Unit 1: Atomic Structure and Properties

Periodic table

An organized chart of elements that summarizes atomic structure and helps predict chemical behavior based on patterns in rows (periods) and columns (groups).

Element symbol

The one- or two-letter abbreviation used to represent an element in formulas and equations (e.g., C for carbon).

Atomic number

The number of protons in an atom’s nucleus; it defines the element and equals the number of electrons in a neutral atom.

Average atomic mass

The weighted-average mass of an element’s naturally occurring isotopes; shown on the periodic table and usually not a whole number.

Atomic mass unit (amu)

A unit used for atomic-scale masses; periodic-table atomic masses are commonly expressed in amu.

Molar mass

The mass of 1 mole of a substance (g/mol); for an element it matches the periodic-table atomic mass numerically.

Period

A horizontal row on the periodic table; elements in the same period have valence electrons in the same principal energy level.

Group

A vertical column on the periodic table; main-group elements in the same group share similar valence electron patterns and properties.

Alkali metals

Group 1 elements; very reactive metals that commonly form +1 ions.

Alkaline earth metals

Group 2 elements; reactive metals that commonly form +2 ions.

Transition metals

Elements in groups 3–12 (d-block); metals that often form cations with multiple possible charges.

Halogens

Group 17 elements; reactive nonmetals that commonly form −1 ions.

Noble gases

Group 18 elements; largely unreactive gases with filled valence shells (He has 2 valence electrons; others have 8).

Proton

A positively charged subatomic particle located in the nucleus; the number of protons equals the atomic number.

Neutron

A neutral subatomic particle located in the nucleus; changing the number of neutrons creates different isotopes.

Electron

A negatively charged subatomic particle located in regions around the nucleus; involved in bonding and chemical behavior.

Isotope

Atoms of the same element (same protons) that have different numbers of neutrons and therefore different masses.

Mass number

The total number of protons plus neutrons in a specific isotope (an integer).

Weighted average (atomic mass)

A calculation that multiplies each isotope’s mass by its fractional abundance and sums the results to get average atomic mass.

Mass spectrometry

An analytical technique that measures isotopic masses and relative abundances by separating ions based on mass-to-charge ratio.

Mass-to-charge ratio (m/z)

The quantity used in mass spectrometry to separate ions; determines where peaks appear in a mass spectrum.

Relative abundance

The proportion of each isotope present in a natural sample; in a spectrum it corresponds to peak height/area.

Mole

A counting unit equal to exactly 6.02214076 × 10^23 representative particles.

Avogadro’s number

6.02214076 × 10^23; the number of representative particles in 1 mole.

Mole ratio

The ratio of moles of substances indicated by coefficients in a balanced chemical equation.

Formula unit

The smallest repeating unit of an ionic compound (e.g., 1 NaCl unit); not called a “molecule” for ionic substances.

Ideal gas law (PV = nRT)

An equation relating pressure (P), volume (V), moles (n), and temperature (T) for an ideal gas.

Standard temperature and pressure (STP)

Common reference conditions for gases: 1 atm and 273 K.

Percent composition (of an element)

The mass percent of a given element in a compound: (mass of that element in 1 mol compound ÷ molar mass) × 100%.

Mass percent (mixture)

A mixture concentration: (mass of component ÷ total mass of mixture) × 100%.

Parts per million (ppm)

A very dilute concentration unit: (mass of solute ÷ mass of solution) × 10^6 (often approximated as mg/kg for dilute aqueous solutions).

Mole fraction

The fraction of total moles contributed by a component: χA = nA / ntotal; mole fractions sum to 1.

Molarity (M)

A solution concentration unit: M = (moles of solute) / (liters of solution); bracket notation like [Na+] indicates molarity.

Empirical formula

The simplest whole-number ratio of elements in a compound (e.g., CH2O).

Molecular formula

The actual number of each type of atom in a molecule; a whole-number multiple of the empirical formula (e.g., C6H12O6).

Electron configuration

A notation that shows how electrons are distributed among orbitals and subshells in an atom.

Noble-gas shorthand

A shortened electron configuration that uses the previous noble gas in brackets to represent core electrons (e.g., [Ne] 3s1).

Orbital

A region of space with high probability of finding an electron at a specific energy within an atom.

Aufbau principle

Electrons fill lower-energy orbitals before higher-energy orbitals (e.g., 4s fills before 3d in the common filling order).

Pauli exclusion principle

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

Hund’s rule

Electrons occupy equal-energy orbitals singly (with parallel spins) before pairing up.

Photoelectron spectroscopy (PES)

A technique that uses high-energy light to eject electrons and measure how tightly electrons are held in different subshells.

Binding energy (PES)

The energy required to remove an electron from a particular subshell; in PES, peak position corresponds to binding energy.

Coulomb’s law

Electrostatic force relationship: F = k(q1q2)/r^2; attraction increases with greater charge and decreases with greater distance.

Shielding

Reduction of nuclear attraction on valence electrons due to repulsion by inner (core) electrons.

Effective nuclear charge (Zeff)

The net positive charge felt by an electron after accounting for shielding; generally increases across a period.

Atomic radius

A measure of atomic size; generally decreases across a period (higher Zeff) and increases down a group (more shells/shielding).

Ionic radius

The size of an ion; cations are smaller than their neutral atoms, and anions are larger due to electron loss/gain effects.

Electronegativity

An atom’s tendency to attract shared bonding electrons; generally increases across a period and decreases down a group (fluorine is highest).

Ionization energy

The energy required to remove an electron from a gaseous atom (e.g., M(g) → M+(g) + e−); generally increases across a period and decreases down a group.