bigtrevs98
Harmless
Posts: 35
Registered: 14-4-2015
Member Is Offline
Mood: No Mood
|
|
Problems with Ammonium Chloride
So I was trying to make ammonium chloride. I put a 1:1 ratio of 30% HCl in with Pure NH3. I saw a couple white crystals form in the bottom but other
that nothing... what did I do wrong?
|
|
|
gdflp
Super Moderator
Posts: 1320
Registered: 14-2-2014
Location: NY, USA
Member Is Offline
Mood: Staring at code
|
|
What is "pure NH3"? Ammonia is a gas at room temperature. Also did you evaporate the water? If not, then the ammonium chloride is simply dissolved
in the water present in the hydrochloric acid and ammonia. Boil or evaporate to remove the water and you should be left with ammonium chloride
crystals.
|
|
|
bigtrevs98
Harmless
Posts: 35
Registered: 14-4-2015
Member Is Offline
Mood: No Mood
|
|
What I meant by pure NH3 was the store bought "pure ammonia" and are you talking about evaporating the water from the HCl? And if I don't evaporate
it... would I be able to evaporate the water from the ammonium chloride? Or is it basically ruined by the water?
And thanks gdflp I know it may get annoying answering questions that may partially make sense
But I appreciate you help. I'm learning lol
[Edited on 14-4-2015 by bigtrevs98]
|
|
|
ISCGora
Harmless
Posts: 47
Registered: 10-3-2015
Member Is Offline
Mood: No Mood
|
|
He meant evaporate water from final reaction because there was 70% water in your HCL.And you will get crystals because as you get rid of water stuff
starts falling out of the solution.and you end up with dry stuff.
[Edited on 14-4-2015 by ISCGora]
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Quote: Originally posted by bigtrevs98  | | So I was trying to make ammonium chloride. I put a 1:1 ratio of 30% HCl in with Pure NH3. I saw a couple white crystals form in the bottom but other
that nothing... what did I do wrong? |
Probably nothing. NH4Cl is very water soluble (Wiki: nearly 400 g/L water). Without boiling off much of the water, you won't get much solid NH4Cl.
[Edited on 14-4-2015 by blogfast25]
|
|
|
gdflp
Super Moderator
Posts: 1320
Registered: 14-2-2014
Location: NY, USA
Member Is Offline
Mood: Staring at code
|
|
Yes your ammonia is likely only a 5% solution of ammonia in water. You likely have a large excess of HCl in your solution, but that's fine as its a
gas as well. The solution you have from mixing the two is a solution of ammonium chloride dissolved in water with excess HCl, simply evaporate or
boil down the water to get crystals, the excess HCl will simply be expelled as a gas into the air so don't do this near any metal tools or in a living
area.
|
|
|
bigtrevs98
Harmless
Posts: 35
Registered: 14-4-2015
Member Is Offline
Mood: No Mood
|
|
Well its " LA's totally awesome, Pure Ammonia" is that a dilute? Or is it pure ammonia like it says it is?
|
|
|
gdflp
Super Moderator
Posts: 1320
Registered: 14-2-2014
Location: NY, USA
Member Is Offline
Mood: Staring at code
|
|
It can't be pure ammonia unless it is currently -33°C or below, which I can guarantee it's not, because ammonia is a gas at room temperature. I
assume that this is the stuff used for cleaning windows, if so then it is likely dilute, around 5% as I previously stated.
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Pure ammonia here means 'a solution of ammonia in water' (and nothing else: no surfactants, no odourants etc). Often the bottle will tell you the w%
NH3 in the product. 'Domestic' products are usually 5 - 10 w%. The theoretical maximum is about 33 % (with a density of about 0.9 kg/L).
Using really dilute NH3 is far from ideal to prepare NH4Cl because you'll need to boil off large amounts of water, after you've reacted the NH3 with
HCl.
|
|
|
bigtrevs98
Harmless
Posts: 35
Registered: 14-4-2015
Member Is Offline
Mood: No Mood
|
|
Well I looked all over the bottle. And it doesn't say... is there some kind of way to tell the concentration of the ammonia? Kinda also wondering this
for personal knowledge
[Edited on 14-4-2015 by bigtrevs98]
|
|
|
Metacelsus
International Hazard
Posts: 2539
Registered: 26-12-2012
Location: Boston, MA
Member Is Offline
Mood: Double, double, toil and trouble
|
|
Many.
Titration, density measurements, index of refraction measurements, etc.
