Keras
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Copper (I) Chloride + Bleach → Copper (II) oxide?
Folks,
I was trying to make a small quantity of copper (I) chloride for experimenting with a Sandmeyer reaction.
Copper (I) chloride I made from copper (II) sulphate
CuSO₄ + Na₂CO₃ → Cu₂(CO₃)(OH)₂
Cu₂(CO₃)(OH)₂ + HCl → CuCl₂
Then reduction of CuCl₂ by sodium hydrosulfite into copper (I) chloride.
Well, I got a small amount of white powder. When, to test it, I tried to react it with bleach hoping to form copper (II) chloride back, I got
apparently copper (II) oxide?
NaClO + CuCl → CuO + Cl₂
Is that expected?
(So yeah, when I added HCl it dissolved instantly into that yellow/green solution typical of copper (II) chloride).
(And the smears of white powder are turning blue in contact with air).
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RU_KLO
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It seems so, my understanding is that you are removing chlorine from the equation, because of gas.
Did you smell chlorine gas when perfoming the experiment?
Maybe Cl- + Cl+ -> Cl2(g) is "more" thermodinamically favorable than Cu+ (/Cu++) + Cl- -> CuCl (/ CuCl2)
So Chlorine will be removed from the mix, and only Cu++ + O-- will be there.
Someone who understands chemistry thermodinamics can confirm?
Go SAFE, because stupidity and bad Luck exist.
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Keras
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It was very small scale, so I suppose the chlorine either stayed in solution or there was so few produced that I didn't smell it.
Sodium hydrosulphite doesn’t seem to be the best reducing agent in this case. I got a miserable yield, probably because most of the copper was
reduced directly as copper metal rather than copper (I) chloride. I’ll retry with a somewhat less aggressive reductant (Ascorbic acid? Sodium
metabisulphite? Ethanol?).
[Edited on 12-5-2025 by Keras]
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RU_KLO
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from wiki:
Copper(I) chloride can also be prepared by reducing copper(II) chloride with sulfur dioxide, or with ascorbic acid (vitamin C) that acts as a reducing
sugar: 2 CuCl 2 + SO 2 + 2 H 2O → 2 CuCl + H 2SO 4 + 2 HCl.
2 CuCl 2 + C 6H 8O 6 → 2CuCl + 2HCl + C 6H 6O.
Many other reducing agents can be used.
https://en.wikipedia.org/wiki/Copper(I)_chloride
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Keras
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Yeah, I read that.
I’m a bit sceptical because it is known that ascorbic acid can reduce copper sulphate all the way to copper (0) nanoparticles. See here for example.
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bnull
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Quote: Originally posted by Keras  | Well, I got a small amount of white powder. When, to test it, I tried to react it with bleach hoping to form copper (II) chloride back, I got
apparently copper (II) oxide?
NaClO + CuCl → CuO + Cl₂
Is that expected? |
Not exactly. You must include the sodium hydroxide also in solution (always present in bleach). Like this:
$$NaClO(aq)+2CuCl(s)+2NaOH(aq) \rightarrow 2CuO(s)+3NaCl(aq)+H_2O(l).$$
You usually get copper(ii) hydroxide in place of the oxide but it depends on the amount of NaOH in solution.
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Keras
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| Quote: Originally posted by bnull |
You usually get copper(ii) hydroxide in place of the oxide but it depends on the amount of NaOH in solution. |
Interesting. Albeit I used quite an amount of bleach, I didn't get any hydroxide. It was clearly a dark oxide powder.
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woelen
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The best way to make copper(I) chloride is adding a lot of copper metal to a concentrated solution of copper(II) chloride (or copper(II) oxide) in
concenctrated HCl. You can even use CuSO4, but this brings in undesirable ions of another type than chloride.
When you add copper, the solution first turns very dark brown, almost black. This is due to formation of a mixed oxidation state complex, a Cu(I, II)
chloride, something like Cl-Cu-Cl-Cu-Cl. It cannot be isolated, it only exists in conc. HCl. If more copper dissolves (and air is excluded), then the
solution turns lighter again, it almost becomes colorless. Due to impurities, usually it remains pale olive green.
The solution, obtained in this way can be dumped in water, which was boiled for a while and allowed to cool down again, with exclusion of air. A lot
of snow-white CuCl precipitates under a dilute HCl solution, which may be very pale blue. Just store it like that. Any attempt to dry it will lead to
aerial oxidation and makes it dirty green and you get contamination with a mixed oxide/chloride of copper(II).
For copper, you can use electrical wire. The flexible wire is best, it contains many thin strands, which have a nice large surface area and dissolve
quite fast.
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DraconicAcid
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Agreed. You can wash it with de-aerated ethanol.
I seem to recall reading that you can also use acetonitrile as a solvent for the reaction of CuCl2 with Cu, as the Cu(I) ion is more stable in that
solvent. Presumably, you can precipitate out the halide using dry alcohol or acetone as an antisolvent.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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Keras
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Thanks Woelen 
In hindsight, I wonder why I used this rather intricate path, while copper sulphate probably reacts with calcium chloride to directly form copper (II)
chloride and insoluble gypsum.
This is the same path I plan to exploit to make copper (II) nitrate from calcium nitrate and copper sulphate.
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Keras
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| Quote: Originally posted by woelen | The best way to make copper(I) chloride is adding a lot of copper metal to a concentrated solution of copper(II) chloride (or copper(II) oxide) in
concenctrated HCl.
[…]
For copper, you can use electrical wire. The flexible wire is best, it contains many thin strands, which have a nice large surface area and dissolve
quite fast. |
That’s what I did today. I made copper (II) chloride by double displacement from copper (II) sulphate and calcium chloride, then I used your method
(found a reference to it somewhere) and even excluded air during the reduction using an erlenmeyer covered with a watch glass and in which I had blown
argon. It went from dark brown to medium yellow.
