Chapter 4

What is meant by the term orbital overlap?

Which of the following orbitals could overlap? (a) \(2 s+2 s\) (b) \(2 p_{x}+2 p_{y}\) (c) \(2 p_{y}+2 p_{y}\) (d) \(2 p_{z}+2 p_{z}\) (e) \(2 s+2 p_{z}\)

Draw valence bond pictures showing the overlap of atomic orbitals in each of the following molecules. (a) \(\mathrm{NH}_{3}\) (b) \(\mathrm{H}_{2} \mathrm{~S}\) (c) \(\mathrm{NO}_{2}\) (d) \(\mathrm{C}_{2}\)

Determine the hybridization of the central atom in the following molecules or ions. (a) \(\mathrm{NH}_{3}\) (b) \(\mathrm{NH}_{4}^{+}\) (c) NO (d) \(\mathrm{NO}_{2}\) (e) \(\mathrm{NO}_{2}^{+}\)

Determine the hybridization of the central atom in the following molecules or ions. (a) \(\mathrm{CH}_{4}\) (b) \(\mathrm{H}_{2} \mathrm{CO}\) (c) \(\mathrm{HCO}_{2}^{-}\)

Determine the hybridization of the central atom in the following molecules or ions. (a) \(\mathrm{SF}_{4}\) (b) \(\mathrm{BrO}_{3}^{-}\) (c) \(\mathrm{XeF}_{3}^{+}\) (d) \(\mathrm{Cl}_{2} \mathrm{CO}\)

Write the Lewis structures for molecules of ethylene, \(\mathrm{C}_{2} \mathrm{H}_{4},\) and acetylene, \(\mathrm{C}_{2} \mathrm{H}_{2} .\) Use hybrid atomic orbitals to describe the bonding in these compounds.

Write the Lewis structures for carbon monoxide and carbon dioxide. Use hybrid atomic orbitals to describe the bonding in the compounds.

Describe the difference between \(\sigma\) and \(\pi\) molecular orbitals.

Describe the difference between bonding and antibonding molecular orbitals.

Describe the molecular orbitals formed by the overlap of the following atomic orbitals. (Assume that the bond lies along the \(z\) axis of the coordinate system.) (a) \(2 s+2 s\) (b) \(2 p_{x}+2 p_{x}\) (c) \(2 p_{y}+2 p_{y}\) (d) \(2 p_{z}+2 p_{z}\) (e) \(2 s+2 p_{z}\)

Write the electron configuration for the following diatomic molecules and calculate the bond order in each molecule. (a) \(\mathrm{H}_{2}\) (b) \(\mathrm{C}_{2}\) (c) \(\mathrm{N}_{2}\) (d) \(\mathrm{O}_{2}\) (e) \(\mathrm{F}_{2}\)

Use molecular orbital theory to predict whether the \(\mathrm{H}_{2}^{+}, \mathrm{H}_{2}^{-}\), and \(\mathrm{H}_{2}{\underline{\phantom{xx}}}^{2-}\) ions should be more stable or less stable than a neutral \(\mathrm{H}_{2}\) molecule.

Use molecular orbital theory to explain why the oxygen-oxygen bond is stronger in the \(\mathrm{O}_{2}\) molecule than in the \(\mathrm{O}_{2}^{2-}\) (peroxide) ion.

Use molecular orbital theory to predict whether the bond order in the superoxide ion, \(\mathrm{O}_{2}^{-}\), should be larger or smaller than the bond order in a neutral \(\mathrm{O}_{2}\) molecule.

Use molecular orbital theory to predict whether the peroxide ion, \(\mathrm{O}_{2}{\underline{\phantom{xx}}}^{2-}\), should be paramagnetic.

Write the electron configuration for the following diatomic molecules. Calculate the bond order in each molecule. (a) \(\mathrm{HF}\) (b) \(\mathrm{CO}\) (c) \(\mathrm{CN}^{-}\) (d) \(\mathrm{ClO}^{-}\) (e) \(\mathrm{NO}^{+}\)

Classify the following molecules as paramagnetic or diamagnetic. (a) HF (b) \(\mathrm{CO}\) (c) \(\mathrm{CN}^{-}\) (d) NO (e) \(\mathrm{NO}^{+}\)