Leonard Klevan, Joseph Peone, Jr., and Stanley K. Madanl
SUNY ot Binghomton Binghamton, New York 13901
Molecular Oxygen Adducts of Transition Metal Complexes Structure a n d mechanism
The natural oxygen carriers of biological systems are transition metal complexes which are able to bind reversihlv to molecular oxwen. While most of these natural complexes (such as the ferrous protoporphyrin group in hemoglobin) ( I ) tend to be quite intricate, there are a number of synthetic systems whose relative simplicity has led to intensive investigation of their chemical properties. The importance of these model systems and their novel chemistry should be of general interest. This review is not intended to he a thorough and complete one on oxygen adducts. I t is designed to provide a general, and in some representative cases, an in-depth study of molecular oxygen adducts of transition metal complexes. Synthesis of the oxygen carrier, chlorocarbonylbis(tripheny1phosphine)iridium(I), [IrCI(CO)(Ph3P)z], by Vaska (2) in 1963 bas led to intensive investigation of numerous iridium and rhodium analogs. Cobalt complexes with a cobalt oxygen ratio of 1:1 (monomeric) and 2 : l (dimeric) have been investigated as have complexes of iron, ruthenium, nickel, palladium, and platinum. An early honding scheme which has found application in the analysis of some of these complexes will first he considered and some recent coordinated oxygen adducts of these metals will tben be examined.
Similarly, ionization of an electron from a (ruZp,) orbital was calculated by Griffith using molecular spectra approximations with a resulting value of I (ru2pJ = 16.6 eV. Therefore, based on considerations such as these, the bonding was determined to he a donation of a-electrons from oxygen to a dzsp3 orbital on Fe(I1). This orientation of the oxygen molecule parallel to the heme group also allows for back donation of electron density from the filled d,, orbital on the iron to the empty (rg2pJ antibonding orbital on the oxygen ligand. The resulting structure would tben have the nature of a double bond (Fig. 1).For clarity, the coordinate system describing the oxygen molecule has been changed to that of the Fe(I1) group (different systems were used in the original article).
Bonding
In an attempt to explain the various magnetic properties of all hemoglobin (Hh) compounds, Griffith (1) proposed a bonding scheme which has since been extended (3) to include complexes of the irridium and rhodium analogs mentioned above. In this model, the oxyhemoglobin (HbOz) is composed of submolecules of hemoglobin and oxygen, each of which is in its "valence" or prepared state (all electrons paired). For Hb, this would correspond to a low spin Fe(I1) ion with six electrons in the tz, orbitals. The oxygen submolecule may be represented by the following molecular orbital representation in which the degeneracy of the a-antibonding orbitals have been removed(2).
This structure is isoelectronic with that of ethylene, and the consequent spatial orientations of the spZ hybrid may be applied to the diamagnetic oxygen molecule. The regions of maximum electron density lie between the oxygen atoms and along the lone pair orbitals. This implies that a coordinated u-bond must form in either a 90 or 60" angle from the internuclear oxygen-oxygen axis (x-axis). Lone pair honding was suggested by Pauling in 1949 (corresponding to the spz 60" angle), but calculations by Griffith favor the a-electron bonding scheme (90"). Using the approximation that a spz orhital is nne-third "s" character and two-thirds "p" character, we have for one particular orbital Author t o whom inquiries should he addressed. 670
/ Journal of Chemical Education
Figure 1 . The nafure of (Hb-02) bonding in oxyhemoglobin (from ref. ( 1 ) ) . (a) Mode of o-bonding (a,2pz d'sp'). (b) a-bond formation ( 3 6 , ~ ~&'Pz).
-
-
Griffith's mode of bonding is found to he applicable to Vaska's complex [IrCI(CO)(P(CaH&)z] and similar rhodium coordination compounds. Ir(1) is a ds metal which may be considered to exist in an initial dspZ configuration but which forms a dsp3 hybrid (containing an empty orbital) upon assuming the "prepared" excited state (3, 4). The oxygen then forms a u-bond with this orbital through use of its a-electron density, and this sets u p a partial dipole in the oxygen molecule. The dipole is then reduced through back donation (=-bond formation) from the metal ion. The extent of hack donation which occurs will depend upon the relative symmetry and degree of overlap between the ligand and metal orbitals to determine the resulting oxygen-oxygen bond order (5). This effect will he explained in more detail in the next section. Complexes of Iridium and Rhodium
The compound [IrCI(CO)(Ph3P)2] (Ph = C6H5) which was cited previously as the first iridium compound to undergo extensive investigation is insoluble in most polar solvents but undergoes a color change from yellow to orange when oxygenated in benzene (2). X-ray diffraction analysis has shown that all such diamagnetic oxygen complexes display equal bond lengths between the oxygen and iridium atoms with the oxygen assuming a triangular
structure (6). Both of these facts may he rationalized in terms of Griffith's bonding scheme. This triangular structure may be visualized as either of the two forms
unidentate isosceles where "M" is central metal. By comparison of the oxygen-oxygen stretching force constants (a larger force constant indicates shorter hond lennth and nreater enerev) -.. (7) . . for the excited states of the fie; molecure and the molecule in a coordinated complex, the two possibilities above m a s he distinnuished (8). Sienificant hack honding results in a decrease in hond order since the electrons are being fed into an antibonding orhital. This delocalized honding would he the isosceles model and should correspond to a decrease in the force constant. In the unidentate model we would expect very little deviation from the value of 11.5 mdyn1.A for the uncomplexed oxygen molecule due to the retention of the double hond. The reported values for a series of dioxygen-transition metal complexes (for example, [Pt(02)(P(CsH5)&] show a range of 3.0-3.5 mdynlA which concurs with the notion of back *-donation and the reduction of hond order displayed in the isosceles model. This is close to the value ohsewed for oxygen in an excited state (3Z,+or 32,-) and less than that of 0 2 2 - (force constant = 5.4) which corresponds to a bond order of one. It is also pointed out that the infrared vibrations urn-, and u,-, are orthogonal in the unidentate model and would therefore he expected to he mutually independent (8). As shown in Tahle 1, there
oxygen and form the original compound) may be related to the low iridium-02 hond order (one), owing to an "insufficient" hack donation of only one unit of electron density from the metal to the oxygen. This prevents further reduction of the metal. Even though the adduct can he 0 2 - = -2). formally viewed as an Ir(II) moiety (Clthe diamagnetism of the starting complex and the adduct makes such an assignment questionable. The reason for this hesitancy is that Ir(I1) is a d7, paramagnetic entity as is the odd electron superoxide, 02-, species. One would therefore predict that the crystallographic data should he paramagnetic (unless the Ir(I1) and 0 2 - odd electrons share a common molecular orbital-this is pure rationalization). As of this time, it remains an open question as to how the compound can he diamagnetic and contain the superoxide species (12). Nevertheless, the structure may be considered as intermediary between five- and six-coordinate. Tahle 2 further illustrates this point by comparing the geometries of the six-coordinate hydrogen adduct and the five-coordinate carbon monoxide adduct to that of the dioxygen adduct of [IrCI(CO)(PhP)z].
+
Table 2. Geometry of the Iridium Complex with Simple Gaseous Molecules (from ( 9 ) )
IrCI(CO)(PPh3)2 + XY =IrCI(XY)(CO)(PPh3)2 Bond Angle Bond Angle XY (1r.X-Y) (Cl-Ir-CO) Geometry CO
Hz
45" 180"
0 2
72"
90" 120' 100'
d2sp3 (trig. bypyramid) dsp3 (octahedral)
Intermediate between
Table 1. Halogen Substituent Effects in the Iridium-09 ComDlex
Bond- T.enpths -~---
~~
5-0 (A)
Complex [IrCI(CO)(Ph3P)*]
1.30
[IrI(CO)(Ph3P)21
1.51
I
(A)
2.09 2.04 2.035 ? na?
