J. Phys. Chem. 1995, 99, 7172-7179
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Spin-Lattice Relaxation Enhancement of Water Protons by Ferric Porphyrins Complexed with Cyclodextrins and Fluoride Ions Sandip K. Sur and Robert G. Bryant* Chemistry Department, University of Virginia, McCormick Road, Charlottesville, Virginia 22901 Received: December I , 1994; In Final Form: February 20, 1995@
The addition of a-, /3-, and y-cyclodextrins to ferric tetrakis(4-sulfonatopheny1)porphine complex, [Fe(III)TPPS4], is characterized by optical and magnetic resonance spectroscopy. The formation of the cyclodextrin complex with the metalloporphyrin inhibits dimerization reactions and increases the electron spin relaxation time of the iron(II1) center. The water proton nuclear magnetic spin-lattice relaxation rates are reported for the iron(II1) porphyrins as a function of the magnetic field strength. The relaxivities for the (cyc1odextrin)[Fe(III)TPPS4] complexes are higher than those of the cyclodextrin-free [Fe(III)TPPS4] at pH 2 and 7. The observed relaxivities for the (cyc1odextrin)-[Fe(III)TPPS4] complexes at pH 7 are lower than the values at low pH, presumably because the formation of hydroxo complexes reduces the number of first coordination sphere protons. The electron spin resonance spectra of the (cyc1odextrin)- [ ( F ~ F z T P P S ~ )complex ~-] were measured at 77 K and at room temperature. The spectrum at 77 K displays a superhyperfine coupling to fluorine of 4.17 mT in the g 2 region. The difluorocomplex produces a water proton relaxation rate of 14.73 rnh4-l s-l at a magnetic field strength corresponding to 0.01 MHz 'H Larmor frequency. The shapes of the magnetic relaxation dispersion profiles are not Lorentzian at pH 7 and are not well described by the standard theories for paramagnetic relaxation; however, these results demonstrate that it is possible to construct iron(II1) based relaxation agents that produce efficient relaxation of the water protons.
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Nuclear magnetic relaxation of water protons is central to control of image contrast and information content in magnetic resonance imaging.' In a previous paper, we reported the high water proton relaxivities for cyclodextrin-bound paramagnetic manganese@) and manganese@) complexes.2 Both Fe(III) and Mn(I1) are d5 ions and in the high spin configuration may have long electron spin relaxation times, which are required for efficient water proton spin relaxation. However, the nuclear relaxation induced by iron(II1) is often much less efficient than that induced by Mn(II).3-8 The water proton relaxivities of iron(111) porphyrins are usually strongly pH dependent9-" and are large only in strongly acidic solutions.* However, we report here a method for dramatically increasing the iron(III) relaxation efficiency by prohibiting dimerization and the consequent deleterious magnetic coupling between iron atoms. Porphyrin complexes have been used in recent years for magnetic resonance imaging because they localize in tumors and provide sharp anatomical images of the tumor among the background tissues.12-18 Ferric tetrakis(4-sulfonatopheny1)porphine complexes, [Fe(III)TPPS4], are soluble in water, stable in air,and exist in several ionization states over a wide pH range in solution. These complexes are not usually considered as effective in vivo contrast agents because, at pH values near 7, they form p-oxo bridged dimer, [(FeTPPS4)20]8-,9 The antiferromagnetic coupling of the iron centers in [(FeTPPS4)20I8reduces the effective magnetic moment and decreases the electron relaxation times; therefore, iron(II1) porphyrins are inefficient as contrast agents in this configuration. At near neutral pH, aside from the predominant p-oxo bridged dimer, the formation of [FeTPPS4(OH)(H20)]4- and [FeTPPS4(0H)2I5species has been suggested based on proton NMR and optical However, at low pH, the diaquo species, [FeTPPS4(H20)2I3-, is formed, and efficient water proton spin relaxation is observed.8
* To whom correspondence should be addressed. @
Abstract published in Advance ACS Abstracts, April 15, 1995.
0022-3654/95/2099-7172$09.00/0
The cyclodextrins (CD), also called Schardinger dextrins, are a class of cyclic oligosaccharides and contain 6 to 11 a-1,4linked D-glucose molecules. a - , b-, and y-cyclodextrins are produced from enzymatic degradation of starch and correspond to cyclohexa-, cyclohepta-, and cyclooctaamyloses. They have molecular weights of 972.9, 1135.0, and 1297.1 and a barrel structure with internal diameters of 4.5, 7.0, and 8.5 A, respectively. These six-, seven-, and eight-membered species form stable host-guest clathrates with several compound^'^-*^ that have drawn widespread attention in recent years because they serve to mimic biological systems and may be used to deliver drugs.26927 The formation of a complex between cyclodextrins and the porphyrin should increase the rotational correlation time for the metal center and decrease the water translational diffusion coefficient in the vicinity of the complex, both of which should increase the magnetic relaxation efficiency of the complex. Further, the steric constraints imposed by the cyclodextrin inhibit dimer formation and, thus, antiferromagnetic electron spinelectron spin coupling.
