of 10

Reaction of O 2 with the Hydrogen Atom in Water up to 350 °C

33 views
All materials on our website are shared by users. If you have any questions about copyright issues, please report us to resolve them. We are always happy to assist you.
Share
Description
Reaction of O 2 with the Hydrogen Atom in Water up to 350 °C
Transcript
  Reaction of O 2  with the Hydrogen Atom in Water up to 350  ° C Ireneusz Janik and David M. Bartels*  Notre Dame Radiation Laboratory, Notre Dame, Indiana 46556  Timothy W. Marin Chemistry Department, Benedictine Uni V  ersity, Lisle, Illinois 60532 Charles D. Jonah Chemistry Di V  ision, Argonne National Laboratory, Argonne, Illinois 60439 Recei V  ed: August 9, 2006; In Final Form: October 27, 2006  The reaction of the H • atom with O 2 , giving the hydroperoxyl HO 2 • radical, has been investigated in pressurizedwater up to 350  ° C using pulse radiolysis and deep-UV transient absorption spectroscopy. The reaction ratebehavior is highly non-Arrhenius, with near diffusion-limited behavior at room temperature, increasing to anear constant limiting value of   ∼ 5  ×  10 10 M - 1 s - 1 above 250  ° C. The high-temperature rate constant is innear-perfect agreement with experimental extrapolations and ab initio calculations of the gas-phase high-pressure limiting rate. As part of the study, reaction of the OH • radical with H 2  has been reevaluated at 350 ° C, giving a rate constant of (6.0  (  0.5)  ×  10 8 M - 1 s - 1 . The mechanism of the H • atom reaction with theHO 2 • radical is also investigated and discussed. I. Introduction The reaction of the H • atom with aqueous O 2 , forming thehydroperoxyl radical,is of central importance in combustion chemistry and has beencarefully studied both experimentally 1,2 and via quantumchemistry. 3 The reaction in very-high-temperature water is of importance for supercritical water oxidation processes. 4 In water-cooled nuclear reactor chemistry, the reaction is an importantstep in “hydrogen water chemistry”, whereby a slight over-pressure of H 2  is used to reduce the radiolytically produced O 2 and H 2 O 2  back to water. 5 It is important to be able to predictthe minimum H 2  concentration needed to accomplish this kinetictrick, because excess H 2  may result in undesirable hydridingcorrosion.In the gas-phase formation of HO 2 • , there is no barrier toO - H bond formation for approach of the H • atom at about 45degrees with respect to the O - O axis. 3 In room-temperaturewater, the measured reaction rate is consistent with nearlydiffusion-limited behavior 6 - 9 and a barrierless reaction. Thequestion is whether this remains true at elevated temperature,and whether the gas-phase potential is strongly perturbed bythe aqueous environment. If not, the gas-phase measurementscan be simply transferred to the aqueous-phase simulationproblem.In the present investigation, the reaction of H • atoms and O 2 in pressurized water has been studied at temperatures up to 350 ° C using pulse radiolysis and transient deep-ultraviolet absorp-tion spectroscopy. The reaction demonstrates strongly non-Arrhenius behavior above 100  ° C, with the rate constantapparently decreasing above 325  ° C. In the discussion below,we show that the behavior for this reaction is entirely consistentwith the gas-phase high-pressure limit. The mechanism forreaction of H • atom with HO 2 • radical is also discussed. II. Experimental Section Electron pulse radiolysis/transient absorption experimentswere carried out using 4 - 20 ns pulses from the ArgonneChemistry Division’s 20 MeV electron linac accelerator. Thesample cell, flow system, and basic experimental setup andcharacteristics were described in previous publications. 10,11 However, for the current measurements some changes wereapplied. Analyzing light from a pulsed 150 W xenon lamp(Osram XBO-150W/S) was selected using an ISA-545 double-grating monochromator. With this monochromator, there wasno detectable scattered light from the visible or near-UV todistort the deep-UV absorption measurements. The detector wasa five-stage Hamamatsu photomultiplier wired to deliver linearphotocurrent up to 2 mA. To minimize the scattering loss of deep-UV analyzing light encountered over a potentially longoptical path, the detector system was placed roughly 50 cm fromthe sample cell and shielded from the linac radiation with leadbricks.The sample was mixed from three separate syringe pumps(ISCO-260C) working in constant flow mode. The total flowrate was 4 mL/min. The first pump contained water saturatedwith N 2 O at room temperature, giving an N 2 O concentration of 0.024 molal ( m ). The N 2 O (AGA gas, Ultrahigh Purity) wasfirst bubbled through a sparging vessel filled with a highly basicsolution of pyrogallol to remove any traces of oxygen. 12 Thesecond pump contained water saturated with O 2  (AGA gas,Ultrahigh Purity) at room temperature, giving an O 2  concentra-tion of 0.0013  m . The third pump contained water pressurized * Author to whom correspondence should be addressed. E-mail: bartels@hertz.rad.nd.edu. Phone: (574) 631-5561. Fax: (574) 631-8068. H • + O 2 f  HO 2 • (1) 79  J. Phys. Chem. A  2007,  111,  79 - 8810.1021/jp065140v CCC: $37.00 © 2007 American Chemical SocietyPublished on Web 12/14/2006  with H 2  (AGA gas, Ultrahigh Purity) at room temperature.Pressurized H 2  samples were prepared in our laboratory-builthigh-pressure gas - liquid saturator. 13 The pressure of hydrogenin the saturator was constant for given samples of water usedto refill the syringe pump but decreased in the course of severaldays of experiment. Typically, the hydrogen pressure rangedfrom 82 to 149 bar, giving aqueous hydrogen concentrationsof 0.100 - 0.055  m . 14 The hydrogen concentration in the samplemix was constant for a given experiment and was always equalto 50% of the corresponding hydrogen concentration in thepressurized water saturator. The oxygen concentration in thesample was changed by changing the flow ratio between pumpsdelivering N 2 O- and O 2 -saturated water and was in the range2.5  ×  10 - 5 to 2.0  ×  10 - 4 m . Correspondingly, the N 2 Oconcentration was varied between 1.2  ×  10 - 2 and 5.0  ×  10 - 3 m . The total pressure in the system was adjusted with a back pressure regulator to 250 ( 0.1 bar. Normal temperature stabilitywas  ( 0.3  ° C.Pulse radiolysis of water creates predominantly hydratedelectrons (e aq - ) and OH • radicals, which are quickly convertedto hydrogen atoms via reactions 2 and 3.In the limit of high oxygen concentration, reaction 1 becomespseudo-first-order, and thus the reaction 1 rate constant,  k  1 , canbe established by monitoring the increase of hydroperoxylradical absorbance at 230 nm. 15 At room temperature, the p K  a for the HO 2 • radical is 4.8, and virtually all of the radicalpopulation converts to the basic O 2 •- form via reaction 4, givinga product that absorbs at 250 nm. Because all of the measure-ments were performed in water of neutral pH, and the UVabsorption bands are of typical condensed-phase width for bothHO 2 • and O 2 •- , the resulting absorption was a sum of both HO 2 • and O 2 •- radicals. At elevated temperature, the fraction of O 2 •- in the measured signals is lower due to changes in equilibrium(eqs 4,-4) with increasing temperature. 16 Above 300  ° C, theequilibrium lies almost entirely to the left, giving essentiallyonly HO 2 • .Measurement of the reaction 1 rate under the chosenconditions is limited by the lowest oxygen concentration thatcan be applied before the oxygen is depleted by repeated electronpulses. However, with too high an oxygen concentration, thesecondary reaction (5) becomes increasingly important as itgenerates O 2 •- without ever reacting with H • atoms, thusobscuring the observed reaction 1 rate. Furthermore, it decreasesthe amount of H atoms produced by competition with reactions2 and 3. The experimental limits placed on the O 2  concentrationeffectively minimized this pathway so that nearly all the e aq - were eventually converted to H • . To overcome the limitationof too low an O 2  concentration, the doses applied in theexperiment were the lowest possible to obtain usable UV signalsafter averaging of 30 consecutive traces. However, even for thelowest applied doses, H • atom recombination cannot be ignoredin the overall reaction scheme. Reaction 6 slightly decreasesthe H • atom concentration, especially at high doses due to itssecond-order nature. It becomes increasingly important at highertemperatures where this reaction becomes very fast. 17 The collected UV traces were fitted to a complex waterradiolysis model that includes a set of all the known waterradiolysis reactions and all the extinction coefficients for speciesabsorbing at the experimental wavelength. The occurrence of reaction 6 can be nicely accounted for when the dose depen-dence of the fitted kinetics is included. Therefore, all measure-ments were carried out using two different doses of roughly 2and 10 Gy and at least three different oxygen concentrations.Fits to the experimental data must take into account therelative dose delivered in each experiment, as well as changesin the absorbed radiation due to decreasing water density withincreasing temperature. The relative dose delivered by theaccelerator was measured by integrating charge on a thick brass“shutter” inserted between the beam port and sample cell. It isassumed that the absorbed dose is simply proportional to thewater density, as long as the sample is thin enough to avoidsevere scattering of the electrons. The water density wascalculated using the IAPWS-IF97 formulation for light waterPVT relations. 18 III. Results and Analysis The presence of more than one absorbing species (HO 2 • , O 2 •- ,OH • , H • , etc . ) reacting simultaneously leads to kineticallycomplex absorbance - time profiles that can only be resolvednumerically using a computer code. In turn, this means that asmany parameters as possible should be accurately known apriori; these include molar absorptivities of all species,  G  values,rate constants and radiation dose. A. G Values.  G  values (escape yield of radicals in molesper unit energy of radiation absorbed) used in the data analysiswere based on the compilation of previous measurements of Lin et al., 19 and our unpublished results for the yields of hydratedelectron, H • atom, and H 2 . 20 From the Lin et al. results we takethe total yield  G (e aq - + OH • + H • ) reported up to 350  ° C andby subtraction of our experimental  G (H • ) + G (e aq - ) values weestimate  G (OH • ) numbers in the studied temperature range. The G  values for H 2 O 2  and H 2  were based on previous data suppliedby Elliot et al. 21,22 B. Molar Absorption Coefficients.  OH  • . The product of    Gfor the OH • radical in N 2 O-saturated water was measured inour apparatus at constant dose at 230 and 250 nm fortemperatures in the range 25 - 350  ° C. Recombination inoxygen-free water is slow enough that a simple measurementof absorption at 200 ns after the 4 ns electron pulse is sufficientfor this purpose. We confirmed previous observations that theOH • radical spectrum does not change up to 200  ° C 23 if a densitycorrection is applied for the absorbed dose. However, consider-ing the increase in the spur escape yields of initial transientspecies with temperature, 19 the absorption coefficient of OH • radicals at 230 and 250 nm must become lower with temper-ature. This is apparently related either to the spectrum shiftingtoward deeper UV or to depleting of the 230 nm band with thetemperature increase (more detailed analysis of the OH • spectrum is the subject of a future paper 24 ). Applying themeasured   G and calculated  G (OH • ) at a given temperature,the OH • radical extinction coefficients for 230 and 250 nm havebeen determined. Figure 2 shows the temperature dependenceof these parameters. The OH • radical extinction coefficientcontinuously decreases with temperature from room temperature e aq - + N 2 O f  OH • + N 2 + OH - (2)OH • + H 2 f  H • + H 2 O (3)HO 2 • T  O 2 •- + H + (4,-4)e aq - + O 2 f  O 2 •- (5)H • + H • f  H 2  (6) 80  J. Phys. Chem. A, Vol. 111, No. 1, 2007   Janik et al.  up to 350  ° C. The values used for the data analysis aresummarized in Table 1.  HO 2 •  /O 2 •- . Initially, the molar absorptivities for O 2 •- andHO 2 • were determined on the basis of the   max G  changes vstemperature reported by Christensen and Sehested up to 250 ° C for hydrogenated and oxygenated solutions of variable pH. 16 To estimate the  G  values after the 1  µ s electron pulse for theseexperimental conditions (i.e., H 2  and O 2  concentrations, and pH),we applied the total  G (e aq - + OH • + H • ) recently obtained frommethyl viologen measurements by Lin et al. 19 However, fromall our fitting attempts it was evident that the extinctioncoefficient values so derived for HO 2 • and O 2 •- are too high.The increase of temperature implied an increase of the givenextinction coefficients. Our kinetics observations indicate theopposite trend.Consequently, it was decided to record new HO 2 • and O 2 •- spectra as a function of temperature. Spectra were measuredusing a 50:50 mixture of H 2 -pressurized water and O 2 -saturatedwater, thus the final room-temperature concentrations of the H 2 and O 2  solutions upon mixing of the two solutions wereapproximately 5.0 × 10 - 2 and 6.5 × 10 - 4 m , respectively. Toachieve acidic or basic pH, perchloric acid or potassiumhydroxide were respectively added to the O 2 -saturated water.Just as in the experiments of Christensen and Sehested, 16 thechemistry ensures that all the primary OH • , e - aq , and H • radicalsformed upon radiolysis are scavenged to form HO 2 • or O 2 •- .The applied dose was held constant from day to day, varyingon different days by no more than 6%. We were able to registerHO 2 • spectra at pH ) 2 and O 2 •- spectra at pH ) 8 up to 200 ° C. Figure 1 shows a comparison of HO 2 • and O 2 •- spectracollected at room temperature and at 200  ° C. The extinctioncoefficients presented in Figure 1 represent the new experimental  G divided by  G (e aq - +  OH • +  H • ) obtained by Lin et al. 19 From the results presented in Figure 1 for the HO 2 • and O 2 •- spectra, one can see that a temperature increase up to 200  ° Cslightly lowers the HO 2 • extinction coefficient. This observationagrees with results obtained by Buxton et al., for HO 2 •  /O 2 •- spectra recorded up to 175  ° C. 25 It also roughly agrees with thebehavior of the gas-phase spectrum recorded by Kijewski andTroe. 26 In contrast, the O 2 •- spectrum does not change up to200  ° C within experimental error. Though these results con-tradict Buxton’s report, it agrees with the observations of Christensen and Sehested 16 that the product   max G  for O 2 •- increases more than for HO 2 • with increasing temperature. Itshould be noted, however, that the current result is quantitativelydifferent from either of the previous reports.Above 200  ° C, reliable measurements could not be performedin acidic conditions as substantial corrosion in the metal flowsystem was observed. In addition, hot alkaline solution in thepresence of oxygen etches the sapphire windows 27 and makesthe UV measurements impossible due to excess light scattering. Figure 1.  Temperature dependence of the HO 2 •  /O 2 •- radical spectrumat various pH. Circles: O 2 •- radicals at room temperature (solid) and200  ° C (open). Squares: HO 2 • radical at room temperature (solid) and200  ° C (open). Triangles: effective extinction coefficients of HO 2 •  / O 2 •- at 250  ° C (inverse, open), 300  ° C (solid), and 350  ° C (open). Allspectra shown were acquired using the same applied dose of 9.5 ( 0.5Gy. Figure 2.  Changes of extinction coefficient of species contributing tothe absorption at the experimentally chosen wavelengths. Symbols:extinction coefficients at 230 nm. Symbols and lines: extinctioncoefficients at 250 nm. Corresponding symbols: squares, HO 2 • ; filledcircles, O 2 •- ; triangles, OH • ; open circles, H • . TABLE 1: Extinction Coefficients as a Function of Temperature (M - 1 cm - 1 ) for the Species ExperimentallyObserved (n.d.  )  No Data Available) fit results (av values)Christensen et al. HO 2 • O 2 •- temp( ° C)HO 2 • at  λ max O 2 •- at  λ max 230nm250nm230nm250nmOH • 230nmH • 230nm25 1251 1892 1222 717 1760 1890 582 20100 1274 1950 n.d. 690 n.d. 1850 n.d. n.d.150 1290 1988 1183 673 1735 1813 550 30200 1306 2027 n.d. 665 n.d. 1800 n.d. n.d.225 1314 2046 1160 660 1720 1750 525 106250 1322 2066 1121 650 1715 1730 518 112275 1330 2085 1075 640 1710 1700 492 134300 1338 2105 1050 620 1705 1680 475 176325 1346 2124 1020 610 1700 1660 445 208350 1354 2143 980 600 1680 1600 409 227 Reaction of O 2  with the Hydrogen Atom in Water  J. Phys. Chem. A, Vol. 111, No. 1, 2007   81  Therefore, for temperatures higher than 200  ° C, effectivecombined spectra of HO 2 •  /O 2 •- were recorded at neutral pH.The spectra recorded at 300 and 350  ° C are shown in Figure 1(triangles). The maximum of the HO 2 •  /O 2 •- effective spectrumat neutral pH shifts toward deeper UV and the effectiveextinction coefficients decrease with increasing temperature. Thechange in the effective HO 2 •  /O 2 •- spectrum can be correlatedwith the change of the p K  a  value of the HO 2 • radical.The value of the p K  a  for equilibrium (4, - 4) is 4.8 at roomtemperature 15 and increases to a value of 6.15 at 285  ° C,according to the report of Christensen and Sehested. 16 Extrapo-lating these experimental numbers up to 350  ° C, one can expecta value of p K  a  ∼  7, which is above the pH of water at thistemperature (pH 350 ° C  )  5.98), thus suggesting predominanceof HO 2 • . For the whole range of temperatures between 200 and350  ° C, we have measured partial spectra of HO 2 •  /O 2 •- between230 and 250 nm. A sample of a partial spectrum recorded at250  ° C is superimposed in Figure 1 (reversed triangles). Thecombined extinction coefficients obtained at 230 and 250 nmwere used to estimate separate extinction coefficients for HO 2 • and O 2 •- at these wavelengths by iteration, using the ratio of HO 2 • and O 2 •- concentrations based on the p K  a  value at a giventemperature. For temperatures higher than 285  ° C, the p K  a  wasextrapolated using a third-order polynomial function that fitsthe experimental p K  a  values encountered at lower temperatures.Uncertainty in the extrapolated p K  a  value is probably ( 0.2 p K  units at 350  ° C. The iteration results were used as initial guessesin the analysis of the kinetics. During numerical analysis of thedata, the initial guesses were slightly iterated to find the bestfit. The measured combined absorbance applies in any case tothe final product; separate extinction coefficients are only neededto fit the signal rise for relatively large O 2  concentrations, wherethe observed signal could be due, in part, to product contributionfrom reaction 5. The final best estimates of both extinctioncoefficients and their changes with the temperature are shownin Figure 2.  H  • . The H • atom absorption spectrum was reported previouslyfrom room temperature up to 200  ° C, and no changes in thespectral shape were found over this temperature range. 17 It isgenerally accepted that the absorption must be due to watermolecules in the first solvation shell, because the hydrogen atomLyman Alpha line is encountered far into the vacuum ultraviolet.The shape of the H • atom spectrum decreases exponentiallytoward the red and somewhat resembles the shape of the waterabsorption edge, which also does not change shape withtemperature, although it shifts to the red. 28 (For water, the shapeis probably controlled by the Franck  - Condon envelope for thelowest bound-dissociative transition. 28 ) At 250 nm, the H • atomextinction coefficient is as low as 30 M - 1 cm - 1 . 29,30 For ourdata analysis at 250 nm, the H • atom absorption was very minor.However, for all data recorded at 230 nm it was necessary toincrease the value of the H • atom extinction coefficient,especially for temperatures higher than 200  ° C. The changesof the H • atom extinction coefficient at 230 nm resulting fromfits to the data are summarized in Table 1 and are superimposedin Figure 2. Given the 0.6 eV red shift of the water absorptionedge between room temperature and 400  ° C, 28 it seems quitereasonable to assume that the H • atom spectrum could shift bya similar amount.  H  2 O 2  and e aq - . Both the hydrogen peroxide and the hydratedelectron extinction coefficients were included in the data analysisand fits. However, they did not affect the fitted results as (i)the hydrated electron was converted to OH • radicals in tens of nanoseconds according to reaction 2, basically in the limit of our time resolution, and (ii) H 2 O 2  forms in negligible amountsfrom OH • recombination during the course of the experimentdue to strong scavenging of the radical by reaction 3, and couldexist only as a product of spur reactions with much loweryields 22 and relatively small extinction coefficients in the rangeof interest. C. Data Analysis.  For reaction 1, the change of the totalabsorbance with time after an electron pulse was initiallymeasured for the given temperature range at 250 nm and neutralpH. A study was carried out as a function of temperature up to350  ° C. All data were collected at a pressure of 250 bar. Typicaldata taken at 25 and 200  ° C are shown in Figures 3 and 4, withfitted curves superimposed. The signals track the initial decayof OH • radicals and subsequent formation of HO 2 •  /O 2 •- radicalsby their absorption at 250 nm. The different sets of tracescorrespond to different doses applied, with larger amplitudesignals representing the higher doses. The different tracescorrespond to different oxygen concentrations, where theconcentration was varied between 1.6  ×  10 - 5 and 1.3  ×  10 - 4 m . The concentration of hydrogen was not changed for experi-ments conducted at a given temperature, and the concentrationof nitrous oxide varied only by 20%, being kept in the range of 1.2 × 10 - 2 m . An increase of the oxygen concentration causesan apparent faster rise of the HO 2 •  /O 2 •- product, in agreementwith pseudo-first-order behavior. At temperatures above 200 ° C, the signal amplitude at 250 nm decreases considerably(compare absorption spectra in Figure 1) for the range of applieddoses. An increase of the dose at higher temperatures causestwo unwanted effects: (i) increase of contributions from second-order reactions and (ii) depletion of the oxygen in the systemas a result of the many electron pulses applied to the samesample. Therefore, the experiment was repeated for the rangeof temperatures 150 - 350  ° C at 230 nm where there is a muchhigher contribution from the HO 2 • absorption than from O 2 •- ,and higher overall effective absorption. Using these conditions,we could keep the same range of applied doses for all the desiredtemperatures.The simple pseudo-first-order approach to the kinetics wasnot good enough to provide satisfactory fits to the data. We Figure 3.  Formation of HO 2 • at 25  ° C for two applied doses, (a) and(b), with (a) being a higher dose. The three traces of each set correspondto O 2  concentrations 3.2 × 10 - 5 , 6.5 × 10 - 5 , 1.3 × 10 - 4 m , respectively.The N 2 O concentration varies between 1.18 and 1.00  ×  10 - 2 m , andthe H 2  concentration is a constant 4.82 × 10 - 2 m . Signals were acquiredat 250 nm. Fits are superimposed as solid lines. 82  J. Phys. Chem. A, Vol. 111, No. 1, 2007   Janik et al.  were forced to build a kinetically complex model that includedmany reactions involving all reactive species present. Theobserved kinetics can be modeled by a set of some 50 reactions,using a model described in previous publications, 6 but adaptedto handle UV-absorbing species. Each rate constant in the modelwas tested for sensitivity toward the fit quality. The majorityof the reactions are only minor contributors to the kinetics andcan be ignored. The set of reactions responsible for the observedkinetics follows:The temperature dependence of reaction 2 was previouslydetermined 31 and was fixed for the data fitting.Non-Arrhenius behavior of   k  3  was reported previously 13 andthese results were used up to 325  ° C. However, at 350  ° C asatisfactory fit could not be achieved with the rate constant valuereported previously. 13 It appeared that the reported value wassignificantly lower than was necessary to fit the presentexperimental results. To redetermine  k  3  at 350  ° C, an experimentwas performed to directly monitor the OH • radical decay at 250nm in the presence of N 2 O and various concentrations of H 2 ,but in the absence of O 2 . With these reactants present, the e - aq are scavenged by the N 2 O and converted to OH • . The OH • subsequently reacts via reaction 3. To fit the experimental traces,we used the same fitting model described above, but  k  3  wasfitted instead of   k  1 . The results of this experiment are presentedin Figure 5. With increasing H 2  concentrations, a faster decayof the OH • radical is observed as a result of reaction 3. However,at the same time, reaction 7 restores part of the OH • radicals,giving a contribution to the absorption in the tail of the signal(a chain reduction of the N 2 O). For the highest concentrationof H 2 , the contribution of reaction 7 is less obvious, as theregenerated OH • is quickly removed again by reaction 3. TheH • atom does not absorb significantly at this wavelength, andits second-order decay rate via reaction 6 is diffusion-limited.From the fits, a  k  3  value of (6.0  (  0.5)  ×  10 8 M - 1 s - 1 wasobtained, which is roughly 70% higher than the rate reportedpreviously. (We should note that the previous value was obtainedat the detection limit of the competition method being used,and a large error bar was admitted. 13 ) The fits were found tohave a high sensitivity to  k  7 . It was found that the collectedtraces could be fit only using rate constant values between 1.0and 1.5 × 10 8 M - 1 s - 1 to give an uncertainty in  k  3  of  ( 0.5 × 10 8 M - 1 s - 1 . Therefore, the upper limit for  k  7  was taken to be1.5 × 10 8 M - 1 s - 1 at 350  ° C. On the basis of the reported room-temperature rate constant 32 of 2.1 × 10 6 M - 1 s - 1 , we estimatedthe Arrhenius dependence for reaction 7 with the activationenergy  E  a ) 18.5 kJ mol - 1 and pre-exponential  A ) 3.7 × 10 9 M - 1 s  - 1 over the temperature range studied. Correspondingrate constants for reaction 7 at given temperatures were includedin the reaction 1 analysis as a minor correction.Values for  k  - 4  were calculated assuming diffusion-limitedbehavior on the basis of the Debye - Smoluchowski equationwith Figure 4.  Formation of HO 2 • at 200  ° C and two applied doses (a) and(b), with (a) being a higher dose. The four traces at each dose correspondto following O 2  concentrations 1.625 × 10 - 5 , 3.25 × 10 - 5 , 6.5 × 10 - 4 ,1.3  ×  10 - 4 M, where the higher the concentration, the faster theobserved rise rate. Here, the N 2 O concentration varies between 1.00and 1.18  ×  10 - 2 m , and the H 2  concentration is 4.82  ×  10 - 2 M. Thesignals are acquired at 250 nm. Fits are superimposed as solid lines. e aq - + N 2 O f  OH • + N 2  (2)OH • + H 2 f  H • + H 2 O (3)H • + O 2 f  HO 2  (1)HO 2 • T  O 2 •- + H + (4,-4)H • + H • f  H 2  (6)H • + N 2 O f  OH • + N 2  (7)H • + O 2 •- f  HO 2 - (8)H • + HO 2 • f  H 2 O 2  (9) Figure 5.  Decay of the OH • absorption after addition of H 2  at 350  ° Cfor two applied doses ( a  )  8.5 Gy,  b  )  17 Gy). H 2  concentrations:(a) 0.0, 3.55  ×  10 - 3 , 7.16  ×  10 - 3 , 3.52  ×  10 - 2 m ; (b) 0.0, 3.55  × 10 - 3 , 7.16  ×  10 - 3 , 1.42  ×  10 - 2 , 3.52  ×  10 - 2 m . A faster decaycorresponds to higher H 2  concentration. The N 2 O concentration variesbetween 2.5 and 1.28 × 10 - 2 m . The signals were acquired at 250 nm.Fits are superimposed as solid lines. k  diff  )   4 π   RDF  D  (10) Reaction of O 2  with the Hydrogen Atom in Water  J. Phys. Chem. A, Vol. 111, No. 1, 2007   83
Related Search
Advertisements
Related Docs
View more...
We Need Your Support
Thank you for visiting our website and your interest in our free products and services. We are nonprofit website to share and download documents. To the running of this website, we need your help to support us.

Thanks to everyone for your continued support.

No, Thanks