D10.3 Electronegativity

The polarity of a covalent bond can be estimated by the difference between the electronegativities of the bonded atoms. Electronegativity (EN) is the tendency of an atom in a molecule to attract bonding electron density. Thus, in a bond, the more electronegative atom is the one with the δ charge.

The greater the difference in electronegativity between two bonded atoms, the larger the shift of electron density in the bond towards the more electronegative atom. Greater electronegativity difference, Δ(EN), gives larger partial charges on the atoms. Electronegativity values for most elements are shown in the periodic table in the Activity below; they also are tabulated in the appendix.

Activity: Periodic Trends in Electronegativity

Electronegativity, Electron Affinity, and Ionization Energy

These three properties are all associated with an atom gaining/losing electrons. Electron affinity and ionization energy are experimentally measurable physical quantities.

Electron affinity (EA) is the energy change when an isolated gas-phase atom acquires an electron; it is usually expressed in kJ/mol.

X(g) + e → X(g)          ΔE = EA1

Ionization energy (IE) is the energy that must be transferred to an isolated gas-phase atom to remove an electron; it is also typically expressed in kJ/mol.

X(g) → X+(g) + e          ΔE = IE1

Electronegativity describes how strongly an atom attracts electron density in a bond. It is calculated, not measured, has an arbitrary relative scale, and has no units.

Activity: Electronegativity, Electron Affinity, and Ionization Energy

Electronegativity and Bond Type

The difference in electronegativity, Δ(EN), of two bonded atoms provides a rough estimate of polarity of the bond, and thus of the bond type. When Δ(EN) is very small (or zero), the bond is covalent and nonpolar. When Δ(EN) is large, the bond is polar covalent or ionic. (In a pair of ions, such as Na+Cl, there is nearly complete transfer of valence electrons from one atom to another to produce a positive ion and a negative ion. The Na+ and Cl ions form a dipole with δ+ approximately equal to +1 and δ approximately −1.)

Δ(EN) spans a continuous scale and serves as a general guide; there is no definitive cutoff that defines a bond type. For example, HF has Δ(EN) = 1.8 and is considered a polar covalent molecule. On the other hand, NaI has a Δ(EN) of 1.7 but is considered an ionic compound. When estimating the covalent or ionic character of a bond, you should also take into account the types of atoms involved and their relative positions in the periodic table. Bonds between two nonmetals are usually described as covalent; bonding between a metal and a nonmetal is often ionic.

Some compounds contain both covalent and ionic bonds. For example, potassium nitrate, KNO3, contains the K+ cation and the polyatomic NO3 anion, which has covalent bonds between N and O.

Exercise: Polarity and Electronegativity Difference

Exercise: Bond Polarity and Electronegativity

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