In this video, we will learn about how dynamic equilibria respond to changes in temperature, pressure, volume, and composition. We will see how the equilibrium law, otherwise known as Le Chatelier’s principle, helps us to predict the consequences of any of these changes on the position of equilibrium. Before we look at Le Chatelier’s principle, let’s review a few principles of equilibria.
An equilibrium is the state of a reaction where the concentrations of reactants and products are constant over time. Let’s imagine that we have a forward and a reverse reaction. A reacts with B to produce C, and C decays into A and B. If we start with A and B, they’ll react quickly, producing C. As we generate C, the reverse reaction will start to occur. As we build up C, the reverse reaction accelerates further, until eventually the rate of the forward reaction equals the rate of the reverse reaction.
From this point forward, the system is in equilibrium. The concentrations of A, B, and C remain constant. That’s not to say that the concentrations of A, B, and C need to all be the same. We could have more of A and less of C, or we could have more of C and less of A and B. The key thing is that the concentrations are constant over time.
The equilibrium is established when the rate of the forward reaction equals the rate of the reverse reaction. But the definition of equilibrium is only in reference to the concentrations. So conceivably, the rate could be zero, which brings us to the next definition.
A dynamic equilibrium is an equilibrium where the rates of the forward and reverse reactions are nonzero. All this means is that even though the reactions are balanced in their rate, they’re actually happening. The situation where the rate of the forward reaction and the rate of the reverse reaction are equal to zero is called a static equilibrium.
Now we can move on to our last bit of background information. The position of equilibrium isn’t a well-defined property. It’s only really used when comparing one set of equilibrium conditions to another. But we can think about it like the pivot in a seesaw. It’s the point which balances the amounts of product and reactant. If we have more product and less reactant, we say the position of equilibrium moves to the left to keep them in balance. If we have more product and less reactants than we originally had, we say the position of equilibrium moves to the right. Moving to the left favors the reactants. Moving to the right favors the products.
Now that we know what dynamic equilibria are, we can have a look at what consequences change in the temperature has on a dynamic equilibrium. So let’s have a look at a simple, dynamic equilibrium, A plus B in equilibrium with C plus D. We have a mixture of chemicals A, B, C, and D. And while some of A is reacting with some of B to produce some C and some D, some C and some D are reacting together to produce some A and some B. So what happens when we change the temperature?
Let’s assume that we’re heating up the mixture. All the chemicals have more energy. A and B react together more quickly. But so do C and D. You might imagine that this wouldn’t affect the position of equilibrium, but it depends. When chemicals react, there’s usually an energy change. Energy is either released in an exothermic reaction or absorbed in an endothermic reaction.
In a dynamic equilibrium, there are two reactions happening all the time, the forward and the reverse reaction, because these are opposites to each other. One must be exothermic, and the other must be endothermic. Let’s assume for this example that the forward reaction is endothermic and the reverse reaction is exothermic.
Let’s consider the reaction profile of this reaction. This is our change in energy. The activation energy for the exothermic reaction is relatively small. However, the activation energy for the endothermic reaction is quite big. So the activation energy for our endothermic reaction is greater than the activation energy for our exothermic one. An increase in the temperature will favor the endothermic reaction more than the exothermic one. This means that the rate of the endothermic reaction increases more than the rate of the exothermic reaction.
So if this is how our equilibrium looks before we change the temperature, this is how it looks once equilibrium is reestablished after increasing the temperature. The position of equilibrium shifts in the direction of the endothermic reaction products. In this case, it shifts to the right towards C and D. Meanwhile, decreasing the temperature will favor the exothermic reaction. So an increase in temperature will favor the endothermic reaction, the one that absorbs heat, while a decrease in temperature will favor the exothermic reaction, the one that releases heat. In both cases, the position of equilibrium shifts in such a way that the system resists the change.
Now what about pressure and volume changes? Changes in pressure and volume are very similar. So we can look at them together. By increasing the pressure or decreasing the volume, all the components of our system are more tightly packed together. Their concentration increases. If we decrease the pressure or increase the volume, we do the opposite and decrease the particle concentration. Increasing the concentration will increase the rate of any reactions, while decreasing concentration will have the opposite effect and slow down the reactions.
All this ties back to collision theory. The more tightly packed the particles, the more collisions there will be. As with increasing temperature, increasing the pressure or decreasing the volume will accelerate both the forward and reverse reactions. But what if the reactions change the number of particles?
Let’s imagine we have an equilibrium like this, five A in equilibrium with two B. If it takes five A to make two B, then the forward reaction will reduce the number of particles, while the reverse reaction will lead to the opposite, increasing the number of particles. Let’s imagine our initial equilibrium. If we increase the pressure, both the concentration of A and B will go up. But because we have more particles of A, the increase in the frequency of collisions for the forward reaction will be more than for the reverse reaction. This means, in our new equilibrium, we will have a higher proportion of products than we did at the start. So increasing the pressure favors the reaction that reduces the number of particles.
