Chapter 6 – Ionic and Molecular Compounds: Study Notes

Chapter 6 – Ionic and Molecular Compounds

Introduction

This chapter explores the formation, properties, and nomenclature of ionic and molecular compounds, focusing on the nature of chemical bonds, the behavior of ions, and the rules for writing and naming compounds. Understanding these concepts is fundamental for predicting compound properties and chemical reactivity.

Review: Valence Electrons and the Periodic Table

Valence Electrons for Representative Elements

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

• The number of valence electrons determines an element's chemical properties and its ability to form bonds.

• Examples:

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

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

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

  • Br (Bromine): 7 valence electrons ($4s^2 4p^5$)

Section 6.1: Ionic and Covalent Bonds

Chemical Bonds and the Octet Rule

• Chemical bonds are formed when atoms lose, gain, or share valence electrons to acquire an octet (eight valence electrons), known as the octet rule.

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

Ionic Bonds

• Formed when valence electrons of a metal atom are transferred to a nonmetal atom.

• Result in the formation of ionic compounds.

Covalent Bonds

• Formed when valence electrons are shared between nonmetal atoms to achieve a noble gas configuration.

• Result in the formation of molecular compounds.

Transfer of Electrons and Ion Formation

• Atoms form cations (positively charged ions) by losing electrons.

• Atoms form anions (negatively charged ions) by gaining electrons.

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

• Example: $\text{Na} \rightarrow \text{Na}^+ + e^-$; $\text{Cl} + e^- \rightarrow \text{Cl}^-$

Formation of Ions: Metals and Nonmetals

Positive Ions: Metals Lose Electrons

• Metals (Groups 1A, 2A, 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 eight).

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

• They gain electrons until they have a noble gas configuration.

• 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 Electrons)	Nonmetals (Gain Electrons)
1A	+1 (e.g., Na+)
2A	+2 (e.g., Mg2+)
3A	+3 (e.g., Al3+)
5A		-3 (e.g., N3-)
6A		-2 (e.g., O2-)
7A		-1 (e.g., Cl-)

Practice: Determining Subatomic Particles in Ions

• Given an ion's symbol, you can determine the number of protons, neutrons, and electrons.

• Example: $^{55}\text{Co}^{3+}$

  • Protons: 27 (atomic number)

  • Neutrons: 28 ($55-27$)

  • Electrons: 24 ($27-3$)

Section 6.2: Ionic Compounds

Properties of Ionic Compounds

• Consist of positive and negative ions held together by strong electrical attractions (ionic bonds).

• High melting points.

• Solids at room temperature.

Formulas of Ionic Compounds

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

• The sum of ion charges equals zero: $\text{total positive charge} = \text{total negative charge}$

• The lowest whole-number combination is called a formula unit.

• Examples:

  • $\text{Na}^+ + \text{Cl}^- \rightarrow \text{NaCl}$

  • $\text{Mg}^{2+} + 2\text{Cl}^- \rightarrow \text{MgCl}_2$

  • $2\text{Na}^+ + \text{S}^{2-} \rightarrow \text{Na}_2\text{S}$

Practice: Writing Chemical Formulas

• Combine ions so that the total positive and negative charges balance to zero.

• Examples:

  • $\text{Al}^{3+}$ and $\text{Cl}^-$: $\text{AlCl}_3$

  • $\text{Mg}^{2+}$ and $\text{N}^{3-}$: $\text{Mg}_3\text{N}_2$

  • $\text{Ca}^{2+}$ and $\text{Br}^-$: $\text{CaBr}_2$

Naming Ionic Compounds

Metals with Fixed Charge

• The metal is written first, using its element name.

• The nonmetal is written second, using the first syllable of its name plus the -ide ending.

• Examples:

  • KI: potassium iodide

  • MgBr2: magnesium bromide

  • Al2O3: aluminum oxide

Guide to Naming Ionic Compounds

1. Identify the cation and anion.

2. Name the cation by its element name.

3. Name the anion by using the first syllable of its element name followed by ide.

4. Write the name of the cation first and the anion second.

Example: K2O is potassium oxide.

Metals with Variable Charge (Transition Metals)

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

• A Roman numeral equal to the ion charge is placed in parentheses after the metal name.

• Examples:

  • Fe2+: iron(II)

  • Fe3+: iron(III)

  • Cu+: copper(I)

  • Cu2+: copper(II)

Example: MnF2 is manganese(II) fluoride.

Practice: Naming and Writing Formulas

• SnO2: tin(IV) oxide

• PbS2: lead(IV) sulfide

• Fe2O3: iron(III) oxide

Section 6.4: Polyatomic Ions

Definition and Properties

• A polyatomic ion is a group of atoms with an overall charge that behaves as a single unit.

• Often consist of nonmetals (phosphorus, sulfur, carbon, nitrogen) plus oxygen.

• Usually have a negative charge (1-, 2-, or 3-), except for NH4+ (ammonium).

Common Polyatomic Ions

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

Naming Polyatomic Ions

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

• If the ion has 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., HCO3-: hydrogen carbonate or bicarbonate).

• Halogens (Group 7A) can form four polyatomic ions with oxygen: perchlorate (ClO4-), chlorate (ClO3-), chlorite (ClO2-), hypochlorite (ClO-).

Practice: Writing Formulas with Polyatomic Ions

• Combine the cation and polyatomic anion so that the total charge is zero.

• Example: Al3+ and HCO3- → Al(HCO3)3

Naming Ionic Compounds with Polyatomic Ions

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

• Examples:

  • Na2SO4: sodium sulfate

  • FePO4: iron(III) phosphate

  • Al2(CO3)3: aluminum carbonate

Additional info: These notes cover the foundational aspects of ionic and molecular compounds, including the formation and naming of ions, the writing of chemical formulas, and the identification and naming of polyatomic ions. Mastery of these concepts is essential for further study in chemistry, especially in predicting compound properties and chemical reactivity.