Glossary
Definitions of nuclear physics, atomic structure, and chemistry terms used throughout this reference. Click any term to get a direct link.
A
Atomic Number (Z)
The number of protons in the nucleus of an atom. This uniquely identifies each element — hydrogen has Z=1, helium Z=2, and so on up to oganesson at Z=118. All atoms of the same element share the same atomic number.
Alpha Decay (α)
Emission of a helium-4 nucleus (2 protons + 2 neutrons). The parent loses 2 protons and 2 neutrons: Z → Z−2, A → A−4. Common in heavy nuclei (Z > 82). Example: ²³⁸U → ²³⁴Th + α.
Atomic Radius
The distance from the nucleus to the outermost electron shell. Empirical atomic radii are measured from crystallographic data. Radii generally decrease across a period (more protons pull electrons closer) and increase down a group (more electron shells).
Atomic Mass
The weighted average mass of an element's naturally occurring isotopes, expressed in daltons (Da) or atomic mass units (u). For example, chlorine's atomic mass is ~35.45 Da because it's a mixture of ³⁵Cl (75.8%) and ³⁷Cl (24.2%). For radioactive elements without stable isotopes, the mass of the longest-lived isotope is typically listed.
Atomic Mass Unit (u / Da)
A unit of mass defined as exactly 1/12 the mass of a carbon-12 atom, approximately 1.6605 × 10⁻²⁷ kg. Also called the dalton (Da). One proton or neutron has a mass of roughly 1 u, so a nucleus with mass number A has a mass near A u (slightly less, due to binding energy).
Alkali Metal
Elements in Group 1 (excluding hydrogen): lithium, sodium, potassium, rubidium, cesium, and francium. They have one valence electron, making them highly reactive — they react vigorously with water and must be stored under oil. Reactivity increases down the group. They are soft, low-density metals with low melting points.
Alkaline Earth Metal
Elements in Group 2: beryllium, magnesium, calcium, strontium, barium, and radium. They have two valence electrons and are reactive, though less so than alkali metals. They form +2 ions and produce alkaline (basic) solutions when their oxides dissolve in water — hence the name.
Actinide
The 15 elements from actinium (Z=89) to lawrencium (Z=103), which fill the 5f subshell. All actinides are radioactive. The first four (Ac, Th, Pa, U) occur naturally; the rest are synthetic. Uranium and plutonium are used in nuclear energy and weapons. The heavier actinides have very short half-lives and are produced in particle accelerators.
Aufbau Principle
The rule that electrons fill orbitals starting from the lowest energy level upward. The filling order (1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, ...) follows the n+l rule: orbitals with lower n+l fill first, and for equal n+l, lower n fills first. There are notable exceptions — chromium and copper prefer a half-filled or fully-filled 3d subshell, borrowing an electron from 4s.
Aromaticity
A property of cyclic, planar molecules with a continuous ring of overlapping p-orbitals containing 4n+2 π electrons (Hückel's rule). Aromatic compounds like benzene (6 π electrons) are unusually stable. The delocalized electrons create equal bond lengths and special reactivity patterns. About two-thirds of known organic compounds contain at least one aromatic ring.
Amino Acid
An organic molecule containing both an amino group (–NH₂) and a carboxylic acid group (–COOH). The 20 standard amino acids are the building blocks of proteins, linked by peptide bonds. Each has a unique side chain (R group) that determines its chemical properties. Nine are 'essential' — the human body cannot synthesize them and must obtain them from food.
B
Binding Energy
The energy needed to completely disassemble a nucleus into individual protons and neutrons. Higher binding energy means a more tightly-bound, more stable nucleus. It's often expressed per nucleon (binding energy / A). The peak occurs near iron-56 and nickel-62 — the most stable nuclei.
Beta-Minus Decay (β⁻)
A neutron converts to a proton, emitting an electron and antineutrino. Z → Z+1, A stays the same. Occurs in neutron-rich nuclei. Example: ¹⁴C → ¹⁴N + e⁻ + ν̄ₑ.
