SYNTHESIS of AMMONIA
Chemistry KS4 science GCSE/IGCSE/O level Revision Notes
PART C The HABER SYNTHESIS of AMMONIA - the
FIXATION of NITROGEN
The industrial manufacture of ammonia is described
in terms of the reaction conditions employed in the Haber Synthesis of ammonia
from hydrogen and nitrogen e.g. temperature, pressure, iron catalyst, recycling
unreacted gases, condensing out of ammonia. Equilibrium theory and rates of
reaction factors are explained and discussed to arrive at the optimum (most
efficient) reaction conditions to make ammonia gas. These revision notes on the
Haber synthesis and artificial synthetic nitrogen fixation from air, should
prove useful for the new AQA GCSE chemistry, Edexcel GCSE chemistry & OCR GCSE
chemistry (Gateway & 21st Century) GCSE (9–1), (9-5) & (5-1) science courses.
A Reversible Reactions * B Reversible reactions
* C Haber Synthesis of ammonia (this page)
D(a) The Uses of ammonia-nitric acid-fertilisers
The nitrogen cycle
Advanced A Level Notes - Chemical Equilibrium
Advanced A Level Notes
- Nitrogen & Ammonia
PART C 3a. The
Synthesis of ammonia - The Haber Process of Nitrogen Fixation
The Haber process is used by the
chemical industry to manufacture ammonia which is used to make
acid and fertilisers. The raw
materials for the Haber process are nitrogen from air and hydrogen is
usually obtained from natural gas. The purified gases are passed over a
catalyst of iron at a high temperature (typically 450°C) and a very high
pressure (typically 200 atmospheres). Some of the hydrogen reacts with
nitrogen reacts to form ammonia (nitrogen + hydrogen ==> ammonia).
Unfortunately the reaction is reversible i.e. some of the ammonia
produced decomposes back into nitrogen and hydrogen. On cooling the
reaction mixture, the ammonia liquefies and is removed. To increase
efficiency the remaining unreacted hydrogen and nitrogen are recycled.
The rules on equilibria can be applied to predict the best conditions to
get the highest yield of ammonia but unfortunately there has to be a
trade-off between rate of production and position of equilibrium.
.. which means an equilibrium
will form, so there is no chance of 100% yield even if you use, as you
actually do, the theoretical reactant ratio of nitrogen : hydrogen of 1 : 3 !
In forming ammonia 92kJ of heat energy is given out
exothermic, 46kJ of heat released per mole of ammonia formed).
Also, four moles of 'reactant' gas form two moles of 'product' gas, so there is
a net decrease in gas molecules on forming ammonia.
Actual values of % ammonia in an equilibrium
mixture with nitrogen and hydrogen are shown in the graph above.
Each line of the graph gives you the % of
ammonia in the equilibrium mixture as the pressure is gradually increased at
a constant temperature (in this case 100oC to 800oC at
100 degree intervals)..
These graphs should match your prediction
from applying the equilibrium rules explained in
So applying the
equilibrium rules, the formation of ammonia
should be favoured by ...
- Ammonia gas is synthesised in the chemical industry by
reacting nitrogen gas with hydrogen gas in what is known as the
Haber-Bosch Process, named after two highly inventive and subsequently
- The Haber synthesis of ammonia is important
for agriculture because nitrogen is an important element for plant growth.
- But, it is a very stable molecule and only a
few plants like legumes (peas, beans etc.) can directly 'fix' nitrogen the
from air and incorporate it into protein molecules.
- The Haber synthesis allows the efficient
mass production of 'artificial fertilisers'.
- If the chemical feedstocks for the
Haber Process are nitrogen and hydrogen, where do we get these materials
- The nitrogen was once obtained from
the fractional distillation of
liquified air (80% N2).
- the air is filtered to remove dust and then
compressed under high pressure.
- The filtered air is cooled and water
condensed out, and then carbon dioxide freezes out at -78oC.
- The air further cooled to form liquid air (liquefaction)
at around -200oC.
- The liquid air is fractionally distilled at low temperature to separate
oxygen (used in welding, hospitals etc.), nitrogen (for making ammonia),
Noble Gases e.g. argon for light bulbs, helium for balloons).
- Oxygen and argon are very close in boiling
point and initially come out in the same fraction so a further fractional
distillation is needed to separate them.
- However, nitrogen is also produced by
'deoxygenating' air by combustion with methane.
- The hydrogen is made by (i) reacting methane (natural
gas) and water or (ii) from cracking hydrocarbons (both reactions are done at
high temperature with a catalyst).
