with a 5y0 ammonium citrate solution adjusted to p H 3.8 as eluent. At this pH, strontium remained at the origin, and the complexed yttrium moved to a ne11 defined spot a t R , 0.75. The phosphorylated cellulose paper with an HC1 eluent has been used for numerous cation separations ( 2 ) and provided an excellent separation with stronium mol-ing to R, 0.89. Figure 3 s h o w that the beta particles of the separated Ym can be used to excite K-radiation in target elements-in this case silver. The data for the figure were taken on a Radiation Instrument Development Laboratory Inc. Model 34-12,400 channel pulse height analyzer. X determination of the half life of YgO from decay measurements taken for several lives was made. The decay data were analyzed by the method of least squares, and a half life of 65.0 hours with a standard deviation of 0.3 hour was obtained. By following the decay or growth of the Y90,it is apparent that one can de-
termine how much of either element mas present a t zero time or at any given time after separation. These aspects of the experiments mill be discussed elsewhere.
quantities of material. Another important advantage is that radioactive detection is not interfered n-ith by the paperf. ACKNOWLEDGMENT
The paper-solvent combinations listed in Table I represent an exploratory attempt a t this separation. Since two of the separations were highly successful, no further investigation of such variables as p H , ionic affinities, complexing, or pretreatment of the paper \\-as undertakcn. It seems likely that any ion exchange method of separating strontium and yttrium can be adapted to a paper technique. Such a technique as compared Ivith a method using an ion exchange column requires a shorter time for the separation, eliminates fraction collection, and allows greater ease and safety in handling. It is possible to vork with very small
The author thanks Lucille Conklin for her assistance and Ralph H. Muller for his continuing interest. LITERATURE CITED
(1) Head, A. J., Kemper, iX.F., Xiller, R. P., Wells, R. A , J . Chenz. SOC. 1958,3408. ( 2 ) Knight, C. S., S u t u r e 183, 165 (1959). (3) Muller, R. H., “Progress in Suclear Energy Vol. 2,” C. E. Crouthamel, ed., p. 313, Pergamon Press, 1961. I n
press. (4) Fterson, H. T., ANAL. CHEM. 31, 1219 (1969). RECEIVED for reviex September 18, 1961. Accepted December 12, 1961. Work performed under the auspices of the U. S. -1tomic Energy Commission.
Spectrophotometric Investigation of Vanadium(II), Vanadium(lll), and Vanadium(IV) in Various Media ERNEST L. MARTIN and KENTON E. BENTLEYI Department o f Chemistry, University o f New Mexico, Albuquerque, N.
b
Methods for preparing vanadium
(II), (Ill), and (IV) in 0.5 and 5M sulfuric acid, 0.5 and 5M hydrochloric acid, and 0.5M perchloric acid solutions were studied. Absorption spectra were determined for vanadium(ll), (Ill), and (IV) in the various media. The absorption spectra of all vanadium (11) solutions were qualitatively alike. Spectral curves for vanadium(1V) in the different media varied slightly in the red region of the spectrum. Absorption of vanadium(ll1) in 0.5M sulfuric acid, 5M hydrochloric acid, and 5M sulfuric acid solutions exhibited an increasing bathochromic effect, in that order.
I
of reactions of the loner valences of vanadium with oxidants in sulfuric acid and hydrochloric acid media, in which the reactions were follorved spectrophotometrically, the need arose for methods of preparing solutions having a single valence state of vanadium in a n acid medium of a single anion, and absorption curves for these solutions. The absorption spectra of vanadium(II), N -4 SEPARATE STUDY
Present address, Department of Chemistry, American University of Beirut, Beirut, Lebanon.
354
ANALYTICAL CHEMISTRY
M.
