States and phases of matter

Class 11 Chemistry, unit 5, book pages 96–113

Practise 67 MCQs Take the unit quiz Review 25 cards

Matter exists as a gas, a liquid or a solid, and the kinetic molecular theory explains how the particles behave in each phase. This unit studies the properties of liquids (evaporation, vapour pressure, boiling point, viscosity and surface tension), uses hydrogen bonding to explain the unusual properties of water, and looks at the energy of phase changes and at liquid crystals. It ends with the general properties of gases and of solids, including crystalline and amorphous solids.

Learning outcomes (14)

Describe simple properties of liquids e.g., diffusion, compression, expansion, motion of molecules, spaces between them, intermolecular forces and kinetic energy based on kinetic molecular theory.
Describe types of intermolecular forces.
Explain the strength and applications of dipole-dipole forces, hydrogen bonding and London forces.
Describe physical properties of liquids e.g., evaporation, vapour pressure, boiling point, viscosity and surface tension.
Apply the concept of hydrogen bonding to explain the properties of water (specifically high surface tension, high specific heat, low vapour pressure, high heat of vaporization, and high boiling point).
Define molar heat of fusion and molar heat of vaporization.
Describe how heat of fusion and heat of vaporization affect the particles that make up matter.
Outline the importance of heat of fusion in the study of glaciers and ice sheets (particularly while studying polar ice caps).
Describe the physical properties of gases (including compressibility, expandability, and pressure exerted by gases).
Describe liquid crystals and give their uses in daily life.
Differentiate liquid crystals from pure liquids and crystalline solids.
Describe simple properties of solids e.g., compression, expansion, motion of molecules, inter particle space, intermolecular forces and kinetic energy based on kinetic molecular theory.
Differentiate between amorphous and crystalline solids.
Describe properties of crystalline solids like geometrical shape, melting point, habit of a crystal, cleavage, and crystal growth.

Phases of matter

Phases within a system exist as gas, liquid or solid. The matter within a system contains only a single phase of the gaseous phase, but it may have one or more liquid or solid phases.

All the phases of matter are made up of microscopic particles which differ in their behaviour in the three phases. The unit explains this behaviour with the kinetic molecular theory.

The particles of a solid are close together in fixed positions, those of a liquid are close together but can slide past each other, and those of a gas are far apart and free to move.

5.1Kinetic molecular interpretation of liquidsOften examined

The postulates of the kinetic molecular theory of liquids are given below.

A liquid consists of molecules in contact with each other.
Molecules within a liquid are in constant motion, but the motion of the molecules is limited by their close packing.
The attractive forces of liquid molecules are greater than the attractive forces of gas molecules. However, these attractive forces are not sufficient to hold the molecules in a fixed position. Liquid molecules can slide past each other.
The average kinetic energy of liquid molecules is directly proportional to the absolute temperature.
At constant temperature, the average kinetic energy of the molecules of a liquid is equal to the kinetic energy of the vapours of the liquid.

5.1.1Properties of liquids from kinetic molecular theoryOften examined

There are some common properties of liquids that can be explained by the kinetic molecular theory.

1. Diffusion

Diffusion in liquids occurs because molecules move from one place to another. The restricted movement of a molecule reduces the rate of diffusion. For example, an ink drop added to water spreads slowly because the spaces between the molecules are relatively small. Diffusion in densely packed liquids is slow because there are fewer collisions between the molecules.

2. Compression (effect of pressure)

By raising the pressure, a liquid cannot be significantly compressed, because its molecules are already in close contact with each other. For example, an increase in pressure from one to two atmospheres reduces the volume of water by only 0.0045 per cent, which is negligible. The same pressure reduces the volume of a gas by up to 50 per cent.

3. Expansion (effect of temperature)

Liquids expand when heated because the intermolecular forces between the molecules decrease. In addition, an increase in temperature increases the effective collisions between molecules. As the temperature drops, the volume decreases.

4. Motion of molecules

The molecules of a liquid move with lesser speed because of the larger forces of attraction among them, so they have lesser kinetic energy. However, the kinetic energy increases with an increase in temperature.

5. Spaces between the molecules

The molecules of a liquid are quite close to each other, so there is very little space between them. As a result, the rate of intermolecular collisions is moderate, and therefore the average kinetic energy is also moderate.

6. Intermolecular forces

The forces of attraction between individual particles of matter are called intermolecular forces. The physical properties of liquids, such as boiling point, vapour pressure, surface tension, viscosity and heat of vaporization, depend on the strength of the attractive forces between the molecules.

Intermolecular forces: The forces of attraction between molecules of a substance in any phase.
Dipole-dipole forces: The forces of attraction between the positive end of one molecule and the negative end of another molecule.
Dipole-induced dipole forces (Debye forces): The attractive forces between dipoles and induced dipoles.
London forces (instantaneous dipole-induced dipole forces): The momentary forces of attraction created between an instantaneous dipole and an induced dipole.

7. Kinetic energy based on kinetic molecular theory

According to the kinetic molecular theory, molecular movements and collisions caused by strong intermolecular attraction are minimal. Take water as an example: its molecules are close together and attract each other strongly because of hydrogen bonding, so they have low kinetic energy.

Properties of liquids explained by the kinetic molecular theory
Property	Explanation from kinetic molecular theory
Diffusion	Slow: small spaces between molecules, so few collisions and restricted movement
Compression	Negligible: molecules are already in close contact
Expansion	Occurs on heating: intermolecular forces decrease and effective collisions increase
Motion of molecules	Slow, because of large attractive forces; speed and kinetic energy rise with temperature
Spaces between molecules	Very little space; moderate collision rate and moderate kinetic energy
Intermolecular forces	Strong enough to keep molecules close but not in fixed positions

5.2.1Vapour pressure and boiling pointOften examined

When a liquid is heated, its vapour pressure increases with increasing temperature because the intermolecular forces decrease. As a result, more and more vapour is released into the air. A stage is reached where the liquid begins to boil. The temperature at which the vapour pressure of a liquid is equal to the atmospheric pressure or some external pressure is called the boiling point of the liquid.

Boiling point of water at two pressures (the book's example)
External pressure	Boiling point of water
760 mmHg	$100^{\circ} C$
23.7 mmHg	$25^{\circ} C$

5.2.2EvaporationOften examined

The molecules in a liquid are in motion and move within the liquid. The energy of the molecules is not equally distributed. Molecules with high kinetic energy move faster, while molecules with low kinetic energy move slowly. When the high-speed molecules reach the surface of the liquid, they may escape the attraction of their side or neighbouring molecules and leave the surface of the liquid.

Evaporation: The spontaneous change of a liquid into its vapour at the surface of the liquid at a given temperature.

Evaporation takes place at all temperatures. Evaporation causes cooling. When the high-energy molecules leave the liquid during evaporation, low-energy molecules are left behind, so the average kinetic energy of the molecules of the liquid decreases and the temperature of the liquid falls. Heat then moves from the surroundings to the liquid, and the temperature of the surroundings falls.

Evaporation: the fast, high-energy molecules escape from the surface and the slower molecules are left behind, so the liquid cools.

Factors that control evaporation

Surface area: when the surface area of a liquid increases, evaporation increases, because evaporation takes place from the liquid surface and more molecules can escape.
Intermolecular forces: when the forces are strong, molecules cannot leave the liquid surface easily and evaporation decreases. The rate of evaporation is slow when the intermolecular forces are strong and faster when they are weak.
Temperature: at a high temperature, molecules have greater energy, so the rate of evaporation increases.

Gasoline and ethyl alcohol have molecules with weaker intermolecular forces (London forces of attraction), so their evaporation is much faster than that of water, whose molecules have strong intermolecular forces (hydrogen bonding).

Rate of evaporation and intermolecular forces (as in the book)
Liquid	Main intermolecular force	Rate of evaporation
Gasoline	London forces (weak)	Much faster
Ethyl alcohol	London forces (weak), as the book states	Much faster
Water	Hydrogen bonding (strong)	Slower

5.2.3Vapour pressureOften examined

The liquid molecules having high kinetic energy leave the liquid surface and escape out of the liquid; this process is called evaporation. When the surface of a liquid is open, the liquid molecules mix with the air above the liquid, leave the surface and go away.

If the surface of the liquid is covered and the system is closed, the molecules that evaporate start gathering above the surface of the liquid. These evaporated molecules collide with the surface of the liquid and with the walls of the container. They are captured by the surface of the liquid. This coming back of vapours to the liquid, captured by the liquid surface, is called condensation.

Both processes, evaporation and condensation, continue until a time is reached when the rate of evaporation is equal to the rate of condensation. This state is called the state of dynamic equilibrium.

Liquid Vapours

Dynamic equilibrium in a closed container: the rate of evaporation equals the rate of condensation, so the pressure of the vapour stays constant.

Vapour pressure: The pressure exerted by the vapour of a liquid in equilibrium with the liquid at a given temperature.

At a constant temperature, the number of molecules leaving the surface of the liquid is equal to the number of molecules coming back into the liquid. At one moment a molecule may be in the liquid state, and in the next moment it may be in the vapour state, during evaporation.

The vapour pressure is independent of the amount of liquid or the volume of the container. The surface area has no effect on vapour pressure: a larger surface area has a larger target for evaporation and also a larger target for the return of the molecules.

Factors affecting vapour pressure

The vapour pressure of a liquid depends on the nature of the liquid (the intermolecular forces between its molecules and the size of the molecules, whether small or large) and on temperature.

Temperature: increasing temperature increases the kinetic energy of the molecules. At a high temperature, the molecules of a liquid have high kinetic energy and a greater capacity to break the intermolecular forces and escape from the liquid surface. The rate of evaporation increases, and this causes the vapour pressure to increase. Water shows an increase of vapour pressure from 4.579 torr at $0^{\circ} C$ to 9.209 torr at $10^{\circ} C$ , 527.8 torr at $90^{\circ} C$ and 760 torr at $100^{\circ} C$ .
Intermolecular forces: at a particular temperature, the stronger the intermolecular forces, the lower the vapour pressure, and the weaker the intermolecular forces, the higher the vapour pressure. At the same temperature different liquids have different vapour pressures. At $20^{\circ} C$ glycerol has 0.00016 torr and isopentane has 580 torr.

Table 5.1: Vapour pressure of some liquids at $20^{\circ} C$
Name of compound	Vapour pressure at $20^{\circ} C$ (torr)
Glycerol	0.00016
Mercury	0.012
Water	43.9
Carbon tetrachloride	87
Chloroform	170
Ethyl ether	442.2
Isopentane	580

Vapour pressure of water

100

110

120

130

200

400

600

800

1000

1200

1400

1600

T (^{\circ} C)

p (torr)

6 9^{\circ} C, 323

9 0^{\circ} C, 527.8

9 8^{\circ} C, 700

10 0^{\circ} C, 760

12 0^{\circ} C, 1489

The vapour pressure of water at different temperatures, plotted from the values given in the book. Water boils when its vapour pressure equals the external pressure: at 760 torr (the dashed line) this is

100^{\circ} C

Worked example: why water has a lower vapour pressure than ether and carbon tetrachloride

Water has a lower vapour pressure than

CCl X_{4}

and ether at the same temperature. Explain why, sketch the vapour pressure curves of the three liquids on one graph, and say what the graph shows about their boiling points.

Compare the evidence in the book at $20^{\circ} C$ : ether $442.2$ torr, $CCl X_{4}$ $87$ torr, and water lower still (the book prints $43.9$ torr).
Name the intermolecular forces. Water molecules are held by strong hydrogen bonds. $CCl X_{4}$ has only weak London forces and ether is held by weaker forces than water.
Link force to vapour pressure: the stronger the intermolecular forces, the fewer molecules have enough kinetic energy to escape the surface, so fewer enter the vapour and the vapour pressure is lower.
Sketch: draw three rising curves of the same shape. Order them water lowest, $CCl X_{4}$ in the middle, ether highest at any one temperature.
Read the boiling points. A liquid boils where its curve meets the dashed line at the atmospheric pressure (760 torr). The ether curve meets it first (lowest boiling point) and the water curve last ( $100^{\circ} C$ ).

Vapour pressure against temperature (sketch)

100

120

200

400

600

800

1000

1200

1400

1600

T (^{\circ} C)

p (torr)

760 torr

10 0^{\circ} C

water

ether

CC l_{4}

A sketch, not to scale. At every temperature ether has the highest vapour pressure and water the lowest, so ether reaches 760 torr at the lowest temperature and water at the highest (

100^{\circ} C

Answer in one line: strong hydrogen bonding in water makes evaporation slow, so water has the lowest vapour pressure and needs the highest temperature to reach the external pressure, which is why it has the highest boiling point of the three.

