/*! 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 34 (a) What conditions must be met ... [FREE SOLUTION] | 91影视

91影视

Log out

Learning Materials

Features

Discover

Chapter 9: Problem 34

(a) What conditions must be met if a molecule with polar bonds is nonpolar? (b) What geometries will give nonpolar molecules for \(\mathrm{AB}_{2}, \mathrm{AB}_{3}\), and \(\mathrm{AB}_{4}\) geometries?

Short Answer

Expert verified
(a) For a molecule with polar bonds to be nonpolar, it must have a symmetrical distribution of bond dipoles, resulting in the vector sum of all bond dipoles being zero. (b) The nonpolar geometries for AB_2, AB_3, and AB_4 molecules are linear, trigonal planar, and tetrahedral, respectively.

Step by step solution

01

Identify conditions for a molecule with polar bonds to be nonpolar

For a molecule with polar bonds to be nonpolar, the bond dipoles must cancel each other out. This can happen if the molecule has a symmetrical distribution of these bond dipoles. In other words, the molecule must have a geometry where the vector sum of all the bond dipoles is zero.
02

Determine geometry for nonpolar AB2 molecules

For an AB_2 molecule to be nonpolar, the bond dipoles must cancel each other out. This can happen when the A atom is in the center and the two B atoms are in a linear orientation. The bond dipoles will be equal and opposite, thus canceling each other out. So, the nonpolar geometry for an AB_2 molecule is linear.
03

Determine geometry for nonpolar AB3 molecules

For an AB_3 molecule to be nonpolar, the bond dipoles must cancel each other out. This can happen when the A atom is in the center and the three B atoms are in a trigonal planar orientation. The bond dipoles will be equal in magnitude and in a symmetrical distribution, thus canceling each other out. So, the nonpolar geometry for an AB_3 molecule is trigonal planar.
04

Determine geometry for nonpolar AB4 molecules

For an AB_4 molecule to be nonpolar, the bond dipoles must cancel each other out. This can happen when the A atom is in the center and the four B atoms are in a tetrahedral orientation. The bond dipoles will be equal in magnitude and in a symmetrical distribution, thus canceling each other out. So, the nonpolar geometry for an AB_4 molecule is tetrahedral.

Unlock Step-by-Step Solutions & Ace Your Exams!

  • Full Textbook Solutions

    Get detailed explanations and key concepts

  • Unlimited Al creation

    Al flashcards, explanations, exams and more...

  • Ads-free access

    To over 500 millions flashcards

  • Money-back guarantee

    We refund you if you fail your exam.

Over 30 million students worldwide already upgrade their learning with 91影视!

Key Concepts

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

Polar Bonds
Polar bonds occur when two atoms share a pair of electrons, but the sharing is not even. One atom pulls the electrons closer due to its higher electronegativity, creating a difference in electric charge. This uneven distribution of charge creates what we call a dipole moment. It is essential to remember that just because a molecule has polar bonds, it doesn't necessarily mean the whole molecule is polar. There are other factors at play, like the molecule's shape, that can influence the overall polarity. Understanding polar bonds is crucial for predicting how a molecule will interact with others, influencing properties such as solubility and boiling points.
Molecular Geometry
Molecular geometry describes the three-dimensional arrangement of atoms within a molecule. It is a significant factor in determining whether a molecule with polar bonds is overall polar or nonpolar. Specific geometries allow the bond dipoles to cancel each other out, resulting in a nonpolar molecule despite the presence of polar bonds. For example:
  • Linear geometry (e.g., AB2) can lead to nonpolar molecules if the dipoles are exactly opposite.
  • Trigonal planar geometry (e.g., AB3) can be nonpolar when dipoles are evenly distributed at 120 degrees.
  • Tetrahedral geometry (e.g., AB4) achieves nonpolarity when the dipoles are symmetrically arranged around the central atom.
Thus, understanding molecular geometry helps in predicting whether the molecule's polar bonds cancel out or not.
Bond Dipoles
A bond dipole arises when electrons in a bond are unevenly shared between two atoms, creating a miniature electric "push and pull" along the bond. This dipole can be visualized as an arrow pointing toward the more electronegative atom. It's like having a little battery in each bond where one side is slightly positive, and the other is slightly negative. These bond dipoles can add together to affect the whole molecule's polarity. If bond dipoles in a molecule point in different directions and do not cancel out, the molecule is polar as a whole. In contrast, if they cancel out, the molecule becomes nonpolar overall, regardless of having polar bonds.
Symmetrical Distribution
Symmetrical distribution in molecules with polar bonds is key to achieving nonpolarity. When the bond dipoles distribute uniformly around the central atom, they can effectively cancel each other out. In such arrangements, the vector of each dipole is balanced by another, leading to a net dipole moment of zero. This means the molecule does not have a positive or negative side, appearing neutral in terms of electric charge. Symmetry is crucial in determining the overall nonpolar nature in molecules such as CO2, where linear geometry ensures opposite dipoles counterbalance each other, resulting in no net dipole moment.