(111), and (IV) solutions have been reported by several investigators (1-3, 6, 11), b u t the spectral data for the individual ions are frequently in rather poor agreement. Often a point of doubt is the purity of the valence state of vanadium in the solution measured. Many discrepancies are explained by the spectral studies of the loiver valences of vanadium by Furman and Garner (4) and King and Garner ( 7 ) . These workers reported complexing effects of media on the spectra of vanadium ions. This paper presents the results of a spectrophotonietric investigation of solutions of vanadium(II), (111), and (IV) in 0.5 and 5M sulfuric acid, 0.5 and 5 X hydrochloric acid, and 0.5X perchloric acid media. Sulfuric and hydrochloric acid media were considered because they are commonly used for solutions of vanadium. However, the anions of these acids tend to form complexes with many cations. Since perchlorate ion does not enter into complex formation, perchloric acid solutions of the vanadium ions were studied simultaneously to obtain comparative data for these ions in a noncomplexed form. The acid concentrations chosen, 0.5M and 5 M , differ significantly and make the solutions sufficiently acidic to keep hydrolysis of vanadium ions negligible (4). Thus, any spectrophotometric
or chemical anomalies affecting a single valence state of vanadium and traceable to the effect of the anion used, the concentration of anion, or the acid concentration were readily detectable and could be estimated quantitatively. Methods are described for preparing solutions of vanadium(II), (111), and (IV) in the various media. Absorption spectra are given for the several s o h tions. The data were obtained for use in developing spectrophotometric procedures involving the lower valence states of vanadium. EXPERIMENTAL
Reagents and Apparatus. The source of vanadium for this study was ammonium metavanadate (Electrochemical-Metallurgical Division, Union Carbide Corporation), assayed as 77.53y0 vanadium pentoxide, 0.3% carbon dioxide, 0.2% chloride ion, 0.057, sulfate ion, a n d 0.05% insoluble material. Cesium chloride of u n k n o m purity was recrystallized from water. Cesium sulfate was obtained by evaporating a sulfuric acid solution of t h e chloride to sulfur trioxide fumes twice. Compressed sulfur dioxide gas was used from the tank as supplied. C.P. hydrochloric and sulfuric acids were used. All other chemicals mere analytical reagent grade. Analytical solutions were prepared by the methods of Willard and Furman
( I S ) . il Jones reduct)or (15) and a silver reductor (12) v ere constructed, used, and rejuvenated following recommended procedures. Oxygen-free nitrogen was used to maintain an inert atmosphere chamber for work with oxygen-sensitive solutions. Solutions of reducing agents and solutions which were to be reacted with reductant solutions were made oxygen-free and stored under nitrogen. Absorbance Measurements. -4 Coleman Nodel 11 Universal spectrophotometer T\as used to measure absorption spectra over the visible wavelength rcgion of 400 to 775 mg. Solutions were placed in matched cells (1.6-cm. sample path length) and closed by rubber stoppers. Care was taken that the solution did not come in contact with the stopper. Solutions of vanacliuni(I1) and (111) n-eresealed in a nitrogen atmosphere. Concentrations of solutions were volumetrically adjusted, n-hen necessary, to bring absorbance readings M ithin the accurate range of the instrument. Spectral data are expressed in terms of molar absorptivity, E , defined by the relationship B = A/bc, where A is absorbance, b is sample path length in centimeters, and c is molar concentration of the vanadium ion. Determination of Valence States in Solutions. Vanadium m a y exist in solution in the +2, + 3 , $4, or + 5 valence state, or as a mixture of any t \ \ o contiguous valence states. The individual valence states and their concentrations were determined by titrating t n o aliquots of each solution; one aliquot was titrated to vanadium (IV) with standard ferric solution, the other to vanadium(V) n i t h standard permanganate solution (9). Accuracy of the results was checked by comparing the sum of the concentrations of individual valence states with the total vanadium concentration as determined on a third aliquot [permanganate ouidation-ferrous titration method (10) 1. End points were measured electrometrically using a vacuum-tube voltmeter and a platinum-calomel electrode system. Vanadium(I1) and (111) solutions were titrated under nitrogen. Procedure. Several methods were tried for preparing solutions of vanad i u m ( I I ) , (111),and (IV) in 0.5 and 5 M sulfuric acid, 0.5 and 5111 hydrochloric acid, and 0.5X perchloric acid. The object was to prepare each solution of a single valence state of vanadium in a n acid medium of a single anion. The only allowable impurities were those that would not interfere in the spectrophotometric or wet chemical determinations of valence states of vanadium--e.g., perchlorate ion was not considered to interfere with absorbance readings in sulfate or chloride media. Hydrogen ion and anion concentrations of the medium in final solution were not determined; however, adjustment of these concentrations was attempted to approximate concentrations of the medium under study. Each solution was volumetrically analyzed to determine concentrations of the vanadium ion prepared and any
