Sciencemadness Discussion Board - Ca(NO3)2

Ca(NO3)2

prayser - 24-5-2006 at 08:04

Hi! I have bought a bag of "kalisalpeter" which contains Ca(NO3)2, which i didn't know. Can i do anything to make KNO3? Sad if the bag goes to waste

12AX7 - 24-5-2006 at 09:09

Yeah, grab a bag of potassium chloride (chloride of potash), dissolve as much in boiling water as you can, and do the same for the calcium nitrate. Mix the hot solutions, then cool, to freezing if possible. Filter KNO3 crystals, rinse and dry. Solution still contains some KNO3 and lots of CaCl2.

Tim

ordenblitz - 24-5-2006 at 10:42

I would make a slightly less than saturated solution "at ~99 deg" of your mixed nitrates then place in a well insulated or temperature controlled container "thermos" such that you can lower the temperature very slowly "days" to say ~5 deg. This will form large crystals of the potassium nitrate that will be fairly pure.

The key is to lower the temperature very slowly and steady.

chloric1 - 24-5-2006 at 14:44

I disagree about the large crystals being of high purity. Actually, for example, when potassium nitrate is made by metasythetic reactions and large crystals are formed, more mother liquor is trapped in the crystals. I use a magnetic spin bar and chill to 0 celcius with rapid agitation. This gives me a fine "crystal meal" which contains much less mother liquor.
Oh and since your saltpeter will have other ions besides potassium and nitrate, a second recrystalization from the purest water is most prudent.

ordenblitz - 24-5-2006 at 16:16

It depends on how fast the crystals form.
If you cool fast, impurities are captured. If you do it slowly, actual contiguous crystals are quite pure. A quick rinse with cold water takes most of the rest of the impurities away.

I have been to a medium sized potassium nitrate manufacturer in Valencia Spain that does just this very thing to purify their KNO3. Its made from the reaction of KCl and NaNO3 in solution. Their production is intended for the chemical industry and pyrotechnics. I personally observed a building full of crystallizing vats slowly cooling, forming lovely crystals from the mother liquor loaded with NaCl impurities.

First round yields 97.5% KNO3... second crystallization, 99%+

[Edited on 25-5-2006 by ordenblitz]

Organikum - 24-5-2006 at 18:54

Use it for growing pot.

franklyn - 9-6-2006 at 05:11

I know that HTH pool chlorine ( Calcium Hypochlorite ) added to

Acetone yields chloroform CHCl3. Reaction and intermediates are :

CH3COCH3 + Ca(ClO)2 -> 2CH3Cl + CaCO3

2CH3Cl + 2Ca(ClO)2 -> 2CHCl3 + 2Ca(OH)2

So why not use Acetone on Ca(NO3)2 to get Calcium Nitromethane

CH3COCH3 + Ca(NO3)2 -> 2CH3NO2 + CaCO3 -> Ca(CH2NO2)2 + H2O + CO2

A tantalyzing reference from a material safety data sheet :

Nitrate (NO3-)

Mixtures of the nitrate with organic solvents, after the solvent

has evaporated, the remaining solid may have explosive properties.

also _

The dry alkali salts of nitromethane are shock sensitive and the

sodium salt bursts into flame upon contact with water

_ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _

I know also that Ammonium Hydroxide added to metal perchlorates

produces metal Amine Chlorate.

So why not use Ammonium hydroxide on Ca(NO3)2 to get Calcium Nitroamine

2NH4OH + Ca(NO3)2 -> Ca(HNNO2)2 + 4H2O

Chemical data indicates solubility with the above mentioned reagents

http://www.chemnet.com/show/tyxlchem/eproduct/00029185.html

_ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _ _

First off I have not tried this , and I'm not a chemist.

Those of you who are , can say if this is safe or not.

Despite appearances I still have all my fingers.

So cleaning my glass eyes is not a problem

Potassiumcyanide - 9-6-2006 at 09:07

you can get Potassiumnitrate by putting Potassiumcarbonate into a solution of your Calciumnitrate.

