The atomic radius exhibits distinctive trends across the periodic table: it decreases from left to right across a period and increases as you move down a group.

Specifically, as you progress across a period from left to right, the atomic radius tends to decrease. This phenomenon occurs due to the increasing number of protons in the nucleus, which enhances the positive charge. The greater positive charge exerts a stronger attractive force on the electrons in the outer shell, drawing them closer to the nucleus and thereby reducing the atomic radius. Although there is a slight increase in electron shielding as you move across the period, this effect is overshadowed by the increasing nuclear charge. Consequently, the overall trend is a decrease in atomic radius.

Conversely, when moving down a group in the periodic table, the atomic radius increases. This increase can be attributed to the addition of a new electron shell for each subsequent period. Each new shell is located further from the nucleus than the previous one, leading to a greater atomic radius. In this case, the influence of the additional protons in the nucleus is outweighed by the effect of the newly added electron shell, resulting in an overall increase in atomic radius.

It is important to note that the atomic radius is not a fixed value for any given element; it can vary depending on the element’s chemical environment. For instance, the atomic radius of an element may differ when it is part of a molecule compared to its elemental state. Nevertheless, the trends in atomic radius across periods and down groups remain generally consistent, regardless of the specific chemical context.

In summary, the atomic radius is influenced by two main factors: the number of protons in the nucleus, which affects the nuclear charge, and the number of electron shells, which determines the distance of the outermost electrons from the nucleus. These factors lead to a decrease in atomic radius across a period and an increase down a group within the periodic table.