My boiled water was maybe a bit too hot at first, so nothing happened, but then tiny white crystals started to snow, and the water turned slightly
blue. This was where the crash happened :p I decided to add a pinch of sodium hydrosulphite to keep oxygen out, and.......... it reacted with whatever
was in solution to form a dark powder and turn the solution yellow (colloidal sulphur?) So now I had at the bottom of my beaker a mix of white copper
(I) chloride and something else.
Apparently, copper (I) chloride can be dried. Procedure is thus: first rince with water in which some HCl has been added, then twice with 99% ethanol,
then a final time with diethyl ether. All the liquids must be poured before the previous one has been completely drained to exclude oxygen. After the
final ether washing, the product can be further sucked dry.
Not tried the procedure for the time being, so if you feel bold…
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DraconicAcid
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With sulphite, you can get Chevreul's salt. You need more HCl to turn that into CuCl.
Please remember: "Filtrate" is not a verb.
Write up your lab reports the way your instructor wants them, not the way your ex-instructor wants them.
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woelen
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The method, I presented, requires highly concentrated HCl. So, if you make CuCl2-solution from CaCl2 and CuSO4, then the resulting solution of CuCl2
must be boiled down considerably, until is becomes saturated (dark green). This then must be mixed with the HCl and copper must be added, in excess.
Adding the dithionite was not a smart action  The dark powder is copper(II)
sulfide, which is highly insoluble, even in dilute HCl. Dithionite decomposes in acids, one of the decomposition products being H2S (you can smell
it). You get something like Wackenroder's solution from dithionite in acid. It is a dirty ill-defined solution, containing many sulphur-species, with
a yellow/orange color and in the long run it turns turbid and pale yellow, due to formation of very finely divided impure sulphur.
You can add a pinch of plain sufite or metabisulfite to the acidic solution. This leads to formation of SO2. It does not protect against oxidation,
but it reduces any formed copper(II). If any copper(I) is oxidized to copper(II) by oxygen, then the sulfite (slowly) reduces the copper(II) back to
copper(I).
My personal experience is that drying copper(I) chloride without olive-green discoloration is really hard to achieve. I also tried rinsing with
ethanol, diethylether and acetone, but none could really prevent discoloration. Even tiny amounts of water spoil the snow-white CuCl. E.g. with
diethylether, the solid becomes cold, due to evaporation, and then it attracts a little water vapor from the air and then it turns green/brown. In a
professional lab, all handling of the CuCl is done in a glove box under nitrogen.
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Keras
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| Quote: Originally posted by woelen | The method, I presented, requires highly concentrated HCl. So, if you make CuCl2-solution from CaCl2 and CuSO4, then the resulting solution of CuCl2
must be boiled down considerably, until is becomes saturated (dark green). This then must be mixed with the HCl and copper must be added, in excess.
|
I followed the preparation presented p. 154 of ‘Inorganic Preparations – A Laboratory Manual’ by H.F.Walton (1948). The recipe calls for 43 g of
CuCl₂ in 100 mL of water (to which are added 100 mL of 36% HCl). In this case, with the double displacement reaction, I started from the molar
equivalent of CuSO₄ and CaCl₂, which resulted in a very dark blue solution of (hopefully ~ 43 g) CuCl₂ in 150 mL of water that I gently
evaporated to 100 mL.
| Quote: Originally posted by woelen |
Adding the dithionite was not a smart action The dark powder is copper(II)
sulfide, which is highly insoluble, even in dilute HCl. Dithionite decomposes in acids, one of the decomposition products being H2S (you can smell
it). |
Yes, I immediately suspected so. Good news is that I can probably purify my cuprous chloride by redissolving it in concentrated HCl and filtering. The
copper sulphide should remain in the filter. Then I again dilute the solution.
| Quote: Originally posted by woelen | | You can add a pinch of plain sufite or metabisulfite to the acidic solution. This leads to formation of SO2. It does not protect against oxidation,
but it reduces any formed copper(II). If any copper(I) is oxidized to copper(II) by oxygen, then the sulfite (slowly) reduces the copper(II) back to
copper(I). |
Thanks, I’ll try that. Or I’ll just tolerate the slightly blue solution…
[EDIT: Added the preparation scheme.]
[Edited on 16-5-2025 by Keras]
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Keras
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So I redid it on a smaller scale and this time I didn't add anything to the final diluted solution of HCl. It was funny to see the copper wires
covered in tiny crystal of white cuprous chloride, that turned blue in 3 to 4 hours.
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Keras
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| Quote: Originally posted by Keras |
| Quote: Originally posted by woelen |
Adding the dithionite was not a smart action The dark powder is copper(II)
sulfide, which is highly insoluble, even in dilute HCl. Dithionite decomposes in acids, one of the decomposition products being H2S (you can smell
it). |
Yes, I immediately suspected so. Good news is that I can probably purify my cuprous chloride by redissolving it in concentrated HCl and filtering. The
copper sulphide should remain in the filter. Then I again dilute the solution. |
That worked. I redissolved the mixture of black and white crystals into 36% hydrochloric acid, put some fresh copper wire bits in the mixture, let it
stir for an hour or so at room temperature and filtered it. The black crystals remained on the filter, and the filtrate, once diluted into a litre of
tap water, reverted to pure white crystals. I also added, as you suggested, a bit of sodium sulfite to prevent over-oxidation. At the end, I got like
the equivalent of 40 mL of white powder in a 100 mL bottle. Difficult to estimate a weight, though.
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