Ionic and Molecular Compounds: Structure, Bonding, and Properties

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Chapter 06: Ionic and Molecular Compounds

Overview

This chapter introduces the fundamental differences between ionic and molecular compounds, focusing on how atoms combine through electron transfer or sharing. It covers the formation, naming, and properties of these compounds, as well as the forces that hold them together.

Transfer and Sharing of Electrons

Formation of Compounds

Most elements, except the noble gases, are found in nature as compounds. Compounds form when electrons are transferred or shared between atoms, allowing them to achieve stable electron configurations.

Ionic Bond: Formed when electrons are transferred from a metal to a nonmetal, resulting in positive and negative ions (e.g., NaCl).
Covalent Bond: Formed when atoms of nonmetals share valence electrons (e.g., H2O).
Octet Rule: Atoms tend to gain, lose, or share electrons to achieve eight valence electrons, similar to noble gases.
Exception: Hydrogen achieves stability with two valence electrons.

Types of Particles and Bonds in Compounds

Ionic vs. Molecular Compounds

Ionic compounds consist of ions held together by ionic bonds, while molecular compounds consist of molecules held together by covalent bonds.

Type	Particles	Bond Type	Example
Ionic Compounds	Ions (Na+, Cl-)	Ionic	NaCl
Molecular Compounds	Molecules (H2O, C3H8)	Covalent	H2O, C3H8

Ions: Formation and Stability

Formation of Ions

Cations: Formed when atoms (usually metals) lose electrons, becoming positively charged.
Anions: Formed when atoms (usually nonmetals) gain electrons, becoming negatively charged.
Electron Configuration: Ions achieve the electron configuration of the nearest noble gas.

Writing Ionic Formulas

Charge Balance and Subscripts

The chemical formula of an ionic compound represents the lowest whole-number ratio of ions, ensuring the total positive and negative charges are equal.

Formula Example: For NaCl,
Subscripts: Indicate the number of each ion needed for charge balance.

Naming Ionic Compounds

Rules for Naming

Cation: Named first, using the element name.
Anion: Named by taking the root of the element name and adding "-ide" (e.g., chloride).
Transition Metals: May have variable charges; use Roman numerals to indicate the charge (e.g., iron(III) chloride).

Polyatomic Ions

Definition and Examples

A polyatomic ion is a group of covalently bonded atoms with an overall charge. Most are anions, except for ammonium ().

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

When more than one polyatomic ion is needed, parentheses are used (e.g., Ca(NO3)2).

Molecular Compounds: Sharing Electrons

Covalent Bonding

Atoms of nonmetals share electrons to form molecules.
Molecules have definite proportions and are held together by covalent bonds.

Naming Molecular Compounds

Rules for Naming

The first nonmetal is named by its element name.
The second nonmetal is named using the root and "-ide" ending.
Prefixes (mono-, di-, tri-, etc.) indicate the number of atoms; "mono-" is usually omitted for the first element.
Example: CO2 is carbon dioxide; N2O is dinitrogen oxide.

Lewis Structures for Molecules and Polyatomic Ions

Drawing Lewis Structures

Shows how atoms share electrons to achieve octets (except H, which needs two electrons).
Lone pairs and bonding pairs are indicated.
Multiple bonds (double, triple) may be needed to complete octets.

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

Electronegativity and Bond Polarity

Electronegativity

Electronegativity: The ability of an atom to attract electrons in a bond.
Increases across a period and decreases down a group.

Bond Polarity

Nonpolar Covalent Bond: Electrons shared equally (electronegativity difference < 0.4).
Polar Covalent Bond: Electrons shared unequally (difference 0.5–1.8).
Ionic Bond: Electrons transferred (difference > 1.8).

Bond Type	Electronegativity Difference
Nonpolar Covalent	0.0–0.4
Polar Covalent	0.5–1.8
Ionic	>1.8

Shape and Polarity of Molecules

VSEPR Theory

Valence Shell Electron-Pair Repulsion (VSEPR) Theory: Electron groups around a central atom are arranged as far apart as possible to minimize repulsion.
Common shapes: linear, trigonal planar, tetrahedral, trigonal pyramidal, bent.

Polarity of Molecules

Nonpolar molecules: All dipoles cancel due to symmetrical arrangement (e.g., CO2, CCl4).
Polar molecules: Dipoles do not cancel, resulting in a molecule with a positive and negative end (e.g., H2O).

Molecular Forces in Compounds

Types of Intermolecular Forces

Dipole-Dipole Attractions: Occur between polar molecules.
Hydrogen Bonds: Strong dipole-dipole attractions when H is bonded to N, O, or F.
Dispersion Forces: Weak attractions present in all molecules, especially nonpolar ones.

Force Type	Relative Strength	Example
Ionic Bonds	Strongest	NaCl
Hydrogen Bonds	Strong	H2O
Dipole-Dipole	Moderate	CHCl3
Dispersion Forces	Weakest	CO2

Stronger intermolecular forces result in higher melting and boiling points.

Summary Table: Ionic vs. Molecular Compounds

Property	Ionic Compounds	Molecular Compounds
Particles	Ions	Molecules
Bond Type	Ionic	Covalent
Melting Point	High	Low to moderate
Electrical Conductivity	Conducts when dissolved	Does not conduct

Key Equations

Additional info: Some content and examples have been expanded for clarity and completeness based on standard General Chemistry curriculum.