Polarity

6.3 Bond Polarity and Molecular Polarity

Learning Objective

To calculate the percent ionic character of a covalent polar bond..

So far in this text, especially in section 4.2, we have described the two idealized extremes of chemical bonding: (1) ionic bonding—in which one or more electrons are transferred completely from one atom to another, and the resulting ions are held together by purely electrostatic forces—and (2) covalent bonding, in which electrons are shared equally between two atoms. Most compounds, however, have polar covalent bondsA covalent bond in which the electrons are shared unequally between the bonded atoms., which means that electrons are shared unequally between the bonded atoms. Figure 6.3(a) compares the electron distribution in a polar covalent bond with those in an ideally covalent and an ideally ionic bond. A lowercase Greek delta ( $δ$ ) is used to indicate that a bonded atom possesses a partial positive charge, indicated by $δ^{+},$ or a partial negative charge, indicated by $δ^{-},$ and a bond between two atoms that possess partial charges is a polar bond.

Figure 6.3(a) The Electron Distribution in a Nonpolar Covalent Bond, a Polar Covalent Bond, and an Ionic Bond Using Lewis Electron Structures

In a purely covalent bond (a), the bonding electrons are shared equally between the atoms. In a purely ionic bond (c), an electron has been transferred completely from one atom to the other. A polar covalent bond (b) is intermediate between the two extremes: the bonding electrons are shared unequally between the two atoms, and the electron distribution is asymmetrical with the electron density being greater around the more electronegative atom. Electron-rich (negatively charged) regions are shown in blue; electron-poor (positively charged) regions are shown in red.

Electronegativity

The tendency of an element to gain or lose electrons is important in determining its chemistry. The most important method to describe this tendency uses a measurement called electronegativityThe relative ability of an atom to attract electrons to itself in a chemical compound. (represented by the Greek letter chi, χ, pronounced “ky” as in “sky”), defined as the relative ability of an atom to attract electrons to itself in a chemical compound. Elements with high electronegativities tend to acquire electrons in chemical reactions and are found in the upper right corner of the periodic table. Elements with low electronegativities tend to lose electrons in chemical reactions and are found in the lower left corner of the periodic table.

The Pauling Electronegativity Scale

The original electronegativity scale, developed in the 1930s by Linus Pauling (1901– 1994) was based on measurements of the strengths of covalent bonds between different elements. Pauling arbitrarily set the electronegativity of fluorine at 4.0 (although today it has been refined to 3.98), thereby creating a scale in which all elements have values between 0 and 4.0.

Periodic variations in Pauling’s electronegativity values are illustrated in Figure 6.3(b) and Figure 6.3(c). If we ignore the inert gases and elements for which no stable isotopes are known, we see that fluorine (χ = 3.98) is the most electronegative element and cesium is the least electronegative nonradioactive element (χ = 0.79). Because electronegativities generally increase diagonally from the lower left to the upper right of the periodic table, elements lying on diagonal lines running from upper left to lower right tend to have comparable values (e.g., O and Cl and N, S, and Br).

Figure 6.3(b) A Plot of Electronegativity Values

Linus Pauling (1901–1994)

Pauling won two Nobel Prizes, one for chemistry in 1954 and one for peace in 1962. When he was nine, Pauling’s father died, and his mother tried to convince him to quit school to support the family. He did not quit school but was denied a high school degree because of his refusal to take a civics class.

Figure 6.3(c) Pauling Electronegativity Values of the s-, p-, d-, and f-Block Elements

Values for most of the actinides are approximate. Elements for which no data are available are shown in gray.

Source: Data from L. Pauling, The Nature of the Chemical Bond, 3rd ed. (1960).

Example 6.3-1

On the basis of their positions in the periodic table, arrange Cl, Se, Si, and Sr in order of increasing electronegativity and classify each as a metal, a nonmetal, or a semimetal.

Given: four elements

Asked for: order by increasing electronegativity and classification

Strategy:

A Locate the elements in the periodic table. From their diagonal positions from lower left to upper right, predict their relative electronegativities.

