Mechanism of Cobalt Catalysis of Luminol Chemiluminescence Timothy G. Burdo’ and W. Rudolf Seltz Department of Chemistry, University of Georgia, Athens, Ga.30602
The effects of pH, hydrogen peroxide concentratlon, and lumlnol concentratlon on the lnltlal Intensity of cobalt-catalyred lumlnol chemllumlnescence have been measured In the presence and absence of cltrate under condltlons where cobalt was the llmltlng reactant. The behavlor of lumlnol, which can act as a bldentate chelate, was compared to anthracene hydrazlde, whlch cannot act as a bldentate chelate, to determlne whether a lumlnol-cobalt complex Is an essentlal lntermedlate for lumlnescence as previously proposed. The results Indicate that a cobalt-hydrogen peroxide complex Is the essentlal lntermedlate requlred for chemllumlnescence. Any factor that affects the cobalt-peroxlde complex concentration slmllarly affects lnltlal chemllumlnescence Intensity. The rate-determining step In the cobaltcatalyzed chemllumlnescence reaction Is the oxidation of luminol by one-electron to a radlcal by the cobalt-peroxide complex.
The luminol reaction has attracted considerable attention as a means of determining trace metals because it is capable of detecting sub-part-per-billion levels of several important metal ions with simple and inexpensive instrumentation. In the presence of excess luminol (o-aminophthalhydrazide) and hydrogen peroxide, chemiluminescence (CL) from this reaction is proportional to metal ion concentration over several orders of magnitude ( I ) . Methods using photographic detection have been reported for cobalt, copper, and iron (2-4). Selective methods for Cr(II1) and Fe(I1) have been developed using photoelectric detection and have been demonstrated on both environmental and biological samples (5-7). Because the luminol reaction is inherently nonselective, it has been proposed that it be used as a detection system for ion exchange separation of metal ions ( I , 8). A small volume cell has been developed for chemiluminescence detection of chromatographically separated metals and used to measure Co(I1) and Cu(I1) separated as anionic chloride complexes in concentrated lithium chloride solutions (9). A serious obstacle to further application of the luminol reaction to trace metal analysis is the absence of knowledge as to how the metal ion functions to induce chemiluminescence from basic solutions of luminol and peroxide. This knowledge is necessary to predict which metal ions will catalyze luminol chemiluminescence and to predict and cope with interferences in chemiluminescence analyses for trace metals. Without such knowledge, all analytical applications must necessarily be completely empirical. In addition, elucidating the role of the metal catalyst in luminol chemiluminescence would be an important step toward a more complete understanding of the luminol reaction. I t has been proposed that the metal ion forms a complex with luminol which is then oxidized by hydrogen peroxide; however, the evidence for this is unconvincing (10-12). The basis for this hypothesis is the fact that luminol can act as a bidentate chelating agent, coordinating with a metal ion a t the amino group and one of the hydrazide carbonyls to form a 6-membered ring. Present address, Upjohn Company, Kalamazoo, Mich.
