The physical state of any substance is determined by the balance between the kinetic energy of its molecules and the attractive forces between them. Gases represent the state of matter where molecular movement is at its maximum freedom and speed. A gas is defined by its ability to expand and uniformly fill any container, a direct consequence of the unrestricted motion of its particles.
Comparing Movement in Gases, Liquids, and Solids
The primary difference between the states of matter lies in the degree of freedom the molecules possess to move relative to one another. In a solid, molecules are held tightly in fixed positions by strong intermolecular forces, allowing them only to vibrate or “jiggle” around those fixed points. This lack of translational motion gives solids their definite shape and volume.
Molecules in a liquid have enough kinetic energy to overcome some of the attractive forces, enabling them to slide past one another. This translational motion allows liquids to flow and take the shape of their container, though they maintain a definite volume because the molecules remain relatively close together. The molecules are still in constant contact, meaning there is little empty space between them.
Gas molecules exhibit the highest degree of freedom and speed, moving independently except during collisions. The distance between gas molecules is vastly greater than their size, making the gas mostly empty space. This separation means that attractive forces between gas molecules are negligible, allowing them to move rapidly, randomly, and in straight lines.
The Rules Governing Gas Motion
The scientific model explaining the unique behavior of gas molecules is the Kinetic Molecular Theory (KMT). This theory posits that a gas consists of a large number of particles in continuous, random motion. These particles travel in straight lines and only change direction when they collide with other particles or the container walls.
A foundational assumption of the KMT is that the volume occupied by the gas molecules themselves is negligible compared to the total volume of the container. The theory assumes there are no attractive or repulsive forces acting between the gas molecules. This means the motion of one molecule is entirely independent of its neighbors.
Collisions between gas molecules and the container walls are considered perfectly elastic, meaning there is no net loss of kinetic energy during these interactions. While energy may be transferred between colliding particles, the total average kinetic energy remains constant at a fixed temperature. This average kinetic energy is directly proportional to the gas’s absolute temperature, meaning higher temperatures correspond to faster-moving molecules.
How Molecular Movement Creates Observable Gas Properties
The constant, rapid, and random movement of gas molecules generates the macroscopic, observable properties of a gas. Pressure is a direct result of the force exerted by the constant bombardment of gas molecules colliding with the interior surfaces of the container. Increasing the speed of the molecules by raising the temperature increases the force and frequency of impacts, leading to a higher pressure.
The massive amount of empty space between gas molecules allows for the property of high compressibility. Because the molecules are so far apart, external pressure can easily push them closer together, significantly reducing the gas volume. Liquids and solids are nearly incompressible because their molecules are already closely packed.
Another consequence of their unconstrained motion is the ability of gases to mix completely and spontaneously, a process known as diffusion. When a gas is released, its molecules rapidly move throughout the available space, spreading from an area of high concentration to one of low concentration. The random, straight-line path of each molecule ensures that the gas will eventually distribute itself uniformly throughout the entire volume of the container.