ANALYTICAL CHEMISTRY
520
LITERATURE CITED (1)
Alber, H. K.. Harand. J.. IND.ENO.CKEM..ANAL. Eo. 10, 403 (1938).
(2) Alber, H. K., Harand. J.. J . Fmnklinlnst. 224,729 (1937). (3)
(4) (5)
($1 (7) (8) (9)
(10) (11) (12) (12)
Fipire i.
Section of f i . .
lators can alnays he made with the minimum of difficulty. Figure 7 shows B part of the finished table, ahich ncrommodates nine balances Lighting. Individual fluorescent fixtures of the type described previously (83, $4) are used for illumination. These are eontrolled through the snitch belou- the table t,op (Figures 5 and 7).
(18) (19) (20)
415 (1943). (21)
ACKNOWLEDGMENT Thc authors are indebted to James Thomson for his assistance in preparing Figure 4; to Patricia Hillman for her assistance in preparing Figures 4 and 5 ; to Carmine Aurirchia for the preparation of the photographs; to Jack Harris of thc Korfund Co., Ino., for the oacillograms and analysis of t,hPm; and to Hane Frey and the Rlettler Instrument C o p for supplying the Rlettler balance used in bhe tests.
BoCtius, M., "Uber die Fehlerquellen bei der mikroanaiytisoheo Bestimmung des Kohlenstoffes und Wasserstoffes naoh der RSethode von Frits Pregl." p. 18. Veriaa Chemie. Berlin. 1931. Brush Development Co..Cleveland 14. Ohio. Teoh. Bull. 282. Emieh. F.. "Lehrbuch der hlikrorhemie." 2nd ed.. p. 76. J. F Bergman. Munioh. 1926. Feldman. J. R.. General Foods Corp.. Hoboken, N. J., private eonrmunieation. Fleiseher. K. D.. Sterling-W'inthrop Research Institute. Remselaer, N. Y., private communication. Gage. D. G.. Sullivan. P.. ANAL.CHEM..in press. Grant, J.. "Quantitative Organic nlicroandysis" (based on "Methods of Frits Pregl." 5th English ed.). p. 10, Blakirton Co.. Philadelphia. 1951. Gysel. H.. Strebel. W.. Mikmchim. Acto 1954. 782. Hallett. L. T.,IND.Eno. CKEM..ANAL.Eo. 14.956 (1942). Hnllett. L. T..1st Xletmpolitlrn Mioroohemioal Society Symposium. New York. Msroh 1946 Hnward, H. C...T.Ind. Eny. Chem. 13, 231 (1921). ner. CV. R.. IND.EN(I_CKEM.. ANAL.Eo.9 , 3 0 0 (1937). .fund Co.. Inc., 48-15 32nd Place. LOW Island City 1. 7 . Y., Chart SA-2028-0 and Drawing SC32-2. ck,J. A,. College of the City of New York. American Cyanalid Co.. Stamford, COM., private oommunieation. Ilwreith. C. G.. ANAL.CKBM. 23,688 (1951). Pregl. F.. "Quantitative OrK%dOMicrosnaiysis." p. 10: translated fmm 2nd revised snd enlarged German ed. by E. Fyieman, J. & A. Churchill. London. 1924. Rodden. C. J.. Atomic Energy Commission. New Brunswiek. N. J.. private communication. Rodden. C. J.. Kuek. J. A.. Benedetti-Piohler, A. A.. Corwin. A. H.. Huffman, E. W. D.. IND.ENO.CBEM., ANAL.Eo. 15.
(22) (23) (24) (25)
(26)
Roth. H.. "Quantitative organiaehe Mikmanalyse von Frits Pregl." p. 9, 5th ed.. Julius Springer. Vienna. 1947. Stehr. E., Texas Co.. Beacon. N. Y., private communication. Steyermark. A.. IND.ENO.CKEM..ANAL.Eo. 17, 523 (1945). Steyermark. A.. "Quantitative Organic Microanalysis." PP. 3-4. 20-27. Bialdston Co.. Philadelphia. 1951. Streeter. K. B.. Sharp and Dohme. h e . . West Point. Pa.. private communication. Waldmann, H.. Hoffmsnn-La Roche & Co., Ltd.. Bade. Switzerland private communication..
R ~ r e i v e otor review Ootoher 25, 1855.
