Ranking By Radius Size

In this lesson, we will explore the periodic trends in atomic radius, focusing on how to rank elements based on their size.

Atomic Radius

The atomic radius is defined as the distance from the nucleus to the outermost shell of electrons. It can be influenced by two primary factors:

• Effective Nuclear Charge $({Z}_{eff})$: The net positive charge experienced by electrons in the outer shell. It increases across a period due to increasing protons in the nucleus, pulling electrons closer.
• Electron Shielding: Inner electrons shield outer electrons from the full charge of the nucleus. This effect increases down a group.

Trends in Atomic Radius

1. Period Trend: Atomic radius decreases from left to right across a period. This is due to the increase in ${Z}_{eff}$, which pulls electrons closer to the nucleus.

   Trend: Left to Right: ${R}_{1} > {R}_{2} > {R}_{3}$, where $R$ is the atomic radius.

2. Group Trend: Atomic radius increases down a group. This is because additional electron shells are added, increasing the distance of the outermost electrons from the nucleus.

   Trend: Top to Bottom ${R}_{A} < {R}_{B} < {R}_{C}$, where $R$ is the atomic radius.

Key Questions

• How does effective nuclear charge affect atomic radius?
• Why does atomic radius increase down a group?
• How can we predict the relative sizes of atoms based on their positions in the periodic table?

Conclusion

Understanding periodic trends in atomic radius helps us predict and rationalize the behavior of elements in chemical reactions and bonding.