Coulometric Titration of Iron(l1) with Chlorine PAUL S. FARRINGTON, WILLIAM P. SCHAEFER,’ and JOHN M. DUNHAM2 Department of Chemistry, University o f California at 10s Angeles, 10s Angeles, Calif.
b The coulometric titration of ferrous iron with chlorine in 2F to 6F hydrochloric acid is both precise and accurate for samples of more than 500 p g . of iron in 50 ml. of solution. For satisfactory results, a catalyst such as cupric ion must be present. Zinc, silver, or bismuth reductors were used to reduce small iron samples. Vanadium and molybdenum interfere in the titration while manganese, chromium, and titanium do not. Cuprous ion can be titrated as easily as ferrous. The method has been tested on Bureau of Standards samples.
F
IRON has been titrated coulometrically using Ce(1V) (2, 9, 7), Mn(II1) (8, I 4 ) , and Brp (IS) as intermediates in constant current processes. The most accurate results for m a l l amounts of iron were reported by Cooke, Reilley, and Furman (Z), who used Ce(1V) as their intermediate. -4s is shown in the present work, chloride interferes in this titration. Because hydrochloric acid is useful for dissolving iron samples and because chloride ion is necessary to the operation of certain reductors, the possibility of using chlorine as the intermediate was investigated. The reaction between ferrous iron and chlorine has been used as the basis for a chlorine analyzer (11) in which ferrous ion is generated to titrate amounts of chlorine in the parts per million range. Coulometric analyses of more concentrated chlorine solutions by reaction with electrogenerated ferrous ion have also been reported (12). We have found that the ferrous-chlorine reaction is not instantaneous, but that either large amounts of phosphoric acid (concentration about 6F) or small amounts of cupric ion (concentration about 0.01F) catalyze it (6). The mechanism of this catalysis is not yet known; the chlorine-ferrous reaction could conceivably take place via either a Cu(1) or a Cu(II1) intermediate. Studies are now being initiated to determine which (if either) of these possibilities is correct. This paper reports the procedure developed ERROUS
Present addresa, Divisipn of Chemistry and Chemical Engineenng, California Institute of Twhnology, Pasadena, Calif. ’ Present addrens, Sterling-Winthrop Research Institute, Itensselaer, N. Y. 13 18
0
ANALYTICAL CHEMISTRY
for the determination of iron by chlorine titration, using cupric ion as a catalyst. EXPERIMENTAL
The constant-current generator circuit was similar to the one used by Meier, Myers, and Swift (9); a 10-ma. generating current was used. Two indicating systems were compared: an amperometric circuit using two similar platinum electrodes with a voltage impressed across them and the modified amperometric system described by Cooke et al. ( 2 ) . This latter proved to be easier to use and much more reproducible, although not as sensitive as the former in terms of indicator current diange per second of generation time. The potential of the platinum electrode was set a t +650 mv. us. the saturated calomel electrode, and the current in the circuit was measured with a 0- to 15-fia. meter. The titration cell was similar to the one used by Myers and Swift (IO). Magnetic stirring was used; since the end point was taken a t zero current, the stirring rate did not affect this indication. The generator cathode was contained in a guard tube. The lower end of the tube was closed by a fineporosity, sintered-glass disk; the tube contained 2F hydrochloric acid in all cases, and the level of the hydrochloric acid was kept qbove that of the solution in the cell. The generator anode was a piece of platinum foil of about 50 sq. mm. total area. Samples were introduced with a 0.5-ml. micropipet which was treated with Desicote (Beckman Instruments, Inc.) and calibrated ‘‘to deliver” by weighing water. The reproducibility of this pipet was excellent; the average deviation from the average of 16 calibrations (0.4979 ml.) was *0.05’%. The range of volumes measured was 0.0010 ml. for these same calibrations, made over a 6-month period during which the Desicote was changed three times. Time was measured with a Standard Electric Time Co. timer, Model S-10, connected to the 60cycle mains and actuated by the same switch used to control generation. Reagents. I n determining suitable conditions for the titration, a solution of ferrous ammonium sulfate in 0 1F sulfuric acid was used as the source of iron. This solution was standardized within 24 hours of each use by titration with potassium permanganate and ceric sulfate solutions, which were in turn standardized against Bureau of Standards arsenious oxide and/or sodium oxalate. The result for iron(I1) concentration was not accepted unless over-all agreement to Apparatus.
