3154
R. E. FLORIN, F. SICILIO,AND L. A. WALL
The Paramaglretic Species from Titanous Salts and Hydrogen Peroxide by Roland E. Florin, Fred Sicilio, and Leo A. Wall h7atwnal Bureau of Standards, Washington,D . C. 8088.4
(Received January 88, 1968)
The radical species giving esr spectra on mixing titanous salts and hydrogen peroxide cannot be hydroxyl as formerly supposed. Observed kinetics are inconsistent with simple generation and disappearance schemes. In organic-substrate mixtures the species increases with time, while organic radicals decrease. The two “hydroxyl” species are probably forms of HOZcomplexed with Ti4+. The rate constant kl of the initial reaction, Ti*+ HzOz -+ OH OHTi4+,is estimated as equal to or greater than 200 M-l sec-l from the appearance rate of Ti4+-Hz02and as 800-1800 M-’ sec-’ from indirect analysis of radical concentration-time curves.
+
+
+
Introduction Upon mixing aqueous hydrogen peroxide and titanium trichloride in a flow system, Dixon and Norman observed an esr spectrum which they ascribed tentatively to hydroxyl.’ Organic additives were attacked to yield plausible hydroxyl abstraction products Tia+
+ H2Oz Ti4++ .OH + OH.OH + R H +R * + H2O
(1) (2)
Later investigators2-6 found complicating features and proposed modifications in the structure. We have made further observations bearing on the rate of reaction 1, the stoichiometry and incidental products of the over-all reaction, and the kinetics of organic-substrate mixtures. According to our results, which are related to those of Fischer6 and Turkevich, et u Z . , ~ it is now very unlikely that the observed spectra are those of a primary reactive species such as OH.
Experimental Section Esr Spectra and Kinetics. The general experimental arrangements have been described2 and resemble those used elsewhere.’-6 I n the present work, the only “dead volume” used was 0.10 cm3. Titanium Kinetics. A T-shaped mixer, followed by a uniform cylindrical downstream region housed in a sliding comparator block, was used to observe the rate of appearance of the deep yellow Ti4+-HZ02 complex and also, for concentrated solutions in stoichiometric ratio only, the rate of disappearance of the purple Ti3+ color. Reference solutions, containing known concentrations of the Ti4+ complex in tubes of the same diameter, were moved up and down stream until a visual match was obtained. Note was taken of the downstream distance matching half the final Ti4+concentration, as a function of flow rate. The time corresponding to this distance was computed with the aid of the measured cross section d by the formula t = ?rr2s/F,where s is the downstream distance and F is the flow rate, Although static comparisons could distinguish concentration differences of f l O % of the usual The Journal of Physical Chemistry
final concentration of the yellow complex, it is unlikely that times are more reproducible than f20%, which was the probable error of the apparent rate constant in one series of observations a t different flow rates. Titrations. Approximately 0.05 M hydrogen peroxide and 0.01 M T i c 4 were made up by dilution with 0.1 M sulfuric acid. Titrations were done to the first appearance or final disappearance of a faint yellow peroxide-complex color. The titanium solution was standardized immediately after use with acid permanganate. The peroxide concentrations are relative to an assumed consumption ratio of 2.00 Ti/1.00 HzOzin the usual titration, which agrees with the prepared concentration