chemkid - 28-4-2007 at 06:20
Hello,
I am looking to make boron nitrate. For this purpose i need ammonium chloride. Now i am aware that HCl and Ammonia will easily form ammonium chloride
however it forms as a mist and i cannot get the mist to deposit and form crystals. I cannot order ammonium chloride either. I need either advice as to
making it or where to buy it.
Thank you.
Nicodem - 28-4-2007 at 07:33
Boron nitrate? Are you sure any such thing exists at normal conditions? It would be one of the weirdest mixed acid anhydrides ever and I can't see why
it would not just decompose spontaneously. And how is NH4Cl supposed to be used in any such synthesis of this hypothetical compound? Or do you perhaps
mean boron nitride?
For NH4Cl, just slowly add 1.1 equivalents of ammonia of any concentration to 10% HCl(aq) in an ice bath and then strip off the solvent. Avoid using
higher HCl concentrations if you are unsure how to deal with exothermic neutralizations!
Why did you even try to combine them in the gas phase? NH4Cl is not particularly hygroscopic and can be crystallized from water without problem. I
truly see no point in your obsolete attempt in using NH3 and HCl gas.
Eclectic - 28-4-2007 at 07:46
It sounds like you want to make boron nitrIDE (white powder solid self lubricating graphite like material).
You might be able to make it from urea and boric acid if those materials are easier for you to get.
Urea and muriatic acid held just under boiling for a long time will make NH4Cl. I used an electric ceramic crock pot once to do that and it took
several days, topping up with acid occasionally.
chemkid - 28-4-2007 at 11:51
Sorry, i did mean boron nitride. I wasn't intentionally attempting to mix ammonia and HCl as gasses. However when they mix they form a white vapor.
Your saying all i need to do is mix 10% HCl and Ammonia and then evaporate off the water and the ammonium chloride will cystalize? I apologize for my
general lack of knowledge, i am new to chemistry.
Thank you very much
chemkid - 28-4-2007 at 12:13
However i would need to react the ammonia and HCl on a ice bath. And get the proportions right.
Nicodem - 28-4-2007 at 12:36
Well, yes NH4Cl is a simple salt of a weak base (NH3) and a strong acid (HCl). When you combine the two in water you get a solution containing
solvated ions, namely ammonium cations (NH<sub>4</sub><sup>+</sup>
which are a weak acid and the chloride anions (Cl<sup>-</sup> which are a very, very weak base (read about Acid-base reactions). In order to prepare NH4Cl, a crystalline compound, you need to induce crystallization. This is best done by removing the
solvent that keeps the ions separated by solvation though there are other methods as well. Removing the solvent means rotavaping it, boiling it off, letting it evaporate or whatever
appropriate. This also removes the 10% excess of NH3 used (better to use excess of the reagent that is removed easier; HCl is not as easy to remove as
NH3).
As for calculating the amount of aq. ammonia and hydrochloric acid, you first need to learn the most basic thing there is to know about reactions and
reagent amounts. It is called stoichiometry, so read this and the links cited therein.
vanderDecken - 28-4-2007 at 12:40
Ammonium chloride (aka sal ammoniac) is available in solid blocks from hardware stores, and in arts & craft shops that cater to the stained glass
crowd. It's used for cleaning the tip of a soldering iron.
garage chemist - 28-4-2007 at 12:52
I made my NH4Cl which I needed in fair amounts for platinum metal separations by combining concentrated HCl and NH3 solutions.
In order to avoid losses by evaporation of NH3 due to the heat produced, I intruduce the NH3 solution under the surface of the HCl by means of a
syringe with needle. That way no NH3 is lost.
The heat of neutralization will easily heat the solution so much that you cant touch the beaker with bare hands. Be careful, cool the solution if the
temperature approaches 100°C.
The resulting NH4Cl solution is placed in a flat plastic bowl and on the room heater. In a week at most, it will be completely dry.
Dont dry with stronger heat, NH4Cl tends to sublimate.
not_important - 29-4-2007 at 07:25
Borax and ammonium sulfate heated together will give some BN.
