Unit 5: States and Phases of Matter — Long Questions

Matter exists in four states i.e., solid, liquid, gas and plasma. The simplest form of matter is the gaseous state and most of matter around us in the solid state. Liquids are less common than solids, gases and plasmas. The reason is that the liquid state of any substance can exist only within a relatively narrow range of temperature and pressure.

Properties of gases

(i) Shape and volume: Gases do not have a definite shape and volume. The volume of a gas is the volume of the container

(ii) Ideal gas: Ideal gases have zero particle volume.

(iii) Intermolecular Spaces: The molecules of gases are widely separated from one another and most of the volume of the gas is empty space (nearly 99.9%). So, gases can be compressed easily. When sudden expansion of gases occurs cooling takes place. It is called Joule-Thomson effect.

(iv) Collision: Pressure of a gas is due to the collisions of gas molecules with the walls of the container

(v) Intermolecular forces: There are negligible intermolecular forces in ideal gases.

Properties of liquids

(i) Diffusion: Liquids also diffuse like gases, however the rate of diffusion in liquids is much lower than that in gases.

(ii) Spaces: There are little spaces among liquid molecules because of relatively stronger intermolecular forces.

(iii) Compressible: Liquids are 10⁵ times less compressible than gases, but about 10 times more compressible than solids.

(iv) Motion: The molecules of a liquid are in constant random motion, but their speed of movement is lower than gases. Diffusion among the miscible liquids is an evidence for the free motion of liquid molecules. Similarly, liquids can flow because of the free movement of their molecules.

Ideal gas

It is a matter of common observation that when external conditions of temperature and pressure are changed, the volume of a given quantity of a gas is affected.

Gas Laws The gas laws describe the relationships between volume of a given amount of a gas and the prevailing conditions of temperature and pressure.

In Boyle's and Charles's laws, some of the variables are held constant during the changes produced in the gases. According to Boyle's law.

V ∝ 1/P (when 'n' and 'T' are held constant)

According to Charles's law

V ∝ T (where 'n' and 'P' are held constant)

According to Avogadro's law "the volume of the given gas at constant temperature and pressure is directly proportional to the number of moles".

V ∝ n (when 'P' and 'T' are held constant)

Three relationships can be joined together if no variable is to be kept constant. By combining gas laws:

V ∝ nT/P

V = constant nT/P

The constant suggested is 'R' which is called ideal gas constant. Its value is 0.0821 atm dm³ K⁻¹ mol⁻¹. Its value depends upon units of pressure and volume. The value for 'R' in SI unit is 8.3143 Nm K⁻¹mol⁻¹.

Hence V = R nT/P

PV = nRT

This equation is called the ideal gas equation.

Calculation of Relative Molecular Mass (Mr) of a Gas

General gas equation can be employed to calculate the relative molecular mass of a gas whose P, T, V and mass in grams are known. This is achieved by rearranging equation.

Putting n = m/M in equation.

Where m = mass in g and M = molar mass of the gas

PV = mRT/M

Now rearranging equation

M = mRT/PV

At molecular level, the molar mass becomes relative molecular mass of a compound, therefore, we can write

Mr = mRT/PV

Sample Problem 5.1

The volume of 134g of a gas at -73°C under 10 atm pressure is 5 dm³ Determine the relative molecular mass of the gas.

Solution The molecular mass of the gas can be calculated by using the formula:

M = mRT/PV

M = (134g)(0.0821 atm dm³ mol⁻¹ K⁻¹)(200 K) / (10 atm)(5 dm³)

M = 44 g mol⁻¹

Therefore, the relative molecular mass of the gas is 44 amu.

Quick Check 5.1

(a) Explain why gases can be compressed easily?

Ans. Gases can be compressed easily because they have enough empty spaces among their molecules (about 99%). That's why they are compressible.

(b) What is Joule-Thomson effect?