within 10%. In the more careful work, the Tic13 solution was stored in a deep container and was withdrawn from the bottom into a nitrogen-flushed buret, and the titrations were done in flasks, in a flowing nitrogen stream to guard against the fairly rapid oxidation of titanous chloride by atmospheric oxygen.6 Detection of Chemical Products. Solutions of equal molarity (i5%) in titanous chloride and hydrogen peroxide, containing methanol a t several molar concentrations and 0.1 M sulfuric acid, were passed through a mixer into a receiver at 10 f 5 cm3 sec-’. The product was adjusted to a colorless end point with one of the two solutions, usually 1-4% of the total volume. After a delay of 1 hr, formaldehyde was estimated in a portion of product by a Schiff’s reagent m e t h ~ d . ~ (1) W. T. Dixon and R. 0. C. Norman, Nature, 196, 891 (1962); J . Chem. SOC.,3119 (1963). (2) F. Sicilio, R. E. Florin, and L. A. Wall, J. Phys. Chem., 70, 47 (1966). (3) L. H. Piette, Preprint, Petroleum Division, 147th National Meeting of the American Chemical Society, Philadelphia, Pa., April 1964. (4) Y. S. Chiang, J. Craddock, D. Mickewich, and J. Turkevich, J . Phys. Chem., 70, 3509 (1966). (5) H.Fischer, Ber. Bunsenges. Phys. Chem., 71, 685 (1967). (6) H.A. E. Mackenzie and F. C. Tompkins, Trans. Faraday SOC., 38, 465 (1942). (7) J. E. Walker, “Formaldehyde,” 3rd ed, Reinhold Publishing Corp., New York, N. Y., 1954,p 458; for an evaluation, see C. N. Satterfield, R. E. Wilson, R. M. LeClair, and R. C. Reid, Anal. Chem., 26, 1792 (1954).
3155
THEPARAMAGNETIC SPECIESFROM TITANOUS SALTSAND HYDROGEN PEROXIDE The remainder was neutralized to pH 7 ==! 0.5 with concentrated NaOH. A 5-10-ml portion was separated from salts by 5 hr of vacuum sublimation, room temperature to liquid-nitrogen temperature, followed by 30 min of heating of the residue a t 100”. After storage of the condensate, 1-10-days a t lo”, samples were analyzed by vapor chromatography, using a 10-ft by a/16-in. stainless steel column, containing 20% Carbowax 20M on 60-80 mesh Chromosorb W (acid washed) with a programmed column temperature of 100-BO”, and a dual flame ionization detector. The column was calibrated with a synthetic mixture of 2 M methanol, 0.01 M formaldehyde, 0.01 M glycol and water; however, recovery efficiency in the complete experiments is uncertain. The flame-ionization detector was not sensitive to formaldehyde.8
C,M
-
-
\
5 -
-
e\
-
‘\
-
‘0,
‘20
2-
-
0
10-8
I
I
Results Radical Concentration-Time Curves. When large amounts of an organic reactant are present, the curves of the logarithm of radical concentration against time can be regarded as somewhat distorted straight lines. The slope is primarily governed by the hydrogen peroxide concentration, provided a large excess is present, and the general location is rather insensitive to substrate concentration provided that this is high. This interpretation is illustrated in Figure 1 and will be discussed in greater detail e l s e ~ h e r e . g ~ ’ ~The ~ ~ ~gross * behavior can be rationalized by the reaction sequence
+ H2OZ-+ .OH + OH- + Ti4+ -OH + RH R. + H2O R. + Re -+ products
Ti3+
(1) (2) (3)
From the steady-state condition kl(H20z)(Ti3+)= kR2 the slope of the plot of log [ E ]us. t should approach the value -h(HzO2)/(2 X 2.303) The rather indefinite kl’s thus deduced are in the range 880-1840 M-‘ sec-’. Other interpretations are also possible; the curves of Figure 1suggest that the simple picture needs at least minor modification.
