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Ionic bonding is one of three primary bonding types — alongside covalent (next lesson) and metallic (lesson 4) — and is the model invoked whenever a metal combines with a non-metal. This lesson develops the electron-transfer mechanism that creates cations and anions, the dot-cross representation of that transfer, the giant ionic lattice ions assemble into in the solid state, and the physical properties (melting point, conductivity, brittleness, solubility) that fall out of those lattice forces. It closes with the experimental evidence — electrolysis, ion migration, and X-ray crystallography — that justifies the claim that ionic compounds really contain discrete ions, not neutral atoms. Ionic bonding is the gateway concept for §3.1.4 energetics (lattice enthalpy and Born-Haber, course 3) and for the Group 2 and Group 7 chemistry of §3.2.
Spec mapping (AQA 7405): This lesson maps to §3.1.3.1 (ionic bonding), the first bonding-type subsection in §3.1.3. The energetic deepening — lattice enthalpy and the Born-Haber cycle — is developed in §3.1.4 (course 3). Group 2 ion stability (Mg²⁺ → Ba²⁺) and Group 7 anion behaviour (F⁻ → I⁻) appear in §3.2.2 and §3.2.3; ionic-radius and charge trends developed here are reused there. Electron configuration of stable ions — the reason Na loses one electron but Mg loses two — was established in §3.1.1.2 (course 1, atomic structure). Refer to the official AQA specification document for the exact wording of each section.
Assessment objectives: Definitions of ionic bond, giant ionic lattice, and formula unit are AO1 recall. Drawing dot-cross diagrams and predicting formulae from group numbers (e.g. Al₂O₃ from Al³⁺ and O²⁻) are AO2 application. Linking lattice properties to ionic charge and radius — rationalising why MgO melts at 2852 °C but NaCl at 801 °C, or explaining the melting-point trend across Period 3 chlorides — is AO3 analysis and routinely appears as a 4–6 mark structured question.
Ionic bonding occurs between metals and non-metals. The metal atom loses one or more electrons to form a positive ion (a cation); the non-metal atom gains the same number of electrons to form a negative ion (an anion). The driving force is not, as is often stated, that atoms "want" a full outer shell — atoms have no desires — but that the overall energy change is exothermic. The energy released when the resulting cations and anions assemble into a three-dimensional lattice (the lattice enthalpy, course 3) more than pays for the ionisation energy of the metal and the (much smaller) electron affinity of the non-metal.
Key Definition: An ionic bond is the strong electrostatic attraction between oppositely charged ions, acting throughout a giant lattice rather than between any single pair of ions.
The group number gives a fast rule for likely ionic charges, because the s- and p-block elements achieve a noble-gas configuration by losing or gaining the smallest possible number of electrons:
| Group | Loses/Gains | Charge | Example | Configuration after |
|---|---|---|---|---|
| 1 | loses 1 | 1+ | Na⁺, K⁺ | [Ne], [Ar] |
| 2 | loses 2 | 2+ | Mg²⁺, Ca²⁺ | [Ne], [Ar] |
| 3 | loses 3 | 3+ | Al³⁺ | [Ne] |
| 5 | gains 3 | 3− | N³⁻ | [Ne] |
| 6 | gains 2 | 2− | O²⁻, S²⁻ | [Ne], [Ar] |
| 7 | gains 1 | 1− | F⁻, Cl⁻ | [Ne], [Ar] |
This rule is reliable for s-block metals and for non-metals up to Group 7. It fails for transition metals (variable oxidation states) and for the heavier p-block (Pb²⁺ rather than the predicted Pb⁴⁺, because of the inert-pair effect — A2 extension).
Sodium chloride (NaCl):
Magnesium oxide (MgO):
Calcium fluoride (CaF₂):
Aluminium oxide (Al₂O₃):
Exam Tip: When drawing dot-cross diagrams for ionic compounds, always (i) draw cation and anion separately, not overlapping; (ii) place square brackets around each ion; (iii) write the charge as a superscript outside the brackets; (iv) check that the total positive charge balances the total negative charge. Marks are routinely lost for omitting brackets or for drawing electrons "shared" between Mg and O as in covalent bonding.
