Unit 6
Gases, Phase Changes and Introduction
to Thermochemistry
Part I: Gases
Characteristics of Gases
Pressure
Kinetic-Molecular Theory
The Gas Laws
Partial Pressures
Effusion and Diffusion
Real Gases
Properties of Gases
 Three phases of matter
 solid
Definite shape and volume
 liquid
Definite volume, shape of
container
 gas
Shape and volume of
container
Properties of Gases
 A gas is a collection of molecules that are
very far apart on average.
 In air, gas molecules occupy only 0.1% of
the total volume.
 In liquids, molecules occupy ~ 70% of the
total space.
Properties of Gases
 Gases are highly compressible.
 Volume decreases when pressure is applied.
 Gases form homogeneous mixtures with each
other regardless of the identities or relative
proportions of the different gases.
 Water and gasoline = heterogeneous
mixture.
 Water vapor and gasoline vapor =
homogeneous mixture.
Properties of Gases
 Properties of gases vary depending on their
composition.
 Air:
~ 78% N2 and ~ 21% O2
 CO2: colorless, odorless
 CO:
colorless, odorless, highly toxic
 NO2: toxic, red-brown, irritant
 N2O: colorless, sweet odor (laughing gas)
Pressure
 Four quantities are commonly needed to
describe a gas:
 amount of gas (n)
 Temperature (T)
 Volume (V)
 Pressure (P)
Pressure
 Gases exert pressure on the objects in their
surroundings.
 Pressure is caused by collisions between the
gas molecules and objects with which they are
in contact.
 Pressure: the force exerted on a unit area
P = F
A
Pressure
 Atmospheric pressure: the pressure exerted
by gas molecules in the air on all objects
exposed to the atmosphere
 Atmospheric pressure varies with altitude.
Altitude
(ft above
sea level)
Atmospheric Pressure
in. Hg
Torr
psi
0
29.92
760
14.7
5000
24.9
632.5
12.23
10,000
20.58
522.7
10.1
Pressure
Why does atmospheric pressure
decrease with increasing altitude?



Gravity decreases
Density of gas decreases
Fewer gas molecules
Fewer collisions
Lower pressure
Pressure
 Many different units used to report pressure.
 millimeters of Hg (mm Hg)
 inches of Hg (in. Hg)
 pounds per square inch (psi)
 atmosphere (atm)
 torr (torr)
 pascal (Pa) = SI base unit
 kilopascal (kPa)
Must know
units and
abbreviations!!
Pressure
 Relationships between different pressure units:
1 atm = 760 mm Hg
= 760 torr
= 29.92 in. Hg
= 14.7 psi
= 1.01325 x 105 Pa
Must be able to
interconvert
between units.
Memorize the ones
in red…I’ll give you
the others.
You must know that 1 kPa = 1000 Pa
Pressure
Example: The measured pressure inside the eye
of a hurricane was 669 torr. What was the
pressure in atm?
Pressure
Example: On a nice sunny day in Chicago the
barometric pressure was 30.45 in. Hg. What
was the pressure in Pa?
Pressure
Example: On Titan, the largest moon of Saturn,
the atmospheric pressure is 1.631 Pa. What is
the pressure in atm?
Kinetic Molecular Theory
 The behavior of gases can be described and
explained using kinetic molecular theory.
 the “theory of moving molecules”
 You must know the basic ideas that are part
of kinetic molecular theory.
Kinetic Molecular Theory
 Gases consist of large numbers of molecules
that are in continuous, random motion.
 The combined volume of all the molecules of
the gas is negligible compared to the total
volume in which the gas is contained.
 i.e. the molecules are very far apart on
average
Kinetic Molecular Theory
 Attractive and repulsive forces between gas
molecules are negligible.
 Energy can be transferred between molecules
during collisions, but the average kinetic
energy of the molecules does not change as
long as the temperature remains constant.
 Collisions are perfectly elastic.
