How to quickly distinguish between covalent bonds and ionic bonds?

Most metal elements form ionic bonds with nonmetal elements, while nonmetal elements form covalent bonds with each other. Based on this principle, the electronegativity difference can also be used to roughly determine the nature of a bond.

Why can electronegativity be used to determine covalent bonds? Electronegativity refers to an atom's ability to attract electrons in a chemical bond. The difference in electronegativity describes the type of chemical bond, as the relative strength of attraction for electrons determines the bond type. The stronger the attraction, the more electrons are drawn toward that atom, making it partially negative, while the other atom becomes partially positive.

Here is an empirical criterion:

1. Electronegativity difference > 1.7: The bond exhibits ionic bond characteristics. This can be understood as a significant difference in electron-attracting ability, where one atom completely gains electrons while the other completely loses them (one side "completely wins," and the other "completely loses").

2. 0.4 < Electronegativity difference < 1.7: The bond exhibits polar covalent bond characteristics. This means there is a difference in electron-attracting ability, but it is not large enough for complete electron transfer. Instead, electrons are more inclined toward the atom with stronger attraction but are not fully transferred (a "competitive fight").

3. Electronegativity difference < 0.4: The bond exhibits nonpolar covalent bond characteristics. The difference in electron-attracting ability is negligible, meaning electrons are not significantly drawn to one side (a "balanced battle").

Following this criterion, one can understand why diatomic molecules like H₂ and O₂ form nonpolar covalent bonds, while fluorine and most metals form ionic bonds. For molecules with an electronegativity difference close to 1.7, their chemical bonds exhibit both covalent and ionic characteristics.

Figure 1:Electronegativity Periodic Table