Carbonate and Superoxide Ion-Radicals as Intermediates in the Chemiluminescence of the Ferricyanide — Hydrogen Peroxide Redox System

Zeitschrift fur Physikalische Chemie Neue Folge, Bd. 148, S. 197-214 (1986)© by R. Oldenbourg Verlag, Miinchen 1986

Carbonate and Superoxide Ion-Radicalsas Intermediates in the Chemiluminescence of theFerricyanide — Hydrogen Peroxide Redox SystemBy Jacek Wierzchowskil, Danuta Slawinska1 and Janusz Slawinski21 Department of Physics, Academy of Agriculture, ul. Wojska Polskiego 38/42,60-637 Poznan2 Institute of Physics, Pedagogical University, ul. Podchorazych 2,30-084 Krakow, Poland

(Received August 22, 1985; accepted January 3, 1986)

Chemiluminescence / Oxygenated active species / Ferro-ferri redox systems

A weak chemiluminescence with an average quantum yield of 10 10 in the spectralrange 380 —620 nm with maximum at 460 + 5 nm accompanies the reduction ofFe(CN)3," by H202 in alkaline solutions. Chemiluminescence and fluorescence spectracover the same spectral range and show no red emission bands. The dependence of thelight output on the concentrations of the reactants follows the stoichiometry of thereaction and reaches the maximum at concentrations of the order of 10~2 M andpH 10.5

—

12.0. The studied redox system involves Haber-Weiss and Fenton-like reactionswith close analogy to 02 and H202 activating biochemical processes. The lack of 02(lzfg)- quenching, a strong enhancement of the emission by carbonates and formates andits decrease by the scavenger of 02, tiron, indicate the recombination of HC03 and/orC02 radicals as the chemiexcitation step.

Bei der Reduktion von Fe(CN)|" mit H202 in alkalischer Losung tritt im Spektralbereich380

—

620 nm eine schwache Chemilumineszenz mit einer mittleren Quantenausbeute von10"10 auf, deren Maximum bei 460 + 5 nm liegt. Die Spektren der Chemilumineszenzund der Fluoreszenz liegen im gleichen Spektralbereich und enthalten keine Banden imRoten. Die Abhangigkeit der Lichtausbeuten von der Konzentration der Reaktionsteil-nehmer entspricht der Stochiometrie der Reaktion, sie erreicht ein Maximum im pH-Bereich 10,5

—

12,0 bei Konzentrationen um 10"

2 M. Das Redoxsystem weist Reaktionenvom Typ der Haber-Weiss- und Fenton-Reaktionen auf und zeigt enge Analogie zu denbiochemischen Prozessen, die 02 und H202 aktivieren. Das Fehlen der 02 (MgJ-Loschungund die erhebliche Verstiirkung der Emission durch Karbonate und Formiate sowie ihreAbschwachung durch den 02-Inhibitor Tiron weisen auf eine Rekombination derHC03- und/oder C02-Radikale als dem chemischen Anregungsschritt hin.

198 J. Wierzchowski, D. Slawinska and J. Slawinski

1. IntroductionJust 20 years ago Stauff et al. [1—5] proposed 02 —02 van der Waalscomplexes as possible emitters of chemiluminescence (CL) common tonumerous radical oxidation reactions. Since then the importance ofoxygenated species for the majority of chemiluminescent reactions havebecome increasingly apparent. Among many light-producing reactionsstudied, a weak CL from an apparently simple redox system consisting ofan alkaline solution of K3Fe(CN)6 and H202 was observed by Stauff andSchmidkunz [1], This system has been widely applied for oxidation oforganic compounds and promotion of CL of luminol [6, 7], polyphenols[8, 9] and other compounds [6, 10]. An additional important aspect of theabove redox system stems from the fact that it might be a simple model fornumerous biochemical reactions activating 02 and H202. The mechanismof the opposite reactions:

2Fe(CN)|- + H202 + 2 OH" -»2Fe(CN)S" + 2H20 + 02 (1)and

2Fe(CN)£" + H202 + 2H + -> 2Fe(CN)3,- + 2H20 (2)probably involves the Haber-Weiss and Fenton-like reactions and revealsthe following features: (i) an electron transfer process takes place between2e" donor (l/202 or H202) and 1 e" acceptor molecules [Fe(CN)3"/Fe(CN)g"] similarly as in cytochromes, (ii) the Fe-ligand field forms theoctahedral configuration both in Fe(CN)6~/Fe(CN)6~ and in the haemcenter ofmany haemoproteins, and (iii) both the model and natural systemsproduce reactive oxygenated species which play an important role indetoxication of xenobiotics, DNA defects, carcinogenesis, cytotoxicity andmay other processes [11],

