Why does atomic size decrease across a period but increase down a group?

Atomic size decreases across a period due to increasing nuclear charge, but increases down a group due to added energy levels.

As you move across a period from left to right on the periodic table, the number of protons in the nucleus (the atomic number) increases. This increase in nuclear charge attracts the electrons in the energy levels more strongly, pulling them closer to the nucleus and causing the atomic radius to decrease. This effect is known as the effective nuclear charge. It's important to note that the number of energy levels remains the same across a period, so there's no additional shielding effect from inner electrons to counteract the increased nuclear charge.

On the other hand, as you move down a group on the periodic table, new energy levels are being added. Each new energy level is further from the nucleus than the last, which results in an increase in atomic size. This is because the outermost electrons are located in these higher energy levels, and they are further away from the nucleus. The shielding effect also plays a role here. The inner electrons in the lower energy levels shield the outer electrons from the pull of the nucleus, reducing the effective nuclear charge. This allows the outermost electrons to be held less tightly, and they can move further away from the nucleus, increasing the atomic size.

In summary, atomic size is influenced by two main factors: the effective nuclear charge and the number of energy levels (also known as electron shells). The effective nuclear charge increases across a period, pulling electrons closer to the nucleus and decreasing atomic size. Conversely, the number of energy levels increases down a group, pushing the outermost electrons further from the nucleus and increasing atomic size. The shielding effect of inner electrons also contributes to the increase in atomic size down a group.