Greetings fellow nerds.
So i've made tremendous progress in developing the alcohol catalyzed magnesium reduction approach to making sodium.
As you might know by now i have in fact been able to get it to work with cheap tertiary alcohol.
So in this video i'm going to give you a detailed progress report on what i've found and future directions i'll be going in.
This is not a how-to video, this is a video on how i'm figuring out how to make a how-to video on sodium.
So a brief recap, making sodium metal is considered one of the holy grails of amateur chemistry.
It's very useful for such applications as a reducing agent, a drying agent, and just being lot's of fun.
Making it however is very hard.
Electrolysis requires handling molten sodium hydroxide or sodium chloride.
This is extremely dangerous and requires special equipment.
Last year i showed that sodium could be obtained from high temperature magnesium reduction of sodium hydroxide along with dioxane separation.
I call this the thermochemical dioxane approach.
While this is every bit as dangerous as it sounds,
it could be performed with just a soup can and a brick, no need for special equipment.
And the danger was only present for a few minutes during the initial thermochemical reduction.
After cooling it was considerably safer as a sodium magnesium oxide aggregate.
At which point it could then be separated into sodium metal by boiling with dioxane.
This process brought sodium metal within reach of the well-equipped amateur chemist.
But there was always the question of whether we could do better, could we make sodium with just the heat of the hotplate and nothing more complicated than that?
That's the promise of alcohol catalyzed magnesium reduction approach.
Basically just mix a bunch of chemicals, heat them up on a hotplate, and get sodium without the dangers of molten salts or fireballs.
The process already works for potassium and is well-known in the amateur chemistry community.
But there was a lot of difficulties to get it to work for sodium and many conflicting and inconsistent reports of success.
So i've been trying my hand at it and seeing if i could get an efficient and reproducible process.
First I started with my standard mixture of sodium hydroxide and magnesium metal along with various tertiary alcohols and tetrahydronaphthalene solvent.
I selected tetrahydronaphthalene solvent because it worked for me in the past when making potassium.
But after many weeks of failure i eventually ran some tests on the tetrahydronaphthalene and found out it was the direct cause of the failures.
Apparently tetrahydronaphthalene reacts with sodium metal to produce tar.
This is not entirely unprecedented, as naphthalene, a related compound, will react with sodium as well.
In fact, during my old experiments with potassium, the tetrahydronapthalene also reacted with potassium and coated the spheres with tar but the reaction was slow enough as not to be a problem.
As long as alcohol kept being added the tar would be reacted away and the metal production continued.
I thought sodium would work similarly but it turns out the production of sodium is so slow that the tar formation overwhelms the process.
Bottom line, tetrahydronaphthalene had to be abandoned as a solvent.
At this point in time i was weary of continue a project after months of failure
so i attempted a theory testing experiment using a superheavy alcohol i had specifically engineered for such a test.
This alcohol, 7-hexyl-7-tridecanol, was designed to work as both catalyst and solvent
as well as possess the high boiling point needed to test the upper limits of the reaction.
If there was any chance of this reaction working, it would have to work here.
If it didn't, i would give up, cut my losses, and move onto projects with greater return on investment.
So i prepared the alcohol, jumpstarted it with a tiny amount of sodium and added in magnesium and sodium hydroxide reagents.
Upon heating we had our breakthrough and for the first time we produced sodium by the alcohol catalyzed magnesium reduction approach.
I've posted a link to that video in the video description.
Now this alcohol was far too expensive to actually use in practice.
But it nonetheless proved the process could work.
Knowing something is possible is tremendously important before committing more time, effort and money to it.
Now the work so far gave us some tremendously important leads.
First, the tetrahydronaphthalen failures showed us that aromatic rings are probably detrimental to the reaction.
At these high temperatures they seem to be susceptible to tar formation.
So whatever we do in the future, we have to avoid anything containing aromatic rings.
Second, 7-hexyl-7-tridecanol working shows us that a good solvent to try would be a high boiling solvent consisting entirely of long chain hydrocarbons.
Basically, we need mineral oil also known as paraffin oil.
But we need mineral oil without aromatic rings as shown earlier.
A fair number of commercial mineral oil formulations have small amounts of impurities that contain aromatic rings.
This might be a contributing factor to the failures of myself and previous amateur chemists who tried mineral oils.
The easiest to get commercial source of mineral oil that's pure enough for our purposes is hypoallergenic baby oil.
Despite what the name suggests, baby oil is not made from babies,
it's a mineral oil but it's been specially treated and distilled to remove all traces of aromatic rings and other impurities.
Thus making it hypoallergenic.
It's also very easy to get from domestic sources.
This is going to be our new choice of solvent.
So i tried a new experiment, mixing together hypoallergenic baby oil, magnesium, sodium hydroxide and a milliliter of 3-ethyl-3-pentanol.
This alcohol is also expensive but still considerably cheaper than 7-hexyl-7-tridecanol.
If it worked, i would try even cheaper alcohols.
But it didn't work.
I even tried to jumpstart the reaction by adding in some sodium metal.
I ran the reaction for a day and still nothing seemed to happen.
On a hunch i hooked up a bubbler to try and observe hydrogen formation.
If the reaction was proceeding, it should be producing hydrogen gas as well as sodium.
Interestingly enough I saw bubbling happening.
This wasn't confirmation of success though.
The bubbling could simply be the sodium metal jumpstarter reacting with residual water in the sodium hydroxide.
It could even be magnesium reacting with water.
In either case we would see hydrogen bubbling, but not get new sodium.
Now up to this point i had not stirred the reaction.
This was because we never had to when making potassium.
But if the reaction was proceeding and if it was happening very slowly, then stirring would logically speed it up.
