Physical Characteristics of Gases
Section 1
The Kinetic-Molecular Theory of
Matter
Elemental states at 25oC
H
He
Solid
Liquid
Gas
Li Be
Na Mg
K
Ca Sc Ti
Rb Sr
Y
V
C
N
O
F
Al
Si
P
S
Cl Ar
Ne
Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te
Cs Ba Ls Hf Ta
Fr Ra Ac
B
W Re Os
Ir
I
Xe
Pt Au Hg Tl Pb Bi Po At Rn
Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
5-2
Behavior of Atoms
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
Kinetic-molecular theory  based on the idea
that particles of matter are always in motion
Can be used to explain the properties of solids,
liquids, and gases in terms of the energy of the
atoms and the forces that act between them
K-M Theory of Gases

Theory provides model of ideal gas
Ideal gas  imaginary gas that perfectly fits all
the assumptions of the kinetic-molecular theory

Based on 5 assumptions

Assumption #1




Gases consist of large numbers of tiny particles that
are far apart relative to their size.
Typically occupy volume about 1000 times greater
than volume of liquid or solid
Molecules of gas are much farther apart than
liquid/solid
Accounts for lower densities, and compressibility
The gaseous state
In this state, the particles have sufficient
energy to overcome all forces that attract
them to each other.
Each particle is completely
separated from the others.
This results in low densities
and the fact that gases completely fill the
container that holds them.
5-7
Assumption #2
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

Collisions between gas particles and between particles
and container walls are elastic collisions.
Elastic collision  there is no net loss of kinetic
energy
Kinetic energy transferred but TOTAL kinetic energy
remains the same as long as temperature is constant
Assumption #3

Gas particles are in continuous, rapid, random
motion. They therefore possess kinetic energy,
which is energy of motion.
Gas particles move in all directions
 Kinetic energy overcomes attractive forces
between atoms

Assumption #4



There are no forces
of attraction or
repulsion between
gas particles.
Think of gas atoms
like billiard balls
They hit each other
and bounce off
Assumption #5




The average kinetic energy of gas particles depends
on the temperature of the gas.
Kinetic energy of any moving object shown by
m = mass of particle
v = velocity (speed)
K-M Theory and the Nature of Gases




K-M theory only applies to ideal gases
They do not actually exist
Many gases behave NEARLY ideally if pressure not
very high or temperature not very low
How does K-M theory account for the physical
properties of gases?
Expansion



Gases do not have definite shape OR volume
Completely fill any container and take its shape
K-M theory:
 Assumption
3  gas particles move rapidly in all
directions
 Assumption 4  no significant attraction or repulsion
between them
Fluidity



Because attractive forces
between gas particles
are insignificant
(assumption 4), the
particles easily pass each
other
This ability causes gases
to behave similarly to
liquids
Because liquids and
gases flow, both are
referred to as fluids
Low Density


Density of gases about 1/1000 density of same
substance in liquid or solid state
This is because the particles are so much farther
apart in the gaseous state (assumption 1)
Compressibility


During compression, the
gas particles, which
are initially very far
apart (assumption 1),
are crowded closer
together
The volume of a given
sample of a gas can
be greatly decreased
Diffusion and Effusion




Gases spread out and mix with one
another, even without being stirred
If the stopper is removed from a
container of ammonia in a room,
ammonia gas will mix uniformly with the
air and spread throughout the room
The random and continuous motion of
the ammonia molecules (assumption 3)
carries them throughout the available
space
Such spontaneous mixing of the particles
of two substances caused by their random
motion is called diffusion

1.
2.
3.
Rate of diffusion of one gas through another
depends on three properties of the gas particles
Their speeds
Their diameters
The attractive forces between them
Effusion


Effusion  process by which gas particles pass
through a tiny opening
Rates of effusion directly proportional to the
velocities of the particles
Deviation of Real Gases from Ideal
Behavior


