Video: The Haber Process

In this video, we will learn how to describe the reaction of nitrogen and hydrogen in the Haber Process and explain how the reaction conditions are chosen.

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Video Transcript

In this video, we’ll learn about the Haber process, the production of ammonia from nitrogen and hydrogen. And the production of ammonia is on the order of 200 million metric tons a year. And the majority of it is produced by the Haber process, also called the Haber–Bosch process. Nitrogen and hydrogen gasses react together to form ammonia in an exothermic reaction. The six nitrogen–hydrogen bonds in the products released more energy when they form than the energy required to break the nitrogen≡nitrogen triple bond and the three hydrogen–hydrogen single bonds.

The formation of ammonia from nitrogen and hydrogen is reversible. So in a sealed reactor, an equilibrium is quickly established. For the Haber process, the position of equilibrium depends on the temperature and the pressure. Le Chatelier’s principle states that, for a dynamic equilibrium, if the conditions change, the position of equilibrium will move to counteract the change. A dynamic equilibrium is a state of a reaction where the forward and reverse reactions are both happening and at the same rate so that the concentrations stay constant. If we increase the temperature, the system will respond. But in what way?

We already know that there are two reactions going on, the reaction of ammonia and hydrogen producing ammonia and the decomposition of ammonia producing nitrogen and hydrogen. The forward reaction is exothermic, so it releases energy. The reverse reaction is endothermic, so it absorbs energy. So the exothermic reaction will cause an increase in the temperature and the endothermic reaction will cause a decrease in the temperature. So according to Le Chatelier’s principal, increasing the temperature will favor the endothermic reaction, shifting the position of equilibrium towards the reactants. So increasing the temperature will favor the endothermic reaction and shift the position towards the reactants. And increasing the temperature decreases the yield of ammonia. Of course, decreasing the temperature produces the opposite, favoring the exothermic forward reaction, shifting the position of equilibrium to the right, and increasing the equilibrium yield.

Now let’s have a look at pressure changes. In our reaction, all the components are gases. If we don’t change anything else, adding more particles of gas will increase the pressure. We can make the simplification that the pressure of the gas is proportional to the number of gas particles. If we double the number of particles, we’ll double the pressure. So what happens to our equilibrium when we increase the pressure? According to Le Chatelier’s principle, the position of equilibrium will shift in favor of the reaction that reduces the pressure. In the forward reaction four gas particles ⁠— one nitrogen, three hydrogen ⁠— react to form two gas particles, two ammonia molecules.

In the reverse reaction, two gas particles react to form four gas particles. So the forward reaction reduces the pressure, while the reverse reaction increases the pressure. So increasing the pressure will favor the forward reaction and shift the position of equilibrium towards the products. Increasing the pressure increases the yield of ammonia. Meanwhile, decreasing the pressure shifts the position of equilibrium to the left and decreases the equilibrium yield. Now, let’s have a look at how these factors impact the Haber process.

In the Haber process, nitrogen is mixed with hydrogen at high pressure and temperature in the presence of a catalyst. The catalysts used today are mostly based on iron. Nitrogen is extracted from the air, while hydrogen is produced from small hydrocarbons, like those in natural gas or in light of fractions of crude oil. One lot of nitrogen and three lots of hydrogen are mixed together in the mixer. Once mixed together, that passed into the compressor. Once inside, the mixture is taken up to the pressure closer to what it needs to be to react. The mixture is passed to the reactor, where it’s heated over catalyst in order to produce ammonia. However, what comes out of the reactor still contains a lot of nitrogen and hydrogen.

So the mixture is passed into a cooler. This cools the mixture until the ammonia liquefies. And the remaining nitrogen and hydrogen is recycled back into the compressor. The Haber process is usually performed around 450 degrees Celsius and at 200 atmospheres. Under these conditions, the equilibrium yield for the reaction is about 30 percent. If the nitrogen and hydrogen are recycled back into the process, we can achieve a practical yield of about 97 percent. Now that we’ve had a look at the overall process, let’s consider why these particular conditions are used in industry.

