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UniversityMicixxilms

International3 0 0 N. Z E E B R O A D . A N N A R B O R . Ml 4 8 1 0 6 IB B E D F O R D R O W , L O N D O N WC1 R 4 E J . E N G L A N D

8008802

Z i m m e r , L in n L a w r e n c e

TOTALLY SYNTHETIC IRON(II) HEME-PROTEIN MODELS AND THEIR INTERACTIONS WITH SMALL MOLECULES

The Ohio State University PH.D. 1979

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UniversityMicrdfiims

International2 0 0 \ Z = = = H O. A N N 4 R 3 0 P Ml J 8 1 0 6 ' 3 1 2 1 7 6 1 - 4 7 0 0

TOTALLY SYNTHETIC IRON(II) HEME-PROTEIN MODELS

AND THEIR INTERACTIONS WITH SMALL MOLECULES

DISSERTATION

Presented in Partial Fulfillment of the Requirements for

the Degree Doctor of Philosophy in the Graduate

School of The Ohio State University

By

Linn Lawrence Zimmer, B.S.

* * * * *

The Ohio State University

1979

Reading Committee:

Professor G, G. Christoph

Professor D. H. Busch

Professor A. Wojcicki

Approved By

C7 AdviserDepartment of Chemistry

To Christine Rea

ACKNOWLEDGEMENT

My deepest thanks are extended to all of the members of

Dr. Christoph's and Dr. Busch's research groups whose assistance and

encouragement helped make this work possible. Dr. Mark Beno was

especially helpful in the area of X-ray crystallography and Dr. Joseph

Grzybowski shared many valuable Insights with me in the area of

iron(II) chemistry. But, above all, special thanks must be given to

Professors Busch and Christoph for their guidance and encouragement

throughout the course of this work.

iii

CURRICULUM VITAE

March 4, 1953 Born, Detroit, Michigan

June, 1975 B.S., Wayne State University

Sept., 1975 -June, 1977 . Teaching Assistant, Department of Chemistry, The Ohio State University

June, 1977 -Nov., 1979 Research Associate, Department of Chemistry, The Ohio State University

December, 1979 Ph.D., The Ohio State University

FIELDS OF STUDY

Major Field: Chemistry

Specialization: Biological Aspects of Inorganic Coordination Chemistry.Professors G. G. Christoph and D. H, Busch, Advisers

iv

TABLE OF CONTENTS

PageACKNOWLEDGEMENTS................................................. H i

CURRICULUM V I T A E ................................................. iv

LIST OF T A B L E S ................................................... vii

LIST OF F I GURES................................................... x

ABBREVIATIONS........... xvi

INTRODUCTION ..................................................... 1

Requirements of a Model System ............................. 6Model Systems for Oxygen Binding . . . . .................. 7Kinetic and Equilibrium Studies for Carbon Monoxide

Binding to Model S y s t e m s .......................... 13Structural Studies of CO Adducts of Model Systems ......... 15Infrared Studies of CO Adducts of Natural and

Model Systems ................. 18Cobalt(II) Model Systems .................................... 20The Dry Cave M o d e l .......................................... 21

E X P E R I M E N T A L ..................................................... 29

General Procedures .......................................... 29R e a g e n t s ..................................................... 29Physical Measurements ........................................ 30Synthesis of Unbridged Nickel(II) Complexes ............... 32Synthesis of Monomeric Dry Cave Nickel(II) Complexes . . . 35Synthesis of Mixed Monomeric and Dimeric Nickel(II)

Dry Cave C o m p l e x e s ........................................ 38Synthesis of Dimeric Nickel(II) Dry Cave Complexes . . . . 41Synthesis of Ligand Salts ................................... 44Synthesis of Iron(II) Starting Materials ................. 49Synthesis of Unbridged Iron(II) Complexes ................. 50Synthesis of Iron(II) Dry Cave Chloro Complexes .......... 52Synthesis of Iron(III) Dry Cave Complexes ........ . . . . 55Synthesis of Other Iron(II) Dry Cave Complexes .......... 56Synthesis of Carbon Monoxide Adducts of the

Iron(II) Complexes ........................................ 59Synthesis of Dimeric Iron(II) Dry Cave Complexes ........ 63Synthesis of Copper(II) Dry Cave Complexes ............... 66

v

TABLE OF CONTENTS (Continued)Page

Equilibrium Constant Measurement ........................... 66X-Ray Crystallographic Procedures ........................... 68

RESULTS AND DISCUSSION .......................................... 80

Unbridged Nickel(II) Complexes . . . . ............. . . . 80Synthesis of Nickel(II) Dry Cave Complexes ................ 97Separation of Monomeric and Dimeric Complexes ............. 107Characterization of Monomeric Nickel(II)

Dry Cave C o m p l e x e s ........................................ 110Characterization of Dimeric Nickel(II)

Dry Cave C o m p l e x e s ........................................ 121Characterization of the Complex Derived from

9,lO-Bis(chloromethyl)anthracene 137Removal of Ligands from Nickel(II) ........................ 144Iron(II) Complexes .......................................... 154Monomeric Iron(II) Chloro Complexes of Dry

Cave Ligands, [FeLCl+ ] .................................... 155Attempted Synthesis of Four-Coordinate Iron(II)

Complexes, [FeL]2+ 160Iron(II) Complexes of Unbridged Ligands .................... 162Crystal Structures of Two Iron(II) Chloro Complexes . . . . 165Summary of Crystal Structure Results ...................... 174Reactions of Iron(II) Complexes with Axial Ligands . . . . 178CO Adducts of Iron (II) Dry Cave Complexes.................. 187Crystal Structures of a CO Adduct of an Iron(II)

Dry Cave C o m p l e x .......................................... 198Equilibrium Studies of the Reaction between Carbon Monoxide

and Monomeric Iron(II) Dry Cave Complexes................ 207Correlations Between Physical Properties of CO Adducts

and Equilibrium Constants................................. 224Reactions of Iron(II) Dry Cave Complexes with Oxygen . . . 235Iron(II) Complexes Having Rearranged Ligands ............. 242Iron(II) Complexes Derived from Dimeric Dry Cave Ligands . 251Monomeric Copper(II) Dry Cave Complexes .................... 268

APPENDIXES

A. Rotameter Calibration .................................... 270

B. Final X-ray Positional and Thermal Parametersand Structure Factors ................................. 272

C. Derivation of Chloride EquilibriumConstant Expression . . . . . 306

BIBLIOGRAPHY ..................................................... 310vi

LIST OF TABLES

Table Page1. Carbon Monoxide Binding by Hemoproteins and Model

Iron(II) Porphyrins .......................................... 14

2. Analytical Data for the Complexes rFe((R)Me2 [16]tetraeneN4}(B)(CO)](PF6)2 .................... 61

3. Summary of Crystallographic D a t a ........................... 72

4. E - S t a t i s t i c s ................................................ 73

5. Selected Infrared Frequencies and Molar Conductancesfor Unbridged Nickel(II) Complexes ........................ -82

6 . Proton NMR Data for Unbridged Nickel(II) Complexes . . . . 85

7. Carbon-13 NMR Data for Unbridged Nickel(II) Complexes . . . 87

8 . Electrochemical Data for the Unbridged Nickel(II)Complexes..................................................... 95

9. Selected Infrared Frequencies and Molar ConductanceData for Monomeric Dry Cave Nickel (II) C o m p l e x e s ......... 112

10. Orisager Plot Parameters for Monomeric Dry CaveNickel(II) Complexes ........................................ 114

11. Proton NMR Data for Monomeric Dry CaveNickel(II) Complexes ........................................ 116

12. Carbon-13 NMR Data for Monomeric Dry CaveNickel(II) Complexes ........................................ 118

13. Electrochemical Data for Monomeric Dry CaveNickel(II) Complexes ........................................ 122

14. Selected Infrared Frequencies and MolarConductances for Dimeric Dry Cave Nickel(II) Complexes . . 124

15. Proton NMR Data for Dimeric Dry CaveNickel (II) C o m p l e x e s ................................... 126

vii

LIST OF TABLES (Continued)Table Page16. Carbon-13 NMR Data for Dimeric Dry

Cave Nickel(II) Complexes .................................... 128

17. Electrochemical Data for Dimeric DryCave Nickel(II) Complexes ..................................... 132

18. Bond Distances for Dimeric Nickel(II) Complex ............... 134

19. Bond Angles for Dimeric Nickel(II) Complex . . . . . . . . 135

20. 13C N1IR Data for the Nickel (II) ComplexDerived from 9 ,10-BIs(chloromethyl)anthracene ............... 142

21. Selected Infrared Frequencies for Dry Cave Ligand Salts . . 147

22. Proton NMR Data for Dry Cave Ligand S a l t s ................... 1501323. C NMR Data for Monomeric Dry Cave Ligand S a l t s ........... 152

24. Bond Distances (esd) for a) [Fe{(m-Xylyl(NHEthi)2)-Me2 [16]tetraeneN^}Cl]Cl*2CHÔH and b) [Fe{(m-Xylyl- (MeNEthi)2)Me2 [16]tetraeneN4}Cl](PFg) ........................ 168

25. Bond Angles (esd) for a) [Fe{(m-Xylyl(NHEthi)2)Me2 [16]-tetraeneN^Cl]Cl^CHgOH and b) [Fe{(m-Xylyl(MeNEthi)2)2~ [16]tetraeneN4}Cl](PFg) ........................................ 169

26. Summary of Structural Data for Dry Cave Complexes . . . . 175

27. Carbon-13 NMR Data for CO Adducts of MonomericIron(II) Dry Cave Complexes, [Fe{(R)Me? [16]tetraeneN,}- (B)(CO)](PF6)2 ....................... f .. ....... 7 . . . 196

28. Bond Distances (esd) for [Fe{(l,5-Pent(NHEthi)2)Me2 [16]-tetraeneN4 )(PY)(CO)](PF&) •CHÔH ............................. 200

29. Bond Angles (esd) for [Fe{(l,5-Pent(NHEthi)_)Me„[16]-tetraeneN4>(PY)(CO)](PF6)2-CH3OH ..................... 201

30. Summary of Structural Data for CO Adductsof Model S y s t e m s .............................................. 203

viii

LIST OF TABLES (Continued)Table Page31. CO Equilibrium Constants for [Fe{(m-Xylyl(MeNEthi)2)-

Me2[16]tetraeneN/, }(CH^CN) ](PFg)2 in A c e t o n i t r i l e ....... 210

32. Equilibrium Data for the Reaction of the Para-xyleneBridged Iron(II) Complex with CO and Chloride .............. 216

33. CO Equilibrium Constants for Dry CaveComplexes at 0.0°C in C H ^ C N ................................... 222

1334. Summary of Infrared, Electrochemical and C NMR Data for CO Adducts of Monomeric Iron(II) Complexes, [Fe{(R)Me2 [16]tetraeneN4}(B)(C0)](PF6)2 ...................... 225

35. Carbon-13 NMR Data for Iron(II) Complexesof Rearranged Ligands .......................................... 247

ix

LIST OF FIGURESFigure Page

1. Structure of Fe ProtoporphyrinIX ........................... 3

2. Structure of the "Picket Fence" PorphyrinDioxygen Adduct ............................................ 11

3. Structure of the "Capped" Porphyrin ....................... 12

4. Structure of the Cyclophane Porphyrin .................... 16

5. Crystal Structure of [Ni{(MeOEthi)2Me2[15]-tetraeneN,}](C10,)_ . . . . . . . . . . . . . ............ 244 4 2

6 . ORTEP Drawing of ja-Xylyl Bridged Dry Cave Complex . . . . 26

7. Infrared Spectra of a) [Ni{(n-BuNHEthi)2ê2 [16]— tetraeneN^}](PFg)2 and b) [Ni{ (t-BuNHEthi)2Me2-[16] tetraeneN^.}] (PFg)2 ..................................... 83

8 . 1H NMR Spectra of a) [Ni{(n-BuNHEthi)2Me2 [16]-tetraeneN^}] (PFg)2 and b) [Ni{(_t-BuNHEthi)2Me2 [16]- tetraeneN^}](PFg)2 ......................................... 86

9. 13C NMR Spectrum of [Ni{(Me2NEthi)2Me2 [16]-tetraeneN^}](PFg)2 at a) 300K, b) 278 K, c) 238 K . . . . 89

10. 13C NMR Spectrum of [Ni{(n-BuNHEthi)2Me2 [16]-tetraeneN^}] (PFg)2 at a) 300 K, b) 238 K ................. 90

11. 13C NMR Spectrum of [Ni{(t-BuNHEthi)2Me2[16]-tetraeneN^}](PFg>2 ......................................... 91

12. Cyclic and Rotating Platinum Electrode Voltamagramsfor [Ni{(MeNHEthi)2Me2 [16]tetraeneN4}](PF6)2 ............. 96

13. ORTEP Drawing of meta-Xylyl Bridged Dimer ............... 102

14. Summary of Established Monomeric Nickel(II)Dry Cave Complexes............................................ 105

x

LIST OF FIGURES (Continued)Figure Page

15. Summary of Dimeric Nickel(II) Dry Cave Complexes . . . . 106

16. HPLC Chromatogram Showing Monomer-Dimer Separationfor the meta-xylene Bridge Nickel(II) Complexes . . . . 109

17. Infrared Spectra of a) [Ni{(m-Xylyl(NHEthi)2^Me2tetraeneN4)](PF6)2 and b) [Ni{(m-Xylyl(MeNEthi)2)Me2~ [16)tetraeneN^}](PFg)2 ................................... Ill

18. Onsager Plots for Monomeric Dry Cave Nickel(II)C o m p l e x e s ................................................ 113

19. Proton NMR Spectra of a) [Ni{(jj-Xylyl(NHEthi)2)Me2 t 16]-tetraeneN^)](PFg)2 and b) [Ni{(m-Xylyl(MeNEthi)2)Me2~ [16]tetraeneN^}J (PFg)2 ................. 117

20. 13C NMR Spectrum of [Ni{(m-Xylyl(NHEthi)2)Me2 [16]-tetraeneN^)](PFg)2 ....................................... 119

21. ^3C NMR Spectra of [Ni{ (jv-Xylyl (NHEthi) 2)Me2 [ 16]-tetraeneN^}](PFg)2 and b) [Ni{(m-Xylyl(MeNEthi)2)- Me2 [16]tetraeneN^}](PFg)2 ............................... 120

22. Infrared Spectrum of [Ni{(m-Xylyl(NHEthi)„)Me_[16]-tetraeneN^}]2 (PFg)/. ..................................... 123

23. Proton NMR Spectra of a) [Ni{(m-Xylyl(NHEthi)2)Me2 [16]-tetraeneN4)]2(PF6)4 and b) [Ni{(DURYL(MeNEthi)2Me2~ [16]tetraeneN4)]2(PFg)4 ................................. 127

24. 13C NMR Spectrum of [Nl{(m-Xylyl(NHEthi) j M e £16]-tetraeneN4 }]2 (PFg)4 ..................................... 129

25. 13C NMR Spectrum of [Ni{(DURYL(MeNEthi)oMe2 [16]—tetraeneN4)]2 (PFg)4 ..................................... 130

26. ORTEP Drawings of the Dimeric Nickel(II) Complex . . . . 133

27. Numbering Scheme for Dimeric Nickel(II) Complex . . . . 136

xi


28. Infrared Spectrum of the Anthracene Derivative ........... 138

29. Proton NMR Spectrum of the Anthracene Derivative ......... 1391330. C NMR Spectrum of the Anthracene D e r i v a t i v e ..............141

31. Infrared Spectrum of [(m-Xylyl(NHEthi)2)Me,,[16]- tetraeneN^](PFg)^ .......................................... 148

32. Si NMR Spectra of a) [(m-Xylyl(NHEthi)2)Me2 [16]“ tetraeneN^](PFg)^ and b) [ (£-Xylyl(NHEthi)2)Me2~[16]tetraeneN^](PFg)^ ...................................... 149

33. 13C NMR Spectra of a) [(m-Xylyl(NHEthi)2)Me2 [16]- tetraeneN^](PFg)^ and b) [(£-Xylyl(NHEthi)2)Me2~[16]tetraeneN^](PFg)^ ...................................... 151

34. Infrared Spectra of a) [Fe{(m-Xylyl(NHEthi)2)Me2 [16]-tetraeneNA}Cl]Cl*2CH30H, b) [Fe{(m-Xylyl(NHEthi)2)Me2~ [16]tetraeneN^)Cl](PFg), and c) [Fe{(m-Xylyl- (MeNEthi)2)Me2 [16]tetraeneN4}Cl](PF6) 157

35. Cyclic Voltamagram of [Fe{(m-Xylyl(MeNEthi)2)Me2~[16]tetraeneN4)Cl](PF6) .................................... 159

36. Infrared Spectrum of [Fe{(MeNHEthi)2Me2 [16]-tetraeneN,}(CH«CN) ](PFC)„ ................................. 1634 J X O Z

37. Proton NMR Spectrum of [Fe{(MeNHEthi)_Me_[16]-tetraeneN^}(CH3CN)x l(PFg)2 ................................. 164

38. Cyclic Voltamagram of [Fe{(MeNHEthi) Me_[16]-tetraeneN4>(CH3CN)x ](PF6)2 ......... 1 . 7 .................. 165

39. ORTEP Drawings of [Fe{(m-Xylyl(NHEthi)„)Me„[16]-tetraeneN4}Cl]Cl-2CH30 H ............... 166

40. ORTEP Drawings of [Fe{(m-Xylyl(MeNEthi)2)Me„[16]-tetraeneN4)Cl](PFg) ............. ,......................... 167

41. Molecular Numbering Scheme for Chloro-Iron(II) Complexes . 171

xii


42. Cyclic Voltamagram of [Fe{(1,6-Hex(MeNEthi)2)Me2~ [16]tetraeneN^}Cl](PFg) with a) No Excess Chlorideand b) 100 Equivalents of Excess Chloride . . . . . . . . 180

43. 3H NMR Spectrum of [Fe{(l,6-Hex(MeNEthi)„)Me_[16]- tetraeneN4>Cl] (PFg) ...................... 182

44. ^H NMR Spectrum of [Fe{(m-Xylyl(MeNEthi)9)Me„[16]- tetraeneN4}Cl] ( P F g ) ................................ 183

45. 13C NMR Spectrum of [Fe{(m-Xylyl(MeNEthi)„)Me„[16]-tetraeneN4>Cl](PF6) .....................7 . 7 .............. 184

46. Infrared Spectra of [Fe{(]3-Xylyl(NHEthi)2)Me2 [16]— tetraeneN4}(B)(CO)](PF6)2 , a) B *= CH3CN, b) B = PY,c) B = 1 - M e l m ................................................. 191

47. 13C NMR Spectrum of [Fe{(1,6-Hex(MeNEthi)„)Me„[16]-tetraeneN4)(CH3CN)(CO)](PF6)2 . . . . . . 7 193

48. 13C NMR Spectrum of [Fe{(l,6-Hex(MeNEthi)„)Me„[16]-tetraeneN4}(PY)(CO)](PF6)2 .............. 7 . 7 194

49. 33C NMR Spectrum of [Fe{(l,6-Hex(MeNEthi)„)Me„[16]-tetraeneN4}(l-MeIm)(CO)](PF6)2 .......... 7 . 7 195

50. ORTEP Drawings of [Fe{(1,5-Pent(NHEthi)9)Me„[16]- tetraeneN4 >(PY)(CO)](PF6)2 *CH3OH . . . 7 . 7 ............. 199

51. Molecular Numbering Scheme for [Fe{(1,5-Pent(NHEthi)„)- Me2 [16]tetraeneN4}(PY)(CO)](PF6)2*CH3OH ............. 202

52. Packing Diagram for [Fe{(l,5-Pent(NHEthi)9)Me„[163- tetraene^KCO) (PY) ] (PF6)2 -CH3OH . . . . . . 7 .......... 205

53. Spectral Changes for the Reaction of [Fe{(m-Xylyl- (MeNEthi) 2)Me2 [16] tetraene^} (CH3CN) ] (PFg) 2 with COin CH3CN at 0 ° C ............................................... 209

54. Van't Hoff Plot for [Fe{(m-Xylyl(MeNEthi)-) Me 2 [16]- tetraeneN4}](PF6)2 In CH3CN ............................... 211

xili


55. Spectral Changes for the Reaction of [Fe{(£-Xylyl- (NHEthi)2)Me2 [16]tetraeneN4)Cl](PFg) in CH^CN at0°C, [Cl“ ] » 1 x 10~3 m o l a r ..................................215

56. Plot to Determine Krl for the para-xyleneBridged Complex ............................................. 218

57. Spectral Changes for the Reaction of [Fe{(1,6-Hex- (MeNEthi)2)Me2 [16]tetraeneN4}Cl](PFfi) in CHgCN at0.0°C, [Cl"] = 0.1 m o l a r ...................................... 220

58. Cyclic Voltamagrams for [Fe{(1,6-Hex(MeNEthi)2)Me2[16]-tetraeneN.}(B)(CO)](PF.,) _, a) B = PY, solvent = CH_CN,

h b 2. jb) B = l-Melm, solvent = D M F ................................232

59. Electronic Spectra of the Chloro-iron(III) and Vi-oxo-dimer Derivatives of the (NMe)„Mxyl BridgedS p e c i e s ........................................................ 237

60. Spectral Changes for Reaction of [Fe{(1,4-But- (MeNEthi)2)Me2 (16]tetraeneN4}]I2 with 02 in CI^CN,1.5 H 1 — Melm, - 3 0 ° C ...........................................240

61. Infrared Spectrum of the Iron(II) Complex of theSexadentate Ligand .......................................... 243

62. Proton NMR Spectrum of the Iron(II) Complex of theSexadentate Ligand .......................................... 2451363. C NMR Spectra of a) the Iron(II) Complex of the Sexadentate Ligand and b) the CO Adduct of the Pentadentate Complex ........................................ 246

64. IR Spectrum of the Pentadentate CO A d d u c t .................. 249

65. Infrared Spectrum of the Pentadentatemeta-xylyl Complex .......................................... 251

66. Infrared Spectra of [Fe{(m-Xylyl(NHEthi)2)Me2 [16]-tetraeneN4 >(B)]2 (PF6)4 , a) B = PY, b) B - I m ................ 254

xiv


67. Cyclic Voltamagrams of [Fe{(m-Xyly 1 (NHEthi)2^Me2 t^]- tetraeneN4}(PY)]2 (PFg)4 in DMF with excess PY ............ 257

68. Titration of [Fe{(m-Xylyl(NHEthi)_)Me.[16]tetraeneN,} (PY)]2 (PF6)4 with Py with CH3CN 7 . 7 ................. 259

69. Titration of [Fe{(m-Xylyl(NHEthi)_)Me„[16]tetraeneN,} (Im)]2 (PF6)4 with Im in CH3CN . 7 . ................. 261

70. Titration of [Fe{(m-Xylyl(NHEthi)2)Me2 [16] tetraeneN^(2 - MeIm)]2 (PF6)2 with 2-MeIm in CH3C N .................263

71. Spectral Changes of [Fe{(m-Xylyl(NHEthi)2)Me2 [16]- tetraeneN4K l - M e I m ) 3 2 (PF6)4 with 02 in 50%H20/50%1-Melm . 265

72. Spectral Changes of [Fe{(m-Xylyl(NHEthi)2)Me2 [16]- tetraeneN4)(l-MeIm)]2 (PF6)4 with CO in 50%H20/50%1-Melm . 267

xv

ABBREVIATIONS

Me - Methyl

Mb - Myoglobin

Hb - Hemoglobin

1 - Melm - 1 - Methlimidazole

2 - Melm - 2 - Methylimidazole

Im - Imidazole

PY - Pyridine

TPP - Tetraphenylporphyrin

PpIX - ProtoporphyrinIX

PPp - Deuteroporphyrin

PPIXDME - Protoporphyrin IX Dimethyl Ester

TpivPP - Picket Fence Porphyrin

OMBP - Octamethylbenzoporphyrin

Pc - Phthalocyanine

TAAB - Tetramer of o-aminobenzaldehyde

xv i

INTRODUCTION

The role of inorganic chemistry in the functioning of biologi

cal systems is indeed important and widespread. Nature uses metal

complexes as the active sites in a number of enzymes and other proteins

having a variety of functions. Examples include hemoglobin and myo

globin, the oxygen transport and storage proteins; cytochromes, involved

in electron transport; the oxidase, oxygenase and peroxidase enzymes.

The study of these systems is made more difficult by the very size and

complexity of the molecules. In order to overcome some of the problems

associated with the study of biological systems, the field of bioinor

ganic chemistry has developed. Workers in this area synthesize and

study simple metal complexes which are designed to mimic the active

sites of the naturally occurring proteins. These model complexes

ideally allow for selective variation of structural features and, as a

result, the effects of structural changes on the reaction of interest

can be examined and evaluated in detail. The ultimate goal of model

system studies is two-fold: First, to more thoroughly understand the

natural systems; and second, to apply the knowledge to industrial and

laboratory processes in order to accomplish the same tasks as the living

systems, but on an industrial scale and perhaps with even greater effi

ciency or selectivity than the natural systems.

1

The work described In this thesis involves the design, synthesis

and characterization of a series of totally synthetic iron(II)

heme-protein models and the study of their interactions with the small

molecules dioxygen and carbon monoxide. The complexes have been

designed to mimic the behavior of the oxygen transport and storage pro

teins, hemoglobin (Hb) and myoglobin (Mb). The active site of these

proteins consists of a heme group (shown in figure 1) bound to the

globin at a single coordination site of the iron(ll) center through the*

imidazole group of the "proximal histidine." There are also some 80

hydrophobic interactions which aid in the binding of the heme to the

globin.^ Mb Is a monomer containing a single heme unit per protein,

whereas Hb is tetrameric, composed of two different globins known as

the a and 0 chains. The heme unit is located deep within the globin

resulting in the formation of a highly protected environment in the

vicinity of the metal, which appears to permit only small, rather

non-polar molecules to approach the metal when in the reduced state.

When the metal is in the oxidized state, having an unneutralized

charge, polar molecules can be drawn into the protected area.

The coordination chemistry of the iron(II) center has been2reviewed recently for simple hemes by Hoard and for hemoproteins by

3Perutz. Six-coordinate iron(II) hemes are Invariably diamagnetic

(S = 0). High-spin (S = 2) iron(II) hemes are always five-coordinate

with the metal ion displaced from the plane of the porphyrin toward the 4fifth ligand. A third, intermediate spin state (S *“■ 1) has no prece

dent in biological systems but has been observed in four-coordinate

Fe(TPP).^ In the deoxy Hb and Mb proteins, the heme is five-coordinate

3

H 0 2C(CH2)2

H02C(CH2)2

CH=CH,

i\//c h -c h 2

Figure 1. Structure of Fe ProtoporphyrlnIX

and high spin with the iron(II) displaced 0.55 A from the porphyrin 3plane toward the coordinated proximal imidazole. Upon addition of

oxygen or carbon monoxide, the iron(II) becomes six-coordinate and low

spin with the metal ion moving essentially into the porphyrin plane.

Another important feature of Hb and Mb is the presence of the

distal histidine whose imidazole group is positioned near to the vacant

binding site of the heme iron. The role of this group in the binding

of small molecules has been the subject of much debate and will be

referred to often throughout this work.

Until recently, the manner of binding of 0^ to Hb and Mb wasg

unknown. Pauling proposed a bent, end-on form of binding (structure I)

in which the M-0 bonding Is primarily sigma in character.

Griffith favored a triangular, side-on mode of 0^ coordination (struc

ture II) in which the bonding is primarily tt in character.

0I

•Fe

I- F e

lt

The appropriateness of the Pauling model was strongly supported byg

structures of cobalt-02 analogs and was confirmed by the X-ray crystal

structure analysis of the dioxygen adduct of Collman's "Picket Fence"9porphyrin. Two oxyheme proteins have recently been structurally

characterized^*^ and verify the bent, end-on form of bonding by

dioxygen. In oxy-erythrocruorln^ the Fe-0-0 angle was found to be

170° whereas in oxymyoglobin the angle was 121°.

The nature of the binding of carbon monoxide to Hb and Mb is

still the subject of debate. The X-ray structural analysis of HbCO has12been reported by Heidner et al. and a neutron diffraction study of

13MbCO has been reported by Norvell et al. Both of these groups report

that the CO oxygen atom lies off the heme axis by 0.7 A but that the CO

carbon atom cannot be observed. The cause of the bending of the bound

CO is interaction between the CO oxygen atom and the distal imidazole12of the protein. Heidner et al. calculated that the distances between

a linear CO molecule bound along the porphyrin axis and the distal

imidazole would be unacceptably short and they concluded that the

potential stress is relieved by the observed bending. A question which

remains, however^ is: Does the bending occur at the metal as in struc

ture III, at the CO carbon atom as in structure IV, or is there a com

bination of the effects given by structures III and IV as shown by

structure V?

/-Fe-

IB

in

cf y

-F.e-

B

nz;

/-Fe-

IB

JL

On the grounds that a bent Fe-C-0 conformation as in structures IV and

V has been observed only when the CO serves to bridge two metal centers

in simple organometallic compounds and that a bending similar to that14of HbCO was found in the structure of cyano-methemoglobin, it was

12concluded that structure III is most probable.

Another question which remains concerning the role of the distal

Imidazole is: What effect does the interaction between the bound CO

and the distal imidazole have on the stability of the resulting CO

adduct? In general, the binding constants for CO adduct formation with

iron(II) hemes are much larger than those observed with the heme pro

teins.^ It has been suggested*-'* that the weaker binding in the latter

is due to the interaction between CO and the distal imidazole pre

posed that this weakened binding of CO protects the heooproteins from

CO poisoning without affecting 0^ binding which, because of its intrin

sic angular geometry is not sterically affected by the distal imida

zole. The importance of steric interactions in decreasing the

iron-carbon bond strength is even more clear when one recognizes that

the principle source of CO in biological systems is endogenous. The

biological catabolism of Hb and Mb produces one mole of CO per mole of

heme. If no steric hinderance were present, one would expect approxi

mately 20% of the Hb and Mb to exist as CO bound forms with concommit-

tant impairment of vital bodily functions.

Requirements of a Model System

A synthetic system must meet a number of requirements in order

to serve as a functional model for Hb and Mb. First it must reversibly

form adducts with CO and 0^ (equations 1 and 2).

The second requirement is that the environment in the vicinity of the

CO and O2 binding site resemble that of the proteins. The site should

be well protected by a non-polar structure which inhibits dimerization

reactions and restricts the size and polarity of molecules which can

approach and bind to the metal. Third, the model should have coordina

tion and spin states which are like those of the proteins. The complex

15venting linear binding of the CO molecule. Suslick et al. have pro-

( , ij;Fet.Lj + o2 ^ ^ Q B jF e iL j i c o ;

(B)Fe(L) + CO (B)Fe(L) (CO) (1)

(2)

should change from hlgh-spin, five-coordinate to low-spin,

six-coordinate upon exposure to an appropriate small molecule ligand.

Fourth, because the natural system is an aqueous one, we require that

the model system likewise be water soluble. Fifth, the incorporation of

a structure which resembles the distal imidazole of the natural systems

is necessary if its importance in the binding of carbon monoxide to the

proteins is to be tested. Finally, the system must be sufficiently

synthetically versatile to allow the effects of systematic structural

changes to be examined and quantified.

Model Systems for Oxygen Binding

A wide variety of model dioxygen compounds have been studied by

various researchers using different techniques. The most basic systems

to be examined, the simple iron porphyrins, were found to bind oxygen

reversibly only at reduced temperatures, e.g., at -79°C.^ ^ Two

major difficulties limited the utility of these complexes as model sys

tems for Hb and Mb. At temperatures above -79°C, irreversible oxida

tion occurred, resulting in the formation of y-oxo-dlmers by the 21mechanism of equations 3-7:

FeP(B)2 v FeP(B) + B (3)

FeP(B) + 02 * FeP(B)(02) (4)

FeP(B)(02) + FeP(B) --- * FeP(B)-02~(B)PFe (5)

FeP(B)-02-(B)PFe ~ p-ld-»- 2FeP(B)(0) (6)

FeP(B) (0) + FeP(B) -rap-d->- FeP-O-FeP (7)

A number of oxo-bridged dimers have been studied in detail and are the22subject of a review. There are also a number of other simple oxida-

23tion reactions which irreversibly produce inactive ferric complexes.

The second problem encountered with simple porphyrin models was

the strong preference for iron(II) to be six-coordinate in the pres

ence of nitrogenous base ligands. Five-coordinate complexes are not24readily maintained in solution due to the fact that for the

consecutive equilibria shown in equation 8 , making FeP(B)2 the princi

ple species in solution.

K K2FeP + B * FeP (B) + B ^ ^ FeP(B)2 (8)

As a result, the reaction of simple ferrous porphyrins with CO and 0^

25has as rate determining a dissociative process in which a five—

coordinate intermediate is formed which then reacts rapidly with the 0^

or CO molecule (equations 9-11)

klFe(TPP) (B) „ * = = * Fe(TPP) (B) + B (9)1 -1

k 2Fe(TPP) (B) + 02 ■»-" “ Fe(TPP) (B) (02> (10)

k3Fe(TPP) (B) + CO » Fe(TPP) (B) (CO) (11)-3

Despite these difficulties kinetic and equilibrium studies on26the reaction of Fe(TPP)(B)2 with 02 and CO have been reported with

26some interesting results. It was shown that the affinity of the por

phyrins for CO was greater than for 02 and that the CO affinity was

greater than that observed for the proteins. It was also observed that

there is no large kinetic preference for oxygen compared with nitro

genous bases, a finding which suggests that if the distal imidazole of

the protein could bind to the iron(II), the protein would not function

as an effective oxygen carrier.

In order to overcome the preference of iron(II) to be

six-coordinate, sterically hindered ligands such as 2-methylamidazole

and 1,2-dimethylimidazole were used. These ligands form only

five-coordinate adducts with iron(II) porphyrins at room temperature,27but at low temperatures, six-coordinate species can be obtained.

28Hoard has reported the X-ray crystal structure of Fe(TPP)(2-MeIm)-EtOHO

and found the iron(II) to be high-spin and displaced 0.55 A from the

plane of the porphyrin in good agreement with the value observed for3deoxymyoglobin. It has further been shown that complexes having such

29hindered ligands are capable of reversibly binding dioxygen.

Another type of modification was used in Traylor's "tail base"30porphyrin which had an imidazoyl propyl side chain bonded to the

porphyrin ring through an amide linkage. This five-coordinate complex31was found to reversibly bind dioxygen in dichloromethane at -45°C.

At room temperature, however, irreversible oxidation occurred. "Tail

base" complexes have the advantage that the effects of changes in axial32base on the binding of O2 and CO can be very accurately quantified.

In order to overcome the problems due to dimer formation and

also to inhibit the formation of bis-base complexes, steric control of

the environment at the oxygen binding site is required. Baldwin and 33Huff synthesized an iron(II) macrocyclic species which included bulky

10

9,10-bridged-9,10-dihydroanthracene groups for which reversible oxygena-

tion was observed at -78°C. An Irreversible decomposition occurred at

temperatures above -50°C.

The same sterlc hlnderance idea was applied by Collman in theQ /

synthesis of the "picket fence" porphyrin which was one of the first

examples of an oxygen carrier which was functional at room temperature.

The complex interacted with two moles of l-Melm but only weakly with a

second due to steric interactions with the pivaloyl "picket" groups.33 34 36Complexes having one and two * l-Melm molecules bound to the

iron(II) have been isolated and studied. The most important feature of

the "picket fence" model has been the isolation, characterization and

X-ray crystal structural analysis of the solid dioxygen adducts

Fe(TpivPP) (1-Melm)(02)9 and Fe(TpivPP) (2-MeIm) (C>2) ,37 Although

four-fold disorder of the dioxygen and two-fold disorder of the axial

ligand limit the accuracy of the structural results, the bent, end-on

nature of the dioxygen binding is clearly evident, figure 2. In addi

tion, the structure of the five-coordinate, high-spin precursor complex,37Fe(TpivPP)(2-MeIm), is known, and consequently the structural changes

which occur upon binding of 0^ can be assessed directly. It was found

that the iron(II) moved nearly into the plane of the porphyrin and

changed from high- to low-spin upon oxygenation.

Baldwin has also developed model systems referred to as the

"capped" and "homologous capped" p o r p h y r i n s ^ ( f i g u r e 3) in which

a benzene ring is centered above the metal ion effectively forming a

hydrophobic cavity and inhibiting formation of bis-base and dimeric com

plexes. The compounds were found to reversibly bind at room

ch3h , C s ! / C H 3

H3 C ^ C

ch3CH-

CO ?h3I HgC | .CH3

HN __ CHa ch3H£ U cf>

€

Figure 2. Structure of the "Picket Fence" Porphyrin Dioxygen Adduct

Figure 3. Structure of the "Capped" Porphyrin

temperature, however, with affinities less than those observed for the

"picket fence" porphyrins. Baldwin attributed this behavior to an

increase in the conformational strain energy of the "capped" porphyrins41upon oxygenation, not to steric interaction between the bound 0^ and

the "cap."

A number of other model systems have been developed which

utilize different means for protecting the oxygen binding site. Wang

immobilized simple ferrous porphyrins in solid polymer films and

observed reversible oxygenation. Traylor extended this technique to2 q 42

ferrous porphyrins having appended axial ligands. Tsuchida and44Bayer have incorporated ferrous hemes into polymeric systems and

subsequently achieved reversible oxygenation. The irreversible

13

dimerization reaction was avoided in such systems because the iron(II)

centers are effectively isolated from each other and prevented from the

close approach, in much the same way as the four heme centers of Hb are

isolated from each other in the protein.

Kinetic and Equilibrium Studies for Carbon Monoxide Binding to Model Systems

A number of studies have been reported describing the kinetics

and equilibria of reactions between simple ferrous hemes and CO.45 45James has investigated the binding of CO to Fe(TPP), Fe(PpIX) and

Fe(OMBP),^ Stynes has reported data for Fe(Pc)^ and Fe(TAAB),^® and49Lcugee and Brault studied the Fe(PPD) system. The latter used a flow

technique in which gas mixtures having a known partial pressure of CO

were bubbled through the solutions. The other investigators introduced

a known partial pressure of CO over the solution and allowed time for

equilibration. Equilibrium constants for the reaction

FeP (B) + CO - K * FeP (B) (CO) + B (12)

possessed the following trend in the magnitude of K

PpIX > TPP > PPD > OMBP > Pc > TAAB

For these systems differences on the order of 10^ were observed as a

function of the in-plane ligand. Unfortunately, the usefulness of such

systems as model complexes is limited because the dissociative mecha- 24nlsm that controls the rates and equilibria of CO binding has no

counterpart in biological systems.

1A

Chang and Traylor**^ used flash photolysis methods utilized by

Gibson^ and rapid mixing techniques to study the rates of reaction of

CO and 0^ with the "tail base" porphyrin models in a number of solvents

(including water) and have examined a ferrous protoheme imidazole com-52plex in aqueous solution. Although the equilibrium constants for CO

binding were not reported, oxygenation rates and equilibria were found

to be very similar to those of the proteins.

The available data on carbon monoxide binding by hemoproteins

and model iron(II) porphyrin systems as summarized by Collman are

reproduced in table 1. It is apparent that only the porphyrins having

very weak or sterically hindered axial ligands have equilibrium con

stants similar to those of Hb and Mb. In particular, the binding con

stants for the "picket fence" porphyrin and Fe(PP^)(Im) are very large.

Collman'*' attributes these results to steric structural effects which,

as described above, cause a bending of the bound CO in the hemopro

teins but which are lacking In the model systems.

TABLE l1

CARBON MONOXIDE BINDING BY HEMOPROTEINS AND MODEL IRON(II) PORPHYRINS

Substance Pl/2C0 , torr

Mb, horse 1.8 x 10“ 2Hb, human 3.5 x 10“2Fe(PPD)(Im) 2 .A x 10“4Fe(PPD )(2-MeIm) A x 10"2Fe(PPD)(THF) A x 10~2Fe(TpivPP)(l-Melm) "very small"

15

The only systems containing structural features similar to the

distal imidazole of the proteins whose reactions with CO have been53investigated are the cyclophane porphyrins reported by Traylor. As

shown in figure 4, these complexes contain an anthracene or bridged

anthracene moity over the CO binding site of the metal, effectively

maintaining five-coordination of the metal in the absence of small

molecules. Preliminary kinetic results indicated the rate of CO

association to be 3 to 4 orders of magnitude less than that observed

in comparable systems not having steric hinderance. Thus it was con- 53eluded that steric hinderance is a controlling factor in the binding

of CO in these model systems.54 55Studies reported by Rougee and Wayland have demonstrated

the existence of both mono- and biscarbonyl complexes for Fe(TPP) and

Fe(PPp) in non-coordinating solvents. The consecutive equilibrium con

stants for the reaction (equation 13) differ by approximately two

orders of magnitude in both cases with

K KFeP + CO ' H 5-* FeP (CO) + CO y FeP (CO) 2 (13)

Structural Studies of CO Adducts of Model Systems

Despite the interest in the reactions between model systems and

carbon monoxide, a surprisingly small number of crystal structures of56CO adducts have been reported. Ibers reported the crystal structure

of Fe(TPP)(PY)(CO) in which the Fe-C-0 linkage was observed to be

linear and perpendicular to the porphyrin plane, in contrast to the

bend arrangement reported for the hemoproteins. The iron(II) was found

16

R= (ChÔ^nBu

Figure 4. Structure of the Cyclophane Porphyrin

17O

to be low-spin and displaced 0.02 A from the porphyrin plane towardO O

the CO. Fe-C and C-0 distances were 1.77 A and 1.12 A, respectively.57Although no structural details are available, Collman has

reported that in the structure of Fe(TpivPP)(l-Melm)(CO), the Fe-C-0

linkage is also strictly linear.58Goedkin has reported the structures of five- and

six-coordinate CO adducts of the ferrous complexes of the totally syn

thetic [14] annulene ligand VI. He has also determined the structure59of the ferrous CO adduct of the ligand VII.

YL

HN v K ^ l j K H

N N

JXYK

The range of C-0 bond distances for the complexes reported by GoedkinO O O O

is 1.12 A to 1.17 A and for the Fe-C bond distances, 1.69 A to 1.77 A.O

The distance of the iron from the plane ranges from -0.05 A to

18O

+0.30 A, where a negative value indicates displacement toward the other

axial ligand. In each case the Fe-C-0 angle is essentially linear and

along the macrocycle axis.

No model complex has yet been reported that satisfactorily

reproduces the bent Fe-C-0 linkage of HbCO and Mb CO. The reason is that

none of the structurally characterized models has an appropriately dis

posed functional group that can mimic the steric influence of the distal

imidazole.

