a i t h decrease of pH. When these values ‘we plotted against pH, a straight line is ,,btained with a slope of -0.0769 volt per pH unit. With these daba it could be dediiced that the optimum potential is related to p H according to the quation E (us. S.C.E.) = 1.554 0.0769 pH. Using this equation the optimum potential at any p H in the icid range could be calculated. These re in good agreement. with experimental values. Figure 1 gives the current-time plot foi. a typical run. The current de1:reases smoothly with time and reaches R negligible value in 30 minutes to 1
hour depending on the amount of ascorbic acid present. This indicates nearly 100% current efficiency of the coulometric oxidation process. The initial current in each case is found to be proportional t o the initial concentration of ascorbic acid. At pH 4.37, while the curlwit decreased to a low value, the fall of current with time was lees smooth than a t higher p H d u e s i)ossibly because of dissolution of hydrogen in the body of the solution (6). The separation of anode from the cathode by the use of a porous partition, however, helped to overcome this effect and to yield a smooth current-time curve.
LITERATURE CITED
( 1 ) Bogdanova, V. A., Gigiena i Sanit. 13, No. 10, 31 (1948); Chem. Abstr. 43,
5439 (1949). (2) Gillam, \V. S., IND.ENG.CHEM., ANAL. ED. 17,217 (1945). (3) Jewski, Jedize Mi., Chem. Anal. (Warsaw) 2, 453 (1957‘). (4) Kontio, Pekka, Casagrande, V. M., Suomen Keinistzlehti 18B, 9-10 (1945); Chem. Abstr. 40. 6532 (1946). (5) Lingane, J. ’ J., “Electroanalytical Chemistry,’’ p. 212, Interscience, New York, 1953. (6) Meyers, It. J., Swift, E. H., J . A m . Chem. SOC.70, 1047 (1948). RECEIVEDfor review March 6, 1961 Accepted June 8, 1961.
Stepwise Titration of Iron(lll) and Iron(l1) with EDTA and Ferricyanide The Square W a v e Titrimetric End Point LARRY C. HALL and DAVID A. FLANIGAN Department of Chemistry, Vanderbilt University, Nashville, Tenn.
b Mixtures of Fe(lll) and Fe(ll) in the ratios of 1:1, 5 1 , and 1:5 and varying from 0.001 25 to 0.00625M were titrated, giving a standard deviation of 0.04 to 0.10 ml. and a relative accuracy of 0.94 to 0.12%. Iron(ll1) was first titrated with EDTA at pH 4.7 to 5.0. Excess EDTA was added, the pH was raised to 6.7 to 7.2, and Fe(ll)-EDTA was titrated with ferricyanide. The slope of current-voltage polarization curves a t the null potential was used for end point detection. The square wave technique was used to measure the slope.
B
01” uoii(II1)
and iron(I1) can be determined in the presence of each other in a number of ways. The permanganate titration of Fe(I1) followed by reduction of the total Fe(III), either with standard Ti(II1) ( 0 ) or with a conventional reductor and subsequent reoxidation with a standard oxidant j5), are examples of redox techniques. Thiocyanate can be used for a colorimetric determination of Fe(II1) (4). 2,2’-Bipyridine and dimethylglyoxime (4)will yield a deep red complex with Fe(I1) in the presence of Fe(II1). Many excellent colorimetric agents are available for either (Fe(II1) or Fe(I1) ( 2 , 3, f4). Once the two oxidation states have been separated, their determination is straightforward. Solvent extraction separations are well known for Fe(lI1) (acetylacetone,
8-quinolinol, cupferron, ether-HC1, thiocyanate) and for Fe(I1) (4,7-diphenyl-l10 - phenanthroline, dimethylglyoxime) (IO). However, these methods are time consuming and somewhat complicated. Both Fe(II1) and Fe(I1) form EDTA comylexes stable enough to be titrated directly. The stability constants, log K p e Y - i = 25.1, log K ~ ~ y -= 2 14.3, indicate that at a low p H Fe(lI1) can be titrated n i t h no interference from Fe(I1) and, by subsequently raising the pH, Fe(I1) can be titrated. Koros and Barcza (6) used this principle and determined Fe(II1) at p H 1.8 to 2.2 using sulfosalicylic acid as an indicator. They next added ascorbic acid, adjusted the p H to 5.0 to 6.5, and titrated Fe(I1) using methylthymol blue as the indicator. The role of ascorbic acid is somewhat obscure since it will not reduce the Fe(II1)-EDTA complex. Since Fe(II1) sloidy forms a stable methylthymol blue complex, it was suggested that ascorbic acid reduced any Fe(II1) produced from the Fe(II1)-EDTA equilibrium and thus prevented blocking of the indicator. Lucchesi and Hirn (8) determined the sum of Fe(II1) and Fe(I1) by adding excess itnndard EDTA a t pH 2.0, raising the p H to 9.0, and back-titrating with standard zinc. Eriochrome Black T was used as the indicator, but the solution had to be a t least 50y0ethyl alcohol to prevent formation of the more stable Fe(II1)Erio T complex.
