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Author: Subject: Isolation of Schweizer's Reagent
Vargouille
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[*] posted on 1-8-2012 at 12:46
Isolation of Schweizer's Reagent


Well, I'm back, along with more copper chemistry. I've heard multiple times, on this forum and elsewhere, that Schweizer's Reagent, [Cu(NH3)4(H2O)2](OH)2, only exists in aqueous conditions. I, wanting to ensure it for myself, tried to isolate a solid product of Schweizer's Reagent. My procedure follows, with the additional note that the ambient temperature is approximately 37C.

12.5 grams of CuSO4-5H20 are dissolved in approximately 200mL distilled water. In another vessel, 6 grams NaOH are dissolved in 30mL water. Once both solutions are clear, the NaOH is added to the CuSO4, resulting in a mass of dark blue Cu(OH)2. This is filtered and washed once with distilled water. About half of this is removed, and the remainder is placed into a clean beaker and NH3 is added, resulting in a dark blue-purple solution above the Cu(OH)2 which is surely the Schweizer's Reagent.

Here's where things began to go awry: The supernatant is drawn up in a pipette, and added to an excess of Crown Denatured Alcohol (a mixture of 65-75% MeOH, 20-30% EtOH, <10% IPA, and <10% methyl isobutyl ketone), to precipitate the Schweizer's reagent in the same way as an amount of tetraaminecopper (II) sulfate was isolated previously. This resulted in a white precipitate with a slight blue tinge. More of the supernatant is added, and the precipitate begins to go yellow, eventually becoming a tan color. An exotherm may have occured, but it is unclear due to high ambient temperature. Some white fumes were noted to be evolved. The solution is filtered, and the tan precipitate allowed to dry somewhat. It quickly becomes darker brown, and when added to an amount of water, it dissolved very slightly, creating a yellow solution.

My first thought was that it was being reduced by the MIBK, but this is unlikely unless the Schweizer's Reagent is a strong oxidizer, which it has not been noted to be from a cursory examination.

Any ideas on how to get the solid Schweizer's reagent? I could try again with hexanes, which are unlikely to react with the Schweizer's reagent, but that would require purchasing the hexanes.
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barley81
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[*] posted on 1-8-2012 at 12:51


In Brauer's amazing inorganic compilation of preparations, this is accomplished by evaporating the solution under a stream of dry ammonia gas. Take a look in the sciencemadness library.
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Hexavalent
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[*] posted on 1-8-2012 at 12:55


I have heard of people evaporating solutions slowly under a gentle stream of dry ammonia gas to get the crystals. Evaporation under air simply produces copper hydroxide (then oxide, of course) as the ammonia part of the complex comes off with the water. The crystals then must be stored either under a vacuum, under an inert gas or under ammonia gas itself.

Edit - Barley, you beat me to the post by a few seconds:)

[Edited on 1-8-2012 by Hexavalent]




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Vargouille
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[*] posted on 1-8-2012 at 12:58


Aye, dry ammonia would work to prevent the ammonia from leaving the precipitate, but I'm working in a shed, so my access to dry ammonia is somewhat diminished.
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[*] posted on 2-8-2012 at 05:14


Quote: Originally posted by Vargouille  
Aye, dry ammonia would work to prevent the ammonia from leaving the precipitate, but I'm working in a shed, so my access to dry ammonia is somewhat diminished.
Best I can tell, there's little other way of doing it. Consider two linked reactions: dissociation of the ammonia complex and evaporation of the ammonia. According to Le Chatelier's principle, escape of gaseous ammonia will increase the rate of dissociation of the complex. According to Raoult's law, you'll need (1) at least as much ammonia partial pressure as the vapor pressure of ammonia, and (2) less than the partial pressure of water. Item (1) prevents ammonia evaporation; item (2) allows water evaporation. So dry ammonia it is.

Perhaps the right question is how to build an ammonia generator and a drying train for it.
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barley81
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[*] posted on 2-8-2012 at 07:48


I can think of two ways:

1. Heating of concentrated ammonia solution. Heat the stock solution in a flask (don't boil) and lead the vapors through drying tubes filled with CaO. This will produce dry ammonia gas.

