CHARLESR. MAXWELL AND DOROTHY C. PETERSON
5110
Acknowledgment,-Financial support from the U. S. Public Health Service and from the Louis and Maud Hill Family Foundation is gratefully acknowledged. Preliminary work on the de-
VOl. 79
velopment of some of the procedures has been carried Out by Dr. R. K. Neogy. We are indebted to Dr. R. W. Lumry and Dr. B. R. Willeford for their interest in our work. MINNEAPOLIS,
MI”.
RADIATION BRANCHNATIONaL CANCER INSTITUTE, NATIONAL INSTITUTES HEALTH SERVICE, u. s. DEPARTMENT OF HEALTH, EDUCATION AND WELFARE]
[CONTRIBUTION FROM THE
OF
HEALTH,PUBLIC
A Mechanism for the Reaction of Fe++, H,O, and Glycine in an Aerobic Aqueous Solution BY CHARLES R. MAXWELL AND DOROTHY C. PETERSON RECEIVED APRIL 11, 1957 The kinetics and stoichiometry of the action of Fenton’s reagent in aerobic aqueous solutions of glycine are presented and a mechanism is proposed to describe the observations. The reaction is initiated and the rate controlled by the produc~ The stoichiometry, which varies as the reaction proceeds, is the result of tion of an OH free radical from a H z O molecule. two pairs of competing steps: the first between F e + + and glycine for the OH free radical and the second between F e + + and F e + + +for an intermediate organo-peroxy free radical. The Fe++-peroxy reaction regenerates the H20z used in the initial step and thus accounts for the small AHz02/AFe++ratios observed in the early phase of the reaction. The Fe+++-peroxy reaction regenerates the F e + + used in the initial step and thus accounts for the large AH2OZ/AFe++ratios observed in the late phase of the reaction. HCOCOOH and NHs are formed by both reactions. The effects of H + concentration and of glycine concentration have been measured and are discussed, The ratios of the reaction rate constants of the competing steps have been evaluated for various conditions.
Introduction tions which oxidized the HCOCOOH and HCHO A mixture of H2Oa and Fe++, commonly re- but did not affect the NHs. Merz and Waters7 have proposed a general ferred t o as Fenton’s reagent, has been used since before the turn of the century to study the oxida- scheme for the action of OH radicals from Fenton’s tion of organic compounds in aqueous solution. reagent upon organic substrates in which many More recently, as the mechanism of the reaction be- compounds are oxidized by a chain reaction involvgan to be more fully ~ n d e r s t o o d l -it~ has been used ing the substrate and H202. Glycine and alanine extensively to study the effect of OH free radicals. are listed among these compounds. The rationale We became interested in the action of this reagent for this chain, which is initiated by an OH radical upon glycine as a result of a study4 of the effect of formed from Fe++ and HzOzand terminated by the ionizing radiation upon solutions of this amino acid. reaction of an OH radical with a F e f + ion, appears Since ionizing radiation in aqueous solution reacts to be based solely upon large HzOz consumption/ indirectly5 through a complex mixture of H and OH Fe++ oxidation ratios. No quantitative measurefree radicals and HzOz, it was believed that an in- ments were reported for the yield of products or vestigation of the Fenton reaction would provide consumption of substrate nor of the rate of the rea mechanism which would be useful in the inter- action. pretation of radiation data. Kolthoff and Medaliaa have suggested another Dakina has reported that the products of the scheme for the action of Fenton’s reagent upon oroxidation of amino acids by Fenton’s reagent are in ganic substrates. In the presence of dissolved oxygeneral ammonia, a keto acid, an aldehyde and gen (which Merz and Waters did not consider) this carbon dioxide with the yield of the aldehyde scheme includes a different chain reaction which is usually greater than the yield of the keto acid. also initiated by the action of an OH radical upon However, with glycine he found the yield of HCO- the substrate. This chain which involves oxygen COOH greater than the yield of HCHO. Merz and (not HzOz),substrate and F e f f provides for the Waters,7on the other hand, have listed only HCHO oxidation of many molecules of substrate and F e + + as the product of the oxidation. The a ~ t h o r s , ~ions for each OH radical formed ( H ~ Oconsumed). Z working with ionizing radiation rather than Fen- However, here again the mechanism appears to be based solely upon HzOz consumption/Fe++ oxidaton’s reagent, found HCOCOOH, HCHO, “3, COz, and in addition observed that the ratios be- tion ratios. No data were reported for substrate tween the yields of HCOCOOH, HCHO and NHs oxidation, product yield or rates of reaction. varied with the dose of radiation. These observaFrom the above it is apparent that data as to the tions were interpreted in terms of secondary reac- extent of secondary reactions, the actual substrate oxidation or product yield, and the rate of the reac(1) Iz. Haber and J. Weiss, Proc. R o y . SOC.