Polar Bonds and Electronegativity
When two identical atoms form a covalent bond, as in H2 or Cl2, each has an equal share of the electron pair in the bond. The electron density at both ends of the bond is the same, because the electrons are equally attracted to both nuclei.

When different kinds of atoms combine, as in HCl, the attractions usually are not equal. Generally one of the nuclei attracts the electrons more strongly than the other.

The effect of unequal attractions for the bonding electrons is an unbalanced distribution of electron density within the bond. It has been found that a chlorine atom attracts the electrons more strongly than does a hydrogen. In the HCl molecule, therefore, the electron cloud is pulled more tightly around the Cl, and that end of the molecule experiences a slight buildup of negative charge. The electron density that shifts toward the chlorine atom is removed from the hydrogen, which causes the hydrogen end to acquire a slight positive charge.
In HCl the electron transfer is incomplete. The electrons are still shared, but unequally. The charges on either end of the molecule are less than full +1 and -1 charges; they are partial charges, normally indicated by the lowercase Greek letter delta.
A bond that carries partial positive and negative charges on opposite ends is called a polar bond, or polar covalent bond. The term polar comes from the notion of poles of opposite charge at either end of the bond. Because there are two poles of charge involved, the bond is said to be a dipole.
The polar bond in HCl causes the molecules as a whole to have opposite charges on either end, so we say that HCl is a polar molecule. The HCl molecule as a whole is also a dipole.
The degree to which a covalent bond is polar depends on the relative abilities of bonded atoms to attract electrons. The term that we use to describe this relative attraction of an atom for the electrons in a bond is called the electronegativity of the atom. In HCl the chlorine is more electronegative than hydrogen. The electron pair of the covalent bond spends more of its time around the more electronegative atom, which is why that end of the bond acquires a partial negative charge.
The concept of electronegativity has been put on a quantitative basis - as numerical values have been assigned for each element. Refer to your periodic table for these assigned values. This information is useful because the difference in electronegativity values provides an estimate of the degree of polarity of a bond. In addition, the relative magnitudes of the electronegativities indicate which end of the bond carries the negative charge. For instance, fluorine is more electronegative than chlorine. Therefore, we would expect an HF molecule to be more polar than an HCl molecule. In addition, hydrogen is less electronegative than either fluorine or chlorine, so in both of these molecules the hydrogen bears the positive charge.
     ..       ..
H-F: H-Cl:
    ..       ..
+   -     +    -
The concept of electronegativity shows that there is no sharp dividing line between ionic and covalent bonding.
Ionic bonding and nonpolar covalent bonding simply represent the two extremes. Ionic bonding occurs when the difference in electronegativity between two atoms is very large; the more electronegative atom acquires essentially complete control of the bonding electrons. In a nonpolar covalent bond, there is no difference in electronegativity, so the pair of bonding electrons is shared equally.
The degree of polarity, which is the amount of ionic character of the bond varies in a continuous way with changes in the electronegativity difference.   The bond becomes more than 50% ionic when the electro-negativity difference exceeds 1.7
Within the periodic table, electronegativity varies in a more of less systematic way, and the trends follow those for the ionization energy (IE). Atoms with large ionization energies also have large electronegativities. An atom that has a small IE will give away an electron more easily than an atom with a large IE, just as an atom with a small electronegativity will lose its share of an electron pair more readily than an atom with a large electronegativity.