Kinetic Theory

Low tempurature

At low temperatures, the particles are moving slowly, and the attractions between the particles become significant, this gives a pressure that is smaller than expected because the particles are slowing down as they are drawn to each other (and not bouncing off the side walls as often because they are slowing down)

Effusion and Diffusion

Both effusion and diffusion will happen faster for a lighter molecule which will have a larger velocity at the same temperature

Extreme temperature

Close to the condensation point (boiling point), molecules exert attractive forces on each other - therefore the gas will deviate from ideal behavior; Volume will be less than predicted by ideal gas law, Pressure will be less than predicted by ideal gas law

Postulates

Gases are composed of a large number of particles that behave like hard, spherical objects in a state of constant, random motion. These particles move in a straight line until they collide with another particle or the walls of the container. These particles are much smaller than the distance between particles. Most of the volume of a gas is therefore empty space. There is no force of attraction between gas particles or between the particles and the walls of the container. Collisions between gas particles or collisions with the walls of the container are perfectly elastic. None of the energy of a gas particle is lost when it collides with another particle or with the walls of the container. The average kinetic energy of a collection of gas particles depends on the temperature of the gas and nothing else. Gases consist of particles in constant, random motion. They continue in a straight line until they collide with something—usually each other or the walls of their container. Particles are point masses with no volume. The particles are so small compared to the space between them, that we do not consider their size in ideal gases. No molecular forces are at work. This means that there is no attraction or repulsion between the particles. Gas pressure is due to the molecules colliding with the walls of the container. All of these collisions are perfectly elastic, meaning that there is no change in energy of either the particles or the wall upon collision. No energy is lost or gained from collisions. The time it takes to collide is negligible compared with the time between collisions. The kinetic energy of a gas is a measure of its Kelvin temperature. Individual gas molecules have different speeds, but the temperature and kinetic energy of the gas refer to the average of these speeds. The average kinetic energy of a gas particle is directly proportional to the temperature. An increase in temperature increases the speed in which the gas molecules move. All gases at a given temperature have the same average kinetic energy. Lighter gas molecules move faster than heavier molecules.

Polar Molecules

Polar molecules exert attractive forces on each other - therefore polar gases will deviate from ideal behavior

High pressure

Under high pressure, the particles are squished close together (because higher pressure smaller distance) and the volume of the particles themselves affects the total volume, this results in a measured volume which is higher than expected (but as the pressure of the gas increases, the intermolecular distances become smaller and smaller (Figure 6.23). As a result, the volume occupied by the molecules becomes significant compared with the volume of the container. Consequently, the total volume occupied by the gas is greater than the volume predicted by the ideal gas law. Thus at very high pressures, the experimentally measured value of PV/nRT is greater than the value predicted by the ideal gas law)

Newton's First Law Of Motion

any object with mass (intertia) will move in a straight line at a constant speed unless acted on by a force

Particles

have no appricable volume, have perfect elastic collisions when they collide with the containor and other particles, no forces between the particles (in ideal gas)

Deviations from Ideal gas

if particles are attracked/ interacting with one another than the gas is no longer ideal (n fact , deviations from ideality occur due to both attractive and repulsive forces. The Kinetic Gas theory makes two controversial assumptions: 1.Gas molecules are point sized particles occupying negligible volume 2. They exert negligible forces on each other)

Diffusion

mixing of gases with each other due to random motion of particles (ex: opening the stopcock) (Diffusion is a simple process that can be explained by kinetic theory. When you open a bottle of perfume, it can very quickly be smelled on the other side of the room. This is because as the scent particles drift out of the bottle, gas molecules in the air collide with the particles and gradually distribute them throughout the air. Diffusion of a gas is the process where particles of one gas are spread throughout another gas by molecular )

Gases

mostly epty space; particles seperated by long distances, particles occupy no volume the volume of the gas is a result of the motion of the particles and the collisions with the containor

Elastic Collisions

no net loss of energy in the collision (ex: billiard balls)

Point masses

particles are considered this and have no appreciative volume (because they are so small)

Kinetic energy

tempurature is a meausre of average kinetic energy of the particles (substances with higher molecular weight will move with a lower average velocity at the same temperature), KE = ½ mv2

Effusion

the escape of a gas through a small opening due to random motion of particles