Hydrogen Bond

Question

Water (H₂O) has an unusually high boiling point of 100 °C compared to H₂S (−60 °C), despite both having bent geometry. The primary reason is:

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1. **Compare Molecular Properties**: H₂O (M = 18 g mol⁻¹) actually has a *lower* molecular mass than H₂S (M = 34 g mol⁻¹). If only van der Waals forces mattered, H₂S would have the higher boiling point.
2. **Identify the Anomaly**: The boiling point of H₂O (100 °C) is dramatically higher than that of H₂S (−60 °C), a difference of 160 °C that cannot be explained by dispersion forces alone.
3. **Role of Hydrogen Bonding**: Oxygen is highly electronegative (3.44) and has two lone pairs. Each H₂O molecule can form up to four hydrogen bonds (two as donor via O–H, two as acceptor via lone pairs), creating an extensive 3D hydrogen-bonded network.
4. **Why H₂S Cannot Form Strong H-bonds**: Sulfur has lower electronegativity (2.58) and larger atomic size, making S–H···S hydrogen bonds very weak and negligible.
5. **Eliminate Other Options**: Higher molecular mass would favour H₂S, not H₂O. London forces are weaker in the lighter H₂O. Bond angle affects shape but not boiling point directly.
6. **Conclusion**: The extensive intermolecular hydrogen bonding network in water requires significantly more energy to disrupt, resulting in its anomalously high boiling point.

Hydrogen Bond

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