71-22,508
MERRELL , P hilip Hayden, 1944- SYNTHESIS AND CHARACTERIZATION OF SOME TETRADENTATE MACROCYCLIC COMPLEXES OF IRON.
The Ohio State University, Ph.D., 1971 Chemistry, inorganic
University Microfilms, A XEROX Company , Ann Arbor, Michigan
THIS DISSERTATION HAS BEEN MICROFILMED EXACTLY AS RECEIVED SYNTHESIS AND CHARACTERIZATION OF
SOME TETRADENTATE MACROCYCLIC COMPLEXES
OF IRON
DISSERTATION
Presented in Partial Fulfillment of the Requirements for •the Degree Doctor of Philosophy in the Graduate School of The Ohio State University
By
Philip Hayden M errell, B.S.
* * #
The Ohio State University 1971
Approved hy
A d v ise r Department of Chemistry To Dr. Robert T. Clark (1917-1966), who guided me throughout my undergraduate career not only as a superb coach and a great scientist, but also as a sincere Christian.
i i ACKNOWLEDGEMENTS
The author expresses his gratitude to his parents, Ralph J. and Audrey D. Merrell for the help he has received throughout his life and especially to his father for financial support throughout his undergraduate education and to his mother for her help in the preparation of the initial drafts of this dissertation.
Ke also wishes to thank the members of the ’’Busch team'' for their helpfulness during the many months of the research that have culminated in the writing of this dissertation. Special thanks go to
C. Robert Sperati for his help in the preparation of the many electronic and Mossbauer spectra that appear throughout this dissertation which were made by use of computer techniques. Also to Miss Jacqueline A.
Orwig who helped in countless secretarial ways with the final form of t h i s work.
A very special thanks goes to Dr. John A. Stone who took a complete novice in the wares of Mossbauer spectroscopy and taught him through his patience and expertise to become fairly capable in the use of a fragment of this fruitful field.
Much gratitude is also given to Dr. Daryle H. Busch, who through his many helpful suggestions and his encouragement has guided the author to become a confident research scientist.
i i i VITA
September 29 9 19^4- ...... Born, Akron, Ohio
1966 ...... B.S., Harding College Searcy, Arkansas
1 9 6 6 -1 9 6 7 ...... Teaching Assistant, Department of Chemistry, The Ohio State University, Columbus, Ohio
1967-1970 ...... Research Assistant, Department of Chemistry, The Ohio State University, Columbus, Ohio
PUBLICATIONS
' 'Five-Coordinate High Spin Iron(ll) Complexes of Synthetic Macrocyclic L i g a n d s '', J . Am. Chem. S o c ., in p r e s s .
FIELDS OF STUDY
Major Field: Chemistry
Specialization: Inorganic Coordination Chemistry Professor Daryle H. Busch TABLE OF CONTENTS Page ACKNOWLEDGMENTS...... i i i
VITA ...... iv
LIST OF TABLES ...... v i i
LIST OF ILLUSTRATIONS...... v i i i
INTRODUCTION
Biological Significance of Iron ...... 1 New Macrocyclic L igands ...... k- Preparation and Properties of Iron Complexes ...... 10 Oxidation States • •••••••••«•«•••••• 12 Electronic Structures and Ground States ...... 13 Magnetic Moments ...... IT Mossbauer^Spectroscopy ...... 21 The Mossbauer E ffect ...... 22 The Mossbauer Spectrom eter ...... 26 The Mdssbauer S p ectra ...... 28 The Chemical Isomer Shift ...... 28 The Quadrupole S p littin g ...... 32 EXPERIMENTAL
M a t e r i a l s ...... 35 S y n t h e s e s ...... 35 Physical M easurements ......
RESULTS AND DISCUSSION
Complexes of Fe(ll) and F e(lll) with the Macrocyclic Ligand m-CRH ...... k9 Ligand Synthesis ...... 50 Synthesis of the Iron(ll) Complexes of meso-CRH . . . 51 Synthesis of the Iron(lll) Complexes of meso-CRH . . 55 The Iron(ll) Complexes of meso-CRH ...... 54 Iron(lll) Complexes of m-CRH ...... 69 Complexes of Fe(ll) and F e(lll) with the Macrocyclic Ligand m-1,7-CTH ...... 71 S y n t h e s i s ...... 72 Infrared Spectra ...... 7^ Characterization of Complexes ...... 83
v Table of Contents - continued Page
Complexes of Fe(ll) with the Macrocyclic Ligand 1,7-CT • • 92 Mossbauer Spectra of the Complexes of Fe(ll) and F e(lll) with Macrocyclic Ligands ...... 99 Results of the Mbssbauer Spectra ...... 100 M bssbauer S p e c tra o f Some TAAB C om plexes ...... 121 Mbssbauer Spectra of Some TTP Complexes ...... 125
SUMMARY...... 127
APPENDIX ...... 130
BIBLIOGRAPHY...... 197
vi LIST OF TABLES Page T able I Elemental Analyses for Iron Complexes of m-CRH ...... 40 T ab le I I Elemental Analyses for Iron Complexes of m-1,7~CTH. . . k-3 T able III Some Physical Properties of Fe(m-CRH)2+’3 + Complexes • • 6l T able IV Electronic Spectra of Fe(ll) and F e(lll) m-CRH Complexes *63 T able V Some Physical Properties of Fe(m-1,7-CTH)2+’3+ Complexes. 75 T able VI Electronic Spectra of Fe(ll) and Fe(lll) Complexes of m -1 ,7 -C T H ...... 85 T able V II Properties of Fe(ll) Complexes of the Macrocyclic Ligand 1 , 7 -C T ...... 93 T able V III Mossbauer Spectra of Five-Coordinate High Spin Fe(ll) C o m p le x e s ...... 102 T able IX Mossbauer Spectra of Six-Coordinate High Spin Fe(ll) Complexes •••••••••••••••«••••• 10^ T able X Mossbauer Spectra of Low Spin Fe(ll) Complexes ...... po^ T able XI Mossbauer Spectra of High Spin F e(lll) Complexes .... 105 T ab le XII Mossbauer Spectra of Low Spin F e(lll) Complexes .... 105 T able X III Mossbauer Spectra of Miscellaneous Complexes ...... 105 T able XIV Physical Properties of Some New Five-Coordinate Fe(ll) C o m p le x e s ...... 110 T able XV Values of Field Gradient q and Asymmetry Parameters v f o r 3d Electrons ...... 113 T able XVI Variation of the Axial Ligand in Some Low Spin Fe(ll) C o m p le x e s ...... 119 T able XVII Mossbauer Spectra of Some Iron Complexes of the Macrocyclic Ligand TAAB ...... 122 T able XVH[ Mossbauer Spectral Parameters for the Iron Complexes o f TT P ...... 126
v i i LIST OF ILLUSTRATIONS Page 1. Some Tetradentate Macrocyclic Ligands Containing Nitrogen Donors . 6 2. Tanabe-Sugano Diagram for ds Ion in Octahedral F ield ...... lU 3 . Tanabe-Sugano Diagram for d 5 Ion in Octahedral Field ...... 16 b. d Orbital Diagram for Octahedral d 5 and d 6 Io n s ...... 18 5 . d Orbital Diagrams for Several Geometric Fields ...... 20 6 . Scheme for the Decay of Co 57 to Fe 5 7 ...... 26 7- Block Diagram for Mossbauer Spectrometer ...... 27 8 . Typical Mossbauer Spectrum Fe(m-CRH)C1 2 ...... 29 9. Origin of Isomer Shift ...... 31 10. Origin of Quadrupole S plitting ...... 33 11. Infrared Spectra of m-CRH* H20 and Fe(m-CRH)C1 2 ...... 55 12. Infrared Spectra of Fe(m-CRH) (NCS )2 ...... 58 13. Electronic Spectra of Fe(m-CRH)X 2 X=C1 , Br , 1 ...... 6b lb . Results of Strong Tetragonal Distortion ...... 67 15. Electronic Spectrum of [Fe(m-CRH)C1 2 ]BF4 ...... 70 16 . Infrared Spectra of m-1,7-CTH and Fe (m-1, 7-CTH )ci2 ...... 76 17- Infrared Spectra of Fe(m-1,7-CTH)(NCS )2 and F e(m -1 ,7 _CTH) - (C104 )2 -3CH3CN...... 79 18. Infrared Spectrum of Fe (m-1, 7-CTH) (cn )2 ...... 8 l 19. Electronic Spectra of Fe(m-1,7“CTH)c12, Fe(m-1,7~CTH)Br 2 and Fe(m-1,7-CTH)l2 ...... 87 20. Electronic Spectrum of Fe (m-1,7-CTH) (NCS )2 . . . . . * ...... 90 21. Infrared Spectrum of Fe(l,7-CT)(cH 3CN)2 I 2 ...... 9b 22. Electronic Spectrum of Fe(l, 7-CT)(CH 3CN)2 I 2 ...... 98 2 3 . d-Orbital Diagram for Octahedral and Square Pyramidal (C4V) S t r u c t u r e s ...... 112 2b. The Variation of Isomer Shift with the Degree of Unsaturation of the Macrocyclic Ligands ...... 115
25. The Variation of AEq with the Degree of Unsaturation of the Macro- cyclic Ligands ...... 117
v i i i Page 26. Plot of 6 v s AE q for Different Axial Ligands ...... 120
F ig u re s 27-87 are Mossbauer Spectra of:
27. Fe(m-CRH)Cl2 T= -19*4°C ...... 131 28 . Fe(m-CRH)Br2 ...... 132 29. Fe(m-CRH)l 2 ...... 133 JO. [Fe(l,7-CT)cnci0 4 ...... 13^ 31. [Fe(l,7-CT)C1]C10 4 T= -19*4°C ...... • 135 32. [Fe(1,7-CT)Br]C104 ...... 136 3 3 . [F e ( 1, 7 -C T )l]C 104 ...... 137 3*4. [Fe(l, 3s 7 , 10-CT)C1]C104 ...... 138 35. [Fe(l,3,7,l0-CT)C1]C10 4 T= -1 9 * 4 ° C ...... 139 3 6 . [Fe(l,3,7,10-CT)Br]C10 4 ...... 1*4-0 37. [ F e ( l, 3j Tj 10-C T)l]C 104 ...... 1*4-1 3 8 . [ F e ( l , 3 , 7,lO-CT)ci]BPh 4 ...... 1*4-2 39. Fe (m-CRH) (0Ac )2 ...... 1*4-3 k0. [Fe(m-CRH)0Ac]PFs ...... 1*4*4 Ul. Impure Fe(m-1, 7-CTH)Cl 2 ...... 1*4-5 *4-2 . Pure Fe (m-1,7"CTH)C12 ...... 1*46 *4-3. Impure Fe(m-1,7~CTH)Br 2 ...... 1*4-7 *4-1-4-. Pure Fe(m-l,7-CTH)Br2 ...... 1*48 *45. Im pure Fe (m-1, 7-CTH ) l 2 ...... 1*4-9 *4-6 . Fe (m-CRH) (C104 )2 ...... 4 ...... 150 *47- Impure Fe (m-1, 7-CTH) (NCS )2 ...... 151 *48 . Pure Fe (m-1,7-CTH) (NCS )2 ...... 152 *49. Fe (m-1, 7-CTH) (cn )2 ...... 153 50. Fe(m-1,7-CTH) (C104)2* 3 CH3CN ...... 15*4- 51. Fe(m-1,7-CTH)(C104)2* 2 Im idazole ...... 155 52. F e (1, 7-CT) (NCS)2 ...... 156 53. [ F e ( l, 7 -CT)CN]BPh 4 ...... 157
ix 5l+. [Fe(l,T-CT)(CH 3CN)2 ] ( c l 0 4 )2 ...... 15g 55. [Fe(l,T-CT) (isomer-2)] (CH3CU)2 (C104 )2 ...... 159 56. Fe(l, 7-CT)(imidazole) 2 (BPh 4 )2 160 57. Fe(l, 7~CT)(CH3 CW)2 I 2 ...... l 6 l 58 . [F e(l,5? 7,10-CT)(CH3CN)2 ](C 104 )2 ...... 162 59. [Fe(l, 7j10-CT)(CH3CN)2 1(BFh 4 )2 ...... 163 60. rFe(lJU,8,10CT)(CH 3CW)2 ](BPh 4 )2 164 61. F e ( l, 3? 7S 10-CT) (NCS )2 ...... 165 62. Fe (l,3 , 1 ,10-CT) (im id azo le ) 2 (C104 )2 ...... 166 6 3 . Fe(l,3, 7,10-CT)(imidazole) 2 (BPh 4 ) 2 ...... 167 6k. [Fe(m-CRH)C12 ]BF4 ...... X68 6 5 . [Fe(m-CRH)Br2 ]BF4 ...... 169 6 6 . [Fe(m-1,7"CTH)C12 ]C104 ...... 170 67. [Fe(m-lJ7-CTH)Br2 ]BF4 ...... 171 6 8 . [ F e ( l, 7-C T )ci2 ]C 104 ...... 172 69- [F e(l, 7-CT) (NCS ) 2 ]BPh 4 ...... 173 70. [Fe(l,7-CT)(CH 3CN)2 ](C104 ) 3 ...... 1714. 71. [Fe (l, 7-CT) (NCS )] 2 0(NCS)2 ...... 175 72. F e(l, 7-CT )Fe(NCS )4 M a r o o n ...... 176 73. Fe(l,7-CT)Fe(NCS )4 Exposed to air (Black) ...... 17T 7^. [Fe(l,3s 7S10-CT)(Phen)](C10 4 )2 ...... 178 75- [Fe(l,7-CT)]2 (C104 )3 *H20 ...... 179 76. Fe(TAAB)F(d04 )2 *2H20 ...... '..... l8 0 77. [Fe(TAAB)F](BF4 )2 ...... 181 78. Fe (TAAB ) f (N0 3 )2* 2H2 0 ...... 1 8 2 79. [Fe(TAAB)]2 0 ( d 0 4 )4-lj-H20 ...... 183 80. [Fe (tAAB)] 2 0 (n 63 )4* ^H2 0 ...... 18^ 81. [Fe(TAAB)] 2 0(BPh 4 )4 ...... 18 5 82. {Fe[TAAB(0Me)2]}20 ...... 186 83 . Fe(TAAB)(NCS)2 ...... 18 7
x Page
Sh. Fe(TTP)(BF4 )2 ...... 188 85 . Fe(TTP)(CH3CN)2 (BF4 )2 ...... 189 86 . Sodium Nitroprusside Stan d ard ...... 190 87. o^“Fe203 ...... 191
x i INTRODUCTION
Biological Significance of Iron
The occurrence of iron compounds in biological systems is extensive. Of the transition metals only copper is as abundant in living systems as iron. Iron compounds function in biological transport systems, photosynthesis, nitrogen fixation, electron transfer, and in enzymes for the catalysis of various chemical reactions. Among the naturally occurring compounds of iron are the complexes of the tetra- dentate macrocyclic ligand porphyrin(l) which utilizes nitrogen as its ligating atom. The porphyrin ligand is found throughout nature in complexes of Cu, V, Co, Mg and several other metal ions, as well as Fe.
Iron complexes of porphyrins serve many functions in the chemistry of living systems. Among these functions are the catalysis of peroxide breakdown into less harmful substances (catalase), the catalysis of oxidations by hydrogen peroxide (peroxidases), their use as one electron transfer reagents (cytochrome c), and the reversible transport
(hemoglobin) and storage (myoglobin) of molecular oxygen. Several reviews have been written on these essential products (l-j). Studies on these hemo- and hemi-proteins in the last decade have finally established a foundation for the study of related enzymes.
The macrocyclic complex hemoglobin reversibly transports oxygen to the cells of the body where it is utilized in energy
1 2
production. Hemoglobin consists of a conjugate protein containing a prosthetic group, heme, and the protein, globin. Heme is a complex
of iron(ll) and protoporphyrin IX which is a derivative of porphyrin with methyl, vinyl, and propyl groups substituted onto the backbone.
A hemoglobin molecule consists of four heme units coordinated through
imidazole groups of the histidine proteins in the globin. The heme group is surrounded by the protein so that the iron(ll) is not
attacked by potential ligands that could interfere with its function.
The iron(ll) atom is not bound precisely in the plane of the four nitrogen atoms but is displaced about 0 .5 a out of the plane ( 1+, 5 ).
The increased bond distance and decreased ligand field accounts for the high spin state of the iron(ll) in deoxyhemoglobin. The iron atom existing out of the plane also accounts for the five-coordination number of this molecule. When molecular oxygen is bound, the iron becomes six-coordinate and low spin ( 6 , 7 ).
HN
I Porphyrin 5
There are currently three theories as to the geometry of the hound oxygen molecule. It can be bound parallel to the porphyrin plane, perpendicular to the plane, or at some intermediate angle.
There has been much discussion on this point, however, the mode of bonding is still not clear ( 8 ).
Deoxyhemoglobin has been shown to be a square pyramidal five- coordinate complex ( 9 ) as has the sim ilar complex deoxymyoglobin.
At least one other naturally occurring oxygen-carrier exists,
which is hemocyanine (lO) containing Cu(l). The study of the natural oxygen-carriers has been hindered by the complexities of their structures. However, investigations of synthetic oxygen-carriers as model compounds have aided the elucidation of the physical and chemical properties of the more complex molecules.
Recently elegant experiments have been carried out by Petering and Hoffman (ll) in which a Co(ll) ion replaces the Fe(ll) ion in hemoglobin. Coboglobin, as this new compound is called, is an exceptionally useful model for the study of hemoglobin and myoglobin because of its electronic state which makes available EPR studies.
It has also been found to reversibly carry oxygen molecules. Among the major results is the conclusion from EPR data that a negative charge resides on the 02 molecule and that the Co ion acts as if it were in the low spin +3 s t a t e ( 1 2 ). k
Synthetic oxygen-carrying iron complexes are rare. An example
is found in dioxane solutions of dimethylglyoxime and Fe(ll) which also
contain bases, such as histidine, ammonia, pyridine, or imidazole ( 1 3 ).
In these solutions iron combines with molecular oxygen to yield oxygen
complexes. The reaction can be reversed by bubbling nitrogen through
the solutions.
Fe(DMG)2 *base 5= ^ Fe(DMG)2 02 *base N2
New Macrocyclic Ligands
During the last decade several new tetradentate nitrogen macro-
cyclic ligands have been prepared. These ligands have been studied as
complexes of many metal ions including N i(ll), Co(l,II, III), C u(l,II),
Zn(ll), and several others. Several reviews (lU-1?) have been written
describing their syntheses and their chemical and physical properties.
A number of these synthetic ligands are shown in Figure 1. Abbrevi
ations for the ligands are given with the structures. Although this
set of structures is by no means exhaustive, it serves to show the
great variety of ligands of this class.
Phthalocyanine(ll) complexes have been known for many years because of their use in the dye industry. The ligand, phthalocyanine,
has many properties similar to those of the natural porphyrins. The
main differences are the azido-linkages which replace the methine 5
bridges between pyrrole rings, and the four benzo-groups fused to the
pyrrole rings. Complexes of phthalocyanines have been studied
extensively ( 18 , 1 9 ).
Curtis began his study of synthetic macrocycles with the
accidental discovery of the complexes of nickel with the ligands
I,ij--CT(IV) and 1,7-CT(v) by the reaction of trisethylenediam ine
nickel(ll) perchlorate with acetone (20,2l). Many derivatives and
related compounds have been prepared and studied since that time ( 1 5 )-
Complexes of the fully saturated ligands 1, U-CTH(vi) and 1, T-CTH(vil)
have also been prepared and studied extensively ( 2 2 ). Complexes of
the Curtis ligands have been studied with Ni, Co, Cu, Zn and several
other metal ions ( 1 6 ).
Several groups of researchers began working in this area
simultaneously beginning about I 960 ; and, consequently, the wide variety of ligands shown in Figure 1 exist with varying numbers of
donating groups, varying types of donating groups, varying numbers of
carbons in the backbone, and various amounts of saturation. Develop ments in this area have contributed substantially to the knowledge of
the chemistry of metal complexes.
One of the striking facts is the very small number of complexes
of iron that have been reported with these ligands. It had long been
known that iron forms complexes with the macrocyclic ligands that play
an important role in biology, and this is one of the important reasons NH
N C==N
N
II Phth III TAAB
rvN N-n. "V"~xiTN N CAAN N M IV 1,4-C T V 1,7-CT
NH HN- N H HN.
■NH HN' NH HN
VI 1,4-CTH VII 1 ,7-CTH
Figure 1. - Some Tetradentate ilacrocyclic Ligands Containing Nitrogen Donors 7
h 3c CH,
H—N N H
V III CR IX CRH
„ r i «
C""N Hk^H H W H X Cyclam XI
H Nn N. H u ’N N * X II DIM U X III TIM
Figure 1 continued 8
for the study of these ring systems. These synthetically produced
ligands are similar to porphyrin in that they are all cyclic tetra- dentate ligands and contain four nitrogen donors. They differ from porphyrin in the number of members in the ring and the amount of
conjugation throughout the ring. Some resemble porphyrin more closely
than others.
The only macrocyclic ligands of this class that have previously been used to form complexes with iron salts are phthalocyanine(ll),
TAAB(lIl) and substituted porphyrins.
The phthalocyanine complexes of iron are well known ( 23, 2k)
and have been studied quite extensively (l8). These compounds have
shown a variety of coordination numbers and electronic structures but have been fairly difficult to study because of their insolubilities.
Fe(TAAB)2+,3+ compounds have been studied only recently (25 j 2 6 ).
The F e(lll) compounds show a strong tendency to form Fe-O-Fe oxobridged
species which exhibit spin-spin coupling between the electrons of the
two high spin Fe(lll) atoms.
Tetraphenylporphyrin complexes of iron have been prepared and
studied by several groups of chemists (2T~3l)* The synthetic porphyrin
ligand is preferred for some studies instead of natural porphyrins because of its ease of preparation. The reactions of pyrroles with
the appropriate aldehydes, under pressure and at high temperatures,
yield these substituted porphyrin derivatives. The tetrapyridyl- 9
porphyrin complexes are among the more useful derivatives that have
"been prepared ( 3 2 - 3^ ).
Since the study of the natural oxygen-carrier, hemoglobin, is so difficult, it is thought that if synthetic oxygen-carriers can be produced, much w ill be gained in the understanding of the requirements for reversible oxygen transport. The manner of oxygen binding, require ments for transport and mechanism of 02 release are not yet clear in the case of the natural products.
The objectives of the studies reported in this thesis are related to the fact that very few synthetic iron (il or III) macro- cyclic complexes have been studied. This dissertation reports the synthesis, characterization, and some of the reactions of several new iron complexes with macrocyclic ligands. The two ligands that have been studied to the greatest extent are meso-CKH(lx) and meso-1,T-CTH(vil). Several derivatives of other ligands have been prepared, and these w ill be discussed briefly.
The ligand, meso-CRH, is well characterized in complexes of n ic k e l (35-3 7 ) and cobalt (38 , 3 9 ). The ligand is produced by pre paration of the parent Ni(CR)2+ complex and subsequent reduction with hydrogen and Pt02. There are two isomers formed, and the structures of these isomers which had been determined by Karn were confirmed by
Dewar and Fleischer (^0). This ligand contains a pyridine molecule in its structure. The me so isomer has the two methyl groups on the 10
same side of the plane while the d , 1 isomers have methyl groups on opposite sides of the molecular plane.
The ligand meso-1,7-CTH (structure VTI) is also well known. It is one of two isomers obtained in the reduction of the Curtis macro- cycle, 1,7“CT. Complexes of this ligand have not been discussed as completely; however, its nickel(ll) complexes have been thoroughly characterized ( 2 2 ). These ligands were chosen because of their avail ability as pure ligands, and because they are not subject to rapid hydrolysis.
