My book says that the enthalpy of solution of anhydrous calcium chloride is negative because it produces hydrated calcium chloride. But when hydrated calcium chloride is dissolved in water, the enthalpy turns out to be positive because it is already hydrated. What I don't understand here is: if it is already hydrated and doesn’t want to interact with water anymore, then why can't the enthalpy be zero? Why does it have to be positive? Can it be logically deduced somehow?
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Dissolution enthalpy is kind of scale balancing as being combination of lattice energy ( positive solid matrix breaking enthalpy) and hydration enthalpy(usually negative), following the formal process
$$\ce{CaCl2(s) -> Ca^2+(g) + 2 Cl-(g) ->[H2O] Ca^2+(aq) + 2 Cl-(aq)}$$
The hydration enthalpy is very negative for anhydrous $\ce{CaCl2}$, but just little negative for $\ce{CaCl2.6H2O}$. It leads to negative dissolution enthalpy for the former, but positive one for the latter.
There is more of known $\ce{CaCl2}$ hydrates. Aside of the hexahydrate, the most common $\ce{CaCl2.2 H2O}$ has the dissolution enthalpy much closer to zero than for the above two cases.
For more, see Wikipedia.org: Enthalpy change of solution
