Direction of reversible reactions — AP Chemistry Study Guide

For: AP Chemistry candidates sitting AP Chemistry.

Covers: Prediction of net reaction direction for non-equilibrium reversible systems, comparison of reaction quotient $Q$ to equilibrium constant $K$, application of Le Chatelier's principle, and predicting shifts after concentration, pressure, and temperature disturbances.

You should already know: Definition of reversible reactions and dynamic equilibrium, how to calculate equilibrium constants ($K$) from equilibrium data, how to write reaction quotient expressions from balanced equations.

A note on the practice questions: All worked questions in the "Practice Questions" section below are original problems written by us in the AP Chemistry style for educational use. They are not reproductions of past College Board / Cambridge / IB papers and may differ in wording, numerical values, or context. Use them to practise the technique; cross-check with official mark schemes for grading conventions.

1. What Is Direction of reversible reactions?

Reversible reactions proceed simultaneously in both the forward direction (reactants converting to products) and reverse direction (products converting to reactants). The direction of a reversible reaction refers to which net change will occur in a non-equilibrium system as it progresses toward equilibrium. In the AP Chemistry CED, this topic makes up ~18% of the Unit 7 Equilibrium exam weight, and appears in both multiple-choice (MCQ) and free-response (FRQ) sections, often paired with calculations of $Q$, $K$, or equilibrium concentrations.

It is one of the most frequently tested foundational concepts in equilibrium, as it underpins all further applications including acid-base chemistry, solubility, and reaction thermodynamics. A common misconception is that "direction" means only one reaction occurs: even at equilibrium, forward and reverse reactions proceed at equal rates, so there is no net change. We only describe net direction when the system is not at equilibrium, or after a disturbance to an established equilibrium causes a new net shift to re-establish equilibrium. Standard notation uses "shifts right" for net forward reaction (more products formed) and "shifts left" for net reverse reaction (more reactants formed).

2. Predicting Reaction Direction with Q vs K

The most rigorous, quantitative method to determine net reaction direction is comparing the reaction quotient $Q$ to the equilibrium constant $K$. $Q$ is calculated using the exact same form as $K$, but uses current (non-equilibrium) concentrations or partial pressures of reactants and products, instead of equilibrium values. For a general balanced reaction: $$aA+bB\rightleftharpoons cC+dD$$ The reaction quotient is: $$Q=\frac{[C{]}^c[D{]}^d}{[A{]}^a[B{]}^b}$$ For gas-phase systems, $Q_p$ uses partial pressures of gases instead of molar concentrations, following the same structure.

The core relationship is intuitive if you think of $K$ as the fixed "target ratio" of products to reactants that all systems will naturally reach:

• If $Q<K$: The current product-to-reactant ratio is too low. The system will make more products to reach equilibrium, so net direction is forward (shift right).
• If $Q=K$: The system is already at equilibrium, no net change occurs.
• If $Q>K$: The current product-to-reactant ratio is too high. The system will make more reactants to reach equilibrium, so net direction is reverse (shift left).

Worked Example

For the gas-phase reaction $N_2{O}_4(g)\rightleftharpoons 2N{O}_2(g)$, $K_p=0.36$ at 25°C. The current partial pressures in the tank are $P_{N_2{O}_4}=0.20$ atm, $P_{N{O}_2}=0.15$ atm. Predict the net direction of the reaction as the system approaches equilibrium.

1. Write the expression for $Q_p$ matching the balanced equation: $Q_p=\frac{(P_{N{O}_2}{)}^2}{P_{N_2{O}_4}}$
2. Substitute the current partial pressures into the expression: $Q_p=\frac{(0.15{)}^2}{0.20}=\frac{0.0225}{0.20}=0.1125$
3. Compare $Q_p$ to the given $K_p$: $0.1125<0.36$, so $Q<K$
4. Apply the Q vs K rule: net reaction proceeds forward (shifts right) to produce more $N{O}_2$ and reach equilibrium.

Exam tip: Always confirm you have placed products in the numerator and reactants in the denominator for $Q$, exactly the same as you do for $K$. Swapping these two is the most common error on this topic, which will flip your direction prediction.

3. Predicting Shifts with Le Chatelier's Principle

Le Chatelier's principle is a qualitative rule that predicts the net direction an equilibrium system will shift after a disturbance. It states: 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. Each type of disturbance follows a predictable pattern:

1. Concentration changes: Adding any reactant or removing any product shifts right to consume the extra reactant or replace the removed product. Adding any product or removing any reactant shifts left to consume the extra product or replace the removed reactant. Concentration changes never change the value of $K$.
2. Pressure/volume changes (gas-phase only): Increasing pressure by decreasing volume shifts the system to the side with fewer moles of gas, to reduce total pressure. Decreasing pressure by increasing volume shifts to the side with more moles of gas. If moles of gas are equal on both sides, there is no shift. Pressure/volume changes never change the value of $K$.
3. Temperature changes: Only temperature changes alter the value of $K$. For endothermic reactions ($\Delta H>0$, heat acts as a reactant), increasing temperature shifts right, and $K$ increases. For exothermic reactions ($\Delta H<0$, heat acts as a product), increasing temperature shifts left, and $K$ decreases.