In view of the blocking action of Fe(II1) on the indicators used for the above methods, i t seemed desirable to use an electrometric end point. Mercury electrodes are generally precluded because Fe(II1) oxidizes mercury in most ionic solutions. The method presented here depends upon the measurement of the slope of the currentvoltage polarization curves of the Fe(II1)-Fe(I1) system at the null potential using a pair of platinum microelectrodes. Square wave titrimetry ( 7 ) achieves this measurement and is essentially the alternating current analog of the Foulk-Bawden “dead stop” and the Willard-Fenwick potentiometric methods. Ferric iron was titrated first with E D T A at a p H of 4.7 to 5.0. After this end point, excess EDTA was added, the p H raised to 6.6 to 7.2, and the Fe(I1)-EDTA titrated with ferricyanide. EXPERIMENTAL
Apparatus. T h e apparatus and theory of square wave titrimetry have been given previously (7). Figure 1A shows the schematic arrangement. The Heath Co. (Benton Harbor, Mich.) SQ-1 square wave generator was modified by replacing the 2000-ohm attenuating potentiometer with the arrangement shown in Figure 1B to provide output voltages of 30 volts to 1 mv. Switch SI was provided as a means of conveniently reading iR1 or Eapplled without disconnecting the electrode leads. When the switch is up and the VOL. 33, NO. 11, OCTOBER 1961
1495
Figure
1.
Schematic Arrangement
of Apparatus
R,I 0 to 10,000-ohm decode resistance 0 to 20-megohm pot. Ra, 1800 ohms h, 2000 ohms PI, 2000-ohm pot. P2, 250-ohm pot. Pa, 1 0-ohm pot. Meas. device, 0.c. voltmeter or oscilloscope SI,Switching box S2, 2-System, 5-position, rotating, nonshorting switch R2,
4 s,
fleliible lead plugged in A, the measuring device can read tR1. With the switch down and the flexible lead plugged in B, EaD,,i, can be read. The platinum electrodes were constructed from 26-gage wire and sealed in soft glass, care being taken to anneal the wire-glass interface. The tips were fused into a small ball. The net length of wire exposed to the solution was 8 mm. Each electrode was cathodized in 0.1V perchloric acid to the point of
Table 1.
Titration
hydrogen evolution for 1 to 3 minutes before a titration. The titration cell was a 500-ml. tall form beaker with a side arm through which nitrogen was introduced over the top of the solution. -4 properly drilled rubber stopper contained the platinum electrodes, buret tip, and a glass electrode-calomel pair. -4 magnetic stirrer was used for all titrations. Current-voltage polarization curves were obtained with the Sargent Model
of Fe(lll) and Fe(ll) in Different Ratios
EDTA, M1.O
Ferricyanide, MI.*
25.23 25.26 25.31 25.30 25.20 25.36 25.20 25.27 25.24 0.06 0. 12y0
23.91 23.94 23.82 23.88 23.87 23.63 23.85 23.84 23.88 0.10 0.17%
Initial PH
Second PH
Buffer System
7.0 7.0
G1y.-ac. G1y.-ac.
1:1 Ratio
Av. Ca1cd.c Std. dev. Rel. acc.
Av.
Ca1cd.e Std. dev. Rel. acc.
