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Chapter 11: Problem 16

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.

Short Answer

Expert verified
The Lewis structure for CO2 is \(O = C = O\) showing carbon atom in the center, double bonded to two oxygen atoms. In terms of Valence Bond theory, each double bond contains one \(\sigma\) bond and one \(\pi\) bond. The \(\sigma\) bond is formed by the overlap of an sp hybrid orbital from carbon and a p orbital from oxygen, while the \(\pi\) bond is formed through the sideways overlapping of the p orbitals of both carbon and oxygen atoms. Carbon undergoes sp hybridization in CO2.

Step by step solution

01

Lewis Structure

A Lewis structure shows all the valence electrons in a molecule. Since oxygen has 6 valence electrons and carbon has 4, we total these together and get \(6*2 + 4 = 16\) electrons. These 16 electrons form the structure: \(O = C = O\). That structure indicates carbon atom in the center, double bonded to two oxygen atoms on either side.
02

The Valence Bond Method and Identifying \(\sigma\) and \(\pi\) Bonds

The valence bond method lays the foundation to identify \(\sigma\) and \(\pi\) bonds. Each double bond consists of a \(\sigma\) bond and a \(\pi\) bond. In this molecule, one \(\sigma\) bond is formed by the overlap of an sp hybrid orbital from carbon and a p orbital from oxygen. The \(\pi\) bond is formed by the sideways overlapping of p orbitals.
03

Identifying Hybridization Scheme

In CO2, the carbon atom is the center atom linked to two oxygen atoms through double bonds. Hence, carbon forms 2 sigma bonds with the two oxygen atoms, which requires two hybrid orbitals. Therefore, the hybridization of carbon in CO2 is sp (two hybrid orbitals: one s and one p orbital).

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Most popular questions from this chapter

102\. Which of the following species are paramagnetic? (a) \(\mathrm{B}_{2} ;\) (b) \(\mathrm{B}_{2}^{-} ;\) (c) \(\mathrm{B}_{2}^{+}\). Which species has the strongest bond?

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.

The ion \(\mathrm{F}_{2} \mathrm{Cl}^{-}\) is linear, but the ion \(\mathrm{F}_{2} \mathrm{Cl}^{+}\) is bent. Describe hybridization schemes for the central \(\mathrm{Cl}\) atom consistent with this difference in structure.

In your own words, define the following terms or symbols: (a) \(s p^{2} ;\) (b) \(\sigma_{2 p}^{*} ;\) (c) bond order; (d) \(\pi\) bond.

For each of the following species, identify the central atom(s) and propose a hybridization scheme for those $\operatorname{atom}(\mathrm{s}):(\mathrm{a}) \mathrm{CO}_{2} ;(\mathrm{b}) \mathrm{HONO}_{2} ;(\mathrm{c}) \mathrm{ClO}_{3}^{-} ;(\mathrm{d}) \mathrm{BF}_{4}^{-}$
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