Equilibrium Constant Kp and Gaseous Systems

A set of vocabulary flashcards based on lecture notes regarding gaseous equilibria, partial pressures, mole fractions, and the factors affecting the equilibrium constant Kp.

Partial pressure ($p$)

The pressure that an individual gas in a mixture would exert if it occupied the container on its own; the sum of these pressures equals the total pressure.

Pascal ($Pa$)

The SI unit of pressure, where $1$ pascal is equal to a pressure of $1$ newton per square metre ($1\,Nm^{-2}$).

Mole fraction of a gas A

The number of moles of gas A in a mixture divided by the total number of moles of gas in the mixture.

Calculation of Partial pressure ($p$) of A

$$\text{mole fraction of A} \times \text{total pressure}$$

$$K_p$$

The equilibrium constant for gaseous equilibria, derived from the equilibrium law using partial pressures of reactants and products instead of concentrations.

Equilibrium expression for $aA(g) + bB(g) \rightleftharpoons yY(g) + zZ(g)$

$$K_p = \frac{p^y Y(g)_{eqm} \times p^z Z(g)_{eqm}}{p^a A(g)_{eqm} \times p^b B(g)_{eqm}}$$

Haber process equilibrium

The synthesis of ammonia represented by the equation $3H_2(g) + N_2(g) \rightleftharpoons 2NH_3(g)$, where $K_p$ has units of $Pa^{-2}$.

Temperature effect on exothermic gaseous equilibrium

For reactions where $\Delta H$ is negative (like the Haber process at $-92\,kJ\,mol^{-1}$), increasing the temperature decreases the value of $K_p$ and the yield of products.

Pressure effect on gaseous equilibrium

Increasing the total pressure forces the equilibrium toward the side with fewer molecules to reduce the pressure; it only affects the position if there is a change in the total number of molecules.

Rate of gaseous equilibrium

Increasing the pressure or temperature, or using a catalyst, increases the rate at which equilibrium is attained in a gas phase reaction.

Homogeneous systems

Chemical systems, such as gaseous equilibria, where all reactants and products exist in the same phase.