Chapter 10
Gases
Gas
Gas Model
• Gases are composed of tiny,
widely-spaced particles.
– For a typical gas, the average
distance between particles is about
ten times their diameter.
Gas Model (cont.)
• Because of the large distance between
the particles, the volume occupied by
the particles themselves is negligible
(approximately zero).
– For a typical gas at room temperature
and pressure, the gas particles
themselves occupy about 0.1% of the
total volume. The other 99.9% of the
total volume is empty space. This is
very different than for a liquid for
which about 70% of the volume is
occupied by particles.
Gas Model (cont.)
• The particles have rapid and
continuous motion.
– For example, the average velocity of a
helium atom, He, at room temperature
is over 1000 m/s (or over 2000 mi/hr).
The average velocity of the more
massive nitrogen molecules, N2, at
room temperature is about 500 m/s.
– Increased temperature means
increased average velocity of the
particles.
Gas Model (cont.)
• The particles are constantly colliding
with the walls of the container and
with each other.
– Because of these collisions, the gas
particles are constantly changing their
direction of motion and their velocity. In
a typical situation, a gas particle moves
a very short distance between collisions.
Oxygen, O2, molecules at normal
temperatures and pressures move an
average of 10−7 m between collisions.
Gas Model (cont.)
• There is no net loss of energy in
the collisions. A collision between
two particles may lead to each
particle changing its velocity and
thus its energy, but the increase
in energy by one particle is
balanced by an equal decrease in
energy by the other particle.
Ideal Gas
• The particles are assumed to be
point-masses, that is, particles
that have a mass but occupy no
volume.
• There are no attractive or
repulsive forces at all between
the particles.
Gas Properties and
their Units
• Pressure (P) = Force/Area
– units
• 1 atm = 101.325 kPa = 760 mmHg = 760 torr
• 1 bar = 100 kPa = 0.9869 atm = 750.1 mmHg
• Volume (V)
– unit usually liters (L)
• Temperature (T)
– ? K = --- °C + 273.15
• Number of gas particles expressed in
moles (n)
Gas Law Objectives
• For each of the following pairs of gas
properties, (1) describe the relationship
between them, (2) describe a simple
system that could be used to demonstrate
the relationship, and (3) explain the
reason for the relationship.
• V and P when n and T are constant
• P and T when n and V are constant
• V and T when n and P are constant
• n and P when V and T are constant
• n and V when P and T are constant
Apparatus for Demonstrating Relationships
Between Properties of Gases
Decreased Volume Leads
to Increased Pressure
P α 1/V
if n and T are constant
Relationship
between
P and V
Boyle’s Law
• The pressure of an ideal gas is inversely
proportional to the volume it occupies if the
moles of gas and the temperature are
constant.
Increased Temperature Leads
to Increased Pressure
PαT
if n and V are constant
Relationship between
P and T
Gay-Lussac’s Law
• The pressure of an ideal gas is directly
proportional to the Kelvin temperature of the gas
if the volume and moles of gas are constant.
Increased Temperature Leads
to Increased Volume
VαT
if n and P are constant
Relationship between T and V
Charles’ Law
• For an ideal gas, volume and temperature
described in kelvins are directly proportional if
moles of gas and pressure are constant.
Increased Moles of Gas Leads
to Increased Pressure
Pαn
if T and V are constant
Relationship between n and P
Relationship Between
Moles of Gas and Pressure
• If the temperature and the volume of an ideal gas
are held constant, the moles of gas in a container
and the gas pressure are directly proportional.
Increased Moles of Gas
Leads to Increased Volume
Vαn
if T and P are constant
Relationship between n and V
Avogadro’s Law
• For an ideal gas, the volume and moles of gas
are directly proportional if the temperature and
pressure are constant.
Engines
and
Pressure
Breathing
Ideal Gas Equation
Ideal Gas Equation
Problems
• Tip-off – The usual tip-off that you
can use the ideal gas equation to
answer a question is that you are
given three properties of a sample
of gas and asked to calculate the
fourth. A more general tip-off is that
only one gas is mentioned and
there are no changing properties.
Ideal Gas Equation
Problem Step 1
• Step 1: Assign variables to the
values given and the value that is
unknown. Use P for pressure, V for
volume, n for moles, T for
temperature, g for mass, and M for
molar mass.
Ideal Gas Equation
Problem Step 2
• Step 2: Write the appropriate form of the
Ideal Gas Equation.
– If the number of particles is given or desired in
moles, use the most common form of the ideal
gas equation.
– If mass or molar mass is given or desired, use
the expanded form of the ideal gas equation.
