E. L. LEWISAND F. Srcmo
2590
Electron Spin Resonance Kinetic Studies of Two Dimethoxymethane Radicals in Aqueous Solution by Eileen L. Lewis and F. Sicilio Department of Chemistry, Texas A & M University, College Station, Texas 77845
(Received November 18, 1968)
I n the Ti(III)-H20t flow system containing dimethoxymethane as a substrate, two radicals are produced by abstraction of hydrogen from the substrate by the “.OH” radical. Both of the dimethoxymethyl radicals undergo apparent three-halves-order decays, with k = 6.8 X loa and k = 4.0 X lo4M-’/z sec-1 at 27’. The indicating that subspecific reaction rate constant is independent of [Ti(III)]o and proportional to [HzO~]~, strate radicals are reacting with HzOZ. Also, the substrate radicals are probably undergoing addition of . OH. Absence of secondary radicals and dimer products indicates that substrate radical combination reactions and decomposition play minor roles. The system is complicated by the presence of metal complexes. The main contribution to the apparent activation energies of 14.8 and 13.0 kcal/mol could be associated with the sub-
strate radical-HzOz reactions.
Introduction The reduction of hydrogen peroxide by transition metals to yield the hydroxyl radical was reported by Uri.I Using a flow system, Dixon and Norman2 treated T i c &in acidic media with HzOz arid observed a single resonance which was attributed to the hydroxyl radical. Subsequently, the Ti(III)-H202 system has been studied by a number of investigators. Piette, et aLj3 detected two esr signals and attributed the lowfield peak (g = 2.0132) to HOz. radicals and the highfield peak (g = 2.0119) to .OH radicals. Wall, et aL14 assigned the low-field peak (g = 2.0128) as due to .OH radicals and suggested that the high-field peak (g = 2.0114) could be due to HO,. or -OH radicals complexed with Ti(1V) ions. Turlievich, et u Z . , ~ interpreted the esr spectrum as arising from Ti(1V) complexes involving OH and HOz- radicals and concluded that uncomplexed .OH radicals are not detected by esr spectroscopy. This conclusion is supported by evidence that radiolytically produced hydroxyl radicals decay within tens of microseconds, as reported by Adams and Boag.6 Fischer’ interprets the two esr peaks as not due to .OH or HOz. radicals but presumably to titanium peroxy radical species. Fischer’s suggestions were considered further by Wall, et a1.,* in the attempt to explain features observed a t intermediate substrate concentrations. These authors conclude that the two “hydroxyl” species are probably forms of HOz. complexes with Ti(1V). Takakura and RBnby9 assign the low-field peak to HOz. and the high-field peak to .OH, both radicals being coordinated to Ti(1V) ions or Ti(IV)-H2Oa complexes. Many investigations involving the production of free radicals from organic substrates included in the Ti(II1)HZOZ system have been reported since Norman and Dixon’sz initial work. Recently, Turkevich, et u Z . , ~ have presented evidence that “the formation of free 9
The Journal of Physical Chemistry
radicals of substrate molecules occurs in the complex rather than by direct action of the hydroxyl radical with the organic molecule.” In this work the diether, dimethoxymethane (methylal), was used as a substrate in the Ti(II1)H202 flow system. Kinetic studies were performed to follow the decay of two radicals formed by the abstraction of hydrogen from the dimethoxymethane molecule. Emphasis was focused on the latter stages of decay of the substrate radicals in the hope of observing features to help clarify the behavior of the system.
