Reaction Quotient: Predicting the Direction of Chemical Change
Analogy- You are observing a chemical reaction in progress.
- You know the equilibrium constant $K$, which represents the ratio of products to reactants when the system is at equilibrium.
- But what if the reaction hasn’t yet reached equilibrium?
- How can you predict whether the reaction will move forward to produce more products or backward to regenerate reactants?
This is where the reaction quotient, $Q$, becomes an invaluable tool. By comparing $Q$ to $K$, you can determine the direction of the reaction and understand how the system will adjust to reach equilibrium.
The Reaction Quotient: Definition and Calculation
Definition
Reaction quotient
The reaction quotient, $Q$, measures the relative concentrations of products and reactants in a chemical reaction at a given moment in time, whether or not the system is at equilibrium
- It is calculated using the same formula as the equilibrium constant $K$, but with non-equilibrium concentrations of the reacting species.
- For a general reaction: $$aA + bB \rightleftharpoons xX + yY$$
- The expression for $Q$ is: $$Q = \frac{[X]^x[Y]^y}{[A]^a[B]^b}$$
- Products appear in the numerator, raised to the power of their stoichiometric coefficients.
- Reactants appear in the denominator, also raised to the power of their stoichiometric coefficients.
- Square brackets $[... ]$ denote the concentrations of the species in moles per cubic decimeter ($\text{mol dm}^{-3}$).
Ensure that all concentrations used in the $Q$ calculation are expressed in the same units, typically $\text{mol dm}^{-3}$
Comparing $Q$ and $K$: What It Tells Us
The value of $Q$ in relation to $K$ reveals the direction in which the reaction will proceed to achieve equilibrium:
- If $Q< K$:
- The concentration of reactants is too high (or the concentration of products is too low) compared to the equilibrium state.
