: Intermolecular Forces of Attraction
Intermolecular Forces of Attraction - Forces of
attraction that exist between molecules
Network (Covalent) Bonds - Strongest
In a network solid, atoms are
bonded in a lattice structure.
Network solids are very hard and their strong bonds result in very high melting
and boiling point temperatures.
Network solids have localized electrons which are in fixed positions
in the covalent bonds. This makes network solids poor conductors of
Examples of network solids include
Diamond - a form of carbon in which its electrons have an sp3
hybridization and the atoms are arranged in a tetrahedral network.
Quartz - SiO2
Silicon carbide - SiC
Ionic Bonds (between ionic molecules)
- Medium strong bond
In an ionic solid, adjacent ions have electrostatic attractions and
are arranged in a lattice structure.
Ionic solids have strong bonds resulting in high melting and boiling point
Ionic solids have localized electrons which are in fixed positions
around the atoms. This makes ionic solids poor conductors of
Ionic liquids however (e.g. molten NaCl) do conduct electricity because the
electrons are free to move about.
Coulomb's Law states that the force of attraction between two objects is
equal to the product of their charges (q1 and q2)
divided by the square of the distance between them (r):
Therefore, the charges on the ions in an ionic solid are directly proportional
to the strength of the ionic bonds.
The sizes of the ions in an ionic solid are inversely proportional to the
strength of the ionic bonds
e.g. The ionic bonds in MgO (charges +2 and -2) are stronger than those
in NaF (charges +1 and -1)
e.g. The ionic bonds in LiF are stronger than those in KBr (because
LiF is much smaller than KBr)
Metallic Bonds (bonds between metal
atoms) - Weakest strong bond
Metallic solids are often described as "a group of nuclei surrounded by a
sea of mobile electrons." In other words, the electrons
in metallic solids are delocalized and are free to move about, making
metals good conductors of electricity.
Metallic bonds are still relatively strong forces of attraction and most
metals have high melting and boiling point temperatures.
Mercury is the only metal that is not solid at room temperature.
Though most metals are very hard, the unrigid structure of the electrons
Like in ionic solids, Coulomb's Law can be used to describe the relative
strengths of metallic solids:
The size of the metal atoms is inversely proportional to the strength of
the metallic bonds. In other words, the smaller the atoms, the stronger
the force. Smaller size allows for the positively charged nucleus of
one metal atom to be closer to the negatively charged electrons of another,
increasing the strength of the attraction between them.
Hydrogen Bonding - Strongest weak
In hydrogen bonds, the positively charged hydrogen end of one molecule is
attracted to the negatively charged end of another molecule which must be
an extremely electronegative element (fluorine, oxygen, or nitrogen -
e.g. H2O, HF, and NH3.
also form hydrogen bonds with water in which the O in the OH group of the
alcohol bonds to the positively charged H end of the water molecule and the
H in the OH group of the alcohol bonds to the negatively charged O of the
Hydrogen bonds are the strongest weak bond because the H atom essentially
gives its single electron to form a bond and is therefore left unshielded.
The relatively strength of hydrogen bonds results in higher melting
and boiling point temperatures than those in molecules with other van der
Waals forces of attraction.
Hydrogen bonding can explain why water is less dense in the solid phase than
it is in the liquid phase (contrary to most other substances). The
hydrogen bonds between water molecules in ice to form a crystal structure,
keeping them further apart than they are in the liquid phase.
Dipole-Dipole Forces - Medium weak
Dipole-dipole forces exist between neutral, polar molecules
where the positive end of one molecule is attracted to the negative end of
The greater the polarity (difference in electronegativity of the atoms in
the molecule), the stronger the dipole-dipole attraction.
Dipole-dipole attractions are very weak and substances held together by these
forces have low melting and boiling point temperatures. Generally,
substances held together by dipole-dipole attractions are gases at room
London Dispersion Forces - Weakest
London dispersion forces (LDF) occur between neutral, nonpolar
molecules. LDF occur due to the "random motion of
electrons." At any moment, one atom may be surrounded by an
extra electron from a neighboring atom, resulting in an instantaneous polarity
on the atom. During that instant, the "polarized" atom will act as
a very weak dipole.
Since LDF are dependent upon the random motion of electrons, the more electrons
an atom or molecule has, the greater the LDF between them.
LDF as a whole are extremely weak, so substances held together by these forces
have extremely low melting and boiling point temperatures. These substances
tend to be gases at room temperature.
Next: "Gases: Properties of