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Chapter 9: Problem 37

What is the distinction between a bond dipole and a molecular dipole moment?

Short Answer

Expert verified
A bond dipole is the electric dipole moment within a single bond, resulting from the difference in electronegativity between two bonded atoms, while a molecular dipole moment is the overall dipole moment of the entire molecule, obtained as the vector sum of all individual bond dipoles. Bond dipoles exist due to charge separation within a bond, while molecular dipole moments depend on both the bond dipoles and the molecular geometry, which may result in a net zero molecular dipole moment even when individual bond dipoles are present.

Step by step solution

01

Definition of bond dipole

A bond dipole is created when two atoms in a molecule have a difference in electronegativity. The more electronegative atom attracts the electrons in the bond, creating a partial negative charge, denoted as "未-", while the less electronegative atom has a partial positive charge, represented by "未+". This charge separation leads to the formation of an electric dipole, with the dipole moment \(\vec{\mu}\) pointing from the positive to the negative charge. It is represented as: \[ \vec{\mu} = \delta q \cdot \vec{d} \] Where \(\delta q\) is the magnitude of the partial charge and \(\vec{d}\) is the vector distance between the atoms.
02

Definition of molecular dipole moment

The molecular dipole moment is the vector sum of all bond dipoles in the molecule. In other words, it is the net dipole moment of the entire molecule, considering the geometry of the molecule and the distribution of partial charges. It is calculated as: \[ \vec{\mu}_{molecule} = \sum_{i} \vec{\mu}_{i} \] Where \(\vec{\mu}_{i}\) represents each bond dipole in the molecule.
03

Distinction between bond dipole and molecular dipole moment

Bond dipole refers to the electric dipole moment that exists within a single bond, resulting from the difference in electronegativity between the two bonded atoms. In contrast, the molecular dipole moment is the overall dipole moment of the entire molecule, which is the vector sum of all individual bond dipoles. A molecular dipole moment may end up being null if the bond dipoles present in the molecule cancel each other out due to the molecular geometry, even if individual bond dipoles exist.

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Most popular questions from this chapter

(a) Does \(\mathrm{SCl}_{2}\) have a dipole moment? If so, in which direction does the net dipole point? (b) Does \(\mathrm{BeCl}_{2}\) have a dipole moment? If so, in which direction does the net dipole point?

(a) Starting with the orbital diagram of a sulfur atom, describe the steps needed to construct hybrid orbitals appropriate to describe the bonding in \(\mathrm{SF}_{2}\). (b) What is the name given to the hybrid orbitals constructed in (a)? (c) Sketch the large lobes of these hybrid orbitals. (d) Would the hybridization scheme in part (a) be appropriate for \(\mathrm{SF}_{4} ?\) Explain.

What is the difference between the electron-domain geometry and the molecular geometry of a molecule? Use the water molecule as an example in your discussion. Why do we need to make this distinction?

The \(\mathrm{PF}_{3}\) molecule has a dipole moment of \(1.03 \mathrm{D}\), but \(\mathrm{BF}_{3}\) has a dipole moment of zero. How can you explain the difference?

Predict whether each of the following molecules is polar or nonpolar: (a) \(\mathrm{CCl}_{4},\) (b) \(\mathrm{NH}_{3}\), (c) \(\mathrm{SF}_{4}\), (d) \(\mathrm{XeF}_{4}\), (e) \(\mathrm{CH}_{3} \mathrm{Br}\), (f) \(\mathrm{GaH}_{3}\)
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