Atomic Structure & Properties: every key term you need (+ practice quiz)

More AP Chemistry guides

• Compound Structure & Properties
• Intermolecular Forces & Properties
• Chemical Reactions
• Kinetics

Atom

The smallest unit of an element that retains its properties. Made of a dense nucleus (protons + neutrons) surrounded by electrons.

Proton

A positively charged particle in the nucleus. The number of protons (atomic number, Z) defines the element.

Neutron

A neutral particle in the nucleus. Changing neutron number changes the isotope but not the element.

Electron

A negatively charged particle occupying orbitals around the nucleus. Electrons are involved in bonding and chemical reactions.

Atomic Number (Z)

The number of protons in an atom's nucleus; it identifies the element and equals the electron count in a neutral atom.

Mass Number (A)

The total number of protons plus neutrons in a nucleus. Neutrons = A − Z.

Isotopes

Atoms of the same element (same protons) with different numbers of neutrons, and thus different masses but nearly identical chemistry.

Average Atomic Mass

The weighted average of an element's isotope masses, weighted by their natural abundances. It's the value on the periodic table.

Calculating Average Atomic Mass

Multiply each isotope's mass by its fractional abundance and sum: Σ(mass × fraction). Abundant isotopes pull the average toward their mass.

Mass Spectrometry

A technique that ionizes atoms/molecules and separates them by mass-to-charge ratio, revealing each isotope's mass and relative abundance.

Reading a Mass Spectrum

Peaks show isotope masses (x-axis) and relative abundance (y-axis). Taller peaks = more abundant isotopes; use them to compute average atomic mass.

The Mole

The chemist's counting unit: 6.022 × 10²³ particles (Avogadro's number). It bridges the atomic scale and lab-measurable grams.

Avogadro's Number

6.022 × 10²³ — the number of particles in one mole. Connects number of particles to moles.

Molar Mass

The mass of one mole of a substance (g/mol), numerically equal to the atomic/formula mass in amu. Used to convert grams ↔ moles.

Moles ↔ Grams

grams = moles × molar mass; moles = grams ÷ molar mass. Molar mass is the conversion factor.

Moles ↔ Particles

particles = moles × 6.022 × 10²³. The mole lets you count atoms by weighing.

Percent Composition

The mass percent of each element in a compound: (mass of element ÷ molar mass of compound) × 100. Reveals a compound's makeup.

Empirical Formula

The simplest whole-number ratio of atoms in a compound (e.g., CH₂O). Found from percent composition by converting to moles and simplifying.

Molecular Formula

The actual number of atoms per molecule (e.g., C₆H₁₂O₆). A whole-number multiple of the empirical formula, found using molar mass.

Finding an Empirical Formula

Assume 100 g → grams = percent → convert each to moles → divide all by the smallest → round to whole numbers for the atom ratio.

Pure Substance vs. Mixture

A pure substance has a fixed composition (element or compound); a mixture is two or more substances physically combined in variable proportions.

Analyzing Mixtures

A mixture's mass spectrum or composition data reflects a combination of its components' signatures, weighted by amount — not a single fixed formula.

Coulomb's Law

The force between charges: F ∝ (q₁q₂)/r². Attraction increases with larger charges and smaller distance — the basis of all atomic attractions.

Coulombic Attraction in Atoms

Electrons are held by attraction to the positive nucleus. Closer electrons and higher nuclear charge feel stronger attraction (lower energy).

Energy Levels (Shells)

Electrons occupy discrete energy levels (n = 1, 2, 3…). Higher n means farther from the nucleus, higher energy, and weaker attraction.

Subshells (s, p, d, f)

Each energy level contains subshells: s (1 orbital), p (3), d (5), f (7). They differ in shape and energy.

Orbitals

Regions of space where an electron is likely to be found. Each orbital holds a maximum of 2 electrons with opposite spins.

Electron Capacity

s holds 2, p holds 6, d holds 10, f holds 14 electrons. Level n holds up to 2n² electrons.

Aufbau Principle

Electrons fill the lowest-energy orbitals first, building up the configuration from the ground state.

Pauli Exclusion Principle

No two electrons in an atom can have the same set of quantum numbers; an orbital holds at most 2 electrons with opposite spins.

Hund's Rule

Electrons fill degenerate (equal-energy) orbitals singly, with parallel spins, before pairing up — minimizing repulsion.

Electron Configuration

The arrangement of electrons in orbitals (e.g., oxygen: 1s²2s²2p⁴). Order follows the Aufbau diagram.

Noble-Gas (Condensed) Configuration

Abbreviate core electrons with the previous noble gas in brackets (e.g., Na = [Ne]3s¹). Highlights valence electrons.

Valence Electrons

The outermost-shell electrons that participate in bonding. An element's group number (main group) gives its valence electron count.

