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., CHEII.27,877 (1955).
h A L .
(2) Dainton, F. S.,
1536. (3) Dreisch,
J. Chem. SOC.1952,
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 (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. ,
A
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).
EXPERIMENTAL
Reagents. (Ethylenedinitrilo) tetraacetic Acid, 0.02-If. Dissolve 7.4 grams of disodium (ethylenedinitri1o)tetraacetate dihydrate in water and dilute the resulting solution to 1 liter. Standardize this solution by any of the standard procedures ( 7 ) or by titrating known quantities of thorium by the procedure recommended heriin. Ferrous Ammonium Sulfate Solution, 0 0231. Dissolve 0.8 gram of Fe(KHJ2(SO4)*6H20 in 100 nil. of deaerated 0.01JI HClO,. Buffer solution, pH 2.5. Dissolve 28 grams of chloroacetic acid in water, add 150 ml. of 1-V sodium hydroxide, and dilute the resulting solution to 500 nil. with water. Sodium Acetate Solution, 3 X . Apparatus. Polarograph, recording, ORNL Model &-1160 ( 1 ) or equivalent. Electrode, platinum foil, 1-sq. em. area. Saturated Calomel Electrode, (S.C.E.) together with agar-KC1 salt bridge. Syringe Buret Titrant-Delivery Apparatus, consisting of a 10-ml. precisionbore syringe and a Harvard single speed infusion-n ithdran-a1 pump, was equipped with a 2-r.p.m. synchronous motor. This device delivered about 0.3 ml. per minute The rate of titrant delivery can be varied by the use of syringes of different sizes. A No. 9 rubber stopper, which fits a 50-ml. beaker, was bored to allow insertion of: the platinum electrode, the salt bridge from the S.C.E., a nitrogen-bubbling tube, and the buret tip. An O R S L Model (2-1160 recording polarograph waq used to apply the potential t o the Plectrodes and to record the current. Sufficient damping was applied to remove curlent oscillations. Current sensitii ity Tvas 20 pa. full scale. The solutions were stirred by a magnetic stirrer and Teflon-coated stirring bare. In some cases, titrations were performed manually by replacing the syringe buret with B conventional buret and the polarograph with a Sargent dmpot. RECOMMENDED PROCEDURE
Transfer from the sample solution to a 50-ml. beaker a test portion that contains about 0.025 to 0.05 mmole of the ion to be titrated. For titrations a t p H 4.5, add 2 ml. of 311 sodium acetate solution and sufficient perchloric acid to adjust the p H to 4.5; for titrations a t pH 2.5, use 2 ml. of p H 2.5 buffer and adjust the pH to 2.5. Dilute the resulting solution to about 25 ml. with water. Deaerate the solution with nitrogen. Add 0.2 ml. of 0.0211f ferrous ammonium sulfate solution. Apply a potential of f 0 . 4 volt us. the S.C.E. t o the platinum electrode. Titrate with the standard EDTA solution. Record the current and the corresponding volume of titrant added during the titration. A rapid increase in current indicates that excess titrant has been
-3?;.C +0:8 +0.6 +0:4 +0!2 0 APPLIED P O T E N T I A L v s T H E S.C.E., v o l t s
Figure 1. Current-voltage curves for Fe+2and Fey-?
added. Locate the equivalence point of the titration by extrapolating the two segments of the titration graph to their point of intersection. In addition, titrate in the same way a blank that contains all the reagents. Subtract the titer for the blank from that for the test portion of the sample.