Assigned )
Uptake Properties
858
Reversible
862
Irreversible
v
c
'Extracted from Table VIII, Reference (5). From Reference (3,6). is a change of only 4 cm-1 in the absorption corresponding to an 0-0 hond difference of approximately 0.2 ,& in the iodo and chloro analogs of Vaska's compound. This has been attributed to a mixing of these stretching vihrations (3). This is further evidence for the isosceles model and a resulting Czu symmetry. Reactions involving '80 isotope substitution has demonstrated that "oxygen atom scrambling" does not occur upon reversible oxygenation of the chloro-iridium complex (6). Below is a two dimensional representation of the oxygen addition reaction which normally takes place in benzene. The triphenylphosphine groups are not present hut would he normal to the plane of the paper.
c1
It has been shown (from 3-dimensional X-ray data ohtained from a single crystal) that the oxygen-oxygen hond 0.03 ,& which is longer length in this complex is 1.30 than obsewed in molecular oxygen (1.21 A) yet less than in a peroxide (1.49 A) (10). The bond length corresponds with that of the superoxide 0 2 - (1.28 A) which indicates hack donation in the amount of one electron. Thus the oxygen reversibility of this compound (ahility to give up
*
Other substituted ligands on the complex may affect the degree of overlap hetween the metal and oxygen molecule causing a consequent lengthening or shortening of the oxygen-oxygen hond. For instance, when iodine is suhstituted for the more electronegative chlorine in Vaska's compound, it allows for a greater degree of back donation which increases the oxygen-oxygen hond length to 1.51 A which resembles that of the peroxide 0 z 2 (5)). If an "inert" gas is passed through an [OzIrCl(CO)(Ph3P)2] solution, the reversible properties of this complex become evident and the oxygen free compound is ohtained. The formation of [02IrI(CO)(PhsP)2] however, is irreversible and does not yield the nonoxvnenated product upon similar treatrnpnt ( I , . These d,hervarioni are in agwernent with the pr~\.iousnotion ut' reversihilitv as it relates t o rhp degree of hack donation. Similarly, the irreversible complex dioxygenbis(bis(diphenylphosphino)ethane)iridium(I) hexafluorophosphate [O~IT((C~H~)~PCH~CH~P(C~H~ [PFs] (0-0 hond length 1.625 A) and its reversihle rhodium analog (0-0 hond length 1.418 ,&) have been prepared (5). If the d orbitals in the rhodium complex are considered to he of higher energy than those of iridium (as suggested in reference (5)) which may he due to the lanthanide contraction, then one would expect less efficient overlap with the metal orbitals in the rhodium complex and a reduction in hack donation. The oxygen-oxygen hond strength would then he greater in the rhodium adduct, and it would he expected to exhibit a higher degree of reversibility with oxygen (which i t does). Some hack donation is present, however, since the 0-0 hond distance is still greater than that in uncomplexed oxygen (1.21 A). The kinetics of oxygenation in Vaska's analogs have been studied (9, 13, 14) and they have been found to obey a second-order rate law. The chemical reaction and rate expressions are
Volume 50, Number 10, October 1973 / 671
Table 3. Kinetic and Activation Data tor the Reversible Oxygenation of Some Iridium Complexes in Solution as a Functionof the Anionic Ligand A in IlrA(CO)(Ph,P):I (from ( 1 4 ) )
where L = tertiary phosphine, A = univalent anionic ligand, e.g., F, CI, Br, I . . .and -dr1rA(C01L21 rlt .
