6.2 Electronegativity and Polarity

Compounds that contain covalent bonds exhibit different physical properties than ionic compounds. Because the attraction between molecules, which are electrically neutral, is weaker than that between electrically charged ions, covalent compounds generally have much lower melting and boiling points than ionic compounds. In fact, many covalent compounds are liquids or gases at room temperature, and, in their solid states, they are typically much softer than ionic solids. Furthermore, whereas ionic compounds are good conductors of electricity when dissolved in water, most covalent compounds are insoluble in water; since they are electrically neutral, they are poor conductors of electricity in any state.

Pure vs. Polar Covalent Bonds

If the atoms that form a covalent bond are identical, as in [latex]\ce{H2}[/latex], [latex]\ce{Cl2}[/latex], and other diatomic molecules, then the electrons in the bond must be shared equally. We refer to this as a pure covalent bond. Electrons shared in pure covalent bonds have an equal probability of being near each nucleus.

In the case of [latex]\ce{Cl2}[/latex], each atom starts off with seven valence electrons, and each [latex]\ce{Cl}[/latex] shares one electron with the other, forming one covalent bond:

[latex]\ce{Cl}+\ce{Cl}\rightarrow \ce{Cl2}[/latex]

The total number of electrons around each individual atom consists of six nonbonding electrons and two shared (i.e., bonding) electrons for eight total electrons, matching the number of valence electrons in the noble gas argon. Since the bonding atoms are identical, [latex]\ce{Cl2}[/latex] also features a pure covalent bond.

When the atoms linked by a covalent bond are different, the bonding electrons are shared, but no longer equally. Instead, the bonding electrons are more attracted to one atom than the other, giving rise to a shift of electron density toward that atom. This unequal distribution of electrons is known as a polar covalent bond, characterized by a partial positive charge on one atom and a partial negative charge on the other. The atom that attracts the electrons more strongly acquires the partial negative charge and vice versa. For example, the electrons in the [latex]\ce{H-Cl}[/latex] bond of a hydrogen chloride molecule spend more time near the chlorine atom than near the hydrogen atom. Thus, in an [latex]\ce{HCl}[/latex] molecule, the chlorine atom carries a partial negative charge and the hydrogen atom has a partial positive charge. Figure 6.2.1 shows the distribution of electrons in the [latex]\ce{H-Cl}[/latex] bond. Note that the shaded area around [latex]\ce{Cl}[/latex] is much larger than it is around [latex]\ce{H}[/latex]. Compare this to the even distribution of electrons in a [latex]\ce{H2}[/latex] nonpolar bond.

We sometimes designate the positive and negative atoms in a polar covalent bond using a lowercase Greek letter “delta,” [latex]\delta[/latex], with a plus sign or minus sign to indicate whether the atom has a partial positive charge ([latex]\delta{+}[/latex]) or a partial negative charge ([latex]\delta{-}[/latex]). This symbolism is shown for the [latex]\ce{H-Cl}[/latex] molecule in Figure 6.2.1.

Electronegativity

Whether a bond is nonpolar or polar covalent is determined by a property of the bonding atoms called electronegativity. Electronegativity is a measure of the tendency of an atom to attract electrons (or electron density) towards itself. It determines how the shared electrons are distributed between the two atoms in a bond. The more strongly an atom attracts the electrons in its bonds, the larger its electronegativity. Electrons in a polar covalent bond are shifted toward the more electronegative atom; thus, the more electronegative atom is the one with the partial negative charge. The greater the difference in electronegativity, the more polarized the electron distribution and the larger the partial charges of the atoms.

Figure 6.2.2 shows the electronegativity values of the elements as proposed by one of the most famous chemists of the twentieth century: Linus Pauling (Figure 6.2.3). In general, electronegativity increases from left to right across a period in the periodic table and decreases down a group. Thus, the nonmetals, which lie in the upper right, tend to have the highest electronegativities, with fluorine the most electronegative element of all (EN = 4.0). Metals tend to be less electronegative elements, and the group 1 metals have the lowest electronegativities. Note that noble gases are excluded from this figure because these atoms usually do not share electrons with others atoms since they have a full valence shell. (While noble gas compounds such as [latex]\ce{XeO2}[/latex] do exist, they can only be formed under extreme conditions, and thus they do not fit neatly into the general model of electronegativity.)

