Chemical Effects of Current

 Definition

Current can produce or speed up chemical change, this ability of current is called chemical effect.

When current is passed through an electrolyte, it dissociates into positive and negative ions. This is called the chemical effect of current.

$\textbf{Some Terms}$

$\to$ $\textbf{Electrolysis}$

Electrolysis is a process by which electric current is passed through a substance to effect a chemical change.

$\to$ $\textbf{Electrolyte}$

The liquids which allow the current to pass through them and also dissociates into ions on passing current through them are called electrolytes. E.g. solutions of salts, acids and bases in water, etc.

$\to$ $\textbf{Electrodes}$

Two metal rods or plates which are partially dipped in the electrolyte for passing the current through the electrolyte.

        Anode: Connected to positive terminal of battery  

        Cathode: Connected to negative terminal of battery  

$\to$ $\textbf{Ionisation}$

The process of decomposition of a compound into its constituent ions is called ionization.

$\to$ Anions

The negatively charged ions which move towards the anode during electrolysis is called anions.

$\to$ Cations

The positively charged ions which move towards the cathode during electrolysis is called cations.

REMEMBER: $A \rightarrow A$ (i.e. anions to anode): $C \rightarrow C$ (i.e. cations to cathode)

$\to$ Chemical equivalent

The ratio of the atomic weight of an element to its valency is defined as its equivalent weight.

$\to$ $\textbf{Voltameter}$

The vessel in which the electrolysis is carried out is called a voltmeter. It is also known as an electrolytic cell.

$\textbf{Theory of Electrolysis}$

$\to$ Arrhenius explained the process of electrolysis by his theory of ionic dissociation ($1887$).

$\to$ When an electrolyte is dissolved in a liquid then some molecules of electrolyte dissociate into oppositely charged ions.

$\to$ When no current is passed through the solution the ions move randomly and the solution is electrically neutral.

$\to$ When an electric current is passed then anions and cations move towards their respective electrodes under influence of potential difference.

$\to$ On reaching electrodes the ions get discharged (becomes neutral) and appear as gas molecules or are deposited as thin layers on the electrode.

$\to$ The electrolytes conductivity is very low $10^{-6}$ times) than that of a good conductor because ions are heavier than electrons.

Common Types of Voltameters

Copper Voltameter

$\to$ It consists of $CuSO_4$ as electrolyte and two $Cu$ plates which work as electrodes.

$\to$ Reaction : $CuSO_4$ ionises in its aqueous solution as

$CuSO_4↔Cu^{++}+SO_4^—$

At the cathode:

$Cu^{++}+2e^- \rightarrow Cu$

The copper atoms are deposited on cathode.

At the anode:

$Cu\rightarrow Cu^{++}+2e^-$

$\to$ $Cu$ is lost from anode and deposited on cathode.

$\to$ The $Cu^{++}$ and $SO_4^{--}$ ions carry current from anode to cathode in the electrolyte. In external circuit the current is due to electrons.

$\to$ The concentration of $CuSO_4$ remains constant.

Silver Voltameter

$\to$ It consists of $AgNO_3$ as electrolyte and two $Ag$ plates which work as electrodes.

$\to$ Reaction : $AgNO_3$ ionises in its aqueous solution as

$AgNO_3↔Ag^+ + NO_3^+$

At the cathode:

$Ag^+ + e^- \rightarrow Ag$

	The silver atoms are deposited on cathode.

At the anode:

$Ag \rightarrow Ag^+ + e^-$

$\to$ Ag is lost from anode and deposited on cathode.

$\to$ The $Ag^+$ and $NO_3^-$ ions carry current from anode to cathode in the electrolyte. In external circuit the current is due to electrons.

$\to$ The concentration of $AgNO_3$remains constant.

$\textbf{Important Points}$

$\to$ Back emf(polarization) : The emf set up in water voltameter which opposes the external dc supply to the voltameter.

$\to$ Electrolysis is possible for dc and low frequency AC as at high frequency due to inertia ions cannot follow frequency of ac.

