Lattice Enthalpy: Definition and Significance

Spec Mapping — OCR H432 Module 5.2.1 — Lattice enthalpy, covering the formal definition of lattice enthalpy under the OCR formation convention ($\Delta_{LE} H^\ominus$ΔLE​H⊖ for one mole of ionic solid formed from gaseous ions under standard conditions), the resulting negative sign for all lattice enthalpies, the physical interpretation of lattice-enthalpy magnitude as a measure of ionic bond strength, the inaccessibility of direct calorimetric measurement, and qualitative correlation with the physical properties of ionic solids — melting point, hardness, brittleness, solubility, and electrical conductivity (refer to the official OCR H432 specification document for exact wording).

Lattice enthalpy is the gateway concept of OCR Module 5.2.1 — it underpins Born–Haber cycles (Lessons 2 and 3), the trends in ionic-bond strength across the periodic table (Lesson 4), the energetics of dissolving (Lesson 5), and ultimately the entropy and free-energy treatments that close out Module 5.2.2. Despite its centrality, the OCR definition is subtle: it specifies a direction (gaseous ions to solid lattice, never the reverse), a quantity (one mole of compound, not one mole of ions), and a sign convention (always negative under OCR, the opposite of the convention used by some other awarding bodies and textbooks). Getting the definition exactly right is the single most-marked recall item in the Year 13 thermodynamics paper. This lesson develops the OCR definition with state symbols and worked equation-writing examples, contrasts the OCR sign convention against the rival "lattice dissociation" convention used by AQA and older textbooks, explains why the gaseous-ion starting state makes direct measurement impossible, and traces the chain from lattice-enthalpy magnitude to the macroscopic physical properties (melting point, hardness, solubility, conductivity) that you can observe on the school-laboratory bench. Lesson 2 will then construct the Born–Haber cycle that lets us calculate lattice enthalpies from experimentally accessible data.

Key Equation: the OCR definition of lattice enthalpy is written as a balanced equation with state symbols. For a general ionic compound $M_a X_b$Ma​Xb​ formed from cations $M^{n+}$Mn+ and anions $X^{m-}$Xm−: $a\,M^{n+}(g) + b\,X^{m-}(g) \rightarrow M_a X_b(s) \quad \Delta_{LE} H^\ominus < 0$aMn+(g)+bXm−(g)→Ma​Xb​(s)ΔLE​H⊖<0 where the coefficients $a$a and $b$b are chosen so that exactly one mole of $M_a X_b$Ma​Xb​ appears on the right-hand side, and the enthalpy change refers to this one mole of compound (not to one mole of ions).

---

Learning Objectives (OCR Spec 5.2.1)

By the end of this lesson you should be able to:

• State the OCR definition of lattice enthalpy precisely, including direction, quantity, and standard-conditions clause
• Write a balanced equation with state symbols representing the lattice enthalpy of any binary ionic compound
• Explain why lattice enthalpy is always negative (exothermic) under the OCR formation convention
• Contrast the OCR formation convention against the rival dissociation convention (positive) used by AQA and older texts
• Recognise the magnitude of lattice enthalpy as a measure of ionic bond strength
• Justify why lattice enthalpy cannot be measured directly and must be found via a Born–Haber cycle
• Relate lattice enthalpy to the physical properties of ionic compounds (melting point, hardness, solubility, conductivity)

What is Lattice Enthalpy?

When isolated gaseous ions come together to build an ionic lattice, strong electrostatic attractions between the oppositely charged ions release a large amount of energy. This energy release is what holds the solid crystal together and what we quantify as the lattice enthalpy. The process is the inverse of the imaginary dissociation in which a sodium-chloride crystal would have to vaporise into individual gaseous Na$^+$+ and Cl$^-$− ions — that hypothetical process is endothermic by exactly the same magnitude as lattice enthalpy is exothermic. OCR has chosen, by convention, to define the quantity in the formation direction (gas → solid), so it is always negative.

OCR uses the formation convention for lattice enthalpy. The OCR definition, in the exact wording you should reproduce in any extended-answer question:

Lattice enthalpy ($\Delta_{LE} H^\ominus$ΔLE​H⊖ or $\Delta_{lat} H^\ominus$Δlat​H⊖) is the enthalpy change when one mole of an ionic compound is formed from its constituent gaseous ions under standard conditions.

