Sciencemadness Discussion Board - Convert NaNO3 to KNO3 and purify it

Convert NaNO3 to KNO3 and purify it

PyroPlatinum - 8-1-2019 at 16:28

Hello everyone. I want to make a question that is intriguing me for a while.
So. I bought a pack of saltpeter but it has two components in it: sodium nitrate and potassium nitrate. I suspect that might be a 50:50 mixture or 65:35 based on some data sheets i found in the internet. But I'm not really sure. Ok, so this is the thing, i'm only interested in the potassium nitrate and i want to convert the remaining sodium nitrate in the mix to KNO3 too, but i'm not sure if my plan is gonna work so i want your opinion.
So, i plan to dissolve all the nitrate in hot water and mix potassium chloride and try to precipitate the KNO3 out by cooling the solution in the freezer and then filtrating out the KNO3 crystals. But I'm not sure if that is gonna work since i heard that this double displacement reaction can't occur according to some forums. But I'm not 100% convinced of it.
So, what you guys sugest? I'm not a chemistry expert and don't have access to all chemicals i might need. So i want some suggestion for something i can do at home with no fancy glassware or inaccessible chemicals.
I know i can buy pure KNO3 but i'm not willing to and i'm interested in using this that i already have. But anyways. Thank for any answers and your time.

Sigmatropic - 8-1-2019 at 22:43

Use a really soluble potassium salt, KCl should work but my mind springs up with K2CO3. KOH would also be perfect.

Just be sure to check that your not over the saturation point for the starting potassium salt or the product sodium salt,otherwise your double displacement won't work.

PyroPlatinum - 9-1-2019 at 06:50

Hmm interesting. Thanks. I didn't tought about KOH, but sounds promising. The only problem is that i think i will have to get hold of some KOH because i only have NaOH :/. But anyways. What do you mean about not being over with the saturation point for the potassium salt? Because of the KNO3 already present in the solution? (Sorry. I'm a little slow to understand the things lol)

PyroPlatinum - 9-1-2019 at 09:22

Unfortunately i have no KOH or K2CO3 at my disposal right now. I have only these chemicals:
Potassium Chloride
Sodium Chloride
Iron(III) Chloride
NaNO3+KNO3 mix
Sodium Bicarbonate
Sodium Hydroxide
Diluted Sulfuric Acid
Hydrochloric Acid
Phosphoric Acid
Diluted Acetic Acid (Vinegar lol)
Calcium Oxide and Hydroxide
Magnesium Hydroxide
Sodium Hypochlorite
Calcium Carbonate
Potassium Permanganate
Hexane
And some metals...

r0749547 - 9-1-2019 at 09:58

Well if you want to convert NaNO3 into KNO3 you need a K+ source.
K2CO3 aka potash is a common fertilizer and can be bought at a graden store.
and using potassium permanganate would be a waste in my eyes.

FYI put the KNO3 and NaNO3 mix in a freezer and your KNO3 should crystallize out.

[Edited on 9-1-2019 by r0749547]

PyroPlatinum - 9-1-2019 at 10:29

So, would be wiser to try to crystallize out some of the KNO3 first and then convert the remaining NaNO3? And by the way, i tought about KCl because i already have a full sack of it here that i bought a while ago...

Sigmatropic - 9-1-2019 at 10:45

To get a good yield with double displacement reactions you want everything but your product to be highly soluble.

Do understand that there is such a thing as a solubility constant but for this theoretical excersize we can use the concentration of saturated solutions as a guideline.
KCl will form a saturated solution with 34.2 g/100 mL water
KNO3 at 31.6 g/100 mL

Without taking into account the common ion effect, which is a result of solubility constant, we can already see that KCl will not work since KCl has a solubility comparable to that of KNO3. Meaning that if you use KCl as a potassium Donor, you may remove NaNO3 as an impurity but introduce KCl in its place.

Now if we use K2CO3, a highly soluble potassium salt (111 g/100 mL water), we run into another problem. Namely the low solubility of sodium carbonate (12.5 g/100 mL). Now you run the risk of co-percipitating Na2CO3 in your KNO3.

You can run the same thought experiment with potassium hydroxide or potassium acetate.

So what is the goal? Retrieving as much as possible NO3? Obtaining pure enough KNO3 to see its purple color in reactions?

