Introduction
The classification of matter into three different states, namely solid, liquid and gaseous state is termed as the physical classification of matter. Most properties of solid, liquid and gases that can be observed with our sense organs are called as 'macroscopic' properties. The description of the behaviour of the three states of matter in terms of atomic theory is called 'microscopic' description of matter. From the study of the observable properties of different states of matter one can understand the microscopic nature of matter in terms of the behaviour of constituent particles.
The Gaseous State
Of the three states of matter, the gaseous state is the simplest and shows greatest uniformity in behaviour. Gases show almost similar behaviour irrespective of their chemical nature.
Gas Laws - Boyle's Law
Boyle's Law
Robert Boyle proposed this law in the year 1662, giving the relationship between pressure and volume of given mass of a gas at constant temperature. This law states that volume (V) of a given mass of gas is inversely proportional to the pressure (P) at constant temperature.
Charle's Law
Charles formulated this law in 1787 giving the relationship between volume and temperature of a gas. This law stated that at constant pressure, the volume of a given mass of gas increases or decreases by 1/273 of its volume at 0oC for every one degree rise or fall.
Avogadro's Law
The relationship between the volume of a gas to the number of molecules at constant temperature and pressure is known as Avogadro's law. It states that equal volumes of all gases under similar conditions of temperature and pressure contain equal number of molecules. 22.4 litres of any gas at STP contains 6.023 x 1023 number of molecules irrespective of its nature. Therefore, the volume of a gas is directly proportional to the number of molecules N.
Dalton's Law of Partial Pressures
Dalton proposed this law on the pressure exerted by a mixture of non-reacting gases in an enclosed vessel. The law of partial pressure states that the total pressure exerted by a mixture of two or more non-reacting gases in a definite volume is equal to the sum of the individual pressures, which each gas would exert if it occupies the same volume at a constant temperature. If p1, p2, p3 are the individual partial pressures of the known gases, then the total pressure 'P' of the mixture of gases at the same temperature and pressure is given by the relation:
P = p1 + p2 + p3 + …….
Graham's Law of Diffusion
Gases have the tendency to spontaneously intermix and form a homogenous mixture without the help of external agency. This is due to the presence of large amount of empty space between the gas molecules that makes their movement rapid into each other. The gases move from a region of higher concentration to a region of lower concentration until the mixture attains uniform concentration.
Ideal Gas Equation
By combining Boyle's and Charles' laws, an equation can be derived that gives the simultaneous effect of the changes of pressure and temperature on the volume of the gas. This is known as combined Ideal Gas Equation.
Ideal and Real Gases
An ideal gas is one, which obeys the general gas equation of PV = nRT and other gas laws at all temperatures and pressures. A real gas, does not obey the general gas equation and other gas laws at all conditions of temperature and pressure.
Kinetic Molecular Theory of Gases
In order to explain the observed behaviour of gases, a model was proposed based on the molecular and kinetic concept of gas molecules. This model takes into account the particulate nature of matter and the constant movement of particles.
Liquefaction of Gases
There are large empty spaces (voids) separating the tiny molecules of gases from one another. Each molecule enjoys an almost independent existence. Molecules are in a state of continuous rapid motion and have negligible attractive forces between them due to wide separation. This is particularly so, when temperature is high and pressure is low. When the temperature of the gas is lowered, both the volume of the gas and the kinetic energy of the molecules decrease. The molecular motion becomes slow and molecules become sluggish. The progressive decrease of temperature brings the molecules closer and closer because they are unable to resist the attractive force that starts operating between them. Ultimately, at sufficiently low temperature, the voids between the molecules become less than 10-5cm and the gas changes into liquid state.
Relationship between Critical Constant and Van der Waal's Constants
The relationship between critical constants of the gases and their Van der Waal constants is as follows:
(i) Vc = 3b
(ii) Pc = a/27b2
(iii) Tc = 8a/27Rb
(iv) The critical compressibility factor Zc is given by,
Maxwell's Distribution of Molecular Speeds
On plotting a fraction
of molecules having different speeds against the speeds of the molecules (along x-axis) a curve known as Maxwell's distribution curve is obtained.
The Liquid State
The liquid state lies between the gaseous and the solid state. Liquids are neither completely disordered nor completely ordered. The cohesive forces between the liquid particles are strong enough to keep them together but these particles do not occupy fixed lattice sites and are relatively more free as compared to the particles in solids.
Properties of Liquids - I
Liquids have a definite volume under given temperature and pressure conditions. Though they take the shape of the container, they maintain their volume. The intermolecular forces in liquids are strong and therefore, they do not expand to occupy all the space available (as gases do). A given mass of liquid has a fixed volume. For example, 10 cm3 of water always occupies 10 cm3 whether it is placed in a beaker, a conical flask or a large round bottom flask.
Properties of Liquids - II
Vapour pressure measures the tendency for the molecules to escape from liquid to the gas phase. At lower temperatures the vapour pressure of a liquid is much lower than the pressure on the surface of the liquid. When the temperature of the liquid is gradually increased, its vapour pressure also increases. Ultimately a stage is reached when the vapour pressure of the liquid equals the pressure of the air above it. At this point, molecules and vapours formed within the liquid can easily rise through the liquid in the form of bubbles and escape into the air. This phenomenon is known as boiling and the temperature at which this occurs is known as boiling point.
The Solid State
Solids are substances having definite volume and definite shape. In terms of kinetic molecular model, in solids there is a regular order in the arrangement of the constituting particles (atoms, molecules or ions) i.e. the particles constituting the solids occupy fixed positions These particles are held together by fairly strong forces.
Classification of Solids
The substances whose constituents are arranged in definite orderly arrangements are called crystalline solids. Many naturally occurring solid substances occur in the crystalline form. Some common examples of crystalline solids are sodium chloride, sulphur, diamond, sugar, etc.
Crystalline Solids
Crystalline substances have a definite rigid shape. The shape and size of crystals (even of the same materials) differs depending upon the conditions under which they are grown. Crystals of a given substance are bound by plane surfaces called faces. The angle between any two faces is called interfacial angle. But the angles between the faces of a given form always remains same. This important characteristic feature of a given crystalline substance is known as law of constancy of interfacial angles.
