8.3
Ionic Radii
Ionic radius is the measure used to describe the size of an ion. A cation always has fewer electrons and the same number of protons as the parent atom…
An ionic radius is the radius of a cation or an anion defined by the distance between ions in an ionic compound. Cations are smaller than the parent atom, whereas anions are larger. Similar to atomic radii, ionic radii are determined by the number of electrons, the orbitals holding its valence electrons, and the nuclear charge.
Consider lithium, which has an electron configuration of a helium core and one outermost 2s electron. The 2s electron is shielded from the nuclear charge by two 1s electrons and contributes to the atomic radius of 152 picometers.
The loss of the outermost 2s electron generates a lithium cation, which has fewer electrons but the same number of protons as the parent atom. The two 1s electrons are held closer to the nucleus because they experience a greater effective nuclear charge than the 2s electron did. Thus, the ionic radius of the lithium cation is 60 picometers, which is much smaller than the parent atom.
This trend is generally observed with all metal cations and their parent atoms.
In contrast, anions are larger than their parent atoms. When a fluorine atom accepts an electron, there is an additional outermost electron, but the number of protons, and hence the nuclear charge, remains the same. The increased electron – electron repulsion causes the electrons to spread out more in space. Thus, the fluoride anion has a radius of 136 picometers, which is much larger than the parent atom.
In general, the ionic radius for s and p-block elements increases down the column, as the number of principal energy levels, and hence the number of orbitals, increases.
What about a group of atoms and ions that have the same number of electrons? They are referred to as an isoelectronic series and can be arranged by increasing atomic number.
Each member of the isoelectronic series depicted has 18 electrons. However, they differ in the number of protons. The sulfide ion has 16 protons attracting 18 electrons, while the calcium ion has 20 protons attracting the same number of electrons.
Thus, with more protons, calcium can pull electrons much closer to the nucleus than sulfide can, resulting in the calcium ion being smaller than the sulfide ion.
The greater the nuclear charge, the smaller the radius, although adding an electron shell disrupts this trend. Nevertheless, sulfide is the biggest and calcium is the smallest ion in this series.
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Q1: Why is a cation smaller than its parent atom?
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.
Q2: How does adding an electron affect an anion's size?
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.
Q3: What is an isoelectronic series and how does it determine ion size?
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.
Q4: How do ionic radii change as you move down a group in the periodic table?
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.
Q5: What factors determine ionic radius for any ion?
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.
Q6: Why do cations with different charges have different sizes?
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.
Q7: How does electron shielding affect the relationship between atomic and ionic radii?
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.