Question

Account for the fact that, among the chlorinated derivatives of methane, chloromethane has the largest dipole moment and tetrachloromethane has the smallest dipole moment. $$ \begin{array}{llc} \hline \text { Name } & \begin{array}{l} \text { Molecular } \\ \text { Formula } \end{array} & \begin{array}{c} \text { Dipole Moment } \\ \text { (debyes, D) } \end{array} \\ \hline \text { Chloromethane } & \mathrm{CH}_{3} \mathrm{Cl} & 1.87 \\ \text { Dichloromethane } & \mathrm{CH}_{2} \mathrm{Cl}_{2} & 1.60 \\ \text { Trichloromethane } & \mathrm{CHCl}_{3} & 1.01 \\ \text { Tetrachloromethane } & \mathrm{CCl}_{4} & 0 \\ \hline \end{array} $$

Short answer

Answer: Chloromethane (CH3Cl) has the largest dipole moment, while tetrachloromethane (CCl4) has the smallest dipole moment.

Step-by-step solution

Understand Molecular Geometry and Polarity

The molecular geometry of a molecule dictates how the individual bond dipoles will add up to form the overall molecular dipole. In the case of methane (CH4) and its chlorinated derivatives, all of them have a tetrahedral molecular geometry, meaning that the hydrogens and chlorines are evenly distributed around the central carbon. Keep in mind that in order to create a molecule's dipole moment, the bond dipoles should not cancel each other out.

Determine Electronegativity Differences

The bond dipole moment is determined by the difference in electronegativity between the two elements forming the bond. In the case of methane and its chlorinated derivatives, we will consider the electronegativity differences between carbon (C) and its bonding partners, hydrogen (H) and chlorine (Cl). The greater the electronegativity difference between carbon and its bonding partner, the greater the bond dipole moment. The electronegativity values for carbon, hydrogen, and chlorine are approximately 2.5, 2.2, and 3.2, respectively. Thus, the C-Cl bond has a much larger bond dipole moment than the C-H bond due to the larger difference in electronegativity.

Analyze the Dipole Moment of Each Chlorinated Derivative

Now, we will consider the dipole moments of the chlorinated methane derivatives, taking into account their molecular geometry and the electronegativity differences between carbon and the bonding partner (H or Cl). 1. Chloromethane (CH3Cl) - In chloromethane, there are three C-H bonds and one C-Cl bond. These dipole moments do not cancel each other out completely due to the asymmetry in the molecule, and the dipole moment of the stronger C-Cl bond is the main contributor to the overall larger dipole moment of 1.87 D. 2. Dichloromethane (CH2Cl2) - In this molecule, there are two C-H bonds and two C-Cl bonds. Due to the molecular geometry, the two C-H bonds are opposite to each other and cancel out their bond dipoles. However, the two C-Cl bonds are also opposite but with a much greater bond dipole moment resulting in a dipole moment of 1.60 D. 3. Trichloromethane (CHCl3) - With one C-H bond and three C-Cl bonds, the dipole moments of the C-Cl bonds decrease due to increased symmetry in the molecule, and the C-H bond's dipole moment is not strong enough to create a large overall dipole moment. As a result, the dipole moment is 1.01 D. 4. Tetrachloromethane (CCl4) - In tetrachloromethane, there are four C-Cl bonds, making it completely symmetric. Since the bond dipoles are identical and arrange symmetrically around the central carbon, they cancel each other out, resulting in an overall dipole moment of 0 D. This explanation accounts for the fact that chloromethane has the largest dipole moment among the chlorinated derivatives of methane, while tetrachloromethane has the smallest dipole moment.

Key concepts

Molecular Geometry

Understanding the molecular geometry of a compound is fundamental to grasping why different molecules have varying dipole moments. In the context of chlorinated methane derivatives, they all share a tetrahedral geometry. This means that the atoms attached to the central carbon atom are positioned at the corners of a tetrahedron, with bond angles approximating 109.5 degrees.

However, even with a tetrahedral framework, the distribution of different atoms around the carbon can lead to asymmetry in the molecule. For instance, chloromethane (CH₃Cl) is not symmetric because the presence of one chlorine atom creates a corner with different electronegativity. This asymmetry prevents complete cancellation of the individual bond dipoles, hence resulting in a measurable dipole moment.

In contrast, when all corners are identical as in tetrachloromethane (CCl₄), symmetry is restored, and the molecule does not possess a net dipole moment. Each C-Cl bond has a dipole, but the symmetrical arrangement allows these dipoles to cancel out completely.

Electronegativity Differences

Electronegativity is the tendency of an atom to attract electrons towards itself. When atoms with different electronegativities form a bond, the shared electrons are not distributed equally. The atom with higher electronegativity will have a stronger pull, resulting in a partial negative charge, while the atom with lower electronegativity will have a partial positive charge.

The different electronegativities of hydrogen (approx. 2.2) and chlorine (approx. 3.2) compared to carbon (approx. 2.5) leads to the creation of bond dipoles. The larger the difference in electronegativity, the greater the bond dipole moment. Chlorine being more electronegative than carbon creates a sizable bond dipole in C-Cl bonds. When assessing chlorinated methanes, this difference is why the C-Cl bond has a stronger dipole moment than the C-H bond.

As there are more chlorine atoms, the bond's dipole moments increase, up to the point where symmetry overrides and results in cancellation of these dipoles. Differing numbers of C-Cl and C-H bonds in chloromethane (CH₃Cl), dichloromethane (CH₂Cl₂), and trichloromethane (CHCl₃) display how the molecular dipole moment gets influenced by the varying number and arrangement of the polar bonds.

Bond Dipole Moment

The bond dipole moment is a vector quantity that has both a magnitude, related to how strong the difference in electronegativity between the two bonded atoms is, and a direction, from the positively charged atom to the negatively charged atom.

Given their tetrahedral shape, the chlorinated methane derivatives demonstrate that even with identical molecular geometry, the molecular dipole moment, which is the vector sum of its bond dipoles, can vary greatly. Chloromethane with its three weaker C-H dipoles and one strong C-Cl dipole has the highest molecular dipole moment because the vectors do not cancel out completely. Dichloromethane and trichloromethane have intermediate dipole moments owing to different vector arrangements and a higher number of stronger C-Cl bond dipoles.

Tetrachloromethane stands as a unique case where four strong bond dipoles of equal magnitude and symmetrically arranged completely cancel out, leading to a zero net molecular dipole moment. These examples highlight the intricate balance between molecular geometry and bond dipole moments, as they collectively influence a molecule's electrical polarity.