Each clause carries marks in the OCR mark scheme. Picking them apart:

• Enthalpy change — not "energy" or "heat". OCR explicitly requires the term enthalpy.
• One mole of the ionic compound (not one mole of ions). This is the stoichiometric reference.
• Ionic compound — the species formed is the solid crystalline product, e.g. NaCl(s), MgO(s), Al$_2$2​O$_3$3​(s).
• Formed from — the direction is gas → solid. Reversing the direction gives the lattice dissociation enthalpy, which OCR does not use.
• Constituent gaseous ions — the starting state is the isolated cations and anions in the gas phase, well separated and non-interacting.
• Standard conditions — 100 kPa pressure, stated temperature (usually 298 K), with all species in their standard states except for the gaseous ions (which are by convention treated as an ideal-gas mixture of monoatomic ions at standard pressure).

CRITICAL — OCR CONVENTION: OCR defines lattice enthalpy as an energy release accompanying the formation of the lattice from gaseous ions. It is therefore always negative. This is DIFFERENT from the dissociation convention used by AQA, by older A-Level textbooks, and by some university physical-chemistry texts, in which lattice enthalpy is defined as the energy required to break apart the lattice into gaseous ions (always positive). Both conventions describe the same physical reality; only the sign and direction differ. Under OCR you should ALWAYS write lattice enthalpy as a negative number. If you encounter a positive value in a textbook, that text is using the dissociation convention — flip the sign before substituting into any OCR Born–Haber cycle.

Writing Equations for Lattice Enthalpy

The OCR examiner expects fully balanced equations with state symbols. Three worked examples below illustrate the conventions.

Sodium chloride (one cation, one anion, both singly charged):

$\text{Na}^+(g) + \text{Cl}^-(g) \rightarrow \text{NaCl}(s) \quad \Delta_{LE} H^\ominus = -787\ \text{kJ mol}^{-1}$Na+(g)+Cl−(g)→NaCl(s)ΔLE​H⊖=−787 kJ mol−1

Magnesium oxide (both ions doubly charged):

$\text{Mg}^{2+}(g) + \text{O}^{2-}(g) \rightarrow \text{MgO}(s) \quad \Delta_{LE} H^\ominus = -3791\ \text{kJ mol}^{-1}$Mg2+(g)+O2−(g)→MgO(s)ΔLE​H⊖=−3791 kJ mol−1

Notice the state symbols: ions on the left are (g) to show they are isolated gaseous ions, and the product is (s) because ionic compounds are solids at standard temperature. Writing (aq) on the ions on the left, or (l) on the solid, is a one-mark deduction.

Calcium chloride (stoichiometry matters):

$\text{Ca}^{2+}(g) + 2\,\text{Cl}^-(g) \rightarrow \text{CaCl}_2(s) \quad \Delta_{LE} H^\ominus = -2258\ \text{kJ mol}^{-1}$Ca2+(g)+2Cl−(g)→CaCl2​(s)ΔLE​H⊖=−2258 kJ mol−1

You must balance the equation so that exactly one mole of the ionic compound is formed. For CaCl$_2$2​ that means 2 moles of Cl$^-$− ions are needed on the left, but the enthalpy change is per mole of CaCl$_2$2​, not per mole of ions. A common mistake is to write $\frac{1}{2}\text{Ca}^{2+} + \text{Cl}^-$21​Ca2++Cl− — fractional ions are not permitted on the left because individual ions are discrete particles.

Aluminium oxide (high charges, complex stoichiometry):

$2\,\text{Al}^{3+}(g) + 3\,\text{O}^{2-}(g) \rightarrow \text{Al}_2\text{O}_3(s) \quad \Delta_{LE} H^\ominus \approx -15{,}900\ \text{kJ mol}^{-1}$2Al3+(g)+3O2−(g)→Al2​O3​(s)ΔLE​H⊖≈−15,900 kJ mol−1

The enormous magnitude reflects both the high charges (product $|q_1 q_2| = 6$∣q1​q2​∣=6) and the small ionic radii of Al$^{3+}$3+ and O$^{2-}$2−.

Why is Lattice Enthalpy Exothermic?

Under the OCR convention, gaseous ions (well separated, in the gas phase, non-interacting) come together to form a regular array of ions tightly bound by electrostatic attraction. As the oppositely charged ions approach from infinite separation, the system descends a Coulombic potential-energy well: $E \propto -\frac{q_+ q_-}{r}$E∝−rq+​q−​​. The ionic lattice sits at a much lower potential energy than the separated gaseous ions, and this energy difference is released to the surroundings as heat — typically of the order of hundreds to thousands of kJ per mole of compound.