Ans. When a compressed gas is allowed to expand suddenly, then it creates cooling effect. It is called "Joule-Thomson effect".

(c) The volume of 21g a gas is 2 dm³ at -90°C under 7 atm pressure. Calculate the relative molecular mass of the gas.

Ans.
mass = 21g
Volume = 8 dm³
Temperature = -90°C = 183K
Pressure = 7 atm
R = 0.0821 atm dm³ mol⁻¹ K⁻¹
M. mass = ?

Solution

PV = m/M RT

M = mRT/PV

M = (21g)(0.0821 atm dm³ mol⁻¹K⁻¹)(183K) / (7 atm)(8 dm³)

M = 5.63 gmol⁻¹

Therefore, the relative molecular mass of the gas is 5.63 amu.

Intermolecular Forces

Definition The forces of attraction among the molecules are called inter molecular forces. The intermolecular forces among the molecules of a liquid are reasonably stronger. Three types of such forces are

(i) Instantaneous dipole-induced dipole forces (id-id) (London dispersion forces).

(ii) Permanent dipole-permanent dipole dipole forces.

(iii) Hydrogen bonding.

Permanent dipole-permanent dipole dipole forces (pd-pd)

Definition The force of attraction between the positive end of a polar molecule and the negative end of a nearby polar molecule is called permanent dipole-permanent dipole dipole force.

Examples Whenever the HCl molecules are close to each other (as in liquid state), they tend to line up as in figure. Similar type of forces is present among the molecules of chloroform (CHCl₃).

The molecules of chloroform have a positive center at the H atom and the negative center on the end with Cl atoms.

Instantaneous dipole-induced dipole forces (id-id)

Definition The momentary force of attraction between an instantaneous dipole and an induced dipole is called instantaneous dipole-induced dipole force.

Examples They are present in all molecules but significant in non-polar molecules like F₂, Cl₂ (Halogens) and He, Ne (noble gases) etc.

These are only forces that exist among the molecules of non-polar compounds.

Factors affecting Instantaneous dipole-induced dipole forces (id-id)

The strength of these forces depend upon the following factors.

(i) Molecular mass and size

When the molecular mass (Mr) of a compound is greater, the id-id forces are stronger and vice versa.

Reason A larger atom or molecule the dispersion of electronic cloud and polarization is relatively easier.

Example This factor operates in the halogen and noble gas families.

Physical state The physical state changes in the halogen family from top to bottom due to the increase in size and polarizability. The first two members in the halogen family (F₂ and Cl₂) are gases, but bromine is a liquid and iodine a solid (B.P – 458K). The boiling points of noble gases in group 18 also increase from helium (He) to xenon (Xe) as in table 5.1.

Table 5.1 Molar masses and bonding points of group 17 and 18 of periodic table.

17 (VII-A) | Molecular Mass | Boiling points (K) | 18 (VIII) | Molecular mass | Boiling points (K)
F₂ | 38 | 85 | He | 4.003 | 4.22
Cl₂ | 71 | 239 | Ne | 20.18 | 27.1
Br₂ | 159 | 333 | Ar | 39.95 | 87.3
I₂ | 253.8 | 458 | Kr | 83.8 | 120
| | | Xe | 131.3 | 165

(ii) Surface area (shape of molecule)

A molecule with larger surface area has more number of contacts for a nearby molecule. The forces among the molecules of a compound having larger surface area are stronger.

Saturated Hydrocarbons

• These have chains of carbon atoms linked with hydrogen atoms.
• Compare the length of the chains in C₂H₆ and C₆H₁₄. They have the boiling points -88.6°C and 68.7°C, respectively.
• Molecules with larger chain length experiences stronger attractive forces.

Reason The longer molecules have more places along the length where they can be attracted to other molecules. Boiling points of some hydrocarbons are given in the following table.

