I have recently done an experiment in an attempt to find the relationship between the water ratio in a hydrate and the enthalpy of hydration, and my results showed a correlation, but an anomaly for calcium chloride. My results were:

Sodium carbonate monohydrate: 10.03 kJ
Calcium chloride dihydrate: -19.51 kJ
Sodium acetate trihydrate: 28.98 kJ
Sodium thiosulfate pentahydrate: 47.37 kJ
Sodium sulfate decahydrate: 66.32 kJ

I have tried researching about what makes calcium chloride dihydrate unique, but I can't seem to find anything. And this is not a random error, as each of these salt compounds had 3 trials each, with consistent results.

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    $\begingroup$ In this group, it is unique for two reasons. It is the only species with calcium cations, and the only species with chloride anions. $\endgroup$ – Karsten Theis Nov 13 '20 at 22:43
  • $\begingroup$ Try also different hydrates of the same compound, like anhydrous, dihydrate and hexahydrate of calcium chloride.( Other hydrates exist, but are less common ). $\endgroup$ – Poutnik Nov 14 '20 at 14:29

The heats of hydration of your salts, plus some others, can be obtained by subtracting the heat of formation of the anhydrous from the heat of formation of the hydrated salt, using the numerical values from the Handbook of Chemistry and Physics, 43rd Ed., 1961, pages 1807 - 1831. The so obtained values are rather different from yours, but the tendency is the same. Here are the obtained molar heats of hydration, divided by the number of water molecules. They are strangely similar.

$\ce{Na2CO3·H2O....... 295.3 kJ (mol H2O)^{-1}}$

$\ce{Na2CO3·10 H2O..... 295.3 kJ (mol H2O)^{-1}}$

$\ce{NaCH3COO·3H2O ... 297.6 kJ (mol H2O)^{-1}}$

$\ce{Na2SO4·10H2O ..... 294.3 kJ (mol H2O)^{-1}}$

$\ce{Na2S·4.5H2O ...... 305.0 kJ (mol H2O)^{-1}}$

$\ce{Na2S·5H2O....... 304.6 kJ (mol H2O)}^{-1}$

$\ce{Na2S·9H2O ....... 301.1 kJ (mol H2O)^{-1}}$

$\ce{Na2S2O3·5H2O..... 297.2 kJ (mol H2O)^{-1}}$

$\ce{Ca(NO3)2·4H2O.... 297.7 kJ (mol H2O)^{-1}}$

$\ce{Ca(CH3COO)2·H2O . 291.2 kJ (mole H2O)^{-1}}$

$\ce{CaCl2·2H2O ...... 179.6 kJ (mol H2O)^{-1}}$ ........ (!)

$\ce{CaCl2·6 H2O ......301.8 kJ (mol H2O)^{-1}}$

It is immediately clear that the heat of hydration of one molecule $\ce{H2O}$ in these compounds is $\ce{300 ± 10 kJ (mol H2O)^{-1}}$. But there is one odd exception : $\ce{CaCl2·2H2O}$. Has anybody an explanation ?

I would have been happy to show my previous numerical values in a table, or an array with lines and several columns, showing the different heats of formation. But I have not been able to write it correctly. There was always a message saying : \hline incorrect. Can somebody tell me how to proceed ? Thank you in advance.

  • 2
    $\begingroup$ Try chemistry.meta.stackexchange.com/questions/86/… $\endgroup$ – Poutnik Nov 14 '20 at 20:34
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    $\begingroup$ That is hardly an answer! $\endgroup$ – Mithoron Nov 14 '20 at 20:50
  • $\begingroup$ @Mithoron. I agree : it is not an answer. But I had not enough room in the "Comment" section" to insert my text and its numerous numerical values. $\endgroup$ – Maurice Nov 15 '20 at 10:02
  • $\begingroup$ This is much better suited as an edit to the question. I'm sure op will appreciate this, too. $\endgroup$ – Martin - マーチン Nov 15 '20 at 11:46
  • $\begingroup$ Maybe make it CW for others to put it into shape? $\endgroup$ – Mithoron Nov 15 '20 at 13:02

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