Helen 8. Brooks and F. Sicilio
2544
(6) T. Forster, Z. Nektrochem., 54, 42 (1950). (7) T. Werner and D. Hercules, J. fhys. Chem., 73, 2005 (1969): 74, 1030 (19701. ( 8 ) K. Nishimoto and N. Mataga. Z. Phys. Chem. (Frankfurt am Main), 12, 335 (1957); 13, 140 (1957). (9) J. R. Platt, J. Chem. Phys., 18, 1168 (1950); J. Opt. SOC.Am., 43, 252 (1953). (IO) Note the similarity of this behavior to that of doublet and triplet states of many other molecules, discussed elsewhere: H. M. Chang, H. H. Jaffe, and C. A. Masmanidis, J. fhys. Chem., 79, 1109, 1118 (1975). (1 1) E. Clar, “Aromatische Kohlenwasserstoffe”, 2nd ed, Springer, Berlin, 1952. (12) Y. Hirschberg and R. Norman, Can. J. Res., 278, 437 (1949); R. M. Hochstrasser, Can. J. Chem., 39, 1776, 1853 (1961); V. N. Lisltsyn, Didenko, and Dashevskii, Zh. Org. Khim., 4, 1086 (1968). (13) J. T. D’Agostino and H. H. Jaffb, J. Am. Chem. SOC.,91, 3384 (1969).
$4.50 for photocopy or $2.50 for microfiche, referring to code number JPC-75-2543. References and Notes (1) Author to whom correspondence should be addressed at the Department of Chemistry, University of Cincinnati, Cincinnati, Ohio 45221. (2) E. L. Wehry and L. B. Rogers, J. Am. Chem. SOC.,88, 351 (1966). (3) E. Vander Donckt and G. Porter, Discuss. Faraday Soc.. 64, 3215 (1968). (4) E. Vander Donckt and G. Porter, Trans. Faraday SOC., 64, 3218 (1968). (5) J. Del Bene and H. H.Jaffe, J. Chem. fhys., 48, 1807, 4050 (1968); 49, 1221 (1968); 50, 426 (1969); R. L. Ellis, G. Kuehnlenz, and H. H. Jaffe, Theor. Chim. Acta, 26, 131 (1972).
Electron Spin Resonance Study of the TiF2+-H202 Reaction System Helen B. Brooks and F. Sicllio* Department of Chemistry, Texas A&M University, College Station, Texas 77843 (Received June 25, 1975) Publication costs assisted by The Robert A. Welch Foundation
The Ti(III)-HzO2-substrate reaction system has been used extensively to generate organic free radicals in aqueous media. The Ti(II1) decay in this system is not normally observed by ESR. We have studied the reaction of TiF2+ with H202 in which decay of the reactant TiF2+ can be followed by ESR. Inclusion of an organic substrate to the TiF2+-H202 system leads to the formation of organic radicals as in the case of the Ti(III)-H202 system. The TiF2+-H202 system has the distinct advantage that both a reactant and the organic radical can be followed simultaneously, leading to an obvious delineation of mechanism,
Introduction The reaction of H202 with Ti(II1) in the presence of an organic substrate has been used to generate organic radicals in aqueous media.1,2Kinetic studies have been made with a continuous flow system monitored by ESR.2,3 However, the reactants cannot be monitored by ESR. The kinetics of Ti(II1) decay have been studied by a stopped-flow spectrophotometric t e ~ h n i q u e . However, ~ these studies must be carried out with excess Ti(III)* and the ESR studies must be carried out in excess H ~ 0 2 . ~ In this work fluoride was added to the Ti(II1) reactant solution to form the complex species TiF2+ which can be monitored by ESR.5 This system allows a reactant and also radical intermediates to be observed simultaneously so that the mechanism may be delineated more critically. Experimental Section Reagents. Titanium(II1) solutions were prepared from W. H. Curtin & Co. technical grade 20% Tic13 and were standardized by titrations with KMn04 (KyGz04 was the primary standard). H202 solutions were prepared from Baker Analyzed Reagent grade 30% solutions and standardized with KMn04. Baker reagent grade NaF and Fisher Certified ACS NaF were both used for preparing solutions. Methanol, sulfuric acid, and hydrochloric acid were reagent grade. All solutions were deaerated by bubbling nitrogen for 20 min. Equilibrium Studies. Two mole ratio studies were done The Journal of Physical Chemist!y, Vol. 79, No. 23, 1975
on a Cary 14 spectrophotometer. Concentrated Ti(II1) and F- solutions were mixed and diluted to volume with deaerated water. Two wavelengths, 4350 and 4860 8,were selected, and additional data for mole ratio and continuous variation experiments were collected with a Beckman DU. The ESR data were developed using a Varian 4502-15 spectrometer and a Varian V-4556 flat cell. Each set of data was collected continuously to minimize errors. The flat cell was removed between solutions, but the instrument was retuned by adjusting the flat cell position rather than the instrument settings. Known runs gave agreement within f5% for this technique. Two sets of data were collected a t different pW, using both ESR and a Beckman DU on the same solutions. The pH was monitored with a Corning Model 101 digitaI electrometer and HCl was added to maintain constant pH. Kinetic ESR Procedures. The flow system and ESR apparatus has been described previously.6 Solutions were prepared from deaerated water and stored in 10-1. glass reservoirs. Under nitrogen pressures of 2 atm, the separate streams were forced through a standard Varian 4547 mixer and a Varian 4548 quartz flat cell via polyethylene tubing. The flow rate was regulated by a needle valve on the exit side of the flat cell. An average flow rate was measured by timing the collection of 100 ml of product solution. The system was set up so that the flow rates from each reservoir were identical. The “dead times”, t, for the kinetic data were calculated by the relationship t = V / r where V is the dead volume and r is the flow rate. Spacers were inserted