Why the pressure outside matters

Water boils at $100^{\circ} C$ at 760 torr pressure at sea level and boils at $120^{\circ} C$ at 1489 torr pressure. It boils at $25^{\circ} C$ at 23.7 torr pressure. At Murree hills, where the external pressure is 700 torr, water boils at $98^{\circ} C$ , while water boils at $69^{\circ} C$ at the top of Mount Everest, where the external pressure is 323 torr.

A liquid reaches its boiling point when its vapour pressure becomes equal to the atmospheric pressure at sea level, 760 torr. When different liquids are heated, they show different vapour pressures because of the difference in their intermolecular forces, and so their boiling points are different.

5.2.4Boiling pointOften examined

By heating a liquid, its vapour pressure increases. When heating is continued, a stage is reached when the vapour pressure of the liquid becomes equal to the external pressure or atmospheric pressure. The temperature at which the vapour pressure of a liquid becomes equal to the external pressure or atmospheric pressure is called the boiling point or boiling temperature.

At this temperature boiling starts. The kinetic energy of the molecules increases by heating, and the temperature of the liquid also increases. When the kinetic energy of the liquid molecules becomes maximum, the further heat supplied does not increase the temperature (the temperature remains constant at the boiling point), because the further heating breaks the intermolecular forces and the liquid is converted into vapour.

Temperature against heat supplied (a sketch, not to scale): the temperature stays constant during melting and during boiling, because the heat supplied is used to break the intermolecular forces.

Molar heat of vaporization: The amount of heat required to convert one mole of a liquid into vapours at its boiling point.

For water the molar heat of vaporization is 40.6 kJ per mole.

5.2.5Boiling point and atmospheric pressureOften examined

The boiling point of a liquid depends on its intermolecular forces: the greater the intermolecular forces, the higher the boiling temperature, and vice versa. The boiling point also depends on the external or atmospheric pressure, because it is the temperature at which the vapour pressure of the liquid becomes equal to the external pressure.

Effect of external pressure on boiling point
When the external pressure is	The liquid needs	So the boiling point is
High	A larger amount of heat to produce a vapour pressure equal to it	High
Low	A smaller amount of heat to boil	Low

Boiling point of water at different external pressures (values from section 5.2.3)
Place or process	External pressure	Boiling point of water
Top of Mount Everest	323 torr	$69^{\circ} C$
Murree hills	700 torr	$98^{\circ} C$
Sea level	760 torr	$100^{\circ} C$
Pressure cooker	1489 torr	$120^{\circ} C$

5.2.6ViscosityOften examined

A common observation is that water can be poured from one container to another very quickly compared with honey and glycerine. The resistance to flow of a liquid is called viscosity. The higher the viscosity, the slower the liquid flows. Viscosity measures how easily molecules slide past each other.

Viscosity: The resistance to flow of a liquid.

To understand viscosity, think of a fluid flowing in a pipe as consisting of a series of concentric circular layers. Resistance is caused by internal friction between the layers of molecules. The layer next to the walls has the lowest speed. Each layer attracts the other and thus causes resistance to flow.

Layers of a liquid flowing in a pipe: the layer next to the wall is slowest and each layer drags on the next, which is the resistance to flow (viscosity).

Units

The SI unit of viscosity is the pascal second ( $Pa s$ ). The non-SI unit of viscosity is the poise.

1 poise = 0.1 kg m^{- 1} s^{- 1} = 1 g c m^{- 1} s^{- 1}

1 Pa s = 1 kg m^{- 1} s^{- 1} = 10 poise

Factors affecting viscosity

Size and shape of the molecule: liquids such as water, acetone, benzene and methanol, whose molecules are small and compact, have a low viscosity. In liquids with large and irregular molecules, such as honey and glycerine, the molecules tend to tangle with each other. This hinders the flow of the liquid and leads to high viscosity.
Intermolecular forces: the stronger the intermolecular force, the higher the viscosity. Liquids whose molecules form hydrogen bonds are more viscous than others without hydrogen bonds. For example, water is more viscous than methanol mainly because of extensive hydrogen bonding.
Temperature: molecules move faster when the temperature rises, because an increase in temperature decreases the intermolecular forces. This dependence is quite visible for very viscous liquids such as honey and syrup. These liquids are easier to pour when hot than cold.

5.2.7Surface tensionOften examined

Surface tension is the property of the surface of a liquid to behave as if a membrane were stretched over it. All the molecules below the surface of a liquid are surrounded by other molecules in all directions, so the force exerted by such molecules is balanced in all directions. At the surface of a liquid there are molecules beside and below, but none above. This results in an unbalanced force that pulls the surface molecules inward. The molecules on the surface therefore feel an inward pull, creating surface tension.

Surface tension: A force in dynes acting at right angles to a unit length of the liquid surface (the book's definition). It can also be defined as the amount of energy required to expand the surface of a liquid per unit area.

The molecule inside a liquid is pulled equally in all directions, but a molecule at the surface has no neighbours above it and feels a net pull inward.

For a molecule to reach the surface, it must overcome the downward force of attraction. This means that work must be done to bring it to the surface. Therefore increasing the surface area of a liquid requires input energy.

Molecules on the surface of a liquid are less stable than those inside it, so a liquid is stable if it has the fewest molecules on its surface. This happens when the surface of the liquid is the smallest. A sphere has less surface area per unit volume than any other shape, so small liquid droplets are usually spherical.

Factors affecting surface tension

Intermolecular forces: the surface tension of a liquid depends directly on the strength of the intermolecular forces. The stronger the forces between the liquid molecules, the higher the surface tension, and vice versa. For example, water has a higher surface tension than many other liquids, such as alcohols, ethers and benzene, because of the strong hydrogen bonding between water molecules.
Temperature: the surface tension of a liquid decreases with increasing temperature, because the increased kinetic energy of the molecules reduces the strength of the intermolecular forces. It is different for different liquids because of the different types of intermolecular forces.

Worked example: why the surface tension of water is about three times that of carbon tetrachloride

From the table, the surface tension of water is

7.275 \times 1 0^{- 2} N m^{- 1}

and of

CCl X_{4}

2.70 \times 1 0^{- 2} N m^{- 1}

. Compare them numerically and explain the difference. Also explain why a falling drop of water is spherical.

Compare the numbers.
$\frac{7.275}{2.70} = 2.7 \approx 3$
Name the forces. Water molecules attract each other by strong hydrogen bonds. $CCl X_{4}$ molecules attract each other only by weak London forces.
Relate force to surface tension. A surface molecule is pulled inward by its neighbours below. The stronger the attractions, the bigger this inward pull, so more work is needed to bring molecules to the surface and the surface tension is higher.
Falling drop. Surface molecules are less stable than those inside, so the liquid is most stable when it has the fewest molecules at the surface, which means the smallest surface area. A sphere has the least surface area for a given volume, so the drop is spherical.

Surface tension of some liquids (values from the book's table). Water, with hydrogen bonding, is far above the rest.

γ_{H X_{2} O} \approx 3 γ_{CCl X_{4}} because of hydrogen bonding

Units

The SI unit of surface tension is joule per square metre ( $J m^{- 2}$ ), or newton per metre ( $N m^{- 1}$ ).

5.2.8Hydrogen bonding and the properties of waterOften examined

Hydrogen bonding in water contributes to its unique properties, which are as follows.

Hydrogen bonds (dashed) between water molecules: the hydrogen of one molecule is attracted to the oxygen of the next. These strong intermolecular forces explain the properties of water.

i. High surface tension

Surface tension is affected by strong intermolecular attractive forces. Water has hydrogen bonding, which is a strong intermolecular force. This creates a high surface tension.

Surface tension of water compared with some other liquids
Solvent	Surface tension ( $γ \times 1 0^{- 2}$ , $N m^{- 1}$ )
Water	7.275
Methanol	2.26
Ethanol	2.28
Benzene	2.888
Hexane	1.84
$CCl X_{4}$	2.70

ii. Low vapour pressure

The vapour pressure of a liquid is affected by intermolecular attractive forces. Strong forces of attraction make evaporation slow. Because of hydrogen bonding in water, which is a strong intermolecular attractive force, evaporation decreases and so water has a low vapour pressure.

iii. High heat of vaporization

The heat of vaporization of water is high compared with other liquids because of the strong intermolecular forces of hydrogen bonding in water. More heat is required to overcome these attractive forces to vaporize water, and as a result water shows a high heat of vaporization.

iv. High boiling point

The high heat of vaporization of water results in a high boiling point. Because of strong hydrogen bonding, water has a low vapour pressure, so more heat is required to enhance evaporation and to produce a vapour pressure equal to the external pressure. As a result water has a high boiling point.

Hydrogen bonding explains four properties of water
Property of water	Reason
High surface tension	Hydrogen bonds are strong intermolecular forces
Low vapour pressure	Strong forces make evaporation slow
High heat of vaporization	More heat is needed to overcome the hydrogen bonds
High boiling point	Low vapour pressure means more heat is needed to make the vapour pressure equal to the external pressure

5.3Energetics of phase changesOften examined

Physical and chemical changes are accompanied by a change in energy in the form of heat. A physical energy change is a quantitative measure of the strength of the intermolecular forces. The change in energy at constant pressure is called the change in enthalpy, denoted by $Δ H$ . It is expressed in $kJ mo l^{- 1}$ .

When a substance undergoes a phase change (a change of state), its temperature remains constant even when heat is added.

5.3.1Molar heat of fusion and molar heat of vaporizationOften examined

Molar heat of fusion ( $Δ H_{f}$ ): The amount of heat required to convert one mole of a solid into its liquid state at its melting point.
Molar heat of vaporization ( $Δ H_{v}$ ): The amount of heat required to convert one mole of a liquid into its vapours at its boiling point.

For example, the molar heat of fusion of ice is $+ 6.02 kJ mo l^{- 1}$ .

H X_{2} O (s) H X_{2} O (l) Δ H_{f} = + 6.02 kJ mo l^{- 1}

For water, the molar heat of vaporization is $+ 40.7 kJ mo l^{- 1}$ .

H X_{2} O (l) H X_{2} O (g) Δ H_{v} = + 40.7 kJ mo l^{- 1}

Molar heat of fusion and molar heat of vaporization compared
	Molar heat of fusion	Molar heat of vaporization
Symbol	$Δ H_{f}$	$Δ H_{v}$
Change of state	Solid to liquid	Liquid to vapour
Temperature	At the melting point	At the boiling point
Amount of substance	One mole	One mole
Value for water	$+ 6.02 kJ mo l^{- 1}$	$+ 40.7 kJ mo l^{- 1}$

Worked example: heat needed to melt ice and to boil water

Use

Δ H_{f} = + 6.02 kJ mo l^{- 1}

and

Δ H_{v} = + 40.7 kJ mo l^{- 1}

for water. Find (a) the heat needed to melt

36 g

of ice at its melting point, (b) the heat needed to vaporize

36 g

of water at its boiling point, and (c) the mass of ice that

30.1 kJ

can melt.

Given

$m = 36 g$
$M_{H X_{2} O} = 18 g mo l^{- 1}$
$Δ H_{f}, Δ H_{v} = 6.02, 40.7 kJ mo l^{- 1}$

To find

$q_{melt}$
$q_{vap}$
mass melted by $30.1 kJ$

Moles of water.
$n = \frac{36}{18} = 2.0 mol$
(a) Heat to melt: moles times molar heat of fusion.
$q = 2.0 \times 6.02 = 12.04 kJ$
(b) Heat to vaporize.
$q = 2.0 \times 40.7 = 81.4 kJ$
Compare the two. Vaporization needs about 6.8 times more heat, because the molecules move from close contact to far apart, a much bigger change in distance and potential energy than in melting.
$\frac{81.4}{12.04} = 6.76$
(c) Moles melted by $30.1 kJ$ , then the mass.
$n = \frac{30.1}{6.02} = 5.0 mol, m = 5.0 \times 18 = 90 g$

q_{melt} = 12.04 kJ; q_{vap} = 81.4 kJ; m = 90 g

5.3.2Energy changes and intermolecular forcesOften examined

As a result of the melting of a solid, a small change in intermolecular distance and potential energy takes place in the atoms, molecules or ions. On the other hand, on evaporation of a liquid, the atoms, molecules or ions undergo large changes in their intermolecular distance and potential energy. Therefore, the heat of vaporization is much greater than the heat of fusion.

Energy levels of one mole of water: the step from liquid to gas (+40.7\,\mathrm{kJ\,mol^{-1}}) is much bigger than the step from solid to liquid (+6.02\,\mathrm{kJ\,mol^{-1}}). The heights are drawn in proportion.

