Vanadium(III) Sulfate as a Reducing Agent for Determination of Perchlorate DAVID A. ZATKO and BYRON KRATOCHVIL Department o f Chemistry, University o f Wisconsin, Madison, Wis.
b Air-stable, solid V2(SO& quantitatively reduces perchlorate to chloride upon 5 to 10 minutes' refluxing in 7 to 8M HzS04 in the presence of Os04 as catalyst. The determination is completed by potentiometric titration of the chloride formed or by spectrophotometric measurement of the VO+2 formed. Accuracies of 0.1 to 0.2% are obtained for 0.01- to 3-mmole samples by chloride determination. The spectrophotometric method gives accuracies of 1 to 270, with some interference by V(III) absorption. A correlation between the effectiveness of a metal ion as a reductant for perchlorate and the affinity of the oxidized metal for oxygen is noted.
P
analytical methods for perchlorate generally involve gravimetric precipitation ( I @ , reduction to chloride by fusion with a salt (?ri&Cl, Na2C03, Ya202) ( I I ) , or reduction to chloride with lower oxidation states of transition metals. Of these methods, the one using transition metals appears to have the widest applicability (1, 2 ) . Rothmund in 1909 observed that solutions of the lower oxidation states of Cr, Mo, Ti, V, and W reduce perchlorate, and reported Ti(II1) to be the most effective of this group (9). As a result, investigations of metals other than titanium have involved primarily kinetic studies, although some have been studied as catalysts-for example, molybdenum catalysis of perchlorate reduction with zinc or cadmium amalgam ( 5 , 8). =Ilthough Ti(II1) reduction is more rapid than most other methods its use requires provision for oxygenfree storage and handling of solutions. Recently this disadvantage has been mitigated by the use of titanium hydride as a solid reductant ( I ) . I n connection with work on lower oxidation states of vanadium, it was found t h a t perchlorate reduction by V(II1) under the proper conditions is as rapid and as efficient as by Ti(II1). This paper reports conditions whereby fast, quantitative reduction of perchlorate is effected using air-stable, solid v2(so4)3 as thz source of V(II1). I n addition to speed and convenience, vanadium(II1) sulfate has these further advantages: Excess V(II1) does not RESENT
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ANALYTICAL CHEMISTRY
have to be remoyed before chloride titration with AgN03, and a wide range of sample sizes can be accommodated.
Frederick Smith Chemical Co., Columbus, Ohio) and standardized both acidimetrically and by reduction with titanium hydride ( I ) . Ammonium perchlorate (G. Frederick Smith Chemical Co.), vanadium pentoxide (J. T. EXPERIMENTAL Baker Chemical Co., Phillipsburg, Preparation of VP(S04)a. V2(S04)3 N. J.), and titanium hydride powder (Metal Hydrides, Inc., Beverly, Mass.) was prepared by the method of Claunh were used as received. All other chemand Jones (3) with some modifications. icals were reagent grade. Reagent grade VZOs (0.2 mole) and Perchlorate reductions were carried sulfur (0.3 mole) were added with out in a 150-ml. (50-ml. for microstirring to 400 ml. of concentrated titrations) Erlenmeyer flask connected (98%) sulfuric acid in a 500-ml., to a 20-cm. water-cooled condenser by round-bottomed three-necked flask a ground-glass joint. An electric hot equipped with a thermometer, overhead plate was used as heat source for paddle stirrer, and outlet