Unit 1 Foundations: How Electron Arrangement Explains Atomic Behavior

Atom

The smallest unit of an element that retains that element’s identity in chemical reactions.

Valence electrons

The outermost electrons (typically in the highest n level for main-group elements) that primarily determine bonding and chemical reactivity.

Subatomic particles

The three main particles in atoms: protons and neutrons in the nucleus, and electrons in regions of space around the nucleus.

Atomic number (Z)

The number of protons in the nucleus; it defines the element.

Neutral atom

An atom with equal numbers of protons and electrons (net charge 0).

Ion

A charged species formed when an atom gains or loses electrons; the number of protons stays the same.

Isotope

Atoms of the same element (same Z) that have different numbers of neutrons, leading to different masses.

Mass number (A)

The total number of protons plus neutrons in the nucleus.

Weighted average atomic mass

The periodic table “atomic mass” value, calculated from the natural abundances of an element’s isotopes; usually not a whole number.

Coulomb’s law

The electrostatic force relationship: attraction/repulsion increases with charge and decreases with the square of distance (F = k(q1q2)/r^2).

Nuclear charge

The positive charge of the nucleus (set by the number of protons), which generally increases electron–nucleus attraction as it increases.

Quantum mechanical model

The modern model of the atom in which electrons are described by probabilities (orbitals) rather than fixed planetary paths.

Orbital

A region of space where there is a high probability of finding an electron; a probability distribution, not a physical track.

Principal quantum number (n)

The number that labels principal energy levels (shells); higher n generally means higher energy and larger average distance from the nucleus.

Sublevel (subshell)

A division within a principal energy level labeled s, p, d, or f, containing one or more orbitals.

Electron capacity of sublevels

Maximum electrons per sublevel: s holds 2, p holds 6, d holds 10, f holds 14 (based on number of orbitals).

Aufbau principle

Electrons occupy the lowest-energy orbitals available first when building ground-state electron configurations.

Pauli exclusion principle

An orbital can hold at most two electrons, and they must have opposite spins.

Hund’s rule

Within a set of equal-energy orbitals, electrons occupy orbitals singly (with parallel spins) before pairing.

Electron configuration

A shorthand description of where electrons are located in an atom, written by sublevel and electron count (e.g., 1s² 2s² 2p⁴).

Orbital diagram

A box/line-and-arrow representation of orbitals and electron spins used to show pairing and unpaired electrons and apply Pauli and Hund correctly.

Noble-gas notation

A condensed electron configuration that replaces core electrons with the symbol of the previous noble gas (e.g., Na: [Ne] 3s¹).

Core electrons

Inner electrons that generally do not participate in bonding and are typically held more tightly than valence electrons.

Transition-metal cation electron removal

For many transition metals, electrons are removed from the highest principal energy level first (often the ns electrons before (n−1)d) when forming cations.

Photoelectron spectroscopy (PES)

A technique that measures energies of electrons ejected by high-energy light, providing evidence for quantized energy levels and electron configurations via binding energies and peak intensities.