[MUSIC]
Lighthouse Scientific
Education presents a lecture
in the Gases series
The topic; The Basics of Gas.
Material in this lecture
relies on an understanding of
the previous lectures
Fundamental
Properties of Matter;
Dimensional Analysis;
Temperature
and The Mole as a Quantity.
Gas is a broad subject.
Prior to the beginning of
this lecture, an overview
of the subject matter covered
in the 'Gas' series
of lectures is given.
This particular lecture begins
with a general overview of
gases with an introduction to
the concept of an ideal gas.
And then it is on to
a discussion of gas
as particles, with
an introduction
to the kinetic
molecular theory.
The 4 basic properties
of gases are presented
and then explained
in more detail.
The properties are
amount; temperature,
volume, (along with
some conversions between
volume unit types) and
pressure (along with
some conversions between
pressure unit types).
And finally, we'll
talk a bit about
the standard temperature
and pressure.
There are four lectures in
the 'Gas' series of lectures
including this one
'The Basics of Gas'.
Its contents were just listed.
The second lecture in the
series covers the
'Fundamental of Gas Law'.
Expect to find named laws like
Boyle's law, Charles'
Law as well as
the Combined Gas law.
The third lecture is
also a gas law lecture.
It is called 'Gas
Law: One Condition'.
The all-important Ideal
Gas law is discussed as
well as Dalton's Law
of Partial Pressure.
The final lecture is
'Gas: Advanced Concepts'.
Topics in this lecture may
or may not be not covered
too much depth in your course.
The lecture is presented
for those students whose
instructor does
delve deeper into the
kinetic molecular theory
or expect some understanding
of diffusion and effusion.
There is one more advanced
topic involving gases that
is not covered in this
series of lectures.
That is stoichiometry
involving gases.
It is presented in the lecture
'Stoichiometry: Advanced
Concepts' which is found
in the in Stoichiometry
Series of lectures.
It covers stoichiometry
using molar volume
and the Ideal Gas law. Now
to an overview of gases.
By now we are familiar with
gas being a state of matter.
And that gas particles
are in random motion
and that motion
is fast and fluid.
Gas particles speeds
up when heated.
Unlike solids, gases
have no defined shape.
The gas particles themselves
occupy very little volume
but travels within
all available space.
Gas expands into
its surroundings.
Gas particles exert pressure
on their surroundings.
Gas is readily compressible
unlike liquids and solids.
As such, its density is
a small fraction of that
of a solid or a liquid.
Gases form
homogenous solutions.
That is they mix completely
with other gases and the
do so in all proportions.
Gases are miscible.
What about
intermolecular forces?
Do gas particles
form dipole dipole or
experience London
dispersion forces such as
those seen in
liquids and solids?
Well, gases can
undergo chemistry so the
answer has to be yes.
On the other hand
when dealing with the
properties of a gas,
gas particles can be
treated like rapidly moving,
and colliding, tiny
balls that have few to no
intermolecular interactions.
So if they are not undergoing
chemical reactions they are
modeled like colliding balls.
There is a model for
treating gases like that.
It is called an ideal gas.
Ideal because it is theoretical
and not a completely
accurate portrayal of a gas.
Still, in many instances it
does stands in quite well for
a real gas and it is quite
useful for developing a basic
understanding of gases.
An ideal gas consists of
randomly moving
particles that collide but
do not interact.
More specifically
the model has 3
basic assumptions.
The first is that a gas
is composed of particles.
These particles can
be compounds or atoms.
They are thought of as
small hard spheres that have
no volume and do
not participate in
intermolecular interactions.
An analogy is billiard
balls on a pool table.
The balls collide with
each other but there is
no sticking together
or alteration to their
size or shape.
The second assumption
is that particles in a
gas move rapidly and are
in constant random motion.
They move in straight
lines and only change
their direction
upon collisions.
After collisions they will
move in a new straight line
until their next collision.
Collisions of a gas
are said to be
perfectly elastic.
That is, upon collision,
the energy of motion
is transferred without loss.
In the billiard ball analogy,
when a 'break' occurs
at the start of a game
the energy of the cue ball
is transferred into the motion
of the other balls.
No energy is lost, it is only
communicated to
the other balls.
These three assumptions
of an ideal gas
form the basis of a
powerful model for reasoning
out gas behavior.