The magnitude of the lattice enthalpy is a direct measure of how strongly the ions attract each other in the lattice. A more negative (more exothermic) lattice enthalpy means a stronger ionic bond — and as we will see in Lesson 4, this depends predictably on ionic charge and inter-ionic distance.

| Compound | Lattice enthalpy / kJ mol$^{-1}$−1 | Charge product $|q_+ q_-|$∣q+​q−​∣ | Comment | |----------|----------------------------------:|:--------------------------:|---------| | NaCl | $-787$−787 | 1 | reference 1+/1- benchmark | | NaBr | $-751$−751 | 1 | larger anion, slightly weaker | | NaI | $-705$−705 | 1 | even larger anion | | LiF | $-1031$−1031 | 1 | smallest ions in Group 1 + Group 17 | | CsI | $-604$−604 | 1 | largest ions in Group 1 + Group 17 | | MgCl$_2$2​ | $-2526$−2526 | 2 | doubly charged cation | | MgO | $-3791$−3791 | 4 | both ions doubly charged, small radii | | CaO | $-3401$−3401 | 4 | larger cation than Mg$^{2+}$2+ | | Al$_2$2​O$_3$3​ | $\approx -15{,}900$≈−15,900 | 6 | trebly + doubly charged, small radii |

Lattice enthalpies for simple 1+/1- salts are around $-600$−600 to $-1000$−1000 kJ mol$^{-1}$−1; for 2+/2- compounds they are typically $-3000$−3000 to $-4000$−4000 kJ mol$^{-1}$−1; and for 3+/2- compounds like Al$_2$2​O$_3$3​ they exceed $-15{,}000$−15,000 kJ mol$^{-1}$−1. The factor-of-4 jump in going from NaF to MgO (both ions doubly charged) follows directly from the $|q_+ q_-|$∣q+​q−​∣ product in Coulomb's law — a foretaste of the quantitative trends developed in Lesson 4.

Why Can't Lattice Enthalpy Be Measured Directly?

Unlike the enthalpy of combustion (which can be measured cleanly in a bomb calorimeter — see Lesson 2 of the rates-and-equilibrium course) or the enthalpy of neutralisation (a simple cup-calorimeter experiment), lattice enthalpy cannot be measured in any calorimeter. The reason is the starting state: the OCR definition specifies gaseous ions, but isolated Na$^+$+(g) and Cl$^-$−(g) at standard pressure are not experimentally accessible. To create a flask containing 1 mole of Na$^+$+(g) and 1 mole of Cl$^-$−(g) under standard conditions, you would have to:

1. Take 1 mole of solid sodium and vaporise it to Na(g) (atomisation).
2. Ionise each Na(g) atom by stripping its outer electron — at extremely high temperature or under intense electric discharge.
3. Take half a mole of Cl$_2$2​(g), dissociate it to 1 mole of Cl(g) atoms, and attach an electron to each.
4. Keep all these species apart so they don't immediately recombine — by separating them in a vacuum chamber over macroscopic distances.

Even if all this were feasible, the moment Na$^+$+(g) and Cl$^-$−(g) meet they recombine violently to form NaCl(s), releasing $\approx 787$≈787 kJ mol$^{-1}$−1 as heat plus a flash of light. The reaction is uncontrolled and the resulting solid is not in thermal equilibrium with the surroundings, so calorimetric measurement is impossible.

Instead, lattice enthalpy is found indirectly using a Born–Haber cycle — an application of Hess's law that links the lattice enthalpy to experimentally measurable quantities such as enthalpies of formation (from calorimetry), enthalpies of atomisation (from spectroscopy and vapour-pressure data), ionisation energies (from photoelectron spectroscopy), and electron affinities (from ion-pair production). Lessons 2 and 3 of this course develop the Born–Haber cycle in full detail.

Physical Significance of Lattice Enthalpy

The lattice enthalpy governs many of the macroscopic properties that make ionic compounds distinctive in the laboratory and in industry:

Melting and boiling points. Breaking up the lattice into a liquid (or eventually a gas) requires supplying energy comparable in magnitude to the lattice enthalpy. MgO, with its very exothermic lattice enthalpy ($-3791$−3791 kJ mol$^{-1}$−1), melts at 2852 °C and is used to line industrial furnaces and the bricks of high-temperature kilns. NaCl melts at 801 °C, much lower than MgO because its lattice enthalpy is roughly five times smaller in magnitude.