Why heat of vaporization is greater than heat of fusion

Melting: the particles stay close together, so the intermolecular distance changes only slightly.
Vaporization: the particles move from close contact to far apart, so the intermolecular distance changes a lot and so does the potential energy.
More energy is needed to overcome the attractions over a larger distance, so \Delta H_v > \Delta H_f.
\Delta H_v = 40.7 > \Delta H_f = 6.02\ \ (\mathrm{kJ\,mol^{-1}}\ \text{for water})

5.3.3Liquid crystalsOften examined

Pure liquids have sharp melting points, and at the melting point the temperature remains constant. Generally, when solids are heated, they are converted into liquids. In 1888 Friedrich Reinitzer (spelt 'Reimitaz' in the book) discovered an unusual property while studying cholesteryl benzoate, an organic compound. When this compound is heated it turns into a milky liquid at 145\,^\circ\mathrm{C} and then into a clear liquid at 179\,^\circ\mathrm{C}. When it is cooled, the reverse occurs. The turbid (milky) liquid state is called the liquid crystal state.

There are many crystalline solids which melt to a turbid liquid and then to a clear liquid phase. The turbid liquid phase can flow and has the properties of liquids such as viscosity and surface tension. The molecules of these turbid liquids have some degree of order and resemble crystals in properties like optical properties.

Liquid crystals: The liquid crystalline state between the melting temperature and the clearing temperature.

\text{Solid crystal} \rightleftharpoons \text{liquid crystal} \rightleftharpoons \text{clear liquid}

Liquid crystal state: the rod-like molecules are aligned in one direction, as in a crystal, but can still flow like a liquid. In a solid crystal they are fixed in order, and in a clear liquid they point in random directions.

Substances made of long rod-like molecules form these liquid crystals. The molecules are oriented in random directions in the normal liquid phase, and in the liquid crystal form they are in order. On the basis of the order of the molecules, liquid crystals are of three types: nematic, smectic and cholesteric (spelt 'choloestrinic' in the book).

The liquid crystal is intermediate between a solid crystal and a clear liquid, so its properties are also between the properties of crystals and those of isotropic liquids. Liquid crystals have crystal properties such as optical properties, and they have the fluidity of liquids.

Nematic
Smectic
Cholesteric

Worked example: states of cholesteryl benzoate on heating and cooling

Cholesteryl benzoate turns milky at 145\,^\circ\mathrm{C} and clear at 179\,^\circ\mathrm{C}. State what it is at 130\,^\circ\mathrm{C}, 160\,^\circ\mathrm{C} and 185\,^\circ\mathrm{C}, what happens as a clear sample cools from 185\,^\circ\mathrm{C} to 160\,^\circ\mathrm{C}, and which of the three states has both order and fluidity.

Place each temperature on the scale. Below 145\,^\circ\mathrm{C} the substance is a solid crystal. Between 145 and 179\,^\circ\mathrm{C} it is a milky liquid crystal. Above 179\,^\circ\mathrm{C} it is a clear liquid.
\text{130 °C: solid},\quad \text{160 °C: liquid crystal},\quad \text{185 °C: clear liquid}
Cooling reverses the change. The clear liquid turns milky at 179\,^\circ\mathrm{C}, so at 160\,^\circ\mathrm{C} it is the milky liquid crystal again.
Order and fluidity. The solid has order but no flow. The clear liquid flows but its molecules point in random directions. Only the liquid crystal has the rod-like molecules lined up in one direction and also flows like a liquid.

The liquid crystal state of cholesteryl benzoate lies between 145 and 179 °C: the rod-shaped molecules keep their common direction but are free to flow.

\text{solid},\ \text{liquid crystal},\ \text{clear liquid};\ \text{liquid crystal has order and fluidity}

5.3.4Uses of liquid crystals

Liquid crystals are used in the display of electrical devices like digital watches, laptop computers, calculators and so on.
Because of the electromagnetic optic characteristic and photoelectric properties of nematic and cholesteric types of liquid crystals, different laboratories, universities and industries began to focus on their applications.
Cholesteric liquid crystals applied to the skin are used to locate veins, arteries, infections, tumours and so on, which are warmer than other tissue.
Nematic liquid crystals are used in magnetic resonance. Molecules dissolved in nematic liquid crystal solvents show a resolved NMR spectrum, and intermolecular dipole-dipole fine structures are obtained.
Liquid crystals are used in chromatographic separation to study molecular arrangements and kinetics, and as an anisotropic fluid for visible, UV and IR spectroscopy.
Liquid crystals are used in oscillography systems and television displays, and liquid crystal screens have also been developed.
Liquid crystal polymers are used in industries. Polyester liquid crystals were developed for fire resistance and are used as a coating for multifibre optical cables and so on.

Summary of uses of liquid crystals
Use	Type or field
Displays of watches, laptops, calculators	Electrical devices
Locating veins, arteries, infections and tumours (warmer than other tissue)	Cholesteric, applied to skin
Resolved NMR spectrum of dissolved molecules	Nematic, magnetic resonance
Chromatography and spectroscopy (visible, UV, IR)	Anisotropic fluid
Oscillography, television and liquid crystal screens	Display systems
Fire-resistant polyester coating for optical cables	Liquid crystal polymers

5.3.5Differentiation of liquid crystals from pure liquids and crystalline solidsOften examined

In liquid crystals, the molecules are arranged as in a crystal but with certain degrees of freedom, so that they can flow like liquids. They have both fluidity and some long-range molecular order. They have the ability to change their molecular structure in response to external factors such as temperature, electric field or pressure.

Pure liquids have a random arrangement of molecules with no long-range order. They can flow freely and take the form of a vessel. Pure liquids show no significant changes in molecular arrangement in response to external stimuli, other than changes in temperature or pressure.

In solids, the molecules or atoms have a well-ordered and repeating three-dimensional arrangement. A solid maintains a certain shape and volume. Some crystalline solids can change phase under the influence of temperature.

Liquid crystals compared with pure liquids and crystalline solids
Property	Liquid crystal	Pure liquid	Crystalline solid
Arrangement of molecules	As in a crystal but with some degrees of freedom	Random, no long-range order	Well-ordered, repeating three-dimensional arrangement
Flow	Can flow like a liquid (fluidity)	Flows freely and takes the form of a vessel	Keeps a certain shape and volume
Order	Some long-range molecular order	No long-range order	Long-range order
Response to external factors	Changes molecular structure with temperature, electric field or pressure	No significant change except with temperature or pressure	Some crystalline solids can change phase with temperature

5.3.6Glaciers and ice caps

The study of glaciers, particularly polar ice sheets, relies heavily on understanding the significant role that heat of fusion plays in their behaviour.
The melting and freezing of ice are closely tied to the heat of fusion. Hence it holds great significance in anticipating the response of glaciers and ice sheets to fluctuations in temperature.
The increasing global temperatures caused by global warming make it necessary to consider the heat of fusion as a crucial element in evaluating the susceptibility of polar ice sheets. By using this information, scientists are able to make predictions about the extent of ice melting and its impact on sea levels as well as other environmental transformations.
The quantity of water released during the melting of ice is directly proportional to the heat of fusion. Through the analysis of the heat absorbed during the melting process of ice cores, researchers are able to obtain valuable information about previous climatic conditions.

5.4General properties of gasesOften examined

Gases do not have a definite shape and volume. Gas molecules are free to move. The molecules of a gas are in random motion. The volume of a gas is the volume of its container, and a gas takes the shape of its container.
The intermolecular space in a gas is large, so gases have low densities. Gases bubble and rise up through a liquid because their densities are lower than those of liquids and solids.
The intermolecular spaces in gases are large enough to accommodate the molecules of other gases, so gases show the properties of diffusion and effusion.
When gases are heated, the velocity of the molecules increases and the volume of the gas increases. Similarly, if volume is available, gases can expand and occupy the available volume.
Gas molecules are in random motion and collide with each other and with the walls of the container. Gases exert pressure because of the number of collisions.
According to the kinetic molecular theory, the space between gas molecules is more than in solids and liquids, and the molecules of a gas are free to move. The kinetic energy of the molecules of a gas depends on temperature. Gases are free to flow and fill the available volume or container.

1. Compressibility

Gases are highly compressible because the particles in gases are far apart and large spaces exist between them. So gases can be compressed into a small volume by applying pressure. Due to this property, a large volume of a gas can be placed in a small cylinder by compression. For example, oxygen gas in cylinders is supplied to hospitals, compressed natural gas (CNG) is used in vehicles, and liquefied petroleum gas (LPG) is used for cooking in kitchens.

The state of a substance is determined by the balance between the kinetic energy of the individual particles (atoms or molecules) and the intermolecular forces between them. The kinetic energy keeps the particles apart and moving around and determines the temperature of the substance, while intermolecular forces draw the molecules together. In gases, the average kinetic energy of the molecules is larger than the average energy of attraction between the particles or molecules of the gas.

2. Expandability

Gases expand to fill their container. The expandability of gases is spontaneous, and gases occupy the available volume. The gas molecules are in random motion, having kinetic energy, and there is empty space between the gas molecules. The air contained in a balloon fills up the whole balloon. If the gas in the balloon is let into the room, it will spread in the room. In the balloon the gas has a high pressure, and in the room the gas has a low pressure. Compressed gas can also expand if volume is available. Compression and expansion of gases are also used in the liquefaction of gases.

3. Pressure exerted by gases

The molecules of a gas are in constant random motion and collide against the walls of the container. Because of the collisions with the walls of the container, the walls experience a continuous force. Because of this force, gases exert pressure. Pressure is force per unit area.

P = \dfrac{F}{A}

Gas pressure: the molecules move at random and strike the walls of the container. The continuous force of the collisions on the wall, divided by the area of the wall, is the pressure.

Three properties of gases
Property	Reason (kinetic molecular theory)	Example given in the book
Compressibility	Particles are far apart with large spaces between them	Oxygen cylinders, CNG, LPG
Expandability	Random motion with empty space between molecules	Air fills a balloon; gas spreads in a room
Pressure	Collisions of molecules with the walls cause a continuous force	Gas in a balloon is at a higher pressure than in the room

5.5General properties of solidsOften examined

A solid has a definite mass, volume and shape. It has a compact arrangement of constituent particles. The distance between the molecules is short and the forces of attraction between the particles are very strong. For this reason solids are non-compressible and cannot diffuse into each other. Solid particles have only vibrational motion.
A solid is a collection of atoms or molecules held together in a definite shape under constant conditions. Substances which are rigid and hard, cannot flow and have a definite volume under constant conditions are solids.
They have a definite size under constant conditions because of the strong forces of attraction between their atoms or molecules. The constituent particles (atoms, molecules or ions) cannot move randomly. The particles only show vibrational motion.
Solid particles are held together by strong cohesive forces.
The constituent particles (atoms, molecules or ions) are closely packed with each other. There is a well-ordered arrangement in solids.

5.5.1Kinetic molecular interpretation of solidsOften examined

Solids show the maximum attractive forces among their particles (atoms, molecules or ions) because of the very close packing of the arrangement.
The particles of solids are closely packed, so they occupy the minimum volume and show a high density.
Solid particles do not show translational motion, therefore there is no collision between the molecules.
The particles of a solid can only vibrate and possess only vibrational kinetic energy.
Crystalline solids have a definite, distinctive geometric shape because of the orderly arrangement of the particles (atoms, molecules or ions) in three dimensions.

5.5.2Properties of solids based on kinetic molecular theoryOften examined

1. Diffusion

Diffusion depends on the velocity of molecules. The movement of molecules in solids is vibration only, so solids show negligible diffusion.

2. Motion of molecules

The particles of solids (atoms, molecules or ions) are closely packed with each other and show only vibrational motion. The particles can only vibrate about their mean position.

3. Intermolecular forces

Intermolecular forces are maximum in solids. The molecules of solids are held together in fixed positions by strong forces of attraction and can only vibrate about their fixed positions.

4. Compression

The particles of solids are closely packed and have no space between them. There is a minimum effect, or no effect, of pressure on solids.

5. Expansion

Solids expand when heated (the volume increases). This is due to the fact that an increase in temperature decreases the intermolecular attractive forces, and space is created between the molecules, so the volume increases.

