Anomalous Properties of Water

Learning Objectives (OCR Spec 2.2.2)

By the end of this lesson you should be able to:

• Identify the anomalous physical properties of water compared with other Group 16 hydrides
• Explain each anomalous property in terms of hydrogen bonding
• Describe the structure of ice and explain why it is less dense than liquid water
• Appreciate the biological importance of water's properties

Why Water Is Anomalous

Water (H2O) has several physical properties that appear contradictory if you only consider its small molar mass (Mr = 18):

• Unusually high melting and boiling points
• Solid is less dense than liquid (ice floats)
• High specific heat capacity and high latent heats
• High surface tension
• Universal solvent for polar substances

Each of these is explained by the extensive hydrogen bonding in liquid and solid water.

Comparing H2O with Other Group 16 Hydrides

Consider the Group 16 hydrides: H2O, H2S, H2Se, H2Te. Based on molecular mass alone, one would expect boiling points to rise steadily down the group as London forces increase. The actual data tell a very different story:

Hydride	Mr	Boiling point (degC)
H2O	18	+100
H2S	34	-61
H2Se	81	-41
H2Te	130	-2

If water followed the trend of the other hydrides, extrapolation gives a predicted boiling point of around -80 degC. The observed +100 degC is ~180 degC higher than expected.

This enormous difference is because liquid water contains extensive hydrogen bonding networks, which H2S and the heavier hydrides lack (S, Se, Te are less electronegative and larger, so they do not support hydrogen bonding). Water molecules need far more energy to escape into the gas phase.

Anomaly 1: High Melting and Boiling Points

Each H2O molecule can form up to four hydrogen bonds: the O has two lone pairs (acceptors) and two H atoms (donors).

       H
        \
         O . . . H
        /           \
       H             O -- H
       .            /
       .           H
        .
        O -- H
       /
      H

The result is a 3D network in which each molecule is H-bonded to several neighbours. To melt ice or boil water you must supply enough energy to break or disrupt these H bonds - hence the high m.p. (0 degC) and b.p. (100 degC).

Numerically: a hydrogen bond in water is about 20 kJ mol-1; there are on average about 3-4 per molecule in liquid water, contributing roughly 40 kJ mol-1 to the enthalpy of vaporisation of 41 kJ mol-1.

Anomaly 2: Ice Is Less Dense than Liquid Water

For almost all substances, the solid is denser than the liquid because particles pack more closely in the ordered solid. Water is the striking exception: ice at 0 degC has a density of 0.917 g cm-3, while liquid water at 0 degC is 1.000 g cm-3. Ice therefore floats.

Why?

In ice, each water molecule forms 4 hydrogen bonds to its neighbours in a tetrahedral arrangement around O. This rigid geometry gives a very open, hexagonal lattice with large gaps between the molecules.

        H             H
        |             |
    H - O . . . H - O - H
        |             :
        :             :
        H             H
        |             |
    H - O . . . H - O - H
        |             |

When ice melts, the tetrahedral lattice partially collapses - about 10% of the H bonds break - allowing water molecules to pack more efficiently. The liquid is therefore denser than the solid.

Density and Temperature

Liquid water actually reaches its maximum density at ~4 degC, not at 0 degC. Between 0 and 4 degC, collapse of residual ice-like structure dominates, increasing density. Above 4 degC, ordinary thermal expansion dominates, decreasing density.

Biological Importance

• Ice floats on ponds and lakes, forming an insulating layer that keeps the water beneath from freezing, enabling aquatic life to survive winter
• If ice were denser than water, lakes would freeze from the bottom up, killing most aquatic life

Anomaly 3: High Specific Heat Capacity