2. Liberation of ammonia from ammonium salts. Ammonium sulfate (cheap fertilizer grade, ammonium chloride/nitrate can probably be used) and sodium hydroxide are put in a plastic bottle with a tube attached to the cap. Smuv did this with great results:
http://www.sciencemadness.org/talk/viewthread.php?tid=16061
He initiated the reaction by adding some aqueous ammonia and shaking.
If you lead the gas through drying tubes with CaO, you will obtain dry ammonia gas.
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[*] posted on 2-8-2012 at 08:37


Heating strong ammonia solution is probably best for this purpose: 12 -15 M ammonia solution contains massive amounts of ammonia gas. By contrast 1 mol of, say ammonium sulphate only contains two mol of ammonia gas.

And no ammonia needs to be lost: just lead it through water at the end of your apparatus. It's so soluble in water that the last time I made some pure ammonia solution from fertiliser ammonium sulphate + NaOH the gas didn't even bubble through the recipient water: it just formed one small static bubble at the end of the outlet tube!




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Vargouille
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[*] posted on 2-8-2012 at 08:45


Agreed, dry ammonia would be the best way to go about it. To go about it another way, perhaps carrying out the reaction at low temperatures and crashing out in ice-cold hexanes would slow down the evaporation of ammonia enough that the solid Schweizer's reagent may be obtained, contaminated with Cu(OH)2 thought it would be. After that, it would be a matter of storing it in a cool place to diminish decomposition.
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watson.fawkes
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[*] posted on 2-8-2012 at 09:47


Quote: Originally posted by Vargouille  
To go about it another way, perhaps carrying out the reaction at low temperatures and crashing out in ice-cold hexanes would slow down the evaporation of ammonia enough that the solid Schweizer's reagent may be obtained, contaminated with Cu(OH)2 thought it would be.
Where would the water go? Is there any reason to believe that the water would spontaneously separate from the copper complex?

Back to the reaction dynamics. If there were a solvent in which water was soluble but ammonia was insoluble, you would be able to draw away the water by a solution phase change (as opposed to an evaporation phase change). But if that same solvent also dissolves ammonia, then it also drives decomposition of the complex just like evaporation would. On the other hand, if ammonia were unable to dissolve, there would be no phase change tug toward decomposition. Perhaps such a solvent exists with such differentiating solubilities, but given how polar both water and ammonia are, nothing comes to mind.
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[*] posted on 2-8-2012 at 10:32


I was calling to mind the precipitation of [Cu(NH3)4H2O]SO4 in ethanol. That complex is more stable as a solid than Schweizer's reagent, to be sure. The solid Schweizer's Reagent wouldn't be anhydrous, from what I understand. In retrospect, however, perhaps cold hexanes wouldn't be the best solvent to do the crashing out in. The point of using a cold solvent, perhaps ethanol, perhaps a longer-chain alcohol, should still stand.
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[*] posted on 2-8-2012 at 11:09


Quote: Originally posted by Vargouille  
In retrospect, however, perhaps cold hexanes wouldn't be the best solvent to do the crashing out in.

Perhaps you are not aware that alkanes are totally and utterly immiscible with water?

All the hexanes are alkanes.




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[*] posted on 2-8-2012 at 12:23


At the time, I was considering them to prevent the side-reactions with the Schweizer's Reagent that I noted in the OP. It cannot be so simple as loss of ammonia ligands, which would result in Cu(OH)2. Are there organic reactions that would explain the white and brown precipitates?
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[*] posted on 2-8-2012 at 13:38


As I understand the preparation, you can substitute NH4OH for NaOH, being a more preferable path (as support, see comments at: https://docs.google.com/viewer?a=v&q=cache:EyoCPBE_HxUJ:... ). This would also avoid Sodium ions contamination (but leave sulfate). Add more ammonia to dissolve (see video at http://www.youtube.com/watch?v=oLbnHEQsXNo ). Note, on loss of ammonia, Cu(OH)2 can precipitate out.

To store the deep blue needle-like crystals of tetraamminediaquacopper dihydroxide, [Cu(NH3)4(H2O)2](OH)2, an ammonia atmosphere has been recommended (see Wiki: http://en.wikipedia.org/wiki/Schweizer's_reagent ).