24.95 24.71 24.87 24.82 24.95 24.85 24.93 24.88 24.96
4.67 4.69 4.74
0.09
0.06
0.3270
5.0 5.0 .5 .0
6 8
5.0 4.7 5.0
6 6 6.9 7.0 7.2
3 .o
iic .
Gly. G1y.-ac. G1y.-ac.
Gly.-ac. PROCEDURE
5:l Ratio 4.8 5.0 5.0
7.1 6.9 7.1 7.3 7.1 7.0 7.2
G1y.-ac. G1y.-ac. AC. Gly. Gly .-ac. G1y.-ac. G1y.-ac.
7.4 7.3 7.4 7.3 7.4 7.2 7.2
Gly-ac. G1y.-ac.
4.73 4.77
0.8470 1:5 Ratio
5.31 5.36 5.34 5.41 5.30 5.40 5.35 5.35 5.30
23.76 23.62 23.81 23.86 23.86 23.79 23.77 23.78 23.88
Av. Calcd.0 Std. dev. 0.04 0.08 Rel. acc. 0.9470 0.4270 EDTA = 0.1001M. * Ferricyanide = 0.1048M. 0 Fe(II1) = 0.1010M; Fe(I1) = 0.1002M.
1496
ANALYTICAL CHEMISTRY
XXI polarograph using a rotated platinum electrode a t 600 r.p.m. LS. the saturated calomel electrode. Solutions and Chemicals -1stock solution of F e ( I I ) , prepared from ferrous a m m o n i i m sulfate (Tho111 Smith, 98.87% assay), was standardized with Ce(1T’) sulfate (G. Frederick Smith Chemical Co., reagent grade, 0.1019X) which, in turn, had been checked against primary standard grade arsenious oxide (Thorn Smith, Thi, Fe(I1) 99.99Cc w a y , 0.10291;). m-as stoied under nitrogen and its strength checked at the beginning of each set of titrations. -1 standard stock solution of Fe(TI1) \vas prepared from iron nire (RIerck, 99.8$;, 0.09972 M). A solution of EDTA (J. T. Baker, 0.1001M) as prepared determinately according to the procedure of l3laedel and Kniiht (I). Pbtassiuni ferricyanide (J. T. Ruker, 0.1048.1/1 was prepared from the rccrystallized and dried salt and further checked iodonietriually with standard .is(III) ( 5 ) .
5.0 5.0 5.0
5.0
5.0 5.0 5.0
Glp.-ac.
Gly.,-ac. G1y.-ac. Gly.-ac. G1y.-ac.
lliutures of Fe(II1) to Fe(1I) in the ratios of 1:1, 5:1, 1:s were titrated. Either 25 or 5 nil. of each mere pipetted into the tall form beaker and diluted The concentrations of to 400 ml. Fe(II1) or Fe(I1) ranged from 0.00625 to O.O0125J/, depending upon the particular ratios studied. 9 stream of nitrogen flowed over the surface of this solution at all times. Concentrated hydrochloric acid, 5 ml., was added, followed by enough glycine to raise the p H to 4.7 to 5.0. Alternatively, sodium acetate can be added in place of glycine. Also, glycine can be added to pH 3.0 followed by acetate to p H 5.0. If the solutions a t pH 5.0 or 7.0 (see below) turned from yellow to amber brown, the p H was loirered gradually by adding hydrochloric acid until a clear yellowish solution was obtained. The square wave generator output was adjusted so that the applied voltage to the platinum electrodes was 30 mv. A frequency of 20 c.p.s. was used, R1 was set at 500 ohms and R, a t 0 ohm. After each incremental addition of EDTA, the SQ-1 output nas adjusted so that EsPp,,was 30 mv. before iR1was read. After the Fe(II1) end point, enough excess EDTA was added to
complex all of the Fe(I1). The Fe(I1) end point was indicated when the sharp rise in iR1began to level off and become constant. After this point, 10 ml. of EDTA provided a sufficient excess. The p H was raised to 6.6 to 7 . 2 by adding sodium hydroxide, and titration was continued with ferricyanide. Both end points are V-shaped. DISCUSSION
2A, Fe (Ill)
28, Fe(II) 32 9
I
76-
-
28 24
-
5-
20
5:
Solution Conditions. I n the absence of foreign complexing agents, the minimal p H for the titration of 0.01M Fe(II1) or Fe(I1) to within 0.1% relative accuracy is about 1.8 and 5.2, respectively (11, 12). When, in addition to pH, the side effect of other complexing agents such as those furnished by buffer materials is taken into account, the conditional stability constants of the Fe(II1) and Fe(I1)EDTA complexes are lowered. This trend can be offset by reducing the p H side effect, Le., by raising the pH of the titration solution. However, the formation of hydroxyl mixed ligand species and the beginning of polynuclear complex formation impose a restriction on how high the p H can be adjusted. There is also the obvious severe limitation of iron hydroxide precipitation before the E D T A complexation begins. Ringbom ( I S ) has shown that the EDTA stability constants of Fe(II1) and Fe(,II) reach their maximum values a t p H 5 and 10, respectively, with the side effects of pH, mixed hydrosyl and hydrogen ligand species, and polynuclear complexes taken into account. If the stability constants of the iron glycinate and acetate complesrs were knljwn accurately, the essct pH conditions could be cniculat,ed for optimum t,itrations. Hew i s v ni t,he lack or' this inforniation, one can only specify the rangts of pH to be 1.8 to 5.0 for Fe(1IIj and 5.2 t,o I(! for Fe(I[,!, For thc. results listed in ' ~ a h kI! titration was cwried out just below che p H of the point of incipient (lariit.tiing of the sulutioiw. Khtr:her g1yc.int. and x & t e :trc used d o n e or together had no rffect upon the results. Srpnrste rsperimentr shon c>Li t h n c 0.006253/ Fe(II1) wuid he kept in sollition for periods m e r 3 hours a t R pII of 1.7 to 5.0 using glycine or acetate buffers. Similar results were found for FctIII) and Fe(II)-EI)'TA solutions a t a !)I1 of 6 . i to 7.2. End Point Detection. T h e rnti points are established by the variou? Fc(lI1)-Fe(I1) couples. The buffer. 1.; [)'I'A, and cyanide systems n x :3hhrevixted as F ~ . ~ . I I I ~ R - ~ F P I I T I,'e!R, (IIIjk--Fe(II)Y, arid ferricyanideferrocyanide. Initially, Fe! I [ I iB.Fi:ilI)B exists as the potent'ial :ictcrrnin!ng couple. As Fe(t1IIR 1s 'r)nipiexes: w t h fi:D'r,\* tiit. T I I I , ~ 1 o : n t 3ioi:e continually decreases < i h i d,
E 4-
-
165
[L-
[L
3-
12-
2-
8 I
I-
4
'b I
0-
I
I
I
I
I
Figure 2) until all Fe(I1I)B has been converted to Fe(II1)Y. A t this point (see b in Figure 2) there is no reversible potential determining couple in solution and the slope is a t a minimum. When excess E D T A is added, enough Fe(I1)B is complexed so that the second reversible potential determining couple, Fe(1II)Y-Fe(II)Y, is established and the slope wdl increase sharply (line c, Figure a), thereby causing a V-shaped end point. If tJhe Fe(I1)EDTA complex were stable enough at the p H of 4.7 to 5.0, the null point dope would reach a maximum and then level off giving .z ,;-'-shaped second end point. A t pH 4.7 to 5.0 or 6.7 to 7 . 2 this end point was unreliable. The ferricyanide-ferrocyanide system (E" = +N45 wit, pfI 7.0) can oxidize Fe(II)-ErJTA P O f7e !:I) -Er)'r2i(E' t11.136 volt). is fr,iricyanide is added, Fe(1IjY is renioveci d i t s dope decreases (line ti, Figure 21 until a niinimum is reached (point e'i, a t which time only Fel1II)Y and ferrocyanide exist. .After ~ x u e s s ferricyanide is Added, the third reversible potential deterniining coupie is ?staiJiished, Le., ferrii~SLinide-Fei.rocysriifi~.3rLrl the null y i n t slorr rapidly i:icrwses again. The use of ferric:yanitIe for the FeiII) *ic:teriiiination gave much sharper end points arid higher :tccuracy than E D T A :ind the ;--. sh:cptd end point. Currentmitage pnlarization curves were obfatineti iising n. rotating platiniirti e1t.i.crode for sdrit,ions corrvspoiiding to ijtiints a t,o f in Figure 2 . The curves , onfirmed the changes in null point ~ l o p r s f m i t i r i experinientn dg with the q u a r e W R V Q wchnique. T Lhe ! . [ , t i .'clints art' ; r . t + z o f inter2y
