Page 150 - Chemical equilibria Volume 4

126 Chemical Equilibria
NOTE 4.3.– it is important not to confuse the bond energy, which is an
average value, with the energy of dissociation of that bond, which is the
energy needed to break the bond in a given compound. For example, the
energy of dissociation of the C-H bond in methane is different depending on
whether we are talking about breaking the first C-H bond or breaking the
second, with the first already having been broken.
4.2.7.6. Estimation of the enthalpies of formation on the basis of the
bond energies
The enthalpy of formation of a compound can quickly be evaluated on
the basis of the bond energies; we merely need to add together all of the
bond energies in the compound, which gives us the opposite of the enthalpy
of the reaction by which the compound is synthesized. Based on that
enthalpy of synthesis, we are able to work out the standard energy of
formation (see section 4.2.7.1).

4.3. Reaction entropies

For the enthalpies, we have also managed to assign to each compound a
standard enthalpy by the standard enthalpies of formation, by arbitrarily giving
the value of zero to the enthalpy of formation of the simple substances at any
temperature. In the case of entropies, we shall construct a scale by deciding
that for all solid substances, the entropy has the same value at a certain
temperature. Thus, we shall define an absolute scale of the entropies.

4.3.1. Planck’s hypothesis – calculation of the calorimetric
entropies

In 1911, Planck hypothesized that the entropies of all crystalline
substances tend toward zero as their temperature tends toward zero. Thus,
we write:
s → 0 as T → 0 [4.32]
T
0
We deduce the same property for the standard entropy associated with the
reaction:

0 ∑
Δ s = ν k s 0( ) k = 0 [4.33]
0
0
r