To determine which of the given compounds has the lowest boiling point, we will analyze the hydrides of the elements in Group 15: NH3 (ammonia), PH3 (phosphine), AsH3 (arsine), and SbH3 (stibine).
### Step-by-Step Solution:
1. **Identify the Compounds**: The compounds given are NH3, PH3, AsH3, and SbH3. These are hydrides of nitrogen, phosphorus, arsenic, and antimony, respectively.
2. **Understand the Trend in Boiling Points**: As we move down the group from nitrogen to antimony, the size of the atoms increases. This increase in atomic size leads to an increase in the strength of Van der Waals forces (dispersion forces) between the molecules.
3. **Analyze Van der Waals Forces**: As the size of the hydride increases (from NH3 to SbH3), the Van der Waals forces also increase, which typically results in higher boiling points. Therefore, we would expect SbH3 to have the highest boiling point, followed by AsH3, then PH3, and finally NH3.
4. **Consider Hydrogen Bonding**: However, NH3 exhibits hydrogen bonding due to the presence of a highly electronegative nitrogen atom. Hydrogen bonding is a strong intermolecular force that significantly raises the boiling point of NH3 compared to what would be expected based solely on molecular size.
5. **Compare Boiling Points**:
- NH3 has a boiling point that is higher than expected due to hydrogen bonding.
- PH3, being larger than NH3 and lacking significant hydrogen bonding, has a lower boiling point than NH3.
- AsH3 and SbH3, while larger, have boiling points that are also higher than PH3 due to their larger molecular size and corresponding Van der Waals forces.
6. **Conclusion**: Based on the analysis, PH3 has the lowest boiling point among the given hydrides.
### Final Answer:
The compound with the lowest boiling point is **PH3** (Option B).
