## Solubility of copper sulphate compared to its pentahydrate

SnailsAttack - 5-3-2023 at 01:32

Despite being an apparently simple topic, the solubility behavior of hydrated salts compared to their anhydrates (and associated anomalies) remain undocumented and unexplained.

The results of a previous thread I made on this topic were inconclusive and ultimately spiraled into a mess, so I'm narrowing the scope of this post to only look at copper sulphate and its pentahydrate.

Part 1 - Theory

The solubility in water of any hydrate salt should theoretically be proportional to the total dry mass of salt and the total mass of water, including the water of solvation and the water contributed by the salt's water of crystallization.

I devised the following formula to compute the solubility of any hydrate (s ₕ) from the solubility of its anhydrate (sₐ) and their molar mass ratio (Mᵣ):

s ₕ = sₐ/(Mᵣ + sₐ(Mᵣ - 1)/1000)

Where:
s ₕ is the solubility of the hydrated salt in g/L at a given temperature.
sₐ is the solubility of the anhydrous salt in g/L at a given temperature.
Mᵣ is the anhydrate/hydrate molar mass ratio.

The point of this formula is that the mass ratio between the water and salt should remain constant regardless of the salt's hydration state.

 Quote: Solved example for copper sulphate: The molar mass of anhydrous copper sulphate (CuSO₄) is 159.61 g/mol. The molar mass of copper sulphate pentahydrate (CuSO₄·5H₂O) is 249.69 g/mol. Therefore the molar mass ratio (Mᵣ) is 159.61/249.69 = 0.63923. Let's assume that the solubility of anhydrous copper sulphate (sₐ) is 200 g/L at 0°C. According to the following formula, the solubility of the pentahydrate (s ₕ) at this temperature should be: s ₕ = sₐ/(Mᵣ + sₐ(Mᵣ - 1)/1000) s ₕ = 200/(0.63923 + 200(0.63923 - 1)/1000) s ₕ = 352.7 g/L 200 grams of anhydrous copper sulphate dissolved in 1 liter, or ~1,000 grams of water is equivalent to a water:salt mass ratio of 1,000/200 = 5.00. 352.7 grams of copper sulphate pentahydrate (CuSO₄·5H₂O) consists of 225.5 grams of copper sulphate (CuSO₄) and 127.2 grams of water (H₂O). If this mass of copper sulphate pentahydrate is dissolved in 1 liter, or ~1,000 grams of water, the water:salt mass ratio is again equal to (1,000+127.2)/225.5 = 5.00.

Part 2 - Empirical data

Copper sulphate was chosen as the subject for this thread because of its relatively simple hydration behavior (as far as I know), stability at high temperatures when dehydrated, and the abundance of (highly incongruent) data on its solubility at different temperatures and hydration states, given below from various internet sources:

 Quote: Wikipedia solubility table and Sciencemadness wiki CuSO₄·5H₂O 231 g/L in water at 0°C 320 g/L in water at 20°C 1140 g/L in water at 100°C Sciencemadness wiki CuSO₄·0H₂O 143 g/L in water at 0°C 205 g/L in water at 20°C 754 g/L in water at 100°C Sciencemadness wiki (physical properties desc.) CuSO₄·5H₂O (assumed) 316 g/L in water at 0°C 2033 g/L in water at 100°C Wikipedia page on copper sulphate CuSO₄·0H₂O (assumed) 201 g/L in water at 20°C PubChem source 1 CuSO₄·5H₂O (assumed) 243 g/L in water at 0°C PubChem source 2 CuSO₄·0H₂O (assumed) 754 g/L in water at 100°C PubChem source 3 CuSO₄·0H₂O (assumed) 203 g/L in water at 20°C NPIC technical sheet CuSO₄·0H₂O (pentahydrate stated, anhydrate assumed) 148 g/L in water at 0°C 736 g/L in water at 100°C INCHEM safety data sheet CuSO₄·5H₂O (assumed) 317 g/L in water at 0°C Sigma Aldrich solubility table CuSO₄·0H₂O (pentahydrate stated, anhydrate assumed) 148 g/L in water at 0°C 208 g/L in water at 20°C 736 g/L in water at 100°C Sigma Aldrich solubility table CuSO₄·5H₂O (anhydrate stated, pentahydrate assumed) 255 g/L in water at 0°C 362 g/L in water at 20°C 830 g/L in water at 100°C Crystal growing wiki CuSO₄·0H₂O 142 g/L in water at 0°C 200 g/L in water at 20°C 770 g/L in water at 100°C Crystal growing wiki CuSO₄·5H₂O 231 g/L in water at 0°C 326 g/L in water at 20°C 1150 g/L in water at 100°C USDA data sheet CuSO₄·5H₂O 316 g/L in water at 0°C 2033 g/L in water at 100°C USDA date sheet CuSO₄·5H₂O (anhydrate stated, pentahydrate assumed) 243 g/L in water at 0°C 320 g/L in water at 20°C 1140 g/L in water at 100°C

Of these 15 sources, 6 of them fail to specify the copper sulphate's hydration state and 4 of them appear to have swapped the data for the anhydrate and pentahydrate (based on their dissonance with the other 11 sources).

The data from these sources is analyzed below.

 Quote: // list of all solubility data for the pentahydrate, and the average solubility CuSO₄·5H₂O at 0°C 231 316 243 317 255 231 316 243 // 269 ± 43 g/L CuSO₄·5H₂O at 20°C 320 362 326 320 // 332 ± 21 g/L CuSO₄·5H₂O at 100°C 1140 2033 830 1150 2033 1140 // 1388 ± 602 g/L - // list of all solubility data for the anhydrate, and the average solubility CuSO₄·0H₂O at 0°C 143 148 142 148 // 145 ± 3 g/L CuSO₄·0H₂O at 20°C 205 201 208 200 // 204 ± 4 g/L CuSO₄·0H₂O at 100°C 754 736 770 736 // 749 ± 17 g/L // computed solubility for the pentahydrate based on the avg. solubility of the anhydrate CuSO₄·5H₂O at 0°C 247 g/L // 8.9% below avg. CuSO₄·5H₂O at 20°C 361 g/L // 8.7% above avg. CuSO₄·5H₂O at 100°C 2030 g/L // 46.3% above avg. - // computed solubility for the anhydrate based on the avg. solubility of the pentahydrate CuSO₄·0H₂O at 0°C 156 g/L // 7.6% above avg. CuSO₄·0H₂O at 20°C 189 g/L // 7.9% below avg. CuSO₄·0H₂O at 100°C 591 g/L // 26.8% below avg.

The solubility data for anhydrous copper sulphate is in very strong agreeance, with an average percent deviation of only around 2% between sources. By contrast, the average percent deviation for the pentahydrate is over 20%.

The discrepancy for the pentahydrate solubility measurements at 100°C is especially strange, with an average deviation of 43% and values ranging from 2033 g/L down to 830 g/L.

However, what's interesting is that there are 2 sources that corroborate the computed solubility of 2030 g/L at 100°C for the pentahydrate almost exactly, at 2033 g/L.

Judging by the huge range in the solubility values measured for the pentahydrate (particularly at 100°C), I suspect that the researchers had trouble maintaining the copper sulphate at a stoichiometric pentahydrate.

I intend to test this, and try to come up with a reliable way to measure solubility values myself.

-

If there's errors in this post please let me know. Discussion is appreciated and I'd really like to get to the bottom of this, because it's a subject that we should've had straightened out 200 years ago.

[Edited on 3/5/2023 by SnailsAttack]

Sulaiman - 5-3-2023 at 10:55

A good protocol would be required - Lots of procedures.

Purification and analysis of reagents might be a collaborative first step?

An optional initial collaboration could be to choose one temperature (eg 50C) and members could submit their own results.
Each to their own capabilities with estimates of error.
I'd probly have a go at it..... if someone organised it

B(a)P - 5-3-2023 at 12:43

Hi SnailsAttack, good to see the continuation of this after your last thread. I agree with Sulaiman sentiments and would be happy to contribute, if you decide you would like experimental input from other members.

Rainwater - 5-3-2023 at 17:27

If you can write a step-by-step, I'm in.
I've always assumed that the solubility of a salt is equal to a molar ratio with water for a given temperature.

### Just some ideas

Sulaiman - 5-3-2023 at 19:50

Someone (if no volunteer then I'll volunteer, if before May) could provide, via post,
'reference samples' for those interested,
maybe in return for p&p and misc. costs.
This would enable a degree of cross-checking our own results.
Ideally the samples should be of high purity etc.
But shared 'reference samples' of lower quality may serve the same purpose. (?)

I'm thinking small packets (e.g. 10g?) just to be used for 'calibration'
but if postal rates are acceptable then possibly enough for all requirements ?