Dot-cross diagrams record the outer-shell electrons of each ion after transfer. Electrons originating on one atom are drawn as dots (•); electrons originating on the other are drawn as crosses (×). Electrons are physically indistinguishable — the dot/cross convention is a pedagogical accounting device, not a claim about real differences.
The Na atom has 1 outer electron (drawn as •). It loses this electron, leaving Na⁺ with the neon-like core [Ne] — drawn with empty brackets or with the inner-shell dot indicating "no outer electrons remaining":
[Na]⁺ [×× ×× Cl ×× ×ו]⁻
The chloride ion in the diagram has 8 outer electrons: 7 originally on Cl (×) plus 1 transferred from Na (•). The total charge on each species is shown as a superscript outside the bracket.
Mg has 2 outer electrons; both are transferred to one O atom (which needs exactly 2 to reach octet). After transfer:
[Mg]²⁺ [×× ×× O ×× ••]²⁻
The oxide ion has 8 outer electrons (6 from O drawn ×, 2 from Mg drawn •). Charge balance: (+2) + (−2) = 0.
Ca has 2 outer electrons; each is transferred to a different Cl atom (each Cl needs only 1 to reach octet). The diagram shows one Ca²⁺ cation and two Cl⁻ anions, each chloride showing 7 of its own electrons (×) plus 1 of the calcium's (•):
[Ca]²⁺ [×× ×× Cl ×× ×ו]⁻ [×× ×× Cl ×× ×ו]⁻
Charge balance: (+2) + 2(−1) = 0. The 1:2 stoichiometry is explicit in the diagram.
Al has 3 outer electrons; O needs 2. The least-common-multiple of 3 and 2 is 6, so 2 Al atoms (donating 6 electrons total) pair with 3 O atoms (accepting 6 electrons total):
2 × [Al]³⁺ 3 × [O]²⁻ (each oxide showing 6× and 2•)
Charge balance: 2(+3) + 3(−2) = 0.
Two Na atoms each donate one electron to a single O atom:
2 × [Na]⁺ [O ×××ו•]²⁻
Charge balance: 2(+1) + (−2) = 0.
Pedagogical Note: Draw each ion separately with brackets and charges — never as overlapping circles (covalent convention). For compounds with multiple anions/cations (CaCl₂, Na₂O, Al₂O₃), drawing all ions makes the stoichiometry visible and protects against formula errors.
For an ionic compound to form, the overall energy balance must be favourable. The relevant terms are:
| Factor | Effect on ionic formation |
|---|---|
| Ionisation energy of the metal | Must be relatively low. Group 1 and Group 2 metals form cations easily; transition metals require more energy but the lattice enthalpy compensates |
| Electron affinity of the non-metal | First electron affinities are typically exothermic for Group 6 and Group 7 elements; second electron affinities (O⁻ → O²⁻) are endothermic, paid for by lattice enthalpy |
| Lattice enthalpy | Large negative value (energy released on lattice formation). Dominant term in the Born-Haber cycle — without it, ionic compound formation would not be feasible |
| Ionic charge product q⁺q⁻ | Larger product → more exothermic lattice enthalpy → more stable lattice. MgO (q⁺q⁻ = 4) far more stable than NaCl (q⁺q⁻ = 1) |
| Sum of ionic radii r⁺ + r⁻ | Smaller sum → more exothermic lattice enthalpy. LiF lattice enthalpy more negative than CsI |
The combined effect of charge and radius is captured in the proportionality:
Lattice enthalpy ∝ q⁺q⁻ / (r⁺ + r⁻)
This single relationship explains why MgO melts at 2852 °C whereas NaCl melts at only 801 °C: MgO has a fourfold-larger charge product (2 × 2 = 4 versus 1 × 1 = 1) and the Mg²⁺ ion is smaller than Na⁺.
Common Misconception: Students often write that ionic bonds form because atoms "want" a full outer shell. Atoms have no desires. The formation of an ionic compound is energetically favourable because the lattice enthalpy released exceeds the sum of the ionisation energies and the (sometimes endothermic) electron affinities — the Born-Haber cycle in course 3 of this catalogue makes the bookkeeping explicit.
