Rhodanide
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What happened?
I've been trying to make Copper (I) Chloride for the past few years now and it's never worked.
I've tried:
Sodium metabisulfite - nothing; crap
Sulfur dioxide - not only is it incredibly inefficient for me to make it, nothing results
CuCl2 + Cu metal - takes way too long and doesn't produce any product.
So I tried Oxalic acid, because it's on a list of common reducing agents. Mine is pure, and in the Dihydrate form. I dissolved it in water to make a
saturated solution and prepared a syringe full of highly concentrated CuCl2 solution in water with residual HCl. I then added a drop of the chloride
solution to the Oxalic acid solution which yielded for a short time, a fine white precipitate of what seemed to be CuCl. I then got excited and added
all of it at once, which produced a baby blue, extremely fine precipitate. I thought that the blue was from just extra dissolved CuCl2, so I ran,to go
filter it. It wasn't just some residual CuCl2. It was the precipitate. It goes right through the filter - TWO actually - and it's also extremely
possible that it's CuC2O4. I'M REALLY GETTING FED UP WITH THIS.
What's a commonly available, easy to buy, simple to find reducing agent? Please, help a brother out here. I'm losing hope with this one!
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Texium
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Ascorbic acid.
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Melgar
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When I was dissolving copper into hydrochloric acid with H2O2, the solution would get dark and opaque when it ran out of chloride ions, meaning I had
to add more HCl, which would turn it back to transparent green. Adding H2O2 would just release oxygen if I tried that. If this wasn't Cu(I)Cl, what
was it? I assumed that the CuCl2 was oxidizing Cu metal, which was acting as the reducing agent, but I take it it's not that easy?
[Edited on 6/4/17 by Melgar]
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Geocachmaster
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I have a book with a few ways on how to prepare cuprous chloride (solution). One uses copper sulfate, sodium chloride, sodium bisulfite, and sodium
hydroxide. Another uses copper (II) chloride, hydrochloric acid, and copper metal. If you want later I can post the procedures here or send it to you,
right now I'm supposed to be doing homework. 
From Unitized Experiments in Organic Chemistry (fourth edition), experiment 55, page 368.
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UC235
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Download Brauer (Handbook of Preparative Inorganic Chemistry) from the forum library, pg 1005. Several preparations are available through orgsyn as
part of Sandmeyer reactions.
I don't have specifics, but I have in the past had success with sodium metabisulfite reduction of concentrated cupric sulfate in the presence of
excess sodium chloride. Pouring the solution on ice is seemingly needed for decent yield. Keep volumes as small as possible, it is somewhat soluble in
water (presumably through complex formation). The still-wet product is exceptionally prone to oxidation and needs to be dried as quickly as possible.
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Rhodanide
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Forgot to add that worked likely the worst. I wasted over 100 grams for nothing.
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Rhodanide
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I also forgot to add that I would love to know where to get Sodium Bisulfite that isn't a wine maker's shop or whatnot, because there's nothing like
that anywhere around here.
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ninhydric1
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You could prepare sodium bisulfite by bubbling sulfur dioxide gas into a saturated solution of sodium carbonate.
Na2CO3 + 2 SO2 + H2O => 2 NaHSO3 + CO2
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Waffles SS
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42 g of copper(II) chloride are dissolved in 100 ml of hot water containing 32 g metallic copper (granulated or turnings). 200 ml of hydrochloric acid
(d=1.175g/ml) is added and the reaction mixture is boiled gently under reflux until the solution becomes colorless. During an experiment additionally
a small amount of concentrated hydrochloric acid is added to ensure that solution becomes colorless. This operation lasts one to two hours and when is
complete the colorless solution is decanted from the metallic copper into a cylinder filled with cold distilled water. The dilution with water
decomposes H3CuCl4 complex and cheesy precipitated copper(I) chloride is allowed to settle, decanted immediately, quickly filtered, washed with
alcohol, ether, and dried in a vacuum desiccator. Copper(I) chloride forms heavy white masses, insoluble in water, and oxidizes easily in the air
yielding green oxidized product.
Laboratory manual of inorganic preparations, by H. T. Vulte, 73-74, 1895
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unionised
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Treatment of Copper(I) chloride with sodium hydroxide solution will convert it to the oxide.
