What happens if I add an inert gas like argon to the container?

If you add an inert gas at a constant volume, the total pressure increases, but the partial pressures (and concentrations) of the reacting gases do not change. Therefore, the equilibrium does not shift.

Does a shift mean the equilibrium constant (K) changes?

Only a change in temperature will change the actual value of the equilibrium constant, $K$. Changes in pressure, volume, or concentration shift the reaction temporarily, but once the new equilibrium is reached, the ratio of products to reactants (the value of $K$) remains exactly the same as before.

Why don't pure solids affect the equilibrium?

The concentration of a pure solid is determined by its density, which doesn't change regardless of how much of the solid is present. Since its concentration is constant, adding more of it doesn't unbalance the reaction's rates.

What does Le Chatelier's principle predict?

Le Chatelier's principle predicts how a chemical system at equilibrium responds to an outside disturbance. If a reaction at equilibrium experiences a change in concentration, temperature, volume, or pressure, the principle predicts that the system will shift its chemical makeup to counteract the change and establish a new equilibrium.

Think of a chemical equilibrium like a perfectly balanced seesaw. If you add extra weight to the left side, the seesaw tips. To get it back to being level, weight must slide over to the right side. Chemical reactions do the exact same thing when their balance is disturbed.

How concentration changes the balance

When you add more reactant to a system, the seesaw is unbalanced. To relieve this stress, the system reacts by converting some of that extra reactant into product. This is called shifting to the right. Conversely, if you remove a product as it forms, the system will keep shifting right to try and replace it. This is a common trick used in factories to maximize the amount of product a reaction makes.

Pressure and volume shifts

For reactions involving gases, changing the pressure or volume also disturbs the balance. If you decrease the volume of a container, the pressure goes up. The system responds by shifting toward the side of the reaction with fewer moles of gas, effectively reducing the pressure. If you increase the volume, the pressure drops, and the system shifts toward the side with more moles of gas to fill the space.

Temperature acts like a reactant or product

To predict the effect of temperature, it helps to treat heat as a physical part of the reaction. In an endothermic reaction (which absorbs heat), heat is like a reactant. In an exothermic reaction (which releases heat), heat is like a product. If you raise the temperature of an exothermic reaction, it is like adding a product, so the system shifts to the left (making more reactants).

Where students slip up

A common mistake is thinking that any addition to a system causes a shift. Adding a pure solid or a pure liquid does not change the equilibrium position because their concentrations remain constant. Another common trap is catalysts. Adding a catalyst speeds up both the forward and reverse reactions equally. It gets the system to equilibrium faster, but it does not shift the final balance of products and reactants.

Worked through

Consider the exothermic Haber process for synthesizing ammonia: $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g) + ext{heat}$ . Predict which way the equilibrium will shift if you: (a) add more $N_2$ gas, (b) increase the temperature, and (c) increase the volume of the container.

(a) Adding $N_2$ gas increases the concentration of a reactant. The system will shift to the right to consume the extra $N_2$ , producing more $NH_3$ .

(b) Because the reaction is exothermic, heat is treated as a product. Increasing the temperature is like adding more product. The system shifts to the left to consume the extra heat, producing more $N_2$ and $H_2$ .

(c) Increasing the volume decreases the overall pressure. The system will shift toward the side with more moles of gas to restore pressure. The left side has 4 moles of gas (1 mole $N_2$ + 3 moles $H_2$ ), and the right side has 2 moles of gas ( $NH_3$ ). Therefore, the system shifts to the left.

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Where this comes from: OpenStax Chemistry 2e, Chapter 13: Fundamental Equilibrium Concepts · Khan Academy: Le Chatelier's principle