Worked Example

The reaction $2S{O}_2(g)+O_2(g)\rightleftharpoons 2S{O}_3(g)$ $\Delta H=-198$ kJ/mol is at equilibrium. Predict the net direction shift after each disturbance: (a) add more $O_2$, (b) increase the temperature, (c) decrease total pressure by increasing container volume.

1. For (a): $O_2$ is a reactant. Adding a reactant causes the system to shift right (forward) to consume the added $O_2$, per Le Chatelier's principle.
2. For (b): $\Delta H$ is negative, so the reaction is exothermic, and heat is a product of the forward reaction. Increasing temperature adds heat, so the system shifts left (reverse) to consume the excess heat.
3. For (c): Count moles of gas on each side: 2 + 1 = 3 moles on the reactant side, 2 moles on the product side. Decreasing pressure shifts to the side with more moles of gas, so net shift left (reverse).

Exam tip: On FRQ questions asking for justification, always explicitly connect your prediction to Le Chatelier's principle and the type of disturbance. For example, do not just say "shifts left" — say "shifts left because increasing temperature adds heat to an exothermic reaction, so the system shifts left to consume excess heat".

4. Special Cases of Pressure/Volume Disturbances

Many students struggle with special cases where a total pressure change does not actually cause a shift, because it does not change the partial pressures of reactants and products. The core rule to remember is: a pressure change only causes a shift if it changes the partial pressures (or concentrations) of the reacting species. Two of the most commonly tested special cases are:

1. Adding an inert gas (unreactive gas) at constant volume: Adding inert gas increases total pressure, but the volume and moles of reacting gases stay the same, so their partial pressures do not change. $Q$ remains equal to $K$, so no net shift.
2. Adding an inert gas at constant total pressure: To keep total pressure constant when adding inert gas, the container volume must increase. This decreases the partial pressure of all reacting gases, which is identical to decreasing pressure by increasing volume, so the shift follows the mole rule: shift to the side with more moles of gas.

Worked Example

The reaction $H_2(g)+I_2(g)\rightleftharpoons 2HI(g)$ is at equilibrium in a rigid 1.0 L container at 25°C. Predict the net direction shift when 0.5 mol of argon (inert gas) is added to the container at constant temperature.

1. The container is rigid, so volume stays constant after adding argon.
2. The moles and partial pressures of $H_2$, $I_2$, and $HI$ do not change, so $Q$ does not change and still equals $K$.
3. Conclusion: the system remains at equilibrium, no net shift occurs.

Exam tip: Never automatically assume inert gas addition causes no shift. Always first check if volume or pressure is held constant, and confirm whether partial pressures of reactants/products change.

5. Common Pitfalls (and how to avoid them)

• Wrong move: Mixing up the Q vs K rule, claiming $Q<K$ means shift left. Why: Students confuse the product-to-reactant ratio, or mix up which ratio corresponds to which direction. Correct move: Use the mnemonic "Q smaller than K = need more Products = shift Right" (Q<K→P→R) to avoid mixing up.
• Wrong move: Claiming any change that causes a shift also changes $K$. Why: Students confuse temperature changes with concentration/pressure changes, all of which cause shifts. Correct move: Memorize that only temperature changes change the value of $K$; all other disturbances leave K unchanged.
• Wrong move: Counting moles of solid/liquid when predicting pressure shift direction, for example claiming $CaC{O}_3(s)\rightleftharpoons CaO(s)+C{O}_2(g)$ shifts right with increasing pressure because there are more moles of product. Why: Students count all moles instead of only gaseous moles. Correct move: Only count moles of gaseous reactants and products when predicting shifts from pressure/volume changes.
• Wrong move: Claiming adding an inert gas always causes a shift. Why: Students forget that constant volume inert gas addition does not change partial pressures. Correct move: For any inert gas addition, first confirm if volume is constant (no shift) or pressure is constant (shift follows mole count) before answering.
• Wrong move: Claiming increasing temperature shifts exothermic reactions right. Why: Students forget heat is a product for exothermic reactions. Correct move: Always write heat into the reaction equation (e.g. $2S{O}_2+O_2\rightleftharpoons 2S{O}_3+\text{heat}$ for exothermic) before applying Le Chatelier's principle.