contaminate vanadium ion present. An absorption spectrum n as obtained for each solution. Determinations were made on each vanadium(I1) and (111) solution immediately after preparation and again after absorbance measurements to learn whether the solution had been altered appreciably by oxidation since first analyzed; oxidation was negligible. Volumetric and spectrophotometric results were subjected to a comparative evaluation, and data were correlated to establish qualitative and quantitative specifications for the solutions in terms of absorption spectra. Spectral curves for solutions of the same valence state in the same media but prepared by different methods were compared. Solutions of the same valence state but in different media were compared for shape of curve as well as for characteristic wavelengths. The most satisfactory methods for preparing vanadium solutions of single valence states in the various acid media were selected on the criteria of purity of the particular valence state obtained and ease of manipulation of the procedure. The latter criterion was naturally more important for the oxygen-sensitive valence states. Preparation of Solutions of Vanadium(II), (111),and (IV). All solutions of vanadium(I1) and (111) were prepared in a nitrogen atmosphere using oxygen-free reagents because these ions are readily oxidized by air. Solutions of vanadium(I1) were obtained using a Jones reductor. Stock solutions approximately 0.05X in ammonium metavanadate in 0.5M sulfuric, hydrochloric, and perchloric acids were prepared. Each solution was passed through the reductor a t room temperature, a t a rate of 25 to 50 ml. per minute. Portions of the reduced sulfuric acid and hydrochloric acid solutions were made 511.1 in the acid medium by addition of the concentrated acid. Vanadium(II1) solutions were prepared from the sIightly soluble, cesium)~, (8). vanadium alum, C S V ( S O ~12H20 The alum was obtained by precipitation from a very dilute sulfuric acid solution of \-anadous and cesium sulfates. The precipitate was allowed to form slowly, over several hours, as vanadium (11) was oxidized by air to vanadium (111). The rate of precipitation was increased by partial oxidation of the vanadium(I1) by hydrogen peroxide or vanadium(V). The slower the precipitation, the more coarse and darker blue was the crystalline precipitate. The supernatant solution varied in color from almost colorless to shades of green or blue, depending upon the amount and valence of vanadium remaining in solution. conditions favorable to precipitation and quality of precipitate were low acidity, a large excess of cesium ion, low temperature, and slow precipitation. A high concentration of chloride inhibited pre-
cipitation. There was no precipitate with cesium(1)-vanadium(II1) from either hydrochloric or perchloric acid solution nithout sulfate prment. S o r was there a precipitate from solutions of vanadium(I1) or (IY) unless vanadium(II1) was present. Any adsorbed or coprecipitated vanadium(I1) or (11') or chloride was removed by subsequent washing and recrystallization. The precipitate n-as recrystallized from 0.1X sulfuric acid, then washed with 0.1M sulfuric acid until several consccutive washings vere the blue-green of r-anadium(II1) in dilute sulfuric acid solution. The dried, light blue alum was stored under nitrogen for 3 months. During this time the surface became green, but the body of the solid remained unchanged. Solutions of fresh alum and the 3-month-old solid in like media had the same spectral characteristics. Solution of the alum in 0.5 and 531 sulfuric, 0.5 and 5 N hydrochloric, and 0.5.U perchloric acids was carried out in an inert atmosphere with oxygenfree reagents. I n hydrochloric acid media, the sulfate of the alum wis precipitated by addition of a saturated barium chloride solution and was filtered a t approximately 80" C. using fine filter paper. A saturated solution of barium perchlorate was added t o the perchloric acid solution of the alum, precipitating barium sulfate and cesium perchlorate. The precipitates were removed by filtration at room temperature through medium-porosity, sinteredglass filtering crucibles. For greater solubility of the alum in sulfuric acid media, n-hen desired, cesium was precipitated as the perchlorate. Solutions of vanadium(1V) in sulfuric acid and perchloric acid media were obtained by sulfur dioxide reduction of vanadium(V) ( 5 ) . Sulfur dioxide gas \\-as passed through the solution for 30 minutes a t room temperature. Excess sulfur dioxide was expelled by boiling while a stream of nitrogen m s slvept through the solution; espulsion was considered complete when effluent gas ceased to reduce permanganate ion. Solutions 0.05X in ammonium metavanadate in 0.5 and 5M sulfuric acid and 0.5X perchloric acid media were reduced to vanadium(1V). A slight excess of a saturated solution of barium perchlorate n-as added to the vanadyl perchlorate solution to precipitate sulfate formed during the reduction. A vanadyl chloride solution was prepared by passing a solution 0.05M in ammonium metavanadate and 0.131 in hydrochloric acid through a silver reductor a t room temperature, a t a rate of approximately 25 ml. per minute. The reduced solution was divided into two portions. Sufficient concentrated hydrochloric acid was VOL. 34, NO. 3, MARCH 1962
355
Table
I.
