BRCC
CHM101
Chapter 6 Notes
(Chapter 5 in older text versions)
Page 1 of 9
Chapter 6: Gases, Liquids, and Solids
There are 3 states of matter: gases, liquids, and solids

The physical properties of solids and liquids differ greatly within each state. ex. solids include
rubber, steel, and wood. Liquids include water, maple syrup, and gasoline.

Gases strongly resemble each other in physical properties.
Kinetic Molecular Theory - explains the behavior of gases
1.
A gas consists of atoms or molecules that are in rapid, random motion. (Think of kids playing on a
playground.)
2.
As temperature increases, gas particles move faster. (You use Kelvin degrees when working with
gases.)
3.
Collisions between gas particles are completely elastic. Energy is transferred from one particle to
another, but total energy is unchanged. (think of pool balls)
4.
The volume occupied by the gas particles themselves is negligible compared to the total volume
occupied by the gas.
5.
Gas particles behave as if they neither attract nor repel one another.
6.
Gas particles collide with the walls of the container. This is the pressure of the gas. The greater the
number of collisions, the greater the pressure is.

A gas described by these 6 assumptions is called an ideal gas.
Pressure, P - is the amount of force exerted by a gas.
How do you measure pressure?
A simple barometer works like this:
Measured as mm of Hg
1 torr = 1 mm Hg
pressure at sea level is usually 760 mm Hg
760 mm Hg is defined as an atmosphere
1 atm = 760 mm Hg
How much force is one atmosphere?
BRCC
CHM101
Chapter 6 Notes
(Chapter 5 in older text versions)
Page 2 of 9
manometer - device for measuring pressure of a gas in a container.
3
Volume, V - is the amount of space a gas occupies, expressed as liters, mL, or cm
Amount - how much gas is present, expressed as moles (n) or as grams (g)
Temperature, T - how hot or cold the gas is, expressed as Kelvin (K) remember K = C°
+
273.15
Gas Laws
Boyle’s Law - as P increases, V decreases with constant T
P1 V1
=
P2 V2
Charles’s Law - as T increases, V increases with constant P (hot air balloons are less dense than air
so they float)
V1
---T1
=
V2
---T2
Gay-Lussac’s Law - as T increases, P increases with constant V
Combined Gas Law -
example:
P1 V1
--------T1
=
P2 V2
--------T2
P1
---T1
=
P2
---T2
MEMORIZE THIS
A gas occupies 3.00 L at 1.50 atm pressure. What is its volume at 10.00 atm at the same
temperature?
remember the combined gas law - since T is constant in this example, you can drop it
P1 V1
=
P2 V2
P1 = 1.50 atm
P2 = 10.00 atm
V1 = 3.00 L
V2 = ?
Use algebra to rearrange the equation to solve for V2
V2
P1 V1
= --------
=
1.50 atm
----------------
3.00 L
x
=
0.45 L
BRCC
example:
CHM101
Chapter 6 Notes
(Chapter 5 in older text versions)
Page 3 of 9
P2
10.00 atm
An autoclave is 100 °C at 1.00 atm. The pressure increases to 1.13 atm. What is the
new temperature?
volume is constant, so you can cancel it out of the combined gas law
P1
---T1
=
P2
---T2
T1 = 100 °C = 100 °C + 273 = 373 K
T2 = ?
P1 = 1.00 atm
P2 = 1.13 atm
rearrange the equation to solve for T2
T2
example:
P2
T1
= ------- x
P1
1.13 atm
------------1.00 atm
=
373 K
x
=
421 K
The gas in a balloon has a volume of 0.50 L and a pressure of 1.0 atm and a temperature
of 393 K. When the gas is heated to 500 K, the volume expands to 3.0 L. What is the
new pressure?
P1 V1
--------T1
=
P2
P1
=
P2 V2
--------T2
x
you need to solve for P2
V1
-----V2
x
T2
-----T1
1.0 atm
=
x
0.50 L
--------3.0 L
500 K
x -------393 K
= .21 atm
Avogadro’s Law - equal volumes of a gas at the same T and P contain equal numbers of molecules.
V1 = V2
At the same Temperature
and Pressure there are the
same number of molecules
in these containers
He
CO2
V1
V2
standard temperature and pressure, STP - a pressure of 1 atm and temperature of 273 K
What volume of gas at STP contains 1 mole of molecules?
22.4 L
BRCC
CHM101
Chapter 6 Notes
(Chapter 5 in older text versions)
Page 4 of 9
Ideal Gas Law - also called the universal gas law
PV = nRT
MEMORIZE
P
V
n
R
T
=
=
=
=
=
pressure, atmospheres
volume, Liters
number of moles, mols
Ideal Gas Law constant, .0821 L atm / K mol
temperature, Kelvin

Holds true for all ideal gases and mixtures of gases at any temperature, pressure, volume, or
amount.

