Chemguide: Support for CIE A level Chemistry

Learning outcome 24: Electrochemistry

24.2: Standard electrode potentials

Learning outcome 24.2.6

This statement looks at how electrode potentials tell you about oxidising and reducing ability. Read the statement before you go on.

Now is the time to re-read three pages about E° values and oxidising and reducing ability. Start with the page The Electrochemical Series.

When you are happy with that, go on to the section "Including these new redox potentials in the electrochemical series" on the page Redox potentials for non-metal and other systems.

And then re-read the final page in the sequence called Making predictions. . .

The equivalent statement in the previous syllabus asked you to relate this to the reactions of the halogens. That isn't specifically mentioned in the current statement, but there is no reason why CIE couldn't set a question on it.

What follows is taken from what I previously wrote. It is important that you can understand it, but not necessarily learn it.

You will remember that the oxidising ability of the halogens decreases as you go down the Group in the Periodic Table.

The E° values of the four halogens from fluorine to iodine are:

Remember that the electrode potentials give a measure of the positions of the equilibria. The more positive, the further the equilibrium lies to the right.

In the fluorine case, the E° value is almost as positive as they get. That means that fluorine will very readily pick up electrons to make fluoride ions.

Fluorine will therefore remove electrons from other things extremely well. Taking electrons away from something is oxidising it. So fluorine is a very powerful oxidising agent indeed.

As you go down the rest of the group, the E° values become less positive, and so the oxidising ability decreases.

Three simple examples of this

Why does chlorine oxidise iodide ions to iodine?

The two E° values are:

When you couple two of these equilibria together in a test tube, the more positive one will tend to move to the right, and the more negative one (or less positive one) to the left.

That is exactly what you want to happen to turn iodide ions into iodine. The chlorine E° is more positive, and so chlorine molecules take electrons from the iodide ions to turn them into iodine.

Why won't bromine oxidise chloride ions to chlorine?

The two E° values are:

The chlorine equilibrium lies further to the right because it is the more positive. That means that if you couple the two equilibria together, you would expect the chlorine one to move to the right and the bromine one to the left.

But if you start with bromine and chloride ions, the two equilibria are already as far in those directions as possible. To get a reaction, they would have to move in a direction opposite to that predicted by the E° values. That can't happen.

Which halogens could you use to oxidise Fe2+ to Fe3+?

The E° value for the iron(II) / iron(III) system is

. . . and what you want to do is to drive it to the left to turn Fe2+ into Fe3+. To do that you would need to couple it with something with a more positive E° value.

If you look at the halogen list above, you will see that fluorine, chlorine and bromine are all capable of oxidising iron(II) to iron(III), but iodine isn't.

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© Jim Clark 2020