AP Chem Semester Review

1. Mass Spectrometry

  • Purpose: Separates isotopes of an element to determine isotopic composition.

  • Process:

    1. High-energy electrons ionize atoms.

    2. Ions are separated by their mass-to-charge ratio (m/z) using an electric field.

    3. Data is compiled into a mass spectrum.

  • Key Points:

    • Lighter isotopes are deflected more.

    • The mass spectrum displays isotopic deflection and abundance.

2. Empirical and Molecular Formulas

  • Empirical Formula: Smallest whole number ratio of atoms in a compound.

  • Molecular Formula: Actual number of atoms in a molecule.

  • Steps to Solve Empirical Formulas:

    1. Convert percentage to mass (e.g., 42% = 42g).

    2. Convert mass to moles.

    3. Divide all mole values by the smallest mole value.

    4. Adjust to whole numbers if necessary (e.g., multiply by 2 for 0.5 fractions).

  • Molecular Formula Calculation:

    1. Find empirical formula.

    2. Divide molecular mass by empirical mass.

    3. Multiply empirical subscripts by the resulting multiplier.

3. Hydrate Formulas

  • Definition: Ionic compounds containing water molecules in their structure.

  • Steps to Solve:

    1. Convert the mass of anhydrous compound and water to moles.

    2. Divide moles of water by moles of salt to find the formula.

4. Electron Configurations

  • Orbitals:

    • Shapes: s (sphere), p (dumbbell), d, and f.

    • Each orbital holds 2 electrons with opposite spins.

  • Energy Levels:

    • Higher orbitals (e.g., d, f) have more energy.

    • Electrons fill lower energy orbitals first (Aufbau principle).

5. Coulomb's Law

  • Formula: Attractive force (F) ∝ (charge of particles) / (distance between particles)².

  • Effective Nuclear Charge (Z_eff):

    • Pull of the nucleus on valence electrons.

    • Calculated as: Zeff=Protons−Core ElectronsZ_\text{eff} = \text{Protons} - \text{Core Electrons}.

6. Periodic Trends

  • Atomic Radius:

    • Decreases across a period (higher nuclear charge).

    • Increases down a group (more energy levels).

  • Ionization Energy:

    • Energy required to remove the outermost electron.

    • Increases across a period, decreases down a group.

  • Electron Affinity:

    • Energy change when an electron is added.

    • More negative across a period; exceptions in Groups 2, 15, and 18.

7. Bonding

  • Ionic Bonds: Transfer of electrons between a metal and a nonmetal.

  • Covalent Bonds:

    • Polar: Unequal sharing of electrons (partial charges).

    • Nonpolar: Equal sharing of electrons.

  • Metallic Bonds: Delocalized electrons in a “sea of electrons.”

8. Intermolecular Forces (IMFs)

  • Types:

    1. London Dispersion Forces: Present in all substances; stronger in larger, more polarizable molecules.

    2. Dipole-Dipole Interactions: Between polar molecules.

    3. Hydrogen Bonding: Strong dipole interaction; occurs with H bonded to N, O, or F.

    4. Ion-Dipole Forces: Between ions and polar molecules.

  • Impact: IMFs influence boiling point, melting point, and solubility.

9. Gas Laws

  • Boyle’s Law: P1V1=P2V2P_1V_1 = P_2V_2 (Pressure inversely proportional to volume).

  • Charles’s Law: V1/T1=V2/T2V_1/T_1 = V_2/T_2 (Volume proportional to temperature).

  • Dalton’s Law: Total pressure = sum of partial pressures.

  • Graham’s Law: Lighter gases diffuse faster than heavier ones.

10. Solutions

  • Definitions:

    • Solute: Substance dissolved.

    • Solvent: Substance doing the dissolving.

  • Concentration: Measured in molarity (M=moles of solute/liters of solutionM = \text{moles of solute} / \text{liters of solution}).

  • Dilution: M1V1=M2V2M_1V_1 = M_2V_2.

11. Chemical Reactions

  • Types:

    1. Synthesis: A+B→ABA + B \to AB.

    2. Decomposition: AB→A+BAB \to A + B.

    3. Single Replacement: A+BC→AC+BA + BC \to AC + B.

    4. Double Replacement: AB+CD→AD+CBAB + CD \to AD + CB.

    5. Combustion: Hydrocarbon+O2→CO2+H2OHydrocarbon + O_2 \to CO_2 + H_2O.

  • Balancing Equations: Conservation of mass ensures equal atoms on both sides.

12. Titrations

  • Definition: Analytical method to determine unknown concentration.

  • Key Point: At the equivalence point, moles of titrant = moles of analyte.

13. Oxidation-Reduction (Redox) Reactions

  • Oxidation: Loss of electrons.

  • Reduction: Gain of electrons.

  • Balancing Steps:

    1. Assign oxidation numbers.

    2. Write half-reactions for oxidation and reduction.

    3. Balance atoms and charge.

    4. Combine half-reactions.

14. Spectroscopy

  • Purpose: Studies how matter interacts with electromagnetic radiation.

  • Types:

    • UV-Vis: Measures electronic transitions.

    • Infrared: Identifies molecular vibrations.

    • Microwave: Measures rotational transitions.

  • Beer-Lambert Law: A=εlcA = \varepsilon lc, relates absorbance to concentration.