Parakeet
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Pure CuCl2 from dirty HCl
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]
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Sir_Gawain
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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.
“Alchemy is trying to turn things yellow; chemistry is trying to avoid things turning yellow.” -Tom deP.
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bnull
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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:
Top layer: copper (II) chloride in ethanol; bottom layer: sodium sulfate and excess sodium chloride.
[Edited on 1-5-2024 by bnull]
Quod scripsi, scripsi.
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Parakeet
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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.
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Precipitates
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| 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]
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Precipitates
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| 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.
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Texium
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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?
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Parakeet
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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.
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EF2000
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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.
Wroom wroom
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bnull
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| 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?
Quod scripsi, scripsi.
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Parakeet
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| 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.
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.
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Texium
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| 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?
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EF2000
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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]
Wroom wroom
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clearly_not_atara
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I vote for heating the CuCl2 crystals to 400 C, then cool and recrystallize. Should be pretty good after that.
[Edited on 04-20-1969 by clearly_not_atara]
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bnull
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| 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.
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The CRC Handbook 9th says almost the same. I wish they had given the values.
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DraconicAcid
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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.
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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Precipitates
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| 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.
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The CRC Handbook 9th says almost the same. I wish they had given the values.
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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.
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