=
h2[IrA(CO)L,I[0,1
Table 3 summarizes some of the kinetic and thermodynamic data for A = halogen and L = (CeH5)sP. The data were obtained ir: chlorobenzene solution a t 40"C, and the results are in agreement with previously reported trends. The oxygenation reactions are exothermic [AHzo's are negative (AHzo = AHz* -AH-I*)], and AHzo's range from -10 for F to -23 kcal mole-' for I. As is expected for simple addition reactions, the entropy changes, ASP, are highly negative (A&" = ASz* - AS-I*). They range from -24 for F to -47 kcal mole-* for I. The entropies of activation, ASz* are also negative and their magnitudes are close to those of the corresponding A S P . This implies that the transition state configurations resemble the structure of the adducts. However, since oxygen is usually present in large excess and its concentration varies only slightly, the reactions may be considered pseudo-firstorder. The high negative entropy factor may be accounted for by either polarity in the transition state or stereochemical factors in the formation of the transition state (13). When these reactions are carried out in dimethylformamide (a more polar solvent than benzene) an increase in reaction rate is noted which would tend to favor the notion of induced transition state polarity in accord with Griffith's bonding ideas. Finally, Vaska has recently reported the synthesis of a compound [M(2 = phos)~]Awhere M may be cobalt, rhodium, or iridium; A is CI, I, BF4, or B(CsHd4; and 2 = phos is cis-1,2-bis(diphenylphosphino)ethylene (15). Electronic and magnetic measurements have shown that the complexes of all three ds metals are square planar prior to activation with oxygen and assume the same type of configuration as the triphenylphosphine complex upon oxygenation. This represents the first synthesis of molecular oxygen coordinating compounds from three metals in the same family of the periodic table and they display the reA correlation is noted between the nrtivitv C- o- > Ir -- > Rh. "electron excitation energy" of the complexes and the enthalnv of activation due to oxveen ." comdexation. but the thermodynamics are reported to be uncertain. he "electron excitation energy" (El) is considered to be related to the energy involved in the xy = xZ - y2 electron transition. This, reportedly, is the lowest absorption in the spectra of dS systems. For the [M(2 = p h o s ) ~ ]complexes the energies are 13.5, 24.7, 19.1 kK for M = Co, Rh, 11, respectively. The respective enthalpies of activation (&Hz*) are 3.4, 11.6 and 6.5 kcal mole-l. This agreement between E l and AHz* suggests that the reactivity of the ds.. MU) . . com~lexesis directlv related to their liaand field stabilization energies (15). In the analoes IIrA(CO)(PhsP)zl (where A = F, CI, Br, I) a relationship has been noted between the free energy of TASz* (Table 3) and the lowactivation AGz* = AHz* ; est energy band in the visible absorption spectra (14). This relationship is illustrated in Figure 2, uiz., that the free energy of activation is directly related to the "electron excitation energy" (El)of the parent, ds, [IrA--d
."
672
/
Journal of Chemical Education
Excitation
Energy Et (k K)
Figure 2. Free energy of activation for the oxygenation of (IrA(CO)(PhsP)2)(Table 3) as a function of the electronic excitation energy (El) COrreSpOnding to the absorption band in the visible spectrum of (IrAICO)(Ph3P)~) (from ref. ( 1 4 ) ) .
