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© Jim Clark 2002 (modified April 2013)

The Haber Process

Core concepts.

In this article, you will learn about the Haber Process and its importance, as well as its kinetics, thermodynamics, and mechanisms.

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What is the Haber Process?

The Haber Process is an industrial method of reacting nitrogen and hydrogen gas to produce ammonia. Stoichiometrically , this involves one mole of diatomic nitrogen and three moles of diatomic hydrogen per two moles of ammonia.

Haber Process Equation: N 2 + 3H 2 → 2NH 3

Many chemists call the Haber Process a form of “nitrogen fixation,” meaning the conversion of nitrogen gas to a more chemically useful form. In nature, microorganisms called diazotrophs perform nitrogen fixation , which helps cycle nitrogen through ecological systems. Nitrogen, of course, serves as one of the fundamental elements of biological life, forming an important part of proteins and nucleic acids . 

Nitrogen, in the form of ammonia, also has importance in research and industrial applications. However, despite diatomic nitrogen serving as the most abundant atmospheric gas, its lack of reactivity made it hard for chemists to harness. 

In 1909, German chemist Fritz Haber made scientific history when he demonstrated synthesizing ammonia from atmospheric nitrogen and hydrogen. Later in 1911, the method was significantly improved by another German chemist Carl Bosch, which made the reaction more efficient for industry. As a result, when performed today, many refer to the method as the Haber-Bosch Process. 

Due to the work of Haber and Bosch, ammonia can now be readily synthesized for a variety of important purposes, such as fertilizer, explosives, refrigeration technology, and pharmaceuticals. 

But what exactly makes the Haber Process so efficient? Let’s take a closer look at the physical chemistry behind the reaction.

The Kinetics of the Haber Process

If nitrogen gas is so inert, how did Fritz Haber manage to get it to react? To get past the unreactivity of nitrogen, the Haber Process involves two important factors. The first factor is an effective catalyst . 

Transition metals are almost always used to catalyze the Haber Process. Most often, industrial chemists use magnetite, an iron oxide (Fe 3 O 4 ), with smaller amounts of potassium hydroxide (KOH). Other common catalysts include chromium and osmium oxides.

Mechanistically, the catalyst must bind to nitrogen gas. This allows hydrogen to add across the nitrogen bonds until the nitrogen fully saturates, yielding ammonia. Without this crucial binding of the catalyst, nitrogen remains inert and unreactive, producing virtually no ammonia.

The second factor to get nitrogen to react relates to pressure . Specifically, higher pressures indicate more molecular collisions, which increases the rate of reaction. As a result, industries tend to perform the Haber Process at the highest pressures possible, often higher than 200 atmospheres.

The Thermodynamics of the Haber Process

With or without a catalyst, the reaction producing ammonia is spontaneous (∆G < 0). This is due to a fundamental aspect of physical chemistry: the kinetics and thermodynamics of a reaction remain often independent of each other. Though we may intuit that a highly-thermodynamically favored reaction occurs faster than one less favored, there exist exceptions. The Haber Process serves as one such exception; a reaction with favorable thermodynamics but incredibly slow kinetics. 

Aside from the spontaneity, we can also observe the reaction has a negative change in entropy (∆S < 0). Specifically, we see four moles of reactant gas producing two moles of product gas.

N 2 + 3H 2 → 2NH 3

A reduction in the moles of gas indicates that our system has fewer possible microstates, indicating decreasing entropy. We can also justify the negative change by saying fewer gas molecules means more “disorder”.

For our reaction to be spontaneous despite the decreasing entropy, our reaction must have decreasing enthalpy . This is indeed the case: the Haber Process is exothermic (∆H < 0). According to the definition of Gibbs free energy , spontaneity must increase as temperature decreases. Put differently, decreasing temperature makes ∆S less significant, making ∆G more negative:

∆G = ∆H – T∆S

(In this equation, T indicates the temperature at which the reaction takes place)

This has important implications for the ideal temperature at which to perform the Haber Process, as we observe in the next section.

The Equilibrium Properties of the Haber Process

As we just covered, the Haber Process is exothermic and increases spontaneity at lower temperatures. With this information, chemists can increase the yield of the reaction by lowering temperatures. The explanation for this phenomenon involves the principles of chemical equilibrium.

Due to exothermicity, we could say that heat serves as a “product” of the reaction. According to L’Chatelier’s Principle , we can shift the equilibrium of a reaction mixture by placing stress on the system. The reaction then responds in the direction to counteract the stress. By cooling the mixture, we “remove” heat, and the system counteracts this stress by producing more products. Thus, we get more ammonia under lower temperatures. 

