Contributed by Ryan Canning (Y13/December 2002).
Introduction
– Who was Fritz Haber?
Fritz Haber was
the German scientist who developed an efficient way of producing ammonia from
hydrogen and atmospheric nitrogen.
His discovery was
a breakthrough. It meant that ammonia could be produced efficiently and cost
effectively.
Ammonia is the starting chemical used in the production of nitrate fertilisers and warfare explosives. For this reason Fritz Haber will go down in history as one of science’s greatest heroes and greatest villains.
He was born
December 9, 1868 in Breslau, Germany. He
was the son of a merchant. He did
many chemical experiments in school as a child – even at a young age he had a
deep interest for chemistry.
For five years from 1886, at the age of 18, he studied chemistry at three different universities in Germany. After his studies at university he volunteered to work in his father’s chemical business for a while.
Uncertain whether
to devote himself to physics or chemistry, he accepted an apprenticeship with
Hans Bunte, professor of Chemical Technology, Karlsruhe.
He worked alongside him until 1911 during which time he focused on
combustion chemistry and the study of petroleum.
Haber studied the
decomposition and combustion of hydrocarbons in his own free time and in 1906
was appointed a Professor at Karlsruhe to study these subjects.
In 1933 Nazi laws
insisted that all Haber’s staff at the institute resigned.
Instead of agreeing to this, Haber himself resigned.
On invitation he worked in Cambridge for some time until his health
deteriorated, not helped by the English winter.
In 1898 he
published his book on Electrochemistry and in the same year he reported his
results on electrolytic oxidation and reduction.
From 1899 to 1917
Haber studied many areas of science including the electrolysis of solid salts,
Bunsen flames and energy lost during operation of steam engines, motors and
turbines. He always tried to use
his knowledge of science for a purpose.
But the greatest
accolade of Haber’s career came in 1918 when he received the Nobel Prize in
Chemistry for successfully devising a method of obtaining ammonia from reacting
nitrogen fixed from the air and hydrogen obtained from methane under certain
optimum conditions of temperature and pressure.
Read on to find out more.
Haber invented
such things as the glass cathode used in electrolysis, the firedamp whistle used
for the protection of miners and the quartz thread manometer used for recording
low gas pressures.
When
the First World War broke out Haber was appointed as consultant to the German
War Office. He was in charge of
organising poison gas attacks and defending against ally gas attacks.
The poison gas used was formulated by Haber; another invention of his!
So naturally enough he was the man the German government called in to
help.
Haber died on
January 29, 1934, at the age of 66.
He is remembered
not just as a man of science but of economics, politics and industry.
The Institute of
Physical and Electrochemistry where he had worked for most of his life was
renamed The Fritz Haber Institute after his death.
Here
is a picture of it:
The
Fritz Haber Institute
Fritz
Haber and Albert Einstein
Two
German Postal Stamps
in
remembrance of Haber
The
Nitrogen Problem
From 1850 onwards
the demand for nitrogenous fertilisers rapidly increased.
With world populations growing the demand for food increased, putting
pressure on natural nitrate sources such as manure.
The demand for
nitrogen compounds for making nitric acid in the chemical industry grew around
the same time, as Germany was about to enter world war one.
Nitric acid was used to make explosives such as TNT and dynamite amongst
others, to fuel the war. Natural
reserves of nitrates such as Peruvian guano, obtained from sea-bird droppings,
and Chilean sodium nitrate dwindled.
The Solution – The Haber Process
Fritz Haber, as you already know, was the man who came up with the solution to the world’s problem. The production of nitric acid requires ammonia (NH3) as a starting material. He realised that atmospheric air was rich in nitrogen, approximately 78% by volume. But air is a mixture of gases, 20% oxygen and under 2% of others such as carbon dioxide, carbon monoxide, sulphur dioxide, argon, helium, etc.
Pure nitrogen is
required for the manufacture of ammonia. It
is obtained by the fractional distillation of liquid air (under pressure).
Clean air is fed
into a compressor and cooled by refrigeration. The cold air expands through a
nozzle and is cooled still further; enough to cause it to liquefy. The liquid
air is filtered to remove solid CO2 and hydrocarbons and then
distilled. Liquid air enters the top of the column where nitrogen, the most
volatile component (lowest boiling point), passes off as a gas. In the middle of
the column gaseous argon is removed. Liquid oxygen, the least volatile
component, collects at the bottom. The normal boiling points of nitrogen, argon,
and oxygen are 77.4, 87.5, and 90.2 K, respectively.
This diagram
simplifies what happens:
As well as
nitrogen, hydrogen is required for the manufacture of ammonia.
Hydrogen is usually obtained from either natural gas (methane) or naphtha
(a fraction of crude oil containing 5-9 carbon atoms), by either catalysis with
steam or by partial oxidation with oxygen:
Catalysis with
steam (in the presence
of a nickel catalyst).
C6H14(g) + 6H2O(g) ®
6CO(g) + 13H2(g)
Naphtha
CH4(g) + H2O(g) ®
CO(g) + 3H2(g)
Methane
Partial oxidation with oxygen
C6H14(g) + 3O2(g) ®
6CO(g) + 7H2(g)
Naphtha
CH4(g) + 1/2O2(g) ®
CO(g) + 2H2(g)
Methane
The
reaction between nitrogen and hydrogen
The
reaction between magnesium ribbon and oxygen when the magnesium ribbon is
heated, is an example of an irreversible reaction. All of the magnesium will change to white, powdered magnesium
oxide.