Electronegativity versus Electron Affinity

We must be careful not to confuse electronegativity and electron affinity. The electron affinity of an element is a measurable physical quantity, namely, the energy released or absorbed when an isolated gas-phase atom acquires an electron, measured in kJ/mol. Electronegativity, on the other hand, describes how tightly an atom attracts electrons in a bond. It is a dimensionless quantity that is calculated, not measured. Pauling derived the first electronegativity values by comparing the amounts of energy required to break different types of bonds. He chose an arbitrary relative scale ranging from 0 to 4.

Portrait of a Chemist: Linus Pauling

Linus Pauling, shown in Figure 6.2.3, is the only person to have received two unshared (individual) Nobel Prizes: one for chemistry in 1954 for his work on the nature of chemical bonds and one for peace in 1962 for his opposition to weapons of mass destruction. He developed many of the theories and concepts that are foundational to our current understanding of chemistry, including electronegativity and resonance structures.

Pauling also contributed to many other fields besides chemistry. His research on sickle cell anemia revealed the cause of the disease—the presence of a genetically inherited abnormal protein in the blood—and paved the way for the field of molecular genetics. His work was also pivotal in curbing the testing of nuclear weapons; he proved that radioactive fallout from nuclear testing posed a public health risk.

Electronegativity and Bond Type

The absolute value of the difference in electronegativity (ΔEN) of two bonded atoms provides a rough measure of the polarity to be expected in the bond and, thus, the bond type. When the difference is very small or zero, the bond is covalent and nonpolar. When it is large, the bond is polar covalent or ionic. The absolute values of the electronegativity differences between the atoms in the bonds [latex]\ce{H-H}[/latex], [latex]\ce{H-Cl}[/latex], and [latex]\ce{Na-Cl}[/latex] are 0 (nonpolar), 0.9 (polar covalent), and 2.1 (ionic), respectively. The degree to which electrons are shared between atoms varies from completely equal (pure covalent bonding) to not at all (ionic bonding). Figure 6.2.4 shows the relationship between electronegativity difference and bond type.

A rough approximation of the electronegativity differences associated with covalent, polar covalent, and ionic bonds is shown Table 6.2.1. This table is just a general guide, however, with many exceptions. For example, the [latex]\ce{H}[/latex] and [latex]\ce{F}[/latex] atoms in [latex]\ce{HF}[/latex] have an electronegativity difference of 1.9, and the [latex]\ce{N}[/latex] and [latex]\ce{H}[/latex] atoms in [latex]\ce{NH3}[/latex] a difference of 0.9, yet both of these compounds form bonds that are considered polar covalent. Likewise, the [latex]\ce{Na}[/latex] and [latex]\ce{Cl}[/latex] atoms in [latex]\ce{NaCl}[/latex] have an electronegativity difference of 2.1, and the Mn and I atoms in [latex]\ce{MnI2}[/latex] have a difference of 1.0, yet both of these substances form ionic compounds.

The best guide to the covalent or ionic character of a bond is to consider the types of atoms involved and their relative positions in the periodic table. Bonds between two nonmetals are generally covalent; bonding between a metal and a nonmetal is often ionic.

Some compounds contain both covalent and ionic bonds. The atoms in polyatomic ions, such as [latex]\ce{OH-}[/latex], [latex]\ce{NO3-},[/latex] and [latex]\ce{NH4+},[/latex] are held together by polar covalent bonds. However, these polyatomic ions form ionic compounds by combining with ions of opposite charge. For example, potassium nitrate, [latex]\ce{KNO3}[/latex], contains the [latex]\ce{K+}[/latex] cation and the polyatomic [latex]\ce{NO3-}[/latex] anion. Thus, bonding in potassium nitrate is ionic, resulting from the electrostatic attraction between the ions [latex]\ce{K+}[/latex] and [latex]\ce{NO3-},[/latex] as well as covalent between the nitrogen and oxygen atoms in [latex]\ce{NO3-}[/latex].

Glossary

covalent bond: bond formed when electrons are shared between atoms

electronegativity: tendency of an atom to attract electrons in a bond to itself

polar covalent bond: covalent bond between atoms of different electronegativities; a covalent bond with a positive end and a negative end

pure covalent bond: (also, nonpolar covalent bond) covalent bond between atoms of identical electronegativities