1. General Chemistry I 1 BONDING IN TRANSITION METAL COMPOUNDS AND COORDINATION COMPLEXES 8.1 Chemistry of the Transition Metals 8.2 Introduction to Coordination Chemistry 8.3 Structures of Coordination Complexes 8.4 Crystal Field Theory: Optical and Magnetic Properties 8.5 Optical Properties and the Spectrochemical Series 8.6 Bonding in Coordination Complexes 8 CHAPTER General Chemistry I

2. 2 347 Emerald 3BeO∙Al 2 O 3 ∙ 6SiO 2 with some Al 3+ replaced by Cr 3+

3. General Chemistry I 3 8.1 CHEMISTRY OF THE TRANSITION METALS 348

4. General Chemistry I 4 - Decreasing radii for small Z transition atoms → Increase in Z eff - Increasing radii for large Z transition atoms → Increase in electron-electron repulsion 349

5. General Chemistry I 5  Lanthanide contraction: bad shielding by 4f orbitals → the radii of the 6 th period ~ the 5 th period → decrease in atomic and ionic radii by increasing Z along the 6 th period 349

6. General Chemistry I 6 350  melting point: function of the bond strength in solids - roughly correlated with the number of unpaired e -

7. General Chemistry I 7 351  Enthalpy of hydration of M 2+ ions M 2+ (g) → M 2+ (aq):  H hyd =  H o f (M 2+ (aq)) –  H o f (M 2+ (g)) Lowering of  H hyd from a line → due to crystal field stabilization Anomalies of Mn → due to the stable half-filled d shell

8. General Chemistry I 8  Oxidation states more common oxidation state  Increasing tendency toward higher oxidation states among heavier transition elements in the same group: Fe (2,3) → Ru (2,3,4,6,8), Ni(2,3) → Pd(2,4) 351

9. General Chemistry I 9  Hard and Soft Acids and Bases  Pearson (1963) ~ Extension of Lewis’ definition – electron pair acceptor (acid) and donor (base) – by adding categories ‘hard’ and ‘soft.’ ~ 'Hard' species: small, high charge states, low electronegativities, weakly polarizable ~ 'Soft' species: large, low charge states, high electronegativities, strongly polarizable ~ ‘Borderline’ species Ralph Pearson (US, 1919 - ) 353

10. General Chemistry I 10 354

11. General Chemistry I 11 HgF 2 (g) + BeI 2 (g) → BeF 2 (g) + HgI 2 (g) s/h h/s h/h s/s 354  Prediction of chemical reactivities of inorganic reactions ~ Preferred direction: hard acid/hard base or soft acid/soft base AgBr(s) + I – (aq) → AgI(s) + Br – (aq) s/b s s/s b EXAMPLE 8.2 Predict whether the following reactions will occur. (a)CaF 2 (s) + CdI 2 (s) → CaI 2 (s) + CdF 2 (s) (b)Cr(CN) 2 (s) + Cd(OH) 2 (s) → Cd(CN) 2 (s) + Cr(OH) 2 (s) NO YES

12. General Chemistry I 12 8.2 INTRODUCTION TO COORDINATION CHEMISTRY 355  Formation of Coordination Complexes  Werner’s investigation: Compound 1: CoCl 3  6NH 3 (orange-yellow) Compound 2: CoCl 3  5NH 3 (purple) Compound 3: CoCl 3  4NH 3 (green) Compound 4: CoCl 3  3NH 3 (green) Alfred Werner (Swiss,1866-1919) Nobel prize in chemistry(’13)  Treatment with HCl → did not remove NH 3 AgNO 3 + Cl - → AgCl(s) in the ratio of 3 : 2 : 1 : 0

13. General Chemistry I 13  Conductivity measurements: Compound 1: [Co(NH 3 ) 6 ] 3+ (Cl – ) 3 → Conductivity of Al(NO 3 ) 3 Compound 2: [Co(NH 3 ) 5 Cl] 2+ (Cl – ) 2 → Conductivity of Mg(NO 3 ) 2 Compound 3: [Co(NH 3 ) 4 Cl 2 ] + (Cl – ) → Conductivity of NaNO 3 Compound 4: [Co(NH 3 ) 3 Cl 3 ] → Nonelectrolyte → Concept of “coordination sphere” around the central metal ion inner and outer coordination sphere → Formation of an octahedral complex 356 In the above complexes, NH 3 and Cl - that are attached to Co are called LIGANDS