Properties of solids explained by the kinetic molecular theory
Property	Solids	Reason
Diffusion	Negligible	Only vibration, no movement from place to place
Motion of particles	Vibration about the mean position	Closely packed particles
Intermolecular forces	Maximum	Particles held in fixed positions
Compression	Minimum or no effect of pressure	No space between the particles
Expansion	Volume increases when heated	Intermolecular forces decrease and space is created

The three states compared (collected from sections 5.1, 5.4 and 5.5)
Property	Gas	Liquid	Solid
Particles	Far apart, free to move	Close, slide past each other	Closely packed, fixed positions
Intermolecular forces	Weakest	Strong, but not enough to fix molecules	Maximum
Motion	Random, with many collisions	Limited by close packing	Vibration only
Compression	Highly compressible	Negligible	Minimum or none
Diffusion	Shown (diffusion and effusion)	Slow	Negligible

5.5.3Types of solidsOften examined

Solids are classified on the basis of the regular arrangement of their particles (atoms, ions or molecules). There are two types of solids.

1. Crystalline solids

The solids which have a definite regular and three-dimensional geometric shape are called crystalline solids, for example \ce{NaCl}, diamond, ice, glucose and \ce{ZnSO4}. In these solids the arrangement of atoms, ions or molecules is in a definite three-dimensional pattern. This recurring geometrical pattern of structure extends in three dimensions.

2. Amorphous solids

Amorphous solids are those whose constituent particles do not possess a regular, orderly arrangement. These solids have no definite geometric shape, for example glass, plastic and rubber. They have the properties of the solid state but do not have the ordered crystalline state.

Crystalline solids can be converted into amorphous solids by melting and cooling the molten mass rapidly, so the particles do not find time to arrange themselves. Amorphous solids can have some small regions of orderly arrangement. These crystalline parts of amorphous solids are known as crystallites.

Amorphous solids have no sharp melting points. The particles of glass soften over a high temperature range and can be moulded.

In a crystalline solid the particles form a regular pattern that repeats in three dimensions. An amorphous solid has no regular pattern, apart from small ordered regions called crystallites.

Crystalline and amorphous solids compared
Property	Crystalline solids	Amorphous solids
Arrangement of particles	Definite, regular, repeating three-dimensional pattern	No regular, orderly arrangement (small ordered regions called crystallites may exist)
Shape	Definite, distinctive geometric shape	No definite geometric shape
Melting point	Sharp melting point	No sharp melting point; soften over a range of temperature
Examples	\ce{NaCl}, diamond, ice, glucose, \ce{ZnSO4}	Glass, plastic, rubber

5.5.4Properties of crystalline solidsOften examined

1. Geometrical shape

Crystals have a geometric shape. The particles of solids (ions, atoms or molecules) are arranged in three-dimensional space. The internal angles of a crystal (where surfaces intersect) are always the same, no matter in what form the crystal is grown.

2. Melting point

Crystalline solids have sharp melting points. When a solid is heated, its particles start vibrating at a higher temperature and transfer their kinetic energy throughout the solid. At the melting point, their vibration energies become so high that they start leaving their fixed positions simultaneously and the solid turns into a liquid.

3. Cleavage planes

Crystalline solids break up along certain planes called cleavage planes. When pressure is applied to a crystal along a cleavage plane, it changes into small crystals of the same size and shape as the original crystal. The angles between the cleavage planes vary from one crystalline solid to another.

4. Habit of a crystal

Crystals form slowly when a saturated solution is cooled. Crystals form by growing in different directions. If the growth conditions remain the same, crystals of the same shape will form. The shape of a crystal in which it usually grows is called the habit of a crystal. For example, cubic crystals of sodium chloride are obtained from its aqueous solution. But if the conditions are changed, the shape of the crystals also changes. For example, if 10% urea is present in the aqueous solution of sodium chloride, octahedral crystals of sodium chloride are obtained.

Four properties of crystalline solids
Property	Meaning
Geometrical shape	Definite shape; the angles between the faces of a crystal are always the same
Sharp melting point	All the particles leave their fixed positions at once at the melting point
Cleavage planes	Definite planes along which a crystal breaks into smaller crystals of the same shape
Habit	The shape in which a crystal usually grows; it changes if the growth conditions change

Formula sheet

Viscosity unit conversion

1\ \mathrm{Pa\,s}=1\ \mathrm{kg\,m^{-1}\,s^{-1}}=10\ \text{poise}

The pascal second is the SI unit of viscosity and the poise is the non-SI unit; 1\ \text{poise}=0.1\ \mathrm{kg\,m^{-1}\,s^{-1}}.

Units: \mathrm{Pa\,s} or poise

Molar heat of fusion of water

\ce{H2O(s) -> H2O(l)}\qquad \Delta H_f=+6.02\,\mathrm{kJ\,mol^{-1}}

\Delta H_f is the heat needed to convert one mole of a solid into its liquid at the melting point.

Unit: \mathrm{kJ\,mol^{-1}}

Molar heat of vaporization of water

\ce{H2O(l) -> H2O(g)}\qquad \Delta H_v=+40.7\,\mathrm{kJ\,mol^{-1}}

\Delta H_v is the heat needed to convert one mole of a liquid into its vapour at the boiling point.

Unit: \mathrm{kJ\,mol^{-1}}

Dynamic equilibrium of a liquid and its vapour

\ce{Liquid <=> Vapours}

At equilibrium the rate of evaporation equals the rate of condensation; the pressure of the vapour is then the vapour pressure.

Units: pressure in torr or mmHg

Pressure of a gas

P=\dfrac{F}{A}

F is the force of the colliding molecules on the wall and A is the area of the wall.

Unit: \mathrm{N\,m^{-2}} (pascal)

Key terms

Phase: One of the three forms in which matter exists in a system: gas, liquid or solid.
Intermolecular forces: The forces of attraction between molecules of a substance in any phase.
Dipole-dipole forces: The forces of attraction between the positive end of one molecule and the negative end of another molecule.
Debye forces (dipole-induced dipole forces): The attractive forces between dipoles and induced dipoles.
London forces: The momentary forces of attraction created between an instantaneous dipole and an induced dipole (instantaneous dipole-induced dipole forces).
Evaporation: The spontaneous change of a liquid into its vapour at the surface of the liquid at a given temperature.
Condensation: The coming back of vapours to the liquid, when they are captured by the liquid surface.
Dynamic equilibrium: The state in which the rate of evaporation is equal to the rate of condensation.
Vapour pressure: The pressure exerted by the vapour of a liquid in equilibrium with the liquid at a given temperature.
Boiling point: The temperature at which the vapour pressure of a liquid becomes equal to the external pressure or atmospheric pressure.
Viscosity: The resistance to flow of a liquid.
Surface tension: The property of the surface of a liquid to behave as if a membrane were stretched over it; the book defines it as a dyne force acting at right angles to a unit length of the liquid surface.
Molar heat of fusion (\Delta H_f): The amount of heat required to convert one mole of a solid into its liquid state at its melting point.
Molar heat of vaporization (\Delta H_v): The amount of heat required to convert one mole of a liquid into its vapours at its boiling point.
Liquid crystals: The liquid crystalline state between the melting temperature and the clearing temperature.
Crystalline solids: Solids with a definite, regular, three-dimensional geometric shape because of the orderly arrangement of their particles, for example \ce{NaCl}, diamond and ice.
Amorphous solids: Solids whose particles have no regular, orderly arrangement and which have no definite geometric shape, for example glass, plastic and rubber.
Crystallites: The small crystalline (orderly) regions that an amorphous solid can contain.
Cleavage planes: The definite planes along which a crystalline solid breaks.
Habit of a crystal: The shape of a crystal in which it usually grows.

Multiple choice questionsp. 109–111

Van der Waals' forces are effective:

AAt long distance
BBoth at long as well as short distance
COnly at short distanceAnswer
DIndependent of distance

Why

Van der Waals' forces are short range. They act only when molecules are very close together, and their strength falls off steeply as the distance between molecules increases. So they are effective only at short distance, option (C).

Note. The book prints no key; the answer was worked out. The book does not define van der Waals forces, so the answer rests on the general short-range behaviour of the dipole and London forces.

Which one of the following forces are also called London forces?

AIon-dipole forces
BDipole-induced dipole forces
CDipole-dipole forces
DDispersion forcesAnswer

Why

London forces are the instantaneous dipole-induced dipole forces, which are also called dispersion forces, option (D). Dipole-induced dipole forces (B) are the separate Debye forces, so (B) is the tempting wrong answer. Ion-dipole (A) and dipole-dipole (C) forces are different types again.

iii

Which of the following two halogens are gases at room temperature?

AFluorine and iodine
BChlorine and bromine
CFluorine and chlorineAnswer
DIodine and bromine

Why

Fluorine (\ce{F2}) and chlorine (\ce{Cl2}) are gases at room temperature, option (C). Bromine is a volatile liquid and iodine is a solid, so every pair that contains either of them (A, B and D) is wrong.

The intermolecular forces are of:

ATwo types
BThree typesAnswer
CFour types
DFive types

Why

The book's key points name three intermolecular forces: dipole-dipole forces, dipole-induced dipole (Debye) forces and instantaneous dipole-induced dipole (London) forces, so option (B). Hydrogen bonding is treated in the book as a strong dipole-dipole attraction, not as a separate type. Options (A), (C) and (D) give the wrong count.

Note. The book prints no key. Some texts count hydrogen bonding as a fourth type, which would give (C). The answer follows the book's three named forces.

Thermostat is an instrument which:

AIncreases the temperature
BDecreases the temperature
CMaintains the temperatureAnswer
DFluctuates the temperature

Why

A thermostat senses temperature and switches heating or cooling on and off so that the temperature stays at a set value, option (C). Options (A) and (B) describe a heater and a cooler on their own, and (D) is the opposite of what a thermostat does.

The scientist who discussed the phenomenon of viscosity is:

APoiseuille
BNewtonAnswer
CFritz
DVan der Waals

Why

Isaac Newton is the scientist credited with defining viscosity and its law (the shearing stress is proportional to the velocity gradient between layers), option (B). Poiseuille (A) also studied viscous flow, but Newton is the name most closely linked with the phenomenon. Options (C) and (D) are not associated with viscosity.

Note. The book prints no key and does not name a scientist. Poiseuille (A) is a defensible alternative, since he measured viscous flow in tubes. The answer chosen is Newton because he is the scientist who first described viscosity. The printed spellings 'Poisuelle' and 'Vander Wall' were corrected in the options.

vii

The distillation under reduced pressure is called:

AFractional distillation
BVacuum distillationAnswer
CSteam distillation
DPressure distillation

Why

Distillation under reduced pressure is called vacuum distillation, option (B). Lowering the pressure lowers the boiling point, so heat-sensitive liquids such as glycerine can be distilled without decomposing. Fractional (A) and steam (C) distillation are different processes, and pressure distillation (D) is not a named technique.

viii

The unit of surface tension is:

ANewton per metreAnswer
BNewton per square metre
C760\,\mathrm{mmHg}
DNewton square per metre

Why

Surface tension is a force per unit length, so its SI unit is the newton per metre, \mathrm{N\,m^{-1}}, which is equal to \mathrm{J\,m^{-2}} (option A). Newton per square metre (B and D) is a unit of pressure, and \mathrm{mmHg} (C) is also a pressure unit.

The intermediate phase lying between the solid phase and the normal liquid phase is called:

ACrystalline solid
BLiquid crystalsAnswer
CMesogens
DCrystal lattice

Why

Liquid crystals are the intermediate state between the crystalline solid and the normal liquid, as the book's key points define them, option (B). A crystalline solid (A) is one of the two outer phases. Mesogens (C) are the molecules that form liquid crystals, and a crystal lattice (D) is an arrangement of particles, not a phase.

In which of the following are the particles the most disordered?

AWater at 100\,^\circ\mathrm{C}
BSteam at 100\,^\circ\mathrm{C}Answer
CImpure water at 102\,^\circ\mathrm{C}
DWater at 10\,^\circ\mathrm{C}

Why

Steam is a gas. Its particles are the most widely spaced and move most freely, so they are the most disordered, option (B). Water at 100 ^\circC (A) and impure water at 102 ^\circC (C) are liquids, and water at 10 ^\circC (D) is a liquid with the lowest kinetic energy of the four.

Which of these statements best supports the idea that matter is made up of particles?

ALiquids always fill the space available to them
BLiquids are easily compressible
C1\,\mathrm{cm^3} of water produces nearly 1700\,\mathrm{cm^3} of steam
DIf a bottle of perfume is opened, the smell spreads quicklyAnswer

Why

The spreading of a perfume smell through a room shows that tiny particles of the perfume move and mix with the air. This supports the particle idea most directly, option (D). Statement (C) is also consistent with particles, because steam has large spaces between particles, but it mainly shows the expansion of a gas. Statement (A) describes a liquid without referring to particles, and (B) is untrue because liquids are almost incompressible.