One App. One Place for Learning.

All the tools & learning materials you need for study success - in one app.

Get started for free

Most popular questions from this chapter

How many nonbonding electron pairs are there in each of the following molecules: (a) \(\left(\mathrm{CH}_{3}\right)_{2} \mathrm{~S} ;\) (b) \(\mathrm{HCN}\); (c) \(\mathrm{H}_{2} \mathrm{C}_{2} ;\) (d) \(\mathrm{CH}_{3} \mathrm{~F}\) ?

For both atoms and molecules, ionization energies (Section 7.4) are related to the energies of orbitals: The lower the energy of the orbital, the greater the ionization energy. The first ionization energy of a molecule is therefore a measure of the energy of the highest occupied molecular orbital (HOMO). See the "Chemistry Put to Work" box on Orbitals and Energy. The first ionization energies of several diatomic molecules are given in electron-volts in the following table: $$ \begin{array}{ll} \hline \text { Molecule } & I_{1}(\mathrm{eV}) \\ \hline \mathrm{H}_{2} & 15.4 \\ \mathrm{~N}_{2} & 15.6 \\ \mathrm{O}_{2} & 12.1 \\ \mathrm{~F}_{2} & 15.7 \\ \hline \end{array} $$ (a) Convert these ionization energies to \(\mathrm{kJ} / \mathrm{mol}\). (b) On the same plot, graph \(I_{1}\) for the \(\mathrm{H}, \mathrm{N}, \mathrm{O}\), and \(\mathrm{F}\) atoms (Figure 7.11) and \(I_{1}\) for the molecules listed. (c) Do the ionization energies of the molecules follow the same periodic trends as the ionization energies of the atoms? (d) Use molecular orbital energy-level diagrams to explain the trends in the ionization energies of the molecules.

(a) Explain why \(\mathrm{BrF}_{4}^{-}\) is square planar, whereas \(\mathrm{BF}_{4}^{-}\) is tetrahedral. (b) Water, \(\mathrm{H}_{2} \mathrm{O}\), is a bent molecule. Predict the shape of the molecular ion formed from the water molecule if you were able to remove four electrons to make \(\left(\mathrm{H}_{2} \mathrm{O}\right)^{4+}\).

Ethyl acetate, \(\mathrm{C}_{4} \mathrm{H}_{8} \mathrm{O}_{2}\), is a fragrant substance used both as a solvent and as an aroma enhancer. Its Lewis structure is (a) What is the hybridization at each of the carbon atoms of the molecule? (b) What is the total number of valence electrons in ethyl acetate? (c) How many of the valence electrons are used to make \(\sigma\) bonds in the molecule? (d) How many valence electrons are used to make \(\pi\) bonds? (e) How many valence electrons remain in nonbonding pairs in the molecule?

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]
See all solutions

Recommended explanations on Chemistry Textbooks

Chemical Analysis

Read Explanation

Chemistry Branches

Read Explanation

Nuclear Chemistry

Read Explanation

Chemical Reactions

Read Explanation

Physical Chemistry

Read Explanation

Ionic and Molecular Compounds

Read Explanation
View all explanations

What do you think about this solution?

We value your feedback to improve our textbook solutions.

Study anywhere. Anytime. Across all devices.