|
Chemguide: Support for CIE A level Chemistry Learning outcomes 9.4(c) - part 1 This statement is about the reactivity of the halogens as oxidising agents. Before you go on, you should find and read the statement in your copy of the syllabus. You will find that a part of the statement is in bold type. That means that it won't be examined until the very end of the course. I have separated that part out onto another page. Reactions between halogens and other halide ions This is covered on the page halogens as oxidising agents. You only need the first part of this page - the descriptions, the equations and the trend. You do not need the explanations. You can also ignore the reference to fluorine. It is important that you remember what you have read on this page. Other redox reactions illustrating the same trend The important reactions are those you have already read about. However, the teachers' support material also mentions two other sets of reactions as well. Since these aren't specifically mentioned by the syllabus, you probably don't need to remember these, but they could come up in a question asking you to draw some conclusions from given information. Reactions with iron If you pass chlorine gas over hot iron, the iron burns to form iron(III) chloride.
This is a rapid reaction (the iron burns), and the iron has been oxidised to an iron(III) compound - in other words, from an oxidation state of zero in the iron metal to an oxidation state of +3 in the iron(III) compound. | |
|
Note: If you have forgotten about oxidation states (oxidation numbers), it is essential that you revise them before you go on with this topic. There are several places where you will need them in Group VII chemistry. You will find them discussed in learning outcome 6(a). | |
|
If you pass bromine vapour over hot iron, a similar but slightly less vigorous reaction happens, this time producing iron(III) bromide.
Again the iron has been oxidised from an oxidation state of zero to +3. However, the reaction between hot iron and iodine vapour only produces iron(II) iodide, and is much less vigorous.
This time, the iodine is only capable of oxidising the iron as far as the +2 oxidation state. Reactions with sodium thiosulphate solution These are much more complicated! You may have come across the reaction between sodium thiosulphate solution and iodine as a titration.
It isn't easy to see that an oxidation has happened here, but something must be being oxidised because the iodine is being reduced from an oxidation state of zero to -1. (If you can't see this, then you must go to the page mentioned in the last green box.) The element (apart from iodine) whose oxidation state is changing is the sulphur. You need to work out the oxidation states of the sulphur in the sodium thiosulphate and the sodium tetrathionate. In the sodium thiosulphate, the oxidation state of the sulphur is +2. In the sodium tetrathionate, it is +2.5. There has been a small oxidation of the sulphur. | |
|
Note: If you can't get these answers, then go back and and look again at learning outcome 6(a). Don't worry about the fractional oxidation state in the sodium tetrathionate. This an average of the oxidation states of the various sulphur atoms involved. | |
|
The reactions between sodium thiosulphate and chlorine or bromine are totally different, and the sulphur-containing product this time is sodium sulphate, Na2SO4. The oxidation state of the sulphur in this is +6 - a much more major oxidation than that with the iodine. I am not going to give you the equations for these other two reactions, because some students will waste time trying to learn them. Because these reactions are not specifically mentioned by the syllabus, and are just some out of many possible other reactions, you shouldn't be asked to give them in an exam.
© Jim Clark 2011 (modified August 2013) |
|
