NA from Ca(NO3)2 - Archive File

[megalomania]

the_wingman
Frequent Poster
Posts: 49
From:
Registered: JAN 2001
posted 03-31-2001 03:26 PM
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I'm not able to obtain nitric acid so I'll try this soon:
Ca(NO3)2 + H2SO4 => 2 HNO3 + CaSO4
Calcium nitrate is freely available as fertilizer.
Has sb tried this before? Is CaSO4 ( which is gypsum) insoluble in nitric acid or will the reaction give HSO4(-)ions (which will stay in solution) ?



pete
Frequent Poster
Posts: 56
From: u.k
Registered: JAN 2001
posted 04-02-2001 02:01 PM
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without the distillation step the nitric acid that you get will be quite impure, and not that concentrated. However the answer to your question is that no the sulphate will not cpmplex in the nitric acid, however it will be difficult to remove fully, and there will be the formation of an equalibrum in the reaction, so yeild will also be affected.


PHILOU Zrealone
Frequent Poster
Posts: 479
From: Brussels,Belgium,Europe
Registered: SEP 2000
posted 04-06-2001 05:48 AM
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Ca(NO3)2 is very hygroscopic so it means you can dissolve a lot in a little water. Add H2SO4 conc and distil.
------------------
"Life that deadly disease sexually transmitted".
"Chemistry is all what stinks and explode; Physic is all what never works! ;-p :-) :o)"



the_wingman
Frequent Poster
Posts: 49
From:
Registered: JAN 2001
posted 04-06-2001 11:13 AM
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I think this is not completely correct. If you mix Ca(NO3)2 with a little water you won't get a solution but a hydrated form of it ( for example Ca(NO3)2.2H20 )

[gliper]

Do's anyone know if Ca(N03)2 + heat = NO2 + CaO? If so you would have pure HNO3 bubling NO2 through H2O.

also CaO + cole + heet = CaC2 and CaC2 +H2O = welding gas and CaO

[Hang-Man]

CaO + cole + heet = CaC2 and CaC2 +H2O = welding gas and CaO

lets just go over a few equation pointers shall we.

if by cole you mean coal, write 'C' coal is carbon
use '-->' instead of '="
don't write X=Y=Z
Balance your equations
don't use generic terms like 'welding gas' I assume that is CO, but you need to write that in formula.
Explain odd statements like this: 'CaC2 and CaC2 + H20' ? do you mean its hydrated?

[Anthony]

CaC2 is calcium carbide. Reacted with water, it releases acetylene, which is often used for flame welding.

[YayItGoBoom!]

Ahhhhh, my friend and I have been wanting to know how to make CaC2, never thought I could learn something from someone with poor grammer and even worse chemistry annotations ;) . Hangman, what gliper is trying to say is:

2Ca0 + 4C --(heat)--> 2CaC2 + O2

then

2CaC2 + 2H20 --> 2Ca + 2C2H2 + O2

Though I find it very interesting that Carbon will replace out Oxygen, it seems like CaO would be more stable than CaC2 but I guess not. Anyways thanks that has been bothering me for a while.

[gliper]

First CaO is more stable.
I still don't know if Ca(NO3)2 + heat = Cao +NO2 or not. so don't get your hopes up
and what I read was that KNO3 + heat = K2O + NO2. This thed was about Ca(NO3)2 and geting

old so I thew that idea in.

Can anyone conferm (X)(NO3)x + heat = (X)xO?

[Tuatara]

Well, I know for certain that heating Cu(NO3)2 will release NO2, having found out by accident while trying to crystalise my newly made copper nitrate! At that point I had molten nitrate, not a solution, without realising. (Stupid)

[Guerilla]

Found these two
Ca(NO3)2 => CaO + 2 NO2 + 1/2 O2 + 369 kJ/mol
Ca(NO3)2 => CaO + 2 NO2

The latter is probably more reliable. Anyway, the nitrate should start decompose somewhere at 500C degrees. I think I will try this out soon, first drying the Ca(NO3)2 so that there wont be no water reacting with NOx as I'm only interested in CaO, and then heating the nitrate in a flask till no more NOx is released.