Cobalt was chosen for this study because it is the most efficient metal catalyst and it has a straightforward redox chemistry. Because of the low solubility of metal catalysts a t the basic pH’s required for luminol chemiluminescence, the number of techniques that can be applied to the study of the role of the metal is limited. The approach taken here was to measure initial chemiluminescence intensitv. .., integrated over the first five and a half seconds after mixing, as a function of several variables, pH, peroxide concentration, luminol concentration, and complexing agent concentration, and to see how these variables interact. T o test the hypothesis that a metal ion-luminol complex is an essential intermediate for chemiluminescence, the behavior of lumino1 was compared to anthracene hydrazide. Anthracene hydrazide chemiluminesces similarly to luminol, but, unlike luminol, it is not capable of forming a stable bidentate complex with a metal ion because it lacks the amino group, as shown in the structures below:
0
0
luminol
anthracene hydrazide
EXPERIMENTAL Apparatus. Except for the dark room experiments, CL was measured using a flow system previously described in detail (5, 6). Luminol (or anthracene hydrazide) in 0.1M KOH-HzB03 buffer is mixed with hydrogen peroxide in the flow system. This reagent is mixed with background solution (similar to the sample but containing no Co(I1)) directly in a cell positioned in front of a photomultiplier. A sampling valve is used to insert slugs of Co(I1) solution in the background flow line. Steady state CL is measured as the Co(I1) solution passes through the cell. Nitrogen bubbling is used to stir the cell. The steady-state CL measured in the flow system is essentially the same as CL from the time of mixing for the length of time an increment of solution remains in the cell, 5.5 seconds for the flow rates and cell volume used in this study. Because 5.5 seconds is much shorter than the total duration of Co(I1)-catalyzed CL (except in the presence of ethylenediamine), the measured steady-state CL is essentially an initial intensity. Absorption spectra for luminol and anthracene hydrazide were measured on a Cary 15 Recording Spectrophotometer. The fluorescence spectrum of anthracene-2,3-dicarboxylicacid was recorded on a Turner Spectro Model 210 Spectrofluorometer. Chemicals. Luminol from Eastman Organic Chemicals was dissolved in warm base and recrystallized by cooling the resulting solution. Anthracene hydrazide, purified by thin layer chromatography, was supplied by Karl Gundermann of the Technical University of Clausthal. All other reagents were used as received. Solutions were prepared using deionized water from a Continental Water Conditioning Company. Procedures. The effect of reaction pH on chemiluminescent response for both luminol and anthracene hydrazide was studied by adjusting the pH of a 1 X 10-4M solution of hydrazide dissolved in KOH-H3BO3 buffer with either 2M HC1 or 2M KOH before dilution. The reaction pH was then measured as it exited the chemiluminescence cell. Hydrazide concentration effects on chemiluminescent response were determined by filling the syringe with the appropriate concentrations of luminol or anthracene hydrazide in ANALYTICAL CHEMISTRY, VOL
47, NO. 9, AUGUST 1975
1639
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1
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a
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IO
I2
II
13
PH -6
-7
-6
-5
-4
3
Chemiluminescence intensity of anthracene hydrazide vs. reaction pH
Figure 3.
LOG LIGAND CONCENTRATION
Effect of ligand concentration on Co(I1) catalyzed chemiluminescence
Figure 1.
(A)Citrate ( 0 )Ethylenediamine Conditions 1 X 10-3M H202, 1 X 10-4M luminol, 1 X lO-’M Co(ll), reaction pH 10 6, and 4 41 ml/min /syringe flow rate
I
I
I
i ,
0
I2
I1
I3
PH
Figure 2. Chemiluminescence intensity minol vs. reaction pH
and fluorescence yield of Iu-
(H) Fluorescence yield (29).
(A)1 X 10-7M Co(ll). ( 0 ) 1 X lO-’M Co(ll), 1 X 10-4M citrate ( X 10). Conditions: 1 X 10-3M 10-4M lurninol, 4.41 ml/min./syringe flow rate
H202,
1 X
a KOH-H3B03 buffer solution. The effects of ligand and H202 concentration on chemiluminescent response were studied by making standard additions of ligand and Hz02 to both sample and background.