Determination of Valence States in Solutions of Vanadium(ll), Vanadium (Ill), and Vanadium(1V)
Vanadium, Mg./25 M1. of Solution Valences Found Sum of Total (11) (111) (IV) Valences Determined 62.6 0 62.6 f 0.22 62.6 f 0.11 37 6 0.01 37.6 f 0.18 37.5 f 0.09 6 3 . 8 f 0.22 63.8 f 0.11 37.6 f 0.09 37.9 f 0.18 60.5 f 0.11 60.1 f 0.22 (111) 0 . 5 M HzSO4 58.8 0.51 59.3 f 0.49 59.0 f 0.11 5M 0.31 59.1 59.4 f 0.49 59.5 f 0.11 0 . BM HC1 42.2 0.20 42.4 f 0.52 42.4 f 0.09 j@-HCib:c94.8 15.0. 109.8 f 0.62 103.2 f 0.17 0.5MHClO4 0.46 41.5 42.0 f 0.52 42.3 f 0.09 63.6 63.6 f 0.11 63.4 f 0.11 (IV)’ 0.5M 61.9 61.9 f 0.11 62.0 f 0.11 5M HISO, 65.1 65.1 f 0.11 65.3 f 0.11 39.3 39.3 f 0.09 38.7 f 0.09 61.0 61.0 f 0.11 60.3 f 0.11 a This negative value is assumed to be result of titration error. A high concentration of chloride ion can effect incomplete oxidation of the lower states of vanadium to vanadium(V) by permanganate. Consequently, results of the permanganate titration would tend to be high, while those of ferrous titration would be low. These results would, in turn, lead t o low values for total vanadium and vanadium (11)and high values for vanadium(II1) and (IV). This solution was believed to be pure vanadium(II1) because the spectral data did not indicate contamination by vanadium(1V). The value of “Vanadium( 111) Found” is assumed t o be correct because it is the result of ferric titration only. d Analyses of vanadium(1V) solutions included only permanganate titration and determination of total vanadium. Solution Prepared Valence Medium (11) 0.5134 HzSOa 5144 HzSOi
~~
~
added to the separate portions to produce vanadium(1V) solutions in 0.5 and 5BI hydrochloric acid media. RESULTS AND DISCUSSION
Results of the determinations of valence states in solutions of vanadium (11), (111), and (IV), prepared by the methods described above, are given in Table I. The limits of values in Table I are the propagated errors for the analytical procedures and apparatus. These limits do not account for uncertainties in the methods of preparation. The accuracies of titrations in 5 X hydrochloric acid media are frequently outside the limits used in calculating propagated error. However, absorbance-wavelength curves for the
Table II.
5M hydrochloric acid solutions of vanadium(II), (111), and (IV) were analogous t o those in other media, allowing the explanation that permanganate titrations of vanadium in 5 M hydrochloric acid are not accurate ( 1 4 ) . The absorption spectra of the solutions listed in Table I are shown in Figures 1 and 2, and the characteristic wavelengths are given in Table 11. Values from the “Valences Found” column for the valence of the solution prepared, in Table I, were used to obtain molar concentrations of vanadium ions for calculation of molar absorptivities. Absorption spectra of all vanadium (11) solutions were qualitatively the same. The wavelengths of maximum absorption for vanadium(1V) in dif-
Qualitative Spectral Data for Solutions of Vanadium(ll), Vanadium(lll), and Vanadium(1V) Mean Wavelengths of Absorption, mp
Medium Vanadium( 11) Eacha Vanadium( 111)
Visual Color
0 . 5 ~ %HzSO4 f
5M HISOA
Violet Blue-green Green Blue
Green Blue
Minima
Maxima
465, 720
565
495 520 485 505 485
400, 415. 400; 400, 400,
605 645 590 620 590
Vanadium( IV) 445 750-770‘ Each0 Blue a 0.5 and 5M HISO?,0.5 and 5M HCl, and 0.5M HC104. Spectrophotometric data were not well defined a t wavelengths greater than 750 mp.
