2834
Anal. Chem. 1984,56,2834-2836
are represented by the straight line
In R =
+ 2.505 f 4.0 X
The uncertaintiy term in this equation represents the standard deviation. The activation energy, E,, for the overall reaction is 20 kJ. Due to the complex nature of the reaction rate this temperature dependence may not be valid for concentrations of Cu2+and allyl alcohol other than those employed in this series of experiments.
Registry No. Cu, 7440-50-8; 107-18-6.
14808-79-8;allyl alcohol,
LITERATURE CITED (1) Ullbarri, D., G.; Ogura, T.; Tzeng, J.-W.; Scott, N.; Fernando, Q.Anal. Chem. 1982, 5 4 , 2307. (2) Yanagihara, N.; Ulibarri, D., G.; Ogura, T.; Scott, N.; Fernando, Q. Anal. Chem. 1983, 5 5 , 1873. (3) Funkhouser, J. G. Corrosion (Houston) 1961, 17, 283. (4) Poling, G. W. J . Electrochem. SOC. 1967, 114, 1209. (5) Tedeschi, R. J.; Natali, P. W.; McMahon, H. C. Proc., Conf. Nafl. Assoc. Corros. Eng., 25fh, 1969 1970, 173.
RECEIVED for review May 25,1984. Accepted August 2, 1984.
Kinetic Studies of the Promoting Effect of Glutaraldehyde on the Oxidation of p -Phenylenediamine by Hydrogen Peroxide James C. Thompsen and Horacio A. Mottola* Department of Chemistry, Oklahoma State University, Stillwater, Oklahoma 74078
Glutaraldehyde accelerates the oxldatlon of p-phenylenedlamine by hydrogen peroxide, and this effect has been used for the “catalytlc” deteotlon of aldehyde groups and for the determination of these groups on the surface of modified (by glutaraldehyde) amlnosllane glass to be used for protein immobllizatlon. A reexamination of the effect of glutaraldehyde (25 ‘C, 1.00 M Ionic strength, and pH 5.00) has demonstrated that the unprotonated form of p-phenyienedlamine Is the reactive specles, that gldaraldehyde acts as a promoter and not as a true catalyst, and that once formed the product of the reaction (Bandrowski’s base) is destroyed by further reaction with hydrogen peroxide. An emplrlcal reactlon model of the Oxidation was analyzed by computer, and satlsfactory reproduction of experlmental data suggests three regions In the overall reaction profile: (1) early periods of reactlon In whlch the predomlnant process is the converslon of p phenylenedlamlne to Brandrowskl’s base by a complex formed between glutaraldehyde and hydrogen peroxide, (2) an intermediate region in which the uncatatyzed oxldatlon also becomes a slgnlflcant contrlbutor to the overall rate, and (3) late reaction tlmes when the destruction of Bandrowskl’s base by hydrogen peroxide dominates.
The use of immobilized reagents in analytical chemistry is constantly increasing. Enzymes are commonly used today in immobilized form in many enzymatic procedures, and covalent binding to inert supports offers attractive characteristics for analytical applications, particularly in continuous-flow sample processing ( I ) . The choice of a synthetic route to immobilization depends on the stability of the enzyme for the pH at which coupling is performed, the stability of the linkage for the pH at which the enzyme preparation will be used, and the carrier or inert support. One of the most widely used methods involves alkylamino glass and glutaraldehyde (GA) coupling. Although the detailed chemistry of this reaction is not well understood, the individual steps include the reaction of the silica framework with an aminosilane, modification of the product of this reaction with GA, and finally immobilization of an enzyme. The coupling of the aldehyde to the alkylamino glass is of importance because the amount of protein immo-
bilized depends on the availability of free aldehyde groups. Several reactive pathways give rise to various numbers of terminal, reactive aldehyde groups within the same synthetic route. Quantitative information concerning the number of reactive aldehyde groups is then of capital importance. In answer to this need, Shapilov (2) has proposed a kinetic determination inspired by Feigl’s spot test for aldehydes in solution ( 3 ) based in turn on Woker’s observation that aldehydes accelerate the oxidation of p-phenylenediamine (PD) by hydrogen peroxide (4). The product of oxidation was first described by Bandrowski (5) and is consequently known as Bandrowski’s base (BB). The following overall stoichiometry for the oxidation reaction was proposed by Bandrowski 3CeHgNz + 3H202 + CIgHigNB 6H20 (1) but controversy exists over the exact structure of the resulting base. Attempts to utilize Shapilov’s “catalytic” determination to determine reactive aldehyde groups in open-tube coil reactors and in single-bead-string reactors ( I ) , however, resulted in apparent contradictory behavior in a continuous-flow/stopped-flow system (6). Particularly, the catalytic response was observed to decrease with repetitive determinations using the same coil with attached GA, and the response for the same coil showed a decreasing catalytic activity from day to day. These observations have prompted a reexamination of the overall chemistry involved; the present paper reports on an empirical reaction model (7) describing experimental observations of the effect of GA in homogeneous, aqueous medium, which point to a ”promoting” effect (8)rather than a catalytic one. EXPERIMENTAL SECTION Apparatus. A Beckman Model 25 UV/vis spectrophotometer with a Beckman recorder/controller unit as the readout was used to monitor the progress of the reaction at 485 nm (wavelength of maximum absorption of BB) and 25 “C. The individual reaction profiles were subsequently entered into an Apple 11+ microcomputer for numerical analysis. Adjustment of pH was performed by using an Orion Research (Cambridge, MA) Model 601A digital pH meter equipped with an epoxy-body combination electrode (Sensorex, Westminster, CA). Reagents and Solutions. All chemicals used were of AR grade, The water used for solution preparation was deionized