6. Practice Questions (AP Chemistry Style)

Question 1 (Multiple Choice)

For the reaction $AgCl(s)\rightleftharpoons A{g}^{+}(aq)+C{l}^{-}(aq)$, $K_{sp}=1.8\times 10^{-10}$ at 25°C. A solution has $[A{g}^{+}]=1.0\times 10^{-5}$ M and $[C{l}^{-}]=2.0\times 10^{-5}$ M. Which of the following correctly describes the net reaction direction? A) Precipitation of solid AgCl occurs, net direction is reverse. B) Dissolution of solid AgCl occurs, net direction is forward. C) The system is at equilibrium, no net change. D) Precipitation of solid AgCl occurs, net direction is forward.

Worked Solution: First calculate $Q$ (the ion product, equivalent to reaction quotient for solubility): $Q=[A{g}^{+}][C{l}^{-}]=(1.0\times 10^{-5})(2.0\times 10^{-5})=2.0\times 10^{-10}$. Compare $Q$ to $K_{sp}$: $2.0\times 10^{-10}>1.8\times 10^{-10}$, so $Q>K$. For this reaction, forward direction is dissolution of solid AgCl to ions, and reverse direction is precipitation of solid AgCl from ions. A $Q>K$ means we have too many ions, so net reverse reaction (precipitation) occurs. Correct answer: A.

Question 2 (Free Response)

The Haber process for ammonia synthesis follows the reaction: $$N_2(g)+3H_2(g)\rightleftharpoons 2N{H}_3(g)\Delta H=-92\text{ kJ/mol}$$ The system is initially at equilibrium. For each disturbance below, predict the net direction of shift, state whether $K$ increases, decreases, or stays the same, and justify your answer. (a) More $N_2$ gas is added to the system at constant volume and temperature. (b) The temperature of the system is increased at constant volume. (c) The total volume of the container is decreased at constant temperature.

Worked Solution: (a) Net shift: right (forward). $K$ stays the same. Justification: Adding $N_2$, a reactant, disturbs equilibrium, so the system shifts right to consume the added $N_2$ per Le Chatelier's principle. Temperature is unchanged, so $K$ does not change. (b) Net shift: left (reverse). $K$ decreases. Justification: The reaction is exothermic ($\Delta H<0$), so heat is a product of the forward reaction. Increasing temperature adds heat, so the system shifts left to consume the excess heat. Shifting left reduces the equilibrium amount of product, lowering the product/reactant ratio, so $K$ decreases. (c) Net shift: right (forward). $K$ stays the same. Justification: Decreasing volume increases the partial pressure of all gases. There are 4 moles of gas on the reactant side and 2 moles on the product side, so the system shifts right to reduce the total number of gas moles, counteracting the pressure increase. Temperature is unchanged, so $K$ does not change.

Question 3 (Application / Real-World Style)

Human blood maintains a bicarbonate buffer system that regulates pH, with the overall reaction: $C{O}_2(g)+H_{2}O(l)\rightleftharpoons H^{+}(aq)+HC{O}_3^{-}(aq)$ with $K=1.7\times 10^{-3}$. When a person exercises, their cells produce excess $C{O}_2$ that diffuses into the blood. After $C{O}_2$ addition, the measured concentrations are $[C{O}_2(aq)]=0.0025$ M, $[H^{+}]=0.00040$ M, $[HC{O}_3^{-}]=0.0024$ M. Calculate $Q$, predict the net direction of shift, and explain what this means for blood pH.

Worked Solution: Write the reaction quotient for the overall reaction: $$Q=\frac{[H^{+}][HC{O}_3^{-}]}{[C{O}_2]}=\frac{(0.00040)(0.0024)}{0.0025}=3.84\times 10^{-4}$$ Compare $Q$ to $K$: $3.84\times 10^{-4}<1.7\times 10^{-3}$, so $Q<K$, meaning net shift is forward (right). In context, adding excess $C{O}_2$ (a reactant) causes net forward reaction that produces more $H^{+}$, lowering blood pH (a condition called mild respiratory acidosis during heavy exercise). The body compensates by increasing breathing rate to expel excess $C{O}_2$, shifting the reaction back left and raising pH back to normal.

7. Quick Reference Cheatsheet

8. What's Next

This topic is the foundational prerequisite for all further equilibrium topics in AP Chemistry. Next, you will apply direction prediction rules to build ICE tables for calculating equilibrium concentrations, which requires correctly assigning signs to concentration changes based on the net direction of shift. Without mastering direction prediction here, you will get the wrong signs for changes, leading to incorrect equilibrium results and lost points on both MCQ and FRQ problems. This topic also feeds into acid-base equilibrium, solubility equilibrium, and thermodynamics, where you will use the Q vs K comparison to predict reaction spontaneity and precipitation behavior.

• Calculating equilibrium concentrations (ICE tables)
• Le Chatelier's principle for equilibrium shifts
• Solubility product and precipitation reactions
• Gibbs free energy and equilibrium