Chemistry · Unit 7: Equilibrium · 14 min read · Updated 2026-05-11
Direction of reversible reactions — AP Chemistry
AP Chemistry · Unit 7: Equilibrium · 14 min read
1. Core Concept: Net Direction of Reversible Reactions★★☆☆☆⏱ 3 min
Reversible reactions proceed simultaneously in both forward (reactants converting to products) and reverse (products converting to reactants) directions. Net direction describes which overall change will occur when a system is not at equilibrium, or after a disturbance to an existing equilibrium. At equilibrium, forward and reverse reaction rates are equal, so there is no net change in concentrations.
2. Quantitative Prediction: Comparing Q and K★★★☆☆⏱ 4 min
The most rigorous method to determine net direction is comparing the reaction quotient ($Q$), calculated from current non-equilibrium concentrations, to the equilibrium constant $K$, which is fixed at a given temperature. $Q$ uses the exact same expression form as $K$, only substituting current instead of equilibrium values.
aA + bB \rightleftharpoons cC + dD \\ Q = \frac{[C]^c[D]^d}{[A]^a[B]^b}
If $Q < K$: Current product ratio is too low → net forward (shift right)
If $Q = K$: System is at equilibrium → no net change
If $Q > K$: Current product ratio is too high → net reverse (shift left)
Exam tip: Always confirm your expression matches the balanced reaction stoichiometry
3. Qualitative Prediction: Le Chatelier's Principle★★★☆☆⏱ 4 min
Le Chatelier's principle is a qualitative rule that predicts how an equilibrium system shifts after a disturbance. It states that when a system at equilibrium is disturbed by a change to concentration, pressure/volume, or temperature, the system will shift in the net direction that counteracts the disturbance.
**Concentration changes**: Adding a reactant/removing a product shifts right; adding a product/removing a reactant shifts left. $K$ never changes.
**Pressure/volume changes (gases only)**: Increasing pressure (decreasing volume) shifts to the side with fewer moles of gas; decreasing pressure (increasing volume) shifts to the side with more moles of gas. Equal moles → no shift. $K$ never changes.
**Temperature changes**: Only temperature changes alter the value of $K$. For endothermic ($\Delta H > 0$, heat = reactant): increasing T shifts right, $K$ increases. For exothermic ($\Delta H < 0$, heat = product): increasing T shifts left, $K$ decreases.
4. Special Cases: Inert Gas Pressure Disturbances★★★★☆⏱ 3 min
A pressure change only causes an equilibrium shift if it changes the partial pressures of reactant and product gases. Two commonly tested special cases involve adding inert (unreactive) gases that do not participate in the reaction:
**Inert gas at constant volume**: Total pressure increases, but partial pressures of reacting gases do not change → $Q$ remains equal to $K$ → no shift.
**Inert gas at constant total pressure**: Volume must increase to keep total pressure constant, which decreases partial pressures of all reacting gases → shift follows the mole rule (to the side with more moles of gas).
5. AP Style Worked Practice Problems★★★★☆⏱ 6 min
Common Pitfalls
Why: Confusion over the product-to-reactant ratio leads to flipped direction predictions
Why: Confusing temperature changes with concentration/pressure changes, which all cause shifts but do not alter K
Why: Students incorrectly count all moles instead of only gaseous moles when applying the pressure shift rule
Why: Students forget that constant volume inert gas addition does not change partial pressures of reacting species
Why: Students forget that heat is a product for exothermic forward reactions