B Arrange the elements in order of increasing electronegativity.

C Classify each element as a metal, a nonmetal, or a semimetal according to its location about the diagonal belt of semimetals running from B to At.

Solution:

A Electronegativity increases from lower left to upper right in the periodic table (See Figure 6.3(c)). Because Sr lies far to the left of the other elements given, we can predict that it will have the lowest electronegativity. Because Cl lies above and to the right of Se, we can predict that χ_Cl > χ_Se. Because Si is located farther from the upper right corner than Se or Cl, its electronegativity should be lower than those of Se and Cl but greater than that of Sr. B The overall order is therefore χ_Sr < χ_Si < χ_Se < χ_Cl.

C To classify the elements, we note that Sr lies well to the left of the diagonal belt of semimetals running from B to At; while Se and Cl lie to the right and Si lies in the middle. We can predict that Sr is a metal, Si is a semimetal, and Se and Cl are nonmetals.

Exercise

On the basis of their positions in the periodic table, arrange Ge, N, O, Rb, and Zr in order of increasing electronegativity and classify each as a metal, a nonmetal, or a semimetal.

Answer: Rb < Zr < Ge < N < O; metals (Rb, Zr); semimetal (Ge); nonmetal (N, O)

Note the Pattern

Electronegativity values increase from lower left to upper right in the periodic table.

Bond Polarity

The polarity of a bond—the extent to which it is polar—is determined largely by the relative electronegativities of the bonded atoms. We have defined electronegativity as the ability of an atom in a molecule or an ion to attract electrons to itself. Thus there is a direct correlation between electronegativity and bond polarity. A bond is nonpolar if the bonded atoms have equal electronegativities. If the electronegativities of the bonded atoms are not equal, however, the bond is polarized toward the more electronegative atom. A bond in which the electronegativity of B (χ_B) is greater than the electronegativity of A (χ_A), for example, is indicated with the partial negative charge on the more electronegative atom:

A_{less electronegative}^{δ^{+}} – B_{more electronegative}^{δ^{-}}

One way of estimating the ionic character of a bond—that is, the magnitude of the charge separation in a polar covalent bond—is to calculate the difference in electronegativity between the two atoms: Δχ = χ_B − χ_A.

To predict the polarity of the bonds in Cl₂, HCl, and NaCl, for example, we look at the electronegativities of the relevant atoms: χ_Cl = 3.16, χ_H = 2.20, and χ_Na = 0.93 (see Figure 7.14 "A Plot of Periodic Variation of Electronegativity with Atomic Number for the First Six Rows of the Periodic Table"). Cl₂ must be nonpolar because the electronegativity difference (Δχ) is zero; hence the two chlorine atoms share the bonding electrons equally. In NaCl, Δχ is 2.23. This high value is typical of an ionic compound (Δχ ≥ ≈1.5) and means that the valence electron of sodium has been completely transferred to chlorine to form Na⁺ and Cl⁻ ions. In HCl, however, Δχ is only 0.96. The bonding electrons are more strongly attracted to the more electronegative chlorine atom, and so the charge distribution is

\overset{δ^{+}}{H} – \overset{δ^{-}}{Cl}

Remember that electronegativities are difficult to measure precisely and different definitions produce slightly different numbers. In practice, the polarity of a bond is usually estimated rather than calculated.

Note the Pattern

Bond polarity and ionic character increase with an increasing difference in electronegativity.

Dipole Moments

The asymmetrical charge distribution in a polar substance such as HCl produces a dipole moment. When a molecule with a dipole moment is placed in an electric field, it tends to orient itself with the electric field because of its asymmetrical charge distribution (Figure 6.3(d).

Figure 6.3(d) Molecules That Possess a Dipole Moment Partially Align Themselves with an Applied Electric Field

In the absence of a field (a), the HCl molecules are randomly oriented. When an electric field is applied (b), the molecules tend to align themselves with the field, such that the positive end of the molecular dipole points toward the negative terminal and vice versa.

Video: A stream of acetone is bent by an electric field while a stream of cylohexane remains unchanged by the electric field. Video Credit: Highland Community College Please attribute this work as being created by Highland Community College. This work is licensed under Creative Commons CC-BY https://creativecommons.org/licenses/by/3.0/ via panopto

More Complicated Molecular Dipole Moments

The dipole moments of simple diatomic molecules, like HCl, are straightforward. With only two atoms in the molecule, the more electronegative atom is going to pull electron density toward itself making that atom more negative, leaving the other atom more positive, and therefore creating a dipole. In more complex molecules with polar covalent bonds, the three-dimensional geometry and the compound’s symmetry determine whether there is a net dipole moment. Mathematically, dipole moments are vectors; they possess both a magnitude and a direction. The dipole moment of a molecule is therefore the vector sum of the dipole moments of the individual bonds in the molecule. If the individual bond dipole moments cancel one another, there is no net dipole moment. Such is the case for CO₂, a linear molecule (part (a) in Figure 6.3(e)). Each C–O bond in CO₂ is polar, yet experiments show that the CO₂ molecule has no dipole moment. Because the two C–O bond dipoles in CO₂ are equal in magnitude and oriented at 180° to each other, they cancel. As a result, the CO₂ molecule has no net dipole moment even though it has a substantial separation of charge. In contrast, the H₂O molecule is not linear (part (b) in Figure 6.3(e)); it is bent in three-dimensional space, so the dipole moments do not cancel each other. Thus a molecule such as H₂O has a net dipole moment. We expect the concentration of negative charge to be on the oxygen, the more electronegative atom, and positive charge on the two hydrogens. This charge polarization allows H₂O to hydrogen-bond to other polarized or charged species, including other water molecules.

Figure 6.3(e) How Individual Bond Dipole Moments Are Added Together to Give an Overall Molecular Dipole Moment for Two Triatomic Molecules with Different Structures

(a) In CO₂, the C–O bond dipoles are equal in magnitude but oriented in opposite directions (at 180°). Their vector sum is zero, so CO₂ therefore has no net dipole. (b) In H₂O, the O–H bond dipoles are also equal in magnitude, but they are oriented at 104.5° to each other. Hence the vector sum is not zero, and H₂O has a net dipole moment.

Other examples of molecules with polar bonds are shown in Figure 6.3(f). In molecular geometries that are highly symmetrical (most notably tetrahedral and square planar, trigonal bipyramidal, and octahedral), individual bond dipole moments completely cancel, and there is no net dipole moment. Although a molecule like CHCl₃ is best described as tetrahedral, the atoms bonded to carbon are not identical. Consequently, the bond dipole moments cannot cancel one another, and the molecule has a dipole moment. Due to the arrangement of the bonds in molecules that have V-shaped, trigonal pyramidal, seesaw, T-shaped, and square pyramidal geometries, the bond dipole moments cannot cancel one another. Consequently, molecules with these geometries always have a nonzero dipole moment.

Figure 6.3(f) Molecules with Polar Bonds

Individual bond dipole moments are indicated in red. Due to their different three-dimensional structures, some molecules with polar bonds have a net dipole moment (HCl, CH₂O, NH₃, and CHCl₃), indicated in blue, whereas others do not because the bond dipole moments cancel (BCl₃, CCl₄, PF₅, and SF₆).

Note the Pattern

Molecules with asymmetrical charge distributions have a net dipole moment.

Example 6.3-2

Which molecule(s) has a net dipole moment?

H₂S
NHF₂
BF₃

Given: three chemical compounds

Asked for: net dipole moment

Strategy:

For each three-dimensional molecular geometry, predict whether the bond dipoles cancel. If they do not, then the molecule has a net dipole moment.

Solution:

The total number of electrons around the central atom, S, is eight, which gives four electron pairs. Two of these electron pairs are bonding pairs and two are lone pairs, so the molecular geometry of H₂S is bent (Figure 6.2(f)). The bond dipoles cannot cancel one another, so the molecule has a net dipole moment.
Difluoroamine has a trigonal pyramidal molecular geometry. Because there is one hydrogen and two fluorines, and because of the lone pair of electrons on nitrogen, the molecule is not symmetrical, and the bond dipoles of NHF₂ cannot cancel one another. This means that NHF₂ has a net dipole moment. We expect polarization from the two fluorine atoms, the most electronegative atoms in the periodic table, to have a greater affect on the net dipole moment than polarization from the lone pair of electrons on nitrogen.
The molecular geometry of BF₃ is trigonal planar. Because all the B–F bonds are equal and the molecule is highly symmetrical, the dipoles cancel one another in three-dimensional space. Thus BF₃ has a net dipole moment of zero:

Exercise

Which molecule(s) has a net dipole moment?

CH₃Cl
SO₃
XeO₃

Answer: CH₃Cl; XeO₃

Summary

Compounds with polar covalent bonds have electrons that are shared unequally between the bonded atoms.The electronegativity (χ) of an element is the relative ability of an atom to attract electrons to itself in a chemical compound and increases diagonally from the lower left of the periodic table to the upper right. The polarity of such a bond is determined largely by the relative electronegativites of the bonded atoms. The asymmetrical charge distribution in a polar substance produces a dipole moment, which is the product of the partial charges on the bonded atoms and the distance between them.

Molecules with polar covalent bonds can have a dipole moment, an asymmetrical distribution of charge that results in a tendency for molecules to align themselves in an applied electric field. Any diatomic molecule with a polar covalent bond has a dipole moment, but in polyatomic molecules, the presence or absence of a net dipole moment depends on the structure. For some highly symmetrical structures, the individual bond dipole moments cancel one another, giving a dipole moment of zero.

Key Takeaway

Bond polarity and ionic character increase with an increasing difference in electronegativity.

Numerical Problems

Predict whether each compound is purely covalent, purely ionic, or polar covalent.
1. RbCl
2. S₈
3. TiCl₂
4. SbCl₃
5. LiI
6. Br₂
Based on relative electronegativities, classify the bonding in each compound as ionic, covalent, or polar covalent. Indicate the direction of the bond dipole for each polar covalent bond.
1. NO
2. HF
3. MgO
4. AlCl₃
5. SiO₂
6. the C=O bond in acetone
7. O₃
Based on relative electronegativities, classify the bonding in each compound as ionic, covalent, or polar covalent. Indicate the direction of the bond dipole for each polar covalent bond.
1. NaBr
2. OF₂
3. BCl₃
4. the S–S bond in CH₃CH₂SSCH₂CH₃
5. the C–Cl bond in CH₂Cl₂
6. the O–H bond in CH₃OH
Predict whether each molecule has a net dipole moment. Justify your answers and indicate the direction of any bond dipoles.
1. NO
2. HF
3. PCl₃
4. CO₂
5. SO₂
6. SF₄
Predict whether each molecule has a net dipole moment. Justify your answers and indicate the direction of any bond dipoles.
1. OF₂
2. BCl₃
3. CH₂Cl₂
4. TeF₄
5. CH₃OH
6. XeO₄
Of the molecules Cl₂C=Cl₂, IF₃, and SF₆, which has a net dipole moment? Explain your reasoning.
Of the molecules SO₃, XeF₄, and H₂C=Cl₂, which has a net dipole moment? Explain your reasoning.

Answers

1. purely ionic
2. purely covalent
3. purely ionic
4. polar covalent
5. purely ionic
6. purely covalent
1. ionic
2. polar covalent
3. polar covalent
4. covalent
5. polar covalent
6. polar covalent
Cl₂C=CCl₂: Although the C–Cl bonds are rather polar, the individual bond dipoles cancel one another in this symmetrical structure, and Cl₂C=CCl₂ does not have a net dipole moment.

IF₃: In this structure, the individual I–F bond dipoles cannot cancel one another, giving IF₃ a net dipole moment.

SF₆: The S–F bonds are quite polar, but the individual bond dipoles cancel one another in an octahedral structure. Thus, SF₆ has no net dipole moment.