The atomic radius of a chemical element is a measure of the size of its atoms. It represents the mean distance from the nucleus to the boundary of the surrounding cloud of electrons.
Atomic radii vary in a predictable manner across the periodic table. Radii generally decrease along each period (row) of the table from left to right and increase down each group (column). These trends in atomic radii (as well as trends in various other chemical and physical properties of the elements) can be explained by considering the structure of the atom.
Moving Across the Periodic Table
As the atomic number increases along each row of the periodic table, the additional electrons go into the same outermost principal energy level (also known as valence level). This can be predicted to lead to
- an increase in atomic size because of additional repulsions between electrons,
- a decrease in size because of the additional protons in the nucleus,
- no effect at all as the two opposing tendencies of electron repulsion and nuclear attraction balance each other out.
Experiments have shown that the first case is what happens: the increase in nuclear charge overcomes the repulsion between the additional electrons in the valence level. Therefore, the size of atoms decreases as one moves across a period from left to right in the periodic table.
Moving Down the Periodic Table
The principal energy levels hold electrons at increasing radii from the nucleus. In a noble gas, the outermost level is completely filled; therefore, the additional electron that the following alkali metal (Group I) possesses will go into the next principal energy level, accounting for the increase in the atomic radius. Therefore, atomic size, or radius, increases as one moves down a group in the periodic table.