Table Boiling points and physical states of some hydrocarbons

Name | B.P.°C (1 atm) | Physical state | Name | B.P.°C (1 atm) | Physical state
Methane (CH₄) | -164 | Gas | Pentane (C₅H₁₂) | 36.1 | Liquid
Ethane (C₂H₆) | -88.6 | Gas | Hexane (C₆H₁₄) | 68.7 | Liquid
Propane (C₃H₈) | -42.1 | Gas | Decane (C₁₀H₂₂) | 174.1 | Liquid
Butane (C₄H₁₀) | -0.5 | Gas | Isodecane (C₂₀H₄₂) | 327 | Solid

Isomers

• Three isomers of pentane have the same relative molecular mass (Mr), however, they differ in their surface areas.
• Straight chain pentane has largest surface area so its boiling point is the highest among the three isomers of pentane. 2-Methylbutane (Isopentane) has smaller surface area than n-pentane but larger than 2,2-Dimethylpropane (Neopentane).
• The forces among 2-Methylbutane molecules are intermediate and the least strong forces occur in 2,2-Dimethylpropane.

The isomers of pentane alone with their boiling points are given in figure.

Hydrogen Bonding

Definition The forces of attraction between partially positively charged hydrogen atom and highly electronegative elements (F, O, N) are called hydrogen bonding.

Conditions of hydrogen bond A hydrogen bond is formed when the following conditions are fulfilled:

(i) The hydrogen atom is connected to a highly electronegative atom, such as F, O, or N

(ii) The electronegative atom must have a lone pair of electrons on it.

Explanation

• Hydrogen bond is a special type of dipole-dipole force.
• It is the strongest force among intermolecular forces, but is weaker than all the major types of bond, i.e. ionic, metallic, and covalent bonds.
• The covalent bond between a highly electronegative atom and hydrogen is highly polarized. It means the hydrogen atom carries a sufficient δ⁺ charge and the electronegative atom δ⁻ charge.
• The presence of lone pair on the electronegative atom allows the partially positively charged hydrogen atom to make a bond with F, O, or N called a hydrogen bond.
• This bond is represented by a dotted line (....) and its strength is one tenth of an ordinary covalent bond.

Examples of hydrogen bonding

Liquid water (H₂O)

• A water molecule contains two hydrogen atoms and two lone pairs on oxygen.
• It can form two bonds on average.
• Water molecules are extensively hydrogen bonded with one another because of two hydrogen bond per molecules.
• This unique feature of water is responsible for its characteristics properties, like high boiling point of water (100°C).
• The water molecules in liquid form are arranged three dimensionally as shown in figure.
• Hydrogen bonding acts as a bridge between two electronegative oxygen atoms.

Ammonia (NH₃)

• Ammonia can form only one hydrogen bond per molecule despite having three hydrogen atoms. Because of availability of only one lone pair on the N atom as shown in figure.
• The hydrogen bonding in liquid ammonia is much weaker than water So, its boiling point is much lower (-33°C).

Hydrogen fluoride (HF)

HF is even more distinct than water and ammonia.

• It has three lone pairs on the F atom which allows it to form three H-bonds.
• However, it has only one H atom and due to this reason, it is restricted to form only one bond again.
• Its hydrogen bond is (B.P =19.9°C) is much higher than ammonia (B.P = -33.5°C).
• The molecules of HF join with one other in a zig-zag manner in solid form due to the presence of the hydrogen bonds.

Hydrogen halides

• The exceptional low acidic strength of HF molecule as compared to HCl, HBr and HI is attributed to this strong hydrogen bonding, because the partial positively charged hydrogen is entrapped between two highly electronegative atoms.