(CO)(Ph3P)2]. The electronic spectra of these complexes show three absorption hands (427-451 nm (El), 378-398 nm (Ez), and 328-356 nm ( E d ] , assignable to d-d transitions in the d8 planar [IrA(CO)(Ph3P)z]. In Figure 2 the lowest energy transition (El, d,, - d,z-,z) is used. This increase in excitation energy parallels an increase in halopen electroneeativitv. Therefore. the observed trend mav Ybe related touthe &eater e~ectrbnwithdrawing ability i f the haloeen substitutents and an increase of oxidation state in t i e metal ion. Complexes of Cobalt
In contrast to the iridium and rhodium com~lexes which combine with oxygen in a 1:l (metal:oxygen) ratio. dimeric (2:l) complexes of colialt have been known for some time. The brown diamagnetic ion [(NC)5CoOzCo(CN)sI6- when treated with bromine in basic solution will form the red ion [(NC)5C002Co(CN)5]5- which is paramagnetic (16). The diamagnetic ion may be considered an oxygen bridged complex in which each cobalt (a d7 metal) is in the Co(II1) oxidation state, and the oxygen has assumed a peroxide structure. In the paramagnetic case, however, the oxygen must either be a snperoxide with Co(II1) or in the unlikely form of a peroxide with one cobalt atom as Co(II1) and the other as Co(IV) (first suggested by Werner, 1898) (17). The molar magnetic moment of the compound [(NH~)~COOZCO(NH~)SI(SO~)Z(HS04).3Hz0 has been measured as 1.63 BM, and paramagnetic tesonance absorption has demonstrated that the unpaired electron is delocalized over two equivalent comolecular orbital calhalt atoms (18). . . VlEek has andied .. culations and group theoretical procedures to the problem of hondine in (Co-02-COP+ complexes and has assigned oneUunpaired el&ron into a molecular orbital encompassing the entire molecule (17). He has derived the highest energy molecular orbital in the diamagnetic ion (CO-OZ-CO)~+ to be an antibonding orbital whose suhsequent oxidation yields the more stable paramagnetic ion. Bondine involves the formation of molecular orbitals with the oxygen molecule oriented perpendicular to the cobalt-cobalt internuclear axis. This is similar to the proposed orientation of Fe(I1) in Griffith's description of oxyhemoglobin (see section on Bonding). More recent investigation of the compound decaam-
through the solution rather than from any type of interaction with the solvent molecules. Reactions have also been observed in which hoth the oxygen and a different ligand form a dihridged species which may exist in either a diamagnetic or paramagnetic form. Infrared examination of compounds such as fi-amido-p-peroxo dicohalt ammines show cohalt atoms which have similar electronic distrihutions as those found in the oxygen single-bridged compounds. This is in accord with the idea that in reactions such as (26).
-
-
[(NH,)4Co(NH~)(0,)Co(NH3)I13* [(NH,),CO(NH,)(O~)CO(NH~)~]~+ + e-'
Figure 3. The structure of ( N H ~ ) ~ C O O ~ C O ( N H (from J ) P rel. ( 1 9 ) )
mine-p-peroxodicohaltmonosulfatetris(hisulfate) [(NH3)5COO~CO(NH~)~](SO~)(HSO~)~ has provided results (19) which cast doubt on VlEek's theory. The cation is reported to exist as two almost regular octahedra which are o-bonded to an oxygen through each cohalt ion. The NH3-Co hond length is 1.95 A and the 0-0 hond length is 1.31 A, only slightly greater than that of the superoxide. The resulting structure is depicted in Figure 3. The oxygen molecule is not orientated perpendicular to the cobalt--cobalt internuclear axis as predicted by VIEek, and the oxygen molecule exists in the form of a supemxide. This is in agreement with earlier predictions of 1.28 0.07 b, for the 0-0 hond length based upon molecular orbital calculations of a molecule with D4h symmetry. It is noted that the complex "should not he appreciably destabilized by departure from linearity" (20) and oxidation would involve removal of an electron from a molecular orbital which is concentrated on the oxveen molecule. Re"" duction of this compound to form the diamagnetic ion should involve formation of the ~eroxidewith an anoroximate hond length of 1.48 A. TI& has been ohse&d for the compound decaammine-pperoxo-dicohalt disulfate tetrahydrate in which the actual oxygen-oxygen hond length is 1.47 A with a torsion angle of 146' (approx. 90' in H202) due to steric factors (21). Finally, examination of the paramagnetic tetrahydrate complex in terms of valence hond theory gives good agreement for a single and "three electron hond" between the oxygen atoms (1.31 A) and for the cobalt-oxygen and cobalt-nitrogen hond lengths of 1.89 A and 1.95 A, respectively (22). Kinetic studies on decomposition and rate of oxygen complex formation of Co(lI) in ammonia solution have sho& the pentaammine to he a far more reactive species than the hexaammine (23). Decomposition reactions in aqueous ammonia solutions of varying concentration have eliminated such adducts as ((NH3)4(H2O)Co)zOz since the same species was found to he present in solutions of very high and very low ammonia concentrations. The important reactions (from ref. (23)) are
which forms a paramagnetic green ion, the electron is lost from a molecular orhital rather than a hybridized cobalt orbital. The first monomeric Co(lI) species to he crystallized and positively identified was reported in 1969 to he the pyridine oxygenation product of a methoxy ring-suhstituted form of N,N'-ethylenebis(salicylideneiminato)cohalt(ll)called Co-salen (27).
