The Haber Process
Nitrogen Fixation, Fertiliser, and the Chemistry That Fed and Armed the World
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Chapters explain how iron catalysts make the reaction practical, while others cover ruthenium alternatives and what poisons can stop the process. The book traces the development from Fritz Haber's early work through large-scale industrial implementation, including economic costs and environmental impacts.
Anyone interested in how chemistry feeds the world while also making weapons will find this thorough explanation of nitrogen fixation worth their time.
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During the 19th century, people needed more nitrogen fertilizers to grow crops, and they got it mainly by mining niter deposits and collecting guano from tropical islands. By the beginning of the 20th century, those supplies were running low, and scientists started looking for new ways to make ammonia. Atmospheric nitrogen makes up about 78% of the air, but it's very stable and doesn't react easily with other substances.
Haber worked with his assistant Robert Le Rossignol to build the high-pressure equipment and catalysts needed to prove the process in the lab. In the summer of 1909, they successfully made ammonia from air, producing about 125 milliliters an hour. The German company BASF bought the idea and assigned Carl Bosch the job of making it work on a large scale. Bosch managed to do this by 1910. Haber was later awarded the Nobel Prize in 1918, and Bosch received his in 1931, both for solving the tough problems of turning the lab process into a working industrial system.
In 1913, the Haber process began producing ammonia on an industrial scale at BASF's Oppau plant in Germany, reaching 20 tonnes a day by 1914. During World War I, Germany needed vast quantities of nitrate for making munitions. The Allies controlled the sea lanes and had access to major sodium nitrate deposits in Chile and India, both under British influence. Even though German companies may have held nominal legal rights to these resources, the blockade prevented them from reaching Germany. The Haber process became so vital to the German war effort that it’s widely believed Germany would have been defeated within months without it. Synthetic ammonia from this method was used to make nitric acid, which was essential for producing explosives.
In the early days of the Haber–Bosch process, the reaction chambers used osmium as a catalyst, but it was so rare that it was nearly impossible to obtain in any meaningful quantity. Haber himself observed that uranium worked almost as well and was far easier to get hold of. Then, in 1909, a researcher at BASF named Alwin Mittasch made a breakthrough with an iron-based catalyst that was much more affordable. Gerhard Ertl played a major role in the development of this catalysis. Today, the most widely used catalysts still rely on iron, enhanced with oxides of potassium, calcium, silicon, and aluminum.
During the interwar years, several new methods for producing ammonia were developed. The Casale process and the Claude process appeared alongside the Mont-Cenis process, which was created by the Friedrich Uhde Ingenieurbüro. Luigi Casale and Georges Claude recommended increasing the pressure in the synthesis loop to 80–100 megapascals, raising the single-pass ammonia conversion and making it possible to liquefy ammonia at ambient temperatures. Claude also suggested using three or four converters in series with liquefaction steps, eliminating the need for recycling. Although some plants have moved away from the original Haber process, most continue to use it, operating at 20 megapascals and 500 degrees Celsius, though with improved efficiency and reduced energy consumption thanks to better catalysts and process design.
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Ammonia production uses a lot of energy, about 1 to 2 percent of all global energy use, and it sends out 3 percent of carbon emissions worldwide. The process needs hydrogen and nitrogen from the air, and most of that hydrogen comes from natural gas through a method called steam reforming. In 2012, 72 percent of ammonia globally was made this way. But in China, by 2022, coal was responsible for 75 percent of hydrogen production, while natural gas accounted for only 20 percent. Hydrogen can also be made from water and electricity using electrolysis, as was done at the Hydro plant in Vemork, which supplied most of Europe’s ammonia at one time. Steam reforming remains the cheapest way to make large amounts of hydrogen right now.
In 2012, Hideo Hosono's group reported ammonia synthesis using a catalyst called Ru-loaded calcium-aluminium oxide C12A7:e− electride, which they described as an electron donor and reversible hydrogen store. Following up on that work, researchers studied how nitrogen dissociates and how the catalyst functions in this process. This method is now being used in a small plant in Japan for ammonia production. Then, in 2019, Hosono’s team discovered another catalyst—a perovskite oxynitride-hydride called BaCeO3−xNyHz—that operates at lower temperatures and doesn’t require the costly ruthenium used in earlier versions.