Hardness and brittleness. Strong ionic attractions give hard crystals (NaCl scratches gypsum and talc, MgO scratches steel), but the same regular alignment of charges means a small shear displacing one layer by half a lattice spacing brings like charges face-to-face, generating strong electrostatic repulsion that splits the crystal along the cleavage plane. This is why ionic crystals are hard yet brittle — quite unlike metals, which deform plastically.

Solubility in water. Whether an ionic compound dissolves depends on whether the energy released in hydrating the ions can compensate for the energy required to break apart the lattice (Lesson 5). Compounds with very large $-\Delta_{LE} H^\ominus$−ΔLE​H⊖ tend to be insoluble (or sparingly soluble) — MgO is essentially insoluble in cold water, BaSO$_4$4​ is medically used as an X-ray contrast agent precisely because it does not dissolve.

Electrical conductivity. Solid ionic compounds do not conduct electricity because the ions are locked in place in the lattice. But once the lattice is broken — by melting (high $T$T needed for large $-\Delta_{LE}$−ΔLE​) or by dissolving in water — the ions become mobile and the substance becomes a strong electrolyte. The cost of breaking the lattice is again set by the magnitude of $-\Delta_{LE} H^\ominus$−ΔLE​H⊖.

Visualising the Energy Change

graph TD
  A["Gaseous ions<br/>Na+(g) + Cl-(g)<br/>Higher enthalpy"] -->|Lattice enthalpy<br/>ΔLE H = -787 kJ/mol<br/>EXOTHERMIC| B["Ionic solid<br/>NaCl(s)<br/>Lower enthalpy"]

The arrow points downward in energy: enthalpy is released going from the scattered gaseous ions to the ordered solid. The "length" of the arrow (the magnitude of $\Delta_{LE} H^\ominus$ΔLE​H⊖) is proportional to the strength of ionic bonding in the crystal.

Worked Example 1: Writing the Correct Equation

Question: Write an equation, with state symbols, to represent the lattice enthalpy of magnesium fluoride, MgF$_2$2​.

Answer: MgF$_2$2​ contains 1 Mg$^{2+}$2+ and 2 F$^-$− ions per formula unit. To form one mole of MgF$_2$2​(s) we need 1 mole of Mg$^{2+}$2+(g) and 2 moles of F$^-$−(g):

$\text{Mg}^{2+}(g) + 2\,\text{F}^-(g) \rightarrow \text{MgF}_2(s) \quad \Delta_{LE} H^\ominus = -2957\ \text{kJ mol}^{-1}$Mg2+(g)+2F−(g)→MgF2​(s)ΔLE​H⊖=−2957 kJ mol−1

The lattice enthalpy is large because the cation is doubly charged and small ($r_{\text{Mg}^{2+}} = 72$rMg2+​=72 pm) and the anion is small ($r_{\text{F}^-} = 133$rF−​=133 pm).

Worked Example 2: Interpreting a Lattice Enthalpy Value

Question: The lattice enthalpy of CaO is $-3401$−3401 kJ mol$^{-1}$−1. Interpret this value in words.

Answer: When 1 mole of solid CaO is formed from 1 mole of gaseous Ca$^{2+}$2+ ions and 1 mole of gaseous O$^{2-}$2− ions under standard conditions, $3401$3401 kJ of energy is released to the surroundings as heat. The negative sign shows the process is exothermic. The large magnitude reflects the strong electrostatic attraction between the doubly charged ions at relatively short inter-ionic distance ($r_{\text{Ca}^{2+}} + r_{\text{O}^{2-}} = 100 + 140 = 240$rCa2+​+rO2−​=100+140=240 pm).

Worked Example 3: Spotting a Wrong Equation

Question: A student writes the lattice enthalpy of K$_2$2​O as:

$2\,\text{K}^+(aq) + \text{O}^{2-}(g) \rightarrow \text{K}_2\text{O}(s) \quad \Delta_{LE} H^\ominus = +2238\ \text{kJ mol}^{-1}$2K+(aq)+O2−(g)→K2​O(s)ΔLE​H⊖=+2238 kJ mol−1

Identify three errors.