( L o n d o n ) , 8 1 4 7 , 332 (1934) tion is needed. We, therefore, first studied the (2) A. I. Medalia and I. 51. Kolthoff, J . Polymer Sci., 4, 377 (1949). yields of products as a function of “OH exposure” (3) W. G. Barb, J. H. Baxendale, P. George and K. R. Hargrave, and differentiated the initial and secondary prodTrans. Faraday Soc., 47,462, 591 (1951). ucts. The study of the reaction was then limited (4) C. R. Maxwell, D. C. Peterson and S . E. Sharpless, R a d . Reto conditions which eliminated or minimized the s e n r c h , 1, 530 (1954). ( 5 ) A. 0.Allen, ibid., 1, 85 (1954). (6) H. D.Dakin, J . B i d . Chcm., 1, 171 (1905). (7) J. H. Merz and W. A. Waters, J . Chem. Soc., Sl5, 2427 (1949).
(8) I. R I . Kolthoff and A. I Medalia, THISJ O U R N A L , 71, 3777, 378% (1949).
Oct. 5, 1957
REACTION OF FE++, HzO AND GLYCINEIN AEROBICAQUEOUSSOLUTION
secondary reactions. I n this study the concentrations of Fe++, H202 and HCOCOOH (a product of the OH-glycine reaction) were each followed as a function of time. The AHZOz/AFe++ ratios were observed to vary widely and in a range and manner similar to the observations of other investigators for other compounds. The observed kinetics is, however, incompatible with either of the general schemes heretofore proposed and a new mechanism is presented to describe our observations. Although tested only for glycine, this mechanism suggests an alternate general scheme for the explanation of the stoichiometric observations of the other investigators. Experimental Procedure.-Two beakers of glycine, one containing H202 and the other Fe++, were rapidly poured together and aliquots for analysis removed by pipet. The reaction was stopped by blowing each aliquot into a solution of either o-phenanthroline or tetravalent titanium. The o-phenanthroline stops the reaction by complexing the ferrous ion; the titanium stops the reaction by complexing the H202.g In some cases where the NHI yield was to be measured the reaction was stopped in the late slow stages of the reaction by the addition of catalase which destroyed the H 2 0 2without effect upon the substrate. I n general th: reaction was followed until 75-90y0 of the Fe++ was OXIdized with time intervals between pairs of aliquots in the early rapid phase of the reaction of the order of 15 seconds. pH.-The PH of the reaction mixture was adjusted by the addition of 1 N H2S04 to the glycine solution and was measured with a Beckman p H meter and glass electrode. The kinetic study was limited to pH levels between 2.8 and 4.6. A t levels greater than 4.6 the reaction was too fast to follow with accuracy and a t levels less than 2.8 the reaction was not stopped immediately by addition to o-phenanthroline. Temperature.-When the effect of temperature upon the rate of reaction was being studied the reaction mixture was immersed in a regulated water-bath, otherwise the reaction was carried out on the laboratory bench a t room temperatures between 22 and 25’. Analytical.-The F e + + and Hz02 concentrations were determined by the method described e1sewhere.O This method restricted the investigation to mixtures in which the HnOaconcentration was greater than one-half the F e + + concentration and, with no buffer in the o-phenanthroline reagent, to pH levels above 2.8. Ammonia was determined colorimetrically with Nessler reagent after an overnight microdiffusion in a Conway dish into dilute H2S04 from an aliquot of sample made basic Z the sample was destroyed with with NaOH. Any H ~ O in catalase prior to the addition of NaOH in order to prevent a “dark” reaction which produced additional ammonia. HCOCOOH was determined colorimetrically by a modification of the Bamburger testlo for keto-acids, using ethyl acetate as the organic solvent and filtering the final basic solution to remove turbidity caused by traces of T i + + + + carried (mechanically?) through the procedure. The procedure was standardized with twice recrystallized SaOOCCHO . H 2 0 prepared by the periodic oxidation of tartaric acid according to a method described by Weissbach and Sprinson.” HCHO was determined colorimetrically with chromotropic acid using a modification of the method described by Kleinert and Srepel.12 Appropriate knowns and controls were measured simultaneously in order to minimize the interference of large amounts of F e + + +and HCOCOOH. HCOOH was also determined colorimetrically with chromotropic acid, after reduction t o HCHO with Mg and (9) C. R. hIaxwell and D. C. Peterson, t o be published. (10) H. J. Fister, “hlanual of Standardized Procedures for Spectrophotometric Chemistry,” Standard Specific Supply Corp., New York, N. Y.,1950. (11) A. Weissbach and D. B. Sprinson, J . Biol. Chem., 203, 1023
,(1953). (12) T. Kleinert and E. Srepel, Mikvochcm. Mikvochim. Acta, 93,328 i1948).