Experimental Section Reagents. Titanium trichloride (20% solution, which analyzed to 1.6 M TiCh) was obtained from W. H. Curtin & Co. Dimethoxymethane was obtained from Distillation Products Industries. Potassium nilrosodisulfonate, NO(S03K)2 or FrBmy’s salt, was prepared by the method of Harvey and Hollingshead.lo Other reagents were ACS reagent grade. Apparatus. The esr spectra were obtained with a Varian 4502-15 spectrometer system equipped with V-4560 100-IcHz field modulation and Fieldial units. (1) N. Uri, Chem. Rev., 50, 375 (1962). (2) (a) W. T. Dixon and R. 0. C. Norman, Arature, 196, 891 (1962); J . Chem. Soc., 3119 (1963). (3) L. H. Piette, G. Bulow, and K. LoefIler, Preprint, Division of
Petroleum Chemistry, American Chemical Society, Washington,
D. C., April 1964. (4) F. Sicilio, R. E. Florin, and L. A. Wall, J. Phys. Chem., 70, 47 (1966). (6) Y. S. Chiang, J. Craddock, D. Michewich, and J. Turkevich, ibid., 70, 3509 (1966). (6) G . E. Adams and J. W. Boag, Proe. Chem. SOC.,113 (1964). (7) H.Fischer, Ber. Bunsenges. Phys. Chem., 71, 685 (1967). (8) R. E. Florin, F. Sicilio, and L. A. Wall, J . Phys. Chem., 72, 3154 (1968). (9) K. Takakura and B. RHnby, ibid., 72, 164 (1968). (10) G. Harvey and R. G . Hollingshead, Chem. I n d . (London), 244 (1953).
259 I
ESRKINETICSTUDIES OF Two DIMETHOXYMETHANE RADICALS Kinetic and spectral measurements were made a t about 30-mW microwave power and 0.2-G modulation amplitude. At these settings, no saturation or modulation broadening was noticeable. The Varian 4548 quartz aqueous solution cell and V 4549 liquid flow mixing chamber used in this work have been described previously.6 The remainder of the flow system is described in the following section. Procedure. Normally, for each run, two aqueous solutions (0.01 M TiCI,, 0.1 M HzSO4; 0.1 M HzOz, 0.50 M dimethoxymethane) were prepared, deaerated with nitrogen, and stored in 10-1. glass reservoirs. Flow was effected by using nitrogen, up to several atmospheres, to force the solutions, independently, through glass condenser coils which were immersed in a water bath. Via insulated polyethylene tubing, the separate flows entered the mixing chamber. The mixed single stream traversed a hold-up volume of 0.16 ml before reaching the sensitive portion of the quartz flat cell in the resonant cavity. Reaction temperatures, &0.5", were measured close to the flat cell. Flow rates, up to 13 ml/sec, were monitored by timing the collection of a given volume at the exit. The flow rate of each solution was adjusted independently by flow metering valves and Teflon stopcocks. Equal flow rates were maintained for all kinetic runs, except for those in which the eff'ects of variation of initial concentrations were studied qualitatively. Reaction times (in seconds) were calculated from the ratio of hold-up volume (milliliters) to flow rate (milliliters per second).4 The composite spectra of the two dimethoxymethane radicals were recorded at various flow rates and temperatures. For kinetic measurements, the peak heights of the first derivative central hyperfine peaks (low-field signal of the doublet and central. signal of the triplet) were monitored as a function of time. For ultimate calibration, peak height was related to area under the total absorption curve, computed by graphical double integration, and thence to the low-field hyperfine peak of M manganous sulfate. Calibration was performed at 27, 15, and 3" since temperature dependency was noted for line width and sensitivity for Mn(II)4 and the ratio of peak height/area for the dimethoxymethane radicals. Linearity of the spectrometer signal level at several flow rates indicated that control of flow was good and that kinetic data could be reduced to a standard signal level with validity. Magnetic field intervals were calibrated with the nitrosodisulfonate anion (a = 13.0 f 0.1 G; g = 2.00550 f 0.0000511). To measure g values, an aqueous solution of potassium nitrosodisulfonate was introduced into the flat cell after obtaining the first low-field peak of the triplet, without interruption of the scan.12
Results and Discussion The spectrum in Figure 1 shows that two radicals, a
Y
'
15G
'
R
Figure 1. Esr spectrum of dimethoxymethane when included as a substrate in the Ti(III)-H202 flow system. Magnetic field increases to right. (a) Experimental spectrum: initial concentrations prior to mixing: 0.01 M TiCla, 0.1 M HSSO,; 0.1 M H202, 0.50 M CHaOCHgOCH3. (b) Computer-simulated spectrum (cf. program by B. D. Faubion, unpublished).