RESULTS Dark Room Experiments. When equal volumes of 10-2M H202, 10-6M Co(I1) a t p H 11, and 10-4M luminol in 0.1M KOH-HBBO~were mixed in a dark room, CL was observed visually for over an hour gradually decreasing in intensity. Further additions of Co(I1) led to an increase in CL intensity. The intensity was independent of the order of reagent addition except when the basic Co(I1) and H202 solutions were mixed first and allowed to stand before adding luminol. In this case, CL intensity was diminished by an amount proportional to the time interval between mixing Co(I1) and H202 and adding luminol. When the Co(II)-H~O~-luminol solution was made 1.6 X 1640
(A)1 X 10-7MC~(ll).( 0 )1 X lO-’MCo(ll), 1 X 10-4Mcitrate. Conditions: 1 X 10-3M H202,1 X 10-4Manthracene hydrazide, 4.41 ml/min./syringe flow rate
10-6M in ethylenediamine, CL was instantly quenched. When the Co(II)-H~O~-luminolsolution was made 1.6 X 10-4M in citrate, CL intensity was diminished but the duration of visible CL was unchanged. Co(II1) Catalysis. Steady state CL intensity catalyzed by 1 X 10-’M [(NH3)jCoCl]C12 was less than 5% of the intensity catalyzed by 1 X 10-7M Co(I1) in 5 X 10-7M NH4C1 as measured in the flow system in the presence of 10-3M H202 and 10-4M luminol. Effect of Organic Complexation. The effects of addition of organic ligands on light emission in the flow system are shown in Figure 1. At ethylenediamine concentrations greater than 10-’M, the steady state CL decreases as a function of increasing ethylenediamine concentration. Replacing ethylenediamine with citrate, a slight rise in emission intensity was observed as the concentration of citrate rose from to 10-6M. With citrate concentrations above lO-jM, the intensity of light emission recorded decreases as a function of increasing citrate concentration. Effect of pH. The effect of reaction p H on CL intensity for luminol is seen in Figure 2. At a Co(I1) concentration of 1 X 10-’M, the p H profile shows a maximum a t pH 10.9. Under the same conditions with the addition of lO-*M citrate to the 1 X 10-7M Co(I1) sample and background, a tenfold decrease in peak height was observed along with a shift in the maximum to higher pH. Figure 2 also shows thk fluorescence yield of luminol as a function of pH. If a fluorescence yield correction were made, the p H shift observed with citrate would have been even greater. The effect of reaction pH on CL intensity for anthracene hydrazide is seen in Figure 3. Using 1 X 10-7M Co(I1) as the sample, the pH profile shows a pH 11.0 maximum. On either side of the maximum, a sharp decrease in light emission is observed. With the addition of 10-4M citrate to the Co(I1) sample and background a significant decrease in peak height is observed a t pH’s below 12.0. This new maximum is also seen a t pH 11.0, but there is no steady decline in CL response between pH 11.0 and 12.0 as seen with the sample containing Co(I1) alone. Effect of Hydrazide Concentration. The effect of luminol concentration on CL intensity is seen in Figure 4. Using 1 x 10-’M Co(I1) as the sample, the maximum light emission is observed a t 2 X 10-4M luminol. Replacing the M the intensity sample with 1 X 10-6M C O ( I I ) - ~ O - ~citrate, of light emission per unit Co(I1) decreases. The maximum
ANALYTICAL CHEMISTRY, VOL. 47, NO. 9, AUGUST 1975
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Chemiluminescent intensity vs. luminol concentration
(A)1 X 10-7M Co(ll). (0) 1 X 10-6M Co(ll). 1 X lOP4Mcitrate. (m) 1 X 10-6MCo(ll), 1 X 10-2Mcitrate ( X 40). Conditions: 1 X 10-3MH202. reaction pH 10.6, and 4.41 ml/min./syringe flow rate
4s-
2
---_-
_ I 5c3
*&3
WAiELENGTh qrr
LOG LUMINOL CONCENTRATION
Figure 4.
I
4 3
\ \
Chemiluminescence and absorbance spectra of anthracene hydrazide
Figure 6.
