fable 11.
Sample No.
Results of the Combustion-Coulometric Determinah’on of Sulfur in Iron and Various Steels
Sulfur, p.p.m. Electrolytic iron 8.0 9.6 36, 38 14, 12
4
High-Carbon steel (no alloying elements) 5 6
7
8 9 10
27 24, 25 28,27 21, 21 31, 32 30, 26
Steel, 5% Nil 0.5% Cr, 0.5% Mo 11 12
26, 26 28, 26
of 5 to 55 pap.m. sulfur by adding various weights of NBS Standard 55e to various weights of low-sulfur iron so t h a t the sample weight was 0.5 gram. A peak area of 0.08 square inch, representing the blank obtained from the combustion of the low-sulfur iron in a preignited boat with 0.13 gram of copper, was subtracted from all peakarea measurements. The remaining peak area was plotted against the amount of sulfur in the weight of NBS Standard 55e burned. The curve was a 23.4 y sulfur straight line with a slope of (inched2 * A similar plot of amount ‘sulfur‘ us. corrected peak area was obtained for
Sample No.
Sulfur,p.p.m. Maraging steel
13 14 15 16 17 18 19 20 21 22 23 24 25 26
14, 13, 12 10, 11 12 11, 12 14, 12 15, 15 18, 18 41, 32, 32 10, 11 21) 21 18, 18 11, 11 18, 18 14, 13
Steel, 57’ Nil 0.757’ Cr, 0.570 Mo 27
27, 27
NBS Standard 101e, with 0.5 gram of sulfur-free iron and 0.13 gram of copper added to each sample to accelerate the combustion. The similarity indicated that the same recovery could be obtained with high-alloy steels providing proper preheat and combustion times are used and the proper accelerators are added to the sample. A calibration curve for the range of 1 to 5.5 p.p.m. sulfur was obtained in the same manner as previously described except that the sensitivity was changed to 6.4 ohms and a bias setting of 120 was used. The slope of this calibration curve was 5.87 y sulfur (inches)2 ’
Sulfur results were obtained for a number of low-sulfur steels (Table 11). The deviation from the average of duplicate results is no more than 1 p.p.m. for all but two samples. The average deviation of the points on the calibration chart is also about 1 p.p.m. From theoretical calculations of the peak area based on 80% recovery, the estimated accuracy of results between 2 and 50 p.p.m. sulfur is &lo’%. LITERATURE CITED
(1) Abresch,
K., Claassen, I., “Die Coulometric Analyse,” Verlag Chemie, Weinheim, Germany, 1961. (2) “ASTM Chemical Analysis of Metals,” Philadelphia, Pa., 193, p. 25
(1965). (3) Ibid., p. 27. (4) Ibid., p. 94. (5) “British Standards Institution, BS1121, part lA,” British Standards Institution, London, 1957. ( 6 ) Burke, K. E., Davis, C. M., ANAL. CHEM. 34, 1747 (1962). (7) Fulton, J. W., Fryxell, R. E., Ibid., 31, 401 (1959). ( 8 ) Hibbs, L. E., Wilkins, D. H., Anal. Cham. Acta. 20,344 (1959). (9) Luke, c. L., ANAL. CnEM. 29, 1227 (1957). (10) Nydahl, F., Ibid., 26, 580 (1954). (11) Pigott, E. C., “Ferrous Analysis,”
pp. 442-68, Wiley, New York, 1953. (12) Rooney, R. C., Scott, F., J. Iron Steel Inst. (London)195, 417 (1960).
RECEIVEDfor review May 18, 1966. Accepted August 7 , . 1966. Pittsburgh Conference on Analytical Chemistry and Applied Spectroscopy, Pittsburgh, Pa., 1966.
CouIometric Titration of +3 Iron with Electrogenerated Chlorocuprous Ion JAMES J. LINGANE Department of Chemistry, Harvard University, Cambridge, Mass.
b The factors which govern the accuracy and precision of the coulometric titration of $3 iron with electrogenerated chlorocuprous ion in a hydrochloric acid-sulfuric acid medium have been identifled. Using potentiometric end point detection, 5to 50-mg. quantities of iron were titrated with a mean error of +O.OOl mg., and an average deviation of ? O . O l B mg. By reversing the polarity of the platinum generator electrode, 4-2 iron can be titrated anodically with electrogenerated chlorine, and thus mixtures of + 3 and +2 iron can be analyzed. Relatively few other metallic elements interfere.