Aqua_Fortis_100% - 26-1-2008 at 05:50
I made mine some months ago using ammonium bicarbonate and muriatic acid.. Ammonium bicarbonate is very good choice because that is very cheap and
amke you boil much less water. All what you need is add NH4HCO3 slowly on muriatic until no CO2 foam is observed.. You should be careful because if
you add to much NH4HCO3 it will overflow..During the additions the solution will cool, because of the NH4+ solvatation. Also when you finishes adding
bicarbonate, you will get a red ppt from the solution.. This is Fe carbonate, I assume, which is very hard to remove with one or even two coffee
filter paper.. I've added some ammonium hidroxide to get rid off any acidity and filtrated many times and with some cotton on funnel bottom until
clear.
Then I put in a paint can and heated until almost all liquid was gone.. I didn't recrystallized that (I think that my product is still somewhat acidic
(HCl) , but I think which dont I really need that.. since I will not going to use in any smoke composition)
Yesterday I tried a slight modified "procedure" which I found in old THORP , inorganic chemical preparations for preparing sodium bicarbonate (
laboratory method for Solvay process), but I actually want is extract the NH4Cl ...
NaCl + NH4HCO3 ---> NaHCO3 + NH4Cl
58,5..........79...................84..........53.5
I've used 60g of NaCl in 180mL of water , 80 of NH4HCO3 in 170mL of water , mixed both solutions (the NH4HCO3 was freshly prepared (very cold)) and
after some minutes a large crop of ppt was obtained.. About 55g of NaHCO3 which is about 65% of theoretical. This means that I still have in about
350mL of solution approx.: 20,2g NaCl , 35g NH4Cl and residual ammonium (bi)carbonate which I assume will decompose when I boil the solution..
I'm only in trouble in how much water should I boil off and how much crystalisations should I do to get a purer product..(?) Has anyone any table or
so with mutual solubility for these compounds?
Thanks
[Edited on 26-1-2008 by Aqua_Fortis_100%]
not_important - 26-1-2008 at 06:35
It's going to be tough to clean up the ammonium chloride from the sodium salts with it. Several recrystallisations will be needed. The simplest way
may be to sublime the NH4Cl off the mixed salts.
An alternative to that NH4HCO3 + HCl is to put the ammonium bicarb in a glass or plastic bowl/beaker and cover it with distlled water, put the HCl in
another glass/plastic bowl, put both side by side in a plastic bucket, and cover it with a think plastic sheet of PE or PVC. The non-polar plastic
allows CO2 to diffuse through it fairly readily, while the more polar HCl has a more difficult time.
After a few days most of the HCl will have diffused over and reacted with the ammonium bicarb (or carbonate). The iron stays behind in the original
HCl container. If you have a slight excess of ammonium bicarbonate, heating the NH4Cl crystals to 100 C or so will drive off the bicarb but is low
enough that the chloride doesn't disassociate.
Aqua_Fortis_100% - 26-1-2008 at 07:07
Thank you very much n.i.!!! Great idea!
'Subliming' the NH4Cl from NaCl/NH4Cl mix is really promising! Producing it from very cheap NaCl and NH4HCO3 may be better (less $) than wasting HCl..
Although hard to separate.
Very good also the idea of plastic bucket... I will surely try that soon and post results latter..
But about the urea mentioned earlier in this thread:
How urea reacts with HCl to form NH4Cl , if its really a feasible reaction
The probable way should be mixing urea and Ca(OH)2 , heating and bubbling cautiously (to avoid suckback) the generated NH3 in HCl...
Any thinkings?
[Edited on 26-1-2008 by Aqua_Fortis_100%]
not_important - 26-1-2008 at 08:57
urea (NH2)2CO + H2O ==(heat+acid)==> 2 NH3 + CO2
then NH3 + HCl
urea is fairly stable in water below 60 C or so, hydrolysis to ammonia and carbon dioxide becomes more rapid with increasing temperature. The acid
will speed things up, too. Boiling helps drive the CO2 off, but also increases the loss of acid unless done under reflux.
