70
A. R. ASDERSON ASD ED%-IN J. ZLLRT
Vol. 66
RADIATION CHEMISTRY OF WATER WITH PULSED HIGH INTENSITY ELECTRON BEAMS’ BY4.R. ANDERSON AND EDWIN J. HART Chemistry Division, Argonne National Laboratory, Argonne, Illinois Received J u l y 6, 1861
The radiation chemistry of water, aqueous ferrous sulfate, hydrogen peroxide and formic acid has been studied using pulsed electron beams from the Argonne linear electron accelerator. Wit,h a pulse length of -1.4 fisec. and an electron energy of -15 MeV., the dose rate received by the solutions during a pulse is equivalent to -2 x 102s e.v. ml.+ set.-+. Under these conditions, -10-6 M hydrogen atoms (and hydroxyl radicals) are generated throughout the irradiated volume. All irradiations are monitored with the Fricke dosimeter, for which G(Fea*) is 11.4 k 0.5 under these irradiation conditions. Preliminary studies with neutral water and with 0.8 N sulfuric acid solutions show that the initial yields of both hydrogen and hydrogen peroxide are about 80% higher in the acid solution but that, in each case, G(H2) = G( H2Q2). Oxygen is not an initial product. Yields of hydrogen have been measuredin 0.8 N sulfuric acid up to 8 M hydrogen peroxide and up to 0.1 M in neutral solution. Hydrogen peroxide exerts a much stronger scavenging action in neutral than in acid solution. In each case, but a t different concentrations of hydrogen peroxide, the scavenging curves eventually coincide -xith those determined using Coeoy-radiation a t a dose rate of e.v. mI.+ sec.-l. Although difficulties remain, these data are consistent with the transient existence of two reducing species designated as the solvated electron, e-ast predominant in neutral solutions, and the H atom, predominant in acid solution. Hydrogen peroxide is 50 times more effective in suppressing hydrogen formation i n neutral solutions than in 0.8 11‘ sulfuric acid solutions. Some relative hydrogen atom (or e- ) and hydroxyl radical rate constants consistent Li-ith the G(H2)and G(HzQ2) are reported. In contrast to the lower G(FeS$l obtained with y 0 2 ) give the same ferrous sulfate, measurements Rith the formic acid dosimeter (0.01 M HCQOH, 0.001 N H2S04,0.001 l yields as with (2060 y-radiation a t much lower dose rates. G(H202)for oxygen saturated neutral water is higher than that for y-rays. This greater yield can be explained on the basis of competing radical-radical reactions.
Introduction Current theory holds that the molecular products, hydrogen and hydrogen peroxide) formed in the radiolysis of water originate from H and OH radicals produced initially in high local concentrations along the track of the ionizing radiation. With yradiation or with fast electrons the radicals foim in localized clusters or spurs, with the intei-radical or radical-solute reaction taking place during the expansion of the spurs. At the dose rates presently available from y-ray sources or from X-ray and van de Graaff machines, alteration of the dose rate produces no quantitative chemical changes as the distance between individual spurs is so large that the radicals formed in any one spur react with solute molecules before they diffuse to the reaction zone of a neighboring spur. With electron linear accelerators, however, the dose rate during a pulse can be as much as lo6 greater than that from the most powerful y-ray source and a condition is produced in the solution where inter-track reactions can be observed. The general chemical manifestation of the higher dose rate will be the enhancement of products from inter-spur reactions a t the expense of the radical-solute reactions. Previous observations of such effects have been made by Brasch, Huber and Waly,2 by Hutchinby Rotblat and Sutton4and by Glazunov and Pikayev. In our work we have investigated the radiolysis of pure water, of 0.8 N HzS04, and of solutions containing oxygen, formic acid, hydrogen and hydropen peroxide, a t dose rates during a pulse corresponding to 2 x 10z5 e.v. rnl.-l sec.-l. The irradiations were carried out with the Argonne elec(1) Based on work performed under the auspices of the U. S. Atomic Energy Commission. (2) A. Brasch, I&’. Huber and A. Waly, Arch. Baochem. Baophys., 39, 245 (1952); A. Braach and W. Huber, Sczence, 105, 112 (1947). (3) F. Hutohinson, Radaatzon Research, 9, 13 (1957). (4) (a) J. Rotblat and H. C. Sutton, Nature, 180, 1332 (1967); (b) Proe. R o y . Soe. (London), A 265, 490 (1960). ( 5 ) P. Ya. Glasunov a n d A. I*. Pikayev, Dokladv d k a d . N a u k S.S.S.R., 130, 1051 (1900).
