APFIVE Which of the following conclusions is best supported by the evidence that for the weak acid equilibrium $$HA \rightleftharpoons H^+ + A^-$$, the reaction quotient $$Q = \frac{[H^+][A^-]}{[HA]}$$ calculated from the given concentrations (yielding $$Q = 1.0 \times 10^{-3}$$) is greater than the acid dissociation constant $$K_a = 1.8 \times 10^{-5}$$?
| Species | Concentration (M) |
|---|---|
| $$H^+$$ | $$1.0 \times 10^{-3}$$ |
| $$A^-$$ | 0.10 |
| $$HA$$ | 0.10 |
| $$K_a$$ | $$1.8 \times 10^{-5}$$ |
Because $$Q > K_a$$, the reaction must shift to the right in order to produce more $$H^+$$ and $$A^-$$.
Since $$Q > K_a$$, the equilibrium will shift to the left, favoring the formation of undissociated acid $$HA$$.
Despite $$Q$$ exceeding $$K_a$$, the system’s shift direction cannot be determined because common ion effects might override the equilibrium.
A reaction quotient greater than the dissociation constant indicates that the system is at equilibrium with no net directional shift.
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