However, if we decrease the temperature too much, the kinetics of our reaction decrease since the molecules don’t move as fast. To compromise, industrial chemists usually set the reaction temperature of the Haber Process at 450-500°C.

The Full Haber Process

Now that we know the physical chemistry, let’s take a look at the full process. 

One thing you may notice is that dihydrogen gas is not directly inputted into the system. Instead, methane (1) and water (2) react to produce hydrogen within the system (3). Later, oxygen and nitrogen gas enter the system (4), which further generates hydrogen from methane (5). In both reactions, carbon monoxide forms as a by-product.

Further water vapor enters the system, oxidizing carbon monoxide to carbon dioxide (6). This carbon dioxide then leaves (8), producing a mixture of nitrogen and hydrogen gas. 

This gas mixture becomes compressed to more than 200atm (7) and pre-heats to ~500°C (9). Then, the heated and pressurized mixture enters the reaction chamber, fitted with the transition metal catalyst (10). A certain amount of ammonia is produced, which is distilled by running a cold water jacket over the reaction mixture (11). This condenses the ammonia gas to a liquid, which drains out of the system (13).

Importantly, on the first run, only about 15% of the nitrogen and hydrogen will have reacted. The unreacted gas separates from the liquid ammonia (12) and enters the reaction chamber again. Eventually, about 98% of the original gas mixture reacts to form ammonia after enough runs.

Edexcel GCSE Chemistry

The Haber Process

The Haber process is a reversible chemical reaction between nitrogen and hydrogen gas to produce ammonia. The chemical equation for this process is:

N 2 (g) + 3H 2 (g) ⇌ 2NH 3 (g)

Ammonia is a very important fertiliser. Before the Haber process was invented, fertiliser was mostly obtained from seabird droppings and by mining potassium nitrate. However, as the global population grew, the demand for fertiliser began to exceed the supply from these natural sources.

The Haber process, developed in the early 1900s by German chemists Fritz Haber and Carl Bosch, revolutionised agriculture. This method allows for the production of ammonia — an effective fertiliser — using just nitrogen and hydrogen, two abundant gases.

Today, the Haber process generates over 200 million tons of fertiliser per year, which is used for growing the crops that feed millions worldwide.

• Ammonia also serves as a key raw material in the explosives and dyes industries.

Nitrogen, which makes up 78.1% of the air we breathe, is sourced directly from the atmosphere for the Haber process. Hydrogen is typically extracted from natural gas or other sources. The production of ammonia is a complex process, as nitrogen gas is highly stable and not very reactive.

Steps in the Ammonia Production Process

Let’s look at each step of the process:

Conditions for the Haber Process

Ammonia is a product of high commercial interest, so optimising its production is important to save costs, time, and resources. This involves regulating temperature, pressure, and catalysts used. The following conditions have been found to be optimal:

Producing ammonia, a valuable commercial product, involves optimising several conditions to save costs, time and resources. The optimal conditions are:

• Temperature : 450ºC
• Pressure : 200 atmospheres
• Catalyst : Iron

Let’s look at each factor.

Temperature

The ideal temperature for the Haber process is 450°C. The forward reaction is:

N 2 (g) + 3H 2 (g) → 2NH 3 (g)

This reaction is exothermic, meaning it releases heat. According to Le Chatelier’s Principle, more ammonia could potentially be produced at lower temperatures since the equilibrium would shift to absorb the excess heat. However, lower temperatures slow the reaction rate. Therefore, a balance is necessary between maximising ammonia yield and maintaining a practical rate of reaction.

A temperature of 450°C provides this balance, optimising the rate of reaction while maximising the yield of ammonia at dynamic equilibrium.

The pressure used in the Haber process is approximately 200 atmospheres (atm). Higher pressure speeds up the reaction because the gases are compressed into a smaller space, increasing their chances of interacting. However, creating high pressures is costly and can be dangerous if equipment fails.

A pressure of 200 atm is an optimal compromise, allowing for efficient ammonia production while minimising costs and safety risks.

Iron catalyst

The catalyst in the Haber process is typically made of iron. Iron is effective at accelerating the rate of reaction. It is ground into small, porous particles to maximise the surface area available for the reaction. This increased surface area improvesthe efficiency of the process.

