JefferyH
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Washing an organic solvent of potassium bromide with water?
I have an unknown organic solvent that is not miscible with water. Dissolved in 1L of it are 600g of potassium bromide.
How effective will it be if I wash this organic solvent with more than enough water to dissolve all of the KBr? KBr's solubility is greater in water,
but does this mean it also has a higher affinity for the water, and would rather be in the water than the organic solvent?
I'm not exactly sure what makes potassium salts soluble in organic solvents, as opposed to sodium salts.... any tips for the best way to rid this
organic solvent of the salt, assuming distillation isn't an option?
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The Volatile Chemist
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Quote: Originally posted by JefferyH  | I have an unknown organic solvent that is not miscible with water. Dissolved in 1L of it are 600g of potassium bromide.
How effective will it be if I wash this organic solvent with more than enough water to dissolve all of the KBr? KBr's solubility is greater in water,
but does this mean it also has a higher affinity for the water, and would rather be in the water than the organic solvent?
I'm not exactly sure what makes potassium salts soluble in organic solvents, as opposed to sodium salts.... any tips for the best way to rid this
organic solvent of the salt, assuming distillation isn't an option? |
I was thinking about a similar problem recently. First off, I don't think you could ever get all of it out by washing. What is your said solvent? How
about it's freezing point? I assume you've already investigated the idea of crystallizing the KBr out of solution, but I think it might work if the
freezing point's low enough.
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JefferyH
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| Quote: Originally posted by The Volatile Chemist | | Quote: Originally posted by JefferyH | I have an unknown organic solvent that is not miscible with water. Dissolved in 1L of it are 600g of potassium bromide.
How effective will it be if I wash this organic solvent with more than enough water to dissolve all of the KBr? KBr's solubility is greater in water,
but does this mean it also has a higher affinity for the water, and would rather be in the water than the organic solvent?
I'm not exactly sure what makes potassium salts soluble in organic solvents, as opposed to sodium salts.... any tips for the best way to rid this
organic solvent of the salt, assuming distillation isn't an option? |
I was thinking about a similar problem recently. First off, I don't think you could ever get all of it out by washing. What is your said solvent? How
about it's freezing point? I assume you've already investigated the idea of crystallizing the KBr out of solution, but I think it might work if the
freezing point's low enough. |
I'm not sure on solvent, and do not have any adequate ways to cool the temperature below the temperature of ice, (no dry ice).
I do not need to get all of the KBr out, though I am looking for a way to estimate how much will leave the organics per wash. Not sure this can be
done without knowing the properties of why KBr is soluble in some organics. For all I know it may have such a strong affinity for water that it may
all come out on the first wash, I am not sure. Not at the lab right now, and won't be for a few days, so I have some time to think on this.
What I was planning to do if I can't find any definitive answers is to wash with water a few times and see how much of a weight increase the water
gains, and try to calculate what % of KBr is coming out each time. Though, I'd like to find a way to better understand its properties before doing
this.
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The Volatile Chemist
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| Quote: Originally posted by JefferyH |
I'm not sure on solvent, and do not have any adequate ways to cool the temperature below the temperature of ice, (no dry ice).
I do not need to get all of the KBr out, though I am looking for a way to estimate how much will leave the organics per wash. Not sure this can be
done without knowing the properties of why KBr is soluble in some organics. For all I know it may have such a strong affinity for water that it may
all come out on the first wash, I am not sure. Not at the lab right now, and won't be for a few days, so I have some time to think on this.
What I was planning to do if I can't find any definitive answers is to wash with water a few times and see how much of a weight increase the water
gains, and try to calculate what % of KBr is coming out each time. Though, I'd like to find a way to better understand its properties before doing
this. |
Huh. I doubt there'll be any mathematical way to figure this out. It's a pretty variable situation. I recommend a first wash of 4 times the solvent's
volume or more, shaken, then drawn out. This should get most of it out.
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JefferyH
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Quite a lot of water....! But yes, this does sound like a viable idea.
