An ideal gas is a theoretical construct used in chemistry and physics to model the behavior of gases under common conditions. This model simplifies the complex reality of gas molecules by making two primary assumptions: the particles occupy no volume and they exert no forces on one another. The ability of this theoretical gas to be compressed—the reduction of its volume under the application of external pressure—is a direct consequence of these foundational assumptions. Compressibility is a defining characteristic of the gaseous state.
The Vast Empty Space Between Particles
The fundamental reason an ideal gas is highly compressible lies in the immense amount of empty space separating the constituent particles. The Kinetic Molecular Theory assumes that the actual volume occupied by the gas molecules themselves is negligible compared to the total volume of the container. In a typical gas at standard atmospheric pressure, the particles take up far less than one percent of the total space. Since the container is overwhelmingly empty space, applying external pressure simply pushes the widely separated particles closer together. Compression is the process of reducing this unoccupied volume, which can continue until the particles are forced into close proximity.
Particle Movement and Resulting Pressure
The continuous, random motion of gas particles is directly linked to the concept of pressure and the response to compression. Ideal gas particles move at high speeds in straight lines until they collide with another particle or the walls of the container. Pressure is the macroscopic result of the constant, perfectly elastic collisions of these particles with the container walls. When an external force compresses the gas, the total volume of the container is reduced. This means particles have a shorter distance to travel before striking a wall, leading to an increase in the frequency of collisions and a proportionate increase in pressure.
Negligible Attractive Forces
A third assumption of the ideal gas model is that the particles experience no intermolecular forces of attraction or repulsion. This absence of attractive forces is necessary for the gas to be compressed without immediately changing its state. If the gas particles strongly attracted one another, they would resist being forced into closer quarters, and the overall behavior would be more similar to a liquid. The lack of any cohesive force means that compression can proceed unimpeded by molecular resistance until the physical size of the particles becomes a factor.
Where Real Gases Deviate from the Ideal
While the ideal gas model is useful, real gases only follow these rules closely under conditions of low pressure and relatively high temperature. Real gases deviate from ideal behavior when these conditions are altered, specifically at very high pressures or very low temperatures. These deviations occur because the two main assumptions of the ideal model begin to break down in the physical world.
High Pressure Deviation
At high compression, the volume of the gas particles themselves, which was assumed to be negligible, starts to occupy a significant fraction of the total container volume. This finite particle volume makes the gas less compressible than the ideal model predicts.
Low Temperature Deviation
Conversely, at very low temperatures, the particles move slower, allowing the weak attractive forces that do exist between all real molecules to become dominant. These forces pull the molecules closer together, causing the pressure to be lower than the ideal calculation suggests.