Ionic Bonding

Overview

Ionic bonding occurs when electrons are transferred from one atom to another, creating oppositely charged ions that attract each other. This typically happens between metals (which lose electrons) and nonmetals (which gain electrons).

Formation of Ionic Bonds

Process

1. Metal atom loses electron(s) → forms cation (+)
2. Nonmetal atom gains electron(s) → forms anion (-)
3. Electrostatic attraction holds ions together

Example: Sodium Chloride (NaCl)

$$\text{Na} \rightarrow \text{Na}^+ + e^- \quad \text{(loses 1 electron)}$$ $$\text{Cl} + e^- \rightarrow \text{Cl}^- \quad \text{(gains 1 electron)}$$ $$\text{Na}^+ + \text{Cl}^- \rightarrow \text{NaCl} \quad \text{(ionic compound)}$$

Lattice Energy

The energy released when gaseous ions combine to form a solid ionic compound.

$$\text{M}^+(g) + \text{X}^-(g) \rightarrow \text{MX}(s) \quad \Delta H = -\text{Lattice Energy}$$

Coulomb's Law

$$E = k\frac{q_1 \times q_2}{r}$$

Where:

• $E$ = electrostatic energy
• $k$ = Coulomb's constant
• $q_1, q_2$ = ion charges
• $r$ = distance between ions

Factors Affecting Lattice Energy

Factor	Effect on Lattice Energy
Higher charges	Increases
Smaller ions	Increases

Comparison

$$\text{NaCl} < \text{MgO} < \text{Al}_2\text{O}_3$$ $$(+1,-1) < (+2,-2) < (+3,-2)$$

Born-Haber Cycle

A thermochemical cycle used to calculate lattice energy indirectly.

Steps for NaCl

1. Sublimation of Na(s): Na(s) → Na(g)
2. Ionization of Na: Na(g) → Na⁺(g) + e⁻
3. Dissociation of Cl₂: ½Cl₂(g) → Cl(g)
4. Electron affinity of Cl: Cl(g) + e⁻ → Cl⁻(g)
5. Lattice formation: Na⁺(g) + Cl⁻(g) → NaCl(s)

$$\Delta H_f° = \Delta H_{sub} + IE + \frac{1}{2}D + EA + (-U)$$

Common Ions

Cations (Positive)

Ion	Name	Electron Config
Na⁺	Sodium	[Ne]
Mg²⁺	Magnesium	[Ne]
Al³⁺	Aluminum	[Ne]
Fe²⁺	Iron(II)	[Ar] 3d⁶
Fe³⁺	Iron(III)	[Ar] 3d⁵

Anions (Negative)

Ion	Name	Electron Config
Cl⁻	Chloride	[Ar]
O²⁻	Oxide	[Ne]
S²⁻	Sulfide	[Ar]
N³⁻	Nitride	[Ne]

Properties of Ionic Compounds

Property	Explanation
High melting/boiling points	Strong electrostatic forces
Hard and brittle	Rigid crystal lattice
Conduct electricity when molten/dissolved	Free ions can move
Do not conduct as solids	Ions fixed in lattice
Soluble in polar solvents	Ion-dipole interactions

Ionic Formulas

Ionic compounds must be electrically neutral.

Writing Formulas

1. Write the cation first, then the anion
2. Balance charges so total = 0
3. Use subscripts to show ratios

Examples

$$\text{Ca}^{2+} + \text{Cl}^- \rightarrow \text{CaCl}_2 \quad \text{(2 Cl⁻ balance 1 Ca²⁺)}$$ $$\text{Al}^{3+} + \text{O}^{2-} \rightarrow \text{Al}_2\text{O}_3 \quad \text{(cross-multiply charges)}$$ $$\text{Fe}^{3+} + \text{SO}_4^{2-} \rightarrow \text{Fe}_2(\text{SO}_4)_3$$

Naming Ionic Compounds

Binary Compounds

Metal name + Nonmetal root + "-ide"
NaCl = Sodium chloride
MgO = Magnesium oxide

Transition Metals (multiple oxidation states)

FeCl₂ = Iron(II) chloride
FeCl₃ = Iron(III) chloride

Polyatomic Ions

Na₂SO₄ = Sodium sulfate
Ca(NO₃)₂ = Calcium nitrate

Predicting Ionic Compound Formation

Ionic bonds typically form when:

• Electronegativity difference > 1.7
• Metal reacts with nonmetal
• Elements are far apart on the periodic table