Accepted February 6. 1956.
Coulometric Titration of Iron(lll) with Electrolytically Generated Iron(l1)-Ethylenediamine Tetraacetate R. W. SCHMID and CHARLES N. REILLEY Department of Chemirrry, University of North Carolina, Chapel Hill,
Conditions for coulometric t i t r a t i o n s w i t h electrolytically generated iron(lI)-ethylenediamin~ tetraacetate are described. A u t o m a t i c and manual titrat i o n s of iron(II1) gave satisfactory results. The redox s y s t e m imn(III/lI)-EDTA and iron(III/Il) was investigated potentiometrically in o d e r to select the optimum conditions.
s
CHWARZENBACH and Heller ( I f ) pointed out that the strikingly ' . strong reduction p o m r of mixtures of ferrous salts with (ethy1enedinitrilo)tetrsaretic acid (ethylenediaminetetraacetic acid, EDTA, or H,Vj may lead to practical applirations. They determined the redox potential over a pH range
N. C.
from 2.4 to 8.3, by means of the potentiometric titration curves of ferrous ethylenediamine tetraacetate with chlorine or bromine, and found it to have B constant d u e of +117.2 mv. (us. normal hydrogen eleclrodej from p H 4 to 6 . At higher and loa.er p H the redox potential depends strongly on pH (Figure 1). The reduction power is similar to titanous (E' = ea. 0.1 volt vs. normal hydrogen electrode) but spplicilhle over a different pH region. Kolthoff and Auerhaeh (6) investigated the same system polarographically in a pH range from 1 to 11 and ahtained half~ a v potentials e deviating hy only a few millivolt8 from the values reported hy Sehwarzenbach and Heller. In anslytical ehemist.ry, iron(1Ij-EDTA has hem used for det,erminations of iodine (7), silver (6j, and copper ($). Coulometric titrations are esperially valuahle when the tit,rsnt
+
V O L U M E 28, NO. 4, A P R I L 1 9 5 6
521
is an unstable substance. As iron(I1)-EDTA is very sensitive to air, it seemed desirable to work out condit,ions for the electrolytic generation of this reagent. Iron( 11)-EDTA w-aa generated cathodically at a platinum electrode from iron(II1)-EDTA by the reaction:
FeY-
+e
-+
Fey--
The resulting iron(I1)-EDTA titrant then reacts in the bulk of the solution n-ith the iron(II1) sample: Fey--
+ Fe+++
-P
Fey-
+ Fe+?
h more detailed study was made on the effect of pH on this titration reaction in order to find the optimnni p H condition.
6 ui
w
:I .3
.2
titratcd. Acetate ions complex ferric ions more tightly than ferrous ions, and therefore shift the redox potential of this system to more negative values, as shown in curve B, Figure 1. The curve represents the potential of an iron( 111)-iron(I1) solution containing excess acetate as a function of pH. At p H values greater than 3, basic acetates precipitate. Cheng, Bray, and Kurtz (3) found that a t p H values greater than 3 the recovery of iron as determined by titration with EDTA was no longer 100%. Even though a stock solution of ferric ion in the present work appeared clear at p H 4 (a weak Tyndall beam, however, existed), incomplete recovery was found. In summary, Figure 1 shows that in order to obtain a large potential break a t the end point (corresponding to the difference bct,ueen curves A and B in Figure 1) and to avoid the formation of basic precipitates, the titration is carried out best a t pH 2 t o 3. The iron(II1)-EDTA stock solution \%-asconsequently neutralized with sodiuni acetate to the desired pH of 2.5 (final concentration 0.131). Iron(I1)-EDTA is very easily osidized by air (4, 5 ) and therefore air must be excluded carefully, Iron(II1)-EDTA is decomposed by sunlight ( 4 ) and care has to be taken in this respect.
OF
-DATA
SCHWARZENBACH AND
EXPERIMENTAL
HELLER
Cell. Thc cell shoivn in Figure 2 was used. OUTLET
AND
GENERATING
I
I
,
,
,
2
3
4
5.
