1
KINETIC THEORY OF MATTER
- molecules in matter are always in motion
- speed of molecules is proportional to the temperature
THE STATES OF MATTER
1. Gas
a) ideal gas
- molecules move freely
- molecules have no interactions with each other
- no attraction or repulsion
- like ping-pong balls
b) real gas
- molecules move freely
- molecules have small attractions for each other
- like “sticky” ping-pong balls
- attractions decrease pressure of gas
- molecules have a volume
- inclusion of the molecular volume increases the pressure of gas (more crowding)
2. Liquid
- molecules move but movement is heavily influenced by other molecules
- molecules must move through a crowd of molecule
- ideal gas can never condense to be a liquid.
3. Solid
- molecules (or ions) vibrate about fixed position in crystal lattice
- motion is not great enough to overcome attractions
- molecules (or ions) stay in a specific arrangement
CHARGE DISTRIBUTIONS OF MATTER
1. Ion
2. Dipole
- A polar molecule has a dipole.
- The value of a dipole depends on
a) polarity of bonds
b) molecular geometry
H
H
O
H
C
N
2
3. Induced Dipole
- A molecule can have its electron cloud polarized by neighboring charge.
- Consider F2.
F
F

F
F
- Fluorine molecule has no dipole.
- Consider F2 next to sodium ion, Na+.
Na+
F
F
+
- Negative electron cloud adjusts to move closer to positive ion.
- Fluorine molecule now has a dipole!
- Sodium ion induces dipole in fluorine.
- Thus fluorine has an induced dipole.
Some atoms in molecules have electron clouds that adjust easier than others.
Na+
Na+
F
F
I
I
- Sodium ion induces larger dipole in iodine molecule than in fluorine.
The ability of an atom or molecule to adjust its electron cloud in response to an outside
charge is called its polarizability.
The less tightly bound to the nucleus the electrons are, the more polarizable they are.
- as a periodic trends, polarizability increases as row increases.
Example: Arrange the following atoms in order of increasing polarizability: Ar, He, Kr, Ne
and Xe.
Answer: He, Ne, Ar, Kr, Xe
Xe – very polarizable
He – barely polarizable
3
INTERMOLECULAR FORCES
Overriding Principles
1. Opposite charges attract.
2. The greater the charge, the greater the attraction.
Types of Intermolecular Forces
1. Ion – Dipole
Example: salt water
O
H
H
O
H
H
H
O
H
H
H
O
+
-
O
H
H
O
H
H
H
O
O
H
H
H
- strongest type of intermolecular force
- as the charge increases, the strength of the force increases
Ca2+ --- H2O forces are stronger than Na+ --- H2O forces
2. Dipole – Dipole
Example: Formaldehyde
+
H
H
H

-
O
-
O
+
C
H
H
H
C
dipole – dipole
forces
C
O
H
H
C
O
Molecules with greater dipole moments have greater attractions to each other.
4
Intermolecular forces (such dipole-dipole forces) can occur between different molecules.
+
H
+
C C N
H
H
HH
C
O
3. Hydrogen bonding
- – dipole bond.
- A hydrogen bond is a very strong dipole
- Molecules with N, O, F covalently bonded to H can form hydrogen bonds.
- N, O, F atoms have high electronegativity
- Electrons in covalent bonds are greatly polarized.
O

