Chapter 3: Ionic Compounds – Structure, Formation, and Properties

Chapter 3: Ionic Compounds

3.1 Ions

Ions are charged particles formed when atoms gain or lose electrons. They are essential in biological and chemical processes, such as nerve signal transmission.

• Cation: A positively charged ion formed by the loss of electrons.

• Anion: A negatively charged ion formed by the gain of electrons.

• Example: Sodium (Na) loses one electron to form Na+; chlorine (Cl) gains one electron to form Cl-.

• Properties of Ionic Compounds: High melting points, stable crystalline solids, water solubility, and electrical conductivity in solution.

3.2 Ions and the Octet Rule

The octet rule states that main group elements tend to react to achieve eight valence electrons, resembling noble gas configurations.

• Metals: Lose electrons to form cations with noble gas configurations before them.

• Nonmetals: Gain electrons to form anions with noble gas configurations after them.

• Example: Magnesium (Mg, Z=12) loses two electrons to form Mg2+ with a neon configuration.

• Electron Configuration: $1s^2\ 2s^2\ 2p^6\ 3s^2$ (Mg); after losing two electrons: $1s^2\ 2s^2\ 2p^6$ (Mg2+)

3.3 Ions of Some Common Elements

The periodic table helps predict ion charges for main group elements.

• Group 1A: +1 cations (e.g., Na+)

• Group 2A: +2 cations (e.g., Ca2+)

• Group 6A: -2 anions (e.g., O2-, S2-)

• Group 7A: -1 anions (e.g., Cl-)

• Transition Metals: Form cations with variable charges; octet rule not always followed.

• Charge Prediction: For metals, cation charge = group number; for nonmetals, anion charge = 8 – group number.

3.4 Periodic Properties and Ion Formation

Periodic trends influence ion formation.

• Ionization Energy: Energy required to remove an electron; lower for metals, higher for nonmetals.

• Electron Affinity: Energy released when an atom gains an electron; highest for halogens.

• Trend: Metals (left) lose electrons easily; nonmetals (right) gain electrons easily; noble gases are unreactive.

3.5 Naming Monoatomic Ions

Monoatomic ions are named based on their element and charge.

• Main Group Cations: Element name + "ion" (e.g., potassium ion).

• Transition Metal Cations: Old system: -ous (lower charge), -ic (higher charge); New system: Roman numeral (e.g., iron(II) ion).

• Anions: Element name ending replaced with -ide + "ion" (e.g., chloride ion).

Transition Metal	Old Name	New Name
Fe2+	Ferrous ion	Iron(II) ion
Fe3+	Ferric ion	Iron(III) ion
Cu+	Cuprous ion	Copper(I) ion
Cu2+	Cupric ion	Copper(II) ion
Hg22+	Mercurous ion	Mercury(I) ion
Hg2+	Mercuric ion	Mercury(II) ion

Anion	Name
Cl-	Chloride ion
O2-	Oxide ion
S2-	Sulfide ion
N3-	Nitride ion

3.6 Polyatomic Ions

Polyatomic ions are ions composed of multiple atoms covalently bonded, carrying a net charge.

• Examples: Nitrate (NO3-), sulfate (SO42-), ammonium (NH4+).

• These ions are common and should be memorized.

Polyatomic Ion	Formula	Charge
Nitrate	NO3	-1
Sulfate	SO4	-2
Phosphate	PO4	-3
Ammonium	NH4	+1
Hydroxide	OH	-1
Carbonate	CO3	-2
Acetate	CH3COO	-1

3.7 Ionic Bonds

Ionic bonds are electrostatic attractions between oppositely charged ions, forming ionic compounds.

• Formed by transfer of electrons from metals to nonmetals.

• Result in crystalline solids with repeating patterns.

• Example: NaCl is formed from Na+ and Cl- ions.

3.8 Formulas of Ionic Compounds

Ionic compound formulas reflect the ratio of cations to anions needed for electrical neutrality.

• Rule: Total positive charge equals total negative charge.

• Formula unit: Smallest neutral unit (e.g., NaCl, Ca(NO3)2).

• Writing Formulas:

  • Cation first, anion second.

  • No charges shown in the formula.

  • Parentheses used for polyatomic ions with subscripts.

• Example: Calcium nitrate: Ca2+ and NO3- combine as Ca(NO3)2.

3.9 Naming Ionic Compounds

Ionic compounds are named by stating the cation first, then the anion.

• Type I: Main group cations; charge does not vary (e.g., sodium chloride).

• Type II: Transition metals; charge specified with Roman numerals (e.g., iron(III) chloride).

• Do not: Use prefixes like di- or tri- for ionic compounds.

• Example: FeCl2 is iron(II) chloride; FeCl3 is iron(III) chloride.

Compound	Name	Application
NaCl	Sodium chloride	Table salt
CaCl2	Calcium chloride	De-icing agent
FeCl3	Iron(III) chloride	Water treatment
AgNO3	Silver nitrate	Antiseptic
NaOH	Sodium hydroxide	Drain cleaner

3.10 Properties of Ionic Compounds

Ionic compounds exhibit characteristic physical properties due to strong ionic bonds.

• High melting and boiling points: e.g., NaCl melts at 801°C.

• Electrical conductivity: Conduct electricity when dissolved in water.

• Brittleness: Ionic solids shatter when struck.

• Solubility: Many, but not all, ionic compounds dissolve in water.

3.11 H+ and OH- Ions: Introduction to Acids and Bases

Hydrogen cations and hydroxide anions are central to acid-base chemistry.

• Acid: Provides H+ ions in water; H+ is a proton, often forms hydronium ion (H3O+).

• Base: Provides OH- ions in water; hydroxide is a polyatomic ion.

• Examples:

  • Hydrochloric acid (HCl): provides one H+ per molecule.

  • Sulfuric acid (H2SO4): provides two H+ per molecule.

  • Phosphoric acid (H3PO4): provides three H+ per molecule.

  • Sodium hydroxide (NaOH): provides one OH- per formula unit.

  • Barium hydroxide (Ba(OH)2): provides two OH- per formula unit.

Acid	Formula	Anion Derived
Hydrochloric acid	HCl	Cl- (chloride)
Sulfuric acid	H2SO4	SO42- (sulfate)
Phosphoric acid	H3PO4	PO43- (phosphate)
Acetic acid	CH3COOH	CH3COO- (acetate)

Additional info: Polyatomic ions and acid-base concepts are foundational for later chapters on solutions, equilibrium, and biological chemistry.