1. Intermolecular Forces
The state of matter depends upon
intermolecular forces.
Intermolecular forces are the electrostatic attractions
that molecules feel for one another.
To understand intermolecular forces,one must first
understand intramolecular forces.
Intramolecular force is interpreted using the
electronegativity of the atoms combining to form
compounds and molecules.

2. Electronegativity Electronegativity is the same property as electron
affinity, except it is for atoms in compounds or
molecules.
Electronegativity is the desire of atoms to draw
electrons when they combine with other atoms.
The larger the atomic radius, the harder it becomes to
attract electrons to an atom.
The larger the atomic radius, the lower the
electronegativity of an atom.

3. Electronegativity
Atomic Radii Decrease
Electronegativity Increase
Electronegativity Decrease
Atomic Radii Increase

4. Intramolecular Forces – A Review
Intramolecular forces are the forces that bind atoms
together resulting in molecular bonding.
Intramolecular forces, along with molecular geometry
will be used to interpret intermolecular forces.
Intramolecular forces fall into one of three categories:
1. Ionic bonding
2. Covalent bonding
3. Polar covalent bonding

5. Ionic Bonding Ionic bonds are formed when the combining atoms
have large differences in electronegativity.
Ionic bonds are formed when the combining atoms
are separated by large distances in the periodic table.
Elements from the left lose electrons (becoming
cations) and elements from the right gain electrons
(becoming anions).
The electrostatically opposites attract to form an
ionic compound.

6. The Covalent Bond Covalent bonds are formed by combining atoms with
exactly the same electronegativities.
Because only elements of the same type have exactly
the same electronegativity, covalent bonds are
formed by elements combining with themselves.
Because each element in the bond has exactly the same
electrongegativity, covalent bonds are non-polar.
Covalently bound materials exist as individual
entities called molecules. They do not repeat in a
lattice like ionic compounds.

7. The Polar Covalent Bond
The polar covalent bond is an intermediate form of
bonding.
Polar covalent bonds are formed between elements
with similar electronegativities.
Polar covalent bonds are formed between elements
that are nearby each other in the periodic table.
Ionic bonds are formed between elements with large
differences in electronegativity (far apart in the table).
Covalent bond are formed between elements with the
same electronegativity (combining with themselves).
Polar covalent bonds exhibit some degree of electron
sharing with some degree of polarity.

8. Classifying the Type of Bonding
N KBr PCl LiNO H SO2
polar
covalent
covalent
ionic
ionic
covalent
polar
covalent

9. The Bond Dipole In polar covalent bonds, electrons are not shared
evenly between the two atoms.
The more electronegative element in the bond
acquires more of the electron density than the
less electronegative element.
This leads to an imbalance in charge distribution
called a bond dipole.

10. F Cl N O Si S C H The Bond Dipole
Bond dipoles in a molecule are indicated by
drawing an arrow pointing from the less
electronegative element towards the more
electronegative element.
F Cl
N O
Si S
C H
Atoms bonded to themselves are covalent bonds with
no polarity and no bond dipole!

11. The Dipole Moment In polyatomic polar covalent molecules, there
are multiple bond dipoles.
When these bond dipoles are in opposition,
they cancel leading to a net non-polar molecule.
When these bond dipoles are not in opposition,
or are in opposition but are not of equal strength,
it leads to a net polarity called a dipole moment.

12. Non-Polar Molecules 120o 180o 109.5o Trigonal planar molecule
Linear molecule
120o
180o
Tetrahedral molecule
Octahedral Molecule
90o
109.5o

13. Polar Molecule Examples- +
O is more electronegative than S - +
Total dipole moment is directed
tail at more positive head at more negative
No bond dipoles along
S-S bonds + - - +

14. Total Dipole Examples d- d+
non-polar d- d+ d- d+

15. Intermolecular Force in Non-Polar Molecules
Molecules with symmetric charge distribution have
very little attraction for one another.
This is the reason that most of them are gaseous at
room temperature.
However, they must possess some mechanism for
attraction, otherwise they would never condense.
Their mechanism of attraction is called
Dispersion Forces.

16. Dispersion Forces Non-polar molecules have symmetric charge
distribution.
Nuclei
Electronic charge distributed
evenly over the nuclear framework.
The net result is a non-polar molecule with little
mechanism for intermolecular force

17. Dispersion Forces
In the influence of an electric field, this symmetric
charge distribution can be induced to distort.
Temporary dipole moment + -
Positive charge
(nucleus) orients
toward negative pole.
Electron density
piles up near
positive pole.

