Heat and Thermodynamics
Temperature
Temperature and pressure are macroscopic properties of gases. Thses properties are related to molecular motion, which is a microscopic phenomenon. The theory which correlates between macroscopic properties and microscopic phenomena is called kinetic theory of gases. Kinetic means motion, and in this case motion of gas molecules.
The individual molecules prossess the standard physical properties of mass, momentum and energy. The density of a gas is simply the sum of the mass of the molecules divided by the volume which the gas occupies.
As the gas molecules collide with the walls of a container, the molecules impact momentum to the walls, producing a force that can be measured. The force divided by the area is defined to be the pressure.
The temperature of a gas is a measure of the mean kinetic energy of the gas. The molecules are in constant random motion and there is an energy associated with that motion. The higher the temperature, the greater will be the motion.
In a solid, the location of the molecules relative to each other remains almost constant. But in a gas, the molecules can move around and interact with each other and with their surroundings in different ways, i.e. the molecules of gases are in constant random motion and frequently collide with each other and with the walls of the container in which it is placed.
When the space for expansion of a gas is very large the molecules are far apart and thus, their mutual attraction is almost neglible and rarelyt collide with each other. Such a state of gas is called perfect or ideal gas.
On the other hand, if the space available for expansion of a gas is comparatively small, the molecules are closer and are under the influence of their mutual interaction, they collide with each other freely. Such a state of gas is called the real gas.
The properties of the gases have been studied experimentally and these experiments yielded three important gas laws. These gas laws are the study of the relationship between any two of the physical quantities (pressure, temperature and volume) specifying the state of gas, when the third is kept constant.
Boyle's law
In 1662, the Irish scientist Robert Boyle discovered a basic law of gases that at a constant temperature, the volume and the pressure of a given mass of a gas are inversely proportional.
or PV = K. (a constant)
This relationship is known as Boyle's law. K is a constant of proportionality.
Charles' Law
This law was discoverd by A. C. Charles in 1787. According to this law, at constant pressure, the volume of a given mass of a gas is directly proportional to its absolute temperature.
or 
= K (a constant)
Note: This constant K should not be confused with Kelvin temperature or thermal conductivity. It is a constant of proportionality.
Law of pressure:
At constant volume, the pressure of a given mass of a gas is directly proportional to its absolute temperature. This is called law of pressure.
According to Avogadro's law under the same condition of pressure of temperature equal of volume of different gases contain the same number of molecules.
Perfect gas equation:
Experimental study of gases show that there exists a relation between the pressure, volume and temperature of a given mass of a gas. This relation is called gas equation.
PV = nRT
Where, P is pressure, V is volume, T is temperature, n the number moles of gas and R the ideal gas constant.
In SI, units of pressure, volume and temperature are 
, 
and K respectively.
Therefore units of gas constant R will be R = 
= 
= 
i.e. unit of R = joule/mole - K.
R = 8.35 J/mol K is called universal gas constant.
= 
= 
is called Boltzman constant.