*
1:2: 1 triplet of triplets, R1 (a, = 17.3 0.2 G, a, = 0.73 f 0.03 G, g = 2.0033 f 0.0004), and a 1:1 doublet of septets, Rz (a, = 13.1 f 0.2 G, ay = 0.71 i0.03 G, g = 2.0032 f 0.0004) are formed. The parameters are not significantly different from those reported recently.I2 The radicals are derived from the parent (11) J. Q.Adams, S. W. Nicksic, and J. R. Thomas, J . Chem. Phgs., 45, 654 (1966); J. S. ISyde, personal communication. (12) R. E. Florin, F. Sicilio, and L. A. Wall, S.Res. Nat. Bur. Stand., A , 72, 49 (1968).
Volume 73, Number 8 August 1969
2592
E. L. LEWISAND F. SICILIO
+ .OH +HzOz .OH + Ti(II1) +OH- + Ti(1V) *OH
(2)
(3)
This simple mechanism is in agreement with the 2:1 consumption ratio of Ti(III)/H202 in the absence of reactive s u b ~ t r a t e s . ~ J -It~ was recognized that the .OH could be complexed with Ti(IV),4-6tQ or that two titanium peroxy radical species develop.'^* Evidently the uncomplexed .OH species is not o b s e r ~ e d . ~ - ~
0
A
8
12
16
0.5
lO~l(r.c)
Figure 2. Reciprocal square root of peak height (P-1'2) us. reaction time, at 27O, for R1. Initial concentrations prior to mixing: 0.01 M Ticla, 0.1 M HzSO~;0.1 M HzOz, 0.50 iM CHsOCHzOCH3.
0.4
compound, CH30CH20CH3, by abstraction of hydrogen. For convenience, a and y refer to protons at the site of the unpaired electron and two positions removed, respectively. Only five of the seven y-hyperfine peaks *
CH20CH20CH3
CH30cHOCH3
Ri
0.2
R2
in R2 are observed, corresponding to the intensity ratios (1) :6: 15:20: 15:6: (1). The kinetic data were treated by graphical methods. A regression analysis computer program was applied to all curves to determine which plot yielded the best straight line and to calculate the slope of that line. Kinetic plots for various orders were made for each temperature to determine which order would best describe the decays of RI and R2. The plot yielding the best straight line a t long times, in each instance, was (peak height) -'/'(P-'/') os. time (seconds), corresponding to three-halves order. However, firstorder plots appeared to fit the data fairly well also. Figures 2 and 3 are representative three-halves-order plots of the kinetic data. The slopes of each plot were used to calculate the specific reaction rate constants, k , presented in Table I. The only definite indication of formation of a substrate radical in this system was at 3", a t which temperature a maximum in concentration for R1 was observed at about 40 msec. The rate of formation is sufficiently rapid to preclude the direct analysis of the rate of formation with this experimental setup. The following mechanism has been proposed by Wall, et aL14on the basis of esr kinetic studies to explain the reaction of Ti(II1) with H202 in acidic media. Ti(II1)
+ HzO2 -+
0.3
Ti(1V)
The Journal of Physical Chemistry
+ OH- + .OH
(1)
10~t(r.c)
Figure 3. Reciprocal square root of peak height us. reaction time, a t 15', for R1. *' Initial concentrations prior to mixing: 0.01 M TiCls, 0.1 M HzSOr; 0.1 M HzOZ,0.50M CHsOCHzOCHs.
