A gas is a state of matter characterized by widely separated particles that move randomly to fill any container they occupy. The starting point for understanding gas behavior is the “ideal gas,” a theoretical construct that simplifies molecular interactions. A real gas is any actual gas found in nature, such as oxygen or nitrogen, and its behavior deviates from the ideal model. This deviation occurs because the ideal model ignores physical properties inherent to all actual molecules.
The Ideal Gas Model and Its Assumptions
The theoretical framework for the ideal gas is built upon the Kinetic Molecular Theory, which makes two fundamental simplifying assumptions. First, the theory assumes that gas particles occupy a negligible volume compared to the total volume of the container. Individual molecules are treated essentially as point masses, meaning their size is considered zero.
The second major assumption is that gas particles exert no attractive or repulsive forces on one another. Collisions between particles or with the container walls are considered perfectly elastic, resulting in no loss of kinetic energy. These two assumptions allow for the derivation of the simple Ideal Gas Law (\(PV = nRT\)), which provides a baseline for gas behavior. However, this reliance on theoretical negations means the model cannot perfectly describe the behavior of actual matter.
The Physical Factors That Define a Real Gas
The defining characteristic of a real gas is its possession of the two properties the ideal model ignores: finite molecular volume and intermolecular forces. Real gas molecules, whether simple atoms or complex molecules, physically occupy space. This finite volume reduces the actual free space available for the particles to move within the container.
When gas is compressed, the space occupied by the molecules becomes a significant fraction of the total volume. This effect leads to a measured volume that is larger than the ideal prediction, as the molecules are incompressible. The second factor is the presence of intermolecular forces, which are the weak attractions and repulsions, often called Van der Waals forces, that exist between all molecules.
These forces are predominantly attractive, causing a slight pull between neighboring molecules that slows them down before they impact the container wall. This molecular attraction reduces the force and frequency of wall collisions. Consequently, the measured pressure is lower than what the ideal gas equation would predict. A real gas acknowledges and accounts for the volume of the molecules and the forces between them, leading to deviations from simplified ideal behavior.
Conditions That Cause Significant Deviation
The difference between ideal and real gas behavior becomes most pronounced under specific conditions: high pressure and low temperature. The effect of high pressure forces the gas molecules much closer together, dramatically amplifying the significance of their finite volume. Under normal atmospheric pressure, the space between molecules is vast, making molecular size negligible.
When pressure is increased, the molecules are packed tightly, and the volume they occupy can no longer be ignored. This causes the gas to be less compressible than the ideal model suggests. Similarly, low temperature causes the molecules to slow down significantly, reducing their average kinetic energy. When molecules move slowly, the weak intermolecular attractive forces have a longer time to exert influence on neighboring particles.
The reduced speed allows these forces to pull the molecules toward one another, which is the necessary precursor for a phase change, such as condensation. Consequently, at low temperatures, the attractive forces cause the gas pressure to be measurably lower than predicted, as the molecules strike the container walls with less force. Real gases generally only behave like an ideal gas at the opposite extremes: very high temperatures and very low pressures.
Modeling Real Gas Behavior
To accurately model the behavior of actual gases under non-ideal conditions, scientists modify the simple Ideal Gas Law using equations that incorporate the physical factors of real molecules. The most common and historically significant of these is the Van der Waals equation, developed in 1873. This equation adjusts the Ideal Gas Law by applying correction factors to both the pressure and volume terms.
The first correction, represented by the constant b, accounts for the finite volume of the gas molecules. This constant is subtracted from the total volume (V) to represent the volume unavailable for compression, effectively reducing the available volume for molecular motion. The second correction, represented by the constant a, accounts for the attractive intermolecular forces. This term is added to the measured pressure (P), compensating for the pressure lost due to internal attractions slowing the molecules.
Both the a and b constants are unique to each specific gas, reflecting inherent differences in molecular size and intermolecular attraction. Another method for quantifying non-ideal behavior is the Compressibility Factor (\(Z\)). This factor is the ratio of the real gas’s molar volume to the ideal gas’s molar volume under the same conditions. For an ideal gas, \(Z\) is exactly 1; therefore, the difference between the measured \(Z\) value and 1 directly indicates the degree of deviation.