The behavior of gases, liquids, and solids is explained by the kinetic molecular theory, which describes matter in terms of moving particles. According to this theory, the particles of a gas are in constant, random motion, have negligible volume compared to the space between them, and undergo elastic collisions in which no kinetic energy is lost. The average kinetic energy of these particles is directly proportional to the absolute temperature of the gas.
Phase changes occur at specific temperatures. The melting point is the temperature at which a solid becomes a liquid at standard pressure, while the boiling point is the temperature at which a liquid's vapor pressure equals the surrounding atmospheric pressure, allowing it to transition into a gas. Vapor pressure is the pressure exerted by the vapor of a liquid in equilibrium with its liquid phase at a given temperature. These transitions are strongly influenced by intermolecular forces, the attractions that exist between molecules. The three main types are hydrogen bonding, dipole-dipole interactions, and London dispersion forces. Substances with stronger intermolecular forces generally have higher melting and boiling points.
A solution is a homogeneous mixture formed when one substance dissolves in another. The solute is the substance being dissolved, usually present in a smaller amount, while the solvent is the dissolving medium, typically present in greater quantity. Concentration is commonly expressed as molarity, defined as the number of moles of solute per liter of solution, written as \(M = \text{mol/L}\). Solubility depends on several factors, including temperature, which generally increases the solubility of solids but decreases that of gases; pressure, which affects gas solubility according to Henry's law, meaning the solubility of a gas in a liquid is directly proportional to its partial pressure above the liquid; and the chemical nature of the solute and solvent, summarized by the principle that "like dissolves like."