Download States of Matter and Intermolecular Forces Chapter 11

States of Matter and
Intermolecular Forces
Chapter 11
11-1 States and State Changes
Solids
• Particles have an orderly,
fixed arrangement
• Fixed volumes and shapes
Liquids
• Particles move easily
past one another (have
more energy)
• Fixed volume, no fixed
shape
Viscosity
• Ability to Flow
• Honey is very viscous
Surface Wetting
Cohesion
• Stick to each other
Adhesion
• Stick to something else
Capillary Action
• The movement of
water up through a
tube – because of
adhesion and
cohesion
Surface Tension
• Cohesive forces
• Causes liquids to minimize surface
area
• That’s why water drops are round
Gas
• Particles are
independent
• Far apart
• No fixed volume or
shape
• Gases and liquids are
fluids
Changing State
• Freezing – liquid becomes a solid
• Melting – solid becomes a liquid
• Evaporation – liquid becomes gas
• Condensation – gas becomes liquid
• Sublimation – solid becomes gas
• Deposition – gas becomes solid
Temperature, Energy, and State
Evaporation
• High energy particles change to gas
• Causes the substance to cool
Boiling Point
• The temperature at which bubbles
of vapor rise to the surface
• Also depends on atmospheric
pressure
Intermolecular Forces
11-2
Attraction between Particles
• Takes energy to separate particles (change state)
• The stronger the force, the more energy it takes
• The boiling and melting
point is a good measure of
the strength of the force
• Strong force of attraction =
high boiling point
Force of attraction in Ions
• Higher force of attraction then between molecules
• High melting points
• Smaller ions  larger force (NaCl > KCl)
• Larger charge  larger force (CaF2 > NaCl)
Intermolecular Forces
• The Force of Attraction between molecules
Types of Intermolecular Forces
• Dipole-Dipole Forces
• Hydrogen Bonds
• London Dispersion Forces
• All are short range
• Little effect on gases
• Many gases have low
boiling point (that is why
they are gases)
Polar
Molecule
• A molecule that has an unequal distribution of charge
• One end slightly positive, One end slightly negative
• Caused by difference in electronegativity of the atoms
Dipole-Dipole Forces
• Interaction between polar molecules
• Positive end of one molecule attracts the negative end of another
Dipole-Dipole Forces and Boiling Point
• The more polar the molecules, the stronger the force
between them, the higher the boiling point
Hydrogen Bonds
• When a hydrogen atom of one molecule is attracted to an atom of a different
molecule
• Water
• Can create a larger
difference in
electronegativity
• Also hydrogen is small
and has only 1
electron
• Which increases the
bond strength
• Which increases the
boiling point
Hydrogen Bonds
Hydrogen Bonds
and Water
• Water has unique
properties, because of
hydrogen bonds
• Can form multiple
hydrogen bonds 
Strong intermolecular
forces
Solid water is less dense than liquid water
• Ice Floats
• Ponds freeze
from top down
• Expanding ice
cracks rocks
and concrete
London
Dispersion
Forces
• The force that hold nonpolar molecules together
• The weakest of the
intermolecular forces
• Explains why some nonpolar molecules are not
gases
London Dispersion Forces
• Nonpolar molecules can become temporary dipoles (electrons move from side
to side)
• Causes molecules to attract each other
• Nearby
molecules
always
attract
• The more
electrons, the
stronger the
force
London Dispersion Forces
Energy of State Changes
11-3
Enthalpy
• The total energy of a system
Entropy
• A measure of
system’s
disorder
Enthalpy of Fusion
• The energy added during melting or removed during
freezing
• AKA the heat of fusion
Entropy of Fusion
• The increase of entropy when a solid melts
Enthalpy of Vaporization
• The energy added during evaporation
Entropy of Vaporization
• The increase of entropy when a liquid evaporates
• Much larger than entropy of fusion
The molar enthalpy of fusion
• The heat energy needed
to melt 1 mol of a
substance
For water it is
6.01 kJ/mol
The molar enthalpy of vaporization
• The heat energy
needed to evaporate 1
mol of a substance
For water it is
40.67 kJ/mol
Phase Equilibrium
11-4
System
• A set of components that are being studied
Phase
• A region that has the same
composition and properties
throughout
Lava lamp – Two phases of liquid
- Different chemical compositions
Phase
• Water – Two phases, same
chemical composition Different States
Dynamic Equilibrium
• The net amount of
substance in a given phase
stays the same
• Eg. The rate of evaporation
equals the rate of
condensation
Which of these?
Vapor Pressure
• The pressure exerted by a gas in equilibrium with a liquid
• Boiling point – The
temp at which
vapor pressure
equals the external
pressure
As temperature increases,
vapor pressure increases
• Normal Boiling Point – when
vapor pressure equals the
atmospheric pressure
Phase Diagrams
• A graph of the
relationship
between the state
of a substance and
its temperature
and pressure
Phase Diagrams
• 3 lines
• Vapor pressure for
liquid-gas equilibrium
A-B
• Liquid-solid
equilibrium A-D
• Solid gas equilibrium
A-C
Triple Point
• The
temperature
and pressure
at which all
three states
are in
equilibrium
Critical Point
• The temperature and pressure at which the gas and liquid states become
identical
• Called a
supercritical
fluid
Supercritical Fluid