I
in the coulometric titration of $ 3 iron to the + 2 state with electrogenerated chlorocuprous ion only partly because i t has BECAME INTERESTED
02 7 38
not been described heretofore, but chiefly because it appeared a prim* as an unfavorable reaction in a thermodynamic sense, and thus presented an interesting challenge. It is unfavorable thermodynamically because its equilibrium constant is only about 10+6, which is much smaller than one usually considers necessary for a precise titration. . However, the ferric-ferrous and cupric-cuprous couples both function reversibly in chloride media, so that reproducible and precise results are easily attained when the titration is monitored potentiometrically. Arthur and Donahue (1) were the first t o describe a reductometric coulometric titration of + 3 iron, and they employed electrogenerated 3 titanium. Lmgane and Kennedy (4) and Malmstadt and Roberts (6) also used electrogenerated + 3 titanium for this purpose.
+
+
The use of electrogenerated 5 uranium for the titration of + 3 iron was described by Edwards and Kern ($), while Schmid and Reilley (IO)employed the ferrous EDTA complex generated by the electroreduction of the ferric EDTA complex, and Dunham and Farrington (.@ utilized electrodeposited metallic copper. The characteristics of these methods, and their relative virtues, have been discussed in reference (6), which also describes previous coulometric titrations employing electrogenerated chlorocuprous ion. EXPERIMENTAL
As shown in Figure 1, the titration cell (capacity 100 comprised a platinum generator electrode and a silver auxiliary electrode, the latter being separated from the test solution VOL 38, NO. 1 1 , OCTOBER 1966
1489
REFERENCE Ag’AgCI
I I
Ag/AgCI AUXILIARY
1
It mg Fe / 100 cc.
0.7
i = 78.20 ma.
Pt INDICATOR
MAGNETIC STIRRER Figure 1.
Coulometric titration cell
Volume ca. 100 c m 8 The platinum generator electrode (2 X 1 2 cm.) was bent in a 300’ semi-circle to nearly surround the auxiliary electrode compartment.
in the central tube whose bottom was closed by a fine-porosity sintered-glass disk. The generator electrode was a 2- x 12-cm. platinum sheet, bent into a 300’ semi-circle. A silver wire electrode (area ca. 15 cm.2) was chosen as auxiliary electrode because in the chloride medium used the reaction a t it (Ag C1- = AgCl e) either simply removed C1from the solution when it functioned as anode, or added C1- when it acted as cathode, without producing interlering products. The level of solution in the auxiliary electrode compartment was kept above that in the test solution, a precaution which efficiently prevents mass flow of the test solution into the auxiliary electrode compartment. The solution in the auxiliary electrode compartment had the same composition as the test solution, except that it contained no cupric sulfate. The course of the titrations was followed potentiometrically by means of a platinum indicator electrode and Ag-AgC1-saturated KCl reference electrode. The lower end of the Perley type salt bridge tube (available from Leeds and Korthrup Co.) had a fine crack, formed by sealing a soft glass bead into the borosilicate glass tube, which provided a flowing junction of very small outflow rate. The potential of the Ag-AgC1, KCl (S) reference electrode is f0.197 volt with respect to the standard hydrogen electrode, or -0.045 volt with respect to the saturated calomel electrode. Air (oxygen) was removed from both comnartments of the cell with nitrogen. purified by passage through a “copper heater. ” The electrolysis current was auto-
+
1490
+
ANALYTICAL CHEMISTRY
Figure 2.
