PHYSICAL CHARACTERISTICS
OF GASES
Chapter 10
Section 10-1
Objectives:
State the kinetic-molecular theory of matter, and describe how it
explains certain properties of matter.
¤  List the 5 assumptions of the kinetic-molecular theory of gases.
Define the terms ideal gas and real gas.
¤  Describe these characteristic properties of gases:
¤ 
Expansion
n  Density
n  Fluidity
n  Compressibility
n  Diffusion
n  Effusion
n 
¤ 
Describe the conditions under which a real gas deviates from
“ideal” behavior.
Terms in this Chapter
Expansion
¨  Pressure
¨  Barometer
¨  Manometer
¨  Density
¨  Fluidity
¨  Diffusion
¨  Effusion
¨  Condensation temperature
¨ 
Kinetic-Molecular Theory
¨ 
Particles of matter are always in motion!
¤  For
gases, the theory provides a model of what is
called an “ideal gas”
n  An
imaginary gas that perfectly fits all the assumptions of
the kinetic-molecular theory.
Background – Kinetic-Molecular Theory
¨ 
Johann van Helmont – 1662
¤  Invented
the term, “gas”, to describe the most energetic
phase of matter. The term, “gas”, was derived from the
Greek “chaos”, meaning original matter of the earth.
¤  Vapor
– a gas produced from a substance that is a
solid or liquid under normal conditions. It is formed if
the vapor pressure is equal to that of the atmosphere.
Kinetic-Molecular Theory of Gases
Based on 5 assumptions
1. 
Gases consist of large numbers of tiny particles
that are far apart relative to their size.
¤  Most
2. 
of the volume occupied by a gas is empty space
Collisions between gas particles and between
particles and container walls are elastic collisions.
¤  An
elastic collision is one in which there is NO net loss of
kinetic energy. Kinetic energy is transferred between 2
particles during collisions and is conserved as long as
temperature is constant.
K-M Theory of Gases (assumptions)
3. Gas particles are in continuous, rapid, random
motion. Therefore, they possess kinetic energy, which is
energy of motion.
4. There are no forces of attraction or repulsion
between gas particles. (behave as billiard balls)
5. The average kinetic energy of gas particles is
constant and depends on the temperature of the gas.
KE = ½ mv2
m=mass
v=speed
Kinetic-Molecular Theory and the
Nature of Gases
¨ 
“The kinetic-molecular theory applies only to ideal
gases. Although ideal gases do not actually exist,
many gases behave nearly ideally if pressure is not
very high or temperature is not very low.”
Expansion
¨ 
Gases have no definite shape or
volume. They completely fill any
container and take its shape.
¤  A
gas transferred from a 1-liter
vessel to a 2-liter vessel will quickly
expand to fill the entire 2-liter
volume.
¤  Gas particles move rapidly in all
directions, without significant
attraction or repulsion between
them.
Fluidity
¨ 
Gas particles glide easily past one another. This
ability to flow causes gases to behave similarly to
liquids. Because liquids and gases flow, they are
both referred to as fluids.
Low Density
The density of a substance in the gas state is about
1/1000 the density of the same substance in the
liquid or solid state.
¨  For gases, density is a measure of mass (number of
molecules) per volume, g/L.
¨ 
Compressibility
During compression, the gas particles, which are
initially very far apart, are crowded close together.
The volume of a gas can be greatly decreased.
¨  Steel cylinders containing gases under pressure are
widely used in industry. When they are full, they
may contain 100 times as many particles of gas as
would be contained under non-pressurized
conditions.
¨ 
Diffusion and Effusion
¨ 
¨ 
Gases spread out and mix with one another, even without
being stirred. Such spontaneous mixing of the particles of
two substances caused by their random motion is called
diffusion.
Rate of diffusion depends on 3 things
¤ 
Speed, diameter, and the attractive forces between the particles
n 
¨ 
Hydrogen gas would diffuse faster because its particles are smaller
and move faster than molecules of other gases.
Effusion is a process by which gas particles pass through a
tiny opening. Rates of effusion are directly proportional to
the velocities of their particles. Molecules of low mass effuse
faster than molecules of high mass.
Condensation Temperature
¨ 
The temperature (point) at which gas molecules form
a liquid.
