101
Chemiluminescence from Hypochlorite and Luminol Reaction (16) F. W. Daiby. Can. J. Phys., 36, 1336 (1958); J. L. Bancroft, J. M. Hoilas, and D. A. Ramsay, 1 6 4 40,322 (1962). (17) E. A. Guggenheirn,-Mo/.Phys., 10, 401 (1966); 11, 403 (1966). (18) R. L. Scott, Mol. Phys., 11, 399 (1966). (19) In the previous papers we gave smaller radiation cross sections. The orlgins for the wrong estimations came from the value assumed for Snz: d7. Namely, 10 cm3 assumed In the previous papers for the reaction volume was too large In comparison wlth the exact value calculated in the present work. For the H 4- N202 system, we moreover made a mistake In a simple calculatlon.
(20) (21) (22) (23)
The increasing of the background light on beam-on was consideredIn the present estimations. The detail treatment will be described in the Doctoral thesis (Kyoto University) of Ibaraki. M. Bodenstein, Heiv. Chim. Acta, 18, $45 (19351. L. D'or, A. delattre, and P. Tarte, J. Chem. Phys., 19, 1064 (1951). 0. K. Rice, J. Chem. Phys., 4,367 (1936). D. Neuberger and A. B. F. Duncun, J. Chem. Phys., 22, 1963 (1954); K. Sakurai and G. Capelle, IbM., 53,3764 (1970); S. E. Schwartz and H. S. Johnson, IbM., 51, 1286 (1969); L. F. Keyser, S. 2. Levine, and F. Kaufman, l6M.. 54, 355 (1971).
Chemiluminescence from the Reaction between Hypochlorite and Luminol W. Rudolf Seilz Department of Chemlstry, University of Georgis, Athens, &org/a 30602 (ReceivedJune
IO, 1974)
The oxidation of luminol by hypochlorite to generate chemiluminescence has been studied in aqueous base using a flow system to mix the reactants. In the presence of excess of luminol, the variation of steady-state chemiluminescence intensity as a function of hypochlorite concentration can be resolved into first- and second-order components. The first-order component requires oxygen while the second-order component does not. These data are explained in terms of an azaquinone intermediate formed when hypochlorite oxidizes luminol. The azaquinone can react with oxygen or hypochlorite to generate the first- and secondorder processes, respectively. The rate constant for the reaction between hypochlorite and luminol decreases as the pH goes from 9.2 to 11.8. The chemiluminescence efficiency for the reaction between azaquinone and hypochlorite is greater than for the reaction between azaquinone and oxygen.
Introduction The oxidation of luminol (5-amino-2,3-dihydrophthalazine-l,4-dione) is accompanied by intehe blue chemiluminescence (CL): "2
0
"2
Although this reaction has been the subject of many studied there is still no agreement as to what intermediates are formed when the reaction is carried out in aqueous media. Both the azaquinone formed by two-electron oxidation of luminol and the radical formed by one-electron oxidation of luminol have been proposed as possible intermediated-6 and strong evidence has been presented for each intermediate. It is possible that both intermediates can occur depending on what oxidizing agent is used to oxidize luminol. This paper reports a study of chemiluminescence from the reaction between hypochlorite and luminol. Since hypochlorite oxidizes by two electrons, the azaquinone inter-
mediate seems most likely for this reaction. Azaquinones have been previously proposed as intermediates inthe oxidation of luminol by persulfate3 and iodine.2 Also, azaquinones have been shown to chemiluminesce upon oxidation bv hvdroeen ~eroxide'aDroducine: the same emission spectium as observed for the CL of ihe corresponding hydrazide. The study reported here shows that the properties of the OC1--1uminol reaction are consistent with an azaquinone intermediate. The azaquinone intermediate can react either with oxygen or with hypochlorite to generate CL.
