Oxidation occurs at the anode
\({ Pb }_{ (s) }\rightarrow { Pb }_{ (aq) }^{ 2+ }+{ 2e }^{ - }\)
The Pb²⁺ions combine with \({ SO }_{ 4(aq) }^{2-}\) to from PbSO₄ precipitate.
\({ Pb }_{ (aq) }^{ 2+ }+{ SO }_{ 4(aq) }^{ 2- }\rightarrow { PbSO }_{ 4(s) }\)
Reduction occurs at the cathode
\({ PbO }_{ 2(s) }+{ 4H }_{ (aq) }^{ + }+{ 2e }^{ - }\rightarrow { Pb }_{ (aq) }^{ 2+ }+{ 2H }_{ 2 }O(l)\)
The Pb²⁺ ions also combine with \({ SO }_{ 4(aq) }^{2-}\) ions from sulphuric acid to form PbSO₄ precipitate.
\({ Pb }_{ (aq) }^{ 2+ }+{ SO }_{ 4(aq) }^{ -2 }\rightarrow { PbSO }_{ 4 }\)

(i) Oxidation occurs at anode:

(ii) Reduction occurs at cathode:

(iii) Overall reaction:
\({ Zn }_{ (s) }+{ HgO }_{ (s) }\rightarrow { ZnO }_{ (s) }+Hg(l)\)
(iv) Cell emf: about 1.35V.

The overall redox reaction in Leclanche cell is
\( \mathrm{Zn}_{(\mathrm{s})}+2 \mathrm{NH}_{4 \text { (aq) }}^{+}+2 \mathrm{MnO}_{2(\mathrm{~s})} \longrightarrow \mathrm{Zn}_{(\mathrm{aq})}^{2+}+\mathrm{MnO}_{2} \mathrm{O}_{3(\mathrm{~s})}+\mathrm{H}_{2} \mathrm{O}(l)+2 \mathrm{NH}_{3} \)
The ammonia produced at the cathode combines with Zn²⁺ to form a complex ion [Zn (NH₃)₄]²⁺_(aq). As the reaction proceeds the concentration of NH₃ will decrease and the aqueous NH₃ will increase which lead to the decrease in the emf of cell.

(i) Anode: Zinc container
(ii) Cathode: Graphite rod in contact with MnO₂
(iii) Electrolyte: Ammonium chloride and zinc chloride in water
(iv) Oxidation at anode:
\( { Zn }_{ (s) }\rightarrow { Zn }_{ (aq) }^{ 2+ }+2e^{ - }\)
(v) Reduction at cathode:
\({ 2NH }_{ 4(aq) }^{ + }+{ 2e }^{ - }\rightarrow { 2NH }_{ 3(aq) }+{ H }_{ 2(g) }\)

The maximum work that can be obtained from a galvanic cell is
(W_max)_cell = - nFE_cell      ...(1)
Here the (-) sign is introduced to indicate that the work is done by the system on the surroundings. According to Second Law of thermodynamics, the maximum work done by the system is equal to the change in the Gibbs free energy of the system.
i,e, W_max= ∆G  .......(2)
From (1) and (2),
∆G = - nFE_cell   ....(3)
For a spontaneous cell reactions, the ∆G should be negative. The above expression (3) indicates that E_cell should be positive to get a negative ∆G value.
When all the cell components are in their standard state, the equation becomes
∆G^o = -nFE^o_cell

The galvanic cell is represented by a cell diagram, for example, Daniel cell is represented as
\({ Zn }_{ (s) }|{ Zn }_{ (aq) }^{ 2+ }||{ Cu }_{ (aq) }^{ 2+ }|{ Cu }_{ (s) }\)
(i) A single vertical bar (|) represents a phase boundary
(ii) The double vertical bar (||) represents the salt bridge.
(iii) The anode half cell is written on the left side of the salt bridge and the cathode half cell on the right side.
(iv) The anode and cathode are written on the extreme left and extreme right, respectively.
(v) The emf of the cell is written on the right side after cell diagram.

The galvanic cell is represented as
\({ Zn }_{ (s) }|{ Zn }_{ (aq) }^{ 2+ }||{ Cu }_{ (aq) }^{ 2+ }|{ Cu }_{ (s) }\)
The anode half cell is \({ Zn }_{ (s) }|{ Zn }_{ (aq) }^{ 2+ }\)
The cathode half cell is \({ C }u_{ (Aq) }^{ 2+ }|{ Cu }_{ (s) }\)

Zinc is oxidised to Zn^{2+ i}ons and the Cu²⁺ ions are reduced to metallic copper. The half reactions are represented as below.
\( { Zn }_{ (s) }\rightarrow { Zn }_{ (aq) }^{ + }+2e^{ - }\) (oxiation)
Loss of election oxidation
\({ Cu }_{ (aq) }^{ 2+ }+{ 2e }^{ - }\rightarrow { Cu }_{ (s) }\) (reduction)
Gain of electron oxidation.

Electrochemical cells are mainly classified into the following two types.
(i) Galvanic Cell ( Voltaic cell) : It is a device in which a spontaneous chemical reaction generates an electric current i.e., it converts chemical energy into electrical energy. It is commonly known as a battery.
(ii) Electrolytic cell : It is a device in which an electric current from an external source drives a nonspontaneous reaction i.e., it converts electrical energy into chemical energy.

For a weak electrolyte, at high concentration, the plot is almost parallel to concentration axis with slight increase in conductivity as the dilution increases. When the concentration approaches zero, there is a sudden increase in the molar conductance and the curve is almost parallel to \(\Lambda \)_m axis. This is due to the fact that the dissociation of the weak electrolyte increases with the increase in dilution (Ostwald dilution law)

Oxidation half cell : A metallic zinc strip that dips into an aqueous solution of zinc sulphate taken in a beaker.
Reduction half cell : A copper strip that dips into an aqueous solution of copper sulphate taken in a beaker.

Specific conductance k = \(C(\frac{l}{a})\)
k = \(\frac{1}{R}(\frac{l}{a})\)
That is k α \(\frac{1}{R}\)

Equivalent conductance is defined as the conductance of 'V' m³ of electrolytic solution containing one gram equivalent of electrolyte in a conductivity cell in which the electrodes are one meter apart.
\(\Lambda =\frac { k({ Sm }^{ -1 })\times { 10 }^{ -3 }(\text{ gram equivalent) }^{ -1 }{ m }^{ 3 } }{ N } \)

The conductivity cell in which the electrodes are separated by 1m and having V m³ of electrolytic solution which contains 1 mole of electrolyte. The conductance of such a system is called the molar conductance (\({ \Lambda }_{ m }\))
\({ \Lambda }_{ m }=\frac { k({ Sm }^{ -1 })\times { 10 }^{ -3 } }{ M } { mol }^{ -1 }{ m }^{ 3 }\)

(i) Specific resistance is denoted by ρ(rho).
(ii) Specific resistance depends on the nature of the electrolyte.
(iii) Reciprocal is specific resistance \(\frac{1}{ρ}\) is specific conductance or conductivity denoted by K (Kappa).
(iv) ohm metre.