If I can't get it all out, it won't be the worst scenario. The end fate of this mixture will be to precipitate the salt out into water and then
evaporate the water. Though, I guess the dual-solubility of potassium salts can work in my favor this way, of being very easy to wash out if the other
compound is only soluble in one solvent.
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blogfast25
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The so-called partition coefficient may help you understand your problem a bit better:
http://en.wikipedia.org/wiki/Partition_coefficient
To remove all KBr from the organic phase, simply wash repeatedly with water, each time separating the 'leachate' from the organic phase. After a
sufficient number of washes the organic phase will be free of KBr.
Industrially such a phase transfer could be done with so-called continuous counter current extraction: a continuous flow of organic phase (say moving
up) is in constant contact with a continuous flow of water (say moving down) in a column. If long enough the water phase will have absorbed all the
KBr from the organic phase when it flows out of the bottom of the column.
[Edited on 17-5-2014 by blogfast25]
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BromicAcid
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The issue is that your organic solvent has a very similar solubility for potassium bromide as water. The affinity for the potassium bromide then is
similar so your partitions will not be very efficient. Remember, several small washes do much better than a few large washes. Despite that fact that
the affinity is similar, if you keep using fresh water you will bring the system to equilibrium and keep removing KBr. But it will take plenty of
water.
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unionised
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Does anyone know what the solvent might be?
The reason I ask is that if it's that good a solvent for KI it must be fairly polar, and if it's that polar it might be rather more soluble in water
than you expect.
You might need to combine the water extracts and wash them with another solvent- perhaps toluene- to remove the last of the original solvent.
Also, it may be easier to oxidise the bromide to bromine, and then extract that with sodium hydroxide solution then reduce that extract to get sodium
bromide.
On the other hand, if the solvent is something hydrolysable, it might be easier to remove the solvent by degrading it.
Where did this mystery solution come from?
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JefferyH
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If I understood why potassium salts are soluble in some organics this would be a lot easier. If I could figure out what the organics are, then perhaps
reacting them in some way might change their properties and make the KBr immediately fall out. I know KBr isn't soluble in all organic solvents, but
there's something about it that makes it soluble in some.
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JefferyH
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Since the KBr is fairly soluble in this solvent is it reasonable to rule out the solvent as being an aromatic molecule? Since they are stabilized and
less polar than other organics?
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deltaH
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I would love to know what solvent not miscible with water can dissolve potassium bromide at the level of 600g/l?!
You might be able to precipitate the salt by adding in a non-polar solvent. The pressence of this other solvent may be tolerable for you or not.
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The Volatile Chemist
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| Quote: Originally posted by deltaH | I would love to know what solvent not miscible with water can dissolve potassium bromide at the level of 600g/l?!
You might be able to precipitate the salt by adding in a non-polar solvent. The pressence of this other solvent may be tolerable for you or not.
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Look up common solvents on Wikipedia, and read through their solubility and miscibility.
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HgDinis25
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Am I missing something, or distilling the solvent to precipitate the bromide is a bad idea?
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JefferyH
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Distillation would probably work but I would rather not distill. I am interested in learning whether or not there is a certain washing technique to
effectively wash something like this in these circumstances.
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HgDinis25
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How about distilling a few mL of the solvent and then run some tests to find out what it is?
You could determine the boiling point, melting point, density, etc.
After knowing what solvent you're working with it will be much easier to come up with a wash technique.
[Edited on 18-5-2014 by HgDinis25]
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unionised
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| Quote: Originally posted by The Volatile Chemist | | Quote: Originally posted by deltaH | I would love to know what solvent not miscible with water can dissolve potassium bromide at the level of 600g/l?!
You might be able to precipitate the salt by adding in a non-polar solvent. The pressence of this other solvent may be tolerable for you or not.
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Look up common solvents on Wikipedia, and read through their solubility and miscibility. |
I have been playing with chemicals for many years: I carry a lot of that information in my head- so I don't need to look it up and I can't think of
any solvent that would dissolve that muchKI, but not mix with water.
Of course, there are solvents I have never heard of.