Concentration-Time Curves at Lower Substrate Concentration. At methanol concentrations typically below 0.05 M but dependent upon peroxide concentration, both the organic radical and the radical previously identified as OH are present, Figure 2. In the range of shorter times, the “OH” concentration grows while the CH2OH concentration is falling. Others have reported the same change^.^ A possible explanation is that the substrate becomes exhausted, which, in view of the relative amounts of Ti3+and C&OH, would necessitate a chain reaction. A chain reaction is not too unlikely per se,ll but an objection to the exhaustion hypothesis
TIME, rnstc
Figure 2. Radicals and “-OH” a t intermediate substrate concentrations: temperature, 22”; 0.05 M HzOz; 0.005 M Ticla; (A) 0, “*OH;” (B) 0, “.OH;” (C) A, CHzOH from 0.025 M CHsOH; (D) A, CHzOHfrom 0.017 M CHsOH. One concentration unit equals 3.9 X low8M ‘CHIOH or 3.3 X 10-0 M “-OH.” (8) Thanks are extended to Dr. J. M. Antonucci, National Bureau of Standards, and Dr. E. Ragelis, Food and Drug Administration, for the chromatographic determinations. (9) R. E. Florin, F. Sicilio, and L. A. Wall, J. Res. Nut. Bur. Stand., A72, 49 (1968). (10) R. E. Florin, F. Sicilio, and L. A. Wall, unpublished data. (loa) NOTEADDEDIN PROOF. In all figures concentrations must beincreased by the factor 35/3, related to quantum mechanics of the calibration with aqueous manganese. This does not affect the derived values of b ~ .
Volume 7% Number 9 September 1068
3156
R. E. FLORIN, F. SICILIO, AND L. A. WALL
I
0
is shown in Figure 4. Results are not directly comparable with those of ref 2 because of the incompatible time scales, as discussed there. The most striking result is the much higher concentration reached in the sulfate system, curve A, although spectra are qualitatively similar. Unsuccessful attempts were made to produce radicals by the reaction of titanium trichloride with oxygen M ) in 0.1 M sulfuric acid. No (saturated, 1.3 X trace of the “hydroxyl” spectrum was seen and no cH20H spectrum was seen when methanol was present, although hydrogen peroxide at similar concentrations, low3M , produced spectra several times above noise level. The reaction
I
+ 02 + H+ +Ti4+ + €IOz.
Ti3+
I
2.0
I
I
I
I
I
)
is a conceivable path for the obviously rapid consumption of Ti3+exposed to oxygen and could be a factor in the nonappearance of O2 when Ti3+ and HzOz react. If the reaction occurs, either it is too slow for detection of the HOz. produced or else the spectrum of the latter is broader than those seen with Ti3+ HzOz. Some related reaction is likely, however, since dilute solutions of polyvinyl alcohol with Tic13 are rapidly degraded when oxygen is bubbled through them.lO Titanium Consumption Rates and kl. The kl values computed from the downstream distance for half color change, flow speed, and tube diameter, are assembled in Table I. The uncertainty is great, but the order of magnitude agrees with the old instrumental determination by Chance,lZwhose result, 430-750 M-l sec-l, is probably the one to be preferred. The rate of color formation was apparently the same in the presence or absence of substrate, but the result is not very firmly established because of the large uncertainty. The significance of these values and of the value given by Chance is ambiguous in several ways. The reactions in question are
+
20
40
60
80
TIME,rnrec
Figure 4. Concentration of the supposed “.OH” radicals produced a t small or zero organic-substrate concentration: temperature] 22’; (A) 0.05 M HZOz,titanous sulfate 0.005 M in Ti3+; (B) 0.1 M HzOZ,0.05 M Ticla, no substrate; (C) 0.005 M CH30H, 0.05 M HzOZ,0.005 M TiCl3; (D) 0.05 M HzO,, 0.005 M Ticla, no substrate; (E) 0.01 M CHsOH, 0.05 M HzOz, 0.005 M Ticla; (F) 0.025 M HzOz,0.005 M TiCI, no substrate; (G) 0.015 M CHsOH, 0.05 M HzOS,0.005 M T i c l a ; (H) 0.05 M H z O ~titanous , sulfate 0.114 M in Ti3+.