14. General Chemistry I 14 357 anhydrous CuSO 4 CuSO 4 ∙5H 2 O → [Cu(OH 2 ) 4 ]SO 4 ∙H 2 O

15. General Chemistry I 15  Monodentate ligands mono “one” and dens “tooth” 357

16. General Chemistry I 16  Bidentate ligands  Chelating ligands: chelate (G. chele, “claw”) [Co(EDTA)] – ~ 1 hexadentate [Pt(en) 3 ] 4+ ~ 3 bidentates 358 (‘en’)(‘ox’)

17. General Chemistry I 17  Naming coordination compounds 1) Single word for a coordination complex ~ [prefix-ligand-metal] 2) Cation first followed by anion ~ K[…] or […]Cl 3) Ending with the suffix “-o” for anionic ligand, chlorido (Cl), no change for neutral ligands except aqua (H 2 O), ammine (NH 3 ), carbonyl (CO). Note: “chloro” for Cl in a compound ligand 4) Prefixes for the number of ligands ~ di-, tri-, tetra-, penta-, hexa-, … (bis-, tris-, tetrakis-, … for ligands with di- (etc) in their names) 5) Alphabetical ordering for many ligands 6) Roman numeral (oxidation state) in (..) after the name of metal ~ […cobalt(III)]Cl or K[…ferrate(III)] anionic complex ions: the ending “-ate” 359

18. General Chemistry I 18 359

19. General Chemistry I 19  Ligand substitution reactions [Ni(OH 2 ) 6 ] 2+ (aq) + 6 NH 3 (aq) → [Ni(NH 3 ) 6 ] 2+ (aq) + 6 H 2 O 360

20. General Chemistry I 20  ‘Inert’ coordination complex: thermodynamically unstable, kinetically stable (inert)  ‘Labile’ coordination complex: thermodynamically unstable, kinetically unstable (labile) takes a week takes a matter of seconds 361 Difference between ‘inert’ and ‘labile’

21. General Chemistry I 21 8.3 STRUCTURES OF COORDINATION COMPLEXES 361 Octahedral complexes with geometrical isomers (complexes of type MA 2 B 4 (or MA 2 B 2 ; B is bidentate) cis-[Co(NH 3 ) 4 Cl 2 ] + trans-[Co(NH 3 ) 4 Cl 2 ] + cis-[CoCl 2 (en) 2 ] + trans-[CoCl 2 (en) 2 ] +

22. General Chemistry I 22  Octahedral complexes with mer / fac isomers (Complexes of type MA 3 B 3 ) mer-isomer: Similar ligands define a meridian of the octahedron fac-isomer: Similar ligands define a face of an octohedron -- all three groups are 90° apart. mer-Co(NH 3 ) 3 (Cl) 3 fac-Co(NH 3 ) 3 (Cl) 3 362

23. General Chemistry I 23  Tetrahedral complexes ~ Dominant for four-coordinate complexes ~ No geometrical isomers for tetrahedral complexes of MA 2 B 2 363  Square planar complexes ~ Au 3+, Ir +, Rh +, Ni 2+, Pd 2+, Pt 2+ ~ cis-[Pt(NH 3 ) 2 Cl 2 ] (anticancer drug, ‘cisplatin’) ~ trans-[Pt(NH 3 ) 2 Cl 2 ]  Linear geometry ~ Ions with d 10 configuration: Cu +, Ag +, Au +, Hg 2+

24. General Chemistry I 24  Optical isomers are molecules that rotate plane polarized light  Enantiomers (Gk.  ά τιος, “opposite”, and μέρος, “part or portion”) are optical isomers whose structures are non- superimposable mirror images (they lack reflection-rotation symmetry)  Chiral center (chirality [G. χειρ (kheir), "hand"] ~ handedness) is a central atom around which enantiomers are formed  A racemic mixture has equal amount of enantiomers (net rotation of plane polarized light = 0)  Chiral Structures 366

25. General Chemistry I 25 E.g. enantiomers of the [Pt(en) 3 ] 4+ ion E.g. enantiomers of all-cis [Co(NH 3 ) 2 (H 2 O) 2 Cl 2 ] + 366 Octahedral complexes of type MA 3 (A is bidentate) Octahedral complexes of type MA 2 B 2 C 2

