Chapter #12
States of Matter
Inter-particle Forces
Chapter #12
Intermolecular Forces (IMF) Topics
• Molecular interactions
• Properties of Liquids and Solids
• IMF Force Properties
• Evaporation and Condensation
• Melting Freezing and Sublimation
• Types of IMF
• Types of Crystalline Solids
12.1 Intramolecular Forces
• “Within” the molecule and are called
covalent and ionic bonds.
• Molecules are formed by sharing
electrons between the atoms.
• Covalent bonds hold the atoms of a
molecule together.
• Ionic bonds hold oppositely charged
ions together in a formula unit
12.1 Intermolecular Forces
• Forces that occur between molecules.
• Intramolecular (ionic and covalent)
bonds are stronger than intermolecular
forces.
12.6 Intermolecular Forces
• There are two kinds of intermolecular
forces that occur between molecules.
 Dipole–dipole forces (between polar molecules)
 Hydrogen bonding (between hydrogen of one polar
molecule and O, N, F of another polar molecule)
 London dispersion forces (between nonpolar
molecules)
12.6 Dipole–Dipole Attraction
Polar molecules contain different
atoms which have different
electronegativity's and can be
thought to be like a small bar
magnet. Oppositely charged
ends attract and like ends repel.
Molecules organize themselves
to maximize the attractive forces
and minimize repulsive forces
giving molecules unique shapes.
12.6 Dipole-Dipole Forces
• Dipole moment – molecules with polar bonds often
behave in an electric field as if they had a center of
positive charge and a center of negative charge.
• Molecules with dipole moments can attract each other
electrostatically. They line up so that the positive and
negative ends are close to each other.
• Only about 1% as strong as covalent or ionic bonds.
12.6 Hydrogen Bonding
• Strong dipole-dipole forces.
• Hydrogen is bound to a highly
electronegative atom – nitrogen,
oxygen, or fluorine.
12.6 Hydrogen Bonding in Water
• Blue dotted lines
are the
intermolecular
forces between
the water
molecules.
12.6 Hydrogen Bonding
• Affects physical properties
 Boiling point
12.6 London Dispersion Forces
• Instantaneous dipole that occurs
accidentally in a given atom induces a
similar dipole in a neighboring atom.
• Significant in large atoms/molecules.
• Occurs in all molecules, including
nonpolar ones.
12.6 London Dispersion Forces
• Become stronger
as the sizes of
atoms or
molecules
increase.
12.6 London Dispersion Forces
Nonpolar Molecules
Ion Dipole Forces
Ion Dipole forces are typically found in aqueous
solutions of ionic compounds, such as salt water. In
the figure below the nonbonding electrons found on the
oxygen atom are strongly attracted to the sodium ion.
Overview of Particle Forces
Type of Force
Dispersion Force
(London Force)
Dipole-dipole
force
Hydrogen bond
Ion-dipole
Relative Strength
Weak, increasing
with size
moderate
Strong
Very strong
Present in
Examples
atoms/molecules
H2 (g)
Polar
molecules
HCl
Molecules
with hydrogen
bonded to
N,O, or F
Mixtures of ionic
and polar
compounds
HF, NH3,
H 2O
Na+,Cl-,
and H2O
12.2 The Liquid State
In the liquid state the particles are randomly
arranged (like in a gas).
The particles are closer to each other than in gases
so the density of liquids is greater than that of
gases, thus attractive forces between liquid particles
is stronger than in gases.
Liquids adopt the shape of the container into which
they are placed.
Unlike gases liquids have a fixed volume and
density. This is because they have greater
attractive forces between particles holding them
close together.
12.2 The Liquid State
As the particles in liquids are very close to one
another they have small compressibility.
When particles of a liquid are heated the particles
will move around more rapidly.
As they have many close neighbors they may only
travel a short distant before undergoing a collision
and bouncing back in the opposite direction. For
this reason liquids have little thermal expansion.
12.2 Gas/Liquid Comparison
Liquid state
Gaseous state
Particles are close together
Not held in fixed positions
Take the shape of container
Particles are far apart
Have fixed volume
Completely fill container
Little compressability
Easily compressed
Small thermal expansion
Moderate thermal expansion
12.2 Solid State
In the solid state particles are held in fixed lattice
positions.
Cohesive forces are much more dominant for solids
than dispersive forces.
