SWIM tried making calcium sulfate CaSO4 from calcium chloride CaCL2 and magnesium sulfate heptahydrate MgSO4. These were sourced from the garden supply store, and SWIM verified that their msds listed >99% composition of each desired reagent.

SWIM ran the math for what to mix to get 1kg of CaSO4 and dissolved the reagents in RO treated water.

MgSO4(aq) + CaCl2(aq) ===> CaSO4(s) + MgCl2(aq) 
246.47g/mol + 110.98g/mol >>162.17g/mol + ......

The reagents were combined in a bucket and stirred to form a white clay with particles in suspension that SWIM left to separate for 3 days. SWIM then filtered it through cotton shirt to remove water, then washed it with several times it’s volume in new ROwater while stirring, and finally filtered again and removed as much water as possible before baking it in a Pyrex tray at 250c in convection oven until dry (~3h).

The material was crumbled and packed in jars in a consistency that kinda looks like Drierite but is more porous. However, samples of it where dissolved in RO water and show pH10.

SWIM read that CaSO4 must have pH7.7 in solution, so SWIM must have made some basic contaminants, and adding new washes of water with MgSO4 (acid) yields no change to the pH.

Why do y’all think the ph is so high?
It doesn’t look like this batch can be used to dry acetone, which was the end goal of this experiment.

Did you measure the pH of your starting materials? Probably your CaCl2 was basic, if it was the anhydrous stuff.

Possibly your t-shirt was basic from washing powder residue, or your pH measurement equipment/method is off.

The CaCl2 was basic. It was the anhydrous stuff as you inferred. I also checked that the ph probe was calibrated properly at pH4 and 7.

I imagine the excess of hydroxide must have reacted to form some magnesium and calcium hydroxides as contaminants of my CaSO4

I decided to neutralize the hydroxide with a small amount of phosphoric acid, and then realized that this could potentially create other insoluble contaminants as phosphate salts of Mg and Ca. However no precipitation was observed during the acid addition, and moreover, dissolution of some of the crystals already present in my reaction vessel was observed. The pH also dropped to around 5 after thorough mixing, but repeated washes of the precipitate brought the pH to about 6.8-6.9. Mission accomplished?

Would you consider this an acceptable CaSO4 for preparation of anhydrous acetone without risking Aldol self-condensation?

If you were to neutralise the excess hydroxide with anything it should be HCl, then you're not introducing any further anions to confuse matters. HCl also has the advantage of being fully volatile, so a small excess is not problematic.

Thus: dissolve crude CaCl2 in dilute HCl, check pH (should be </= 7), then add your magnesium sulfate solution. Don't add solid MgSO4 as the crystals will become coated in CaSO4. To increase precipitation warm the solution - calcium sulfate has a retrograde solubility curve.

All that said, the aldol condensation thing is a bit overstated (even if I've raised it as a potential issue in the past ) If you only dry as much acetone as you'll be needing just using MgSO4 shouldn't present any problems. Warming acetone with MgSO4 for long periods would be the thing to avoid. But CaSO4 will be a more effective drying agent so do as you please.