What sort of temps should be used in the CaO + C process, and would it be better to be done in an oxygen free atmosphere?

[gliper]

Thanks everyone! I first read that CaO burnt with coal = CaC2 but later some one said he tried that and it diden't work. Doesen't sound likely because the O2 should just oxadize it back to CaO so O2 redused seems more resonable.

Three notes 1 this will still release CO and CO2 so you can't just seal it. 2 I remember that a normal fame is hot enough. 3 After CaC2 reacts with water it reverts back to CaO so the reaction can be rerun with more coal and water.

The only remaning question I can think of is, will the acetylene ignite from the heat of its formation?

[gliper]

Update: Having looked aroumd CaC2 + H2O aparently will not ignite.

[YayItGoBoom!]

Update: Having looked aroumd CaC2 + H2O aparently will not ignite.

Yes, if you try and light a wet pile of CaC2, it won't do crap. But if you collect the gas from this reaction, this is acetylene gas, and sure as hell is flamable! I know this is going way off topic but I don't want to start a new topic on this.

If you have a burning pile of CaO and carbon, combustion would prevent all oxygen from bonding with the calcium. To keep it from recombining with oxygen, you would have to stuff out the rest of the oxygen. I am thinking this would be done in a crucible in a metal foundry set up, and quickly cover the top of the crucible with a lid.

[rikkitikkitavi]

US3572991...

During preliminary feasibility test with decomposing CaN (in an open steelpan, outdoors) by heating with a blowtorch , first to anhydrous CaN (about 300 C ) and then further , I got only a little of NO2. I probably used to much heat since the
Ca(NO)3 => CaO + 2 NO2 + O2 continues with 2NO2=> 2NO + O2.

The NO will quickly react with O2 forming NO2 again , especially if done in a closed vessel where O2 content is higher than in air. (from the decomposition)

Since I had only a crude method of heating and decomposition starts to catch up first when the salt is molten I assume to much heat decomposed the NO2 and combined with a low decomposition rate diluted the NO in the surrounding air so much that I couldnt see the NO2 forming . But there was an obvious NO2 smell (careful boy, careful :)

As comparison I tried Cu(NO3)2 *5 H2O . First it melted in its own crystalwater, then at about 200 C heavy NO2 smoke formed (which I didnt tried to smell) in the same pan, same day e t c.

My conclusions are:

It is possible to produce NOx by decomposing CaN

It will take some time to think up of some sort of steel cruicible that is sealed , preventing gases from escaping. It should be easy to recharge with fresh CaN.

Some sort of thermostatic control of heat is needed.

There should be enough gas volume in the system allowing NO react to NO2 (O2 is on excess anyway)

the gas should be cooled (promotes NO2 forming and further N2O4) and absorbed in H2O to yeild HNO3.



BTW I have tried CaN(dried at 200 C ) with H2SO4. What a fucking mess... easier with KNO3 or NaNO3...

/rickard

[Rosco Bodine]

Another way in which Calcium Nitrate or Calcium Nitrate / AN hydrated
double salt ( Norsk Hydro manufacture ) might be used is to mix
a hot fairly concentrated solution of it with a similar solution of
Epsom Salt , to form a solution of Magnesium Nitrate and a precipitate
of Calcium Sulfate which is filtered out . The Mg Nitrate solution
could be boiled down to a concentrate , useful as a dehydrating agent
for breaking the azeotrope of 68 per cent nitric acid mixed with the
Mg Nitrate and distilled . This Mg Nitrate "syrup" could be used the same way
as sulfuric acid is used for the same purpose , and recycled by heating
it further to drive off the water before it is used for a new distillation
of nitric .