It is called the
Kinetic Molecular
Theory or KMT.
The purpose of the KMT is
to provide an understanding
of gases as particles.
That understanding comes
from a few basic
attributes of the model
which follow straight
from an ideal gas. At the
foundation of the model is
the reasonable proposition
that a gas consists of
large number of particles.
The particles move
in constant, random,
straight-line motion.
They speed up when heated.
Particles are separated by
large distances which means
that gases are mostly empty
space and the particles
can be treated as essentially
taking up none of that space.
That is, the volume of
the gas particle is so much
smaller than the volume of
the container holding the gas
that the volume of the
particles can be neglected.
The gas particles are
moving so fast that there is
essentially no
intermolecular interaction.
And finally, particles make
elastic collisions with each
other and with the
walls of their container.
The KMT is an ideal
gas theory and a lens to
bring gas behavior into focus.
We can use these
considerations to build
an understanding of the 4
basic properties of gases.
Specifically, we will
use the considerations
when asking the broad question
'What properties are
needed to describe a gas?'
and the follow up question
"what would increase the
number of collisions
that a gas makes with the
sides of its container?"
Both of these questions
help us form a behavioral
understanding of gases.
The most obvious answer
to 'what would increase the
number of collisions'
is an increase in the
number of particles.
That is the property of
amount as in the number of
moles of gas (lower case n).
Quite simply more particles
leads to more collisions.
More cars on the road
leads to more collisions.
Another factor that increases
the number of collisions
is the speed of the particle:
That property is temperature
(capital T). Increasing the
temperature (hotter particles)
means faster moving
particles and that leads to
more collisions.
And then there is the
factor of the size of the
container: the property
is volume (capital V).
With a smaller container
the sides of the wall are
going to be closer together.
It takes less time for the
gas particles to reach the side
and that means they will
collide with the side
more often. Larger
containers see less collisions
because particles
have farther to travel.
The last property is found
with a third question.
What is the measure
of 'collision with
the side of the container'?
It is called pressure; more
collisions is higher pressure.
As we will see later,
harder collisions are also
higher pressures.
Pressure is the only one of
the properties that is a
result of the other properties.
By focusing on pressure we
can reason out the roles of
the other properties
in gas behavior.
So these are our 4
properties, amount,
temperature, volume and
pressure: n, T, V, and P.
Understanding these
properties is essential
to understanding gas.
Let's look a little closer
at each one in turn
beginning with amount.
Like most areas in chemistry,
the study of gas uses
the mole as the unit of amount.
Amount is either given
in moles or converted
into moles.
Remember, 1 mole of
particles is 6.022 times
10 to the 23rd particles.
Some textbooks do not
use 6.022 but rather 6.02.
Not a big deal. Use the
one from your textbooks.
Moles can be gotten
from mass or weight with
the atomic or molar mass.
Making this type of
conversion is a necessary
skill at this point
in our studies.
The variable 'n' is used to
represent the number of moles.
The next property
is temperature, T.
Let's take a step
back; re-familiarize
ourselves with what we
learned in the 'Temperature'
lecture in the Introduction to
Chemistry series of lectures.
It stated that the
components of matter,
be they atom or compound,
are in constant motion.
That is true for
solids, liquids or gases.
Importantly, motion is energy
and motion energy
is a form of heat.
They are linked,
more motion comes
from more heat.
Now to temperature,
temperature is a
measurement of
the amount of heat.
Higher temperatures
means more heat and
faster gas particles.
From a collision
perspective, the faster
a particle moves the
more often it will collide
and the harder it collides.
Faster moving give rise to
higher pressure. Conversely,
cooling a gas down
will slow the particles
and reduce the temperature
and pressure (all
things being equal).
Also from the temperature
lecture we saw that there are
different temperature scales.
These scales all measure the
same heat but do so with
different size increments.
The most common scale in
the United States is the
Fahrenheit scale. This older
scale has the freezing point
of water at 32 degree F
and boiling point of water
at 212 degree F.
The most widely used scale
is Celsius. It was designed
so that temperature at which
water freeze is 0 degree C
and the temperature of
boiling will be 100 degree C.
The temperature scale
for science is Kelvin.