Ionic compounds do not exist as discrete "NaCl molecules" — that is a covalent picture inappropriate to ionic bonding. Instead, every ionic compound in the solid state is a giant ionic lattice: a regular three-dimensional array in which each ion is surrounded by oppositely charged neighbours. The chemical formula (NaCl, MgO, CaF₂) is a formula unit — the simplest whole-number ratio of cations to anions — and not a description of any single molecular entity.
The NaCl lattice has a face-centred cubic arrangement:
The NaCl structure is adopted by most Group 1 halides (LiCl, NaCl, KCl, etc.), by silver halides (AgCl, AgBr), and by Group 2 oxides (MgO, CaO).
The CsCl lattice has a body-centred cubic arrangement:
The decision between 6:6 (rock-salt) and 8:8 (CsCl) packing is governed by the radius ratio r⁺/r⁻. When the cation is small relative to the anion (r⁺/r⁻ < 0.732), 6:6 packing is preferred; when the cation is larger (r⁺/r⁻ > 0.732), 8:8 packing becomes geometrically feasible.
Because there are twice as many anions as cations in CaF₂, the coordination numbers are unequal:
The fluorite structure is adopted by CaF₂, SrF₂, BaF₂, and (in inverted form) by Li₂O and Na₂O (the antifluorite structure).
The properties of ionic compounds fall out of the lattice-enthalpy framework: anything that requires breaking the lattice (melting, boiling, dissolving) is energetically demanding; anything that requires moving the ions while preserving the lattice (electrical conduction in the solid) is forbidden.
| Compound | Mᵣ | Melting pt / °C | Solubility in water | Lattice charges |
|---|---|---|---|---|
| NaCl | 58.5 | 801 | High (~360 g dm⁻³) | 1+ / 1− |
| KCl | 74.6 | 770 | High (~340 g dm⁻³) | 1+ / 1− |
| MgO | 40.3 | 2852 | Very low | 2+ / 2− |
| CaCO₃ | 100.1 | decomposes ~825 | Very low | 2+ / 2− (polyatomic anion) |
The pattern is consistent. MgO has the highest melting point because q⁺q⁻ = 4 and r⁺ + r⁻ is small. NaCl and KCl have similar melting points (charges identical; radii differ only modestly). CaCO₃ never melts at atmospheric pressure — it decomposes to CaO and CO₂ before the lattice can liquefy.
The energy required to melt an ionic compound is closely related to its lattice enthalpy. Because lattice enthalpy ∝ q⁺q⁻/(r⁺+r⁻), melting points rise sharply with ionic charge: MgO (2,2 charges) > CaO > NaCl (1,1 charges) > KCl > CsI. They rise more gently with decreasing ionic radius along a group.
A solid ionic compound has its ions locked in fixed lattice positions — they cannot migrate, so the solid does not conduct electricity. On melting, the lattice collapses and the ions become free to move; molten ionic compounds therefore conduct (this is exploited in the industrial extraction of aluminium from molten Al₂O₃). Aqueous solutions of soluble ionic compounds also conduct, because dissolution liberates the ions as hydrated cations and anions.
| State | Ions mobile? | Conducts electricity? |
|---|---|---|
| Solid | No (lattice positions fixed) | No |
| Molten | Yes (lattice destroyed) | Yes |
| Aqueous | Yes (hydrated ions) | Yes |
Ionic crystals are hard (the lattice resists small deformations) but brittle (they shatter under shock loading). The mechanism is the like-charge layer slip:
Before stress: + − + − + − − + − + − + + − + − + −
After a shear displacement of one ionic radius:
+ − + − + −
− + − + − + ← layer shifted
+ − + − + −
Now like charges face each other across the slip plane. The resulting electrostatic repulsion pushes the layers apart and the crystal cleaves. This is the reason an NaCl crystal can be split cleanly along its (100) crystallographic plane with a chisel — and the reason it shatters rather than deforming plastically.
Exam Tip: When explaining brittleness, always (i) name the layer-slip mechanism, (ii) state that ions of like charge come into alignment, (iii) state that the electrostatic repulsion is what shatters the crystal. "The bonds break" earns no marks.
Many ionic compounds dissolve in water but are insoluble in non-polar solvents (hexane, cyclohexane). Water dissolves NaCl through ion-dipole interactions: the δ⁻ oxygen of water orients toward Na⁺ (cation hydration); the δ⁺ hydrogens orient toward Cl⁻ (anion hydration). The total energy released on hydration (the hydration enthalpy) typically matches or exceeds the lattice enthalpy, so dissolution is exothermic or only modestly endothermic and is driven entropically.
