JoeyJoystick
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Behaviour of various hydrous states of a compound.
Hello all,
A small introduction where I am coming from with the actual question.
I wanted to show my son how to make some crystals. I figured I would do this with Copper Sulfate. So the first step was making the Copper Sulfate
(Could have bought it, but would have been less fun). So here I got my precipitate. And I dried it. I figured I should dry it more, so I moved it into
an oven for some time. Not too hot. I then noticed that the was a drastic colour change from a deep blue to an almost turquoise lightish blue. I knew
I'd done something wrong. Went to check and I think the following happened. The is a few forms of Copper Sulfate.
CuSO4 (Anhydrous)
CuSO4*H2O (Monohydrate)
CuSO4*3H2O (Trihydrate)
CuSO4*5H2O (Pentahydrate)
There may be more, but this is where my knowledge stops.
I think when I dried it in the over it converted from the pentahydrate to the trihydrate. Since I never measured the temperature of the oven I can
never be sure, but I had put it on the lowest temperature and it was not hot to get seriously hurt. I strongly believe that the temp was somewhere
between 70 and 100C. And as Wikipedia shows at these temperature it looses 2 water molecules.
I decide to proceed anyway. I grow some starter crystals and used these to make a few larger crystals. And then I got this beautiful deep blue colour
all back again.
So my question is the following. How do the different hydrates behave when in solution? I am not a chemist. I can well imagine that once in solution
it is simply a bunch of water with a bunch of Copper Sulfate. But due to my lack of knowledge I can also imagine that it would be quite different. Do
the various hydrates only apply to there dry states? Or does this also impact a solution made of it?
In line with the assumption that the lighter CuSO4 was indeed not the pentahydrate but the trihydrate, and the Crystals produced are indeed the
pentahydrate, can I put these crystals back in the oven and dry them to make a trihydrate crystal, or any other hydrate for that matter?
And does you answer only apply to CuSO4 or is this general information that applies to all hydrates from all compounds?
Mentioned it a few times. Not a chemist here as you can see from the above writing. My apologies if I have used the wrong terminology. Please correct
me if you have this uncontrollable urge. It will help me put the questions down better a next time. Also, the above may not be very practical, but it
is just about curiosity and gaining more knowledge.
hmmm, not exactly 1 question... Sorry.
Thank you so much and hope to hear from you,
Joey
[Edited on 27-4-2020 by JoeyJoystick]
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DraconicAcid
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Hydrates only exist in the solid form- once you dissolve them in water, you have aqueous copper(II) sulphate, regardless of what form you added to the
water. Crystallizing a solution of copper(II) sulphate will give the pentahydrate under most conditions.
Heating the pentahydrate will cause it to decompose to the trihydrate, and heating it more will give the monohydrate. I believe it's difficult to get
the monohydrate without accidentally overheating it and getting the oxide, but I could be wrong.
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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JoeyJoystick
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Quote: Originally posted by DraconicAcid  | Hydrates only exist in the solid form- once you dissolve them in water, you have aqueous copper(II) sulphate, regardless of what form you added to the
water. Crystallizing a solution of copper(II) sulphate will give the pentahydrate under most conditions.
Heating the pentahydrate will cause it to decompose to the trihydrate, and heating it more will give the monohydrate. I believe it's difficult to get
the monohydrate without accidentally overheating it and getting the oxide, but I could be wrong. |
Thanks for clearing that up. And am I correct when I say that the various hydrates of a compound may have or can have very different properties?
Joey
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ScienceBum
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Hi Joey,
You would be correct in stating that hydrates can have very different properties! I asked about this a few years ago:
http://www.sciencemadness.org/talk/viewthread.php?tid=33772#...
I'll be talking about ligands and crystal field theory in an upcoming post, but the short answer is what DraconicAcid said - with a salt like copper
(II) sulfate when it's dissolve in water it'll behave the same.
