= Solution
Approximate the reaction <enthalpy> and <entropy> as constant over the measured interval, so $\Delta_rG^\circ=\Delta_rH^\circ-T\Delta_rS^\circ$. A <linear least squares> fit of all three <free energies> against <temperature> gives
$$
\Delta_rG^\circ\simeq-478580.96+183.9664T\quad\mathrm{J\,mol_{O_2}^{-1}},
$$
and hence
$$
\boxed{\Delta_rH^\circ\simeq-478.6\,\mathrm{kJ\,mol_{O_2}^{-1}},\qquad\Delta_rS^\circ\simeq-184.0\,\mathrm{J\,K^{-1}\,mol_{O_2}^{-1}}}.
$$
The negative <entropy> change is reasonable because <oxygen> gas is incorporated into solids. The fit residuals are $+0.129,-0.257,+0.129\,\mathrm{kJ\,mol_{O_2}^{-1}}$, so the data support this approximate straight-line description. Extrapolating its intercept to zero <temperature> is a way of estimating a constant local <enthalpy>, not evidence that the <heat capacities> vanish or that the same phases persist to zero <temperature>. More <temperatures> and repeated heating/cooling measurements would test curvature and equilibration.
Back to article page