5111
HCl by a modification of the method described by Grant.18 Prior to reduction, the HCOOH was removed from the reaction mixture by steam distillation and concentrated as the sodium salt. Reagents. Water.-All water was distilled from a tin lined still, redistilled in Pyrex from basic KMnO4 and again from dilute H2S04and stored for short periods in Pyrex bottles. Hz02.-99.9Yo inhibitor-free HZOZwas used in all experiments reported here. Merck 30y0 Superoxal gave identical results in other experiments not reported. Fe++.-Both reagent grade FeSO4 and reagent grade Fe( NH4)2(S04)2.6H20were used without further purification. Fe +++: reagent grade FeNHa(,S04)2~12H20. Glycine .-Glycine from Pu’utritionalBiochemicals Corporation was recrystallized twice; once from distilled water by the addition of C.P. absolute methanol followed by vacuum drying a t 50’ and again thermally from distilled water with a second vacuum drying.
Results The Reaction.-Ammonia, glyoxalic acid, formic acid and formaldehyde have been found in glycine solutions treated with Fenton reagent where the initial Fe++ and HzOz concentrations were of the order of to M . However, the ratios among the Fe++ oxidation, HzOz consumption and the yields of the various products varied with the initial concentrations of Fe++ and H20Z and with the time the reaction was allowed to proceed. Pertinent portions of the data are shown in Table I. The lack of a constant ratio among the products has been interpreted as the result of secondary reactions which oxidize some of the primary products. TABLE I STOICHIOMETRY OF FENTON’S REACTION IN 0.1 M GLYCISE-~H 3.&AIR SATURATED Reactants,
p M
Feat+
H201~ A F e + + AH202
HCOCOOH
450 920 1800 120 480 480 1920 1920 132 200 330 400
1030 440 190 1810 910 760 5700 1790 3410 320 1280 1280 5120 5120 113 45 83 104 176 68 167 300 130 210 360 170
300 480 490 145 260 305 560 500 88 124 197 236
Products,
p M
HC- Time. N H I HCHO OOH mln.
500 940 1980
.. .. .. 0 =2 =3
=5 =4 93 135 231 292
19 60 110
5 7 12 10 2 10 2 10 10 10 12 12
In order to determine the initial ratios before secondary reactions or oxygen depletion became important the data were analyzed on the assumption that ammonia is a stable product which is a measure of the glycine consumption. A plot of the HCOCOOH/NH3 ratio against the NH3 yield for the last four lines of data in Table I is a straight line which extrapolates t o a ratio of one a t zero NHa yield. This indicates that initially NHI and HCOCOOH are formed in equal amounts but that a secondary reaction or oxygen depletion has distorted the ratio by the time the NH3 yield is large enough to be measured accurately. The fact that this distortion is due to a secondary reaction rather than oxygen depletion has been clearly demonstrated by measuring the ratio in an oxygen-free (13) W. M. Grant, Anal. Chcm., 20, 267 (1948).
