Chapter 11

Represent bonding in the carbon dioxide molecule, ${\mathrm{CO}}_2,$ by $(\mathrm{a})$ a Lewis structure and $(\mathrm{b})$ the valencebond method. Identify $\sigma$ and $\pi$ bonds, the necessary hybridization scheme, and orbital overlap.

The structure of the molecule allene, ${\mathrm{CH}}_2{\mathrm{CCH}}_2$, is shown here. Propose hybridization schemes for the $C$ atoms in this molecule.

Explain the essential difference in how the valencebond method and molecular orbital theory describe a covalent bond.

Describe the bond order of diatomic carbon, ${\mathrm{C}}_2,$ with Lewis theory and molecular orbital theory, and explain why the results are different.

${\mathrm{N}}_2(\mathrm{g})$ has an exceptionally high bond energy. Would you expect either ${\mathrm{N}}_2^{-}$ or ${\mathrm{N}}_2^{2-}$ to be a stable diatomic species in the gaseous state? Explain.

The paramagnetism of gaseous ${\mathrm{B}}_2$ has been established. Explain how this observation confirms that the ${\pi }_{2p}$ orbitals are at a lower energy than the ${\sigma }_{2p}$ orbital for ${\mathrm{B}}_2$

Is it correct to say that when a diatomic molecule loses an electron, the bond energy always decreases (that is, that the bond is always weakened)? Explain.

For the following pairs of molecular orbitals, indicate the one you expect to have the lower energy, and state the reason for your choice. (a) ${\sigma }_{1s}$ or ${\sigma }_{1s}^{\ast };$ (b) ${\sigma }_{2s}$ or ${\sigma }_{2p}$ (c) ${\sigma }_{1s}^{\ast }$ or ${\sigma }_{2s};$ (d) ${\sigma }_{2p}$ or ${\sigma }_{2p}^{\ast }$

For each of the species ${\mathrm{C}}_2^{+},{\mathrm{O}}_2^{-},{\mathrm{F}}_2^{+},$ and ${\mathrm{NO}}^{+}$ (a) Write the molecular orbital diagram (as in Example $11-6)$ (b) Determine the bond order, and state whether you expect the species to be stable or unstable. (c) Determine if the species is diamagnetic or paramagnetic; and if paramagnetic, indicate the number of unpaired electrons.

Write plausible molecular orbital diagrams for the following heteronuclear diatomic species: (a) $\mathrm{NO};$ (b) ${\mathrm{NO}}^{+}$ (c) $\mathrm{CO};$ (d) $\mathrm{CN};$ (e) ${\mathrm{CN}}^{-};$ (f) ${\mathrm{CN}}^{+};$ (g) BN.