Infrared Studies of CO Adducts of Natural and Model Systems

The CO stretching frequency (V^q ) and in some cases the band

width at half-height ôr a numêr heme-proteins and model

systems have been measured.^ Alben and Caughey^ reported on the

inductive and resonance effects of substituents on both the porphyrin

(cis effects) and the pyridine coordinated to the iron(II) (trans

effects) for some simple heme derivatives. They also examined solvent

effects and compared the results with those for HbCO. It was shown

that changes in the pK of the porphyrin caused by substitution at the

2 and A positions is negatively correlated with the CO stretching fre

quency. In a similar way, an increase in the pK of the coordinated

A-substituted pyridine was correlated with a decrease in V ^ . These

data can be explained in terms of the simple resonance contributors

VIII and IX. As the electron donating ability of a pyridine substitu

ent Increases, structure IX becomes more important because of increased

TT backbonding from the iron(II) center to the CO anti-bonding orbitals.

19

As a result, the strength of the Fe-C bond increases and that of the

C-0 bond decreases, resulting in a lower CO stretching frequency.

R— — Fa— c= ° R—(0 /N— Fe==c= 0

YEL IX

63Isotopic labelling experiments allowed the CO stretching fre

quency of human hemoglobin in aqueous solution to be accurately

assigned at 1951 cm It was observed that the half-band width in

HbCO was much narrower (8 cm than in simple heme carbonyls (27 cm ^).

A recent study on the interaction of CO with ferrous cytochrome P-450

concluded that an increase in ^v^/2 corresPonc s to "exposure of the

bound CO to a less homogeneous environment which may include some60external solvent." This same study concluded that the bending of the

bound CO is the cause of the decrease in vpn from that for the linearLUFe-C-0 of ferrous hemes.

For simple heme systems, is quite solvent dependent, as was

shown for Fe(PPIXDME)(B)(CO), B B PY or l-Melm.^ With pyridine as the

base, decreases from 1986 cm ^ in CCl^ (dipole moment = 0) to

20

1964 cm ^ in pyridine (dipole moment = 2.3). A similar trend was62observed when l-Melm was the base. It is therefore concluded that a

polar environment increases the strength of the Fe-C bond and causes a

decrease in63Caughey and Alben also examined some abnormal hemoglobins in

which the distal histidine residue was replaced with a tyrosine or

arginine residue. The effects on v_n for these proteins were large,

showing an increase in the value of and confirmed the importance61of the histidine residue in the binding of CO. It was also concluded

that there was no correlation between and the overall affinity of

the iron(II) in the protein for CO. V is determined to a greatuUextent by the strength of the Fe-C and C-0 bonds and does not reflect

ligand structural changes which occur upon formation of the adduct.

Thus caution must be exercised in relating CO stretching frequencies to

overall equilibrium constants for the binding of CO.

Cobalt(II) Model Systems

A large number of model systems have been developed in which

cobalt(II) was used as the metal ion Instead of iron(II). Reversible

oxygenation has been observed for non-macrocyclic complexes (e.g.,

salen derivatives^), for simple and modified porphyrin ligands^ and

for some totally synthetic macrocyclic c o m p l e x e s . A s an extensive

review^ has recently been published, only a few general comments about

the cobalt(II) systems will be mentioned here. In general, the

cobalt(II) systems for specific ligands are much more reversible in

their reactions with 0^ than are their iron(II) counterparts.

21

Cobalt(II) complexes offer the advantage that bis-base adducts do not

readily form. Irreversible oxidation processes similar to those of the

iron(II) complexes including peroxo-bridged dimer formation do occur

however. In general, the binding constants for Co(II) models are

several orders of magnitude smaller than those of comparable iron(II)

complexes. Cobalt(II) complexes do not form carbon monoxide adducts68except at high CO pressures and thus this aspect of the natural sys

tem chemistry cannot be examined with Co(II) models.

It Is apparent from this summary of model systems that a

totally satisfactory ligand structure has not yet been produced which

satisfactorily mimics the behavior of Hb and Hb in all critical

respects. In each of the examples described, one or more of the

requirements for a functional model system have not been met. Accord

ingly, a completely new synthetic ligand model system, the dry cave

complexes, was conceived and constructed.

The Dry Cave Model

In 1973, Busch et a l . ^ reported the synthesis, reactivity and

crystal structure of a novel methoxyethylidene complex of a fifteen69membered macrocycle (X). HIpp demonstrated the reactivity of this

complex with nucleophiles to introduce substituent groups on the peri

phery of the molecule. S c h a m m e l ^ * ^ extended the range of nucleophiles

used. The unusual methyl vinyl ether starting material (X) was shown to

react with primary and certain secondary amines according to equation

14. The significance and Importance of XI lies in the terminal nitro

gen functionalities which suggest and In fact permit a wide variety of

22

OCH

OCH

+ 2RNH2 CH3Cli (14)

substituents which can confer specific electronic and steric properties

on the complex. For example, good electron donating or withdrawing

substituents can be used to inductively control the oxidation potential

of the metal. Contributing resonance forms for these complexes are

shown by structures XII, XIII, and XIV, which together suggest signi

ficant electron delocalization through the tt system.

Certain steric properties can be imparted on the complex

through careful selection of the R group. Large B. groups such as

benzyl or naphthyl can create a non-polar environment in the vicinity of

one or both of the axial coordination sites, thus effectively blocking

the approach of ligands to the axial sites.

23

This class of compounds is of particular interest because they

serve as precursors to the bicyclic, dry cave type molecules as will be

shown in this work. The functional group R can be incorporated into

the structure at this stage of complex synthesis and its steric and

electronic effects will necessarily be present in the resultant dry

cave molecules.

Having demonstrated the general reactivity of the novel methyl

vinyl ether complex (X) with various mononucleophiles, Schammel extended

the reaction to a series of dinucleophiles.^ The crystal structure

analysis of [Ni{ (MeOEthi)2 162^ ^ ^ tetraene^4^ ^ 4^2 ^ ^ 6 ure showed the molecule to be in a distinct saddle shape and suggested that a

diamine of appropriate length could bridge the gap between the two

Figure 5. Crystal Structure of [Ni{(MeOEthi)_Me.[15]tetraeneN,}](ClO,)„L l if if 2

2569functional groups of the macrocycle. This reaction (equation 15)

resulted in the synthesis of the first dry cave compounds.

R

?* (15)

X V

72Christoph et al. reported the crystal structure of the

nickel(II) complex which had been bridged using para-xvlvlenediamine.

As shown in figure 6 , the structural analysis confirmed the bridged

nature of the complexes and gave important information about the size

and shape of the desired hydrophobic cavity.

The xylyl group is above the coordination site of the metal,

generating a rather hydrophobic environment and mimicking the steric

presence of the distal imidazole of the natural systems.

Schammel demonstrated the generality of this reaction using a

variety of dinucleophiles. He found that the ligands could be removed

intact from nickel(II) and chelated to cobalt(II) or the biologically

Figure 6 . ORTEP Drawing of jj—Xylyl Bridged Dry Cave Complex

27

more interesting iron(II) ion. Although it has recently been shown that

he was working with a dimeric molecule, as was discovered in this work,

Schammel demonstrated the reversibility of carbon monoxide binding to

one of the iron(II) complexes and also had promising results from

studies involving C^.

Stevens^ has synthesized and characterized a number of

cobalt(II) complexes derived from the dry cave ligands. The reactions

of these complexes with oxygen were studied in detail and a systematic

dependence of the oxygen binding constants as a function of bridging

group was found. As the bridging group was made smaller, the oxygen

affinity of the complex decreased, indicating repulsive interactions

between the bound and the bridge. The complexes proved to be excel

lent models for Hb and Mb, reversibly binding oxygen at ambient tem

peratures in aqueous solution with binding constants as large as or66 73greater than those of the natural proteins. *

It is apparent from the molecular design that the dry cave

model should be able to fulfill all of the above stated requirements of

a model oxygen carrier. The oxygen binding site is protected from the

solvent and from other complexes by the bulk of the bridging group so

that irreversible diraerization on that side of the molecule is pre

vented. The resultant cavity is small enough that five-coordination is

maintained in the absence of small ligand molecules. The coordination

numbers and spin states of the iron(II) center are identical to those

of the proteins. Since the complexes are salts, water solubility can

be readily achieved. Appropriate anion selection gives good solu

bility in polar solvents. The bridging group is designed in such a way

28

that it performs the suggested steric function of the distal imidazole

of the proteins. This model system is very versatile allowing major

and minor structural changes and thus facilitates systematic study of

the physical and chemical properties as a function of these variations.

There is one additional advantage of the dry cave model over

most other systems. The basic macrocycle is not a porphyrin but rather

is totally synthetic. As a result, reversible behavior with CO and 0^

does not simply derive from some characteristic of a porphyrin ligand,

but results from conditions we have newly created.

With the established need for improved Hb and Mb model systems

and the encouraging results of Schammel, the preliminary studies of

iron(II) dry cave complexes which he initiated were extended and pur

sued in much greater depth and constitute the bulk of this thesis.

One of the particular goals of this work was the examination of the

extent of steric contributions by the distal imidazole to the CO

binding. Through detailed spectroscopic and equilibrium studies of CO

adducts of dry cave complexes, and an X-ray structural study of a key

CO adduct, the exact nature of bonding and location of bending of the

bound CO molecule has been elucidated.

EXPERIMENTAL

General Procedures

All reactions and syntheses of nickel(II) complexes were carried

out in the open atmosphere (except if the exclusion of water vapor was

required, when a nitrogen blanket was used). All synthetic procedures

and manipulations of iron(II) complexes were performed in a Vacuum

Atmospheres glove box under an atmosphere of dry nitrogen gas con

taining less than 5 ppm of oxygen to prevent oxidation of air sensitive

materials.

Reagents

Solvents used in the reactions and characterization of

nickel(II) complexes were reagent grade and were used without further

purification. The solvents used in the synthesis and characterization74of iron(II) complexes were purified by the conventional means, dis

tilled under nitrogen and degassed under vacuum before use. Pyridine

was distilled from potassium hydroxide and N-methydimidazole was dis

tilled from barium oxide. Imidazole was recrystallized from hot ben

zene. All other reagents and chemicals employed were used as received

without further purification.

29

30

Physical Measurements

Elemental analyses were performed by Galbraith Laboratories,

Inc., Knoxville, Tennessee. Conductance measurements were obtained

using an Industrial Instruments, Inc., Model RC 16B Conductivity Bridge

and a cell having cell constant of 0.110 cm ^ at 1000 cycles per second.

Routine electronic spectra were measured on a Cary 17D Recording

Spectrophotometer using matched one cm quartz cells. Infrared spectra

were measured In the solid state asnujolmulls pressed between potas

sium bromide plates and in solution using matched demountable cells

having Irtran II windows and 0.5 mm pathlength using either a Perkin

Elmer Model 457 or 283B Infrared Spectrophotometer in the region from

4000 to 400 cm

Proton NMR spectra were obtained with a Varian Associates 360-L13spectrometer operating at 60 MHz. C NMR spectra were recorded on

either Bruker WP-80 or HX-90 spectrometers operating in the Fourier13transform mode at 20 or 22 MHZ, respectively. C NMR spectra spectra

were generally obtained using broadband proton decoupling as well as

off-resonance (CW) decoupling. Deuterated solvents were used through

out and chemical shifts were always assigned relative to an internal

TMS standard.

Magnetic susceptibilities were determined at room temperature by

the Faraday method using N i C e n ^ ^ O ^ and Hg[C0(SCN)^] as standards.

Pascal's constants^ were used to correct the molar susceptibilities for

ligand diamagnetism. Electrochemical measurements were provided by

Drs. Katherine Holter and Joseph Grzybowski of these laboratories and by

the author. The apparatus used was a Princeton Applied Research Corp.,

Potentiostat/Galvanostat Model 173 equipped with a Model 175 Linear

Programmer and a Model 179 Digital Coulometer. Current versus potential

curves were measured on a Houston Instruments Model 2000 X-Y Recorder.

All measurements were performed in a Vacuum Atmospheres Glove Box under

an atmosphere of dry, oxygen-free nitrogen. The working electrode for

voltametric curves was a platinum disk electrode, with potentials

measured versus a silver wire immersed in an acetonitrlle solution of

0.1 molar silver nitrate as reference. The working electrode was spun

at 600 rpm by a synchronous motor for the rotating platinum electrode

(RPE) voltamagrams. Peak potentials (Ep) were measured from cyclic

voltamagrams measured at 50 mV/sec scan rate. Half-wave potentials

(E1/2) were taken as the potential at one-half the height of the volta-

magram obtained using the RPE. The value of JE3 was usec* as a

measure of the reversibility of the couple and is also obtained using

the RPE. For reversible one electron couples,

Controlled potential electrolysis was performed using a plati

num gauze working electrode and a silver wire coated with silver

chloride as the reference. For all measurements, the solution con

tained 0.1 molar tetra-n-butylammonium tetrafluoroborate as supporting

electrolyte.

High performance liquid chromatography was performed using a

Du Pont Instruments Model 841 unit with a 50 cm C-18 reverse phase

column operating at 1000 psl. Ten microliters of an acetonitrlle solu

tion of the complex was eluted with a solution of 80% water/20% ace

tonitrlle. Chromatograms were recorded on a Varian Instruments Model

A-25 Strip Chart Recorder.

Synthesis of Unbridged Nickel(II) Complexes

(3,ll-Diacetyl-4,10-dimethy1-1»5,9,13-tetraazacyclohexadeca-

l >3,9,ll-tetraenatoN^)nickel(II), [Mi(Ac„He„[16]tetraenatoN^)]. This

complex was synthesized according to the published procedure.^

(2,12-Dimethyl-3,11-bis[1 - methoxyethylidene]-l,5,9,13-

tetraazacyclohexadeca-1,4,9,12-tetraenelQnickel(II) Hexafluorophos-

phate, [Hi{(MeOEthi)^Me^[16]tetraeneN^}] (PF^.)„.

(2,12-Dimethyl-3,11-bis [1-(amino) ethylidenel-1.5.9,13-tetraaza

cy clohexadeca-1 ,4,9,12-tetraenetQnickel(II) Hexafluorophosphate,

[Ni{(NHpEthi)2Me2 [16]tetraeneN,}](PFC)„.

(2,12-Dimethyl-3»11-bis F1-(dimethylamino)ethylidene]-l,5,9,13-

tetraazacyclohexadeca-l,4,9,12-tetraeneN^)nickel(II) Hexafluorophos

phate. [Nl{ (Me^NEthi) nMe2 ] tetraeneN^} ] ) 2 • The above three com-70plexes were prepared according to the procedures of Schammel.

(2,12-Dimethy1-3,11-bis[1-(methylamino)ethylidene3-1.5,9,13-

tetraazacy clohexadeca-1 ,4,9, 12-tetraeneN,)nickel(II) Hexafluorophos

phate, [Ni{ (MeNHEthi)nMe„[16]tetraeneN^}] (PF^.)„ . Methylamine gas was

bubbled through an acetonitrlle solution containing 14.0 g (19.7 mmoles)

of [Ni{(ME0Ethi)2Me2[16]tetraeneN^}](PFg)2 in 300 ml. The color of the

solution changed Immediately from yellow-green to deep red-orange.

After about 15 min, the gas bubbling was stopped and the solvent volume

reduced on a rotary evaporator to 50 ml. Methanol was added and the

volume was reduced further until crystals began to form. The solution

33

was refrigerated overnight to yield a yellow, highly crystalline pro

duct which was isolated and dried in vacuo. Yield: 13.2 g, (95%).

Anal. Calc, for N±c2oH34N6P2F121 C* 33,97» H * 4.85; N » 11.88. Found:C, 33.36; H, 5.10; N, 11.73.

(2,12-Ditnethy 1-3,11-bis [1- (tert-butylamino) ethylidene]-

1,5,9,13-tetraazacyclohexadeca-l.4,9,12-tetraeneN^)nickel(II) Hexa-

f luorophosphate, [Ni{ (t-BuNHEthi) ,Me„ [ 16 ] tetraeneEL } ] (PF^.) „. To a

solution of 3.0 g (4.2 nnnole) of [Ni{(MeOEthi)2Me2 [16]tetraeneN^}]-

(PFg)2 dissolved in 50 ml of acetonitrile was added 0.65 g (8.9 mmole)

of tert-butylamine. The color of the solution changed from yellow-green

to deep red. After stirring for one hour, the volume was reduced to

15 ml and methanol was added. Further volume reduction yielded the

yellow-orange powdery product which was isolated and dried under

vacuum. Yield: 1.7 g (51%). Anal. Calc, for N i C „ , H , , N , P „ F , C , 26 46 6 2 1239.46; H, 5.86; N, 10.62. Found: C, 38.79; H, 5.77; N, 10.28.

(2,12-Dlmethyl-3,11-bis[1-(n-butylamino)ethylidene]-1,5.9,13-

tetraazacy clohexadeca-1 ,4 ,9 , 12-tetraeneN^)nickel(II) Hexafluorophos

phate, [Ni{(n-BuNHEthi)2Me2 [16]tetraeneN^}](PF^)^. To a solution of

10.0 g (14.1 mmole) of [Nl{(Me0Ethi)2Me2 [16]tetraeneN^}](PFg)2 dis

solved in 250 ml of acetonitrile was added 2.3 g (31.4 mmole) of

n-butylamine. The color changed immediately from yellow-green to

orange. After stirring for 30 min, the volume was reduced to 50 ml and

methanol was added. Further volume reduction yielded the yellow micro

crystalline product which was isolated and dried in vacuo.

Yield: 8.8 g (79%). Anal. Calc, for Nic26H46N6P2F12: C, 39.46; H,5.86; N, 10.62. Found: C, 39.61; H, 5.88; N, 10.68.

(2,12-Dimethyl-3»11-bis[1-(benzylamino)ethylidene]-1,5,9.13-

tetraazacyclohexadeca-1,4,9,12-tetraeneN^.) nickel (II) Hexafluorophos-

phate, fNi{ (BZNHEthi^Me^f^ltetraeneN^}] (PF^.)„. To a solution of

3.0 g (4.2 mmole) of [Ni{(MeOEthi)2Me2[16]-tetraeneN^}](PFg)^ dis

solved in 250 ml of acetonitrile was added 1.0 g (9.3 mmole) of benzy-

lamine. The color of the solution changed from yellow-green to orange.

After stirring for one hour, the volume was reduced to 50 ml and 100 ml

of ethanol was added. The volume was again reduced until the powdery

yellow product had formed. This was isolated and dried in vacuo.

Yield: 2.8 g (77%). Anal. Calc, for N i C ^ H ^ N ^ F ^ : C, 44.73; H,

4.93; N, 9.78. Found: C, 45.07; H, 5.09; N, 9.96.

(2,12-Dimethyl-3.11-bis[1-n-butylaminoethylidene]-1,5,9,13-

tetraazacy clohexadeca-1 ,4,9,12-tetraeneN^)nickel(II) Iodide,

[Ni{(n-BuNHEthi)^He^[163tetraeneN^} ] . To a solution of 9.0 g

(11.4 mmole) of [Ni{ (n-BuNHEthi) 2 ^ e 2 [16] tetraeneN^}] (pFg) dissolved

in 250 ml of acetone was added dropwise a solution of 15.0 g

(40.7 mmole) of tetra-n-butylammonium iodide in 50 ml of acetone. The

yellow powder which formed immediately was isolated and dried in vacuo.

Yield: 8.45 g (98%). No analytical data were obtained for this

unbridged intermediate.

(2,12-Dimethyl-3,11-bis[1-(benzylamino)ethylidene-1,5.9,13-

tetraazacyclohexadeca-l,4,9,12-tetraeneN^)]nickel(II) Iodide,

[Ni{(BZNHEthi)„Me„[16]tetraenaN„}jl„. To a solution of 2.0 g -------- ;------ i-- z-------------- —Z

35

(2.3 mmole) of [Ni{(BZNHEthi)2^62[16]tetraeneN^}](PF^^ dissolved in

30 ml of acetone was added dropwise a solution of 3.5 g (9.5 mmole) of

tetra-rv-butylammonium iodide dissolved in 10 ml of hot acetone. The

powdery yellow precipitate was collected and dried in vacuo. Yield:

1.8 g (94%). Analytical data were not obtained for this unbridged

intermediate.

(2,12-Dimethyl-3,11-bis[1-(amino)ethylidene]-1.5.9.13-

tetraazacy clohexadeca-1 »4»9,12-tetraeneN^)nickel(II) Iodide,

[ N i { ( N H ^ E t h i ) [16]tetraeneN^}]1^. To a solution of 7.1 g (10.5

mmole) of [Ni^NJÊthi^f^tlS]tetraeneN^}](PFg)^ dissolved in 150 ml

of acetone was added dropwise a solution of 13.0 g (35.2 mmole) of

tetra-n-butylammonium iodide in 50 ml of acetone. The powdery yellow

product was collected and dried in vacuo. Yield: 6.6 g (98%). No

analytical data were obtained for this unbridged intermediate.

Synthesis of Monomeric Dry Cave Nickel(II) Complexes

(2,3,11,12,14.20-Hexamethyl-3.11,15,19,22,26-hexaazatricyclo-

[11.7.7.1^*^]octacosa-1,5,7,9(28),12,14,19,21,26-nonaeneN^)nickel(IX)

Hexafluorophosphate. [Ni{ (m-Xvlvl (MeNEthi^Hfen 1 tetraeneN^} ] (PF .) 2 •

To a solution of 5.0 g (7.1 mmole) of [Ni{(MeNHEthi)2Me2[16]tetraene-

N4 }](PF6)2 dissolved in 500 m3, of acetonitrile was added a solution

prepared by the reaction of 0.34 g (14.9 mmole) of sodium metal with

10 ml of methanol. The solution was heated to reflux and a solution

containing 1.87 g (7.1 mmole) of a fa''-dibromo-m-xylene dissolved in

250 ml of acetonitrile was added dropwise over a period of 4 h.

36

After the solution was cooled and filtered, the volume was reduced to

20 ml and the product was chromatographed on a column of neutral Woelm

alumina (25 cm x 2.5 cm) eluting with acetonitrile. The yellow hand was

collected, the volume reduced and ethanol added to yield the yellow

crystalline product which was collected and dried in vacuo. Yield:

4.3 g (75%). Anal. Calc, for N i C ^ H ^ N ^ F ^ : C, 41.55; H, 4.98;

N, 10.38. Found: C, 41.48; H, 5.02; N, 10.40.

(3,ll-Di-n-butyl-2.12,14,20-tetramethy1-3,11,15,19,22,26-

hexaazatricyclo[11.7.7. r* *9 ]octacosa-1,5,7,9.(28),12.14.19.21,26-

nonaenelQnickel(II) Hexafluorophosphate, [Ni{(m-Xylyl(n-BuNEthi)„)Me„-

[16]tetraeneN^}] (PF^.)„ . Three grams (4.0 mmole) of [Ni{ (nBuNHEthi)2ê2~

[16]tetraeneN^}]l2 was dissolved in 500 ml of methanol and heated to

reflux. To this solution was added a solution prepared by the reaction

of 0.19 g (8.3 mmole) of sodium metal with 10 ml of methanol. A solu

tion of 1.26 g (4.8 mmole) of a,a‘'-dibromo-m-xylene dissolved in 250 ml

of methanol was added very slowly to the first solution over a period of

5 h during which time the color changed from deep red to yellow-orange.

When the addition was complete, the volume was reduced to 100 ml and

3.5 g (21.5 mmole) of ammonium hexafluorophosphate in 25 ml of methanol

was added. An orange product formed and was collected. This product

was dissolved in 25 ml of acetonitrile and passed through a column of

neutral Woelm alumina (2.5 cm x 25 cm) eluting with acetonitrile. The

major yellow band was collected, the volume reduced and ethanol added.

Further volume reduction yielded the yellow-orange product which was

37

isolated and dried in vacuo. Anal. Calc, for N i C ^ H ^ N g P 2*12:

45.71; H, 5.87; N, 9.41. Found: C, 45.05; H, 5.97; N, 9.22.

(3,ll-Dibenzyl-2.12.14,20-tetramethyl-3.11,15,19,22.26-

hexaazatricyclo [11.7.7.1~**^ ]octacosa-1 ,5,7,9 (28) , 12,14,19,21,26-

nonaeneN,)nickel(II) Hexafluorophosphate, [Ni{(m-Xylyl(BZNEthi)„)-/s 2Me„[16]tetraeneN^}] (PF^.)„ . To a solution of 1.8 g (2.2 mmole) of

[Ni-CCBZNHEthiJ^MetlBJtetraeneN^}]]^ in 400 ml of methanol was added a

solution prepared by the reaction of 0.11 g (4.8 mmole) of sodium metal

with 10 ml of methanol. The solution became deep red in color and was

heated to reflux. A solution of 1.1 g (4.2 mmole) of a,a‘*-dibromo-m-

xylene dissolved in 250 ml of methanol was added dropwise to the above

solution. During the addition, the color of the solution became

yellow-orange. When the addition was complete, the solution was

cooled and the volume was reduced to 100 ml. Three and one-half grams

(21.5 mmole) of ammonium hexafluorophosphate in methanol was added to

yield a yellow-precipitate. The product was dissolved in 10 ml of

acetonitrile and passed through a column of neutral Woelm alumina

(2.5 cm x 10 cm) eluting with acetonitrile. The yellow band was col

lected, the solution volume was reduced and ethanol was added to yield

a yellow crystalline product which was isolated and dried in vacuo.

Yield: 1.0 g (47%). Anal. Calc, for NiC40H48N6P2F12: C, 49.97;H, 5.03; N. 8.74. Found: C, 49.63; H, 5.29; N. 8.71.

(2,3,10,11,13,19-Hexamethyl-3,10,14,18,21,25-hexaazablcyclo-

[10.7.7]hexacosa-l.11,13,18,20,25-hexaeneN^)nickel(XI) Hexafluonophos-

phate, [Ni((l,6-Hex(MeNEthi)„)Men[16]tetraeneN^}](PF^)„. This complex

38

was prepared in high yield by the method of Olszanski and Busch^

through the reaction of N,N'’-diraethyl-l,6-hexanediamine and the methyl66vinyl ether starting material as reported by Stevens.

2,3,11,12,14,20-Hexamethyl-3,11,15.19.22,26-hexaa2abicyclo-

[11.7.7]heptacosa-l,12.14.19,21,26-hexaeneN,]nickel(II) Hexafluorophos

phate , [Ni{ (1,7-Hept (MeNEthi) ,,)Me,, [16 ] tetraeneN^} ] (PF -) n » To a solution

of 8.0 g (11.3 mmoles) of [Ni{(MeNHEthi)_Me„[l6]tetraeneN.}](PF,)_ dis-Z / H D Zsolved in 500 ml of acetonitrile was added a solution prepared by the

reaction of 0.55 g (23.9 mmole) of sodium metal with 10 ml of methanol.

The solution was brought to reflux and 5.0 g (11.3 mmole) of 1,7-bis-

(para-toluenesulfonato)heptane dissolved in 250 ml of acetonitrile was

added dropwise over a period of six hours. When the addition was com

plete, the solution was evaporated to dryness and the residue was dis

solved in 25 ml of acetonitrile. After filtering through celite, the

solution was applied to a column (2.5 cm x 25 cm) of neutral Woelm

alumina and eluted with acetonitrile. The single yellow band was col

lected and the volume of the solution was reduced. Addition of ethanol

followed by further volume reduction resulted in formation of the yellow

orange product which was collected, washed with ethanol and dried in

vacuo. Yield: 5.4 g (59%). Anal. Calc, for NiC^yH^^NgP^F^ ^ •

C, 40.37; H, 5.77; N, 10.46. Found: C, 40.70; H, 6.10; N, 10.52.

Synthesis of Mixed Monomeric and Dimeric Nickel(II) Dry Cave Complexes

(2,12,14,20-Tetramethyl-3,11,15,19,22,26-hexaazatricyclo-

[11.7.7.1~* * ] octacosa-1,5.7,9 (28) , 12,14,19,21,26-nonaeneN^) nickel (II) -

39

Hexaf luorophosphate, [Nl{ (m-Xylyl (frHEthi) „)He^ [ 16 ] tetraeneN^, } ] (PF^ )

and (2,12,14.20.22,32,34,40-0ctamethyl-3.il,15.19.23,31.35,39.42,46.-50,54-dodecaazapentacyclo [31.7.7. 7 ^ 1~* Îhexapentaconta-

1,5,7,9(56),12,14,19,21,25,27,29(48).32,34,39.41,46,49,54-octadecaene)

dlnlckel(II) Hexaf luorophosphate, [Ni{ (m-Xyly (NHEthl) ,,)He^[16j tetraene-

Nj,}]„(PFc.)j, . Ten grams (14.1 mmole) of [Ni{(MeOEthi)2Me2 [16]tetraene-

N^}](PFg)2 was dissolved in 750 ml of acetonitrile. To this solution

was added a solution of 1.92 g (14.1 mmole) of m-xylylenediamine dis

solved in 750 ml of acetonitrile over a period of six hours. The yel

low solution was rotary evaporated to 25 ml while maintaining a

temperature of no greater than 30°C in the heating bath. The solution

was applied to a column (2.5 cm x 18 cm) of neutral Woelm alumina and

was eluted with acetonitrile giving a single yellow band. The solution

volume was again reduced at 30°C and ethanol was added to yield the

yellow product which was isolated and dried in vacuo. Yield: 6.8 g

(62%).

Separation of Monomeric and Dimeric Products. The above pro

duct was shown by HPLC to be a mixture of two products (later shown to

be monomer and dimer) which were separated as follows: Six grams (7.7

mmole) of the mixture was dissolved in 100 ml of acetonitrile. To this

solution was added 500 ml of ethanol causing a slight cloudiness. The

solution was refrigerated for 30 h to yield a yellow precipitate which

was collected and dried in vacuo. Yield: 3.21 g (54%). This function

was shown by HPLC to be primarily the dimeric species. Anal. Calc, for

40

[NiC26H36N6P2F12]2 i C, 39.97; H, 4.64; N, 10.76. Found: C, 39.85;

H, 5.00; N, 10.83.

The volume of the filtrate from above was reduced by approxi

mately one-third to yield a yellow precipitate which was collected

after cooling in a refrigerator overnight. The product was washed with

ethanol and dried in vacuo. Yield: 2.5 g (42%). This was shown by

HPLC to be a very pure sample of the monomeric species. Anal. Calc, for

NiC26H36N 6P2F12: C, 39.97; H, 4.64; N, 10.76. Found: C, 40.29;

H, 4.84; N, 10.84.

If the above reaction is carried out in refluxing acetonitrile

rather than at room temperature, the product is essentially pure monomer

based on HPLC analysis.

(2,11,13,19-Tetramethyl-3,10,14.18,21,25-hexaazatricyclo

[10.7 ♦ 7.2~* * octacosa-1,5,7,11,13,18,20,25,27-nonaeneN^) nickel (II)

Hexaf luorophosphate, [Nj{ (p-Xylyl (NHEthl) »)ê2 £3-61 tetraeneN^} ] (pF^) ~

and (2,11,13.19,21,30,32,38-Octamethyl-3.10,14,18,22,29,33.37.40,44.

47,51-dodecaazapentacyclo [29.7.7.7^ * 2~* * . 2 ^ * ] hexapentaconta-

1,5,7,11,13,18.20,24.26.30.32,37,39.44,46,51.53,55-octadecaene)

dinickel(II) Hexafluorophosphate, [Nl{(p-Xylyl(NHEthl)2)Me„[16]

tetraeneN^}] 2C ^ ^ ) ^ » These complexes were prepared by the method of78Callahan and Busch. To a suspension of 2.36 g (11.3 mmole) of

ja-xylylenediaminedihydrochloride in 100 ml of ethanol was added 0.55 g

(23.9 mmole) of sodium metal. After stirring for two hours, the sodium

chloride was filtered off and the filtrate was diluted to 500 ml with

acetonitrile. A second solution was prepared containing 8.0 g

(11.3 mmole) of [Ni{(Me0Ethi)2Me2 [16]tetraeneN^}l(PFg)2 dissolved in

41

500 ml of acetonitrile. These two solutions were added simultaneously

to 500 ml of refluxing acetonitrile over a period of 7 h. The resulting

yellow solution was evaporated to dryness and the residue was dissolved

in 30 ml of acetonitrile. The solution was applied to a column of

neutral Woelm alumina (2.5 cm x 25 cm) and was eluted very slowly with

acetonitrile. The product was collected as two fractions from the

column, the first as a yellow-green band and the second as a broad

yellow band. The volumes of both fractions were reduced and ethanol

added to yield the solid yellow products. Proton NMR showed fraction 1

to be primarily the dimeric species and fraction 2 the monomer. Yield

of dimer: 0.7 g (8%). Yield of monomer: 3.1 g (35%). Anal: Calc.

for n 1C26H36N6P2F12: C* 39*97; H> 4 *64; N » 10*76* Found: C, 40.15;H, 4.50; N, 10.96. When this synthesis was carried out at room tem

perature rather than at reflux, the yields of monomer and dimer were

more nearly equal but the overall yield was about the same.

Synthesis of Dimeric Nickel(II) Dry Cave Complexes

(2.3,10.11.13.19.21,22,29.30,32,38-Dodecamethyl-3,10,14,18,22,-

29,22,37,40,44.47.51-dodecaazapentacyclo[29.7.7.712,20.25,8.224>271-

hexapentaconta-1,5,7.11,13.18.20.24.26.30.32,37,39.44.46.51.53.55-octa-

decaene)dinickel(II) Hexafluorophosphate, [Ni{(p-Xylvl(MeNEthi)„)Me„[161

tetraeneN^}]„ (FF^)•• a solution of 6.70 g (9.5 mmole) of

[Ni{(MeNHEthi)„Me_[16]tetraeneN.}](PF,)_ dissolved in 500 ml of ace- Z Z H b Ztonitrile was added a solution prepared by the reaction of 0.46 g

(20.0 mmole) of sodium metal with 10 ml of methanol. A second solution

of 2.50 g (9.5 mmole) of a.a^-dibromo-ja-xylene dissolved in 500 ml of

42

acetonitrile was prepared. The two solutions were added simultaneously

to 500 of refluxing acetonitrile over a period of 7 h. The resulting

solution was treated in the same manner as in the preparation of

[Ni{(m“Xylyl(MeNEthi)2)Me2[16]tetraeneN^}3(PFg)2 to yield the yellow

product which was recrystallized from acetone. Yield: 5.2 g (68% .

Anal: Calc, for fN±c28H40N6P2F12]2: C * A1'56* H » 4 *98* N * 10.38.Found: C, 41.90; H, 5.37; N, 10.18.

(2,6.7.11.13,19,21.25,26,30.32,38,53,54,55,56-Hexadecamethy1-

3,10.14,18,22,29.33,37.40144,47.51-dodecaazapentacyclo 12 20 S ft 2 A 27[29.7.7.7 > .2 lhexapentaconta-1.5,7.11.13.18.20,24,26.30.-

32,37.39,44,46,51,53.55-octadecaene)dinickel(II) Hexafluorophosphate,

[ N i { ( D u r y l ( N H E t h i ) ^ t e t r a e n e N ^ }] 0 ^ ) ^ • To a solution of 2.0 g

(3.1 mmole) of [Ni{(NH2Ethi)2Me2[16]tetraeneN^}]l2 dissolved in 250 ml

of methanol was added a solution prepared by the reaction of 0.16 g

(7.0 mmole) of sodium metal with 10 ml of methanol. The solution was

heated to reflux and a solution of 0.80 g (3.5 mmole) of 3,6-bis-

(chloromethyl)durene (Pfaltz and Bauer) dissolved in 100 ml of

dichloromethane and 50 ml of methanol was added dropwise over a period

of 4 h. When the addition was complete the volume of the solution was

reduced to 100 ml and 5.0 g (30.7 mmole) of ammonium hexafluorophos

phate in methanol was added dropwise to yield the yellow product.

Addition of ethanol resulted in the formation of additional product

which was collected and dried in vacuo. Yield: 0.9 g (35%). Anal.

Calc, for [NiC30H44N6P2Flo]: C, 43.03; H, 5.30; N, 10.04. Found:

C, 42.70; H, 5.55; N, 10.14.

A3

(2.3.6.7.10.11.13.19.21.22.25.26.29.30.32,38.53.54.55.56-

Eicosamethy1-3,10,14,18»22.29,33,37.40.44.47,51-dodecaazapentacyclo 12 2Q 5 8 24 27[29,7.7.7 * .2 * .2 » ]hexapentaconta-l,5,7,11.13,18,20,24,26.30,

32,37,39.44.46,51,53,55-octadecaene)dinickel(II) Hexafluorophosphate,

[Ni{ (Duryl(MeNEthi)^)Me„ [16]tetraeneN^}] (PF^.)^. To a solution of

3.7 g (5.2 mmole) of [Ni{(MeNHEthi)2Me2 [16]tetraeneN4)] dissolved

in 250 ml of acetonitrile was added a solution prepared by the reac

tion of 0.24 g (10.4 mmole) of sodium metal with 10 ml of methanol. The

solution was heated to reflux and a solution of 1.20 g (5.2 mmole) of

3,6-bis(chloromethyl)durene dissolved in 250 ml of acetonitrile was

added dropwise over a period of 12 h. The solution was cooled and the

white sodium chloride was filtered off. The volume of the solution was

reduced to 100 ml and 4.0 g (24.5 mmole) of ammonium hexafluorophos

phate in 100 ml of ethanol was added. Further volume reduction resulted

in formation of the yellow product. Recrystallization from an aceto

nitrile ethanol mixture yielded the yellow crystals which were collected

and dried in vacuo. Yield: 4.2 g (93%). Anal. Calc, for

[NiC32H48N 6P2F12]2 : C, 44.41; H, 5.59; N, 9.71. Found: C, 44.57;

H, 5.72; N, 9.53.

Product of the Reaction Between [Ni{(MeNHEthi)2Me2 [16]tetraene

N4 )](PF .)2 and 9,10-Bis(chloromethyl)anthracene. To a solution of

2.0 g (2.83 mmole) of [Ni{(MeNHEthi)2Me2 [16]tetraeneN3}](PF6>2 dis

solved in 1 £ of acetonitrile was added 0.14 g (6.09 mmole) of sodium

which had been reacted with 10 ml of methanol. The solution was heated

to reflux and 0.85 g (3.1 mmole) of 9,10-bis(chloromethyl)anthracene

44

(Pfaltz and Bauer) was put into a soxhlet extractor above the solution.

The extraction was continued for 12 h. The resultant solution was

rotary evaporated to dryness and the residue was taken up in 25 ml of

acetonitrile and filtered. This solution was applied to a column of

neutral Woelm alumina (7.0 cm x 8.0 cm) and eluted with acetonitrile.

Volume reduction and addition of ethanol resulted in product formation.

Large red crystals were obtained by recrystallization from an acetoni

trile ethanol mixture. Anal: Calc, for NIC, „H,-„NrtP„F„ „: C. 48.45;— — tfU o / ±zH, 5.08; N, 11.30. Found: C, 48.88; H, 5.22; N, 11.06.

The two acetonitrile molecules of crystallization were

removed by grinding .ne crystals to a powder and drying in vacuo.

Anal: Calc, for N i C ^ H ^ N ^ F ^ : C, 47.55; H, 4.88; N, 9.24. Found:

C, 47.49; H, 5.08; N, 9.31.

Synthesis of Ligand Salts

(2,3,11,12.14,20-Hexamethyl-3,11,15,19,22,26-hexaazatricyclo- 5 9[11.7.7.1 * loctacosa-1.5.7,9(28),12.14,19,21,26-nonaene) Tetrachloro-

zincate, [(m-Xylyl(MeNEthi)„)Me„ [16]tetraene](ZnCl,,)„. Hydrogen

chloride gas was bubbled through a solution of 1.5 g (1.9 mmole) of

[Ni{(m-Xylyl(MeNEthi)2)Me2 [16]tetraeneN^}](PFg)2 dissolved in 50 ml of

acetonitrile until the solution turned blue. Slow addition of a solu

tion of tetrachlorozlncate anion (prepared by reaction of 0.75 g (11.5

mmole) of granular zinc metal with hydrogen chloride gas in 50 ml of

acetonitrile) resulted in formation of a white precipitate. This ligand

salt was filtered, washed with acetonitrile and ether and dried

45

in vacuo. Yield: 1.40 g (78%). Anal: Calc, for C ^ H ^ N g Z n ^ C l g :

C, 38.26; H, 5.04; N, 9.56. Found: C, 39.34; H, 5.63; N, 10.33.

(2,12,14,20-Tetramethyl-3,11,15,19,22,26-hexaazatricyclo-

[11.7.7.1~** octacosa-1,5,7,9(28).12,14,19,21,26-nonaene) Tetrachloro-

zincate, [ (m-Xylyl(NHEthi)QMen [I6]tetraenel (ZnCl^)„. The same pro

cedure was used as described in previous syntheses using 3.0 g (3.8

mmole) of [Ni{ (m-Xylyl(NHEthi)2Me2 [16] tetraeneN^ ] (PF6)2 and 1.5 g

(23.1 mmole) of zinc metal. No precipitate formed upon addition of the

tetrachlorozincate solution. The solution volume was reduced and fresh

acetonitrile added several times resulting in the formation of the

white, granular ligand salt which was collected, washed with acetoni

trile and dried in vacuo. Yield: 2.25 g (69%). Anal: Calc, for

C26H40N6Zn2C18 : C ’ 36*7°5 H » 4 *74 N * 9 ’88* Foundl C, 36.55;H, 5.07; N, 9.91.

(2,12,14,20-Tetramethyl-3,11,15,19,22,26-hexaazatricyclo- 5 9[11.7.7.1 * ]octacosa-l,5,7,9(28),12,14,19,21,26-nonaene) Hexafluoro-

phosphate, [(m-Xylyl(NHEthi)2)Me2 [16]tetraene](PF^)^. One gram (1.2

mmole) of [(m-Xylyl(NHEthl)2)Me2 [16]tetraene](ZnCl^)2 was dissolved in

50 ml of water and filtered. A solution of 3.0 g (18.4 mmole) of

ammonium hexafluorophosphate dissolved in 20 ml of water was added

dropwise to yield the white granular solid which was collected and

dried in vacuo. Yield: 0.74 g (72%). Anal: Calc, for C ^ H ^ N g P g F ^ :

C, 35.87; H, 4.52; N, 9.65. Found: C, 35.72; H, 4.75; N, 9.89.