Influence of Hydrogen bonding on the properties of water

(i) Structure and low density of ice

• The molecules of water have tetrahedral structure.
• Two lone pairs of electrons on oxygen atom occupy two corners of the tetrahedron.
• When the temperature of water is decreased and ice is formed, the molecules become more regular and this regularity extends throughout the whole structure.
• Empty spaces are created in the extensively hydrogen bonded structure as shown in the following diagram figure.
• When water freezes, it occupies 9% more space and its density decreases.
• Due to decrease in density ice floats on water.
• The fish and other living creatures survive under the frozen lakes and oceans, due to this so called anomalous behavior of ice.
• The ice blanket covers the water surface due to its low density and the water beneath is insulated from the environment.

(ii) High heat capacity

Water has a high specific heat capacity. This is due to its unique molecular structure which allows strong hydrogen bonds.

(iii) Anomalous heat of vaporization and boiling point

The general trend of enthalpy change of vaporization of group 16 hydrides is a regular increase from H₂S to H₂Po figure. This is because the molecules of these compounds are bound by weak London dispersion forces.

• The van der Waals forces become stronger, with the increasing atomic size down the group, from S to Po.
• If water had followed this trend, the heat of vaporization of water would be the lowest among the group members.
• Water has the highest enthalpy change of vaporization (41 kJmol⁻¹) in the group.
• Water is exceptionally hydrogen bonded, which makes it difficult to break forces among water molecules and vaporize it.
• The boiling point of water is also remarkably higher than the rest of group 16 hydrides.
• The energy required to break the water molecules apart is very high due to extensive hydrogen bonding in it. It needs high temperatures to provide so much energy as to boil water into its vapor.

(iv) Surface tension and viscosity

Surface tension The surface tension and viscosity of water are also very high.
• Surface tension is a downward pull of water molecules at the surface, making the water surface stretched and strained.
• The downward pull is due to the attraction of inner molecules through hydrogen bonds.

Viscosity Viscosity is the resistance by a liquid to its flow.
• Water has higher viscosity than many of liquids including lower hydrocarbons and alcohols (hexane and ethanol).
• The high viscosity of water is also due to strong hydrogen bonds among the molecules. Alcohols can make hydrogen bonds, but the strength and extent of hydrogen bonding is far less than that in water.
• Hydrocarbons lack the ability to make the hydrogen bond.

Surface tension of liquids

Definition Surface tension is property of liquids that describes the force acting along the surface of a liquid, causing it to behave like an elastic sheet.

Explanation

(iv) Force arises due to the intermolecular forces between the molecules at the surface.

(v) Molecules at the surface experience a net force inward due to intermolecular forces, leading them to "sink" on the surface.

(vi) Liquids tend to minimize their surface area due to surface tension, which explains that the droplets form spheres.

(vii) Surface tension influences various phenomena, such as, capillary action, wetting of surfaces and formation of droplets and bubbles.

Factors affecting Surface tension

(i) Temperature: Surface tension typically decreases with increasing temperature.

(ii) Intermolecular forces: Stronger intermolecular forces lead to higher surface tension.

Table 5.2 Surface tension of some liquids at 20°C in Nm⁻¹

Sr. No. | Liquid | Surface tension (10⁻³ Nm⁻¹)
1. | Water H₂O | 72.75
2. | Ethyl alcohol C₃CH₂OH | 22.75
3. | Methyl alcohol CH₃OH | 22.61
4. | Acetone (CH₃)₂CO | 23.70
5. | Benzene C₆H₆ | 28.85
6. | Carbon tetrachloride CCl₄ | 26.95

Viscosity of liquids

Definition Viscosity is the resistance by a liquid to its flow.

Explanation

• The property that describe thickness or stickiness of a liquid.
• Liquid with higher viscosity are thicker and more resistance to flow.
• Liquid with low viscosity are thinner and less resistant to flow.
• Honey is more viscous than H₂O.
• Stronger intermolecular forces results in high viscosity.

Units of viscosity In international system (SI), unit of viscosity is in Kg m⁻¹ s⁻¹. In CGS system, the unit of viscosity is poise (P).