Beta-Plus Decay (β⁺)
A proton converts to a neutron, emitting a positron and neutrino. Z → Z−1, A stays the same. Occurs in proton-rich nuclei. Only possible when the mass difference exceeds 1.022 MeV (twice the electron rest mass). Example: ²²Na → ²²Ne + e⁺ + νₑ.
Beta-Delayed Neutron Emission (β⁻n)
A two-step process: the nucleus first undergoes beta-minus decay, and the resulting daughter is left in such a highly excited state that it promptly emits one or more neutrons. Very common in neutron-rich nuclides — over 600 known cases. Critical in nuclear reactor control, where delayed neutrons slow the chain reaction enough to allow human intervention.
Beta-Delayed Proton Emission (β⁺p)
The nucleus undergoes beta-plus decay or electron capture, and the excited daughter immediately emits a proton. Common in proton-rich nuclides far from stability. Over 270 known cases. The net effect is Z → Z−2, A → A−1.
Block (s, p, d, f)
Groups of elements classified by the type of orbital their outermost electrons occupy. s-block: groups 1-2 + He, p-block: groups 13-18, d-block: transition metals (groups 3-12), f-block: lanthanides and actinides.
Boiling Point
The temperature at which a substance transitions from liquid to gas at standard pressure (1 atm). Expressed in kelvin (K). Rhenium has the highest boiling point at 5,903 K. Noble gases have very low boiling points because they only interact through weak van der Waals forces.
Branching Ratio
The fraction (as a percentage) of decays that follow a particular decay mode when multiple modes are possible. For example, ²¹²Bi decays by beta-minus 64% of the time and by alpha 36% of the time. Branching ratios must sum to 100% for a given nuclide.
Binding Energy Curve
A plot of binding energy per nucleon vs. mass number (A). It peaks near A ≈ 56–62 (iron and nickel), meaning these nuclei are the most tightly bound. Lighter nuclei can release energy by fusion (moving up the curve), while heavier nuclei release energy by fission (also moving toward the peak). This curve explains why stars fuse elements up to iron and why uranium can power fission reactors.
C
Charge Radius
The root-mean-square radius of the nuclear charge distribution, measured in femtometers (fm, 10⁻¹⁵ m). Roughly, nuclear radii follow R ≈ 1.2 × A^(1/3) fm. Charge radii reveal nuclear structure details like deformation.
Cluster Decay (Exotic Decay)
Emission of a fragment heavier than an alpha particle but lighter than a fission fragment — typically ¹⁴C, ²⁰Ne, ²⁴Ne, ²⁸Mg, or ³⁴Si. Extremely rare, with branching ratios of 10⁻⁹ to 10⁻¹⁶ relative to alpha decay. First observed in 1984 when ²²³Ra was found to emit ¹⁴C nuclei.
Covalent Radius
Half the distance between two bonded atoms of the same element in a covalent bond. Typically smaller than the atomic radius because bonding pulls atoms closer together.
Crystal Structure
The geometric arrangement of atoms in a solid. Common types include FCC (face-centered cubic, e.g. copper, gold), BCC (body-centered cubic, e.g. iron at room temperature), and HCP (hexagonal close-packed, e.g. titanium). The structure affects physical properties like density and ductility.
Covalent Bond
A chemical bond formed when two atoms share one or more pairs of electrons. Single bonds share one pair (σ bond), double bonds share two pairs (one σ + one π), and triple bonds share three pairs (one σ + two π). Bond strength increases with bond order: C–C (346 kJ/mol) < C=C (614) < C≡C (839). Most bonds in organic molecules are covalent.
Chirality
The property of a molecule that is non-superimposable on its mirror image, like left and right hands. Usually caused by a carbon atom bonded to four different groups (a stereocenter). Enantiomers have identical physical properties except they rotate plane-polarized light in opposite directions. In biology, chirality is critical — enzymes distinguish between enantiomers, which is why (S)-ibuprofen works but (R)-ibuprofen doesn't.
