# Le Châtelier’s Principle Hey guys, welcome to this video on Le Châtelier’s Principle.

Le Châtelier’s Principle deals with how chemical reactions behave when they have reached equilibrium. So before we dive into Le Châtelier’s Principle, let’s review what we mean by equilibrium.

Consider a simple chemical reaction where a moles of A and b moles of B react to form c moles of C.

$$aA + bB -> cC$$

The study of equilibrium is concerned with how much of A and B will be converted to C. In other words, the extent of the reaction. It’s important to keep this concept separate from how fast the reaction happens, which is the study of kinetics.

When we study equilibrium, we give the reaction an unlimited amount of time; just however long it takes for the concentrations of A, B, and C to no longer change. Because all reactions are reversible to some extent (represented by the double-headed arrow), at equilibrium, some A and B are still reacting to form C and some C is reacting to form A and B.

$$aA + bB ⇌ cC$$

Equilibrium is reached when the rate of the forward and the backward reactions are equal.

Rate of $$aA + bB -> cC = Rate of cC -> aA+ bB$$

To reflect the activity taking place in the reaction vessel, we call this a dynamic equilibrium.

This state is described by the equilibrium coefficient, Keq, which reflects the ratio of products and reactants. The system arrives at this ratio of reactants and products by no mistake, but rather through a careful balancing act between enthalpy, entropy, and temperature, simplified as a relationship with the standard free energy change of the system, known as Gibbs free energy.

Once the system arrives at equilibrium, it will stay at this happy balance, unless we disturb it. How the system reacts to disturbance is what Le Châtelier’s principle answers.

Le Châtelier’s principle states that when a system is disturbed, it undergoes a net reaction to reduce the effect of the disturbance and re-attain a state of equilibrium. Hopefully, that feels somewhat intuitive, as we know the equilibrium is the energetically preferred balancing point of the system. In this context, a disturbance is a change in concentration of a product or reactant, a change in pressure, or a change in temperature. Let’s see how a change in concentration affects our theoretical system A, B, and C.

Imagine we’ve mixed A and B and allowed them to react long enough for the reaction to reach equilibrium. But now, instead of leaving the reaction vessel alone, we add in a bit more A, thus throwing the system out of equilibrium. The additional A will react with B to create more C, thus favoring the forward reaction towards the product. The concentrations of A, B, and C will not be the same as before extra A was added, but the system maintains the same equilibrium constant, Keq. This same process would take place if B were added. Conversely, if extra C were added, the equilibrium would favor the reverse reaction and generate more A and B.

Let’s apply this to a real chemical reaction, the production of ammonia (NH3) from hydrogen and nitrogen gases, commonly known as the Haber process.

$$3H2 (g) + N2 (g) ⇌ 2NH3 (g) ΔH = – 92 kJ/mol$$

Our goal is to make as much ammonia as possible. Applying Le Châtelier’s principle, we could add more hydrogen or more nitrogen to achieve this, but in practice, chemists remove ammonia as it is produced to continually push the system to produce more ammonia.

Chemists also shift the equilibrium of this reaction by adjusting the pressure and temperature of the system.

Let’s consider how we could use pressure first. Notice that there are 4 moles of reactants- 3 moles of hydrogen and 1 mole of nitrogen, while there are only 2 moles of ammonia. Thus, as the reaction progresses towards equilibrium at a fixed volume, the pressure will naturally drop because there are fewer moles of gas. Once equilibrium is reached, there will be an associated pressure that reflects the preferred balance of the system. Thus, if we increase the pressure by decreasing the volume of the reaction vessel, the system will work to maintain the equilibrium pressure by producing more ammonia, which reduces the number of moles in the reaction vessel and thus reduces the pressure of the system.

Lastly, let’s consider how we might use temperature to increase the production of ammonia. You’ll notice that the reaction is exothermic (ΔH < 0), so we can think of heat as a product. Therefore, to shift the equilibrium position toward the products, we could remove heat, which is done by keeping the reaction cold. However, we also need to consider the kinetics, or the speed of the reaction. By cooling the reaction down, we shift the equilibrium to create more ammonia, but it slows the reaction down to a prohibitive rate. Thus, in practice, adjusting the temperature isn’t always a straightforward method for shifting the equilibrium. Let’s wrap up by summarizing what we’ve learned. First, we reviewed dynamic equilibrium, which is the natural extent of a chemical reaction. The system reaches a balance of products and reactants, represented by the equilibrium constant, Keq. This is a special balance and reflects the energetics of the system. Le Châtelier’s principle states that if we disturb the system, it will counteract the disturbance to return the system to equilibrium. Not only is this an important concept theoretically, as we saw with the production of ammonia, but practically chemists also use Le Châtelier’s principle to manipulate reactions to yield more of the desired product. Thanks for watching and happy studying!