S2.3.3 Transition elements and delocalized d-electrons (Higher Level Only)

The Role of d-Electrons in Transition Metals and Trends in the d-Block

Imagine holding a piece of copper wire in your hand. It’s durable, conducts electricity with ease, and has a remarkably high melting point. But why is copper—and other transition metals—so different from metals like sodium or potassium?

The answer lies in the unique behavior of d-electrons and their role in metallic bonding.

Metallic Bonding and the Role of d-Electrons

What Makes Metallic Bonding Unique?

In metals, valence electrons don’t belong to individual atoms. Instead, they form a shared "sea of delocalized electrons" that move freely throughout the metallic lattice. These electrons interact with positively charged cations, generating strong electrostatic forces known as metallic bonds.

For most metals, these delocalized electrons originate from the s-orbital of the outermost energy level. However, in transition metals, electrons from partially filled d-orbitals also contribute to this electron sea. This additional electron density significantly strengthens metallic bonds, setting transition metals apart.

The Impact of d-Electrons on Properties

The delocalized d-electrons in transition metals give rise to several key properties:

1.Higher Melting Points

The strength of metallic bonds depends on the density of delocalized electrons and their attraction to the metal cations.

Transition metals, with their additional d-electrons, form stronger metallic bonds than metals like sodium or magnesium.

This explains their generally higher melting points.

2.High Electrical Conductivity

The mobility of delocalized electrons makes metals excellent conductors of electricity.

In transition metals, the greater number of delocalized electrons (from both s- and d-orbitals) enhances this conductivity.

When a potential difference is applied, these electrons move freely, allowing electric current to flow.

Trends in the d-Block

The d-block elements (transition metals) exhibit fascinating trends in their physical properties, but these trends are often more complex than those observed in the s- or p-block elements. Let’s explore these patterns.

Melting Point Trends Across the d-Block

You might expect melting points to increase steadily across a period in the d-block as more delocalized electrons strengthen metallic bonds. However, this trend is not always straightforward. Why?

1.Addition of d-Electrons

As we move across a period in the d-block, more d-electrons become delocalized, which should strengthen metallic bonds and increase melting points.

2.Electron-Electron Repulsion

However, adding d-electrons also increases electron-electron repulsion within the d-orbitals. This repulsion can partially offset bonding strength, leading to irregularities in the melting point trend.

Comparison with Group 1 and Group 2 Metals

To appreciate the unique properties of transition metals, let’s compare them to group 1 and group 2 metals:

• Group 1 Metals (e.g., Sodium): These metals have only one delocalized s-electron per atom, resulting in weaker metallic bonds and lower melting points.
• Group 2 Metals (e.g., Magnesium): These metals have two delocalized s-electrons, leading to stronger bonds and higher melting points than group 1 metals.
• Transition Metals (e.g., Iron): The combination of s- and d-electrons results in even stronger metallic bonds, giving rise to higher melting points and greater hardness.

Key Relationship: Bond Strength and Metallic Properties

The strength of metallic bonds—and the resulting properties of metals—can be summarized using the following relationship:

$$\text{Bond Strength}\propto \frac{\text{Charge of Cation}}{\text{Radius of Cation}}$$

• Higher Charge: Increases the electrostatic attraction between cations and delocalized electrons.
• Smaller Radius: Reduces the distance between cations and electrons, strengthening the bond.

Reflection and Connections