Coordination Chemistry: Structures & Isomers

1. Coordination Chemistry: Structures & Isomers • The history of coordination chemistry goes back only about 100 year, beginning with a dispute between Alfred Werner and Sophus Jørgensen. • This was a time of great strides in Organic bonding theory and many of these ideas were applied in general to transition metal compounds. • For example, bonding was thought to be directly related to an element’s “valence” (now: oxidation state). • Interestingly, this latter compound reacts with ammonia to form a complex containing four ammonia molecules: PtCl2(NH3)4 • Jørgensen thought that this complex would form like an Organic compound: Na+ + Br- → NaBr valence = 1 (so, 1 bond) Fe+3 + 3Cl- → FeCl3 valence = 3 (so, 3 bonds) Pt+2 + 2Cl- → PtCl2 valence = 2 (so 2 bonds) NH3—NH3—Cl Pt NH3—NH3—Cl called “Chain Theory” maintains the valence of 2

2. Coordination Chemistry • Werner examined a series of cobalt chloride complexes, noticing that their colors and reactivity with AgNO3 were very different. • Remember that Ag+ reacts only with chloride and not covalent chlorine. • So, in the first complex, the chlorines all act as chloride ions, in the second only two of the three do, while in the third and fourth only one behaves as chloride. The rest, presumably, would be covalent. • Jørgensen looked at these as Organic-like compounds: Color Complex Yellow CoCl3(NH3)6 + xsAg+ → 3AgCl ↓ Purple CoCl3(NH3)5 + xsAg+ → 2AgCl ↓ Green CoCl3(NH3)4 + xsAg+ → 1AgCl ↓ Violet CoCl3(NH3)4 + xsAg+ → 1AgCl ↓ NH3+−Cl- Co—NH3−NH3−NH3−NH3+−Cl- NH3+−Cl- NH3+−Cl- Co—NH3−NH3−NH3−NH3+−Cl- Cl all ionic not ionic

3. Werner Complexes • Werner recognized that in this series of compounds, cobalt could be viewed as maintaining a constant coordination of 6, rather than 3. • He formulated the salts as: • This could all have been explained uning Chain Theory. But, where this new coordination theory proved to be more effective was in predicting isomers. • Werner postulated that the bonds between metal and ligand were fixed in space in some symmetrical orientation. • For coordination of six, there are three possible arrangements, planar hexigon, trigonal prism, and octahedral. Color Complex Werner’s Prediction Yellow CoCl3(NH3)6+ xsAg+ → 3AgCl ↓ [Co(NH3)6] Cl3 Purple CoCl3(NH3)5+ xsAg+ → 2AgCl ↓ [Co(NH3)5Cl] Cl2 Green CoCl3(NH3)4+ xsAg+ → 1AgCl ↓ [Co(NH3)4Cl2] Cl Violet CoCl3(NH3)4+ xsAg+ → 1AgCl ↓ [Co(NH3)4Cl2] Cl

4. Werner Complexes. • Say that we have a complex, MA4B2. There are different possible numbers of isomers depending on the geometry. • Werner predicted that for coordination number = 6, octahedral geometry would be present; we now know this to be correct. • He also predicted that PtA2B2, for which two isomers were known, would be square planar and not tetrahedral.

5. Coordination Chemistry Nomenclature • In coordination complexes, the metal acts as a Lewis Acid (electron pair acceptor) and the surrounding species are called ligands and act as Lewis Bases (electron pair donors). ● ● M L

6. Survey of Coordination Numbers • Very uncommon; generally limited to low-valent metals. • e.g. Cu(I), Ag(I), Au(I), Hg(II); note filled d-subshell • always linear, e.g. [H3N—Ag—NH3]+ ; [Cl—Au—Cl]- • most easily react with additional ligands to form higher-coordinate complexes: [AuCl2]- + 2Cl- → [AuCl4]-3 • No transition metal examples. • many compounds appear to be 3-coordinate, but are actually of higher coordination. • e.g. FeCl3 is really Fe(sol)3Cl3 in solution and a dimer in the gas phase: Cl Cl Cl Fe Fe Cl Cl Cl

7. Survey of Coordination Numbers • Quite common; either tetrahedral or square planar. • Td; e.g. FeCl4-2, CoCl 4-2, CoX 4-2 • mainly anionic or neutral complexes; few cations adopt Td coordination. • Co(III) and Cr(III) are NEVER Td • Td coordination favored by steric requirements; large ligands and small metals. • do not exhibit geometric isomerism, but can be optically active. • Square Planar; characteristic of Pd(II), Pt(II), Rh(I), Ir(I), Au(III) and sometimes Cu(III) and Ni(II); almost no other examples. • normally d8 systems • Square Planar less favored sterically than Td complexes, especially when metal is small and ligand are large. • heavier metals are large enough to overcome steric requirements.

