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Basic Concepts of Thermodynamics
2.1. Introduction to kinetic theory of gases. 2.2. Definition of thermodynamics.
2.3. Thermodynamic systems—system, boundary and surroundings—closed system—open
system—isolated system—adiabatic system—homogeneous system—heterogeneous
system. 2.4. Macroscopic and microscopic points of view. 2.5. Pure substance.
2.6. Thermodynamic equilibrium. 2.7. Properties of systems. 2.8 State. 2.9. Process.
2.10. Cycle. 2.11. Point function. 2.12. Path function. 2.13. Temperature. 2.14. Zeroth law of
thermodynamics. 2.15. The thermometer and thermometric property—introduction—
measurement of temperature—the international practical temperature scale—ideal gas.
2.16. Pressure—definition of pressure—unit for pressure—types of pressure measurement
devices—mechanical-type instruments—liquid manometers—important types of pressure
gauges. 2.17. Specific volume. 2.18. Reversible and irreversible processes. 2.19. Energy,
work and heat—energy—work and heat. 2.20. Reversible work—Highlights—Objective Type
Questions—Theoretical Questions— Unsolved Examples.2.1. Introduction to Kinetic Theory of Gases
The kinetic theory of gases deals with the behaviour of molecules constituting the gas.
According to this theory, the molecules of all gases are in continuous motion. As a result of this
they possess kinetic energy which is transferred from molecule to molecule during their collision.
The energy so transferred produces a change in the velocity of individual molecules.
The complete phenomenon of molecular behaviour is quite complex. The assumptions are
therefore made to simplify the application of theory of an ideal gas.
Assumptions :
- The molecules of gases are assumed to be rigid, perfectly elastic solid spheres, identical
in all respects such as mass, form etc. - The mean distance between molecules is very large compared to their own dimensions.
- The molecules are in state of random motion moving in all directions with all possible
velocities and gas is said to be in state of molecular chaos. - The collisions between the molecules are perfectly elastic and there are no intermolecu-
lar forces of attraction or repulsion. This means that energy of gas is all kinetic. - The number of molecules in a small volume is very large.
- The time spent in collision is negligible, compared to the time during which the mol-
ecules are moving independently. - Between collisions, the molecules move in a straight line with uniform velocity because
of frictionless motion between molecules. The distance between two collisions is called
‘free path’ of the molecule, the average distance travelled by a molecule between succes-
sive collision is known as ‘mean free path’. - The volume of molecule is so small that it is negligible compared to total volume of the gas.
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