Chemistry Glossary: 150+ Terms & Definitions

General Chemistry

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Foundational concepts that underpin all branches of chemistry, including atoms, molecules, solutions, and acid-base theory.

Acid

A substance that donates protons (H⁺) or accepts electron pairs in chemical reactions, producing a pH below 7 in aqueous solution.

Atom

The smallest unit of an element that retains the chemical properties of that element, consisting of a nucleus surrounded by electrons.

Avogadro's Number

The number of constituent particles in one mole of a substance, approximately 6.022 × 10²³, a fundamental constant in chemistry.

Base

A substance that accepts protons or donates electron pairs, producing a pH above 7 in aqueous solution and often yielding hydroxide ions.

Colligative Properties

Physical properties of solutions that depend on the number of dissolved solute particles rather than the identity of the solute.

Compound

A substance formed when two or more elements are chemically bonded together in a fixed ratio, with properties distinct from its constituent elements.

Concentration

The amount of solute present in a given quantity of solvent or solution, commonly expressed as molarity, molality, or mass percent.

Element

A pure substance consisting entirely of one type of atom, defined by its atomic number and unable to be broken down by chemical means.

Ion

An atom or molecule that carries a net electric charge due to the loss or gain of one or more electrons.

Mixture

A combination of two or more substances that are not chemically bonded and can be separated by physical methods.

Mole

The SI unit for amount of substance, equal to exactly 6.02214076 × 10²³ elementary entities such as atoms, molecules, or ions.

Molecule

A group of two or more atoms held together by covalent bonds, representing the smallest fundamental unit of a chemical compound.

Mole Fraction

The ratio of the number of moles of a particular component to the total number of moles of all components in a mixture.

Osmotic Pressure

The minimum pressure that must be applied to a solution to prevent the inward flow of solvent through a semipermeable membrane.

Oxidation

The loss of electrons by an atom, ion, or molecule during a chemical reaction, resulting in an increase in oxidation state.

pH

A logarithmic scale measuring the acidity or basicity of an aqueous solution, defined as the negative logarithm of the hydrogen ion concentration.

Raoult's Law

A law stating that the partial vapor pressure of each component in an ideal solution is proportional to its mole fraction multiplied by the vapor pressure of the pure component.

Reduction

The gain of electrons by an atom, ion, or molecule during a chemical reaction, resulting in a decrease in oxidation state.

Solubility

The maximum amount of a solute that can dissolve in a given quantity of solvent at a specified temperature and pressure to form a saturated solution.

Solution

A homogeneous mixture in which a solute is uniformly dissolved in a solvent at the molecular level.

Valence

The combining power of an element, equal to the number of hydrogen atoms it can combine with or displace, determined by the number of electrons in the outermost shell.

Vapor Pressure

The pressure exerted by a vapor in thermodynamic equilibrium with its condensed phase at a given temperature.

Atomic Structure

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Terms related to the internal architecture of atoms, including subatomic particles, orbitals, and quantum mechanical models.

Atomic Radius

A measure of the size of an atom, typically defined as half the distance between the nuclei of two bonded identical atoms.

Aufbau Principle

The principle stating that electrons fill atomic orbitals from lowest to highest energy level, following the order 1s, 2s, 2p, 3s, 3p, 4s, 3d, and so on.

De Broglie Wavelength

The wavelength associated with a particle, calculated as Planck's constant divided by the particle's momentum, demonstrating wave-particle duality.

Effective Nuclear Charge

The net positive charge experienced by a valence electron after accounting for the shielding effect of core electrons, often denoted Z_eff.

Electron

A negatively charged subatomic particle that orbits the nucleus and determines chemical bonding behavior and reactivity.

Electron Affinity

The energy change that occurs when a neutral gaseous atom gains an electron, indicating the atom's tendency to accept electrons.

Electron Configuration

The distribution of electrons among the orbitals of an atom, written using notation such as 1s² 2s² 2p⁶.

Excited State

A higher energy state of an atom in which one or more electrons have absorbed energy and moved to a higher orbital.

Ground State

The lowest energy state of an atom in which all electrons occupy the lowest available energy levels.

Heisenberg Uncertainty Principle

The principle stating that it is impossible to simultaneously know both the exact position and momentum of a subatomic particle.

Ionization Energy

The minimum energy required to remove the outermost electron from a gaseous atom or ion in its ground state.

