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Chapter 9: Problem 16

Describe the bond angles to be found in each of the follow- ing molecular structures: (a) trigonal planar, \((\mathbf{b})\) tetrahedral, (c) octahedral, (d) linear.

Short Answer

Expert verified
The bond angles for each molecular structure are: (a) Trigonal planar: \(120^{\circ}\) (b) Tetrahedral: \(109.5^{\circ}\) (c) Octahedral: \(90^{\circ}\) and \(180^{\circ}\) (d) Linear: \(180^{\circ}\)

Step by step solution

01

Analyze Trigonal Planar Structure

A trigonal planar structure has three atoms bonded to the central atom, with all of them placing in one plane and the electron pairs arranging themselves as far apart as possible. In this structure, the bond angles are equal.
02

Calculate Bond Angle for Trigonal Planar Structure

As there are 3 atoms bonded to the central atom in a trigonal planar structure, and they are all in the same plane, the bond angle between them is \(120^{\circ}\).
03

Analyze Tetrahedral Structure

A tetrahedral structure has four atoms bonded to the central atom. The electron pairs arrange themselves as far apart as possible, creating a three-dimensional structure.
04

Calculate Bond Angle for Tetrahedral Structure

In a tetrahedral structure, the bond angle between the atoms is approximately \(109.5^{\circ}\).
05

Analyze Octahedral Structure

An octahedral structure has six atoms bonded to the central atom. The electron pairs arrange themselves as far apart as possible to minimize repulsion, resulting in a regular three-dimensional structure.
06

Calculate Bond Angle for Octahedral Structure

In an octahedral structure, there are two possible bond angles: 1. The angle between two atoms that are adjacent to each other, which is \(90^{\circ}\). 2. The angle between two atoms opposite each other (linear arrangement), which is \(180^{\circ}\).
07

Analyze Linear Structure

A linear structure has two atoms bonded to the central atom. The electron pairs arrange themselves as far apart as possible, resulting in a linear arrangement.
08

Calculate Bond Angle for Linear Structure

In a linear structure, the bond angle between the atoms is \(180^{\circ}\) as they are in a straight line. To summarize, the bond angles for each molecular structure are as follows: (a) Trigonal planar: \(120^{\circ}\) (b) Tetrahedral: \(109.5^{\circ}\) (c) Octahedral: \(90^{\circ}\) and \(180^{\circ}\) (d) Linear: \(180^{\circ}\)

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Most popular questions from this chapter

(a) If the valence atomic orbitals of an atom are sp hybridized, how many unhybridized \(p\) orbitals remain in the valence shell? How many \(\pi\) bonds can the atom form? (b) Imagine that you could hold two atoms that are bonded together, twist them, and not change the bond length. Would it be easier to twist (rotate) around a single \(\sigma\) bond or around a double \((\sigma\) plus \(\pi)\) bond, or would they be the same?

Place the following molecules and ions in order from smallest to largest bond order: $\mathrm{N}_{2}{ }_{2}^{2+}, \mathrm{He}_{2}{ }^{+}, \mathrm{Cl}_{2} \mathrm{H}_{2}^{-}, \mathrm{O}_{2}{ }^{2-}$.

Shown here are three pairs of hybrid orbitals, with each set at a characteristic angle. For each pair, determine the type of hybridization, if any, that could lead to hybrid orbitals at the specified angle.

Antibonding molecular orbitals can be used to make bonds to other atoms in a molecule. For example, metal atoms can use appropriate \(d\) orbitals to overlap with the \(\pi_{2}^{*}\), orbitals of the carbon monoxide molecule. This is called \(d-\pi\) backbonding. (a) Draw a coordinate axis system in which the \(y\) -axis is vertical in the plane of the paper and the \(x\) -axis horizontal. Write ${ }^{4} \mathrm{M}^{\prime \prime}$ at the origin to denote a metal atom. (b) Now, on the \(x\) -axis to the right of \(\mathrm{M}\), draw the Lewis structure of a CO molecule, with the carbon nearest the \(\mathrm{M}\). The CO bond axis should be on the \(x\) -axis. (c) Draw the \(\mathrm{CO} \pi_{2 p}^{*}\) orbital, with phases (see the "Closer Look" box on phases) in the plane of the paper. Two lobes should be pointing toward M. (d) Now draw the \(d_{x y}\) orbital of \(\mathrm{M}\), with phases. Can you see how they will overlap with the \(\pi_{2 p}^{*}\) orbital of $\mathrm{CO} ?\( (e) What kind of bond is being made with the orbitals between \)\mathrm{M}$ and \(\mathrm{C}, \sigma\) or \(\pi ?\) (f) Predict what will happen to the strength of the CO bond in a metal-CO complex compared to CO alone.

How many nonbonding electron pairs are there in each of the following molecules: (a) $\mathrm{N}\left(\mathrm{CH}_{3}\right)_{3},(\mathbf{b}) \mathrm{CO},(\mathbf{c}) \mathrm{BF}_{3},$ (d) \(\mathrm{SO}_{2} ?\)
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