S2.1
The ionic model
4 SL (SL 2.1.1-2.1.3)
Unverified
← View on the mapParent topic: Models of bonding and structure
Guiding question
Guiding questionWhat determines the ionic nature and properties of a compound?
The one big idea
An ionic bond is electrostatic attraction between oppositely charged ions arranged in a 3-D lattice. The bond has no preferred partner — each ion is surrounded by all its oppositely charged neighbours.
What this rests on
Our editorial judgement.The dependencies in this section are this site’s own assessment of what a sub-topic rests on. They have notbeen verified by a chemistry teacher and are not part of the official IB guide. By contrast, section 8 (Linking questions) reproduces the IB’s own wording verbatim.
- EditorialS1.3 Electron configurationsPredicting ionic charge is done from electron configuration (explicit skill in 2.1.1).
What rests on this
- EditorialS2.2 The covalent model
- EditorialS2.4 From models to materials
- EditorialR1.2 Energy cycles in reactions
- EditorialR3.1 Proton transfer reactions
- EditorialR3.2 Electron transfer reactions
Core concepts that must be mastered
- Ion formation from electron configurations(2.1.1)Metals lose electrons to form cations; non-metals gain electrons to form anions. You predict the ionic charge from the atom's electron configuration: a group 1 atom (ns¹) loses one electron to achieve a noble-gas core, a group 17 atom (ns²np⁵) gains one. The charge is chosen to reach the nearest stable configuration — but this is a pattern, not a cause. The bond is electrostatic, and the "stable configuration" is a consequence of the energy landscape, not a driving force. The idea that "atoms want a full outer shell" is a GCSE simplification that the IB retires: the full shell is a result, not a reason. Transition elements can form ions with different charges (Fe²⁺ and Fe³⁺, Cu⁺ and Cu²⁺) because their 3d and 4s sublevels are close in energy, so different numbers of electrons can be removed without crossing into a much deeper core shell. You must be able to predict ionic charge from electron configuration for main-group elements and recognise that transition elements can have variable charges. Common ionic charges follow group number for main-group elements: group 1 → +1, group 2 → +2, group 13 → +3, group 15 → −3, group 16 → −2, group 17 → −1.
- The ionic bond and lattice structure(2.1.2)The ionic bond is electrostatic attraction between oppositely charged ions. In a solid ionic compound, ions are arranged in a 3-D lattice — not as discrete molecule-like pairs. In the rock-salt (NaCl) structure, each Na⁺ is surrounded by six Cl⁻ ions, and each Cl⁻ by six Na⁺. There is no privileged partner and no memory of which electron came from which atom: NaCl is an empirical formula, not a molecular one. You must be able to deduce the formula and name of an ionic compound from its component ions, including polyatomic ions. The polyatomic ions you must know by name and formula are: NH₄⁺ (ammonium), OH⁻ (hydroxide), NO₃⁻ (nitrate), HCO₃⁻ (hydrogencarbonate), CO₃²⁻ (carbonate), SO₄²⁻ (sulfate), PO₄³⁻ (phosphate). In binary ionic compounds, the cation is named first and the anion takes the "-ide" suffix (NaCl = sodium chloride, MgO = magnesium oxide). For compounds with polyatomic ions, the name of the polyatomic ion is used as-is (NaNO₃ = sodium nitrate, CaCO₃ = calcium carbonate).
- Physical properties and lattice enthalpy(2.1.3)Ionic compounds form 3-D lattices represented by empirical formulas. Their physical properties follow from the lattice: high melting and boiling points (strong electrostatic forces must be overcome throughout the lattice — every ion is bonded to all its neighbours, not just one), electrical conductivity only when molten or dissolved (in the solid, ions are locked in place and cannot move; when molten or in solution, ions are free to move and carry charge), brittleness (if the lattice layers are displaced, ions of like charge are brought next to each other and repel, causing the crystal to cleave), and solubility in polar solvents (water stabilises the separated ions through ion–dipole interactions — the δ+ end of water attracts anions, the δ− end attracts cations). Lattice enthalpy is introduced as a measure of ionic bond strength, influenced by ion radius and charge: smaller ions and higher charges give stronger electrostatic attraction and higher lattice enthalpy. This is a qualitative treatment here; the quantitative use of lattice enthalpy in Born–Haber cycles comes at HL in Reactivity 1.2. The volatility of ionic compounds is low — they do not evaporate easily because the lattice energy must be supplied to separate the ions.
