Did you know? Magnesium oxide and sodium chloride are both ionic solids, but the magnitude of their lattice energies is dramatically different. The explanation is not simply that magnesium oxide has more ions: it is the strength of attraction between highly charged, closely separated ions.
JC1 H1 Chemistry Tutor · Hougang, Singapore · 2027 8873. JC1 H1 Chemistry students can use ionic charge and ionic size to predict lattice-energy trends before being shown a numerical value. The numbers in worked examples are labelled as illustrative; this is not an official SEAB examination paper, classroom performance report or claim about available lessons.
Twelve chapters · Worked example · Try a new question · Questions and answers · Curriculum sources
Hougang Chemistry learning library · 2027 JC1 H1 Chemistry syllabus
Chemistry 234 · Continue across JC levels: JC1 H2 · Ethanol versus methoxymethane boiling points · JC2 H2 · Cyanohydrin addition and chirality · JC2 H3 · IR carbonyl and hydroxyl bands
Chemistry 234 — further original worked tutorials: JC2 H2 | Chlorinated ethanoic acid and induction · JC2 H3 | Conjugation and UV–visible red shifts
Learn the Chemistry in twelve connected chapters
- 1. 1. Define lattice energy before looking at a salt’s melting point
- 2. 2. Read the ionic formulae and calculate the actual ion charges
- 3. 3. Use a Coulombic model without pretending it is a complete crystal calculation
- 4. 4. Explain why ionic radius also matters
- 5. 5. Predict the direction of the enthalpy change
- 6. 6. Compare MgO and NaCl in one physical-reasoning table
- 7. 7. Use an explicitly illustrative energy comparison
- 8. 8. Relate strong ionic attractions to melting while respecting other factors
- 9. 9. Explain why larger lattice energy does not alone determine water solubility
- 10. 10. Transfer to NaF versus NaCl when the charges are the same
- 11. 11. Find the first error in five plausible wrong answers
- 12. 12. Finish with three independent lattice-energy rankings
1. Define lattice energy before looking at a salt’s melting point
In the formation convention used in this article, lattice energy is the enthalpy change when one mole of an ionic crystalline solid is formed from its gaseous constituent ions. For sodium chloride, Na⁺(g) + Cl⁻(g) → NaCl(s). For magnesium oxide, Mg²⁺(g) + O²⁻(g) → MgO(s).
Forming a stable ionic lattice from oppositely charged gaseous ions releases energy, so both formation lattice enthalpies are negative. This does not mean the attractive forces have a negative strength; the sign says that the system releases energy as the ions approach, arrange and stabilise one another.
The opposite process, pulling an ionic crystal apart into infinitely separated gaseous ions, requires energy input and is assigned a positive lattice dissociation enthalpy. Some books use the name “lattice energy” for the dissociation process. Always state your convention before comparing signs.
2. Read the ionic formulae and calculate the actual ion charges
Sodium, a Group 1 metal, commonly forms Na⁺; chlorine forms chloride Cl⁻. One of each balances charge, giving NaCl. Magnesium commonly forms Mg²⁺ and oxygen forms oxide O²⁻; one of each also balances to MgO. Both solids have a 1:1 formula-unit ratio, but the ionic charges are not the same.
That is the decisive contrast. Each ion in NaCl carries a charge magnitude of one elementary unit. Each ion in MgO carries a magnitude of two. Stronger charges lead to stronger electrostatic attraction, provided the other relevant structural conditions are comparable.
Do not write Mg⁺O⁻ simply because the empirical formula contains one atom of each element. A chemical formula’s subscripts establish the neutral numerical ratio, not the magnitude of the charged species. A formula unit of MgO contains Mg²⁺ and O²⁻, and that charge assignment must enter any lattice-energy explanation.
3. Use a Coulombic model without pretending it is a complete crystal calculation
In a very simplified electrostatic picture, attraction between opposite charges becomes stronger as the product of their charge magnitudes increases and as the distance between them decreases. At a fixed separation, the rough magnitude scales with |z₊z₋|, where z₊ and z₋ are the charge numbers.
For NaCl, |(+1)(−1)| = 1. For MgO, |(+2)(−2)| = 4. Under the deliberately unrealistic condition that two ion pairs had identical separation and surroundings, the charge term would make the MgO-like interaction approximately four times larger in magnitude.
Real lattice energies describe whole three-dimensional solids, not a single isolated pair. The crystal arrangement, ionic separation, repulsion and other contributions matter. Therefore the charge-product comparison is an explanatory trend, not an exact predicted fourfold ratio for these two real materials.
4. Explain why ionic radius also matters
Coulombic attraction increases as oppositely charged particles can approach more closely. A smaller ionic radius often permits a smaller nearest-neighbour ion separation, strengthening attraction. Magnesium’s +2 charge and small Mg²⁺ radius, combined with the oxide ion, produce a strongly attracting ionic lattice in the usual model.
