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The Haber process,^[1] also called the Haber–Bosch process, is the main industrial procedure for the production of ammonia.^[2][3] The German chemists Fritz Haber and Carl Bosch developed it in the first decade of the 20th century. The process converts atmospheric nitrogen (N₂) to ammonia (NH₃) by a reaction with hydrogen (H₂) using an iron metal catalyst under high temperatures and pressures. This reaction is slightly exothermic (i.e. it releases energy), meaning that the reaction is favoured at lower temperatures^[4] and higher pressures.^[5] It decreases entropy, complicating the process. Hydrogen is produced via steam reforming, followed by an iterative closed cycle to react hydrogen with nitrogen to produce ammonia.

The primary reaction is:

${\displaystyle {\ce {N2 + 3H2 <=> 2NH3}}\qquad {\Delta H^{\circ }=-92.28\;{\ce {kJ}}}\ ({\Delta H_{298\mathrm {K} }^{\circ }=-46.14\;\mathrm {kJ/mol} })}$

Before the development of the Haber process, it had been difficult to produce ammonia on an industrial scale,^[6][7][8] because earlier methods, such as the Birkeland–Eyde process and the Frank–Caro process, were too inefficient.

History

During the 19th century, the demand for nitrates and ammonia for use as fertilizers and industrial feedstocks rapidly increased. The main source was mining niter deposits and guano from tropical islands.^[9] At the beginning of the 20th century these reserves were thought insufficient to satisfy future demands,^[10] and research into new potential sources of ammonia increased. Although atmospheric nitrogen (N₂) is abundant, comprising ~78% of the air, it is exceptionally stable and does not readily react with other chemicals.

Haber, with his assistant Robert Le Rossignol, developed the high-pressure devices and catalysts needed to demonstrate the Haber process at a laboratory scale.^[11][12] They demonstrated their process in the summer of 1909 by producing ammonia from the air, drop by drop, at the rate of about 125 mL (4 US fl oz) per hour. The process was purchased by the German chemical company BASF, which assigned Carl Bosch the task of scaling up Haber's tabletop machine to industrial scale.^[7][13] He succeeded in 1910. Haber and Bosch were later awarded Nobel Prizes, in 1918 and 1931 respectively, for their work in overcoming the chemical and engineering problems of large-scale, continuous-flow, high-pressure technology.^[7]

Ammonia was first manufactured using the Haber process on an industrial scale in 1913 in BASF's Oppau plant in Germany, reaching 20 tonnes/day in 1914.^[14] During World War I, the production of munitions required large amounts of nitrate. The Allies had access to large deposits of sodium nitrate in Chile (Chile saltpetre) controlled by British companies. India had large supplies too, but it was also controlled by the British.^[15] Moreover, even if German commercial interests had nominal legal control of such resources, the Allied powers controlled the sea lanes and imposed a highly effective blockade which would have prevented such supplies from reaching Germany in any case. The Haber process proved so essential to the German war effort^[7][16] that it is considered virtually certain Germany would have been defeated in a matter of months without it. Synthetic ammonia from the Haber process was used for the production of nitric acid, a precursor to the nitrates used in explosives.

The original Haber–Bosch reaction chambers used osmium as the catalyst, but it was available in extremely small quantities. Haber noted uranium was almost as effective and easier to obtain than osmium. In 1909, BASF researcher Alwin Mittasch discovered a much less expensive iron-based catalyst that is still used. A major contributor to the elucidation of this catalysis was Gerhard Ertl.^{[17][18][19][20]} The most popular catalysts are based on iron promoted with K₂O, CaO, SiO₂, and Al₂O₃.

During the interwar years, alternative processes were developed, most notably the Casale process, Claude process, and the Mont-Cenis process developed by Friedrich Uhde Ingenieurbüro.^[21] Luigi Casale and Georges Claude proposed to increase the pressure of the synthesis loop to 80–100 MPa (800–1,000 bar; 12,000–15,000 psi), thereby increasing the single-pass ammonia conversion and making nearly complete liquefaction at ambient temperature feasible. Claude proposed to have three or four converters with liquefaction steps in series, thereby avoiding recycling. Most plants continue to use the original Haber process (20 MPa (200 bar; 2,900 psi) and 500 °C (932 °F)), albeit with improved single-pass conversion and lower energy consumption due to process and catalyst optimization.

Process

Combined with the energy needed to produce hydrogen and purified atmospheric nitrogen, ammonia production is energy-intensive, accounting for 1% to 2% of global energy consumption, 3% of global carbon emissions,^[22] and 3% to 5% of natural gas consumption.^[23] Hydrogen required for ammonia synthesis is most often produced through gasification of carbon-containing material, mostly natural gas, but other potential carbon sources include coal, petroleum, peat, biomass, or waste. As of 2012, the global production of ammonia produced from natural gas using the steam reforming process was 72%.^[24] Hydrogen can also be produced from water and electricity using electrolysis: at one time, most of Europe's ammonia was produced from the Hydro plant at Vemork. Other possibilities include biological hydrogen production or photolysis, but at present, steam reforming of natural gas is the most economical means of mass-producing hydrogen.