- How this sub-topic connectsThe ionic model is one of three bonding models (alongside covalent S2.2 and metallic S2.3) that Structure 2.4 unifies into a continuum. The electron-transfer language connects directly to redox (R3.2): forming an ionic compound from its elements is a redox reaction — the metal is oxidised, the non-metal is reduced. Lattice enthalpy returns in HL thermodynamics (R1.2, Born–Haber cycles, where you interpret the cycle to extract a single value such as lattice enthalpy or electron affinity). The polyatomic ions learned here reappear in acid–base chemistry (R3.1: CO₃²⁻ and HCO₃⁻ are bases that react with acids to release CO₂; SO₄²⁻ is the conjugate base of HSO₄⁻). Bonding is not a binary — the continuum model (S2.4) shows that most compounds have a mixture of ionic and covalent character, and the position in the bonding triangle determines the dominant properties.
Quantitative non-negotiables
No quantitative non-negotiables specified.
Common failure modes
M-08 — Atoms bond because they want a full outer shell.
Confidence: verified
Why it’s wrong: Circular and causally backwards. This is Taber's 'full outer shells explanatory principle'.
Correction: Bonding is electrostatic attraction; full-shell configurations correlate with low energy, they don't cause bonding.
M-09 — Sodium gives its electron to chlorine, so Na⁺ and Cl⁻ are bonded to each other.
Confidence: verified
Why it’s wrong: The molecular framework for ionic bonding — treating NaCl as discrete ion pairs.
Correction: An ionic solid is a 3D lattice in which each Na⁺ is attracted to all surrounding Cl⁻ (six nearest neighbours in rock salt); there is no privileged partner and no memory of which electron came from where. NaCl is an empirical formula, not a molecular one.
M-10 — Electrons remember which atom they came from.
Confidence: verified
Why it’s wrong: The 'ownership of electrons' conception.
Correction: Electrons are indistinguishable; once transferred or shared, the electronic history is chemically irrelevant.
M-11 — A compound is either ionic or covalent.
Confidence: verified
Why it’s wrong: Bonding is not a binary classification.
Correction: Bonding is a continuum; use Δχ and χ̄ and the bonding triangle. Most bonds are polar covalent to some degree.
What "HL standard" actually looks like
No HL extension. The ionic model is shared SL and HL. Lattice enthalpy is introduced qualitatively as a measure of bond strength; its quantitative treatment comes at HL in Reactivity 1.2 (Born–Haber cycles).
Linking questions
Official IB.The questions in this section are the IB’s own linking questions, reproduced verbatim from the guide. They are not this site’s editorial judgement — see section 3 for that distinction.
- Official IB
Why are some changes of state endothermic and some exothermic?
- Official IB
How does the position of an element in the periodic table relate to the charge of its ion(s)?
- Official IB
How does the trend in successive ionization energies of transition elements explain their variable oxidation states?
- Official IB
Why is the formation of an ionic compound from its elements a redox reaction?
- Official IB
How is formal charge used to predict the preferred structure of sulfate?
- Official IB
Polyatomic anions are conjugate bases of common acids. What is the relationship between their stability and the conjugate acid's dissociation constant, Ka?
- Official IB
How can lattice enthalpies and the bonding continuum explain the trend in melting points of metal chlorides across period 3?
- Official IB
Why do ionic bonds only form between different elements while covalent bonds can form between atoms of the same element?
- Official IB
What properties of ionic compounds might be expected in compounds with polar covalent bonding?
- Official IB
What are the limitations of discrete bonding categories?
- Official IB
Why do composites like reinforced concretes, which are made from ionic and covalently bonded components and steel bars, have unique properties?
- Official IB
How do differences in bonding explain the differences in the properties of metal and non-metal oxides?
- Official IB
What are the factors that influence the strength of lattice enthalpy in an ionic compound?
- Official IB
What are the conjugate acids of the polyatomic anions listed in Structure 2.1?
- Official IB
Under what conditions can ionic compounds act as electrolytes?
- Official IB
How is it possible for a radical to be an atom, a molecule, a cation or an anion? Consider examples of each type.
1 further official linking question target a Tool, Inquiry or Nature of Science strand and are not drawn as edges on the map:
Editorial The decision to surface these off-graph questions here is our editorial judgement — the IB does not prescribe where they should appear.
- Official IB
What experimental data demonstrate the physical properties of ionic compounds?
Bridge: GCSE → IB HL
Derived. These are the bridge items tagged to this sub-topic— places where the GCSE model gets redefined, genuinely new territory, or carry-over strengths. The tagging is this site’s editorial judgement.