P.S. I may be biased because I too expect the saturated concentration of ions will be independent of whether the anhydride or pentahydrate is used.
If any other result then, I would first suspect my own procedures
but... who KNOWS ?

For the initial collaboration it may be better for each participant to measure solubility at their local room temperature
Just for practical reasons.

[Edited on 6-3-2023 by Sulaiman]

unionised - 6-3-2023 at 03:22

It's actually quite difficult to measure solubility accurately (especially far from room temperature).
How accurate an answer do you actually ever need?

If you use molarity, the hydration doesn't matter.

teodor - 6-3-2023 at 03:54

Near 95C is the crytical point after which water solution gives CuSO4*3H2O crystalls. And below this point probably it would be mixture of penta- and tri- hydrate. That's why solubility data at high temperature can disagree with each other.
Another challenge is hydrolisis to basic copper sulfate, I guess for this reason some authors can measure solubility of CuSO4 in acidic solution and interpolate the results.

If you define the solubility as an "ammount of solid which can be dissolved" it's better to use the data for anhydrous sulfate.
As I already mentioned in the previous thread, the classical definition is equillibrium, so it depends also on crystallisation properties.
Also, it should be mentioned that "amount of solid which can be dissolved" in the real life is never possible to dissolve without either raising the temperature or adding some additional water, that's why it is never used as a definition.

[Edited on 6-3-2023 by teodor]

[Edited on 6-3-2023 by teodor]

yobbo II - 6-3-2023 at 08:02

There is also the question of the meta stable state. When you are dissolving (or crystallizing) you must wait for quite a long time for ALL
of the stuff that can go into solution (or come out) to go into solution at the given temperature. Most studies use a temperature controlled set up and stirr for 24 hours.
No solvent must be let escape as well.

Then there is the problem of when you extract the crystals, ( either extra crystals of stuff that you have put there to allow saturation to occur (or crystals that have come out of solution, if you are crystallizing)) the water stuck to these crystals will have a certain weigh and will also have a certain amount of the stuff dissolved in it. How do you measure this?
The method of wet residues is used (there may be other methods) to measure this liquid.
See https://pubs.acs.org/doi/10.1021/je00103a002
What I am trying to say is that it is quite a task to measure the whole thing accurately.
The higher the measuring temperature the more the hassel. You need preheated filters, etc etc

With very soluble (not CuSO4) compounds the whole thing is a nightmare. Some substances will dissolve in their own water of crystallization.

Yob

Texium - 6-3-2023 at 08:35

As the last three posters have pointed out, solubility is not as easy to measure as one would think. I’d also argue that it’s hardly important. The only time I look up solubility data for a compound is if I’m making a saturated solution of some common salt to keep on hand, and even then I overshoot by a few grams/L to ensure that it will stay saturated with potential temperature fluctuations. So what if there’s some undissolved solid at the bottom of the bottle.

So my question is, why even bother? I won’t accept “for fun” as an answer, because frankly there is nothing fun about measuring the solubility of salts to a high degree of accuracy, and I can guarantee that there are much more fulfilling ways you could be spending your precious lab time!

SnailsAttack - 6-3-2023 at 11:29

 Quote: Originally posted by Sulaiman I'd first want to test my reagents carefully, Are they actually anhydrous, and/or exactly pentahydrate?

Good suggestion. I've found that some salts which exhibit hydration behavior behave very nicely while others are a complete mess.

 Quote: Originally posted by Sulaiman Maybe decompose and compare to copper oxide weight if >650C available, or (zinc or Iron etc.?) precipitation of copper, dry, weigh etc. Maybe Iodometric titration for greater accuracy. Something like that.

I probably can't maintain a sample at 650°C+ on the stovetop but I could convert the copper sulphate to copper oxyhydroxide and then pyrolyze that and weigh it.

 Quote: Originally posted by Sulaiman Maybe a first calculation would be what accuracy of results is required? For example, in absolute terms, I can weigh to great accuracy, but volumes are only accurate to about +/- 0.1%, on a good day and temperature accuracy is even worse at about +/- 1C.

Solubility measurements extrapolated from individual mass measurements rather than volume measurements would be ideal, yes.

According to this graph (which appears to be for the anhydrate), a difference of 1°C at 20°C affects the solubility by 3.5 g/L, equivalent to an error of 1.73%.

B(a)P - 6-3-2023 at 11:44

 Quote: Originally posted by Texium As the last three posters have pointed out, solubility is not as easy to measure as one would think. I’d also argue that it’s hardly important. The only time I look up solubility data for a compound is if I’m making a saturated solution of some common salt to keep on hand, and even then I overshoot by a few grams/L to ensure that it will stay saturated with potential temperature fluctuations. So what if there’s some undissolved solid at the bottom of the bottle. So my question is, why even bother? I won’t accept “for fun” as an answer, because frankly there is nothing fun about measuring the solubility of salts to a high degree of accuracy, and I can guarantee that there are much more fulfilling ways you could be spending your precious lab time!

The 'why even bother' is that there still seems to be an unexplained phenomenon here and just because three posters have said solubility is hard to measure is not a good reason to not undertake an experiment. I do not want to be putting words in the mouth of the OP so please correct me if I am wrong, but this is not about being precise when making up a saturated solution.

In all of the posts on this thread and the other relating to this matter, the observed (based on published solubilities) difference in solubility between anhydrous and hydrated salts (taking into account the additional water in the hydrate) has still not been explained.

SnailsAttack - 6-3-2023 at 11:46

The first step to this is to see whether copper sulphate actually crystallizes from solution to form an ideal stoichiometric pentahydrate.

CuSO₄·5H₂O(s) -> CuSO₄(s) + 5H₂O(v)

Incredibly, it does, it behaves perfectly. I pyrolyzed ~15 grams of hydrated copper sulphate and measured a mass reduction of 0.6386x compared to the theoretical of 0.6392x. That's an error of 0.10%.

The anhydrous copper sulphate seems to be very hygroscopic, and pulls water from the air on standing. I'm not sure what hydration state it tends to. Doesn't matter for the purpose of this thread.

Texium - 6-3-2023 at 12:58

 Quote: Originally posted by B(a)P The 'why even bother' is that there still seems to be an unexplained phenomenon here and just because three posters have said solubility is hard to measure is not a good reason to not undertake an experiment. I do not want to be putting words in the mouth of the OP so please correct me if I am wrong, but this is not about being precise when making up a saturated solution. In all of the posts on this thread and the other relating to this matter, the observed (based on published solubilities) difference in solubility between anhydrous and hydrated salts (taking into account the additional water in the hydrate) has still not been explained.
A discrepancy between different data sets is not an “unexplained phenomenon” it just means that the result can be affected by a multitude of variables, and controlling one may make it nearly impossible to control others. I have read the other thread, and I think teodor’s comment that solubility is a state of equilibrium is important to highlight again. We like to talk about the solubility of a given substance at a given temperature as though it is an immutable physical constant, but it is not. Some salts behave better than others, but it’s always an equilibrium, and even if you use a very consistent method, you’re bound to get slightly different results sometimes. With salts that can form multiple hydrates, have a tendency to supersaturate, and/or are deliquescent (e.g. sodium sulfate, sodium acetate, calcium chloride) calculating a meaningful value becomes even more futile, hence the severe discrepancies highlighted for those in the other thread.

To be clear, I’m not trying to discourage experimentation on this simply because it’s hard to do, but because it’s very hard to do in a way that will actually yield meaningful results. What is the best outcome from this? Another result that may agree partially with some of the existing measurements, or possibly none of them at all? I don’t think the CRC would be rushing to revise their handbooks.

Sulaiman - 6-3-2023 at 19:45

 Quote: Originally posted by Texium We like to talk about the solubility of a given substance at a given temperature as though it is an immutable physical constant, but it is not.
Really?
Other than the chemical composition and the environment of our hypothetical copper sulphate solutions,
what could cause the equilibrium to not be a constant ?
(this is precisely what the OP question implies)

Not trying to rewrite the CRC handbook,
just settle the copper sulphate anhydrous vs pentahydrate anomalous (or not) solubility question.

SnailsAttack - 7-3-2023 at 00:19

 Quote: Originally posted by Texium As the last three posters have pointed out, solubility is not as easy to measure as one would think. I’d also argue that it’s hardly important.

Solubility might be difficult to measure, but I think ballpark-accuracy solubility data is valuable to have.

 Quote: Originally posted by Texium The only time I look up solubility data for a compound is if I’m making a saturated solution of some common salt to keep on hand, and even then I overshoot by a few grams/L to ensure that it will stay saturated with potential temperature fluctuations. So what if there’s some undissolved solid at the bottom of the bottle.

Online measurements of the solubility of copper sulphate at room temperature range anywhere from 200 g/L to 320 g/L. A solution in that range of accuracy is fine for analytical chemistry, but if you knew the exact solubility you could infer the molarity of a saturated solution, which could be really useful.