0003-2700/84/0356-2834$01.50/00 1984 American Chemical Society
+
ANALYTICAL CHEMISTRY, VOL. 56, NO. 14, DECEMBER 1984
and further purified by distillation in an all-borosilicate glass still with a quartz immersion heater. The ionic strength of all solutions was kept at 1.00 M by appropriate additions of NaC104. Phosphate buffer solutions were prepared by mixing appropriate volumes of 0.10 M solutions of NaH2P04and Na2HP04(both in 1.00 M NaC104) until the desired pH was obtained. Solutions of PD were prepared every 3 days. Hydrogen peroxide solutions were prepared from stock 30% HzOz and stored at 4 O C in a refrigerator. A 25% stock solution of GA (Eastman Kodak, white label) was diluted as required for preparation of the solutions of glutaraldehyde. Procedure. Reagents were added in the same order and quantity in repetitive runs in order to maintain reproducible experimental conditions. Manual mixing was employed since it allowed adequate time for simultaneous monitoring of four trials of a given experimental condition in separate cells. Since at least the first 30 min are significant, the 90 s required to mix the four solutions did not cause any significant loss of information. The order of addition was 1.00 mL of buffer solution, 1.00 mL of H202 solution, and 150 p L of GA solution (or the same volume of 1.00 M NaC104for reaction in the absence of GA). This mixture was prepared directly in the spectrophotometric cell and used to set the absorbance value to 0 at time 0. The reaction was initiated by addition of 1.00 mL of the PD solution to each cell. Addition of all reagents was performed with the aid of an automatic pipet (“Quickpette”, Helena Labs., Beaumont, TX). After addition of the PD solution, the cells were closed with Teflon caps and inverted several times to ensure complete mixing. The spectrophotometer was set to plot one point every 2 min for each cell; this allowed 20 s to cycle from one cell to the next. The scale expansion for absorbance depended upon the concentration of the reagents added and varied from 0.100 to 2.000 absorbance units full scale.
RESULTS AND DISCUSSION Initial experiments conducted to ascertain the reaction times involved indicated that the concentration of PD solution to be added should be below 8.75 X M. Higher concentrations resulted in precipitation of BB. In accordance with the available information in the literature, the rate equation was initially assumed to be of the form d[BB]/dt = ku[PD]x[HzO,]”
+ k,[PD]x[GA]2
(2)
where k, is the rate coefficient for the *uncatalyzed* path and k, is the rate coefficient for the “catalyzed” path. Reaction Order with Respect to p-phenylenediamine. Reactions with P D concentrations of 7.00 X M ([H202] = 0.34 M, pH 5.00, and [GA] = 9.51 X lo4 M) were conducted to ascertain the order with respect to the amine reagent. I t was assumed that the uncatalyzed reaction did not contribute to the overall rate when initial rates were calculated since very little, if any, change in absorbance was observed in the absence of GA. The slope of the log-log plots of initial rates vs. initial reagent concentration provided the desired reaction order in accordance with log (d[BB]/dt) = log k,
+ log [GA] + x log [PD],
(3)
where x , the reaction order for PD, was observed to be equal to 1. Effect of Hydrogen Ion Concentration. The rate of the overall reaction is dependent upon pH (2, 3), the reaction proceeding faster a t higher pH. Kinetic data were collected in the presence and in the absence of GA a t pH 4.00, 5.00, 6.00, and 7.00. From a K, value for P D of 7.14 X lo-’ (9), the concentration of the unprotonated form of the reactant was calculated at each pH. Initial rates were found to be directly proportional to the equilibrium concentration of the unprotonated form of PD, indicating it is the reactive form; this was taken into account in subsequent studies. Effect of Glutaraldehyde Concentration. Figure 1 shows the change of absorbance with time obtained a t different GA concentrations. The relationship between initial
2835
l i m e , min Figure 1. Absorbance vs. time curves for the oxidation of W by H,O in the presence and absence of GA. (a) No GA added. (b) 1.9 X 10-B M GA. (c) 3.8 X lo-’ M GA. PD (unprotonated form) = 5.57 X M; H,Oz = 0.34 M; pH 5.00; ionic strength = 1.00 M; 25 OC. Solid llnes correspond to computer-generated curves for the proposed mechanism in the text. For Interpretation of regions 1, 2, and 3, see the text.
