Polar Covalent Bonds
- No true dividing line exists between ionic and covalent bonds (better
to rely on general observations).
- The ionic character of a bond (how ionic it is) is also determined by
ΔEN (eg. bond with ΔEN = 1.7 has 50 % ionic character)
- electrons in bonding pair spend more time closer to the more
electronegative element
- atom with higher electronegativity (EN) will attract bonding electrons
more strongly, this atom therefore has a partial negative charge(δ-),
the less electronegative element has a partial positive charge (δ+)
- The slight difference in charge within a covalent molecule is called a
dipole (two poles)
- The arrow indicates the direction of the ‘dipole moment’ or ‘partial
negative charge’ in the direction of the chlorine atom
e.g. HCl
e.g. H2O
- the negative end of the molecule tends to attract the positive end of
a molecule, resulting in higher mp and bp for that molecule
- molecules which have polar bonds can result in polar molecules,
but not always it depends on the shape of the molecule
Molecular Shapes
- polar molecule results when a molecule contains polar bonds in an
unsymmetrical arrangement. (NH3, H2O)
- a symmetrical molecule can contain polar bonds(the bond dipoles to
cancel each other) e.g CO2, CCl4
- Compounds that are made up of non-polar
molecules generally have lower melting
points and boiling points than compounds that are made up of
polar molecules.
Guidelines for predicting polar and Non-polar Molecules
(for your information)
Polar
Non
polar
Type
AB
HAx
AxOH
OxAy
NxAy
Description
- Diatomic compounds of different
elements
- Molecules with a single H
- Any molecule with an OH at one
end
- Any molecule with an O at one end
- Any molecule with an N at one end
Ax
- All molecules of a single element
- Many carbon compounds
CxAy
including organic solvents, fats and
oils)
Examples
CO
HCl, HF
C2H5OH
H2O
NF3, NH3
Cl2, N2
CO2, CH4