Sciencemadness Discussion Board

Pure CuCl2 from dirty HCl

Parakeet - 1-5-2024 at 06:57

Is there a way to make copper (II) chloride from copper (I) chloride without using hydrochloric acid? I don’t want to use chlorine gas either.

It might seem like a weird question, so let me explain:
I need some copper(II) chloride, and the only hydrochloric acid I have is some really dirty and dillute (10%) muriatic acid that was sold as detergent. It can be used to dissolve copper slowly, but it has some inorganic impurities and a strong perfume. It's also slightly viscous. If I tried making copper(II) chloride directly from it, its scent would definitely remain in the product, and I don't want copper(II) chloride crystals that smell like floor detergent. Apparently, copper(II) chloride is fairly soluble even in cold water, so I doubt whether recrystallizations would be effective.
So my idea was to first make copper(I) chloride, which is insoluble and can be washed easily, and then convert back to copper(II) chloride to get a pure product. I've already made sure that the first step of producing copper(I) chloride is possible using the stinky acid.

Yeah, it's a roundabout method but I can probably manage because I don't need a large amount. Although if anyone have a better idea, that would also be appreciated.

[edit] typo
[edit2] Changed title

[Edited on 2024-5-2 by Parakeet]

[Edited on 2024-5-2 by Parakeet]

Sir_Gawain - 1-5-2024 at 07:36

Place an open container of the dirty acid and an open container of water inside a larger sealed container. HCl will diffuse from the acid into the water until they’re roughly the same concentration.

bnull - 1-5-2024 at 11:15

If sodium chloride is not an issue, react copper sulfate and sodium chloride, with methanol or ethanol as solvent. The solubility of copper (II) chloride in water, methanol, and ethanol is almost the same. Here's a picture:

20240501_160103.jpg - 58kB


Top layer: copper (II) chloride in ethanol; bottom layer: sodium sulfate and excess sodium chloride.

[Edited on 1-5-2024 by bnull]

Parakeet - 1-5-2024 at 17:52

Quote: Originally posted by Sir_Gawain  
Place an open container of the dirty acid and an open container of water inside a larger sealed container. HCl will diffuse from the acid into the water until they’re roughly the same concentration.


You mean the so-called two-container technique? I've actually tried that, but the perfume also diffused, and the concentration of HCl I got was really dilute.


Quote: Originally posted by bnull  
If sodium chloride is not an issue, react copper sulfate and sodium chloride, with methanol or ethanol as solvent. The solubility of copper (II) chloride in water, methanol, and ethanol is almost the same.


That's an interesting method! Yes, small amounts of sodium contamination won't be a problem.
I cannot buy copper sulfate: it's a restricted compound in my country but maybe I can dissolve copper in sodium bisulfate and make crude copper sulfate that way.

Precipitates - 1-5-2024 at 18:10

Quote: Originally posted by Sir_Gawain  
Place an open container of the dirty acid and an open container of water inside a larger sealed container. HCl will diffuse from the acid into the water until they’re roughly the same concentration.


At low concentrations this process is ridiculously slow. I've tried with 22% HCl before and after two weeks the concentration of HCl in my other solution was barely 1%.

It probably doesn't work due to the azeotrope. At 20.2% HCl and below the vapour will contain water. It's a great technique if it's concentrated acid, but not if it isn't.

Edit: Just see Parakeet beat me to it!

[Edited on 2-5-2024 by Precipitates]

Precipitates - 1-5-2024 at 19:33

Quote: Originally posted by Parakeet  
If I tried making copper(II) chloride directly from it, its scent would definitely remain in the product, and I don't want copper(II) chloride crystals that smell like floor detergent.


You may be able to get rid of detergent smells by heating the product - copper (II) chloride decomposes at quite a high temperature (400°C+) so you can heat strongly to boil off volatile organics.

