Understanding the Decrease in Atomic Size as Atomic Number Increases Across a Period

The size of an atom, more specifically the radius of the valence electron, decreases as the atomic number increases across a period in the periodic table. This phenomenon is governed by several factors related to the atomic structure, including nuclear charge, electron shielding, and effective nuclear charge. In this article, we will delve into these factors to provide a comprehensive understanding of this intriguing behavior.

Key Factors Affecting Atomic Size

Nuclear Charge

As you move from left to right across a period in the periodic table, the atomic number increases. Consequently, the number of protons in the nucleus also increases, leading to a greater positive charge in the nucleus known as the nuclear charge. This increased positive charge exerts a stronger pull on the valence electrons, which in turn decreases the atomic radius.

Electron Shielding

Despite the addition of more electrons as you move across a period, these electrons are added to the same energy level or shell. This means that the additional electrons do not significantly shield the outer electrons from the pull of the nucleus. As a result, the nuclear charge remains relatively unshielded, leading to a stronger attractive force between the nucleus and the outer electrons.

Effective Nuclear Charge (Z_eff)

The effective nuclear charge (Z_eff) is the net positive charge experienced by an electron in a multi-electron atom. It is calculated by considering the total nuclear charge and subtracting the shielding effect of inner-shell electrons. As the atomic number increases, the effective nuclear charge also increases. This means that the valence electrons experience a stronger attraction to the nucleus, causing them to be pulled closer, and thus decreasing the atomic size.

A Real-World Example: Hydrogen and Helium

To better illustrate the influence of nuclear charge and electron shielding on atomic size, let's consider the hydrogen and helium atoms. The hydrogen atom, with an atomic number of 1, has a diameter of approximately (1.10 times 10^{-10}) meters. In contrast, the helium atom, with an atomic number of 2, has a significantly smaller diameter of approximately (0.62 times 10^{-10}) meters.

These differences in size are primarily due to the following reasons:

Nuclear Charge: The helium nucleus has twice the positive charge of the hydrogen nucleus, leading to a stronger attractive force on the valence electrons. Electron Shielding: While both atoms have a single electron in their outer shell, the shielding effect in helium is minimal. This results in the valence electron in helium experiencing a more effective nuclear charge, contributing to its smaller atomic size.

The Role of Valence Shell Filling

Once the valence shell has been filled, the next electron to join the atom will move into the next energy level, typically the 2s orbital. This transition is evident in atoms like lithium, where the third electron fills the 2s orbital. Such a change in the electron configuration can affect the shielding and therefore the effective nuclear charge experienced by the valence electrons, influencing the atomic size.

For example, lithium has a larger atomic size compared to helium because the additional electron is added to the 2s orbital, which is farther from the nucleus and thus experiences less effective nuclear charge compared to the 1s electron in helium.

Conclusion

In summary, the decrease in atomic size as the atomic number increases across a period is a result of the increased nuclear charge and the relatively poor shielding effect of additional electrons. Understanding these factors helps us predict and explain the periodic trends in atomic size, which is a fundamental concept in chemistry and physics.