Warning: Nitric acid and sodium bisulfate are corrosive, wear gloves when handling them.

These reactions also produce toxic nitrogen dioxide gas and should be performed outside or in a fume hood.

Greetings fellow nerds.

In this video, we're going to revisit making nitric acid.

Nitric acid is one of the top three acids used in chemistry along with hydrochloric and sulfuric acid.

It also has the very useful property of being an oxidizing acid even at room temperature,

and so can dissolve difficult metals like copper.

So almost every general chemist has at least a little somewhere.

Now the most common and still the most effective way to make it is to react a nitrate salt with sulfuric acid and distill off the nitric acid.

Unfortunately for some amateurs, sulfuric acid is very difficult to acquire.

I've shown many ways to make sulfuric acid on my channel, but the labor cost is still high so an alternative to sulfuric acid is desirable.

One of the most common alternative acids is sodium bisulfate, which is often sold as a pH lowering chemical for swimming pools.

It's cheap and easy to acquire in large quantities, and it's not nearly as restricted as actual sulfuric acid.

So in this video we're going to focus on making nitric acid using sodium bisulfate.

Now there are a couple of popular methods for using sodium bisulfate.

The wet method where the nitrate salt and the sodium bisulfate are both dissolved in water first.

And the dry method where the sodium bisulfate is directly heated until molten with the nitrate salt.

I'm actually not sure which one produces higher yields so i'm going to try both and observe the results.

Skip to the end if you want to know which method is best.

The first method i'm going to try is the dry method.

We start with 43g of sodium nitrate and add to it a slight molar excess of 75g of sodium bisulfate monohydrate,

the most commonly sold form of sodium bisulfate.

Then we thoroughly mix the chemicals since the later molten stage doesn't mix very well on its own.

Now we set up a distillation apparatus and directly heat it on the hotplate to the highest temperature we can reach.

In my case that's 300 celsius.

Overall, it's actually quite a simple process.

What's happening is the sodium nitrate is reacting with the sodium bisulfate to produce nitric acid and sodium sulfate.

The orange brown gas you're seeing is nitrogen dioxide as at high temperatures the nitric acid decomposes into nitrogen dioxide, water and oxygen.

Don't worry too much about it though, the amount that decomposes is small and does not significantly impact yield.

But you do need to perform this reaction outside or in a fume hood for safety.

Nitrogen dioxide is quite toxic to inhale.

Some amateurs will use sealed apparatus and try to collect the nitrogen dioxide and lead it into water to produce slightly more nitric acid.

You can do that if you like, just be careful of suckback as nitrogen dioxide is very soluble in water.

Anyway, we keep heating until no more liquid distills over.

I'm going to set this run aside and try some other methods.

Now because sodium bisulfate is very cheap, some procedures use a large molar excess of sodium bisulfate to improve yield.

So I'm trying that here using 43g of sodium nitrate and 150g of sodium bisulfate, which is a little over two molar equivalents to the sodium nitrate.

Hopefully this will drive the reaction forward and improve yield.

I said before that I'm using sodium bisulfate monohydrate and not anhydrous sodium bisulfate.

There is a good reason for that.

Strictly speaking, this method isn't entirely dry, it requires that one of the reactants melt and react.

Anhydrous sodium bisulfate melts at 315 celsius.

This is a very high temperature and most lab heating equipment strains to get that high.

While it's possible to achieve, it would be much easier if we could use more modest temperatures.

Sodium bisulfate monohydrate melts at just 58.5 celsius.

So we just need to reach the boiling point of nitric acid, which is modestly higher at around 120 celsius.

We don't need to reach 315 celsius.

So try and use sodium bisulfate monohydrate for this procedure

and if you only have anhydrous sodium bisulfate then it might be desirable to add in one mole equivalent of water to bring down the melting temperature.

If you're unsure of which type you have, then just try melting some sodium bisulfate on the hotplate by itself.

If it requires more than a hundred celsius then you know you'll need to add water.

Anyway, looks like our second procedure is done.

I'm going to set that aside now.

Now i'm going to try the wet chemical method.

To 150mL of water, we add 43g of sodium nitrate and 75g of sodium bisulfate.

And then we stir until dissolved.

Now we set up a distillation apparatus and distill over the nitric acid solution.

By dissolving the reactants first in water we ensure they have a chance to completely react,

and so we should get a higher yield than the dry method.

At least that's the idea, dry methods sometimes work just as well as wet methods.

You might be wondering where I got my sodium nitrate.

I actually made it by reacting ammonium nitrate with sodium hydroxide.

You can use ammonium nitrate directly and it will work,

but if the chemicals aren't perfectly pure then a decomposition reaction may occur where the ammonia reacts with the nitric acid at high temperatures.

You can end up destroying a huge amount of your yield and also belch out copious quantities of nitrogen dioxide gases.

You can even have violent decomposition that's so fast it blows apart your apparatus.

So for better yield and safer handling, I recommend converting ammonium nitrate to sodium nitrate if ammonium nitrate is all you have.

Alternatively, if you have potassium nitrate or calcium nitrate, you can use those directly as long as you recalculate the appropriate stoichiometry.