I would agree with r0749547 and simply recrystalize the crude material without a potassium donor. Perhaps you can save your mother liquors and use them in a future preparation of nitric acid (by distillation after acidification with phosphoric acid).

If anyone feels free to complete the approaches to double displacement of NaNO3 with a soluble potassium salt by taking solubility constants into account, feel free to correct me.

hissingnoise - 9-1-2019 at 10:54

Historically, Chilean nitrate (NaNO₃) in boiling soln. was converted by mixing into a boiling soln. of KCl.

A solubility table comes in handy...

With a bit of calculation you can use it to precipitate KNO₃ from the mix and finish the metathesis using KCl.

PyroPlatinum - 9-1-2019 at 10:59

Quote: Originally posted by Sigmatropic

To get a good yield with double displacement reactions you want everything but your product to be highly soluble.

Do understand that there is such a thing as a solubility constant but for this theoretical excersize we can use the concentration of saturated solutions as a guideline.
KCl will form a saturated solution with 34.2 g/100 mL water
KNO3 at 31.6 g/100 mL

Without taking into account the common ion effect, which is a result of solubility constant, we can already see that KCl will not work since KCl has a solubility comparable to that of KNO3. Meaning that if you use KCl as a potassium Donor, you may remove NaNO3 as an impurity but introduce KCl in its place.

Now if we use K2CO3, a highly soluble potassium salt (111 g/100 mL water), we run into another problem. Namely the low solubility of sodium carbonate (12.5 g/100 mL). Now you run the risk of co-percipitating Na2CO3 in your KNO3.

You can run the same thought experiment with potassium hydroxide or potassium acetate.

So what is the goal? Retrieving as much as possible NO3? Obtaining pure enough KNO3 to see its purple color in reactions?

I would agree with r0749547 and simply recrystalize the crude material without a potassium donor. Perhaps you can save your mother liquors and use them in a future preparation of nitric acid (by distillation after acidification with phosphoric acid).

If anyone feels free to complete the approaches to double displacement of NaNO3 with a soluble potassium salt by taking solubility constants into account, feel free to correct me.

Yes. I'm interested in obtaining as much as Nitrate as possible but at same time without too much of contamination from sodium.
I think that the potassium acetate option sounds interesting. What you suggest about using it? And thanks so far.

PyroPlatinum - 9-1-2019 at 11:01

Quote: Originally posted by hissingnoise

Historically, Chilean nitrate (NaNO₃) in boiling soln. was converted by mixing into a boiling soln. of KCl.

A solubility table comes in handy...

With a bit of calculation you can use it to precipitate KNO₃ from the mix and finish the metathesis using KCl.

Hmm interesting

PyroPlatinum - 9-1-2019 at 11:38

After thinking a little about it i might just crystallize out the KNO3 and after i buy some KOH use it to convert the remaining nitrate. The NaOH should remain in solution right? Since potassium nitrate has low solubility in water at low temperatures. What do you think?

hissingnoise - 9-1-2019 at 12:16

KCl is more convenient than KOH ─ and its conc. solution won't etch your glass...

MrHomeScientist - 9-1-2019 at 12:50

This is similar to my recent thread on making KNO₃ from ammonium nitrate and KCl: http://www.sciencemadness.org/talk/viewthread.php?tid=124565

Sigmatropic's mention of saturated solutions is somewhat misleading, because that is only for one temperature - he used 20C. What you really need to look at is the full solubility curves for everything in your solution. Take a look at the spreadsheet I posted in that thread. (I pulled that info from Wikipedia's solubility table.) While KNO₃ and KCl are indeed very similar at room temperature, KNO₃'s solubility drops considerably as the solution is cooled. This is why you can recover it so cleanly from solutions.

In your case, you'd need to make a similar spreadsheet with KCl, KNO₃, and NaNO₃. A quick glance and you see that sodium nitrate is far more soluble than the other two at any temperature, so really you're just concerned with the other two. You then need to take a look at the solubility at the temperature of your fridge, and use that to calculate how much water you need for a known amount of salts.
(Edit: Whoops, forgot about NaCl. That is also above the solubilities of the potassium salts, though, so it shouldn't be a problem either.)

In my case, my fridge is at 3C and I calculated KCl's solubility to be 28.96 g/100mL at that temperature. If you used, say, 100g of KCl this means you need a minimum of 345mL of water to ensure no KCl drops out. Your potassium nitrate will crystallize out and everything else will remain in solution. If you used more soluble potassium salts like the hydroxide or carbonate (as others mentioned), you could get away with less water and recover more nitrate per run.

