- [Narrator] Life on Earth needs nitrogen to grow
and plants are no exception.
The good news is our atmosphere is full of nitrogen molecules.
The bad news?
This molecule isn't something plants can easily use to make new molecules.
Compounds such as urea or ammonia are far easier to convert into proteins
and are chemicals commonly found in decaying matter or animal waste.
Throughout history, farmers have given their crops a boost
by grinding up bones or dumping sewage on their fields.
This works just fine for small farms.
But with only so much waste to go around,
more fertiliser is needed to feed large populations.
One method to make large amounts of ammonia for fertiliser
is the Haber process.
Let's see how it works.
In theory, a single molecule of nitrogen
can react with three molecules of hydrogen
to produce two molecules of ammonia.
In practice, the nitrogen molecule has a triple bond that locks it up tight,
making this reaction a real challenge to carry out.
Even if you can get the reagents to combine,
there's the issue of equilibrium.
Any ammonia they make will simply react
to form hydrogen and nitrogen in a reverse reaction.
Make enough ammonia and provide enough energy,
the product will turn back into the starting materials,
like a cookie that unbakes itself.
The Haber process relies on Le Chatelier's principle
to make this process a little more efficient.
Le Chatelier's principle says by changing conditions,
you can move the position of equilibrium to favour one of the reactions.
Nitrogen and hydrogen molecules lose energy when they combine,
releasing it in the form of heat.
Since this reaction is exothermic, cooling the equilibrium mixture down
favours the reaction that produces ammonia
over the reverse one that turns ammonia into nitrogen and hydrogen.
This is a win-lose situation though.
For one thing, Le Chatelier's principle comes with payback.
Increase the forward reaction, you just give more heat to ammonia,
providing it with the energy it needs to react
and turn back into nitrogen and hydrogen.
Even if you keep your cool, a cold reaction is a slow one.
It takes a long time for the ammonia to build up,
even if the ammonia isn't quickly reverting back into the reagents.
In 1909, a German chemist named Fritz Haber
found a more efficient way to combine nitrogen from the air
with hydrogen that could be taken from methane.
The first step is to use a catalyst.
This speeds up a reaction without contributing to the product.
Haber suggested using osmium or uranium,
but later chemists used iron with something like potassium oxide added.
This isn't quite a victory though, not yet.
These catalysts require a temperature
of close to five hundred degrees Celsius to work.
So while the forward reaction picks up speed,
the ammonia also crumbles the moment it's made,
making it hard to extract fast enough.
Applying Le Chatelier's principle to another stress provides a solution.
The forward reaction is made of two different molecules sticking together
while the reverse reaction is a single large molecule splitting apart.
Increasing the pressure on a system at equilibrium
favours the reaction that produces fewer molecules.
So in this case, applying a few hundred atmospheres of pressure
favours the forward reaction.
With the catalyst speeding things up,
the pressure can help keep the ammonia together
until it's cooled down and removed as a liquid.
Of course, forcing two gases together leaves room,
which drops the pressure again.
Le Chatelier's principle always comes at a cost,
but balancing the reagents going in with the product coming out
ensures it all stays in balance.
Haber's process was scaled up by another chemist named Carl Bosch.
The Haber-Bosch process revolutionised not just agriculture,
but the weapons industry with a surplus of new nitrogen-based chemicals.
Taming the ammonia equilibrium literally changed our world.