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The Haber Process: A Simple Discovery that Changed the World

Mukhtar Quraishi, Year 12, Clitheroe Royal Grammar School, Lancashire

The Haber Process seems like a simple discovery. The process takes nitrogen and hydrogen to form ammonia and is seemingly unimportant. It feels like a backyard experiment gone wrong. After all, ammonia is a toxin to sea life and is even a common by-product of reactions discarded by our bodies. So why is it that the Haber Process is such a well-known reaction, how has it changed our lives and how does it affect the world today?

During the early 1900s, ammonia was in high demand for fertilisers across the globe. Although nitrogen is highly abundant, forming almost 80% of the air around us, it is an extremely stable gas and does not readily react with other chemicals. Converting N2 into ammonia posed a challenge for chemists globally. The Haber-Bosch process, named after its German inventors Fritz Haber and Carl Bosch, was able to produce ammonia on a large scale at a higher efficiency than ever before. The process was ground-breaking, and the duo earned two Nobel Prizes, in 1918 and 1931, for their work on large scale, chemical processes. The Haber Process now produces 450 million tonnes of ammonia, most of which is used to produce ammonium nitrate for fertiliser. The abundance of synthetic fertiliser has led to a decrease in guano harvesting around the world, consequently leading to fewer seabirds becoming extinct and less damage to cave ecosystems (since guano trade often caused bats to leave their roost). Another consequence of the Haber Process is how efficient agriculture has become.

“With average crop yields remaining at the 1900 level, the crop harvest in the year 2000 would have required nearly four times more land, and the cultivated area would have claimed nearly half of all ice-free continents, rather than under 15% of the total land area that is required today.” (Smil, Vaclav – 2011). Although this seems like a wonderful thing, the ability to supply so much produce has led to the ‘population boom’ over the past century. Shockingly, the population has increased by almost 5x since the 1900s. In fact, “nearly 50% of the nitrogen found in human tissues originated from the Haber-Bosch Process” (Solomon, P. M. – 2004). The Haber Process has directly led to the situation the world is in, from the greenhouse emissions due to the much larger carbon footprint of our generation, to the laws put in place limiting childbirth and (arguably) the rights of humans all around the world.

However, ammonia isn’t just used for fertiliser. During the First World War, Germany required high deposits of nitrates to form explosives. The Allies had access to large sodium nitrate deposits in Chile but Germany had nothing. The Haber Process allowed the Germans to produce weapons from thin air, arguably helping Germany with their journey to the Second World War as well as directly killing many during the first. During World War 2, the Nazis used hydrogen cyanide (Zyklon B) to murder minorities in the gas chambers and used ammonia to safely neutralise the toxic gas. These examples may be from long ago, but ammonium-nitrate based explosives are still used all around the world today. They were used in the Sterling Hall bombing in Wisconsin, 1970; the Oklahoma City bombing, 1995; the Delhi and Oslo bombings, 2011; and even as recently as 2013 during the Hyderabad blasts. The Haber Process and the ability to easily produce ammonia has not only resulted in cheaper yields but has directly caused deaths all around the world ever since its discovery.

The Haber Process has also contributed to problems in the environment, many of which affect us today and which we have to face the direct consequences of. For example, the Haber Process uses tonnes of energy, consuming around 1-2% of the world energy’s supply. It also contributes to a build-up of reactive nitrogen in the biosphere, causing potential harm to the nitrogen cycle. As well as this, runoff from the very popular nitrogen fertilisers can disrupt biological habitats with nitrates leaching into rivers, ponds and lakes alongside expanding dead zones in oceans and affecting natural ecosystems. It is now the third most important greenhouse gas, only succeeded by carbon dioxide and methane.

To conclude, the Haber-Bosch Process most certainly affects us. It has helped extensively with our crop production and allowed us to use less land and destroy fewer habitats. Nevertheless, ammonium-based explosives are in abundance due to the Haber Process and ammonia production has had many detrimental effects on the environment, leaving us to clean up the Haber Process’ problems today.

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Chemistry 30

Chemical equilibrium.

Assignment 4: Applications of Chemical Equilibrium: The Haber Process This worksheet is also available in the following formats:  Word | RTF | PDF

Teacher Resources

Other links, the haber process.

For this assignment you will research the Haber Process, an important industrial application of equilibrium.

Begin by finding at least five different sources of information about this process. You may use textbooks, the Internet, or library books. List these references on your assignment in a Bibliography section.

You have several options regarding how you submit your research. Your teacher may wish to assign different marks depending on your method of presentation. For example, simply answering the questions posed below may not earn you as many marks as a formal report.

