Anomalous Properties of Water
A concise revision guide to the anomalous properties of water: its high melting point and boiling point, the open structure and low density of ice, surface tension and the Group 4 to 7 hydride graph, all explained by hydrogen bonding, for AQA A Level Chemistry.
Why Water Is Anomalous
Water has a much higher melting point and boiling point, a solid that is less dense than its liquid, and a higher surface tension than any molecule of its size should have. Every one of these anomalous properties is caused by hydrogen bonding. A molecule with only 10 electrons would be expected to boil below -80 °C if induced dipole-dipole forces were the only force present; water boils at 100 °C.
Each water molecule can form up to four hydrogen bonds, more than any other small molecule, so a very large amount of energy is needed to separate the molecules. This is the single explanation behind all the anomalies on this page.
Key idea: Two O-H bonds and two lone pairs per molecule mean water forms the maximum number of hydrogen bonds, and every anomalous property follows from that.
High Melting point and Boiling point
The hydrides of Group 6 show the effect most clearly. Going down the group from H₂S to H₂Te the boiling point rises steadily because the number of electrons rises and the induced dipole-dipole forces strengthen. Extending that trend upwards predicts a boiling point for water of about -80 °C. The real value is 100 °C, because hydrogen bonds between water molecules are far stronger than the induced dipole-dipole forces and permanent dipole-dipole forces that hold the other hydrides together.
| Hydride | Electrons | boiling point / °C | Strongest intermolecular force |
|---|---|---|---|
| H₂O | 10 | 100 | hydrogen bonds |
| H₂S | 18 | -60 | permanent dipole-dipole forces |
| H₂Se | 36 | -41 | permanent dipole-dipole forces (induced dipole-dipole forces growing) |
| H₂Te | 54 | -2 | induced dipole-dipole forces dominate |
NH₃, H₂O and HF sit far above the trend of their groups because of hydrogen bonding; CH₄ does not, because carbon is not electronegative enough for C-H to hydrogen bond.
The same graph shows that methane is not anomalous: carbon is not electronegative enough for a C-H hydrogen to form a hydrogen bond, so the Group 4 hydrides follow a smooth induced dipole-dipole forces trend. Ammonia and hydrogen fluoride are anomalous but less so than water, because each forms on average only one hydrogen bond per molecule.
Exam sentence: Water has a much higher boiling point than H₂S because water molecules form hydrogen bonds, which are much stronger than the permanent dipole-dipole forces and induced dipole-dipole forces between H₂S molecules, so much more energy is needed to separate water molecules.
The Density of Ice
Ice floats on water because it is less dense: 0.92 g cm⁻³ against 1.00 g cm⁻³ for water at 4 °C. Almost every other substance is denser as a solid than as a liquid, because the particles pack more closely when they stop moving. Water is different because of the geometry of its hydrogen bonds.
In ice each water molecule is held by four hydrogen bonds to four neighbours arranged tetrahedrally around it, and the hydrogen bonds are fixed in place. This tetrahedral arrangement builds an open hexagonal lattice with large empty spaces, like a honeycomb, so the molecules are held further apart than they would be if they could pack freely. When ice melts, some of the hydrogen bonds break and the molecules can move into the spaces, so the liquid is denser than the solid. Water reaches its maximum density at 4 °C; above that, thermal expansion takes over.
The Open Hexagonal Lattice of Ice, and Why It Collapses on Melting
Drag to rotate, scroll or pinch to zoom. Every molecule in the ice lattice is held by four hydrogen bonds in a tetrahedral arrangement, which forces open hexagonal channels through the solid. Switch to liquid water to watch the channels collapse.
© Dr. Mohammed Al-Fatah – onlinelearningsystem.net
This has consequences beyond the exam: ice on a pond insulates the water below it, so aquatic life survives winter, and water expanding as it freezes shatters rocks and bursts pipes.
Exam sentence: Ice is less dense than water because hydrogen bonds hold the water molecules in a fixed, open tetrahedral lattice with large spaces; when ice melts some hydrogen bonds break and the molecules move closer together.
