GIDEOXCZAPSKI AKD B m o x H. J. BIELSKI
2 180
(4) The standard electrode potential, E O , for the 2e- is 0.95 T-. z's. the standard reaction Tec2 = TeO hydrogen electrode in contrast with previous estimates of 1.14 1 7 . ~ and 1.1v.' (5) A F o for the disproportionation reaction Tez? S Te Te-2 is 5.1 kcal., in contrast to the previous published value of 14.0 kcal. Acknowledgments.-The author is pleased to acknowledge helpful discussions with many colleagues at
+
+
Vol. 67
the Westinghouse Research Laboratories. He is especially grateful to Dr. Joseph Veissbart for stimulating discussions. Professor Paul Delahay's suggestion that a test of diffusion control be made is gratefully acknowledged. bfr. Paul Sassoon gave able assistance in the experimental work. The author wishes to thank Dr. Robert Nadalin for the loan of his polarograph instrument. This sTork was supported in part by a S a v y Bureau of Ships Contract.
THE FORMATION AXD DECAY OF H203 L4ND HO, I N ELECTRON-IRRADIATED AQEEOUS SOLUTIONS1 BY GIDEON CZAPSKI~ AND BESOXH. J. BIELSKI Chemistry Department, Brookhaven h-ational Laboratory, Upton, Long Island, S e u : York Receicecl March $7. 2963 The kinetics of the reaction between perhydroxyl radicals, HOZ,have been studied a t 23". The radicals were generated in rapidly flowing water containing dissolved oxygen by an electron beam from a Van de Graaff accelerator, and the quantity remaining at any time after irradiation was determined by rapid mixing with a solution containing tetranitromethane, which reacts with HOa radicals. The rate constants found were k ~ - o~0~~ = 2.2 x 106 M-1 see.-' and koa- + 02- = 1.5 X 10' -1f-l sec.-l, and the p K of HOa was found to be 4.4 2c 0.4. hnother intermediate, believed to be HZ03 on kinetic grounds, was found to exhibit pseudo-first-order disappearance, the first-order rate constant being a function of acidity. This intermediate was found using ferrous sulfate as the scavenger solution. The half-life of H203 reaches a maximum of 2 sec. in 0.02 M (H+),decreasing to less than 0.1 8ec. in 1 and 10-4 A' (H-), Tlit kinetics are shown to agree well with acid- and baee-catalyzed 0.. decomposition of HZ03 to form H,O
+
Introduction Rate studies on two unstable hydrogen-oxygen coinpounds are reported in this paper. The first compound is the hydroperoxy radical, H02,and its basic form Os-. The hydroperoxy radical is believed to react with itself to give hydrogen peroxide aiid oxygen. The rate constant of the reaction between 02-and 02- was measured recently by Schmidt using a conductivity rnethodu3 Saito and Bielski generated H02 in a flow system by mixing ceric ions with an excess of HzOzand followed the decay of the radical by e.8.r. measuremeiit~.~ The rate constants mere 1.45 x lo7 M-l sec.-l and 2.4 X lo6 J1-I sec.-l for 02- and HOz, respectively, when defined by -d(HOz),'dt = ~ ? C ( H O ~ ) ~ . We have generated HOn and 02-by flowing oxygensaturated water past a 1.6 RIev. electron beam. The HOz radical was detected by reaction with a solution of tetranitromethane in a mixing chamber shortly after the water had left the electron beam. In the course of these experinleiits another interniediate, longer-living, was found which exhibits a firstorder decay and which reacts with ferrous sulfate solutions. Experimental General Description of Apparatus and its Operation.-The experimental device mas a combination of a flom- system end a Tan de Graaff accelerator. The flom s j stem >vas similar to that emploied by Sutin in fast reaction studies.B T m o motor-driven syringes force two solutions through glass tubing to a mixing chamber and the mixed solutions flow into a beaker One SJ ringe contains the solution to be irradiated. This solution flows (1) Research performed under the auspices of the U. 5 . A t o m i c Energy Commission. (2) The H e b r e n Unirersity, Jerusalem, Israel. (3) K. Schmidt, 2. S a t u T f o r s c h . , 16B,206 (1961). (4) B. H. J. Bielski and E. Saito, J . P h y s . Chem., 6 6 , 2266 (1962). ( 5 ) E. Saito a n d H. H. J. Hielski, J . Am. Chem. Soc.. 83,4467 (1961). (6) N. %tin and B. XI. Gordon, ;hid., 83, 70 (1961).
