Video Transcript
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.
