Romain
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Synthesis of calcium nitrate using calcium hydroxide and ammonium nitrate
Good morning everyone,
Today I came up with a way to make calcium nitrate by boiling a solution of stoichiometric amounts of calcium hydroxide with ammonium nitrate. Ammonia
would be generated and evaporate, thus driving the reaction to completion. I would be left with a calcium nitrate solution given enough time. Since
calcium hydroxide has very low solubility in boiling water (0.77g/l according to Wikipedia.fr on "hydroxide de calcium"), I was wondering: how slow
would the reaction be if we have 2 moles of ammonium nitrate and 1 mole of calcium hydroxide in, say, 500 ml of water? If it takes a year for them to
react, it's not a viable route.
The reason is seek such a complicated way to produce calcium nitrate is because: I don't have access to calcium nitrate fertilizer, I don't have
access to nitric acid, I can't synthesize it, I only have CaCl2, NaOH, and NH4NO3. The first two are combined to make the Ca(OH)2.
The reaction are:
1) CaCl2 + 2NaOH -> 2NaCl + Ca(OH)2, which is filtered to keep the Ca(OH)2.
Then, in boiling water:
2) Ca(OH)2 + 2NH4NO3 -> 2NH4OH + Ca(NO3)2, ammonia boils off the solution.
3) NH4OH -> NH3 + H2O
Tell me what you think!
R.
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Oscilllator
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I can tell you from experience that filtering the Ca(OH)2 will be nigh on impossible unless you use dilute, hot solutions to maximize the particle
size. Even then , it will take ages. One way to achieve this without using lots of water is to get boiling, plain water and then add both solutions
dropwise. That way the concentration of both reagents will always be very low and so you will achieve the same effect as using many litres of water.
Be sure to rinse the precipitate with lots of water to try and get out all of the NaCl, as I imagine the Ca(OH)2 will absorb it fairly well, given its
small particle size.
Just out of curiosity, what are you going to use the Calcium Nitrate for? I can't think of any uses for it off the top of my head.
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Romain
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Ok, thanks for your advice, I'll try adding it dropwise, that will save me some trouble!
I'll be using the calcium nitrate to make copper nitrate: CuSO4 + Ca(NO3)2 -> Cu(NO3)2 + CaSO4.
I'm using copper nitrate to make other nitrate such as lead, iron, via simple metathesis reactions: Pb + Cu(NO3)2 -> Pb(NO3)2 + Cu.
You just have to filter the solution to separate the copper powder and your metal nitrate solution. Again, I do all this just because I don't have
nitric acid or copper nitrate. Maybe someday I'll be lucky enough to find some in an hardware store (not gonna happen though).
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blogfast25
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Quote: Originally posted by Romain  | Ok, thanks for your advice, I'll try adding it dropwise, that will save me some trouble!
I'll be using the calcium nitrate to make copper nitrate: CuSO4 + Ca(NO3)2 -> Cu(NO3)2 + CaSO4.
I'm using copper nitrate to make other nitrate such as lead, iron, via simple metathesis reactions: Pb + Cu(NO3)2 -> Pb(NO3)2 + Cu.
You just have to filter the solution to separate the copper powder and your metal nitrate solution. Again, I do all this just because I don't have
nitric acid or copper nitrate. Maybe someday I'll be lucky enough to find some in an hardware store (not gonna happen though). |
Calcium hydroxide isn't hard to get: slaked lime is sold for building purposes, as a powder or as a slurry/paste.
I did prepare Ca(OH)2 once from CaCl2 and NaOH and used several decantations to get rid of the NaCl. Not very hard to do.
Calcium nitrate is a fairly useless chemical and very hygroscopic, so hard to dry and keep dry.
[Edited on 17-7-2014 by blogfast25]
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Romain
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I agree that I could buy calcium hydroxide, I'm sure I can find some, but I don't have any at the moment and I'm too lazy to look for Ca(OH)2.
Is calcium nitrate harder to keep dry than ammonium nitrate? But honestly I don't need it as a dry powder since I'm going to make copper nitrate from
it. Maybe an hour in the oven at 120°C will do though.
