Electrochemistry

Electrolytic Cells

An external source of voltage is used to bring a chemical reaction in an electrolytic cell. One of the simplest electrolytic cell is made up of two copper strips dipped in aqueous solution of copper sulphate. When DC voltage is applied to the two electrodes then following reactions take place:

At Cathode: `Cu^(2+)(aq)+2e^(-)→Cu(s)`

At Anode: `Cu(s)→Cu^(2+)+2e^(-)`

Here, copper is dissolved (oxidized) at anode and deposited (reduced) at cathode. This is the basis of producing pure copper from impure copper. Large scale production of many metals is carried out using this process.

Faraday’s Law of Electrolysis:

First Law: The amount of chemical reaction occuring at any electrode during electrolysis is proportional to the quantity of electricity passed through the electrolyte.

Second Law: The amounts of different substances liberated by the same quantity of electricity passed through the electrolytic solution are proportional to their chemical equivalent weights (Atmoic Mass of Metal + Number of electrons required to reduce the cation).

The amount of electric charge required for oxidation or reduction depends on stoichiometry of the electrode reaction. Let us take following example:

`Ag^(+)(aq)+e^(-)→Ag(s)`

In this reaction, one mole of electron is required for reduction of one mole of silver ions.

We know that charge on one electron `=1.6021xx10^(-19)` C

Hence, charge on one mole of electrons

`=N_A\xx1.6021xx10^(-19) C``=6.02xx10^23 mo\l^(-1)xx1.6021xx10^(-19)`

Or, `C=96487` C mol^-1

This quantity is called Faraday and is represented by ‘F’

Approximate value of F = 96500 C mol^-1

Products of Electrolysis

Products of electrolysis depend on the nature of material being electrolysed and type of electrodes being used. If the electrode is inert, it does not participate in chemical reaction, but reactive electrode participates in the chemical reaction. Product of electrolysis also depend on different oxidizing and reducing agents present in the electrolytic cell and their standard electrode potentials.

Batteries

There are two main types of batteries, viz. primary batteries and secondary batteries.

Primary Batteries

In primary batteries, the chemical reaction occurs only once. The battery becomes dead after use over a period of time, and cannot be reused. Dry cell (Leclanche cell) is the most common primary battery. It consists of a zinc container which also acts as anode. The cathode is a carbon (graphite) rod surrounded by powdered magnesium oxide and carbon. The space between the electrodes is filled by a moist paste of ammonium chloride (NH₄Cl) and Zinc chloride (ZnCl₂)

Following is the simplified version of reactions which take place in dry cell:

Anode: `Zn(s)→Zn^(2+)+2^(-)`

Cathode: `Mn\O_2+NH_4+2e^(-)→Mn\O(OH)+NH_3`

Mercury Cell: It is used for low current devices. It consists of zinc-mercury amalgam as anode and a paste of HgO and carbon as cathode. A paste of KOH and ZnO is used as electrolyte. Following reactions take place in mercury cell.

Anode: `Zn(Hg)+2OH^(-)→Zn\O(s)+H_2O+2e^(-)`

Cathode: `Hg\O+H_2O+2e^(-)→Hg(l)+2OH^(-)`

Secondary Batteries

A secondary cell can be recharged by passing current through it in the opposite direction, and thus can be reused. Lead storage battery is the most common secondary cell. It consists of a lead anode and a grid of lead packed with lead oxide (PbO₂) as cathode. A 38% solution of sulphuric acid is used as electrolyte. Following reactions take place in this battery.

Anode: `Pb(s)+SO_4^(2-)(aq)→PbSO_4(s)+2e^(-)`

Cathode: `Pb\O_2(s)+SO_4^(2-)(aq)+4H^(+)(aq)+2e^(-)→Pb\SO_4(s)+2H_2O(l)`

Overall reaction is as follows:

`Pb(s)+Pb\O_2+2H_2SO_4(aq)→2Pb\SO_4(s)+2H_2O(l)`

On charging the battery, the reaction is reversed and PbSO₄ on anode and cathode is converted respectively into Pb and PbO₂`

Fuel Cells

Galvanic cells which are designed to convert the energy of combustion of fuels directly into electrical energy are called fuel cells.

One of the most successful fuel cells uses the reaction of hydrogen to form water. In the cell, hydrogen and oxygen are bubbled through porous carbon electrodes into concentrated aqueous solution of sodium hydroxide. Catalysts (palladium or platinum) are incorporated into electrodes to hasten the reaction. Following reactions take place in this cell.

Cathode: `O_2(g)+2H_2O(l)+4e^(-)→4OH^(-)(aq)`

Anode: `2H_2(g)+4Oh^(-)(aq)→4H_2O(l)+4e^(-)`

Overall Reaction: `2H_2(g)+O_2(g)→2H_2O(l)`

Corrosion

In corrosion, a metal is oxidized by loss of electrons to oxygen and formation of oxides. The chemistry of corrosion is quite complex but it can be considered as an electrochemical phenomenon. Oxidation takes place at a particular spot of iron object, and that spot behaves as anode. The reaction at anode can be written as follows:

Anode: `2Fe(s)→2Fe^(2+)+4e^(-)`

`E_((Fe^(2+))/(Fe))^Θ=-0.44V`

Electrons released at the anodic spot move through the metal and go to another spot on the metal. These electrons reduce oxygen in the presence of H⁺ . This spot behaves as cathode where following reaction takes place.

Cathode: `O_2(g)+4H^(+)(aq)+4e^(-)→2H_2O(l)`

`E^Θ=1.23V`

Overall Reaction:

`2Fe(s)+O_2(g)+4H^(+)(aq)→2Fe^(2+)(aq)+2H_2O(l)`

`E_((cell))^Θ=1.67 V`

Ferrous ions are further oxidized by atmospheric oxygen to ferric ions and come out as rust in the form of hydrated ferric oxide (Fe₂O₃.`x`H₂O)