amaurer
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Dehydrating a hydrate?
Do all hydrates dehydrate by simply heating them?
And what determines the temperature required for dehydration - it is always just above 212F or do some hydrates hold onto their water more strongly
than others?
As you might guess based on my recent posts, I'm most interested in sodium bisulphate and the dehydration thereof.
Thanks for your patience with me... 
[Edited on 4-8-2012 by amaurer]
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Vargouille
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Not all. Some decompose, like aluminum nitrate nonahydrate.
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amaurer
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Using that as an example, wikipedia says it decomposes at 275F... so staying under that temperature won't dehydrate it?
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weiming1998
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Quote: Originally posted by amaurer  | Do all hydrates dehydrate by simply heating them?
And what determines the temperature required for dehydration - it is always just above 212F or do some hydrates hold onto their water more strongly
than others?
As you might guess based on my recent posts, I'm most interested in sodium bisulphate and the dehydration thereof.
Thanks for your patience with me...
[Edited on 4-8-2012 by amaurer] |
Not all hydrates dehydrate just by heating. Some (like AlCl3) decomposes completely to HCl and Al(OH)3. Whether the hydrate of a compound decomposes
on heating is due to the volatility of the acid and the Lewis acidity of the metal ion. AlCl3 is the salt of HCl, and HCl is very volatile (gas at
room T)
The Al3+ ion is very Lewis acidic, and thus this compound decomposes instead of dehydrates. But Al2(SO4)3 can be dehydrated by heating because H2SO4
is much less volatile.
Not all compounds dehydrate at 100 degrees celsius. Copper sulfate, for example, will only lose its water of hydration completely at 200 degrees
celsius.
Anyway, you should be able to dehydrate sodium bisulfate via heating. Further heating even eliminates a molecule of water from two molecules of sodium
bisulfate, forming sodium pyrosulfate (Na2S2O7). But don't heat it past 500 degrees celsius, or it will start emitting SO3.
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amaurer
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Do we know what temperature range NaHSO4 would require to dehydrate? I am drying it with a hair dryer and I *think* its working... but I'm surprised
a hair dryer works.
Here's what I'm doing... I make a NaHSO4 solution with 2.0g of what I believe is anhydrous NaHSO4 (I'm buying it in the form of a pool chemical so
they . Then I evaporate it and dry the NaHSO4 with a hair dryer and weigh it - I end up right back at 2.0g.
If I leave the stuff alone in the open air, in about an hour it will gain ~0.55g. and look moist. But the hydrate is only 15% heavier than the
anhydrous... so I'd have expected only 0.30g gain.
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woelen
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The extra weight is not due to formation of a nice well-defined hydrate, but due to absorption of extra moisture. Compounds which absorb extra
moisture from air are called 'hygroscopic'. Many salts are hygroscopic. Some are even so hygroscopic that they absorb so much water that they dissolve
in it and change into a liquid. A nice example is NaOH, another example is Fe(NO3)3.9H2O. Both of these liquefy on exposure to air, especially in
somewhat more humid climates.
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Hexavalent
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This property (the ability of a chemical to absorb enough water to form an aqueous solution) is called deliquescence.
"Success is going from failure to failure without loss of enthusiasm." Winston Churchill
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amaurer
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I am familiar with deliquescence... and have noticed that my NaHSO4 does seem to behave that way. Funny thing is that I can't find a single reference
that confirms its deliquescent - they all say its merely hygroscopic.
The only hit I'm getting is from this older text, which says the anhydrous form is NOT deliquescent but the monohydrate is. Strange distinction to
make.
link
That reference also says that the hydrate is decomposed by the water it takes up... ??
[Edited on 4-8-2012 by amaurer]
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MR AZIDE
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Copper (II) nitrate, Iron(III) chloride, Cobalt(II) Chloride are some very deliquescent salts. FeCl3 I have, has absorbed so much water its formed a
saturated solution, and then the rest of the FeCl3 cannot further dissolve and has formed a mass of crystals at the bottom of the container....
According to one of my Old inorganic text books, Lithium Chloride, is stated as one of the most deliquescent salts known.
Heating Copper(II) sulphate pentahydrate, to 100'C removes 4 molecules of water, to leave the salt pale blue. Heating above 100'C is needed to fully
dehydrate it, and the resulting salt is now off white / greyish. You could use it as a nice dessicant in a dessicator, which as it absorbs the water
again will act as its own indicator when it needs to be ' recharged'
The opposite property is called efflorescence ( which is much less common), where a salt can lose some of its water or crystallization. Sodium
carbonate can do this, forming a lower hydrate, starting from the decahydrate.
Zinc chloride cannot be dehydrated by heating, I think its decomposed into ZnOHCl, with liberation of HCl when heated. Further heating decomposes it
to ZnO.
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jamit
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| Quote: Originally posted by amaurer | I am familiar with deliquescence... and have noticed that my NaHSO4 does seem to behave that way. Funny thing is that I can't find a single reference
that confirms its deliquescent - they all say its merely hygroscopic.
The only hit I'm getting is from this older text, which says the anhydrous form is NOT deliquescent but the monohydrate is. Strange distinction to
make.
link
That reference also says that the hydrate is decomposed by the water it takes up... ??
[Edited on 4-8-2012 by amaurer] |
The difference is that deliquescent chemicals are an extreme form of hygroscopic substances. Sodium hydroxide and copper II nitrate are two examples
of a deliquescent chemical. If you leave them out in open air when its humid, it will absorb enough water to dissolve itself.
Hygroscopic chemicals also absorb water molecules from the air but doesn't dissolve itself. Examples are calcium sulfate and silica gel.
Sodium bisulfate sold as a pH lowering chemical is not deliquescent nor is it hygroscopic, unless its the anhydrous form.
[Edited on 5-8-2012 by jamit]
[Edited on 5-8-2012 by jamit]
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amaurer
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| Quote: Originally posted by jamit |
Sodium bisulfate sold as a pH lowering chemical is not deliquescent nor is it hygroscopic, unless its the anhydrous form.
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That conflicts strongly with what I am observing.
Both the anhydrous form and the monohydrate gain mass over time if left in the open air. And moreover they can collect so much water as to become a
gel-like substance.
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jamit
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| Quote: Originally posted by amaurer | | Quote: Originally posted by jamit |
Sodium bisulfate sold as a pH lowering chemical is not deliquescent nor is it hygroscopic, unless its the anhydrous form.
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That conflicts strongly with what I am observing.
Both the anhydrous form and the monohydrate gain mass over time if left in the open air. And moreover they can collect so much water as to become a
gel-like substance. |
You're right bisulfate is hygroscopic.
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