Sciencemadness Discussion Board - Sodium nitrate synthesis from salt

Sodium nitrate synthesis from salt

servo - 28-12-2021 at 01:46

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
Servo

Microtek - 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.

servo - 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]

Sulaiman - 28-12-2021 at 05:01

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

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

unionised - 28-12-2021 at 06:56

The mixture will have a lower melting point.
And the stuff will decompose when you melt it.

servo - 28-12-2021 at 11:02

Yup I can vouch for that the mixture decomposed at 210 c

caterpillar - 29-12-2021 at 00:45

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

servo - 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

Texium - 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.

clearly_not_atara - 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.

Rainwater - 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]

Syn the Sizer - 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]

AJKOER - 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]

S.C. Wack - 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.

servo - 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

Amos - 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 :/

Texium - 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?

Bedlasky - 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.

macckone - 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]

Rainwater - 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.

S.C. Wack - 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.

servo - 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

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

unionised - 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

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.

AJKOER - 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]

Bedlasky - 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]

AJKOER - 2-1-2022 at 07:06

Quote: Originally posted by Bedlasky

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

[Edited on 2-1-2022 by Bedlasky]

Well, perhaps a bit faster than the reportedly fast action of Al with KNO3 in a dilute acidified solution with added NaCl and a trace of Cupric ion.

However, the latter is definitely cheaper and further benefits as readily available reagents.

Also, AgNO3 is a light sensitive reagent, which adds 'insult' to the 'injury' just suffered by your bank account. Albeit, one can further purchase NaOH to recycle your AgCl to Ag (but this is still not AgNO3 which requires yet another purchase, or preparation, of HNO3).

[Edited on 2-1-2022 by AJKOER]

Bedlasky - 2-1-2022 at 19:29

Quote: Originally posted by AJKOER

Well, perhaps a bit faster than the reportedly fast action of Al with KNO3 in a dilute acidified solution with added NaCl and a trace of Cupric ion.

And then what? How do you find how much nitrate do you have?

Quote: Originally posted by AJKOER

However, the latter is definitely cheaper and further benefits as readily available reagents.

Also, AgNO3 is a light sensitive reagent, which adds 'insult' to the 'injury' just suffered by your bank account. Albeit, one can further purchase NaOH to recycle your AgCl to Ag (but this is still not AgNO3 which requires yet another purchase, or preparation, of HNO3).

[Edited on 2-1-2022 by AJKOER]

I just pointed out, that rather than bothering with precipitation, filtration, washing and drying light-sensitive AgCl, it is better and much quicker to perform titration. And AgNO3 isn't that sensitive like you suggest. If it would be so sensitive, why would be argentometry so widely used method?

Cheaper alternative is bromatometry. You can precipitate CuCl, wash it, dissolve in HCl and titrate it with KBrO3 solution. However, this particular mixture also contain nitrate, which will oxidize CuCl. There is other option, to precipitate Cu2O instead of CuCl. Nitrate won't oxidize Cu(I) in alkaline solution. But there are also ammonium ions, which will be converted to ammonia in alkaline solution, which dissolve your Cu2O to form soluble [Cu(NH3)2]+. So you see why this method cannot be used.

[Edited on 3-1-2022 by Bedlasky]

Fyndium - 3-1-2022 at 01:28

Big processes benefit from large amounts. When you load 10 000 tons of each in a reactor pool, and add that every day, you can keep the Solvay cycle running and precipitate the desirables in plenty amounts. But you still keep that base amount dissolved in there.

So, for amateur who wants a few kg of something once, it would be detrimental to get 25kg of each precursor and eventually lose most of them in the process.

Also, sodium nitrate is not "one of the easiest" to synthesize in practice. Nitrates may not be directly available in pure form at all without certain troubles.

Anyway, I usually classify stuff in 4 categories, which are 1: always bought, 2: may be bought but could be made if necessary, 3: always made, and category 4: stuff that is just too hard to obtain, dangerous, expensive or illegal to work with.

AJKOER - 3-1-2022 at 15:35

Bedlasky:

You may be right as here is a supporting SM thread https://www.sciencemadness.org/whisper/viewthread.php?tid=82... noting the presence of impurities is likely a key factor for the induced light sensitivity of AgNO3.

Further agreeing comment at https://www.largeformatphotography.info/forum/showthread.php...

