AP Chemistry Exam Cram Session by Jeremy Kug

Introduction

• Final days/hours before the AP Chemistry exam
• Video is a long but valuable cram session
• Print out guided notes PDF for following along
• Visit ultimaterreviewpacket.com for comprehensive review resources

Unit 1: Basic Chemistry Concepts

Mass Percent

• Definition: Percentage by mass of each element in a compound
• Example 1: Magnesium Chloride (MgCl₂)
  • Atomic masses: Magnesium (24.31 amu), Chlorine (35.45 amu)
  • Calculation: Total molar mass = 95.21 amu, Mg = 25.53%, Cl = 74.47%
• Example 2: Calcium Chloride (CaCl₂)
  • Atomic masses: Calcium (40.08 amu), Chlorine (35.45 amu)
  • Calculation: Total molar mass = 110.98 amu, Ca = 36.11%, Cl = 63.89%
• Patterns: Lighter metals in chlorides result in a higher chloride percentage

Mass Spectrometry

• Graph Interpretation: Peaks represent isotopes
• Example: Element with isotopes at 85 and 87 amu
  • Abundance: 75% and 25% respectively
  • Average atomic mass: Approx. 85.50 amu, close to Rubidium

Electron Configurations

• Writing Configurations: Use periodic table sections to write electron configurations
• Examples:
  • Chlorine (Cl): 1s² 2s² 2p⁶ 3s² 3p⁵
  • Chloride ion (Cl⁻): 1s² 2s² 2p⁶ 3s² 3p⁶
  • Aluminum (Al): 1s² 2s² 2p⁶ 3s² 3p¹
  • Aluminum ion (Al³⁺): 1s² 2s² 2p⁶
• Periodic Trends: Atomic radius increases downwards and leftwards due to electron shells and effective nuclear charge

Unit 2: Molecular Structures and Bonding

Lewis Electron Dot Diagrams

• Example: SF₂
  • Central atom: Sulfur
  • Electron distribution: Fluorine atoms have 7 valence electrons, Sulfur has 6
  • Final structure: Bent molecular geometry

Sigma and Pi Bonds

• Definitions:
  • Single bond = 1 sigma bond
  • Double bond = 1 sigma + 1 pi bond
  • Triple bond = 1 sigma + 2 pi bonds
• Hybridization:
  • Determined by electron regions/domains around central atom
  • Examples:
    • sp³ (4 electron regions)
    • sp² (3 electron regions)

Molecular Geometry and Bond Angles

• Example: SF₂
  • Geometry: Bent
  • Bond angle: ~104.5°
• Geometries: Vary based on electron regions—tetrahedral, trigonal planar, bent

Polarity

• Concept: Polarity determined by unbalanced regions of negative charge
  • Bent structures are generally polar
  • London dispersion forces and dipole-dipole forces are present in polar molecules

Unit 3: Gas Laws and Spectrophotometry

Gas Law Example

• Equation: PV = nRT
  • Use given pressure, volume, and temperature to solve for moles

Spectrophotometry

• Beer-Lambert Law: Plot absorbance vs. concentration to determine unknown concentrations
• Outliers: Points deviating from the calibration curve indicate contamination

Unit 4: Reaction Stoichiometry

Net Ionic Equations

• Example: Mg + Cu²⁺ → Mg²⁺ + Cu
  • Exclude spectator ions like Cl⁻
  • Balance charges with electrons

Stoichiometry Problems

• Steps:
  1. Convert to moles
  2. Use mole ratio
  3. Convert to final unit (usually grams)

Unit 5: Kinetics

Rate Laws and Reaction Order

• Determine Order: Compare trials where only one reactant changes
• Rate Law Writing: Rate = k[A]ⁿ[B]ᵐ
• Rate Constant: Use experimental data to calculate, noting units

Unit 6: Thermochemistry

Specific Heat and Calorimetry

• Equation: Q = mcΔT
  • Calculate heat transfer between substances (e.g., copper and water)

Heating and Cooling Curves

• Phases: Solid, liquid, gas with phase changes at melting and boiling points

Enthalpy Calculations

• Equation: ΔH = Σ(products) - Σ(reactants)
  • Use bond energies, enthalpies of formation

Unit 7: Equilibrium

Equilibrium Expressions

• KC and KP: Write expressions using products/reactants, excluding solids and liquids

ICE Box Method

• Application: Solve equilibrium problems with initial, change, equilibrium concentrations

Le Chatelier’s Principle

• Responses: System's adjustment to changes in concentration, pressure, temperature

Unit 8: Acids and Bases

pH and pOH

• Relationships: pH + pOH = 14, [H₃O⁺][OH⁻] = 1 x 10⁻¹⁴ at 25°C

Strong Acids and Bases

• Strong Acids: HCl, HBr, HI, HNO₃, H₂SO₄, HClO₄
• Strong Bases: Group 1 and 2 hydroxides

Weak Acids and Bases

• Equilibrium Approach: Use ICE box for pH calculation
• Salt Analysis: Determine acidic/basic nature of salts using ionic equations

Unit 9: Thermodynamics and Electrochemistry

Entropy

• Concept: Measure of disorder, higher in gases than solids

Gibbs Free Energy

• Calculations: ΔG = ΔH - TΔS
  • Negative ΔG indicates a thermodynamically favorable reaction

Electrochemistry

• Galvanic Cells: Calculate cell potential, identify anode/cathode
  • Anode to cathode electron flow

Conclusion: Review key concepts, practice problems, and tackle the AP exam with confidence. Good luck!