Phosphorous hydride's bond angle is closer to 90° while the bond angle of ammonium hydride is closer to 104.5°. Among the following, which one explains this structural feature correctly?
Correct answer: C. Due to the significant energy difference between 3s and 3p orbitals, the lone pair on phosphorous prefers to occupy an unhybridized s-orbital rather than a hybridized sp3 orbital.
- A. Nitrogen bond pair electron cloud concentrates near the central atom because of its higher electronegativity, thus the bond pair - bond pair repulsion increases which in turn decreases the bond angle in NH3.
- B. Due to larger size of the lone pair electron cloud, there is larger lone pair - bond pair repulsion in PH3 compared to NH3.
- C. Due to the significant energy difference between 3s and 3p orbitals, the lone pair on phosphorous prefers to occupy an unhybridized s-orbital rather than a hybridized sp3 orbital.
- D. Phosphorous forms pπ - dπ bonds while nitrogen doesn't.
Explanation
The correct explanation for the difference in bond angles between PH3 and NH3 lies in the high energy difference between the 3s and 3p orbitals in phosphorus. This causes the lone pair to reside in the unhybridized 3s orbital, leading to a bond angle closer to 90° in PH3. In contrast, nitrogen forms sp3 hybridized orbitals, resulting in larger bond angles due to greater hybridization and repulsion in NH3. The other options incorrectly attribute the differences to factors such as electronegativity or the availability of d-orbitals, which do not accurately reflect the primary reason for the observed bond angles.
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