Page 1 : at rest and occupied fixed positions, then, a gas would have had a fixed shape which, is not observed., , * Particles of a gas move in all possible, directions in straight lines. During their, random motion, they collide with each, other and with the walls of the container., Pressure is exerted by the gas as a result, of collision of the particles with the walls, of the container., , * Collisions of gas molecules are perfectly, elastic. This means that total energy of, molecules before and after the collision, remains same. There may be exchange of, energy between colliding molecules, their, individual energies may change, but the, sum of their energies remains constant., If there were loss of kinetic energy, the, , , , frequently thus exerting more pressure., , Kinetic theory of gases allows us to derive, theoretically, all the gas laws studied in the, previous sections. Calculations and, predictions based on kinetic theory of gases, agree very well with the experimental, observations and thus establish the, correctness of this model., , 5.8 BEHAVIOUR OF REAL GASES:, DEVIATION FROM IDEAL GAS, BEHAVIOUR, , Our theoritical model of gases corresponds, very well with the experimental observations., Difficulty arises when we try to test how far, the relation pV = nRT reproduce actual, pressure-volume-temperature relationship of, gases. To test this point we plot pV vs p plot, , 144 C:\ChemistryX]\Unit-S\Unit-5(4)-Lay-2.pmd_ 14.1.6 (Final), 17.1.6, 24.1.6, , STATES OF MATTER, , of gases because at constant temperature,, pV will be constant (Boyle's law) and pV vs p, graph at all pressures will be a straight line, parallel to x-axis. Fig. 5.8 shows such a plot, constructed from actual data for several gases, al. 273 K., , , , p ——, , Fig. 5.8 Plot of pV vs p for real gas and, ideal gas, , It can be seen easily that at constant, temperature pV vs p plot for real gases is not, a straight line. There is a significant deviation, from ideal behaviour. Two types of curves are, seen.In the curves for dihydrogen and helium,, as the pressure increases the value of pValso, increases. The second type of plot is seen in, the case of other gases like carbon monoxide, and methane. In these plots first. there is a, negative deviation from ideal behaviour, the, pV value decreases with increase in pressure, and reaches to a minimum value, characteristic of a gas. After that pV value, star(s increasing. The curve then crosses the, line for ideal gas and after that shows positive, deviation continuously. It is thus, found that, real gases do not follow ideal gas equation, perfectly under all conditions., , Deviation from ideal behaviour also, becomes apparent when pressure vs volume, plot is drawn. The pressure vs volume plot of, experimental data (real gas) and that, , , , 145, , theoretically calculated from Boyle's law (ideal, gas) should coincide. Fig 5.9 shows these, plots. It is apparent that at very high pressure, the measured volume is more than the, calculated volume. At low pressures,, measured and calculated volumes approach, each other., , , , , Real gas, , Ideal gas, , Pressure ———>, , , , Volume ———>, , Fig. 5.9 Plot of pressure vs volume for real gas, and ideal gas, , It is found that real gases do not follow,, Boyle’s law, Charles law and Avogadro law, perfectly under all conditions. Now two, questions arise., , (i) Why do gases deviate from the ideal, behaviour?, , (ii) What are the conditions under which, gases deviate from ideality?, , We get the answer of the first question if, we look into postulates of kinetic theory once, again. We find that two assumptions of the, kinetic theory do not hold good. These are, (a) There is no force of attraction between the, , molecules of a gas., , (b) Volume of the molecules of a gas is, negligibly small in comparison to the space, occupied by the gas., , If assumption (a) is correct, the gas will, never liquify. However, we know that gases, do liquify when cooled and compressed. Also,, liquids formed are very difficult to compress.