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Chapter 8: Q45E (page 451)

A friend tells you that the \({\rm{2s}}\)orbital for fluorine starts off at a much lower energy than the \({\rm{2s}}\) orbital for lithium, so the resulting \({{\rm{\sigma }}_{{\rm{2s}}}}\) molecular orbital in \({{\rm{F}}_{\rm{2}}}\) is more stable than in \({\rm{L}}{{\rm{i}}_{\rm{2}}}\). Do you agree?

Short Answer

Expert verified

Yes, it is agreed.

Step by step solution

01

Definition of diatomic molecule

A diatomic molecule is a molecule that consists of only two atoms of the same or different chemical elements

02

 Evaluating \({{\bf{F}}_{\bf{2}}}\) is more stable than \({\bf{L}}{{\bf{i}}_{\bf{2}}}\)

Yes, i agree. The \(2s\) Fluorine orbits start at much lower energies \(2s\) this is because fluorine is smaller than lithium and is orbital lithium \(2s\) The orbit of fluorine is closer to the nucleus than lithium and is more stable. Therefore, the result is \({\sigma _{2s}}\)molecular orbital in \({F_2}\)is more stable than in \({\rm{L}}{{\rm{i}}_{\rm{2}}}\).

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Most popular questions from this chapter

Why are bonding molecular orbitals lower in energy than the parent atomic orbitals?

Strike-anywhere matches contain a layer of \({\rm{KCl}}{{\rm{O}}_{\rm{3}}}\) and a layer of \({{\rm{P}}_{\rm{4}}}{{\rm{S}}_{\rm{3}}}\). The heat produced by the friction of striking the match causes these two compounds to react vigorously, which sets fire to the wooden stem of the match. \({\rm{KCl}}{{\rm{O}}_{\rm{3}}}\) contains the \({\rm{Cl}}{{\rm{O}}_{\rm{3}}}^{\rm{ - }}\) ion. \({{\rm{P}}_{\rm{4}}}{{\rm{S}}_{\rm{3}}}\) is an unusual molecule with the skeletal structure.

  1. Write Lewis structures for \({{\rm{P}}_{\rm{4}}}{{\rm{S}}_{\rm{3}}}\) and the \({\rm{Cl}}{{\rm{O}}_{\rm{3}}}^{\rm{ - }}\) ion.
  2. Describe the geometry about the \({\rm{P}}\) atoms, the \({\rm{S}}\) atom, and the \({\rm{Cl}}\) atom in these species.
  3. Assign a hybridization to the \({\rm{P}}\) atoms, the \({\rm{S}}\)atom, and the \({\rm{Cl}}\) atom in these species.
  4. Determine the oxidation states and formal charge of the atoms in \({{\rm{P}}_{\rm{4}}}{{\rm{S}}_{\rm{3}}}\) and the \({\rm{Cl}}{{\rm{O}}_{\rm{3}}}^{\rm{ - }}\) ion.

A typical barometric pressure in Redding, California, is about 750mm Hg. Calculate this pressure in atm and kPa.

Give the shape that describes each hybrid orbital set:

(a) \({\rm{s}}{{\rm{p}}^{\rm{2}}}\)

(b) \({\rm{s}}{{\rm{p}}^{\rm{3}}}{\rm{d}}\)

(c) sp

(d) \({\rm{s}}{{\rm{p}}^{\rm{3}}}{{\rm{d}}^{\rm{2}}}\)

Sketch the distribution of electron density in the bonding and antibonding molecular orbitals formed from two S orbitals and from two P orbitals.

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