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The elements of Group 7 (halogens) (Topic 4B)
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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.
Electronegativity refers to the ability of an atom to attract a shared pair of electrons in a covalent bond.
As you go down group 7, the electronegativity decreases.
This occurs because the atomic radius increases with the addition of electron shells, which causes greater shielding of the outer electrons from the positive charge of the nucleus.
Moving down the Group 7 elements, reactivity decreases due to:
- increasing atomic radius
- increasing shielding effect from inner shells of electrons
- weaker electrostatic attraction between the positive nuclei and the outer shell.
Therefore, the ease of gaining an electron decreases down Group 7.
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.
Displacement reactions are chemical reactions in which a more reactive element displaces a less reactive element from a compound.
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.
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.
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:
A disproportionation reaction of chlorine occurs when chlorine is reacted with hot alkali at temperatures above . It is important to note that the products are different from those formed at room temperature.
The ionic equation for this reaction is:
The atoms of chlorine are in an oxidation state of 0 in elemental chlorine. They are;
- oxidised to +5 in the sodium chlorate, .
- reduced to −1 in the sodium chloride, .
The reduction half-equation is:
The oxidation half equation is:
In reactions with concentrated sulfuric acid, the halide ion donates electrons to a proton, to produce a gaseous hydrogen halide, which immediately escapes from the system.
If the hydrogen halide is exposed to moist air, steam fumes are formed. The gases produced are toxic, so a fume cupboard must be used.
The general reaction of halide ions, with concentrated sulphuric acid is:
Halide ions can act as reducing agents and donate electrons to another atom.
The reducing power of halide ions with sulfuric acid increases down the group. This is due to the ionic radii and number of filled subshells, increasing down the group. The valence electrons are further away and more shielded from the positive nucleus, and are therefore held less tightly.
This results in the halide ion more easily losing electrons making it a better reducing agent.
When concentrated sulfuric acid is reacted with sodium fluoride, hydrogen fluoride and sodium hydrogen sulfate are produced.
Key observations: white fumes of gas.
No further reaction takes place as is not a strong enough reducing agent
When concentrated sulfuric acid is reacted with sodium chloride, hydrochloric acid and sodium hydrogen sulfate are produced.
Key observations: White fumes of gas.
No further reaction takes place as is not a strong enough reducing agent.
When concentrated sulfuric acid is reacted with sodium bromide, hydrogen bromide and sodium hydrogen sulfate are produced.
The concentrated sulfuric acid then oxidises the hydrogen bromide to produce bromine and water.
The sulfuric acid is reduced to sulfur dioxide gas.
Sulfur reduced from OS +VI to OS +IV.
Key observations: The bromine is observed as a brown gas.
Hydrogen iodide and sodium hydrogen sulfate are produced when concentrated sulfuric acid is reacted with sodium iodide.
Subsequent reactions:
Sulfur is reduced from OS +VI to OS +IV.
Key observations: The iodine is seen as a purple vapour.
Sulfur is reduced from OS +VI to OS 0
Key observations: The iodine is seen as a purple vapour. Sulfur is seen as a yellow solid.
Sulfur is reduced from OS +VI to OS -2
Key observations: The iodine is seen as a purple vapour. Hydrogen sulfide has a strong smell of bad eggs.
As you go down group 7, the hydrogen halides become easier to oxidise and therefore stronger reducing agents.
The reactions between the halide salt and sulfuric acid demonstrates this trend, as the degree of oxidation observed for sulfur increases down the group. These reactions are shown in the table below.
When a hydrogen halide reacts with ammonia gas, an ammonium halide is produced.
General Equation:
Where:
- , or .
Ammonia acts as a base, accepting a from the hydrogen halide.
Hydrogen halides dissolve in water to produce acidic solutions because they dissociate into and ions.
General Equation:
, , and are strong acids and dissociate almost completely in solution.
is a weak acid due to strong hydrogen bonding, and it only partially dissociates in solution.
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:
