/*! This file is auto-generated */ .wp-block-button__link{color:#fff;background-color:#32373c;border-radius:9999px;box-shadow:none;text-decoration:none;padding:calc(.667em + 2px) calc(1.333em + 2px);font-size:1.125em}.wp-block-file__button{background:#32373c;color:#fff;text-decoration:none} Problem 81 The \(\mathrm{PF}_{3}\) molecule... [FREE SOLUTION] | 91影视

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Chapter 9: Problem 81

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?

Short Answer

Expert verified
The difference in dipole moments between PF3 and BF3 molecules can be explained by their molecular structures and electronegativity differences. PF3 has a trigonal pyramidal structure, and its bond dipoles do not cancel out, resulting in a dipole moment of \(1.03 D\). In contrast, BF3 has a trigonal planar structure, where the bond dipoles are symmetrically distributed and cancel each other out, leading to a dipole moment of zero.

Step by step solution

01

Understand Dipole Moment

A dipole moment is a measure of the separation of positive and negative charges in a molecule, resulting from the difference in electronegativity between the atoms. If the electronegativity difference between two bonded atoms is significant, the bond will have a dipole moment, and if the molecule has an asymmetric distribution of these dipoles, the molecule as a whole will have a nonzero dipole moment.
02

Consider Molecular Structure

The molecular structures of PF3 and BF3 are both based on a central atom (P and B, respectively) with three F atoms bonded to it. PF3 has a trigonal pyramidal structure, while BF3 has a trigonal planar structure. In the trigonal pyramidal structure of PF3, the three F atoms all lie below the plane of the central P atom. In the trigonal planar structure of BF3, the three F atoms lie in the plane with the central B atom, at 120-degree angles from each other.
03

Analyze Electronegativity Differences

In a molecule, the difference in electronegativity between the central atom and the outer atoms results in an uneven distribution of electron density, creating bond dipoles. In both PF3 and BF3, the electronegativity of the F atoms is significantly greater than that of the central atoms, P and B, respectively. This will result in bond dipoles pointing towards the F atoms.
04

Determine the Molecular Dipole Moment

Now we can analyze the dipole moment for each molecule. In PF3, the molecular shape is trigonal pyramidal, and the bond dipoles do not cancel each other out. Therefore, PF3 has a nonzero dipole moment, which is found to be \(1.03 D\). In BF3, the molecular shape is trigonal planar. The bond dipoles are symmetrically distributed, with each bond dipole oriented at a 120-degree angle from the other two. In this case, the bond dipoles cancel each other out, resulting in a net dipole moment of zero.
05

Explain the Difference in Dipole Moments

The difference in dipole moments between the PF3 and BF3 molecules can be explained by their respective molecular structures and the electronegativity differences between their atoms. The trigonal pyramidal structure of PF3 leads to a nonzero dipole moment, while the trigonal planar structure of BF3 causes the bond dipoles to cancel each other out, resulting in a dipole moment of zero.

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Most popular questions from this chapter

Consider the \(\mathrm{H}_{2}{ }^{+}\) ion. (a) Sketch the molecular orbitals of the ion, and draw its energy-level diagram. (b) How many electrons are there in the \(\mathrm{H}_{2}{ }^{+}\) ion? (c) Draw the electron configuration of the ion in terms of its MOs. (d) What is the bond order in \(\mathrm{H}_{2}{ }^{+}\) ? (e) Suppose that the ion is excited by light so that an electron moves from a lower-energy to a higher-energy MO. Would you expect the excitedstate \(\mathrm{H}_{2}{ }^{+}\) ion to be stable or to fall apart? Explain.

You can think of the bonding in the \(\mathrm{Cl}_{2}\) molecule in several ways. For example, you can picturethe Cl- -Cl bond containing two electrons that each come from the \(3 p\) orbitals of a \(\mathrm{Cl}\) atom that are pointing in the appropriate direction. However, you can also think about hybrid orbitals. (a) Draw the Lewis structure of the \(\mathrm{Cl}_{2}\) molecule. (b) What is the hybridization of each Cl atom? (c) What kind of orbital overlap, in this view, makes the \(\mathrm{Cl}-\mathrm{Cl}\) bond? (d) Imagine if you could measure the positions of the lone pairs of electrons in \(\mathrm{Cl}_{2}\). How would you distinguish between the atomic orbital and hybrid orbital models of bonding using that knowledge? (e) You can also treat \(\mathrm{Cl}_{2}\) using molecular orbital theory to obtain an energy level diagram similar to that for \(\mathrm{F}_{2}\). Design an experiment that could tell you if the MO picture of \(\mathrm{Cl}_{2}\) is the best one, assuming you could easily measure bond lengths, bond energies, and the light absorption properties for any ionized species.

(a) Draw Lewis structures for ethane \(\left(\mathrm{C}_{2} \mathrm{H}_{6}\right)\), ethylene \(\left(\mathrm{C}_{2} \mathrm{H}_{4}\right)\), and acetylene \(\left(\mathrm{C}_{2} \mathrm{H}_{2}\right)\). (b) What is the hybridization of the carbon atoms in each molecule? (c) Predict which molecules, if any, are planar. (d) How many \(\sigma\) and \(\pi\) bonds are there in each molecule? (e) Suppose that silicon could form molecules that are precisely the analogs of ethane, ethylene, and acetylene. How would you describe the bonding about Si in terms of hydrid orbitals? Does it make a difference that Si lies in the row below \(\mathrm{C}\) in the periodic table? Explain.

(a) What is the difference between a localized \(\pi\) bond and a delocalized one? (b) How can you determine whether a molecule or ion will exhibit delocalized \(\pi\) bonding? (c) Is the \(\pi\) bond in \(\mathrm{NO}_{2}^{-}\) localized or delocalized?

The molecule shown here is difluoromethane \(\left(\mathrm{CH}_{2} \mathrm{~F}_{2}\right)\), which is used as a refrigerant called \(\mathrm{R}\) -32. (a) Based on the structure, how many electron domains surround the \(C\) atom in this molecule? (b) Would the molecule have a nonzero dipole moment? (c) If the molecule is polar, in what direction will the overall dipole moment vector point in the molecule? [Sections \(9.2\) and 9.3]
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