- (i) methane + water (steam) ==> hydrogen +
- CH4 + H2O ====> 3H2
- The 'deadly' carbon monoxide can be
reacted with water in a 2nd stage to make more hydrogen.
- CO + H2O ====> CO2
- and the carbon dioxide is
removed to give the desired hydrogen gas.
- These are called 'reforming'
- or (ii) from cracking an alkane
hydrocarbon from crude oil e.g.
- AND, but just in passing and convenient ....
- Other uses of hydrogen
- Hydrogen-oxygen fuel cells to make electricity on
- Hydrogenation vegetable oils to make margarine.
- Reducing metal oxides to free the metal.
- An atomic hydrogen-oxygen welding torch.
- Inflating weather balloons.
- HABER PROCESS: The Haber process is based on the
reversible reaction between nitrogen and hydrogen to form ammonia and using
an iron catalyst.
- The balanced equation for the Haber
Synthesis reversible reaction is ...
(plus 92 kJ of heat energy given out,
N2(g) + 3H2(g)
In industry pressures of 200 - 300 times normal atmospheric pressure are used
in line with the theory (200-300 atm).
Theoretically a low temperature would give a high yield of ammonia BUT ...
- (a) Using high pressure because you are
going from 4 to 2 gas molecules, so high pressure favours the forward
reaction to give fewer gas molecules.
- The high pressure also speeds up the
reaction because it effectively increases the concentration of the gas
- but, the higher pressure means more dangerous and more costly
engineering, so a compromise needed.
- Does this prediction match the graph?
- Note: Increasing the pressure also increases the rate
of the reactions, because increasing pressure is effectively increasing the
concentration of the reactants.
- So increasing pressure has two positive effects in
terms of yield and getting to the maximum possible % equilibrium of ammonia.
- (b) Carrying out the reaction at a
low temperature, favouring the forward reaction,
- because it is an
exothermic reaction favoured by lowering the temperature,
- Does this prediction match the graph?
- However, this may produce too
slow a rate of reaction, so a compromise is needed.
- Therefore, the idea is to use a set of optimum conditions to get the
most efficient yield of
ammonia and this involves getting a low % yield (e.g. 8% - 15% conversion)
- Described below are the conditions to give the most economic
production of ammonia.
- These arguments make the point that
yield of an equilibrium reaction depends on
the conditions used.
- The word 'yield' means how much
product you get compared to the theoretical maximum possible if the
forward reaction goes 100% that way.
- For more on chemical economics see
Industrial Chemistry page.
reference to the HABER SYNTHESIS chemical plant DIAGRAM
is very stable molecule and not very
reactive i.e. chemically inert, so the rate of reaction is too slow
at low temperatures.
- To speed up the reaction
an iron catalyst is used as well as a
higher temperature (e.g.
- The higher temperature is an economic compromise,
i.e. it is more economic to get a low yield fast, than a high yield slowly!
- Note: a catalyst does NOT affect the yield of a
reaction, i.e. the equilibrium position BUT you do get
to the equilibrium position a lot faster!
At the end of the process, when the
gases emerge from the bottom of the iron catalyst reaction chamber, the gas mixture is cooled under high pressure, when
only the ammonia liquefies and is so can be removed, tapped off from
the cooled compression chamber and stored in cylinders for use e.g. making
Because the reaction is reversible, not
all the nitrogen and hydrogen are converted to ammonia.
Any unreacted nitrogen and hydrogen (NOT liquified),
is recycled back through the reactor chamber, very little is wasted!
- Hydrogen and nitrogen gases are mixed
in the ratio 3:1 (to fit in with the molecular equation mole ratio) and the gaseous mixture fed
into the top of the reaction chamber.
- The gases are pumped down through the
reaction chamber filled with lots of beds ('shelves') coated in the iron
- The hydrogen and nitrogen gases react on the
surface of the iron catalyst to form ammonia.
N2(g) + 3H2(g)
The atom economy is 100%.
The initial yield is 6%-8%, but
unreacted gases are recycled to raise this to nearer 100% eventually ...
read on ...
To sum up: A low % yield of ammonia is produced quickly at
moderately high temperatures and pressure in the presence of an iron catalyst, and is more economic than getting a higher %
equilibrium yield of ammonia at a more costly high pressure and a slower lower
- Nitrogen (-196oC) and hydrogen (-252oC) have much lower boiling
points than ammonia (-33oC) and stay as gases.
