States and Phases of Matter

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

Properties of gases (book 5.1): (i) Gases do not have a definite shape and volume — the volume of a gas is the volume of its container. (ii) Ideal gases have zero particle volume. (iii) The molecules of gases are widely separated from one another and most of the volume of the gas is empty space (nearly 99.9%) — that is why gases can be compressed easily. When sudden expansion of gases occurs, cooling takes place; it is called the Joule-Thomson effect. (iv) Pressure of a gas is due to the collisions of gas molecules with the walls of the container. (v) There are negligible intermolecular forces in ideal gases.

📌 Book note: kinetic molecular theory of gases helps us understand these properties — widely spaced, fast-moving particles with negligible attractions explain why a gas fills its container, presses on its walls and squeezes into a smaller volume.

Gas laws. When external temperature and pressure change, the volume of a given quantity of gas is affected. The gas laws describe the relationships between volume of a given amount of gas and the prevailing conditions of temperature and pressure. Boyle's law: V ∝ 1/P (when 'n' and 'T' are held constant). Charles' law: V ∝ T (when 'n' and 'P' are held constant). Avogadro's law: V ∝ n (when P and T are held constant) — the volume of a given gas at constant temperature and pressure is directly proportional to the number of moles.

Combining the three: V ∝ nT/P, so V = constant × nT/P. The constant is 'R', the ideal gas constant, whose value is 0.0821 atm dm³ K⁻¹ mol⁻¹. Hence V = R × nT/P, i.e. PV = nRT — the ideal gas equation.

Calculation of relative molecular mass (M_r) of a gas. If P, T, V and mass in grams of a gas are known, put n = m/M in PV = nRT: PV = mRT/M, so M = mRT/PV. At molecular level the molar mass becomes the relative molecular mass, so M_r = mRT/PV.

Sample Problem 5.1 (book): The volume of 134 g of a gas at −73 °C under 10 atm pressure is 5 dm³. Determine the relative molecular mass. Solution: M = mRT/PV = (134 g)(0.0821 atm dm³ mol⁻¹ K⁻¹)(200 K) / (10 atm)(5 dm³) = 44 g mol⁻¹. Therefore the relative molecular mass of the gas is 44 amu. (Note: −73 °C = 200 K.)

Properties of liquids (book 5.3): (i) Liquids also diffuse like gases, however the rate of diffusion in liquids is much lower than that in gases. (ii) There are little spaces among liquid molecules because of relatively stronger intermolecular forces. (iii) Liquids are 10⁵ times less compressible than gases, but about 10 times more compressible than solids. The book's Figure 5.1 shows it: doubling pressure from 1 atm to 2 atm halves the volume of a gas, but changes the volume of a liquid (water) by only about 0.0045%. (iv) 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 evidence for the free motion of liquid molecules. Similarly, liquids can flow because of the free movement of their molecules. (v) Liquids expand when their temperature is increased. However, this expansion is negligible compared to that in gases. (vi) According to kinetic molecular theory, molecules of a liquid are in constant motion, therefore they possess a certain amount of kinetic energy at any temperature. The kinetic energy of the molecules may increase or decrease by the rise or fall in temperature respectively. For example, a liquid is converted into a solid by decreasing its kinetic energy (on cooling) and converted into vapour by increasing kinetic energy (on heating).

Three types of intermolecular forces among the molecules of a liquid are: (i) Instantaneous dipole-induced dipole forces (id–id), (ii) Permanent dipole-permanent dipole forces (pd–pd), (iii) Hydrogen bonding. The intermolecular forces among the molecules of a liquid are reasonably stronger than in gases.

How id–id forces arise. Bromine is a liquid under ordinary conditions, but very volatile and gives dense fumes even at room temperature. It shows that its molecules are attracted by weak forces that can be easily broken. This force is actually London dispersion force. When two molecules of bromine come close, the electron cloud of one molecule influences the moving electrons of the other. The electrons of one molecule are pushed away while its nucleus is attracted by the electrons of the nearby molecule. As a result, this molecule becomes a dipole at any instant, called an instantaneous dipole. This molecule then disturbs the electron density of the molecule next to it and induces a polarity in it. A temporary dipole created due to the instantaneous dipole is called the induced dipole. A force of attraction develops between these two dipoles and holds them together temporarily. The momentary force of attraction between an instantaneous dipole and an induced dipole is called instantaneous dipole-induced dipole force. These are the only forces that exist among the molecules of non-polar compounds.