Kinetic Molecular Theory
 The average kinetic energy of the molecules is
proportional to the absolute temperature.
 At any given temperature all molecules of a
gas have the same average kinetic energy.
 As T (in K) increases,
KE increases.
Gas Laws
 Four variables are needed to define the
physical condition or state of any gas:
 Temperature (T)
 Pressure (P)
 Volume (V)
 Amount of gas (moles: n)
 Equations relating these variables are known as
the gas laws.
Gas Laws
Consider a fixed amount of gas that is confined
to a container with a certain volume.
P
At a specific temperature,
the gas sample will exert
a certain pressure on the
container.
Gas Laws
What will happen to the pressure if the
volume is decreased?
P
Volume
decreases
P
Gas Laws
 As the volume of a fixed quantity of gas
decreases, the pressure increases because:
 gas molecules are more tightly packed
together
 i.e. denser
 more collisions between gas molecules and
the container
 greater pressure
Gas Laws
 Boyle’s Law:
 The volume of a fixed quantity of gas
maintained at constant temperature is
inversely proportional to the pressure.
 Mathematically,
V = k x 1
or
PV = k or P1V1 = P2V2
P
at constant temperature and quantity of gas
Gas Laws
 As liquid nitrogen (-196oC) is poured over a
balloon, the volume of the balloon decreases.
Gas Laws
 Charles’ Law:
 The volume of a fixed amount of gas
maintained at constant pressure is directly
proportional to its absolute temperature.
V = k x T
or
V = k or
T
V1 = V2
T1
T2
At constant pressure and quantity of gas
Remember: T must be in Kelvin
Gas Laws
 On a molecular level, as the temperature of a
gas maintained at constant pressure
decreases,
 KE decreases
 fewer collisions between gas molecules and
the environment (i.e. container)
 volume decreases in order to maintain
constant pressure
Gas Laws
What happens when
you “blow up” a
balloon?
Gas Laws
What happens when
you “blow up” a
balloon?
Gas Laws
What happens when
you “blow up” a
balloon?
Gas Laws
What happens when
you “blow up” a
balloon?
Gas Laws
What happens when
you “blow up” a
balloon?
– the number of moles of
gas (n) increases
and
– the volume of the gas
(balloon) increases
Gas Laws
 Avogadro’s Law:
 The volume of a gas maintained at constant
temperature and constant pressure is
directly proportional to the number of moles
of the gas.
 Mathematically,
V = constant x n
At constant temperature and pressure
Gas Laws
 At any given temperature and pressure, as the
amount of gas increases,
 the number of gas molecules increases
 the number of collisions between gas
molecules and the environment (container)
increases
 the volume must increase in order to
maintain constant pressure
Gas Laws
 In a chemical reaction, we use the coefficients
to tell us how many moles or molecules are
used or produced in a chemical reaction.
N2 (g) + 3 H2 (g)  2 NH3 (g)
 1 mole of nitrogen reacts with 3 moles of
hydrogen to produce 2 moles of ammonia
Gas Laws
 Since the volume of a gas is directly
proportional to the number of moles of gas at
constant temperature and pressure, we can
also use the coefficients to represent the
volume of a gas involved in a reaction.
(Avogadro’s Hypothesis)
N2 (g) + 3 H2 (g)  2 NH3 (g)
 1 liter of nitrogen reacts with 3 liters of
hydrogen to produce 2 liters of ammonia
Gas Laws
 Boyle’s Law, Charles’ Law, and Avogadro’s Law
can be combined to make a more general gas
law:
 Ideal Gas Law:
PV = nRT
where
P = pressure
V = volume
n = moles
T = temperature (K)
R = gas constant
Gas Laws
 The value of the gas constant (R) depends on
the units of P, V, n, and T.
 T must always be in Kelvin
 n is usually in moles
 If P (atm) and V (L),
 then R = 0.08206 atm.L
mol.K
 If P (torr) and V (L),
 then R = 62.36 L.torr
mol.K
I will give you
these on the
test.