,
I
6
7
8
9
SAMPLE
OPENING
CAT"ODE
-
IO
PH Figure 1.
pH-poten tial diagram GENERATING ANODE
A . Iron(III/II)-ethylenedismine tetraacetate redox system B . Iron(III/II)iedox eysteni (in excess acetate)
For this purpose the redos potential of the iron(II1)-EDTAiron(I1)-EDTA has been measured potentiometrically as a function of pH, aa is shown in curve A , Figure 1. The values agree rvell with the points reported by Schwarzenbach and Heller ( 1 1 ) . The reduction potential of the iron(II1)-EDTA-iron(I1)-EDTA system becomes more negative in alkaline solution, but solutions more alkaline than p H 10 are unstable and slorvly forni precipitates of iron(II1) hydroxide (11). At p H values lower than 2 the redox potential changes rapidly in a positive direction (230 mv. per pH). Although the slope of this rise corresponds fairly closely to the participation of 4 protons in the potential-determining equilibrium, the process is actually a mixed one, in view of the two equilibria (IO):
+ +
HFeY -+ H + F e y H,Y * H + H,Y-
pK = 1.4 pK = 2.0
In order to obtain a strong reducing power it is desirable to work at pH values higher than 2. Therefore, partial neutralization of the free acid formed in the preparation of iron(II1)-ethylenediamine tetraacetate is necessary. Fe+++
+ HsY---+
Fey-
+ 2H+
Partial neutralization with sodium hydroxide, however, leads to an opalescence, attributed to the precipitation of iron(II1) hydrosides in the vicinity mherc the drops of base are added. This opalescence clears only very slowly after stirring, but could be entirely avoided by neutralizing the solution with sodium acetate. On the other hand, an equally important consideration is the effect of p H on the reduction potential of the substance to be
Ne
I
TUBE
FRITTED
DISCS
Ne
TUBE
DETAIL
[SIDE VIEW)
Figure 2.
CoulomeLric titration cell
A 2 X 1 cm. platinum foil was used as working cathode and a small platinum foil as an anode. The anode compartment, filled with 1Jf sodium sulfate solution, was separated from the main part of the cell by means of a 1.U sodiurn sulfate liquid junction and two fine glass frits. This wits necrssary in order to avoid diffusion of the oxygen-saturated anodic solution into the cathode compartment. A 1-cm. platinum wire served as an indicator electrode, and a silver-silver chloride electrode as the reference electrode. To minimize the internal resistance drop between the indicator and reference electrodes, especially for the automatic titrations, the electrodes in the indicator circuit were kept close together and away from the generator electrode circuit. Constant Current Supply. For small currents, a supply previously described (9) was used which furnished currents from 0.005 to 4 ma. Current measurements with an accurate potentiometer circuit (Leeds & Northrup student potentiometer) showed that current fluctuations did not exceed 0.1%. For the titration of the lnrger amounts of iron, a supply ( 8 ) which gave one defined curreut of larger magnitude (43.30 0.10 ma.) was used.
*
REAGENTS
Iron(II1)-EDTA Solution. To a solution of iron(II1) chloride (Merck) an exactly equivalent arnount of a disodium dihydrogen ethylenediamine tetraacetate solution was added and the resulting solution was partially neutralized to pII 2.5 by an equivalent amount of sodium acetate. The exact titer of the respective solutions was determined relative t o each other by electronietric
ANALYTICAL CHEMISTRY
522
Table I. Current, Ma.
Titration of Iron(II1) by Electrolytically Generated Iroii(I1)-EDTA Iron Taken,
Iron Found,
Mg.
Mg.
