Solubility of MgSO47H2O in water

[JT Smith]

TL;DR

What is the solubility of epsom salt in water?

It seems like a simple enough question: what is the solubility of epsom salt in water at 20°C? A graph or table showing how it varies with temperature would be a bonus. But upon searching the internet I have been unable to determine this with confidence. Wikipedia gives the value of 113g/100ml. But other sources disagree and I can't find a definitive source for the information. I even asked chatgpt but it couldn't be sure either. I thought, naively, that this would be easy to look up without enrolling in university in order to gain access to a library with scientific journals. Will I simply have to measure it myself?

 

[renormalize]

I thought, naively, that this would be easy to look up without enrolling in university in order to gain access to a library with scientific journals.

You don't have to enroll in order to access the journals in an academic library. Simply visit your nearest university and ask library staff for assistance.

 

[willyengland]

A quick search point me immediately to:

https://www.sigmaaldrich.com/DE/de/.../solubility-table-compounds-water-temperature

https://www.researchgate.net/figure...bility-at-various-temperatures_fig4_325459625

 

[JT Smith]

You don't have to enroll in order to access the journals in an academic library. Simply visit your nearest university and ask library staff for assistance.

Thank you, I'll have to try that. I used to be able to go to one of the libraries at Stanford but then one time they turned me away. That was a long time ago. And there are other universities.

 

[JT Smith]

A quick search point me immediately to:

https://www.sigmaaldrich.com/DE/de/.../solubility-table-compounds-water-temperature

https://www.researchgate.net/figure...bility-at-various-temperatures_fig4_325459625

I had already seen the first one. I found a number of sources but they don't all agree. The second one, the article from researchgate, says it's 3 moles per kg of water at 20°C. A mole of epsom salt is 246g so that would mean about about 74g per 100g of water. The first source says it's 35.6g/100g.

When it dissolves it adds more water -- am I supposed to include that in what g/100g means? That is, if the number is, say 35.6g/100g does that mean I can only dissolve 35.6g of epsom salt in 100g of water or is it more than that because adding the salt also adds water?

I really should have taken more than just that one chemistry class in college...

 

[renormalize]

Thank you, I'll have to try that. I used to be able to go to one of the libraries at Stanford but then one time they turned me away. That was a long time ago. And there are other universities.

Try UC Berkeley. Public university libraries generally have a mission to serve the interests of the public.

 

[Chestermiller]

https://www.sigmaaldrich.com/US/en/...LxWEVYh6G9RL0gi9uEg95IcpF7DQmoPcuCs4Wdy45Kaj6

 

[JT Smith]

That's the same page as was posted previously. It says 35.6g per 100g of water. Last night I dissolved over 70g of epsom salt in 100g of water at roughly 20°C before I got tired of stirring.

What am I doing wrong?

 

[Chestermiller]

That's the same page as was posted previously. It says 35.6g per 100g of water. Last night I dissolved over 70g of epsom salt in 100g of water at roughly 20°C before I got tired of stirring.

What am I doing wrong?

Maybe they counted the water of hydration as part of the total water.

 

[JT Smith]

That's what I was wondering.

35.6g of epsom salt includes 18.2g of water. So if I started with 100g of water I should be able to dissolve proportionally more, 43.5g. But I was able to dissolve even more than that.

Suppose they don't really mean heptahydrate in that table. What if it's actually the value for the anhydrous form, just MgSO4? I read that a saturated solution at 20°C has a density of 1.36g/ml. That's consistent with the 35g/100ml value that Wikipedia gives for the anhydrous form. That would in turn imply that one could dissolve 73g of epsom salt in 100ml of total solution. I keep coming across the value 71g/100ml, which is close. Perhaps a coincidence? Wikipedia says it's 113g/100ml for the heptahydrate form which could be an error.


l will just measure it. At this point I believe it's the most expedient way to get the information I want. I'm still curious though...

 

[Mayhem]

Hydration will only aid water dissolution (for example, anhydrous CrCl₃ is practically insoluble in neutral water but its hexahydrate can allow 2 mol/L solutions)
Unless otherwise stated (or apparent from context), assume solubility data is reported for the anhydrous salt.