26. General Chemistry I 26  EDTA (ethylenediaminetetraacetate) ion Hexadentate ligand, sequestering metal ions Antidote for lead poisoning, preserves freshness of oil 367

27. General Chemistry I 27 8.4 CRYSTAL FIELD THEORY: OPTICAL AND MAGNETIC PROPERTIES 367  Crystal Field Theory ~ Ionic description of metal-ligand bonds ~ Ligands are treated as point charges approaching the central metal ion Octahedral coordination complexes  Degeneracy of d-orbitals lifted into two groups :

28. General Chemistry I 28  Crystal Field Theory Ligands such as a halide or oxide are regarded as an electrostatic, point charge, or point dipole type, which set up an electrostatic field. A B  o = crystal field splitting energy metal d orbitalsspherical charges octahedral environment Cr 3+ 367

29. General Chemistry I 29 368

30. General Chemistry I 30 369 Fig. 8.17 An octahedral crystal field increases the energies of all five d orbitals, but the increase is greater for the d z and d x - y orbitals. 222

31. General Chemistry I 31 Electron configuration of octahedral complexes d 1 -d 3 by Hund’s rule 370

32. General Chemistry I 32 -From d 4 to d 7 octahedral complexes there are two possibilities, illustrated for d 4 (E.g. Mn(III) complexes) e-e repulsion low-spin configuration ligand-ligand repulsion If  o is large (strong-field ligands), t 2g 4 has a lower energy. : low-spin complex, minimum number of unpaired e - If  o is small (weak-field ligands), t 2g 3 e g 1 has a lower energy. : high-spin complex, maximum number of unpaired e - 370 high-spin configuration

33. General Chemistry I 33 Fig. 8.18. Electron configuration for (a) high spin (large  o ) and (b) low spin (small  o ) octahedral crystal field splitting energies for Mn(III) complexes Weak field configuration Strong field configuration H 2 O weak field ligand CN – strong field ligand 369 - Example: d 4 octahedral complexes of Mn(III)

34. General Chemistry I 34 370  Crystal Field Stabilization Energy (CFSE) The amount by which the (otherwise equal) energy levels for the d electrons of a metal ion are split by the electrostatic field of the surrounding ligands in a coordination complex.  Energy difference between electrons in an octahedral crystal field and those in the hypothetical spherical crystal field.

35. General Chemistry I 35  Square planar crystal field 370  sp > 1.6  0

36. General Chemistry I 36  Tetrahedral crystal field 371  t = 4/9  o

37. General Chemistry I 37 Fig. 8.20. Correlation diagram showing the relationships among d-orbital energy levels in crystal fields of different symmetries. 372

38. General Chemistry I 38  Magnetic properties  Paramagnetic compounds ~ One or more unpaired electrons ~ Large, positive magnetic susceptibility ~ Attracted by the magnetic field → “weigh” more ~ Prevalent among transition-metal complexes  Diamagnetic compounds ~ All of the electrons are paired ~ Small, negative susceptibility ~ Repelled by the magnetic field 373  Magnetic susceptibility ~ Strength of a sample’s interaction with a magnetic field

39. General Chemistry I 39 8.5 OPTICAL PROPERTIES AND THE SPECTROCHEMICAL SERIES 374  Transition-metal complexes ~ absorb visible light → colorful E.g. [Co(NH 3 ) 5 Cl] 2+ ion absorbs greenish yellow light (~530 nm) Only red and blue light transmitted → purple (complementary color)  Wavelength of the strongest absorption, max d 10 complex ~ colorless (no absorption, all d-levels are filled) High-spin d 5 complex ~ weak absorption (spin flip required)

40. General Chemistry I 40 Cr(CO) 6 [Co(NH 3 ) 5 (OH 2 )]Cl 3 K 3 [Fe(C 2 O 4 ) 3 ] K 3 [Fe(CN) 6 ] [Co(en) 3 ]I 3 Colors of the hexaaqua complexes of metal ions prepared from their nitrate salts. E.g. [Co(H 2 O) 6 ] 2+ 375

