Radiation Chemistry - ACS Publications

9 Pulse Radiolysis of Alkaline Solutions JOSEPH R A B A N I

Downloaded by YORK UNIV on November 9, 2012 | http://pubs.acs.org Publication Date: January 1, 1968 | doi: 10.1021/ba-1968-0081.ch009

Department of Physical Chemistry, The Hebrew University, Jerusalem, Israel

The absorption spectrum of O , with a peak at 2400A. (ε = 240M cm. ) is reported. Extinction coefficients at 2600A. of intermediates produced in alkaline solutions are reported. εO 2600 = 200, εO 2600 = 1000 at neutral pH, 1100 in alkaline solutions containing both N O and H O , and 1900 and 1600 in aerated and oxygenated solutions respectively (all at pH = 13, in units of M cm. ). The possible forma tion of species equivalent to O2- and having different ab sorptions is discussed. The effect of pH on the apparent reaction rate constant of hydroxyl radicals with hydrogen peroxide was studied. We find: k = 2.74 X 10 M sec. for zero ionic strength, and (k + 1.42 k ) = 1.18 X 10 M sec. . The error in the above rate constants and extinction coefficients is 10-15%. -

-1

-1

-

-

2

-1

2

2

-1

8

(O¯+HO2¯)

-1

-1

(O¯+H2O2¯)

10

(OH+HO2¯)

-1

-1

"TQulse radiolysis of N 0 alkaline solutions has not yet been interpreted. Rabani and Matheson (33) reported an optical absorption at 2600A. in N 0 aqueous solutions at both neutral and alkaline pH's. They suggested that this absorption might be caused by O H i n neutral and O " i n alkaline solutions. In neutral solutions, the absorption decayed away and the rate could be accounted for by reactions between the products listed on the right hand side of Equation 1 and Reaction 2. 2

2

H 0 — M — > h , e- , O H , H 0 \ OH", H 0 , H 2

aq

3

2

2

2

H 0 e'm + 2 θ -> N + O H + O H "

(1)

2

Ν

2

(2)

The absorption at p H = 13 lasted several seconds, and four forma tion and decay steps could be separated. Demonstration of the observa tions at p H = 13, at both 2600 and 4300A. has been presented i n a previous paper (30). 131 In Radiation Chemistry; Hart, E.; Advances in Chemistry; American Chemical Society: Washington, DC, 1968.

132

RADIATION CHEMISTRY 1

The present work has been carried out with the aim of obtaining more data on possible intermediates i n pulsed alkaline solutions. Data concerning optical absorptions and reactivities of these intermediates is essential for the future interpretation of irradiated alkaline N 0 aqueous solutions. 2


Experimental The pulse radiolysis apparatus (15, 17, 27) and syringe technique (17, 20, 39) for cell filling has been previously described. A 450-w Osram xenon lamp used to produce the analyzing light beam (18). Irradiation temperature was 24 °C. A four-cm. cell with an 8 cm. light path was used, instead of the multiple reflection cell (35). This was done to avoid interference from transients formed i n the cell house water. This is important when relatively small optical absorptions are investigated, especially below 3000A. Below 2800A., scattered light passed through the monochromator and was collected by the photomultiplier. The amount of scattered light depended on the wavelength and on the optical align ment. Thus, at 2600A. it was usually between 3 - 6 % , rising to about 8 % at 2500A., 10-15% at 2400A., 20-25% at 2300A., and 50% at 2220A. The scattered light was checked using a Corning glass filter which is supposed to cut all the light below the wavelength region, which is just higher than the wavelength used. Alkaline solutions, carbonate free, have been prepared using either B a ( O H ) - N a O H mixtures, or from 2 0 M N a O H stock solutions i n which carbonate is relatively unsoluble. Unless otherwise stated, [ B a ( O H ) ] / [ N a O H ] = 0.038. N 0 has been purified as described before (34). For the spectrum of O " the light beam was split into two, each beam passed through a Baush and L o m b mono chromator. One of the monochromators has been set on 2600A. and used as a monitor for the electron beam intensity, while the other was used to obtain the absorption of the pulsed solution at various wavelengths (41). 2