 Quote: Originally posted by Texium So my question is, why even bother? I won’t accept “for fun” as an answer, because frankly there is nothing fun about measuring the solubility of salts to a high degree of accuracy, and I can guarantee that there are much more fulfilling ways you could be spending your precious lab time!

No yeah I completely agree, this is super boring. B(a)P describes my exact reason for taking interest nonetheless:

 Quote: Originally posted by B(a)P The 'why even bother' is that there still seems to be an unexplained phenomenon here and just because three posters have said solubility is hard to measure is not a good reason to not undertake an experiment. In all of the posts on this thread and the other relating to this matter, the observed (based on published solubilities) difference in solubility between anhydrous and hydrated salts (taking into account the additional water in the hydrate) has still not been explained.

Other observed phenomena include:

- Woelen's claim that vanadyl sulphate is virtually insoluble in water while its hydrate dissolves readily. I'd like to test this myself with the vanadyl you sent me.
- Various strange behavior concerning the solubility of hydrated sodium sulphates discussed by Teodor and Tsjerk.
- B(a)P noted that online solubility measurements for anhydrous sodium sulphate at 20°C range anywhere from 139 to 445 g/L.
- According to wikipedia, anhydrous sodium acetate is ~3x more soluble in water than the trihydrate, but the data in this post suggests that the hydrated salt should always be more soluble.

SnailsAttack - 7-3-2023 at 00:49

 Quote: Originally posted by Texium A discrepancy between different data sets is not an “unexplained phenomenon” it just means that the result can be affected by a multitude of variables, and controlling one may make it nearly impossible to control others.
 Quote: Originally posted by Texium I think teodor’s comment that solubility is a state of equilibrium is important to highlight again. We like to talk about the solubility of a given substance at a given temperature as though it is an immutable physical constant, but it is not.

There can't be that many variables, can there? i would think that the hydration state of emerging salts and their equilibrium with the solution would be controlled solely by the temperature.

 Quote: Originally posted by Texium With salts that can form multiple hydrates, have a tendency to supersaturate, and/or are deliquescent (e.g. sodium sulfate, sodium acetate, calcium chloride) calculating a meaningful value becomes even more futile, hence the severe discrepancies highlighted for those in the other thread.

Depending on how the solubility measurements are performed, supersaturation and deliquescence aren't necessarily a problem. The main showstopper would be the salts that evaporate down to a goo instead of crystallizing (e.g. magnesium acetate, nickel acetate), since the line between "solid" and "liquid" is indistinguishable.

 Quote: Originally posted by Texium What is the best outcome from this? Another result that may agree partially with some of the existing measurements, or possibly none of them at all? I don’t think the CRC would be rushing to revise their handbooks.

not to be conceited, but I trust our observations over the institutional ones

SnailsAttack - 7-3-2023 at 02:02

 Quote: Originally posted by Sulaiman A good protocol would be required - Lots of procedures. Purification and analysis of reagents might be a collaborative first step? An optional initial collaboration could be to choose one temperature (eg 50C) and members could submit their own results. Each to their own capabilities with estimates of error. I'd probly have a go at it..... if someone organised it
 Quote: Originally posted by B(a)P Hi SnailsAttack, good to see the continuation of this after your last thread. I agree with Sulaiman sentiments and would be happy to contribute, if you decide you would like experimental input from other members.

Glad to hear you guys are interested. I'm down for organizing a collaboration on solubility measurements once I've got a reliable procedure figured out.

 Quote: Originally posted by Sulaiman Someone (if no volunteer then I'll volunteer, if before May) could provide, via post, 'reference samples' for those interested, maybe in return for p&p and misc. costs. This would enable a degree of cross-checking our own results. Ideally the samples should be of high purity etc. But shared 'reference samples' of lower quality may serve the same purpose. (?) I'm thinking small packets (e.g. 10g?) just to be used for 'calibration' but if postal rates are acceptable then possibly enough for all requirements ?

I've only got like 20 grams of copper sulphate but I have a whole bunch of sodium acetate and magnesium sulphate I could send that I think would behave well for solubility measurements.

 Quote: Originally posted by Rainwater If you can write a step-by-step, I'm in.

Working on it.

teodor - 7-3-2023 at 03:25

 Quote: Originally posted by Sulaiman Other than the chemical composition and the environment of our hypothetical copper sulphate solutions, what could cause the equilibrium to not be a constant ?

It could not ever exist, like equilibrium between CuSO4 (anhydrous) crystalls and water under 100C or CuSO4 * 5H2O and water above 96C.

But I have my own interest to compose the table of solubility of salts of mono- and di- carboxylic acids with different metals up to C6, so I will be happy to read the experiment results and will hope it will be possible to extend your method for measuring those solubilities which are probably completely unknown.

Amos - 7-3-2023 at 04:48

This is, uhh, pretty silly. Examine solubility from the perspective of crystallizing a compound. It’s all to common to encounter supersaturation of a solution, and in general it’s an incredibly hard-to-predict process. Then examine every instance of a YouTube video where more than the theoretical amount of water or another solvent is needed to dissolve something. Then look at how incredibly slow further dissolution is as you approach the solubility limit of a given compound in a given volume of solvent. It’s all very cryptic and inconsistent. There are far too many variables at play, clearly many that we don’t even know or have words to describe, as this thread has pointed out. You’re not going to crack the case in a home lab with reagents you cannot properly purify or analyze, in a non-controlled setting for variables like temperature, humidity, etc. This isn’t a problem that needs solving and it’s not one you won’t figure it out standing at a lab bench with basic equipment.
SnailsAttack - 7-3-2023 at 14:59

please watch me nail this entire project
yobbo II - 7-3-2023 at 15:46

If there is a difference in solubility between the Cu and SO4 ion pair that comes from anhydrous Copper Sulphate and the ion pair that comes from the hexahydrate (or whatever hydrate) then the water has some way of 'distinguishing' between the different source of ion pairs.
This is impossible? (surley).
One 'set' from a particular source may dissolve more quickly that the other but if given enough time both must have equal solubility.

Copper sulphate has hydrates from 1 to 7 according to the net.

The solubility according to Seidell is below.

Perhaps this data is included in the first post.

Yob

Attachment: dehyd_OF_cu_sulfate.pdf (1.8MB)

Texium - 7-3-2023 at 17:38

 Quote: Originally posted by yobbo II One 'set' from a particular source may dissolve more quickly that the other but if given enough time both must have equal solubility.
Yes.

I was typing out a longer response last night, but I lost it by accident, and teodor and Amos pretty much covered what I was going to say.

### Help prevent my brain from breaking!

Sulaiman - 7-3-2023 at 18:55

What would happen if anhydrous copper sulphate is added to an already saturated solution?

Meanwhile....
I think that Iodometric titrations will be required to measure copper sulphate concentrations.
Or is there a better (cheaper, easier, more accurate etc.) method.

[Edited on 8-3-2023 by Sulaiman]

Texium - 7-3-2023 at 18:57

 Quote: Originally posted by SnailsAttack please watch me nail this entire project
Good luck bud
SnailsAttack - 7-3-2023 at 23:00

 Quote: Originally posted by Texium Good luck bud

 Quote: Originally posted by teodor Near 95C is the crytical point after which water solution gives CuSO4*3H2O crystalls. And below this point probably it would be mixture of penta- and tri- hydrate. That's why solubility data at high temperature can disagree with each other.

Maybe. Could be tough to maintain a 100°C bath anyways.

 Quote: Originally posted by teodor Another challenge is hydrolisis to basic copper sulfate, I guess for this reason some authors can measure solubility of CuSO4 in acidic solution and interpolate the results.

Noted.

 Quote: Originally posted by yobbo II There is also the question of the meta stable state. When you are dissolving (or crystallizing) you must wait for quite a long time for ALL of the stuff that can go into solution (or come out) to go into solution at the given temperature. Most studies use a temperature controlled set up and stirr for 24 hours.

24 hours? I give it five minutes max before a sample of solid copper sulphate emersed in water is effectively fully equilibriated with the solution.

 Quote: Originally posted by yobbo II Then there is the problem of when you extract the crystals, ( either extra crystals of stuff that you have put there to allow saturation to occur (or crystals that have come out of solution, if you are crystallizing)) the water stuck to these crystals will have a certain weigh and will also have a certain amount of the stuff dissolved in it. How do you measure this? The method of wet residues is used (there may be other methods) to measure this liquid. See https://pubs.acs.org/doi/10.1021/je00103a002 read the first few lines.