In summary, if I were doing this I would (1) make a saturated room temperature solution of your mixed nitrates, (2) chill the solution to recover as much KNO₃ as possible, (3) warm back to room temperature and dissolve a calculated amount of KCl to convert the remaining sodium nitrate, (4) chill again to recover a second crop of KNO₃. If you want to ensure purity, recrystallize the KNO₃ once or twice. But it should be quite pure as is. A flame test is a very sensitive indicator for sodium contamination, if you want to confirm.

[Edited on 1-9-2019 by MrHomeScientist]

PyroPlatinum - 9-1-2019 at 12:59

Quote: Originally posted by hissingnoise

KCl is more convenient than KOH ─ and its conc. solution won't etch your glass...

Well. That's a nice point because we don't want to dissolve our glassware, right? lol but we can use plastics. But i see three reasons why it's more difficult for me.:
1 - Here where i live Potassium Hydroxide is expensive and difficult to come by; (Only buying online)
2 - I think would be better to minimize the risks by not using really caustic compounds;
3 - And since i only have about 500g of nitrate don't know if it's viable and i have a ton of KCl too! Lol

But it's a interesting alternative in it's own right and if my experiments with KCl and Nitrates don't wield a satisfactory result i might consider trying it out. At least for the sake of learning some more chemistry.

PyroPlatinum - 9-1-2019 at 13:24

Quote: Originally posted by MrHomeScientist

This is similar to my recent thread on making KNO₃ from ammonium nitrate and KCl: http://www.sciencemadness.org/talk/viewthread.php?tid=124565

Sigmatropic's mention of saturated solutions is somewhat misleading, because that is only for one temperature - he used 20C. What you really need to look at is the full solubility curves for everything in your solution. Take a look at the spreadsheet I posted in that thread. (I pulled that info from Wikipedia's solubility table.) While KNO₃ and KCl are indeed very similar at room temperature, KNO₃'s solubility drops considerably as the solution is cooled. This is why you can recover it so cleanly from solutions.

In your case, you'd need to make a similar spreadsheet with KCl, KNO₃, and NaNO₃. A quick glance and you see that sodium nitrate is far more soluble than the other two at any temperature, so really you're just concerned with the other two. You then need to take a look at the solubility at the temperature of your fridge, and use that to calculate how much water you need for a known amount of salts.
(Edit: Whoops, forgot about NaCl. That is also above the solubilities of the potassium salts, though, so it shouldn't be a problem either.)

In my case, my fridge is at 3C and I calculated KCl's solubility to be 28.96 g/100mL at that temperature. If you used, say, 100g of KCl this means you need a minimum of 345mL of water to ensure no KCl drops out. Your potassium nitrate will crystallize out and everything else will remain in solution. If you used more soluble potassium salts like the hydroxide or carbonate (as others mentioned), you could get away with less water and recover more nitrate per run.

In summary, if I were doing this I would (1) make a saturated room temperature solution of your mixed nitrates, (2) chill the solution to recover as much KNO₃ as possible, (3) warm back to room temperature and dissolve a calculated amount of KCl to convert the remaining sodium nitrate, (4) chill again to recover a second crop of KNO₃. If you want to ensure purity, recrystallize the KNO₃ once or twice. But it should be quite pure as is. A flame test is a very sensitive indicator for sodium contamination, if you want to confirm.

[Edited on 1-9-2019 by MrHomeScientist]

Your strategy is really close to what i was considering doing! And yes, i was wondering why that would not be possible since the KNO3 solubility at low temperatures is poor, like 14g/100g of water at 0°C if i recall it correctly.
Sure i will take a look at your thread and make a list of the solubility of all the components involved in this reaction and then the calculations.
Unfortunately I'm not 100% sure about the percentage of the nitrate so i think i will have to guess it...
And about the chilling.. I could make a ice bath too with salt to cool it even further but that's not really necessary right?
Anyways. Thank you very much so far.

Felab - 9-1-2019 at 14:10

Quote: Originally posted by hissingnoise

KCl is more convenient than KOH ─ and its conc. solution won't etch your glass...

Potassium hydroxide doesn't etch glass at a significant rate unless it is extremely concentrated and boiling hot.

