Combination Reactions – Other Ionic Reactions

00:01 Combination reactions.

00:02 So, having done RedOx, let’s move on to combination reactions in which two substances, usually two elements, combine to form a third substance. And this is what we would see if, for example, we had some sodium on a desk, the sodium would react with the oxygen around it and form, in the first instance, sodium oxide. If we were to take sodium metal and we were to react it with chlorine, apart from having to stand well back, because the reaction is very exothermic, we would also find a rather expensive way of generating sodium chloride.

00:39 And these are some of the reactions that occur every day. Calcium oxide plus sulphur dioxide can form calcium sulphide, as shown here in the below equation. And this is a way, for example, for mediating harmful sulphur oxide gases from waste, from fuel combustion.

01:03 Decomposition reactions. A decomposition reaction, as the name suggests, is where a single compound breaks down to give two or more other substances. And this is a particularly interesting one because here, we have something called ammonium dichromate.

01:18 Now, it looks rather complicated to look at. We have our two equivalents of ammonium and here, we have a transition metal that you perhaps haven’t encountered before, chromium.

01:28 Chromium, as you can see here, exists in a highly oxidised state. Indeed it exists in its +6 oxidation state, as you can tell. Now, the reason you can tell or you should be able to tell is from what I’ve said before. Chromium, sorry, oxygen, in this case, exists in an oxidation state of -2, right? So, that means that there are 7 oxygens, 7 x -2 = -14.

02:02 Ammonium ions as cations have a formal charge of +1. There is two of them. This results in an overall charge of -12. To counter that charge, each chromium must have an oxidation state of +6, in order for the oxidation states to be even and evenly matched.

02:22 Now, let’s look what happens. We move to the products for this decomposition reaction.

02:28 We see that the chromium, which originally had an oxidation state of +6, has been reduced.

02:35 And the reason we know it’s been reduced i.e. it has gained electrons is because, going from what we said before, that each oxygen having an oxidation state of -2, 3 x -2 = -6, each chromium must have a formal oxidation state or charge of +3 for the oxidation states to be even. And we can see what’s been oxidised, can’t we? Because what’s been happening is that ammonium, as effectively the fuel for this oxidation reaction, is being converted into water and nitrogen gas. And actually, to view this reaction, it looks a lot like a volcano, you just set fire to it, it contains its own fuel, its own oxidising agent and you produce a rather nice green fog.

03:26 Then we come on to displacement reactions. These are also called single replacement reactions which where element reacts with a compound displacing another element from it. And here, we have an example of this. So, zinc, not unlike, for example, magnesium, many of the alkaline metals or alkaline earth metals, will react with acids to form a salt and to produce hydrogen. This is relatively straightforward, it’s a very common reaction that you see in schools. So, you take zinc or magnesium ribbon, you add hydrochloric acid and you generate a bit of hydrogen gas which you can detect by the squeaky pop test.

04:09 And finally, combustion reactions. So, this is something that we’ve been doing since time immemorial where a substance itself reacts with oxygen usually with a rapid release of heat to produce a flame. So, here, we have, for example, the reaction of butane, which is an alkane, and we’ll come on to that in the next module, reacting with equivalents of oxygen to generate carbon dioxide and water. This is exothermic and obviously produces an amount of energy given as obviously light and heat.