46

(2,11,13,19-tetramethyl-3,10,14.18,21,25-hexaazatricyclo 5 8[10.7.7.2 * ]octacosa-l,5.7,ll,13,18,20,25,27-nonaene) Hexafluorophos-

phate, [ (p-Xylyl(NHEthl) n)Me„ri6]tetraene] (PF^.)^. This ligand salt was78prepared by the method of Callahan and Busch. Hydrogen chloride gas

was bubbled through a solution of 4.0 g (5.1 mmole) of [Ni{ (£-Xylyl-

(NHEthi^)!^ [16]tetraeneN^}] (PF^)2 dissolved in 300 ml of acetonitrile.

The flask was stoppered and set aside for 2 days during which time the

solution turned green in color. A solution of tetrachlorozincate anion

(prepared by the reaction of 4.0 g (61.5 mmole) of zinc metal with

hydrogen chloride gas in 200 ml of acetonitrile) was added to yield the

white tetrachlorozincate ligand salt which was collected, washed with

acetonitrile and dried in vacuo. Yield: 4.1 g (91%). The above pro

duct was dissolved in 100 ml of water and filtered. Dropwise addition

of 10,0 g (61.3 mmole) of ammonium hexafluorophosphate dissolved in

50 ml of water resulted in formation of an off-white precipitate which

was collected and dried in vacuo. Yield: 3.8 g (91%). Anal: Calc.

for C26H39N6P3F18: C ’ 35,87i H > 4*52» N > 9 *65* F°und: C, 33.94;H, 4.17; N, 9.21.

(2,3,10,11,13,19-Hexamethyl-3,10,14,18.21,25-hexaazabicyclo-

[10.7.7]hexacosa-l,11,13,18,20,25-hexaene) Hexafluorophosphate,

[ (1,6-Hex(MeNEthi))Me„ [16 ]tetraene] (PF^)^* This ligand salt was pre-77 66pared by the method of Olszanski and Busch as reported by Stevens.

47

(2,3,8,9,11,17-Hexamethyl-3 ,8,12,16,19, 23-hexaazabicyclo[8.7.71

tetracosa-l,9,ll,16,18,23-hexaene) Hexafluorophosphate, [(1,4-But(MeNEthi)„)Me„[161tetraenel(PF,)„' t- £----------- — 6—3

(2,3,9,10,12,18-Hexamethyl-3,9,13,17,20,24-hexaazabicyclo

[9.7.7]-pentacosa-l,10,12,17,19,24-hexaene) Hexafluorophosphate,

[(l,5-Pent(MeKEthi) O M e ^ [16]tetraene] (PF .) .

The above two ligand salts were prepared in the method

described by Stevens.^

(2.12,14,20,22,32,34,40-Octamethyl-3,11,15,19,23,31.35,39,42,

46,50,54-dodecaazapentacyclo[31.7.7. 7 ^ * 1~**^. 1 ^ * hexapentaconta-

1,5,7,9(56),12,14,19,21,25,27,29(48),32,34,39,41,46,49,54-octadecaene)

Tetrachlorozincate, [ (m-Xylyl(NHEthi) n)Me,, F^^ tetraene] ,, (ZnCl^) • This

ligand salt was prepared in the same way as previous examples using 4.0 g

(2.6 mmole) of dimeric nickel(II) complex derived from m-xylylenedi-

amine and 2.0 g (3.1 mmole) of zinc metal. The HC1 gas was passed

through the solution at 0°C and addition of the tetrachlorozincate

anion resulted in immediate precipitate formation. The product was

collected, then resuspended in 100 ml of fresh acetonitrile and stirred

for 1 h. The product was again collected, washed with acetonitrile and

ether and dried in vacuo. Yield: 3.7 g (85%).

48

(2.3.11.13,19.21.22.29.30.32.38-Dodecamethyl-3 ,10,14,18,22.29.

33,37,40,44,47,51-dodecaazapentacyclo [29.7.7.712 » 20.25 >8 .224’27 ]

hexapentaconta-1,5,7,11.13,18.20.24,26,30,32,37,39,44.46,51,53,55-

octadecaene) Tetrachlorozincate, [(p-Xylyl(MeNEthl),,)Me,, [16] tetraene],,

(ZnCl^)^. This ligand salt was prepared by the method described above

using 3.5 g (4.3 mmole) of [Nl-CC^-XylylCMeNEthi^Jt^tlbJtetraeneN^}^-

(PFg)^ and 1.5 g (23 mmole) of zinc metal. The off-white precipitate

was collected and dried in vacuo. Yield 3.0 (79%). No analytical data

were obtained for this ligand salt.

(2,6,7,11,13,19,21,25,26,30,32,38,53,54,55,56-Hexadecamethyl-

3,10,14,18,22,29,33,37,40,44,47,51-dodecaazapentacyclo

[29.7.7.712,2Q.25 »8 .224>27]hexapentaconta-l,5,7,ll,13,18,20,24,26,30,

32.37.39.44.46.51.53.55-octadecaene) Tetrachlorozincate, [(Duryl-

(NHEthi) ,)Me^[16]tetraenel^(ZnCl^)^. This ligand salt was prepared by

the same method as those reported above using 1.5 g (1.8 mmole) of the

nickel complex derived from the duryl bridging group and 0.75 g (11.5

mmole) of zinc metal. The precipitate was collected, washed with

acetonitrile and ether and dried in vacuo. Yield: 1.0 g (61%). No

analytical data were obtained for this impure ligand salt.

(2.3,6.7.10,11,13.19.21,22,25,26,29,30,32,38,53,54,55,56-

Eicosamethyl-3,10,14,18,22,29,33,37,40,44,47,51-dodecaazapentacyclo

[29.7.7.712 *20.25 *8.224 *27]hexapentaconta-l,5,7,11,13.18.20,24,26,30,

32.37.39.44.46.51.53.55-octadecaene) Tetrachlorozincate, [(Duryl-

(MeNEthi)„)Me^[16]tetraene],, (ZnCl^) . This ligand salt was prepared by

the method reported above using 2.0 g (2.3 mmole) of the nickel methyl

49

substituted durenyl complex and 1.0 g (15 mmole) of zinc metal. Yield:

1.73 g (80%). No analytical data were obtained for this impure ligand

salt.

(2»12-Dimethy1-3,11-bis[1-(dimethylamino)ethylidene]-!,5,9,13-

tetraazacyclohexadeca-l,4,9,12-tetraene) Hexafluorophosphate,

[ (Me^NEthi) ,Me„ [ 16] tetraene] CPF^-)^. This ligand salt was synthesized79and characterized by Dr. R. A. Wilkins.

(2,12-Dimethyl-3,11-bis[1-(amino)ethylidene]-1,5,9,13-

tetraazacy clohexadeca-1 ,4,9,12-tetraene) Tetrachlorozincate„

[ (NHÊthi)^Me^. [16] tetraene] (ZnCl^K •

(2,12-Dimethy1-3,11-bis[1-(methylamino)ethylidene J-1,5,9,13-

tetraazacyclohexadeca-1,4,9,12-tetraene) Tetrachlorozincate,

[ (MeNHEthi)^Me^[18]tetraene] (ZnCl^.)„.

The above two ligand salts were prepared in good yield in the

same way as [(m-Xylyl(MeNEthi)2)Me2 [16]tetraene)ZnCl^ ) ^

Synthesis of Iron(II) Starting Materials

Bis-Acetonitrileiron(II) Chloride. To 100 ml of acetonitrile

was added commercial grade anhydrous iron(II) chloride until the solu

tion was well saturated. Iron filings were added and the solution was

stirred under reflux for 4 h. The solution was filtered through celite

while hot. Upon cooling, the off-white precipitate formed. This was

recrystallized from hot acetonitrile to yield the white crystalline

product which was collected and dried in vacuo. Analytical data

50

Indicated the stoichiometry of the compound to be Fe(CH-CN).. ,(H„0)n<3 1 1 / Z (J> JAnal. Calc.: C, 20.23; H, 2.82; N, 11.79. Found: C, 20.28; H» 2.86;

N, 11.96.

80Hexakis-Acetonitrileiron(II) Hexafluorophosphate. To a sus

pension of 2.0 g (11.4 mmole) of nitrosyl hexafluorophosphate in 30 ml

of acetonitrile was added 0.64 g (11.4 mmole) of iron filings. Immedi

ate foaming of the solution occurred and a vacuum was applied to the

flask until the reaction was complete. The volume of the solution was

brought to 75 ml and the solution was heated to boiling. After fil

tering the hot solution through celite, its volume was reduced to 25 ml,

at which point a considerable amount of precipitate formed. This was

collected and recrystallized from hot acetonitrile and dried in vacuo.

Anal. Calc, for FeCi2Hl8N6P2F12: C * 24.34; H, 3.06; N, 14,19; Fe, 9.43.Found: C, 24.18; H, 3.05; N, 14.03; Fe, 9.33.

Synthesis of Unbridged Iron(II) Complexes

Acetonitrile(2.12-Dimethy 1-3,11-bis [l-(methylamino) ethylidene-]-

1,5,9,13-tetraazacyclohexadeca-l,4,9.12-tetraenelQ iron(II) Hexafluoro

phosphate, [Fe{ (MeNHEthi)nMe„[16]tetraenetQ(CH^CN) ] (PF^)„. To a sus

pension of 2.0 g (2.6 mmole) of [(MeNHEthi)2M©2 [16]tetraene](ZnCl^)^ in

50 ml of acetonitrile was added 0.54 g (2.6 mmole) of bis-acetonitrile-

iron(II) chloride and 1.04 g (10.3 mmole) of triethylamine. The solu

tion was refluxed for 10 min then stirred overnight. The solvent was

removed and the residue dissolved in methanol and filtered through

celite. Addition of excess ammonium hexafluorophosphate dissolved in

51

ethanol followed by volume reduction yielded the orange product which

was collected and dried in vacuo. Yield: 1.6 g (82%). Analytical

data and spectroscopic measurements are consistent with a mixture of

the product labelled above and [Fe{Me_(MeIMEt)„[16]tetraeneN.}](PF,)„Z Z h 6 2in a ratio of 3:1. Calc, for PeC21 5H36 25N & 75P2F12: C * 35,13»H, 4.97; N, 12.87; Fe, 7.60. Found: C, 35.06; H, 5.18; N, 12.95;

Fe, 7.71.

A Sexadentate Isomer of the Unbridged Iron(II) Complex:

[(3,11-bis(l-methylaminoethyl)-2,12-dimethy1-1,5,9,13-tetraazacyclohexa-

deca-l,3,9,ll-tetraenelQiron(II) ] Hexafluorophosphate, [Fe{Me2~

(MelMEt) „ [ 16] tetraeneN^} ] (PF .) „. To a slurry of 1.2 g (1.5 mmole) of

[ (MeNHEthi)2Me2 [ 16] tetraeneN^. (ZnCl^)3 in 20 ml of methanol was added

0.33 g (1.6 mmole) of bis-acetonitrileiron(II) chloride and 0.94 g

(9.3 mmole) of triethylamine. The solution was refluxed for 45 minutes

during which time a red-orange precipitate formed. Methanol was added

to bring the volume to 50 ml and the solution was quickly filtered

through celite. To the warm solution was added dropwise 2.0 g (12.3

mmole) of ammonium hexafluorophosphate dissolved in a minimum volume of

methanol. The red-orange crystalline product formed upon standing over

night. The crystals were collected and washed with methanol and ether

and dried in vacuo. Yield: 0.75 g (69%). Anal. Calc, for

FeC20H34N6P2F12: C ’ 34,1°i H » 4.87; N, 11.93. Found: C, 33.78;H, 4.91; N, 11.72.

Synthesis of Iron(II) Dry Cave Chloro Complexes

52

Chloro(2,12.14.20-tetramethyl-3,11,15,19,22,26-hexaazatricyclo- 5 9[11.7.7.1 * loctacosa-1,5.7.9,(28),12,14<19t21,26-nonaeneH^)iron(II)

Chloride dimethanol, [Fe{ (m-Xylyl(NHEthi)2)Me2 [16] tetraeneN,. }ci]~Cl* 2CHjOH. To a solution of 1.0 g (1.2 tnmole) of [ (m-Xylyl(NHEthl)

Me2 [16]tetraene](PFg)^ dissolved in 50 ml of acetonitrile was added

0.24 g (1.2 mmole) of bis-acetonitrileiron(II) chloride and 0.35 g

(3.5 mmole) of triethylamine to yield a deep red solution which was

filtered through celite. The solution was stirred overnight during which

time an orange precipitate formed which was collected and dried in vacuo.

Yield: 0.5 g (78%). Anal. Calc, for FeC26H36N6Cl2 : C, 55.83;

H, 6.49; N, 15.07. Found: C, 55.73; H, 6.53; N, 15.16. This product

was recrystallized from methanol to yield large crystals of the

dimethanol solvate. Anal. Calc, for FeC„ftH,,N„0„C1„: C, 53.94; 28 44 6 2 2H, 7.11; N, 13.48; Cl, 11,37; Fe, 8.96. Found: C, 53.99; H, 6.94;

N, 13.77; Cl, 11.32; Fe, 8.75.

Chloro(2,12,14,20-tetramethyl-3,11,15,19,22,26-hexaazatricyclo-

[11.7.7.1***^]octacosa-l,5,7,9(28), 12,14,19,21,26-nonaenelQiron(II)

Hexafluorophosphate, [Fe{(m-Xylyl(NHEthi)2)Me2 [16]tetraeneN^}cl] (PF^) •

To a solution of 0.5 g (0.89 mmole) of [Fe{(m-Xylyl(NHEthi)2)Me2[16]-

tetraeneN^}Cl]Cl*2CH^0H dissolved in 75 ml of methanol was added an

excess of ammonium hexafluorophosphate dissolved in methanol to yield a

red-orange microcrystalline precipitate which was collected and dried

53

in vacuo. Yield: 0.5 (85%). Anal. Calc, for Fe C ^ H ^ N g C l P F ^ :

C, 46.69; H, 5.43; N, 12.56; Cl, 5.30. Found: C, 46.53; H, 5.48;

N, 12.57; Cl, 5.26.

Chloro(2.3,11,12,14,2Q-hexamethyl-3,11,15,19,22,26-hexaazatri-

cyclo [11.7.7.1~* * ] octacosa-1,5,7,9(28) 12.14,19,21,26-nonaeneN^) iron (II)

Hexafluorophosphate, [Fe{ (m-Xylyl(MeNEthi)n)He^[16ItetraeneN^}Cl] (PF .) .

To a suspension of 2.0 g (2.3 mmole) of [(m-Xylyl(MeNEthi)2)Me2 [16]-

tetraene](ZnCl^)g in 50 ml of methanol was added 0.48 g (2.3 mmole) of

bis-acetonitrileiron(II) chloride and 0.92 g (9.1 mmole) of triethyl-

amlne. The solution was refluxed for 10 minutes then filtered through

celite. After addition of 3.0 g (18.4 mmole) of ammonium hexafluoro

phosphate dissolved in a minimum volume of methanol the solution was

allowed to stand overnight. The large red crystals which formed were

collected and dried in vacuo. Yield: 1.15 g (72%). Anal. Calc, for

FeC28HA()N 6ClPF6 : C, 48.26; H, 5.78; N, 12.06; Cl, 5.09. Found:

C, 48.22; H, 5.93; N, 12.00; Cl, 5.01.

Chloro(2,11,13,19-tetramethy1-3,10,14,18,21,25-hexaazatricyclo-

[10.7.7.2~* *^]octacosa-l,5,7,11,13,18,20,25,27-nonaeneN^) iron(II) Hexa

fluorophosphate, [Fe{ (p-Xylyl (NHEthi)2)Me2 [16] tetraeneN^Cl] (PF .) . This78complex was prepared by the method of Callahan and Busch. To a solu

tion of 1.0 g (1.1 mmole) of [(j5-Xylyl(NHEthi)2)Me2 [16]tetraene] (PFg)^

dissolved in 20 ml of acetonitrile was added 0.48 g (2.3 mmole) of

bis-acetonitrileiron(II) chloride. To this mixture was added dropwise

a solution of 0.35 g (3.5 mmole) of triethylamine in 5 ml of acetoni

trile. The solution immediately turned deep red and an orange

54

precipitate began to form. After stirring overnight, the precipitate

was collected and dried in vacuo. The product was dissolved in 75 ml

of methanol and 2.0 g (12.3 mmole) of ammonium hexafluorophosphate dis

solved in ethanol was added to yield the yellow-brown microcrystalline

product which was dried in vacuo. Recrystallization from an

acetonitrile-ethanol mixture yielded large orange crystals. Yield:

0.55 g (75%). Anal. Calc, for Fe C ^ H ^ N g C l P F g : C, 46.69; H, 5.43,

N, 12.56; Cl, 5.30. Found: C, 46.59; H, 5.77; N, 11.93; Cl, 4.67.

Chloro(2,3,10,11,13,19-hexamethyl-3,10,14,18,21,25-hexaaza-

bicyclo[10.7.7]hexacosa-l,11,13,18,20,25-hexaenelQ iron(11) Hexafluoro

phosphate, [Fe{ (!,6-Hex(MeNEthi)„)Me„[16]tetraeneN^}Cl] (PF^) • This com-77plex was prepared according to the method of Olszanski and Busch. To

a solution of 2.0 g (2.3 mmole) of [(1,6-Hex(MeHEthi)2)Me2 [16]tetraene]-

(PFg)^ dissolved in 50 ml of acetonitrile was added 0.46 g (2.3 mmole)

of bis-acetonitrileiron(II) chloride and 0.70 (6.9 mmole) of triethyl-

amine to yield a deep red solution. After refluxing for ten minutes,

the solution was filtered through celite and allowed to stir for 2 h.

The solvent was removed and the residue dissolved in 40 ml of methanol.

Dropwise addition of a solution of 2.0 g (12.2 mmole) of ammonium hexa

fluorophosphate dissolved in 20 ml of ethanol yielded the deep red pro

duct which was recrystallized from an acetonitrile, ethanol mixture to

yield large red crystals. Anal. Calc, for FeCggH^NgClPFg: C, 46.13;

H, 6.55; N, 12.41; Cl, 5.24. Found: C, 45.93; H, 6.77; N, 12.24;

Cl, 5.14.

Synthesis of Iron(III) Dry Cave Complexes

55

Chloro (2.3,11,12,14.20-hexamethyl-3,11.15,19,22,26-hexaazatri- 5 9cyclo[ll.7.7«1 * ]octacosa-l,5.7.9(28),12,14.19,21.26-nonaeneN,,)-

iron(III) Hexafluorophosphate, [Fe{Cm-XylvI(MeNEthi)»;Me„[16]tetraene-

N^}C1](PF^.)„. To a solution of 0.2 g (0.29 mmole) of [Fe{(m-Xylyl-

(MeNEthi)2)Me2 [l6]tetraeneN^}ci](PF^) dissolved in 25 ml of methanol

and 5 ml of acetonitrile was added 0.16 g (0.29 mmole) of ammonium

hexanitratocerate(IV) causing a color change from red to deep blue.

Dropwise addition of a solution of 0.28 g (1.7 mmole) of ammonium

hexafluorophosphate dissolved in 10 ml of methanol yielded a deep blue

mlcrocrystalllne precipitate which was recrystallized from an acetoni

trile, ethanol mixture. Yield: 0.18 g (74%). Anal. Calc, for

FeC28H40N6ClP2F12: C, 39.95; H, 4.79; N, 9.98; Cl, 4.21. Found:

C, 38.12; H, 4.99; N, 9.44; Cl, 4.03.

U-Oxo-bis r(2,3.11,12,14.20-hexamethyl-3.11.15.19,22,26-hexa- 5 9agatricyclo[11.7.7.1 * ]octacosa-l,5,7,9(28),12,14,19,21,26-nonaeneN^)-

iron(III)] Hexafluorophosphate trihydrate, {[Fe[(m-Xylyl(MeNEthi)2)-

Me„[l6 ]tetraeneN^]]20}(PF^)^* 3H20. A solution of 0.5 g (0.7 mmole) of

[Fe{(m-Xylyl(MeNEthi)2)Me2[16]tetraeneN^}Cl](PFg) dissolved in 20 ml of

acetonitrile was exposed to the air for 3 days. The solution changed in

color from deep red to brown. Upon addition of water, a brown preci

pitate formed which was collected and washed with water and ether.

Anal. Calc, for Fe2C5&H9oNl2°4F4F24: C * 3^.87; 5.38; N, 9.96.Found: C, 40.19; H, 4.79; N, 9.54.

Synthesis of. Other Iron(II) Dry Cave Complexes

56

A Pentadentate Isomer of a Bridged Iron(II) Complex: [2.12,14.

20-tetramethyl-3,11,15,19,22,26 hexaazatrlcyclo til * 7.7.1~* * loctacosa-

2,5,7,9,(28)12,14,19.21,26-nonaenelQIron(II) Hexafluorophosphate,

[Fe{ (m-Xylyl(NHEthl) (lE)Me„ [16]tetraeneN^}] (PF .) .. To a solution of

1.0 g (1.1 mmole) of [(m-Xylyl(NHEthi)2)Me2 [16]tetraene](PFg)^ dissolved

in 20 ml of acetonitrile was added 0.67 g (1.1 mmole) of

hexakis-acetonitrileiron(II) hexafluorophosphate and 0.35 g (3.4 mmole)

of triethylamine. The solution was refluxed for 10 minutes then fil

tered through celite. The solution volume was reduced to 10 ml and

20 ml of methanol was added. An orange crystalline product formed over

night and was collected. Yield: 0.4 g (47%). Anal. Calc, for

FeC26H36N6P2F12: C, 40.12; H, 4.66; N, 10.80. Found: C, 46.39;

H, 5.90; N, 12.48.

(2,3.11,12,14,20-hexamethy1-3,11,15,19.22,26-hexaazatricyclo

[ 11.7.7.1~* * 1 octacosa—1,5,7,9(28) , 12.14119,21, 26-nonaeneN^, ) iron(II)

Hexafluorophosphate, [Fe{ (m-Xylyl(MeKEthi)2)Me2 [16]tetraeneN^}] *

To a solution of 0.3 g (0.4 mmole) of [Fe{(m-Xylyl(MeNEthi)2)Me2 [16]-

tetraeneN,}C1](PF,) dissolved in 5 ml of acetonitrile and 50 ml ofH D

ethanol was added 0.3 g (4.2 mmole) of pyridine. Introduction of CO to

the solution produced a color change to yellow-orange. A vacuum was

then applied to the flask resulting in a deepening of the color . An

excess of ammonium hexafluorophosphate dissolved in ethanol was added

dropwise to yield a red mlcrocrystalline product which was collected.

Upon drying in vacuo overnight, the color of the solid changed from red

57

to brown. Analytical data indicate the product is a mixture of the

starting material and the desired product in a ratio of 6:1. Calc. C,

42.52; H, 5.10; N, 10.57. Found: C, 42.60; H, 5.04; N, 10.57.

An alternate synthesis of this complex follows:

To a suspension of 0.5 g (0.56 mmole) of [(m-Xylyl(MeNEthi)2)-

Me2 [16]tetraene](PF^)^ was added 0.32 g (0.56 mmole) of

hexakis-acetonitrileiron(II) hexafluorophosphate and 0.1 g (0.99 mmole)

of triethylamine. The solution was refluxed for 10 min and filtered

through celite. After addition of excess ammonium hexafluorophosphate

in ethanol, the solution volume was reduced until the precipitate

formed. The dark brown powder was collected and dried in vacuo.

Analytical data were not obtained but spectroscopic data match those of

the analyzed sample.

(2,3,11,12,14,20-Hexamethyl-3,11.15,19,22,26-hexaazatricyclo-

m . 7 . 7 . 1 5,9 ] octacosa-1,5,7,9 (28) 12,14,19,21,26-nonaeneN^) iron(II)

Iodide«Q.5 acetone, [Fe{(m-Xylyl(MeNEthi)^)Me„[16]t e t r a e n e N ^ *0.5-

CHjCOCHq To a solution of 0.5 g [Fe{(m-Xylyl(MeNEthi) Me [16]tetraene-

N^}](PFg)2 dissolved in 30 ml of acetone was added an excess of

tetra n-butylammoniura iodide dissolved in acetone. A brown precipitate

formed overnight which was collected and dried in vacuo. Anal. Calc.

for FeC2g g H ^ N ^ 5I2 : C, 44.33; H, 5.42; N, 10.51; I, 31.75. Found:

C, 44.53; H, 5.43; N, 10.55; I, 31.40.

An alternate synthesis of this complex is described below:

To a solution of 1.0 g (1.1 mmole) of [(m-Xylyl(MeNEthi)2)Me2~

[16]tetraene)](PF£)_ dissolved in 50 ml of acetone was added 0.65 g o 3(1.1 mmole) of hexakls-acetonitrileiron(II) hexafluorophosphate and

58

0.45 g (4.4 mmole) of trlethylamine. The solution was filtered through

celite and a solution of 2.0 g (5.5 mmole) of tetra-ii-butylainmonium

iodide dissolved in 10 ml of acetone was added. Reduction in volume

yielded a brown precipitate which was collected. Analytical data were

not obtained but spectroscopic measurements match those of the analyzed

sample.

Attempt to prepare (2,3,8,9,11,17-Hexamethy1-3,8,12,16,19,23-

hexaazabicyclo[8.7.7]tetracosa-1,9,11,16,18.23-hexaeneN^)iron(XI)

Hexafluorophosphate, [Fe{ (1,4-But (MeNEthi) j M e „ [ 16 ] tetraeneN^} ] (PF .) „ .

In the same way as described for the alternate synthesis of

[Fe{ (m-Xylyl(MeNEthi)2)Me2 [16] tetraeneN^, } ] (PFg)^, 1.03 g (1.8 mmole) of

hexakis-acetonitrilelron(II) hexafluorophosphate and 0.70 g (6.9 mmole)

of trlethylamine were mixed. An orange, partially crystalline product

was Isolated and dried in vacuo. Satisfactory elemental analyses were

not obtained. Calc, for F e C ^ H ^ N ^ F ^ : C, 38.01; H, 5.32; N, 11.08.

Found: C, 41.42; H, 6.39; N, 12.01.

Attempt to prepare (2,3,9,10,12,18-Hexamethyl-3,9,13,17,20,24-

hexaazabicyclo(9.7.7]pentacosa-l,10,12,17,19,24-hexaeneN^)iron(II)

Hexafluorophosphate, [Fe{ (1,5-Pent (MeNEthi) )Me^ [ 16 ] tetraeneN^ } ] (PF .) 2 ’

In the same way as described for the preceding complex, 1.5 g (1.7

mmole) of [ (1,5-Pent (MeNEthi) 2 ^ &2 tetraene] (PFg)^, 1.03 g (1.8

mmole) of hexakis-acetonitrileiron(II) hexafluorophosphate and 0.70 g

(6.9 mmole) of trlethylamine were mixed. A red-orange crystalline pro

duct was collected. Satisfactory elemental analyses were not obtained.

59

Calc, for FeC25H 42N 6P2F12: C ’ 38*87» H » 5 *48» N, 10.88. Found: C,41.65; H, 6.51; N, 12.06.

Attempt to prepare (2,3,10,11,13,19-Hexamethyl-3,10,14,18,21,

25-hexaazablcyclo[10.7.7] hexacosa-1,11,13,18.20,25-hexaeneN^)Iron(II)

Hexafluorophosphate, [Fe{ (1,6-Hex (MeNEthi) )Me„ [ 161 tetraeneN^ } 3 (PF -) 2 •

In the same way as described for the two preceding complexes, 1.5 g

Cl.7 mmole) of [(l,6-Hex(MeNEthi)2)Me2 [16]tetraene](PFg)^, 1.03 g (1.7

mmole) of hexakis-acetonitrilelron(II) hexafluorophosphate and 0.7 g

(6.9 mmole) of trlethylamine were reacted. The brown product was iso

lated and dried in vacuo. Satisfactory elemental analyses were not

obtained. Calc, for FeC26H44N6P2F12: C * 39*71» H, 5.64; N, 10.69.Found: C, 42.55; H, 6.47; N, 11.82.

Synthesis of Carbon Monoxide Addurts of the Iron(II) Complexes

All of the six coordinate irc.ull) carbon monoxide complexes

were synthesized in the same way; exceptions are noted below. Approxi

mately 0.250 g of the appropriate iron(II) chloro complex and an equal

mass of the appropriate base were dissolved in 5 ml of acetonitrile and

30 ml of ethanol. The solution was filtered into a Schlenk flask and

removed from the dry box. The flask was connected to a cylinder of

Matheson Research Grade carbon monoxide and the flask and all hoses were

evacuated under high vacuum. Carbon monoxide was introduced at a pres

sure slightly above atmospheric, causing the solution color to change

from deep red-orange to a much less intense orange color. The flask

was sealed and returned to the dry box whereupon 0.5 g of ammonium

hexafluorophosphate dissolved in a minimum volume of ethanol was added

60

dropwise to yield a red-orange, generally crystalline precipitate which

was collected and washed with ethanol and ether. Yields were generally

greater than 65%. Analytical data for the CO adducts are summarized

in table 2. The following complexes were synthesized by this method;

[Carbonyl(B)(2,12,14,20-tetramethyl-3,11,15,19,22,26-hexaazatricyclo

[11.7.7.15 *9]octacosa-1,5,7,9(28),12,14,19,21,26-nonaeneN4> iron(II)]

Hexafluorophosphate, [Fe{(m-Xylyl(NHEthi)2)Me2 [16]tetraeneN^}(B)(CO)]

(PFg)^, B = pyridine, N - methylimldazole, imidazole, 4-arainopyridine.

[Carbonyl(B)(2,3,11,12,14,20-hexamethyl-3,11,15,19,22,26-hexaazatri

cy clo [11.7.7.1^ *9]octacosa-1,5,7,9(28)12,14,19,21,26-nonaeneN^)iron(II)]

Hexafluorophosphate, [Fe{ (m-Xylyl(MeNEthi)2)Me2 [1®] tetraeneN^}(B) (CO)]

(PFg)^, B = acetonitrile, N - me ■ 'limidazole, imidazole,

4-aminopyridine.

[Carbonyl(B)(2,11,13,19-tetramethy1 - , _ 3,14,18,21,25-hexaazatricyclo

[10.7.7.2^*®]octacosa-1,5,7,11,13,18,20,25,27-nonaeneN^)iron(II)]

[Fe{ (£-Xylyl(NHEthi)2)Me2 [16] tetraeneN^}(B) (CO) ] (PFg)2> B = acetoni

trile, pyridine, N - methylimldazole.

[Carbonyl(B)(2,3,10,11,13,19-hexamethyl-3,10,14,18,21,25-hexaazabicyclo

[10.7.7]hexacosa-l,11,13,18,20,25-hexaeneN^)iron(II)] Hexafluorophos

phate, [Fe{(l,6-Hex(MeNEthl)2)Me2 [16]tetraeneN4}(B)(CO)](PF6)2 , B =

acetonitrile, pyridine, N - methylimldazole.

61

TABLE 2ANALYTICAL DATA FOR THE COMPLEXES

[Fe{(R)Me2[16]tetraeneN^>(B)(CO))(PFg)2

R BC H N Fe

Calc. Fd. Calc. Fd. Calc. Fd. Calc. Fd.

m-Xylyl(NHEthl) 2 PY 43.41 42.83 4.67 4.93 11.07 10.72 6.31 5.80BrXylyl(NHEthi)2 1-MeIM 41.91 40.99 4.76 4.85 12.61 12.64 6.29 5.61.m-Xylyl(NHEthi)2 Im 41.21 41.33 4.61 4.84 12.81 12.36m-Xylyl(NHEthi)2 A-NH2-Py 42.68 42.57 4.70 4.70 12.44 12.57 6.20 6.14a-Xylyl(MeNEthi)2 1-MelM 43.24 43.76 5.06 5.01 12.23 12.92m-Xylyl(MeNEthi)2 CHjCN 42.53 43.20 4.95 5.33 11.20 12.24m-Xylyl(MeNEthi)2 Im 42.59 41.37 4.91 5.08 12.42 12.37 6.19 5.90m-Xylyl(MeNEthi)2 «-NH2-Py 43.98 43.51 4.99 5.27 12.07 12.70 6.01 5.89j^-Xylyl(NHEthi)2 CH3CN 41.10 40.05 4.64 4.97 11.57 11.01£-Xylyl(NHEthi)2 Py 43.41 43.07 4.67 4.76 11.07 11.09£-Xylyl(NHEthi)2 1-MeIM 41.91 41.50 4.76 5.12 12.61 12.571,6-Hex(MeNEthi) 2 ch3cn 41.20 40.97 5.59 5.66 12.08 12.251,6-Hex(MeNEthi)2 Py 43.01 41.91 5.53 5.73 10.97 11.29 6.25 5.901,6-Hex(MeNEthi) 2 1-MeIM 41.53 41.80 5.62 5.68 12.50 12.98 6.23 5.78(Me2NEthi)2 ch3cn 37.47 36.10 5.16 5.45 12.23 12.02(MeNHEthi)(NE) ---- 34.44 33.78 4.68 4.48 11.48 11.35

62

The CO Adduct of a Pentadentate Unbrldged Iron(II) Complex:

[Carbonyl(3-[1-(methylamino)ethylldenel-11-(l-methyl-iminoethyl)-2,12-

dlmethy1-1,5,9,13-tetraazaeyclohexadeca-1,3,9,11-tetraeneN^)Iron(II)]

Hexafluorophosphate, [{Fe(MeNHEthl)(MeImEt)Me„ri6]tetraeneN^}(C0)]

(FF^.)„. To a solution of 0.3 g (0.4 mmole of [Fe{(MeNHEthi)2Me£ [16]—

tetraeneN^}(CH^CN) H P F g ^ dissolved in 3 ml of acetonitrile was added

an equal mass of pyridine, N-methylimidazole or acetonitrile and 50 ml

of ethanol. After exposure to CO, there was a very slow color change

from red to light orange. The solution was allowed to stand overnight.

After addition of excess ammonium hexafluorophosphate followed by

volume reduction, a red orange precipitate formed which was isolated

and dried in vacuo. The same product was obtained regardless of the

base used as shown by analytical, infrared and NMR data.

[(Acetonitrile)(carbonyl)(2,12-Dimethyl-3,ll-bis-[l-(dimethyl-

amino)ethylidene]-1,5,9,13-tetraazacyclohexadeca-1,4,9,12-tetraeneN^)

iron(II)] Hexafluorophosphate, [Fe{(Me„HEthi)^Me^,[16]tetraeneH^}

(CHjCN) (CO) ] ( P F ^ K . To a solution of 1.0 g (1.03 mmole) of

[(Me2NEthi)2Me2[16]tetraene](PF^)^ dissolved in 15 ml of acetonitrile

was added 0.25 g (1.2 mmole) of bis-acetonitrileiron(II) chloride and

0.5 g (4.9 mmole) of trlethylamine. After stirring for 10 minutes the

solution was filtered through celite and 20 ml of ethanol was added.

Upon addition of CO, the solution changed in color from deep red to

yellow-orange and an orange crystalline precipitate began to form. The

solution was returned to the dry box and the product was isolated and

dried in vacuo. Yield: 0.75 g (90% based on ligand salt).

[Carbonyl(B)(2,3,10,11,13,19-hexamethyl-3,10,14,18.21,25-

hexaazabicyclo[10.7.7]hexacosa-l, 11,13,18,20,25-hexaeneN^)iron(II)]

Iodide 1.5 acetone, [Fe[(l,6-Hex(MeNEthi)„)Me„[16]tetraeneN^](B)(CO)]

I^,»1.5 acetone, B = pyridine, N - methylimldazole. To a solution of

0.3 g (0.44 mmole) of [Fe{(l,6-Hex(MeNEthi)_)Me0[16]tetraeneN. Cl}(PF,)Z Z H O

dissolved in 40 of acetone was added an equal mass of the appropriate

base. CO was introduced causing the color to change from deep red to

light orange. The solution was returned to the dry box and excess

tetra-n-butylaramonium iodide in acetone was added dropwise. The solu

tion was allowed to stand overnight during which time red-orange

crystalline clusters of the product formed. The crystals were col

lected and washed with acetone and ether. Anal. Calc for

FeCsg 5^58^7®2 5^2* 46.45; H, 6.20; N, 10.35. Found: C, 46.80;

H, 5.97; N, 10.78; and Calc, for C^^ s^gNgOj 5*2: 45.00; H, 6.28;

N, 11.83. Found: C, 45.21; H, 6.37; N, 11.95.

Synthesis of Dimeric Iron(II) Dry Cave Complexes

Starting Material for Dimeric Iron(II) Complexes. To a suspen

sion of 1.0 g (0.6 mmole) of [(m-Xylyl(NHEthi)2)Me2 [16]tetraene]2“

(ZnCl^)^ in 25 ml of acetonitrile was added 0.78 g (1.8 mmole) of70tetrakis-(pyridine)iron(II) chloride and 0.71 g (7.0 mmole) of

trlethylamine to yield a deep red solution. This solution was filtered

rapidly through celite and allowed to stir overnight during which time

a red precipitate formed. The product was collected and washed with

acetonitrile and ether and dried in vacuo. Yield: 0.65 g.

64

(Pyridine) {2_r l 2 ,14,20,22.32.34.40-0c tame thy 1-3.11.15.19.23.31.35.39.42.46.50.54-dodecaazapentacyclo [31.7.7. 7 ^ ^25 291 * ]hexapentaconta-1,5.7.9(56).12.14.19.21.25.27.29(48).32.34.39.

41.49.43-octadecaene)diiron(II) Hexafluorophosphate. [Fe{(m-Xylyl-

(NHEthi)„)Me„[l6]tetraeneN^}(Py)]„(PF^)^. One gram of the crude

starting material was dissolved slowly in 100 mL of methanol and the

solution was filtered through cellte. To this solution was added 3.0 g

(38.0 mmole) of pyridine and the solution volume was reduced to 50 mL.

A solution of 2.0 g (12.3 mmole) of ammonium hexafluorophosphate dis

solved in methanol was added dropwise to yield a red precipitate. The

precipitate was washed with methanol and ether and dried in vacuo.

Yield: 0.90 g (58% based on ligand). Anal. Calc, for

FeC31H41N7P2F12: C, 43.42; H, 4.82; N, 11.43; Fe, 6.51,; Zn, 0.0.

Found: C, 43.51; H, 4.96; N, 11.33; Fe, 6.23; Zn, 0.24.

(Imidazole) (2,12,14,20,22,32,34,40-Qctamethvl-3.11.15.19.23.31.35.39.42.46.50.54-dodecaazapentacyclo T31.7.7. 7^ . 1~* * . -25 291 * ]hexapentaconta-l,5.7.9(56).12.14.19.21.25.27.29 f48),32.34.39.

41.49.43-octadecaene)diiron(II) Hexafluorophosphate. [Fe{(m-Xylyl-

(NHEthi)2)Me„ [161 tetraeneN^} (Im) ]„(PF^)^. To a solution of 0.275 g of

the crude starting material dissolved in 40 mL of methanol was added

0.3 g (5.5 mmole) of imidazole causing the solution to turn deep red.

Slow addition of excess ammonium hexafluorophosphate dissolved in

methanol, without stirring, yielded a red crystalline product upon

standing overnight. The product was filtered, washed with methanol and

dried in vacuo. Yield: 0.22 g (52% based on ligand). Anal. Calc, for

^29^4^8^2*12: 41.15; H, 4.76; N, 13.24; Fe, 6.60; Zn, 0.0. Found:C, 41.01; H, 5.11; N, 13.50; Fe, 6.51; Zn, 0.11.

(2-Methy1imidazole)(2,12,14,20,22,32,34,4Q-0ctamethYl-3«11.15.

19.23.31.35.39.42.46.50.54-dodecaazapentacyclo T31.7.7.7^* . 1~**^.- 25 291 * ]hexapentaconta-l,5.7.9(56).12.14.19.21.25.27.29(48).32.34.39.

41.49.43-octadecaene)dliron(II) Hexafluorophosphate. [Fe{(m-Xylyl-

(NHEthi) „)Me„ [16] tetraeneN^}(2-MeIm) 3^(PF^.)^. This complex was pre

pared in the same way as the imidazole derivative using 0.4 g of the

crude starting material and 0.5 g (6.1 mmole) of 2-methylimidazole.

Yield: 0.31 g (50% based on ligand). Anal. Calc, for [FeC^QH^Ng-

P2Fi2)2 : C, 41.87; H, 4.92; N, 13.02; Fe, 6.49. Found: C, 41.12;

H, 5.37; N, 12.79; Fe, 6.24.

(1-Methylimidazole) (2.12,14,20.22,32.34.40-0ctamethyl-3.11.15.

19.23.31.35.39.42.46.50.54-dodecaazapentacyclo f31.7.7.7^*^.!***^.- 25 291 * ]hexapentaconta-1,5,7.9(56).12.14.19.21.25.27.29 f 48).32.34.39.

41.49.43-octadecaene)diiron(IX) Hexafluorophosphate. [Fe{(m-Xylyl-

(NHEthi) )Me^ f 16 ] tetraeneN^} (l-Melm) ] (PF^) » This complex was prepared in the same way as the imidazole derivative using 0.5 g of the crude

starting material and 0.3 g (3.6 mmole) of N-raethylimidazole. Yield:

0.35 g (45% based on ligand). Anal. Calc, for [FeC3QH42^8*2*12^2:C, 41.87; H, 4.92; N, 13.02; Fe, 6.49; Zn, 0.0. Found: C, 41.31;

H, 5.15; N, 13.48; Fe, 6.48; Zn, 0.24.

Synthesis of Copper(II) Dry Cave Complexes

66

(2,3,10,11,13,19-Hexamethyl-3,10,14,18,21,25-hexaazabicyclo

[10.7.7]hexacosa-l,11,13,18,20,25-hexaeneN^)copper(IX) Hexafluorophos-

phase, [Cu{(l,6-Hex(MeNEthi)n)Me^[16J t e t r a e n e N ^ } ] * ®ne 8ram

(1.1 mmole) of [(1 ,6-Hex(MeNEthi)2 ^ e 2 [16]tetraene](PFg)^ was slurried

in 20 ml of refluxing methanol. To this solution was added a solution

of 0.25 g (1.3 mmole) of copper(II) acetate hydrate and 0.46 g (3.4

mmole) of sodium acetate trihydrate dissolved in 15 ml of hot methanol.

The solution turned brown immediately with formation of a red crystal

line product. The product was recrystallized from an acetonitrile,

ethanol solution to yield large red crystals. Yield: 0.74 g (85%).

Anal. Calc, for CuC26H44N6P2F12: C ’ 39*32; H » 5 *58; N > 10.58; Cu, 8.00.Found: C, 39.42; H, 5.49; N, 10.49; Cu, 7.95.

