AP Chemistry Unit 1 (Atomic Structure & Properties): Learning Moles, Mass Data, and Composition

Mole (mol)

A counting unit in chemistry defined as the amount of substance containing exactly 6.022 × 10^23 entities.

Avogadro’s number (N_A)

6.022 × 10^23 entities per mole (mol⁻¹); used to convert between moles and particles.

Molar mass (M)

The mass of 1 mole of a substance, typically in g/mol; for a compound, it is the sum of atomic masses in the formula.

Formula unit

The “particle” counted for an ionic compound; the lowest whole-number ratio of ions represented by the chemical formula (e.g., NaCl).

Grams-to-moles conversion

Conversion using molar mass: n = m/M, where n is moles, m is mass (g), and M is molar mass (g/mol).

Moles-to-particles conversion

Conversion using Avogadro’s number: N = nN_A, where N is the number of entities and n is moles.

Mass-to-particles conversion

Direct conversion combining molar mass and Avogadro’s number: N = (m/M)N_A.

Mass spectrometry

An experimental technique that determines particle masses and the relative abundances of isotopes, helping explain average atomic mass values.

Ionization (mass spectrometry step)

The process of converting atoms into positive ions (often +1) so they can be manipulated by electric/magnetic fields.

Mass-to-charge ratio (m/z)

The quantity used to separate ions in a mass spectrometer; ions are deflected based on mass divided by charge.

Mass spectrum

A graph produced by a mass spectrometer showing signal intensity versus m/z values (peaks correspond to isotopes/components).

Relative abundance

How common an isotope is in a sample; shown by peak height on a mass spectrum (often scaled so the tallest peak is 100).

Average atomic mass

The weighted average mass of an element’s naturally occurring isotopes; this is the non-integer value listed on the periodic table.

Weighted average

An average that accounts for different contributions: average atomic mass = Σ(fi mi), where fi are fractional abundances and mi are isotope masses.

Isotope

Atoms of the same element (same number of protons) with different numbers of neutrons, giving different masses.

Mass number

An integer equal to protons + neutrons for a specific atom/isotope (not the same as measured isotopic mass).

Isotopic mass

The experimentally measured mass of a specific isotope; usually not a whole number.

Percent composition by mass

For element X in a compound: % by mass = (mass of X in 1 mol compound / molar mass of compound) × 100.

Empirical formula

The simplest whole-number ratio of atoms in a compound; represents the lowest “recipe ratio,” not necessarily the actual molecule.

Empirical formula determination procedure

Common steps: assume a sample size (often 100 g), convert each element’s grams to moles, divide by the smallest mole value, then multiply all ratios to get whole numbers if needed.

Molecular formula

The actual numbers of each type of atom in a molecule; it is a whole-number multiple of the empirical formula.

Molecular formula multiplier

Multiplier = (molar mass) / (empirical formula mass); multiply all empirical subscripts by this (near-integer) value to get the molecular formula.

Hydrate

An ionic compound containing water in its crystal structure, written as salt · xH2O; x is found from mole ratios using mass loss upon heating.

Mass percent (mixture)

For component A: % by mass = (mA / mtotal) × 100; uses total mixture mass in the denominator.

Parts per million (ppm)

A very small mass fraction: ppm = (msolute / msolution) × 10^6; often treated as mg solute per kg solution.