Catalyst
A substance that increases the rate of a chemical reaction without being consumed. It works by providing an alternative reaction pathway with lower activation energy. Enzymes are biological catalysts — they can accelerate reactions by factors of 10⁶ to 10¹⁷. Industrial catalysts (platinum, palladium, zeolites) are essential for petroleum refining, fertilizer production, and emissions control.
D
Decay Mode
The mechanism by which a radioactive nuclide transforms. The primary modes are alpha, beta-minus, beta-plus, electron capture, and spontaneous fission. Which mode occurs depends on the nuclide's position relative to the valley of stability.
Double Beta Decay (2β⁻ / 2β⁺ / 2EC)
Two simultaneous beta decays, changing Z by 2 while A stays the same. Occurs when single beta decay is energetically forbidden but double decay is allowed. Extremely rare — half-lives are typically 10¹⁸ to 10²⁴ years. Neutrinoless double beta decay, if observed, would prove neutrinos are their own antiparticles.
Drip Line
The boundary beyond which adding one more proton (proton drip line) or neutron (neutron drip line) results in immediate emission. Nuclides beyond the drip line are unbound — the last nucleon 'drips' off. The neutron drip line has been experimentally reached only for the lightest elements.
Density
Mass per unit volume, typically in g/cm³ for solids and liquids, or g/L for gases. Osmium is the densest naturally occurring element at 22.59 g/cm³. Density generally increases toward the middle-bottom of the periodic table and reflects how tightly atoms pack in the solid or liquid state.
Decay Chain
A sequence of radioactive decays where each unstable nuclide transforms into another until a stable nuclide is reached. The most well-known chains start from ²³⁸U (uranium series, 14 steps to ²⁰⁶Pb), ²³⁵U (actinium series), and ²³²Th (thorium series). Each step in the chain has its own decay mode and half-life.
Diatomic Nonmetal
Nonmetals that naturally form diatomic molecules (two-atom pairs) in their elemental state: hydrogen (H₂), nitrogen (N₂), oxygen (O₂), and the halogens fluorine (F₂), chlorine (Cl₂), bromine (Br₂), and iodine (I₂). The diatomic bond is often very strong — N₂ has a triple bond with a dissociation energy of 945 kJ/mol.
E
Electron Capture (EC)
An inner orbital electron is captured by the nucleus, combining with a proton to form a neutron plus a neutrino. Z → Z−1, A unchanged. Competes with β⁺ decay and is the only option when the mass difference is below 1.022 MeV.
Electronegativity
A measure of how strongly an atom attracts electrons in a chemical bond. The Pauling scale is most common, ranging from ~0.7 (cesium) to 3.98 (fluorine). It increases across a period (left to right) and decreases down a group.
Electron Affinity
The energy change when a neutral atom gains an electron to form a negative ion. More negative values mean the atom more readily accepts electrons. Noble gases and some alkaline earth metals have near-zero or positive electron affinities (they resist gaining electrons).
Electron Configuration
The arrangement of electrons in atomic orbitals, written as subshell notation like 1s² 2s² 2p⁶. The order follows the Aufbau principle (filling lowest energy first), with some exceptions for transition metals and lanthanides. Shorthand uses the preceding noble gas: [Ar] 3d¹⁰ 4s² for zinc.
Electron Shell
A set of electron orbitals with the same principal quantum number n. The first shell (n=1) holds up to 2 electrons, the second (n=2) up to 8, and so on following the 2n² rule. Each period of the periodic table corresponds to filling a new shell. The outermost occupied shell is the valence shell, which determines chemical behavior.
Electronvolt (eV / keV / MeV)
A unit of energy equal to the kinetic energy gained by an electron accelerating through 1 volt: 1 eV = 1.602 × 10⁻¹⁹ J. Nuclear binding energies are typically in MeV (millions of eV) or keV (thousands of eV) per nucleon. Particle masses can also be expressed in eV/c² via E = mc².
