Final answer:
The atomic radius of elements decreases when moving across a period due to increased effective nuclear charge, despite more electrons being added. Down a group, the radius increases as electrons are added to higher energy levels further from the nucleus. The radius decreases for cations and increases for anions compared to their parent atoms.
Step-by-step explanation:
When examining the radius of elements on the periodic table, it's observed that although the number of electrons increases as we move from left to right across a period, the atomic radius decreases. This is due to the concept of effective nuclear charge (Zeff), which is the net positive charge exerted by the nucleus on the valence electrons after subtracting any electron-electron repulsions due to shielding by inner electrons. The increased nuclear charge attracts the outermost electrons more strongly, pulling them closer to the nucleus and resulting in a smaller atomic radius.
Down a group, on the other hand, the principal quantum number, n, increases, and electrons are added to energy levels that are further from the nucleus. This results in the increase of the atom's size, since the outermost electrons are located further from the nuclear pull.
However, when an atom becomes a cation by losing electrons, its radius becomes smaller compared to the neutral atom. Conversely, when an atom becomes an anion by gaining electrons, its radius enlarges due to increased electron-electron repulsion.