It has been postulated by Stauff et al. [12] that the weak CL observedduring the reaction of Fe3+ or Mn2+ with peroxides is due to side reactionsof free radicals generated in the system with contaminating carbonateradical-ions. Recent papers on photochemically generated carbonate andformate ions and radicals support this view [13]. Also the CL occuringduring the oxidation of H202 by periodate has been ascribed tocontaminating carbonates and their radicals [14].

The purpose of the present investigation is to characterize the CL ofthe Fe(CN)3,"/Fe(CN)t5_-H202 system, to provide evidence for or againstthe 'Of participation and to elucidate the role of 02, OH - and carbonatein the chemiexcitation steps.

2. Materials and methodsCommercially available compounds were of analytical grade: l,2-dihydroxybenzene-3,5-disulfonic acid (Tiron), mannitol and thiourea from Merck, 1,4-diazabiscyclo 2.2.2-oktane (DABCO), ethanol, methanol and 1,3-diphenylisobenzofurane (DPBF) from

Carbonate and Superoxide Ion-Radicals 199

Aldrich, ascorbate from Polfa, Krakow, H202 of special purity for semiconductors(POChem, Gliwice, Poland). Other reagents were of analytical grade from the last firmand additionally recrystallized from water or water-alcohol mixtures. All solutions were

freshly prepared and kept in a dim light. Redistilled water from the all-glass apparatuswas used.

The intensity and kinetics of the CL were measured using an RCA 6655 A photo-multiplier with an S-ll photocathode, sensitive in the spectral range 340

—

650 nm. Thesignal was recorded with a K-201 Zeiss integrating recorder. Solutions were rapidlyinjected (0.2 s) into a cuvette externally aluminized, thermostated and mounted on thefront of the photomultiplier. The temperature of reacting solutions was 20

—

23(±0.5)°CThe reproducibility of the CL peak-height was about 10—16%, and eachmeasurement was repeated 3

—

6 times. The system was calibrated for the absolute numberof photons emitted by means of the low level chemiluminescent luminol standard [15].The detailed description is given elsewhere [16]. The spectral distribution of the CL wasmeasured combining the single

—

photon counting mode with the cut-off filter method.An EMI 9558 QB photomultiplier sensitive in the range 180

—

800 nm (S-20) was cooledto

—

70°C. Measurements were performed in the flow system with a constant delivery offresh solutions to the observation cell. The thickness of the cell was decreased to 4 mm

to diminish the internal filter effect, while the active volume of the cell was increased to20 ml. A conical light-guide joined optically the cell with the photomultiplier. Theefficiency of photon counting, t\, measured as the number of counts per the number ofphotons emitted r\ = Nc/Nhv was (6 ± 2)

•

10"3 counts//iv. The emission spectra werecorrected for the spectral sensitivity of the photocathode and for the filters' transparencyusing the method described in reference [17], The absolute number of photons emitted,Nh„ was calculated with the luminol chemiluminescent standard reaction [15],

Fluorescence and phosphorescence spectra were recorded with a Perkin-Elmer MPF-3and an Aminco-Bowman SPF instrument. For phosphorescence measurements, samplesof K3Fe(CN)6, K4Fe(CN)6, buffers and H202 were deoxygenated by bubbling with pureAr for 20 min and then frozen to 77 K in liquid nitrogen. Methanol-water solutions (4:1,v/v) were used. Because of the weak photoluminescence intensity with a quantum yieldof the order of ~ 10~3, and for kinetic reasons, samples with a high optical density (0.5

—1.5) were used. In order to calculate the internal filter effect, the formula:

F0_

2.303 D (di-d2)~F ~

\o M<-10'Dd2

was applied [18]. Here F0 and Fare "normal" and weakened photoluminescence in-tensities, di and d2 optical paths to edges of the exit slit, and D is an optical density of a

sample. This correction factor of the internal filter effect is, however not quantitativeabove an absorbance of about 0.3. Therefore the corrected values of the fluorescenceintensity may be underestimated.

Absorption spectra were measured with a Zeiss Specord UV-VIS spectrophotometerusing 1 cm-path quartz cuvettes.

