AP Chemistry Unit 1 Review: Atomic Structures and Properties

Introduction

• Instructor: Cara
• Main topics: Moles, molar mass, atomic structures, isotopes, Dalton’s laws, electron configurations, periodic trends

Basic Concepts

• **Atoms: Protons (positive), Electrons (negative), Neutrons (neutral)
• **Periodic Table: Protons = Atomic Number, Atomic Mass in amu

Moles and Molar Mass

• Mole (mol): 6.02×10²³ (Avogadro’s number)
  • Converts atomic mass unit (amu) to grams (g)
  • E.g., 14 g of Nitrogen equals 1 mole of Nitrogen (14 amu → 14 g x 1 mole)
• Key formula: grams/mole = atomic mass in amu
• Example Calculation:
  • 25 g of calcium (Ca) = 25 g / 40.08 g/mol = 0.624 moles
  • Convert moles to molecules: 0.624 moles x 6.02×10²³ molecules/mole

Isotopes

• Definition: Atoms with same protons, different neutrons
• Atomic Mass: Weighted average of all natural isotopes
• Relative Abundance: Example with Carbon atoms (C-12, C-13)
• Mass Spectrometry: Tool to find isotopes’ mass and abundance
  • Atom ionized, accelerated, deflected in a magnetic field, and detected
  • Produces charge-to-mass ratio graph

Dalton’s Laws

• Key Points: (Dalton’s Atomic Theory)
  • Atoms are smallest unit
  • Mass is conserved
  • Compounds maintain a specific ratio of atoms
• Compound Notation:
  • Structural Formula
  • Chemical Formula
  • Empirical Formula

Ions

• **Atoms and Electrons: Positively charged nucleus, negatively charged electrons in orbitals
• Valence Electrons: Outermost shell (importance in bonding)
  • Cation: Positive ion (e.g., Na+)
  • Anion: Negative ion (e.g., Cl-)
• Ionic Bonding: Exchange of electrons (e.g., NaCl)
• Naming Compounds:
  • Cation + Anion (with “-ide” suffix)
  • For transition metals, specify charge (e.g., Iron(III) Chloride)**

Electron Configurations

• Quantum Numbers (n, l, m, s):
  • n = Principal energy level
  • l = Sublevel (s, p, d, f)
  • m = Orbital orientation
  • s = Spin (+½, -½)
• Electron Configuration Notation: Indicate the number of electrons in each sublevel (e.g., 1s², 2s², 2p⁶)
  • Use periodic table to determine fill order
  • Example: Calcium (Ca): [Ar] 4s²
• Hund’s Rule: Electrons fill orbitals singly before pairing
• Exceptions: e.g., Cr ([Ar] 4s¹ 3d⁵), Cu ([Ar] 4s¹ 3d¹⁰)

Periodic Trends

• Atomic Radius: Decreases across a period, increases down a group
• Ionization Energy: Increases across a period, decreases down a group
• Electron Affinity: Energy change when an electron is added; highest for elements like Cl
• Electronegativity: Tendency to attract electrons; highest in F, decreases from there
• Metallicity: Decreases across a period, increases down a group

Photoelectron Spectroscopy

• Basic Concept: Shines light on electrons, measures energy needed to remove them
  • Graph Interpretation: Higher peaks ≠ closer electrons, show electron number and energy levels

Conclusion

• Be aware of periodic table trends and definitions
• Understand isotope significance and usage
• Practice electron configurations for various elements
• Familiarize with Dalton’s laws for compounds
• Recognize and interpret mass spectrometry and photoelectron spectroscopy data

Final Notes

• Review naming conventions, especially for ionic compounds
• Understand electron configuration exceptions (e.g., Cr, Cu)
• Grasp basic periodic trends: radius, ionization energy, electronegativity, metallicity
• Remember key concepts for AP Chem success!