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J. Electroanal Chew 198 (1986) 319-330Elsevter &quota S.A., Lausanne - Prmted m The Netherlands

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THE REDUCTION OF CHLORATE AND PERCHLORATE IONS AT ANACTIVE TITANIUM ELECTRODE

GILBERT M. BROWNChemutry Diorsron, Oak Ridge Natzonal Laboratory Oak Ridge. TN 37831 (U.S.A.)(Recetved 3rd July 1985; in revised form 2nd September 1985)

ABSTRACTPerchlorate ion is electrochemically reduced to chloride ion at an active tttanium electrode in aqueous

1.0 M HClO,. The reductton occurs by direct reaction at the surface rather than a pathway involvingcatalysis by soluble titanium corrosion products. The reaction occurs by oxygen atom transfer to thetttanium surface, and the implicattons of this mechanism for the surface composition of active tttaniumelectrodes are discussed. Chlorate ion is also reduced at titanium, and the rate constant for chloratereduction is at least lo5 times greater than that for perchlorate reduction.

INTRODUCTIONAlthough perchloric acid is thermodynamically a powerful oxidizing agent, the

direct reduction of perchlorate ion at an electrode surface is unusual. In fact, theperchlorate salts of tetraalkylammonium ions or alkali metal ions are frequentlyemployed as supporting electrolytes for non-aqueous electrochemistry. The cathodicprocess limiting the usefulness of alkali metal perchlorates as supporting electrolytein acetonitrile is the reduction of the alkali metal ion rather than ClO; [l].Nevertheless there are a few reports [2-111 in the literature of the direct reduction ofperchlorate at an electrode, although most of these are poorly characterized. Ad-ditionally, only a few examples of the direct reduction of chlorate ion have beenpublished [7,12,13]. A number of relatively well characterized examples of the metalion catalyzed electrochemical reduction of ClO, and ClO; have been reported[14-171.

The reduction of perchlorate ion has been reported at platinum [2,3], tungstencarbide [4], ruthenium [5], carbon impregnated with Cr,O, or Al,O, [6], aluminum[7,8], and titanium [9-111 electrodes. The reactions at Pt, WC, and Ru are thereduction of adsorbed ClO; by a surface hydride species. The reduction of ClO;was observed during oxide film breakdown at Al or Ti with a high anodic potentialapplied to the interface. There are also reports of the oxidation of Ti in non-aqueousmedia in which perchlorate is the source of oxygen for film growth, The directreduction of chlorate was observed at platinized titanium [12], stainless steel [12],

0022-0728/86/$03.50 0 1986 Elsevier Sequoia SA


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Armco Iron [13], and aluminum [7]. The reduction of chlorate is catalyzed bynumerous metal ions including those of MO, W, V, and Ti. The catalysis bymolybdenum-catechol complexes has been discussed recently [17]. There are also afew reports of the catalyzed reduction of perchlorate ion by metal ions [14]. Themetal ion catalyzed pathway should be available any time a metal ion can bereduced at an electrode surface to generate a species which will reduce perchlorate orchlorate. The reduction of perchlorate ions by metal ions in solution has beendiscussed recently by Taube [18].

The results reported in this paper are an outgrowth of a study of the influence ofanions on the anodic dissolution of titanium [19]. Perchlorate ions are generallyconsidered to be weakly complexing in solution and to be weakly adsorbed atelectrode surfaces. Perchlorate ions do not appear to assist the dissolution oftitanium; however, at low concentrations of activating anion (chloride, sulfate, oroxalate) in a perchlorate electrolyte the stationary-state current-potential behaviorof titanium changes. This change in behavior is attributed to the reduction ofperchlorate at the titanium surface.EXPERIMENTAL

