The concept of an ideal gas is a theoretical construct used in chemistry and physics to simplify gas behavior. This model assumes perfect conditions that real-world gases, or real gases, do not naturally meet. A real gas most closely mirrors an ideal gas when subjected to conditions of high temperature and low pressure.
The Theoretical Blueprint of Ideal Gas Behavior
The ideal gas model is founded on the Kinetic Molecular Theory, which makes several simplifying assumptions about gas particles. The primary assumption is that the volume occupied by the gas molecules themselves is considered negligible compared to the total volume of the container they fill. In this theoretical scenario, the particles are treated essentially as point masses occupying no space.
The theory also assumes that there are no significant intermolecular forces acting between the gas molecules. This means the particles neither attract nor repel one another as they move and collide with the walls of the container. These two assumptions form the basis for predicting gas behavior using simple mathematical relationships like the Ideal Gas Law, which provides a straightforward way to relate the macroscopic properties of pressure, volume, temperature, and the amount of gas.
Why Real Gases Deviate from the Ideal Model
All gases encountered in the physical world are considered real gases and possess physical properties that cause them to deviate from the ideal model. Unlike the theoretical point masses, real gas molecules occupy a definite, measurable volume that cannot be zero. This finite molecular size means that the total available free space for molecular movement is slightly less than the container volume, especially when the gas is compressed into a smaller space.
Furthermore, real gas molecules exhibit weak, short-range intermolecular forces, such as London dispersion forces, dipole-dipole interactions, or hydrogen bonding. These slight attractive forces cause the molecules to spend a fraction of time near one another, which influences the pressure they ultimately exert on the container walls. The degree of deviation from ideal behavior depends directly on the magnitude of these two physical realities, becoming most pronounced when the gas is cooled toward its condensation point or heavily compressed.
The Role of Low Pressure in Minimizing Molecular Volume
Applying a low external pressure to a gas is the primary mechanism by which the volume assumption of the ideal model is successfully approximated. Low pressure allows the gas to expand significantly, spreading the molecules far apart from one another. This expansion results in an extremely large total volume for the container compared to the total volume of the molecules themselves.
When the container volume is vast, the finite volume occupied by the actual gas molecules becomes insignificant in comparison to the total empty space. This condition is what allows the real gas to approximate the zero-volume point masses of the ideal gas model. The reduction in pressure effectively increases the available free space for molecular motion, which minimizes the error introduced by the real gas’s finite size in calculations using the Ideal Gas Law.
As the space between molecules increases, the total available volume overwhelmingly dwarfs the physical volume of the individual particles. This dynamic ensures that the volume term correction, often represented by the ‘b’ term in the van der Waals equation, becomes practically zero. Therefore, maintaining a low-pressure environment creates such large distances between particles that the molecules behave as if they truly had negligible volume, aligning perfectly with the theoretical blueprint.
The Role of High Temperature in Overcoming Attractive Forces
The application of high temperatures is the necessary condition to address the second major deviation: the presence of intermolecular attractive forces. Temperature is a direct measure of the average kinetic energy of the gas molecules. Raising the temperature significantly increases the speed at which all the gas particles are constantly moving, according to the Maxwell-Boltzmann distribution.
When molecules are traveling at very high velocities, their kinetic energy dramatically increases, reaching a state sufficient to overcome the weak cohesive forces that naturally exist between them. The molecules simply do not spend enough time near one another for the weak attractive forces to significantly influence their movement or collision dynamics. The high kinetic energy essentially negates the potential energy associated with the intermolecular attractions, ensuring the particles continue their straight-line paths.
This means that even the strongest of the instantaneous dipole-induced dipole interactions, like London dispersion forces, become insignificant when the molecules possess enough energy to break free instantaneously. The high-speed collisions are too brief for the forces to take hold and slow the particle down before it impacts the container wall. By keeping the gas at a high temperature, the real gas satisfies the ideal gas assumption of having no attractive or repulsive forces between particles. This condition makes the pressure exerted by the real gas match the pressure predicted by the ideal model, eliminating the need for the pressure correction term, or ‘a’ term, in the van der Waals equation.