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Chapter 7: Q96E (page 404)

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. \({\rm{SiO}}_{\rm{4}}^{{\rm{4 - }}}\)

Short Answer

Expert verified

Electron-pair geometry considers the placement of all electrons. Molecular structure considers only the bonding-pair geometry.

Step by step solution

01

Concept Introduction

The three-dimensional arrangement of atoms in a molecule is dictated by the VSEPR theory. Minimum repulsion is preferred by this theory.

02

Find the electron pair geometry and the molecular structure

Therefore, the arrangement of all electrons is taken into account in electron-pair geometry. Only the bonding-pair geometry is taken into account in molecular structure.

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Most popular questions from this chapter

Which is the most polar bond? (a) \({\rm{C - C}}\) (b) \({\rm{C - H}}\) (c) \({\rm{N - H}}\) (d) \({\rm{O - H}}\) (e) \({\rm{Se - H}}\) .

Use the simulation (http://openstaxcollege.org/l/16MolecPolarity) to perform the following exercises for a real molecule. You may need to rotate the molecules in three dimensions to see certain dipoles. (a) Sketch the bond dipoles and molecular dipole (if any) for O3. Explain your observations. (b) Look at the bond dipoles for NH3. Use these dipoles to predict whether N or H is more electronegative. (c) Predict whether there should be a molecular dipole for NH3 and, if so, in which direction it will point. Check the molecular dipole box to test your hypothesis.

From its position in the periodic table, determine which atom in each pair is more electronegative: (a)\({\rm{Br or Cl}}\)(b)\({\rm{N or O}}\)(c)\({\rm{S or O}}\)(d)\({\rm{P or S}}\)(e)\({\rm{Si or N}}\)(f)\({\rm{Ba or P}}\)(g)\({\rm{N or K}}\).

Which compound in each of the following pairs has the larger lattice energy? Note: \({\rm{B}}{{\rm{a}}^{{\rm{2 + }}}}\) and \({{\rm{K}}^{\rm{ + }}}\) have similar radii; \({{\rm{S}}^{{\rm{2 - }}}}\) and \({\rm{C}}{{\rm{l}}^{\rm{ - }}}\) have similar radii. Explain your choices.

(a) \({{\rm{K}}_{\rm{2}}}{\rm{O}}\) or \({\rm{N}}{{\rm{a}}_{\rm{2}}}{\rm{O}}\)

(b) \({{\rm{K}}_{\rm{2}}}{\rm{S}}\) or \({\rm{BaS}}\)

(c) \({\rm{KCl}}\) or \({\rm{BaS}}\)

(d) \({\rm{BaS}}\) or \({\rm{BaC}}{{\rm{l}}_{\rm{2}}}\)

The lattice energy of \({\rm{LiF}}\) is \({\rm{1023 kJ/mol}}\), and the \({\rm{Li - F}}\) distance is \({\rm{200}}{\rm{.8 pm}}\). \({\rm{NaF}}\) crystallizes in the same structure as \({\rm{LiF}}\) but with a \({\rm{Na - F}}\) distance of \({\rm{231 pm}}\). Which of the following values most closely approximates the lattice energy of \({\rm{NaF}}\): \({\rm{510, 890, 1023, 1175,}}\) or \({\rm{4090 kJ/mol}}\)? Explain your choice.
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