Aqueous solutions of perchloric and hydrochloric acids were prepared from BakerUltrex concentrated acids. The 72% perchloric acid has a stated chloride impurity of< 0.5 ppm so that 1.0 M HClO, solutions should have < 2 X low6 M Cl-. Houseline distilled water was further purified by triple distillation: first from acidchromate, then from alkaline permanganate, and finally from an all quartz still andcollected in quartz containers. Sodium chlorate was Fisher certified grade, andsolutions were prepared by adding weighed amounts to the desired acid solutions.Chlorate ion reacts with chloride ion to produce intermediate oxidation states ofchlorine. On the basis of literature data [20], dilute solutions of chlorate in 1 M HClshould decay with a half life of about 4 x lo5 s at 30C. Accordingly, solutions ofNaClO, in HCl were prepared immediately before use and kept in an ice bath whilebeing degassed. Solutions were analyzed for chloride ion by Volhards method asoutlined by Vogel [21]. Titanium was estimated by converting it to the peroxycomplex and analyzing spectrophotometrically [22].The electrochemical cell was constructed of glass and teflon and had threecompartments (reference, test, auxiliary) separated by ungreased precision-groundglass stopcocks. The test and reference compartments were jacketed, and thetemperature was maintained at 3O.OC with a circulating water bath. The testcompartment had a flat bottom, and solution agitation was achieved with a magneticstirrer and a teflon-coated spin bar. Dissolved oxygen was removed and protectionfrom atmospheric contamination was maintained by continuously purging the cellsolutions with hydrogen gas. The source of this gas was a GE model 15EHGhydrogen generator. A PAR Model 173 or PAR Model 371 potentiostat was used forpotential control and current measurements. The current output was monitored witha Keithley Model 171 or Model 177 digital multimeter and simultaneously recorded


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on a Hewlett-Packard Model 7100B strip chart recorder. In solutions containingperchlorate ion, commercial SCE reference electrodes could not be used becauseKClO, precipitated in the ceramic or fiber junction, and potential measurementsbecame erratic. Therefore the reference electrode used in this work was a saturatedsodium chloride calomel electrode (NaSCE). The saturated KC1 fill solution of acommercial ceramic junction SCE was removed and replaced with saturated NaCl.The potential of this electrode is 9 mV positive of the SCE at 30C.

Titanium electrodes were fabricated from polycrystalline 6.4 mm diameter rods.The source of titanium was MRC (zone refined, MARZ grade) or Alfa Inorganics(99.98% pure); the stationary-state current-voltage behavior of these materials wasidentical. Internally threaded titanium samples were held on the end of a glass orKel-F electrode holder with a threaded stainless steel rod, and the titanium to holderseal was made leak tight with a teflon washer. A Kel-F nut and teflon washer sealedthe other planar surface of 12.7 mm long electrodes such that only the cylindricalsurface was exposed to solution. Larger samples (63.5 mm long) had the cylindricalsurface and the bottom planar surface in contact with solution. Titanium electrodeswere pretreated to insure reproducible current-voltage behavior. Samples weremechanically and/or chemically polished and then held at constant potential in 0.5M H,SO, until the current reached a stationary state and the parameters were inreasonable agreement with those reported by Kelly: potential maximum, E,,, =-530 mV vs. SCE; current at potential maximum, i,,, = 54 PA cmp2; and opencircuit corrosion potential, E,, = -740 mV vs. SCE [23,24]. Several days wereusually required for a new or freshly polished electrode to come to a stationary state.The electrodes were then removed from 0.5 M H,SO, solution, rinsed well withtriply distilled water, and put in the medium of interest or stored in an air filledbottle. With subsequent use, electrodes would reach a stationary state current at anactive potential in a few hours.RESULTS

Stationary-state current-voltage data for titanium inand mixtures of HCl and HClO, at a total acidity of 1.0

1.0 M HCl, 1.0 M HClO,,M are shown in Fig. 1. The

current-voltage behavior for the hydrogen evolution reaction (HER) on titanium isincluded in this figure, using the data of Kelly and Bronstein [24] and corrected tothe acidity of this experiment (1.0 M H+). The potential range shown in Fig. 1encompasses the hydrogen evolution reaction (HER), the active dissolution state oftitanium, and the active-passive transition region [23]. In 1.0 M HCl, as thepotential is made more positive from open circuit, the current increases withpotential until a critical potential is reached. The current then decreases withpotential. This critical potential (E,,,) divides the active and the active-passivetransition potential regions. The anodic dissolution current density in the active andactive-passive transition regions decreases as the concentration of chloride iondecreases at constant acidity. At a chloride ion concentration of less than approxi-mately 0.1 M, the current in this entire potential region is net cathodic. The


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-400 -500 - 600 -700 -800E /mV vs. NaSCE

