Sciencemadness Discussion Board - Seperation of NH4Cl and urea

Seperation of NH4Cl and urea

Sedit - 26-3-2009 at 11:36

I have a mixture of Urea and ammonia chloride that I would like to seperate. There is 45 grams total in this mix and I dont know the ratio of the two chemicals which complicates things a little more.

The compound I want to purify out is Urea and fractional crystalization seems like it would be the most straight forward method at first glance. If any one knows of a way to precipitate just the ammonia chloride it would be greatly appreciated because that seems like it would require less work then recrystalizing the solution.

What would be a better solvent to recrystalise this mix for seperation alcohol or water? IIRC urea is semi soluble in alcohol where as ammonia chloride is very sparingly so. I am in the process of trying to find solubility data on the two compounds in methanol but havent had to much luck just yet.
I have how ever been able to aquire the solubility data for H2O and given the wide margin of differnce in there solubility this would quite possibly be a good solvent for seperation.

Ammonia Chloride [ NH4Cl ] 29.7 g/100 water at 0 °C
Urea [(NH2)2CO ] 108.0 g/100 water at (20 °C)

Does anyone have any ideas on how to go about seperation of the two compounds?

Thanks for any help,

~Sedit

not_important - 26-3-2009 at 13:57

Crystallization from water might work, but you'd have to know the relative amounts to figure out how to do it.

Download a copy of Seidell's "solubilities of inorganic and organic compounds", get the 2nd edition. he Internet Archives and Google Books both have it.

In it you will find tables both the solubilities of both NH4Cl and CO(NH2)2 in water and several alcohols, and I believe in water+alcohol mixes.

As H4Cl is ionic it will have low solubility in less polar solvents, while urea is less affected. As I recall, hot isopropyl alcohol will dissolve a few tenths of a gram of NH4Cl per 100 ml, and close to 25 g of urea.

Finely grind the mixture first. Heat the alcohol, add some to the solids, and stir while maintaining temperature. Keep adding potions of alcohol until it seems little more of the solids dissolve. Stir briefly, filter quickly, allow to cool and then chill while not allowing water to condense into the mixture. Urea reacts slowly with alcohol to form urethanes, so prolonged heating is to be avoided.

A second recrystallisation from water may be desirable to really clean up the product. It should remove NH4Cl mechanically trapped in the urea crystals, as there won't be much of it in the hot alcohol and only a fraction of that caught in the urea, all of it will remain in the water solution.

Common 70% isopropyl alcohol might be OK, 90% would be better if you can't get the 99+% grade. You'll want to check the solubility tables first, and certainly do the 2nd crystallization.

Obviously this won't do much if there is only a few tenths of a gram of NH4Cl in the mix, but the water recrystallisation will take care of that. If there's only a small amount of urea, then enough NH4Cl can dissolve so as to still be a bother. In that case doing a second dissolution and crystallisation should do the trick, as the ratio of urea to NH4Cl would be much higher and you've a better idea on the amount of solvent to use.

[Edited on 26-3-2009 by not_important]

Sedit - 26-3-2009 at 14:58

Thanks for your responce N_I thats pretty much the same conclusion I have come to since posting this. Where as I dont have anything but 70% Isopropyl alcohol I am going to attempt to use dry methanol in the manner you discribed of slowly adding the powdered mix to hot Alcohol. After this I was going to filter it and cool it to precipitate any remaining NH4Cl. A trace amount of ammonium chloride shouldnt hurt me to much but I want to remove it as best I can.

~Sedit

[edit]
Ok given the data you provided for solubility in Isopropyl I just went out and ground up 25 grams and mixed this with 100 ml of methyl alcohol, everything went into solution. So am I to assume that this is relativly pure urea or that the more polar MeOH has a much wider range of solubility for NH4Cl?

I got it from an instant cold pack because I use them to get ammonia nitrate for sodium nitrate/ammonia synthesis but this new kind said it containted Urea and ammonium chloride. Im going to keep going with adding more till I can get an idea of how much dissolves in MeOH.

[Edited on 26-3-2009 by Sedit]

not_important - 26-3-2009 at 16:14

see chart from aforementioned book. 100% MeOH = 2.75 g NH4Cl / 100 g MeOH, EtOH or nPrOH = 0.18 (the propyl column is mislabed as EtOH+PrOH when it's MeOH+PrOH.

Even a small amount of water results in a noticeable increase in the solubilities, especially in MeOH.

Sedit - 26-3-2009 at 16:50

In total I got 28grams of the mixture to dissolve in an unknown amount of MeOH(By the time I went back to check some of the MeOH from the original 100ml had evaporated). On cooling the entire solution was taken over by large needle crystals that I have yet to weigh. I am pretty confident that the majority of this mix is urea.