tube. The refluxing. mixture was continuously stirred a t Chloride titrations with silver ni190' to 210' C. for 24 to 36 hours. trate were done Dotentiometricallv with Filtration of the yellow product was a Beckman Model 72 p H meter&nd a very slow, so centrifugation was emsilver billet-saturated calomel electrode ployed. After separation, the solid was washed first with 200 ml. of 9M H2S04, pair. A saturated potassium sulfate salt bridge was used to isolate the then with 200 ml. of water. calomel electrode from the sample Excess sulfur was removed by sussolution. pending the compound in a mixture of A Model 14 Cary recording spectro100 ml. of 1 to 1 ethanol-water and 100 photometer was used for spectroml. of CSZ, stirring vigorously for 30 photometric measurements. minutes, centrifuging, and washing with Procedure. Place a sample consuccessive 100-ml. portions of CSz, taining approximately 1 mmole of absolute ethanol, and ether. The perchlorate in a 150-ml. reaction Vz(so4)3 was air-dried for 3 hours, dried flask. Add 2.5 grams of Vz(SO4)3, 2 in a 120' C. oven overnight, and stored drops of 0.01X osmium tetroxide in in a tightly capped bottle. Yields of 0.131 H2S04, and sufficient sulfuric approximately 95% based on the vaacid to give a final volume of 25 to nadium were obtained. 30 ml. of a solution which is 7 . 5 M The vanadium content of several in HzS04. Attach a water-cooled conbatches was determined by ignition denser, heat the flask contents to boiling to V20s a t 720' C. in a muffle furnace, on a hot plate, and reflux for 5 to 10 and was found to be 94 to 96% of minutes. Remove the hot plate and theoretical [based on pure Vz(SO&]. cool the solution to room temperature. The difference is attributed to the presTransfer the contents to a 250-ml. ence of H2S04or HSO4- in the product. [A compound, ~ V Z ( S O ~ ) ~ . H Z Shas O ~ , beaker and titrate potentiometrically with 0.1-21 -4gNOs. been reported (10) which has a vaFor determinations of perchlorate in nadium content 9401, that of pure quantities on the order of 0.01 mmole, V2(S0&.] The product obtained in use a 50-ml. flask, add sufficient HzS04 the above preparation is more soluble to give a total volume of 10 ml. of 7.5M than pure Vz(SOa)3 and therefore is H2S04,add 0.2 gram of vz(SO4)3, and preferred for perchlorate reduction. reflux as before. Titrate potentioWhen necessary, its solubility may be metrically with 10-3Jf AgN03. increased by exposure to moist air for 2 or 3 hours. Pure V2(so&(100% based RESULTS A N D DISCUSSION on vanadium analysis) was obtained With some of the V~(soi)3preparaby washing the above product with concentrated sodium bicarbonate solution tions scattered results for perchlorate until the filtrate was basic and no were obtained. This scatter was not longer green. It is essentially insoluble reduced by varying the reflux time, acid in boiling water or dilute H2S04and so concentration, or the amount of excess is not satisfactory for perchlorate vanadium(II1) present. Other prepreduction. arations gave quantitative results with Other Reagents and Apparatus. excellent precision. I n all cases, howPerchloric acid solutions, approxiever, quantitative results were obtained mately 0.2M, were prepared from by the addition of a trace of osmium doubly distilled perchloric acid (G.
mo
-
Bo80 70
-
8 aou
y
5040-
01 4
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'
4
3
'
6
'
a %SO,
'
7
"
a
e
'
'
loll
I?.