Frg. 1. Plot of the logarithm of the stationary-state current vs. potential for tttanium at an acidity of 1.0 Mm HCl. HClO,. and mixtures of the two acids. (0) 1.0 M HCl; (0) 0.5 A4 Cl-; (A)0.2 M Cl-: (A) 0.1 MCl-; (0) 1.0 M HClO,. In solutions containing Cl-. the solid lines are the calculated net currentassuming a partial cathodic process with a 120 mV per decade current-potential dependence and a partialanodic process given by Kellys equatton [23] for the anodtc dissolution of titanium (fitted by a non-hnearleast squares routine). (- - - ) Calculated current for the HER [24].

influence of chloride and other anions on the dissolution of titanium will bediscussed in a future publication [19b]. The cathodic current density in 1.0 M HClO,is nearly a factor of 7 larger than the current for the HER in 1.0 M HCl. Kelly andBronstein found that the current density for the HER reaction is dependent onsolution pH and independent of the electrolyte medium (1.0 M chloride or 0.5 Msulfate) [24]. Therefore it does not seem reasonable to expect the HER kinetics tochange by such a large factor in going from chloride or sulfate to perchlorate. Atpotentials greater than -700 mV vs. NaSCE, where the contribution to the netcurrent by the anodic dissolution reaction is minimized, the observed cathodiccurrent density (or that extrapolated from more positive values) increases withincreasing concentration of perchlorate. Although the dependence is non-linear, itseems reasonable to attribute this increase in cathodic current density over that ofthe HER alone to perchlorate reduction.

Controlled potential electrolysis experiments in 1.0 M HClO, with analysis of thesolution for chloride ion indicate that perchlorate is reduced. Although these


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TABLE 1Current den&es and yields of chlonde ion m 1.0 M HCIO, (30C). Electrode area, A =12.73 cm;solution volume. V = 100 mlE/(mV vs. NaSCE) 103[Cl_ J/M 10-%/s r/pAcrn-* I,,,,./~A cm- b- 625 1.27 3.47 22 21- 600 1.61 7.7 13 17- 525 i 0.1 8.53 _ -3a Determmed from the coulometnc relationship. I = (nFV[CI- I)/( At ).b Difference between the observed current density and that measured for the HER m ref. 24.

experiments used a relatively large electrode, the electrolysis required a few days tobuild up sufficient product. The yield of chloride ion as a function of potential in 1.0M HClO, is reported in Table 1. The current was recorded continuously and did notvary more than f4%. The difference between the observed current density and thatcalculated for the HER alone is also given in Table 1. The current density forperchlorate reduction was calculated from a Coulombs law expression, assuming thestoichiometry is given by eqn. (1).ClOi + 8 e--t 8 H+-t Cl-+ 4 H,O (1)The current density, determined from chloride analysis, is in good agreement withthe difference in the observed current density and that calculated for the HER alone.The polarization curves in Fig. 1 can be accounted for as the summation of threepartial reactions: anodic dissolution of titanium, hydrogen evolution, and reductionof perchlorate to chloride.

Several groups have reported [23,25,26] that the current density for titaniumcorrosion in H,SO, or HCl solutions is independent of stirring rate. It was thereforesomewhat surprising to find a slight stirring dependence of the current in HClO,solutions. At potentials between - 500 and -600 mV vs. NaSCE, a net cathodiccurrent increased with decreasing stirring rate. This stirring dependence becomesinsignificant, however, in the presence of a sufficient concentration of activatinganion such as chloride, sulfate, or oxalate. Preliminary experiments [27] with atitanium rotating disk electrode indicate the dependence of the current on rotationalfrequency saturates at several hundred rpm. The mass transport provided by arotating spin-bar at cylindrical electrodes was sufficient to obtain a limiting current.As will be discussed later, this influence of mass transport on the current densityseems to be due to the mass transport limited reaction of an intermediate in thereduction of perchlorate [27a].

The influence of perchloric acid on the passivation of titanium electrodes hasbeen reported by Romanushkina et al. [15], in HClO, + HCl mixtures. Their studyindicated that perchlorate was not reduced directly at titanium, but rather theinfluence of increasing the concentration of perchloric acid is to increase the anodicdissolution of titanium. The corrosion product is aquo Ti(II1) ions which areoxidized to Ti(IV) by perchlorate.