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The process involves hydrogen reacting with nitrogen—extracted from air through separation—to create anhydrous liquid ammonia. It's a challenging and costly procedure. Lower temperatures slow down the reaction rate, making it less efficient. At the same time, the need for high pressure means that the equipment must be built from materials strong enough to handle the stress without becoming brittle from hydrogen exposure.
Nitrogen gas, or diatomic nitrogen, is incredibly stable because the two nitrogen atoms are held together by a strong triple bond. This makes it hard to break apart and use in chemical reactions. To make this happen, scientists use catalysts that help speed up the process of breaking those bonds. Since not all the gases react completely, the unreacted materials are sent through the system again to get the most out of the process.
The process uses a specific type of catalyst, often made of magnetite and promoted with other elements, to help nitrogen and hydrogen react into ammonia. The reaction mixture passes through four separate beds of this catalyst. Between each pass, the gas is cooled to keep the reaction at an effective equilibrium. Because the catalyst is sensitive, the gases must be very pure—especially free of oxygen-containing compounds like carbon monoxide, carbon dioxide, and water vapor. Nitrogen can be separated from air, but extra steps may be needed to remove any remaining oxygen.
On each pass through the reactor, only about 15% of the gases convert into ammonia, so the product must be separated out and the remaining gases sent back through the system to continue reacting. Even with this low conversion rate, the process eventually achieves a 97% overall conversion because the unused reactants are recycled. But that constant recycling causes inert gases to build up in the mixture over time.
The Haber Process hinges on two key factors: where the chemical equilibrium lies and how fast the reaction happens. At room temperature, the balance naturally favors ammonia formation, but the reaction is so slow it's practically undetectable because of the high energy barrier it must overcome. Since the process releases heat—meaning it’s exothermic—increasing the temperature pushes the equilibrium back toward the original reactants, per Le Châtelier’s principle. By around 150 to 200 degrees Celsius, the equilibrium constant drops to one, meaning the forward and reverse reactions proceed at equal rates.
The Van 't Hoff equation shows that at temperatures higher than a certain point, the balance of the reaction shifts in a way that makes ammonia production less effective when pressure remains at normal levels. Yet reducing the temperature isn't helpful either, since the catalyst used in the process only works well if it reaches at least 400 °C.
The Haber Process uses high pressure to boost ammonia production, since four moles of reactants turn into two moles of product, shifting the balance toward ammonia. The pressure applied—between 15 and 25 megapascals, or roughly 150 to 250 bar—is strong enough to change how much ammonia forms. This pressure makes the reaction more efficient, increasing yield significantly. The equilibrium relationship shows why this works, though the exact numbers aren’t listed here. Still, it’s clear that pressure plays a key role in making ammonia on an industrial scale.
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The Haber Process operates under specific conditions that are crucial to its success. The earlier steps—steam reforming, shift conversion, carbon dioxide removal, and methanation—all function at pressures between 25 and 35 bar. Then, the actual ammonia synthesis happens at temperatures ranging from 300 to 500 degrees Celsius, with pressure levels between 60 and 180 bar, depending on the method used. After the reaction, the ammonia must be separated from leftover hydrogen and nitrogen gases, which requires cooling down to minus 20 degrees Celsius.
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The Haber–Bosch process uses catalysts to speed up the reaction that combines nitrogen and hydrogen gases. These catalysts are solid materials that don’t get used up in the process. They help make ammonia from these gases more quickly than it would happen on its own. The catalysts work by letting the gas molecules stick to their surface, which changes how they react. This method was developed to make fertilizers on a large scale. It also had major consequences for both food production and warfare during the twentieth century.
The catalyst in the Haber Process is built from iron that’s been ground into tiny pieces and bonded to a base of iron oxide. That base might also include extra components called promoters. These promoters can be made up of several substances: aluminium oxide, potassium oxide, calcium oxide, potassium hydroxide, molybdenum, or magnesium oxide.
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The iron catalyst used in the Haber Process starts as finely ground iron powder, typically made by reducing high-purity magnetite. The iron is then oxidized to form particles of magnetite or wüstite, which are partially reduced to create a core-shell structure. At the center lies magnetite, surrounded by a shell of wüstite, and an outer layer of metallic iron. This structure retains most of its volume during reduction, forming a highly porous material with a large surface area that boosts its effectiveness. Support materials like calcium and aluminium oxides are added to help maintain the catalyst’s surface area, and these oxides—along with those of potassium and silicon—are unreactive to hydrogen reduction.
