Chapter 6 – Ionic and Molecular Compounds: Structure, Bonding, and Nomenclature

Chapter 6 – Ionic and Molecular Compounds

Introduction

This chapter explores the formation, structure, and properties of ionic and molecular compounds. It covers the nature of chemical bonds, how to predict and write chemical formulas, and the rules for naming compounds. Understanding these concepts is fundamental for studying chemical reactions and the behavior of substances.

Review from Chapter 4: Valence Electrons for Representative Elements

• Valence electrons are the s and p electrons in the highest energy level of an atom.

• Group "A" elements (main group) have predictable numbers of valence electrons:

  • Li (Lithium): 1 valence electron ($1s^2 2s^1$)

  • Mg (Magnesium): 2 valence electrons ($1s^2 2s^2 2p^6 3s^2$)

  • S (Sulfur): 6 valence electrons ($1s^2 2s^2 2p^6 3s^2 3p^4$)

  • Br (Bromine): 7 valence electrons ($1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^5$)

Section 6.1: Ionic and Covalent Bonds

Chemical Bonds and the Octet Rule

• Chemical bonds form when atoms lose, gain, or share valence electrons to achieve a stable octet (eight valence electrons), known as the octet rule.

• Eight valence electrons correspond to a noble gas configuration (e.g., Argon: $3s^2 3p^6$).

Ionic Bonds

• Formed when valence electrons of a metal atom are transferred to a nonmetal atom, resulting in an ionic compound.

Covalent Bonds

• Formed when valence electrons are shared between nonmetal atoms, resulting in a molecular compound.

Transfer of Electrons

• Atoms form cations (positively charged ions) by losing electrons and anions (negatively charged ions) by gaining electrons.

• Ionic bonds are the strong attractive forces between positive and negative ions.

• Example: $\text{Na}^+$ (loses 1 electron), $\text{Cl}^-$ (gains 1 electron)

Positive Ions: Metals Lose Electrons

• Metals in Groups 1A, 2A, and 3A have low ionization energies and readily lose electrons to form cations.

• They lose electrons until they have the same number of valence electrons as the nearest noble gas (usually 8).

• Example: $\text{Na}: 1s^2 2s^2 2p^6 3s^1$ loses one electron to become $\text{Na}^+$ ($1s^2 2s^2 2p^6$)

Negative Ions: Nonmetals Gain Electrons

• Nonmetals in Groups 5A, 6A, and 7A have high ionization energies and readily gain electrons to form anions.

• They gain electrons until they have the same number of valence electrons as the nearest noble gas.

• Example: $\text{Cl}: 1s^2 2s^2 2p^6 3s^2 3p^5$ gains one electron to become $\text{Cl}^-$ ($1s^2 2s^2 2p^6 3s^2 3p^6$)

Ionic Charges and Group Numbers

The group number in the periodic table helps determine the charge of ions for representative elements (Group A).

Group	Metals Lose Valence Electrons	Nonmetals Gain Valence Electrons	Noble Gases
1A	Li+		He
2A	Mg2+		Ne
3A	Al3+		Ar
5A		N3-
6A		O2-
7A		F-
8A			Ne, Ar, Kr

Practice: Determining Subatomic Particles in Ions

• Given $^{55}\text{Co}^{3+}$:

  • Protons: 27

  • Neutrons: 28

  • Electrons: 24

Practice: Determining Symbol and Charge of an Ion

• Example: An ion with 16 protons and 18 electrons is $\text{S}^{2-}$ (sulfide ion).

Section 6.2: Ionic Compounds

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

• Properties:

  • High melting points

  • Solids at room temperature

Formulas of Ionic Compounds

• Formulas are written in the lowest whole-number ratio of ions.

• The sum of ion charges equals zero:

$\text{Total positive charge} = \text{Total negative charge}$

• Example: $\text{NaCl}$: $1(+1) + 1(-1) = 0$

• Example: $\text{MgCl}_2$: $1(+2) + 2(-1) = 0$

• Example: $\text{Na}_2\text{S}$: $2(+1) + 1(-2) = 0$

Practice: Writing Chemical Formulas for Ionic Compounds

• Al3+ and Cl-: $\text{AlCl}_3$

• Mg2+ and N3-: $\text{Mg}_3\text{N}_2$

• Ca2+ and Br-: $\text{CaBr}_2$

Naming Ionic Compounds: Metals with Fixed Charge

• Name the metal first (same as the element name).

