This document discusses different types of intermolecular and intramolecular forces. It begins by defining intramolecular forces as those within a molecule, such as covalent, ionic, and metallic bonds. Intermolecular forces are defined as those between different molecules, including ion-dipole, hydrogen bonding, dipole-dipole, and dispersion forces. The document then discusses how these different intermolecular forces relate to physical properties like boiling points, solubility, surface tension, and phase changes. It emphasizes that stronger intermolecular forces lead to higher boiling points, freezing points, and surface tension as well as differences in solubility between polar and nonpolar substances.

1. Intermolecular forces
1

2. Intramolecular forces among
atoms or ions within a ‘molecule’
• covalent forces
• ionic forces
• metallic bonding
• extended covalent bonding
(graphite, diamond, graphene)
• coordinate covalent (transition metal complexes;
Lewis acid-base)
2

3. examples of phenomena that depend on intermolecular forces
now what about Intermolecular forces among differing
molecules
• physical states (phases) and phase changes
( solid  liquid  gas )
secondary and tertiary structure of biologically important
molecules
(how differing parts of a large molecule interact to form its
full 3-D structure)

4. physical states and intermolecular forces
Intermolecular forces
weak
moderate
strong
4

5. types of intramolecular (bonding) and intermolecular force
• intramolecular
 ionic
 covalent
 metallic
 coordinate covalent (transition metal
complexes; Lewis acid-base)
• intermolecular
 ion-dipole
 hydrogen bonding
 dipole-dipole
 ion-induced dipole
 dipole-induced dipole
 dispersion (London)

6. energies of intramolecular (bonding) ‘forces’
6

7. Ion-Dipole Interactions
• ion-dipole interactions are an important force in
solutions of ions.
• The strength of these forces are what makes it
possible for ionic substances to dissolve in polar
solvents.
ion-dipole intermolecular forces

8. ion-dipole intermolecular forces: ion (polar) ↔ polar
- H +
Na+ O
H +
7

9. Intermolecular Forces
Dipole-Dipole Forces
Attractive forces between polar molecules
Orientation of Polar Molecules in a Solid
11.2

10. Cl ─ I
I ─ Cl I ─ Cl
dipole-dipole intermolecular forces: polar ↔ polar
H H
H C O kJ/mol
H
− +
+ − + −
8

11. Dipole-Dipole Interactions
The more polar the molecule, the higher is its
boiling point.

12. ordering by dipole-dipole forces
Lower T Higher T 9

13. Intermolecular Forces
Dispersion Forces
Attractive forces that arise as a result of temporary
dipoles induced in atoms or molecules
11.2
ion-induced dipole interaction
dipole-induced dipole interaction

14. ion-induced_dipole and dipole-induced_dipole (polar ↔ nonpolar)
isolated
He
kJ/mol
kJ/mol
10

15. dipole
more dipole – induced dipole
nonpolar
dipole induced
11

16. dispersion forces (instantaneous dipoles): (non-polar ↔ non-polar)
kJ/mol
12

18. dispersion forces (instantaneous dipoles; figure 16.5)
13

19. 11.2
Polarizability is 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
Factors Affecting London Forces

20. Factors Affecting London Forces
• The shape of the molecule affects
the strength of dispersion forces:
long, skinny molecules (like n-
pentane tend to have stronger
dispersion forces than short, fat ones
(like neopentane).
• This is due to the increased surface
area in n-pentane.

21. boiling points and intermolecular forces
greater molecular surface  greater dispersion forces 
higher boiling points
24

22. Intermolecular Forces
Hydrogen Bond
11.2
The hydrogen bond is a special dipole-dipole interaction
between they hydrogen atom in a polar N-H, O-H, or F-H bond
and an electronegative O, N, or F atom.
A H…B A H…A
or
A & B are N, O, or F

23. Hydrogen bonding, interaction involving
a hydrogen atom located between a pair of other
atoms having a high affinity for electrons; such
a bond is weaker than an ionic bond or
covalent bond but stronger than van der Waals
forces
A hydrogen bond is a primarily electrostatic force of
attraction between a hydrogen (H) atom which is
covalently bound to a more electronegative atom or
group, particularly the second-row elements nitrogen
(N), oxygen (O), or fluorine (F)