Influence of chloride ion concentration on the titration
curve matically held constant to *O.Ol% by the amperostat described on p. 503 et seq. of reference (6). The current was measured with a WoH potentiometer via the ohmic drop across a precisely known resistance. The currentmeasunng resistor was calibrated against a certified resistor from the Leeds and Northrup Co., and its value (10.067 ohms) was known t o somewhat better than =tO.Ol%. A recently certified, unsaturated Weston cell from the Eppley Laboratories was used to calibrate the potentiometer. Time was measured with a Standard Electric Time Co. Model S-10 electric stopclock (1000-second range, but repeatable). Although this clock is precise to the order of 1 0 . 0 1 second with respect to start-stop error and scale calibration, its accuracy depends on the constancy of the frequency of the 110volt a x . supply which drives it. As demonstrated elsewhere (5) short term variations in the frequency of nominal 60-cycleJ 110-volt commercial ax. lines can cause errors of as much as &0.15%. Therefore, the timing clock was driven from a frequency-regulated power supply accurate to +O.OOl% (American Time Products, Inc., Type 2005). Cupric sulfate pentahydrate was used as the source of cupric ion from which the chlorocuprous ion was electrogenerated. I n the early phase of this study a small “post-wave” was ob-
served, very close to the equivalence point of the iron titration, which impaired the precision with which the iron e.p. could be detected. The characteristics of this post-wave suggested that it was due to a n impurity reducible by chlorocuprous ion, and the impurity was finally traced to the cupric sulfate. When the cupric sulfate (originally hlallinckrodt Analytical Reagent grade) was recrystallized twice from water the post-wave disappeared. The hydrochloric acid used was redistilled, and the constant-boiling fraction was collected. The distilled water used had a specific conductance of less than 2 X 10-6 ohm-’ em.-’ The iron standard was Mallinckrodt Analytical Reagent iron wire, with a purity factor of 99.90%. This value was established by assay against specially purified potassium dichromate as follows. The sample of iron wire (about 1 gram) was dissolved in dilute sulfuric acid, under carbon dioxide to prevent air-oxidation. A few tenths of a per cent less t h m the theoretical quantity of Dhe purified potassium dichromate was then weighed out (calibrated weights and correction for air bouyancy) , dissolved in oxygen-free water, and added to the iron solution. The titration was finally finished with 0.01M ceric sulfate solution, using potentiometric equivalence point detec-
tion. I n two trials the observed purity factor of the iron wire was 99.89% and 99.90%. A standard iron solution was prepared by weight by dissolving 2.0361 grams of the iron wire in dilute hydrochloric acid in g. fused silica beaker, with care to avoid "spray loss." Nitric acid was then added to oxidize the iron to the ferric state, and the solution was evaporated to dryness three times with excess hydrochloric acid to remove nitrogen compounds. The residue of ferric chloride was finally dissolved in dilute hydrochloric acid, and the solution was diluted to 1002.82 grams. I n view of the purity factor of 99.90%, this solution contained 2.0284 mg. of iron per gram, and it was approximately 0.1111 in respect to hydrochloric acid. Samples of the standard iron solution were weighed out from a weight buret with a precision of 1 mg., corresponding to *0.002 mg. of iron. The technique of the coulometric titrations was as follows. The iron sample solution was first weighed into the titration cell, and diluted to approximately 100 cm.3 with the hydrochloric acid-cupric sulfate supporting electrolyte. The cell was then assembled, and dissolved air was removed from both the test solution and auxiliary electrode solution with nitrogen before starting the titration. Electrogeneration was performed incrementalwise, and the potential of the platinum indicator electrode was observed after each increment. The end point was recognized in the usual way from the maximal value of A E / A t .
The corresponding half-reactions are
+ e = Fe+Z+ C1C U ++ ~ 2 C1- + e CuC12FeC1+2
=
(2) (3)
These representations ere consonant with the observed facts that the formal potential of the ferric-ferrous couple becomes more reducing (less oxidizing or less positive) with increasing chloride ion concentration, whereas the formal potential of the cupric-cuprous couple shifts in the opposite direction. Therefore, with increasing concentration of chloride ion the degree of completeness of the overall titration reaction at the equivalence point becomes smaller. As a consequence, the rate of potential change dE/dt becomes less, and the precision of e.p. detection is impaired. This effect of the Concentration of chloride is demonstrated by the titration curves in 0.1 and 0.5M hydrochloric acid shown in Figure 2. I n both cases, 11 mg. of iron was titrated, and the concentration of cupric sulfate was 0.01F. Evidently, it is preferable to employ a relatively small concentration of chloride ion. However there is another side t o this coin, because the chloride ion concentration has to be kept large enough to stabilize the +1 copper as