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2.1.3Additional lessons

Ionic Lattices — IB Chemistry SL Revision Notes & Topical Questions

5 min read 6 parts Lesson 21 of 98

📚 IB Chemistry SL Hub / Topical Revision
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Curriculum Equation Key قوانين الرقم الهيدروجيني والاتزان الأيوني - pH & Acid-Base Equilibrium FormulasFast Exam Solution
pH = -log[H⁺] | pOH = -log[OH⁻] | pH + pOH = 14 | [H⁺][OH⁻] = 10⁻¹⁴
Word Equation & Conditions: الرقم الهيدروجيني = -لوغاريتم تركيز أيون الهيدروجين | pH + pOH = 14
Location & Specifications: الاتزان الأيوني، كيمياء 3 ثانوي، التحصيلي، و AP Chemistry.
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Part 1 of 6The Lattice Structure

Ionic solids do not contain separate molecules; the whole crystal is one continuous structure.

Regular arrangement
  • Regular arrangement Oppositely charged ions pack in a regular three-dimensional pattern called a giant ionic lattice.
Alternating ions
  • Alternating ions Each ion is surrounded by ions of opposite charge, which maximises attraction and minimises repulsion.
Sodium chloride
  • Sodium chloride In NaCl each Na⁺ is surrounded by six Cl⁻ ions and each Cl⁻ by six Na⁺ ions.
Strong bonding
  • Strong bonding Ionic bonding acts in all directions throughout the crystal, so many strong attractions must be broken to melt it.

Part 2 of 6Explaining the Properties

Explaining the Properties
High melting point
  • High melting point Large amounts of energy are needed to overcome the strong electrostatic attractions,
    • so melting points are high; NaCl melts at 801 °C.
Brittle
  • Brittle If layers shift, ions of the same charge become aligned and repel, so the crystal splits when struck.
Electrical conductivity
  • Electrical conductivity Solids do not conduct because ions are fixed, but molten and aqueous ionic compounds conduct since ions are free to move.
Solubility
  • Solubility Many ionic compounds dissolve in water because polar water molecules attract the ions and pull them from the lattice.

Part 3 of 6Strength of the Ionic Bond

1
The force between two ions increases as the product of their charges increases, so higher charges give stronger attraction.
2
The force also increases as the distance between ion centres decreases, so smaller ions give stronger attraction.
3
Compare NaCl (Na⁺, Cl⁻) with MgO (Mg²⁺, O²⁻): doubling both charges gives a much stronger lattice.
4
The stronger lattice requires more energy to break, so MgO melts at about 2852 °C compared with 801 °C for NaCl.
5
Therefore melting point trends in ionic compounds can be predicted from ionic charge and ionic radius.
CompoundIon chargesMelting point / °C
NaF1+, 1−996
NaCl1+, 1−801
MgO2+, 2−2852

Part 4 of 6Conductivity in Different States

Solid
  • Solid Ions are held in fixed positions so there are no mobile charge carriers and the solid does not conduct.
Molten
  • Molten Heat frees the ions from the lattice, and they carry current when a voltage is applied.
Aqueous
  • Aqueous Dissolved ions are surrounded by water molecules and move freely, so the solution conducts.
Electrolysis
  • Electrolysis Conduction in liquids is accompanied by chemical change at the electrodes, unlike metallic conduction.

Part 5 of 6Worked Example

Question Explain why magnesium oxide has a much higher melting point than sodium chloride.

  1. Both have giant ionic lattices held by electrostatic attraction between oppositely charged ions.
  2. Mg²⁺ and O²⁻ have charges of magnitude 2, while Na⁺ and Cl⁻ have charges of magnitude 1, so the force of attraction is greater.
  3. Mg²⁺ is smaller than Na⁺ and O²⁻ is smaller than Cl⁻, so the ions are closer together and the attraction is stronger still.

Answer More energy is needed to separate the ions in MgO because its ions have higher charges and smaller radii, so the melting point is higher.

Part 6 of 6Exam Checklist & Tips

2.1.3 – Key Points for the Exam

  • Ionic lattice A giant regular three-dimensional arrangement of oppositely charged ions.
  • Electrostatic attraction The force between oppositely charged particles.
  • Brittle Tending to shatter when force is applied rather than bending.
  • Aqueous Dissolved in water.

Common Mistakes

  • Saying that sodium chloride molecules melt; there are no molecules, only ions.
  • Explaining high melting points by strong covalent bonds, when the bonds in ionic compounds are electrostatic.
  • Claiming that solid ionic compounds conduct because of free electrons; the ions are fixed in place.
  • Forgetting to mention both charge and ion size when comparing lattice strength.

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