Answer:

1. The cation state symbol is (aq) — it should be (g) because lattice enthalpy starts from gaseous ions, not aqueous ions.
2. The sign of $\Delta_{LE} H^\ominus$ΔLE​H⊖ is positive — under OCR convention lattice enthalpy is always negative (the value should be $-2238$−2238 kJ mol$^{-1}$−1).
3. The equation is balanced for "one mole of compound" correctly with 2K$^+$+, so the stoichiometry is fine — but the student has invoked the dissociation convention rather than the OCR formation convention, which is why both the state symbol and the sign are wrong.

Worked Example 4: Comparing Two Lattice Enthalpies

Question: Without doing a calculation, predict whether NaF or NaCl has the more exothermic lattice enthalpy, and justify your choice.

Answer: Both compounds have ions of the same charges (+1 and $-$−1), so the charge product is identical. The cation (Na$^+$+) is the same in both. The anion differs: F$^-$− ($r = 133$r=133 pm) is smaller than Cl$^-$− ($r = 181$r=181 pm). Coulomb's law gives $E \propto |q_+ q_-|/r$E∝∣q+​q−​∣/r, so the smaller inter-ionic distance in NaF gives a stronger electrostatic attraction, and hence a more exothermic lattice enthalpy. Indeed NaF is $-918$−918 kJ mol$^{-1}$−1 versus NaCl at $-787$−787 kJ mol$^{-1}$−1.

Worked Example 5: Sign-Convention Conversion

Question: A chemistry textbook quotes the lattice enthalpy of KBr as "+670 kJ mol$^{-1}$−1". What value should you use in an OCR Born–Haber cycle, and why?

Answer: The textbook is using the dissociation convention (lattice breaking apart into gaseous ions, always positive). To convert to the OCR formation convention (gaseous ions forming the lattice, always negative), reverse the sign:

$\Delta_{LE} H^\ominus(\text{KBr, OCR}) = -670\ \text{kJ mol}^{-1}$ΔLE​H⊖(KBr, OCR)=−670 kJ mol−1

Always check the source's sign convention before substituting a tabulated lattice enthalpy into an OCR calculation. The Data Booklet values supplied in the OCR exam paper itself will already be in the formation convention.

---

Synoptic Links

Synoptic Links — Connects to:

• ocr-alevel-chemistry-enthalpy-rates-equilibrium / enthalpy-changes-and-standard-conditions (standard-state definitions, sign conventions, the meaning of $\Delta H^\ominus$ΔH⊖ — all directly applicable to lattice enthalpy).
• ocr-alevel-chemistry-enthalpy-rates-equilibrium / hess-law (state-function reasoning that justifies the use of Born–Haber cycles to extract lattice enthalpy from experimentally accessible quantities).
• ocr-alevel-chemistry-atomic-structure-periodicity / ionic-bonding (electrostatic model of ionic bonding, ionic radii, the formal source of ions from atoms via ionisation and electron capture).
• ocr-alevel-chemistry-energetics-electrode / born-haber-cycles (Lesson 2 — the cycle that uses this lesson's definition to calculate lattice enthalpy from measurable data).
• ocr-alevel-chemistry-energetics-electrode / factors-affecting-lattice-enthalpy (Lesson 4 — quantitative correlation of $\Delta_{LE} H^\ominus$ΔLE​H⊖ with ionic charge and radius).

Practical Activity Group anchor: PAG 3 — Enthalpy determination. PAG 3 has students measure $\Delta_{sol} H^\ominus$Δsol​H⊖ for an ionic salt by simple cup calorimetry, then combine this with tabulated hydration enthalpies in a Hess cycle to extract the lattice enthalpy indirectly. Lesson 5 develops this cycle in full; the present lesson provides the conceptual definition of the lattice-enthalpy term that PAG 3 ultimately measures.

---

Specimen question modelled on the OCR H432 paper format

Question (6 marks): (a) State the OCR definition of lattice enthalpy. (b) Write an equation, with state symbols, representing the lattice enthalpy of calcium chloride, CaCl$_2$2​. (c) Explain why lattice enthalpy cannot be measured directly by calorimetry, and identify the alternative method.

AO breakdown

Mark	AO	Awarded for
1	AO1	"Enthalpy change when one mole of ionic compound..."
2	AO1	"...formed from its gaseous ions under standard conditions"
3	AO2	Correct balanced equation: $\text{Ca}^{2+}(g) + 2\text{Cl}^-(g) \rightarrow \text{CaCl}_2(s)$Ca2+(g)+2Cl−(g)→CaCl2​(s)
4	AO2	Correct state symbols throughout (gas, gas, solid)
5	AO3	Reasoning: isolated gaseous ions cannot be prepared in a calorimeter
6	AO3	Identification of the Born–Haber cycle as the indirect method

AO split: AO1 = 2, AO2 = 2, AO3 = 2.