Isotope

Atoms of the same element that have the same number of protons but different numbers of neutrons, giving them different masses.

Mass Number

The total number of protons and neutrons in an atom's nucleus, represented by the symbol A.

Neutron

An electrically neutral subatomic particle in the nucleus that contributes to atomic mass and accounts for isotope variation.

Orbital

A mathematical function describing the region of space where an electron is most likely to be found around an atom.

Pauli Exclusion Principle

The quantum mechanical principle stating that no two electrons in an atom can have the same set of four quantum numbers.

Proton

A positively charged subatomic particle found in the atomic nucleus whose number defines the element's identity.

Quantum Number

A set of four numbers (n, l, mₗ, mₛ) that describe the energy, shape, orientation, and spin of an electron in an atom.

Shielding Effect

The reduction in effective nuclear charge experienced by outer-shell electrons due to the repulsive effect of inner-shell electrons.

Spin Quantum Number

A quantum number (ms = +½ or −½) that describes the intrinsic angular momentum of an electron, determining its magnetic orientation within an orbital.

Chemical Bonding

21

Concepts covering the forces that hold atoms together in molecules and crystals, from ionic to covalent to metallic interactions.

Bond Energy

The amount of energy required to break one mole of a specific type of bond in a gaseous substance under standard conditions.

Bond Length

The average distance between the nuclei of two bonded atoms in a molecule, typically measured in picometers.

Bond Order

The number of chemical bonds between a pair of atoms, calculated in molecular orbital theory as half the difference between the number of bonding and antibonding electrons.

Covalent Bond

A chemical bond formed by the sharing of one or more pairs of electrons between two atoms.

Dipole Moment

A quantitative measure of the net polarity of a molecule, equal to the product of the charge magnitude and the distance between the positive and negative charge centers.

Electronegativity

A measure of the tendency of an atom to attract a shared pair of electrons in a chemical bond, as quantified by the Pauling scale.

Hybridization

The mixing of atomic orbitals to form new hybrid orbitals of equal energy, such as sp, sp², and sp³, used to explain molecular geometry.

Hydrogen Bond

A relatively strong intermolecular force between a hydrogen atom bonded to an electronegative atom and a lone pair on a nearby electronegative atom.

Intermolecular Force

Attractive or repulsive forces between molecules, including dipole-dipole interactions, London dispersion forces, and hydrogen bonds, which determine bulk physical properties.

Ionic Bond

A chemical bond formed by the electrostatic attraction between oppositely charged ions, typically between a metal and a nonmetal.

Lattice Energy

The energy released when gaseous ions combine to form one mole of an ionic solid crystal, serving as a measure of the strength of ionic bonds in a compound.

Lewis Structure

A diagram showing the bonding between atoms and the lone pairs of electrons in a molecule using dots and lines.

Metallic Bond

A bond resulting from the attraction between metal cations and a delocalized sea of valence electrons shared among all atoms.

Molecular Orbital Theory

A bonding model in which atomic orbitals combine to form molecular orbitals that are delocalized over the entire molecule, producing bonding and antibonding states.

Nonpolar Bond

A covalent bond in which electrons are shared equally between two atoms of the same or very similar electronegativity.

Pi Bond

A covalent bond formed by the lateral overlap of parallel p orbitals above and below the internuclear axis.

Polar Bond

A covalent bond in which electrons are unequally shared between atoms due to a difference in electronegativity.

Polarizability

A measure of how easily the electron cloud of an atom or molecule is distorted by an external electric field, influencing the strength of London dispersion forces.

Resonance

A way of describing delocalized electrons using two or more Lewis structures, where the true structure is a hybrid of all contributing forms.

Sigma Bond

A covalent bond formed by the head-on overlap of atomic orbitals along the internuclear axis, allowing free rotation.

Van der Waals Forces

Weak intermolecular forces arising from temporary or permanent dipoles, including London dispersion forces and dipole-dipole interactions.

Thermodynamics

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Terms describing energy changes in chemical processes, including enthalpy, entropy, and the laws governing spontaneous reactions.

Calorimetry

The experimental technique of measuring heat changes during chemical reactions or physical processes, typically using an insulated calorimeter and a temperature probe.

Chemical Potential

The partial molar Gibbs free energy of a component in a mixture, representing the change in free energy when an infinitesimal amount of the substance is added at constant temperature and pressure.

