Chemistry of the Transition Metals and Coordination Compounds

1. Chemistry of the Transition Metals and Coordination Compounds Valentim M. B. Nunes Engineering Unit – Chemistry Section

2. Introduction The transition elements are those that belong to the periods of the Periodic Table where the d and forbitals are progressively filled. They are characterized by possessing incomplete d orbitals or to give origin to ions with the d orbital not completely filled. This characteristic is responsible for the peculiar properties of these metals, including the great tendency to form complex ions.

3. Transition Elements

4. General Physical Properties Consider, as an example, the period from scandium to copper: Sc  Ti  V  Cr  Mn  Fe  Co  Ni  Cu We verify that an increase in the nuclear charge is compensated by the shielding effect, and as a consequence the atomic radius does not reduces must and the electronegativity and ionization energies are approximately constants. These elements possess a strong metallic bond ( high density, high melting and boiling points, and also high transition phase temperatures).

5. Electronic configuration Sc: Ar 4s2 3d1 Ti: Ar 4s2 3d2 . . . Cu: Ar 4s1 3d10 When ions are formed the e- are firstly removed from 4s orbitals and only then from d orbitals. Fe2+ : Ar 3d6

6. Oxidation States The oxidation states can change between +1 and +7 The higher values correspond to the oxides: V2O5, Mn2O7, ....

7. Metallic complexes The transition metals have a great tendency to form complexes. A metallic complex is an ion containing a central metallic cation, bonded to one or several molecules or ions. Co2+(aq) + 4 Cl-(aq)  CoCl42-(aq)

8. Stability constant The tendency to form a complex is measured by the correspondent formation constant, Kf, or stability constant, that is the equilibrium constant for the complexation reaction.

9. Example 0.2 moles of CuSO4 are dissolved in a liter of a solution of ammonia [NH3] = 1.20 M. What is the concentration of Cu2+ free in solution? The equilibrium constant is: The concentration is practically zero!

10. Coordination Compounds Complex ions can itself combine with simple ions or other complex ions to form coordination compounds. A coordination compound is a neutral specie that must have, at least, one complex ion. The knowledge about the chemical bond in this compounds is due to Alfred Werner (Nobel prize in Chemistry) who stated that this elements possess two types of valence: primary and secondary. This corresponds to the oxidation state and the coordination number of the element. Primary valence: 3 Co(NH3)6Cl3 Secondary valence: 6

11. Ligands We call ligands to the molecules or ions that surround the central metal in a complex ion. Ligands possess at least one valence electron pair not shared. In this sense they are Lewis bases. The chemical bonds are coordinate covalent bond, also known as dative bond.

12. Examples of ligands

13. Polydentate ligand (EDTA)

14. Some definitions The atom of a ligand that is directly connected to the metallic atom is called donor atom. We define coordination number as the number of donor atoms that surround the central metallic atom in a complex ion. The more common are 4 and 6 (but also 2 and 5). Fe(CN)63- : coordination number = 6 Ag(NH3)2+ : coordination number = 2 The ligands can be monodentate, bidentate or polydentate, accordingly with the number of donor atoms present. Bidentate and polydentate ligands are also called as chelating agents.

15. Oxidation states of metals in coordination compounds The oxidation state of the central metallic atom is important and can be calculated from the global charge of the complex ion. Ru(NH3)5H2OCl2 : oxidation state of Ru = +2 K4Fe(CN)6 : oxidation sate of Fe = +2 Contra ion is K+!

16. Practicing R:

17. Next slides about the nomenclature in coordination compounds only in Portuguese

18. Nomenclatura (Nomenclature) A designação sistemática dos compostos de coordenação obedece às seguintes regras: >> O nome do anião surge antes do catião (tal como nos compostos iónicos). >> Dentro do ião complexo, o nome dos ligandos surge em primeiro lugar, por ordem alfabética, e no final o nome do metal. >> os nomes dos ligandos têm terminação o se forem aniões ou não têm designação especial se forem neutros ou catiões, excepto H2O (aquo), CO (carbonilo) e NH3 (amino). >> Quando os compostos contêm vários ligandos iguais utilizamos os prefixos di, tri, tetra, penta e hexa. >>Se o ligando possui ele próprio um prefixo grego utilizamos os prefixos bis, tris e tetraquis. >> O número de oxidação é indicado em numeração romana a seguir ao nome do metal. Se o complexo é anião a terminação do nome é o nome do metal seguido de ato.

