This is probably one of the first neutralisation reactions your students will explicitly meet.
And certainly when it comes to ionic equations we've just seen an example of the first one
that they should be able to explain.
But how do we know that the ions are there?
Sure you can add an indicator, but we first have to sell that this tells us about the concentration
of hydrogen ions and part of the charm of the way that indicators and pH work
is that they respond in a logarithmic rather than a linear way as the reaction proceeds.
One thing that does scale in a linear fashion, at least at lower concentrations,
is the conductivity. I fine-tuned this demo specifically to show what happens
to the ions in a neutralisation reaction. And there are lots of tips along the way for getting the best possible
results on any budget so watch carefully and you'll see how to get beautiful results like this.
Before we get going, let's clear up some linguistics.
I recognise that it would be more correct to talk about the hydronium ion H3O+ than it would be
to talk about H+ or protons floating around in solution, but, especially given the level that
we're working at, and that exam specifications and textbooks commonly use the same simplification,
I hope everybody's cool with me just saying H+ for today.
At least until next time when
we'll need a more subtle model to explain some of the effects that we'll see.
What we want students paying attention to at this level is that acids are solutions containing hydrogen ions and
alkaline solutions are those containing hydroxide ions. Specifically when they write the formula for HCl,
we want them to be able to visualise the ions dissociated into solution, not as molecules.
So let's just reinforce that for a moment by hunting for ions with the CLEAPSS conductivity tester.
Brought up with warnings of electrocution hazards, students will often instinctively feel
like deionised water will conduct, but in reality with an 'electricity' as movement of charge model,
it's not hard to see that while our water molecules can move,
if they're neutral we have no current. Meanwhile in solutions of acids and alkalis we have ions,
which when they move through the solution will be able to carry a charge conducting electricity.
But, we have a problem with the other spectator ions in solution.
Again, once we've got the basic
set-up working here we can leave them in, but for now, let's use the reaction between barium
hydroxide and sulfuric acid to take care of that issue.
At this level it's possible students
will already know how a solution of barium ions can test for sulfate ions by forming the white
barium sulfate precipitate. This removes those ions from solution allowing us to focus just
on the hydrogen and the hydroxide ions.
I've used barium hydroxide octahydrate making up my solution.
Now the classic recipe for making up barium hydroxide solutions can be a bit of a faff
due to its propensity for forming carbonates, but I've found there's no need to go through
the rigmarole of boiling water and using guard tubes for this reaction to work nicely.
Any insoluble solid can be filtered off and you'll still get a solution close in concentration to the
bench acids you have to hand, which, let's be honest, will themselves vary a little in their accuracy.
Do be aware that depending on your situation it's possible that the main use of barium hydroxide
might be to show its endothermic reaction with ammonium salts. We've already covered this demo in
a previous column (link below).
As we'd already had a bottle that had been open for a while before the pandemic,
making up a solution for me here now in late 2021 with that was a real problem initially.
On the plus side, you do get to see with this reaction just how bad any old stock actually is,
because cloudiness before you filter can make it seem like your hydroxide is a lot more carbonated
than it actually is. We'll see that in our results later on.
We use these squeezy dropper bottles
holding about 100 ml of solution. They're fantastic for microscale work but are large enough
that you don't have to constantly refill them, which if you use the tiny supplied funnels is
a bit of a nightmare. Shout out to Mark Cooper, my former technician, for finding these for us.
Luckily the Matson-sized syringes that we use a lot recently can help us to load them up in a jiffy.
I've left some solution out for about 20 minutes exposed to the air here and
this is what it looked like. Meanwhile the dropper bottles without any guard
tubes have lasted for well over a month so far with no signs of precipitation.
Once you have your solutions, you need to think about how you're going to show their conductivity.
The classic Nuffield technique suggests a 12V bulb and 1M sulfuric acid.
Even then the bulb's going to be pretty dim. I'm going to suggest some alternatives, because
while I recognise the idea was to use an AC supply to minimise electrolysis, it really
doesn't have that much of a huge effect on this demonstration. It's also a lot of stuff to set up.
The CLEAPSS conductivity indicator is remarkably sensitive and there's not
a huge amount of gas production when you consider the volumes that we're going to be working with.
If you want to be quantitative, then you can use a multimeter on resistance mode. You'll probably get
values in the tens of kilohms for this demo, but a scale switch is sometimes needed mid-demonstration.
