1. Physical States of Matter
CHAPTER 7
Physical States of Matter
General, Organic, & Biological Chemistry
Janice Gorzynski Smith

2. 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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3. Smith. General Organic & Biolocial Chemistry 2nd Ed.
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:
London dispersion forces
Dipole–dipole interactions
Hydrogen bonding
4. Ion-Dipole & Ion-Ion interactions
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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+).
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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
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(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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Smith. General Organic & Biolocial Chemistry 2nd Ed.

12. Gas < Liquids < Solids
Matter
Intermolecular Attractions in
Gases, Liquids, & Solids
London Dispersion Forces
Dipole-Dipole Forces
Hydrogen Bonds
Ion-Dipole Forces
Weakest
Strongest
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Gas < Liquids < Solids
Increasing Average Kinetic Energy
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13. Smith. General Organic & Biolocial Chemistry 2nd Ed.
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.
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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).
Force F
Pressure = =
Area A
Use modeling kit!
760. mm Hg
760. torr
1 atmosphere (atm) =
14.7 psi
101,325 Pa
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Gases
Gas Laws
Boyle’s Law
T1 = T2
P1V1 = P2V2
Charles’ Law
P1 = P2
Gay-Lussac’s Law
V1 = V2
Avagadro’s Law
Combined
Gas
Law
Use modeling kit! V V1 V2
= k = n n1 n2
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16. Smith. General Organic & Biolocial Chemistry 2nd Ed.
Gases
The Combined Gas Law
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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Gases
Ideal Gas Law
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:
L • atm
mol • K
For atm: R =
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L • mm Hg
mol • K
For mm Hg: R =
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18. Standard Temperature & Pressure
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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19. Smith. General Organic & Biolocial Chemistry 2nd Ed.
Gases
Dalton’s Law of Partial Pressures
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
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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.
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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.
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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).
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).
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.
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Examples include rubber, glass, and plastic.
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Matter
Summary of Properties
Property of s, l, g
Increases
Decreases
Example
Boiling Point
increasing total intermolecular forces
decreasing total intermolecular forces
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 .
Melting Point
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)
Retention of V & Shape
Increasing intermolecular forces and decreasing T & P
Decreasing intermolecular forces, and increasing kinetic 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
with decreasing 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
increasing intermolecular forces and decreasing temperature
decreasing intermolecular forces and decreasing temperature
Not just a property of liquids, also gases and solids. Amorphous solids change shape over time because of their viscosity.
Vapor Pressure
Decreasing intermolecular forces and increasing temperature
Increasing intermolecular forces and decreasing temperature
Ether has weaker intermolecular forces than water and a higher vapor pressure, so it evaporates much faster then water.
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Phase Change
Summary of Phase Changes
fusion
evaporation
SOLID LIQUID GAS
freezing
condensation
deposition
sublimation
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endothermic
System absorbs energy from surrounds in the form of heat
Requires the addition of heat
exothermic
System releases energy into surrounds in the form of heat or light
Requires heat to be decreased
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Phase Change
Solid  Liquid
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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
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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
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gaseous CO2
solid CO2
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Phase Change
Heating Curve
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A heating curve shows how a substance’s temperature changes as heat is added.
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Phase Change
Cooling Curve
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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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Smith. General Organic & Biolocial Chemistry 2nd Ed.