Experimental Section The iron(II1) tetrakis(4-sulfonatopheny1)porphine was purchased from Porphyrin Products, Inc. (Logan, UT); the cyclodextrins from Aldrich Chemical Co.; potassium fluoride from Kodak Chemical Co.; and deuterium oxide from Cambridge Isotope Laboratory. Stock solutions were prepared in 0.10 M phosphate or 50 mM HEPES buffer. Water used in all experiments was routinely taken from a Barnstead Millipore Filtration System with both ionic and organic sections that used house deionized water as the feed. The pH was measured on an Orion Model 720A or a Coming Model 240 pH meter. All electron paramagnetic resonance (EPR) measurements were carried out at 77 K or at ambient laboratory temperature using a Bruker ESP-300 spectrometer at X-band with 100 kHz magnetic field modulation with modulation amplitude of 0.1 to 0 1995 American Chemical Society
Spin-Lattice Relaxation of Ferric Porphyrins
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J. Phys. Chem., Vol. 99, No. 18, 1995 7173
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Figure 2. The 500 MHz proton NMR spectra of 10.0 mM j3-cyclodextrins in the presence of (a) 0.0, (b) 2.8, and (c) 5.0 mM [Fe(III)TF'PS4] at 298 K.
Figure 1. The UV-visible absorption spectra from 8.60 x [Fe(III)TPPS4] in the presence and absence of 0.99 x
M M j3-cyclodextrin in 0.20 M aqueous sodium perchlorate solution at pH 2.0 and room temperature, showing a decrease in the 394 nm band from cyclodextrin binding.
jJ( Q
0.5 mT and microwave power level of 20 mW. 2,2'-Diphenylpicrylhydrazyl (DPPH) was used to calibrate the magnetic field assuming a g-factor of 2.0036.28 The nuclear magnetic relaxation rates were measured over a range of magnetic field strengths corresponding to proton Larmor frequencies between 0.01 and 30 MHz using a fieldcycling spectrometer described e l s e ~ h e r e .The ~ ~ W-visible absorption spectra were obtained on a Varian Cary-4 spectrophotometer. The 'H and 13CNMR measurements were made on a Varian Unity-Plus 500 NMR spectrometer operating at a proton Larmor frequency of 500 MHz. The 2D-ROESY experiments on the cyclodextrin- [Zn(II)TPPS4] complex were carried out by the small-angle pulse method30 at 499.88 MHz with an effective spin-lock field of 5.2 kHz during the mixing period of 200 ms; 32 scans were acquired for each t l value, and the 2048 x 512 data set was converted to a 2048 x 2048 data set by zero-filling.
Results and Discussion All three cyclodextrins form host-guest complexes with the metal porphyrin, [Fe(III)TPPS4], as demonstrated by the observation of optical absorption spectra from the porphyrins and the proton or carbon NMR spectra of the complexes. The effects are most pronounced for the P-cyclodextrin, which is consistent with the relative order of the binding constants.2 W-Visible Absorption Spectra. The W-visible absorption spectra of 8.6 x M [Fe(III)TPPS4] at pH 1.5 in the absence and in the presence of 1 mM p-CD in 0.2 M aqueous sodium perchlorate solution are shown in Figure 1. No new absorption bands are observed in these spectra when cyclodextrins are added, which is consistent with the absence of aggregation in acidic ferric porphyrin solutions. At low pH, the ferric porphyrin exists primarily as the monomeric [FeTPPS4(H20)2I3- species, but addition of 1 mM P-CD causes the Soret band intensity to decrease, implying complex fonnation between ferric porphyrin and cyclodextrin.lO.ll Further addition of p-CD up to 6 mM did not appreciably change the spectrum. The absorbance data of the Soret band of Figure 1 were analyzed assuming two independent and equivalent binding sites for p-cyclodextrin and yielded a microscopic binding constant of 315 L M-' for each. Similar spectra were obtained from both a-CD and p C D , although the spectral changes are smaller. These spectral data demonstrate a significant binding interaction between the cyclodextrins and the [Fe(III)TPPS4] complex.
55
45
35
25
15 ppm
Figure 3. The 500 MHz proton NMR spectra of [Fe(III)TPPS4] at 298 K: (a) 5.0 mM, pH 12.0; (b) 5.0 mM, pH 2.0; (c) 2.8 mM with 10.0 mM j3-cyclodextrin, pH 2.0; (d) 5.0 mM with 10.0 mM j3-cyclodextrin, pH 2.0; (e) 5.0 mM with 10.0 mM P-cyclodextrin, pH 7.0; (0 5.0 mM with 10.0 mM j3-cyclodextrin, pH 12.0; and (g) 5.0 mM with excess j3-cyclodextrin, pH 12.0. The broad B-pyrrole peak at 81 ppm is shown separately in spectrum f.
Although Mosseri et al. claim 1:4 complexes for both [Zn(II)TPPS41 and [Fe(III)TPPS4] with ,8-CD,l0 molecular modeling studies suggest that only two cyclodextrins can bind simultaneously to one porphyrin molecule, which was concluded by Manka and Lawrence and which is consistent with the present data.31 N M R Spectroscopy. The proton Nh4R spectra of solutions of ,8-CD and [Fe(III)TPPS4] in D20 at ambient laboratory temperature are shown in Figures 2 and 3. The spectra of 10 mM p-CD solutions broaden with increasing [Fe(III)TPPS4] concentration as shown in Figure 2. Scheidt et al? have reported that three peaks at 5 1.4 @-pyrrole protons), 14.3 (orthoprotons), and 10.3 (meta protons) ppm in the proton NMR spectra of 5 mM [Fe(III)TPPS4] at pH 2 (Figure 3, spectrum b) are characteristic of the monomeric [(FeTPPS4(H20)zI3- species.
