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Chapter 9: Problem 69

(a) What is the difference between hybrid orbitals and molecular orbitals? (b) How many electrons can be placed into each MO of a molecule? (c) Can antibonding molecular orbitals have electrons in them?

Short Answer

Expert verified
(a) Hybrid orbitals are formed by mixing atomic orbitals in an atom to create new orbitals with different energy levels and shapes, while molecular orbitals are formed by combining atomic orbitals from multiple atoms within a molecule, resulting in shared electrons among the atoms. (b) Each molecular orbital can hold a maximum of two electrons, following the Pauli exclusion principle. (c) Yes, antibonding molecular orbitals can have electrons in them, but this destabilizes the bond and decreases bond order. Typically, electrons fill bonding orbitals first, following the aufbau principle.

Step by step solution

01

(a) Definition of hybrid orbitals

Hybrid orbitals are formed by the mixing of atomic orbitals in an atom to form new orbitals with energy levels and shapes different from the original atomic orbitals. They help explain molecular geometry and bonding patterns in molecules, such as sp, sp2, and sp3 hybridization in carbon atoms.
02

(a) Definition of molecular orbitals

Molecular orbitals are formed by the combination of atomic orbitals from multiple atoms within a molecule, resulting in the electrons being shared and distributed among the atoms. Molecular orbitals can be bonding, non-bonding, or antibonding orbitals, depending on their energy levels and electron distribution.
03

(b) Number of electrons in molecular orbitals

In each molecular orbital, a maximum of two electrons can be placed, following the Pauli exclusion principle. This principle states that no two electrons in an atom or molecule can have the same set of quantum numbers, meaning that there can only be two electrons with opposite spins in each molecular orbital.
04

(c) Antibonding molecular orbitals and electron occupation

Yes, antibonding molecular orbitals can have electrons in them. Antibonding molecular orbitals result from the out-of-phase combination of atomic orbitals, which leads to a decrease in electron density between the nuclei of the atoms. This destabilizes the bond, effectively counteracting the stabilizing effect of bonding molecular orbitals. However, if electrons are present in the antibonding orbitals, the bond order will decrease and, in some cases, no bond will form. Typically, electrons will fill the bonding orbitals before filling the antibonding orbitals, following the aufbau principle.

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Most popular questions from this chapter

Propylene, \(\mathrm{C}_{3} \mathrm{H}_{6}\), is a gas that is used to form the important polymer called polypropylene. Its Lewis structure is (a) What is the total number of valence electrons in the propylene molecule? (b) How many valence electrons are used to make \(\sigma\) bonds in the molecule? (c) How many valence electrons are used to make \(\pi\) bonds in the molecule? (d) How many valence electrons remain in nonbonding pairs in the molecule? (e) What is the hybridization at each carbon atom in the molecule?

Predict whether each of the following molecules is polar or nonpolar: (a) \(\mathrm{IF}_{3}\) (b) \(\mathrm{CS}_{2}\), (c) \(\mathrm{SO}_{3}\), (d) \(\mathrm{PCl}_{3}\), (e) \(\mathrm{SF}_{6}\), (f) \(\mathrm{IF}_{5}\)

Draw the Lewis structure for each of the following molecules or ions, and predict their electron-domain and molecular geometries: (a) \(\mathrm{AsF}_{3}\), (b) \(\mathrm{CH}_{3}^{+}\), (c) \(\mathrm{BrF}_{3}\), (d) \(\mathrm{ClO}_{3}^{-}\), (e) \(\mathrm{XeF}_{2}\), (f) \(\mathrm{BrO}_{2}^{-}\).

The \(\mathrm{O}-\mathrm{H}\) bond lengths in the water molecule \(\left(\mathrm{H}_{2} \mathrm{O}\right)\) are \(0.96 \AA\), and the \(\mathrm{H}-\mathrm{O}-\mathrm{H}\) angle is \(104.5^{\circ}\). The dipole moment of the water molecule is \(1.85\) D. (a) In what directions do the bond dipoles of the \(\mathrm{O}-\mathrm{H}\) bonds point? In what direction does the dipole moment vector of the water molecule point? (b) Calculate the magnitude of the bond dipole of the \(\mathrm{O}-\mathrm{H}\) bonds. (Note: You will need to use vector addition to do this.) (c) Compare your answer from part (b) to the dipole moments of the hydrogen halides (Table 8.3). Is your answer in accord with the relative electronegativity of oxygen?

(a) What is meant by the term orbital overlap? (b) Describe what a chemical bond is in terms of electron density between two atoms.
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