CHAPTER 7
Physical States of Matter
General, Organic, & Biological Chemistry
Janice Gorzynski Smith
CHAPTER 7: Physical States of Matter
Learning Objectives:
 States of matter: Gas, Liquid, Solid
 Intermolecular forces:
 London dispersion
 Dipole-dipole
 Hydrogen bonding
 Gas behavior
 Combined gas law, Ideal gas law, Dalton’s law of partial P
 Liquid behavior
 Viscosity, surface tension, vapor pressure
 Solid behavior
 Crystalline vs Amorphous Solids
 Phase changes
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Matter
Gas, Liquid, Solid
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Matter
Gas, Liquid, Solid
Existing as a gas, liquid, or solid depends on:
•The balance between the kinetic energy of its
particles.
•The strength of the interactions between the particles.
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Matter
Intermolecular Forces
Intermolecular forces are the attractive forces that
exist between molecules.
In order of increasing strength, these are:
1. London dispersion forces
2. Dipole–dipole interactions
3. Hydrogen bonding
4. Ion-Dipole & Ion-Ion interactions
The strength of the intermolecular forces determines
whether a compound has a high or low melting point
and boiling point, and thus whether it is a solid,
liquid, or gas at a given temperature.
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Matter
London Dispersion Forces
More e− density
in one region
creates a partial
negative charge (δ−).
Less e− density
in one region
creates a partial
positive charge (δ+).
London dispersion forces are very weak
interactions due to the momentary changes in
electron densityin a molecule.
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Matter
Dipole-Dipole Interactions
Dipole–dipole interactions are the attractive forces
between the permanent dipoles of two polar
molecules.
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Matter
Hydrogen Bonding
Hydrogen bonding occurs when a hydrogen atom
bonded to O, N, or F is electrostatically attracted
to an O, N, or F atom in another molecule.
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Matter
Hydrogen Bonding
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Matter
Ion-Dipole interactions
Attractions between ion and charged end of
polar molecules
(a) Negative ends of water dipoles surround cation
(b) Positive ends of water dipoles surround anion
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Matter
Strength of Intermolecular Attractions
•The boiling point is the temperature at which a
liquid is converted to the gas phase.
•The melting point is the temperature at which a
solid is converted to the liquid phase.
•The stronger the intermolecular forces, the higher
the boiling point and melting point.
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Matter
Intermolecular Attractions in
Gases, Liquids, & Solids
London Dispersion Forces
Weakest
Dipole-Dipole Forces
Hydrogen Bonds
Ion-Dipole Forces
Strongest
Gas < Liquids < Solids
Increasing Average Kinetic Energy
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Gases
Kinetic Molecular Theory
•A gas consists of particles that move randomly and
rapidly.
•The size of gas particles is small compared to the
space between the particles.
•Gas particles exert no attractive forces on each other.
•The kinetic energy of gas particles increases with
increasing temperature.
•When gas particles collide with each other, they
rebound and travel in new directions.
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Gases
Pressure
•When gas particles collide with the walls of a
container, they exert a pressure.
•Pressure (P) is the force (F) exerted per unit area (A).
Pressure
=
Force
=
Area
F
A
760. mm Hg
1 atmosphere (atm) =
760. torr
14.7 psi
101,325 Pa
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Gases
Gas Laws
Boyle’s Law
T1 = T2
Charles’ Law
P1 = P2
P1V1 = P2V2
V1
T1
Gay-Lussac’s
Law
Avagadro’s
Law
P1
V1 = V2
T1
V
n
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=
k
V1
n1
=
=
=
V2
T2
P2
Combined
Gas
Law
P1V1
T1
=
P2V2
T2
T2
V2
n2
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The Combined Gas Law
Gases
P1V1
=
P2V2
T1
T2
initial conditions
new conditions
This equation is used for determining the effect
of changing two factors (e.g., P and T) on the
third factor (V).
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Ideal Gas Law
Gases
All four properties of gases (i.e., P, V, n, and T) can
be combined into a single equation called the ideal
gas law.
PV = nRT
•R is the universal gas constant:
For atm:
For mm Hg:
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R = 0.0821
R
=
62.4
L • atm
mol • K
L • mm Hg
mol • K
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Gases
Standard Temperature & Pressure
•Often amounts of gas are compared at a set of
standard conditions of temperature and pressure,
abbreviated as STP.
•STP conditions are: 1 atm (760 mm Hg) for pressure
273 K (0 oC) for temperature
•At STP, 1 mole of any gas has a volume of 22.4 L.
•22.4 L is called the standard molar volume.