(2,3,11,12,14,20-Hexamethyl-3,11,15,19,22,26-hexaazatricyclo 5 9[11.7.7.1 * 1octacosa-1,5,7,9(28),12,14,19,21,26-nonaeneN^)copper(II)

Hexafluorophosphate acetonitrile, [Cu[(m-Xylyl(MeNEthi)^)Me„ [16]-

tetraeneN^] 1 (PF^.)^*CH^CN. This complex was synthesized in the same way

as [Cu[(l,6-Hex(MeNEthi)2)Me2[16]tetraeneN^]](PFg)2 « Yield: 0.4 g

(45%). Anal. Calc, for Cu C30H43N7P2F12: C, 42.13; H, 5.07; N, 11.47;

Cu, 7.43. Found: C, 42.10; H, 5.10; N, 11.22; Cu, 7.37.

Equilibrium Constant Measurement

Axial base equilibrium constants were determined in the manner

described in detail by Stevens.^ A solution of the iron(II) complex

was titrated with the appropriate axial base while monitoring spectral

changes. Carbon monoxide binding constants were obtained using the

67

method of Stevens.^ A brief outline of the method and changes in the

system are described below. Equilibrium constants are determined by

monitoring electronic spectral changes as a function of carbon monoxide

partial pressure. The partial pressure of CO is regulated through the

use of four precision rotameters which control the flow of analyzed

mixtures of CO and nitrogen. Tank 1 contains prepurified nitrogen.

Tanks 2 and 3 contain Matheson primary standard grade CO/N2 mixtures

analyzed to contain 0.251% CO and 5.002% CO, respectively. Tank A con

tains pure CO. To remove residual traces of 0^, an L. C. Company, Inc.

high capacity oxygen trap was installed in each gas line. Rotameters

1, 2, and 3 were calibrated for Ng. (The calibration curve is included

in appendix A.) The partial pressure of CO is determined using equa-3 -1tion 16, where f__ and f„ are the flow rates in cm «mln of the CO CO n2

(or CO/N2 mixtures) and pure N2 streams, %C0 is the analyzed percentage

of CO in the CO/N2 mixture, Patm is ti.e atmospheric pressure and P^ is

the vapor pressure.

P = ------- — ^£2 (P - P ) (16)C0 fC0 + fN2 100 3tm V

Accessible CO partial pressures are from 0.05 to "*760 torr. The

absorbance data are collected at several wavelength for different par

tial pressures of CO at a controlled temperature.

The equilibrium constant, K c q , for the formation of a 1:1

iron-CO adduct, as expressed by:

FeLB + CO FeLBCO (17)

68

was determined using a non-linear least squares technique developed by

Dr. E. V. Dose of these laboratories. The equation to be fit was

. . (GFeLBCO“ EFeLB5KCO[FeLB'lO (PCOJ .A = Ao + ~ a + - Co (pCo » (18)

The non-linear least squares program minimized the function:

<19)

where the weight w is inversely proportional to the variance of

The program calculates estimated standard deviations (esds) for the

refined parameters Kc0 and eFeLBC0*

X-Ray Crystallographic Procedures

The following is a general description of the procedure used in

the single-crystal X-ray structural analyses. Specific details for

each structure will be described individually. A crystal was selected,

mounted on a quartz fiber and coated with several thin layers of epoxy.

After mounting the fiber in a goniometer head, the crystal was optically

centered on a Syntex PI four-circle computer controlled diffractometer.

The radiation used was graphite monochromatized MoKa. From a rotation

photograph 15 reflections were located and the setting angles automati

cally optimized. These reflections were then indexed using the auto

indexing program. After the approximate cell constants were obtained

by a least squares calculation, approximately 1000 reflections in the

range 10° < 20 < 30° were rapidly collected and about 30 moderately

69

Intense reflections were selected and used to obtain more precise cell

constants after optimizing their setting angles.

Intensity data were collected using the to-29 scan technique.

The scan rate was varied linearly according to the intensity of the

reflection; weak reflections were collected slowly, intense reflections

were measured rapidly. Backgrounds were counted for a total of one-half

the time spent in each scan. Ten check reflections were monitored

periodically to follow the condition of the crystal in the X-ray beam.2The estimated standard deviations in the reduced intensities (F ) were

calculated using

«J(F2) = (r/Lp)[S + G2 (B1 + B2) + (pl)2 ]1/2 (20)

where r is the scan rate, 1/Lp is the Lorentz-polarization correction,

^1* ^1* ^2 are t*ie scan anc* background counts, G is the ratio of scantime to total background counting time, I is the net intensity, p is a

factor chosen as 0 .0 2, included in a term presumed to represent that

component of the total error expected to be proportional to the dif-8Xfracted intensity. Systematic absences were deleted from the data

set and multiply measured reflections were averaged. The R factors

for multiply measured reflections were defined as

R1 = E | |Fj I - | Fav | | /£ | Fav | (21)

and

R2 - {lj[(Fj2 -Fav 2)2/a2 (Fj2 -Fav2)]/Z[FavA/a2 (Fj2 -Fav 2)]}1/2 (22)

70

The E-statistics were examined to verify the space group as centric or

acentric.

The phase problem was solved in each case using the heavy atom

Patterson method. Host calculations were carried out using the CRYM82crystallographic computing system. Some calculations for

[Fe{(m7Xylyl(NHEthi)2)Me2 ll6]tetraeneN^}Cl]Cl*2CH^0H were performed83using the program ORFLSE. Refinement was carried out using standard

Fourier and least square techniques. Convergence was considered

achieved when no shifts in the final least-squares cycle exceeded

0.25 esd. The discrepancy indices were defined as

R = 2 | | Fo | - | Fc | | /I | Fo | (23)

R = {Zw(F, 2 - F 2)2/XwF 4}1/2 (24)W O C O

2 2where w = 1/a (F ) o

GOF = {Zw(F 2 - F 2)2/(n - n )}1/2 (25)o c o p

The atomic form factors for all atoms except hydrogen were taken from84the International Tables while that of hydrogen was taken from

85Stewart et al. The function minimized in the least squares refine-

2 2 2 2 ments was Ew(F - F ) except for ORFLSE in which Etf(F - F ) was o c o cminimized (for intensities above 3a(I) only).

[Ni{(m-Xylyl(NHEthi)„)Me„[16]tetraeneN,}]„(PF,),-4CH„C0CH„. ---------■*— *--------- 2 Z---------------4— 2--- 6—4---- 3---- 3Crystals of the complex were grown by slow evaporation of an acetone

solution. It was apparent upon examination of the crystals that two

distinct forms were present; one form was in the shape of a square

71

pyramid, the other form was needles. A square pyramidal crystal of

dimensions 0.25 x 0.25 x 0.35 mm was mounted, centered, and indexed

giving an orthorhombic cell having the parameters listed in table 3.

The space group was uniquely identified as Pbca by examination of sys

tematic absences (hkO, h = odd; h0£, 2, = odd, Okf,, k ** odd). The cal-3culated density of 1.49 g/cm (based on four acetone molecules of

crystallization per molecule of complex) for Z = 4 is considerably3greater than the observed density of 1.44 g/cm measured by the flota

tion method in a benzene-bromoform mixture. These values are in dis

agreement because the observed density is an average of the densities

of the two different crystalline forms.

The intensity data were measured in two shells; the first,

inner, shell contained 8776 reflections having 4° < 20 < 60°, while the

second shell remeasured all reflections in the positive hklt octant

between 30** and 50° in 20". The check reflection 313 diminished greatly

in intensity but all others remained essentially constant throughout the

data collection. The total number of reflections measured was 13057

(excluding check reflections) of which 7104 were unique. A total of22123 reflections had intensities greater than 3a(F ) above background.

The R-factor for multiply measured reflections was 0.113, unusually

high due to the high percentage of weak data. After application of the

Wilson plot, the average temperature factor, 2B, was 5.36 and the scale

factor, K, was 3.9518. The E-statistics were consistent with the cen

tric space group as shown in table 4 .

In early stages of refinement, only the data having I > 2a(I)

were used. The position of the nickel(II) was found from a sharpened

72

TABLE 3SUMMARY OF CRYSTALLOGRAPHIC DATA

FeC28H44N6°2Cl2 FeC28H40N6C1PF6 FeC39H47N702P2F12 NiC32H48N6°2P2!Space Group P21 pT Plj/m Pbca

a 10.607(2) 8.575(1) 10.115(3) 16.353(5)b 14.656(3) 13.788(2) 16.468(3) 20.173(5)c 10.132(2) 14.383(2) 11.926(3) 24.157(6)a 89.99(1) 84.15(1) 90.02(2) 89.97(2)B 79.79(2) 76.32(1) 112.78(2) 89.95(2)

y 90.00(1) 104.80(1) 90.02(2) 90.02(2) 'V 1550.21(50) 1573.64(38) 1831.71(75) 7969.25(373)

pcalc 1.34 1.47 1.60 1.49

pobs 1.34 1.46 1.59 1.44

p(Mo-0.71069 A) 6.95 6.76 5.96 6.54T °C 20(1) -73(3) -40(2) 20(1)Scan technique e)-20 w-20 e>-20 tij-28Region measured 4* < 26 < 50* 4'<20 <60* 4° < 20 < 55° 4° < 20 < 60°Reflections measured 6418 12956 6826 13057Independent reflections 2874 9230 4964 7104Reflectionsi 1 > 3oI 1509 3042 2332 2123Parameters 353 428 499 496D:P (30) 8.1(4.3) 21.6(7.1) 99(4.7) 14.32(4.28)R (30) .103(.057) .154{.066) .105(.054) .218(,135)Rw (3c) .Q92(.079) .112(.0B5) .120(.104) .243(.214>GOF (30) 1.33(1.58) 1.21(1.60) 1.17(1.40) 2.09(3.21)

TABLE 4

E-STATISTICS

Centric <NiC32H48N6°2P2FI2>2 FeC28H44N6°2C12 FeC28H40N6C1PF6 FeC30H47N7°2P2F12 Acentric

<|e |> 0.798 0.778 0.871 0.802 0.811 0.886

<E2> 1.000 1.016 1.003 1.021 1.017 1.000

< f E2-11> 0.968 0.964 0.731 0.920 0.910 0.736

|e | > i , % 32.0 33.8 38.9 36.1 34.4 36.8

|e| > 2, % 5.0 4.1 1.0 3.3 3.9 1.8

|e | > 3, % 0.3 0.3 0.0 0.2 0.1 0.01

u>

74

Patterson map. A structure factor calculation based solely on the

nickel(II) coordinates yielded an R-factor of 0.626. (For a random

distribution of atoms in the centric cell, this value was calculated to 86be 0.83.) Five successive Fourier syntheses and structure factor cal

culations resulted in the location of all other atoms in the molecule.

After several cycles of least squares refinement, the two acetone mole

cules of crystallization were located from a different Fourier map.

The populations of the acetone molecules were allowed to vary and con

verged at populations near 1.0. Therefore each acetone was included as

having full occupancy. Further least squares refinement of coordinates

and isotropic temperature factors converged at R B 0.192. Least squares

refinement was continued using anisotropic temperature factors for all

non-hydrogen atoms until convergence was attained at R a 0.161,

R = 0.223, and G0F = 2.89 for the data having I > 2a(I). Three cycles wof least squares were calculated using all of the data converged at

R = 0.218, R^ = 0.243, and G0F = 2,09. A final difference map contained* °3a number of peaks of approximately 1.5 electrons/A , particularly in the

regions of the hexafluorophosphate anions. Attempts to assign atoms to

these peaks and variation of the fluorine populations did not signifi

cantly improve the quality of the structure. Appendix B contains a sum

mary of the final positional and thermal parameters as well as a listing

of the calculated and observed structure factors.

It is apparent from the discrepancy indices that refinement did

not converge to within normally acceptable limits. The most obvious

reasons for such poor refinement Is the large number of weak data col

lected. Disorder in the hexafluorophosphate anions and high thermal

75

motion in the molecule also contributed to the refinement difficulties.

Despite these difficulties, significant important information was

obtained from the structural analysis as regards the constitution of the

molecule which had been previously incorrectly assumed. A discussion of

the structural details begins on page 131.

fFe{(m-Xylyl(NHEthi)jMe,, [16]tetraeneN^}C1]Cl* 2CH3OH. Crystals

of the complex were obtained by slow evaporation of a methanolic solu

tion. The red crystals appeared to be stable in air for several hours

before cracking and turning dark brown. A crystal of dimensions

0.25 mm x 0.25 mm x 0.15 mm was selected and mounted using epoxy cement.

Precession photographs indicated a monoclinic crystal system and the

systematic absences (OkO, k = odd) were consistent with the space groups

P2^ and P2^/m. After centering and indexing, the cell constants were

obtained (table 3). The calculated and experimental densities agree

at 1.34 g/cm assuming Z = 2. Three dimensional intensity data were

collected for reflections having 4° < 20 < 50° for all positive h and

k and all positive and negative £jof the 6047 reflections measured

(excluding check reflections), 2874 were unique and 1509 had I > 3a(I)

above background. The R-factor for multiply measured reflections was

0.086. After scaling of the data the value of 2B was 5.62 and K was

1.1188. The E-statistics listed in table 4 were consistent with an

acentric space group and thus P2^ was selected.

The coordinates of the iron(II) ion and one of the chlorine

atoms were determined from a sharpened Patterson map and the y coordi

nate of the iron atom was fixed at 0.25 due to the polar nature of the

76

space group. At this point, the R-factor was 0.416. From successive

structure factor calculations and Fourier maps, the remainder of the

non-hydrogen atoms of the molecule, except for the methanol solvate

molecules, were found, reducing R to 0.235. Several cycles of least

squares refinement were executed and a difference Fourier map was

generated. Electron density corresponding to the methanol molecules

was found, but the refinement resulted in high thermal parameters which

indicated disorder. The populations of the methanol molecules were

allowed to vary and approached values of 1.0. Refinement was continued

using anisotropic temperature factors for all non-hydrogen atoms of

the cation, anion, and solvent molecule. Hydrogen positions were cal

culated after each least squares cycle and were included in structure

factors calculations but were not refined. At convergence, the dis

crepancy indices were R = 0.103, Rw = 0.092, and GOF = 1.33 based on

all 2874 reflections. The corresponding indices based on the 1509

reflections having I > 3a(I) were R ** 0.057, R^ = 0.079, and GOF = 1.58.

The data to parameter ratios were 8.1 and 4.3 respectively for the two

data sets. A final difference Fourier map indicated some residual elec

tron density in the vicinity of the methanol molecules. Appendix B con

tains a listing of the final positional and thermal parameters as well

as the calculated and observed structure factors.

The y-coordinates of all atoms were inverted about y = 0.25 to

determine if there was a handedness to the cell contents. After

several least squares calculations, the model showed no Improvement

over the original one. The unresolved disorder in the methanol mole

cules is undoubtedly responsible for the relatively high estimated

77

standard deviations in the bond lengths and angles. The model can be

considered as a good one for the complex, however, and bond distances

and angles can be interpreted with confidence. A detailed discussion of

the structural features begins on page 165.

[Fe{(m-Xylyl(MeNEthi)„)Me„[161tetraeneN^lCl](PF^)• Crystals of

the complex were grown slowly from a solution of acetonitrile and

ethanol. A crystal of dimensions 0.11 mm x 0.21 nun x 0.29 mm was

mounted and centered. Precession photographs indicated a triclinic

crystal system. The crystal was cooled to -73 + 3°C and maintained at

that temperature throughout measurement of the cell constants and data

collection. After centering and indexing the cell constants were

determined (table 3).

Three-dimensional intensity data were collected for reflections

having 4.0° < 20 < 60.0° for all positive h and all positive and nega

tive k and A. The intensities of ten check reflections were monitored

every 190 reflections and they remained essentially constant throughout

the data collection. Of 12956 reflections measured, 9230 were unique2and 3042 had intensities greater than 3a(F ) above background. Psi

scans were measured and these indicated an absorption correction was

necessary. This correction was made using the program of Beno and 87Christoph using psi scans for eight sets of indices. The R-factor for

multiply measured reflections was 0.075. After Wilson scaling of the

data the values of 2B and K were 3.67 and 1.2893 respectively. The

E-statistics listed in table 4 were not conclusively centric or acen

tric but agreed more closely with the centric values and thus the space

group PI was selected. (Refinement was also attempted in space group PI

78

with unsatisfactory results.) The position of the iron(II) was found

in a sharpened Patterson map and the rest of the molecule, Including

the ordered hexafluorophosphate anion, was found from two Fourier maps.

The structure was refined in the usual manner and finally converged

with all non-hydrogen atoms anisotropic and refined hydrogen positional

parameters. The final discrepancy indices were R = 0.154, Rw = 0.112,

and GOF ** 1.21 based on all 9230 reflections. Using the 3042 reflec

tions with I < 3(7(I) , R = 0.066, Rw “ 0.085, and GOF = 1.60. The data

to parameter ratios were 21.6 and 7.1 for the two data sets. A final

difference map showed no peaks having electron density greater than 0.6 °3electrons/A . Appendix B contains a listing of all positional and

thermal parameters as well as the calculated and observed structure fac

tors. A detailed discussion of the structural features begins on page

165.

[Fe{(1.5-Pent(NHEthi)2)Me„ T16]tetraenelQ (PY)(CO)1(PFC) C H ^ OH.

Crystals of the complex were grown by Dr. J. J. Grzybowski from a

methanolic solution. The red crystals grew together as clusters and a

small single crystal fragment approximately 0.10 mm x 0.15 mm x 0.25 mm

was selected for the X-ray structure analysis. The crystal was mounted

and cooled to -40 + 2°C. After centering, the crystal was indexed in

the monoclinic system with the cell parameters listed in table 3.

Three dimensional intensity data were measured for all positive h, all

k between -3 and +32 and all positive and negative I having

4f)° < 20 < 55.0°. Of the 6291 reflections measured (excluding check

reflections) 4964 were unique and 2329 had I > 3a(I). The R-factor for

79

multiply measured reflections was 0.067. Based on the systematic

absences (OkO, k *= odd) and centric E-statistics (table 4) the space

group P2^/m was chosen. Based on the known crystalline density and

cell volume, it was necessary for Z to equal 2. After Wilson scaling,

the values of 2B and K. were 4.12 and 1.2769. The molecule was there

fore restricted to lie on a crystallographic mirror plane at y = 0.25.

The coordinates of the iron(II) were determined from a sharpened

Patterson map. A structure factor calculation based solely on the

iron(II) coordinates yielded a value of R = 0.611. Standard refinement

procedures yielded all other atomic coordinates including the methanol

of crystallization and the ordered hexafluorophosphate anion. Aniso

tropic refinement of all non hydrogen atoms and refinement of hydrogen

positional parameters (including methyl groups) converged with discrep

ancy indices R = 0.105, R^ = 0.120, and GOF = 1.17 based on all 4964

reflections. The corresponding values based on the 2329 reflections

having I > 3a(I) were R = 0.054, R^ = 0.104, and GOF = 1.48. The data

to parameter ratios were 9.9 and 4.7 respectively for the two data

sets. A final difference Fourier map contained no peaks having elec-°3tron density greater than 0.6 electrons/A . Because the molecule was

on a crystallographic mirror plane, the populations of all atoms on the

plane were set at 0.5 occupancy. Thus only one-half of the other atoms

needed to be located with symmetry related atoms being generated auto

matically. Appendix B contains a listing of the final positional and

thermal parameters as well as the calculated and observed structure fac

tors. A detailed discussion of the structural features begins on page

198.

RESULTS AND DISCUSSION

The work described in this thesis involves the design, syn

thesis and characterization of a series of totally synthetic iron(II)

heme-protein models derived from dry cave ligands. A number of

unbridged nickel(II) precursor complexes were studied in some detail

and were used to form dry cave ligands through the development of a new

method of bridging. Monomeric and dimeric complexes were separated and

characterized by a number of methods including the single-crystal X-ray

structural analysis of one dimeric species. The ligands were removed

intact from nickel(II) and several were studied in detail. The ligands

were then chelated to iron(II) and the properties and reactivity of the

resultant complexes were examined. Two five-coordinate iron(II) com

plexes were subjected to X-ray structural analyses. Carbon monoxide

adducts were prepared and characterized, again through the use of X-ray

crystallography. Some preliminary equilibrium studies with carbon

monoxide were performed and reactions with dioxygen were examined.

Finally, two copper(II) complexes derived from dry cave ligands were

synthesized and characterized.

Unbridged Nickel(II) Complexes

In order to study some of the unbridged nickel(II) complexes in70further detail, three compounds first synthesized by Schammel were

prepared; [Ni{(MeNHEthi)2Me2 I ]tetraeneN^}](PFg)^,

80

81

[Nl{(Me2NEthi)2Me2 [16]tetraeneN4}3 (PF6)2 and [Ni{ (NHEthi) 2Me2 [16]-

tetraeneN^}](PFg)2 . In addition, three new complexes of this class were

prepared: [Nl{(n-BuNHEthi)2Me2 [16]tetraeneN4}] [Nl{(t^BuNHEthi)^-

Me2 [16]tetraeneN4}] 0PFg)2 and [Ni{(BZNHEthi)2Me2 [16]tetraeneN4 }]<PF^> £ *

The synthesis of the methylamine complex was modified by using an

excess of methylamine gas instead of methylamine hydrochloride. The

new compounds were prepared by adding a slight excess of two equiva

lents of n-butylamine, tert-butylamine or benzylamine to a solution of

[Ni{(MeOEthi)2Me2 [l6]tetraeneN4}](PFg)2 in acetonitrile to yield the

respective monoamine complexes.

The new complexes were characterized by elemental analysis and

all of the complexes were characterized by conductivity, IR, and 13C NMR spectroscopy and electrochemistry. The infrared spectra are

consistent with the proposed formulations and are very similar to those

observed by Scharamel. Some selected frequencies are summarized in

table 5. The IR spectra of n-butylamine and _t-buty lamine derivatives

are shown in figures 7a and b. The significant features are the N-H

stretching frequencies in the vicinity of 3400 cm ^ and the C = N and

C ■* C stretching frequencies in the 1500-1600 cm ^ region. The

benzylamine and ammonia derivatives show N-H regions more complex than

expected in the solid state spectra. In acetonitrile solution however,

the benzylamine complex shows the expected single stretch and the

ammonia complex shows two absorptions due to the symmetric and

anti-symmetric N-H stretching modes of the primary amine.

Molar conductance data for the complexes are listed in Table 5

and are consistent with their formulations as 2:1 electrolytes

TABLE 5

SELECTED INFRARED FREQUENCIES3 AND MOLAR CONDUCTANCES15 FOR UNBRIDGED NICKEL(II) COMPLEXES

Compound V (cm N-H } VC=C,C=N(cin 5 A ’. -1 n -I 2 onm mole cm

[Ni{ (NH2Ethi) 2Me2 [16] tetraeneN^ ] (PFfi) 2 3280,3 3395? 34753 1535, 1611, 1658 2713250,° 3360C

[Ni{(MeNHEthi)2Me2[16]tetraeneN4}](PF6)2 3397 1542, 1582 303

[Ni{ (Me2NEthi) 2Me2 [16] tetraeneN^}] (FFg) - ~ - 1528, 1563, 1602 292

[Ni((n-BuNHEthi)2Me2[16]tetraeneN^}](PF&)2 3373 1573 234

[Ni{(b-BuNHEthi)2Me2[16]tetraeneN^}](PF^) 2 3379 1562, 1613 258

[Ni{(BZNHEthi)2Me2 [16]tetraeneN4)](PF6)2 3360,3 3392,3 3298C 1565, 1608 275

0Spectra obtained fromnujolmulls on KBr plates.b “3Acetonitrile solutions of -1 x 10 molar at ambient temperature.cAcetonitrile solution.

00ho

83

4000 3000 2000 1500cm 1000II I tII

500

Figure 7. Infrared Spectra of a) [NiRn-BuiraEthiKMeJIfiltetraeneN.}] (PF,)„ and b) [Ni{(trBuNHEthi)2He2 [16]tetraenSN45j(PF6)2 * 6 2

84

in acetonitrile. (The acceptable range for 2:1 electrolytes in acetoni

trile is 220-300 ohm ^mol ^cra^.) NMR spectra were recorded for all

of the complexes in deuteroacetonitrile and are summarized in table 6 .

They demonstrate no unusual features and are in support of the struc

tural assignments made by Schammel. As examples, the spectra of the

n-butylamine and fr-butylamine are presented in figures 8a and b,

respectively.

Carbon-13 NMR spectra were measured for each of the complexes

and the data are summarized in table 7. The resonances have been

assigned through the use of broadband and off-resonance proton

decoupling techniques according to structure XVI. All of the complexes

show mirror symmetry which includes

the metal and the two central car

bon atoms of the saturated tri-

methylene linkages of the

macrocycle. There are several

features of the basic macrocycle

which are common to all of the

species. Two methyl resonances

are observed between 15 and 21 ppm

from TMS due to carbon atoms f and

h. Methylene resonances observed

between 29 and 31 ppm are assigned

to carbon atoms b and b^, those

between 51 and 57 ppm are due to

a and a"*, the methylene carbon

— N i - V

TABLE 6

PROTON NMR DATA3, FOR UNBRIDGED NICKEL(II) COMPLEXES

Compound Methyl Methylene Aromatic N-H

[Ni{(NHjEthi)2Me2 [16]tetraeneN4>](PFg)% 2.18, 2.32 3.05, 3.40C 7.50 6.75b

[Ni{(MeNHEthi)2Me2 [16]tetraeneN^}](PFfi) 2.10, 2.27, 3.03 3.12, 3.35c 7.47 6.95b

[Ni{(Me2NEthi)2Me2[16]tetraeneN4}](PFfi)2 1.85, 2.33, 3.20 3.02, 3.35° 7.40 -----

[Ni((n-BuNHEthi)2Me2[16]tetraeneN4}](PF^)2 0.93, 2.12, 2.30 1.03, 1.55,b 7.48 6.73b3.04, 3.38

[Nit(t-BuNHEthi)2Me2[16]tetraeneN4}](PF6> 1.67, 2.13, 2.47 3.07, 3.38c 7.53 6.63b

[Nit (BZNHEthi) 2Me2[16] tetraene^} ] (PF&) % 2.08, 2.32 1.97, 3.03, 7.35, 7.52 7.02b3.38,C 4.58

aChemical shifts given in ppm from TMS, concentrated solution in CD^CN.

bBroad.

CMultiplet.

8 6 4 2 t OP P m

Figure 8 . NMR Spectra of a) [Ni{(n-BuNREthi),Me ri6]tetraeneN.}](PF,)_ ajtd b) [Ni{(trBuNHEthi)2He2[16]tetraeneN^}](PF6)2 A 6 2

87

TABLE 7CARBON-13 NMR DATA FOR UNBRIDGED NICKEL(II) COMPLEXES

3 bCompound Chemical Shifts

[Ni{(NH2Ethi)2Me2 [16]tetraeneNA>]2+ 169.4 , 167.5, 161.3, 111 •5,56.2, 51.3, 30.4, 29. 9, 20.5,19.8

[Ni{ (MeNEthi) 2Me2 [16] tetraeneN^,} ]2+ 170.0,,c 168.3, 159.9, 112.4,°56.2, 51.5, 31.8,c 30 .6 , 30.2,20.9, 15.5C

[Ni{ (MegNEthl) 2Me2 [16]tetraeneN^}]2"** 173.7,, 168.0, 159.5, 111 .4,56.2, 51.1, 44.4, 30. 6, 30.5,20.8 , 19.4

[Ni{(n-BuNHEthi)2Me2 [16]tetraeneN^}]2+ 168.7.,° 168.1, 159.9, 112.3,°56.2, 51.3, 45.8 ,c 31 .7, 30.5,30.1, 20.6, 15.3,C 13 .8

[Ni{(BZNHEthi)2Me2 [16]tetraeneN^}]2+ 168.6,, 168.2, 160.3, 137 .2 ,129.9,, 129.1, 112.4,C 56 .2 ,51.3, 49.6 ,C 30.4, 30 .0 , 21.0,16.3C

[Ni{(t-BuNHEthi) 2Me2 tetraeneN^ } ] 2+ 167.9,, 167.4, 159.6, 112 .7,56.5, 55.9, 50.9, 30. 1 , 29.7,20.3, 16.2

aAs PFg— Salts.

^CD^CN solution, ppm relative to TMS.

cBroad.

88

atoms bonded to nitrogen. A single resonance near 110 ppm is due to

the carbon atom d. The resonance near 160 ppm splits into a doublet in

the off-resonance spectrum and is therefore assigned to carbon atom c.

Two resonances between 167 and 174 ppm are due to carbon atoms e and g.

The resonances due to the amine substituents are found in the

expected regions. The N-methyl resonances are observed at 31.8 and

44.4 in the monomethyl- and dimethylamine complexes respectively. The

benzylic carbon atom of the benzylamine complex is found at 49.6 ppm

and three aromatic resonances occur between 129 and 138 ppm. Appar

ently there is accidental overlap of two aromatic peaks since only

three of the expected four resonances are observed. The n-butylamine

complex shows the methylene carbon atom bonded to nitrogen at 45.8 ppm.

The other two methylene carbon atoms are in the same region as the cen

tral carbon atoms of the macrocycle side chains and are not clearly *

resolved. The alkyl methyl group resonance occurs at 13.8 ppm. The

tert-butylamine complex exhibits a quarternary carbon resonance near

56 ppm and a single methyl resonance at 29.7 ppm.

As shown in figure 9, the spectrum of the dimethylamine deriva

tive at 300 K exhibits only a single resonance due to the amine methyl

groups at 44.4 ppm. At 238 K however, two distinct resonances appear

at 46.4 and 42.4 ppm. In contrast to the above example, the spectrum

of the n-butylamine complex shows the presence of two, or perhaps

three, components in solution at 238 K. Multiple resonances appear in

the regions corresponding to carbon atoms d, g, i and f or h, as shown

in figure 10. The spectra of the benzylamine and methylamine complexes

resemble that of the n-butylamine derivative. In the case having the

89

Figure 9

Mu, hhihsK'110 0

. NMR Spectrum of [Ni{(Me2NEthi)2Me2 [16]tetraeneN,}](PFfi)2at a) 300K, b) 278 K, c) 238 K.13,

t

aJ \* Myl1 Al***,

b P L .I80 40ppm

1 1Figure 10. C NMR Spectrum of [Ni{(n-BuNHEthi)9Me9 [lbJtetraeneN.}](PF,)0

at a) 300 K, b) 238 K. L 1 * 6 2 u>o

0120 8 0 4 01 8 0 ppm

13Figure 11. C NMR Spectrum of [Ni{(t-BuNHEthi)„Me9 [16]tetraeneN.}](PF )i- L 4 6 2

92

bulky tert-butyl substituent, all of the resonances are sharp with no

indication of broadening (figure 11).

The above observations indicate that there is rotation about

the C-N bond and that the barrier to rotation is dependent upon the

nature of the nitrogen substituents. This suggests that there is a

favored orientation of a single nitrogen substituent relative to the

macrocycle. As shown in structures XVII and XVIII, the R group can be

oriented in a direction perpendicular to the macrocycle plane, XVII or

parallel to the planej XVIII. When the nitrogen substituents are

H

X5zn XVTTT

identical, these orientations are indistinguishable and the only

resonances which should be affected by C-N bond rotation are those due

to the amine methyl carbon atoms. When only one substituent is present,

however, structure XVII would be preferred since interaction between

the macrocycle and the smaller hydrogen atom would be less than that

interaction between the macrocycle and the nonhydrogen R group. When R

is very bulky as in the tert-butyl case, structure XVII is highly

93

favored and only that single conformer is present in significant

amounts at ambient temperatures.

The presence of structure XVII is accompanied by methyl13resonances which are separated by about 5 ppm in the C NMR spectrum.

When structure XVIII becomes important the methyl resonances are13separated by about 1 ppm. This feature of the C NMR spectra will be

of importance in discussions of the bridged "dry cave" complexes.

As mentioned above, the low temperature spectrum of the

nybutylamine complex indicates the presence of two or perhaps three

species in solution. These can be accounted for on the basis of struc

tures XVII and XVIII. It is possible for both amines of the macrocycle

to have structure XVII or for both to have structure XVIII. In these

cases the molecule has a mirror plane of symmetry. It is also possible

for one amine to have structure XVII and the other to have structure

XVIII. In such a case, mirror symmetry would be lost and each carbon

atom of the molecule would be unique.

It is important to note that structures XVII and XVIII are dis

tinguishable when there is a single nitrogen substituent. This clearly

accounts for the fact that several resonances are affected by the C-N

bond rotation in the complexes described above.13Variable temperature C NMR spectra also show that at reduced

temperatures the rate of boat-chair interconversion of the six-membered

rings containing the saturated trimethylene linkages of the macrocycle

decreases, causing the resonances of the methylene carbons to broaden.

Electrochemical studies performed on these unbridged nickel(II)

complexes demonstrate some very interesting features. All of the

complexes exhibit a reversible oxidation between +0.77 and +0.90 V ver

sus Ag/Ag"** reference. This oxidation has been assigned to the 2+ 3+ 9 0Ni /Ni couple and is the oxidation of interest in this work.

A second oxidation which is reversible in only three of the complexes90occurs between +1.1 and +1.4 V and is assigned to a ligand oxidation.

A third oxidation occurring in the vicinity of +1.6 V is irreversible.

In each case an ill-defined, irreversible reduction wave occurs at

<-1.8 V. Electrochemical parameters were obtained from rotating elec

trode voltametry and cyclic voltametry using a platinum disk electrode

and are summarized in table 8 . The behavior of [Ni{ (MeNHEthi^l^-

[16]tetraeneN^}](PFg)^ is typical and the cyclic voltamagram and

rotating electrode voltamagram are shown in figure 1 2.

From the E ^ ^ values for the first oxidation of these com

plexes it is apparent that the electron donating ability of the nitrogen

substituent directly influences the oxidation potential at the metal

center. The ammonia derivative has the most positive E ^ ^ since hydro

gen is the least donating substituent in the series. Substitution by

one methyl group lowers by 70 mV as the metal center responds to

the substituents’ inductive effect. Substitution by a second methyl

group results in further lowering of E ^ ^ by 55 mV. Although the dif

ferences are small, E a p p e a r s to change in the following sequence:

CH3 < n-Bu < BZ < _t-Bu < H .

From these results it is apparent that careful selection of

substituents on the amine nitrogen atoms can be used to control the

polarity and steric environment in the vicinity of one of the metal

TABLE 8

ELECTROCHEMICAL DATA FOR THE UNBRIDGED NICKEL(II) COMPLEXES

Compound El/2 ’ V lE3/4“ El/4l' mV Ep, V

[Ni{(Me2NEthl)2Me2[16]tetraeneN^}](PF&) +0.745 70 +0.775+1.080 70 +1.085

[Ni{(MeNHEthi)2Me2 [16]tetraeneN4}](PFfi)2 +0.800 80 +0.835+1.260

[Ni{(^BuNHEthi)2Me2[16Jtetraenr" }](PFg) 2+0.810 75 +0.835+1.240

[Ni{ (BZNHEthi) 2Me2 [ 16 ] tetraene^) ] (PFg) 2 +0.830 85 +0.835+1.305

[Ni{(t-BuNHEthl)2Me2[16]tetraeneN^}](PF6)2+0.845 75 +0.870+1.190+1.400

[Nl{(NH2Ethi)2Me2 [16]tetraeneN^}](PF6)2 +0.870 80 +0.900+1.420

VOUi

02 000.8 0.41.0

0.50.7 030.9

Figure 12. Cyclic and Rotating Platinum Electrode Voltamagrani3 for [Ni{(MeNHEthi)^Me^[16]tetraeneN^}](PF^)^ VOov

97

axial coordination sites as well as provide a tool for "fine tuning" of

the oxidation potential at the metal center.

Synthesis of Nickel(II) Dry Cave Complexes

The nickel(II) dry cave complexes were prepared by one of two

synthetic routes. The first was that developed by Schammel in which an

appropriate diamine was reacted with the methyl vinyl ether complex as70shown in scheme I. The specific diamines used in this study were

m-xylylenediaraine, £-xylylenediamine and N,N'-dimethy1-1,6-hexanedi-

amine.

Scheme 1

+ V < "H

R

An alternate synthetic method was developed to expand the num

ber of available dry cave compounds. This route is shown in scheme II.

RScheme H

-+• 2 RNHe _,CH?CN

o o

+ 2 CH3OH + 2 No X

CH

\oCO

99

An unbrldged Ni(II) complex having the desired nitrogen substituent is

first synthesized as described in the previous section. The amine

derivative is then deprotonated and combined with an appropriate

dielectrophile under conditions of high dilution in refluxing acetoni

trile. This reaction was first demonstrated for several instances in

which a xylyl dihalide was used as the dielectrophile. Stevens extended66the reaction to include aklyl ditosylates as the bridging reagents.

The alternate route described above offers several advantages

over the original method. First it allows for the incorporation of a

wide variety of bridgehead nitrogen substituents having widely varying

electronic and steric properties. Second, the availability of desired

dielectrophiles is much greater than that of diamines which must often

be prepared using the tedious Gabriel synthesis. Third, the effective

length of the bridging group is two atoms shorter since the bridge

nitrogens have already been incorporated into the macrocyclic struc

ture. Through the combined use of these two synthetic routes many

structural variations are accessible.

The structure of the dry cave complexes was first confirmed by

the solution of the crystal structure of [Ni{(p-Xylyl(NHEthi)2^Me2 [16]-72tetraeneN^}](PFg)^ by Christoph and Mertes as described in the intro-

13duction. However, it soon became apparent from C NMR and HPLC results

that isomers existed for several of the complexes. One possible type of

isomerism is similar to that described for the unbridged complexes. As2shown in structure IXX the bridge can be bonded to the planar sp

hybridized bridge nitrogen in such a way that the bridge projects in a

direction parallel to the macrocycle plane. This results in the

E X XX

bridge being forced away from a position centered over the metal; this

is called the "lid-off' isomer. Alternatively, the bridge can bond to

the nitrogen in such a way that its linkage projects perpendicular to

the plane (structure XX). This centers the bridge almost directly

over the metal, producing a taller, narrower cavity; such isomers are

denoted as "lid-on." To date there are no complexes known which exist

in both "lid-on" and "lid-off" conformations and there is no evidence

to indicate that interconversion between these two isomeric forms is

an important process.

Although "lid-on", "lid-off" isomerism has not been observed,

several of the complexes do exhibit both monomeric and dimeric forms.

Because spectroscopic, analytical, and physical data did not distin

guish between monomers and dimers, considerable confusion as to the

identity of the species of identical chemical composition hindered the

progress of research on these compounds. Considerable chemical evi

dence supporting the existence of dimers had been accumulating for some

time, but it was not until the solution of the herein described crystal

101

structure of a dimeric nickel(II) complex that concrete evidence for

dimers was finally obtained. The dimer consists of two macrocyclic

groups joined by two meta-xylene bridges as shown in figure 13. The

detailed structural information will be discussed in later sections of

this work.88Dr. J. Grzybowski has verified the generality of the dimeric

species by the selective synthesis of a dimer linked by pentamethylene

bridging groups. This was accomplished by first reacting the methyl

ated starting material with an excess of 1 ,5-diaminopentane to yield

the unbrldged complex as shown in scheme III. Reaction of this complex

with a second equivalent of the methyl vinyl ether complex gives the

desired dimeric complex which is identical in all properties to the

dimer prepared under normal bridging conditions.

* Bri dge Nitrogen

Pla ne

Figure 13. ORTEP Drawing of meta-Xylyl Bridged Dimer.

Scheme IHo c h 3

+ 2H 2N(CH2)6NH2 ►

c h 3

(CH2)5N H |

(CH^gNHg

Aru

c h 3c n

2N qOCH3

103

104

The following notation will be used for describing the bridged

complexes. As shown in structures XXI and XXII the plane refers to

the chelating nitrogen atoms of the macrocycle. The bridge nitrogens

are those which are external to the parent macrocycle. (NH)2 and

(NMe)2 refer to the bridge nitrogen substituent as hydrogen or methyl,

respectively. The rest of the acronym refers to the linkage between

the bridge nitrogen atoms. (Refer to Table of Abbreviations).

The combined efforts of several workers have resulted in the

preparation of complexes, consisting of monomers or dimers, with a wide91variety of bridging groups. Both monomeric and dimeric species can be

generated with the following bridging groups: (NH)2Pxyl, (NH)2Mxyl,

(NH)2 (CH2)^, (NH)2 (CH2)3 and (NH^Fl, whereas only monomeric complexes

have been achieved with ( N H ^ C C i y ^ (NMe)2 (CH2)6, (NMe)2 (CH2)5 »

(NHe)2 (CH2)7 , <NMe)2 (CH2)8 and (NR)„Mxyl (R = -CH3 , -(CH^CIhj,

-CHgCgH^) as the bridging agents. L.. .usively dimeric complexes result

when (NMe)2Pxyl, (NH)2Duryl, (NMe)2buryl, or (NMe)2 (CH2)3 are used as

bridging groups. The above information is summarized in figures 14 and

15. The complex generated from 9,lO-bis(chloromethyl)anthracene is

unusual in many respects and is not readily classified as a monomer or

dimer on the basis of present evidence, and warrants separate discus

sion.

It is not yet entirely clear which factors control whether mono

mers or dimers will be formed, but some trends are apparent. The only

complexes which show both monomeric and dimeric products as a result of

the bridging reaction are those formed using primary diamines. As

expected, elevated temperatures and high dilution reaction conditions

105

R

R= -H , -CH3, -(CH2)3 CH3, *CHz C6H5 ,

CH2".

_ (C H 2JrT’ n= 4 - 8 ,

Figure 14. Summary of Established Monomeric Nickel(II) Dry Cave Complexes

favor the formation of monomeric complexes, particularly with the m- and

£-xylyl bridges. Since the probability that two singly condensed macro

cycle plus diamine fragments will persist and encounter each other is

reduced under these conditions, closure of the bridge to the remaining

vinyl ether site on the macrocycle predominates.

Steric factors are important for cases in which the substituent

on the nitrogen is larger than hydrogen. The fact that only one or the

other of monomer or dimer is formed exclusively suggests a favored

orientation of both the R substituent and the bridge. Interactions

between the bridging group and macrocycle appear to be important for

cases which have a particularly bulky bridging group; i.e., durene.

R = -Hf —CH3 .

’ — H ^ — C H 2—I I

“i 3 - c h 2~

—(CH2)n- . n-3,4,5.

ch2— *

Figure 15. Summary of Dimeric Nickel(II) Dry Cave Complexes

106

107

Although these trends are consistent with the observed results, it is

still difficult to predict which forms will result for a given bridging

group. The fact that the (NH^CCI^)^ bridge forms only a monomer

whereas the (NH^CCH^)^ a**cl (NH^CCI^)^ bridges form both monomers and

dimers clearly illustrates this difficulty.

