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I have calcium chloride in the form of small (1-2 mm) round white pellets, originally bought at a home-brew shop. The packaging gives no indication what hydrate it is. Since each water molecule adds ~16% of the weight of anhydrous CaCl2 this introduces a big error when using it to treat water to a specified amount (in ppm).

Naturally I figured I could throw it in the toaster-oven for a while at 350°F and weigh it every 15 minutes, to figure out how much water was driven off. Results:

0 min.  -  1.15 g.
15 min. -  0.92 g.
30 min. -  0.90 g.
45 min. -  0.90 g.

I originally assumed this would be anhydrous, but after reading that calcium chloride dihydrate decomposes at 347°F, and calcium chloride monohydrate decomposes at 500°F, I'm not so sure.

I couldn't find any other information on figuring out the hydrate online or indicating what state calcium chloride is typically in (at least that I could understand). Am I correct in thinking I'm left with the monohydrate? Or is it actually anhydrous?

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  • $\begingroup$ I don't think your answer is there, but have you seen this? $\endgroup$ – Molx Apr 2 '15 at 2:05
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The mass of your sample decreased by 22%.

If you originally had $\ce{CaCl2.4H2O}$ and quantitatively dehydrated that to $\ce{CaCl2.H2O}$, then you should have seen a decrease in mass of 27%.

If you originally had $\ce{CaCl2.4H2O}$ and quantitatively dehydrated that to $\ce{CaCl2.2H2O}$, then you should have seen a decrease in mass of 18%.

Although it's tempting to say that you just have some mix of mono-and-dihydrates, I think that we just can't ignore the stability of your mass measurements and the temperature of your sample, which due to the radiative heating from the toaster heating coils was probably even greater than $\mathrm{350^oF}$. The most reasonable option remaining is that you originally had primarily $\ce{CaCl2.4H2O}$, but also some $\ce{CaCl2.6H2O}$ mixed in, and now you have just the monohydrate remaining.

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