Preparation and Properties of Iron Complexes
Iron complexes are generally more difficult to prepare than are those of nickel and cobalt. Consequently, when an interesting new ligand becomes available, it is often convenient to postpone the study of the iron complexes until the complexes of other metal ions have been studied thoroughly. The preparation of complexes of iron with nitrogen donor ligands is difficult because of the basic properties of the ligands, and the ease of oxidation of Fe(ll) to Fe(lll) in wet basic media. Thus, hydrolysis and oxidation often result in the formation of oxides and hydrous oxides of iron. Iron(ll) complexes with simple nitrogen ligands must be made and handled under an inert atmosphere and under otherwise carefully controlled conditions because of their sensitivity to oxygen and moisture. Failure to maintain rigorously oxygen-free and water-free conditions (although water does not cause 11
problems in the absence of oxygen in many of these complexes, e.g. m-CRH complexes) results in formation of a red-brown gelatinous or colloidal suspension instead of the desired product.
Iron complexes have been found to exhibit several different coordination polyhedral structures. The predominant geometry is the six-coordinate octahedron, and most of the known iron compounds have this structure. Although more rare in occurrence, a number of compounds are known to have five-and four-coordinate structures. Most of the five-coordinate iron compounds have the trigonal bipyramidal geometry; e.g. the terpyridyl derivatives (1+1,1+2), or square pyramidal geometry; e.g. complexes with phthalocyanine (1+3)? salicylaldimine (1+1+, 1+5)? and other ligands with strong in-plane fields (i+6,i+T)•
In the case of iron the tetrahedral structure most often occurs with anionic ligands in such substances as FeX42_ and FeX4 where X is a halide or pseudohalide. However, there are a few dithioferrates(m )
(1+8) and carbonyls (1+9) that are also tetrahedral.
A few complexes of higher coordination number have been reported. One example involves a seven-coordinate iron(lll) complex which contains a pentadentate macrocyclic ligand ( 5 0 ). A second example involving eight-coordination is the complex tetrakis-(l,8-naphthyridine) i r o n ( l l ) ( 5 1 ). These are only possible because of the specific 12
structures of the ligands. Complexes of iron.with ethylenediaminetetra- acetic acid and propylenediaminetetraacetic acid also have coordination numbers of seven and eight ( 5 2 ).
Oxidation States - Iron compounds exist in oxidation states from -2 to +6 , however, iron is normally found in either the +2 o r +3 oxidation state. This results in a ds or a d 5 electronic configuration respectively. With common ligands bound to iron(ll), it is easily oxidized to the higher state by oxygen; however, several ligands stabilize Fe(ll), and their Fe(lll) complexes tend to reduce back to
Fe(ll), e.g. trisphenanthroline complexes of iron ( 5 3 )* Although many of the iron(ll) compounds of interest are sensitive to oxygen, they can be stored indefinitely in the absence of moisture.
Oxides of iron in the +4 and +6 states are known and are made by air oxidation at high temperatures. The +6 state is also known in ferrate salts Fe042~ and the +b state has been claimed in a few com plexes.
A nother common o x id a tio n s t a t e i s th e Fe(o) state. Iron organo- m etallic and carbonyl compound are well-known and normally exist in the zero oxidation state, however, some organometallic complexes form stable complexes of -1 and -2 oxidation state (^ 9 )*
Almost all of the coordination complexes of iron exist in the
+2 o r +3 state. However, recently a complex in the +J+ state has been 15
synthesized. The complexes of diarsine(xrv) have been shown to be d4
cases with two unpaired electrons ( 5^). These are prepared by
oxidation of the corresponding F e(lll) diarsine complex in 15M nitric
a c id .
XIV
Electronic Structures and Ground States. - The ground term
state for a d6 configuration is 5D, and this is the only quintet term for this configuration. The higher terms are all triplet or singlet
states. In weak octahedral ligand fields, the 5D term splits, and the ground state is 5T2g while the only excited quintet state is the 5Eg also originating from the 5D term. See Figure 2. Ik
70
i ’ t ,
60
30
40
30
20
O q/8
Figure 2. -Tanabe-Sugano Diagram for ds Ion in Octahedral Field 15
For Fe(ll) this means that the visible spectra of octahedral high spin complexes have only one d-d transition, that being the
5T2g—=>5Egs which in the hexaquo species occurs around 1 0 ,0 0 0 cm 1.
This band is expected to be a doublet because of Jahn-Teller distortion; however, this splitting is usually difficult to detect.
In strong octahedral fields the ground state is 1Aig arising from one of the free ion states which has dropped to lower energy than th e 5 T2g state. (See Figure 2.) This configuration corresponds to the low spin Co(lll) ion and gives rise to three spin allowed and two spin forbidden d-d transitions. The spin allowed transitions are
1A ig —» 1T ig , 1A ig —► 1T2g , and 1A ig — *• 1Eg and the spin forbidden transitions are 1Aig —*3Tig and- 1A ig — >,3T2g* Spectra of this type are not well-documented, however, they are comparable to the spectra of low spin Co(lll) complexes which are quite well know n.
Although the octahedral structure is often assumed for the six-coordinate structure, it is, in fact usually the exception. More often is the case where at least one or two of the ligands is different from the rest. In most of the cases discussed here the macrocyclic ligand w ill be considered planar, thus bonding four nitrogen atoms in a plane. However, the axial ligands are quite different from the nitrogen atoms in the plane which gives a tetragonal distortion. The amount of distortion depends upon the difference in strengths of the axial and in-plane ligands. This distortion causes a splitting of the 16
higher energy 5 T?g into a sEg and 5 B2g, and the low energy sSg into
5A ig and 5 Big. This tetragonal distortion is manifested in the electronic spectrum as a broadening and sometimes an actual split of th e 5E g — * 5 T2g band. In macrocyclic complexes of nickel(ll), this phenomenon is quite well known ( 1 5 ).
For tetrahedrally coordinated Fe(ll) the ground state is SE, and a single transition from SE —»5T is expected.
The electronic ground state of Fe(lll)d 5 i s s S, and i t i s the only sextet state in the energy scheme of the configuration d5. (See
F ig u re 3 ). There are several quartet states to which transitions take
70 -/■
60
50
40 V.m
; *29 '*A„ 1 *T<» i I 1 I
D q/B
Figure 3. -Tanabe-Sugano Diagram for d5 Ion in Octahedral Field 17
place; however, in Fe(lll) spectra charge transfer hands are normally sufficiently strong to obscure almost all d-d transitions. These transitions are weak also because they are spin forbidden. Since
Fe(lll) is isoelectronic with Mn(ll), the same transitions are expected in the spectra of the two ions. The d-d transitions are.observed fairly readily in the Mn(ll) case because of their low energy.
In sufficiently strong octahedral fields the ground state becom es T 2g. The spectra associated with this ground state are not discussed in great detail in literature; however, there are several quartet states and doublet states to which transitions can be made.
These are probably not normally seen because of their low extinction coefficients for the charge transfer bands occurring in the same spectral region have extremely high extinction coefficients.
Magnetic moments. - Figure ^ presents the d orbital diagrams for octahedral fields on both the d 5 and th e d 6 ions in both high spin and low spin states.
From these diagrams the number of unpaired electrons is obvious. In the d 6 high spin case there are four unpaired electrons which give magnetic moments in the range from 5*0 to 5*6 B. M. for Fe(ll).
The expected moment from spin only considerations is b .9 B. M. T his is much smaller than that actually measured because of the large orbital contribution of the 5T2g s t a t e . 18
11 J_ j.
J_L_L JL1L1L JLJ_± JLiLJ_ d6 high spin d6 low spin d5 high spin d5 low spin
Figure 4. d Orbital Diagram for Octahedral d5 and d6 Ion.
The low spin complexes of Fe(ll) should be diamagnetic because all the electrons are paired; however, the susceptibilities of Fe(ll) low spin complexes normally are not zero because of the second order
Zeeman effect. The susceptibilities are independent of temperature
(T. I. P. ) and values on the order of 50x10 6 e.g. s./mole are commonly observed for Fe(ll) with even greater values being found for Co(lll).
This accounts for the occurrence of magnetic moments between 0.5 and
1.0 B. M. Also, because of the difficulty of obtaining pure low spin
Fe(ll) complexes without a small amount of high spin impurity, some contributions to the susceptibility often result from impurities. 19
The magnetic moments of high spin d 5 Fe(lll) compounds fall in the range from 5 *6 to 6.2 B.M. There are five unpaired electrons, and the spin only value would he 5* 92 B. M. Most of the compounds exhibit moments close to this value because there is no orbital contribution (SA ground state), and there is no T.I.P.
The low spin d 5 state shows only one unpaired electron, and the spin only value for jj, eff would be 1.73 B. M. j however, because of the 2T ground state, a large orbital contribution is obtained. This orbital contribution usually raises the moment into the range from
1 .7 to 2 .k B.M.
Interesting examples have been found wherein the magnetic moment is intermediate between the predicted ranges. For several years compounds that had temperature independent magnetic moments in the range from 3 .3 to 3*5 B.M. have been discussed, e.g. Fe(ll) complexes of phthalocyanine and terpyridine ( 53j 55j 5 6 ). C>ne plausible explanation of this phenomenon is presented in Figure 5 (^3)* This model indicates that two unpaired electrons are to be expected for low spin Fe(ll) compounds of strong tetragonal fields and of square pyramidal and trigonal bipyramidal geometry. — — dxz > dyz Oh Tetragonal ®?Uare„ , Trigonal Pyramidal Bipyramidal
Figure 5. d Orbital Diagrams for Several Geometrical Fields
Another interesting case is discussed by Goodwin and Sylva in which an Fe(ll) complex of 2,i4--bis(2-pyridyl)thiazole goes from
S=2 to S=1 (u e f f 5. 3^ 298° K to 3 .2+5 at 80° k) upon lowering the temperature (^7). This complex is thought to be trigonal bipyramidal in structure.
In the Fe(lll) case the splitting diagrams for the tetragonal, trigonal bipyramidal, and the square pyramidal structures are the same as shown in Figure 5 except that there is one less electron yielding an
S=3/2 case. This magnetic state has also been discussed quite 21
extensively (^-5, 57, 58, 59,6 o ) j t h i s s t a t e b e in g common f o r F e ( l l l ) complexes of tetraphenylporphyrin, phthalocyanine and several others
lig a n d s .
The phenomena of Fe(ll) and F e(lll) complexes having normal high spin moments (that follow Currie-Weiss laws) at high temperatures and low spin moments at low temperatures is well documented ( 6 l - 6 l ) .
This concept is relatively new but is thought- to be understood now.
Another interesting phenomenon that exists quite abundantly among iron(lll) compounds is that of superexchange. Certain compounds were observed to have lower moments than expected. These compounds were found to be dimers or polymers and spin-spin coupling occurs through the oxygen bridges. The first case of such spin-spin coupling was copper acetate which is reviewed by Kato, Johnson and Fanning ( 6 8 ).
Later Earnshaw and Lewis developed the theory that governs the change of magnetism with temperature from the relatively abundant compounds containing Fe-O-Fe bridges (l9).
In the last decade there have been many complexes synthesized that have been found to contain Fe-O-Fe bridges and to follow the theory of Earnshaw and Lewis quite well.
• • Mossbauer Snectroscopy
• • About fifteen years ago Rudolph L. Mossbauer discovered a phenomenon that has greatly enhanced the study of compounds of iron. 22
This phenomenon is discussed here in a fairly complete form so that the principles and applications of this tool w ill he understood.
The Mffssbauer Effect. - Since the discovery of this effect by Rudolph L. Mossbauer in 1957 (70-72), it has been used quite extensively in chemistry and physics with most of the earlier work done in physics. In the field of chemistry Mossbauer spectroscopy has proven to be an extremely useful tool for the study of compounds containing Mossbauer-active elements. The purpose of this portion of the introduction is to present the important principles of the
Mossbauer effect and to outline how this effect applies in obtaining information about structure and bonding in the chemistry of iron complexes. Several reviews and books have been written during the last few years that present in great detail the applications of the
Mossbauer effect in Inorganic Chemistry (73“80).
The Mossbauer effect is defined as the recoil-free emission and subsequent absorption of gamma-rays. The lifetim es of the nuclear excited states that produce Mossbauer gamma rays are between 10 7and
10* 10 seconds, and the resulting (Heisenberg uncertainty principle) lin e w id th s a re 10 7 -1 0 10eV ( 8 l). Linewidth is often expressed in terms of the ratio of the width to the total energy of the gamma rays u sed ( l0 4 - l 0 5ev), this ratio being in the range 10 12 to 10 14. T h is indicates that such measurement must be extremely precise; i.e .,
Mossbauer lines are approximately 10 9 times as sharp as the sharpest 25
infrared lines. It is this precision that makes Mossbauer spectro
scopy such a sensitive method of studying the influence of external
forces on the electric and magnetic properties of the nucleus.
There are two reasons that gamma rays emitted from a radio active nucleus normally cannot be resonantly reabsorbed by a sim ilar nucleus. The first reason is that thermal motion of the nuclei within the source give rise to Doppler broadening so that the gamma rays are not emitted with their precisely natural energy. The second reason is that the gamma emission (being of high energy) imparts a nuclear recoil, so that the resultant gamma ray energy is somewhat smaller than the full transition energy, by the amount of recoil energy producing the recoil. Consequently, the resonance condition is destroyed.
However, if the emitting nucleus is bound in a crystalline lattice at very low temperatures, there is a finite probability that gamma ray emission w ill occur without energy loss due to recoil—the crystalline lattice absorbing the recoil energy as a single unit. The resulting quantized lattice-vibrational energy is termed the ’'phonon energy’*.
However, for a large number of events a small fraction w ill occur without lattice excitation, and this fraction of recoil-free emissions w ill give rise to gamma rays of natural line width which possess the energy of the nuclear transition. These gamma rays can now be used to cause resonant recoilless absorption with nuclei which have identical energy levels. 2k
Whenever the electronic environment of the absorbing nuclei is not precisely the same as the emitting nuclei (different oxidation state or different ligand environment), maximum resonance w ill not occur because the energy difference between the ground and the excited states of the two nuclei w ill not normally be the same. How ever, by using the Doppler effect, the energy of the emitted gamma ray can be changed so as to achieve maximum resonance. These differences in energy of the nuclear states only amount to about 10 12 o f th e total gamma energy, so that the relative velocities are only about 1 cm per second.
Mossbauer active elements are numerous. There are several properties that an isotope should possess if it is to be active.
Several of these properties are listed below:
1. The isotope should be heavy to reduce the effects
o f r e c o i l .
2. The gamma ra y en erg y sh o u ld be sm a ll (Ey< lOOKeV) a ls o
to reduce recoil effects.
J. The first excited state should have a half-life in
the range of 10 7 seconds so that the line width is
neither too broad nor too narrow.
4. The internal conversion coefficient (a) should be low
(i.e. few interactions with the electrons of the emitting 25
atom which lower theSy) because each internal conversion
event lowers the number of usable gamma rays.
5. Nuclear spin states should not be too complex or else
very complicated spectra w ill result.
6. The natural abundance of the active isotope should be
high, and it should be stable (or have a long half-life).
These factors make the number of nuclei which can be useful for
Mossbauer investigation quite small. Iron probably has the best combi nation of all of these factors^ and, therefore, it has been most extensively studied. Tin also is a good element for study. The
Mossbauer spectra of compounds of europium, xenon, gold, iodine, nickel, tellurium, ruthenium, tungsten, iridium, and several other elements have been studied to a much smaller extent because of the difficulties involved in obtaining good spectra.
The Mossbauer active nucleus of interest here is iron 57 and attention w ill be confined to iron in the discussions that follow.
F ig u re 6 shows the decay scheme of Co 57 from which excited Fe 57 n u c le i a re form ed. As can be seen th e re a re gamma ra y s o f en erg y 137KeV,
1 2 2 .6 KeV, and l!+. Ij-KeV. (There is also a 7KeV X-ray resulting from internal conversion.) The 1^. kKeV gamma ra y i s th e one s e le c te d f o r use in iron Mossbauer spectroscopy, and the others are removed by filters or selected against by the adjustment and design of the counting instrument. 26 570 270 D a y C o 57
Electron Capture (~0.6 MeV)
137 s e c s .
> 4) W a >* &o u
F ig u r e 6 . Scheme for the Decay of Co to Fe 57
The Mossbauer Spectrometer. - Figure 7 shows a block diagram of the Mossbauer spectrometer built and operated at the U. S. A. E. C.
Savannah River Laboratory in Aiken, South Carolina, by Dr. W. L. PiUinger and J. A. Stone. It is to be noted that the raw data is punched out onto IBM cards which then can be fed into a computer and graphed into normal spectra. A more detailed discussion of the instrument is in the experimental section. Channel-Advance Pulses
P ic k u p M u lti- O utput D riv e C h a n n e l D riv e Electronics A n a ly z e r
/ / \ Synchonization Pulses
.A b so rb er i O ± L Proportional H C o u n te r | K e y p u n c h
HOI ource V e lo c ity T r a n s d u c e r pt-Iigh- J V o lta g e S c a la r f S upply
S in g le- >j P r e - L in e a r ^ Channel § A m p lifie r! A m p lifie r I A n a ly z e r j L ------. L. L— .J.l'rT-.-J
Figure 7. Block Diagram for Mossbauer Spectrometer. 28
The Mossbauer Spectra. - A normal Mossbauer spectrum is a plot of the number of gamma rays passing through the absorber versus the velocity of the source relative to the absorber. A typical spectrum is shown in Figure 8. There are three important characteristics of a
Mossbauer spectrum.
1. The number o f peaks
2. The shape of the peaks
3. The position of the peaks along the velocity scale
The normal number of peaks in a spectrum of iron (in the absence of external or internal magnetic fields) is two. These peaks should be sharp and symmetrical if only one iron species is present.
The two peaks arise because when a non-zero electric field gradient
(efg) interacts with the nuclear quadrupole moment ' ' quadrupole splitting'' occurs. This w ill be discussed later.
The position of the peaks along the velocity scale is called the isomer shift and is a measure of the s-electron density at the n u c le u s.
The Chemical Isomer Shift ( 5 ). - If the atoms in the absorber exist in a state different from that of the source, then, to assure maximum resonance, the absorber must be moved relative to the source to change the energy of the gamma rays (Ey). Therefore the energy of the levels of the absorber relative to the energy of the source nucleus is indicated by the corresponding source velocity. These differences RELATIVE TRANSMISSION 97.00 97.75 98.50 99.25 100.00 -U .00 -3 .0 0 -2 .0 0 -1 .0 0 0.00 1.00 2.00 3.00 3.00 2.00 1.00 0.00 0 .0 -1 0 .0 -2 0 .0 -3 .00 -U gure 8 Mcssbauer r f ( CRH) y, ly )C H R -C e(m F of tra c e p S r e u a b s s c M 8. e r u ig F EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY t % + * * + + + 4> <► + + * / V + 4 - + h *■++*#++ -h U.00 5.00 vo 30
in values are called the isomer shifts (§) and are compared to some
standard compound (in this case, sodium nitroprusside, the currently
accepted standard) ( 8 2 ).
The isomer shift arises because of the interaction of the gamma
ray with electronic charge. The greater the difference in size between
the nuclear excited state and the nuclear ground state, the greater the
transition energy w ill be. The potential energy of the system is also
increased by an increase in electronic charge in close proximity to
the nucleus. Since only s-electrons have a finite probability of being at the nucleus, they are the only electrons that directly involve
the potential energy of the system. Other types of electrons must be
considered, however, because they shield the s-electrons from the
n u c le u s.
The expression which relates the isomer shift to the s-electron
density at the nucleus and the change in nuclear radius is shown below ( 7 7 ):
(1 ) 6 - f t / 5 ) « K=a2 (4E/R){M,s(0 ) ) 2 absorl)er-(*s(0)^ souroe}
where R = radius of the ground state nucleus
e = charge on an electron
R = change of nuclear radius upon excitation (i.e. Rex-Rgr)
fi|;s (o ) ] 2 = s-electron density at the nucleus 31
This expression is based on the fact that when the source and absorber are chemically different, the s-electron density of nuclei of the two w ill differ in both their excited states and their ground states. This is represented diagrammatically in Figure 9-
Excited State n I (Eex_Egr)absorber (E ex ^gr/sourceE a r )
Ground State JL J
Figure 9. Origin of Isomer Shift
Equation 2 shows this relationship in mathematical form:
( 2 ) 5 = (Eex-Egr) . (Eex-Egr) B abs source
From equation 1 it is observed that the isomer shift is directly proportional to the difference in the s-electron density between the absorber and the source, and to the ratio AR/R. If AR/R is zero, i.e. the ground state and excited state radii are equal, no isomer shift is observed and the technique could have no chemical applications. How ever, in all cases AR/R is not equal to zero. Consequently, the chemical isomer shift yields much chemical information. The quantity 32
ar/ r is negative for iron. This indicates that the excited state radius is smaller than the ground state radius.
There are several factors that influence the isomer shifts.
Some o f them a re as fo llo w s :
1. The oxidation state of the iron
2. The coordination number
3« The nature of the bonding, i.e ., the type of groups
to which it is bound
if. The orbitals used in bonding
Much information about the above points can be deduced from isomer shift data.
Isomer shifts for octahedral iron compounds normally fall in the following ranges:
1. Fe(ll) h. s. 1 .1 - 1 .3 mm/ sec
2. Fe(ll) 1. s. 0. 1+ - 0 .6 mm/sec
3 . F e ( l l l ) h. s. 0. 5 ~ 0. 7 mm/sec
If. Fe(lll) 1. s. 0 .3 - 0. 7 mm/ sec
The Quadrupole Splittings (aEq). - Normally the Mossbauer spectra of iron compounds exhibit two peaks. These peaks arise from a nonzero electric field gradient which interacts with the nuclear quadrupole moment. This is detailed as follows. For nuclei with spin quantum number I > 1, the nuclear charge distribution is nonspherical and the 33
nucleus has a quadrupole moment. If some outside gradient surrounds this nucleus, the 21+1 fold degeneracy is partially lifted and a quadrupole splitting of states arises.
For Fe57, the excited state has a spin quantum I = J>/2.j th u s , there are four degenerate levels corresponding to nuclear magnetic quantum numbers Mj = 3/2, l/2, -l/2, "3/2. When an efg is present, part of this degeneracy is lifted. In the ground state I = l/2 so in the presence of an efg, the levels remain degenerate. This is shown diagrammatically in Figure 10.
Figure 10. Origin of Quadrupole Splitting
In general anything that causes an unsymmetrical distribution of electrostatic charge around the nucleus causes an efg. There are two major reasons for the existence of an efg around the nucleus of an iron atom. The first and most important is the distribution of the d-electrons themselves. For example, in octahedral Fe2+ high spin
complexes there is an ' ' odd’' electron in the T2g level which causes
an efg. While in the low spin case all the electrons are paired, and
no efg is present.
The other factor contributing to the efg is that of the
arrangement of the ligands attached to the iron atom. Cubic arrange
ments, such as octahedral or tetrahedral, w ill not generate an efg
while departure from these geometries w ill cause an amount to be added
to the efg.
These two factors are additive; however, difficulty often
arises because the algebraic signs of the contributing terms are not
known.
The AEq values for normal octahedral complexes of iron fall in
the following ranges:
1. Fe(ll) h. s. 2 . 0 - 3 .5 mm/sec
2. Fe(ll) 1. s. 0 - 1. 0 mm/ sec
3. Fe(lll) h. s. 0 - 1. 0 mm/sec
Ik Fe(lll) 1. s. 1 . 0 - 2 . 0 mm/sec
Thus, the ground state of normal octahedral complexes of iron
can often be inferred by a comparison of isomer shift and quadrupole
splitting values. From Mossbauer data it is difficult to distinguish between low spin Fe(ll) and high spin Fe(lll); however, by use of
magnetic moments, such difficulties can be resolved. EXPERIMENTAL
M a te ria ls
The 2 , 6 -Diacetylpyridine was obtained from Aldrich Chemical
Co. and recrystallized once from abs. ethanol. Ethylenediamine and
3, y -Diaminodipropylamine were obtained from Matheson, Coleman and
Bell Chemical Co. and used without further purification. Platinum
oxide catalyst was obtained from Matheson, Coleman and Bell Chemical
Co.; and sodium borohydride was obtained from the City Chemical Co.
The ferrous salts were obtained from Alfa Inorganics, Inc. and were used as obtained except for the bromide, which was refluxed in tetra-
hydrofuran over iron filings under nitrogen, filtered, and the THF
evaporated off under vacuum, affording the anhydrous ferrous bromide.