Core Electrons

Inner-shell electrons not involved in bonding; they shield valence electrons from the full nuclear charge.

Ground State vs. Excited State

Ground state = lowest-energy arrangement (Aufbau). Excited state = an electron has absorbed energy and jumped to a higher orbital.

Photoelectron Spectroscopy (PES)

A technique that measures the energy needed to remove electrons from each subshell, mapping an atom's electron energy structure.

Reading a PES Spectrum

Each peak = a subshell. Peak position (binding energy) shows how tightly electrons are held; peak height shows how many electrons are in that subshell.

PES and Nuclear Attraction

Higher binding energy (peaks farther left, higher energy) means electrons closer to the nucleus, held more tightly (e.g., 1s > 2s > 2p).

Effective Nuclear Charge (Zeff)

The net positive charge a valence electron actually feels after inner electrons shield it. Higher Zeff pulls electrons in more strongly.

Shielding

Core electrons repel and 'block' outer electrons from the full nuclear charge, reducing the attraction valence electrons feel.

Zeff Across a Period

Zeff increases left to right (more protons, same shielding), pulling electrons closer — the driver of periodic trends.

Periodic Law

When elements are arranged by atomic number, their properties repeat periodically — the basis of the periodic table's structure.

Groups vs. Periods

Groups (columns) share valence-electron count and similar chemistry; periods (rows) fill the same principal energy level.

Atomic Radius

The size of an atom. It decreases across a period (rising Zeff pulls electrons in) and increases down a group (more energy levels).

Atomic Radius Trend Explained

Across a period, more protons pull the same shell tighter (smaller). Down a group, electrons occupy higher shells farther out (larger).

Ionization Energy

The energy required to remove an electron from a gaseous atom. Higher when electrons are held more tightly.

Ionization Energy Trend

Increases across a period (higher Zeff) and decreases down a group (valence electrons farther out, easier to remove).

Successive Ionization Energies

Each electron removed requires more energy; a huge jump occurs when you start removing core electrons — evidence of shell structure.

Electron Affinity

The energy change when an atom gains an electron. More negative (favorable) toward the upper right (nonmetals want electrons).

Electronegativity

An atom's tendency to attract shared electrons in a bond. Increases across a period and up a group; fluorine is highest.

Electronegativity Trend

Rising Zeff and smaller radius toward the top-right make those atoms pull bonding electrons hardest (excluding noble gases).

Metallic Character

The tendency to lose electrons. Increases down and to the left; metals have low ionization energies and electronegativities.

Ionic Radius

Cations (lost electrons) are smaller than their atoms; anions (gained electrons) are larger. Charge changes electron-electron repulsion and Zeff per electron.

Valence Electrons and the Periodic Table

Group 1 = 1 valence e⁻, Group 2 = 2, Groups 13–18 = 3–8. Valence count predicts bonding behavior and periodic trends.

s, p, d, f Blocks

The periodic table divides into blocks by which subshell fills last: s-block (Groups 1–2), p-block (13–18), d-block (transition metals), f-block (lanthanides/actinides).

Quantized Energy

Electrons can only have specific allowed energies. Transitions between levels absorb or emit photons of exact energies — the origin of line spectra.

Line Spectra

Elements emit light at specific wavelengths as excited electrons fall to lower levels — a 'fingerprint' revealing quantized energy levels.

Photon Energy (E = hν)

A photon's energy is proportional to its frequency (E = hν = hc/λ). Higher frequency = shorter wavelength = more energy.

Bohr Model

Pictures electrons in fixed circular orbits of quantized energy. It explains hydrogen's spectrum but is superseded by the quantum (orbital) model.

Quantum Mechanical Model

Describes electrons as probability clouds (orbitals) rather than fixed paths. The modern, accurate model of the atom.

Conservation of Mass

Matter is neither created nor destroyed in a reaction; total mass of reactants equals total mass of products — the foundation of stoichiometry.

Amu (Atomic Mass Unit)

A unit of atomic mass defined as 1/12 the mass of a carbon-12 atom. Numerically equal to molar mass in g/mol.

Why Isotopes Have the Same Chemistry

Chemical behavior depends on electrons, which depend on protons (Z). Isotopes differ only in neutrons, so their chemistry is essentially identical.

Trends Summary (Toward F)

Toward the upper-right (excluding noble gases): smaller radius, higher ionization energy, higher electronegativity — all driven by increasing Zeff.

Trends Summary (Toward Fr)

Toward the lower-left: larger radius, lower ionization energy, more metallic character — electrons farther out and more shielded.

Coulomb's Law and Periodic Trends

Every periodic trend traces back to Coulomb's law: the balance of nuclear charge, distance, and shielding determines how tightly electrons are held.

Dimensional Analysis

The method of using unit conversion factors (like molar mass and Avogadro's number) to solve quantitative problems — the workhorse of chemistry math.