18. Systems with Dispersion Force Interactions
Group VIIIA Elements (Noble Gases)
He Ne Ar Kr Xe Rn
Homonuclear diatomics
O O
Other homonuclear diatomics: H2 N2 F2 Cl2 Br2 I2
Non-polar
polyatomics

19. Dispersion Force and Size
In symmetric charge distributions, dispersion forces
are the only mechanism of attraction.
The charge distributed around larger systems feels
less electrostatic attraction, and is more polarizable.
Dispersion forces increase with the size of system
Increasing Strength of Dispersion forces
He < Ne < Ar < Kr < Xe < Rn
F2 < Cl2 < Br2 < I2

20. The Hydrogen Bond Because of their relatively small size and
extremely high electronegativity, the elements
F, O, and N bond to H and form an extremely
polar bond.
The resulting polarity creates unusually strong
intermolecular force between molecules.
This intermolecular force is called the
Hydrogen Bond.

21. Water is a “really bent”
Hydrogen Bonding in Water
Water is a “really bent”
molecule
Lone pairs + -

23. Hydrogen Bonding and Intramolecular Force

24. Hydrogen Bonding in the Alpha - Helix
The alpha-helix is a corkscrew strucutre held in place by
hydrogen bonding between N-H groups and C=O groups in
the next turn of the helix four amino acids away. This
creates a “telephone cord” affect.

25. Hydrogen Bonding in the Beta-Pleated Sheet
In a beta-pleated sheet, strands of polypeptide chains are held
together side by side by hydrogen bonds between the
peptide chains.

26. Hydrogen Bonding in the Double Helix
A double helix is formed by intertwining
strands of polypeptide chains that hydrogen bond.
DNA is a double helix

27. Hydrogen Bonding in Starch
Hydrogen bonding between
glucose monomers makes
the conformation of the amylose
polymer coil into a helical
structure known as starch.

28. Hydrogen Bonding and the Triple Helix
Collagen makes up about 1/3
of all protein in vertibrates
Collagen is a triple helix that
forms the connective tissue
(ligands and tendons, etc)
in animals.

29. Hydrogen Bonding and Tertiary Structure
The tertiary structure of a protein involves attractions and
repulsions between the side chain groups of the amino acids
in the polypeptide chain.
This creates twists and
bends in the in the protein
until it acquires a specific
three-dimensional shape
that determines its
biological function (or
structure-activity
relationship).

30. Summary of Intermolecular Forces
There are various types of distinct intermolecular force
which molecules can employ:
1. Ionic Forces - Forces exerted by cations and anions
in ionic compounds
2. Hydrogen bonding – Exerted by polar molecules
which contain H bonded to F, O, or N.
3. Dipole forces – Exerted by polar covalent molecules
with unsymmetric charge distributions such that they
possess a permanent dipole moment.
4. Dispersion forces – Exerted by non-polar molecules
when they acquire temporary dipoles.

31. Hierarchy of Intermolecular Forces
Strongest Weakest
Ionic
Forces
Hydrogen
Bonding
>
Dipole
Forces
>
Dispersion
Forces
>

32. Summary of Intermolecular Forces
There are various types of distinct intermolecular force
which molecules can employ:
1. Ionic Forces - Forces exerted by cations and anions
in ionic compounds
2. Hydrogen bonding – Exerted by polar molecules
which contain H bonded to F, O, or N.
3. Dipole forces – Exerted by polar covalent molecules
with unsymmetric charge distributions such that they
possess a permanent dipole moment.
4. Dispersion forces – Exerted by non-polar molecules
when they acquire temporary dipoles.

33. Hierarchy of Intermolecular Forces
Strongest Weakest
Ionic
Forces
Hydrogen
Bonding
>
Dipole
Forces
>
Dispersion
Forces
>

34. Classifying Intermolecular Forces
The interaction in the ionic lattice is
characterized as ion-ion interaction.
Na+
Cl-
The interaction in pure water is best
characterized as hydrogen bonding
interaction.
The interaction between water molecules
and ions in solution is best characterized
as ion - dipole interaction.
Na+

35. Classifying Intermolecular Forces
Ammonia - Water
Solution
Hydrogen Bonding
PCl3 - Water
Solution
Dipole-Dipole
Interaction
Gas mixture (such as air) - Dispersion Forces

36. Consequences of Polarity
The more polar a molecule the more intermolecular
forces exerted among a bulk sample of molecules.
Physical properties of bulk sample of these molecules
are affected by their degree of attraction to one another.
These properties include:
Bulk Phase
Vapor Pressure
Boiling Point
Melting Point
Solubility

37. Intermolecular Force and Bulk Phase
Because of their highly polar bonds, ionic compounds
are solids at ambient temperatures.
CaCl2
Na2S
MgSO4
LiNO3
K3PO4

38. Intermolecular Force and Bulk Phase
Because of their non-polar charge distributions the
covalent diatomics only exert dispersion forces and
are mostly gases at ambient temperatures.
H2 (g)
N2 (g)
O2 (g)
F2 (g)
Cl2 (g)
Br2 (l)
Larger size, higher dispersion forces.
Even larger size, even more dispersion
Forces.
I2 (s)

39. Intermolecular Force and Bulk Phase
In polar covalent molecules, the geometry and dipole
moment play a role in the polarity and phase of matter.
Of course, temperature and pressure play significant
roles in the bulk phase of matter as well.
At atmospheic pressure, water becomes a solid at
temperatures of 0 oC and below.
At pressures over the atmospheric value, the freezing
point of water is lowered below 0 oC.
At atmospheic pressure, water exists as a vapor at
temperatures of 100 oC and above.
At pressures over the atmospheric value, the boiling
point of water is raised above 100 oC.

