Unit Quiz: Bonding


Chapter 8: Ionic Compounds

Key Terms

  • Chemical bonds – The force that holds atoms together.

  • Ions

    • Anion – Negatively charged ion.

    • Cation – Positively charged ion.

  • Electrolyte – A substance that dissolves in water to conduct electricity.

  • Ionic bond – Bond formed between a metal and a nonmetal by transferring electrons.

  • Crystal lattice – Repeating pattern of ions in an ionic compound.

  • Monatomic ion – Ion consisting of one atom.

  • Polyatomic ion – Ion composed of multiple atoms.

  • Oxyanion – A polyatomic ion containing oxygen.

Key Concepts

  • Identifying Ion Charges

    • Atoms gain or lose electrons to achieve a stable noble gas electron configuration.

    • Charge is determined by the number of electrons lost/gained (usually follows group trends).

  • Octet Rule

    • Atoms tend to gain, lose, or share electrons to achieve 8 valence electrons.

  • Chemical Formulas for Ionic Bonds

    • Combine two ions to create a neutral compound.

    • Example: Na⁺ + Cl⁻ → NaCl

  • Naming Ionic Compounds

    • Metal name + Nonmetal root + "-ide" (for monatomic ions).

    • Polyatomic ions retain their specific names (e.g., sulfate, nitrate).

  • Writing Formulas from Names

    • Use ion charges to balance the overall charge of the compound.

    • Example: Calcium chloride → CaCl₂ (Ca²⁺ and Cl⁻ balance out).

Chapter 9: Covalent Bonding

Key Terms

  • Covalent bond – Bond where atoms share electrons.

  • Lewis structure – Diagram showing valence electrons and bonding.

  • Molecule – A group of covalently bonded atoms.

  • Polar covalent bond – Unequal sharing of electrons.

  • Sigma bond (σ bond) – Single covalent bond formed by direct orbital overlap.

  • Structural formula – Representation of a molecule showing bonding.

  • VSEPR model – (Valence Shell Electron Pair Repulsion Model) Determines molecular shape.

  • Acids

    • Binary acid – H + one other element (e.g., HCl).

    • Oxyacid – Acid with hydrogen, oxygen, and another element (e.g., H₂SO₄).

Key Concepts

  • Differences Between Ionic and Covalent Bonds

    • Ionic: Transfer of electrons, forms ions.

    • Covalent: Sharing of electrons, forms molecules.

  • Single, Double, and Triple Bonds

    • Single: 2 shared electrons.

    • Double: 4 shared electrons.

    • Triple: 6 shared electrons.

  • Naming Molecular Compounds

    • Use prefixes (mono-, di-, tri-, etc.) to indicate the number of atoms.

    • Example: CO₂ → Carbon dioxide.

  • Writing Molecular Formulas

    • Convert name to formula using prefixes.

  • Drawing Molecular Structures

    • Use Lewis structures to show bonds and lone pairs.

  • Exceptions to the Octet Rule

    1. Molecules with an odd number of valence electrons (e.g., NO).

    2. Molecules that form with fewer than 8 valence electrons (e.g., BF₃).

    3. Expanded octet molecules (more than 8 valence electrons) (e.g., SF₆).

  • VSEPR Shapes

    • Shape depends on bonded atoms & lone pairs.

    • Example:

      • Linear (CO₂)

      • Trigonal Planar (BF₃)

      • Tetrahedral (CH₄)

      • Bent (H₂O)

  • Electronegativity and Bond Type

    • Ionic: Large difference in electronegativity.

    • Polar Covalent: Moderate difference in electronegativity.

    • Nonpolar Covalent: Little to no difference in electronegativity.

  • Dipole Polarity

    • The more electronegative atom pulls electrons closer, creating a dipole (partial charges).

    • Example: H₂O has a negative pole near oxygen and a positive pole near hydrogen.

  • Acids in Water

    • Acids donate H⁺ ions when dissolved in water.

    • Example: HCl → H⁺ + Cl⁻

  • Naming Acids

    • Binary Acids: "Hydro-" + root + "-ic acid" (e.g., HCl → Hydrochloric acid).

    • Oxyacids:

      • "-ate" → "-ic acid" (H₂SO₄ → Sulfuric acid).

      • "-ite" → "-ous acid" (H₂SO₃ → Sulfurous acid).