is that the same behavior occurs with massive concentrations of the less reactive substrate acetic acid, which could not change appreciably during reaction, Figure 3. At the lowest substrate concentrations, the ‘‘OH” alone is seen. Its concentration grows more slowly and to lower maximum levels than in pure HzOz, Figure 4. These results are difficult to reproduce. It seems established that a t some substrate concentrations the concentration of “OH” can reach higher maximum levels than in pure hydrogen p e r o ~ i d e . ~ Titanium and Peroxide without Substrate. The growth of the hydroxyl-like radical for several conditions The Journal of Physical Chemistry
Ti4+
+ H20z
----f
Ti4+-HzO2
(4)
and
OH
+ Ti3+
4Ti4+
+ OH-
(5) If eq 4 is always rapid, then the result measures kl provided eq 5 does not occur. If eq 1 is always followed rapidly by eq 5 , at least up to half-consumption, then the kl of Table I really represents 2 k l ; on more careful investigation, the titanium consumption might show a change of regime, being governed initially by kl (OH low), soon after by 2kl (Ti3+ and OH both high), and finally again by kl (Ti3+ low). If eq 4 is comparable with or slower than eq 1, the result merely states that kl (or 2kl) is at least as great as the tabulated value. It should be noted that the kl of Table I are only about (11) J. H. Merz and W. A. Waters, J . Chem. SOC.,615 (1949). (12) B. Chance, J . Franklin Ifinst., 229, 737 (1940).
THEPARAXAGNETIC SPECIES FROM TITANOUS SALTSAND HYDROGEN PEROXIDE Table 111: Chemical Product Recovered from Methanol Oxidation"
Table I: Titanium Color Kinetics cross section (mixer), om2
Temp, "C
0.036 0,048 0.048 0.036 0.036 0.036 0.036 0,129 0.129
22 22 48 22 22 22 22 22 22
(HzOz), M
0.05 0.0125 0.0125 0.05 0.05 0.05 0.05 0.025 0.025
(CHsOH), M
0 0 0 0 1.0 0 2.0 0 2.0
Half-life, seo
0.048-0.070 0.260-0.310 0.070-0.087 0.036-0.076 0.051-0.070 0.025-0.040 0.038-0.049 0,106-0.130 0.129-0.129
hi,
280-580 200-240 780-1 180 186-386 198-272 350-555 284-370 210-260 210
Table 11: Titration of Titanous Chloride with Hydrogen Peroxide (g-atomti of Tis+ per mole of HzOz) Peroxide into titanium
Titanium into peroxide
2.00 1.24
0.96
...
Concn of titanous ohloride and hydrogen peroxide (prepared) 0.03 0.03 0.01 0.01 0.01 ---Conon of methanol (prepared)-----? 2.0 2.0 2.0 2.0 0.67 r Conon of product (found)---
M-1
sec-I
one-fifth as great as those deduced from slopes in Figure 1, which range from 880 to 1840 M-l see-'. If the interpretation of Figure 1 is correct, then reaction 4 is slow and rate determining. Stoichiometry. A l-mol amount of hydrogen peroxide consumes roughly 2 g-atoms of titanium when no substrate is present and consumesnearly 1 g-atomwithahigh concentration of substrate present, as shown in Table 11. These results were obtained in a nitrogen atmosphere; in air the titanium consumption tends to be higher. The low result indicated in the second column of the first row was accompanied by a strong smell of chlorine in the effluent nitrogen stream. I n 0.5 M polyethylene oxide, titrated in air, an approximate 1: 1 consumption rakio was also obtained.
Pure (in 0 . 1 Id sulfuric acid) Methanol added (2 M ) Polyethylene oxide added (0.5 M ) , titration in air
3157
b
Formaldehyde
All concent.rations are in moles per liter of the final mixture volume (0.1 M HzS04). Not determined.
Possible complexing of the OH in metal ion-hydrogen peroxide systems is also suggested by the variations of the relative reactivity of a- and &hydrogen atoms, e.g., in ethanol, as one changes the source of the hyd r o ~ y l . ~ ~ ~However, '*~s several features of the kinetics reported here are incompatible with the identification of the spectrum with any form of reactive hydroxyl. The steady-state interpretation of radical concentration-time curves at high substrate concentrations, advanced in connection with Figure 1, agrees conceptually with that of Fischer,lg although the latter does not need to consider explicit time dependence. Second-order curves taken individually, l / c os. t, actually make a better fit to the data, but they are mutually inconsistent, unlike those of Shiga1*with ferrous systems. Other interpretations of Figure 1 are difficult to exclude. For example, other reactions beside eq 3 could consume radicals
+ Ti4+-+ Q (nonradical) R . + Ti4+ R+ + Tia+ R+ + HzO +ROH + H"
R.
----f
1.50 1.o
Chemical Products. Ethylene glycol and formaldehyde were recovered from the rapid mixing of equimolar amounts of titanous chloride and hydrogen peroxide, containing methanol substrate. The amounts reported in Table I11 may be low by a large factor if recovery was inefficient.
Discussion Previous measurements2 indicated a possible disappearance rate constant of 4.5 X 105 M-l sec-l for the supposed "hydroxyl," which is very much less than the value for ordinary OH deduced from pulsed and steadystate radiolysis.'a-16 The discrepancy of rates alone is not completely conclusive, since it could be brought about by complexing and consequent charge effects.
0.0038,
b 0.0012 b 0.0054 0,0023 0.0013 0.0005 0.0004 0.0005
Glycol
+ HO-OH Q (nonradical) R . + HO-OH +ROH + *OH
R.
4
(6)
(7) (8) (9)
(10)
Several mechanisms involving one or more of reactions 6-10 as well as eq 1-3 yield closed-form equations reproducing the behavior shown in Figure 1. I n most of these cases the slope and the color-appearance rates of Table I lose their simple significance. Reaction 10, once advocated for the Fenton's-reagent system, was (13) H.A. Schwartz, J . Phys. Chem., 66, 255 (1962). (14) G. Czapski and L. M. Dorfman, ibid., 68, 1169 (1964). (15) J. H. Baxendale, "Pulse Radiolysis," Academic Press Ino., New York, N. Y.,1965,pp 17,25. (16) M.Anbar and P. Neta, Int. J . Appl. Radiat. Isotopes, 16, 227 (1965). (17) R.Livingston and H. Zeldes, J. Chem. Phys., 44, 1245 (1966). (18) T.Shiga, J . Phys. Chem., 69, 3805 (1965). (19) H.Fischer, Makromoz. Chem., 98, 179 (1966).
R. E. FLORIN, F. SICILIO, AND L. A. WALL
3158 abandoned in part by Watersz0in favor of a combination involving analogs of eq 1 and eq 7, e.g. Re
mechanisms involving OH can be shown to be inadequate
+ Fe3+ -+ Fez+ + R + +
+eG R : -
(E+,) By observing the occurrence or failure of reduction in reference systems such as Fe3+-Fez+, Sn4+-Sn2+, and methylene blue, it was estimated that E-, was less than 0.4 V for a cyanoalkyl radicalz2and perhaps -0.2 V for a radical RcHOH.21 Since the value for Ti02+-Ti3+ is near 0.1 V,23at p H 0, by the convention in which more positive values indicate stronger oxidants, reaction 7 seems thermodynamically possible for some radicals.24 However, if eq 7 or 10 occurs, the rate cannot be a large multiple of the rate of reactions 1 and 3. Both reactions would increase the consumption of peroxide relative to titanium beyond the limits of Table 11. The glycol yield, Table 111, demonstrates the occurrence of the combination reaction (eq 3) to some extent, possibly much greater than indicated; the recovery of formaldehyde, a possible primary product of eq 7 or 10, is no greater. Thus although the reactions 7 and/or 10 may modify the set of eq 1-3 appreciably, they do not seem to be constituents of a long chain reaction. Unlike curves of radical concentration, our previous OH”-disappearance curves2 cannot describe a steadystate concentration of the species, since in their time range the titanous ion is almost completely consumed and the radical formation rate is negligibly small. When both the “OH” and R are present, the changes with time, Figure 2, are contrary to reasonable expectations for the true OH. The reactions involving radicals can be consolidated as R.