oxygen, must be excluded from the test solution before the Fe+? is added. By titrating a reagent blank a t the same time the test portion of the sample is titrated, corrections can be made for any Fe+3 in the indicator solution. An Fe L2 solution prepared by dissolving ferrous ammonium sulfate in a deaerated acid solution is stable for seJeral weeks and has a very small blank titer, usually less than 0.001 mmole of EDTA per milliliter of Fe+2 solution. Other strong oxidizing agents such as dichromate can be reduced with hydroxylamine hydrochloride before the Fe+2 is added. Hydroxylamine hydrochloride is a sufficiently strong reducing agent to reduce most oxidants and is not itself oxidized a t the platinum electrode a t t 0 . 4 volt. Hydrogen peroxide, which will effectively reduce dichromate, acts as an oxidizing agent tom-ard FeY -2. Hydrazine hydrochloride is also an effective reductant; however, it is oxidized a t the platinum electrode giving a very high background current. OPTIMVM COSCESTRrlTION O F IRON
(11). Several factors must be considered in choosing the concentration of Fe+* indicator. If too much Fe+2 is present, it may compete for the EDTA with the ion being titrated, with a resultant increase in current before the equivalence point is reached. Too little FeT2 on the other hand results RESULTS AND DISCUSSION in only a small increase in current. The choice of the concentration of Iron(I1) a s Indicator Ion. CURR E N T - T.-OLTAGE CHARACTERISTICS. Fe +*, therefore, depends on the concentration of the ion being titrated, The current-voltage curves for the the stability of its EDTA complex, oxidation of Fe+2 and of the Fe+2and the sensitivity of the currentEDTA complex ( F e y w 2 )a t a platinum measuring device. As a general rule, electrode a t each of two p H values are in the titration of ions that form very shown in Figure 1. Complexation stable EDTA chelates, the moles of of Fe+2 with EDTA results in a change Fe+2 present a t the beginning of the in the standard reduction potential titration should be about lo%, of the from +0.77 to +0.12 volt (8).When moles of the ion being titrated. a potential of f 0 . 4 volt is applied to -4PPLICATION TO TITRATIOK OF VARthe platinum indicator electrode, the IOUS CATIONS. The results of the current is small because Fe+2 is not titrations of various cations in soluoxidized appreciably a t this potential. tions of pH 2.5 and 4.5 are presented However, when EDTA is added, Fey-2 in Table I. Representative titration is formed and is oxidized with a recurves are shown in Figure 2. Most sulting increase in current. The curions that form EDTA chelates having rent is a linear function of the amount a stability constant ( K ) greater than of EDTA added until all the Fe+2 has 1OI8 can be titrated a t pH 4.5; the been titrated. An applied potential end points of the titrations are sharply of f 0 . 4 volt rather than a more positive defined. These ions include Th+4, potential was chosen to keep the Bi+3, Fe+3, S C + ~ Y+3, , Luf3 YbL3, residual current small before Fey-2 Tm+3, Er+3, HO-~,DyT3, Td+3, and formed and to avoid possible interferPb+2. I n addition, some ions that ences by oxidizable substances. form somewhat less stable EDTA EFFECTOF PH. The Fe+2 indicator chelates (log K = 16 to 18), such as can be used successfully over the p H Gd+3, Cd+2, Co+2, and Zn+*, can also range from about 2 to 5. At pH 1.5, be titrated to reasonably well defined does not form. At p H greater end points. When attempts were than 5, hydrolysis of many ions, inmade to titrate ions for which log K 6 cluding Th+4, is likely to occur; therefore, higher pH values were not tested. 16, such as Li+, Mgf2, Mn+2, U02+2, EFFECTOF OXIDIZING AGENTS. OxiLa+3, Cef3, and Al+3, the added dizing agents, including atmospheric EDTA reacted immediately with the VOL. 34, NO. 3, MARCH 1962
359
