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Author: Subject: Copper sulfate has become "chalky white" after prolonged air exposure
RogueRose
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[*] posted on 10-12-2018 at 03:39
Copper sulfate has become "chalky white" after prolonged air exposure


I had some really nice CuSO4x5H2O crystals that I allowed to sit out for about 2-3 months and now they have turned a chalky white. They look more like gypsum that had some CuSO4 mixed in with it but I'm certain they are pure CuSO4.

I don't have any tests of weight before and after (when deep blue) and I haven't tried dissolving these in water to see if they dissolve more easily than the pentahydrate.

The crystals have been in room temp with mildly dry air, so I would doubt that the humidity has been less than ~20% or so and temp about 68-72.
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[*] posted on 10-12-2018 at 04:21


Now you know that hydration gives CuSO4 its blueness...

You'll likely find some blue within the crystals.

[Edited on 10-12-2018 by hissingnoise]
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[*] posted on 10-12-2018 at 06:27


now you know why they suggest to apply a coating of nail varnish on the crystals if you are planning on storing them




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[*] posted on 10-12-2018 at 07:05


The odd thing is that the penta hydrate has a decomp point at about 230F. I didn't think it was possible to loose the water below this temp even if the air is dry. Will the air pull the water from the crystal even at ~70F? That seems very odd from my experience.
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[*] posted on 10-12-2018 at 07:24


It will indeed. It's happened to my crystals, too. It's an equilibrium thing; it doesn't need to be hot to lose water, just be in an environment that has less water in it than the crystal does. It's slow but it does happen, as you observed.
I did an awesome experiment in one of my videos a while back: you can also suck the water out of the crystal by immersing it in concentrated sulfuric acid. It almost immediately starts turning white, and after some time it disintegrates into powder. I waited a month or two before poking at it, but it might happen significantly faster. I just never messed with it because the crystal never lost its shape, despite being reduced to powdery consistency!
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[*] posted on 10-12-2018 at 07:43


The extraction of water from crystals indeed happens. Slowly by dry air, quickly with strong dehydrating agents.
I have done the same experiment as MrHomeScientist. With CuSO4.5H2O you see that the blue crystals quickly become much lighter when added to concentrated H2SO4.

Even more impressive is adding solid CuCl2.2H2O (which has a nice bright cyan color) to conc. H2SO4. As soon as the crystals are added, they become brown like chocolate. This is due to formation of anhydrous CuCl2, which is brown (yellow/brown like mustard, when ground to a fine powder).




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[*] posted on 10-12-2018 at 08:23


Thanks, I'll have to try that one!
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[*] posted on 10-12-2018 at 09:07


Quote: Originally posted by RogueRose  
The odd thing is that the penta hydrate has a decomp point at about 230F. I didn't think it was possible to loose the water below this temp even if the air is dry. Will the air pull the water from the crystal even at ~70F? That seems very odd from my experience.


The decomposition point of a hydrate is like the boiling point of a liquid- it will happen very quickly at that temperature, but it will also happen slowly at much lower temperatures.

At any temperature, you will have the equilibrium reaction:
CuSO4*5H2O(s) = CuSO4*H2O(s) + 4 H2O(g)

At the decomposition point, the eq'm constant is 1. At lower temperatures, you will have a smaller equilibrium constant, but if the partial pressure of water is lower than the vapour pressure of the hydrate, it will dry out.




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[*] posted on 10-12-2018 at 09:53


Apperently the vapour pressure of the pentahydrate is 7.3 mmHg at 25oC (Reddit), how that translates to a humidity in which it is stable I don't know, but it is about 5.5x lower than water at that temperature. Would that mean a humidity of 100/5.5 = 18% at 25 degrees? (wild guess)

Edit: I remember a classmate wondering why his copper sulfate turned white and wouldn't dissolve in the 0.05 molar sulfuric acid he "prepared"... His volumetric flask was very heavy... He didn't dilute the acid but just poured it from the bottle 96% sulfuric acid.

[Edited on 10-12-2018 by Tsjerk]
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[*] posted on 10-12-2018 at 11:11


Will storing hydrated crystals in a humid environment (eg living in a humid area/put a small cup of water beside the crystal in the same airtight container) keep the crystals hydrated?



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[*] posted on 10-12-2018 at 11:23


Well I live in Florida and mine dry out, so "choking humidity" doesn't seem to be humid enough.
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[*] posted on 10-12-2018 at 13:57


Quote: Originally posted by Tsjerk  
Would that mean a humidity of 100/5.5 = 18% at 25 degrees? (wild guess)

[Edited on 10-12-2018 by Tsjerk]

Yes.
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[*] posted on 10-12-2018 at 13:59


I guess why it is so odd is b/c I had a large container of CuSO4 crystals (about 7-8lbs) that had recrystalized in to one large block of pentahydrate and it sat in the same house for about 2 years, open to the air. After it sat out for that long, I recrystalized them using distilled water and a very fine filter (and activated carbon), and I ended up with these crystals, of which I left about 200g out in the air. This time they turned chalky white.