$\to$ In electrolysis electrical energy is converted to chemical energy.

$\to$ In electrolyte the current is due to directed motion of ions. Current due to positive and negative ions are not equal due to different mobilities.

$\to$ Insoluble electrode voltameters mass of cathode increases while that of anode decreases and concentration of electrolyte remains constant.

$\to$ The conductivity of electrolytes increases with rise in temperature.

$\to$ Mercury is a liquid which conducts electricity but does not dissociate into ions.

Faraday’s Law of Electrolysis

Published by Michael Faraday in $1834$ .

Gives the quantitative (mathematical) relationships that describe the above electrolysis.

$\textbf{First Law of Faraday}$

$\to$ It states that the mass of substance deposited or liberated at the electrode during electrolysis is directly proportional to the quantity of electricity (total charge) passed through the electrolyte.

i.e. m ∝q

or, $m= Zq=ZIt$ where,

Z is the electrochemical equivalent (ECE) of substance.

$\to$ If $q =1$ coulomb, then we have m = z \times 1 or $z = m$ . Hence, the electrochemical equivalent of substance may be defined as the mass of its substance deposited or liberated at the electrode, when one coulomb of charge passes through the electrolyte.

$\to$ S.I. unit ECE is $kilogram coulomb^{-1} (kg-C^{–1})$ but generally expressed in $ram coulomb^{-1} (g-C^{–1})$.

$\to$ Dimension of ECE is $M^1 A^{-1} T^{-1}$.

$\textbf{Second Law of Faraday}$

$\to$ If the same quantity of electricity is passed through different electrolytes, masses of the substance deposited at the respective cathodes are directly proportional to their chemical equivalents.

i.e. $m ∝E$

or,$\dfrac{m}{E}=constant$

$\textbf{Important Points}$

$\to$ $E=FZ$

Where F is faraday’s constant. Faraday is charge of $1 g$ equivalent ions.

If $p$ is the valency and $N$ is Avogadro’s number then,

$F=Charge in an ion \times no.of ions in 1g equivalent=(p \times) \times (\dfrac{N}{p}=Ne$.

or,$F=6.023×10^23×1.6 \tiimes10^{-19} ≈ 96500 Cmol^{-1}$

NOTE: $96500C$ charge is required to liberate or deposit $1.008g$ of hydrogen or $31.5g$ of Cu or $108g$ of silver on cathode during electrolysis.

$\to$ In general Faraday’s law can be written as $m=Zq=\dfrac{E}{q} Q=E(\dfrac{Q}{F})=\dfrac{M}{p}.\dfrac{Q}{F}=\dfrac{atomic mass}{valency} \times charge in faraday$

Applications of Electrolysis

Electroplating

The process of depositing a thin and uniform layer of metal on any conducting surface is electroplating. The articles to be electroplated are made cathode and metal to be deposited is made anode.

A soluble salt of anodic metal is used as electrolyte.

NOTE : If ρ is the density of the material deposit and A is the area of deposition, then thickness(d) of the layer of the material deposited in the electroplating process is

$d=\dfrac{m}{ρA}=\dfrac{zIt}{ρA}$

Purification of metals

Electrolysis is used in the refining of metals like copper, zinc, thin etc. The anode is made of impure metal and the cathode is made of pure metal. On passing the current, the pure metal is deposited on the cathode.

Manufacture of

Nonmetals like hydrogen, oxygen, chlorine, etc. Chemicals like $NaOH$,$Na_2CO_3$,$KClO_3$,$KMnO_4$.

Precise measure of charge or current.

Determination of equivalent weight and atomic weight.

For construction of electrolytic capacitors.

Used to cure rheumatism.

ElectroChemical Cell

An electrochemical cell is a device that can generate electrical energy from the chemical reactions occurring in it or use the electrical energy supplied to it to facilitate chemical reactions in it.

The cells which are capable of generating an electric current from the chemical reaction are also called Galvanic or Voltaic cells which are opposite of electrolytic cells.

A diagram detailing the different parts of an electrochemical cell is provided below.