 

[JT Smith]

Hydration will only aid water dissolution (for example, anhydrous CrCl₃ is practically insoluble in neutral water but its hexahydrate can allow 2 mol/L solutions)

So one can't assume, as I did, that the bound water simply adds to the total solution as if it were poured in separately? My chemistry ignorance again...

Unless otherwise stated (or apparent from context), assume solubility data is reported for the anhydrous salt.

It was stated in every case I looked up. Should I also assume that even when explicitly listed otherwise that it is still data for the anhydrous form?

 

[Mayhem]

The water content does contribute to the final amount of water.

But when you prepare a 1000 ml solution, you first dissolve the salt in a small amount of water and then adjust the total volume to 1000 ml. This accounts for any solvent contribution from the salt.

If you are doing w/w% calculations you simply account for it in your calculations.

 

[JT Smith]

If you are doing w/w% calculations you simply account for it in your calculations.

What does that mean? If I find that I am able to dissolve a maximum of, say, 70g of epsom salt in an initial 100g of water what is the solubility in g/100g? Is it 70g/100g? Or since the epsom salt adds 36g of water is the solubility 70g/136g = 51g/100g? Or do I just consider the 29g of MgSO4 part of the epsom salt so that the solubility is 29g/136 = 21g/100g?

 

[Mayhem]

What does that mean? If I find that I am able to dissolve a maximum of, say, 70g of epsom salt in an initial 100g of water what is the solubility in g/100g? Is it 70g/100g? Or since the epsom salt adds 36g of water is the solubility 70g/136g = 51g/100g? Or do I just consider the 29g of MgSO4 part of the epsom salt so that the solubility is 29g/136 = 21g/100g?

Since your solvent and your cocrystallized solvent are the same, you can just choose what you want to do.

When specifying w/w%, writing the exact formulation of your solute is mandatory. If you specify a solution of MgSO₄ * 7 H₂O and actually use this salt to create your solution, then do not account for the water molecules in your solvent mass. If you specify just MgSO₄, but actually use the heptahydrate to prepare your solution, then you need to count the water molecules as they are part of the solvent.

All this confusion is avoided by using molarity!

 

[JT Smith]

I think it's possible that all of the sources I found were correct in their own way. Maybe.

If one assumes that the Millipore Sigma value of 35.6g/100g refers to MgSO4 then that is equivalent to 0.29 moles/100g. The ResearchGate graph appears to cross 20°C just shy of 3.0 moles/1kg. So that fits.

35.6g of MgSO4 is what 73g of epsom salt contains. I found several references to 71g/100ml. If the density is really 1.36g/ml then 35.6g/100g of MgSO4 in water is approximately 36g/100ml. Hence 71g/100g of epsom would be 71g/100ml.

If Wikipedia meant 100ml of *initial* water instead of the total solution then the value of 113g/100ml fits as well.


I haven't finished my crude measurement yet as I'm iteratively adjusting the mixture twice a day. One thing that occurred to me is that epsom salt may be hygroscopic enough that I'll end up with something significantly less than the textbook number.

 

[James Demers]

In reply to JT Smith:
Most of the time, a chemist will be making up a solution having a desired concentration. This involves weighing out the proper mass of solute, dissolving it, and in the final step, diluting to reach the appropriate volume. Any water of hydration is automatically taken into account in that final step. Concentrations, and solubilities in g/l, are conventionally referred to the anhydrous salt, e.g., there's a single value recorded for ferrous chloride that applies whether you're working with the dihydrate, tetrahydrate, or hexahydrate.
Experiments to measure solubility are pretty rare in modern practice. An accurate value can be obtained by preparing a saturated solution (excess, undissolved solute is present) at the desired temperature, accurately measuring out a volume, and evaporating it to complete dryness. Weighing the residual solid gives you the desired data.

 

[JT Smith]

Thank you for that post.

I was thinking of drying some of the epsom salt since I suspect it picks up moisture readily. Then it occurred to me to do what you're saying. It makes far more sense than what I've been doing. But while I think it's safe enough to heat MgSO4 in my kitchen oven to a little over 100°C for a couple of hours I'm not absolutely sure. Ignorance demands a cautious approach.

