Getting Started
A chemical system at equilibrium appears static on a macroscopic level, with constant color, pressure, and concentrations. However, at the atomic scale, it is a scene of intense activity, where forward and reverse reactions occur at precisely the same rate. The core challenge this section addresses is predicting how this dynamic balance responds when it is disturbed by an external change, ensuring the system can adjust to find a new state of equilibrium.
What You Should Be Able to Do
After completing this section, you should be able to:
Define dynamic equilibrium and Le Châtelier’s principle.
Predict the direction an equilibrium will shift in response to a change in the concentration of a reactant or product.
Predict how an equilibrium will respond to changes in temperature, based on whether the reaction is endothermic or exothermic.
For reactions involving gases, predict the effect of a change in pressure or volume on the equilibrium position.
Connect these equilibrium shifts to observable macroscopic changes, such as a change in the color or pH of a solution.
Key Concepts & Analysis
Baseline Condition: The State of Dynamic Equilibrium
A reversible reaction reaches dynamic equilibrium when the rate of the forward reaction equals the rate of the reverse reaction. At this point, the concentrations of all reactants and products become constant. This state is described by the equilibrium constant (K), a specific ratio of product concentrations to reactant concentrations. We can measure the state of a system at any moment using the reaction quotient (Q), which has the same mathematical form as K.
If Q < K, the ratio of products to reactants is too small. The forward reaction will dominate to produce more products, and the reaction will proceed to the right.
If Q > K, the ratio of products to reactants is too large. The reverse reaction will dominate to produce more reactants, and the reaction will proceed to the left.
If Q = K, the system is at equilibrium. There is no net change in concentrations.
The "baseline" for any system we analyze is this state of equilibrium, where Q = K.
The Process or Stress: Disturbing the Equilibrium
What happens if we take a system at equilibrium and change the conditions? This imposed change is called a stress. Le Châtelier’s principle provides the framework for our analysis: When a chemical system at equilibrium is subjected to a stress, it will adjust by shifting in the direction that partially counteracts the stress to re-establish equilibrium.
The primary stresses we will consider are:
Change in Concentration: Adding or removing a reactant or product.
Change in Temperature: Adding or removing heat.
Change in Pressure or Volume: For gaseous systems, compressing or expanding the container.
Dilution: For aqueous systems, adding more solvent.
Applying a stress forces the system out of equilibrium, such that Q is no longer equal to K. The system must then react to return to a state of equilibrium.
The Resulting Change: The Shift to Re-establish Equilibrium
The "shift" is the net reaction (either forward or reverse) that occurs to bring Q back to the value of K.
1. Stress: Change in Concentration
If you add a chemical species, the system will shift to consume what was added. If you remove a species, the system will shift to produce more of what was removed.
Example: The Haber-Bosch process for synthesizing ammonia:
N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
Add N₂: The concentration of a reactant increases. To counteract this, the system shifts to the right, consuming N₂ and H₂ to produce more NH₃ until equilibrium is restored.
Remove NH₃: The concentration of the product decreases. To counteract this, the system shifts to the right, producing more NH₃ to replace what was removed.
2. Stress: Change in Temperature
Temperature is the only stress that changes the value of the equilibrium constant, K. To analyze its effect, we can treat heat as a reactant or a product.
Endothermic Reactions (ΔH > 0): Heat is absorbed, so we can think of it as a reactant.
Reactants + heat ⇌ Products
Increasing T (adding heat) will shift the equilibrium to the right to consume the added heat. K increases.
Decreasing T (removing heat) will shift the equilibrium to the left to produce more heat. K decreases.
Exothermic Reactions (ΔH < 0): Heat is released, so we can think of it as a product.
Reactants ⇌ Products + heat
Increasing T (adding heat) will shift the equilibrium to the left to consume the added heat. K decreases.
Decreasing T (removing heat) will shift the equilibrium to the right to produce more heat. K increases.
Example: The equilibrium between nitrogen dioxide (brown gas) and dinitrogen tetroxide (colorless gas):
2NO₂(g) ⇌ N₂O₄(g) ΔH = -57.2 kJ/mol (Exothermic)
Heating this system (adding heat) shifts the equilibrium to the left, favoring the formation of brown NO₂. The color of the gas mixture becomes darker.
Cooling this system (removing heat) shifts it to the right, favoring the formation of colorless N₂O₄. The color of the mixture fades.
3. Stress: Change in Pressure or Volume (Gaseous Systems)
Changes in pressure affect gaseous equilibria where the number of moles of gas on the reactant side differs from the product side. Pressure and volume are inversely related.
Decrease Volume (Increase Pressure): The system will shift to the side with fewer moles of gas to reduce the pressure.
Increase Volume (Decrease Pressure): The system will shift to the side with more moles of gas to increase the pressure.
Example: Again, the Haber-Bosch process:
N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
(1 + 3 = 4 moles of gas) ⇌ (2 moles of gas)
- If the volume is decreased (pressure increases), the equilibrium will shift to the right, toward the side with fewer moles of gas (2 moles of NH₃), to relieve the pressure. This is a key principle used to maximize ammonia yield in industrial settings.
Note: If the number of moles of gas is the same on both sides of the equation, a change in pressure or volume will have no effect on the equilibrium position.
4. Stress: Dilution (Aqueous Systems)
Adding a solvent (e.g., water) to an aqueous equilibrium lowers the concentration of all dissolved species. The system will shift to the side with the greater number of dissolved particles (ions or molecules) to counteract this dilution.
Example: The equilibrium of acetic acid in water:
CH₃COOH(aq) ⇌ H⁺(aq) + CH₃COO⁻(aq)
(1 dissolved particle) ⇌ (2 dissolved particles)
- If water is added, diluting the solution, the equilibrium will shift to the right to produce more ions, partially counteracting the decrease in overall ion concentration. This results in a higher percent ionization of the acid, though the overall [H⁺] will still be lower than before dilution.