f(
d(R)/dt = kz(RH)(OH)
- h ( R ) ( O H ) - L3R2
which includes R
+ OH = ROH
(11)
The steady state in R involves
The ratio on the left should give a linear plot against (R) with slope -k11/k3 and intercept kz(RH)/ka. Without the need for a plot, the data of Figure 2 are enough to show that this relation is incorrect; with increasing (R) (shorter time), (R)2/(OH) does not decrease but increases. The behavior at intermediate concentrations, Figure 2, seems too complex to account for a t present; in the example just given, the supposed (OH) and (R) are related very strangely for any simple mechanism. I n the pure Tia+-H20zsystem, two of the simpler The Journal of Physical Chemistry
+ H20z-+-OH + OH- + Ti4+ OH + Ti3+ OH- + Ti4+
TiS+
Reactions such as eq 7 have been discussed in terms of redox potentials21z22 Re R + e (IL)
(1) (5)
4
and Ti3+
+ HzOz+OH + OH- + Ti4+ OH + OH +HzOz
(1) (12)
These two cases have been developed completely in the literature.26~26 Assuming that we can identify the radical species of ref 1-5 with OH, the experimental data to be accommodated are: the maximum concentration, 1.6 X 10+ M in chloride or 10-6 M in sulfate; kl(HzOz), 12.5 sec-I colorimetric or 44 sec-l by logplot slope; disappearance rate constant, 4.5 X lo6M-I sec-l at 28” or 3.2 X 106 M-l sec-l at 20’; and the initial concentrations, (Tia+) = 0.005 M and (HzOZ)= 0.05 M . The three major types of data-maximum, kl(HzO2), and disappearance-cannot simultaneously be fitted to the Chien mechanismz6
2BxC
M , “ LI” = kl(HzOz) = Taking (A)O = (Ti)o = 5 X 50 sec-i (cf. 44 sec-l), “kz” = 4.5 X 106 M-l sec-’, the intermediate parameter 7 = (Ao)kz/kl has a value of 45. A solution for the choice is shown in Figure 5 , where a normalized reciprocal concentration is plotted for easier comparison with ref 2. Some qualitative features are similar. The theoretical rate of rise to the maximum is too great relative to the experimental rate. The theoretical concentration at the maximum is much M . I n the approach too high, approximately 6 X of l / c vs. t to second-order behavior, the slope theoretically approaches the final within 50% at Ll(Hz0z)t = 5.7 and within 20% at Ll(HzOz)t = 9.2. The times are long, not absolutely but relative to the time at maximum, and they correspond to very small fractions of and rethe initial titanium remaining, 3 X spectively. If we retain the same (A)oand kl(Hz0z) but reject the experimentally derived kz in favor of more normal values, e.g., 109 M-l sec-’, as found in pulsed radiolysis,la-16 then the concentration at maximum is lower, 1.6 X 10-5 M, the time to reach maximum becomes very short, 1.6 X sec, and the second-order slope is approached within 50% at Icl(H2Oz)t = 10 ( i e . , (20) W. A. Waters, “Vistas in Free-Radical Chemistry,” Pergamon Press Inc., New York, N. Y.,1959,pp 155-160. (21) D.J. Mackinnon and W. A. Waters, J . Chem. SOC.,323 (1963). (22) R. M.Haines and W. A. Waters, ibid., 4256 (1955). (23) W. M. Latimer, “Oxidation Potentials,” Prentice-Hall, Inc., Englewood Cliffs, N. J., 1938. (24) The authors are indebted to a reviewer for suggesting the possibility of reactions 6-10, (25) J. Y. Chien, J . Amer. Chem. Soc., 70, 2256 (1948). (26) 8. Benxon, J . Chem. Phys., 20, 1605 (1962).
