Overview
This lecture reviews key concepts from AP Chemistry Unit 1, covering atomic structure, properties, calculations with moles, electron configuration, periodic trends, and compound formulas.
Mole & Mass Calculations
- To convert grams to moles, divide mass by molar mass (sum of atomic masses in the compound).
- Example: 10.00g CO₂ × (1 mol / 44.01g) = 0.2272 mol.
- To convert moles to particles, multiply by Avogadro’s number (6.022 × 10²³ particles/mol).
- Example: 0.2272 mol CO₂ × (6.022 × 10²³ / 1 mol) = 1.368 × 10²³ molecules.
Mass Spectrometry & Isotopes
- Mass spectrometry graphs show relative abundance of isotopes for an element.
- The weighted average from isotopic masses and abundances gives the element’s atomic mass.
Empirical & Molecular Formulas
- Empirical formula is the simplest whole-number ratio of atoms in a compound.
- Convert mass percent to grams, then to moles, then divide by the smallest mole value to get subscripts.
- Law of definite proportions: a compound always has the same percent composition by mass.
Mixtures vs Pure Substances
- Mixtures contain impurities, not all of the sample mass is the desired compound.
- Calculate percent composition to determine purity.
Electron Configuration & Atomic Structure
- Electron configuration lists the arrangement of electrons in energy levels and sublevels.
- Valence electrons are in the outermost shell and determine reactivity.
- Sublevels are labeled as s, p, d, etc.
Coulomb’s Law & Atomic Forces
- Coulomb’s Law: force increases with greater charge, decreases with greater distance.
- Valence electrons (farther from nucleus) are more easily removed than core electrons (closer).
Photoelectron Spectroscopy (PES)
- PES graphs show electron energy levels as peaks corresponding to sublevels.
- Peak heights show number of electrons in each sublevel.
Periodic Trends
- Ionization energy and electronegativity increase up and to the right on the periodic table; atomic radius decreases.
- Effective nuclear charge increases across periods, causing atomic size to decrease.
- Down groups, increased distance from nucleus causes larger atomic radius.
Ions & Compound Formation
- Positive ions (cations) are smaller; negative ions (anions) are larger due to electron/proton ratio.
- The group number predicts typical ion charge (e.g., Group 1: +1, Group 17: -1).
- Compound formulas balance charges to sum to zero (e.g., MgCl₂, Al₂S₃).
Key Terms & Definitions
- Mole — 6.022 × 10²³ particles of a substance.
- Empirical Formula — lowest whole-number ratio of elements in a compound.
- Valence Electrons — electrons in the outermost shell.
- Coulomb’s Law — force between charged particles depends on charge and distance.
- Photoelectron Spectroscopy (PES) — technique to analyze electron energies in atoms.
Action Items / Next Steps
- Practice converting between grams, moles, and particles.
- Review how to interpret mass spectrometry and PES graphs.
- Memorize periodic trends and typical ion charges for main group elements.