Question

The highest dipole moment is of: (a) \(\mathrm{CF}_{4}\) (b) \(\mathrm{CH}_{3} \mathrm{OH}\) (c) \(\mathrm{CO}_{2}\) (d) \(\mathrm{CH}_{3} \mathrm{~F}\)

Short answer

CH_{3}OH has the highest dipole moment.

Step-by-step solution

Understanding Dipole Moment

A dipole moment occurs when there is a separation of electric charge between two atoms in a molecule. It's a vector quantity having both magnitude and direction and is represented as eddbrairedarrow02.jpg. Polar molecules usually have a net dipole moment, while non-polar molecules do not.

Identify Molecular Geometry

Examine the molecular geometry of each compound. For instance, CF_{4} is tetrahedrally symmetric, CO_{2} is linear, CH_{3}  and CH_{3}F are similar in shape but CH_{3}  is symmetric while CH_{3}F has a more electronegative fluorine atom at one end.

Calculating Net Dipole Moment

Evaluate the symmetry and electronegativity of atoms in each molecule. CF_{4} is non-polar and has no net dipole moment, CO_{2} is also non-polar despite having polar bonds due to linear shape, CH_{3}OH has a significant net dipole due to the difference in electronegativity between O and H in the OH group, and CH_{3}F has a dipole moment because of the polar C-F bond.

Compare and Conclude

CH_{3}OH and CH_{3}F both have significant dipole moments. However, CH_{3}OH has a higher dipole moment due to the highly polar OH group, which contributes significantly to the molecule's overall polarity

Key concepts

Polar and Non-Polar Molecules

A molecule’s polarity determines whether it is polar or non-polar. This is based on how the electrons are distributed across the molecule. A polar molecule has an uneven distribution of electrons, which creates partial positive and negative charges within the molecule. This results in a net dipole moment, where the molecule is positively charged on one side and negatively charged on the other.

In contrast, a non-polar molecule either has no electron density difference or its individual dipole moments cancel each other out. For example, in carbon dioxide (\(\mathrm{CO}_2\)), even though the individual \(\mathrm{C=O}\) bonds are polar, the molecule's linear shape means the dipoles cancel out, resulting in a non-polar molecule.

• Polar Molecules: Have a net dipole moment. Examples are \(\mathrm{CH}_3\mathrm{OH}\) and \(\mathrm{CH}_3\mathrm{F}\).
• Non-Polar Molecules: Have no net dipole moment. Examples include \(\mathrm{CF}_4\) and \(\mathrm{CO}_2\).

This distinction is crucial in determining how molecules interact with each other, affecting properties such as solubility, boiling point, and intermolecular forces.

Molecular Geometry

The geometry of a molecule plays a vital role in its polarity and properties. Molecular geometry is determined by the arrangement of atoms in a molecule and can influence how dipoles within the molecule add up to form a net dipole moment.

Tetrahedral geometry like in \(\mathrm{CF}_4\) often results in non-polar molecules if all peripheral atoms are identical, as the individual dipole vectors cancel out due to symmetry. Linear molecules like \(\mathrm{CO}_2\) can also be non-polar if the dipole moments across bonds are symmetrical and cancel each other out.

• Tetrahedral: Common when there are four groups of electrons around a central atom. Example: \(\mathrm{CF}_4\).
• Linear: Bonds or lone pairs are 180° apart. Example: \(\mathrm{CO}_2\).
• Trigonal Pyramidal: Results when a central atom is bonded to three atoms and one lone pair, often leading to a polar molecule. Example: \(\mathrm{NH}_3\).

A molecule like \(\mathrm{CH}_3\mathrm{OH}\) and \(\mathrm{CH}_3\mathrm{F}\), which have different atoms bonded asymmetrically, often present a net dipole moment due to their geometry.

Electronegativity and Bond Polarity

Electronegativity measures an atom's ability to attract electrons within a chemical bond. When two atoms with different electronegativities are bonded together, the electrons tend to be closer to the more electronegative atom, creating a polar bond.

This difference in electronegativity dictates whether a bond is polar. For instance, in \(\mathrm{CH}_3\mathrm{F}\), fluorine is much more electronegative than carbon, resulting in a strong dipole moment pointing towards the fluorine. This makes \(\mathrm{CH}_3\mathrm{F}\) polar.

• High Electronegativity Difference: Greater than 0.5 is usually considered polar.
• Fluorine, Oxygen, and Nitrogen are among the most electronegative elements.
• Polar bonds contribute to the overall dipole moment of the molecule.

Bond polarity along with molecular geometry together contribute to determining the dipole moment of the entire molecule. This, in turn, affects how the molecule behaves in chemical reactions and interactions with other molecules.