Table I : Specific Reaction Rate Constants for R1 and RZ
k, M - I / a sea -1
6.8 X 3.1 x 8.0 X 4.0 x 1.2 x 6.3 x
RI
R2
IO* 103 IOa 104 104 103
Maximum observad concentration, M
Temp,
1.1 x 10-4 6.9 X 5 . 5 x 10-5 1.5 X 6 . 1 X 10-6 5.5 x 10-6
O C
27 15 3 27 15 3
The following additional reactions should be considered in describing the system after inclusion of dimethoxymethane as a dissolved substrate in the HzOz stream
+ CHgOCH20CHg + .R1 + H2O *OH+ CHaOCH20CH3 *R2 + H2O
.OH
----t
(4) (5)
ESRKINETICSTUDIESOF Two DIMETHOXYMETHANE RADICALS
2593
Table I1 : Dependency of k for R1 on Initial Concentration" [HzOzlo = 0 . 1 M ; [CHaOCHzOCH3]a T = 27'
0.5
M;
k,
M
[Ti(III)lo
sec -1
-'/z
x
7.4
0.005
103
7 . 5 x 103 6 . 8 X 108
0.009
0.10
[Ti(III)]o = 0.01 M ; [CHaOCH20CHa]o= 0 . 5 M; T = 27' M
[HzOzlo
The '(.OH" in these equations is symbolic and does not necessarily represent a unique species. No evidence for dimeric products was found by gas chromatographic analysis of reacted streams, even when [Ti(III)l0was made 0.2 M in a titration-type experiment. Therefore, radical combination reactions 6, 7, and 8 are relatively insignificant. Reactions 9 and 10 are probably minor contributors to the decay of R1 and Rz,since no secondary radicals could be detected by esr spectroscopy. The products of (12) and (14) are probably unstable, and gas chromatographic analyses, though consistent with reactions 11-14, could not furnish a stoichiometric balance of predicted products for these latter reactions. The detection of final products in this system is complicated by reactions, such as hydrolysis, occurring after the time scale emphasized in the esr studies. Formaldehyde, formic acid, methanol, and a product in the molecular weight range of dimethoxymethane were detected, within an hour after mixing the reactant solutions. The gas chromatographic analyses were thus restricted to qualitative utility. Different intermediates, with different specificities for abstraction, did not seem to develop for the range of [Ti(III)]/[H202] ratios used in these studies. This is supported, albeit inconclusively, by observation of no change in kinetic order or-esr spectral detail as the ratio of flow rates for the two reactant streams, Ti(III)/HZO2, was varied from 0.4 to 2.6, a procedure which is effectively an alteration of initial concentrations. As can be seen from the data on actual variation of initial concentrations in Table 11, k is relatively independent of [Ti(III)lo. Quantitative evaluation of the data indicates that k is closely proportional to the first power of [H2Oz]o, indicating the probable involvement of Hz02 in the decay of the substrate radicals as represented in reactions 13 and 14. Complications due to complexation of Ti(1V) with HzOz could exert a modifying influence. The qualitative aspects of these observations agree with the Ti(III)-Hz02-CH30H systemn8 The abstractive processes represented by (4) and (5) are very rapid relative to the decay reactions for the substrate radicals. The [R1] and {Rz] observed
IC, -I/*
sec-1
3 . 0 X loa 7 . 5 x 102 13.9 x 103
0.05 0.10 0.25
All initial concentrations are prior to mixing; ratio of flow rates maintained 1: 1. Q
lO'/T
Figure 4. Arrhenius plots: log k (M-lIa sec-l) us. 1/T.
0 , RI; A,
Rz;
initially are at least several orders of magnitude greater than the concentration of complexed OH, supporting Fischer's' interpretation that uncomplexed OH is the abstracting agent. This also suggests that the rate of formation of complexed .OH is slower than the rates of the abstraction reactions 4 and 5 , as noted by Turke ~ i c h the ; ~ rate of decay of complexed .OH is also slow relative to (4) and ( 5 ) . It has previously been noted that the rate of disappearance of a substrate radical is, in general, much more rapid than that of the complexed .OH and that a quasi-steady-st,ate concentration of the latter can exist.12 Such a steady-state condition has been observed in this system with [Ti(III)l0= 0.026 M , [HzOz]o = 0.1 M , and [CH30CH20CH8] = 0.5 M .
