Atomic Radius in the Periodic Table
The atomic radius refers to the size of an atom, typically measured from the nucleus to the outermost electron shell. Its trends in the periodic table are influenced by two key factors: atomic structure and the arrangement of electrons.

Trends in Groups
- **Increase Down a Group**: As you move down a group in the periodic table, the atomic radius increases. This is due to:
- **Addition of Electron Shells**: Each successive element in a group has an additional electron shell, which increases the distance between the nucleus and the outermost electrons.
- **Shielding Effect**: Inner electrons shield the outer electrons from the nuclear charge, reducing the effective attraction between the nucleus and the outermost electrons.

Trends in Periods
- **Decrease Across a Period**: As you move from left to right across a period, the atomic radius decreases. This can be attributed to:
- **Increasing Nuclear Charge**: With each added proton in the nucleus, the positive charge increases, pulling the electrons closer to the nucleus.
- **Constant Electron Shells**: Electrons are added to the same shell, but the increased nuclear charge results in a greater attraction, leading to a smaller atomic radius.

Summary of Trends
- **Group Trend**: Atomic radius increases down a group due to additional electron shells and increased shielding.
- **Period Trend**: Atomic radius decreases across a period due to increasing nuclear charge and constant electron shells.
Understanding these trends helps explain the chemical properties and reactivity of elements, providing insights into their behavior in various chemical reactions.