>> host: Welcome back
to "Chemistry Matters" Unit 9
our Kinetics and Gases Unit.
We learned in the first video
on our Unit 9 playlist
that rates of reaction
can be explained
on a molecular level
using the Collision Theory.
In this video,
we turn our attention
to understanding
how molecular motion
also explains the unique
properties of gases.
Have you ever wondered about
the incredible properties
of gases?
For example, how can hot air
make a 600 pound balloon
rise high above the earth?
And what is air pressure?
There's approximately
one ton of air
pressing down on us,
assuming we're at sea level.
That's about equal
to the weight of a small car!
Why can't we feel it?
It turns out that air
exerts an equal force
in all directions.
Since the same pressure of air
is all around us,
as well as inside of us,
the forces are balanced,
and we feel no net pressure.
We would begin to feel
air pressure if the air
outside of us was at
a different pressure
than the air inside of us.
The effects of a change
in air pressure
can be very apparent.
Air pressure differences
are responsible
for creating wind
and for making your ears pop
while on an airplane.
We can see the dramatic effect
of a change in air pressure
when we create a temporary
vacuum inside a container
like this soda can.
I'll put my goggles on for
safety and show you what I mean.
There's just a little bit
of water in this soda can.
I've boiled that water,
replacing the air inside
with hot water vapor.
When we invert the can
into cold water,
we trap the water vapor,
which condenses
into liquid water,
lowering the pressure
inside the can.
The atmospheric air pressure
outside the can
is now much greater
than the air pressure
inside the can,
which caused the can to crush.
This is an impressive
demonstration.
And you can even
do this yourself.
Just make sure to take
safety precautions,
such as using tongs
and wearing
eye protecting goggles.
Another interesting
property of gases
is their ability to diffuse.
Gases spread out to fill
any sized container
they're put into.
That's why we can smell
an odor created on
another side of a room.
In fact, in order to smell
anything at all,
the odor creating molecule
must be in gas form.
Only a gas can diffuse its way
into the olfactory receptors
in our nose.
By the 19th century,
scientists understood
the macroscopic properties
of gases as described
by pressure, volume,
temperature and quantity,
but they weren't able
to explain why gases behave
in the unique way they do.
Over a period of about
100 years, with contributions
by many scientists,
gases came to be viewed
as tiny molecules
spaced far apart,
moving independently
and rapidly through space,
until they collided
with something.
This microscopic view of gases,
collectively called
the Kinetic Molecular Theory,
explained properties of gases
such as their compressibility
and ability to fill
any container they occupy.
Today, scientists consider
the Kinetic Molecular Theory
a fundamental model
for the behavior of atoms
and molecules in a gas.
The Kinetic Molecular Theory
has five components.
>> host:: Gases are made
of particles that are
very small and far apart.
The vast amount of space
between particles explains
why gases are compressible.
Gas particles are in constant,
random motion.
This random motion allows for
the spreading out
of gas particles until
they are uniformly distributed
in their container.
Gas pressure is caused
by collisions between
gas particles
and the container walls.
The more frequent
the collisions,
the higher the pressure.
As you inflate a tire,
you're introducing
more gas particles,
causing more collisions,
and therefore
increased pressure.
There are no forces
of attraction or repulsion
between gas particles,
which explains why
they're so far apart.
This lack of interaction
also explains why all gases
mix spontaneously
with one another
regardless of their identity.
The average kinetic energy
of gas particles
is proportional
to the temperature in Kelvin.
An increase in temperature
causes an increase
in the kinetic energy
of the particles, which in turn
causes more collisions
with container walls
and higher pressure.
Now that you know the theory
behind the behavior of gases,
let's go back to our classroom
to see gases in action!
>> teacher: Ideal gases
are gases that follow
the behavior described by
the Kinetic Molecular Theory.
The two most important
assumptions about
an ideal gas are that
there are no intermolecular
forces between gas molecules,
and the volume of
the gas molecules themselves
are so small
as to be negligible
compared to the volume
between the molecules.
Of course real gases
do have intermolecular forces
and their molecules
do have volumes.
A perfectly ideal gas
doesn't really exist,
but many gases approximate
the behavior of ideal gases,
which greatly simplifies
the calculations we can use
to describe these gases.
Most gases will behave ideally
under normal conditions,
although under
extreme conditions
of pressure or temperature,
gases begin to deviate
from behavior described by
the Kinetic Molecular Theory.
As pressure increases
or temperature decreases,
gases tend to condensate
to liquids,
and non-ideal behavior
begins to dominate.
As gas molecules condense,
they draw closer together,
causing intermolecular forces
to increase and the volume
between the molecules
to decrease, resulting in
non-ideal behavior.
Gases that do not follow
the ideal behavior predicted by
the Kinetic Molecular Theory
are called real gases.