356
ANALYTICAL CHEMISTRY
WAVE
LENGTH,rnp
Figure 1. Absorption spectra of vanadium(ll), (Ill), and (IV) II. Vonodium(l1) in 0.5 and 5M HzSO4, 0.5 and 5M HCI, and 0.5M HClOa media 111. Vanodium(lll) in 0.5M HCI and 0.5M HClOd media IV. Vanadium(lV) in 0.5M H C I 0 4 media [Absorption spectra of all vanadium(lV) solutions were qualitatively alike, except for slight, indefinite variations from 750 to 770 m p ]
ferent media varied slightly, but indefinitely, from 750 to 770 mp. The spectrophotometer was not sensitive enough for better definition in this region of the spectrum. Figure 2 illustrates the bathochromic effect on vanadium(II1) absorption spectra of the different media. Such an effect is usually assumed to be the result of complex-ion formation. There F-as no shift of the spectrum of a solution in which the niolar chloride ion concentration was 15 times the molar vanadium(II1) concentration. There was about a 15-mp shift for a solution in which the sulfate ion concentration was approximately 10 times the vanadium(II1) concentration; a 30-mp shift when the chloride ion concentration was 70 times the vanadium(II1) concentration; and a 55-mp shift when the sulfate ion concentration was 100 times the vanadium(II1) concentration The so-called temperature broadening of the absorption band was observed on solutions of vanadium(II1) in 0.5114 hydrochloric and 0.5M perchloric acids. When heated to below boiling temperature, these solutions appeared green, and vihen cooled again to room temperature, reverted to their original blue. Furman and Garner (4)suggest that there is a larger degree of hydrolysis of vanadium(II1) a t higher temperatures. This increases the absorptivity, the effect being more pronounced in the region 400 to 500 mp than at 590 mp because of the larger absorption of hy-
d
c
b 0
WAVE L E N G l H , r n p
Figure 2. Absorption spectra of vanadium(ll1) in HzS04, HCI, and HClO4 media a. b. C.
d.
77-day period. Both precipitates appeared to be the cesium-vanadium alum. The precipitate formed more readily in the solution of lower acid concentration. However, vanadium remaining in the 0.5M sulfuric acid solution was largely oxidized to vanadium(JV), favoring dissolution of the precipitate 1% ith the continued oxidation of vanadium(II1) to (IT') as the precipitate redissolved. The rate and extent of this oxidation indicate that the cell was not sealed against the atmosphere. The vanadium remaining in the 551 sulfuric acid solution was mostly vanadium(II1). Changes in the vanadium(I1J) spectra in 0.5 and 5M sulfuric acid media are sh0R-n in Figures 3 and 4, respectively. I n preparing vanadium(I1J) in sulfuric
0.5M HCI and 0.5M HCIO, media 0.5M H&Oa medium 5M HCl medium 5M HzSO, medium
drated vanadium(II1) a t the shorter wavelengths. Stabilities of Vanadium(I1) and (111) Solutions. T o study the relative stabilities of vanadium(I1) and (111) solutions under the conditions of this investigation, rates of oxidation of these ions in the various acid media were compared. Oxidation of vanadium(I1) to (11, 111) and vanadium (111) to (111, IV) was followed spectrophotometrically. T'anadium(1J) solutions n ere prepared by the Jones reductor method, and vanadium( JJJ) solutions were prepared by the cesiuai-vanadium-alum method. Vanadium concentrations n ere approximately 0.03V to 0.OLlP. These solutions were sealed in Coleman cells before being removed from the inert atmosphere chamber. Absorbancewavelength curves for the vanadium( 11) solutions were determined on the day prepared, 7 days later, and 28 days later; for vanadiuni(TI1) solutions, curves were determined on the day prepared and a t 14 and 77 days after preparation. The stoppered cells containing the vanadiurn(I1) and (111) solutions were not returned to the inert atmosphere chamber during these intervals but were kept in the laboratory, a t room temperature, exposed t o fluorescent light and sunlight. The only phenomenon observed, other than color changes, as the solutions stood was the formation of precipitates in the 0.5 and 5 X sulfuric acid vanadium(II1) solutions. I n the 0.534 solution, a fine, crystalline precipitate was present a t the 14-day period, but redissolved by the time the 77-day period had elapsed. I n 5iM sulfuric acid, a small, single crystal, which appeared a t the 14-day period, grew to a large single crystal by the end of the
f--
4
m
Table 111. Relative Stabilities of Vanadium(l1) and Vanadium(lll) in Various Media Oxidation Oxidation
Rate,
Rate,
V(II),
Medium 0 . 5 M His04 5144 HiSOi 0.5'7 HCi 5BF HC1 0 55fHC10,
k(lI),
Dags-'
V(III), k(III),
Days-'
0.036
0.13 0.023 0.002 0 049
0.12 0 013 0 031 0.021 0.059
where t is time elapsed after preparation of the solution, in days; A ( I I ) ~ is absorbance of vandium(I1) solution, read less than 1 day after preparation; A ( I I , I I I ) ~ is absorbance of vanadium (11,111)in the solution, read t days after preparation; A(II,t is absorption due to vanadium(I1) in the solution t days after preparation; ~ ( 1 1 1 and E ( I I I ) are molar absorptivities of vanadium(I1) and (111), respectively. Similarly, for vanadium(II1) in Table 111
I
cc 0
0.60
-a
0.K
0 2:
400
500 NAVE
6CC
700
830
L E U G T ? , rnf
Figure 3. Stability study of vanadium (111); absorption spectra of vanadium (111) and (111, IV) in 0.5M &Sod.