2

2

The stock alkali solutions have been stored in borosilicate glass bottles as long as seven days. The stock ferrocyanide solutions were stocked in the dark for as long as 30 hours in a refrigerator and as long as 12 hours at room temperature before use. Hydrogen peroxide was un stable in our alkaline solutions, the decomposition was enhanced by alkali up to the highest alkali concentrations used ( 0.3M ). The rate of decom position was not reproducible. Therefore, a neutral stock solution of H 0 was prepared ( ^ 1 χ 1 0 " M ) . Measurements showed that only 3 % of the stock H 0 decomposed during the 4-5 hours of storage of this solu tion. The microsyringe injection technique (20) was used to obtain alkaline solutions with the desired H 0 concentrations immediately before pulsing. Analysis of the H 0 concentrations was carried out immediately before and after pulse, for each sample. Neutral H 0 solutions were analyzed by mixing with 3 X 10~ M F e S 0 in 0.8ZV H S 0 . The F e formed by the oxidation of F e was determined spectrophotometrically. Alkaline solutions were made acid by mixing with 0.8N H S 0 and then transferred to a volumetric flask containing 3 X 10" M F e S 0 in 0.8N H S 0 . Since all the alkaline solutions used for the investigations 2

2

2

2

2

2

2

2

2

2

3

3+

4

2

2

4

2+

2

3

2

4

In Radiation Chemistry; Hart, E.; Advances in Chemistry; American Chemical Society: Washington, DC, 1968.

4

4

9.

RABANI

Alkaline Solutions

133

of H 0 contained B a ions, B a S 0 precipitated. After several hours (usually overnight) the solution could be separated from B a S 0 by décantation. The H 0 concentrations were calculated from the ferric sulfate absorption at 3040A., taking e = 2190 at 24 °C. (36). A n electron pulse of 1 μ-sec. duration was used i n all the experiments. In the following, time zero w i l l always be chosen at the beginning of the electron pulse, unless otherwise stated. 2

2 +

2

4

4

2

2


The Absorption Spectrum of O " Presentation of Data. In Figures l a and l b , we present the optical absorption spectrum obtained in pulsed 0.105M N a O H solutions in the presence of 0.05 atmospheres N 0 . Experiments in a solution of both N a O H and B a ( O H ) , [ B a ( O H ) ] / [ N a O H ] = 0.035, but with the same [ O H ] gave similar results. During the electron pulse, optical ab sorption is formed in the region 2200-2800A. Within the first 100 ^sec. after the pulse the optical density increases (half-life ^ 2 5 /xsec. under the conditions of Figure 1) by 25% or less. Then, part of the optical density decays away within e-

(3)

m

W h e n N 0 is present, the hydrated electrons formed by Equation 1 and Reaction 3 are converted into (13, 24) O H or O", k (14,17, 25) = 8.6 X WMsec." . 2

4

1

1

+ N 0 -> N + Ο"

e~

2

m

(4)

2

Ο" + H 0 ^± O H + O H "

(5)


2

Thus, in 0.05 atms. N 0 , at p H = 13, only O H , O", and H 0 " contribute significantly to the optical absorption below 3000A. The absorption spectrum of H 0 (23), H 0 " (23), and O H (30, 40), have been studied previously. In Figure 2, we present the optical absorptions of hydrogen peroxide ( H 0 in equilibrium with H 0 ) , O H and O , the sum of which equals the absorption spectrum of Figure 1 Curve a. The absorption attributed to hydrogen peroxide was calculated from the spectrum taken with a Cary spectrophotometer in a solution containing 0.105M N a O H and 1.34 X 1 0 M hydrogen peroxide. Our results were in agreement with those obtained by Jortner and Stein (23). The total concentration of hydrogen peroxide ( H 0 and H 0 " ) in the pulse irradiated solutions was calculated to be 2.36 X 10" M, based on G (hydrogen peroxide total) = 0.7. 2

