Final answer:
As you move across a period, atomic radii decrease due to the increased attraction between the growing number of protons and inadequately shielded electrons, increasing the effective nuclear charge (Zeff) and drawing electrons closer to the nucleus.
Step-by-step explanation:
The trend in atomic radii as you move across a period is due to increased attraction between a greater number of protons and electrons. As you move from left to right across a period, the atomic number (Z) increases, as does the positive charge of the nucleus. However, the increase in shielding is only slight, leading to an increase in the effective nuclear charge (Zeff). This stronger pull experienced by the electrons draws them closer to the nucleus, resulting in smaller covalent radii.
The concept of effective nuclear charge is central to understanding this trend. Although electrons are added to the same principal energy level, they do not shield each other effectively from the increasing nuclear charge. Hence, each successive element across a period has a slightly greater Zeff than the last, leading to a decrease in the atomic radius.