It is the absolute
scale of temperature
with the lowest possible
temperature being 0 K
(no degree sign with Kelvin).
All chemistry
calculations use
the Kelvin scale but
temperature is almost
always measured in Fahrenheit
or Celsius.
Conversion between
temperature scales is
covered in detail in the
'Temperature' lecture.
Briefly, the inter-conversion
between Celsius and
Kelvin uses the following
2 equations.
To convert from
Celsius into Kelvin, take
the temperature in Celsius
and add 273.15
To convert from Kelvin
into Celsius, take the
temperature in Kelvin
and subtract 273.15.
The next property is the
well known quantity; volume.
Its scientific definition
is that volume is the
quantity of three-dimensional
space enclosed by
a boundary.
The boundary can be
real or made up (arbitrary).
The volume of a gas is not
the volume of the gas
particles but rather the
volume available to the
particles as they zip around
in random motion.
Within any volume the
actual gas particles will
occupies very little of
that available volume but the
particles travels within all
available space.
Gas expands to fill the volume
containing it. It is
important to make the
distinction between the
volume of the particles and
the volume of the available
space. KMT has the volume
of particles as 0 so any
discussion involving volume
will be referring to
the available space.
Volume is discussed
in more detail in the
'Density: In the Lab'
lecture in the Introduction to
Chemistry series. Briefly,
volume has 3 dimensions.
For a rectangular box
those are labeled
length, width and height.
If each is measured in
meters then the units of
volume are cubed meters;
3-dimensional.
The SI, or standard
international, unit for volume
is cubic meters.
But there are other
commonly used types of
units besides cubic meters.
There is cubic centimeters
and cubic decimeters.
These are all length based
measurements and are used
when they are appropriate
for the size of the
volume of interest.
Another common type of
volume unit is the liter.
It is given the
unit designation l.
It can be an upper
or lower cased l and
sometimes it is italicized.
Here is a smaller liter
unit, the milliliter.
All the units are 3
dimensional representations
of volume.
Sometimes it will be
necessary to move from
one type of volume unit
to another type. Fortunately,
there are well established
relationships between
most of these unit types.
Perhaps most useful is that
1 ml is equal to 1 cm cubed.
This allows for the
conversion between volume
measurements one might get
from a graduated cylinder
to those from a ruler. And
there are metric conversions
within a unit type such
as 1 liter equals 1000 ml.
Combining these two
equalities gives another very
helpful conversion:
1 liter equals 1000 cubic cm.
It should also be noted
that metric conversions with
cubic length measurements
can be a bit tricky, since its
really a 3 step dimensional
analysis problem.
It takes a million
cubic centimeters to make
1 cubic meter even though
it only takes 100 cm to make
a meter. We'll get to
that issue momentarily.
This list of is
nowhere near complete.
There are many other kinds
of relationships between
different volume unit types
including all the common
US variety (such as ounces,
pints, quartz and gallons).
Still, all conversions between
different unit types use
the standard conversion
practice as covered in the
'Dimensional Analysis' lecture.
If a student is given a volume
of 1 cm cubed and asked to
convert that to volume in
liters a relationship is needed
between the current or have
units of centimeters cubed
and the desired or
wanted units of liters.
When inserted into
the equation as a
conversion factor with
current or have units in the
denominator the
conversion can be made.
Another slightly more
complicated example has
that same one cubic
centimeter as the giving unit
but cubic meters as the
desired or wanted units.
A relationship is needed
between the have units of
cubic centimeters in the
wanted units of meters.
inserting this relationship
as a conversion factor with
cubic centimeters
in the denominator
and canceling units will
complete the conversion.
But what happened if this
particular relationship
was not available
to the student?
The answer is that a simpler,
more common relationship
can be used. Such as
the metric conversion
of 1 m to 100 cm.
The problem here is that
the giving units are
in cubic centimeters.
As written this conversion
factor will not completely
cancel out with the
cubic centimeters.
That can be remedied by the
addition of two more of the
same conversion factors.
Taken as a whole,
the 3 cm units of the
conversion factor cancels
out the 3 cm units of the
cubed centimeter
in the giving units.
This is why a metric
conversion within the
cubic length measurement
can be a bit tricky.
The fourth property, and the
one that is really just of
combination of the first
3 properties, is pressure.
