero. It consists of a platinum
wire dipped in HCl of 1 molar concentration. Hydrogen gas at 1 atm. is passed
through the solution. The electrode can be represented as Pt, H2 /H+(1M).
Nernst Equation For A Cell Reaction:-
aA + bB → cC + dD
E = E0cell
–
2.303 RT * log [C]c
[D]d
nF [A]a [B]b
Where E0 is the standard
electrode potential,
R
is the gas constant (8.314 JK–1 mol–1)
F
is Faraday constant (96500 C mol–1)
T
is temperature in Kelvin.
Nernst equation can be written as:
On Putting the above values
Nernst equations can be written as:
Equilibrium Constant
from Nernst Equation:
Electrochemical Cell and Gibbs Energy of the
Reaction:
Electrical work done in one second is equal to electrical
potential multiplied by total charge passed. The reversible work done by a
galvanic cell is equal to decrease in its Gibbs energy and therefore, if the
emf of the cell is E and nF is the amount of charge passed and ∆G is the Gibbs energy of the reaction,
then
∆G = – nFEcell
If the concentration of all the
reacting species is unity, then Ecell = E0cell
and we have
∆G0 = – nFE0 cell
Thus, from the measurement of E0cell,
we can calculate the standard Gibbs energy of the reaction.
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Electrochemistry - Part 3
Electrochemistry - Part 2
Electrochemistry - Part 1
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