and Sm +3 are apparently not quantitative. The titration curves of Cuf2 (Figure 3) are unusual. At p H 4.5 the titration curve has a peculiar shape, the apparent equivalence point occuring a t too high a value, whereas a t pH 2.5 a two-step titration curve is obtained, the equivalence point occuring at the second current rise. Of the other ions, at p H 2.5, only T h f 4 in addition to C u A 2could be titrated successfully. Zirconiurn(1V) formed a precipitate when EDTA was added to solutions a t pH 2.5. I n titrations of solutions containing two or more titratable cations, only k0.34 the sum total can be determined. It EDTA A D D E D , T I may be possible under other conditions Figure 2. Titration curves of various to titrate successfully, by this general elements in solutions at pH 4.5 method, ions that cannot be titrated under the conditions described herein. For example, the acetate buffer forms stable complexes with many ions and, Fe+2, and indicator current \vas obtherefore, competes with the EDTA. served. Another type of buffer solution would Some ions that form stable EDTA be preferable for the titration of these chelates, such as ZrA4,Cr+3, and Nif2, ions. could not be titrated. Titration curves PRECISION.The precision of the of ?;d +3, Pr +3, and Hg + 2 are so rounded amperometric titration mas evaluated that it is not possible to determine the equivalence points. Recoveries of E u + ~ by titrating replicate test portions of standard solutions of thorium. The data are presented in Table 11. I n general, the relative standard deviation of the results of both manual and inTable I. Titrations of Various Cations strumental titrations is about 1% at Cation. M n . pH either 2.5 or 4.5. Recovered Determination of Thorium(1V). At, At. .-_ EFFECTS OF ANIONS. T h e effect of Log PH each of various anions on the determiElement K (9) Added 2.3 nation of T h t 4 was evaluated by addBit3 26.5 6.40 6.20 a ing the anion, as its sodium salt, to Fe+3 25.1 1.50 1.55 d Zr+4 25.0 2.74 solutions of Th+* as the perchlorate Th+4 23.2 5.20 5.18 5.20 and then titrating the Th-4 a t p H 4.5 by Sc+3 23.1 1.00 1.01 .., the recommended procedure. The data Crt3 23.1 1.30 obtained in these titrations are preHg+z 21.8 6.00 c Luf3 19.8 6.74 6.68 sented in Table 111. Perchlorate, Ybf3 19.5 3.48 3.58 ... chloride, nitrate, and sulfate in a 10:l Ti+4 19.4 5.00 ... mole ratio to T h + 4 can be tolerated. Tm+3 19.3 4.26 4.56 ... On addition of sodium fluoride to the Er+3 18.9 3.G7 3.70 ... Cu+2 18.8 1.97 2.00 T h + 4 solution, a precipitate formed Ho+3 18.7 4.97 4.95 ... that dissolved slowly as the titration &-if2 18.6 1.76 c d proceeded. If the mole ratio of fluoDyf3 1 8 . 3 7.42 7.39 ... ride to Th+4 exceeded 1:1, the Y+3 18.1 4.00 3.94 ... Pb+2 18.0 6.21 6.11 ... precipitate did not dissolve completely Tb+3 17.9 4.00 3.93 ... in the time normally required for the Gd+3 17.4 3.75 3.61 ... titration, and the recovery of Eu+3 17.3 4.21 3.56 ,.. was low. If phosphate was added to Smt3 17.1 4.00 3.32 ... Nd+3 16.6 4.00 d ,.. the Th+4 solution, the thorium phosZnt2 16.5 1.80 1.80 ... phate precipitate that formed did not Cdf2 16.5 2.81 2 65 ... dissolve when EDTA was added. The Prf3 16.4 4.00 ... addition of molybdate caused the preCof2 16.3 1.77 1.70 ... cipitation of thorium molybdate, which AIf3 16.1 0.67 C ... Cef3 16.0 4.20 c ... did not dissolve in the presence of La+3 15.5 3.00 ... EDTA. When 5 ml. of glacial acetic R1ntZ 14.0 1.37 ... acid was added to the thorium molybh4gt2 8 . 7 0.61 ... date solution, the precipitate did disLi+ 2.8 1.00 c ... UO?+' , , . 5.48 e ... solve during the titration; however, after the equivalence point was reached, a Cation hydrolyzed. b High current throughout titration; the current slowly decreased instead no apparent end point. of rapidly increasing. Apparently, Immediate indicator current. molybdate (EO = +0.4 volt) is not an d End point poorly defined. oxidant for Fe+2 ( E o = +0.77 volt),
/ I,
I
.