Understanding how
these other 3 properties
affect pressure will give
us a handle on the more
advanced topic of gas law.
Broadly speaking, pressure
is a force exerted by a
substance on another substance.
This force is over
some specific area.
The formula of pressure is
force over area; F over A.
Pressure is directly
proportional to force.
A stronger force generates
a stronger pressure.
A weaker force produces
a weaker pressure.
In the context of a gas,
the force is generated by
gas particles colliding with
the side of its confinement
or container or really
anything they strike.
Relating this definition
of pressure, the gas is the
first substance that exerts
a force on 'another substance'
which is the sides
of the container.
So, a gas can generate a
higher pressure though a
change in the
other 3 properties.
This can be viewed from
the perspective of how
changing these properties
alters the force exerted
by the gas or the area
over which the gas exerts
that force. Looking at
the role of force first.
According to the equation
if force goes up then
pressure goes up.
Force will go up if the
number of particles
is increased.
If there are more
particles there are more
collision with the wall.
More collisions is more force.
More force is more pressure.
And then there is
temperature. If heat is
added, the temperature
will go up and the particles
will have more motion.
The particles will
be moving faster.
Faster particles
make more collisions.
Faster particles also
strike the side of
the container harder. Harder
collisions exert more force.
Taken together, adding heat
increases the temperature;
increases the force;
increases the pressure.
Then there is area.
According to the
pressure equation,
area and pressure are
inversely proportional.
That is if one goes
down the other goes up.
The lone property that
affect area is volume.
As volume goes down
area goes down and
pressure goes up.
Decreasing the volume
shrinks the area of the
container and decreases
the distance between the walls.
With less distance to travel
before reaching the walls
the particles will in effect
strike the wall more often.
More collisions-more force;
more force-higher pressure.
Pressure can be thought of
as a reflection of the number
of collisions with the
walls of the container.
For completion sake we will
ask the opposite question.
How can pressure be lowered?
Starting again with force
and noting that according
to the equation if the
force goes down the then
pressure goes down. Force
will go down if the number
of particles is decreased.
If there are less
particles then there will be
less collision with the wall.
When the temperature is
decreased, that is heat is
removed, the particles will be
moving slower and there will
be less collisions with the
walls of the container. Those
fewer collisions will also be
'softer' for lack of a
better word than those of
faster moving particles.
Think of being struck by a
baseball. Which one would
hurt more; the ball moving
fast or the ball moving slow?
Slower particles
have less force.
As for area the pressure
equation says that as the
area goes up the
pressure will go down.
Area increases with
an increasing volume.
A large volume means that
the walls of the container are
farther away and gas
particles have further to
travel before striking the wall.
That decreases the number
of collisions with the wall.
Less collisions
is a lower force
which produces
a lower pressure.
The take home lesson for
pressure is that amount,
temperature and volume
affect the number of
collisions and that in
turn affects the pressure.
Getting a good grasp of
pressure is so important to
understanding of gas that we
will further our exploration.
A common way to
demonstrate pressure is
with a cylinder and piston.
Shown here is a cylinder that
is sealed at the bottom.
Toward the top
is a disc or piston that
snugly fits into the cylinder.
It is assumed that any gas
beneath the piston cannot
escape out of the cylinder.
In this cartoon representation
there is some volume
of gas in the cylinder
under the piston.
The gas particles in the
cylinder are rapidly bouncing
around and exerting a
pressure on the sides
of their confinement and,
importantly, on the
underside of the piston.
In this example we will say
that that pressure exerted
by the gas is 1 atmosphere.
We will get to the
units of pressure shortly.
Question, why doesn't that
1 atmosphere of pressure
push the piston
up the cylinder?
The answer is that while
there is a force on the piston
from the gas in the cylinder
There is also air from the
atmosphere pushing down on
the cylinder from above.
With gases the pressure
of the air often
has to be considered.
The bottom line is that the
piston doesn't move because
the force on top is matched
or is equal to or is countered
by the force exerted by the
gas below in the cylinder.
To keep the example as
simple as possible we will
pretend that the piston itself
has no mass and therefore
no weight. Weight would
also be a force pushing down.
Now, if the external
pressure was to increase
(it's shown as 2 atmospheres
here) the piston would move
because the top pressure
would be greater that the
pressure inside the cylinder.
