Question

Among the following, the molecule with the highest dipole moment is: (a) \(\mathrm{CH}_{3} \mathrm{Cl}\) (b) \(\mathrm{CH}_{2} \mathrm{Cl}_{2}\) (c) \(\mathrm{CHCl}_{3}\) (d) \(\mathrm{CCl}_{4}\)

Short answer

(b) CH2Cl2 has the highest dipole moment.

Step-by-step solution

Identify the Presence of Polar Bonds

A molecule's dipole moment depends on both bond dipoles and molecular geometry. Begin by identifying the polar bonds in each option. All these molecules have C-Cl bonds, which are polar because of the difference in electronegativity between C and Cl. However, the magnitude and direction of these bond dipoles will affect the net dipole moment of the entire molecule.

Consider Molecular Geometry

The net dipole moment also depends on the molecule's geometry. - In CH4Cl, it has tetrahedral geometry, and the presence of three hydrogen atoms does not cancel out the chlorine's pull completely. - CH2Cl2 has a symmetric tetrahedral shape, but the atoms are not the same, so it retains a net dipole moment. - CHCl3 divides the electronegative force between three chlorines, reducing the net dipole moment. - CCl4, despite having polar C-Cl bonds, has a perfectly symmetric tetrahedral shape, leading to dipole cancellation.

Analyze Dipole Cancellation

In tetrahedral molecules, if all surrounding atoms are the same or symmetrically arranged, their bond dipoles cancel. Check each molecule: - CH3Cl: Asymmetrical, leading to a significant net dipole moment. - CH2Cl2: Asymmetrically substituted, leading to a significant net dipole moment, more so than CHCl3. - CHCl3: Asymmetrical, but with three chlorines pulling in different directions, reducing its overall dipole moment. - CCl4: Symmetrical, resulting in the dipole moments canceling out completely to zero.

Determine the Highest Dipole Moment

Compare the net dipole moments based on molecular geometry and bond dipoles: - CH2Cl2's geometry creates the highest net dipole moment as it has more polar bonds without perfect cancelation, compared to CH3Cl and CHCl3. Thus, CH2Cl2 has the highest dipole moment of the given options.

Key concepts

Molecular Geometry

Molecular geometry is like the blueprint of a molecule. It tells us how atoms are arranged in a three-dimensional space. This arrangement plays a pivotal role in determining how molecules behave, especially in terms of polarity and interactions. For instance, the shape influences how individual bond dipoles (the electric charge distribution between two bonded atoms) add up.
In tetrahedral molecules, such as those discussed in the original exercise, each carbon atom is at the center, with four bonds arranged symmetrically around it. This geometry can lead to either canceling or reinforcing dipole moments, depending on the types of atoms attached.

• Tetrahedral Geometry: Molecules like \( \text{CH}_4\text{Cl} \) have a tetrahedral shape. However, the presence of different atoms (like hydrogen and chlorine) means the symmetrical arrangement is disrupted, leading to a potential overall dipole moment.
• Asymmetry Effect: In molecules like \( \text{CH}_2\text{Cl}_2 \), the different atoms cause a lack of symmetry, preventing the dipoles from canceling out completely, and contributing to a higher dipole moment.

Understanding molecular geometry helps predict if a molecule will have a dipole moment and how strong it might be.

Polar Bonds

Polar bonds occur when there is an unequal sharing of electrons between two atoms in a covalent bond. This asymmetry is often due to a difference in electronegativity between the bonded atoms. The greater the difference, the more polar the bond.
Considering the molecules in the original exercise, all of them contain carbon-chlorine bonds which are inherently polar. This is because chlorine is much more electronegative than carbon, pulling the shared electrons closer to itself and creating a dipole moment.

• Significance of Polar Bonds: In general, the presence of polar bonds does not guarantee a polar molecule. It all depends on the molecular geometry and whether these individual dipole moments cancel each other out or add up.
• Examining Polar Bonds: For instance, in \( \text{CCl}_4 \), despite having polar C-Cl bonds, the overall molecule is nonpolar due to the symmetric tetrahedral shape which leads to the dipole moment canceling out.

Thus, even though polar bonds are the starting point for discussing molecular polarity, understanding their orientation within the molecule is crucial for determining the net dipole moment.

Electronegativity

Electronegativity is like the "tug-of-war" strength of an atom for electrons in a chemical bond. It's a measure of how strongly an atom can attract and hold onto shared electrons. Atoms with high electronegativity, like chlorine, exert a strong pull on electrons, influencing the molecule's properties.
When discussing molecules such as \( \text{CH}_2\text{Cl}_2 \), electronegativity differences between carbon and chlorine are what create polar bonds. Chlorine, with its higher electronegativity, attracts the electrons more, making the bonds polar, which can lead to a net dipole moment depending on the molecular geometry.

• Understanding Electronegativity Differences: The greater the difference in electronegativity between two bonded atoms, the more polar the bond will be.
• Practical Implications: The electronegativity of atoms involved in a molecule determines the polarity and ultimately the behavior and interaction of molecules. In solvents like \( \text{CH}_2\text{Cl}_2 \), this affects solubility and reactivity.

Electronegativity is the cornerstone for understanding why molecules have dipole moments and influences many chemical properties of the substances we encounter daily.