Besides the difference between σ and π bonds, there is one more factor important in chemical bonding, which is how equally the electrons are shared between the pair of atoms. This is the distinction between ionic and covalent bonding. The difference between the type of bonding is normally based on the differences in electronegativity (χ) of the atoms. There are several different ways to define the numerical value of χ, but it basically corresponds to the ability of an atom to attract bonding electrons to itself. A partial Periodic Table with the Pauling electronegativities is shown below. The general trend is that F has the highest electronegativity, and the numeric values decrease either down the group or across the period (to the left).

When two atoms are bonded together, the more electronegative atom will "acquire" the electrons more than 50% of the time. For example, if O and H are bonded together, the more electronegative O atom will have the two bonding electrons closer to it, on average, than the H atom. This results in a permanent partial negative charge on the O atom (O^δ-) and a corresponding partial positive charge on the H atom (H^δ+). This "asymmetric" charge distribution will produce a bond (dipole) moment. The vector sum of the bond moments of the molecule is the molecular dipole moment. The unsymmetric charge distribution is also important in explaining many physical and chemical properties of molecules such as boiling points, melting points, acidity, and chemical reactivity to name a few.

The difference in electronegativity (Δχ) of the two atoms determines the type of chemical bond:

•    0 ≤ Δχ < 0.4 ⇒ covalent bonding
0.4 ≤ Δχ < 1.7 ⇒ polar covalent bonding
• 1.7 ≤ Δχ          ⇒ ionic bonding

H 2.20

He

Li 0.98

Be 1.57

B 2.04

C 2.55

N 3.04

O 3.44

F 3.98

Ne

Na 0.93

Mg 1.31

Al 1.61

Si 1.90

P 2.19

S 2.58

Cl 3.16

Ar

K 0.82

Ca 1.00

Sc 1.36

Ti 1.54

V 1.63

Cr 1.66

Mn 1.55

Fe 1.83

Co 1.88

Ni 1.91

Cu 1.90

Zn 1.65

Ga 1.81

Ge 2.01

As 2.18

Se 2.55

Br 2.96

Kr

Rb 0.82

Sr 0.95

Y 1.22

Zr 1.33

Nb 1.60

Mo 2.16

Tc 1.90

Ru 2.20

Rh 2.28

Pd 2.20

Ag 1.93

Cd 1.63

In 1.78

Sn 1.96

Sb 2.05

Te 2.10

I 2.66

Xe

Pauling Electronegativity Scale

The differences in electronegativity leading to polar covalent bonds is extremely import in chemistry. The different electronegativities results in a permanent asymmetric charge distribution in the molecule which result in a permanent dipole moment. This can be seen in the following images.

The Jmol images are the Electrostatic Potential Energy (ESP) mapped onto the Electron Density surface. The electron density surface is the boundary within which is contained 95% of the electron density of the molecule. It basically corresponds to the Van der Waals surface of the molecule. The ESP is color coded onto this surface. The coloring indicates areas of high electron density (red) - medium (green) - to low electron density (blue).

HF Δχ = 1.78	HCl Δχ = 0.96	HBr Δχ = 0.76	HI Δχ = 0.46

Notice how the trend Δχ matches the trend in electron density on the halide atom (indicated by the red to green shift in color).