13.1 Faraday's first law

2.1 Difference between metallic conduction and electrolytic conduction

3.1 Difference Between Electrolytic Cell and Galvanic Cell

5.1 Salt Bridge

9.1 Characteristics of electrochemical series
9.2 Applications of electrochemical series

10.1 Reducing power of Metals
10.2 Oxidising Nature of Non-metals

13.1 Faraday's first law
13.2 Farady's Second Law

15.1 Condition of Equilibrium
15.2 Spontaneity of the Reaction

17.1 Primary Battery or Voltaic Cells (Dry Cells):
17.2 Secondary Cells

This law states that the mass of substance liberated or deposited at the electrode depends on the quantity of electric charge passed through it or mass deposited or liberated at the electrode is directly proportional to the amount of electricity passed through it.

Mathematically,$$W\propto Q$$

where

$W$ is the mass of the ions deposited and

$Q$ is the quantity of electricity in coulombs, which is equal to product of current (in amperes) and time (in seconds)

$$W = Z \times I \times t$$

$Z$ is the constant, known as electrochemical equivalent (ECE) of the ion deposited.

When $1$ ampere current is passed for $1$ second then $W=Z$ so we can define $Z$ (called electrochemical equivalent) as mass of ion deposited by passing current of $1$ ampere for$1$ second (i.e., by passing $1$ coulomb of charge). Its unit is $Kg{C^{ - 1}}$

$1\ F$ of charge $=$ charge on one mole of electron

$$\begin{equation} \begin{aligned} = {N_A} \times e \\ = 6.023 \times {10^{23}} \times 1.602 \times {10^{ - 19}} \\ = 96514.8C \approx 96500C \\\end{aligned} \end{equation} $$

$=$ the charge which discharges equivalent weight of ion.

so $$F = {N_A} \times e$$

therefore, $$1\ C\ of\ charge = \frac{E}{{96500}}g\ \ of\ ion\ =\ Z$$

Now substitute the value of $Z$ in the reaction

$$\begin{equation} \begin{aligned} W = ZIt = \frac{{EIt}}{{96500}} \\ \frac{W}{E} = \frac{{It}}{{9650}} = \frac{Q}{{96500}} = \frac{Q}{F} \\\end{aligned} \end{equation} $$