AP Chemistry: Unit 1 – Atomic Structure and Properties

Introduction

• Lecturer: Jeremy Krug
• Covers: Conversion between moles and grams, particles and moles, interpreting mass spectrometer graphs, empirical formulas, mixtures vs. pure substances, electron configurations, Coulomb’s Law, periodic trends, ions and valence electrons.

Converting Moles and Grams

Using Atomic Mass

• Convert grams to moles using atomic mass.
• For compounds, sum of atomic masses is used.
• Example: 10.00g CO₂ → moles via dimensional analysis with conversion factor (1 mole / 44.01g) = 0.2272 moles.

Converting Particles to Moles

• 6.022 x 10²³ particles in a mole.
• Example: 0.2272 moles CO₂ → molecules (1 mole / 6.022 x 10²³ molecules) = 1.368 x 10²³ molecules.

Interpreting Mass Spectrometer Graphs

• Shows relative abundance of isotopes.
• Example: Graph with isotopes 107 amu and 109 amu for silver (Ag).
• Calculate average atomic mass from percentages to identify element.

Determining Empirical Formulas

Simplest Whole Number Ratio

• Empirical formula = simplest ratio.
• Example: H₂C₂O₄ (molecular) → HCO₂ (empirical).

Composition Data

• Example: Substance with 40.05% sulfur and 59.95% oxygen.
• Convert to grams, then moles, divide by smallest mole value.
• Empirical formula = SO₃.

Mixtures vs. Pure Substances

• Sample analysis for specific substances and impurities.
• Example: Potassium chloride (analysis of potassium and chloride content).
• Compare mass percent to pure samples to determine impurity levels.

Electron Configurations

Definitions

• Valence electrons: Outer energy level electrons.
• Sublevels/subshells: s, p, d (in scandium).

Writing Electron Configurations

• Example: Scandium (Sc) = 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹.

Coulomb’s Law

• Force is dependent on charge and distance.
• Greater charge → stronger force; greater distance → weaker force.
• Inner (core) electrons are harder to remove than valence (outer) electrons.

Photoelectron Spectroscopy (PES)

• Interpreting PES graphs by labeling peaks from left to right.
• Example: Peaks heights correspond to electron numbers in sublevels.
• Identify element based on electron configuration from PES data.

Periodic Trends

Trends

• Ionization energy and electronegativity increase right and upward.
• Atomic radius increases left and downward.

Explanations

• Right: greater nuclear charge; left: lower nuclear charge.
• Upper atoms: lower distance from nucleus; lower atoms: greater distance.

Ions

• More positive ion → smaller size; more negative ion → larger size (Coulomb’s Law).
• Example: Sodium chloride (NaCl) and its mass percent analysis for purity.
• Patterns in groups indicate valence electrons.

Octet Rule

• Predict charges for ions: Group 1 (+1), Group 2 (+2), etc.
• Examples:
  • Magnesium chloride (MgCl₂) = +2 (Mg) and -1 (Cl).
  • Aluminum sulfide (Al₂S₃) = +3 (Al) and -2 (S).

Conclusion

• Summary of Unit 1 topics.
• Next review: AP Chemistry Unit 2.