# 3.3 Polar and Nonpolar Covalent Bonds 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 ($\delta^-$). The less electronegative atom consequently acquires a partial positive charge ($\delta^+$). 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: 1. It consists of only nonpolar bonds (e.g., $H_2$, $Cl_2$). 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 ($CO_2$) 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 ($EN = 2.5$) and Oxygen ($EN = 3.5$). This makes each Carbon-Oxygen double bond (C=O) polar. - **Molecular Geometry:** $CO_2$ 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, $CO_2$ 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:** $F = 4.0$, $Cl = 3.0$, $Br = 2.8$, $S = 2.5$, $C = 2.5$, $H = 2.1$ --- **a. Chlorine, $Cl_2$** - **Analysis:** The bond is between two chlorine atoms, which have the same electronegativity. - **Electronegativity Difference ($\Delta EN$):** $$\Delta EN = 3.0 - 3.0 = 0$$ - **Conclusion:** The Cl-Cl bond is a pure (nonpolar) covalent bond. The molecule is **nonpolar**. **b. Hydrogen fluoride, $HF$** - **Analysis:** The bond is between hydrogen and fluorine. - **Electronegativity Difference ($\Delta EN$):** $$\Delta EN = 4.0 - 2.1 = 1.9$$ - **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, $SCl_2$** - **Analysis:** The molecule has two S-Cl bonds. - **Electronegativity Difference ($\Delta EN$):** $$\Delta EN = 3.0 - 2.5 = 0.5$$ - **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, $CH_3Cl$** - **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 ($\Delta EN$):** - C-Cl: $$\Delta EN = 3.0 - 2.5 = 0.5$$ (Polar) - C-H: $$\Delta EN = 2.5 - 2.1 = 0.4$$ (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, $CBr_4$** - **Analysis:** The molecule has four C-Br bonds arranged in a tetrahedral geometry. - **Electronegativity Difference ($\Delta EN$):** $$\Delta EN = 2.8 - 2.5 = 0.3$$ - **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**.