How does the addition of a pure solid reactant affect the equilibrium position of a system?

It has no effect on the equilibrium position. Pure solids have constant activities and do not appear in the equilibrium expression, so changing their amount does not create a stress.

What is the difference between the effect of a catalyst and a temperature change on equilibrium?

A catalyst increases the rate of both forward and reverse reactions equally without shifting the equilibrium position or changing $K$. A temperature change shifts the equilibrium position and is the only factor that changes the value of $K$.

In a gas-phase reaction where the moles of reactant gas equal the moles of product gas, how does increasing pressure affect the shift?

There is no shift in the equilibrium position. Since both sides have the same number of gas moles, the system cannot counteract a pressure change by shifting in either direction.

If a system is diluted with water, in which direction will the equilibrium shift?

The system will shift toward the side with the greater number of aqueous particles. This shift attempts to counteract the decrease in concentration by producing more dissolved species.

Why does increasing the temperature of an exothermic reaction decrease the value of $K$?

In an exothermic reaction, heat is a product. Adding heat shifts the equilibrium toward the reactants (left), decreasing the numerator (products) and increasing the denominator (reactants) in the $K$ expression.

What happens to the reaction quotient $Q$ when a reactant is added to a system at equilibrium?

Adding a reactant increases the denominator of the $Q$ expression, making $Q < K$. To restore equilibrium, the system must shift forward to increase products until $Q$ again equals $K$.

What is the 'Common Ion Effect' in the context of Le Châtelier’s Principle?

It is a specific case of concentration stress where adding an ion already present in the equilibrium mixture (from another source) causes the system to shift to consume that ion, often decreasing the solubility of a salt.

Does adding an inert gas at constant volume shift the equilibrium?

No, it does not. While the total pressure increases, the partial pressures (and thus concentrations) of the reacting gases remain unchanged, so the reaction rates are unaffected.

How do you identify if a reaction is endothermic or exothermic using $K$ values at different temperatures?

If $K$ increases as temperature increases, the reaction is endothermic. If $K$ decreases as temperature increases, the reaction is exothermic.

What is the 'stress' in a volume increase for a gaseous system?

The stress is a decrease in the partial pressures and concentrations of all gaseous species. The system responds by shifting toward the side with more gas moles to restore pressure.

Le Châtelier’s Principle | College Board AP Chemistry

Q: What is the 'Common Ion Effect' in the context of Le Châtelier’s Principle?

It is a specific case of concentration stress where adding an ion already present in the equilibrium mixture (from another source) causes the system to shift to consume that ion, often decreasing the solubility of a salt.

Q: Does adding an inert gas at constant volume shift the equilibrium?

No, it does not. While the total pressure increases, the partial pressures (and thus concentrations) of the reacting gases remain unchanged, so the reaction rates are unaffected.

Q: How do you identify if a reaction is endothermic or exothermic using $K$ values at different temperatures?

If $K$ increases as temperature increases, the reaction is endothermic. If $K$ decreases as temperature increases, the reaction is exothermic.

Q: What is the 'stress' in a volume increase for a gaseous system?

The stress is a decrease in the partial pressures and concentrations of all gaseous species. The system responds by shifting toward the side with more gas moles to restore pressure.

1. Definition & Core Concepts

Le Châtelier’s Principle is a qualitative tool used to predict the direction in which a chemical reaction will shift when subjected to a change in concentration, pressure, volume, or temperature.
A stress is any change that alters the rate of either the forward or reverse reaction disproportionately, momentarily moving the system away from its equilibrium state.
The system 'shifts' by increasing the rate of one reaction (forward or reverse) until a new equilibrium is reached where the rates are again equal, though the concentrations may differ from the original state.

2. Concentration and Dilution Effects

Addition/Removal: Adding a reactant or removing a product causes the system to shift toward the product side (forward) to consume the excess or replace the loss. Conversely, adding a product or removing a reactant shifts the system toward the reactant side (reverse).
Pure Solids and Liquids: Changes in the amount of pure solids or pure liquids do not affect the equilibrium position because their 'effective concentration' (activity) remains constant and they do not appear in the equilibrium expression.
Dilution: In aqueous systems, adding solvent (dilution) decreases the concentration of all species. The system counteracts this by shifting toward the side with the greater number of aqueous particles to increase the total particle count.

Graph showing a concentration spike (stress) followed by a gradual shift to a new equilibrium level.

3. Pressure and Volume in Gas-Phase Systems

Inverse Relationship: According to Boyle's Law, decreasing the volume of a container increases the total pressure of a gas-phase system.
Volume Decrease (Pressure Increase): The system shifts toward the side of the balanced equation with the fewer moles of gas to reduce the pressure.
Volume Increase (Pressure Decrease): The system shifts toward the side with the greater moles of gas to increase the pressure.
If the number of moles of gas is equal on both sides, changes in volume or pressure have no effect on the equilibrium position.

4. Temperature and the Equilibrium Constant

Heat as a Reactant/Product: In endothermic reactions ( $\Delta H > 0$ ), heat is treated as a reactant. In exothermic reactions ( $\Delta H < 0$ ), heat is treated as a product.
Shifting: Increasing temperature adds heat, shifting the reaction in the endothermic direction. Decreasing temperature removes heat, shifting it in the exothermic direction.
The K Constant: Temperature is the only stress that changes the numerical value of the equilibrium constant ( $K$ ). For an exothermic reaction, increasing $T$ decreases $K$ ; for an endothermic reaction, increasing $T$ increases $K$ .

5. Key Distinctions: Catalysts and Inert Gases

Factor	Effect on Equilibrium Position	Effect on $K$
Catalyst	No Shift (increases forward/reverse rates equally)	No Change
Inert Gas (at constant V)	No Shift (total pressure increases, but partial pressures stay same)	No Change
Temperature	Shifts toward endothermic (if $T$ increases)	Changes
Adding Solid	No Shift	No Change

6. Exam Strategy & Tips

Check States of Matter: Always identify $(s)$ , $(l)$ , $(g)$ , and $(aq)$ before predicting shifts. Ignore solids and liquids when evaluating concentration or pressure changes.
Count Gas Moles: For pressure/volume questions, sum the coefficients of only the gaseous species on each side of the arrow.
Q vs K Comparison: If given concentrations, calculate the reaction quotient $Q$ . If $Q < K$ , the system must shift right (forward) to reach equilibrium. If $Q > K$ , it shifts left (reverse).
Common Mistake: Do not assume a catalyst shifts equilibrium; it only helps the system reach the same equilibrium state faster.

Le Châtelier’s Principle

Summary

1. Definition & Core Concepts

2. Concentration and Dilution Effects

3. Pressure and Volume in Gas-Phase Systems

4. Temperature and the Equilibrium Constant

5. Key Distinctions: Catalysts and Inert Gases

6. Exam Strategy & Tips

Le Châtelier’s Principle

Summary

1. Definition & Core Concepts

2. Concentration and Dilution Effects

3. Pressure and Volume in Gas-Phase Systems

4. Temperature and the Equilibrium Constant

5. Key Distinctions: Catalysts and Inert Gases

6. Exam Strategy & Tips