*
k,
[Co(NH,),O2I2++ [CO(NH~)~H~O]*+ CL k-i
+
[(NH:,),COO~CO(NH,),]~H20 Generally, except in very strong ammonia solutions (29 M) in which the effects of the hexaammine are uncertain, it is found that the product of kz and the molarity of [ C O ( N H ~ ) ~ ( H Z Ois ) ] ~much + greater than k-I and a second-order dependence is noted. I t is difficult to determine a reason for the rate a t which a cohalt(I1) complex will react with molecular oxygen (24). If a complex contains an electron withdrawing group then this could interfere with hack donation to the oxygen molecule, and an effect on reaction kinetics mav he noted. Molecular oxygen isotope studies (25, 26) have shown that in cohalt complexes (as has been shown in complexes of iridium and rhodium) hoth oxygen atoms in the bridge arise from the original gaseous sample which was bubbled
Co-salen (27) Co(aeacen) (28) The monomeric adduct [Co(3-methoxysalen)(py)O~]is reversible and shows a magnetic moment indicative of one unpaired electron, which may he due to the presence of a 1:2 adduct with a lower perf. If, however, the magnetic moment is correct then this would correspond to a low d7 state for cohalt with hack donation to the oxygen antibonding orbitals (see section on Bonding). The oxygen molecule would then he orientated with an intemuclear oxygen-oxygen axis 90' to the molecular plane containing the cohalt(I1) atom. Ir data support this structure hut as will he shown in the next paragraph, monomeric cohalt adducts may assume other structures as well. The compound [Co(acacen)] where acacen = ((CH3C(O-)=CHC(CH3)=NCHz-)z) has been investigated and shown to he a low spin cohalt(Il) complex with a magnetic moment of 2.16 BM (28). When 0.1 mmole of this complex is oxygenated in toluene in the presence of 4-aminopyridine (NH2-py), it is found that 0.099 mmole of molecular oxygen is complexed which indicates a cohalt to oxygen ratio of 1.02 (29). The resulting monomeric complex has a magnetic moment of 1.49 BM, so [Co(acacen)(NHz-py)Oz] still contains an unpaired electron. It has been observed that the percentage of oxygencomplex formation is proportional to the strength of the pyridine substituted hase as evidenced by a cohalt to oxygen ratio of 0.91 for CN-py @ K , = 1.86) and 1.02 for NHz-py @ K , = 9.3). The purpose of the ligand hase is nrohahlv to allow formation of an octadedral environment and provide for more efficient overlap to the oxygen molecule from the low snin cohalt comolex (29). Electron naramagnetic resonance (epr) studies identify the lone'electron with the oxygen molecule to form the now familiar superoxide 0 2 - , with the cohalt left in a d6 state (28). Three possible structures are n
Structure 111 may he eliminated since there must he a reduction in symmetry of the oxygen molecule to cause the oxygen electrons to pair in a r,2p orhital as is shown by Volume 50, Number 10, October 1973
/
673
the magnetic moment (29). Ir stretching data and epr measurements (28) which suggest a lack of axial symmetry both favor structure Il. A comparison between the structure of these monomeric cohalt complexes and the structure of oxyhemoglohin shows many similar features. The four oxygen or nitrogen ligands coordinated to the central atom in these complexes is replaced by the porphyrin structure of the heme group. The substituted pyridine base on the cobalt has its counterpart in the iron adduct which is "coordination of a histidine ~ ~ residue ~ ~ in ~ a trans ~ o s i t i o nto the coordinated oxygen" (30). Before leavine the t o ~ i cof cobalt comolexes, it is interesting to noteuwhy iridium and rhodium do not form bridged oxygen compounds yet cobalt does. The cobalt atom in such bridged compounds is in the form of Co(II1) which is a higher oxidation state than exhibited by the other elements of the same periodic group in their oxygen adducts. Consequently, there is less efficient overlap with the a-antibonding orbitals due to contraction of the metal d orbitals (31). In such a case, the sigma style of honding will occur. Complexes of Other Metals Molecular oxygen will react with tetrakis(tripheny1phosphine)M where M may be nickel, platinum, or palladium (32). The reaction with nickel to form the unstable species [(Ph3P)2NiOz] is believed to involve the [Ni(PhaP)z] complex which is present in equilibrium amounts. The reaction in ether or toluene would he (32)
The platinum and palladium complexes were reported to be formed in a similar manner from benzene (32)
+ 0,
[M(Ph3P),I
+
+
[(PhBP)IMOz] 2PhsP (in part being Ph,POj Ph3P0 = (CeH,XPO Infrared analysis of (Ph3P)zPdOz and (PhaP)iPtOz show hands at 880 cm-1 and 830 cm-1, respectively, which has been related to the oxygen-oxygen stretching vibration. The single crystal structural analysis of the platinum complex (33) has yielded the structure [PtOz(PhsP)z]-1.5 C6H6 with both phosphorus atoms, the oxygen atoms and the platinum atom almost situated in a plane. The oxygen-oxygen bond length is 1.45 A which is close to that in Vaska's iridium iodo complex. All three complexes were expected to decompose in excess Ph3P according to the equation (32)
It has been found that by either dissolving Pt(PhsP)+ in benzene or combining one mole of P(C6H5)3 and [Pt(P(CsH5)3)z(CzH4)] per liter of solution (to form Pt(Ph3P)j) equal kinetic rates are evident upon reaction with molecular oxygen to form Pt(Ph3P)zOz. Therefore the oxygenation of the platinum compounds may involve the tris(triphenylphosphine)platinium(O) species, i.e. (34)
The following equation is for decomposition in presence of excess Ph3P (34)
with an associated rate equation d(Pt(PhaP)a)/dt = k(Pt(Ph3P)zOz)(Ph3P).A proposed mechanism is (34) 674
/ Journal of Chemical Education
PPh,
PPh,
Another molecular oxygen adduct of nickel, NiOz(tBuNC)z, is formed by passing oxygen at -20°C through an ether solution of Ni(t-BuNC)a (35). Nmr measurements have shown the equivalency of all the tert-butyl hydrogens, and a low magnetic moment of 0.13 BM was attributed to temperature independent paramagnetism (36). 180isotope substitution of '602, 160-1s0, and I 8 0 2 (58.3, 32.4, and 9.3 mole 7%) produces three oxygen stretching frequencies (898, 873, and 848 cm-') which upon comparison with an isotope study on Vaska's chloro-iridium complex (856, 832, and 809 cm-') seems to indicate a similar type of oxygen-metal CzU symmetry. It should be remembered, however, that as in Vaska's compounds, these values may not represent pure oxygen-oxygen stretch due to metal-oxygen interaction. Summary I t has been shown that a variety of compounds exist which may serve as model systems for reversible oxygen carriers. Elucidation of the structural and chemical properties of these synthetic compounds has increased our understanding of the processes that occur in natural transition metal complexes which bind molecular oxygen. Hoffman and Petering (37) have prepared a paramagnetic