The size of an atom isn’t an easy property to measure because atoms are very small and their electron shell is more of a cloud than a spherical shell. Atomic radius and ionic radius are two of the most common atom size measurements. Here are the definitions of atomic and ionic radius, the difference between them, and their periodic table trend.
Atomic Radius
The atomic radius is the average distance from the center of the nucleus of a neutral atom to the outer boundary of its electron shell. For isolated neutral atoms, the atomic nucleus ranges from 30 picometers (trillionths of a meter) and 300 pm. The largest atom is cesium, while the smallest atom is helium. Most of the size of an atom comes from its electrons. The atomic radius is over 10,000 times larger than the radius of the atomic nucleus (1 to 10 femtometers). To put it another way, the atomic radius is less than one-thousandth the wavelength of visible light (400 to 700 nm).
The edge of the electron shell isn’t well-defined, so you’ll find different values for each atom, depending on the reference. But, the actual numbers aren’t as important as the relative sizes of atoms.
Ionic Radius
While the atomic radius measures the size of a neutral atom, the ionic radius gauges the size of an electrically charged atom. The ionic radius is the radius of a monatomic ion of an element within an ionic crystal or half the distance between two bonded gas atoms. Ionic radius values range from 31 pm to over 200 pm.
Ionic radius isn’t a fixed property, so the value for an ion of an element depends on the conditions. Coordination number and spin state are the main factors that affect ionic radius measurements. X-ray crystallography yields empirical ionic radius measurements. Pauling used effective nuclear charge to calculate ionic radius. Tables of ionic radii usually indicate the method used to determine the values.
Periodic Table Trend
Electron configuration determines the organization of elements on the periodic table, so atomic and ionic radius display periodicity:
- Atomic and ionic radius increase moving down a group or column of the periodic table. This is because atoms gain an electron shell.
- Atomic and ionic radius generally decrease moving across a period or row of the periodic table. This is because the increasing number of protons exerts a stronger attraction to the electrons, drawing them in more tightly. Noble gases are the exception to this trend. The size the noble gas atom is larger than the halogen atom that precedes it.
Atomic Radius vs Ionic Radius
Atomic radius and ionic radius follow the same trend on the periodic table. But, the ionic radius may be either larger or smaller than the atomic radius of an element, depending on the electrical charge. Ionic radius increases with negative charge and decreases with positive charge.
- Cation or positive ion: An atom loses one or more electrons when it forms a cation, making the ion smaller than the neutral atom. Metals typically form cations, so their ionic radius tends to be smaller than their atomic radius.
- Anion or negative ion: An atoms gains one or more electrons to form an anion, making the ion larger than the neutral atom. Nonmetals often form anions, so their ionic radius tends to be larger than their atomic radius. This is particularly noticeable for the halogens.
Atomic and Ionic Radius Homework Questions
Students are often asked to order the size of atoms and ions based on the difference between atomic and ionic radius and the periodic table trends.
For example: List the species in order of increasing size: Rb, Rb+, F, F–, Te
You don’t need to know the sizes of the atoms and ions to order them. You know the rubidium cation is smaller than the rubidium atom because it had to lose an electron to form the ion. At the same time, you know rubidium lost an electron shell when it lost an electron. You know the fluorine anion is larger than the fluorine atom because it gained an electron to form the ion.
Next, look at the periodic table to determine the relative size of the atoms of the elements. A neutral tellurium is smaller than a neutral rubidium atom because atomic radius decreases as you move across a period. But, the tellurium atom is larger than the rubidium cation because it has an additional electron shell.
Putting it all together:
F < F– < Rb+ < Te < Rb
Other Atomic Radius Measurements
The atomic and ionic radii aren’t the only ways to measure the size of atoms and ions. Covalent radius, van der Waals radius, metallic radius, and Bohr radius are more appropriate in some situations. This is because the size of an atom is affected by its chemical bonding behavior.
- Covalent radius: The covalent radius is the the radius of atoms of an element that are covalently bonded to other atoms. It is measured as the distance between atomic nuclei in molecules, where the distance between atoms or length of their covalent bond should equal the sum of the covalent radii.
- van der Waals radius: The van der Waals radius is half of the minimum distance between the nuclei of two atoms of an element that are bound in the same molecule.
- Metallic radius: The metallic radius is the radius of an atom of an element that is connected to other atoms by metallic bonds.
- Bohr radius: The Bohr radius is the radius of the lowest energy electron orbit, calculated using the Bohr model. The Bohr radius is only calculated for atoms and ions that have a single electron.
Isoelectronic Ions
Isoelectronic ions are cations or anions of different elements that have the same electronic structure and same number of valence electrons. For example, K+ and Ca2+ both have the [Ne]4s1 electron configuration. S2- and P3- both have 1s2 2s2 2p6 3s2 3p6 as their electron configuration. Isoelectronicity may be used to compare ionic radii of different elements and to predict their properties based on their electron behavior.
References
- Basdevant, J.-L.; Rich, J.; Spiro, M. (2005). “Fundamentals in Nuclear Physics”. Springer. ISBN 978-0-387-01672-6.
- Bragg, W. L. (1920). “The arrangement of atoms in crystals”. Philosophical Magazine. 6. 40 (236): 169–189. doi:10.1080/14786440808636111
- Cotton, F. A.; Wilkinson, G. (1998). “Advanced Inorganic Chemistry” (5th ed.). Wiley. ISBN 978-0-471-84997-1.
- Pauling, L. (1960). “The Nature of the Chemical Bond” (3rd ed.). Ithaca, NY: Cornell University Press.
- Wasastjerna, J. A. (1923). “On the Radii of Ions”. Comm. Phys.-Math., Soc. Sci. Fenn. 1 (38): 1–25.