Position Electronic Config

Introduction: The Metals That Built Civilisation

Iron, copper, silver and gold — four metals that shaped human history — share something invisible to the eye: each carries a partly filled set of d-orbitals. Throw in titanium for jet engines, vanadium for catalysts, chromium for steel, and you have the d-block elements. Add the radioactive uranium and thorium that power reactors, and you reach the f-block.

This chapter unlocks why a single block in the periodic table — Groups 3 to 12 — produces metals so versatile they appear in every cell of your body (Fe in haemoglobin), in every electronic device (Cu wiring), and in every modern catalyst.

8.1 Position in the Periodic Table

The d-block sits in the wide central section of the periodic table — flanked on the left by the s-block and on the right by the p-block. Within these elements, it is the d-orbitals of the penultimate (n-1) energy level that progressively fill, giving four horizontal series:

• 3d series — Sc to Zn (Z = 21 to 30) in Period 4.
• 4d series — Y to Cd (Z = 39 to 48) in Period 5.
• 5d series — La and Hf to Hg (Z = 57, 72 to 80) in Period 6.
• 6d series — Ac and Rf to Cn (Z = 89, 104 to 112) in Period 7.

The two inner-transition (lanthanoid 4f and actinoid 5f) series sit in a separate panel below the main table.

8.2 Electronic Configurations of the d-Block

The general outer configuration of a d-block element is

The (n-1) tells us electrons enter the inner d-orbitals of the second-from-last shell, while one or two electrons sit in the outer ns. Because the energy difference between (n-1)d and ns is very small, the simple Aufbau prediction sometimes fails — and we get exceptions driven by the extra stability of half-filled (d⁵) and fully filled (d¹⁰) sub-shells.

8.2.1 The Cr and Cu Anomalies

Naively, Cr (Z = 24) should be 3d⁴4s² and Cu (Z = 29) should be 3d⁹4s². Instead, both promote one 4s electron into 3d, giving:

The added exchange energy from a half-filled or fully filled d-shell more than pays for the small extra cost of unpairing the 4s electrons.

8.2.2 Outer Configurations Across All Four Series

Notice the unique Pd exception: it skips its 5s entirely (4d¹⁰ 5s⁰) — a consequence of two paired-up exchange contributions plus relativistic effects.

8.3 General Properties of the Transition (d-Block) Elements

Almost every transition metal is hard, lustrous, ductile, malleable, with high tensile strength and excellent thermal/electrical conductivity. The exceptions you must remember are the soft, low-melting Zn, Cd, Hg (no unpaired d-electrons available for metallic bonding) and the unusual Mn (complex bcc structure).

8.3.1 Lattice Structures

bcc = body-centred cubic · hcp = hexagonal close-packed · ccp = cubic close-packed · X = atypical metal lattice.

8.3.2 High Melting Points and Enthalpy of Atomisation

Transition metals melt at very high temperatures because both the (n-1)d and ns electrons participate in metallic bonding. In each row the melting point peaks roughly at d⁵ (one electron per d-orbital, all spins parallel) — the textbook example is W (3683 K, the highest of any metal). Mn and Tc dip anomalously, and the curve falls again towards Zn/Cd/Hg.

8.4 Why Are They Called "Transition" Elements?

The historical name comes from their position: they bridge the highly electropositive s-block and the largely non-metallic p-block. Chemically, they are "transitional" in two senses — they share metallic character with the s-block but display the variable oxidation states and complex-forming tendency that some p-block elements show only weakly.

The presence of partly filled d-orbitals gives the d-block elements five hallmark features that we will explore in Part 2:

1. Variable oxidation states (e.g. Mn from +2 to +7).
2. Coloured ions (d-d transitions).
3. Paramagnetism (unpaired d-electrons).
4. Complex-ion formation (vacant d-orbitals + small size).
5. Catalytic activity (multiple oxidation states + adsorption sites).