19. Names of common ligands (Portuguese)

20. Names of anions that contain metallic atoms (Portuguese)

21. Examples (Portuguese) K4Fe(CN)6 : hexacianoferrato(II) de potássio Co(NH3)4Cl2Cl : cloreto de tetraaminodiclorocobalto(III) Cr(en)3Cl3 : cloreto de tris(etilenodiamina)crómio(III) Cr(H2O)4Cl2Cl : cloreto de tetraaquodiclorocrómio(III) Hexanitrocobaltato(III) de sódio? Nitrato de diclorobis(etilenodiamina)platina(IV)?

22. Geometry The geometry of coordination compounds depends on the number of ligands surrounding the central atom. The more commons for monodentate ligands are: 2 4 4 6 linear tetrahedral square planar octahedral

23. Isomers The coordination number and geometry are determined by the size of the metallic ion, the size of the ligands and by electronic factors (electronic configuration). These compounds can have several isomers (an aspect that we will not develop).

24. Crystal Field Theory There are several theories to explain the chemical bond in coordination compounds. None of them is completely satisfactory. The more complete is the Crystal Field Theory, that allow us to explain the color and magnetic properties of many coordination compounds. The crystal field theory tries to explain the chemical bond in complex ions in terms of electrostatic forces: attraction between the metallic ion with positive charge and the ligands with negative charge (or polar ligand!) and the repulsion between the isolated pairs localized in the ligands and the electrons that occupy the d orbitals of the metal.

25. Octahedral compounds The octahedral geometry is the more common. The central metallic atom is surrounded by 6 pairs of e- (localized in the ligands): all the five d orbitals are repelled by the e- of the ligands. But the repulsion depends on the spatial orientation of the d orbitals.

26. Biological Systems: Hemoglobin The Fe2+ ion is coordinated with the nitrogen atoms of the heme group. The ligand bellow porphyrin is the histidine group connected to the protein. The sixth ligand can be the oxygen, O2. Octahedral coordination

27. Interaction metal-ligand The orbitals of the central atom that are directed to the ligands increase the energy (less stable) while the remaining decrease in energy (more stable).

28. Crystal Field Splitting The metal-ligand interaction causes that the five d orbitals in an octahedral complex suffers a splitting in two levels of energy. We call to this difference of energies the crystal field splitting energy, .

29. Color The color of a substance results from the fact that it absorbs light in a specific range of  in the electromagnetic spectrum corresponding to the visible region (400 to 700 nm) Cu(H2O)62+: absorbs light in the orange region, with máx = 51014 s-1 or   600 nm. The transmitted light is predominantly blue.

30. Relation between color and 

31. Calculation of  When the energy of a photon, given by h (Planck’s equation), its equal to the difference between the higher and lower levels of energy of the d orbitals:  = h Absorption occurs, i.é., promotion of an electron from the lower level to the higher level.

32. Spectrophotometry A máx = 498 nm 400 500 600 700 A = log I0/I = lc /nm

33. Instrumentation

34. Example The Ti(H2O)63+ ion absorbs light in the visible region (see slide 32) of the electromagnetic spectrum. The maximum absorption corresponds to the wavelength of 498 nm. Calculate the crystal field splitting energy, in kJ/mol.

35. Spectrochemical series With spectroscopic data from different complexes, with the same metal but different ligands, it is possible to establish a spectrochemical series: list of ligand by crescent order of capacity to produce the splitting of energies in the d orbitals. I- < Br- < Cl- < OH- < F- < H2O < NH3 < en < CN- < CO CO and CN- are stronger ligands , as they produce a high splitting of energies. Halogens and OH- are weaker ligands.

36. Magnetic Properties Fe3+   FeF63- Fe(CN)63- The value of  also determines the magnetic properties of a complex. For instance, Ti(H2O)63+ has only one valence electron so its always paramagnetic. However when several d electron exists: high spin low spin F- is a weak ligand , but CN- is a strong ligand (see figure 22.22 of Chang)

37. Tetrahedral and square planar complexes The crystal field theory allow also to explain the splitting of energies in other complexes with 4 ligands. In tetrahedral complexes the splitting is precisely the opposite of the octahedral compounds. The majority of complexes have high spin. The splitting in square planar complexes is more complicated: dx2-y2 dxy dz2 dxz, dyz

38. Applications Therapeutic applications as chelating agents in cancer treatment: cis-diaminodicloropaltina(II) Metallurgy: extraction of silver and gold by forming complexes with cyanates. Chemical analysis: EDTA Detergence: tripoliphosphates as chelating agents (environmental issues!).