How does atomic radius change across a period?

Atomic radius decreases across a period from left to right due to an increase in nuclear charge.

As you move across a period in the periodic table from left to right, the atomic radius, or size of the atom, tends to decrease. This is primarily due to an increase in the positive charge of the nucleus, which pulls the electrons in the outer shell closer to the nucleus, thus reducing the size of the atom.

In each period, the number of protons in the nucleus increases. This increase in nuclear charge attracts the electrons more strongly, pulling them closer to the nucleus. At the same time, the number of energy levels (shells) remains the same as you move across a period. This means that the increase in nuclear charge is not counterbalanced by an increase in the distance between the nucleus and the outer shell. As a result, the outer shell is pulled closer to the nucleus, causing the atomic radius to decrease.

Furthermore, the increase in nuclear charge also leads to an increase in the effective nuclear charge experienced by the electrons. The effective nuclear charge is the net positive charge experienced by an electron in a multi-electron atom. It is affected by both the actual nuclear charge and the degree of shielding by other electrons. As the nuclear charge increases across a period, the effective nuclear charge also increases, further pulling the electrons closer to the nucleus.

In summary, the atomic radius decreases across a period due to an increase in nuclear charge, which pulls the electrons closer to the nucleus, and an increase in effective nuclear charge, which further enhances this effect. This trend is a fundamental aspect of atomic structure and is crucial for understanding the properties of elements and their reactivity.