# Hydration enthalpy of cations

Can someone provide reliable data on the solvation / hydration enthalpies of $\ce{Ti^{4+}}$ and $\ce{Cr^{3+}}$ ions, or explain which is more stable qualitatively? According to this article1

CFSE in $\ce{Cr^{3+}}$ is around 59.7 kcal = 249.7848 kJ

and according to the abstract of an article by Uudsemaa and Tamm,2 the estimated hydration energy of $\ce{Ti^{4+}}$ is around 7800 kJ/mol

Is this data correct? can someone explain qualitatively as well?

## References

(2) Uudsema, M.; Tamm, T. Calculations of hydrated titanium ion complexes: structure and influence of the first two coordination spheres. Chem. Phys. Lett. 2001, 342 (5–6), 667–672. DOI: 10.1016/S0009-2614(01)00617-0.

• Your first link is broken. Anyway CFSE is merely one small contribution to hydration enthalpy (which are ion-dipole interactions). You also need to explain very clearly what you mean by "stability", because that is not a very clear-cut concept. – orthocresol Mar 29 '17 at 14:09
• Logically one would expect the hydration enthalpy of Ti(4+) (smaller, more highly charged cation) to be much more exothermic than that of Cr(3+). However there's a reason why it's an "estimated" hydration energy and there's a reason why it's a computational study. Ti(4+) doesn't exist in water, probably not even as an aquo complex. The paper writes "The Ti(4+) cations have too large charge to be stable as simple hydrates, and usually exist in hydrolysed forms, e.g. [Ti(OH)2](2+) or [Ti(OH)3]+." Reference given is Holleman-Wiberg. – orthocresol Mar 29 '17 at 14:13
• In order to help you out on this question, it would be good if you could include human readable citations in your post. – Martin - マーチン Mar 29 '17 at 14:14
• More the charge density more the hydration energy...You can find the stability this way... – Mitchell Mar 29 '17 at 14:50