All other chemicals were reagent grade. Solvents were heated to boiling
and nitrogen was bubbled through to deoxygenate them.
S y n th eses
Preparation of 2.12-Dimethyl-3.7.11.17-tetraazobicyclor11. 3.11-
heptadeca-l(lT)a 2,11, 1 3 ,15-pentaenenickel(ll)Perchlorate,Ni(CR) (ciO^Jp.
- Equimolar amounts of 2, 6 -Diacetylpyridine and 3> 3' “diaminodipropyl
amine were condensed in the presence of nickel(ll) ion, as described
elsewhere ( 3 6 ). Calc, for NiCi 5H22N4Cl2 08: C, 34.91; H, 4.26;
N, 10.86. Found: C, 3 4 . 83 ; H, 4 .3 2 ; N, 10.70.
35 36
Preparation of 2 , 12-Dimethyl-^, T-11, 17-tetraazabicycloril. 3.1")- heptadeca-l(l7), 13?l^-trienenickel(ll) Perchlorate, meso isomer, ______
Ni(m-CRH)(C1 0 4 )g. - The complex Ni(m-CKH)(C10 4 )2 was prepared by catalytic hydrogenation of Ni(CR)(ci0 4 )2 using platinum oxide as cata lyst as described elsewhere ( 3 6 ). Anal. Calcd. for NiCisH2 sN4 Cl2 08 :
C, H, 5-00; N, 10.77. Found: C, 3^.70; H, 5. 17j N, 1 0 .85 .
Preparation of meso-2.12-dimethyl-3,7,11,17-tetraazabicyclo-
[11. 3 . 1]heptadeca-l(l7), 13; 13~triene monohydrate. meso-CRH'Hg0. -
Ni(m-CRH)(C104 )2 was dissolved in water and a 6:1 molar ratio of cyanide ion was added as is described elsewhere ( 3 6 ). The separation of dl and meso isomers is also the same as described by Karn and Busch. Anal.
Calcd. for Ci 5 H28 W4 0 : C, 6h. 3; H, 10.0; N, 20.0. Found: C, 6 3 .1 ;
H, 1 0 .3 ; N, 1 8 .5 .
Preparation of Fe(m-CRH)x2, X = Cl ,Br ,1 , OAc , and C10 4 . -
Under an inert atmosphere, 0. 01 mole of the appropriate iron(ll) salt was dissolved in 50 ml of absolute ethanol or acetonitrile at i+O0, and the solution was filtered if it was not initially clear. In another container, 0.01 mole of the ligand (meso-CRH* HoO) was dissolved in a small portion of the same solvent. The colorless ligand solution was added slowly to the ferrous salt solution. After stirring the solution for several minutes, the temperature was allowed to return to room temperature. A crystalline product formed after several hours. This 37
product was recrystallized from either ethanol or acetonitrile. The approximate yields and the analyses are in Table 1.
Preparation of Fe(m-CRH)(NCS)g. - In an inert atmosphere,
O.OO5 m of Fe(NCS) 2 *6H 2 0(l.Ug) was dissolved in 50 ml of CH3CN and the solution was filtered. In a separate container, 0. 01 m of m-CRH*H 2 0(2. 8 g) was dissolved in 25 ml of CH 3CN. The ferrous salt solution was added dropwise, and the color of the resulting solution became deep green as the ferrous solution was added. The solution was then filtered and its volume reduced. After several days, dark green crystals were removed by filtration and recrystallized from aceto n itrile. See Table 1 for analysis.
Preparation of Fe(m-CRH)(OAc)(PF6). - 1. Og ( 3 . 6mmole) of meso-
CRH*K20 was dissolved in 75 ml of abs. ethanol and 0. 6g (3 .6 mmole) of
F e(0A c)2 was added to the solution, and it was warmed to *J±0° C. The solution turned pale green and 2. Og ( l 0 . 3 mmole) of NH4 H’e was added.
The solution began to darken and finally turned brown. It was heated for an additional hour at 40°, filtered and allowed to cool to room temperature. A green crystalline product was obtained and recrystal lized from absolute ethanol. See Table 1 for analysis and approximate y ie ld .
Preparation of Fe(m-CRH)(N3)2. - Under an inert atmosphere,
0. 7g (3 .5 mmole) of FeCl 2 *^H20 was dissolved in 50 ml of warm absolute ethanol and filtered. To the pale green salt solution, was added 58
1 . Og (5 .5 mmole) of meso - CRH* H20 dissolved in .20 ml of abs. ethanol.
The resulting solution immediately turned yellow. At this point,
O .Jg (9 mmole) of NaK 3 was added and the solution became cloudy and
turned to a yellow-green color. The solution was filtered and allowed
to cool to room temperature. An orange crystalline product was
obtained. It was recrystallized from absolute ethanol. See Table 1
for analysis and yield.
Preparation of [Fe(m-CRH)x 2 ]Y, X= Cl and Br and Y = C10 4 and
BF4 . - 1 mmole of the appropriate complex Fe(m-CRH)x 2 was dissolved
in warm (T = ^0°c) absolute ethanol while N 2 was constantly passed over
the solution. To the resulting yellow solution, 0 .5 ml of conc. HY was added dropwise. Air was bubbled through the clear yellow solution,
and an immediate color change was observed. The solution of dichloro-
complex became deeper yellow while that of the dibromo-comp lex became
red brown. The solutions were allowed to cool to room temperature and
were filtered, yielding crystalline products. Analyses are in Table 1.
Preparation of [Fe(m-CRH)X2] X, X = Br and I . - Under
N2, 1 mmole of the appropriate Fe(m-CRH)x 2 was dissolved in warm
(T = 1j-0°C. ) absolute ethanol. To this yellow solution 0. 5 mmole o f the
appropriate halide (Br 2 or I2) was added and the solution was stirred.
The addition of the halide immediately changed the appearance of the
solution; and, after several minutes, the solution was cooled. At 39
this point, a dark brown powder separated and was isolated by filtering.
See Table 1 for analysis.
Preparation of 5. 7, 7.12. lA, 1^-Hexa.methyl-l, U. 8 ,11-tetraazacyclo- tetradeca-^,ll-dienenickel(ll) fluoroborate, Ni(l, 7~CT)(bF4)2. - The lig a n d , 1, 7~CT*2HBF4 , was p re p a re d by a method s im ila r to t h a t developed by Curtis ( 8 3 ). In this procedure, 0. 5 moles of ethylenediamine is added to 250 ml of absolute ethanol, and 1.0 mole ( 88 g) o f HBF4 i s added slowly while the solution is kept in an ice bath. To the 0. 5 mole o f en 1 2HBF4 produced, another 0 .5 mole of en and an excess over
2 moles of acetone is added in an ice bath. The white powder produced on cooling is the ligand 1,7~CT*2HBF 4 (yield ~30^).
A 52. 5g (0. 21m) sample of Ni(0Ac) 2 ’1+H20 was dissolved in 2 1. of warm (~A0°) methanol and 95g (0. 21m) of 1,7-CT*2HBF 4 was added slowly. After several hours of stirring, the yellow crystals of
Ni(l,7-CT)(BF 4 )2 were isolated by filtratio n and approximately lOOg
(pkfj) of the product was obtained.
Preparation of meso-5.7.7.12,14,lU-Hexamethyl-1, k*8,11-tetraaza- cyclotetradecanenickel(ll) fluoroborate, Ni(meso-l,7~CTH)(BF4)p. - The reduction of the amine groups of Ni(l,7-CT)(BF 4 )2 was accomplished by a method sim ilar to that developed by Curtis ( 8 ^-) and the separation of isomers was accomplished as reported by Warner and Busch ( 2 2 ). A
30g (0.059m) sample of Ni(l,7~CT)(BF 4 )2 was dissolved in 800 ml of
H20 at 70°, and 3«6g (0.095m) of NaBH 4 and 0.95g of Borax were mixed TABLE I
Elemental Analyses and Approximate Yields of Iron meso-CBH Complexes
Theory Found
Compound Approx. Fe C H N X Fe C H N X Yields ($ )
Fe(m-CRH)C12 70 Hf.lf U6 .3 6 .U3 Ilf. if 1 8 .2 Ilf. 1 if5. 6 o 6 . lf8 Ilf. 10 18 .99
Fe(m-CRH)Br2 70 1 1 .7 3 7 .6 5 .2k 11.7 33. if 11.5 3 7.31 5.16 1 1 .6 8 33.25
Fe(m-CRH)I2 70 1 0 .0 3 1 .6 ^•33 9 .8 2 if3 - 0 1 0 .1 3 2 . 0lf If. 32 10 .0 5
Fe(m-CRH)(NCS)2 ko 12.9 If 7.0 5-99 19-3 if 7.2 9 6 . Ilf 1 9 .0 2 3 0 00 [Fe(m-CRH)OAc]PF6 70 V 5.56 10.72 39.25 5-91 1 1 .1 2
Fe(m-CRH)(OAc)2 50 51. If 7 .0 1 2 .6 if 7.9 2 6 .7 2 1 1 .9 6
Fe(m-CRH)(C104 )2 80 3k. 8 k,8k 1 0 .8 32.1 5 If. 77 1 0 .8 1
Fe(m-CRH)(N3 )2 IfO Iflf. 75 6.h6 31+. 80 ifif. 11 6 .3 2 35.53
[Fe(m-CRH)C12 ]C104 80 3 6 . 83 5-12 11. If6 3 6 . if9 5-17 10.97
[Fe(m-CRH)C12 ]BF4 80* 3 7 .8 0 5-35 11.76 37. 88 5 .3 8 1 1 .5 8
[Fe(m-CRH)Br2 ]BF4 80 3 1.87 lf.lf3 9 .9 2 3 1 .8 7 If. 65 9 .8 9 t- H H [Fe(m-CRH)Br2 ]C104 80 • if. 33 9 .7 0 31. IfO If. If7 9-37
Fe(m-CRH)I3 50 2 5 .8 3-72 8 .0 1 5if.5 2 5 .8 5 3 .6 7 8 .2 9 56.89 kl
together and added slowly (over a lj -5 minute period) to the solution, which was stirred vigorously. If the addition was performed rapidly, a black insoluble substance (nickel boride) appeared. After all of th e NaBH 4 was added, the solution was heated and stirred at 7 0° f o r an additional hour. The pH was then adjusted to 2 or 3 with k&J0 HBF4 , and 300 ml of acetone was added. The solution was filtered while hot, and the total volume was reduced to 600 ml. Yellow-orange crystals were obtained. The first crop of crystals (~10g) was pure
Ni(meso-l, 7“CTH)(BF4)2. After the first crop was obtained, the filtrate was boiled and made basic (pH ~10) with NaOH. A 5g sample of
Na2 C2 04 was added, and the solution was boiled for an hour. The solution was then allowed to cool, and the insoluble purple
{(Ni(rac-1,7-CTH))2 C2 04 }(BF4 )2 was removed, leaving a yellow filtrate.
The excess oxalate was precipitated by boiling the solution and adding
CaCl2 until no additional white CaC 2 04 precipitated. Filter aid
(Super-Cel) was added, and the white CaC 2 04 was removed by filtration.
The filtra te now contained pure Ni(meso-l, 7-CTH)(BF 4 )2 and was obtained by adding several ml of b&j0 HBF4 and reducing the volume of the solution. The total yield was 15~l6g (50-60^).
Preparation of meso-5, 7,7,12,1^-, 1^-Hexamethyl-l, U, 8 ,11-tetra- azacyclotetradecane dihydrate, meso-1,7~CTH* 2H 2 0. - The ligand was abstracted from the Ni2+ ion after the manner of Warner ( 8 5 ). A lOg (.019m) sample of pure Hi(meso-1,7-CTH)(bF^Jp was dissolved in k2
I4-OO ml of 1:1 ethanolrwater solution. A large excess of NaCN(~30:l,
NaCN:Ni2+) was added slowly to the solution. The solution undergoes
several color changes, beginning with a deep orange and proceeding to an almost colorless color. The solution was made extremely basic b y adding 20-h0g of NaOH, and a white precipitate formed and was isolated by filtration. The volume of the filtrate was reduced, and more of the white ligand was obtained upon refrigeration. Total yield, 3~5g (65-95$).
Preparation of Fe(m-1, T~CTH)xg, where X=C1 ,Br , I , and NCS . -
The preparation of all Fe(ll) complexes was carried out in a dry oxygen-free box. A sample containing 0.01m of FeX 2 *6H20 was d is s o lv e d
in 50 ml o f CH3CN while 0.01m of m-1, 7-CTH* 2H2 0(3. 2g) was slurried in a separate portion of 25 ml of CH3CN. The two were mixed. The halide
slurries soon formed brown precipitates. These precipitates were isolated by filtration, placed into fresh acetonitrile and the mixture was heated. These slurries were filtered, and pale green solutions were obtained. The volume of these solutions were reduced and, after
standing, very pale green or blue precipitates were obtained. In the thiocyanate system the complex is soluble and a solid residue was discarded. The filtrate is a deep green and a dark green crystalline product separated slowly from this filtrate. Yields were on the order
o f hOf0. Analytical data are presented in Table 2. TABLE I I
Elemental Analyses for the Iron meso-1,7-CTH Complexes
Theory Found
Compound CH NX CH NX
Fe(m-1,7-CTH)C12 U6- 71 8.76 1 3 .6 2 1+7 .0 1 9.03 13.77
Fe(m-l,7-CTH)Br2 38 . 1*0 7 .2 0 11. 20 3 2 .0 0 3 8 . 21 7.1+1+ 1 1 .1 0 3 2 .3 0
Fe(m -1, 7 “CTH)i2 32. 31* 6 .0 6 9.1*3 32.32 6 .2 6 9-1+8
Fe(m-1,7-CTH)(NCS)2 1+7-37 7.89 18 . 1*2 1+7.6 2 7.91+ 18 . 1+1+
Fe(m-1,7-CTH)(C104 )2 * 3CH3CN 39-89 6.7 9 Ik. 81 1+0. 51+ 6 .8 0 H+.68
Fe(m-1,7-CTH)(CN)2 55.10 9 .1 8 21.1+3 ll*. 23(Fe) 55-20 8 .75 21. 1+1+ ll+. 23 (Fe)
Fe(m-1,7-CTH)(C104 )3* 2CH3CN 3 3 .3 1 5.82 1 1 .6 6 33-89 6 .1 1 1 1 .1 0
[Fe(m-1 ,7-CTH) C12 ] BF4 38.55 7.23 11. 2l* 35.05 6 .35 1 0 .1 9
[Fe(m-1,7-CTH)C12 1C104 . 3 7.62 7.05 10.97 37.58 7.1+8 10 .9 9 Preparation of Fe(m-1, T-CraKciC^)?’3GH 3 CN. - A 0 .0 1 m sample of Fe(ci0 4 )2 *^H 2 0(3. 2g) was dissolved in 30 m l o f CH3CN, y i e l d ing a pale green solution. A 0. 01m sample of m-1, 7-CTH* H 2 0(3-2g) was added slowly and a dark green slurry resulted. This slurry was filtered, after stirring for several minutes, and the pruple filtrate was retained. The volume of this solution was reduced and deep purple crystals formed; these were isolated by filtration. The crystals were recrystallized from CH3CN. Y.ield ~JLOaj0 (0. 5 g ).
Preparation of Fe (m-1, 7-CTH) (CN)2. - A 0.005m («33g) sample of Fe(m-1,7-CTH) (C10 4 )2*3CH3CN was dissolved in 50 ml of EtOH and excess
NaCN was added. After stirring for several minutes, the solution turned purple. The volume of the solution was reduced, and a dark brown crystalline material was filtered out. The yield was not deter mined; however, it was probably at least 80 ^.
Preparation of Fe(m-1,7-CTH)(C104)3- 2CH 3 CN. - To the filtrate from the preparation of Fe(m-1,7~CTH)(C10 4 )2*3CH 3 CN, a few drops (5~T) of 50/, HC104 were added, and air was bubbled through the solution.
Ether was then added slowly to the solution to the cloudpoint. The solution was cooled in the refrigerator. A tan precipitate was filtered out. Yield unknown, because the amount of starting material was unknown. ^5
P r e p a r a tio n o f [Fe(m -1, 7~CTH)C121Y, w here Y=C104 ~ o f BF4 ". -
A 1 mmole sample of Fe(m-1, 7"CTH)C12(0. kg) was dissolved in 50 ml of
CH3CN and approximately 5 drops of HY was added. Air was bubbled through the solution and ether was added to the cloudpoint. The solution was then cooled and yellow crystalline material was filtered out. Yields of 80-90^ were obtained.
Preparation of 9 , J, 7, 12,14,1^-Hexamethyl-l, k, 8,11-tetraazacyclo- tetradeca-l+j 11-diene dihydrogen iodide (1, 7-CT* 2Hl). - The ligand,
1, 7-CT* 2HI, was prepared using a method sim ilar to that of Curtis ( 8 3 )*
A 0. 5 mole sample of ethylenediamine was dissolved in 2^0 m l o f absolute ethanol, and 1.0 mole of HI was added to the solution in an ice bath. To the 0. 5 moles of en* 2HI produced, another 0. 5 mole o f ethylenediamine, and an excess of 2 moles of acetone was added. The white powder produced was 1,7_CT* 2HI.
Preparation of Fe(l,7~CT)lg*2CH 3CN. - A slurry of 1. 0g ( 0. O57 mole) of 1,7-CT*2HI was prepared in 100 ml of warm (~^5°) acetonitrile.
Neither of the reactants is very soluble in acetonitrile. During the stirring of the warm slurry for a day, it became maroon and was filtered.
Maroon crystals of [Fe(l, 7-CT)(ch 3 CN)2 ']I2 were formed upon cooling.
Y ield 10- 30^. Anal. Calculated for FeC2 oH3 aNeI2 : C, 35- 8 ,* H, 5 - 6 6 ;
N, 12.5; I, 37.5. Found: C, 35-6,* H, 5-73; N, 12. Uj I, 37-1. U6
Preparation of Fe(l,7~CT)l2. - A sample of rFe(l, T^CTj-
CcHsCNJalla was put into a drying pestle for one day at 100° u n d er vacuum, and a brown compound was o b ta in e d . Y ield ~10Cy,. A nal. C alcu lated for FeCisH 32N4 l2: C, 33* 2j H, 5*^5j N, 9.50* Found: C, 32. Tj
H, 5* 67; N, 9 .^ 2 .
Preparation of Fe(l, 7~CT)Fe(HCS)4' 2CH 3CN. - The ligand
I, 7-CT*2HC10 4 was made as reported by Curtis ( 8 3 ). In the dry box
1. Og (3 . 6mm) of Fe(sCN) 2 *6h20 was dissolved in 5° ml of acetonitrile and 1.65g (3* 6mm) CT\2HC104 was added. After stirring for several hours at !+0°C, maroon crystals formed, and upon cooling they were filtered and recrystallized from acetonitrile. Anal. Calculated for
Fe2 C24 H38 N icS4 : C, 1+0.8; H, 5-38; N, 19-8. Found: C, 1+1.2; H, 5*21;
N, 1 9 .8 .
Physical Measurements
Visible and near-infrared absorption spectra were obtained on a Cary Model ll+ recording spectrophotometer. Many of the spectra were obtained with methanol, nitromethane, and acetonitrile solutions.
Some of the spectra were run on solid samples as Nujol mulls impreg nated on filter paper. Infrared spectra were obtained on a Perkin-
Elmer Model 337 recording spectrometer, using (mainly) Nujol mulls and potassium bromide pellets.
The conductances of the complexes were obtained using an
Industrial Instruments Model RCl 6B2 conductivity bridge and a conduct- hi
ance cell with a constant of 2. Il 6 cm 1. Conductances were determined at 25° at 1000 cps on samples ~10 3M in concentration. Methanol was dried by refluxing several hours over magnesium turnings and then distilling. Acetonitrile and nitromethane were purified according to the procedure of Fieser and Fieser ( 8 6 ).
The magnetic susceptibilities of solid samples were determined at room temperature by the Faraday method on a balance set up by
L. F. Lindoy and Vladimir Katovic (87).
Elemental analyses were performed by Chemalytics, Inc., Bern hardt, and Crotian Laboratories. Some of the analyses were performed by Mr. Pete Kovi in this department on a Coleman Model 29 nitrogen analyzer. A Kewanee Scientific Equipment controlled atmosphere glove box was used in the preparation of the Fe(ll) complexes. Baker analyzed ultra-pure nitrogen was passed continuously through columns of magnesium perchlorate, molecular sieves, and copper turnings at
500°C to remove oxygen and water.
The Mossbauer spectra were obtained on an electromechanical constant-acceleration spectrometer built by W. L. Pillinger and
J.A. Stone at the U. S. A. E. C. Savannah River Laboratory, Aiken, South
Carolina (see Figure j ) . The gamma radiation source, which was obtained from the New Englander Corporation, 575 Albany Street, Boston,
Massachusetts 02118, consisted of 5mC of 57Co diffused into a thin co p p er d is c . A R e u te r-S to k e s, RSG-JOA k ry p to n -m eth an e f i l l e d p ro - kQ
portional counter was used to obtain the transmission spectra. Other components were used to lim it the range of detection to lk. 37KeV. A
RIDL Model 3k-12B kOO Channel multi-channel analyzer operating in time mode and synchronized with the velocity modulated radiation source was used to count the transmitted radiation. The National Bureau of
Standards Standard Reference M aterial 725, a single large crystal of sodium nitroprusside, Na 2Fe(CN)5N0’ 2H20 ( 8 2 ), was used as the velocity calibrator along with a pure sample of o:Fe 2 03. The calibration was checked periodically and was 0. Ok33mm/sec per channel.
Thin disc-shaped absorbers were made by cutting a hole approxi mately l|cm. in diameter in an aluminum strip. The crystalline samples were packed into the discs and covered with thin (ca. 1 mil) Mylar tape. The absorber thickness was normally less than 75fflg/cm2.
The spectra were obtained by counting until the peaks were statistically well above the background. The minimum difference between the baseline and the peak obtained was on the order of 1 0 ,0 0 0 c o u n ts.
Low temperature measurements were made using a special Israeli cryostat equipped to attach to a liquid dewar. Liquid nitrogen was circulated through the sample compartment in the cryostat by applying air pressure. The temperature was measured by a germanium resistance thermometer. The spectra were observed at the boiling point of the liquid nitrogen. The cryostat was fitted with a heater which could be used to vary the temperature, however, this was not used. RESULTS AND DISCUSSION
A number of new Fe(ll) and F e(lll) complexes of the macro-
cyclic ligands meso- 2 ,12-Dimethyl-3, 7, 11,17-tetraazabicyclo[ 11.3 * 1]
heptadeca-l(l7),13,15-triene(m-CRH, Structure IX), me_so.-5, 7, 7,12, ll+, lU-
H exam ethyl-1, it-, 8 , 11-tetraazacyclotetradecane (m-1, 7-CTH, structure VIl),
and 5, 7, 7,12, lU, l^-Hexamethyl-l, k, 8 , ll-tetraazacyclotetradeca-4,11-
d ie n e ( l, 7_CT, s tr u c tu r e v) have been prepared and characterized. Some have unusual structures which merit detailed description. The complexes were characterized using elemental analysis, infrared spectroscopy, molar conductivity, magnetic susceptibility, electronic spectroscopy,
and Mossbauer spectroscopy. The Mossbauer spectra of several additional compounds were measured and compared to those of the complexes
characterized here. Several interesting relationships were observed and are discussed here.
Complexes of Fe(ll) and F e(lll) with the Macrocyclic Ligand m-CRH
Several Fe(ll) and Fe(lll) complexes of the ligand meso-2,12-
d im e th y l- 3 , 7 ,1 1 ,17_tetraazabicyclo[ll. 3 . Ilheptadeca-l(l7),13> 15"
triene (m-CRH, structure IX) have been prepared, and their synthesis
and characterization are discussed here. The complexes with the
stoichiometry Fe(m-CRH)x2, where X~=C1 ,Br” and I~, are a ll high spin,
square pyramidal, five-coordinate complexes in the solid state. When
X=N3 or OAc , the complex has a high spin, six-coordinate, pseudo-
k9 50
octahedral structure. When X=NCS , the complex has a low spin, pseudo- octahedral structure. The F e(lll) complexes rFe(m-CRH)x2]Y where
X=C1 o r E r and Y=C10 4 o r BF4 , are high spin, six-coordinate, pseudo-octahedral complexes. Infrared spectra, electronic spectra, molar conductances, magnetic susceptibilities, and Mossbauer spectra are used to characterize these new complexes.