What about the process of making it via boiling? Do you think it will work?
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blogfast25
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Probably, yes.
Bear in mind that the displacement reaction Ca(NO3)2/CuSO4 isn't perfect: CaSO4 is poorly soluble but not completely insoluble.
Cu(NO3)2 is another nitrate that is difficult to dry. There a thread on it in 'Beginnings'.
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violet sin
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what about the possible dry reaction of the the two with a little heat? I seem to remember a vid on youtube using NaOH and ammonium nitrate both dry
and mixed in a pyrex brownie tray. ya know 12" x 10" with about 2" lip height. with a drop or two of water and a little heat, the ammonium nitrate
reacted with the sodium hydroxide to make ammonium hydroxide and sodium nitrate. it was pushed towards completion by the heat kicking ammonia gas
out. so you were already left with some water, no need to add any. or further more, calcium carbonate and ammonium nitrate? probably need some water
in that one, not dry. unstable ammonium carbonate decomposes pushing it to the right. the carbonate would be pretty easy to get, just a thought, on
lunch break so no time to look into these.
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Romain
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I thought about calcium carbonate and ammonium nitrate but calcium carbonate is very insoluble so I would probably take ages to react. I may give it a
try though someday.
blogfast, you pointed that the reaction won't be complete, do you know of any way to push it toward completion? I'd try to concentrate the solution to
precipitate a maximum of calcium sulfate (knowing that copper nitrate won't crystallize since it's hygroscopic). The concern is that if I use the
copper nitrate/calcium sulfate solution to make lead nitrate, lead sulfate will also form and it will passivate the surface of the lead metal. Or it
could also (ideally) precipitate with the copper metal and I'd be left with pure lead nitrate, so maybe it's not a problem.
I would experiment but I'm not at home at the moment (holiday).
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blogfast25
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| Quote: Originally posted by Romain | | blogfast, you pointed that the reaction won't be complete, do you know of any way to push it toward completion? |
No, not really. Displacement reactions involving calcium sulphate always lead to a reaction product containing small amounts of CaSO4. You could try
and recrystallize the Ca(NO3)2 several times: unfortunately that's hard to do because of the very high solubility of calcium nitrate: 121 g / 100 g of
water @ 20 C, acc. Wiki.
Try boiling in your calcium nitrate solution until the first crystals appear, the cool and chill. Most of the calcium sulphate should now be in the
supernatant liquor.
[Edited on 18-7-2014 by blogfast25]
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AJKOER
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One could also compare yield by reacting Mg(NO3)2 with Ca(OH)2, and filtering out the Mg(OH)2.
One can prepare Magnesium nitrate, not pure, by reacting readily available KNO3 with MgSO4 and cooling to separate out the K2SO4.
[Edited on 22-7-2014 by AJKOER]
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violet sin
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would carefully adding barium nitrate drop most of the sulfate contamination with a far less soluble ppt, help here?
BaSO4 solubility 0.0002448 g/100 mL (20 °C)
CaSO4....."....... 0.24......... g/100ml at 20 °C (dihydrate)
and at the same time leaving a less bothersome addition in minor proportion?
Ca(NO3)2aq + CuSO4aq --> Cu(NO3)2aq + CaSO4ppt + tiny bit CaSO4aq
then
CaSO4aq + Ba(NO3)2 --> Ca(NO3)2 + BaSO4ppt ( with Cu(NO3)2
spectating )
leaving, for all intents and purposes, an all nitrate solution with only minor calcium content.
[Edited on 22-7-2014 by violet sin]
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AJKOER
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This paper may be of interest as it explores solubility changes in CaSO4 in a multi-salt solution, "Modelling of calcium sulphate solubility in
concentrated multi-component sulphate solutions", by G.Azimia,V.G.Papangelakisa,J.E.Dutrizacb, August2007, available online at www.sciencedirect.com, Fluid Phase Equilibria 260(2007)300–315, link to full text: https://www.google.com/url?sa=t&source=web&rct=j&...
