Chapter 6: Ionic and Molecular Compounds

6.1 Ions: Transfer of Electrons

Atoms achieve stability by acquiring an octet of eight valence electrons, either through losing, gaining, or sharing electrons. This process leads to the formation of ionic or covalent bonds.

• Ionic Bonds: Formed when valence electrons are transferred from a metal atom to a nonmetal atom.

• Covalent Bonds: Formed when nonmetal atoms share electrons to attain a noble gas arrangement.

• Cations: Positively charged ions formed by metals losing electrons. Named after the element (e.g., sodium ion).

• Anions: Negatively charged ions formed by nonmetals gaining electrons. Named by adding 'ide' to the root of the element name (e.g., chloride).

Example: Sodium (Na) loses one electron to form Na+; Chlorine (Cl) gains one electron to form Cl-.

6.2 Ionic Compounds

Ionic compounds are composed of positive and negative ions held together by strong electrical attractions called ionic bonds.

• Properties: High melting points, solid at room temperature, consist of positive and negative ions.

• Charge Balance: The total positive charge must equal the total negative charge in the compound.

Example: Sodium chloride (NaCl) is formed from Na+ and Cl-.

6.3 Naming and Writing Ionic Formulas

The correct naming and formula writing for ionic compounds follows specific rules:

• The metal (cation) is named first, followed by the nonmetal (anion) with 'ide' ending.

• For transition metals with variable charges, a Roman numeral indicates the charge (e.g., iron(III) chloride).

• Subscripts are used in formulas to balance the charges.

Example: K2O is named potassium oxide.

6.4 Polyatomic Ions

Polyatomic ions are groups of covalently bonded atoms with an overall ionic charge. Most are negatively charged except ammonium (NH4+).

• Common Polyatomic Ions: sulfate (SO42-), nitrate (NO3-), phosphate (PO43-), ammonium (NH4+).

• Names ending in 'ate' have more oxygens than those ending in 'ite'.

• Formulas are written using charge balance rules, with parentheses if more than one polyatomic ion is needed.

Example: Magnesium nitrate: Mg(NO3)2.

6.5 Molecular Compounds: Sharing Electrons

Molecular compounds are formed when nonmetal atoms share electrons, resulting in covalent bonds. Prefixes are used in naming to indicate the number of atoms.

• Prefixes: mono-, di-, tri-, tetra-, penta-, etc.

• The first nonmetal is named by its element name; the second nonmetal uses the root plus 'ide' and a prefix.

Example: NCl3 is named nitrogen trichloride; P2O5 is diphosphorus pentoxide.

6.6 Lewis Structures for Molecules and Polyatomic Ions

Lewis structures represent molecules by showing valence electrons arranged to give octets. Shared pairs (bonding pairs) are shown as lines or dots; lone pairs are placed on the outside.

• Central atom is usually the least numerous in the formula.

• Double and triple bonds occur when atoms share two or three pairs of electrons.

• Exceptions: Hydrogen needs only two electrons; some elements can have expanded octets.

Example: Methane (CH4) has a central carbon atom bonded to four hydrogens.

6.7 Electronegativity and Bond Polarity

Electronegativity is the ability of an atom to attract shared electrons. It increases across a period and decreases down a group.

• Bond Types:

  • Nonpolar covalent: Electronegativity difference 0.0–0.4

  • Polar covalent: Electronegativity difference 0.5–1.8

  • Ionic: Electronegativity difference >1.8

• Dipole: Separation of charges in a polar bond, indicated by δ+ and δ-.

Example: HCl is a polar molecule; NaCl is ionic.

6.8 Shapes and Polarity of Molecules

VSEPR theory predicts the three-dimensional structure of molecules based on electron group repulsion.

• Linear: 2 electron groups, bond angle 180°

• Trigonal planar: 3 electron groups, bond angle 120°

• Tetrahedral: 4 electron groups, bond angle 109.5°

• Bent: 2 bonds and 1 or 2 lone pairs

• Trigonal pyramidal: 3 bonds and 1 lone pair

Polarity depends on both bond polarity and molecular shape. If dipoles cancel, the molecule is nonpolar; if not, it is polar.

6.9 Intermolecular Forces in Compounds

Intermolecular forces determine physical properties like melting and boiling points.

• Ionic Bonds: Strongest, highest melting points

• Hydrogen Bonds: Strong dipole-dipole attractions between H and F, O, or N

• Dipole–Dipole: Attractions between polar molecules

• Dispersion Forces: Weak attractions between nonpolar molecules

Example: Water exhibits hydrogen bonding; KCl is ionic.

Tables and Data

Table: Types of Particles and Bonds in Compounds

Table: Electronegativity Differences and Bond Types

Table: Common Polyatomic Ions

Key Equations

• Charge Calculation:

• Lewis Structure Octet:

Relevant Images

Textbook Cover: The cover image below is directly relevant as it identifies the source textbook for these notes, "General, Organic, and Biological Chemistry: Structures of Life."