the soluble chlorocuprous ion. If the chloride ion concentration is too small, reduction of the cupric ion a t the generator cathode can yield insoluble CuC1, or even metallic copper with very RESULTS AND DISCUSSION small chloride ion concentrations. Although both CuCl and Cu reduce ferric Because several chloro complex iron, these reactions, being heterospecies of copper and iron are present, geneous, are intrinsically slower than and their relative proportions depend the homogeneous reduction by CuC12-, on the chloride ion concentration, no and can cause inconvenient slowness in single equation uniquely represents the the establishment of equilibrium near titration reaction. From data for the the equivalence point. A hydrochloric formation constants of the various acid concentration of 01.111' is a satischloro complex ions quoted by Meites factory compromise, and it was used in ( 8 ) ,it follows that a t a chloride ion conall subsequent titrations. centration of 0.1M the predominant In addition to titrating $3 iron with species of + I copper is CuC12-. The electrogenerated CuC12-, any $2 iron major + 2 copper species (in 0.1X originally present can be determined by chloride ion and with 10 m F total reversing the polarity of the generator copper) is the aquo cupric ion, C U + ~ , electrode to titrate the + 2 iron with (about 90%)) with about 10% CuCI+, electrogenerated chlorine-e.g., analysis and a very small proportion of CuC12. of a mixture of +2 and + 3 iron. BeWith 1 to 10 m F iron in 0.131 hydrw cause the iron in both ovidation states is chloric acid, the predominant ferrous present as cationic species, it tends to species is F e f 2 (about 4/5), with about be transferred through the sintered1/5 as FeCl+, and only a few per cent glass disk into the auxiliary electrode FeCL The chief ferric species in 0,LlI' compartment during anodic titrationhydrochloric acid iq FeC1+2, with someLe., when the silver auxiliary electrode what smaller and approximately equal functions as cathode. With only 0.1-11 proportions of FeCl,+ and Fe+3, and a hydrochloric acid the transference very small amount of FeC13. Acnumbers of the iron species are large cordingly, ui) t o a chloride ion conenough so that this electrical migration centration of about 0.5.21, the titration can cause a significant error. The reaction may be represented approxitransference numbers of the iron species mately by can be decreased very nearly to zero, and thus this error can be prevented, by FeC1+2 CuC12- = adding an additional supporting Fe+2 C U + ~ 3 CI- (1) electrolyte of high cation transference
*
+
+
+
number-Le., a strong acid, For this purpose sulfuric acid is suitable, and the supporting electrolyte finally adopted was 0.1M hydrochloric acid containing 2 cmUS of concentrated sulfuric acid per 100 cmS3(0.36M HQSO~). From the observed potentials at various stages of titration in this 0.1 M HCI 0.36M HzSOl supporting electrolyte the formal potential of the ferric-ferrous couple-Le., of reaction 2 above-was found to be 0.483 f 3 volts, and that of the cupric-cuprous 3 volts, couple (reaction 3) is 0.127 both us. the Ag/AgCl, KCl (S) reference electrode. Hence the formal potential of the titration reaction (Equation 1) is only 0.356 volt in this supporting electrolyte, or the formal equilibrium constant is only 1 x 10+6,
+
*
viz.
Consequently, the rate of potential change a t the e.p. is relatively small. In spite of this, the e.p. can be recognized precisely, because both couples function nearly reversibly, and thus yield reproducible potential readings. Prior to the equivalence point the potential becomes constant very quickly. At, and slightly beyond, the equivalence point the potential rapidly drifts downward (more reducing) for about 2 minutes after generation is interrupted. It then remains steady for several minutes, but finally begins a slow upward drift. Most likely this eventual upward drift is caused by a trace of oxygen, which oxidizes the CuC12-. Removal of thc last vestiges of oxygen from a solution by bubbling nitrogen through it is a very slow process, of uncertain completeness. The minimum potential, attained after about 2 minutes, wab employed. Although the titration reaction itself (Equation 1) is symmetrical, in the sense that the + 3 iron and 1 copper react. in a 1: 1 mole ratio, the titration curve is more or less asymmetrical because an excess of Cu"2 necessarily is present. Such asymmetry causes the end point (defined by the maximal value of dE/dt) to occur slightly in advance of the true equivalence point. The resulting negative tit.ration error can be kept negligibly small (order of a few hundreths of a per cent) if the concentration of excess C U + ~ is not greater than about ten times the concentration of iron. Furthermore, an unnecessarily large excess of Cu+* is undesirable because it increases the degree of inconipietrne?- of the titration reaction a t the e.p., and thus decreases the rate of potent,ial change and impairs the precision of e.p. det'ection. On the other hand, the concentration of cupric ion must be large enough, with respect to the current density used and the stirring efficiency, so that its reduction proceeds
+
VOL 38, NO. 1 1 , OCTOBER 1966
B
1491
Table I. Comparison of Observed and Theoretical Equivalence Point Potentials