... how do you envision the measurement process to work? I guarantee it doesn't require supping water off the crystals and using a mathematic model to calculate how much is left on them.

teodor - 8-3-2023 at 02:06

SnailsAttack, I put here a link to a scan from "Palmer, Experimental physical chemistry", the chapter about experiments with solid solibilities. I hope some practical hints from this book could be useful to your experiments.

SnailsAttack - 8-3-2023 at 02:37

 Quote: Originally posted by teodor SnailsAttack, I put here a link to a scan from "Palmer, Experimental physical chemistry", the chapter about experiments with solid solibilities. I hope some practical hints from this book could be useful to your experiments. https://drive.google.com/file/d/1wzUomQlu6dqQWDhW2UloAqPizGy...

Thanks, I looked it over. The procedures for measuring solubility in part A are a little weird, but the titration method is interesting.

In part B the author suggests that only one hydrate can exist in solution at a given temperature, which is what I hoped would happen, since it simplifies things a lot.

yobbo II - 8-3-2023 at 04:24

SnailsAttack
"24 hours?
I give it five minutes max before a sample of solid copper sulphate emersed in water is effectively fully equilibriated with the solution."

How do you actually know it is equilibriated.
If you are going up the the solubility limit (at a certain temperature) it can take ages to equilibriate.
The next dude may feel he has to give it 4 minutes (or 10 mins. or hours) to be fully equilibriated.

I would imagine that with some substances (not copper suplhate, but I don't really know) it could take weeks!

Sulaiman

"
Help prevent my brain from breaking!

What would happen if anhydrous copper sulphate is added to an already saturated solution?

"
The anhydrous stuff may actually suck some water from the solutions and become a hydrate of some value. It depends on whether or not the system (at it's temperature and pressure) at that point is congruent or incongruent.

From Wiki:
congruent/Incongruent dissolution
Many substances dissolve congruently (i.e. the composition of the solid and the dissolved solute stoichiometrically match). However, some substances may dissolve incongruently, whereby the composition of the solute in solution does not match that of the solid.

Trust me.
I am no expert.
Perhaps someone here can put it a better way.
I don't really understand what is meant by '(i.e. the composition of the solid and the dissolved solute stoichiometrically match)'

A phase diagram can explain a system (assuming you can read/interpret) well.
One has been produced and may appear over in Ref's.

Yob

[Edited on 8-3-2023 by yobbo II]

teodor - 8-3-2023 at 08:09

 Quote: Originally posted by Texium I was typing out a longer response last night, but I lost it by accident, and teodor and Amos pretty much covered what I was going to say.

Oh, yes, it happend for me several times already. Now I always copy responses to an external text editor when they become long and continue there.

I think there is a nice demo illustrating your and Amos idea "there are a lot of variables some of which are unknown".

https://en.wikipedia.org/wiki/Storm_glass

Sulaiman - 8-3-2023 at 11:21

 Quote: Originally posted by yobbo II How do you actually know it is equilibriated.

You only need to know that it has equilibriated to within your experimental error limits.

A one-time experiment could be used as a reference for 'time required for a given error range'

Texium - 8-3-2023 at 11:44

Quote: Originally posted by Sulaiman
 Quote: Originally posted by yobbo II How do you actually know it is equilibriated.

You only need to know that it has equilibriated to within your experimental error limits.

A one-time experiment could be used as a reference for 'time required for a given error range'
Not sure what you mean exactly by “one-time experiment.” Generally if one wants to obtain accurate data, running experiments in triplicate is standard, especially if you’re trying to establish an error range. Plus, you’d have to control for temperature, particle size of salt that you’re dissolving, and temperature. As in, making sure that the temperature never gets above a given temperature during the course of the experiment, not just that the final temperature is the same.
teodor - 8-3-2023 at 14:08

Did somebody observe the phenomenon when on long standing some saturated solutions of salts in water deposit crystalls and this is impossible to explain by temperature just because the deposit doesn't disappear during further temperature fluctuations?
The standard method of measuring solubility includes agitation of crystalls and this is also some energy which also participates in the equilibrium even if we keep temperature constant.
The mechanism of crystallization from saturated solution goes through some step of creating pre-crystal film with different (lower) solution concentration. So, in normal equilibrium there are always 2 zones with different concentrations of solution - on the surface of undissolved salt and in the rest of solvent volume. This is a kind of barrier with very complex physics if you would look on the equations. So, in the reality there is a difference between actual concentration of solution which is in equilibrium with crystalls and concentration you try to measure (in the main volume of the solvent).
That's why we can expect some difference from solubility results depending on the method of measurement.

[Edited on 8-3-2023 by teodor]

[Edited on 8-3-2023 by teodor]

yobbo II - 8-3-2023 at 16:10

Regarding the storm glass.
The solubility of stuff will depend on pressure, temperature and the pressure of water vapour (the humidity).
The stuff in the storm glass must be very sensitive to these changes.

I think.

Rainwater - 8-3-2023 at 16:42

I want to provide an example i think is revelant to this discussion, which lead me to a lot of good studies reguarding solubility.
While back I made some silver chloride with unusual results.
A perfectly clear solution. It stumped me for a few days.
Quantitative Chemical Analysis
by: Harris, Daniel C.
ISBN: 9781429218153
If i really wanted to figure out what happened.

This book corrected my understanding of what really happens when a compound dissolves in another.
This book also contains step by step procedures and mathematics required for the experiments the OP suggest.

My AgCl Notes:
Reaction:
$$KNO_3(aq) + HCl(aq) + Ag(s)\leftrightarrow K^+ + H^+ + NO_3^- + Cl^- + Ag$$
$$3 Ag(s) + 4 HNO_3(aq) \rightarrow 3 AgNO_3(aq) + 2 H_2O + NO(g)$$
$$AgNO_3 + Cl^- \rightarrow AgCl(s) + NO_3^-$$

Lab notes
 Code:  Test tube was weighed dry. 18.48 grams 2.31 grams of dry KNO3 added to test tube About 1/3rd of di water was added to the test tube. (5mL) Light heating was applied A dry graduate cylinder was weighed 15.15 grams 2.61g of 32% hcl was weight in. 2.3ml was the volume The test tube was heated until all the solids dissolved and a clear solution was obtained T1500 2.40g of silver metal was added, no reaction observed The hcl was added. T1505 the metal begain to darken. The solution is gently boiling T1535 an additional 2.30g of KNO[sub]3[/sub] and 2.3ml of HCl was added T1555 the smell of no[sub]2[/sub] is coming from the end of the test tube T1605 scraped this crap. Found my bottle of nitric acid 3 days later, the test tube contained a clear solution, which turned white when I picked it up. 

[Edited on 9-3-2023 by Rainwater]

Texium - 8-3-2023 at 18:08

 Quote: Originally posted by teodor Did somebody observe the phenomenon when on long standing some saturated solutions of salts in water deposit crystalls and this is impossible to explain by temperature just because the deposit doesn't disappear during further temperature fluctuations? The standard method of measuring solubility includes agitation of crystalls and this is also some energy which also participates in the equilibrium even if we keep temperature constant. The mechanism of crystallization from saturated solution goes through some step of creating pre-crystal film with different (lower) solution concentration. So, in normal equilibrium there are always 2 zones with different concentrations of solution - on the surface of undissolved salt and in the rest of solvent volume. This is a kind of barrier with very complex physics if you would look on the equations. So, in the reality there is a difference between actual concentration of solution which is in equilibrium with crystalls and concentration you try to measure (in the main volume of the solvent). That's why we can expect some difference from solubility results depending on the method of measurement.
Excellent point! That’s something I’ve definitely observed but didn’t know exactly how to describe. IIRC, making a saturated ammonium chloride solution gave me a really hard time because of that issue, though the resultant crystals were very beautiful.
SnailsAttack - 9-3-2023 at 19:04

I've observed that anhydrous copper sulphate rehydrates instantly on contact with water, which raises more questions about the semantics of what the solubility of anhydrous copper sulphate actually means if it can't coexist with water.

Presumably it's based on the formula in my original post, although I can't find any sources with data that reflect it very well, except for the Crystal Growing wiki and Sigma Aldrich table, which come pretty close except for the 100°C data point and the fact that Ligma Aldrich's hydrate and anhydrate data appear to be swapped.

DraconicAcid - 9-3-2023 at 19:28

 Quote: Originally posted by SnailsAttack I've observed that anhydrous copper sulphate rehydrates instantly on contact with water, which raises more questions about the semantics of what the solubility of anhydrous copper sulphate actually means if it can't coexist with water.

It just means that it's the amount of copper(II) sulphate in a saturated solution. The degree of hydration it came in with just isn't relevant.

SnailsAttack - 9-3-2023 at 19:52

 Quote: Originally posted by DraconicAcid It just means that it's the amount of copper(II) sulphate in a saturated solution. The degree of hydration it came in with just isn't relevant.

If you can find any data online where the solubility of the hydrate and anhydrate are correlated by an actual formulaic expression, please share.