Strength depends on: (i) Molecular mass and size: when the molecular mass (M_r) of a compound is greater, the id–id forces are stronger and vice versa — in a larger atom or molecule the dispersion of electron cloud and polarisation is relatively easier. This factor operates in the halogen and noble gas families: the physical state changes in the halogen family from top to bottom due to the increase in size and polarizability. The first two members from the halogen family (F₂ and Cl₂) are gases, but bromine is a liquid and iodine is a solid (b.p. 458 K). The boiling points of noble gases in group 18 also increase from helium (He) to xenon (Xe) (book Table 5.1: F₂ M_r 38, b.p. 85 K; Cl₂ 71, 239 K; Br₂ 159, 333 K; I₂ 253.8, 458 K; He 4.003, 4.22 K; Ne 20.18, 27.1 K; Ar 39.95, 87.3 K; Kr 83.8, 120 K; Xe 131.3, 165 K). (ii) Surface area (shape of molecule): a molecule with larger surface area has higher number of contacts for a nearby molecule, so the forces among the molecules of a compound having larger surface area are stronger. Compare saturated hydrocarbons: molecules with larger chain length experience stronger attractive forces, because longer molecules have more places along the length where they can be attracted to other molecules. Three isomers of pentane have the same relative molecular mass, but they differ in their surface areas. Straight-chain pentane has the largest surface area and 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), so the forces among 2-methylbutane molecules are intermediate and the least strong forces occur in 2,2-dimethylpropane. Their boiling points (book Figure 5.3): pentane (n-pentane) 36.2 °C, 2-methylbutane 28 °C, 2,2-dimethylpropane 9.5 °C.

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 force. Whenever the HCl molecules are close to each other (as in liquid state), they tend to line up as in the book's Figure 5.4: δ⁺ H–Cl δ⁻ ··· δ⁺ H–Cl δ⁻ ··· δ⁺ H–Cl δ⁻, positive end next to negative end, again and again.

Similar type of forces is present among the molecules of chloroform (CHCl₃). The molecule of chloroform has a positive center at the H atom and the negative center on the end with Cl atoms, so its molecules also attract each other positive-to-negative.

How pd–pd differs from id–id: pd–pd forces exist only between molecules that are permanently polar (they have a built-in δ⁺ and δ⁻ end), while id–id (London) forces arise in all molecules — even non-polar ones — from momentary shifts of the electron cloud. Polar molecules still have London forces too, acting alongside their dipole-dipole forces; hydrogen bonding, covered next, is a special, stronger kind of dipole-dipole force.

What a hydrogen bond is. 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. A hydrogen bond is formed when: (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. The covalent bond between a highly electronegative atom and hydrogen is highly polarized: 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. The number of hydrogen bonds a molecule can form on average depends on: the number of hydrogen atoms bonded to the electronegative atom, and the number of lone pairs present on F, O, or N.

Water, ammonia, hydrogen fluoride. A water molecule contains two hydrogen atoms and two lone pairs on oxygen. Therefore, it can form two bonds on average. This is why water molecules are extensively hydrogen bonded with one another. This unique feature is responsible for its characteristic properties, e.g. high boiling point of water (100 °C). The water molecules in liquid form are arranged three dimensionally. This bonding acts as a bridge between two electronegative oxygen atoms. On the other hand, ammonia can form only one hydrogen bond per molecule despite having three hydrogen atoms. This is due to the availability of only one lone pair on the N atom. The hydrogen bonding in liquid ammonia is much weaker than water. This is why its boiling point is much lower (−33.5 °C). 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. Nevertheless, its hydrogen bond is strong enough compared to NH₃, which owes to its high electronegativity. Thus, its b.p. (19.9 °C) is much higher than ammonia (−33.5 °C).