I
I
I
ferences from other metal ions if these do not form interfering potential determining systems or if they do not reduce Fe(III)B, Fe(III)Y, and ferncyanide or olddize Fe(II)B, Fe(II)Y, and ferrocyanide. The presence of ions such as Ba(II), Ca(II), Mg(II), Sr(II), Zn(II), Ni(II), .4s(III), Pb(II), rare earths (111), etc., would need only be corrected for the amount of E D T A required to complex them. Other electrometnc methods can be iised to detect the end points besides the one described. It is the experience qf the authors that either conventional amperometry or direct and alternating current null point slope measurements give sharper end points at lower conventrations than rmventional potenticmetric titrations. Results. The stock solutions 11; Fe(I1) made from ferrous amn1oni:int sulfate contained some Fe(II1). In ittldition, air oxidat'ion gpaduaily b u l t more Fe(II1). Therefore, the solutions of Fe(I1) nil1 make the F e l I I I ) end points appear to be high unless them factors are recognized. ?'he molarity of Fe(II1) used in Tablc I to determine the theoretical end points wis based lipon the concentration x3f the FeiIIli stock solution, the purity < r E XIohr's s i l t , +,he ;hanging titer cf Fe(II), ant1 the mi; of thc i"t.!I[i I iiqunt .
YO1
:: r 4 0
I
,
OCTOBER 1961
1497
LITERATURE CITED
(1) Blaedel,
W.J., Knight, H . T., .4NAL. CHEM.26,741 (1954). (2) Diehl, H., Buchanan, E. B., Smith, G. F., Ibicl., 32, 1117 (1960). (3) Diel4 H., Smith, G. F., “The Iron Reagents,” G. Frederick Smith Chemical Co., Columbus, Ohio, 1960. (4) Feigl, F., “Spot Tests, I, Inorganic Applications,” 4th ed., pp. 153-8, Elsevier, New York, 1954. ( 5 ) Kolthoff, I. M., Belcher, R.: “Volumetric Analyses,’’ Vol. 111, Interscience, New York, 1957.
( 6 ) Iioros, E., Barcza, L., Chemist Analyst 48,69 (1959). (7) Laitinen, 11. A., Hall, L. C., ANAL.
CHEX29,1390 (1957).
(8) Lucchc.si, C. .4,> Hirn, C. F., Ibid., 32, 1191 (1960). (9) Lundell, G. F. F., Bright, A. A., Hoffman, I. F., “Applied Inorganic .4nalyses,” 2nd ed., Chap. 21, John Wileg BE. Sons, New York, 1953. (10) Morrison, G. H., Freiser, H., “Solvent Extractions in Analytical Chemistry,”
John \%ley K: Sons, New York, 1957. (11) Reilley, C. N., Schmid, R. W., -\SAL. CHEM. 30,947 ( 1 958).
(12) Ringborn, >i., J . Chem. Ecluc. 35, 282 (1958). (13) Ringborn, A., “Treatise on Analytical Chemistry, Theory and Practice,” Part I, Vol. I, I. M. Kolthoff, P. J. Elving, Eds., Chap. 14, Interscience, New York, 1959. (14).Sandell, E. B., “Colorimetric Determlnation of Traces of Metals, 3rd ed., Chap. 22, Interscience, Ken- York, 1959. RECEIVEDfor review l l n y 4, 1961. Accepted July 13, 1961. Southeastern Regional Meeting, ACS, Birmingham, Ala., Kovember 1960.
Chronopotentiometry of Iron(ll) and Iron(III) Adsorbed on Platinum Electrodes FRED C. ANSON California lnsfifufe of Technology, Pasadena, Calif.
b In sulfuric and perchloric acid sobtions Fe(ll) and Fe(lll) are adsorbed on the surface of platinum electrodes. The adsorption is strong enough so that the electrodes can b e washed free of any unadsorbed iron without removing the adsorbed iron. Chronopotentiograms recorded with such electrodes demonstrate the presence of the adsorbed iron. The chronopotentiograms obey the potential-time relation to be expected for adsorbed reactants. Adsorption isotherms for reactants adsorbed on platinum electrodes can b e determined with the aid of chronopotentiometry. The isotherms allow an estimate to be made of the effect of adsorption of reactants on the chronopotentiometric constant.