+
Sup rtin electrolyte waa 0.1OM HC1 0.313% HJO,, with 10 mF cupric sulfate DifIron, E..,. V. ference, d Obsd. Theor. mv. - 13 0.307 9.1 0.307 +7 8.8 -5 0.312 6.1 -1 0.317 4.0 -8 0.317 3.8 -2 0,328 1.7 0.332 +1 1.2 -1 0.334 1.0 Av. f 5
only to CuC12- and not to metallic copper. With the particular platinum genem tor electrode (area of one side 24 cm.2), and stirring condition used, a cupric sulfate concentration of 10 mF was adequate with generating currents up t o about 90 ma. This concentration of cupric sulfate is suitable for titrating 5 to 50 mg. of iron-Le., about 1 to 10 mF iron-in about 100 to 1100 seconds with a %ma. generating current. The stirring in the cell of Figure 1 is only moderately efficient, and by increasing the stirring efficiency-e.g., by use of a rapidly rotated generator electrodethe maximal permissible current density for a given concentration of cupric ion doubtless could be increased several fold, and the titration time could thus be shortened correspondingly. Since both couples behave reversibly the potential a t any point in the titration can be expressed in terms of either one of them, and the appropriate formal potentials a t a given chloride ion concentration, by
From Equation 4 the aonaentration of CuCh- at any point is expreasible in
necessarily is as large m, or larger than, CP., and is constant, so the equivalence
terms of the concentrations of +2 and +3 iron and +2 copper by
point potential depends on the quantity of iron titrated. Table I shows the equivalence point potentials actually observed with variom quantities of iron (expressed as milliformal concentration) compared with the values mlculated from Equation 14, using E ~ F , + I . F=, +0.483 ~ volt, K = 1 X lo*, and CcU = 10 mF. The values are referred to the Ag/AgCl, KCl(S) reference electrode. The average agreement to A5 mv. is a realistic indication of the degree of reproducibility of the potentials in the immediate vicinity of the equivalence point. The observed values of E..,. in Table I pertain to end points defined by the maximal value of A E / A t . Actually, because of the asymmetry of the titration curve, these end points slightly precede the true equivalence point. However, as shown by the theoretical titration data in Table 11, this titration error is very small. The data in Table I1 were calculated by means of Equations 10, 13, and 14, for the relatively unfavorable case of only 1 mF +3 iron with 10 mF cupric ion. I n this case the maximal value of A E / A t occurs a t a “percentage titrated” of 99.9%, so that the titration error is only -0.1% (-0.005 mg. of iron). With larger quantities of iron-Le., a larger CP,/CC~ ratio) the titration error rapidly becomes much smaller. In actual practice, of course, one selects a concentration of cupric ion, so that the ratio CF./CC, is not smaller than about 0.1. The accuracy and precision of the titration of 5- to 50-mg. quantities of ferric iron, in the 0.10M HC1 f 0.36M HzS04 0.010M CuSOa supporting electrolyte are shown by the results in Table 111. All of these titrations were performed with the same constant generating current, adjusted to such a value (86.39 f 0.01 ma.) that 1-second generation time corresponded to exactly 0.05 mg. of iron. In the vicinity of the e.p. the generation current was applied in equal increments, each equivalent to about 0.5% of the total titration time, and the end point was detected in the manner shown in Table 11. In preliminary trials it was discovered that the standard ferric iron solution contained a considerable amount of ferrous iron. This was evident from the initial potential, which was always close to 0.62 volt, whereas a fully oxidizediron solution has a potential of about 0.75 volt. Comparisons of the initial potentials with the formal potential of the ferric-ferrous couple in this medium (the latter evaluated from pbtential readings in the middle, well-poised sections of titration curves), and, more precisely, direct anodic titration to a potential of 0.75 volt, both showed the presence of 0.60 f 0.04% ferrous iron. Correction
(CuC12-) =
(Fefz) ( C U + ~ ) (~ec1+9)K
(1’
where K is the equilibrium constant of the titration reaction. Until the equivalence point is greatly exceeded, the concentration of Cu+a is so much greater than that of CuC12- that it remains constant, and is equal to the total formal concentration, CcU,of the copper species. Combining these relations with Equation 5, the general equation of the titration curve up to, and including, the equivalence point is
E
- 0.059 log
[
= EfFe+a.ge+r
& 7
i,
(Fe+2) Ccu (10)
- E + (FeC1+2)K
Beyond the equivalence point the potential is expressed most directly in terms of the cupric-cuprous couple (Equation 6). Representing the excess generation time beyond the e.p. by tz,, the concentration of CuC12- is (CuC12-)
=
it,, E
+ (FeClf2)
(11)
The concentration of C U +will ~ be CcU (CuC12-), and the concentration of FeC1+2 will be (FeCl+2)
=
(Fef2) ( C U + ~ ) K (CUClz-)
or, since (Fe+Z) beyond the e.p. is very nearly equal to C F ~ ,
E= 0.059 log
~
(Fe+2) (FeC1+2) (5)
or
Hence the potential beyond the e.p. is given by
E
= Efcu+z,cucIs- - 0.059 log
E = 0.059 log
~
(CuC12-) (Cu+2)
(6)
At any time t (sec.), with a constant current i (amp.), in a solution volume V (dm.a), the sum of the concentrations of CuC12- and Fe+2is (CuCh-)
