General Chemistry I
Chapter 10
Gases
Gases You Have Encountered
Characteristics of Gases
• Gases always form homogeneous mixtures with other
gases
• Gases are highly compressible and occupy the full
volume of their containers. (Chapter 1)
• When a gas is subjected to pressure, its volume
decreases.
• .
Pressure
• Pressure is the force acting on an object per unit area:
F
P
A
• Gravity exerts a force on the earth’s atmosphere
• A column of air 1 m2 in cross section exerts a force of 105
N.
• The pressure of a 1 m2 column of air is 100 kPa.
• SI Units: 1 N = 1 kg.m/s2; 1 Pa = 1 N/m2.
Atmosphere Pressure and The
Barometer
• If a tube is inserted into a container of mercury open to
the atmosphere, the mercury will rise 760 mm up the
tube.
• Atmospheric pressure is measured with a barometer.
• Standard atmospheric pressure is the pressure required to
support 760 mm of Hg in a column.
• Units: 1 atm = 760 mmHg = 760 torr = 1.01325  105 Pa
= 101.325 kPa.
Class Guided Practice Problem
• (a) Convert 0.527 atm to torr
• (b) Convert 760 torr to kPa
Class Practice Problem
• (c) Convert 147.2 kPa to (1) atm and (2) torr
Atmosphere Pressure and
The Manometer
• The pressures of gases not open
to the atmosphere are measured
in manometers.
• A manometer consists of a bulb
of gas attached to a U-tube
containing Hg:
– If Pgas < Patm then Pgas + Ph2 = Patm.
– If Pgas > Patm then Pgas = Patm + Ph2.
(See example problem 10.2)
Defining States of Gases
• Gas experiments revealed that four variables will affect
the state of a gas:
• Temperature, T
• Volume, V
• Pressure, P
• Quantity of gas present, n (moles)
• These variables are related through equations know as the
gas laws.
The Ideal Gas Equation
• Consider the three gas laws.
• Boyle’s Law:
V
1
(constant n, T )
P
• Charles’s Law: V  T (constant n, P)
• Avogadro’s Law: V  n (constant P, T )
• We can combine these into a general gas law:
nT
V
P
The Gas Laws: Boyle’s Law
•
•
•
•
•
The Pressure-Volume Relationship:
Weather balloons are used as a practical consequence to
the relationship between pressure and volume of a gas.
As the weather balloon ascends, the volume increases.
As the weather balloon gets further from the earth’s
surface, the atmospheric pressure decreases.
Boyle’s Law: the volume of a fixed quantity of gas is
inversely proportional to its pressure.
Boyle used a manometer to carry out the experiment.
Boyle’s Law
The Pressure-Volume Relationship
• Mathematically:
1
V  constant 
P
PV  constant
• A plot of V versus P is a hyperbola.
• Similarly, a plot of V versus 1/P must be a straight line
passing through the origin.
• The Value of the constant depends on the temperature
and quantity of gas in the sample.
Charles’s Law
The Temperature-Volume Relationship:
• We know that hot air balloons expand when they are
heated.
• Charles’s Law: the volume of a fixed quantity of gas at
constant pressure increases as the temperature increases.
• Mathematically:
V  constant  T
V
 constant
T
Plotting Charles’s Law
• A plot of V versus T is a straight line.
• When T is measured in C, the intercept on the
temperature axis is -273.15C.
• We define absolute zero, 0 K = -273.15C.
• Note the value of the constant reflects the assumptions:
amount of gas and pressure.
All gases will solidify or liquefy before reaching zero volume.
Avogadro’s Law
The Quantity-Volume Relationship:
• Gay-Lussac’s Law of combining volumes: at a given
temperature and pressure, the volumes of gases which
react are ratios of small whole numbers.
Avogadro’s Law
• Avogadro’s Hypothesis: equal volumes of gas at the
same temperature and pressure will contain the same
number of molecules.
• Avogadro’s Law: the volume of gas at a given
temperature and pressure is directly proportional to the
number of moles of gas.
Expressing Avogadro’s Law
• Mathematically:
V  constant  n
• We can show that 22.4 L of any gas at 0C contain 6.02  1023
gas molecules.
The Ideal Gas Constant
• If R is the constant of proportionality (called the gas
constant), then
 nT 
V  R

 P 
• The ideal gas equation is:
PV  nRT
L  atm
J
R  0.08206
 8.314
mol  K
mol  K
Applying The Ideal Gas Equation
• We define STP (standard temperature and pressure) =
0C, 273.15 K, 1 atm.
• Volume of 1 mol of gas at STP is:
nRT 1 mol 0.08206 L·atm/mol· K 273.15 K 
V

 22.41 L
P
1.000 atm
Class Guided Practice Problem
• For an ideal gas, calculate the following quantities: (a)
the pressure of the gas if 1.04 mol occupies 21.8 L at 25
oC; (b) the volume occupied by 6.72 x 10-3 mol at 265 oC
and pressure of 23.0 torr;
Class Practice Problem
• (c) the number of moles in 1.50 L at 37 oC and 725 torr;
(d) the temperature at which 0.270 mol occupies 15.0 L at
2.54 atm.
Relating the Ideal-Gas Equation and
the Gas Laws
• If PV = nRT and n and T are constant, then PV = constant and we
have Boyle’s law.
• Other laws can be generated similarly.
• In general, if we have a gas under two sets of conditions, then
P1V1 P2V2

n1T1 n2T2
Class Guided Practice Problem
• A sample of argon gas is confined to a 1.00-L tank at 27.0
oC and 1 atm. The gas is allowed to expand into a larger
vessel. Upon expansion, the temperature of the gas drops
to 15.0 oC, and the pressure drops to 655 torr. What is
the final volume of the gas?
Two ways to work this problem.
Molar Mass
• Density has units of mass over volume.
• Rearranging the ideal-gas equation with M as molar mass
we get
PV  nRT
n
P

