IB Chemistry S3.1.3 Periodicity of Properties Notes

Trends Across a Period and Down a Group

All periodic trends can be explained by two key factors:

1. Effective Nuclear Charge ($Z_\text{eff}$):
   1. Higher $Z_\text{eff}$ pulls electrons closer, decreasing atomic and ionic radii, and increasing ionization energy and electronegativity.
2. Shielding Effect:
   1. Inner electrons shield outer electrons from the nucleus, reducing $Z_\text{eff}$ and making it easier to remove or add electrons.

Atomic Radius: The Size of an Atom

While we cannot measure this directly (since electrons exist in a cloud), we approximate it based on the distances between two bonded atoms.

Trend Across a Period

1. As you move across a period from left to right, the atomic radius decreases.
2. Each successive element adds a proton to the nucleus and an electron to the same energy level.
3. This increases the effective nuclear charge ($Z_\text{eff}$), pulling the electrons closer to the nucleus.

Trend Down a Group

1. As you move down a group, the atomic radius increases.
2. This is because each successive element adds a new energy level (or shell), increasing the distance between the nucleus and the outermost electrons.
3. Although the nuclear charge also increases, the additional inner shells provide shielding, reducing the pull on the valence electrons.

Ionic Radius: The Size of Ions

Cations (positively charged ions) are smaller than their parent atoms, while anions (negatively charged ions) are larger.

Trend Across a Period

1. Across a period, the cations (the left-hand elements) form one series in which ionic radius decreases as nuclear charge rises, for example $Na^+ > Mg^{2+} > Al^{3+}$.