As you mentioned even a small amount of water will change the ratio of solubility so there is no reason for me to continue trying to purify the rest because even though my MeOH was freshly opened dry MeOH the humidity right now is 100% and I would rather wait for a better time to finish the other 100 grams of mix(dont ask me why I put 45 grams total in the first post).

Thanks for the chart N_I im sure that will come in handy when I do the rest and I think that pure Iso is the way to proceed here.

~Sedit

PHILOU Zrealone - 27-3-2009 at 01:15

Maybe another possibility would be to play on the solubility difference between ammonium nitrate vs urea nitrate by adding some HNO3 to the mix... but you would end up with NH2-CO-NH2.HNO3

or

to play on the volatility difference between NH3 and NH2-CO-NH2 by addition of stoechiometric amount of Ca(OH)2 and gentle warming...
2 NH4Cl + Ca(OH)2 + NH2-CO-NH2 --> 2 NH3(g) + 2H2O(l) + CaCl2(s) + NH2-CO-NH2

not_important - 27-3-2009 at 02:31

While adding Ca(OH)2 will let you drive off the ammonia, it leaves you with CaCl2 which is fairly soluble in alcohol; in this case fractional crystallisation from water might be the best choice . NaOH would result in NaCl being formed, it has only limited solubility in alcohols and so may be a better choice than Ca(OH)2.

Sedit - 27-3-2009 at 07:20

Would the NaOH react with urea also?

Iv been thinking last night and was wondering if anyone know of a quantitive test for analysis that would be selective for ammonium chloride? If I could figure out the exact ratio the seperation could become easyer and more complete.
While I am pretty confident that the vast majority is Urea I would like some kind of confirmation thru test to prove to my self that its not just excess moister in the air causing all the excess material to go into solution.

~Sedit

not_important - 27-3-2009 at 15:11

The NaOH will react much more quickly with the ionic NH4Cl than the di-amide urea, so if just enough NaOH is added to react with the NH4Cl in the mix the urea will be basically unaffected.

Try testing for chloride - add a bit of a soluble lead or silver salt in solution to a few drops of the product in fairly concentrated solution. Even checking the conductivity/salinity should do the job, urea is non-ionic and has little effect on the conductivity of water while ammonium chloride will significantly increase it.

Sedit - 27-3-2009 at 19:06

Thanks N_I those are some thoughts although I dont have any lead salts I have been considering making some so atlest now I have a reason. What is to be expected with the addition of the Pb salts to the solution precipitation I assume? This sounds very useful on determining the concentration by weight of any precipitate that may form.

I have something else to note also. As I stated above the MeOH was allowed to cool and large needle crystals filled the entire container that it rested in all the way to the top of the solution. This is all fine and well and what is expected to happen with a saturated solution so I broke the crystals up and filtered the remaining MeOH off. After sitting in another container for some time now there are small crystals unlike the first in the MeOH. These crystals look more like Sodium chloride then anything else yet the MSDS stated the product contains NH4Cl and Urea which I thought both formed needle crystals. Am I mistaken on this or is there some other material mixed with this they are not telling me?

not_important - 28-3-2009 at 17:25

Crystal forms often differ with the solvent used. Finding this information for inorganics can be difficult, it's not commonly given in the CRC/Lange's type references.

Also compounds may take different forms when crystallised in the presence of appreciable amounts of other substances, or even quite small amounts of the other substance in some cases. While traces of the 2nd compound may be included in the crystals, this is not a case of mixed crystals such as cobalt ammonium sulfate, just a modification of form. For most of these a recrystallisation will result in the normal form, as the 2nd substance's concentration will have been greatly reduced.

Temperature at the time of crystallisation may also affect the crystal form.

In you case this may answer your question

Quote:

On the Crystal Structure of Ammonium Chloride; by Ralph W. G. Wyckoff

Clear optically isotropic crystals in the form of rectangular prisms several millimeters on one side and usually from one to two millimeters thick were obtained by desiccating a solution of ammonium chloride containing urea.

The American Journal of Science Vol IV No. 19 July 1922

PbCl2 would precipitate, however it does have a slight solubility - around 10 g/liter at 20 C; still it would do for a quick check to set a limit on chloride concentration. Conductivity measurement might be better, although you'd need to create calibration data using known pure NH4Cl or find a table of that data.

Sedit - 28-3-2009 at 20:44

Yes I understand completely what you mean N_I. Just recently when I was looking into ammonia nitrate the text went into the five, IIRC, different molecular forms that it took on at various temperatures. It would only seem logical that urea or ammonia chloride would behave in a simular fashion as this.
Odds are these are the same as the needle crystals that formed as the temperature slowly droped in the initial recrystalization but these are forming at a slow and steady pace at a lower temperature.

I think I have experienced the same phenomenon while trying to recrystallize oxalic acid and at first I was convinced that the flatter plate like crystals where another substance but the same results kept appearing time after time where it would begin as fine needles that would get covered in flat plates of sorts after it had been sitting for a while.

~Sedit