Reactions with halogens

​​In a nutshell

Halogens are oxidising agents and can oxidise Group $$1$$​ and $$2$$​ metals and some halogens can oxidise iron(II). Halogens can also undergo disproportionation reactions where they are both oxidised and reduced in the same reaction.

Halogens as oxidising agents

The reactions of halogens involve the halogen gaining an electron to form a full outer shell. Halogens are found in Group $$7$$​ and have seven electrons in their outermost shell. When they react they gain an electron (reduction) to form a full outer shell. Halogen reactivity decreases as you go down the group.

Reactions with Group 1 and 2 metals

Halogens react with Group $$1$$​ and $$2$$​ metals to form ionic compounds which are metal halide salts. The metal is oxidised and the halogen is reduced.

Example 1

​$$2Na(s)\ +\ Cl_2(g)\ \rightarrow\ 2NaCl(s)$$

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Sodium is oxidised from an oxidation number of $$0$$​ to $$+1$$​ whilst chlorine is oxidised from $$0$$​ to $$-1$$​. 

Example 2

​$$Mg(s)\ +\ Cl_2(g)\ \rightarrow\ MgCl_2(s)$$

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Magnesium is oxidised from an oxidation number of $$0$$​ to $$+2$$​ whilst chlorine is oxidised from $$0$$​ to $$-1$$​. 

Reactions with iron(II) 

Chlorine and bromine can oxidise iron(II) to iron(III). 

​$$Cl_2(g)\ +\ 2Fe^{2+}(aq)\ \rightarrow\ 2Cl^-(aq)\ + \ 2Fe^{3+}(aq) \\Br_2(g)\ +\ 2Fe^{2+}(aq)\ \rightarrow\ 2Br^-(aq)\ + \ 2Fe^{3+}(aq)$$

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Iodine is not strong enough to oxidise iron(II) and is itself oxidised by iron(III).

Disproportionation reactions

Halogens undergo disproportionation reactions with alkalis. In these reactions halogens are both reduced and oxidised simultaneously, this is known as a disproportionation reaction. 

​$$X_2\ +\ 2NaOH\ \rightarrow\ NaXO\ +\ NaX\ +\ H_2O$$

Where $$X$$​ is a halogen 

Example 1

Chlorine and cold sodium hydroxide are used to make bleach $$(NaClO)$$. This is an example of a disproportionation reaction.

​$$2NaOH(aq)\ +\ Cl_2(g)\ \rightarrow\ NaClO(aq)\ +\ NaCl(aq)\ +\ H_2O(l)\\$$

Example 2

Mixing chlorine with water causes the chlorine to undergo a disproportionation reaction.

$$Cl_2(g)\ +\ H_2O(l)\ \rightarrow\ HCl(aq)\ +\ HClO(aq)$$

The chlorate ion $$(ClO^-)$$ kills bacteria, this is why chlorine is added to water to make it safe to drink or swim in. ​

Reactions with hydrogen 

The trend in reactivity of halogens is illustrated by how they react with hydrogen, as you go down the group the reaction becomes less vigorous. Hydrogen halide molecules are formed when hydrogen reacts with halides.   

​$$X_2\ +\ H_2\ \rightarrow\ 2HX$$ 

Where $$X$$​ is a halogen

The thermal stability of the hydrogen halides decreases as you go down the group. This is because the covalent bond between the hydrogen and halide becomes weaker, due to an increase in size and distance between the bond. Therefore, the temperature required to overcome this bond decreases.