MagicJigPipe - 26-1-2008 at 09:01
For some reason when I mix household ammonia and 35% HCl I never get enough heating to boil the liquid. In fact, I can still handle the beaker with
my bare hands. Maybe you guys are using .880 ammonia or some other higher concentration. Maybe, my HCl isn't 35%! I sure hope not. I will have to
weigh some.
chemkid - 26-1-2008 at 10:23
From memory, i actually don't ever recall having anything boil either.
Chemkid
chloric1 - 26-1-2008 at 16:22
I got more violent reaction adding ammonia to HCl as opposed to adding HCl to ammonia. Order of operations may be important here.
Aqua_Fortis_100% - 26-1-2008 at 18:54
Thank you again not_important!
I completely forgot of water reacting with urea.. It seems we have many cheap options to choose. Just Nice!
Hey , what about replacing NaCl in home scale solvay process by KCl? Theoretically we will got then KHCO3 , which can be very useful since it isn't
always OTC..
I was wondering this until now I searched for some KHCO3 solubility data and encountered an only problem: The solubility of KHCO3 is very high
comparing with NaHCO3..
Just check out Frogfot's pages:
K salts http://www.frogfot.com/stuff/sol-k.png
Na salts http://www.frogfot.com/stuff/sol-na.png
solubility for NaHCO3 @ 0°C is close to 6,9g/100g water whereas at the very same temp the potassium salt will be 22,1g/100g water (slight less than
KCl ( 28,1g/100g water)). So at least tinny amounts of KHCO3/K2CO3 will be reliable to made with this process and converting the KHCO3 to K2CO3 (or
not) , without bothering on extracting K2CO3 from wood ashes or so..
Maybe some extra CO2 added to that solution will be good to avoid K2CO3 formation since K2CO3 will have a very large solubility and will defeat that
process.. Also using a salted-ice bath can give even higher yields, although I dont know if the solubility curve for KCl at -20°C is "lower" than
KHCO3..
Any ideas?
[Edited on 26-1-2008 by Aqua_Fortis_100%]
JohnWW - 26-1-2008 at 23:51
Boron nitride, BN? More valuable than the graphite-like form of it is its diamond-like form, borazon, which is as hard as diamond and has similar
optical properties. It could be used as a substitute for gem and industrial purposes.
Also, by having a very slight excess of either B or N in it, diamond-like p-type and n-type semiconductors could be made, and this would be easier
than doping artificial diamond with B or N. However, because of the charges on the B and N atoms, slightly higher voltages would probably be required
to use them than for diamond semiconductors, for which the required voltages are themselves 5 times those for silicon ones.
[Edited on 27-1-08 by JohnWW]
Perfect Ammonium chloride process
SAM4CH - 14-2-2008 at 11:15
I made ammonium chloride from concentrated sodium chloride solution and added powder of ammonium bicarbonate "stoichemistry" and note you should add
NH4HCO3 to cold NaCl solution, stirr for 5min then suction filtration and the filtrate solution is boild for 1/5 of orginal volume then let it to
cool.. filter ur NH4Cl
16MillionEyes - 16-2-2008 at 19:28
Perhaps dumb of me to ask but how can you synthesize NH4HCO3? Would adding a solution of NaHCO3 and bubbling some ammonia do the job? Or perhaps
bubbling CO2 through an ammonia solution at cold temperatures? NH4HCO3 sounds like a very convenient compound to make ammonium salts and I certainly
like to get my hands on some.
not_important - 16-2-2008 at 22:16
Ammonium bicarbonate is used in baking, but may be difficult to find retail in the US.
Ammonia plus NaHCO3 would leave the sodium in what form?
CO2 passed into strong aqueous ammonia will make NH4HCO3, but it decomposes at 60 C or so, isolating it can be difficult. It can lose CO2 to give the
normal carbonate, and just plain decompose back into NH3 + CO2 + H2O. Continuing to pass CO2, and adding NH3 gas, might work well - getting a high
concentration of the bicarbonate some much of it precipitates.
This looks like it might be useful
http://www.sciencedirect.com/science?_ob=ArticleURL&_udi...