through a thin-walled, small diameter tube and is irradiated with an electron beam a known distance before the mixing chamber. The other solution contains an excess of reagent which will react with the intermediate surviving a t the mixing chamber (tetranitromethane for the perhydroxyl radical and ferrous sulfate for the H203). The mixing chamber consists of two jets for each solution, the four jets entering tangentially to the discharge tube. To begin the run, a sxitch starts discharging the syringes and turns on the electron beam and a timer simultaneously. When the syringes are empty, the plunger trips a microswitch which stops all three. This use of a flow system differs from the normal use in that the reaction being studied occurs between the electron beam and the mixing chamber instead of occurring after the mixing chamber. The scavenger solution quenches intermediates surviving a t the mixing chamber and a proportionate color change is produced. Analysis is done a t the operator's leisure with a spectrophotometer. The syringes, flow tubes, and the mixing chamber were made of Pyrex glass, the stopcock plugs were Teflon. No grease was used. The solution to be irradiated was saturated with oxygen which was purified by bubbling it through 2 A!! NaOH and then twice through distilled water. Flow rates of 2 to 5 cc./sec. were used. The thin-walled flow tubes were 0.1 or 0.175 em. in diameter. Linear flow rates Tvere 1.5 to 6 m./sec. The Tan de Graaff beam was collimated by an aluminum block with a 3-mm. hole mounted in front of the flow tube. The scavenger solution was shielded from stray radiation by aluminum. Intervals between irradiation and mixing were altered by changing the flow rate, the capillary diameter, or the distance between the beam and the mixing chamber. The flow time was calculated from the flow rate, the capillary cross section, and the distance between the mixing chamber and the electron beam. Relative radiation intensities were measured by collecting the electron beam current and putting it through a current integrator. The current integrator readings were calibrated by flowing 5 X A' FeSOl in On-saturated, 0.8 '3' H&Y.Id solution past the beam and measuring the ferric ion produced. The ferric yield was taken as 15.5 ions oxidized per 100 e.v. absorbed. All experiments were a t room temperature (23 =t2'). Three runs were made for each experiment: (1) a blank run in which the solutions were passed through the system without being irradiated; ( 2 ) an "infinite time" run in which the scavenger was not added to the irradiated solution until a few minutes after the irradiation was stopped (at that time the short-lived inter-
FORMATION AND DECAY OF HzO3AND HOn IN AQUEOUSSOLUTIONS
Oct., 1963
mediates had reacxted, leaving only stable products which might react with the scavenger); (3) an experimental run in which both solutions were flowing, thus allowing the scavenger to react with the intermediates which reached the mixing chamber. In some rune, a gravity-flow system was used; the irradiated water passed into a well stirred flask containing scavenger solution. Materials.-Tetranitramethane (TSR.1) was extracted with a small amount of distilled water to remove some of the waterM Hi304 was then satusoluble decomposition products; rated with TXM ( 7 X X ) and this served as the scavenger solution. Hydroquinone was recrystallized and sublimed. All other reagents were analytical grade. The water was triply distilled. Analysis.--4 Beckman DU spectrophotometer was used with I-, 5-, or 10-cm. optical path quartz cells. Ferric ion was determined a t 305 mp using an extinction coefficient of 2180. Kitroform was measured a t 350 mp using an extinction coefficient of 14,800.