If you're using calcium nitrate in particular then it will precipitate calcium sulfate upon mixing

so you may need to mix it and the sodium bisulfate with water separately,

then combine them and filter out the calcium sulfate before performing the distillation.

Anyway, continue distilling until no more nitric acid comes over.

Now if you look at the boiling flask, you can see a large hunk of sodium sulfate.

The problem with all that sodium sulfate is that it may trap reactants inside of it.

I'm not sure if it's a significant problem, but one of the methods I've seen for making nitric acid is to crystallize out the sodium sulfate before distillation.

I'm going to try that and mix another batch of 43g sodium nitrate and 75g sodium bisulfate in 150mL of water.

Sodium sulfate is very soluble in water at room temperature but becomes dramatically less soluble at cold temperatures.

So we're going to remove a large portion of the sodium sulfate by taking our mixture and putting it into the freezer for a few hours.

And here we are, so much of the sodium sulfate has crystalized out that it's almost like a slush.

Now we suction filter the mixture as fast as possible before it has a chance to melt again.

Interestingly enough, the crystals of sodium sulfate contain ten mole equivalents of water.

Crystallizing sodium sulfate decahydrate removes a tremendous amount of water and that's why the volume of crystals here is so large compared to the original chemicals we had earlier.

This also conveniently concentrates our nitric for the subsequent distillation.

And there is our filtrate.

It still has some sodium sulfate in it so we'll still need to distill our nitric acid.

Now distill until we reach half the volume and stop.

When it cools put it back into freezing and again filter out the crystals of sodium sulfate.

There is still a lot of sodium sulfate in the solution and we need to keep removing it.

Once again we distill the remaining filtrate.

But since there is such a small amount this time we're not going to freeze it again and instead distill all the way to dryness.

And here we are, our four products from our four attempts.

Now we can't directly compare the yields because they're going to have different concentrations.

Even the dry methods don't produce pure nitric acid as the sodium bisulfate monohydrate will release some water when heated and this will not be consistent between runs.

Fortunately we can crudely calculate the concentrations by finding the densities of the solutions and checking them against tables of known nitric acid concentrations.

And the results are very interesting.

It seems the yields are not dependent on which method is used but more on the quantity of sodium bisulfate.

The wet methods seem like a non-starter, they produce slightly higher yield but are too dilute.

Even using the labor intensive freezing and filtering process only improved concentration, but actually hurt yield.

While i'm sure they can be improved by adjusting the water added, it seems the dry method using two mole equivalents of sodium bisulfate is the best.

I'm actually quite surprised just how high the yield is, at around 95% this is competitive with using sulfuric acid.

Since sulfuric acid is somewhat valuable for the amateur, if you don't need pure anhydrous acid and can tolerate 75% concentration,

then using very cheap and easy to acquire sodium bisulfate may be preferable.

A minor bonus is that both reagents are dry and considerably safer to handle than concentrated sulfuric acid.

I think I might use this procedure from now on to make my nitric acid.

I didn't bother testing if the wet method would give similarly high yields with increased sodium bisulfate because at 95%, the most improvement a new method could give is 5%.

The wet method, even without the extra labor of freezing, is much slower as all the extra water needs to be distilled as well.

It's not worth it for just an extra 5%.

So overall, I would use the dry method, and use the huge stoichiometric excess of sodium bisulfate,

since it gives the best yield, the highest concentration and also uses the least labor.

Now some additional notes before I go.

A simple way to test if you have nitric acid and didn't monstrously blunder and grab the wrong chemicals off the shelf is to react it with copper metal.

The copper metal will dissolve and produce the characteristic orange brown haze of nitrogen dioxide.

While some other chemicals will also dissolve copper, extremely few will also generate nitrogen dioxide other than nitric acid.

Moving on, the orange brown haze you're seeing in the higher concentrations of acid is nitrogen dioxide gas that resulted from the decomposition of nitric acid at high temperatures.

While it also occurs at lower concentrations, the nitrogen dioxide reacts with the water to reproduce nitric acid.

At high concentrations though the reversible chemical reaction doesn't go to completion.

Anyway, the actual amount is very small and there is little point in removing it since when you use nitric acid for oxidizing reactions like dissolving copper,

you're going to make nitrogen dioxide regardless.

But if you really want to remove it because you have a special reaction or just want nice clear acid for your collection,

then you can redistill the acid at lower pressure to reduce the boiling temperature and so reduce the decomposition.

You can also remove the nitrogen dioxide by directly stirring in small amounts of either urea, ammonia or hydrogen peroxide.

I'm using hydrogen peroxide here.

Just keep adding small portions until the nitrogen dioxide is gone.

Hydrogen peroxide forces the nitrogen dioxide to convert into nitric acid.

Ammonia and urea outright destroy the nitrogen dioxide and convert it into nitrogen gas.

The drawback of these methods is that ammonia and hydrogen peroxide will add water and dilute your acid.

And if you overshoot and add too much ammonia or urea, the excess reagent will become ammonium nitrate that remains in your acid.

So you'll need to decide if that's acceptable to you.

Anyway, so that's how you make nitric acid using sodium nitrate and sodium bisulfate.

Thanks for watching.

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