+ (Fe+? = F~it
(7)
where F is the faraday constant (96,487 abs. coulombs). Representing the total formal concentration of f 2 and + 3 iron by C F e , the concentration of + 3 iron at any point is (FeClf2)
1492
=
Cp,
- (Fe+2) = C F ~-
ANALYTICAL CHEMISTRY
A t the equivalence point (CuClZ-) is negligible with respect to (Fe+2),but IS equal to (FeC1+2). Therefore, from Equation 10 the equivalence point POtential is =
EJF~+~ -, F ~ + ~ 0.059 C F (14) ~ -log 2
CC“
If + 3 iron were titrated with a standard CuC12- solution in the classical volumetic manner CF, would be equal to Ccu,and E.,*. would be a constant. In the coulometric titration, however, CcU
+
for this has been applied to the data in Table 111. On first thought, one is tempted to attribute this ferrous iron to incomplete oxidation during the preparation of the standard iron solution. However, a generous excess of nitric acid had been used to oxidize the iron (excess evident from the evolution of oxides of nitrogen and chlorine) so incomplete oxidation is very unlikely. Another possibility is that the ferrous iron was produced when the solution was brought into contact with the large platinum generator electrode, because the oxidation of platinum by ferric iron in chloride media is familiar. From chronopotentiometric measurements with the same platinum electrode and supporting electrolyte used in the coulometric titrations, the quantity of electricity required t o oxidize the platinum generator electrode, to form a film of PtC12, was found to be only 10 millicoulombs, or 0.22 millicoulombs/cm.Z since the area of the electrode was 48 cm.2 This value agrees with those observed by Peters and Lingane (9). However, with 5 mg. of ferric iron per 100 0111.3 (about lo-' mole or 10 coulombs) it would account for only about one sixth of the observed amount of ferrous iron, and t o still less with the larger quantities of iron. Furthermore, the percentage of ferrous iron was independent of the quantity of the standard iron solution used, so that reduction of the ferric iron by the platinum electrode can account for only a minor fraction of the observed amount of ferrous iron. I believe that most of the ferrous iron was formed by the reduction of the ferric iron by chloride ion when, in the preparation of the standard iron solution, the solution was evaporated to dryness three times with excess hydrochloric acid after oxidation with the nitric acid. T o be sure, this reaction has a very small equilibrium constant,
Table II. Theoretical Titration Data
0.1 dm.8 +=0.36M Ha04
f, %
+ 2 C1-
= 2 Fe+*
+ Clz;
= 1 mF, Ccu = 10 mF, V Supporting electrolyte 0.1M HCl ( Fe +l) (FeCl+S) E , mv.
99.0
viz.
2 Fe+3
than that of the ferric-ferrous couple in this particular medium, the find end point is defined very sharply between about 0.7 and 1.0 volt. I n contrast to the highly accurate cathodio titrations, these anodic titrations all show a positive error, which increases as the quantity of ferrous ion titrated is decreased. This error probably stems from less than 100% current efficiency for chlorine generation, which is not surprising with only 0.1M chloride ion. With the larger amounts of iron a considerable fraction undergoes direct electrooxidation, but this fraction decreases as the concentration of ferrous ion decreases, and with the smallest quantities of iron most of the oxidation proceeds via electrogenerated chlorine. Hence, inefficient chlorine generation causes a larger error with the smaller quantities of iron. This anodic error can be suppressed by increasing the chloride ion concentration, which increases the current efficiency for chlorine generation. This is demonstrated in the case of the 6.8-mg. sample. I n this case the solution was first back-titrated t o the CuC12- end point. Then 6.7 grams of solid potassium chloride was added to raise the chloride ion concentration to 1M before completing the
are compatible with the quantity of +2 iron actually observed. From Table 111,the cathodic titration of ferric iron is seen to be remarkably accurate, and the mean error of +0.001 mg. (&0.018 mg.) indicates that there is no significant systematic error. The sharp eyed reader will have noticed that the error with the 48.9-mg. sample looms abnormally large (although still only -O.l%), and if this trial is disregarded the mean error is +0.007 mg. (*0.010 mg.). Either way, i t is evident that the accuracy of the titration is limited only by the precision of e.p. detection, and, of course, by the skill of the operator in avoiding accidental errors. After each cathodic titration (except with the 9.6 mg. sample) the solution was back-titrated anodically, first to a n end point for the excess CuC12- which had been generated, and then to the end point for the re-oxidation of the ferrous iron. I n these anodic titrations, some of the ferrous iron underwent direct electrooxidation at the platinum electrode, and the remainder was oxidized by the chlorine produced by the electrooxidation of the chloride ion. Because the formal potential of the chlorinechloride couple is about 0.6 volt greater
90.4
99.5
152
100.0
324
100.5
653
101.0
1093
Maximal A E / A t at f
= 99.5
Af
367
> > 334 > 316 >
13
354
20
> >
+7
>
--A
-2
18 13
303
+ (f
AIE -
AE Af, mv.
)
X 0.5 .= 99.9%
K = 10-11 However, the volatility of the chlorine, and the high chloride ion activity in constant boiling hydrochloric acid, both conduce to a significant amount of reduction of the + 3 iron. With a relatively pure ferric chloride solutione.g., (Fe+3)/(Fe+2) = 104, and at a chloride ion concentration of 6M, it follows from the above equilibrium constant that the initial partial pressure of chlorine will be 27 mm. of mercury, a very considerable value. After reduction of 0.6% of the + 3 iron, the ratio (Fe+$)/(Fe+*)becomes 167, the partial pressure of chlorine decreases to about 0.008 mm., and further reduction would be expected to occur very slowly. Evidently, both the magnitudes and range of these chlorine partial pressures
Table 111.
Performance Data
I n all cases the generating current was 86.39 ma., and 1-sec. generation time corresponded to exactly 0.05 mg. of iron. The su portin electrolyte was 0.10M HCI 0.36M H&O4 0.010M CuS04. The value of t i e fara8ay constant used was 96,487 absolute coulombs, and the atomic weight of iron on the carbon-12 scale is 55.847.
+
Fe taken, mg. 50.838 48.947 34.225 22.180 21.020 9.615 6.825 5.705
a
Fe found, mg. Anodic 50.858 50.903 48.898 48,927 34.230 34.380 22.160 22.415 21.041 21,255 9.618 6.845 6.870. 5.715 5.955
Cathodic
+
Error, mg.
Cathodic
Anodic
+0.020 -0.049 +0.005 -0.020 f0.021 4-0.003 +0.020 +o. 010 Mean +0.001 Av. dev. f0.018
4-0.065 -0.020 + O . 155 +0.235 + O . 235 +0:645a
4-0.250
6.7 g. KC1 added before anodic titration.
VOL 38, NO. 1 1 , OCTOBER 1966
1493
anodic titration. The error of +0.045 mg. is only about one fifth of what it would have been with only 0.1M chloride. Any substance reducible by chlorocuprous ion will, of course, interfere with the titration of ferric iron. Actually, however, most of the other elements usually associated with iron-e.g., Co, Ni, +3 Cr, + 2 Mn, +4 Ti, +6 U, and +6 Mc-should not interfere. Vanadium, if originally present in the +4 state should not interfere. However, as shown by Meier, Meyers, and Swift (6,7), pentapositive vanadium is reduced to the + 4 state by chlorocuprous
ion. The potential of the VOa+fVO+2 couple is about 0.3 volt more positive (more oxidizing) than that of the ferricferrous couple, and therefore it should be possible to analyze mixtures of +5 vanadium and + 3 iron by the present method. LITERATURE CITED
(1) Arthur, P., Donahue, J. F., ANAL. CHEM.24, 1612 (1952). (2) Dunham, J. M., Farrington, P. S., Zbzd., 28, 1510 (1956). (3) Edwards, K. W., Kern, D. M., Zbid., 28, 1876 (1956). (4) Lingane, J. J., Kennedy, J. H., And. Chzm. Acta. 15, 465 (1956).
( 5 ) Lingane, ,,J. J., “Electroanalytical
2nd ed., Interscience, gkm?ix, 1958. (6) Malmstadt, H. V., Roberts, C. B.,
ANAL. CHEM. 28, 1408,’ 1412, 1884 (1956). (7) Meier, D. J., Meyers, R. J., Swift, E. H., J. Am. Chem. SOC. 71, 2340 (1949). (8) Meites, L., “Handbook of Analytical Chemistry,” McGraw-Hill, New York, 1963. (9) Peters, D. G., Lingane, J. J., J . Electroanal. Chem. 4, 193 (1962). (10) Schmid, R. W., Reilley, C. N., ANAL. CHEM.28, 520 (1956).
RECEIVEDfor review June 27, 1966. Accepted August 12, 1966.
Voltammetry and Related Studies of Uranium in Molten Lithium Fluoride-Beryllium FluorideZirconium Fluoride GLEB MAMANTOV’ and D. L. MANNING Analytical Chemistry Division, Oak Ridge National Laboratory, Oak Ridge, Tenn.
b Electrochemical reduction and oxidation of U(IV) in LiF-BeFz-ZrFd (64-34-1.8 mole % and 65-29-5 mole %; the latter composition is used as the solvent for UFa in the molten salt reactor experiment) was investigated by voltammetry, chronopotentiometry, and chronoamperornetry in the temperature range 480-620’ C. Platinum, graphite, molybdenum, tungsten, and tantalum working electrodes have been employed. At 500” C. the reduction of uranium(lV) at platinum is a reversible, one-electron process. This conclusion is based on Nernstian log plots and the diagnostic criteria of linear sweep voltammetry. Well defined and reproducible voltammograms and chronopotentiograms have been obtained at concentrations UFI. The as high as 0.8 mole diffusion coefficient of U(IV) at 500’ C. was calculated as 2 X cm.2/second; the activation energy for diffusion, obtained from a plot of log D vs. 1 / T I was found to be 1 1 kcal./ mole. The results obtained provide the basis for an in situ electroanalytical method for the determination of uranium in molten fluorides.
yo
A
number of electrochemical reactions have been studied in molten salt solvents during the past twenty years. As evidenced by recent reviews (8, 16-18), the most common solvents have been the molten nitrates (for example, KaKOrKN08, 50-50 inole yo)and the molten chlorides (for example, LiC1-KC1 eutectic, 59-41 LARGE
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ANALYTICAL CHEMISTRY
mole %). Few fundamental investigations have been carried out in molten fluorides, although the fluorides are of considerable interest from the applied point of view, for example, the well known Hall process for aluminum. Other recent applications include the use of molten fluorides (38) as fuel and coolant in the molten salt reactor (RISRE) currently under development at the Oak Ridge National Laboratory, and the electrodeposition of coherent deposits of refractory metals, such as tantalum, niobium, and zirconium, from molten fluoride baths (26, 27, 36, 37‘). The most likely reason for the apparent lack of interest in electrochemical studies in molten fluorides is the inability to use (on a long-term basis) glass or quartz as insulating or container materials, because of attack by molten fluorides. We have been concerned with electroanalytical measurements in molten fluorides (81-26). Pizzini, Sternheim, and coworkers (30-39) have been interested in overvoltage measurements by the galvanostatic method in molten fluoride solvents. The electrochemistry of uranium has been studied in molten LiC1-KC1 and MgCITNaC1-KC1 eutectics a t 450” C. by Hill, Perano, and Osteryoung (12). Hill et al. determined the standard potentials for the couples U(II1)-U(0) and U(1V)-U(II1) in MgClz-NaCl-KCl and U(1V)-U(II1) and COZ(VI)-UO~ (IV) in LiC1-KC1. Slow-scan voltammetric studies of U(III), U(IV), and U02(VI) were performed and a coulo-
metric titration procedure with electrogenerated Pt(I1) was developed for the determination of U(II1). The reduction of UOz(VI) was shown to be a twoelectron process. No oxidation waves of U(1V) were observed. Stromatt (39) studied the electroreduction of U02(VI) in molten equimolar NaC1-KC1 at 716’ C., by chronopotentiometry and electrode impedance measurements. His potential-time curves indicated that the reaction occurred in two steps a t nearly the same potential. Pseudo capacity as a function of the potential of ths platinum electrode showed two peaks, substantiating chronopotentiometric results. Stromatt concludes that the electrode reaction proceeds by a reduction of UOz(VI) to soluble VOz(V), followed by the reduction of U02(V) to UOs which forms an insoluble product on the electrode. Other electrochemical investigations of uranium in molten salt media include the work by Inman, Hills, Young, and Bockris (13) and Gruen and Osteryoung (11). References to other work are given by Hill et al. (12) and Stromatt (39). The analytical chemistry of uranium has been reviewed by Booman and Rein (2) and by several Russian workers (29). In this work, we were primarily interested in the reduction of U(IV) to provide the basis for an in situ electroanalytical method for uranium in molten 1 Department of Chemistry, University of Tennessee, Knoxville, Tenn.
