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Liquids: Class 11 Chemistry
Intermolecular forces, hydrogen bonding, evaporation, vapour pressure, boiling point, surface tension, viscosity, capillary action, liquid crystals and heats of fusion and vaporisation, with related calculations.
Key points
A liquid has a fixed volume but no fixed shape. Its particles are close together but can slide past one another. Liquids are almost incompressible, have higher density than gases and diffuse slowly. Their properties depend on the strength of the forces between the molecules.
Intermolecular forces are the attractions between molecules. They are much weaker than chemical bonds. The main types are:
- London dispersion forces: temporary dipoles in all molecules, even non-polar ones. They become stronger as molar mass and molecule size increase, which is why boiling points of the noble gases and halogens rise down the group.
- Dipole-dipole forces: attraction between the opposite ends of permanent dipoles in polar molecules such as HCl and CH₃Cl.
- Hydrogen bonding: a strong dipole-dipole attraction between hydrogen bonded to N, O or F and a lone pair on another N, O or F atom.
- Ion-dipole forces: attraction between ions and polar molecules, as when NaCl dissolves in water.
Hydrogen bonding explains many unusual properties. Water, HF and NH₃ boil at much higher temperatures than the hydrides of the other elements in their groups. Ethanol (78 °C) boils higher than dimethyl ether (−24 °C) although both have the same formula. Hydrogen bond energy is roughly 10 to 40 kJ mol⁻¹, far less than a covalent bond but much more than dispersion forces. Water molecules each form up to four hydrogen bonds.
Ice has an open hexagonal structure held by hydrogen bonds, with empty spaces. When ice melts, some bonds break and molecules pack closer, so water is denser than ice and ice floats. The density of water is greatest at 4 °C. This protects aquatic life because lakes freeze from the top.
Evaporation is the escape of molecules from the surface of a liquid at any temperature. Only molecules with enough kinetic energy to overcome the attractions escape, so the average energy of the remaining liquid falls and it cools. Evaporation is faster when temperature and surface area are larger, when intermolecular forces are weaker and when the air above is moving or dry. Sweating cools the body in this way.
Vapour pressure is the pressure exerted by the vapour in dynamic equilibrium with its liquid in a closed container, where the rate of evaporation equals the rate of condensation. It increases rapidly with temperature, is lower for liquids with stronger intermolecular forces, and does not depend on the amount of liquid or the size of the container. Diethyl ether has a high vapour pressure and is volatile, while glycerol has a low one.
The boiling point is the temperature at which the vapour pressure of a liquid equals the external pressure. At this point bubbles of vapour form throughout the liquid. The normal boiling point is the boiling point at 1 atm (101.325 kPa). At high altitude the external pressure is lower, so water boils below 100 °C. A pressure cooker raises the pressure, so water boils above 100 °C and food cooks faster.
Heat changes at a change of state occur at constant temperature. The molar heat of vaporisation (ΔHvap) is the heat needed to change one mole of liquid into vapour at its boiling point. For water it is 40.7 kJ mol⁻¹. The molar heat of fusion (ΔHfus) is the heat needed to melt one mole of solid at its melting point. For ice it is 6.01 kJ mol⁻¹. Heat supplied during a change of state does not raise the temperature. It overcomes intermolecular forces.
Surface tension is the energy required to increase the surface area of a liquid by one unit, or the force acting along a unit length of the surface. Its SI unit is N m⁻¹. Molecules in the bulk are pulled equally in all directions, but those at the surface are pulled inward, so the surface acts like a stretched skin. It makes drops spherical and lets small insects walk on water. Surface tension decreases as temperature rises and when detergents are added.
Viscosity is the resistance of a liquid to flow. Its SI unit is the pascal second (Pa s), and the older unit is the poise. Liquids with strong intermolecular forces or long chain molecules, such as honey, glycerol and engine oil, are more viscous. Viscosity of a liquid decreases when the temperature increases because molecules gain energy to overcome the forces.
Capillary action is the rise or fall of a liquid in a narrow tube. Water wets glass and rises because adhesive forces (liquid to glass) are greater than cohesive forces (liquid to liquid), giving a concave meniscus. Mercury does not wet glass, so cohesive forces dominate and it is depressed with a convex meniscus. Capillary action carries water up in plants and oil up a wick.
Liquid crystals flow like liquids but their molecules keep some order, like crystals. They are made of long rod-like molecules. The main types are nematic, smectic and cholesteric. Their orientation changes with temperature and with an electric field, which is used in liquid crystal displays (LCDs) of watches, calculators and screens.
Study tip: Learn the order of strength: covalent bond, then hydrogen bond, then dipole-dipole, then dispersion. Use it to explain every boiling point question. Always link a property (boiling point, viscosity, surface tension) to intermolecular forces, and for calculations convert grams to moles before using kJ mol⁻¹ values.