Real gases behave most like ideal gases at low pressures and high temperatures. Why?
example:
How many moles of CO2 are in a 125 mL flask if the pressure is 755 mmHg and the
temperature is 20 °C?
V = 125 mL = 0.125 L
rearrange P V = n R T to solve for n
P = 755 mm Hg = 0.993 atm
T = 20 °C + 273 = 293 K
PV
(0.993 atm) (0.125 L)
-3
n = ----- = ------------------------------------------- = 0.00516 mol = 5.16 x 10 mol
RT
(0.0821 L-atm/K-mol) (293 K)
Moles can be related to Mass (m) and Molecular Weight (MW) by this relation:
m
Substitute for n in the Ideal Gas Law Equation to get: P V = ------- R T
MW
example:
n
=
m
--------MW
If 1.0 L of a gas has 1.15 atm pressure at 40 °C and the mass of this gas is 3.3 g, then what
is the molecular weight of the gas?
RT
(0.0821 L-atm/K-mol) (313 K)
m ------- = (3.3 g) ------------------------------------------ = 74 g/mol
PV
(1.15 atm) (1.0 L)
rearrange to MW =
Dalton’s Law of Partial Pressures - the total pressure of a mixture of gases is the sum of the partial
pressures of the individual gases.
Partial pressure - the pressure the gas would exert if it were alone in the container.
Ptotal
=
example:
P gas 1
+
P gas 2
+
P gas 3
+
...
simple to use - you add or subtract
A tank contains N2 at 2.0 atm and O2 at 1.0 atm. If you add CO2 until the total pressure is
4.6 atm, what is the partial pressure of CO2 ?
P total = P N2
P CO2 =
+
P O2
+
P CO2
solve for P CO2
P total - P N2
-
P O2 =
4.6 atm - 2.0 atm - 1.0 atm = 1.6 atm
BRCC
CHM101
Chapter 6 Notes
(Chapter 5 in older text versions)
Page 5 of 9
Graham’s Law - gas molecules with smaller masses diffuse (and effuse) faster than gas molecules
with larger masses
diffusion - process by which a gas mixes.
effusion - passing through a small opening
Dalton’s Law and Breathing
 A gas will tend to flow from a region where its partial pressure is high to a region where its partial
pressure is low. How does this work in the body?
Lungs :
Tissues:
Blood in veins
Blood in alveoli
Gas flow
P O2 = 40 mm Hg
P O2 = 159 mm Hg
O2 moves from alveoli to blood
P CO2 = 46 mm Hg
P CO2 = 0.3 mm Hg
CO2 moves fro blood to alveoli
Blood in arteries
Blood in tissues
Gas flow
PO2 = 100 mm Hg
P O2 = 30 mm Hg
O2 moves from blood to tissues
P CO2 = 40 mm Hg
P CO2 = 60 mm Hg
CO2 moves from tissues to blood
Intermolecular Forces
Intermolecular forces are attractive forces between molecules which hold them together in the liquid
and solid phases.
In gases, Kinetic Energy is sufficient to overcome intermolecular attractions - except under high
pressure or low temperature. Usually, gases behave as if there are no intermolecular attractions
between gas particles.
As a gas is cooled, KE decreases and intermolecular attractions hold the gas particles together. When
this happens, the gas condenses into a liquid.
Dipole - Dipole Attractions - polar molecules attract one another because the positive portion of one
molecule will attract the negative portion of another molecule.
∂+ ∂∂+ ∂ex. H - Cl ..........H - Cl
Hydrogen Bonding - an especially strong type of dipole-dipole attraction considered to be a 3rd type of
chemical bond. A hydrogen bond will form if:
1. One molecule has a H atom covalently bonded to an O, N, or F atom.
2. The other molecule has an O, N, or F atom.
ex.
H
\ ∂∂+
O ......... H - F
/
H
H-bonds are ten times weaker than covalent bonds.
BRCC
CHM101
Chapter 6 Notes
(Chapter 5 in older text versions)
Page 6 of 9
London Dispersion Forces - weakest intermolecular force, which is formed by temporarily induced
dipoles in non-polar molecules. This happens for a small part of the time where even a non-polar
molecule may have a partial negative and positive charge due to how the electron cloud moves.
Explains why non-polar gases like He will liquefy at low temperature and high pressure.
Relative Strengths of Intermolecular Forces
Covalent Bond
H-Bond
Dipole-Dipole
London Dispersion
50 - 100 kcal
5 - 10 kcal
.1 - 1 kcal
.001 - .2 kcal
Liquids
In liquids, molecules are close together, randomly arranged, and in constant motion. Intermolecular
attractions are strong enough to hold molecules together, but not strong enough to hold molecules in
fixed positions like in a solid.
properties: Molecules of a liquid are close together, so liquids are impossible to compress. a given
sample of liquid has a constant volume. Liquids are fluid so they can flow.
evaporation:

Particles in a sample of matter have a range of energies like that of a bell-shaped curve.

In a liquid, a high energy molecule near the
surface can leave the liquid phase and
become a gas by overcoming the
intermolecular forces holding it to the liquid.
This is called evaporation.

As T increases, more molecules have enough
energy to leave the liquid and go into the gas
phase.

If a container is left open, eventually all the
liquid will evaporate.

