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Chapter 9: Problem 35

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

Short Answer

Expert verified
The distinction between a bond dipole and a molecular dipole moment lies in their scope and relationship to molecular polarity. A bond dipole refers to the separation of charges within a single chemical bond due to differences in electronegativity between two atoms, while a molecular dipole moment is the vector sum of all the bond dipoles in a molecule, representing the overall polarity of the molecule. The presence of bond dipoles does not necessarily guarantee a molecular dipole moment, as bond dipoles can cancel each other out if they have opposite directions and equal magnitudes.

Step by step solution

01

Definition of Bond Dipole

A bond dipole occurs when there is a separation of charges between two atoms within a chemical bond. This is due to the difference in electronegativity between the two atoms. The more electronegative atom attracts the shared electrons, creating a partially negative charge on that atom, and leaving the other atom with a partially positive charge. The bond dipole is represented by an arrow pointing towards the more electronegative atom, with a cross at the tail representing the positive charge.
02

Definition of Molecular Dipole Moment

The molecular dipole moment is the vector sum of all the bond dipoles in a molecule. It represents the overall polarity of the molecule, accounting for the magnitude and direction of all bond dipoles. The molecular dipole moment is expressed in Debye units (D). A molecule with a net molecular dipole moment is said to be polar, whereas a molecule with no net molecular dipole moment (i.e., the bond dipoles cancel each other out) is considered nonpolar.
03

Comparing Bond Dipole and Molecular Dipole Moment

The distinction between bond dipole and molecular dipole moment lies in their scope and relationship to molecular polarity. A bond dipole considers the separation of charges within a single bond, while a molecular dipole moment takes into account the polarity of the entire molecule, accounting for the impact of all bond dipoles. It is important to note that the presence of bond dipoles does not necessarily guarantee a molecular dipole moment, as bond dipoles can cancel each other out if they have opposite directions and equal magnitudes.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Bond Dipole
A bond dipole arises within a chemical bond when two atoms have different tendencies to attract electrons. This tendency is known as electronegativity. When one atom is more electronegative than its neighbor, it pulls the shared electrons closer, acquiring a partial negative charge. Meanwhile, the less electronegative atom ends up with a partial positive charge. To visualize a bond dipole, think of an arrow pointing from the positive atom to the negative atom, indicating the direction of the electron pull. At the tail of this arrow, a '+' symbolizes the positive atom. This representation helps us understand the partial charges within the bond, but it doesn't explain how these charges might affect the entire molecule's polarity.
Electronegativity
Electronegativity is a chemical property that explains how an atom differs in its ability to attract electrons in a bond. When examining the periodic table, electronegativity increases across a period from left to right and decreases down a group. This means that atoms like fluorine and oxygen are highly electronegative, while those like cesium and francium are not. When two bonded atoms have different electronegativities, the more electronegative atom attracts the shared electrons more strongly. This difference leads to the formation of a bond dipole. A significant electronegativity difference means there will be a noticeable partial charge on both atoms, as one becomes partially negative and the other partially positive. Understanding electronegativity is key to predicting where bond dipoles form and how strong they are.
Polarity
Polarity in molecules depends on both the individual bond dipoles and the overall shape of the molecule. A molecule can be polar or nonpolar based on how the bond dipoles interact. For instance, if all the bond dipoles point in a similar direction and don鈥檛 cancel each other out, the molecule exhibits an overall molecular dipole moment, making it polar. However, if bond dipoles are symmetrically arranged and cancel each other, the molecule remains nonpolar despite having individual bond dipoles. The net molecular dipole moment is the vector sum of all bond dipoles in the molecule. It's important to consider molecular geometry when predicting polarity because certain symmetrical shapes can lead to a nonpolar molecule even if the bonds themselves are polar.

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Most popular questions from this chapter

(a) Is the molecule \(\mathrm{BF}_{3}\) polar or nonpolar? (b) If you react \(\mathrm{BF}_{3}\) to make the ion \(\mathrm{BF}_{3}^{2-}\), is this ion planar? (c) Does the molecule \(\mathrm{BF}_{2} \mathrm{Cl}\) have a dipole moment?

(a) Boron trichloride \(\left(\mathrm{BCl}_{3}\right)\) and the carbonate ion \(\left(\mathrm{CO}_{3}^{2-}\right)\) are both described as trigonal. What does this indicate about their bond angles? (b) The \(\mathrm{PCl}_{3}\) molecule is trigonal pyramidal, while \(\mathrm{ICl}_{3}\) is T-shaped. Which of these molecules is flat?

The following is part of a molecular orbital energy-level diagram for MOs constructed from 1 s atomic orbitals. (a) What labels do we use for the two MOs shown? (b) For which of the following molecules or ions could this be the energy-level diagram: $$ \mathrm{H}_{2} \mathrm{He}_{2}, \mathrm{H}_{2}^{+}, \mathrm{He}_{2}^{+}, \mathrm{or} \mathrm{H}_{2}^{-} ? $$ (c) What is the bond order of the molecule or ion? (d) If an electron is added to the system, into which of the MOs will it be added? [Section 9.7\(]\)

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

The highest occupied molecular orbital of a molecule is abbreviated as the HOMO. The lowest unoccupied molecular orbital in a molecule is called the LUMO. Experimentally, one can measure the difference in energy between the HOMO and LUMO by taking the electronic absorption (UV-visible) spectrum of the molecule. Peaks in the electronic absorption spectrum can be labeled as \(\pi_{2 \mathrm{p}}-\pi_{2 \mathrm{p}}{ }^{*}\), \(\sigma_{2 s}-\sigma_{2 s}{ }^{*},\) and so on, corresponding to electrons being promoted from one orbital to another. The HOMO-LUMO transition corresponds to molecules going from their ground state to their first excited state. (a) Write out the molecular orbital valence electron configurations for the ground state and first excited state for \(\mathrm{N}_{2}\). (b) Is \(\mathrm{N}_{2}\) paramagnetic or diamagnetic in its first excited state? (c) The electronic absorption spectrum of the \(\mathrm{N}_{2}\) molecule has the lowest energy peak at \(170 \mathrm{nm}\). To what orbital transition does this correspond? (d) Calculate the energy of the HOMO-LUMO transition in part (a) in terms of \(\mathrm{kJ} / \mathrm{mol}\). (e) Is the \(\mathrm{N}-\mathrm{N}\) bond in the first excited state stronger or weaker compared to that in the ground state?
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