We know that aqueous electrolyte solutions have a lower heat capacity compared to pure water. For example, the heat capacity of a saturated CaCl2 solution at 20℃ (74.5 g/100 g H2O) has a specific heat capacity (Cp) of ~2.4 kJ/kg∙K, much lower than that of water (4.18 kJ/kg∙K).
My question is, how can we explain this phenomenon on a molecular or even quantum perspective?
I understand that, at such a high concentration, there are very few "free" water molecules. The majority of them are "trapped" in the hydration shells of the Ca2+ and Cl- ions. These water molecules from dative covalent bonds with the ions, thus unable to have free translational or rotational movement (i.e. their degrees of freedom are decreased). The water-ion ensemble must now move together.
But how does that explain the lower heat capacity?
Thank you Paul Duineveld for the paper. It elaborates the difference between partial and apparent heat capacity and their uses very clearly.
However, it does not explain the reason why lower heat capacity is observed on the solutions.
Hello Chang-Ting Wang ,
it is, mainly, due to the
"H-bond(s)",
in ('pure & liquid') water.
The lower sp. heat capacity (Cp) in aqueous electrolyte solutions is due to the lower percentage of the Hydrogen bonding (due to the lower[1] percentage of the remaining 'pure & liquid' water).
1. It is a lower percentage because some quantity of (smaller "H-bonding" number/ing) this water, e.g. the rest percentage, is, actually, frozen/trapped near/by the (hydrated-)ions of the electrolyte.
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