Ligand Synthesis. - The ligand, meso-2,12-dimethyl-3>7 311317“ tetraazabicyclo; 11. 3 . I]heptadeca-l(l7)j 13j 15- t r i e n e 1-h y d ra te
(m-CRH*H 2 0) is synthesized by using N i(ll) as a template. One mole of
2, 2 -diacetylpyridine is condensed with one mole of 3 -diaminodipro- pylamine in the presence of one mole of N i(ll) ion as reported by
Karn and Busch ( 3 6 ). The imine functions are then hydrogenated by using a heterogeneous catalyst, Pt02. It has been shown by Karn and
Busch that once the imine functions are reduced, the nickel can be abstracted from the ligand by use of strong coordinating groups such as the cyanide ion ( 3 6 ).
Ni(CR)2+ Ni(CRH)2+
Ni(CRH)2+ + 6 c t " ------>■ Ni(CN)s 4 " + m-CRH* H20 51
Since the hydrogenation yields two isomers, the meso with both methyl groups on the same side of the plane of coordination and the racemic for which the methyl groups are on opposite sides of the plane, care must be taken to separate the two. The predominance of the more insoluble red meso isomer (~ 90 ^) makes it easier to obtain in large amounts; therefore, the meso isomer has been used throughout this study. The total yield of pure ligand is quite small, in part because of the number of steps involved in the synthesis.
Synthesis of the Iron(ll) Complexes, of meso-CKH. - The iron(ll) complexes are all prepared in a protective nitrogen atmosphere to pre vent hydrolysis and oxidation of the iron(ll). When the solutions of the complexes are allowed to stand in air, oxidation occurs, as evidenced by gross changes in the electronic spectra of the solutions.
In all of the solutions Fe(ll) seems to be oxidized in Fe(lll) within a few minutes.
In general the Fe(ll) salt hydrates are used, as starting materials in the syntheses, and the water of hydration does not inter fere with the reactions. A variety of dried, oxygen-free solvents are used in these syntheses. Both protic and aprotic solvents seem to work well. Iron hydroxides are easily formed, however, if the solvents are not deoxygenated properly.
The synthetic procedure requires dissolution of the Fe(ll) salt and the ligand in separate portions of the warmed (~l+ 0°) solvent and 52
then the mixing of the two solutions. The mixing sequence does not
seem to affect the course of the reaction. The mixture changes color
immediately and is allowed to cool to room temperature, whereupon the
product crystallizes. The crystals are collected and recrystallized
from the same solvent. Recrystallization of the halide complexes is
often useless because the recrystallized product sometimes appears to be less pure than the original product. This is believed to be be
cause of the ease of contamination with iron hydroxides.
In the preparation of the Fe(ll) dithiocyanato-complex, the
reaction sequence is important. The ferrous salt solution should be
added to the ligand solution to prevent formation of the Fe(NCS )4
species. A 2:1 mole ratio of ligand to iron is used also to prevent
Fe(NCS)4 anion formation.
The solids are normally stored under nitrogen; however, most
of the complexes are stable as solids in air over periods as long as
several months. Slow decomposition (thought to be because of the water
in the air) is observed in all samples that have been exposed to the
air for extended periods of time. The thiocyanate complex seems to decompose as quickly as one day, while the others are stable for much
longer periods. All the complexes are sufficiently stable in the air to permit the normal measurement of infrared spectra and magnetic 53
susceptibilities. However, some of the Nujol mulls for the infrared spectra must be prepared in the dry box since a slight darkening of the complex is observed upon grinding in air.
Synthesis of the Iron(lll) Complexes of meso-CKH. - The iron(lll) complexes are made from the previously prepared iron(ll) complexes in several ways. The procedure used in the preparation of the majority of the iron(lll) complexes begins with the dissolution of the appropriate iron(ll) compound in dry, 02-free ethanol in a nitrogen atmosphere. Acid is added to prevent solvolysis and to provide a counteranion, and air is bubbled through the solution. The solution darkens immediately. Some time later, crystals begin forming, and these are filtered from the solution. When acid is not added, the dark red-brown color of iron(lll) hydroxide is seen, and the product is a mixture of this and the white ligand.
Other oxidations are performed by adding elemental halogen to the halide complex. Iodine is added to an ethanolic solution of the iodide complex under a nitrogen atmosphere, and the solution immediately turns a deep brown. A dark brown crystalline iron(lll) triiodide complex is obtained. Bromine seems to oxidize the bromide complex in the same way; however, an acceptable analysis has not been obtained for the tribromide complex. Chlorine seems to be too strong am oxidizing agent in ethanol and white decomposition products are o b ta in e d . The Iron(II) Complexes of meso-CRH. - Iron(II) forms three different types of complexes with meso-CRH. Some are normal high spin, six-coordinate pseudo-octahedral complexes, others involve high spin five- coordinate structures (generally rare), and one appears to have a planar four-coordinate structure.
The six-coordinate complexes include the azide, the thiocyanate
(which is low spin), and the two acetate derivatives. The infrared spectra of these complexes are sim ilar to a ll other meso-CRH compounds and w ill be discussed here. The anions afford the only major differ ences in the several spectra and the position of these w ill be discussed where necessary.
A few bands from the infrared spectra of these complexes along with those of other appropriate compounds are listed in Table 3* The infrared spectrum of the free ligand is characterized by a broad N-H stretching band around 3200 cm 1, a doublet in the double bond region at 1590 and 1570 cm 1, and a fingerprint region (lOOO-l^OO cm x) with many sharp bands. (See Figure ll). The reason for the broad N-H stretch ing mode in the spectrum of the free ligand is thought to be because of hydrogen bonding since sharp bands are observed in the metal complexes of this ligand. The double bond region ( 155O-I65O cm x ) i s complicated by the pyridine ring-breathing modes which occur as a characteristic doublet at 1570 and 1590 cm x. These bands are also present in all the complexes. Karn and Busch used intensity relation- .50 .30b .60 Absorbance.20 .40 gure 1. nfar r f m- C12 upper) p p (u 2 1 )C H R -C (m e F of tra c e p S d re fra In 11. e r u ig F .70 . 3.0 2.5 3000 *vv 4.0 aeegh (microns) Wavelength CRH •H0 l er) w (lo • H20 H R -C m 2000 5.0 cm 6.0 1500 7.0 1400 8.0 "loo 1200 -.30 -.60 -.50 -70 .10 .20 00 .40 gure 1. con') CRH •H0 upper) p p (u •H20 H R -C m 't) n o (c 11. e r u ig F .60 .50 Absorba.30 nee .20 0 .40 .70 . 0 1300 - op , 890 , , loop 10.0 em_CHC2 l er) w (lo _-CRH)C12 Fe(m 12.0 aeegh (microns) Wavelength 14.0 16.0 1 I 0 c 500 cm 600 18.0 » (microns) __ i__ _
1 ___ 1 1 1 1 20.0
22.0 24.0 .70 .50 .60 40 30 20 1 J V ON 57
ships to demonstrate that the band at 1570 cm 1 contains the imine stretching mode in the nickel(ll) complexes of both the reduced and the unreduced ligands ( 3 6 ).
The spectra of all derivatives are similar except in the region where bands due to anions appear. The N-H stretching vibrations appear as a sharp doublet upon complexation. This doublet is due to the two different types of N-H groups in the ligand structure. These bands are all sim ilar from complex to complex, changing in a slight but unpredictable fashion.
The characteristic doublet in the imine region also changes slightly from complex to complex, however, the reason for these differences is not known. Neither the N-H nor the C=N regions appear to have much dependence upon oxidation state, coordination number, or ligand field strength of the anion.
The spectral region from 1000 to 1^00 cm 1 is quite complex but a ll of the complexes show sim ilar absorptions except when the anion has features that appear in this region, e.g., the perchlorate and the tetrafluoroborate anions. With very pure compounds, extremely sharp peaks are obtained in this region. The infrared spectra show that the ligand is present, and that it is unchanged during the various reactions.
The thiocyanate in the complex, Fe(m-CRH)(NCS )2> is bonded to the iron through the nitrogen of the thiocyanate as evidenced by its infrared spectrum. (See Figure 1 2 ). According to Burmeister ( 8 8 ), gure 1. nfar r f m- ( )2 S C )(N H R -C (m e F of m tru c e p S d re fra In 12. e r u ig F .50 .60
Absor bcnce.20 0 4 . odr 4000 .70 col . 3.0 2.5 ------L_ 3000 4.0 aeegh (microns) Wavelength 2000 5.0 cm 6.0 1500 7.0 1400 8.0 .30 .10 oo .60 .20 .70 40 50 00 gure 1. t) Fe( CRH)NCS)2 S C )(N H R -C (m e F 't.) n o c ( 12. e r u ig F .60 .50 .301 Absorba nee .20 o.or .40 .70 00 1300
8.0 1000 10.0 12.0 aeegh (microns) Wavelength 4 0 60 18.0 16.0 14.0 20.0 m 22.0 24.0 oo .60 .50 .40 .20 70 30 60
N-bonded thiocyanate complexes have bands in the following regions:
C sN str 2085-2120 cm " 1
C-Sstr 780-860 cm 1
The CsN stretch for this complex is a doublet at 2101 and 2114 cm 1 while the C-S stretch is also a doublet at 815 and 801 cm 1, therefore, it is concluded that the thiocyanate is N-bonded.
The thiocyanate complex is a low spin six-coordinate complex.
Its low conductivity ( 26mhos) in nitromethane (89) suggests that it is a six-coordinate species in nitromethane but that it may react to a small extent with the solvent. The visible spectrum of this compound is similar to those of other low spin six-coordinate Fe(ll) complexes with bands at 1 5 .6 and 20.4 kK in the mull spectrum. These bands are assigned to the transitions 1Aig 1Tig and 1Aig XT 2g. In view of the fact that the lower energy band is quite weak, the alternate possibility that it may be due to the spin forbidden transition
1Aig -> 3Tig cannot be eliminated (90). The first assignment seems the more probable because low spin Fe(ll) spectra should be sim ilar in appearance to low spin Co(ill) spectra which normally show two or three d-d bands and the Dq value should be considerably lower than that of Co(lll) (90). Using the equations derived by Wentworth and
Piper (9l), which are discussed on page 89 3 and using a C value of
2500 cm 1 which is approximated from data in the Fe(m-1,7~CTH)x 2 com plexes, a Dq o f 1810 cm 1 is obtained for this complex. (Table III). TABLE I I I
Some P h y s ic a l P r o p e r tie s o f Fe (m-CRH)2+;> 3+ Complexes
Compound C olor A(mhos) N -H strb C=Nstrb Me OH ch 2no 2
Fe(m-CRH)C12 y e llo w 5.23 110 66 3215 1597 3155 1575 Fe(m-CRH)Br2 y ello w 5- 11 107 68 3215 1597 3185 1575 Fe(m-CRH)l2 y ello w 5.05 11U 55 3185 1597 3155 1575 Fe(m-CRH)(WCS)2 b la c k 0 .8 (2 7 3 c ) 100 27 3195 1597 1575 Fe(m-CRH)(N3 ) 2 orange 5 .5 0 92 50 3 2 3 6 (b r) 1597 1580 Fe(m-CRH)(0Ac)2 orange 5-31 80 36 3289 1600 )tjr 3221 1570) 3165 3 O8 O TFe (m-CRH) 0Ac]PFs p a le 5-11 9*4- 75 3279 l 600b r g re e n 3175 Fe(m-CRH)( 01 04)2 brown 0. 9 6 ( 314-0° ) 210 3185 1600 1582 [Fe(m-CRH)Cl2 ] C104 y e llo w 5.79 115 3226 1603 3155 1580 [Fe (m-CRH) C12 ]BF4 y ello w 5.85 103 3 2 to 1605 3195 1580 [Fe(m-CRH)Br2 lC104 re d 5-73 116 3215 1605 1580 [Fe(m-CRH)Br2 ]BF4 re d 5-76 117 "3216 I 6 O5 1580
Fe(m-CRH)l3 brow n 105 3215 1592 3185 1570 aAll solutions ~ 10 3 m olar. b U n its cm 1. CMagnetic susceptibility at 25°C (units 10 6 cgs units). ^The conductivities in nitromethane were repeated twice and the results were the same within experimental error. 62
The remaining six-coordinate Fe(ll) complexes (the azide, acetate, and the hexafluorophosphate acetate) are high spin with magnetic moments of 5*3 to 5*5 B. M. The diacetate and diazide complexes are uni-univalent electrolytes in methanol and exhibit conductance values between the ranges indicative of nonelectrolytes and uni-univalent electrolytes in nitromethane. The value for the diacetate is quite low indicating a predominant six-coordinate structure. The diazido-complex has a particularly high value (50mhos)
(89). This suggests that an equilibrium occurs wherein one anion is either displaced by solvent or lost outright to produce a mixture of five-coordinate and six-coordinate structures. These complexes have rather low solubilities (~ 10- 3m), and this limits the utility of conductivity studies in nitromethane. The complex^’ [Fe(m-CRH)OAc]PFe is a uni-univalent electrolyte as expected^ however, its value ( 75mhos) is slightly lower than would normally be expected for such an electro ly te in CH3N02. The structural assignments are strongly supported by the nature of the Mossbauer spectra for these complexes. This is discussed in detail in a later section.
The visible spectra of these compounds correspond to those of normal high spin pseudo-octahedral complexes of Fe(ll) (See Table IV).
They exhibit a broad d-d band in the region 10, 000-11, 500 cm 1 w hich shows none of the splitting expected to derive from the tetragonal distortion caused by the anions. (See Figure 1 3 ). This band is 63
TABLE IV
Electronic Spectra of Fe(ll) and F e(lll) m-CRH Complexes
Compound d-d Bands (kK) Charge Transfer (kK)
Fe(m-CRH)Cl2 10. 7a (<==6) 3 1 . Osh 11.3
Fe(m-CRH)Br2 10. 7a (€=7) 26 . Ij-sh 11. 3b
Fe(m-CRH)l2 1 1 .3 a ( Fe(m-CRH)(NCS)2 10. j a> C(f=l0 ) 25. ^t-sh 15. 6b 20. 11-° Fe(m-CRH)(OAc)2 1 0 .7 (€ = 8 )a 30.7(6=660) [Fe (m-CRH) OAclEFs 11. 6 (€=1 0 )a 18 . 5 sh Fe (m-CRH) (N3 )2 Fe(m-CRH)(C104 ) I2 . 0a ,c 3 0 . 7sh 17. 8 sh 20. 6 sh 27. 8 sh [Fe(m-CRH)C12 ]C104 30.7(6=2050) [Fe(m-CRH)C12 ]BF4 30.7(6=2180) [Fe(m-CRH)Br2]C104 30.7(6=2000) [Fe(m-CRH)Br2 ]BF4 30.7(6=1800) Fe(m-CRH)l3 2 8 .2 (s tro n g ) & In methanol solution. ^Reflectance spectra in Nujol mull. c Solvated species. LOG (€) 3 El roni ra o Fe(m-CRH)X2 = I f " r B , " l C X= 2 X ) H R C - m ( e F of a tr c e p S ic n o tr c le E 13. e r u g i F 0 3 . . 5 7 - - 350 27 AEUBR CKKD WAVENUMBERS WAVELENGTH CNMD CNMD WAVELENGTH 2335 31 915 19 550 orde rid lo h C Iodide omi e id m ro B 750 11 1250 & 65 shifted to higher energy in the solid state spectra. The transition is probably due to a slightly split 5T2g -» 5Eg tran sitio n (See page 81)-). The spectra of all the complexes also have a band in the region be tween 25,000 and 30,000 cm 1, except for rFe(m-CRH)0Ac]PF 6 whose spectrum has a shoulder at 18,500 cm 1. This transition.is assumed to involve metal to ligand charge transfer because of the ligand depend ence of the energy of this band. The unique position of the shoulder in the spectrum of [Fe(m-CRH)0Ac]PFs is consistent with the presumed unusual structure of the complex. The acetate anion probably acts as a bidentate chelat ing agent in this case. This should lead to some folding of the macro- cyclic ligand. Examples of such structures are well known (37,59). The difference in this charge transfer band from that of the diacetato- complex is attributable to the differences in the manner of acetate coordination. The three halide complexes Fe(m-CRH)x2, where X=C1 ,Br , and I , are yellow crystalline solids. They all exhibit molar conductances (Table III) that are normal for uni-univalent electrolytes in methanol; however, in nitromethane their conductance values are slightly lower than is normal for uni-univalent electrolytes. The five-coordinate structure is assumed to occur in solution as well as in the solid, and. the low conductances may indicate that some ion pairing occurs in nitromethane solution. Consequently, both the five-coordinate complexes 66 and the six-coordinate complex Fe(m-CRH) (n 3 )2 act similarly in the weakly coordinating solvent nitromethane. The magnetic moments (Table III) of the halide complexes are rather distinct in a way that supports the possibility that the com pounds are pentacoordinate. The values fall in the range from 5*^5 to 5.23 B.M. The presence of only one axial ligand in a square pyramidal coordination sphere would generate a greater splitting of the 5T2 s t a t e (of octahedral parentage) than that due to the pair of axial ligands in a six-coordinate pseudo-octahedral structure ( 9 2 ). The enhanced splitting should produce a reduction of the orbital contribution to Heff> as has been observed. A second and more diagnostic result of the electric dissymmetry represented by five-coordination is found in the Mossbauer spectra of these compounds. Although this provides the strongest evidence for assigning the coordination number five to these compounds, its discussion is reserved for a separate section. The electronic spectra (solid state) show a single d-d band in th e 1 0 ,7 0 0 cm 1 to 1 1 ,3 0 0 cm 1 range with molar extinction co efficients of less than 10. Five-coordinate compounds of this type have been shown to have a second absorption near 5000 cm 1 (See Page 10]) due to the strong tetragonal splitting (See Figure 1^) (93)* The band in the 5000 cm 1 region is attributed to the transition 5E 5A while the one at the higher energy is attributed to 5E 5B. T h is assignment is based on the fact that the band affected by the in-plane Ojj Tetragonal Figure 14. Results of Strong Tetragonal Distortion macrocyclic ligand (5E 5b) should remain constant from complex to complex while the transition due to the axial ligand (5E 5a ) sh o u ld vary as the axial ligand is changed. The low energy band was not observed in the spectra of these compounds because they are not soluble enough in nitromethane to observe absorptions with such small extinction coefficients. The band that is observed is fairly constant in energy and is therefore assigned to the transition 5E -► 5B. In th e s o lid state the region around 5 ,0 0 0 cm 1 is obscured by overtones of infrared spectral bands. The band at 10,000 cm 1 increases in energy consider ably on going from a solution spectrum to the corresponding solid state 68 spectrum. This could indicate that the six-coordinate species [Fe(m-CRH)X*MeOH] e x i s t s in s o lu tio n . The spectra of all these five-coordinate halide complexes show charge transfer bands between 25, 000 cm 1 and 31 j 000 cm 1. These bands occur as shoulders which are fairly intense. The solvent-free perchlorate salt [Fe^m-CRH)](C104)2 appears to provide an example of a four-coordinate, low spin complex. It is a brown amorphous solid of very low solubility in common solvents. The value of its magnetic moment (0.96 BM) proclaims it to be of the low spin configuration while the unsplit, relatively sharp bands at 1000 cm 1 and 650 cm 1 in the infrared spectrum indicate that perchlorate is not coordinated. From its molar conductance, the compound clearly is a di-univalent electrolyte in methanol. However, from its electronic spectrum in methanol (See Table IV), the solvated form Fe(m-CRH)(ch30h)22+ is present. The mull spectrum shows the presence of low spin Fe(ll) species with bands at 17,800 cm 1g 2 0 ,6 0 0 cm 1 and a t 27,800 cm 1. These bands a re p ro b a b ly due to th e tetragonally split 1Tig manifold, i.e ., the transitions to the ^Aig *These compounds are of special interest because they are among the first synthetic complexes which have the high spin, five-coordinate structure that has been assigned to deoxyhemoglobin and deoxymyoglobin. Prior to the discovery of the new high spin pentacoordinate complexes, there was some tendency to treat the natural products as being novel in coordination structure. It now appears that such configurations are not rare. 69 and xEg. The third band is due to the 1Aig -..1T2g transition (9^). The compound is a dangerous explosive being sensitive both to mechanical shock and to heat. Iron(lll) Complexes of m-CKH. - The iron(ill) complexes of m-CRH have the high spin electronic configuration as evidenced by magnetic moments in the characteristic range from 5*T to 5*9 B. M. (See Table III). These complexes are all assumed to be six-coordinate both in the solid state and in solution. In methanol they are uni univalent electrolytes which is consistent with the existence of the cationic species (Fe(m-CRH)X2) . The infrared spectra of the Fe(lll)m-CRH complexes are sim ilar to those of the other complexes with this ligand as can be seen in Table III. The N-H stretching bands generally occur at slightly higher frequencies for the F e(lll) complexes which may indicate a slightly weaker bond than in the iron(ll)complexes. The shift is on the order of 10 cm 1. The perchlorate and tetrafluoroborate anions exhibit their characteristic infrared bands in the 1000 cm 1 and 6 5 O cm 1 re g io n s . Since high spin Fe(lll) is a d5 ion, its spectrum should be similar to that of Mn(ll). However, since all d-d transitions are spin forbidden and should possess small extinction coefficients, it is not surprising that the transitions are very difficult to observe, in fact, none are observed here (See Figure 1 5 )- Each complex shows only LOG (€) 0 5 ect c Spect f JBF4 F B lJ C ) H R C - m ( e F [ of m u tr c e p S ic n o tr c le E 15 e r u g i F 3 . . 