Results Reaction of Hydroperoxy Radicals.-H02 and 02are known to react with T N M yielding the colored product7 nitroform (NF-), according to the following reaction: C ( S 0 2 ) 4 HOz + C(KO~)C H + SO2 02. Keither T S M nor XF- reacts with HzOz but in acid solutions below pH 2, hydrogen peroxide and nitrogen dioxidle form pernitrous acid which does react with TNLI to form NF-.8 Hence the use of T K M a,s scavenger was limited to pH values above 2. The analysis was made immediately after each irradiation in order to reduce any post-irradiation effect resulting from the reaction of TlYbl with pernitrous acid. The HOz radical disappeared in time periods of the order of 10 to 20 msec., which is a t the lower end of the time scale available in our flom apparatus. Therefore it was difficult to measure the concentration of €IO2 a t constant intensities by changing the time, t , between mixing and irradiation. Instead, the radiation intensity was varied and the HOZ concentration was determined, keeping a definite time interval between irradiation and mixing. This procedure was then repeated using another constaiit time interval at varying intensity. Since HOz disappears in a second-order reaction
+
+
HO:,
-t HOs
the rate equation is4z5
+
-
H,Oz
+
0 2
+
2181
14 13
12 I I
IO T 5:
Y f 0 % N -
9
3 7
0
I \ -
v
6
5 4 3 2
I
I
2 I/I.
4
3
Fig. 1.-The variation of l/(HO$) with the reciprocal of the electron beam intensity a t pH 4.0 ( I = 1 corresponds to about 1Oz0 e.v. c c - 1 sec.-l). The time between irradiation and mixing with TNM solution was 10.7 msec., 0 , and 20 msec., 0 ; rate constants, from the intercepts ( 2 k o b s d t ) , are 1.15 X 107 and 1.25 X lo7A1-l see.-', rebpectively. I
1
I
I
I/
r.
I
I
I
I
6
7
8
9
\\
Jcobsd
or 2
3
4
5 PH
Since (H02)o= GHO,IX(where I is the intensity and X is the time spent in the electron beam)
If
is independent of intensity, a plot of 1,1'(H02)~ us. 1/1 at constant t sliould yield a straight line, with 2k,bsdt as intercept and (G1-10~ A) as slope. Figure 1 shows typical results. The deviation from a straight line in Fig. 1 indicates that GKO*decreases somewhat with increasing intensity. This decrease GH0,
( 7 ) A. Henplein and J. Jaspert, 2. pkgszk. Chem., 12, 324 (19.57) (8) 13. H. J. lheliikl and A. 0.Bllen, unpublished d a t a .
Fig. 2.-The pH dependence of the rate constant for reaction between HOz radicals. The curves are calculated from eq. 3 for the conditions: lower curve, ~ I & L = 0, pK = 3.3; middle l ~ 16, pK = 4.4; upper curve, klz/kll = 30, pK = curve, k ~ / k = 4.7.
probably is due to back reactions between the product and the radical intermediates. The intercept, 2 k o b s d t , is found to be proportional to t for the two times used, which confirms second-order kinetics. The dependence of kobsd on pH is shown in Fig. 2. Above pH 5, kobsd is constant a t 1.5 X lo7 M-l sec.-l, in excellent agreeineiit with Schmidt's3 value, 1.45 X lo7 L U - ~set.-'. At pH 2 /++,sd = 2.2 X lo6 lU-l sec.-l compared to 2.4 X loo ~ l i sec.-l - ~ measured by Bielski and Saito4
GIDEOXCZAPSICI AND BENON H. J . BIELSKI
2182
I
.2
.I
3
.4
.5
T I M E , SEC.
Fig. 3.-Example of first-order disappearance of new intermediate in 0.24 1l.I HC104 (ordinateis logarithmicacale). (Fe(II1))t - (Fe(II1))- is proportional to intermediate concentration. Electron beam current is 1.5 x 10-1 amp., 0; and 6 X amp., a. ( F e ( I I I ) ) mwas proportional to beam current. 9
,
'
I
7k .