Note. The book prints no key. Statement (C) is a reasonable alternative, but (D) is the more direct evidence for moving particles.

xii

Which of these processes involve a weakening of the attraction between particles?

ACondensation
BFreezing
CCrystallization
DEvaporationAnswer

Why

Evaporation involves particles gaining kinetic energy so that they can break free of the attraction of their neighbours, option (D). Condensation (A), freezing (B) and crystallization (C) all bring particles closer together, which strengthens the attraction between them.

xiii

A liquid is thought to be pure ethanoic acid (acetic acid). Which of the following is the best way to test its purity?

AMeasure its B.PAnswer
BReact it with ethanol
CBurn it completely in oxygen
DDehydrate it with concentrated \ce{H2SO4}

Why

A pure liquid has a sharp, fixed boiling point at a given pressure, whereas impurities change the boiling point and make it boil over a range. Measuring the boiling point is therefore the best test of purity, option (A). Reacting with ethanol (B), burning (C) or dehydrating with \ce{H2SO4} (D) would not show whether the liquid is pure.

Short questionsp. 111

i(a)3 marks

Give the general properties of liquids as to diffusion.

Solution

Diffusion in a liquid happens because molecules move from one place to another. Their movement is restricted, however, because the molecules are packed closely together, so each molecule travels only a short distance before it collides with its neighbours.

Diffusion is therefore much slower in liquids than in gases. An ink drop added to water spreads slowly through it because the spaces between the molecules are small and there are relatively few collisions between molecules in a densely packed liquid.

i(b)3 marks

Give the general properties of liquids as to compression.

Solution

Liquids cannot be significantly compressed. Their molecules are already in close contact, so there is almost no empty space for them to move into when pressure is applied.

For example, raising the pressure on water from one to two atmospheres reduces its volume by only about 0.0045\,\%, which is negligible. The same change of pressure reduces the volume of a gas by up to 50\,\%, because gas molecules are far apart.

Note. The book's wording says the volume is reduced 'to 0.0045 per cent'. The sense of the sentence, a negligible change, is used here.

ii3 marks

What are the types of intermolecular forces? Give examples.

Solution

There are three types of intermolecular forces, listed here in the order in which the book introduces them:

Dipole-dipole forces: attraction between the positive end of one polar molecule and the negative end of another. Example: hydrogen chloride, \ce{HCl}.
Dipole-induced dipole (Debye) forces: attraction between a polar molecule and a nonpolar molecule whose electron cloud it distorts, so that the nonpolar molecule acquires an induced dipole. Example: water, \ce{H2O}, with oxygen, \ce{O2}.
London (dispersion) forces: momentary attraction between instantaneous dipoles and the dipoles they induce in neighbouring molecules. They act between all molecules, including nonpolar ones. Example: chlorine, \ce{Cl2}, bromine, \ce{Br2}, and iodine, \ce{I2}.

iii3 marks

What is hydrogen bonding? Give particular examples.

Solution

Hydrogen bonding is a strong dipole-dipole attraction. It occurs between a hydrogen atom that is bonded to a highly electronegative atom (fluorine, oxygen or nitrogen) and a lone pair of electrons on a fluorine, oxygen or nitrogen atom of another molecule. The bonded atom pulls electrons away from the hydrogen, so the hydrogen carries a large partial positive charge.

Examples are hydrogen fluoride (\ce{HF}), water (\ce{H2O}), ammonia (\ce{NH3}) and alcohols such as ethanol. In water the link is written \ce{O-H}\cdots\ce{O}: the hydrogen of one water molecule is attracted to the oxygen of a neighbouring molecule.

iv2025 Annual I, Set B, Q.2 (iii) OR3 marks

What are the applications of H-bonding?

Solution

The main application in this unit is to explain the unusual physical properties of water. The strong hydrogen bonds between water molecules account for:

High surface tension: water has a surface tension of 7.275\times10^{-2}\,\mathrm{N\,m^{-1}}, much higher than methanol, ethanol and benzene in the book's table.
Low vapour pressure: the strong attractions slow evaporation, so less vapour is produced at a given temperature.
High heat of vaporization: about 40.7\,\mathrm{kJ\,mol^{-1}}, because much energy is needed to break the hydrogen bonds before molecules can escape.
High boiling point: more heat is needed to raise the vapour pressure of water to the external pressure, so it boils at 100\,^\circ\mathrm{C} at 760\,\mathrm{torr}.

Hydrogen bonding also makes water more viscous than methanol.

v3 marks

What are the different types of physical properties of liquids?

Solution

The physical properties of liquids covered in this unit are:

Evaporation: the spontaneous change of a liquid into its vapour at the surface, at a given temperature.
Vapour pressure: the pressure exerted by the vapour in equilibrium with the liquid at a given temperature.
Boiling point: the temperature at which the vapour pressure of the liquid equals the external pressure.
Viscosity: the resistance to flow of a liquid.
Surface tension: the force in dynes acting at right angles to a unit length of the liquid surface.

vi3 marks

Define vapour pressure. What are the factors affecting the vapour pressure?

Solution

The vapour pressure of a liquid is the pressure exerted by its vapour when the vapour is in equilibrium with the liquid at a given temperature. At equilibrium the rate of evaporation equals the rate of condensation.

Temperature: a rise in temperature gives more molecules enough kinetic energy to escape, so the vapour pressure rises. Water rises from 4.579\,\mathrm{torr} at 0\,^\circ\mathrm{C} to 760\,\mathrm{torr} at 100\,^\circ\mathrm{C}.
Nature of the liquid: stronger intermolecular forces give a lower vapour pressure. At 20\,^\circ\mathrm{C} glycerol (strong attractions) has 0.00016\,\mathrm{torr}, while isopentane (weak attractions) has 580\,\mathrm{torr}.

The amount of liquid, the surface area and the volume of the container have no effect on the vapour pressure.

vii(a)3 marks

What is viscosity?

Solution

Viscosity is the resistance to flow of a liquid. A liquid flows as a series of layers that slide over one another, and each layer attracts the layers next to it. This internal friction between layers resists the flow, so the higher the viscosity, the more slowly the liquid flows.

The SI unit of viscosity is the pascal second, \mathrm{Pa\,s}. Since 1\,\mathrm{Pa\,s}=1\,\mathrm{kg\,m^{-1}\,s^{-1}}=10\,\mathrm{poise}, the poise is the non-SI unit.

vii(b)3 marks

What is surface tension?

Solution

Surface tension is the property of a liquid surface to behave as if a stretched membrane covered it. It is the force, in dynes, acting at right angles to a unit length of the liquid surface. Molecules at the surface are pulled inward by their neighbours, so the surface contracts to the smallest area it can.

The SI unit of surface tension is the newton per metre, \mathrm{N\,m^{-1}}, which is equal to \mathrm{J\,m^{-2}}.

viii3 marks

Define molar heat of fusion and molar heat of vaporization.

Solution

Molar heat of fusion (\Delta H_\mathrm{f}) is the amount of heat required to convert one mole of a solid into its liquid state at its melting point. For ice it is +6.02\,\mathrm{kJ\,mol^{-1}}: \ce{H2O(s) -> H2O(l)}.

Molar heat of vaporization (\Delta H_\mathrm{v}) is the amount of heat required to convert one mole of a liquid into its vapour at its boiling point. For water it is +40.7\,\mathrm{kJ\,mol^{-1}}: \ce{H2O(l) -> H2O(g)}.

Note. The book gives the molar heat of vaporization of water as 40.6 kJ/mol in section 5.2.4 and as +40.7 kJ/mol in section 5.3.1. The value 40.7 kJ/mol is used here.

ix3 marks

How will you differentiate liquid crystals from pure liquids?

Solution

Pure liquids have a random arrangement of molecules with no long-range order. They flow freely, take the shape of their container, and their molecular arrangement changes only with temperature or pressure.

Liquid crystals also flow like liquids, but their long rod-like molecules are partly ordered, lying in a common direction. Their molecular arrangement changes in response to temperature, electric field or pressure, and they have optical properties like those of crystals, which pure liquids do not have.

Note. Section 5.3.3 of the book says 'Liquids crystals are always isotropic while crystalline solid may be isotropic or anisotropic'. This is a misprint: liquid crystals are anisotropic, which is why they have crystal-like optical properties. The answer does not repeat the printed statement.

x3 marks

Why distillation under reduced pressure is often used in the purification of chemicals?

Solution

Distillation under reduced pressure (vacuum distillation) lowers the pressure on the liquid. The liquid then reaches its vapour pressure, and so boils, at a lower temperature. Less heat is needed, and substances that would otherwise decompose can be distilled.

Glycerine, for example, boils at 290\,^\circ\mathrm{C} at 760\,\mathrm{torr} and can decompose at that temperature. At 50\,\mathrm{torr} its boiling point falls to 210\,^\circ\mathrm{C}, and it distils without decomposition.

Long questionsp. 111–113

3(a)5 marks

Give the simple properties of liquids with special reference to the following: diffusion, compression, expansion, intermolecular forces and kinetic energy.

Solution

Diffusion

Molecules of a liquid move from place to place, but they are packed closely together, so their movement is restricted. Diffusion is therefore slow. An ink drop spreads slowly through water.

Compression

Liquids are almost incompressible, because their molecules are already in close contact and there is little empty space between them. Raising the pressure on water from one to two atmospheres changes its volume by a negligible amount.

Expansion

Liquids expand when they are heated. The intermolecular forces weaken as the temperature rises, and the molecules collide more effectively, so the molecules move further apart and the volume increases. The volume falls again when the liquid is cooled.

Intermolecular forces

The attractive forces between liquid molecules are greater than those between gas molecules, but they are not strong enough to hold the molecules in fixed positions. So the molecules can slide past one another. The strength of these forces decides the boiling point, vapour pressure, surface tension and viscosity of a liquid.

Kinetic energy

According to kinetic molecular theory, the average kinetic energy of liquid molecules is directly proportional to the absolute temperature. Because strong attractions limit the molecular movements and collisions, the kinetic energy of liquid molecules is moderate. In water, where hydrogen bonding is strong, it is low.

3(b)5 marks

Explain on the basis of kinetic molecular theory. Why the boiling point of a liquid remains constant although heat is continuously supplied to the liquid?

Solution

Kinetic energy and temperature

The temperature of a substance is a measure of the average kinetic energy of its molecules. When a liquid is heated below its boiling point, the kinetic energy of its molecules rises, and so does its temperature.

Start of boiling

When the vapour pressure of the liquid becomes equal to the external pressure, the molecules throughout the liquid can escape into vapour. Bubbles of vapour form inside the liquid and rise to the surface. This temperature is the boiling point.

Why the temperature stays constant

At the boiling point the molecules already have the maximum average kinetic energy that the liquid can hold at that temperature. Any extra heat cannot raise it further.
The heat that is still supplied is used to break the intermolecular forces, so that the molecules can separate and leave as vapour. This heat is the molar heat of vaporization.
Because the supplied energy goes into breaking the attractions (potential energy) and not into kinetic energy, the temperature stays at the boiling point until all the liquid has vaporized.

Heating a liquid at constant external pressure

The temperature rises while the liquid is heated, stays constant while it boils and rises again only after all the liquid has become vapour.

Result

The temperature stays at the boiling point while the heat is being used to convert the liquid into vapour. Only after the liquid has vaporized does the temperature of the vapour rise.

4(a)5 marks

Define and explain evaporation.

Solution

Definition

Evaporation is the spontaneous change of a liquid into its vapour at the surface of the liquid, at a given temperature. It takes place at all temperatures.

Explanation

The molecules of a liquid move with a range of kinetic energies. A molecule near the surface with high kinetic energy can overcome the attraction of its neighbours and escape into the air. Because the most energetic molecules leave, the average kinetic energy of the molecules left behind falls.

Cooling effect

Since the liquid loses its most energetic molecules, its temperature falls. Heat then flows from the surroundings into the liquid, so the surroundings cool as well.

4(b)5 marks

What are the factors affecting evaporation?

Solution

Evaporation depends on three main factors:

Surface area: a larger liquid surface gives more molecules that can escape, so evaporation is faster.
Intermolecular forces: strong attractions hold molecules in the liquid, so evaporation is slow. Weak forces, such as London forces, let molecules escape easily, so evaporation is fast.
Temperature: at a higher temperature the molecules have more kinetic energy, so more of them can escape and evaporation is faster.

4(c)5 marks

Different liquids have different rates of evaporation. Explain with reference to ether and alcohol.

Solution

The rate of evaporation depends on how strongly the molecules attract each other. Ethyl ether and ethyl alcohol (ethanol) evaporate much faster than water because their intermolecular forces are weaker than those of water.

