3.3 Polar and Nonpolar Covalent Bonds | Chemical Bonding | Chemistry 11

This section explores the distinction between polar and nonpolar covalent bonds and how the overall polarity of a molecule is determined by both bond polarity and molecular geometry.

Polar Covalent Bonds

A polar covalent bond is formed when there is a significant difference in electronegativity between the two atoms sharing electrons. The more electronegative atom attracts the shared electron pair more strongly, acquiring a partial negative charge ( $δ^{-}$ ). The less electronegative atom consequently acquires a partial positive charge ( $δ^{+}$ ).

This separation of charge creates a dipole, which is a pair of equal and opposite electric charges separated by a distance. Molecules with a net dipole are called polar molecules.

Nonpolar Covalent Molecules

A molecule can be nonpolar for two main reasons:

It consists of only nonpolar bonds (e.g., $H_{2}$ , $C l_{2}$ ).
It has polar bonds, but the molecular geometry is symmetrical, causing the bond dipoles to cancel each other out.

Symmetrical Geometries:

Symmetrical polyatomic molecules are often nonpolar. Key symmetrical shapes include:

Linear
Trigonal Planar
Tetrahedral

In these arrangements, the individual bond dipoles exert equal and opposite effects, resulting in a net dipole moment of zero. Therefore, the molecule has no overall charge and is considered nonpolar.

Example: Carbon Dioxide ( $C O_{2}$ )

The carbon dioxide molecule provides a classic example of a nonpolar molecule with polar bonds.

Bond Polarity: There is a significant electronegativity difference between Carbon ( $E N = 2.5$ ) and Oxygen ( $E N = 3.5$ ). This makes each Carbon-Oxygen double bond (C=O) polar.
Molecular Geometry: $C O_{2}$ is a linear molecule ( $O = C = O$ ). The two C=O bond dipoles are equal in magnitude but point in opposite directions.

Linear geometry of CO2 — Figure 3.3.1: Linear geometry of CO₂ showing bond dipoles canceling out.

As shown in the diagram, the dipoles cancel each other out, leading to a zero net dipole moment. Thus, $C O_{2}$ is a nonpolar molecule.

Concept Assessment Exercise 3.1

Determine if the following molecules are polar or nonpolar, providing a reason for each.

Given Electronegativity Values: $F = 4.0$ , $C l = 3.0$ , $B r = 2.8$ , $S = 2.5$ , $C = 2.5$ , $H = 2.1$

a. Chlorine, $C l_{2}$

Analysis: The bond is between two chlorine atoms, which have the same electronegativity.
Electronegativity Difference ( $Δ E N$ ): $Δ E N = 3.0 - 3.0 = 0$
Conclusion: The Cl-Cl bond is a pure (nonpolar) covalent bond. The molecule is nonpolar.

b. Hydrogen fluoride, $H F$

Analysis: The bond is between hydrogen and fluorine.
Electronegativity Difference ( $Δ E N$ ): $Δ E N = 4.0 - 2.1 = 1.9$
Conclusion: The H-F bond is highly polar. Since it is a diatomic molecule, there is a net dipole moment. The molecule is polar.

c. Sulfur dichloride, $S C l_{2}$

Analysis: The molecule has two S-Cl bonds.
Electronegativity Difference ( $Δ E N$ ): $Δ E N = 3.0 - 2.5 = 0.5$
Conclusion: The S-Cl bonds are polar. The molecule has a V-shaped (bent) geometry, which is asymmetrical. The bond dipoles do not cancel each other out, resulting in a net dipole moment. The molecule is polar.

d. Chloromethane, $C H_{3} C l$

Analysis: The molecule has three C-H bonds and one C-Cl bond arranged in a tetrahedral geometry around the central carbon atom.
Electronegativity Difference ( $Δ E N$ ):
- C-Cl: $Δ E N = 3.0 - 2.5 = 0.5$ (Polar)
- C-H: $Δ E N = 2.5 - 2.1 = 0.4$ (Slightly polar)
Conclusion: The molecule is asymmetrical because the atoms bonded to the central carbon are not identical. The strong C-Cl bond dipole is not canceled by the C-H bond dipoles. The molecule is polar.

e. Tetrabromomethane, $C B r_{4}$

Analysis: The molecule has four C-Br bonds arranged in a tetrahedral geometry.
Electronegativity Difference ( $Δ E N$ ): $Δ E N = 2.8 - 2.5 = 0.3$
Conclusion: The C-Br bonds are slightly polar. However, the molecule has a symmetrical tetrahedral shape. The four C-Br bond dipoles are arranged symmetrically and cancel each other out. The molecule is nonpolar.