3159
THEPARAMAGNETIC SPECIES FROM TITANOUS SALTS AND HYDROGEN PEROXIDE
I
I
I
I
I
I
3I-
u‘
-
II
I
Io I
0
I
I
I
I
t,msec
Figure 6. Comparison of “~OH”-radicalkinetics with Benson’s k scheme,*&Ti3+ HzOt .% =OHand -OH Ti3+2 Ti2+ OH-: (A) theoretical, (A), = (Tia+), = 0.0014 M , K E (.OH),., = 2.75 X 10-6 M from sulfate data, (A)o/K = 500; (B) experimental, Figure 4.
0.2 sec), corresponding to concentrations near 3 X M . I n this latter case, an approximate first-order plot holds within a few per cent up to kl(H202)t = 7. Choices of parameters matching a low maximum concentration will demand a high k2/kl and, therefore, a prolonged steady state, essentially first order and governed by kl. Apparently in Chien’s system a secondorder regime can always be reached in principle, but for a larger 7 the region recedes to impractically long times and low concentrations. The limiting equations are obtained by applying to Chien’s solution the properties of the Bessel and Hankel functions a t small values of the argument.
1 B=
- N
Here E is a small quantity, and
is a parameter used by Chien, and the other symbols have been defined already. I n the Benson meohanism26
AAB A
+ B *C k2
(ki/k2 = K E Bmar)
K is close to 10-6 M , making (A),,/K = 5000. Calculations after Benson’s method show (Figure 6) that the high steady level is 92% attained at kl(H202)t = 0.005 and 98% attained a t kl(H202)t = 0.008, i.e., in a very small fraction of a reagent half-life. With the present kl(H202), the corresponding .times would be immeasurably small. The disagreement would be greater for
+
+
+
K = 1.6 X 10-6 M . More general objections are that the high level should remain forever (not serious, since it actually falls off very gradually compared with the growth rate) and that it should be nearly independent of the initial titanium concentration. A modification of Benson’s scheme involving HOY differs only in the identification of the constants, provided that OH is present in a low steady-state concentration
+ Hz02 -% .OH .OH + H202 -% H02. HO2. + Ti3+ + H + -% H202 Ti3+
(13) (14)
More complex systems can be devised, some of which are tractable. The observed species evidently has a low disappearance rate when alone, but the low maximum level suggests that something present during earlier stages of reaction either destroys the species or competes with its formation. If Ti3+ reacts with undetected .OH as well as observed complexed H02.,the low ultimate level, lowered formation rate, and dependence of level upon the initial (Ti3+), can be accommodated. The difference between the sulfate and chloride results can be explained along two lines. (1) Ionic species in the Tic13 system differ generally in reactivity from those in the Tiz(S04)3 system; to illustrate, perhaps a TiC1(H20)52+ ion reacts more slowly with H202 and more rapidly with .OH than does Ti(S04)(H20)4+. (2) The chloride ion content of TiCl3 acts as a competitive substrate or inhibitor, according to the sequence
R + + *OH + C1-
----zc C1
+ H20
(15)
+
c1 c1Cl2 (16) Reaction 15, reported by Taube and Bray, is discussed in ref 21. Its effect will be somewhat offset by resemblances between OH and C1 reactions. This second Volume 76,Number 9 September 1968
R. E. FLORIN, F. SICILIO,AND L. A. WALL
