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Author: Subject: NaCl (s) + NH3 (aq.) = ????
chemkid
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[*] posted on 13-2-2008 at 16:18
NaCl (s) + NH3 (aq.) = ????


I tried this out today and got a white precipitant. It was very fine and dissolved in additional water. I used distilled water, kosher salt, and household ammonia.

Could this perhaps be sodium hydroxide? I know ammonia with a aluminum ions yields aluminum hydroxide, but it doesn't make sense that the NaOH would precipitate considering it is 111g/100mL of water soluble.

Sodium amide was the other answer i came up with, but after looking up some information that didn't sound right at all.

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[*] posted on 13-2-2008 at 16:31


Aluminum hydroxide is less of a base (amphoteric IIRC) than the so called ammonium hydroxide.

I doubt you would be forming sodium hydroxide, let alone sodium amide!

Probably some sort of low solubility impurity in the salt.

[Edited on 13-2-2008 by Fleaker]




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chemkid
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[*] posted on 13-2-2008 at 16:38


The sodium chloride dissolves cleanly in water. Perhaps the pH change precipitates it from solution.

[Edited on 13-2-2008 by chemkid]




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[*] posted on 13-2-2008 at 16:47


I would more suspect some sort of insoluble stuff, possibly that's insoluble in base, or some sort of gas which forms a fine dispersion.

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[*] posted on 13-2-2008 at 19:40


How much salt did you add? Is it possible you just added salt past the saturation point? Other than that a salt impurity is most likely.

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[*] posted on 13-2-2008 at 23:44


I think it is plain salt which you precipitate. It might well be that solubility of salt in ammonia-water is much less than solubility in pure water. A similar effect is created by adding alcohol to a solution of table salt, or adding conc. hydrochloric acid.



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[*] posted on 14-2-2008 at 02:03


do you know if this salt has Magnesium carbonate (free-flow agent) in it?

I`m wondering if it might be Mg(OH)2 you`r seeing?




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chemkid
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[*] posted on 14-2-2008 at 12:41


I used very little salt no were near the point of saturation, at least in water, for NaCl. The salt probably does not have magnesium carbonate in it for it is kosher and has sodium ferrocyanide (yes from that sodium ferrocyanide otc thread) listed as the free flowing agent. Insolubility in ammonia or base sounds right, because it redissolves in additional water.

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[*] posted on 15-2-2008 at 06:55


I think either of them not pure, u have no ppt.ed salt or complex to be formed;)
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[*] posted on 15-2-2008 at 08:24


Now you can show that you are a chemist and not just a dabbler.

First you would have looked up the solubilities of NaCl and NaOH and seen that NaOH is much more soluble than NaCl; commercially 50% solutions of NaOH are sold in drums and tanker cars. There are several books of solubilities online, but I suspect you will not find listing for NaCl in aqueous ammonia, but you might have luck.

Then you would have looked up the Kb of each, and seen that NaOH is a much stronger base than "NH4OH"., a reaction would be more likely to go from a salt of NH4+ and NaOH to NH4OH and a salt of Na+ than the other way round unless there was some other driving force and you'd just checked that any such driving force is not solubility.

The repeat the experiment on a larer scale, and allow the solid to fully settle out. Carefully decant off the liquid leaving the solid with as little liquid on it as possible. Sometimes you can remove a little more water by carefully touching the edge of a paper towel to the clear liquid covering the solid.

Compare the volume of the solid to the amount of salt used. Add a little bit of distilled/deionised water (you did use that to dissolve the salt in the first place, right?) and see what happens Carefulkly warm, say to about 50 C, the mixture and see what happens. The add more water if needed to get solution at 50 C, setting some limit on how much water you do add - not more than you started with to dissolve the NaCl.

If everything dissolves, put the solution in a wide mouthed container and cover it with a paper towel held on with string or a rubber band. Set it aside to slowly evaporate, watching to see crystals forming. Do the same with the original clear liquid from the salt + ammonia mix, also watching the crystals forming.

If all of the solid failed to dissolve, save it by decantation once again. The start figuring out how you might determine what it is. Check solubilities of hydroxide to see what might fit. If you just used tap water to make the solution, look at Mg and Ca for example.

When those solutions have evaporated enough that crystals have formed, pour off and collect the liquid when it is maybe equal in volume to the crystals. Put in new containers, cover with paper towels again, and let evaporate to dryness. Look at all four sets of crystals, compare them - how much of each is there, do they look the same or different. Think on how you might test them. Do they have any ammonium salts in them?

Part of being a chemist is figuring things out. You should have checked solubilities before asking someone else, as that's a readily available resource. After that it's acceptable to ask "so, what is this stuff?" But expect that you may have to do some work to find out, if no one can come up with an answer quickly.
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[*] posted on 15-2-2008 at 12:39


Thank you not_important for that short course on solubility's etc. I don't have time to thoroughly read/try out now but i will soon! Thank you,

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[*] posted on 16-2-2008 at 17:21


Quite honestly solubility curves seem more of a luxury to me than a basic tool. I have unsuccessfully looked for them various times and til this day I am yet to find a free access site with a wide variety of chemicals in various media. Perhaps if you don't mind, would you list those sources as it seem you know quite a few of them? It'd be very useful.
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microcosmicus
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[*] posted on 16-2-2008 at 18:05


Here are some free online references which have lots of solubility data:

Solubilities of Inorganic and Organic Substances
A. Seidell 1907

http://books.google.com/books?id=7Y8AAAAAMAAJ&printsec=f...

A Dictionary of Chemical Solubilities: Inorganic
A. M. Comey, D. A. Hahn 1921
http://books.google.com/books?id=TrdCAAAAIAAJ&pg=PA467&a...

IUPAC-NIST Solubility Database
http://srdata.nist.gov/solubility/sol_main_search.asp

[Edited on 16-2-2008 by microcosmicus]
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[*] posted on 23-6-2011 at 06:49


Here is another route. Assume the presence of dissolved CO2 in the water or in the air.

Then:

NaCl (s) +NH4OH (aq) + H2CO3 (aq) --> NH4Cl (aq) + NaHCO3 (s) + H2O (l)

This is a quite famous (and commercially significant) reaction, our let the author of this tread do the research.
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[*] posted on 23-6-2011 at 10:29


I think you are referring to http://en.wikipedia.org/wiki/Solvay_process

This is unlikely, as the sodium bicarbonate only precipitates out due to very concentrated solution, I would think.
But this most important part is the fact that the concentration of CO2 will be very low in his solution. According to wikipedia, CO2 has a solubility of 1.45 g/L at STP, or about a .o329 M solution
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