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Chapter 9: Problem 12

(a) Methane \(\left(\mathrm{CH}_{4}\right)\) and the perchlorate ion \(\left(\mathrm{ClO}_{4}^{-}\right)\) are both described as tetrahedral. What does this indicate about their bond angles? (b) The \(\mathrm{NH}_{3}\) molecule is trigonal pyramidal, while \(\mathrm{BF}_{3}\) is trigonal planar. Which of these molecules is flat?

Short Answer

Expert verified
(a) The bond angles in both methane (CH₄) and the perchlorate ion (ClO₄⁻) are 109.5°, because they have a tetrahedral geometry. (b) The BF₃ molecule is flat, while the NH₃ molecule is not, as BF₃ has a trigonal planar geometry and NH₃ has a trigonal pyramidal geometry due to the presence of a lone pair of electrons.

Step by step solution

01

Tetrahedral Geometry

Both CH₄ and ClO₄⁻ are described as having tetrahedral molecular geometries. This means that they have central atoms (C in CH₄ and Cl in ClO₄⁻) surrounded by four bonded atoms, all positioned at the vertices of a regular tetrahedron.
02

Tetrahedral Bond Angles

The bond angles of a tetrahedral molecule are always 109.5°. Since both CH₄ and ClO₄⁻ are tetrahedral, the bond angles in both of these molecules will be 109.5°. Answer to (a): The bond angles in both methane (CH₄) and the perchlorate ion (ClO₄⁻) are 109.5°.
03

Understanding Trigonal Pyramidal and Trigonal Planar Geometry

NH₃ has a trigonal pyramidal geometry, which means it has a central atom (N) surrounded by three bonded atoms (H) and a lone pair of electrons. The atoms are positioned at the vertices of a triangle, while the lone pair occupies the position above the central atom, forming a pyramid shape. BF₃, on the other hand, has a trigonal planar geometry - it has a central atom (B) surrounded by three bonded atoms (F) arranged at the vertices of an equilateral triangle, with all atoms lying in the same plane.
04

Determine Flat Molecules

From the geometries described in Step 3, we can see that NH₃ is not flat because it has a pyramidal shape resulting from the presence of the lone pair of electrons. Conversely, BF₃ is considered flat because the central atom and all the bonded atoms lie in the same plane. Answer to (b): The BF₃ molecule is flat, while the NH₃ molecule is not.

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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 higherenergy MO. Would you expect the excited-state \(\mathrm{H}_{2}^{+}\) ion to be stable or to fall apart? Explain.

(a) Draw a picture showing how two \(p\) orbitals on two different atoms can be combined to make a sigma bond. (b) Sketch a \(\pi\) bond that is constructed from \(p\) orbitals. (c) Which is generally stronger, a \(\sigma\) bond or a \(\pi\) bond? Explain. (d) Can two \(s\) orbitals combine to form a \(\pi\) bond? Explain.

What are the electron-domain and molecular geometries of a molecule that has the following electron domains on its central atom? (a) three bonding domains and no nonbonding domains, (b) three bonding domains and one nonbonding domain, (c) two bonding domains and two nonbonding domains.

The azide ion, \(\mathrm{N}_{3}^{-}\), is linear with two \(\mathrm{N}-\mathrm{N}\) bonds of equal length, \(1.16 \AA\). (a) Draw a Lewis structure for the azide ion. (b) With reference to Table \(8.5,\) is the observed \(\mathrm{N}-\mathrm{N}\) bond length consistent with your Lewis structure? (c) What hybridization scheme would you expect at each of the nitrogen atoms in \(\mathrm{N}_{3}^{-}\) ? (d) Show which hybridized and unhybridized orbitals are involved in the formation of \(\sigma\) and \(\pi\) bonds in \(\mathrm{N}_{3}^{-}\). (e) It is often observed that \(\sigma\) bonds that involve an \(s p\) hybrid orbital are shorter than those that involve only \(s p^{2}\) or \(s p^{3}\) hybrid orbitals. Can you propose a reason for this? Is this observation applicable to the observed bond lengths in \(\mathrm{N}_{3}^{-}\) ?

(a) What does the term paramagnetism mean? (b) How can one determine experimentally whether a substance is paramagnetic? (c) Which of the following ions would you expect to be paramagnetic: \(\mathrm{O}_{2}^{+}, \mathrm{N}_{2}^{2-}, \mathrm{Li}_{2}^{+}, \mathrm{O}_{2}^{2-} ?\) For those ions that are paramagnetic, determine the number of unpaired electrons.
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