Experimental Section Apparatus. The apparatus used in this study has been described in detail elsewhere.9 Luminol in 0.1 M KOHH3B03 buffer is continuously mixed with water in a cell positioned in front of a photomultiplier that measures CL intensity. The flow system is driven by an infusion pump. A sampling valve is used to insert 2-3-ml slugs of OC1- solution into the water line. As the OC1- passes through the cell, steady-state CL is measured and recorded. The waterhypochlorite and buffered luminol are mixed by the bubbling oxygen through the cell at 80-100 cc/min. This also ensures that an excess of dissolved oxygen is available to react. Procedures and Solutions. A 0.1 M OC1- standard solution was prepared from Ca(OC1)Z. This solution was kept in an opaque brown plastic bottle and stored in a refrigerator to minimize OC1- decomposition. Dilute standards ( loB3,and M ) were prepared fresh daily in volumetric The Journal of Physica, Chemistry, Vol. 79, No. 2, 1975
W. Rudolf Seitz
102
flasks covered with aluminum foil to keep out light. Grunbaum pipets (50- and 1OO-pl) were used'to make standard additions to a 500-ml polyethylene sample bottle from which sample slugs were withdrawn. The sample bottle was conditioned in hypochlorite solutions and covered with black plastic to keep out light. All additions were made immediately before running the sample. A stock solution of 1M KOH-HSB03 and 4 X low3M luminol was prepared using sodium luminol purified by recrystallization from base. Solutions of 0.1 M buffer and 4 x M luminol were prepared by dilution from this stock solution. Either solid H3B03 or KOH was added to adjust the pH. Kinetic Measurements. The kinetics of the OC1-luminol reaction were measured from the rate at which steady state is achieved in the flow system. Steady-state CL is observed when the rate at which hypochlorite enters the cell is equal to the rate at which hypochlorite reacts plus the rate at which hypochlorite leaves the cell. If the flow rate through the cell is slow enough, virtually all the hypochlorite reacts before leaving the cell. In this case the rate at which steady state is achieved depends on the kinetics of the reaction. For a first-order reaction C L ( ~ )= CL,,(1
- rkt)
(1)
where CL(t) equals chemiluminescence as a function of time, CL,, equals steady-state chemiluminescence, and k is the rate constant for the first-order reaction. In practice, the CL from a slug of hypochlorite passing through the cell was recorded on an expanded time scale. The difference between steady-state CL and observed CL was measured as a function of time after the slug of sample initially entered the cell. This is illustrated in Figure 1 which shows a CL signal on an expanded time scale. The log of the difference between observed and steady-state CL was plotted us. time and the rate constant was obtained from the slope. The kinetics were measured at slow flow rates, 0.08 to 0.3 ml min-l syringe-l for all but the fastest reactions. The smaller rate constants were measured at more than one flow rate to confirm that the experimental rate constant was independent of flow rate. If the rate constant is independent of flow rate, this indicates that the hypochlorite is reacting before leaving the cell. Air bubbles separated the slug of hypochlorite from the background water so that no mixing occurred before the hypochlorite entered the CL cell. The rate constants were measured at room temperature which was 2 1 O . No device for temperature control was used. Results Effect of Hypochlorite Concentration on Chemiluminescence Intensity. Figure 2 shows steady-state chemiluminescence as a function of hypochlorite concentration at four different pH's. To get the data in Figure 2, five-point data sets were run over the concentration ranges 1-5 X M OC1-, M OC1-, 4-20 x 10-8 M OC1-, 2-10 X 1-5 X lop7 M OC1-, etc. The lower concentration ranges were run a t least twice. The individual curves were combined to give the data in Figure 2. This was done by arbitrarily assigning a value of 1.00 to the steady-state CL for 1 x IOqs M OCl-. The steady-state intensity assigned to 2 X M OC1- was then determined from the experimental ratios of steady-state intensities for 1 and 2 X M OC1-. The value for 2 X 10-8 M was then used as a stanThe Journal of Physical Chemistry, Vol. 79, No. 2, 1975
I
0
f tZ0 TIME
-
Figure 1. Dlagram Illustrating method for determining kinetics of the OCI--1umlnol reaction. The time scale of the CL peak is expanded and the difference between steady-state CL and observed CL is measured as a function of time after the slug of OCI- has entered the CL cell.
