In chemical reactions, equilibrium is the state where the forward and reverse reaction rates are equal, and the net change in concentrations of reactants and products is zero. The equilibrium constant quantifies this state. We have two main types: $K_p$ and $K_c$.
$K_c$ is the equilibrium constant expressed in terms of the concentrations of reactants and products at equilibrium. It's used when dealing with solutions or reactions in the liquid phase.
$K_p$ is the equilibrium constant expressed in terms of the partial pressures of gaseous reactants and products at equilibrium. It's used when dealing with reactions in the gas phase.
| Feature | $K_c$ | $K_p$ |
|---|---|---|
| Definition | Equilibrium constant using concentrations | Equilibrium constant using partial pressures |
| Units | Varies depending on the reaction stoichiometry (mol/L) | Varies depending on the reaction stoichiometry (atm or Pa) |
| Applicable Systems | Primarily used for reactions in solution | Primarily used for reactions involving gases |
| Equation | $K_c = \frac{[C]^c[D]^d}{[A]^a[B]^b}$ | $K_p = \frac{(P_C)^c(P_D)^d}{(P_A)^a(P_B)^b}$ |
| Relationship | $K_p = K_c(RT)^{\Delta n}$, where $\Delta n$ is the change in the number of moles of gas | |
Please log in to post your answer.
Log InEarn 2 Points for answering. If your answer is selected as the best, you'll get +20 Points! 🚀