j = O . l % was obtained. All other cheinicals were reagent grade. Titrations. These were carried out to determine the minimum amount of iron that could be titrated by this method, the permissible range of hydrochloric acid and cupric ion concentrations, and possible interferences. About 50 ml of hydrochloric acid of the desired concentration were placed in the cell, the desired amounts of cupric chloride and ferric chloride (as solutions in 2F hydrochloric acid) were added, and nitrogen bubbled through to remove oxygen. Chlorine was generated on both the anode and the indicator electrode for 20 seconds to sensitize the indicator and oxidize any impurities, and the excess removed by again bubbling nitrogen through the cell. The potential applied to the indicator electrode was such that a small excess of chlorine (about 10-8 equivalent) was present a t zero current, and this was generated after the flow of nitrogen had been switched to pass over the surface of the solution. The indicator circuit was then opened, the sample added, and generation of chlorine begun. The indicator current was observed only a t intervals until it dropped below 15 pa. because the indicator electrode reaction consumed ferrous iron. The circuit was left closed after the current had dropped below 15pa. The titration was interrupted a t a current of 2 to 5 ha. and continued stepwise to zero current. As many as 10 samples of 600 pg. each could be run in thesame solution, one after the other, without any further treatment of the cell or indicator electrode, so long as each titration was stopped a t zero current. Even if the end point was passed the excess chlorine could be removed with N2and more titrations run after adjusting the indicator current to zero. RESULTS AND DISCUSSION
Choice of Indicator Electrode Potential. Titrations were made with
various voltages applied to the indicator electrode and the resulting indicator current plotted Y.S. time (Figure 1). From these data i t appeared that a potential between f 6 0 0 and +700 mv. us. S.C.E. could be used. The slope of the indicator current-titration time curve is greater after the equivalence point than before; an end point with some chlorine present is, therefore, indicated for greater sensitivity. Since the chlorine is volatile, this amount should be as small as possible. Because the titration is, in effect, from
a preset potential of the solution back to the same potential, the exact value chosen should cause no error. This will be strictly true only if conditions a t the start and end of a titration are nearly the same. To ensure this, ferric iron was added to all solutions in an amount at least 10 times larger than the sample to be titrated. The potential commonly applied to the indicator electrode was +650 mv. us. S.C.E.; titrations made with 625 or 675 mv. applied gave equally accurate results. Hydrochloric Acid Concentration, The reaction between ferrous iron and chlorine was somewhat slow near the end point in 1F hydrochloric acid; adding cupric ion did not increase the speed as much as was desired. In 2F hydrochloric acid with cupric ion present equilibrium was attained rapidly; the indicator current stabilized in 2 to 5 seconds after an addition of chlorine. The data of Table I show that a maximum error of 1 0 . 2 % is obtained in titrating about 600 pg. of iron in 2F to 6F hydrochloric acid, but that above a concentration of 6F large positive errors result, presumably from oxidation of the platinum generator anode (6); the use of 2 to 3F hydrochloric acid is, therefore, recommended. Copper(I1) Concentration. I n the absence of copper in the titration solution the indicator current was slow to stabilizc when small amounts of chlorine were added near the equivalence point. If the solution was made 10-6F in Cu(II), the indicator current responded noticeably more rapidly, but the response was not judged fast enough for good titration unless the Cu(I1) concentration was a t least 0.03F. No improvement in response time was observed if the Cu(I1) concentration was increased above 0.05F. Large concentrations of copper (0.5F; see Table I) do not affect the accuracy of the titration. On the other hand, the presence of more than about 0.05F Fe(II1) reduces the sensitivity of the indicator system. The current does not respond t o the generation of several microequivalents of chlorine when the ferric concentration is about 0.5F, even if the cupric concentration is also 0.5F. By changing the indicator potential to +850 mv. vs. S.C.E., however, a satisfactory titration can be carried out in the presence of 0.5F Fe(II1). The increased amount of copper i3 necessary in this case; response of the indicator current is slow when 0.05F cupric ion is used, even a t the higher indicator electrode potential. There ought to be, of course, an optimum indicator electrode potential for each iron concentration; the value recommended previously (+650 mv. vs.