0 W Lo
/i
6k 5
1
- L O G (H')
Fig. 4.-The decomposition rate constant of the new intermediate (H20d) as a function of the acidity: 0, HzSO4; 0 , HClOa; I,XaHSOd.
a t pH 0.5 and 1.5. The pH effect is likely due to the ionization of HOz, the reaction between 0 2 - radicals being faster than between HOz radicals. The Formation of a New Intermediate.-When ferrous sulfate n-as used as a scavenger instead of TNM, another species was found which lasted much longer than H02. Hydrogen peroxide is formed directly by radiation and so some oxidation of Fe+2 due to the hydrogen peroxide was observed even when the ferrous reagent was added to the irradiated solution several minutes after the radiation was stopped and the intermediates had decayed. The ferric iron concentration produced by this H202 is denoted as (Fe(III))-, while (Fe(II1)) is the ferric iron observed when the reagent is added at time t . Thus (Fe(III))t - (Fe(III)), is proportional to the intermediate concentration a t time t. A plot of log [(Fe(III))(- (Fe(III))m]/(Fe(III))m as a function of t is shown in Fig. 3. The coilcentrations were varied by more than fourfold b y varying the beam intensity. The linearity demonstrates first-order decay of this intermediate. Similar determinations were made for different acidities from to 1 JI adjusting the acidity by using perchloric acid, sulfuric acid, and sodium bisulfate.
Vol. 67
The decay was first order over the whole acid range. The results are presented in Fig. 4. Hydrogen ion concentrations in solutions of sulfuric acid and sodium bisulfate were calculated from the data of B a e ~ . The ~ rate constants depend only on the hydrogen ion concentration and not on the particular acid used. The intermediate has a maximum half-life of about 2 sec. at 0.02 M acid, while the half-life decreases to about 0.1 M acid. sec. a t 1and The intermediate is only produced in the presence of oxygen. l17hen nitrogen saturated solutions were irradiated, there was no difference between experimental runs and infinite time runs. Reaction with Other Reagents.-The following reagents showed no sign of oxidation or reduction by the new intermediate, i.e., they exhibited the same yield if mixed with the irradiated mater 0.1 sec. or 1 min. after leaving the beam: (a) ceric sulfate; (b) ceric, cerous mixtures; (c) ceric sulfate, thallous sulfate solutions; (d) ceric sulfate, copper sulfate solutions; (e) FeX3+? complexes where X is 1,lO-phenanthroline, 3,4,7,8-tetramethyl-l,lO-pheiianthroline, and 5-nitro-1,lO-phenanthroline; (f) 2,5-dichlorohydroquinone; (g) neutral potassium iodide solution (with ammonium molybdate catalyst added) ; (h) TNJI. Solutions a through f were in 0.4 X HzS04. The addition of copper sulfate to the ferrous sulfate solutions decreased the Fe(II1) yield. 2,5-Dimethylhydroquinone did react with the intermediate to form the quinone but the yield mas dependent 011 the hydroquinone concentration up to a saturated solution. The hydrogen peroxide yield in solutions in which the hydroquinone was mixed with the irradiated solution in the mixing chamber was considerably larger than when the hydroquinone was added later, suggesting that the reaction of the intermediate with 2,5-dimethylhydroquinone produces hydrogen peroxide. Discussion First-Order Decay.-The intermediate which exhibits the first-order decay is composed of hydrogen and oxygen as is shown by the fact that it is produced in oxygen-saturated mater containing either sulfuric acid or perchloric acid. We believe the new intermediate to be hydrogen sesquioxide,10H2O3,formed in the reaction
