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Author: Subject: CuCl2 workup
njl
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[*] posted on 3-12-2020 at 13:34
CuCl2 workup


I recently made some calcium chloride via the following procedure (that I made up as I went along based on a basic understanding of inorganic chem):

Oxidizing solution:

1. A 1 L beaker was charged with a lot of calcium hypochlorite. I used In the swim brand pool shock.

2. Approximately 350 mL of room temperature (~15 C) tap water was poured into the beaker.

3. The beaker was brought just to a boil, swirled, and then taken off heat to cool back down to room temperature. The solution was stirred during both heating and cooling in an attempt to bring as much hypochlorite into solution as possible.

Acid solution:

1. In a separate container 300 mL of concentrated HCl was added to 200 mL of room temperature tap water. The solution was agitated until no more swirls of HCl were present (you know what I mean).

2. The solution was set aside until the oxidizing solution cooled to room temperature.

Copper:

1. Rectangular pieces of copper roughly the same height as the 1 L beaker were hastily wiped clean of any dirt. The copper strips were then folded and compressed so that they would easily fit into the 1 L beaker. The strips were set aside until the oxidizing mixture cooled to room temperature.

Reaction:

1. The above materials were brought outside into a well ventilated area and placed atop a flat brick surface.

2. The copper was then placed into the oxidizing solution, with about half of the metal submerged. No reaction was noted except a small amount of effervescence.

3. Using a 1 and 10 mL syringe, acid solution was added to the oxidizing solution and copper. At first an attempt was made to drip the acid directly onto the exposed metal surface so that the introduction of HCl to the oxidizing solution would be slowed. Upon dissolution of the acid into the oxidizing solution, a large amount of bubbling began and died down within a few seconds. With the bubbling came green chlorine gas that largely dissipated into the air. No significant exotherm was noted until the very end of the acid addition.

4. The acid solution is added over the course of a few minutes. The 1 L beaker is swirled periodically to expose fresh copper to the oxidizing solution, but the copper strips prevented effective mixing. The chlorine gas generated during the course of the reaction did oxidize some of the exposed metal surface before blowing away. This led to some quite pretty copper chloride forming above the oxidizing mixture on the metal surface. I thoroughly enjoyed squirting acid solution onto this copper chloride, as it would dissolve with continuous addition and expose fresh metal.

5. Once the acid solution was added, the beaker was swirled again. The copper strips were then used to agitate the solution. At this point a reasonable exotherm was noted, although the beaker was never too hot to hold. The solution was allowed to cool outside until no more bubbling occurred. Once the beaker had cooled, some solids had precipitated out.

6. The now room-temperature solution was put onto a hotplate and a stirbar dropped in. Medium stirring was turned on followed by heat. The copper strips once again impeded effective stirring but there was enough agitation to suspend some precipitated solids in solution.

7. The solution was again brought to a boil, except this time heating was stretched over about an hour. With heating the solution cleared up and took on a stronger turquoise color. Once the solution just began to boil, heat was turned down to maintain that temperature and the solution left stirring for a few hours.

8. Once the solution was transparent but still strongly colored with minimal solids suspended, heating was turned off and the solution was left to stir overnight.

9. After ~20 hours of stirring, the copper strips were removed from the solution, rinsed with room temperature tap water, and set aside on paper towels to dry. The solution measured only slightly acidic to pH papers. Some solids formed a layer at the bottom of the beaker. The solution was then covered with plastic wrap and brought inside for workup.

Notes:

Oxidizing solution:

1. I don't know how much Ca(OCl)2 was added.
2. RT for me is probably colder than average.
3. I am aware of the dangers of this procedure. Bringing a concentrated hypochlorite solution to a light boil is probably a bad idea. I'm not saying anyone should repeat this exactly, but it worked for me.

Acid solution:

1. I don't know off the top of my head how to describe the HCl swirls but as far as I know it's normal. I think it happens because the refractive index of tap water is different from that of the HCl solution. I used a disposable HDPE bottle and threw it out afterwards, although I'm sure you could use anything resistant to HCl.

Copper:

1. I got my copper from some siding I had laying around. Looking back I should have checked for solder before starting.

Reaction:

All I have to say about the reaction itself is that it is obviously very dangerous. I am aware of this. I did this outside on a cold, windy day. I was only ever bothered by fumes when the beaker was heated for the 2nd time.

I am currently waiting to start workup. All in all, the solution consists of water, Ca2+, Cu2+, and Cl-. After the copper was removed, I noticed some silvery metal (I believe it may be solder of some kind) on the edges of 2 of the strips. I didn't notice it before since it was covered in a dark oxide layer. I assume the metal was exposed after HCl dissolved away the oxide layer. These edges were not submerged in the oxidizing solution, but I may have missed some on other parts of the copper. Therefore it's reasonable to assume some unknown metal ions are present, in some unknown but likely trivial concentration.

I am stuck at this point as I don't know the best way to separate the CuCl2 from CaCl2 and other impurities. I don't have any appreciable amount of other solvents on hand, and what I do have I can't afford to use here. My current idea is to evaporate off all or a majority of the water, add a more reasonable volume to redissolve the products, concentrate the solution, and cool in a refrigerator to collect a first crop of crystals. However, CaCl2 and CuCl2 have very similar solubilities in water at most temperatures so I doubt this method of separation will work as intended.

Suggestions?
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DraconicAcid
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[*] posted on 3-12-2020 at 14:07


That's going to be a pain to separate.



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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AgCollector
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[*] posted on 3-12-2020 at 19:02


Quote: Originally posted by njl  


1. In a separate container 300 mL of concentrated HCl was added to 200 mL of room temperature tap water. The solution was agitated until no more swirls of HCl were present (you know what I mean).


The swirls you mention sound like 'schlieren'

Quote: Originally posted by njl  

I am stuck at this point as I don't know the best way to separate the CuCl2 from CaCl2 and other impurities. I don't have any appreciable amount of other solvents on hand, and what I do have I can't afford to use here. My current idea is to evaporate off all or a majority of the water, add a more reasonable volume to redissolve the products, concentrate the solution, and cool in a refrigerator to collect a first crop of crystals. However, CaCl2 and CuCl2 have very similar solubilities in water at most temperatures so I doubt this method of separation will work as intended.

Suggestions?


Perhaps precipitate the calcium as calcium sulfate?
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MidLifeChemist
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[*] posted on 3-12-2020 at 19:13


Quote: Originally posted by AgCollector  

Perhaps precipitate the calcium as calcium sulfate?


This. A little H2SO4 should do the trick.
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DraconicAcid
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[*] posted on 3-12-2020 at 21:07


Isn't CaSO4 really difficult to filter out?



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njl
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[*] posted on 4-12-2020 at 06:25


Thank you for the replies. I think the sulfate will be the way to go. I wouldn't say it's difficult, but it is pretty slow. The only sulfate source I have is some grossly impure drain cleaner acid which is probably between 80 and 90 percent pure. Based on how goopy it is I think there's probably a good amount of junk in there. Since I don't want to introduce more contamination if possible, is there a way I can titrate the solution to find the calcium ion concentration?
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