7174 J. Phys. Chem., Vol. 99, No. 18, 1995
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Sur and Bryant
h
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240 200 160 120 80 40 0 PPm Figure 5. The 125.7 MHz I3CNMR spectra of 5.0 mM [Fe(III)TPPSd] at pH 2.0 and 298 K. The peaks marked with a * are from internal acetone.
110 100 90 80 70 60 PPm Figure 4. The 125.7MHz 13CNMR spectra of 10.0 mM P-cyclodextrin with added [Fe(III)TPPSd] at 298 K: (a) 0.0 mM [Fe(III)TPPSd], pH 7.0; (b) 2.8 mM [Fe(III)TPPS4],pH 2.0; (c) 5.0 mM [Fe(III)TPPS4], pH 2.0; (d) 5.0 mM [Fe(III)TPPSh],pH 7.0; and (e) 5.0 mM [Fe(III)TPPS41, pH 12.0
At neutral and higher pH values (Figure 3, spectrum a), the p-oxo bridged dimer, [(FeTPPS4)20Is-, with corresponding peaks at 13.8, 8.3 or 8.2, and 7.7 ppm is primarily formed. The proton NMR spectra, however, do not provide a distinction between the dimers containing axial water or axial hydroxyl ligands. When cyclodextrins are added at neutral pH, the dimers are broken to form monomers. Mosseri et ~ 1 . ' ~ reported the formation of (cyclodextrin)-[FeTPPS4(OH)n (H20)l4- complex at neutral pH and (cyclodextrin)-[FeTPPS4(OH)$ species at higher pH. At pH 12 (Figure 3, spectra f and g), the proton NMR spectrum of 5 mM [Fe(III)TPPS4] with p-CD consists of both dimeric and monomeric (cyc1odextrin)[FeTPPS4(0H)2I5- contributions. As shown in spectrum g, excess p-CD does not completely suppress the formation of dimers at pH 12 because of the limited solubility of the P C D . (a) -.-----.---.-In fact, in the 5 mM femc porphyrin solutions used here, both 270 250 230 210 190 170 150 130ppm dimers and monomers are detected by proton Nh4R at all pH Figure 6. The 125.7 MHz I3C NMR spectra of [Fe(III)TPP&] at 298 values, although the concentration of the monomers at high pH K: (a) 5.0 mM [Fe(III)TPPS4], pH 2.0; (b) 2.8 mM [Fe(III)TPPSd], and concentration of the dimers at low pH are both small in 10.0 mM P-cyclodextrin, pH 2.0; (c) 5.0 mM [Fe(III)TPPS4],10.0 mM the absence of cyclodextrin. P-cyclodextrin, pH 2.0; (d) 5.0 mM [Fe(III)TPPSd], 10.0 mM P-cyFormation of the cyclodextrin complex with [FeTPclodextrin, pH 7.0; (e) 5.0 mM [Fe(III)TPPSd], 10.0 mM P-cyclodextrin, PS4(H20)2I3- at pH 2 is accompanied by the broadening and pH 12.0; and (05.0 mM [Fe(III)TPPS4],pH 12.0. shift in the p-pyrrole and ortho protons. However, cyclodextrin complex formation with [FeTPPS4(H20)(0H)l4- or [FeTP12, all carbon peaks are broadened significantly, which is consistent with more efficient nuclear spin relaxation that results PS4(0H)215- at pH 7 or 12 dramatically broadens the proton NMR spectra, which is consistent with the longer electron from a longer electron relaxation time of the metal center. relaxation time of the monomers and is supported by the EPR Figure 5 shows the carbon NMR spectrum of the free femc spectra described later. Similar line broadening was also porphyrin at pH 2, which is 300 ppm wide with observable observed in the carbon NMR spectra. The P-pyrrole proton resonances at 181.0, 144.5, 129.0, 127.9, 124.9, 124.1, and -1.3 shifts from 56.2 to 81 ppm as the pH of the 5 mM [Fe(III)The I3C NMR spectra of the [Fe(III)TPPS4] complex TPPS41 solution containing 10 mM p-CD is raised from 7 to are shown as a function of pH and cyclodextrin concentration 12. The corresponding meta-proton peak of the monomers shifts in Figure 6. In the absence of cyclodextrins, the carbon NMR only from 10.3 to 10.9 ppm between pH 2 and 12. spectra at pH 2 (Figure 6a) and pH 12 (Figure 6f) are from the monomeric and p-oxo dimeric complexes, respectively. When The cyclodextrin resonances in the I3C NMR spectrum of the p-~yclodextrin-[Fe(III)TPPS4]complex are shown in Figure cyclodextrin is added at pH 2 (Figure 6, b and c), the carbon 4. The carbon resonances of the free p-cyclodextrin are peaks of the porphyrin broaden, which is consistent with broadened and shifted when [Fe(III)TPPS4] is present in the complex formation. At pH 7 (Figure 6d), the carbon NMR solution at pH 2 (see Figure 4, b and c). The carbon resonances spectrum is quite broad, although the presence of the dimer peak in this solution with p-CD to [Fe(III)TPPS4]molar ratio of 2: 1 at 144.0 ppm is seen in spectrum f of Figure 6. The concentration of the dimeric species is still higher at pH 12 appear at 103.6, 82.8, 77.9, 73.8, 72.6, and 62.0 ppm. The prominent feature of the binding is a shift of the C-3 resonance, (Figure 6e), which is consistent with the proton NMR observations in the ferric TPPS4 at pH 12. When cyclodextrin binds which becomes deshielded by about 3.7 ppm at the same time it is considerably broadened (Avll2 = 246 Hz). At pH 7 and to opposite phenyl groups of the TPPS4 complex, the phenyl
Ik
!
