# Video: Choosing Equilibrium Conditions

In this video, we will learn how to predict the effect of changing the temperature, concentration, or pressure can have on the yield of industrial reactions and the financial consequences.

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### 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.