3160 explanation is in accord with the reported unfavorable effect of added chloride in this system’ and with the pronounced smell of chlorine observed when the reagents are mixed in a beaker. An oxidation-reduction cycle on chloride may play some part
H+
+ *OH+ C1C1 + C1-
Ti3+
.--)
C1
+ HzO
d(oH) = 0 = kl(HzOz)(Ti3+)dt
(OH)
=
(15)
Clz
~ ( H z O (Ti3+) Z) kzo(S’) klg[(Ti3+)0 - (Ti3+)]
+
(16)
+ Clz * Ti4+ + C1 + C1-
(17)
or Ti3+
+ C1+
+ C1-
Ti4+
(18)
It seems unprofitable to attempt an exact kinetic account. Relying on Fischer’s strongly supported suggestions that one kind of “.OH” is Ti-O-Oe3+ and is formed by the reaction 02+
.OH
/I
+ Ti
\0
-% OH-
+ Ti-O-Oq3+
(19)
and allowing for some inhibitor action by such substances as the C1- content of TiC13 as well as by added substrates, the following outline accounts for some features at intermediate substrate concentration Ti3+
+ HzOz -% .OH + OH- + Ti4+ 0 2 +
/I
Ti4+ + HzOz -% Ti
\0
rapid; nearly all Ti4+ present as Ti
\
d(R’) -= dt
OH 0
0; kzo(S’)(OH) = kzl(R’)(Ti-O-0.3+)
The Journal of Physical Chemistry
kzz(Ti3+)(Ti-0-0. 3+) If no reaction with inhibitor or substrate (kzo(S’), kz(S),or k15(C1-)) is allowed, the initial rate of growth becomes k1(HzO2)(Ti3+),which appears too rapid, as discussed under the Benson scheme above. The kzo(S’) term in the denominator permits a lower initial rate of growth. NIoreover, with a reactive substrate in the right intermediate amount, the growth rate can be zero initially but large later because of the increasing term l ~ ~ g [ ( T i-~(Ti3+) + ) ~ 1, dependent upon oxidized titanium. This is consistent with certain growth curves of Turkevich, et ~ l . as , ~well as with the present Figures 2 and 3. The acceleration of growth by the smallest substrate concentrations must still be explained along the lines of Turkevich, et al. ;4 i.e., the reactivity of Ti4+, its peroxide complex, and Ti-0-0 3 + varies with the additional ligands such CHaOH or Sod2-. Numerical values of many of the rate constants above, or of close analogs, have been estimated by various Thus Anbar tabulates rate constants (in M-’ sec-l) for eq 2, 6 X lo*; eq 5, 3 X 108 (assumed the same as for Fez+); eq 12, 6 X 109; eq 13,101; eq 15,2 X lo7 and 4 X lo9 (divergent); and s ~ ~ ~ ~ rate ~ eq 23, 4 X lo6, and other ~ o r k e r estimate
+ HS04-
=
S04.-
+ HzO
(23)
constants for eq 3 of logM-’ sec-l. Reactions 11 and 16 may be similar to reactions 3 and 12, reaction 14 and 18 may he similar to reaction 5, and reaction 19 may be similar to reaction 13. According to these values, competition for OH among stable components of our mixtures would go as in Table IV. Barring interference from reactions 7-10 or unidentified reactions, it seems that at high concentrations the organic substrate should compete successfully (reaction 2). This would involve low (OH) and the possible applicability of the simple reaction scheme outlined in reference to Figure 1. I n the absence of an organic substrate, reaction 15 could dominate if the high value of kls applies. The previously reported low values for the disappearance rate2 are only an order of magnitude lower than (27) S. J. Rand and R. L. Strong, J . Amer. Chem. SOC.,82, 6 (1960). (28) I. A. Taub and L. M. Dorfman, ibid., 84, 4063 (1962).
3161
THESURFACE PROPERTIES OF PHOSPHORIC ACIDIMPREGNATED SILICA Table IV : Competition for OH Reaction no.
2 4 13 15
k,
Reactant
M-1 eec-1
CHaOH Tis+
HzOz C1- high {low
Present concn, M
X 10’ 2.0 X 10sa
