>> host: Welcome back to
Unit 9 of "Chemistry Matters."
By now you've created
models that demonstrate
the relationship between
pressure and volume in gases.
So let's go back
to our classroom, where we've
mixed some of our students,
but they're ready
with the models our teacher
asked the students to make
in the last video
on the Unit 9 playlist.
You can see how your models
compare to the models created
by our students.
>> professor: Who'd like
to show their model first?
>> student: We will.
We made a model
of the marshmallow.
There are little air bubbles
in a marshmallow, so when
the vacuum gets turned on
and air gets sucked out,
the air bubbles expand.
As they expand, it makes
the marshmallow expand.
>> professor: That's a good
description.
Another way to think of it
is that when there isn't
as much external air pressure
pushing on the surface
of the marshmallow,
the air bubbles inside
the marshmallow can expand.
How about you two,
Dory and Colin,
Can you describe
your molecular model?
>> student: We explained
the balloon kind of like
the other team did.
Without a vacuum, there's
equal pressure inside
and outside the balloon.
>> student: Yeah, it's kind
of like a shoving match
between the air outside
and inside of the balloon.
When the vacuum's turned on,
the outside air is removed,
but the inside air is trapped
inside the balloon.
>> student: So the air
molecules inside the balloon
push out against the balloon,
and the balloon expands.
>> professor: Great explanation!
You all did well making
models to illustrate
the relationship between
pressure and volume,
as described by
the Ideal Gas Law.
Let's look at
another relationship
in the ideal gas law:
if we increase the temperature
of a gas, what will happen
to the volume?
Kayla, what do you think?
>> student: Well, looking
at the ideal gas law,
PV = nRT,
temperature is on one side
and volume is on the other,
so they're directly related.
>> professor: So you're saying
if temperature increases,
volume will increase?
>> student: Yes.
>> professor: All right,
let's see what happens!
This demonstration will
explore the relationship
between temperature
and volume of a gas.
Everyone put on
your safety goggles.
I have in front of me
an Erlenmeyer flask containing
about 25 milliliters
of water on a hot plate
with a balloon over the opening.
We'll heat the flask of water
to boiling and see what happens
to the balloon.
>> student: The balloon
is inflating!
Why is it doing that?
>> student: It has to be
the water vapor!
>> professor: Let's stop this
lab before the balloon pops.
Does this demonstration
confirm your prediction?
>> student: Yes, when
the temperature goes up,
the volume definitely increases.
>> professor: So what's going
to happen when I cool down
the water vapor in this flask?
>> student: I think
the balloon will shrink!
>> professor: Okay, let's see
what happens when I put
this Erlenmeyer flask
into this ice cold water.
>> student: I was right,
the balloon is definitely
shrinking!
>> professor: Exactly!
Let's leave it in there
a little longer.
What do you think will happen?
>> student: When the temperature
goes down, the volume will
get it even smaller, right?
>> professor: Good prediction.
Let's see.
>> student: That was amazing!
>> professor: Yes, it was.
And depending on conditions,
sometimes the balloon
will even collapse
right into the flask.
Does anyone know why
that would happen?
>> student: It must be
because the temperature of
the gas changes so much,
because of the cold temperature
of the ice water.
>> student: Yeah, so the volume
of the gases in the balloon
shrunk even more.
>> professor: I think you've
figured out a significant
characteristic of gas behavior.
When gases are cooled,
the volume they occupy
always gets smaller.
>> student: So, when temperature
increases, volume increases.
And when temperature goes down,
the volume goes down too.
>> professor: Exactly.
Who thinks they can explain
what we just observed
at a molecular level?
>> student: I'll try.
The kinetic molecular theory
says that kinetic energy
and absolute temperature
are directly related.
>> student: So lower temperature
means lower kinetic energy.
But I'm not sure how
that relates to volume.
>> student: I think I get it.
Slower moving molecules
will have fewer collisions
with the balloon.
Will that cause the volume
to decrease?
>> professor: Yes, it will.
Good teamwork!
Remember that gas pressure
is caused by collisions
of gas particles
against a surface.
If the particles
are moving more slowly,
they'll collide less frequently
against the balloon,
and the lower pressure
inside the balloon will cause
the balloon to deflate.