Now I do think the previous CuSO4 that sat out for 2+ years and stayed the same dark/bright blue, there might have been a little H2SO4 on the surface of them where the crystals that turned white still had about 200g CuSO4 liquid left after the crystals formed, so these were more pure and didn't have any H2SO4 (even in small amounts) left on them. I wonder if that could have been the difference.

On another note, I did find that my FeSO4 (heptahydrate??) has also turned a chalky color like a very light/whitish green w/ a little yellow. These were left out at the same time pretty close to each other. When I did this before with FeSO4 I ended up with brownish/orange coating on all the beautiful green crystals.

I'm wondering if my air isn't much dryer this year even though both sat open over last winter when it was dry as well.

[Edited on 12-10-2018 by RogueRose]
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[*] posted on 10-12-2018 at 14:01


Quote: Originally posted by RogueRose  
I'm certain they are pure CuSO4.


I'm pretty certain they will be more or less the monohydrate.
Removing the last mole of water is much more difficult than the first 4.
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[*] posted on 10-12-2018 at 14:16


So here are some pics of the copper and iron sulfates that have air dried. I squeezed both between my fingers and both crumbled like chalk and they weigh significantly less, like 1/2 the weight I would guess, especially for the iron sulfate

The red circle is a crystal I squeezed with my fingers and it shows that it has dried completely through the whole thing. It has turned completely white. The iron sulfate looks kind of like popcorn!

CuSO4 dried.jpg - 158kB

FeSO4 air dried 2.jpg - 296kB

[Edited on 12-10-2018 by RogueRose]
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[*] posted on 10-12-2018 at 14:46


At 25C the partial pressure of water in air at 100% humidity is 30mmHg. Going off the vapor pressure of CuSo4(H2O)5 at 7.5mmHg that would lead me to believe that dehydration to the monohydrate wouldn't occur at relative humidities greater than 25%. The dehydration clearly takes place at higher humidties than 25%, I wonder what accounts for the discrepancy?

[Edited on 10-12-2018 by walruslover69]
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[*] posted on 11-12-2018 at 14:44


The dehydration treatment of the salt hydrate may have created some mesoporous structures (see, for some background, https://epub.ub.uni-muenchen.de/22683/1/oa_22683.pdf and also relating to copper catalyst https://lib.dr.iastate.edu/cgi/viewcontent.cgi?referer=https...).

The structures loaded with copper ions and exposed to air, water vapor, dust particles rich in other metals, with residual acid are not likely inert in my speculation.

So, in time, they literally breakdown.

[Edited on 11-12-2018 by AJKOER]
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[*] posted on 11-2-2019 at 18:27


exposed to Sun?
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[*] posted on 12-2-2019 at 04:47


Quote: Originally posted by pneumatician  
exposed to Sun?


if you tried to read the posts you would understand how the sun has nothing to do in this case;)





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[*] posted on 12-2-2019 at 05:48


I've had the opposite happen. I have some anhydrous CuSO4 in a sealed bottle which (over about 10 years) has turned blue.
It's a glass bottle with a plastic top so I suppose the plastic is pourous to water vapour. At the moment we are in the middle of a severe dry season and I'm measuring the humidity at 60% (26 °C).
The decahydrate can be dehydrated very easily using methanol. The decahydrate inintially dissolves in the methanol (to give a blue solution) then over about an hour a white solid precipitates out and can be filtered off.
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[*] posted on 12-2-2019 at 06:06


I'm sure copper sulfate has an equilibrium hydration level that it wants to be at relative to its environment. So the pentahydrate loses water and the anhydrous gains it until a state somewhere in the middle is reached.

That's very interesting about using methanol to dry it. I'll have to try that!
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[*] posted on 12-2-2019 at 06:18


Wondering how hydrates can dry below their composition point is like being surprised that water can evaporate from a cup at room temperature.

Apparently America had a very cold spell recently as did the UK. When its cold outside even at 100% humidity there is not much water in that cold air.

When that outside air gets inside and warmed the humidity drops. That inside air has the same water contents as the outside air but its humidity is less and it can dry damp cloths and some hydrated salts too.








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[*] posted on 12-2-2019 at 09:24


Why would any decent chemists think that CuSO4 would be decomposed by sunlight?



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[*] posted on 12-2-2019 at 15:29


Quote: Originally posted by fusso  
Why would any decent chemists think that CuSO4 would be decomposed by sunlight?

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