Diethyl ether Ethyl ether: no O-H group, so no hydrogen bonds.

Ethanol Ethyl alcohol: has an O-H group, but its intermolecular forces are weaker than those of water.

Ethyl ether

Ether has no O-H bond, so its molecules cannot form hydrogen bonds. They are held together only by weak dipole-dipole and London forces. Its molecules escape easily from the surface, so it evaporates quickly. At 20\,^\circ\mathrm{C} its vapour pressure is 442.2\,\mathrm{torr}.

Ethyl alcohol and water

Water molecules are held together by strong hydrogen bonds. Ethanol forms weaker attractions, so its molecules escape more readily and it evaporates much faster than water.

Result

The rate of evaporation is ether > alcohol > water. The order follows the strength of the intermolecular forces, which is weakest in ether and strongest in water.

5(a)5 marks

Define and explain vapour pressure. How equilibrium is established between evaporation and condensation?

Solution

Definition

The vapour pressure of a liquid is the pressure exerted by its vapour in equilibrium with the liquid at a given temperature.

Establishing equilibrium

In a closed container the molecules that escape from the liquid surface build up as vapour above the liquid.
Some vapour molecules strike the liquid surface and are captured again. This return of vapour to the liquid is called condensation.
As the vapour builds up, condensation becomes faster. Evaporation stays the same at a fixed temperature, because it depends only on the liquid and its temperature.
When the rate of evaporation equals the rate of condensation, the amount of vapour no longer changes. The system is then in dynamic equilibrium.

Liquid and vapour in a closed container. At equilibrium the two arrows have equal rates.

\ce{H2O(l) <=> H2O(g)}

dynamic equilibrium

Points to note

The vapour pressure does not depend on the amount of liquid, the volume of the container or the surface area of the liquid. It depends on the nature of the liquid and on the temperature.

5(b)5 marks

What are the factors affecting vapour pressure of a liquid?

Solution

Vapour pressure depends on two main factors:

Temperature: when the temperature rises, more molecules have enough kinetic energy to break the intermolecular forces and escape, so the vapour pressure rises. Water rises from 4.579\,\mathrm{torr} at 0\,^\circ\mathrm{C} to 9.209\,\mathrm{torr} at 10\,^\circ\mathrm{C} and to 760\,\mathrm{torr} at 100\,^\circ\mathrm{C}.
Nature of the liquid: the strength of the intermolecular forces and the size of the molecules. At a given temperature stronger forces give a lower vapour pressure, and weaker forces give a higher one.

The table below compares three liquids at 20\,^\circ\mathrm{C}. Glycerol has strong attractions, while isopentane and ethyl ether have weak ones.

Vapour pressure of three liquids at 20 °C (from Table 5.1 of the book).
Liquid	Vapour pressure at 20\,^\circ\mathrm{C} (\mathrm{torr})
Glycerol	0.00016
Ethyl ether	442.2
Isopentane	580

The amount of liquid, the surface area and the volume of the container have no effect on vapour pressure.

Note. Table 5.1 of the book gives water as 43.9 torr at 20 °C. The accepted value is about 17.5 torr, so the printed figure looks wrong, and the value for mercury (0.012 torr) appears about ten times too high (about 0.0012 torr). The table above uses only the three liquids that are consistent with the book's text. The book's text gives 527.8 torr for water at 90 °C, while the accepted value is about 526 torr.

5(c)5 marks

Kinetically, how will you explain the effect of temperature on vapour pressure?

Solution

Kinetic explanation

Temperature is a measure of the average kinetic energy of the molecules. When the temperature rises, the kinetic energy of all the molecules rises, and a larger fraction of them have enough energy to break free of the attractions of their neighbours.
More molecules leave the surface each second, so the rate of evaporation increases.
More vapour molecules build up above the liquid, so condensation also speeds up. Equilibrium is then reached at a higher pressure of vapour.
The vapour pressure therefore rises with temperature. The rise is steep at high temperatures, as water shows between 90\,^\circ\mathrm{C} and 100\,^\circ\mathrm{C}.

Result: vapour pressure increases with temperature because a greater fraction of molecules has enough kinetic energy to overcome the intermolecular forces.

Note. The book's text gives 527.8 torr for water at 90 °C, which looks slightly high (the accepted value is about 526 torr). The statement about the steep rise is an explanation of the book's values, not a new value.

6(a)5 marks

Define and explain boiling point of a liquid.

Solution

Definition

The boiling point of a liquid is the temperature at which its vapour pressure becomes equal to the external (atmospheric) pressure. It is also called the boiling temperature.

Explanation

Below the boiling point the vapour pressure is less than the external pressure, so vapour forms only at the surface. At the boiling point vapour forms throughout the liquid, bubbles rise and escape, and boiling begins. Water at 760\,\mathrm{torr} boils at 100\,^\circ\mathrm{C}. At 23.7\,\mathrm{torr} it boils at 25\,^\circ\mathrm{C}.

Once boiling starts, further heat is used to break the intermolecular forces, so the temperature stays constant. The amount of heat needed to convert one mole of liquid into vapour at its boiling point is the molar heat of vaporization.

6(b)5 marks

How will you explain the effect of pressure on the boiling point of a liquid?

Solution

The boiling point of a liquid depends on the external pressure. A liquid boils when its vapour pressure equals the external pressure. If the external pressure is higher, the vapour pressure must be higher too, so the liquid has to be heated to a higher temperature before it boils. If the external pressure is lower, the liquid boils at a lower temperature.

Boiling point of water at different external pressures (values from the book).
External pressure	Boiling point of water	Place or example
23.7\,\mathrm{torr}	25\,^\circ\mathrm{C}	Low pressure
323\,\mathrm{torr}	69\,^\circ\mathrm{C}	Top of Mount Everest
700\,\mathrm{torr}	98\,^\circ\mathrm{C}	Murree hills
760\,\mathrm{torr}	100\,^\circ\mathrm{C}	Sea level
1489\,\mathrm{torr}	120\,^\circ\mathrm{C}	Raised pressure

Result: the boiling point rises when the external pressure rises and falls when the external pressure falls.

6(c)(i)5 marks

Practically, how will you explain the effect of increase of pressure on boiling point?

Solution

A pressure cooker is a closed container from which vapour cannot escape. The following steps show the effect of raising the pressure:

The lid is closed, so the vapour formed by the boiling water cannot escape and it builds up pressure above the water.
The external pressure on the water rises above 760\,\mathrm{torr}.
The water must now reach a higher vapour pressure before it boils, so its boiling point rises. At 1489\,\mathrm{torr} water boils at 120\,^\circ\mathrm{C}.
Food cooks in the hotter water and is cooked quickly.

Result: an increase of pressure raises the boiling point of the liquid.

6(c)(ii)5 marks

Practically, how will you explain the effect of decrease of pressure on boiling point?

Solution

Lowering the pressure above a liquid lowers its boiling point. This is used in vacuum distillation, where the pressure over the liquid is reduced so that it boils at a lower temperature.

Glycerine, for example, boils at 290\,^\circ\mathrm{C} at 760\,\mathrm{torr} and can decompose at that temperature, so it cannot be distilled at normal pressure. By reducing the pressure to 50\,\mathrm{torr} its boiling point falls to 210\,^\circ\mathrm{C}, and it can be distilled without decomposition.

The same principle explains water boiling at 98\,^\circ\mathrm{C} at 700\,\mathrm{torr} on the Murree hills and at 69\,^\circ\mathrm{C} at 323\,\mathrm{torr} on Mount Everest.

Result: a decrease of pressure lowers the boiling point.

7(a)5 marks

Define and explain the term viscosity of a liquid. How does the resistance to the layers cause viscosity?

Solution

Definition

Viscosity is the resistance to flow of a liquid. It measures how easily one layer of molecules slides past another.

Layers in flow

A fluid flowing in a pipe can be thought of as a series of concentric layers moving at different speeds. The layer touching the wall moves slowest, and the layers further from the wall move faster.

Speed of the layers of a liquid flowing in a pipe. The layer at the wall is slowest.

How the resistance causes viscosity

Each layer attracts the layers next to it. A faster layer is slowed down by its slower neighbour, and a slower layer is pulled forward by the faster one. This internal friction between the layers is the resistance to flow, which we call viscosity. The more strongly the molecules attract each other, the greater the friction and the higher the viscosity.

7(b)5 marks

What are the factors affecting the viscosity of a liquid?

Solution

Size and shape of molecules

Small, compact molecules such as those of water, acetone, benzene and methanol give a low viscosity. Large, irregular molecules such as those of glycerine tangle with one another, which hinders their flow, so the viscosity is high.

Intermolecular forces

The stronger the intermolecular forces, the higher the viscosity. Liquids whose molecules form hydrogen bonds are more viscous than liquids without them. For example, water is more viscous than methanol mainly because of its extensive hydrogen bonding.

Temperature

Viscosity falls as the temperature rises. The higher kinetic energy weakens the intermolecular forces, and the molecules move faster. This is why honey and syrup are easier to pour when they are hot than when they are cold.

7(c)(i)2026 Annual I, Set D, Q.2 (xiii)5 marks

Use the concept of hydrogen bonding to explain the following properties of water: high surface tension.

Solution

The surface tension of a liquid depends on the strength of its intermolecular forces. Water molecules form hydrogen bonds, which are strong attractions. The molecules at the surface of water are therefore pulled strongly inward by their neighbours, and the surface resists being stretched. This gives water a high surface tension.

Surface tension of some liquids at 20 °C (from the book).
Solvent	Surface tension (\gamma\times10^{-2}\,\mathrm{N\,m^{-1}})
Water	7.275
Methanol	2.26
Ethanol	2.28
Benzene	2.888
Hexane	1.84
\ce{CCl4}	2.70

The surface tension of water is far higher than that of the other liquids in the table, and this is explained by its hydrogen bonding.

7(c)(ii)5 marks

Use the concept of hydrogen bonding to explain the following properties of water: high heat of vaporization.

Solution

To vaporize water, its molecules must be separated from one another, which means breaking the hydrogen bonds that hold them together. Much energy is needed for this, so the heat of vaporization of water is high.

The molar heat of vaporization of water is +40.7\,\mathrm{kJ\,mol^{-1}}. This is much greater than the molar heat of fusion of ice, +6.02\,\mathrm{kJ\,mol^{-1}}, which shows how strongly hydrogen bonding resists the escape of the molecules into the vapour.

7(c)(iii)5 marks

Use the concept of hydrogen bonding to explain the following properties of water: high boiling point.

Solution

The boiling point of water is the temperature at which its vapour pressure equals the external pressure. Water's hydrogen bonds keep its molecules from escaping, so its vapour pressure is low at any given temperature. More heat must therefore be supplied before the vapour pressure reaches atmospheric pressure, which is why water boils at 100\,^\circ\mathrm{C} at 760\,\mathrm{torr}.

The high heat of vaporization shows that a large amount of energy is needed to turn water into vapour. This energy is supplied as heat at the boiling point, so the boiling point is high compared with the other liquids in the book.

Result: the hydrogen bonds of water make its boiling point high.

8(a)5 marks

Define and explain the phenomenon of surface tension.

Solution

Definition

Surface tension is the property of a liquid surface to behave as if a stretched membrane covered it. It is the force, in dynes, acting at right angles to a unit length of the liquid surface. It can also be defined as the energy needed to expand the surface by one unit of area.

Explanation

A molecule in the bulk of the liquid is pulled equally in all directions by its neighbours, so the forces balance. A molecule at the surface has no neighbours above it, so it is pulled inward and the surface is placed under tension.

Molecules at the surface of a liquid feel an unbalanced inward pull.

Why a drop is spherical

Molecules at the surface are less stable than those inside, so a liquid tends to have the fewest molecules on its surface. A sphere has the least surface area per unit volume, so small droplets are spherical.

8(b)5 marks

What are the factors affecting surface tension?

Solution

Surface tension depends on the intermolecular forces and on the temperature:

Intermolecular forces: the stronger the forces between the molecules, the higher the surface tension. Water, with its hydrogen bonds, has a higher surface tension than alcohols, ethers and benzene.
Temperature: surface tension falls as the temperature rises. The increased kinetic energy weakens the intermolecular forces that pull the surface molecules inward.

Surface tension is different for different liquids because the types of intermolecular forces differ.

Define dynamic equilibrium between two physical states.

Solution

Dynamic equilibrium between two physical states is the condition in which a change of state and its reverse change take place at the same rate in a closed system. Because the two rates are equal, the amount of each state stays constant, although molecules are continually changing from one state to the other.

