Why does the atomic radius decrease as you move from left to right across a period?

The atomic radius decreases from left to right across a period primarily because of the increase in the number of protons in the nucleus as the atomic number increases. Each element in a period has an additional proton compared to the element to its left. This increase in positive charge within the nucleus enhances the effective nuclear charge experienced by the electrons surrounding the nucleus.

As the effective nuclear charge increases, it exerts a stronger attractive force on the electrons, pulling them closer to the nucleus. This results in a reduction in the size of the electron cloud, thereby decreasing the atomic radius. Additionally, while more electrons are added as you move across the period, they enter the same principal energy level and do not shield each other very effectively from the increasing nuclear charge. Therefore, the overall effect is a decrease in atomic size due to the stronger attraction between the nucleus and the electrons.

The other options presented do not accurately explain this trend. The increase in neutrons or the notion of decreasing electron numbers do not play a role in determining atomic size in this context. Similarly, the addition of new orbitals pertains more to changes in energy levels as you move down a group rather than across a period.