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Dalton’s Law of Partial Pressures
Gases
•Dalton’s law: The total pressure (Ptotal) of a gas
mixture is the sum of the partial pressures of its
component gases.
•For a mixture of three gases A, B, and C:
Ptotal
=
PA
+
PB
+
PC
partial pressures of A, B, and C
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Liquids
Vapor Pressure
•Vapor pressure is the pressure exerted by gas
molecules in equilibrium with the liquid phase.
•Vapor pressure increases with increasing
temperature.
•The boiling point of a liquid is the temperature
at which its vapor pressure = 760 mmHg.
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Liquids
Vapor Pressure
•The stronger the intermolecular forces, the lower
the vapor pressure at a given temperature.
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Liquids
Viscosity
Viscosity is a measure of a fluid’s resistance to flow
freely.
•Compounds with strong intermolecular forces tend to be
more viscous than compounds withweaker forces.
•Substances composed of large molecules tend to be more
viscous, too, because large molecules do not slide past each
other as freely.
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Liquids
Surface Tension
Surface tension is a measure of the resistance of a
liquid to spread out.
•Interior molecules in a
liquid are surrounded by
intermolecular forces on
all sides.
•Surface molecules only
experience intermolecular
forces from the sides and
from below.
Stronger intermolecular
forces the higher the
surface tension
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Solids
Crystalline Solids
•An ionic solid is composed
of oppositely charged ions
(NaCl).
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•A molecular solid is
composed of individual
molecules arranged
regularly (H2O).
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Solids
Crystalline Solids
•A network solid is composed
of a vast number of atoms
covalently bonded together
(SiO2).
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•A metallic solid is a lattice
of metal cations
surrounded by a cloud of
e− that move freely (Cu).
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Solids
Amorphous Solids
•Amorphous solids have no
regular arrangement of
their particles.
•They can be formed when liquids cool too quickly
for regular crystal formation.
•Very large covalent molecules tend to form
amorphous solids, because they can become
folded and intertwined.
•Examples include rubber, glass, and plastic.
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Matter
Property of s, l, g
Summary of Properties
Increases
Decreases
Example
Water has a high boiling point because it has H-bonding,
dipole, and dispersion forces. It is close to heptane
(C7H16), a heavier molecule that only experiences
dispersion forces .
The melting point of ionic solids is extremely high
compared to water which experiences all other
intermolecular forces, but not ion-dipole forces. (NaCl is
1074 K and water is 273 K)
Boiling Point
increasing total
intermolecular forces
decreasing total
intermolecular forces
Melting Point
increasing total
intermolecular forces
decreasing total
intermolecular forces
Retention of V &
Shape
Decreasing
Increasing intermolecular intermolecular forces,
forces and decreasing T & and increasing kinetic
P
energy of particles or T &
P
Gases will fill the volume and shape of the container that
holds them, while solids will retain their own shape and
volume regardless of the container.
Surface Tension
with increasing
intermolecular forces
The molecules on the surface have less neighbors (and
therefore less stabilizing intermolecular forces) and so have a
higher potential energy, which the material will try to reduce
with its shape (sphere): water beading.
Viscosity
Vapor Pressure
with decreasing
intermolecular forces
increasing intermolecular decreasing intermolecular Not just a property of liquids, also gases and solids.
Amorphous solids change shape over time because of their
forces and decreasing
forces and decreasing
viscosity.
temperature
temperature
Decreasing intermolecular Increasing intermolecular Ether has weaker intermolecular forces than water and a
higher vapor pressure, so it evaporates much faster then
forces and increasing
forces and decreasing
water.
temperature
temperature
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Phase
Change
Summary of Phase Changes
fusion
SOLID
evaporation
LIQUID
freezing
GAS
condensation
deposition
sublimation
endothermic
exothermic
System absorbs energy from surrounds in the form of heat
o Requires the addition of heat
System releases energy into surrounds in the form of heat or light
o Requires heat to be decreased
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Phase
Change
Solid  Liquid
solid water
liquid water
The amount of energy needed to melt 1 gram of a
substance is called its heat of fusion.
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Phase
Change
Liquid  Gas
liquid water
gaseous water
The amount of energy needed to vaporize 1 gram of
a substance is called its heat of vaporization.
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Phase
Change
Solid  Gas
solid CO2
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gaseous CO2
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Phase
Change
Heating Curve
•A heating curve shows how a substance’s temperature
changes as heat is added.
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Phase
Change
Cooling Curve
•A cooling curve shows how a substance’s temperature
changes as heat is removed.
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Phase
Change
Energy & Phase Changes
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