rates and [GA] gave an experimental order of 2/3 for the rate dependence on GA. Analysis of the rate profiles in this portion of the study provided the first evidence that the effect of GA was not truly catalytic but of a promoting nature in which the catalytic cycle is interrupted by destruction (or inactivation) of the promoting species (8),GA in this case. Experimental findings did not agree with the expected behavior for a catalyzed reaction of first-order plots, with slopes directly proportional to the “catalyst” concentration, and the absorbance at equilibrium was not the same for different amounts of GA with constant PD. Actually it was noted that after several hours the absorbance began to decrease. Effect of Hydrogen Peroxide Concentration. No changes in the intial portions of the reaction rate curves were observed when the H202concentration was changed. Deviations, however, did appear closer to the absorbance maximum for each curve. This, again, pointed to a promotion effect. In agreement with the postulation of an H202.GAcomplex as the catalytic species (4), the following, fast, preequilibrium step was adopted:
H202
+ GA + [GA*Hz02]
(4)
Since Hz02was present in approximately lo5 molar excess, it is valid to assume that all of the GA in solution was present in the form of the complex [GA.H202]. Formation of this complex would explain why changes in H202concentration do not affect the initial portions of the rate curves. The changes observed at later times, however, are to be interpreted as a result of the reaction of hydrogen peroxide with the product of the main reaction (BB). Destruction of Bandrowski’s Base. To further clarify the effect of GA, experiments were performed under conditions in which the GA concentration was in large excess with respect to PD. Typical absorbance vs. time profiles are shown in Figure 2. Kinetic analysis of the curves showed, as expected, first-order dependence on P D and maximum absorbance directly proportional to the initial concentration of PD. The reaction, in all of these cases, reached completion in a few minutes because of the excess of GA. From the maximum absorbance value in Figure 2, the stoichiometry of eq 1, and the acid-base equilibrium of p-phenylenediamine, a value of approximately lo5 M-’ cm-l can be estimated for the molar absorptivity of BB at 485 nm. The gradual decrease in absorbance with time after reaching the maximum value is indicative of the disappearance of BB from the system. The rate of decrease in absorbance was similar to the case observed with much lower GA concentrations and was not affected by changes in the concentration of GA to any appreciable extent
ANALYTICAL CHEMISTRY, VOL. 56, NO. 14, DECEMBER 1984
2836
Table I. Values of Rate Coefficients for Steps of the Mechanism ReDresented by the EmDirical Model of Promotion Discussed in the Text exptl conditions [PD] = 5.57 X [HZOZ] = 0.34 M [GA] =
kl, M-’ s-’ M
none
Y
I
I
1
I
I
100
50
I
1.9 X 10” M 3.8 X 10” M
I
150
Time, min
Absorbance vs. time curves for the oxidation of PD by H,O, in the presence of a large excess of GA and various concentrations of PD (unprotonated form). (a) 2.32 X lo-’ M PD. (b) 3.48 X lo-’ M PD. (c) 4.64 X lo-’ M PD. GA = 4.76 X lo4 M; H20, = 0.34 M; pH 5.00; ionic strength = 1.00 M; 25 O C . Solid lines correspond to computer-generated curves by use of the mechanism proposed in the text. Flgure 2.