e0
-4
-2
T I M E I N SECONDS ( Z E R O I S EOUIVALENCE POINT)
Figure 1. Variation of indicator current during titrations at selected indicator electrode potentials.
S.C.E.) is satisfactory for solutions whose ferric ion concentration is less than 0.01F. The addition of any strong complexing agents which will alter the potential of the ferrous-ferric couple will also necessitate a change of indicator electrode potential. Attempts to catalyze the bromineiron(I1) reaction by addition of copper(I1) were unsuccessful. The reaction rate was not sufficiently accelerated to permit stable end points. Lower Limit of Titration. An a b tempt was made to titrate 50-pg. quantities of iron with chlorine, but these titrations gave consistently positive errors of about 1%. To extend this method to smaller amounts of iron, a study is currently being made of the factors influencing the current efficiency in the generation of chlorine as a coulometric intermediate. The method as reported here is accurate to *O.l% only for samples of 500 pg. or more of iron. Interferences. Because it was thought that this titration might be
No. of Detns. 8 5 10 4 2 2 2 11 4 4 4 4
useful for the determination of iron in ores and alloys, the effects of some elements which might interfere were investigated. Many substances not found in these alloys will reduce chlorine and would, of course, interfere in the titration of iron. Solutions of a known amount of iron in 2F HC1 with added manganese, chromium, or titanium were reduced with a silver reductor (16, 16), and the subsequent titration of the ferrous iron showed no abnormality. If vanadium or molybdenum were present, the titration was slow and no stable end point was reached. Solutions of iron and copper, when passed through a silver reductor, always gave results smaller than the sum of iron and copper. For this reason, a solution of copper nitrate was prepared and standardized by weighing the electroplated metal. Aliquots of this solution were reduced to cuprous ion by either silver or bismuth reductors (4). Although Birnbaum and Edmonds (1) obtained ~ 0 . 1 %accuracy on large samples of copper, our results were low when the silver reductor was used. Even though the bismuth reductor had a high blank (after extensive washing 2 ml. of wash liquid would introduce about lo-' equivalent of reducing material into the cell), copper was quantitatively reduced by it and recovered from it. Chromium and titanium solutions without added iron were passed through a zinc reductor and reduced to the plus 2 and plus 3 states, respectively. These ions appeared to reduce an equivalent quantity of the ferric ion originally present in the titration cell, and the resulting ferrous iron could be titrated easily. Further experimentation would be necessary to confirm this as a method for either chromium or titanium. Cooke et al. (2) used ceric ion as an intermediate in titrating small quantities of iron, but they did not have chloride present in their solutions. To determine if chloride would interfere, a solution 2F in hydrochloric acid and saturated with cerous sulfate was made up and several samples of ferrous iron titrated in it. The reaction was slow a t the end point, indicating that it
Table 1. Results of Titrations F e W hg.1 R~~~~ % E~~~~ Taken Found (av.) (pg.) of Av. 609.3 609.3 609.3 610.4 609.3 609.3 609.3 609.3 609.3 609.3 609.3 509.3