• Prepare a formal report. All reports should be typed, double-spaced. Use the questions provided below to guide your research and presentation.
• Answer the following questions regarding the Haber Process, but not as a formal report. This assignment must be typed.
• Create a web page about the Haber process. The questions asked below should be answered in your web pages. Be sure to include a list of resources used on your web page.
• A poster or series of posters that provide answers to the following questions.
• A multimedia presentation (Powerpoint) about the Haber Process.
• A short video, podcast, or radio play explaining the Haber Process.
• Who developed the Haber Process? When? What country was he from?
• Write a balanced chemical equation for this reaction, including the energy term. Is it an endothermic or exothermic reaction?
• Use Le Châtalier’s Principle to explain the conditions that favour the forward reaction.
• Temperature is important in regulating this reaction. Is the reaction carried out at high or low temperatures? How does this relate to part of your answer for Question 3?
• Under what temperature and pressure conditions is this reaction typically carried out?
• What catalyst is used for this reaction?
• Provide a short paragraph providing some historical background. Why was this an important chemical process?
• Is the Haber Process still important today? Explain.

• General Principles and Processes of Isolation of Elements

Haber Process

The Haber process provides a good case study to illustrate how industrial chemists use their knowledge of the factors that affect chemical equilibria to find the best conditions needed to produce a good yield of products at a reasonable rate.

Table of Contents

• What is Haber Process?

Explaining the Process and Conditions

Reaction rate and equilibrium, uses of ammonia, what is haber process.

Haber–Bosch process or just Haber process is basically one of the most efficient and successful industrial procedures to be adopted for the production of ammonia. German chemists Fritz Haber along with his assistant in the 20th century developed high-pressure devices and catalysts to carry out the process on a laboratory scale.

Later, Carl Bosch in the year 1910 took the design and created a machine for industrial-level production. This was indeed an important development in the field of science.

Let us look at and understand the process below.

In the Haber process, “ the atmospheric nitrogen (N 2 ) is converted to ammonia (NH 3 ) by reacting it with hydrogen (H 2 )”. Here a metal catalyst is used and high temperatures and pressures are maintained.

The raw materials for the process are

• Air, which supplies the nitrogen.
• Natural gas and water supply the hydrogen and the energy needed to heat the reactants.
• Iron is the catalyst and does not get used up.

Let us take a look at the diagram below.

• As per the diagram, in the Haber process, we take nitrogen gas from the air and combine it with hydrogen atoms obtained from natural gas in the ratio 1:3 by volume.
• The gases are passed through four beds of catalyst, with cooling taking place in each pass. This is done to maintain equilibrium constant.
• While different levels of conversion occur in each pass where unreacted gases are recycled.
• Normally an iron catalyst is used in the process, and the whole procedure is conducted by maintaining a temperature of around 400 – 450 o C and a pressure of 150 – 200 atm.
• The process also involves steps like shift conversion, carbon dioxide removal, steam reforming, and methanation.
• In the final stage of the process, the ammonia gas is cooled down to form a liquid solution which is then collected and stored in storage containers.

The Haber process for the synthesis of ammonia is based on the reaction of nitrogen and hydrogen. The chemical reaction is given below. Notably, in this process, the reaction is an exothermic reaction one where there is a release of energy.

N 2 (g) + 3H 2 (g) → 2NH 3 (g)

Nitrogen in the reaction is obtained by separating nitrogen from the air through liquefaction and hydrogen is obtained from natural gas by steam reforming.

CH 4 (g) + H 2 O → H 2 (g) + CO(g)

According to Le Chatteleir principle, the production of ammonia is favoured by high pressure and low temperature. The Haber process is typically carried out at pressures between 200 and 400 atmospheres and at  temperature of 500 o C. In the commercial production of ammonia, NH3 is continuously removed as it is produced. Removing the products causes more nitrogen and hydrogen to combine according to Le Chatelier’s principle.

The reaction is a reversible reaction. However, the reaction is affected by changes in temperature, pressure and catalyst used mainly in the composition of the equilibrium mixture, the rate of the reaction and the economics of the whole process.

Ammonia which is produced is one product that is essential in many areas. Some uses of ammonia include;

• Agricultural uses: Production of ammonia is important as it is one of the main components in making fertilizers.
• Explosives: Ammonia produced is used in making nitro-based explosives including TNT, RDX, etc.
• Pharma: It is used in manufacturing certain types of drugs such as sulfonamide, antimalarials, and vitamins such as thiamine and nicotinamide.
• Refrigeration: It is also used in large-scale refrigeration plants, air-conditioning units in Buildings, etc.
• Consumer Products: Ammonia is used in various cleaning products and acts as an effective cleaning agent.