Apparently, there is some increase in solubility of CaSO4 in the presence of an acid in a multi-salt solution as compared to a pure water solutions of
CaSO4.
[Edited on 22-7-2014 by AJKOER]
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Refinery
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I was gonna make nitric acid by this method. My route was following:
Mix molar rations of ammonium nitrate and calcium hydroxide and bring to boil to evaporate all ammonia gas, which is collected into cold water.
Resulting mixture is brought to room temp, decanted and filtered and then evaporated to concentrate. Formed calcium nitrate, containing calcium
hydroxide impurity, is then heated sufficiently to drive off the rest of the water, and then moved into SS reactor, where temp is brought to over
650C, when decomposition will take place, releasing NO2 fumes. Fumes are ran through tube into cold water, where hydrolysis will take form, resulting
in HNO3 and NO fumes. Finally HNO3 is concentrated to 68% and distilled over magnesium nitrate to obtain red/white fuming product.
Calcium hydroxide is sold as slaked lime, construction lime or lime milk at very low rates. If this is for some cause available, get some calcium
carbonate sold as garden chaulk and heat it up to 900-1000C with gas or (char)coal fire to make it calcium oxide and then mix it with water to make
lime. It should be very easily available. I am a lot more worried about availability of ammonium nitrate, it is really gone off counter in my country,
as well as most of other nitrates. Some customer stuff is available yes, but it is diluted with all kind of nasty additives, including urea and other
nitrates which are very water soluble, and the rest containing ammonium nitrate yields some 10-20% of mass max. This will come extremely expensive..
[Edited on 22-7-2014 by Refinery]
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AJKOER
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Per a recent thread, it was observed that heating dry Potassium nitrate (in excess) with sugar in an opened container (a confined space would likely
produce an explosion) followed by the addition of HCl produced a gas resembling NO2. To quote precisely from thread http://www.sciencemadness.org/talk/post.php?action=reply&...
| Quote: Originally posted by SimplyChem16 | I had recently performed the "smoke bomb" reaction between Potassium Nitrate, and sugar. After the reaction had completed, a thick paste of
yellowish/white material covered the reaction vessel. I had read this was Potassium Carbonate, so I decided to try to neutralize it simply using
Hydrochloric Acid. Well, much to my surprise, lots of brown looking gas was evolved, and frankly, moved my butt away from that thing as quickly as I
could. Could anybody out there provide an explanation to what happened? 
(Nitrogen Dioxide Perhaps?) |
The likely reaction as described by Zyklon in that thread was presented as:
2 KNO3 + H2CO (relative empirical formula for sucrose) → CO2 + H2O + 2 KNO2
2 KNO2 + 2 HCl (aq) → 2 KCl + 2 HNO2 ↔ 2 HNO2 → NO2 + NO + H2O
and, adding air or O2, forms more NO2:
2 NO + O2 → 2 NO2
Since you are starting with a nitrate as well, you may wish to explore this path.
Note, you should not attempt this reaction with NH4NO3 (explosion hazard and the nitrite is not the likely final product, see http://en.m.wikipedia.org/wiki/NH4NO2 ). As such, I would boil aqueous NH4NO3 with NaHCO3, and harvest the NaNO3 for the above synthesis.
[Edited on 24-7-2014 by AJKOER]
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hissingnoise
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| Quote: | | Finally HNO3 is concentrated to 68% and distilled over magnesium nitrate to obtain red/white fuming product. |
Anhydrous Mg(NO3)2 is hard to find and it's a one-shot desiccant as the hydrated form cannot be dehydrated . . .
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PHILOU Zrealone
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CaCl2 and Ca(OH)2 may complexate NH3 to form the thermolabile (NH3)2CaCl2...under heating it frees NH3 gas...
So Ca(NO3)2 maybe also does complexate two NH3 molecules to form (NH3)2Ca(NO3)2
PH Z (PHILOU Zrealone)
"Physic is all what never works; Chemistry is all what stinks and explodes!"-"Life that deadly disease, sexually transmitted."(W.Allen)
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