24. Hydrogen Bond
11.2

25. Why is the hydrogen bond considered a
“special” dipole-dipole interaction?
Decreasing molar mass
Decreasing boiling point
11.2

26. small electronegative atom hydrogen bonded to
─B: ······ H─A─
H ─ F..: H ─ O.
. O..= H─N─
─ or :N≡
kJ/mol
hydrogen bonds (very important !!)
- + -
with lone pair (N, O, F) electronegative atom
N,O,F
H-bond
smal
.l
.electronegative
.a
.toms:
.. ..
H2O
15

27. hydrogen bonds in biological molecules (RNA and DNA)
17

29. hydrogen bonds in biological molecules (protein secondary structure)
18

30. summary (Silberberg: table 12.2)
strong
moderate
weaker
depends stay tuned
19

32. polarizability: strength of induced and spontaneous dipoles
polarizability: how “free” the electrons in an atom or
molecule are to ‘slosh around’
induced and spontaneous dipoles are larger if atom or
molecule is more polarizable
• periodic trends in polarizability:
 increases down a group (outer electrons further away)
 decreases across a period (higher Zeff, more tightly held)
 anions are more polarizable than parent neutral atom (lower Zeff)
 cations are less polarizable than parent atom (higher Zeff)
22

33. increased polarizability
increased freezing point
boiling point in Kº
23
boiling points, melting points, vapor pressure and intermolecular forces
(nonpolar compounds)
greater polarizability 
greater intermolecular forces 
higher melting (freezing) and boiling points, lower vapor pressure
melting point, strength of
intermolecular forces
] ~
a. highest boiling point
HBr, Kr, or Cl2 HBr > [Cl2>Kr]
c. lowest vapor pressure at 25ºC
Cl2, Br2, or I2 I2 < Br2 < Cl2

34. boiling points and intermolecular forces
molecules with equivalent “molecular weight”
(ie ‘size’ and polarizability and intermolecular dispersion forces)
polarity (dipole moment)
and boiling point 28

35. surface tension
intermolecular forces differ for molecules at surface and in bulk
extra: molecules at surface have higher
energy than those in ‘bulk’; liquids
form spherical droplets to minimize
surface area
29

36. surface tension
greater intermolecular forces  greater surface tension
IMF
30

37. concave vs convex meniscus
H2O greater forces with glass than H2O  concave and high capillarity
Hg greater forces with Hg than glass  convex
31

38. giving ice a lower density than H2O liquid. ICE FLOATS!!
why does ice float
• H2O is polar and can form hydrogen bonds
• High surface tension and capillarity
• Hydrogen bonds form very open structure in solid H2O (ice)
32

39. solubility and intermolecular forces
NaCl(s) → Na+(aq) + Cl- (aq)
C2H5OH + H2O → C2H5OH (aq)
C6H14 + H2O → C6H14 + H2O → C6H14 (aq)
C6H14 + CCl4 → solution
34

40. solubility and intermolecular forces
whether a substance dissolves in ‘solvent’ (solubility),
or two liquids mix (miscibility) is determined by two
factors:
• things like to get ‘mixed up’, Solutions Happen
unless too endothermic (entropy)
• things like to give off heat (stability of
‘products’, interparticle forces in products vs
those in reactants
35

41. solubility and intermolecular forces (ionic solids + polar solvent)
NaCl(s) → Na+(aq) + Cl- (aq)
[ion-ion] [ion-dipole]
36

42. solubility and intermolecular forces (two polar liquids)

43. solubility and intermolecular forces (nonpolar + polar)
C6H14 + H2O → C6H14 (aq)
hexane
H H H H H
H─C─C─C─C─C─C─H
H H H H H
only weak dispersion and dipole-induced dipole forces
among hexane and water molecules immiscible
38

44. nonploar molecules: hydrophobic
39

45. solubility and intermolecular forces (nonpolar + nonpolar)
C6H14 + CCl4 → solution
does dissolve
40

46. solubility and intermolecular forces
‘ in general’ (likes dissolve in likes)
 polar molecules will form solutions with polar
molecules
 nonpolar molecules will form solutions with nonpolar
molecules
 polar and nonpolar substances will not form solutions