J+-' l-
! liL
J. Phys. Chem., Val. 99, No. 18, 1995 7175
Spin-Lattice Relaxation of Ferric Porphyrins
,
160
150
140
130
'
,
I
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120 ppm
Figure 7. The 125.7 MHz I3C NMR spectra of 10.0 mM [Zn(II)TPP&] at pH 8.0 and 298 K without (a) and with (b) 10.0 mM
P-cyclodextrin. groups are enclosed within the cyclodextrin cavity, and the protons and carbons at the P-pyrrole and ortho positions are affected more than those at the meta position. This observation is consistent with the larger broadening observed in the proton and carbon NMR resonances of the P-pyrrole and ortho position compared to the meta position. The carbon NMR lines of zinc@) porphyrin are also broadened in the presence of cyclodextrin. The carbon NMR spectra of the 10 mM [Zn(II)TPPS4] in the presence and absence of equimolar P-cyclodextrin are shown in Figure 7. The carbon peaks of the free [Zn(II)TPPS4] appear at 150.3 (a-pyrrole), 142.6 (para carbon attached to meso carbon), 146.6 (para carbon with SO3-), 135.8 (ortho carbon), 132.8 @-pyrrole), 124.6 (meta carbon), and 120.3 ppm (meso carbon).32 Our assignments of carbon peaks of the [Zn(II)TPPS4] differ from those reported by Goff and Morgan.32 By using the NMR line broadening observed on the carbon NMR spectra of the [Zn(Il)TPPS4] by P-CD,. we have reassigned the carbon peaks at 142.6 and 146.6 ppm to the para carbon attached to the meso carbon and to the para carbon attached to the SO3- group, respectively (see Figure 7). Intermolecular binding interactions are directly demonstrated by observation of intermolecular nuclear Overhauser effects between the binding partners. The rotating frame NOE (ROESY) spectra of the 2 mM [Zn(II)TPPS4] and 8 mM p-CD solution yielded intermolecular cross-peaks between the H-3 and H-5 protons of the cyclodextrin and the P-pyrrole and phenyl protons on the porphyrin demonstrating the character of the complex. Electron Spin Resonance. The EPR spectrum of [Fe(III)TPPS41 taken at X-band is shown in Figure 8b for a frozen aqueous solution at pH 2 without added cyclodextrin. The modest line width is consistent with a relatively long electron spin relaxation time for the axial spectrum at 77 K. The EPR spectra show two signals at g l = 5.9 and gll = 2.0 without any hyperfine splittings, although the signal at gll= 2.0 is probably broadened and is not usually detected at pH 2.2,33 The EPR spectra with added a-, P-, and y-cyclodextrins are similar to those without added cyclodextrin except that the EPR line at g l = 5.9 is narrower when cyclodextrin is present as shown in Figures 8 and 9. The EPR line widths (AH,) of [Fe(III)TPPS4] complexed with P-cyclodextrin measured as the distance between the extrema of the derivative mode spectrum at pH 2 and 7 are 16.2 and 7.4 mT, respectively, while the observed line width is 27.3 mT for the cyclodextrin-free ferric porphyrin at pH 2. The EPR line at g l = 5.9 is narrower at pH 11 (Figure 8d), and the line width in this case is 6.8 mT. Unlike the EPR spectra at pH 2, the signals at g l = 2.0 are clearly observable at pH 7 and 11. The EPR spectrum measured without added cyclodextrin at pH 7 (Figure 9c) is consistent with the dimer
formation at higher pH, although a very small amount of monomer EPR signal at g l = 5.9 is seen. Dramatic changes in the intensity of the EPR spectra and splittings are observed when potassium fluoride is added to the solution containing both [Fe(III)TPPS4] and P-CD, as shown in Figure 9, trace e. The fluorine superhyperfine splitting of 4.17 mT in the gl = 2.0 signal clearly identifies the species as (~yclodextrin)-[FeF2TPPS4]~-complex with two equivalent axial fluorine ligands. In the absence of cyclodextrin, however, the EPR spectrum is practically undetected as shown in Figure 9b, implying that cyclodextrin is needed to disrupt dimer formation. The EPR line at g l = 5.9 with AHppof only 4.2 mT of the fluoro complex measured at pH 7 and 77 K is even narrower than that from the (cyc1odextrin)- [FeTPPS4(0H)2I5complex. The decrease in EPR line width observed at 77 K with added cyclodextrin is consistent with the observation of the EPR spectrum at approximately 296 K as shown in Figure 10. Both (cyclodextrin)-[FeF2TPPS4I5- and (cyc1odextrin)[Fe(III)TPPS4] complexes at pH 7 have observable EPR spectra at room temperature, although no EPR spectrum is observed at room temperature from [Fe(III)TPPS4] solution at pH 2 (Figure loa). The observed room temperature EPR absorption at gl = 5.9 is anisotropic and the narrow line widths of (cyc1odextrin)[FeF2TPPS4I5- and (cyclodextrin)-[Fe(III)TPPS4] complexes indicate an increase in the electron relaxation time (T2e)in these complexes compared to that of the cyclodextrin-free [Fe(III)TPPS41 complex. These spectra suggest that the zero-field splitting is sufficiently large that the rotational motion of the complex is not rapid enough to average the anisotropy, and a powder pattern type spectrum results. Thus, the slow motion models for the spin-lattice relaxation induced in the water protons are more nearly appropriate than the more standard Solomon, Bloembergen, and Morgan (SBM) equation^.^.^^-^^ We note that the EPR lines of the (cyclodextrin)-[FeF2TPPS4I5- complex are quite narrow compared with the [FeTPPS4(H2O)2l3-, suggesting a longer electron relaxation time than that found in the diaquo complex. Both FeF63- and [FeF2TPPS4I5- have narrow EPR lines compared to complexes with coordinated water molecule^.^^^^* A possible explanation is that in the aquo case a distribution of the metal-proton distances and electronic structure may arise because of a distribution in the bonding orientation of the coordinated water molecule. A distribution of the angle that the plane of the water molecule makes with a symmetry axis of the metal center may then cause a distribution in EPR transitions at both the gl = 5.9 and gll = 2.0 regions of the EPR spectrum. In the cyclodextrin-bound porphyrin complexes, the narrower EPR spectrum may result from a narrower distribution of the water molecule orientations that is caused by steric constraints imposed by the cyclodextrins. In fluoro complexes, however, there is no change in metal-fluorine geometry, and consequently line broadening from this mechanism is not possible. The narrow EPR line width observed at pH 12 from the (cyc1odextrin)[Fe(0H)2TPPS4l5- complex is also consistent with this explanation because, in this case also, the number of possible orientations of the metal-proton vector in the hydroxyl ligand is expected to be small. Altematively, the difference between the EPR line widths of the aquo and the fluoro complexes may result from differences in the electron relaxation rates in these complexes. The sizes of the fluoro and aquo complexes are very similar so that rotational motion, if important at all, will be essentially the same for both types of complex. If the different water molecule orientations discussed above were labile, then these librational motions may modulate the zero-field splitting of the iron(II1)
7176 J. Phys. Chem., Vol. 99, No. 18, 1995
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0.1 0.2 0.3 0.4 Magnetic Field (T) +
I
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Figure 8. The X-band electron paramagnetic resonance spectra of 2.0 mM [Fe(III)TPPS4]at 77 K: (a) solid DPPH, (b) [Fe(III)TPPS4]at pH 2.0, (c) [Fe(III)TPPS4}at pH 2.0 with 40.0 mM P-cyclodextrin, and (d) [Fe(III)TPPS4]at pH 11.O with 40.0 mM /3-cyclodextrin. The peaks marked with a * at g 2 are from an impurity in the cavity.
-
I t
t
0.0
V
0.1 0.2 0.3 0.4 C Magnetic Field (T) +
Figure 9. The X-band electron paramagnetic resonance spectra of [Fe(III)TPPS4] at pH 7.0 and 77 K: (a) empty cavity with no sample, (b) 2.0 mM [Fe(III)TPPS4]with 50.0 mM potassium fluoride, (c) 2.0 mM [Fe(III)TPPS4], (d) 10.0 mM [Fe(III)TPPS4]with 40.0 mM B-cyclodextrin, and (e) 10.0 mM [Fe(III)TPPS4]with 40.0 mM /3-cyclodextrin and 100 mM potassium fluoride.