After conducting
many experiments with gases
under a variety of conditions,
scientists developed
a simple mathematical equation
that predicts the behavior
of an ideal gas.
This equation, called the
Ideal Gas Law,
describes the behavior of a gas
using four variables:
pressure-- P, volume-- V,
temperature-- T, and amount,
or number of moles
of the gas-- n.
R, called the Universal
Gas Constant, 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.
Let's analyze the relationship
between these four variables.
For example, what do you think
will happen to pressure, P,
if we increase n,
the number of moles?
Assume that the other variables,
volume and temperature,
are held constant.
>> student: Oh, okay,
so if moles are increased,
pressure will go up, too.
>> teacher: That's right.
But how would you explain that
on a molecular level?
Think back to the
Kinetic Molecular Theory
to help with this explanation.
>> student: Well, I remember
that pressure is caused
by collisions of gas particles
on the container.
>> student: And more particles
means more collisions
and more pressure, right?
>> teacher: That's
absolutely right.
So, Bella, can you think of
a real-life example showing
the relationship between
the amount of gas
and the pressure it exerts?
>> student: I think I can.
When I use hair spray,
I'm using up the gas inside,
so the pressure inside
the container goes down, right?
>> teacher: That's right, Bella.
And when the pressure gets
too low, you won't be able
to spray anything else out.
Okay, now let's see if you can
make more predictions based on
the ideal gas law.
I'm going to conduct
a number of demonstrations
with gases, and I'd like you
to do two things for each
of the demonstrations...
before each demonstration,
predict what will happen
based on what you've learned
about the ideal gas law.
After each demonstration,
explain the observations
made at a molecular level,
using what you've learned about
the Kinetic Molecular Theory.
Our first demonstration
will explore the relationship
between pressure
and volume of a gas.
Remember that the other
variables described
in the ideal gas law,
in this case temperature
and the moles of gas,
will be held constant.
For this demonstration,
I have in front of me
a vacuum pump, a bell jar,
an inflated balloon
and a marshmallow.
When I place the balloon,
containing trapped air,
into a container,
and remove some of the air
using a vacuum pump,
predict what you think
will happen to the volume
of air inside the balloon.
Remember to use
the ideal gas law, PV = nRT,
to help you make
this prediction.
And be prepared
to explain your prediction.
Now Dory and Abrianna,
what do you think?
>> student: A vacuum
reduces pressure,
doesn't it?
>> teacher: Yes, it does.
>> student: Then I think
when pressure goes down,
the volume will go down, too.
>> teacher: Abrianna,
do you agree?
>> student: Well, yeah, because
if there's less pressure
in the balloon,
then it'll shrink.
>> student: I think
the opposite.
When pressure goes down,
volume has to go up.
>> student: Right.
The ideal gas law shows
that pressure and volume
are inversely proportional
to each other.
So if pressure goes down,
volume has to go up.
>> teacher: Nice predictions.
Let's watch the demonstration
and observe what happens.
[whirring]
>> student: Woah!
>> student: Wow!
Look at it grow!
>> student: That's really cool!
>> student: I guess
I was wrong.
Volume actually does increase
when pressure goes down.
>> teacher: Correct.
Just as Kayla described,
pressure and volume
are inversely related.
[hissing]
Now, let's see if we can
verify this relationship
with the marshmallow.
To make our data easier
to analyze, I've drawn a face
on the marshmallow.
[giggling]
Now, Bella, what do you think
will happen to the marshmallow
when we remove a lot
of the air from the bell jar.
>> student: I think
it'll get bigger
like the balloon did.
>> teacher: Let's try and see.
[whirring, hissing]
>> student: Whoa.
>> student: Wow!
Oh, my gosh!
>> teacher: Now, what do you
think will happen if I let air
back into the bell jar.
>> student: I think
it will shrink back down
to it's normal size.
[hissing]
>> student: Wait.
Why did the marshmallow
shrink down to almost nothing
when you let the air back in it?
>> teacher: The vacuum
removed the air
inside the marshmallow,
as well as the air
surrounding it,
causing the marshmallow
to expand.
But once the vacuum was removed,
air returned to the bell jar,
exerting pressure on
the outside of the marshmallow,
causing it to collapse.
Now that we've verified
the relationship between
pressure and volume,
let's see if each team
can draw a model
that explains what we just saw
at a molecular level.
I'll give you a few minutes
to develop your model,
and then I'll ask each team
to explain their model
to the rest of us.
>> host: Let's take a break here
to give you time to create
your own models demonstrating
the relationship between
pressure and volume in gases.
Like our teacher said,
you'll need to be able
to explain your model.
So make sure you understand
why the gases are behaving
the way you think they will.
When you're done,
rejoin us at the top
of the next video
on the Unit 9 playlist.