Example: in a closed container at constant temperature, \ce{Liquid <=> Vapour}. Molecules evaporate from the liquid and condense back at equal rates.

8(d)(i)3 marks

Define molar heat of fusion.

Solution

Molar heat of fusion (\Delta H_\mathrm{f}) is the amount of heat required to convert one mole of a solid into its liquid state at its melting point. For ice, \ce{H2O(s) -> H2O(l)} and \Delta H_\mathrm{f} = +6.02\,\mathrm{kJ\,mol^{-1}}.

8(d)(ii)3 marks

Define molar heat of vaporization.

Solution

Molar heat of vaporization (\Delta H_\mathrm{v}) is the amount of heat required to convert one mole of a liquid into its vapour at its boiling point. For water, \ce{H2O(l) -> H2O(g)} and \Delta H_\mathrm{v} = +40.7\,\mathrm{kJ\,mol^{-1}}.

Note. Section 5.2.4 of the book gives 40.6 kJ/mol for water and section 5.3.1 gives 40.7 kJ/mol. The later value is used.

9(a)3 marks

Define a liquid crystal.

Solution

A liquid crystal is the state that exists between the melting temperature and the clearing temperature of certain substances. It is a turbid liquid that flows like a liquid, but whose long rod-like molecules have some order, so that it shows optical properties like those of crystals.

Example: cholesteryl benzoate melts to a milky liquid crystal at 145\,^\circ\mathrm{C} and becomes a clear liquid at 179\,^\circ\mathrm{C}.

9(b)5 marks

What are the uses of liquid crystals in daily life?

Solution

Liquid crystals are used in the following ways, as described in the book:

Displays: digital watches, laptop computers, calculators, television displays, LCD screens and oscilloscope displays.
Medicine: cholesteric liquid crystals applied to the skin locate veins, arteries, infections and tumours, which are warmer than the surrounding tissue.
Magnetic resonance: molecules dissolved in nematic liquid crystal solvents give resolved NMR spectra, showing fine dipole-dipole structure.
Analysis: liquid crystals are used in chromatographic separation to study molecular arrangements and kinetics, and as an anisotropic fluid for visible, UV and IR spectroscopy.
Industry: liquid crystal polymers, such as polyester liquid crystals, are used as fire-resistant coatings for multifibre and optical cables.

9(c)5 marks

How will you differentiate liquid crystals from pure liquids and crystalline solids?

Solution

Comparison of pure liquids, liquid crystals and crystalline solids.
Property	Pure liquid	Liquid crystal	Crystalline solid
Arrangement of molecules	Random	Arranged as in a crystal, with some freedom to move	Regular, repeating three-dimensional pattern
Long-range order	None	Some	Complete
Flow	Flows freely and takes the shape of its vessel	Flows like a liquid	Does not flow; fixed shape and volume
Response to outside influence	Changes only with temperature or pressure	Molecular arrangement changes with temperature, electric field or pressure	Changes phase on heating, at its melting point

Liquid crystals therefore lie between the crystalline solid and the pure liquid. Like crystals they have some molecular order, and like liquids they can flow.

103 marks

What are the energetics of phase changes?

Solution

A change of state takes place at constant temperature, and the heat that is absorbed or released is a measure of the strength of the intermolecular forces. The change in enthalpy at constant pressure is written \Delta H, in \mathrm{kJ\,mol^{-1}}.

Melting (fusion): \ce{H2O(s) -> H2O(l)}, \Delta H_\mathrm{f} = +6.02\,\mathrm{kJ\,mol^{-1}}. The change in intermolecular distance and potential energy is small.
Vaporization: \ce{H2O(l) -> H2O(g)}, \Delta H_\mathrm{v} = +40.7\,\mathrm{kJ\,mol^{-1}}. The molecules move far apart, so the change in intermolecular distance and potential energy is large.

The heat of vaporization is therefore much greater than the heat of fusion, because separating all the molecules of a liquid requires far more energy than loosening the arrangement of a solid.

11(a)5 marks

Define and explain the boiling point of a liquid.

Solution

Definition

The boiling point of a liquid is the temperature at which its vapour pressure becomes equal to the external pressure. At this temperature vapour forms throughout the liquid and boiling begins. At the normal pressure of 760\,\mathrm{torr}, the boiling point of water is 100\,^\circ\mathrm{C}.

Explanation

Strong intermolecular forces make molecules harder to release, so the liquid must be heated to a higher temperature before its vapour pressure reaches the external pressure. A liquid with strong forces therefore has a high boiling point.

Once boiling begins, the temperature stays constant. The heat supplied is used to break the intermolecular forces and convert the liquid into vapour.

11(b)5 marks

How will you explain the two practical applications regarding the effect of pressure on the boiling point of a liquid?

Solution

1. Pressure cooker (raised pressure)

In a closed pressure cooker the pressure above the water rises, so the water boils only at a higher temperature. At 1489\,\mathrm{torr} water boils at 120\,^\circ\mathrm{C}. Food cooks faster in this hotter water.

2. Vacuum distillation (lowered pressure)

Some liquids decompose at their normal boiling points. Glycerine boils at 290\,^\circ\mathrm{C} at 760\,\mathrm{torr}. Lowering the pressure to 50\,\mathrm{torr} lowers its boiling point to 210\,^\circ\mathrm{C}, so it can be distilled without decomposition.

123 marks

How can you interpret the anomalous behaviour of water?

Solution

Water shows unusual behaviour compared with other liquids: it has a high surface tension, a low vapour pressure, a high heat of vaporization, a high boiling point, and it is more viscous than methanol.

All these properties are explained by hydrogen bonding. Each water molecule forms several hydrogen bonds with its neighbours. These strong attractions keep the molecules together, so the surface is pulled inward strongly (high surface tension), fewer molecules escape (low vapour pressure), more energy is needed to vaporize water (high heat of vaporization and high boiling point), and the flow is hindered (high viscosity).

133 marks

Evaluate the impact of temperature on surface tension of liquids.

Solution

Temperature has a clear effect: surface tension decreases as the temperature rises. At a higher temperature the molecules have more kinetic energy, which weakens the intermolecular forces that pull the surface molecules inward. The surface is then pulled less strongly and its surface tension falls.

The fall is different for different liquids, because it depends on how strong their intermolecular forces are. Temperature is therefore an important factor, but it changes surface tension less than the type of intermolecular force does: liquids with strong forces, such as water, have a much higher surface tension at any temperature.

14(a)5 marks

Interpret the influence of intermolecular forces on the following properties of liquids: vapour pressure.

Solution

The vapour pressure of a liquid depends on how strongly its molecules attract one another. Strong forces hold the molecules in the liquid, so only a few of them have enough energy to escape, and the vapour pressure is low. Weak forces let the molecules escape easily, so the vapour pressure is high.

Influence of intermolecular forces on vapour pressure (values from the book).
Liquid (at 20\,^\circ\mathrm{C})	Intermolecular forces	Vapour pressure (\mathrm{torr})
Glycerol	Strong (extensive hydrogen bonding)	0.00016
Ethyl ether	Weak (no O-H group)	442.2
Isopentane	Weak (London forces only)	580

At the same temperature, stronger intermolecular forces give a lower vapour pressure.

14(b)5 marks

Interpret the influence of intermolecular forces on the following properties of liquids: boiling point.

Solution

The boiling point of a liquid is the temperature at which its vapour pressure equals the external pressure. A liquid with strong intermolecular forces has a low vapour pressure at a given temperature, so it has to be heated to a higher temperature before its vapour pressure reaches atmospheric pressure. Strong forces therefore give a high boiling point.

Glycerol, with strong attractions, boils at 290\,^\circ\mathrm{C} at 760\,\mathrm{torr}. Isopentane, with only weak London forces, has a much higher vapour pressure at the same temperature and so boils at a much lower temperature.

155 marks

Evaluate the importance of H-bonding in understanding physical properties of water.

Solution

Properties that hydrogen bonding explains

High surface tension, because the strong attractions pull the surface molecules inward.
Low vapour pressure, because the molecules are held together and escape slowly.
High heat of vaporization (+40.7\,\mathrm{kJ\,mol^{-1}}), because the hydrogen bonds must be broken before the molecules can vaporize.
High boiling point (100\,^\circ\mathrm{C} at 760\,\mathrm{torr}), which follows from the low vapour pressure and high heat of vaporization.
High viscosity, greater than that of methanol, because of extensive hydrogen bonding.

Evaluation

Hydrogen bonding is the most important factor in explaining why water behaves differently from other liquids. Each of its unusual properties can be traced to the same strong attractions between molecules, so one concept explains a whole set of observations.

The explanation has limits, however. London forces are also present in water, and hydrogen bonding explains trends and differences rather than exact numerical values. Its importance is greatest for water because hydrogen bonding there is both strong and extensive.

163 marks

Differentiate between molar heat of fusion and molar heat of vaporization.

Solution

Molar heat of fusion and molar heat of vaporization compared.
Feature	Molar heat of fusion (\Delta H_\mathrm{f})	Molar heat of vaporization (\Delta H_\mathrm{v})
Change of state	Solid to liquid	Liquid to vapour
Temperature at which it occurs	Melting point	Boiling point
Example (water)	+6.02\,\mathrm{kJ\,mol^{-1}} (ice to water)	+40.7\,\mathrm{kJ\,mol^{-1}} (water to steam)
Change in intermolecular distance and potential energy	Small	Large
Intermolecular forces overcome	Partly loosened, as the solid arrangement breaks	Almost all of them overcome, as molecules separate

The heat of vaporization is much greater than the heat of fusion, because a liquid's molecules must be fully separated to form a vapour.

FBISE 2025 and 2026 questions

Questions from the FBISE HSSC-I 2026 (1st Annual) and FBISE HSSC-I 2025 (1st Annual) papers that belong to this unit, answered the way an examiner expects.

2026 Annual I, Set B, MCQ 161 mark

London forces occur in ______ molecules.

AOnly polar
BOnly ionic
CCovalentAnswer
DOnly non-polar

Why

London (dispersion) forces are the momentary attractions between an instantaneous dipole and an induced dipole. They act between all molecules, polar and non-polar, so they occur in covalent molecules. They are the only intermolecular force between non-polar molecules such as \ce{Cl2}, \ce{Br2} and \ce{I2}.

Option D (only non-polar) is the tempting wrong answer: London forces are the only force in non-polar molecules, but polar molecules have them as well. Ionic compounds (B) are held by ionic bonds, not by London forces between molecules.

2026 Annual I, Set B, Q.2 (ii) OR3 marks

State three properties of crystalline solids.

Solution

A crystalline solid has a definite, regular, three-dimensional arrangement of its particles (atoms, ions or molecules), as in \ce{NaCl}, diamond, ice and glucose. Three properties of crystalline solids are:

1. Definite geometrical shape (1 mark)

Crystals have a definite, distinctive geometric shape because of the orderly arrangement of their particles in three dimensions. The angles between the faces of a crystal are always the same, however the crystal is grown.

2. Sharp melting point (1 mark)

On heating, the particles vibrate more and more. At the melting point their vibrational energy becomes high enough for all the particles to leave their fixed positions together, so the solid melts completely at one sharp temperature.

3. Cleavage planes (1 mark)

Crystalline solids break along certain definite planes, called cleavage planes. When pressure is applied along a cleavage plane, the crystal splits into smaller crystals of the same shape as the original. The angles between the cleavage planes differ from one crystalline solid to another.

2025 Annual I, Set B, Q.2 (iii) OR3 marks

Describe any three properties of water which are due to hydrogen bonding.

Solution

A water molecule forms hydrogen bonds with its neighbours: the \ce{H} of one molecule is attracted to the \ce{O} of the next. These are strong intermolecular forces, and they explain the following properties of water.

Hydrogen bonds (dashed) between water molecules: the hydrogen of one molecule is attracted to the oxygen of the next. These strong intermolecular forces explain the properties of water.

1. High surface tension (1 mark)

The strong attractions pull the molecules at the surface strongly inward, so the surface resists being stretched. Water has a surface tension of 7.275\times10^{-2}\,\mathrm{N\,m^{-1}}, much higher than methanol, ethanol or benzene.

2. Low vapour pressure (1 mark)

The strong hydrogen bonds make evaporation slow, because few molecules have enough energy to escape. Less vapour is present above the liquid, so water has a low vapour pressure.

3. High heat of vaporization and high boiling point (1 mark)

The hydrogen bonds must be broken before molecules can escape, so much heat is needed: \Delta H_v=+40.7\,\mathrm{kJ\,mol^{-1}}. Because the vapour pressure is low, water must be heated to a higher temperature before its vapour pressure equals the external pressure, so its boiling point is high (100\,^\circ\mathrm{C} at 760\,\mathrm{torr}).