Of the six crystal structures which have been solved for mono

meric dry cave complexes, only one example of a "lid-on" structure has

been observed. That structure is of the chloro-meta-xylene bridged

iron(II) chloride complex which will be discussed in detail in later

sections. Based on information obtained from the other crystal struc

tures and information obtained from physical measurements and reac

tivities, it is reasonable to conclude that the configuration for all

monomeric bridged complexes are of the "lid-off" type except for

(NH^Mxyl and (NHJ^FI, which are "lid-on".

The "lid-on", "lid-off" r.:lature can also be applied to the

orientation of the linking groups about the bridge nitrogen atoms of

the dimeric complexes. The structure of the dimeric meta-xylene bridged

nickel(II) complex shows a "lid-on" type of arrangement, i.e., the

xylene groups rise in a direction essentially perpendicular to the

macrocycle N. plane. Steric considerations, as well as some features 13of the C NMR spectra of other dimeric complexes, suggest that all of

the dimers can be expected to have a similar "lid-on" configuration.

Separation of Monomeric and Dimeric Complexes

In order to properly characterize the monomeric and dimeric

complexes when they were formed concurrently, it was necessary to

108

develop special separation techniques. In the cases of the (NH)2 (0112)41

(NH)2(0112)5 * (NH^Mxyl, and (NH)gFluorene bridged materials the most

efficient separation was achieved by fractional crystallization from

mixed solvents (ethanol/acetonitrile). The dimer precipitates from the

solution first, the monomer is then isolated from the filtrate by fur

ther addition of ethanol followed by volume reduction.

Fractional crystallization was unsuccessful in the case of the

(NH^Pxyl derivative. However, the two species were reasonably well

separated on a column of neutral Woelm aluminia upon elution with ace-

tonitrile. The dimer elutes first as a dark yellow-brown band followed

by the monomer. This technique was also applied with good success to

the meta-xylene system.

The purity of the separated materials was monitored by means of

high performance liquid chromatography on a C-18 reverse phase column

operating at 1000 psi using a solvent mixture of 20% acetonitrile in

water. The conditions for separation of the species were optimized by 92Jackels for the (NH^CCn^)^ system and were as well satisfactory for

the (NH^Mxyl and (NH^Pxyl cases. The monomer eluted first as a very

sharp band followed by the dimer as a very broad band. A representative

chromatogram for an equal mixture of the (NH^Mxyl monomer and dimer is

shown in figure 16. This separation technique would be ideal if a pre

parative scale instrument were available.1 13The separation can also be monitored by use of *H and C NMR

spectroscopy. This will be described in the following sections.

109

30 20

Figure 16. HPLC Chromatogram Showing Monoraer-Dimer Separation for the meta-xylene Bridge Nickel(II) Complexes.

110

Characterization of Monomeric Nickel(II)Dry Cave Complexes

Characterization data for the complexes having polymethylene66bridges are summarized in Stevens' dissertation and will not be dis

cussed further here except for the case of the (NMe^CCI^Jg bridged

derivative.

Satisfactory elemental analyses were obtained for all of the

complexes as shown in the experimental section. Infrared spectra are

rather uninformative but do have a few characteristic features. Repre

sentative spectra of the (NH^Mxyl and (NMe^Mxyl bridged complexes are

shown in figures 17a and b. A sharp N-H stretching absorption is

observed in the vicinity of 3400 cm ^ for the (NH^Mxyl and (NH^Pxyl

bridged complexes. As expected, this band is absent . , complexes having

a nonhydrogen substituent on the bridge nitrogen. Luuds due to C = C

and C = N stretches are found in the vicinity of 1500-1600 cm \ Some

of these results are shown in table 9. In all of the spectra of the

xylyl bridged complexes there are bands due to the aromatic system in

the region from 600-800 cm \

Molar conductances were measured in acetonitrile and are con

sistent with the assignment as 2:1 electrolytes. These data are sum

marized in table 9. For the complexes having (NH^Mxyl, (NH^Pxyl* and

(NMe^Mxyl bridges, the conductance was measured as a function of con-93centration. Onsager plots were generated from these data as shown in

figure 18. The relevant data are collected in table 10 and show a slope

ranging from about 450-500 ohm ^ mole c m ^ and a value of Aq of

Figure

111

I i4000 3000 2000 cm 1500-i 1000 500

17. Infrared Spectra of a) [Ni{(m-Xylyl(NHEthi)2)Me2 [16]tetraene- V ] ( P F 6)2 and b) [Nit<HrXylyl(MeNEthi)2)Me2 [16]tetraeneN4}]<PF6>2

TABLE 9

SELECTED INFRARED FREQUENCIES AND MOLAR CONDUCTANCE DATA* FOR MONOMERIC DRY CAVE NICKEL(II) COMPLEXES

Compound V H (cm_1) VC=C,C=N(cm 5 A, ohm ‘'mole

[Ni{(m-Xylyl(NHEthi) [16] tetraeneN^}] (PFg) 2 3422 1575 271

[Ni{ (mrXylyl (MeNEthi) 2 ^ e 2 tetraeneN^ } 1 ) 2 ----- 1615, 1555 300

[Ni{(m-Xylyl(n-BuNEthi)2)Me2 [l6]tetraenc j (PFg)2 ----- 1610, 1550 259

[Ni{(ra-Xylyl(BZNEthi)2)Me2[16]tetraeneN4 j](PFg)2 ----- 1615, 1588, 1535 277

[Ni{ (£-Xylyl(NHEthi)2)Me2 [16]tetraeneN^}](PF6> 2 3405 1628, 1549 286

[Ni{(1,7-Hept(MeNEthi)2)Me2[16]tetraeneN^}](PFg)2 ----- 1608, 1550 292

3> -3Acetonltrile solutions of "*1 x 10 molar at ambient temperature.

113

70

01oE

teJCom0<1<*

60

50

40

30

20

o <NH)2Mxyl + (NMe)2Mxyl o (NH)2Pxyl

eo

e e

+

o«

to

.02 .04 .06 .06 .10 .12 .14

Figure 18. Onsager Plots for Monomeric Dry Cave Nickel(II) Complexes.

TABLE 10

ONSAGER PLOT PARAMETERS FOR MONOMERIC DRY CAVE NICKEL(II) COMPLEXES

Compound Aq , ohm ^mole . -1 . -1.5 3.5 Slope, ohm mole cm

[Ni{ (m-Xylyl(NHEthi) 2)Me2[16] tetraeneN^] (PF6) 2 153 449 + 53[Nl{(mrXylyl(MeNEthi)2)Me2[16]tetraeneN^}](PFg)2 171 497 + 45

[Ni{ (£-Xylyl (NHEthl) 2)Me2 [16 ] tetraeneN^,} ] (PFg) 2 171 466 + 35

115-1 -1 2150-170 ohm mole cm which are in the expected range for 2:1 elec-

93trolytes.

*H NMR spectra for all of the complexes were measured and the

data are summarized in table 11. All of the complexes have the

expected mirror symmetry as has already been described for the unbridged

complexes. The spectra of the complexes having the (NHj^Pxyl and

(NMe^Mxyl bridges are shown in figures 19a and b respectively. The

general assignments have been described in detail by Stevens.^ The

complexes shown here differ primarily in terms of the aromatic reso

nances observed due to the xylyl bridging group and also in terms of

the benzyl or n-butyl substituent on the bridge nitrogen.13C NMR spectra were also measured for the complexes and

assigned through the use of the off-resonance technique. Again the

detailed assignments have been described by Stevens^ with the major

differences being due to the prest. of aromatic resonances due to the13xylyl bridge. The C resonances are listed in table 12 and the spectra

of the (NH^Mxyl, (NH^Pxyl, and (NMe^Mxyl bridged complexes are shown

in figures 20 and 21a and b, respectively.

It is interesting to note that in all cases except the

(NlO^Mxyl derivative the resonances due to the methyl groups directly

bonded to the macrocycle are less than one ppm apart, whereas in the

spectrum of the (NH^Mxyl derivative, the only definitely established

"lid-on" isomer, they are about 6 ppm apart. It seems appropriate to

conclude that methyl resonances which are close together (<1 ppm separa

tion) indicate a "lid-off" structure and methyl resonances which are

widely separated (-5-6 ppm) indicate a lid-on structure. (See sections

TABLE 11PROTON NMR DATA3 FOR HONOMERIC DRY CAVE NICKEL(II) COMPLEXES


[Hl{(a-Xylyl(NHEthl)2)He2[16]tetraeneNAJ](PFfi)2 2.20, 2.38 3.3,b 4.74c 7.32, 7.43, 7,69 6.5b[Nl{(mrXylyl(MeNEthi)2)He2[16]tetraeneN^}](PFfi)2 2.08, 2.45, 3.63 3.15,C 4.68 7.00, 7.08, 7.22 ----[Ni{(w-Xylyl(n-BuNEthi)2)He2[16]tetraeneNAJ](PFfi) 0.97, 2.05, 2.47 3.12,b 3.83,C 4.68 7.10,C 7.17 ----

[Nit(m-Xylyl(BZNEthl)2)Me2[16]tetraeneNAJ](PFfi)2 1.95, 2.67 3.07,b 4.58,d 5.07d 6.90, 7.00, 7.30, 7.47

----

[Nit(B-Xylyl(NHEthl)2)Me2[16]tetracne^}](PFg)2 2.02, 2.43 3.23,c 4.51 7.03, 7.18, 7.43 notseen

[Nit(1,7-Hept(HeNEthl)2)He2[16]tetraeneN4J](PFfi)2 1.88, 2.37, 3.27 1.32,b 3.03, 3.5C 7.38 ----

aAs salts, run on concentrated solution in CD^CN, chemical shifts given in ppm from TMS.bBroad.ûltiplet.dQuartet. eCD^N02 solution.

rAw8 " T

6x* iop p m

Figure 19. Proton NMR Spectra of a) [Nl{(£-Xylyl(NHEthi)2)Me2[16]tetraeneN4}](PF6>2 and b) [Nl{(mrXylyl(MeNEthi)2)Me2[16]tetraeneN4}](PF6)2

117

118

TABLE 12CARBON-13 NMR DATA FOR MONOMERIC DRY CAVE NICKEL(II) COMPLEXES

0Compound Chemical Shifts*5

[Ni{(m-Xylyl(NHEthi)2)Me2[16]tetraeneNA>] 2+ c 170.7, 164.3, 161.8, 140.5, 130.0, 127.3, 124.0, 114.5, 55.2, 50.7,47.7, 30.5, 30.1, 20.7, 14.4

[Nl{(mrXylyl(MeNEthi)2)Me2[16]tetraeneN^}]2+ 174.5, 166.9, 159.2, 137.7,130.6, 126.9, 125.3, 113.6, 62.2, 56.4, 52.0, 45.8, 30.5, 29.8, 20.9, 20.6

[Ni{(o~Xylyl(n-BuNEthi)2)Me2[16]tetraeneN^}]2+ 174.0, 166.3, 158.9, 137.4,130.4, 126.8, 125.2, 113.9, 59.8, 57.2, 56.6, 52.0,30.4, 29.9, 20.9, 20.6, 13.9

[Ni{(m-Xylyl(BZNEthi)2)Me,[16]tetr*H }]2+ 173.5, 167.1, 159.7, 137.1,135.2, 130.4, 130.3, 130.2, 129.9, 126.7, 125.2, 114.4,60.2, 58.9, 56.7, 52.1,30.3, 29.7, 21.1, 20.8

[Nl{<£-Xylyl(NUEthI)2)Me2tl6]tetTaeneNA}]2+ 168.2, 165.6, 161.4, 135.7, 128.9, 127.4, 111.9, 56.4, 51.7, 50.6, 30.1, 29.8, 23.6, 22.2

[Nl{(1,7-Hept(MeNEthi)2)He2[16]tetraeneN^}]2+ 175.7, 166.9, 159.7, 110.8, 57.1, 56.6, 51.8, 39.6, 30.9, 29.9, 27.1, 26.4, 25.4, 20.5, 20.4

aAs PFg- aalts.1>CD3CN solution, ppm relative to TMS. cCD2N02 solution.

1

L+160 ~\20

----T"80ppm

i40

T “0

13Figure 20. C NMR Spectrum of [Ni{(m-Xylyl(NHEthl)2)Me2 [16]tetraeneN/.}] (PF6)2

119

u J L . vvJ^rA^* i*J L _

160 120 I80ppm

” 14 0

13Figure 21. C NMR Spectra of a) [Ni{(jJ-Xylyl(NHEthi)2)Me2 [16]tetraeneN^}] (PFg)^ and b) tNi{(rarXylyl(MeNEthi)2)Me2 [16] tetraeneNA)](PF6)2

120

121

on unbridged nickel(II) complexes, page 93 and dimeric nickel(II) com

plexes, page 131.

The electrochemical studies were also performed on these com

plexes and reveal some very interesting features. As can be seen from

the list of half-wave potentials for the first oxidation in table 13,

the same trend is observed for the bridged complexes as for the

unbridged. For a given bridging group, i.e., (NlO^Mxyl, the oxidation

potential decreases as the bridge nitrogen substituent is changed from

hydrogen to benzyl to n-butyl to methyl, which is exactly the same

order as that observed for the unbridged complexes. That the same

trend is observed in both bridged and unbridged species confirms the

theory that it is the substituent electronic effects and not

simply the bridge that control oxidation potential of the metal center.

As observed by Schammel,^ the potential for the first oxida

tion of (NH^Pxyl derivative is e: onally positive. The reasons

for this peculiarity are as yet unknown.

Oxidations observed at more positive potentials are pre

sumably associated with the ligand and are irreversible. The only

reductions observed are at very negative potentials (<-1.8 V) and are

irreversible.

Characterization of Dimeric Nickel(II)Dry Cave Complexes

78 92The dimeric species derived from 1,5-diaminopentane, *92 74 781,4-diaminobutane and 9,9-bis(3-aminopropyl)fluorene * have been

characterized by others and will not be discussed in detail here.

TABLE 13

ELECTROCHEMICAL DATA FOR MONOMERIC DRY CAVE NICKEL(II) COMPLEXES

Compound El/2’ V lE3/4 “ El/41’ mV Ep, V

[Nl{(m-Xylyl(NHEthl)2)Me2 [16]tetraeneN4 }3(PF6)2 +0.925 67 +0.970+1.140 130 +1.300

[Ni{(mrXylyl(MeNEthi)2)Me2 [16]tetraeneN^}](PF^)2 +0.780 70 +0.830+1.025

[Ni{ (m-Xylyl (n-BuNEthi) 2^Me2 [16] tetra^ne^ } ] (PF^) „ +0.790 2 60 +0.820+0.980

[Ni{(m-Xylyl(BZNEthi)2)Me2[16]tetraeneN^}](PFg)2 +0.820 70 +0.850+1.020

[Nit (£-Xylyl(NHEthi)2)Me2 [16]tetraeneN^}](PFg)2 +1.05 100 +1.055[Nit(1,7-Hept(MeNEthi)2)Me2[16]tetraeneN^}](PF6)£ +0.740 70 +0.775

+1.110 70 +1.135

122

123


complexes (see experimental section). Infrared spectra of the dimeric

complexes have the same general features as those of the related mono

mers with the only major differences appearing in the fingerprint

region. In general, the dimeric complexes show fewer sharp absorptions

than the monomers. The IR spectrum of the dimer derived from

m-xylylenediaraine Is shown in figure 22 and selected data are listed in

table 14.

500100015003000 20004000 cm"

Figure 22. Infrared Snectrum of[Nl{(m-Xylyl(NHEthi)2)Me2 [16]tetraeneN^}]2 (PFfi)

Molar conductances were measured in acetonitrile and appear to

be in the range of 2:1 electrolytes (table 14), as expected since con

ductance measurements at a single concentration do not distinguish

between monomeric and dimeric species. An Onsager plot for the

TABLE 14SELECTED INFRARED FREQUENCIES AND MOLAR CONDUCTANCES3

FOR DIMERIC DRY CAVE NICKEL(II) COMPLEXES

Compound ^ C=C,C=N (cm b) A, ohm ^mole ^cn?

[Ni{ (m-Xylyl (NHEthi) 2)Me21 tetraeneN^ ) ] 2 4 3415 1618, 1580 247

[Ni{(p-Xylyl(NHEthi)2)Me2 [16]tetraeneN^}]2(PF6)4 3400 1610, 1570 233b

[Ni{ (p-Xy lyl (MeNEthi) 2)Me2 [16] tetraeneN^ ] 2 (PFg) 4 ----- 1608, 1542 241

I Nl{ (DURYL) (NHEthi) 2)Me2 [16 ] tetraeneN^ } ] 2 (PFg) 4 3390 1660, 1620, 1570 223

[Ni{(DURYL(MeNEthi)2Me2[16]tetraeneN^ ]2(PFfi)4 ----- 1615, 1540 230

a -3Acetonitrile solution of ~1 x 10 molar at ambient temperature.

Reference 78.

125

(NH^Mxy1 bridged dimer has a slope of 602 + 27 ohm-^ mole*"^'^ cm^**’

which is about 100 units greater than that of the corresponding monomer,

but not nearly as large as expected for a 4:1 electrolyte. These con

ductance results are very similar to the observations of Jackels and

Grzybowski on the related pentamethylene and fluorene systems. It is

thus apparent that Onsager plots are not a useful tool for distin

guishing between monomeric and dimeric species.

Proton NMR data for the complexes are listed in table 15 and

are similar to those of the corresponding monomeric species. The

NMR spectra of the (NH^Mxyl and (NMe^Duryl bridged dimeric complexes

are representative and appear in figures 23a and b.13C NMR spectra were also measured for the dimers and, in all

cases except the dimeric (NH^Mxyl derivatives, the same number of

resonances were observed as would be expected for the corresponding

monomer, indicating the presence of two symmetry elements within the

molecule. As will be described in the discussion of the structural

analysis of the dimeric (NH^Mxyl bridged species, these elements are

probably an Inversion center relating the two metal coordinating macro

cycles and a mirror plane through the macrocycles containing the metal

centers as appears in the monomeric complexes. The resonances are

listed in table 16 and the spectra of the dimeric (NH^Mxyl and

(NMe^Duryl bridged complexes appear in figures 24 and 25. As men

tioned above, the spectrum of the dimeric (NH^Mxyl complex is not as

simple as the others. This is likely due to the presence of several

Isomeric species in solution which are not interconverting on the NMR

time scale.

TABLE 15

PROTON NMR DATA3 FOR DIMERIC DRY CAVE NICKEL(II) COMPLEXES


[Ni{(m-Xylyl(NHEthi)2)Me2 I161tetraeneN^}]2(PFg)^ 2.08,2.41

2.252.52 3.0C,4.6C 7.42,7.55 notseen

[Ni{(p-Xylyl(NHEthi)2)Me2I^JtetraeneN^ }]2 (PFfi) 2.03, 2.37 2.80, 3.5,b 7.43, 7.55 not4.60 seen

[Ni{(p-Xylyl(MeNEthi)2)Me2 [16]tetraeneN^}]2(PFg) 2.07, 2.50, 3.05,° 4.72 6.98, 7.15, -----3.52 7.57

[Ni{ (DURYL) (NHEthi) 2)Me2 [16] tetraeneN^ } ] 2 (PFg) 4 1.73, 2.23, 3.05,° 4.88 7.58 not2.37, 2.50 seen

[Ni{ (DURYL(MeNEthi) 2Me2[16] tetraeneN^, >] 2(PF&) 1.80, 2.10, 3.0,C 4.95 7.53 -----

2.47, 2.57

£Run on concentrated solution in CD^CN, chemical shifts given in ppm from TMS.

bBroad.

CMultiplet.

° ° ppm * *Figure 23. Proton NMR Spectra of a) [Ni{(m-Xylyl(NHEthi)_)Me0[16]tetraeneN.}]„(PF-)

and b) [Ni{(DURYL(MeNEthi)2Me2[16]tetraeneN4}]2(PF6)4

127

128

TABLE 16CARBON-13 NHR DATA FOR DIMERIC DRY CAVE NICKEL(II) COMPLEXES

A bCompound Chemical Shifts

[Hl{(m-Xylyl(HHEthl)2)He2tl6]tetraeneNi,}I24+ C 170.1, 169.3, 169.2, 168.9, 168.7,168.2, 160.7, 139.5, 138.5, 137.9,131.A, 131.1, 130.8, 130.3, 129.2,128.2, 128.0, 127.7, 113.2, 56.9,51.9, 50.0, 49.0, 31.0, 30.4, 20.8, 15.4

lNi{(p-Xylyl(NH£thl)2)Me2[16]tetraeneNA}]^+c,d168.6, 168.2, 160.7, 136.9, 131.5,130.3, 112.9, 56.8, 51.7, 50.3,30.9, 30.4, 20.7, 15.4

[Nl{(p-Xylyl(MeNEthi)2)He2[16]tetraeneN4}]24+ 167.9, 167.4, 161.0, 135.0, 127.8,127.5, 113.8, 62.0, 56.0, 51.7,45.5, 30.2, 29.3, 23.0, 19.3

[Nl{(DURYD(NHEthC2)Me2[16]tetraeneN^.}]24+ 168.1, 167.9, 160.6, 136.4, 135.0,133.0, 112.3, 66.3, 56.5, 51.5,30.3, 29.9, 19.9, 17.1, 16.8, 15.7

[Nl{(DURYL(MeNEthl)2Me2[16]tetraeneN }]24+ ;: .4, 168.5, 159.8, 136.6, 136.0,132.3, 112.6, 66.3, 56.3, 52.7,50.6, 43.2, 31.0, 30.3, 20.8,19.3, 18.0, 17.2, 15.6

aAs FFg- salts.^CD^CN solution, ppm relative to TM5.CCD^N02 solution.^■Reference 78.

viw V f W f f~T~0160 120 -----1

8 0 ppm

4 0

13Figure 2 k . C NMR Spectrum of [Ni{(m-Xylyl(NHEthi)2)Me2 [16]tetraeneNA}]2(PFg)^

129

40 0(60 120 80ppm

13Figure 25. C NMR Spectrum of [Nl{(DURYL(MeNEthi),Me„[16]tetraeneN.}]_(PF£).c L 4 Z 6 4

130

13113One outstanding feature of the C NMR spectra of these dimeric

complexes is the fact that the macrocycle methyl resonances are sepa

rated by approximately 5 ppm. Based on previous arguments, this obser

vation is interpreted to mean that the dimers are all of the "lid-on*1

type with respect to the orientation of groups around the bridge nitro

gens.

Electrochemical studies have shown that the redox behavior of

the dimeric complexes is very similar to that of the monomers with lit

tle dependence on the nature of the linking group (table 17), Changing

the nitrogen substituent from hydrogen to methyl in the case of the

duryl bridging group causes a shift of 110 mV more negative, as was

observed in the monomeric examples. It is interesting to note that the

dimeric (NH^Pxyl species has a half-wave potential of +0.880 V, in the

normal range for this type of macrocyclic nickel(II) complex. This is

in marked contrast to the exceptional behavior of the corresponding

monomer which has the anamolously high potential of >1.04 V.

As mentioned previously, the complex [Ni[(m-Xylyl(NHEthi)Me2~

[16]tetraeneN^] ( P F g ^ ^ C H ^ C O C H ^ was subjected to a complete single

crystal X-ray diffraction structural analysis. Specific details of the

solution and refinement procedures are found in the experimental sec

tion of this work. ORTEP drawings of the molecule appear in figure 26.

The numbering scheme for the molecule is shown in figure 27 and the

relevant bond distances and angles are listed in tables 18 and 19.

Despite numerous difficulties encountered in the structural

analysis and the resultant high discrepancy indices, a number of

TABLE 17

ELECTROCHEMICAL DATA FOR DIMERIC DRY CAVE NICKEL(II) COMPLEXES

Compound El/2- V lE3 / 4 " El/J’ mV Ep, V

[Ni{(m-Xyly1(NHEthi)2)Me2 I^tetraeneN^} ] 2(PFg)^ +0.880 85 +0.925+1.230

[Ni{(p-Xylyl(NHEthi)2)Me2[16]tetraeneN^}]2(PFg) +0.880 88 +0.930+1.230

[Ni{(p-Xyly1(MeNEthi)2)Me2[16]tetraeneN^}]2(PF6)4 - - - - — +0.875+1.175

[Ni{ (DURYL) (NHEthJ) 2)Me2 [16] tetraene^} ] 2 (PFg) +0.890 70 +0.960+1.190 100 +1.350

[Ni{(DURYL(MeNEthi)2Me2[16]tetraeneN^}]2(PFg)^ +0.780 85 +0.830+1.06 55 +1.130

Figure 26. ORTEP Drawings of the Dimeric Nickel(II) Complex

133

134

TABLE 18BOND DISTANCES FOR DIMERIC NICKEL(II) COMPLEX

N i-N l 1.877N1-N2 1.902N1-N3 1.868N1-N4 1.851N l - C l 1.33N1-C12 1.46N2-C3 1 .26N2-C4 1 .4 7N3-C6 1.49N3-C7 1.30N4-C9 1.29N 4 -C 10 1.53C1-C2 1 .47C1-C13 1.51C2-C3 1 ,44C2-C15 1.38C4-C5 1.52

C5-C6 1 .5 7C7-C8 1 .43C8-C9 1 .47C8-C16 1.42C9-C14 1 .55C10-C11 1.48C11-C12 1.51C15-C17 1.52C15-N5 1.34C16-C18 1.52C16-N6 1.35C19-N5 1.44C19-C20 1.54C20-C21 1 .45C20-C25 1.37C21-C22 1.42C22-C23 1.34

C23-C24 1.44C24-C25 1.38C24-G26 1.52C26-N6 1.52P l - F l 1.578P 1 -F 2 1.598P 1 -F 3 1.564P 1 -F 4 1.594P 1 -F 5 1.590P 1 -F 6 1.558P 2 -F 7 1.574P 2-F 8 1.544P 2-F 9 1.541P 2-F 10 1.567P2-F 11 1.491P2-F 12 1.524

135

TABLE 19BOND ANGLES FOR DIMERIC NICKEL(II) COMPLEX

N2-N1-N1 8 9 .2N4-N1-N1 8 9 .9N3-N1-N2 9 1 .3N4-N1-N3 8 9 .3C l - N i - N i 120.3C12-N1-N1 119.4C12-N1-C1 120.3C3-N2-N1 117.1C4-N2-N1 122.2C4-N2-C3 120.3C6-N3-N1 122.1C7-N3-N1 119.6C7-N3-C6 117.2C9-N4-N1 120.2C 10-N4-NI 120.6C10-N4-C9 119.1C 2-C1-N1 121 .7C13-C1-N1 119.9C13-C1-C2 117. 1C3-C2-C1 112.9C15-C2-C1 122.7C15-C2-C3 124.4C2-C3-N2 127.9C5-C4-N2 111.9C6-C5-C4 111.3C5-C6-N3 111.0C 8-C 7-N 3 120.8C 9-C 8-C 7 117. 1

C 16-C 8-C 7 118.6C16-C8-C9 124.0C 8-C 9-N 4 120.6C14-C9-N4 118.8C 14-C9-C8 120.0C11-C10-N4 108.7C12-C11-C10 111.4C11-C12-N1 111.5C17-C15-N5 118.7C17-C15-C2 120.7N5-C15-C2 120.4C18-C16-N6 119.7C18-C16-C8 121.9N6-C16-C8 118.4C19-N5-C15 126.2C26-N6-C16 125.8C 20-C19-N5 114.8C21-C20-C19 118.2C21-C20-C25 121.3C25-C20-C19 120.4C22-C21-C20 117.0C23-C22-C21 119.9C24-C23-C22 123.2C25-C24-C23 117.5C26-C24-C23 120.2C26-C24-C25 121.9C24-C25-C20 121.1C24-C26-N6 109.3

136

C 23' C22

Figure 27. Numbering Scheme for Dimeric Nickel(II) Complex

significant observations can be made regarding the complex. The most

striking feature is the dimeric nature of the molecule. Two macro-

cyclic species are linked by meta-xylene bridges with the two halves of

the molecule related by an inversion center through the middle of the

resultant cavity. The xylyl groups are constrained by symmetry to be

parallel as are the macrocyclic planes. The nickel(II) ions are 0

0.07 A out of the N^ planes (away from the molecular center) with anO

average nickel-nitrogen distance of 1.86 A. These values compare

favorably with those of the monomeric (NH^Pxyl bridged nickel complex,o o ^

0.04 A and 1.87 A, respectively. The six-raembered rings containing

the macrocyclic side chains below the xylene groups are in boat confor

mations whereas the other six-membered rings are in the chair form.

137

This is consistent with expected steric interactions between the

bridging group and the macrocycle. The dihedral angle between the

macrocycle plane and the plane of the xylyl ring is 54°.O

The nickel(II) ions are 13.6 A apart from center to center, ofO O

which 12.8 A is vertical and 4.6 A is horizontal (parallel to the N,

plane) displacement. The point of nearest approach across the cavity isO

between C19 and C19‘", the benzylic carbons, at a distance of 3.87 A.

The distances between non-bridgehead carbons (C15-C16) and bridge nitro-O O

gens (N5-N6) are 6.18 A and 5.74 A, respectively. These distances will

be compared with those of other structures in the Structure Summaries

section of this work.

One final feature of this structure which should be mentioned is

the "lid-on" orientation of the meta-xylene groups at the bridge nitro

gen atoms. This is the predicted orientation for dimers based on

steric and NMR considerations as described earlier.

The concrete evidence for dimeric complexes supplied by theIsolution of this crystal structure resulted in the reexamination and •

interpretation of a great deal of chemical information. As will be

described later, many of the previously unexplained properties of some

iron(II) complexes were readily interpreted once the dimeric nature of

the ligand structure was confirmed.

Characterization of the Complex Derived from 9.10-Bis(chloromethyl)anthracene

The reaction between [Ni{ (MeNHEthi)2*162 [16] tetraeneN^KPFg^ and

9,10-bis(chloromethyl)anthracene proceeded smoothly and in good yield to

produce a complex having stoichiometry of one macrocycle per

138

anthracene group. The complex crystallized as large red crystals from

a solvent mixture of acetonitrile and ethanol and analyzed well as a

diacetonitrile solvate or without solvent (after mulling and vacuum

drying). The infrared spectrum appears in figure 28 and shows no

unusual features. The absorption near 2250 cm ^ is due to the acetonl-

trile of crystallization and the absorptions between 1500 and 1700 cm

are assigned to C - C and C = N stretches and aromatic overtones.

50010001500200030004000 cm

Figure 28. Infrared Spectrum of the Anthracene Derivative

The NMR spectrum is shown in figure 29 and is quite unusual

in many respects. There are a number of aromatic resonances due to the

anthracene group. It is apparent that there is a separate resonance for

every methyl group indicating a lack of mirror symmetry in solution.

The resonances between 5.0 and 6.5 ppm are unlike any observed in other

if'

ppm

Figure 29. Proton NMR Spectrum of the Anthracene Derivative

139

140

monomeric or dimeric dry cave complexes. It is possible that they are

due to the benzylic hydrogens of the anthracene, but this assignment

has not been confirmed.13The broadband proton decoupled C NMR spectrum appears in

figure 30. There is a resonance corresponding to every carbon atom in

the molecule (assuming a small amount of overlapping in the aromatic

region and that two resonances coincide at 173.3 ppm). The resonances

are listed in table 20 along with the multiplicity of the resonance as

obtained from the off-resonance spectrum. These results again clearly

demonstrate the lack of mirror symmetry in solution. The resonances at

113.6 and 69.1 ppm are both singlets in the off-resonance spectrum and

must therefore be assigned to the y carbon of the macrocycle. It is not

at all clear what factors are responsible for causing these resonances

to be separated by 45 ppm.

Molar conductances were measured for this complex in acetoni--3trile as a function of concentration. The conductance of a 1 x 10

-1 -1 2molar solution was 273 ohm mole cm , consistent with the formula

tion as a 2:1 electrolyte. The slope obtained from the Onsager plot

was 440 + 36 ohm ^ mole cm^'^ which is closer to the value observed

for monomers than for dimers.

The half-wave potential for the first oxidation of this species

occurs at +0.460 V versus Ag/Ag+ and is irreversible. This is likely

due to a ligand oxidation and serves as further evidence for the

unusual nature of this complex.

In order to clarify the structure of this species, an attempt

was made to perform a single crystal X-ray structural analysis.

1---------1---------1--------- 1---------1--------- 1--------- r200 T

8 0T0160 40

ppm

13Figure 30. C NMR Spectrum of the Anthracene Derivative in

TABLE 2013C NMR DATA FOR THE NICKEL(II) COMPLEX DERIVED

FROM 9,10-BIS(CHLOROMETHYL)ANTHRACENE

Shift3 (mult^) Shift (mult) Shift (mult)

179.2(s) 127.9 56.4(0

173.4(s) 127.7 55.1(t)

173.4(d) 126.9 50.4(t)

166.9(s) 126.8 45.3(q)

165.5(b ) 125.8 39.7(q)

158.1(d) 125.4 30.7(t)

132.6 125.2 29.3(0

132.2 123.7 25.6(0

131.4 113.6(s) 24.0 (q)

131.0 69.8(s) 22.5(q)

130.8 58.9(t) 20.8 (q)

128.6 57.8(t) 18.9(q)

^ D ^ C N solution, ppm relative to TMS.

^Multiplicity; s q = quartet.

« singlet, d = doublet, t = triplet,

143

Four different crystals were mounted and examined, but each showed dif

fraction patterns characteristic of twinned crystals. Attempts to

obtain suitable crystals from other solvent systems and using different

anions met with little success. As a result the actual structure of the

species is as yet unknown.

One can speculate as to the origin of the peculiar properties

observed for this species. An obvious possibility is the cocrystalliza

tion of two different species, each of which has the symmetry of normal

dry cave complexes. For example, these might be a monomer and dimer.13This could account for the NMR spectra. The C NMR data indicate the

presence of a very unusual complex which gives rise to the singlet at

69.1 ppm. The frequency indicates the carbon is quarternary, having no

double bonds associated with it. Structures consistent with such an

atom are structures XXIII and XXIV in which the bridge nitrogen is now

an imine and either the bridge or methyl group is bonded to the y car

bon. There is no precedent for such structures, however, and thus they

must be viewed as purely speculative.

CH

B r i d g e

N \

ridge

CH

X X I I I

144

Removal of Ligands from Nlckel(II)

The central metal ion in all of the complexes discussed up to

this point has been nickel(II). The complexes were synthesized with

nickel(II) as the central metal ion for a number of reasons: 1) The

synthesis of the parent macrocycle requires the use of nickel(II) as a

template for macrocyclic ring closure. 2) The nickel(II) complexes are,

for the most part, air stable and relatively easy to handle. 3) The

nickel(II) complexes are readily characterized, particularly through

the use of NMR techniques. 4) The ligands can be removed intact from

nickel(II) and coordinated to the biologically important metal ion,

iron(II). The ligand salts of all of the complexes were synthesized

according to the basic method of Schammel^ (equation 26). The tetra-

chlorozincate salts of the ligands which result from this treatment

were in many cases impure and very difficult to characterize.

I IZ n Zt HCI

CH^CN R'(ZnCL)

(26)

145

Other workers in this research group**** have not reported the charac

terization of ligand salts in any detail; however, the characterization

of several salts is included in this work. Good analytical data were

obtained for the monomeric (NH)^Mxyl and (NMe^Mxyl bridged ligands as

tetrachlorozincate salts. There are several disadvantages inherent in

the use of tetrachlorozincate as a counterion. The compounds in

general are water sensitive, difficult to recrystallize and pose synthe

tic difficulties in further reactions as the zinc ion can compete with

iron(II) for coordination by the macrocycle. Because of these prob

lems, several of the ligand salts were metathesized using ammonium

hexafluorophosphate. The resulting compounds were generally white to

off-white granular materials yielding good analytical and spectroscopic

data. In particular, the ligands derived from the monomeric (NH^Mxyl,

(NH^Pxyl and (NMe^CCHj)^ bridges were studied in detail.

The ligands generally protonated readily in acetonitrile with

hydrogen chloride gas, the solution showing the deep blue color of the

concommittantly formed tetrachloronickelate anion within 30 minutes.

The (NH^Pxyl monomer had to be left in a solution saturated with hydro

gen chloride for two days before ligand removal could be successfully

accomplished.^ Another Interesting exception involves the monomeric

and dimeric (NH^Mxyl species. The monomeric ligand salt remained in

solution upon addition of the tetrachlorozincate solution whereas the

dimeric product precipitated immediately. (This is the ligand salt 70used by Schammel and his report therefore must refer to dimeric

iron(II) complexes.) The monomeric ligand salt was isolated by rotary

evaporation of the solvent and HC1 gas. All other complexes studied by

146this author precipitated immediately upon addition of the tetrachloro

zincate solution.

Attempts were made to prepare ligand salts from the (NnBu^Mxyl

bridged complex and the complex prepared from the anthracene group. In

each case there was no color change to Indicate ligand removal and the

only products isolated were tetrachlorozincate salts of the nickel(II)

complexes.

As mentioned above, several ligands were studied in some detail.

Acceptable elemental analyses were obtained for the ligands derived

from the monomeric (NH^Mxyl and (NMe^Mxyl bridged complexes as tetra

chlorozincate salts and for the ligands derived from the monomeric

(NH^Mxyl and (NH^Pxyl bridged complexes as hexafluorophosphate salts.

Infrared spectra show the presence of broad N-H stretches which are

undoubtedly due to hydrogen bonding either to solvent molecules or to

other ligands. The rest of the IR spectrum shows few features except

in the vicinity of 1600 cm ^ and the typical bands due to hexafluoro-

phosphate. The infrared spectrum of the (NH^Mxyl derivative is shown

in figure 31 and some selected frequencies are listed in table 21.

Proton NMR spectra were measured for the hexafluorophosphate

ligand salts of the monomeric (NH^Mxyl, (NH^Pxyl and (NMe^CC^)^

bridged species. The NMR spectra for the (NHjgMxyl and (NH^Pxyl

derivatives are shown in figures 32a and b, and the data are summarized

in table 22. The most Interesting feature of these spectra is the N-H

resonance in the vicinity of 10-11 ppm. Such resonances have also been74 75observed in other ligand salts of this class of macrocycle. * The

remainder of the spectra are shifted relative to the respective

TABLE 21SELECTED INFRARED FREQUENCIES3 FOR DRY CAVE LIGAND SALTS

Compound VN-H(co’1) VC=C,C=N (cm-1)

[ (m-Xylyl (NHEthi) 2 ) ^ 2 [16 ] tetraeneN^ ] (ZnCl^,) 2 3190,b 3450b 1645, 1585

[ (m-Xylyl (NHEthi) 2 ) ^ 2 [16]tetraeneN^] (PFg) g 3358 1645, 1580

[(m-Xylyl(MeNEthi)2)Me2[16]tetraeneN^](ZnCl^)2 3160,b 3500b 1645, 1600

[(p-Xylyl(NHEthi)2)Me2[16]tetraeneNA ](PFg) 3370b 1638, 1585, 1543

[ (1,6-Hex (MeNEthi) 2 ) ^ 2 [16]tetraeneN^ ] (?Fg) g 3330b 1645, 1600

clSpectra obtained fromnujolmulls on KBr plates.

bBroad.

148

4000 3000 2000 1500 1000 500cm

Figure 31. Infrared Spectrum of[(m-Xylyl(NHEthi)2)Me2 [16]tetraeneN4 ](PFfi)

nickel(II) complexes but their assignments are in support of the

structures.

Carbon-13 NMR spectra were also recorded for several of the

ligands. The spectra of the (NH)2Mxyl and (NH)2Pxyl species are shown

in figures 33a and b, and the data are summarized in table 23. The

general patterns of the spectra are very similar to those of the

corresponding nickel(II) complexes with each exhibiting mirror sym

metry. The range of resonances observed for a given carbon atom with

different bridges is much larger than that observed for the nickel(II)

complexes, indicating that in the absence of structural constraints

imposed by the metal ion the ligand distortions become more pronounced

with changes in the bridging group.

p p m

Figure 32. 1H NMR Spectra of a) [(m-Xylyl(NHEthi)9)Me [16]tetraeneN.](PFfi)and b) [(p-Xylyl(NHEthi)2)Me2[16]tetraeneN4 ](PF6)3

149

TABLE 22

PROTON NMR DATA* FOR DRY CAVE LIGAND SALTS


[(m-Xylyl(NHEthi)2)Me2[16]tetraeneN^](PFg)3 2.05, 2.37 3.75,C 4.15,C 6.37, 7.30C 7.82b4.75

t (p-Xylyl(NHEthi) 2)Me2[16] tetraeneN^, ] ( ^ > 3 2.28, 2.60 3.6,b 4.40, 6.95, 7.20, 9.05,b 10.554.70c 7.78

I(1,6-Hex(MeNEthi)2>Me2[16]tetraeneN^](PF^)3 2.20, 2.58, 1.30,b 3.75b 9.70, 9.95 8.3,b 11.lb3.62

aChemical shifts given in ppm from TMS, run on concentrated solutions in CD^CN.

bBroad.

CMultiplet. 150

w> J * w\m

4 0160 120 80ppm13,Figure 33. ^ C NMR Spectra of a) [(m-Xylyl(NHEthi)2)Me2 [16)tetraeneN^](PF&)3and b) E(p-Xylyl(NHEthi)2)Me2 [16]tetraeneN^](PF^)3 151

TABLE 2313C NMR DATA FOR MONOMERIC DRY CAVE LIGAND SALTS

Compound Chemical Shifts3

t(m-Xylyl(NHEthi)2)Me2 [16]tetraeneN4 ](PFg) 176.2, 172.4, 156.9, 138.3, 130.6, 127.2, 120.5, 106.2,54.1, 52.4, 48.2, 27.9,24.2, 20.7, 15.1

[(p-Xyly1(NHEthi)2)Me2[16]tetraeneN4](PF6) 3 174.9, 160.2, 134.6, 127.8, 126.2, 104.0, 52.3, 50.9, 49.5, 27.6, 24.3, 23.4,22.3

((1,6-Hex(MeNEthi)2)Me21 1tetraeneN4](PFg)3 183.0, 174.2, 156.3, 100.0,60.0, 52.1, 47.8, 43.4, 28.4, 26.5, 26.1, 24.8, 23.8, 20.3

aCD3CN solution, ppm relative to TMS.

153

Several interesting conclusions can be drawn from these results.

The fact that these ligand salts precipitate as +4 cations from ace-

tonitrile solution with tetrachlorozincate anion and as +3 cations from

aqueous solution with hexafluorophosphate anion indicates that the

charge on the resultant cationic species is strongly dependent on

solubility. The mirror symmetry exhibited by the ligand salts shows

that the protons are in rapid exchange with the ligand resulting in a

sharp averaged spectrum.

