# Changes in entropy and Gibbs free energy for sublimation at triple point [closed]

The standard enthalpy of sublimation of dry ice (solid carbon dioxide) is $$\Delta H=6.03\ \mathrm{kJ/mol}$$. The triple point of $$\ce{CO2}$$ is at $$p=5.1\ \mathrm{atm}$$, $$T=-56.7\ \mathrm{^\circ C}$$.

I am trying to find the $$\Delta S_\text{sublimation}$$ and $$\Delta G_\text{sublimation}$$ at the triple point. I do not know which formula to use.

## closed as off-topic by Todd Minehardt, Melanie Shebel♦May 19 at 6:27

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• Welcome to Chemistry.SE! We have a site policy for homework questions. Please edit your question to include your attempt at the problem, where you got stuck, and let us know where you're finding difficulty so we may best help you. – Melanie Shebel May 19 at 6:26

$$\Delta_\mathrm{sub} G = 0$$
$$\Delta_\mathrm{sub} S_\mathrm{m} = \frac{\Delta_\mathrm{sub} H_\mathrm{m}}{T_\mathrm{sub}}$$
Note however that the value of $$\Delta_\mathrm{sub} H_\mathrm{m}$$ that you provide seems a bit off$$^\dagger$$. NIST gives a value at $$\pu{207 K}$$ of $$\Delta_\mathrm{sub} H_\mathrm{m} = \pu{26.1 kJmol^{-1}}$$. Values at other temperatures in the range $$\pu{167-195 K}$$ are similar. Other internet sources can helpfully confirm this.
$$\dagger$$ The value you post may be in $$\pu{kcal mol-1}$$, then it would translate into $$\pu{25.3 kJ mol-1}$$, in reasonable agreement with the values on the NIST data page.