ionic model

Ionic Model – IB SL Chemistry Notes

1. Formation of Ions

Atoms form ions to achieve a stable noble gas electron configuration.

• Metals (Groups 1–3) lose electrons → form positive ions (cations)

• Non-metals (Groups 15–17) gain electrons → form negative ions (anions)

• Noble gases (Group 18) do not form ions (already stable)

Cations (Positive Ions)

• Formed by losing valence electrons

• Smaller than parent atom (loss of outer shell = reduced radius)

• Charge = Group Number

• Example: Na → Na⁺ + e⁻ (Group 1 → 1+)

• Example: Ca → Ca²⁺ + 2e⁻ (Group 2 → 2+)

• Example: Al → Al³⁺ + 3e⁻ (Group 13 → 3+)

Anions (Negative Ions)

• Formed by gaining electrons

• Larger than parent atom (extra electrons = increased repulsion)

• Charge = 18 - Group Number

• Example: F + e⁻ → F⁻ (Group 17 → 1−)

• Example: O + 2e⁻ → O²⁻ (Group 16 → 2−)

• Example: N + 3e⁻ → N³⁻ (Group 15 → 3−)

Isoelectronic Species

• Ions that have the same number of electrons as a noble gas

• Example: O²⁻, F⁻, Ne, Na⁺, Mg²⁺ (all have 10 electrons)

2. Ionic Bonding

Definition

Ionic bonding is the electrostatic attraction between oppositely charged ions.

Formation of Ionic Bonds

• Occurs between a metal and a non-metal

• Electron transfer happens from metal → non-metal

• Example: NaCl

• Na (Group 1) loses 1 e⁻ → Na⁺

• Cl (Group 17) gains 1 e⁻ → Cl⁻

• Na⁺ and Cl⁻ attract each other → NaCl

Key Features

✔ Strong bond due to oppositely charged ions

✔ Forms a lattice structure (not molecules)

✔ Overall charge is neutral (total +ve and -ve charges must balance)

3. Structure of Ionic Compounds

Giant Ionic Lattice

• Ions are arranged in a regular 3D repeating pattern

• Each positive ion is surrounded by negative ions and vice versa

• Examples:

• NaCl: Each Na⁺ is coordinated to 6 Cl⁻

• MgO: Stronger attraction (Mg²⁺ and O²⁻ have greater charge than Na⁺ and Cl⁻)

Properties of Ionic Compounds

✔ High melting & boiling points

• Strong electrostatic forces require a lot of energy to break

✔ Hard but brittle

• When force is applied, like charges align → repulsion → lattice shatters

✔ Soluble in water

• Water molecules surround ions and separate them → hydration

✔ Do not conduct electricity as solids

• Ions are fixed in place → no movement = no conductivity

• Conducts when molten or dissolved → free-moving ions carry charge

✔ Strength of Ionic Bonding Depends on:

1. Charge of ions → Higher charge = stronger bond

• Example: MgO (Mg²⁺ and O²⁻) has stronger bonds than NaCl (Na⁺ and Cl⁻)

2. Size of ions → Smaller ions = stronger bond (closer attraction)

• Example: LiF (Li⁺ smaller than Na⁺) has stronger attraction than NaCl

4. Lattice Enthalpy & Bond Strength

Lattice Enthalpy (ΔHₗₑ)

Definition: The energy released when 1 mole of an ionic compound forms from its gaseous ions.

✔ More negative lattice enthalpy = stronger ionic bonding

✔ Depends on:

• Charge of ions → Higher charge = stronger attraction

• Size of ions → Smaller ions = closer distance = stronger bond

✔ Trends:

• NaCl (moderate ΔHₗₑ) vs MgO (very negative ΔHₗₑ → stronger)

• CsCl (weaker, large Cs⁺) vs LiCl (stronger, small Li⁺)

5. Writing Ionic Formulas

• Ionic compounds must be neutral (charges balance)

• Steps to deduce formula:

1. Identify ions and charges

2. Balance charges to ensure total = 0

3. Write formula without charges

Examples

✔ Na⁺ + Cl⁻ → NaCl

✔ Ca²⁺ + O²⁻ → CaO

✔ Al³⁺ + O²⁻ → Al₂O₃ (Criss-cross method: Al³⁺, O²⁻ → swap charges → Al₂O₃)

Common Polyatomic Ions (Must Memorize!)

✔ Example with polyatomic ions:

• Mg²⁺ + NO₃⁻ → Mg(NO₃)₂

• Al³⁺ + SO₄²⁻ → Al₂(SO₄)₃

Summary Table