of 9

Carbonate Radical Formation in Radiolysis of Sodium Carbonate and Bicarbonate Solutions up to 250 °C and the Mechanism of its Second Order Decay

29 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
Carbonate Radical Formation in Radiolysis of Sodium Carbonate and Bicarbonate Solutions up to 250 °C and the Mechanism of its Second Order Decay
Transcript
  Carbonate Radical Formation in Radiolysis of Sodium Carbonate and BicarbonateSolutions up to 250  ° C and the Mechanism of its Second Order Decay Kyle S. Haygarth, Timothy W. Marin, Ireneusz Janik, Kotchaphan Kanjana, Christopher M. Stanisky, and David M. Bartels*  Radiation Laboratory, Uni V ersity of Notre Dame, Notre Dame, Indiana 46556  Recei V ed: No V ember 3, 2009; Re V ised Manuscript Recei V ed: December 15, 2009 Pulse radiolysis experiments published several years ago (  J. Phys. Chem. A ,  2002 ,  106  , 2430) raised the possibilitythat the carbonate radical formed from reaction of   · OH radicals with either HCO 3 - or CO 32 - might actually existpredominantly as a dimer form, for example,  · (CO 3 ) 23 - . In this work we re-examine the data upon which thissuggestion was based and find that the srcinal data analysis is flawed. A major omission of the srcinal analysisis the recombination reaction  · OH  + · CO 3 - f HOOCO 2 - . Upon reanalysis of the published data for sodiumbicarbonate solutions and analysis of new transient absorption data we are able to establish the rate constant forthis reaction up to 250  ° C. The mechanism for the second-order self-recombination of the carbonate radical hasnever been convincingly demonstrated. From a combination of literature data and new transient absorptionexperiments in the 1 - 400 ms regime, we are able to show that the mechanism involves pre-equilibrium formationof a C 2 O 62 - dimer, which dissociates to CO 2  and peroxymonocarbonate anion: · CO 3 - reacts with the product peroxymonocarbonate anion, producing a peroxymonocarbonate radical · O 2 COO - , which can also recombine with the carbonate radical: · CO 3 - + ·  CO 4 - f C 2 O 72 - I. Introduction The carbonate anion radical  · CO 3 - is one of the mostextensively studied inorganic radicals due to its ubiquitousnature in the environment, its relatively long lifetime, andits accessible absorption at 600 nm. 1 - 15 Typically the speciesis generated by  · OH radical reaction with either CO 32 - orHCO 3 - ion as in reactions 1 and 2:Since direct measurement of the weak UV absorption of   · OHis difficult, the carbonate ion is particularly useful as ascavenger for  · OH radicals in laboratory studies. A study of the carbonate radical yield in pulse radiolysis of bothcarbonate and bicarbonate solutions at elevated temperature 15 suggested several years ago that the radical actually existsin a dimeric form, for example,  · (CO 3 ) 23 - . In this work were-examine the evidence cited for this claim and show it tobe based on a flawed analysis. However, we also confirmthe observation 15 that decay of the radical deviates signifi-cantly from simple second-order kinetics when the radicalis generated in bicarbonate solutions. We review the evidenceavailable for the recombination reaction and suggest amechanism that can explain the data.For such a well-known radical, the structure, acid/baseproperties, and recombination mechanism of   · CO 3 - haveremained controversial for a very long time. The firstcomprehensive study of the carbonate radical using pulseradiolysis was accomplished by Weeks and Rabani, 1 whorecorded the transient spectrum and investigated the decaykinetics as a function of pH. They found the decay of thecarbonate radical in alkaline solution was second order over80% of the decay, but deviated from second-order kineticsat longer time. A shift in the nature of the deviation couldbe seen between pH 11.5 and 13. The authors provided noexplanation for the deviation from second order. Two setsof reaction products were postulated as consistent with thedata:Behar et al. generated the carbonate radical in oxygenatedsolutions both by radiolysis and by photolysis of hydrogen * Corresponding author. Phone:  + 1 574 631 5561; Fax:  + 1 574 6318068; E-mail: bartels@hertz.rad.nd.edu. · CO 3 - + ·  CO 3 - T C 2 O 62 - f CO 2 + O 2 COO 2 - · CO 3 - + CO 42 - f  ·  CO 4 - + CO 32 - · OH  +  CO 32 - f  · CO 3 - +  OH - (1) · OH  +  HCO 3 - f  · CO 3 - +  H 2 O (2) · CO 3 - + · CO 3 - f CO 2  +  CO 42 - (3a) f 2CO 2  +  H 2 O 2  +  2OH - (3b)  J. Phys. Chem. A  2010,  114,  2142–2150 2142 10.1021/jp9105162  ©  2010 American Chemical SocietyPublished on Web 01/15/2010  peroxide solutions. 2 They measured the reaction rates of carbonate radical with both H 2 O 2  and HO 2 - , finding a muchfaster rate for HO 2 - :They explained the deviation from second-order kineticsobserved by Weeks and Rabani 1 as the reaction of carbonateradical with the peroxide generated as a product of therecombination, according to reaction 3b. The pH dependencearound pH ) 12 results from the acid - base equilibrium of H 2 O 2 and HO 2 - , and the very different rates of reactions 4 and 5. 2 Behar et al. discounted the possibility that the carbonate radicalitself might be present in an acid - base equilibrium 6 in the pHrange studied:Lilie et al. 4 applied simultaneous transient absorption andconductivity detection to the study of   · CO 3 - reactions. On thebasis of the conductivity transients, these authors ruled out thesecond-order recombination mechanism 3b in favor of 3a. Theyprovided no explanation for the pH-dependent deviation fromsecond-order kinetics seen by other workers.Chen, Cope, and Hoffman 3 also found pH-dependent devia-tion from second-order decay kinetics in alkaline solution. Fromkinetic measurements in the presence of added solutes theauthors concluded that the p K  a  of   · HCO 3  in reaction 6 is 9.6 ( 0.3. However, their measured rate constants of recombinationat high pH were systematically higher than those reported inother studies.Eriksen et al. used the  · CO 3 - radical as a one-electron oxidantin studies of cyclic hydrazines in the pH region reported byChen et al., 3 but found no indication of a possible protonationof the  · CO 3 - radical. 5,6 In further investigations 16 they reporteda range of 7.0 - 8.2 for the p K  a  based on an increase in the initialabsorbance of the radical with pH. This study also first reporteda negative temperature coefficient for the recombination reaction3, although no data was presented. It was proposed that thereaction involves a pre-equilibrium with the short-lived inter-mediate C 2 O 62 - .The negative temperature coefficient was laterconfirmed by Ferry and Fox 12 who found the rate constantincreases again dramatically above 300  ° C.Bisby, et al. 10 investigated the structure of the carbonateradical using time-resolved resonance Raman spectroscopy.They found no change of the radical’s Raman spectrum betweenpH 7.5 and 13.5, and concluded that the p K  a  for reaction 6 mustbe less than 7.5. Using a stop-flow/pulse radiolysis technique,Czapski et al. 11 made measurements of the (relatively slow)reaction of   · OH radical with carbonic acid in acidic solution.They found the same spectrum of carbonate radical with 600nm maximum that is found in neutral and alkaline solutionsand concluded that the radical  · HCO 3  is a strong acid with p K  a <  0. At roughly the same time, Zuo et al. 13 investigated the pHdependence of reactivity of the (bi)carbonate radical in mixedcarbonate/SCN solutions and concluded the  · HCO 3  radical musthave a p K  a  of 9.6. Lymar, Schwarz, and Czapski 14 published arebuttal of the p K  a  claims made in Zuo et al., 13 Chen et al., 3 and Eriksen et al., 16 showing how each of these studies hadreached an incorrect conclusion, and reiterated their earlierconclusion that p K  a  for reaction 6 is less than 0.It would seem that the work of Lymar et al. 14 and the Ramanstudy of Bisby et al. 10 had settled the issue of   · CO 3 - radicalstructure and p K  a . However, the U. Tokyo group 15 thenpublished a survey of carbonate radical yield and second-orderdecay in a wide concentration range of carbonate and bicarbon-ate solutions up to 350  ° C. To explain apparent anomalies inthe yield and decay, Wu et al. 15 proposed that the carbonateradical is normally present in the form of a dimer anion, either · (CO 3 ) 23 - or  · H(CO 3 ) 22 - formed via the following reactionprocesses:which are converted between the acid and base forms via:Using this model and equilibrium constants determined fromtheir data, Wu et al. 15 calculated the p K  a  of the dimer in reaction9 to be 9.30  (  0.15. The authors speculate that the p K  a  of 9.5 - 9.6 previously determined for the assumed  · HCO 3  at roomtemperature is actually the value for  · H(CO 3 ) 22 - and that thedimer model is quite reasonable by analogy with other radicalshaving the  · X 2 - formula.Upon considering the Tokyo group’s evidence and kineticmodel, we found two serious flaws. The first of these is theomission, in considering the acid - base equilibria of thecarbonate/bicarbonate/carbonic acid system, of the additionalequilibrium 12 to form aqueous CO 2  from carbonic acid:At elevated temperature this is a serious flaw: at 250  ° C, lessthan half of the expected scavenger concentration is actuallypresent in the form of bicarbonate ion (see below). Second, forlow concentrations of the bicarbonate or carbonate scavengerfor  · OH radicals, the  · OH radical persists for a relatively longtime in the presence of   · CO 3 - radicals. In this situation thecross-recombination of   · CO 3 - radicals with  · OH cannot beignored as it was by Wu et al. 15 In the following sections we describe and evaluate a kineticmodel using the ideas laid out in the last paragraph and fit theTokyo group’s data to estimate the  · CO 3 - + · OH reaction rate.We use the model to fit pulse radiolysis/transient absorptionexperiments and explain carbonate radical yields, without theneed to postulate a dimer radical form. Then, we consider thecurious observations of second-order decay rate constant of the carbonate radical in NaHCO 3  and Na 2 CO 3  solutions. Themechanism for the second order recombination has never beenfully and convincingly worked out. We propose a mechanismthat is fully consistent with the data. · CO 3 - +  H 2 O 2 f  · HO 2  +  HCO 3 - (4) · CO 3 - +  HO 2 - f  · O 2 - +  HCO 3 - (5) · HCO 3 - T H + + · CO 3 - (6) · CO 3 - +  HCO 3 - T  · H(CO 3 ) 22 - (in NaHCO 3  solution)(7) · CO 3 - +  CO 32 - T  · (CO 3 ) 23 - (in Na 2 CO 3  solution) (8) · H(CO 3 ) 2 • 2 - T H + + · (CO 3 ) 2 • 3 - (9)HCO 3 - T H + +  CO 32 - (10)H 2 CO 3 T H + +  HCO 3 - (11)H 2 CO 3 T CO 2  +  H 2 O (12) Carbonate Radical Formation  J. Phys. Chem. A, Vol. 114, No. 5, 2010  2143  II. Experimental SectionExperiments Performed at Argonne National Laboratory. Pulse radiolysis/transient absorption experiments were carriedout using 4 - 40 ns pulses from the Argonne ChemistryDivision’s 20-MeV electron linac, where longer pulse widthcorresponds to higher applied dose. The differing pulse widthsused have negligible impact on the observed kinetics, whichoccur on a time scale of several microseconds. The high-temperature/pressure sample cell, flow system, and basicexperimental setup and characteristics were described in previ-ous publications. 17 - 24 Normal temperature and pressure stabili-ties were ( 0.2  ° C and ( 0.1 bar, respectively. Analyzing lightfrom a pulsed 75 W xenon lamp (Photon Technology Interna-tional) was selected using a 40 nm bandwidth interference filter(Andover Corporation) centered at 600 nm, the wavelengthcorresponding to the maximum absorption of   · CO 3 - (  max  at600 nm ) 1970 M - 1 cm - 1 ). 14 The same filter was used for alldata collection, as the maximum absorption is known not toshift with temperature changes. 15 A silicon photodiode (FFD-100, EG&G) was used for detection. Kinetics were measuredat 100, 200, 225, and 250  ° C.Sodium bicarbonate (NaHCO 3 ) solutions were prepared at aconcentration of 0.0200 M (sodium hydrogen carbonate, Aldrich,99.99% + used as received) in deionized water (18.2 M Ω -cm,Barnstead Nanopure cartridge system). Water was degassed withdry nitrogen gas for ∼ 30 min prior to adding sodium bicarbonateto avoid contamination by carbonate ions arising from possiblecarbon dioxide absorption. For experiments, the NaHCO 3 concentration was controlled by mixing with a separate deion-ized water sample using two separate HPLC pumps (Alltech301). The water sample was also previously degassed withnitrogen. Immediately before performing kinetic experiments,both the NaHCO 3  solution and deionized water samples werepurged with N 2 O gas for ∼ 30 min, giving an N 2 O concentrationof 0.024 m. Samples were then kept under an N 2 O atmospherethroughout experiments.Two or three concentrations of NaHCO 3  were used at eachtemperature, as diluted by mixing with deionized water. At 100 ° C, concentrations were 0.00188, 0.00313, and 0.00625 m(molal); at 200  ° C, concentrations were 0.00177, 0.00313, and0.00625 m; and at 225 and 250  ° C, concentrations were 0.00313and 0.00625 m. Total system flow rate was 1.6 mL/min, andthe experimental pressure was 250 bar. Multiple doses fromthe linac were used at each temperature in order to extract rateinformation due to influence of the second-order chemistry. Fourto six doses were applied at each temperature, producing final · CO 3 - concentrations between  ∼  0.7 and 5  ×  10 - 6 m. Experiments Performed at Notre Dame Radiation Labo-ratory.  Pulse radiolysis/transient absorption experiments werecarried out using 4 - 50 ns pulses from the Notre Dame RadiationLaboratory’s 8 MeV electron linac, where longer pulse widthcorresponds to higher applied dose. An OLIS RSM-2000 time-resolved spectrometer was used to record the kinetics, with dualbeam fixed wavelength detection at 600 nm.Sodium bicarbonate (NaHCO 3 ) solutions were prepared at aconcentration of 0.0400 M (sodium hydrogen carbonate, Aldrich,99.998%  +  used as received) in purified deionized water(resistivity 18 M Ω -cm, total organic carbon  < 5 ppb as CO 2 ,Serv-A-Pure Co. cartridge system.)Sodium carbonate (Na 2 CO 3 ) solutions were prepared at aconcentration of 0.0400 M (sodium carbonate, Aldrich, 99.98% +  used as received.)Water was sparged with N 2 O gas for ca. 30 min prior toadding sodium bicarbonate or sodium carbonate to avoidbubbling out CO 2 . Final concentrations were generated bydiluting the stock with N 2 O-bubbled deionized water. A 1cm 3 portion of solution was then measured into a disposableUV-grade plastic (methyl methacrylate) cuvette and sealedwith a Parafilm cover. The sample was very briefly spargedwith N 2 O again using syringe needles inserted into thesample. Care was taken not to change the carbonate orbicarbonate concentration and pH by bubbling out CO 2 . Thesample volume was carefully limited to 1 cm 3 in order to becertain the entire solution obtained a uniform dose from theelectron beam: on time scales of hundreds of millisecondsand longer, convection can be expected to relax anyconcentration gradients produced, which would distort kineticanalysis.Each sample was typically irradiated with five linac pulseswith a transient absorption recorded for each. The first fourwere intended to record the kinetic decay out to baseline atroughly one-half of a second. It was possible to observe thebuildup of reactive product (e.g., hydrogen peroxide) asthe decay became shorter on each pulse. On the fifth pulsethe time base was set to the maximum 1  µ s/point resolutionin order to more precisely set the initial amplitude and timezero for kinetic analysis. A fresh cuvette was used for eachsample set. The empty cuvettes displayed no transientabsorption or fluorescence. III. Results and DiscussionFormation Rate and Yield of the Carbonate Radical.  InN 2 O-saturated solutions of carbonate and bicarbonate ions, thecarbonate radical is formed by water radiolysis in the followingsequence of reactions:Assuming a sufficient concentration of carbonate orbicarbonate scavenger, all of the radicals initially formed byradiolysis can be converted to  · CO 3 - , and this method hasbeen used to estimate the total radiolysis yield up to hightemperature (in alkaline solution). 25 Buxton, et al. 7 measuredthe rise of the  · CO 3 - radical absorbance in carbonate/ bicarbonate solutions up to 250  ° C. They deduced thatreactions 1 and 2 have nearly identical Arrhenius activationenergy, but reaction 2 has 50 times smaller pre-exponentialfactor than reaction 1. Their analysis strongly relied on thetemperature-dependent acid - base equilibrium constants forthe bicarbonate ion:The equilibrium in their experiment was controlled by theaddition of NaOH. H 2 O  +  radiation f  · H, (e - ) aq , · OH, H 2 O 2  (13) · H  +  OH - f (e - ) aq  +  H 2 O (14)(e - ) aq  +  N 2 O f N 2  +  OH - + · OH (15) · OH  +  CO 32 - f  · CO 3 - +  OH - (1) · OH  +  HCO 3 - f  · CO 3 - +  H 2 O (2)HCO 3 - T H + +  CO 32 - (10) 2144  J. Phys. Chem. A, Vol. 114, No. 5, 2010  Haygarth et al.  In “natural pH” solutions of carbonate or bicarbonate ions atelevated temperature, the equilibrium strongly shifts away fromcarbonate ion toward bicarbonate and carbonic acid. In fact,the equilibrium shifts even further toward the CO 2  product of reaction 12.Table 1 indicates, for natural pH solutions of sodium bicarbonateof various concentrations, the fraction that actually remains inthe form of bicarbonate ion. At 250  ° C, roughly 50% of thebicarbonate has converted to dissolved CO 2  and carbonic acid.As demonstrated by Czapski et al., 11 the reaction rate of   · OHradical with H 2 CO 3  is 75 times smaller than reaction 2 at roomtemperature and can probably be ignored.Wu et al. 15 published an extensive survey of the yield of carbonate radical versus temperature in a series of natural pHcarbonate and bicarbonate solutions of widely varying concen-tration (0.1 - 0.001 m). Their data for the initial absorbance of bicarbonate solutions is replotted in Figures 1 and 2, for samplesirradiated with 30 Gy/pulse, and assuming extinction coefficient1700 M - 1 cm - 1 at 633 nm independent of temperature. It isvery clear that as concentration of bicarbonate is lowered from 0.1 m, the yield of   · CO 3 - radicals decreases, that is, scavengingof the  · OH radicals is incomplete. A similar but morecomplicated figure was published for the case of CO 32 - solutions, where due to the higher scavenging rate of reaction1, incomplete scavenging occurs at lower solute concentrations. 15 In this situation of incomplete scavenging of   · OH, a quantitativedescription of the formation kinetics requires the followingadditional reactions: · OH  + · CO 3 - f HOOCO 2 - (17) · H  + · OH f H 2 O (18) · H  + · CO 3 - f HCO 3 - (19) Wu et al. 15 simulated their experimental results in bothcarbonate and bicarbonate solutions from 25 to 250  ° C butreported they were only able to obtain quantitative agreementat the highest concentration, that is, in the limit of completescavenging. To better explain their results, Wu et al. postulatedthat the carbonate radical actually exists as a dimer, that is, Figure 1.  Absorbance of the carbonate radical in NaHCO 3  at 633 nmvs temperature. The markers represent data of Wu et al., 15 and the linesrepresent simulations, as described in this work. TABLE 1: Percentage of Bicarbonate Remaining inSolution % of bicarbonate remaining T   ( ° C) 0.001 M 0.002 M 0.02 M 0.1 M0 98.6 98.6 98.6 98.625 97.1 97.1 97.1 97.150 96.5 96.6 96.6 96.675 95.3 95.7 95.9 95.9100 93.3 94.2 94.8 94.8125 90.1 91.8 93.2 93.3150 85.5 88.1 90.8 90.8175 79.1 82.9 86.9 87.0200 70.5 75.4 80.7 80.9225 59.3 65.0 70.7 70.9250 45.5 50.9 55.7 55.8 H 2 CO 3 T H + +  HCO 3 - (11)H 2 CO 3 T CO 2  +  H 2 O (12) Figure 2.  Comparison of the simulated and measured yields for thecarbonate radical in NaHCO 3 . • Wu, et al. 15 experiment;  0  kineticmodel, this work. · OH  + · OH f H 2 O 2  (16) Carbonate Radical Formation  J. Phys. Chem. A, Vol. 114, No. 5, 2010  2145  · CO 3 - reacts with the parent compound to give  · (CO 3 ) 23 - and/ or  · H(CO 3 ) 22 - . The p K  a  around 9.5 attributed to  · HCO 3  radicalin other studies would then correspond to the dimer radical,and the yield behavior illustrated in Figures 1 and 2 can beapproximately explained in terms of the dimerization equilibriaand the difference in absorbance of the dimer and monomerforms.Objections to this proposal have already been mentioned inthe introduction. The apparent p K  a  at 9.5 was already shown tobe an experimental artifact. 14 Resonance Raman data shows thatthe absorbing radical species has  C  2 V  symmetry regardless of pH, with no proton frequency apparent. 10 The most compellingobjection is that the data can be explained with a more completemechanism. Although Wu et al. 15 made use of the carbonate/ bicarbonate/carbonic acid equilibria 10 and 11, they failed toinclude equilibrium 12 between CO 2  and carbonic acid. Table1 demonstrates the importance of this equilibrium. (Solution of the coupled equilibrium equations is given in the SupportingInformation.) Of greater importance still, they failed to includein their (monomer) simulation the radical recombination reaction17 between  · OH radical and  · CO 3 - . As we now show, the dataof Wu et al. can be fit to extract this rate constant as a functionof temperature.The carbonate radical absorbance at 633 nm versustemperature for irradiation of sodium bicarbonate solutionswas extracted from Figure 5a presented by Wu et al. 15 andwas replotted and then smoothed with a third-degree poly-nomial in order to obtain the absorbance at 25  ° C intervals.These numbers are plotted in Figure 1. For bicarbonateconcentrations of 0.001, 0.002, 0.005, 0.01, 0.02, and 0.1 m,a system of kinetic equations corresponding to reactions 1,2, and 13 - 19 was integrated for the temperature range upto 250  ° C. (A summary list of the reactions can be found inTable 2.) Integration was performed with a fifth-orderRunge - Kutta algorithm with adaptive step size in the IGORsoftware package of Wavemetrics, Inc. Yields of the freeradicals vs temperature in eq 13 were taken from Elliot. 26 Reaction 14 was taken from Marin et al. 27 and Han et al. 28 The rate constants for scavenging of the OH radical, reactions1 and 2, were taken from Buxton et al. 7 Reaction 16 rateconstants were published recently by Janik et al. 29 Reactions 18 and 19 determine the fate of the  · H atom,which is the primary uncertainty of the mechanism (plus orminus 10% of the yield) in N 2 O-saturated bicarbonatesolutions. The temperature-dependent rate constants forreaction 18 were taken from the review of Elliot. 26 For lack of any data, the same numbers were used for reaction 19,which we assume will occur with similar rate constant. Theeffect of the two reactions is to reduce the yield of   · CO 3 - by the yield of   · H atoms. In high temperature alkalinesolutions, reaction 14 may convert  · H atoms to (e - ) aq  beforereactions 18 or 19 occur. In this limit, the yield of   · CO 3 - may be augmented by the yield of   · H atoms.The experimental  · CO 3 - radical yields of ref 15 cannotbe fit when reaction 17 is not included. The self-recombina-tion rate of   · OH radicals provides insufficient competitionto produce the low radical yields observed. We found a bestvalue for the reaction 17 rate constant by slightly adjustingthe dose to agree with the highest bicarbonate concentrationand then adjusting the rate of reaction 17 to agree with the TABLE 2: Kinetic Parameters of Carbonate Radical Growth and Decay solution [HCO 3 - ] [CO 32 - ] pH ionic strength  k  3  ( ( 10%)  k  4  +  k  5  ( ( 50%)  k  22  +  k  23  ( ( 50%)bicarbonate0.10 M 0.097 1.50 × 10 - 3 8.4 0.10 M 7.5 × 10 6 8 × 10 5 3.0 × 10 5 0.04 M 0.0388 5.99 × 10 - 4 8.4 0.04 M 6.3 × 10 6 8 × 10 5 3.0 × 10 5 0.02 M 0.0194 2.99 × 10 - 4 8.4 0.02 M 5.7 × 10 6 8 × 10 5 3.0 × 10 5 carbonate0.0050 M 0.94 × 10 - 3 4.06 × 10 - 3 11.0 0.014 M 5.4 × 10 6 3.5 × 10 6 2.0 × 10 6 0.0025 M 0.64 × 10 - 3 1.86 × 10 - 3 10.8 6.85 × 10 - 3 M 5.1 × 10 6 2.0 × 10 6 1.4 × 10 6 0.0010 M 0.63 × 10 - 3 0.37 × 10 - 3 10.6 2.63 × 10 - 3 M 4.8 × 10 6 1.3 × 10 6 1.3 × 10 6 reaction rate constant (M - 1 s - 1 ) sourceCarbonate Growth Kinetics1  · OH  +  CO 3 - f  · CO 3 - +  OH - 4.0 × 10 8 ref 72  · OH  +  HCO 3 - f  · CO 3 - +  H 2 O  1.0 × 10 7 ref 714  · H  +  OH - f (e - ) aq  +  H 2 O  2.5 × 10 7 ref 2815  (e - ) aq  +  N 2 O f N 2  +  OH - + · OH  9.0 × 10 9 ref 3816  · OH  + · OH f H 2 O 2  4.3 × 10 9 ref 2917  · OH  + · CO 3 - f HOOCO 2 -  6.5  (  1.5 × 10 9 fitting, this work 18  · H  + · OH f H 2 O 9.7 × 10 9 ref 2619  · H  + · CO 3 - f HCO 3 - 2.5 × 10 9 estimate, this work Carbonate Decay Kinetics3  · CO 3 - + · CO 3 - f CO 2  +  CO 42 - k  0  )  4.25  (  0.4 × 10 6 fitting, this work 4  · CO 3 - +  H 2 O 2 f  · HO 2  +  HCO 3 - 8 × 10 5 ref 25  · CO 3 - +  HO 2 - f  · O 2 - +  HCO 3 - above fitting, this work 20  · CO 3 - + · O 2 - f CO 52 - k  0  )  2.0  (  0.2 × 10 8 fitting, this work 22  · CO 3 - +  HCO 4 - f HCO 3 - + · CO 4 - above fitting, this work 23  · CO 3 - +  CO 42 - f CO 32 - + · CO 4 - above fitting, this work 24  · CO 3 - + · CO 4 - f C 2 O 72 - k  0  )  1.0 × 10 9 estimate, this work  2146  J. Phys. Chem. A, Vol. 114, No. 5, 2010  Haygarth et al.
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