A cation is an atom or group of atoms that has lost one or more electrons, giving it a net positive charge; in AP Chemistry, metals typically form cations with charges you can predict from their position on the periodic table (EK 1.8.A.3).
A cation is an ion with a positive charge. It forms when an atom (almost always a metal) loses one or more valence electrons. Sodium loses one electron to become Na⁺, magnesium loses two to become Mg²⁺, aluminum loses three to become Al³⁺. Notice the pattern. The charge matches the number of valence electrons the atom gives up, and that number is predicted by the element's column on the periodic table (EK 1.8.A.3). Group 1 makes +1 cations, Group 2 makes +2, and so on.
Why do metals lose electrons in the first place? Metals on the left side of the periodic table have low ionization energies and low electronegativities, so their valence electrons are weakly held. When a metal meets a nonmetal with high electronegativity, the electrons transfer instead of being shared, producing a cation and an anion held together by electrostatic attraction. That attraction, described by Coulomb's law, is the whole basis of ionic bonding in Unit 2. One more thing worth memorizing now: a cation is always smaller than its parent atom, because it lost an entire shell of electrons (or at least shed electron-electron repulsion) while keeping the same nuclear charge pulling everything inward.
Cations sit at the intersection of Unit 1 (Atomic Structure and Properties) and Unit 2 (Compound Structure and Properties). In Topic 1.8, LO 1.8.A asks you to connect periodic trends to reactivity, and EK 1.8.A.3 says typical ionic charges are governed by valence electrons and predictable from periodic table position. In Topic 2.1, LO 2.1.A has you use electronegativity differences to explain why a metal-nonmetal pair forms ions rather than sharing electrons. In Topic 2.3, LO 2.3.A requires you to represent ionic solids as a 3-D array of cations and anions arranged to maximize attraction and minimize repulsion (EK 2.3.A.1). If you can't identify which particle is the cation, what its charge is, and how big it is, you can't reason through Coulomb's law comparisons of lattice energy, melting point, or solubility. Those comparisons show up constantly on the exam.
Keep studying AP Chemistry Unit 2
Coulomb's Law (Unit 2)
Coulomb's law is the math behind why cations matter. Attraction between a cation and anion grows with bigger charges and shrinks with bigger distance. That's why MgO (with 2+ and 2- ions) melts far hotter than NaCl (1+ and 1-). Every 'which ionic compound has the stronger attraction' question is really asking you to compare cation charge and size.
Ionization Energy (Unit 1)
Ionization energy is literally the cost of making a cation. It's the energy needed to remove an electron. Metals form cations because their ionization energies are low, and the big jump in successive ionization energies tells you when an atom stops giving up electrons (which is why Mg forms Mg²⁺ but never Mg³⁺).
Ionic Radius (Unit 1)
Cations are always smaller than their neutral parent atoms. Na⁺ drops a whole electron shell compared to Na, and the remaining electrons get pulled in tighter by the unchanged nuclear charge. Smaller cation means shorter interionic distance, which means stronger Coulombic attraction in the solid.
Electrolyte (Unit 4)
When an ionic solid dissolves in water, its cations and anions break free and move independently, which is what lets the solution conduct electricity. Free-floating cations like Na⁺ and Ca²⁺ are the mobile positive charge carriers in any electrolyte solution.
You won't get a question that just says 'define cation.' Instead, cations are the working parts of Coulomb's law reasoning. Multiple-choice questions give you pairs of ionic compounds (like CaF₂ vs. NaCl) and ask you to explain differences in structure, melting point, or lattice energy using ion charge and ionic radius. Others test whether you can connect ionization energy and electron affinity data to which element forms a cation versus an anion. On FRQs, expect to justify a property comparison in writing. The winning answer names the specific ions, states their charges and relative sizes, and ties it back to Coulombic attraction. One thing you do NOT need is memorized crystal structures. The CED's exclusion statement for 2.3.A says specific crystal geometries won't be assessed, only the general idea that cations and anions arrange to maximize attraction and minimize repulsion.
A cation is positive because it LOST electrons; an anion is negative because it GAINED them. Metals (low ionization energy, left side of the table) form cations, while nonmetals (high electronegativity, right side) form anions. A memory trick that works: 'cat-ion' has a 't' that looks like a plus sign, and 'anion' = 'A Negative ION.' Also remember the size flip. Cations shrink relative to their parent atoms, anions swell.
A cation forms when an atom loses electrons, and its positive charge equals the number of electrons lost, which you can predict from the element's group on the periodic table (EK 1.8.A.3).
Metals form cations because they have low ionization energies and low electronegativities, so their valence electrons transfer easily to nonmetals.
A cation is always smaller than its parent atom because it loses electrons while keeping the same nuclear charge, and that smaller radius means stronger Coulombic attraction in ionic solids.
In an ionic crystal, cations and anions sit in a repeating 3-D array that maximizes cation-anion attraction and minimizes like-charge repulsion (EK 2.3.A.1), but you don't need to memorize specific crystal structures for the exam.
When comparing properties like melting point or lattice energy, argue from cation charge and cation size using Coulomb's law, since higher charge and smaller radius both mean stronger attraction.
A cation is an atom or group of atoms that has lost one or more electrons, giving it a positive charge. Examples include Na⁺ (lost 1 electron), Ca²⁺ (lost 2), and the polyatomic ion NH₄⁺.
A cation is positive because it lost electrons (typically a metal like Na⁺ or Mg²⁺), while an anion is negative because it gained electrons (typically a nonmetal like Cl⁻ or O²⁻). In an ionic compound, cations and anions attract each other through electrostatic forces.
Smaller, always. Na⁺ is much smaller than Na because it loses its entire outer shell, and the remaining 10 electrons feel the full pull of 11 protons. This size difference matters for Coulomb's law comparisons of lattice energy and melting point.
Almost never in AP Chem. Nonmetals have high electronegativities and high ionization energies, so they gain electrons and form anions instead. The main exception you'll see is the polyatomic ammonium ion, NH₄⁺, which is a cation made of nonmetal atoms.
Use the periodic table. Group 1 metals form +1 cations, Group 2 form +2, and aluminum forms +3, because each loses all its valence electrons to reach a stable noble-gas configuration. EK 1.8.A.3 says exactly this: typical ionic charges are governed by valence electrons and predicted by location on the table.