3. Results and discussion3.1. Absorption spectraThe absorption spectrum of ~ 1 mM aqueous solution of K3Fe(CN)6 showsthree high intensity bands at 260,304 and 425 nm which are to be attributed


1.5

12- v-W'Ucm'1]

Fig. 1. Absorption spectra of the reaction solution of 1.6 mM K3Fe(CN)6 + 0.5 mAfH202 + 10 mM phosphate buffer, pH = 11.6. OD = optical density in arbitrary units.Numbers denote the reaction time

to ligand-metal charge transfer (LMCT) transitions involving the r2g"metal" orbitals. Ligand field bands appear as shoulders at 396, 322 and282 nm in agreement with literature data [19, 20]. In the alkaline reactionsolution the lowest energy band at 418 nm is slightly shifted to 423 nm

(Fig. 1). The absorption spectrum of the reaction product, Fe(CN)g~,shows two moderately weak bands at 320 and 270 nm. The bands can beattributed to the spin-allowed ligand field transitions [19]:

Mlg-1rIgand 1axm^1tu. (3)At longer wavelengths (422 nm) there is a very low absorption assigned tothe lAlg -* 3Tlg spin-forbidden ligand field transition [20]. Absorptionspectra do not depend on pH in the range 6

—

13. Changes of the absorptionspectra during the redox reaction between Fe(CN)3,- and H202 are pre-sented in Fig. 1. The decrease of absorbancy at 425 and 304 nm, the increaseat 260 nm and the isosbestic point at 283 nm indicate the transformationof Fe(CN)6~ to be Fe(CN)4," according to the overall reaction (1). It was

very important to check if this reaction is reversible, i.e. to exclude side-reactions such as e.g. exchange of CN-ligands for H20 or H202. The


following data: (i) the isosbestic point at 283 nm, (iii) the equivalence ofthe intensity increase at 260 nm to the intensity decrease at 425 nm, and(iii) the observation that the "recycling" of the redox reaction by consecu-

tive acidification of Fe(CN)g"" + H202 and the alkalization of Fe(CN)3,"+ H202, does not change the position and intensity of the absorptionmaxima, lead to the conclusion that there is a full reversibility within thelimits of sensitivity of the spectrophotometric method.

3.2. Chemiluminescence spectraWhen the concentrations of the reactants and the rate flow are optimizedwith respect to maximum CL intensity, 7max, the steady-state CL intensityequals about 6 • 103 cps • cm-3. This count rate corresponds to about9 • 105 photons • s~1 and is high enough to measure reliably CL spectra (thedark count rate is 90 cps). As it is seen from Fig. 2, the emissions covers abroad spectral range from 380 to 620 nm with maximum at 460 + 5 nm.

There are no significant differences between spectra measured for thereaction with NaOH, phosphate buffer at pH = 9.0, or carbonate bufferpH = 10.4 (not shown). The spectra presented give the first evidence thatthere is no red emission at 634 and 703 nm, typical of the radiative transi-tions in 02-dimols. The short-wavelength limit of the spectrum (380 nm)corresponds to an energy of 315 kJ

mol-1. This implies an enthalpy of\(

—

AH)\ > 315 kJ • mol-1 of an elementary exergonic reaction, sinceAH > Nhc/L

3.3. Photoluminescence spectraAttempts have been undertaken to measure fluorescence from Fe(CN)33"in neutral and basic aqueous solutions. No fluorescence has been foundusing whole the available 2.exQ and Xem range as well as the maximumsensitivity. The measured signals did not exceed a fluorescence blankoriginating from traces of impurities in water or buffers (Fig. 3). The lackof ferricyanide's fluorescence agrees with the generally accepted opinionthat the Fe-ion acts as an efficient internal quencher in Fe-complexes [21].However, the reaction mixture shows a very weak emission within thespectral range 380

—

560 nm with maximum at 460 nm. Fluorescence emis-sion spectra, Fig. 3, agree with these of CL. Corrections for the internalfilter effect calculated from Parker's formula show that the observed fluo-rescence intensity of the system K3Fe(CN)35~ + H202 + OH" (buffer,NaOH) is distinctly greater than that from a blank solution withoutK3Fe(CN)6. This means that the absorption band of Fe(CN)|" with 2max =

425 nm overlaps the fluorescence emission to a significant degree and thatthe formation of a fluorescent product is Fe(CN)6~-dependent.

A very weak phosphorescence from K3Fe(CN)6 in methanol-waterglass has a maximum at 470 nm that varies with the excitation wavelength.