The catalyst used in the process is made through a specific melting method, where the raw materials must be free of anything that could poison the catalyst. The promoter elements need to be evenly mixed within the magnetite melt. The mixture is heated to around 3500 degrees Celsius and then rapidly cooled. This quick cooling creates the necessary starting material for the final product. However, this fast-cooling step leads to a catalyst that doesn't stand up well to wear and tear. Even with this drawback, the rapid cooling technique is frequently used.
In the production plant, magnetite is reduced directly to α-iron using synthesis gas. The process begins with magnetite forming a shell of wüstite around its core. Further reduction leads to α-iron, which combines with promoters to form the outer layer. These reactions are complex and depend on temperature: at lower temperatures, wüstite breaks down into iron and magnetite phases; at higher temperatures, the reduction of wüstite and magnetite to iron becomes dominant.
The α-iron structures themselves are made up of primary crystallites that are roughly thirty nanometers in size. These crystallites give rise to a bimodal pore system. About ten nanometers in diameter, those pores come from the reduction of the magnetite phase. The larger pores, ranging from twenty-five to fifty nanometers, originate instead from the reduction of the wüstite phase. Except for cobalt oxide, the promoters stay in their original state throughout the process.
When iron oxide is reduced using synthesis gas, water vapor is produced, and this must be carefully managed to maintain catalyst quality. If the water comes into contact with finely divided iron, it causes premature aging through recrystallization—especially under high temperatures. To prevent this, the water vapor pressure in the gas mixture during catalyst production is kept as low as possible, with target values below three grams per cubic meter. That’s why the reduction process happens at high gas exchange rates, low pressure, and relatively low temperatures. The reaction itself is exothermic, meaning it releases heat, which leads to a gradual temperature rise during ammonia formation.
It takes four to ten days for a new catalyst to reach full production capacity. The wüstite phase reduces more quickly and at lower temperatures than magnetite. Studies using kinetics, microscopy, and X-ray spectroscopy revealed that wüstite first reacts to form metallic iron. This creates a gradient of iron(II) ions that move from magnetite through wüstite to the surface, where they form iron nuclei. A highly active new catalyst based on this process was discovered in the 1980s at Zhejiang University of Technology and later commercialized by 2003.
Pre-reduced, stabilized catalysts make up a large portion of the market. They come ready with their full pore structure already formed, but after being made, they're oxidized again on the surface, which means they’re no longer pyrophoric. Reactivating these catalysts only takes thirty to forty hours instead of several days. Besides the quick start-up time, they also offer benefits like better water resistance and less weight.
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Many attempts have been made to improve the Haber–Bosch process, with various metals tested as catalysts. The key requirement is that the catalyst must allow nitrogen molecules to split into atoms when they stick to the surface. If the binding is too strong, the catalyst gets blocked—a problem called self-poisoning—and this happens with elements left of iron on the periodic table. Chromium is one example where forming surface nitrides makes it ineffective. Metals right of iron, however, don’t hold nitrogen strongly enough for ammonia synthesis. Haber initially tried osmium and uranium-based catalysts; uranium forms its nitride during the reaction, while osmium oxide is rare.
Studies, both theoretical and practical, show that boosting the performance of iron alone has its limits. Still, adding cobalt to iron catalysts does raise their effectiveness.
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Ruthenium is used to make highly effective catalysts, enabling chemical reactions to proceed under less extreme conditions of pressure and temperature. These are classified as second-generation catalysts, and one method of preparing them involves decomposing triruthenium dodecacarbonyl on a graphite surface. When ruthenium is supported on activated carbon, a challenge arises: the carbon can react with hydrogen to produce methane. The efficiency of such catalysts is influenced significantly by both the support material and any additives included. A variety of materials have been used as supports, from carbon and magnesium oxide to aluminium oxide, zeolites, spinels, and even boron nitride.
Since 1992, the KBR Advanced Ammonia Process, or KAAP, has used carbon-based catalysts activated with ruthenium. The carbon carrier in these catalysts breaks down into methane, but this issue can be reduced by treating the carbon at 1500 degrees Celsius, which helps extend the catalyst's life. Still, the finely dispersed carbon creates an explosion risk. Because of that, along with ruthenium’s low acidity, magnesium oxide has emerged as a better carrier choice. Carriers with acidic traits pull electrons away from ruthenium, reducing its reactivity and causing ammonia to stick to the surface instead of reacting freely.
