Atomic radius generally decreases across a period (left to right) and increases down a group (top to bottom). This pattern stems from the increasing nuclear charge within a period and from the increasing number of electron shells down a group.

Across a period: As the atomic number increases across a period, more protons are added to the nucleus, which increases the positive charge of the nucleus and the attractive force exerted on the electrons. This causes the electrons to be pulled closer to the nucleus, decreasing the atomic radius.

Down a group: As you go down a group, new electron shells are added, which increases the distance between the outermost electrons and the nucleus. The additional electron shells act as barriers that shield the outermost electrons from the nucleus's positive charge, reducing the attractive force and increasing the atomic radius.

Exceptions: There are a few exceptions to this general pattern, mainly due to the irregular shielding effect of the inner electron shells. For example, the atomic radii of the transition metals in the fourth period (Y, Zr, Nb, etc.) are slightly larger than expected, while the atomic radii of the group 13 elements (B, Al, Ga, etc.) are slightly smaller than expected. These exceptions can be attributed to specific electron configurations and the interactions between electrons within the atoms.

In summary, the atomic radius periodic pattern reflects the interplay between the number of protons and electrons in an atom, as well as the shielding effect of inner electron shells, leading to the observed trends across periods and down groups.