Summary
Highlights
Elements are the simplest pure substances composed of atoms, unable to be broken down chemically or physically. Compounds consist of molecules where different elements are chemically bonded in fixed ratios, such as water (H2O). Mixtures are substances that are not chemically bonded and can have variable compositions. Mixtures can be homogeneous (uniform, like saltwater) or heterogeneous (non-uniform, like chicken noodle soup). Notably, elements and compounds can also be heterogeneous if their state of matter is not uniform, such as ice water.
All solutions are homogeneous mixtures, composed of a solute (what is dissolved) and a solvent (what does the dissolving). For example, in saltwater, salt is the solute and water is the solvent. Water is a common solvent and has a special state symbol, 'aq' (aqueous), indicating a substance dissolved in water. Air is also a solution, where nitrogen gas is the solvent, and other gases like oxygen and carbon dioxide are solutes.
Mixtures can be separated based on the differing properties of their components. Filtration uses solubility differences, like separating salt from sand. Distillation separates components based on boiling points, useful for separating saltwater to get pure water or separating liquids like alcohol and water. Magnets can separate magnetic substances, such as iron from cereal. Chromatography separates components based on their absorption tendencies, where different substances adhere to a surface at different rates, often due to polarity.
The states of matter are solid (s), liquid (l), gas (g), and plasma (p), with aqueous (aq) indicating a substance dissolved in water. Solids have the lowest energy, with strong interparticle forces. Adding energy causes particles to move more, leading to liquids, then gases. Even more energy causes gas atoms to ionize, forming plasma. The state of matter depends on the balance between interparticle forces (stickiness) and the kinetic energy of the particles.
States of matter can change by adding or removing energy. Melting (solid to liquid) and freezing (liquid to solid) are reversible processes. Evaporation (liquid to gas) and condensation (gas to liquid) are also reversible. Sublimation is the direct transition from solid to gas (e.g., dry ice), and deposition is the direct transition from gas to solid, both skipping the liquid phase and are reversible.
Heating curves illustrate phase changes by plotting temperature against time (assuming constant heat input). When a solid is heated, its temperature increases (kinetic energy). At the melting point, the temperature stabilizes as added energy converts solid to liquid (potential energy changes due to altered particle position, not kinetic energy). Once all solid melts, the liquid's temperature rises again. At the boiling point, temperature again stabilizes as liquid turns to gas. After all liquid boils, the gas's temperature increases.
The Kelvin scale is preferred in chemistry. To convert from Celsius to Kelvin, add 273.15. The Kelvin scale is advantageous because it has an absolute zero (0 Kelvin), representing zero energy, making negative temperatures physically impossible as energy cannot be negative. This accuracy is crucial for many chemical calculations, especially in gas laws where temperature refers to average kinetic energy.