Mid-band response (4/6):

(a) Lattice enthalpy is the enthalpy change when one mole of an ionic compound is formed from gaseous ions.

(b) Ca$^{2+}$2+(g) + 2Cl$^-$−(g) $\rightarrow$→ CaCl$_2$2​(s).

(c) You cannot measure lattice enthalpy directly because the gaseous ions are too reactive. You have to use a Born–Haber cycle instead.

Examiner commentary: M1 (AO1) for the formation direction; M3 and M4 (AO2) for the correctly written balanced equation with state symbols; M6 (AO3) for naming the Born–Haber cycle. To reach top-band: the AO1 definition is missing "under standard conditions" (lose M2), and the AO3 explanation does not say why gaseous ions cannot be prepared — "too reactive" is vague. The candidate needs to explain that isolated gaseous ions at standard pressure cannot be assembled in a calorimeter because as soon as they meet they recombine violently to form the solid lattice, releasing energy uncontrollably.

Top-band response (6/6):

(a) Lattice enthalpy is the enthalpy change when one mole of an ionic compound is formed from its constituent gaseous ions under standard conditions (100 kPa, stated temperature). Under the OCR formation convention, it is always negative because the process is exothermic.

(b) $\text{Ca}^{2+}(g) + 2\,\text{Cl}^-(g) \rightarrow \text{CaCl}_2(s) \quad \Delta_{LE} H^\ominus = -2258\ \text{kJ mol}^{-1}$Ca2+(g)+2Cl−(g)→CaCl2​(s)ΔLE​H⊖=−2258 kJ mol−1

The stoichiometric coefficient 2 ensures one mole of CaCl$_2$2​ is formed; the enthalpy change refers to one mole of CaCl$_2$2​, not one mole of ions.

(c) The starting state — isolated Ca$^{2+}$2+(g) and Cl$^-$−(g) ions at standard pressure — cannot be prepared in any calorimetric vessel. Even if individual gaseous ions could be generated (by atomisation, ionisation, and electron capture from the gas-phase elements), the moment they were brought together they would recombine violently to form CaCl$_2$2​(s) with the release of $\sim$∼2260 kJ mol$^{-1}$−1 as heat and light, well outside any condition of thermal equilibrium needed for calorimetry. Instead, lattice enthalpy is determined indirectly via a Born–Haber cycle: a Hess-law cycle linking the lattice-formation step to experimentally measurable quantities ($\Delta_f H^\ominus$Δf​H⊖, $\Delta_{at} H^\ominus$Δat​H⊖, ionisation energies, electron affinities), all of which can be measured calorimetrically or spectroscopically.

Examiner commentary: All six marks awarded — M1 and M2 for a complete OCR-style definition (formation direction + standard conditions); M3 and M4 for the balanced equation with state symbols and explicit one-mole stoichiometry; M5 and M6 for a full AO3 explanation of both why direct measurement fails and how the Born–Haber cycle resolves the problem. The "violent recombination" mechanistic insight and the explicit "Hess-law cycle" framing are A* discriminators.

---

Common errors and mark-loss patterns

Pedagogical observations from OCR examiner reports:

• Writing lattice enthalpy as a positive value — this is the AQA/dissociation convention, not the OCR formation convention. OCR awards zero for a positive value.
• Forgetting the (g) state symbol on ions, or writing (aq) instead. The OCR mark scheme is strict on state symbols.
• Writing the equation the wrong way round (solid → gaseous ions) — that is lattice dissociation enthalpy, which OCR does not use.
• Failing to balance the stoichiometry: e.g. writing $\text{Mg}^{2+}(g) + \text{Cl}^-(g) \rightarrow \text{MgCl}_2(s)$Mg2+(g)+Cl−(g)→MgCl2​(s) without the 2 in front of Cl$^-$−. The equation must form exactly one mole of compound.
• Saying "energy needed" instead of "energy released" — under OCR lattice enthalpy is energy released.
• Confusing lattice enthalpy with enthalpy of formation ($\Delta_f H^\ominus$Δf​H⊖). Formation starts from elements in standard states; lattice enthalpy starts from gaseous ions. The two differ by atomisation + ionisation + electron affinity (the rest of the Born–Haber cycle).
• Quoting "energy" rather than "enthalpy change" in the definition — OCR mark schemes require the term enthalpy.
• Forgetting to state "under standard conditions" in the definition — a one-mark deduction.