There are cheap pen style conductivity probes available online for under a tenner.
These are sometimes sold as water testers and I've chosen concentrations here for the demonstration
that we're going to do in the final version to let any of the above cheap options work,
because the 10 microsiemens per centimetre resolution is plenty for this demo, as well as the one that we're going to do next time.
If you have a little bit more room in the budget though, a data logger probe
with one microsiemens per centimetre resolution is a worthy investment as you
can produce a graph on screen to analyse and take all of the concentrations describing here,
divide by 10 and you'll get even better results with solutions that are even safer to handle,
easier to dispose of and behave in a more ideal way without the ions getting in each other's way.
Whatever you decide to use, demonstrating the conductivity of the solutions as we did earlier is
a great opportunity to adjust your meter settings so you don't get a false start, because your sensor
is set up incorrectly like happened here.
Now anyone who's generated a pH curve by opening a
burette over a pH probe in solution will be tempted to just do the same thing as they always do
for that with this demo. I know I was when I first trialled it so let's take a look at what you get
if you run the experiment in lazy mode.
I've got 25 cm3 of 0.1 M sulfuric acid here and I'm about to empty 50 cm3 of barium hydroxide into it of the same concentration.
In theory we'll see the conductivity drop as we progress to a minimum at the
equivalence point and then it will increase on the other side again.
The conductivity probe gets dimmer and brighter the way that we'd like, but you can see that the
graph can be a bit of a mess. Incidentally if you have more precise kit that tenfold dilution will
largely fix this issue for you as well. We do get a broadly v-shaped graph the way that we would like,
but notice the way it curves. Some students might already be able to suggest what's going on here.
We're diluting the solution so that at the end point the ion concentration has actually dropped threefold.
This isn't a problem when we generate pH curves, because the logarithmic nature
compresses out most of the dilution. The other thing that you might notice
is the minimum lagging after the addition of the equimolar quantity of reactants.
Indeed here we saw actually the graph was a real mess before we got to that point.
This can be because poor mixing and also it takes a few moments for the precipitation to complete.
It would be better perhaps if we had more volume for us to play around with.
If you're using some electrodes you'll also notice that they get submerged to a different level as well.
So let's see if we can get the demo working a little bit better.
This time we're going to run the tap a little slower and we're going to begin by adding some water to the acid.
It doesn't matter how much as we still have the same number of moles reacting this way.
I'm going to go for a tenfold dilution so let's top it up to around 250 cm3.
We also get bonus points for being able to get the probe fully into solution and
a comparatively small rise in the level of the liquid. This hugely reduces graphing artifacts like we saw before.
Just look at that beauty.
We still get a bit of a stretch from the dilution, but now the slowing tap on the burette as the liquid
runs out is the major issue.
Before we talk about a final improvement, let's briefly discuss order of addition.
Some people right now are going to be watching this and weeping as I put alkali into the burette.
It really doesn't matter though. There are advantages and disadvantages whichever way you do it.
I've chosen to do it this way to fall into line with the demo that I'm doing next time as well
and at higher concentrations you might actually get some precipitate in the burette, but with
the 0.1 M hydroxide that I've been using here I've never seen any tap or jet blockages.
By having the hydroxide in the burette you avoid stirring in a load of atmospheric CO2 or any
CO2 that's already dissolved in your deionised water at the start. And if you add in the opposite
direction you do benefit from having all the barium precipitated out so if you do it
this way, it is worthwhile adding a little bit more acid before filtering off the solid for disposal.
So it's swings and roundabouts really.
Most demos would finish here.
There is a wealth of detail
revealed in the shape of this graph, but of course, you can go further and plot against
volume instead of plotting against time. This is going to be more of a practical than a
demo because of the time constraints but, just look at the beautiful curve that you can get.
By applying a volume correction factor this gets even nicer.
It's at this point that I'll remind you
all of these values were collected with my speedy dissolution method mentioned before.
Boiling deionised water and guard tubes are overkill for this reaction. Beady eyes will have spotted the curvature
around the equivalence point associated with solubility effects and perhaps more interesting
the gradients aren't identical on either side of the equivalence point, but this is where I
have to leave you on a cliffhanger.
In each column I only get 700 words to play with and this script
is nearly three times that length already so tune in next time to have your minds blown by how much
detail we can see in the curves produced from the neutralisation of strong and of weak acids.