dard for assigning numerical values to the steady-state intensities observed for higher OC1- concentrations. This procedure was used all the way to the highest OC1- concentration. When the concentrations being ratioed appeared in more than one of the five-point data sets, the average ratio was taken. This method smooths out deviation in individual, short-range data sets and reduces errors from OC1- decomposition and changing activity of the luminol solution. After the CL US. [OCL-] data were analyzed, the arbitrary units were adjusted so that the relative units for CL intensity are the same at all pH's. The CL intensity us. [OCl-] data were resolved into first- and second-order components using a computer programlo to calculate the leastsquares values of the coefficients b 1 and b 2 in the equation: steady state CL = bl[OC1']
+
b2[OC1'I2 (2) The least-squares values of b 1 and b2 were first calculated for the lowest five concentrations of hypochlorite. The calculations were then repeated including the next higher hypochlorite concentration for each calculation, until the variance of the fit started increasing indicating that deviations were occurring. The least-squares fits shown in Figure 2 are based on the maximum number of points before deviations started to occur. Figure 2 shows the calculated CL us. [OCl-] data based on the least-squares values of bz and b 2. The figure indicates which data points were used in the calculation. At all four pH's, the calculated fit closely matches the experimental data for low hypochlorite concentrations. At the three lower pH's, the experimentally measured steady-state CL is less than that calculated from the least-squares fit for higher hypochlorite concentrations. The extent of deviation decreases with increasing pH. At the highest pH, 11.7, the experimental data do not substantially deviate from the calculated least-squares fit up to 2 X 10-5 M oci-. At pH 11.7, the data quality is poor. Steady state is established slowly. Steady-state CL is not very reproducible, showing a tendency to increase when the same concentration of hypochlorite is run through the cell more than once. The reason for this is not known. Effect of Oxygen. The first-order process requires oxygen while the second-order process does not. This was demonstrated by measuring steady-state CL us. [OCl-] in the presence and absence of oxygen. Oxygen was removed by bubbling nitrogen through the solutions and using nitrogen rather than oxygen to stir the cell. When oxygen is removed steady-state CL is reduced by an amount linearly proportional to [OCl-1. Data for pH 9.4 are shown in Figure 3. Failure to observe pure second-order response in the absence of oxygen is attributed to residual oxygen not re-
103
Chemiluminescence from Hypochlorite and Luminol Reaction
P
I
-8
-7
4
I
I
-6
-5
LOG [OCl-]
Flgure 2. Steady-state chemiluminescence intensity vs. hypochlorite concentration at (a) pH 9.6, (b) pH 10.1, (c) pH 10.7, (d) pH 11.7; (0)experimental points used in calculating first- and second-order fit; (0) other experimental points; (-) calculated fit to first- and second-order components; (- -) first-order component of calculated fit; (-) seconborder component of calculated fit. The relative units for CL intensity are the same for 2a through 2d.
--
moved by the purging process. Similar data although of poorer quality were obtained at pH 11.7. Kinetics of the Hypochlorite-Lurninol Reaction. Figure 4 plots the rate constant for the reaction between lob7 M hypochlorite and 4 X loV4M luminol. The kinetics are first order in OC1-. The rate constant decreases with increasing pH up to about 11.8. At this point the reaction rate increases rapidly with increasing pH until above pH 1 2 it is too fast to be measured with the flow system used in this study. The behavior of the OCl--luminol reaction a t pH’s above 12 was not further investigated. In the data of Figure 2, the most intense steady-state CL for a given OC1- concentration is observed a t pH 10.1. However, because the reaction rate decreases with increasing pH a larger fraction of the CL occurs after the OC1-luminol solution has left the cell a t the higher pH’s. For
this reason, it is probable that maximum quantum efficiency occurs a t pH’s higher than 10.1. This is consistent with earlier work in which the maximum quantum efficiency for the hypochlorite-luminol reaction was reported to occur at pH ll.ll Discussion Table I lists the proposed steps in the OCI--luminol reaction. The first step is the reaction between OC1- and luminol to produce an azaquinone. The azaquinone then reacts either with oxygen or OC1- to generate CL. Reaction 2 is responsible for the first-order component of steadystate CL us. [OCI-] data while reaction 3 causes the second-order component. As expected the first-order component of steady-state CL is sensitive to oxygen concentration while the second-order component is not. The Journal of Physical Chemistry, Vol. 79, No. 2, 1975
104
W.
Rudolf Seitz
TABLE I: Intermediate Reactions in the Overall Reaction between Hypochlorite and Luminol
0
0
coo-
[OCI-]
0
Figure 3. Effect of deoxygenatlng reagents on steady-state CL intensity vs. [OCI-] at pH 9.4: (H) steady-state CL in the presence of oxygen; (0)steady-state CL wlth deoxygenated reagents; (A)reduc-
&: NH2 0
tion in steady-state CL due to deoxygenation.
