testimento
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Citric acid + nitrate = HNO3?
Since reaction between hydrochloric acid and a nitrate will produce nitric acid and a chloride, I thought could the HCl be substituted with another
acid, like citric acid, which is not watched by the terror office in the extent it does with common acids.
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bismuthate
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Citric acid is way too weak. You should just use phosphoric or sulfuric acid.
[Edited on 10-11-2013 by bismuthate]
[Edited on 10-11-2013 by bismuthate]
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Agricola
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Write down the reaction of citric acid and nitrate giving nitric acid and citrate, and also the reactions of each acid with water. See how the
equilibrium constant (EC) of the first reaction is related to the EC's of the other two reactions.
The equilibrium favors citric acid and nitrate not nitric acid and citrate. If you try to distill the reaction mixture to remove nitric acid and shift
the equilibrium to make more nitric acid, I think the citric acid will be destroyed by decomposition or oxidation.
What I wrote above is true if the cation present doesn't give any insoluble citrate. In that case, you would have a possibility of success. Write down
the reaction of solution of the insoluble salt and see how it changes the EC of our main reaction. Use the search engine to find a paper by Thompson
on the production of sulfuric or phosphoric acids from oxalic acid and calcium sulfate or phosphate.
The Merck Index says 1 part calcium citrate is soluble in 1050 parts cold water and is insoluble in alcohol.
[Edited on 10-11-2013 by Agricola]
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Random
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Agricola would that mean calcium nitrate would form nitric acid and calcium citrate with citric acid?
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Fantasma4500
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NaHSO4 + Ca(NO3)2 have been talked about working.. plante1999 talked about this on the skype call connected to sciencemadness, said he might be able
to dig up the message
NaHSO4 can be made from NaHCO3 or Na2SO4 with H2SO4 of pretty much any concentration
otherwise you could get H2SO4 to boil down, and then add Ca(NO3)2 to it and let it stand, then decant, and the rest you pour through a nitrocellulose
filter, as it wont react with the HNO3 supposedly up to 99% concentration
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Agricola
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Quote: Originally posted by Random  | | Agricola would that mean calcium nitrate would form nitric acid and calcium citrate with citric acid? |
Maybe, some experiment along the lines of Thompson's paper would have to be carried out to check that.
An excess of citric acid would help to push calcium citrate out of solution due to the common ion effect.
[Edited on 10-11-2013 by Agricola]
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Random
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| Quote: Originally posted by Agricola | | Quote: Originally posted by Random | | Agricola would that mean calcium nitrate would form nitric acid and calcium citrate with citric acid? |
Maybe, some experiment along the lines of Thompson's paper would have to be carried out to check that.
An excess of citric acid would help to push calcium citrate out of solution due to the common ion effect.
[Edited on 10-11-2013 by Agricola] |
Wouldn't that also mean H2CO3 would drive nitric out yet it doesn't happen?
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Agricola
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| Quote: Originally posted by Random | | Quote: Originally posted by Agricola | | Quote: Originally posted by Random | | Agricola would that mean calcium nitrate would form nitric acid and calcium citrate with citric acid? |
Maybe, some experiment along the lines of Thompson's paper would have to be carried out to check that.
An excess of citric acid would help to push calcium citrate out of solution due to the common ion effect.
[Edited on 10-11-2013 by Agricola] |
Wouldn't that also mean H2CO3 would drive nitric out yet it doesn't happen? |
You are thinking of the common ion H+, I suppose. Nitric acid is a very soluble, non-gaseous, strong acid. Carbon dioxide is the opposite of all that.
Its a rigged equilibrium fight.
Everything that was said about citric acid is also qualitatively true for carbon dioxide. Unfortunately, carbon dioxide is much less soluble than
citric acid. An interesting experiment is to put a piece of dry ice (or pump CO2) in a calcium nitrate solution and see if calcium carbonate
precipitates.
[Edited on 10-11-2013 by Agricola]
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Random
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I could try it with citric acid but I only have ammonium nitrate. I can mix that with ca hydroxide to make ca nitrate but there will also be ammonium
ion present in solution. Then I could add citric.