>> host: I hope
the demonstrations our students
just performed gave you
a better understanding of
the unique behavior of gases.
These behaviors can
be predicted using
the Ideal Gas Law
and can be explained on
a molecular level with
the Kinetic Molecular Theory.
If you want more information
about how to conduct
the balloon demonstration
our teacher just showed us,
you can find instructions in
our Chemistry Matters Toolkit.
Now that we've investigated
the rates of reactions
and the behavior of gases,
you're ready to
challenge yourself.
Let's apply what we've
learned to understand
a real-life problem.
We have an engineering-design
challenge for you
that uses your new knowledge
about the rates of reaction
and about the behavior of gases.
We're going to develop
our own airbags.
And detailed instructions
for this activity
can be found in our
Chemistry Matters Toolkit.
The use of airbags
in automobiles has reduced
serious head injury
from accidents by 85%.
The concept of airbags
was discovered back
in the 1950s,
but wasn't actually
used in automobiles
for another 30 years due
to the complex requirements
for a successful airbag.
An airbag must inflate to just
the right volume and pressure
within 40 milliseconds
following impact in order
to provide protection
during an automobile accident.
And the gas produced must
be harmless both to people
and the atmosphere.
So scientists not only
had to identify a safe,
gas-producing reaction,
but also had to optimize
the rate of the reaction as well
as the pressure and volume
of the gas produced.
It's no wonder that it
took scientists a few decades
to identify and optimize
our current airbag system.
Airbags save many lives
each year, but the chemicals
used to generate the gas
that fills the airbag
are toxic until they react.
We're going to investigate
the suitability of
an alternate chemical reaction
for use in an airbag
so that we don't have
to handle toxic chemicals.
Acetic acid reacts
with sodium bicarbonate
to produce carbon dioxide gas,
water, and sodium acetate.
These reactants are safe
to use and easily obtained
at a grocery store.
Acetic acid is present
in vinegar at a 5%
concentration by mass.
Sodium bicarbonate
is simply baking soda.
Your goal in this
engineering design challenge
will be to optimize
the conditions for
the chemical reaction
between the acetic acid in
vinegar and sodium bicarbonate
so that a plastic bag
is fully inflated as quickly
as possible without rupturing.
In order to do this,
you'll need to measure
the volume of the plastic bag
and use that value
to determine how much
of each reactant you'll
need to fill the bag
with carbon dioxide.
Let's head back to our classroom
to learn more about
our investigation of airbags.
>> professor: We're ready to
apply what you've learned
about kinetics and gases
to a real-life problem,
designing airbags.
We'll need to figure out
how to determine the amount,
or moles, of gas needed
to fill the bag.
Now we can easily measure
the volume of the plastic bag.
But what's the relationship
between volume and the amount
of reactants we need?
>> student: I think
I can answer that.
The more volume
of a gas you have,
the more moles you have.
More gases inflate
a bigger bag.
>> professor: That's right,
Kayla.
How do you know that?
>> student: Well...
it just makes sense.
>> student: I think I know
what you're asking!
We just learned about
the ideal gas law,
PV=nRT,
and that has moles
and volume in it.
>> professor: That's exactly
what I wanted you to think of,
the ideal gas law.
We haven't used
the ideal gas law
to do any calculations,
but now we can.
>> student: But what values
do you use for pressure
and temperature?
>> student: Yeah,
I wondered that too,
and I don't know what number
to use for R either.
>> professor: Those are
all good questions.
Let's look at
the ideal gas law again
one step at a time.
This time we'll be thinking
of how to use it to calculate
the amount of gas in moles.
Let's start by dividing
both sides of the equation
by RT so that we can isolate n,
the number of moles.
RT cancels out on the right
side of the equation,
leaving the equation
expressed in terms of moles,
n equals PV divided by RT.
Since we're doing
the experiment right here
in our lab, under normal
everyday conditions,
the value we'll use
for pressure is simply
atmospheric pressure,
or 1 atmosphere.
And the value we'll use for
temperature is room temperature,
which is about
25 degrees Celsius,
or 298 Kelvin.
Gas laws calculations require
that temperature be in Kelvin.
R is called
the Universal Gas Constant,
and is equal to
0.08206 Liters
times atmospheres per moles
times Kelvins.