One must avoid treating “which ion is smaller” as an independently sufficient answer when comparing compounds containing different ions. The nearest-neighbour separation depends on the combined ionic sizes and on the structural arrangement. A simple size comparison is most secure when charges and crystal type are otherwise similar.
Ionic radius is not a hard, perfectly defined solid-sphere boundary. Its numerical value can depend on coordination and conventions. The JC1 aim is a qualitative explanation: charges and distances both influence the electrostatic stabilisation of the lattice.
5. Predict the direction of the enthalpy change
Start with separated gaseous ions. When they assemble into a solid, the overall potential energy of the system decreases as attractive interactions dominate at the stable arrangement. The environment receives energy; thus the lattice-formation enthalpy is negative.
If the same stable lattice is separated into gaseous ions, energy must be supplied. The reverse reaction has a positive enthalpy of equal magnitude under equivalent state conditions. A question asking for “enthalpy of lattice dissociation” therefore expects the opposite sign from “enthalpy of lattice formation”.
This is an easy mark to lose even when a student can correctly say “magnesium oxide has stronger attractions”. Always label your equation and reactant states, then decide the sign. Memorising a negative number without the corresponding process is unreliable.
6. Compare MgO and NaCl in one physical-reasoning table
| Feature | Sodium chloride | Magnesium oxide |
|---|---|---|
| Cation | Na⁺ | Mg²⁺ |
| Anion | Cl⁻ | O²⁻ |
| Absolute charge product | 1 | 4 |
| Formula ratio | 1:1 | 1:1 |
| Attraction, qualitative | Weaker than MgO | Stronger than NaCl |
| Formation lattice energy | Negative | More negative in magnitude |
The phrase “more negative” refers to the enthalpy of formation from gaseous ions. For example, a change of −3800 kJ mol⁻¹ has a larger absolute magnitude than a change of −780 kJ mol⁻¹, even though −3800 is a smaller number on the usual number line.
A student must distinguish “larger in magnitude” from “larger signed numerical value”. A common misconception is to compare two negative enthalpies as if more negative meant a smaller physical attraction. In formation language, the larger released amount corresponds to greater stabilisation of the ionic lattice.
7. Use an explicitly illustrative energy comparison
Suppose an invented classroom data table gives a lattice-formation enthalpy of −780 kJ mol⁻¹ for NaCl and −3800 kJ mol⁻¹ for MgO. These are rounded teaching figures of the expected order of magnitude, not an instruction to quote exact reference values in every assessment.
The absolute magnitudes are 780 and 3800. Their illustrative ratio is 3800 ÷ 780 ≈ 4.87. That is not exactly four, although the charge-product term alone changed by a factor of four. The difference reminds us that ionic separation and crystal-level contributions also matter.
| Model quantity | NaCl | MgO |
|---|---|---|
| Illustrative ΔH lattice formation | −780 kJ mol⁻¹ | −3800 kJ mol⁻¹ |
| Magnitude |ΔH| | 780 kJ mol⁻¹ | 3800 kJ mol⁻¹ |
| Ion charge product | 1 | 4 |
One independent check: the energy required for complete lattice dissociation under the opposite convention would be approximately +780 versus +3800 kJ mol⁻¹. Students should explain the sign reversal without changing the physical ranking of attraction strength.
8. Relate strong ionic attractions to melting while respecting other factors
To melt an ionic crystal, enough energy must be supplied to disrupt its ordered lattice and permit ions to move relative to one another. Strong electrostatic interactions generally contribute to high melting points. Magnesium oxide is a well-known refractory material, consistent with very strong ionic attraction.
However, melting temperature is not equal numerically to lattice enthalpy. It depends on the balance of enthalpy and entropy at phase equilibrium, and the bonding, structural and thermal properties of the material. Quoting the lattice-energy magnitude as though it were a melting temperature in °C would mix unlike physical quantities.
Also avoid saying an ionic solid melts by converting every ion into a neutral atom. A melt still consists of charged species. The mobility changes dramatically; it is not a reversal of ionic electron transfer. This links the discussion to conductivity in solid, molten and aqueous states.
9. Explain why larger lattice energy does not alone determine water solubility
Dissolving an ionic solid in water involves more than separating ions. The crystal lattice must be disrupted, but water molecules can stabilise separated ions through hydration. Whether dissolving is thermodynamically favourable depends on competing enthalpy and entropy effects.
A simplistic claim that “higher lattice energy always means lower solubility” is unsafe. Strongly charged ions tend to interact strongly with water as well as with other ions. To predict solubility quantitatively, the full free-energy balance and actual equilibrium conditions are relevant.
For the H1 comparison, it is sufficient to state that strong lattice attraction makes separating the ions energetically demanding. But a separate solubility problem needs appropriate solvation and other data. Good Chemistry answers stop where their given information stops.
10. Transfer to NaF versus NaCl when the charges are the same
Now compare sodium fluoride and sodium chloride. Both contain Na⁺ with a halide ion of charge −1, so their ionic charge products are equal. The difference is that F⁻ is smaller than Cl⁻ under comparable coordination conventions. The Na⁺ and F⁻ ions can generally approach more closely in the ionic structure.
