Unit 18: Transition Metals — Short Questions

A transition element is defined as a d-block element that forms at least one stable ion with an incompletely filled d-subshell. Zinc forms only the Zn²⁺ ion, whose configuration is [Ar]3d¹⁰4s⁰ — a completely filled 3d subshell. Since zinc has no ion with a partially filled d-subshell, it is excluded from the transition metals even though it lies in the d-block.

The 3d and 4s subshells of transition metals have very similar (comparable) energies, so 3d electrons, like 4s electrons, can be removed relatively easily. A metal atom can lose only its 4s electrons (giving the common +2 state) or its 4s electrons plus a variable number of 3d electrons (giving higher oxidation states). This closeness in energy between 3d and 4s is what allows transition elements to show several different, variable oxidation states.

Transition metals catalyse reactions in two main ways. In homogeneous catalysis, they readily switch between stable oxidation states, forming an unstable intermediate with a reactant that then decomposes to give the product while regenerating the catalyst in its original oxidation state, lowering the activation energy — e.g. nickel salts catalyse the preparation of acrylic acid from acetylene via Ni(0)/Ni(II). In heterogeneous catalysis, their energetically accessible vacant d-orbitals allow reactant molecules to adsorb onto the metal surface in a favourable orientation, forming a temporary activated complex that lowers activation energy — e.g. solid iron catalyses the reaction of N2 and H2 to form NH3 in the Haber process.

en (1,2-diaminoethane) is a bidentate (chelating) ligand, forming two dative bonds per ligand molecule, whereas H2O is monodentate. Complexes formed with polydentate/chelating ligands have much higher stability constants than those formed with monodentate ligands (the chelate effect), so [Cr(en)3]3+, built from three chelating en ligands, is more stable than [Cr(H2O)6]3+, built from six monodentate water ligands.

A metal complex (coordination compound) is a molecule or ion formed when a central metal atom or ion is surrounded by a number of electron pair donor species (ligands). The bond between the metal and each ligand is a coordinate covalent (dative) bond, in which the ligand (acting as a Lewis base) donates a lone pair of electrons into an empty d-orbital of the metal ion (acting as a Lewis acid).

Coordination number is the total number of dative (coordinate covalent) bonds formed between ligands and the central metal atom/ion in a complex. In [Fe(CN)6]4-, Fe2+ is bonded to six CN- ligands, so its coordination number is 6.

A bidentate ligand is a species with two donor atoms, each carrying a lone pair of electrons, so it can donate two electron pairs and form two dative covalent bonds to the central metal atom/ion. Examples: 1,2-diaminoethane (en), where both N atoms donate, and the ethanedioate (oxalate) ion, where one O atom from each carboxyl group donates.

The stability constant (Kstab) is the equilibrium constant for the formation of a complex ion from its constituent metal ion and ligands in solution (water, being the solvent/initial ligand, is omitted from the expression). It is usually quoted as log10 Kstab. A higher log10 Kstab means a more stable complex, with the equilibrium lying far to the right (ligand substitution is favourable); a lower log10 Kstab means a less stable complex whose ligands can be easily displaced by other ligands.

For coordination number 4, the actual shape depends on the metal's d-electron configuration and the ligands involved. Metal ions with a d8 configuration (e.g. Ni2+, Pt2+) with ligands such as NH3 or CN- typically form square planar complexes (four ligands coplanar with the metal, 90° bond angles), e.g. [Ni(CN)4]2-, [Pt(NH3)4]2+. Metal ions with d7, d8 or d10 configurations commonly form tetrahedral complexes (four ligands at the corners of a tetrahedron, 109.5° bond angles), e.g. [CoCl4]2-, [MnCl4]2-. So the same coordination number can give either geometry depending on the metal's electronic configuration and ligand type.

[Pt(NH3)2Cl2] is a square planar complex that shows cis-trans (geometrical) isomerism. In cis-[Pt(NH3)2Cl2] (cisplatin), the two Cl ligands are adjacent to each other at 90°; the bond dipoles do not cancel, making the molecule polar. In trans-[Pt(NH3)2Cl2] (transplatin), the two Cl ligands are opposite each other at 180°; the identical bond dipoles cancel, making the molecule non-polar. Cisplatin is used as an anticancer drug, while transplatin shows no such activity.

Cisplatin (cis-diamminedichloroplatinum(II)) is used as an anticancer drug. Inside rapidly dividing cancer cells, its chloride ligands undergo ligand exchange with water; the 'activated' drug then binds to guanine bases in DNA, forming cross-links between the two strands of the double helix. This damage prevents DNA replication and triggers apoptosis (programmed cell death) of the cancer cell.

When excess ammonia solution is added to the blue cobalt(II) hydroxide precipitate [Co(H2O)4(OH)2](s), a ligand exchange reaction occurs: ammonia replaces the water and hydroxide ligands, forming soluble hexaamminecobalt(II) ions, and the colour changes from blue (precipitate) to yellow (solution):
[Co(H2O)4(OH)2](s) + 6NH3(aq) → [Co(NH3)6]2+(aq) + 4H2O(l) + 2OH-(aq)

Yes. Ligand exchange reactions are equilibria, so diluting a complex solution with a large (massive) excess of water shifts the equilibrium, by Le Chatelier's principle, back towards the left, re-forming the original hydrated (aqua) complex ion. For example, diluting a yellow [CuCl4]2- solution with a large amount of water regenerates the blue [Cu(H2O)6]2+ complex.

½I2(aq) + e⁻ → I⁻(aq), E° = +0.54 V
Cu²⁺(aq) + e⁻ → Cu⁺(aq), E° = +0.15 V

Comparing standard E° values, the I2/I- couple (E° = +0.54 V) is more positive than the Cu2+/Cu+ couple (E° = +0.15 V), meaning I2 is a stronger oxidising agent than Cu2+, and Cu+ is a better reducing agent than I-. This means the forward reaction (Cu2+ oxidising I- to I2, being itself reduced to Cu+) has a negative overall E°cell, so based on E° values alone this reaction appears not to be thermodynamically feasible as written.

Although the E° comparison predicts the reaction is unfavourable, it proceeds forward because the Cu+ formed combines immediately with I- to precipitate as insoluble CuI(s). This continuously removes Cu+ (and effectively drives consumption of Cu2+) from the solution, which, by Le Chatelier's principle, shifts the equilibrium to the right, pulling the reaction forward to completion despite the unfavourable calculated E° value.