In this case, according to the described chemical reaction, which takes place between carbon monoxide and hydrogen to produce methanol at 260 °C and 40 atm:

It is possible to calculate the pressure-based equilibrium constant via:

Whereas the change in the Gibbs free energy for the reaction is calculated with the following, assuming these changes can be assumed constant for the temperature range (25°C to 260°C):

Whereas the change in both enthalpy and entropy are based on enthalpies of formation and standard entropies of both carbon monoxide and methanol respectively (exclude hydrogen as it is a single molecule of the same atom):

Thus:

Hence, the pressure-based equilibrium constant will be:

Next, we calculate the concentration-based equilibrium constant:

After that, we calculate the volume for us to get concentrations for the involved species at equilibrium:

![[CO]_0=\frac{200mol}{601.6L}=0.332M](https://tex.z-dn.net/?f=%5BCO%5D_0%3D%5Cfrac%7B200mol%7D%7B601.6L%7D%3D0.332M)
![[H_2]_0=\frac{350mol}{601.6L}=0.582M](https://tex.z-dn.net/?f=%5BH_2%5D_0%3D%5Cfrac%7B350mol%7D%7B601.6L%7D%3D0.582M)
Then, the equilibrium expression and solution according to the ICE chart:

Whose physically-consistent solution would be x = 0.29 M, it means that the equilibrium conversions are:

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