I guess higher pressure for co2 would shift equilibrium.
[Edited on 10-11-2013 by Random]
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testimento
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In case of calcium nitrate, in my opinion it will be cheaper to bring up the temperature to 650C to drive out the NO2 and dissolve it into water to
get HNO3. It's a waste of citric acid.
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Random
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| Quote: Originally posted by testimento | | In case of calcium nitrate, in my opinion it will be cheaper to bring up the temperature to 650C to drive out the NO2 and dissolve it into water to
get HNO3. It's a waste of citric acid. |
Citric acid is like 1.5$ per half a kilo.
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Agricola
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| Quote: Originally posted by Random | I could try it with citric acid but I only have ammonium nitrate. I can mix that with ca hydroxide to make ca nitrate but there will also be ammonium
ion present in solution. Then I could add citric.
I guess higher pressure for co2 would shift equilibrium.
[Edited on 10-11-2013 by Random] |
Try it, Random. I don't think the presence of ammonium ion will make any difference.
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testimento
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I'm again very sorry to tell that I don't live in United States of OTC where you can just go and buy everything at dirt cheap rate from your local
OTC-market. I'm so unfortunate to tell that in my country commodities like most acids and ALL chemical reagents as they are, are unavailable for
common public. I have found citric acid from few stores which sell it for minimum 3 times higher rate you stated, and the more expensive one is nearly
6 times. But it's yet better than non-available-at-all.
[Edited on 12-11-2013 by testimento]
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Random
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| Quote: Originally posted by Agricola | | Quote: Originally posted by Random | I could try it with citric acid but I only have ammonium nitrate. I can mix that with ca hydroxide to make ca nitrate but there will also be ammonium
ion present in solution. Then I could add citric.
I guess higher pressure for co2 would shift equilibrium.
[Edited on 10-11-2013 by Random] |
Try it, Random. I don't think the presence of ammonium ion will make any difference. |
Actually it is going to make it, when I add citric acid to ammonium hydroxide and calcium nitrate solution, first it will form ammonium citrate and
then react with calcium nitrate precipitating calcium citrate before acidic conditions. IMO only way would be using pure calcium nitrate. I'm not able
to heat the solution with ammonia to evaporate it due to lack of apparatus. Now if there is another way to destroy ammonia without heating I would be
interested.
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Agricola
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| Quote: Originally posted by Random | | Quote: Originally posted by Agricola | | Quote: Originally posted by Random | I could try it with citric acid but I only have ammonium nitrate. I can mix that with ca hydroxide to make ca nitrate but there will also be ammonium
ion present in solution. Then I could add citric.
I guess higher pressure for co2 would shift equilibrium.
[Edited on 10-11-2013 by Random] |
Try it, Random. I don't think the presence of ammonium ion will make any difference. |
Actually it is going to make it, when I add citric acid to ammonium hydroxide and calcium nitrate solution, first it will form ammonium citrate and
then react with calcium nitrate precipitating calcium citrate before acidic conditions. IMO only way would be using pure calcium nitrate. I'm not able
to heat the solution with ammonia to evaporate it due to lack of apparatus. Now if there is another way to destroy ammonia without heating I would be
interested. |
An excess of ammonium nitrate and later an excess of citric acid will guarantee acidic conditions.
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violet sin
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Would the ammonium nitrate/ calcium hydroxide method work like nurdrage NH4NO3 + NaOH --> NaNO3 + NH4OH ?
http://m.youtube.com/watch?v=hQJhf_24-QM
Meaning fast and favored reaction. He added a fair amount of water but i seem to remember a different YouTube vid that just mixed the two as dry
pellets and added a few ml of water to kick it off( or just stirr till it goes). the ammonia escapes Leaving water afterward, and allowing more dry
reagents to react. So no heat was actually required. Though it helped. And absolutely min water( only from reaction). Would mean minimal ammonia
trapped
http://m.youtube.com/watch?v=bvVozS7A8io
This vid is as close as I could find to the one I referred to.
