Electonegativity, Polar
Bonds, and Polar Molecules
Some Definitions
• Electronegativity: the ability of an
atom to attract bonding electrons to
itself.
• Intramolecular forces: the
attractive force between atoms and
ions within a compound.
• Intermolecular forces: the
attractive force between molecules.
Polar Covalent Bonds
• When a chemical bond is formed, it
is not always exclusively ionic or
covalent.
• In the case of shared electrons
between two identical atoms, the
electrons are shared equally.
• However, this is not the case for a
compound like hydrogen chloride,
where electrons are shared between
two different elements.
• In this situation, the sharing is
unequal, as the bonding electrons
spend more time near one atom
than near the other.
• Due to this unequal sharing of
electrons, one atom will have a
slightly positive charge while the
other will have a slightly negative
charge.
• We indicate these slight charges by
+ and -. (The Greek letter delta, ,
indicates “a small difference.”)
• The bond is somewhere between an
ionic bond and a covalent bond and
is called a polar covalent bond.
• Does HCl have a polar covalent
bond? NH3?
• To determine if the bond is ionic,
covalent or polar covalent, measure
the difference in electronegativity of
the atoms.
• EN > 1.7 - ionic bond
• EN < 1.7 - covalent bond
• What type of bond do the following
compounds have:
– CsF, HF, and F2
Polar Molecules
• A molecule that is slightly positively
charged at one end and slightly
negatively charged at the other
because of electronegativity
differences is classified as a polar
molecule.
• Not all molecules containing polar
covalent bonds are polar molecules.
• Take a look at CH4 and HCl.
• polar molecule: a molecule in
which the uneven distribution of
electrons results in a positive charge
at one end and a negative charge at
the other end
• non-polar molecule: a molecule in
which the electrons are equally
distributed among the atoms,
resulting in no localized charges
Intermolecular Forces
• There are three kinds of
intermolecular forces.
• Two of these are classified as van
der Waals forces.
Dipole-Dipole Forces
• Dipole-dipole forces: an attractive
force acting between polar
molecules
• As a polar
molecule, hydrogen
fluoride has a
negatively charged
fluoride end and
a positively charged
hydrogen end.
• When two hydrogen fluoride
molecules are next to one another,
the positive end of one molecule is
attracted to the negative end of the
next molecule.
dipole-dipole force
London Dispersion Forces
• London dispersion forces: an
attractive force acting between all
molecules and unbonded atoms,
including nonpolar molecules.
Hydrogen Bonds
• Water, a polar molecule, consists of one
atom of oxygen bound by single covalent
bonds to two hydrogen atoms.
• Its structure is simple, but water exhibits
some rather unusual properties:
– higher than expected melting and boiling
points,
– high vapour pressure,
– high surface tension, and
– the ability to dissolve a large number of
substances.
• To explain these properties, we must
consider the intermolecular forces that
exist between water molecules.
• As a result of the large difference in
electronegativity between the hydrogen
and oxygen, the O-H bonds in a
molecule of water are highly polar
covalent.
• As a result, the hydrogen atoms of one
water molecule exert a strong force of
attraction on the oxygen atom of
neighbouring water molecules.
• This is known as a hydrogen bond.
• Hydrogen bonds occur among highly
polar molecules containing F-H, O-H and
N-H bonds.
• Although a hydrogen bond is similar to a
dipole-dipole force, it is stronger than any
of the van der Waals forces.
Special Properties of Water
• Unusually high Melting and Boiling
Points
• Low Density of Ice
• High Surface Tension (a phenomenon
that leads to the formation of a skin-like film on
the surface of a liquid)
• High Specific Heat Capacity (the
quantity of the energy that a certain mass of a
substance can absorb, and warm up by 1ºC)
Hydrogen Bonds in
Biochemistry
• Hydrogen bonds play a significant
role in determining the shape and
function of large, biologically
important molecules.
• Proteins consist of hundreds or even
thousands of atoms. The chain of atoms
folds into very specific three-
dimensional structures because of
attractions between different parts of the
chain.
• A DNA molecule is
made up of two long
chains of structures
called nucleotides.
• Hydrogen bonds
readily form
between the two
helixes, holding
them together like
the rungs of the
ladder