center and, consequently, produce a significant broadening of the EPR spectrum for the aquo complexes compared with the fluoro complexes in which the ligand has no dipole moment. The importance of water molecule librational motion as a factor determining electron spin relaxation time has been suggested earlier for simpler iron(II1) c ~ m p l e x e s . ~ ~ . ~ ~ Nuclear Magnetic Relaxation Dispersion. The longer electron relaxation times of the cyclodextrin-bound iron(II1)porphyrins are consistent with the observed enhancement in the water proton relaxivities from these complexes. The water
proton relaxivities for aqueous solutions of iron(II1) porphyrins are shown as a function of the proton Larmor frequency in Figures 11, 12, and 13, obtained with 20-fold molar excess of the cyclodextrins. The water proton relaxivity data obtained without cyclodextrin addition agree with those reported earlier by Koenig et a1.,8but included in Figure 11 for comparison. At pH 2 (Figure 1l), the water proton relaxivities of the solution with cyclodextrin-free [Fe(III)TPPS4] increase when cyclodextrin is added. Addition of more than 20-fold molar excess of P-CD caused no additional increase in the water proton relaxation rate. These higher proton relaxation rates at low magnetic field strengths are consistent with an increase in the electron relaxation time on cyclodextrin binding. We29 and others8J3have attempted quantitative analysis of Mn(II), Mn(111), and Fe(II1) porphyrin relaxation dispersion profiles based on the SBM equations, but these analyses are approximations at best. For example, Koenig and collaborators* reported an analysis that included a rotational correlation time of 78 ps for the [Fe(III)TPPS4] complex, but our measurements on the zinc analogue demonstrate that a rotational correlation time 3.5 times larger than that is more appropriate. Hernandez and Bryant29 reported an analysis of relaxation dispersion data on [Mn(II)TPPS41 that required an unusually large outer-sphere contribution, or delocalization of the electron. While these exercises may appear to be useful, they may also be misleading. Nevertheless, the qualitative features of the SBM analysis are usually preserved, and it is fruitful to address the major factors that contribute to the proton relaxation dispersion profile. Measurements of the 13C relaxation times for the [Zn(II)TPPS41 complex2 yield a rotational correlation time for the cyclodextrin free solutions of 275 ps, which casts doubt on an earlier analysis that assumed a rotational correlation time of 78 P S . ~ Kowalewski and Widmalm have measured rotational correlation times of 1.6 and 2.4 ns for the a- and y-cyclodextrins in porphyrin free solutions,39 suggesting that the rotational correlation time for P-cyclodextrin of about 2 ns is reasonable. The rotational correlation times for the cyclodextrin-bound [Fe(III)TPPS4] complexes studied here must, therefore, be larger than 2 ns. Were the effective correlation time for the electronnuclear coupling as long as several nanoseconds, the relaxation efficiency of the complex would be much higher than observed. Therefore, the effective correlation time is limited either by the
J. Phys. Chem., Vol. 99, No. 18, 1995 7177
Spin-Lattice Relaxation of Ferric Porphyrins
c
7
I
7
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6
Y
2
E;
H
c 0
5
0.00 0.05 0.10 0.15 0.20 0.25
Magnetic Field (T) + Figure 10. The X-band electron paramagnetic resonance spectra of 10.0 mM [Fe(III)TPPS4]at room temperature: (a) at pH 2.0, (b) at pH 7.0 with 44.0 mM j?-cyclodextrin, and (c) at pH 7.0 with 44.0 mM j?-cyclodextrin and 216 mM potassium fluoride.
I
0.01 0.10 1.00 10.00 100.00 Proton Larmor Frequency, MHz Figure 12. The water proton spin-lattice relaxation rate of 0.5 mM [Fe(III)TF'PS4]in the presence of 20-fold molar excess of j?-cyclodextrin at pH 2.0 plotted against the magnetic field strength expressed as the proton Larmor frequency at various temperatures: (0)308 K, (A) 298 K, (0)288 K.
-in 7
5
'"..I.
E
c
0
e a Y
'"1
5il
0
10 i
,
,
,
(
0.10 1.00 10.00 100.00 Proton Larmor Frequency, MHz
Figure 11. The water proton spin-lattice relaxation rate of [Fe(III)TPPS41 in the presence of 20-fold molar excess of cyclodextrins at pH 2.0 plotted against the magnetic field strength expressed as the proton Larmor frequency at 298 K: (+) a-cyclodextrin, (A)j?-cyclodextrin, (v)y-cyclodextrin. The water proton spin-lattice relaxation rate of [Fe(III)TPPSd]in the cyclodextrin-free (0)solution at pH 2.0 is also shown for comparison.
water proton exchange rate with the iron atom or by the electron spin relaxation rate. Measurements of the water proton relaxation rates at different temperatures have been reported8 for aqueous cyclodextrin-free [Fe(III)TPPSd] solutions at pH 2, and these suggest that the water proton relaxivity is not limited by exchange rate of water molecules or protons with the metal center. The proton relaxivity is reported to decrease with increasing temperature; if the exchange rate limited the relaxation efficiency, the proton relaxation rate would increase with increasing temperature, which is not observed. These observations are consistent with the 1 7 0 NMR measurements that have demonstrated an oxygen exchange rate of 1.4 x lo7 s-l at 298 K.40 Similar data are shown in Figure 12 for [Fe(III)TPPS4] in 20 mM B-cyclodextrin solution. Although at 288 and 298 K the relaxivities are nearly the same, the proton relaxivity decreases at 308 K, which is inconsistent with an exchange rate limitation. We conclude that the system is at or very near the rapid exchange limit for mixing the relaxation rates in the iron(II1) first coordination sphere and the bulk solution.
I
0
go I
0.01
I
I
0.01 0.10 1.00 10.00 100.00 Proton Larmor Frequency, MHz Figure 13. The water proton spin-lattice relaxation rate of [Fe(III)TPPS41 in the presence of 20-fold molar excess of j?-cyclodextrin at pH 7.0 plotted against the magnetic field strength expressed as the proton Larmor frequency at 298 K: (0)no added potassium fluoride, (H) with 50.0 mM potassium fluoride.