2

2

2

2

2

_

2

-

2

_3

2

2

2

6

The optical density attributed to O H radicals ( Figure 2 ) was calcu lated from the previous work (40). Above 2400A., the results obtained in H 0 , N 0 , or acid solutions were consistent (30, 40). Below 2400A. the results obtained in H 0 and N 0 differed from those obtained in H C 1 0 (24, 25). The spectrum obtained in N 0 and H 0 represents the absorption of O H radicals (25). The absorption of O H in Figure 2 was obtained from these results, after correcting for H 0 formation or destruction in the N 0 and H 0 solutions respectively. ( This correction is quite small and was neglected previously). 2

2

2

2

2

2

4

2

2

2

2

2

2

2

2

( G ( O H ) + G ( 0 " ) ) = G O H + G + G = 5.8 and G o = 0.7 were used for the calculation of Figure 2, together with our values for the extinction coefficient of O H , €OH = 4 1 0 M c m . and ^hydrogen peroxide = 177 ( p H = 13). The pulsed solutions used to obtain Figures 1 and 2, contained 1.96 X 1 0 ' M ( O H + O") and 2.36 X 10" M hydrogen peroxide. Of the oxidizing radicals, 1.29 X 10" M were O H , and 1.83 X 1 0 - M were O " radical ions. The optical density scale in Figure 2 can be easily converted to an extinction coefficient scale for O", using e - = 2 0 0 M cm." at 2600A. Later we shall provide the data used for the evaluation of € - at 2600A. c

H

H 2

1

5

2

1

6

6

0

1

0


-1

136

RADIATION CHEMISTRY

DxlO

2

ο

3

>\ —


1

•J

0"

Γ

. H 0 +HOJ 2

2

OH.

2200

1

1 2400

1 2600

1

2800

3000 λ(Α)

Figure 2.

The absorption spectra of O', hydrogen peroxide, and OH radicals atpH = 13

The absorption of O " differs considerably from that of O H . A peak at 2400A. is obtained for O", with e - = 240M" cm." . As to O H , no peak is observed down to 2100A. (30, 40), although a broad peak between 2100 and 2400A. cannot be excluded. The difference i n absorption at 2600A. between O H and O " has been used for the evaluation of pKoH (32). 1

0

1

Evidence for the Identification of the Spectrum as Owing to O". Figure 2 shows that O H , H 0 " , and H 0 cannot account for the entire optical absorption formed i n the ultraviolet i n pulsed alkaline solutions containing N 0 . Of the species on the right hand side of Equation 1, H 0 , O H " , and H are ruled out because their optical absorptions are known to be negligible i n the range used i n Figures 1 and 2, Η atoms have a life time of about 0.35 /xsec. at p H = 13, and hydrated electrons are converted into O H and O " according to Equation 2 i n less than a microsecond. Thus, of the species on the right hand side of Equation 1 (including their alkaline forms), only O " remains as qualified to account for the major part of the absorption. 2

2

2

2

3

+

2

The change of initial optical density with p H , i n pulsed aqueous N 0 solutions, agrees with the conversion of O H into O " i n alkaline solutions (32). 2


9.

