Electronic Structure and Configuration of Elements – Key Vocabulary

A collection of key vocabulary terms and concise definitions covering historical atomic models, quantum numbers, electron-configuration principles, periodic trends, and basic bonding concepts.

Dalton’s Atomic Theory

Early 19th-century model stating that atoms are indivisible, identical for each element, conserved in reactions, and combine in fixed whole-number ratios.

Indivisibility of Atoms (Dalton)

Postulate that atoms are the smallest particles of matter and cannot be broken down by chemical means.

Fixed Ratios in Compounds

Dalton’s idea that atoms of different elements combine in simple whole-number ratios to form compounds.

Thomson’s Plum Pudding Model

1897 model in which negatively charged electrons are embedded in a diffuse sphere of positive charge.

Electron (Discovery)

Subatomic particle with negative charge discovered by J. J. Thomson using cathode-ray tubes.

Rutherford’s Gold Foil Experiment

1911 scattering experiment that revealed a small, dense, positively charged nucleus at the atom’s center.

Nuclear Model of the Atom

Rutherford’s concept of a central nucleus surrounded by electrons in mostly empty space.

Bohr’s Model

1913 planetary-like model with electrons in fixed, quantized energy levels orbiting the nucleus.

Quantized Energy Levels

Discrete electron orbits in Bohr’s model where energy is absorbed or emitted only during level transitions.

Limitations of Bohr’s Model

Fails for multi-electron atoms, ignores sublevels, fine structure, and uncertainty; explains only hydrogen well.

Schrödinger’s Wave Equation

Mathematical description treating electrons as waves; its solutions define atomic orbitals.

Quantum Mechanical Model

Current atomic model in which electrons occupy probabilistic orbitals rather than fixed paths.

Heisenberg Uncertainty Principle

Statement that the exact position and momentum of an electron cannot be known simultaneously.

Orbital

Region of space around a nucleus where the probability of finding an electron is highest.

Principal Quantum Number (n)

Quantum number indicating main energy level; higher n means higher energy and larger orbital size.

Azimuthal Quantum Number (l)

Quantum number defining orbital shape (s, p, d, f) and ranging from 0 to n − 1.

Magnetic Quantum Number (mₗ)

Quantum number specifying the orientation of an orbital; ranges from −l to +l.

Spin Quantum Number (mₛ)

Quantum number denoting electron spin direction, either +½ (spin-up) or −½ (spin-down).

Aufbau Principle

Electrons fill the lowest-energy orbitals available before occupying higher ones.

(n + l) Rule

Energy ranking rule: orbitals with lower n + l fill first; if equal, lower n fills first.

Hund’s Rule of Maximum Multiplicity

Electrons occupy degenerate orbitals singly with parallel spins before pairing up.

Pauli Exclusion Principle

No two electrons in an atom can share the same set of four quantum numbers; max two per orbital with opposite spins.

Orbital Filling Sequence

Order of increasing energy: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s …

Noble-Gas Core Notation

Shorthand electron configuration that replaces inner-shell electrons with the symbol of the nearest noble gas.

Valence Electrons

Electrons in the atom’s outermost shell responsible for bonding and chemical reactivity.

Octet Rule

Tendency of atoms to gain, lose, or share electrons to achieve eight valence electrons like a noble gas.

Ionic Bond

Chemical bond formed by transfer of electrons from a metal to a non-metal, producing cations and anions.

Covalent Bond

Bond created by sharing electron pairs between non-metal atoms.

Coordinate (Dative) Bond

Covalent bond in which both bonding electrons originate from the same atom (donor).

Metallic Bond

Attractive force between positive metal ions and a sea of delocalized electrons in a metal lattice.

Ligand

Ion or molecule with a lone pair that donates electrons to a central metal ion in a coordination complex.

Coordination Number

Number of ligand donor atoms directly bonded to the central metal ion in a complex.

Complex Ion

Charged species consisting of a central metal ion bonded to one or more ligands, e.g., [Fe(CN)₆]³⁻.

Atomic Radius Trend

Atomic size decreases across a period (↑ nuclear pull) and increases down a group (↑ energy levels).

Ionization Energy Trend

Energy required to remove an electron increases across a period and decreases down a group.

Electron Affinity

Energy change when an atom gains an electron; generally becomes more negative across a period.

Electronegativity

Ability of an atom to attract shared electrons; increases across periods and decreases down groups.

s-Block Elements

Group 1 and 2 elements in which the outermost electrons occupy s orbitals.

p Orbital Shape

Dumbbell-shaped orbital corresponding to l = 1; three orientations (px, py, pz).

d Subshell

Set of five orbitals (l = 2) with complex cloverleaf shapes; can hold 10 electrons.

f Subshell Orbitals

Seven orbitals (l = 3) of very complex shape that together hold up to 14 electrons.

Standard Electron Configuration Notation

Listing of occupied orbitals with superscripts showing electron counts, e.g., 1s² 2s² 2p⁶.