A
OF ~ N ~ T A N C E have S been reported by Lorenz and coworkers in which a substance reacting a t an electrode is adsorbed substantially (IW14). It has been suggested that such reactant adsorption may play an important role in the kinetics of electrode reactions (4, 6-9, 15, IO), so that the nature and properties of adsorbed reactants are of general interest. Previous studies of adsorbed reactants have usually been indirect because of the necessity of separating effects due to the adsorbed species from the (generally predominating) effects due to the unadsorbed reactant in the bulk of the solution (9, 12-14). Recently the ovidation and reduction of reactants adsorbed on platinum electrodes were studied directly by chronopotentiometry (5’). This report presents the results of further evperiments with Fe(I1) and Fe(II1) adsorbed on platinum dectrodeq. SUMBER
1498
ANALYTICAL CHEMISTRY
EXPERIMENTAL
Apparatus. T h e chronopotentiometric circuitry followed standard practice (IO). A Moseley autograph recorder (Model 3 3 ) was used to record the ehronopotentiograms. The potential difference between the working and reference electrodes was supplied to the input of a high-impedance follower amplifier having a gain of unity The amplifier was constructed from plug-in analog computer amplifiers following a circuit of DeFord ( 5 ) . The output of the amplifier was supplied to the input of the recorder. The working electrode was a piece of 0.030inch platinum wire sealed in soft glass. The area of the electrode was 0.15 sq. em. ‘The auxiliary electrode was a 100-sq em. cylindrical platinum gauze electrode The working electrode was positioned in the center of the cylindrical auxiliary electrode to ensure a uniform current density. Electrode geometries that provided less uniform current densities resulted in chronopotentiograms with distorted shapes. Potentials were measured with respect to a calomel electrode saturated with sodium chloride, but are reported us. the usual potassium chloride-saturated calomel electrode. It is necessary to pretreat electrodes prepared from commercially available high-purity platinum wire in order to obtain significant adsorption. The procedure consists of very light platinizing of the electrode by making i t the cathode in a 0.2M chloroplatinate solution and passing 0.1 to 1 coulomb per sq. em. of electricity through the electrode. The rewlting electrode surface should not be visibly darkened by the deposited platinum The potential of a properly prepared electrode can be forced from f0.6 to +0.1 volt us. S.C.E. in deaerated 1F sulfuric acid in less than 0.5 second with a constant current of 100 ,ua. per sq. cm. In
cases where the electrode response is not this sharp because of excessive platinization, the response can be improved by partially dissolving the platinized platinum in aqua regia. Procedure. T h e procedure for adsorption experiments consisted of immersing t h e electrode in solutions of Fe(I1) and/or Fe(II1) for measured time intervals, removing the electrode and rapidly washing i t thoroughly with distilled water, placing i t in an iron-free solution of deaerated perchloric or sulfuric acid, and recording chronopotentiograms corresponding to the adsorbed Fe(I1) and Fe(II1). Reagents. Solutions of Fe(I1) and Fe(II1) perchlorate were prepared and standardized as described previously ( I ) . Solutions were prepared from triply distilled water and measurements were performed in oxygen-free solutions. The temperature of all solutions remained within 2’ of 25’ C. RESULTS A N D DISCUSSION
A platinum electrode that is exposed to solutions of Fe(I1) and Fe(II1) adsorbs both oxidation stafes of iron on its surface. The electrode may he removed from the iron solutions :ind washed free of unadsorbed iron without removing the adsorbed iron. The washed electrode may then he placed in an iron-free perchloric or sulfuric acid solution and chronopotentiograms corresponding to the absorbed iron obtained. That the ehronopotentiograins actually correspond to adsorbed iron is readily demonstrated. For one thing the transition times arc independent of whether or not the solution is stirred during the recording of chronopotentiograms. I n addition, if the direction of the chronopotentiometric current is reversed a t the t,ransi-
