So, it's been a while since we've burnt anything here at Exhibition Chemistry,
time to rectify that.
Yes this month we're going to be returning to the home of all the best things that go
bang, pop and whizz and that's of course the reactivity series.
So this demonstration is going to be showing something to do with
basically the true reactivity of aluminium. When I normally show this I tend to be
showing it
to students who have already seen reactions of metals with
metal compounds and also reactions of metals with acid.
So they've got an idea of the fact that a
reactive enough metal is going to produce hydrogen with some
acid. They've got the idea that
different acids tend to behave in the same way, as long as they're the same
concentration
and they've got the basic idea of
displacement reactions and of course one of the main problems with the displacement
reactions when you want to talk about
aluminium is the fact that it's protected by this coating of aluminium
oxide and that prevents
us from seeing it's true reactivity. So the first thing I'm going to do is
just put some aluminium in the bottom of these 250 ml
flasks. I've got about 20 cm here by
about 8 cm, roughly the diameter of the
flask and I'm just loosely
rolling that up just enough to help me to get it through,
but I'm not crushing it too much so that I'm
preventing the solutions from getting in and doing their business,
and a splint or something can just
help you get that so it's flat down on the bottom. It doesn't have to be pressed
but down
towards the bottom of the flask there. Okay so we've got
two lots of aluminium here and to my left and to my right
we've got some acid solutions
okay and we've got fifty ml of each. On my left we've got sulfuric acid
and that's going to be paired up with copper sulfate, we've got the sulfate element common there,
and then here we've got hydrochloric acid and that's going to be paired up with copper
chloride,
again the chloride element being common and just trying to keep things
fair, we've got 0.5 M sulfuric acid
whereas we have 1 M hydrochloric acid so
to all intents and purposes when I'm explaining this to students, we've got the idea that
both of these contain the same amount in concentration of
acid particles effectively. So,
if we dissolve up the
copper sulfate. There's about 3.5 g here, it doesn't
really matter too much
it's just a ballpark figure and
just give that a swirl. I'll leave that for a second while I do the same thing for the hydrochloric
acid and the copper chloride.
Chuck it in there, it's all good.
Give that a clean.
You might want to point out or just spend a minute or two explaining the
slight difference in the colours of the solutions.
But the key thing to point out with both is that they are
coloured, they have blueish or bluey-green colour,
and that's telling us that we have these copper ions
in solution. That's going to become more important later on.
I'm going to just add these in here, starting off with the copper sulfate.
So this is going to be our control, our benchmark for the students
to think about. So you may want to ask them what do they think is going to
happen.
We might expect for a fairly reactive metal that we might see
hydrogen bubbles. We might see some kind of a displacement reaction coming along.
What will that look like in terms of the products that are being
formed? Exciting... or not.
Okay so of course here we have the
disappointing reaction that
students might see that it's not really going to help them to see the fact that
aluminium is actually pretty reactive.
If this were magnesium we'd be seeing quite a lot of hydrogen being produced certainly
and the metal dissolving, none-of-the-above happening here.
What we need is something that's going to help us to break down
that aluminium oxide layer. So in here we have,
the only difference being, we have chloride ions in with the acid
instead of the sulfate ions and the chloride ions are going to be able to
negate the aluminium ions on the aluminuium oxide surface
and just help to break down that protective layer exposing the aluminium underneath
and we'll get an idea of just how reactive it is.
Similar kind of expectations about what we might expect to see,
maybe we might see copper being produced from the displacement, maybe you'll see
bubbles of gas and of course when you do this you'll know that
there will be gas being produced. So I tend to
have a lit splint ready
so that I can chuck this in at the point where I think
we're reaching the peak amount of gas and
the result is actually really pleasing. Let's have a look.
Okay. Splints ready. If I hold this up here
you can probably hear that there's already gas being produced.
I'm going to give that a few more seconds so that it's definitely going, there we go
and we get this really cool
hydrogen flame picking up of course the green colour of the
dissolved copper and that will
keep going for about, on a good day, at least 20 seconds or so.
Keep the splint handy because you may
get a second go out of this
If you find,
if you find that the flame is self- extinguishing and it's just sort of burning inside
the flask, you're not going to get anything
else coming from this. If you have that as a problem
from the off with your flask it's worth trying
a little bit more copper chloride in the mixture to start off with,
but you may if you're quick about this get a couple of
goes before the reaction reaches a point where it's no longer going to be able to
sustain that flame.
So things to point out of course,
we now have a flask that's very hot.
It's okay to handle for a few seconds but there's definitely a
lot of heat being produced here, both from the flame,
but it's not just from the flame it's the
base of the flask is pretty warm as well. Okay,
and we've got these really obvious colour changes here, as we produce
copper in the displacement reaction, therefore inferring that our
aluminium is dissolving into solution, as aluminium chloride
in one of a number of forms, it's likely, and of course
the fact that the solution is going colourless
and after a minute or two it will be almost completely
colourless, if not colourless. So a really nice,
simple demonstration to work with. If you dim the lights you get
such a nice Willo the Wisp type effect, definitely worth trying with your
students.
Give it a go.