chemrox
International Hazard
Posts: 2961
Registered: 18-1-2007
Location: UTM
Member Is Offline
Mood: LaGrangian
|
|
facile CuBr prep
I found the following:
http://lanthanumkchemistry.over-blog.com/article-how-to-make...
Copper(I) bromide: React copper(II) sulfate with an equal amount of sodium bromide. Add ascorbic acid or sodium metabisulfite. White copper(I) bromide
will precipitate.
What do you think he means by"equal amount?" equal weight or mole/mole? I have the materials to make this (in excess) but before I invest any
reagents I'd like to hear from anyone who has done this. btw- I searched before posting..
Thanks,
crx
"When you let the dumbasses vote you end up with populism followed by autocracy and getting back is a bitch." Plato (sort of)
|
|
|
woelen
Super Administrator
Posts: 7988
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
One mole of CuSO4 uses up one mole of NaBr, the Cu and the Br combine 1 : 1. Use a slight excess of NaBr, to be sure that enough bromide is present
and all copper precipitates as bromide and not as something else, making your product impure. Do not use a large excess, because that could keep
copper in solution as CuBr2(-), especially when you work with concentrated solutions.
The S2O5(2-) ion in theory is capable of reducing 4 copper ions. Just use a fairly strong excess amount of this, to assure that the reaction
completes. I would take 1 mole of copper sulfate for half a mole of S2O5(2-), which is a two-times as much amount as theoretical. Bisulfate is dirt
cheap, so do not save on that too much.
I do not know how many electrons ascorbic acid produces per molecule when it acts as reductor, so first try to find that and then use a slight excess
of ascorbic acid, based on that. Only use a slight excess amount, because ascorbic acid is good at making complexes with copper and these may keep
part of the copper in solution.
I would prefer the bisulfite for this reaction. It is much cheaper and you can more safely use a large excess amount of this to drive the reaction to
completion quickly.
|
|
|
chemrox
International Hazard
Posts: 2961
Registered: 18-1-2007
Location: UTM
Member Is Offline
Mood: LaGrangian
|
|
Thanks so much for clearing this up and providing the detail on how to do this. I have bisulfite, NaBr and CuSO4 so it seems better than ordering
something that runs $100/1/4# when I need it as a catalyst.
Sooo.. that appeared to have worked rather well. The white ppt is the CuBr and the aqua color is the hydrated Cu++ ions ?
[Edited on 11-5-2012 by chemrox]
and I like this too:
The submitter used commercial cuprous bromide. The checkers prepared cuprous bromide by dissolving 600 g. (2.4 moles) of commercial copper sulfate
crystals and 350 g. (3.4 moles) of sodium bromide in 2 l. of warm water; the solution was stirred while 151 g. (1.2 moles) of solid sodium sulfite was
added over a period of 10 minutes. Occasionally a little more sodium sulfite was required to discharge the blue color. The mixture was cooled, and the
solid collected on an 8-in. Büchner funnel, washed once with water, pressed nearly dry, and then dried in the air overnight. The yield of cuprous
bromide was 320 g. (93%).
http://www.orgsyn.org/orgsyn/orgsyn/prepContent.asp?prep=cv3...
[Edited on 11-5-2012 by chemrox]
"When you let the dumbasses vote you end up with populism followed by autocracy and getting back is a bitch." Plato (sort of)
|
|
|
DJF90
International Hazard
Posts: 2266
Registered: 15-12-2007
Location: At the bench
Member Is Offline
Mood: No Mood
|
|
The issue is not with making it, but getting it out of the reaction in a dry and pure state. Whilst wet it oxidises very quickly (at least the
chloride does, and I'm certain the bromide does too), so filtering it out and getting it dry without it going green is likely to be a big challenge.
See Inorganic syntheses vol. 2 for the prep of CuCl.
|
|
|
chemrox
International Hazard
Posts: 2961
Registered: 18-1-2007
Location: UTM
Member Is Offline
Mood: LaGrangian
|
|
Thanks DF, check above the edit I made adding an orgsyn note. Seems like the br is less sensitive than the Cl
That said, I'd like to start with some cleaner CuSO4.6H2O. Mine has visible crap in it. I think they kind of exaggerated the definition of
technical..
[Edited on 11-5-2012 by chemrox]
"When you let the dumbasses vote you end up with populism followed by autocracy and getting back is a bitch." Plato (sort of)
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Quote: Originally posted by woelen  |
The S2O5(2-) ion in theory is capable of reducing 4 copper ions. Just use a fairly strong excess amount of this, to assure that the reaction
completes. I would take 1 mole of copper sulfate for half a mole of S2O5(2-), which is a two-times as much amount as theoretical. Bisulfate is dirt
cheap, so do not save on that too much.
|
I recently tested that hypothesis when I was recovering some gold from a mobile phone (sodium metabisulphite is the reducing agent used to precipitate
the gold from solution). Since as my solution also contained copper, and concerned the copper would co-precipitate, I added quite a bit of Na2S2O5 to
a separately prepared copper solution: no reaction or precipitation ensued.
I then added metabisuphite to my copper/gold solution (somewhat neutralised first) and obtained a precipitate of gold powder. On re-dissolving this
(tiny bit of) gold in (a tiny bit of) Aqua Regia it was clear it contained no copper whatsoever (solution was slightly yellowish, not a hint of blue
or green).
It seems metabisulphite isn’t capable of reducing Cu (II) to Cu (0).
[Edited on 12-5-2012 by blogfast25]
|
|
|
woelen
Super Administrator
Posts: 7988
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
My excuse for not being very clear in my previous post. I meant reduction of copper(II) to copper(I), not reduction to metallic copper.
Indeed, metabisulfite (or just sulfite if pH is higher) is not capable to reduce copper(II) to copper(0), but it can reduce to copper(I). Just a
simple experiment:
- dissolve some copper sulfate in water
- add some sodium chloride
- add some solution of sodium sulfite or metabisulfite
The color changes from greenish/blue to pale yellow. All copper is reduced to copper(I) and remains in solution as a coordination complex.
You can precipitate the copper(I) by acidifying somewhat and pouring the solution in water. It precipitates as CuCl. I expect this works with NaBr as
well as with NaCl and then you get CuBr.
Isolating CuCl (or CuBr) is very hard. As soon as it comes in contact with air it turns green due to partial oxidation and the material becomes
contaminated with a basic copper(II) chloride.
[Edited on 12-5-12 by woelen]
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Woelen:
Ah yes, it’s all coming back to me, I made some CuCl that way years ago. I guess those who recover gold from ewaste solutions in Aqua Regia are
lucky that when they add the sodium bisulphite the copper is reduced and complexed to Cu (I) and doesn't precipitate as CuCl...
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