- Boiling points increase with pressure,
but these normal atmospheric pressure values offer a fair comparison and the
higher the boiling point of the liquid, the higher condensation point of
- The temperature in the lower chamber is
never low enough to condense out the unreacted hydrogen or nitrogen so
only the desired product, ammonia gas condenses out, then the liquid ammonia is drained off at the
bottom of the.
- Since the hydrogen and nitrogen are
still gases above the liquid ammonia, they are easily pumped around and
mixed with new hydrogen and nitrogen and hence recycled through the
- This means non of the original
hydrogen and nitrogen reactants is wasted, despite the reaction
being an equilibrium.
- In fact the yield of ammonia can be
as little as 6% conversion, but FAST, and the other 94% of reactant
gases is recycled FAST.
Detailed notes on "Rates
of Reaction" for further reading.
AND there are some more general notes on
Economics on the Industrial Chemistry page.
- Using an effective iron catalyst can
the reduce the cost of manufacturing ammonia by increasing the rate of
reaction (more efficient) and lowering the energy requirements if the
process can be done at lower temperatures (activation energy reduced).
- Increasing the rate of reaction saves time
and operating at a lower temperature saves energy and therefore saves money.
- However, catalysts can be very specialised
and expensive to produce and they get contaminated ('poisoned') and become
less efficient, in this case sulfur compounds contaminate the iron catalyst.
- So the iron catalyst might have to be
extracted and cleaned up, but if a true catalyst (and it is), this'
refurbishment' should enable the iron catalyst to be reused.
- Remember, theoretically catalysts take part in
the reaction, but are not consumed in the reaction and can be reused over
and over again.
- See 'chemical
economics' for other commercial aspects of chemical production.
PART C 3b. NITROGEN FIXATION -
Nitrogen fixation is the
process of turning atmospheric nitrogen molecules (N2)
into useful nitrogenous compounds like ammonia and ammonium salts
(fertilisers) and ammonia is used in the chemical industry in the
manufacture of explosives, pharmaceutical products and plastics.
produced ammonia is used to make fertilisers, so important to increase crop yields to feed
the world's growing population. However the overuse of these
synthetic-artificial fertiliser can cause
The fixation of nitrogen
by e.g. legume plants is quite slow, but using modern chemical
technology e.g. high pressure & temperature reactors and an iron
catalyst, the whole process is speeded up enormously.
Ideally it would be
great to develop catalyst as powerful as enzymes to 'fix' the
nitrogen at room temperature and atmospheric pressure, thereby
saving on energy and more costly reactor technology making the
overall process cheaper and more efficient.
Since huge quantities of
ammonia are manufactured world-wide, using lots of energy AND oil
based hydrocarbons, how sustainable is the Haber process?
The hydrogen is made
from methane, a non-renewable fossil fuel resource. This is a finite
resource, though the nitrogen comes in abundance from the air and
would never run out!
The atom economy is very
good, theoretically it is 100% and in practice the yield is made
very high by recycling any unreacted hydrogen and nitrogen gases.
There are no waste
products because there is only one product possible (no by-products)
and even the unreacted gases are recycled.
The process is costly in
terms of energy needs to maintain a reactor temperature of 200-450oC
and 200-400 atmosphere pressure and these process conditions come
with a range of health and safety issues to do with high
temperatures and pressures - chemical plants are always potentially
dangerous unless managed very carefully.
The process itself does
not damage the environment directly, but misuse/overuse of
fertilisers causes problems, but is this not a separate issue from
the Haber process itself? unless you get into debate about natural
organic fertilisers versus synthetic fertilisers from ammonium
Despite, some negative
issues mentioned, making ammonia is very profitable because of the
huge and increasing world-wide demand for food, especially growing
it rapidly with high crop yields.
Keywords & formulae: The Haber Synthesis of ammonia
H2 + N2 ==> NH3 N2 + 3H2 ==> 2NH3 reaction conditions temperature
pressure iron catalyst recycling unreacted gases
A Reversible Reactions * B Reversible reactions and Equilibrium
* C The
Haber Synthesis of ammonia * D(a) The Uses of ammonia-nitric acid-fertilisers
The nitrogen cycle
Foundation tier (easier) multiple choice QUIZ on ammonia,
nitric acid and fertilisers etc.
Higher tier (harder) multiple choice QUIZ on
ammonia, nitric acid and fertilisers etc.
Advanced A Level Notes - Equilibrium
Advanced A Level Chemistry Notes p-block nitrogen & ammonia
how is nitrogen from the
atmosphere fixed? Revision notes on Haber synthesis
of ammonia reaction conditions KS4 Science
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