'
MOLARITY
Figure 1 . Effect of sulfuric acid concentration on per cent chloride recovery
0
5-minute retlux time 10-minute reflux time
tetroxide. The catalytic effect of os04 in perchlorate reduction has been reported with titanium(I1I) also (%). The effect of sulfuric acid concentration on the ratio of chloride found to perchlorate taken (per cent recovery) is shown in Figure 1. The shape of this curve is determined by two factors: the dissolution rate of V2(S04)3 and the temperature of the refluxing solution. Increasing the sulfuric acid concentration of the reflux mixture increases the temperature of reflux and accelerates the reaction of V(II1) in solution with perchlorate. However, a t high acid concentrations Vp(so4)3 is slow to dissolve, This causes the vanadium(II1) concentration in solution during the initial 15 to 30 seconds of boiling to be so low that a portion of the perchlorate is reduced only to elemental chlorine. Chlorine evolution was detected under these conditions by bubbling the condenser gases through a solution of potassium iodide. V2(S04)3 will not completely dissolve in H2S04 solutions 10M or higher. I n fact, a t sulfuric acid concentrations above 12.5M the dissolution process is reversed and hydrated V(1II) is dehydrated according to the equation 2V(HzO),+3
than 3 mmoles are not generally recommended because of the large amount of vanadium(II1) sulfate required. Determinations of small amounts of perchlorate by this method are very accurate; the limiting factor under the conditions used was the accuracy of the chloride titration. I n general a 50% excess of VZ(S04)a over the stoichiometric amount is recommended [approximately 2.4 grams of V2(S0J3 per mmole of perchlorate], but for very small amounts of perchlorate this ratio must be increased. For the reduction of 0.01 mmole of perchlorate 0.2 gram of V~(so4)3was found satisfactory; smaller amounts gave low results. Removal of excess vanadium(II1) prior to titration with AgNO3 is not required, as the reduction potential of the vanadium(1V)-(111) couple is higher thant hat of the AgC1-Agc,) couple. Also, it was observed that vanadium (111)present a t the start of the titration is oxidized by nitrate from the AgN03 titrant. Under the titration conditions this oxidation to vanadium(1V) occurs prior to the chloride equivalence point, so that vanadium(II1) interference with the Ag+-Ag(,) couple is not encountered. Experiments run on the vanadium(II1)-nitrate system gave no definite oxidation-reduction stoichiometry, but indicated that a coupled or induced reaction involving vanadium (111), nitrogen oxides, and oxygen may be occurring. Table I1 compares NH4CI04 reduction by vz(s04)3and TiH2. The results obtained by the two methods are essentially identical. h similar comparison was done for samples of HC1O4; five determinations on a 1.161-mmole sample resulted in an average recovery of 100% with a relative standard devia-
tion of 0.1% for both the vp(so4)aand the TiH2 methods. However, the vanadium procedure has several advantages over the titanium method: The filtration and ammonium persulfate oxidation steps with TiHz are not required; H2S is not. formed during the reduction step as it is with TiHz; and the silver electrode does not become fouled during titrations] as occasionally happens with TiHz reduction mixtures. Completion of Determination by Spectrophotometric Measurement of Vanadium(1V) Formed. A spectrophotometric investigation of the solutions after perchlorate reduction by vanadium(II1) showed t h a t the vanadium was oxidized quantitatively to vanadium(1V). The possibility of using the color of the vanadyl ion for completion of the determination was therefore considered. I n this procedure, the solution after refluxing was cooled, transferred to a volumetric flask, and diluted to volume with 7 M H2S04,and the absorbance was measured a t 750 mp. Excess vanadium(II1) present after reduction interfered to the extent of 2 to 3% by absorbing a t the vanadyl peak (750 mp). A correction was made by reading the vanadium(II1) absorbance a t 400 mp and subtracting 0.0585 [the ratio of vanadium(II1) absorbance a t 750 mp to that a t 400 mp] times this value from the absorbance reading a t 750 mp, Vanadium(1V) absorbance at 400 mp is negligible under the conditions used (4). Results of several spectrophotometric determinations of perchlorate in perchloric acid samples were as follows: mmoles of perchlorate taken, 1.165; found (average of 5 determinations) 1.17; relative standard deviation] 0.5; mmole taken, 0.5836; found (average of
Table I.
Determination of Perchlorate over a Range of Sample Sizes Mmoles bImoles % No. of Rel. std. dev. HCIOl taken C1 found recovery detns. 99.8 3 0.04 2.879 2.873 100.0 7 0.1 1,151 1,151 100.1 4 0.2 0.01151 0.01152
+ 3S04-2 e V2(S04)3(s)+ xH~O
Figure 1 also shows the effect of reflux time. Increasing the reflux time from 5 to 10 minutes gives some improvement in chloride recovery a t low acid concentrations, but has essentially no effect throughout the rest of the acid range. The fact that the recovery a t high acidities is not affected by longer reflux times supports the conclusion that chlorine is lost only during the initial period of reflux. Results of analyses of perchloric acid samples ranging from 0.01 to 3 mmoles are given in Table I. Samples larger
Table 11.