"By itself, or mixed in distilled water, silver nitrate is not light-sensitive. Combined with halide salts, it becomes very light-sensitive, and combined with organic materials, somewhat light-sensitive."

However, a mass of others with a diverging opinion, for example, per Google, the first provided answer at https://www.vedantu.com/chemistry/silver-nitrate (which, albeit, I will admit is likely, in my opinion, not #1 for technical accuracy alone), to quote:

"Silver Nitrate is very sensitive to light. This means, that the chemical will react when exposed to light. So, when this is left exposed to sunlight or any bright light, it will start to hydrolyze. This will result in the formation of black or brown colored silver oxide and nitric acid."

Also, someone once asked, and provided answers at https://www.quora.com/Why-is-silver-nitrate-stored-in-black-... :

"Why is silver nitrate stored in black bottle in laboratory?"

Clearly a switch is turned to account for this difference in narrative.

As to what causes the issue, I suspect, that it is first, the liberation of Ag metal (via light), and also the creation of Ag2O/AgOH (via transition metal impurities and oxygen exposure), which together can apparently form a powerful photocatalyst: "The formation of visible light-driven Ag/Ag2O photocatalyst with excellent property of photocatalytic activity and photocorrosion inhibition" here https://www.sciencedirect.com/science/article/abs/pii/S00219... .

[Edited on 4-1-2022 by AJKOER]

AJKOER - 4-1-2022 at 19:34

Just found a supporting reference to my suggested path to account for the lack of stability in (or following) light exposure of AgNO3, here https://www.researchgate.net/post/How_does_silver_nitrate_be... "Photochemical reduction of silver nitrate to silver nanoparticles in the presence of stannous chloride was interestingly described in the attached article." where the cited article notes the present of select organics can result in a stable solution of nano-particles. Also, see "The extra cellular synthesis of gold and silver nanoparticles and their free radical scavenging and antibacterial properties".

I would note that the formation of nano-particles can foster the creation of hydroxyl radicals, see for example "Use of mine waste for H2O2-assisted heterogeneous Fenton-like degradation of tetracycline by natural pyrite nanoparticles: Catalyst characterization, degradation mechanism, operational parameters and cytotoxicity assessment" at https://www.sciencedirect.com/science/article/abs/pii/S09596... , to quote:

"Radical scavenger tests demonstrated that •OH was the main oxidizing agent generated by both solution and surface phase reactions."

And I would claim:

Ag -> Ag+ + e-
•OH + e- -> OH-

Net: Ag + •OH -> AgOH (See Eq 2 here https://pubs.acs.org/doi/pdf/10.1021/nn503459q )

2 AgOH (unstable) -> Ag2O + H2O

So, in summary, light exposure of AgNO3 leads to Ag which in the presence of an organic leads to stable nano-particles. The electrostatic properties of such suspension in the presence of dust with transition metals promotes/recycles a Fenton or Fenton-like reaction in the presence of oxygen (see https://www.ncbi.nlm.nih.gov/pmc/articles/PMC2626252/) forming associated •OH radicals. Per above this leads to Ag2O and with Ag a possible powerful visible light photocatalyst (Ag/Ag2O). The photocatalyst leads to electron holes converting water into a source of •OH radicals, etc, leading to the degradation of the AgNO3.

My suggested explanation accounts for the role of sunlight, organics, metal impurities, air exposure and a newly induced photo sensitivity of the AgNO3.

[Edited on 5-1-2022 by AJKOER]

Bedlasky - 4-1-2022 at 23:15

I really don't understand where is problem. If you store silver nitrate in brown closed bottle in dark place, it is stable for prolonged time. And if some silver is formed in solution, you can standardize it with NaCl anytime. There are much more problematic solutions in volumetry (like FeSO4, TiCl3, CrCl2 etc.) and they are still used for analysis.

And I don't really get, how colloidal silver is related to storage of AgNO3 solution. Do you read that paper? They maintain silver colloidal using cetyl trimethyl ammonium bromide. But cetyl trimethyl ammonium bromide (or another surfactant) isn't present in AgNO3 solution for analysis.

AJKOER - 5-1-2022 at 03:10

Bedlasky:

My understanding is that AgNO3 is, somewhat generally, reported as less stable in the presence of light and organic and "air" exposure (where the latter includes not just ambient air but dust particles).

My presumption is that photochemistry, properties of nano particle suspension and air interface could possibly, with associated induced radical formations, therein provide a possible path to an understanding of why there could be a change in the photosensitivity of AgNO3.

