Introduction to Chemistry

Periodic Table

• Understand names of elements, groups, periods
  • Group 1 (alkali metals): Hydrogen (H), Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Cesium (Cs)
    • Exception: Hydrogen is a non-metal; others are highly reactive metals
    • Have 1 valence electron, form +1 ions
  • Group 2 (alkaline earth metals): Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba)
    • Reactive metals that form +2 ions
    • Have 2 valence electrons
• Transition Metals:
  • Variable charges (e.g., Iron Fe^2+/Fe^3+, Copper Cu^1+/Cu^2+, Zinc Zn^2+)
• Inner Transition Metals: Lanthanides, Actinides
• Representative Elements:
  • Group 13 (3A): Boron (B), Aluminum (Al), Gallium (Ga), Indium (In), Thallium (Tl)
  • Group 14 (4A): Carbon (C), Silicon (Si), Germanium (Ge), Tin (Sn), Lead (Pb)
  • Group 15 (5A): Nitrogen (N), Phosphorus (P), Arsenic (As), Antimony (Sb), Bismuth (Bi)
    • Often form -3 anions (e.g., Nitride N^3-, Phosphide P^3-)
  • Group 16 (6A): Oxygen (O), Sulfur (S), Selenium (Se), Tellurium (Te), Polonium (Po)
    • Often form -2 anions (e.g., Oxide O^2-, Sulfide S^2-)
  • Group 17 (7A, Halogens): Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), Astatine (At)
    • Highly reactive non-metals, often form -1 anions (e.g., Fluoride F^-, Chloride Cl^-)
  • Group 18 (8A, Noble Gases): Helium (He), Neon (Ne), Argon (Ar), Krypton (Kr), Xenon (Xe), Radon (Rn)
    • Inert gases with complete valence shells

Chemical Bonds and Compounds

• Ionic Bonds: Electrons are transferred from metals to non-metals (e.g., NaCl)
• Covalent Bonds: Electrons are shared between non-metals (can be polar or non-polar)
• Polyatomic Ions: Know names and charges (e.g., sulfate SO4^2-, nitrate NO3^-)
• Naming Conventions:
  • Acids:
    • -ate ions become -ic acids (e.g., H2SO4 = sulfuric acid)
    • -ite ions become -ous acids (e.g., H2SO3 = sulfurous acid)
    • -ide ions become hydro-ic acids (e.g., HCl = hydrochloric acid)
  • Ionic Compounds (e.g., sodium chloride NaCl)
  • Molecular Compounds: Use prefixes (e.g., sulfur dioxide SO2)

Periodic Trends and Properties

• Electronegativity: Increases across a period (left to right) and up a group
• Metallic Character: Increases down a group and to the left
• Conductivity:
  • Metals: Good conductors
  • Non-metals: Insulators
  • Metalloids: Semi-conductors
• Diatomic Elements: H2, N2, O2, F2, Cl2, Br2, I2
• Allotropes: Different forms of the same element (e.g., Graphite and Diamond for Carbon)

Atomic Structure and Isotopes

• Atom’s Nucleus: Contains protons and neutrons
• Electrons: Occupy energy levels; valence electrons are in the outermost shell
• Ions: Atoms that have gained or lost electrons
• Isotopes: Atoms with the same number of protons but different numbers of neutrons
• Mass Number: Sum of protons and neutrons

Chemical Reactions

• Combustion Reactions: Hydrocarbon + O2 → CO2 + H2O
• Redox Reactions: Involve the transfer of electrons
• Acid-Base Reactions: Produce salt and water
• Precipitation Reactions: Formation of a solid from two aqueous solutions
• Single Replacement Reactions: An element in a compound is replaced by another element
• Double Replacement Reactions: Exchange of ions between two compounds

Stoichiometry

• Mole Concept: 1 mole = 6.022 x 10^23 particles
• Molar Mass: Mass of one mole of a substance (g/mol)
• Conversions: Grams to moles, moles to atoms, atoms to grams
• Percent Yield: (Actual Yield / Theoretical Yield) x 100%
• Limiting Reactants: Determines the amount of product formed

Key Equations and Conversions

• Density: Mass/Volume
• Molarity (M): Moles of solute/Liters of solution
• Gas Laws: Ideal Gas Law (PV = nRT), Boyle's Law, Charles's Law
• Unit Conversions: Distance, Volume, Time
  • 1 km = 1000 m
  • 1 m = 100 cm
  • 1 mile = 5280 feet
  • 1 inch = 2.54 cm
  • 1 liter = 1000 mL = 1 cubic decimeter
• Significant Figures: Rules to determine how many digits are meaningful
• Balancing Equations: Ensuring the same number of atoms for each element on both sides