Along a period on the periodic table, the atomic radius generally decreases from left to right.
The atomic radius is defined as the distance from the nucleus to the outermost electron shell of an atom. As you move across a period, the number of protons and electrons in the nucleus increases, resulting in a greater positive charge in the nucleus.
The increased positive charge in the nucleus exerts a stronger pull on the electrons, causing the electron cloud to be pulled closer to the nucleus. As a result, the atomic radius decreases. This effect is known as effective nuclear charge.
Additionally, electron-electron repulsion plays a role in decreasing atomic radius across a period. As more electrons are added to the same energy level, they experience repulsion from each other, causing the electron cloud to shrink.
There are a few exceptions to this trend, such as the irregularity in atomic radius between the elements in the transition metals. However, the general trend along a period is a decrease in atomic radius from left to right.