Testing for Halides & the Chemistry of Chlorine (AQA A Level Chemistry): Video

Hi, I'm Eleanor with 3 years of experience teaching Chemistry, and this video is about testing for halide ions and the chemistry of chlorine.

Both of the notes behind this video are about Group 7 reactions we deliberately put to work in real life: some let us identify an unknown halide from what we see in a test tube, and others let us make water safe to drink.

Let's take that connecting idea in full. Everything in this video is a reaction used for a purpose. On one side, the reactions of halide ions give us a way of finding out which halide is present in an unknown solution — and the answer comes from what you can see: the colour of a precipitate, the colour of a fume. On the other, the reactions of chlorine are what make water treatment possible.

So we'll take four things in turn. First the silver nitrate test, the standard way of identifying a halide. Then the reactions of halides with concentrated sulfuric acid, which bring us back to the reducing-power trend. Then disproportionation, the reaction type chlorine undergoes in both water and an alkali. And finally what that chemistry is used for in water treatment.

The standard test for a halide ion starts by acidifying the unknown solution with dilute nitric acid. This step matters: the acid stops carbonate ions precipitating out with the silver and giving a false positive.

Then add silver nitrate solution dropwise. If a halide ion is present you will get a precipitate of the silver halide — silver ion plus halide ion gives solid silver halide.

The colour tells you which one. Silver chloride is white, silver bromide is cream, and silver iodide is pale yellow.

But those three colours can look very similar, so ammonia is used as a follow-up. If the precipitate dissolves in dilute ammonia, it is chloride. If it survives dilute ammonia but dissolves in concentrated ammonia, it is bromide. And if it dissolves in neither, it is iodide.

There is a second set of reactions that also separates the halides, and this one is a redox reaction. Chloride, bromide and iodide ions all react with concentrated sulfuric acid to give toxic gases, so this is a fume-cupboard reaction.

In this reaction, every halide gives the hydrogen halide first: sulfuric acid plus the solid sodium halide gives HX gas plus sodium hydrogensulfate. With chloride that is where it stops — misty fumes of hydrogen chloride, and nothing more.

Bromide goes further. The concentrated sulfuric acid oxidises the hydrogen bromide to bromine, and is itself reduced to sulfur dioxide, so you get orange-brown bromine alongside choking sulfur dioxide.

Iodide goes furthest of all. It reduces the sulfuric acid not only to sulfur dioxide but then on to sulfur, a yellow solid, and further still to hydrogen sulfide, with its smell of rotten eggs.

That progression is the reducing-power trend made visible: the further down the group, the more strongly the halide reduces.

Now to chlorine itself, and a reaction type worth naming. Disproportionation is a reaction in which the same species is both oxidised and reduced.

Chlorine does exactly this in water. Chlorine plus water gives hydrochloric acid and chloric(I) acid, HClO. Look at the oxidation numbers: chlorine starts at zero, and ends up at minus one in the HCl and plus one in the HClO. One element, going both ways in the same equation.

It does the same in cold, dilute alkali, at around fifteen degrees celcius. Chlorine plus sodium hydroxide gives sodium chloride, sodium chlorate(I) and water. Ionically, chlorine plus hydroxide gives chloride, chlorate(I) and water — again zero to minus one and zero to plus one in the same reaction.

That is the signature to look for: one species, two oxidation states, one reaction.

So what is that for? Chlorine is used to clean water and make it drinkable, and the disproportionation reaction is why it works.

The chloric(I) acid sterilises the water by killing bacteria. It can dissociate further to give the chlorate(I) ion, and that acts as a sterilising agent too, so both species do the same job.

Chlorine is used in drinking water and in swimming pools, in small controlled amounts. There is a judgement to make here: chlorine is toxic, and there is a possible risk from chlorinated hydrocarbons forming — but the benefit to health from killing bacteria outweighs those risks.

One practical note. In shallow pools the water is more exposed to sunlight, and in sunlight chlorine reacts with water to give hydrochloric acid and oxygen. The chlorine is lost quickly and has to be topped up, which adds to the cost.

Two things to remember, and two things that cost marks.

To remember: it gets easier to oxidise the hydrogen halides as you go down Group 7 — the halides become stronger reducing agents. And in water treatment, both HClO and the chlorate(I) ion act as sterilising agents.

What costs marks: acidify with dilute nitric acid, never hydrochloric, because hydrochloric acid introduces chloride ions and ruins the test. And be precise about observations — examiners penalise students who write that the solution or the ions dissolve in ammonia. It is the precipitate that dissolves, and mark schemes want white, cream or pale yellow.

To pull it together. Silver nitrate followed by ammonia identifies an unknown halide: white, cream or pale yellow, then dilute and concentrated ammonia to tell them apart.

Concentrated sulfuric acid gives a second test, this one redox, and the further down the group you go the more strongly the halide reduces.

Chlorine disproportionates — oxidised and reduced in the same reaction — both in water and in cold dilute alkali.

And that is the thread: every reaction in this video is one we put to use, either to identify something or to make water safe to drink.

For the full set of equations and observations, open the two notes behind this video.