Basic Concepts of Thermodynamics

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 THERMODYNAMICAL TERMS:

(i) Thermodynamic system:

 A thermodynamical system is an assembly of large number of particles which can be described by thermodynamic variables like pressure (P), volume (V),temperature (T).

The any thing under consideration is called thermodynamic system.

(ii) Surroundings: 

Everything outside the system which can have a direct effect on the system is called surroundings.

The gas cylinder in the kitchen is the thermodynamic system and the relevant part of the kitchen is the surroundings.

(iii) Boundary:

It is defined for Separating it from surroundings is called boundary.

    

 Universe = system + surroundings

Type of system:    

Closed and open system:

 In a closed system, energy may transfer the boundaries of system but mass does not cross the boundary, while in open system, both mass and energy transfer across the boundary of the system.

An isolated system: 

In this type of system neither the mass nor the energy can be exchanged with the surroundings.

(a)open system (b)close system (c)isolated system

Equation of state: The relationship between the pressure,

volume and temperature of the thermodynamical system is called equation of state.

Properties :

A property of a system is any abusable characteristic of the given system various properties of the system depend on the state of the system not on how that state have been reached.

1.Intensive property of a system or those properties whose values does not depend upon the mass of the system. Eg: Pressure,temperature, viscosity etc., while

2.extensive properties depend upon the mass of the system. Eg: Length, volume etc.

Equilibrium: 

A system is said to be in thermodynamic equilibrium when it does not lead to change its properties(macroscopic) and make balance with its surroundings. There,a system in mechanical, thermal and chemical equilibrium is said to be in thermodynamic equilibrium.

THERMODYNAMICAL PROCESSES

Any process may have own equation of state, but each thermodynamical process must obey  PV = nRT.

REVERSIBLE AND IRREVERSIBLE PROCESSES:

Reversible Process:

Any process which can be made to proceed in the reverse direction by variation in its conditions such that any change occurring in any part of the direct process is exactly reversed in the corresponding part of reverse process is called a reversible process.

Examples:

(i) An infinitesimally slow compression and expansion of an ideal gas at constant temperature.

(ii) The process of gradual compression and extension of an elastic spring is approximately reversible.

(iii) A working substance taken along the complete Carnot’s cycle.

(iv) The process of electrolysis is reversible if the resistance offered by the electrolyte is negligibly small. A complete reversible process is an idealised concept as it can never be realised because dissipative forces cannot be completely eliminated.

Irreversible Process:

Any process which cannot be retraced in the reverse direction exactly is called an irreversible process. Most of the processes occurring in the nature are irreversible processes.

Examples:

(i) Diffusion of gases.

(ii) Dissolution of salt in water.

(iii) Rusting of iron.

(iv) Sudden expansion or contraction of a gas.

processes

Quasi-static process:

A process which is carried out in a vary small manner with small gradient. Almost in rest condition or infinite slowness is the characteristics of quasi-statics process .quasi-static process is reversible process.

quasi-static process

BEHAVIOUR OF IDEAL AND REAL GASES

Behaviour of Ideal Gases

The behaviour of ideal gases is based on the following assumptions of kinetic theory of gases :

(1) All the molecules of a gas are identical. The molecules of different gases are different.

(2) The molecules are rigid and perfectly elastic spheres of very small diameter.

(3) Gas molecules occupy very small space. The actual volume occupied by the molecule is very small compared to the total volume of the gas. Therefore volume of the gas is equal to volume of the vessel.

(4) The molecules of gases are in a state of random motion, i.e.,they are constantly moving with all possible velocities lying between zero and infinity in all possible directions.

(5) Normally no force acts between the molecules. Hence they move in straight line with constant speeds.

(6) The molecules collide with one another and also with the walls of the container and change there direction and speed due to collision. These collisions are perfectly elastic i.e., there is no loss of kinetic energy in these collisions.

(7) The molecules do not exert any force of attraction or repulsion on each other except during collision. So, the molecules do not posses any potential energy. Their energy is wholly kinetic.

(8) The collisions are instantaneous i.e., the time spent by a molecule in a collision is very small as compared to the time elapsed between two consecutive collisions.

(9) Though the molecules are constantly moving from one place to another, the average number of molecules per unit volume of the gas remains constant.

(10) The molecules inside the vessel keep on moving continuously in all possible directions, the distribution of molecules in the whole vessel remains uniform.

(11) The mass of a molecule is negligibly small and the speed is very large, there is no effect of gravity on the motion of the molecules. If this effect were there, the density of the gas would have been greater at the bottom of the vessel.

Equation of State or Ideal Gas Equation

The equation which relates the pressure (P), volume (V) and temperature (T) of the given state of an ideal gas is known as ideal gas equation or equation of state.

i.e., PV = nRT

where R = universal gas constant Numerical value of R = 8.31 joule mol–1 kelvin–1 n = no. of moles of gas

Behaviour of Real Gases

The gases actually found in nature are called real gases.

1. Real gases do not obey gas laws

2. These gases do not obey the ideal gas equation

PV = nRT

3. A real gas behaves as ideal gas most closely at low pressure

and high temperature.

4. Equation of state for real gases is given by Vander waal’s

equation

Here a and b are Constant called Vander waal’s constant.

ANALYSIS OF THERMODYNAMIC CYCLES

RELATED TO ENERGY CONVERSION

According to first law of thermodynamics, heat given to a system(DQ) is equal to the sum of increase in its internal energy (DU) and the work done (DW) by the system against the surroundings.

i.e., DQ = DU + DW

Heat (DQ) and work done (DW) are the path functions but internal energy (DU) is the point function.

Cyclic Process and Non-cyclic Process

If a system having gone through a change, returns to its initial state then process is called a cyclic process. If system does not return to its initial state, the process is called non-cyclic process.

Work done in Cyclic Process

Suppose gas expands from initial state A to final state B via the path AXB.The work done in this expansion

         WX= + area A XBCDA

Now gas returns to its initial state B via path BYA. Work done during this compression

       WY= – area BYADCB 

The net work done

   W = WX+ WY

      = area AXBCDA – area BYADCB = + area AXBYA

Thus for a cyclic process

(i) Work done in complete cycle is equal to the area of the loop representing the cycle.

(ii) If the closed loop is traced in the clockwise direction, the expansion curve lies above the compression curve. (WX >WY), the area of loop is positive.

(iii) If the closed loop is traced in the anticlockwise direction, the expansion curve lies below the compression curve (WX <WY), the area of the loop is negative.


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