Chemistry Fundamentals: Atoms, Reactions, and Bonding

Posted on Mar 4, 2026 in Chemistry

Chapter 1: Atomic Structure and Matter

  • Atomic models:
    • Dalton: Solid sphere model
    • Thomson: Plum pudding model, cathode ray experiment
    • Rutherford: Gold foil experiment leading to the nucleus
    • Bohr: Planetary energy levels
    • Modern: Wave mechanical or electron cloud model
  • Subatomic particles:
    • Proton: +1 charge, ~1 amu, located in the nucleus
    • Neutron: 0 charge, ~1 amu, located in the nucleus
    • Electron: −1 charge, ~0 amu, located in the electron cloud
  • Isotopes: Atoms of the same element with a different number of neutrons.
  • Electron arrangement:
    • Ground state: Lowest energy level
    • Excited state: Higher energy level
  • Ions:
    • Metallic ions: Lose electrons to become positive cations
    • Nonmetallic ions: Gain electrons to become negative anions
  • Types of matter:
    • Pure substances: Elements and compounds
    • Mixtures: Homogeneous and heterogeneous

Chapter 2: Chemical Language and Reactions

  • Chemical symbols: Represent elements using 1–2 letters; the first letter is always capitalized and the second is lowercase (e.g., Na, Cl).
  • Chemical formulas: Show the types and numbers of atoms in a compound. Subscripts apply only to the element directly preceding them.
  • Compounds: Substances made of two or more different elements chemically bonded in fixed ratios with properties different from their constituent elements.
  • Polyatomic ions: Groups of atoms that act as a single ion. Use parentheses when more than one is needed, such as Ca(NO3)2.

Writing and Naming Chemical Formulas

  • Ionic compounds: Metal + nonmetal or polyatomic ion. The total charge must equal zero; use the criss-cross method if needed.
  • Covalent (molecular) compounds: Nonmetal + nonmetal. Prefixes indicate the number of atoms.
  • Naming ionic compounds: Name the metal first, then the nonmetal with an –ide ending. If the metal has multiple charges, use the Stock system (Roman numerals). Example: FeCl2 is iron(II) chloride; FeCl3 is iron(III) chloride.
  • Naming with polyatomic ions: Name the metal followed by the name of the polyatomic ion. Do not change the polyatomic ion’s name.
  • Naming covalent compounds: Use prefixes (mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-). Do not use mono- on the first element; the second element ends in –ide.

Understanding Chemical Reactions

  • Chemical reactions: A chemical change where substances rearrange to form new substances.
  • Chemical equations: Reactants → Products. Coefficients show the number of molecules or moles.
  • Balancing equations: Must obey the Law of Conservation of Mass. Ensure the same number of each atom exists on both sides. Change coefficients only, never subscripts.
  • Types of reactions:
    • Synthesis (Combination): A + B → AB
    • Decomposition: AB → A + B
    • Single Replacement: A + BC → AC + B (check activity series)
    • Double Replacement: AB + CD → AD + CB (often forms a precipitate, gas, or water)

Chapter 3: The Mole and Calculations

  • Formula mass: The sum of atomic masses in a formula.
  • Gram formula mass: The mass of 1 mole (g/mol).
  • Percent composition: (Mass of element ÷ total mass) × 100.
  • Hydrates: A compound containing water (e.g., ·xH2O).
  • Percent water in a hydrate: (Mass of water ÷ total mass) × 100.
  • The mole: Represents 6.02 × 1023 particles.
  • Mole conversions:
    • Grams to moles: Divide by molar mass
    • Moles to grams: Multiply by molar mass
  • Mole ratios: Determined by the coefficients from balanced equations.

Chapter 4: Energy and Phases

  • Phases: Solid, liquid, and gas.
  • Phase changes: Melting, freezing, vaporization, condensation, sublimation, and deposition.
  • Kinetic energy: The energy of motion.
  • Potential energy: Stored energy.
  • Temperature: The average kinetic energy of particles.
  • Temperature scales:
    • Celsius (°C)
    • Kelvin (K = °C + 273)
  • Heat: Energy transfer.
  • Heat flow: Always flows from hot to cold.
  • Heating and cooling curves: Graphical representations showing phase changes and energy changes.

Chapter 5: The Periodic Table

  • History: Developed from Mendeleev’s work to the modern table.
  • Arrangement: Organized by increasing atomic number.
  • Classifying elements:
    • Metals: Shiny, conductive, and malleable.
    • Nonmetals: Dull, brittle, and poor conductors.
    • Metalloids: Possess mixed properties of both metals and nonmetals.
  • Groups (columns): Elements with similar chemical properties.
  • Periods (rows): Elements with the same number of energy levels.
  • Trends:
    • Electronegativity: Increases across a period and up a group.
    • Ionization energy: Increases across a period and up a group.
  • Allotropes: Different forms of the same element (e.g., Carbon as diamond or graphite; Oxygen as O2 or O3).

Chapter 6: Bonding and Polarity

  • Energy and chemical bonds: Bonds form when atoms lower their potential energy. Energy is released when bonds form and absorbed when bonds break.
  • Endothermic reactions: Absorb energy from surroundings; products have higher energy than reactants.
  • Exothermic reactions: Release energy to surroundings; products have lower energy than reactants.
  • Lewis dot diagrams: Show valence electrons only. Used to predict bonding and polarity.
  • Octet rule: Atoms tend to gain, lose, or share electrons to achieve 8 valence electrons. Exceptions include Hydrogen (2) and Helium (2).
  • Ion formation: Metals lose electrons to form cations; nonmetals gain electrons to form anions.
  • Isoelectronic particles: Different atoms or ions that have the same number of electrons.

Types of Chemical Bonds

  • Ionic bonds: Form between metals and nonmetals via the transfer of electrons. Strong electrostatic attraction between opposite charges.
  • Covalent bonds: Form between nonmetals via the sharing of electrons. Can be single, double, or triple bonds.
  • Electronegativity differences:
    • Ionic: Large difference
    • Polar covalent: Moderate difference
    • Nonpolar covalent: Little or no difference
  • Polar covalent bonds: Unequal sharing of electrons resulting in partial charges (δ+ and δ−).
  • Nonpolar covalent bonds: Equal sharing of electrons with no partial charges.

Molecular Polarity and Intermolecular Forces

  • Determining molecular polarity: Check bond polarity and molecular shape. Symmetrical molecules are nonpolar, while asymmetrical molecules are polar.
  • Intermolecular forces (IMF): Attractions between molecules (weaker than chemical bonds).
    • Hydrogen bonding: Strongest IMF (occurs when H is bonded to N, O, or F).
    • Dipole-dipole: Attraction between polar molecules.
    • Van der Waals (London dispersion): Weakest IMF, present in all molecules.
  • Metallic bonding: Positive metal ions surrounded by a “sea” of delocalized electrons. This explains the conductivity, malleability, and ductility of metals.