So far we have calculated the equilibrium of gases using their concentrations (Keq). However, gaseous systems can also be expressed in terms of pressure (Kp).
Pressure can have several different units:
atmospheres (atm)
kilopascal (kPa)
millimetres of Mercury (mmHg)
Any unit can be used, as long as ALL the values have the same unit.
This is used to determine the ratio of each species in a mixture.
EXAMPLE:
1 mole of Chlorine gas is mixed with 5 moles of phosphorous trichloride. Therefore, there is a total of 6 moles of gas present.
mole fraction of Cl2 = number of moles of Cl2 / number of moles of gas present
mole fraction of PCl3 = number of moles of PCl3 / number of moles of gas present
When gases collide with the walls of their container they increase pressure (therefore pressure is due to collisions).
EXAMPLE:
Let's assume the vessel is under pressure of 120 atmospheres (atm). The partial pressure is the proportion of the pressure that is due to the collisions for a particular gas present. Based on the mole fractions above, phosphorus trichloride is 5x more likely to make contact with the walls of the vessel than chlorine.
We calculate partial pressure as:
Partial pressure of gas A = mole fraction of gas A x total pressure of the system
We use the partial pressure figures, instead of concentration to calculate Kp.
Some things to remember:
Square brackets mean concentration, so they cannot be used here.
Only gases can be used to calculate Kp
Kp DOES NOT EQUAL Keq (there is a relationship between them, but an understanding of it is not required)