The d-block elements occupy Groups 3 to 12 of the periodic table, sitting between the s-block and the p-block. They are called transition elements because they mark a transition between the most electropositive s-block metals and the more electronegative p-block elements. By the strict IUPAC definition, a transition element has an incompletely filled d subshell either in its elemental state or in one of its common oxidation states. There are three complete series: the 3d series (Sc–Zn), the 4d series (Y–Cd) and the 5d series (La, Hf–Hg).
The general outer electronic configuration is $(n-1)d^{1-10}\,ns^{1-2}$. Two configurations break the simple pattern because half-filled and fully-filled d subshells carry extra stability: chromium is $3d^5\,4s^1$ (not $3d^4\,4s^2$) and copper is $3d^{10}\,4s^1$ (not $3d^9\,4s^2$).
General characteristics
Transition metals are hard, lustrous, high-melting solids that conduct heat and electricity well. They are less reactive than s-block metals, form alloys readily, and act as catalysts. These properties stem from the involvement of $(n-1)d$ electrons in bonding alongside the $ns$ electrons.
Atomic and ionic radii
Across a series the radius first decreases (increasing nuclear charge pulls electrons in), then stays nearly constant in the middle (the added d electron screens the nucleus, balancing the rise in charge), and rises slightly at the end (electron–electron repulsion in the filling d subshell). The 4d and 5d elements of a group have almost identical radii because of the lanthanoid contraction.
Ionisation enthalpies
Ionisation enthalpies increase gradually across a series, but far less steeply than in the s- or p-block, because the added electron enters the inner $(n-1)d$ shell and the increase in nuclear charge is partly offset by screening. Irregularities (e.g. the high $\text{IE}_3$ of Mn and Zn) reflect the stability of $d^5$ and $d^{10}$ configurations.
Oxidation states
The hallmark of transition metals is variable oxidation states, because the $(n-1)d$ and $ns$ electrons have similar energies and can both be lost. Manganese shows the widest range (+2 to +7). The highest states appear in oxides and fluorides (e.g. $\text{Mn}_2\text{O}_7$, $\text{KMnO}_4$).
Metallic character and melting points
All transition elements are metals. Melting points rise to a maximum near the middle of each series (greatest number of unpaired d electrons available for metallic bonding) and fall towards the ends. Manganese and technetium dip because of their stable half-filled configurations.
Why Zn, Cd and Hg are not typical
Zinc, cadmium and mercury have a fully filled $d^{10}$ configuration in both the atom and their common +2 ion. With no partially filled d subshell, they are not coloured, are weakly bonded (low melting points; Hg is liquid), are poor catalysts and form few complexes — so they are not regarded as typical transition elements.