Chapter 6 1-4 sections

Chapter 6: Chemical Bonding

Preview

  • Overview of chemical bonding concepts and objectives.

Lesson Starter

  • Analogy of crowded elevator to explain atomic interactions.

  • When atoms approach, outer electrons repel while being attracted to surrounding nuclei.

  • The attraction degree influences the type of chemical bond formed.

Objectives

  1. Define chemical bonds.

  2. Explain why most atoms form bonds.

  3. Describe ionic and covalent bonds.

  4. Explain the nature of bonding as usually neither purely ionic nor purely covalent.

  5. Classify bonding types based on electronegativity differences.

Importance of Chemical Bonds

  • Atoms generally have high potential energy when independent.

  • Stable atomic configurations lower potential energy.

  • Bonding leads to more stable arrangements of matter, reducing potential energy.

Ionic Bonding

  • Involves transfer of electrons from metal to non-metal.

  • Ionic bonds create charged particles called ions:

    • Metal (Na) loses an electron, becoming Na+.

    • Non-metal (Cl) gains an electron, becoming Cl-.

Example of Ionic Bonding

  • Lithium (Li) loses 1 electron (Li+) to bond with Fluorine (F) that gains 1 electron (F-).

Covalent Bonding

  • Occurs primarily between non-metals.

  • Defines electron sharing, depicted through dot-and-cross diagrams (Lewis Structures).

Nature of Chemical Bonds

  • Bonds between different elements rarely fall into purely ionic or purely covalent categories.

  • Bond type depends on electronegativity, a measure of an atom's ability to attract electrons.

Electronegativity & Bond Classification

  • Metals (Groups 1-3) have low electronegativity;

  • Non-metals (Groups 5-7) have high electronegativity;

  • Ionic bonds likely if electronegativity difference exceeds 1.7.

Sample Problem

  • Determining bond type between sulfur (S) and other elements using electronegativity:

    • H: 2.1, Cl: 3.0, Cs: 0.7.

    • Bond types identified:

      • H-S: 0.4 (polar-covalent)

      • S-Cs: 1.8 (ionic)

      • Cl-S: 0.5 (polar-covalent)

Covalent Bonding Overview

Objectives of Section 2

  • Define molecules and molecular formulas.

  • Explain potential energy relationships in bond formation.

  • Establish the octet rule and its significance.

  • Detail steps for constructing Lewis structures and their application in bonding.

Definition of Molecular Compounds

  • Molecules: Neutral groups of atoms in covalent bonds.

  • Molecular compounds: Compounds with simplest units as molecules.

Chemical Formulas

  • Chemical formulas represent relative atom counts in a compound.

  • Molecular formulas specify types and counts of atoms in a single molecule.

Comparison of Molecules

  • Monatomic: 1 atom (e.g., Sodium, Na)

  • Diatomic: 2 atoms (e.g., Fluorine, F₂)

  • Polyatomic: More than 2 atoms (e.g., Water, H₂O)