Hydrogen Bond
Easychemistry
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:
Select the correct option:
Solution
Incorrect! Answer:
Extensive intermolecular hydrogen bonding in H₂O
- 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.
- 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.
- 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.
- 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.
- 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.
- Conclusion: The extensive intermolecular hydrogen bonding network in water requires significantly more energy to disrupt, resulting in its anomalously high boiling point.
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About This Question
- Subject
- chemistry
- Chapter
- chemical bonding and molecular structure
- Topic
- hydrogen bond
- Difficulty
- Easy
- Year
- 2025
Solution
Correct Answer:
Extensive intermolecular hydrogen bonding in H₂O
- 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.
- 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.
- 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.
- 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.
- 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.
- Conclusion: The extensive intermolecular hydrogen bonding network in water requires significantly more energy to disrupt, resulting in its anomalously high boiling point.
This easy difficulty chemistry question is from the chapter chemical bonding and molecular structure, covering the topic of hydrogen bond. It appeared in the 2025 exam.
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