Le Châtelier's Principle and Its Applications
Analogy
- Consider that you’re sitting in a crowded room when someone opens a window, letting in a cold draft.
- You might instinctively reach for a jacket or move to a warmer spot to counteract the chill.
- This natural response to restore comfort mirrors how chemical systems respond to disturbances.
When a system at equilibrium experiences a change in concentration, pressure, or temperature, it adjusts to minimize the disturbance and reestablish equilibrium.
The Principle: Systems Resist Change
Definition
Le Châtelier's Principle
Le Châtelier's Principle states:
"If a dynamic equilibrium is disturbed by a change in the reaction conditions, the system adjusts to counteract the disturbance and restore a new equilibrium."
The “disturbances” can include:
- Concentration changes: Adding or removing reactants or products.
- Pressure changes: Altering the volume or overall pressure in systems involving gases.
- Temperature changes: Heating or cooling the system.
Effect of Concentration Changes
How Does Concentration Affect Equilibrium?
When the concentration of a reactant or product changes, the system shifts to reduce the impact of this disturbance.
- Adding Reactants: The system shifts toward the products to consume the excess reactants.
- Removing Reactants: The system shifts toward the reactants to replenish the loss.
- Adding Products: The system shifts toward the reactants to reduce the excess products.
- Removing Products: The system shifts toward the products to compensate for the removal.
Example
- Consider the synthesis of ammonia: $$
\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g)
$$ - At equilibrium, if more hydrogen ($\text{H}_2$) is added, the forward reaction is favored, producing more ammonia ($\text{NH}_3$) to counteract the disturbance.
- The equilibrium shifts to the right.
Common Mistake
- Many students incorrectly assume that adding more of a substance will always increase its equilibrium concentration.
- Remember, the system adjusts to partially counteract the change, so the final concentration may still differ from the initial equilibrium state.
Effect of Pressure Changes
How Does Pressure Affect Equilibrium?
Pressure changes only affect systems with gaseous reactants or products.
Hint
The direction of the shift depends on the number of gas molecules on each side of the reaction.
- Increasing Pressure: The system shifts toward the side with fewer gas molecules to reduce pressure.
- Decreasing Pressure: The system shifts toward the side with more gas molecules to increase pressure.
Example
- In the same ammonia synthesis reaction: $$
\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g)
$$ - The left-hand side has 4 moles of gas ($1\text{N}_2 + 3\text{H}_2$), while the right-hand side has 2 moles of gas ($2\text{NH}_3$).
- Increasing the pressure shifts the equilibrium to the right, favoring the formation of ammonia.
Tip
- Pressure changes have no effect on systems where the number of gas molecules is the same on both sides of the equation.
- For example, in $\text{H}_2(g) + I_2(g) \rightleftharpoons 2\text{HI}(g)$, pressure changes do not shift equilibrium.
Effect of Temperature Changes
How Does Temperature Affect Equilibrium?
Temperature changes affect both the position of equilibrium and the value of the equilibrium constant ($K$) because temperature influences reaction spontaneity.
- Exothermic Reactions ($\Delta H< 0$):
- Heat is a product.
- Increasing temperature shifts equilibrium toward the reactants, reducing $K$.
- Endothermic Reactions ($\Delta H >0$):
- Heat is a reactant.
- Increasing temperature shifts equilibrium toward the products, increasing $K$.
Example
- For the exothermic synthesis of ammonia: $$
\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g) + \text{Heat}
$$ - Raising the temperature causes the equilibrium to shift left (toward reactants), reducing ammonia yield.
- Conversely, lowering the temperature favors the forward reaction, increasing ammonia yield.
Note
- Temperature is the only factor that changes the value of $K$.
- Concentration and pressure changes only shift the equilibrium position but leave $K$ unchanged.
Summary of Effects on Equilibrium
| Change in Condition | Shift in Equilibrium | Effect on $K$ |
|---|---|---|
| Increase in Reactant Concentration | Toward products | No change |
| Increase in Product Concentration | Toward reactants | No change |
| Increase in Pressure (gases) | Toward fewer gas molecules | No change |
| Increase in Temperature (Exothermic) | Toward reactants | Decreases |
| Increase in Temperature (Endothermic) | Toward products | Increases |
Applications of Le Châtelier's Principle
Industrial Example: The Haber Process
- The Haber process for ammonia synthesis is a classic example of Le Châtelier's Principle in action: $$
\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g) + \text{Heat}
$$ - To maximize ammonia yield:
- High Pressure: Favors the side with fewer gas molecules (ammonia).
- Moderate Temperature: Balances the need for a reasonable reaction rate (higher at higher temperatures) with the equilibrium position (favoring ammonia at lower temperatures).
- Removal of Ammonia: Continuously removing ammonia shifts equilibrium to the right, producing more ammonia.
Key Insight: $K$ Depends Only on Temperature
- Changes in concentration or pressure shift the equilibrium position but do not change $K$.
- Changes in temperature alter $K$ because they affect the relative favorability of the forward and reverse reactions.
Example
- For the ammonia synthesis reaction at 475 K, $K = 0.59$.
- If the temperature is increased, the equilibrium shifts left (favoring reactants), and $K$ decreases.
- However, doubling the pressure does not change $K$, even though the equilibrium position shifts toward the products.