Non-polar solvents cannot stabilise the separated ions — there is no dipole to align — so dissolution is endothermic and unfavourable. Ionic compounds are therefore essentially insoluble in hexane.
The exceptions — sparingly soluble ionic compounds such as AgCl, CaCO₃, BaSO₄ — have lattice enthalpies that exceed the available hydration energy. These exceptions are the basis of qualitative inorganic chemistry (precipitation tests in §3.2.4 and §3.2.5).
The claim that solid NaCl really contains discrete Na⁺ and Cl⁻ ions — rather than neutral Na and Cl atoms held together by some other force — rests on three independent lines of experimental evidence.
When molten ionic compounds are subjected to a direct-current potential, metal is deposited at the cathode (the negative electrode) and the non-metal is liberated at the anode (the positive electrode). For molten NaCl: Na metal at the cathode, Cl₂ gas at the anode. This demonstrates that the molten compound contains positive particles that move toward the cathode and negative particles that move toward the anode — i.e., it contains ions. Sir Humphry Davy first isolated sodium and potassium by electrolysing molten sodium and potassium hydroxide in 1807, providing the original experimental evidence for the ionic model.
A more direct demonstration uses ions that are themselves coloured. A strip of moist filter paper is laid across a flat surface and the ends connected to a d.c. supply. A small crystal of copper(II) sulfate (blue Cu²⁺ ions) or potassium manganate(VII) (purple MnO₄⁻ ions) is placed in the centre. Over minutes to hours, the colour visibly migrates: Cu²⁺ moves toward the cathode; MnO₄⁻ moves toward the anode. This direct observation of charge-bearing coloured species is hard to explain without ions.
X-ray diffraction yields the three-dimensional electron density of the crystal at near-atomic resolution. For NaCl, the electron density map shows two distinct features: high-density spherical regions centred on each lattice site (the ions), with very low electron density in the regions between sites. There is no bridge of electron density linking adjacent sites — in contrast to a covalent solid like diamond, where the electron density forms a bonding ridge between adjacent carbon atoms. The measured cation and anion radii (Na⁺ ≈ 95 pm, Cl⁻ ≈ 181 pm) agree with the radii derived from gas-phase ionisation experiments. This structural agreement — and the absence of any inter-ion electron-density bridge — is the strongest evidence that the ionic model is correct (and signposts forward to the structure-determination techniques developed in §3.3.16 chromatography and instrumental methods).
Independent of structure determination, the ionic-radius pattern is in itself indirect evidence for ion formation:
| Species | Radius / pm | Explanation |
|---|---|---|
| Na | 186 | Neutral atom |
| Na⁺ | 95 | Lost outer 3s electron; increased effective nuclear charge contracts remaining shells |
| Cl | 99 | Neutral atom (covalent radius) |
| Cl⁻ | 181 | Gained one electron; increased electron-electron repulsion expands the n=3 shell |
| Mg | 160 | Neutral atom |
| Mg²⁺ | 65 | Lost both 3s electrons; large increase in Z_eff |
| O | 73 | Neutral atom (covalent radius) |
| O²⁻ | 140 | Gained two electrons; strong electron-electron repulsion |
Cations are always smaller than the parent atom; anions are always larger. The pattern is precisely what one expects if electrons have been transferred — and not what one would expect from any "neutral atom" model of the solid.
Ionic bonding sits at the centre of a web of A-Level chemistry topics; the following are the most important connections.
Question 1. [12 marks total]
(a) Define an ionic bond and state one physical property of ionic compounds that is a direct consequence of this type of bonding. [2 marks]
(b) Draw a dot-cross diagram for magnesium nitride, Mg₃N₂. Show all outer-shell electrons, brackets, and charges. [3 marks]
(c) Magnesium oxide has a melting point of 2852 °C, whereas sodium chloride has a melting point of 801 °C. Explain this difference by reference to ionic charges, ionic radii, and the lattice forces involved. [4 marks]
(d) Solid sodium chloride is composed of discrete Na⁺ and Cl⁻ ions, not neutral Na and Cl atoms. Evaluate two pieces of experimental evidence that support this claim. [3 marks]
(a) Definition of ionic bond and one consequent property [2 marks, AO1]
Common error: defining ionic bonding as "transfer of electrons" — that is the mechanism of formation, not the bond itself. The bond is the electrostatic attraction after transfer.