609.3 609.3 609.3 610.4 609.9 608.8 608.2 609.3 608.2 609.3 611.6 616.6
2.8 5.6 3.4 1.1 1.7 0.6 0.6 3.9
a.2
b.6 3.4 5.0
0.0 0.0 0.0 0.0 +0.1 -0.1 -0.2 0.0 -0.2 0 .o +0.4 +1.3
Concentration, F HC1 Cu(I1) 2
2 2 2 2 2 2
3 4 6 8 10
VOL. 33, NO. 10, SEPTEMBER 1961
0.01 0.02 0.03 0.03 0.06 0.15 0.5 0.03 0.03 0.03 0.03 0.03
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was proceeding through the chlorine intermediate as one would predict. The addition of cupric ion (final concentration, 0.03F) caused the reaction to speed up satisfactorily; no adverse effects of the cerous salt were noticed. Thus large amounts of chloride interfere in the ceric titration of iron, but one can easily use chlorine as titrant in the presence of cerous if the requisite amount of copper is present as a catalyst. If the sample originally contains a large amount of chloride, then there would be no point in adding cerous ion to the solution a t all. Test Method Using Standard Samples. With the above-mentioned interferences in mind, four Bureau of Standards samples were chosen, and their iron content determined by this method. The silver reductor was used and a parallel blank run in all cases; the blank was negligible except in the case of the ferrosilicon. FLUORSPAR. Standard Sample No. 79. Composition: CaFt, 94.83%; CaCOs, 2.25%; SiOz, 1.88%; Zn, 0.35%; Pb, 0.23%; S, 0.13%; Fe208, 0.15%; MgO, 0.13%; AlzOa, 0.02%; BaO, 0.07%; NatO, 0.06%; K20, 0.01%; PtOr, 0.005%; TiOt; 0.003%; MnO, 0.003%. Method: Two 0.7-gram samples were heated with nitric acid and perchloric acid to fumes of perchloric acid, cooled, and washed through the reductor into the cell with 2F hydrochloric acid. The ferrous iron was titrated and ferric oxide calculated as 0.149 i 0.003’%0. MAQNETITEORE. Standard Sample No. 29a. Composition: SiOz, 2.86%; AlpOs, 0.46%; TiOz, 0.15%; Fe, 69.54%; CrrOg,0.002%; VzOg, 0.002%; MnO, 0.03%; CaO, 0.096%; MgO, 0.095% : P105, 0 007%.
iMethod: Four 0.3-grain samples were dissolved in 4F hydrochloric acid, the silica was filtered off and fused with sodium carbonate. The melt was dissolved in 4F hydrochloric acid, added to the original solution, and the whole made up to 100 ml; 0.5-ml. aliquots were reduced and titrated, giving 69.48 =t 0.03% iron. Two other samples were titrated without recovering the iron in the silica residue; these gave 69.3 i 0.1% iron. These determinations (without the fusion) were made in 20 minutes, including weighing the samples and calculating the results. NICKEL-CHROMIUM CASTINGALLOY. Standard Sample No. 161. Composition: Ni, 64.3%; Cr, 16.9%; Fe, 15.0%; Si, 1.56%; Mn, 129%; Co, 0.47%; C, 0.34%; P, 0.012%; S, 0.005%; Cu, 0.04%; V, 0.03%; Mo, 0.005%. Method: Four 1-gram samples were fumed with perchloric acid, 2 ml. of hydrofluoric acid added, and the samples fumed again. After cooling, the samples were transferred to 100-ml. flasks and made up to the mark with 2F hydrochloric acid; 0.5-ml. aliquots were reduced and titrated, the copper and vanadium subtracted, and iron computed as 14.93 f 0.03%. FERROSILICON, 50% GRADE. Standard Sample No. 59. Composition: Si, 50.0%; Fe, 48.4%; Al, 0.93%; Mn, 0.31%; Ni, 0.125%; Ti, 0.105%; Cu, 0.10%; C, 0.015%; P, 0.035%; S, 0.008%; Cr, 0.08%; V, 0.004%; Zr, 0.01%; Ca, 0.04%. Method : Three 0.3-gram samples were heated with perchloric acid and hydrofluoric acid until dissolved ; the excess hydrofluoric acid was fumed off, the samples transferred to 100ml. flasks and made up to the mark with
