S3.1.3 Periodicity of properties

Trends Across a Period and Down a Group

All periodic trends can be explained by three key factors:

• Effective Nuclear Charge ($Z_{\text{eff}}$):

Higher $Z_{\text{eff}}$ pulls electrons closer, decreasing atomic and ionic radii, and increasing ionization energy and electronegativity.

• Shielding Effect:

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. For cations, the ionic radius decreases across a period due to increasing nuclear charge.
2. For anions, the same trend holds, but anions are always larger than cations within the same period.

Trend Down a Group

Ionic radius increases down a group for both cations and anions, following the same reasoning as atomic radius (additional energy levels and shielding).

Ionic Radius for Isoelectronic Ions

What Are Isoelectronic Ions?

Despite having the same electron configuration, their ionic radii vary due to differences in nuclear attraction on the electron cloud.

Trend Across Isoelectronic Series

1. In a series of isoelectronic ions (e.g., $O^{2-}$, $F^{-}$, $N{a}^{+}$, $M{g}^{2+}$), the ionic radius decreases as nuclear charge increases.
2. This is because a greater positive charge pulls the electrons closer to the nucleus, reducing the radius.

Key Factors Affecting Isoelectronic Radius:

• Nuclear Charge: Higher positive charge reduces ionic radius.
• Electron Repulsion: More negative ions experience greater electron repulsion, increasing radius.

Ionization Energy: The Energy to Remove an Electron

It can be expressed as:

$$X^{n+}(g)\to X^{(n+1)+}(g)+e^{-}$$

Trend Across a Period

1. Ionization energy increases across a period.
2. As the effective nuclear charge increases, electrons are held more tightly, requiring more energy to remove one.

Trend Down a Group

1. Ionization energy decreases down a group.
2. The outermost electrons are farther from the nucleus and experience greater shielding, making them easier to remove.

Electron Affinity: The Energy of Gaining an Electron

$$X^{n-}(g)+e^{-}\to X^{(n-1)-}(g)$$

Trend Across a Period

1. Electron affinity becomes more exothermic (more negative) across a period.
2. As nuclear charge increases, atoms more readily attract additional electrons.

Trend Down a Group

1. Electron affinity becomes less exothermic (less negative) down a group.
2. Increased shielding and distance from the nucleus reduce the attraction for an added electron.

Electronegativity: The Pull in a Bond

It is a dimensionless value, with fluorine being the most electronegative element (3.98 on the Pauling scale).

Trend Across a Period

1. Electronegativity increases across a period as atomic radius decreases and nuclear charge increases.
2. Smaller atoms with higher nuclear charge pull bonding electrons more strongly.

Trend Down a Group

1. Electronegativity decreases down a group.
2. Larger atoms with more shielding are less able to attract bonding electrons.

Reflection