Haber process

Overview

The Haber process, also called the Haber–Bosch process, is the reaction of nitrogen and hydrogen, over an iron substrate, to produce ammonia. The Haber process is important because ammonia is difficult to produce on an industrial scale. 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 and nitrites essential for the production of nitrate fertilizer and munitions.

History

The process was first patented by Fritz Haber. In 1910 Carl Bosch, while working for 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 ammonium nitrate (for use in explosives) at a time when supply of Chile saltpetre from Chile could not be guaranteed because this industry was then almost 100% in British hands. It has been suggested that without this process, Germany would almost certainly have run out of explosives by 1916, thereby ending the war.

The process

The bulk of the chemical technology consists in getting the hydrogen from methane or natural gas using heterogeneous catalysis and then reacting it with the atmospheric nitrogen.

Synthesis gas preparation

First, the methane is cleaned, mainly to remove sulfur impurities that would poison the catalysts. This is done by turning sulfur into hydrogen sulfide:

CH₃SH + H₂ → CH₄ + H₂S

and then reacting with zinc oxide to form zinc sulfide:

H₂S + ZnO → ZnS + H₂O

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. The carbon monoxide formed is also cause catalyst poisoning, which would react with the iron catalyst, forming an iron compound, thus affecting the reaction. The air added during this step also serves as a nitrogen source for ammonia synthesis:

CH₄ + 1/2O₂ → CO + 2H₂

CH₄ + 2O₂ → CO₂ + 2H₂O

Then occur two’’shifts’’ which convert CO to CO₂ by reaction with steam, one at high temperature, then one at low temperature:

CO + H₂O → CO₂ + H₂ high temperature

→ the catalyst here is a mixture of iron, chromium and copper

CO + H₂O → CO₂ + H₂ low temperature

→ the catalyst here is a mixture of copper, zinc and aluminium

Then the carbon dioxide is removed by reaction with potassium carbonate.

K₂CO₃ + H₂O + CO₂→ 2KHCO₃

The gas mixture is now passed into a methanator which converts any remaining CO₂ into methane for recycling:

CO₂ + 4H₂ → CH₄ + 2H₂O

We now have a gas mixture containing nitrogen and hydrogen in the correct ratio of 1:3.

Ammonia synthesis

The final stage is the crucial synthesis of ammonia using promoted magnetite, iron oxide, as the catalyst:

N₂(g) + 3H₂(g) → 2NH₃(g), ΔH^o = -92.4 kJ/mol

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 will be recycled, so that eventually an overall yield of 98% can be achieved.

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:

<math>K_\mathrm{eq} = \mathrm{\frac{[NH_3]^2}{[N_2][H_2]^3}}</math>

Variation in K_eq for the Equilibrium
N₂ (g) + 3H₂ (g) ↔ 2NH₃ (g)
as a Function of Temperature

Temperature (°C)	Keq
300	4.34 x 10–3
400	1.64 x 10–4
450	4.51 x 10–5
500	1.45 x 10–5
550	5.38 x 10–6
600	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, 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.

Catalysts

The catalyst has no effect on the position of 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 and uranium 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:

1. N₂(g) → N₂(adsorbed)
2. N₂(adsorbed) → 2N(adsorbed)
3. H₂(g) → H₂(adsorbed)
4. H₂(adsorbed) → 2H(adsorbed)
5. N(adsorbed) + 3H(adsorbed)→ NH₃(adsorbed)
6. 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.

See also

• Chemical kinetics
• Reaction rate
• Rate equation

External links

• What is the Haber-Bosch Process?
• Haber-Bosch process
• Fertilizer,agriculture and the production of food
• Food Crises and the Role of Agriculture: Past and Current
• Britannica guide to Nobel Prizes: Fritz Haber
• Nobel e-Museum - Biography of Fritz Haber
• Uses and Production of Ammonia
• ”CIEC Catalysis”, [1]
• Haber Process for Ammonia Synthesis