tron linear accelerator, using 1.4 ,usee. pulses of 15 MeV. electrons. Under these conditions we generate throughout the irradiated volume a uniform concentration of 12 to 15 ,ufVof hydrogen atoms and of hydroxyl radicals in a time that is short compared to appreciable radical-radical reaction, which is assumed to proceed with a rate constant of 6 X lo9 1. mole-1 sec.-l (10-11 cc. molecule-l sec.-l>.* Spur reactions, on the other hand, are substantially over in 10-9 second,’ so they are complete a t the end of the pulse. Experimental Preparation of Solution.-A modification of the “gyrin.ge technique” described previously8 was used in our irradlations. Syringe A of Fig. 1 containing the degassed solution was the actual irradiation cell. After evacuation, the solution from chamber “B” was forced by argon pressvre into the syringe which previously had been flushed wlth argon. By this technique the concentration of oxygen in solution could be maintained readily a t less than 1 &f. Solutions equilibrated with oxygen or hydrogen in the evacuation chamber were forced into the syringe using a pressure of the same gas. Triply distilled water was used for all the solutions; hydrogen peroxide solutions were made from 98% HzOz (Buffalo Electro Chem. Corp.) without further purification; foimic acid was purified as desciibed elsewhereg; oxygen was purified by a partial fractional distillation while high purjty hydrogen and argon were used directly from the gas cylinders after passage through a liquid nitrogen trap. The sohtions for dosimetry all were aerated (0.001 M FeS04,0.001 i$f NaCl in 0.8 hi HZSO,). Analysis.-Hydrogen was analyzed chromatographically on a Molecular Sieve column (No. 13X, Wilkens Instrument and Research Go., Walnut Creek, California), using a technique developed in this Laboratory.10 The gas was extracted from solution in a Van Slyke gas analysis apparatus, argon added and the gas mixture forced into a reservoir from where it was carried in a stream of argon to the column. This technique was checked using the mass spectrometer and the t v o methods agreed to within 2%. Oxygen also could be estimated in a similar way although the sensitivity of the (6) H. Fricke, Ann. N . Y . Acad. Sci., 59, 567 (1955). (7) P. J. Dyne and J. h l . Kennedy, Can. J . Chem., 36,1518 (1958). (8) E. J. Hart, S. Gordon and D. A. Hutchinson, J . Am. Chem. Soc., 76, 6165 (1953). (9) E. J. Hart, ibid., 7 6 , 4312 (1954). (10) 8. Gordon and G. E. Adams, private oommuniaation.
Jan., 1962
71
RADIATION CHEMISTRY OF WATERWITH PULSED ELECTRON BE~MS
method using argon as a carrier gas was only about 10% of that for hydrogen. Hydrogen calibrations were carried out periodically during each series of hydrogen analyses and the sensitivity, oxpressed in terms of recorder peak area, was 200 cm.2 per micromole, Thus 0.025 micromole of hydrogen greater could be analyzed readily. At concentrations of HZOZ than 0.1 M , the gas was extracted from solution on a vacuum line, without the solution contacting mercury, and toeppled into an evacuated gas bulb. The oxygen was extracted bsr absorption in alkaline sodium hydrosulfite in the van Slyke apparatus and the hydrogen determined as described above. Gas analysis for formic acid solutions was carried out by extracting the gas in the van Slyke apparatus, absorbing the COz in potassium hydroxide solution, absorbing the oxygen in alkaline sodium hydrosulfite and determining the residual hydrogen by gas chromatography. Hydrogen peroxide in neutral solution and in weakly acid formic acid solution was analyzed using the triiodide method"; the molar extinction coefficient a t a wave length of 3500 A. is taken as 2.59 x lo4, a t 25'. In acid solution the hydrogen peroxide was analyzed by oxidizing a 3 mM FeSO, solution and determining the ferric ion produced by its absorption at 304 mp; the molar extinction coefficient 1s 2225 at 25' and the temperature coefficient of 0.7% per degree. For hydrogen peroxide concentrations k ~ kB, , k4. The initial slope of the hydrogen and hydrogen peroxide dose curves will be independent of dose rate, but the concentration of molecular products built up before deviation from linearity is observed will depend not only on the relative concentration of radicals and molecular products in the bulk of the solution but also on the relative rate constants for reaciions 2, 3, 4, 5, 7 and 8. As the concentration of radicals is proportional to dose rate, increasing the dose rate ill simply extend the linear part of the productdose curve to higher concentrations of molecular products. This argument implies that the linear part of the dose curve exists at the low y-ray dose rates but this linearity cannot be observed for yrays as deviation occurs at molecular product concentrations less than 1 H M . At very high dose rates, however, the important initial linear portion of the curve is extended so that reliable initial yields can be measured. Table IV gives the theoretical yields obtained (19) H. A. Sahwarz, J. M, Caffrey, Jr., and G. Scholes, J. Am. Chsm. Soe., 81, 1801 (1959). (20) A. R. dnderson and E. J. Hart, Radiation Research, 14, 688 (1961). (21) ( a ) N. F. Barr and R. H. Schuler, {bid., 7, 302 (1957); (b) J . Phya. Chem.. 68,808 (1958).
