Energy and phases
All matter can undergo
changes in its state.
These changes have to do
with the amount of
energy in the particles of
matter.




Kinetic theory of
matter1.All matter is made
of particles
2.These particles are
in constant motion.
More energy causes
the particles to move
faster.
Between 0°C and
100 °C, water is a
liquid. In the
liquid state, water
molecules are
close together, but
can move about
freely.
Below 0°C, water
solidifies to become
ice. In the solid state,
water molecules are
held together in a
rigid structure.
At 100°C, water
becomes water
vapor, a gas.
Molecules can
move randomly
over large
distances.



Matter has five
states or phases
Solid : A definite
shape and
volume
Lower energy




Liquid: A definite
volume but it takes
the shape of its
container
Higher energy
Gas : No definite
shape or volume
Yet even higher
energy


Plasma : No definite
shape or volume and
whose particles have
broken apart
Bose-Einstein
Condensate: Gases
near absolute zero
forming a super fluid
Plasma is by far the most common form
of matter. Plasma in the stars and in the
tenuous space between them makes up
over 99% of the visible universe and
perhaps most of that which is not
visible.




Liquids: Are not very
compressible
Useful in hydraulics
Viscosity: Liquids
resistance to flow
Surface tension:
holds the liquid
together



Gases: Fill their container
and can change pressure
http://www.stolaf.edu/
people/giannini/flashani
mat/transport/project.s
wf
Diffusion: Spreading of
particles through an area
until they are uniformly
distributed


Changes of state: When matter changes form, it
is a physical change and has to do with the
energy of the material
Ex. Boiling, melting, freezing, condensing –all
require an energy change


Heat of fusion:
Energy required to
go from solid state to
liquid state (For
water 334kJ/kg)
Melting point –
different for
substances







helium−269
hydrogen−253
Iron-2887
Graphite (carbon)3900
Diamond (carbon)4827
Tungsten-5660
Gold-3080

Heat of Vaporization:
Energy required to go
from liquid to gas. (For
water 2260kJ/kg)



Condensation: A gas
changes to a liquid when
cooled to or below its
boiling point
Vaporization is at boiling
point or below
Evaporation: A liquid
changes to a gas without
reaching its boiling point


Sublimation:
Changing from a
solid to a gas without
existing as a liquid
Deposition :
changing from a gas
to a solid without
being a liquid