6 0 36 - - 300 228 32 fVNMES EKK3 WflVENUMBERS AEEGH ND T0 CNMD WAVELENGTH 400 500 20 16 700 71 one electronic spectral band, and it is in the charge transfer region. This band is strong and would obscure any d-d band that may exist above 20,000 cm 1. The colors of the complexes are due to the tails of these bands which reach into the visible region of the spectrum. These compounds are all crystalline complexes; the dichloro- complexes are yellow, and the dibromo-complexes are red. The tribromo- and triiodo-complexes are dark brown and also crystalline. These complexes are all stable in air. They are soluble in most common polar solvents and can be readily recrystallized from alcohols. Complexes of Fe(ll) and Fe(lll) with the Macrocyclic Ligand m-1, 7-CTH Several complexes of Fe(ll and III) have been prepared with the ligand meso- 5. 1 ,1 ,12, 1^ , 1^-hexamethyl- 1,!+, 8 , 11-tetraazacyclotetra- decane, meso-1,7-CTH (structure VIl). Three high spin six-coordinate pseudo-octahedral Fe(ll) complexes (cl ,Br , and I ) and three low spin six-coordinate pseudo-octahedral Fe(ll) complexes (CN ,NCS and the bisacetonitrile adduct) have been characterized by use of infrared spectroscopy, electronic spectroscopy, molar conductances, magnetic susceptibilities and Mossbauer spectroscopy. One Fe(lll) complex ([Fe(m-1,T-CTH)c1 2 ]C104) has been fully characterized using the same techniques. The synthesis, characterization, and some of the reactions of these complexes are discussed here. 72 Synthesis.- Because of the ease of the formation of oxides and hydrous oxides, a ll the Fe(ll) complexes must he made under a controlled oxygen-free atmosphere. Water, which does not seem to affect the synthesis of the Fe(ll) complexes of m-CRH as long as oxygen is not present (See page 51 ), presents a major problem in the synthesis of these compounds. The halide complexes, Fe(m-1,7“CTH)X 2 where X=C1 , Br , and I , are synthesized by mixing the ligand with hydrated iron salts in CH3CW solutions. The water of hydration present in the salts causes the in itial products to be contaminated with oxides and to appear as tan precipitates. These initial precipitates give fairly good analyses, but they contain variable amounts of ferro magnetic impurities. They often give absurdly high values for their magnetic moments, e. g ., the crude diiodide gave a /ieff of 9.6 B. M. When these compounds are redissolved in acetonitrile, a solid gelatinous residue remains on the filter. From the filtrate, pale green or blue powders are readily obtained which exhibit normal high spin magnetic moments. The synthesis of the thiocyanate complex and the solvated complex, Fe(m-1, 7-CTH)(c10 4 )2 *3CH3CN proceed sim ilarly except in these cases the complex is soluble enough to form pure complexes in the initial filtrate. The usual ferrous salts are so wet that most of the iron produces iron hydroxides when the basic amine ligands are added, thus, yields are very small. This difficulty can be greatly alleviated 73 if the surplus water is removed by use of 2 , 2 -dimethoxypropane or some other appropriate dehydrating agent. However, because all attempts at preparing these compounds in alcohols or other hydroxy- solvents have failed, care must be taken to lim it the amount of dehydrating agent used. Although only preliminary experiments have been conducted, it appears that a better method involves the use of anhydrous ferrous salts and the dry ligand. All of the pure high spin Fe(ll) halide complexes are almost colorless, showing only very pale green or blue colors. In contrast, the low spin Fe(ll) complexes (the cyanide, the thiocyanate, and the acetonitrile adduct) are deep colors. Dry solid samples of all the compounds are fairly stable in the air; however, after several weeks the halide complexes tend to turn tan, probably because of the reaction with the moisture in the air. In acetonitrile solution, all compounds undergo some type of reaction upon exposure to air as evidenced by rapid color changes. Preliminary observations indicate that the acetonitrile adduct in acetonitrile solution tends to undergo ligand oxidative dehydrogenation when exposed to the air. The positions and number of the C=N groups introduced in this manner have not been determined. The halide complexes, Fe(m-1 ,T _CTH)x2, seem to react with oxygen in a much different manner at least in acidic solution. Fe(lll) complexes are formed and may be isolated frcm acetonitrile solutions. The reasons for the distinct differences in reactions, i.e., formation of Fe(lll) compounds in the case of the high spin halides and ligand oxidative dehydrogenation for the low spin compounds, is not yet fully understood. Infrared Spectra. - A summary of the N-H stretching frequencies for the iron complexes of the ligand m-1, 7-CTH are shown in Table V. Since the infrared spectra of these complexes are essentially due to the macrocyclic ligand, and, in some cases, the anion or axial ligand, they are discussed collectively. The infrared spectrum of the pure ligand shows a characteristic peak in the N-H stretching region at 3279 cm 1 (See Figure 16). A second band appears at 32^7 cm 1 in the dihydrate, m-1,7-CTH* 2H 2 0. The ligand has no absorption in the region characteristic of imine bands ( 15OO-16OO cm *) indicating that it is free of starting material, but it does have many sharp bands in the so-called skeletal region between lOOO-llj-OO cm 1. This yields a pattern that is useful in detecting the ligand. The halide complexes all show sim ilar infrared spectra with the N-H band occurring at 3188 cm 1 (See Figure l 6 ); this represents a shift from the free ligand position 3279 cm 1. The skeletal regions of the spectra are all similar to those of the free ligand. The thiocyanate compound displays a markedly higher energy N-H stretch (3215 cm 1), but the skeletal region of its spectrum is 75 TABLE V Some Physical Properties of Fe(m-1, 7"CTH)2+s3+ Complexes Compound C olor £ jeff (B. M. ) ^nhos N -H strd Fe(m-1, 7-CTH)C12 p a le g re e n 5.58 22a 3188 Fe(m-1,7-CTH)Br2 p a le b lu e 5-47 28a 3188 Fe(m-1, 7-CTH)l2 p a le b lu e 5.37 20a 3188 Fe(m-1, 7-CTH)(CN)2 brown 1.1(450)® 40* 3188 Fe(m-1, 7-CTH)(NCS)2 g re e n 0. 4 3 (8 6 )® 35* 3215 Fe(m-1, 7-CTH) (C104 )2* 3 CH3 CN p u rp le 1. 3 ( 761 )® 316® 3247 (Fe(m-1, 7-CTH)c12 )BF4 y ello w 5-97 120b 3215 (F e(m -1 ,7-CTH)ci2)ci04 y ello w 5.95 I33b 3215 aMolar conductance of 10 3 molar nitromethane solution. bMolar conductance of 10 3 molar methanol solution. °Molar conductance of 10 3 molar acetonitrile solution. ^ U n its cm x. SMagnetic susceptibility at 25°C (units 10 6cgs units). .50 .60 .20 Absorbance.40 .70 iue 6 Ifae Seta f em1 -T)l (upper) 7-CTH)Cl2 Fe(m-1, of Spectra Infrared 16. Figure . 3.0 2.5 3000 . 5.0 4.0 aeegh (microns) Wavelength 2000 17CT (lower) TH -1,7-C m cm 6.0 1500 . 8.0 7.0 1400 .40 .20 .30 .10 .70 60 50 iue 6 (o'. F( 17CHC2 (upper) -1,7-CTH)Cl2 Fe(m (con't.) 16. Figure 0 .20 .30 .60 .50 Absorba nee .40- .70 .10 00 . 0 1300 - 8.0 00 t 890 t , 1000 10.0 12.0 -,7CH (lower) 7-CTH m-1, aeegh (microns) Wavelength 4 0 60 18.0 16.0 14.0 T I I I_L I I I I T » 0 c 500 cm 600 20.0 22.0 24.0 .70 .40 .50 .20 .60 30 78 similar to that of the free ligand (See Figure 17). The infrared spectra suggest that the thiocyanate is bonded to the iron(II) through its nitrogen atom. This is indicated by the positions of the C=N stretching frequencies, 20 66 cm 1 and 2096 cm 1, and the C-S stretching mode, 792 cm 1 ( 8 8 ). The infrared spectrum of the acetonitrile adduct (Figure 17) resembles that of the thiocyanate, having its N-H stretch at 32^7 cm 1. A weak, fairly broad acetonitrile peak is observed near 2020 cm 1 and a stronger, sharp one due to the uncoordinated CH3CN occurs at 2257 cm x. The characteristic perchlorate bands are also present at approximately 1 1 0 0 cm 1, and 625 cm 1. The skeletal region indicates that the ligand is present. The infrared spectrum of the dicyanide complex (Figure 18) is sim ilar to those of the halide derivatives with an N-H stretching mode a t 3188 cm 1 and a backbone region similar to that of the free ligand. The ChN stretching region shows three peaks, two of which are weak (2058 and 2008 cm 1) and one of which is strong ( 201+ij- cm 1 ). The N-H stretching band appears at 3 2 1 5 cm 1 in the infrared spectrum of (Fe(m-1,7-CTH)C1 2 )C104. The characteristic bands for ionic perchlorate also appear. None of these spectra have bands in the 0-H or C=N stretching regions. This indicates that no hydrous oxide impurities are present and that no sites of unsaturation exist in the ligands in these .60 0.0- .50 Absorba nee.20 40- .70 30 00 00 00 m 10 10 1200 1400 1500 1 Cm 2000 3000 4000 . 10 iue 7 Ifae Seta f em-,-T)NSz (upper) -1,7-CTH)(NCS)z Fe(m of Spectra Infrared 17. Figure . 304.0 3.0 2.5 - aeegh (microns) Wavelength em1 -T)C0) -H0 (lower) -3H20 7-CTH)(C104)2 Fe(m-1, 5.0 6.0 7.0 8.0 .50 .60 .20 40 0.0 70 30 .50 .60 .20 Absorbo nee . O.Gi .40- .70 30 00 iue 7 (o'. F(n1 -T)NS2 (upper) 7-CTH)(NCS)2 Fe(rn-1, (con't.) 17. Figure 1300 8.0 i i iii. O O 800 IOOO 10.0 12.0 em1 -T)C0) *H0 (lower) *3H20 7-CTH)(C104)2 Fe(m-1, aeegh (microns) Wavelength 14.0 6 018.0 16.0 I —— I L I I I—1—J I I 0 c 5C0 cm 600 20.0 22.0 24.0 .70 oo .50 .40 .60 .20 1.0 30 00 o .50 .60 .30- . Absorbance.20 0 0 iue 8 Ifae Setu o Fem-1,7-CTH)(CN)2 e(m F of Spectrum Infrared 18. Figure 0 4 00 00 00 m ' Cm 2000 3000 4000 .70 . 3.0 2.5 - - 4.0 aeegh (microns) Wavelength Wave 5.0 6.0 1500 40 1200 1400 7.0 8.0 .50 00 00 .10 .40 .30 70 60 20 0 .20 .60 Absorba nee .50 Figure 18. (con't.) Fe(m-1, 7-CTH)(CN)z Fe(m-1, (con't.) 18. Figure .70 30 . 0 I30C - IOOO 10.0 12.0 aeegh (microns) Wavelength 4 0 16.0 14.0 18.0 20.0 22.0 24.0 - .20 oo .50 .60 1.0 70 4040 30 83 compounds. However, the infrared spectrum of a product which is obtained from a solution of the acetonitrile complex after exposure to air gives a band at 16^2 cm 1 while the rest of the spectrum is identical to that of the starting material. This suggests that sites of unsaturation are produced by air oxidation of this complex. Sub sequent studies by others ( 9 5 ) have provided other evidence for this oxidative dehydrogenation. Characterization of Complexes. - The Fe(ll) complexes with the macrocyclic ligand m-1, 7-CTH may be divided into two categories on the basis of their magnetic moments. The three halo-complexes, F e(m -1 ,7 _CTH)x 2 where X=C1 , Br , and I , are six-coordinate pseudo- octahedral and high spin complexes while the thiocyanate, cyanide, and the acetonitrile adduct are six-coordinate, pseudo-octahedral, low spin complexes. The colors of the high spin complexes are very pale, almost colorless, blue or green, while those of the low spin complexes are very deep colors. The halo-complexes have low molar conductances showing that they approach nonelectrolytes in nitromethane (89) (See Table v). F o r this reason these complexes are judged to be six-coordinate in solvents of low polarity and are assumed to be six-coordinate in the solid state as well. The compounds are quite different from the Fe(ll) halide complexes of the macrocyclic ligands m-CKH (structure IX), 1 , T-CT(v) and 1 ,3 j 7 , lO-CT(xv). The latter all contain five-coordinate high spin iron(ll) in the solid state. (See page 107). .N N. N N XV. 1,3, 7, 10-CT The magnetic moments (5-58,5-^7, and 5-37 B. M. ) of the high spin complexes Fe(m-1,7 ~CTH)x2, where X=C1 , Br and I , respectively are also higher (on the order of 0. U B. M. higher) than those (between 5.0 and 5.1 B. M.) of the complexes of the other ligands (five-coordinate species). This may arise because of the larger orbital contribution from the more weakly split trip let ground level in the pseudo-octahedral complex as compared to the five-coordinate species (92). The visible spectra of the high spin compounds are shown in Table VI. They exhibit d-d bands at 12. OkK and a weak shoulder in the 15 to 17kK range. The band in the 12. OkK region is assigned to the transition to the 5Aig? and the weak shoulder is assigned to the TABLE VI Electronic Spectra of Fe(ll) and F e(lll) Complexes of m-1,7-CTH Charge Compound d-d Bands (kK) T ra n s fe r S o lv e n t (kK) Fe(m-1,T-CTH)Cl2a 12. 0 ( 6=8 ) 17. 2 sh 2 9 .4sh LfeOH Fe(m-1,7-CTH)Br2a 12.0(6= 4) 17. 5sh 2 9 .4sh MeOH Fe(m-1,7-CTH)I2a 11.9 (€=5) 16.4(6= 5) 29-4sh MeOH F e(m -1,7-CTH)(cn) * 17-3 2 3 .8 MeOH 1 2 .0 17.3 2 3 .2 2 6 .0 m u ll Fe(m-1,7-CTH)(NCS)2a 12. 2 ( 6=1 2 ) 17.7(€=77) 25.3(€=104) ch 3no 2 1 6 .8 2 3 .8 m ull Fe(m-1,7-CTH)(C104 )2* 3CHaCNa 12.0(e= 5) 17.7(6=90) 27- Osh 3 0 . 8 sh CH3CN 12.3 1 8 .0 2 6 . Osh 2 9 .4sh m u ll (Fe(m-1,7-CTH)Cl2 )bF 4 No peaks MeOH (Fe (m-1 ,7“CTH)C12 )C104 • No peaks MeOH & Solutions made in an inert atmosphere 86 transition to the 5Big level. (See Figure 19). This would give a tetragonal splitting for these complexes on the order of 3 to 5kK which is not at all unlikely (9^). In methanol solution a charge transfer hand is observed in the 3 ®kK range in a ll these compounds. The spectral region between 15 and 27kK is particularly troublesome since contact with traces of oxygen causes shoulders to appear in this region. The weak shoulder around 1J. OkK is intensified considerably when the solution is exposed to oxygen. However, this new band appears at a slightly different position, and thus, is not considered to be the same band that was assigned to 5Eg sAig, but is thought to be a new, more intense band which is a result of exposure to oxygen. The appearance of these shoulders when the halide complexes are exposed to air may be due to oxidative dehydrogenation which could produce sites of unsaturation in the ligand. The metal to ligand charge transfer of these new unsaturated macrocyclic complexes is quite distinctive and different sites of unsaturation have been shown to yield charge transfer bands ( 9 5 )* The low spin complexes are the thiocyanate, cyanide and the acetonitrile adduct. The complexes do not seem to be very sensitive to oxygen in the solid state. They are six-coordinate in nitromethane solution (Table v) and are also assumed to have this coordination number in the solid state. The acetonitrile adduct is a di-univalent iue 9 Eetoi Seta f ( ,7CHX X C7 Br", Cl7 = X 7-CTH)X2 1, - e(m F of Spectra Electronic 19. Figure LOG U) 3.5- 0 MO 0 70 1000 700 500 MOO 300 30 fVNJBR CKKD HfiVENlJHBERS WAVELENGTH CNIO CNIO WAVELENGTH 82 18 22 28 oie . _ Iodide Bromide. Chloride . A V 10 st 88 electrolyte in acetonitrile, as expected, while the thiocyanate and cyanide derivatives are essentially monelectrolytes. The residual paramagnetism of the cyanide and the acetonitrile adduct (expressed in Table V as magnetic moments) are somewhat higher than is normal for low spin Fe(ll). The thiocyanate derivative has a residual moment of 0.1+2 B. M. It is assumed that small amounts of ferromagnetic impurities are responsible for the high values of this quantity for the other two complexes. However, repeated attempts to purify these materials by recrystallization did not alter the magnetic susceptibility appreciably. The residual magnetic moments of many low spin Fe(ll) complexes have been reported to be as high as 1.0 B. M. (96). This is usually attributable to either a high spin or a ferro magnetic impurity. In a few cases this signals the more interesting possibility of spin state equilibrium (53j62). The high values here could also be due to an Fe-O-Fe oxobridged species (97,99) j however, this is probably not the case here for three reasons: the infrared spectra show no Fe-O-Fe stretching band around 83 O cm x, th e v i s i b l e spectra are that of normal low spin Fe(ll) complexes, and the Mossbauer spectra are also normal for low spin Fe(ll) compounds. The visible spectra of low spin pseudo-octahedral F e(ll) complexes are expected to be analogous to those of low spin C o(lll) complexes. The spectra consist of two or three d-d bands with ex tinction coefficients as large as 100, sometimes accompanied by a 89 spin forbidden band at low energies ( 9 0 ,1 0 0 ) . The singlet singlet transitions are (in order of increasing energy) are xAig -> xTig, 1Aig - 1T2g, and xAig xT 2g,' and the spin forbidden transitions are 1Aig - 3Tig and 1Aig - 3T2g- The third singlet singlet and the second singlet triplet transition are not normally seen because they are obscured by other bands. Three d-d bands are observed in these complexes (See Figure 2 0 ); however, the first band of each complex is very weak and broad. Using the equations derived by Wentworth and Piper (see below) (9l)s if the weak band at low energy is considered to be the spin forbidden band 1Aig -» 3Tig and the other two bands are assigned to the xAig -* 1Tig and xAig -> 1T2gj> the spectral parameters for the thiocyanate complex are Dq=l995 cm x, B=^75 cm x, and C=2250 cm x; those of the bisacetonitrile adduct are Dq=2060 cm x, B=500 cm x, and C=2600 cm x; and those of the cyanide are Dq=2000 cm x, B-14-06 cm 1 and 0=26^0 cm x. The equations used by Wentworth and Piper are the following: E( xT2 ) -E ( xA i) =10Dq + 16B-C E(XTi)-E(xA i)=lODq-C E(3T2 ) -E ( xA i) =lODq + 8 B-3C E(3T i)-E(xAi)=l0Dq-3C gur 0 El toni cr f ,7CH)NCS)2 C )(N 1, - 7-CTH m ( e F of m ectru p S ic n ctro le E 20. re u ig F L0GU?) 0.5- 2 2.5- . 0 228 32 - AEUBR CKKO WAVENUMBERS 0 50 0 1000 700 500 400 AEEGH CNMI] WAVELENGTH 24 20 16 12 90 91 The parameters B and C are the electron-electron repulsion parameters and are always smaller in compounds than in the free ion state where B=1058 cm 1 and C=3901 cm 1 (lOl). The values calculated for the spectral parameters of the acetonitrile and the thiocyanate complex are in the expected range (l 0 2 ); however, there is some question as to why the Dq value for the cyanide complex is not higher because cyanide is a much stronger ligand than either acetonitrile or thiocyanate (l 0 2 ). Proceeding on this expectation, the cyanide spectrum might be inter preted in terms of a tetragonal model (9l)- However, the assumption that the 1Tig level is split into the 1A2g and the 1Eg levels leads to too large a separation between these states so that this model has been discarded. Also the tetragonal model produces a Dq value for the ligand of only about 1700 cm x. Exposure of acidified acetonitrile solutions of Fe(m-1, 7~CTH)x2, where X=C1 , Br , and I , to air leads to the formation of the corres ponding F e(lll) complexes [Fe(m-1, 7“CTH)X2]Y where X=C1 , Br , and I and Y=C10 4 and BF 4 . The dichloro-complex has been isolated and purified, but the dibromo- and diiodo-complexes were not obtained as pure solids. These complexes show no C=N stretches in the infrared spectra and have the characteristic N-H stretching modes of the m-1,7“CTH ligand. The fingerprint region is somewhat obscured by the a n io n s (BF 4 and C10 4 ) but several of the sharp bands are s till . - visible indicating that the ligand is still unaltered. 92 [Fe(m-1, 7-CTH)c1£ ]C104 is a yellow crystalline salt, whose magnetic moment has a value typical of high spin iron(lll) (5.97 B. M. ). This salt is a uni-univalent electrolyte in methanol indicating that both halides are bound even in solvents of moderate polarity. The visible part of the electronic spectrum shows no bands in the range from 8,000 to 30,000 cm 1. This is not unexpected as there are no spin allowed d-d transitions in high spin Fe(lll) spectra and the charge transfer bands, which fall above 3 0 ,0 0 0 cm 1, tend to obscure the weak spin forbidden transitions. Complexes of Fe(ll) with the Macrocyclic Ligand 1.7-CT The macrocyclic ligand 5, 5, 7 , 12,12, 1^-hexam ethyl-l, it-, 8 ,1 1 - tetraazacyclotetradeca-l(lit-), 7~diene(l, 7~CT, structure v) is synthesized in a manner similar to that developed by Curtis and Hay ( 8 3 ) in which ethylenediamine dihydroiodide and etheylenediamine are mixed with acetone producing 1,7~CT*2HI. To this macrocyclic ligand is added ferrous acetate in acetonitrile under an inert atmosphere. Ferrous acetate is used because it is readily available in a pure form and the acetate anion is basic enough to deprotonate the acid form of the ligand. However, since anhydrous ferrous acetate is fairly insoluble in acetonitrile, the reaction mixture is a slurry and several hours are required for the reaction to take place. The final product is the complex [Fe(l,7-CT)(cH 3CN)2 ]I 2 from which the acetonitrile can be re moved by warming to 60°C in vacuo forming the complex Fe(l,7_CT)l2. 