5.4.4 Influence on 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. This is why, when water freezes, it occupies 9% more space and its density decreases. For this reason, ice floats on water. Due to this so-called anomalous behavior of ice, the fish and other living creatures survive under the frozen lakes and oceans. 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 (4.18 J/g°C). This is due to its unique molecular structure which allows strong hydrogen bonds. Due to this high heat capacity of water, it can absorb and release a large amount of heat with small temperature changes. (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. This is because the molecules of these compounds are bonded by weak London dispersion forces. With the increasing atomic size down the group, from S to Po, the van der Waals forces also become stronger. If water had followed this trend, the heat of vaporization of water would be the lowest among the group members. However, water has the highest enthalpy change of vaporization (41 kJ mol⁻¹) in the group. This is because 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. Simply, the energy required to break the water molecules apart is very high due to the extensive hydrogen bonding in it. It needs high temperatures to provide so much energy to boil water into its vapour. (iv) Surface tension and viscosity: 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. This downward pull is due to the attraction of inner molecules through hydrogen bonds. Viscosity is the resistance by a liquid to its flow. Water has higher viscosity than many of liquids including lower hydrocarbons and alcohols (e.g. 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.

What surface tension is. Surface tension is a property of liquids which depends upon the strength of intermolecular forces. The molecules located at the free surface of the liquid are pulled inward by the molecules below the surface of the liquid. This downward pull tends to contract the surface of the liquid. Molecules at the surface experience a net force inward due to intermolecular forces, leading to a "skin" on the surface. Liquids tend to minimize their surface area due to surface tension, which explains why droplets form spheres. 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. The book's Table 5.2 (surface tension of some liquids at 20 °C, in 10⁻³ N m⁻¹): water H₂O 72.75, ethyl alcohol CH₃CH₂OH 22.75, methyl alcohol CH₃OH 22.61, acetone (CH₃)₂CO 23.70, benzene C₆H₆ 28.85, carbon tetrachloride CCl₄ 26.95. Water's value towers over the rest — the book's own picture caption makes the same point: mosquitoes can float on the surface of water due to high surface tension.

What viscosity is. Viscosity is a measure of a liquid's resistance to flow. It is a property that describes the thickness or stickiness of a liquid. Liquids with higher viscosity are thicker and offer more resistance to flow, while those with lower viscosity are thinner and are less resistant to flow. For example, honey is more viscous than water. Stronger intermolecular forces result in higher viscosity. Viscosity decreases with rise in temperature.

5.6.1 Factors affecting viscosity: (i) An increase of temperature increases the average kinetic energy of molecules. This allows the molecules to overcome the attractive forces. This lowers the viscosity. (ii) The stronger the intermolecular forces, the greater the viscosity of liquids because liquid molecules cannot move around each other freely, so the resistance to flow increases. That is why water (strong hydrogen bonds) is more viscous than hexane or ethanol, and why glycerine — with three OH groups per molecule and extensive hydrogen bonding — is far more viscous than hexane, a non-polar hydrocarbon with only London forces.

5.7 Evaporation. Evaporation is the spontaneous conversion of a liquid into vapour at any temperature. 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 the surrounding also falls. This phenomenon helps to understand that evaporation causes cooling. 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. 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 the energy required for this comes from the liquid. Thus, the average kinetic energy of remaining water molecules decreases which results in decrease in the temperature of the liquid. Hence, earthenware vessels keep water cool.

5.8 Vapour pressure. The molecules of a liquid which leave the open surface are mixed up with the air above the liquid. If the vessel is open these molecules go on leaving the surface of liquid. But 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 (book Figure 5.9: (a) initial state with evacuated space above the liquid, (b) intermediate state, (c) equilibrium state). The pressure exerted by the liquid vapour in equilibrium with its liquid at a given temperature is called vapour pressure: Liquid ⇌ Vapour (evaporation / condensation). 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 of liquids having stronger intermolecular forces are lower and vice versa. With increase in temperature, the vapour pressure of a liquid also increases.

5.9 Boiling point. 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 stream of bubbles comes out at the boiling point of the liquid. 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. This 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 book's Table 5.3 (e.g. acetic acid 118.50 °C, acetone 56.00 °C, benzene 80.15 °C, carbon disulphide 46.30 °C, carbon tetrachloride 76.50 °C, ethanol 78.30 °C, naphthalene 218.00 °C, water 100.00 °C). The book's Figure 5.10 shows the variation of vapour pressure of water, ethyl alcohol, ethylene glycol and diethyl ether with temperature: the liquids reach up to their boiling points when their vapour pressures are equal to 760 torr at sea level (diethyl ether 34.8 °C, ethyl alcohol 78.3 °C, water 100 °C).