Alternately , by treatment with sulfuric acid , and distilling the acid / nitrate mixture in the usual way , nitric acid could be produced
directly from the Mg Nitrate , and the byproduct would be Mg Sulfate ,
which , after being rehydrated , is the same Epsom Salt
as is used in the beginning of the process ,
being regenerated for reuse in such a synthetic scheme .

If the presence of Ammonium Nitrate in the original CN / AN double salt
should complicate the process , then boiling the solution with an excess
of lime until no more ammonia is evolved , should convert any AN to CN
so that the remaining solution is entirely a CN solution . Excess lime
could be filtered out before use and reused in subsequent conversions .

[simply RED]

I've made conc HNO3 using the hydrated AN/CN fertilizer + H2SO4. And distilling. It works.
It does not work if you don't crush the prills.

[gliper]

If you just want CaO it can also be prepared from CaCo3, easyer to find as chalk, lime and soil sweetener than Ca(NO3)2. Just add heat, bypoduct is CO2.

[markgollum]

If all you want is CaO then just buy it, I think that Home depot stocks it as lime.
Lime is used to keep the smell down in outhouses. There really isint a way for a backyard hobbyist to make CaC2 via CaO(L) + 3C(S) --> CaC2(L) + CO(g) as the temp needed is very high, (the reaction starts to happen at 1600 degC).

[gliper]

All the lime I have ever seen is CaCO3 (or that and MgCO3 in Dolomite lime), I have heard of slack lime as Ca(OH)2. I recall store bought charcoal gets 2,000 degs C.

[Mumble]

Calcium Oxide and Calcium Hydroxide are kind of interchangeable for the term lime. There is Magnesium impurity in the comercial product. CaCO3 probably too, as its formed from CO2 from the atmosphere acting on CaO. I think all they do is crush lime stone, filter out the rocks and bag it up. I think the general comercial product is around 90% Ca(OH)₂. If you want CaO specifically try burnt lime or quicklime. Hydrated or slaked lime for Calcium Hydroxide.

[Centimeter]

I plan to use calcium nitrate to make iron (III) nitrate by mixing solutions of calcium nitrate and iron sulfate. I beleive this will precipitate out calcium sulfate leaving iron (III) nitrate in solution. Iron sulfate can be purchased at almost any gardening store including home depot. I then plan on using the iron (III) nitrate to produce 68% HNO3 by decomposing it and bubbling the NOx through water. Of all the metal nitrates I looked up, iron (III) nitrate had the lowest decomposition temperature; ~100*C. Furthermore, the byproducts of the reaction can be disposed of w/o problems of contaminating your water. The reason I would go through all this trouble is so that I can distill fumming nitric acid w/o waisting the H2SO4 as it can then be recycled. Does this sond like it will work?

[Centimeter]

Update: I made a sample batch of calcium nitrate by adding calcium hydroxide to excess nitric acid then evaporating. I then used this to try and make some metal nitrates. I first tried making iron (III) nitrate by mixing a solution of iron sulfate (probably in the form of ferrous sulfate) and a solution of Ca(NO3)2. A calcium sulfate precipitate clearly formed, however when I filtered and evaporated off the water it formed an orange syrup rather than a purple solid. I attempted to decompose anyway. It began to smell very much like nitric acid fumes however there was no visible NO2. Iron (III) nitrate is a bad choice as its low decomposition point makes it difficult to dry properly and it stains one's glassware. I next tried making aluminum nitrate by mixing a solution of alum of potash and a solution of calcium carbonate. A calcium sulfate precipitate clearly formed and was filtered off. After drying in an oven, the substance was heated to decomposition. This reaction once again produced a strong nitric acid smell (including what looked like HNO3 vapor) however no visible NO2 was produced. At this point, I don't think that decomposition of metal nitrates is a very viable method of producing HNO3. I will however make some more aluminum nitrate and try to collect the observed vapors in my soon to arive distillation kit. I will titrate it and see just how concentrated the acid vapor is. I guess I will just have to accept that H2SO4 is going to have to be used in a non-recyclable way. But why am I whinning when I just got a job at the hardware that I buy my H2SO4 from so now I get a 20% discount. :D

[ProdigyChild]

It began to smell very much like nitric acid fumes however there was no visible NO2. Iron (III) nitrate is a bad choice as its low decomposition point makes it difficult to dry properly and it stains one's glassware.