In my personal experience, growing the pentahydrate crystals takes a lot longer (like weeks depending on how much you're growing and how slowly the
water evaporates) but you can get some pretty huge ones. I've tried heating up the solution to drive off the water and no matter how slowly I go I
always get some level of decomposition into the trihydrate form.
A fun experiment might be trying to grow crystals at different levels of heat evaporation and comparing size and differences in color. Let me know if
you guys wind up doing that ^^
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JoeyJoystick
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| Quote: Originally posted by ScienceBum | Hi Joey,
You would be correct in stating that hydrates can have very different properties! I asked about this a few years ago:
http://www.sciencemadness.org/talk/viewthread.php?tid=33772#...
I'll be talking about ligands and crystal field theory in an upcoming post, but the short answer is what DraconicAcid said - with a salt like copper
(II) sulfate when it's dissolve in water it'll behave the same.
In my personal experience, growing the pentahydrate crystals takes a lot longer (like weeks depending on how much you're growing and how slowly the
water evaporates) but you can get some pretty huge ones. I've tried heating up the solution to drive off the water and no matter how slowly I go I
always get some level of decomposition into the trihydrate form.
A fun experiment might be trying to grow crystals at different levels of heat evaporation and comparing size and differences in color. Let me know if
you guys wind up doing that ^^ |
I should have searched better... Thanks for the tip. I would like to try that, but I do not have possession over an oven with which I can control the
temperature well enough. Besides, it would be the cooking oven. And the misses will kill me if I occupy that for days in a row. lol.
Anyways, it's been growing for 5 days now. They are not the prettiest crystals, but it serves its purpose for now. It is here in Thailand obviously
pretty warm most of the time. But evaporation is still going relatively slow. A lot slower than what I remember from when I was young back in Holland.
I guess it's because the rainy season has just started and the humidity is quite high. I think his makes the evaporation relatively slow.
If I manage to pull it off at different temperatures, I will surely let you know.
Joey
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Bedlasky
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Copper sulfate (no matter what hydrate) dissolves in water to form [Cu(H2O)6]2+ and SO42- ions.
[Cu(H2O)6]2+ are blue. Reality is little bit more complicated because at higher concentrations of CuSO4
sulphates displaced some waters in this complex.
Copper sulfate crystalize as pentahydrate - four waters are directly bonded at copper atom while fifth water is bonded by nonbonding interactions
(that means that this water is bonded due to physical forces, not by chemical bond).
Copper sulfate is really good starting point for making crystals. I made two crystals and one of them is really big. But it crystalize slowly, it
takes months of growing. When you stop growing, don't forget varnish crystal with colourless nail polish. If you don't do this, crystal slowly
dehydrates and it will crumbling.
Alums are also very good for growing crystals. I should recommend to you ferric ammonium sulfate dodecahydrate. It can be made from common chemicals
and forms beautiful violet octahedral crystals. This compound have high solubility (which is uncommon between alums) and grow really fast. Be careful
about changing temperatures, it have steep curve of solubility. I had really nice, big and nearly reagular crystal of this compound in solution.
During one day temperature rose about a few degrees and this crystal all dissolved. One crystal I store in mineral oil, because it dehydrates at the
air. Maybe I try nail polish as protective layer on some small sample. But sometimes nail polish isn't enough good protection (for example
FeSO4 crystals oxidized even under few layers of nail polish).
Look on this website, there is lot of useful info about growing crystals.
| Quote: Originally posted by DraconicAcid |
Heating the pentahydrate will cause it to decompose to the trihydrate, and heating it more will give the monohydrate. I believe it's difficult to get
the monohydrate without accidentally overheating it and getting the oxide, but I could be wrong. |
Why it should be problem? Copper sulfate can be easily dehydrated to anhydrous salt without decomposition. It is used as a test for the presence of
water and as a dehydrating agent.
https://en.wikipedia.org/wiki/Copper(II)_sulfate
Wiki says that trihydrate is formed at 63°C, monohydrate at 109°C and anhydrous salt at 200°C. Anhydrous CuSO4 decompose at 650°C to
CuO and SO3.