The Haber process is performed at about 450 degrees Celsius. That’s about 200 degrees Celsius more than a typical oven. The pressure for the Haber process is 200 atmospheres. That’s 200 times the pressure in Earth’s atmosphere. The choice of conditions used in the Haber process is a complex compromise between lots of factors. Companies want to make ammonia as quickly as possible as cheaply as possible. Increase in the temperature increases the rate of the forward and reverse reactions, so equilibrium is achieved more quickly. That’s good because ammonia is produced more quickly. However, increasing the temperature decreases the equilibrium yield. That’s bad because that limits the maximum amount of ammonia we can produce in one cycle.

Increasing the pressure also reduces the time it takes to get to equilibrium. That’s great, so increasing the pressure increases the rate of production of ammonia. And unlike temperature, increasing the pressure actually increases the equilibrium yield. That’s excellent, so increasing the pressure means that ammonia is produced more quickly and increases the maximum possible amount per cycle. Using a catalyst accelerates both reactions equally. So the equilibrium yield stays the same. But the time it takes to get to equilibrium is lower. If we only cared about the rate of production ammonia, we’d run the Haber process at a medium temperature at super high pressure and with the most effective catalyst. However, the other factor that needs to be considered is cost.

Heating up a reaction is relatively cheap. But we’ve seen that the best temperature will be a compromise between reducing the time it takes to reach equilibrium and still having a decent equilibrium yield. High pressures are much more expensive. The higher the pressure, the stronger the reactors and pipes need to be and the more dangerous it is. Hydrogen is flammable and potentially explosive. Imagine having hydrogen at many times atmospheric pressure. That’s going to be quite a hazard.

There are better catalysts in the iron-based catalyst used today. But modern catalysts do an excellent job for the cost. All these factors balance out with the ideal conditions, economically speaking, being about 450 degrees Celsius, 200 atmospheres, and iron-based catalyst and cooling of the product mixture and recycling of the reactants. Now that we’ve learned all about the Haber process, the equilibrium, and all the factors that affect it, let’s have some practice.

In a model of the Haber process, 28 grams of nitrogen gas is mixed with an excess of hydrogen gas and the mixture is heated under high pressure with an iron-based catalyst. If the yield of the reaction is 12 percent, what is the mass of product produced?

The question tells us that we’re dealing with something like the Haber process, which makes ammonia, NH₃. Our starting material is nitrogen gas, N₂, which is being combined with hydrogen gas, H₂. These are reacting to produce ammonia which is probably going to be in the gas phase, although it doesn’t really matter for this question. So our first task for this question is to balance the equation. We can see we have two equivalence of nitrogen atoms on the left-hand side and one equivalent on the right, with two equivalents of hydrogen on the left and three on the right. If we double up the amount of ammonia, we’ll have two equivalents of nitrogen on both sides. And we can balance the hydrogen atoms by adding two more hydrogen molecules to the reactants.

Now, you may have noticed that I’m dealing with a reaction arrow, not an equilibrium arrow. That’s deliberate because we’re looking at the yield of the reaction. And I don’t want to confuse this with the equilibrium yield. The yield of a reaction is the amount of product divided by the maximum amount, which is the amount we would get if the reaction had a yield of 100 percent. This is all multiplied by 100 percent which in this case results in a value of 12 percent. The process we’re going to go through is to calculate the number of moles that ammonia would produce at 100 percent yield, then calculate the number of moles of product at 12 percent yield, and then convert that amount into the mass of the product.

We start off by working out the amount of nitrogen in 28 grams of nitrogen gas. We do this by taking the mass and dividing it by the molar mass. Let’s start by working out the molecular mass of a nitrogen gas molecule. This is equal to the atomic mass of nitrogen multiplied by two. Using the periodic table, we can see that the atomic mass of nitrogen is 14.007 unified atomic mass units, so the molecular mass of nitrogen is 14.007 times two u. This gives us a molecular mass of 28.014 u. This means that the molar mass of our nitrogen molecule is 28.14 grams per mole. So the amount of nitrogen gas we have is equal to its mass, 28 grams, divided by its molar mass. 28.014 grams per mole. This gives us about 0.1 moles of nitrogen gas.

