Charles Boyle
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CuSO4 from CuCl2 + NH42SO4
I believe I have created a 5 mL solution of both Cu(II)SO4 and NH4Cl.
I took 0.3 grams of reagent grade NH42SO4 and mixed well with 0.3 grams of CuCl2 (created with an excess of HCl and copper wire which was oxidized
with some 3% H2O2)
I then poured the reactants in a beaker and added enough water to completely dissolve the reactants. (approximately 4 mLs)
I am now evaporating the solution, which is a darker shade of light blue than regular CuCl2 solution.
My first question is: Have I created what I think I have?
If not, would heat be required to break the ionic bonds of the initial reactants?
If I have, how would I go about separating the NH4Cl from the solution, as I'm just interested in the CuSO4?
I know the solubility of CuSO4 is 36.1g/100mL at 0 degrees C, and the solubility of NH4Cl is 29.7g/100mL at 0 degrees C.
Could I lower the temperature so that the NH4Cl is no longer soluble, or is there an easier way?
Thank you for your time.
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bbartlog
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Those are all rather soluble... off the top of my head I'm not sure what would crystallize first. Anyway, you haven't created any particular salt:
you've created a solution of ions, copper and ammonium and sulfate and so on. What comes out of the solution will depend on what you do.
It's possible that heating to dryness and then heating further to 350C or so for a while would drive off the ammonium chloride, leaving just CuSO4
(anhydrous). But I doubt the result would be all that pure. If you have other solvents, you could try precipitating out the copper sulfate
selectively... adding 5ml of methanol or ethanol should precipitate most of the CuSO4 while leaving NH4Cl in solution. Given the volumes involved, you
could add it dropwise with swirling and stop once the blue color was almost gone from solution.
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Charles Boyle
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Would I be correct in assuming the concentration of ethanol doesn't matter?
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symboom
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now if you wanted the pure copper sulfate could you just boil off all the ammonium chloride shifting it to the right.
i know ammonium chloride is very soluble but it does decompose at 338 C
copper sulfate at decomposes to SO3 and CuO at 650 C
copper sulfate is insouble in ethanol but ammonium chloride is a little souble.
but i would go for boiling it of to decomposition forming HCl and ammonia gas will be given off. if you want the ammonium chloride.
condense the gasses into ice water then you will have a solution of close to pure ammonium chloride with may be depending on pH. slight contamination
of ammonium hydroxide or HCl acid
[Edited on 17-12-2011 by symboom]
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bbartlog
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Well, the final concentration and volume of your solution will certainly matter. Neither CuSO4 nor NH4Cl are entirely insoluble in ethanol, let alone
50% ethanol. I had suggested 5ml of ethanol (100% or 95%), which would give you a final volume of 10ml of 50% ethanol. According to Atherton Seidell,
the solubility of CuSO4 in 40% ethanol is 0.25g per 100g solvent, so of your roughly 0.4g of CuSO4 only about 0.02g should stay in solution. Meanwhile
the NH4Cl should be soluble to the tune of 13g per 100g solvent in 50% alcohol, so about 1.3g in the resulting 10ml of solution, which gives you lots
of room for error.
Adding ethanol that isn't 95-100% will change things around some (mainly, it will reduce the percentage of CuSO4 you can hope to precipitate, from 95%
to something lower).
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Charles Boyle
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Okay.
So I heated the solution until there was just enough water to barely cover the bottom of the beaker.
I let it sit and evaporate, the crystals created greatly resembled CuCl2 crystals, as in they were long and thin and had a green colour.
I added some 80% ethanol and the crystals eventually dissolved and the resulting solution was again, a light blue colour.
After heating yet again, I shook the solution and an interesting thing happened.
After agitation, microscopic crystals settled out at the bottom.
Right now they are drying off.
They seem light blue in color, but they could be white and tinted by the light blue solution still left over.
I wish I could take some pictures, however I lost the USB cord for my camera.
They may be crystals of NH4Cl and the CuSO4 could still be in solution...
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kavu
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A formation of a double salt might also be a possibility. Quite general lab experiment is to prepare ammonium copper(II) sulphate,
[(NH4)2SO4·CuSO4·n H2O], from aqueous solutions of copper sulphate and ammonium sulphate.
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symboom
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Quote: Originally posted by kavu  | | A formation of a double salt might also be a possibility. Quite general lab experiment is to prepare ammonium copper(II) sulphate,
[(NH4)2SO4·CuSO4·n H2O], from aqueous solutions of copper sulphate and ammonium sulphate. |
oh forgot about copper's abillity to form the double salt
and it being stable futher heating should give off the ammonia from the complex without decomposing the copper sulfate and keeping below 600C
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Charles Boyle
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So, after three days of evaporation I have left a small amount of green solution left, the previously mentioned microscopic crystals are now a
greenish colour, and there are also some dark blue crystals on top of and in between the "bed" of green crystals.
My thoughts now are that I have indeed created a very small amount of CuSO4 and a larger, albeit still small, amount of NH4Cl.
If the balanced equation of NH42SO4 + CuCl2 --> 2NH4Cl + CuSO4 is correct, I should have twice as much NH4Cl than CuSO4, which appears to be the
case.
I will let it keep evaporating and hopefully the green tinge disappears from what are possibly the NH4Cl crystals.
EDIT:
After discovering my camera's USB cord, I will soon bring you some crap pictures (taken before I discovered my camera's macro option) and one decent
picture.
[Edited on 21-12-2011 by Charles Boyle]
[Edited on 21-12-2011 by Charles Boyle]
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Charles Boyle
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Since I cannot edit previous posts, I will have to add the images in another post.
Here is the product after the addition of water and heating, notice how the structure of the crystals resembles CuCl2.
Here is the light blue solution after the addition of ethanol and agitation.
And here is the latest picture, after 3 days of evaporation, the, what I assume are, dark blue CuSO4 crystals sitting on the "bed" of NH4Cl which is a
green tinge due to solution that has not evaporated yet.
[Edited on 22-12-2011 by Charles Boyle]
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Charles Boyle
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Happy Holidays ScienceMadness!
I come bearing visual gifts 
Here are some pretty Chanukah colored pictures of some CuSO4 crystals on top of some snow-white NH4Cl crystals.
<a href="http://www.mediafire.com/imageview.php?thumb=4&quickkey=iy7tiu241cwmu6h" target="_blank"><img
src="http://www.mediafire.com/imgbnc.php/a512de014dca841f029bf953ad6916e38d7f5531cef13b02556d15bf293839364g.jpg" border="0" alt="Unlimited Free Image
and File Hosting at MediaFire" /></a>
The plan of action now is to physically separate the CuSO4 crystals from the NH4Cl (as I don't really have access to a suitable solvent) and then
dissolve them and let the solution recrystallize out into one larger crystal.
As for the NH4Cl, I'm just going to store it until I find a project for it.
My only questions coming out of this experiment are:
Why didn't the CuCl2 react with the NH42SO4 when in a plain water solution, and why did it take the introduction of ethanol to create a reaction?
Could it have been impurities in regular tap water that prevented the reaction?
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