In the last episode, we learned that stoichiometry, the mathy part of chemistry that looks at
the ratios between moles in a reaction, works really well with cooking.
But its real-life applications can go far beyond that.
We use stoichiometry in tons of industries worldwide, including fermentation, metallurgy,
and agriculture.
Lots of us use it in our jobs every day.
In fact, one of the reasons stoichiometry is so important is that it helps solve a bunch
of diverse problems.
Today, we’ll be investigating some of the really cool places we might expect stoichiometry
to pop up in,
using some real-world examples to get more comfortable with the stoichiometry steps we
covered last time.
We’ll go over three different kinds of stoichiometry: mass to mass, volume to volume, and volume
to mass, and find out just how helpful multi-step stoichiometry can be.
Hi, I’m Will Comar, and welcome back to Study Hall Chemistry, presented by Arizona
State University and Crash Course.
Let’s dive in!
Let’s start with some history.
In the early 1900s, two German chemists won Nobel Prizes in Chemistry.
In 1918, the first chemist, Fritz Haber, won a Nobel Prize for developing a way to synthesize
ammonia from nitrogen and hydrogen, making it cheap to produce.
And in 1931 the other chemist, Carl Bosch, won his Prize for refining Haber’s process
to produce ammonia in larger amounts in order to create an inexpensive way to make fertilizer.
These advances might not sound all that special, but they in fact played a huge role in the
evolution of industrial agriculture.
The Haber-Bosch Process, as it’s now called, made waves around the world by increasing
crop yields, which in turn created a population boom around the same time.
It also had some unintended consequences outside the realm of just agriculture.
In the first World War, the Haber-Bosch process also lowered the cost of production of ammonia
needed for explosives in Germany, making the war last even longer.
In fact, Haber’s chemical weapon research during the war is why he is considered a war criminal.
It's some pretty bad stuff!
His scientific contributions definitely proved to be a double-edged sword.
Over time, people have started to wonder if producing ammonia with the Haber-Bosch Process
is really that sustainable anyway,
because it requires a lot of energy to produce the amount of hydrogen needed to complete the reaction.
To see what they mean, let’s use mass-to-mass stoichiometry to calculate how many grams
of hydrogen we’d need to produce 321 grams of ammonia.
This process involves using three conversion factors: we convert from mass to moles, then
use a mole ratio to convert to moles of the substance we want,
and then finally convert back to mass.
Starting with our givens, we have 321 grams of ammonia, and the molar mass of both ammonia
and hydrogen in grams per mole.
Then we should convert from grams of ammonia to moles by dividing it by its molar mass.
Our grams will cancel, and we’ll end up with just moles.
Next, and most importantly, we need our mole ratio.
Like I mentioned last time, the mole ratio is a fraction that relates the moles of substances
to each other using their reaction coefficients.
Consulting our balanced reaction, it looks like there are 3 moles of hydrogen for every
2 moles of ammonia.
We’d write this as our next conversion factor, with moles of ammonia on the bottom so that
they’ll cancel out, leaving us with just moles of hydrogen.
All that’s left now is for us to convert from moles of hydrogen to grams of hydrogen
using its molar mass.
When we calculate our answer, we get 56.647 grams of hydrogen, or, rounded to two sig figs, 57 grams.
So it takes that much hydrogen to make 321 grams of ammonia.
To put this in perspective, a little over half the total amount of hydrogen produced
annually in the world is used for ammonia production.
That’s a lot!
Stoichiometry shows us that the Haber-Bosch process does use up a lot of hydrogen, which
makes it pretty unsustainable in the long term.
There are alternatives to fertilizers, though.
African American agricultural chemist, George Washington Carver, developed and popularized
the crop rotation technique, using specific plants at different times to replenish the soil naturally.
Meanwhile, we could use ammonia as a fuel in a step to clean energy.
There’s been talk worldwide among green energy advocates and scientists of using ammonia
as a replacement for fossil fuels.
When we burn ammonia as fuel, we produce nitrogen gas and water, which makes it a carbon-free
and a more sustainable alternative to the emissions of most fuels these days.
It’s also easier to store and has a higher energy density than hydrogen, so it’s powerful
enough to accomplish what we’d need it for.
To investigate this reaction further, let’s move on to the next kind of stoichiometry problem,
where we need to convert from volume to volume.
These work like mass-to-mass problems — we convert the amount of a substance from a unit
of volume to moles,
swap it for the substance we want, and convert back to volume again.
Let’s say we wanted to find out how much gaseous ammonia we’d need to combust as
a fuel with 12.5 liters of oxygen at standard temperature and pressure, or STP.
