Last time on Exhibition chemistry we were left with some unfinished business.
We were able to get as much nuance as was needed from a neutralisation reaction to help students
visualise the ionic equation in action, following the disappearing ions with a conductivity probe.
Today we are going to do a deep dive on strong and weak acids and see just how much
a conductivity graph can reveal to us about the make-up of solutions at different stages of the reaction.
Now if you’re coming to this video without having seen part 1 in last month’s column,
you really need to go there first. I explain in detail why I have chosen the concentrations that
I have and much of that won’t be repeated in detail here. I’m going to assume that you’re up to speed with
the basic set-up, the choice of equipment and the logic of the dilutions that I’ve used.
We looked at sulfuric acid reacting with barium sulfate and we precipitated out the spectator
ions to hone in on the H++OH- makes neutral H20 aspect of the reaction. I apologised in
advance for talking about acidic solutions as containing hydrogen ions, so now let’s
apply a little more complexity and nuance for an older audience. I’m gonna repeat the set-up
here with a solution of 0.1 M hydrochloric acid. Don’t forget the important dilution step.
The first thing to point out is the brightness of the bulb. Our acid is fully dissociating in
solution. The conductivity is high. This time, instead of adding barium hydroxide,
we’ll track the far more common sodium hydroxide titration. Opening the tap all
of the way is fine this time because we don’t need to wait for the precipitation to complete
as we did before. In this early part of the titration, a pH curve will barely budge, but by
tracking conductivity without those logarithmic effects we get an immediate and linear drop.
This time, we won’t get to 0 as there are sodium ions and chloride ions remaining in solution.
Just like before, as we add more sodium hydroxide to the neutralised solution, the conductivity
rises once more, but not at the same rate. We saw this in the barium hydroxide titration.
Before we talk about gradients, though, there’s a more immediate question that we ought to be
asking ourselves: why is the conductivity dropping at all? I mean we’re starting with a solution of
positive and negative ions from the hydrochloric acid. At the equivalence point, we have a solution
of positive and negative ions from the sodium chloride. The chloride ions remain in solution
while the hydrogen ions are replaced as they neutralise out one-to-one with the sodium ions.
The trick is, of course, that there aren’t naked hydrogen ions in solution at all.
Last time we parked this hydronium problem and now it’s time to revisit it.
When a potential difference is applied across solutions of aqueous ions like sodium or chloride,
the whole ion needs to drag itself through the solution, commonly bringing with it the hydration
shell which adds significantly to the bulk. We saw in the movement of ions demo from November
21 that this bulk motion is deceptively slow. But acids don’t form hydrogen ions in solutions
divorced from the surrounding water molecules. Rather they form hydronium ions – the conjugate acid of water.
Instead of the whole ion moving around, the extremely labile hydrogen ions hop
from one water molecule to another allowing for a vastly superior conductivity than other ions.
In the same way, the hydroxide ions could be thought of as the mirrors of this process,
hydrolising neighbouring water molecules so that an absence of a proton hops through the solution.
So solutions containing hydronium or hydroxide ions will conduct much better than other ions
of the same charge. Leading to the decrease towards the equivalence point. The spectator
ions themselves will contribute in different ways to the gradient as even for ions of similar charge
differences in size will impact the hydration shell and lead to corresponding differences in
the gradient of the lines, coming in to and out from the equivalence point.
This was particularly noticeable in the titration that we did last time. Just one curve here has hugely developed the model
that we have for the behaviour of the particles in solution from this simple titration.
But, post-16 courses and even some 14–16 courses will also require that students differentiate between
strong and weak acids in terms of their degrees of ionisation or dissociation in solution. So
let’s see if we can see those differences reflected in our conductivity.
I’ve set up the same titration as before, but this time with 0.1 M ethanoic acid. If you concentrate on the timing you can compare both lines on the same axis. But, distracted by the filming process,
I didn’t perfectly time it here. But it’s good enough to see what we want to see.
The first thing to point out is the dim light on the conductivity tester at the start.
The value of the conductivity of the solution is correspondingly much smaller for the same
concentration as it was with the strong acid. If you’re doing pH curves, your students might expect
an initial jump up in pH as a strong base is added to a weak acid. Here we see a drop in the highly
conductive hydronium ions reflected in a drop in conductivity, but the graph is much more shallow
and quickly switches to a positive gradient. The drop in the already low concentration of hydronium
ions is now being more than compensated for by not only the addition of the sodium ions, which we had
last time, but by locking in the ethanoate ions which originally would have been overwhelmingly
floating around, associated in solution with the hydrogen of the carboxyl group attached.
The conductivity rises to the equivalence point, where once we run out of acid to mop up the added
hydroxide ions, the conductivity rises more rapidly for the reasons explained before.
Notice how the gradients of the lines after the equivalence point match each other so beautifully,
demonstrating the equivalent role of the hydroxide ion in the solution.
There’s a weird cognitive trap waiting that I might suggest you probe with your students if
doing this. What do you think is going to happen to the conductivity of the solution
now if we add some more acid to it, neutralising the excess sodium hydroxide?
Let’s give it a try with the undiluted ethanoic acid. Your gut instinct and that of your students
might be to think, well the conductivity will go down again. But, of course, all we’re doing
is making more ions. The conductivity is going to continue in fact to rise, just not as steeply as it did before.
So this is where I pull everything together and show you the pièce de résistance.
Just like I mentioned in the previous video, you can, with a bit more patience,
and probably the input of some student-led enquiry, rather than as a demonstration,
collect data against volume, rather than against time. Applying a volume correction to the
conductivity values, you can get these curves for the three reactions that we’ve looked at.
These data were collected with concentrations 10 times less than those outlined before
as I have the resolution for it with my conductivity probe.
It makes for an easier clean-up and I also benefit from minimising ion-ion effects that
cause conductivity to change in a non-linear fashion as you go to higher concentrations.
Notice the way that the sulfuric acid is almost twice the conductivity of the monoprotic HCl
which is, in turn, 10 times more conductive than the weak acid.
This is from a rough titration at home over the holidays with bench solutions using cheap jewellery balances – thanks
Bob Worley – for gravimetric dilutions, and I still had beautiful trend lines crossing within
a couple of microsiemens per centimetre of the conductivity of my deionised water alone, and
really close to the halfway point of the burette.
One of my favourite Bob Worley-isms is the phrase, ‘in a little, you can see a lot’.
I was sick to death of titrations by the time I finished my
second year of uni. Who would have thought, 20 years later, I would be geeking out over them.
I’ve found the journey to this graph deeply satisfying, so I hope that some of you
might feel the same way and use it to stimulate some cool discussions with your classes too.