Centimeter, chemistry sometimes demands a bit of patience!

I've done e very similar experiment (described yesterday in the 'concentrating n
itric by decomposition...' thread). I dissolved a screw in dilute HNO3 (over a d
ay). Then I poured the dirty brown liquid into a flat plastic container and put
it on the heating (50° at most) for over 24h. The Fe(NO3)3 is then completel
y dry with no significant decomposition. I dried indoors although I must have had some HNO3 traces left!

Decomposition of the final product is pure fun! In a test tube obvious brown gas after heating with a candle!

[Centimeter]

What color was your product; mine was definately not purple like the MSDS says it should be. My main problem with the iron method is that it stained all of my glass. Now I have to boil all of my glass in HCL. I must say that the aluminum nitrate method is a much easier synthesis and I was able to easily get NOx to evolve. It is also easily cleaned up. The only problem is finding alum that's not ammonium alum. The only alum I can find in the market is ammonium alum. I was however able to utilize a large bag of potash alum that I acquired from a crystal growing kit. I do admit that I am a sucker for haste and aluminum nitrate is conducive to this tendency. I guess I will give the iron (III) nitrate method another shot.

[tmp]

Centimeter, let me get this straight. You used HNO3 to make a nitrate so that
you can make HNO3 ? Why not just take your existing HNO3 and vacuum
distill it ? On the question of Ca(NO3)2 - melting point for the anhydrous
version is 561 C according to both the 52nd and 81st editions of the CRC.

[Centimeter]

Yup, you got that right. My source of calcium nitrate fertilizer is a long ways away so I have to save up a lot of money before I go there so that I can make the trip worth while. I have plenty of nitric acid, I just wanted to do some experimenting. After all, it will take me a couple weaks to save up enough money. It all sounds logicle in my head! :p

[lucas]

NOTE: All distillations done under reduced pressure obtained by water-jet vacuum pump. Most were conducted using 500ml round bottom flask.

When distilling potassium nitate with sulphuric acid, I typically use 1:1 mass ratio. 100g Sulphuric acid and 100g potassium nitrate has yielded typically approx 60g nitric acid, which is a high yield. This is close to 1:1 molar ratio. Using 1:1 mass will result in a thick pasty mixture of potassium nitrate and sulphuric acid. This mixture will distill but doesn't allow easy movement of heat throughout the mixture. It results in the mixture acting like a solid, "melting" at the source of heat before distilling. It provides a pale-bright yellow nitric acid, depending on particular conditions. The first vapours are deep red, but this soon clears and quickly fades as nitric acid becomes the main gas flowing. The late stages of the distillation involes a molten mass of KNO3, H2SO4 KHSO4 which is over well 100 degrees C. Since the HNO3 vaporises as it forms, it has little time to decompose and the acid distilled at this stage is nearlyas pure as in the earlier stages of distillation. The final waste product, mainly composed of KHSO4, can be poured out while molten, into a very large amount of cold water, using heat resistant gloves, but beware, it will still be evolveing nitric acid fumes at this stage. Cleaning with water is effective in removing the waste material. Adding cold water to a very hot glass vessel is not advised as cracking may occur!

Use of excess KNO3 is effective in maximising yield of nitric acid with respect to sulphuric acid used. I have obtained 112% yield in moles HNO3 per mole H2SO4, without any more trouble than otaining 100% I used approximately 20% excess of nitrate. This is done by distilling excess potassium nitrate and forcing, with heat, a second proton to move from the hydrogensulphate ion to the nitrate which will then vaporise as nitric acid. This means the mixture used to start with will be a particularly thick paste. Even when using reduced pressure the temperature of vapour passing into the stillhead will be 90 deg C or more later in the distillation. The colour of the condensate is still pale yellow with this method.

