Atomic radius is a term used to describe the size of an atom, however, there is no standard definition for this value. An atomic radius may refer to the ionic radius, covalent radius, metallic radius, or van der Waals radius.

No matter what criteria you use to describe the atomic radius, the size of an atom is dependent on how far out its electrons extend. The atomic radius of an element tends to increase the further down you go in an element group. That's because the electrons become more tightly packed as you move across the periodic table, so while there are more electrons for elements of increasing atomic number, the atomic radius may decrease. The atomic radius moving down an element period or column tends to increase because an additional electron shell is added for each new row. In general, the largest atoms are at the bottom lefthand side of the periodic table.

The atomic and ionic radius is the same for atoms of neutral elements, such as argon, krypton, and neon. However, many atoms of elements are more stable as atomic ions. If the atom loses its outermost electron, it becomes a cation or positively charged ion. Examples include K+ and Na+. Some atoms might lose multiple outer electrons, such as Ca2+. When electrons are removed from an atom, it might lose its outermost electron shell, making the ionic radius smaller than the atomic radius.

In contrast, some atoms are more stable if they gain one or more electrons, forming an anion or negatively charged atomic ion. Examples include Cl- and F-. Because another electron shell isn't added, the size difference between the atomic radius and ionic radius of an anion isn't as much as for a cation. The anion ionic radius is the same as or slightly larger than the atomic radius.

Overall, the trend for the ionic radius is the same as for the atomic radius: increasing in size moving across and decreasing moving down the periodic table. However, it's tricky to measure the ionic radius, not least because charged atomic ions repel each other.