As we have pointed out several times in the preceding sections, the Ideal Gas Laws (Chapter 10) tell us that the partial pressure of a gas and the molar concentration of that gas are directly proportional. We can show this simply by beginning with the combined gas law:
If we divide both sides by the volume, V, and state that V must be expressed in liters, the right side of the equation now contains the term . Realizing that the number of moles of gas (n) divided by the volume in liters is equal to molarity, M, this expression can be re-written as:
Using this expression, molar concentrations can easily be substituted for partial pressures, and visa versa.
- For the reaction shown below, if the molar concentrations of SO3, NO and SO2 are all 0.100 M, what is the equilibrium concentration of NO2?
- For the reaction between carbon monoxide and chlorine to form phosgene, the equilibrium constant calculated from partial pressures is K = 0.20. How does this value relate to the equilibrium constant, KC, under the same conditions, calculated from molar concentrations?
CO (g) + Cl2 (g) ⇄ COCl2 (g)