MTLE Chemistry (Grades 9-12) key terms

The amount of substance containing exactly 6.022 × 10²³ representative particles. It links the microscopic count of atoms or molecules to measurable laboratory quantities such as mass and volume.

The value 6.022 × 10²³, equal to the number of particles in one mole of a substance. It is used to convert between particle counts and moles in calculations.

The mass in grams of one mole of a substance, numerically equal to its formula mass in atomic mass units. It converts between moles and grams in stoichiometric problems.

The quantitative relationship among reactants and products in a balanced chemical equation. It uses mole ratios to predict masses, volumes, or particle amounts consumed or produced.

The reactant that is entirely consumed first, capping the maximum amount of product that can form. Identifying it requires comparing available moles to the equation's mole ratios.

The ratio of actual product mass obtained to theoretical maximum product mass, multiplied by 100. It measures reaction efficiency and reveals losses from side reactions or handling.

The simplest whole-number ratio of atoms of each element in a compound. It is derived from percent composition data and may differ from the actual molecular formula.

The mass percentage of each element in a compound, calculated by dividing that element's total mass contribution by the molar mass. It provides the data needed to determine empirical formulas.

The electrostatic attraction between oppositely charged ions formed when one atom transfers electrons to another. It typically occurs between metals and nonmetals and produces crystalline solids with high melting points.

A bond in which two atoms share one or more pairs of valence electrons. It occurs mainly between nonmetals and can be nonpolar or polar depending on electronegativity differences.

A measure of an atom's tendency to attract shared electrons within a bond. Differences in this value determine bond polarity and help classify bonds as nonpolar, polar covalent, or ionic.

A model predicting molecular geometry from the repulsion of electron domains around a central atom. Lone pairs occupy more space than bonding pairs, distorting bond angles from ideal values.

The mixing of atomic orbitals to form equivalent new orbitals that explain observed molecular geometries and bonding. Common types such as sp, sp², and sp³ correspond to specific electron-domain arrangements.

Attractions between separate molecules, including hydrogen bonding, dipole-dipole forces, and London dispersion forces. They are much weaker than covalent bonds but govern boiling points, solubility, and phase behavior.

A strong dipole-dipole attraction between a hydrogen atom bonded to N, O, or F and a lone pair on a nearby electronegative atom. It explains water's unusually high boiling point.

A measure of the separation of positive and negative charge in a polar molecule, expressed in debye units. It reflects both bond polarity and molecular geometry and influences intermolecular attractions.

A rule stating that a system at equilibrium shifts to partially offset a change in concentration, pressure, or temperature. It predicts the direction of shift but not the new equilibrium constant.

The ratio of product concentrations to reactant concentrations, each raised to its stoichiometric coefficient, at equilibrium. Its value depends on temperature and indicates whether products or reactants are favored.

The same ratio as the equilibrium constant but calculated with current, non-equilibrium concentrations. Comparing it to the equilibrium constant predicts whether the reaction proceeds forward or reverse.

The negative base-10 logarithm of hydrogen ion concentration in a solution. Values below 7 are acidic, above 7 are basic, and each unit represents a tenfold change in concentration.

A solution of a weak acid and its conjugate base, or a weak base and its conjugate acid, that resists large pH changes on adding acid or base. It works best near its pKa.

The equilibrium constant for the ionization of a weak acid, denoted Ka, with larger values indicating a stronger acid. Its negative logarithm, pKa, is convenient for comparing acid strengths.

A reaction involving the transfer of electrons, in which one species is oxidized (loses electrons) and another is reduced (gains electrons). Oxidation states change for the reacting elements.

A number assigned to an atom reflecting its apparent charge if bonding electrons were fully assigned to the more electronegative atom. Tracking its changes identifies redox processes and balances equations.

The separate oxidation or reduction portion of a redox reaction, showing electrons explicitly. Combining balanced half-reactions yields the overall balanced equation for complex redox systems.

The minimum energy that colliding particles must possess for a reaction to occur. It determines reaction rate and is lowered by catalysts, which provide an alternative pathway.

A substance that increases reaction rate by lowering activation energy without being consumed. It does not change the equilibrium position or the thermodynamic favorability of the reaction.

The change in concentration of a reactant or product per unit time. It depends on concentration, temperature, surface area, and the presence of a catalyst.

The equation PV = nRT relating pressure, volume, moles, and absolute temperature of a gas. It assumes negligible particle volume and no intermolecular attractions.

The pressure a single gas in a mixture would exert if it alone occupied the container. Dalton's law states that the total pressure equals the sum of the component partial pressures.

A model describing gases as particles in constant random motion whose average kinetic energy is proportional to absolute temperature. It explains pressure, diffusion, and deviations from ideal behavior.

The experimental measurement of heat exchanged during a physical or chemical process, often using a calorimeter and the equation q = mcΔT. It yields enthalpy changes for reactions in solution.

A thermodynamic quantity representing heat content at constant pressure, with changes denoted ΔH. Negative ΔH indicates an exothermic process, positive indicates endothermic.

The principle that the enthalpy change of an overall reaction equals the sum of enthalpy changes for its individual steps, regardless of pathway. It allows indirect calculation of ΔH from known reactions.

The heat energy needed to raise the temperature of one gram of a substance by one degree Celsius. It appears in calorimetry calculations and reflects a material's thermal inertia.

A thermodynamic function, ΔG = ΔH − TΔS, whose sign predicts reaction spontaneity at constant temperature and pressure. Negative values indicate spontaneous processes.

A measure of energy dispersal or the number of accessible microstates in a system, denoted S. It increases for spontaneous processes in isolated systems and drives many dissolution and phase changes.

The time required for half of a radioactive sample to decay, a constant characteristic of each isotope. It supports radiometric dating and calculations of remaining activity.

The spontaneous emission of particles or radiation from an unstable nucleus, transforming it into a more stable one. Common modes are alpha, beta, and gamma emission.

The splitting of a heavy nucleus into lighter nuclei, releasing large amounts of energy and neutrons. It underlies nuclear power generation and can sustain a chain reaction.

A teaching approach in which students investigate questions through their own observations and data collection rather than receiving conclusions directly. The teacher facilitates questioning, experimental design, and evidence-based explanation.

Ongoing checks of student understanding used to adjust instruction during learning, such as exit tickets, whiteboard work, or questioning. It differs from summative assessment, which grades final achievement.

The behaviors scientists and engineers use, including asking questions, planning investigations, analyzing data, and constructing explanations. Instruction should embed these practices in chemistry content rather than teach them in isolation.

A persistent, plausible but incorrect idea students hold, such as believing atoms expand when heated. Effective instruction surfaces these ideas and provides evidence-based conflict to promote conceptual change.

A document listing a chemical's hazards, handling precautions, and emergency procedures. Teachers must review it before laboratory activities and maintain accessible copies.

Practices and rules that prevent injury in the lab, including goggles, fume hoods, waste disposal, and emergency equipment access. Instruction must include student training and documented safety contracts.