Properties Trends

Why d-Block Chemistry Is Endlessly Varied

What makes a transition metal a transition metal? Five interlocking properties — variable oxidation states, paramagnetism, coloured ions, catalysis and complex formation — all flow from a single architectural fact: the (n-1)d sub-shell is partly filled and lies energetically close to the ns sub-shell. In Part 2 we examine how that single fact ripples through the periodic trends of the 3d series.

8.5 Variation in Atomic and Ionic Sizes

As we move across a transition series, the nuclear charge rises by one each step, and a new electron enters the inner d sub-shell. Inner d-electrons shield the outer ns electrons rather poorly, so the effective nuclear charge felt by the outermost shell rises — and the atomic / ionic radius shrinks. The shrinkage, however, is much smaller than across a typical s- or p-block period because the d-electrons themselves screen each other to some extent.

8.5.1 The "Plateau" After Mn — and the d-Block Contraction

From Sc to Cr the radius drops smoothly. Then from Mn to Cu the values flatten near 125–129 pm. Why? In the second half, electrons start to pair up in the d-orbitals; the rising electron–electron repulsion partly cancels the rising nuclear pull, so the radius barely changes. At Zn, the now-full 3d¹⁰ shell repels the 4s electrons outwards, pushing the radius back up to 137 pm.

8.5.2 Comparing the Three Series — Lanthanoid Contraction

If you compare the 3d, 4d and 5d series at the same group, you find that the radii grow on going 3d → 4d (more shells), but then 4d ≈ 5d (Zr 160 pm vs Hf 159 pm; Nb 146 pm vs Ta 146 pm). The reason is a one-time event embedded between the second and third series — the filling of the entire 4f sub-shell of the lanthanoids, called the lanthanoid contraction. The 4f electrons shield very poorly, so the 5d series begins after a hidden 14-element shrinkage that almost exactly cancels the size increase you would expect on going from Period 5 to Period 6.

8.6 Ionisation Enthalpies

The first ionisation enthalpy of the 3d series rises slowly from 631 kJ mol⁻¹ (Sc) to 906 kJ mol⁻¹ (Zn) — a much shallower climb than across a typical s/p period. Why? Because the new electrons enter the inner 3d sub-shell, which screens the outer 4s electron quite well; the effective nuclear charge experienced by 4s rises only slightly.

The trend is irregular at d⁵ (Mn) and d¹⁰ (Zn): ionising Mn (which would expose its half-filled 3d⁵) costs less than expected after an inflection at Cr, while ionising Zn (4s² → 4s¹ leaving 3d¹⁰) costs more than expected.

Anomalies: the ΔᵢH₂ of Cr and Cu is unusually high — removing the second electron destroys the freshly-stable d⁵ (Cr⁺) and d¹⁰ (Cu⁺) configurations. The ΔᵢH₃ of Mn is unusually high — removing it from Mn²⁺ wrecks the half-filled d⁵.

8.7 Variable Oxidation States

Because (n-1)d and ns electrons are very close in energy, transition metals can lose different numbers of electrons in different reactions — giving an enormous variety of oxidation states.

8.7.1 Patterns to Remember

• The greatest range of oxidation states occurs in the middle of the series — Mn shows +2 to +7. Ends of the series show fewer states (too few electrons in Sc/Ti; too many in Cu/Zn for higher valence).
• Maximum oxidation state up to Mn equals (s + d) electrons: Ti⁴⁺ (3d⁰), V⁵⁺ (in VO₂⁺), Cr⁶⁺ (in CrO₄²⁻), Mn⁷⁺ (in MnO₄⁻).
• Beyond Mn the maximum oxidation state drops abruptly: Fe(II,III), Co(II,III), Ni(II), Cu(I,II), Zn(II).
• Within a group of d-block elements, heavier members favour the higher oxidation state — opposite to the p-block. Cr(VI) is a strong oxidant; Mo(VI) and W(VI) are stable and weakly oxidising.
• Very low oxidation states (0, -1) appear when ligands like CO, NO can π-accept electrons. Example: Ni(CO)₄ — Ni is in the 0 state.