Ans. Evaporation
Definition: Evaporation is the spontaneous conversion of a liquid into vapour at any temperature.
Evaporation causes cooling
• When high energy molecules leave the liquid and low energy molecules are left behind, the temperature of the liquid falls.
• The heat moves from the surrounding to the liquid and then the temperature of surrounding also falls.
• This phenomenon helps to understand the evaporation causes cooling.
One can feel cool after bath
• A person after bath feels a sense of cooling due to evaporation of water from his body when exposed to air.
• The molecules of H₂O take away the energy of body.
• Earthenware vessels keep water cooler under the same phenomenon.
• As a result earthenware vessels are porous.
Earthenware vessels keep water cool
• The water molecules having sufficient kinetic energy to overcome inter molecular forces come out of these pores and evaporate.
• This process of evaporation keeps on taking place and energy required for this process comes from the liquid.
• The average kinetic energy of remaining water molecules decreases which results in decreases in temperature of the liquid. As a result earthenware vessels keep water cool.
Vapour pressure
Definition: The pressure exerted by the liquid vapour in equilibrium with its liquid at a given temperature is called vapour pressure.
Explanation
• The molecules of a liquid which leave the open surface are mixed up with air above the liquid.
• If the vessel is open these molecules go on leaving the surface of liquid.
• If we close the system, the molecules of liquid start gathering above the surface.
• These molecules collide with the walls of the container, and also with the surface of the liquid as well.
• There are chances that these molecules are recaptured by the surface of liquid. This process is called condensation.
• The two processes, i.e., evaporation and condensation continue till a stage reaches when the rate of evaporation becomes equal to the rate of condensation figure.
Liquid ⇌ (evaporation/condensation) Vapour
Factors affecting vapour pressure
• The magnitude of vapour pressure does not depend upon the amount of liquid in the container or the volume of container.
• It does not depend on surface area of a liquid. The larger surface area presents a larger target for returning the molecules, so the rate of condensation also increases.
• The vapour pressure decreases with increase in intermolecular forces of compounds and vice versa.
• The vapour pressure of liquid increases by increasing temperature.
Fig: Attainment of equilibrium when the evaporation of liquid is carried in an evacuated closed vessel (a) Initial state, with evacuated space above the liquid, (b) intermediate state, and (c) equilibrium state, when the rate of evaporation is equal to the rate of condensation.

Ans. Boiling Point
Definition: The temperature at which the vapour pressure of a liquid becomes equal to its external pressure is called boiling point.
Explanation
• When a liquid is heated, its vapour pressure goes on increasing.
• A stage reaches when the vapour pressure of the liquid becomes equal to the atmospheric or external pressure. At this temperature the liquid starts boiling.
• The reason for this is that the bubbles of vapour which are formed in the interior of the liquid have greater internal pressure than atmospheric pressure on the surface of liquid.
• This makes the bubbles to come out of the liquid and burst at the surface. A constant of bubbles comes out at the boiling point of the liquid.
Graph of boiling point
• When a liquid is heated, the kinetic energy of its molecules increases. This causes the increase of the temperature.
• At the boiling point, the kinetic energy of the molecules becomes maximum.
• Any further heating at this stage will not increase the temperature.
• The further heating is utilized to break the intermolecular forces and convert the liquid into its vapour.
The boiling points of some commonly available liquids at one atmospheric pressure are shown in the table 5.3.
• The variation of vapour pressure of water, ethyl alcohol, ethylene glycol and diethyl ether with temperature is shown in figure.
• The graphs show that the liquids reach upto their boiling points at when their vapour pressures are equal to 760 torr at sea level.
Factors affecting boiling points of liquids
The boiling point of a liquid is affected by the factors given below.
(i) Strength of intermolecular forces
Stronger the intermolecular forces, lower will be the vapour pressure and higher will be the boiling point. Higher boiling point of H₂O indicates stronger intermolecular forces than that of ethanol and methanol, table 5.3.
Table 5.3: Boiling points of some common liquids at 760 torr (1 atm)
Liquids | Formula | B.P (°C) | Liquids | Formula | B.P (°C)
Acetic acid | CH₃COOH | 118.50 | Carbon tetrachloride | CCl₄ | 76.50
Acetone | CH₃COCH₃ | 56.00 | Ethanol | C₂H₅OH | 78.30
Benzene | C₆H₆ | 80.15 | Naphthalene | C₁₀H₈ | 218.00
Carbon disulphide | CS₂ | 46.30 | Water | H₂O | 100.00
(ii) External pressure
• When vapour pressure of a liquid becomes equal to the external pressure then the liquid boils.
• It means that when external pressure is changed, its boiling point will also change.
• When the external pressure is high the liquid requires greater amount of heat to equalize its vapour pressure to external pressure, so, boiling point is raised.
• At a lower external pressure, a liquid absorbs less amount of heat and it boils at a lower temperature.
• Water shows boiling point of 120°C at 1489 torr pressure and boils at 25°C at 23.7 torr.
• Water boils at 98°C at Murree hills due to external pressure of 700 torr while at the top of Mount Everest water boils at only 69°C at 323 torr.
• Increase the external pressure artificially on the surface of boiling water by using a pressure cooker.
• Pressure cooker is a closed container. The vapour of water formed is not allowed to escape.
• It exerts more pressure on the water surface in the cooker and the boiling temperature increases.
• As more heat is absorbed in water, so food is cooked quickly under increased pressure.
Fig: Vapour pressures (torr) of four common liquids shown as a function of temperature

