Sciencemadness Discussion Board
Not logged in [Login ]
Go To Bottom

Printable Version  
Author: Subject: Copper (I) Chloride + Bleach → Copper (II) oxide?
Keras
International Hazard
*****




Posts: 1062
Registered: 20-8-2018
Location: (48, 2)
Member Is Offline


[*] posted on 12-5-2025 at 04:37
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).
View user's profile View All Posts By User
RU_KLO
Hazard to Others
***




Posts: 269
Registered: 12-10-2022
Location: Argentina
Member Is Offline


[*] posted on 12-5-2025 at 05:38


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.
View user's profile View All Posts By User
Keras
International Hazard
*****




Posts: 1062
Registered: 20-8-2018
Location: (48, 2)
Member Is Offline


[*] posted on 12-5-2025 at 07:33


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]
View user's profile View All Posts By User
RU_KLO
Hazard to Others
***




Posts: 269
Registered: 12-10-2022
Location: Argentina
Member Is Offline


[*] posted on 12-5-2025 at 09:12


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




Go SAFE, because stupidity and bad Luck exist.
View user's profile View All Posts By User
Keras
International Hazard
*****




Posts: 1062
Registered: 20-8-2018
Location: (48, 2)
Member Is Offline


[*] posted on 12-5-2025 at 11:36


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.
View user's profile View All Posts By User
bnull
National Hazard
****




Posts: 699
Registered: 15-1-2024
Location: East Woods
Member Is Offline

Mood: Editorial

[*] posted on 13-5-2025 at 06:14


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.




Quod scripsi, scripsi.

B. N. Ull

We have a lot of fun stuff in the Library.

Read The ScienceMadness Guidelines. They exist for a reason.
View user's profile View All Posts By User
Keras
International Hazard
*****




Posts: 1062
Registered: 20-8-2018
Location: (48, 2)
Member Is Offline


[*] posted on 13-5-2025 at 08:15


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.
View user's profile View All Posts By User
woelen
Super Administrator
*********




Posts: 8107
Registered: 20-8-2005
Location: Netherlands
Member Is Offline

Mood: interested

[*] posted on 13-5-2025 at 12:26


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.




The art of wondering makes life worth living...
Want to wonder? Look at https://woelen.homescience.net
View user's profile Visit user's homepage View All Posts By User
DraconicAcid
International Hazard
*****




Posts: 4434
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline

Mood: Semi-victorious.

[*] posted on 13-5-2025 at 14:00


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.
View user's profile View All Posts By User
Keras
International Hazard
*****




Posts: 1062
Registered: 20-8-2018
Location: (48, 2)
Member Is Offline


[*] posted on 13-5-2025 at 21:44


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.
View user's profile View All Posts By User
Keras
International Hazard
*****




Posts: 1062
Registered: 20-8-2018
Location: (48, 2)
Member Is Offline


[*] posted on 15-5-2025 at 11:50


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… :)
View user's profile View All Posts By User
DraconicAcid
International Hazard
*****




Posts: 4434
Registered: 1-2-2013
Location: The tiniest college campus ever....
Member Is Offline

Mood: Semi-victorious.

[*] posted on 15-5-2025 at 12:13


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.
View user's profile View All Posts By User
woelen
Super Administrator
*********




Posts: 8107
Registered: 20-8-2005
Location: Netherlands
Member Is Offline

Mood: interested

[*] posted on 15-5-2025 at 23:19


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.




The art of wondering makes life worth living...
Want to wonder? Look at https://woelen.homescience.net
View user's profile Visit user's homepage View All Posts By User
Keras
International Hazard
*****




Posts: 1062
Registered: 20-8-2018
Location: (48, 2)
Member Is Offline


[*] posted on 16-5-2025 at 02:07


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]

Screenshot 2025-05-16 at 12.00.28.png - 886kB
View user's profile View All Posts By User
Keras
International Hazard
*****




Posts: 1062
Registered: 20-8-2018
Location: (48, 2)
Member Is Offline


[*] posted on 16-5-2025 at 11:17


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.
View user's profile View All Posts By User
Keras
International Hazard
*****




Posts: 1062
Registered: 20-8-2018
Location: (48, 2)
Member Is Offline


[*] posted on 22-5-2025 at 23:40


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.
View user's profile View All Posts By User

  Go To Top