Electronegativity summarizes the tendency of an element to gain, lose, or share electrons when it combines with another element. But there are limits to the success with which it can be applied. BF₃ (EN = 1.94) and SiF₄ (EN = 2.08), for example, have electronegativity differences that lead us to expect these compounds to behave as if they were ionic, but both compounds are covalent. They are both gases at room temperature, and their boiling points are -99.9^oC and -86^oC, respectively.

The source of this problem is that each element is assigned only one electronegativity value, which is used for all of its compounds. But fluorine is less electronegative when it bonds to semimetals (such as B or Si) or nonmetals (such as C) than when it bonds to metals (such as Na or Mg).

This problem surfaces once again when we look at elements that form compounds in more than one oxidation state. TiCl₂ and MnO, for example, have many of the properties of ionic compounds. They are both solids at room temperature, and they have very high melting points, as expected for ionic compounds.

TiCl₄ and Mn₂O₇, on the other hand, are both liquids at room temperature, with melting points below 0^oC and relatively low boiling points, as might be expected for covalent compounds.

The principal difference between these compounds is the oxidation state of the metal. As the oxidation state of an atom becomes larger, so does its ability to draw electrons in a bond toward itself. In other words, titanium atoms in a +4 oxidation state and manganese atoms in a +7 oxidation state are more electronegative than titanium and manganese atoms in an oxidation state of +2.

As the oxidation state of the metal becomes larger, the difference between the electronegativities of the metal and the nonmetal with which it combines decreases. The bonds in the compounds these elements form therefore become less ionic (or more covalent).

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