Ionic Radii

A cation forms when an atom loses electrons, reducing the total electron count while the number of protons remains constant. The remaining core electrons experience greater effective nuclear charge and are pulled closer to the nucleus. For example, lithium's ionic radius is 60 picometers compared to its parent atom's 152 picometers due to loss of the outermost 2s electron.

When an atom gains an electron to form an anion, increased electron-electron repulsion causes electrons to spread out more in space. Additionally, the effective nuclear charge per electron decreases. These effects combine to make anions larger than their parent atoms. For instance, fluoride anion has a radius of 136 picometers versus fluorine's smaller atomic radius.

An isoelectronic series consists of atoms and ions with identical electron configurations but different numbers of protons. Within such a series, the number of protons determines size: greater nuclear charge pulls electrons closer, resulting in smaller radius. For example, sulfide ion with 16 protons is larger than calcium ion with 20 protons, despite both having 18 electrons.

Ionic radii increase down a group because successive elements have additional principal energy levels and orbitals. Both cations and anions follow this trend: as the principal quantum number increases, electrons occupy shells farther from the nucleus, resulting in larger ionic radii for elements lower in the periodic table.

Ionic radius depends on three key factors: the number of electrons, the orbitals holding valence electrons, and the nuclear charge. Ions with higher positive charges are smaller than those with lower charges because more protons pull electrons closer. The effective nuclear charge experienced by electrons is the primary driver of ionic size differences.

Cations with larger charges have more protons attracting the same number of electrons, creating greater effective nuclear charge. This stronger attraction pulls electrons closer to the nucleus, making highly charged cations significantly smaller. For example, V3+ with three positive charges is smaller than V2+ with two positive charges.

Electron shielding occurs when inner electrons reduce the effective nuclear charge experienced by outer electrons. When an atom loses electrons to form a cation, fewer shielding electrons remain, so the remaining electrons experience stronger nuclear attraction and move closer to the nucleus, dramatically reducing ionic radius compared to the parent atom.