Warning: This experiment handles corrosive acids, high temperatures and pyrophoric substances.
Wear gloves and have appropriate fire safety protocols.
Greetings fellow nerds.
In this video we’re going to make ferrous oxalate and in turn use that to make pyrophoric iron.
Now as you know iron reacts with air as easily seen by rust which is iron oxide.
This process is very slow but if given enough time will completely destroy the iron.
Now bulk iron doesn’t burn but much finer iron like steel wool does.
The greater surface allows it to react air fast enough that the exothermic reaction is self sustaining.
Now extremely fine iron particles also burn and this can be seen if you grind iron.
The particles spontaneously ignite on air.
Now we’re going to try and ramp that up that demonstration by making such particles in bulk through a chemical process.
First we start with a source of iron. In this case i’m using 5.5g of iron nails.
The nails are in a flask and i’ve attached a condenser as a reflux column.
This is because we’re going to boil the mixture in a minute.
Now we add 30 mL of 10mL hydrochloric acid. The exact concentration isn’t critical.
We just want enough acid to dissolve our iron.
Now we add 70 mL of water.
Now I often get people complaining you should never add water to acid.
That rule is really only important for anhydrous acids and doesn’t really apply to a solution of hydrochloric acid.
Anyway, once the water is added, turn on the heating, the cooling water to the condenser and the stirring.
A great thing about using nails for this experiment is that they are magnetic and behave as their own stir bar.
So you can use a magnetic hotplate stirrer and not need to put in an extra stir bar.
As for the heating, you just need to get it up to a light boil.
If you see water refluxing in the condenser then you’re at the right temperature.
We installed the reflux column earlier to minimize losses of the hydrochloric acid as it boils.
What’s happening here is that we’re reacting the hydrochloric acid with iron to make hydrogen gas and iron (II) chloride, also known as ferrous chloride.
If you already have an iron (II) salt like ferrous sulfate or ferrous chloride
then you can skip this part and just dissolve the 0.1 moles of the salt in 100 mL of water to perform the next steps.
I didn’t have a salt on hand so i’m making my own here.
Also, since some ferrous salts can oxidize in air under certain conditions,
making them fresh like we are here tends to work better for this experiment.
This ensures we have only iron (II) ions rather than iron (III) ions that we don’t want.
Anyway, continue heating and stirring until all the iron has completely dissolved.
At that point, turn off the heating and let it cool.
Now we get a beaker containing 13 g of oxalic acid dihydrate.
On top of that put in a funnel with filter paper and filter our ferrous chloride solution directly into it.
The filtering is optional as the undissolved impurities are usually very minor.
But I’m the kind of chemist that likes to work with pure chemicals.
Now what’s happening in the beaker is that the iron (II) chloride solution is reacting with oxalic acid to produce hydrochloric acid and iron (II) oxalate,
also known as ferrous oxalate.
Now ferrous oxalate has a very low solubility and precipitates out as a bright yellow powder that you can see in the beaker.
Now when the filtering is complete,
drop in a stir bar and stir and heat the mixture on the hotplate stirrer until it just starts boiling.
We’re doing this to force the reaction forward.
Now you might be asking, can you use iron (III) chloride or ferric chloride for this procedure?
I actually tried that and nothing precipitated even after boiling.
Perhaps ferric oxalate is very soluble and to recover it i would have to let the solution evaporate.
I didn’t have enough time to wait for that so I went with making iron (II) chloride instead.
Anyway, once the mixture starts to boil, turn off the heating and let it cool.
Here we are after it’s cooled.
You can see the settled ferrous oxalate.
You can decant off the supernatant but i recommend filtering it for higher purity.
Be sure to wash down any residues with water.
After it’s filtered, wash the residue again with water to insure we have removed all soluble impurities.
Then let the residue dry overnight.
Here we are the next day. Collect the residue in a clean container.
And there you have it, ferrous oxalate.
It’s best to store it as this form as it’s relatively stable even on air.
To make pyrophoric iron we take a sample in a glass container and seal it with a long exhaust tube.
This tube leads outside or to a fume hood.
Now we get a propane torch and start heating our mixture.
We have to get it hot enough to decompose and I found a hotplate simply wasn’t strong enough.
What’s happening is first the ferrous oxalate is liberating its water of crystallization
You can see condensing at the top of the flask and in the exhaust tube.
After it’s completely dry, the anhydrous ferrous oxalate is being converted to iron (II) oxide,
also known as ferrous oxide, and releasing carbon dioxide and carbon monoxide.
This is why we have the exhaust tube lead outside or a fume hood to protect ourselves from the carbon monoxide gas.
The exhaust tube also prevents oxygen from reaching the mixture and oxidizing it before we’re able to finish our experiment.
Now the iron (II) oxide produced is very reactive and some it disproportionates to pure iron and magnetite.
The iron particles are extremely small and highly reactive and will spontaneously ignite on air.
This property is called pyrophoricity.
When the ferrous oxalate has completely turned black the reaction is complete
and we can demonstrate the pyrophoricity by scattering the iron on air.
I’m going to do so over this sheet of aluminum foil.
Now turn off the lights and quickly open the container and dump out the iron.
You can see the bright red and orange shower of sparks from the iron reacting with air.
The larger chunks will glow for several seconds as the air diffuses into the interior.
Afterward we have this ash of iron (III) oxide, also known as ferric oxide.
You know, this iron oxide would make excellent thermite.
It’s pure and extremely fine, even finer than my laboratory grade iron oxide.
Anyway, there you have it. Pyrophoric iron.
Thanks for watching.
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