5.9.1 Factors affecting boiling points. (1) 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. (2) External pressure: as already explained, that 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. In this way boiling point is raised. Similarly, at a lower external pressure a liquid absorbs less amount of heat and it boils at a lower temperature. For example, water shows B.P 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. We can 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. In this way, 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.

Energy always accompanies a physical change. Whenever matter undergoes a physical change, it is always accompanied by an energy change. This change in energy is the quantitative measurement of the difference in the strength of intermolecular forces.

Molar heat of fusion (ΔH_f). 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: H₂O_(s) → H₂O_(l), ΔH_f = 4.6 kJ mol⁻¹ (the book's printed value). 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. Also, it causes accumulation of ice on the earth's 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 (ΔH_v). 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. The molar heats of vaporization depends upon the strength of intermolecular forces. When one mole H₂O is converted to vapour at 100 °C, then heat absorbed is 40.6 kJ mol⁻¹: H₂O_(l) → H₂O_(g), ΔH_v = 40.6 kJ mol⁻¹. NH₃ and HCl only need 21.7 and 15.6 kJ mol⁻¹ respectively to become vapours at their boiling points. This 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. Due to this, the particle move away from one another and change the state of the substances from solids to liquids and liquids to gases respectively. Particles gain more freedom to move and rotate in the liquid and gas phases.

📌 Book note: the book prints ΔH_f for water as 4.6 kJ mol⁻¹ on page 104 and ΔH_v as 41 kJ mol⁻¹ for group 16 hydrides on page 99 and 40.6 kJ mol⁻¹ here — both values are kept as the book prints them in their own places.

5.11 Solids. 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.

5.11.1 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. So, when such particles are arranged, they can give birth to hexagonal and cubic close-packed structures. Here 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. No doubt, 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 rotational motion. They only vibrate about their mean positions. These vibrations become intense at higher temperature. (iv) Inter-particle space in solids: the interparticle spaces in solids are far less than liquids. This is due to stronger forces among their particles. 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: 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 motion of constituent particles of a solid. Solid particles have only vibratory motion and they do not have translational or rotational motion. Therefore, the only kinetic energy that solids possess is vibrational kinetic energy.

5.12 Types of solids. There are two types of solids on the basis of the way their particles are arranged: (i) Crystalline solids, (ii) Amorphous solids.

5.12.1 Crystalline solids — crystal lattice: the regular arrangement of ions, atoms, or molecules in three dimensional space is called the crystal lattice. The solids which have definite regular and three dimensional geometrical shapes are called crystalline solids. For example, diamond, sodium chloride, ice, etc.

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 given crystalline solid. The value of this angle varies from one solid to another. (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. For example, the solubility of sodium thiosulphate (Na₂S₂O₃) in water at 100 °C is 231 g/100 cm³. At room temperature, the solubility is 50 g/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. For example, a cubic crystal of NaCl becomes needle like when 10% urea is present in its solution as an impurity.

5.12.2 Amorphous solids. Amorphous solids, contrary to the crystalline solids, do not possess regular three dimensional geometrical shapes. The examples of 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. Comparison (book Table 5.4): geometry — crystalline solids show characteristic geometrical shapes, amorphous solids generally appear in lump or fine powder form; melting — crystalline solids melt sharply at their melting points, amorphous solids do not have sharp melting points and melt over a range of temperature; directional character — crystalline solids are anisotropic (properties depend on the direction along which the measurements are made), amorphous solids are isotropic (properties do not depend upon the direction of measurement); order — crystalline solids have long-range order, in amorphous solids long-range order is absent.

What liquid crystals are. We know that the distinction between liquids and solids is clear-cut. Moreover, the phase transition between them is always sharply defined. However, at a 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.

5.13.1 General properties of liquid crystals: liquid crystals (i) have parallel ordered arrangement, (ii) are elongated, rod-like and linear, (iii) flow like liquids, (iv) show viscosity like liquids, (v) show optical properties like crystals, (vi) are somewhat rigid, (vii) are always anisotropic. So in the properties where they resemble solids: parallel ordered arrangement, optical properties like crystals, somewhat rigid, anisotropic; and where they resemble liquids: they flow like liquids and show viscosity like liquids.

5.13.2 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 are used in thermometers. (iii) Liquid crystal displays (LCDs): the most common application of liquid crystal technology is liquid crystal displays (LCDs). LCDs are used in the screens of T.Vs, computers, mobiles, etc.