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Catalyst poisons reduce the effectiveness of catalysts used in the Haber process, and they usually come as unwanted impurities in the synthesis gas. These poisons can be permanent or temporary. Permanent ones, like sulfur compounds, phosphorus compounds, arsenic compounds, and chlorine compounds, cause a lasting drop in catalytic activity. Temporary poisons, such as water, carbon monoxide, carbon dioxide, and oxygen, only lower activity while they're present. Understanding how these substances affect the process is key to maintaining efficient nitrogen fixation for fertiliser production.
In the Haber Process, certain components of the gas mixture don’t technically count as poisons, but they still mess with the reaction. Noble gases and methane, for example, build up over time as the process recycles its gases. Because these inert parts accumulate, they reduce how much space the actual reacting gases have. That drop in pressure slows down how fast the reaction moves forward.
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The reaction involved in the Haber Process is an exothermic equilibrium reaction where the gas volume decreases. The equilibrium constant K for this process can be determined from data presented in a table. This chemical balance plays a key role in how nitrogen and hydrogen combine to form ammonia, under specific conditions of temperature and pressure that affect the outcome of the reaction.
Because the reaction gives off heat, lowering the temperature pushes the equilibrium toward making more ammonia. The reaction also involves a decrease in volume—four parts of reactants yield two parts of product—so increasing pressure helps drive the formation of ammonia, as described by Le Chatelier’s principle. High pressure is essential to keep the catalyst surface covered with nitrogen. To achieve this balance, engineers use a ratio of one part nitrogen to three parts hydrogen, operate at pressures between 250 and 350 bar, maintain temperatures around 450 to 550 degrees Celsius, and employ α iron as the catalyst.
The catalyst used in the process is ferrite, or α-Fe, which forms inside the reactor when magnetite reacts with hydrogen. It works best at around 400 to 500 degrees Celsius. Even though it helps break the strong nitrogen triple bond, the reaction still needs high temperatures to happen fast enough. In industry, the temperature is set between 450 and 550 degrees Celsius to balance how well the catalyst works against the risk of ammonia breaking down again. As the ammonia forms, it's constantly taken out of the system. The final gas mixture contains about twenty percent ammonia.
In the Haber Process, it’s important to keep the amount of inert gases like argon low, because too much of them reduces the pressure of the reacting nitrogen and hydrogen. To manage this, some of the gas mixture is continuously removed from the system. The argon is then separated using a gas separation plant that employs what's known as the Linde process. This method allows for the extraction of pure argon from the circulating gas, helping to maintain optimal conditions for ammonia production.
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Modern ammonia plants can produce over 3000 tons of ammonia per day from a single production line. The design shown here comes from a "single-train" Haber–Bosch plant built in the early 1960s by Kellogg. This setup represents one of the largest-scale chemical processes ever developed, turning nitrogen and hydrogen into fertiliser on an industrial scale.
The synthesis gas used in the process comes from various sources, and before it can be used, it must be cleaned of impurities like hydrogen sulfide and organic sulfur compounds, which interfere with the reaction. When the gas contains high levels of hydrogen sulfide—often found in synthesis gas from carbonization coke—it’s treated in a wet cleaning stage, such as the sulfosolvan process. Lower concentrations are removed instead by passing the gas through activated carbon. Organosulfur compounds are separated using pressure swing adsorption, often after carbon monoxide has been converted to carbon dioxide.
To make hydrogen for the Haber process, methane gas reacts with water vapor in a reactor called the primary reformer. A nickel oxide-alumina catalyst helps the reaction happen. The result is carbon monoxide and hydrogen gas. This reaction needs energy—specifically, it takes 206 kilojoules of enthalpy per mole to complete.
The methane gas doesn’t fully react in the first reactor. To get more hydrogen and keep the methane levels low, the leftover methane gets converted in a second step using oxygen inside a secondary reformer. Air provides the oxygen needed for this process. At the same time, nitrogen is added to the gas mixture, which will be used later in making ammonia.