2545
ESR Study of TiFZ+-HzOz
TABLE I: TiFZf Decay without Substrateu [H20210,
M
[H2S04107
M
[F-I,, M
0.05 0.05 0.05 0.25 0.25
0.035 0.010 0.03 5 0.060 0.022
kaw
sec''
kl, sec-'
-
., c
0.25 0.25 0.10 0.10 0.10
E
- IO
.-
Q C
.-c
c
::
1 - 5
/([Ti(D.)I,
24 42
15 28 30
45
40
[Ti(III)]o = 0.01 M ,before mixing. All initial concentrations are halved upon mixing. a
5
+ CF-1,)
Figure 1. Continuous variation method: temperature 25', pH 1.28, [Ti(lll)]o= 0.01: (El)visible absorption, 4350 A; (0)ESR absorption, in arbitrary units, for the same solutions.
IO0
80
P .I-
s
36
W r
[F-],
51 50
6C
E = c
z 4(
21
ted in Figure 1. The spectrophotometric data show conclusively that a 1:l complex, TiF2+,is present and has a large extinction coefficient. The ESR data show that the sharp peak monitored in kinetic studies is that for a 2:l complex, TiF2+. Also, mole ratio experiments were performed with both spectrophotometric and ESR techniques. Quantitative calculations of equilibrium constants were inconsistent at different pH. However, calibration of the maximum TiF2+ ESR intensity with a V02+ standard and a Mn2+ standard shows a region of 100%TiF2+, within experimental error. A typical plot of a mole ratio experiment is shown in Figure 2. Since the primary concern of this research is the kinetics of the TiF2+ decay, the rate at which equilibrium between F- and Ti(II1) is established is important. Ti(II1) solutions were mixed with F- solutions in the same continuous flow system used to perform the peroxide experiments. Several initial concentrations of Ti(II1) and F- were used. The first point was always a t maximum [TiF2+] for the system, indicating that equilibrium is established very rapidly compared to the peroxide reactions. Therefore, a condition of equilibrium between F- and Ti(II1) is assumed. Kinetics without Substrate Present. The decay of TiF2+ was observed as a function of time. Plots of In [TiF2+]vs. t were curved sharply. The slopes of these curves during the first two half-lives are tabulated as kapp in Table I. The data were replotted assuming that Ti(1V) competes with Ti(II1) for .OH in the reaction scheme TiFz+
-
+ H202 ki TiF20H+ + .OH
(1)
CF-1, /[TI (lU0,
Figure 2. Mole ratio ESR experiment; percent Ti as TiF2+ determined by comparison of areas of ESR absorptions of unknown Ti-F system with standard VOz+ and Mn2+ solutions: [Ti(lll)]o = 0.01; (13) 0.25 M HCI; (0)0.05 M HCI. between the mixer and flat cell for some slower reactions. The volume of a spacer was determined from the weight of water required to fill the spacer. A run of a slow reaction with three different dead volumes produced data which overlapped well. Calibrations of concentration were made by comparison with standard Mn2+ solution. When the concentrations were of special interest, both Mn2+ and V02+ standard solutions were used for comparison. Results Equilibrium. Continuous variation plots were made to determine the species present in the titanium(II1)-fluoride solution. Both spectrophotometric and ESR data are plot-
TiF20H+
+ .OH 2 SI
(3)
(SIis a free-radical intermediate and is observed in the sysThe data fit well out to four half-lives when k2 is set equal to k3. Slopes of these plots are tabulated as kl in Table I. The decay rate shows no significant changes with variation of concentration of F- even when only 10%of the Ti(II1) is present as TiF2+. Hence, TiF2+ and TiF2+ must have similar decay rates. Kinetics with Substrate Present. Figure 3 shows typical kinetic plots for the decay of TiF2+ and sCH20H. All of our In [TiF2+] plots show some curvature. The slopes of the best straight lines for up to 90% reaction are denoted kapp in Table 11. The slopes of In [.CH20H] vs. t plots are denoted kR. Initial conditions were varied to determine a rate law for the reaction. Variation of k, with [H202]0, plotted in Figure 4, indicates a first order dependence on H202. VariaThe Journal of Physical Chemistry, Vol. 79, No. 23, 1975
2546
Helen B. Brooks and F. Sicilio
1 zxio-'
0.10
005
2
fsec)
Flgure 3. Kinetic plot of typical data: solution A ([TiCls]~= 0.01 M, [NaF]o = 0.03 M, [CH30H]o = 0.50 M, [HCl]o = 0.05 M) mixed with solution B ([H202]0 = 0.25 M, [HCI]o = 0.05 M): (A) TiF2+ ESR absorption; (0) .R ESR absorption.
