Group 7 Trends: Oxidising & Reducing Power (AQA A Level Chemistry): Video

Hi, I'm Eleanor with 3 years of experience teaching Chemistry, and this video is about the trends in the Group 7 elements: the halogens, and how their oxidising and reducing power changes as you go down the group.

Here is the idea that ties the whole video together. Going down Group 7, the atoms and the ions get larger — and that single fact explains both of the trends we are about to cover.

Let's state that connecting idea in full. Going down Group 7 the atoms and the ions get bigger, and their outer electrons sit further from the nucleus with more shielding. That is why a halogen gets worse at accepting an electron, so its oxidising power falls, and why a halide ion gets better at giving an electron up, so its reducing power rises. One cause, two trends that run in opposite directions.

One thing to note straight away is the boiling points also rise down the group, but for a separate reason. That trend is about the size of the molecules and how many electrons they contain, not about attracting or releasing an electron. Keep the two explanations apart.

So we will take three things in turn. First the physical trends: colours, boiling points and bond strength. Then the halogens acting as oxidising agents. Then the halide ions acting as reducing agents.

Start with what you can actually see. All the halogens have distinct colours, and those colours get darker going down the group: fluorine is a pale yellow gas, chlorine is a green-yellow gas, bromine is an orange-brown liquid, and iodine is a grey-black solid.

Next, boiling point. The boiling points increase going down the group, which means the halogens become less volatile — volatility being simply how easily a substance evaporates. Fluorine is the most volatile, iodine the least.

Why? The halogens are diatomic molecules held together by weak van der Waals' forces. The more electrons a molecule contains, the stronger those forces become, so the larger the molecule, the harder it is to pull the molecules apart and the higher the boiling point. Notice that this is about the forces between molecules, not the bond inside one.

The bond inside a halogen molecule does the opposite. Going down the group the atoms get bigger, the bonding pair of electrons sits further from the nuclei and is less strongly attracted to them, so the covalent bond gets weaker and the bond enthalpies fall. Fluorine is the exception to this rule and it has the weakest bond. This is due to it being a very small atom which brings the lone pairs close together causing repulsion and weakening the fluorine-fluorine bond.

That is our connecting idea already doing work: bigger atoms mean a weaker hold on electrons.

Now the first of the two main trends. When a halogen reacts with a metal, it accepts an electron from the metal atom to become an ion with a one-minus charge. That makes the halogen an oxidising agent: it oxidises the metal by removing an electron, and is itself reduced as it gains one.

The oxidising power of the halogens decreases going down the group — they get less reactive. And this is exactly the connecting idea again. Going down the group the atomic radii increase, so the outer shell sits further from the nucleus and an incoming electron experiences more shielding from the positive nuclear charge. The halogen's ability to accept an electron therefore falls. The same pattern shows up in electronegativity, which also decreases down the group.

You can see the order directly in displacement reactions. A more reactive halogen displaces a less reactive one from a solution of its halide: chlorine added to sodium bromide gives sodium chloride and bromine, because chlorine is the more reactive of the two.

Now turn it round. Halide ions can act as reducing agents: they donate electrons to another atom, and in doing so they are themselves oxidised.

The reducing power of the halide ions increases going down the group — the opposite direction to the oxidising trend, and for the very same underlying reason. Going down the group the halide ions get larger. Their outermost electrons sit further from the nucleus and experience more shielding from the inner electrons, so they are held less tightly to the positive nuclear charge. A larger halide ion therefore loses an electron more easily, and its reducing power is greater.

So one fact pushes the two trends in opposite directions. A bigger atom is worse at pulling an electron in; a bigger ion is better at letting one go.

Two quick reminders before we finish.

First, when you write about the forces between halogen molecules, remember that instantaneous dipole-induced dipole forces are a type of van der Waals' force, so calling them van der Waals' forces is correct.

Second, a memory aid: halogen comes from a Greek word meaning salt-maker — a useful reminder of their reaction with metals to form salts.

To pull it all together. Going down Group 7, the atoms and the ions get bigger. The oxidising power of the halogens falls, because a bigger atom with more shielding is less able to accept an electron. The reducing power of the halide ions rises, because a bigger ion holds its outermost electrons less tightly and gives it up more easily.

That is the link we opened with, and it is the one worth carrying away: size explains both trends. The boiling points rising down the group is a separate story, about van der Waals' forces between molecules.

For the detail — the bond enthalpy data, the displacement equations and the full explanation of volatility — open the two notes this video is based on.