Molecular polarity describes an uneven distribution of electrical charge within a molecule. This internal charge separation occurs when electrons are not shared equally among the atoms, which results in the molecule having a slightly positive end and a slightly negative end. This separation creates an electric dipole moment, essentially giving the molecule two poles. A molecule must possess this permanent, non-zero dipole moment to be classified as polar. Understanding why dichloromethane (\(text{CH}_2text{Cl}_2\)) has such a moment requires examining its atomic composition and three-dimensional structure.
The Role of Electronegativity in Bond Polarity
The first step in determining a molecule’s polarity involves analyzing the individual bonds connecting the atoms. Polarity in a covalent bond is a consequence of electronegativity, the measure of an atom’s tendency to attract shared electrons toward itself. When two atoms with different electronegativities bond, electrons spend more time closer to the atom with the greater pull, creating partial negative (\(delta^-\)) and positive (\(delta^+\)) charges.
Dichloromethane contains a central carbon atom bonded to two hydrogen atoms and two chlorine atoms, resulting in two distinct types of bonds: \(text{C}-text{H}\) and \(text{C}-text{Cl}\). The electronegativity values for carbon (C), hydrogen (H), and chlorine (Cl) are approximately 2.55, 2.20, and 3.16, respectively. This difference in electron-pulling strength gives each bond a specific degree of polarity.
The difference in electronegativity between carbon and chlorine is \(3.16 – 2.55 = 0.61\) units, making the \(text{C}-text{Cl}\) bonds significantly polar. The highly electronegative chlorine atoms pull electron density away from the central carbon, making the chlorine side partially negative. In contrast, the difference between carbon and hydrogen is \(2.55 – 2.20 = 0.35\) units, which makes the \(text{C}-text{H}\) bonds only weakly polar, with the carbon atom being the slightly negative pole in these bonds. The presence of these polar bonds is necessary for overall molecular polarity, but it is not the sole determining factor.
Understanding the Tetrahedral Shape
The three-dimensional arrangement of atoms is the second determining factor for molecular polarity, as it dictates how individual bond polarities interact. The central carbon atom in \(text{CH}_2text{Cl}_2\) is bonded to four surrounding atoms and has no lone pairs of electrons, forcing the molecule into a tetrahedral geometry. This arrangement positions the four surrounding atoms at the corners of a tetrahedron, maximizing the distance between electron groups.
The ideal bond angle in a perfect tetrahedron, such as methane (\(text{CH}_4\)), is \(109.5^{circ}\). However, \(text{CH}_2text{Cl}_2\) is built with two different types of atoms—larger, more electronegative chlorine atoms and smaller hydrogen atoms—resulting in a slightly distorted tetrahedron. The differing repulsion between the two large chlorine atoms and the two small hydrogen atoms leads to slight variations in the bond angles.
For instance, the angle between the two chlorine atoms (\(angle text{Cl}-text{C}-text{Cl}\)) and the angle between the two hydrogen atoms (\(angle text{H}-text{C}-text{H}\)) are both approximately \(112^{circ}\). Meanwhile, the \(angle text{Cl}-text{C}-text{H}\) angle is slightly reduced to about \(108^{circ}\). This small distortion from the perfect symmetry of methane is a physical manifestation of the molecular asymmetry, which plays a direct role in the final polarity of the molecule. The three-dimensional structure is what prevents the individual forces of polarity from canceling each other out.
Why Dipoles Do Not Cancel Out
The reason \(text{CH}_2text{Cl}_2\) is polar lies in the combination of its two different bond types and its asymmetrical three-dimensional structure. The polarity of each individual bond is represented by a bond dipole moment, a vector quantity having both magnitude (strength of polarity) and direction (pointing toward the more electronegative atom). To find the overall molecular polarity, the individual bond dipole moment vectors must be added together.
In a highly symmetrical molecule like carbon tetrachloride (\(text{CCl}_4\)), which is also tetrahedral, all four \(text{C}-text{Cl}\) bonds are identical. Their dipole moment vectors are arranged perfectly symmetrically, causing them to cancel out completely. This results in a net dipole moment of zero and an overall nonpolar molecule.
However, \(text{CH}_2text{Cl}_2\) is asymmetrical because it has two strong \(text{C}-text{Cl}\) dipoles and two much weaker \(text{C}-text{H}\) dipoles. When the two strong \(text{C}-text{Cl}\) dipoles are added vectorially, they point generally toward the chlorine atoms, concentrating the negative charge on that side of the molecule. The two weaker \(text{C}-text{H}\) dipoles point toward the carbon, leaving the hydrogen atoms on the opposite side slightly positive. Because the four surrounding atoms are not identical, the individual bond dipoles do not balance or cancel each other out, leading to a permanent net dipole moment of \(1.67\) Debye. This non-zero net dipole confirms \(text{CH}_2text{Cl}_2\) is a polar molecule.
Consequences of Being a Polar Molecule
The polarity of dichloromethane affects its physical properties, most notably its ability to act as a versatile solvent. Chemical solubility is governed by the “like dissolves like” principle, meaning polar substances dissolve other polar substances, and nonpolar substances dissolve other nonpolar substances. Dichloromethane’s net dipole moment allows it to effectively interact with and dissolve many other polar compounds, such as certain organic salts and polar organic molecules.
Despite its measurable polarity, the overall dipole moment of \(1.67\) Debye is not as strong as that of highly polar solvents like water (1.85 Debye). This allows it to also dissolve many nonpolar organic compounds. This dual solubility profile means dichloromethane is often used in chemical processes like liquid-liquid extraction to separate compounds from complex mixtures that contain both polar and nonpolar components.