Chemistry of transition elements

📚The chemistry

28.1 What makes a transition element

A transition element is a d-block element that forms one or more stable ions with incomplete d orbitals. In the first row that means titanium to copper. Scandium and zinc are in the d-block but are not transition elements: scandium’s only ion, Sc³⁺, has no d electrons (3d⁰), and zinc’s only ion, Zn²⁺, has a full 3d sub-shell (3d¹⁰). Copper qualifies because Cu²⁺ is 3d⁹.

d orbital shapes. The 3d_xy orbital has four lobes lying in the xy plane, pointing between the x and y axes (like a four-leaf clover set at 45° to the axes). The 3d_z² orbital has two lobes along the z axis with a ring (doughnut) around the middle in the xy plane.

The characteristic properties, and why:

• Variable oxidation states. The 3d and 4s sub-shells are close in energy, so different numbers of electrons can be lost (or used in bonding) for similar energy costs. Manganese shows every state from +2 to +7; vanadium +2 to +5; chromium commonly +3 and +6; iron +2 and +3.
• Catalysis. Transition elements and their compounds are good catalysts because they have more than one stable oxidation state (so they can accept and donate electrons, being oxidised and reduced in turn, as Fe²⁺/Fe³⁺ does in the I⁻/S₂O₈²⁻ reaction, topic 26) and vacant d orbitals that are energetically accessible, which can form dative bonds with reactant molecules and hold them on a surface.
• Complex ions. The same vacant, energetically accessible d orbitals can accept lone pairs from ligands, forming dative bonds.
• Coloured compounds — explained in 28.3 below.

28.2 Complexes and ligands

• A ligand is a species that contains a lone pair of electrons that forms a dative covalent bond to a central metal atom or ion.
• A complex is a molecule or ion formed by a central metal atom or ion surrounded by one or more ligands.
• The coordination number is the number of dative bonds from ligands to the central metal ion.

Ligands are classified by how many dative bonds each one forms:

Shapes

Large ligands such as Cl⁻ often give four-coordinate complexes, because fewer fit around the metal. With bidentate ligands, count the bonds, not the ligands: [Ni(en)₃]²⁺ has three ligands but coordination number 6.

Formula and charge: overall charge = charge on the metal ion + the sum of the ligand charges. Fe³⁺ with six CN⁻: 3 − 6 = −3, so [Fe(CN)₆]³⁻. Co³⁺ with six NH₃ (neutral): [Co(NH₃)₆]³⁺.

Ligand exchange: copper(II) and cobalt(II)

A ligand exchange reaction replaces some or all of the ligands in a complex with different ones. In aqueous solution the starting complex is the hexaaqua ion.

Copper(II), starting from blue [Cu(H₂O)₆]²⁺:

• + OH⁻ (NaOH): pale blue precipitate, insoluble in excess. \[ \mathrm{[Cu(H_2O)_6]^{2+} + 2OH^{-} \rightarrow Cu(OH)_2(H_2O)_4 + 2H_2O} \]
• + NH₃: a little gives the same pale blue precipitate (ammonia solution contains OH⁻); excess dissolves it to a deep (dark) blue solution: \[ \mathrm{[Cu(H_2O)_6]^{2+} + 4NH_3 \rightarrow [Cu(NH_3)_4(H_2O)_2]^{2+} + 4H_2O} \]
• + Cl⁻ (concentrated HCl): a yellow tetrahedral complex; mixtures with the blue aqua ion look green. Reversible on adding water. \[ \mathrm{[Cu(H_2O)_6]^{2+} + 4Cl^{-} \rightleftharpoons [CuCl_4]^{2-} + 6H_2O} \]

Cobalt(II), starting from pink [Co(H₂O)₆]²⁺:

• + OH⁻: a blue precipitate of cobalt(II) hydroxide, which turns pink on standing and slowly brown in air as it is oxidised. \[ \mathrm{[Co(H_2O)_6]^{2+} + 2OH^{-} \rightarrow Co(OH)_2(H_2O)_4 + 2H_2O} \]
• + NH₃: a little gives the blue precipitate; excess concentrated ammonia dissolves it to a pale yellow-brown (straw) solution, which darkens in air as Co(II) is oxidised to Co(III). \[ \mathrm{[Co(H_2O)_6]^{2+} + 6NH_3 \rightarrow [Co(NH_3)_6]^{2+} + 6H_2O} \]
• + Cl⁻ (concentrated HCl): a blue tetrahedral complex. Reversible; the pink colour returns on dilution. \[ \mathrm{[Co(H_2O)_6]^{2+} + 4Cl^{-} \rightleftharpoons [CoCl_4]^{2-} + 6H_2O} \]

Notice the change in coordination number and shape with chloride: six water ligands (octahedral) are replaced by four larger chloride ligands (tetrahedral).