---

Going further

The lattice-enthalpy concept extends well beyond A-Level into the heart of crystal-energetics research. The Madelung constant $M$M — a geometrical factor characterising the long-range electrostatic summation over an infinite crystal lattice — multiplies the Coulombic pair-interaction $-\frac{q_+ q_-}{4\pi\epsilon_0 r_0}$−4πϵ0​r0​q+​q−​​ to give the lattice energy of any specific crystal structure (rock-salt structure $M = 1.7476$M=1.7476, caesium-chloride structure $M = 1.7627$M=1.7627, zinc-blende structure $M = 1.6381$M=1.6381). The Born–Landé equation incorporates a short-range repulsive term (the "Born exponent" $n$n, typically 5–12, capturing the Pauli repulsion of closed shells) to give theoretical lattice enthalpies accurate to a few percent for purely ionic crystals. The Kapustinskii equation provides a quick estimate when the Madelung constant is unknown, requiring only the number of ions per formula unit, their charges, and the sum of their radii. Beyond classical electrostatics, ab-initio quantum-chemistry methods (DFT with periodic boundary conditions) now compute lattice enthalpies to $\sim$∼10 kJ mol$^{-1}$−1 accuracy for hundreds of ionic compounds — but the OCR formation-convention definition we have used here remains the unambiguous experimental reference. Recommended reading: Atkins, Physical Chemistry, ch. 19; Burdett, Chemical Bonding in Solids, ch. 4; Born and Haber's joint 1919 paper that introduced both the lattice-enthalpy concept and the eponymous cycle (Born and Haber, Verhandlungen der Deutschen Physikalischen Gesellschaft, 21, 13–24, 1919 — paraphrased).

---

A-Level misconceptions

The errors that distinguish A from A*:

1. Lattice enthalpy is always negative under OCR. Some textbooks use a positive (dissociation) convention. Always check the source's convention before substituting a value into an OCR Born–Haber cycle.
2. Lattice enthalpy refers to one mole of compound, not one mole of ions. For CaCl$_2$2​ the lattice enthalpy is $-2258$−2258 kJ mol$^{-1}$−1 per mole of CaCl$_2$2​ — not per mole of Cl$^-$− ions.
3. The starting state for lattice enthalpy is gaseous ions, not elements. The journey from elements in standard states to gaseous ions involves atomisation, ionisation, and electron capture — the rest of the Born–Haber cycle.
4. Lattice enthalpy is not the same as lattice energy. Strictly, lattice energy refers to $\Delta U$ΔU (internal energy change), and lattice enthalpy refers to $\Delta H = \Delta U + \Delta(pV)$ΔH=ΔU+Δ(pV). At ordinary temperatures the difference $\Delta(pV) = -\frac{1}{2}RT$Δ(pV)=−21​RT (typical, in kJ mol$^{-1}$−1 a small correction) is negligible compared with the magnitude of the lattice term, and A-Level treats them as interchangeable — but the distinction is real at the undergraduate level.
5. More negative lattice enthalpy means stronger ionic bonding, not weaker. A common slip is to confuse "larger numerical value" (which could mean either +2000 or $-$−2000) with "more exothermic" — always reason in terms of sign and magnitude separately.
6. The OCR formation convention treats the gaseous ions as the high-energy reference state, the lattice as the low-energy product. Reversing this is the single most common conceptual error in the topic.

---

Summary

OCR lattice enthalpy is the exothermic enthalpy change accompanying the formation of one mole of an ionic compound from its constituent gaseous ions under standard conditions. It is always negative under the OCR formation convention, and its magnitude is a direct measure of ionic bonding strength. Because gaseous ions are not an experimentally accessible starting material, lattice enthalpy cannot be measured directly and must be determined indirectly through a Born–Haber cycle — the topic of Lesson 2. The magnitude of $\Delta_{LE} H^\ominus$ΔLE​H⊖ correlates predictably with ionic charge and radius (Lesson 4), and it ultimately governs the macroscopic physical properties (melting point, hardness, solubility, conductivity) of ionic solids.

---

Reference: OCR A-Level Chemistry A (H432) Module 5.2.1 (refer to the official OCR H432 specification document for exact wording).