F
Functional Group
A specific group of atoms within a molecule that determines its chemical reactivity and properties. Common examples: hydroxyl (–OH, alcohols), carboxyl (–COOH, acids), amine (–NH₂), carbonyl (C=O, ketones/aldehydes), ester (–COO–), amide (–CONH–). The same functional group behaves similarly regardless of the rest of the molecule, making organic chemistry systematic.
G
Group
A vertical column in the periodic table. Elements in the same group share similar chemical properties because they have the same number of valence electrons. For example, Group 1 (alkali metals) all have one valence electron and are highly reactive.
H
Half-Life
The time it takes for half of a sample of radioactive atoms to decay. After one half-life, 50% remain; after two, 25%; after three, 12.5%, and so on. Half-lives range from attoseconds (10⁻¹⁸ s) for the most unstable nuclides to billions of years for nearly-stable ones like uranium-238 (4.5 billion years).
Halogen
Elements in Group 17: fluorine, chlorine, bromine, iodine, astatine, and tennessine. They have seven valence electrons and eagerly gain one more to form −1 ions, making them the most reactive nonmetals. Reactivity decreases down the group. Fluorine is the most reactive element of all. Halogens exist as diatomic molecules (F₂, Cl₂, etc.).
Hund's Rule
Within a subshell, electrons occupy empty orbitals singly (with parallel spins) before pairing up. This minimizes electron-electron repulsion and results in maximum total spin. For example, nitrogen's 2p subshell has three electrons occupying three separate orbitals (↑ ↑ ↑) rather than pairing in one (↑↓ ↑ _).
Hydrogen Bond
A strong intermolecular attraction between a hydrogen atom bonded to N, O, or F and a lone pair on another N, O, or F atom. Roughly 10–40 kJ/mol — much weaker than covalent bonds but strong enough to determine protein folding, DNA base pairing, and water's anomalous properties (high boiling point, ice floating, surface tension).
I
Isotope
Atoms of the same element (same Z) that have different numbers of neutrons (different A). For instance, deuterium (²H) and tritium (³H) are isotopes of hydrogen. Some isotopes are stable; others are radioactive and undergo decay.
Isomeric Transition (IT)
A nucleus in a metastable excited state drops to a lower energy state by emitting a gamma ray. Z and A don't change. The excited state (isomer) can have a measurable half-life, sometimes hours or years.
Ionization Energy
The energy needed to remove an electron from a neutral gaseous atom. The first ionization energy removes the outermost electron. Successive ionization energies increase as each electron is harder to remove. Large jumps between values reveal electron shell structure.
Ionic Bond
A bond formed by the electrostatic attraction between oppositely charged ions. Typically occurs when a metal (low electronegativity) transfers electrons to a nonmetal (high electronegativity). Ionic compounds like NaCl form crystal lattices, have high melting points, conduct electricity when dissolved or molten, and are generally water-soluble.
Isomer (Chemical)
Molecules with the same molecular formula but different structural arrangements. Structural isomers differ in atom connectivity (butane vs. isobutane). Stereoisomers have the same connectivity but different 3D arrangement — including enantiomers (mirror images, like left/right hands) and diastereomers. Isomers can have dramatically different properties and biological activities.
InChI
International Chemical Identifier — a textual identifier for chemical substances developed by IUPAC and NIST. Unlike SMILES, InChI produces a unique, canonical representation for each molecule through layered encoding of connectivity, charges, stereochemistry, and isotopes. InChIKey is a fixed-length 27-character hash of InChI used for database searching.
K
Kelvin (K)
The SI unit of temperature and the scale used throughout scientific data. Zero kelvin (0 K = −273.15°C) is absolute zero, where all thermal motion ceases. Water freezes at 273.15 K and boils at 373.15 K. To convert: K = °C + 273.15. Unlike Celsius or Fahrenheit, kelvin has no degree symbol and starts at a true physical zero.