RABANI

137

Alkaline Solutions

The formation of absorption during the electron pulse shows that the absorption is caused by a primary product or a product formed from N 0 . The possibility of reaction with impurities is unlikely to produce absorption within the 1 /xsec. pulse. Only O " and N 0 " (or N O H ) remain as qualified species. There is strong evidence against N 0 " (or N O H ) , and those have been summarized i n previous publications (32, 34, 41). W e do not have an alternative satisfactory interpretation for each of the previous studies which suggested N 0 " (or N O H ) being relatively stable intermediates in N 0 irradiated solutions (5, 6, 7). However, in our opinion, the evidence against a long life time of N 0 " (or N O H ) in both neutral and alkaline solutions is very strong, and it is highly unlikely that N O H or N 0 ~ will live more than 1 /xsec. It can be argued that the decay of the optical absorption i n alkaline solutions does not agree with the expected rate of O " decay. Indeed, as seen from Figure 1, there is only a partial decay within the first 4 msec, and the time dependency of the optical density is not i n agreement with the assumption that O " absorbs initially and then decays to form nonabsorbing products. The complexity of the kinetics was demonstrated in a previous publication (32) (see Figure 2). As w i l l be discussed later, the mechanism for O " decay involves several possible intermediates, some of which are known to have optical absorptions in the ultraviolet. Thus, the long life time of the absorption may be caused by the conversion of O" into other, more stable, absorbing species. 2

2

2

2

2

2

2

2

2


2

2

2

In a solution containing 0.1M C 0 and 5 X 10~ atm. N 0 at p H = 13, the initial optical density after a 1 /xsec. electron pulse was 0.035. This is about 15 per cent higher than the absorption obtained with a similar solution containing no carbonate. However, in the carbonate solution, at least 85 per cent of the absorption decayed away within 100 msec, and assuming the reaction was second order, 2k = 5 X 10~ M sec." has been obtained, i n agreement with previous results (2,41). The small change i n initial optical density, when 0.1M C 0 is added to the alkaline solution, shows that C 0 " and O " have comparable absorptions at 2600A. Since carbonate ions react as specific scavengers for O H radicals (2, 41) and adding the carbonate eliminated the complicated kinetic sequence which was found in carbonate free solutions, we conclude that O H (or O") are the precursors of the above long living species. Therefore, the absence of a fast decay of the optical density corresponding to k -+o-) = 1 0 M s e c (34) is not evidence against the assignment of the major part of the initial absorption in pulse irradiated N 0 alkaline solutions because of O " radical ions. In Figure 3, we present oscilloscope traces obtained i n the presence (Figures 3c and 3d) and absence (Figures 3a and 3b) of N 0 . These traces were photographed on Polaroid pictures, and then projected and 3

2 _

2

2

7

3

_1

1

2 _

3

i0

9

- 1

-1

2

2


138

RADIATION CHEMISTRY

1

plotted on optical density paper which enables the direct reading of optical densities. The electron beam intensity was the same for Figures 3a and 3b and Figures 3c and 3d. The initial ( extrapolated to the middle of the electron pulse) optical density i n the N 0 solutions is one half of that obtained i n N 0 free solutions under the same conditions. This is owing to the hydrated electrons (e = 600M" c m . at 2600A.) (19). 2

2

1


.00 .01 .02 .03 .04

.00

ë

.01

.

/ 2600 A

1

_L

1

-\

I

I

I

(b)

1 •4msec

50 usee

.03 .04

(a)

20 Hjjsecl

2600 A

> .oo U) .01

0 - + H 0 2

2

When a sufficient concentration of hydrogen peroxide is initially present, Reactions 6 to 9 are the only important reactions taking place in the


140

RADIATION CHEMISTRY

1

microsecond range. Reactions such as 10, 11, and 12 have only a small contribution to the overall rate of hydroxyl radicals decay. OH + O H ^ H 0

2

(10)

O H + Ο" -> H 0 -

(11)

2

2

H 0 2

Ο" + Ο"

H 0 - or 0 " 2

2

(12)