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Rhodanide
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Quote: Originally posted by ninhydric1  | You could prepare sodium bisulfite by bubbling sulfur dioxide gas into a saturated solution of sodium carbonate.
Na2CO3 + 2 SO2 + H2O => 2 NaHSO3 + CO2
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What I said may have been missed, but I don't have an efficient way to generate SO2.
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Rhodanide
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| Quote: Originally posted by Waffles SS | 42 g of copper(II) chloride are dissolved in 100 ml of hot water containing 32 g metallic copper (granulated or turnings). 200 ml of hydrochloric acid
(d=1.175g/ml) is added and the reaction mixture is boiled gently under reflux until the solution becomes colorless. During an experiment additionally
a small amount of concentrated hydrochloric acid is added to ensure that solution becomes colorless. This operation lasts one to two hours and when is
complete the colorless solution is decanted from the metallic copper into a cylinder filled with cold distilled water. The dilution with water
decomposes H3CuCl4 complex and cheesy precipitated copper(I) chloride is allowed to settle, decanted immediately, quickly filtered, washed with
alcohol, ether, and dried in a vacuum desiccator. Copper(I) chloride forms heavy white masses, insoluble in water, and oxidizes easily in the air
yielding green oxidized product.
Laboratory manual of inorganic preparations, by H. T. Vulte, 73-74, 1895 |
While I don't have much time to do that, I don't either have a generously available source of pure copper metal. I make my CuCl2 from Cu Carbonate and
HCl.
Thanks for all of the help, btw guys!
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Texium
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Did you use NaCl as a
source of chloride ions? Cause if not then nothing will happen. Check out the SM wiki page. Chemistry doesn't just decide to work for some people and
not for others. If a reaction that should work doesn't work for you, clearly you're doing something wrong.
http://www.sciencemadness.org/smwiki/index.php/Copper(I)_chloride
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ELRIC
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| Quote: Originally posted by Tetra | | Quote: Originally posted by Waffles SS | 42 g of copper(II) chloride are dissolved in 100 ml of hot water containing 32 g metallic copper (granulated or turnings). 200 ml of hydrochloric acid
(d=1.175g/ml) is added and the reaction mixture is boiled gently under reflux until the solution becomes colorless. During an experiment additionally
a small amount of concentrated hydrochloric acid is added to ensure that solution becomes colorless. This operation lasts one to two hours and when is
complete the colorless solution is decanted from the metallic copper into a cylinder filled with cold distilled water. The dilution with water
decomposes H3CuCl4 complex and cheesy precipitated copper(I) chloride is allowed to settle, decanted immediately, quickly filtered, washed with
alcohol, ether, and dried in a vacuum desiccator. Copper(I) chloride forms heavy white masses, insoluble in water, and oxidizes easily in the air
yielding green oxidized product.
Laboratory manual of inorganic preparations, by H. T. Vulte, 73-74, 1895 |
While I don't have much time to do that, I don't either have a generously available source of pure copper metal. I make my CuCl2 from Cu Carbonate and
HCl.
Thanks for all of the help, btw guys! |
Do you live anywhere near new residential construction?
That's how I get my copper, from the scrap wire the electricians leave laying around
Might want to ask the contractor first though
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RogueRose
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Looking at the Wiki page CuCl, it states the folllowing:
| Quote: |
Copper(I) chloride can also be prepared by reducing copper(II) chloride, e.g. with sulfur dioxide:
2 CuCl2 + SO2 + 2 H2O → 2 CuCl + H2SO4 + 2 HCl
Many other reducing agents can be used. |
In another thread about making Cu2O there was talk about using monosaccharide sugars as a reducing agent (glucose & fructose - honey is a good
source, all monosaccharides I believe and also molasses will work which is a 2:1:1 sucrose:glucose:fructose - as well as one thing that mentioned
sugar invert which is table sugar mixed with something like citric acid to break down into glucose/fructose). Could the sugar be used to replace the
SO2 in the above equation
On another note, I was playing with some CuSO4 crystals and 12.5% NaClO. The two were mixed (about 4g of CuSO4 x 5H2O & 30ml NaClO). The
reaction was instant and started bubbling pretty furiously while emitting some SO2 (and possibly a slight amount of Cl2). The mixture turned an
aqua-green.