Clausius-Clapeyron Equation

An equation relating the vapor pressure of a substance to temperature through the enthalpy of vaporization, used to describe phase transition boundaries on a phase diagram.

Endothermic

Describing a process that absorbs heat from the surroundings, resulting in a positive change in enthalpy (ΔH > 0).

Enthalpy

A thermodynamic quantity equal to the internal energy of a system plus the product of its pressure and volume, commonly measured as heat flow at constant pressure.

Entropy

A measure of the disorder or randomness of a system, which tends to increase in spontaneous processes according to the second law.

Equilibrium Constant

A numerical value (K) expressing the ratio of product concentrations to reactant concentrations at chemical equilibrium, each raised to their stoichiometric coefficients.

Exothermic

Describing a process that releases heat to the surroundings, resulting in a negative change in enthalpy (ΔH < 0).

First Law of Thermodynamics

The law of conservation of energy stating that the total energy of an isolated system is constant: energy can be converted but not created or destroyed.

Free Energy of Mixing

The change in Gibbs free energy when two or more pure substances are combined to form a homogeneous mixture, driven primarily by the entropy of mixing.

Gibbs Free Energy

A thermodynamic potential (G = H − TS) that predicts whether a process will occur spontaneously at constant temperature and pressure.

Heat Capacity

The amount of heat energy required to raise the temperature of a substance by one degree, expressed per unit mass or per mole.

Hess's Law

The principle that the total enthalpy change of a reaction is independent of the pathway and equals the sum of enthalpy changes of the individual steps.

Internal Energy

The total kinetic and potential energy of all particles in a system, a state function fundamental to the first law.

Second Law of Thermodynamics

The law stating that the total entropy of an isolated system always increases over time, defining the direction of spontaneous processes.

Spontaneous Process

A process that occurs naturally without external intervention, characterized by a negative Gibbs free energy change (ΔG < 0).

Standard State

The reference condition for a substance defined as 1 bar pressure (or 1 atm in older conventions) at a specified temperature, used as the baseline for thermodynamic data.

Thermodynamic System

A defined region of the universe under study, classified as open, closed, or isolated depending on whether it exchanges matter and energy with its surroundings.

Third Law of Thermodynamics

The law stating that the entropy of a perfect crystal approaches zero as the temperature approaches absolute zero.

Chemical Kinetics

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Concepts related to how fast reactions proceed, including rate laws, activation energy, catalysis, and reaction mechanisms.

Activated Complex

A transient, high-energy arrangement of atoms that exists at the peak of the energy barrier along the reaction coordinate, also called the transition state complex.

Activation Energy

The minimum energy that reacting molecules must possess for a collision to result in a chemical reaction.

Arrhenius Equation

The equation k = Ae^(−Ea/RT) that relates the rate constant to temperature, activation energy, and a pre-exponential factor.

Autocatalysis

A reaction in which one of the products serves as a catalyst for the reaction itself, resulting in a sigmoidal rate curve that accelerates before slowing.

Catalyst

A substance that increases the rate of a chemical reaction by lowering the activation energy without being consumed in the process.

Catalyst Poisoning

The partial or total deactivation of a catalyst caused by exposure to chemical compounds that bind irreversibly to active sites, reducing catalytic performance.

Collision Theory

A theory stating that reactions occur when molecules collide with sufficient energy and proper orientation to break and form bonds.

Elementary Step

A single molecular event in a reaction mechanism that describes the actual collision or rearrangement of molecules.

Enzyme

A biological catalyst, typically a protein, that dramatically accelerates specific biochemical reactions with high selectivity.

Half-Life

The time required for the concentration of a reactant to decrease to half of its initial value.

Molecularity

The number of reactant molecules or atoms that participate in a single elementary step of a reaction mechanism, classified as unimolecular, bimolecular, or termolecular.

Order of Reaction

The sum of the exponents to which reactant concentrations are raised in the rate law, indicating how rate depends on concentration.

Pseudo-First Order

A simplification of reaction kinetics in which one reactant is present in such large excess that its concentration effectively remains constant, reducing the apparent reaction order.

Rate Constant

The proportionality constant k in a rate law that is specific to a given reaction at a given temperature.

Rate-Determining Step

The slowest elementary step in a reaction mechanism, which limits the overall rate of the reaction.

Rate Law

A mathematical expression relating the rate of a reaction to the concentrations of reactants raised to experimentally determined powers.

