Gases Part 1
1
Properties of Gases
• Gases have very low densities, and may be compressed
or expanded easily: in other words, gases expand or
compress to completely fill whatever container they are
in.
• Gases also tend to form homogeneous mixtures, that is
they tend to mix completely to form uniform mixtures.
Solids and liquids often form heterogeneous mixtures.
• Gases exert a pressure or a force where the gas
particles collide with a surface. Remember, gas
particles are moving randomly with high kinetic
energy. Everywhere they collide, the force of the
collision is the pressure that the gas exerts.
2
Properties of Gases
3
Properties of Gases
4
Pressure Units of Gases
• Scientists typically use atmospheres, atm,
where sea level atmospheric pressure (called
standard atmospheric pressure) is exactly 1
atm.
• Although atm is quite commonly used, the SI
unit of pressure is Pascals, Pa, where 1 atm =
101.325kPa.
• In the US, our barometers (what we use to
measure relative atmospheric pressure),
usually read in mm Hg or in torr, where 1 atm
= 760 mm Hg = 760 torr.
5
Measuring Pressure of Gases
• Scientists use manometers and barometers to
measure pressure.
• Barometers measure atmospheric pressure.
• Manometers measure the pressure of a rxn
vessel.
• Manometers can be open or closed.
6
Measuring Pressure of Gases
7
Kinetic Molecular Theory of Gases
•
The special properties of gases have been
explained by a series of assumptions, called
the Kinetic-Molecular Theory:
1. Gases consist of large numbers of particles that are
in continuous, random, straight-line motion.
2. The volume of the particles of gases is negligible
compared to the volume of “empty” space between
the particles. In fact, gas particles typically take
up only about 0.1% of the total volume of their
containers, the rest being empty space.
8
Kinetic Molecular Theory of Gases
3. Attractive and repulsive forces between gas
particles is negligible. Therefore, there are no
appreciable intermolecular forces and all of the gas
particles act essentially independently of one
another.
4. For a given temperature, the average kinetic
energy of the gas particles is constant, as is the
total kinetic energy of all the gas particles.
Therefore, gas particle collisions are perfectly
elastic.
9
Kinetic Molecular Theory of Gases
5. The kinetic energy of gas particles depends
on temperature. The higher the
temperature, the higher the kinetic energy
of the particles. At a given temperature,
the average kinetic energy of ALL gases is
the same. So the average kinetic energy of
O2 is the same as He is the same as CO2 so
long as they are at the same temperature.
Because KE = 0.5mv2, and they have
different masses, they must have different
velocities, BUT THEY HAVE THE SAME
KINETIC ENERGY!
10
Gas Laws
• All gases have very similar physical
behavior.
• What are the gas variables? What can
we change about a gas?
• These variables and how they relate to
one another are described in the Gas
Laws: Boyle’s; Charles’; Gay-Lussac’s;
Avogadro’s; Combined Law; Ideal Gas
Law; and Dalton’s Law of Partial
Pressure.
11
Boyle’s Law
• Boyle’s Law describes the relationship
between the pressure of a gas sample
and its volume (n and T constant).
• What do you think happens to the
volume of a gas when we increase the
pressure in an expandable and
compressible container?
• Why?
12
Boyle’s Law
13
Boyle’s Law: An Inverse Relationship
14
Boyle’s Law
• Boyle’s Law is stated as:
P1V1 = P2V2
• where 1 are the initial conditions and 2
are the final conditions.
• Note that we use Boyle’s Law when we
change the volume and pressure.
•
So if we change the volume, we can calculate
the new pressure (or vice versa).
15
Charles’ Law: V and T
• Charles’ Law states the relationship
between Volume and Temperature (P, n
constant).
• If we have a balloon filled with gas and
we start to heat it, what will happen to
the volume of the balloon?
• Why?
• Charles’ Law is stated as:
V1/T1 = V2/T2
16
Charles’ Law: A Linear Relationship
• It is important to remember that in all
of the Gas Laws, the temperature
MUST be in Kelvin!
• Why?
17
Gay-Lussac’s Law: P and T
•
•
•
•
Gay-Lussac’s Law states the relationship
between Pressure and Temperature (V, n
constant).
If we have a closed steel canister filled with
gas and we throw the canister in a fire, what
will happen to the pressure of the gases inside
the canister and why?
Gay-Lussac’s Law is stated as:
P1/T1 = P2/T2
It is a linear relationship.
18
Avogadro’s Law: n and V (and P as well)
•
•
•
•
•
Avogadro’s Law states the obvious
relationship between amount and volume.
If we have a balloon and we blow air into it,
what happens to the volume?
Avogadro’s Law is stated as:
V1/n1 = V2/n2
It is a linear relationship.
The pressure and amount are also a linear
relationship.
19
Combined Gas Law
•
•
•
•
In real life, think weather balloons, more than
2 gas variable changes.
For example, as a weather balloon ascends,
the P, V, and T all change!
So the Combined Gas Law combines all of the
prior Laws into one:
P1V1/n1T1 = P2V2/n2T2
or more commonly:
P1V1/T1 = P2V2/T2
20
Ideal Gas Law
•
•
•
•
But what if we just want to know what the
pressure of a sample is currently?
If we know the temp, volume, and amount of
the gas sample, we can do this using the Ideal
Gas Law:
PV = nRT
where R is the Ideal Gas Constant
Now you know 3 of 4 variables and solve for
the 4th. No change is involved!
21