Electronegativity and Bond Polarity
If the bonding electrons are not shared equally in a covalent bond, then the bond will be polar. To quantify how polar a bond will be, we introduce the concept of Electronegativity.
Electronegativity - the ability of an atom in a molecule to attract electrons to itself.
For example, in the HF molecule, the bonding electrons spend more time on the fluorine than the hydrogen because the fluorine has a higher electronegativity than hydrogen. We indicate this slight excess of negative charge on the fluorine with the symbol δ-:

As a result of this unequal sharing of bonding electrons, the HF molecule will have an electric dipole moment. The electric dipole moment is a vector quantity and is represented by the symbol μ. In the case of HF, μ will lie along the direction of the H-F bond:

The strength of the electric dipole moment across a bond will be proportional to the difference in electronegativity of the two atoms forming the bond. There are a number of ways to quantify atom electronegativities. Below are a few numbers based on an approach by Linus Pauling.
H: 2.1 | |||||||
Li: 1.0 | Be: 1.5 | B: 2.0 | C: 2.5 | N: 3.0 | O:3.5 | F: 4.0 | |
Na: 0.9 | Mg: 1.2 | Al: 1.5 | Si: 1.8 | P: 2.1 | S: 2.5 | Cl: 3.0 |
The general trend of Electronegativity in the periodic table is

Homework from Chemisty, The Central Science, 10th Ed.
8.7, 8.9, 8.11, 8.29, 8.35, 8.37, 8.39