2

The apparent rate of reaction of hydroxyl radicals with hydrogen per oxide can be obtained from the increase of optical density with time at 2600A. because of the formation of 0 " by Reactions 6 to 9. The absorption of 0 ~ in the ultraviolet light has been reported by Czapski and Dorfman ( I I ) . Typical oscilloscope traces, projected and replotted on O . D . paper and a plot of D i eaii — D vs. t, using a semilogarithmic scale are shown in Figures 4 and 5 respectively. ( D i teau ~ the optical density obtained, after all the O H and O " have reacted with the hydrogen peroxide, D optical density at time t ). The straight line in Figure 5 is typical of a first order reaction, and when the slope is divided by ( [ H 0 ] + [ H 0 ~ ] ), the apparent rate constant of the reaction between hydroxyl radicals and hydrogen peroxide is obtained. Corrections which were made to Obtain a Precise Rate Constant for the Reaction of Hydroxyl Radicals and Hydrogen Peroxide. (1) Reac tions such as 10, 11, 12, 13, and 14 compete with Reactions 6, 7, 8, and 9. 2


2

P

at

t

p a

t

2

2

2

OH + 0 - -» 0

2

+ OH"

(13)

O-

2

+ 20H-

(14)

2

+ 0 - -> 0 2

or 0 " 3

2

or H 0 - + O H 3

The nature and the reactivity of some of the intermediates towards hydroxyl radicals and radical ions are not known. However, the effective overall rate constants for O H (and O") reactions in a similar system containing ferrocyanide ions instead of hydrogen peroxide have been investigated (34). Using the results of the ferrocyanide work (34), we corrected the apparent rate constants for the reaction of hydroxyl radicals with hydrogen peroxide by adding the calculated first half-life of hydroxyl radical's decay in the absence of added hydrogen peroxide, to that measured in our hydrogen peroxide solutions. (2) A correction was made to account for the decomposition of H 0 before pulsing the alkaline solutions. W e assumed that over the short time periods in which hydrogen peroxide was kept alkaline ( usually less than 15 minutes), the decomposition was linear with time. In most cases, the differences between the concentrations of hydrogen peroxide found by analysis and those calculated from the concentrations of the 2

2


9.

RABANI

141

Alkaline Solutions

neutral stock solutions, did not exceed 10%. In all cases it was less than 20%. Thus, the concentration of hydrogen peroxide was chosen as the average between the values obtained by analysis before and after the electron pulse, after correcting for the change in hydrogen peroxide con centration by the pulse.

.00 .01

(α)


.02

\J

2600 A

μ sec

I

.03

if) Ζ

/ fc(o-+Ho -) = 30 (for zero ionic strength) may be compared with & 6-6 reported by Adams et al. for p H = 13 ( J ) . (Both H 0 " and C N S " have the same negative charge.) Although cor rections for the ionic strength effect at p H = 13 cannot be precise, if we make the same corrections for O" + C N S " as made for O " + H 0 " , we obtain &(OH CNS-)/&(O- + CNS-) = 12 for zero ionic strength, which is comparable with the analogous value of 30 found in the H 0 " case. 8

2


+

2

2

1

2

2

1

2

2

2

7

(0

+

2

1

1

(0

+

2

a

2

=

2

2

+

2

Acknowledgment The author wishes to thank M . S. Matheson for helpful discussions during the early phases of this work. Literature Cited (1) Adams, G. E., Boag, J. W., Currant, J., Michael, B. D . , "Pulse Radiolysis," p. 117, Acad. Press, London, 1965. (2) Adams, G . E., Boag, J. W., Michael, B. D., Proc. Roy. Soc. 289, 321 (1966). (3) Adams, G. E., Boag, J. W., Michael, B. D., Nature 205, 898 (1965). (4) Adams, G. E., Boag, J. W., Michael, B. D., Trans. Faraday Soc. 61, 492 (1965). (5) Anbar, M . , Munoz, R. Α., Rona, P., J. Phys. Chem. 67, 2708 (1963). (6) Asmus, K. D., Henglein, Α., Beck, G., Ber. Bunsengesellschaft 70, 459 (1966). (7) Baxendale, J. H . , Fielden, Ε. M . , Keene, J. P., "Pulse Radiolysis," p. 207, Acad. Press, London, 1965. (8) Benson, S. W., "The Foundations of Chemical Kinetics," McGraw-Hill Book Co., Inc., New York, Ν. Y., 1960.


