I'm in need of a small amount of CuCl, and am unsure of what would be the best method of producing a product free of Cu(II). Right now, the best two
options are:
1 - Redox reaction between CuCl2 and Cu metal - Using stoichiometric amounts of both should give pure CuCl, but I have found no data regarding whether
or not the reaction goes to completion.
2 - Electrolytic reaction of Cu metal to Cu2O, followed by reaction with HCl to form CuCl...Not sure how well this would work, though.
So, anyone care to chime in on which reaction would yeild a purer product? Anyone got a better method that doesn't require any exotic reagents?garage chemist - 22-6-2007 at 14:53
Boil a solution of copper(II)sulfate in nearly saturated NaCl solution with an excess of metallic copper until it is completely colorless (may take
many hours, depending on the surface area of the copper! Works best with powdered copper. Also, protect the solution at least somewhat against aerial
oxygen, e.g. by refluxing in a round bottom flask with a loosely stoppered dimroth condenser). The forming CuCl is held in solution as the unstable
complex ion CuCl2(-) due to the high concentration of chloride ions.
Upon lowering the chloride concentration by pouring the solution into a large amount of water, the complex decomposes and CuCl precipitates as a white
powder.
The water used for this operation must be made air-free by previous boiling and should contain dissolved sulfur dioxide as an additional antioxidant.
An alternative method for CuCl preparation is to reduce a copper(II)sulfate solution with added NaCl with sodium metabisulfite or sulfite. Keep adding
reducing agent until it is colorless. Then precipitate the CuCl like above.
CuCl must be stored in completely airtight vials and be protected from light. It is very prone to aerial oxidation, it becomes green when that
happens.guy - 22-6-2007 at 16:01
Sufite is the fastest. You can find solid sodium sulfite in Iron Out which is cheaply available in Home Depot as rust remover.garage chemist - 22-6-2007 at 16:38
Though dont use sodium dithionite or products containing it, it is too strong a reducing agent.not_important - 22-6-2007 at 16:46
An old standard method was to make a solution of CuCl2 or CuSO4 and fairly strong HCl, then store this over bare copper wire or sheet (usually scraps
of either). As already said, protect from air access.
Sometimes they'd use copper and copper oxide, and just pour strong HCl onto it, as in the days when labs made a lot of their own equipment they end up
with copper scale.
Anyway, the solution was left to sit until it turned a pale yellow-green, this might take several weeks. If all the copper dissolved first, more was
added. The solution was then decanted from the remaining copper through a filter into ice cold, previously boiled water, sometimes specified as
being saturated with SO2 or with a bit of sodium bisulfite dissolved in it. This cause the CuCl, which has a low solubility, to precipitate out. The
CuCl would be collected on a filter and washed with more cold, boiled water - again containing SO2 or bisulfite. Care must be taken to keep the CuCl
covered with water and not draw air through it. If it needed to be dry it would then be washed with alcohol and then ether, acetone works well in
place of ether, and quickly transfered to a storage bottle. The dry CuCl is much less reactive to O2 than when moist, and the vapours of ether or
acetone would drive the air out of the bottle.
The copper and remaining HCl+copper salt solution would have fresh HCl+copper salt added, to start the next batch going.
Slow, but works well when you're not needing it right away. If you use CuCl2 as the starting salt and SO2 instead of a bisulfite, you're only dealing
with Cu, Cl, and a little SO4 ions, so decent purity of CuCl is easy to obtain.guy - 22-6-2007 at 17:15
Quote:
Originally posted by garage chemist
Though dont use sodium dithionite or products containing it, it is too strong a reducing agent.
Just use less of it. I will produce Cu(0) but that will quickly react with CuCl2 to form CuClIntergalactic_Captain - 22-6-2007 at 17:22
Hmm...The CuSO4/NaCl/Cu(s) method GC gave seems to be the most economical at the moment. For copper, I'll most likely use copper shavings from a
filed down piece of old pipe...I think I'll get started on that now, actually.
Any idea if this would work with NaBr as well? I believe someone made reference to this in another thread, but it was an off-topic post and I have no
idea which one it was in.
Am I correct to assume that the CuCl precipitate will be reasonably pure? As long as all of the Cu(s) has reacted, CuCl should be the only insoluble
material, correct?The_Davster - 22-6-2007 at 17:29
I don't know if you also need the iodide, but making the iodide is very simple
Cu2+ + 3I- -->I2 +CuI12AX7 - 22-6-2007 at 18:34
I discovered a good deposit of it when plating copper in an unusually high chloride concentration. Preusmably, the mechanism is chloride ions
attracted to the anode, where copper is oxidized progressively, from metal to Cu(I) to Cu(II). The intermediate, being in a chloride solution too
weak to form the complex and dissolve, remains as a slippery white coating on the anode.
The same happens in a solution of Cu(II), it would appear.
And of course, reacting copper and chloride will go through the green and brown intermediates, giving the soluble complex which can be hydrolysed. It
will be of good purity, but moist and of fine size. You'll have to dry it out quickly (or in a glove box) to keep it from oxidizing.
I seem to recall Brauer gives a pyrolytic prep of CuCl2 + Cu. Probably a tube furnace and vacuum or inert gas filled vycor involved.
Timgarage chemist - 22-6-2007 at 21:48
You should also be able to find a detailed CuCl preparation on Orgsyn, as this substance finds use in organic chemistry (Sandmeyer reaction of
aromatic diazonium compounds to haloarenes). I think it was o-chlorotoluene that had the preparation of CuCl included in the synthesis.
