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Chapter 13 “Gases”
What are the properties of gases?
Chapt. 13 OBJECTIVES
 Describe gas behavior & explain how the
Kinetic-Molecular (K-M) Theory accounts for it.
 Explain what ‘gas pressure’ means & describe
how it is measured.
 State the “Gas Laws” and use them to solve
problems.
 Understand the significance of the “Ideal Gas
Law.”
 Compare ideal & real gases.
 Relate gas density to temperature and molar
mass.
What are the Physical Properties of
Gases?
 Gases have mass.
 It is easy to compress gases.
 Gases fill their containers completely.
 Different gases quickly move through each
other. (This mixing is called diffusion.)
 Gases exert pressure.
 The pressure of a gas depends on its
temperature.
How do we explain gas behavior?
 Gas properties are explained by a
model, called the kinetic-molecular
model, that describes the behavior of
the subatomic particles that make up a
gas.
 This is known as “The Kinetic Molecular
Theory” or “K-M” Theory.
The Kinetic-Molecular (K-M) Theory
 Gases are composed of very small particles (atoms or






molecules) having mass.
Gas particles have low volume, but are spaced widely,
with lots of empty space between them, making gases
easy to compress.
There is no attraction or repulsion among the gas
particles.
Gas particles move in constant, rapid, random motion,
so a gas has no definite shape, but it fills its container
completely.
The gas particles travel in short straight lines, with each
particle moving independently.
When colliding with walls, other particles, etc., the gas
particles rebound elastically, so the total kinetic energy
(KE) is constant at a given temperature.
The average kinetic energy of the molecules of a gas is
proportional to the temperature of the gas in kelvins.
13-2 Measuring Gases
 To completely describe a gas sample and
to make predictions about its behavior
under new conditions, we must deal with
four variables:
 Amount
of gas (n, in moles)
• n = mass/molar mass = m(g)/M(g/mol)
 Volume
of gas (V, in liters)
• 1 L = 1000mL = 1000 cm3
 Temperature
(T, in kelvins)
• T(K) = T(°C) + 273
 Pressure
(P, in atm, torr, kPa, etc.)
How to Measure Pressure
 Pressure is the force exerted by gas particles on
the walls of the container.
 Pressure = force/area (newtons/meter2 =
pascals)
 Atmospheric pressure results from the mass of
the air being attracted by the Earth’s gravity.
 Conversions:
 1 atmosphere (atm) = 101.3 kPa
 1 atmosphere = 760 mm Hg (same as torr)
 1 atmosphere = 14.70 lb/in2
 (See Fig. 13-12, pg 426 for inter-conversions.)
Tools for Measuring Pressure
 Barometer
 Atmospheric pressure used to predict the weather
 Pressure gauge
 Tire pressure gauges or gas cylinder regulators
 Manometer
 Laboratory tools for working with gases in closed space
 Examples: reading manometers
 Do you remember STP?
 STP = Standard Temperature & Pressure
 STP = 273 K (0°C) & 1 atm (760 torr, etc.)
 It is important for measuring gas properties.
13-3 The Gas Laws (1)
 What happens when you put pressure on a gas?
 Examples: Small balloon being squeezed, books
piled on a small cylinder.
 Robert Boyle (1600s) studied the effect of
pressure (P) on the volume (V) of a gas.
 Worksheet class activity

What is the relationship between V and P?
 Boyle’s Law: The pressure and volume of a
sample of gas are inversely proportional to each
other at constant temperature.
Boyle’s Law
Pressure (inches
Hg)
Sam ple Data for Boyle's Law
60
40
20
0
0
50
100
Volum e (cubic inches)
150
13-3 The Gas Laws (2)
 What happens when you change the temperature
of a gas?

Examples: Hot air balloon inflating, automobile tires in
very cold weather.
 Jacques Charles (1700s) studied the effect of
temperature (T) on the volume (V) of a gas.
 Worksheet class activity

What is the relationship between V and T?
 Charles’ Law: The volume of a sample of gas is
directly proportional to the kelvin temperature at
constant pressure.
 Lab Activity: Predicting Absolute Zero
Charles’ Law
Data for Charles' Law
Volume (liters)
6
5
4
3
2
1
0
0
100
200
300
400
500
600
Tem perature (kelvins)
700
13-3 The Gas Laws (3)
 What happens to pressure of a fixed volume of
gas if you change the temperature?

Example: “Empty” aerosol can thrown into an
incinerator. (Read the warning on the label!)
 Gay-Lussac (1700s) studied the effect of
pressure (P) and temperature (T) on a fixed
volume (V) of gas.
 Worksheet class activity

What relationship do you see between P and T?
 Gay-Lussac’s Law: The pressure and
temperature of a fixed volume of gas are directly
proportional to each other.
Gay-Lussac’s Law
Pressure (nn Hg)
Data for Gay-Lussac's Law
6
4
2
0
0
200
400
600
Tem perature (kelvins)
800
13-3 The Gas Laws (4)
 How is the number of gas particles related to its
volume under constant conditions?