• Name the nonmetal using the first syllable + ide ending.

• Example: $\text{K}_2\text{O}$ is potassium oxide.

Compound	Metal Ion	Nonmetal Ion	Name
KI	K+	I-	Potassium iodide
MgBr2	Mg2+	Br-	Magnesium bromide
Al2O3	Al3+	O2-	Aluminum oxide

Naming Ionic Compounds: Metals with Variable Charge

• Transition metals (except Zn2+, Cd2+, Ag+) can form more than one cation.

• A Roman numeral indicates the ion charge in parentheses after the metal name.

• Examples:

  • Fe2+: iron(II)

  • Fe3+: iron(III)

  • Pb2+: lead(II)

  • Pb4+: lead(IV)

Determining Charge of a Metal with Variable Charge

• Use the charge on the anion and charge balance to calculate the metal's charge.

• Example: $\text{MnF}_2$

$1(x) + 2(-1) = 0 \implies x = +2$

Name: manganese(II) fluoride

Section 6.4: Polyatomic Ions

• Polyatomic ions are groups of atoms with an overall charge that behave as a single unit.

• Often consist of nonmetals (P, S, C, N) + oxygen (e.g., $\text{NO}_3^-$ nitrate).

• Usually have a negative charge (1-, 2-, or 3-), except $\text{NH}_4^+$ (ammonium).

Nonmetal	Formula of Ion	Name of Ion
Hydrogen	OH-	Hydroxide
Nitrogen	NH4+	Ammonium
Carbon	CO32-	Carbonate
Sulfur	SO42-	Sulfate
Phosphorus	PO43-	Phosphate

Naming Polyatomic Ions

• Most common polyatomic ions end in -ate (e.g., sulfate, phosphate, nitrate).

• If one less oxygen, the name ends in -ite (e.g., sulfite, nitrite).

• Exceptions: CN- (cyanide), OH- (hydroxide).

• Adding H+ to a polyatomic ion increases its charge by +1 and adds "hydrogen" or "bi-" to the name (e.g., $\text{HCO}_3^-$ hydrogen carbonate).

• Halogens (Group 7A) form four polyatomic ions with oxygen:

  • ClO4-: perchlorate

  • ClO3-: chlorate

  • ClO2-: chlorite

  • ClO-: hypochlorite

Writing Formulas with Polyatomic Ions

• Identify the cation and polyatomic ion (anion).

• Balance charges using the criss-cross method.

• Write the formula, cation first, using subscripts as needed.

• Example: Al3+ and HCO3- → $\text{Al(HCO}_3)_3$

Naming Ionic Compounds with Polyatomic Ions

• Name the positive ion (usually a metal) first, then the polyatomic ion.

• No prefixes are used.

• Examples:

  • Na2SO4: sodium sulfate

  • FePO4: iron(III) phosphate

  • Al2(CO3)3: aluminum carbonate

Section 6.5: Molecular Compounds – Sharing Electrons

• Molecular compounds (molecules) form when atoms of two or more nonmetals share electrons, creating covalent bonds.

• Valence electrons are shared to achieve stability (octet rule).

• Examples: HCl, H2O, NH3, CH4

Ball-and-Stick and Space-Filling Models

• Ball-and-stick models show atoms and bonds explicitly.

• Space-filling models show the overall molecular shape but not explicit bonds.

Summary Table: Common Polyatomic Ions

Ion	Formula	Name
Hydroxide	OH-	Hydroxide
Ammonium	NH4+	Ammonium
Nitrate	NO3-	Nitrate
Sulfate	SO42-	Sulfate
Phosphate	PO43-	Phosphate

Additional info: This summary covers the first half of Chapter 6, focusing on ionic and molecular compounds, their formation, nomenclature, and the role of polyatomic ions. For a complete understanding, students should also study Lewis structures, molecular geometry, bond polarity, and intermolecular forces, which are typically covered in the latter part of this chapter.