Unit 1

Unit 1 of AP Chemistry covers the foundational concepts of atomic structure and properties. Here’s a summary of the key ideas:

1. Atomic Structure:

• Subatomic Particles: Atoms consist of protons (positively charged), neutrons (neutral), and electrons (negatively charged).

  • Protons and neutrons are found in the nucleus, while electrons orbit around the nucleus in various energy levels (shells).

• Isotopes: Atoms of the same element with different numbers of neutrons, resulting in different mass numbers.

  • Example: Carbon-12 and Carbon-14 are isotopes of carbon.

2. Atomic Mass and Atomic Number:

• Atomic Number: The number of protons in an atom, which defines the element.

• Mass Number: The total number of protons and neutrons in an atom’s nucleus.

• Average Atomic Mass: A weighted average of the isotopes of an element based on their natural abundance.

3. The Mole and Avogadro’s Number:

• Mole (mol): A fundamental unit in chemistry representing 6.022 × 10²³ particles (atoms, molecules, etc.).

  • Used for quantifying substances in terms of atoms or molecules.

• Molar Mass: The mass of one mole of a substance (g/mol). It is numerically equal to the atomic or molecular mass in atomic mass units (amu).

4. Mass Spectrometry:

• Mass Spectrometry is used to determine the relative abundances of isotopes of an element.

  • A mass spectrometer separates ions based on their mass-to-charge ratio, allowing the calculation of an element's average atomic mass.

5. The Periodic Table:

• Periodic Trends: Important trends include atomic radius, ionization energy, electronegativity, and electron affinity.

  • Atomic Radius: Decreases across a period (left to right) and increases down a group (top to bottom).

  • Ionization Energy: Increases across a period and decreases down a group.

  • Electronegativity: The ability of an atom to attract electrons in a bond; it increases across a period and decreases down a group.

6. Electron Configuration:

• Describes the arrangement of electrons in an atom’s energy levels, subshells (s, p, d, f), and orbitals.

  • Aufbau Principle: Electrons fill lower energy orbitals first.

  • Hund’s Rule: Electrons fill degenerate orbitals singly before pairing up.

  • Pauli Exclusion Principle: No two electrons can have the same set of quantum numbers (no more than two electrons per orbital).

7. Quantum Mechanical Model:

• Electrons have specific energy levels, but their exact location is described by probability distributions rather than fixed orbits (as in the Bohr model).

• Orbitals: Regions around the nucleus where electrons are likely to be found (spherical for s orbitals, dumbbell-shaped for p orbitals, etc.).

8. Ions and Ionic Compounds:

• Ions form when atoms gain or lose electrons to achieve a full outer electron shell.

  • Cations: Positively charged ions (formed by losing electrons).

  • Anions: Negatively charged ions (formed by gaining electrons).

• Ionic Compounds: Formed by electrostatic attraction between cations and anions, resulting in ionic bonds (e.g., NaCl).

9. Empirical and Molecular Formulas:

• Empirical Formula: The simplest whole-number ratio of atoms in a compound.

• Molecular Formula: The actual number of atoms of each element in a molecule.

10. Chemical Reactions and Stoichiometry:

• Balancing Equations: Ensuring the number of atoms for each element is the same on both sides of a chemical equation.

• Stoichiometry: The calculation of reactants and products in chemical reactions based on balanced equations.

  • Limiting Reactant: The reactant that limits the amount of product that can be formed.

  • Percent Composition: The percentage by mass of each element in a compound.

11. Mixtures and Pure Substances:

• Mixtures: Composed of two or more substances that are not chemically combined (e.g., solutions).

  • Heterogeneous Mixture: Components are not evenly distributed.

  • Homogeneous Mixture (Solution): Components are uniformly distributed.

• Pure Substances: Elements or compounds with a fixed composition.