So i started a new mixture with the same conditions but this time began stirring and i did not put in a sodium metal jumpstart.
I also hooked up a bubbler in the back to observe any gas formation.
This time things went much better.
Gas formation was much more pronounced.
But i didn't see any trace of sodium even after several hours.
Potassium normally formed much faster.
Because of this experience with potassium, most of us amateur chemists, including myself, gave up on sodium because we thought it should have produced something by now.
But watching the bubbler made me realize something was happening, i just had to be patient.
i left the reaction for a few days, checking back occasionally to see if anything had progressed.
Gas kept coming out of the bubbler but it wasn't until a couple of days later that i observed liquid sodium mixed with lots of magnesium.
It had worked! Excited i stopped heating and poured out the slurry.
Looking inside i found the sodium and magnesium had solidified.
It was a lot of work to get it out but eventually i did.
Now this sodium is grossly impure being covered in magnesium oxide and solidified with magnesium metal inside.
I tried to separate it by heating it in baby oil but all this did was remove the magnesium oxide.
The magnesium metal bits still remained in the sodium.
Fortunately we can draw on our experience with making sodium by the thermochemical dioxane process.
I placed the sodium magnesium aggregate in dioxane and refluxed the mixture with stirring.
It slowly worked and as they seperated, the magnesium metal sank while sodium metal floated.
It coalesced into a single large sphere of sodium metal.
And that was beautiful.
Interestingly enough there was still a small amount of magnesium stuck to the bottom of the sodium.
This was a very small amount though and easily cut off if desired.
But I didn't bother since it was minor and not a problem for future experiments.
I wanted to keep this as a trophy but i had to measure its mass.
Unfortunately it was too big to fit through the neck of the flask.
So i poured off the dioxane and smashed the flask to get the sodium out.
The flask was cracked anyway so it wasn't a significant loss.
The sodium weighs in at 7.8 grams for a yield of about 33% on an initial quantity of 40 g of sodium hydroxide.
Now if we look over here there is still a lot of magnesium left over.
I presume the unreacted sodium hydroxide is still in the slurry i poured off earlier.
I think the low yield is for two reasons.
One is impurities in the sodium hydroxide.
Brand new sodium hydroxide is fairly pure but has a significant amount of water left over from manufacture.
It's also highly hygroscopic and will absorb water right out of the air.
Unfortunately this water will react with the magnesium and deplete it.
Reducing the amount of magnesium available for making sodium.
This can be mitigated somewhat by using excess magnesium but magnesium tends to be more valuable than sodium hydroxide.
So that is only viable up to a point.
Another impurity is sodium carbonate.
As the sodium hydroxide sits it absorbs carbon dioxide from the air, making sodium bicarbonate.
At the temperatures of the reduction reaction the sodium bicarbonate decomposes into sodium carbonate and water.
The water goes on to destroying more magnesium while the sodium carbonate isn't thermodynamically favorable for making sodium.
To confirm this last point, i actually ran a test where i used sodium carbonate instead of sodium hydroxide.
After a week of heating I got no visible sodium or obvious sodium level reactivity.
Bottom line, sodium carbonate doesn't seem to work and commercial grade sodium hydroxide will have some sodium carbonate impurities reducing yield.
Another contributing factor to reduced yield is simply because i probably didn't react the mixture long enough.
So i'm currently attempting longer runs.
Now on a separate front i was using 3-ethyl-3-pentanol as it was cheaper than 7-hexyl-7-tridecanol.
But it's still not that easily accessible to most amateurs.
My success with hypoallergenic baby oil lead me to believe that even simpler tertiary alcohols should have worked.
So i attempted a new experiment using the same conditions but this time using t-amyl alcohol, which is a relatively common, cheap, and easier to obtain.
It can even be made by a diligent amateure chemist using domestically available chemicals.
Now this particular mixture had been used before by other amateurs including myself but there weren't any reports of success.
But i think none of us tried to react it long enough with stirring.
I cannot recall any amateur chemist who tried the process for longer than a single day let alone three.
This time, it worked on the second day and i got sodium.
Now i didn't bother measuring yield since it was more important to me to prove that simple tertiary alcohols worked.
This is a very important finding as it brings the alcohol catalyzed magnesium approach within reach of many more amateur chemists
and not just the ones with enough time and money to make 3-ethyl-3-pentanol.
Now moving on, in addition to time and labour, there are other drawbacks of the alcohol catalyzed magnesium reduction approach.
First liquid sodium is destructive to PTFE coated stir bars as seen here.
While stir bars wearing down is nothing new to the hard working chemist, this is far greater than normal and presents a small but additional cost to making sodium.
The second, and i think bigger problem, is that sodium hydroxide at these temperatures is extremely destructive to glassware.
I've shown in previous videos that hot molten sodium hydroxide eats through glass and as seen here we're getting the same problem even with solid sodium hydroxide.
Borosilicate laboratory glassware is expensive and continually destroying it to make sodium is a rather severe cost.
Even greater than the magnesium, sodium hydroxide, alcohol and baby oil directly used in the process.
But I'm currently testing ways to mitigate the problem like preheating the mixture before stirring to dry the sodium as much as possible.
Anyway, that's my progress so far.
I'm currently optimizing for yield and trying to find ways to mitigate the drawbacks.
Eventually i'll have an actual how-to video on making sodium that focuses just on what works.
Such a video should be a lot more streamlined.
But let me know in the comments if you liked this progress report format that also discusses what doesn't work,
and more importantly discusses how the observations from both failures, success and test experiments are combined into theories.
Sometimes a very simple experiment had extremely complicated inception before being cut down into just the parts that worked.
We almost never see the test experiments that looked at edge cases and helped build the theory of the main experiment.
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.