Is there a list on WIKi of "solvents Unionised has never heard of" which I can look at?
If not, how do I go about searching for them?
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blogfast25
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You keep mentioning KI but the OP's post is about KBr. Granted that the problem is much the same for either. I didn't see the '600 g per liter' bit...
[Edited on 18-5-2014 by blogfast25]
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JefferyH
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| Quote: Originally posted by blogfast25 |
You keep mentioning KI but the OP's post is about KBr. Granted that the problem is much the same for either. I didn't see the '600 g per liter' bit...
[Edited on 18-5-2014 by blogfast25] |
From what I've gathered KI and KBr have very different solubilities, but why, I am not exactly sure. Both Ionic. The potassium cation seems to play
some role in solubilizing in organic solvents, but Iodide seems to counter that, from what I've seen, ie, KBr has greater organic solubility than KI
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blogfast25
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But you still have no idea as to what the mystery solvent could be?
I'm thinking (very tentatively) maybe a halogenated alkane?
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deltaH
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| Quote: Originally posted by unionised | | Quote: Originally posted by The Volatile Chemist | | Quote: Originally posted by deltaH | I would love to know what solvent not miscible with water can dissolve potassium bromide at the level of 600g/l?!
You might be able to precipitate the salt by adding in a non-polar solvent. The pressence of this other solvent may be tolerable for you or not.
|
Look up common solvents on Wikipedia, and read through their solubility and miscibility. |
I have been playing with chemicals for many years: I carry a lot of that information in my head- so I don't need to look it up and I can't think of
any solvent that would dissolve that muchKI, but not mix with water.
Of course, there are solvents I have never heard of.
Is there a list on WIKi of "solvents Unionised has never heard of" which I can look at?
If not, how do I go about searching for them? |
DITTO!
In my experience, as for simple inorganic salts and not counting special cases, glycerine generally tends to be pretty darn good, however, even
glycerine saturates at a 'mere' 25% KBr on a mass basis and is off course fully water miscible. Even neat 18-crown-6 ether can only complex 31% KBr on
a mass basis, though I have no idea if the resulting complex is an ionic liquid, nor miscible with water?
That said, I do know that many ionic liquids are immiscible with water, however, can these manage to dissolve 60% KBr? Anyhow, if it is an ionic
liquid, it cannot be distilled because ionic liquids generally have negligible vapour pressures AFAIR, so that would explain that bit at least 
Could this be an ionic liquid?
Anyhow, whatever it is, it's affinity for KBr is gigantic, so it equilibrium with water is not likely to be particularly favourable, hence many
washings would be required.
But that said, I get the feeling that you're going to have to 'bribe' it with another ionic to get it to give up it's KBr in fewer wash steps.
[Edited on 18-5-2014 by deltaH]
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S.C. Wack
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Nonexistent, obviously.
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HgDinis25
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Don't be too quick assuming something doesn't exist. Believe me, there may be solvents you couldn't even dream of that may do what the OP described.
I still think that a few testes to the solvent should be conducted.
PS: KBr solubiliy in water is about 678g/L, acording to wiki. Hope that's not what we're talking about here.
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JefferyH
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| Quote: Originally posted by HgDinis25 |
Don't be too quick assuming something doesn't exist. Believe me, there may be solvents you couldn't even dream of that may do what the OP described.
I still think that a few testes to the solvent should be conducted.
PS: KBr solubiliy in water is about 678g/L, acording to wiki. Hope that's not what we're talking about here. |
Perhaps using solvent was the wrong word. The compound probably isn't a solvent in the strictest sense, but a liquid oily compound that has in it
dissolved salts. From previous experiments the boiling point is 180+. I do not think it is a typical low boiling point solvent, or even one like DMSO
or DMF. We'll probably try again at distilling this Monday with a stronger vacuum and see what the result is.
[Edited on 19-5-2014 by JefferyH]
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deltaH
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This is sounding a lot like an ionic liquid or deep eutectic or something like that. If your distillation fails at high vaccuum, I think you would
have pretty strong evidence along those lines.
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