1. Ammonia is industrially synthesized from nitrogen and hydrogen through the Haber process. It is produced on a large scale, exceeding 130 million tonnes annually. 2. The synthesis of ammonia is an equilibrium reaction that is exothermic and favored by high pressure and low temperature. Industrial production occurs at 400-500°C and 150-300 atmospheres using an iron catalyst. 3. Haber optimized the process by implementing a recycling system to increase space-time yield, improving economic production. His recycling design is still used as the basis for industrial ammonia synthesis.

1. C.U.EBONG
INDUSTRIAL SYNTHESIS OF AMMONIA GAS
Synthetic ammonia industrially, is produced from the reaction between nitrogen and hydrogen and it is the base
from which virtually all nitrogen-containing products are derived. The worldwide production ofammonia
exceeds 130 million tonnes and is the sixth largestchemical produced (Chemical and Engineering News, 1996).
Ammonia is a major raw material for industry and agriculture.
CHEMICAL REACTION AND EQUILIBRIUM
Ammonia synthesis from nitrogen and hydrogen is a reversiblereaction in a ratio 1:3 and can be described by
the overall reaction – (1) .
-------------------------(1)
And the equilibrium constant for the reaction is defined as;
Keq =
[NH3]2
[N2][H2]3
A flow scheme for the Haber Process
Nitrogen from air
4000C at 200atm
Iron Catalyst
Hydrogen from
Natural gas
Unreacted gases
Recycled Gases are cooled &
NH3 turns to liquid
Liquid NH3
The formation of ammonia is an exothermic reaction with considerablerelease of heat. The reaction is a
reversible reaction, that is, itcan proceed both in forward direction (ammonia synthesis) and backward direction
(ammonia decomposition). The reaction is accompaniedby decrease in volume because there is a decrease
innumber of moles of gas according to Le Chatelier’s Principle.
Le Chartelier’s Principle
1:3 by volume
Nitrogen gas N2 and Hydrogen
gas H2

2. In 1884, the French Chemist Henri Le Chatelier suggested that when a system at equilibrium is disturbed, the
equilibrium position will shiftin the direction which tends to minimise, or counteract, the effect of the
disturbance. For a reversible chemicalreaction like ammonia (NH3), Chatelier’s principle has several important
implications as given below:
1. If the concentration of a reactant is increased, the equilibrium position shifts to use up the added
reactantsby producing more products.
2. For gaseous reactions, gas pressure is related to the number of gas particles in the system; more gas
particlesmeans more gas pressure. Consider a reaction which is accompanied by decrease in number of
moles, such as,ammonia synthesis (a). Increasing the pressure on this equilibrium system will result in
the equilibrium positionshifting to reduce the pressure, that is, to the side that has the least number of
gas particles.
3. In an endothermic reaction, energy can be considered as a reactant of the reaction while in an
exothermicreaction, energy can be considered as a product of the reaction. Consider an exothermic
reaction which isaccompanied by release of heat, such as ammonia synthesis;
 Reducing the temperature of this equilibriumsystem (removing heat) will result in the equilibrium
position shifting to increase the temperature (producingmore heat), that is, to shift the equilibrium
position to the right.
YOU CAN HAVE A HIGHER YIELD OF AMMONIA BY;
 Increasing the pressure whch causes the equilibrium to shift tothe right resulting in a higher yield of
ammonia since there is apressure drop accompanying the transformation;
 Decreasing thetemperature which also causes the equilibriumposition to move to the rightagain
resulting in a higher yield of ammonia since the reaction isexothermic (releases heat).
EFFECT OF TEMPERATURE AND PRESSURE ON AMMONIA SYNTHESIS
It can be seen that the ammonia mole fraction decreases as the temperature is increased while it
increases as the pressure is increased.
Temperature
(oC)
Keq
25 6.4 x 102
200 4.4 x 10-1
300 4.3 x 10-3
400 1.6 x 10-4