Deviations of Real Gases from Ideal
Behavior
¨ 
Real gas – a gas that does not behave completely
according to the assumptions of kinetic-molecular
theory
¤  1.
in real gases, a weak intermolecular attraction occurs
between some gas molecules
¤  2. London dispersion forces
Johannes Van Der Waals -1873
n  Caused
by motion of e-, and increases
as number of e- increases
¤ 
3. dipole-dipole interactions –
electrostatic attractions based on
polarity of molecules.
Van der Waal forces
¨ 
¨ 
Particles of real gases occupy space and exert attractive
forces on each other. At very high pressures and low
temperatures, the deviation may be considerable. Under
these conditions, particles are closer together and their
kinetic energy will be insufficient to completely overcome the
attractive forces.
Noble gases show ideal gas behavior. They are monatomic
and nonpolar.
¤ 
¨ 
N2 and H2 also exhibit ideal gas behavior.
The more polar a gas’s molecules are, the greater the
attractive forces between them and the more the gas will
deviate from ideal gas behavior. (NH3 and H2O vapor)
TO DESCRIBE A GAS FULLY,
YOU MUST STATE 4
MEASURABLE QUANTITIES:
VOLUME
TEMPERATURE
NUMBER OF PARTICLES
PRESSURE
Units to Know!
1. 
2. 
3. 
4. 
5. 
6. 
7. 
Newton (force) = kg-m/sec2
Pressure as pascal = N/m2
Kilopascal = 1000 pascal
Pressure units include: inches Hg, mm Hg, torr, atm, kPa
Standard = 29.92 in Hg = 760 mm Hg =760 torr = 1
atm = 101.325 kPa
STP – standard temperature (00C, 273 K) and
pressure (101.325 kPa)
K = 0C + 273
Unit of Pressure
Symbol
Units of Pressure
Pascal
¨ 
Pa
Millimeters of mercury- mm Hg
¤  A
Definition/
relationship
SI pressure unit
1 Pa = 1 N/ m2
pressure of 1 mm Hg is now called “1 torr”
Millimeter of mercury
¨ 
Torr ¨ 
mm Hg
Pressure that supports a 1 mm
Atm – atmosphere
of pressure; defined
as equal to 760
mercury column in a
mm Hg
barometer
Pascal (Pa) SI unit
Pascal,
torrof pressure; named 1for
torr Blaise
= 1 mm Hg
French mathematician and philosopher who studied
pressure during the 17th century.
Atmosphere
¤  1
Pascal (Pa) = atm
the pressure exered by aAverage
forceatmospheric
of 1 N pressure
at sea level and 0 0C
acting on an area of 1 m2.
1 atm = 760 mm Hg
¤  Kilopascals (kPa) – The standard atmosphere =(1760
atm)
torr =
= 1.01325 x 105Pa
1.01325 x 105 Pa, or 101.325 kPa.
= 101.325 kPa
Conversion of Units-Practice
¨ 
Calculate kPa for an 800 N force over a 0.1 meter
by 0.2 meter area.
¤  N/m2
¨ 
Convert 850 mm Hg to atm.
¤  760
¨ 
Convert 123.5 kPa to torr.
¤  760
¨ 
mm Hg = 1 atm
torr = 101.325 kPa
Convert -240C to K
¤  K
= 0C + 273
Pressure
¨ 
The force per unit area on a surface. (P)
PRESSURE = FORCE/AREA
¨ 
SI Unit for FORCE is the Newton. (N) It is the force that
will increase the speed of a 1-kg mass by 1 meter per
second each second it is applied.
¤  1
¨ 
Newton = 1kg x m/s2
At Earth’s surface, each kg of mass exerts 9.8 N of
force, due to gravity.
¤  Ex.
A mass of 51 kg exerts a force of 500 N (51 x 9.8) on
Earth’s surface.
Atmospheric Pressure
The atmosphere, because of its mass, exerts pressure
on Earth.
¨  At sea level, atmospheric pressure is about equal to
1.03 kg mass per cm2 of surface, or 10.1 N/cm2 or
760 mm Hg at 0 degrees Celsius.
¨  Since the atmosphere is made up of about
¨ 
¤  78%
N2, 21% O2, and 1% other gases such as CO2 and
Argon
n  Atmospheric
pressure is the sum of all the individual pressures
of the various gases.
Measuring Pressure
¨ 
¨ 
A barometer is a device used
to measure atmospheric
pressure.
1st type of barometer was
introduced by Evangelista
Torricelli in the early 1600s.
