10 Hybridization and Molecular Orbital Theory Online Quiz Questions

An atom X has one double bond, one single bond, and one lone pair. Using the introductory electron-region method, determine its common hybridization and approximate arrangement. Explain how you counted the regions, and state one reason the actual shape or bond angles might differ from the ideal arrangement.

For ethene, C₂H₄, explain how each carbon's hybridization accounts for both the σ and π components of the C=C bond. Include the role of the unhybridized orbital.

Use the electron-region model to explain why water is commonly described as having sp³-hybridized oxygen while its H–O–H angle is about 104.5° rather than the ideal tetrahedral angle of 109.5°.

For methane, connect the number of electron-density regions around carbon to its common hybridization, the bonds formed by its hybrid orbitals, and the ideal arrangement and angle.

In the LCAO approach, a pair of atomic orbitals does not necessarily combine effectively just because both orbitals are present. Identify the three compatibility conditions that favor effective combination and explain why an unfavorable match in any of them matters.

Two atomic orbitals combine in an introductory MO description. Describe the two molecular orbitals produced, comparing their relative energies, electron-density patterns between the nuclei, and effects on bonding.

Compare the relative energies of the π(2p) and σ(2p) molecular orbitals in N₂ and O₂. Explain why using the correct ordering is necessary when interpreting an MO diagram.

An MO diagram has two orbitals at the same energy and two electrons to place in them. State the arrangement that follows Hund's rule, then explain how the Pauli exclusion principle constrains any later pairing in one of those orbitals.

A hypothetical diatomic molecule has 10 electrons in bonding MOs and 6 in antibonding MOs. Calculate its MO bond order and explain what that value generally suggests about bond strength and bond length when compared with a molecule of the same atoms that has a lower bond order.

Suppose one electron is transferred from a bonding MO to an antibonding MO in a diatomic molecule. Using the MO bond-order definition, determine how much the bond order changes and explain the likely qualitative effect on bond strength and length for the same pair of atoms.

Neutral O₂ has eight bonding electrons and four antibonding electrons. In O₂⁺, one electron has been removed from an antibonding π* orbital, leaving eight bonding electrons and three antibonding electrons. Calculate the bond order of each species and explain the change.

Use MO electron occupancy to explain why O₂ is paramagnetic while N₂ is diamagnetic. Include the role of the two equal-energy π* orbitals in O₂ and explain what this comparison reveals about the limitation of a basic Lewis structure for oxygen.

Compare the orbital overlaps in a double bond and a triple bond. State how many $\sigma$ and $\pi$ bonds each typically contains, and describe the overlap that forms each bond type.

Compare the carbon atom in $\mathrm{CO_2}$ with the carbon atom in $\mathrm{CH_4}$. For each, count the electron-density regions, give the common hybridization and approximate arrangement, and explain why the multiple bonds in $\mathrm{CO_2}$ do not count as extra regions.

What does hybridization describe in valence-bond theory, and what is one important limitation on how it should be interpreted as a physical description of bonding?

A student knows that a multiple bond contains more than one shared electron pair but cannot identify its orbital overlaps. Compare the overlap that forms sigma and pi bonds, then state the usual sigma/pi composition of a double bond and a triple bond.

A proposed hybridization scheme combines one s orbital and two p orbitals but claims to produce four hybrid orbitals. Assess the proposal using the orbital-count rule and give the correct hybrid set.

Compare valence-bond theory and molecular-orbital theory in terms of where their orbitals are associated and what each model helps explain. Use oxygen's magnetic behavior to identify one explanatory advantage of the molecular-orbital model over a basic Lewis structure.

When two atomic orbitals combine, explain how the resulting bonding and antibonding molecular orbitals differ in energy, electron distribution between the nuclei, and effect on bonding. What does the asterisk in σ* or π* indicate?

Three electrons are to be placed in two molecular orbitals of equal energy. Describe their most appropriate distribution, state how many electrons remain unpaired, and explain how Hund's rule and the Pauli exclusion principle constrain the filling.

A student uses one 2p-derived molecular-orbital ordering for every second-period homonuclear diatomic molecule. Explain why this is not reliable, and state the relative order of π(2p) and σ(2p) for B₂ through N₂ and for O₂ and F₂.

In a hypothetical molecule, the occupied molecular orbitals contain 10 bonding electrons and 4 antibonding electrons. Calculate the bond order using $\text{bond order}=\frac{N_\text{bonding}-N_\text{antibonding}}{2}$, and explain what the result predicts about net bonding in this introductory model.

O₂ has eight bonding electrons and four antibonding electrons. In forming O₂⁺, one electron is removed from an antibonding π* orbital. Calculate the bond order of each species, identify which has the larger bond order, and state the general bond-strength and bond-length implication for these same atoms.

The occupied molecular orbitals of $\mathrm{O_2}$ contain eight bonding electrons and four antibonding electrons. In $\mathrm{O_2^+}$, one electron is removed from an antibonding $\pi^*$ orbital. Calculate the bond order of each species and explain what the change generally predicts about their relative bond strength and length.

A student claims that one s orbital and two p orbitals produce four hybrids in a literal physical transformation that every atom must complete before bonding. Identify and correct the errors in this claim, including what the orbital combination produces and how hybridization should be understood as a model.

For the oxygen atom in water, determine its common hybridization from the electron-density regions. Explain how those hybrid orbitals are used and why the H–O–H angle is about $104.5^\circ$ rather than the ideal tetrahedral angle of $109.5^\circ$.

In ethene, $\mathrm{C_2H_4}$, use the electron-density regions around each carbon to identify its common hybridization. Then explain which orbitals form the C=C bond's sigma and pi components, and count the total sigma and pi bonds in the molecule.

In the linear combination of atomic orbitals approach, two atomic orbitals can produce a bonding MO and an antibonding MO. Compare their relative energies and electron distributions between the nuclei, and state what properties of the atomic orbitals allow them to combine most effectively.

For $\mathrm{N_2}$, the two equal-energy $\pi(2p)$ orbitals lie below the $\sigma(2p)$ orbital. Consider the six electrons in these $2p$-derived orbitals. Describe their occupancy in order of increasing energy and use it to explain whether these electrons are paired or unpaired, and what that implies about the molecule’s magnetic behavior.

A diatomic species has nine electrons in bonding MOs and three electrons in antibonding MOs. Calculate its MO bond order and explain what that result implies about the bond. In your interpretation, state the general relationship between bond order, bond strength, and bond length for the same pair of atoms.