SnailsAttack - 9-3-2023 at 20:20

I came up with a method of measuring the solubility of copper sulphate (anhydrate or hydrate) using only mass measurements. The procedure is as follows:
 Quote: Step 1. Soak a ~10 gram sample of copper sulphate (anhydrate or hydrate) in ~25 mL of water for ~10 minutes, stirring occasionally. Approximately half the copper sulphate will dissolve at room temperature. Step 2. Decant most of the copper sulphate solution and record the solution's temperature and mass (m ₜ ₜ ₗ). Step 3. Evaporate the copper sulphate solution and weigh the resulting pentahydrate crystals (m ₕ). To ensure they're fully dry, break them up occasionally, taking multiple measurements over time to ensure they stop losing mass.

This method relies on my observation that copper sulphate solution evaporates to form a stoichiometric pentahydrate, and will have to be modified to work for more finicky salts.

I devised the following formula (based on the one in the original post, not backed up by empirical data) to compute the solubility of any hydrate (s ₕ) from the mass of a saturated solution (m ₜ ₜ ₗ) and the mass of dissolved hydrate (m ₕ):

s ₕ  =  m  ₕ·d ₛ /(m ₜ ₜ ₗ + m ₕ(Mᵣ - 2))

Where:
s ₕ is the solubility of the hydrated salt in g/L of solvent at a given temperature.
m ₕ is the mass of the solvated hydrated salt in grams.
d ₛ is the density of water at a given temperature.
m ₜ ₜ ₗ is the mass of the saturated solution at a given temperature.
Mᵣ is the anhydrate/hydrate molar mass ratio.

I performed this procedure and did the calculations for a sample of copper sulphate pentahydrate, and recorded the results below.

 Quote: = Copper sulphate pentahydrate solubility test Mass of solution (m ₜ ₜ ₗ): 18.579 ± 0.010 g Temperature of solution: 19.0 ± 0.5°C Mass of dissolved CuSO₄·5H₂O (m  ₕ): 4.469 ± 0.005 g The density of water at 19°C is 998 kg/L. The molar mass ratio of anhydrous to hydrous copper sulphate is 159.61/249.69 = 0.63923. s ₕ  = m  ₕ·d ₛ /(m ₜ ₜ ₗ + m  ₕ(Mᵣ - 2)) s ₕ = 4.469·998/(18.579 + 4.469(0.63923 - 2)) s ₕ = 4460/(18.579 + 4.469(-1.36077)) s ₕ = 4460/12.498 s ₕ = 356.9 ± 6.1 g/L of solvent Predicted hydrous solubility based on data: 351 ± 7 g/L (extrapolated from anhydrate data) Deviation from theoretical: 1.7%

The solubility formula and measurement procedure appear to work, but it's kind of hard to say for sure since I don't have a reliable source to reference against.

The total list of hydrate solubility equations that I've come up with are given below. At the moment they're not backed up with empirical data, except for the formula correlating m ₕ to mₐ which is basic stoichiometry.

s ₕ = sₐ/(Mᵣ + sₐ(Mᵣ - 1)/1,000)

m ₕ = mₐ/Mᵣ

sₐ  =  mₐ·d ₛ /(m ₜ ₜ ₗ - mₐ/Mᵣ)  =  m  ₕ·Mᵣ·d ₛ /(m ₜ ₜ ₗ - m  ₕ)

s ₕ  =  m ₕ·d ₛ /(m ₜ ₜ ₗ + m ₕ(Mᵣ - 2))  =  mₐ·d ₛ /(m ₜ ₜ ₗ·Mᵣ + mₐ(Mᵣ - 2))

Where:
s ₕ is the solubility of the hydrated salt in g/L of solvent at a given temperature.
sₐ is the solubility of the anhydrous salt in g/L of solvent at a given temperature.
Mᵣ is the anhydrate/hydrate molar mass ratio.
--
m  ₕ is the mass of the solvated hydrated salt in grams.
mₐ is the mass of the solvated anhydrous salt in grams.
--
d ₛ is the density of water at a given temperature.
m ₜ ₜ ₗ is the mass of the saturated solution at a given temperature.

Sulaiman - 9-3-2023 at 20:43

 Quote: Originally posted by SnailsAttack I've observed that anhydrous copper sulphate rehydrates instantly on contact with water, which raises more questions about the semantics of what the solubility of anhydrous copper sulphate actually means if it can't coexist with water.
so,,,
if anhydrous is added to a saturated solution,
it will precipitate as pentahydrate?

SnailsAttack - 9-3-2023 at 20:46

 Quote: Originally posted by yobbo II Regarding the storm glass. The solubility of stuff will depend on pressure, temperature and the pressure of water vapour (the humidity). The stuff in the storm glass must be very sensitive to these changes.

Not possible. It's a sealed, rigid container, only the external temperature is capable of affecting the solubility.

 Quote: Originally posted by Rainwater I want to provide an example i think is revelant to this discussion, which lead me to a lot of good studies reguarding solubility. While back I made some silver chloride with unusual results. A perfectly clear solution. It stumped me for a few days.

In the past I've observed some salt mixtures that should result in a metathesis reaction but inexplicably fail to produce a precipitate unless the solution is evaporated to an absurd concentration.

What do you know about how this works? I've been calling it 'quasi-stability'.

SnailsAttack - 9-3-2023 at 20:53

Quote: Originally posted by Sulaiman
 Quote: Originally posted by SnailsAttack I've observed that anhydrous copper sulphate rehydrates instantly on contact with water, which raises more questions about the semantics of what the solubility of anhydrous copper sulphate actually means if it can't coexist with water.
so,,,
if anhydrous is added to a saturated solution,
it will precipitate as pentahydrate?

Yeah, I read your original question about this and I'm pretty sure the anhydrate would pull water from solution to form the pentahydrate, but fail to dissolve due to insufficient water. I think some of the solvated copper sulphate would also crash out as pentahydrate since the anhydrate will have absorbed a portion of the water that was keeping it in solution.

It's worth noting that the hydration of anhydrous copper sulphate is very exothermic (to the point that small quantities of water actually boil on contact), so that would briefly raise the solubility.

Rainwater - 10-3-2023 at 16:20

starting on page 121
of the book i references above, some useful google terms are
Solubility product = Ksp
Disproportionation
Common ion effect
Coprecipitation
Complex formation

Simple version.
When you have "Pure" H2O + CuSO4
You have every combination of cations and anions in solution.
H+, OH-, CuOH, H2SO4, etc etc etc. Lots of combination.
But each combination is interacting in its own little area of the solution until an
equilibrium is reached.

Applied to original post, from dry to wet
Starting with anhydrous CuSO4 and adding 1 molar of water, one would think you
would end up with the monohydrate. But without the ability to distribute the water
evenly on an atomic level, a variety of hydrates will be formed.

From wet to dry, as in measuring a percipitate, separating the hydrate from pure
water is a combination of which compounds percipitate first and how much.

Now lets leave the ideal behind and now consider all the N2, O2, CO2 and everything
else found in water. The combinations start adding up, and interactions get complex

Now the context of these example is from the prospective of reducing
measurement errors to as close to zero as possible, then using a hot mess of
formulas to know how much will remain in solution, so that measurements can be
corrected and total masses calculated.

And all that very improperly sums up the first 6 chapters, 142 pages
Edit:
 Quote: Maybe. Could be tough to maintain a 100°C bath anyways.

Double boiler. Will get you close, a pressure vessel will get you perfect, but that sets the difficulty bar up to max by introducing another variable no one has mentioned yet. Pressure.
Higher pressure will decrease the entropy, empathy damnit, .... disorder ... of the system, effecting the equilibrium.

[Edited on 11-3-2023 by Rainwater]

yobbo II - 10-3-2023 at 16:54

Sounds simply enough!

There is a paper here

showing the system CuSO4 water from 0 to 100C (APPROX).
It may be of use.

The more I read the less I know.

Yob

Tsjerk - 11-3-2023 at 02:56

Quote: Originally posted by SnailsAttack
 Quote: Originally posted by DraconicAcid It just means that it's the amount of copper(II) sulphate in a saturated solution. The degree of hydration it came in with just isn't relevant.

If you can find any data online where the solubility of the hydrate and anhydrate are correlated by an actual formulaic expression, please share.

You mean a formula that tells you how soluble the anhydrate is compared to the hydrate? You either divide or multiply by the ratio between the molecular weights.

If 10 grams of the anhydrate dissolves in 100 ml solution you divide by the molecular weight of the anhydrate and multiply by the molecular weight of the hydrate (multiply by 1.56). The other way around would be division by 1.56.

I think the question whether the starting hydration of a salt influences the final solubility can be answered by the fact there are no sources stating different solubilities for the two. Just different solubilities from different sources. Doesn't that tell you solubility is just hard to measure?