Kilojoule per Mole (kJ/mol)
A unit of energy per amount of substance, used for ionization energies, electron affinities, and bond energies. One kJ/mol is the energy needed to affect one mole (6.022 × 10²³) of atoms or molecules. For reference, a typical covalent bond has a strength of 150–500 kJ/mol. To convert: 1 eV/atom = 96.485 kJ/mol.
L
Lattice Constant
The physical dimension of a unit cell in a crystal lattice, typically given in ångströms (Å, 10⁻¹⁰ m) or picometers (pm). For cubic crystals, a single value defines the cell edge length. For hexagonal crystals, two values (a and c) are needed. The lattice constant determines interatomic spacing and affects mechanical and electronic properties.
Lanthanide
The 15 elements from lanthanum (Z=57) to lutetium (Z=71), which fill the 4f electron subshell. Also called rare earth elements (along with scandium and yttrium), though most aren't actually rare. They have similar chemical properties because the 4f electrons are deeply buried and don't participate much in bonding. Used in magnets, lasers, and catalysts.
Lipinski's Rule of Five
A set of guidelines for predicting oral bioavailability of drug candidates, proposed by Christopher Lipinski in 1997. A compound is likely to be orally active if it has: molecular weight ≤ 500, LogP ≤ 5, hydrogen bond donors ≤ 5, and hydrogen bond acceptors ≤ 10. About 90% of approved oral drugs satisfy these rules. Notable exceptions include natural products and antibiotics.
M
Mass Number (A)
The total number of protons plus neutrons in a nucleus: A = Z + N. Different isotopes of the same element have different mass numbers. For example, carbon-12 has A=12 (6 protons + 6 neutrons) while carbon-14 has A=14 (6 protons + 8 neutrons).
Mass Excess
The difference between an atom's actual mass and its mass number A (in atomic mass units). It's defined as Δ = (M − A) × 931.494 MeV/c². Stable nuclei have negative mass excesses (they're lighter than the sum of their parts due to binding energy).
Magic Numbers
Specific numbers of protons or neutrons (2, 8, 20, 28, 50, 82, 126) that result in unusually stable nuclei. These correspond to complete nuclear shells, analogous to closed electron shells in atoms. Doubly-magic nuclei (magic Z and magic N) like ⁴He, ¹⁶O, ⁴⁰Ca, and ²⁰⁸Pb are exceptionally stable.
Melting Point
The temperature at which a substance transitions from solid to liquid at standard pressure. Expressed in kelvin (K). Tungsten has the highest melting point of any element at 3,695 K. Melting points generally increase with stronger interatomic bonding — metals with d-orbital bonding (transition metals) tend to have higher melting points.
Molar Heat Capacity
The amount of energy (in joules) needed to raise the temperature of one mole of a substance by one kelvin, measured at constant pressure. Written as J/(mol·K). The Dulong–Petit law predicts ~25 J/(mol·K) for most solid elements, with significant deviations for light elements like beryllium and carbon.
Metalloid
Elements with properties intermediate between metals and nonmetals: boron, silicon, germanium, arsenic, antimony, tellurium, and sometimes polonium. They are semiconductors — their electrical conductivity increases with temperature, opposite to metals. Silicon and germanium are the foundation of modern electronics.
Molecular Weight (MW)
The sum of the atomic weights of all atoms in a molecule, expressed in grams per mole (g/mol) or daltons (Da). Water has MW = 18.015 g/mol. Molecular weight affects physical properties: higher MW generally means higher boiling point, lower vapor pressure, and slower diffusion. In drug design, MW below 500 g/mol is associated with better oral absorption (Lipinski's rule).
N
Neutron Number (N)
The number of neutrons in a nucleus: N = A − Z. Neutrons provide the nuclear binding force that holds the nucleus together. Heavier elements need proportionally more neutrons to remain stable.
