1: Its a old 250g container where my CuCl is in and its had become a solid block (Seriously , I can hardly get anything out). How do I safely break
the block without risking bottle breakeage or fine dust in the air.
2: I used some in a testube and I cant wash the test tube anymore. I tried hot water and a green residu remains. I tried acidified (acetic acid)
hydrogen peroxide to get the CuCl to copper in the 2+ oxidation state. A brown residu remains (probably anhydrous CuCl2. This solid does not dissolve.
I tried complexing it with ammonia. Didnt work as well . How the hell do I clean it?
joris-jeffB - 17-1-2008 at 06:16
It's been ages since I tried anything with cuprous compounds, but my go-to reagent for attacking anything copper-related was always nitric acid.
Outdoors, if I wasn't being absent-minded.SecretSquirrel - 17-1-2008 at 06:54
Regarding your first question: I always attack clumped chemicals with a screwdriver. This method proved to be the most successful. Just be patient in
order not to damage the container if it is made of glass.Jor - 17-1-2008 at 08:10
yes but a screwdriver is made of steel. So it contains Iron wich would mean iron would react with the cuprous chloride forming iron(II)cloride and
iron(III)chloride contaminets in the container. I already tried plastic but the stuff is too hard too break with plastic.not_important - 17-1-2008 at 08:51
Try hot 10% hydrochloric acid, let the acid filled tube sit in a waterbath for some time.
A good hardware store will carry copper and/or bronze nails in fairly large sizes. Use one to make holes in the block of CuCl, hammering it in as
needed. Get several holes and the block should split.
The green colour comes from Cu(II), likely a basic chloride similar to CuCl2.3Cu(OH)2, simple CuCl2 would dissolve in water.microcosmicus - 17-1-2008 at 09:07
CuCl2 is quite soluble in water, so that can't be your insoluble residue. Anhydrous CuCl2
will readily absorb moisture, even from air, to become green and hydrous. Maybe what
you have is a copper oxide, hydroxide, some Cu+ compound, or even precipitated metal .
Try using an acid. (Whatever acid you have handy and consider expendable for cleaning.)
If your acid is not oxidizing and doesn't do the trick alone, add some of
your H2O2 to the acid.
As for the stuff caked in the bottle, maybe find or improvise some sort of rotary rasp
out of copper, glass, ceramic, or some other substance which shouldn't contaminate
your chemical, lower that into the bottle and gring the stuff back to powder right in the bottle.woelen - 17-1-2008 at 10:43
What color has the CuCl? CuCl is white, white like fresh snow. On storage, CuCl quickly becomes green, due to oxidation. I also have some CuCl, but
now it is a green solid, which contains quite some Cu(Cl/½O), some basic copper chloride, or copper oxychloride.
This material is very insoluble, but in warm conc. HCl it certainly dissolves. Consider your old CuCl not as CuCl anymore, but as some oxychloride of
Cu(I) and Cu(II) at the same time. Still, this can be interesting, it can be used as a nice source of copper-solutions in conc. HCl.
Breaking down the big lump of your chemical can be done by buying a piece of copper tubing (used for water pipes and plumbing) and strongly hammering
on the tube. You most likely will break the container, so keep another container prepared for this, and also put the container with CuCl in a CLEAN
plastic bucket, just in case the container breaks.
[Edited on 17-1-08 by woelen]Jor - 17-1-2008 at 11:26
ok, but if the top layer is oxidised then the bottom is not, right? Indeed the top layer is green, but I cant see whats under it, because the solid is
in a very brown Baker bottle, where it is impossible to look through.
I will try a to break it with copper metal. I have a HPDE container ready!YT2095 - 17-1-2008 at 11:36
empty all you can out the test tube, and then heat it to near red heat, until the copper contam turns black.
then after it`s cooled, add some dilute sulphuric acid to it (enough to cover it all) and reheat it, it Should all dissolve to leave a light clear
blue soln of copper sulphate.
edit: and if after 10 mins you see no change, add a little H2O2, and try again.
