Chapter 5

How many \(\sigma\) and \(\pi\) bonds are present in the molecule HCN?

Draw the Lewis structures for \(\mathrm{CO}_{2}\) and \(\mathrm{CO}\), and predict the number of \(\sigma\) and \(\pi\) bonds for each molecule. (a) \(\mathrm{CO}_{2}\) (b) CO

Why is the concept of hybridization required in valence bond theory?

Give the shape that describes each hybrid orbital set: (a) \(s p^{2}\) (b) \(s p^{3} d\) (c) \(s p\) (d) \(s p^{3} d^{2}\)

Explain why a carbon atom cannot form five bonds using \(s p^{3} d\) hybrid orbitals.

Two important industrial chemicals, ethene, \(\mathrm{C}_{2} \mathrm{H}_{4},\) and propene, \(\mathrm{C}_{3} \mathrm{H}_{6}\), are produced by the steam (or thermal) cracking process: \(^{2} 2 \mathrm{C}_{3} \mathrm{H}_{8}(g) \longrightarrow \mathrm{C}_{2} \mathrm{H}_{4}(g)+\mathrm{C}_{3} \mathrm{H}_{6}(g)+\mathrm{CH}_{4}(g)+\mathrm{H}_{2}(g)\) For each of the four carbon compounds, do the following: (a) Draw a Lewis structure. (b) Predict the geometry about the carbon atom. (c) Determine the hybridization of each type of carbon atom.

Analysis of a compound indicates that it contains \(77.55 \%\) Xe and \(22.45 \%\) F by mass. (a) What is the empirical formula for this compound? (Assume this is also the molecular formula in responding to the remaining parts of this exercise). (b) Write a Lewis structure for the compound. (c) Predict the shape of the molecules of the compound. (d) What hybridization is consistent with the shape you predicted?

Consider nitrous acid, HNO_(HONO). (a) Write a Lewis structure. (b) What are the electron pair and molecular geometries of the internal oxygen and nitrogen atoms in the \(\mathrm{HNO}_{2}\) molecule? (c) What is the hybridization on the internal oxygen and nitrogen atoms in \(\mathrm{HNO}_{2}\) ?

Strike-anywhere matches contain a layer of \(\mathrm{KClO}_{3}\) and a layer of \(\mathrm{P}_{4} \mathrm{S}_{3}\). The heat produced by the friction of striking the match causes these two compounds to react vigorously, which sets fire to the wooden stem of the match. \(\mathrm{KClO}_{3}\) contains the \(\mathrm{ClO}_{3}^{-}\) ion. \(\mathrm{P}_{4} \mathrm{S}_{3}\) is an unusual molecule with the skeletal structure. (a) Write Lewis structures for \(\mathrm{P}_{4} \mathrm{S}_{3}\) and the \(\mathrm{ClO}_{3}^{-}\) ion. (b) Describe the geometry about the \(P\) atoms, the \(S\) atom, and the Cl atom in these species. (c) Assign a hybridization to the \(P\) atoms, the \(S\) atom, and the Cl atom in these species. (d) Determine the oxidation states and formal charge of the atoms in \(\mathrm{P}_{4} \mathrm{S}_{3}\) and the \(\mathrm{ClO}_{3}^{-}\) ion.

Write Lewis structures for \(\mathrm{NF}_{3}\) and \(\mathrm{PF}_{5}\). On the basis of hybrid orbitals, explain the fact that \(\mathrm{NF}_{3}, \mathrm{PF}_{3}\), and \(\mathrm{PF}_{5}\) are stable molecules, but \(\mathrm{NF}_{5}\) does not exist.

The bond energy of a \(\mathrm{C}-\mathrm{C}\) single bond averages \(347 \mathrm{kJ} \mathrm{mol}^{-1}\); that of a \(\mathrm{C} \equiv \mathrm{C}\) triple bond averages \(839 \mathrm{kJ}\) mol - \(^{-1}\). Explain why the triple bond is not three times as strong as a single bond.

For the molecule allene, \(\mathrm{H}_{2} \mathrm{C}=\mathrm{C}=\mathrm{CH}_{2}\), give the hybridization of each carbon atom. Will the hydrogen atoms be in the same plane or perpendicular planes?

