2.1 Molecular Model of an Ideal Gas

The ideal gas law relates the pressure and volume of a gas to the number of gas molecules and the temperature of the gas.
A mole of any substance has a number of molecules equal to the number of atoms in a 12-g sample of carbon-12. The number of molecules in a mole is called Avogadro’s number $N_{A},$
$N_{A} = 6.02 \times 10^{23} {mol}^{−1} .$
A mole of any substance has a mass in grams numerically equal to its molecular mass in unified mass units, which can be determined from the periodic table of elements. The ideal gas law can also be written and solved in terms of the number of moles of gas:
$p V = n R T,$
where n is the number of moles and R is the universal gas constant,
$R = 8.31 J/mol \cdot K .$
The ideal gas law is generally valid at temperatures well above the boiling temperature.
The van der Waals equation of state for gases is valid closer to the boiling point than the ideal gas law.
Above the critical temperature and pressure for a given substance, the liquid phase does not exist, and the sample is “supercritical.”

2.2 Pressure, Temperature, and RMS Speed

Kinetic theory is the atomic description of gases as well as liquids and solids. It models the properties of matter in terms of continuous random motion of molecules.
The ideal gas law can be expressed in terms of the mass of the gas’s molecules and $\overset{–}{v^{2}},$ the average of the molecular speed squared, instead of the temperature.
The temperature of gases is proportional to the average translational kinetic energy of molecules. Hence, the typical speed of gas molecules $v_{rms}$ is proportional to the square root of the temperature and inversely proportional to the square root of the molecular mass.
In a mixture of gases, each gas exerts a pressure equal to the total pressure times the fraction of the mixture that the gas makes up.
The mean free path (the average distance between collisions) and the mean free time of gas molecules are proportional to the temperature and inversely proportional to the molar density and the molecules’ cross-sectional area.

Every degree of freedom of an ideal gas contributes $\frac{1}{2} k_{B} T$ per atom or molecule to its changes in internal energy.
Every degree of freedom contributes $\frac{1}{2} R$ to its molar heat capacity at constant volume $C_{V} .$
Degrees of freedom do not contribute if the temperature is too low to excite the minimum energy of the degree of freedom as given by quantum mechanics. Therefore, at ordinary temperatures, $d = 3$ for monatomic gases, $d = 5$ for diatomic gases, and $d \approx 6$ for polyatomic gases.

The motion of individual molecules in a gas is random in magnitude and direction. However, a gas of many molecules has a predictable distribution of molecular speeds, known as the Maxwell-Boltzmann distribution.
The average and most probable velocities of molecules having the Maxwell-Boltzmann speed distribution, as well as the rms velocity, can be calculated from the temperature and molecular mass.