Nuclide
A specific combination of protons and neutrons — identified by both Z and A. While 'isotope' refers to variants of one element, 'nuclide' is the general term for any specific nuclear species. There are roughly 3,300 known nuclides.
Natural Abundance
The fraction of a particular isotope found in naturally occurring samples of that element. For example, about 99.76% of natural carbon is ¹²C, 1.1% is ¹³C, and trace amounts are ¹⁴C. Only stable (or very long-lived) isotopes have measurable natural abundances.
Neutron Emission (n)
Direct emission of a neutron. Z unchanged, A → A−1. Rare in ground states but common in excited states (delayed neutron emission after beta decay).
Nuclear Isomer
A metastable excited state of a nucleus that has a measurable half-life before decaying to the ground state (usually by gamma emission). Denoted with 'm' after the mass number, e.g. ⁹⁹ᵐTc. Some isomers are remarkably long-lived — ¹⁸⁰ᵐTa has never been observed to decay and may be effectively stable.
Nonmetal
Elements that lack metallic properties — they are poor conductors of heat and electricity, and tend to gain electrons in reactions. Includes hydrogen, carbon, nitrogen, oxygen, phosphorus, sulfur, selenium, and the halogens. Most are gases at STP; some are solids (carbon, sulfur). Nonmetals make up the upper-right region of the periodic table.
Noble Gas
Elements in Group 18: helium, neon, argon, krypton, xenon, radon, and oganesson. They have completely filled outer electron shells, making them extremely unreactive — earning them the name 'noble' (chemically aloof). They are colorless, odorless gases used in lighting, welding shielding, and cryogenics.
Nuclear Fusion
The process of combining light nuclei into heavier ones, releasing energy because the products have higher binding energy per nucleon. Powers all main-sequence stars — the Sun fuses hydrogen into helium at about 15 million K. On Earth, fusion research aims to harness this process for clean energy, most commonly using deuterium-tritium reactions.
Nuclear Fission
The splitting of a heavy nucleus into two or more lighter fragments, typically triggered by neutron absorption. Releases energy because the fragments are closer to the binding energy peak. Uranium-235 and plutonium-239 are the primary fissile materials. Each fission event releases about 200 MeV and 2–3 neutrons, enabling chain reactions.
Nuclear Shell Model
A model of nuclear structure where protons and neutrons occupy quantized energy levels (shells) within the nucleus, analogous to electron shells in atoms. Complete nuclear shells occur at the magic numbers (2, 8, 20, 28, 50, 82, 126), producing nuclei of exceptional stability. The model explains why certain nuclides have unusually high binding energies and natural abundances.
Nuclide Chart (Segrè Chart)
A plot of all known nuclides with neutron number (N) on the x-axis and proton number (Z) on the y-axis. Each square represents a specific nuclide, typically color-coded by half-life or decay mode. Stable nuclides trace the valley of stability along the diagonal. It's the nuclear physicist's equivalent of the periodic table.
O
Oxidation State
The hypothetical charge an atom would have if all its bonds were purely ionic. For example, iron commonly has oxidation states +2 (ferrous) and +3 (ferric). Transition metals often have multiple oxidation states due to accessible d-orbital electrons.
Orbital
A mathematical function describing where an electron is likely to be found. Orbitals come in types: s (spherical, holds 2e⁻), p (dumbbell, holds 6e⁻), d (clover, holds 10e⁻), and f (complex, holds 14e⁻). Each subshell fills according to the Aufbau principle and Hund's rule.
P
Proton Emission (p)
Direct emission of a proton from the nucleus. Z → Z−1, A → A−1. Occurs in extremely proton-rich nuclei beyond the proton drip line.
Period
A horizontal row in the periodic table. Elements in the same period have the same number of electron shells. There are 7 periods: period 1 has 2 elements, periods 2-3 have 8 each, periods 4-5 have 18, and periods 6-7 have 32.