A reversible reaction is one which can take place in both directions, forward and reverse and does not go to completion.
The
reaction between nitrogen and hydrogen is a reversible one.
When they are mixed under appropriate conditions nitrogen and hydrogen
will react. The forward reaction
rate will be greatest at first when reactant concentrations are greatest.
As more ammonia is formed the reverse reaction rate will increase and the
forward reaction rate will decrease until they are of the same magnitude (same
rate but in opposite directions). At
this point we can say the reaction is in a state of dynamic equilibrium.
A dynamic equilibrium is where the forward and reverse changes are occurring at the same rate thereby cancelling each other out and creating no overall change.
A
dynamic equilibrium can only exist in a closed system.
A
closed system is a system where there is no interchange of particles between the
reacting mixture and the outside environment.
An
unopened bottle of coke is an example of a closed system.
Nitrogen and hydrogen will not react under normal circumstances. Special conditions are required for them to react together at a decent rate forming a decent yield of ammonia. These conditions are:
A catalyst is a substance which speeds up the rate of a chemical reaction but remains unchanged itself at the end of the reaction.
A
promoter is a substance which increases the efficiency of the catalyst.
But why are these specific temperatures and pressures favoured?
Le Chatelier was a French chemist. This is his principle:
The position of the equilibrium of a system changes to minimise the effect of any imposed change in conditions.
Put
simply, this means that if a change is made in the conditions of a system in
equilibrium, the system will try to counteract the imposed change and therefore
the conditions will become more like they originally were, before any change was
imposed.
This
is the balanced equation of the reaction producing ammonia:
N2(g) + 3H2(g) = 2NH3(g) DH = - 92kJmol-1
Firstly, the
enthalpy change in the forward reaction is negative.
The enthalpy of a
chemical is the energy contained within its structure.
All chemicals
want to lower their enthalpy. In
this case the negative enthalpy change indicates that the forward reaction
producing ammonia is an exothermic reaction.
An exothermic reaction is a reaction which gives out heat energy to the surroundings.
By collision
theory increasing the temperature of the reactants, nitrogen and hydrogen, would
increase the forward and reverse reaction rates to the same extent and
equilibrium would be reached quicker. However
the yield of the product, ammonia, at equilibrium would be reduced because by
increasing the temperature, the reverse reaction rate would increase to minimise
the effect of the increase in temperature by taking heat in.
So increasing the
temperature does allow equilibrium to be reached faster but the yield of the
product, ammonia, is reduced.
400°C is the best compromise temperature.
From the balanced
equation it can be seen that 4 moles of reactants produce 2 moles of product.
A mole is the amount of a substance in grams which has the same number of particles as there are atoms in 12g of carbon-12.
Avogadro’s law
states that equal amounts in moles of all gases occupy the same volume under the
same conditions of temperature and pressure.
Therefore the 4
moles of reactants will occupy twice the volume of the 2 moles of product.
So by increasing
the pressure of the gases the concentration of ammonia at equilibrium will
increase because by Le Chatelier’s principle the system will attempt to
minimise the effect of the imposed increase in pressure by decreasing the
pressure. This happens by increasing the forward reaction rate.
By
collision theory increasing the pressure of the reacting gases increases the
forward and reverse reaction rates meaning equilibrium is reached faster.
By increasing the pressure the gas molecules are pushed closer together
meaning collisions are more frequent. Any
factor which increases the frequency of collisions always increases the reaction
rate.
However
the capital cost of producing an ammonia synthesis plant with machinery capable
of withstanding high pressures is great. High
pressures wear down machinery prematurely meaning maintenance costs are also
high.
200
atmospheres is a compromise pressure.
The
iron catalyst does not effect the position of the equilibrium, so it has no
effect on the concentration of ammonia product.
What it does is it increases the rate of the forward and reverse reaction
rates to the same extent meaning equilibrium is reached faster.
In big business time is money so the quicker the ammonia is produced the
better.
The
iron catalyst is finely divided for maximum surface area, increasing reaction
rate.
The
proportion of hydrogen to nitrogen in the reaction mixture is monitored to
ensure they are in the 3:1 ratio demanded by the balanced equation:
N2(g)
+ 3H2(g)
= 2NH3(g)
DH
= - 92kJmol-1
If
the proportion of hydrogen to nitrogen in an ammonia plant’s reaction mixture
was say, 5:1, then there would be an excess of hydrogen, much of which could
never react to produce ammonia as there would not be enough nitrogen available
to react with it. This would be a
serious waste of the space available inside the reaction chamber.
For this reason the ratio of reactants is carefully monitored and
adjusted.
On each pass
through the reaction chamber the conversion to ammonia is about 15%.
The mixture of nitrogen, hydrogen and ammonia is cooled in a refrigerator
to about -50°C.
Nitrogen and hydrogen have boiling points of about -190°C so they remain in the gas phase and are pumped back
to the reaction chamber for another chance to react.
But ammonia has a boiling point of –30°C.
It condenses and liquefies and is run off as the product.
Summary
 |
Contributed
by Ryan Canning (Y13/December 2002).
Email Ryan: bigtrig99@hotmail.com