Ca(NO3)2+K2CO3 --> CaCO3+2KNO3

Since Calciumcarbonate is not soloutable in water, you will only have to filter the product to get Potassiumnitrate.

If you want to get Sodiumnitrate just use Sodiumcarbonate instead of Potassiumcarbonate.

Nicodem - 9-6-2006 at 09:45

Quote:

Originally posted by franklyn
First off I have not tried this , and I'm not a chemist.

That is quite obvious from your post. Not a single one of your equations makes any sense whatsoever. They are all completely wrong and quite depressing for chemists. Reading books is advised. Study, study, study ... or just stay away from chemical experiments.
And using MSDS sheets as a source of chemical properties is truly one of the worst choices.

DeAdFX - 9-6-2006 at 19:52

Quote:

Originally posted by franklyn

So why not use Ammonium hydroxide on Ca(NO3)2 to get Calcium Nitroamine

2NH4OH + Ca(NO3)2 -> Ca(HNNO2)2 + 4H2O

You get Ammonium Nitrate and Calcium Hydroxide... Nitroamine exsists but there is no calcium nitroamine.

franklyn - 10-6-2006 at 11:27

Quote:

Originally posted by Nicodem
Not a single one of your equations makes any sense whatsoever. They are all completely wrong and quite depressing for chemists. Reading books is advised. Study, study, study ... or just stay away from chemical experiments.
And using MSDS sheets as a source of chemical properties is truly one of the worst choices.

HMmmm yes that really clears thing up for me. It's remarkable how

I passed the term Without such enlighted guidance huh. Trouble is

I don't have a second lifetime to learn the secret handshake.

I can always send order to hire a consultant who has all the

answers like yourself.

U P D A T E _

I don't usually troll but I'll have you know that I am not alone

in being uninformed. I plucked this gem from the

Argonne National Laboratory's Newton site "Ask a Scientist".

The person asking the question knows more than the "scientists"

who answered it. But the Ph.D got it almost right , Duh

http://www.newton.dep.anl.gov/askasci/chem00/chem00776.htm

Chloramine and Hydrazine are clearly indicated as products when

mixed with ammonia on every container of bleach I've seen.

Oh sure MSDS are the worst source of relevant chemical data why

else would they be made available ? Much better to get the smart

academic answer.

Being a chemist you should know that any conceivable arrangement

of the reagent atoms is possible , and actual , in infinitesimal

amounts , due to things like hydrated electrons from cosmic rays.

The dominant reaction is only the most probable. So in this respect

the two learned gentlemen referenced above can be said to be

correct , at least in part. It is inconsistent to hold me to a

different standard.

The haloform reaction I guessed at , is much more complex than it

needs to be. I'll have to have a word with god about that and see

if it can be changed.

.

[Edited on 12-6-2006 by franklyn]

MSDS

MadHatter - 10-6-2006 at 12:05

Franklyn, can you point me to that particular MSDS sheet ? I was wondering if the manufacturer
made an error or the reactions you stated occur under other than normal conditions. It's known
that nitrates don't mix well with some metals because of the unstable amines formed. I'm not
quick to dismiss what could be an unintended side reaction.

Pyrovus - 11-6-2006 at 00:29

Calcium nitrate is very nice stuff, so I most definitely wouldn't throw it out. Owing to the fact that calcium carbonate and calcium sulphate are both highly insoluble, you can use it to prepare all sorts of nitrates, provided you have a soluble carbonate or sulphate of the cation you want the nitrate of.

For instance, if you need ammonium nitrate because the government's banned it, you can make it thusly:
(NH4)2SO4 + Ca(NO3)2 -> 2NH4NO3 + CaSO4

Another interesting possibility is copper nitrate:
CuSO4 + Ca(NO3)2 -> Cu(NO3)2 + CaSO4

Or why not make some nitric acid:
H2SO4 + Ca(NO3)2 -> 2HNO3 + CaSO4

[Edited on 11-6-2006 by Pyrovus]

neutrino - 11-6-2006 at 06:56

Don't expect that much from that last reaction. There was a lot of debate about how to prepare dry nitric acid like that and they eventually found that it was just impossible. I doubt even dilute acid could be prepared without significant Ca2+ and HSO4- contaminations.