Here, remember the LOAN rule.

L = left

O = oxidation

A = anode

N = negative, which means left-hand side represents anode (this is general conversion) which oxidizes and is negative and opposite for right, reduction , cathode, positive.

The total amount of energy that can be provided by this cell is limited and depends upon the amount of reactants. Electrochemical cells are of two types.

Primary Cell

$\to$ Primary cells are basically use-and-throw galvanic cells.

$\to$ Chemical reaction is irreversible.

$\to$ This cell cannot be recharged but the chemicals have to be replaced after a long use.

$\to$ The commonly used primary cells are

Cell	Positive electrode	Negative electrode	Electrolyte	EMF	Main reaction
Voltaic cell	Cu rod	Zn rod	$dil H_2 SO_4$	$1.08V$	$H_2 SO_4\rightarrow 2H^+ + S4—$ || $Zn\rightarrow Zn^{++} + 2e^-$
Daniel cell	Cu rod	Zn rod	$dil H_2 SO_4$	$1.1V$	$Zn+H_2 SO_4\rightarrow ZnSO_4+2H^++2e^-$ || $2H^++CuSO_4 \rightarrow H_2 SO_4+Cu^{++}$
Leclanche cell	Carbon rod	Zn rod	$NH_4 Cl$	$1.45V$	$Zn+2NH_4 Cl\rightarrow 2NH_3+ZnCl_2+2H^++2e^-$ || $2H^++2MnO_2\rightarrow Mn_2 O_3+H_2 O$
Dry cell	Carbon rod and brass cap	Zn vessel	Paste of $NH_4 Cl$ and saw dust	$1.5V$	Similar to leclanche cell

Secondary Cell

$\to$ A secondary cell is that cell in which the electrical energy is first stored up as a chemical energy and when the current is taken from the cell, the chemical energy is reconverted into electrical energy.

$\to$ Chemical reactions are reversible.

$\to$ The cell can function as a Galvanic cell as well as an Electrolytic cell.

$\to$ They are also called storage cells or accumulators.

$\to$ The commonly used secondary cells are

In charged	Lead accumulator	Alkali accumulator
Positive electrode	Perforated lead plates coated with $PbO_2$	Perforated steel plates coated with $Ni(OH)_4$
Negative electrode	Perforated lead plates coated with pure lead	Perforated steel plates coated with $Fe$

$\textbf{During charging - Lead accumulator}$

Chemical reaction

At cathode: $PbSO_4+2H^++2e^-\rightarrow Pb+H_2 SO_4$

At anode: $PbSO_4+SO_4^{--}+2H_2 O-2e^-\rightarrow PbO_2+2H_2 SO_4$

• Specific gravity of $H_2 SO_4$ increases and when specific gravity becomes $1.25$ the cell is fully charged.

• EMF of cell: When cell is full charged then,$E=2.2V$.

$\textbf{During discharging -Lead accumulator}$

• Chemical reaction

At cathode :

$Pb+SO_4^{--}-2e^-\rightarrow PbSO_4$

At anode:

$PbO_4+2H^+-2e^-+H_2 SO_4\rightarrow PbSO_2+H_2 O$

• Specific gravity of $H_2 SO_4$ decreases and when specific gravity falls below $1.18$. The cell requires recharging.

• EMF of cell: When emf of cell falls below 1.9V The cell requires recharging.

$\textbf{During Charging - Alkali accumulator}$

• Chemical reaction

At cathode:

$Fe(OH)_2+2OH^--2e^-\rightarrow Ni(OH)_4$

At anode :

$Fe(OH)_2+2K^+\rightarrow Fe+2KOH$

• EMF of cell: When cell is fully charge then, $E=1.35 V$.|

$\textbf{During Discharging - Alkali accumulator}$

• Chemical reaction

At cathode:

$Fe+2OH^-\rightarrow Fe(OH)_2$

At anode :

$Ni(OH)_4+2K^++2e^-\rightarrow Ni(OH)_2+2KOH$

• EMF of cell : When emf of cell falls below 1.1 V it requires charging. |