Deviation
Me. -0.08
0.00
t o . 10
43,30
3.862 3.864 3.880 2.886 2,002 2.001 2.002 2,002 1.997 1.997 1.995 1.993 1.1424 1.1420 0.9992 0.9992 0.9991 0,9991 0.5008 0.5022 0.5027 0.5022
7.01
2.63
1.402
1.402
0.701 0.701
0.280
7.04 7.09 7.02 7.02 2.64 2.67 2.67 2.04 1.404 1.398 1.404 1.406 1.406 1.400 1.405 1.406 0.700 0.703 0.689 0.697 0.709 0.694 0.271 0.272 0.725 0.275
-0.18 +0.05 -0.05 + O . 03 t0.08 f0.01
+0.01 $0.01 +0.04 so.04
+O.Ol +0.002
% -0.5 0.0 +0.6 -0.6 +0.3 -0.3 +1.1
to.1 +o. 1 +0.4 4-1.5 +l.5 +0.4
+0.004
-0.001 +0.002 -0.012 -0,004 +0.008
1-0.3 -1.7 -0.6 i l . 1
+ O . 002 + O , 004 +0.001 -0.002 +0.003
-0.007 -0.009 -0.008 -0.005 -0 005
0.01M iron(I1) ammonium sulfate 0.01M iron(II1) chloride 0 . 2 X sodium acetate
4-0.4
$0.1 -0.3 +o. 1 +0.3 +0.3 -0.1 +0.2 +0.3
-0.004
ethylenediamine tetraacetate, and 0.04111sodium acetate. The potential of this solution as a function of p H was measured with a platinum electrode, using a Leeds & Northrup student potentiometer. The p H was adjusted by addition of cqncentrated per chloric acid or sodium hydroxide and measured with a glass electrode. The procedure for the iron(II1)-iron(I1) system was analogous, the measured solution containing:
-0.1
-1.0
-3.3 -2.9 -1.8 -1.8
titrations. The solution was then diluted to give a final coiicentration of 0.1Hiron(II1)-EDTA. The solution should be allowed t o stand overnight before use, because its blank value changes in the beginning. Nitrogen. For bubbling oxygen from the solutions, oxygenfree nitrogen (Airco, Seaford grade) waa used, being bubbled first through miter to saturate the nitrogen with water vapor. Ferric Chloride Solutions. The titer of the iron(II1) chloride stock solution was determined by potentiometric titration with a standard EDTA solution ( 1 ) . The solutions used for the coulometric titrations were subsequently made up by dilution of this stock solution. DETERMINATION OF END POINT
The potential of the platinum indicator electrode was measured against a saturated calomel electrode or silver-silver chloride with a Leeds &. Northrup p H meter (Type 7664), which in turn was connected to a Brown recorder in order to register the titration curves automatically. A chart speed of 0.5 inch per minute was used and the current so chosen that the length of chart paper for a titration curve was 8 to 10 inches. \$’hen a current of 43 ma. was used, the titration curves were determined manually. For each point the current was temporarily interrupted, and the time (Grs Lab Micro timer, Type 202) and potential were noted. The potential in the vicinity of the end point required 5 to 10 seconds to reach equilibrium. PROCEDURE
The cell mas filled with 25 to 30 ml. of the iron(III)-EDTAi solution (0.05U for titrations a t any current up t o 4 ma., O . l d 1 for the current of 43 ma.), and deaerated 10 minutes with nitrogen. A pretitration was carried out next after addition of a small but unmeasured amount of the solution to be titrated. The sample was then added and titrated. The distance between these two end points was then measured and taken to represent the titer of the sample. I17iththe prescribed amount of iron(II1)EDTA background electrolyte four 5-ml. samples could be titrated one after the other before the cell was refilled with fresh background solution. During the titration the solution was vigorously stirred with a magnetic stirrer and a slorv flus of nitrogen bubbles was maintained. iUEASURER3ENT OF R E D O X POTENTIALS
A solution was prepared containing 0.01M iron(I1) ammonium sulfate, O.OliZ1 iron(II1) chloride, 0.02111 disodium dihgdrogen
RESULTS
The results obtained are listed in Table I. The outstanding sources of error were estimated to be the following: Automatic Titration,
%
Graphic determination of end point Reproducibility of pipet Accuracy of Fe titer Constant current supply + + +
zt0.5 bO.1
Manual Titration,
%
*o.
15
z!co.1
*0.3
10.3
ztl.0
10.8
10.25
A comparison of this error estimation with the experimental results shows that the deviations of a single titration in general are within these error limits, with exception of the titration of 0.28 mg. According to their reduction potentials, permanganate, chromate, vanadate, and uranyl should be reduced by iron(I1)ethylenediamine tetraacetate. However, experiments showed that these ions could not be titrated under the conditions used. Chromate and permanganate destroyed the iron(II1)-ethylenediamine tetraacetate by oxidation of the organic part of the ion, whereas the reaction \vith uranyl and vanadate proceeds too slowly for practical direct titration. A back-titration procedure might be effective, but was not tried. ACKNOWLEDGMENT
This research was supported by the United States Air Force, through the Office of Scientific Research of the Air Research and Development Command. LITERATURE CITED
(1) (2) (3) (4) (5)
Blaedel, W. J., Knight, H. J., A N ~ LCHEAI. . 26, 741 (1954). Cheng, I