Comparison of TiH2" and Vz(SO& Reduction Methods for Perchlorate in NH4ClOd
TiHl
hrz(so4)3
a
NH4C104, g. taken
Mmoles C1 found
0.1347 0.1286 0.1220 0.1319 0.1280
1.146 1.095 1.038 1.123 1.089
% recovery
99.5 99.5 99.5 99.7 99.6 Av. 99.6 Rel. std. dev., yo 0.1
NH,CIOL, g. taken
Mmoles C1 found
0.2514 0.2765 0.2581 0.2335 0.2493
2.132 2,342 2.192 1.979 2.110
% recovery
99.6 99.5 99.8 99.6 99.4 Av. 99.6 Re]. std. dev., yo 0 . 1
Procedure of (1).
VOL. 37, NO. 12, NOVEMBER 1965
1561
3 determinations) 0.580; relative standard deviation, 0.3; mmole taken (1 determination), 0.0461; found, 0.0475. Air oxidation of vanadium(I11) under the acidic conditions present is slow (7). S o special precautions are required, other than taking the measurements within 20 to 30 minutes of completion of the reduction. This spectrophotometric procedure may be convenient where the perchlorate samples contain substances interfering with silver nitrate titration and accuracy of 1 to 2% is acceptable. For example, chloride if present with perchlorate is titrated with AgNOa in the potentiometric procedure and will cause a positive error, as will other anions forming insoluble silver salts. The spectrophotometric method avoids this difficulty. Chlorine in oxidation states other than chloride, however, such as chlorate or hypochlorite, will be reduced by vanadium (111) and will cause positive errors by either procedure. Rothmund first noted that the effectiveness of a metal in perchlorate reduction is not directly related to the metal reduction potential. However, those metal systems which reduce perchlorate a t an appreciable rate tend to form stable metal-oxygen bonds in their oxidized forms [V(IV)-(III), Ti(1V)-
(111), Mo(V)-(IV), etc.]. In contrast, systems or couples which form less stable metal-oxygen bonds in the oxidized form, such as Zn(I1)-(0), Sn(1V)-(11), and Cr(II1)-(11), are much slower despite generally lower reduction potentials. This suggests that the mechanism may involve an oxygen transfer from perchlorate to the metal. The kinetics of perchlorate oxidation of vanadjum(I1) and (111) a t 50’ have been investigated by King and Garner (6). They were not able t o formulate a satisfactory mechanism for the vanadium(II1) reaction because of the uncertain nature of the hydrogen ion dependence. For the vanadium(I1) reaction two mechanisms were considered, one involving a rate-determining step of oxygen transfer from perchlorate to vanadium and the other an electron transfer from vanadium to perchlorate. Of these mechanisms, that involving oxygen transfer was considered more probable. It is possible then that removal of the first perchlorate oxygen is favored by the formation of a stable bond between the oxygen and the metal reductant. Unfortunately, sufficient data are not available to permit calculation of the metal-oxygen bond energies in solution, so a quantitative comparison cannot be made at the present time.
