1. Electrochemical Cells:

For any electrode $\to$ Oxidation potential $= -$ Reduction potential.

$E_{c e l l} =$ Reduction potential of cathode $-$ Reduction potential of anode.

$E_{c e l l} =$ Reduction potential of cathode $+$ Oxidation potential of anode.

$E_{c e l l}$ is always a positive quantity for a spontaneous reaction; Anode will be electrode of low reduction potential and Cathode will be of high reduction potential.

${E^{\circ}}_{c e l l} =$ SRP of cathode $-$ SRP of anode. (SRP = Standard Reduction Potential)

The greater is the SRP value, the greater will be oxidising power.

2. Different types of Electrodes:

(i) Metal-metal ion electrode $(M^{+ n} | M)$

$M^{n +} + {n e}^{-} \to M (s)$

(ii) Gas-ion Electrode $P t |H_{2} (P a t m)| H^{+} (X M)$

As reduction electrode $, H^{+} (a q) + e^{-} \to \frac{1}{2} H_{2} (P a t m), E = E^{°} - 0.0591 l o g \frac{P_{H_{2}}^{\frac{1}{2}}}{[H^{+}]}$

(iii) Oxidation-reduction electrode $P t | {F e}^{2 +}, {F e}^{3 +}$

As reduction electrode, ${F e}^{3 +} + e^{-} ⟶ {F e}^{2 +} E = E^{\circ} - 0.0591 \log \frac{[{F e}^{2 +}]}{[{F e}^{3 +}]}$

(iv) Metal-metal insoluble salt electrode $e . g ., A g | A g C l, {C l}^{-}$

As reduction electrode, $A g C l (s) + e^{-} ⟶ A g (s) + {C l}^{-}$ .

$E_{{C l}^{-} |A g C l| A g} = E_{{C l}^{-} |A g C l| A g}^{0} - 0.0591 \log [{C l}^{-}]$

3. Gibbs Free Energy Change:

$Δ G = - n F E_{c e l l}$

4. Nernst Equation:

(Effect of concentration and temperature on emf of cell)

$Δ G = Δ G^{°} + R T l n Q$ (where $Q$ is reaction quotient)

$Δ G^{°} = - R T l n K_{e q}$

$E_{c e l l} = {E^{\circ}}_{c e l l} - \frac{2.303 R T}{n F} l o g Q$

$E_{c e l l} = {E^{\circ}}_{c e l l} - \frac{0.0591}{n} l o g Q [A t 298 K]$

At chemical equilibrium: $Δ G = 0$ ; $E_{c e l l} = 0$

$l o g K_{e q} = \frac{n E_{c e l l}^{0}}{0.0591}$

$E_{c e l l}^{0} = \frac{0.0591}{n} l o g K_{e q}$

5. Concentration Cell:

A cell in which both the electrodes are made up of the same material.

For all concentration cell,

${E^{\circ}}_{c e l l} = 0$

(i) Electrolyte Concentration Cell:

$e . g ., Z n (s) | {Z n}^{2 +} (c_{1}) ‖ {Z n}^{2 +} (c_{2}) | Z n (s) E = \frac{0.0591}{2} l o g \frac{C_{2}}{C_{1}}$

(ii) Electrode Concentration Cell:

e.g., $P t, H_{2} (P_{1} a t m) | H^{+} (1 M) | H_{2} (P_{2} a t m) | P t E = \frac{0.0591}{2} l o g (\frac{P_{1}}{P_{2}})$

6. Calculation of different Thermodynamics Function of Cell Reaction:

$Δ G = - n F E_{c e l l}$

$∆ S = - {[\frac{d G}{d T}]}_{P}$ (At constant pressure)

$Δ S = - {[\frac{d (Δ G)}{d T}]}_{P} = n F {(\frac{d}{d t} (E_{c e l l}))}_{P}$

${[\frac{\partial E}{\partial T}]}_{P} =$ Temperature coefficient of emf of the cell.

$Δ H = n F [T {(\frac{\partial E}{\partial T})}_{P} - E]$

$Δ C p$ of cell reaction:

$C_{P} = \frac{d H}{dT}$

$Δ C_{P} = \frac{d}{d T} (Δ H)$

$Δ C_{P} = n F T \frac{d^{2} E_{c e l l}}{d T^{2}}$

7. Electrolysis:

(i) $\frac{K^{+}, {C a}^{+ 2}, {N a}^{+}, {M g}^{+ 2}, {A l}^{+ 3}, {Z n}^{+ 2}, {F e}^{+ 2}, H^{+}, {C u}^{+ 2}, {A g}^{+}, {A u}^{+ 3}}{I n c r e a s i n g o r d e r o f d e p o s i t i o n .} >$

(ii) Similarly, the anion which is stronger reducing agent (low value of $S R P$ ) is liberated first at the Anode.

