Warning: Sodium is a highly reactive and flammable metal.
This reaction also uses high temperatures and generates flames.
Fire safety protocols must be in place.
Greetings fellow nerds.
In this video we're going to make a sodium magnesium oxide aggregate useful as a substitute for pure sodium metal in most cases.
First get a steel container.
I'm using a steel soup can.
Now we get 30g of magnesium metal turnings.
Magnesium powder is much better but all I had was these turnings.
Now we're also going to need 40g of sodium hydroxide.
This is available as drain cleaner.
Thoroughly mix the two chemicals and add them to the steel container.
And there we are.
Now for the fun part, setting it on fire.
Insert a sparkler or fuse and place on top a heavy lid like a brick or steel plate.
Then light it.
What's happening is the magnesium is reacting with the sodium hydroxide in a thermite like reaction to produce magnesium oxide, sodium metal and hydrogen gas.
The hydrogen ignites and is the source for the massive flames you see here.
This process makes sodium without electrolysis, so it's much more accessible to the amateur chemist.
You need the heavy lid to reduce the exposure of air.
If you leave it open while it cools then the most of the sodium produced will be oxidized by the air.
The lid needs to be heavy so the hydrogen flame thrower doesn't blow it off.
Now i left open a tiny crack to fit the sparklers in but this is a bad move since more air can get back in.
I later realized you could just cut the sparklers, ignite them, and then cover the container completely.
I recommend doing it this way for your runs.
Although make sure the sparklers are long enough that you have time to go a safe distance.
Shooting flames is not fun.
Actually correction, shooting flames is LOTS of fun but only if you're safe while doing it.
Muauahahahaů.
Now it might seem odd that we're using a less reactive metal like magnesium to make a highly reactive metal like sodium.
Usually it's the other way around.
This works because the enthalpy of formation for magnesium oxide is very high.
Higher than the enthalpy of formation for sodium hydroxide.
So overall it's energetically favourable to rip apart the sodium hydroxide to make magnesium oxide
and leave behind sodium metal and hydrogen gas.
Okay, let that cool before handling it.
And now we have our mixture of magnesium oxide and sodium metal.
The sodium is extremely fine and interspersed throughout the magnesium so it's almost impossible to see, but it's there.
Now using a chisel or metal strip, break out the aggregate of sodium magnesium oxide.
And there it is.
Now it doesn't look sodium, but it's very reactive.
Let me show you.
I have some water here and here is a chunk of the aggregate.
Whoaů Let's try that again.
And there we go.
Muahahaha!
It's almost like sodium in its reactivity.
And in fact, it's actually slowly being oxidized on air as it sits.
So i recommend storing it under mineral oil or toluene.
Otherwise the activity will rapidly decay until it no longer reacts with water.
Now you're probably asking, what good is this stuff if it's not pure sodium metal.
Well it can be used in place of sodium for many reactions.
I used sodium to dry triethylamine in a previous video but if you didn't have that, you could use the aggregate instead.
The magnesium oxide is not a problem since it doesn't distill.
In fact in my video on dioxane I actually used chunks of this aggregate rather than actual sodium metal.
As long as your application can tolerate magnesium oxide this aggregate can be used.
Now you might be wondering how to isolate pure sodium metal.
Unfortunately the aggregate is so well mixed with magnesium oxide that it's very difficult to separate the two.
The actually sodium appears to stick to the magnesium oxide.
So even if it's boiled under an inert solvent like xylene to melt it the sodium still remains in the aggregate.
You can get extremely crude separation by dropping the chunks in a mixture of mineral oil floating on water.
This method was pioneered by youtuber nighthawkinlight and i've linked his video in the description.
Some the sodium floats into the mineral oil allowing it to be retrieved,
but unfortunately most of it gets destroyed as it falls back into the water.
Also the sodium that is collected is still extremely impure with lots of magnesium oxide leftover.
For the most part it's not worth it and if you're just going to use the sodium for solvent drying i say don't bother.
The aggregate is sufficient for that.
The problem of getting sodium without electrolysis has been plaguing amateur chemists for decades.
But it can be done.
It's an extremely involved process though so I'll be showing how to obtain pure sodium metal from the aggregate in an upcoming video.
Anyway, that's how you make sodium magnesium oxide aggregate, a useful substitute for pure sodium metal.
Thanks for watching.
Special thank you to all of my supporters on patreon for making these science videos possible
with their donations and their direction.
If you are not currently a patron, but like to support the continued production of science videos like this one,
then check out my patreon page here or in the video description.
I really appreciate any and all support.