What are examples of dipoles?

A dipole moment refers to slight opposite charges on opposite sides of a bond. The resulting bond is said to be polar; it has a positive pole and a negative pole, much like a bar magnet.

In order to determine if a particular bond is polar or not, one must look for the electronegativity of each atom. Pauling's electronegativity is a measure of how strong a particular atom pulls electrons towards it in a bond. The value of the difference between their electronegativities (#DeltaEN#) determines how polar a bond is.

If:
#0<= DeltaEN <=0.4#, the bond is nonpolar.

#0.4<##DeltaEN##<=1.8#, the bond is polar

#DeltaEN>1.8#, the bond is ionic

Consider the bonds in #H_2O#,

There is one oxygen bonded to two hydrogens in one water molecule. Based on the difference in electronegativites for the bonds, it is clearly a polar molecule
#EN_O=3.44#
#EN_H=2.20#
#DeltaEN=3.44-2.20=1.24#

In the figure above, the #delta# symbol indicates an area of partial charge on the atom. Note that they are not full charges as in ions, but partial charges due to a difference in electron density at each "pole". The arrow in the figure indicates the direction of electron density and the slight negative charge #delta^-#and the cross indicates an area of electron deficiency and the slight positive charge #delta^+#. This difference in charges is called a dipole moment and it is a vector quantity; it has magnitude and direction.

Notice that the water molecule has an overall dipole moment that points straight up towards the oxygen. This is because a dipole moment of a molecule depends on the vector sum of the bond dipoles.

Consider #CO_2#,

As you can see, the #DeltaEN# for the #C-O# bond is within the polar range. However, since #CO_2# is a linear molecule, the dipoles point in opposite directions and the vector sum of the two is equal to zero. #CO_2# is nonpolar.