Titration is probably the most accurate, and isn't too hard. Density measurements would be second best, but they require an accurate scale to be
useful.
|
|
|
bigtrevs98
Harmless
Posts: 35
Registered: 14-4-2015
Member Is Offline
Mood: No Mood
|
|
So what... could I freeze distill the ammonia? Or what would I do to go about purifying it?
|
|
|
gdflp
Super Moderator
Posts: 1320
Registered: 14-2-2014
Location: NY, USA
Member Is Offline
Mood: Staring at code
|
|
Here is a link to the purification and concentration of 5% ammonia on woelen's website, there are a lot of other interesting experiments on this
website as well(woelen is one of our moderators on Sciencemadness). http://woelen.homescience.net/science/chem/exps/concentratin...
[Edited on 4-15-2015 by gdflp]
|
|
|
MrHomeScientist
International Hazard
Posts: 1806
Registered: 24-10-2010
Location: Flerovium
Member Is Offline
Mood: No Mood
|
|
Don't buy into over-the-counter (OTC) products' claims of 'pure' or 'organic' - they are meaningless from a chemistry perspective. Always check the
labels. Household ammonia is, as blogfast said, a solution of NH<sub>3</sub> gas in water somewhere around 5% - 10%. If it really was pure
ammonia, the bottle would have to be at -33C or be a pressurized gas cylinder.
To determine concentration of your ammonia solution, titration or density measurement would be simplest. Since you already have acid of a known(ish)
concentration, you can add this to your ammonia of unknown concentration until the solution has neutral pH. Then you can calculate how strong the
ammonia is. You need pH paper or pH indicator for this - if I remember right, methyl red is the appropriate one for this situation. Try measuring
density too, by weighing a precise amount of the liquid and dividing the two numbers (g/mL). Then find a density vs. concentration chart online -
ammonia being an industrial chemical, there should be a bunch.
Woelen's site is an excellent resource. He's got a few pages on starting up a new lab that would also be helpful.
As for your original question, your ammonium chloride is just dissolved in all the water from your original solutions. Boil this off and you'll be
left with your white crystals. At least one recrystallization wouldn't be a bad idea, too.
|
|
|
Amos
International Hazard
Posts: 1406
Registered: 25-3-2014
Location: Yes
Member Is Offline
Mood: No
|
|
Before buying a product, it's a good idea to google search "(product name) MSDS"; MSDS sheets often describe at least relative proportions of the
chemical constituents.
|
|
|
Zyklon-A
International Hazard
Posts: 1547
Registered: 26-11-2013
Member Is Offline
Mood: Fluorine radical
|
|
Amonium chloride is not easy to make in solution.
But here's how I'd do it. a) find out how much ammonia (aq) is needed to bring the pH of 10ml of 30% hyrdrochloric acid to around 5-6.5.
b) Add, on a much larger scale, the reagents in the ratio discovered in step a.
c) Boil off 95% of the water or until crystals apear. It may have a high viscosity, which seems to be the trend for highly soluble compounds in
concentrated solution. Chill and filter.
I've purified this compound via sublimation before.
|
|
|
Zyklon-A
International Hazard
Posts: 1547
Registered: 26-11-2013
Member Is Offline
Mood: Fluorine radical
|
|
Or just buy some from me, link in sig.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
I recently made NH4Cl in good yield by reacting aqueous MgCl2 (a topical skin treatment and magnesium supplement) and ammonia in excess. Filter out
the Mg(OH)2, or decant the clear solution after cooling.
One can prepare impure MgCl2 also by mixing aqueous MgSO4 and NaCl, and freezing out the Na2SO4.
In place of MgCl2, one could also try CaCl2, ZnCl2, AlCl3, CuCl2...if available.
[Edited on 21-4-2015 by AJKOER]
|
|
|
chloric1
International Hazard
Posts: 1135
Registered: 8-10-2003
Location: GroupVII of the periodic table
Member Is Offline
Mood: Stoichiometrically Balanced
|
|
By mixing hot concentrated calcium chloride and ammonium sulfate, you get a thick gypsum slurry that you suction filter. You might have to filter
twice but you get a clear relatively concentrated ammonium chloride solution that needs minimal work up. Some calcium sulfate will be solublized by
the ammonium salt but no worries. Get that ammonium chloride saturated in respect to 80-90C and add ammonia soln to raise pH. This should solve
calcium sulfate dissolution.
Fellow molecular manipulator
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