Ideal Gas Equation
Problem Steps 3-6
• Step 3: Rearrange the equation to
isolate the unknown.
• Step 4: Plug in the known values,
including units. Be sure to use kelvin
temperatures.
• Step 5: Make any necessary unit
conversions and cancel your units.
• Step 6: Calculate your answer and
report it to the correct significant figures
and with the correct unit.
Combined
Gas Law
Equation
Combined Gas Law
Equation Problems
• Tip-off – The problem requires
calculating a value for a gas
property that has changed. In
other words, you are asked to
calculate a new pressure,
temperature, moles (or mass), or
volume of gas given sufficient
information about the initial and
other final properties.
Combined Gas Law
Equation Problem
Steps 1 and 2
• Step 1: Assign the variables P, T, n, and
V to the values you are given and to the
unknown value. Use the subscripts 1 and
2 to show initial or final conditions.
• Step 2: Write out the combined gas law
equation, but eliminate the variables for
any constant properties. (You can
assume that the properties not mentioned
in the problem remain constant.)
Combined Gas Law
Equation Problem
Steps 1 and 2
• Step 3: Rearrange the equation to
isolate the unknown property.
• Step 4: Plug in the values for the
given properties.
• Step 5: Make any necessary unit
conversions and cancel your units.
• Step 6: Calculate your answer and
report it with the correct units and
significant figures.
Equation Stoichiometry
Dalton’s Law of Partial
Pressures
Dalton’s Law of Partial
Pressures Equations
nA RTA nB RTB nA RT nB RT
PT = PA + PB =
+
=
+
VA
VB
V
V
RT
RT
or PT = ( ∑ neach gas )
PT = (nA + nB )
V
V
nA RT
PA
nA
V
=
=
= X A = mole fraction A
RT ( ∑ neach gas )
PT ( ∑ n
each gas )
V
PA = X A PT
Dalton’s Law of Partial
Pressures Problems
• Tip-off – The problem involves a
mixture of gases and no chemical
reaction. You are asked to
calculate a value for one of the
variables in the equations below,
and you are given (directly or
indirectly) values for the other
variables.
Dalton’s Law of Partial
Pressures Problem
Steps 1 & 2
• Step 1: Assign variables to the values that are
given and the value that is unknown.
• Step 2: From the following equations, choose
the one that best fits the variables assigned in
Step 1.
PT = PA + PB
RT
PT = ( ∑ neach gas )
V
PA = X A PT
Dalton’s Law of Partial
Pressures Problem Steps 3-6
• Step 3: Rearrange the equation to solve
for your unknown.
• Step 4: Plug in the values for the given
properties.
• Step 5: Make sure that the equation
yields the correct units. Make any
necessary unit conversions.
• Step 6: Calculate your answer and
report it with the correct units and
significant figures.
Collecting
Gas Over
Water
Vapor
Pressures
of Water
T
mmHg
T
mmHg
-10
2.15
40
55.3
0
4.58
60
149.4
5
6.54
80
355.1
10
9.21
95
634
11
9.84
96
658
12
10.52
97
682
13
11.23
98
707
14
11.99
99
733
15
12.79
100
760
20
17.54
101
788
25
23.76
110 1074.6
30
31.8
120
37
47.07
200 11659
1489
Ideal Gas
Assumptions
• The particles are assumed to be
point-masses, that is, particles
that have a mass but occupy no
volume.
• There are no attractive or
repulsive forces at all between
the particles.
Real Gases
• The ideal gas assumptions are
generally valid for real gases at
relatively low concentrations (leading
to low pressures) and at temperatures
well above the condensation
temperature of the substance.
• Decreased temperature and/or
increased concentration of gas leads
to greater differences between the
properties predicted by gas law
calculations and the real, measured
properties.
Attractions and
Gas Pressure (1)
• Attractions between the particles in
a gas decrease the average force
of their collisions with the walls,
leading to a measured pressure
that is lower than that predicted by
the ideal gas equation calculations,
which are based on the
assumption that there are no
attractions between the particles.
Attractions and Gas
Pressure (2)
Attractions and
Gas Pressure (3)
Preal = Pmeasured
nRT
Pideal =
V
Preal < Pideal
Concentration
Effect (1)
• Increased concentration of gas (due to
increased number of gas particles
and/or decreased volume) leads to
decreased average distance between
the particles, increased attraction
between the particles, decreased
average force of collisions with the
walls, and decreased pressure. This
means a greater deviation of the real,
measured pressure from the ideal,
calculated pressure.
Concentration Effect (2)
Temperature Effect
• Decreased temperature leads to
decreased velocity, which leads to
a decreased chance of having
particles break attractions among
them.
