NY Regents Chemistry key terms

The counting unit of chemistry: 6.02 × 10²³ representative particles (atoms, molecules, or formula units) of a substance. One mole of any element has a mass in grams equal to its atomic mass.

The mass in grams of one mole of a substance, found by summing the atomic masses of all atoms in its formula. It converts between grams and moles in stoichiometry problems and is expressed in g/mol.

The value 6.02 × 10²³, the number of particles in one mole of a substance. It links the microscopic count of atoms or molecules to the macroscopic amount measured in the laboratory.

The mass percentage of each element in a compound, calculated by dividing the total mass of that element in one mole by the molar mass and multiplying by 100. Used to check formulas and identify unknowns.

The simplest whole-number ratio of atoms of each element in a compound, such as CH₂O for glucose. It is found from percent composition or experimental mass data and may differ from the actual molecular formula.

The actual number of atoms of each element in one molecule of a covalent compound, such as C₆H₁₂O₆ for glucose. It is a whole-number multiple of the empirical formula, found using the molar mass.

The reactant that is used up first in a chemical reaction and therefore determines the maximum amount of product that can form. The other reactant is said to be in excess.

The actual mass of product obtained divided by the theoretical mass predicted from stoichiometry, multiplied by 100. It measures reaction efficiency and is always less than or equal to 100% in practice.

The quantitative relationship between amounts of reactants and products in a balanced chemical equation, based on mole ratios from the coefficients. It is used to predict masses, volumes, and particle counts in reactions.

A symbolic representation of a reaction in which the number of atoms of each element is equal on both sides, satisfying conservation of mass. Coefficients are adjusted, never subscripts, to achieve the balance.

The principle that matter is neither created nor destroyed in a chemical reaction, so total mass of reactants equals total mass of products. It justifies balancing equations and explains why coefficients matter.

A change in a substance's form, size, or state of matter without altering its chemical identity, such as melting ice or dissolving salt. The original substance can usually be recovered by physical means.

A process in which one or more substances are converted into new substances with different properties, such as burning or rusting. Evidence includes color change, gas production, precipitate formation, and energy release.

A combination of two or more substances that are not chemically bonded and can be separated by physical means. It may be homogeneous (uniform, like salt water) or heterogeneous (non-uniform, like sand and iron).

A substance formed when two or more different elements are chemically bonded in fixed proportions, with properties unlike those of its elements. It can be broken down only by chemical means, not physical ones.

The digits in a measurement that carry meaningful information about its precision, including all certain digits plus one estimated digit. Rules govern how they are counted and how results of calculations are rounded.

Precision describes how closely repeated measurements agree with each other; accuracy describes how close a measurement is to the true value. A measurement can be precise but inaccurate if a systematic error is present.

The absolute difference between an experimental value and the accepted value, divided by the accepted value and multiplied by 100. It quantifies accuracy rather than precision of a measurement.

The mass of a substance per unit volume, calculated as mass divided by volume and usually expressed in g/cm³ or g/mL. It is an intensive property useful for identifying substances and predicting floating or sinking.

An atom of an element with the same number of protons but a different number of neutrons, giving a different mass number. Isotopes have identical chemical behavior but different masses and stability.

The number of protons in the nucleus of an atom, which uniquely identifies the element and equals the number of electrons in a neutral atom. It determines position on the periodic table.

The sum of protons and neutrons in the nucleus of a particular atom. It is written as a superscript before the element symbol and distinguishes isotopes of the same element.

The weighted average of the masses of all naturally occurring isotopes of an element, using their relative abundances. It is the value shown on the periodic table and reported in atomic mass units.

The arrangement of electrons among the energy levels and orbitals of an atom, written in notation such as 1s²2s²2p⁶. It explains chemical properties and periodic trends.

The electrons in the outermost principal energy level of an atom, which are involved in bonding and determine chemical reactivity. Main-group elements have a number matching their group number pattern.

The energy required to remove the most loosely bound electron from a neutral gaseous atom. It generally increases across a period and decreases down a group as atomic radius grows.

A measure of how strongly an atom attracts shared electrons in a chemical bond. It generally increases across a period and decreases down a group, and differences in it predict bond type.

Half the distance between the nuclei of two bonded atoms of the same element, representing atomic size. It decreases across a period and increases down a group as energy levels are added.

The electrostatic attraction between oppositely charged ions formed when a metal transfers electrons to a nonmetal. The resulting compound is a crystalline solid that conducts electricity when melted or dissolved.

A chemical bond in which two atoms share one or more pairs of valence electrons, typically between nonmetals. It may be nonpolar or polar depending on the electronegativity difference between the atoms.

A molecule with an unequal distribution of charge because it contains polar bonds arranged asymmetrically, creating partial positive and negative ends. Water is the classic example; symmetrical molecules like CO₂ are nonpolar.

The relatively weak attractions between molecules, including dipole-dipole forces, London dispersion forces, and hydrogen bonds. They determine melting and boiling points and are weaker than ionic or covalent bonds.

A strong intermolecular 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 and surface tension.

The minimum energy colliding particles must possess for a reaction to occur, corresponding to the energy barrier between reactants and the transition state. Catalysts lower it, increasing reaction rate.

A substance that increases the rate of a reaction by providing an alternate pathway with lower activation energy, without being consumed. It speeds both forward and reverse reactions equally and does not change equilibrium position.

The state of a reversible reaction in which the forward and reverse reaction rates are equal, so concentrations of reactants and products remain constant. It is dynamic, and both reactions continue at the molecular level.

When a stress such as a change in concentration, temperature, or pressure is applied to a system at equilibrium, the system shifts to partially offset that stress. It predicts the direction of the resulting change.

The ratio of product concentrations to reactant concentrations, each raised to its coefficient in the balanced equation, at equilibrium. Its value indicates whether products or reactants are favored at that temperature.

The enthalpy change for a chemical reaction, written as ΔH. It is negative for exothermic reactions that release heat and positive for endothermic reactions that absorb heat.

A graph showing the potential energy of a system as reactants convert to products, including the activation energy peak and the net energy change. It reveals whether a reaction is exothermic or endothermic.

A carbon-based compound, usually containing carbon-hydrogen bonds, whose properties depend on its functional groups and carbon skeleton. Examples include alkanes, alkenes, alcohols, and carboxylic acids.

An organic compound containing only carbon and hydrogen atoms, such as methane, ethene, or benzene. Hydrocarbons are classified as saturated (single bonds only) or unsaturated (containing double or triple bonds).

One of two or more compounds with the same molecular formula but different arrangements of atoms and therefore different properties. Structural isomers differ in connectivity; others differ in spatial orientation.

A specific group of atoms within an organic molecule that gives the compound characteristic chemical properties, such as hydroxyl (alcohols), carboxyl (acids), or amino groups. It determines reactivity and naming.

The time required for half of the radioactive nuclei in a sample to decay. It is constant for a given isotope and is used to calculate remaining mass or to date materials.

The spontaneous emission of particles or radiation from an unstable nucleus, changing it into a different element. Alpha decay emits helium nuclei, beta decay emits electrons, and gamma emission releases energy only.

The splitting of a heavy nucleus into lighter nuclei with the release of a large amount of energy and neutrons. It powers nuclear reactors and can sustain a chain reaction.

The combining of light nuclei to form a heavier nucleus, releasing far more energy than fission per unit mass. It powers stars and requires extremely high temperatures to overcome electrostatic repulsion.