dt
= -2k7(OH)2
- ks(H) (OH) + ez(t)
In these equations,22 el(t) and ez(t) represent the rate of H and OH radical generation during an electron pulse assumed to be rectangular.
and (HzOz) =
k.i(OH)*
These equations have been solved on the analog computer for the radiation yields2lbVZ3 g(H), g(OH), g(H2) and g(H202), given in Table IV applying to neutral and acid solutions, The free radical rate constants used vary about the zero diffusion controlled activation energy rate constant of lo-" cc. molecule-' see.-1 calculated by Fricke.* The agreement between experiment and theory is close enough to conclude that the relative rate constants selected for kg, k7 and k8 are satisfactory for acid solution, I n neutral solution, however, the large divergence betmeen experimental and calculated yields indicates that the pertinent hydrogen forming reactions represented by (5) are very slow or that reaction 8 specifically written as e-aq
+ OH +OH-
(8')
is very high. This conclusion is derived from results in Table IV. Here the low experimental yields in neutral solutions can be approximated by assuming k6 equal to zero and ks' equal to 6k7. Alternatively if k8' is 1Ok5, G ( H 2 ) a a l o d . approaches G(J%)exp.
Measurements with neutral solutions of 0.001 M give G(Hz02) = 1.98 compared to 1.23 for Go6" y-rays a t much lower dose rates. I n the presence of oxygen and low hydrogen peroxide we have reactions 9, 10 and 11 0 2
H
OH
+
0 2
+ HOz
+ HOz +HzO + 2HOz
--ic
(9) 0 2
HzOz
(10)
(11)
in addition to 7 and 8. By assuming that all II disappears by (9), reaction 5 can be eliminated and the equations for hydrogen peroxide formation become : d(oH) dt
c
_
dt
- ~ ~ ? ( o H-) kz l o ( ~ o(OH) z ) + ez(t)
=
=
-kl0(HOp)(OH) - 2kll(H02)2
+ edt)
and (H202) =
[k.i(OH)'
+ k~~(HOz)~ldt
Since the dosage curves are linear, secondary reactions of hydrogen peroxide are unimportant. (22) D. L. Phillips, Argonne National Laboratory, has developed an analytical solution for these equation8 for an infinitely short pulse, Applied Mathematios Division Memorandum No. 17, Feb. 8, 1961. (23) (a) P. V. Phung and M. Burton, Radiation Eeaearcfi, 71 199 (1957); (b) E*J. Hart, J. Ami Cfiema $OS,, 76,4198 (1954);
Jail,, 1962
75
2-AMIXOETHAKOL AND ETHYLENEDI-4MINE COMPLEXES
TABLEIV CONPARISOX OF EXPERIMENTAL AXD CALCULATED G( &Os) AKD G(&) FOR PVI43ED ELECTRON IRRADIATIONS I N PTATER AND 0.8 N SULFURICACID g(H)
a
g(0H)
3.66
2.95
3.16 3.16 3.05
2.64 2.64 2.91
3.05 3.05
2.91 2.91
g(Hn)
0.45 a45 .39 .39 .41 .41 .41 .41 .41 .41 .39
~(HZOE)
0.80alb .80 .652a -65 .48@ .48 .48
.48 .48 .48 .65
bib
6 6 6 6 6 0.0, 0 6 0 6 6
h b
ksb
6 3 6 3 6 6 6 6 6 6 6
.41 .48 6 6 G(Hz)oalod. = G(H~Oz)oalod. k's in units of 1. mole-' sec.-l X
Solving t,hese six equations for G(H20a)on t'he analog computer gives a value of 1.70 and 1.58 for the conditions detailed in Table IV. I n this calculation we arbitrarily assume kg = kg = kIo and that lcll
