MCAT General Chemistry Review

Introduction

• Review of general chemistry for the MCAT.
• Presenter achieved a 100 percentile score using the Miles Down review sheets.
• Key study strategy: pause the video, take notes, and ensure understanding by answering questions.

Basic Atomic Structure

• Symbols:
  • A: Mass number = number of protons + neutrons
  • Z: Atomic number = number of protons
• Isotopes: Atoms with the same number of protons but different neutrons (e.g., Carbon-12, Carbon-13, Carbon-14).
• Periodic Table Insights:
  • Weighted average of isotopes determines atomic mass.
  • Atomic mass example for Carbon: 12.011.

Historical Models

• Rutherford Model (1911): Discovered the nucleus using gold foil experiment.
• Bohr Model (1913): Electrons in energy levels; light emitted when electrons move to lower energy.
• Heisenberg Uncertainty Principle: Cannot know position and momentum of an electron simultaneously.

Quantum Mechanics

• Hund's Rule: Electrons fill orbitals singly before pairing.
• Pauli's Exclusion Principle: No two electrons can have the same set of quantum numbers.
• Constants to Know:
  • Avogadro's Number: 6.022 x 10^23
  • Planck's Constant: 6.626 x 10^-34
  • Speed of Light: 3 x 10^8 m/s

Atomic Theory

• Quantum Numbers: Define electron positions and energy levels:
  • Principal (n): Energy level
  • Azimuthal (l): Shape of orbital
  • Magnetic (ml): Orientation of orbital
  • Spin (ms): Electron spin direction
• Electron Configurations: Follow Aufbau principle for filling order.

Periodic Table Trends

• Effective Nuclear Charge (Z_eff): Increases across a period.
• Ionization Energy: Energy required to remove an electron; increases across a period.
• Electron Affinity: Energy change when an electron is added; highest near halogens.
• Electronegativity: Attraction in a bond; fluorine is the most electronegative.
• Atomic Size: Increases down a group, decreases across a period.

Chemical Bonds

• Covalent Bonds: Sharing of electrons. Polar if difference in electronegativity is 0.5-1.7.
• Ionic Bonds: Transfer of electrons; typically between metals and nonmetals.
• Intermolecular Forces:
  • Hydrogen bonds, dipole-dipole, and London dispersion forces.
• Sigma and Pi Bonds:
  • Sigma: First bond between atoms.
  • Pi: Additional bonds in double/triple bonds.

Molecular Geometry

• VSEPR Theory: Shapes of molecules based on electron pair repulsion.
• Hybridization Types: sp, sp2, sp3, etc., depending on the number of electron groups.

Reaction Types and Kinetics

• Reaction Rates: Zero, first, and second order kinetics.
• Gibbs Free Energy: Determines reaction spontaneity.
• Arrhenius Equation: Relates temperature and activation energy to reaction rate.

Solutions and Solubility

• Solution Terminology: Solvent, solute, molarity, molality.
• Solubility Product (Ksp): Equilibrium constant for salts in solution.
• Colligative Properties: Depend on the number of particles in solution (boiling point elevation, freezing point depression).

Acids and Bases

• Definitions:
  • Arrhenius, Bronsted-Lowry, Lewis definitions.
• pH and pOH: Measure of acidity/basicity; pH + pOH = 14.
• Buffers: Resist changes in pH; Henderson-Hasselbalch equation.
• Titrations: Determine concentration of an unknown solution.

Electrochemistry

• Galvanic vs. Electrolytic Cells:
  • Galvanic: Spontaneous redox reactions.
  • Electrolytic: Non-spontaneous, driven by external voltage.
• Nernst Equation: Relates cell potential to concentration of ions.

Conclusion

• Comprehensive review covering essential chemistry concepts for the MCAT.
• Emphasis on understanding and memorizing key formulas and trends.