UNIT 2 CHEMISTRY · PEARSON AAQ · B1.8
Hydrogen Bonding and the Properties of Water
FREE CHEMISTRY REVISION · TOPIC 17
Explain water’s unusual properties from its hydrogen-bond network.
Link molecular polarity and network structure to phase changes, floating ice, density and surface tension.
Pearson BTEC Level 3 AAQ Applied Science
Unit 2: Principles and Applications of Chemistry · B1.8
What you will learn
Use water’s polar, bent molecules and hydrogen-bond network to explain its melting and boiling points, why ordinary ice is less dense than liquid water, and why water has relatively high surface tension. Distinguish molecular geometry from network geometry.
1. Why water forms an extensive hydrogen-bond network
Oxygen is more electronegative than hydrogen, so each O–H bond is polar: O is δ− and H is δ+. Two lone pairs on oxygen give a bent H₂O molecule with a bond angle of about 104.5°. The bond dipoles do not cancel; the molecule has a permanent dipole.
A hydrogen bond forms between an H attached to O in one molecule and an oxygen lone pair in another: O–H···O. The dotted H···O link is intermolecular; the solid O–H bond is covalent. Water also has London and permanent-dipole attractions. Hydrogen bonding provides a particularly important additional contribution to its unusual properties.
Each H₂O can donate two hydrogen bonds through its two hydrogens and accept two through its two lone pairs. In ideal ordinary ice, each molecule has four hydrogen-bonded neighbours in an open, approximately tetrahedral network. In liquid water bonds continually break and reform; the number and arrangement fluctuate. Do not assume every liquid molecule has four permanent links at every instant.
2. Melting and boiling points
At approximately 101.3 kPa (one atmosphere), pure water melts at about 0°C and boils at about 100°C. These values depend on pressure and purity. Boiling at 100°C is not a universal rule at every altitude or pressure.
Compared with similar small molecular hydrides such as H₂S, water has unusually high melting and boiling points. Water molecules can form an extensive hydrogen-bond network. More energy is needed to disrupt and overcome these attractions, so higher temperatures are required for the phase changes than would be expected from its small molecular size alone.
Melting disrupts the ordered ice network sufficiently to allow molecules to move past one another. Many hydrogen bonds persist in the liquid, where they continuously rearrange. It is wrong to say that every hydrogen bond has been removed on melting.
Boiling separates molecules widely into the gas, overcoming most of the intermolecular cohesion. The O–H covalent bonds within the H₂O molecules largely remain intact. Heating water into steam is a physical change, not decomposition into H₂ and O₂. The energy supplied during a phase change increases intermolecular potential energy rather than simply increasing temperature while the phase change proceeds at fixed pressure.
3. Why ice is less dense than liquid water
Density = mass ÷ volume. In ordinary ice, directional hydrogen bonding maintains an open network with spaces between molecules. On melting, some of this ordered structure collapses and molecules can pack closer together. The same mass therefore occupies a smaller volume as liquid water near the melting point, giving a higher density.
Approximate useful values are 0.917 g cm⁻³ for ordinary ice near 0°C and 1.00 g cm⁻³ for liquid water near 4°C. Because ice is less dense, it floats. The molecules themselves have not lost atoms or become lighter; the average arrangement and volume have changed.
An ice layer forms at the surface of a lake and helps insulate liquid water below. Expansion on freezing can damage a full, sealed water container because its contents require more volume. This explanation applies to ordinary ice at everyday pressures; water has other high-pressure solid forms with different densities.
4. The density maximum near 4°C
Liquid water reaches its maximum density near 4°C at ordinary pressure. Cooling from a higher temperature initially reduces ordinary thermal expansion, allowing closer packing. Below about 4°C, increasing local open hydrogen-bond order becomes important enough to increase volume as the liquid cools towards freezing.
Consequently warming water from 0°C towards 4°C makes it denser, whereas warming it above approximately 4°C generally makes it less dense. Do not claim that liquid water always contracts on heating or always expands on cooling. The anomaly is a balance between thermal motion and changing network structure.
5. Surface tension and cohesive attraction
Surface tension is the tendency of a liquid surface to resist an increase in area. It can be expressed as force per unit length (N m⁻¹), or equivalently as energy per unit area. Hydrogen bonding contributes to strong cohesion between water molecules and relatively high surface tension compared with many common molecular liquids.