• This leads to more attractions
between particles at a given time,
decreased average force of
collisions with the walls, and
decreased pressure.
Real
(Measured)
Pressure
Different
from Ideal
(Calculated)
Pressure
Real and Ideal Volume
• The ideal gas equation was derived
based on the assumption that the
particles in a gas occupy no volume,
therefore, the volume calculated from the
ideal gas equation is the volume of empty
space.
• The real volume is the volume of empty
space plus the volume occupied by the
particles. Therefore, it is greater than the
calculated ideal volume.
Videal = Vempty
Vreal = Vempty + Vparticles
Vreal > Videal
Concentration Effect
on Volume
• Increased concentration of gas
leads to an increase in the
percentage of the volume occupied
by particles and therefore, a
greater difference between the
calculated ideal volume and the
measured real volume.
Ozone, O3, as
Oxidizing Agent
• Used to sanitize hot tubs
• Used in industry to bleach waxes, oils,
and textiles.
• Strong respiratory irritant that can lead to
shortness of breath, chest pain when
inhaling, wheezing, and coughing
• Damages rubber and plastics, leading to
premature deterioration of products
made with these materials.
• Ozone damages plants.
Ozone as Pollutant
ƒ Highest concentrations found in large
industrial cities with lots of cars and lots
of sun.
→
N2(g) + O2(g)
2NO(g) + O2(g)
2NO(g)
→
2NO2(g)
λ < 400 nm
NO2(g)
O(g) + O2(g)
NO(g) + O(g)
→
O3(g)
National Ozone Concentrations
The Earth’s Atmosphere
Ultraviolet Light
• UV-A
320-400 nm
– reaches the surface of the Earth
– helps create Vitamin D
• UV-B
290-320 nm
– some reaches the surface of the Earth
– leads to sunburn, skin aging, and skin cancer
• UV-C
40-290 nm
– mostly removed in upper atmosphere
– alters DNA (≈260 nm) and protein (≈280 nm)
Removal of UV in
Stratosphere
Ozone Destruction Paul Crutzen (1970)
NO(g) + O3(g) → NO2(g) + O2(g)
NO2(g) + O(g) → NO(g) + O2(g)
net reaction
NO catalyst
O3(g) + O(g)
2O2(g)
Chlorofluorocarbons
(CFCs)
• CFC-11
CFCl3
– average lifetime in atmosphere is ≈ 50 years
• CFC-12
CF2Cl2
– average lifetime in atmosphere is ≈ 102
years
• used as propellants in aerosol cans,
solvents, blowing agents for foams,
coolant in refrigerators, and other uses
• very stable, nontoxic, and can be
liquefied with minimal pressure
Aerosol Can Propellants
Refrigeration
• The refrigerant is a substance that is a gas at
normal pressures but one that can be
converted into a liquid at slightly greater than
normal pressures.
• Outside the refrigerator, gas is compressed to
liquid. Increased attractions leads to increased
stability, lower PE, and the release of energy
into the room.
• Inside the refrigerator, the liquid is allowed to
form a gas. Decreased attractions leads to
decreased stability, higher PE, and energy is
absorbed. This decreases the temperature
inside the refrigerator.
CFC Threat to Ozone (1974)
Mario Molina and F. Sherwood
Rowland
λ < 215 nm
CF2Cl2(g)
CF2Cl(g) + Cl(g)
Cl(g) + O3(g) → ClO(g) + O2(g)
ClO(g) + O(g) → Cl(g) + O2(g)
net reaction
O3(g) + O(g)
Cl catalyst
2O2(g)
Another Possible
Mechanism for the CFC
Threat to Ozone
λ < 215 nm
CF2Cl2(g)
CF2Cl(g) + Cl(g)
2Cl(g) + 2O3(g) → 2ClO(g) + 2O2(g)
2ClO(g) → ClOOCl(g)
ClOOCl(g) → ClOO(g) + Cl(g)
ClOO(g) → Cl(g) + O2(g)
net reaction
2O3(g)
Cl catalyst
3O2(g)
Inactive Chlorine
CH4(g) + Cl(g) → CH3(g) + HCl(g)
ClO(g) + NO2(g) → ClONO2(g)
1985 - Ozone Hole Reactions on the surface
of ice crystals
ClONO2(g) + HCl(s) → Cl2(g) + HNO3(s)
ClONO2(g) + H2O(s) → HOCl(g) + HNO3(s)
HOCl(g) + HCl(s) → Cl2(g) + H2O(s)
radiant energy
HOCl(g)
Cl(g) + OH(g)
radiant energy
Cl2(g)
2Cl(g)