Frequently Asked Questions – FAQs

How is ammonia manufactured by haber’s process.

Production of ammonia by the cycle of Haber. It is widely provided by the nitrogen (N2) and hydrogen (H2) Haber cycle. The Haber process takes nitrogen gas from the atmosphere and combines it to form ammonia gas with molecular hydrogen gas.

Why is iron catalyst used for Haber process?

Iron is used in the Haber cycle as a cheap catalyst. It allows in acceptable time to reach a reasonable yield. State three conditions of reaction regulated in industrial reactions.

How do we get hydrogen for Haber process?

Methane from natural gas is the main source of hydrogen. In a high-temperature and -pressure pipe inside a reformer with a nickel catalyst, the process, steam reforming, is carried out by separating the carbon and hydrogen atoms in the natural gas.

What factors affect the Haber process?

The yield of ammonia can be changed by increasing the pressure or temperature of the reaction because the Haber cycle is a reversible reaction. Increasing the reaction pressure increases ammonia yield.

Why is Haber process important?

Today, the Haber process is still necessary because it produces ammonia, which is vital for fertilizer and many other purposes. Every year, the Haber cycle produces around 500 million tons of fertilizer (453 billion kilograms). This fertilizer helps feed about 40% of the population of the world.

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5.4: The Haber Process

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The Haber Process is used in the manufacturing of ammonia from nitrogen and hydrogen, and then goes on to explain the reasons for the conditions used in the process. The process combines nitrogen from the air with hydrogen derived mainly from natural gas (methane) into ammonia. The reaction is reversible and the production of ammonia is exothermic.

\[\ce{ N2(g) + 3H2(g) <=> 2NH3 (g)} \label{eq1}\]

with \(ΔH=-92.4 kJ/mol\).

A flow scheme for the Haber Process looks like this:

General Conditions of the Process

• The catalyst: The catalyst is actually slightly more complicated than pure iron. It has potassium hydroxide added to it as a promoter - a substance that increases its efficiency.
• The pressure : The pressure varies from one manufacturing plant to another, but is always high. You cannot go far wrong in an exam quoting 200 atmospheres.
• Recycling: At each pass of the gases through the reactor, only about 15% of the nitrogen and hydrogen converts to ammonia. (This figure also varies from plant to plant.) By continual recycling of the unreacted nitrogen and hydrogen, the overall conversion is about 98%.

Composition

The proportions of nitrogen and hydrogen: The mixture of nitrogen and hydrogen going into the reactor is in the ratio of 1 volume of nitrogen to 3 volumes of hydrogen. Avogadro's Law says that equal volumes of gases at the same temperature and pressure contain equal numbers of molecules. That means that the gases are going into the reactor in the ratio of 1 molecule of nitrogen to 3 of hydrogen. That is the proportion demanded by the equation.

In some reactions you might choose to use an excess of one of the reactants. You would do this if it is particularly important to use up as much as possible of the other reactant - if, for example, it was much more expensive. That does not apply in this case. There is always a down-side to using anything other than the equation proportions. If you have an excess of one reactant there will be molecules passing through the reactor which cannot possibly react because there is not anything for them to react with. This wastes reactor space - particularly space on the surface of the catalyst.

Temperature

• Equilibrium considerations : You need to shift the position of the equilibrium (Equation \(\ref{eq1}\)) as far as possible to the right in order to produce the maximum possible amount of ammonia in the equilibrium mixture. The forward reaction is exothermic with \(ΔH=-92.4 kJ/mol\). According to Le Chatelier's Principle , this will be favored if you lower the temperature. The system will respond by moving the position of equilibrium to counteract this - in other words by producing more heat. To get as much ammonia as possible in the equilibrium mixture, you need as low a temperature as possible. However, 400 - 450° C is not a low temperature!
• Rate considerations : The lower the temperature you use, the slower the reaction becomes. A manufacturer is trying to produce as much ammonia as possible per day. It makes no sense to try to achieve an equilibrium mixture which contains a very high proportion of ammonia if it takes several years for the reaction to reach that equilibrium. You need the gases to reach equilibrium within the very short time that they will be in contact with the catalyst in the reactor.
• The compromise: 400 - 450°C is a compromise temperature producing a reasonably high proportion of ammonia in the equilibrium mixture (even if it is only 15%), but in a very short time.

