Electrochemistry ...
it really has the power
to get your heart beating
This month's demonstration can be used to illustrate simple energy transformations
to younger students with an incredibly weird kind of effect that's really captivating
or for older students you can use it to test their understanding of the electrochemical series
and their abilities to make the predictions that they need to be able to do
in order to work out what's going on with redox reactions.
Now, in a previous demonstration, the 'Nailing corrosion' demo, we used our simple understanding
of metallic bonding to test out what might be going on when we put a piece of metal into water.
So let's imagine for example you stick a piece of iron into some water
At the macroscopic scale, initially you don't really see anything taking place
but at the microscopic scale we might begin to visualise
the odd metal ion dissolving into solution.
Now that's going to leave a negative charge behind on the surface
to which the dissolved ion will be attracted.
Now very quickly obviously an equilibrium will be reached here
because each time we try to dissolve a positive metal ion,
we're having to pull it away from an increasingly negatively charged surface.
Essentially, this is what our standard electrode potentials are describing to us.
For metals whose ions are poorer oxidizing agents than water is, a stalemate will quickly be reached.
Something else needs to be able to accept those electrons from the metal if it's going to continue to dissolve.
But is there some way of seeing the effect of what's going on here on the metal?
Now not very easily if you've got a solid piece of metal, but if you've got a liquid then it's possible
that that metal shape can actually respond to these differing forces that are acting upon it.
Let's imagine a droplet of liquid metal.
To keep things simple, we'll assume any forces with surrounding molecules and negligible.
At its surface, we have attractive forces pulling towards the middle and around the edge of the droplet
but none acting away from the droplet.
This leads to surface tension.
The droplet will tend to hold itself together when we place it onto a surface with which it has poor compatibility.
Compare that with a droplet with an excess of negative charge surrounded by positive ions.
An increased repulsion will weaken those forces while also increasing forces acting away from the center of the droplet.
Therefore, we might expect the droplet to sag more under its own weight than it did before.
Now, if we were to give those excess electrons somewhere to go to,
we might see that metal droplet pulling itself back together again.
This is where gallium and mercury can come in handy.
Now, in a previous column we've already had a look at the mercury beating heart
but obviously mercury isn't the most pleasant of materials to use and dispose of,
so I decided to see whether or not could I come up with a gallium modification.
I'm going to contain the reaction in the sample vial. This means that cleanup is super easy,
you reduce the number of surfaces which come into contact with the gallium
and you can easily reuse it for future demonstrations.
Glass or plastic vials work fine,
but never put liquid gallium into a glass vessel if it has the chance of wetting the complete base
because gallium expands upon freezing and that could lead to breakages.
Now, literature descriptions of this experiment call for the use of six molar sulfuric acid,
sometimes higher temperatures, even the use of additional oxidizing agents,
and all of those will certainly get your gallium heart throbbing.
However, I found that actually it's not necessary if you want to show this process taking place.
You can go for higher concentrations, but i'm using simple 1.4 molar sulfuric acid,
the sort of stuff that you'd have around on the bench and is still within the classification of being irritating rather than corrosive.
As usual, do check out the Education in Chemistry website for all of the additional safety information, technician sheets, additional resources, etc.
In terms of the gallium, you might be able to come across gallium chips like these
or it sometimes comes in little bottles like this.
Obviously the advantage of buying it in chips is that you can actually extract it from here without having to melt it ahead of time.
I'm finding that actually one to two grams, which is about two of these chips
is about right for this demonstration to get going.
If you are going to be potentially coming into contact with the gallium, you should avoid obviously skin contact
and if skin contact is likely, make sure you're wearing gloves.
Otherwise, make sure you're wearing eye protection for this demonstration.
The gallium melts around 29 degrees celsius
so I've got a water bath around 40 degrees celsius set up here to keep the metal well above that.
Again, the temperature isn't critical.
No, I want to add enough sulfuric acid to cover the metal
and if you're going to leave this on show for a while, you might want to add a little bit extra to keep the reaction going.
So right now we have our saggy metal droplet forming in the base of our vial
and the hydrogen ions in the solution are making for a better oxidizing agent than water to really emphasise the effect for us.
In theory, we ought to be able to see the effect of the repulsive build-up of negative charge
by giving it somewhere else to go.
And this is where my iron nail comes in.
I figured out that a syringe without its plunger makes for a great holder for an iron nail like this.
You can get a good grip on it with a clamp and move it precisely to a distance where the gallium can come into contact with the iron,
dump its excess electrons there, and retract away before repeating the process.
Notice how the reaction proceeds the fastest where there is the highest charge density
on the rough surface of and pointed end of the nail
rather than the rounded surface of the gallium droplet, even though the gallium is more reactive.
The droplet is almost acting like the rounded surface on top of a Van de Graaf generator.
Now, there are a lot of variables that you can play around with here
if you want to get your gallium heart beating in different ways
Confusingly, some of the descriptions in the literature also call for the use of acidified dichromate.
Now, with the addition of the nail as an electron sink, then that's not really necessary.
However, other oxidising agents like dichromate certainly can get this beating effect to take place without a nail kicking around.
Now, I found that actually hydrogen peroxide is also a good enough oxidising agent
for us to see that effect taking place
and obviously comes without the additional environmental and other hazards associated with the use of dichromate.
In the 'Nailing corrosion' demonstration we saw that when reacting with a solution
there tended to be distinct sites of oxidation and reduction with electron transfer happening through the metal to the cathodic site.
So when the metal can change shape and respond to the resulting changes in the forces acting across it,
exacerbated by the production of gas bubbles,
some really intriguing effects can be produced.
Perhaps if you get some really cool effects with your gallium beating heart,
you might be able to share them with us on any of the Education in Chemistry social feeds.
We really enjoy seeing what you've been getting up to.
Otherwise, see you again in a couple of months.