Chapter 6-annotated

Ionic and Molecular Compounds

Chapter Overview

• Focus on the formation and naming of ionic and molecular compounds.

Key Topics

• Forming ions

• Writing and naming ionic compounds

• Polyatomic ions

• Writing and naming molecular compounds

• Lewis structures

• Electronegativity and bond polarity

• Shapes of molecules

• Attractive forces between molecules

The Octet Rule

• Atoms are stable with 8 electrons in their valence shell, mimicking noble gases.

• Atoms achieve stability by sharing or exchanging electrons:

  • Covalent Bonds: Atoms share electrons.

  • Ionic Bonds: Atoms lose (metals) or gain (nonmetals) electrons.

Positive Ions

• Metals easily lose electrons, forming cations (e.g. Na⁺, Mg²⁺).

• Once they lose electrons, they are no longer referred to as atoms.

Negative Ions

• Nonmetals gain electrons to form anions (e.g. Cl⁻, O²⁻).

• Named with the suffix -ide.

Ionic Charges from Group Numbers

• Group 1 (1A): +1

• Group 2 (2A): +2

• Group 16 (6A): -2 (16 - 18)

• Group 15 (5A): -3 (15 - 18)

• Group 4 (4A): Typically do not form ions.

Ionic Charge of Main Group Elements

Common Ions

Ionic Compounds

• Composed of cations and anions held by ionic bonds.

• High melting and boiling points; solids at room temperature (e.g. NaCl).

Charge Balance in Ionic Compounds

• The sum of positive and negative charges must equal zero.

• The metal symbol precedes the nonmetal.

• Example 1: Na⁺ + O²⁻ → Two Na⁺ ions bond with one O²⁻ ion: Na₂O.

• Example 2: Al³⁺ + Cl⁻ → One Al³⁺ ion bonds with three Cl⁻ ions: AlCl₃.

Naming Ionic Compounds

• Metal name remains the same; nonmetal name changes to -ide.

• Example: SrBr₂ is named Strontium bromide.

Practice Naming Compounds

1. CaO - Calcium oxide

2. KBr - Potassium bromide

3. Al₂O₃ - Aluminum oxide

4. MgCl₂ - Magnesium chloride

Transition Metals

• Transition metals can have multiple positive charges; use Roman numerals for distinction.

• Example: Fe²⁺ and Fe³⁺ = Iron(II) and Iron(III).

• CuCl₂ = Copper(II) chloride.

Polyatomic Ions

• Charged groups of atoms, often with oxygen.

• Common endings: -ate (more oxygen) and -ite (fewer oxygen).

• Example: Chlorine derivatives:

  • Cl⁻: chloride

  • ClO⁻: hypochlorite

  • ClO₂⁻: chlorite

  • ClO₃⁻: chlorate

  • ClO₄⁻: perchlorate

Naming with Polyatomic Ions

• Cation first, anion second; no prefixes used.

• Example: NaNO₃ = Sodium nitrate, Ca₃(PO₄)₂ = Calcium phosphate.

Practice Naming Polyatomic Compounds

• Cu(NO₃)₂ = Copper(II) nitrate

• Fe(OH)₃ = Iron(III) hydroxide

• CaCO₃ = Calcium carbonate

• NaNO₂ = Sodium nitrite

Molecular Compounds

• Formed between two or more nonmetals through covalent bonds.

• Naming format: Prefix + first element + prefix + second element (-ide).

• Example: CO₂ = Carbon dioxide.

Common Prefixes for Molecular Naming

Lewis Structures

• Illustrate electron sharing between atoms to achieve octets.

• Shared pairs: Electrons involved in bonds; lone pairs: non-bonding electrons.

• Diatomic elements: H₂, N₂, O₂, etc.

Drawing Lewis Structures

• Example: Methane (CH₄) - Carbon needs 4 electrons, surrounded by 4 hydrogens.

Electronegativity and Bond Polarity

• Measure of an atom's ability to attract electrons in a bond.

• Increases across the periodic table; decreases down a group.

• Fluorine is the most electronegative element.

Bond Polarity

• Non-polar covalent: No difference in EN (ΔEN = 0).

• Polar covalent: Small difference in EN.

• Ionic: Large difference in EN.

Shapes and Polarity of Molecules

• VSEPR Theory (Valence Shell Electron Pair Repulsion) predicts molecular shapes.

• The geometry depends on the number of electron groups: linear (2), trigonal planar (3), tetrahedral (4).

Molecular Shapes Summary

Polarity of Molecules

• Nonpolar: Symmetrical molecules cancel out polar bonds (e.g., CO₂).

• Polar: Asymmetrical charge distribution leads to polarity.

Attractive Forces in Compounds

• Strength of forces varies:

  • Ionic bonds: Strongest.

  • Covalent bonds: Moderate.

  • Intermolecular forces: Weaker (e.g., hydrogen bonding, dipole-dipole, dispersion forces).

Summary of Attractive Forces

Review

• Ions form through the gain or loss of electrons.

• Naming conventions vary between ionic and molecular compounds.

• Lewis structures depict valence electron arrangements.

• Electronegativity assesses bond polarity, influencing molecular shapes and intermolecular forces.