A gas is a state of matter characterized by the absence of a fixed shape or definite volume. Unlike solids or liquids, a gas readily expands to completely fill any container it occupies. To understand the macroscopic behavior of gases, scientists examine the microscopic actions of the individual particles (atoms or molecules) that compose them. These particles govern all observable properties of the gas. This microscopic perspective provides a framework for predicting how gases respond to changes in pressure, temperature, and volume.
Spacing and Volume
A defining characteristic of a gas is the vast distance separating its component particles. The space between molecules is often ten times greater than the diameter of the molecules themselves. This large separation means the actual volume occupied by the particles is vanishingly small compared to the total volume of the container, which is predominantly empty space.
This significant distance explains why gases are easily compressed and lack a fixed volume. Since the particles are far apart, they can be pushed much closer together without strong repulsive forces. Therefore, in models describing gas behavior, the individual volume of the particles is disregarded. The total volume of the system is defined by the boundaries of the vessel holding the gas.
Constant, Random Motion
Gas particles are characterized by continuous, rapid movement in a state of perpetual disarray. This motion is entirely random, meaning there is no preferred direction or predictable path for any single particle. A particle travels in a straight line until it encounters another particle or the container wall. This chaotic movement ensures the gas quickly and uniformly disperses throughout the entire available volume.
The speeds achieved by these particles are remarkably high, contributing to their ability to rapidly fill a space. For example, nitrogen molecules travel at an average speed of about 515 meters per second at room temperature. This velocity is significantly faster than the speed of sound in air. The constant, unpredictable changes in direction result from the frequent collisions occurring within the gas.
This perpetual state of kinetic energy means gas particles never settle or cease their movement under normal circumstances. The overall effect of this microscopic activity is the macroscopic phenomenon known as pressure. Pressure is the cumulative force exerted by these constantly moving particles as they strike the interior surfaces of their container.
Energy and Temperature Relationship
The energy possessed by gas particles is directly related to the observable temperature of the system. Temperature is a direct measure of the average kinetic energy of the particles within the gas. Kinetic energy is the energy an object possesses due to its motion, meaning faster-moving particles possess higher kinetic energy. Therefore, increasing the temperature of a gas causes its particles to move with greater average velocity.
Not every particle moves at the same speed at any given moment. At a specific temperature, the particles exhibit a wide distribution of speeds, ranging from very slow to extremely fast. The temperature reflects the calculated average speed of this entire population of molecules. This relationship is linear, meaning doubling the temperature measured on the Kelvin scale will double the average kinetic energy.
This concept explains why a gas expands when heated. The increased kinetic energy leads to more forceful and frequent collisions. The faster-moving particles push harder against the container walls, attempting to occupy a larger volume.
Interactions and Collisions
A general truth about gas particles is that the attractive or repulsive forces between them, known as intermolecular forces, are negligible. Due to the large distances separating the particles, these forces are too weak to influence particle motion significantly. The particles act independently of one another, unlike the cohesive forces that hold liquids and solids together. This assumption simplifies the study of gases, allowing for accurate prediction of their behavior.
When gas particles collide with each other or the container walls, these interactions are considered perfectly elastic. An elastic collision is one where the total kinetic energy of the system remains unchanged before and after the impact. While one particle might lose energy and slow down, the other particle gains an equivalent amount of energy and speeds up. This conservation of energy means the gas as a whole does not lose kinetic energy over time due to internal friction or heat generation.
These elastic collisions are the mechanism by which pressure is maintained and energy is distributed throughout the gas volume. The particles constantly exchange energy and momentum without any net loss to the system. This behavior contrasts sharply with the inelastic collisions commonly observed in the macroscopic world.