3.
N
0
0
2
9
10
I1
12
PH Flgure 4. Rate constants for reaction between OCI- and excess Iuminol as a function of pH measured at 21'.
This mechanism is analogous to that proposed previously for iodine and luminol.2 However, in the iodine-luminol system, a combination of first- and third-order processes was observed at pH's above 11where the rate of hypoiodite disproportionation is fast. There is no direct evidence for the azaquinone intermediate. However, azaquinones have been previously proposed as intermediates in hydrazide chemiluminescence. In the aueous persulfate-luminol-hydrogen peroxide system, kinetic evidence suggests that the first step is the oxidation of luminol to an azaquinone which then rapidly reacts with hydrogen peroxide to produce fairly intense chemiluminescence.3 Azaquinones have been prepared that produce the same chemiluminescence upon oxidation as the corresponding hydrazides. 7,8 The radical formed by one-electron oxidation of luminol has also been proposed as an intermediate in luminol chemiluminescence.4-6 However, a one-electron oxidation with hypochlQrite would require chlorine atoms or some other thermodynamically unlikely chlorine species. The The Journal of Physioal Chemistry, Vol. 79, No. 2, 1975
HzO
+
+ N, + hv
2"
+ OC1- + 2OH- 5
coo-
$3
HzO
+
+ C1- + hv
subsequent reactions of the one-electron oxidation product to produce first- and second-order processes would also require some unlikely intermediates. The expected variation of steady-state CL with [OCl-] concentration entering the cell can be calculated for low [OCl-] concentrations, as a function of the rate constants for the reactions in Table I and the CL efficiencies for reactions 2 and 3. At steady state the rate at which hypochlorite enters the cell equals the rate at which hypochlorite reacts plus the rate at which hypochlorite leaves the cell
Re = k,[OCl'] + k,[Aza][OCl'] + R , (3) where R e is the rate at which hypochlorite enters the cell ( M sec-l) k l is the pseudo-first-order rate constant for reaction 1 in the presence of 2 x 10-4 M luminol (concentration in the luminol-buffer solution is diluted by 50% in the cell), [Aza] and [OCI-] are the steady-state concentrations of azaquinone and hypochlorite, respectively, k 3 is the rate constant for reaction 3 ( M sec-l), and R1 is the rate a t which hypochlorite leaves the cell. A t low concentrations of hypochlorite entering the cell, [Aza] will be small, and the term k 3 [Aza][OCl-] can be neglected. This is equivalent to assuming that all of the hypochlorite entering the cell reacts with luminol and none reacts with azaquinone. As the concentration of hypochlorite entering the cell increases, the steady-state concentration of azaquinone increases and a greater fraction of the hypochlorite participates in reaction 3. This reduces the extent of reaction 1, and is one, if not the only, factor causing the observed CL for high hypochlorite concentrations to deviate from the expected CL calculated from the least-squares fit to firstand second-order components. At steady state the rate at which azaquinone is formed equals the rate at which azaquinone reacts
k,[OCl-] = k2[AzaT + k,[OCl-][Aza] (4) where kz is the pseudo-first-order rate constant for reaction 2 in a solution saturated with oxygen. The kS[Aza][OCl-] term can also be neglected in eq 4 a t small hypochlorite concentrations entering the cell. A t steady state the rate at which product, aminophthalate, is formed is constant with time
105
Chemiluminescence from Hypochlorlte and Lumlnol Reaction
P = k2[Aza]
+ k3[OC1'][Aza]
(5)
TABLE 11: Values Used to Calculate K242/K343
where P is the rate at which product forms ( M sec-l). From eq 4, neglecting k 3 [OCl-] [Aza]:
[Aza] = k,[OCl']/k2
(61
while from eq 3
[OCl'] = (R, - R,)/k,
(7) Substituting for [Aza] and [OCl-] in eq 5 leads to the expression
The observed steady-state CL intensity depends on the rate at which the product is formed in reactions 2 and 3 multiplied by 42 and 43, CL efficiencies for those reactions, respectively:
This equation shows that the reactions in Table I should result in CL us. hypochlorite data consisting of first- and second-order components at low hypochlorite concentrations. The values of k 1 and R e in eq 9 are known as well as the relative magnitudes of the first- and second-order terms:
where CL1 is the intensity of the first-order component of steady-state CL while CL2 is the intensity of the secondorder component. Rearranging eq 10 so that the unknown terms are on the left side of the equation