8. Survey of Coordination Numbers • b) Square Planar Geometic Isomers. • MA2B2 and MA2BC can exhibit cis- and trans- isomerization. • MABCD have three possible isomers. • bridged complexes. A B M B A A B M A B A B M C D A B M D C A D M C B Cl Cl P(Et)3 Fe Fe (Et)3P Cl Cl (Et)3P Cl Cl Fe Fe Cl Cl P(Et)3 note nomenclature cis bis-μ-chloro trans bis-μ-chloro

9. Survey of Coordination Numbers • Uncommon in general for metals. • square pyramidal geometry; often the metal lies above the square plane. • trigonal bipyramidal; most often all five ligands are the same. • on the basis of ligand repulsions alone, the trigonal bipyramidal structure would be favored. This is why non-metals adopt this geometry.

10. Survey of Coordination Numbers • Most common coordination number for metals. • Co(III), Cr(III), Fe(III) are always Oh. • Can have degrees of distortion to octahedral geometry. Trigonal distortion. Compression or elongation along 3-fold axis (i.e. face). Tetragonal distortion. Compression or elongation along 4-fold axis (Jahn-Teller). Twist to Trigonal Prism. Least common; observed for Re, Mo, W, V, Zr & Nb with ligands of the type: S R C C S R

11. Survey of Coordination Numbers • Geometric Isomers. MA4B2 MA3B3 Cl H3N NH3 Co H3N NH3 Cl Cl H3N Cl Co H3N NH3 NH3 cis trans Familiar? Two of Werner’s complexes. Cl Cl NH3 Co H3N NH3 Cl Cl H3N Cl Co H3N Cl NH3 mer meridinal fac facial Very few fac-mer isomer pairs known.

12. Survey of Coordination Numbers. • Quite rare. Usually only with large metal ions like Lanthanides and Actinides. • Three “Common” Structures: • Third most common geometry, but still uncommon. Usually large metals with small ligands. • major coordination for heavy metals; fairly common for Zr, Hf, Nb, Ta, Mo, W. • most structures based on the cube: monocapped octahedron capped trigonal prism pentagonal bipyramid dodecahedron (Chrysler symbol) square antiprism

13. Other Types of Isomerization. • Ligand Isomerization. • due strictly to metal coordinating to ligands that can have different isomers. • Ionization Isomerization. • occurs when two compounds with same chemical composition yield different ions in solution. e.g. [Co(NH3)4Cl2]NO3 vs. [Co(NH3)4ClNO3]Cl • subgroup: hydrate isomerization. NH2 NH2 CH2CH2CH2 NH2 NH2 CH2CH−CH3 1,3-diaminopropane 1,2-diaminopropane [Co(H2 O)6] Cl3 [Co(H2O)5Cl] Cl2 .H2O [Co(H2O)4Cl2] Cl.2H2O

14. Other Types of Isomerization. • Linkage (ambidentate) Isomerization. • occurs when a particular single ligand can bond multple ways. e.g. NO2- can bind either at the N or an O. [Co(NH3)5NO2]+2 [Co(NH3)5ONO]+2 e.g. SCN- either through the S or the N. [R2Pd(SCN)2] [R2Pd(NCS)2] N-bonded = nitro O-bonded = nitrito S-bonded = thiocyanato N-bonded = isothiocyanato

15. The Chelate Effect. • Compare the stability of a monodentate system with a bidentate one. Ni+2 + 6NH3 → [Ni(NH3)6]+2 K = 108.61 Ni+2 + 3en → [Ni(en)3]+2 K = 1018.28 • So, system with chelate rings is 1010 times more stable than monodentate system. But, why? • Entropy Effect. Ni(H2O)6+2 + 6NH3 → [Ni(NH3)6]+2 + 6H2O no change in entropy Ni(H2O)6+2+ 3en → [Ni(en)3]+2 + 6H2O large change in entropy ΔG = ΔH – TΔS; so, increase in entropy will result in increased stability. • Dissociation Effects • for [Ni(H2O)6]+2 say that a molecule of ammonia dissociates momentarily from the complex. What happens? It is carried away into the solvent. • for [Ni(en)3]+2 if one amine dissociates there is no way it can be washed away; it is still bound through second amine.

16. The Chelate Effect • Size of the chelate ring. • Compare: Cu+2 + 2en → [Cu(en)2]+2 K = 1020.03 Cu+2 + 2tn → [Cu(tn)2]+2 K = 1017.17 • tn ~1000 times less stable than en. • why? en more rigid, tn can “wag” • larger chelate ring reduces stability: 5 > 6 >>7 NH2 CH2 CH2 NH2 NH2 M CH2 NH2 CH2 CH2 M