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VCE Chemistry (Year 11) · Unit 1 — Bonding and Materials · 11m read

Ionic Bonding and Compounds

Ever wondered why salt doesn't behave like sodium metal at all, even though it contains sodium? That's because of ionic bonding — when atoms transfer electrons to each other, creating charged particles called ions that attract and stick together. You'll learn why this electron transfer happens, how it forms the crystals and compounds we use every day, and why ionic compounds have such useful properties like melting at high temperatures and dissolving in water. Understanding this is crucial for VCE chemistry and helps explain everything from the salt in your food to the minerals in rocks.

Learning objective: Describe ionic bonding and the properties of ionic compounds.

By the end of this lesson you can:

  • I predict ions formed by elements.
  • I write ionic formulas.
  • I explain ionic compound properties.

Ionic bonding is a give-and-take relationship. Metals give electrons; non-metals take them. The opposite charges attract strongly, forming a lattice.

Electron transfer and lattices

IONIC BONDING:
• Between METAL and NON-METAL.
• Metal LOSES electrons (becomes cation, +).
• Non-metal GAINS electrons (becomes anion, −).
• Attraction between opposite charges = ionic bond.

COMMON IONS:
• Group 1 → +1 (Na⁺, K⁺).
• Group 2 → +2 (Mg²⁺, Ca²⁺).
• Group 13 → +3 (Al³⁺).
• Group 15 → −3 (N³⁻).
• Group 16 → −2 (O²⁻).
• Group 17 → −1 (Cl⁻).

TRANSITION METALS:
• Variable charges: Fe²⁺, Fe³⁺; Cu⁺, Cu²⁺.

POLYATOMIC IONS (memorise):
• NH₄⁺ (ammonium).
• OH⁻ (hydroxide).
• NO₃⁻ (nitrate).
• SO₄²⁻ (sulfate).
• CO₃²⁻ (carbonate).
• PO₄³⁻ (phosphate).

FORMULA WRITING:
1. Write cation first.
2. Balance charges using CRISS-CROSS.
3. Simplify ratios if possible.
4. Example: Al³⁺ and O²⁻ → Al₂O₃.

NAMING:
• Metal first.
• Non-metal ends in -IDE (chloride, oxide).
• Transition metals use Roman numerals (Iron(II) chloride = FeCl₂).

IONIC LATTICES:
• 3D repeating structure of alternating cations and anions.
• Electrostatic forces in every direction.

PROPERTIES EXPLAINED BY LATTICE:
• HIGH melting / boiling points (strong forces to overcome).
• HARD but BRITTLE (shifting layers → like charges align → repulsion → cleavage).
• SOLID state at room temp.
• ELECTRICAL CONDUCTIVITY:
– Solid: NO (ions fixed).
– Molten / dissolved: YES (ions mobile).
• OFTEN WATER-SOLUBLE (polar water separates ions).

LATTICE ENERGY:
• Energy released forming 1 mol of ionic solid from gaseous ions.
• Large + means strong lattice.
• Determines melting point, solubility.

DOT-CROSS DIAGRAMS:
• Show electron transfer.
• Example: Na → Cl shows one electron from Na to Cl.

REAL-WORLD EXAMPLES:
• NaCl (table salt).
• CaCO₃ (limestone).
• MgO (fire-bricks).

EXAM TIP:
• Always balance charges to determine formula.
• Watch for polyatomic ions in brackets when multiple.

Ionic rules

  • Metal + non-metal: Usually ionic.
  • Criss-cross: Balance charges.
  • Conductive: Only when molten/dissolved.

Quick reference

| Property | Ionic compounds |
|---|---|
| Structure | 3D lattice |
| Melting point | High |
| Conductivity | Only molten/dissolved |
| Brittleness | Brittle |
| Formula | Balance charges (e.g. MgCl₂) |

Worked example (Medium)

Write the formula for magnesium nitrate.