Identify the hybridization of the central atom in each of the following molecules and ions that contain multiple bonds: (a) CINO (N is the central atom) (b) \(\mathrm{CS}_{2}\) (c) \(\mathrm{Cl}_{2} \mathrm{CO}\) (C is the central atom) (d) \(\mathrm{Cl}_{2} \mathrm{SO}\) (S is the central atom) (e) \(\mathrm{SO}_{2} \mathrm{F}_{2}(\mathrm{S}\) is the central atom) (f) \(\mathrm{XeO}_{2} \mathrm{F}_{2}\) (Xe is the central atom) (g) \(\mathrm{ClOF}_{2}+(\mathrm{Cl}\) is the central atom)

Describe the molecular geometry and hybridization of the \(\mathrm{N}, \mathrm{P},\) or \(\mathrm{S}\) atoms in each of the following compounds. (a) \(\mathrm{H}_{3} \mathrm{PO}_{4},\) phosphoric acid, used in cola soft drinks (b) \(\mathrm{NH}_{4} \mathrm{NO}_{3},\) ammonium nitrate, a fertilizer and explosive (c) \(\mathrm{S}_{2} \mathrm{Cl}_{2}\), disulfur dichloride, used in vulcanizing rubber (d) \(\mathrm{K}_{4}\left[\mathrm{O}_{3} \mathrm{POPO}_{3}\right]\), potassium pyrophosphate, an ingredient in some toothpastes

Sketch the distribution of electron density in the bonding and antibonding molecular orbitals formed from two s orbitals and from two p orbitals.

If molecular orbitals are created by combining five atomic orbitals from atom A and five atomic orbitals from atom B combine, how many molecular orbitals will result?

Can a molecule with an even number of electrons ever be paramagnetic? Explain why or why not.

Why are bonding molecular orbitals lower in energy than the parent atomic orbitals?

Calculate the bond order for an ion with this configuration: \(\left(\sigma_{2 s}\right)^{2}\left(\sigma_{2 s}^{*}\right)^{2}\left(\sigma_{2 p x}\right)^{2}\left(\pi_{2 p y}, \pi_{2 p z}\right)^{4}\left(\pi_{2 p y}^{*}, \pi_{2 p z}^{*}\right)^{3}\)

Explain why an electron in the bonding molecular orbital in the \(\mathrm{H}_{2}\) molecule has a lower energy than an electron in the 1 s atomic orbital of either of the separated hydrogen atoms.

Determine the bond order of each member of the following groups, and determine which member of each group is predicted by the molecular orbital model to have the strongest bond. (a) \(\mathrm{H}_{2}, \mathrm{H}_{2}^{+}, \quad \mathrm{H}_{2}^{-}\) (b) \(\mathrm{O}_{2}, \mathrm{O}_{2}^{2+}, \quad \mathrm{O}_{2}^{2-}\) (c) \(\mathrm{Li}_{2}, \mathrm{Be}_{2}^{+}, \mathrm{Be}_{2}\) (d) \(\mathrm{F}_{2}, \mathrm{F}_{2}^{+}, \mathrm{F}_{2}^{-}\) (e) \(\mathrm{N}_{2}, \mathrm{N}_{2}^{+}, \quad \mathrm{N}_{2}^{-}\)

For the first ionization energy for an \(\mathrm{N}_{2}\) molecule, what molecular orbital is the electron removed from?

Which of the period 2 homonuclear diatomic molecules are predicted to be paramagnetic?

A friend tells you that the 2s orbital for fluorine starts off at a much lower energy than the 2s orbital for lithium, so the resulting \(\sigma_{2 s}\) molecular orbital in \(\mathrm{F}_{2}\) is more stable than in \(\mathrm{Li}_{2}\). Do you agree?

True or false: Boron contains \(2 s^{2} 2 p^{1}\) valence electrons, so only one \(p\) orbital is needed to form molecular orbitals.

Using the MO diagrams, predict the bond order for the stronger bond in each pair: (a) \(\mathrm{B}_{2}\) or \(\mathrm{B}_{2}+\) (b) \(\mathrm{F}_{2}\) or \(\mathrm{F}_{2}^{+}\) (c) \(\mathrm{O}_{2}\) or \(\mathrm{O}_{2}^{2+}\) (d) \(\mathrm{C}_{2}^{+}\) or \(\mathrm{C}_{2}^{-}\)