Reaction Coordinate

A one-dimensional coordinate representing the progress of a reaction from reactants through the transition state to products.

Reaction Mechanism

The step-by-step sequence of elementary reactions by which an overall chemical change occurs.

Reaction Rate

The change in concentration of a reactant or product per unit time, indicating how fast a chemical reaction proceeds.

Transition State

A high-energy, unstable configuration of atoms at the peak of the activation energy barrier in a reaction, existing only momentarily as reactants convert to products.

Organic Chemistry

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Terms covering carbon-based compounds, their structures, functional groups, stereochemistry, and polymerization.

Aldehyde

An organic compound containing a carbonyl group (C=O) bonded to at least one hydrogen atom, with the general formula RCHO.

Alkane

A saturated hydrocarbon containing only single carbon-carbon bonds, with the general formula CₙH₂ₙ₊₂.

Alkene

An unsaturated hydrocarbon containing at least one carbon-carbon double bond, with the general formula CₙH₂ₙ.

Alkyne

An unsaturated hydrocarbon containing at least one carbon-carbon triple bond, with the general formula CₙH₂ₙ₋₂.

Aromatic Compound

A cyclic organic compound containing a planar ring of atoms with delocalized pi electrons, such as benzene and its derivatives.

Carboxylic Acid

An organic compound containing a carboxyl group (−COOH), which can donate a proton, making it a weak organic acid.

Chirality

The geometric property of a molecule whose mirror image cannot be superimposed on itself, typically arising from a carbon atom bonded to four different substituents.

Condensation Reaction

A reaction in which two molecules combine to form a larger molecule with the simultaneous loss of a small molecule, typically water or an alcohol.

Electrophile

An electron-deficient species that accepts an electron pair to form a new bond, commonly encountered as a reactive intermediate in organic reaction mechanisms.

Enantiomer

A pair of stereoisomers that are non-superimposable mirror images of each other, related like left and right hands.

Functional Group

A specific grouping of atoms within a molecule that determines the molecule's characteristic chemical reactions and properties.

Grignard Reagent

An organomagnesium halide of the general formula RMgX, widely used in organic synthesis for forming new carbon-carbon bonds by reacting with carbonyl compounds.

Hydrocarbon

An organic compound consisting entirely of hydrogen and carbon atoms, serving as the fundamental framework of organic chemistry.

Isomer

One of two or more compounds with the same molecular formula but different structural arrangements of atoms.

Ketone

An organic compound containing a carbonyl group (C=O) bonded to two carbon atoms, with the general formula RCOR'.

Monomer

A small molecule that can be bonded to other identical molecules to form a polymer through a polymerization reaction.

Nucleophile

An electron-rich species that donates an electron pair to form a new covalent bond with an electrophilic center, driving substitution and addition reactions.

Polymer

A large molecule composed of many repeating structural units called monomers, linked together by covalent bonds.

Polymerization

A chemical process in which small molecules called monomers combine to form a large macromolecule called a polymer, through either addition or condensation mechanisms.

Stereoisomer

Isomers that have the same connectivity of atoms but differ in the three-dimensional spatial arrangement of those atoms.

Tautomerism

A dynamic equilibrium between two structural isomers called tautomers that readily interconvert by migration of a proton and rearrangement of bonding electrons.

Inorganic Chemistry

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Concepts related to non-carbon-based compounds, coordination chemistry, transition metals, and crystal field theory.

Actinide

Any of the 15 metallic elements with atomic numbers 89–103, characterized by the filling of 5f orbitals and significant radioactivity.

Allotrope

One of two or more structural forms of the same element in the same physical state, such as diamond and graphite for carbon.

Bioinorganic Chemistry

A field studying the role of metal ions in biological systems, including metalloenzymes, oxygen transport proteins, and metal-based drugs.

Chelation

The formation of a ring structure when a single polydentate ligand bonds to a metal ion through two or more donor atoms, producing a more stable complex than equivalent monodentate ligands.

Complex Ion

An ion formed when a metal cation bonds to one or more ligands through coordinate covalent bonds.

Coordination Compound

A compound in which a central metal atom or ion is surrounded by a set of bound molecules or ions called ligands.

Coordination Number

The number of ligand donor atoms directly bonded to the central metal atom or ion in a coordination complex.

Crystal Field Theory

A model that explains the electronic structure and color of transition metal complexes by the effect of ligands on d-orbital energy levels.

