Chemical Thermodynamics of Tin - Volume 12 - OECD Nuclear ...

VIII.1 Halide compounds
147
SnCl 2 (cr) + 2 H 2 O(l) → SnCl 2·2H 2 O(cr)
an entropy change of – 61.59 J·K –1·mol –1 .
(VIII.1)
From calorimetric dissolution experiments of anhydrous tin(II) chloride and its
dihydrate in aqueous HCl containing H 2 O 2 Vasiliev et al. [1973VAS/VAS2] determined
the standard enthalpies of formation:
anhydrous tin(II) chloride:
f
ο
m
tin(II) chloride dihydrate:
f
Δ H (SnCl 2 , cr, 298.15 K) = − (327.9 ± 2.2) kJ·mol –1
ο
m
Δ H (SnCl 2 .2H 2 O, cr, 298.15 K) = − (918.26 ± 1.42) kJ·mol –1 .
These values are selected.
This yields:
Δ G (SnCl 2·2H 2 O, cr, 298.15 K) = − (760.68 ± 1.49) kJ·mol –1 .
f
ο
m
From the difference of the formation enthalpies using the same value of the
standard enthalpy of formation of water as the authors (= − 285.829 kJ·mol –1 ) for
Reaction (VIII.1) an enthalpy change of − 18.7 kJ·mol –1 is obtained. Combining the
entropy and enthalpy change the standard Gibbs energy of hydration becomes
− 0.34 kJ·mol –1 . However, this value means that regardless of speciation the dihydrate
would never exist as an equilibrium phase in aqueous solutions.
According to Section VIII.2.2.1, a solution saturated with tin(II) chloride
dihydrate is 12.3 molal at 25 °C. Assuming that this solution would be in equilibrium
also with the anhydrous phase the equilibrium constant of Reaction (VIII.2) could be
calculated.
SnCl 2 (cr) + 2 H 2 O(l) SnCl 2·2H 2 O(s).
(VIII.2)
From the value of the Gibbs energy of hydration of − 0.34 kJ·mol –1 the
equilibrium water activity of Reaction (VIII.2) would be 0.934, which is much too high.
With the mole fraction of water x(w) = 0.819 in the saturated solution and even in case
of an ideal solution a Δ rGm
= 2RT ln(0.819) = − 0.99 kJ·mol –1 would be nessecary to
enable formation of the dihydrate. Since the real solution has water activity coefficients
below 1 the nessecary Δ r G m must be more negative. Also the equilibrium relative
humidity above the solid hydrate would be 93.4% according to the Gibbs energy of
hydration, which seems to be too high for a bivalent metal halide.
A later calorimetric work of Vasiliev et al. [1976VAS/KOK] yields even a less
negative value of the hydration enthalpy (approx. – (16.7 ± 1.7) kJ·mol –1 ) comparing
directly the dissolution enthalpies of the anhydrous salt and the dihydrate obtained by
formal integration of the Cp
( T ) function from the measurements of [1974MAT/OGU].
In conclusion the entropy of the dihydrate is maybe too negative by at least 3 to
CHEMICAL THERMODYNAMICS OF TIN, ISBN 978-92-64-99206-1, © OECD 2012