$\frac{{S O}_{4}^{2 -}, {N O}_{3}^{-}, {O H}^{-}, {C l}^{-}, {B r}^{-}, I^{-}}{I n c r e a s i n g o r d e r o f d e p o s i t i o n} >$

8. Faraday's Law of Electrolysis:

(i) First Law:

$w = Z q w = Z i t$ Where $Z =$ Electrochemical equivalent of substance.

(ii) Second Law:

$W α E or \frac{W}{E} = c o n s t a n t \frac{W_{1}}{E_{1}} = \frac{W_{2}}{E_{2}}$

$\frac{W}{E} = \frac{i \times t \times c u r r e n t e f f i c i e n c y f a c t o r}{96500}$

$C u r r e n t e f f i c i e n c y = \frac{A c t u a l m a s s deposited / produced}{T h e o r e t i c a l m a s s deposited / produced} \times 100$

9. Conductance in Electrolytic Solutions:

(i) $C o n d u c t a n c e = \frac{1}{R e s i s t a n c e}$

(ii) Specific conductance or conductivity:

(Reciprocal of specific resistance) $κ = \frac{1}{ρ} κ =$ specific conductance

specific conductance $=$ conductance $\times \frac{l}{a}$

(iii) Equivalent conductance:

$Λ_{E} = \frac{κ \times 1000}{Normality} unit : {ohm}^{- 1} {cm}^{2} {eq}^{- 1}$

(iv) Molar conductance:

$Λ_{m} = \frac{κ \times 1000}{Molarity} unit : {ohm}^{- 1} {cm}^{2} {mole}^{- 1}$

10. Kohlrausch's Law:

Variation of $λ_{e q} & λ_{M}$ of a solution with concentration:

(i) Strong electrolyte $λ_{M}^{c} = λ_{M}^{\infty} - b \sqrt{c}$

(ii) For both strong and weak electrolytes:

$λ_{\infty} = n_{+} λ_{+}^{\infty} + n_{-} λ^{\infty}$ where $λ$ is the molar conductivity.

$n_{+} =$ No. of cations obtained after dissociation per formula unit.

$n_{-} =$ No. of anions obtained after dissociation per formula unit.

11. Application of Kohlrausch’s Law:

(i) Calculation of $λ_{m}^{0}$ of weak electrolytes:

$λ_{M ({C H}_{3} C O O H)}^{o} = λ_{M ({C H}_{3} C O O N a)}^{0} + λ_{M_{(H C l)}}^{0} - λ_{M (N a C l)}^{0}$

(ii) To calculate degree of dissociation of a weak electrolyte

$α = \frac{λ_{m}^{}}{λ_{m}^{0}}; K_{e q} = \frac{c α^{2}}{(1 - α)}$

(iii) Solubility $(S)$ of sparingly soluble salt & their $K_{s p}$

$λ_{M} = λ_{M}^{o} = κ \times \frac{1000}{s o l u b i l i t y} K_{s p} = S^{2}$

(iv) Ionic mobility: It is the distance travelled by the ion per second under the potential gradient of $1$ volts per $c m$ . Its unit is ${c m}^{2} s^{- 1} v^{- 1}$ .

(v) Absolute ionic mobility:

$λ_{c}^{o} {α μ}_{c}; λ_{a}^{0} \propto μ_{a}$

$λ_{c}^{0} = F \times μ_{c}^{0}; λ_{a}^{0} = F \times μ_{a}^{0}$

$I o n i c M o b i t i t y μ = \frac{S p e e d}{P o t e n t i a l g r a d i e n t} = \frac{v}{\frac{V}{l}}$

(vi) Transport Number:

$t_{c} = [\frac{μ_{c}}{μ_{c} + μ_{a}}], t_{a} = [\frac{μ_{a}}{μ_{a} + μ_{c}}]$

Where $t_{c} =$ Transport Number of cation & $t_{a} =$ Transport Number of anions.

KCl can be used in salt bridge as the electrolyte in which of the following cells?

Important Questions on Electrochemistry

Learn and Prepare for any exam you want !

$KCl$ can be used in salt bridge as the electrolyte in which of the following cells?