The main problem with the dry acid route is that the calcium sulfate precipitate is very voluminous and virtually impossible to filter out, forming a moist cake at best. This route of preparing nitric acid is just a waste of calcium nitrate and nitric acid, nothing more.

The other problem is that at very acidic conditions the following happens:

CaSO4 (s) + H+(aq) <---> Ca2+(aq) + HSO4-(aq)

Some calcium sulfate dissolves in the nitric acid formed, reducing the yield of nitric acid as well as its purity.

12AX7 - 11-6-2006 at 07:29

Ok, so double the acid amount.

I bet if you add the acid (or nitrate) slowly and with heating, you can get a larger crystal precipitate.

If nothing else, you can do it in dilute solution, wash the precipitate (easier since the solution isn't going to instantly nitrate your filter, although I can't imagine the filter will much appreciate it anyway), then evaporate or distill it down, or react with something else, e.g. potash, to get a non-hydrated salt to work with (from which you can do the usual H2SO4 + KNO3 = HNO3 + KHSO4 distillation or whatever). (Of course such a roundabout route begs the question why you can't just go directly to, say, NaNO3(aq) + CaCO3(s) in one step.)

Tim

Here's another great idea _

franklyn - 20-7-2006 at 21:27

Methyl alchohol and Oxalic acid will give you Methyl Oxalate

adding Calcium Nitrate to this should precipitate out Calcium

oxalate leaving a fairly pure solution of Methyl nitrate. If this

is dangerous I'm sure someone here will point it out.

.

woelen - 20-7-2006 at 23:16

Franklyn, this is not going to work. Making methyl oxalate may work, but I'm quite sure that does not give methyl nitrate with a solution of calcium nitrate. You mix up the properties of salts and esters.

not_important - 21-7-2006 at 19:52

Problem with ammonium sulfate and calcium nitrate is that calium salts tend to get more soluable in solutions of ammonium salts; plus CaSO4 is slightly soluable to start with.

We did this many decades ago, as kids. Ended up evaorating the filtrate to dryness, and then dissovling it in boiling alcohol and filtering. A lot of trouble, but you could use the NH4NO3 in place of HNO3 when making nitrates for fireworks colours, just boil a solution of it with the oxide, hydroxide, or carbonate of the metal you want the nitrate of.

Later on, when we got old enough that we could buy battery acid, we did go the Ca(NO3)2 route to HNO3. Started out with slightly diluted battery acid, heated the solution for awhile to get a little tighter ppt, decanted; washed the ppt to get water to make up the next batch. Reduced pressure distallation to get to c.b. HNO3, just never distill to dryness. And we weren't after 100% HNO3, constant boiling was OK.

The methyl nitrate thing won't work the way you want it to.

neutrino - 22-7-2006 at 07:32

Quote:

Problem with ammonium sulfate and calcium nitrate is that calium salts tend to get more soluable in solutions of ammonium salts; plus CaSO4 is slightly soluable to start with.

Calcium sulfate is close enough to insoluble.

How do calcium salts become more soluble in the presence of ammonium salts? I could see some salts becoming a bit more soluble in the presence of free ammonia, but ammonium ions don't have that complexing ability.

not_important - 22-7-2006 at 11:57

When we were making NH4NO3, we were getting a few grams of CaSO4.aq per liter, we guessed 5 to 15 grams. Didn't know exact amount because this was long before cheap, sensitive electronic balances; our homebrew for handling production amounts wasn't very sensitive.

Lange's Handbook : "CaSO4 2H2O s a, gly, Na2S2O3, NH4 salts" If you've a decent scale, you can find out for youself the exact amount of enhancement

When you're making nitrates for coloured flames, doesn't take too much calcium to mess up other colours. And when we were making HNO3, the solubility was high enough to be a bother during distillation.