1.0 i

0.5

2. 00)I 1.0

rH

W^*rh+^-._i

0.5

400 500 600 700

A, nm

Fig. 2. Chemiluminescence spectra from the reaction of 10 mM K3Fe(CN)6 + 10 mMH202 in phosphate buffer, pH = 9.0 (above) and with 10 mM NaOH (below). The heightof the rectangles is equal to the mean value of the relative spectral intensity calculatedfrom 5 parallel measurements. The length of the vertical segments shows the mean

experimental error. The width of the rectangles corresponds to the difference inwavelengths between two succesive cut-off filters. All spectra are corrected for the spectralresponse of the photomultiplier and filters transmission. Temperature 23 °C

There is no correspondence between the phosphorescence excitation spec-trum and the absorption spectrum of K3Fe(CN)6, and the intensity ofphosphorescence diminishes after a repetitive recrystallization of the samplethat indicates a mixture.

3.4. The quantum yield of the chemiluminescenceThanks to the absolute calibration of the system it was possible to evaluate<P from the formula:

27/0 =

-V-c-N


400Atnm]

Fig. 3. Fluorescence emission spectra of the K3Fe(CN)6 —H202 —OH" system. Initialconcentrations of reagents: 4 mM K3Fe(CN)6 + 20 mM K2HP04 with addition of:a-

1 mM H202 + 2 mM KOH; b-

1 mM H202 + 2 mM KOH; c-

10 mM H202+ 2 mM KOH; d

-

1 mM H202 + 0.5 mM KOH; e-

1 mM H202 + 0.125 mMKOH; f

-

1 mM + 0.25 mM KOH; 2exc = 320 nm for a, and 350 nm for b-f. Thecurve g shows fluorescence from solutions c, e and f without K3Fe(CN)6. The spectrawere measured about 2 min after mixing the substrates

where V is the total volume of the light-emitting solution in the cell, c isthe concentration of Fe(CN)|" and N is Avogadro's number. The valueof the integrated light intensity 27/ was calculated from the kinetic curve ofthe CL, / = f{t), for t = 0 to t = 5 min (/ < 2% of /„).

The calculated values of <Z> range from 6 • 10~n to 3

10"9 (± 60%)and strongly depend on the concentration of Fe(CN)g" and OH". Thesmall value of $ results from two factors: (i) a low value of <Pt, since thefluorescence from the reaction mixture is extremely weak ((^ < 10~3),and (ii) the fact that about 25% of photons emitted are reabsorbed becauseof the CL and the 425 nm-absorption band overlapping. The probabilityof the excited state generation, <Pexc, roughly estimated from the abovedata, should not exceed 10"5 (for the chemical yield <PC assumed to be<0.1).


Table 1. The effect of 02('/lg)-quenchers on the maximum light intensity, /max, and thedecay time, f1/2, of the chemiluminescence of the system: Fe(CN)36"

—

H202—OH"

Quencher Q [Q], mM /max, a .11. ^1/2? S

Without (control) 0 19.8 ± 0.6 16.5 ± 0.3

Azide 0.1 20.0 + 0.6 17.3 + 0.40.4 22.6 + 0.8 18.6 + 0.41.0 22.7 ±0.6 20.0 + 1.29.0 24.8 + 0.9 17.3 + 0.3

16.0 28.1 + 1.6 17.8 +0.9Control 0 17.8+0.2 10.8 + 0.8DABCO 1.0 16.5 + 0.3 5.0 ± 0.8

3.0 20.5 + 0.9 2.9 + 0.23.0* 20.3 + 0.8 2.5 + 0.7

Control 0 16.3+0.5 16.3 + 2.0DPBF in 2.6 M MeOH 0.02 0.5 + 0.1 18.4 + 3.2MeOH alone 0 2.4 + 0.1 14.2 + 0.7

DABCO-l,4-diazabiscyclo 2.2.2-octane, DBBF-l,3-diphenylisobenzofurane, MeOH-methanol, *in 160 mM phosphate buffer

3.5. The problem of the emission from oxygen dimols

Further data have been gained regarding the possible participation of 02('dg'dg) in the reaction studied and the CL. The ratio of CL intensitymeasured at the wavelengths where the emission from i(02)* does occur

to X = 595 nm (no oxygen emission) has been found as: /700//595 = 0.8,and I626/IS9S = 1.1. Instead, the ratio/476//595 = 90 convincingly indicatesthat there is no emission at 634 and 703 nm, typical of the radiative tran-sitions:

(02)2(1dg1dg)„ = 0^202(32;-32:-)„ = o 2 = 634 nm (4)(O2)2(1dg1dg)u = o^ 202(327g-32:g-)1)=1 A = 703 nm. (5)Several well-known quenchers and traps of 02(1dg) were used to exam-

ine additionally its participation. The results are summarized in Table 1. Ina wide range of concentrations, azide does not quench the CL, on thecontrary, it enhances the light intensity and does not influence the kineticsof the CL decay (;1/2-values). Another quencher of 'O*, DABCO, has no

effect on 7max, but strongly decreases t1/2. This compound as a relativelystrong base increases [OH-] in aqueous solutions, thus possibly stimulatingthe rate of Fe(CN)3," reduction and the CL. However, experiments withDABCO repeated in more concentrated buffers (* in Table 1) did notshow any significant effect. DPBF has to be used in methanolic solutions.Compared with the effect of methanol alone, the net influence of DPBF is


30 24 18 12- v- 10"3 [cm"1!

Fig. 4. Absorption spectra of the reaction solution in the presence of 50 mM Na2C03.Concentration of reagents: 1.6 mM K3Fe(CN)6, 0.5 mM H202, 10 mM phosphatebuffer, pH = 11.6. Numbers denote the reaction time

also not significant. All these results provide strong evidence that theemission from excited oxygen dimols does not contribute to the observedCL.

3.6. The influence of carbonates

Taking into consideration the observation of Stauff et al. [12] regardingthe influence of HCOJ ions on the Cl of the Fe2 +-system, the effect ofNa2C03 was examined. As it is seen from Fig. 4, carbonates do not changethe absorption spectra but very distinctly increase the rate of theFe(CN)3,- reduction. Kinetic curves measured under identical experimentalconditions are presented in the common semilogarithmic coordinatesystem, Fig. 5. Comparison of the CL kinetics with the rate of theFe(CN)e~ reduction clearly shows that CO2- ions accelerate both proc-esses, the CL being more enhanced. It is also evident that the CL decay ismuch faster (k = 0.0094 s"1) than the reduction rate (k = 0.0022 s"1),this relationship is observed in all experiments performed.

The influence of CO2- and HCOJ on the maximum CL intensity ispresented in Fig. 6. Since the kinetics of the CL did not vary with the


/[min]

Fig. 5. Comparison of the change with time of the chemiluminescence intensity (/) andthe rate or reduction of Fe(CN)|~, measured as absorbance at 420 nm (A), in a

semilogarithmic plot. The kinetics of the chemiluminescence and the reduction ofFe(CN)3," were measured under identical experimental conditions and initial concentra-tions of substrates, given in Fig. 1

carbonate concentration—

/1/2 being nearly constant and equal to about5

—

7 s—

the same applies to the light sum (integrated intensity SI).However, the proportion of carbonate to bicarbonate ions, and/or the pHof the medium influenced both the intensity of the signal and its dependenceon the buffer concentration, Fig. 6a. A greater signal at higher pH valuesmay suggest a greater role of bicarbonate, in comparison with carbonate,in the chemiexcitation mechanism. In order to prove this hypothesis, thedata are replotted in such a manner, that only the concentration of bi-carbonate is accounted for, Fig. 6 b. From this figure it is seen that onlythe concentration of bicarbonates, but not the pH itself, determines theintensity of the CL. It should be emphasized that the concentration of the


0 0.1 IMl 0.2 0 0.1 IM]

lHC0jl + ICOj'l [HCO3IFig. 6. (a): Relative chemiluminescence intensity during the reduction of Fe(CN)6~ byH202 as a function of the carbonate concentration, (b): The same points replotted as a

function of the concentration of bicarbonate ions in 10 mM phosphate buffer. Pointsdenoted by crosses were obtained at pH = 11; by circles at pH = 10.2. The initialconcentration of Fe(CN)6~ was 0.8 mM, and of H202 8 mM

buffer did not markedly influence the course of the reaction, as followedspectrophotometrically. This once again indicates that side reactions are

responsible for the observed light emission.