So if this is our starting equilibrium, if we increase the pressure, the position of equilibrium will shift such that the number of particles goes down. If we decrease the pressure, the position of equilibrium shifts such that we have more particles. Generating particles raises the pressure. So in both cases, the position of equilibrium shifts in such a way that the system resists the change. These effects would only be seen where a change in pressure or volume affects the concentration, for instance, with gases. Also, equilibria where the number of particles does not change would not be affected.
Now we can have a look at concentration. So we’ve seen how changing the pressure or volume affects the concentration of all the components of our reaction. But what if we change the concentration of only one component at a time? Let’s go back to our four-component equilibrium. Let’s imagine our system in perfect balance with equal amounts of A, B, C, and D.
Now that we’ve added A, we’ve unbalanced our system. So it’s no longer in equilibrium. However, the extra A accelerates the forward reaction. The initially faster forward reaction generates more C and D. Once the concentration of C and D are at sufficient level, the position of equilibrium is reestablished closer to the products. This means that some of the extra A has been consumed. Adding reactant or product will shift the position of equilibrium to the other side. So adding reactant or removing product will shift the equilibrium in favor of the products. And adding product or removing reactant will shift the equilibrium in favor of the reactants. In both cases, the position of equilibrium shifts in such a way that the system resists the change.
We now have three scenarios which demonstrate the change in the conditions on equilibrium or cause the system to resist the change. The position of equilibrium shifts in whichever direction achieves this effect. This is the foundation of Le Chatelier’s principle.
Le Chatelier’s principle is that, for a dynamic equilibrium, if the conditions change, the position of equilibrium will move to counteract the change. Le Chatelier’s principle is a simple tool that allows us to predict the consequences of a change on an equilibrium without going into the details of the mechanism. So we can see that increasing or decreasing the temperature produces the opposite, and the same for pressure and the concentration of any one component. Now that we’ve learned about Le Chatelier’s principle and its origins, let’s do some practice.
A suspension contains solid Mg(OH)₂ in equilibrium with dissolved Mg²⁺ and OH⁻ ions. If solid MgCl₂ is added to the suspension, which of the following effects is not observed? A) Concentration of Mg²⁺ increases. B) The amount of solid Mg(OH)₂ increases. C) The concentration of OH⁻ increases. D) The concentration of Cl⁻ increases. Or E) the pH decreases.
Suspension refers to fine particles suspended in solvent, usually water. So what we have is an aqueous suspension of magnesium hydroxide. This solid is in equilibrium with dissolved magnesium and hydroxide ions. So here we have our equilibrium equation, solid magnesium hydroxide in equilibrium with magnesium ions and hydroxide ions, both dissolved in solution. The question asks, which of the following options would not happen if we added solid magnesium chloride?
Magnesium chloride is significantly more soluble than magnesium hydroxide. So the magnesium chloride would dissolve. Adding the magnesium chloride would raise the concentration of magnesium ions in solution. We can ignore the chloride ions because they don’t participate in our equilibrium.
Le Chatelier’s principle tells us that, for a dynamic equilibrium, like that between magnesium hydroxide and its ions in solution, the position of equilibrium will shift to counteract a change in the conditions. In this case, the change is an increase in the concentration of magnesium two plus ions. In this case, the position of equilibrium will shift in favor of magnesium hydroxide. Therefore, by adding magnesium chloride, we would expect the amount of solid magnesium hydroxide to increase. Therefore, this is not a correct answer.
What Le Chatelier’s principle misses out is that a shift in the position of equilibrium will never completely counteract the change. So while the position of equilibrium will move away from the reactants, the concentration of magnesium hydroxide will increase overall. Therefore, our correct answer is not that the concentration of magnesium hydroxide increases. We’re looking for something that is not observed. And the concentration of magnesium ions would increase compared to the pure magnesium hydroxide suspension.
What about an increase in hydroxide ion concentration? This shift in the position of equilibrium is in the wrong direction. We are using up hydroxide ions to produce more magnesium hydroxide. Therefore, the hydroxide concentration would actually decrease when we added magnesium chloride. Since an increase in the hydroxide ion concentration would not be observed, this is our correct answer. But let’s have a look at the other options just in case.
The chloride ion concentration would increase when adding magnesium chloride because the magnesium chloride would dissolve. And the last option, the pH decreasing, is a little bit of a tricky one. But because we have less hydroxide ions, we will have a greater concentration of H⁺. Therefore, we’d expect the acidity of the solution to increase and the pH to decrease. So of the five options given, when magnesium chloride is added to a suspension of magnesium hydroxide, the one we would not observe is an increase in the hydroxide ion concentration.
Now that we’ve done a little bit of practice, let’s have a look at the key learning points. Firstly, Le Chatelier’s principle is that, for a dynamic equilibrium, if the conditions change, the position of equilibrium will move to counteract the change. An increase in temperature favors the endothermic reaction. A decrease in temperature favors the exothermic reaction.
Regarding pressure and volume, an increase in pressure or a decrease in volume favors the reaction that reduces the pressure. A decrease in pressure or an increase in volume favors the reaction that increases the pressure. And finally, an increase in the concentration of 𝑥, where 𝑥 is a reactant or product, favors the reaction that decreases the concentration of 𝑥. A decrease in the concentration of 𝑥 favors the reaction that increases the concentration of 𝑥.