(b) Dot-cross diagram for Mg₃N₂ [3 marks, AO2]
Common error: drawing one Mg²⁺ and one N³⁻ only, ignoring stoichiometry; or putting all 8 anion electrons as × (failing the dot-cross convention).
(c) MgO vs NaCl melting-point difference [4 marks, AO3]
Common error: invoking only the charge effect and ignoring radius; or invoking "stronger bonds" without identifying the underlying lattice-enthalpy relationship.
(d) Evidence for discrete ions in NaCl [3 marks, AO3]
Common error: giving two examples of the same type of experiment (two electrolysis variants) rather than two independent lines of evidence; or describing evidence without evaluating its strength.
The three responses below cover the meaningful A-Level range: Grade C (borderline-pass), Grade B (solid mark-scheme coverage), and Grade A* (top-band synthesis). Commentary after each response (editorial, not a real examiner report) names the marks earned and the specific moves that differentiate adjacent bands.
(a) An ionic bond is the electrostatic attraction between positive and negative ions. One consequent property is that ionic compounds have high melting points because a lot of energy is needed to overcome these strong attractions.
(b) Mg₃N₂ dot-cross: three [Mg]²⁺ ions drawn with empty outer shells, and two [N]³⁻ ions, each with 8 outer electrons in square brackets (5 × from N and 3 • from Mg). Charges: 3(+2) + 2(−3) = 0.
(c) MgO has a higher melting point than NaCl because the ions have higher charges. Mg²⁺ and O²⁻ carry charges of 2+ and 2− respectively, whereas Na⁺ and Cl⁻ carry only 1+ and 1−. The product q⁺q⁻ is 4 in MgO but only 1 in NaCl. The ions in MgO are also smaller (Mg²⁺ is 65 pm, O²⁻ is 140 pm) than in NaCl (Na⁺ 95 pm, Cl⁻ 181 pm). Both factors make MgO's lattice enthalpy more exothermic, so more energy is needed to break the lattice and melt the solid.
(d) Two pieces of evidence: (i) electrolysis of molten NaCl produces sodium metal at the cathode and chlorine gas at the anode, showing that the molten compound contains positive and negative charged particles. (ii) X-ray diffraction studies of solid NaCl show electron density localised in spherical regions around each lattice site, consistent with discrete ions rather than neutral atoms.
Editorial commentary (Grade C): All four parts are correct and would secure the mark-scheme points. Part (c) names both the charge and radius effects, which is enough for full marks. To progress to B, the response would tie the charge/radius observation to the explicit proportionality lattice enthalpy ∝ q⁺q⁻/(r⁺+r⁻) and would evaluate the relative strength of the two pieces of evidence in (d) rather than merely listing them.
(a) An ionic bond is the strong electrostatic attraction between oppositely charged ions that operates throughout a giant ionic lattice rather than between any single pair of ions. A direct consequence is the high melting point of ionic compounds: the lattice contains many such attractions in three dimensions, so a great deal of thermal energy is required to overcome them and produce the mobile-ion liquid state.
(b) Mg₃N₂ dot-cross diagram: three magnesium ions drawn as [Mg]²⁺ with empty outer shells (the two 3s electrons have been transferred); two nitride ions drawn as [N]³⁻ in square brackets, each containing 8 outer-shell electrons — 5 × from nitrogen's own outer shell and 3 • from the three magnesiums that contributed to that nitride. Overall charge balance: 3 × (+2) + 2 × (−3) = +6 − 6 = 0.
(c) Lattice enthalpy is approximately proportional to q⁺q⁻ / (r⁺ + r⁻). In MgO the charges are 2+ and 2−, giving q⁺q⁻ = 4; in NaCl they are 1+ and 1−, giving q⁺q⁻ = 1. This is a factor of 4 in the numerator. The ionic radii are also smaller in MgO (Mg²⁺ 65 pm, O²⁻ 140 pm; sum 205 pm) than in NaCl (Na⁺ 95 pm, Cl⁻ 181 pm; sum 276 pm), making the denominator smaller in MgO. Both effects make the MgO lattice enthalpy substantially more exothermic than NaCl's, and the melting point — controlled by the energy required to disrupt the lattice — is correspondingly much higher.