 Quote: Originally posted by SnailsAttack Despite being an apparently simple topic, the solubility behavior of hydrated salts compared to their anhydrates (and associated anomalies) remain undocumented and unexplained.

The difference in behaviour is undocumented because there is no difference, the differences between sources is explained by the difficulty of measurement.

When you think about a system that contains copper sulfate both dissolved and undissolved and is in equilibrium (saturated solution) all there is is water, copper sulfate in solution and fully hydrated solid copper sulfate. How would the hydration state of the starting copper sulfate influence the final concentration?

In chemistry, when something is not there, it can't influence the state of what is there. The hydration state of the starting material is not there anymore, so it can't influence the equilibrium, can it?

[Edited on 11-3-2023 by Tsjerk]

Texium - 11-3-2023 at 13:36

 Quote: Originally posted by Tsjerk The difference in behaviour is undocumented because there is no difference, the differences between sources is explained by the difficulty of measurement.
At this point, I have given up on trying to argue reason in this thread, and I figure we may as well just sit back and watch SnailsAttack attempt to reinvent the wheel. It could be entertaining.
teodor - 11-3-2023 at 16:01

I think this is the correct table (and it is published in 1894):
https://gallica.bnf.fr/ark:/12148/bpt6k34902q/f552.item

There are 2 reasons why some sources contain incorrect data for CuSO4.
1. Solubility curve of CuSO4 has 3 anomalies in the pentahydrate range: 28.9, 34.85, 53.72C. It is like spikes where properties of crystalls are changing very fast in very narrow range (fractions of degree) and comes back to the curve:

There is a reason for this behaviour: water molecules in crystalls change behaviour from oscilation to rotation. But those anomalies of CuSO4 solubility were discovered in the end of 19 centure. Try to evaluate the quality of your in-house methods. And probably what you can observe at those points is "there is no constant equilibrium".

Solubility table in classical references are very often just interpolations, for example Bronsted (1928) gives the full table built by 4 actual measurement, the rest is calculated by generic formula.

2. After 95.9C it doesn't crystallize as pentahydrate, so those table which gives solubility at 100C can incorrectly substract 5 mols of water instead of 3.

So, for me there is no any miracle here.

[Edited on 12-3-2023 by teodor]

Sulaiman - 11-3-2023 at 18:09

Maybe not a miracle but certainly a revelation.

Two obvious questions:

1 is this common behaviour for hydrated ions/molecules?
(eg MgSO4, Cu(NO3) 2 etc)

2 can I use this effect to create a hyper-saturated copper sulphate solution ?

Texium - 11-3-2023 at 18:28

1. Probably.

2. Doubtful, because it requires the temperature to be held at an exact value with a precision of 0.1 degrees.

[Edited on 3-12-2023 by Texium]

SnailsAttack - 12-3-2023 at 08:27

 Quote: Originally posted by Tsjerk You mean a formula that tells you how soluble the anhydrate is compared to the hydrate? You either divide or multiply by the ratio between the molecular weights. If 10 grams of the anhydrate dissolves in 100 ml solution you divide by the molecular weight of the anhydrate and multiply by the molecular weight of the hydrate (multiply by 1.56). The other way around would be division by 1.56.

If you do the math, your formula doesn't make any sense. I've got it written as follows:

s ₕ = sₐ / Mᵣ

Where:
s ₕ is the solubility of the hydrated salt in g/L of solvent at a given temperature.
Mᵣ is the anhydrate/hydrate molar mass ratio.
sₐ is the solubility of the anhydrous salt in g/L of solvent at a given temperature.

Let's assume the solubility of the anhydrous salt at 20°C is 205 g/L. Therefore the solubility of the hydrate is: 205/0.63923 = 320.70 g/L.

 Quote: 320.70 grams of CuSO₄·5H₂O dissolves in 1 liter of solvent (998 grams of water). Total CuSO₄: 205.00 g Total water: 115.70 + 998 g Water/salt ratio: 5.433 or 184.1 g/L 205.00 grams of CuSO₄·0H₂O dissolves in 1 liter of solvent (998 grams of water). Total CuSO₄: 205.00 g Total water: 998 g Water/salt ratio: 4.868 or 205.0 g/L

The solubility of the anhydrate comes out 10% higher, which would suggest that it's somehow intrinsically more soluble than the hydrate, which we both agree isn't possible. The water/salt ratio is balanced in the formula I already discussed in the original post.

 Quote: Originally posted by Texium At this point, I have given up on trying to argue reason in this thread, and I figure we may as well just sit back and watch SnailsAttack attempt to reinvent the wheel. It could be entertaining.

Your guys' wheel is a friggin octagon at best!

fusso - 12-3-2023 at 08:33

Snailsattack, I think you forgot to include the extra volume when calculating solubility.
SnailsAttack - 12-3-2023 at 08:46

 Quote: Originally posted by teodor There are 2 reasons why some sources contain incorrect data for CuSO4. 1. Solubility curve of CuSO4 has 3 anomalies in the pentahydrate range: 28.9, 34.85, 53.72C. It is like spikes where properties of crystalls are changing very fast in very narrow range (fractions of degree) and comes back to the curve: [] There is a reason for this behaviour: water molecules in crystalls change behaviour from oscilation to rotation. But those anomalies of CuSO4 solubility were discovered in the end of 19 centure. Try to evaluate the quality of your in-house methods. And probably what you can observe at those points is "there is no constant equilibrium".

Why would it even do this? There's only 15 data points on that graph. Can the anomalous 3 spikes be attributed to some sort of weird error? Were the measurements repeatable? Each data point is between 1 and 3 degrees apart, so how wide are the spikes?

I've never seen a graph showing this behavior or read anything about it. It's also not reflected in the table you cited.

SnailsAttack - 12-3-2023 at 09:03

 Quote: Originally posted by fusso Snailsattack, I think you forgot to include the extra volume when calculating solubility.

What extra volume? Solubility is supposed to be measured as grams per liter of solvent rather than grams per liter of solution.

If you measure in g/L of solution you have to know the density of the solution and that's a pain in the butt. I don't think that reframing the measurements in g/L of solution fixes the problem.

teodor - 12-3-2023 at 09:11

 Quote: Originally posted by Sulaiman Maybe not a miracle but certainly a revelation. Two obvious questions: 1 is this common behaviour for hydrated ions/molecules? (eg MgSO4, Cu(NO3) 2 etc) 2 can I use this effect to create a hyper-saturated copper sulphate solution ?

1. This discovery is connected with the name of Linus Pauling. He wrote several excelent books by the way, for example: "Quantum Mechanics", "The nature of the chemical bond" which are highly recommended though require refereshing of mathematical knowledge.
And really, following his prediction, that behaviour was discovered in many compounds.

Attachment: taylor1936.pdf (686kB)

2. I agree that it would be interesting to utilize somehow this behaviour in chemical demonstration. This definitely can affect crystall grow. As for supersaturated solution I think it is easier to try some hardly crystallisable anions.

[Edited on 12-3-2023 by teodor]

teodor - 12-3-2023 at 09:17

 Quote: Originally posted by SnailsAttack It's also not reflected in the table you cited.

It is. Look at 2 different values printed at 20C and 54C. That probably because author was unable to get one defined value.

[Edited on 12-3-2023 by teodor]

fusso - 12-3-2023 at 09:26

Why "Solubility is supposed to be measured as grams per liter of solvent rather than grams per liter of solution" and not the reverse?
The solute is gonna dissolved in solvent to form the final solution. Why is molaRity unimportant? That's what we ultimately want to know, isnt it?

SnailsAttack - 12-3-2023 at 09:51

 Quote: Originally posted by teodor This discovery is connected with the name of Linus Pauling. He wrote several excelent books by the way, for example: "Quantum Mechanics", "The nature of the chemical bond" which are highly recommended though require refereshing of mathematical knowledge. And really, following his prediction, that behaviour was discovered in many compounds.

Alright, the graph in that paper is reasonably convincing. The spike at 53.7°C is about 3°C across, which might be wide enough to test for. It could be referenced against the solubility at ~47°C and ~61°C.

yobbo II - 12-3-2023 at 10:09

What are the units on the left hand side of the graph?
I presume it's degrees centigrade along the bottom (as Teodor said)

Yob

SnailsAttack - 12-3-2023 at 10:28

 Quote: Originally posted by yobbo II What are the units on the left hand side of the graph? I presume it's degrees centigrade along the bottom (as Teodor said) Yob

which graph

unionised - 12-3-2023 at 10:33

Quote: Originally posted by teodor
 Quote: Originally posted by Texium I was typing out a longer response last night, but I lost it by accident, and teodor and Amos pretty much covered what I was going to say.

