Reactions of Halogens

Did you know that we can use bromine to reduce mercury emissions from coal power plants? Scientists have already discovered that they can use something called activated carbon to trap mercury particles released when coal is burnt. However, bromine has recently come into the spotlight, because it can help the activated carbon oxidise mercury and trap more of the harmful element. This is just one example of the many reactions of halogens.

• This article looks at many of the different reactions of halogens.
• You'll explore the trends in various halogen redox and displacement reactions, including the reactions of halogens with halides, hydrogen, metals, sodium hydroxide and organic molecules.
• You'll learn how reactivity varies as you move down the group.

What are halogens?

Fig. 1 - The halogens

Halogens can react in many ways. They can:

• Displace other halogens.

• React with hydrogen.

• React with metals.

• React with sodium hydroxide.

• React with alkanes, benzene and other organic molecules.

We'll explore these reactions throughout the rest of this article.

Reactions of halogens as oxidising agents

Halogens can act as oxidising agents.

You may already know the acronym OIL RIG. It helps you remember the movement of electrons in redox reactions:

Fig. 2 - The OIL RIG acronym

Similarly, there is another acronym, RAD OAT. It helps you remember the actions of oxidising and reducing agents, once again in terms of electrons:

Fig. 3 - The RAD OAT acronym

This means that an oxidising agent takes electrons from another species and gains them itself. It is reduced.

In general, the oxidising power of the halogens decreases as you go down the group. In fact, fluorine is one of the most potent oxidising agents out there!

Displacement reactions of halogens

An example of an oxidation reaction involving halogens is a displacement reaction.

In principle, these reactions are pretty simple. A more reactive halogen displaces a less reactive halide from an aqueous solution:

• The more reactive halogen atoms oxidise the less reactive halide ions.
• The halide ions lose electrons and form halogen atoms. They are oxidised.
• The halogen atoms gain electrons to form halide ions. They are reduced.

Suppose you have a sodium bromide solution to which you add chlorine. The halogens get less reactive as you go down the group in the periodic table, meaning that chlorine is more reactive than bromine.

Fig. 4 - Reactivity of halogens

So chlorine atoms displace bromide ions from an aqueous solution. The bromide ions are oxidised and lose electrons, whilst the chlorine atoms are reduced and gain electrons. The bromide ions form bromine, and the chlorine forms chloride ions. Here's the equation:

${\mathrm{Cl}}_2\left(\mathrm{aq}\right)+2{\mathrm{Br}}^{-}\left(\mathrm{aq}\right)\to 2{\mathrm{Cl}}^{-}\left(\mathrm{aq}\right)+{\mathrm{Br}}_2\left(\mathrm{aq}\right)$

You can also formulate this as two half equations:

${\mathrm{Cl}}_2\left(\mathrm{aq}\right)+2{\mathrm{e}}^{-}\to 2{\mathrm{Cl}}^{-}\left(\mathrm{aq}\right)$

$2{\mathrm{Br}}^{-}\left(\mathrm{aq}\right)\to {\mathrm{Br}}_2\left(\mathrm{aq}\right)+2{\mathrm{e}}^{-}$

We know this reaction occurs because we can see a colour change. The bromide ions turn the solution orange-brown.

However, if you add bromine to a solution containing chloride ions, nothing will happen – bromine is less reactive than chlorine. It isn't a potent enough oxidising agent to oxidise the chloride ions.

The following table shows the displacement reactions between different combinations of chlorine, bromine, iodine and their aqueous ions alongside any observable colour changes. Notice that fluorine is not inclluded. The reason is that fluorine is too strong of an oxidising agent. In fact, it manages to oxidise water into oxygen, complicating the reaction further. Likewise, we haven't included astatine because it is both extremely rare and highly reactive, and you'll probably never come across it.

Fig. 5 - Halogen displacement reactions

Reactions of halogens with hydrogen

Halogens react with hydrogen in another example of a redox reaction. They form a hydrogen halide, HX. Like in the displacement reactions that we looked at above, the halogens act as oxidising agents, and so their reactivity falls as you move down the group.

For example, fluorine and hydrogen react explosively to give hydrogen fluoride gas:

${\mathrm{F}}_2\left(\mathrm{g}\right)+{\mathrm{H}}_2\left(\mathrm{g}\right)\to 2\mathrm{HF}\left(\mathrm{g}\right)$

Hydrogen is oxidised and loses electrons, whilst fluorine is reduced and gains electrons.

However, iodine and hydrogen only partially react. The mixture forms an equilibrium:

${\mathrm{I}}_2\left(\mathrm{g}\right)+{\mathrm{H}}_2\left(\mathrm{g}\right)\rightleftharpoons 2\mathrm{HI}\left(\mathrm{g}\right)$

Reactions of halogens with metals

Let's now look at a few examples of reactions between halogens and metals. All of these reactions form salts. Again, the reactions are redox reactions and reactivity decreases as you move down the group.