When their particles are far enough apart and
have enough kinetic energy, most gases behave
ideally
Real gas  gas that does not behave completely
according to the assumptions of the kineticmolecular theory
1873 – Johannes van der Waals



Accounted for movement away from ideal behaviour by
pointing out that particles of real gases occupy space
and apply attractive forces on each other
At very high pressure and low temperatures, the
deviation may be significant
Under such conditions, the particles will be closer
together and their kinetic energy will be insufficient to
completely overcome the attractive forces
(a) Gas molecules in a car engine cylinder
expand to fill the cylinder. (b) As pressure is
exerted on them, the gas molecules move closer
together, reducing their volume. The closer they
are together, the more the attractive forces act
on the particles.


K-M theory more likely to hold true for gases whose
particles have little attraction for each other
Ex. – noble gases
They are monoatomic
 They are nonpolar


The more polar a gas’s molecules are, the greater the
attractive forces between them and the more they will
stray from ideal gas behavior
Section 2 - Pressure
Suppose you have a one-liter bottle of air. How much air do
you actually have? The expression a liter of air means little
unless the conditions at which the volume is measured are
known. A liter of air can be compressed to a few milliliters. It
can also be allowed to expand to fill an auditorium.
To describe a gas fully, you need to state four measurable
quantities: volume, temperature, number of molecules, and
pressure. You already know what is meant by volume,
temperature, and number of molecules. We will examine the
mathematical relationships between volume, temperature,
number of gas molecules, and pressure.
Pressure and Force



If you blow air into a rubber balloon, the balloon
will increase in size
The volume increase is caused by the collisions of
molecules of air with the inside walls of the balloon
The collisions cause an outward push, or force,
against the inside walls
Pressure



Pressure (P)  the force per unit area on a surface
SI unit for force = Newton (N)  force that will
increase the speed of a one kilogram mass by one meter
per second each second it is applied
At Earth’s surface, each kilogram of mass exerts 9.8 N
of force, due to gravity
Gas pressure
Gases exhibit pressure on any container they
are in.
Pressure is defined as a force per unit of area.
Pressure = Force / Area
Several common units
1.00 atm =
760 torr
760 mm Hg
29.9 in Hg
14.7 lb/in2
1.01 x 105 Pa
force
area
5 - 27


Gas molecules exert
pressure on any
surface with which they
collide
The pressure exerted
by a gas depends on
volume, temperature,
and the number of
molecules present
Measuring Pressure


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

Barometer  a device used to measure
atmospheric pressure
Torricelli sealed a long glass tube at
one end and filled it with mercury
Held open end with his thumb, he
inverted the tube into a dish of mercury
without allowing any air to enter the
tube
When he removed his thumb, the
mercury column in the tube dropped to
a height of about 760 mm above the
surface of the mercury in the dish
He repeated the experiment with tubes
of different diameters and lengths
longer than 760 mm
In every case, the mercury dropped to a
height of about 760 mm
Units of Pressure





Many units used to
measure pressure
mmHg 
millimeters of
mercury
1 mmHg = 1 torr
1 atm 
atmosphere of
pressure = 760
mmHg
SI unit  Pascal 
the pressure exerted
by a force of one
Newton (1N) acting
on an area of one
square meter
Standard Temperature and Pressure



To compare volumes of gases, it is necessary to
know the temperature and pressure at which the
volumes are measured
For purposes of comparison, scientists have agreed
on standard conditions of exactly 1 atm pressure and
0°C
These conditions are called standard temperature
and pressure and are commonly abbreviated STP
Practice Problems
1. The average atmospheric pressure in Denver,
Colorado, is 0.830 atm. Express this pressure
(a) in mm Hg and (b) in kPa.
631 mm Hg
84.1 kPa
2. Convert a pressure of 1.75 atm to kPa and to
mm Hg.
177 kPa, 1330 mm Hg
3. Convert a pressure of 570. torr to atmospheres
and to kPa.
0.750 atm, 76.0 kPa