The choice of catalyst is important for synthesizing ammonia. In 2012, Hideo Hosono's group found that Ru-loaded calcium-aluminum oxide C12A7:e⁻ electride works well as a catalyst and pursued more efficient formation.^[25][26] This method is implemented in a small plant for ammonia synthesis in Japan.^[27][28] In 2019, Hosono's group found another catalyst, a novel perovskite oxynitride-hydride BaCeO_3−xN_yH_z, that works at lower temperature and without costly ruthenium.^[29]

Hydrogen production

The major source of hydrogen is methane. Steam reforming extracts hydrogen from methane in a high-temperature and pressure tube inside a reformer with a nickel catalyst. Other fossil fuel sources include coal, heavy fuel oil and naphtha.

Green hydrogen is produced without fossil fuels or carbon dioxide emissions from biomass, water electrolysis and thermochemical (solar or another heat source) water splitting. However, these hydrogen sources are not economically competitive with steam reforming.^[30][31][32]

Starting with a natural gas (CH
₄) feedstock, the steps are as follows;

• Remove sulfur compounds from the feedstock, because sulfur deactivates the catalysts used in subsequent steps. Sulfur removal requires catalytic hydrogenation to convert sulfur compounds in the feedstocks to gaseous hydrogen sulfide (hydrodesulfurization, hydrotreating):

${\displaystyle {\ce {H2 + RSH -> RH + H2S}}}$

• Hydrogen sulfide is adsorbed and removed by passing it through beds of zinc oxide where it is converted to solid zinc sulfide:

${\displaystyle {\ce {H2S + ZnO -> ZnS + H2O}}}$

• Catalytic steam reforming of the sulfur-free feedstock forms hydrogen plus carbon monoxide:

${\displaystyle {\ce {CH4 + H2O -> CO + 3 H2}}}$

• Catalytic shift conversion converts the carbon monoxide to carbon dioxide and more hydrogen:

${\displaystyle {\ce {CO + H2O -> CO2 + H2}}}$

• Carbon dioxide is removed either by absorption in aqueous ethanolamine solutions or by adsorption in pressure swing adsorbers (PSA) using proprietary solid adsorption media.
• The final step in producing hydrogen is to use catalytic methanation to remove residual carbon monoxide or carbon dioxide:

${\displaystyle {\ce {CO + 3 H2 -> CH4 + H2O}}}$

${\displaystyle {\ce {CO2 + 4 H2 -> CH4 + 2 H2O}}}$

Ammonia production

The hydrogen is catalytically reacted with nitrogen (derived from process air) to form anhydrous liquid ammonia. It is difficult and expensive, as lower temperatures result in slower reaction kinetics (hence a slower reaction rate)^[33] and high pressure requires high-strength pressure vessels^[34] that resist hydrogen embrittlement. Diatomic nitrogen is bound together by a triple bond, which makes it relatively inert.^[35][36] Yield and efficiency are low, meaning that the ammonia must be extracted and the gases reprocessed for the reaction to proceed at an acceptable pace.^[37]

This step is known as the ammonia synthesis loop:

${\displaystyle {\ce {3 H2 + N2 -> 2 NH3}}}$

The gases (nitrogen and hydrogen) are passed over four beds of catalyst, with cooling between each pass to maintain a reasonable equilibrium constant. On each pass, only about 15% conversion occurs, but unreacted gases are recycled, and eventually conversion of 97% is achieved.^[3]

Due to the nature of the (typically multi-promoted magnetite) catalyst used in the ammonia synthesis reaction, only low levels of oxygen-containing (especially CO, CO₂ and H₂O) compounds can be tolerated in the hydrogen/nitrogen mixture. Relatively pure nitrogen can be obtained by air separation, but additional oxygen removal may be required.

Because of relatively low single pass conversion rates (typically less than 20%), a large recycle stream is required. This can lead to the accumulation of inerts in the gas.

Nitrogen gas (N₂) is unreactive because the atoms are held together by triple bonds. The Haber process relies on catalysts that accelerate the scission of these bonds.

Two opposing considerations are relevant: the equilibrium position and the reaction rate. At room temperature, the equilibrium is in favor of ammonia, but the reaction does not proceed at a detectable rate due to its high activation energy. Because the reaction is exothermic, the equilibrium constant decreases with increasing temperature following Le Châtelier's principle. It becomes unity at around 150–200 °C (302–392 °F).^[3]

Above this temperature, the equilibrium quickly becomes unfavorable at atmospheric pressure, according to the Van 't Hoff equation. Lowering the temperature is unhelpful because the catalyst requires a temperature of at least 400 °C to be efficient.^[3]

Increased pressure favors the forward reaction because 4 moles of reactant produce 2 moles of product, and the pressure used (15–25 MPa (150–250 bar; 2,200–3,600 psi)) alters the equilibrium concentrations to give a substantial ammonia yield. The reason for this is evident in the equilibrium relationship: