Molecular Polarity and Dipole Moment

They can occur between two ions in an ionic bond or between atoms in a covalent bond; dipole moments arise from differences in electronegativity. The larger the difference in electronegativity, the larger the dipole moment. … The dipole moment is a measure of the polarity of the molecule.

Not only bonds can have dipole moments, however. A whole molecule may also have a separation of charge, depending on its molecular structure and the polarity of each of its bonds. If such a charge separation exists, the molecule is said to be a polar molecule (or dipole); otherwise the molecule is said to be nonpolar. The dipole moment measures the extent of net charge separation in the molecule as a whole. We determine the dipole moment by adding the bond moments in three-dimensional space, taking into account the molecular structure.

For diatomic molecules, there is only one bond, so its bond dipole moment determines the molecular polarity. Homonuclear diatomic molecules such as Br₂ and N₂ have no difference in electronegativity, so their dipole moment is zero. For heteronuclear molecules such as CO, there is a small dipole moment. For HF, there is a larger dipole moment because there is a larger difference in electronegativity.

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. The dipole moment of a molecule is 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 (Figure below). 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 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.

Figure7.103 Above: 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.

Reference: https://courses.lumenlearning.com/suny-mcc-organicchemistry/chapter/dipole-moments/

In the following OCS molecule has a structure similar to CO₂, but a sulfur atom has replaced one of the oxygen atoms. To determine if this molecule is polar, we draw the molecular structure. VSEPR theory predicts a linear molecule:

Figure 7.104 Molecular Polarity of  OCS

Chloromethane, CH₃Cl, is another example of a polar molecule. Although the polar C–Cl and C–H bonds are arranged in a tetrahedral geometry, the C–Cl bonds have a larger bond moment than the C–H bond, and the bond moments do not completely cancel each other. All of the dipoles have a upward component in the orientation shown, since carbon is more electronegative than hydrogen and less electronegative than chlorine:

Figure 7.105 Molecular Polarity of  CH₄

Other examples of molecules with polar bonds are shown in the following Figures. 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.

Figures 7.106  showMolecules 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).

In brief, to summarize, to be polar, a molecule must:

1. Contain at least one polar covalent bond.
2. Have a molecular structure such that the sum of the vectors of each bond dipole moment does not cancel.

Properties of Polar Molecules

Polar molecules occur when there is an electronegativity difference between the bonded atoms. Nonpolar molecules occur when electrons are shared equal between atoms of a diatomic molecule or when polar bonds in a larger molecule cancel each other out. Following Phet Simulation activity may help to understand better.

Molecule polarity is illustrated in a Phet simulation activity below:

https://phet.colorado.edu/sims/html/molecule-polarity/latest/molecule-polarity_en.html

In the simulation above, the students can explore:

Predict how changing electronegativity will affect the bond polarity.

Explain the relationship between the bond dipoles and the molecular dipole.

Determine if a non-polar molecule can contain polar bonds.

Describe how the ABC bond angle effects the molecular dipole.

Compare the behavior of non-polar and polar molecules in an external electric field.

Two Atoms Screen

Change the electronegativity of the atoms, view the resulting electrostatic potential or electron density, and predict the bond polarity.

Three Atoms Screen

Explore the relationship between the bond dipoles and the molecular dipole, and observe the molecule in an electric field.

Simulation conclusions:

The electronegativity slider ranges from 2 to 4, but the value is never displayed. The resulting electronegativity difference between two bonded atoms varies from 0 to 2.

Bond dipoles are parallel to the bond axis, and their length is linearly proportional to the difference in electronegativity. Note that this is a simplification; in reality, the dipole is not influenced solely by electronegativity.

The molecular dipole is the vector sum of the bond dipoles. In the Two Atoms screen, the molecular dipole is not shown, as it is equivalent to the bond dipole. In the Three Atoms screen, manipulating electronegativity results in an understanding of summing vector magnitudes, while manipulating bond angles results in an understanding of summing vector angles.

The magnitude of an atom’s partial charge is linearly proportional to the electronegativity difference between the bonded pair. If an atom has a higher electronegativity than the atom at the other end of the bond, then the partial charge’s sign is negative; otherwise it is positive. For atoms that participate in more than one bond (e.g., atom B in the “Three Atoms” screen), net partial charge is the sum of the partial charges contributed by each bond.

The electrostatic potential and electron density are linearly proportional to the electronegativity difference set by the sliders. These surfaces are not implemented for the triatomic molecule in the Three Atoms screen, because the manipulation of bond angles results in undefinable surfaces.

The Three Atoms screen allows for students to change the bond angle between the outer atoms C). The    AB and BC bonds are treated independently, and the model does not allow for these atoms to    repel     each other. To explore how atoms would repel one another when the bond angles are changed.

Follow Up Questions:

1.       Explain how the polarity of a molecule is related to the electronegativity of the atoms within the molecule. Use your knowledge gained from the previous chapters.

•           In general, in chemistry, polarity refers to the distribution of electric charge around atoms, chemical groups, or molecules.

•           Polar molecules occur when there is an electronegativity difference between the bonded atoms.

•           Nonpolar molecules occur when electrons are shared equal between atoms of a diatomic molecule or when polar bonds in a larger molecule cancel each other out.

Examples of polar molecules include:

Water – H₂O                           Hydrogen sulfide – H₂S

Ammonia – NH₃                     Ethanol – C₂H₆O

Sulfur dioxide – SO₂

Examples of nonpolar molecules include:

• Any of the noble gasses: He, Ne, Ar, Kr, Xe (These are atoms, not technically molecules.)
• Any of the homonuclear diatomic elements: H₂, N₂, O₂, Cl₂ (These are truly nonpolar molecules.)
• Carbon dioxide – CO₂
• Benzene – C₆H₆
• Carbon tetrachloride – CCl₄
• Methane – CH₄
• Ethylene – C₂H₄
• Hydrocarbon liquids, such as gasoline and toluene
• Most organic molecules

Reference: Pauling, L. (1960). The Nature of the Chemical Bond (3rd ed.). Oxford University Press. pp. 98–100. ISBN 0801403332.

Reference: https://www.youtube.com/watch?v=SiZXRScxbl0&t=2s

Reference: https://youtu.be/4ykSzYl_4vI