The ionic radius measures an ion’s size, representing the distance from the nucleus to the outermost electron shell within a crystal lattice. Atoms become ions by gaining or losing electrons, a process that dramatically alters the balance between the positive nuclear charge and the negative electron cloud. This change in electron count causes a significant difference in the ion’s physical size compared to its neutral parent atom. Understanding the ionic radius is fundamental for predicting the properties and structures of ionic compounds, such as salts.
Defining Ionic Radii and Measurement
The ionic radius is distinct from the atomic radius because it is measured when the atom is charged and chemically bonded within a solid structure. Unlike a neutral atom, an ion’s size is not a fixed property but varies slightly depending on its neighbors in the crystal lattice, known as the coordination number. Ionic radii are very small, typically ranging from about 30 picometers (pm) to over 200 pm.
Measuring the radius of a single, isolated ion is impossible because the electron cloud lacks a sharp, defined boundary. The size is determined indirectly using X-ray crystallography on solid ionic compounds. This method involves firing X-rays at a crystalline sample and analyzing the resulting diffraction pattern to determine the precise distance between the nuclei of adjacent ions.
The measured distance between a positively charged ion (cation) and a negatively charged ion (anion) is the sum of their individual ionic radii. To determine the size of each ion separately, scientists partition this total distance based on established standards derived from carefully selected compounds. This methodology treats the ions as rigid spheres that are just touching, allowing for the calculation of their estimated radii.
How Ion Formation Changes Size (Cations and Anions)
Ion formation, whether by losing or gaining electrons, leads to a pronounced change in size relative to the neutral parent atom. This size change is driven by the alteration in electron-electron repulsion and the corresponding change in the effective nuclear charge. The effective nuclear charge is the net positive charge experienced by an electron, which is always less than the total nuclear charge due to the shielding effect of inner-shell electrons.
When a neutral atom loses electrons to form a positively charged cation, its radius dramatically shrinks. The loss of electrons decreases electron-electron repulsion among the remaining electrons. The fixed number of protons now pulls the fewer remaining electrons more strongly, increasing the effective nuclear charge. In many cases, the atom loses all the electrons in its outermost shell, resulting in the complete removal of a principal energy level, which causes a substantial decrease in size.
Conversely, when an atom gains electrons to form a negatively charged anion, its radius expands, becoming larger than the parent atom. The addition of electrons increases the total number of electrons without changing the number of protons in the nucleus. This results in greater electron-electron repulsion within the valence shell, pushing the electron cloud further outward. The increased repulsion effectively spreads the electrons out, making the ion physically larger.
Predicting Ion Size Using the Periodic Table
Ionic radii follow predictable patterns across the periodic table, though the trends must be considered separately for cations and anions. When comparing ions of the same charge, the size increases as you move down a group (column). This is because each step down adds a new principal electron shell, placing the outermost electrons progressively farther from the nucleus and increasing the ion’s overall size.
Moving across a period (row), the ionic radius decreases for ions of the same charge. This is a consequence of the increasing nuclear charge; as more protons are added from left to right, the stronger positive pull draws the electrons closer. For example, comparing Li+ to B3+ (both cations), the increased nuclear charge of boron causes its ion to be smaller than the lithium ion.
This relationship between nuclear charge and size is especially clear in an isoelectronic series, which is a group of ions and atoms that all possess the same number of electrons. For instance, the ions N3-, O2-, F-, Na+, and Mg2+ all have ten electrons. In such a series, the size is determined entirely by the number of protons, with the species having the highest nuclear charge (Mg2+) being the smallest, and the one with the fewest protons (N3-) being the largest.