An ionic bond is the electrostatic attraction between oppositely charged ions (a cation and an anion), with strength predicted by Coulomb's law, so larger charges and smaller ions make stronger bonds, higher lattice energies, and higher melting points.
An ionic bond forms when a metal transfers one or more valence electrons to a nonmetal, creating a positive cation and a negative anion that stick together through electrostatic attraction. The charges each ion takes on aren't random. They're governed by valence electrons and predictable from position on the periodic table (EK 1.8.A.3), which is why Na is always +1 and O is almost always -2 in ionic compounds.
Here's the part AP Chem actually cares about. Ionic bond strength is a straight application of Coulomb's law. Bigger charges and smaller distance between ion centers mean a stronger attraction. That one relationship explains why MgO (charges of +2 and -2) melts way hotter than NaCl (+1 and -1), and why compounds with small ions out-bond compounds with big floppy ones. Ionic compounds don't form little molecules, either. They build a repeating 3D crystal lattice where every cation is surrounded by anions and vice versa, which is why we talk about lattice energy instead of a single bond energy.
Ionic bonds show up in four different units, which makes them one of the most cross-tested concepts in the course. In Topic 1.8 (LO 1.8.A), you predict ionic charges from valence electrons and periodicity. In Topic 2.2 (LO 2.2.A), you use Coulomb's law and potential energy vs. internuclear distance graphs to rank bond strengths and lattice energies. In Topic 3.1 (LO 3.1.A), you have to keep ionic bonds separate from intermolecular forces and recognize ion-dipole interactions when salts dissolve. And in Topic 4.4 (LO 4.4.A), the dissolution of an ionic salt is the textbook gray area between physical and chemical change, because ionic bonds break while ion-dipole interactions form. If you can run the Coulomb's law logic (charge up, distance down, attraction up), you can answer questions in all four places.
Keep studying AP Chemistry Unit 2
Coulomb's Law (Unit 2)
Coulomb's law is the engine behind every ionic bond question. Attraction scales with the product of the charges and drops off with distance, so a +2/-2 pair of small ions beats a +1/-1 pair of big ones every time. When the exam asks you to rank melting points or lattice energies, it's really asking you to apply this equation in words.
Crystal Lattice (Unit 2)
Ionic compounds aren't molecules. NaCl is a giant repeating lattice of alternating Na⁺ and Cl⁻ ions, not a pair of atoms holding hands. That lattice structure is why ionic solids have high melting points, are brittle, and only conduct electricity when melted or dissolved (when the ions are free to move).
Valence Electrons and Periodicity (Unit 1)
You can predict the formula of an ionic compound before you ever calculate anything. Group 1 metals lose one electron, Group 2 lose two, halogens gain one, and elements in the same column form analogous compounds (EK 1.8.A.2). The periodic table is your cheat sheet for ionic charges.
Dissolution of Salts (Units 3-4)
When NaCl dissolves in water, ionic bonds break and ion-dipole interactions form between the ions and water molecules. The CED explicitly says you could argue this is either a physical or chemical process (EK 4.4.A.2), and that argument itself is fair game on the exam.
Ionic bonds are mostly tested as reasoning, not recall. A classic MCQ stem hands you melting point or lattice energy data and asks which Coulomb's law factor explains it. For example, you might explain why BaO has a higher lattice energy than KCl even though Ba²⁺ is bigger than K⁺ (the 2+/2- charges in BaO dominate over the size difference). Other questions test whether you can predict relative melting points from ion charge and radius, or explain solubility differences like CaSO₄ vs. CaCl₂ using attraction strength between ions. On FRQs, expect to write a sentence-level justification that names both factors, charge magnitude and ionic radius, and ties them to the property being asked about. A bare answer like "stronger bonds" won't earn the point; "greater charges and smaller internuclear distance produce stronger Coulombic attraction" will.
Ionic bonds are interparticle attractions that hold a compound itself together, and they're far stronger than intermolecular forces like dipole-dipole or hydrogen bonding, which act between separate molecules. The trap shows up in Unit 3, where ranking boiling points means asking what kind of attraction you're breaking. Melting an ionic solid breaks ionic bonds (very hard); boiling water breaks hydrogen bonds, not covalent or ionic bonds (much easier). If you call an ionic bond an IMF on an FRQ, you lose the point.
An ionic bond is the electrostatic attraction between a cation and an anion, formed when a metal transfers valence electrons to a nonmetal.
Coulomb's law predicts ionic bond strength, so larger ion charges and smaller ionic radii mean stronger attraction, higher lattice energy, and a higher melting point.
Charge usually beats size in these comparisons, which is why MgO and BaO (2+/2- charges) have much higher lattice energies than NaCl or KCl (1+/1- charges).
Ionic compounds form crystal lattices rather than discrete molecules, which explains their high melting points and why they conduct electricity only when molten or dissolved.
Ionic charges are predictable from the periodic table, since elements in the same group have the same number of valence electrons and form analogous compounds.
Dissolving a salt in water breaks ionic bonds and forms ion-dipole interactions, and the CED says you can defend classifying it as either a physical or chemical process.
It's the electrostatic attraction between oppositely charged ions, formed when a metal transfers valence electrons to a nonmetal. Its strength follows Coulomb's law, so it increases with bigger ion charges and smaller ionic radii.
No. Ionic bonds are interparticle attractions within the compound itself and are much stronger than IMFs like dipole-dipole forces or hydrogen bonding. Mixing these up is one of the most common ways to lose points in Unit 3.
Ionic bonds involve electron transfer and the attraction between resulting ions, typically between a metal and a nonmetal. Covalent bonds involve shared electron pairs between nonmetals, and their strength depends on bond order and atom size rather than ion charges.
Coulomb's law. MgO has 2+ and 2- charges while NaCl has 1+ and 1-, so the attraction in MgO is roughly four times stronger from charge alone. Stronger attraction means more energy is needed to pull the lattice apart.
The CED (EK 4.4.A.2) says you can argue either way, and that's the point. Ionic bonds in the lattice break (sounds chemical) but no new compound forms, just ion-dipole interactions with water (sounds physical). On the exam, explain the bond changes rather than just picking a label.
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