Ans. Energetic of phase changes
Physical change of matter is always accompanied by an energy change. The change in energy is the quantitative measurement of the difference in the strength of intermolecular forces.
Molar heat of fusion (ΔHf)
Definition: The amount of heat absorbed by one mole of a solid to melt it into the liquid at its melting point at 1 atmospheric pressure is called molar heat of fusion.
Example:
H₂O₍ₛ₎ ⟶H₂O₍ₗ₎ ΔHf =4.6kJmol⁻¹
Explanation
• Heat of fusion plays a significant role in the study of glaciers and ice sheets.
• Heat of fusion determines the energy required for ice to melt.
• As water has high heat of fusion, ice requires a lot of energy to melt.
• Due to the low melting, glaciers gather mass which limits their flow.
• It causes accumulation of ice on the earth poles, known as polar ice caps.
• Polar ice caps are important in regulating the earth's temperature by reflecting the sunlight.
• These ice caps also help regulate sea levels by storing fresh water ice.
• Both the ice caps and icebergs are stable due to the high heat of fusion of water.
• If the heat of fusion of water were low, melting of glaciers and ice caps would result in drastic changes in the sea and earth's ecosystems.
Molar heat of vaporization (ΔHv)
Definition: The amount of heat absorbed by one mole of a liquid to convert it into one mole of vapour at its boiling point at 1 atmospheric pressure is called molar heat of vaporization.
Example: When one mole H₂O is converted to vapour at 100°C, then heat absorbed is 40.6 kJ mol⁻¹
H₂O₍ₗ₎ ⟶H₂O₍ₐ₎ ΔHᵥ =40.6kJmol⁻¹
Dependence on inter molecular forces
• The molar heats of vaporization depends upon the strength of intermolecular forces.
• NH₃ and HCl only need 21.7 and 15.6 kJ mol⁻¹ respectively to become vapour at their boiling points. The difference is due to the stronger intermolecular forces in water.
Effect of molar heat of vaporization and fusion on matter particles
• The molar heat of fusion and vaporization affect the particles that make up matter by providing them with enough energy.
• The particle move away from one another and change the state of the substances from solids to liquids to gases respectively.
• Particles gain more freedom to move and rotate in the liquid and gas phase.

