During the first decade of the twentieth century the world-wide demand for ammonia for use in fertilisers (in the form of nitrates) and in the production of explosives for use in mining and warfare could only be satisfied on a large scale from deposits of guano in Chile (2). Though this deposit was of huge size (approximately five feet thick and 385 kilometres long) it represented a rapidly depleting resource when compared to world-wide demand. As a result of this there was much research into how ammonia could be produced from atmospheric nitrogen.
The problem was eventually solved by Fritz Haber (1868 – 1934) in a process which came to be known as the “Haber Process” or the “Haber – Bosch Process”.
Haber developed a method for synthesising ammonia utilising atmospheric nitrogen and had established the conditions for large scale synthesis of ammonia by 1909 and the process was handed over to Carl Bosch for industrial development (1). the reaction is a simple equilibrium reaction which occurs in gaseous state as follows;
N2 (g) + 3H2 (g) = 2NH3 (g) heat of enthalpy = -92.
In predicting how to obtain the highest yield from this reaction we must refer to Le Chatlier’s Principle. This states that for an equilibrium reaction the equilibrium will work in the opposite direction to the conditions forced upon it. The conditions most pertinent to the above reaction are temperature and pressure.
The pressure exerted by any gas or mixture of gasses in an enclosed space is directly proportional to the number of atoms or molecules of gas regardless of their size or molecular mass. Reference to the above reaction shows that, as the reaction moves to the right the number of molecules and hence the pressure decreases. Therefore the reaction moving to the right (i.e. towards the product required) is favoured by an increase in pressure.
With regard to temperature, the reaction moving to the right is exothermic i.e. it gives off energy (in the form of heat). Therefore reference to Le Chatlier’s Principle shows that the reaction to the right is favoured by low temperatures.
However, when Haber placed the reactants together under these conditions it was shown that the rate of reaction was so slow as to render the process unfeasible as an industrial process. This is because of an unusually high activation energy.
The activation energy of a reaction is the energy required by the reactants to achieve an intermediate state required before they form the products. In the case of the above reaction the intermediate state requires the dissociation of diatomic gaseous nitrogen. The triple bond found between two nitrogen atoms when they form diatomic nitrogen is amongst the strongest chemical bonds known. this leads to an extremely high activation energy.
At extremely high temperature the nitrogen molecule will dissociate and so, as the temperature approaches this point the rate at which the reaction to the right occurs and therefore the speed with which equilibrium is reached increases rapidly. Unfortunately experimentation showed that, as temperature approached the point at which the speed of the reaction was sufficient to produce a viable reaction the amount of ammonia produced was so low that the reaction was still unfeasible on as an industrial process.
Faced with this failure to find conditions suitable for an industrial process Haber began to experiment to find a catalyst that would facilitate the reaction. A catalyst is a substance that reduces the activation energy of a reaction, thus increasing the speed at which the reaction occurs, or in the case of equilibrium reactions the speed at which equilibrium is reached. After hundreds of experiments Haber discovered that a fast enough reaction with a high enough yield of ammonia would occur with a pressure between 200 and 400 atmospheres and at a temperature between 670K and 920K in the presence of a catalyst of iron (in the form of iron filings to increase its active surface area) plus a few percent of oxides of potassium and aluminium. This process was first demonstrated in 1909 and patented as the Haber Process in 1910 (3).
Experiments aimed at finding the most efficient conditions for the reaction have since resulted in the process described by the flow diagram in Appendix 1.
The Haber process has been used since its discovery to produce ammonia which has been used predominately to produce fertilisers which have helped to feed a rapidly growing world population and has been one of the main props used to avoid world-wide famine. The increase in the use of nitrogen based fertilisers is demonstrated in Appendix 2.
Unfortunately there are consequences to such a high level of use of this industrial process.
In 1998 the Haber Process accounted for 29% of the atmospheric nitrogen fixed in the form of nitrates used by vegetation world-wide (4). If this reliance on artificial fertiliser is continued and the world population increases as expected (with the attendant increase in the number of crops being grown) then by the year 2050 160,000,000 tons of nitrogen will need to be manufactured per annum requiring the burning of 270,000,000 tons of coal or its equivalent to feed this energy – hungry process with all of the attendant environmental problems (5). Further to this the use of chemical fertilisers also affects the global nitrogen cycle, pollutes groundwater and increases the level of atmospheric nitrogen dioxide – a potent “greenhouse” gas.
As a result of this work is now underway to both try to solve the problem of the high energy consumption of the Haber Process and to reduce our reliance on chemical fertilisers.
The Unit of Nitrogen Fixation at Sussex University has now identified the reaction with the metal molybdenum within the enzyme nitrogenase which allows bacteria to fix atmospheric nitrogen at soil temperatures. This has enabled research to commence on low energy methods of producing ammonia.
With regard to reducing our reliance on chemical fertilisers, funding is now being allocated to experiments into ways in which the amount of biological nitrogen fixation occurring can be encouraged the growth of nitrogen fixing microbes in the soil (7).
The current method of production of nitrates via the production of ammonia in the Haber Process has been identified as being destructive to the environment despite its beneficial effects in helping to feed the world population. As a result funding is now being allocated to finding alternatives to this process. Though both of the above projects are far from complete they do demonstrate a commitment to making the Haber Process redundant and it is fairly certain that even if these avenues of research prove to be unsuccessful others will be explored until an alternative is found. it therefore seems that the days of one of the most widespread industrial processes in the world are now numbered.
1.Encyclopaedia Britannica – 3 June 2000
2.University of Wisconsin Web site – “Science is Fun” – 3 June 2000
3.Raffles Institute Media Networking Club – Web page – 4 June 2000
4.Micro-organism’s in Action. J M Lynch & J E Hobbie. Blackwell Publication 1998
5.Biological Nitrogen Fixation – National Research Council . National Academic Press 1994
6.Article – New Scientist – 10 May 1986
7.The Microbial World. J Deacon. The University of Edinburgh 2000
Cite this The Haber Process
The Haber Process. (2018, Jul 04). Retrieved from https://graduateway.com/the-haber-process-essay/