Ionic and Molecular Compounds: Structure, Naming, and Properties

Ionic and Molecular Compounds

Introduction

This study guide covers the foundational concepts of ionic and molecular compounds, including their structure, naming conventions, and properties. These topics are essential for understanding chemical bonding and the behavior of substances in introductory college chemistry.

Atoms, Ions, and Molecules

What is a Molecule?

• Molecule: A group of two or more atoms held together by chemical bonds.

• Examples:

  • Oxygen (O2)

  • Nitrogen (N2)

  • Hydrogen (H2)

  • Methane (CH4)

  • Ammonia (NH3)

  • Glucose (C6H12O6)

• Molecular Shape: Influences physical properties such as boiling point, solubility, and reactivity.

Atoms vs. Ions

• Atom: Electrically neutral; same number of protons and electrons.

• Ion: Electrically charged; different number of protons and electrons.

  • Cation: Positively charged ion (fewer electrons than protons).

  • Anion: Negatively charged ion (more electrons than protons).

Chemical Bonding

Octet Rule

• Atoms tend to gain, lose, or share electrons to achieve a stable configuration of eight valence electrons (octet).

• Exceptions: Hydrogen (2 electrons), Boron (6 electrons), elements in period 3 and beyond can have expanded octets.

Ionic Bonds

• Formed by the transfer of electrons from a metal to a nonmetal.

• Example: Sodium chloride ($NaCl$) forms when sodium donates an electron to chlorine.

• Properties: High melting points, conduct electricity when dissolved in water.

Covalent Bonds

• Formed by the sharing of electrons between nonmetals.

• Example: Water ($H_2O$) forms when hydrogen and oxygen share electrons.

• Properties: Lower melting points, do not conduct electricity in solution.

Naming and Writing Formulas

Naming Ionic Compounds

• Name the cation (metal) first, then the anion (nonmetal).

• Anion name ends with “-ide.”

• For transition metals, use Roman numerals to indicate charge (e.g., Iron(III) chloride: $FeCl_3$).

Naming Molecular Compounds

• Name the first nonmetal by its element name.

• Name the second nonmetal using the root and “-ide.”

• Use prefixes to indicate the number of atoms: mono-, di-, tri-, tetra-, penta-, etc.

• Example: $CO_2$ is carbon dioxide; $N_2O_3$ is dinitrogen trioxide.

Polyatomic Ions

• Groups of covalently bonded atoms with an overall charge.

• Examples:

  • Nitrate ($NO_3^-$)

  • Sulfate ($SO_4^{2-}$)

  • Ammonium ($NH_4^+$)

Common Polyatomic Ions Table

Name	Formula	Charge
Nitrate	$NO_3^-$	-1
Sulfate	$SO_4^{2-}$	-2
Phosphate	$PO_4^{3-}$	-3
Ammonium	$NH_4^+$	+1

Lewis Dot Structures

Drawing Lewis Structures

• Show valence electrons as dots around element symbols.

• Shared pairs (bonds) are shown as lines or pairs of dots.

• Lone pairs are placed on the outside of atoms.

• Central atom is usually the least electronegative.

• Follow the octet rule for most atoms.

Single, Double, and Triple Bonds

• Single bond: One pair of shared electrons ($-$).

• Double bond: Two pairs of shared electrons ($=$).

• Triple bond: Three pairs of shared electrons ($riplebond$).

Molecular Geometry (VSEPR Theory)

Predicting Molecular Shape

• Electron groups around a central atom arrange themselves as far apart as possible.

• Common shapes:

  • Linear: 2 electron groups

  • Trigonal planar: 3 electron groups

  • Tetrahedral: 4 electron groups

  • Bent: 2 bonding pairs + 2 lone pairs

Polarity and Electronegativity

Electronegativity

• Ability of an atom to attract shared electrons in a chemical bond.

• Increases across a period, decreases down a group.

• Fluorine is the most electronegative element.

Bond Polarity

• Nonpolar covalent bond: Electrons shared equally.

• Polar covalent bond: Electrons shared unequally; creates dipole.

• Ionic bond: Electrons transferred.

• Bond polarity increases with greater difference in electronegativity.

Molecular Polarity

• Depends on both bond polarity and molecular shape.

• Symmetrical molecules (e.g., $CO_2$) can be nonpolar even with polar bonds.

• Asymmetrical molecules (e.g., $H_2O$) are polar.

Intermolecular Forces

Types of Intermolecular Forces

• Dispersion forces: Weak attractions between nonpolar molecules.

• Dipole-dipole interactions: Attractions between polar molecules.

• Hydrogen bonds: Strong dipole attractions involving H bonded to F, O, or N.

• Ionic bonds: Strongest, between ions in ionic compounds.

Relative Strengths Table

Type	Relative Strength
Dispersion	Weak
Dipole-dipole	Moderate
Hydrogen bonding	Strong
Ionic	Very strong

Medically Important Ions

• Sodium ions ($Na^+$): Regulate osmotic pressure and water content, transmit nerve signals, contract muscles.

• Potassium ions ($K^+$): Transmit nerve signals, contract muscles.

• Chloride ions ($Cl^-$): Regulate osmotic pressure and water content, transmit nerve signals, secrete stomach acid.

• Calcium ions ($Ca^{2+}$): Form bones and teeth, contract muscles, clot blood.

Summary Table: Key Skills in Ionic and Molecular Compounds

Skill	Description
Write	Symbols for simple ions
Using	Charge balance to write correct formula
Write	Name and formula for ionic compounds and polyatomic ions
Draw	Lewis structures for molecules and polyatomic ions
Use	Electronegativity to determine bond polarity
Predict	Three-dimensional structure and polarity
Describe	Intermolecular forces

Applications and Lab Context

• Understanding ionic and molecular compounds is essential for predicting chemical behavior, biological functions, and designing experiments such as building a solar battery.

Additional info: Some context and examples were inferred to clarify fragmented notes and ensure completeness.