CHARLES R. MAXWELL AND DOROTHY c. PETERSON
5112
system. The data, which will be presented in a subsequent paper, yields HCOCOOH/NH1 ratios which also extrapolate to one a t zero NH3 yield and which incidentally are much less distorted by secondary reactions than the ratios in air. A similar examination of the HCOOH yields indicates that HCOOH is a secondary product. It is almost certainly the result of the oxidation of HCOCOOH. The low yield of HCHO has prevented a detailed study of this product and its role in the initial reaction has been assumed to be insignificant. I n order to minimize the distortion resulting from secondary reactions the investigation of the kinetics of the glycine oxidation was made using initial Fe++ and H2Oz concentrations of less than M which produced final Concentrations of HCOCOOH of less than 50 p M . Under these conditions only Fe++, HzOa and HCOCOOH could be measured with precision. The analysis of data from many experiments in this region has led to the following general observations. (1) Throughout the course of any experiment the rate of glyoxalic acid production is proportional to the product of the ferrous and hydrogen peroxide concentrations. (2) The initial rate of Fe++ oxidation is equal to twice the initial rate of the HCOCOOH production. (3) The rate of Fe++ oxidation, AFe++/[Fe++] [H202]At,approaches zero as the reaction proceeds (4) The initial rate of HzOz consumption, A[H20z]/[Fe++][H20z]At, is small but the rate increases as the reaction proceeds and approaches the rate of HCOCOOH production. ( 5 ) The presence of F e + + + reduces the initial rate of Fe++ oxidation. ( 6 ) The presence of both F- and F e + + +produces no change in the znitzal rate of Fe++ oxidation. (7) The removal of dissolved oxygen increases the initial rate of HzOz consumption. The Proposed Mechanism.-Equations 1 through 6 are proposed as a mechanism for the action of Fenton's reagent on glycine as describecl above. Equation G is not a truly mechanistic equation but an over-all equation which for convenience is used in lieu of eq. 11 and 12 developed later in the section on the effect of H + concentration. The species reacting with the F e + + + is believed to be the anion of the ionized peroxy radical rather than the peroxy radical itself. However, in experiments where the H + concentration is essentially constant the concentration of the anion will be a constant fraction of the concentration of the radical and it is equivalent and more expedient mathematically to work with the effective equation as presented. Fe++
+ HZOZ F e + + + + OH- + OH OH + Fe++---t F e + + + + OH----t
+ KHzCHzCOOH + HzO + NHzCHCOOH NH~CHCOOH+ --+ NH2CH(Oz)COOH NHlCH(02)COOH + F e + + --+ F e f f + + (NHZCH(O~)COOH)-
k1
(1)
ki
(2)
ks
(3)
k4
(4)
ks
(5)
VOl. 79
(NHZCH(OZ)COOH)-+ HN=CHCOOH HOzHOzH + +Hi02 HN=CHCOOH HpO --f NHa HCOCOOH
+
+
+
+
+
+ HzO + + NHs + H +
I~HsCH(OZ)COOH F e + + + F e + + 4- O2 HCOCOOH
+
(5a)
(5b) (5~)
ke (6)
The reaction is initiated and the rate controlled by reaction 1 which produces an OH free radical. Fe++ and glycine then compete for the OH free radical according to eq. 2 and 3 . However, as is shown later, a t [Glycine)/[Fe++] > 2000 reaction 2 occurs less than 5y0of the time and may be neglected in a gross analysis. Reaction 3 produces a glycine free radical which reacts in the presence of dissolved oxygen by 4 to form a peroxy free radical. The reaction is then terminated by either 5 or 6 where Fe++ and F e + + + compete for the peroxy radical. The net reactim measured a t any moment is thus the sum of two reaction sequences 1, 3 , 4 and 5 and 1, 3 , 4 and ti whose relative contribution to the sum changes from moment to moment as the Fe ++/Iie+++ratio changes. The first sequence, summed as A, describes the initial phase of the reaction when the 2Fe+' Hi02
+ 02 f
+
S H L H L C O O Hf 2 H + --+ 2 F e + T + HCOCOOH XHS HzO (A) SH,CHLCOOH---+ HCOCOOH f XHg €120 (B)
+
+ +