-
-) Absorbance spectrum of 5 X 10-4M anthracene hydrazide in 0 1M NaOH-HsB04 buffer (--) Chemiluminescence spectrum of anthracene hydrazide
(-
II
2-
,'A
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, W
I
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LOG ANTHRACENE HYDRAZIDE CONC
0
Chemiluminescent intensity vs. anthracene hydrazide concentration
Figure 5. (
--
) Uncorrected light emission observed. ( A -
- - - - - A ) Emission cor-
rected for self absorbance. Corrected data point for 4 X 10-3M is approximately 6 times the corrected peak height at 1 X 10-3M. Conditions: 1 X 10-3MH202, 1 X lO-'MCo(ll), pH 10.6, 4 41 ml/min./syringe flow rate
is observed to be slightly skewed toward higher luminol concentrations. A new solution of 1 X 10-6M C0(11)-10-~M citrate shows a 40-fold decrease in measured CL compared t o 1 X 10-6M C O ( I I ) - ~ O - ~ M citrate. A definite shift of the CL maximum to higher luminol concentrations is observed. Figure 5 shows a plot of CL as a function of anthracene hydrazide concentration. Under conditions similar to the luminol concentration study with 1 X 10-7M Co(I1) as the sample, the maximum light emission was observed a t an anthracene hydrazide concentration of about 4 X 10-4M. The data for anthracene hydrazide need to be corrected for the absorption of CL by the unreacted hydrazide, as can be seen from Figure 6 which shows both the absorption spectrum and the CL emission spectrum for anthracene hydrazide. Instead of measuring the CL spectrum for anthracene hydrazide directly, the fluorescence spectrum of anthracene-2,3-dicarboxylicacid was measured. I t has been shown that hydrazide CL spectra generally match the fluorescence spectra of the corresponding dicarboxylic acids (13-15). Figure 5 includes steady state CL corrected for absorption by anthracene hydrazide. A simplified self absorption correction was performed by calculating the average extinction coefficient of anthracene hydrazide from 410 to
-3
-2
I
LOG H202 CONCENTRATION
Figure 7.
Chemiluminescent intensity vs. hydrogen peroxide concen-
tration (m)
1 X 10-7M Co(ll). (A)1 X 10-6M Co(ll), 1 X 10-4M citrate. (0) 1 X Co(ll), 1 X 10-'M citrate. Conditions: 1 X 10-4M luminol, pH 10.6, 4.41 ml/min./syringe flow rate
455 nm. The average path length was assumed to be the radius of the cylindrical chemiluminescence reaction cell, about 0.5 cm. A Beer's law calculation was then applied for the average absorbance a t each anthracene hydrazide concentration and the corrected intensity calculated. The correction factor for self absorption is large a t 1 X 10-3M and extremely large a t 4 X 10-3M. Consequently, the exact positioning of these corrected data points is associated with some error; however, a trend toward increasing chemiluminescence intensity with increasing anthracene hydrazide concentration can be seen. For luminol, the absorption and CL spectra do not overlap so there is no need for this correction. Effect of Hydrogen Peroxide Concentration. The results of the study of CL intensity vs. hydrogen peroxide concentration are seen in Figure 7 . With a sample solution of 1 X 10-jM Co(II), a steady rise in intensity with increasing H202 concentration is observed a t low concentrations of H202. Maximum intensity is observed between 1 X 10-3M and 2 X H202. After the maximum has been reached, a steady decline in emission is observed with increasing H202 concentration. When the concentration of ANALYTICAL CHEMISTRY, VOL. 47, NO. 9, AUGUST 1975
