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If I have a hydrate, such as copper (II) sulfate pentahydrate, $\ce{CuSO_4 \cdot 5H_2O}$, what is the amount of heat to required to dehydrate it? I assume that it is equal to or greater than the amount of heat required to evaporate the water molecules. Does the required heat vary for every hydrate?

The formula for dehydrating copper (II) sulfate pentahydrate is:

$$\ce{CuSO_4 \cdot 5H_2O ->[\Delta] CuSO_4 +5H_2O}$$

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According to Wikipedia, copper(II) sulfate will dehydrate from penta- to trihydrate at 63 °C, then to monohydrate at 109 °C and then finally to the anhydrate at 200 °C. It melts at 150 °C, but won't decompose into ions until 650 °C which is beyond most benchtop rigs (your Pyrex flask would give way at about 500 °C).

The required heat of decomposition, as you can see, does vary for each hydrate (especially for the same salt), but doesn't always have to be greater than the boiling point; water has a significant vapor pressure at nearly all ambient Earth temperatures (even at freezing), and so in a dry environment, water will evaporate rapidly at temperatures well below boiling.

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    $\begingroup$ Pyrex melts at 500°C? Huh? Softens, maybe, but I hardly believe any glass actually melts at such a temperature. $\endgroup$ Apr 13 '19 at 0:56
  • $\begingroup$ Depends on your definition of "melt". Glass, and many polymers, transition through a "putty" phase before becoming truly liquid. Practically every lab scientist of the last half-century has learned to form Pyrex rods or tubing over a Bunsen flame. Borosilicate at 500*C might not meet the dictionary definition of a liquid, but I wouldn't rely on it to hold the weight or pressure of a reaction it was containing at that temperature. $\endgroup$
    – KeithS
    Apr 22 at 17:55

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