General Properties of Transition Elements

What Counts as a Transition Element

A transition element is defined as one whose atom has an incompletely filled d subshell, or which can give rise to cations with an incomplete d subshell. The d-block proper spans Groups 3 to 12 across the 3d, 4d and 5d series. Scandium and zinc sit at the boundaries and illustrate the definition sharply: $\ce{Zn}$ ($\ce{[Ar]3d^{10}4s^2}$) and its ion $\ce{Zn^2+}$ ($\ce{3d^{10}}$) have a fully filled d shell, so zinc is not regarded as a typical transition element even though it appears in the block.

Because the general properties below flow almost entirely from partly filled d orbitals, it pays to keep the 3d configurations in view. They are the data set on which NEET trend questions are built.

Metallic Character and Physical Properties

Nearly all transition elements display the classic metallic signature: high tensile strength, ductility, malleability, high thermal and electrical conductivity, and metallic lustre. With the exceptions of $\ce{Zn}$, $\ce{Cd}$, $\ce{Hg}$ and $\ce{Mn}$, they crystallise in one or more of the typical close-packed metallic structures (bcc, hcp or ccp) at ordinary temperatures.

The root cause is the strength of the metallic bond. Transition metals can throw both their ns electrons and a number of their (n−1)d electrons into the delocalised bonding lattice. The more unpaired d electrons available, the stronger the binding — which is why hardness, low volatility and high melting points all peak near the middle of a series and ease off towards the ends.

The same logic explains the soft, volatile outliers. $\ce{Zn}$, $\ce{Cd}$ and $\ce{Hg}$ have a complete $\ce{d^{10}}$ shell with no unpaired d electrons to contribute; mercury is liquid at room temperature, and these three are the softest, lowest-melting metals in the block.

Atomic and Ionic Radii Trends

Two trends matter here: how radius changes across a series, and how it changes down a group.

Across a series, ions of the same charge show a progressive decrease in radius with increasing atomic number, and atomic radii follow the same direction. Each added electron enters an inner (n−1)d orbital. A d electron shields the outer electrons rather poorly, so the effective nuclear charge felt by the outermost electron rises and the size contracts. Crucially, though, the variation within a series is small: the d electrons shield the 4s electrons well enough that the radius falls only gently and is nearly constant through the middle of the row.

Down a group the surprise is that the 4d and 5d elements have almost the same radii. On going from the 3d to the 4d series the size increases as expected, but the 5d radii are virtually identical to their 4d partners. The reason is the intervention of the 4f orbitals, which must fill before the 5d series begins. The poor shielding by 4f electrons produces a steady contraction — the lanthanoid contraction — that cancels the expected size increase. The textbook pairing is $\ce{Zr}$ (160 pm) and $\ce{Hf}$ (159 pm), nearly identical in radius and remarkably alike in chemistry.

Density, Melting and Boiling Points

Transition metals (except $\ce{Zn}$, $\ce{Cd}$, $\ce{Hg}$) are very hard, of low volatility, and have high melting and boiling points. Across the 3d series the melting points rise to a maximum around $\ce{d^5}$ and then fall regularly, except for the anomalously low values of $\ce{Mn}$ and $\ce{Tc}$. The driving factor is again metallic bonding: the more (n−1)d and ns electrons engaged in interatomic bonding, the higher the melting point. One unpaired electron per d orbital — the situation near the middle of the series — is especially favourable for strong interatomic interaction.

Two further generalisations from the enthalpy-of-atomisation data deserve attention. First, the enthalpies of atomisation peak in the middle of each series, confirming that one unpaired d electron per orbital maximises metal–metal bonding. Second, the 4d and 5d metals have greater enthalpies of atomisation than their 3d counterparts, which is why metal–metal bonding is far more common in compounds of the heavier transition metals.

Density rises across the 3d series. The metallic radius shrinks while atomic mass climbs, so mass is packed into a smaller volume. From $\ce{Ti}$ (Z = 22) to $\ce{Cu}$ (Z = 29) the increase is marked — densities run from about 4.1 to nearly 8.9 g cm⁻³.

Ionisation Enthalpies

Along each series the ionisation enthalpy generally increases left to right, because nuclear charge rises as the inner d orbitals fill. But two features set transition metals apart from main-group elements. The increase is far less steep than across a typical period of non-transition elements, and the first-ionisation trend across the 3d row is distinctly irregular rather than smooth.

