AP Chemistry · Unit 1 · Topic 1.8
Valence Electrons and Ionic Compounds
The formula of an ionic solid is a charge-balance statement. It tells you the ratio of ions in a bulk array, not the membership of a tiny isolated molecule. Electron configurations help predict common charges, but the stability of the solid depends on the energy of the whole process.
Identify the electrons that change
Main-group valence electrons occupy the outermost principal shell. Many common main-group cations result from losing those electrons, while common anions often complete the valence shell by gaining electrons. Magnesium commonly forms Mg2+; chlorine commonly forms Cl−. These patterns are useful predictions, not a claim that every element must form a monatomic ion with a noble-gas configuration.
Neutrality determines the ratio
A macroscopic ionic compound has zero net charge. Choose the smallest whole-number counts that make the total positive and negative charges cancel.
Subscripts count ions or atoms; superscripts indicate charge. Altering an ion's charge to make arithmetic easier changes the chemical species and is not a valid way to write a formula.
Common combinations, different ratios
| Cation | Anion | Smallest neutral formula | Charge check |
|---|---|---|---|
| K⁺ | O²− | K₂O | 2(+1) − 2 = 0 |
| Mg²⁺ | N³− | Mg₃N₂ | 3(+2) + 2(−3) = 0 |
| Al³⁺ | SO₄²− | Al₂(SO₄)₃ | 2(+3) + 3(−2) = 0 |
Parentheses preserve a polyatomic ion as a group. Three sulfate ions contain three sulfur atoms and twelve oxygen atoms, while the internal sulfate bonds are covalent.
Worked example: magnesium nitride
Magnesium starts with [Ne]3s², so Mg2+ has the [Ne] electron configuration. Nitrogen starts with [He]2s²2p³, so N3− also has ten electrons. Three magnesium ions provide +6 total charge and two nitride ions provide −6, giving Mg3N2. Equal final electron counts do not imply equal ion charges because the nuclei have different proton counts.
Where the energetic benefit comes from
Removing electrons from isolated gas-phase metal atoms costs energy. Ionic-solid formation can nevertheless be favorable because other steps, especially bringing opposite charges together throughout a lattice, release energy. The phrase “electron transfer makes both atoms stable” misses this energy accounting. Neither ion should be treated as existing alone in the solid without interactions with many neighbors.
Evidence from conductivity
A typical ionic crystal does not conduct well as a solid because its ions are fixed in position. Melting it, or dissolving a soluble sample in water, allows charged particles to move. A comparison of solid and dissolved conductivity can support the ionic model. However, low solution conductivity could also reflect poor solubility, so include a solubility observation before rejecting the model.
Know where simple predictions stop
Transition metals can have several common oxidation states. Specify the ion, for example iron(II) versus iron(III), before calculating a formula. Noble-gas similarity is not a universal law of ionic charge. Also distinguish nitride, N3−, from nitrate, NO3−; similar names do not imply similar charge or composition.
Practice, then check your reasoning
What formula combines Ca²⁺ with PO₄³⁻?
Answer and reasoning: Ca3(PO4)2; +6 and −6 cancel.
Why is MgO written rather than Mg₂O₂?
Answer and reasoning: The bulk ion ratio is reduced to the smallest integers, one Mg²⁺ for each O²⁻.
Scientific references: OpenStax: ionic bonding; OpenStax: lattice-energy reasoning. OpenStax Chemistry 2e authors: Paul Flowers, Klaus Theopold, Richard Langley, and William R. Robinson. Consult source pages for their license terms. Accessed September 13, 2026. Explanations, comparison tables, and practice tasks were independently written for this site; numerical teaching datasets are illustrative unless explicitly identified otherwise.
Official curriculum reference: College Board AP Chemistry course page. The explanation and worked example are independently written for this study site.