Ch. 6: Ionic and Molecular Compounds

Introduction

This chapter covers the fundamental principles of ionic and molecular (covalent) compounds, including their formation, properties, and systematic naming conventions. Understanding these concepts is essential for grasping the basics of chemical bonding and compound classification in GOB Chemistry.

Octet (8-) Rule

Definition and Application

• Octet Rule: When elements form compounds, atoms will have the same number of electrons as the closest noble gas (usually 8 for main group elements).

• Atoms can achieve an octet by gaining, losing, or sharing electrons.

• Only valence electrons are involved in bonding.

Example: Sodium (Na) loses one electron to achieve the same electron configuration as neon (Ne).

Types of Chemical Bonds

Ionic Bonds vs Covalent Bonds

• Ionic Bonds: Formed through the transfer of electrons from a metal to a nonmetal, resulting in positive (cation) and negative (anion) ions that attract each other.

• Covalent Bonds: Formed through the sharing of electrons between two nonmetals.

Comparison Table:

Formation of Ionic Compounds

Process and Properties

• Ionic compounds are formed when metals lose electrons to become cations and nonmetals gain electrons to become anions.

• The resulting compound is electrically neutral (overall charge is zero).

• Example: Sodium (Na) and chlorine (Cl) form sodium chloride (NaCl).

Equation:

Diatomic Elements

Definition and List

• Seven elements exist in nature as diatomic molecules (two atoms bonded together): H2, N2, O2, F2, Cl2, Br2, I2.

• These should be memorized for naming and formula writing.

Naming Ionic Compounds

Rules and Examples

• First word: Cation (positively charged ion) name.

• Second word: Anion (negatively charged ion) name (ends in "-ide" for monatomic ions).

• For compounds with polyatomic ions, use the ion's name (e.g., sulfate, nitrate).

• Parentheses are used when more than one polyatomic ion is present.

Example Table:

Writing Chemical Formulas for Ionic Compounds

Steps and Examples

• Balance the charges so the total positive and negative charges are equal.

• Use subscripts to indicate the number of each ion needed.

• For polyatomic ions, use parentheses if more than one is present.

Examples:

• Al3+ and Cl-:

• Sn4+ and SO42-:

• Mg2+ and P3-:

Monatomic and Polyatomic Ions

Definitions and Examples

• Monatomic ion: An ion containing only one atom (e.g., Na+, Cl-).

• Polyatomic ion: An ion containing more than one atom (e.g., OH-, CO32-, SO42-).

Naming Covalent Compounds

Rules and Prefixes

• Composed of only nonmetal atoms.

• First word: Name of the first element.

• Second word: Name of the second element ending in "-ide".

• Prefixes indicate the number of each atom (mono-, di-, tri-, tetra-, penta-, etc.).

• Omit "mono-" for the first element.

Example Table:

Multiple Bonds in Covalent Compounds

Single, Double, and Triple Bonds

• Atoms can share more than one pair of electrons, forming double or triple bonds.

• Double bond: Two pairs of shared electrons (e.g., O2).

• Triple bond: Three pairs of shared electrons (e.g., N2).

Electronegativity (EN)

Definition and Trends

• Electronegativity: The ability of an atom to attract electrons in a chemical bond.

• Increases across a period (left to right) and decreases down a group (top to bottom) in the periodic table.

Example: Fluorine (F) has the highest electronegativity.

Polar vs Nonpolar Covalent Bonds

Bond Polarity and Electronegativity Difference

• Polar covalent bond: Electrons are shared unequally due to a significant difference in electronegativity between atoms.

• Nonpolar covalent bond: Electrons are shared equally; electronegativity difference is small or zero.

Equation:

If is large (> 0.4), the bond is polar; if small (< 0.4), the bond is nonpolar.

Summary Table: Types of Compounds

Additional info: Some context and examples have been expanded for clarity and completeness, including systematic naming rules and electronegativity trends.