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Co(I1) is increased 10-fold to 1 X 10-6M and made 10-4M in citrate, a similarly shaped curve is observed with a slight shift in the maximum toward higher HzOz concentrations. Repeating with 1 X 10-6M Co(I1) and making the sample solution and background 10-2M in citrate, a large decrease in light intensity is seen a t HzOz concentrations below 2 X 10-3M. The maximum CL intensity is observed a t 9.5 x 10-3M H202, a noticeable shift from the previous two curves. Again, concentrations greater than the maximum show a continuous decrease in chemiluminescence. DISCUSSION Catalyst Identification. Co(I1) catalyzes the reaction while Co(III), a t least the complex tested here, does not. Under the conditions of all the experiments in this study, large excesses of hydrazide and peroxide, the kinetics of CL appear to reflect the rate a t which Co(I1) is oxidized to inactive Co(II1). The reduction in CL intensity when Co(I1) and H202 are premixed in base and allowed to stand is due to the slow oxidation of Co(I1) to Co(II1) by peroxide. In Co(II1) is inbase, the oxidation potential for Co(I1) creased so that the reaction is thermodynamically feasible (16):
-
Co(OH)z(s)
+ OH-
= Co(OH)&)
+ e-
E o = -0.17 V (1)
The fact that Co(I1) is the limiting reactant is shown by the observation that new additions of Co(I1) restore the initial CL intensity. Ethylenediamine forms a Co(I1)-complex that is rapidly oxidized to the Co(111)-complex. The immediate extinction of CL upon adding ethylenediamine reflects this reaction, since some residual CL would be expected if ethylenediamine's only effect were to tie up Co(I1) as citrate does. Possible Mechanisms. Because hydrogen peroxide and hydrazides do not react in basic solution in the absence of a catalyst, it can be assumed that the catalyst interacts with one of the reactants to produce a species that can then react with the other reactant. Three possible interactions can be envisioned. The metal ion can form a complex with the hydrazide as previously proposed, or it can interact with the hydrogen peroxide, either by complexing with the peroxide or by being oxidized to Co(II1) with the production of an OH radical which could then react with luminol. There are several difficulties with the proposed metal ion-luminol complex. First, it would be expected that lumino1 would behave very differently from anthracene hydrazide because there will be a large difference in their formation constants. However, the CL intensity is about the same for anthracene hydrazide and luminol. Second, it is difficult to explain why there is a decrease in CL a t high luminol concentrations. The explanation that this is due to fluorescence quenching is contradicted by the observation that with other oxidizing systems such as Fez+0 2 and Fe(CN)e3--H2O2, the maximum CL is observed a t luminol concentrations greater than 10-2M (6, 17). Third, one would expect CL intensity to be linearly proportional to luminol concentration since the complex concentration would be proportional to the amount of luminol; if the metal is the limiting reagent. In fact, CL is linearly proportional both to metal ion concentration and peroxide but not to hydrazide. A second possibility is that the following reaction is occurring Co(I1)
+ H202
+
Co(II1)
+ OH- + OH.
(2)
The OH- could then react with hydrazide to initiate the sequence of reactions leading to CL. This reaction is known to occur with some metals in acidic solutions. However, 1642
* ANALYTICAL CHEMISTRY, VOL. 47, NO. 9, AUGUST 1975
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Table I. Reaction Scheme for the Cobalt-Peroxide Complex CO-HO,' + cit3- i Co-citl- + H 0 2 ' Co-HO,' A H' ZZ? CO" + H,O, CO-HO,' + OH- ZI? CO4H' + HO,HZO, f H' + HO2K, = lo-'? Co-citl- - luml- e Co-lum" - cjt3[Co-HO,'] = K, [HO?-I (~co-[C~-HO?+]} IC, = formation constant f o r cobalt-peroxide complex C, = total cobalt concentration i n all forms CO-HO~' i HO2CO" - 0, + 2 OH[CO-HO,'] = K e g [~it3-l "21 { C~,-[Co2'-HO~-]~-
K,,
(12)
(13)
= equilibrium constant for Reaction 6.