3.7. The effect of radical scavengersVarious compounds were tested as possible inhibitors of the CL and as

selective scavengers of free radicals. The results of experiments withcarbonate buffer are summarized in Table 2. The most striking observationwas that typical scavengers of the hydroxyl radical, like /-butanol andmannitol, as well as other alcohols, did not inhibit the light emission, even

in concentrations of several mol/liter. Formate ions, also known as HO'scavengers, appreciably enhanced the observed emission, so did 10% (v/v)tert-butyl alcohol, and, to a less extent, other alcohols (data not shown).The above compounds affected only minimally the kinetics of the CL,but some of them (ascorbate, mannitol, ethanol) markedly retarded thereduction of Fe(CN)6~ by H202. Considerable quenching of the CL,together with shortening of its duration, was observed with thiourea. Inthis instance quenching occured in relatively low concentration of the


Table 2. The influence of free radical scavengers on the total emission (integral intensity),observed in the reaction of Fe(CN)6~ (1 mM)- with H202 (7 mM) in 45 mM carbonatebuffer, pH = 11.0 at 25 °C

Scavenger Molarity II, % of control

Ethanol 1.7 (10% v/v) 120 ±20Tert4outanol 0.65 135 + 20

1.30 250 + 30Acetone 0.40 110+15Mannitol 0.17 150+18

Thiourea 0.0021 39 ± 90.0083 20 ± 40.10 15+ 3

Ascorbate 0.01 58+40.02 15+3

Methanol* 2.70 49 + 10Ethanol* 1.80 31 ± 6Sodium formate 0.45 170 + 22

0.90 220 + 501.81 600 + 82

Sodium formate 2.00 300 + 45+ tert-butanol 8% v/v

* In ammonium acetate buffer pH = 9.0

scavenger, Table 2, but seemed to reach a saturation level without completeabolishing the light emission. Simultaneously, a marked acceleration of themain reaction, i.e. the reduction of Fe(CN)g~, has been observed.

Quenchers having its normal redox potential, E'0, lower than that ofFe(CN)|" (E'a = 0.36 V), e.g. ascorbic acid (£"0 = 0.06 V for n = 1 e",[22]) or radicals originating from methanol and/or ethanol (R —CH —O",R-CH-OH, R-H, CH3, E'a = -1.3 to —0.7 V, [23]) reduceFe(CN)!", thus retarding the reaction progress and changing (usuallylowering) the CL intensity.

In slightly alkaline ammonium acetate, phosphate or NaOH solutionsin the pH range 9.0

—

10.0, quenching of the CL is observed by methanoland ethanol. This renders some doubt with regard to the specificity of thesescavengers (and of thiourea too). Nevertheless, most facts give evidenceagainst involvement of OH' radicals in the chemiexcitation step of thesystem (but not in the reaction as a whole). For all the above reasons, a

possible action of superoxide 02 was tested.A significant result has been obtained by the comparison of the effect

of formate on the CL of the Fe(CN)g"—

H202-system and a chemi-luminogenic probe-luminol (Table 3). With increasing concentration offormate. 7max of the Fe(CN)e~

—

H202 system strongly increases in different


Table 3. The effect of formate ions on the chemiluminescence intensity of the Fe(CN)JH202-system and on the luminol chemiluminogenic probe

System [HCOONa] Maximum or integratedM light intensity 27/, a.u.

50 mM Na2C03 + 2 mMFe(CN)i~ + 10mMH2O2,pH = 10.4

50 mM CH3COONH4 + 2 mM Fe(CN)£"+ 10mMFJ2O2, pH = 10.4

50 mM Na2C03 + 10 mM H202+ 1 nM luminol, pH = 0.5

50mMNa2CO3 + 10mMH2O2+ 2mMFe(CN)|~ + 1 nM luminol,pH = 10.5

00.10.210.621.2300.10.210.621.2300.6200.62

28 +162 +206 +

1090 +

41 ±149 +236 ±872 +

1000 +

43 +283 +

61926

1901390+ 122

0.95183

218224

498

2645+ 35313580 + 1470

/max = maximum intensity of the chemiluminescence (peak height), arbitrary units (a.u.),•^Aomin = integrated light intensity (light sum) measured from t = 0 to t = lOmin ofthe chemiluminescent reaction, a.u., pH of the formate solutions was adjusted to that ofbuffers

buffers, approaching saturation at [HCOO~] st 1.23 M. Formate enhancesalso the integral light intensity, 27/, in alkaline solutions of H202 withoutK3Fe(CN)6, where a slow decomposition of H202 and concomitant CLmay be mediated only by active oxygen species like OFF, 02, 'Of andcarbonate radicals. Chemiluminescence of luminol has been proposed as

an assay for 02 [14, 24]. Recently, the reduction of luminol by 02 has beenused as a specific test for 02 [25]. Although there is still a controversyconcerning the detailed mechanism of light-producing reactions fromluminol and oxidants, a strong enhancement of CL by formate (Table 3)might suggest the involvement of 02 in the Fe(CN)|"