Another interesting conclusion arises from the fact that the

monomeric (NH^Mxyl ligand spectrum shows simple mirror symmetry and

yet the spectrum of the nickel(II) compound regenerated from it is

rather complex (figure 20). The cause of this must be Isomerization in

the nickel(II) complex having a rate of interchange which is slow on

the NMR time scale. Such an isomerization could be between the

"lid-on1* and "lid-off" conformers, but this has not been verified. It

Is also possible that the xylyl ring can maintain two distinct orienta

tions, one centered over the metal, the other centered over the side

chains of the macrocycle.

It is reassuring to note that in the case of the (NH^Mxyl

bridge, the monomeric and dimeric species retain their integrity and do

not interconvert during ligand removal. This has been confirmed by

experiments in which the ligand was removed from and put back onto

nickel(II) with no change in properties or spectra.

Iron(II) Complexes

The major goals of this work involve the synthesis and study

of the iron(II) complexes of dry cave ligands in order to model the heme

proteins of natural systems. It was intended that through equilibrium

and structural studies, the nature of the binding of CO to the iron(II)

center would be elucidated and the importance of bending of the Fe-C-015linkage would be assessed. It has been suggested that the bending of

the Fe-C-0 linkage in natural systems plays an important role In con

trolling the stability of CO adducts of Hb and Mb. Therefore a series

of dry cave complexes were prepared and their physical and chemical

properties were examined in detail. Through the use of single crystal

X-ray structural analysis, the changes in the size and shape of the

cavity which accompany changes in the bridging group were clearly

shown, demonstrating the wide variety of structural types which are

available. Equilibrium constants for the reaction between the iron(II)

complexes and CO were determined; these decrease as the size of the

cavity decreases. The crystal structure of a key CO adduct was deter

mined demonstrating the first bent Fe-C-0 linkage observed for a model

system. These structural and equilibrium data were correlated with the

CO stretching frequencies and the resonances of the CO carbon atoms in 13the C NMR spectra. From these experiments, the importance of the

interaction between the bridging group and bound CO is made clear. It

is postulated that similar interactions are critical in the functioning

of natural systems.

The next sections of this work describe the synthesis and

characterization of three classes of monomeric iron(II) complexes:

five-coordinate chloro complexes, four-coordinate species and an

unbridged complex. The structures of two chloro derivatives will be

discussed, followed by a summary of crystallographic data for dry cave

structures. In addition, the Interactions of the iron(II) compounds

with axial ligands will be described.

Monomeric Iron(II) Chloro Complexes of Dry Cave Ligands, [FeLCl'H

Having synthesized and characterized a number of ligand salts

of monomeric dry cave compounds, the synthesis and study of their

iron(II) complexes could now be undertaken. [(m-Xylyl(NHEthi)^Me^-

[16]tetraene](PF£)„ was combined with one equivalent ofO J

bis-(acetonitrile)iron(II) chloride and three equivalents of tri-

ethylamine (to deprotonate the ligand) in acetonitrile under an atmos

phere of dry nitrogen according to equation 27.

R'<PF6>3 + Fa (C H jC N jg C ^

+ 3 E*S N*

cr^

156

Upon addition of the base, the color of the solution immediately changed

to deep red and after a few minutes an orange precipitate formed. This

compound was identified as [Fe{(m-Xylyl(NHEthi^jMe^[16]tetraeneN^}-

C1]C1. Recrystallization from methanol yielded large red-orange

crystals of the complex containing two methanol molecules of crystalli

zation. Addition of ammonium hexafluorophosphate to a methanolic solu

tion of the complex resulted in the formation of red crystals of the

hexafluorophosphate salt of the complex.

In a similar way, [ (m-Xylyl(MeNEthi)2)Me2 [16] tetraene] (ZnCl^,^

was combined with one equivalent of bis-(acetonitrile)iron(II)

chloride and four equivalents of triethylamine in acetonitrile. No

precipitate formed, however, so the acetonitrile was removed and

replaced with methanol. Upon addition of ammonium hexafluorophosphate

the red product [Fe{(m-Xylyl(MeNEthi)2^Me2 [16]tetraeneN^JCl] (PFg)

formed. Large red crystals were obtained by recrystallization from an

acetonitrile/ethanol solvent mixture. The synthesis was carried out in

high yield by using methanol as the solvent throughout. Hydrogen

bonding between the N-H of the macrocycle bridging group and the

anionic chloride appears to significantly reduce the solubility of

(NH^Mxyl derivative. This opportunity is not available for the methyl

substituted complex.


complexes described above, verifying the stoichiometry. The infrared

spectra of [Fe{(m-Xylyl(NHEthi)2>Me2 [16JtetraeneN^}Cl]Cl*2CH30H,

[Fe{ (m-Xylyl(NHEthi) 116] tetraeneN^}ci] (PFg) > and [Fe{ (m-Xylyl-

(MeNEthi)2)He2 [16]tetraeneN^}Cl](PFg) are shown in figures 34a, b, and

157

■500

l1000I

1500I

2000I3000

l4000 cm

Figure 34. Infrared Spectra of a) [FeKm-XylylCNHEthiJîMe^tlS]-tetraeneN4 }Cl]Cl-2CH30H, b) [Fe{(m-Xylyl(NHEthi)2)Me2 [16]- tetraeneN^Jci] (PFg), and c) [Fe{ (m-Xyly1(MeNEthl) 2)Me2 H ® 3 ” tetraeneN^JCl](PF^).

158

c, respectively. Of particular Interest in these spectra Is the N-H

region which has a broad absorption at 3205 cm ^ for the chloro-(NH^-

Mxyl chloride complex due to the extensive hydrogen bonding among the

ligand, solvent, and chloride ions. In the spectrum of the PF,- saltbof the same complex there is a sharp N-H stretch at 3420 cm \ showing

no indication of hydrogen bonding. The spectrum of the (NMe^Mxyl com

plex contains no absorptions at all above 3100 cm The region

between 1550 cm ^ and 1650 cm ^ is typical of bridged complexes with

the (NH^Mxyl chloro chloride complex having bands at 1620 cm ^ and

1575 cm The bands appear at 1615 cm ^ and 1565 cm ^ for the

chloro-(NH) Mxyl PF^ species and at 1620 cm ^ and 1550 cm"" * in the

(NMe^Mxyl species. In each case, the higher energy band is sharper

and less intense than the lower energy band.

Solid state room temperature magnetic moments were determined by

the Faraday method and all are consistent with the presence of high-spin

iron(II). The magnetic moment values for the chloro-(NH)2*kcyl

chloride, chloro (NH) 2Mxyl PFg, and chloro (NMe^Mxyl PFg bridged com

plexes are 5.183, 5.303, and 5.383, respectively. The spin only value

for high-spin iron(II) is 4.903. Molar conductances were measured In_3

acetonitrile for 1 x 10 molar solutions of the chloro- (NH^Mxyl and

chloro-(NMe^Mxyl PFg bridged complexes yielding values of 127.3 and -1 -1 2142.7 ohm mole cm which are in the range found for 1:1 electro

lytes. (The acceptable range for a 1:1 electrolyte in acetonitrile is -1 -1 2120-160 ohm mole cm .) These data indicate that the chloride ion

remains essentially coordinated to the iron(II) center in solution.

159

Electrochemical studies were performed on the complexes in

acetonitrile. The chloro-(NH)2Mxyl PFg derivative has two oxidations

with an^ '^3/4 ~ ^1/41 ~0.345 V, 62 mV and +0.935 V, 65 mV ver

sus 0.1 M Ag/Ag+ . As in the case of the nickel(II) complexes, the

reductions occur at very negative potentials and are irreversible.

Controlled potential electrolysis was performed on the complex by oxi

dizing at a potential of -0.1 V. A value of 1.04 electrons was

obtained for two separate experiments. The E^ 2 ant* 1 3/4 " ^ 1/4 1 ôr

the reduction of the oxidized product are -0.330 V and 63 mV, indi

cating that simple oxidation of the metal occurred during electrolysis,

Similar electrochemical studies were performed for the (NMe)^-

Mxyl bridged complex. As shown in figure 35, there are two oxidations

having anc* ^3/4 ” *1/4 values “0.390 V, 60 mV, and +0.900 V,

-0.8 - 1.0-0 .2 -0 .4 - 0.60.0i.2 0.4 0.20.8 0 61.0V. vs A

Figure 35. Cyclic Voltamagram of[Fe{(m-Xylyl(MeNEthi)2)Me2 [16]tetraeneN^ >C1](PF6)

160

70 mV. Again an irreversible reduction wave is observed at -2.12 V.

Controlled potential electrolysis was performed on three separate

samples by first oxidizing at -0.1 V, then reducing the same solution

at -0.8 V. In each case, both processes appeared to be completely

reversible yielding an average n value of 1.03 electrons. Voltama-

grams of the oxidized species verified that the only process occurring

during electrolysis was simple oxidation of the metal. It is very

significant that changing the bridge substituent from hydrogen to

methyl causes a shift in the oxidation potential of about 45 mV (more

negative), consistent with the observations for the related nickel(II)

complexes.

Through the efforts of several workers, eight different iron(II)

chloro complexes have been synthesized and characterized to varying

degrees. The complexes studied having hydrogens on the bridge nitrogens78 92 88 92are those with meta-xylene, para-xylene, (CH^)^, (CHg)^, *

77 78 88and fluorene * bridging groups. The complexes with methyl

substituents on the bridge nitrogens are those with meta-xylene and 77 bridging groups. The chemistry of the complexes not described

in this work was developed concurrently with that of the meta-xylene

bridged complexes. All of the complexes have properties very similar

to those described above, and several of the complexes were studied in

further detail by this author, as will be described below.

Attempted Synthesis of Four-Coordinate Iron(II) Complexes, [FeL]^+

Because of the lack of an active chloro iron(II) analog in

natural systems, attempts were made to synthesize four- or

161

five-coordinate iron(II) complexes which did not contain chloride.

Direct syntheses were attempted using the hexafluorophosphate salts of

the ligand and iron(II) according to equation 28 for the (NMe)

(NMe)2 (CH2)5 ,

CH CNH3L(PFfi)3 + Fe(CH3CN)6 (PF6)2 + 3Et3N --- ►

[FeL(CH3CN)x ](PF6)2 , x = 0,1 (28)

and (NMe)2 (CH2)g bridged species. Analytical data were unsatisfactory

for either of the above formulations in each case. Infrared spectra

were sharp and typical of monomeric dry cave complexes. NMR spec

tra were similar to those of the paramagnetic chloro complexes as will

be described below. The electrochemical behavior was essentially iden

tical to that of the chloro complexes with values of E^ 2 and

l E ^ - E ^ J of -0.435 V, 70 mV and -0.440 V, 65 mV for the

(NMe)2 (CH2)^ and (NMe)2 (CH2)3 species, respectively. The (NMe)2(CH2)3

species had very ill defined voltamagrams. These complexes are prob

ably a mixture of four- and five-coordinate species, perhaps containing

a significant amount of chloro complex. Although the presence of

chloride has not been confirmed, Dr. N. Herron has suggested that the

ligand salt may have the composition H^LfPF^^Cl and therefore act as94a source of chloride.

An analytically pure sample of the four-coordinate iron(II)

complex derived from the (NMe)2Mxyl bridged ligand was synthesized by

the above method. The only unusual properties for this species are its

dark brown color and rather positive oxidation potential of -0.120 V,

consistent with the absence of chloride in the coordination sphere.

162

Iodide salts of this complex were prepared either by metathesis

of the hexafluorophosphate salt in acetone or by direct synthesis of

the iron(II) complex in acetone. These complexes analyze well and are

water soluble. Infrared spectra are typical of dry cave complexes.

The spin state of the iron(II) Is unknown because the solid complexes

are too air sensitive to permit magnetic moment determination by the

Faraday method.

Iron(II) Complexes of Unbridged Ligands

In order to more thoroughly evaluate the effects of the

bridging group on the reactions of interest, the syntheses of some

unbridged iron(II) complexes were attempted. In the case of the

unbridged ligand derived from methylamine, the ligand salt was mixed

with one equivalent of bis-(acetonitrlle)iron(II) chloride and four

equivalents of triethylamine to deprotonate the ligand as described

above for the chloro complexes. The product formed upon addition of

ammonium hexafluorophosphate. The expected product was

[Fe{(MeNHEthi)2Me2 [l6]tetraeneN^}(CH2CN)x ](FFg)2 » x = 1,2 and the

infrared spectrum, shown in figure 36, is in support of this formula

tion, having a sharp N-H stretch at 3430 cm ^ and bands at 1580 cm ^

and 1625 cm ^ which resemble those of the corresponding nickel(II)

complex. Absorptions due to coordinated acetonitrile were not

observed, but this is not unusual for acetonitrile bound to

Iron(II) The Ijj NMR spectrum, shown in figure 37, is in support

of the proposed structure, but contains doublets centered at 5.58 and

8.67 ppm which did not appear in the spectrum of the

163

tooo15004000 3000 2000

Figure 36. Infrared Spectrum of[Fe{(MeNHEthi)2Me2 [16]tetraeneN^)(CH3CN)x ] (FF^)^

nickel(II) complex. It will be shown later that this pattern is due to

a novel rearrangement of the ligand structure.

Electrochemical studies of the complex were performed in ace

tonitrile, showing two reversible one-electron oxidations having E^/2

and !e3/4 - Ei/zJ values -0.045 V, 70 mV, and +1.185 V, 70 mV versus+ 2+ 3+Ag/Ag , figure 38. The first oxidation is attributed to the Fe /Fe

couple and the second is a ligand oxidation as was found in the

nickel(II) complex. An additional small wave at +0.4 V is due to the

presence of the rearranged product. Analytical data were consistent

with a mixture of the desired product having one molecule of acetoni

trile and the rearranged species in a ratio of 3:1.

ppm

Figure 37. Proton NMR Spectrum of [Fe{(MeNHEthi)9Me7[16]tetraeneN.}(CH CN) ](PFC)

164

165

- 1.0-0.8- 0.6-0.4- 0.2 V. vs Ag/Ag

0.4 0.00.6 0.2

Figure 38. Cyclic Voltamagram of[Fe{(MeNHEthi)2Me2 [l6]tetraeneN4 >(CH3CN)x ](PF6)2

Attempts were made to synthesize [Fe{ (Me2NEthi) 2*Ie2 [16] —

tetraeneN^}(CH^CN)^](PFg)2 , x = 1,2, using a variety of solvents,

iron(II) sources and counterions. In all cases, the product was much

too soluble to isolate in a pure form and was therefore never crystal

lized or characterized.

Crystal Structures of Two Iron(II)Chloro Complexes

It was of great importance to demonstrate the size and shape of

the cavities of the dry cave ligands in the iron(II) complexes. There

fore two of the complexes were subjected to single crystal X-ray

structural analysis to demonstrate that the ligands had been removed

intact from nickel(II) and chelated to iron(II). In addition, the

Figure 39. ORTEP Drawings of [Fe{(m-Xylyl(NHEthi)2)Me2 [16]tetraeneN^Cl]Cl*20^011

Figure 40. ORTEP Drawings of [Fe{(m-Xylyl(MeNEthi)2)Me2[16]tetraeneN4}ci](PFg)

167

168

TABLE 24BOND DISTANCES (esd) FOR a) [Fe{(m-Xylyl-

(NHEthi)2)Me2tl6]tetraeneNA}Cl]Cl*2CH3OH ANDb) [Fe{(m-Xylyl(MeNEthi)2)Me2[l6]tetraeneN4}Cl](PFfi)

a b a b

Fe-Cl 2.307(3) 2.326(1) C10-C11 1.54(2) 1.503(7)Fe-Nl 2.131(8) 2.123(3) C11-C12 1.54(1) 1.522(8)Fe-N2 2.079(7) 2.107(4) C15-C17 1.54(1) 1.49B(8)Fe-N3 2.074(7) 2.069(4) C15-N5 1.33(1) 1.352(6)Fe-N4 2.163(7) 2.170(4) C16-C18 1.50(1) 1.504(8)Nl-Cl 1.25(1) 1.289(5) C16-N6 1.32(1) 1.345(5)N1-C12 1.49(1) 1.472(8) C19-N5 1.49(1) 1.474(6)N2-C3 1.28(1) 1.291(4) C19-C20 1.50(1) 1.518(6)N2-C4 1.45(1) 1.468(7) C20-C21 1.39(1) 1.382(9)N3-C6 1.50(1) 1.486(7) C20-C25 1.39(1) 1.394(6)N3-C7 1.31(1) 1.296(5) C21-C22 1.34(2) 1.384(7)N4-C9 1.28(1) 1.301(5) C22-C23 1.39(2) 1.376(7)N4-C10 1.43(1) 1.464(8) C23-C24 1.37(1) 1.398(8)C1-C2 1.52(1) 1.472(8) C24-C25 1.42(1) 1.383(6)C1-C13 1.56(1) 1.519(7) C24-C26 1.51(1) 1.508(7)C2-C3 1.44(1) 1.429(7) C26-N6 1.46(1) 1.481(5)C2-C15 1.43(1) 1.405(5) N5-C27 ---- 1.461(5)C4-C5 1.50(1) 1.515(7) N6-C28 ---- 1.472(5)C5-C6 1.52(1) 1.493(8) P-Fl ---- 1.575(4)C7-CB 1.45(1) 1.437(7) P-F2 ---- 1.543(5)C8-C9 1.49(1) 1.461(8) P-F3 ---- 1.578(4)C8-C16 1.41(1) 1.401(5) P-F4 ---- 1.500(4)C9-C14 1.50(1) 1.518(8) P-F5 ---- 1.540(5)

P-F6 _ _ _ 1.567(4)

169

TABLE 25BOND ANGLES (esd) FOR a) [Fe{(m-Xylyl(NHEthi)2)

Me2[16]tetraeneNA>Cl]Cl*2CH3OH AND b) [Fc{Cm-Xylyl(MeNEthi)2)Me2 [16]tetraeneN^}Cl](PF^)

(a) (b) (a) (b)

N2-Fe-Nl 85.1(3) 84.9(2) C15-C2-C1 121.7(8) 120.3(4)

NA-Fe-Nl 83.5(3) 87.4(1) C15-C2-C3 117.9(8) 118.2(5)

N3—Fe-N2 86.8(3) 89.3(2) C2-C3-N2 127.1(8) 126.8(5)

N4-Fe-N3 83.6(3) 84.1(2) C5-C4-N2 113.6(8) 113.2(4)

Cl-Nl-Fe 128.7(8) 129.4(4) C6-C5-C4 117.7(9) 116.6(5)

C12-Nl-Fe 111.3(8) 110.1 (2) C5-C6-N3 112.0(9) 112.7(3)

C12-N1-C1 119.9(9) 120.2(4) C8-C7-N3 124.9(8) 127.7(5)

C3-N2-Fe 120.2 (6) 124.4(4) C9-C8-C7 121.3(7) 120.8(4)

C4-N2-Fe 118.3(6) 118.5(2) C16-C8-C7 116.0(8) 116.1(5)

C4-N2-C3 117.3(7) 116.2(4) C16-C8-C9 122.4(9) 122.3(4)

C6-N3-Fe 116.5(6) 116.9(3) C8-C9-N4 118.3(8) 121.5(5)

C7-N3-Fe 123.9(7) 126.7(3) C14-C9-N4 124.1(8) 120.8(5)

C7-N3-C6 116.8(8) 115.2(4) C14-C9-C8 117.1(8) 117.2(4)

C9-N4-Fe 129.1(7) 129.5(4) C11-C10-N4 113.3(8) 113.8(4)

C10-N4-Fe 111.0(5) 109.4(3) C12-C11-C10 116.5(9) 115.4(5)

C10-N4-C9 119.8(8) 119.6(4) C11-C12-N1 108.5(8) 110.1(4)

C2-C1-N1 118.3(9) 120.6(4) C17-C15-N5 116.5(8) 116.5(3)

C13-C1-N1 125.2(8) 122.2(5) C17-C15-C2 122.9(8) 124.1(4)

C13-C1-C2 116.2(8) 116.7(4) N5-C15-C2 120.5(8) 119.3(5)

C3-C2-C1 120.2(7) 121.4(3) C18-C16-N6 117.0(7) 115.1

170

TABLE 25 (continued)

(a) (b) (a) (b)

C18-C16-C8 123.1(8) 122.9(A) C21-C20-C25 118.2(9) 118.6(A)

N6-C16-C8 119.9(8) 122.0(A) C25-C20-C19 121.5(8) 117.7(5)

C19-N5-C15 12A.A(7) 121.0(3) C22-C21-C20 119.9(9) 120.3(5)

C27-N5-CX5 _ _ _ 122.5(A) C28-C22-C21 122.A (10) 121.1 (6)

C27-N5-C19 - - - 115.5(A) C2A-C23-C22 120.2(9) 119.A(A)

C26-N6-C16 125.3(8) 122.7(3) C25-C2A-C23 117.2(9) 119.2(A)

C28-N6-C16 - - - 121.7(3) C26-C2A-C23 121.9(9) 122.1(A)

C28-N6-C26 - - _ 115.2(3) C26-C2A-C25 120.7(8) 118.A(5)

C20-C19-N5 111.0 (8) 113.7(5) C2A-C25-C20 122.0(9) 119.2(A)

C21-C20-C19 120.3(9) 123.6(A) C2A-C27-N6 112.5(7) 11A.7(5)

171

five-coordinate, high-spin nature of the iron(II) was clearly shown.

The details of the determinations and refinements of the structures of

[Fe{ (m-Xylyl(NHEthl) 2^ e2 [16] tetraeneN^}Cl] Cl* 20^011 and

[Fe{ (m-Xylyl(MeNEthi)2)Me2[16]tetraeneN^}Cl] (PFg) were discussed in the

experimental section of this work.

ORTEP drawing for the two complexes are shown in figures 39 and

40 and the molecular numbering scheme appears in figure 41. Relevant

bond distances with estimated standard deviations (esds) are listed in

table 24 and bond angles are listed in table 25.

It is apparent from the ORTEP drawings and the bond distances

and angles that these two complexes have many features in common.

Considering first the coordination sphere of the iron(II), it can be

seen that the metal ion is five-coordinate in a square pyramidal typee7

C17,

CI3C3

20N2

22

C6 C23CIO

,C26C9,Cll CI4

CI6

cm28

Figure 41. Molecular Numbering Scheme for Chloro-Iron(II) Complexes

172

of structure with the iron(II) drawn out of the macrocycle planeO

toward the axially bound chloride. The displacement is 0.65 A in theO

N-H derivative and 0.54 A in the N-Me derivative and is very similar toO

the displacement of 0.55 A from the porphyrin plane estimated for

deoxymyoglobin.^ The iron-chlorine bond responds to this difference byO O

increasing in length from 2.307 A in the former complex to 2.326 A inO

the latter. The average iron-nitrogen distances are 2.112 A andO

2.117 A for the (NH^Mxyl and (NMe^Mxyl derivatives, respectively,O

slightly longer than the distances of 2.07-2.09 A observed for com-4parable porphyrins, but definitely consistent with the high-spin state

of the iron(II). The six-membered rings of the macrocycle in both

molecules are in the chair form since steric interactions with the

axially bound chloride would result if they were in the boat form.

Examination of the bond distances starting at the metal center

and continuing the planar bridge nitrogen atoms indicates an extensive

delocalization of electron density throughout these segments of the

molecules and is wholly consistent with the previously discussed

(page 160) large electronic effects at the metal center caused by sub

stitution at the bridge nitrogen atom.

In both molecules, the xylene ring is analogously disposed as

that of the distal imidazole in Hb and Mb. It is apparent from

figures 39 and 40 that the xylene ring is more directly centered over

the metal in the hydrogen derivative than in the methyl derivative.

This is a consequence of the geometrical isomerism about the bridge

nitrogen atoms according to which the former is "lid-on" and the

latter is "lid-off." The profound effects of this isomerism on the

173

size and shape of the cavities is clearly illustrated in figures 39 and

AO. In the "lid-on" structure, the bridge rises in a direction perpen

dicular to the macrocycle plane, resulting in a tall, narrow cavity.

In the "lid-off" structure, the bridge lies approximately parallel to

the plane, yielding a shorter, wider cavity. The cavity in the0 O

(NH^Mxyl bridged complex is 7.57 A tall In the front (C22) and 5.A6 A

at the rear (C25), while in the methyl substituted species, the cor-O 0

responding dimensions are 5.02 A and 3.9A A, respectively (measured as

the perpendicular distance from the carbon atom center to the macro

cycle N, plane). The cavity widths as measured between bridge nitrogenO O

atoms are 5.05 A for the N-H species and 7.3A A for the N-Me complex.

When these distances are corrected for the van der Waals radii of theO

atoms involved, the maximum cavity height is A.22 A and the width isO

2.05 A in the "lid-on" species. The corresponding values for theO O

"lid-off" complex are 1.67 A and A.3A A, respectively. The dihedral

angle between the xylene and macrocycle planes is 53° in the

(NH^Mxyl derivative and 30° in the methyl substituted complex. In

solution, it is expected that this angle would be quite variable within

certain sterically controlled limits. For example, the xylene ring in

the "lid-on" species could conceivably rotate to put C25 in the front of

the molecule and C22 in the rear. Such a rotation is prevented in the

"lid-off" structure due to interactions between the hydrogen on C22 and

the macrocycle.

One further feature of the hydrogen derivative which should be

noted is the extensive hydrogen bonding within the lattice. Although

the hydrogen atoms were not located in the structure determination, the

174

distances between nonhydrogen atoms clearly show the extent of inter-O

action. The anionic chloride is 3.26 A from one bridge nitrogen atom,o o

3.01 A from one methanol oxygen atom and 3.32 A from the second methanolO

oxygen atom. The bound chloride is 3.60 A from the second methanolO

oxygen atom and the oxygen atoms are 3.35 A from each other. All of

these distances are well within normal hydrogen-bonding limits for the

involved atoms.

The five-coordinate, high spin nature of the iron(II) center and

the hydrophobic ligand structure clearly demonstrate that these com

plexes are good structural models for deoxymyoglobin. The number of

structural variations which are accessible makes these systems attrac

tive for systematic study of the interaction of the iron(II) center

with small molecules.

Summary of Crystal Structure Results

The crystal structures of six monomeric dry cave complexes

have now been determined (including the structure of [Fe{(1,5-Pent-

(NHEthi)2)Me2[16]tetraeneN^}(C0)(PY)](PFg)2*CH^0H, to be described

later in this work). These structures give a good indication of the

wide range of cavity sizes and shapes which are available. Some rele

vant data are summarized in table 26 for the structures reported in72this work, the nickel(II) para-xylene bridged complex and two cobalt

65complexes having bridges derived from N,N"-l,6-hexanediamine. The

only example of a ’'lid-on’1 structure is that containing the

N-H-meta-xylene bridge; all others are "lid-off." The cavity widthO O

varies from 5.05 A to 7.34 A between bridge nitrogen atoms (2.05 to

TABLE 26

SUMMARY OF STRUCTURAL DATA FOR DRY CAVE COMPLEXES

Width•

L , AHeight, A

(to metal atom)Height, A ( plane)

Slope of Iaomer Bridge Type

Displacement of Metal from

Plane, AComplex a b Front Rear Front Rear

[Ni{(2 -Xylyl(NHEthi)2)Me2~ [16]tetraeneN^}](PF6)2

6.78(3.78)e

7.20(4.20)C

4.55 4.27 4.03 3.08 23" Lid-off --

[Fe{(m-Xylyl(HHEthi)2)He2-[16]tetraeneH^}Cl]Cl*2CHÔH

5.20(2.20)C

5.05(2.05)e

8.22(7.57)C

6.11(5.46)C

7.57 5.35 53* Lid-on 0.65

[Fe{(m-Xylyl(HeNEthi)2)He2~ [16]tetraeneN^}Cl](PF^)

7.37(4.37)*

7.34 (4.34)e

5.54 (5.02)c

4.44(3.94)C

4.96 3.57 30" Lid-off 0.54

[Fe{(1,5-Pent(NHEthi)2)Me2-[16]tetraeneN^}(PY)(CO)](PFfi)2'CH3OH

6.70(3.70)e

6.28(3.28)e

4.72 5.88 3.83 5.31 — — Lid-off 0.046

tCot(1,6-Hex(MeNEthl)2)Me2- [16]tetraeneN^}(PFg)^

6.65(3.65)e

6.76(3.76)e

5.60 4.83 — — — — *“ — Lid-off

[Cot(l,6-Hex(MeNEthi)2)Me2- [16]tetraeneN4}(NCS)2]Cld

7.09(4.09)e

6.92(3.92)e

6.17 4.80 Lid-off

Measured between ethylldene carbon atoms. ^Measured between bridge nitrogen atoms.°Measured to dummy atom at center of plane. ^Reference 65.eCorrected for van der Waals radii.

176O

4.34 A when corrected for van der Walls radii). The cavity height (as

measured from the center of the plane to the bridge) ranges fromo o o o

4.55 A to 7.57 A in the front and 4.27 A to 5.88 A in the rear, where

front refers to the atoms extending furthest over the metal and rear

refers to-the atoms extending over the trimethylene groups of theD O

macrocycle. The maximum cavity height ranges from 2.53 A to 3.99 A

when the distances are corrected for van der Waal radii. As structure

XXVII shows, it is estimated that the cavity must be large enough toO O

accommodate a cylinder of radius 1.40 A and height of 4.35 A if CO is

to bind in a linear manner along the macrocycle axis. It is clear from

the ORTEP drawings that such dimensions are not accessible for the

xylene bridged complexes. It is conceivable however that longer

aliphatic bridges in the "lid-off11 conformation could be flexible

N-

0 -IC-

1.40

1.15

F.e'1.80 — N

B

m m

177

enough to have minimal sterlc interactions with a bound CO molecule by

extending upward toward the rear of the molecule. It has in fact been

shown that a hexamethylene bridged complex has a sufficiently large65cavity to accommodate a thiocyanate ligand, although the Co-N-C bond

is bent 148° due to strong interactions between the bridge and the car

bon and sulfur atoms of the thiocyanate. It will be shown in this work

that for the more restrictive pentamethylene bridged complex there are

strong interactions between the oxygen atom of a bound CO and the

bridge. The fact that the longer bridge interacts most strongly with

the third atom of the bound ligand, NCS, and the shorter bridge inter

acts with the second atom of the bound CO suggests that bridge size

effects should have a very significant role in determining the

stability of adducts which form by the coordination of a small mole

cule to the metal center at the protected axial site. In particular,

variation of the bridging group should be manifest in the equilibrium

constants for the binding of CO to the iron(II) center.

It Is significant that in the nickel(II) para-xylene derivative

and the four coordinate cobalt(II) species the cavity Is wider between

the bridge nitrogen atoms than between the ethylidene carbon atoms. The

opposite is true for all other cases. Furthermore, this inversion is

observed even when the bridging group remains the same, as in the caseO

of the two cobalt complexes. An increase of 0.44 A between theO

ethylidene carbon atoms is accompanied by a decrease of 0.16 A between

the bridge nitrogen atoms. This demonstrates that rotation about the

bond between the bridgehead and ethylidene carbon atoms is one way that

Internal stress is relieved in the bridging part of the molecule.

178

Reactions of Iron(ll) Complexes with Axial Ligands

Before equilibrium studies with carbon monoxide could be under

taken, it was necessary to study the behavior of the iron(II) complexes

with axial ligands. Specifically, the reactions of the iron(II) chloro

and four-coordinate complexes with chloride, pyridine, and

1-methylimldazole were examined. These results comprise the following

discussion.

It was found that upon addition of l-Melm to acetonitrile solu

tions of the complexes there were no observable changes in the elec

tronic spectra, the oxidation potentials remained essentially the same,

and the molar conductances remained consistent with a 1:1 electrolyte

designation. From these observations It was concluded that the chloride

ion remains tightly bonded to the iron(II) center in solution and that

the iron(II) remains five-coordinate. This indicated that the bridge

was effectively functioning to block the approach to the protected site

of potential sixth ligands and five-coordination was maintained as is

required of a good model system.

In order to learn more about the binding of chloride to the

metal center, some electrochemical studies were undertaken. It was

found that for [Fe{(m-Xylyl(NHEthi)2)Me2 [16]tetraeneN^}ci]+

in DMF, the oxidation peak potential shifted from -0.360 V versus + —Ag/Ag when PFg was the anion to -0.435 V when chloride was the anion.

Such a shift in potential is consistent with partial chloride ion dis

sociation in solution or with interaction of a second chloride Ion with

the metal center. Carefully controlled experiments were performed for

for the methyl substituted meta-xylyl and hexamethylene bridged

179

complexes in which the oxidation potential was measured as a function of

chloride ion concentration in acetonitrile. For the (NMe^Mxyl complex,

excess chloride shifted the potential only 30 mV more negative, and no

further changes were evident beyond 100-fold excess chloride. The

results were not as simple for the (NMe)„(CH„), Bpecies however. AsZ Z D

shown in figure 42, in the absence of excess chloride there was a single

oxidation wave with a peak potential of -0.410 V. Upon addition of 100

equivalents of chloride, two waves appeared at -0.365 V and -0.475 V.

The shift to -0.475 is consistent with that observed in the (NMe)Mxyl

example but the position shift remains unexplained. Even in the pres

ence of the 100-fold excess of chloride ion, the ratio of the currents

of the two waves was still changing. Although it is expected that a

second chloride ion would interact more readily with longer and more

flexibly bridged hexamethylene species than with the meta-xylene

bridged one, the data do not conclusively support such a process.

Electronic spectra were measured as a function of chloride ion

concentration for the (NMe^Mxyl and (NH^Pxyl complexes in acetonitrile

and demonstrated only very slight changes indicating no major modifica

tion in the ligand field of the metal ion. This can be explained in two

possible ways. First, a second chloride ion is not interacting with the

metal ion at all or second, either chloride or solvent is always bound

to the sixth coordination site with each yielding identical spectra.

If a sixth ligand were to interact with the iron(II) center, the2 ispin state should change from high- to low-spin. Therefore, XH NMR

spectra were measured in CD^CN for the methyl substituted hexamethylene

and meta-xylene bridged iron(II) chloro species as shown in

180

0.0 - 0.2 -0.4 - 0.8 - 1.0- 0.6

Figure 42. Cyclic Voltamagram of [Fe{(1,6-Hex(MeNEthi^JMe^ 16]- tetraeneN^)Cl](PFg) with a) No Excess Chloride and b) 100 Equivalents of Excess Chloride

181

figures 43 and 44, respectively. It is clear that there are numerous

resonances which are broadened and outside of the normal shift ranges of

diamagnetic species. The implication is that the major component in

solution is a five-coordinate high-spin iron(II) complex. There are

also resonances which can be attributed to a diamagnetic species in the130-10 ppm range for both complexes. The C NMR spectrum of the

(NMe^Mxyl species is shown in figure 45 and confirms the presence of a

diamagnetic species in solution. There are also several broad reso

nances due to the presence of a paramagnetic component. The large num-13ber of scans required to obtain a good C NMR spectrum on a highly

concentrated solution indicates that the diamagnetic species is only a

minor component in the solution.

The above data combine to clarify the solution behavior of the

iron(II) chloro complexes. Scheme IV includes the various equilibria

which can exist simultaneously in solution. Conductivity measurements

are consistent with [FeLCl] as the principle species in solution.

Steric arguments suggest that [FeLC^]0 and [FeLCl(S)]+ are only very

minor components due to interaction between the chloride and ligand

structure. The diamagnetic component in the NMR spectra indicates that

the rate of exchange of the various axial ligands is slow on the NMR

time scale. This is not surprising for iron(II) when spin state

changes associated with changes in coordination numbers and subsequent

ligand modifications must occur in all of the equilibria. Electro-2+chemical observations demonstrate that significant amounts of [FeL] ,

[FeL(S)]^+ , or [ F e M S ^ ] exist in solution in addition to [FeLCl]+.

Addition of excess chloride drives the equilibria toward [FeLCl]

40 20 -20 -40ppm

Figure 43. NMR Spectrum of [Fe{ (1,6-Hex(MeNEthi) 0)Me_ [16] tetraeneN. }cl] (PF,)L L 4 6

182

40 20 -20 -400ppm

Figure 44. XH NMR Spectrum of [Fe{(m-Xylyl(MeNEthi)0)Me„[16]tetraeneN.}cl](PF,)1 L h 6

183

160 120 80ppm 40

Figure 45- 13C NMR Spectrum of lFe{(m-Xylyl(MeNEthi)2)Me2[16]tetraeneN4}Cl](PF6)

184

SCHEME T3ZF e L SL O WSPIN

2 +* ► F e L S 2 + < — — » F e L S C I(HIGH LOWSPIN SPIN

L IF e L Z + 4 = t F e L C ! + 5 = * F e L C I 2 °

INTERMEDIATE HIGH LOWSPIN SPIN SPIN

185

186

which has a more negative potential. The conclusion drawn from these

observations is that the two principle species in solution are the+ 2 paramagnetic [FeLCl] and diamagnetic [FeLCS)^] with the former

species as the major component.

Because the equilibria involving chloride ion were not simple,

several attempts were made to remove the coordinated chloride from the

iron(XI) center. Addition of silver ion to methanolic solutions of

the (NH^Mxyl and (NMeJ^Mxyl iron(II) chloro complexes resulted In the

formation of deep blue iron(III) complexes which were very difficult to

purify. Similar products were obtained when silver ion was added to

chloro complexes of iron(HI). In each instance, analytically pure

samples were not obtained and the electrochemical behavior was very

complex and Ill-defined. Direct synthesis of chloride-free species

met with some success as was described In an earlier section of this

work.

The equilibria for axial base adduct formation with

[Fe{(m-Xylyl(MeNEthi)2)Me2[16]tetraeneN^,}]l2 (equation 29) in aqueous

solution were examined at ambient temperatures.

[FeL]I2 + B <5^ [FeL(B)]I2 ’ (29)

66Stevens has described in detail the method for measuring such

equilibrium constants and for evaluating errors. Total spectral

changes were on the order of 0.25-0.3 absorbance units over the range

of the titration which was adequate to allow good refinement of the

equilibrium constants. K was determined to be 8.48 + 0.58 & mole”'1'

B « Py and 81 + 7 JZ. mole ^ for B ** l-Melm. The value for the pyridine

187

case was essentially independent of wavelength, but for B «= l-Melm the

value was somewhat wavelength dependent resulting in a higher but still

reasonable standard deviation. The stronger Lewis basicity of 1 - Melm

compared with Py is clearly reflected in these data, consistent with49observations for related porphyrin systems.

CO Adducts of Iron(II) Dry Cave Complexes

It was of paramount importance to structurally characterize a

CO adduct of an iron(II) dry cave complex to elucidate the nature of

binding and bending of the CO molecule at the metal center. It was

also necessary to prepare solid CO adducts in order to verify the com

position of the products and to measure the spectroscopic properties of

the complexes. Therefore, a series of six-coordinate iron(II) carbon

monoxide complexes of dry cave ligands were synthesized and charac

terized. In the following discussions, the synthesis and properties of

the complexes are described. The crystal structure of a CO adduct is

then reported followed by equilibrium studies between CO and the

iron(II) complexes.

Synthesis and Characterization. The iron(II) chloro complexes

were found to undergo a very novel reaction with CO as shm- v.

tion 30. The high-spin, five-coordinate iron(II) complex reacts with

a molecule of CO and a molecule of base to produce a six-coordinate,

low-spin iron(II) CO adduct and a chloride ion. The complexes having

(NH^Mxyl, (NMe^Mxyl, (NH^Pxyl, and (NMe^CCHj)^ bridging groups were

synthesized and studied by this author. Complexes derived from the

(NH)2 (012)5 , ( ^ ^ ( O ^ ) ^ , and (NH)2Fluorene bridged ligands were

R

+ xs B + CO

B =

.CO

2 +

(30)

188

18988prepared simultaneously by Dr. J. J. Grzybowski of these laboratories

and will be included in discussions in this work due to their obvious

relevance. In general, the reactions were carried out in methanol or

acetonitrile/ethanol mixtures in the presence of an excess of the

desired axial ligand. Addition of CO caused the solution to change in

color from deep red to orange. Subsequent addition of ammonium hexa-

fluorophosphate yielded the red-orange CO addition products which were

usually highly crystalline. The resulting solids are generally air

stable for days but, In solution, reaction with air occurs within

minutes or hours depending on the particular bridging group and axial

ligand.

Iodide salts of CO adducts were prepared by the reaction of the

iron(II) chloro complex with an appropriate base and CO in acetone.

Addition of excess (nBu)^NI produced the crystalline iodide salts in

high yield. Attempts to produce chloride salts in the same way

resulted in the loss of CO yielding only chloro-chloride species.

CO adducts were also prepared from some of the four-coordinate

iron(II) complexes. The properties of such species matched those of

the adducts prepared from the corresponding chloro complexes, verifying

that the ligand remains intact in the conversion to and from the

four-coordinate species. Again, for purposes of comparison an

unbridged CO adduct was synthesized. Because it was not possible to

Isolate a simple iron(II) complex of the ligand derived from dimethyla-

mine, the carbon monoxide adduct was synthesized directly from the

ligand and iron(II) starting materials in acetonitrile. The product,

identified as [Fe { (Me2NEthi) 2 ^ 2 [16 ] tetraeneN^ } (CH^CN) (CO) ] (PFg ) 2 > was

190

Isolated and characterized. This complex was designed to be similar to

the CO adducts of most other model systems because unhindered linear

binding of the CO to the iron(II) center is expected, thus allowing the

bridged complexes to be compared with an unbridged structure having the

same basic macrocycle.

Analytical data for the CO adducts demonstrated that, in some

cases, mixtures of the desired product and either a chloro complex or

four-coordinate species were obtained. In other cases, some excess sol

vent was present due to the fact that the complexes were not rigorously

dried for fear of removing some of the bound CO. Despite these diffi

culties, the analyses are in support of the proposed formulations.

(See Experimental section.)

The infrared spectra of the CO adducts contain a great deal of

useful information. The N-H and C = C , C = N regions are typical of the

mr.crocyclic species. Of particular interest however are the absorp

tions due to the bound CO between 1900 cm ^ and 2000 cm \ which are

discussed separately later, and those due to the nitrogenous base which

occupies the axial site trans to the CO. In order to eliminate solid

state effects present in nujol mulls, spectra were also measured in

acetonitrile solution and, for some complexes, in other solvents.