If it worked and not much ammonia was present then a litle excess of citric acid may be enough. If your worried about it, let the Ca(NO3)2 sol sit
out overnight or in the sun tomorrow. Can't imagine ammonia will stick around too long in an open container. The. Add the citric acid.
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papaya
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A short question to aforementioned post:
NH4NO3 + NaOH --> NaNO3 + NH4OH
Is heating needed at all? Can one preserve NH4OH as well (when intended to get NaNO3 as a main product) by leading gases to ice water and capture NH3,
will you be able to get at least 10% solution ? Also, when bicarbonate is used instead of base it forms solid ammonium (bi?)carbonate, which will
decompose on heating - if the gases CO2 and NH3 are lead to ice water, can one obtain ammonium (bi)carbonate or only hydroxide?
Thanks.
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AJKOER
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Those acquainted with hydrometallury are acquainted with the use of Citric acid which is apparently becoming more popular as it is environmentally
more green. To quote a recent source (August 2011, Volume 13, Issue 2, Journal of Material Cycles and Waste Management, pp 118-126, "Leaching of heavy
metals by citric acid from fly ash generated in municipal waste incineration plants", by Kai Huang,.. at http://link.springer.com/article/10.1007%2Fs10163-011-0001-5... ):
"From the results of screening tests of leaching lixiviants, citric acid was found to be the most effective leaching agent"
This is quite an interesting statement as among the competing lixiviants where the inorganic acids HCl, H2SO4 and even HNO3.
Now, as to why, I would add some advice from very old chemistry textbooks that note, if I recall correctly, the role of volatility and solubility
issues, in addition to acid strength, in working with acids. Other explanations (see http://books.google.com/books?id=QmbfLX4TgGEC&pg=PA249&a... ) include its ability to chelate metals owing to the presence of three carboxyl
groups and one hydroxyl group in its structure. Other, more generally cited characteristics of acids include thermal stability and dipole movements
(see, for example, http://www.transtutors.com/chemistry-homework-help/s-and-p-b... ).
[Edited on 14-11-2013 by AJKOER]
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Random
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So I did the experiment and as i already predicted, nothing happened at all. Excess of citric acid and copper did not appear to be anywhere near
dissolving.
I did it in a hurry though and didn't manage to filter the solution before adding citric.. But a very small bit of copper should dissolge if there was
nitric present.
I doubt any calcium citrate formed. Solution did heat up a little bit after adding citric though.. Probably from reaction with ammonia.
[Edited on 14-11-2013 by Random]
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Agricola
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| Quote: Originally posted by Random | So I did the experiment and as i already predicted, nothing happened at all. Excess of citric acid and copper did not appear to be anywhere near
dissolving.
I did it in a hurry though and didn't manage to filter the solution before adding citric.. But a very small bit of copper should dissolge if there was
nitric present.
I doubt any calcium citrate formed. Solution did heat up a little bit after adding citric though.. Probably from reaction with ammonia.
[Edited on 14-11-2013 by Random] |
What were the amounts of reactants used and what exactly was the procedure employed? Why do you say you didn't filter the solution before adding
citric acid, not all the calcium hydroxide had dissolved? Maybe not enough ammonium nitrate and stirring time were used. Did you check the pH of the
solution at each step?
If you have time, please evaporate the final solution down at low heat until crystals appear then let it cool.
Copper should not dissolve even if nitric acid was present. You need concentrated nitric acid, the solution at that point was at best a very dilute
one.
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violet sin
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@papaya
I'm not sure heating is needed, but it facilitates the loss of NH3 so you only get the nitrate. Which was the point of the vids. I am sure you could
trap the ammonia in cold water, which would be nice. Not the most expensive reagent, but waste not want not right.
It wasn't until I read this thread I remembered I have a couple pounds of barium and strontium carbonate. Could be fun to get their nitrates for
colorfull amusement and sulphate sequestering with the Ba salt. Only have a little Ba(NO3)2 left so I was saving it for something special.