Examples: Two balloons of different size.
 Amedeo Avogadro (1800s) studied different
gases to determine the relationship between the
number of gas particles and the volume at a
given pressure (P) and temperature (T).
 Worksheet class activity

What is the relationship between V and n?
 Avogadro’s Law: Equal volumes of gases at
the same pressure and temperature contain an
equal number of particles.
Avogadro’s Law
Equal volumes of different gases have the same
number of particles under the same conditions.
13-3 The Gas Laws (Combined)
These Gas Laws seem complicated, but fortunately we
can simplify things by combining the relationships into two
simple expressions, the one first being:
P1V1
T1
=
P2V2
T2
This equation is used to solve “Combined Gas Law” problems, by simply
“plugging in” the numbers!
Just remember two things:
“Go Kelvin!” (Convert temperature to kelvins.)
Be sure all units are consistent for P and V.
13-3 The Gas Laws (5)
 Another relationship about gases was proposed by John
Dalton. (Remember him? Why?)
 He proposed that gas particles in a mixture of gases act
independently to exert pressure on the container.
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Each gas in the mixture exerts the same pressure that it would if it
was alone in the container.
This proposal was made before the Kinetic Molecular Theory was
developed, so he didn’t really have its ideas to help him.
 Dalton’s Law of Partial Pressure: the sum of the partial
pressures of all components of a gas mixture is equal to
the total pressure of the gas mixture, or...
 PT = P1 + P2 + P3 + .….
 See sample problems.
13-4 The “Ideal Gas Law”
 The second equation that may be derived from the
various gas laws is the “Ideal Gas Equation” ---
PV = nRT
 This describes the relationship among the four variables
(P,V,n,T) of an ideal gas, where R is the gas constant.


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R = 0.0821 atm-L/mol-K
R = 8.314 Pa-m3/mol-K
R = 8.314 J/mol-K
 Problems may be solved by simple substitution, but be
careful to use consistent units.
But what is an “Idea Gas”?
 Ideal Gas: one that is described by the
postulates of the Kinetic-Molecular Theory.
 No such gases exist! The gases we
encounter are “real” – not “ideal”!
 However, most gases behave like ideal
gases under ordinary conditions of
temperature and pressure.
 At
low temperature and high pressure real
gases behave in non-ideal ways. (Why is this
so? See the next slide!)
Deviations from Ideal Gas Behavior
Occur Because:
 Kinetic-Molecular Theory makes two simplifying,
but WRONG, assumptions:
Gas particles have no volume of their own.
 Gas particles have no attraction for each other.

 But, as pressure increases…
 gas particles get closer together.
 At very high pressure, the volume of the gas particles
themselves become a significant part of the total
volume, contrary to K-M Theory.
 As temperature decreases…
 gas particles slow down.
 At very low temperature, attractive forces between
gas particles become significant, contrary to K-M
Theory.
Nevertheless…
 K-M Theory is valid for studying gas
behavior under ordinary conditions.
 We are safe to use the Ideal Gas
Equation (PV = nRT) to solve problems
for real gases, except at low
temperature and high pressure
conditions where gases start to behave
in a non-ideal manner.
13-5 How Gases Work: Lift
 Lifting Power
 This is the result of low density of a gas,
but the density must be lower than that of
the air.
 Low molar mass gases (H2, He, NH3, CH4)
may be used to gain lift.
 Or the gas, usually air, may be heated.
 Higher mass gases (CO2, Kr) are too
dense, so they sink in air.
• Can have disastrous effects.
13-5 How Gases Work: Effusion
 Effusion
 Related to ‘diffusion’ (one gas moving through
another).
 Effusion is the gas movement through a tiny hole
one particle at a time.
 Small, light gases have greater speeds than large,
heavy gases at a given temperature, so they
effuse faster.
 Graham’s Law of Effusion relates effusion to
mass.
Graham’s Law of Effusion
Gases will effuse at a rate that is inversely proportional
to the square root of the Molecular Masses, MMx, or:
rate1
rate2
See worksheet problems.
=
√
MM2
MM1
Did we meet the Chapt. 13
OBJECTIVES?
 Describe gas behavior & explain how the
Kinetic-Molecular Theory accounts for it.
 Explain what ‘gas pressure’ means & describe
how it is measured.
 State the “Gas Laws” and use them to solve
problems.
 Understand the significance of the “Ideal Gas
Law.”
 Compare ideal & real gases.
 Relate gas density to temperature and molar
mass.
Gases are Fun!
The Kinetic-Molecular (K-M) Theory (A)
 A gas consists of very small particles having mass.
 The distances separating gas particles are very large
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compared to the size of the particles. (The volume of
the gas particles themselves are negligible compared
with the total volume of gas.)
Gas particles exert no attractive force on each other.
Gas particles are in constant, rapid, random motion.
Collisions of gas particles with each other or the walls
of the container are perfectly elastic.
The average kinetic energy of gas particles is
proportional to the kelvin temperature of the gas.
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