The data of Figure 12 show that, at low temperatures, the water proton relaxivity is a weak function of temperature. This result is consistent with the electron spin relaxation rate dominating the effective correlation time. This conclusion is supported by two additional observations. (1) The EPR spectrum is not detectable at room temperature in the aquo complexes, which requires that the electron spin T2 be less than approximately 100 ps. (2) The rotational correlation time for the complex when cyclodextrin binds exceeds several nanoseconds, and, if this slow rotational motion dominated the correlation time for nuclear relaxation, the 'H relaxation rate would be much larger than the observed values. The electron spin relaxation is complex. In some cases the electron relaxation rate is dependent on the temperature and viscosity, as expected if rotational or collisional processes contribute to the relaxation event.38 In other cases, the relaxation rate is independent of the t e m p e r a t ~ r e ,and ~ ~ ,for ~ ~molecules like [Gd(DTPA)]- the electron relaxation rate becomes independent of the solution dynamics as measured by the viscosity once the relative viscosity has increased sufficiently!2 The data of Figure 12 are similar to those reported by Koenig et al.' for the [Fe(III)TPPS4] in that, at lower temperatures, a temperature
7178 J. Phys. Chem., Vol. 99, No. 18, 1995 increase has a small effect and, at higher temperatures, the 'H relaxivity decreases. The factors that control the electron spin relaxation rate are not fully understood for complex systems like these, but we may make several useful observations. In some cases where the electron relaxation rate dominates the correlation time for the electron-nuclear coupling, the magnetic field dependence of the electron spin relaxation time may lead to a substantial increase in the electron spin relaxation time at higher fields that may cause an increase in the 'H relaxation rate at higher fields. The [Mn(III)TPPS4] complex is a good e ~ a m p l e , ~but , ~ this . ~ ~response is not observed at any temperature in Figure 12 over the field strengths sampled. This observation suggests that there is no dispersion in the electron spin relaxation time over this field range, which requires that the correlation time, zv, for the process controlling the electron spin relaxation is short compared to l/w, at the highest field sampled, i.e., short compared to 8 ps. The electron spin relaxation time is proportional to the product of zv and the square of the effective zero-field splitting. A decrease in either of these with increasing temperature will cause an increase in the electron relaxation time that, in tum, will cause an increase in the 'H relaxation rate, which is not observed. Since we generally expect the correlation time, zv, to decrease with increasing temperature, changes in tv apparently do not dominate the observed temperature effects in Figure 12. However, an increase in the effective size of the zero-field splitting with increasing temperature would decrease the electron spin relaxation time and, therefore, the 'H relaxivity, as observed. An additional factor that may affect the temperature dependence of the 'H relaxivity and the EPR line width is an equilibrium between high- and low-spin forms of the iron(II1) complex that shifts toward low spin with increasing temperature. In metmyoglobin and methemoglobin, the EPR spectra of single crystals have been interpreted to suggest that the aquo complex exists with a small admixture of the low-spin form that interconverts to the high-spin form at a rate of the order of lo9 s-1?3-45 The porphyrin and the local environment are different in the present case from these protein cases, but such an interconversion may affect the electron spin relaxation rate, broaden the EPR line, and reduce the 'H relaxivity. The EPR spectrum b of Figure 8 shows a predominance of the high-spin iron(II1) complex at 77 K, but studies on metmyoglobin suggest that the low-spin fraction increases with increasing temperature. Were similar effects important in the complexes studied here, the electron spin relaxation time and 'H relaxivity would decrease with increasing temperature. Although the iron(1II) EPR spectrum of the difluoro cyclodextrin complex is readily observable as shown in Figure 10, trace c, the relaxation times are small enough to obscure the fluorine coupling. Our proton and carbon NMR experiments at low pH show the existence of a very small amount of the p o x 0 dimer in addition to the highspin monomer. However, there is no trace of the low-spin monomer in our NMR spectra which should be easier to detect by NMR, suggesting the existence of monomer-dimer equilibrium even at low pH. Thus, although we have no firm evidence that an electron spin state equilibrium dominates these observations, we may not rule out the possibility. The water proton relaxivity for the [Fe(III)TPPS4] at pH 7 in 10 mM P-cyclodextrin solution is shown in Figure 13. Unlike the cyclodextrin-free solution, which has very low relaxivity, this essentially monomeric complex relaxes the water protons well. Nevertheless, the relaxivity is lower than it is at pH 2 by almost a factor of 2. This decrease is expected if the coordinated axial water molecules dissociate to form the hydroxo comple~.~J~
Sur and Bryant The water proton relaxivities of monomeric @-cyc1odextrin)[Fe(III)TPPS4] complex increase when axial water molecules are replaced by fluoride ions at pH 7, as shown in Figure 13. The observed relaxation rates are very low, whether the fluoride ion is present or not, when cyclodextrins are absent in the solution. These observations are consistent with the hypothesis that, at pH 7, the EPR silent dimeric species is predominantly formed. The changes in the water proton relaxation dispersion profile of the [Fe(III)TPPSd] observed at neutral pH on addition of 50 mM fluoride ion are dramatic, and the relaxation rate of 14.73 mM-' s-' is comparable to that obtained at pH 2 in fluoride ion free solutions. The EPR measurements discussed earlier establish the presence of two axial fluoride ligands; therefore, the increased proton relaxation rate must derive from outer sphere effects or exchange with the water molecules hydrogen-bonded to the axial fluoride ligands as observed in fluor~methemoglobin.~~ The observation that an iron(II1) porphyrin complex may be constructed that is highly efficient for water proton relaxation suggests several synthetic and structural strategies for building magnetic relaxation or magnetic resonance imaging contrast agents using iron(II1) rather than lanthanides, which may be advantageous because of reduced toxicity.
