[curiouscat]

Suppose I try to precipitate out Sodium Sulfate by evaporating out the water from its solution (say at 70°C) will the crystals so formed be the anhydrous form (Na2SO4) or the hydrated form (Na2SO4.10H2O)?

The melting point of the decahydrate i.e. Glaubers salt is listed as 32°C. Does this mean that if I crystallize at any Temp. above 32°C I will get the anhydrous crystals?

[Arkcon]

Generally yes, you will always get the hydrated salt from water solution.  Whih one you get, if more than one is possible, is sometime hard to be sure.

[curiouscat]

Generally yes, you will always get the hydrated salt from water solution.  Whih one you get, if more than one is possible, is sometime hard to be sure.

Hmm but what does the 32°C melting point even mean then? So if you crystallize out of a saturated solution at 70°C technically you are above the MP of the hydrated form, right?

So it ought to immediately melt & give you anhydrous + saturated solution. Or at least as soon as you remove from contact with the solution?

[Borek]

The melting point of the decahydrate i.e. Glaubers salt is listed as 32°C.

While it is listed as a melting point, it is more of a dissolution in the substance own hydration water.

Does this mean that if I crystallize at any Temp. above 32°C I will get the anhydrous crystals?

As Arckon said, this can be hard to predict without experimenting.

There exist another stable hydrate, Na₂SO₄·7H₂O.

[curiouscat]

The melting point of the decahydrate i.e. Glaubers salt is listed as 32°C.

While it is listed as a melting point, it is more of a dissolution in the substance own hydration water.

Just to clarify my understanding: If you took 1000 gm of Glaubers salt (Na2SO4.10H2O) at 20°C & started heating it to above 32°C then you'd end up with 164 gm of crystals of anhydrous Na2SO4 + 838 gm of saturated solution (with 33 % w/w Na2SO4)

Solubility at 32°C of Na2SO4 is approx 50 gm / 100 gm water. i.e. 33% w/w

There exist another stable hydrate, Na2SO4·7H2O.

Yes, but I think it is only a metastable form. At least at temps. above 20 C.

[unsu]

That is correct, at temperatures above 32 °C it crystallizes as anhydrous sodium sulfate. The decahydrate melts incongruently in its own water of hydration at 32 °C giving the anhydrous crystals and the saturated solution.

[curiouscat]

That is correct, at temperatures above 32 °C it crystallizes as anhydrous sodium sulfate. The decahydrate melts incongruently in its own water of hydration at 32 °C giving the anhydrous crystals and the saturated solution (when it is at equilibrium).

Thanks for the info.

[unsu]

I am not sure if the simple heating would work (if the system can reach the phase equilibrium). You can also try to prepare a hot saturated solution of sodium sulfate, and when crystallize the anhydrous salt. Make sure that the temperature is always above 32 °C, even during filtration, and then dry the crystals in vacuum.

[curiouscat]

I am not sure if the simple heating would work (if the system can reach the phase equilibrium). You can also try to prepare a hot saturated solution of sodium sulfate, and when crystallize the anhydrous salt. Make sure that the temperature is always above 32 °C, even during filtration, and then dry the crystals in vacuum.

Well from what I can find out it (empirical observations) seems you always end up with the hydrated salt during crystallization. Even if you operate above 32 C.

And then you must dry the salt to anhydrous form if you need. And that's the bit that confuses me. I assumed that if you crystallize above 32 C you would have gotten anhydrous directly. Apparently not?

[Arkcon]

The only way that would work is if the trick specified by unsu: works:  you collect crystals that form from a hot saturated water solution.  Otherwise, as the crystals cool, they're simply going to "steal" water from solution and other crystals, and at least coat the dehydrated crystals with some hydrated salt.

[unsu]

The anhydrous crystals will be absorbing moisture from the air anyway, because it is not a stable form at ambient conditions. If you need to use it as a drying agent, you can heat it in vacuum to remove the crystallized water.

[curiouscat]

The anhydrous crystals will be absorbing moisture from the air anyway, because it is not a stable form at ambient conditions. If you need to use it as a drying agent, you can heat it in vacuum to remove the crystallized water.

No this is merely an effluent disposal application. But the anhydrous form leads to less landfill costs.

The other very confusing part in  this solubility literature is how solubilities are reported. It seems we all agree that the crystals will be hydrated  even above 32C so long as they are crystallized in contact with mother liquor.

Now see the solubility curve in the Fig. below. This seems a standard curve reported without controversy in many sources. Reading the curve, solubility at approx. 70°C is roughly 45 gm of Na2SO4 in 100 gm water.

Let's say I take 100 gm pure H2O at 70°C & slowly start adding anhydrous salt to it. Technically I should be able to add 45 gm & then I will see crystals. Say I added 50 gm anhydr. Na2SO4 (MW=142) we are 5 gm above the solubility limit so I expect 5 gm crystals & 105 mother liquor which ought to be a saturated soln. of Na2SO4 in H2O (31% w/w) .

On a dry basis. But if these are indeed Na2SO4.10H2O (MW=322) crystals, I'd be seeing 11.3 gms of decahydrate crystals. So would I be left with mother liquor that is 48% solution? i.e. technically above the saturation solubility at that temperature?  Or does it mean I'd be getting less than 5 gm of dry crystals out?

Perhaps this seems pedantic, but in reporting stoichiometry an accurate way of communication seems essential. I'm only wondering if this field has a standard convention of how these things are reported & I'm not getting it?

Besides a difference of only a few percent might mean tons of extra solids at scale. Finally I could always do the whole experiments myself but Na2SO4 seems such a commonly studied crystallization that it seems a waste to ignore the prior work & re-invent the wheel.

[unsu]

No you can't have Na₂SO₄·10H₂O crystals in water at 70 °C because this compound (the decahydrate) decomposes at temperatures above 32 °C to Na₂SO₄(anhydrous), it can't exist at 70 °C. The crystals will not be hydrated even if they are in contact with water. When you cool it down below 32 °C, it grabs the water molecules from the solution and becomes the decahydrate, because below 32 °C the decahydrate is the stable form.

If you find the Temperature-compositon Na₂SO₄/H2O phase diagram, you can see that the phase of the decahydrate salt lies below 32 °C. It undergoes the phase change to anh. Na₂SO₄ above that temperature and becomes the anhydrous salt. Also, there is no true melting point for Na₂SO₄·10H₂O, it melts incongruently (with decomposition)

(If you just want to prepare the anhydrous salt, and you have some solid Na₂SO₄ contaminated with the decahydrate, it can be simply heated in vacuum to remove the crystallized water. Sulfates of alkali metals melt at T above 1000 °C without decomposition, they are very stable.)