Why Is Water a Liquid at Room Temperature? Role of Hydrogen Bonding

A hydrogen bond forms when a hydrogen atom covalently bonded to a highly electronegative atom (O, N, or F) interacts electrostatically with the lone pair on another electronegative atom.

In water (H₂O):

• Each oxygen atom is electronegative (electronegativity = 3.44)
• The O–H bonds are highly polar (O has partial negative charge δ⁻, H has partial positive δ⁺)
• Each water molecule can form up to 4 hydrogen bonds:
  • 2 as a hydrogen bond donor (through its 2 O–H bonds)
  • 2 as a hydrogen bond acceptor (through its 2 lone pairs on O)

Hydrogen bond strength in water: ~20 kJ/mol per H-bond (Much weaker than covalent bonds ~400 kJ/mol but much stronger than van der Waals forces ~1–5 kJ/mol)

In liquid water, molecules form a dynamic, fluctuating network of 3–4 hydrogen bonds per molecule.

To convert a liquid to a gas, enough energy must be supplied to overcome intermolecular forces.

For water at room temperature (25°C):

• The thermal energy (kT ≈ 2.5 kJ/mol) is insufficient to break the extensive H-bond network (~20 kJ/mol per bond × multiple bonds per molecule)
• Therefore, water molecules cannot escape the liquid phase at room temperature
• Water remains liquid until 100°C, when sufficient thermal energy overcomes H-bonding

Comparison with H₂S (hydrogen sulphide):

• H₂S has almost no H-bonding (S is not electronegative enough)
• Boiling point of H₂S = −60°C → gas at room temperature
• Molecular mass of H₂S (34 g/mol) > H₂O (18 g/mol)
• Despite heavier molecules, H₂S is a gas because it lacks H-bonding

This comparison shows clearly that hydrogen bonding, not molecular mass, determines water's liquid state at room temperature.

Water's ability to form hydrogen bonds arises from its polar nature:

• H₂O is a bent/V-shaped molecule (bond angle ≈ 104.5°)
• The two lone pairs on oxygen push the hydrogen atoms, creating an asymmetric shape
• This gives H₂O a permanent dipole moment (1.84 D — Debye)
• Oxygen carries a partial negative charge (δ⁻) and each hydrogen a partial positive charge (δ⁺)

Molecular properties:

• Shape: Bent (V-shaped)
• Bond angle: 104.5° (less than tetrahedral 109.5° due to lone pair repulsion)
• Hybridisation: sp³
• Dipole moment: 1.85 D
• Molecular weight: 18 g/mol

The polar nature and bent shape of water make it an effective hydrogen bond former, which in turn makes it liquid at room temperature.

Hydrogen bonding explains several unusual properties of water:

1. High boiling point (100°C): Far higher than predicted from molecular weight. H₂S bp = −60°C, H₂Se bp = −41°C, H₂Te bp = −2°C, but H₂O bp = 100°C (anomaly)

2. High specific heat capacity (4.18 J/g·K): Water absorbs a lot of heat before its temperature rises — important for climate regulation

3. High latent heat of vaporisation (2257 J/g): A lot of energy needed to convert liquid water to steam — important for sweating and cooling

4. Density maximum at 4°C: Ice is less dense than liquid water (ice floats) — H-bonds form a more open tetrahedral lattice in ice. This protects aquatic life in winter.

5. Universal solvent: Polar H₂O dissolves many ionic and polar compounds (important for biochemistry)

6. Surface tension: High surface tension due to H-bonding — water striders can walk on water