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Chapter 7: Q107E (page 405)

A molecule with the formula\({\rm{A}}{{\rm{B}}_{\rm{2}}}\), in which A and B represent different atoms, could have one of three different shapes. Sketch and name the three different shapes that this molecule might have. Give an example of a molecule or ion for each shape.

Short Answer

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A molecule with molecular formula \({\rm{A}}{{\rm{B}}_{\rm{2}}}\) can have linear shape for example\({\rm{C}}{{\rm{O}}_{\rm{2}}}\), bent (\({\rm{12}}{{\rm{0}}^{\rm{^\circ }}}\)) shape for example \({\rm{S}}{{\rm{O}}_{\rm{2}}}\) and tetrahedral electron pair geometry \(\left( {{\rm{10}}{{\rm{9}}^{\rm{^\circ }}}} \right)\)shape for example \({{\rm{H}}_{\rm{2}}}{\rm{O}}\)shapes shown as follows:

Step by step solution

01

Definition of Concept

There are two types of bonds in trigonal bipyramidal structures: axial bonds and equatorial bonds. In total, there are five bonds in the trigonal bipyramidal geometry.

02

Sketch and name the three different shapes that this molecule

One A atom and two B atoms make up a molecule with the molecular formula\({\rm{A}}{{\rm{B}}_{\rm{2}}}\). There are two atoms that are bonded together. As a result, this molecule can have three molecular shapes: linear electron pair geometry and linear shape when there are only two electron density regions and no lone pair on the central A-atom, trigonal planar electron pair geometry and bent\({\rm{12}}{{\rm{0}}^{\rm{^\circ }}}\)shape when there are two bond pairs on the central A-atom, and trigonal planar electron pair geometry and bent\({\rm{12}}{{\rm{0}}^{\rm{^\circ }}}\)shape when there are two bond pairs on the central A-atom.

Therefore, the required shapes are below:

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Most popular questions from this chapter

Which of the following atoms would be expected to form negative ions in binary ionic compounds and which would be expected to form positive ions:\({\rm{P, I, Mg, Cl, In, Cs, O, Pb, Co}}\)?

As a general rule, \({\rm{M}}{{\rm{X}}_{\rm{n}}}\) molecules (where \({\rm{M}}\) represents a central atom and \({\rm{X}}\) represents terminal atoms; \({\rm{n = 2 - 5}}\)) are polar if there is one or more lone pairs of electrons on \({\rm{M}}\). \({\rm{N}}{{\rm{H}}_{\rm{3}}}\) (\({\rm{M = N, X = H, n = 3}}\)) is an example. There are two molecular structures with lone pairs that are exceptions to this rule. What are they?

Calculate the formal charge of chlorine in the molecules \({\rm{C}}{{\rm{l}}_{\rm{2}}}\), \({\rm{BeC}}{{\rm{l}}_{\rm{2}}}\), and \({\rm{Cl}}{{\rm{F}}_{\rm{5}}}\).

Predict the electron pair geometry and the molecular structure of each of the following:

  1. \({\rm{IO}}{{\rm{F}}_{\rm{5}}}\)(I is the central atom)
  2. \({\rm{POC}}{{\rm{l}}_{\rm{3}}}\)(P is the central atom)
  3. \({\rm{C}}{{\rm{l}}_{\rm{2}}}{\rm{SeO}}\)(Se is the central atom)
  4. \({\rm{ClS}}{{\rm{O}}^{\rm{ + }}}\)(S is the central atom)
  5. \({{\rm{F}}_{\rm{2}}}{\rm{SO}}\)(S is the central atom)
  6. \({\rm{N}}{{\rm{O}}_{\rm{2}}}^{\rm{ - }}\)
  7. \({\rm{SiO}}_{\rm{4}}^{{\rm{4 - }}}\)

Use the Molecule Shape simulator (http://openstaxcollege.org/l/16MolecShape) to build a molecule. Starting with the central atom, click on the double bond to add one double bond. Then add one single bond and one lone pair. Rotate the molecule to observe the complete geometry. Name the electron group geometry and molecular structure and predict the bond angle. Then click the check boxes at the bottom and right of the simulator to check your answers.
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