Oh, yes, it happend for me several times already. Now I always copy responses to an external text editor when they become long and continue there.

I think there is a nice demo illustrating your and Amos idea "there are a lot of variables some of which are unknown".

https://en.wikipedia.org/wiki/Storm_glass

The important thing about storm glasses is that they never actually worked.

yobbo II - 12-3-2023 at 11:00

The graph at the above link.
What are the units on the left hand side of the graph?
I presume it's degrees centigrade along the bottom (as Teodor said)

There is another paper on H2O Copper sulphate system attached.
There is no talk of strange spikes?
It's all rather technical.

Yob

Attachment: CHEM_Sibarani_et_al_Critical_Evaluation_2022_Chemical_Engineering_Science.pdf (3.1MB)

Texium - 12-3-2023 at 13:55

 Quote: Originally posted by SnailsAttack Your guys' wheel is a friggin octagon at best!
Before I recuse myself from this thread to preserve my sanity, let me point out a couple assumptions you've made that will make roundin' off those corners hard if not impossible for you...

 Quote: Originally posted by SnailsAttack 24 hours? I give it five minutes max before a sample of solid copper sulphate emersed in water is effectively fully equilibriated with the solution.
yobbo II already addressed this, but you never replied, so I feel it's worth repeating. This is going to be a big problem for you, and it makes me wonder if you've ever actually tried to make a saturated copper sulfate solution before, cause it tends to take a lot longer than 5 minutes to get everything dissolved, even with heating and stirring, not just immersing the solid in water. You're not going to get close to saturation if you only wait 5 minutes. It'll depend on how large the crystals you're trying to dissolve are, too.

 Quote: Originally posted by SnailsAttack I've observed that anhydrous copper sulphate rehydrates instantly on contact with water, which raises more questions about the semantics of what the solubility of anhydrous copper sulphate actually means if it can't coexist with water.
Yes, it turned blue, wonderful, but do you know for sure that it's immediately forming the pentahydrate? Based on observations that I have made before, I don't think that it is. If you add 5 equivalents of water to 1 equivalent of anhydrous copper sulfate and thoroughly mix it, it doesn't fully absorb all the water and immediately become a nice homogeneous pentahydrate. This won't affect the final outcome, but it will change how long it takes X moles of anhydrous copper sulfate to dissolve in Y+5X moles of water, compared to how long it would take X moles of pentahydrate to dissolve in Y moles of water.
Tsjerk - 12-3-2023 at 14:36

Quote: Originally posted by SnailsAttack
 Quote: Originally posted by fusso Snailsattack, I think you forgot to include the extra volume when calculating solubility.

What extra volume? Solubility is supposed to be measured as grams per liter of solvent rather than grams per liter of solution.

If you measure in g/L of solution you have to know the density of the solution and that's a pain in the butt. I don't think that reframing the measurements in g/L of solution fixes the problem.

Do some more reading and some more calculations until you get it, because as fusso pointed out: that calculation where you get a 10% difference on page 2 is wrong.

Do the calculations again and you will see "reframing" solves your problem.

Solubility is measured in grams per total volume, not grams per amount of solvent. Have you ever made a solution in a chemistry class? You will add an amount of compound to a volumetric flask and add solvent till you reach a certain volume.

And why would you have to know any density to calculate or express any solubility this way? You were the only one in this thread who needed a density in any calculation. You only need density of you would want to know the total weight of the solution. Which is not relevant for solubility.

[Edited on 12-3-2023 by Tsjerk]

SnailsAttack - 12-3-2023 at 16:13

 Quote: Originally posted by Texium Before I recuse myself from this thread to preserve my sanity,

Quote: Originally posted by Texium

 Quote: Originally posted by SnailsAttack 24 hours? I give it five minutes max before a sample of solid copper sulphate emersed in water is effectively fully equilibriated with the solution.
yobbo II already addressed this, but you never replied, so I feel it's worth repeating. This is going to be a big problem for you, and it makes me wonder if you've ever actually tried to make a saturated copper sulfate solution before, cause it tends to take a lot longer than 5 minutes to get everything dissolved, even with heating and stirring, not just immersing the solid in water. You're not going to get close to saturation if you only wait 5 minutes. It'll depend on how large the crystals you're trying to dissolve are, too.

Yeah, in retrospect I think 5 minutes isn't long enough. Maybe I'll do a separate test to see how long it takes before the saturation hits a limit.

 Quote: Originally posted by SnailsAttack Yes, it turned blue, wonderful, but do you know for sure that it's immediately forming the pentahydrate? Based on observations that I have made before, I don't think that it is. If you add 5 equivalents of water to 1 equivalent of anhydrous copper sulfate and thoroughly mix it, it doesn't fully absorb all the water and immediately become a nice homogeneous pentahydrate. This won't affect the final outcome, but it will change how long it takes X moles of anhydrous copper sulfate to dissolve in Y+5X moles of water, compared to how long it would take X moles of pentahydrate to dissolve in Y moles of water.

You're right, it's possible that it doesn't rehydrate immediately to the pentahydrate even in direct contact with water. For instance, I've observed that calcium sulphate is reluctant to form the dihydrate. That's another matter though.

SnailsAttack - 12-3-2023 at 16:24

 Quote: Originally posted by Tsjerk Have you ever made a solution in a chemistry class?
nope
 Quote: Originally posted by Tsjerk Solubility is measured in grams per total volume, not grams per amount of solvent.

Really? This wiki page suggests that the standard units are grams per 100 milliliters of water. Very few sources actually specify whether it's grams per unit solvent or grams per unit solution, however.

SnailsAttack - 12-3-2023 at 20:44

Quote: Originally posted by SnailsAttack

 Quote: Originally posted by Tsjerk Solubility is measured in grams per total volume, not grams per amount of solvent.

Really? This wiki page suggests that the standard units are grams per 100 milliliters of water. Very few sources actually specify whether it's grams per unit solvent or grams per unit solution, however.

Yeah, all the readings I can find on solubility say that “g/mL” or “g/L” refer to grams per unit of solvent, not grams per unit of solution.

Tsjerk - 12-3-2023 at 23:41

Well, that Wiki page is a bit ambiguous, probably because the purpose of that sentence is explaining water is the solvent, not what is meant by g/L. Molarity and g/L are interchangeable by knowing only molecular weight.

Care to share some of those readings?

https://www.omnicalculator.com/chemistry/molarity

[Edited on 13-3-2023 by Tsjerk]

SnailsAttack - 13-3-2023 at 01:33

 Quote: Originally posted by Tsjerk Well, that Wiki page is a bit ambiguous, probably because the purpose of that sentence is explaining water is the solvent, not what is meant by g/L. Molarity and g/L are interchangeable by knowing only molecular weight. Care to share some of those readings?

Yeah, here's some I found by just googling: 1 2 3 4 5

measuring in grams per liter solvent is easier than grams per liter solution because, like I said, maths with the latter requires you to know the density of the solution which requires additional measurements.

SnailsAttack - 13-3-2023 at 02:08

 Quote: Originally posted by yobbo II There is another paper on H2O Copper sulphate system attached. There is no talk of strange spikes? It's all rather technical.

The graph on page 11 looks.. very clean. I might convert that to a table.

Looks like the authors investigated a ton of different sources. If they came across any that included the spikes, they were apparently rejected.

Admittedly I'd love to just ignore that study for the sake of simplicity, but that would be bad science.

Texium - 13-3-2023 at 05:54

 Quote: Originally posted by SnailsAttack Yeah, here's some I found by just googling: 1 2 3 4 5 measuring in grams per liter solvent is easier than grams per liter solution because, like I said, maths with the latter requires you to know the density of the solution which requires additional measurements.
Here’s the thing: that method of measurement is only useful for knowing the maximum solubility of a substance at a given temperature, i.e. what you are trying to do. In nearly every other application, it is more useful to know g/mL solution, or molarity, because you can measure out any volume of the solution and know how much mass is dissolved in that portion. If you had made up the solution as grams/mL solvent, you’d either have to measure the total volume of the solution to be able to do the same, or weigh the solution and measure everything by mass, but either way it’s a pain in the ass and completely unnecessary. So yes, for YOUR purpose, g/mL solvent is easier because you aren’t actually doing anything with the solution you make. If you were using the solution in a reaction or for some spectroscopy experiment, then you would be the one doing extra measurements and math, while someone who prepared using g/mL solution would be all set.

I will reiterate that most chemists simply do not care about the maximum solubility of a substance because it is completely irrelevant unless you are growing crystals or doing some esoteric physical chemistry like some of the papers posted in this thread.

yobbo II - 13-3-2023 at 06:27

At snailsAttack.