Phase (State of Matter)
The physical state of a substance — solid, liquid, or gas. At standard temperature and pressure (STP: 0°C, 1 atm), most elements are solid, eleven are gases (H, He, N, O, F, Ne, Cl, Ar, Kr, Xe, Rn), and two are liquid (bromine and mercury).
Pauling Scale
The most widely used electronegativity scale, devised by Linus Pauling in 1932. Values range from 0.7 (francium) to 3.98 (fluorine). The scale is relative and dimensionless — values are derived from bond dissociation energies. A difference greater than ~1.7 between two bonded atoms generally indicates an ionic bond.
Post-Transition Metal
Metallic elements located between the transition metals and the metalloids in the periodic table: aluminium, gallium, indium, thallium, tin, lead, bismuth, nihonium, flerovium, moscovium, livermorium. They are softer, have lower melting points, and are poorer conductors than transition metals. Sometimes called 'poor metals' or 'other metals'.
Polyatomic Nonmetal
Nonmetals that form molecules or extended structures with more than two atoms in their standard state. Examples include sulfur (S₈ rings), phosphorus (P₄ tetrahedra), carbon (diamond or graphite networks), and selenium (Se₈ or helical chains). Their varied bonding arrangements give them diverse physical properties.
Pauli Exclusion Principle
No two electrons in an atom can have the same set of four quantum numbers (n, l, mₗ, mₛ). In practice, this means each orbital holds at most two electrons, and they must have opposite spins (↑↓). This principle is the reason electron shells have limited capacity and explains the structure of the periodic table.
Picometer (pm)
A unit of length equal to 10⁻¹² meters, commonly used for atomic and ionic radii. A typical atom is 100–300 pm across. One ångström (Å) equals 100 pm. For comparison, covalent bond lengths range from about 74 pm (H–H) to over 250 pm (Cs–I).
Periodic Trend
A predictable pattern in element properties that repeats across each period or down each group. Key trends include: electronegativity and ionization energy increase left-to-right and decrease top-to-bottom; atomic radius decreases left-to-right and increases top-to-bottom; metallic character increases toward the bottom-left. These trends arise from changes in nuclear charge and electron shielding.
Polar Molecule
A molecule with an uneven distribution of electron density, creating a permanent dipole moment. Arises when bond dipoles don't cancel due to molecular geometry — water (bent) is polar but CO₂ (linear) is not, despite both having polar bonds. Polar molecules dissolve in polar solvents (like dissolves like) and are deflected by electric fields.
Peptide Bond
A covalent bond between the carboxyl group of one amino acid and the amino group of another, formed by a condensation reaction (losing H₂O). The bond has partial double-bond character due to resonance, making it planar and rigid. Chains of amino acids linked by peptide bonds form peptides (< 50 residues) and proteins (≥ 50 residues). The ribosome catalyzes peptide bond formation during translation.
pH
A logarithmic scale measuring the acidity or basicity of a solution: pH = −log₁₀[H⁺]. Ranges from 0 (very acidic) to 14 (very basic), with 7 being neutral. Each unit represents a 10× change in H⁺ concentration. Blood pH is tightly regulated at 7.35–7.45 — deviations of just 0.2 units can be life-threatening.
Polymer
A large molecule made of repeating structural units (monomers) linked by covalent bonds. Natural polymers include proteins (amino acid monomers), DNA (nucleotide monomers), cellulose (glucose monomers), and rubber (isoprene). Synthetic polymers include polyethylene, nylon, PVC, and polystyrene. Properties depend on chain length, branching, and monomer chemistry.
Q
Q-Value
The energy released (or absorbed) in a nuclear reaction or decay. A positive Q-value means the reaction releases energy and can occur spontaneously. Qα is the Q-value for alpha decay; Qβ⁻ for beta-minus decay; Qε for electron capture.
