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Chapter 10: Problem 71

Comment on the similarities and differences in the molecular structure of the following triatomic species: \(\mathrm{CO}_{2}, \mathrm{NO}_{2}^{-}, \mathrm{O}_{3},\) and \(\mathrm{ClO}_{2}^{-}\).

Short Answer

Expert verified
CO2 has a linear shape, 180° bond angle, and is non-polar. NO2- and O3 have bent shapes with <120° bond angles and are polar. ClO2- is also bent but with a <109.5° bond angle, and is polar.

Step by step solution

01

Draw Lewis Structures

Begin by drawing Lewis structures. For CO2, there are two double bonds between C and each O atom. For NO2-, there is one single bond and one double bond between N and O atoms, and the remaining O atom carries a negative charge. The O3 molecule has a double bond between one pair of O atoms and a single bond with a lone pair on another O atom. Lastly, the ClO2- molecule has one double bond with an O atom, a single bond with another O atom, and three lone pairs on Cl atom, with an overall negative charge.
02

Determine Molecular Geometry using VSEPR Theory

Each molecule has a different shape based on the VSEPR theory. CO2 is linear as it has two bonded and no non-bonded electron pairs around the central C atom. NO2- is bent as it has two bonded and one non-bonded electron pairs around the N atom. O3 is bent as it also has two bonded and one non-bonded electron pairs around the central O atom. Lastly, ClO2- is bent due to three bonded and one non-bonded electron pairs around the Cl atom.
03

Comment on Bond Angles and Polarity

As per the VSEPR theory, CO2 has bond angles of 180° and is non-polar due to the symmetry of the molecule. NO2- and O3 have bond angles of less than 120° and are polar molecules. ClO2- has bond angles of less than 109.5° and is also polar.

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Most popular questions from this chapter

The concept of formal charge helped us to choose the more plausible of the I ewis structures for \(\mathrm{NO}_{2}^{+}\) given in expressions (10.14) and \((10.15) .\) Can it similarly help us to choose a single Lewis structure as most plausible for \(\mathrm{CO}_{2} \mathrm{H}^{+} ?\) Explain.

Use VSEPR theory to predict the geometric shapes of the following molecules and ions: (a) \(\mathrm{N}_{2}\); (b) HCN; (c) \(\mathrm{NH}_{4}^{+} ;\) (d) \(\mathrm{NO}_{3}^{-} ;\) (e) NSF.

By means of Lewis structures, represent bonding between the following pairs of elements: (a) Cs and \(\mathrm{Br} ;\) (b) \(\mathrm{H}\) and \(\mathrm{Sb} ;\) (c) \(\mathrm{B}\) and \(\mathrm{Cl} ;\) (d) \(\mathrm{Cs}\) and \(\mathrm{Cl}\); (e) Li and O; (f) Cl and L. Your structures should show whether the bonding is essentially ionic or covalent.

Which of the following molecules would you expect to be polar: (a) \(\mathrm{HCN} ;\) (b) \(\mathrm{SO}_{3} ;\) (c) \(\mathrm{CS}_{2} ;\) (d) OCS; (e) \(\operatorname{SOCl}_{2} ;\) (f) \(\operatorname{SiF}_{4} ;\) (g) \(\operatorname{POF}_{3}\) ? Give reasons for your conclusions.

Explain why it is not necessary to find the Lewis structure with the smallest formal charges to make a successful prediction of molecular geometry in the VSEPR theory. For example, write Lewis structures for \(S O_{2}\) having different formal charges, and predict the molecular geometry based on these structures.
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