[Edited on 27-4-2020 by Bedlasky]
[Edited on 27-4-2020 by Bedlasky]
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DraconicAcid
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| Quote: Originally posted by Bedlasky |
Why it should be problem? Copper sulfate can be easily dehydrated to anhydrous salt without decomposition. It is used as a test for the presence of
water and as a dehydrating agent.
https://en.wikipedia.org/wiki/Copper(II)_sulfate
Wiki says that trihydrate is formed at 63°C, monohydrate at 109°C and anhydrous salt at 200°C. Anhydrous CuSO4 decompose at 650°C to
CuO and SO3.
[Edited on 27-4-2020 by Bedlasky]
That was probably my mistaken impression based on watching students doing it with a Bunsen burner, in which precise temperature control is impossible.
[Edited on 27-4-2020 by Bedlasky] |
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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JoeyJoystick
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I would like to thank you all for clarifying. And the links provide even more info.
I saw mention of putting a filter on top of the jar/beaker into which you grow the crystal(s) in order to avoid dust falling and causing a wild growth
of small crystals. I think this is a great idea. I am assuming that due to the very low evaporation that it will still breath through the filter. Can
someone confirm this? Because I really do not like to assume.
I am now thinking of a small low temperature chamber inside which I can maintain the temperature. Low temperatures though. No more than 100C. An oven
to pull this off is not an option in Covid-19 times with limited financial resources. Damn this virus. But I can make a chamber like this myself I
hink. Small wooden box with a lightbulb inside which is controlled by a PID Temp controller. Glass pane in the front for luring inside at times.
Shouldn't cost too much I recon. Put a filter over the jar to avoid dust and keep a second flask next to it with additional solution to fill up the
beaker/flask in which the crystals grow. A small tube can be pressurised to transfer additional liquid. Can blow by mouth if chemicals are not
dangerous. An additional flask with water to keep fumes away from my mouth with some chemicals and if it is too dangerous I will think of something
else. lol. Sounds like a fun exercise. This way there is no interaction whatsoever other than a slight disturbance when topping up the solution.
Needless to say, i will be back and show the results when done. I will take a while, because ordering this from China now is the cheapest but far from
the quickest solution. The PID controller that is. Checked this while ago for something else, and they go for less than 20USD.
Again, Thank you all for the info.
| Quote: Originally posted by DraconicAcid | | Quote: Originally posted by Bedlasky |
Why it should be problem? Copper sulfate can be easily dehydrated to anhydrous salt without decomposition. It is used as a test for the presence of
water and as a dehydrating agent.
https://en.wikipedia.org/wiki/Copper(II)_sulfate
Wiki says that trihydrate is formed at 63°C, monohydrate at 109°C and anhydrous salt at 200°C. Anhydrous CuSO4 decompose at 650°C to
CuO and SO3.
[Edited on 27-4-2020 by Bedlasky]
That was probably my mistaken impression based on watching students doing it with a Bunsen burner, in which precise temperature control is impossible.
[Edited on 27-4-2020 by Bedlasky] |
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So you were both right. lol. It is easy, provided you have the correct equipment. I am sure that most professional labs have this equipment, but I am
also convinced that few people have these capabilities in a home lab... It can still be done of course, but is a different ball game compared to
growing a crystal on a string in a jar.
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Bedlasky
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| Quote: Originally posted by JoeyJoystick | | It is easy, provided you have the correct equipment. I am sure that most professional labs have this equipment, but I am also convinced that few
people have these capabilities in a home lab... It can still be done of course, but is a different ball game compared to growing a crystal on a string
in a jar. |
For making anhydrous CuSO4 ordinary oven is just fine. Or gently heating in test tube above candle.
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