Unlike gases solids have fixed shape, volume and
density.
The particles may only move a small amount around
their fixed positions so solids have little thermal
expansion.
In solids the particles are close together and so they
have high density.
12.2 Solid State
 Strong cohesive forces
 Particles in fixed lattice positions
 Constant shape
 Constant density
 Constant volume
 Minimal compressibility
 Little thermal expansion
12.3 Surface Tension
Viscosity
Viscosity is the resistance to flow. Viscose liquids slowly flow
such as syrup, while non viscose liquids flow rapidly, such as
water. The stronger the antiparticle forces the more viscose
the liquid
12.3 Cohesive and Adhesive Forces
Beading and Wetting
Beading is desired for cars, while wetting is desired for the
dishwasher. Beading in a dishwasher produces spots on the
glassware. Beading (strong cohesive forces) and wetting
(strong adhesive forces) are illustrated below.
12.4 Changes in States of Matter
When matter takes energy from its surroundings
(endothermic processes) the kinetic energy of the
particles increases resulting in greater dispersive
forces and the particles moving away from each
other.
i.e. Processes in which particles move away from
each other (solid to liquid change of state) are
endothermic.
12.4 Changes in States of
Matter
When matter releases energy to its surroundings
(exothermic processes) the kinetic energy of the
particles decreases resulting in greater cohesive
forces and the particles moving closer to each
other. i.e. Processes in which particles move closer
to each other (liquid to solid change of state) are
exothermic.
H2O (s)
H2O (l)
H2O (g)
12.4 Changes in States of Matter
When matter releases energy to its surroundings
(exothermic processes) the kinetic energy of the
particles decreases resulting in greater cohesive
forces and the particles moving closer to each other.
i.e. Processes in which particles move closer to each
other (liquid to solid change of state) are exothermic.
H2O (s)
melting
H2O (l)
H2O (g)
When matter releases energy to its surroundings
(exothermic processes) the kinetic energy of the
particles decreases resulting in greater cohesive
forces and the particles moving closer to each other.
i.e. Processes in which particles move closer to each
other (liquid to solid change of state) are exothermic.
H2O (s)
melting
H2O (l)
H2O (g)
When matter releases energy to its surroundings
(exothermic processes) the kinetic energy of the
particles decreases resulting in greater cohesive
forces and the particles moving closer to each other.
i.e. Processes in which particles move closer to each
other (liquid to solid change of state) are exothermic.
H2O (s)
melting
H2O (l)
H2O (g)
evaporation
When matter releases energy to its surroundings
(exothermic processes) the kinetic energy of the
particles decreases resulting in greater cohesive
forces and the particles moving closer to each other.
i.e. Processes in which particles move closer to each
other (liquid to solid change of state) are exothermic.
H2O (s)
melting
H2O (l)
H2O (g)
evaporation
sublimation
When matter releases energy to its surroundings
(exothermic processes) the kinetic energy of the
particles decreases resulting in greater cohesive
forces and the particles moving closer to each other.
i.e. Processes in which particles move closer to each
other (liquid to solid change of state) are exothermic.
condensation
H2O (s)
melting
H2O (l)
H2O (g)
evaporation
sublimation
When matter releases energy to its surroundings
(exothermic processes) the kinetic energy of the
particles decreases resulting in greater cohesive
forces and the particles moving closer to each other.
i.e. Processes in which particles move closer to each
other (liquid to solid change of state) are exothermic.
condensation
H2O (s)
melting
H2O (l)
H2O (g)
evaporation
sublimation
When matter releases energy to its surroundings
(exothermic processes) the kinetic energy of the
particles decreases resulting in greater cohesive
forces and the particles moving closer to each other.
i.e. Processes in which particles move closer to each
other (liquid to solid change of state) are exothermic.
freezing
H2O (s)
melting
condensation
H2O (l)
H2O (g)
evaporation
sublimation
When matter releases energy to its surroundings
(exothermic processes) the kinetic energy of the
particles decreases resulting in greater cohesive
forces and the particles moving closer to each other.
i.e. Processes in which particles move closer to each
other (liquid to solid change of state) are exothermic.
deposition
freezing
H2O (s)
melting
condensation
H2O (l)
H2O (g)
evaporation
sublimation
The specific heat of a solid, liquid, or gas is the amount of
energy required to increase the temperature of one gram by
one degree C.