Is MgSO4 volatile to any significant degree?

 

[Borek]

Drying hydrates is (almost) never simple, some of these salts have water molecules that can't be removed without destroying the compound itself, some will lose final water molecules at much higher temperatures than 100 °C. To make things worse actual number of water molecules in the sample can be non stoichiometric and depends on the temperature, humidity and the sample history (hydration can be a slow process, so you can be limited by kinetics). There is no simple universal approach to keep the composition of the hydrate exact (then there are some hydrates that are particularly stable, these are often candidates for primary substances in analytical chemistry).

 

[James Demers]

MgSO4 is not volatile at all. The anhydrous salt is a white powder that is quite hygroscopic - chemists routinely use it to remove water from organic solvents. Epsom salts tend to LOSE water at the drop of a hat. Best strategy is to dry it in a pre-weighed glass dish, cover it while it cools, and immediately re-weigh. The challenge is that you need to get it up to almost 500°F to completely remove the water. (Put it under the broiler if you're using a typical household oven.)
https://pubchem.ncbi.nlm.nih.gov/source/hsdb/664#section=Density&fullscreen=true

 

[JT Smith]

Very interesting.

What I really want is a solution of consistent concentration that is as close to maximum concentration as reasonable given a minimum temperature of say 15°C. If I make saturated solution with undissolved epsom salt on the bottom the concentration will vary with temperature. I want a solution that doesn't change when it's warmer or colder.

I thought I could just look up the solubility for a given temperature and then use a bit less than that and be good to go. So then I'd have a simple formula: X grams of epsom salt in W grams of water. But it seems that the better approach is to make a saturated solution (solution + crystals) at a slightly lower temperature than what I expect to be the minimum and then pour off the liquid to get what I really want.

 

[Mayhem]

Very interesting.

What I really want is a solution of consistent concentration that is as close to maximum concentration as reasonable given a minimum temperature of say 15°C. If I make saturated solution with undissolved epsom salt on the bottom the concentration will vary with temperature. I want a solution that doesn't change when it's warmer or colder.

I thought I could just look up the solubility for a given temperature and then use a bit less than that and be good to go. So then I'd have a simple formula: X grams of epsom salt in W grams of water. But it seems that the better approach is to make a saturated solution (solution + crystals) at a slightly lower temperature than what I expect to be the minimum and then pour off the liquid to get what I really want.

Saturate at 40 C, cool to r.t. under lid. Add water until dissolved or simply just filter and add 5% more water.

 

[JT Smith]

Saturate at 40 C, cool to r.t. under lid. Add water until dissolved or simply just filter and add 5% more water.

That's a reasonable approach although I think to be sure it won't saturate at 15°C I'd need to add a bit more than 5%.

I still think the easiest recipe is simply add X grams of salt to Y grams of water. Then there's no need to heat, wait for cooling, and decant. It's just weigh, mix, screw on the cap, and give it a little time. But if the moisture content of epsom salt is variable then perhaps that won't work well.

I'm using this stuff for ceramic glazes. It causes the clay in the glaze to flocculate which increases the viscosity of the glaze and helps keep the materials in suspension longer. Most people just make a saturated solution and determine by experimentation how much is needed. Then when making the same glaze later they repeat the experiment, adding a little at a time: "salt to taste". My thought was to keep a record of what I've added to make that simpler and more repeatable. But then I'd need a reliable concentration since a saturated solution changes significantly with temperature, something like 30% between 10 and 30°C. So off on this side track I go...

 

[JT Smith]

I still haven't figured this out entirely.

I put too much epsom salt in some water and kept it in a water bath at 20°C for about half an hour. Then I took out 12ml and dehydrated that in our kitchen oven at 215°F for five hours. The dried portion was way more than predicted by the data on the Millipore Sigma website. I had 15.111g of solution and 5.257g of dried material. That works out to 43.8g/100g of water versus 35.6g/100g (Millipore Sigma). I know there are sources of experimental error but this is way beyond that. I'm not sure what's wrong.

I also dried some of the epsom salt, hoping to get a sense for how much moisture it might pick up from the atmosphere. It reduced by 19% in weight but also changed in appearance significantly, from a translucent crystal to an opaque powdery white substance. So I'm not sure how to interpret that.