LITERATURE CITED
(1) Alley, B. J., Dykes, H. W. H., ANAL. CHEM.36, 1124 (1964). (2) Burns, E. A., Muraca, R. E”., Ibid., 32, 1316 (1960). (3) Claunch, R. T., Jones, M. M., “Inorganic Syntheses,” J. Kleinberg, ed.,
Vol. VII, p. 92, McGraw-Hill, New York. 1963. (4) Furman, S. C., Garner, C. S., J . Am. Chem. SOC.72, 1785 (1950). (5) Haight, G. P., Jr., ANAL. CHEM.25, 642 (1953). (6) King, W. R., Jr., Garner, C. S., J . Phys. Chem. 58, 29 (1954). (7) Ramsey, J. B., Sugimoto, R., DeVorkin, H., J . Am. Chem. SOC. 63, 3480 (1941). (8) Rechnit;,’ G. A., Laitinen, H. A,, ANAL.CHEM.33, 1473 (1961). (9) Rothmund, V., 2. Allgem. Chem. 62,
108(1909). (10) Schulek, E., Pais, I., Pataki, L.,
Ann. Univ. Sci. Budapest, Roland0 Eotvos Nominatae, Sect. Chim. 2, 583 (1960); C. A . 56, 6868 (1961). (11) Schumacher, J. C., ed., “Perchlorates,’, pp. 105-9, Reinhold, New York,
1960.
(12) Willard, H. H., Smith, G. M., ANAL. CHEM.11, 186 (1939).
RECEIVEDfor review May 3, 1965. Accepted August 27, 1965. Pittsburgh Conference on Analytical Chemistry and Applied Spectroscopy, Pittsburgh, Pa., March 1965. Work supported in part by the Research Committee of the Graduate School from funds supplied by the Wisconsin Alumni Research Foundation.
Reduction of Oxygen to Superoxide Anion in Aprotic Solvents D. L. MARICLE
and W. G. HODGSON
Research Service Department, Central Research Division, American Cyanamid Co., Stamford, Conn. The reduction of oxygen in aprotic solvents has been studied b y electroanalytical and electron spin resonance techniques. In contrast to the normal 2-electron reduction observed in aqueous systems for the first oxygen reduction process, a 1 -electron reduction of O2 to OzA, the superoxide anion, was found. The addition of a proton source (phenol) converted the process to the normal 2-electron reduction. Preparative scale reductions produced both tetrabutylamrnonium superoxide and potassium superoxide. The latter was shown to release equivalent amounts of oxygen and peroxide when treated with water (a known reaction of KOz). Electron spin resonance confirmed that the products of the preparative scale electrolysis were paramagnetic, as expected for species containing the superoxide ion.
T
electrolytic reduction of oxygen in nonaqueous organic solvents has received relatively little attention, conHE
1562
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ANALYTICAL CHEMISTRY
sidering the extensive literature pertaining to aqueous systems. Coetzee and Kolthoff (2) studied the polarography of oxygen in acetonitrile, employing sodium perchlorate or potassium thiocyanate as supporting electrolyte. On the basis of a calculated diffusion coefficient they reported a2-electron change for the first wave, analogous to the usual aqueous reduction. The polarographic reduction of oxygen has also been reported in dimethylsulfoxide (DMSO) (8) and nmethylacetamide (6), but with no attempt to determine the product or the number of electrons involved. I n the course of conducting cyclic voltammetric experiments in aerated DMSO (12) we observed a quasireversible reoxidation for the first oxygen reduction product. It is very unlikely that any reduction involving the consumption of protons would occur reversibly in a solvent with such low proton availability; therefore, it was concluded that the usual 2-electron reduction to HOz- was not operative.
A 1-electron reduction to the superoxide anion, 02”,not involving protons was postulated as the most reasonable explanation of the cyclic voltammetric results. Russell has also suggested 02’ as the product formed during basecatalyzed reactions of O2 with organic mono- and dianions in DMSO (14). Superoxide salts have also been found as one of the products of electrolytic oxygen reduction in liquid ammonia (10).
This paper reports the detailed investigation of the electrolytic reduction of oxygen in aprotic organic solvents. Since the superoxide ion is paramagnetic, electron spin resonance was used in support of the electrochemical measurements to establish with greater certainty the nature of the initial reduction product. EXPERIMENTAL
The cyclic voltammetry and polarography were carried out using an
