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.
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.
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.
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.
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.
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.
Now, let’s look what happens. We move to
the products for this decomposition reaction.
We see that the chromium, which originally
had an oxidation state of +6, has been reduced.
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.
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.
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