Intermetallic Compound

A solid-state compound of two or more metals with a definite stoichiometry and ordered crystal structure, distinct from a simple alloy.

Isomerism

The phenomenon in which coordination compounds have the same formula but different spatial arrangements or connectivity of ligands.

Jahn-Teller Effect

A geometric distortion of a non-linear molecular system that reduces its symmetry and energy when the electronic configuration leads to a degenerate ground state.

Lanthanide

Any of the 15 metallic elements with atomic numbers 57–71, characterized by the progressive filling of 4f electron orbitals.

Ligand

An ion or molecule that donates one or more electron pairs to a central metal atom or ion to form a coordination bond.

Ligand Field Theory

An extension of crystal field theory that incorporates molecular orbital concepts to explain bonding, electronic spectra, and magnetic properties of coordination compounds.

Main Group Element

An element in the s-block or p-block of the periodic table whose chemistry is primarily governed by valence s and p electrons.

Oxidation State

The hypothetical charge an atom would have if all bonds were completely ionic, used to track electron transfer in reactions.

Spectrochemical Series

An empirical ordering of ligands by the strength of the crystal field splitting they produce when coordinated to a metal ion, ranging from weak-field to strong-field ligands.

Trans Effect

The tendency of certain ligands in square-planar complexes to direct substitution of the ligand opposite (trans) to themselves, following a well-established strength series.

Transition Metal

An element in the d-block of the periodic table (Groups 3–12) characterized by partially filled d-orbitals and variable oxidation states.

Werner's Theory

Alfred Werner's theory that transition metals exhibit primary and secondary valences, with secondary valence directed in space to form coordination complexes.

Analytical Chemistry

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Terms describing techniques used to determine the composition and structure of matter, from spectroscopy to chromatography.

Calibration Curve

A graph of instrument response versus known concentrations of a standard, used to determine the concentration of an unknown sample.

Chromatography

A technique for separating mixtures by passing them through a medium in which the components move at different rates due to differential affinity.

Detection Limit

The lowest concentration of an analyte that can be reliably distinguished from the background noise by an analytical method.

Electrophoresis

A separation technique that uses an electric field to migrate charged molecules through a gel or solution, commonly applied to proteins and nucleic acids.

Fluorescence Spectroscopy

An analytical method that measures the intensity and wavelength of light emitted by a substance after it absorbs photons and undergoes electronic excitation.

Gas Chromatography

A separation technique in which volatile compounds are partitioned between a gaseous mobile phase and a liquid or solid stationary phase inside a heated column.

Gravimetric Analysis

A quantitative method in which the amount of an analyte is determined by measuring the mass of a precipitate or residue.

High-Performance Liquid Chromatography

A chromatographic technique that uses high pressure to push a liquid mobile phase through a column packed with fine stationary phase particles, achieving rapid and efficient separation of non-volatile compounds.

Indicator

A substance that changes color at or near the equivalence point of a titration, signaling the completion of the reaction.

Internal Standard

A known compound added to all samples and standards at a constant concentration to correct for variations in sample preparation, injection volume, and instrument response.

IR Spectroscopy

A technique that measures the absorption of infrared light by molecules, revealing characteristic functional group vibrations.

Limit of Detection

The lowest concentration of an analyte that can be reliably distinguished from a blank sample with a stated confidence level, typically defined as three times the standard deviation of the blank.

Mass Spectrometry

An analytical technique that measures the mass-to-charge ratio of ions to identify and quantify molecules in a sample.

NMR Spectroscopy

A technique exploiting the magnetic properties of certain nuclei to determine the physical and chemical environment of atoms in a molecule.

Qualitative Analysis

Analytical methods that identify what substances are present in a sample without measuring their exact amounts.

Quantitative Analysis

Analytical methods that determine how much of a substance is present in a sample, yielding a numerical concentration or mass.

Spectroscopy

The study of the interaction between matter and electromagnetic radiation, used to identify substances and determine their structure.

Standard Solution

A solution of precisely known concentration used as a reference for quantitative analysis, typically prepared from a primary standard.

Titration

A quantitative analytical method in which a solution of known concentration is added to a solution of unknown concentration until the reaction is complete.

Titration Curve

A graphical plot of pH or potential versus the volume of titrant added during a titration, used to identify the equivalence point and the buffering regions of the analyte.