93 Another complex is prepared by mixing ferrous thiocyanate hexahydrate with the ligand 1, 7-CT* 2HC10 4 which was prepared in a manner sim ilar to the 1, 7-CT*2HI. From this reaction mixture a deep maroon compound is obtained which analyzes for Fe(l,7~CT)Fe(NCS) 4 *2CH^CH This complex is extremely unstable when exposed to the air turning black within seconds. Several of the physical properties of the iodide complexes are shown in Table VII. TABLE V II Properties of Fe(ll) Complexes of the Macrocyclic Ligand 1,7~CT j'T-.----r7=r-j--a.-sjj.—* 1 ■■ cur v.r = ...... - Compound Electronic Spectrum(kK) ;ie f f Anho (B.M.) [Fe(l,7-CT)(cH 3 CN)2 ]I 2 0. 62 2^ all. 9(6=5) 19.6(€=8l)a29. ^(€= 5 2 0 0 ) (Fe(l,7-CT)l2) 5-25 90b 12.2(€“13)° a10 3M acetonitrile solution b10 3M nitromethane solution Cethanol solution The infrared spectra (Figure 21) of the two compounds are comparable to the spectra of similar nickel complexes (l 6 ). The C=N stretches of these complexes occur at l 6 k l cm 1 which is identical to those .50 .60 0 3 . Absorbance.20 .40 .70 Figure 21. Infrared Spectrum of Fe(l, 7-CT)(CH3CN)2I2 Fe(l, of Spectrum Infrared 21. Figure . 304.0 3.0 2.5 3000 aeegh (microns) Wavelength Wave 2000 5.0 cm 6.0 501400 1500 7.0 8.0 -.60 -.50 -.70 oo .20 .40 .30 .10 Figure Figure .60 .50/ Absorba.30 nee .20 .7 Or 00 40 1300 21. (con't.) (con't.) 21. 10.0 1000 Fe(l, Fe(l, 12.0 522 7-CT)(CH3CN)2I2 aeegh (microns) Wavelength 14.0 16.0 —I 1 I I— — 0 c 500 cm 600 18.0 ------1 _____ I ____ 20.0 I I 22.0 24.0 oo .70 .60 .50 0 4 .20 30 96 reported for the nickel complexes of 1,7~CT (l6). The regions between 1000 cm 1 and 1500 cm 1 are identical to the spectra of the two compounds, and there is a weak peak at 2260 cm 3 in the spectrum of the acetonitrile complex. This arises from the C=N stretching mode of the acetonitrile. The N-H stretches appear at quite different positions. In the acetonitrile complex the N-H stretch occurs at 3135 cm 1 while in the diiodo-complex the peak in this region appears at 3^85 cm 1* From the magnetic moments given in Table VII, the maroon acetonitrile complex is a low spin Fe(ll) complex while the diiodo- complex is a high spin Fe(ll) complex. This indicates that the acetonitrile acts as an axial ligand in the former case. From the conductance data in Table VII, it is observed that the acetonitrile adduct is a di-univalent electrolyte in acetonitrile while the diiodide is a uni-univalent electrolyte in nitromethane. Thus, the acetonitrile complex is assumed to be six-coordinate while the diiodide is five-coordinate in solution. The visible spectra of these two complexes shows that the diiodo-complex is a high spin complex with a d-d band at 12. 2kK. The lower energy band which appears in the spectra of five-coordinate complexes (see page 1 0 7 ) is not observed here and this region in its solid state spectrum is very unclear because of overtones of infrared b a n d s. 97 The acetonitrile adduct shows a low spin Fe(ll) spectrum with two d-d hands, one at 11. 9kK and the other at 19. 6 kK (See F ig u re 2 2 ). These hands are assumed to he due to the xAig - XTig transitions (See page 89 ). The second hand is higher in energy in this complex than in the sim ilar complex Fe(m-1,7“CTH)(C104)2. 3CH3CN which indicates that the macrocyclic ligand 1,7~CT is stronger than m-1, 7-CTH. Using the equations of Wentworth and Piper (9l), the Dq of the 1, 7-CT compiex is 2350 cm 1 compared to 2060 cm 1 for that of the similar m-1,7_CTH complex. The second singlet singlet transition (xAig -» ^ag) is prohahly not observed because of the intense charge transfer hand at 29- ^kK. These complexes tend to oxidatively dehydrogenate upon exposure of their solutions to air. In this way they are similar to the Fe(ll) complexes of the m-1,7-CTH ligand. Bands begin to grow in the region betw een 15 and 3 0 kK when they are exposed to air indicating that double bonds are being formed. The synthesis of several other Fe(ll) compounds of 1,7-CT was attempted with little success. Compounds of the dibromide and dichlo ride were obtained, but their analyses were never acceptable. These complexes were light brown amorphous powders with infrared spectra showing the presence of the ligand. The complex Fe(l,7-CT)Fe(NCS) 4 *2CH3CN is an interesting complex. Its Mossbauer spectrum shows the presence of the Fe(NCS)42“ anion LOG («) F igure 22. E lectronic Spectrum of Fe(l, 7-CT)(CH3CN)2I2 Fe(l, of Spectrum lectronic E 22. igure F 0.5- 2 2.5- . 0 0 40 0 70 1000 700 500 400 300 - 026 30 REUBR CKK3 WRVENUMBERS REEGH CNM3 WRVELENGTH 22 1418 99 (See page 115 )• It is interesting to note that upon exposure to air the complex turns from maroon to black; however, this change in color is not accompanied by a change in the Mossbauer spectrum. Consequently, the reaction that takes place is thought to be a surface reaction. Mossbauer Spectra of the Complexes of Fe(ll) and Fe(lll) with Macrocyclic Ligands r—■ « • • In the last few years the Fe Mossbauer effect has become a useful tool for the study of iron compounds. It has been used very successfully in the determination of spin state and structure and also in the correlation of electronic properties of iron compounds. The spectra of many iron complexes that had been previously prepared have been measured in the last seven or eight years in order to try to under stand how the Mossbauer effect can be applied to chemistry. The chemistry of iron has been advanced by the following applications of the effect: compounds of tetrahedral coordination ( 103, 101+) and o f five-coordination ( 1+2, 58 , 6 0 ) have been identified; distinctions have been made between cis and trans isomers ( 1 0 5 ); compounds of intermediate s p in s (l+ 3 > 53s 1 0 7) have been recognized; spin-state equilibria ( 6 3 - 6 5 , 108 -in ) have been elucidated; and the effects of both cations ( 1+8 ) and anions (57*112) on the Mossbauer parameters have been elaborated. The Mossbauer spectra of many biological macrocyclic iron complexes have been utilized to help understand the role of the iron in these necessary compounds ( 3 1 ,113* The Mossbauer spectra of 100 one synthetic tetradentate macrocyclic iron compound, the prophyrin- like phthalocyanine iron has been discussed in literature ( 1+3 , 107, lip) as have the substituted porphyrins (31,57)* The Mossbauer spectra of a few new iron macrocyclic complexes are discussed here. The spectra of many different compounds have been measured and such additional physical data as their electronic spectra and magnetic data are avail able to help relate the several structural characteristics of the complexes (oxidation state, spin state, and coordination number). The literature clearly shows that the Mossbauer spectra of many different compounds can be related ( 79 ~8 l ) . The Mossbauer spectra of iron complexes of several tetradentate macrocyclic ligands have been measured and are discussed here (See Appendix). The parameters from the Mossbauer spectra (6 and AEq) y ie ld much information relevant to coordination number, comparative strengths of the macrocyclic ligands, ground states of the iron, oxidation states of the iron, conformation of the ligands, and comparative strengths of the axial ligands. Results of the Mossbauer Spectra. - A typical two line spectrum is shown in Figure 8 . A Mossbauer spectrum is centroid. The width at half height of a given peak was normally between 0. 3 - 0. 1+ mm/sec with several being less than 0 . 3 mm/sec. These give values that are accurate to within + 0.01 mm/sec. A few poor spectra had line widths o f O.5 mm/sec or greater. These are designated in the tables. Tie $ 101 absorbance of the 57Fe nuclei ranged from 0. 5^ to 7. 5$ with most falling in the range from 2 to k<{0. The $ absorbance is related to the number of 57Fe nuclei in the sample. Several other nuclei w ill absorb 14. U Kev gamma ra y s , and th u s low er t h i s v a lu e ; e . g . , I w i l l ab sorb a large number of gamma rays, and thus the $ absorbance of complexes with I in them w ill be low. A compilation of the Mossbauer spectra appears in the appendix. Tables VIII, IX and X show the Mossbauer parameters for the Fe(ll) complexes. The data are grouped according to the spin state and coordination number of the complexes. Tables XI and XII show the parameters for the high spin and low spin Fe(lll) complexes reported here. Table XIII shows two complexes that are different; the first, by virtue of its having an Fe-O-Fe linkage; and the second containing a tetrahedral iron anion. From the entries in Tables VIII, IX and X, it is seen that high spin and low spin Fe(ll) complexes are quite easily distinguished from each other on the basis of the values of their isomer shifts ( 5 ). The high spin complexes have a 6 value of greater than 1. Imm/sec with respect to nitroprusside while the low spin complexes range between 0. 5O-O. 8 Qmm/sec (ll 6 ). The Fe(lll) complexes are not so easily distinguished by their isomer shifts. However, the high spin Fe(lll) compounds, because of the symmetry of their d-electrons, normally have quadrupole splitting 102 TABLE V III Mossbauer Spectra of Five-Coordinate High Spin Fe(ll) Complexes a / Compound ue f f B.M. 6 mm/sec AFnmm/ sec Fe(m-CRH)C12 5*21 1 .1 1 3.72. 1 .23b 3 . 08° 9 .6 Fe(m-CRH)Br2 5-14 1 .1 2 3 . 81+ l . 6 Fe(m-CRH)l2 5-05 1. 08 3 . 8 ^ 1 .9 [Fe(l,T “CT)Cl]C104 5-11 l . ! 5 b 3 - 7 \ 2. 6, 1 .2 9 3*98 3*o T F e(l, 7-CT)Br]C104 5-1^ 1. 1^ 3 .8 1 3 .7 [Fe(l,7-CT)l]C10 4 5 .1 6 1 .1 2 3 . 3^ 5 -7 5 .0 0 3 . 6 a 2 . 8 . [ F e ( l , 3 , 7j10-CT)Cl]C104 L 15b 1 .2 9 3 .8 1 6 .? [Fe(1,3,7,10-CT)Br]C104 5 .1 k 1. lb 3-78 1 .8 [Fe (1,3, 7 ,10-CT) I] C104 5 .1 0 1 .1 0 3-79 2 .3 [F e(l,3, 7,10-CT)ci]BPh 4 1 .1 1 3 .6 2 1 .2 €1 A ll 6 values are with respect to sodium nitroprusside. ^Spectra taken at -19^°C. 1 0 3 TABLE IX Mbssbauer Spectra of Six-Coordinate High Spin Fe(ll) Complexes Compound j^eff B.M. 6amm/sec AE mm/sec <$, e f f e c t 4 Fe(m-CKH)(0Ac)2 5-32 1. 36 2 . 27 2 .2 [Fe(m-CRH)(OAc)]PF6 5.H 1.19 2 . lt-0 3-3 Fe (m-CRH) (N3 ) 2 5 .5 1 1. 2k 1 .6 8 6 . l Fe(m -1, 7-CTH)ci2 5 .5 8 1 .1 7 2 .1 1 k A F e (m -l3 7-CTH)Br2 5-V7 1. l 6b 1. ^ o .k Qt A ll 6 values are with respect to sodium nitroprusside. bBroad Spectrum. 10U TABLE X Mossbauer Spectra of Low Spin Fe(ll) Complexes u e f f B.M. 6amm/sec A£Qmm/sec ^ effect Fe(m-CRH)(WCS)2 o. 80 o. 71 0.67 Fe(m-CRH)(C104 )2 o. 96 0.62 o. 68 3-1 Fe(m-1,7-CTH)(NCS)2 o. 1+3 0. 81 0.29 5-^ Fe (m-1,7-CTH)(cn )2 1.12 O.65 1.28 4.4 Fe(m-1, 7-CTH)(CH3CN)3 (C104 )2 1. 33 O.76 0.55 1.7 ■u Fe (m-1, 7~CTH) ( im id azo le ) 2 (C104 )2 0. 93 0.54 1.54 1. 0 F e ( 1 ,7~CT) (NCS)2 diam ag 0.70 0.77 0.5 Fe(l,7-CT)(CN)BPh 4 diam ag 0.51 0.57 1. 05 (Fe (1, 7-CT) (CH3CN)2 ) (C104 )2 diam ag 0.67 1.07 1.9 ( F e ( l, 7-CTi-2)(CH3CN)2 (ci04 )2C diam ag 0.66 I.03 2.6 (Fe(l,7~CT)(imidazole) 2 )(BPh 4 )2 diam ag 0.66 1.07 1.28 Fe(1, 7-CT)(CH3 CN)2 I 2 0.62 .67 1.06 Fe(1,5, 7,10-CT)(CH3CN)2 (C104 )2 diam ag 0.60 1.36 2.5 F e ( l, 5, 7 , 10-CT)(CH3CN)2 (BPh 4 )2 diam ag 0.60 1. 50 2.2 F e (1 ,4 , 8 , 10-CT) (CH3 CW)2 (BPh 4 )2 diam ag 0.56 1.86 0.9 Fe (1 ,3 ,7 ,10-CT) (NCS )2 0.60 1.21 7.8 Fe(l, 3, 7 ,10-CT) (im idazole) 2 (C104 )2^ 0.55 - 1.64 1.7 Fe (1,3, 7 ,10-CT) (imidazole )2 (BPlxj. )2 0.60 1.38 1.8 & All 5 values are with respect to sodium nitroprusside. ^Analytical data suspect. Thought to be a different isomer. 105 TABLE XI Mossbauer Spectra of High Spin F e(lll) Complexes Compound y e f f B.M. 5amm/sec AE I^mm/ sec $ e f f e c t [Fe(m-CRH)C12]BF4 5.86 0. 59 0.66 3 .3 [Fe(m-CRH)Br2"]BF4 5. 75 0 .6 l 0 . 51+ 1-3 [Fe(m-1,T-CTH)C12 ]BF4 5- 98 0. 53 1.^5 1-3 [Fe(m-1, T~CTH)Br2 ]BF4 0 .7 ^ 0 1 .1 TABLE X II Mossbauer Spectra of Low Spin F e(lll) Complexes Compound ^ e f f B.M. 6amm/sec AE mm/sec $ effect [Fe(l,7-CT)C1 2 ]C104 2 .3 0 0. 1+8 2 .6 6 k .6 Fe(1, 7-CT) (NCS)2 JflPh* 2 . 2 0. 82 2 .1 6 0 .9 Fe(l, 7-CT)(cH3CN)2 )(c 1 0 4 )2, 2 .1 9 0 .8 3 1 . k6 1 .5 TABLE X III Mossbauer Spectra of Miscellaneous Complexes & / Compound jLieff B. M. 5 mm/ sec AE-rnm/ sec [Fe(1,7-CT)(NCS)12 0(NCS)2 2 .3 0 . 52b 0. 50b 3-T Fe(1 ,7-CT)Fe(SCN)4- 2CH3CN 0 .6 6 l . 00 2 .3 1 .1 3 1.29 aA ll 6 values are with respect to sodium nitroprusside. I'Broad Spectrum 106 values (AEq) of less than 1. 5mm/sec, while thcpse of low spin F e(lll), Which has a nonspherical d-electron distribution, normally have AEn values greater than 1. 5mm/sec (ll 6 ). Eecause of the great difference in magnetic moment between high spin and low spin F e(lll), no complica tions arise here. However, it is difficult to distinguish between low spin compounds of Fe(ll) and Fe(lll). The 5 values for these compounds have the same ranges, but the AEq for low spin Fe(ll) is normally be low 1. Qtnm/sec because of its spherical d-electron symmetry, while, as already mentioned, AEq values of low spin Fe(lll) are greater than 1. 5™i/sec (ll 6 ). Consequently, unless the ligand field produces a large amount of dissymmetry, these two can be distinguished fairly readily on the basis of their Mossbauer parameters. Several Fe(ll) and Fe(lll) complexes with several different macrocyclic ligands are discussed here. The ligands that have been used are the following: meso- 2 . 1 2 -d im e th y l- 5 . 7 , 1 1 , 1 7~tetraazabicyclo- [11. 3 . l]heptadeca-l(lT), 15>15-triene(m-CRH, structure IX); meso-5. 7.7.- 12,1^, lH-hexamethyl-1, It-, 8 , 11-tetraazacyclo-tetradecane (m-1, 7-CTH, structure VII); 5,7, 7,12, l^, lij--hexamethyl-l, k} 8 ,11-tetraazacyclotetra- deca-1,7~diene(l, 7-CT, structure v); 5, 7 ,7 ,1 2 ,1^, l^-hexamethyl-l,^, 8, - 11-tetraazacyclotetradeca-1,3? 7,10-tetraene(1, 3 ,7*10-CT, structure XV); 107 5, 5, 7, 12, 1^ , 1^-hexamethyl-l,^, 8 , 1 1 -tetraazacyclotetradeca-l, l<-, 8 , 10- te tr a e n e (1, *4-, 8 , 10-CT, structure XVI); and tetrabenzo(b,f, j,n)- (l,5,9,13)tetraazacyclohexadecine(TAAB, structure III). These ligands form both high spin and low spin Fe(ll) and Fe(lll) complexes in a variety of coordination numbers. XVI. 1,4,8, 10-CT Table IX summarizes the rather interesting parameters for the most unusual group of materials studied here. The quadrupole splitting values (aEq) for these structures are among the highest values ever recorded for iron compounds. The AEq values are in the range from 3*3 to 3.9mm/sec, with all but two being larger than 3- 6mm/sec. The corresponding isomer shift values are in the range normally found for high spin Fe(ll) compounds (ca. 1.1-1. 3mm/sec). Normal six-coordinate high spin Fe(ll) compounds have quadrupole splittings in the range 108 2. 0~3. Omm/sec (ll6), while those of so-called ionic salts are normally between 2.5 and 3*5mm/sec (llT). There have been very few compounds reported with AEa v a lu e s *=4 greater than 3 . 5mm/sec. Values of this magnitude have been reported for high spin Fe(ll) pyrazolylborate (ll 8 ) and for a five-coordinate, low spin Fe(lll) phthalocyanine complex ( 1 1 5 ) which has a geometric structure similar to those of the complexes reported here. There are two factors (ll9) that contribute to the quadrupole splitting. The first is the dissymmetry of the electron density around the iron nucleus. The more nonspherical the electron density, the larger is the electric field gradient (efg), and this separates the nuclear energy levels giving rise to a quadrupole splitting in the Mossbauer spectrum. The other factor, which usually has a smaller e f f e c t on A E q , is due to the dissymmetry of the ligand field. The more symmetric the ligand field, the smaller is the electric field gradient. Consequently, the nuclear splitting is smaller and AE q is smaller. Chemical evidence supports a five-coordinate structure as being the reason for this large quadrupole splitting in the complexes [Fe(l, 7-CT)x]C10 4 and [Fe(l,3, 7,10-CT)X]C10 4 where X=Cl‘,Br" and I". This evidence involves their stoichiometries, magnetic moments, molar conductivities, infrared spectra, and d-d electronic spectra. The infrared spectra indicate the absence of water or other hydroxylic species and show only free perchlorate ions. Conductivity data 109 indicates uni-univalent electrolytes in nitromethane (89). (Table XIV). The best evidence is the d-d electronic spectra which show that the strong tetragonal distortion (c4v symmetry) present in these five- coordinate complexes produces a very large splitting of the 5E spectro scopic term state (derived from octahedral parentage) into 5B i and 5Ai states (corresponding to the d ^ .^ and d^ orbitals of the one- electron orbital description). This is consistent with the behavior of other metal ions (l20). Careful spectroscopic measurements on these complexes (solid state and nitromethane solution; 3 . 0 to 25kK region) reveal two broad weak d-d transitions near 5»0 and 12. 5kK f o r each compound (Table Xiv). The higher energy absorption occurs at the same energy for each complex and is assigned to the transition to th e B i(d x 2 _y2 ) state. The position of the lower energy absorption (transition to Ai(dz2)) is dependent upon the field strength of the axial ligand. For the complexes Fe(m-CRH)x 2 where X=C1 ,Br , and I much of this chemical evidence is not possible for assigning a five-coordinate structure. Much effort was put into attempting to observe the lower d-d absorption in the electronic spectra (See page 66 ), however, it was never observed because the complexes are not soluble enough in nitromethane. However, since these complexes have the same Mossbauer parameters as the other complexes in Table VIII, they are assumed to be structurally similar. 110 TABLE XIV Physical Properties of Some New Five-Coordinate Fe(ll) Complexes Compound /\amhos u e ff d-d Electronic Spectra (B.M. ) (Fe ( l , 7 -C T )d )C 1 0 4 95 5-05 I*. 7(7); 12.5(5) (Fe(l, 7-CT)Br)d04 106 5 . U 5(~5); 12.2(5) (Fe(l,7-CT)I)C104 93 5.15 5 ( J 4); 1 2 .2 ( 14-) [Fe(1,3, 7 ,10-CT)Cl]C104 97 5 .0 0 *4. 6 5 (3 ); 11. 6 ( 3 ) [Fe(1,3,7,10-CT)Br]C104 99 5 .1 ^ *4. 7 6 (3 ); 1 1 .7 (3 ) [Fe(l,3,7,10-CT)l]C104 102 5 .1 0 5 ; 11. 7 8. Molar conductance obtained in purified nitromethane under N 2 atm os phere; concentrations were in the range 1. 5x l 0~3 to 0. 