Ions and Ionic Bonds

The driving force. Atoms are most stable when they have a FULL OUTER SHELL of electrons (a noble-gas configuration). They lose, gain, or share electrons to reach that state.

Metals lose electrons to become positive cations.

• Sodium ($2, 8, 1$) loses 1 e⁻ → Na⁺ ($2, 8$, like neon).
• Magnesium ($2, 8, 2$) loses 2 e⁻ → Mg²⁺ ($2, 8$).
• Aluminium ($2, 8, 3$) loses 3 e⁻ → Al³⁺ ($2, 8$).

Non-metals gain electrons to become negative anions.

• Chlorine ($2, 8, 7$) gains 1 e⁻ → Cl⁻ ($2, 8, 8$, like argon).
• Oxygen ($2, 6$) gains 2 e⁻ → O²⁻ ($2, 8$, like neon).
• Nitrogen ($2, 5$) gains 3 e⁻ → N³⁻ ($2, 8$).

Group rules.

Group	Charge formed
1	$+1$
2	$+2$
13	$+3$
15	$-3$
16	$-2$
17	$-1$
18	already full — usually no ions

Worked. Sulfur ($2, 8, 6$) → S²⁻ (gain 2 to fill outer shell).

Ionic bond. The strong electrostatic ATTRACTION between oppositely charged ions.

Forming sodium chloride.

1. Na atom loses 1 e⁻ → Na⁺.
2. Cl atom gains that e⁻ → Cl⁻.
3. Na⁺ and Cl⁻ attract each other electrostatically.
4. Many Na⁺ and Cl⁻ ions arrange themselves in a 3D pattern — a giant ionic lattice — alternating + and − ions.

Drawing dot-and-cross diagrams. Show outer-shell electrons only.

• For NaCl, you'd draw Na surrendering its outer × to Cl, then Cl shown with 8 dots (full shell). The two ions are drawn next to each other in square brackets, with their charges.

Stoichiometry from charges. Total positive charge = total negative charge in the formula.

• Na⁺ and Cl⁻ → NaCl (one of each).
• Mg²⁺ and Cl⁻ → MgCl₂ (one Mg²⁺ balances two Cl⁻).
• Al³⁺ and O²⁻ → Al₂O₃ (lowest common multiple is 6: 2×3 = 3×2).

Worked. Calcium reacting with bromine.

• Ca: $2, 8, 8, 2$ → Ca²⁺.
• Br: gains 1 e⁻ → Br⁻.
• Need 2 Br⁻ to balance one Ca²⁺ → CaBr₂.

1. High melting and boiling points. Strong electrostatic forces in the lattice need lots of energy to break. NaCl m.p. = $801\degree\text{C}$.

2. Crystalline. Regular lattice → regular shape → crystals.

3. Often soluble in water. Polar water molecules pull ions out of the lattice.

4. Electrical conductivity.

• Solid: NO — ions can't move (locked in the lattice).
• Molten: YES — heating breaks the lattice; ions free to move and carry charge.
• Dissolved: YES — water-separated ions carry charge.

5. Brittle. Force in one direction shifts ions; like ions repel; lattice cracks along the cleavage plane.

Cambridge tip. When asked "why does sodium chloride conduct when molten but not when solid?" — the answer is about ION MOBILITY, not about the existence of ions. Ions exist in both states; they just can't move in the solid.

Worked. Why is the m.p. of MgO ($2852\degree\text{C}$) higher than NaCl ($801\degree\text{C}$)? Both are ionic. MgO has $2+$ and $2-$ ions — the electrostatic attraction is stronger (charges 2 vs 1) → harder to break the lattice → higher m.p.