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
Metals | Nonmetals | Cation Name | Anion Name | |
|---|---|---|---|---|
1A | 5A | Li⁺ | N³⁻ | |
Na⁺ | P³⁻ | |||
K⁺ | O²⁻ | |||
2A | 6A | Mg²⁺ | S²⁻ | |
Ca²⁺ | F¯ | |||
Ba²⁺ | Cl⁻ | |||
3A | 7A | Al³⁺ | Br¯ | |
I¯ |
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
CaO - Calcium oxide
KBr - Potassium bromide
Al₂O₃ - Aluminum oxide
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
Prefix | Number |
|---|---|
mono | 1 |
di | 2 |
tri | 3 |
tetra | 4 |
penta | 5 |
hexa | 6 |
hepta | 7 |
octa | 8 |
nona | 9 |
deca | 10 |
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
Electron Groups | Geometry | Example | Bond Angle* |
|---|---|---|---|
2 | Linear | CO₂ | 180° |
3 | Trigonal planar | H₂CO | 120° |
4 | Tetrahedral | CH₄ | 109° |
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
Type | Strength |
|---|---|
Ionic bond | Strong |
Covalent bond | Moderate |
Hydrogen bond | Strongest |
Dipole-dipole | Moderate |
Dispersion forces | Weak |
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.