So when I was doing my teacher training, I showed up to a job interview
where the requirement was to deliver a lesson on the formation of ionic
compounds and, as part of the lesson, I decided to demonstrate the reaction of sodium with chlorine,
and maybe that was a bit ambitious for an interview lesson at the time but
now I've come to a realization that neither the way that I chose to do it, nor the reason that
I chose to do it were great ideas. So, let's have a chat about that shall we.
First of all, here's how not to do this demonstration. With the use of a deflagrating spoon.
You will, of course, get a reaction but the problem is that what is actually reacting. Most likely, what
you're going to see is a lot of brown smoke as the chlorine reacts with iron in the spoon
to produce iron chloride. When your students have a clear idea that table salt is supposed to be white,
this can put a bit of a dampener upon things.
I've even had broken gas jars in the past where I've
switched schools and hadn't really noticed that the diameter of the gas jars was a little bit more
narrow and that deflagrating spoon just tilted across made contact with the glass and produced
a crack, so definitely not something that I'd be looking to do again. I actually tend not to do this
reaction above key stage 3 because it can lead to all kinds of misconceptions about ionic bonding
being somehow the result of electron transfer rather than simply an electrostatic attraction
between oppositely charged ions. But, it can make for an excellent demonstration to illustrate how
different the properties of a compound can be from the properties of its constituent elements.
So, if you want to see those properties, here's a better way of doing that demonstration.
First of all, I recommend working in a fume cupboard. You're going to need to wear splash-proof
goggles and your students will need to be wearing eye protection. The nice thing about working in a
fume cupboard, where possible, is that with the sash lowered, no additional safety screens are required.
You're going to need to check that there is a distinct green colour to the contents of the gas
jar to ensure that the partial pressure of the chlorine is sufficient for this demonstration.
And rather than using a deflagrating spoon, we're going to use a brick. Before getting your sodium out
do make sure that the diameter of the gas jar is less than the width of the brick to ensure
that the chlorine remains inside during the reaction. Cut yourself a six to eight millimetre
cube of sodium, pat it dry with an absorbent towel or lab roll, and point out to students
the lustre as it fades as well as the malleability of the material. There's absolutely no reason to go
any larger than this size because, as you'll see later on, usually only the surface of the sodium
reacts substantially with the chlorine, leaving the middle unreacted, for disposal later on.
Have the gas jar of chlorine nearby already inverted so that you only need to slip off the
cover and place it on top when the time is right. Heating can be a little bit tricky. You need to
have the flame hot enough to get the sodium to ignite, but you'll also need to ensure that the
gas flowing out from the burner does not cause the molten ball of sodium to roll around too much.
You'll need to play around a little bit with angles to get this going just right
Once the sodium is lit, simply place the gas jar on top.
An extra pair of hands can be handy here. And then lower the sash and enjoy the show.
Once the reaction is complete, you may wish to wipe the residue with isopropyl alcohol
because almost certainly there will be unreacted sodium on the brick the whole brick can then be
submerged in water. Now for this video, I had to do five different takes and so by the time I cleaned
up at the end, there was definitely enough sodium to ignite, even after a good scraping
and a wipe over with the isopropyl alcohol. As such, it's worthwhile doing the clean-up
and the submersion step in the fume cupboard still, so that the sash can be lowered if there
is enough sodium there to ignite. Obviously, do keep wearing those splash-proof goggles.
I never do this demonstration anymore when I'm teaching about ionic bonding. In the accompanying
article link below, I do suggest some alternative routes for teaching structure and bonding which
aim to avoid misconceptions and emphasise the point that ionic bonding is simply an
attraction between oppositely charged ions. I mean, if you think about it, most of the salt that we're
going to add to our food has simply come from sodium ions which were themselves in other ionic
compounds moved over the years through the water cycle between rocks and oceans, perhaps at some
point involved in the firing of a nerve in the brain of a sentient being as it tries to explain
where this stuff is coming from. It's entirely possible, in fact, that since that sodium nuclide
was generated in the death of some high mass star billions of years ago and shortly picked up
afterwards some electrons, it may never actually have been a sodium atom in its entire existence.
Still, I do think this reaction is worthwhile
for demonstrating the properties of the white stuff that we season our food with every day
and how that can be so different from a toxic gas or an explosive metal.