Chemistry B11
Chapter 5
Gases, Liquids, and Solids
States of matter: the physical state of matter depends on a balance between the kinetic
energy of particles, which tends to keep them apart, and the attractive forces between them,
which tend to bring them together. The attractive forces between molecules are the same in all
three states (gas, liquid, and solid). However, in the gaseous state, the kinetic energy of the
molecules is great enough to overcome the attractive forces between them. Kinetic energy
increases with increasing temperature.
Kinetic molecular theory (properties of an ideal gas):
1. Gases consist of particles, either atoms or molecules, constantly moving through space in
straight lines, in random directions.
2. The average kinetic energy of gas particles is proportional to the temperature in kelvins.
The higher temperature, the faster they move, and the greater their kinetic energy.
3. Molecules collide with each other, much as billiard balls do, bouncing off each other and
changing directions (they may exchange kinetic energies).
4. Gas particles have no volume.
5. There are no attractive forces between gas particles (they do not stick together after a
collision occurs).
6. Molecules collide with the walls of the container, and these collisions constitute the
pressure of the gas.
Pressure: the force per unit area.
P=
Force (F)
Area (A)
Pressure has the different units: atmosphere, torr, mm Hg, in. Hg, Pascal, bar.
1 atm = 760 torr = 760 mm Hg = 101,325 Pascals = 29.92 in. Hg = 1 bar
Note: we use a barometer to measure atmospheric pressure (see Figure 5.2 on page 143).
Note: we use a manometer to measure the pressure of a gas in a container (Figure 5.3 on
page 144).
Gas laws:
Boyle’s law: for a fixed mass of an ideal gas at a constant temperature, the volume of the gas
is inversely proportional to the applied pressure.
P1V1 = P2V2
P1 V1
P2 =
V2
V2=
P1 V1
P2
Charles’s law: for a fixed mass of an ideal gas at a constant pressure, the volume of gas is
directly proportional to the temperature in Kelvin (K).
Dr. Behrang Madani
Chemistry B11
Bakersfield College
V1 V 2
=
T1 T 2
V2=
V1 T2
T1
T2 =
T 1 V2
V1
Gay-Lussac’s law: for a fixed mass of an ideal gas at constant volume, the pressure is
directly proportional to the temperature in Kelvin (K).
P1 P 2
=
T1 T2
P2 =
P1 T 2
T1
T2 =
T1 P2
P1
Combined gas law: the three gas laws can be combined:
P1 V1 P2 V2
=
T1
T2
Avogadro’s law: equal volumes of gases at the same temperature and pressure contain equal
numbers of molecules (regardless of their identity).
Ideal gas law: under most experimental conditions, real gases behave sufficiently like ideal
gases.
PV = nRT
n: amount of the gas in moles (mol)
R: a constant for all gases (ideal gas constant)
V: volume of the gas in liters (L)
T: temperature of the gas in Kelvins (K)
P: pressure of the gas in atmospheres (atm)
Note: at the Standard Temperature and Pressure (STP) (T = 0°C (273K) and P = 1 atm), one
mole of any gas occupies a volume of 22.4 L.
R=
PV (1.00 atm) (22.4 L)
L. atm
=
= 0.0821
nT (1.00 mol) (273 K)
mol. K
Dalton’s law: in a mixture of gases, each molecule acts independently of all the others. The
total pressure, PT, of a mixture of gases is the sum of the partial pressures of each individual
gas.
PT = P1 + P2 + P3 + …
Partial pressure: the pressure that a gas in a mixture of gases would exert if it were alone in
the container.
Dr. Behrang Madani
Chemistry B11
Bakersfield College
Attractive forces:
London dispersion forces: extremely weak attractive forces between atoms or molecules
caused by the electrostatic attraction between temporary induced dipoles. London dispersion
forces exist between all molecules (polar or nonpolar). However, these forces are the only
forces of attraction between nonpolar molecules.
Ne
Ne
Dipole-Dipole interactions: the attraction between the positive end of a dipole of one
molecule and the negative end of another dipole in the same or different molecule. DipoleDipole interactions exist between polar molecules.
Hydrogen bonds: a noncovalent force of attraction between the partial positive charge on a
hydrogen atom bonded to an atom of high electronegativity (most commonly oxygen or
nitrogen) and the partial negative charge on a nearby oxygen or nitrogen.
H
H2O
H2O
O
Note: hydrogen bonds increase the surface tension and the boiling points of the liquids
(higher boiling point than dipole-dipole interactions and London dispersion forces).
Vapor pressure: the partial pressure of a gas in equilibrium with its liquid form in a closed
container.
Equilibrium: a condition in which two opposing physical forces are equal. In a closed
container with an equilibrium condition, the number of vapor molecules reentering the liquid
equals the number escaping from it (as long as the temperature does not change).
Dr. Behrang Madani
Chemistry B11
Bakersfield College
Boiling point: the temperature at which the vapour pressure of a liquid is equal to the
atmospheric pressure. At this temperature, a liquid starts boiling.
Normal boiling point: the temperature at which a liquid boils under a pressure of 1 atm. For
example, 100°C is the normal boiling point of water because that is the temperature at which
water boils at 1 atm pressure.
The boiling point of covalent compounds depends on three factors:
1. Intermolecular forces:
higher attractive forces → higher boiling point
London dispersion forces < Dipole-Dipole interactions < Hydrogen bonds
2. Number of sites for intermolecular interaction (surface area):
Larger surface area of the molecule (more electrons) → more sites for London dispersion
forces → higher boiling point.
For example: C6H14 (hexane) has a larger surface area and it has more electrons than CH4
(methane). Because of its larger surface area, there are more sites for London dispersion
forces. Therefore, hexane has the higher boiling point than methane.
3. Molecule shape: shape of the molecules can affect the boiling point. A linear molecule has
a higher boiling point than a spherical molecule (if both compounds have the same
intermolecular forces and the same molecular weight). Because, a linear molecule has a larger
surface area than a spherical molecule. As surface area increases, contact between adjacent
molecules, the strength of London dispersion forces, and boiling points all increase.
Fusion (melting): changing phase from the solid state into the liquid state.
Sublimation: a transition from the solid state directly into the vapour state without going
through the liquid state (example: sublimation of dry ice (CO2))
Amorphous solid: its atoms are arranged randomly (examples: wax and glass). They have
much lower melting points than network solids.
Network solid or network crystal: in these solids, a very large number of atoms are
connected by covalent bonds. These atoms are arranged in a symmetric form (crystalline
form). They are hard and they have very high melting points (examples: diamond and quartz).
Heating curve: it shows us the phase changes (a change from one physical state (gas, liquid,
or solid) to another).
Note: during the phase changing (evaporation, melting, sublimation, condensation, freezing),
the temperature of the sample stays constant.
Dr. Behrang Madani
Chemistry B11
Bakersfield College
Heat added (cal)
Heating curve of ice
Phase diagram: by using this diagram, we can show all phase changes for any substance.
Temperature is plotted on the x-axis and pressure on the y-axis. The lines separating the
different states and contain all the boiling points (A-C), all the melting points (A-B) and all
the sublimation points (A-D). A phase diagram illustrates how one may go from one phase to
another by changing the temperature and the pressure.
Triple Points: At this unique point on the phase diagram, all three phases coexist.
Dr. Behrang Madani
Chemistry B11
Bakersfield College