As the atoms get bigger, the new electrons find themselves further from the nucleus, and more and more screened from it by the inner electrons (offsetting the effect of the greater nuclear charge). How easily the element forms its ions depends on how strongly the new electrons are attracted. This is normally given for the trend in oxidising ability of chlorine, bromine and iodine, and goes like this: We'll deal with this first before giving a proper explanation. Unfortunately, this is often over-simplified to give what is actually a faulty and misleading explanation. The reason that the hydrated ions form less readily as you go down the Group is a fairly complicated mixture of several factors. Looking at all four of the common halogens:Īs you go down the Group, the ease with which these hydrated ions are formed falls, and so the halogens become less good as oxidising agents - less ready to take electrons from something else. Whenever one of these halogens is involved in oxidising something in solution, the halogen ends up as halide ions with water molecules attached to them. This all means that oxidising ability falls as you go down the Group. Bromine can remove electrons from iodide ions to give iodine - and the iodine can't get them back from the bromide ions formed. Similarly bromine is a more powerful oxidising agent than iodine. That means that chlorine is a more powerful oxidising agent than either bromine or iodine. Bromine and iodine can't get those electrons back from the chloride ions formed. That means that they are all potentially oxidising agents.įluorine is such a powerful oxidising agent that you can't reasonably do solution reactions with it.Ĭhlorine has the ability to take electrons from both bromide ions and iodide ions. Each of the elements (for example, chlorine) could potentially take electrons from something else to make their ions (e.g. Iodine won't oxidise any of the other halide ions (unless you happened to have some extremely radioactive and amazingly rare astatide ions - astatine is at the bottom of this Group). (You have just seen exactly the reverse of that happening.)Ī red solution of iodine is formed (see the note above) until the bromine is in excess. It isn't a strong enough oxidising agent to convert chloride ions into chlorine. If the chlorine is in excess, obviously there isn't anything left for the iodine to react with, and so it remains as a dark grey precipitate.īromine can only oxidise iodide ions to iodine. Note: The reason for the red solution is that iodine dissolves in potassium iodide (or other soluble iodides) by reacting to give a red ion, I 3. The iodine appears either as a red solution if you are mean with the amount of chlorine you use, or as a dark grey precipitate if the chlorine is in excess. The bromine appears as an orange solution.Īs you have seen above, chlorine can also oxidise iodide ions (in, for example, potassium iodide solution) to iodine: For example, chlorine can oxidise the bromide ions (in, for example, potassium bromide solution) to bromine: In each case, a halogen higher in the Group can oxidise the ions of one lower down.
Fluorine oxidises water to oxygen and so it is impossible to do simple solution reactions with it. We'll have to exclude fluorine from this descriptive bit, because it is too strong an oxidising agent. This is obviously a redox reaction in which chlorine is acting as an oxidising agent. The chlorine molecules have gained electrons to form chloride ions. The iodide ions have lost electrons to form iodine molecules. In the chlorine and iodide ion case, the reaction would be: The sodium or potassium ions will be spectator ions, and are completely irrelevant to the reaction. The iodide ions will be in a solution of a salt like sodium or potassium iodide. We are going to look at the reactions between one halogen (chlorine, say) and the ions of another one (iodide ions, perhaps). Note: If you aren't comfortable with terms like oxidation and oxidising agent in terms of electron transfer, then you should explore the area of the site dealing with redox reactions before you go on. We are going to look at the ability of one halogen to oxidise the ions of another one, and how that changes as you go down the Group. This page explores the trend in oxidising ability of the Group 7 elements (the halogens) - fluorine, chlorine, bromine and iodine. THE OXIDISING ABILITY OF THE GROUP 7 ELEMENTS (THE HALOGENS)
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