F e + + + concentration is small. The second sequence, summed as B, describes the late phases of the reaction when the F e + + concentration is small. Thus, on the basis of this model there should be no simple constant ratio between any two of the measured quantities, Fe++ oxidation, H202 consumption, and glyoxalic production but there should he a simple relation between the three. The glyoxalic acid production should be equal to the sum of the HIOl consumption and one-half the Fe++ oxidation. Figure 1 shows how the observed data from two experiments carried out under widely different conditions satisfy this prediction. Inherent in this model are other more complex relationships between Fe++, H202 and HCOCOOH concentrations and time which provide methods for the evaluation of kl, the ratio ke/ks and the ratio kz/k3. These relationships provide additional methods for testing the model against the observed data. The development of such relationships and the testing of them by the observed data follow. Evaluation of ks/k-.--hn equation describing the H202 concentration as a function of the F e + + concentration, the initial concentrations of the reactants and the ratio ks/k5 for the reaction rate constants of 5 and G may be derived in the following manner: (I) follows from the fact that the rate of Fe++ oxidation by A is twice the rate of 5 , and (11) follows from the fact that the rate of H202 consumption by B is equal to the rate of 6. \Then the F e + + +
OH
0 2
concentration is expressed as a function of the F e + + concentration and I is divided by I1 to re-
REACTIONOF FE++, HzO AND GLYCINEIN AEROBICAQUEOUSSOLUTION
Oct. 5, 1957 50
[[Fe+++]" I I= 300,Ml I
200
I
I
I
0
40
A
s; I E, 30 c
I
5113
[HrOzlO = 7 8 5 " *
3
150
I
/
T
&
x'
T 100
0 0 20 u 0
II N
0
50
-1
10 0 Y I 0
I
I
I
I
I
I
I
I
I
I
1
, I
10 20 30 40 50 60 I/*AFe++ AH202, fimoles/l. Fig. 1.-The HCOCOOH yield as a function of the F e + + oxidation and H*0* consumption. See equations A and
0
+
B.
move the unknown peroxy radical concentration and the time factor dt, a differential equation is obtained which, when integrated yields I11 or the equivalent form 111'. A plot of the HzOg consumption against Z should yield a straight line with slope of k0/2k5. Figure 2 shows such a plot for the data previously used in Fig. 1. The difference in the ka/ ks ratio for the two sets of conditions is, as will be
evaluation of the ratio of the reaction rate constants ka and ks by equation 111.
Fig. 2.-The
The differential equation X follows mathematically (2B
I n principle this can be integrated between limits to yield X I and kl evaluated as the slope of the plot of At vs. A I . In practice we were unable to ki
(to
- t ) = Io - I
(XI)
find an analytical expression for I and had to resort to numerical integrations. Plots of such values of A I vs. the observed time are shown in Fig. 3 for data from the two experiments previously shown in Figs. 1 and 2 . The difference in the values of kl is believed to be caused primarily by the difference in
PH. I
l
/
-*
l
4
(IV) +
I
/
/
I
0
(VI)
1
k'ip6 6 = Ips -
(VI11 (VIII)
1= kl [Fe++I [H2021
(1x1
(14) Unpublished data. Measurements of the F e + + formation a n d HCOCOOH depletion as a function of time in both water and 0.1 A4 glycine brought to a pH of 3.5 with HzSOh have shown the following reaction sequence to proceed in the absence of HtOx.
+ +
= [Fe++]
x
+ [Fe++]o
HCOCOOH. Evaluation of kl.-With I11 and a value for the ratio k6/k6 i t is now possible to derive an expression for the Fe++ concentration as a function of time and to evaluate k1. With an underlined integer defined as the rate of the corresponding reaction and $6 and p6 defined as the probabilities of any particular peroxy radical reacting by either 5 or 6, formulae IV + IX follow from the proposed mechanism.
pS =
(X)
dx x2
A = [Fe++]o B = K"021a C = [Fe+++]o K = k6/ka
shown later, primarily due to the difference in H + concentration. The deviations from linearity observed in the late stages of the reaction in the presence of a large Fe+++ concentration are almost surely the result of a secondary "dark" reaction14 which we have observed between Fe++f and
+
+ KA
where
I
PS
- K)x - Kx - K C l n 4)
K C + (1 k1 dt
- [H2021= & J. ([Fe+++Io + . [ Fe + +lo [Fef+lo) In [~e+il - ([Fe++10- [Fe++])
-
30
AH~OS, pmoles/l.