Standard Temperature & Pressure
Necessary to know the temperature and pressure at
which volumes of gases are measured in order to
compare volumes of gases.
¨  STP - Standard conditions (determined by scientists)
are exactly 1 atm pressure and 0 0C.
¨ 
GAS LAWS
Chapter 10
Gas Laws
¨ 
Simple mathematical relationships between the
volume, temperature, pressure, and amount of a
gas.
¤  Boyles’s
Law
¤  Charles’s Law
¤  Gay-Lussac’s Law
¤  Combined Gas Law
¤  Dalton’s Law of Partial Pressures
Boyle’s Law – Pressure-Volume
Relationship
¨ 
1662-Robert Boyle discovered that gas pressure and
volume are related mathematically.
¤  Doubling
the pressure on a sample of gas at constant
temperature reduces its volume by ½.
¤  Likewise, reducing the pressure on a gas by ½ allows the
volume of the gas to double.
¨ 
Boyle’s Law-a volume of a fixed mass of gas varies
inversely with the pressure at constant temperature.
n  V
= k x 1/P
OR
PV = k
n  P1V1 = P2V2 ( used to compare changing conditions for a gas
where P1 and V1 stand for initial conditions, and P2 and V2
stand for new conditions.)
Boyle’s Law example
¨ 
A sample of oxygen gas has a volume of 150. mL
when its pressure is 0.947 atm. What will the
volume of the gas be at a pressure of 0.987 atm if
the temperature remains constant?
Charles’s Law: Volume-Temperature
Relationship
¨ 
1787-French scientist Jacques
Charles showed that all gases
expand to the same extent when
heated through the same
temperature interval. He found that
the volume changes by 1/273 of
the original volume for each Celsius
degree.
Raising the temp. to 10C causes
the gas volume to increase by 1/273
of the volume it had a 00C.
¤  Ex.
Charles’s Law
¨ 
The volume of a fixed mass of gas at constant
pressure varies directly with the Kelvin temperature.
V1 and T1 represent initial conditions
¨  V2 and T2
“ new conditions
¨ 
¨ 
Absolute zero = -273.15 0C or 0 Kelvin.
¤  K
= 273.15 + 0C
Sample Problem
¨ 
Problem – A sample of neon gas occupies a volume
of 752 mL at 250C. What volume will the gas
occupy at 500C if the pressure remains constant?
Gay-Lussac’s Law: Pressure-Temperature
Relationship
Joseph Gay-Lussac recognized in 1802 that for
every kelvin of temperature change, the pressure of
a confined gas changes by 1/273 of the pressure
at 00C.
¨  Gay-Lussac’s Law – the pressure of a fixed mass of
gas at constant volume varies directly with the
Kelvin temperature.
¨ 
Sample Problem
¨ 
The gas in an aerosol can is at a pressure of 3.00
atm at 250C. Directions on the can warn the user not
to keep the can in a place where the temperature
exceeds 520C. What would the gas pressure in the
can be at 520C?
¤  Remember
to convert from 0C to K.
Combined Gas Law
All of the formulas we have learned assume that
either temperature, pressure, or volume remains
constant depending on the conditions and the type
of problem we are trying to solve.
¨  But a gas sample often undergoes changes in
temperature, pressure, and volume all at the same
time, causing us to deal with 3 variables at the same
time.
¨  CGL – expresses the relationship between pressure,
volume, and temperature of a fixed amount of gas.
¨ 
Combined Gas Law
Sample Problem
¨ 
A helium filled balloon has a volume of 50.0 L at
250C and 1.08 atm. What volume will it have at
0.855 atm and 10.0C?
Dalton’s Law of Partial Pressures
¨ 
¨ 
¨ 
John Dalton found that in the absence of a chemical
reaction, the pressure of a gas mixture is the sum of
the individual pressures of each gas alone.
Dalton’s Law of Partial Pressures – the total
pressure of a mixture of gases is equal to the sum
of the partial pressures of the component gases.
PT = P1 + P2 + P3 + ….
Gases Collected by Water
Displacement
¨ 
¨ 
Patm = Pgas + PH2O
Sample Problem – Oxygen gas from the
decomposition of potassium chlorate, KClO3, was
collected by water displacement. The barometric
pressure and the temperature during the
experiment were 731.0 torr and 20.00C,
respectively. What was the partial pressure fo the
oxygen collected?