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AP Chemistry Study Guide

1. Introduction to AP Chemistry

AP Chemistry focuses on the principles and concepts of chemistry, including atomic structure, chemical bonding, stoichiometry, thermodynamics, and kinetics. The course emphasizes problem-solving, scientific inquiry, and laboratory skills, while preparing students for the AP exam.

Exam Format:

  • Multiple-choice questions: Test your understanding of chemistry concepts and problem-solving ability.
  • Free-response questions: Assess your ability to apply chemical knowledge to novel situations, analyze data, and explain concepts.

2. Atomic Structure and Periodicity

Atomic Structure:

  • Atoms: The smallest unit of an element, composed of protons, neutrons, and electrons.
  • Subatomic Particles:
    • Protons have a positive charge and are found in the nucleus.
    • Neutrons have no charge and are found in the nucleus.
    • Electrons have a negative charge and are found in orbitals around the nucleus.
  • Isotopes: Atoms of the same element with different numbers of neutrons (e.g., Carbon-12, Carbon-14).

Electron Configuration:

  • Electrons are arranged in energy levels (shells) around the nucleus. The Pauli exclusion principle and Hund’s rule govern the arrangement of electrons in orbitals.
  • Periodic Table: Elements are arranged by atomic number and share similar properties in vertical columns (groups/families).

Periodic Trends:

  • Atomic Radius: Decreases across a period and increases down a group.
  • Ionization Energy: Increases across a period and decreases down a group.
  • Electron Affinity: Becomes more negative across a period and less negative down a group.
  • Electronegativity: Increases across a period and decreases down a group.

3. Chemical Bonding and Molecular Structure

Ionic Bonding:

  • Occurs when electrons are transferred between atoms, creating positive and negative ions that are held together by electrostatic forces.
  • Example: NaCl (sodium chloride).

Covalent Bonding:

  • Occurs when two atoms share electrons to achieve a full outer shell.
  • Polar Covalent Bonds: Unequal sharing of electrons, leading to partial positive and negative charges (e.g., H₂O).
  • Nonpolar Covalent Bonds: Equal sharing of electrons (e.g., Cl₂).

Bond Strength and Length:

  • Bond Energy: The energy required to break a bond.
  • Bond Length: The distance between the nuclei of two bonded atoms.

Lewis Structures:

  • Use electron-dot symbols to represent atoms and their valence electrons.
  • Resonance: Some molecules cannot be represented by a single Lewis structure, and multiple structures are required to show the possible configurations.

4. Stoichiometry and Chemical Reactions

Stoichiometry:

  • The calculation of reactants and products in a chemical reaction.
  • Mole Concept: 1 mole = 6.022 × 10²³ particles (Avogadro’s number).
  • Molar Mass: The mass of one mole of a substance (in g/mol).

Balancing Chemical Equations:

  • Chemical reactions must be balanced to conserve mass and charge.
  • Example: 2H₂ + O₂ → 2H₂O.

Types of Chemical Reactions:

  • Synthesis: Two or more substances combine to form a product (e.g., 2H₂ + O₂ → 2H₂O).
  • Decomposition: A compound breaks down into simpler substances (e.g., 2H₂O → 2H₂ + O₂).
  • Single Replacement: One element replaces another in a compound (e.g., Zn + 2HCl → ZnCl₂ + H₂).
  • Double Replacement: Two compounds exchange ions to form new compounds (e.g., AgNO₃ + NaCl → AgCl + NaNO₃).
  • Combustion: A hydrocarbon reacts with oxygen to produce CO₂ and H₂O (e.g., CH₄ + 2O₂ → CO₂ + 2H₂O).

5. Thermochemistry and Thermodynamics

Enthalpy (ΔH):

  • The heat content of a system at constant pressure.
  • Exothermic Reactions: Release heat (ΔH is negative).
  • Endothermic Reactions: Absorb heat (ΔH is positive).

First Law of Thermodynamics:

  • Energy cannot be created or destroyed, only transferred or converted.

Entropy (ΔS):

  • A measure of disorder or randomness in a system.
  • Spontaneous Processes: Processes that increase the total entropy of the universe.

Gibbs Free Energy (ΔG):

  • ΔG = ΔH - TΔS, where T is the temperature in Kelvin.
  • Spontaneous Reactions: Occur when ΔG is negative.

6. Kinetics and Chemical Equilibrium

Reaction Rate:

  • The speed at which a chemical reaction occurs.
  • Influenced by concentration, temperature, and the presence of a catalyst.

Rate Law:

  • The rate of a reaction is proportional to the concentration of reactants raised to a power (e.g., rate = k[A]²[B]).

Activation Energy (Ea):

  • The minimum energy required for a reaction to occur.

Catalysts:

  • Substances that speed up a reaction by lowering the activation energy without being consumed.

Chemical Equilibrium:

  • Occurs when the rate of the forward reaction equals the rate of the reverse reaction.
  • Le Chatelier’s Principle: If a system at equilibrium is disturbed, it will shift in a direction that counteracts the disturbance.

7. Acids and Bases

Bronsted-Lowry Definition:

  • Acid: A proton (H⁺) donor.
  • Base: A proton (H⁺) acceptor.

Strong vs. Weak Acids and Bases:

  • Strong acids (e.g., HCl, H₂SO₄) dissociate completely in water.
  • Weak acids (e.g., CH₃COOH) partially dissociate in water.

pH and pOH:

  • pH = -log[H⁺], pOH = -log[OH⁻].
  • pH + pOH = 14.

Buffer Solutions:

  • Solutions that resist changes in pH when small amounts of acid or base are added.
  • Composed of a weak acid and its conjugate base or a weak base and its conjugate acid.

8. Electrochemistry

Redox Reactions:

  • Oxidation: The loss of electrons.
  • Reduction: The gain of electrons.
  • In a redox reaction, one species is oxidized and the other is reduced.

Electrochemical Cells:

  • Galvanic (Voltaic) Cells: Spontaneous reactions that generate electrical energy (e.g., batteries).
  • Electrolytic Cells: Non-spontaneous reactions that require an external power source.

Cell Potential (E°):

  • The difference in potential energy between the anode and cathode in an electrochemical cell.
  • A positive E° indicates a spontaneous reaction.