It would decrease the volume
available to our sample of gas
which we just saw would
effectively increase
the pressure of the gas.
The piston would move until
the pressure on top of
the piston is equal to the
pressure exerted by the
gas in the cylinder on the
bottom of the piston.
This example would not
work if there was a liquid
or solid in the cylinder.
Gases are compressible.
We can take this understanding
of pressure and bring
in the mercury barometer
The barometer is a device
that is used to measure
atmospheric pressure.
In fact the suffix 'baro'
means pressure.
It should be point out here
that our air is a collection
of gases that exerts
a force on us (and
everything else)
as if we were the walls
of a container. We do not
sense that pressure but
it most certainly is there
and the barometer will show
us just how strong it is.
A basic barometer has
a pool mercury, Hg.
Modern barometers use an
alcohol but the principle
is the same. Inserted
into the pool is an
evacuated tube. All the
air has been removed.
It is empty or close to it.
Exposing the barometer
to the atmosphere has
air molecules colliding
with the pool of mercury just
like they collide with
everything else they
come in contact with.
The force of these collisions
has the effect of pushing
down on the surface
of the mercury so that some
of the mercury gets pushed
up the empty tube.
The larger the atmospheric
pressure the greater the
push and the higher the
mercury will climb up the tube.
At some point
the weight of the
mercury, due to gravity,
equals the pressure
exerted by the atmosphere
and the mercury stops rising.
Therefore, the distance from
the top of the pool of mercury
to the top of the column of
mercury is an indicator of
the atmospheric pressure.
Usually the distance is given
in millimeters. In
essence the barometer is
equating the force due
to weight (of the mercury
as measured by the height
of the column of Hg) to the
force exerted by the collision
of gas particles on the
pool of mercury.
More collisions means
more force which is
compensated by more weight
from more mercury in the tube.
The atmospheric pressure
at sea level pushes the
column of mercury to 760 mm.
760 mm Hg is called
the 'standard pressure'.
Sea level is the standard
altitude for pressure and
like all standards it is
just an agreed up amount.
Noting that the atmospheric
thins with increasing
altitude it follows that
the atmospheric pressure at
the top of a mountain is
less than it is at sea level.
Like most types of
measurements there
are different units that
can be used to describe
the measurement. For
example, that 760 mm Hg
is equal to 1
atmosphere of pressure.
And that brings to our next
topic regarding pressure.
Units. There are a number
of different units that are
used to describe pressure.
The different units are all
measuring the same
pressure but they are
presenting it in
different ways.
Let's compare the
more common unit types.
We just saw the
unit mm of mercury.
The units are a height
measurement in millimeters.
As a pressure unit it is
mm Hg. Standard sea level
pressure for this
unit is 760 mm.
A similar unit is torr.
It is name after Torricelli the
co-inventor of the barometer.
It is abbreviated as torr
and it shares the same
increment as mm of mercury.
Standard pressure:
760 torr is equal to
760 mm of mercury.
The two only differ in name.
Another prominent scientist
in the development of the
barometer is Pascal. He too
gets a unit named after him.
Its units are Newtons
over meters square. It is a
force (Newtons) per area
(squared meters) scheme.
It is abbreviated Pa or, if
the scale is multiplied by
1000, kPa. The k is for
kilo. And as we see here
that same standard
pressure is given
with a more complicated
series of number.
Pascals and kPa are not
particularly common unit
at this level of chemistry.
What is common
is the atmosphere.
It is abbreviated 'atm' and
it is a unit defined as being
1 standard pressure.
That is convenient.
The last unit type of pressure
is the most commonly used
scale in the US. It is
pounds per square inch.
It too has is a force (pounds)
per area (square inch) unit
and its abbreviation is 'psi'.
The standard pressure
in psi is 14.1.
The nice thing about this
table is that since each
type of pressure unit
is given at that same
standard pressure all
values in this column
are equal to each other.
760 mm of mercury is
equal to 1 atm. And when
an equality is made a
unit factor can be made.
760 mm Hg over 1 atm is a
conversation factor between
denominations of
the pressure unit.
This ratio can just as
easily have atmospheres in the
numerator and mm of
Hg in the denominator.
Any pair of values in the
column can be made into
an equality and a unit factor.