2F hydrochloric acid; 0.5-ml. aliquots were titrated after being reduced. Possibly because of the heavy etching of the flasks, the blank time (1.1 seconds) was high, and this time was subtracted from the titrating time (about 120 seconds) to get the sum of iron and copper. The 0.10% copper was subtracted, leaving 48.4 f 0.1% iron. LITERATURE CITED
(1) Birnbaum, N.,
Edmonds, S. M., IND. ENG.CHEM.,ANAL. ED. 12, 155
(1940). (2) Cooke, R. D., billey, C. N., Furman, N. H., ANAL.CHEM.23,1662 (1951). (3) Dilts, R. V., Furman, N. H., Zbid., 27, 1596 (1955). (4) Dunham, J. hl., Ph.D. thesis, University of California at Los Angeles, 1956. (5) Farrington, P. S., Dunham, J. M., Ramsev. W. J.. Division of Analvtical Chemistry, 132nd Meeting, ACS,“New York, N. Y., 1957.
Farrington, P. S., Swift, E. H., ANAL. CHEM.22,889 (1950). (7) Furman, N. H., Cooke, W. D., Reilley, C. N., Ibid., 23,945 (1951). f8) Horn. H.. Ph.D. thesis. Northwestern University,’1954. (9) Meier, D. J., Myers, R. J., Swift, (6)
j
,
E. H., J. Am. Chem. SOC.71, 2340 (1949) (10) Myers, R. J., Swift, E. H., Zbid., 70, 1047 (1948). (11) Phillips, J. P., “Automatic Titrations,” p. 201, Academic Press, New York, 1959. (12) Takahashi, T., Sakurai, H., Japan Analyst 7,636 (1958). (13) Takahashi, T., Sakurai, H., Talanta 5.205 f 1960). (14) TutundZik, P. S., MladenoviC, S., Anal. Chim. Acta 12,390 (1955). (15) Walden, G. H., Jr., Hammett, L. P., Edmonds, S. M., J . Am. Chem. Soc.
56,350 (1934). (16) Wells, I. C., ANAL. CHEM.23, 511 (1951).
RECEIVED for review March 6, 1961.
Accepted May 22, 1961.
Stepwise Reactions in Chronopotentiometry A. C. TESTA and W. H. REINMUTH Department o f Chemistry, Columbia Univerdty, New York, N. Y.
b A theoretical discussion is given of the chronopotentiometric behavior of systems in which the reaction mechanism consists of three successive firstorder steps, each of which may b e chemical or charge transfer. In many cases the theory reduces to that for simpler two- or one-step mechanisms. Rigorous equations are given for some cases in which this is not so. Methods for distinguishing between possible mechanisms are discussed with particular emphasis on the difficulty of assigning unambiguously the location of slow chemical steps.
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ANALYTICAL CHEMISTRY
C
has been applied by many workers- to the qualitative and quantitative elucidation of electrochemical reaction mechanisms. This work has been reviewed recently (11). In the simplest cases only a single (charge-transfer) step is involved. Twostep mechanisms having one chemical step and one charge transfer were considered originally by Rosebrugh and Miller (19) for first-order chemical steps and by others (7-9) for higher order kinetics. Berzins and Delahay (1) studied cases of two-step mechanisms in which both steps were charge transfers. HRONOPOTENTIOMETRY
Recently it was shown that the reduction of o-nitrophenol takes place by a three-step mechanism in which a chemical step is interposed between two charge transfers (17). The relation describing the behavior of transition time as a function of current was given without derivation for that case. The aim of the present work is a more detailed and general theoretical consideration of three-step reaction mechanisms. For brevity, in the following discussion we designate three-step reaction schemes by three letters signifying the