41. General Chemistry I 41  Spectrochemical series ~ An ordering of ligands according to their ability to cause crystal field splittings.  Spectrochemical series for ligands  Spectrochemical series for metal ions Mn 2+ < Ni 2+ < Co 2+ < Fe 2+ < Fe 3+ < Co 3+ < Mn 4+ < Pd 4+ < Ir 3+ < Pt 4+  Crystal field theory cannot explain the spectrochemical series! 376

42. General Chemistry I 42 8.6 BONDING IN COORDINATION COMPLEXES 377  Valence bond theory  dsp 3 hybrid orbitals ~ linear combination of one s, three p atomic orbitals and the d z2 atomic orbital ~ five equivalent new hybrid orbitals ~ trigonal bipyramid, PF 5, CuCl 5 –

43. General Chemistry I 43  d 2 sp 3 hybrid orbitals ~ linear combination of one s, three p atomic orbitals and d z2, d x2-y2 orbitals ~ six new hybrid orbitals ~ octahedron, SF 6 378

44. General Chemistry I 44  Molecular orbital theory  Ligand field theory ~ Failure of CFT and VB theories to explain the spectrochemical series ~ MO description for ligands  Construction of  MOs for octahedral complexes (of 1st row D-block metals) ~ Interaction between the metal 4s orbital with six ligand orbitals →  s and  s * orbitals ~ Interaction between three metal p orbitals with three ligand orbitals → triply degenerate  p and  p * orbitals ~ Interaction of the d z2 and d x2-y2 orbitals with ligand orbitals → a pair of  d and  d * orbitals 378

45. General Chemistry I 45 Fig. 8.27. Formation of  bonding MOs from overlap of metal and ligand orbitals. 379

46. General Chemistry I 46 Bonding MOs Nonbonding MOs 380 MO correlation diagram for octahedral Cr(III) complex ([CrCl 6 ] 3- ):  bonding only Antibonding MOs

47. General Chemistry I 47  Formation of  and  * bonds (1) Interaction between an empty metal d orbital with a filled atomic ligand p orbital. E.g. 3p orbitals of Cl – (2) Interaction between a filled metal d orbital with an empty ligand  * antibonding molecular orbital. E.g. CO, CN – → metal-to-ligand (M-L)  donation or  backbonding -  and  * MOs: M d orbital - L p orbital or M d orbital - L  * orbital 381

48. General Chemistry I 48 (3) Overlap of each of the metal nonbonding d xy, d yz, and d xz orbitals with four ligand p orbitals → Formation of three pairs of bonding and antibonding MOs, t 2g and t 2g *. Fig. 8.30. Bonding  MO by constructive overlap of a metal d xy orbital with four ligand p orbitals. 382

49. General Chemistry I 49  Order of bonding strengths for different ligands  Weak-field ligands (small  o ) → Overlap between occupied p(  ) bonding orbitals of ligands (Br –, Cl –, CO) with t 2g orbitals of metal → Increase in energy of t 2g and decrease in  o  Strong-field ligands (large  o ) → Overlap between unoccupied  * antibonding orbitals of ligands (CO, CN – ) with t 2g orbitals of metal → Lowering of energy of t 2g orbitals by  back-bonding (M→L)  Intermediate-field ligands ~ H 2 O, NH 3 383

50. General Chemistry I 50 Fig. 8.31. (a) (M  L) [or (b) (M  L)]  donation showing a reduction (or increase) in Δ o compared with that from  bonding alone. (a) Slight increase in energy of t 2g electrons (in t 2g * MOs) (b) Significant lowering in energy of t 2g electrons due to  back-bonding → Electrons of t 2g MOs are delocalised into unoccupied  *(L) 383 (a) (b)

51. General Chemistry I 51 Fig. 8.32. Effect of  bonding on the energy-level structure for octahedral coordination complexes.  Summary of the MO picture (Ligand Field Theory) of bonding in octahedral coordination complexes Cl -, Br - ligands e.g. [CrCl 6 ] 3– IIlustrated for V 2+,Cr 3+,Mn 4+ (d 3 ) CO, CN –, NO + Ligands e.g. Mn(CN) 4 384 H 2 O, NH 3 ligands e.g. [V(H 2 O) 6 ] 2+

52. General Chemistry I 52 10 Problem Sets For Chapter 8, 2, 8, 18, 26, 32, 44, 46, 58, 64, 66