The Haber Process

1. The Haber Process
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 = -92.6 kJ/mol
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.

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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.


The Future of the Haber 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.



References
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
MANCAT: Dept. Access ; Cont. Ed.Andrew Bates
Access Env. ; Integrated Sciences Assignment
Unit: Chemistry – Chemical Equilibrium/KineticsMay 15 2000
Appendix 1
MANCAT: Dept. Access ; Cont. Ed.Andrew Bates
Access Env. ; Integrated Sciences Assignment
Unit: Chemistry – Chemical Equilibrium/KineticsMay 15 2000
2. Acid Rain
Rainwater is acidic. The presence of nitrogen dioxide and sulphur dioxide in the atmosphere lead to low concentrations of nitric acid and sulphuric acid in rainfall. If anthropomorphic effects on the atmosphere are excluded the estimated pH of rainfall would be between 5.6 and 5.8.


Unfortunately man-made changes in the composition of the atmosphere have had an effect on this. The introduction of fossil fuels burning in power stations and other industrial processes has increased the level of sulphur dioxide in the atmosphere to such an extent that it is estimated that more than half of the sulphur dioxide in the atmosphere is there as a result of burning fossil fuels. Concern over this has lead recently to the introduction of technology to reduce the sulphur emissions from plants burning fossil fuels and the introduction of legislation amongst the majority of industrial nations enforcing its use. However, as the level of sulphur dioxide in he atmosphere has been reduced (or at least its growth curtailed) there has been an increase in the emissions of nitrogen dioxide mainly from the increased use of vehicle powered by fossil fuels.


The effects of these gasses has been exacerbated by the introduction of a number of other chemicals into the upper levels of the troposphere such as iron and manganese particles and troposheric ozone (also from vehicle exhausts) all of which act as catalysts for the acidification of rainfall.
The net result has been that the average pH of rainfall across inland Europe is now 4.1 i.e. over ten times more acidic than the “natural” level of acidity of rainfall.


Within an urban environment this has lead to the destruction of stonework such as statues and had accelerated the decay of all objects that are exposed to the elements. However the most worrying effects of acid rain are the three main effects on the non-urban environment (both rural and “natural”) as follows;
Leeching of Soil
Minerals held in soil and used by plants usually exist as complex ions or salts. An increase in the acidity of rainfall results in the rain reacting chemically with the complex ions and forming soluble salts which then “leech” into rivers, lakes and the sea – resulting in nutrient deficient soils no longer capable of sustaining their previous levels of vegetation.


Poisoning of Lakes and Rivers
the most obvious result is the lowering of the pH of the water resulting in the death of those species of plant and animal which are intolerant of low pH values.


However, far more destructive than this has been the poisoning of lakes by aluminium. Aluminium is the most abundant metal in the upper crust of the Earth and is a major constituent of clay. Acidic water reacts with the layers of silicates that hold the aluminium based salts in place, releasing them in a colloidal state. When this enters the water it blocks the gills of the fish present, suffocating them and blocks the feeding mechanisms of filter feeders that occur near the base of the food chain. Once the aluminium has settled to the bottom of the water, taking the sediment in the water with it, it leaves behind extremely clear waters which are almost completely devoid of animal life!
Any reintroduction of fish or other animal life into the waters and it is very likely that the next season of rainfall will replay the process due to the large abundance of aluminium in the soil.


Deforestation
The mechanisms by which acid rain causes or accelerates the destruction of forests are complex and many processes are involved. It is therefore likely that no two forests are affected in the same way. The two main mechanisms are the reduction of photosynthesis and poisoning by aluminium and other metals.


Some trees have leaves that are susceptible to acidic rain – most notably pine needles – with acidic water inhibiting the process of photosynthesis. This causes the leaves most exposed to die and be lost. This effect continued over time leads to “die back” with the outermost tips of twigs and branches continually dying off until the tree itself is no longer viable.


The other main problem is the effect of aluminium released as described above. Tree roots specialise in the uptake of minerals in the soil. In the presence of high counts of minerals, especially heavy metals this results in the tree actively taking up high enough concentrations of the metals concerned for them to act as a poison.