Notice that there are 4 molecules on the left-hand side of Equation \(\ref{eq1}\), but only 2 on the right. According to Le Chatelier's Principle, if you increase the pressure the system will respond by favoring the reaction which produces fewer molecules. That will cause the pressure to fall again. In order to get as much ammonia as possible in the equilibrium mixture, you need as high a pressure as possible. 200 atmospheres is a high pressure, but not amazingly high.

• Rate considerations: Increasing the pressure brings the molecules closer together. In this particular instance, it will increase their chances of hitting and sticking to the surface of the catalyst where they can react. The higher the pressure the better in terms of the rate of a gas reaction.
• Economic considerations: Very high pressures are very expensive to produce on two counts. You have to build extremely strong pipes and containment vessels to withstand the very high pressure. That increases your capital costs when the plant is built. High pressures cost a lot to produce and maintain. That means that the running costs of your plant are very high.
• The compromise : 200 atmospheres is a compromise pressure chosen on economic grounds. If the pressure used is too high, the cost of generating it exceeds the price you can get for the extra ammonia produced.
• Equilibrium considerations: The catalyst has no effect whatsoever on the position of the equilibrium. Adding a catalyst doesn't produce any greater percentage of ammonia in the equilibrium mixture. Its only function is to speed up the reaction.
• Rate considerations: In the absence of a catalyst the reaction is so slow that virtually no reaction happens in any sensible time. The catalyst ensures that the reaction is fast enough for a dynamic equilibrium to be set up within the very short time that the gases are actually in the reactor.
• Separating the ammonia : When the gases leave the reactor they are hot and at a very high pressure. Ammonia is easily liquefied under pressure as long as it is not too hot, and so the temperature of the mixture is lowered enough for the ammonia to turn to a liquid. The nitrogen and hydrogen remain as gases even under these high pressures, and can be recycled.

By mixing one part ammonia to nine parts air with the use of a catalyst, the ammonia will get oxidized to nitric acid .

\[\begin{align*} \ce{4 NH_3} + \ce{5 O_2} &\rightarrow \ce{4 NO} + \ce{6 H_2O} \\[4pt] \ce{2 NO} + \ce{O_2} &\rightarrow \ce{2 NO_2} \\[4pt] \ce{2 NO_2} + \ce{2 H_2O} &\rightarrow \ce{2 HNO_3} + \ce{H_2} \end{align*}\]

Haber process

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• Q 1 / 16 Score 0 What is the name of the compound produced in the Haber process? 29 ammonia

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• Q 1 What is the name of the compound produced in the Haber process? ammonia 30 s
• Q 2 What makes this process difficult is that the reaction is ................? reversible 30 s
• Q 3 The two reactants needed for the Haber process are nitrogen and ...............? hydrogen 30 s
• Q 4 The main use for ammonia is in the production of artificial ....................? fertilisers 30 s
• Q 5 We need artificial fertilisers to increase crop .............? yield 30 s
• Q 6 The nitrogen needed for the Haber process is obtained by fractional distillation of ............? air 30 s
• Q 7 The hydrogen needed for the Haber process is obtained from ...............? methane 30 s
• Q 8 Methane is obtained from crude oil, which cannot be replaced, so it is non .....................? renewable 30 s
• Q 9 The balanced equation for the Haber process reaction shows that one mole of nitrogen is needed to make 2 moles of ammona, NH3. How many moles of hydrogen react with the one mole of nitrogen? 3 30 s
• Q 10 The reaction conditions need to be controlled carefully to maintain a high yield of ammonia. A catalyst is used to reduce the time taken for ammonia to appear. The catalyst used is ..............? iron 30 s
• Q 11 A very high pressure would favour the production of ammonia, because the equilibrium would shift to the side with less particles. Very high pressures can be expensive because the reactor walls have to be very ..............., to reduce the risk of explosion. thick 30 s
• Q 12 High temperature favours a fast reaction, but the reverse reaction in the Haber process is endothermic. This would lower the yield of ammonia as the heat is absorbed by the reaction. We describe the actual temperature used as a ....................? compromise 30 s
• Q 13 If the reverse reaction in the Haber process is endothermic, the forward reaction must be ...................? exothermic 30 s
• Q 14 In a dynamic equilibrium, the forward and reverse reaction rates are ................? equal 30 s
• Q 15 In a closed system or container, a dynamic equilibrium can be achieved. We would expect, at this point, the ................................. of reactants and products to be equal. concentration 30 s

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