Redox and titrations

Use E^⦵ values (topic 24) to predict whether a transition-metal redox reaction is feasible: it is if E^⦵_cell is positive. Three titrations are named in the syllabus:

• Manganate(VII) and iron(II), in acid; purple MnO₄⁻ is decolourised until the end-point, when a permanent pale pink appears (self-indicating). Ratio 1 : 5. \[ \mathrm{MnO_4^{-} + 8H^{+} + 5Fe^{2+} \rightarrow Mn^{2+} + 5Fe^{3+} + 4H_2O} \]
• Manganate(VII) and ethanedioate, in acid, warmed to about 60 °C (the reaction is slow at first, then catalysed by the Mn²⁺ it produces — autocatalysis). Ratio 2 : 5. \[ \mathrm{2MnO_4^{-} + 16H^{+} + 5C_2O_4^{2-} \rightarrow 2Mn^{2+} + 10CO_2 + 8H_2O} \]
• Copper(II) and iodide: Cu²⁺ oxidises I⁻ to iodine, forming a white precipitate of copper(I) iodide; the iodine is then titrated with sodium thiosulfate, using starch near the end-point (blue-black → colourless). \[ \mathrm{2Cu^{2+} + 4I^{-} \rightarrow 2CuI + I_2} \] \[ \mathrm{I_2 + 2S_2O_3^{2-} \rightarrow 2I^{-} + S_4O_6^{2-}} \] Overall ratio Cu²⁺ : S₂O₃²⁻ = 1 : 1.

Manganate(VII) titrations are acidified with dilute sulfuric acid, never hydrochloric acid, because MnO₄⁻ can oxidise Cl⁻ (E^⦵ +1.52 against +1.36 V).

28.3 Colour

In an isolated transition metal ion, the five 3d orbitals have the same energy: they are degenerate. When ligands approach, their lone pairs repel the d electrons, but not equally: orbitals pointing towards the ligands are raised in energy more than those pointing between them. The d orbitals split into two non-degenerate sets separated by an energy gap, ΔE:

• Octahedral complexes: two orbitals higher (d_z² and d_x²−y², which point at the ligands on the axes) and three lower.
• Tetrahedral complexes: the reverse — three higher and two lower, with a smaller gap.

Why complexes are coloured: an electron in a lower d orbital can absorb a photon and be promoted to a higher d orbital. The photon’s energy must equal ΔE, so only light of one particular frequency (ΔE = hf) is absorbed. The remaining light is transmitted or reflected, and we see the complementary colour of what was absorbed. [Cu(H₂O)₆]²⁺ absorbs in the red-orange, so it looks blue.

Ions with no d electrons (Sc³⁺, Ti⁴⁺) or a full d sub-shell (Zn²⁺, Cu⁺) have no possible d–d transition, so their compounds are colourless (white as solids).

Different ligands split the d orbitals by different amounts, so they change ΔE, the frequency absorbed, and the colour seen. Replacing water by ammonia around Cu²⁺ increases ΔE: a higher frequency is absorbed and the colour changes from pale blue to deep blue. Changes of coordination number and shape (octahedral to tetrahedral with chloride) change ΔE even more — pink to blue for cobalt. That is why every ligand exchange above comes with a colour change.

28.4 Stereoisomerism in complexes

Geometrical (cis/trans) isomerism occurs when two identical ligands can be next to each other (cis, 90° apart) or opposite (trans, 180°):

• Square planar [Pt(NH₃)₂Cl₂]: the cis isomer is the anticancer drug cisplatin; the trans isomer is inactive.
• Octahedral [Co(NH₃)₄(H₂O)₂]²⁺: the two water ligands cis or trans.
• Octahedral with bidentate ligands [Ni(en)₂(H₂O)₂]²⁺: the two water ligands cis or trans.

Optical isomerism occurs in complexes with no plane of symmetry, which exist as non-superimposable mirror images. The examples involve bidentate ligands, which wrap around the metal like propeller blades:

• [Ni(en)₃]²⁺ — three en ligands; two enantiomers.
• cis-[Ni(en)₂(H₂O)₂]²⁺ — the cis form is chiral and has two enantiomers; the trans form has a plane of symmetry and is not. So this complex has three stereoisomers in total.

Polarity: a complex is polar if the bond dipoles do not cancel. In trans isomers identical ligands are opposite each other, so their dipoles cancel and the complex is non-polar (trans-[Pt(NH₃)₂Cl₂]). In cis isomers they are on the same side, so the complex is polar (cisplatin).

28.5 Stability constants

The stability constant, K_stab, of a complex is the equilibrium constant for the formation of the complex ion in a solvent from its constituent ions or molecules. Water is the solvent, so [H₂O] is not included:

A large K_stab means the equilibrium lies far to the right: the complex is stable. Values are often quoted as log K_stab. In ligand exchange, a ligand forming a complex with a larger K_stab replaces one with a smaller K_stab. For copper(II): the ammine complex (log K_stab about 13) is far more stable than the chloro complex (about 5.6), and EDTA complexes are more stable still (about 18.8). Adding EDTA to almost any aqua or ammine complex therefore replaces the other ligands, which is why EDTA is used to treat metal poisoning and to measure metal ions by titration.