Consequently, a simple charge-and-size model predicts a greater magnitude of lattice attraction for NaF than NaCl, provided other relevant lattice-structure factors are comparable. This example isolates the effect of ionic size much more cleanly than comparing MgO with NaCl, where both charges and ion identities change.
Do not claim identical crystal structures without checking. The qualitative radius reasoning is the learning target; a complete numerical lattice-energy calculation requires structural and thermodynamic details.
11. Find the first error in five plausible wrong answers
- “MgO and NaCl have the same lattice strength because both are one cation plus one anion.” The ion charges differ.
- “Lattice formation requires energy because breaking a salt crystal requires energy.” Formation and dissociation are opposite processes.
- “Four times the charge product proves exactly four times the lattice enthalpy.” Ion separation and whole-lattice structure also matter.
- “The larger negative number means weaker attraction.” Lattice-formation magnitude and numerical ordering of negatives are different ideas.
- “A higher lattice energy alone determines exact water solubility.” Hydration and entropy contributions are missing.
A tutor should point to the original equation first. If the student chose the wrong enthalpy sign, fix that before teaching more electrostatics. If they identified the sign correctly but made an exact fourfold prediction, the missing concept is the dependence on separation and lattice arrangement.
One successful answer can be concise: Mg²⁺ and O²⁻ carry greater charge magnitudes than Na⁺ and Cl⁻ and have strongly interacting ions, leading to a larger magnitude of electrostatic lattice stabilisation in MgO.
12. Finish with three independent lattice-energy rankings
Case A: Compare MgO and CaO. Both contain a +2 metal and O²⁻. Mg²⁺ is smaller than Ca²⁺, so, with comparable lattice considerations, MgO is expected to have the greater magnitude of lattice energy.
Case B: Compare NaF and NaCl. Charges match; fluoride is smaller, so NaF is expected to have the larger magnitude. Case C: Compare lattice formation and lattice dissociation for the same NaCl crystal. Their enthalpies are equal in magnitude and opposite in sign for the exactly reversed processes under the same standard conditions.
- Write a balanced equation with explicit gaseous ions and solid product.
- Identify actual formal charges of both ions.
- Compare charge products when compositions differ.
- Compare nearest-neighbour separation and ionic radius with appropriate caution.
- State whether the sign convention is formation or dissociation.
- Give a qualitative conclusion without inventing an exact ratio or solubility.
A student who can transfer the reasoning across those three cases understands charge, size and energy direction rather than repeating one remembered sentence about magnesium oxide.
Frequently asked questions
Why is MgO’s lattice-energy magnitude much larger than NaCl’s?
MgO has Mg²⁺ and O²⁻ with larger ion charges and strong attraction at their stable separation. NaCl contains singly charged ions, producing weaker electrostatic lattice stabilisation.
Is lattice energy positive or negative?
It depends on the convention. Formation of solid from gaseous ions is negative; dissociation into gaseous ions is the opposite positive enthalpy.
Does a charge-product ratio of four mean exactly four times the lattice energy?
No. A real crystal’s nearest-neighbour distances and full lattice interactions also matter.
Can lattice energy alone tell me whether a salt dissolves?
No. Hydration, entropy and other equilibrium considerations affect dissolution.
Is this in JC1 H1 Chemistry?
Yes. H1 Chemistry 8873 includes lattice energy and qualitative effects of ionic charge and radius on its magnitude.
Turn the explanation into a productive tutorial
A strong Chemistry tutorial starts with the student’s own prediction, not the displayed worked answer. Ask them to distinguish what is known, what is assumed and which chemical representation is relevant. Then let the learner draw the species or structure, complete the working, check units and explain a changed-context problem without copying the example. The first mistaken assumption is a more useful correction target than simply marking the last number wrong.
The unchanged Secondary 1 Mathematics Tutor Clementi small-group reference illustrates eduKateSG’s separate Mathematics teaching approach of close attention and sequenced practice. It is an immutable format and teaching reference, not a claim that this Chemistry tutorial takes place in Clementi or has an available seat. For actual arrangements contact eduKateSG.
Syllabus, scope and related reading
The authoritative course reference is the 2027 Singapore-Cambridge JC1 H1 Chemistry 8873 syllabus, listed in SEAB’s 2027 A-Level school-candidate directory. The H1 Chemistry 8873 chemical energetics outcomes include lattice energy for gaseous ions forming an ionic solid, and qualitative effects of ionic charge and ionic radius on lattice-energy magnitude. The comparison here is qualitative first; any numerical values are explicitly labelled illustrative rather than presented as data-booklet requirements. Full Born–Haber cycles belong to a broader energetics discussion.
Related explanations: Born–Haber cycle sign audit · Ionic solids and mobile ions. Worked numbers are explicitly educational models, not measured results. If a question or syllabus supplies different values, its specified data and conditions take precedence.
Original eduKateSG Chemistry educational article · reviewed 30 October 2026 · first principles, clear examples, transparent limits, independent transfer.