Just as a reminder, R serves
as a proportionality constant,
which is simply
a multiplying factor that makes
pressure times volume,
P times V, equal to
moles times temperature,
n times T.
This is a really handy number
to use, but you don't have
to worry about memorizing it.
Now you have all of the values
you'll need to calculate
the numbers of moles of gas
for your airbag.
Colin?
>> student: So do we plug
the values you just gave us
into the ideal gas law
and solve for n, moles?
>> professor: That's right.
The ideal gas law is
a very common equation.
Once you practice,
it'll become familiar to you.
Let's try to do
a sample calculation now
and see what we get
for moles of gas.
We'll pretend that
the volume of the airbag
is 500 milliliters.
Take a few minutes to work
with your partner and see
what value you can calculate
for moles of gas
that this bag would hold.
Okay, you've had
a few minutes to do
the calculations.
Who'd like to share
their results?
>> student: I think I got it.
I calculated that if
the bag holds 500 milliliters,
which is 0.5 liters,
the number of moles of gas
it would hold is 0.02.
>> professor: That's right.
Well done! Can you describe
how you arrived at that value?
>> student: Sure, we started
with the ideal gas law,
rearranged to solve
for number of moles,
just like we all did together
a few minutes ago.
Then we had to find
the right values for pressure,
volume, the gas constant,
and temperature.
>> student: And we just talked
about using normal room pressure
and temperature, so those
are the values we used:
P equals 1 atmosphere
for pressure, and T equals
298 Kelvin for temperature.
>> student: You have to be
careful to use the gas constant
correctly so that all the units
match and cancel out
except for moles.
The gas constant, R,
is 0.08026, and has units
of liters atmospheres
over moles Kelvin.
According to this example,
the bag holds 500 milliliters,
which is the same as 0.5 liters,
so that's the volume in liters.
>> student: So we just made sure
that those were the values
we used for pressure, volume,
and temperature.
We canceled out ATM on the top
and ATM on the bottom.
We also canceled units
of Liters and Kelvin,
leaving "moles" as
the remaining unit.
>> student: Then we did
the math.
In the numerator,
1 times 0.5 equals 0.5.
In the denominator,
we used the value for R,
0.08026,
times 298
and got 23.9.
Then we divided 0.5
by 23.9,
which equals 0.02 moles of gas
needed to fill this airbag.
>> professor: Well done!
I think we're ready to start
the experiment.
>> host: Before our students
begin this lab, let me
remind you of a few things.
We're going to investigate
the suitability of
an alternate chemical reaction
for use in an airbag
so that we don't have
to handle toxic chemicals.
Again, your goal is
to optimize the conditions
for the chemical reaction
between the acetic acid in
vinegar and sodium bicarbonate
so a plastic bag
fully inflates as quickly as
possible without rupturing.
Begin by writing
an experimental procedure,
which must be approved
by your teacher before
starting your experiment.
The procedure should include
proper safety precautions,
details on how to measure
the volume of the bag,
instructions on how
to calculate the amount
of baking soda and vinegar
reactants needed,
as well as moles of
carbon dioxide needed
to fill the bag,
and an explanation
of your plan for timing
how long it takes
to fully inflate each bag.
We'll assume that
carbon dioxide gas
behaves as an ideal gas,
so that the ideal gas law,
PV=nRT, can be used.
And of course, make sure
that you prepare a data table
to record relevant information
during your experiment.
And make sure you take
certain precautions.
For instance, be prepared
for the possibility
of your bag bursting,
which can make a mess.
Wear safety goggles
and an apron.
Carefully consider how you'll
introduce the two reactants
inside the plastic bag.
The reaction occurs
very quickly!
It may be necessary to
encapsulate the baking soda
in a tissue so you have time
to seal the bag
before the reaction begins.
No chemical reaction
is 100% efficient.
The reaction will likely
produce less carbon dioxide
than is calculated
theoretically.
Be prepared to adjust
the amounts of reactants
to compensate for less
than 100% yield.
Once again we'll take
a break in our unit playlist
to give you time
to write a procedure
and prepare a data table
for your experiment.
When you're finished,
come back and we'll compare
your procedures with those
of our students to make sure
you're ready to begin this
engineering-design challenge.