(d) Electrolysis of molten NaCl yields Na metal at the cathode and Cl₂ gas at the anode, which directly demonstrates that the molten material contains mobile cations and anions. Independently, X-ray crystallography of solid NaCl shows spherical regions of electron density localised around each lattice site, with low density in between — a pattern characteristic of ions rather than neutral atoms or shared-pair covalent bonding.
Editorial commentary (Grade B): The response is now A-level-rigorous: part (a) distinguishes the bond from the mechanism of its formation; part (c) cites the explicit proportionality and gives numerical values; part (d) gives two independent pieces of evidence. To progress to A*, the response would link the lattice-enthalpy proportionality forward to the Born-Haber cycle as the quantitative deepening, would discuss why ionic bonding is rarely 100% ionic (Fajans's rules / electronegativity), and would evaluate the relative epistemic strength of electrolysis vs X-ray crystallography in part (d).
(a) An ionic bond is the strong electrostatic attraction between oppositely charged ions operating throughout a giant ionic lattice. This defines the bond itself (an attractive force), distinct from the mechanism of formation (electron transfer). A direct consequence is the high melting point: melting requires enough thermal energy to disrupt the lattice. For NaCl the enthalpy of fusion is ~28 kJ mol⁻¹ — small compared to the lattice enthalpy of −786 kJ mol⁻¹ because melting only liberates ion translational motion, not full atomisation.
(b) Mg₃N₂ dot-cross: three [Mg]²⁺ ions (each [Ne] core, no outer electrons shown) and two [N]³⁻ ions, each drawn in square brackets containing 8 outer-shell electrons — conventionally 5 × originating from nitrogen plus 3 • originating from the three transferring magnesiums. Charge balance: 3(+2) + 2(−3) = 0. The 3:2 stoichiometry is fixed by the requirement that the total electrons donated (6 from 3 Mg) equal the total electrons accepted (6 by 2 N).
(c) Lattice enthalpy ΔH_LE ∝ q⁺q⁻/(r⁺+r⁻). MgO: q⁺q⁻ = 4, r⁺+r⁻ ≈ 205 pm. NaCl: q⁺q⁻ = 1, r⁺+r⁻ ≈ 276 pm. Predicted |ΔH_LE| ratio ≈ 4 × (276/205) ≈ 5.4; experimental lattice enthalpies are −3791 (MgO) and −786 kJ mol⁻¹ (NaCl), ratio 4.8. Melting point tracks lattice enthalpy, so MgO (2852 °C) far exceeds NaCl (801 °C). The Born-Haber cycle (course 3) is the quantitative tool for extracting ΔH_LE from measurable thermochemical data.
(d) Three independent lines support the discrete-ion model. (i) Electrolysis of molten NaCl: Na at the cathode, Cl₂ at the anode — mobile charged species. (ii) Migration of coloured ions (Cu²⁺, MnO₄⁻) on moist filter paper under a d.c. potential — direct visual evidence. (iii) X-ray crystallography of solid NaCl: spherical electron-density features localised on lattice sites with no inter-site bonding ridge, contrasting sharply with the bridging electron density seen in covalent diamond. X-ray crystallography is epistemically the most direct because it images the electron distribution itself; electrolysis and migration demonstrate ion mobility but cannot, alone, rule out transient ion generation. The combination of three lines is therefore much stronger than any one.
Editorial commentary (Grade A):* Full synoptic command. Part (a) distinguishes bond from mechanism; (c) gives the lattice-enthalpy proportionality, quantifies it, compares to experimental data, and signposts Born-Haber in course 3 as the quantitative deepening; (d) gives three independent pieces of evidence and evaluates their relative epistemic weight. The connection of (c) to q⁺q⁻/(r⁺+r⁻) and forward to Born-Haber is the synoptic move that lifts the response into A*.
Three undergraduate-adjacent extensions:
This content is aligned with the AQA A-Level Chemistry (7405) specification.