Texium - 1-5-2024 at 20:30

You could also wash the product with acetone or another organic solvent that won’t dissolve the salt to remove the organic impurities.

If you have sodium bisulfate though, why not just make your own HCl from that and NaCl?

Parakeet - 2-5-2024 at 01:02

Quote: Originally posted by Texium  
If you have sodium bisulfate though, why not just make your own HCl from that and NaCl?
You are absolutely right, that’s one option. I was just wondering if detergent grade HCl could be used, because they are really cheap and I already have a lot of them and wanted to use preferentially. (Sodium bisulfate on the other hand, has many uses, so I didn’t want to waste it.)

Now I’m thinking to first make crude copper(II) chloride crystals, and then try cleaning or roasting it.

Unfortunately, copper(II) chloride is soluble in many organic solvents, including acetone, EtOH, MeOH and even ethyl acetate! I can try petroleum ether though.

P.S. Considering these discussion so far, maybe I should change the title.

EF2000 - 2-5-2024 at 02:27

Since you are concerned with organic impurities, maybe you can add hydrogen peroxide to your muriatic acid to oxidize them?
I don't know how it works with dilute HCl, but it's known method to purify sulfuric acid.
After that mixture of hydrochloric acid and hydrogen peroxide (and inorganic impurities) will more easily dissolve copper, since it acts as oxidizing acid.

bnull - 2-5-2024 at 05:06

Quote: Originally posted by EF2000  
Since you are concerned with organic impurities, maybe you can add hydrogen peroxide to your muriatic acid to oxidize them?
I don't know how it works with dilute HCl, but it's known method to purify sulfuric acid.
After that mixture of hydrochloric acid and hydrogen peroxide (and inorganic impurities) will more easily dissolve copper, since it acts as oxidizing acid.

If iron is one of the impurities, peroxide will oxidise it to iron (III) chloride, which is unfortunately soluble in many organic solvents. And the perfume may be oxidised to something worse (dead skunk on the backseat in a 1000-mile trip, for example) or even dyestuff. Without knowing what the perfume is, who knows.

Let's see: you cannot buy copper sulfate because it is regulated; your hydrochloric acid is full of wossnames, has a scent, and is quite viscous; you have bisulfate but would rather save it for something else. Am I missing something?

Another roundabout method (which unfortunately uses sodium bisulfate): (1) make copper (II) chloride from that acid; (2) precipitate copper (II) hydroxide with a soluton of sodium carbonate in water; (3) wash the precipitate with water and whatever organic solvent you have until the scent is gone; (4) react the precipitate with the smallest amount possible of sodium bisulfate so that the precipitate dissolves; (5) crystallise the salts and perform the reaction with sodium chloride I described earlier.

If iron is present in the acid (I'm quite sure it is), it will precipitate as hydroxide in step (2) and convert to sulfate in step (4). Iron sulfates are practically insoluble in ethanol (iron chlorides, on the other hand, are quite soluble), so the only impurity will be sodium. The amount of bisulfate used will be less than if you made HCl from it and NaCl (no losses from evaporation of HCl, for example).

Do you happen to have any other acids, such as nitric acid?

Parakeet - 2-5-2024 at 06:34

Quote: Originally posted by EF2000  
Since you are concerned with organic impurities, maybe you can add hydrogen peroxide to your muriatic acid to oxidize them?
I don't know how it works with dilute HCl, but it's known method to purify sulfuric acid.
After that mixture of hydrochloric acid and hydrogen peroxide (and inorganic impurities) will more easily dissolve copper, since it acts as oxidizing acid.

Adding hydrogen peroxide didn't change anything. I think it needs to be concentrated and hot enough, but I only have 3% H2O2, and heating the solution will release HCl gas.


The method of using copper hydroxide and sodium bisulfate is probably possible.
However, judging from the color, I don't think that Fe is present in it, bnull. It's nearly colorless and transparent, even after adding a few drops of hydrogen peroxide. Iron(III) ions have a strong color, so I'm quite sure about this.