Table II. Precision of Replicate Titrations of Th+* pH, 4.5 SumRelative ber Standard of DeviaTh +4, Repli- -4verage tion, Mg. cates Titer % Manual Titration Milliliters 2 6.50 4 1.49 0.5 13.00 4 2.97 26.00 4 5.85 1 5 26.0Oa 5.92 1
5.20
Instrumental Titration Millimole 9 0.0223
1
pH 2.5.
Figure 3. Titration curves for Cu+* (1.97
.
me.)
(I
0
0
360
0
ANALYTICAL CHEMISTRY
Table 111.
Effects of Anions on Determination of Thorium Anion Th +4 Iden- Added, ddded, Recovered tity mmolea mmole Mmole % ClO41.0 0.102 0.102 100 0.103 101 c11.o 1.0 0.102 100 NO,1.0 0.103 101 s04'F-
0.05
0.10 0.20
Pod3-
0.10
Cr2072-
0.008 0.0224
0.008b
h10042- 0.03
0.03~ 0 .03btC
0
b c
No end point.
0.102
0.101 0.089 0.029 a
0.0231 0.007
100
99 87 28
io3 31
0
0.0223
io0
Hydroxylamine hydrochloride added. 5 ml. of glacial acetic acid added.
SULFATE
2
PERCHLORATE
E
CHLORIDE
I I 008
0
002 004 006 Z I R C O N U V ADDED, m II moles
I 0+0
Figure 4. Effect of zirconium on determination of thorium (0.0224 mmole) in presence of various anions
but i t is an oxidant for (Eo= +0.12 volt) and oxidizes the before the reaches the electrode. When both hydroxylamine hydrochloride and acetic acid were added, a n indicator current was observed after the equivalence point, and the recovery of T h f 4 1%-asquantitative. EFFECTS OF EXTRANEOUS CATIONS. The cations that interfere in the determination of T h f 4 because they are also titrated are indicated in Table I. Other cations encountered that do not interfere at a 1:1 mole ratio level include Mn+*, Ce+3, .41+3, and Cr+3. Uranium(V1) present in a mole ratio as high as 5 : l can be tolerated. However, the slope of the titration curve beyond the equivalence point decreases as the concentration of U+6increases. Zirconium, under some conditions, interferes seriously in the titration of
Table IV. Results of Determination of Thorium in Thorium Oxide-Uranium Oxide Slurries
Sample 1
2
4 5
6 7 8 9 10
11
Th+4, Mg. per Gram of Slurry hmperometric Visual 519 513 226 229 217 223 281 284 265 2 75 143 139 27.4 131 135 45.8 81.3 79.3 13.2
514 509 23 1 238 222 220 281 280 277 269 148 149 27.0 129 134 48.7 73.8 74.4 12.3
Th+4 (Figure 4). Zirconyl chloride, perchlorate, and sulfate caused low recovery of Th+4. However, when airconyl acetate was added, the recovery of Th+4 was quantitative. The interference of zirconium can therefore be eliminated by adding sufficient acetic acid. The data in Figure 5 mere obtained by adding glacial acetic acid to solutions that contained 0.0224 mmole of Th+4 and various amounts of zircony1 perchlorate. When more than 10 ml. of glacial acetic acid was present in the final 25-ml. volume, the background current became erratic. This effect limits the tolerable mole ratio of zirconium to thorium to about 5:l. Studies of the effect of the order of addition of the reagents showed that the acetic acid must be added before the p H is adjusted. APPLICATIONTO SAMPLES. Samples of thorium oxide-uranium oxide slurries were dissolved, and the thorium n-as titrated at p H 4.5 as described in the recommended procedure. The results are presented in Table IV and are compared with EDTA titrations using xylenol orange as the indicator. These samples also contained corrosion products of stainless steel as