—

H202 system.Spectroscopic measurements indicate that formate ions do not change

distinctly the absorption spectrum of the reaction system. They increasethe rate of the Fe(CN)3,- reduction by a factor of about 2 (Fig. 7). Underthe same experimental conditions, the increase of the CL intensity is 4 to

6 times higher (compare with Table 3). This result again indicates that theCL is not directly related to the reduction of Fe(CN)g~.

Additionally, tiron oxidation was used as a diagnostic probe for 02.Tiron reacts specifically and sufficiently rapidly with 02 [26, 27]:

TH2 + 02 i*TH' + Or + H +

k = (1.5-5)

108 NT's-1 .

(6)


Fig. 7. The influence of 0.1 M HCOONa on the reduction rate of 1 mM K3Fe(CN)6 with10 mM H202. Solutions buffered with 50 mM CH3COONH4, pH = 10.2. a, b

-intensity of the absorption band of Fe(CN)3, ~ at 420 nm for various reaction timeswithout and with HCOONa, respectively, c-optical density at 420 nm vs. reaction time

In the system of H202 + Fe(CN)6~ + OH", tiron undergoes a rapidoxidation by Fe(CN)e~ alone with a concomitant CL flash. Therefore, theprobe could only be used in the partial system, consisting of H202 + OH ".From the data in Table 4 it is evident that 02 participates in the observedvery low CL accompanying the slow decomposition of alkaline H202. This


Table 4. The effect of the Ol-scavenger tiron, l,2-dihydroxybenzene-3,5-disulfonic acid,on the chemiluminescence intensity of the H202 —OH "-system

Solution Stationary intensity of the weakchemiluminescence, cps

Effect%

Control With 1 mM tiron

Water0.1 MH20250 mM H202+ 100 mM Na2C03100 mMH202+ 200 mM Na2CQ3

13.7 ± 2.114.7+ 1.679.1 + 3.9

165.1 ± 11.0

14.1 +3.012.0 + 2.327.0+1.0

45.3 + 4.4*

3.0-18.4-65.8

-72.6*

cps = counts per second, measured with a single photon counting mode photoelectricarrangenent, * concentration of tiron 2.5 mM

effect additionally supports the idea of the participation of carbonateradicals in the chemiexcitation step.

4. Conclusions: Proposed chemiexcitation mechanismThe findings of this study clearly exclude the emission from 'Of -dimols onthe basis of spectral data and the effect of02('zlg)-quenchers. Side reactionsof the CN-ligand with H20 or H202 (e.g. formation of aquo- or peroxy-complexes) or with oxygenated species like HO', 02, 02H etc. are hardlyprobable in the light of the spectroscopically confirmed reversibility of thereduction of Fe(CN)e~.

It is obvious that the CL rate-limiting step is the interaction betweenFe(CN)g~ and HOO~. This pH-dependent reaction is not directly relatedto the light emission, but it initiates chain reactions leading to chemiexcita-tion. The second order rate constants ofall elementary steps are not known.For the reaction of 02 with Fe3 + -cytochrome c, the rate constant is about2 • 106 M_1s_1 [28]. Since the haem iron of cyt c is buried deeply insidethe protein molecule, this rate constant should be lower than that ofFe(CN)i~ and Oi. Oxidation of Fe(CN)£" with OH' proceeds at thediffusion-limiting rate with A: = 10loM-1s_1 [29],

The stimulating effect of carbonate and formate ions on the CL providean attractive explanation for the chemiexcitation reactions. Formate ionsare known to react with OH' radicals and, subsequently, with 02 to formOi and C02 [30]:

HCOO" + OH•

4 COOT + H20k = 2.7- 109 M'V1 atpH4-10 .

coor + o2 - co2 + oi.

(7)

(8)


In this way the efficiency of 02 generation can be doubled. Reactions ofOH' and 02 radicals with carbonates/bicarbonates/lead to the generationof carbonate radicals [12]:

OH' + HCO3- A*OH" + HCO3 (9)with the second order rate constant k = 3.2- 107 M_1s_1, and

02 + HCO3- -»Or + HCOj.