8.8 Standard Electrode Potentials and Reactivity

The key data for the 3d series are tabulated below.

• Cu has positive E°(Cu²⁺/Cu) = +0.34 V — Cu cannot liberate H₂ from non-oxidising acids. The high ΔₐH and modest ΔₕydH of Cu are not balanced enough to give a negative potential.
• The Mn²⁺/Mn and Zn²⁺/Zn potentials are more negative than the smooth trend predicts — the products Mn²⁺ (d⁵) and Zn²⁺ (d¹⁰) are unusually stable.
• Cr²⁺ is a strong reducing agent: changing d⁴ (Cr²⁺) → d³ (Cr³⁺) gives a half-filled t₂g level. Mn³⁺ is a strong oxidising agent: changing d⁴ (Mn³⁺) → d⁵ (Mn²⁺) gives the half-filled d⁵.

8.9 Magnetic Properties

A substance is paramagnetic if it is attracted into a magnetic field — caused by unpaired electrons (each electron is a tiny magnet). It is diamagnetic if it is feebly repelled — all electrons are paired. Ferromagnetism (Fe, Co, Ni at room temperature) is an extreme cooperative form of paramagnetism.

For most 3d ions the orbital contribution is "quenched", so the spin-only formula gives the magnetic moment to good accuracy:

One unpaired electron gives μ = √3 = 1.73 BM. Five unpaired electrons (high-spin d⁵ like Mn²⁺) give μ = √35 = 5.92 BM.

8.10 Coloured Ions of the d-Block

Most d-block ions are coloured in the solid state and in solution. The colour arises because the five d-orbitals of a transition-metal ion split into two sets of slightly different energies under the influence of surrounding ligands. An electron in the lower set absorbs a photon of visible light to jump to the upper set — a "d-d transition" — and the colour we see is the complementary colour of the wavelength absorbed.

Configurations with no d-electrons (d⁰: Sc³⁺, Ti⁴⁺) and configurations with all d-orbitals filled (d¹⁰: Zn²⁺, Cu⁺, Cd²⁺) cannot undergo d-d transitions — they are colourless.

8.11 Catalytic Properties

Transition metals and their compounds are master catalysts. Two factors are responsible: (i) they can shuttle between multiple oxidation states and (ii) they offer surface sites at which reacting molecules can adsorb. Some industrially crucial examples:

A homogeneous example: Fe³⁺ catalyses the oxidation of I⁻ by S₂O₈²⁻:

The Fe³⁺/Fe²⁺ couple shuttles electrons faster than the direct (uncatalysed) reaction.

8.12 Complex Formation

Transition metals form an extraordinary range of complex compounds with ligands such as CN⁻, NH₃, H₂O, Cl⁻, CO, en (ethylenediamine). Three structural features make this possible:

• Small ionic size + high ionic charge – strong electric field draws ligands in.
• Vacant low-energy d-orbitals – room to accept lone pairs from ligands.
• Variable oxidation states – different ligand fields stabilise different states.

Examples: [Fe(CN)₆]³⁻, [Fe(CN)₆]⁴⁻, [Cu(NH₃)₄]²⁺, [PtCl₄]²⁻. The detailed treatment is in Chapter 9 (Coordination Compounds).

8.13 Interstitial Compounds and Alloy Formation

Interstitial compounds form when small atoms (H, B, C, N) get trapped in the holes between metal atoms in the lattice. Examples: TiC, Mn₄N, Fe₃H, VH₀.₅₆, TiH₁.₇. They are typically non-stoichiometric. Properties:

• Higher melting points than the parent metals.
• Extremely hard — some borides rival diamond (e.g., TiB₂).
• Retain metallic conductivity.
• Chemically inert.

Alloys are solid solutions of two or more metals. Because transition-metal radii sit within ~15% of each other and they share similar bonding, they form alloys readily. Examples: stainless steel (Fe + Cr + Ni), brass (Cu + Zn), bronze (Cu + Sn), nichrome (Ni + Cr).