Surface Tension
Water has a high surface tension because the molecules at the surface are pulled inwards and sideways by hydrogen bonds to their neighbours, with no molecules above to balance the pull. The surface therefore behaves like a stretched skin that resists being broken or increased in area. This is why water forms near-spherical droplets, why a steel needle can be floated on water and why pond skaters walk on it.
Liquids held only by induced dipole-dipole forces, such as hexane, have a much lower surface tension because far less energy is needed to create new surface. Adding a detergent to water lowers its surface tension by placing molecules at the surface that cannot hydrogen bond as strongly.
Exam focus: Surface tension is not named in the AQA specification, but it appears in data-response questions; explain it with the same hydrogen-bonding argument.
Other Consequences of Hydrogen Bonding in Water
Water also has an unusually high specific heat capacity and a high enthalpy change of vaporisation, because energy supplied to water is used to break hydrogen bonds before the molecules can move faster or escape as a gas. These properties make water an excellent coolant and moderate the climate of coastal regions. Water is also an excellent solvent for ionic compounds and for molecules that can hydrogen bond, which the Choosing Solvents page explains.
| Property | Value for water | Comparison | Cause |
|---|---|---|---|
| Boiling point | 100 °C | H₂S: -60 °C | four hydrogen bonds per molecule |
| Density of solid | 0.92 g cm⁻³ | liquid: 1.00 g cm⁻³ | open hydrogen-bonded lattice |
| Surface tension | 72 mN m⁻¹ | hexane: 18 mN m⁻¹ | surface molecules pulled inwards by hydrogen bonds |
| Specific heat capacity | 4.18 J g⁻¹ K⁻¹ | ethanol: 2.44 J g⁻¹ K⁻¹ | energy breaks hydrogen bonds before raising temperature |
Key idea: Whatever the anomalous property, the answer begins “because water molecules form hydrogen bonds” and ends with what those bonds do.
Common Exam Points
Explain why ice floats
Hydrogen bonds hold molecules in an open tetrahedral lattice with spaces; on melting some bonds break and molecules move closer, so the liquid is denser.
Explain the boiling point of water compared with H₂S
Hydrogen bonds between water molecules are stronger than the permanent dipole-dipole forces and induced dipole-dipole forces between H₂S molecules.
Why is CH₄ not anomalous?
Carbon is not electronegative enough for C-H hydrogens to form hydrogen bonds.
Do not say
“Ice has stronger hydrogen bonds than water” (the bonds are the same strength; ice has more of them held rigidly); “covalent bonds break when ice melts”.
Check Your Understanding
Apply the hydrogen-bonding explanation to data and substances that are not discussed above.
FAQs
Use these quick answers to check the explanations for water that AQA questions ask for.
Why is the boiling point of water so high?
Each water molecule can form up to four hydrogen bonds, which are far stronger than the permanent dipole-dipole forces and induced dipole-dipole forces holding together hydrides such as H₂S. Extending the Group 6 trend would predict about -80 °C; water boils at 100 °C.
Why does ice float on water?
In ice the hydrogen bonds hold each molecule to four neighbours in a fixed tetrahedral arrangement, building an open hexagonal lattice with large spaces. When ice melts some hydrogen bonds break and the molecules move closer together, so the liquid is denser.
Why is methane not anomalous?
Carbon is not electronegative enough for the hydrogen atoms in CH₄ to carry a large δ+ charge, so methane cannot hydrogen bond. The Group 4 hydrides follow a smooth trend set by their number of electrons.
Why does water have a high surface tension?
Molecules at the surface have no neighbours above them, so the hydrogen bonds pull them inwards and sideways. Making new surface therefore costs energy, which is why water forms beads and supports a floating needle.
Are the hydrogen bonds in ice stronger than in water?
No, they are the same strength. Ice simply has more of them and holds them in fixed positions, which is what produces the open structure.
Copyright and author footprint: This OLS revision page was written for Online Learning System by Dr. Mohammed Al-Fatah. It is designed for A Level Chemistry revision and should not be copied or redistributed without permission.
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