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.

Effective Nuclear Charge (Z_eff)

The net positive charge felt by a valence electron: roughly Z minus the number of core (shielding) electrons. Z_eff rises across a period because added protons are not shielded by same-shell electrons.

Shielding vs. Penetration

Core electrons shield valence electrons from the nucleus; s electrons penetrate closer to the nucleus than p or d in the same level, so 4s fills before 3d in neutral atoms.

Ionization Energy Trend

IE increases across a period (higher Z_eff, same shell) and decreases down a group (electron farther away, more shielding). Removing an electron from a filled shell costs a huge jump.

Successive Ionization Energies

IE1 < IE2 < IE3… because each removal is from an increasingly positive ion. A very large jump reveals a core electron is being removed, exposing the number of valence electrons.

IE Exceptions: Be→B and N→O

B has lower IE1 than Be because its 2p electron is higher in energy than Be's 2s. O has lower IE1 than N because the paired 2p electron in O experiences extra repulsion.

Electron Affinity

The energy change when a gaseous atom gains an electron. Halogens release the most energy (most negative EA); noble gases and group 2 have near-zero or positive values.

Atomic Radius Trend

Radius decreases across a period (rising Z_eff pulls the same shell inward) and increases down a group (added shells). Anions are larger than parents; cations are smaller.

Isoelectronic Series

Ions with the same electron count (e.g., O²⁻, F⁻, Na⁺, Mg²⁺). Radius decreases as nuclear charge increases because the same electrons are pulled by more protons.

Transition-Metal Ion Configurations

When forming cations, transition metals lose ns electrons before (n−1)d electrons: Fe = [Ar]4s²3d⁶ but Fe²⁺ = [Ar]3d⁶ and Fe³⁺ = [Ar]3d⁵.

Cr and Cu Exceptions

Cr is [Ar]4s¹3d⁵ and Cu is [Ar]4s¹3d¹⁰; half-filled and fully filled d subshells are extra stable, so one 4s electron shifts into 3d.

PES Peak Interpretation

In a photoelectron spectrum, peaks at higher binding energy are core electrons; a peak's height is proportional to the number of electrons in that subshell (a 2:6 ratio marks s vs. p).

Comparing PES Across Elements

The same subshell (e.g., 1s) appears at higher binding energy in an element with more protons, because the increased nuclear charge holds every electron more tightly.

Mass Spectrum of a Molecule

For a compound like Cl₂, peaks appear at each combination of isotopes (70, 72, 74 amu) with heights set by the product of isotope abundances.

Limiting Precision in Molar-Mass Problems

Molecular formula = empirical formula × n, where n = molar mass ÷ empirical mass. Round n to the nearest whole number; a non-integer means an arithmetic or data error.

Combustion Analysis

Burning a C/H/O compound gives CO₂ and H₂O; moles of C come from CO₂, H from H₂O, and O by mass difference, from which the empirical formula follows.

Hydrate Formulas

Heating a hydrate drives off water; the mole ratio of lost water to the anhydrous salt gives x in the formula salt·xH₂O.

Paramagnetism

Atoms or ions with unpaired electrons are attracted into a magnetic field. Hund's rule predicts which species (e.g., O with two unpaired 2p electrons) are paramagnetic.

Coulomb's Law Quantitatively

Potential energy between charges scales as E ∝ (q₁q₂)/r. Doubling one charge doubles the energy; halving the separation doubles it as well, so charge magnitude and distance trade off predictably.

Why 2p Is Higher Than 2s

A 2s electron penetrates closer to the nucleus, so it experiences a larger Z_eff and lower energy than a 2p electron in the same shell. PES confirms this: the 2s peak sits at higher binding energy.

Slater-Style Z_eff Estimates

Approximate Z_eff = Z − (number of core electrons), treating same-shell electrons as weak shielders. For sulfur, Z_eff ≈ 16 − 10 = 6 for a valence electron.

PES Peak Height Ratios

Relative peak heights equal the ratio of electrons in each subshell. A 2:2:4 pattern of peak heights identifies oxygen (1s² 2s² 2p⁴) without any labels.

Distinguishing Two Elements by PES

Compare the highest-binding-energy (1s) peak: the element with more protons holds its 1s electrons more tightly, so its 1s peak lies at greater energy even if valence peaks look similar.

Mass Spectrum of Diatomic Bromine

Br₂ shows three molecular peaks (158, 160, 162) in roughly a 1:2:1 ratio because Br-79 and Br-81 are nearly equally abundant and combine randomly in pairs.

Back-Calculating Isotopic Abundance

If average mass and two isotope masses are known, solve x·m₁ + (1−x)·m₂ = M_avg for the fraction x. For chlorine, 35x + 37(1−x) = 35.45 gives x ≈ 0.775.

Combustion Analysis Logic

All carbon in the sample ends up as CO₂ and all hydrogen as H₂O. Convert product masses to moles of C and H; oxygen mass is found by difference from the original sample mass.

Empirical vs. Molecular Multiplier

n = (molar mass) ÷ (empirical formula mass), rounded to the nearest integer. If the ratio is 3.0 for CH₂O (30 g/mol) with M = 90 g/mol, the molecular formula is C₃H₆O₃.

Non-Integer Mole Ratios

Ratios near 1.5, 1.33, or 1.25 must be multiplied by 2, 3, or 4 rather than rounded. A 1:1.5 C:O ratio becomes C₂O₃, not CO₂.

Photon Energy and Ionization

A photon can eject an electron only if hν exceeds the binding energy. Excess energy becomes the electron's kinetic energy: KE = hν − BE, the basis of PES measurement.

Rydberg-Type Reasoning

Energy gaps between hydrogen levels shrink as n increases, so transitions to n = 2 from high n converge toward a series limit — the visible Balmer lines crowd together at short wavelength.

Interpreting Successive IE Jumps

A jump of roughly a factor of four or more between IE_n and IE_(n+1) marks the transition from valence to core electrons, revealing the element's group number.

Why Group 2 IE Exceeds Group 13

Removing a Group 13 electron takes it out of a higher-energy p orbital that is partly shielded by the filled s subshell, so it is easier to remove than a Group 2 s electron despite the higher Z.

Ionic Radius in Isoelectronic Sets

Among species with identical electron counts, radius falls as nuclear charge rises: N³⁻ > O²⁻ > F⁻ > Na⁺ > Mg²⁺ > Al³⁺.

Photon Absorption vs. Emission Spectra

An absorption spectrum shows dark lines where a cool gas removes photons; an emission spectrum shows bright lines at those same wavelengths. Both map identical quantized energy gaps.