Video Transcript
Choosing Equilibrium Conditions
Changes to the environment will
affect how our reaction proceeds. In industrial settings, making the
right choices about a reaction surroundings can help increase yields, eliminate
waste, and save money. In this video, we will learn how to
predict the effect that changing the temperature, concentration, or pressure will
have on the yield of industrial reactions and the financial consequences.
Heres the chemical equation for the
synthesis of ammonia, an important industrial reaction. The double arrow in the middle
means that the reaction is reversible. The compounds on the left-hand side
combine to form the compound on the right. This is called the forward
reaction. But since its a reversible
reaction, at the same time, the reverse reaction occurs where produced ammonia from
the right-hand side of the equation breaks down to form the nitrogen and hydrogen
gases on the left. If we start with only nitrogen gas
and hydrogen gas, we would expect the forward reaction to occur, producing
ammonia. Once the ammonia is produced, we
might then see more of the reverse reaction to produce the initial gases. The concentration of our gases and
our ammonia will change until the reaction reaches equilibrium.
Equilibrium means that the
concentrations of the reactants and products dont tend to change over time. For example, if we combine our
gases and let them sit, we might find that, after a while, theres an extremely high
concentration of ammonia and an extremely low concentration of the original
gases. If these concentrations stayed
constant for a period of time, we would say that the equation has reached
equilibrium. However, just because a reaction
has reached equilibrium does not mean that the reaction has stopped. In situations like this, were
referring to a dynamic equilibrium. There are constant concentrations
of the products and the reactants.
However, thats not because the
reaction isnt proceeding, but rather because the forward reaction and the reverse
reaction are proceeding at the same rate. In other words, at any given
moment, for each nitrogen molecule and three hydrogen molecules that combine to form
two ammonia molecules, somewhere else there are two ammonia molecules that are
breaking down to form one nitrogen molecule and three hydrogen molecules.
As an analogy, we can think of a
dynamic equilibrium sort of like a beach where some people sunbathe and some people
go swimming. As the sunbathers get hot, they may
move to the water to cool off, and as the swimmers get cold, they may move to the
beach to warm up. If the beach has reached dynamic
equilibrium, then the rate of sunbathers entering the water will equal the rate of
swimmers going back to the beach. As a result, the concentration of
sunbathers and swimmers will not change, even though the individual people are
moving back and forth between the two areas.
In this video about choosing
equilibrium conditions, were going to learn about how changes the environment can
shift the equilibrium to the right or to the left. Shifting to the right means
favoring the forward reaction and increasing the concentration of the products. Conversely, a shift left favors the
reverse reaction, resulting in increased concentration of the reactants. In our beach analogy, the Sun
beating down at hot temperatures might cause a shift to the right, driving some of
the sunbathers into the water. Conversely, the presence of a
windsurfing class might make the water more crowded and shift the equilibrium to the
left.
For a chemical reaction, if we
change the concentration, temperature, or pressure of the reaction, we will shift
the equilibrium right or left. Lets take a look at how changing
these conditions can cause the equilibrium to shift. The equilibrium can shift right,
favoring the forward reaction with an increased presence of the products. Or it can shift left, favoring the
reverse reaction with an increased presence of the reactants.
In order to determine which way the
equilibrium will shift, we can look at Le Chateliers principle. This principle says that the
equilibrium of a system will shift to counteract the effect of any applied
constraint. So if there is, say, an increase in
a concentration, the temperature, or the pressure, the equilibrium will shift in the
way that decreases that concentration, the temperature, or the pressure in order to
counteract the effect.
Lets take a look at what this might
mean more specifically. First, if theres an increase in the
concentration of the elements or compounds on one side of the equation, the
equilibrium will shift to the other side. Thinking of our beach analogy, if
the water gets too crowded, some of the swimmers might leave to come to the
beach. In this equation, we can shift the
equilibrium to the right by adding more nitrogen and hydrogen, resulting in the
additional production of ammonia. We could shift the equilibrium to
the left by adding more ammonia, which would then break down to nitrogen and
hydrogen.
If there is a decrease in
concentration, then the equilibrium will shift to the side of the decrease. If we, say, remove ammonia from the
reaction, there will be a lower concentration of ammonia that can be counteracted by
shifting the equilibrium to the right, favoring the forward reaction and producing
more ammonia. For a change in temperature, the
equilibrium shift depends on whether the reaction is endothermic or exothermic. The enthalpy change for this
reaction is negative 92 kilojoules. This negative number means that the
chemicals lose energy, which is then released into the surroundings. A reaction with a negative Δ𝐻 that
loses energy to the surroundings is called an exothermic reaction.
