The fact that particles of matter are in constant motion can be explained by the following:

1. Brownian motion: Robert Brown in his pollen grain experiment observed the behavior of pollen grains in a drop of water under a microscope. He observed that the pollen grains were moving about in a haphazard, zig-zag manner. This motion was due the bombardment of solid pollen grains by its surrounding liquid environment. This haphazard and irregular motion is called Brownian motion.

2. Diffusion: Diffusion is the spontaneous movement of particles of an (atoms, ions, or molecules) from an area of higher concentration to an area of lower concentration. This process occurs due to the random motion of particles, resulting in a uniform distribution of particles throughout the medium. Diffusion is fastest in gases due to weak intermolecular forces. Example is the spreading of the scent of a flower

3. Osmosis: Osmosis is the movement of solvent molecules across a selectively permeable membrane from an area of lower solute concentration to an area of higher solute concentration. This process aims to equalize the concentration of solute on both sides of the membrane, resulting in the establishment of a dynamic equilibrium.

1. Gas consists of particles: According to the kinetic theory, gases are composed of tiny particles, either atoms or molecules, in constant, random motion.
2. Negligible volume and no intermolecular forces: The volume occupied by gas particles is considered negligible compared to the total volume of the gas. Additionally, the kinetic theory assumes that there are no significant attractive or repulsive forces between gas particles, except during collisions.
3. Collisions are elastic: Collisions between gas particles and with the walls of the container are perfectly elastic, meaning there is no net loss or gain of kinetic energy during these collisions.
4. Continuous, random motion: Gas particles are in constant motion, moving in straight lines until they collide with each other or with the walls of the container. The direction of motion is random.
5. Temperature is related to kinetic energy: The temperature of a gas is directly proportional to the average kinetic energy of its particles. As temperature increases, so does the average kinetic energy of the gas particles.

1. Volume: In the context of gases, volume refers to the amount of space occupied by the gas. According to the kinetic theory of gases, gas particles themselves occupy negligible volume, and the total volume is determined by the space between particles and the size of the container. $$ 1dm³ = 1000cm³ $$
2. Pressure: Pressure is the force exerted per unit area. In gases, pressure arises from the collisions of gas particles with the walls of the container. The kinetic theory explains pressure in terms of the continuous, random motion and elastic collisions of gas particles.
$$ 1 atm = 760mmHg $$ $$ 1atm = 1.0135 × 10^5pa $$
3. Temperature: Temperature in gases is related to the average kinetic energy of particles. As temperature increases, gas particles move faster, leading to an increase in kinetic energy. This relationship is a fundamental aspect of the kinetic theory of gases. For gas laws, it is important to use kelvin scale. To convert Celsius scale to kelvin scale we use: $$ T(K) = t(°C) + 273 $$
4. Number of moles: The number of moles represents the quantity of substance in a given sample. In the context of gases, the kinetic theory incorporates the concept of moles to describe the amount of gas present, influencing its behavior and properties.
$$ 1 mole = 22.4dm³ $$ $$ 1 mole = 6.023 × 10^{23} particles $$
5. Standard Temperature and Pressure (STP): STP is a standard set of conditions used for comparing gas properties.
$$ \text{standard temperature} = 273k $$ $$ \text{standard pressure} = 760mmHg $$