Unit 18: Transition Metals — Long Questions

Stereoisomers have the same molecular formula and the same bonds, but differ in the spatial arrangement of the ligands. Two types occur in transition metal complexes:

1. Geometrical (cis-trans) isomerism: arises when identical ligands can occupy either adjacent (cis, 90°) or opposite (trans, 180°) positions around the central metal ion. Example: square planar [Pt(NH3)2Cl2]. In cis-[Pt(NH3)2Cl2] (cisplatin) the two Cl ligands are adjacent (90°); their individual bond dipoles do not cancel, so the molecule is polar. In trans-[Pt(NH3)2Cl2] (transplatin) the two Cl ligands are opposite (180°); the identical bond dipoles point in opposite directions and cancel, making the molecule non-polar. This structural difference changes their biological activity — cisplatin is an anticancer drug, transplatin is not.

2. Optical isomerism: arises in octahedral complexes built with bidentate (chelating) ligands, where the molecule and its mirror image are non-superimposable, forming a pair called enantiomers. Example: [Ni(en)3]2+, where en = 1,2-diaminoethane (NH2CH2CH2NH2) is bidentate, each en coordinating through both its N atoms, giving coordination number 6 (3 en × 2 N) and octahedral geometry. The two non-superimposable mirror-image forms rotate the plane of plane-polarized light in opposite directions: the dextrorotatory (+) isomer rotates it to the right, the laevorotatory (−) isomer rotates it to the left.

Key structural difference geometrical isomerism results from different relative positions (cis/adjacent vs trans/opposite) of identical monodentate ligands on a square planar or octahedral centre, with no chirality involved; optical isomerism results specifically from the chirality created when bidentate/chelating ligands wrap around an octahedral centre, producing non-superimposable mirror images.

The general rule overall charge of a complex ion = (charge/oxidation state of the central metal ion) + (number of ligands × charge on each ligand), with the metal placed first in the formula, followed by the ligand(s) with their number as a subscript, all inside square brackets.

For [Co(NH3)6]3+ metal ion is Co3+ (oxidation state +3); ligand is NH3 (ammine), neutral, charge 0; coordination number = 6. Overall charge = (+3) + [6 × 0] = +3. This gives the cationic complex [Co(NH3)6]3+.

For [Fe(CN)6]3- metal ion is Fe3+ (oxidation state +3); ligand is CN- (cyano), anionic, charge −1; coordination number = 6. Overall charge = (+3) + [6 × (−1)] = +3 − 6 = −3. This gives the anionic complex [Fe(CN)6]3-.

Role of each factor the metal's oxidation state sets the starting positive charge; the ligand charge (0 for neutral ligands like NH3, H2O, CO; negative for anionic ligands like CN-, Cl-) is multiplied by the coordination number (how many ligands are attached) and added to the metal's charge. Neutral ligands leave the overall charge equal to the metal's charge (cationic complex), while sufficiently many anionic ligands can make the complex neutral or, as here, overall anionic.

Transition metals commonly exist in multiple oxidation states, and the feasibility of a redox reaction is predicted by comparing standard electrode potentials (E°) of the relevant half-reactions: E°cell = E°reduction − E°oxidation; the reaction is feasible if E°cell is positive. The species with the more positive E° is reduced; the species with the more negative (less positive) E° is oxidized. Three characteristic examples:

1. MnO4-/C2O4 2-: 2MnO4-(aq) + 5C2O4 2-(aq) + 16H+(aq) → 2Mn2+(aq) + 10CO2(g) + 8H2O(l), E°cell = +1.03 V (feasible). MnO4- (E° = +1.52 V) is reduced to Mn2+; oxalate (E° = +0.49 V as written for oxidation) is oxidized to CO2.

2. MnO4-/Fe2+: MnO4-(aq) + 5Fe2+(aq) + 8H+(aq) → Mn2+(aq) + 5Fe3+(aq) + 4H2O(l), E°cell = +2.29 V (feasible). MnO4- (E° = +1.52 V) is reduced to Mn2+; Fe2+ (Fe3+/Fe2+, E° = +0.77 V) is oxidized to Fe3+.

3. Cu2+/I-: 2Cu2+(aq) + 4I-(aq) → 2CuI(s) + I2(aq). Comparing standard E° values of Cu2+/Cu+ (+0.15 V) and I2/I- (+0.54 V) suggests this reaction is not feasible as written, since Cu2+ appears to be a weaker oxidising agent than I2. However, the reaction proceeds forward because insoluble CuI precipitates continuously, removing Cu+ from solution and shifting the equilibrium to the right (Le Chatelier's principle). The I2 produced can be quantified by titration with thiosulfate.

These examples show that redox feasibility is normally predicted from E° values, but can also be driven forward by removal of a product (e.g. precipitation) even when the simple E° comparison is unfavourable.

Stereoisomers are compounds with the same chemical formula and the same bonds, but differing in the spatial arrangement of their ligands. There are two types shown by transition metal complexes:

1. Geometrical (cis-trans) isomerism: shown by square planar and octahedral complexes containing two or more kinds of ligands. Identical ligands can be adjacent to each other (cis, 90°) or opposite each other (trans, 180°). Examples: square planar [Pt(NH3)2Cl2] — cis isomer (cisplatin: polar, used as anticancer drug) and trans isomer (transplatin: bond dipoles cancel so it is non-polar, not anticancer active); and octahedral complexes such as [Co(NH3)4(H2O)2]3+, where the two H2O ligands can be cis (adjacent, 90°) or trans (opposite, 180°) to each other.

2. Optical isomerism: shown mainly by octahedral complexes containing bidentate (chelating) ligands, e.g. [Ni(en)3]2+ where en = 1,2-diaminoethane. The complex and its mirror image are non-superimposable; such a pair is called enantiomers. Each enantiomer rotates the plane of plane-polarized light in opposite directions: the dextrorotatory (+) isomer rotates it to the right, and the laevorotatory (−) isomer rotates it to the left.