Chemistry Regents Review Notes

Unit 1 & 2: Math, Measurement & Matter

• Particle Diagrams:
  • Element: Ex. Mg, H2
  • Compound: Ex. HF (two or more different elements chemically combined)
  • Mixture: Ex. Mg & Zn & NaCl (two or more different substances not chemically combined)
• Separation Techniques:
  • Filtration: Insoluble solid from liquid (e.g., sand from salt water).
  • Distillation: Separates liquids based on boiling points.
• Phases of Matter:
  • Melting/Freezing
  • Boiling/Condensing
  • Sublimation: $CO<em>2(s) \rightarrow CO</em>2(g)$
  • Deposition: $CO<em>2(g) \rightarrow CO</em>2(s)$
  • Endothermic (add heat), Exothermic (remove heat)
• Changes:
  • Physical: Does not change chemical properties (e.g., phase changes, cutting, dissolving).
  • Chemical: New substance formed (e.g., burning, rusting, reacting).
• Significant Figures:
  • Start counting at first nonzero digit (e.g. 0.000789 has 3 sig fig).
  • Zeros trapped between non zero digits are significant (ex. 3006 has 4 sig fig).
  • Trailing zeros after the last non zero digit are significant if decimal point is present (ex. 0.0005600 has 4 sig fig but 5600 has only 2 sig fig).
• Rounding:
  • Addition/Subtraction: Least precise place value.
  • Multiplication/Division: Least number of sig figs.
• Unit Conversions:
  • milli- to base unit & kilo- to base unit.
  • Larger unit (m → km), smaller number.
  • Smaller unit (m → mm), larger number.
• Temperature Conversions:
  • Freezing/melting point of water = $0°C$
  • Boiling/condensation point of water = $100°C$
  • $K = °C + 273$
• Density: $D = \frac{m}{v}$
• Percent Error: $\% error = \frac{\left| accepted - measured \right|}{accepted} \times 100$

Unit 3: Moles, Stoichiometry, Naming and Formula Writing

• Chemical Formulas: Counting Atoms
  • Coefficient: moles of compound
  • Subscript: moles of each atom
• Molar Mass: Gram formula mass (GFM), mass of 1 mole.
• Calculating Moles: $Moles = \frac{given \ mass(g)}{gfm}$
• Types of Reactions:
  • Synthesis: A + B → AB
  • Decomposition: BA → A + B
  • Single Replacement: A + BC → B + AC
  • Double Replacement: AB + CD → AD + CB
• Mole Ratios: Comparing moles in a question.
• Conservation of Mass: Reactants = Products.
• Balancing Reactions: Keep polyatomics together.
• Molecular to Empirical Formula: Divide by greatest common factor.
• Empirical to Molecular Formula: Find ratio using molecular mass.
• Percent Composition: From Table T
• Naming Binary Compounds: First element name + second element "-ide".
• Writing Chemical Formulas: Drop and swap charges.
• Multiple Charges: Use Roman numerals to indicate charge.
• Naming Tertiary Compounds: Use Table E for polyatomic ions.

Unit 4: Atomics

• Atomic Theory:
  1. Solid sphere.
  2. Uniform positive charge with embedded electrons.
  3. Small positive nucleus, mostly empty space.
  4. Electrons orbit in energy levels.
  5. High probability of finding electrons in orbitals.
• Thomson vs. Rutherford:
  • Thomson: discovered electrons
• Rutherford:
  • positive charge in nucleus, electrons outside
  • Gold Foil Experiment:
    1. Small dense positive nucleus.
    2. Mostly empty space.
• Current Model: Wave mechanical model (orbitals).
• Subatomic Particles:
  • Atomic number = # of protons
  • Protons = electrons in neutral atom
  • Neutrons = mass # - protons
• Atomic mass: weighted average of isotopes.
• Ions: Cations (+), Anions (-).
• Isotopes: Same protons, different neutrons.
• Bohr Diagrams: Electron arrangement.
• Excited vs. Ground State: Electron configuration.
• Light Production: Electrons from excited to ground state.
• Lewis Dot Diagrams:
  • Atoms: Valence electrons as dots.
  • Ions: Metals (no dots), Nonmetals (8 dots).

Unit 5: Periodic Table

• Organization:
  • Mendeleev: by mass
  • Mosley: by atomic number
  • Periods: Same electron shells.
  • Groups: Same valence electrons, similar properties.
• Groups:
  • 1: Alkali Metals (most reactive metal)
  • 2: Alkaline Earth Metals
  • 3-12: Transition Metals (colored solutions)
  • 17: Halogens (most reactive nonmetal)
  • 18: Noble Gases (inert, full valence shell)
• Properties:
  • Metals: Malleable, ductile, conductors, luster, lose electrons, solid (except Hg).
  • Sea of mobile electrons: conducts electricity
  • Metallic Character: Closer to Fr (Francium) is more metallic.
  • Metalloids: Semiconductors, luster, brittle.
  • Nonmetals: Poor conductors, brittle, dull, gain electrons.
• Periodic Trends:
  • Atomic radius, electronegativity, ionization energy.
  • Across period: increased proton pull.
  • Down group: more electron shells.
  • Atomic Radius: Decreases across, increases down.
  • Electronegativity: Attraction for electrons, F (Fluorine) is most electronegative.
  • Ionization Energy: Energy to remove electron.

Unit 6: Bonding

• Bonding:
  • Breaking Bonds: Absorbs energy.
  • Making Bonds: Releases energy.
• Octet Rule: Full valence shell (8 electrons, 2 for H (Hydrogen)).
• Types of Bonds:
  • Ionic: Metal and nonmetal, transfer of electrons.
  • Covalent: Nonmetal and nonmetal, sharing electrons.
  • Metallic: Metals, sea of mobile electrons.
• Electrolytes: Conduct electricity when dissolved in water.
• Lewis Dot Diagrams:
  • Ionic: Metals (no dots), Nonmetals (8 dots).
  • Covalent: Shared electron pairs.
• Covalent Bonds:
  • Nonpolar: Equal sharing, E.N.D. 0-0.4.
  • Polar: Unequal sharing, E.N.D. > 0.4.
• Molecular Polarity:
  • Nonpolar: Symmetrical charge.
  • Polar: Asymmetrical charge.
• Molecular Geometry: Linear, bent, pyramidal, tetrahedral.
• Intermolecular Forces:
  • Weak forces between molecules.
  • Hydrogen Bonding: H and FON (Fluorine, Oxygen, Nitrogen).
  • Stronger IMF, higher melting/boiling point.

Unit 7: Heat

• Heat: Flows from hot to cold.
• Average kinetic energy = temperature
  • Endothermic: absorb heat, +∆H.
  • Exothermic: release heat, -∆H.
• Phases:
  • Solid: Definite shape/volume, strongest attraction.
  • Liquid: Definite volume, takes container shape, moderate attraction.
  • Gas: No definite shape/volume, weakest attraction.
• Specific Heat: Heat to raise 1g of substance by 1°C. Water has high specific heat.
• Heat of Fusion: Solid to liquid.
• Heat of Vaporization: Liquid to gas.
• Heating/Cooling Curve:
  • Kinetic energy increases during phases; potential energy stays the same.
  • Potential energy increases during phase changes; kinetic energy stays the same.
• Calculating Heat:
  • Change in temperature: $q = mc\Delta T$
  • Melting/Freezing: $q = mH_f$
  • Boiling/Vaporizing/Condensing: $q = mH_v$