HOZ
+ OH
fast, in the electron beam
HzOa
(A)
The perhydroxyl radical is formed by the reaction of hydrogen atoms with oxygen. Previously, it was believed that reaction A was a disproportionation reaction yielding water and oxygen. The yield of H2O3is consistent with reaction A. The stoichiometry of the reaction between H20sand ferrous sulfate is assumed to be given by the reaction 4Fe(II) HzOa 4Hf = 4Fe(III) 3H20, while the H203 is 0 2 , assumed to decay by the reaction H2Oz = HzO which yields no peroxide. The yields of H202 and Hz03 (in terms of G, molecules per 100 e.v. absorbed) are then given by G(H202) = 1/2G(Fe(III))mand G-
+
+
+
+
(9) C. F.Baes, J r , J . Am Chem S o c , 79, 5611 (1957) (10) HsOa u a s suggested b y Bray [W G. Bray, z b d , 60, 82 (193811 as an intermediate in the reaction of 0 3 w l t h H20,. Hlgher oxldes of hs drogen (Hz03 and HzOi) were also suggested t o be formed b> dlscharae ln uater i a p o r and to be stdbilued b~ trapylng at lou telrlperatrlrea 'le? I 4 G h u r m l e ~ z, b z d , 79, 18fi2 (1957)
FORMATIOX ASD DECAY OF HzO3AND HOz IN AQUEOUSSOLUTIONS
Oct., 1963
(Hz03) = l/4[G(Fe(lII)), - G(Fe(III)),], where G(Fe(III))ois the iron yield extrapolated to zero time and G(Fe(III)), it3 the iron yield in “infinite time” runs. Hydrogen peroxide yields are given in Table I and H203 yields in Fig. 5 . Similar results for G(H202) were obtained using I - and Ce(1V) methods of analysis. TABLEI THEHYDROQEN PEROXIDE Y I E L D IN OXYGEN-SATURATED TIOKS t l i c DIFFEREXT ACIDITIES
SOLU-
(Intensities around lozoe.v. cc.-I sec.-l) Concentration, M
solution
HClOL
G(Hz0z)
1.05
10-8 0.26 .50 .78 10-3 .04
1.4 1.6 1.8 1.07 1 33
The hydroglen peroxide yields are about the same as have been observed at much lower iiitensitiesll which suggests that the combination of hydroxyl radicals is unimportant. Hence most of the hydroxyl radicals react with HOz. The lack of importance of hydroxyl radical combination is to be expected under the conditions of our experiments The rate of production of H 0 2 radicals exceeds the rate of production of hydroxyl radicals by about 2070.12 The hydroxyl radical lifetime, with respect to reaction with HO2, is about 1% of the irradiation time, so that the hydroxyl radical concentration will be close to the value calculated by steady-state assumptions. The H 0 2 radical will not approach a steady-state concentration, however, since it is produced in excess and the combination rate of two HOz radicals is low. Since the rate constants for reaction of O H with OH and HOz should be about the same, the large excess of H 0 2 should cause reaction A to predominate. The yield of H203 between pH 2 and 3 is constant at 1.7 which is somewhat less than GOH in this region, 2.5.lZ It mould appear that 70% of the reaction proceeds as in rextion 4 while the remaining 30% disproportionate t,o produce water and oxygen. The yield of H203 is lowered by H2S04and NaHS04 and increased by H(21O4. This is in agreement with the proposed mechanism, as the hydroxyl radical is known to react with bisulfate ion to give the bisulfate radical, l3 while acid increases the yield of hydroxyl radicals. I n another experiment, H203 was not produced when the irradiated solution contained bromide, in accordaiice with &he reaction of hydroxyl radicals with bromide ion. (Another intermediate was fornied which decayed over a period of several minutes and probably contained bromine.) The hrst-order decay of H203can be explained according to the rnechaniism
-ki
+ r3+ + o2 Z z HOs- + H+ HOs- +OH- +
Hdh
k.
E203
ka
ha
0 2
I A m Chem Suc 76, 4687 (1934) (12) A 0 Allen “ T h e Radiation Cheinlstrl of Watei arid lilueous ‘oliitions ’ I 1 Van Kostrand C o Neu York Y Y 1961 1, 47 (1 4 ) I I i u o i - k i J A m (“hem S u e , 7 8 , 1768 (1056) (11) T J h o r s k i
2
2183
-\ 0
-P 0
0
/ 0
0
0.
0
oo-
0
I -
Fig. 5.--STariation of H208 yield with hydrogen ion concentration: 0:I&SO,; 0 , HC104; . , NaHSOa.
The rate equation according to this mechanism, assuming a steady state for HOa-, will be
(Hz03)
=
k’(Hz03)
(2)
This expression predicts first-order kinetics over the whole pH range. A graph of l/k’ us. the hydrogen ion concentration between and 2 X lop3 M yields a 187. straight line with the equation l/k’ = 0.109 (H+). From this equation, kz is 9.2 sec.-l and k3/k4 is 1720 M-l. At high acid concentration k1 is found to be 6 M - l sec.-l. The curve in Fig. 4 is calculated from eq. 2 with these constants. Alternative 1nterpretations.-(1) One may assume that the intermediate is a long-lived excited mater molecule, H20*. However, since oxygen niust be present in the solution iii order to form the intermediate, it cannot be H20* alone. The reaction H20* O2 4 H203would agree with our analysis; however, large yields of HzO* are unsupported by any other evidence. ( 2 ) A higher oxide, H204, could be postulated which might react in a scheme analogous to H203
+
+
fast
2H02 +Hz04
+ H T +Hz02 + + H + Hz04 JJ H+ + Hodkl
HzOa
0 2
kz
ks k4
HOC +HOz-
+
0 2
This mechanism is kinetically indistinguishable from the Hz03 mechanism, yielding exactly the same equations. The two mechanisms do differ in the decomposition products Hz03 + HzO 0 2 , H204 + . H202 02. H204 could oxidize six ferrous ions, while liydrogen peroxide is known to oxidize two ferrous ions. Thus the maximum ratio of ferrous oxidation as the solution leaves the beam to that a t infinite time is
+
+
+ 6G(Hz04) < + 2(;.(H204) -
(Fe(III))o - 2G(H~02) (Fe(III)), ZG(Hz02)
while our experimental results iii 0.24 ;I1 HC104 givc the value (Fe(III))“/(Fe(III))m = 3 f. 0.4. (Thc i l l -
GIDEOXCZAPSKI ASD BENON 13. J. BIELSKI
2184
terccpt in Fig 4 is (17e(III))o/(1~e(III))m - 1.) There is 110 such restrictioii on tlie ferric yields in the case of 11203, as the deconiposition products are IIzO and 0 2 which do not oxidize ferrous. The effect of lowering tlic ferric yield by adding copper iori can be explaiiied if the rraction of 1 3 2 0 8 with fcrrous ion produces ferric ion and a perhydroxyl radical.
+ E'e(I1) +€ 1 0 2 + 011- + p'c(II1)
IIL03
VOl. 67
iaed form, 02-, mas demonstrated by Schmidt, who followed reaction 22 by measuring condu~tivity.~ The sum of the two radical concentrations, (1-102)4(02-) = R, was measured in our experimeiits. Thus
or
?'he IT02 will react with fcrrous or with copper ion as proposed by IIart.lL
+ l+(II) + 11' +1 1 2 0 2 + Fc(II1) 1102 + C U ++ ~ H+ + + Cut C'uf + E'c(II1) + C U + + ~ Fe(I1)
€TO?
li',bsd
0 2
Reartion with C'u+? thus produces no net oxidation of fcrrous ion. The pK of H 2 0I.---Thc pK of H201would be given by -log h2,'L I a i d /i? vas fouild to hr 9 6 ~ ~ . - 1 , The rcwrse rate constant, X