In the work by Hongxai Yu et al, it is noted that nano silver is a photocatalyst. Also, nitrate is a weak photocatalyst (UV). So, if nano silver is introduced via associated light exposure, the solution becomes more photo sensitive. As whether organics increase photo sensitivity, likely yes (see my extended comment below). Air exposure suggests to me surface chemistry inducing eventually H2O2 and Ag2O formation (more details below). Together with photo liberated Ag, a yet more powerful visible light photocatalyst construct (Ag/Ag2O).

As per this source https://www.sciencedirect.com/science/article/abs/pii/S00456... Natural organic matter (NOM) "promoted the formation of hydroxyl radical, induced the generation of triplet-excited state NOM and thus greatly enhanced the indirect photolysis" . With respect to Dissolved Organic Matter (DOM), per a source https://infoscience.epfl.ch/record/223453?ln=en to quote: " is a complex mixture of thousands of organic molecules ubiquitously present in surface waters. Its influence on the photodegradation of organic contaminants is a complicated topic. For decades DOM had been mostly assumed to have a positive effect (i.e., promoting the degradation of contaminants), but more recently it was discovered that for some compounds, DOM can also have negative effects (i.e., inhibiting the photodegradation of contaminants) in surface waters". Lastly, a source https://www.researchgate.net/publication/51379919_Production... notes: "Under UV irradiation, an important primary photochemical reaction of colored dissolved organic matter (CDOM) is electron ejection to produce hydrated electrons (e-aq)." As such, light (especially rich in UV) with added organics, in general, likely leads to the accelerated photo decomposition of AgNO3.

Note, with respect to alluded to surface chemistry, apparently, the near UV illumination of a suspension of ZnO, for example, forms measurable surface amounts of H2O2 (see https://pubs.acs.org/doi/10.1021/j100585a011). As to mechanics, Zinc oxide, which is a photocatalyst, releases electrons (e-) at the gas water interface with air. So I suspect, e- + O2 -> .O2- , the creation of the superoxide radical anion, which is transformed in a medium of air/water vapor to .HO2 (or at pH < 5). Further action at the interface: .HO2 + e- -> HO2- which can apparently rapidly borrows a H+ from even water to create the claimed surface H2O2. Then per Hongxai Yu et al, Eq (1) and (2), a product formation of AgOH ( --> Ag2O) in the likely related case of a suspension of nano Ag (a photocatalyst).

Note, my suggested chemistry is an attempt at providing a paradigm that could possibly explain the progression in AgNO3 from "little" light sensitivity to "high".

[Edited on 5-1-2022 by AJKOER]

servo - 14-1-2022 at 10:26

Quote: Originally posted by Fyndium

Big processes benefit from large amounts. When you load 10 000 tons of each in a reactor pool, and add that every day, you can keep the Solvay cycle running and precipitate the desirables in plenty amounts. But you still keep that base amount dissolved in there.

So, for amateur who wants a few kg of something once, it would be detrimental to get 25kg of each precursor and eventually lose most of them in the process.

Also, sodium nitrate is not "one of the easiest" to synthesize in practice. Nitrates may not be directly available in pure form at all without certain troubles.

Anyway, I usually classify stuff in 4 categories, which are 1: always bought, 2: may be bought but could be made if necessary, 3: always made, and category 4: stuff that is just too hard to obtain, dangerous, expensive or illegal to work with.

Well I never said I am making few grams, I am already making about 30 tons a month using Caustic soda and AN , I see an obvious benefit in making it from salt , I say I am an amateur because I have no formal education in chemistry but I have a chemical manufacturing Unit and I make chemicals for a living

I did say I want to challenge myself because it is a challenge but I am not doing for the fun of it but because i see obvious cost savings

yobbo II - 14-1-2022 at 13:00

What do you do with the Ammonia?

Yob

SWIM - 14-1-2022 at 19:15

I have an idea, but it may not be economically viable.

What is your cost per ton for ammonium nitrate, and what do you get per ton for sodium nitrate?

servo - 14-1-2022 at 21:45

@yobo

There is ample demand for liqour ammonia so we dissolve ammonia in water and sell it

@swim

Yes what is the idea ? it may give me some lead

cost per ton of AN is 228 USD
and Sodium Nitrate is about 680 USD currently costs me

[Edited on 15-1-2022 by servo]