-
Volume 73, Number 8 Auguat 1069
2594
JOHNF. ENDICOTT
It has been suggested that mixing may be incomplete even after the stream reaches the esr cavity4 and that reaction 1 cannot necessarily be assumed to be complete within the mixing chamber.13 A reasonable value8 for k~ is -lo3 M-’ sec-I. Thus, reaction 1 could very well be a minor source for -OH up to several hundred milliseconds after mixing. The pure Ti(II1)H z 0 2 system without substrate is complicated,8 arid additional sources for .OH could easily be present. The possible contribution of Ti(1V) or its complexes to the decay of substrate radicals has been mentioned p r e v i ~ u s l y . ~ Substrate ~J~ complexes in the abstraction process were considered by Turkevich;5 (4)and (5) may therefore involve complexed CH30CHSOCH3. Similarly, the substrate radicals in (11)-(14) may be serving as ligands in metal complexes, and this possibility is suggested by the low values of the velocity constants
for over-all decay relative to those expected for diffusion-controlled reactions of neutral molecules. Arrhenius plots (Figure 4) of log k vs. T-I gave apparent activation energies of 14.8 and 13.0 kcal/mol and preexponential factors of 4 X l O I 4 and 10“ M-’’2 sec-l for RI and RP, respectively. The apparent activation energies are high for purely diffusioncontrolled reactions. Because of the low activation energies for radical combination reactions, the major parts of the E A values could presumably be associated with reactions 13 and 14,respectively. Acknowledgments. This research was supported by Robert A. Welch Foundation, Grant A-177. The esr spectrometer was made available by National Science Foundation Grant GP-3767. (13) P. Smith and P. B. Wood, Can. J . Chem., 45, 649 (1966). (14) J. K. Kochi, Science, 155, 415 (1967).
The Effects of Magnetic Exchange Interactions on the Rates of Electron-Transfer Reactions
by John F. Endicott Department of Chemistry, Boston University, Boston, Massachusetts OR216 (Received November 21, 1968)
Magnetic exchange interactions between adjacent paramagnetic metals in solids and in known binuclear complexes are often very large even at room temperature. Similar interactions would be expected to occur in the “activated complexes” of at least some electron-transfer reactions. Since such interactions would define the spin alignment of donor and acceptor orbitals (at the reductant and oxidant, respectively), some magnetic restrictions on the probability of electron transfer are to be expected. A simplified application of the theory proposed for magnetic interactions in solids leads to the conclusion that at least some very slow reactions may involve a magnetic restriction on the electron-transfer probability.
Introduction Electron-transfer reactions between metal ions in solution are simple enough that very detailed mechanistic information may be obtained and that reasonably sophisticated theoretical treatments of the reaction rates have been developed. Many of the theoretical discussions have found qualitative and semiquantitative experimental justification (for pertinent reviews and discussion see ref 1-4). Despite the great deal of analytical thought and discussion, very large reactivity differences in some of the seemingly simplest systems have not yet been satisfactorily accounted for. Probably the most remarkable instance of this frustration is the -I06-fold variation in The Journal of Physical Chemistry
reactivity observed for the simple isotope-exchange rates between aquo ions in solution4 M3+ + *M2+ 142+ + *M3+ A common starting point in most theoretical discussions is the formation of a binuclear ‘‘intermediate’’ in which the reactant metal centers are near enough that the interaction of donor and acceptor orbitals can lead (1) R.A. Marcus, Ann. Rev. Phys. Chem., 15, 155 (1964). (2) W. L. Reynolds and R. W. Lumry, “Mechanisms of Electron Transfer,” Ronald Press, Inc., New York, N. Y., 1966. (3) J. Halpern and L. E. Orgel, Discussions Faraday Soc., 29, 32 (1960). (4) A. G. Sykes, Advan. Inorg. Chem. Radiochem., 10, 163 (1967).