97xlO“3M. Obtained in nitromethane under N 2 atmosphere, absorption maxima in kK, molar extinction coefficients are given in parentheses. The one electron d-orbital diagram for octahedral and square pyramidal structures (c 4 v ) are shown for the d 6 case in Figure 20. It can be seen that there is an odd electron which causes the iarge electric field gradient normally seen in high spin Fe(ll) complexes. Since the complexes in Table VIII are the first high spin five- coordinate and tetragonal pyramidal complexes of Fe(ll) reported, their Mossbauer parameters are without precedent; however, because of the electron distribution shown in Figure 20, a large electric field I l l gradient is to be expected. There are several, other high spin Fe(ll) complexes that are five-coordinate; [e.g., Fe(terpy )x2 w ith AE q v a lu e s o f ~ 3. Omm/sec (U2 )]; however, they are trigonal bipyramidal (4l,M<-). There are several other five-coordinate Fe(ll) square pyramidal complexes but they are of the so-called intermediate spin' (s=l), e.g., the phthalocyanines (U 6 , 1+7 , 12 l), N,NT salicylaldemines (M+), and dithiocarbamates (55?56), and others (1+6,1+7) with AE q values between 2 and Jmm/sec. Previously reported five-coordinate complexes of Fe(ll) have generally involved soft donors (P,As,S and imines of porphyrin and phthalocyanines), and this is assumed to be the reason the tetra gonal pyramidal complexes have all involved triplet (s=l) rather than quintet (s=2) ground states. Both the strong ligand fields and the large nephelauxetic effects of these donors favor spin pairing. The compounds reported here may have weaker in-plane field strengths than those studied earlier, and they certainly have smaller nephelauxetic e f f e c t s . Since the two effects causing the electric field gradient are additive (ll 9 )? the signs of their contribution to the net electric field gradient are important. These signs have not been determined in the present studies so that more than one union is possible between them. The most probable reason for the large AE q is that it is the sum of a large AEq due to the *'odd'’ electron and a sizeable but 112 1 dx * - y 2 1 1 dx 2 - y 2 » ^ z2 dz 2 dxy 1 1__ d X y , dx z , d y Z 11 1 uxdx z ’> uyzd> O h -'4 V Figure 20. d6 Ion in Octahedral and Square Pyramidal Fields. smaller contribution from the ligand field dissymmetry the two effects being of the same sign. If such a relationship holds for both five- coordinate (c 4v) and six-coordinate pseudo-octahedral cases, the dissymmetry due to the ligand field is much larger in the C4v case. This is true because in the C4v case, one side of the molecule contains no ligand while the other extremity of the same axis does contain a weak ligand. This produces a large electric field gradient. The asymmetry due to the d-electrons should be equal in the two structures. From the equation relating En to orbital parameters, '=1 113 both structures give the same value for Eq (l22). Table XIV gives the values of q and 77 for the different 3d electrons ( 1 2 3 ). Eq = + e2 Ob (1 + - ^ 3 where: e = charge on electron Q, = nuclear quadrupole moment q = electric field gradient 71 = asymmetry parameter TABLE XV Values of Field Gradient q and Asymmetry Parameter 77 f o r 3 d Electrons <1 77 dx^-y2 < r 3 > 0 k dz2 < r 3 > 0 ‘ T dxy < r"3 > 0 ♦* 2 dxz < r 3 > " T + 3 2 dyz < r 3 > - 3 " 7 From calculations of the E^, it can be seen that both T2g ground state for a six-coordinate structure and the E ground state of the five-coordinate structure yield the same value for E_. Consequently, Ill* the difference in AE^ between the five-coordinate and the six- coordinate structure must be due to the dissymmetry of the ligand f i e l d . The low spin Fe(ll) complexes yield several interesting com parisons. Since low spin Fe(ll) compounds have spherical d-electron symmetry, electronic assymmetry does not contribute to AEq, and a close look at the ligands, and how they affect the iron nuclei is in order. Since low spin Fe(ll) compounds are formed by three related macrocyclic ligands with varying degrees of unsaturation in their structures, useful comparisons can be made. To discuss Figures 2 k ,25, and 26 a new set of parameters w ill be defined. It is convenient to assume that the four nitrogen atoms of the macrocyclic ligand are in a plane, and that the four N-Fe bonds of the macrocycle are equivalent. The related bond moments are called xy. The axial bond moments w ill be called z. The r a t i o ~ w ill be used. In the low spin Fe(ll) complexes the nearer the ratio, — , is to one, the smaller w ill be the AEq xy 'i value, while the farther this ratio is from one (in either direction), the larger the AEQ value w ill be. Presumably z and xy w ill be very sim ilar when the ligands on the z-axis have bonding characteristics (to the metal ion) very similar to those of the macrocycle. F ig u re 2k shows the variation of isomer shift with degree of macrocycle unsaturation for two series of compounds, each series having Fe(CTH)(NCS) .8 0 Fc(l, 7-CT)(NCS) .70 .75 Fe(CTH)(CH3CN)2(C104) .60 % Pe(l, 3, 7, 10-CT)(CH3CN)2(C104) Fe(l, 3,7, 10-CT)(NCS)2 Fe(l, 7-CT)(CH3CNUC104) .55 0 2 4 Double Bonds Figure. 24. The Variation of Isomer Shift with the Degree of Unsaturation of the Macrocyclic Ligands. a constant axial ligand (z). As the degree of unsaturation increases, 6 decreases. This is explainable if one assumes that iron d-orbital reverse donor pi-bonding increases with the degree of unsaturation of the ligand. Since the d-orbitals normally shield s-electron density from the nucleus, the greater their participation in pi-bonding, the greater is the s-density at the Fe nucleus. This in turn decreases the isomer shift value. F ig u re 2k also points out that the 5 value for 1,3, 7,10-CT complexes is nearly independent of the axial ligand. Table 10 shows that all of the 6 values for the 1,3,7,10-CT complexes are 0. 60mm/sec. This invariance must imply that the s-electron density around the Fe nucleus is unchanged by changing the axial ligand in these cases. This in turn may imply only weak or moderate axial ligand—iron atom bonding. Table 10 also shows that the Fe(ll) complex of the ligand 1,1+,10-CT has a low er 6 (O. 56mm/sec) than the complexes of 1,3,7,10-CT. This suggests that the a-diimine group is a stronger pi-bonder than the 0-d iim in e group. Figure 25 shows the variation of AEq with the number of double bonds present in the macrocyclic ligand. As the unsaturation of the ligand increases (holding the axial ligand constant), AE q in c re a s e s . This trend accords with the conclusion about the strengths of bonding of axial ligands implied above (MAC=tetraimine > diimine > tetramine). 4 Fe(l, 3, 7 ,10-CT)(NCS)2 5 Fe( 1, 3, 7, 10-CT)(CH3CN)2(C104), a o & 2 Fe(CT)(CH3CN)2(C104), w o 13 &. cn Fe(CTH)(CH3CN)2(C104) 0 .4 8 1 . 2 1.6 AE q F ig u r e 25. The Variation of AEq with the Degree of Unsaturation of the Macrocyclic Ligands. 118 The increase in AEn values is not necessarily to he expected because the signs of the electric field gradients are not known, and thus only the magnitudes (absolute values) of the efg are plotted. The thio- cyanates almost form a straight line so that their efgs may all have the same sign (i.e., they do not go through zero). In any case the low AEa of the Ee(m-1, 7-CTH)(NCS )2 complex may be attributed to the fact that xy and z are nearly equal. (This AEq is so small that the peaks of the doublet overlap and Moon's correction has to be applied ( l 2U). ) Another interesting point is made in Table XVI. Here the macro- cyclic ligand is held constant (m-1, 7~CTH, xy-constant) while the axial ligand (z) is varied. The isomer shift values are ordered imid < CN « CH3CN < NCS . One can correlate the 6 values with the relative effective strengths of the axial ligands. This is considering a ratio of pi-bonding to sigma-bonding effects. Pi-bonding w ill tend to lower the isomer shift while sigma-bonding tends to raise the 5 value. Con sequently, this isomer shift data gives a relative value of pi-bonding versus sigma-bonding. This order of large pi-bonding to sigma-bonding ratio for the Fe(m-1,7-CTH)2+ species is: imidazole > CN » CH3CN > NCS . A similar ordering of this type is discussed by Konig with Fe(ll) bisphenanthroline complexes (53)* 1 1 9 TABLE XVI Variation of the Axial Ligand in Some Low S p in F e ( l l ) Complexes a / Compound 6 mm/sec ASAmm/sec Fe(m-1, T'CTH)(NCS)2 0 .8 1 0.29 Fe(m-1, 7-CTH)(cH3 CN)2 (c104 )2 0 .7 6 O.5 5 Fe(m-1, 7-CTH)(cn )2 0 .6 5 1. 28 Fe(m-1,7"CTH)(imidazole) 2 (C104 )2 0. 5^ 1. 3k a Isomer shift with respect to sodium nitroprusside. Again interpretation of the AEA values is limited in this series because the signs of electric field gradients are not known; however, it is interesting to note that as the donor strengths of the axial ligands decrease, the AEa decreases. This suggests that the axial ligands are stronger than the in-plane donors and continue in a smooth fashion evidenced by the near linearity of the plot in Figure 26. The compound rFe(m-CRH)OAc]PFs does not have the large AEa value similar to those of the other five-coordinate compounds; there fore, it is believed to be effectively six-coordinate with the acetate anion acting as a bidentate ligand. It is obvious from Table IX that the diacetate is quite different from the monoacetate as witnessed by 120 • Fe(m -1, 7-CTH)(imid)2(C104)2 • Fe(m-1, 7-CTH)(CN)2 1.2 ■ .8 - Fe(m-1, 7-CTH)(CH3CN)3(C104)2 • .4 . F e ( m - 1, 7C T H )(N C S)2 0 55 JO J5 i70 J5 F igure Plot of 6 vs AE q for Different Axial Ligands. 321 both its isomer shift ( 1. 56mm/sec vs. 1 . 19 mm/sec) and its AE ( 2 .27mm/sec vs. 2.2Qmm/sec) for the diacetate and the monoacetate respectively. The diacetato-complex may involve trans acetato-groups. The monoacetato-complex would have the macrocyclic ligand in a folded structure which is a well documented occurrence ( 3 7 , 3 9 )* The complex Fe(l,7~CT)Fe(NCS) 4 *2CH3CN is interesting because it contains the tetrathiocyanatoferrite anion. The peaks with the v a lu e s 6 =1. 13mm/sec and AEq = 1. 29 ran/sec correspond closely to the values reported for other Fe(NCS)42~ derivatives by Edwards, Johnson, and Williams ( 1O3 ). Mossbauer Spectra of some TAAB Complexes. - Several iron complexes of the macrocyclic ligand TAAB (structure III) have been pre pared in this laboratory by Dr. Vladmir Katovic. These complexes have a tendency to form oxobridged Fe-O-Fe dimers, however, it has been shown that this bridge can be broken by HF forming the corresponding monofluoride complex ( 2 6 ). The Mossbauer data for a series of these TAAB complexes is shown in Table XVII along with the room-temperature magnetic moments. The magnetic moments of oxobridged species are found to be lower than normal because of a superexchange mechanism. This phenomenon is well- known and well-documented in the literature ( 6 2 ,97~99)* The temperature dependence of these complexes is known ( 2 6 ) to follow the equation developed by Earnshaw and Lewis ( 6 9 ). 122 TABLE XVII • • M ossbauer S p e c tra o f Some Iro n Complexes of the Macrocyclic Ligand TAAB Compound ueff(BM ) 6amm/ sec AEqinm/ sec Fe(TAAB)F( C104 )2‘ 2H2 0 2 .1 .6 5 .8 6 2 .8 • 00 F e (TAAB)f (BF4 )2* 2H20 2 .k • 63b - y Fe(TAAB)F(N03 )2- 2H20 2 .2 .6 2 .8 8 0 .8 [Fe (TAAB) ] 2 0(C104 )4* 1^H20 2 .0 o . 6 l o .6 k 1 .9 [Fe(TAAB)]2 0(N03 )4 -^H20 1.89 0 .6 0 0 .7 0 k.'? [Fe(TAAB)]2 0(BPh 4 )4 2 . 1J+ 0. 59 0 .8 3 0 .7 [Fe(TAAB)(0Me)2]20 1 .9 0 o. 53 1 .0 5 1 .8 Fe(TAAB)(NCS)2 1 .3 1 o . 6 o 0 .2 7 1-3 SL Isomer shifts are with respect to sodium nitroprusside. Broad Spectrum. 123 The Mossbauer data are quite interesting^ however, they are inconclusive because of the number of compounds studied. It seems as though the oxobridged species fall into a range different from that of the nonbridged fluorides. The shifts of the oxobridged species are in the range from 0. 53 to 0 .6 lmm/sec while the fluorides are from 0.62 to 0. Spmm/sec. These values are in the range of high spin or low spin iron(ill) complexes. The oxobridged species contain high spin Fe(lll) atoms interacting through the oxobridge, thereby producing lower moments. The isomer shift values are lower in the oxobridged compounds than the fluoride compounds evidently because the oxobridge pulls more d-electron density away from the iron nucleus than does the fluoride, and thus more s-density is found at the nucleus. Fluoride is small and hard and does not tend to participate in bonding to as large an extent as does oxygen. Other oxobridge iron complexes give Mbssbauer parameters in the same range of values (97~99). The quadrupole splitting values of the oxobridged species are also lower than those of the fluorides. This indicates that there is less symmetry in the fluorides than in the oxobridges. The values for the oxobridged species are 0. 6^ - 0. 83 mm/sec while those of the fluorides a re 0. 86 - 0. 88 mm/sec. The complex rFe(TAAB)] 2 ° ( BPh 4 )4 has a AEQ value of 0. 83mm/sec which is much greater than those of the other oxobridged species. This is probably due to the bulkiness of the phenyl groups which make the 12k complex pack in a strained manner in the crystal lattice. This is also seen for other BPhi complexes. For example, the complexes Fe(1,3, 7 ,10-CT)(CH3 CN)2 (c104 )2 and F e ( l , 3 , 7 ,10-CT)(CH3CN)2 (BPh 4 )2 (Table X) demonstrate this effect. Both complexes should have the same quadrupole splitting because their inner spheres are identical. How ever, the AFn for the BPhi complex is 1. 50mm/sec while that of the perchlorate is 1.3Smm/sec, indicating less symmetry for the BPh 4 com plex. Another interesting feature of these compounds is the low isomer shift value and high quadrupole splitting value for the dimethoxy derivative. The structure of this ligand is shown below (structure XVII). This ligand forms the oxobridged complex but does not have water or anions to bond to the iron atoms in their sixth positions. Thus, these compounds are five-coordinate, which accounts for the high AEq value and also for the low isomer shift value. It has already been seen that five-coordinate complexes have high AEq values. The reason for the low 6 value in this case is that the ligand is now aromatic and can withdraw much d-electron density from the iron. The complex Fe(TAAB)(NCS )2 is a newly prepared compound for which the stoichiometry is uncertain. From the Mossbauer data it is thought to be a low spin Fe(ll) complex. The low value of the AEq is the main parameter which gives this impression. This value is the same as that for the low spin complex Fe(m-1, 7-CTH)(NCS)2. x c . OM e CH—N / M eO C XVII TAAB(OMe)2 Mossbauer Spectra of some Fe(TTP)2+ Compounds. - The Mossbauer spectra of iron(ll) complexes of a tetradentate macrocyclic ligand with sulfur being the ligating species (structure XVIIl) were also measured. Their parameters are presented in Table XVIII. The complexes [Fe(TTP)(CH3CN)2](BF4 )2 and Fe(TTP)(BF4)2 show typical low spin F e(ll) isomer shift values. The AE q values enforce the idea that these complexes are low spin F e(ll). This is also confirmed by the low values of the magnetic moments. It is very interesting to note that there is no quadrupole splitting for the complex [Fe(TTP)(CH3CN)2](BF4)2. Since this case normally arises only in pure cubic symmetry, octahedral or tetrahedral, 126 TABLE X V III Mossbauer Spectral Parameters for the Iron Complexes of TTP u e f f 8 mm/ se e 8. AEq [Fe(TTP)(CH3CN)2 ](BF4 )2 0 .7 6 0. 6k 0 [Fe(TTP)](BF4 )2 0. 82 0 .6 8 0.14-2 Ql Relative to sodium nitroprusside s s k k X V III T T P it is fortuitous that the acetonitrile ligands attach in such a way that they exert the same electronic influence on the iron nucleus as do the sulfur donors of the macrocycle. The complex [Fe(TTP)](BF4)2, however, shows a small quadrupole splitting indicating that the symmetry of the complex is reduced when the acetonitrile molecules are removed from the coordination sphere. SUMMARY The complexes of Fe(ll) and F e(lll) have been prepared and characterized with several tetradentate macrocyclic ligands with nitrogen donors. The Fe(ll) complexes were prepared in an inert atmos phere to prevent oxidation of the Fe(ll) to Fe(lll). The water of hydration from the iron(ll) salts did not react to a large extent in the synthesis of the complexes of Fe(ll) with m-CRH while in the preparation of the Fe(ll) complexes of m-1, 7“CTH, the presence of water led to the formation of undesirable hydroxides of iron. The Fe(lll) complexes were prepared by air oxidation of the Fe(ll) complexes in ethanol or acetonitrile solutions. This reaction takes place clearly under acidic conditions while in basic or neutral solutions oxidative dehydrogenation of the ligand was the end result at least in the case of the m-1, 7-CTH and 1,7-CT complexes of Fe(ll). Hydrolysis easily occurs in neutral or basic solutions, and thus desirable products were not obtained. The Fe(ll) and Fe(lll) complexes are for the most part stable toward oxygen in the solid state. A few of the compounds are sensitive to the air in the solid state such as the Fe(l, 7 -CT)Fe(NCS)4 * 2CH3CN which turns from maroon to black within minutes after exposure to the air. In solution all of the Fe(ll) complexes react rapidly with air (seconds) as evidenced by color changes and changes in the electronic s p e c tra . 127 128 The complexes are characterized by elemental analyses, infrared spectra, electronic spectra, magnetic susceptibilities, molar conduct ivities, and Mossbauer spectra. Elemental analyses are given in Tables I and II. The infrared spectra served as a tool to determine the presence of the macrocyclic ligand and of the anion. It also served to determine the absence of the Fe-O-Fe linkage. The electronic spectra served to determine the ground state and oxidation state of the iron. The high spin complexes of m-1,7-CTH complexes show strong tetragonal distortion, but for those of m-CRH it is quite small. The low spin complexes all show only a small amount of tetragonal distortion except for the complex Fe(m-CRH)(C10 4 )2 w hich does show a large tetragonal distortion in the solid state. The mag netic moments all fall in the expected ranges for high spin or low spin complexes except for the complexes Fe(m-1, 7-CTH)(cn )2 and Fe(m- 1 ,7“CTH)(C104 )2 * 3 CH3CN which have questionable moments of greater than one Bohr Magneton. The conductivity serves to demonstrate some of the solution properties of the complexes. Mossbauer spectra have been used in studying the coordination number in the solid state, and the spin state, and in correlating electronic properties of the new complexes and of a number of complexes prepared by other investigators. These iron complexes appear to have several coordination structures. Four-, five-, and six-coordinate structures are observed. Six-coordinate pseudo-octahedral is the most common structure for iron 129 complexes; however, several five-coordinate, presumed square pyramidal structures and a four-coordinate, presumed planar structure were also found. The five-coordinate complexes are unique in the fact that they are the first synthetic complexes that have high spin Fe(ll) in a square pyramidal structure similar to the biological complexes deoxy hemoglobin and deoxymyoglobin. These structures were considered rare, but now may be considered to fit properly into the systematic chemistry o f F e ( l l ) . Oxobridged species of iron were not found in the characterization of the three series of macrocyclic complexes of iron described here; however, the Mossbauer spectra of several Fe-O-Fe complexes of the macrocyclic ligand TAAB were measured and discussed. The Mossbauer spectra of the macrocyclic iron complexes show that as the unsaturation of the macrocyclic ligand increases, the strength of its influence on the d-orbitals of the iron increases. This implies a concommitant increase in their strength of bonding. In those cases where a large enough series of complexes of the same spin state and coordination geometry are available, spectrochemical series have been obtained for axial ligands with respect to a given macrocyclic ligand. APPENDIX RELRTIVE TRANSMISSION 90.00 92 50 95.00 97.50 100.00 y0 -.0 20 -.0 .0 .0 .0 .0 10 5.00 *1.00 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 -y.00 gur 2, Mosa r pcrm o Fe( CHC? T - 9° C 194° - T= -CRH)C1?, (m e F of Spectrum er ossbau M 27,. re u ig F | | I | | EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY V V % * + + + + + + + + + 4 - 4- * + + | I | | ] ■ * + * + 4 - ++ + + 4 »* ■H- VjJ H RELRTIVE TRANSMISSION 98.00 98.50 99.00 99.50 100.00 40 -.0 20 -.0 .0 .0 .0 .0 .0 5.00 4.00 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 -4.00 g e 8 Mosa r pcrm o Fem- Br2 )B H R -C e(m F of Spectrum er ossbau M 28. re igu F I I EOIY RLTV T NTGRSIE CMM/SECD NITRGPRUSSIDE) TO (RELATIVE VELOCITY | | I I I | | | | | + + 4 - ■*»- 4 * 4- + + 4 - + vS ro RELATIVE TRANSMISSION 98.00 98.50 99.00 99.50 100.00 I.Q 30 -.0 10 00 10 2Q 30 U0 5.00 U.00 3.00 2.Q0 1.00 0.00 -1.00 -2.00 -3.00 -IJ.QQ g e 9 Mosa r pcrm o Fem- )j, H R -C e(m F of Spectrum er ossbau M 29. re igu F I ] I | | I ] I EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY + . + + + 4 + - + + + + + * ■ + «■ + .+. ** * " + + + * 1 I | I | + + + + VJJ V>4 RELATIVE TRANSMISSION 97.00 97.75 98.50 99.25 100.00 I I I I I I I I I I I 1 M0 -.0 20 -.0 .0 .0 .Q .0 U.OQ 3.00 2.QQ 1.00 0.00 -1.00 -2.00 -3.00 -M.00 gur 3. fsae Setu f Fel 7CT) ] 04 l]Q )C T 7-C e(l, [F of Spectrum tfssbauer M 30. re u ig F EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY * » * * . 4- * + + + + + + + + + .00 RELRTIVE TRANSMISSION 94.00 95.50 97.00 98.50 100.00 40 -.0 20 -.0 .0 .0 .0 .0 4.00 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 -4.00 gur 3. sbue Setu f Fel 7C)lC0 T— 194°C — T 7-CT)Cl]C104 e(l, [F of Spectrum er ossbau M 31. re u ig F + J i i i r i i i i i i i ______EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY I ______+ + ♦ * v * * I + ______* 4 ++ - 4 4 , - - - ♦ + + + + I ______I ______I ______\ L • + *• + \ 4- + + . + .00 VJJ VJl RELRTIVE TRRNSMISSION 96.00 97.00 98.00 99.00 100.00 .0 30 -.0 10 00 10 20 30 M.00 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 U.00 g e 2 Mosa r pcrm o [ (, - )rC104 T)Br]C 7-C e(l, [F of Spectrum er ossbau M 32. re igu F | | | | EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELRTIVE VELOCITY % * * * * +* * + * * * * . * | | I | | I | [ + + 4 4 4 4 + ♦ • * * * * * */♦ H ^ 4 + t .00 RELRTIVE TRRNSMISSI ON 98.00 98.50 99.00 99.50 100.00 M0 -.0 20 -.0 .0 .0 .0 .0 >3.00 3.00 2.00 3.00 0.00 -1.00 -2.00 -3.00 -M.00 gur 3. sbue S crm o [ (, 7-CT)l]C104 e(l, [F of ectrum Sp er ossbau M 33. re u ig F * I I i I i i I I I I I * * * ** * * EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY v * + * ♦ v <• ■v + + + + + + / * * ** * * ■ * + ? * # t ^ } i + + + + + +• + 00 .0 RELRTIVE TRANSMISSION 97.00 97.75 98.50 99.25 100.00 40 -.0 20 -.0 .0 .0 .0 .0 .0 .00 S 4.00 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 -4.00 g e 4 Mosa r petu f Fel ,7,10-CT)C1]C104 , 7 3, e(l, [F of ectrum Sp er ossbau M 34. re igu F * + + * i i r i i i i i i EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY 4t- ^^ ++ + *v v * +* * + * + + ♦ ^ v * 4 4 4 * ♦ + + + + + 4 ♦ + 4 + 4 + ++ * * & ■ * + * ' V ^ - . v * * ^ ,*V + + * V ++ + + $ H + -H- + G 00 RELRTIVE TRRNSMISSI ON 93.00 94.75 96.50 90.25 100.00 40 -.0 20 -.0 .0 .0 .0 .0 4.00 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 -4.00 g e 5 Mosa r pcrm o [ (,3 0C)1C0 T - 194°C - T= 10-CT)C1]C104 , 7 3, e(l, [F of Spectrum er ossbau M 35. re igu F I I I | I I I I I I I EOIY RLTV T NT0R5IE CMM/SEC3 NITR0PRU5SIDE) TO (RELATIVE VELOCITY ++ + 4 4 4 T + + 4 4 & 4 + .++ ** * * * * * + + . ♦+ * * * v ♦ ♦ ♦ / v / A 4 4 4 4 5 4 5.00 RELRTIVE TRANSMISSION 97.00 97.75 98.50 99.25 100.00 ¥0 -.0 20 -.0 .0 .0 .0 .0 .0 5.00 1.00 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 -¥.00 F ig u re 36. M ossbau er Spectrum of [F e(l, 3, 7 , 10-C T)Br ]C104 T)Br 10-C , 7 3, e(l, [F of Spectrum er ossbau M 36. re u ig F ------t + + + + + + EOIY RLTV T NT0RSIE CMM/SECD NITR0PRUSSIDE) TO (RELATIVE VELOCITY 1 ------1 ------T + •**. V J + , T . + i ------1 ------1 ------1 ------+ + “ 4 W - J + t r > •p* o RELRTIVE TRANSMISSION 97.00 97.75 98.50 99.25 100.00 y0 -.0 20 -.0 .0 .0 .0 .0 U.OQ 3.00 2.00 3.00 0.00 -1.00 -2.00 -3.00 -y.00 gur 3. sbue S crm o [ (,3 , 10-CT)I]C104 7, 3, e(l, [F of ectrum Sp er ossbau M 37. re u ig F I 1 I I I I ” 1 I I I EOIY RLTV T NTQRSIE CMM/SECD NITRQPRUSSIDE) TO (RELRTIVE VELOCITY * ^ , H ^ + 4^ + ^ f 4 4 - * + - 4 * 4 + - + + + 4 *4 5.00 ,v • F ig u re 38. M ossbau er Spectrum of [F e(l, 3, 7, lO-CTjCllBPh* 7, 3, e(l, [F of Spectrum er ossbau M 38. re u ig F • RELRTIVE TRRNSMISSI ON 98.00 98.50 99.00 99.50 100.00 0 0 . 3 - EOIY RLTV T NTGRSIE CMM/SECD NITRGPRUSSIDE) TO (RELATIVE VELOCITY - 2.00 - 1.00 + + + 0.00 1.00 2.00 3.00 IJ.OQ 5.00 ro E- RELRTIVE TRRNSMI SSI ON 97.00 97.75 98.50 99.25 100.00 1 I I I I I I I I I I I "1 40 -.0 20 -.0 .0 .0 .0 .0 4.00 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 -4.00 gur 3. sbue Setu f ( CRH)OAc A )(O H R -C e(m F of Spectrum er ossbau M 39. re u ig F J ______EOIY RLTV T NTORSIE CMM/SECU NITROPRUSSIDE) TO (RELATIVE VELOCITY I ______I ______v + * * / 4- k l i + + * I ______# + 4 4 + 4 2 +■ + + «H d - - d «H I ______? ) T + + + * * \ + / \ * i I ------+ + 4 1 .00 RELRTIVE TRANSMISSION 96.00 97.00 98.00 99.00 100.00 40 -.0 20 -.0 .0 .0 .0 .0 .0 5.00 U.00 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 -4.0Q gur 40. sbue Setu f F( CRH) c]PF6 A )O H R -C [Fe(m of Spectrum er ossbau M . 0 4 re u ig F i i i i r i i i i i i EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY + + ♦ / ♦ * * 4 » *4 » +4 4 * 4 + + + ++ 4- + / £ H RELATIVE TRANSMISSION 8U.00 95.50 97.00 98.50 100.00 U0 -.0 20 -.0 .0 .0 .0 .0 .0 5.00 y.00 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 -U.00 gur 4. sbue Setu f mpr Fem-, - Cl2 )C H T 7-C -1, e(m F pure Im of Spectrum er ossbau M 41. re u ig F t i i r i i i i i EOIY RLTV T NTORSIE CMM/SECU NITROPRUSSIDE) TO (RELATIVE VELOCITY \ ; ; v AA + + + + f s j + + + \ ♦ + 5 V £• RELATIVE TRANSMISSION 95.00 96.25 97.50 98.75 100.00 40 -.0 20 -.0 .0 .0 .0 .0 .0 5.00 U.00 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 -4.00 gur 4. sbur pcrm o Pr Fe(,7CTH) l2 )C H T 7-C (l, e F Pure of Spectrum ossbauer M 42. re u ig F i i i i r i i i i i i EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY * + + + + 4 4 t 4 4 4 4 4 4 ' 4 4 4 . + 4. •$*■ * W + 4 4 + 4 4 £ CT% RELRTIVE TRRNSMISSION 99.00 99.25 99.50 99.75 100.00 U0 -.0 20 -.0 .0 .0 .0 .0 .Q 5.00 U.OQ 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 -U.00 gur 4. sbue Spetu f mpr Fem-, - Br2 )B H T 7-C -1, e(m F pure Im of ectrum p S er ossbau M 43. re u ig F i | | r | | | | | i i EOIY RLTV T NTORSIE CMM/SECU NITROPRUSSIDE) TO (RELATIVE VELOCITY +, i 7v +7 o ++ + + ++ + + + +++++ --3 E- RELRTIVE IRRNSMISSI ON 97.00 97.75 9 8 .SO 99.25 100.00 g e 4 Mosa r pcrm o Pr Fe( 1 - Brz )B H T 7-C -1, (m e F Pure of Spectrum er ossbau M 44. re igu F -3.00 EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY * H- -H 2.00 - 1.00 0.00 1.00 2.00 3.00 5.00 £ 00 RELRTIVE TRRNSMISSI ON 99.00 99.25 99.50 99.75 100.00 gur 4. sbue Setu f mpr Fe( 1 - )l2 H T 7-C -1, (m e F pure Im of Spectrum er ossbau M 45. re u ig F -3.00 EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY - 2.00 - 1.00 0.00 1.00 2.00 3.00 00 .0 4 5.00 RELRTIVE TRRNSMISSI ON 96.00 97.00 98.00 99.00 100.00 40 -.0 20 -.0 .0 .0 .0 .0 .0 5.00 4.00 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 -4.00 g e 6 Mosa r pcrm o F( CRH)C104)2 )(C H R -C Fe(m of Spectrum er ossbau M 46. re igu F t i i i r i i i i i i EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY " JL 4" + ■& + ■h 4 % ■*•- + + + +• vn H O RELRTIVE TRRNSMISSI ON 95.00 96.25 97.50 98.75 100.00 gur 4. sbue Setu f mpr Fe( 1 7CT (CS)2 )(NC TH 7-C -1, (m e F pure Im of Spectrum er ossbau M 47. re u ig F -3.00 EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELRTIVE VELOCITY - 2.00 - 1.00 0.00 * + 1.00 2.00 3.00 5.00 vn H RELATIVE TRANSMISSION 99.00 9 5 .SO 97.00 98.50 100.00 40 -.0 20 -.0 .0 .0 .0 .0 .0 5.00 4.00 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 -4.00 gur 4. sbue Setu f ue ( 1 - H)N S)2 )(NC TH 7-C -1, e(m F Pure of Spectrum er ossbau M 48. re u ig F EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY + + + + + + 4* 4* 4 - + + + + + + + \S\ to RELRTIVE TRRNSMISSION 95.00 96.25 97.50 98.75 100.00 40 -.0 20 -.0 .0 .0 .0 .0 .Q 5.00 U.OQ 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 -4.00 iue4. osae pcrmo em1 -T)C) H MossbauerSpectrum of Fe(m-1, 7-CTH)(CN)z Figure49. I I I I I I I I I I I EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY * *. ? 4 . J* + 4 + 4 4 4 ""*X *+ + + +*‘+ 4 4 4 + 4 ^ VJ4 \J1 RELATIVE TRANSMISSION 98.00 98.50 99.00 99.50 100.00 40 -.0 20 -.0 .0 .0 .0 .0 .0 5.00 4.00 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 -4.00 ______g e 0 Mosa r pcrm o Fe( l -T)C0) •3CH3CN 7-CTH)(C104)2 -lt (m e F of Spectrum er ossbau M 50. re igu F I I I I I I I I I I I I ______EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY L______I ______* \ I ______4 * - + M- + + 4+4 + + 4 + 4 4 * * 4* 4 4 + + 4 4 I ______4 I ______I ______I ______vn ■r=~ RELRTIVE TRRNSMISSI ON 98.00 9 8 .SO 99.00 99.50 100.00 40 -.0 20 -.0 .0 .0 .0 .0 4.00 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 -4.00 gur 51. re u ig F J I I I I I I I I I I I ______sbue Setu f m-,7CH(142 * 7-CTH)(C104)2 -1, (m e F of Spectrum er ossbau M EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY I ______I ______-H* + v ^ I ______■»». + + ■»». I ______Z midazole Im I ______I ______I 0 0 . vn vn RELRTIVE TRANSMISSION 1 I l I I r I I I I l l I I “199.00 99.25 99.50 99.75 100.00 40 -.0 20 -.0 .0 .0 .0 .0 .Q 5.00 U.OQ 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 -4.00 gur 5. sbue S crm o Fe(, - ( S)2 C )(N T 7-C (l, e F of ectrum Sp er ossbau M 52. re u ig F + + EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY ' * * J % \ 4 -K«- 4 ■H- + * ++ * 4 + w RELRTIVE TRRNSMISSI ON 98.00 98.50 99.00 99.50 100.00 .0 30 -.0 10 00 10 20 30 40 5.00 4.00 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 4.00 F ig u re 53. M ossbau er Sp ectrum of [F e(l, 7-CT)CN]BPh4 e(l, [F of ectrum Sp er ossbau M 53. re u ig F j i i i i r i i i i i i i I ______EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY : ______i ______+ + > ++ i ______* * ' * * 4 * - + ** * + i * ______i ______i ______i relrtive transmission 97.00 97.75 9:1. 99.35 100.00 40 -.0 20 -.0 .0 .0 .Q .0 .Q 5.QQ U.OQ 3.Q0 2.QQ 1.00 0.00 -1.00 -2.00 -3.00 -4.00 gur 5. sbue Setu f Fei - )CH3CN)2](ClQ4)? H T)(C 7-C e(i, [F of Spectrum er ossbau M 54. re u ig F *+++ ♦ * \ * V * 4 ♦ + + + * v + J I I I I I I | I I I I ______EOIY RLTV T NTORSIE CMM/SEC3 NITROPRUSSIOE) TO (RELATIVE VELOCITY I ______I ______♦ * ♦ ♦♦ ♦ + ♦ ♦ I ______+ t 4 ♦ - + + +• I ______t ■ t +■ ++ I ______I ______I NJ1 M 00 RELRTIVE TRANSMISSION 97.00 97.75 90.50 99.25 100.00 40 -.0 20 -.0 .0 .0 .0 .0 4.00 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 -4.00 F igu re 55. M ossbau er Spectrum of [F e(l, 7 -C T (iso m e r-2 ) ](CH3CN)2(C104)2 ) r-2 e m (iso T -C 7 e(l, [F of Spectrum er ossbau M 55. re igu F J j j I | J I | j | j | ______EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY I ______I ______+ \ + + + * 4 + + I + - ______+ 4 - 4 4 - + - + + + r + / ++ I 4 ______4 * 4 * 4 -+ 4 - - t I ______I ______I .00 \J1 VO H RELRTIVE TRRNSMISSI ON 98.00 98.50 99.00 99.50 100.00 40 -.0 20 -.0 .0 .0 .0 .0 U.00 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 >4.00 I I I I I I I I I I I I I gur 5. sbue S crm o Fe(, CT)i daz le)2(BPh4)2 zo a id )(im T -C 7 (l, e F of ectrum Sp er ossbau M 56. re u ig F EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY *** + % * «•** 4 - + + + 4 4 - 4 + - + * - * 4- 4 + - 4 -+ 4 - K .00 o RELRTIVE TRRNSMISSI ON 99.00 99.25 99.50 99.75 100.00 40 -.0 20 -.0 .0 .0 .0 .0 4.00 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 -4.00 g e 7 Mosa r pcrm o Fe(,7CT( 3CN)2l2 H T)(C 7-C (l, e F of Spectrum er ossbau M 57. re igu F + + * + + *+ I | | | J | J | j I p EOIY RLTV T NTORSIE EMM/SEC3 NITROPRUSSIDE) TO (RELATIVE VELOCITY f. + + ■ft. 4 A - + * + -*L ^ * * x ■H* + ^ - H ^ + ; * * * + + + *+ % * / v / * % *+ *+ q * y .00 RELRTIVE TRRNSMISSI ON 97.00 97.75 98.50 99.25 100.00 40 -.0 20 -.0 .0 .0 .Q .0 .0 5.00 U.00 3.00 2.0Q 1.00 0.00 -1.00 -2.00 -3.00 -4.00 gur 5. sbue Setu f Fel , 10-CT)(CH3CN)2](C104)2 , 7 3, e(l, [F of Spectrum er ossbau M 58. re u ig F I I I I I I I I I I I EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY + ++ ++ + ■ - 4 + 4 4- +■ ♦ ♦ ♦ 4 + 4 + * 4 4 + * * 4 4 + 4 + RELATIVE TRANSMISSION 97.00 97.75 98.50 99.25 100.00 40 -.0 20 -.0 .0 .0 .0 .0 .0 5.00 4.00 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 -4.00 F igu re 59. M ossbau er Spectrum of [F e(l, 3, 7, 10-CT)(CH3CN)2](BPh4)2 7, 3, e(l, [F of Spectrum er ossbau M 59. re igu F j j j j j j I j EOIY RLTV T NTORSIE CMM/SEC3 NITROPRUSSIDE) TO (RELATIVE VELOCITY w \ \ * > * * * * * ♦ ♦ + * ♦ +* % + 4* * * * ♦ RELATIVE TRANSMISSION 99.00 99.25 99.50 99.75 100.00 40 -.0 20 -.0 .0 .0 .0 .0 4.00 3.00 2.00 3.00 0.00 -1.00 -2.00 -3.00 -4.00 g e 0 Mosa r petu f Fel ,7 - )CH3CN)2](BPh4)2 H T)(C 9-C 7, 4, e(l, [F of ectrum Sp er ossbau M 60. re igu F I I I I I I I I I I I + EOIY RLTV T NT0RS1E CMM/SECD NITR0PRUSS1DE) TO (RELATIVE VELOCITY + % + % + ♦ + ■+ ++ +■++ ■& + + + + ++ + + + 1 V f ^ 4 + *w+ w +* + A + V + .++T ♦+ + + 4+i * « ■ -** 4 V + + 5T0Q RELRTIVE TRANSMISSION 93.00 99.00 96.00 98.00 100.00 40 -.0 20 -.0 .0 .0 .0 .0 .0 5.00 U.00 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 -4.00 iue6. MossbauerSpectrum of Fe(l, 3, 7,10-CT)(NCS)2 Figure61. | | | EOIY RLTV T NTGRSIE CMM/SECU NITRGPRUSSIDE) TO (RELATIVE VELOCITY * X + * + % + * * + - ■ +■ +■ 4- + 4 * 4 4 J | | | | J | + - - + 4 4- + - ++ 4 - VJ1 o> h RELATIVE TRRNSMISSI ON 98.00 98.50 99.00 99.50 100.00 40 -.0 20 -.0 .U .0 .0 .0 4.00 3.00 2.00 1.00 O.OU -1.00 -2.00 -3.00 -4.00 g e 2 Mos ue Setu f l ,7 1- (mi azole)2(C104)2 id )(im T 10-C 7, 3, (l, e F of Spectrum er au ossb M 62. re igu F | | | | | | EOIY RLTV T NTGRSIE CMM/SECD NITRGPRUSSIDE) TO (RELATIVE VELOCITY + * ¥ + + ¥ ¥ + +¥ j * + .+ * | | | | | + ♦ + ¥ ¥ ¥ 5.00 RELRTIVE TRANSMISSION 98.00 98.50 99.00 99.50 100.00 M O -.0 20 -.0 .0 .0 .0 .0 y.00 3.00 2.00 3.00 0.00 -1.00 -2.00 -3.00 -M.OQ gur 6. sbue Setu f l ,7 1- (mi azole)2(BPh4)2 id )(im T 10-C 7, 3, (l, e F of Spectrum er ossbau M 63. re u ig F ;++ v J I I I I I I I I I I I ______EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY I ______I ______4 + 4 4 I ______ I ______4 4 I ______I ------1 .00 -5 ON RELRTIVE TRANSMISSION 96.00 97.00 98.00 99.00 100.00 40 -.0 20 -.0 .0 .0 .0 .0 >1.00 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 -4.00 gur 64. re u ig F i i i i i i i i i i i sba r petu f F( CRH) 12]BF4 )C H R -C [Fe(m of ectrum Sp er au ossb M EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY ■ J^* ^ J +■. t + + + ,4- + + 4 > ; 4 4 44 4 + 4 4 4 4 4 4- 4 4* .00 CD O « o o _ -k r+ 4^ ■**■++■*■ 4 O ++ + ++4 o in cn I— O) _ O) LU I— CD I CD _ UJ o> CC CD O CO m 3.00 - 2.00 - 1.00 0.00 2.00 3.00 5.00 VELOCITY (RELATIVE TO NITROPRUSSIDE) CMM/SECD Figure 65. Mossbauer Spectrum of ifFei(m- 2 ]C10 4 H RELATIVE TRANSMISSION 98.00 98.50 99.00 99.50 100.00 40 30 -.0 .0 O J0 20 30 40 5 4.00 3.00 2.00 J.00 OU U 1.00 -2.00 3.00 -4.00 gur 6. sbue S crm o [ ( 1 - Br )B H T 7-C -1, e(m [F of ectrum Sp er ossbau M 67. re u ig F i i i i i i i i i i i EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY * V " \ . V / \ + + + + ^ ^ +■ £ * ' * J 2 ]BF 4 .00 171 RELATIVE TRANSMISSION 95.00 95.25 97.50 98.75 100.00 U0 -.0 20 -.0 .0 .0 .0 .0 .0 .00 S M.00 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 -U.00 F ig u re re u ig F * + + * 68 Mosa r pcrm o [ (, CT) l )C T -C 7 e(l, [F of Spectrum er ossbau M . EOIY RLTV T NTORSIE CMM/SECU NITROPRUSSIDE) TO (RELATIVE VELOCITY V * V W V L+- + -L T V * * + '^ - « f 4’+ + + “*HH 4’+ f « - '^ + 2 ]C10 4 **** 172 RELATIVE TRANSMISSI ON 99.00 99.25 99.50 99.75 tOO.OO 40 -.0 20 -.0 .0 .0 .0 .0 .O 5.00 U.QO 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 -4.00 F ig u re 69- M ossbau er Spectrum, of [F e(l, e(l, [F of Spectrum, er ossbau M 69- re u ig F * * *&% & %* - A + f V + t + 1 i i i r i I i i i i v * % s EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY * - * + 7 CT)NCS) C )(N T -C 2 ]BPh 4 + 4 ^ - T . +■*+ + . + 4 4 + *• + + * • * ++ + RELATIVE TRRNSMISSI ON 98.00 98.50 99.00 99.50 100.00 F igu re 70. M ossbau er Spectrum of [F e(l, e(l, [F of Spectrum er ossbau M 70. re igu F -3.00 EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY - 2.00 - 1.00 0.00 ++-** + 7 CT( H T)(C -C 1.00 3 CN) + + + 2 ](C 104)3 2.00 3.00 5.00 RELRTIVE TRRNSMISSI ON 95.00 95.25 97.50 98.75 100.00 yO -.0 20 -.0 .0 .0 .0 .0 U.00 3.00 2.00 3.00 0.00 -1.00 -2.00 -3.00 -y.OQ gur 7. sbue Setu f Fel - ( S)] C )(N T 7-C e(l, [F of Spectrum er ossbau M 71. re u ig F I I I I I I I I I I I EOIY RLTV T NTORSIE CMM/SEC3 NITROPRUSSIDE) TO (RELATIVE VELOCITY + V t * 4 ^ *■ *■ 4 +4 : * 4 4 4 4 4 * 4 * * t 4 * 2 0(NCS w )2 .00 RELATIVE TRANSMISSION 97.00 97.75 98.50 99.25 I00.00 40 -.0 20 -I.0 00 10 20 30 40 5.00 4.00 3.00 2.00 1.00 0.00 00 . I - -2.00 -3.00 -4.00 gur 7. sbue Setu f l - FeNCS C e(N )F T 7-C (l, e F of Spectrum er ossbau M 72. re u ig F EOIY RLTV T NTORSIE EMM/SEC3 NITROPRUSSIDE) TO (RELATIVE VELOCITY t i i r i i i i i * 4 + 4 4 + * + * * *+ + 4 + 4 \ 4 + ♦ * * )4 V + + V aroon M + RELRTIVE TRANSMISSION 98.00 98.50 99.00 99.50 100.00 40 -.0 20 -.0 .0 .0 .0 .0 .0 5.00 4.00 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 -4.00 gur 73, re u ig F + V + + + + V + + + +* , ++ +4. 4 + + + , * + , * v H - + v + / sbue Spetu f l - FeNCS C e(N )F T 7-C (l, e F of ectrum p S er ossbau M EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY * u % +4 i & * )4 psd oAi Black) ( ir A to xposed E \ + * * & * ? * RELATIVE TRRNSMISSI ON 36.00 96.50 99.00 99.50 100.00 1 I I I I i I I >1.00 3.00 I 2.00 1.00 I 0.00 I -1.00 -2.00 I -3.00 -U.00 “1 F ig u re 74. M ossbau er Sp ectrum of [F e(l, 3, 7, 10-CT)(Phen)](C10 7, 3, e(l, [F of ectrum Sp er ossbau M 74. re u ig F EOIY RLlIE O IRPUSD) CMM/SEC3 NITROPRUSSIDE) (RELflTIVE TO VELOCITY 4 4- 4 * 4 4 + + * 4)2 .00 RELRTIVE TRANSMISSION 95.00 95.25 97.50 98.75 100.00 40 -.0 20 -.0 .0 .0 .0 .0 .0 5 4.00 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 -4.00 gur 7. sbue S crm o [ (,7- )] T -C 7 e(l, [F of ectrum Sp er ossbau M 75. re u ig F I I I I I I I I I I I EOIY RLTV T NTORSIE CMM/SEC3 NITROPRUSSIDE) TO (RELATIVE VELOCITY \ + + * + * + A + 4 4 - - + ♦ + 2 C0)‘ H20 (C104)3‘ AE 6 q 11 mm/ c e /s m m 1.16 = % 3. sec e /s m m 4 .4 3 = * * 4 + - 4 - ++ .00 RELATIVE TRANSMISSION 96.00 97.00 96.00 99.00 100.00 U0 -.0 20 -.0 .0 .0 .0 .0 U.00 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 -U.00 gur 7. sbue Spetu f Fe(TAAB)F(C10 of ectrum p S er ossbau M 76. re u ig F + + + 4 i i i i i i i i i i i 4 4.4 ++ 4 4 ++ 4.4 4 EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY ' * * + ^■+++4 + *"V 4++ * ++^ + + + ■ X +■ ; * v +4 4)2 «* •2H20 * ^ + + + + 5.00 H 00 o RELRTIVE TRANSMISSION 98.00 98.50 99.00 99.50 100.00 U0 -.0 20 -.0 .0 .0 .0 .0 .0 5.00 y.00 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 -U.00 gur 7. sbue Spetu f FeTAAB) ( F ](B )F B A A e(T [F of ectrum p S er ossbau M 77. re u ig F + + * + i i i i r n i i i i i i + + ++ ++ + + EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY \ + \ 4)2 . + + 4 +■< -4^ .+ 4- + ++ + + + + ++ H H oo RELATIVE TRANSMISSION 98.00 90.50 99.00 99.50 100.00 gur 7. sbue Spetu f eTAAB)( Q )F(N B A A Fe(T of ectrum p S er ossbau M 78. re u ig F 3.00 EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY - 2.00 1.00 0.00 1.00 3)2 2 20 -2H 2.00 3.00 + % + 4" * \ + + + \ * 4" + % + 5.00 to oo RELRTIVE TRRNSMISSI ON 97.00 97.75 98.50 99.25 100.00 I I I I I I I I I I I 1 yO -.0 20 -.0 .0 .0 .0 .0 4.00 3.00 2.00 3.00 0.00 -1.00 -2.00 -3.00 -y.OQ gur 79. re u ig F sbue Setu f FeTAAB)] B A A e(T [F of Spectrum er ossbau M EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY \ * % * * * * % 4 * ^ + 4 2 O(C10 ■w 4 4 4 + 4)4 4 20 -4H + 5.00 RELRTIVE TRANSMISSION 95.00 95.25 97.50 98.75 100.00 40 -.0 20 -.0 .0 .0 .0 .0 .0 5.00 4.00 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 -4.00 0 Mossbauer Spectum, f[ TAAB)2{C) - 20 H -4 )]20{NC^)4 B A A (T e [F of , m tru c e p S r e u a b s s o M 80. e r u g i F 4 4 .1% .1% 4 4 I I I I I I I I I I I EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY •** - * V * * ■*■ 4 * > A + 4 4 4 4 4 4 4 4 4 4 4 4 4 4 4 4 4 4 4 4 + l g H RELRTIVE TRANSMISSION 99.00 99.25 99.50 99.75 100.00 gur 8. sbue Setu f FeTAAB)] B A A e(T [F of Spectrum er ossbau M 81. re u ig F 3.00 EOIY RLTV T NTORSIE CMM/SECU NITROPRUSSIDE) TO (RELATIVE VELOCITY - 2.00 - 1.00 0.00 +*■ 20 (BPh 3.00 4)4 2.00 3.00 ... • * > 5.00 RELRTIVE TRRNSMISSI ON 98.00 98.50 99.00 99.50 100.00 I 1 I I I I I 1 I I I 1 40 -.0 a0 -.0 0.00 -1.00 -a.00 -3.00 -4.00 ______gur 8. sbue S crm o { [ 0 e)2]}20 (0M B A A e[T {F of ectrum Sp er ossbau M 82. re u ig F * + * ♦ * ♦ * + ■** + + ♦ J ______* EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELRTIVE VELOCITY * * + + I ______♦ A + v + * ♦ +♦ + < + ♦ + ♦+ *** + +v + + +A +♦ I ______* 4 4 4 t i I ______4 4 4 ______V * 4 4 1.00 4 I 4 ______4 4 4 4 4 .0 .0 4.00 3.00 2.00 I ______I ------L +++ +++ -ft*.+ 4 4 5.00 RELATIVE TRANSMISSION 96.00 98.50 99.00 99.50 100.00 40 -.0 20 -.0 .0 .0 .0 .0 4.00 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 -4.00 gur 8. sba r pcrm o F(AAB( S C B)(N A Fe(TA of Spectrum er au ossb M 83. re u ig F I I I I I I I I I I I EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY + * *+ * \ * ♦ * s - * *■ * \ *■ + + + + + + + + V £ / )2 5.00 00 b* RELRTIVE TRANSMISSION 96.00 96.50 99.00 99.50 100.00 gur 8. sbue Setu f ( ( F )(B P T e(T F of Spectrum er ossbau M 84. re u ig F -3.00 EOIY RLTV T NTORSIE CMM/SECD NITROPRUSSIDE) TO (RELATIVE VELOCITY - 2.00 - 1.00 0.00 4)2 1.00 2.00 3.00 5.00 RELATIVE TRANSMISSION 97.00 97.75 98.50 99.25 100.00 .0 30 -.0 0 00 10 20 30 4.00 3.00 2.00 1.00 0.00 .00 -2.00 -3.00 4.00 gur 8. sbue Spetu f eTTP( H P)(C T Fe(T of ectrum p S er ossbau M 85. re u ig F I I I I I I I I I I I EOIY RLTV T NT0RSIE CMM/SECD N1TR0PRUSSIDE) TO (RELATIVE VELOCITY * + * i * 4 4 4 4 4 + 4 4 ■H- 3 $ CN) 2 BF (B 4)2 L .00 RELRTIVE TRRNSMISSI ON 84.00 88.00 92.00 96.00 100.00 40 -.0 20 -.0 .0 .0 .0 .0 4.00 3.00 2.00 1.00 0.00 -1.00 -2.00 -3.00 -4.00 gur 8. sbue Setu f oim Nir usi Standard e ssid ru p itro N Sodium of Spectrum er ossbau M 86. re u ig F I I I I I I I I I I I EOIY RLTV T NT0RSIE CMM/SEC3 NITR0PRUSSIDE) TO (RELRTIVE VELOCITY 4 + 4r 4 + 4 4 4 + 4 - 4 4- + + 4 4 4 4 + + + 4 4 4 4 4 4 .00 RELRTIVE TRRNSMISSI ON 68.00 76.00 84.00 92.00 100.00 40 -.0 20 -.0 .0 .Q .0 .0 .0 5.00 4.00 3.00 2.00 l.QQ 0.00 -1.00 -2.00 -3.00 -4.00 gur 87. re u ig F sbue Setu f 2Qj e F - a of Spectrum er ossbau M EOIY RLTV T NT0RSIE CMM/SECD N1TR0PRUSSIDE) TO (RELATIVE VELOCITY Mv+ WMwv y BIBLIOGRAPHY 1. 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