Final answer:
The atomic radius decreases from left to right across a period of the periodic table, mainly due to the increased nuclear charge that pulls electrons closer to the nucleus. Minor exceptions may occur due to electron-electron repulsions. In contrast, atomic radius increases from top to bottom in a group because additional energy levels are added.
Step-by-step explanation:
The trend in atomic radius across the periodic table from left to right involves a decrease in atomic size. This is due to the increase in positive charge from additional protons in the nucleus, which strengthens the electrostatic pull on the electrons. Consequently, these electrons are drawn closer to the nucleus, resulting in a smaller atomic radius. While the valence electrons remain in the same principal energy level across a period, the growing nucleus charge dictates this trend.
It is important to note that there may be minor exceptions to this trend, where electron-electron repulsions can impact atomic size. Nonetheless, the atomic radius generally decreases from left to right in a period. On the other hand, as we move down a group, the atomic radius increases due to the addition of valence electrons in higher energy levels, or shells. Therefore, the largest atoms can be found in the lower-left corner of the periodic table, while the smallest are in the upper right.