The three common states of matter—solid, liquid, and gas—exhibit distinct behaviors based on the arrangement and movement of their constituent particles. Solids maintain a fixed shape and volume, while liquids maintain a fixed volume but change shape to match their container. Gases are fundamentally unique because they expand without limit and mix completely with any other gas. Understanding this behavior begins with identifying the macroscopic observation that defines the gaseous state.
The Defining Feature: Indefinite Shape and Volume
The most defining observable property of a gas is that it has neither a fixed shape nor a fixed volume. A sample of gas will always fully expand to occupy the entire volume of its container, simultaneously taking on the container’s exact shape. This contrasts sharply with the other states of matter.
A solid, like an ice cube, retains its shape and volume regardless of the vessel it is placed into. A liquid, such as water, will conform to the shape of its container but maintains a constant, measurable volume, creating a distinct upper surface if the container is not full.
Gases are different because they always spread out to fill the entire volume available. If a gas is moved from a five-liter container into a ten-liter container, the gas will fill the full ten liters. This means the volume of a gas is entirely dependent on the capacity of the vessel that holds it.
The Microscopic Cause: Constant Motion and Vast Empty Space
The reason for this expansive behavior lies in the microscopic world, explained by the principles of the Kinetic Molecular Theory. Gas particles, whether atoms or molecules, are separated by extremely large distances relative to their size, meaning a gas is mostly empty space. The actual volume occupied by the particles themselves is considered negligible compared to the total volume the gas fills.
These particles are in a state of rapid, continuous, and random straight-line motion. They travel until they collide with another particle or the walls of the container, at which point they change direction in an elastic collision that conserves the total kinetic energy of the system. Furthermore, the particles in a gas exert virtually no attractive or repulsive forces on one another.
This lack of strong attractive forces and the presence of empty space allow the particles to move independently. Because nothing restricts their movement, they naturally spread out until the sample is uniformly distributed throughout the available volume. The average kinetic energy of these particles is directly proportional only to the absolute temperature of the gas.
Key Resulting Behaviors: Compressibility and Diffusion
The microscopic structure of gases leads directly to two highly observable and practical behaviors: compressibility and diffusion. Compressibility is the ability of a substance to significantly decrease in volume when external pressure is applied.
Because gas particles are far apart, external pressure can easily force them closer together, greatly reducing the gas’s volume. For example, the air in a scuba tank is compressed to hundreds of times its normal pressure, allowing a large amount of gas to be stored in a small volume. Liquids and solids, which have closely packed particles, are nearly incompressible.
Diffusion is the spontaneous process where gas particles mix completely and spread out until they are uniformly distributed. This occurs rapidly because the constant, random motion of the particles drives them to move into and occupy any empty space. The scent of an open bottle of perfume quickly filling a room is a common demonstration of this principle.