The reverse reaction is just the
same reaction going the other way, so it has the exact opposite Δ𝐻, 92
kilojoules. The reverse reaction absorbs energy
from the surroundings with a positive Δ𝐻, making it endothermic. For all chemical reactions, one
direction will be endothermic and one direction will be exothermic. If we know which direction is
which, we can know how a change in temperature will shift the equilibrium. In this case, an increase in
temperature will shift the equilibrium to the left, favoring the endothermic reverse
reaction. A decrease in temperature shifts
the equilibrium to the right, favoring the exothermic forward reaction.
And how about a change in
pressure? How will that shift the
equilibrium? Well, since different gases take up
the same volumes under the same conditions, we care not so much about the identity
of the gas molecules, but rather the number of gas molecules. An increase in pressure shifts the
equilibrium to the side with fewer gas molecules. In our example equation here, the
left side of the equation has four gas molecules, one nitrogen molecule and three
hydrogen molecules. The right-hand side of the equation
has just two ammonia molecules. An increase in pressure can be
counteracted by having the equilibrium shift to the right. When the four gas molecules on the
left-hand side of the equation combine resulting in two gas molecules, the fewer
number of molecules lowers the gas pressure to counteract the initial increase in
gas pressure.
The opposite is also true; if
theres a decrease in the pressure, then the equilibrium will shift to the side with
more gas molecules. The additional gas molecules
increase the pressure to counteract the initial decrease in pressure. If there are an equal number of gas
molecules on either side of the equation, changing the pressure will not affect the
equilibrium. It is also worth knowing that an
increase in the concentration, temperature, or pressure will increase the rate of
the reaction. With more particles, more pressure,
or more energy, the reactants are more likely to interact and combine to form
products.
Note that the reaction rate and the
equilibrium are two different characteristics. If the equilibrium shifts to the
right and the reaction rate is fast, the product will be produced quite quickly. If the equilibrium shifts to the
right and the reaction rate is very slow, it will still favor the forward reaction,
but much more gradually. Another way to increase the rate of
reaction is by using a catalyst. A catalyst is like a reaction
helper. It helps the reaction proceed with
a lower activation energy without itself changing over the course of the
reaction. Catalysts only speed up the
reaction. They do not affect the
equilibrium.
Now that weve learned the building
blocks of how to affect the equilibrium and the reaction rate, lets take a look at
some applied situations to industrial settings. In industrial applications,
companies can change the equilibrium and reaction rates intelligently to alter the
reactions taking place in the factory. However, they must balance
practical and financial considerations in order to come up with a system that is not
only productive but also not too expensive or difficult to maintain. Each step of the process is
maximized to meet these goals. So taking nitrogen from the air and
hydrogen from natural gas is a relatively inexpensive way to start this
reaction.
Because the production of ammonia
requires three times as many hydrogen atoms as nitrogen atoms, the nitrogen and
hydrogen are combined in a one-to-three ratio by volume and added to a reactor. The reactor is an expensive piece
of equipment as it must be large and be able to withstand high temperatures and
pressures. However, it makes sense for the
factory to spend lots of money on reactors as they are one-time cost that will
continue to produce ammonia day after day. It is more effective to cut costs
on the repeated costs, for example, finding the cheapest way to acquire nitrogen and
hydrogen.
The reaction inside the reactor is
carried out at a very high temperature between 400 and 450 degrees Celsius. This high temperature increases the
rate of the reaction, allowing us to produce ammonia more quickly. However, it does shift the
equilibrium to the left as the reverse reaction is endothermic and the high
temperature favors the endothermic reaction. Although this is counterproductive,
its worth it to increase the rate. And as we will see, the other
conditions will help shift the equilibrium back to the right.
For example, the high pressure of
200 atmospheres increases the rate of the reaction and the yield of the
reaction. We say that something increases the
yield of the reaction when it shifts the equilibrium towards the desired
product. Yield basically means how much is
produced, so one of our goals in this reaction is to increase the yield. Increasing the pressure increases
the yield because it favors the side of the reaction with fewer gas molecules, the
right side of the reaction. So when we increase the pressure,
the equilibrium will shift to the right and the yield will increase.