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The kinetics for reaction (2) have been reported [28], and the indicated rateexpression is given by eqn. (5).- 2 [Ti(III)] = 8-$[Cll] = k,[Ti(III)] [ClO;] (5)The rate of formation of Ti(II1) in solution by reaction (3) is i,A/3FV where i, isthe anodic current density for the dissolution of titanium (reaction 3), V is thesolution volume, and A is the electrode area. The rate of disappearance of Ti(II1) byreaction (2) is k,,,[Ti(III)] where k,, is the pseudo-first-order rate constant forreaction (2) (kobs = k,[ClO;]). The rate of appearance of Ti(II1) in solution byreaction (4) is (k,A/V)[Ti(IV)] where k, is the electrochemical rate constant forreduction of Ti(IV) at the electrode surface. The rate of appearance of Ti(IV) byreaction 2 is k,,,[Ti(III)]; the rate of disappearance of Ti(IV) by reaction 4 is(k,A/V)[Ti(IV)]. These considerations lead to the coupled differential eqns. (6) and(7).z [Ti(III)] = (i,A/3FV) - k,,,[Ti(III)] + (k,A/V)[Ti(IV)]2 [Ti(IV)] = k,,,[Ti(III)] - (k,A/V)[Ti(IV)]These coupled differential equations were solved by the method of Laplace transfor-mation, with the boundary condition [Ti(III)] = [Ti(IV)] = 0 at time zero.S[=E(III>] = (i,A/3FV)l/s - kobs[~(III)] +(k,A/V)[E(IV)] (8)s[Z(IV)] = kobs[~(III)] -(k4A/V)[n(IV)] (9)


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where s is the Laplace variable and a superscript bar indicates the concentration inthe Laplace plane. Equation (9) is solved for [z(IV)] as a function of [E(III)].[n(Iv)] = { kobs/(S + k,A/lr)} [Ti(III)] (10)Equation (8) with [z(IV)] given by eqn. (lo), is solved for [Ti(III)].

Inverse Laplace transformation of eqn. (11) was obtained from a tabulated standardform [32].LP[(s+a)/s2(.s+~)] =(1/a-a/cy)(l-e-*)+(a/cy)t (12)Therefore the concentration of Ti(II1) ions as a function of time is given by eqn.(13).

kobstkobs + k,W) [1+exp[-(k,,,+k,A/V)1]]

(b%V) t+ (kobs M/V) (13)The exponential term in equation 13 becomes insignificant at the time of observationin Table 1, and it can be neglected. This approximate expression for [Ti(III)] as afunction of time is substituted in eqn. (14) and then integrated using the boundarycondition [Cl-] = 0 at time zero to give an expression for [Cl-] which can becompared with the experimental values.$[Cll] = ik,,,[Ti(III)] (14)Lc1- I = i,k,tdt

ikoba

24FV ( kobs + k,A/V)*+ 1 (k,A/V)t

2 (kotx + k,A/V) (15)The current density for the anodic dissolution of Ti is less than 1 x 10-h A cm-

at -600 mV vs. NaSCE (see Fig. 1). The rate constant for reduction of Ti(IV) toTi(II1) is dependent on the concentration of complexing anions. as noted in afollowing section, and a conservative estimate is 1 X 1O-6 cm SK. The pseudo-first-order rate constant for the reaction of Ti(II1) with perchlorate in 1.0 M HCIO, is3 x lo- s- [18,28]. Using the other parameters appropriate for the experiments inTable 1, the calculated concentration of chloride ion after 8 X lo5 s is - 5 X lo-M, a value far lower than the observed concentration. Thus, perchlorate reduction inthese experiments occurs by direct reduction at the electrode rather than by acatalyzed path involving Ti solution species. It should be noted that the catalyzedpathway is still present, but it is too slow to account for the product yields in these


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experiments. If the electrode area to solution volume ratio (A/V) were increased bya factor of 100, the catalyzed pathway would be significant.

