Ionic & Molecular Compounds, Acids, Bases, Reactions

Posted on Mar 14, 2025 in Chemistry

Ionic and Molecular Compounds

  1. Atoms, Ions, and Molecules

    • Atoms: The fundamental building blocks of matter, consisting of a nucleus containing protons and neutrons, surrounded by electrons.
    • Ions: Atoms or molecules that have gained or lost electrons, resulting in a net charge. Cations are positively charged, while anions are negatively charged.
    • Molecules: Groups of two or more atoms chemically bonded together. They can be elements (O₂) or compounds (H₂O).

  2. Seven Diatomic Molecular Gases

    • Names: Hydrogen, Nitrogen, Oxygen, Fluorine, Chlorine, Bromine, Iodine.
    • Formulas: H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂.
    • These gases exist naturally as molecules rather than individual atoms due to their need for stable electron configurations.

  3. Naming and Writing Formulas for Ionic Compounds

    • Monoatomic Ions: Formed from single atoms gaining or losing electrons.
      • Examples: Na⁺ (Sodium ion), Cl⁻ (Chloride ion), Ca²⁺ (Calcium ion).
    • Polyatomic Ions: Groups of atoms that act as a single charged unit.
      • Examples: NO₃⁻ (Nitrate), SO₄²⁻ (Sulfate), NH₄⁺ (Ammonium).
    • Formula Writing: Combine cations and anions in the simplest whole-number ratio to form neutral compounds.
      • Example: NaCl (Sodium chloride), CaSO₄ (Calcium sulfate).

  4. Naming and Writing Formulas for Covalent (Molecular) Compounds

    • Use prefixes to indicate the number of atoms in the compound:
      • Mono- (1), Di- (2), Tri- (3), Tetra- (4), Penta- (5), Hexa- (6), etc.
    • Example: CO₂ (Carbon dioxide), N₂O (Dinitrogen monoxide).

  5. Single, Double, and Triple Bonds

    • Single Bonds: Two shared electrons (H₂, CH₄).
    • Double Bonds: Four shared electrons (O₂, CO₂).
    • Triple Bonds: Six shared electrons (N₂, C₂H₂).

  6. Polar and Nonpolar Bonds

    • Nonpolar Bonds: Equal electron sharing, no charge difference (O₂, N₂).
    • Polar Bonds: Unequal electron sharing due to electronegativity differences (H₂O, NH₃).

  7. Molecular Geometry & Bond Angles (VSEPR Theory)

    • Tetrahedral (109.5°): CH₄, NH₃.
    • Trigonal Planar (120°): BF₃, CO₃²⁻.
    • Linear (180°): CO₂, BeCl₂.

  8. Determining Molecular Polarity

    • Consider molecular shape and dipole movement.
    • Label δ+, δ-, and net dipole movement to determine overall polarity

Acids, Bases, and Equilibrium

  1. Define Acids and Bases

    • Arrhenius: Acids produce H⁺, bases produce OH⁻.

    • Brønsted-Lowry: Acids donate H⁺, bases accept H⁺.

  2. Identify Conjugate Acid-Base Pairs

    • Example: HCl (acid) → Cl⁻ (conjugate base).

  3. Recognizing Acids and Bases

    • Acids: H at the beginning (HCl, H₂SO₄).

    • Bases: OH at the end (NaOH, KOH).

  4. Distinguish Between Strong and Weak Acids

    • Strong Acids: Fully dissociate in water (HCl, HNO₃, H₂SO₄, HBr, HI, HClO₄).

    • Weak Acids: Partially dissociate (CH₃COOH, HF).

  5. Distinguish Between Strong and Weak Bases

    • Strong Bases: Fully dissociate in water (LiOH, KOH, NaOH, Ca(OH)₂).

    • Weak Bases: Partially dissociate (NH₃).

  6. Write Equations for Acid and Base Dissociation

    • Example: HCl → H⁺ + Cl⁻ (Strong Acid).

    • Example: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻ (Weak Base).

  7. Name and Write the Formulas for Strong Acids and Bases

    • Strong Acids: HCl, HNO₃, H₂SO₄, HBr, HI, HClO₄.

    • Strong Bases: LiOH, KOH, NaOH, Ca(OH)₂.

    • Common Weak Base: NH₃ (Ammonia).


Chemical Quantities and Reactions

  1. Avogadro’s Number

    • 1 mole = 6.022 × 10²³ particles (atoms, molecules, or formula units).

  2. Molar Mass

    • The molar mass of a substance is the mass (in grams) of one mole of its particles.
    • Found by summing the atomic masses of the elements in a compound.
      • Example: H₂O = (2 × 1.008) + (1 × 16.00) = 18.02 g/mol.

  3. Grams to Moles Conversion

    • Formula: Moles = grams / molar mass.
    • Example: 36.04 g of H₂O → 2.00 moles of H₂O.

  4. Balancing Chemical Equations

    • Adjust coefficients to ensure equal numbers of atoms on both sides of the equation.
      • Example: 2H₂ + O₂ → 2H₂O.

  5. Types of Reactions

    • Combination: A + B → AB.
    • Decomposition: AB → A + B.
    • Single Replacement: A + BC → AC + B.
    • Double Replacement: AB + CD → AD + CB.
    • Combustion: Hydrocarbon + O₂ → CO₂ + H₂O.

Solutions

  1. Definitions

    • Solute: Substance being dissolved.
    • Solvent: Substance doing the dissolving.
    • Solution: Homogeneous mixture of solute and solvent.

  2. Hydrogen Bonding

    • Strong intermolecular force occurring when H is bonded to N, O, or F.

  3. Formation of Solutions

    • Solutes dissolve in solvents based on intermolecular forces.

  4. “Like Dissolves Like” Principle

    • Polar solutes dissolve in polar solvents (NaCl in H₂O).
    • Nonpolar solutes dissolve in nonpolar solvents (Oil in gasoline).

  5. Electrolytes and Nonelectrolytes

    • Strong Electrolytes: Fully dissociate in water (NaCl).
    • Weak Electrolytes: Partially dissociate (CH₃COOH).
    • Nonelectrolytes: Do not conduct electricity (C₆H₁₂O₆).

  6. Concentration Calculations

    • Molarity (M) = moles of solute / liters of solution.
    • Percent Concentrations: Volume, mass, volume/mass calculations.

  7. Dilution Calculations

    • M₁V₁ = M₂V₂ (Initial and final concentrations and volumes).

Acids, Bases, and Equilibrium

  1. Distinguish Between Acidic, Basic, and Neutral Solutions

    • Acidic: pH < 7.

    • Neutral: pH = 7.

    • Basic: pH > 7.

  2. Complete and Balance Neutralization Reactions

    • Example: HCl + NaOH → NaCl + H₂O.

  3. Define Buffer Systems and Explain Their Function

  • A buffer consists of a weak acid and its conjugate base or a weak base and its conjugate acid.

  • Buffers resist changes in pH by neutralizing added acids or bases.

  1. Perform pH Calculations

  • pH = -log[H₃O⁺].

  • Kw = [H₃O⁺][OH⁻] = 1.0 × 10⁻¹⁴.

  • [H₃O⁺] = 10⁻ᵖᴴ, [OH⁻] = 10⁻ᵖᴼᴴ.