Latent heat: heat
absorbed without a
change in
temperature (stored
until a phase change)
Is water a solid or
liquid at 0C? Why?
Honor only
Intermolecular Forces
A phase is a homogeneous part of the
system in contact with other parts of the
system but separated from them by a welldefined boundary.
2 Phases
Solid phase - ice
Liquid phase - water
11.1
Intermolecular forces are attractive forces between
molecules.
Intramolecular forces hold atoms together in a
molecule.
Intermolecular vs Intramolecular
• 41 kJ to vaporize 1 mole of water (inter)
• 930 kJ to break all O-H bonds in 1 mole of water
(intra)
“Measure” of intermolecular
Generally,
force
intermolecula
boiling point
r forces are
melting point
much weaker
than
DHvap
intramolecula
DHfus
11.2
r forces.
Intermolecular forces are feeble; but
without them, life as we know it would
be impossible. Water would not
condense from vapor into solid or
liquid
forms if its molecules didn't attract
each other. Intermolecular forces are
responsible for many properties of
molecular compounds, including
crystal
structures (e. g. the shapes of
snowflakes), melting points, boiling
points,
heats of fusion and vaporization,
surface tension, and densities.
Intermolecular forces pin gigantic
molecules like enzymes, proteins, and
DNA into the shapes required for
biological activity.
http://www.nationmaster.com/encyclopedia/Image:Myoglobin.png
Intermolecular Forces
1. London Forces (Dispersion Forces)
2. Dipole-Dipole Interactions
3. Ion-Dipole Interactions (Salt dissolving in solution
4. Hydrogen Bonding
Dispersion Forces
Occur between every compound and arise from the net attractive forces
amount molecules which is produced from induced charge imbalances
Figure 10-8 Olmsted Williams
The magnitude of the Dispersion Forces
is dependent upon how easily it
is to distort the electron cloud.
The larger the molecule the greater
it’s Dispersion Forces are.
Figure 10-9 Olmsted Williams
The boiling point of alkanes increase with the length
of the carbon chain. Long-chain alkanes have larger
dispersion forces because of the increased polarizability
of their larger electron cloud.
Olmsted Williams
Fig 10-10
Pg 437
How molecular shape affects the strength of the dispersion forces
The shapes of the molecules also matter. Long thin molecules can develop
bigger temporary dipoles due to electron movement than short fat ones
containing the same numbers of electrons.
Long thin molecules can also lie closer together - these attractions are at their
most effective if the molecules are really close.
For example, the hydrocarbon molecules butane and 2-methylpropane both
have a molecular formula C4H10, but the atoms are arranged differently. In
butane the carbon atoms are arranged in a single chain, but 2-methylpropane
is a shorter chain with a branch.
Butane has a higher boiling point because the dispersion forces are greater.
The molecules are longer (and so set up bigger temporary dipoles) and can lie
closer together than the shorter, fatter 2-methylpropane molecules.
http://www.chemguide.co.uk/atoms/bonding/vdw.html
Polarizability
the ease with which the electron distribution in the
atom or molecule can be distorted.
Polarizability increases with:
• greater number of electrons
• more diffuse electron cloud
Dispersion
forces usually
increase with
molar mass.
11.2
Is the Molecule Polar?
We have already talked about diatomic molecules. The more
Electronegative atom will pull the electron density of the bond
Closer to itself giving it a partial negative charge leaving the other
Atom with a partially positive charge. Thus giving the molecule
A dipole moment.
But what about molecules made up of more than two molecules?
Dipole-Dipole Forces
Attractive forces between polar
molecules
Orientation of Polar Molecules in a Solid
11.2
Dipole Forces occur between molecules containing a dipole moment.
The positive end of the dipole moment on one mole is attracted to the
Negative end of the dipole moment on a nearby molecule.
Consider 2-methyl propane
(left) and acetone (right)
Both compounds are about
Equal in size and shape therby
Having similar dispersion forces,
But Acetone contains an
Oxygen (red) and causes the
Molecule to have a dipole
Moment allowing it to have
Dipole forces and thus a
Higher boiling point
Figure 10-11
Olmsted Williams
Ion-Dipole Forces
Attractive forces between an ion and a polar
molecule
Ion-Dipole Interaction
The larger the charge the stronger the force
11.2
Olmsted Williams Fig 10-34
A molecular picture showing the ion-dipole
Interaction that helps a solid ionic crystal dissolve
in water. The arrows indicate ion-dipole interactions.
What type(s) of intermolecular forces
exist between each of the following
molecules?
HBr
HBr is a polar molecule: dipole-dipole forces. There
are also dispersion forces between HBr molecules.
CH4
CH4 is nonpolar: dispersion forces.
SO2
S
SO2 is a polar molecule: dipole-dipole forces. There
are also dispersion forces between SO2 molecules.
11.2
Hydrogen Bond
The hydrogen bond is a special dipole-dipole
interaction between they hydrogen atom in a polar NH, O-H, or F-H bond and an electronegative O, N, or F
atom.
or
H… B
H… A
A
A
A & B are N, O, or F
11.2
Intermolecular Forces
1. London Forces (Dispersion Forces)
2.
Dipole-Dipole Interactions
3. Ion-Dipole Interactions (Salt dissolving in solution)
4. Hydrogen Bonding
These forces affect how molecules will interact with each other and
As a general rule as the strength of the force increases the boiling
Point of the compound increases
Liquids and Surface Tension
Surface tension is the amount of energy required to
stretch or increase the surface of a liquid by a unit
area.
Strong
intermolecul
ar forces
High
surface
tension
11.3
Properties of Liquids
Cohesion is the intermolecular attraction between like
molecules
Adhesion is an attraction between unlike molecules
Adhesion
Cohesion
11.3
Condensation
Evaporation
T2 > T1
Least
Order
Greatest
Order
11.8
The equilibrium vapor pressure is the vapor
pressure measured when a dynamic equilibrium exists
between condensation and evaporation
H2O
(l)
H2O
(g)
Dynamic Equilibrium
Rate of
condensation
=
Rate of
evaporation
A substance with a high
Vapor pressure is considered
To be volitile therefore, the lower
The boiling point the higher the
Vapor pressure and the weaker
The intermolecular forces
11.8
The boiling point is the temperature at which the
(equilibrium) vapor pressure of a liquid is equal to
the external pressure.
The normal boiling point is the temperature at
which a liquid boils when the external pressure is 1
atm.
11.8
H2O
(l)
The melting point of a
solid or the freezing
point of a liquid is the
temperature at which the
solid and liquid phases
coexist in equilibrium
Freezing
(s)
Melting
H2O
11.8
11.8
H2O
(g)
Molar heat of
sublimation (DHsub) is
the energy required to
sublime 1 mole of a solid.
Deposition
(s)
Sublimation
H2O
DHsub = DHfus + DHvap
( Hess’s Law)
11.8
A phase diagram summarizes the conditions at
which a substance exists as a solid, liquid, or gas.
Phase Diagram of Water
11.9
11.9