The irregularity comes from the shifting balance of three terms: the attraction of each electron to the nucleus, inter-electronic repulsion, and exchange energy. Exchange energy stabilises configurations with many parallel spins, so half-filled ($\ce{d^5}$) and fully filled ($\ce{d^{10}}$) sets resist losing an electron more than a smooth nuclear-charge argument would predict. This is why the trend breaks at points such as the second ionisation enthalpy of $\ce{Cr}$ (its $\ce{Cr+}$ would be $\ce{d^5}$) and $\ce{Cu}$ (its $\ce{Cu+}$ is $\ce{d^{10}}$), which are unusually high.

The successive enthalpies also behave characteristically: the second and third ionisation enthalpies rise much more sharply than the first because, once the 4s electrons are gone, every further electron is pulled from a $\ce{d^n}$ core where d–d shielding is poor. The high third ionisation enthalpies of $\ce{Cu}$, $\ce{Ni}$ and $\ce{Zn}$ are precisely why oxidation states above +2 are hard to reach for these elements.

Standard Electrode Potentials (E°)

The lowest common oxidation state of these metals is +2. Forming $\ce{M^2+(aq)}$ from the gaseous atom needs the enthalpy of atomisation plus the first and second ionisation enthalpies, offset by the hydration enthalpy of $\ce{M^2+}$. The $\ce{E^\circ}$(M²⁺/M) value is set by how these terms balance, so it is a thermochemical composite rather than a single quantity.

The general trend is towards less negative $\ce{E^\circ}$ values across the series, mirroring the increasing sum of the first and second ionisation enthalpies. But the curve is not smooth — $\ce{Mn}$, $\ce{Ni}$ and $\ce{Zn}$ are more negative than the trend predicts:

Copper is the standout: its $\ce{E^\circ}$ is positive (+0.34 V). The high energy needed to convert $\ce{Cu(s) -> Cu^2+(aq)}$ — a large enthalpy of atomisation plus high ionisation enthalpies — is not balanced by its hydration enthalpy, so copper cannot liberate hydrogen from non-oxidising acids. Only oxidising acids attack it:

$\ce{3Cu + 8HNO3(dilute) -> 3Cu(NO3)2 + 2NO ^ + 4H2O}$

The $\ce{E^\circ}$(M³⁺/M²⁺) values tell a parallel story driven by ion stability. The low value for $\ce{Sc}$ reflects the noble-gas stability of $\ce{Sc^3+}$; the high value for $\ce{Mn}$ shows that $\ce{Mn^2+}$ ($\ce{d^5}$) resists oxidation; the comparatively low value for $\ce{Fe}$ reflects the extra stability of $\ce{Fe^3+}$ ($\ce{d^5}$). These same energetics make $\ce{Cr^2+}$ a reducing agent (it moves to the half-filled $t_{2g}$ $\ce{d^3}$) while $\ce{Mn^3+}$ is oxidising (it falls to $\ce{Mn^2+}$, $\ce{d^5}$).

Formation of Coloured Ions

Most transition metal ions are coloured in the solid state and in solution — one of the most visible signatures of the block. The mechanism is a d–d transition: when the ion has partly filled d orbitals, an electron can be promoted from a lower-energy d orbital to a higher-energy d orbital. The energy gap corresponds to a frequency in the visible region, so visible light is absorbed and the ion appears in the complementary colour of the light absorbed. The size of the gap, and hence the colour, is governed by the nature of the surrounding ligand (in aqueous solution, water).

The mechanics of how the d orbitals split into two energy sets under a ligand field — and how the gap controls colour and magnetism — belongs to crystal field theory and the magnetism discussion. For the energetic basis of paramagnetism and the spin-only formula, see the dedicated note linked below.

Complex Formation

Transition metals form an enormous number of complex compounds — species in which a metal ion binds several anions or neutral molecules (ligands). Familiar examples include $\ce{[Fe(CN)6]^3-}$, $\ce{[Fe(CN)6]^4-}$, $\ce{[Cu(NH3)4]^2+}$ and $\ce{[PtCl4]^2-}$.

Three features of the metal ions explain this readiness to complex: their comparatively small size, their high ionic charge, and the availability of vacant d orbitals of suitable energy to accept lone pairs from ligands. Together these give a high charge density and the right orbitals for coordinate bonding. The detailed bonding, geometry, isomerism and nomenclature of these species are taken up fully in Coordination Compounds.