(10)Recombination of 03 or COO' ion-radicals may generate energy-rich pre-cursors of electronically excited molecules or directly electronic excitation:

otto 0||02hco3 -+o=c-o-o-c=o^o= cuc=o <-» o=cnc=o

HO\OH Ha lib

H202 2C02 + hv (11)Peroxalate I dissociates to H202 and dioxetanes IIa and/or lib, which

are considered to be excited triplet dimers of two C02 molecules [12]. Spinreversion can be achieved by the emission of radiation hv. It can be regardedas a T —S transition, enhanced by the neighbourhood of the radicalelectrons. As calculated from the energy balance of the covalent bonds'formation and splitting, the last step of reaction (11) liberates AE =

E2c=0—

(Ec-c + E0-0) = 860 kJ • mol"1. Another possible mode of therecombination and decomposition involves a Russel-like tetraoxide tran-sient complex:

0 = C^> C = 0 - 0 = C = 0 + 02 + H-C = 0 (12)X0-H-0/ IXH OH

For this mode the energy difference is AE = Ec = 0 + Ec = 0 + Ec-H—

(E2c-0+ -Eo_h) = 305 kJ • mol"1 that corresponds to the shortest wavelength386 nm.

The recombination of formate ion-radicals, e.g. from reaction (7) maygenerate directly excited oxalate dianions:

2COO" -> [(COO)2"]* (13)with AE = 360 kJ • mol"1 high enough to promote emission with thewavelength ^ 325 nm.

Another possibility which cannot be entirely excluded is the formationof the peroxide-ferri/ferrocyanide complex in the conformational out-of-equilibrium state:

Fe3+ + OOH" - [Fe2+OOH] [Fe3+OOH"] -> Fe2 + + H+ + 02 .

(14)


This state would result from the perturbation of the rhombic symmetry ofthe octahedral Fe(CN)|~/4~ -ion, similar to the Jahn-Teller distortion. Be-cause of the repulsive interaction between two negative ions, the transientcomplex would be unstable. This can facilitate an electron transfer fromthe filled n* orbital of H202 to the empty 3d orbital of Fe3 +

.

Finally, thecomplex dissociates to Fe(CN)4,- and an 02 radical. The change of theGibbs' free energy, AG, of the non-equilibrium complex to the equilibriums|ate of products should be greater than that calculated from theequilibrium-redox potential (AG0 = 132 kJ • mol-1). As it was shownpreviously (Fig. 1) the absorption spectrum of Fe(CN)|~ displays a veryweak band at 422 nm (at pH = 7 at 418 nm) assigned to the spin-forbiddenligand transition lAlg -» 3Flg

d

= 282 kJ-mol-1

metal orbitals molecular orbitals of the octahedral complexThe AE value is very close to that observed in the CL and fluorescerJce

spectra (260 kJ•

mol-1, corresponding to /lmax = 460 nm). Therefore it ispossible that the out-of-equilibrium state of the peroxy-complex producedin reaction (14) is simultaneously an electronic excited state, i.e. the emitterof the CL. The measured low values of the quantum efficiency of the CL,0 ~ 10"10, and that of luminescence, <PL < 10"3, would agree well witha very low intensity of the photoluminescence and the suppression of the422 nm absorption band. The formation of out-of equilibrium conforma-tional states was observed in the fast electron-transfer reaction betweenferri-complexes and suitable donors [31]. Recently, a weak emission in thenear infrared part of the spectrum from the reaction of catalase and H202has been alternatively explained by the forbidden transition in the haemcomplex [32].

In conclusion, the pH-dependent, reversible oxidation/reduction ofFe(CN)t"/3" reveals many features analogous to the Haber-Weiss andFenton reactions. However, it is not directly related to the chemiexcitationand light emission. The most adequate explanation of the chemiexcitationmechanism is the 02-mediated formation ofcarbonate and formate radicalswith their subsequent recombination. A 25 fold increase of 7max of the CLup to saturation with increasing concentrations of HCOO" ions cannothowever entirely account for the contribution of the reactions (7) and (8)(see Table 3). It cannot be excluded that traces of impurities affect the CL.This emphasizes the importance of using especially pure reagents whenworking with systems generating weak luminescence.


AcknowledgementsThe research was supported by the Polish Academy of Sciences underproject MR 1.9.

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Carbonate and Superoxide Ion-Radicals as Intermediates in the Chemiluminescence of the Ferricyanide — Hydrogen Peroxide Redox System