If a container is sealed, condensation may
also take place. Gas particles of lower
energy are attracted to the liquid and join the
liquid phase. This is called condensation.
BRCC
CHM101
Chapter 6 Notes
(Chapter 5 in older text versions)
Page 7 of 9
dynamic equilibrium - when evaporation and condensation occur at the same rate.
vapor pressure, VP - the pressure exerted by the vapor above a liquid at equilibrium - VP increases as
T increases.

VP at a specific temperature is an intensive property of a liquid.

Evaporation takes energy. It is an endothermic process. The fastest particles are the warmest.
These leave the liquid phase and enter the gas phase, taking the heat with them. This is why
evaporation of sweating cools us.
boiling point, BP - the T at which the VP of a liquid is equal to the atmospheric pressure.
normal boiling point - the T at which the VP of a liquid equals exactly 1 atmosphere (760 mmHg)
What determines BP?
1. MW and shape - If you look at molecules with similar intermolecular attractions, you see that BP
increases with MW. If intermolecular attractions and MW are similar, then you see that shape
decides the differences in BP. Long, straight chain molecules have more surface area to attract
that molecules of another shape.
2. Type of Molecular Attraction - The stronger the intermolecular attractions, the higher the BP.
ex. CH4 MW = 16
H2O MW = 18
CH4 BP = -161 °C
H2O BP = 100 °C
London dispersion
H-bonding
Surface Tension

The surface of a liquid behaves differently from its interior. A molecule on the interior will experience
equal attractions from other molecules on all sides. At the surface, a molecule is attracted by other
liquid molecules from below, but is not attracted by air molecules from above. There is a net
downward pull toward the interior called surface tension. Water has a high surface tension
because of strong H-bonds.
Densities of Ice and Water

Most pure substances contract when they freeze to bring particles into closest possible proximity so
as the maximize intermolecular attractions. Water is an exception.

The intermolecular attractions in water are H-bonds. The arrangement of water molecules in which
H-bonds are maximized is an arrangement where the molecules are farther apart than in the liquid
state. Water expands when it freezes. Ice is less dense than water so ice floats on top of water.
Solids

When a liquid is cooled, molecules lose energy. If they lose enough energy, they come to occupy
fixed positions, usually in a regular, orderly pattern called a crystal. The process where a substance
changes from the liquid state to the solid state is called crystallization or freezing.

Some elements exist in more than one form, called allotropic forms. ex. Carbon exists as diamond,
graphite, and fullerenes.
BRCC
CHM101
Chapter 6 Notes
(Chapter 5 in older text versions)
Page 8 of 9
Types of solids:
amorphous solids - molecules are not arranged regularly - arranged like the molecules of a liquid, but
occupy fixed positions.
ionic solids - solids composed of ions like NaCl - held by ionic bonds and are hard and brittle with
high melting points.
molecular solids - solids composed of molecules held by intermolecular attractions - softer with lower
melting points than ionic solids.
network solids - a single gigantic molecule held together by covalent bonds - extremely hard with high
melting point - ex. diamond
metallic solids - metals in the solid phase - metal ions surrounded by a sea of valence electrons.
polymeric solids - huge molecules called polymers - molecular solids partly crystalline and partly
amorphous. ex. nylon
Freezing and Boiling Phase Changes
|
|
/
|
/ gas warming
T |
________________/ ____ 100 C
|
/
|
/
|
/ liquid warming
|
/
|
__/ ___ 0 C
| __/ solid warming
|__________________________________
heat added --------------->
heat of fusion - amount of energy which must be absorbed in order to melt one gram of ice.
H2O (s) ------------------------> H2O (l)
heat of fusion = 79.71 cal / g
freezing is the reverse - ex. one gram of H2O gives up 79.71 cal / g of heat to freeze.
heat of vaporization - the amount of energy which must be absorbed to boil one gram of liquid. ex. H2O
requires 539.4 cal / g to boil.
BRCC
CHM101
Chapter 6 Notes
(Chapter 5 in older text versions)
Page 9 of 9
phase diagram - a graph with pressure plotted along the Y axis and temperature plotted along the X axis with
curves indicating the equilibria between the solid, liquid, and gas phases of a substance.
Figure 1: Phase diagram of water.
A. Triple Point – the point where the 3 phase curves of a substance meet – all 3 phases coexist.
B. Normal Freezing point – freezing point at 760 mm Hg or 1 atmosphere.
C. Normal Boiling point – boiling point at 760 mm Hg or 1 atmosphere.
Things you can do with a phase diagram:
1. Find the melting point or boiling point at a given pressure.
2. Find the triple point
3. Find the vapor pressure at a given temperature.
4. Find the conditions where a phase can or cannot exist.
sublimation - when a substance passes directly from the solid to the gas phase because the pressure of the
vapor has been reduced below the triple point.
melting point - the temperature at which the solid and liquid phases of a substance are in equilibrium. The
melting point is the same as the freezing point.
What happens when:
Move from point E (95 °C & 660 mm Hg) to point F (70 °C & 660 mm Hg)? (constant pressure)
Move from point E (95 °C & 660 mm Hg) to point G (95 °C & 760 mm Hg)? (constant temperature)
Move from point E (95 °C & 660 mm Hg) to point H? What are the temperature & pressure for point H?