"They had trouble getting or identifying a pure pentahydrate......."

edit:

Paper attached on the dehydration of Copper Sulphate.
Extract:
A study of the thermal transitions in copper sulphate pentahydrate has been made from
warming curves obtained by a differential thermocouple method. Small transitions were
observed at 29°C, 35°C, 53. 7°C, and dehydrations at 96.5°C, 102°C, and at 113°C. The first
three of these may be interpreted as transitions from oscillation to rotation of the water molecules
in the crystal. The last three are associated with the stepwise dehydration to the tetra-,
trio, and monohydrates, respectively. The existence of the tetrahydrate, not previously known,
has been demonstrated. A discussion of the dehydration in the light of its crystal structure is
given. A mechanism of dehydration by heat, based on the concept of molecular rotation, has
been suggested.

Graph from above is below:

It has nothing to do with actual solubilities IMHO

See attached file with similar bumps on a graph of heat absorbtion.

Another edit:

Having done some reading on sci.mad. this subject has been doing the rounds for quite some time.

Would there be a hysteresis error?

One thing that may be causing a difference in stating the solubility of different hydrates (I don't know myself) is the way the solution is actually taken to saturation.
You can add your salt (whatever hydrate) to the water and heat up to the measureing temperature and wait for everything to dissolve and then measure the liquid part for m/l or grams per mole water (or whatever).

You could also heat up the water + salt untill everything is dissolved and then slowly cool untill you see the first crystal form (a messy method IMO).
You are now approaching saturation 'from the other side' as it were.

I never heard this mentioned anywhere so it is probably not a problem.

Yob

Attachment: dehyd_OF_cu_sulfate.pdf (1.8MB)

[Edited on 13-3-2023 by yobbo II]

[Edited on 13-3-2023 by yobbo II]

teodor - 13-3-2023 at 10:59

" It is like spikes where properties of crystalls are changing very fast in very narrow range (fractions of degree) and comes back to the curve". I didn't say it is the solubility.

The solubility is dependent on the parameter depicted on the graph.
The necessary condition of thermodinamic equilibrium is "stability against small perturbations" - https://en.wikipedia.org/wiki/Thermodynamic_equilibrium#Stab...

But we can expect the steepness of the graph is related to steepness of solubility changes if we would able to measure it in this small temperature range.

The concept of existing (=constant) equilibrium implies we could find a method how to put the system in this state otherwise we are talking about some virtual property of the model which has no relation to the real life.

Tsjerk - 13-3-2023 at 14:41

I should have listened to Texium. I don't care whether it is gram per liter solvent or per volume anymore. Good luck with it. All I know is that the initial hydration doesn't influence the final solubility.

My god, just use molarity. I have worked in labs for years, never ever have I doubted grams not being per volume. Never have I seen people adding a volume of solvent to a weight of a to be solute.

Saturated solutions you have on hand by having a bottle with a thick layer of solid on the bottom of the bottle, and even then; who cares how much is actually dissolved.

[Edited on 13-3-2023 by Tsjerk]

yobbo II - 13-3-2023 at 15:11

 Quote: Originally posted by teodor " .... I didn't say it is the solubility.

I understand now.
Probably everyone in the thread knew that except Yobbo!

SnailsAttack - 14-3-2023 at 07:43

 Quote: Originally posted by yobbo II You probably answered your own question in the opening question, thus: "They had trouble getting or identifying a pure pentahydrate......."

I found that copper sulphate forms a stoichiometric pentahydrate when crystallized from solution under normal conditions, so its use shouldn't present an issue, although a measurement procedure involving the crystallization of other hydrates at high temperature would get messy.

 Quote: Originally posted by yobbo II Would there be a hysteresis error?

The dissolution process probably doesn't lag too much (solid crystals should equilibrate with solution relatively quickly), but the crystallization process could definitely lag, resulting in supersaturation, hence why I based my method of measuring solubility on the former process.

 Quote: Originally posted by yobbo II One thing that may be causing a difference in stating the solubility of different hydrates (I don't know myself) is the way the solution is actually taken to saturation. You can add your salt (whatever hydrate) to the water and heat up to the measureing temperature and wait for everything to dissolve and then measure the liquid part for m/l or grams per mole water (or whatever).

Yeah, pretty much exactly my process.

 Quote: Originally posted by yobbo II You could also heat up the water + salt untill everything is dissolved and then slowly cool untill you see the first crystal form (a messy method IMO). You are now approaching saturation 'from the other side' as it were.

Yep.

SnailsAttack - 14-3-2023 at 08:02

 Quote: Originally posted by Tsjerk My god, just use molarity. I have worked in labs for years, never ever have I doubted grams not being per volume.

Like I said though, the internet suggests that solubility is usually measured per unit solvent.

 Quote: Originally posted by Tsjerk Never have I seen people adding a volume of solvent to a weight of a to be solute.

Molality has its advantages, namely being that I don't have to bother with volume measurements.

 Quote: Originally posted by Tsjerk Saturated solutions you have on hand by having a bottle with a thick layer of solid on the bottom of the bottle, and even then; who cares how much is actually dissolved.

You're thinking too inside the box. Yes, for most lab purposes you can just take the solubility of 320 g/L at 20°C given by wikipedia, but depending on what source you look at, the solubility of the pentahydrate ranges anywhere from 208 to 360 grams per liter at room temperature. That level of ambiguity isn't good enough, I want to get this stuff figured out so I can start to look at other unsolved problems.

I'm also interested in reliable solubility values for use in geochemistry (e.g. ocean salinity) and purifying salts by solvent extraction and fractional dissolution/recrystallization.

SnailsAttack - 14-3-2023 at 08:35

Solubility of CuSO₄·0H₂O and CuSO₄·5H₂O in grams per kilogram of water
 Code:  0°C 138.5 235.0 10°C 170.8 295.7 20°C 207.8 368.3 30°C 247.4 449.8 40°C 291.9 546.7 50°C 341.6 662.0 60°C 399.7 807.4 70°C 466.4 990.3 80°C 543.2 1225 90°C 637.2 1557 100°C 751.0 2039 

I converted the graph to a table from the paper that Yob posted. The paper's based on an extensive meta-analysis, so I trust it pretty well.

Each value is accurate to probably ± 1.7%. The solubility values for the pentahydrate are calculated from the anhydrate using the formula I came up with in the original post.

I think it's worth mentioning that the mean solubility I measured earlier is accurate to an astounding 0.25% with respect to this table, which might suggest that the error range is lower than what I proposed.

yobbo II - 16-3-2023 at 11:41

From Kirk Othmer:

The pentahydrate slowly effloresces in low humidity or above 30.6C. Above
88C dehydration occurs rapidly.

Anhydrous copper(II) sulfate [7758-98-7] is a gray to white rhombic crystal
and occurs in nature as the mineral hydrocyanite. CuSO 4 is hygroscopic. It is
produced by careful dehydration of the pentahydrate at 250C.

Yob

SnailsAttack - 9-4-2023 at 20:06

I found a pdf copy of a book on solubility that might be valuable:

https://ia904709.us.archive.org/12/items/2ndsolubilitieso00s...

It was published in 1907 and includes data from as far back as the mid 1800s, which is how you know you're in some s***.

There's solubility data for copper sulphate at the bottom of page 300, and the error with respect to the table I posted above ranges from 3.2% to as low as 0.4%, which suggests that it's most likely reliable unless there's some circular referencing going on.

The book actually makes the distinction between solubility as measured by 'unit per unit solution' compared to 'unit per unit solvent', which, if this thread has taught me anything, is extremely important.

In the book's foreword, the author reaches pretty much all the same conclusions as I have:
 Quote: In those cases ... where the amount of solvent is expressed in volume instead of weight, it is first necessary to multiply by the specific gravity of the solvent in order to find the weight corresponding to the given volume.
 Quote: One of the forms of presenting solubility data for which especial care is needed in converting the values to a different basis is in the case of results for salts with water of crystallization.
 Quote: ...if, for instance, the grams of hydrated salt per 100 grams of saturated solution or of water have been given, then it will be necessary to add the weight of water present as water of crystallization in the salt, to the weight of water present as solvent. The total weight of solvent is, therefore, made up of the weight of water used for preparing the solution and that carried by the salt as that of crystallization.

More testing is needed to verify that the theories and formulas I've written in this thread actually work, but at the moment I'd say things are looking pretty good (aside from the fact that I'm falling back to the findings of a textbook written 20 years before the invention of sliced bread)

Texium - 9-4-2023 at 20:37

Everything you’ve quoted is stuff we all already know and fully agreed about, so I’m having a little trouble seeing what new insight you got out of it.
SnailsAttack - 10-4-2023 at 04:38

 Quote: Originally posted by Texium Everything you’ve quoted is stuff we all already know and fully agreed about, so I’m having a little trouble seeing what new insight you got out of it.

it's some confirmation, plus a source of potentially reliable data on the solubility of a whole buttload of salts