where the terms are analogous. Results for vanadium(I1) were calculated from measurements taken shortly after preparation of vanadium(I1) solutions and 7 days later; and for vanadium(III), shortly after preparation and 14 days later. Absorbance and molar absorptivity values were taken a t 400 mp. Spectral data a t this wavelength were related to the concentrations of va-
a. Spectrum determined less than 1 d a y after preparation of the vanadium(lll) solution b. Spectrum determined 14 days after preparation of the solution e. Spectrum determined 77 days after preparation of the solution
acid media by solution of cesiumvanadium-alum in later work, precipitation of the alum was circumvented by removing cesium as the slightly soluble perchlorate. Concentrations of vanadium ions in each solution were related using absorbance measurements and previously determined molar absorptivities. Oxidation rates of vanadium(I1) in the various media are compared in Table 111. Values for this expression were calculated using the rate relationship
690
500 WALE
600 700 LE\GTH,mp
800
Figure 4. Stability study of vanadium (111); 'absorption spectra of vanadium (111) and (111, IV) in 5M HzS04 Spectrum determined less than 1 d a y after preparation of the vanodium(lil) solution b. Spectrum determined 14 days after preparation of the solution e. Spectrum determined 77 days after preparation of the solution a.
VOL. 34, NO. 3, MARCH 1962
357
nadium ions in solution with sufficient accuracy for these studies. In general, highly acid media or media having a bathochromic effect on the absorption spectrum of vanadium(II1) have an adverse effect on the stability of vanadium(I1) and enhance the stability of vanadium(II1). Increased acidity and complexing of vanadium(IT1) favor the rate of the forward reaction of Equations 3 and 4, but not of Equation 5. 1 + H+ + -HP (31 2 1 2V+z + 70, + 2H+ 2V+3 + H?O ( 4 ) 1 2VA3+ 20’ + H?O = 2VO+* + ~ H T
V+*
V+3
=
perchlorate ion is reduced by vanadium (II), forming considerable amounts of yanadium(II1) and chloride. The results for vanadium(II1) correspond to the general order with the exception of 0.5AlIsulfuric acid. The high value of I Z ( I ~ I ) for the latter medium is not valid for tn-o reasons: first, some of the vansdium(II1) came out of solution in a precipitate; and, second, there was excmsive oLidation of the vanadium in solution to vanadyl ion. The oxidation probably occurred because the cell was not completely sealed. Precipitation of the alum in the 5-11 sulfuric acid solution of vanadium(II1) was slight at the l4-day period and therefore had little effrct on the value of k(111) for the incdium.
(5)
From this generalization, the stability of vanadium(I1) should decrease, and the stability of vanadium(II1) should increase in the various media, in the following order: 0.5X perchloric acid, 0.5M hydrochloric acid, 0.5M sulfuric acid, 5J1 hydrochloric acid, 5.11 sulfuric acid. The results for vanadium(I1) follow this order with the one exception, 0.5M perchloric acid. This exception is readily explained by the fact that
ACKNOWLEDGMENT
The authors thank H. H. Willard, University of Michigan, for his assistance and helpful counsel. This work was performed under contract SC-5 with the Vniversity of California. LITERATURE CITED
(1) Canning, R. G., Dixon, P.,
h A L .