Use of just below, ie 20%, molar ratio of potassium nitrate to sulphuric acid means starting with a viscous fluid which tends to foam on heating. This yields a slightly better product than the above mixture, but lower final yield, wrt sulphuric acid consumed. I do not enjoy distilling such a mixture as the control of the heat source is a troublesome affair, trying to maintain vapour production while minimising foaming to a level which won't cause the mixture to overflow.

On one occation I used a separatory funnel instaed of a round bottomflask to recieve the distillate. I tapped off fractions of the acid. Definate differences in colour were noticed. The acid distilled soon after the initial red fumes cleared were the least coloured. Colouration intensified only slowly as the distillation continued.

All the acid batches produced by this method gave off large amounts of white fumes, even when chilled to 0 deg or -5 deg C. The measured density was always around 1.50 g.cc^-1

[The Order]

A journal article exists, studying the solubility of calcium sulfate dihydrate, or gypsum, in nitric acid:

Abstract The solubility of gypsum in 0-30 wt % nitric acid solutions at 20°C was studied. The data were compared with the results obtained at 60°C.

The journal article appears in the Russian Journal of Applied Chemistry, and can be purchased here:

http://www.springerlink.com/content/mm823x123582j158/

Although I am too cheap to buy this article at the moment, I would be quite interested to know the results. From everything I've gathered so far, calcium sulfate and nitric acid do not form any sort of matrix of crystals when existing together in solution/slurry.

If equimolar parts of Ca(NO3)2 and H2SO4 were combined (calcium nitrate in saturated solution), I wonder if, after cooling to discourage solubility, the gypsum could then be filtered through diatomaceous earth (filter sand) placed in a piece of fiber-glass cloth (the kind used for autobody, boats ), and this inside of a glass/plastic funnel which is directed into a glass bottle.

This method, if practical, could offer an easy OTC method for nitric. If one were to start with say, canning lime (Ca(OH)2) and instant cold-paks (NH4NO3), they could then evaporate off the ammonia water, and be left with relatively pure Ca(NO3)2, especially if excess lime was added, then filtered before evaporating. Of course, the aqua ammonia could also be distilled off and salvaged, but of course if such equipment to do so were present in the first place, then distilling nitric would be the first course of action.

Any thoughts?

[megalomania]

Well, look at the handy picture I just found. What a serendipitous coincidence...

Solubility of Calcium Sulfate Dihydrate in Nitric Acid at 20oC
Russian Journal of Applied Chemistry, Vol. 76, No. 1, 2003, pp. 156 3157. Translated from Zhurnal Prikladnoi Khimii, Vol. 76, No. 1, 2003, pp. 162-163.

http://img139.imageshack.us/img139/5523/journalmr3.th.gif (http://img139.imageshack.us/my.php?image=journalmr3.gif)

[Enkidu]

A journal article exists, studying the solubility of calcium sulfate dihydrate, or gypsum, in nitric acid:

If equimolar parts of Ca(NO3)2 and H2SO4 were combined (calcium nitrate in saturated solution), I wonder if, after cooling to discourage solubility, the gypsum could then be filtered through diatomaceous earth (filter sand) placed in a piece of fiber-glass cloth (the kind used for autobody, boats ), and this inside of a glass/plastic funnel which is directed into a glass bottle.

Any thoughts?

You're right, gypsum isn't very soluble in concentrated nitric acid. Unfortunately, there's no gypsum in a calcium nitrate / sulfuric acid solution.

First, you'll have very little water in your solution, assuming you're using the theoretical 98% conc. sulfuric acid. So, CaSO4.(H2O)2 isn't likely.

You won't even have any calcium sulfate (CaSO4) in your solution. You'll have calcium bisulfate - Ca(HSO4)2.

Look at the pKa's for nitric acid vs. a bisulfate. Unless you begin removing nitric acid from the equation, you're not getting any sulfates.