40. Bulk Phase of Polar Covalent Molecules
Carbon Dioxide is a gas
at room temperature d-
Water is a liquid
at room temperature d+
Sucrose is a solid
at room temperature

41. Vapor Pressure Molecules that are loosely held in the liquid or solid
phase can break free to the gaseous phase and exert a
vapor pressure.
The lower the intermolecular force between molecules
in a sample, the more escapable its molecules become.
The lower the intermolecular force between molecules
in a sample, the higher its vapor pressure.

42. Vapor Pressure H2O(l) H2O(g) H2O(g)  H2O(l) (Evaporation)
Without a lid, water
vapor can escape
A lid prohibits water
vapor from escaping
H2O(l) H2O(g)
(Evaporation)
H2O(g)  H2O(l)
(Condensation)
When the rate of
evaporation matches
the rate of condensation
we reach equilibrium: H2O(l) H2O(g)

43. Vapor Pressure and Intermolecular Force
Because of the high polarity of their bonds, ionic
compounds have high intermolecular force and
therefore exhibit no vapor pressure.
Nonvolatile substance- A material with no vapor
pressure.

44. Vapor Pressure and Intermolecular Force
Because of the non-polar nature of their bonds,
covalent molecules have low intermolecular force and
therefore exhibit high vapor pressure.
Although Br2 is a liquid and I2 a solid at room
temperature, they still exhibit very high vapor
pressures.
Volatility - measures the vapor pressure exerted by a
substance.

45. Vapor Pressure and Intermolecular Force
Because of the slight polarity of their bonds, polar
covalent molecules may or may not exhibit strong
intermolecular forces, depending on their geometry.

46. Vapor Pressure of Polar Covalents
Isomers - molecules with the same chemical
formula, but different structures.
Polar d- d+
Non-polar
Lower Vapor Pressure Higher Vapor Pressure

47. Temperature and Vapor Pressure
As temperature increases, molecules move more
rapidly, so that more can escape the liquid phase.
As temperature increases, vapor pressure increases.
Vapor Pressure of H2O at Various Temperatures
0 o C mm Hg
60 o C mm Hg
20 o C mm Hg
80 o C mm Hg
40 o C mm Hg
100 o C mm Hg

48. The Boiling Point The boiling point of a substance is the
temperature where is vapor pressure matches
external pressure.
The normal boiling point of a substance is
when its vapor pressure matches sea level
atmospheric pressure (1 atm or 760 mm Hg).
Substances at high altitude boil at lower
temperatures.
Substances under high pressures boil at higher
temperatures.

49. Boiling Point and Intermolecular Force
The higher the intermolecular force between
molecules in a bulk sample, the more
“stickability” of their molecules
The higher the intermolecular force between
molecules in a bulk sample, the higher its
boiling point.

50. Boiling Points of the Noble Gas Elements
Lowest Highest
He < Ne < Ar < Kr < Xe < Rn
Source of intermolecular forces – Dispersion
forces.
Relative size of Group VIIIA elements:
Smallest Largest
He < Ne < Ar < Kr < Xe < Rn

52. Boiling Points of Hydrides
Lowest
Highest
Group IVA: CH4 < SiH4 < GeH4 < SnH4
Group VA: PH3 < AsH3 < SbH3 < NH3
Hydrogen bonding – higher intermolecular force
Group VIA: H2S < H2Se < H2Te < H2O
Group VIIA: HCl < HBr < HI < HF

53. Boiling Points of Polar Covalent Isomersd- d+
Polar Non-Polar
Higher
boiling point
Lower
boiling point

54. Melting Points X(s)  X(l)
Melting point is the change of phase from solid
to liquid – Scientists call this fusion.
Melting point trends follow the same pattern
as boiling point trends.
Materials with high intermolecular force will
exert strong attractions to one another, and
therefore melt at higher temperatures.

55. Solubility of Polar Covalent Isomers
Polar – Higher
solubility in water d- d+
Non-polar – Lower
solubility in water

56. Solubility Examples Ammonia (NH3) d- d+ Soluble in water,
ethanol, acetic acid
Soluble in toluene,
benzene, CCl4
Insoluble in toluene,
benzene, CCl4
Insoluble in water,
ethanol, acetic acid

57. Soaps and Detergents Polar region
Fats and oils are generally non-polar materials
that will not dissolve in water.
Soaps and Detergents are designed to make
these substances water soluble.
Sodium Lauryl Sulfate
Non-polar region
dissolve in grease
Polar region
dissolves in water

58. Detergent Action Polar end of detergent points to polar water
molecules
Non-Polar
end of
detergent
points toward non-polar grease
Polar
Water