Topic 1: Atomic Structure and Properties

Unit 1 builds the foundation of AP Chemistry by connecting the invisible structure of atoms to measurable laboratory quantities. You will learn to count particles with the mole, read mass and photoelectron spectra, and predict properties from electron arrangement. Mastering these ideas makes every later unit easier.

Moles and Molar Mass

Atoms are far too small and too numerous to count one at a time, so chemists group them into moles. One mole contains Avogadro's number of particles, 6.022 x 10^23, whether those particles are atoms, molecules, ions, or electrons. The molar mass of a substance, expressed in grams per mole, is numerically equal to its average atomic or formula mass read from the periodic table. To convert between grams and moles, divide a sample's mass by its molar mass; to find the number of particles, multiply moles by Avogadro's number. These conversions are the backbone of nearly every quantitative problem in the course, so practice moving fluidly among grams, moles, and particles.

Mass Spectrometry of Elements

A mass spectrometer ionizes atoms, accelerates the ions, and separates them by their mass-to-charge ratio, producing a spectrum of peaks. For a pure element, each peak represents a different isotope, and the height of each peak reflects the relative abundance of that isotope. Because isotopes share the same number of protons but differ in neutrons, they have nearly identical chemistry but distinct masses. To calculate an element's average atomic mass, multiply each isotope's mass by its fractional abundance and add the results. This weighted average explains why periodic-table masses are rarely whole numbers, such as chlorine's value near 35.45.

Percent Composition and Empirical Formula

Percent composition tells you what fraction of a compound's mass comes from each element. You find it by dividing the total mass of one element in a formula by the compound's molar mass and multiplying by 100. Working in reverse, experimental mass percentages let you determine an empirical formula, the simplest whole-number ratio of atoms. The standard procedure is to assume a 100 gram sample, convert each element's grams to moles, divide every mole value by the smallest, and round to small whole numbers. If the molar mass of the actual molecule is known, dividing it by the empirical formula mass gives the multiplier that yields the molecular formula.

Pure Substances Versus Mixtures

A pure substance has a fixed, uniform composition and includes both elements and compounds. A mixture combines two or more substances that are not chemically bonded, so its composition can vary. Mixtures are classified as homogeneous, where the composition is uniform throughout like saltwater, or heterogeneous, where distinct regions are visible like sand in water. Pure substances show consistent physical properties such as a single sharp melting point, while mixtures often melt or boil over a range. On the AP exam, this distinction supports reasoning about separations, spectra, and whether observed data reflect one substance or several.

Electron Configuration and the Aufbau, Pauli, and Hund Rules

Electrons occupy orbitals according to three guiding principles. The Aufbau principle says electrons fill the lowest-energy orbitals first, following the order 1s, 2s, 2p, 3s, 3p, 4s, 3d, and so on. The Pauli exclusion principle states that no orbital can hold more than two electrons, and those two must have opposite spins. Hund's rule requires electrons to occupy degenerate orbitals singly, with parallel spins, before any orbital is doubled up, which minimizes repulsion. Writing a configuration like 1s2 2s2 2p6 captures how many electrons sit in each subshell and underlies bonding, magnetism, and periodic trends.

Photoelectron Spectroscopy (PES)

Photoelectron spectroscopy uses high-energy light to eject electrons from an atom and measures the energy required to remove them, revealing the energy of each subshell. A PES spectrum shows peaks plotted against binding energy, usually with higher energy on the left. The position of a peak indicates which subshell the electrons came from, with core electrons closer to the nucleus appearing at higher binding energies. The height of a peak is proportional to the number of electrons in that subshell, so a 2p peak is typically taller than a 2s peak. PES gives direct experimental evidence for the shell model and electron configurations.

Periodic Trends: Atomic Radius, Ionization Energy, and Electronegativity

Periodic trends arise from two competing factors: effective nuclear charge and the number of occupied shells. Atomic radius decreases across a period as growing nuclear charge pulls electrons inward, and increases down a group as new shells are added. First ionization energy, the energy to remove the outermost electron, increases across a period and decreases down a group, the opposite pattern of radius. Electronegativity, an atom's tendency to attract bonding electrons, also rises across a period and falls down a group, peaking near fluorine. Recognizing the underlying cause, not just the arrow direction, lets you justify trends on free-response questions.

Valence Electrons and Ionic Compounds

Valence electrons are the outermost electrons that determine an element's chemical behavior, and for main-group elements their count matches the group number pattern. Atoms tend to gain, lose, or share electrons to reach a stable, noble-gas-like arrangement. Metals lose valence electrons to form positive cations, while nonmetals gain electrons to form negative anions. When oppositely charged ions attract, they assemble into an ionic compound held together by strong electrostatic forces in a repeating lattice. The ratio of ions in the formula reflects the charges balancing to a neutral whole, as in the one-to-one pairing of sodium and chloride.