When the difference between the electronegativities of the elements in a compound is relatively large, the compound is best classified as ionic.

Example: NaCl, LiF, and SrBr₂ are good examples of ionic compounds. In each case, the electronegativity of the nonmetal is at least two units larger than that of the metal.

NaCl			LiF			SrBr2
Cl	EN = 3.16		F	EN = 3.98		Br	EN = 2.96
Na	EN = 0.93		Li	EN = 0.98		Sr	EN = 0.95
	¯¯¯¯¯¯¯¯¯			¯¯¯¯¯¯¯¯¯			¯¯¯¯¯¯¯¯¯
	EN = 2.23			EN = 3.00			EN = 2.01

We can therefore assume a net transfer of electrons from the metal to the nonmetal to form positive and negative ions and write the Lewis structures of these compounds as shown in in the figure below.

These compounds all have high melting points and boiling points, as might be expected for ionic compounds.

They also dissolve in water to give aqueous solutions that conduct electricity, as would be expected.

When the electronegativities of the elements in a compound are about the same, the atoms share electrons, and the substance is covalent.

Example: Examples of of covalent compounds include methane (CH₄), nitrogen dioxide (NO₂), and sulfur dioxide (SO₂).

CH4			NO2			SO2
C	EN = 2.55		O	EN = 3.44		O	EN = 3.44
H	EN = 2.20		N	EN = 3.04		S	EN = 2.58
	¯¯¯¯¯¯¯¯¯			¯¯¯¯¯¯¯¯¯			¯¯¯¯¯¯¯¯¯
	EN = 0.35			EN = 0.40			EN = 0.86

These compounds have relatively low melting points and boiling points, as might be expected for covalent compounds, and they are all gases at room temperature.

Inevitably, there must be compounds that fall between these extremes. For these compounds, the difference between the electronegativities of the elements is large enough to be significant, but not large enough to classify the compound as ionic. Consider water, for example.

H2O
O	EN = 3.44
H	EN = 2.20
	¯¯¯¯¯¯¯¯¯
	EN = 1.24

Water is neither purely ionic nor purely covalent. It doesn't contain positive and negative ions, as indicated by the Lewis structure on the left in the figure below. But the electrons are not shared equally, as indicated by the Lewis structure on the right in this figure. Water is best described as a polar compound. One end, or pole, of the molecule has a partial positive charge (+), and the other end has a partial negative charge (-).

As a rule, when the difference between the electronegativities of two elements is less than 1.2, we assume that the bond between atoms of these elements is covalent. When the difference is larger than 1.8, the bond is assumed to be ionic. Compounds for which the electronegativity difference is between about 1.2 and 1.8 are best described as polar, or polar covalent.

Covalent:		EN	< 1.2
Polar:		1.2 <	EN	< 1.8
Ionic:		EN	> 1.8

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