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Chapter 9: Problem 59

(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 from the combination of atomic orbitals within a single atom, while molecular orbitals are formed from the overlapping atomic orbitals of two or more atoms, describing electron distribution in a molecule. b) Each molecular orbital can hold a maximum of two electrons, due to the Pauli Exclusion Principle. c) Yes, antibonding molecular orbitals can have electrons in them, although electrons preferentially fill lower-energy bonding orbitals before occupying higher-energy antibonding orbitals.

Step by step solution

01

a) Difference between Hybrid Orbitals and Molecular Orbitals

Hybrid orbitals are formed from the combination of atomic orbitals of a single atom, which results in a new set of orbitals with equivalent energy levels and optimized shapes for forming chemical bonds with other atoms. On the other hand, molecular orbitals are formed when atomic orbitals from two or more atoms combine and overlap, resulting in a new set of orbitals that describe the distribution of electrons in a molecule. In summary, hybrid orbitals deal with the rearrangement of a single atom's atomic orbitals, while molecular orbitals describe the distribution and behavior of electrons in a molecule formed by two or more atoms.
02

b) Number of Electrons in each Molecular Orbital

Each molecular orbital (MO) can hold a maximum of two electrons. This is due to the Pauli Exclusion Principle, which states that no two electrons in an atom or molecule can have the same set of quantum numbers. When two electrons occupy the same orbital, they must have opposite spins (one up and one down), which ensures that their quantum numbers are different.
03

c) Antibonding Molecular Orbitals with Electrons

Yes, antibonding molecular orbitals can have electrons in them. These orbitals result from the out-of-phase combination of atomic orbitals from two or more atoms, leading to a node between the atoms with minimal electron density. When electrons are in antibonding orbitals, they counteract the bonding effect of electrons in bonding molecular orbitals, leading to a weaker overall bond or even an unstable molecule. However, it is important to note that electrons will preferentially fill lower-energy bonding orbitals before occupying higher-energy antibonding orbitals.

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Most popular questions from this chapter

Many compounds of the transition-metal elements contain direct bonds between metal atoms. We will assume that the \(z\) -axis is defined as the metal-metal bond axis. (a) Which of the \(3 d\) orbitals (Figure 6.24) can be used to make a \(\sigma\) bond between metal atoms? (b) Sketch the \(\sigma_{3 d}\) bonding and \(\sigma_{3 d}^{*}\) antibonding MOs. (c) With reference to the "Closer Look" box on the phases of orbitals, explain why a node is generated in the \(\sigma_{3 d}^{*}\) MO. (d) Sketch the energy-level diagram for the \(\mathrm{Sc}_{2}\) molecule, assuming that only the \(3 d\) orbital from part (a) is important. (e) What is the bond order in \(\mathrm{Sc}_{2} ?\)

Propylene, \(\mathrm{C}_{3} \mathrm{H}_{4}\) 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?

The nitrogen atoms in \(\mathrm{N}_{2}\) participate in multiple bonding, whereas those in hydrazine, \(\mathrm{N}_{2} \mathrm{H}_{4}\), do not. (a) Draw Lewis structures for both molecules. (b) What is the hybridization of the nitrogen atoms in each molecule? (c) Which molecule has a stronger \(\mathrm{N}-\mathrm{N}\) bond?

The Lewis structure for allene is Make a sketch of the structure of this molecule that is analogous to Figure 9.27. In addition, answer the following three questions: (a) Is the molecule planar? (b) Does it have a nonzero dipole moment? (c) Would the bonding in allene be described as delocalized? Explain.

The three species \(\mathrm{NH}_{2}^{-} \mathrm{NH}_{3}\), and \(\mathrm{NH}_{4}{ }^{+}\) have \(\mathrm{H}-\mathrm{N}-\mathrm{H}\) bond angles of \(105^{\circ}, 107^{\circ}\), and \(109^{\circ}\), respec- tively. Explain this variation in hond
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