Haber Bosch 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.
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.
Bibliography
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