Draw the Lewis structure for HF.

Draw the Lewis structure for HF. Identify the bond order and discuss the polarity of the molecule based on electronegativity: Draw the molecular orbital (MO) diagram for HF (see Figure 2.14). Compare and contrast the Lewis and MO descriptions of HF. For any orbitals which are nonbonding; describe why they are nonbonding: Explain how an electrophile and a nucleophile would attack HF_

The Correct Answer and Explanation is:

Correct Answer:

1. Lewis structure of HF:

rCopyEditH : F

Hydrogen has one valence electron, fluorine has seven. They share one pair (single bond), giving hydrogen two electrons (full shell) and fluorine eight (full octet), with three lone pairs left on fluorine.

2. Bond order:
   Since there is one single bond, the bond order is 1.
3. Polarity:
   Fluorine is much more electronegative than hydrogen (F ≈ 3.98, H ≈ 2.20), causing the shared electrons to be pulled toward fluorine. This creates a partial negative charge on fluorine and a partial positive charge on hydrogen. Therefore, HF is a highly polar molecule with a dipole pointing toward fluorine.
4. Molecular Orbital (MO) diagram:

• The main interaction involves H’s 1s orbital overlapping with F’s 2p_z orbital (assuming z is the bond axis).
• The 2p_x and 2p_y orbitals of fluorine do not overlap with H’s 1s orbital, so they remain nonbonding.
• The filled 2s orbital of fluorine is much lower in energy than hydrogen’s 1s and does not mix significantly, also staying nonbonding.
• The MO diagram has:
  • A bonding sigma orbital from the combination of H 1s and F 2p_z.
  • An antibonding sigma* orbital.
  • Nonbonding orbitals corresponding to F’s 2s, 2p_x, and 2p_y.

5. Nonbonding orbitals:
   The 2s of fluorine and the 2p_x and 2p_y orbitals remain nonbonding because they have no suitable counterpart in hydrogen’s 1s to interact with (either they are too low in energy like 2s or have no directional overlap like 2p_x and 2p_y).
6. Lewis vs. MO description:
   The Lewis structure shows a single shared pair of electrons completing octets, but does not account for the nonbonding orbitals explicitly. The MO diagram shows both the bonding interaction and the presence of nonbonding fluorine orbitals, offering a more detailed picture of electron distribution.
7. Electrophile and nucleophile attack:
   Because fluorine carries a partial negative charge, an electrophile (electron-pair acceptor) would attack the fluorine atom. Conversely, a nucleophile (electron-pair donor) would be attracted to the partially positive hydrogen atom, attacking it because it is electron-deficient.

Explanation:

The Lewis structure is the simplest way to represent bonding by showing atoms sharing electrons to achieve stable configurations. HF’s single bond satisfies the duet rule for hydrogen and the octet for fluorine. However, it does not explain energy levels or orbital interactions. The molecular orbital approach shows that bonding in HF involves the mixing of H’s 1s with F’s 2p_z orbital, creating bonding and antibonding orbitals, while fluorine’s 2s and the two 2p orbitals perpendicular to the bond remain nonbonding because they do not have symmetry or suitable energy match to overlap with hydrogen’s 1s orbital. The polarity of HF, revealed both in Lewis and MO descriptions, arises from fluorine’s high electronegativity, which pulls electron density toward itself. This polarity makes HF highly reactive with both electrophiles and nucleophiles: nucleophiles attack the proton (hydrogen), while electrophiles can accept electrons from the fluorine’s lone pairs. This combination of Lewis and MO perspectives gives a comprehensive understanding of HF’s structure, bonding, and reactivity.