We’ll use STP for this example to simplify it a bit, so we can get the hang of this process
without making it too complicated on our first go.
In this situation, we’d need to use both stoichiometry and the Ideal Gas Law, which
you might remember is PV=nRT.
This law will be helpful in our example because it shows the relationship between moles of
a gas and its volume.
Specifically, we used the Ideal Gas Law in an earlier episode to find that the volume
of one mole of a gas at STP is 22.4 liters.
We’ll use this ratio, the standard gas volume, as a conversion factor not once but twice
in STP volume-to-volume problems like this one.
It helps us convert from liters to moles and back again.
Let’s walk through an ammonia combustion reaction and talk about a shortcut with some
help from TestTube.
Let’s set up our calculation with our conversion factors.
We know we have 12.5 liters of oxygen as a given, and we need to end up with liters of ammonia.
To do this, we’ll first need to convert from liters of oxygen to moles of oxygen,
using our standard gas volume conversion factor.
We can divide 12.5 liters of oxygen by 22.4 liters per mole of oxygen, and end up with
units of moles of oxygen.
From here, we’ll need our mole ratio to convert from moles of oxygen to moles of ammonia.
In our reaction, we have four moles of ammonia for every three moles of oxygen, so we’ll
write this with moles of oxygen on the bottom in order to cancel them out.
We’re left with moles of ammonia as our units, and now we just need to get that into liters.
We’ll multiply by the same 22.4 liters per mole of ammonia gas, because it can be used
at STP for any gas.
Now, if you look closely at the last calculation, you'll see we divide and multiply by the standard gas volume —
that means technically, we could just cancel the numbers out and use the mole ratio as a liter ratio!
This shortcut works anytime we calculate a volume-to-volume problem at STP.
When we calculate this problem, our answer becomes 16.7 liters of ammonia when we round
to three sig figs.
This is the amount of ammonia we’d be able to burn as fuel based on our given amount of oxygen.
We said earlier that the Haber-Bosch process for producing ammonia requires a lot of energy,
and it’s generally accomplished by burning fossil fuels.
However, chemists at some industrial companies have also been able to produce ammonia for
fuels using wind power, which would make ammonia a more sustainable fuel source.
Smelly, but sustainable.
Let’s move on to the third and final way we can use stoichiometry to convert a substance:
from mass to volume.
Here, we’ll convert from mass to moles like always, then go on to convert moles to volume.
To see this in action, let’s use an example from wine-making, or fermentation.
Fermentation is the breakdown of carbohydrates in an environment where there isn’t oxygen.
To make wine, fermentation takes glucose sugar from grapes and breaks it down to carbon dioxide
gas and ethanol.
Say we wanted to find out the volume of carbon dioxide produced during winemaking if we start
with 325 grams of glucose at STP.
Since we’re at STP, we can count the standard gas volume, along with the molar mass of glucose,
as our conversion factors.
325 grams of glucose somehow needs to become liters of carbon dioxide, so let’s start
with first converting grams to moles of glucose.
We’ll need this in order to convert to moles of carbon dioxide using our mole ratio.
From there, if we know there’s 22.4 liters of carbon dioxide for every one mole
of gas at STP, we can eventually get to our answer in liters.
Knowing this is our path forward, let’s fill in our numbers and do the calculation.
We end up with 80.8 liters of carbon dioxide produced in our wine making, and that’s our answer.
Cheers!
We're going to stop here, but making wine — or really, anything in chemistry — isn't perfect.
So we can also use stoichiometry to calculate how much wine we expected to make and how
much we actually did.
And we'll calculate the theoretical and percent yield next episode!
In multiple step stoichiometry problems, we use conversions to link our known numbers
together to reach our unknowns.
It’s something that can get overwhelming, but one way to streamline this is to compartmentalize.
We’ve said before that taking things one step at a time always helps, and it’s the
same for stoichiometry.
It might help to write everything out, or use a spreadsheet to automate your calculations
to save some time.
No matter how you decide to write it down, hopping from unit to unit with mole ratios
helps us clearly see both where we’ve been and where we’re going.
And chemistry puts tons of different paths and applications for stoichiometry on the map.
Next time, we’ll look at another way stoichiometry fits into the real world: when we make mistakes.
Thanks for watching Study Hall: Chemistry, which is produced by Arizona State University
and the Crash Course team at Complexly.
If you liked this video and want to keep learning with us here in Study Hall, be sure to subscribe.
You can learn more about ASU and the videos produced by Crash Course in the links in the description.
See you all next time!