Adsorption is the process by which atoms, ions, or molecules from a gas, liquid, or dissolved solid adhere to a surface.

In the ideal gas equation, R is the universal gas constant, which relates the pressure, volume, temperature, and number of moles of a gas.

'a' accounts for the attractive forces between particles, while 'b' accounts for the volume occupied by the gas particles in the van der Waals equation.

Partial pressure is the pressure that a single gas in a mixture would exert if it occupied the entire volume alone at the same temperature.

According to the Arrhenius equation, an increase in temperature results in an exponential increase in the rate constant, thus increasing the reaction rate.

The rate of diffusion of a gas is primarily affected by its molar mass; lighter gases diffuse faster than heavier gases according to Graham's law.

Boyle's Law states that for a given mass of gas at constant temperature, the pressure of the gas is inversely proportional to its volume (P1V1 = P2V2).

The equilibrium constant (K) expresses the ratio of the concentrations of products to reactants at equilibrium, indicating the extent of the reaction.

Gibbs free energy (G) indicates the spontaneity of a process; a negative change in G (ΔG < 0) suggests a spontaneous reaction.

In an isothermal process for an ideal gas, the temperature remains constant, thus the change in internal energy (ΔU) is zero.

For an ideal gas, the change in internal energy (ΔU) is given by ΔU = nC_vΔT, where C_v is the molar heat capacity at constant volume.

The triple point on a phase diagram is the unique set of conditions at which all three phases (solid, liquid, gas) coexist in equilibrium.

The Kinetic Molecular Theory explains gas behavior by describing how gas particles are in constant motion and how this motion relates to temperature and pressure.