TABLE 11: Rate Dependence Studiesa
[HzOz]o, M
Anion
kaDD,sec-'
k,, sec-'
A. Peroxide Variationb
c1-
0.40 0.40 0.25 0.22 0.10 0.10 0.087 0.05
46 54 26 23 12 17 15 6
HS06'
c1c1c'1 HS04' HS04HS04-
23 25 14 11 6.1 7.4 7.0 3.6
[Ti(IIDIo,iZI [HzOz]O, M [F-],M kaDp,sec-' k,, sec-' B. Ti(I1I) VariationC
0.05 0.01
0.40 0.40
0.15 0.03
55 46
22 23
C . Ti(IV) Variation'
0 .oo
16 17 18 17 19
0.01 0.015 0.03 0.04 Acid
[Acid],, M
HC1
0.25 0.10 0.05
6.7 7.2 6.O 7.6 6.7 kQD'
[H'), M
8ec-l
kR, sec-'
D. Acid Variatione HZSOI H C1
0.125 0.06 0.025
27 16 12
16 6.7 6.1
a Concentrations listed are prior to mixing. Upon mixing the concentrations are halved. [Ti(III)]o= 0.01 M; [HClIo = 0.05 M; [H&304]0= 0.05 M, [F-lo = 0.030 M in C1- mediaand 0.035-_"i"n i'; [HzOzlo = HS04- media. [HClIo = 0.05 M. [Ti(III)]o= 0.01 h 0.10 M , [F-Io = 0.04 M ; [H2S04]o= 0.10M; [CH~OHIO = 0.5 M .
The Journal of Physical Chemistry, Vol. 79, No. 23, 1975
Figure 4. Peroxide dependence; [Ti(lll)]o= 0.01 M, [F-] = 0.030 M in CI- media and 0.035 M in H2S04 media: (0) kapp from TiF2+ decay, [HClIo = 0.05 M, slope 110; (El) k~ from radical decay, [HCI]o = 0.05 M, slope 55: (a)kap from TiF2' decay, [H2S04]0 = 0.05 M (A)k~ from radical decay, fH2S0410 = 0.05 M. tion of [Ti(III)]oindicates that a first-order dependence on Ti(II1) is correct since ksPp is within experimental error. Variation of [Ti(IV)]oindicates there is no significant effect from Ti(1V) on either the TiF2+ reaction or the substrate reaction. Even the [.R] at any time did not change significantly with change in [Ti(IV)]. Acid variation indicates an acid-independent step and an acid-catalyzed step. The data reported in Table IID follow the equation kapp
+
= 1.6 X 102[H202] 3.0 X 1O3[Hz02][H+]
Data between different runs at different acid strength would not be good enough to discern a more complex acid dependence. Methanol concentration was also varied. Identical data were obtained for 0.5 and 1.0 M methanol. A t 0.1 M the slopes of the curve did not change significantly, but the [-R] at any time was lower than [.R] in the experiments at 0.5 and 1.0 M. The intermediate SI was not observed in any of these experiments with substrate. The fluoride ion dependence is more complex. At IF-]/ [Ti(III)] ratios greater than unity the TiF2+ decay curves have the same slopes and curvatures. However, the radical decay curves are the same only where TiF2+ is the predominant species present. Excess F- shows no effect on the radical decay. Even [.RJ at any time is the same for 0.04 and 0.09 M F- ([Ti(III)] = 0.01 M ) . Hence, the fluoride ion and HF are not involved in the radical reactions. For a small amount of F- (0.002 M) and 0.01 M TiCl3, the TiF2+ concentration remains nearly constant at 2 X lom6 M for what would be about one half-life for TiF2+, which would be equivalent to 2.5 half-lives for Ti(II1). Then the concentration decays with the normal half-life of TiF2+. The radical decay is the same as that in the absence of F-.
