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The halogens (3.1.3)
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The non-metal elements in group 7 are called halogens.
All of the elements in group 7 covalently bond with themselves to produce diatomic molecules.
Figure: Lewis diagram showing the formation of a molecule of chlorine (Cl2) from two atoms of chlorine
As you go down group 7, the melting points and boiling points of the halogens increases.
At room temperature and pressure, flourine and chlorine are gasses, bromine is a liquid, and iodine is a solid.
Figure: Graph showing the boiling points and melting points of the group 7 elements
The halogen molecules increase in molecular size as you move down the group. This results in larger and more polarisable electron clouds, as shown in the diagram below.
London dispersion forces, which are caused by temporary fluctuations in electron distribution, become stronger in larger, more polarisable molecules.
The larger halogen molecules exhibit stronger London dispersion forces compared to the smaller halogen molecules. This means more energy is required to overcome these forces, leading to higher boiling and melting points for the larger halogen molecules.
The halogens are in group 7 and hold seven outer shell electrons.
The outer shell electron configuration of the halogens is .
For example, chlorine’s electron configuration is:
The ‘noble gas configuration’ for an element is quicker to write than the full electron configuration. The symbol of the preceding noble gas is given in squared brackets, followed by the remaining electron configuration.
For example, chlorine can be written as instead of .
To achieve a full outer shell of electrons, halogens are reduced, gaining one electron in redox reactions to form -1 ions.
When halogens react with metals, the halogen is reduced and the metal is oxidised.
The halogens act as oxidising agents, which means they oxidise other species, and are themselves reduced.
The ability of halogens to act as oxidising agents decreases down the group.
Displacement reactions are chemical reactions in which a more reactive element displaces a less reactive element from a compound.
Reactivity decreases as you go down group 7.
The displacement equations below show elemental halogens displacing halide ions below them in the group.
Half equations can be used to show the transfer of electrons.
The half equations for the reaction between chlorine () and potassium bromide () are:
In this reaction, elemental chlorine gains electrons and is reduced to chloride ions. It acts as an oxidising agent, as it oxidises the bromide ions to elemental bromine.
The reactivity of halogens can be observed in aqueous conditions through the reactions (or lack of reactions) between halide salts and elemental halogens.
The observations in the table below show that an elemental halogen will only displace a halide ion below it in the group.
This shows the reactivity decreases as you go down group 7.
These colour changes are different in an organic solvent.
The reactivity of halogens can be observed in cyclohexane through the reactions (or lack of reactions) between halide salts and elemental halogens.
The observations in the table below show that an elemental halogen will only displace a halide ion below it in the group.
This shows reactivity decreases as you go down group 7.
These colour changes are different in aqueous conditions.
Group 7 elements gain one electron to form -1 ions. They are reduced and act as oxidising agents. Their oxidising power depends on how easily they attract electrons.
Oxidising power decreases down the group because atomic radius increases meaning the outer shell is further from the positive nucleus.
Increased electron shielding also reduces the nuclear attraction for incoming electrons, making it harder for the halogen atom to gain an electron.
This means the oxidising power and the reactivity of halogens both decrease as you go down the group.
A disproportionation reaction is a reaction in which the same element is simultaneously oxidised and reduced.
The disproportionation reaction of chlorine is used to sterilise drinking water.
The reaction below shows the atoms of chlorine are in an oxidation state of in elemental chlorine. They are oxidised to in hypochlorous acid, , and reduced to in .
can dissociate further:
Chlorate ions, produced behave as sterilising agents, and will kill bacteria present in the water.
Sodium hypochlorite, , is used to disinfect, by killing bacteria, and is the main active component in bleach.
A disproportionation reaction of chlorine occurs at room temperature in the manufacture of bleach:
The chlorine atoms are in an oxidation state of in elemental chlorine.
They are oxidised to in the sodium hypochlorite, , and reduced to in the sodium chloride.
The ionic equation for this reaction is:
The reduction half-equation is:
The oxidation half equation is:
Chlorine is a powerful disinfectant that effectively kills harmful bacteria, viruses, and other pathogens present in water.
Its ability to react with and destroy the cellular structures of microorganisms makes it an essential part of maintaining water safety.
The use of chlorine provides a reliable and cost-effective solution to ensuring that drinking water is free from waterborne diseases, safeguarding public health on a large scale. Water that is safe to drink is sometimes called potable water.
Despite its effectiveness, the use of chlorine comes with inherent risks.
One significant concern is the potential release of toxic chlorine gas during the chlorination process. Exposure to chlorine gas can cause respiratory problems and other health issues.
Additionally, the formation of chlorinated hydrocarbons, resulting from the reaction of chlorine with organic matter in water, raises environmental and health concerns. Some chlorinated hydrocarbons are known to be carcinogenic and may pose other long-term risks to both human and environmental health.
Halide ions can be identified with simple test tube reactions.
The unknown solution is first acidified with nitric acid.
Silver nitrate is then added dropwise to produce a silver halide precipitate.
The reaction follows the following general ionic equation:
If a solution contains halide ions it will form a silver halide precipitate, .
The colour of the precipitate formed depends on the halide ion present:
Silver chloride : white
Silver bromide : cream
Silver iodide : yellow
Silver carbonates also precipitate in neutral solutions, but the nitric acid added reacts with any carbonate ions that might be present, preventing a false positive result.
Hydrochloric acid should never be used to acidify the silver nitrate in this test as it is a source of chloride ions. These would react with silver ions, giving a false positive result.
Further testing of precipitates of silver halides, formed by the addition of acidified silver nitrate to a halide ion solution, can be carried out by adding dilute ammonia, followed by concentrated ammonia to confirm the silver halide present.
The table shows the observations of reactions of halide ions with silver nitrate, and the subsequent addition of ammonia solutions:
