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[*] posted on 28-12-2021 at 01:46
Sodium nitrate synthesis from salt


Hey everyone,

I did not know where to put this so I thought sodium nitrate is used in pyro so I put it here anyways quick to topic

Sodium nitrate has to be one of the easiest chemical to synthesize and I have done it various way but recently I have been trying to get to this kind of impossible route from Amonium Nitrate and Sodium chloride , well reaction does occur and I have tested in final product sodium nitrate is present but it's contaminated with Amonium chloride in such big quantity that it's of no use

The solubilities of Amonium Chloride and Sodium nitrate are rather different , at ,0 c sodium nitrate is about 75g/100 ml water and Amonium Chloride being about 29 g /100 ml yet I have not been able to separate the two , there is no info on solvent extraction as both are insoluble in the common solvents available

My question is has anyone been able to separate the two or does anybody has any idea how can it be done

Many thanks
Cheers
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[*] posted on 28-12-2021 at 03:09


Most of the solvents I can think of off the top of my head would tend to precipitate NaCl along with the others. I would just use NaOH with your AN and collect the ammonia.
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[*] posted on 28-12-2021 at 04:11


Thanks for the reply

But it will defeat the purpose because I can simply add NAOH to NH4NO3 and make Sodium nitrate and take Ammonia and yes I have done that previously I want to challenge myself in into making it with NACL and NH4NO3

I have found one way dunno if it works, Flotation process so what we do is add an oil of wetting agent and charge them in a flotation cell NH4CL is taken up by oil and Nitrates emain in water and we get two layer aquas and Oily , will try it out and se e if it works

does anyone has anyother lead or anyone knows about this flotation?


[Edited on 28-12-2021 by servo]
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[*] posted on 28-12-2021 at 05:01


You could try dehydrating the solution then heating to greater than 337oC
which should decompose/sublimate the ammonium chloride.

Sodium nitrate melts at 308oC
and decomposes at 380oC so keep temperature below that.




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[*] posted on 28-12-2021 at 06:56


The mixture will have a lower melting point.
And the stuff will decompose when you melt it.
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[*] posted on 28-12-2021 at 11:02


Yup I can vouch for that the mixture decomposed at 210 c
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[*] posted on 29-12-2021 at 00:45


Strange idea. Use Na2CO3 plus AN. Mix them, add water, and boil.



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[*] posted on 29-12-2021 at 07:14


I appreciate the answers but I have already synthesized Sodium Nitrate with Sodium carbonate and Sodium Hydroxide its so easy , its not even a challenge but the best route to make Sodium Nitrate is to make calcium nitrate first and than add soda ash , this gives best yield and purity
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[*] posted on 29-12-2021 at 09:36


“I could make it the easy way that I know works but I want to make it a highly impractical, probably impossible way instead!”

Is an attitude among many amateur chemists that I’ve never really understood. There’s nothing wrong with trying alternative methods, but when you’ve already tried it and found that it doesn’t work, why do you still insist there must be some way to make it work? Why not take that determination and apply it to something more interesting than trying to make a really cheap chemical even more cheaply? Not everything has to be a challenge. There’s enough challenges in chemistry already without creating additional silly ones.




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[*] posted on 29-12-2021 at 10:06


This simply isn't a good idea; you won't find a reasonable solvent that precipitates what you want. Heating mixtures with ammonium and nitrate in them is a good way to end up on the news. You can easily convert NaCl to the nitrate by reaction with AgNO3 but that's not any different from what you did before.

You can convert sodium chloride to the bicarbonate by the Solvay process. This can be extended in theory to convert NaCl to NaNO3 by calcium catalysis. But you risk an explosion if you heat NH4NO3 with CaCO3 to accomplish the ammonia distillation step. So instead you should convert NaCl to NaHCO3 and use that to make sodium nitrate the normal way.

You learn more about chemistry by studying a variety of processes rather than trying to learn every process that reaches a particular target.




[Edited on 04-20-1969 by clearly_not_atara]
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[*] posted on 29-12-2021 at 16:21


Purification by standard crystallization should be rather easy. Wiki requires some conversions so I may have messed this up.

NaNO3 -
180 g/100 g water (100 °C)
73 g/100 g water (0 °C)

NH4Cl -
74 g / 100mL - 100c
29.4g / 100mL - 0c

if you know how much NaNO3 you have I would add exactly as much water was needed to dissolve it at 100c temp. then decant. remaining solids should be favor NH4Cl.
then by slowly cooling the decanted solution, NH4Cl should appear first.


if a challenge is what your looking for. may I suggest growing a crystal. get the biggest crystal of NaNO3 you can and tie it to a string. then suspend it in a saturated solution. the sodium should be more attracted to the seed crystal then the ammonia your trying to filter out. but after reading the solubility the reverse might work better. the process will take a lot of time, temperature control, and control of the rate of evaporation of the solvent.

Say you could increase the concentration of your desired product, as time progressed. either by evaporating the solvent or adding more of the compound , the impurity will get forced out of the solution. In Theory

and if were going way out there as a challenge, you can use acetone. sodium nitrate is insoluble in acetone.
Please dont kill yourself. because you are dealing with compounds that are made entirety out of gas.
any decomposition will be fast and violent.

[Edited on 30-12-2021 by Rainwater]




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[*] posted on 29-12-2021 at 17:31


Quote: Originally posted by Texium  
“I could make it the easy way that I know works but I want to make it a highly impractical, probably impossible way instead!”

Is an attitude among many amateur chemists that I’ve never really understood.


I agree with you, I used to have this mentality, but I realized that wasting my time on boring chem was not my thing, and as I have mentioned in a few other posts, the cost is usually not much less if less at all then either doing the easy hassle free route, or just buying it outright.

Now I understand with nitrates not everybody can just buy it so synthesis is necessary. But even then I would just go the easy route, unless I had no plans for the nitrates so purity was not a concern and I just wanted to screw around.

Though that being said, I am also more into Organic Chem so maybe that is why I would consider this boring chem, so my opinion is biased.

Syn'

[Edited on 30-12-2021 by Syn the Sizer]
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[*] posted on 29-12-2021 at 18:03


More interesting path to NaNO3, add Na2CO3 (Washing Soda) to H2O2 for the treatment of NH3. Acidic stabilized hydrogen peroxide will liberate some CO2 from the action of Washing Soda.

Apparently, it is claimed that HNO2 can be created by the careful step wise addition of H2O2 to NH3 (see https://www.researchgate.net/publication/317692348_Effects_o...) with an approximate reaction (which I would use as a working estimate) given by:

NH3 + 3 H2O2 -> HNO2 + 4 H2O

And further:

2 HNO2 + Na2CO3 -> 2 NaNO2 + H2O + CO2

Now, in the presence of Na2CO3, the product is NaNO2. With more H2O2 (or a reputedly slow oxidation in air, see https://www.sciencedirect.com/topics/agricultural-and-biolog...), one has NaNO3.

If you happen to experience a sudden massive ampunt of gas (N2 actually), the created salt NH4NO2 has decomposed, which would limit your yield.

[Edited on 30-12-2021 by AJKOER]
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[*] posted on 29-12-2021 at 18:04


Quote: Originally posted by Texium  
There’s enough challenges in chemistry already without creating additional silly ones.


The question is why. He got you to waste your time responding in any case. The account will go dormant and another opened in 18-19 will soon make its first post, then it will start threads.




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[*] posted on 30-12-2021 at 06:03


Quote: Originally posted by Rainwater  
Purification by standard crystallization should be rather easy. Wiki requires some conversions so I may have messed this up.

NaNO3 -
180 g/100 g water (100 °C)
73 g/100 g water (0 °C)

NH4Cl -
74 g / 100mL - 100c
29.4g / 100mL - 0c

if you know how much NaNO3 you have I would add exactly as much water was needed to dissolve it at 100c temp. then decant. remaining solids should be favor NH4Cl.
then by slowly cooling the decanted solution, NH4Cl should appear first.


if a challenge is what your looking for. may I suggest growing a crystal. get the biggest crystal of NaNO3 you can and tie it to a string. then suspend it in a saturated solution. the sodium should be more attracted to the seed crystal then the ammonia your trying to filter out. but after reading the solubility the reverse might work better. the process will take a lot of time, temperature control, and control of the rate of evaporation of the solvent.

Say you could increase the concentration of your desired product, as time progressed. either by evaporating the solvent or adding more of the compound , the impurity will get forced out of the solution. In Theory

and if were going way out there as a challenge, you can use acetone. sodium nitrate is insoluble in acetone.
Please dont kill yourself. because you are dealing with compounds that are made entirety out of gas.
any decomposition will be fast and violent.

[Edited on 30-12-2021 by Rainwater]


Well Exactly what I thought and it works, I have been able to make Sodium Nitrate and I tested with sugar burning ad it burnt really well but how pure it is I dont know, I could not find a titration method for Sodium NItrate so instead I will titrate amonium chloride

Thanks for the response Highly appreciated
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[*] posted on 30-12-2021 at 07:12


Great, you've succeeded at making sodium nitrate of unknown purity in a more expensive and time-consuming way :/
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[*] posted on 30-12-2021 at 07:51


Quote: Originally posted by S.C. Wack  
Quote: Originally posted by Texium  
There’s enough challenges in chemistry already without creating additional silly ones.


The question is why. He got you to waste your time responding in any case. The account will go dormant and another opened in 18-19 will soon make its first post, then it will start threads.
Oh, you’re still on that weird conspiracy?



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[*] posted on 30-12-2021 at 09:50


Quote: Originally posted by servo  

I could not find a titration method for Sodium NItrate so instead I will titrate amonium chloride


Titration by ferrous sulfate (or ammonium ferrous sulfate) in conc. sulfuric acid media. Water content at the end of the titration shouldn't be more than 25% (water is formed in the reaction and you also introduce water with FeSO4 solution). Nitrate is reduced to nitrosyl cation according to equation:

2Fe2+ + NO3- + 4H+ = 2Fe3+ + NO+ + 2H2O

When you add sulfuric acid to nitrate solution, solution become hot, so you must cool it back to room temperature. Titration must be performed in cold water bath (water precooled from fridge) to keep temperature as low as possible (addition of FeSO4 solution to sulfuric acid media will cause heating the solution). If solution become too hot, nitrate will be partialy reduced to NO.

End-point is first permanent pink colour caused by Fe(II) nitrosyl complex. I never tried visual indication, just potentiometric with Pt electrode. I probably visit this determination again and try visual indication. More info about this method in article below:

https://sci-hub.se/https://pubs.acs.org/doi/abs/10.1021/ie50...

There is also mention about phosphoric acid media, but I never tried that. Sulfuric acid media 100% works.




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[*] posted on 30-12-2021 at 12:00


I am unaware of a practical solvent for separating ammonium chloride, ammonium nitrate, sodium chloride and sodium nitrate
All four compounds will exist in solution as the solvated ions.
In a melt you have the same problem.
Heating will cause the ammonium nitrate to decompose, potentially violently at 210C.

There may be an uncommon solvent that will work but the common ones will not work.

On a side note potassium nitrate has a much steeper solubility curve and can be separated using water.

If you feel the need to recover your mixture I suggest adding potassium chloride and precipitate the potassium nitrate.

Once you have removed the nitrate you can heat it to volatilize the ammonium chloride leaving a mixed chloride salt and then
utilize the fact sodium chloride has a flat solubility while potassium chloride does not.

Now a method that probably will work. The mixture should melt below 210C.
If you add sodium bicarbonate you will offgas ammonium carbonate.
That will leave you with a mix of sodium chloride, sodium carbonate and sodium nitrate that will be easier to separate.
It will take several rounds to get a pure product but that is how it is done commercially to separate sodium nitrate and potassium nitrate.




[Edited on 30-12-2021 by macckone]
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[*] posted on 31-12-2021 at 14:24


Quote: Originally posted by servo  
[I could not find a titration method for Sodium NItrate so instead I will titrate amonium chloride


Nitrates and organic solvents are not forgiving.
This is not to been done lightly.
Tho this could be used to purify with a soxhlet extractor.
I must say it is not worth the risk given how cheap the material being purified is.
And the fact that unknown/unexpected products will likely explode and damage your equipment/lab/self or anyone near by

If i wanted to see just how pure i got it. This is a method which i would use

First you need an accurate scale. More significant digits the better
Second you need DRY compound. Not a solid but truly dry.
Best amateur method is slight heat and hard vacuum.

1) Weigh your flask. Record (a)
2) Weigh out ~1g of your dry salts in the flask. Record (b)
3) Weight out ~50g of acetone in the same flask. Record (c)
4) Stopper, shake. Vent, Repeat,
5) Decant as much of the solution as possible without losing the solids.
Repeat step 4-5 as needed.
6) then dry the solids same as before
7) Weigh the flask again. Record (d)
b-d = weight of compounds extracted.
d ÷ b = % by mass of extracted compounds
100 - % = how pure you got your NaNO3 sample

NH4Cl is only slightly soluble in acetone. Sodium nitrate is not soluble. By using a large surplus of acetone a small amount NH4Cl will dissolve and be decanted.




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[*] posted on 31-12-2021 at 16:59


Quote: Originally posted by macckone  
I am unaware of a practical solvent for separating ammonium chloride, ammonium nitrate, sodium chloride and sodium nitrate


The answer is right above this. Ask yourself how one would make ammonium nitrate from ammonium chloride and sodium nitrate.

Carry on with the game.




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[*] posted on 1-1-2022 at 00:29


Quote: Originally posted by Rainwater  
Quote: Originally posted by servo  
[I could not find a titration method for Sodium NItrate so instead I will titrate amonium chloride


Nitrates and organic solvents are not forgiving.
This is not to been done lightly.
Tho this could be used to purify with a soxhlet extractor.
I must say it is not worth the risk given how cheap the material being purified is.
And the fact that unknown/unexpected products will likely explode and damage your equipment/lab/self or anyone near by

If i wanted to see just how pure i got it. This is a method which i would use

First you need an accurate scale. More significant digits the better
Second you need DRY compound. Not a solid but truly dry.
Best amateur method is slight heat and hard vacuum.

1) Weigh your flask. Record (a)
2) Weigh out ~1g of your dry salts in the flask. Record (b)
3) Weight out ~50g of acetone in the same flask. Record (c)
4) Stopper, shake. Vent, Repeat,
5) Decant as much of the solution as possible without losing the solids.
Repeat step 4-5 as needed.
6) then dry the solids same as before
7) Weigh the flask again. Record (d)
b-d = weight of compounds extracted.
d ÷ b = % by mass of extracted compounds
100 - % = how pure you got your NaNO3 sample

NH4Cl is only slightly soluble in acetone. Sodium nitrate is not soluble. By using a large surplus of acetone a small amount NH4Cl will dissolve and be decanted.


Thanks I'm gonna heed your warning nothing is worth risking your life..
I guess the only thing I'm left with is the method of precipitation, the both compounds have different solubilities and the last one to settle is Sodium nitrate and I purified it with sitcho of soda needed for Amonium chloride present, ammonia was liberated and salt was formed which is rather easy to remove and there u have it, I found a rather pure product
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[*] posted on 1-1-2022 at 05:25


Quote: Originally posted by Bedlasky  
Quote: Originally posted by servo  

I could not find a titration method for Sodium NItrate so instead I will titrate amonium chloride


Titration by ferrous sulfate (or ammonium ferrous sulfate) in conc. sulfuric acid media. Water content at the end of the titration shouldn't be more than 25% (water is formed in the reaction and you also introduce water with FeSO4 solution). Nitrate is reduced to nitrosyl cation according to equation:

2Fe2+ + NO3- + 4H+ = 2Fe3+ + NO+ + 2H2O

When you add sulfuric acid to nitrate solution, solution become hot, so you must cool it back to room temperature. Titration must be performed in cold water bath (water precooled from fridge) to keep temperature as low as possible (addition of FeSO4 solution to sulfuric acid media will cause heating the solution). If solution become too hot, nitrate will be partialy reduced to NO.

End-point is first permanent pink colour caused by Fe(II) nitrosyl complex. I never tried visual indication, just potentiometric with Pt electrode. I probably visit this determination again and try visual indication. More info about this method in article below:

https://sci-hub.se/https://pubs.acs.org/doi/abs/10.1021/ie50...

There is also mention about phosphoric acid media, but I never tried that. Sulfuric acid media 100% works.

I strongly suspect that chloride ions would interfere with the reaction.
An "old school" approach would be the reduction of nitrate to nitrite using cadmium sponge and titration of the nitrite with permanganate.
I'm not sure if playing with cadmium compounds is considered more or less dangerous than doing titrations in conc sulphuric.

But precipitating silver chloride and weighing it looks a much better option.

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[*] posted on 1-1-2022 at 15:44


On titration of nitrate, a possible path is to convert the nitrate, albeit usually slowly, to say N2 gas. A source https://www.sciencedirect.com/science/article/abs/pii/S00108...