Here is what it looked like during the reaction and after it settled. I'm curious as to whether it has formed CuCl, CuCl2 or Cu(OH)2. It looks like
all of them can have the same color in various solutions (if mixed with other compounds)so it is difficult to determine ATM.
 
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Melgar
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Well, the green indicates copper chloride; the copper probably just catalyzed the hypochlorite decomposition into oxygen and chloride. The pale green
stuff is probably just a mix of sodium sulfate (white), copper hydroxide (light blue), copper carbonate, (light green) and copper chloride (green).
SO2 and Cl2 production was probably insignificant, and the bubbles would have been primarily oxygen.
OP could possibly try using copper as a reducing agent, since that would be the easiest of all of them to purify. It might take a while though.
edit: Oh. Right, that was just suggested. It seems really odd to not be able to find copper though. Don't old scraps of wire exist pretty much
everywhere? And the same with electronics that are broken beyond repair? The metal in electrical wires is always quite high purity, since pure
metals tend to conduct electricity best.
[Edited on 6/5/17 by Melgar]
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Rhodanide
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| Quote: Originally posted by Melgar | Well, the green indicates copper chloride; the copper probably just catalyzed the hypochlorite decomposition into oxygen and chloride. The pale green
stuff is probably just a mix of sodium sulfate (white), copper hydroxide (light blue), copper carbonate, (light green) and copper chloride (green).
SO2 and Cl2 production was probably insignificant, and the bubbles would have been primarily oxygen.
OP could possibly try using copper as a reducing agent, since that would be the easiest of all of them to purify. It might take a while though.
edit: Oh. Right, that was just suggested. It seems really odd to not be able to find copper though. Don't old scraps of wire exist pretty much
everywhere? And the same with electronics that are broken beyond repair? The metal in electrical wires is always quite high purity, since pure
metals tend to conduct electricity best.
[Edited on 6/5/17 by Melgar] |
Where'd the Carbonate come from, lol
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Rhodanide
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| Quote: Originally posted by zts16 | Did you use NaCl as a
source of chloride ions? Cause if not then nothing will happen. Check out the SM wiki page. Chemistry doesn't just decide to work for some people and
not for others. If a reaction that should work doesn't work for you, clearly you're doing something wrong.
http://www.sciencemadness.org/smwiki/index.php/Copper(I)_chloride |
Nope, I used the reaction of Copper (II) Carbonate and HCl. It gave a dark green solution, as expected albeit a small amount of light green
precipitate which settled on the flask bottom.
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Texium
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| Quote: Originally posted by Tetra | | Quote: Originally posted by zts16 | Did you use NaCl as a
source of chloride ions? Cause if not then nothing will happen. Check out the SM wiki page. Chemistry doesn't just decide to work for some people and
not for others. If a reaction that should work doesn't work for you, clearly you're doing something wrong.
http://www.sciencemadness.org/smwiki/index.php/Copper(I)_chloride |
Nope, I used the reaction of Copper (II) Carbonate and HCl. It gave a dark green solution, as expected albeit a small amount of light green
precipitate which settled on the flask bottom. |
I don't think you understood my post. The NaCl is necessary
in the reduction using ascorbic acid. How you produced your copper(II) chloride is irrelevant. Read the wiki page that I linked, please, that's what
it's there for.
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Melgar
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| Quote: Originally posted by Tetra | | Quote: Originally posted by Melgar | Well, the green indicates copper chloride; the copper probably just catalyzed the hypochlorite decomposition into oxygen and chloride. The pale green
stuff is probably just a mix of sodium sulfate (white), copper hydroxide (light blue), copper carbonate, (light green) and copper chloride (green).
SO2 and Cl2 production was probably insignificant, and the bubbles would have been primarily oxygen.
OP could possibly try using copper as a reducing agent, since that would be the easiest of all of them to purify. It might take a while though.
edit: Oh. Right, that was just suggested. It seems really odd to not be able to find copper though. Don't old scraps of wire exist pretty much
everywhere? And the same with electronics that are broken beyond repair? The metal in electrical wires is always quite high purity, since pure
metals tend to conduct electricity best.
[Edited on 6/5/17 by Melgar] |
Where'd the Carbonate come from, lol |
Bleach pretty much always has a small amount of sodium carbonate in it, from CO2 from the air or otherwise. It wouldn't be much, but there's always
going to be at least a little, unless you got it from Sigma Aldrich or something.