cobalt analog of hemoglobin from cohalt(1II-porphyrin and animal glohins. Electron paramagnetic resonance (epr) studies of these cohoglobins indicate that the hinding of the oxygen molecule is independent of the coordinating protein although the presence of the protein may be related to the stability of the complex. The oxygen is considered to exist in the form of a superoxide and is thought to have a geometry analogous to that described for [Co(acacen)(NHz-py)Oz]. This would he consistent with the bent geometry suggested by Pauling in 1949, hut inconsistent with the mode of honding demonstrated in Griffith's model (see section on Bonding). A comparison of the two systems, Co(II1)-02- and Fe(II1)-02-, suggests that this may also be the structure present in oxyhemoglohin. Further evidence for the Pauling style of honding in these complexes has been obtained from X-ray studies on a recently synthesized compound [Co(bzacen)(py)Oz] where hzacen is N,N'-ethylenebis(benzoylacetoniminide), (C~H~-C(O-)=C~-C(CH~)=NCHZ-)~ (38). The Co0-0 bond angle is 126", and the bond length is consistent with that of the superoxide species. As more is learned about the functioning of these synthetic compounds, a groundwork is laid for an understanding of the complex processes which occur in hiological systems. The geometry of the oxygen-metal bond proposed by Griffith has been applied to complexes of iridium and rhodium, but does not apply in the case of hemoglobin. The synthesis and investigation of new oxygen adducts will lead to a better understanding of the manner in which honding does occur and will also clarify the mechanism of reversible oxygen uptake. The study of molecular oxygen adducts is a rapidly growing field of transition metal chemistry in which much has been learned in a relatively short time. However, the investigation is far from complete, and a t this point the experimental results have not been organized into a unified approach to the problems under investigation.
Literature Cited Ill Grimth, J. s.. R o c . Roy Soe.. ~236.i3119591. 121 Vaska. L., Science. 140.809 11963). 131 MeGinnety, J.A.,Doedens, R.J., andlbera, J.A.,Inorg. Cham.. 6,2243119671, 141 McGinnety, J.A.,Dcedens, R. J., sndlbers, J.A.,Scienee, m , 7 0 9 ( ~ % 6 7 l . 151 MeGinnety, J . A,, Psyne, N. C., and Ibers, 3. A . J Amor. Chem. Soc., 91, 6301
16) Nakamura, A,. Tawuno. Y., Yamsmoto. M:. and Otsuka, S., J. Amer. Chsm See. 505211971), L..Arcounfs C h m . R e 8. I .,~ 336 I l f f R i . I W ~ a ~ i s eS.s .J.. a n d ~ b e r s , ~ . ~ . . ~them . ~ m esot. r 81.2~11~51. (11) Martell, A. E.. and Calvin, M.. "Chemistry of the Metel Chelate Compounds:' Prentiee Hall, Inc., Englewwd Cliffs, New Jersey. 1912, p. 352. I121 Iben. J.A.. and LaPlacs. S. J..Srimce. 145.920119MI. . . . 113) Ch0ck.P. B., andHalpern, J . , J . A m r r ~ h m Soe.. t 88,3511 11966). 1141 Va8ke.L.. Chcn. L. 8.. sndSenoff, C. V.,Science. 174.58711971). I151 Vasks, L.,Chen, L. S., sndMiller, W. V.. J. Amer Chem S o c . 93,6671 (1971). 116) Basoio. F.. and Pearsan, R. G., "Meehsnism of Inorganic Reactions, A study of Metal Complexes irr Solution" 2nd ed interscience Publirhers. Now York, 1967. p. 662-3.
.I'
\.",",. I291 Crumb1ini.A. L.. andBaaalo. F., J. Amer Cham. Sac.. 92.55 11970). 1301 Crumbliss, A. L., and Basdo. F., Seienre. 164,1168 11969). 1311 Mason. R.. Naturr.. 2l7.54311968) . .~~ , 1321 Wilko, G.. Scott. H.. and Heimbach. P.. Angew. Chem., hl.Ed. E n p i . 6, 92
,."*",