The heat of fusion of a solid is the amount of heat required
to change one gram of solid to liquid.
The heat of vaporization of a liquid is the amount of heat
required to change one gram of liquid to a gas.
Compound Specific Heat
Heating Curve
Calculate the energy in Kj required to change 5.00g of
ice into water at 66°C.
First find energy to melt ice using heat of fusion.
5.00 g 333 j
g
Kj
= 1.665 Kj
3
10 j
Now find the energy to heat water from 0°C to 66°C
4.184 j
g-°C
5.00 g 66 °C
Kj
103 j
= 1.381 Kj
1.665 Kj + 1.381 Kj = 3.05 Kj
Evaporation is the name of the process by which a
liquid becomes a gas.
Considering our previous discussion would you
expect evaporation to be endothermic or exothermic?
Evaporation takes place from the surface of a liquid.
How do you expect the rate of evaporation to be
affected by the surface area of the liquid?
We can define the vapor pressure of a liquid as:
“the pressure exerted by a vapor that is in
equilibrium with its liquid.”
This is a little confusing so lets take sometime to
explain.
If we place a liquid in a sealed
conatiner with some empty space
above the liquid initially there will be
no vapor or gas above that liquid.
Those molecules at the
surface of the liquid with
sufficient energy will leave the
liquid and enter the gas phase.
Some of the vapor molecules
will strike the surface of the
liquid and return to the liquid
phase.
When the rate at which the liquid is entering the gas
phase equals the rate at which the vapor is returning
to the liquid phase we say the system is at
equilibrium. After this time the liquid level will remain
constant. The pressure exerted by the vapor at this
time is called the vapor pressure.
The vapor pressure of a liquid decreases with
molecular mass.
The vapor pressure of a increases with temperature.
The vapor pressure of a liquid depends upon the
chemical nature of the liquid.
Those molecules that have strong intermolecular
attractive forces have lower vapor pressures than
expected for their molecular mass.
Lets consider some examples:
At 20oC H2O (MW = 18 gmol-1) has a vapor pressure of 17.5 torr !!
This due to strong hydrogen bonds between water molecules.
As we increase the temperature the vapor pressure of
a liquid increases.
The temperature at which the vapor pressure equals the
external pressure (atmospheric pressure) is called the
boiling point.
Bubbles of vapor with atmospheric pressure may form
anywhere in the liquid and rise to the surface at the
boiling point of a liquid.
What is the effect of lowering the atmospheric
pressure on the boiling point?
12.5 Melting and Boiling Points
• In general, the stronger the
intermolecular forces, the higher the
melting and boiling points.
As we increase the temperature the vapor pressure of
a liquid increases.
The temperature at which the vapor pressure equals the
external pressure (atmospheric pressure) is called the
boiling point.
Bubbles of vapor with atmospheric pressure may form
anywhere in the liquid and rise to the surface at the
boiling point of a liquid.
What is the effect of lowering the atmospheric
pressure on the boiling point? Lowering of the
boiling point, thus food takes longer to cook at
higher elevations.
Molecular Solids: Solids made of individual molecules
attached by interpartical forces.
Examples: ice, dry ice, and sugar
Properties: Low melting, Non conductors of heat and
electricity, and brittle.
Ice
Covalent Network Solids: Atoms or molecules held together by
covalent bonds.
Examples: Diamond, Bucky balls, Quartz
Properties: High melting, nonconductors of heat and electricity and
brittle
Quartz
Bucky Balls
12.7 Types of Crystalline Solids
Ionic Solids: Solids held together by ionic bonds. They are
never ending arrays of oppositely charged ions.
Examples: Salt (sodium chloride), Copper (II) nitrate
Properties: High melting, brittle, nonconductors of
electricity.
Sodium Chloride
Atomic Solids: Solids consisting of individual atoms.
Examples: Selenium, sulfur, and Neon
Properties: Low melting and boiling points, non conductors of
electricity, and brittle.
sulfur
12.7 Types of Crystalline Solids
Metallic Solids: Solids composed of metals
Examples: Copper, Iron, Silver and Gold
Properties: Variable melting and boiling points.
– 39°C, while Tungsten melts at 3422°C
Electron Sea Model
Eg. Mercury melts at
End Chapter #12