States of Matter and Phase Changes

Compressing or cooling a gas makes it harder for molecules to move away from each other. As molecular motion decreases, weak attractive intermolecular forces become dominant, transforming the gas into a denser, less compressible liquid. This phase transition occurs because reduced kinetic energy allows intermolecular attractions to hold molecules closer together.

During phase transitions, molecules exist in both phases simultaneously while temperature stays constant despite continuous heat influx. The added energy breaks intermolecular bonds rather than increasing molecular motion. Once the bulk transition completes, temperature rises again. This energy change without temperature change is called molar enthalpy of the transition.

Substances with stronger intermolecular forces require more energy to undergo phase transitions, so these transitions occur at higher temperatures than substances with weaker intermolecular interactions. The molar enthalpy of vaporization, for example, depends directly on intermolecular force strength. Stronger attractions demand greater energy input to separate molecules into the gaseous state.

Melting occurs when added heat causes solid particles to vibrate fast enough to overcome intermolecular attraction and move from fixed positions into a liquid state. Vaporization happens at higher temperatures when particles move fast enough to escape into the gaseous state. Both are endothermic processes, but vaporization requires more energy because gas molecules are farther apart than liquid molecules.

Vaporization absorbs energy to overcome intermolecular forces and allow molecules to escape into gas phase, making it endothermic with positive enthalpy. Condensation releases energy as gas molecules lose kinetic energy and form liquid, making it exothermic with negative enthalpy. These processes are thermodynamic opposites: vaporization requires energy input while condensation releases it.

Internal energy—the total kinetic and potential energy of all molecules—is highest in the gaseous state, lowest in the solid state, and intermediate in the liquid state. Gas molecules move rapidly with weak intermolecular attractions, while solid molecules vibrate in fixed positions with strong attractions. Liquid molecules occupy an intermediate state with moderate kinetic and potential energy.

Pressure affects phase transitions by influencing intermolecular interactions and the internal energy of substances. Compressing a gas increases pressure and forces molecules closer together, facilitating the transition to liquid. At a given pressure, phase transition points and their associated energy changes depend on intermolecular force strength, making pressure a critical variable in controlling phase equilibria like distillation vapor ndash liquid equilibria.