[Edited on 17-1-2008 by YT2095]12AX7 - 17-1-2008 at 12:02
Do not heat it, you will only evaporate noxious copper fumes. CuCl2 has a high vapor pressure, and probably half your copper metal will be lost as
fumes if you heat it to dryness.
If you must process it, break it up mechanically and apply aqueous chemistry. You might reduce to Cu(0), treat with base to Cu2O (rusty looking!)
and/or with acid to get a combination of CuCl and CuCl2, or CuSO4 and Cu(0), etc., or with oxidizer to get just Cu(II). Cu2O of course can be roasted
in air to a composition nearly CuO, which can be dissolved in acid, or reduced (with carbon or as a thermite charge) to pure metal.
Chrome plated tools will not react with copper salts very much. Steel will not react with dry salts (I suppose your product may've absorbed
moisture). A piece of titanium, if you have some, will not react with a damn thing. I would not recommend pounding on a ceramic of any kind, even
SiC, etc., if you had such a piece.
Timjamit - 21-6-2011 at 01:55
I need some help with the chemical reaction of copper I chloride synthesis. I reacted copper II sulfate with sodium chloride in one beaker and then
created another beaker with a solution of sodium hydroxide and sodium metabisulfite. I mixed the two solutions and that produced a snow white ppt of
copper I chloride.
Now I need to know what happened. Here's what I see happening. When NaCl was added to CuSO4 5H20, the result was Na2SO4 and CuCl2 (copper II
chloride).
Then I added the metabisulfite to reduce Cu++ to Cu+. But what is the function of NaOH?
And how you would write the chemical reaction of adding sodium metabisulfite and sodium hydroxide?
Metabisulfite is Na2S2O5, not NaHSO3. Reaction with sodium hydroxide leads to sodium sulfite, Na2SO3, if I'm not mistaken.
So are you saying that the reaction is something like this:
NaOH + Na2S2O5 + CuCl2 --> CuCl + Na2SO3.
So is sodium hydroxide converting the metabisulfite into a sulfite in solution and that in turn reduces the Cu++ to Cu+.
What do you think?woelen - 22-6-2011 at 02:51
Your reaction equation is wrong. Metabisulfite has sulphur in oxidation state +4 and sulfite also has. The sulphur does not change redox state in your
equation, while the copper does. This reaction is not possible.DJF90 - 22-6-2011 at 05:42
It is also not a balanced equation.LanthanumK - 22-6-2011 at 06:00
I tried to figure out a reaction for the reduction of copper(II) chloride with metabisulfite and it appears that the reaction cannot be stated as one
balanced equation; there are several possible results.woelen - 22-6-2011 at 06:59
There is exactly one possibility. Metabisulfite is oxidized to sulfate and copper(II) is reduced to copper(I).
In water, S2O5(2-) becomes mostly SO2 + SO3(2-) and both are oxidized to sulfate. The net equation in terms of solid compounds is as follows:
Reply in chemistry mode: Yes, you are right. The discussion in the previous post was a more formal one in order to make things not more complicated
than they already are.
Actually, the situation is even more complicated. Keep inmind that the ion S2O5(2-) has the following structure: [O3S-SO2](2-). There is a direct
S-S bond in this ion. It is not like pyrosulfate, where there is an -S-O-S- core in the ion.
When water is nearby, then it bonds through its somewhat negatively charged oxygen to the SO2-side of the ion, and all charge is repelled towards the
O3S part of the ion. The ion breaks apart and gives an SO3(2-) part and a O2S:OH2 part (O2S:OH2 is not the same as what we write as H2SO3, it only
exists a very brief period of time as an intermediate). Because negative charge is drawn away from the oxygen in the water molecule and distributed
over the entire molecule, the latter particle quickly looses one proton and becomes O2SOH(-) and the proton quickly combines with the SO3(2-) to form
another O2SOH(-) ion.
Secondary, there are equilibria in which H2SO3 and H2O + SO2 are formed. For this reason, a metabisulfite solution always has a smell of SO2, albeit a
weak one.