  1. Ions. Mg²⁺ and NO₃⁻.
  2. Balance. Need two NO₃⁻ per Mg²⁺.
  3. Brackets. Mg(NO₃)₂.

Answer: Mg(NO₃)₂.

Worked example (Hard)

Iron can form Fe²⁺ and Fe³⁺ ions. Write the formula for iron(III) sulfate and iron(II) oxide. Explain how the Roman numeral system resolves naming ambiguity.

  1. Iron(III) sulfate. Fe³⁺ and SO₄²⁻. Criss-cross charges: Fe₂(SO₄)₃. Brackets needed for polyatomic ion with subscript.
  2. Iron(II) oxide. Fe²⁺ and O²⁻. Charges balance 1:1 → FeO.
  3. Why Roman numerals. Transition metals have VARIABLE oxidation states. Without Roman numerals, "iron oxide" could mean FeO or Fe₂O₃ — completely different compounds. The Roman numeral specifies the metal's charge unambiguously.

Answer: Iron(III) sulfate = Fe₂(SO₄)₃. Iron(II) oxide = FeO. Roman numerals eliminate ambiguity for transition metals.

Common mistakes

  • Forgetting to simplify the criss-cross ratio When writing Ca²⁺ and O²⁻, criss-crossing gives Ca₂O₂ — but this simplifies to CaO. Students leave unsimplified formulas. Fix: Always reduce to the simplest whole-number ratio after criss-crossing.
  • Writing ionic formulas with charges shown Ionic compound formulas show the overall NEUTRAL formula, not individual ion charges. Writing NaCl as Na⁺Cl⁻ is incorrect for the formula. Fix: Charges are used to DERIVE the formula but are NOT shown in the final formula: NaCl, not Na⁺Cl⁻.
  • Confusing when to use brackets for polyatomic ions Brackets are only needed when the subscript for a polyatomic ion is greater than 1. Ca(OH)₂ needs brackets; NaOH does not. Fix: Only use brackets when the subscript on the polyatomic ion is 2 or more.

Why ionic compounds are so important in everyday life

Ionic compounds are everywhere — from the salt on your food (NaCl) to the limestone in buildings (CaCO₃) to the fluoride in toothpaste (NaF). Understanding their properties explains why they're used in specific ways.

High melting points make ionic compounds ideal for applications requiring heat resistance. Magnesium oxide (MgO, melting point 2852°C) is used to line furnaces and kilns. Aluminium oxide ($Al_2O_3$, known as alumina) is used in spark plugs and high-temperature ceramics.

Electrical conductivity when dissolved is exploited in electrolysis — passing current through molten or dissolved ionic compounds to extract metals. Aluminium is produced industrially by electrolysing molten $Al_2O_3$ dissolved in cryolite (the Hall-Héroult process). Electrolysis of brine (NaCl solution) produces chlorine gas, hydrogen gas, and sodium hydroxide — three essential industrial chemicals.

Solubility in water makes ionic compounds essential in biological systems. Your body relies on dissolved ions: Na⁺ and K⁺ for nerve impulses, Ca²⁺ for bone and muscle contraction, Fe²⁺/Fe³⁺ in haemoglobin for oxygen transport, and Cl⁻ for stomach acid (HCl).

Lattice energy — the energy released when gaseous ions form a solid lattice — determines many properties. Compounds with small, highly charged ions (like MgO: Mg²⁺ and O²⁻) have very high lattice energies and correspondingly high melting points. Compounds with large, singly charged ions (like CsI) have lower lattice energies and lower melting points. This relationship between ion size, charge, and properties is a favourite VCAA exam topic.

Practice questions

  1. Solid ionic compounds DO NOT conduct electricity because:

    1. They have no electrons.
    2. Ions are fixed in lattice.
    3. They are non-metals.
    4. They dissolve quickly.

    (1 marks)

  2. Explain why ionic compounds are hard but brittle.

    (3 marks)

Summary

  • Predict ions formed by elements.
  • Write ionic formulas.
  • Explain ionic compound properties.

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