since base is one of the products, it should be suppressed in basic solution. A more serious problem with this proposed mechanism is the observation that ethylenediamine completely quenches CL. The presence of ethylenediamine accelerates the kinetics of Co(1I) oxidation and reduces the oxidation potential (16), thus providing more energy for the formation of OH-. An increase in OH- would be expected to give brighter CL, but this is not observed. The failure of the two previously proposed mechanisms to explain the experimental results necessitated the formulation of a new mechanism which would be consistent with all observations. This new mechanism involves the formation of a cobalt-peroxide complex as the important intermediate leading to luminescence (Equation 3). This intermediate then reacts with luminol in basic aqueous media to produce a luminol radical in the rate-determining step (Equation 4). The luminol radical then reacts with H202 to produce the light emission (Equation 5 ) . This reaction presumably occurs in several steps with which this study is not concerned. Co2+ Co2+-H02-
+ HOn-
~i Co2+-H02-
+ H20 Luminol-
-
-
co3+-3 OHLuminole-
+ HOz-
+ Luminol-
Light
(3) (4)
(5)
Since the cobalt-peroxide complex is the limiting reagent in the rate determining step, anything that affects its concentration should produce a similar effect on the intensity of the chemiluminescent emission observed. The remainder of this discussion will be devoted to showing that the observed variations in CL intensity with pH, ligand concentration, peroxide concentration, and hydrazide concentration are qualitatively consistent with the expected effects of these variables on cobalt-peroxide complex concentration. Because of the complexity of the multiple equilibria involved particularly with various mixed cobalt complexes plus the absence of formation constant data for cobalt-peroxide and cobalt-luminol complexes, quantitative support for the proposed mechanism was not attempted. The equilibria presented here are all over-simplified for the purpose of illustrating qualitatively what changes in concentrations of various reactants would have (Table I). Effect of Complexation. Complexing agents such as citrate reduce CL by tying up Co(I1) and reducing the concentration of cobalt-peroxide complex (Reaction 6, Table I). pH Effects on Response. Both high and low pH's should have an effect on the concentration of cobalt peroxide complex. Reaction 7 shows what would be expected a t lower pH's (pH's lower than 11.0).As the pH decreases, in-
creasing amounts of hydronium ion will react with the complex causing its dissociation. The overall effect expected is a decrease in light emission as the pH decreases. Reaction 8 shows what would be expected at higher pH's. With increasing reaction pH, hydroxide ion will increasingly react with the complex to form cobalt hydroxide, a non-light producing species. The overall effect expected with increasing pH should be a decrease in the amount of light observed. Combining both the high and low pH effects, it is expected that a plot of chemiluminescent emission vs. pH would show a maximum at some pH. This is observed experimentally for both luminol and anthracene hydrazide. An alternative explanation for the decrease in CL intensity at low pH's is the possibility that Reaction 4 between the luminol radical and hydrogen peroxide yields CL only with the peroxide anion. The pK, for Hz02 is 1 2 (Reaction 9). Effect of Citrate Complexation on Chemiluminescence vs. pH. Citrate complexation reduces CL a t all pH's. However, the effect is smaller at high pH's. This is reasonable since citrate selectively complexes Co(I1) and thereby reduces the effect of OH- complexation at high pH's. In effect, a t high pH, one masking effect is replacing another masking effect, with little additional decrease in CL (proportional to cobalt-peroxide complex concentration). At low pH, the two effects reducing complex concentration are additive; therefore, there is a greater reduction in CL intensity at low pH's. This is best seen in the data with anthracene hydrazide because, for this hydrazide, the fluorescence efficiency of the emitting dicarboxylic acid is independent of pH. For anthracene hydrazide in the presence of citrate, there is virtually no decrease in CL a t high pH, indicating that the hydroxide effect has been almost totally masked. Effect of Hydrazide Concentration. Since luminol can act as a bidentate chelate, a high concentration of luminol should lead to a decrease in CL because luminol will tie up cobalt and decrease the concentration of cobalt-peroxide complex. This effect explains why there is a decrease in CL at high luminol concentrations. In the presence of citrate, maximum CL intensity is observed a t higher hydrazide concentrations than in the absence of citrate. This is reasonable if the equilibrium between Co-citrate and Co-luminol is considered, Reaction 10. Increasing the citrate concentration means that more luminol is required to complex a significant fraction of the Co(I1). Therefore, a higher luminol concentration is required to cause a decrease in CL due to complexation. In the case of anthracene hydrazide which cannot form a bidentate chelate, CL does not decrease at high hydrazide concentrations when the hydrazide absorption is corrected for. Effect of Peroxide. At low H202 concentrations, the intensity of CL emission should be proportional to H202 concentration. This is shown by Equation 11 in Table I, derived from the expression for the formation of Co-HO2complex. As long as [Co2+-H02]