As for the ammonium bicarbonate, I'm not sure. I have little experience with that, and would not be the one to ask there. I had planned on making
the hydroxides first, but I do remember reading a few things that refered to solutions of ammonium salts and metal carbonates. They were usually
heated to drive off the ammonia, as its carbonate is easily destroyed. I don't think it was done dry though. Maybe some one with more experience
there may be able to answer that
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Random
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| Quote: Originally posted by Agricola |
What were the amounts of reactants used and what exactly was the procedure employed? Why do you say you didn't filter the solution before adding
citric acid, not all the calcium hydroxide had dissolved? Maybe not enough ammonium nitrate and stirring time were used. Did you check the pH of the
solution at each step?
If you have time, please evaporate the final solution down at low heat until crystals appear then let it cool.
Copper should not dissolve even if nitric acid was present. You need concentrated nitric acid, the solution at that point was at best a very dilute
one. |
Actually I mixed KAN fertilizer with Ca(OH)2 and added very little water. Almost all Ca(OH)2 dissolved with time and it smelled of ammonia. Wanted to
filter the solution but there was not enough water, then I added more water and filter paper broke so I used the stuff as is. Added citric in great
excess and added a tiny bit of copper. Overnight nothing changed. Concentrated nitric as far as I know passivates copper while more dilute should
dissolve it. I wish I had pure calcium nitrate so I could check for aure. Would calcium nitrate solution ever evaporate alone by sitting, since it's
hygroscopic.
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Agricola
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| Quote: Originally posted by Random | | Quote: Originally posted by Agricola |
What were the amounts of reactants used and what exactly was the procedure employed? Why do you say you didn't filter the solution before adding
citric acid, not all the calcium hydroxide had dissolved? Maybe not enough ammonium nitrate and stirring time were used. Did you check the pH of the
solution at each step?
If you have time, please evaporate the final solution down at low heat until crystals appear then let it cool.
Copper should not dissolve even if nitric acid was present. You need concentrated nitric acid, the solution at that point was at best a very dilute
one. |
Actually I mixed KAN fertilizer with Ca(OH)2 and added very little water. Almost all Ca(OH)2 dissolved with time and it smelled of ammonia. Wanted to
filter the solution but there was not enough water, then I added more water and filter paper broke so I used the stuff as is. Added citric in great
excess and added a tiny bit of copper. Overnight nothing changed. Concentrated nitric as far as I know passivates copper while more dilute should
dissolve it. I wish I had pure calcium nitrate so I could check for aure. Would calcium nitrate solution ever evaporate alone by sitting, since it's
hygroscopic. |
Thanks for the details, Random.
I think you are mistaken about the reaction of copper with nitric acid. It is the oxidizing properties of concentrated nitric acid that dissolve
copper, not the acidic properties of dilute nitric acid. Compare the reduction potentials:
2H+ + 2e => H2 + 0.00V
Cu2+ + 2 e− => Cu +0.34V
NO3- + 2 H+ + e- => NO2 + H2O +0.80V
http://en.wikipedia.org/wiki/Standard_electrode_potential_(d...
Are you familiar with the cycle of copper experiment? The experiment begins with the dissolution of copper in concentrated nitric acid. See this movie for example. I performed this experiment years ago during graduation. We used zinc powder instead of aluminum foil.
I would dissolve the NH4NO3, filter or decant anything that is left undissolved, then repeat that with Ca(OH)2 in the same solution. I would use
enough NH4NO3 so as to keep the solution strongly acidic.
Then I would add a solution of citric acid. I would evaporate the solution at low heat until a precipitate forms, let it cool, and filter or decant
the solution. The process of adding citric acid, heating, etc., would be repeated until no more precipitate forms. Heating is necessary for
evaporation when the solute is hygroscopic.
Lastly, I would distill the solution or at least evaporate it at low heat until I could check if it behaves like concentrated nitric acid.
Don't forget the equilibrium equations: maybe the ionization of nitric acid beats the insolubility of calcium citrate so badly as to make the process
worthless.
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