Introduction to d-Block Elements
Bridge the gap between s and p blocks. Master the strict definition of transition metals, the Chromium/Copper electronic anomalies, and the physics behind their melting points.
Module Focus: The Transition Metals
The d-block elements occupy the large middle section of the periodic table flanked by s- and p-blocks. Their properties represent a "transition" between the highly reactive metallic s-block and the largely non-metallic p-block. However, in chemistry, the term "Transition Element" has a very strict, electron-based definition that examiners love to test.
1. The Strict Definition & The Zinc Trap
Not all d-block elements are transition elements. IUPAC defines a transition element precisely based on its d-orbitals.
A transition element is defined as the one which has incompletely filled d orbitals in its ground state OR in any one of its common oxidation states.
Zinc ($Zn$), Cadmium ($Cd$), and Mercury ($Hg$) have fully filled $d^{10}$ configurations in their ground state.
Even in their common $+2$ oxidation state, they lose their two s-electrons, remaining $d^{10}$.
Group 11 metals like Copper ($Cu$) also have a fully filled $d^{10}$ configuration in their ground state ($3d^{10} 4s^1$).
However, in its common $+2$ oxidation state ($Cu^{2+}$), the configuration becomes $3d^9$.
2. Electronic Configurations
The general electronic configuration of d-block elements is $(n-1)d^{1-10} ns^{1-2}$. The filling of the inner $(n-1)d$ orbitals causes unique anomalies.
| Element | Atomic No. (Z) | Expected vs Actual Configuration |
|---|---|---|
| Scandium (Sc) | 21 | $[Ar] \ 3d^1 \ 4s^2$ |
| Chromium (Cr) | 24 |
Expected: $[Ar] \ 3d^4 \ 4s^2$ Actual: $[Ar] \ 3d^5 \ 4s^1$ |
| Manganese (Mn) | 25 | $[Ar] \ 3d^5 \ 4s^2$ |
| Copper (Cu) | 29 |
Expected: $[Ar] \ 3d^9 \ 4s^2$ Actual: $[Ar] \ 3d^{10} \ 4s^1$ |
An electron shifts from the $4s$ orbital to the $3d$ orbital in $Cr$ and $Cu$ to achieve exactly half-filled ($d^5$) or fully-filled ($d^{10}$) configurations. These configurations are exceptionally stable due to two factors:
- Symmetrical Distribution: A symmetrical electron distribution minimizes electron-electron repulsion.
- Exchange Energy: Electrons with parallel spins in degenerate orbitals can exchange positions. Maximum exchanges occur in half/fully filled shells, releasing energy and stabilizing the atom.
3. Enthalpy of Atomization & Melting Points
Transition metals generally have high melting and boiling points, and high enthalpies of atomization. This indicates very strong metallic bonding.
In addition to the metallic bond formed by $ns$ electrons, transition metals utilize their $(n-1)d$ electrons for covalent-like interatomic bonding.
Rule: The strength of the metallic bond (and thus the melting point and enthalpy of atomization) is roughly proportional to the number of unpaired d-electrons.
Elements in the middle of each series have the maximum number of unpaired electrons ($d^5$). Therefore, they have the highest melting points.
Manganese ($3d^5 4s^2$) and Technetium have exactly half-filled d-subshells. These configurations are so stable that the d-electrons are held tightly by the nucleus and do not participate effectively in metallic bonding.
Sc < Ti < V < Cr (Max) > Mn (Dip) < Fe > Co > Ni > Cu > Zn (Lowest)
Zinc has $0$ unpaired d-electrons, hence the weakest metallic bonding and lowest MP. Mercury (Hg) is a liquid for the same reason in the 5d series.
NEET Grand Test: d-Block Basics
15 High-Yield Questions testing IUPAC definitions, exact configurations, and thermodynamic anomalies.
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