The Haber process
, also called the Haber–Bosch process
, is the nitrogen fixation
reaction of nitrogen
, over an iron substrate
, to produce ammonia
. The Haber process is important because ammonia is difficult to produce on an industrial scale and the fertilizer generated from the ammonia is responsible for sustaining one-third of the Earth's population. Even though 78.1% of the air
we breathe is nitrogen
, the gas is relatively unreactive because nitrogen molecules are held together by strong triple bonds
. It was not until the early 20th century that this method was developed to harness the atmospheric abundance of nitrogen
to create ammonia
, which can then be oxidized
to make the nitrates
essential for the production of nitrate fertilizer
The process was first patented by German chemist Fritz Haber
. In 1910 Carl Bosch
, while working for German
chemical company BASF
, successfully commercialized the process and secured further patents. Haber and Bosch were later awarded Nobel prizes
, in 1918 and 1931 respectively, for their work in overcoming the chemical and engineering problems posed by the use of large-scale high-pressure technology.
Ammonia was first manufactured using the Haber process on an industrial scale in Germany during World War I
, to meet the high demand for nitric acid
, for use in the manufacturing of explosives, at a time when supply of Chile saltpetre
could not be guaranteed because this industry was then almost 100% in British hands. It has been suggested that without this process, Germany would not have fought in the war.
Prior to the use of natural gas as a hydrogen source, electricity was used to electrolyse water. The Vemork 60 MW hydro electric plant in Norway was constructed purely to produce hydrogen via electrolysis of water as a precursor to ammonia production, and up until the second world war provided the majority of Europe's ammonia.
Nowadays, bulk of the chemical technology consists of isolating hydrogen
) using heterogeneous catalysis
and then reacting it with atmospheric nitrogen, but this is not in fact the Haber Process.
Synthesis gas preparation
First, the methane is cleaned, mainly to remove sulfur impurities that would poison the catalysts.
The clean methane is then reacted with steam over a catalyst of nickel oxide. This is called steam reforming:
- CH4 + H2O → CO + 3H2
Secondary reforming then takes place with the addition of air to convert the methane that did not react during steam reforming.
- 2CH4 + O2 → 2CO + 4H2
- CH4 + 2O2 → CO2 + 2H2O
Then two shift reactions convert CO to CO2 by reaction with steam.
The gas mixture is now passed into a methanator, which converts any remaining CO2 into methane for recycling:
- CO2 + 4H2 → CH4 + 2H2O
Ammonia synthesis - Haber Process
The final stage, which is the actual Haber Process is the synthesis of ammonia using magnetite
, iron oxide, as the catalyst:
- N2(g) + 3H2(g) ⇌ 2NH3(g), ΔHo = −92.4 kJmol-1
This is done at 150–250 atmospheres (atm) and between 300 and 550 °C, passing the gases 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 any unreacted gases are recycled, so that eventually an overall conversion of 98% can be achieved.
The steam reforming, shift conversion, carbon dioxide removal, and methanation steps each operate at absolute pressures of about 25 to 35 bar, and the ammonia synthesis loop operates at absolute pressures ranging from 60 to 180 bar, depending upon which proprietary design is used. There are many engineering and construction companies that offer proprietary designs for ammonia synthesis plants. Haldor Topsoe of Denmark, Lurgi AG of Germany, Uhde of Germany, and Kellogg, Brown and Root of the United States are among the most experienced companies in that field.
Reaction rate and equilibrium
There are two opposing considerations in this synthesis: the position of the equilibrium and the rate of reaction
. At room temperature, the reaction is slow and the obvious solution is to raise the temperature. This may increase the rate of the reaction but, since the reaction is exothermic
, it also has the effect, according to Le Chatelier's Principle
, of favouring the reverse reaction and thus reducing equilibrium constant
, given by:
Variation in Keq for the Equilibrium
N2 (g) + 3H2 (g) ↔ 2NH3 (g)
as a Function of Temperature
| Temperature (°C)
|| Keq |
|| 4.34 x 10–3 |
|| 1.64 x 10–4 |
|| 4.51 x 10–5 |
|| 1.45 x 10–5 |
|| 5.38 x 10–6 |
|| 2.25 x 10–6 |
As the temperature increases, the equilibrium
is shifted and hence, the constant drops dramatically according to the van't Hoff equation
. Thus one might suppose that a low temperature is to be used and some other means to increase rate. However, the catalyst itself requires a temperature of at least 400 °C to be efficient.
Pressure is the obvious choice to favour the forward reaction because there are 4 moles of reactant for every 2 moles of product (see entropy), and the pressure used (around 200 atm) alters the equilibrium concentrations to give a profitable yield.
Economically, though, pressure is an expensive commodity. Pipes and reaction vessels need to be strengthened, valves more rigorous, and there are safety considerations of working at 200 atm. In addition, running pumps and compressors takes considerable energy. Thus the compromise used gives a single pass yield of around 15%.
Another way to increase the yield of the reaction would be to remove the product (i.e. ammonia gas) from the system. In practice, gaseous ammonia is not removed from the reactor itself, since the temperature is too high; but it is removed from the equilibrium mixture of gases leaving the reaction vessel. The hot gases are cooled enough, whilst maintaining a high pressure, for the ammonia to condense and be removed as liquid. Unreacted hydrogen and nitrogen gases are then returned to the reaction vessel to undergo further reaction.
has no effect on the position of chemical equilibrium
; rather, it provides an alternative pathway with lower activation energy
and hence increases the reaction rate, while remaining chemically unchanged at the end of the reaction. The first Haber–Bosch reaction chambers used osmium
catalysts. However, today a much less expensive iron
catalyst is used almost exclusively.
In industrial practice, the iron catalyst is prepared by exposing a mass of magnetite, an iron oxide, to the hot hydrogen feedstock. This reduces some of the magnetite to metallic iron, removing oxygen in the process. However, the catalyst maintains most of its bulk volume during the reduction, and so the result is a highly porous material whose large surface area aids its effectiveness as a catalyst. Other minor components of the catalyst include calcium and aluminium oxides, which support the porous iron catalyst and help it maintain its surface area over time, and potassium, which increases the electron density of the catalyst and so improves its reactivity.
The reaction mechanism, involving the heterogeneous catalyst, is believed to be as follows:
- N2(g) → N2(adsorbed)
- N2(adsorbed) → 2N(adsorbed)
- H2(g) → H2(adsorbed)
- H2(adsorbed) → 2H(adsorbed)
- N(adsorbed) + 3H(adsorbed)→ NH3(adsorbed)
- NH3(adsorbed) → NH3(g)
Reaction 5 occurs in three steps, forming NH, NH2, and then NH3. Experimental evidence points to reaction 2 as being the slow, rate-determining step.
A major contributor to the elucidation of this mechanism is Gerhard Ertl.
Economic and environmental aspects
The Haber process now produces 100 million tons of nitrogen fertilizer
per year, mostly in the form of anhydrous ammonia
, ammonium nitrate
, and urea
. 3-5% of world natural gas production is consumed in the Haber process (~1-2% of the world's annual energy supply). That fertilizer is responsible for sustaining one-third of the Earth's population, as well as various deleterious environmental consequences. Generation of hydrogen using electrolysis of water, using renewable energy, is not currently competitive cost-wise with hydrogen from fossil fuels, such as natural gas, and is responsible for only 4% of current hydrogen production. Notably, the rise of this industrial process led to the "Nitrate Crisis" in Chile
, when the British industrials left the country — since the natural nitrate mines were no longer profitable — closing the mines and leaving a large unemployed Chilean population behind.