Table 11. Water Determinations id Dowex AG 50-W - H + Form Resins by Karl Fischer, NMR, and Drying Methods
Dowex AG 50-W Hf form
K. F.
x2 x4 X8 XI2 X16
74.5 65.3 51.8 45.6 38.4
a
Water (w/w) NMRa Dryingo 74.0 64.6 54.0 44.2 36.0
78.3 66.5 53.0
45.6 37.0
From Ref. (5).
Li+ > Na+ > K+ > Rb+ > Cs+ as expected (7). Qualitative confirmation for this trend is obtained by the integration of in-water peaks in the NMR spectra of the resins in the above ionic forms. Quantitative results could not be obtained since in-water peak overlaps out-water peak (6, 8). In the case of Ca2+and Laa+forms of Dowex AG 50-W (X8) stable end points could be reached in 10-15 minutes provided the resins were taken in “wet” form. For air-dried resins, it took hours before the titrations could be completed. For all ionic forms, the Karl Fischer method is rapid and accurate if the resin is taken in the completely swollen (wet) form, Recently Blasius and Schmitt (9) have determined up to 0.5 (wiw) of water accurately by contacting the resin for 24 hours in excess Karl Fischer reagent and back-titrating the unreacted reagent. In conclusion, direct titration with Karl Fischer reagent can be relied upon to give rapid and accurate estimation of water content for the sulfonated cation exchange resins provided that water content in the resin is not very low ( 210 w/w) and errors up to =t2xin the estimation are tolerable. Where more accurate results are desired, the procedure of Blasius and Schmitt can be adopted using a more elaborate Karl Fischer apparatus.
case did the time required to attain a stable end point exceed 35 minutes. De Villiers and Parrish (5) estimated the water contents of Dowex 50-W resins in the H+ form (X2 to X16) by drying and also by the NMR method. Their values were confirmed in our laboratory by NMR chemical shift measurements and are compared with those obtained by the Karl Fischer method in Table 11. The moisture contents of Dowex AG 50-W (Xl-16) resins in the H+, Li+, Na+, K+, Rb+, and Cs+ forms were determined by us using the titration method. The results show a progressive decrease in the water content in the order H+ >
RECEIVED for review June 16, 1969. Accepted September 25, 1969. The authors gratefully acknowledge the financial support for this work by the National Research Council of Canada.
(5) . _ J. P. de Villiers and J. R. Parrish, J. Polymer Sei., Part A, 2, 1331 (1964). (6) . , H. D. Sharma and N. Subramanian, DeDartment of Chemistry. University of Waterloo, Waterloo, Ontarib, Canada, unpublished data, 1969.
(7) F. Helfferich, “Ion Exchange,” McGraw-Hill, New York, p 103, 1962. (8) J. E. Gordon, J . Phys. Chem., 66, 1150 (1962); Chem. Znd. (London), 1962,267. (9) E. Blasius and R. Schmitt, 2.Anal. Chem., 241, 4 (1968).
(a,
x
Gamma-Ray Titrations Using Potentiometric End Point Detection R. F. Sympson and W. D. Hunter Clippinger Laboratories, Ohio University, Athens, Ohio 45701
CHEMICAL DOSIMETERS in which radiation dosage is measured by the quantity of product formed by a radiation-induced chemical reaction have been in use for many years ( I , 2). The reverse process in which the concentration of a reagent is determined by measuring the radiation dosage required to cause it to react completely has found limited use in analytical chemistry. An aqueous system exposed to ionizing radiation develops a potential that has been called the “equivalent redox potential” or “equivalent reduction potential” and is characteristic of the radiation used and the solvent (pN, oxygen content, etc.) (3,4). The value of this potential was found to be 0.85 to 0.90 volt more positive than the normal hydrogen electrode. Cartledge (5)interpreted this potential to be a mixed potential dependent upon the steady-state concentrations and electrochemical characteristics of all radiolytic products present in (1) H.Fricke and S . Morse, Am. J. Roentgenol., 18,430 (1927). ( 2 ) J. W. T. Spinks and R. J. Woods, “Introduction to Radiation
Chemistry,” Wiley, New York, 1964, Chap. 4.
(3) F. S. Dainton and E. Collinson, Ann. Rev. Phys. Chem., 2, 99 (1951); Discussions Faraday SOC.,12,251 (1952). (4) H. S. Henderson, E. G. Lovering, R. L. Waines, and E. J. Casey, Can. J . Chem., 37, 164 (1959). ( 5 ) G. H. Cartledge, Nature, 186, 370 (1960). 2064
*
the solution. The steady-state concentrations depend upon the rates of formation of the particles and their rates of decay by various paths. If a redox couple is added to the solution, reaction between the added couple and radiolytic products occurs until the activities of the components of the oxidationreduction couple are such that they no longer disturb the steady-state activities of the radiolytic products. If the E a of the added oxidation-reduction couple is more than a few tenths volt from the equivalent redox potential, one component of the couple is converted quantitatively to the other form. Any couple having a standard potential more than 0.2 volt positive to the equivalent redox potential is quantitatively converted to the reduced form, and any couple with a standard potential more than 0.2 volt negative to the equivalent redox potential is quantitatively converted to the oxidized form. It would seem that if a solution containing an oxidationreduction couple were irradiated by gamma-rays or x-rays of constant intensity, the time required to establish the equivalent redox potential would be proportional to the initial concentration of one of the components of the couple. This is not ,generally true, however, because in many radiationinduced reactions the yield of product varies with concentration of the reactant. If the yield is not constant over a con-
ANALYTICAL CHEMISTRY, VOL. 41, NO. 14, DECEMBER 1969
siderable concentration range, linear calibration curves of concentration us. dosage cannot be obtained. If the yield were to change measurably before quantitative reaction of the unknown, it would be analogous to a conventional volumetric titration in which the stoichiometry of the reaction changed during the course of the titration. A limited amount of work using X-radiation or gammaradiation for quantitative determinations has been reported. Dewar and Hentz (6) oxidized iron(I1) by irradiation with X-rays and followed the course of the reaction potentiometrically. The precision of their results was limited by the experimental features of their work. The potential of a Pt electrode was measured periodically with a Leeds & Northrup Model K potentiometer instead of recording the potential continuously. X-radiation came from an industrial X-ray unit which required continual manual adjustment of the tube current to maintain constant dose rate. Musha, Sugimoto, and Tsujino (7) determined iron(I1) by oxidation with cobalt-60 gamma-rays. They followed the course of the oxidation amperometrically using two polarized Pt electrodes. Sugimoto and Musha (8) titrated sodium carbonate with hydrochloric acid generated from the radiolysis of chloral hydrate. The end point was determined by conductometric measurements. In the present work four radiation-induced reactions were followed potentiometrically : oxidation of iron(I1) to iron (111), oxidation of uranium(1V) to uranium(VI), reduction of cerium(1V) to cerium(III), and reduction of dichromate to chromium(II1). When acidic aqueous solutions are irradiated, four primary radiolytic products are formed-H., .OH, HS,and Hz02. Determinations were based on quantitative oxidation or reduction of the substance of interest by various reaction paths involving one or more of these primary products. In preliminary experiments an industrial X-ray unit was used as a source of radiation, but it was not possible to obtain reproducible dose rates, so a cobalt-60 gamma-ray source was used in experiments reported here. Potentials were recorded continuously, so a plot of the entire titration curve was obtained. The principal advantages of such titrations are that once the solution is placed in the radiation source and the recorder started, no operator attention and no standard solution are required. Radiolysis of aqueous solutions produces intermediates such as .OH, H -,H20z,and the solvated electron that are powerful oxidizing or reducing agents but are obviously too unstable to be used as conventional titrants. It might be possible to find solution conditions such that one of these species could be used to cause a reaction that could not be produced by conventional volumetric titrations. In general the method is applicable when only one oxidationreduction couple is present, although there might be a situation in which unfavorable kinetics between the second oxidation-reduction couple and the radiolytic products would permit titration of the desired species. EXPERIMENTAL
Apparatus. The radiation source used was a Gammacell 200 cobalt-60 irradiator from Atomic Energy of Canada, Ltd. The central dose rate was 5.96 X 106 roentgens per hour at (6) M. A. Dewar and R. R. Hentz, Nucleonics, 13, No. 2, 54 (1955). (7) S . Musha, K. Sugimoto, and R. Tsujino, Birnseki Kagaku, 9, 1057 (1960); CA, 57, 4021i. (8) K. Sugimoto and S . Musha, Bull. Unio.Osaka Prefecture, Ser. A, 11,67 (1962); CA, 58,4073~.
the time of purchase. The source was equipped with access tubes which accommodated the electrical lead to the platinum indicator electrode, a salt bridge from an external saturated calomel electrode to the experimental solution, and gas inlets for an air-driven magnetic stirrer and for saturating the solution with oxygen when desired. Potentials were measured with a Sargent Model LS pH meter connected to a Sargent Multirange recorder. The absorbed dose rate in rads was measured with a Fricke dosimeter (9). A Beckman D U 2 spectrophotometer was used for the photometric measurement. Experimental absorbed dosages for titrations were calculated by multiplying the dose rate as determined by a Fricke dosimeter by the time required to reach the end point. Theoretical absorbed dosages were calculated from the appropriate G value and the concentration of reagent to be titrated using well-known conversion factors (9). Irradiation vessels were prepared by cutting off 150-ml borosilicate glass beakers about 21/2 inches from the bottom. These were tightly fitted with an aluminum cap having three holes for the platinum electrode, the salt bridge, and a gasinlet tube. The electrodes were always positioned in the same place and the irradiation vessel was always centered in the irradiation chamber. Electrode Treatment. The preparation or conditioning of the indicator electrode was very critical if reproducible results were to be obtained. The treatment procedure that worked was discovered by trial and error. Each new platinum electrode was stored for 15 minutes in hot, concentrated nitric acid, then made sufficiently cathodic in approximately 0.2M H2S04to evolve hydrogen, and the electrolysis was continued for 45 to 60 minutes. The electrode was then stored for several hours or overnight in an iron(II1) solution if iron(I1) or uranium(1V) was to be determined and in a cerium(1V) solution if cerium(1V) or dichromate was to be determined. Once an electrode was prepared by this method, the only treatment required between titrations was to store it in iron(II1) or cerium(IV), depending on what was to be determined. After being used for several weeks electrodes became poisoned, as evidenced by poorly defined and nonreproducible end points. Once an electrode became poisoned it could not be regenerated to a usable condition. Reagents. All solutions were prepared from triply distilled water with reagent grade chemicals used without further purification, except thallium(1) sulfate which was recrystallized twice from triply distilled water. Iron(P1) solutions were standardized daily by dichromate titration. Cerium(1V) solutions were standardized against the standard iron(II), and stock uranium(V1) solutions were standardized by reduction to uranium(1V) and titration with standard cerium(1V). Solutions of uranium(1V) that were irradiated were prepared by reducing an aliquot of the uranium(V1) in approximately 3.ON sulfuric acid in a Jones reductor and bubbling air through the resulting solution to oxidize all uranium(II1) to uranium(1V). The solvent systems found most suitable were : oxygen-saturated 0.8N HzS04 for iron(I1) determination; 0.8N HzS04and 0.001M thallium(1) sulfate for cerium (IV) determination; 0.4N and 0.001M thallium(1) sulfate for dichromate determination; and oxygen-saturated 3.ON H P S O and ~ 1 X 10-4M FeNH4(S04)2 for uranium determination. Procedure. The cell was filled with the solution to be irradiated and placed in the irradiation chamber with the Gammacell drawer in the up position. A 50-ml aliquot was always used, so the cell would be filled to the same height each time. This is important, because the dose rate varies in different parts of the cavity, and it is necessary that all solutions receive the same average dose rate. The plat(9) J. W. T. Spinks and R. J. Woods, “Introduction to Radiation Chemistry,” Wiley, New York, 1964, pp 108, 109.
ANALYTICAL CHEMISTRY, VOL. 41, NO. 14, DECEMBER 1969
2065
Table I. Effect of Sulfuric Acid Concentration on Iron Analysis Av. end point,
No. of
H2s04
concn, N
mrn.I.rno8e-l
titrations X low6 1.5 6 1.23 1.1 3 1.24 0.8 5 1.26 0.5 3 1.28 0.1 3 1.30 Fez+ concentration = 5.82 X IOw4 M. _
_
Range 1,21-1.24 1,234.25 1,25-1.27 1.27-1.29 1 ,29-1.32
~
Table 11. Effect of Foreign Salts or; Iron(II) End point End point, Fez+concn, M Added substance, M mm.l.mole-l X 0.001 NaCl 1.27, 1.28, 1.26 5.82 x 10-4 0.1 ZllSOe 1.26, 1.24, 1.25 5.79 x 10-4 0.001 MnClz 1.29, 1.28, 1.28 5.79 x 10-4 0.1 MnCh 1.33, 1.35, 1.33 6.08 X I
.-+ 0
0.90.
E
inum electrode and salt bridge were inserted into the solution. Oxygen was connected to the gas-inlet tube through an access tube when iron or uranium was titrated and was bubbled through the solution at least 15 minutes before starting irradiation. The drawer was then lowered, which placed the solution in the radiation source, and the recorder was started. Oxygen was bubbled through the solution continuously during irradiation of iron and uranium solutions. Solutions were stirred continuously with an air-driven magnetic stirrer. Irradiation was continued until an end point break occurred.
e
+ 0
a
0.404
0,30k-r--xT i rn e (min) 30
A. 02-saturated 0.0QlMFe(I1) in 0.8N H2S0, B. O?-saturatedQ.8N H2S0d
The radiation chemistry of aqueous iron(I1) solutions has been thoroughly studied (10). The following reactions OCCUT in an air-saturated iron(I1) solution irradiated by gammarays :
+ Fez+ = OH- + Fe3+ HzOz + Fez+ = OH- + .OH + Fe3+
(1)
A more extensive study of the determination of iron(1I) by gamma irradiation with potentiometric end point detection seemed desirable. Between concentrations of 1.99 X 10-4 and 3.88 X lO-3M iron(I1) a linear relationship between end point time (millimeters of chart paper) and concentration of iron@) was obtained. The curve did not extrapolate through the origin but gave a y-intercept of 24 mm (1-89 min) of chart paper. This represents a blank correction which arises partly from the fact that the recorder is started when the button is pushed to lower the drawer, but it takes about 10 seconds before the cell is lowered completely into the irradiation source. It also takes a finite time for radiolysis products to build up to a sufficient level to initiate reaction with iron(I1) at a constant rate. 'The millimeters of chart to reach the end point corrected for the 24-mm blank divided by the Concentration of iron(1I) was calculated for a series of 33 titrations containing between and 3.88 X 10-3M iron(E1). The average value 1.99 X
(IO) J. W. T. Spinks and R. J. Woods, "Introduction to Radiation Chemistry," Wiley, New York, 1964, p 268. 2066
e
0
Figure 1. Potential of P t electrode during gamma irradiation
RESULTS AND DISCUSSION
.OH
40
was 1.26 X lo5 mm.l.mole-l, and the standard deviation was 0.03 X lo5mm.l.mole-1 or 2.4% relative. The potential-time curves for a Pt electrode in air-saturated 0.8N sulfuric acid with and without iron(I1) are shown in Figure 1. The curve for the iron(I1) solution is very similar to an ordinary potentiometric titration curve. The end point was taken as the inflection point or the midpoint of the linear region of the sharply rising portion of the curve. Because it is necessary to know how critical the solution composition is, a number of experiments were performed in which iron(I1) was oxidized in solutions of varying composition. The effect of varying the concentration of sulfuric acid is shown in Table I. The end point increases slightly with decreasing concentration of sulfuric acid, so in an iron determination the concentration of sulfuric acid would have to be controlled. The variation in end point with change in sulfuric acid concentration is small enough between 0.8N and 1.5N that minor variations in acid concentration would cause no significant error. The potential break at the end point is greater in the solutions of higher sulfuric acid concentration. The effects of sodium chloride, zinc sulfate, and manganese(I1) chloride on the iron end point are summarized in Table 11. Sodium chloride and zinc sulfate do not interfere. Manganese(I1) chloride in large excess gives high results. To see if the method gives satisfactory results under conditions approximating an actual iron ore analysis, two samples of a commercially available student iron ore unknown were weighed out, dissolved, reduced in a Jones reductor, and diluted to the mark in volumetric flasks. Three aliquots of the first sample were titrated and gave an average value of 25.4%:. Four aliquots of the second sample gave an average of 26.5 %.
ANALYTICAL CHEMISTRY, VOL. 41, NO. 14, DECEMBER 1969
The analysis furnished by the manufacturer for this unknown was 26.01 %. At concentrations below 2 X 10-4M iron (11), the calibration curve deviated from linearity. The G(Fe3+) value is known to be constant between 0,0001 and 0.05M ( I I , 1 2 ) , so linearity would not be expected below 2 X lO-4M. Also below 2 X lO-4M iron(I1) the time required to reach the end point is short-1.34 minutes for 4.9 X 10-6Miron(II)-which introduces considerable error in measuring the end point time. There was no indication that 3.88 X 10-3M iron(I1) was an upper limit. Since G(Fe3+) is constant to 0.05M iron(II), the curve should be linear to this value. For the dose rate available in the Gammacell 200 used, the time required for oxidation of iron(I1) present in concentrations above 3.88 X 10-3Mwas too great-40.4 minutes at 3.88 X lO-3M. Reduction of Cerium(1V). In 0.8N sulfuric acid cerium(1V) is reduced by the following sequence of reactions with radicals and molecules produced by interaction of water molecules with gamma radiation (13): H’
+
0 2
=
HO;
+ Ce4+ = Ce3++ H+ + O2 H202+ Ce4+ = Ce3+ + HO; + H+ ‘OH + Ce3+ = Ce4+ + OHHO;
(6) (7) (8)
(9)
Because the molecular yield of cerium(II1) is much lower than iron(III), lower concentrations of cerium(1V) were irradiated. A linear calibration curve was obtained between and 8.08 X 10-‘M, with a concentrations of 1.34 X blank of 6.3 mm. The millimeters of chart per mole per liter of cerium(1V) corrected for blank was calculated for 38 titrations between these concentrations, and an average value of 9.26 X 106 mm-mole-’‘1 with a standard deviation of 0.45 x 106 or 4.9% relative was obtained. The presence of thallium(1) increases the molecular yield of cerium(II1) from 2.4 to 8.2 (14). This improved yield is due to reduction of hydroxyl radicals by thallium(1) which prevents a reoxidation of cerium(II1) to cerium(1V) by the hydroxyl radicals. The thallium(I1) produced reduces a cerium(IV), so the net effect of thallium(1) is to cause one cerium(1V) to be reduced for each hydroxyl radical formed. Thus Reaction 9 is replaced by
+ T1+ = TIZ++ OH-
(10)
+ Ce4+ = Tla++ Ce3+
(11)
‘OH
and T12+
A series of titrations of cerium(1V) in the presence of 0.001M thallium(1) sulfate gave a linear relationship between end point time and concentration between 1.34 X lO-‘M and 2.69 x 10-3M cerium(1V). The average value of millimeters per mole per liter of 38 titrations within this concentration range was 2.55 X 106mm .mole-l.l and the standard deviation was 0.018 X lo6 or 0.71 % relative. The presence of thallium(1) increases not only the cerium yield so that higher concentrations of cerium can be determined but also the precision by a factor of about 7. The presence of 8 X 10-4M lanthanum(III), 1.8 x lO-4M (11) J. Weiss, Nucleonics, 10, No. 7,28 (1952). (12) C. J. Hochanadel and J. A. Ghormley, J. Chem. Phys., 21, 880 (1953). (13) J. W. T. Spinks and R. J. Woods, “Introduction to Radiation Chemistry,” Wiley, New York, 1964, p 273. (14) T. J. Sworski, Radiation R e s , 4,483 (1956).
; ; i 0.78-
-
. I -
0
>
Y
-
- 0.76.-0
. I -
C
al 4-
o 0.74-
a
0.72-
0.700
10
20
T i rn e
30
40
(rnin)
Figure 2. Potential of Pt electrode in dichromate solution during gamma irradiation
Solution composition. 2.5 X lo-‘ A4 K2Cr2B7,0.4N HZS04,0.001MTl~S0~
thorium(IV), 1 X lO-3M neodymium(III), and 7.2 X 10-+M yttrium(II1) caused no interference in the titration of cerium(1V) with thallium(1) sulfate present. Phosphate present in concentration of 8 x lO-5M gave an end point 3.5 high for 0.001M cerium(1V). The effect of the presence of iron (111) on the end point was determined; up to a 10-fold excess of iron(II1) over cerium(1V) no change in time required for the cerium end point occurred. The presence of iron(I1I) caused an extra break in the potential-time curve at times beyond the cerium equivalence point. As the concentration of iron(II1) decreases, the second break shifts toward the cerium end point. The second break is even obtained when no iron(II1) is present in solution, if the electrode has been conditioned in iron(II1) solution prior to being used in the cerium(1V) reduction. The cerium end point is independent of sulfuric acid concentration between 0.8N and 1.6N. The end point depends slightly on the thallium(1) sulfate concentration. Increasing the thallium(1) sulfate concentration to 0.01M caused a 3 increase in end point time. Reduction of Dichromate. Initial experiments in reducing dichromate in acidic solutions were not promising. The G(Cr3+) value is low, 0.78 (13, which limited the method to very dilute solutions; and reproducible potentiometric end points could not be obtained. The mechanism for the reduction of dichromate is identical to that for cerium(IV) reduction, and the presence of thallium(1) has the same effect, increasing the G(Cr3+) value to 2.59 (15). When thallium(1) was added to the solutions, satisfactory results were obtained. This improvement may be due to the increased G(Cr3f) value or to the stabilizing effect of the thallium couple on the potential or a combination of both effects. The reduction of dichromate was performed in solutions 0.4N in sulfuric acid and 0.001M in thallium(1) sulfate. A typical potential-time curve for the reduction of dichromate is shown in Figure 2. The shape of this curve is different from potential-time curves of the other couples studied, but similar ~
(15) T. J. Sworski, J. Phys. Chem., 63, 823 (1959).
ANALYTICAL CHEMISTRY, VOL. 41, NO. 14, DECEMBER 1969
*.
2067
Table 111. Comparison of Experimental Dosage with Theoretical Dosage Required for Titration Substance titrated
Concn., M X 108
Fe aiCe4+ (no TI*) Ce4+(T1” present) Cr(V1)
0.100 1.oo 0.200
1.oo
Experimental dosage, rads
Theor. dosage, rads
67,500 43,500 114,000 70,800
60,800 39,200 115,000 73,600
error +9.93 4-9 89 -0.88 -3.95 I
to the shape of the potentiometric titration curve obtained for the titration of dichromate with iron(I1) (16). The peculiar potential behavior of a Pt electrode in a dichromate solution is believed to be associated with some kind of surface oxide formation (17). The average millimeters of chart per mole per liter corrected for the blank for 12 titrations within the concentration range 5.0 X 10-jM to 5.0 X 10-4Mpotassium dichromate was 1.60 x 106 with a standard deviation of 0.018 x 106 or 1.1 %relative. Increasing the thallium(1) sulfate concentration to 0.005M caused a 1 to 3 % increase in the dichromate end point. At a concentration of 0.01M thallium(1) sulfate an orange-red precipitate, presumably thallium(1) chromate, slowly formed. This caused extremely premature end points for dichromate. Doubling the sulfuric acid concentration caused a 2.4% decrease in end point time, and the potential break at the end point was less distinct. Oxidation of Uranium(1V). When uranium(1V) solutions were oxidized by irradiation, the resulting potential-time curves were impossible to predict or interpret. The potential-time behavior was similar to that of a Pt electrode in an irradiated sulfuric acid solution containing no redox couple other than the radiolytic products of water. Uranium (IV) was oxidized to uranium(V1) but the potential of the Pt electrode did not respond to the variation in concentration of uranium(1V) during the titration in a reasonable manner, presumably owing to the irreversibility of the uranium(1V)uranium(V1) couple. When iron(I1I) was added to the system, the shape of the potential-time curve was virtually the same as that of an iron(I1) oxidation curve. A linear relationship between end point time and uranium(1V) concentration was obtained between 2 x lO-4M and 3.0 X lO+M uranium(IV). All of these solutions contained 1.0 X 10-4M iron(II1). It was necessary to perform the titrations in a 3.0N sulfuric acid solution because at concentrations of sulfuric acid 1.ON or lower there was evidence of some direct oxidation of uranium(1V) by oxygen. A series of 24 titrations containing between 2.0 x lO-*M and 3 x 10-3M uranium(1V) had an average value of 2.78 x 105 mm-mole-1.1 and a standard deviation of 0.048 X lO5or 1.7% relative. Iron(I1I) added to a uranium(1V) solution will oxidize an equivalent amount of uranium(1V) to uranium(V1). When the reaction beaker is lowered into the irradiation source, one of the following two sequences could occur. A . The iron remains in the reduced state until the uranium(1V) is oxidized by radiolytic products, and then the iron(I1) is reoxidized to iron(II1). If the same concentration of iron (111) is used each time, the time required to reach the end point is proportional to uranium(1V) Concentration.
B. The iron(I1) is oxidized by radiolytic products and the resulting iron(II1) rapidly oxidizes uranium(1V) to uranium(V1). The rate-controlling step is the rate of oxidation of iron(I1) to iron(II1) by the radiolytic products. The iron functions like a homogenous catalyst. If the exchange current density of the iron couple is much larger than the uranium couple, the potential of the electrode would be determined primarily by the iron couple, for either path A or B. For either A or B the time required to reach the end point would be proportional to uranium concentration. If the uranium is oxidized by mechanism B, the time required to oxidize a given concentration of uranium(1V) should be twice that required to oxidize the same molar concentration of iron(I1) [2 iron(II1) react with one uranium(IV)], because the time depends on the kinetics of the oxidation of iron(I1) by radiolytic products. The time required to oxidize uranium(1V) by mechanism A would depend on the kinetics for oxidation of uranium(1V) by radiolytic products and would have no theoretical relationship to the time required to oxidize an equal concentration of iron(1I). Dividing the calibration figure for uranium(IV), 2.78 X 10j mmlmole-l, by the corresponding figure for iron(II), 1.26 X 105 mmlmole-1, gives a value of 2.2. This is slightly greater than 2.0, but these numbers can be expected to agree only roughly. The uranium work was done approximately 9 to 10 months after the iron work, so the intensity of the cobalt40 source had decreased by about 10% (18). This alone would give a ratio 10% greater than 2. The sulfuric acid concentration in the uranium work was 3.0.V) and no calibration figure was determined for iron in 3.ON sulfuric acid. It was determined that the value decreases slightly with increasing concentration of sulfuric acid, as shown in Table I. This factor would give a ratio slightly less than 2.0. In view of these factors it is believed that a ratio of 2.2 supports mechanism B. Also the same time was required to oxidize 4 X IO-*M uranium(1V) whether the concentration of iron(II1) was 1 X IO-jM, 1 X lO-dM, 3 x IO-dM, or 1 X 10M3M.The time required to reach the end point would vary with iron concentration 8 path A were followed, but would be independent of iron concentration if oxidation occurs by path B as long as the reaction between iron(II1) and uranium(1V) is fast compared to radiolytic oxidation of iron(I1). One additional reason for believing the oxidation proceeds by path B is that the value of G[U(VI)] varies with uranium(1V) concentration (19),so a linear calibration curve would not be obtained if uraniurn(1V) were being oxidized by radiolytic products. Relationship between Experimental Dosages and G Values. Although for strictly analytical purposes a linear calibration curve of millimeters of chart paper ES. concentration would suffice, it was desirable to relate actual absorbed dosages to the
(16) G. F. Smith and W. W. Brandt, ANAL.CHEM.,21,948 (1949). (17) H. A. Laitinen, “Chemical Analysis,” McGraw-Hill, New York, 1960, p. 335.
(18) J. W. T. Spinks and R. J. Woods, “‘Introduction to Radiation Chemistry,” Wiley, New York, 1964, p 27. (19) M. Haissinsky and M. Duflo, J. Chim. Phys., 53, 970 (1956).
2068
rn
ANALYTICAL CHEMISTRY, VOL. 41, NO. 14, DECEMBER 1 9 6 9
theoretical absorbed dosage required for complete oxidation or reduction of the species of interest based on G values taken from the literature. Comparison was made using the average of several titrations performed at a concentration level in about the middle of the useful range for iron(II), cerium(1V) without thallium, cerium(1V) with thallium, and chromium(V1) with thallium. Uranium was not compared because the mechanism for oxidation of uranium(1V) in the presence of iron(II1) has not been established unequivocally. The results are summarized in Table 111. The agreement
between experimental and theoretical dosages is only fair and emphasizes the desirability of using a calibration curve for each substance to be determined; however, a rough estimate of concentration can be obtained using experimental dosages and known G values. Exceptional agreement would not be expected because different workers commonly report G values that differ by several percentages (2). RECEIVED for review May 19, 1969. Accepted September 16,1969.
Vibrational Spectrometric and Electrochemical Lanthanum(lll)=Nitrate Complexes in Aqueous John Knoeckl Institute for Materials Research, National Bureau of Standards, Washington, D. C. 20234
NUMEROUS ION EXCHANGE (1-3) and solvent extraction (4, 5 ) schemes have been proposed for separating lanthanides in aqueous nitrate media. These mechanisms are usually explained in terms of the formation of nitrate complexes. In addition, nitrate is reduced polarographically to ammonia and hydroxylamine at the DME in the presence of lanthanum (111) and other rare earths at potentials considerably less cathodic than those at which ordinary nitrate reduction occurs (6). Wharton (7) has interpreted similar nitrate reduction in the presence of zirconium(II1) to imply the formation of a zirconium-nitrate complex. In the present work the aqueous lanthanum-nitrate system has been investigated polarographically, with a nitrate ionselective electrode, and by infrared and Raman spectroscopy. The appearance of the normally infrared forbidden y1 (A1') symmetric nitrate stretching mode at about 1050 cm-1 in lanthanum nitrate solutions is consistent with a lowering of nitrate symmetry to Cz, due to coordination. Similarly the appearance of four bands in the 1400-cm-1 region of both the infrared and Raman spectra is consistent with the splitting of the degenerate Y@') asymmetric stretching modes due to coordination and to solvent perturbation of the free nitrate (8).
The Raman depolarization ratio of the 1 4 0 0 - ~ m -spectral ~ envelope is indicative of bidentate nitrate coordination. Nitrate ion-selective electrode data indicate at least two complexes of 1:1 and 1 :3 lanthanum to nitrate ratio exist in aqueous solution. The 1:1 complex appears to be reduced polarographically with diffusion control. Present address, Department of Chemistry, North Dakota State University, Fargo, N. D. 58102 (1) J. Alstad and A. 0. Brunfeldt, Anal. Chim. Acta, 38,185 (1967). ( 2 ) T. Arends and M. L. Gallango, Brit. J . Hoematol., 11, 350 (1965). (3) J. Korksih, I. Hazan, and G. Arrhenius, Talanra, 10,865 (1963). (4) Z. Kolarik, Collection Czech. Chem. Commun., 32, 350 (1965). (5) E. B. Mikhlin and G. V. Korpusov, Zh. Neorgan. Khim., 10, 2780 (1965). (6) J. W. Collat and J. J. Lingane, J. Am. Chem. SOC.,76, 4214 (1954). (7) H. W. Wharton, J. Electroanal. Chem., 9, 134 (1965). (8) D. E. Irish and A. R. Davis, Can. J. Chem., 46,943 (1968).
1 1600
1
1 I500
1 1400
l 13CO
cm-1
I
l l 1200
I 1 1100
l 000
Figure 1. Infrared (top) and Raman (bottom) spectra of 1.5Mlanthanum nitrate 1400-cm-1 region of Raman spectrum resolved into four Gaussian components
EXPERINIEKTAL
Solutions. Lanthanum stock solutions were prepared by dissolving commercially available lanthanum nitrate hexahydrate or lanthanum chloride hexahydrate in distilled water or by dissolving lanthanum oxide in nitric or perchloric acid. The lanthanum content was determined in these solutions by titration with standard sodium fluoride solution. A cornmercial fluoride ion-selective electrode was used for end point detection. Nitrate stock solutions were prepared by dissolving accurately weighed portions of dried reagent grade potassium nitrate in distilled water. All other chemicals were reagent grade and used as received. Apparatus. Commercially available instrumentation was used throughout. Infrared spectra were obtained from samples run as smears between silver chloride plates. Raman spectra were obtained using a multipass reflecting cell and 6328-A He-'Ne laser excitation.
RESULTS AND DISCUSSION Vibrational Nitrate Spectra. Table I correlates the number and assignment of the infrared and Raman active frequencies which should be observed for free nitrate, Dah symmetry, and symmetry. The partial infrared spectrum bound nitrate, C2,,
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