You are probably familiar with ionic bonds and covalent bonds from gcse chemistry.
Ionic bonding involves the transfer of an electron, from a metal
to a non-metal. This results in the formation of ions
with positive and negative charges. In the example shown opposite the metal
sodium reacts with the non-metal chlorine to form sodium chloride.
Here the metal sodium loses its outer shell, 3s1 electron to
chlorine. This means that the sodium atom forms a sodium ion
with a charge of 1+. The non-metal chlorine
gains an electron from the sodium and forms a chloride ion, with an
arrangement of 2,8,8 or 1s22s22p62s23p6.
Most covalent compounds involve bonding between elements with different
electronegativities. This means that the electrons
in any bonds formed will NOT be shared EQUALLY.
In an ionic compound the ions formed
have positive and negative charges
because there is complete transfer of
an electron by the metal to the non-metal. In a
polar covalent bond the electron
is shared unequally between the
two atoms, or we could simply say one atom in the bond has a much larger share of the
electron. This means that
the atoms will have partial charges and not
full positive and negative charges.
The Greek symbol delta (lowercase)
δ is used to show the
resultant partial δ+ and δ- charges that result from
the unequal sharing of the electrons in a polar
A polar covalent bond or simply a polar bond can be thought as a sort of halfway house between covalent and ionic bonding; in that the electrons in the bond are shared, but not equally as is the case in a covalent bond and the polar covalent bond contains charges, similar to those found in an ionic compound, but the charges are only partial charges and not full charges as is the case with ionic compounds. The charges on the atoms are only partial charges simply because the electrons are not completely transferred, just shared unequally.
The larger the difference in the electronegativity between the two atoms in the covalent bond the more polar (polar covalent) the bond will be. If the bond between 2 different atoms has a very large difference in electronegativity, say 1.7-2.0 or more it will be ionic. If the differences in the electronegativity is around 0.5 or less then the bond is likely to be covalent. Click on the link for bond character above or here for more information.
Consider a molecule of hydrogen chloride (H-Cl). Chlorine has an electronegativity value of 3.0 and hydrogen has a value of 2.1; this means that the bond is polarised in such a way that the chlorine atom is electron rich (δ-) and the hydrogen atom is electron deficient (δ+). Since the H-Cl bond has 2 charged ends it is described as having a bond dipole or just a dipole (a dipole is simply a small charge difference across a bond due to differences in electronegativities, which results in one end of the bond containing an atom with a partial positive charge (δ+ ) and at the other end of the bond will be an atom with a partial negative charge (δ-). The charges on both end of the bond will be the same but different in sign, one negative and one positive. So H-Cl is a non-symmetrical molecule with polar covalent bonds. We would also say that this is a polar molecule (that is one end of the molecule has a partial positive charge (δ+ ) and the other a partial negative charge (δ- ), this is shown in the image below:
Are all molecules which have polar covalent bonds polar molecules? The answer to this question is NO. To decide if a molecule is polar you not only have to decide if it contains polar covalent bonds but you MUST also consider the shape of the molecule. Bond dipoles are vector quantities and have both a size and a direction, if the molecule has polar bonds but is highly symmetrical the chances are that the molecule will be non-polar, even though it contains polar covalent bonds. As an example consider a molecule of carbon tetrachloride (CCl4) which is shown below:
You may also see the carbon tetrachloride molecule drawn out using arrows to represent the bond dipoles (see the image opposite). This is often helpful
in deciding if the molecule overall will be a polar one with a dipole moment. The arrows representing the bond dipoles give
an indication of the size of the dipole present in any particular bond. I hope you can see that all the bond dipole arrows will effectively cancel each other out
due to the shape of this molecule. However if you study the examples below you will see that this is not always the case.
Consider each of the examples below which show both polar and non-polar molecules.