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Rhodanide
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| Quote: Originally posted by Melgar | | Quote: Originally posted by Tetra | | Quote: Originally posted by Melgar | Well, the green indicates copper chloride; the copper probably just catalyzed the hypochlorite decomposition into oxygen and chloride. The pale green
stuff is probably just a mix of sodium sulfate (white), copper hydroxide (light blue), copper carbonate, (light green) and copper chloride (green).
SO2 and Cl2 production was probably insignificant, and the bubbles would have been primarily oxygen.
OP could possibly try using copper as a reducing agent, since that would be the easiest of all of them to purify. It might take a while though.
edit: Oh. Right, that was just suggested. It seems really odd to not be able to find copper though. Don't old scraps of wire exist pretty much
everywhere? And the same with electronics that are broken beyond repair? The metal in electrical wires is always quite high purity, since pure
metals tend to conduct electricity best.
[Edited on 6/5/17 by Melgar] |
Where'd the Carbonate come from, lol |
Bleach pretty much always has a small amount of sodium carbonate in it, from CO2 from the air or otherwise. It wouldn't be much, but there's always
going to be at least a little, unless you got it from Sigma Aldrich or something. |
Huh, interesting!
Makes sense now :]
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Rhodanide
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| Quote: Originally posted by zts16 | | Quote: Originally posted by Tetra | | Quote: Originally posted by zts16 | Did you use NaCl as a
source of chloride ions? Cause if not then nothing will happen. Check out the SM wiki page. Chemistry doesn't just decide to work for some people and
not for others. If a reaction that should work doesn't work for you, clearly you're doing something wrong.
http://www.sciencemadness.org/smwiki/index.php/Copper(I)_chloride |
Nope, I used the reaction of Copper (II) Carbonate and HCl. It gave a dark green solution, as expected albeit a small amount of light green
precipitate which settled on the flask bottom. |
I don't think you understood my post. The NaCl is necessary
in the reduction using ascorbic acid. How you produced your copper(II) chloride is irrelevant. Read the wiki page that I linked, please, that's what
it's there for. |
Gotcha, gotcha!
Thanks, buddy. I'll take a look now. Just misunderstood what ya said.
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Lotilko
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I've had the same problems with copper (I) chloride. I have dissolved copper wire in H2O2/HCl. It is the cheapest and most convenient method if one
doesn't mind chlorine and a boiling soution of HCl/H2O2.
Anyways, after dissolving all of the copper, the solution becomes very dark green possibly by the formation of complex ions, and then lightens to an
emerald green color as it cools. Then the ascorbic acid solution is added, and the green color chages to a clear, with a tinge of brown/green. I found
it odd at first that the copper (I) chloride does not precipitate immediately, but I figured that adding a large amount of water precipitates it out.
The white precipitate should be immediately filterrd and dried. I dried mine by putting it in an oven for 15-20 mins at 210 Celsius. It dries to
produce a brownnish green mass, that can be easily broken. After breaking it up became an almost white, yellowish-green powder. I had no problem
dissolving it in ammonia though.
It certainly isn'tthe purest, but the most simple way to make it. Also, copper (I) chloride made by the above method should not be stored form mor
than 3 weeks. It completely decomposes to copper (II) chloride even when stored in an air tight container.
Hope this helps!
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Melgar
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You can make copper (ii) chloride from HCl, H2O2, and copper with little if any chlorine gas formation, as long as you only add a very small amount of
H2O2 initially. Cl2 can combine with Cl- to form Cl3-, much like iodine and bromine can. As the metal chloride level increases, the solution is able
to dissolve much more chlorine this way, and you can add higher levels of peroxide accordingly.
Interestingly, the solution gradually takes on a mild, slightly pleasant smell, which are not words that are normally associated with compounds
composed of chlorine, oxygen, and hydrogen. My only theory is that copper and nickel (which does the same thing) can catalyze the formation of
chloric or perchloric acid, which imparts the smell. In that case, those species would have to be removed or decomposed in order to not have them
oxidize your copper (I) salt.
The first step in the process of learning something is admitting that you don't know it already.
I'm givin' the spam shields max power at full warp, but they just dinna have the power! We're gonna have to evacuate to new forum software!
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