Quantum Numbers
Four numbers that uniquely describe each electron in an atom. Principal (n = 1, 2, 3...) determines the shell and energy. Angular momentum (l = 0 to n−1) determines the subshell shape (s, p, d, f). Magnetic (mₗ = −l to +l) determines the orbital orientation. Spin (mₛ = +½ or −½) determines the electron's intrinsic spin direction.
R
Radioactivity
The spontaneous emission of particles or energy from an unstable atomic nucleus. The three classical types are alpha (helium nuclei), beta (electrons or positrons), and gamma (high-energy photons) radiation. Radioactivity was discovered by Henri Becquerel in 1896. The rate of decay is characterized by the half-life.
S
Spin-Parity (Jᵖ)
A quantum mechanical property describing the nucleus's total angular momentum (spin J) and its behavior under spatial inversion (parity π, either + or −). For example, ⁴He has Jᵖ = 0⁺ (no spin, even parity). These determine what decay modes and reactions are allowed.
Separation Energy
The energy required to remove one nucleon from a nucleus. Neutron separation energy (Sn) is the energy to remove one neutron; proton separation energy (Sp) removes one proton. Low separation energies indicate that the nucleus is near the drip line and weakly bound.
Spontaneous Fission (SF)
The nucleus splits into two (or more) lighter fragments plus neutrons. Dominant decay mode for very heavy elements (Z ≥ 100). Example: ²⁵²Cf undergoes SF about 3% of the time.
Standard Temperature and Pressure (STP)
A reference condition defined as 0°C (273.15 K) and 1 atmosphere (101.325 kPa). Used as the baseline when reporting properties like phase, density of gases, and molar volume. Some references use 25°C instead — always check which convention is being used.
Stable Isotope
An isotope that does not undergo radioactive decay, or has a half-life so long that no decay has been observed. There are about 254 known stable nuclides. Elements up to lead (Z=82) have at least one stable isotope, except for technetium (Z=43) and promethium (Z=61). Tin has the most stable isotopes at 10.
Subshell
A subdivision of an electron shell defined by the angular momentum quantum number l. Within a shell of principal quantum number n, subshells are labeled s (l=0), p (l=1), d (l=2), and f (l=3). Each subshell contains 2l+1 orbitals: s has 1, p has 3, d has 5, and f has 7. The orbital diagram shows how electrons fill these subshells.
SMILES
Simplified Molecular-Input Line-Entry System — a text notation that encodes molecular structure as a string. Atoms are represented by their symbols, bonds by symbols (= for double, # for triple), branches by parentheses, and rings by digits. For example: ethanol is CCO, benzene is c1ccccc1, aspirin is CC(=O)Oc1ccccc1C(=O)O. Widely used in cheminformatics databases and software.
T
Transition Metal
Elements in Groups 3–12 of the periodic table, characterized by partially filled d-orbitals. They typically have multiple oxidation states, form colored compounds, and are good catalysts. Most are hard, dense metals with high melting points. Iron, copper, silver, gold, and platinum are familiar examples.
V
Van der Waals Radius
Half the distance between the nuclei of two non-bonded atoms at their closest stable approach. Represents the effective 'size' of an atom when it's not chemically bonded. Always larger than the covalent radius.
Valley of Stability
The region on the nuclide chart where stable nuclides exist. For light elements (Z < 20), stable nuclides have roughly equal protons and neutrons (N ≈ Z). For heavier elements, stable nuclides become increasingly neutron-rich (N > Z) because extra neutrons help overcome proton-proton electrostatic repulsion.
Valence Electron
An electron in the outermost shell of an atom, responsible for chemical bonding and reactivity. Elements in the same group have the same number of valence electrons, which is why they show similar chemistry. For main-group elements, the group number (1–8) equals the number of valence electrons.
Terms defined here cover nuclear physics, atomic structure, and general chemistry at an introductory level. For rigorous definitions, consult the IAEA Nuclear Data Services or IUPAC Gold Book.