V RT
nM
PM
d 
V
RT
Gas Densities
• The molar mass of a gas can be determined as follows:
dRT
M
P
.
Class Guided Practice Problem
• What is the density of carbon tetrachloride vapor at 714
torr and 125 oC?
Volumes of Gases in Chemical
Reactions
• The ideal-gas equation relates P, V, and T to number of
moles of gas.
• The n can then be used in stoichiometric calculations
Class Guided Practice Problem
• The safety air bags in automobiles are inflated by
nitrogen gas generated by the rapid decomposition of
sodium azide, NaN3:
2 NaN3(s)  2 Na(s) + 3N2(g)
If an air bag has a volume of 36 L and is filled with
nitrogen gas at a pressure of 1.15 atm at a temperature of
26 oC, how many grams of NaN3 must be decomposed?
Gas Mixtures and Partial Pressures
• Since gas molecules are so far apart, we can assume they
behave independently.
• Dalton’s Law: in a gas mixture the total pressure is given by
the sum of partial pressures of each component:
Ptotal  P1  P2  P3  
• Each gas obeys the ideal gas equation:
 RT 
Pi  ni 

 V 
• Combining the equations we get:
 RT 
Ptotal  n1  n2  n3  

 V 
Collecting Gases over Water
Collecting Gases over Water
• It is common to synthesize gases and collect them by
displacing a volume of water.
• To calculate the amount of gas produced, we need to
correct for the partial pressure of the water:
Ptotal  Pgas  Pwater
Class Guided Practice Problem
• A gaseous mixture made form 6.00g O2 and 9.00g CH4 is
placed in a 15.0-L vessel at 0 oC. What is the total
pressure in the vessel?
Kinetic Molecular Theory
• Theory developed to explain gas behavior.
• Theory of moving molecules.
• Assumptions:
– Gases consist of a large number of molecules in constant
random motion.
– Volume of individual molecules negligible compared to
volume of container.
– Intermolecular forces (forces between gas molecules)
negligible.
Kinetic Molecular Theory
• Assumptions:
– Energy can be transferred between molecules, but total kinetic
energy is constant at constant temperature.
– Average kinetic energy of molecules is proportional to
temperature.
• Kinetic molecular theory gives us an understanding of
pressure and temperature on the molecular level.
• Pressure of a gas results from the number of collisions
per unit time on the walls of container.
Kinetic Molecular Theory
• Magnitude of pressure
given by how often and
how hard the molecules
strike.
• Gas molecules have an
average kinetic energy.
• Each molecule has a
different energy.
Kinetic Molecular Theory
• As kinetic energy increases, the velocity of the gas
molecules increases.
• Root mean square speed, u, is the speed of a gas molecule
having average kinetic energy.
• Average kinetic energy, , is related to root mean square
speed:
  mu
1
2
2
Application to Gas Laws
• As volume increases at constant temperature, the average
kinetic of the gas remains constant. Therefore, u is
constant. However, volume increases so the gas
molecules have to travel further to hit the walls of the
container. Therefore, pressure decreases.
• If temperature increases at constant volume, the average
kinetic energy of the gas molecules increases. Therefore,
there are more collisions with the container walls and the
pressure increases.
Real Gases: Deviations from Ideal
Behavior
• From the ideal gas equation, we have
PV
n
RT
• For 1 mol of gas, PV/RT = 1 for all pressures.
• In a real gas, PV/RT varies from 1 significantly.
• The higher the pressure the more the deviation from ideal
behavior.
Real Gases: Deviations from Ideal
Behavior
• From the ideal gas equation, we have
PV
n
RT
• For 1 mol of gas, PV/RT = 1 for all temperatures.
• As temperature increases, the gases behave more ideally.
• The assumptions in kinetic molecular theory show where
ideal gas behavior breaks down:
– the molecules of a gas have finite volume;
– molecules of a gas do attract each other
• .
Real Gases: Deviations from Ideal
Behavior
• As the pressure on a gas increases, the molecules are
forced closer together.
• As the molecules get closer together, the volume of the
container gets smaller.
• The smaller the container, the more space the gas
molecules begin to occupy.
• Therefore, the higher the pressure, the less the gas
resembles an ideal gas.
Real Gases: Deviations from Ideal
Behavior
• As
the
gas
molecules
get
closer
together,
the smaller the
intermolecular
distance.
Real Gases: Deviations from Ideal
Behavior
• The smaller the distance between gas molecules, the
more likely attractive forces will develop between the
molecules.
• Therefore, the less the gas resembles and ideal gas.
• As temperature increases, the gas molecules move faster
and further apart.
• Also, higher temperatures mean more energy available to
break intermolecular forces.
Real Gases: Deviations from Ideal
Behavior
• Therefore, the higher the
temperature, the more ideal the
gas.
The van der Waals Equation
• We add two terms to the ideal gas equation one to correct
for volume of molecules and the other to correct for
intermolecular attractions
• The correction terms generate the van der Waals
equation:
nRT
n2a
P
 2
V  nb V
where a and b are empirical constants.
The van der Waals Equation
nRT
n2a
P
 2
V  nb V
Corrects for
molecular volume
Corrects for
molecular attraction

n2a 
 P  2 V  nb  nRT
V 