2025 Annual I, Set B, Q.2 (v)3 marks

State any three differences between crystalline and non-crystalline solids.

Solution

Solids are of two types. Crystalline solids have a regular, repeating three-dimensional arrangement of particles. Amorphous solids have no regular, orderly arrangement of particles (apart from small ordered regions called crystallites).

	Crystalline solids	Amorphous solids
Arrangement of particles	Definite, regular, repeating three-dimensional pattern	No regular, orderly arrangement
Shape	Definite, distinctive geometric shape	No definite geometric shape
Melting point	Sharp melting point	No sharp melting point; they soften over a range of temperature
Examples	\ce{NaCl}, diamond, ice, glucose, \ce{ZnSO4}	Glass, plastic, rubber

In a crystalline solid the particles form a regular pattern that repeats in three dimensions. An amorphous solid has no regular pattern, apart from small ordered regions called crystallites.

Any three of the differences above are enough for 3 marks. A crystalline solid can be made amorphous by melting it and cooling the molten mass rapidly, so that the particles do not get time to arrange themselves.

2026 Annual I, Set B, Q.2 (v) OR2+1 = 3 marks

Identify intermolecular forces among the following molecules. Also write in descending order of their strength. \ce{C2H2}, \ce{NH3}, \ce{HCl}, \ce{H2O}

Solution

The intermolecular force depends on the type of the molecule: a molecule with \ce{H} bonded to \ce{F}, \ce{O} or \ce{N} forms hydrogen bonds (a strong dipole-dipole attraction); a polar molecule has dipole-dipole forces; a non-polar molecule has only London forces (instantaneous dipole-induced dipole forces). London forces are present in all of them, but the strongest force present decides the order.

Molecule	Polarity	Intermolecular force	Order (1 = strongest)
\ce{H2O}	Polar, \ce{O-H} bonds	Hydrogen bonding (two \ce{O-H} per molecule)	1
\ce{NH3}	Polar, \ce{N-H} bonds	Hydrogen bonding	2
\ce{HCl}	Polar, no \ce{H} on \ce{F}, \ce{O}, \ce{N}	Dipole-dipole forces	3
\ce{C2H2}	Non-polar (symmetrical)	London forces only	4

Descending order of strength of the intermolecular forces (2 marks for the forces, 1 mark for the order):

\ce{H2O} > \ce{NH3} > \ce{HCl} > \ce{C2H2}

Hydrogen bonds are the strongest, dipole-dipole forces are next and London forces are the weakest. Between the two hydrogen-bonded molecules, water forms more hydrogen bonds per molecule than ammonia (two hydrogen atoms and two lone pairs), so the attraction in water is the larger one.

2026 Annual I, Set B, Q.2 (xiii)3 marks

Water has low vapour pressure than \ce{CCl4} and ether. Why?

Solution

The vapour pressure of a liquid at a given temperature depends on the strength of the intermolecular forces: the stronger the forces, the lower the vapour pressure.

Water molecules are held together by hydrogen bonding, a strong intermolecular force. Few molecules have enough energy to break free from the surface, so evaporation is slow and little vapour is present above the liquid. The vapour pressure of water is therefore low.
\ce{CCl4} is non-polar and has only weak London forces, and ether has only weak dipole-dipole and London forces with no \ce{O-H} group to form hydrogen bonds. Their molecules escape easily, so their vapour pressures are high.

Liquid (at 20\,^\circ\mathrm{C})	Main intermolecular force	Vapour pressure (torr)
Water	Hydrogen bonding (strong)	low (the book table gives 43.9, the smallest of the three)
Carbon tetrachloride	London forces (weak)	87
Ethyl ether	Weak dipole-dipole and London forces	442.2

So water, with the strongest forces, has a lower vapour pressure than \ce{CCl4} and ether.

2026 Annual I, Set D, MCQ 161 mark

Dipole forces are present in ______ molecules.

ANon-polar
BNon-metallic
CMetallic
DPolarAnswer

Why

Dipole-dipole forces are the attraction between the positive end of one polar molecule and the negative end of another. A molecule must have a permanent dipole to show them, so they are present in polar molecules (for example \ce{HCl}).

Non-polar molecules (A) have no permanent dipole and show only London forces. Metallic substances (C) are held by metallic bonds, and 'non-metallic' (B) is not a measure of polarity: \ce{Cl2} and \ce{CH4} are non-metallic but non-polar.

2026 Annual I, Set D, Q.2 (ii) OR1+1+1 = 3 marks

State three differences between crystalline and amorphous solids.

Solution

	Crystalline solids	Amorphous solids
Arrangement of particles	Definite, regular, repeating three-dimensional pattern	No regular, orderly arrangement
Shape	Definite, distinctive geometric shape	No definite geometric shape
Melting point	Sharp melting point	No sharp melting point; they soften over a range of temperature
Examples	\ce{NaCl}, diamond, ice, glucose, \ce{ZnSO4}	Glass, plastic, rubber

In a crystalline solid the particles form a regular pattern that repeats in three dimensions. An amorphous solid has no regular pattern, apart from small ordered regions called crystallites.

2025 Annual I, Set D, Q.2 (ii) OR1+2 = 3 marks

What is dynamic equilibrium? How does it exist between two physical states of a substance? Describe by giving an example.

Solution

Dynamic equilibrium (1 mark)

Dynamic equilibrium is the state in which two opposite changes take place at the same rate in a closed system, so that the amounts of the two states stay constant even though the particles continue to change from one state to the other.

Equilibrium between two physical states (2 marks)

Take a liquid in a closed container at constant temperature. Molecules with high kinetic energy escape from the surface (evaporation). The vapour molecules collide with the surface and are captured by the liquid (condensation). At first evaporation is faster, but as the vapour builds up, condensation speeds up until the two rates are equal. This is the state of dynamic equilibrium between the liquid and its vapour:

\ce{Liquid <=> Vapour}

Dynamic equilibrium in a closed container: the rate of evaporation equals the rate of condensation, so the pressure of the vapour stays constant.

At this stage the number of molecules leaving the liquid is equal to the number returning to it, so the pressure of the vapour (the vapour pressure) stays constant at that temperature. At one moment a molecule is in the liquid and the next moment it is in the vapour, which is why the equilibrium is called dynamic and not static. Example: water in a closed bottle, where \ce{H2O(l) <=> H2O(g)}.

2026 Annual I, Set D, Q.2 (v) OR2+1 = 3 marks

Identify intermolecular forces among the following molecules. Also write in descending order of their strength. \ce{C6H6}, \ce{HF}, \ce{H2O}, \ce{HBr}

Solution

Look at the type of intermolecular force in each molecule. The two with an \ce{H} atom on \ce{F} or \ce{O}, \ce{HF} and \ce{H2O}, form hydrogen bonds. \ce{HBr} is polar without hydrogen bonding, and \ce{C6H6} is non-polar.

Molecule	Polarity	Intermolecular force	Order (1 = strongest)
\ce{H2O}	Polar	Hydrogen bonding (up to two per molecule, a network)	1
\ce{HF}	Polar	Hydrogen bonding (strong, but about one per molecule)	2
\ce{HBr}	Polar	Dipole-dipole forces	3
\ce{C6H6}	Non-polar	London forces only	4

\ce{H2O} > \ce{HF} > \ce{HBr} > \ce{C6H6}

Hydrogen bonds are stronger than dipole-dipole forces, which are stronger than London forces. Each \ce{H-F} hydrogen bond is strong because fluorine is the most electronegative atom, but a water molecule can form two hydrogen bonds each (two \ce{H} atoms and two lone pairs), so the overall attraction in water is greater; this is why water boils at 100\,^\circ\mathrm{C} and \ce{HF} at 19.5\,^\circ\mathrm{C}.

2025 Annual I, Set D, Q.2 (v) OR1.5+1.5 = 3 marks

Explain why: (i) is the falling drop of water spherical in shape? (ii) has water high surface tension than ether?

Solution

(i) Why a falling drop of water is spherical (1.5 marks)

A molecule inside a liquid is pulled equally in all directions, but a molecule at the surface is pulled inward by an unbalanced force because there are no neighbours above it (surface tension). Surface molecules are less stable than those inside, so work is needed to bring molecules to the surface, and a liquid is most stable when it has the fewest molecules on its surface, that is the smallest surface area.

The molecule inside a liquid is pulled equally in all directions, but a molecule at the surface has no neighbours above it and feels a net pull inward.

A sphere has less surface area per unit volume than any other shape. So a free-falling drop pulls itself into a sphere.

(ii) Why water has a higher surface tension than ether (1.5 marks)

The surface tension of a liquid depends directly on the strength of the intermolecular forces between its molecules. Water molecules are joined by strong hydrogen bonds, so the inward pull on the surface molecules is large and the surface tension is high (7.275\times10^{-2}\,\mathrm{N\,m^{-1}}). Ether has no \ce{O-H} group and cannot form hydrogen bonds with its own molecules; it has only weak dipole-dipole and London forces, so its surface tension is much lower.

2026 Annual I, Set D, Q.2 (xiii)3 marks

Surface tension of water is about three times more than \ce{CCl4}. Why?

Solution

The surface tension of a liquid depends directly on the strength of its intermolecular forces: the stronger the forces, the higher the surface tension. A surface molecule is pulled inward by its neighbours, and the stronger the pull, the more the surface resists being stretched.

Water has strong hydrogen bonds between its molecules, so the surface molecules are pulled strongly inward. Its surface tension is 7.275\times10^{-2}\,\mathrm{N\,m^{-1}}.
\ce{CCl4} is a non-polar molecule with only weak London forces, so the inward pull is much smaller. Its surface tension is only 2.70\times10^{-2}\,\mathrm{N\,m^{-1}}.

\dfrac{\gamma(\ce{H2O})}{\gamma(\ce{CCl4})}=\dfrac{7.275}{2.70}\approx2.7\approx3

So the surface tension of water is about three times that of \ce{CCl4}, because of hydrogen bonding in water.

Exam focus

Use hydrogen bonding to explain why water has a high surface tension and a low vapour pressure, and a high heat of vaporization and boiling point. (long)
Asked in all four sets: 2026 Set B Q.2 (xiii) (low vapour pressure against \ce{CCl4} and ether), 2026 Set D Q.2 (xiii) (surface tension against \ce{CCl4}), 2025 Set B Q.2 (iii) OR and 2025 Set D Q.2 (v) OR. It is also book short question iv and long questions 7(c), 12 and 15.
Go to answer
Identify the intermolecular forces in given molecules and arrange them in descending order of strength. (short)
Asked in both 2026 sets (Set B Q.2 (v) OR with \ce{C2H2}, \ce{NH3}, \ce{HCl}, \ce{H2O}; Set D Q.2 (v) OR with \ce{C6H6}, \ce{HF}, \ce{H2O}, \ce{HBr}). The rule: hydrogen bonding, then dipole-dipole, then London forces.
Go to answer
State the differences between crystalline and amorphous solids, and the properties of crystalline solids. (short)
Asked in three sets: 2026 Set B Q.2 (ii) OR, 2026 Set D Q.2 (ii) OR and 2025 Set B Q.2 (v). It is a sure 3 marks if you learn the table of arrangement, shape and melting point.
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Define surface tension, explain why a drop is spherical, and give the factors that affect it. (long)
Asked in 2025 Set D Q.2 (v) OR (spherical drop and water against ether). The book's long questions 8(a), 8(b) and 13 cover the definition, the factors and the effect of temperature.
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Define dynamic equilibrium between a liquid and its vapour, and explain vapour pressure and evaporation. (short)
Asked in 2025 Set D Q.2 (ii) OR (moved from unit 8 because the book teaches it in unit 5), and the base of book long questions 5(a) and 8(c).
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State which molecules show London forces and which show dipole-dipole forces. (short)
Two MCQs in 2026: Set B MCQ 16 (London forces occur in covalent molecules) and Set D MCQ 16 (dipole forces in polar molecules). Compare book exercise MCQ ii and iv.
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Explain the effect of external pressure on the boiling point (pressure cooker and vacuum distillation). (long)
Not in the 2025 or 2026 papers, but it is the book's long questions 6(b), 6(c) and 11(b) and short question x, so it is a likely 5 to 6 mark question.
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Define viscosity, liquid crystals, molar heat of fusion and molar heat of vaporization. (short)
Not in the 2025 or 2026 papers. Book short questions vii, viii and ix and long questions 7(a), 9(a) and 16; these are standard 2 to 3 mark definitions.
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