Q.Why do the transition elements exhibit higher enthalpies of atomisation?

The key idea is that transition elements have strong metallic bonding due to the involvement of unpaired $d$-electrons in interatomic bonding.

Reasoning:

1. In transition metals, the $(n-1)d$ and $ns$ electrons participate in metallic bonding.
2. As we move across a series, the number of unpaired $d$-electrons increases up to the middle, leading to stronger covalent character in the metallic bond. …

The Core Idea: What Enthalpy of Atomisation Really Means

Enthalpy of atomisation ($\Delta H_{atom}$) is the energy required to break one mole of a solid metal into isolated gaseous atoms. For transition elements, this value is significantly higher than for s-block metals. The reason lies in the nature of metallic bonding in these elements.

In a transition metal solid, each atom contributes not just its outer $ns$ electrons but also its $(n-1)d$ electrons to the metallic bond. This creates a much stronger "electron sea" that holds the lattice together. More electrons in the bonding pool means more energy is needed to pull the atoms apart.

Step-by-Step Reasoning

1. The unique electron configuration of transition metals

Transition elements have the general electronic configuration $[noble\ gas]\ (n-1)d^{1-10}\ ns^{1-2}$. The key point is that both the $ns$ and $(n-1)d$ orbitals are close in energy and can participate in bonding.

2. How metallic bonding works here

In the solid state, the $ns$ electrons delocalise first, forming the basic metallic bond. But unlike s-block metals, transition elements also allow their $(n-1)d$ electrons to delocalise into the conduction band. This means each atom contributes more than one electron to the metallic bond — typically 2–3 electrons per atom instead of just 1 or 2.

3. The trend across the series

The enthalpy of atomisation generally increases from left to right across a transition series, peaks near the middle, then decreases. This pattern directly mirrors the number of unpaired $d$ electrons available for bonding.

4. Why the middle of the series has the highest values …

Method: Electronic Configuration & Metallic Bonding Analysis

This method explains physical properties of transition elements by linking their electronic structure to the strength of metallic bonding.

Steps

1. Recall the definition of transition elements

   Transition elements are those which have partially filled $d$-orbitals either in their ground state or in any of their common oxidation states.

   Example: Fe ($[Ar]\,3d^6 4s^2$), Cu ($[Ar]\,3d^{10} 4s^1$ — Cu⁺ has $3d^{10}$, but Cu²⁺ has $3d^9$, so it qualifies).

2. Identify the key feature: unpaired $d$-electrons

   Most transition metals have unpaired electrons in their $3d$, $4d$, or $5d$ subshells. These unpaired electrons participate in metallic bonding along with the $ns$ electrons.

3. Link to enthalpy of atomisation

   Enthalpy of atomisation ($\Delta H_{\text{atom}}$) is the energy required to convert one mole of a solid metal into isolated gaseous atoms.

   • Stronger metallic bonds → more energy needed → higher $\Delta H_{\text{atom}}$.

4. Explain the “why”

   • In transition metals, the $d$-electrons (especially unpaired ones) contribute to covalent character in the metallic bond via $d$-$d$ overlap.
   • This creates stronger interatomic attraction compared to $s$-block metals (which have only $s$-electrons in bonding).
   • More unpaired $d$-electrons → stronger bonding → higher enthalpy of atomisation.

5. State the trend …

🔍 The Core Concept First

Enthalpy of atomisation is the energy required to convert one mole of a solid element into isolated gaseous atoms.

For transition elements, this value is higher than for s-block elements because:

• Strong metallic bonding due to unpaired d-electrons participating in bonding.
• More unpaired electrons → stronger interatomic attraction → more energy needed to separate atoms.

✗ Common Mistake #1: Confusing "enthalpy of atomisation" with "ionisation enthalpy"

What students write:

"Transition elements have high atomisation enthalpy because they have high ionisation enthalpy."

Why it's wrong:

Ionisation enthalpy is about removing an electron from an atom — atomisation is about separating atoms from the solid. They are different processes.

✓ How to avoid:

Always define the term before answering. Remember:

• Atomisation = solid → gaseous atoms (breaking metallic bonds)
• Ionisation = gaseous atom → gaseous cation (removing electron)

✗ Common Mistake #2: Saying "d-orbitals are involved" without specifying how

What students write:

"Due to the presence of d-orbitals."

Why it's wrong:

Vague. Every transition element has d-orbitals — you must explain why that leads to stronger bonding.

✓ How to avoid:

Be specific:

• Unpaired d-electrons contribute to covalent character in metallic bonding.
• More unpaired electrons → stronger bonding → higher $\Delta H_{\text{atom}}$.

Example:

V (3 unpaired $d$-electrons, all bonding-active) → very high atomisation enthalpy (515 kJ mol⁻¹).

Zn (0 unpaired electrons) → the lowest (126 kJ mol⁻¹).

Careful: Mn is the famous exception — despite 5 unpaired electrons, its stable half-filled $3d^5$ set participates poorly in bonding, so its value (281 kJ mol⁻¹) is anomalously LOW.

✗ Common Mistake #3: Ignoring the trend across the series

What students write:

"All transition elements have equally high atomisation enthalpies."

Why it's wrong:

The values vary — they rise to a maximum at V (515 kJ mol⁻¹), dip at Cr and especially Mn (stable half-filled $3d^5$), recover over Fe–Ni, and fall to the minimum at Zn (126 kJ mol⁻¹).

✓ How to avoid:

Mention the general trend:

• Rises from Sc to V
• Dips at Cr and (sharply) Mn — their stable $3d^5$ holds electrons back from bonding
• Recovers for Fe–Ni, then falls through Cu to Zn

Reason: what counts is the number of unpaired electrons actually taking part in metallic bonding.

✗ Common Mistake #4: Forgetting to compare with s-block elements …