Also, I would use freshly boiled distilled water to dissolve the CuSO4 to remove any dissolved CO2, as well as limited atmospheric contact, as carbon dioxide appears to readily react to convert the dihydroxide to a carbonate. Source, see critique of an old preparation process combining Cu particles, aqueous ammonia and air at:

http://books.google.com/books?id=Tm4CAAAAYAAJ&pg=PA264&lpg=PA264&dq=schweizer's+reagent+preparation+old&source=bl&ots=qKaaJzyNZj&a mp;sig=CPlvCo0tmZ5EVqZFrUbzWIELeKc&hl=en&sa=X&ei=uQIbUOn5NNOn0AHtn4GQDA&ved=0CEEQ6AEwAA#v=onepage&q=schweizer's%20reagent%20prepar ation%20old&f=false

Do not add alcohol as to quote from one source: "cupro-ammonium hydroxide solution is decomposed by simple addition of alcohol to its ammoniacal solution, a blue substance essentially consisting of hydrated copper oxide being precipitated; the same result ensues on boiling, save that anhydrous black copper oxide is then formed, ammonia being driven off. In presence of a large excess of ammonia, the instability is less marked in all cases;"

Link: http://chestofbooks.com/crafts/mechanics/Workshop-Receipts-4...


[Edited on 3-8-2012 by AJKOER]
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[*] posted on 2-8-2012 at 15:43


Then perhaps acetone or MEK? I'm not sure of the solubility of the ammoniated copper salts in either, but that's something some testing can discover. If all else fails, I suppose I'll just bite the bullet and make the dry ammonia.
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[*] posted on 3-8-2012 at 12:41


Here is another original thought on how to dehydrate Tetraamminediaquacopper dihydroxide, [Cu(NH3)4(H2O)2](OH)2. Well, at least, I think it would be an interesting experiment, but not exactly sure if it can be made to work. First, we note per Wikipedia that Calcium nitrate can be prepared from an aqueous solution of ammonium nitrate and calcium hydroxide:

2 NH4NO3 + Ca(OH)2 → Ca(NO3)2 + 2 NH4OH

If we substitute say aqueous Copper ammonium nitrate, [Cu(NH3)4](NO3)2, for Ammonium nitrate, then possibly:

[Cu(NH3)4](NO3)2 + Ca(OH)2 + 6 H2O → Ca(NO3)2.4H2O + [Cu(NH3)4(H2O)2](OH)2

where I am speculating on the known dehydrating ability of Ca(NO3)2 to help us out. Note, added ammonia may also be necessary. The final step would involve either a mechanical or solvent separation of the salts (?).

Now a point of interest, there is actually nothing sacred about the tetraamminediaquacopper cation as depending on the ammonia concentration in solution, one could have anywhere from [Cu(NH3)(H2O)5]2+ to [Cu(NH3)5(H2O)]2+, with the latter occurring in very concentrated ammonia solutions. See first page at: https://docs.google.com/viewer?a=v&q=cache:IjHK0vuBZhcJ:...

FYI, here is an old reference (see bottom of page 5) on the preparation of Copper ammonium nitrate. To quote:

"35. For the preparation of tetramino cupric nitrate - Cu(NO3)2 .4 NH3, 10 grams of cupric nitrate.6H2 0 was dissolved in 8 cc. of water. To this was added 25 cc. of ammonium hydroxide (Sp.Gr. 0.9), the mixture well stirred and cooled to 5 C. A precipitate formed and the supernatent liquid was filtered on a Buchner funnel and the crystalline salt washed with alcohol, ether and air dried"

Link:
http://www.dtic.mil/dtic/tr/fulltext/u2/629884.pdf

[EDIT] As a final word, upon reading the above source please note that you should avoid the dry Copper (II) ammonium nitrate as it is a sensitive (impact) high explosive (this is not true of the hydrated form of this salt, see http://www.sciencemadness.org/talk/viewthread.php?tid=16220 ). The anhydrous form has been responsible for serious accidents (see, for example, comments at: http://www.epa.gov/oem/docs/chem/ammonitr.pdf and http://laws-lois.justice.gc.ca/eng/regulations/C.R.C.,_c._11...

[Edited on 4-8-2012 by AJKOER]
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[*] posted on 20-7-2013 at 18:36


I just can't resist sharing this, I think that it's quite appropriate to this topic:
<iframe sandbox width="560" height="315" src="//www.youtube.com/embed/jxoHB_sTkI8?rel=0" frameborder="0" allowfullscreen></iframe>

<a href="http://en.wikipedia.org/wiki/Schweizer's_reagent" target="_blank">Schweizer's reagent</a> <img src="../scipics/_wiki.png" />

[Edited on 7/21/13 by bfesser]




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