Within the bulk, a molecule has water neighbours in all directions, so attractions approximately balance. At a water–air surface it has fewer water neighbours above. Creating more surface places more molecules in this less favourable environment and costs energy, so the surface tends to minimise its area. The simple “inward attraction” picture is a useful model; the surface is not a rigid skin made from different molecules.
Small droplets tend towards a rounded shape because a sphere minimises surface area for a given volume; gravity distorts larger drops. Water striders can be supported by surface forces acting along the contact lines of their water-repellent legs. Their support is not explained by assuming their body material is less dense than water.
Detergents contain surfactants that adsorb at the interface and lower surface tension, helping water spread and wet surfaces. They alter interfacial interactions; they do not need to break the covalent O–H bonds. Increasing temperature generally decreases water’s surface tension as thermal motion competes with cohesion.
| Observation | Molecular explanation | Avoid this claim |
|---|---|---|
| Relatively high boiling point | Significant energy to overcome intermolecular hydrogen bonding | Boiling breaks every O–H covalent bond |
| Ice floats | Open solid network gives greater volume per unit mass | Ice molecules weigh less |
| Liquid remains cohesive | Many hydrogen bonds remain and rearrange | Melting removes all hydrogen bonds |
| Rounded small droplets | Making extra surface costs energy; area tends to decrease | Water has a permanent solid skin |
Worked examples
Example 1 · Calculate the volume of ice
Given mass = 100 g and density = 0.917 g cm⁻³, use V = m/ρ. V = 100 ÷ 0.917 = 109 cm³ to 3 significant figures. The same mass of water at density 1.00 g cm⁻³ occupies 100 cm³. The ice volume is about 9% larger; mass is unchanged.
Example 2 · Explain floating using density
Compare equal volumes, 10.0 cm³. Ice mass ≈ 0.917 × 10.0 = 9.17 g; liquid mass ≈ 1.00 × 10.0 = 10.0 g. Ice is less dense because of its open network, so it floats. This equal-volume comparison must not be confused with the equal-mass calculation above.
Example 3 · Explain water versus H₂S
Both are small Group 16 hydrides. In the usual comparison H₂S lacks the extensive strong O–H···O hydrogen bonding found in water. Water therefore has stronger total intermolecular cohesion than its small size alone suggests and a much higher boiling point. Both substances still have London forces.
Example 4 · Distinguish the two angles
The H–O–H angle within a water molecule is about 104.5°. The directions towards four neighbouring O atoms in the ideal ice network are approximately tetrahedral, near 109.5°. These angles describe different sets of atoms; the ice network does not turn each water molecule into H₄O.
Example 5 · Explain detergent spreading
A drop on a suitable surface spreads more readily when surfactant lowers the surface tension. Explain the change through adsorption and altered interfacial interactions. No new ions or broken O–H covalent bonds are required to explain the reduction in the energetic cost of surface area.
Common misconceptions and exam tips
“Ice is denser because it is solid.” Ordinary ice is an exception: its open network gives lower density.
“Ice molecules are lighter.” A molecule remains H₂O; spacing changes.
“Water boils at 100°C at all pressures.” Quote the pressure condition.
“All hydrogen bonds disappear on melting.” The liquid retains a dynamic network.
“The 104.5° angle describes all ice neighbours.” Distinguish intramolecular H–O–H from the network’s O-neighbour geometry.
“Surface tension is a solid skin.” It is an interfacial property of a dynamic liquid.
Link every explanation to the property asked: energy for phase-change temperatures, volume per unit mass for density, and cohesion and surface area for surface tension. Label hydrogen bonds with dotted lines and covalent bonds with solid lines. Use m, V and ρ with consistent units in calculations.
Your revision summary
Polar bent H₂O molecules form an extensive hydrogen-bond network. Its attractions raise phase-change temperatures; its open arrangement in ordinary ice lowers density and allows floating. Many hydrogen bonds persist dynamically in the liquid. Cohesion contributes to high surface tension and the tendency to minimise surface area.
Check your understanding
Choose an answer, select Check Answer and read the scientific explanation. Try Again clears the choice and feedback. Keyboard: Tab between controls, arrow keys within radio groups, Enter or Space on buttons. Answers are not saved.