The spectra of [Fe{(£-Xylyl(NHEthi)2)Me2[16]tetraeneN^}(B)(CO)]-

wêre ® c CH^CN, Py, and 1 - Melm are shown in figures 46a, b, and

c, respectively. The N-H stretch at 3420 cm ^ and the bands between

1550 and 1650 cm ^ are typical of the bridged macrocycles. The CO

stretching mode gives bands at 1985, 1980, and 1970 cm ^ when the base

is CH^CN, Py, and l-Melm, respectively. Two weak absorptions near

u

r \

r

%y r-191

I1000

I500

I1500

I20004000 3000 cm

Figure 46. Infrared Spectra of [Fe{(j>-Xylyl(NHEthi)2)Me2[16]tetraeneN^}- (B)(C0)](PF6)2, a) B = CH3CN, b) B = PY, c) B = 1 - Melm

192

2290 and 2370 cm are due to the bound acetonitrile molecule In the

CH^CN derivative. (Free acetonitrile is observed as a single sharp

absorption at 2250 cm Bound 1 - Melm is identified by the absorp

tion at 3140 cm ^ which is caused by the C-H stretch at the carbon

between the two imidazole nitrogen atoms. Absorptions due to bound

pyridine are obscured by ligand bands in the 1500-1650 cm ^ region, but

extra bands in the aromatic region between 600 and 800 cm ^ are

generally present in pyridine derivatives. The spectra described above

are typical of the solid state spectra observed for monomeric dry cave

complexes.

Carbon-13 NMR spectra were measured for a number of the CO

adducts in CT^CN solution at ambient probe temperatures (~40°C) using

a five second pulse delay. Representative spectra for the (NMe)^”

(CH^)^ bridged CO adducts having CH^CN, Py, and 1 - Melm as the trans13axial ligand are shown in figures 47, 48, and 49, respectively. C

NMR data for a number of monomeric CO adducts are listed in table 27.

Simple mirror symmetry is present for the unbridged complex

derived from dimethylamine and for all bridged complexes containing

CH^CN as the trans axial ligand, as shown in figure 47 for the (NMe^-

^■^2^6 examPê * Similar spectra are anticipated for pyridine deriva

tives but are not observed. Instead, there is a doubling of most

resonances except for those due to pyridine. Careful examination of

the spectra shows that the line doubling is due to the presence of both

pyridine and acetonitrile derivatives in solution as a result of ligand

exchange at the trans axial site. The two species are present in nearly

equal concentrations.

rV Al* Ja ... ,n t..j

240 200 160 120ppm 80 40

13Figure 47. C NMR Spectrum of tFe{(l,6-Hex(MeNEthi)2)Me2[16]tetraeneN4}(CH3CN)(CO)](PFg)2 193

240 200 160 120ppm

60 40

13Figure 48. C NMR Spectrum of [Fe{(l,6-Hex(MeNEthi)„)Me„[16]tetraeneN.}(PY)(C0)](PF,)„Z Z q 6 2

194

240 200 160 120ppm

00 40

13Figure 49. C NMR Spectrum of [Fe{(1,6-Hex(MeNEthi)2>Me2[16]tetraeneN^}(l-Melm)(CO)](FF6>2

195

196

TABLE 27CARBON-13 NMR DATA* FOR CO ADDUCTS OF MONOMERIC IRON(II)

DRY CAVE COMPLEXES, [Fe{(R)Me2 [16]tetraene^}(B) (CO) ] (PFg)2

B Chemical Shifts

l,5-Pent(NHEthi)2 PY

1,6-Hex (MeNEthi) 2 C ^ C N

1,6-Hex(MeNEthi)2 PY

1,6-Hex (MeNEthi) 2 1 - Melm

m-Xyly1(NHEthi)2 Im

m-Xylyl(MeNEthi)2 Im

220.1, 219.1, 172.A, 170.6, 165.1, 163.6,161.9, 152.8, 137.7, 124.6, 109.8, 57.3,49.A, 49.0, 45.9, 30.1, 29.0, 26.9, 19.7,19.2, 16.2

222.7, 172.1, 167.2, 165.5, 111.1, 60.3,56.7, 52.0, 38.3, 31.8, 30.6, 25.6, 24.3,21.9, 16.8

224.0, 174.0, 172.6, 168.8, 167.6, 167.3,166.0, 155.4, 140.5, 127.5, 111.8, 111.6,60.8, 57.6, 57.2, 52.5, 52.2, 39.0, 38.8,31.8, 30.4, 26.1, 24.9, 22.7, 22.3, 17.8,17.3

223.8, 173.2, 172.7, 168.4, 167.9, 166.5,151.0, 131.1, 119.2, 112.0, 111.8, 61.3,57,9, 57.4, 52.7, 52.6, 38.9, 38.8, 32.2,30.7, 26.1, 24.9, 22.7, 18.6, 17.6, 17.5,12.5

218.4, 174.0, 173.8, 168.5, 168.1, 161.8,141.4, 140.1, 130.6, 130.3, 127.5, 119.9,112.7, 112.5, 60.1, 52.1, 51.9, 49.6, 30.9,30.3, 22.6

216.5, 174.3, 174.1, 169.1, 168.8, 164.8,164.6, 141.2, 140.3, 130.9, 130.5, 129.1,126.4, 119.8, 114.3, 60.5, 52.3, 52.1,44.5, 30.5, 30.3, 22.7, 16.8

197

TABLE 27 (continued)

R B Chemical Shifts*1

£-Xylyl(NHEthl)2 CH3CN 213.2, 169..0, 165.0, 156.3, 135.3, 127.1,126.8, 109.,2, 56.9, 48.6, 45.7, 28..6,21.7, 17.4,, 12.0

(Me2NEthi)2 CH3CN 223.1, 171.,0, 170.4, 164.2, 109.3, 65.8,60.1, 52.2,, 43.2, 30 .4, 28.4, 21.6,, 18.2,15.2

Measured on concentrated solutions in CD^CN.

kppm relative to TMS.

198

When an unsymmetrical Imidazole base is used, a doubling of

peaks is again observed but the intensities for related carbon atoms are

essentially the same in all cases. The positions of the resonances

establish the absence of an acetonitrile adduct, thereby eliminating

axial base exchange as the cause of the doubling. The conclusion is that

the observed spectrum is due to the asymmetry of the imidazole ligand,

which rotates and exchanges only slowly on the NMR time scale. The

higher base equilibrium constants for the binding of imidazoles rela

tive to pyridine, as shown earlier, are in support of this explanation.

Crystal Structures of a CO Adduct of an Iron(II) Dry Cave Complex

The importance of the bending of the Fe-C-0 linkage in natural

systems has already been emphasized (page 6). There is no precedent

for such bending in model systems^ ^ and thus it was highly desir

able to structurally characterize a CO adduct of an iron(II) dry cave

complex to determine unambiguously the nature of the bending in the

Fe-C-0 linkage.

Dr. J. J. Grzybowski of these laboratories generously supplied

crystals of the complex [Fe{(1,5-Pent(NHEthi)2)Me2[16]tetraeneN^}-

(CO) (PY)] ( P F g ^ ’CHÔH. The details of the structural refinement were

discussed in the Experimental section of this work. ORTEP drawings of

the molecule are shown in figure 50 and the molecular numbering scheme

appears in figure 51. The bond distances and angles with esds are

listed in tables 28 and 29, respectively. The quality of the structure

is excellent as is indicated by the small esds and the particularly low

Figure 50. ORTEF Drawings of [Fe{(l,5-Pent(NHEthi) )Me [16]tetraeneN.}(PY)(C0)](PFj_*CH OH£■ £■ h 6 2 3

199

200

TABLE 28BOND DISTANCES (esd) FOR [Fe{(l,5-Pent(NHEthi)2)

Me2 [16]tetraeneN4)(PY)(CO)](PFfi)2«CH3OH

Fe-Nl 1.997(3) C16-C17 1.387(5) C10-H103 0.93(4)

Fe-N2 1.986(3) C18-02 1.56(2) Cll-Hlll 0.99(4)

Fe-N4 2.052(3) P-Fl 1.590(3) C11-H112 1.00(3)

Fe-C14 1.792(4) P-F2 1.595(2) C12-H121 0.97(4)

Nl-Cl 1.295(4) P-F3 1.576(4) C12-H122 0.97(4)

N1-C7 1.478(4) P-F4 1.551(3) C13-H131 0.96(4)

N2-C3 1.295(4) P-F5 1.537(4) C13-H132 0.98(5)

N2-C4 1.478(4) P-F6 1.555(4) C15-H151 0.90(4)

N3-C9 1.341(5) C3-H31 0.99(4) C16-H161 0.93(4)

N3-C11 1.467(6) C4-H41 0.97(4) C17-H171 0.83(4)

N4-C15 1.345(4) C4-H42 1.05(3) N3-NH31 0.79(3)

C1-C2 1.465(5) C5-H51 0.95(6)

C1-C8 1.516(6) C5-H52 0.90(4)

C2-C3 1.446(5) C6-H61 1.00(4)

C2-C9 1.402(4) C6-H62 0.95(5)

C4-C5 1.518(5) C7-H71 1.01(3)

C6-C7 1.501(5) C7-H72 1.03(4)

C9-C10 1.487(6) C8-H81 0.95(3)

C11-C12 1.513(5) C8-H82 1.05(4)

C12-C13 1.538(6) C8-H83 0.88(4)

C14-01 1.150(5) C10-H101 1.03(3)

C15-C16 1.370(4) C10-H102 0.93(4)

201

TABLE 29BOND ANGLES (esd) FOR [Fe{(l,5-Pent(NHEthi)2)

Me2 [16]tetraeneN4>(PY)(CO)](PF6)2 ‘CH3OH

Nl'-Fe-Nl 90.2(2) C8-C1-C2 116.4(3)

N2-Fe-Nl 87.8(1) C3-C2-C1 119.2(3)

ClA-Fe-Nl 93.7(1) C9-C2-C1 124.0(3)

N4-Fe-Nl 93.1(1) C9-C2-C3 116.7(3)

N2"-Fe-N2 94.0(2) C2-C3-N2 123.0(4)

Cl4-Fe-N2 83.7(1) C5-C4-N2 110.9(4)

N4-Fe-N2 89.7(1) C4"-C5-C4 115.0(4)

Cl4-Fe-N4 170.3(2) C7 "-C6-C7 114.5(5)

Cl-Nl-Fe 121.2(3) C6-C7-N1 111.4(3)

C7-Nl-Fe 119.1(2) C2-N9-N3 122.8(3)

C7-N1-C1 119.5(4) C10-C9-N3 114.7(3)

C3-N2-Fe 120.3(3) C10-C9-C2 122.5(3)

C4-N2-Fe 121.5(2) C12-C11-N3 111.2(3)

C4-N2-C3 117.1(3) C13-C12-C11 115.5(3)

C11-N3-C9 127.7(3) C12 "-C13-C12 115.3(6)

Cl5-N4-Fe 121.3(2) 01-C14-Fe 170.6(5)

C15"-N4-C15 116.7(4) C16-C15-N4 123.6(4)

C2-C1-N1 120.3(4) C17-C16-C15 118.7(4)

C8-C1-N1 122.8(4) C16'-C17-C16 118.6(5)

C3 qi-

Qt

•ce

/C5

N 2 c

N3.

CI2

202

II

C4

Fe

*•s.-\

C(

/

CI6

C13

CI2

-c17/CI6'

Figure 51. Molecular Numbering Scheme for [Fe{(l,5-Pent(NHEthi)„)Me9- [16]tetraeneN4)(FY)C0)](PF6)2*CH3OH 2 2

GOF value of 1.17. Therefore, bond distances and angles are very well

defined and can be interpreted with confidence.

The ORTEP drawings clearly show the coordination sphere of the

iron(II) with the macrocycle bound equatorially, pyridine bound axially

and carbon monoxide bound axially within the dry cave. The iron(II) isO

low-spin and is displaced but 0.05 A out of the macrocycle plane

toward the pyridine. The average iron-macrocycle nitrogen bond dis-® 56tance of 1.992 A is typical of low-spin iron(II) porphyrin complexes.

OThe iron-pyridine bond distance of 2.052 A is the shortest of the three

reported iron(II) pyridine carbon monoxide adducts. Table 30 contains

a summary of relevant data for Fe(TPP)(CO)(PY)"^ and three CO adducts58reported by Goedkln.

TABLE 30

SUMMARY OP STRUCTURAL DATA FOR CO ADDUCTS OF MODEL SYSTEMS

Complex- ■ s~ -

Fe-H(Mac), A1 * Fe-N(B), A Fe-C, A C-0, A <Fe-C-0,° <C-Fe-N(B) vco’ co'1

[Pet <1,5-Pent(NHEthi),)Me,[16]- 1960dtetraeneNA)(CO)(Py)(PF6)2*CH3OH 1.995(3) 2.059(3) 1.792(4) 1.150(5) 170.6(5)C 170.3(2) 1952e[Fe(TPP)(CO)(PY)]a 2.02(1) 2.10(1) 1.77(2) 1.12(2) 179(2) 177.5(8) 1980f[Fe(C22H22H4)(CO)]b 1.927(4) ---- 1.694(4) 1.157(2) 177.2(3) ---- 1915*tFe(C22H22N4)(FY)(C0)]b 1.941(2) 2.088(3) 1.730(3) 1.146(3) 178.2(3) 177.2(1) 1940f[Fe(C22H22H4)(H2H4)(CO)]b 1.943(4) 2.122(5) 1.751(5) 1.137(6) 177.6(5) 179.2(2) 1930*

Reference 56.^Reference 58.CThls alone does not adequately describe the displacement of the CO from a linear orientation. Âcetonitrile solution.*Nujol mull.Âlthough conditions are not given, these are presumed to be solid state values.

203

204O

The iron-carbon and carbon-oxygen distances are 1.79 A and

1.150 A respectively. This is the longest Fe-C bond of those reported

and is the second longest carbon-oxygen bond. By far the most signifi

cant feature of the structure is the bent Fe-C-0 linkage which has an

angle of 170/>°. All other structures reported for model systems

exhibit an almost linearly bound CO. Furthermore the carbon is dis-O

placed from the macrocyclic axis by 0.76 A thus showing that the carbon

monoxide bends both at the iron center and at the carbon atom. The

cause of this bending is steric interactions between the bound CO and

the pentamethylene bridging group. The cavity is quite restrictiveo D

having a width of 6.28 A between nitrogen atoms (3.28 A when corrected

for van der Waals radii). The bridge is also quite low, with a heighta o

of 3.83 A from Cll and 5.31 A from C13 to the plane. Most impor

tantly, the hydrogens of the bridge point into the cavity at the bound

CO. Thus it appears that there is little room within the cavity toQ

accommodate a ligand. As a result, the CO oxygen atom is 2.62 A fromO

the hydrogen atom on Cll and 2.92 A from the hydrogen atom on C12.

These distances are very close to the van der Walls contact distance ofO

2.72 A indicating strong non-bonded interaction and explain the observed

CO bending. The bridge is in turn distorted, having C13 pushed as farO

as possible from the bound CO. Furthermore, the oxygen atom is 2.77 A

from the hydrogen atom of C17 of the pyridine belonging to the next

molecule in the lattice. As the packing diagram in figure 52 shows,

this interaction is in a direction which should force the CO toward a

more linear orientation. Thus it is expected that the degree of bending

of the Fe-C-0 linkage is possibly even greater in solution where such

205

Figure 52. Packing Diagram for [Fe{(l,5-Pent(NHEthi)2)^6 2[16]tetraene- }(CO)(PY)](PFfi)2 * CH3OH

206

intermolecular interactions are unimportant and intramolecular interac

tions are dominant. From these observations two major conclusions are

drawn: First, the Fe-C-0 linkage is bent both at the iron and at the

carbon atoms. Such an arrangement can reasonably be expected in the

natural systems. Secondly, the dry cave complexes are verified to be

excellent models for demonstrating the interaction of Hb and Mb with

carbon monoxide. The bridging group appears to play a similar steric

role as the distal imidazole functions in the natural systems. This is

the first example of a model system which has been shown through struc

tural analysis to contain a bent Fe-C-0 linkage. It is obvious from

the drawings of this molecule that the degree of bending of the Fe-C-0

linkage can be precisely controlled by careful selection of the

bridging group and thus the effects of this bending on the physical

properties of the complexes can be systematically studied as will be

described in the next section of this work.

The packing diagram in figure 52 also shows that the pyridine

molecule is bent relative to the macrocycle axis. Such a bending is

indeed unusual but is readily explained in terms of packing forces.

(This is accidental in our CO complex, but serendipitously mimics devi

ations from axial alignment of the proximal imidazole in the natural 97systems. ) It has already been mentioned that the pyridine of one

molecule contacts with the CO oxygen atom of a second. The effect of

this interaction on the pyridine is to cause it to bend in the observed

direction. A further interaction is seen between the hydrogen of C16

and one of the hydrogens of CIO of the next molecule which has aO

distance of 2.53 A. This repulsive interaction is in the same direction

207

as the first and together they result in the observed deviation from

axial alignment. In solution this interraolecular interaction does not

exist and can therefore be disregarded in solution studies.

Equilibrium Studies of the Reaction between Carbon Monoxide and Monomeric Iron(II) Dry Cave Complexes

A central goal of this work is the assessment of the effect of

the bridging group of dry cave ligands on the binding of CO to the

iron(II) center. By means of a systematic study of the CO binding con

stants as a function of cavity size, the effects of bending of the

Fe-C-0 linkage can be quantified. Such a study was undertaken in this

work and demonstrates that as the cavity size becomes more restrictive,

the CO binding constant decreases. Equilibrium constants spanning a

range of at least four orders of magnitude were observed due solely to

changes in the nature of the bridging group. In addition, the inhibi

tory effect of chloride ion on the binding of CO was clearly demon

strated for several of the complexes. The method used for equilibrium

measurements utilized the flow system designed and described by

Stevens.^ The basic technique and the modifications which were made to

study reactions with CO instead of 0^ are described in the Experimental

section of this work.

The simplest complex examined in detail was [Fe{(m-Xylyl-

(MeNEthi^H^tlbJtetraeneN^}] (PF^Jg, the well characterized

four-coordinate iron(II) dry cave complex. This complex was selected in

order to avoid the equilibria involving chloride in the chloro deriva

tives. Preliminary results had indicated that the equilibrium constants

were such that they could be determined over a broad range of

208

temperatures using the accessible pressures of CO. To minimize the

number of equilibria in solution, the studies were carried out in ace-

tonitrile with no added axial base. It was assumed that in acetoni-

trlle solution the principle species was five-coordinate with

acetonitrile as the axial ligand. The expected reaction with CO is

that shown in equation 31. The spectral changes

[{FeL}(CH3CN)]2+ + CO =5= ^ [{FeL> (CHgCN) (CO) 2+ (31)

observed upon addition of increasing partial pressures of CO are shown

in figure 53. The absorbance increases above 413 nm, maxima at 360 and

287 nm decrease in intensity and new maxima grow in near 332 and 257 nm

upon introduction of CO. Reasonably sharp isoshestic points are

observed at 413, 342, 302, and 274 nm. The spectral changes are essen

tially reversible upon flushing the solution with pure nitrogen.

The CO equilibrium constants for this complex at several dif

ferent temperatures are listed in table 31. The values listed are the

average values determined at three or more different wavelengths.

Although there was some variation in the calculated values at different

wavelengths, no systematic errors were observed. The standard devia

tions calculated by the refinement program are all less than 10% and

are believed to be accurate estimates of the experimental error.

Since was known over a range of temperatures, a van’t Hoff

plot was generated to calculate the thermodynamic parameters AH and AS

(figure 54). A linear least squares fit to the data yield values of3(-9.8 + .6) x 10 kcal/mole for AH and -38 + 2 eu for AS as obtained

209

0 .9

0,8

0.7

0,6

0.5-

0 .4

0 .3 -

0.2-

3 5 0 4 0 0 4 5 0

Figure 53. Spectral Changes for the Reaction of [Fe{(m-Xylyl-(MeNEthi)2)Me2 [l6]tetraeneN4}(CH3CN)](PF6)2 with CO inCH3CN at 0°C

210

TABLE 31

CO EQUILIBRIUM CONSTANTS FOR [Fe{(m-Xylyl(MeNEthi)2)Me2 [16]tetraeneN4}(CH3CN)](PF6> 2

IN ACETONITRILE

T,° C K, . -1 torr

-29.4 2.3 ± *2-19.1 1.6 •—1 •+1

-10.0 0.84 1 + • o -F-

0.0 0.43 + .03

+ 9.8 0.16 + .01

+20.0 0.086 + .008

211

2.0

0.0

- 1.0

- 2.0R InK

-3.0

-4.0

-ao

- 6 0

-7.0

3 4 3.5 3.6 3.9 4.03.8

Figure 54. Van't Hoff Plot for [Fe{(m-Xylyl(MeNEthi)2)Me2 [16]tetraene- N4 }](PF6)2 in CH3CN

212

from the slope and Intercept, respectively. The correlation coeffi

cient for the line was 0.9893. This value for AH compares favorably98with the value of -11.5 + .5 kcal/mole reported for hemoglobin.

Careful examination of the plot reveals an apparently systematic devia

tion from linearity, however, indicative of a systematic problem which

is temperature dependent. It is possible that an irreversible reaction

is occurring with residual traces of oxygen in the flow system as a

result of leaks of insufficient scrubbing of the CO. Such reactions

are known to be irreversible and temperature dependent and thus the

effect of such a reaction could well be systematic. It is also pos

sible that the equilibrium between acetonitrile and the complex is not

saturated under the experimental conditions. The result of this would

not be a temperature dependent change In the concentration of the reac

tive acetonitrile adduct in solution. would therefore respond to

such an equilibrium in the observed manner. As a result of this, the

values of the thermodynamic parameters must be interpreted with

caution.

In order to evaluate the effects of axial base on the CO

equilibrium, the above complex was studied in acetonitrile solution

containing 10% 1 - Melm at 10°C. Although isosbestic spectral changes

were not observed, an approximate value for of 0.489 + .054 torr ^

was determined which is about 5 times greater than the value observed

at the same temperature in the absence of l-Melm. This observation

confirms the importance of IT backbonding in stabilizing the iron-carbon

bond, since it is enhanced by the good Lewis base l-Melm.

213

The behavior of the complex believed to be the four-coordinate

iron(II) derivative of the (NMe) 2(0112)4 bridged ligand was examined at

-8 .6°C in acetonitrile. was found to be 0.104 + 0.020 torr \

approximately 10 times smaller than the value for the (NMe^Mxyl com

plex under comparable conditions. Although the relative error is

rather large for this case, the observed Kcq is wavelength independent.

The very small value of K is presumed to be a result of strongLUsteric interaction between the bound CO and bridge, causing a weakening

of the iron-carbon bond.

The four-coordinate iron(II) complex having the (NMe^^l^)^

bridge was studied in acetonitrile at 20°C. The spectral changes were

essentially complete after introduction of the lowest accessible par

tial pressure of CO (0.2 torr), indicating a value of too large to

measure under the available experimental conditions. These results

are all in good agreement with expectation that the observed equilibrium

constants should be dependent upon the nature of the bridge and size of

the resulting cavity.

Due to the synthetic difficulties encountered In the synthesis

of four-coordinate complexes and the need to measure as a function

of bridge size, the well characterized chloro-iron(II) complexes were

examined. The studies revealed a behavior that Is quite complex but

yielded some very useful information. Preliminary studies indicated

that as the concentration of chloride in solution increased, the

observed decreased. It is possible that chloride ion is competing

with CO for the vacant coordination site of the metal. It is also pos

sible that chloride ion is involved in another equilibrium with the

214

iron(II) complex which affects the CO binding constant. In order to

determine the role of chloride ion, a detailed study was undertaken

using [Fe{(£^XylylCNHEthi)2)Me2 [16]tetraeneN^,}Cl] (PFg) . This system

was selected because the spectral changes upon addition of CO were iso-

sbestic (figure 55) and the equilibrium constant was of a magnitude

which allowed for accurate measurement of K^q over a wide range of

chloride ion concentrations. was determined as a function of

chloride ion concentration in acetonitrile solution at 0.0°C. The

total ionic strength was maintained at 0.1 molar with (nBu)^NBF^. As

the data in table 32 clearly demonstrate, the value of spans more

than four orders of magnitude as the chloride ion concentration ranges

from 0.0 to 0.1 molar. A model has been developed which is consistent

with the observed data and allows the equilibrium constant for the

binding of chloride to the metal to be determined. The model is based

on the three equilibria defined by equations 32-34.

[FeL(CH3CN)]2+ + CO *5= ^ [FeL(CH3CN) (CO) ]2+ (32)

[FeL(Cl)]+ + CO =5=^* [FeL(Cl)(C0)]+ (33)

[FeL(CH3CN)] 2+ + Cl [FeL(Cl)]+ (34)

The observed equilibrium constant for CO binding is defined as

= [FeL(CO)] OBS [FeL]pC0 (35)

where

[FeL] = [FeL(CH3CN)] + [FeL(Cl)]

[FeL(CO) ] *= [FeL(CH3CN) (CO)] + [FeL(Cl) (CO) ] (37)

(36)

215

09-

0.5-

0.4-

0.2-

0 .1-

— I— --------- 1______________ L400 450 500

Figure 55. Spectral Changes for the Reaction of [Fe{ (£-Xylyl-(NHEthi)2)Me2 [16]tetraeneN4>Cl](FF6) in CHgCN at 0.0°C,[Cl"] = 1 x 10“3 molar

TABLE 32EQUILIBRIUM DATA FOR THE REACTION OF THE

PARA-XYLENE BRIDGED IRON(II) COMPLEX WITH CO AND CHLORIDE

[Cl . M Kco* torr"1K-Kk 2-k

0 7.91 + .88 0

1.0 X 10"3 0.0253 + .0011 353 + 43

3.0 X 10"3 0.0144 + .0008 691 + 79

4.0 X 10" 3 0.00903 + .00047 1308 + 184

7.0 X 10"3 0.00732 + .00043 1827 + 2.6

1.0 X 10-2 0.00547 + .00026 3196 + 549

5.0 X 10"2 0.00280 + .00026 CO

1.0 X 10"1 0.00300 + .00021 oo

217

From these equations, the relationship between K__„ and [Cl ] can beOddderived (Appendix C), to yield

K„(K„-K..)[C1 ]‘3 2 1l + k3 [ci-] (38)

In the absence of chloride ion, Krt__ = K. and at infinite chloride ionOd d 1concentration = Kg. For this system K ^ g was found to be 7.9

determined at chloride concentrations greater than 0.05 molar. When

these two values are known, equation 38 simplfies to a linear form in

which Kg, the chloride ion binding constant is the slope

data to a functional form y = ax + b. It is gratifying that the inter

cept (b term) is well within standard error of being zero, as would be

expected for the functional form, y = ax, appropriate to the model

equation, equation 39. This large value for the binding constant of

chloride ion to the iron(II) center is entirely consistent with the con

ductance, electrochemical, and NMR measurements described earlier for

the iron(II) chloro complexes in acetonitrile.

measured. For the (NH)gPxyl bridged complex, the isosbestlc spectral

torr ^ in the absence of chloride ion and 0.003 torr ^ was a limit

(39)

A plot of the data in table 32 using the values of and Kg given above

yields a straight line (R e 0.9877 ) (figure 56)with a slope of 5 -1(3.1 + 0.2) x 10 mole which is Kg for a least squares fit of the

The model described above is rigorous and should be applicable

to other systems. The model is applicable provided that can be

218

K-KK£K

36001

3200

2 8 0 0

2 4 0 0

2000

1600

1200

8 0 0

4 0 0

.062 .003 .0 0 4 .005 0 0 6[Cl“], M

.001 .0 0 7 .0 0 8 .009 .0(0

Figure 56. Plot to Determine KC1 for the para-xylyl Bridged Complex

219

changes (figure 55) facilitated determination of K___. This isosbesticUJddbehavior was purely fortuitous, however, since up to four absorbing

species may be present in solution at any given time. The data indi

cate that, for this system, the two possible CO adducts and two pos

sible non-adducts have very similar spectral properties, resulting in

apparent isosbestic spectral changes. Alternatively, one or more of

the components may be present in such low concentrations that its con

tribution to the observed spectrum is negligible. For the case at hand,

the spectra of the non-adducts in the presence and absence of chloride

are essentially identical and the spectral changes associated with CO

adduct formation are very similar in the presence and absence of

chloride ion. Thus it appears that the para-xylene bridged species was

ideal for this type of study. Spectral changes for the (NH)2Mxyl and

(NMe)2(CH2)g bridged compounds were complex (figure 57) and demonstrate

the potential difficulties. Knowing the nature of the equilibria

involved, it should be possible to design experiments and interpret

data in a way which allows for the determination of all three equili

brium constants in other related systems.

The large difference between K^ and K2 is explainable in terms

of the Lewis basicity of acetonitrile versus chloride. The better base,

acetonitrile, donates more electron density to the metal, thus

enhancing 7T backbonding from the metal to the bound CO, thereby

strengthening the M-C bond. Sterically, the chloride may prevent the

iron(II) from readily moving into the plane whereas acetonitrile

would not be expected to sterically interact with the macrocycle.

220

o.r

0.6-

oi>-

0.4-

0.3-

0.2-

0.1-

Figure 57.

T o o 450 500

Spectral Changes for the Reaction of [Fe{(1,6-Hex- (MeNEthi)2)Me2 [16]tetraeneN4}Cl](PFg) in CH3CN at 0.0°C,(Cl ] =0.1 molar

221

Having explained the role of chloride in the binding of CO to

iron(II) dry cave complexes, the effect of changes in bridge type on

KC0 cou^ now determined. Several iron(II) chloro complexes were

studied in acetonitrile solution at 0.0°C in the presence of varying

concentrations of chloride at a total Ionic strength of 0.1 molar. The

data are listed in table 33. In all cases examined, the observed

equilibrium constant decreases as the chloride ion concentration

increases. The observed has a range of more than four orders of

magnitude for the (NH)2Pxyl species over a range of chloride ion con

centrations of 0-0.1 molar. The same trend is observed for the other

species but several constants were either too large or too small to

measure within the available temperature and chloride ion concentra

tion ranges. Although difficulties were encountered for many of the

measurements, due to complex spectral changes and interference by

traces of 02 » some important conclusions can be drawn from the data.

When the values at one temperature (0.0°C) and one chloride ion con--3centration (1 x 10 molar) are compared the effects of changing the

bridging group are clearly shown. The relatively unhindered (NMe)^-

(CH-), bridged complex has by far the largest K„_ while the penta- 2 b CUmethylene and tetramethylene complexes have the smallest The

xylyl bridged complexes lie between these extremes, with the ordering in

decreasing magnitude of KCq shown below. This ordering is in full

agreement

(NMe)2 (CH2)6 » (NH)2Pxyl- (NH)2Mxyl > (NMe^Mxyl > <NH)2 (CH2)5

£ (NMe)2 (CH2)4

TABLE 33

CO EQUILIBRIUM CONSTANTS3 FOR DRY CAVE COMPLEXES AT 0.0°C IN CILjCN

""^Bridget c r i T v ^ ^(NH2)Pxyl (NMe) 2Mxyl (NH)2Mxyl (NMe)2(CH2)6 (nh)2 (ch2)5 (NMe)2(CH2)A

0 7.91 4.27 x 10_1 — — — TLTM — — — — — —

8.0 x 10"5 A.75 x 10_1 2.5 x 10"2 0.5 TLTMC - — — — _ _, -3 -2 -3 -2 r b bo•pH x 10 2.53 x 10 1.2 x 10 J 1.0 x 10 TLTM TSTM TSTM

1.0 x 10”2 5.47 x 10“3 TSTMb - - - TLTMC ----- -----

1.0 x 10_1 3.0 x 10"3 TSTMb ----- 1.0 ----- -----

3 -1Torr , ionic strength 0.1 molar.b -3 —Too small to measure (K < 1.0 x 10 torr ).

cToo large to measure (K » 10 torr .

223

with the available structural information. The cavity of the hexa-

methylene bridged complex is by far the largest and most open and

therefore is expected to interact the least with bound CO. The

"lid-off" configuration of the (NH^Fxyl complex displaces the xylene

ring from being directly over the metal, whereas the "lid-off" (NMe^-

Mxyl bridge is more nearly centered over the metal and should therefore

interact more strongly with the bound CO. The "lid-on" (NH^Mxyl

bridge is taller allowing a more nearly linear binding of the CO and a

higher KCQ. The known structure of the CO adduct of the (NH)2 (0112)5

bridged species clearly shows the degree of interaction between the

bound CO and the bridging group. This interaction is expected to be

even greater in the more restrictive tetramethylene bridged complex.66 73The results reported by Stevens ' for cobalt(II) complexes reacting

with oxygen are in full agreement with this expectation.

The importance of these results must be emphasized. Under

identical conditions of study, the observed equilibrium constant varies

in a systematic manner over a range of at least four orders of magni

tude due solely to changes in the nature of the bridging group. The

crystal structure of the CO adduct clearly shows the nature of the

bending and bonding of CO in these model systems and the equilibrium

data demonstrate the effects of bending of the Fe-C-0 linkage on the

affinity of the Iron(II) site for CO.

From the data presented in this section, I conclude that

two major factors control the overall binding of CO to the iron(II)

center of monomeric dry cave complexes: 1. The nature of the bridging

group controls the steric interactions between the ligand

224

superstructure and the bound CO. 2) Chloride ion acts to inhibit the

binding of CO to the metal center. A combination of these two effects

permits systematic variation of the observed CO binding constant over a

range of many orders of magnitude. It is apparent that with such a

wide range of CO equilibrium constants, the values observed for Hb and

Mb can be duplicated in this totally synthetic model system.

Correlations Between Physical Properties of CO Adducts and Equilibrium Constants

The equilibrium studies reported in the previous section

clearly demonstrate the effects of the bridging group of the dry cave

ligand on the stability of the CO adducts which form. It was expected

that some of the physical properties of the CO adducts should also show

trends which correlate with the equilibrium studies. Therefore,

detailed infrared, carbon-13 NMR, and electrochemical studies were

undertaken. The results of these studies comprise the following dis

cussion.

Infrared Studies. Solid state and solution infrared spectra

were measured for all of the CO adducts and the stretching frequencies

are listed in table 34. It is apparent that in general, in ace

tonitrile solution is greater than in the solid state. Due to the role

of packing forces and dielectric constant of the environment in deter

mining the energy of the absorptions in the solid state, comparisons

will be made only among the solution spectra.

The first trend to note is that for any given bridge, occurs

at highest energy when the axial base is acetonitrile and at lowest

energy when the base is l-Melm. Pyridine adducts have CO stretching

TABLE 34 225SUMMARY OF INFRARED, ELECTROCHEMICAL AND C NHR DATA

FOR CO ADDUCTS OF MONOMERIC IRON(II) COMPLEXES, [Fe{(R)Me2[16]tetraeneN4>(B)(CO)](PFfi)2

R B vC0(so1id>a vco(8oln)b EPC

(He2NEthi)2 ch3cn 1961 1975 +0.035 223.1l,5-Pent(NHEthl)2a FT 1952 1961 40.310 220.1l,5-Pent(NHEthl)2e l-Kelm 1934 1950 ---- ----

1,6-Hex(NHEthi)2* FT 1955 1967 +0.3B5 ----

l,6-Hex(NHEthi)2e 1-Helm 1951 1955 ---- ----

1,6-Hex(HeNEthi)2 ch3cn 1964 1965 +0.245 222.71,6-Hex(HeNEthi)2 FT 1959 1967 +0.240 224.0l,6-Hex(MeNEthi)2 l-Melm 1955 1958 +0.140 223.81,6-Hex(HeNEthi)2 l-Kelm 1951f ---- ----

Ff. (NHEthi) 2 FT 1977 1971 40.245 ----

Ff. (NHEthi) 2 l-Melm 1970 1960 +0.145 ----

m-Xylyl(NHEthi)2 FT 1972 1988 +0.375 ----

m-Xylyl{NHEthi)2 l-Melm 1968 197B +0.300 ----

m-Xylyl(NHEthi)2 4-NH2-FT 1980 1978 +0.290 ----

n-Xyly1(NHE thi)2 Im 1967 1979 +0.290 218,4m-Xylyl(NHEthi)2 Im 1973f ---- ----

m-Xylyl(MeNEthi)2 CHjCN 1980 1993 +0.315 ----m-Xylyl(HeNEthi)2 PT -- 1989 +0.300 ----

m-Xylyl(HeNEthi)2 l-Melm 1971 1981 +0.245 ----

m-Xylyl(KeNEthl)2 Im 1975 1981 216.5m-Xylyl(MeNEthi)2 4-NH2-PT 1978 1981 ---- ----

p-Xylyl(NHEthi)2 ch3cn 1985 1996 +0.325 213.2p-Xylyl(NHEthi)2 FT 1980 1996 +0.355 ----

p-Xylyl(NHEthi)2 l-Melm 1970 1987 +0.240 ----

l,4But-(MeNEthi)2 l-Melm -- 1943 ----

aAs nujol mills, an-1.Âcetonitrile solution, cm”*.°V vs. Ag/Ag+ , Acetonitrile solution.^Chemical shift of CO carbon atom resonance, CD.CN solution, ppm

relative to IKS.eReference 88. ^1-Helm solution.

226

frequencies comparable to those of acetonitrile adducts; imidazole and

4-NH2”Py adducts are very similar to 1 - Melm adducts. In all cases,

VCQ for the pyridine derivatives is 9 or 10 cm-1 higher in energy than

for 1 - Melm adducts. This trans ligand effect compares very favorably

with the effect reported by Alben and Caughey^ for porphyrin systems.

The more basic imidazole type ligands enhance ¥ backbonding from the

metal to the CO. This results in a lower energy v for imidazole thanLUfor the less basic pyridine or acetonitrile adducts.

V^0 for the unbridged compound derived from the dimethylamine

ligand is 1975 cm ^ and is of approximately the same energy as reported

for porphyrin model systems.^

The bridged complexes are divided into two groups based on CO

stretching frequencies: those having aliphatic bridges and those

having aromatic bridges. Among the aliphatic bridged complexes, for a

given axial base the order for is

Fluorene > (NMe)2 (CH2)6 ~ (NH)2 (CH2>6 > (NH)2 (CH2)5 > (NMe)2 (CH2)4

This follows the same order as the effective bridge length from 7 car

bons to 4 carbons. As the bridge length is shortened and the cavity

size is reduced, decreases in energy by 17 cm ^ within this series.63A similar trend has been reported for protein systems in which a

decrease in V^0 is correlated with increased bending of the Fe-C-0

linkage due to increased steric interactions with the protein structure.

In a similar way, interactions between bound CO and the 7 carbon bridge

of the fluorene species are minimal, resulting in a value for \>c0 which

is essentially the same as for the unhindered, unbridged complex.

227

As shown by the several crystal structures, decreasing bridge length is

accompanied by decreasing cavity size. This is expected to lead to

increasing steric interactions between the bound CO and the bridge,

resulting in increasing distortion of the Fe-C-0 linkage and thus a

decreasing CO stretching frequency. It has clearly been shown above

that decreasing cavity size causes a decrease in the observed equili

brium constant. In this series of totally synthetic iron(II) model

complexes, then, a completely systematic change is observed in v r n whichuucorrelates directly with equilibrium constants and can be attributed to

the designed interactions between the bound CO and the ligand struc

ture.

It is interesting to note that v^0 for the (NH^CCI^)^ and

(NMe^CNl^)^ are essentially the same, demonstrating that the effect of

a bridge methyl group on this property of the complexes is minor rela

tive to the steric and trans ligand effects.

The data for the aromatic bridged complexes are not as readily

interpreted. Although steric factors are expected to be similar to, or

more restricting than for the aliphatic bridged complexes, all of the

CO stretching frequencies are between 1978 and 1996 cm \ The trans

ligand effect is the same as in the aliphatic complexes. An additional

factor must be present in the aromatic systems which causes such high

energy stretching frequencies. The close proximity of the tt systems of

the CO and xylene (structure XXVIII) offers a possible explanation. Some

type of interaction between these two pi systems apparently strengthens

the CO bond resulting in a higher energy CO stretch. Such an effect

combined with the bending of the Fe-C-0 linkage would greatly complicate

228

N F e N

XXV nr

the interpretation of infrared data. This difficulty is exemplified by

comparing the (NH^Mxyl and (NMe^Mxyl complexes. The "lid-off" nature

of the (NMe)Mxyl bridge is expected to cause strong interaction with the

bound CO resulting in considerable bending of the Fe-C-0 linkage and a

decrease in V ^ . The (NH^Mxyl complex is "lid-on" which should result

in less interaction with the bound CO as verified by equilibrium

studies and a higher CO stretching frequency. The putative tt system

interaction, however, causes v to be essentially the same for the twoLUspecies. The fact that the observed equilibrium constants for the

(NH^Pxyl and (NH^Mxyl are virtually identical yet their values

differ by 9 cm ^ is further evidence in support of this point.

Although there is no precedent for such IT system interactions causing

such effects on VCQ, this theory is consistent with the observed data.

Significant solvent effects on are observed for the

(NH^Mxyl and (NMeJ^CCl^)^ systems. Upon changing from acetonitrile to

l-Melm as solvent, decreases approximately 7 cm”^. This is

22962consistent with the observations of Maxwell and Caughey that more

polar solvents cause a decrease in V--.LUIt is apparent from the above data that a number of factors are

of importance in determining the CO stretching frequency. The degree of

bending of the Fe-C-0 linkage, the nature of the trans axial ligand, tt

systems in the vicinity of the bound CO, and the polarity of the solvent

system all contribute to the observed frequency to varying degrees.

Since the individual contributions of each factor to the observed

stretch has not been fully quantified, infrared data must be inter

preted with caution, particularly when comparisons between model sys

tems and proteins are being made. In particular, the tt systems of the

proteins and the polarity of the environment within the globin

structure are very difficult to simulate in the model system. It is

encouraging to note, however, that several of the dry cave model com

plexes have CO stretching frequencies comparable to 1953 cm ^ observed 63In the proteins,

Carbon-13 NMR Studies. The resonances of the CO carbon in the13C NMR spectra are included in table 34. The CO carbon atom resonance

in the unbridged complex having a linear Fe-C-0 linkage occurs at

223 .1 ppm. The resonance In the (NMe)„(CH„), species having the same2 2 oaxial ligand (CH^CN) occurs at 222*7 ppm. Further restriction of the

cavity size causes an upfield shift for the resonance to 220.1 ppm in

the (NH^CC^)^ species. The shift Is explained in terms of compres

sion effects at the CO carbon atom due to interaction with the bridging

group. A slight rehybridization of the CO carbon atom from pure sp

2302(structure XXIX) In a linearly bound molecule towards sp (structure

XXX) in a slightly bent Fe-C-0 linkage is expected

Sc

N Fe----- N n

B

13As with the infrared data, the C NMR spectra of the xylene

bridged CO adducts must be handled separately from the aliphatic com

plexes. The effects of bending on the shift of the CO carbon atom

resonance have been described above for the aliphatic species. Ring

current effects are observed in the atomatic complexes which result in

an increased shielding of the CO carbon nucleus, shifting the resonance

upfield relative to the non-aromatic systems. It is therefore apparent

that ring current and compression effects both cause shifts in the same

direction. The magnitude of the ring current effect depends entirely

on the spacial orientation of the CO and xylene ring and therefore the

compression and ring current contributions to the observed shift are

difficult to separate. From the information available, however, an

estimate can be made for the (NMe^Mxyl system. Equilibrium studies

have shown that this complex hinds CO only slightly more strongly than

•CI

■ F e -

/

•N

B

231

the (NH)2 (6112)5 species, thus implying similar compression effects in

the two structures. One can therefore estimate that of the 6.6 ppm dif

ference between the xylyl species and unbridged species, 3.0 ppm is due

to compression effects (as in the pentamethylene species) and 3.6 ppm is

due to ring current effects.