I then turned to making solutions that would be crystal free above a certain temperature. But the behavior of these solutions baffles me. For example, I had a solution with crystals at 21.4°C that I decanted the liquid off of. But when I later put this decanted, presumably saturated, solution in the refrigerator (~4°C) no crystallization occurred.

I'm a failure as a chemist.

 

[Borek]

I also dried some of the epsom salt, hoping to get a sense for how much moisture it might pick up from the atmosphere. It reduced by 19% in weight but also changed in appearance significantly, from a translucent crystal to an opaque powdery white substance. So I'm not sure how to interpret that.

Translucent crystal vs opaque powdery is a typical difference between hydrated and anhydrous (doesn't mean it always works this way, and I would never assume something is really anhydrous just because it is white and powdery, but that's a good first guess when you see such an effect).

I then turned to making solutions that would be crystal free above a certain temperature. But the behavior of these solutions baffles me. For example, I had a solution with crystals at 21.4°C that I decanted the liquid off of. But when I later put this decanted, presumably saturated, solution in the refrigerator (~4°C) no crystallization occurred.

There is a chance the solution is supersaturated, try to drop in a small crystal and see what happens.

And don't panic. Yes, basics seem deceptively simple, but there are tons of fine prints everywhere. Basically whole chemistry is like that ;)

 

[James Demers]

dehydrated that in our kitchen oven at 215°F for five hours.

As I noted above, you need to heat to about 500°F to remove all of the water of hydration.

 

[JT Smith]

There is a chance the solution is supersaturated, try to drop in a small crystal and see what happens.

I had added some objects (little pieces of wire) to hopefully act as points of crystallization. In the end what I believed to be a saturated solution did finally crystallize at refrigerator temperature. It just took a while. I had this idea that it would be very quick but I guess that's not how it works. There's a significant time lag.

 

[JT Smith]

As I noted above, you need to heat to about 500°F to remove all of the water of hydration.

I mistakenly thought that only referred to dehydrating the solid. It's also true when epsom salt is dissolved in water?

What happens as it dehydrates at 215°F? What form of salt crystallizes?

I'll try again with my oven set to 550°F and see how it goes.

 

[Borek]

Note: there is always series of hydrates, differing with the number of water molecules, and getting to each next step requires higher temperatures.

Sometimes they are more stable and the thermogravimetric curve (basically you heat the sample very slowly and you register mass of the sample left at a given temperature) looks more like a staircase:

Still, these with lower amount of water typically look more and more whitish and opaque.

 

[JT Smith]

Finally coming back to this...

I dehydrated three different solutions in my oven set to 550°F: two were saturated solutions at 6.6 and 19.6°C; the other was made by adding 90g of epsom salt to 100g of water. I plotted these along with the three data from the Millipore Sigma website:

The 90g/100g solution was arrived at by trial and error. It was the maximum concentration that did not form crystals in my refrigerator. But the theoretical MgSO4 concentration for this solution would be supersaturated. When I dehydrated some of the solution it confirmed that concentration. So why no crystals even when seeded with a crystal and agitated?

Well it turns out I gave up too soon. Days left alone it didn't form crystals. And when seeded nothing appeared to happen. But 24 hours later the crystal seed finally grew just enough to notice. I guess that level of supersaturation wasn't enough to result in the dramatic crystal growth I was imagining would happen.

I also dehydrated a sample of epsom salt crystals and the dried portion matched the theoretical percentage of MgSO4. I had thought that epsom salt would be hygroscopic and would have some free water in addition to the bound water but apparently not. I also left some epsom salt exposed for 24 hours (25°C 65% rh) and there was no measurable change in the weight of the sample (<0.1%).

So it appears that epsom salt is pretty stable in terms of its water content, at least in my environmental conditions. So I can simply measure out a tad less than 90g of epsom salt per 100g of warm water, shake it well, and be confident that I have a solution that won't crystalize above 10°C.


That was a lot of effort for something I had originally assumed I could look up in 30 seconds. The information was actually out there, just not as easily discerned as I assumed it would be.