Keep looking for it, you will eventualy find a good deal.

I recomend you use potassium hydroxide but if you cannot find any there are a few videos wich detail the synthesis of potassium nitrate with ammonium nitrate and potassium chloride.

MrHomeScientist - 9-1-2019 at 18:15

I get mine from Essential Depot, an online store for soap makers. Pretty good deals, usually.

It couldn't hurt to reduce the temperature further; it'll only make more of the nitrate come out! You could probably go a ways below 0C too, with all the dissolved salts pushing the freezing point down. I don't have solubility data below 0, though.

If you can find an MSDS for your mixed nitrate product that might tell you more. Sometimes they're intentionally vague, though. After the first cooling cycle (before adding any KCl), the amount of KNO₃ you get out might give you a clue to the percentage. Good luck!

Herr Haber - 10-1-2019 at 04:09

If you are going to use the KNO3 for pyro uses as your names suggests I would definitely go the KCl route.
This way you dont have to worry too much about purification and KNO3 with a little leftover KCl will probably give you a better violet color.

PyroPlatinum - 10-1-2019 at 13:11

Quote: Originally posted by Felab

Quote: Originally posted by hissingnoise

KCl is more convenient than KOH ─ and its conc. solution won't etch your glass...

Yeah but weirdly i cannot find ammonium nitrate for sale here.. only sodium, potassium, calcium, barium, strontium and some other...
After i get done with this experiment i'm gonna look for more pure chemicals..

PyroPlatinum - 10-1-2019 at 13:30

Quote: Originally posted by MrHomeScientist

I get mine from Essential Depot, an online store for soap makers. Pretty good deals, usually.

It couldn't hurt to reduce the temperature further; it'll only make more of the nitrate come out! You could probably go a ways below 0C too, with all the dissolved salts pushing the freezing point down. I don't have solubility data below 0, though.

If you can find an MSDS for your mixed nitrate product that might tell you more. Sometimes they're intentionally vague, though. After the first cooling cycle (before adding any KCl), the amount of KNO₃ you get out might give you a clue to the percentage. Good luck!

Good to know. I'm gonna look around for this site. And soapmaking is a interesting thing on it's own anyways lol got some dozens of liters of soy oil here. But without getting off topic, i'm gonna experiment with lower temperatures and see what i can do.
And about the safety data sheet.. i couldn't find any information though and even tried to ask to the company that manufacturers it, but i don't think i'm gonna receive a relevant information regarding this.. the companies unfortunately tend to "hide" the composition for their products a little XD.
But as you said, after measuring the weight of all before and afterwards and making some calculations i can have an ideia of the proportion NaNO3/KNO3.

Anyways... Thanks man, you have been very helpful, as well everybody that answered my question here.

Felab - 10-1-2019 at 13:36

Quote: Originally posted by PyroPlatinum

When I started with chemistry, I didn't knew where to buy reagents. Now I know three stores which sale basicaly every normal chemical you can encounter in the average lab with out asking you questions. The problem is that I live in Spain, I don't know if the same thing happens on other countries.

Although a part of the EU, Spain is quite soft when it comes to laws about prohibiting chemicals.

[Edited on 10-1-2019 by Felab]

PyroPlatinum - 10-1-2019 at 13:37

Quote: Originally posted by Herr Haber

If you are going to use the KNO3 for pyro uses as your names suggests I would definitely go the KCl route.
This way you dont have to worry too much about purification and KNO3 with a little leftover KCl will probably give you a better violet color.

Yes, mainly for pyrotechnics but it's a nice useful oxidant that i like to keep for other reactions too. And yeah, gonna try this and see if i get some good results. I'm just waiting my new digital scale arrive so i can do the measurements. But hey, thank for your opinion.

PyroPlatinum - 10-1-2019 at 13:43

Quote: Originally posted by Felab

Quote: Originally posted by PyroPlatinum

Yes.. that's a problem. I was thinking about buying some chemicals from united nuclear but i don't know if they ship to other countries. And there is ebay too XD but some prices are a bit steep.

PyroPlatinum - 10-1-2019 at 13:49

Quote: Originally posted by Felab

Quote: Originally posted by PyroPlatinum

I'm gonna start studying Agronomic engineering this year so i think after i graduate, hopefully, going to be easier for me to buy some chemicals... but there is still a long way to go before this...