The total cathodic currents in Fig. 1 are the sums of the partial currents for theHER and perchlorate reduction. The acid concentration is constant at 1.0 M, andthe partial current for the HER is assumed to be constant. As previously noted. thepartial current for perchlorate reduction is not linearly dependent on perchlorateconcentration. This current seems to be inversely dependent on the concentration ofchloride ion. When the concentration of chloride exceeds about 0.2 M, the partialcathodic current decreases to the value for the HER alone. Adsorption of perchlo-rate on the surface seems to be a necessary condition for its reduction to occur. Asufficient concentration of chloride ion will displace the adsorbed perchlorate fromthe electrode surface. The current-voltage data in Fig. 1 also show that the decreasein the partial cathodic process for perchlorate reduction is paralleled by an increasein the partial anodic current for titanium dissolution with an increase in chlorideconcentration.Reduction of Ti(IV) at a titanium electrode

An estimate of the rate at which Ti(IV) in solution is electrochemically reduced toTi(II1) at an active titanium electrode is needed to gange the effect of the catalyzedpathway for ClOi reduction. This rate constant has been reported in 1.0 M HCland 0.5 M H,SO, at various potentials by Kelly [29,30]. The rate constant wasdetermined by introducing a relatively large concentration of Ti(IV) to the testcompartment of the electrochemical cell, and measuring the current before and afterthe addition of Ti(IV). The difference in current density as a function of concentra-tion of Ti(IV) is linear, demonstrating the reaction is first-order in Ti(IV). In thiswork, solutions of Ti(IV) in 1.0 M HCIO, were too unstable to precipitation ofhydrous oxides of titanium for this method to be successful. Small amounts ofsulfuric acid stabilize dilute solutions of Ti(IV), probably by complexation. and therate constant for Ti(IV) reduction was measured at several concentrations of sulfuricacid. The value reported at 30C in 0.5 M H,SO, is 7.9 x 10K5 cm ss [29] at - 600mV vs. SCE: rate constants measured in this work were 2.4 X 10m5 cm ss and2.8 X lop6 at -600 mV vs. NaSCE in 1.0 M HClO, containing 1 X lo- M H,SO,and 1 X 1O-4 M H,SO, respectively. An estimate of the rate constant, extrapolatedto 10e6 M H,SO, by a log-log plot of the rate constant versus concentration ofH,SO,, is in the range 3 X 10-cm s-l to 1 X 10e6 cm s-.Reduction of chlorate

The good agreement between the current density for perchlorate reduction andthe appearance of chloride in solution indicates that the concentration of theintermediate species ClO, does not build up to a significant level in solution. Thereaction scheme for perchlorate reduction requires that chlorate be reduced at afaster rate. This was confirmed by independent experiments in which solutions of


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TABLE 2Rate constants for the reaction of perchlorate and chlorate Ions at titamum (30C)E/mV vs. NaSCE k/cm s -la

ClO,- c10; c10;(1 M HClO,) (1 M HClO,) (1 M HCl)

- 600 2.2x10-s 1 7x10m3 1.6x 10-j- 550 7.2X10- 1.2x10-3 1.1 x10-3- 500 1.3x1o-9 o.9x1o-3 1.4x 10-7A The reactions are assumed to be first order in [ClO; ] or (ClO; 1, I = nFAkc.NaClO, in 1.0 A4 H+ were added to the test electrode compartment, and thedifference in current was attributed to chlorate reduction. The reductions of chlorateand perchlorate are assumed to be first order. and the stoichiometries are given byeqns. (1) and (16).ClO; + 6 e-+ 6 H++ Cl-+ 3 H,O (16)The rate constants are summarized in Table 2. The current for reduction of ClO; ishighly stirring dependent and only weakly dependent on potential or ionic medium.The dependence of the current on mass transport was probably not completelyremoved in these experiments. The stirring conditions of these experiments producea diffusion layer thickness, 6, of 1O-2 to 1O-3 cm. The diffusion controlled rateconstant is D/S, 10e3 to 1O-2 cm-. Thus the apparent rate constants for ClO;reductions in Table 2 represent a lower limit only. It can be concluded, however, thatthe rate constant for ClO, reduction is at least 5 orders of magnitude greater thanthe rate constant for ClO,- reduction. The reductions of ClO, and ClO,- arepresumed to occur stepwise with the lower oxidation state 0x0 species (ClO_, CIO- )occurring as intermediates.