It is surprising and significant that the frequency of the CO

resonance is quite insensitive to the nature of the axial ligand. This

demonstrates that the compression and ring current effects are of pri

mary importance In determining the frequency of the resonance relative

to other electronic effects.

The observed resonances do not compare favorably with those of99HbCO and MbCO at 206-209 ppm, however the factors discussed above for

the IR data such as solvent and globin effects have not been studied and

may account for the observed differences.

Electrochemical Studies. The unusual electrochemical behavior

of the CO adducts is very informative. The cyclic voltamagram of

[Fe{(l,6-Hex(MeNEthi)2)Me2 [16]tetraeneN^}(CO)(PY)](FFg^ acetonitrile

is shown in figure 58. A voltamagram beginning at 0.0 V and scanning

negatively shows no reductions out to -1.0 V. Reversal of the scan

direction reveals an oxidation peak at +0.270 V. Reversal of the scan

direction results In new reduction waves at -0.175 V, -0.535 V, and

-0.875 V which are coupled to the oxidation. The multiple reduction

waves are due to species which exchanged axial base with solvent while

in the oxidized state. A new oxidation wave sometimes appears near

-0.2 V for some complexes. The apparent complexity of the voltamagrams

-12- 0.6 -0.B - 1.01 -0 .4V. vs Ag/Ag

0.4 0.2 0.0 0.20.6

Figure 58. Cyclic Vo ltama grains for [Fe{ (1,6-Hex(MeNEthi) )Me [16]tetraeneN.}(B)(CO)](FF,)_ 2 2A o 2.a) B ** PY, solvent = CH^CN, b) B = 1 - Melm, solvent = DMF

233

is explained in.the following way. The wave at +0.270 V corresponds to

the oxidation of the iron(II) CO adduct to the Iron(III) state with loss

of CO. The reduction waves are due to the CO-free iron(III) species

being reduced to an iron(II) state The reduced iron(II) complex can

then recombine with the dissolved CO or remain as a non-CO adduct giving

rise to the oxidation wave at -0.2 V. The described reactions are

shown in equation 40.

[Feli:(B)(C0) '*~e w [Fem (B)] + CO [FeI:CL(B)] (40)

Independent experiments have been conducted to prove the expla

nation just described for the (NMe) Mxyl complex. First the Iron(II)

chloro complex was placed in the electrochemical cell and CO was added.

The oxidation wave near -0.4 V disappeared and a new one due to the CO

adduct appeared at +0.315 V. After degassing the solution with nitro

gen, the original voltamagram returned. Identical behavior was

observed when the four-coordinate species was used as the starting

material. Finally, the CO adduct with CH^CN as the axial base was

electrolyzed and the resulting solution degassed. After reduction of

the complex, voltamagrams Identical to those of the iron(II)

four-coordinate species were obtained.

The same general behavior was observed for all of the com

plexes studied and the oxidation peak potentials are listed In table 34.

The only trend which is conclusively shown by these data is that imida

zole adducts oxidize most readily, followed by pyridine and acetonitrile.

This is consistent with all of the other data which have demonstrated

that imidazoles are the best Lewis bases, donating electron density to

234

the metal center. The inductive effect of bridge methyl groups is

again seen in the cases of Mxyl and (CHj)^ bridged complexes, wherein

the (NMe)^ derivatives oxidize at more negative potentials than the

(NH)2 derivatives. Unfortunately, no correlation between oxidation13potentials and XR or C NMR data is apparent.

The electrochemical behavior of the complex [Fe{(l,6-Hex-

(MeNEthi)2)Me2[16]tetraeneN^}(C0)(MIM) ](PF^Jg was also studied in DMF

as shown in figure 58. The cyclic voltamagram strongly resembles that

observed in acetonitrile except that the reduction behavior is much

simpler. The single reduction wave indicates that DMF does not compete

strongly for the metal coordination sites and the principle iron(lll)

species in solution has l-Melm as the fifth ligand.

Summary of Correlations for CO Adducts. Some correlations do

exist between the equilibrium and spectroscopic data. A change in the

trans axial ligand from acetonitrile to 1 - Melm causes an increase in

with a decrease in and Ep, all of which are consistent with the

good Lewis base characteristics of l-Melm. Secondly, for the aliphatic

bridged species, a decrease in corresponds to a decrease in the fre-CO

13quency of the CO carbon resonance in the C NMR spectrum and a

decrease in Vq q * The equilibrium data confirm the suggestion that the

spectroscopic properties of the aliphatic and aromatic bridged species

must be treated separately and that there is no overall correlation

between any of the spectroscopic properties and the observed K _.

235

Reactions of Iron(II) Dry Cave Complexes with Oxygen

The Importance of the reversible reactions between 02 and Hb and

Mb in living systems is obvious. It has already been described in the

Introduction that Iron(II) dry cave complexes have many of the features

required for reversible reactions with oxygen. Therefore, a number of

studies were conducted to examine the potential of these complexes as -

reversible oxygen carriers. Because of their importance as irreversible

oxidation products in porphyrin systems, a chloro-iron(III) complex and

a y-oxo-dimer were synthesized so that they could be recognized if they

formed during spectroscopic studies. The discussion of these complexes

will be followed by a description of the reactions of Iron(II) dry cave

complexes with oxygen.

Iron(III) Complexes. The iron(III) complex [Fe{(m-Xylyl-

(MeNEthi)2Me2 [16]tetraeneN^}Cl](PFg)2 was prepared by the addition of a

slight excess of Ce(IV) to a solution of the iron(II) complex. The pro

duct was deep blue having the composition [FeLCl](PFg)2 and was very

difficult to purify. Alternatively, nitrosyl hexafluorophosphate was

used as the oxidizing agent yielding an Identical iron(III) product.

The Infrared spectrum is similar to that of the comparable iron(II) com--3plex. The molar conductance of a 1 x 10 molar acetonitrile solution

-1 -1 2of the complex was 287.8 ohm mole cm , within the acceptable range for89a 2:1 electrolyte. The solid state magnetic moment, determined by the

Faraday method, was 6.05 8 , consistent with hlgh-spin iron(III). (The

spin only value is 5.92 8 .) The electrochemical behavior was virtually

identical to that of the sample prepared by controlled potential

236

electrolysis as described earlier. Similar coulometric studies on this

complex yielded an n value of 0.96 electrons with the behavior of the

reduced species being identical to that of the starting iron(II) chloro

complex. All of the above data clearly demonstrate that upon oxidation,

the chloride ion remains coordinated to the iron(III) center. The oxi

dation and reduction of the iron center are completely reversible pro

cesses.

The li-oxo-dimer of the (NMe^Mxyl bridged species was prepared

in a number of ways. The most direct synthetic method involved the

exposure of a solution of the iron(II) chloro complex to air in the

presence of an alcohol or water. Analytical data for the brown product

confirmed the absence of chloride ion in the complex. Materials having

identical properties were prepared by the addition of water to an ace

tonitrile solution of the Iron(III) chloro species and also by oxida

tion of the iron(II) chloro species with Ce(IV) in the presence of

water. Chromatography of the iron(III) chloro complex on neutral Woelm

alumina resulted in a color change from blue to brown on the column with

eventual Isolation of the brown y-oxo-dimer.

One of the characteristic reactions of oxo-bridged porphyrin

dimers is cleavage by HC1 as shown in equation 41. Addition of base

reverses this reaction.

PFe-0-FeP + 2 HC1 * 2 FePCl + H20 . (41)

As shown In figure 59, the iron(III) chloro complex has absorption max

ima near 820 nm and 615 nm, whereas the oxo-dimer has peaks at 760 nm

and 502 nm in the visible spectrum. The spectrum of the oxo-dimer was

0.8-

0,7-

0.6-

0.5-

0A-

0.3-

0 . 1-

"1 1----------------------------- 1------------------------------- 1---- T -5 0 0 6 0 0 7 0 0 8 0 0 9 0 0

Figure 59. Electronic Spectra of the Chloro-iron (III) and p-oxo-ditner Derivatives of the (NMe^Mxyl Bridged Species

237

238

generated either by dissolution of a genuine sample, by addition of

water and triethylaraine to a solution of the iron(III) chloro complex,

or by addition of water and triethylamine to a solution of the Iron(II)

chloro species followed by exposure to 0^.

The iron(III) chloro complex was generated In solution from the

y-oxo-dimer by addition of HC1 as confirmed by the electronic spectra.

Although the reaction was complicated by ligand removal at high acid

concentrations, the point is clear that the iron(III) chloro and

y-oxo-dimer complexes interconvert in a manner analogous to the simple

hemes. The ready formation of the undesirable irreversible products

was a discouraging result since it demonstrated a limitation Inherent in

the dry cave model. Despite the presence of the bridge, oxo-dimer

formation occurs through the one unprotected axial site (structure XXXI).

/

A modified dry cave model which Incorporates steric bulk or a "tail

base" in the vicinity of the unprotected site would thus be highly

desirable.

239

Oxygen Uptake Experiments. In spite of the results described

above, a number of experiments were performed in an attempt to deter

mine the conditions required for reversible oxygenation. Stoichio-77metric oxygen uptake experiments performed by Dr. D. J. Olszanski in

pyridine on the iron(II) (NlD^Mxyl chloro complex at room temperature

showed that uptake asymptotically approached 4.0 moles of 0£ per mole

of iron(II) after two days of reaction. This indicates a reaction much

more complex than simple oxidation or oxo-brdlged dimer formation since

these processes require only 0.25 moles of oxygen per mole of iron(II).

At -36°C, however, 1.10 moles of 0^ were taken up per mole of lron(II),

consistent with monomeric adduct formation. The reaction which con

sumes most of the oxygen is apparently inhibited at reduced tempera

tures. Attempts to remove the oxygen after reaction at -36°C using

freeze-pump-thaw techniques resulted in no color changes indicative of

reversal; however, this may not be an adequate criterion on which to

base a conclusion concerning the reversibility of the reaction.

Spectral Studies. The reaction of the complex derived from the

(NMe^CC^)^ bridged ligand with oxygen was examined at -30°C in ace

tonitrile solution containing 1.5 molar l-Melm. Addition of various

partial pressures of oxygen caused spectral changes which were not

isosbestlc (figure 60) but were somewhat dependent upon the oxygen par

tial pressure. Nitrogen was passed through the solution in an attempt

to remove the oxygen but there was no spectral change. The lack of

isosbestlc behavior and Irreversibility of the reaction indicate that a

process other than simple oxygen adduct formation occurred.

240

0.7-

0.6-

0 .5 -

0 .4 -

0 .3 -

0.2-

0.1.-

3 5 0 4 0 0 4 5 0 5 0 0

Figure 60. Spectral Changes for Reaction of [Fe{(l,4-But(MeNEthi) ^ ) ~

Me2[16]tetraeneN^}]^ with O2 in CH^CN, 1.5 M l-Melm, -30°C

241

The reaction of [Fe{(m-Xylyl(MeNEthi)2)^62[16]tetraeneN^}]^

with oxygen in aqueous solution containing 1.5 molar l-Melm (enough to

saturate the axial base equilibrium) was examined at 10°C. At very low

partial pressures of oxygen, the spectral changes were nearly isosbes-

tic. Upon deoxygenation of the solution with nitrogen, however, the

spectrum continued to change in the same direction rather than

returning to that of the original unoxygenated material. This indi

cates that the complex reacted with even trace amounts of oxygen in an

irreversible manner. The isosbestic spectral behavior suggests a simple

process occurred under these conditions with irreversible oxygen adduct

formation as one possibility.

The reaction of 0^ with the (NH^Mxyl and (NMe^Mxyl chloro

complexes was studied spectrally at room temperature. In acetonitrile

solution, the iron(III) chloro complex was formed in the absence of base

and the y-oxo-dimer resulted when base and water or alcohol were

present. In methanolic solution, the principle product was the

y-oxo-dimer. These products were identified by comparison of their

electronic spectra of those of genuine samples.

In pyridine, the reaction of the (NH^Mxyl chloro complex with

oxygen was quite slow, demonstrating non-isosbestic spectral behavior.

This complexity was expected based on the results of the uptake experi

ment under comparable conditions as described above.

None of the reactions showed any reversible character as was

expected based on the nature of the known products. It Is also pos

sible that oxygen adduct formation does occur and that the bound oxy

gen reacts further. Hydrogen atom abstraction by the bound dioxygen

242

followed by a series of radical reactions can ultimately result in for

mation of oxygenated species. Meta-xylene bridges are expected to be

particularly susceptible to such reactions because of the relative ease

of hydrogen atom abstraction and stability of the radical formed. Such

decomposition reactions have been postulated for related cobalt(II)

dry cave complexes.^

Iron(II) Complexes Having Rearranged Ligands

It was mentioned earlier (page 163) that during the synthesis of

the iron(II) complex of the unbridged methylamlne derivative, a rear

ranged product formed. The principle evidence for such a process was

the set of doublets observed in the proton NMR spectrum at 5.6 and

8.7 ppm. Similar patterns have been observed by Riley^^ for

related compounds which contained a rearranged macrocycle and have been

assigned to coupling between hydrogen atoms on the 6 and y carbon atoms

of the macrocycle. A similar rearrangement has occurred In this com

plex, in which a hydrogen atom shifted from the external nitrogen atom

to the y carbon atom with subsequent rehybridization and geometrical

rearrangement to yield the sexadentate ligand in structure XXXII.

The sexadentate species was synthesized directly and in good

yield by the procedure described above except the solution was refluxed

for 45 minutes in the presence of excess triethylamine. The analytical

data confirmed the stoichiometry of the complex. The infrared spectrum

(figure 61) contains no absorptions due to N-H stretching and NMR

spectrum contains the same coupled resonances as observed above at

243

'X X XTT

i500

I1000

I1500

I20004000 3000

Figure 61. Infrared Spectrum of the Iron(II) Complex of the Sexadentate Ligand

244

5.6 and 8.7 ppm (figure 62). The splittings are apparently more complex

than simple doublets due to the lack of any elements of symmetry in the

molecule. This is reinforced by the complex resonance pattern observed

in the methyl region of the spectrum.13The C NMR spectrum of the sexadentate complex is shown in

figure 63a and the resonances are listed in table 35. The spectrum is

in support of the proposed structure with the lack of symmetry apparent

from the uniqueness of each carbon atom. The most striking feature of

the spectrum is the absence of resonances between 105 and 115 ppm and

the appearance of resonances near 65 ppm which split into doublets in

the off-resonance spectrum. These resonances are assigned to the y

carbon atoms of the macrocycle. It is not possible to determine

whether the complex is a cis or trans isomer based on spectral data

because neither isomer contains symmetry elements or any distinguishing

features. The striking similarity between the spectra of this complex

and those of the related clathrochelates which are constrained to be88cis complexes as described by Grzybowski strongly suggests that this

complex is a cis isomer.

Electrochemical studies performed on the sexadentate complex

showed only a single process over the entire accessible potential range

which was a reversible one-electron oxidation having ^ / 2 an<*

JE3/4 — Ei/4 I values of +0.390 V and 65 mV versus Ag/Ag+ . These values88compare favorably with the values obtained for the clathrochelates and

is approximately 90 mV more positive than the values for related hexaene. 103compounds.

A yv/P1 l r t — 1— rrto

Figure 62. Proton. NMR Spectrum of the Iron(II) Complex of the Sexadentate Ligand 245

w

240

Figure 63.

200 160ppm

13C NMR Spectra of a) the Iron(II) Complex of the Sexadentate Ligand and b) the CO Adduct of the Pentadentate Complex

246

247

TABLE 35

CARBON-13 NMR DATA FOR IRON(II) COMPLEXES OF REARRANGED LIGANDS

Compound Chemical Shifts

[FeiCMelmEtJ^MeÊlSltetraeneN^.}] (PFg)2 171.8, 170.6, 170.1, 65.6, 65.2,56.3, 55.4, 49.9, 49.3, 44.9,44.6, 43.2, 27.3, 27.1, 25.0,24.5, 24.3, 23.2

[Fe{(MeNHEthi)(MeImEt)Me2 [16]- 219.2, 172.9, 175.9, 173.5,tetraeneN.}(CO)](PF.)„ 170.7, 169.5, 166.6, 163.3,

H O Z108.5, 65.4, 57.7, 56.8, 54.3,52.7, 49.9, 49.4, 47.1, 46.7, 43.2, 38.8, 31.9, 27.6, 25.6,24.5, 23.4, 21.7, 20.8, 16.1

248

The unbridged complex derived from the methylamine ligand under

goes a very novel reaction with CO in acetonitrile solution in the

presence of Py, l-Melm or no base to yield the same product which is

represented by structure XXXIII.

CH

xxxmThe infrared spectrum of the CO adduct of the rearranged isomer

of the unbridged complex is shown in figure 64. The sharp N-H stretch

at 3410 cm ^ is due to the free external amine and the several absorp

tions between 1600 and 1700 cm ^ are typical of the rearranged imine

structure. The sharp, intense absorption at 1989 cm ^ is due to the-1 13bound CO with the band at 1946 cm attributed to CO.

13The C NMR spectrum (figure 63b, table 35) of the rearranged

pentaene CO adduct has essentially one resonance for every carbon atom,

4000I i I I

3000 2000 1500 1000cm-'i

500

Figure 64. IR Spectrum of the Pentadehtate CO Adduct

250

consistent with the lack of molecular symmetry. The resonances due to

each half of the molecule are readily assigned by comparison with the

spectra of the unrearranged nickel(II) complex and the sexadentate

iron(II) complex.88Grzybowski has shown that in the presence of base at ele

vated temperatures, some of the iron(II) dry cave complexes having

hydrogen as the bridge nitrogen substituent rearrange to yield sexaden

tate clathrochelate ligands. When the bridging group is meta-xylene,

the rigidity of the xylene ring prevents formation of a sexadentate

ligand but permits a pentadentate ligand to form (structure XXXIV).

Such a complex has been synthesized and characterized. The elemental

analysis was consistent with a mixture of this product and the chloro

type starting material. The infrared spectrum (figure 65) differs

251

iI4000 3000 2000 1500 1000 500cm-1

Figure 65. Infrared Spectrum of the Pentadentate meta-xylyl Complex

from that of the chloro complex, particularly in its very sharp N-H

stretches at 3400 and 3415 cm ^ and the additional bands at 1620 and“X 881630 cm which have been attributed to a rearranged ligand structure.

The oxidation potential at -0.250 V is considerably more positive than

that of the chloro complex. It is very likely that equilibrium reac

tions involving the (NH^Mxyl bridged complexes may be quite slow due to

the structural rearrangements which are possible for this ligand.

Iron(II) Complexes Derived from Dimeric Dry Cave Ligands

Synthesis and Characterization. The iron(II) complexes derived

from dimeric dry cave ligands were in general much more difficult to

synthesize and purify than those from monomeric ligands. All of the

252

work described in this section was performed prior to the solution of

the crystal structure of the dimeric nickel(II) complex. Many of the

results were initially difficult to interpret because of the earlier

assumption that these materials were actually monomers.

The dimeric iron(II) complex which was studied in greatest

depth by this author contained the (NH^Mxyl linkages as shown in the

crystal structure described earlier. Other workers have studied dimers88 Q9 Q9prepared from the (NH)2 (CH2)3 , (NH)2 (CH2)4 , NH2(CH2)3 and

78(NH)2Fluorene ligands. The properties of all of the complexes were

very similar in all respects to those which will be described for the

(NH)2Mxyl example.

The complexes were synthesized through the reaction of the

tetrachlorozincate salt of the ligand with an excess of

tetrakis-(pyridine)iron(II) chloride and six equivalents of triethyla-

mine. Upon mixing of the reagents the solution color became deep red

and a red-orange precipitate formed. This product was collected but

was never satisfactorily characterized. It is believed to be a pyri

dine adduct of the iron(II) complex with a mixture of chloride and

tetrachlorozincate anions. Analytically pure complexes were obtained

by dissolving the initial product in methanol, adding an excess of pyri

dine, imidazole, 1 - Melm, or 2-MeIm, then slowly adding an excess of

ammonium hexafluorophosphate in methanol. The resulting red products

were often of variable composition but careful control of conditions

yielded pure samples.

The Infrared spectra are very similar for all of the complexes

with an N-H stretch at 3420 cm ^ and a broad absorption at 1580 cm \

253

The IR spectra of the pyridine and imidazole adducts (shown in figures

66a and b) are representative of the series. Absorptions due to the

axial ligand are typical of those described earlier for the CO adducts

of the monomeric iron(II) complexes. The l-Melm adduct has a band at

3140 cm ^ due to the C-H stretch at the carbon between the imidazole

nitrogens. This band is absent in the 2 - Melm adduct, but a broad N-H

stretch at 3390 overlaps with the N-H stretching band of the macro

cycle. Bands at 3420 and 3150 are present in the imidazole deviation.h

Absorptions due to the bound pyridine are obscured by ligand bands.

It was not possible to obtain solid state magnetic moments for

the complexes, using existing equipment, because of their extreme air

sensitivity. ^H and NMR spectra were measured for all of the above

complexes in CD^CN in the presence or absence of excess axial ligand.

In all cases, the resolution was very poor but the spectra did indicate

the presence of a diamagnetic species in solution. It is unclear

whether the poor resolution was due to air oxidation during the experi

ment or to broadening caused by rapid exchange between paramagnetic and

diamagnetic species.-3Molar conductances were measured In acetonitrile for 1 x 10

molar solutions of two of the complexes. The values obtained with Py-1 -1 2and l-Melm as the axial ligands were 214.2 and 221.2 ohm mole cm ,

respectively. These values are at the lower end of the acceptable89 —1 —1 2range for 2:1 electrolytes in acetonitrile, 220-300 ohm mole cm .

A number of electrochemical experiments were performed In order

to learn more about the equilibria in solution between the iron(II) com

plex and the axial ligands. The complex having 1 - Melm as the axial

254

I4000

\r

3000I I

2000 1500cm-1

i •1000 500

Figure 66. Infrared Spectra of [Fe{(ra-Xylyl(NHEthi)2)Me2 [16]- tetraeneN4 }(B)]2 (PF6)4 , a) B = PY, b) B = Im

2552+ 3+ligand had two oxidations which were attributed to the Fe /Fe couple.

The first had values of E-jy2» |E3/^ ~ ei/4 I» and EP of "°*350 v » 33 mV, and -0.300 V, respectively. The magnitude of the second oxidation wave

depended on the sample used and had a peak potential of +0.030 V. Upon

addition of 1 drop of l- M e l m to the cell, Ep shifted to -0.385 V. Fur

ther addition of l-Melm shifted Ep to -0.425 V. This large dependence

upon the concentration of axial ligand indicates the involvement of the

redox active species in equilibria in solution. The dimeric nature of

the complex makes the potential equilibria very complex since one or

two molecules of 1 - Melm can interact with each metal center. Fur

ther complexity arises from the possibility that the two centers do not

act totally independently of one another, causing the equilibria at one

center to be affected by the coordination status of the second center.

Unfortunately, the exact nature of the equilibria cannot be determined

from electrochemical studies.

The 1 - Melm adduct was also studied by controlled potential

electrolysis in the presence of excess l-Melm. The first oxidation at

+0.30 V versus Ag/Ag+ yielded an n value of 0.82 electrons. Reduction

of the same solution yielded an n value of 0.67 electrons. Another

cycle of oxidation and reduction of the same solution resulted in n

values of 0.58 and 0.60 electrons. It is apparent from these data that

the oxidation and reduction processes involve more than simple revers

ible metal oxidations. It is possible that the oxidized complex reacts

irreversibly with l- M e l m giving rise to a reduction wave at -1.16 V,

but this has not been proven.

256

Similar electrochemical results were obtained when Im was the

axial ligand. In the absence of excess axial ligand, the peak potential

of the first oxidation was -0.240 V. Upon addition of excess Im, this

potential shifted to -0.420. Controlled potential electrolysis was per

formed and showed that in the presence or absence of excess imidazole,

the same average n value of 0.82 electrons was obtained for oxidations

and reductions. There was no indication of a further reaction of the

oxidized product with the excess base, contrary to the results observed

with 1 - Melm.

In acetonitrile it was found that when Py was the base the peak

potential of the first oxidation was much less affected by the pyridine

concentration. The potential shifts from -0.08 V in the absence of

excess pyridine to -0.130 V in the presence of a large excess of pyri

dine. This would suggest that pyridine does not compete very strongly

with acetonitrile for coordination to the iron(II) center as has been

shown for the monomeric complexes. Experiments were also carried out

in DMF in which the complex was carefully titrated with pyridine. The

changes observed are shown in figure 67. In the absence of excess pyri

dine, the peak potential was -0.355 V versus Ag/Ag+ . Addition of one

equivalent of Py shifted the potential to -0.335 V. Further pyridine

additions resulted in a limiting value for Ep of -0.240 V when the

ratio of Py to iron(II) was -60:1. These data indicate that under care

fully controlled conditions some useful equilibrium data may be obtain

able.

The interaction between the iron(II) complexes and axial ligands

was also studied spectrally. The titration of the pyridine adduct with

- 1.0-0.6 - 0.8-0.4- 0.20.2 0.0V vs A g / A g -*"

Figure 67. Cyclic Voltamagrams of [Fe{(m-Xylyl(NHEthi)2)Me2 [l6] tetraeneN^}(PY)]^ (PFg)^ in DMF with excess PY

258

pyridine in acetonitrile resulted in a regular spectral change as shown

in figure 68. The band maximum shifts from 446 nm to 460 nm and the

intensity increases upon addition of Py. Although the spectra were

complex in the region hear 400 nm, the regularity of the spectral

changes suggests a simple equilibrium process. The fact that even in

the presence of 900 equivalents of base the spectrum is still changing

significantly suggests a small value for the observed equilibrium con

stant. The simplicity of the spectrum suggests the observed equili

brium is that described by equations 42 or 43

[FeL(CH3CN)2]2+ + 2Py ^ K * [FeL(CH3CN) (PY) + 2CH3CN (42)

[FeL(CH3CN)(PY)]2+ + 2Py [FeL(Py)2]2+ + 2CH3CN (43)

The above two possibilities are expected to show the same type of spec

tral behavior, thus neither can be ruled out.

In order to eliminate the possibility that the observed spectral

changes were due to ligand deprotonatlon by the pyridine, the same

experiment was performed using the non-coordinating base 2,6-Lutidine.

The spectral changes were negligible relative to those observed with

pyridine thus ruling out the deprotonation possibility.

In order to assess the role of solvent in these equilibrium

studies, the pyridine adduct was allowed to react with excess pyridine

in acetone. The observed spectral changes were much larger than those

observed in acetonitrile and the equilibrium appeared to be saturated in

the presence of a 900-fold excess of pyridine. This result verifies the

suggestion that pyridine competes with the solvent for a binding site on

259

$n

$

o'

■uveT

*1•c

.ft/

260

the iron(II) center. Since acetone is a much weaker ligand than ace

tonitrile, the pyridine competes more effectively, yielding the

observed results. This result is also consistent with the equilibrium

described above in equations 42 and 43.

adduct as shown in figure 69. The spectral changes are very different

from those observed with pyridine, with the maximum shifting from 453 nm

to about 485 nm with a decrease in intensity upon addition of Im. At

low concentrations of axial base, the spectral changes are

non-isosbestic. When the imidazole is present in greater than 100 fold

excess, however, an isosbestic point appears at 482 nm. These results

strongly suggest the occurrance of two equilibria whose constants are

such that the spectral changes due to each overlap at low concentra

tions of base. The first equilibrium appears to be saturated at an_2imidazole concentration of 1 x 10 molar, resulting in the observed

isosbestic behavior due to the second equilibrium at higher base con

centrations.

the observed behavior. The first possibility requires that the metal

centers act independently of each other. A stepwise substitution of

acetonitrile by imidazole is then observed, with the imidazole entering

the dry cave in the second step as shown in equations 44 and 45.

A similar equilibrium study was undertaken for the imidazole

There are two different two-step processes which can account for

[FeL(CH3CN)2]2+ + 21m K [FeL(CH3CN) (Xm) ]2

[FeL(CH3CN)(Im)32+ + 21m [FeL(Im>2]2+ (45)

(44)

Figure 69. Titration of [Fe{(m-Xylyl(NHEthi)_)Me„[16]tetraeneN#}(Im)]„(PF ) with Im in CH_CNL 2. h Z u h J 261

262

This explanation is consistent with the observed data since the second

molecule of imidazole is expected to encounter some steric interac

tions with the ligand structure of the cavity.

The second possible explanation requires that the metal centers

not act independently. In this process, the first step is substitution

of one acetonitrile in the cavity by an imidazole. The second step is

the replacement of the remaining acetonitrile at the second site by

imidazole which sterically interacts with the first molecule of Im

resulting in a smaller value for K2 than for K^. This is shown in

equations 46 and 47.

[(FeL)2 (CH3CN)2(Im)2]4+ + Im [(FeL)2 (CH3CN) (Im)3]4+ (46)

[(FeL)2 (CH3CN)(Im)3]4+ + Im [ (FeL)2 (Im)4]A+ (47)

It Is not possible to distinguish between the two possibilities based

on the available data.

Other investigators have shown that bis-base adducts do not

readily form when the hindered ligand 2 -Helm is used instead of Im.

Because of this a spectral titration was carried out in acetonitrile

with 2 - Helm as the axial ligand (figure 70). Upon addition of excess

base the maximum of 450 nm shifts to 445 nm and the intensity decreases.

Sharp Isosbestic points are observed, indicating a simple equilibrium

in solution which Is nearly saturated at a 2 - Helm concentration of _28 x 10 molar. The conclusion drawn from this result is that the

observed equilibrium Is that given by equation 44 and that the process

described by equations 44 and 45 is applicable when an unhindered

0 .5 -

0 .4 -

0 .3 -

0.2-

0.1-

4 0 0 4 5 0 5 0 0 6 0 0“ I—5 5 0

Figure 70. Titration of [FefOn-XylylCNHEthi^jF^tlSJtetraeneN^} ^ - M e l n O J ^ P F g ^ with 2-MeIm 263

264

imidazole ligand is used. This necessarily requires that the rela

tively large imidazole ligands can fit into the cavity in dimeric

structures and that the iron(II) sites resemble those of the porphyrins

in their coordination behavior.

Reactions with Oxygen and Carbon Monoxide. Before the dimeric

nature of these complexes was determined, a number of studies involving

reactions with oxygen were undertaken. Schammel had shown that in

SOXHLjO/SOXl-Melm, the particular complex with which he was working

apparently reacted in part reversibly with oxygen. This result was

duplicated by this author, verifying that Schammel was indeed working

with the same complex. Further studies in the 50%1-Melm/50%H20 sol

vent system were performed and yielded some interesting results as

shown in figure 71. Upon exposure to 0^, the band at 490 nm disappears

and new bands at 760 nm and 370 nm appear. Application of vacuum to

the system results in reappearance of a band near 490 nm with approxi

mately 1/3 of the intensity. The absorbance at 370 nm is less than in

either of the two previous spectra. After 12 hours under a static

vacuum, the absorbance at 490 nm is essentially unchanged but the

absorbance at 760 nm and 370 nm has decreased. Reexposure to 0^ results

in an Increase at 760 nm and 370 nm and a decrease at 490 nm. The opti

cal density throughout the spectrum is less than that resulting from

the first oxygenation. From these results, it is apparent that at

least two processes are occurring in solution. The first may be oxygen

adduct formation which results in the apparent reversibility. The

second reaction is irreversible and is responsible for the overall

0.0-

0.7-

0.6-

0.S-

0.4-

0.3-

0.2-

0.1-

800400Figure 71. Spectral Changes of [Fe{(m-Xylyl(NHEthi) jMeJieitetraeneN.JU-MeliiOj^PF,.). with 0„ in

50%H20/50%1-Melm . 2 6 4 2

500 700600[Fe{(m-tralSp }(1Changes ■Xylyl(NHEthi)of )Me [16] Helm) (PF )ec tetraeneN with 0,2 2 2 460/50%1 265

266

decrease in absorbance. When a solution is allowed to remain exposed

to O2 for extended periods (e.g., overnight), the initially formed

bands at 760 and 370 nm slowly decrease in intensity due to the slow

irreversible process.

The interaction of the dimeric iron(II) complex with CO at

ambient temperatures in a 50%H20/50%1-Melm solvent system was also

studied. As shown in figure 72, the reaction was essentially revers

ible. Upon exposure to CO, the band at 495 nm disappears and a

shoulder grows in at 380 nm. Very high vacuum and extended periods of

time were required to cause any spectral changes indicative of loss of

CO. However, upon exposure of the solution to sunlight, the band at

495 nm returned quickly. It thus appears that the loss of CO is photo

catalyzed by UV light and that the observed reaction is simple revers

ible CO adduct formation.

This result is particularly important because it suggests that51 52the flash photolysis techniques of Gibson and Traylor can be applied

to this system to study the kinetics and equilibria of reactions of

these iron(II) complexes with oxygen and carbon monoxide.

The equilibrium constant for CO adduct formation is apparently

very large for the dimeric species. A study using the flow system

(described above for the monomeric complexes) showed that the equili

brium was saturated even at the lowest accessible partial pressure of

CO. This is not surprising since the CO is not expected to be steri-

cally hindered by the ligand structure and linear binding should result.

Solid CO adducts of several of the dimeric iron(II) complexes

have been synthesized and the CO stretching frequencies have been

0.8-

0.7-

oe-

0 5 -

0 .4 -

0.3-

0.2-

0.1-

L4 0 0 500 600

Figure 72. Spectral Changes of [Fe{(m-Xylyl(NHEthi)o)Me0[16]tetraeneN.}(1 - Melm)(PF,). with CO in50%H20/50Z1-Melm 4 2 6 4 267

268

measured in acetonitrile. With pyridine as the axial ligand, vrn occurswU—1 flfiat 1973 cm for the (NH),,Mxyl, (NH) Fluorene and (NMe) Dury 1 com-

••1 —1 QOplexes, and at 1970 cm and 1967 cm for the ( N H ^ C C I ^ g and 88(NH)jCCI^)t; complexes. It is apparent that v ^q Is quite insensitive

to the nature of the bridging groups linking the two macrocycles. This

is expected due to the fact that the linking groups are essentially

parallel to the Fe-C-0 linkage and no Interaction between the CO and

bridge is expected. It is reassuring to note that the frequencies

above are very similar to that observed for the unbridged monomeric

compound as an acetonitrile derivative (1975 cm which is also

expected to have a linear Fe-C-0 linkage. It is also of interest to

note that the xylene groups demonstrate no unusual effects on vrn inLUthe dimers, In contrast to the observations for the monomeric complexes,

and thus supporting the conclusion of unusual pi-pi interactions in the

latter.

Further studies of the dimeric complexes were not pursued

because it was apparent with the solution of the dimer crystal struc

ture that they were not the desired model complexes.

Monomeric Copper(II) Dry Cave Complexes94 94Dry cave^ligands had been chelated to Mn(II) and Mn(III),

' • ' fifi fifi RRFe(II) and Fe(III) , Co(II) and Co(III), Ni(II), and Zn(II) but

the copper complexes had not yet been prepared. Because Cu(I) com-67plexes have been known to interact with and CO, it was of interest

to examine the chemistry of copper dry cave complexes. Copper(II)

complexes were prepared by the reaction of the appropriate ligand salt

with one equivalent of CuCOAc^'lÔ and three equivalents of NaOAc'SIÔ

in methanol. The highly crystalline copper(II) complexes were obtained

in satisfactory yields. Analytical and infrared data confirmed the

stoichiometry of the species as four-coordinate copper(II) dry cave com

plexes. The electrochemical characterizations yielded the most useful

information. In acetonitrile solution, a one electron reversible oxida

tion was observed with an<* lE3/4- E l/4^ values of +.580 V, 65 mV

and +.640 V, 65 Mv versus Ag/Ag+ for the (NMe^CCI^)^ and (NMe^Mxyl

bridged complexes, respectively. The cyclic voltamagram for the

(NMe^CCl^g complex is shown in figure 11. An ill-defined reduction

wave occurs at -1.350 V and is irreversible. The return oxidation wave

contains a peak at -0.570 which is due to the stripping of copper metal

from the platinum disk electrode. It is concluded from this result

that the reduction is a two-electron process yielding metallic copper.

Thus the copper(I) oxidation state is not stable for this class of com

pounds, but investigation of the Cu(III) state may prove fruitful.

APPENDIX A

Rotameter Calibration

270

TUBE

SC

ALE

REA

DIN

G-

MIL

LIM

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g g j a Q I f o s s o n . nFLOWMETER CALIBRATION- S c a le R ead ing v * . Flow R ata C a lib r a t io n w ith Both F lo a t s S im u lta n e o u s ly

FLOW RATE /M IN. AT 760 mm. Hg.

./<£>&21*C .

O

TUBE N O .r ^ g j SER.

.£>00

Sisii in in

m m:i itti

iiilrFLOAT

/ 0 FLOW RATE

£0 35" 3 0 W/M IN. AT 760 mm. Hg. A Z l'C .

271

83

84

APPENDIX B

Final X-ray Positional and Thermal Parameters

and Structure Factors

272

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aa*

3443

APPENDIX C

Derivation of Chloride Equilibrium Constant Expression

306

The observed equilibrium constant is defined as:

= [FeL(CO)]OBS [FeL]pCO

where

[FeL] - [FeL(CH3CN)]2+ + [FeL(Cl)]+

[FeL(CO)] = [FeL(CH3CN)(CO)]2+ + [FeL(Cl)(CO)]+

The individual equilibrium constants are defined as:

K1[FeLCCHgCN)] + CO [FeL(CH3CN) (CO) ]

[FeL(CH3CN)(C0)]*4. = [FeL(CH3CN)]pC0

[FeL(CH3CN) (CO) ] = ^ [ F e K C ^ C N ) ] pCO

K2[FeL(Cl) ] + CO ^ [FeL(Cl) (CO) ]

= [FeL(Cl)(CO)]2 [FeL(Cl)]pCO

[FeL(Cl)(CO)] = K2 [FeL(Cl)]pCO

- K 3[FeL(CH3CN)] + Cl FeL (Cl)

K- 0 [FeL(Cl) 3 3 [FeL(CH3CN)][Cl"]

[FeL(Cl)] = K3 [FeL(CH3CN)][Cl"]

Rearrange equation (48)

KOBS[FeL ]pC ° = t F e L <C0>]

307

(48)

(49)

(50)

(51)

(52)

(53)

(54)

(55)

(56)

(57)

(58)

(59)

(60)

308

and substitute for [FeL(CO)]

K0BS[FeLlpC0 = K!tFeL(CH3CN)]pC0 + K2 [FeL(Cl)(CO)]pCO

Cancel pCO

KQBS[FeL] *= K^[FeL(CH3CN) ] + K2 [FeL(Cl)]

K0BS[FeL] = ^ e L - F e L (Cl )] + K2 [FeL(Cl)]

substitute for [FeL(CH3CN)] in equation 58

[FeL(Cl)]3 [FeL- FeL (Cl)] [Cl“ ]

[FeLCl] = K3 [FeL-FeL(Cl)][Cl"]

[FeL(Cl)](l+K3)[Cl]) = K3 [FeL][cr]

K [FeL][Cl"][FeL(cl) 3 - l+iqiciT"

Substitute for [FeL(Cl)] in equation 63

1CK [FeL[Cl"] K K [FeL][Cl"] K0BS[FeL] " Kx[FeLl 1 + K3 [C1] + 1 + K3 [C1]

Cancelling [FeL] yields:

K.K.tCl"] K K [Cl"]K „ „ = K, + , . „ r nt 1 +OBS 1 1 + K3 [C1] 1 + K3 [C1]

and collect terms to obtain

K„(K - K )[Cl"]K = K, +OBS 1 1 + K-[Cl“ ]

(61)

(62)

(63)

(64)

(65)

(66)

(67)

(68)

(69)

(70)

309

When and are known, this rearranges to a linear form.

l°SS I ^ - K-IOl-]2 " OBS 3

from which is determined by the slope.

By propagation of errors, the standard deviation for each

y-value was determined from:

2 = 2 / 1 (KOBS " V ^K0 B S V K2 “ K0BS (K2 - K obs)2

+ °*1 ( K2 -^ O B s )

2 / ~^K0BS " Kl^

K2 V K2 “ K0BS>2j

(71)

(72)

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91. S. C. Jackels, W. A. Schammel, R. W. Callahan, J. C. Stevens,J. J. Grzybowski, D. J. Olszanski, L. L. Zimmer, and D. H. Busch, manuscript in preparation.

92. S. C. Jackels and I). H. Busch, unpublished results.

93. R. D. Feltham and R. G. Hayter, J. Chem. Soc., A, 4587 (1964).

94. N. Herron and D. H. Busch, unpublished results.

95. D. A. Baldwin, R. M. Pfeiffer, D. W. Reichgott, and N. J. Rose,J. Am. Chem. Soc., 95, 5152 (1973).

96. V. L. Goedkin, P. H. Merrell, and D. H. Busch, J. Am. Chem. Soc.,94, 3397 (1972).

97. T. Takano, J. Mol. Biol., 110, 659 (1977).

98. M. Brunori, J. Bonaventura, C. Bonaventura, E. Antonlni, andJ. Wyman, Proc. Nat. Acad. Sci. U.S.A., 69, 868 (1972).

99. M. H. Chisholm and S. Godleski, Progress in Inorganic Chemistry,Vol. 20, S. J. Lippard, ed., Wiley, New York, 1975, p. 299.

100. K. Bowman, D. P. Riley, D. H. Busch, and P.W.R. Corfield, J. Am.Chem. Soc., 97, 5036 (1975).

101. D. P. Riley, J. A. Stone, and D. H. Busch, J. Am. Chem. Soc.,98, 1752 (1976).

102. D. P. Riley, J. A. Stone, and D. H. Busch, J. Am. Chem. Soc.,.99, 767 (1977).

103. M. C. Rakowski, Ph.D. Thesis, The Ohio State University, 1974.

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