minor components. Where two results are reported for the same sample, two separate portions of the original slurry were analyzed. Statistical analysis of the data given in Table IV indicates that there is no significant bias between the tn-o methods. A procedure was also developed for the determination of thorium in molten salt reactor fuels. These fuels are composed of LiF, BeF2, ZrF4, ThF4, and UF4 and are about 10% Li, 5% Be, 6% T h , 10% Zr, and 6% U by weight, Because these fuels are hydroscopic, they were handled in a dry box. The procedure for the analysis of the salt is as follows: SAMPLEDISSOLUTIOX. Place a 1gram sample in a 50-ml. beaker. Add 10 ml. of 1: 1 HCl, 10 ml. of concentrated HC104, and 1 gram of boric acid. Evaporate the solution to fumes of HC104 with moderate heating. Fume the solution until all the sample has dissolved and only 5 ml. of solution remains. Transfer the HClO, solution t o a 25-m1. volumetric flask; then dilute i t to volume with water. DETERhIINATION OF THORIUM. Transfer a 2-ml. aliquot of the solution to a 50-ml. beaker. Add 8 ml. of glacial acetic acid and about 50 mg. of hydroxylamine hydrochloride. Then add 13 nil. of water and 2 ml. of 31M sodium acetate solution. Follow the recommended procedure for the titration. I n Table V are given the results of analyses by this method in which no separations were made, and no attempt was made to eliminate the last traces of fluoride from the sample and by another method whereby thorium was
1
23-
Z
1 2
4 s ~ f i C 1 - L ACE’IC
6
{0
8
A C D , ml
Figure 5. Effect of acetic acid on recovery of thorium (0.0224 mrnole) in presence of zirconium
separated from uranium, zirconium, and fluoride prior to its determination by titration with EDTA using xylenol orange as the indicator. The relative standard deviation of the amperometric titrations, based on duplicates, was 2%. Although the agreement between the methods is not outstandingly good, the average difference (5%) is reasonable considering the difficulties involved in analyzing this type of sample. Table V. Results of Determination of Thorium in Molten Salt Reactor Fuels
Sample
Thorium, 7 0 SeparationAmpero- xylenol metric= orange
6.08 5,82 6.67 6.39 3 5.47 5.19 4 5.69 6.16 5 6 04 5.82 a Average of duplicates. 1 2
Difference, -4 -4 +j
-8 +4
LITERATURE CITED
( I ) Kelley, >I. T., Miller, H. H., Ax.ra~* CHEM.24, 1895 (1952). (2) Kolthoff, I. M., Sightingale, E. R., Jr.. Anal. Chim. Acta 17. 329 (1954). (3) Laitinep, H. A., Sympson,‘ R. F., ANAL.CHEW26, 556 (1954). (4) Reilley, C. N., Schmid, R. IT., Lan-son, D. W.,Ibid., 30, 953 (1958). (5) . , Reillev. C. N.. Scribner. W.G.. Temole. C.. ?bid.. 28. 450 (19g6). ( 6 j Ringbom, ’ -4.;FTilkman, B., dcta Chem. Scand. 3, 22 (1949). (7) Welcher, F. J., “The Analytical Uses of Ethylenediaminetetraacetic -4cid,” p. 14, Van Xostrand, Princeton, N. J., 1958. ( 8 ) Ibid., p. 25. (9) Ibid., p. 7 ; “The EDTA Titration,” p. 30, J. T. Baker Chemical Co., November 1957; Booman, G. L., Phillips Petroleum Co., Idaho Falls, Idaho,
private communication. RECEIVEDfor review October 9, 1961. Accepted January 12, 1962. Combined Meeting of Southeast and Southwest Sections of ACS, New Orleans, La., December 7-9. 1961. VOL. 34, NO. 3, MARCH 1962
361
