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 reactants. 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.