The physical properties of molecular substances result from
different types of forces between their molecules.
 Intermolecular forces
 Temporary/instantaneous dipole
 Induced dipole
 London (dispersion) forces
 Permanent dipole
 Dipole-dipole interaction
 Van der Waal’s forces
 Hydrogen bonding
 INTERmolecular forces: forces between molecules
 Forces where whole molecules interact with each other
 INTRAmolecular forces: forces within molecules
 Forces that hold a single molecule together
 So what are examples of intramolecular forces?
 Examples of intermolecular forces?
 Just as the strength of the intramolecular forces can
determine properties such as melting and boiling
points, so can intermolecular forces
 The increase in intermolecular forces will
 Increase m.p. and b.p. (less volatile)
 Can determine the Solubility (like dissolves like)
 Conductivity (moving charges = conductivity)
 Non-polar molecules can have:
 Weak dipole = temporary/instantaneous dipole
 This can influence a neighbor molecule = induced
dipole
 LDF strength increases as molecular size increases
 More e- in atom = more chances for temporary dipoles
occurring
 LDF are ONLY forces between non-polar molecules
 But exist in all molecules!!!
 Generally have low m.p. and b.p. = Easy to break weak
LDFs
 Polar covalent molecules have a permanent dipole
 These permanent charges orient other molecules the
same way = dipole-dipole attraction
 Strength of the attraction depends on distance &
orientation of dipoles
similar molecular mass
 Note: A polar molecule can induce a dipole in a non-
polar molecule = dipole-induced dipole
 Van der Waal forces: a term that includes both London
dispersion forces and the dipole interactions (both
dipole-dipole and dipole-induced dipole)
 Note: this term does not include H-bonding!
 Is a special case of dipole-dipole attraction
 Hydrogen bonded to:
 Nitrogen
 Fluorine
 Oxygen
 These are the strongest form of intermolecular attraction
(so b.p. and m.p. are much higher than predicted)
 Can you explain why water is less dense in solid than
in liquid form?
 H-bonds are the strongest intermolecular force, but is
still weaker than covalent and ionic bonds
 What happens when CH3CH2OH is boiled?
 H2O and CO2 are released as gases
 CH3CH2OH is released as gas
 CH3CH3 and O2 are released as gases
 H2 and CO2 are released as gases
 So are intermolecular forces broken or intramolecular
forces broken when something is boiled?
 Covalent cmpds have lower values than ionic
 Weak intermolecular forces vs electrostatic attraction
 Higher intermolecular bond strength, polar strength,
and increase in molecular size increases m.p. and b.p.
 “Like dissolves like”
 Non-polar dissolve in non-polar (such as oil)
 Due to the formation of LDF
 Polar covalent molecules are usually soluble in water
and other polar solvents
 Dipole interactions and H-bonding
 The larger the molecule with non-polar areas, the less
polar the overall molecule
 Giant covalent are insoluble in nearly all solvents
 Too much E needed to break the bonds for interaction
 Covalent compounds do not conduct as there are no charges
particles to carry a charge
 However, some polar covalent will conduct when in aqueous
solution
 Giant covalent molecules
 Diamond = no conduction
 Fullerene & silicon = semiconductors
 Graphite & graphene = conductors
Metallic bonds involve a lattice of cations with
delocalized electrons.
 Metals have:
 Few valence electrons (typically)
 Low ionization energy
 Form positive ions by losing electrons when reacted
 When in elemental state, who can accept the e- s?
 These electrons
tend to “wander
off” or be
delocalized
 These e- are no
longer associated
with 1 atom!
 The strength is determined by:
 Number of delocalized e- (more e-, stronger bond)
 Charge on cation (higher charge, stronger bond)
 Radius of the cation (smaller radius, stronger bond)
 Write the full e- configuration for Na and Mg:
 Compare the above points and determine which will
have stronger bonding. Why?
 Melting point of Na = 98 °C
 Melting point of Mg = 650 °C
 Which atom will have a higher melting point:
 Na, K or Rb?
 How did you determine this?
 Melting point of Na = 98 °C
 Melting point of K = 63 °C
 Melting point of Rb = 39 °C
 Transition metals: very strong metallic bonds. Why?
 Many delocalized e- from 3d and 4s sub-shells
 Alloys:
 Solutions of metals
 Enhanced properties
 Typically solid but made when molten
 Metallic bonds can accommodate other cations of
different sizes into lattice
 Bound by the delocalized e- in metal (thus metallically
bonded)
 Have different properties than original components
 More chemically stable
 Often stronger
 Often more resistant to corrosion