Acknowledgment. We thank Dr. Henry Bryant for the Zn(II)TPPS4 sample used in this study. This work was supported by the National Institutes of Health, No. 5 R 0 1 GM39309. Supplementary Material Available: The water proton relaxivity data of free [Fe(III)TPPS4] and [Fe(III)TPPS4] complexed with cyclodextrins and fluoride ions as shown in Figures 11, 12, and 13 (3 pages). Ordering information is given on any current masthead page. References and Notes (1) Mansfield, P.; Moms, P. G. NMR Imaging in Biomedicine; Academic Press: New York, 1982. (2) Sur, S. K.; Bryant, R. G. J . Phys. Chem., in press. (3) Koenig, S. H.; Brown, R. D., 111Mag. Reson. Med. 1984, I , 478. (4) Bertini, I.; Luchinat, C. NMR of Paramagnetic Molecules in Biological Systems; Benjamin Cummings: Menlo Park, CA, 1986. (5) Banci, L.; Bertini, I.; Luchinat, C. Nuclear and Electron Reluxation; VCH Publishers: Weinheim, 1991. (6) Bertini, I.; Capozzi, F.; Luchinat, C.; Xia, Z. J . Phys. Chem. 1993, 97, 1134. (7) Bertini, I.; Briganti, F.; Xia, Z.; Luchinat, C. J . Magn. Reson., Ser. A 1993, 101, 198. (8) Koenig, S. H.; Brown, R. D., KI; Spiller, M. Mag. Reson. Med. 1987, 4 , 252. (9) Ivanca, M. A.; Lappin, A. G.; Scheidt, W. R. Inorg. Chem. 1991, 30,711-718 and references cited therein. (10) Mosseri, S.; Mialocq, J. C.; Perly, B.; Hambright. P. J. Phys. Chem. 1991, 95, 2196, 4659. (11) El-Awady, A. A.; Wilkins, P. C.; Wilkins, R. G. Inorg. Chem. 1985, 24, 2053. (12) Chen, C.-W.; Cohen, J. S.; Meyers, C. E.; Sohn, M. FEBS Lett. 1984, 168, 70. (13) Lyon, R. C.; Faustino, P. J.; Cohen, J. S.; Katz, A,; Mornex, F.; Colcher, D.; Baglin, C.; Koenig, S. H. Magn. Reson. Med. 1984, 4 , 24. (14) Fiel, R. J.; Musser, D. A.; Mark, E. H.; Mazurchuk, R.; Alletto, J. J. Magn. Reson. Imaging 1990, 8, 255. (15) Fiel, R. J.; Button, T. M.; Gilani, S.; Mark, E. H.; Musser, D. A,; Henkelman, R. M.; Bronskill, M. J.; van Heteren, J. G. Magn. Reson. Imaging 1987, 5, 199. (16) Ogan, M. D.; Revel, D.; Brasch, R. C. Invest. Radiol. 1987, 22, 822. (17) Furmanski, P.; Longley, C. Cancer Res. 1988, 48, 4604. (18) Doiron, D. R.; Gomer, C. J., Eds. Porphyrin Localization and Trearment; Liss: New York, 1984. (19) Inoue, Y. Annual Reports on NMR Spectroscopy; Webb, G. A., Ed.; Academic Press: London, 1993; Vol. 27, p 59. (20) Szejtli, J. Cyclodextrins and Their Inclusion Complexes; Akademiiai Kiado: Budapest, 1982. (21) Saenger, W. Angew. Chem., Int. Ed. Engl. 1980, 19, 344.
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Spin-Lattice Relaxation of Femc Porphyrins (22) Bender, M. L.; Komiyama, M. Cyclodextrin Chemistry; Springer-Verlag: Berlin, 1978. (23) Tabushi, I. Inclusion Compounds, Atwood, J. L., Davies, J. E. D., MacNicol, D. D., Eds.; Academic Press: London, 1984; Vol. 3, p 445. (24) James, W. J.; French, D.; Fundle, R. E. Acta Crystullogr. 1959, 12, 385. (25) Flohr, K.; Paton, R. M.; Kaiser, E. T. J . Am. Chem. SOC. 1975,97, 1209- 1217. (26) Tabushi, I.; Kuroda, Y.; Mizutani, T. Tetrahedron 1984,40,545552. (27) Breslow, R.; Greenspoon, N.; Guo, T.; Zarzycki, R. J . Am. Chem. Soc. 1989, 111, 8296. (28) Assenheim, H. M. Introduction to Electron Spin Resonance; Plenum Press: New York, 1967; p 102. (29) Hemandez, G.; Bryant, R. G. Bioconjugate Chem. 1991, 2, 394 and references cited therein. (30) Kessler, H.; Griesinger, C.; Kerssebaum, R.; Wagner, K.; Emst, R. R. J . Am. Chem. SOC. 1987, 109, 607. (31) Manka, J. S.; Lawrence, D. S. J.Am. Chem. SOC.1990,112,2440. (32) Goff, H. M.; Morgan, H. Bioinorg. Chem. 1978, 9, 61. (33) Yonetani, T.; Drott, H. R.; Leigh, J. S., Jr.; Reed, G. H.; Waterman, M. R.; Asakura, T. J. Biol. Chem. 1970, 245, 2998-3003.
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