Next time I have some time, I will make some copper(II) chloride from the detergent, and see how much additives remain.


Quote: Originally posted by bnull  

Do you happen to have any other acids, such as nitric acid?
I have a small amount of nitric acid and sulfuric acid, but that's even more valuable than sodium bisulfate, and the quantity is not enough anyway. I also have oxalic acid, but I doubt if it will be useful here.

Texium - 2-5-2024 at 06:42

Quote: Originally posted by Parakeet  
Unfortunately, copper(II) chloride is soluble in many organic solvents, including acetone, EtOH, MeOH and even ethyl acetate! I can try petroleum ether though.
The anhydrous form may be, but I suspect that hydrated copper(II) chloride would not be very soluble in aprotic organic solvents like acetone or ethyl acetate. I would still give it a try. What’s the worst that can happen, it dissolves and you have to evaporate it down?

EF2000 - 2-5-2024 at 06:50

I have another idea: copper can be dissolved in ammonia solution with ammonium chloride (or even sodium chloride).
Resulting tetraammine copper(II) chloride should decompose when heated to copper(II) chloride and ammonia.

Dissolution in ammonia works better with "ammonia water" (24%), but diluted household grade ammonia also works, but more slowly, I once made tetraammine copper perchlorate from 10% ammonia and lithium perchlorate. Thin copper wire (incompletely) dissolved in several hours. Edit: in 24 hours, actually.

What I'm not sure is whether decomposition will produce copper chloride or copper oxide and HCl.

[Edited on 2-5-2024 by EF2000]

clearly_not_atara - 2-5-2024 at 11:08

I vote for heating the CuCl2 crystals to 400 C, then cool and recrystallize. Should be pretty good after that.

bnull - 2-5-2024 at 11:46

Quote: Originally posted by Texium  
The anhydrous form may be, but I suspect that hydrated copper(II) chloride would not be very soluble in aprotic organic solvents like acetone or ethyl acetate.

They are soluble. The entry "Copper(II) Chloride" in the Encyclopedia of Reagents for Organic Synthesis (EROS) says:
Quote:
Solubility: anhydrous: sol water, alcohol, and acetone; dihydrate: sol water, methanol, ethanol; mod sol acetone, ethyl acetate; sl sol Et2O.

The CRC Handbook 9th says almost the same. I wish they had given the values.

DraconicAcid - 2-5-2024 at 12:00

I've found that the dihydrate is nicely soluble in methanol and ethanol, but not so much in isopropanol. I don't remember if I've tried acetone.

Precipitates - 2-5-2024 at 17:53

Quote: Originally posted by bnull  
Quote: Originally posted by Texium  
The anhydrous form may be, but I suspect that hydrated copper(II) chloride would not be very soluble in aprotic organic solvents like acetone or ethyl acetate.

They are soluble. The entry "Copper(II) Chloride" in the Encyclopedia of Reagents for Organic Synthesis (EROS) says:
Quote:
Solubility: anhydrous: sol water, alcohol, and acetone; dihydrate: sol water, methanol, ethanol; mod sol acetone, ethyl acetate; sl sol Et2O.

The CRC Handbook 9th says almost the same. I wish they had given the values.



It's not that soluble in acetone (solubility data for copper (II) chloride) (1-3 g copper (II) chloride per 100 g acetone), so if you're okay with a small loss of yield for increased purity (and non-smelly crystals), washing the product with acetone may be acceptable.

pesco - 2-12-2024 at 09:47

Quote: Originally posted by clearly_not_atara  
I vote for heating the CuCl2 crystals to 400 C, then cool and recrystallize. Should be pretty good after that.


I was about to suggest this route, but clearly_not_atara be me to it :(:D.

It is a easiest and cheapest way and if it works then you are sorted.
At high heat all the organics should be destroyed.

Try on small amount and see what's the outcome.