The Haber process, also called the Haber–Bosch process, is the main industrial procedure for the production of ammonia. It converts atmospheric nitrogen (N2) to ammonia (NH3) by a reaction with hydrogen (H2) using finely divided iron metal as a catalyst:
This reaction is exothermic but disfavored in terms of entropy because four equivalents of reactant gases are converted into two equivalents of product gas. However, due to the extremely high kinetic barriers to bond cleavage, high pressures and temperatures are needed to drive the reaction forward. The Haber Process was first created by the German Chemist Fritz Haber, then developed after a few years by Carl Bosch.
The German chemists Fritz Haber and Carl Bosch developed the process in the first decade of the 20th century, and its improved efficiency over existing methods such as the Birkeland–Eyde and Frank–Caro processes was a major advancement in the industrial production of ammonia.
The Haber process can be combined with steam reforming to produce ammonia with just three chemical inputs: water, natural gas, and atmospheric nitrogen. Both Haber and Bosch were eventually awarded the Nobel Prize in Chemistry: Haber in 1918 for ammonia synthesis specifically, and Bosch in 1931 for related contributions to high-pressure chemistry.
History
During the 19th century, the demand rapidly increased for nitrates and ammonia for use as nitrogen fertilizers, which can supply plants with extra nitrogen compounds, and for industrial feedstocks. The main source was mining niter deposits and guano from tropical islands. At the beginning of the 20th century these reserves were thought insufficient to satisfy future demands, 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.. They demonstrated their process in the summer of 1909 by producing ammonia from the air, drop by drop, at the rate of about 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.Bosch, Carl (2 March 1908). . 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.
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. During World War I, the production of munitions required large amounts of nitrate. The Allied powers 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. Moreover, even if German commercial interests had nominal legal control of such resources, the Allies controlled the sea lanes and imposed a highly effective blockade which would have prevented such supplies from reaching Germany. The Haber process proved so essential to the German war effort 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 this was available in extremely small quantities. Haber noted that 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 discovery of this catalysis was Gerhard Ertl. 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, the Claude process, and the Mont-Cenis process developed by the Friedrich Uhde Ingenieurbüro. Luigi Casale and Georges Claude proposed to increase the pressure of the synthesis loop to , 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 ( and ), 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 1–2% of global energy consumption, 3% of global carbon emissions, and 3% to 5% of natural gas consumption. Hydrogen required for ammonia synthesis is most often produced through gasification of hydrocarbons, mostly natural gas, but other potential hydrogen 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%; however, in China as of 2022, natural gas and coal were responsible for 20% and 75% respectively. 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 reported ammonia synthesis using Ru-loaded calcium-aluminium oxide C12A7: electride, describing the electride as an electron donor and reversible hydrogen store. Subsequent work studied nitrogen dissociation and catalytic mechanisms on Ru-loaded C12A7:. This method is implemented in a small plant for ammonia synthesis in Japan. In 2019, Hosono's group found another catalyst, a novel perovskite oxynitride-hydride , that works at lower temperature and without costly ruthenium.
Hydrogen production The major source of hydrogen is methane. Steam reforming of natural gas 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, using water electrolysis or thermochemical (solar or another heat source) water splitting.
Starting with a natural gas () 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): 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:
H2S + ZnO -> ZnS + H2O Catalytic steam reforming of the sulfur-free feedstock forms hydrogen plus carbon monoxide: CH4 + H2O -> CO + 3 H2 Catalytic shift conversion converts the carbon monoxide to carbon dioxide and more hydrogen: 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: CO + 3 H2 -> CH4 + H2O CO2 + 4 H2 -> CH4 + 2 H2O
Ammonia production The hydrogen is catalytically reacted with nitrogen (derived from air separation) to form anhydrous liquid ammonia. It is difficult and expensive, as lower temperatures result in slower reaction kinetics (hence a slower reaction rate) and high pressure requires high-strength pressure vessels that resist hydrogen embrittlement.
The primary reaction taking place is the ammonia synthesis loop:
3 H2 + N2 -> 2 NH3 As diatomic nitrogen is bound together by a triple bond, it is relatively inert. To combat this, catalysts are used accelerate the scission of these bonds and unreacted gases are reprocessed.
The reactants are passed over four beds of catalyst, with cooling between each pass to maintain a reasonable equilibrium constant. 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.
On each pass, only about 15% conversion occurs, meaning that the ammonia must be extracted and the gases reprocessed for the reaction to proceed at an acceptable pace. Despite that, due to the recycling of leftover reactants, eventually conversion of 97% is achieved. This extensive recycling can lead to the accumulation of inerts in the gases.
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 .
+ K(T) for + 3 ⇌ 2 Temperatureθ (°C) Equilibrium constantK 300 434 × 10−5 400 16.4 × 10−5 450 4.51 × 10−5 500 1.45 × 10−5 550 0.538 × 10−5 600 0.225 × 10−5 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.
Increased pressure favors the forward reaction because 4 moles of reactant produce 2 moles of product, and the pressure used () alters the equilibrium concentrations to give a substantial ammonia yield. The reason for this is evident in the equilibrium relationship:
where is the fugacity coefficient of species , is the mole fraction of the same species, is the reactor pressure, and is standard pressure, typically .
Economically, reactor pressurization is expensive: pipes, valves, and reaction vessels need to be strong enough, and safety considerations affect operating at 20 MPa. Compressors take considerable energy, as work must be done on the (compressible) gas. Thus, the compromise used gives a single-pass yield of around 15%.
While removing the ammonia from the system increases the reaction yield, this step is not used in practice, since the temperature is too high; instead it is removed from the gases leaving the reaction vessel. The hot gases are cooled under high pressure, allowing the ammonia to condense and be removed as a liquid. Unreacted hydrogen and nitrogen gases are returned to the reaction vessel for another round. While most ammonia is removed (typically down to 2–5 mol.%), some ammonia remains in the recycle stream. In academic literature, a more complete separation of ammonia has been proposed by absorption in metal halides, metal-organic frameworks or zeolites. Such a process is called an absorbent-enhanced Haber process or adsorbent-enhanced Haber–Bosch process.
Pressure/temperature