DISCUSSION

In a 1966 review article, Pearson stated that the perchlorate ion shows nooxidizing properties in solution [31]. This statement is still correct in the sense thatperchlorate ion must be bound to another species or a surface before it can bereduced to a lower valence of chlorine. Taube [18] has recently reviewed the evidencefor reduction of ClOc in solution. Perchlorate has no low lying unfilled electronicorbitals, and thus it cannot easily accommodate an additional electron to form theC10i2 ion. Reduction necessarily occurs via transfer of an oxygen atom. Theprocess in reactions (17) and (18) can occur if some species is available to stabilize02-Cloy + e-j ClO, + 02-ClO; + 2 e-+ ClO; + 02-

(17)(18)


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Protons can stabilize O-. and this may account for the activity of concentratedsolutions of HClO, as oxidizing agents.

Taube [IS] notes that metal ions are the only species known to reduce Cloy atreasonable rates at ambient temperatures. The most active ions are those formingyl ions such as V(IV). Thus the reduction of Cloy by V(II1) and Ti(II1) areunderstandable if oxygen atom transfer occurs.ClO,- + M( H?O);f + + ClO, + MO( H,O);+ + H,O M = Ti.V (19)Reduction can also occur by a two-electron mechanism such as that for U(II1) asreactant.ClO, + U(II1) + Hz0 -ClO, +UO; +2H+ (20)

The reduction of ClO; on an active titanium electrode can be taken as evidencethat surface bound species are responsible for this rate enhancement. This enhancedrate of perchlorate reduction is consistent with the kinetic model proposed by Kelly[23] for the dissolution of titanium. The surface of Ti in the active and active-pas-sive transition potential regions is covered by a monolayer of adsorbed oxy orhydroxy Ti(II), Ti(III), and Ti(IV) species. At potentials negative to E,,,, Ti(I1) isthe predominant species. The surface has an oxygen vacancy at least, and thepotential dependence (120 mV per decade Tafel slope) of perchlorate reduction isconsistent with a rate limiting one-electron transfer reaction, assuming a transfercoefficient of 0.5. The reaction of perchlorate with the adsorbed Ti(I1) species bytransfer of an oxygen atom is indicated asClO; + (Ti(II)),+ + Hi+ (Ti(III)OH),+ + ClO, (21)(Ti(III)OH),:+ + e-+ H++ (Ti(II))I+ + H,O (22)where subscript a indicates an adsorbed species. The net reaction is reduction ofCloy to ClO, catalyzed by the surface.Cloy + 2 H++ em--, ClO, + H,O (23)Either the one-electron reduction pathway mentioned above or a two-electronreduction pathway is possible.ClOL + (TiOH): + H+-, (Ti(OH),)t+ + ClO, (24)(Ti(OH),)z+ + 2 e-+ H+-+ (TiOH): + H,O (25)The relative stability of the intermediate ClO, will determine whether the reactionwill proceed via the one-electron pathway or directly to ClO; via the two-electronpathway. The same mechanistic arguments that have been made for ClO< reductionapply to ClO, reduction. For the one-electron pathway, ClO, is the indicatedproduct, and it is a relatively stable species. This may account for the greatlyenhanced reaction rate of ClO; over ClOL. The stirring dependence of the currentmentioned earlier is probably due to the mass transport-dependent reduction ofeither ClO, or ClO, at the surface. It may be possible to distinguish between the


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one-electron and two-electron reduction pathways by a rotating disk electrodeinvestigation of the ClOL and ClO, reduction reactions [27b]. The rate of perchlo-rate reduction is retarded by a significant concentration of other anions such as Cl-by a competitive adsorption process. Reactions (21) and (24) may involve thepreliminary adsorption of ClO; at a hydroxy Ti(II) surface species in a reactionsuch as eqn. (26).(Ti(OH),) u + ClO; + Ht + (Ti(OH)(ClO,)). + H,O (26)The dependence of the current on stirring disappears with an increase in theconcentration of activating anions. This observation is consistent with an adsorptionprocess in which these anions (chloride. sulfate, or oxalate) compete with perchloratefor active sites.ACKNOWLEDGEMENT

Research sponsored by the Division of Materials Sciences. Office of Basic EnergySciences, U.S. Department of Energy under contract DE-AC05840R21400 with theMartin Marietta Energy Systems, Inc.REFERENCES

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