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Chapter 10: Problem 37

Specify which hybrid orbitals are used by carbon atoms in these species: (a) \(\mathrm{CO},\) (b) \(\mathrm{CO}_{2},\) (c) \(\mathrm{CN}^{-}\).

Short Answer

Expert verified
The carbon atom in each molecule uses the following hybrid orbitals: (a) CO uses sp hybrid orbitals, (b) CO2 also uses sp hybrid orbitals, and (c) CN- uses sp2 hybrid orbitals.

Step by step solution

01

Determine Electron Domains for CO

For molecule CO, Carbon is bonded to a single Oxygen atom and has one lone pair. This means that Carbon has two electron domains.
02

Identify Hybridization for CO

Since there are two electron domains surrounding Carbon in CO, it implies that Carbon uses sp hybrid orbitals.
03

Determine Electron Domains for CO2

For molecule CO2, Carbon is bonded to two Oxygen atoms and carries no lone pairs. This means that Carbon has two electron domains.
04

Identify Hybridization for CO2

Since there are two electron domains surrounding Carbon in CO2, it implies that Carbon uses sp hybrid orbitals.
05

Determine Electron Domains for CN-

For molecule CN-, Carbon is bonded to one Nitrogen atom and has no lone pairs but has a triple bond with nitrogen. It has three electron domains.
06

Identify Hybridization for CN-

Since there are three electron domains around Carbon in CN-, it implies that Carbon uses sp2 hybrid orbitals.

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Most popular questions from this chapter

Draw a molecular orbital energy level diagram for each of these species: \(\mathrm{He}_{2}, \mathrm{HHe}, \mathrm{He}_{2}^{+} .\) Compare their relative stabilities in terms of bond orders. (Treat HHe as a diatomic molecule with three electrons.)

List these molecules in order of increasing dipole moment: \(\mathrm{H}_{2} \mathrm{O}, \mathrm{CBr}_{4}, \mathrm{H}_{2} \mathrm{~S}, \mathrm{HF}, \mathrm{NH}_{3}, \mathrm{CO}_{2}\)

Does the molecule OCS have a higher or lower dipole moment than \(\mathrm{CS}_{2} ?\)

Which of these species is not likely to have a tetrahedral shape (a) \(\operatorname{SiBr}_{4}\) (b) \(\mathrm{NF}_{4}^{+},(\mathrm {c}) \mathrm{SF}_{4}\) (d) \(\mathrm{BeCl}_{4}^{2-}\) (e) \(\mathrm{BF}_{4}^{-},\) (f) \(\mathrm{AlCl}_{4}^{-}\)

Which of these pairs of atomic orbitals of adjacent nuclei can overlap to form a sigma bond? Which overlap to form a pi bond? Which cannot overlap (no bond)? Consider the \(x\) -axis to be the internuclear axis, that is, the line joining the nuclei of the two atoms. (a) \(1 s\) and \(1 s,\) (b) \(1 s\) and \(2 p_{x},\) (c) \(2 p_{x}\) and \(2 p_{y}\) (d) \(3 p_{y}\) and \(3 p_{y},\) (e) \(2 p_{x}\) and \(2 p_{x}\), (f) 1 s and 2 s.
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