PyroPlatinum - 20-1-2019 at 13:55

So. I have found something that might help me to know how much KNO3 and NaNO3 i have here.
In the box (photo attached) says:
"Total Nitrogen (N)..................................................15%"
"Potassium Oxide (K2O) soluble in water..............14%"
And it's NPK "code" is: 15-00-14.

hissingnoise - 21-1-2019 at 03:37

Why do you think it contains NaNO₃?

PyroPlatinum - 21-1-2019 at 07:29

Quote: Originally posted by hissingnoise

Why do you think it contains NaNO₃?

Because it says in the label explicitly that it has. And potassium nitrate too. But unfortunately they never say how much of each.. that make the calculations a little difficult...

[Edited on 21-1-2019 by PyroPlatinum]

PyroPlatinum - 22-1-2019 at 21:29

So. I have done the first recrystalization, and so far i've got about 57,5 grams of (a bit impure KNO3). That's about 12% of the total nitrate i dissolved in the water (471,2 grams).
Tried again to recrystallize but very little crystallized out. (with a bit of NaNO3 i think or it was only ice..)
Anyway. Now i think I'm gonna have to boil it down to about 75% the total volume of the solution and try to recrystallize more again.
If I'm doing something wrong please feel free to tell me.
And thanks for everyone that helped me out so far.

PyroPlatinum - 22-1-2019 at 21:41

And btw, i have found here where i live sacks of Potassium Nitrate im fertilizer form with about 90% purity. (I think).
They are selling it for about $55 for each 25 Kg sack.
But i also found out that it has some Magnesium Sulfate (MgSO₄) as the impurity in it.
So i'm thinking how i would get rid of that sulfate if i buy the nitrate in the future...
I think i would have to convert that magnesium sulfate into chloride before trying to crystallize out the potassium nitrate, because i'm afraid that the sulfate might crystallize out with the nitrate, due to it's relatively low solubility at low temperatures (the double of KNO3 though).
And that's the only option before magnesium iodide, bromide, acetate and chlorate. And i would not want to waste my chlorate with this! lol
Any suggestions?

[Edited on 23-1-2019 by PyroPlatinum]

MrHomeScientist - 23-1-2019 at 13:56

Now that you've made it through the spam, here's an idea: You could approach it from the other direction. Rather than cooling and crystallizing the KNO₃ at low temperature, you could try recovering the MgSO₄ from a hot solution instead. The solubilities are much farther apart at high temperature. So, dissolve the maximum amount of nitrate possible at near-boiling temperature, then filter off any remaining solids. That will be a lot of your sulfate contamination.

Or, perhaps even better: react a solution of your fertilizer with calcium nitrate. That will precipitate out ALL the sulfate, leaving you with the much more soluble magnesium nitrate to work with. Then you're back to the original procedure of cooling and recovering the KNO₃.

PyroPlatinum - 23-1-2019 at 14:50

Quote: Originally posted by MrHomeScientist

Now that you've made it through the spam, here's an idea: You could approach it from the other direction. Rather than cooling and crystallizing the KNO₃ at low temperature, you could try recovering the MgSO₄ from a hot solution instead. The solubilities are much farther apart at high temperature. So, dissolve the maximum amount of nitrate possible at near-boiling temperature, then filter off any remaining solids. That will be a lot of your sulfate contamination.

Or, perhaps even better: react a solution of your fertilizer with calcium nitrate. That will precipitate out ALL the sulfate, leaving you with the much more soluble magnesium nitrate to work with. Then you're back to the original procedure of cooling and recovering the KNO₃.

lol these spams pisses me of so much.

Hmmmm that sounds pretty straightforward... gonna try that once i buy of that nitrate sacks. I think that heating a concentrated solution of it and filter off the sulfate would be much simpler.

Oh and about the Calcium nitrate method... it really makes sense but it would require this additional salt. But couldn't i use Calcium Chloride?
Would end up with Calcium Sulfate that i would filter off, and after that i might just recrystallize the Potassium nitrate and leave the Magnesium Chloride in solution.
And Calcium Chloride is pretty cheap compared to it's nitrate counterpart.
(But i understand that would yield less nitrate)
Anyways. Good suggestions. Thanks.

MrHomeScientist - 28-1-2019 at 14:14

Sure, chloride would work. It's the calcium ion that does the heavy lifting here. I just chose nitrate to minimize the number of extra ions.