Ans. Solids
Definition: Solids are those substances which are rigid, hard, have definite shape and definite volume.
The atoms, ions and molecules that make up a solid are closely packed. They are held together by strong cohesive forces. The constituent atoms, ions or molecules of solids cannot move at random.
General properties of solids
(i) Compression of solids
• The atoms, molecules or ions of a solid substance are closely packed.
• The particles of solids cannot move closer to each other unlike gases.
• It is true for metallic solids that their atoms are spherically symmetrical.
• When such particles are arranged, they can give birth to hexagonal and cubic close-packed structures.
• Outer boundaries of constituent particles touch each other but due to their spherical shapes, they cannot occupy more than 74% space.
• It means the compression of solids is not possible.
(ii) Expansion of solids
• The expansion is the property in which spaces between constituent particles are increased.
• In case of solids the forces of attractions are so strong that increase of temperature hardly affects their relative positions.
• The particles in a solid have vibrational motions about their mean positions.
• Increase in temperature of solids hardly increases their volume.
• There are parameters like coefficient of linear and cubic expansions of various solids, but these are negligible as compared to liquids and gases.
(iii) Motion of particles in solids
• The constituent particles of a solid do not undergo translatory motion, and neither rotational ones.
• They only vibrate about their mean positions.
• These vibrations become more intense at higher temperature.
(iv) Inter-particle space in solids
• The interparticle spaces in solids are far less than liquids.
• The forces among their particles are stronger.
• Due to closely packed atoms, molecules and ions, solids are mostly hard, have high melting points and high stability.
(v) Inter-particle forces in solids
• The particles (atoms, ions, molecules) are held together by ionic, covalent, metallic or van der Waals forces.
• These forces are strong enough to fix the particles at their places thus allowing these particles to just vibrate about their mean positions.
(vi) Kinetic energy based on KMT
• Kinetic energy is due to the motion of constituent particles of a solid. Solid particles have only vibratory motion and they do not have translational or rotational motion.
• The only kinetic energy that solids possess is vibrational kinetic energy.

Ans. Properties of crystalline solids
(i) Geometrical shape
• All the crystalline solids have a definite, distinctive geometrical shape due to definite and orderly arrangement of atoms, ions or molecules in three-dimensional space.
• For a given crystal, the interfacial angles, at which the surfaces intersect, are always the same no matter in which shape they are grown.
• The faces and angles remain characteristic even when the material is ground to a fine powder.
(ii) Melting points
Crystalline solids have sharp melting points. They can be identified from their definite melting points.
(iii) Cleavage planes
• Whenever the crystalline solids are broken they do so along definite planes.
• These planes are called the cleavage planes and they are inclined to one another at a particular angle for a give crystalline solid.
• The value of this angle varies from one solid to another solid.
(iv) Growing of a crystal
• When we have a saturated or super saturated solution of a crystalline material in a suitable solvent, it can give us same types of crystals by arranging atoms, ions or molecules.
• This happens through the process called growth of crystals.
• This can be done by slow evaporation of the solvent or by seeding process from saturated solution.
Example: The solubility of sodium thiosulphate (Na₂S₂O₃) in water at 100°C is 231g/100 cm³. At room temperature solubility is 50g/100 cm³. In case, we have saturated solution of Na₂S₂O₃ at 100°C, on cooling slowly no growth of crystal happens. It means that super saturated solution is not in equilibrium with solid substance. Now, if small crystal of Na₂S₂O₃ is added to super saturated solution, the crystallization happens rapidly.
(v) Habit of a crystal
• The shape of a crystal in which it usually grows is called habit of a crystal.
• If the conditions for growing a crystal are maintained, then the shape of the crystal always remains the same.
• If the conditions are changed the shape of the crystal may change.
Example: A cubic crystal of NaCl becomes needle like when 10% urea is present in its solution as an impurity.