3. Fig. 1: Showing the effect of Pressure Fig. 1: Showing the effect Temperature variation
We can conclude then that ammonia synthesisaccording to equation (1) is an equilibriumreaction that is
favoured by low temperatureand high pressure. Equilibriumconditionsare only a part of the picture, that is,
thermodynamics does not give us any ideaabout the rate of the reaction. The reactiondoes not proceed at
ambient temperature becausenitrogen requires a lot of energy todissociate. In the gas phase this
dissociationoccurs only at around 3000°C. Even the hydrogenmolecule, which has a weaker molecularbond,
only dissociates markedly attemperatures above 1000°C. Thus, the reactioncannot be performed at lower
temperaturebecause it needs high energy, and if we increase the temperatureto a high level, the reverse reaction
predominates. This is clearly avicious circle. This is where the role of the iron catalyst comes in.
The hydrogen and nitrogenmolecules loose their translational degrees of freedom when boundto the catalyst
surface. This reduces the activation energy dramaticallyand thus makes the forward reaction go faster. This
means wecan do away with extremely high temperature conditions. Also, theuse of lower temperature reaction
conditions means there is limitedreverse reaction which is energy saving as well. This does not meanthat we can
go down to ambient temperatures. We still need reasonablyhigh temperatures (250–400°C) to dissociate the N2
and H2reactant molecules, even with the use of a catalyst. The use of acatalyst essentially provides a good
trade-off. It accelerates thereaction sufficiently so that we can obtain ammonia at conditionswhere the
equilibrium conversion is large enough to be useful. Thereaction rate depends on the temperature as well as
conversion ofthe reactants. It can be seen that at lowconversion, higher temperatures can bemaintained to
achieve higherreaction rates. However, as the conversions increase, the temperaturehas to be decreased to
overcome the limitations posed by theequilibrium (rate equals zero). The ammonia converters today combine
the catalyst section and heat exchanger to achievedesired temperature profiles to strike a balance between
higher ratesof reaction and constraints posed by the equilibrium considerations.
CATALYST AND RATE OF REACTION IN AMMONIA SYNTHESIS
Activation energy of a reaction, Ea, is the minimum amount of energy that reactant molecules must possessin
order to form products. In an Energy Profile diagram, the activation energy is the energy differencebetween the
reactants and the peak of the energy profile diagram, which represents the ‘transition state’ of the reaction. The
lower the activation energy, the faster will be the rate of the reaction.Enthalpy change, ∆H, is the amount of
500 1.5 x 10-5

4. energy absorbed or released during the transformation of thereactants to products, and in the energy profile
digram it is depicted as the energy difference between thereactants and products. During an exothermic reaction
energy is released, as the products are lower inenergy than the reactants (∆H is negative), while during an
endothermic one energy is absorbed (∆H ispositive). Catalysts speed up the rate of reaction by lowering the
activation energy without themselvesbeing consumed during the reaction. A catalyst does not alter the net
enthalpy change for the reaction and,therefore, does not alter the equilibrium of the reaction, but merely
increases the rate of reaction.
HABER’S CONTRIBUTIONS/THE RECYCLING PROCESS
By the turn of the 19th century, complete understanding andapplication of the law of mass action, kinetics and
chemical equilibriaenabled chemists to investigate the synthesis of ammonia moresystematically. From the
equilibrium data measured it was obviousthat, at normal pressure, the reaction temperature should be keptwell
below 300°C in order to obtain even a small percentage ofammonia. For this temperature range, however, no
catalyst wasavailable. By increasing the pressure to 75 bar the equilibriumconditions improved, but even at this
pressure, and an operatingtemperature of about 600°C, most known catalysts at that time led toa very low
ammonia concentration. It was Haber who finally overcamehis colleagues’ excessive preoccupation with
unfavorableequilibrium concentrations. Firstly, he recognized that much higherpressures had to be employed
and he constructed a small laboratoryapparatus for the continuous production of ammonia. Secondly,and
perhaps more importantly, he developed the concept of a recycleprocess.
The amount of ammonia formed in a single gas pass is much toosmall to be of interest for the economic
production of ammonia.Haber, therefore, recycledthis gas over the catalyst after separatingthe ammonia formed
by condensation. The gas lost by conversion was compensated with a fresh gas input and the mixture
wasrecycled under pressure. This process became the basis for thetechnical manufacture of ammonia. Since
then, the same principlehas found widespread application for numerous high-pressure reactionsin the organic
chemistry sector. Haber’s recycle idea changedthe previously static conception of process engineering in favor
of amore dynamic approach. For the first time, reaction kinetics as wellas the thermodynamics of the system
were being considered. Inaddition to chemical equilibrium, Haber recognized that reactionrate was a
determining factor in this problem. Instead of simple reactionyield, he concentrated on space-time yield, that is,
the amountof ammonia obtained per unit volume of the catalyst per unit time. In this manner it became apparent
that the real problemwas to find a suitable catalyst so that the maximum amount of productis obtained with
minimum volume of the catalyst in the shortesttime possible, that is, space-time yield needs to be maximized.
REFERENCES
Bradley,David(2004). "A GreatPotential:The GreatLakesas a Regional Renewable EnergySource" (PDF).Archivedfrom
the original on 29 October 2008. Retrieved 2015-03-12.

5. Max Appl "Ammonia" in Ullmann's Encyclopedia of Industrial Chemistry 2006 Wiley-VCH, Weinheim.
doi:10.1002/14356007.a02_143.pub2
A SCHEMATIC DIAGRAM SHOWING THE INDUSTRIALSYNTHESIS OF AMMONIA