760 torr over 14.1 psi
is equal to 1 and it
is a conversion factor.
Same goes for
atmospheres and Pascals.
These unit factors are
central in the conversion of
pressure units. Pressure
conversions are a common
homework or test problem
in the early study of gases.
Say we are asked to
convert 570 mm Hg into an
equivalent amount
of atmospheres.
Our standard
approach is to start
with what the problem
gives as the current units
and through a conversion factor
transform that value in to
the desired or asked for units.
In this particular problem
the 570 mm of mercury
is the current or given units
and atmospheres is the
desired or asked for unit.
The conversion factor will
need to be oriented such that
the current units are in
the denominator and the
desired units in the numerator.
In this orientation the
current units will cancel out
leaving the desired units.
To get the appropriate
unit factor we return to the
list of pressure units.
Find the current and
desired units in the column
of standard pressure.
These values are equal to each
other and can be made
into a unit factor.
Of particular importance is
to arrange the ratio such that
the current units are
in the denominator.
That is not the case with
this unit factor. The
current units in our problem
are mm Hg not atm.
No problem, simply flip the
ratio over and the correct
conversion factor is obtained.
Once this unit factor is
inserted the problem reduces
to some calculator work.
Units of mm Hg cancel out
and 570 divided by 760 is
0.750. 570 mm of Hg
is the same pressure
as 0.750 atm. One more,
What is 100 kPa
expressed in torr?
The current unit is
the 100 kPa and the
desired unit is torr.
What is the unit factor
that allows this conversion?
Returning to the table
of pressure units.
Identify the units in
the conversion in the
standard pressure
column; torr and kPa.
Be sure to get the
appropriated Pascal unit.
With these units create
a unit factor with
the current units of kPa
in the denominator so that
cancel out. 100 kPa
is the same pressure
as 752 torr.
The last topic of the
lecture is something called
standard temperature
and pressure (STP).
Standard temperature and
pressure is an agreed upon
set of common values.
The P in STP
was covered in the
standard pressure column of
the units table.
We saw that 1 atm is an
exact value for
standard pressure.
In fact that is
how it was defined.
It can just as rightly
be given as 760 mm Hg.
The standard temperature
is 0 degree C. That too
is an exact figure.
It corresponds to 273.15 K
to 5 significant figures.
An important relationship
at STP is that 1 mole of gas
occupies of 22.4 liters.
It doesn't matter
what the gas is.
This relationship
between amount and volume
is called 'molar volume'.
To restate, at STP, a gas
has a temperature of 273.15 K,
its pressure is 1 atm
and 1 mole of the gas
occupies 22.4 liters.
A quick note, molar volume
is not limited to the
temperature and pressure
at STP. It can be
thought of as 'moles
per liter' of that gas.
Recapping the lecture.
Gas is a state of matter.
Gas particles are fast moving
and occupy very little volume.
That is they physically
occupy a small amount of the
volume available to them.
Gas form homogenous solutions.
They mix completely
with other gases.
Gases are miscible.
One of the most helpful
ways to consider gas is with
the Kinetic Molecular
Theory (KMT).
It is built of the
assumptions of an Ideal Gas.
1. Gas is comprised of
particles. Those particles
are considered
small hard spheres.
They occupy no volume and
they do not participate
in intermolecular interactions.
2. Particles in a gas
move rapidly in constant
random motion. They move
in straight lines and
only change direction
upon collision.
3. These collisions
are perfectly elastic.
There is no energy
loss in the collision.
Gases can be described
using 4 properties.
1. Amount; this is the
number of particles in moles.
2. Temperature;
the measure of heat.
Heat is motion and
temperature is proportional
to the speed of particles.
Gas chemistry calculations
use the Kelvin scale
for temperature.
3. Volume; that is the
3- dimensional space
available to the gas.
It is the size of the
container holding the gas.
And 4. Pressure; it is the
force exerted by a substance
(per unit area) on
another substance.
For gases it can be
rationalized from the
perspective of collisions
of the gas particle with the
sides of a container.
Finally, standard temperature
and pressure (STP).
It is a gas at 0 degrees
C and 1 atm of pressure.
At STP 1 mole of the
gas occupies 22.4 liters
(molar volume).
And that concludes our lecture.
Remember, gases can
be viewed for the
perspective of collisions.
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