for over a 100-fold change. The experimental order in the “destructive reaction” was determined to be 2 / 3 with respect to BB, as determined from initial rates for the decrease of absorbance. The order with respect to hydrogen peroxide was found to be 1. Numerical Treatment and Fit of Experimental Curves. In agreement with all the experimental observations listed above, the following overall rate equation is proposed d[BB]/dt = kl[PD][H202] k2[PD][GA*H202]2/3- k3[BB]2/3[Hz0,] (5) with the following reaction steps in mind:
+
3PD
+ 3H202 -% BB + 6 H 2 0
(6)
HzOz+ GA + [GA.H202] fast equilibration (7) 3PD
+ [GA.H202] BB
k2
BB
+ 3 H 2 0 + other products
-
+ H202
k3
(8)
products
(9)
Each of these reactions predominates at different times during the course of the process. Initially reaction 8 predominates depleting PD, hydrogen peroxide, and GA (region 1,Figure 1). As the concentration of GA is decreased, the uncatalyzed path (reaction 6) becomes a significant contributor to the overall rate (region 2, Figure 1). As the concentration of BB increases, the reaction of eq 9 becomes predominant (region 3, Figure 1)when the GA has been almost entirely exhausted and the concentration of PD is significantly low. Since all the concentrations in the rate expression are simultaneously changing, with the exception of HzOzwhich can be considered constant since it is in large excess, there is no direct correlation between the reaction curves obtained experimentally and the rate expression of eq 5. To perform a numerical fit, the change in concentration of each species needs to be calculated and substituted in eq 5 to determine the change with time for the concentration of BB, the monitored species. To perform such a task the following individual changes of concentration were computed with the help of the corresponding rate expressions: d[PD] / d t = -3kl[PD] [HZ021 - 3k2[PD] [GA*H202]2/3 (10) d[H,O,]/dt
= -3kl[PD][H202] 3kz[PD] [GA*H202]2/3- k3[BB]2/3[H202] (11) d[GA]/dt = -k2[PD][GA.H202]2/3
(12)
[H,02] = 0.34 M [GA] = 4.76 X [PD] = 4.64 X 10” M 3.48 X 10” M 2.32 X 10” M 1.16 X 10” M
av values
1.02 x 10-2 1.02 x 10-2 1.14 X
162 174
6.00 x 10-3 3.96 x 10-3 4.80 x 10-3
1.20 x 10-2 1.02 x 10-2 1.08 X 1.02 x 10-2
456 486 522 534
3.84 3.84 3.36 3.84
M
1.07 X 10.07 X
389 1173
x x x x
10-3 10-3 10-3
10-3
4.23 x 10-3 10.89 X
A program in BASIC was developed to generate the change in concentration for each species; a listing can be obtained from the senior author. To reconstruct the absorbance vs. time curve, the change in concentration of BB was computed for a given time increment and added to the tabulated value to obtain the concentration at t dt. The experimental and computed values were graphicaly inspected for differences to decide on good initial values for kl, k2, and k3, The sum square differences were used to refine the k’s for the best fit by a successive approximations procedure. Table I lists sample values for the rate coefficients giving the best fit to the experimental data. The larger standard deviation for k2 may result from error in the determination of the initial data points because of the time needed for mixing reactants. Figures 1 and 2 also show the correspondence of computer-generated data with experimental points. It should be mentioned that over 20 other different combinations of possible rate steps for the species involved were subjected to the same fitting procedure without satisfactory results. No mechanistic speculations considering GA as a true catalyst by inclusion of a catalytic cycle provided even a marginal fit. The empirical reaction model proposed above, including a promoting action, the general rate expression (eq 5), and the individual rate expressions (eq 6-9), reproduced the experimental observations in a satisfactory manner. The promoting action, involving GA destruction (eq 8), also provides an explanation for the decrease in signal height in the experiments under continuous flow, since the number of aldehyde groups available for reaction decreases as the reaction is repeated with the same coil. Registry No. H20z, 7722-84-1; GA, 111-30-8; PD, 106-50-3.
+
LITERATURE CITED (1) Mottola, H. A. Anal. Chlm. Acta 1983, 145. 27-39. (2) Shapilov, 0. D. J. Anal. Chem. USSR (Eflgl. Trans/.) 1980, 35, 1429-143:. (3) Feigl, F. Spot Tests in Organlc Analysis”, 7th ed.; Elsevier: New York, 1966; pp 198-199. (4) Woker, Q , Ber. Dbch. Chem. G e s . 1914, 4 7 , 1024-1029. ( 5 ) Bandrowskl, E. Ber. Dtsch. Cbem. G e s . 1894, 27, 480-486. (6) Gnanasekaran, R., Oklahoma State University, Stillwater, unpublished results. (7) Come, G. M. Compr. Chem. Klflet. 1983, 2 4 , 252. (8) Eswara Dutt, V. V. S.; Mottola, H. A. Anal. Cbem. 1974, 4 6 . 1090-1094. (9) Morrison, R. T.; Boyd, R. N. “Organic Chemistry”, 3rd ed.; Allyn and Bacon: Boston, MA, 1973; p 749.
RECEIVED for review May 8,1984. Accepted August 3,1984. This work was supported by a grant from the National Science Foundation (CHE-8312494).