R 1 can be estimated from k 1, the flow rate through the cell, and the cell volume. The average residence time in the cell for an increment of solution is equal to the cell volume divided by the volume flow rate. For the data in Figure 2, the flow rate was 4.41 ml/min for each syringe or 8.8 ml/min altogether. The cell volume is approximately 0.8. ml. This means that the average residence time in the cell is 5.4 sec. For an increment of hypochlorite solution of concentration C 0 entering the cell:
where Ct, is the concentration unreacted after the average residence time in the cell and t , is the average residence time in the cell. If it is assumed that all increments of solution entering the cell remain in it for t,, then Ct,/Co = R 1/R e. Substituting from eq 1 2 in eq 11gives
Table I1 lists the experimental values for CL1/Cl2, Re, k 1, and e - k l t r and the calculated values for k 2 4 2/k 3 43 as a function of pH. The values for R e and CLl/CL2 are taken for a particular hypochlorite concentration which is indicated in the table. Table I1 shows that k2+2/k3$3 is very small. The data analysis presented here cannot determine what fraction of calculated value for k 2 4 2/k 3 43 is due to the difference in rate constants and how much is due to the difference in CL efficiencies. Nevertheless, it can be shown that $2 must be
0.73 1.8 x 5.5 x 1.1 X lo-? 7.3 X 10" 1.7 0.6 1.8X lom7 5.5 X lo-' 2.3 X 4.8 9.2 X a Concentration for which CLl/CLz and Re are measured.
9.6 10.1 10.7 11.7
10 6 10 50
0.10 0.036 0.015 0.005
0.42 0.14 0.076 0.027
smaller than 43. At low hypochlorite concentrations, steady-state CL us. [OCl-] can be resolved into first- and second-order components, without any observable deviations. To account for this it was necessary to assume that essentially all the hypochlorite reacts with luminol as discussed earlier. This is equivalent to saying that essentially all the azaquinone reacts with oxygen rather than OC1-. Under these conditions, a significant fraction of the CL still comes from reaction 3 in Table I (see data of Figure 2). Since reaction 3 contributes a significant fraction of the observed CL while it is contributing a negligible amount to product formation, 4 2 must be significantly smaller than
43.
The calculated value for k 2 4 2/12 3 43 is independent of pH from pH 9.6 to 10.7. However a t pH 11.7, k 2 42/12 3 43 is approximately 40 times larger than at the lower pH's. This is consistent with the observations that the second-order component of response is much smaller relative to the firstorder component at pH 11.7 and that the observed response does not deviate from the calculated response for first- and second-order components at high hypochlorite concentrations. This means that essentially all the OC1- is reacting with luminol rather than azaquinone, and essentially all the azaquinone is reacting with oxygen. When 2 X M hypochlorite is allowed to react with luminol at pH 11.7, the ratio of CL2/CL1 is 8.5. If it is assumed that reaction 3 is contributing less than 2% to P then the ratio 42/43 must be less than 0.0025. Since hypochlorite is a stronger oxidizing agent than oxygen, reaction 3 produces more energy for populating electronically excited states of the aminophthalate. Therefore, the greater efficiency predicted here for reaction 3 would be expected. In summary, this paper makes several contributions to the chemistry of luminol. It confirms earlier work on the iodine-luminol reaction.2 In particular it shows that the results observed with iodine at pH's below 11 are not due to hypochlorite disproportionation since hypochlorite does not disproportionate at room temperature. The theory developed in this work applies also to the earlier work showing that the reactions proposed for the 12-luminol reaction will in fact lead to the observed variation in steady-state CL with 12 concentration. As with hypochlorite, the hypoiodite-azaquinone reaction must be more efficient in generating CL than the oxygen-azaquinone reaction to explain the observations. Finally, this paper provides additional evidence that azaquinone intermediates occur in the chemiluminescent reactions of hydrazides. References and Notes (1) K. D. Gundermann, "Chemilumineszenz Organischer Verbindungen." 1st ed, Springer Verlag, New York, N.Y., 1968, pp 63-90. (2) W. R. Seitz and D. M.Hercules, J. Amer. Chem. Soc., 96, 4094 (1974). (3) M. M. Rauhut, A. M. Semsel, and 6. G. Roberts, J. Org. Chem., 31, 2431 (1966,. The Journal of Physical Chemistry, Vol. 79, No. 2, 7975
4 08 (4)
A.