152

RADIATION CHEMISTRY

1

(9) Czapski, G., A D V A N . C H E M . SER. 81, 106 (1968).

(10) (11) (12) (13) (14) (15)


(16) (17) (18) (19) (20) (21) (22) (23) (24) (25) (26) (27) (28) (29) (30)

Czapski,G.,Bielski, Β. H . J., J. Phys. Chem. 67, 2180 (1963). Czapski, G., Dorfman, L . M., J. Phys. Chem. 68, 1169 (1964). Czapski, G., Behar, D. (private communication). Dainton, F. S., Peterson, D . B., Nature 186, 878 (1960). Dorfman, L . M . , Matheson, M . S., "Progress in Reaction Kinetics," Vol. 3, G. Porter, ed., Pergamon Press, New York, 1965. Dorfman, L . M., Taub, I. Α., Bühler, R. Ε., J. Chem. Phys. 36, 3051 (1962). Felix, W . D., Gall, L . , Dorfman, L . M., J. Phys. Chem. 71, 384 (1967). Gordon, S., Hart, E . J., Matheson, M . S., Rabani, J., Thomas, J. K., Dis cussions Faraday Soc. 36, 193 (1963). Gordon, S., Hart, E . J., Thomas, J. K., J. Phys. Chem. 68, 1262 (1964). Hart, E . J. (private communication). Hart, E . J., Gordon, S., Thomas, J. K., J. Phys. Chem. 68, 1271 (1964). Hochanadel, C. J., Radiation Res. 17, 286 (1962). Jortner, J., Rabani, J., J. Phys. Chem. 66, 2081 (1962). Jortner, J., Stein, G., Bull. Res. Counc. Israel 6A, 239 (1957). Jortner, J., Ottolenghi, M . , Stein, G., J. Phys. Chem. 66, 2037 (1962). Keene, J. P., Radiation Res. 22, 1 (1964). Klotz, I. M . , "Chemical Thermodynamics," p. 21, Englewood Cliffs, New Jersey, 1958. Matheson, M . S., Dorfman, L. M . , J. Chem. Phys. 32, 1870 (1960). Matheson, M . S., Mulac, W . Α., Weeks, J. L . , Rabani, J., J. Phys. Chem. 70, 2092 (1966). Matheson, M . S., Rabani, J., J. Phys. Chem. 69, 1324 (1965). Nielsen, S. O. (private communication).

(31) Rabani, J., A D V A N . C H E M . SER. 50, 242 (1965).

(32) (33) (34) (35) (36) (37) (38) (39)

Rabani, J., Farkas Symp., Jerusalem, Israel, Dec., 1967 (in press). Rabani, J., Matheson, M . S., J. Am. Chem. Soc. 86, 3175 (1964). Rabani, J., Matheson, M . S., J. Phys. Chem. 70, 761 (1966). Rabani, J., Mulac, W . Α., Matheson, M . S., J. Phys. Chem. 69, 53 (1965). Schaft, K., Lee, R. M . , Radiation Res. 16, 115 (1962). Schmidt, Κ. H . , Argonne Natl. Lab. Rept. 7199 (1966). Schwarz, H . Α., J. Phys. Chem. 66, 255 (1962). Senvar, C., Hart, E . J., Proc. 2nd Intern. Conf. Peaceful Uses At. Energy (Geneva) 29, 19 (1958). (40) Thomas, J. K., Rabani, J., Matheson, M . S., Hart, E. J., Gordon, S.,J.Am. Chem. Soc. 70, 2409 (1966). (41) Weeks, J. L . , Rabani, J., J. Phys. Chem. 70, 2100 (1966).

RECEIVED January 25, 1968. The experimental work was carried out at the chemistry division, Argonne National Laboratory, Argonne, Ill., under the auspices of the U . S. Atomic Energy Commission.


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