Ch 6 - Ionic and Molecular Compounds: Structure, Naming, and Properties

Chapter 6: Ionic and Molecular Compounds

Introduction to Compounds

Most elements, except noble gases, are found in nature as compounds. Compounds are formed when atoms of different elements combine through chemical bonds to achieve greater stability, often by attaining a noble gas electron configuration.

• Noble gases: Stable, rarely form compounds under standard conditions.

• Chemical bonds: Attractive forces between atoms, such as ionic or covalent bonds, that involve the transfer or sharing of valence electrons.

Ionic vs. Covalent Compounds

• Ionic compounds: Electrons are transferred from metals to nonmetals, resulting in the formation of ions.

• Covalent (molecular) compounds: Electrons are shared between nonmetal atoms.

Example: Sodium chloride (NaCl) is an ionic compound; water (H2O) is a molecular compound.

Transfer of Electrons and Ion Formation

Atoms form ions by gaining or losing electrons:

• Cations: Positively charged ions formed when atoms lose electrons (typically metals).

• Anions: Negatively charged ions formed when atoms gain electrons (typically nonmetals).

The transfer of electrons creates electrostatic attraction between oppositely charged ions, forming an ionic bond.

Ionic Bonds: Metals Lose Electrons

• Metals lose valence electrons to achieve a stable electron configuration (octet rule).

• Nonmetals gain electrons to form anions with a full valence shell.

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

Negative Ions: Nonmetals Gain Electrons

• Nonmetals in groups 5A, 6A, and 7A gain electrons to achieve a stable octet.

• Ion charge is calculated as: $\text{Ionic charge} = \text{number of protons} - \text{number of electrons}$

Example: Oxygen (O) gains two electrons to form O2-.

Ionic Charges and Group Numbers

The periodic table can be used to predict the charges of ions formed by main group elements:

Group	Typical Ion Charge
1A	+1
2A	+2
3A	+3
5A	-3
6A	-2
7A	-1

Ionic Compounds: Structure and Properties

Ionic compounds consist of cations and anions held together by strong electrostatic forces (ionic bonds). The formula unit represents the lowest whole-number ratio of ions.

• Properties:

  • High melting points

  • Solid at room temperature

  • Form crystalline lattices

Formulas of Ionic Compounds

• Symbols and subscripts are written in the lowest whole-number ratio.

• The sum of positive and negative charges must equal zero: $\text{Total positive charge} = \text{Total negative charge}$

Example: For NaCl, one Na+ balances one Cl-.

Naming and Writing Ionic Formulas

1. Name the metal (cation) first, then the nonmetal (anion) with an -ide ending.

2. For transition metals with variable charges, indicate the charge with Roman numerals in parentheses.

Example: FeCl2 is iron(II) chloride; FeCl3 is iron(III) chloride.

Polyatomic Ions

Polyatomic ions are covalently bonded groups of atoms with an overall charge.

Ion Name	Formula	Charge
Ammonium	NH4+	+1
Nitrate	NO3-	-1
Sulfate	SO42-	-2
Phosphate	PO43-	-3

When naming compounds with polyatomic ions, use the ion name as is.

Writing Formulas for Compounds with Polyatomic Ions

• Balance the total positive and negative charges.

• Use parentheses if more than one polyatomic ion is needed.

Example: Ca(NO3)2 is calcium nitrate.

Molecular Compounds: Sharing Electrons

Molecular (covalent) compounds are formed when two or more nonmetals share electrons to achieve stability.

• Atoms are held together by covalent bonds.

• Prefixes are used to indicate the number of each atom in the formula.

Example: CO2 is carbon dioxide; SO3 is sulfur trioxide.

Naming Molecular Compounds

1. Name the first nonmetal by its element name.

2. Name the second nonmetal with an -ide ending.

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

Example: N2O3 is dinitrogen trioxide.

Lewis Structures for Molecules and Polyatomic Ions

Lewis structures represent the arrangement of valence electrons in molecules and polyatomic ions.

• Shared pairs of electrons (bonding pairs) are shown as lines.

• Unshared pairs (lone pairs) are shown as dots.

Steps to Draw Lewis Structures:

1. Count total valence electrons.

2. Identify the central atom (usually the least electronegative, not H).

3. Connect outer atoms to central atom with single bonds.

4. Distribute remaining electrons to complete octets.

5. Form double or triple bonds if necessary to satisfy the octet rule.

Example: CO2 has two double bonds between C and O.

Exceptions to the Octet Rule

• Hydrogen (H) requires only 2 electrons.

• Boron (B) is stable with 6 electrons.

• Elements in period 3 or higher can have expanded octets.

Electronegativity and Bond Polarity

Electronegativity (EN) is the ability of an atom to attract shared electrons in a bond. The difference in EN between two atoms determines bond type:

• $\Delta EN = |EN_1 - EN_2|$

• If $\Delta EN < 0.5$, bond is nonpolar covalent.

• If $0.5 \leq \Delta EN \leq 1.8$, bond is polar covalent.

• If $\Delta EN > 1.8$, bond is ionic.

Example:

• O–O ($\Delta EN = 0$): Nonpolar covalent

• C–O ($\Delta EN = 1.0$): Polar covalent

• Na–Cl ($\Delta EN = 2.1$): Ionic

Dipoles and Bond Polarity

A polar covalent bond has a dipole, with partial positive and negative charges at opposite ends. The direction of the dipole is indicated by an arrow pointing from positive to negative.

Additional info: These notes cover the essential concepts of ionic and molecular compounds, including structure, naming, and properties, as well as the basics of Lewis structures and bond polarity, as relevant to a GOB Chemistry course.