Ans. Crystalline solids
Definition: The solids which have definite regular and three dimensional geometric shapes are called crystalline solids.
Example: Diamond, sodium chloride, ice, etc.
Crystal Lattice: The regular arrangement of ions, atoms, or molecules in three dimensional space is called the crystal lattice.
Amorphous solids
Amorphous solids, contrary to the crystalline solids, do not possess regular three dimensional geometrical shapes.
Example: The amorphous solids are glass, wood amorphous sulphur (plastic sulphur), charcoal, coal, coke, etc.
Properties of amorphous solids
(i) The amorphous solids can have small regions where orderly arrangement of particles is found, but they do not have long range of regularity. The regions where orderly arrangement of particles is present are known as crystallites.
(ii) Amorphous solids like glass melts over a wide temperature range.
(iii) They can be molded and blown to form different shapes.
(iv) Amorphous solids do not have definite value of heat of fusion.
Difference between crystalline and amorphous solids are given following in the table.
Table 5.4: Comparison of amorphous and crystalline solids
Property | Crystalline solids | Amorphous solids
Geometry of solid | Crystalline solids show characteristic geometrical shapes. | Amorphous solids generally appear in lump or in a fine powdered form.
Melting of solid | Crystalline solids melt sharply at their melting points. | Amorphous solids do not have sharp melting points and they melt over a range of temperature.
Directional character of the properties | Crystalline solids are anisotropic in nature. It means that their properties depend upon the direction along which the measurements are made. | Amorphous solids are isotropic. Their properties do not depend upon the direction of measurement.
Order of particles in solid | Crystalline solids have long-range order. | In amorphous solids, long-range order is absent.

Ans. Liquid crystals
Definition: The liquid crystalline state exists between two temperatures that is melting temperature and clearing temperature.
Example: The cholestryl benzoate forms liquid crystals between 145°C and 179°C. Its melting temperature is 145°C and clearing temperature is 179°C.
Explanation
• The distinction between liquids and solids is clear-cut.
• The phase transition between them is always sharply defined.
• At certain temperature, many substances exist in a phase that is neither fully liquid nor fully solid.
• The molecules in these substances can move around, as in viscous liquids, but have a restricted range of motion, as in solids. These substances are called liquid crystals.
Crystalline solid → Cloudy liquid (liquid crystal) → Clear liquid
• In most liquid crystals, the molecules have a rigid, rod-like shape with a length four to eight times greater than their diameter.
• When packed together, the molecules tend to orient with their long axes roughly parallel, like logs in a stack of firewood.
• Individual molecules can migrate through the fluid.
• They can spin around their long axis, but they can't rotate end over end.
General properties of liquid crystals
The general properties of liquid crystals are as follows:
(i) Parallel ordered arrangement. (ii) Elongated, rod-like and linear.
(iii) Flow like liquids (iv) Show viscosity like liquids
(v) Somewhat rigid (vi) Always anisotropic
Uses of liquid crystals in daily life: The main application area of liquid crystals is in electro-optic devices. These are electrically controlled devices that modulate light in a desired way. Liquid crystals have many uses in daily life.
(i) In diagnostics
• Special liquid crystal devices can be used to diagnose the tumors and infections in the human body.
• This is because often physical problems, such as tumors, have a different temperature than the surrounding tissue.
• When cholesteric liquid crystals are applied to the breast, a tumor is located because of being warmer than the other parts.
• This technique is successfully applied to diagnose breast cancer in the early stage.
(ii) Characteristic higher temperature determination
• Liquid crystal temperature sensors can also be used to find faulty connections on a circuit board by detecting the characteristic higher temperature.
• As temperature sensors, they can be used in thermometers.
(iii) Liquid crystal displays (LCDs)
• The most common application of liquid crystal technology is liquid crystal displays (LCDs).
• Liquid crystal screens are oscillograph and TV, laptops, cell phone displays.
• Liquid crystals are used in TV displays, computer screens, calculators and watches, etc.
• As new properties and types of liquid crystals are investigated and researched, these materials are sure to gain increasing importance in industrial and scientific applications.