Key Models & Representations
The following table summarizes the application of Le Châtelier’s principle for a generic reversible reaction.
| Stress Applied to System | Initial Effect on Q vs. K | Direction of Shift | Effect on the Value of K |
|---|---|---|---|
| Concentration | |||
| Add Reactant | Q < K | Right (toward products) | No change |
| Add Product | Q > K | Left (toward reactants) | No change |
| Remove Reactant | Q > K | Left (toward reactants) | No change |
| Remove Product | Q < K | Right (toward products) | No change |
| Temperature | |||
| Increase T (Endothermic, ΔH>0) | Q < K | Right (toward products) | Increases |
| Increase T (Exothermic, ΔH<0) | Q > K | Left (toward reactants) | Decreases |
| Pressure/Volume (Gases) | |||
| Decrease Volume / Increase Pressure | Depends on moles | Toward side with fewer gas moles | No change |
| Increase Volume / Decrease Pressure | Depends on moles | Toward side with more gas moles | No change |
Key Terms, Quantities, & Concepts
Dynamic Equilibrium: A state in a reversible reaction where the rate of the forward reaction is equal to the rate of the reverse reaction, resulting in no net change in the concentrations of reactants and products.
Le Châtelier’s Principle: States that if a change of condition (a stress) is applied to a system in equilibrium, the system will shift in a direction that partially counteracts the change.
Equilibrium Position: The specific set of reactant and product concentrations that exist at equilibrium. A "shift" refers to a change in this position.
Reaction Quotient (Q): A value calculated using the same expression as the equilibrium constant, but for a reaction mixture that is not necessarily at equilibrium. It is used to predict the direction of a net reaction.
Equilibrium Constant (K): The numerical value of the reaction quotient when a system has reached equilibrium. It is constant for a given reaction at a specific temperature.
Stress: An external change imposed on a system at equilibrium, such as a change in concentration, temperature, or pressure.
Endothermic Reaction: A chemical reaction that absorbs heat from its surroundings (ΔH is positive).
Exothermic Reaction: A chemical reaction that releases heat into its surroundings (ΔH is negative).
Skill Snapshots
Causation:
Cause: Adding HCl to a weak acid equilibrium like CH₃COOH ⇌ H⁺ + CH₃COO⁻. Effect: The concentration of H⁺ (a product) increases, causing the equilibrium to shift to the left to consume the added H⁺.
Cause: Increasing the temperature of the endothermic reaction Co(H₂O)₆²⁺(aq) + 4Cl⁻(aq) ⇌ CoCl₄²⁻(aq) + 6H₂O(l). Effect: The equilibrium shifts to the right, consuming the added heat and causing the solution to change color from pink to blue.
Cause: Decreasing the volume for the reaction 2SO₂(g) + O₂(g) ⇌ 2SO₃(g). Effect: The system shifts to the right, toward the side with fewer moles of gas (2 vs. 3), to reduce the internal pressure.
Comparison:
Changing the concentration of a reactant shifts the equilibrium position, but it does not change the value of the equilibrium constant, K. In contrast, changing the temperature shifts the equilibrium position and changes the value of K.
For a gaseous reaction, decreasing the container volume causes a shift to the side with fewer moles of gas. In contrast, adding a non-reacting inert gas at constant volume increases total pressure but does not cause a shift because the partial pressures of the reacting gases are unchanged.
Adding a reactant causes a shift toward products. In contrast, adding a product causes a shift toward reactants.
Continuity, Change, Over Time (CCOT):
Baseline: A system is at dynamic equilibrium, where Q = K and macroscopic properties are constant.
Change 1: A stress, such as the removal of a product, is applied. The system is now out of equilibrium (Q < K).
Change 2: The forward reaction rate becomes temporarily greater than the reverse rate. This causes a net shift to the right, consuming reactants and forming products until a new equilibrium position is reached where Q once again equals K.
Continuity: Throughout this process, the value of the equilibrium constant K remains unchanged (assuming the temperature was held constant).
Common Misconceptions & Clarifications
Misconception: At equilibrium, the concentrations of reactants and products are equal.
Clarification: At equilibrium, the rates of the forward and reverse reactions are equal. The concentrations are simply constant, and they are rarely equal to each other.
Misconception: A catalyst changes the equilibrium position to favor products.
Clarification: A catalyst speeds up both the forward and reverse reactions equally. It allows a system to reach equilibrium faster but has no effect on the value of K or the final concentrations at equilibrium.
Misconception: Le Châtelier's principle means the system completely reverses the applied change.
Clarification: The principle states the system partially counteracts the stress. If you increase the concentration of a reactant, the system will shift to use some of it up, but the final concentration will still be higher than it was at the initial equilibrium.
Misconception: Adding water to an aqueous equilibrium always shifts it to the right.
Clarification: Dilution shifts the equilibrium toward the side with the greater number of dissolved particles. While this is often the product side (e.g., for weak acid dissociation), it could be the reactant side in other reactions.
One-Paragraph Summary
Chemical equilibrium is a dynamic state of balance, not a static one. Le Châtelier's principle provides a powerful qualitative tool for predicting how this balance responds to disturbances. When a system at equilibrium is stressed by a change in concentration, temperature, or pressure, it will shift its equilibrium position—favoring either the forward or reverse reaction—to partially counteract that stress. This principle explains how manipulating reaction conditions can maximize product yield in industrial processes like ammonia synthesis and why observable properties like the color of a solution can change with temperature. Understanding these shifts is fundamental to controlling and interpreting the behavior of reversible chemical reactions.