The question tells us that hydrogen is in excess. So the reactant that limits the amount of ammonia we can produce is nitrogen gas. So we can move on to calculating the amount of ammonia we produce. We can calculate the amount of ammonia by taking the amount of nitrogen and multiplying it by two ammonia molecules per nitrogen molecule. We get that value from the stochiometric coefficients in the balanced equation. So the amount of ammonia we get is 0.9950 moles of nitrogen multiplied by two ammonia molecules per nitrogen molecule. This gives us about two moles of ammonia if we got 100 percent yield.

Now, we’ve worked out the amount of ammonia we produce at 100 percent yield. We multiply that by 12 percent to get the yield at 12 percent. This gives us a final yield of 0.23988 moles of ammonia. If you want to multiply it by a percentage in your calculator and you don’t have a percentage abundant, you can multiply by 12 over 100 or 0.12. All that remains is to calculate the mass of ammonia at this percentage yield. The mass of ammonia is equal to the number of moles multiplied by its molar mass. With one nitrogen atom and three hydrogen atoms, the molar mass of ammonia is 14.007 plus three times 1.008 grams per mole, which works out at 17.031 grams per mole.

So our massive product is equal to 0.23988 moles of ammonia multiplied by 17.31 grams per mole of ammonia, giving us a final value for our mass of 4.0854 grams. All the values in our question aren’t given to two significant figures. So we should give our answer to the same precision. So if 28 grams of nitrogen gas is reacted with excess of hydrogen gas with a yield of 12 percent, the mass of product produced will be 4.1 grams.

The Haber process involves the reversible reaction of two gaseous reactions. What is the main disadvantage of increasing the temperature at which the reaction is performed?

Haber process is an industrial method to produce ammonia via the reaction of nitrogen and hydrogen. Nitrogen and hydrogen are the two gaseous reactants referred to in the question. And this is the balanced equation for their reaction in the Haber process. The question asks what the main disadvantage of increasing the temperature is. Increasing the temperature of the reaction is always going to cost a little more. But that’s probably not going to be the main disadvantage referenced here. To solve this question, we’re going to need to recall what the entropy change of the reaction is, or at least what its sign is.

The entropy change for the forward reaction is negative, meaning the reaction of nitrogen and hydrogen is exothermic. It might interest you to know that the entropy change for the forward reaction is actually minus 92.4 kilo joules per mole. But we don’t need to know this value in order to answer the question. Le Chatelier’s principle says that, for a dynamic equilibrium, if the conditions change, the position of equilibrium will move to counteract the change. In this case, the change is an increase in temperature. For this reaction, since the forward reaction is exothermic, the reverse reaction must be endothermic. An endothermic reaction is a cooling process, which reduces the temperature. So increasing the temperature favors the reverse reaction.

So if we take our initial equilibrium and heat it up, our position of equilibrium will move to the left, favoring the reactants. So increasing the temperature reduces our equilibrium yield. So we get less ammonia at equilibrium. So the main disadvantage of increasing the temperature is that we reduce the maximum percentage yield. Now, we’ve looked at the whole of the Haber process on done a few examples. Let’s review the key points.

The Haber process is the industrial production of ammonia from nitrogen and hydrogen. The Haber process is done at around 450 degrees Celsius, around 200 atmospheres, using an iron-based catalyst and recycling of reactants to achieve a high yield and a low cost. Because of the nature of the equilibrium reaction, the equilibrium yield increases as temperature decreases or pressure increases. And the time taken to achieve equilibrium decreases as temperature increases, pressure increases, or when a catalyst is used. These factors are all balanced to achieve the highest yield at the lowest cost.

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