Rainis and M. Szwarc
(8) K. D. Gundermann, H. Fiege, and &. Klockenbring, Justus Llebigs Ann. Chem., 738, 140 (1970). (9) W. R. Seitz, W. W. Suydam. and D. M. Hercules, Anal. Chem., 44, 957 (1972). (10) W. Nonidez, private communication. (1 1) H. H. Seliger In “Light and Life.” W. D. McElroy and B. Glass, Ed., Johns Hopkins Press, Baltimore, Md., 1961, pp 200-205.
E. H. White in “Light and Life,” W. D.McElroy and B. Glass, Ed., Johns
Hopkins Press, Baltimore, Md., 1961, pp 183-195. (5)E. Epstein and P. Kuwana, Photochem. Photobiol., 4, 1157 (1965). (6) P. E. Shevlin and H. A. Newfeld, J. Org. Chem., 35, 2178 (1970). (7) I 5 X lo5 M-l sec-l and the electron affinity of pyridine, Py, in DME is lower by at least 0.15 Y than that of triphenylene, Tr. The dimer, Na+(-Py-Py-)Na+, formed by this method is stablle in THF or DME. +
-
-
Pyridine, Py, may be reduced to its radical anion, Pym-, and under appropriate experimental conditions the latter may be identified by its esr spectrum.l-3 However, in contrast to radical anions of aromatic hydrocarbons, Py- salts rapidly dimerize and this reaction often prevents their detection in the investigated system. The formation of the diamagnetic dimer was first postulated by Smith4 and later by Ward5 who proposed for it a covalently bonded structure Na
PJa
Subsequently, Szwarc and his coworkers2 and simultaneously Schmulbach, Hinckley, and Wasmund6 proposed for it the dimeric dianion structure (I).The dimer is readily
(1) dehydrogenated, particularly in the presence of alkali metal, the overall reaction being I
bipyridyl
-4
.t
2NaM
and, indeed, many investigator^^-^ found bipyridyl and its radical anion as the products of pyridine reduction. These consecutive reactions have caused some confusion in the interpretation of experimental data. For example, the 335-nm band observed after brief exposure of a solution of pyridine in tetrahydrofuran (THF) to a sodium mirThe Journal of Physical Chemistry, Voi. 79, No. 2, 7975
-
-
ror was erroneously attributedg to the Py- radical anion. We.shal1 show, however, that it results from the presence of dimer (I),in agreement with previous suggestions.226 The rates of the various reactions consuming Py- are greatly influenced by the experimental conditions. For example, Talcott and Meyers found the half-life of Py- to be about 1 min when the radical was formed electrolytically in liquid ammonia, while Kemp, et aL,3 were unable to detect Py- in static experiments over shorter periods of observation when the base was mixed with a solution of sodium in liquid ammonia. The difference in the nature of the cation, NMe4+ in the electrolytic study and Na+ in Kemp’s investigation, was claimed to be responsible for these divergent observations. Substitution of a methyl group for the 4 hydrogen of pyridine seems to slow the dimerization of the respective radical anion^,^ and the dimerization of 2,6-dimethylpyridine radical anions seems to be even s10wer.~ According to Atherton, et al.,1° the radical anions of 3,5dimethylpyridine, prepared by sodium reduction in dimethoxyethane (DME), do not dimerize at all, presumably due to steric strain induced by the substituents. However, Kemp, et a1.,3as well as Talcott and Meyers,l found them to dimerize in liquid ammonia in less than 15 min. The capricious behavior of Py- is illustrated by the fact that only once was its esr spectrum recorded in HMPA2 (hexamethylphosphoric triamide), subsequent attempts at reproducing this result being unsuccessful. Nevertheless, the sharp and clear spectrum recorded in that experiment leaves no doubt about the identity of the paramagnetic species. We started, therefore, our investigation with the inten-