CHEII.27,877 (1955).
(2) Dainton, F. S., J. Chem. SOC.1952, 1536. (3) Dreisch, T., Kallscheuer, O., 2. physik. Chem. B45, 19 (1939). ( 4 ) Furman, S. C., Garner, C. S., J . Am. Chem. Soc. 72, 1785 (1950). (5) Hillebrand, W. F., Lunde!l, G. E. F., “Applied Inorganic Analysis,” p. 359, Wiley, New York, 1929. (6) Kato, S., Sci. Papers Inst. Phys. Chem. Research ( T o k y o ) 12, 230 (1930); Ibid.. 13.49 11930). (7) King, iV. R., Garner, C. S., J . Phys. Chem. 58, 29 (1954). (8) Latimer, \T. M., Hildebrand, J. H., ”Reference Book of Inorganic Chemistrv.” 3rd ed.. D. 362. Macmillan, New York, 1957. (9) Maass, K., 2. anal. Chem. 97, 241 ,
A
(1934).
(10) Rulfs, C. L., Logomki, J. J., Bahor, R. E., Ax.4~.CHEW28, 84 (1956). (11) Sabatini, R., Hazel, J., hIcNabb, W., Anal. Chim. Acta 6,368 (1952). (12) Kalden, G. H., Hammett, L. P., Edmonds, S. M., J . .4m. Chens. SOC. 56,350 (1934). (13) Killard, H. H., Furman, N. H., Bricker, C. E., “Elements of Quantitative dnalvsis.” 4th ed.. D. 218. Van Nostrand,*Ne&York, 1956.(14) Ibid., p. 216. (15) Ibid., p. 222. RECEIVEDfor review April 20, 1960. Resubmitted December 21, 1960. 4ccepted December 21, 1960. Division of Analytical ChemiFitry, 136th Meeting, ACS, Atlantic City, Tu’. J., September 1959.
Iron(II) as an Indicator Ion for Amperometric Titrations with (Et hylenedinitrilo) te traace tic Acid Application to the Determination of Thorium GERALD GOLDSTEIN, D. L. MANNING, and
H. E.
ZITTEL
Analytical Chemistry Division, Oak Ridge National laboratory, Oak Ridge, Tenn.
b A method for amperometric titrations with EDTA in which Fe+2 is used as an indicator is described. A potential of +0.4 volt vs. the S.C.E. is applied to a platinum electrode. At this potential Fe+2 is not oxidized; however, which is formed after the titration is complete, is oxidized, and the excess titrant is indicated b y a linear increase in current. Most titrations are carried out in an acetate buffered medium a t p H 4.5. This method is applicable to the determination of elements that form EDTA chelates with a stability constant of 10” or greater with several exceptions. Thorium(lV) in the presence of U+6, Zr, and other cations was determined with a relative standard deviation of about 270. Sulfate and small amounts of fluoride do not interfere. 358
ANALYTICAL CHEMISTRY
E
THPLEKEDINITRILO) T E T R A A C E T I C ACID (EDTA) is a n extremely use-
ful titrant for the determination of many elements. Most often, a visual indicator is used to detect the equivalence point. However, a visual indicator cannot be used if test solutions are highly colored or contain substances that interfere in the indicator action. I n the case of the remotely controlled analysis of radioactive solutions, visibility through the shielding may be poor, and the indicator may be damaged by radiation. Under such conditions, other methods for establishing the equivalence point are required. Several electrometric methods are available. Potentiometric end point detection using the Hg/HgY+ indicator electrode (4) has proved to be extremely versatile, particularly for selective titrations in multicomponent solu-
tions. However, the potential break is often small when solutions which are less than 10-3;M are titrated and halide ions can interfere. Amperometric titrations with EDTA have been applied to the determination of elements which are electroactive at a dropping mercury electrode ( 5 ) . Electroinactive cations can also be determined if a suitable electroactive indicator ion is used (3, 6). It was desirable, for this work, to use a platinum electrode rather than a dropping mercury electrode, so that the titrant could be delivered continuously and the full titration curve recorded. Ferrous ion was chosen for study as a possible indicator ion because i t is electroactive at a platinum electrode (2), and its oxidation potential is shifted in the presence of EDTA (8).