[HIOllo
-d [Fe++] = l- + 5_- 6 dt & = kg [Fe++] pe ke [Fe+++]
20
10
0
F e + + + HCOCOOH +F e + + F e + + + X +F e + + X
+
+X
(a) (b)
T h e kinetics of the initiating reaction (a) can be described by a second order rate constant of 0.1 liter mole-' set.-' a t a pH of 3.6.
0
200
400
600
800
1000
Time in seconds. Fig. 3.-The evaluation of the reaction rate constant R1 by a numerical integration of equation X. One group of data has been displaced by 100 seconds along the abscissa.
CHARLES R. MAXWELL AND DOROTHY C. PETERSON
5114
A comparison between the observed concentrations of Fe++ and H202 as a function of time and curves calculated using the values of kl and ks/ks from Figs. 2 and 3 are shown in Figs. 4 and 5. The agreement is considered good in view of the “dark” reaction between F e + f + and HCOCOOH.
,OKr 26 22 18
1
?
-
r
i
i
I
CURVES-CALCULATED POINTS - D B I E R V E D
’
T
2 14 ..-
;12
TABLE I1 THEEVALUATION OF k l / k 3 pH 3.5; [H2O2IO= 60 pM; [Fe++Ia = 60
8
IGlycine], A4
iI
4’
0 1
1
1
,
500 1000 1500 Time in seconds. Fig. 4.-Calculated and observed F e + + concentrations as a function of time. The curves were calculated by a numerical integration of eq. X using the constants evaluated in Figs. 2 and 3. 0
CURVES
- CALCULATED
i
POINTS-OBSERVED
a
[Fe++l~ra,HCOCOOH, P #moles
u
1
0
_,
OH, #moles
p M
kr/kP
0.010 35 15 3 19 0 48 ,003 35 9.0 20.0 73 002 34 6 26 18 9 81 001 31 3.19 I8 3 95 0005 30 1 97 18 8 85 From XI1 using an “average” [Fe++]of 50 p M .
The Influence of Glycine Concentration-The mechanism as proposed in 1 through 6 provides no explanation of the large influence of both glycine concentration and H + concentration upon values for both kl and the ratio kslks and is thus obviously a simplification of a more complex mechanism. Figure 6 portrays the data used in the evaluation of 61 a t various glycine concentrations a t constant PH. The linearity of the data for each of these concentrations indicated the validity of the general mechanism over this concentration range. Column 2 in Table I11 lists the values in Fig. A after a correction for small temperature differences.
__
I
79
was evaluated by a numerical integration of the rate equation for 1 using a value of 50 1. mole-1 set.-' for kl and the measured values of Fe++ and HZOZconcentrations. The results are shown in Table 11. The average value of 80 for kz/k3 is of the same order of magnitude as the 30 reported by lferz and ’Cl’aters.’ A potentially more accurate method of evaluating k2/k3 by complexing the ferric ion with F- so as to suppress 6 and using the Hz02 consumption as a measure of 2 was unsuccessful. The F- complexed the T i + + + + to such an extent that the remaining T i + + + + concentration was too small to complex all of the HzOz and stop 1.
4 10
+‘ +
1701.
7
,
,
’
-I
200
400 600 800 1000 1200 1400 1600 Time in seconds. Fig. 5.-Calculated and observed H202 concentrations as a function of time. The curves were calculated by a numerical integration of eq. X, using the constants evaluated in Figs. 2 and 3 and eq. 111.
Evaluation of k2/k3.--Since HCOCOOH is produced by both 5 and 6 the yield of this product is a measure of 3. There is no direct measure of 2 but i t may be evaluated as the difference betureen the OH production by 1 and the HCOCOOH production. Thus, X I 1 follows from the mechanism kz = (OH-HCOCOOH k3
HCOCOOH
[Glycine]
) ( [Fe++]’)
+ +
&
L
y/