Reaction with sodium

The halogens react vigorously with hot sodium metal to produce a sodium halide. They oxidise the sodium into Na(I) ions with a charge of +1. Sodium fluoride, chloride, bromide and iodide are all white solids.

For example, the reaction between chlorine and sodium:

${\mathrm{Cl}}_2\left(\mathrm{g}\right)+2\mathrm{Na}\left(\mathrm{s}\right)\to 2\mathrm{NaCl}\left(\mathrm{s}\right)$

Reaction with iron

Halogens can oxidise iron into iron(III) ions. The overall reaction produces an iron(III) halide. However, this reaction only happens with fluorine, chlorine and bromine – iodine isn't a potent enough oxidising agent for the reaction to occur.

For example, the reaction between chlorine and iron:

$3{\mathrm{Cl}}_2\left(\mathrm{g}\right)+2\mathrm{Fe}\left(\mathrm{s}\right)\to 2{\mathrm{FeCl}}_3\left(\mathrm{s}\right)$

Reactions of halogens with sodium hydroxide

Another redox reaction involving halogens is their reaction with sodium hydroxide. However, this reaction is slightly different: it is an example of a disproportionation reaction.

Let's look at the reaction between chlorine and cold sodium hydroxide. The two react to produce sodium chloride, sodium chlorate(I) and water:

${\mathrm{Cl}}_2\left(\mathrm{aq}\right)+2\mathrm{NaOH}\left(\mathrm{aq}\right)\to \mathrm{NaCl}\left(\mathrm{aq}\right)+\mathrm{NaClO}\left(\mathrm{aq}\right)+{\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right)$

Look at the oxidation states of the compounds produced:

• In sodium chloride, chlorine has an oxidation state of -1. The chlorine has been reduced.
• In sodium chlorate, chlorine has an oxidation state of +1. The chlorine has been oxidised.

Oxidations states of chlorine is the disproportionation reaction with sodium hydroxide. StudySmarter Originals

Reacting chlorine with hot sodium hydroxide produces a slightly different product: sodium chlorate(V). In this compound, chlorine has an oxidation state of +5:

$3{\mathrm{Cl}}_2\left(\mathrm{aq}\right)+6\mathrm{NaOH}\left(\mathrm{aq}\right)\to 5\mathrm{NaCl}\left(\mathrm{aq}\right)+{\mathrm{NaClO}}_3\left(\mathrm{aq}\right)+3{\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right)$

Bromine and iodine react similarly. However, you can produce sodium bromate(V) at a much lower temperature than needed to make sodium chlorate(V) since bromine is a better reducing agent than chlorine. Likewise, producing sodium iodate(V) is even easier.

In general, whilst halogens become better oxidising agents as you go up the group, they become better reducing agents as you go down the group.

Reactions of halogens with organic molecules

Finally, halogens can react with organic molecules. You'll see these again as part of organic chemistry, but we'll look at them now too.

Reaction with alkanes

If you mix an alkane with a halogen and shine UV light on the mixture, the two molecules will react. Halogen atoms replace some of the alkane's hydrogen atoms to produce a halogenoalkane. This reaction is known as free radical substitution. If the halogen used is chlorine, it's also known as chlorination.

Free radical substitution can produce a variety of products. But to give an example, the reaction between ethane and chlorine could produce chloroethane:

${\mathrm{CH}}_3{\mathrm{CH}}_3+{\mathrm{Cl}}_2\to {\mathrm{CH}}_3{\mathrm{CH}}_2\mathrm{Cl}+\mathrm{HCl}$

Reaction with benzene

Halogens react with benzene in another substitution reaction known as electrophilic substitution. This reaction needs a catalyst – either an aluminium halide, or iron.

For example, reacting benzene and chlorine in the presence of aluminium chloride produces chlorobenzene and hydrochloric acid:

${\mathrm{C}}_6\mathrm{H}{}_6+{\mathrm{Cl}}_2\to {\mathrm{C}}_6{\mathrm{H}}_5\mathrm{Cl}+\mathrm{HCl}$

Regarding reactions with both alkanes and benzene, the reactivity of the halogens decreases as you move down the group. Fluorination of benzene is explosive, whilst iodination doesn't tend to happen. Likewise, you'd be hard-pressed to react iodine with any alkanes. Chlorine and bromine are the preferred reactants of choice. These reactions also aren't redox reactions – no electrons transfer. Instead, they form covalent molecules.

Summary of reactions of halogens

Yes, we know – we've thrown a lot of new information at you! But once you get your head around reactivity and oxidising ability, most halogen reactions aren't too tricky:

• As you move down the group, oxidising power decreases.
• As you move down the group, reducing power increases.