UV-Vis Spectroscopy

A technique measuring the absorption of ultraviolet and visible light by a substance, useful for determining concentration and electronic transitions.

Volumetric Analysis

A quantitative method based on measuring the volume of a solution of known concentration required to react completely with the analyte.

Electrochemistry

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Concepts at the intersection of chemistry and electricity, including galvanic cells, electrolysis, and electrode potentials.

Anode

The electrode at which oxidation occurs in an electrochemical cell, releasing electrons into the external circuit.

Battery

A device consisting of one or more electrochemical cells that converts stored chemical energy into electrical energy.

Cathode

The electrode at which reduction occurs in an electrochemical cell, accepting electrons from the external circuit.

Concentration Cell

A galvanic cell in which both half-cells contain the same electrodes and electrolyte but at different concentrations, generating voltage from the concentration gradient.

Corrosion

The gradual electrochemical degradation of metals through oxidation by environmental agents such as water, oxygen, and acids.

Electrochemical Series

A ranking of elements or half-reactions by their standard electrode potentials, predicting the direction of electron transfer in redox reactions.

Electrode

A conductor through which electric current enters or leaves an electrolyte, a gas, a vacuum, or a solid in an electrochemical cell.

Electrolysis

A process in which electrical energy drives a non-spontaneous chemical reaction, decomposing compounds by passing direct current through an electrolyte.

Electrolyte

A substance that dissociates into ions when dissolved in a solvent, enabling the solution to conduct electric current.

Electrolytic Cell

An electrochemical cell that uses an external electrical energy source to drive a non-spontaneous chemical reaction.

Electroplating

The process of depositing a thin layer of metal onto a surface by electrolysis, used for corrosion protection and decoration.

Faradaic Efficiency

The ratio of the number of moles of a desired product formed to the total moles of electrons transferred in an electrochemical process, expressed as a percentage.

Faraday's Law

The law stating that the amount of substance produced at an electrode during electrolysis is directly proportional to the total electric charge passed.

Fuel Cell

An electrochemical cell that continuously converts the chemical energy of a fuel (such as hydrogen) and an oxidant into electricity.

Galvanic Cell

An electrochemical cell that produces electrical energy from a spontaneous redox reaction occurring at two separate electrodes.

Half-Reaction

Either the oxidation or reduction component of a redox reaction, showing explicitly the electrons lost or gained.

Nernst Equation

An equation relating the electrode potential of a half-cell to the standard potential, temperature, and activities of the chemical species involved.

Overpotential

The additional voltage beyond the thermodynamic equilibrium potential that must be applied to drive an electrochemical reaction at a measurable rate.

Standard Electrode Potential

The voltage of a half-cell measured under standard conditions relative to the standard hydrogen electrode (SHE), set at 0 V.

Nuclear Chemistry

21

Terms related to nuclear reactions, radioactive decay, fission, fusion, and the applications of radioisotopes.

Alpha Decay

A type of radioactive decay in which an atomic nucleus emits an alpha particle (two protons and two neutrons), reducing its mass number by 4.

Beta Decay

A type of radioactive decay in which a neutron converts to a proton (or vice versa), emitting a beta particle (electron or positron) and a neutrino.

Binding Energy

The energy required to disassemble a nucleus into its component protons and neutrons, a measure of nuclear stability.

Breeder Reactor

A nuclear reactor designed to produce more fissile material than it consumes by converting fertile isotopes such as uranium-238 into fissile plutonium-239.

Carbon Dating

A radiometric dating method that uses the known half-life of carbon-14 to determine the age of organic materials up to approximately 50,000 years old.

Chain Reaction

A self-sustaining sequence of nuclear fission events in which each fission produces neutrons that trigger further fissions.

Critical Mass

The minimum amount of fissile material needed to sustain a self-propagating nuclear chain reaction.

Decay Series

A sequence of successive radioactive decays by which an unstable nuclide transforms through intermediate isotopes until reaching a stable end product.

Fission

A nuclear reaction in which a heavy nucleus splits into two or more lighter nuclei, releasing a large amount of energy and additional neutrons.

Fusion

A nuclear reaction in which two light nuclei combine to form a heavier nucleus, releasing enormous energy as in the Sun.

Gamma Radiation

High-energy electromagnetic radiation emitted from an excited nucleus, carrying no charge or mass and having the highest penetrating power.