In the third step of the process, carbon monoxide gets converted into carbon dioxide. This reaction is known as CO conversion or the water–gas shift reaction. It’s a key part of making ammonia on a large scale, which then leads to fertilizers that feed billions. The Haber Process relies on several chemical steps, and this one ensures the right balance of gases. Without it, the production of nitrogen-based fertilizers would be much less efficient. The reaction helps remove impurities so that the final product can be used effectively in agriculture and beyond. It’s a crucial stage in turning raw materials into something that changes the world.
In the process, carbon monoxide and carbon dioxide would react with ammonia to form carbamates, which appeared as solid deposits. These solids quickly clogged up the pipelines and equipment. This problem arose during a later step in the process.
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An energy diagram shows how much energy is needed or released during each step of a reaction. For the Haber Process, the diagram reveals why the homogeneous gas-phase reaction won’t work—because breaking the nitrogen bond requires too much energy. A catalyst solves this by making the process possible through surface binding, where the energy gained from attaching nitrogen atoms outweighs the cost of splitting them. Still, the step where nitrogen breaks apart on the catalyst surface remains the slowest part—not because of high activation energy, but due to a poor pre-exponential factor in the rate equation. Hydrogenation is endothermic, but that energy comes easily from the reaction temperature, around 700 Kelvin.
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When the Haber process was first developed, it had to compete with another method of making fertilizer called the cyanamide process. The cyanamide process needed a lot more electricity to run and required more workers than the Haber process did. That difference in energy use and labor made the Haber process more attractive for large-scale production.
As of 2018, the Haber process was producing two hundred and thirty million tonnes of anhydrous ammonia every year. That ammonia is used mainly as nitrogen fertilizer, either as ammonia itself, in the form of ammonium nitrate, or as urea. The process uses three to five percent of the world’s natural gas production, which equals about one to two percent of global energy use. Combined with improvements in crop breeding, herbicides, and pesticides, these fertilizers have greatly increased how much food can be grown on agricultural land.
If crop yields had stayed at the level they were in 1900, the harvest in the year 2000 would have needed almost four times more land than what was actually used. The area under cultivation would have consumed nearly half of all ice-free land on Earth, instead of the less than 15% of total land that is required today.
The Haber–Bosch process, which makes fertilizer from nitrogen and hydrogen, is extremely energy-intensive. This energy use contributes to climate change and other environmental issues. Nitrogen from this process leaks into groundwater and waterways, causing eutrophication and expanding dead zones in oceans. Atmospheric nitrates and ammonia from the process harm natural ecosystems. The buildup of reactive nitrogen in the biosphere disrupts the nitrogen cycle in ways that are largely human-made. This process also increases emissions of nitrous oxide, a greenhouse gas that ranks as the third largest contributor to global warming after carbon dioxide and methane.
Because nitrogen use efficiency is typically less than 50%, much of the industrial nitrogen applied to farms ends up in runoff. This excess nitrogen disrupts natural biological habitats, altering the balance of ecosystems that depend on careful nutrient levels. The impact isn't just environmental—it affects the long-term health of the land and waterways that support agriculture and wildlife alike.
The Haber–Bosch process set off a chain reaction that reshaped humanity itself. Nearly 50% of the nitrogen in human tissues comes from it, making it central to the dramatic rise in global population. In 1900, there were only 1.6 billion people on Earth. By November 2018, that number had soared to 7.7 billion. This chemical breakthrough didn’t just feed more people—it enabled the explosive growth of civilizations and fundamentally altered the course of history.
Reverse fuel cell technology represents a breakthrough in how ammonia is made, using electric energy, water, and nitrogen together—without needing a separate step to split hydrogen from oxygen first. This method streamlines the process, making it more efficient by removing the need for an additional electrolysis stage. It’s a development that could reshape how we produce this vital chemical, especially in contexts where energy use and resource efficiency are key concerns. The technology builds on existing knowledge of fuel cells but applies it in a new direction to tackle nitrogen fixation directly. By integrating these steps, it offers a potential pathway toward more sustainable production methods.
Using synthetic nitrogen fertilizers takes away the reason for farmers to rely on better methods like rotating crops that include legumes, since those naturally add nitrogen to the soil. When farmers depend on artificial fertilizers, they don’t need to use crop rotation practices that are more sustainable in the long term. This shift toward chemical fertilizers changes how farming is done and can reduce the natural nitrogen-fixing benefits that come from certain plants. It also affects the overall approach to managing farmland and its resources.
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