This section explores the distinction between polar and nonpolar covalent bonds and how the overall polarity of a molecule is determined by both bond polarity and molecular geometry.

Polar Covalent Bonds

This separation of charge creates a dipole, which is a pair of equal and opposite electric charges separated by a distance. Molecules with a net dipole are called polar molecules.

Nonpolar Covalent Molecules

A molecule can be nonpolar for two main reasons:

It consists of only nonpolar bonds (e.g., $H_{2}$ , $C l_{2}$ ).
It has polar bonds, but the molecular geometry is symmetrical, causing the bond dipoles to cancel each other out.

Symmetrical Geometries:

Symmetrical polyatomic molecules are often nonpolar. Key symmetrical shapes include:

Linear
Trigonal Planar
Tetrahedral

Example: Carbon Dioxide ( $C O_{2}$ )

The carbon dioxide molecule provides a classic example of a nonpolar molecule with polar bonds.

Bond Polarity: There is a significant electronegativity difference between Carbon ( $E N = 2.5$ ) and Oxygen ( $E N = 3.5$ ). This makes each Carbon-Oxygen double bond (C=O) polar.
Molecular Geometry: $C O_{2}$ is a linear molecule ( $O = C = O$ ). The two C=O bond dipoles are equal in magnitude but point in opposite directions.

As shown in the diagram, the dipoles cancel each other out, leading to a zero net dipole moment. Thus, $C O_{2}$ is a nonpolar molecule.

Concept Assessment Exercise 3.1

Determine if the following molecules are polar or nonpolar, providing a reason for each.

Given Electronegativity Values: $F = 4.0$ , $C l = 3.0$ , $B r = 2.8$ , $S = 2.5$ , $C = 2.5$ , $H = 2.1$

a. Chlorine, $C l_{2}$

Analysis: The bond is between two chlorine atoms, which have the same electronegativity.
Electronegativity Difference ( $Δ E N$ ): $Δ E N = 3.0 - 3.0 = 0$
Conclusion: The Cl-Cl bond is a pure (nonpolar) covalent bond. The molecule is nonpolar.

b. Hydrogen fluoride, $H F$

Analysis: The bond is between hydrogen and fluorine.
Electronegativity Difference ( $Δ E N$ ): $Δ E N = 4.0 - 2.1 = 1.9$
Conclusion: The H-F bond is highly polar. Since it is a diatomic molecule, there is a net dipole moment. The molecule is polar.

c. Sulfur dichloride, $S C l_{2}$

Analysis: The molecule has two S-Cl bonds.
Electronegativity Difference ( $Δ E N$ ): $Δ E N = 3.0 - 2.5 = 0.5$
Conclusion: The S-Cl bonds are polar. The molecule has a V-shaped (bent) geometry, which is asymmetrical. The bond dipoles do not cancel each other out, resulting in a net dipole moment. The molecule is polar.

d. Chloromethane, $C H_{3} C l$

Analysis: The molecule has three C-H bonds and one C-Cl bond arranged in a tetrahedral geometry around the central carbon atom.
Electronegativity Difference ( $Δ E N$ ):
- C-Cl: $Δ E N = 3.0 - 2.5 = 0.5$ (Polar)
- C-H: $Δ E N = 2.5 - 2.1 = 0.4$ (Slightly polar)
Conclusion: The molecule is asymmetrical because the atoms bonded to the central carbon are not identical. The strong C-Cl bond dipole is not canceled by the C-H bond dipoles. The molecule is polar.

e. Tetrabromomethane, $C B r_{4}$

Analysis: The molecule has four C-Br bonds arranged in a tetrahedral geometry.
Electronegativity Difference ( $Δ E N$ ): $Δ E N = 2.8 - 2.5 = 0.3$
Conclusion: The C-Br bonds are slightly polar. However, the molecule has a symmetrical tetrahedral shape. The four C-Br bond dipoles are arranged symmetrically and cancel each other out. The molecule is nonpolar.