H
+
O
H
- Hydrogen is “bare naked”.
- A hydrogen bond is formed when polarized H is attracted to nonbonding pair of
electrons.
Example: H2O
H
O
H
H
H
hydrogen bond
O
- Hydrogen bond is much stronger than dipole – dipole bond.
- Hydrogen bonding gives water many important and distinctive properties.
- Hydrogen bonding holds together DNA helix and proteins.
4. Induction Forces (Dipole – Induced Dipole)
Example: HCl in I2.
+
I
I
H
I
Cl
I
I
I
- Dipole can be induced by other dipoles.
- Induction forces are weaker than dipole – dipole forces.
I
I
5
5. Dispersion Forces (Induced dipole – induced dipole)
- also known as London dispersion forces
Charge distribution of molecules fluctuates over time.
- i. e., charge distribution is unequal in nonpolar molecules at times
- i. e., nonpolar molecules become polar for a short time, though average dipole is
zero.
Spontaneous fluctuation of electron cloud.
Consider a single bromine molecule.
+
Br
Br
Br
Br
Br
Br
Br
Br
+
-
- dipole appears, then disappears, then appears reversed, then disappears, etc…
Fluctuating dipole induces dipole in another molecule so that they fluctuate together.
Consider two bromine molecules
Br
Br
Br
Br
Br
Br
+
-
Br
+
Br
-
Strength of London dispersion forces depends on strength of induced dipole.
- Molecules with greater polarizability have greater dispersion forces.
Illustration:
Br
Br
Br
stronger bond
Br
Cl
Cl
Cl
Cl
weaker bond
Bond between bromine molecules is stronger than bond between chlorine molecules.
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General order of strength of intermolecular forces
1. Ion – dipole
2. Hydrogen bonding
3. Dipole – dipole
4. Induction
5. Dispersion
PROPERTIES OF LIQUIDS
Viscosity
- resistance to flow
- “thickness” of a fluid
- corn syrup has a higher viscosity than water
Surface Tension
- The ability of a liquid to form film on its surface that can resist puncture.
- Water and Hg are liquids with high surface tension, dish soap has low surface tension.
- Bugs walk on water because water has surface tension.
- Menicus forms in graduated cylinder because of surface tension.
Vapor Pressure
- The pressure that comes from gas molecules escaping from the surface of a liquid.
- Vapor pressure increases with temperature.
- When the vapor pressure equal atmospheric pressure, liquid is boiling.
- As intermolecular forces increase, the vapor pressure of the liquid decreases.
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THE SOLID STATE
Crystalline Solids
- Atoms in structure have a regular repeatable structure.
- Arrangement of atoms is called a lattice.
- Smallest repeatable unit is called a unit cell.
Three general types of crystalline solids
1. Ionic
- Ions are arranged in a lattice.
- Lattice is held together with ionic bonding.
2. Molecular
- Molecules are arranged in a lattice.
- Intermolecular forces such as dipole – dipole or London dispersion hold lattice together.
3. Metallic
- Metal atoms are arranged in a lattice.
- Metallic bonding holds lattice together.
- Metallic bonds are much different from ionic, covalent or other intermolecular
bonds.
- Metal ions have a fixed position within a “sea of electrons”.
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
+
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Amorphous Solids
- amorphous  without form
- Solids where atoms do not have regular structure
Two examples
1.) Network solids
- Atoms are covalently bonded to each other in a network.
- Examples
1. Diamond
2. Graphite
3. Glass (SiO2)
2.) Polymer solids
- Polymer molecules are long and “skinny”.
- Long skinny molecules knot together like a pile of cooked spaghetti.
- Intermolecular forces are also very important.
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INTERMOLECULAR FORCES AND PHASE TRANSITIONS
Recall phase transitions from Chapter 2.
Solid
condensation
Liquid
Gas
evaporation
Phase changes occur when intermolecular bonds are affected by an increase or decrease in
kinetic energy.
Latent Heats
- Each phase transition has a change of energy associated with it, this energy is the latent
heat.
- The latent heat is the energy needed to separate molecules (or push them out of the
solid lattice).
- Melting, boiling and sublimation have positive latent heats.
- Freezing, condensation and deposition have negative latent heats.
- The greater the value of the latent heat, the stronger the intermolecular forces.
- During a phase transition, the temperature remains constant.
- All energy input goes into the latent heat, not temperature increase.
Heating Curve of Water
250
Temperature (C)
200
150
100
50
63
0
60
0
57
0
54
0
51
0
48
0
45
0
42
0
39
0
36
0
33
0
30
0
27
0
24
0
21
0
18
0
15
0
90
12
0
60
30
0
0
Tim e (sec)
liquid water is
heated
water is boiling
steam is getting
hotter
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Estimating the relative values of boiling points and freezing points
Boiling point of substance is temperature where vapor pressure of liquid equals atmospheric
pressure.
The boiling point and freezing point of substances depends on
1. Intermolecular forces.
- The stronger the forces between molecules, the harder it is to separate them
- Therefore the stronger the forces, the higher the boiling points and freezing points of
substances.
2. Kinetic energy of molecules
- The greater the mass of the molecules, the greater the boiling and freezing points.
- The higher the temperature, the faster the molecules move and the more easily the
molecules can break away from each other.
Comparisons of boiling and freezing points can be made between molecules with similar
intermolecular forces.
Example: Order the following according to increasing boiling point: Br2, Cl2, F2, I2.
Recall: London dispersion forces increase as polarizability increases.
Polarizability increases as row of element increases.
Therefore answer is: F2, Cl2, Br2, I2
Example: Order the following according to increasing boiling point.

Acetaldehyde
Acetonitrile
Methyl Ether
Methyl Chloride
CH3CHO
CH3CN
CH3OCH3
CH3Cl
2.7 D
3.9 D
1.3 D
1.9 D
 - strength of dipole
The strength of the intermolecular force is proportional to the strength of the dipole.
As the value of the dipole increases, the strength of the intermolecular force increases.
Therefore answer is: CH3OCH3, CH3Cl, CH3CHO, CH3CN