An iron catalyst is also used in
this reaction in order to increase the rate. As a reminder, catalysts do not
affect the equilibrium and thus do not affect the yield. Therere also choices we can make
outside of the reactor that will affect the rate and yield of the reaction. After leaving the reactor, the
gases can be cooled and separated. The liquid ammonia can be removed
as a successful product of the reaction. Meanwhile, the unreacted hydrogen
and nitrogen gases are fed back into the reactor. Recycling the hydrogen and nitrogen
gas increases the reactant concentration in the reactor, while removing the liquid
ammonia decreases the products concentration. Both of these changes in
concentration have the effect of shifting the equilibrium to the right, increasing
the yield of ammonia.
As we can see, each step of this
process is carefully considered to result in an overall process that safely and
efficiently produces ammonia. This specific process is called the
Haber process, and its used to produce about 100 million tons of ammonia per year
worldwide. This is just one process to produce
one product. In order to minimize waste,
factories will often use multiple overlapping systems where waste products from one
reaction are used as reactants in another reaction or where heat given off by one
reaction is used to heat another reaction. Overall, we can take what we know
about shifting equilibrium and changing reaction rates and couple it with financial
and practical concerns in order to choose the ideal equilibrium conditions for an
industrial process.
Now that weve learned about
choosing equilibrium conditions, lets do some practice problems to review.
In which of the following gas-phase
reactions would the equilibrium yield increase with increasing pressure?
This question is asking about how
the equilibrium will shift when we increase the pressure. When we increase the pressure, the
equilibrium will shift to the side with fewer gas molecules. Le Chateliers principle says that
the equilibrium will shift to counteract an applied constraint. All else equal, fewer gas molecules
will have a lower pressure than more gas molecules. So the equilibrium will shift to
the side with fewer gas molecules in order to lower the pressure to counteract the
initial increase in pressure.
We wanna find the reaction where an
increase in pressure causes the equilibrium yield to increase. Another way of saying an increase
in equilibrium yield is a shift to the right of the equilibrium. When the equilibrium shifts to the
right, the forward reaction is favored and more products are produced. For all of these reactions, when we
increase the pressure, the equilibrium will shift to the side with fewer gas
molecules. For the one choice that is the
correct answer, that side with the fewer gas molecules will be on the right-hand
side. So lets look for the choice with
more gas molecules on the left-hand side of the equation and fewer gas molecules on
the right. Since the question specifically
says that these are gas-phase reactions, we can consider all molecules named as gas
molecules.
The first two choices (A) and (B)
have the same number of gas molecules on the left- and right-hand sides of the
equation, so increasing the pressure will not change the equilibrium. Choices (D) and (E) have more
molecules on the right-hand side and fewer molecules on the left-hand side. For the reactions in choice (D) and
(E), increasing the pressure will cause the equilibrium to shift to the left,
decreasing the equilibrium yield. Finally, choice (C) has three
molecules on the left-hand side of the equation, 2NO2 molecules and a CH3OH
molecule, and only two molecules on the right-hand side. For choice (C), an increase in
pressure will cause the equilibrium to shift to the right-hand side of the equation,
increasing the equilibrium yield as that is the side with fewer gas molecules.
So in which of the following
gas-phase reactions would the equilibrium yield increase with increasing
pressure? Well, thats choice (C) 2NO2 plus
CH3OH reversibly produces CH3ONO plus HNO3.
Now that weve done some practice,
lets review the key points of the video. A chemical reaction reaches dynamic
equilibrium when the rate of the forward reaction equals the rate of the reverse
reaction. According to Le Chateliers
principle, the equilibrium will shift to the right or to the left to counteract a
constraint. For example, if we increase the
temperature, the equilibrium will shift to favor the endothermic reaction. A decrease in temperature will
favor the exothermic reaction. If we increase the pressure, the
equilibrium will shift to the side of the equation with fewer gas particles. Decreasing the pressure shifts the
equilibrium to the side with more gas particles.
An increase in concentration of
some of the compounds or elements on one side of the equation can be counteracted by
shifting the equilibrium to the other side of the equation. An increase in the temperature,
pressure, or concentration also increases the reaction rate. Another way to increase the
reaction rate is with the addition of a catalyst. In order to run an effective
operation, factories must increase rate and yield. But when doing so, they must also
factor in the cost and safety of the choices they make.