ESR Study of Ti(lll)-H202
2547
Discussion Equilibrium data indicate the presence of both TiF2+ and TiF2+ in the system, but a system with primarily TiFZ+ can be attained by adjusting the F- concentration. Small amounts of TiF2+ do not effect the kinetics as shown by variation of [F-1, The decay of TiF2+ with substrate present can be described by (1),followed by
+
*OH RH 2.R
k.5 +
+
-R + HzO
products
(4)
(5)
where (1)is the rate-determining step and (5) is the termination step. This mechanism is in agreement with the firstorder dependences on [TiF2+] and [HzO,] observed in the rate law for TiF2+ decay. Assuming that [.R] and [.OH] are in steady state, then the rate constant for the decay of -R, KR, is calculated to be kR
In a system containing primarily TiF2+ the faster k R may be due to the competition of the analog of reaction 3 with reaction 4. This would change only the observed [.R] and not the TiF2+ decay rate. In conclusion, the TiF2+ substrate system is in steadystate equilibrium where reaction 1 is the rate-determining step and reaction 5 is the terminating step. With substrate in excess reactions 2 and 3 are unimportant. The important sequence is (l),(4), and (5). This fluoride system has the advantage that the initiating reaction can be observed along with the substrate radicals in a single system.
kappI2
where kapp is the slope for the TiF2+ decay. In fact, this was observed for all runs where [F-] is large enough for TiF2+ to be the predominant species. Reactions involving SIand Sz, do not the other titanium-containing compete with substrate in the TiF2+ system.
Acknowledgment. This research was supported by the Robert A. Welch Foundation, Grant No. A-177. References and Notes (1) R. 0.c. Norman and P. R. West, J. Chem. SOC., 13, 389 (1969). (2) R. E. James and F. Sicilio, J. Phys. Chem., 74, 1166 (1970), and references therein. (3) G. Czapski, A. Samunl, and D. Meisel, J. Phys. Chem., 75, 3271 (1971). (4) A. Samuni, D. Meisel, and G. Czapski, J. Chem. Soc., Dalton Trans., 1273 (1972). (5) L. 0. Morgan, "Advances in Chemistry of Coordination Compounds", MacMlllan, New York. N.Y., 1961, p 471. (6) E. L. Lewis and F. Sicilio, J. Phys. Chem., 73, 2590 (1969). (7) R. E. James and F. Siclllo, J. Phys. Chem., 74, 2294 (1970). (8) ti. Fischer, Ber. Bunsenges. Phys. Chem., 71, 685 (1967). (9) H. 8. Brooks and F. Sicilio, J. Phys. Chem., 74, 4565 (1970).
Electron Spin Resonance Study of Effects of Sulfate and Chloride Ions on Kinetics of the Titanium(ll1)-Hydrogen Peroxide Reaction System Helen B. Brooks and F. Slcllio" Department of Chemistry, Texas AbM University, College Station, Texas 77843 (Received June 25, 1975) Publicationcosts assisted by The Robert A. Wejch Foundation
Previous studies of the Ti(III)-H202-substrate reaction system have produced divergent views on the mechanisms of reactions involving organic radicals in this system. Sulfuric acid has been used conventionally as the acidifying medium to prevent the occurrence of precipitation. A critical study of the effects of sulfate and chloride ions indicates conclusively that a 1:l titanium(II1)-sulfate complex is responsible for the anomalous kinetic behavior in the sulfate-containing system. As reported by other investigators, the reaction of Ti(II1) with H202 is rate determining, leading to a secular equilibrium in which the decay behavior of the organic radical is regulated by the reaction of Ti(II1) with H202. The present study shows that the titanium(II1)-sulfate complex is responsible for the much faster decay rates observed at shorter times.
Introduction In ESR-flow studies of the Ti(IWH2O2-alcohol reaction system, the initial reactions are1-3 Ti(II1) + H202 .OH
ki
Ti(1V) + .OH + OH-
(1)
+ H2O
(2)
+ alcohol
-
.R
The concentration of the alcohol substrate radical, [.R], has
been monitored as a function of time and initial concentrations of r e a ~ t a n t s . l -In~ this reaction system, the solutions must be acidified to prevent precipitation of titanium pounds. In previous work, sulfuric acid was chosen as the medium because both HC104 and HCl undergo complicating reactions. At concentrations used, HC104 oxidizes Ti(II1) in about half the time required to complete an ex~ e r i m e n tThe . ~ C1- ion in HCl solutions undergoes the reaction
+
*OH C1-
-+
*C1+ OH-
(3)
The Journal of Physical chemistry, Vo/. 79, No. 23, 1975