Definitions and flashcards
Say the meaning first, then open the card to check.
Intermolecular forces
Weak attractions between molecules, including dispersion, dipole-dipole and hydrogen bonds
Hydrogen bond
Attraction between H bonded to N, O or F and a lone pair on another N, O or F
Evaporation
Escape of molecules from the surface of a liquid below its boiling point
Vapour pressure
Pressure of the vapour in equilibrium with its liquid at a given temperature
Normal boiling point
The boiling point of a liquid at 1 atm pressure
Surface tension
Force per unit length (N m⁻¹) acting along the surface of a liquid
Viscosity
Resistance of a liquid to flow
Cohesion and adhesion
Attraction between like molecules and between unlike molecules
Liquid crystal
A state with the flow of a liquid and partial order of a crystal
Formula sheet
| Name | Formula | What it means |
|---|---|---|
| Heat of vaporisation | q = n × ΔHvap | Heat to vaporise n moles at the boiling point |
| Heat of fusion | q = n × ΔHfus | Heat to melt n moles at the melting point |
| Moles | n = mass / molar mass | Used with the two formulas above |
| Dry gas pressure | P(gas) = P(total) − P(water vapour) | For a gas collected over water |
Short questions with model answers
These are the kind of questions asked for 2, 3 and 5 marks. Try each one before opening the answer.
Why does water have a much higher boiling point than H₂S? (3 marks)
Water molecules form strong hydrogen bonds because oxygen is small and highly electronegative. H₂S forms only weak dipole-dipole and dispersion forces. More energy is needed to separate water molecules, so it boils higher.
Define vapour pressure. Name two factors that affect it. (2 marks)
Vapour pressure is the pressure exerted by the vapour in equilibrium with its liquid in a closed vessel at a given temperature. It increases with temperature and decreases as intermolecular forces become stronger.
Explain why evaporation causes cooling. (3 marks)
The molecules with the highest kinetic energy escape from the surface. The molecules left behind have a lower average kinetic energy, so the temperature of the liquid falls. Heat then flows in from the surroundings, which cools them.
Why does water boil below 100 °C on a high mountain? (3 marks)
A liquid boils when its vapour pressure equals the external pressure. On a mountain the atmospheric pressure is lower, so water reaches this vapour pressure at a lower temperature and boils below 100 °C.
Explain surface tension and name the factors that change it. (5 marks)
Molecules inside a liquid are pulled equally in all directions but surface molecules are pulled inward by the molecules below. The surface therefore tends to have the smallest area, like a stretched membrane. Surface tension is the force per unit length along the surface, measured in N m⁻¹. It decreases when the temperature rises and when detergent or soap is added. Stronger intermolecular forces give a higher value. It explains spherical drops and capillary rise.
Why does ice float on water? (2 marks)
In ice, hydrogen bonding holds molecules in an open hexagonal structure with empty spaces, so ice is less dense than liquid water.
Numericals with worked steps
Find the heat needed to vaporise 36 g of water at 100 °C. (ΔHvap = 40.7 kJ mol⁻¹, H = 1, O = 16)
- Molar mass of water = 18 g mol⁻¹
- n = 36/18 = 2.0 mol
- q = n × ΔHvap = 2.0 × 40.7 = 81.4 kJ
Answer: 81.4 kJ
Find the heat needed to melt 90 g of ice at 0 °C. (ΔHfus = 6.01 kJ mol⁻¹)
- n = 90/18 = 5.0 mol
- q = 5.0 × 6.01 = 30.05 kJ
Answer: 30.05 kJ (about 30.1 kJ)
Hydrogen is collected over water at 25 °C. The total pressure is 760 torr and the vapour pressure of water at 25 °C is 23.8 torr. Find the pressure of dry hydrogen.
- P(total) = P(H₂) + P(water vapour)
- P(H₂) = 760 − 23.8 = 736.2 torr
Answer: 736.2 torr
Find the heat taken to vaporise 23 g of ethanol at its boiling point. (ΔHvap = 38.6 kJ mol⁻¹, C = 12, H = 1, O = 16)
- Molar mass of C₂H₅OH = 46 g mol⁻¹
- n = 23/46 = 0.50 mol
- q = 0.50 × 38.6 = 19.3 kJ
Answer: 19.3 kJ
Common mistakes
- Calling hydrogen bonds covalent bonds. They are weak intermolecular attractions.
- Saying that temperature rises while a liquid boils. It stays constant until all the liquid has changed.
- Believing that vapour pressure depends on the volume of the container or the amount of liquid. It depends only on the liquid and the temperature.
- Using the molar heat with mass in grams. First convert the mass to moles.
- Stating that viscosity increases with temperature for liquids. It decreases.