Uses of chlorine and its compounds

Halogens)

reactions of chlorine with aqueous sodium hydroxide
11.2.1.13 understand the use

of chlorine as a water treatment and understand the balance of risks and benefits in this process

purification (drinking water, pool water)
- considers the formation of

substances with chlorinated water and their effect on humans
- considers the advantages and disadvantages of chlorination of water
- considers the harmful effects of unchlorinated water

of redox reaction called disproportionation when it reacts with alkali.

Disproportionation can be thought of as a ‘self reduction/oxidation’ reaction.
When chlorine reacts with dilute alkali some chlorine atoms are reduced and some are oxidised in the same reaction.
The actual reaction that takes place depends on the temperature.

+ NaClO(aq) + H2O(l)
sodium chlorate(I)
The ionic equation for the

reaction is:
Cl2(aq) + 2OH−(aq) → Cl−(aq) + ClO−(aq) + H2O(l)
0 −1 +1
oxidation number of Cl
The ionic equation for this redox reaction can be split into two half-equations, showing the reduction and oxidation.
The reduction reaction (in which chlorine’s oxidation number is reduced is):
½Cl2 + e− → Cl−
0 −1
The oxidation reaction is: ½Cl2 + 2OH− → ClO− + H2O + e−
0 +1

hot concentrated aqueous sodium hydroxide a different disproportionation reaction takes

place:

small amount of chlorine to a water supply will kill

bacteria and make the water safer to drink.
The chlorine undergoes disproportionation in water:
Cl2(aq) + H2O(l) → HCl(aq) + HClO(aq)
0 −1 +1
HClO is called chloric(I) acid, and it decomposes slowly in solution.
One theory suggests that it produces reactive oxygen atoms that can kill bacteria in water:
HClO → HCl + [O]

and sodium chlorate(I) (NaClO), made from chlorine and cold alkali.

It ‘bleaches’ colours and stains because oxygen atoms from the chlorate(I) ions oxidise dye and other coloured molecules.
They also kill bacteria when toilets are cleaned with bleach.