"The nitrate ion has high chemical stability, especially at low concentrations. Standard reduction potentials indicate that it should serve as an excellent oxidizing agent, but in order to react with suitable reducing agents to form elemental nitrogen or ammonia, special conditions, such as catalysts and high temperature and pressure, are required. A review of the literature on the chemical reduction of nitrate in aqueous systems has found about a hundred articles dealing with nitrate removal from such systems, with the majority having been published over the last decade. The reducing agents which have been examined to the greatest extent for acidic solution are formic acid, iron metal, methanol and the ammonium ion; while for basic solution aluminum, zinc and iron metals, iron(II), ammonia, hydrazine, glucose and hydrogen have been studied."

So, reducing the nitrate by paths noted above leading to a measured volume of N2 could provide a slow estimate of the moles of starting elemental nitrogen.

Here, however, relatedly, I would recommend the reported fast action of Al with KNO3 in a dilute acidified solution with added NaCl and a trace of Cupric ion where the Sodium chloride serves two roles, an electrolyte and to salt out any created N2O. To quote a source https://www.researchgate.net/publication/215904038_Reaction_...

"Metallic aluminium was found not to react with either concentrated or diluted nitric acid. Providing the diluted acid contains dissolved sodium chloride and traces of copper(II) cations, a vigorous reaction occurs. The product is basically nitrous oxide (possibly containing some elemental hydrogen and nitrogen gases), and was identified by its IR spectrum."

To me here, the most interesting aspect is the likely underlying chemistry, in my opinion. In particular, the action of iron metal with dilute nitric acid, for example, is well-known to form N2O (likely as a result of associated radical formation). The Al reaction, itself, is fundamentally a Fenton type-reaction with dilute nitric acid (acting as a substitute for hydrogen peroxide in the normal Fenton system and Al for Fe) creating the associated hydroxyl radical (.OH). The presence of a small amounts Cupric implies depositions of copper metal on aluminum suface forming cathodic zones with favorably small surface area. This thereby forms a galvanic cell with a production of solvated electrons that react with H+ leading to the production of the hydrogen atom radical (.H). The presence of both these radicals appears to result in a rapid decomposition of the nitrate ion. In essence, this can be referred to as a Galvano-assisted Fenton-type reaction for the degradation of the nitrate ion.

Note: The full text of the article is available upon scrolling down. Further, the author appears to agree with me, at least, on the role of the hydrogen atom radical.
===========================

For those interested in my speculative take on the reaction mechanics:

NO3- + .OH = .NO3 + OH-
.H + .NO3 = HNO3 (or HONO2)
.H + HONO2 -> H2O + .NO2
.H + .NO2 = HNO2 (or HONO)
.H + HONO -> H2O + .NO
.OH + .NO -> HONO
.H + .NO = HNO
HNO + HNO = H2N2O2 (Creation of unstable Hyponitrous acid)
H2N2O2 -> H2O + N2O (which finally explains the formation of Laughing Gas)

Possible alternate pathways:

.H + .OH = H2O + Photon (Energy)
HNO3 + Energy = .OH + .NO2 (Hence, sensitivity of Nitric Acid to light inducing color change)
HNO2 + Energy = .OH + .NO (Hence, Nitrous acid is only stable in cold and dark conditions)
HNO2 <=> H2O + N2O3
N2O3 + Energy -> .NO + .NO2 (How Nitrous acid liberates radical gases)

Relatedly, explaining properties of Hypochlorous acid (HOCl):

2 HOCl <=> H2O + Cl2O (Smell of Hypochlorous acid is Cl2O)
Cl2O + hv -> .Cl + .ClO (Source: https://www.researchgate.net/publication/30768788_Primary_an... )
.ClO + .ClO = (ClO)2 (One of many pathways)
(ClO)2 + H2O = HCl + HClO3 (So, HOCl slowly with light exposure leading to chlorate, ClO3-, formation)
.ClO + .ClO + hv -> Cl2 +O2 (Alternate pathway, so HOCl, in strong light, can also liberates oxygen)

[Edited on 2-1-2022 by AJKOER]
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[*] posted on 1-1-2022 at 19:50


Quote: Originally posted by unionised  


But precipitating silver chloride and weighing it looks a much better option.



Titration with silver nitrate is far quicker. You can use K2CrO4 or fluorescein as indicator.

[Edited on 2-1-2022 by Bedlasky]




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