Isotope Enrichment

The process of increasing the proportion of a specific isotope in a sample relative to other isotopes of the same element, commonly applied to uranium-235 for nuclear fuel.

Mass Defect

The difference between the mass of a nucleus and the sum of the masses of its individual protons and neutrons, equivalent to the nuclear binding energy via E=mc².

Nuclear Binding Energy Curve

A plot of binding energy per nucleon versus mass number that peaks near iron-56, explaining why energy is released by both fission of heavy nuclei and fusion of light nuclei.

Nuclear Half-Life

The time required for half of the atoms in a radioactive sample to undergo decay, characteristic of each radioactive isotope.

Nuclear Reactor

A device in which controlled nuclear fission chain reactions are maintained to produce energy for electricity generation or research.

Radioactive Decay

The spontaneous transformation of an unstable atomic nucleus into a more stable configuration by emitting radiation in the form of alpha particles, beta particles, or gamma rays.

Radioactive Isotope

An isotope with an unstable nucleus that undergoes radioactive decay, used in medicine, industry, and scientific research.

Radioactivity

The spontaneous emission of particles or radiation from an unstable atomic nucleus as it transforms into a more stable configuration.

Radiopharmaceutical

A radioactive compound administered to patients for medical imaging diagnostics or targeted radiation therapy, such as technetium-99m for bone scans.

Transmutation

The conversion of one element into another through nuclear reactions, either spontaneous radioactive decay or artificial bombardment.

Biochemistry

20

Concepts covering the chemistry of living organisms, including proteins, nucleic acids, enzymes, and metabolic pathways.

Active Site

The specific region on an enzyme where the substrate binds and the catalytic reaction takes place.

Amino Acid

An organic molecule containing both an amino group (−NH₂) and a carboxyl group (−COOH), serving as the building block of proteins.

ATP

Adenosine triphosphate, the primary energy currency of cells that provides energy for biochemical reactions by hydrolysis of its phosphoanhydride bonds.

Carbohydrate

An organic compound with the general formula Cₙ(H₂O)ₙ, serving as a primary energy source and structural component in living organisms.

Cellular Respiration

The metabolic process by which cells oxidize organic molecules to produce ATP, carbon dioxide, and water.

Coenzyme

A small non-protein organic molecule that assists an enzyme in catalyzing a reaction, often derived from vitamins.

Cofactor

A non-protein chemical compound or metallic ion required for an enzyme's catalytic activity, which may be tightly bound (prosthetic group) or loosely associated (cosubstrate).

Denaturation

The loss of a protein's native three-dimensional structure due to external stress such as heat, pH changes, or chemical agents, resulting in loss of function.

DNA

Deoxyribonucleic acid, the double-helix molecule that carries genetic information encoding the instructions for protein synthesis.

Enzyme (Biochemistry)

A protein that catalyzes biochemical reactions with remarkable specificity, lowering activation energy without being permanently altered.

Glycolysis

A central metabolic pathway that converts one molecule of glucose into two molecules of pyruvate, yielding a net gain of two ATP and two NADH molecules.

Krebs Cycle

A cyclic series of enzymatic reactions in the mitochondrial matrix that oxidizes acetyl-CoA to CO2, generating NADH, FADH2, and GTP for downstream energy production.

Lipid

A broad class of hydrophobic organic molecules including fats, oils, and phospholipids that store energy and form cell membranes.

Metabolism

The complete set of chemical reactions occurring in living organisms to maintain life, divided into catabolism and anabolism.

Michaelis-Menten Kinetics

A model describing the rate of enzyme-catalyzed reactions as a function of substrate concentration, characterized by the parameters Vmax and Km.

Photosynthesis

The process by which green plants and other organisms convert light energy, water, and carbon dioxide into glucose and oxygen.

Protein

A large biological polymer composed of amino acid chains folded into a specific three-dimensional shape that determines its function.

Protein Folding

The physical process by which a polypeptide chain acquires its functional three-dimensional structure, driven by hydrophobic interactions, hydrogen bonds, and disulfide bridges.

RNA

Ribonucleic acid, a single-stranded nucleic acid involved in translating genetic information from DNA into proteins.

Substrate

The specific molecule upon which an enzyme acts, binding at the active site to undergo a chemical transformation.

Materials Science

19

Terms describing the properties and applications of solid-state materials, from semiconductors to nanomaterials and polymers.