When we first start learning chemistry, a way to distinguish between a physical and a chemical
change often emphasizes the production of new material by pointing out that a
process cannot easily be reversed – at least when say, compared to melting or evaporation.
So as students begin to develop a more refined understanding of chemistry and
start thinking about reversible reactions, having some visually appealing changes to
show is always handy, and today we're going to be looking at a great example.
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I've modified this reaction a little from some other variations that are out there to, I hope,
optimize it for a classroom setup. As always, you can find links to technician's notes, more
information about the reaction etc, on the Education in Chemistry website. The link will
be beneath. Like most other variations out there, we are still using 10 vol (3%)
hydrogen peroxide and if you're new to the game let me reinforce with you that it is really
important that you make up 10 vol solutions of peroxide fresh. There's a stabilizer that is added
to stock, like 100 vol solutions that you'll have, and that stabilizer becomes less effective as you
dilute those down. So definitely make it very close to or ideally on the day of the reaction.
I've made up around 100 ml of the 10 vol, because the maths is easy and that's going
to be plenty for many reactions and cycles that we're going to be doing today.
If you're a CLEAPSS member, you should consult RB045 for diluting down your peroxide.
We're going to be using a 1 M solution of potassium sodium tartrate. You can of course make
up a stock solution, but just in case you wanted to make just enough for one demonstration, I'll
show you how to do that too. I've got 8.47 g of the tetrahydrate salt here and I'm going to
dissolve that in 30 ml of water to make up a 1 M solution. Let's talk glassware. The reaction
starts with a small volume of liquid and you're going to add to this over the demonstration to get
it cycling over and over. So that means that – that small volume of liquid – to maximize its visibility
you want it to be as tall as possible. So, if you do have a 250 ml tall form beaker, I would definitely
recommend using that. I don't – I just went for the regular one, and that works totally fine.
To simplify some later steps now, I'm popping the tartrate into a 1 L beaker and I'll
put a white tile underneath to aid the visibility of the reaction. I'll measure out 15 ml of the
10 vol hydrogen peroxide and put that into a test tube that's also in the 1 L beaker.
These two solutions will be warmed up before the demonstration, but keep the rest of the peroxide
aside at room temperature. Finally, if you pour a few ml of 0.5 M copper sulfate solution
into a beaker so that you can easily suck up 1 ml with a dropping pipette or a small syringe
that would be handy. Using 0.5 M as opposed to 1 M copper sulfate means that you don't need
to measure volumes less than a milliliter as you might find if you're working on other variations
at this scale. The copper sulfate solution does strongly affect the colour and the speed of the
demonstration, so working with 1 ml of 0.5 M felt like a good compromise.
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Don't forget to wear eye protection. The reaction works best between 60 to 70 degrees celsius, but
the first change is quite exothermic, so aim to start between 50 and 60 degrees celsius. My way of doing this is
to pop a thermometer into the tartrate solution, add hot water from a kettle to the 1 L beaker
and then when the tartrate has reached 50 degrees celsius, you can then remove the 1 L beaker, placing
the tartrate beaker directly on the white tile, which helpfully has also warmed a little as well.
To initiate the reaction, add the 1 ml of your copper sulfate solution and pay
attention to the colour. We're about to heavily dilute the solution, but ...
... it's still a vibrant blue. We must have made a different, more darkly coloured complex.
Bubbles of gas are being released now, and then, depending on the temperature you began with, after about
10 to 15 seconds, the blue colour fades, and a golden copper(I) complex is produced, while at the same
time the temperature jumps right up to around 70 degrees celsius. Initially, the copper acts as a
catalyst for the decomposition of the hydrogen peroxide. As the peroxide concentration drops
relative to the concentration of the tartrate, the copper itself becomes a reactant and is reduced to
the copper(I) complex. But with the concentrations and volumes chosen only a small amount of
the tartrate is consumed on each round. That means that further additions of peroxide can re-oxidize
the copper back up to copper(II) and start the process all over again. Temperature management
is crucial for a good effect. That's why I kept the rest of the peroxide at room temperature.
If it was heated, then the next exotherm, when it comes along, would potentially heat the solution above 70
degrees celsius, at which points the cycling of the reaction does somewhat become inhibited. We want to
keep it, if possible, somewhere between 60 and 70 degrees celsius by riding those exotherms as we
go along. Adding in 10 ml of the peroxide will allow us to drop the temperature down
from near 70 back into that sort of sweet zone and give us a little bit of buffer from above.
You can repeat this process over and over again with 10 ml additions of peroxide until, if you
haven't got bored before then, eventually the tartrate concentration just drops too low.
Variations on this demonstration call for larger volumes of peroxide. I've reduced
the relative amount of peroxide to reduce the latency period, and that means that, you know
if you get distracted by a class, you're less likely to see your temperature drop
too far below that sweet spot and find that the exotherm never comes along to rescue you
and bring that temperature back up and keep things going at a reasonable pace.
If you work with fresh peroxide, if you get that initial solution when you're adding your tartrate
in with your peroxide between that sort of sweet spot 50 to 60 degrees celsius, this is pretty bomb proof
and very worthwhile doing with a class. You can easily show how a chemical reaction with
colour changes, precipitation and gas evolution can be made seemingly to reverse. And if you're using this
to lead into equilibria, you might then think about asking your students whether or not this
really represents a reaction that has reversed, or is there something else going on, preparing
them to start thinking about the importance of a genuinely closed system. See you next time.
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