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Chapter 9: Problem 14

Describe the bond angles to be found in each of the following molecular structures: (a) planar trigonal, (b) tetrahedral, (c) octahedral, (d) linear.

Short Answer

Expert verified
In each molecular structure, the bond angles are as follows: (a) Planar trigonal: all bond angles are 120°, as the atoms lie in one plane forming an equilateral triangle. (b) Tetrahedral: bond angles are approximately 109.5°, minimizing repulsion between electron pairs. (c) Octahedral: bond angles are 90° between adjacent atoms and 180° between opposite atoms. (d) Linear: the bond angle is 180°, with all atoms aligned in a straight line.

Step by step solution

01

(Planar Trigonal Bond Angles)

In a planar trigonal (or trigonal planar) molecular structure, there are three atoms bonded to a central atom, and all atoms lie in one plane. The bond angles are determined by drawing lines between the atoms and the central atom, forming an equilateral triangle. In this structure, all bond angles are equal to 120°.
02

(Tetrahedral Bond Angles)

In a tetrahedral molecular structure, there are four atoms bonded to a central atom. The structure is shaped like a tetrahedron, where all vertices (atoms) are equidistant from the central atom. The bond angles in this structure are approximately 109.5°, because this angle minimizes repulsion between the electron pairs around the central atom.
03

(Octahedral Bond Angles)

In an octahedral molecular structure, there are six atoms bonded to a central atom. The structure is shaped like an octahedron, with the central atom at its center. The bond angles in this structure are as follows: - Between any two adjacent atoms (vertices), the bond angle is 90°. - Between any two opposite atoms (vertices), the bond angle is 180°.
04

(Linear Bond Angles)

In a linear molecular structure, there are two atoms bonded to a central atom, and all three atoms are aligned in a straight line. In this case, the bond angle between the two atoms and the central atom is 180°.

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Most popular questions from this chapter

(a) Starting with the orbital diagram of a boron atom, describe the steps needed to construct hybrid orbitals appropriate to describe the bonding in \(\mathrm{BF}_{3}\). (b) What is the name given to the hybrid orbitals constructed in (a)? (c) Sketch the large lobes of the hybrid orbitals constructed in part (a). (d) Are any valence atomic orbitals of \(\mathrm{B}\) left unhybridized? If so, how are they oriented relative to the hybrid orbitals?

Consider the bonding in an \(\mathrm{MgH}_{2}\) molecule. (a) Draw a Lewis structure for the molecule, and predict its molecular geometry. (b) What hybridization scheme is used in \(\mathrm{MgH}_{2}\) ? (c) Sketch one of the two-electron bonds between an \(\mathrm{Mg}\) hybrid orbital and an \(\mathrm{H} 1 \mathrm{~s}\) atomic orbital.

From their Lewis structures, determine the number of \(\sigma\) and \(\pi\) bonds in each of the following molecules or ions: (a) \(\mathrm{CO}_{2}\); (b) cyanogen, \((\mathrm{CN})_{2} ;(\mathrm{c})\) formaldehyde, \(\mathrm{H}_{2} \mathrm{CO} ;\) (d) formic acid, \(\mathrm{HCOOH},\) which has one \(\mathrm{H}\) and two \(\mathrm{O}\) atoms attached to \(\mathrm{C}\).

(a) The \(\mathrm{PH}_{3}\) molecule is polar. Does this offer experimental proof that the molecule cannot be planar? Explain. (b) It turns out that ozone, \(\mathrm{O}_{3}\), has a small dipole moment. How is this possible, given that all the atoms are the same?

Determine the electron configurations for \(\mathrm{CN}^{+}, \mathrm{CN},\) and \(\mathrm{CN}^{-}\). (a) Which species has the strongest \(\mathrm{C}-\mathrm{N}\) bond? (b) Which species, if any, has unpaired electrons?
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