V O L U M E 28, N O . 4, A P R I L 1 9 5 6
443
The presence of selenium affects the titration, however. No actual titration was made with selenium compounds present, because it was found that the selenium in solution reduced hypobromite as rapidly as i t was formed. Although the species of selenium involved is not known, experiments made u-ith selenious acid showed that selenite reduced hypobromite. It is evident that if selenium is to be used as a catalyst, some way must be found t o eliminate interference before the ammonia determination is begun. Experiments 11-ere performed in which metallic lead was added to the sulfuric acid solution containing selenium in the hope that the selenium would be reduced to thc metal. Although most of the selenium was removed from the solution in this way, 4 days were required to complete the reduction. Reduction with sulfur dioxide ( 1 ) mas not considered because of the time required for removal of the excess.
tities of ammonia between 14 y and 230 y. Both the relative and absolute accuracies are somewhat better than those reported by Laitinen (9). ACKNOWLEDGiMENT
Preliminary experiments made by Fred C. Anson have been helpful in this work. G. Myron Arcand is indebted to the General Electric Co. for a predoctoral fellowship during the academic year 1954-55. LITERATURE CITED
(1) Adams, C. I., Spaulding, G. H., ANAL.CHEM.27, 1003 (1955). (2) Buck, R. P., Swift, E. H., Ibid., 24, 490 (1952). (3) Clark, P. P., “Semimicro Quantitative Organic Analysis,” p. 40, Academic Press, New York, 1943. (4) Clusius, K., Rechnitz, G., H e h . Chim. A c t a 36, 59 (1953). (5) Farrington, P. S., Meier, D. J., Swift, E. H., ANAL. CHEY. 2 5 , 5 9 1 (1953).
Table IV. No. of
Detn.
Taken
(6) Jones, G., BaeckstroGm, S., J . Am. Chenz. SOC.56, 1517 (1934). (7) Kolthoff, I. AI., Stenger, lr. A., I X D . I h G . CHEW ANAL. E D .
Confirmatory Titrations .4mmonia, y Found
Error
Std. Dev., Y
Confirmatory Titrations. Table IV shows the results of a series of confirmatory titrations made by the procedure described above. I n all cases, the average error was less than 0.2 y, mhile the standard deviation was about the same. Thus, one might expect a maximum error of 0.4 y in a determination involving quanti-
7 , 7 9 (1935).
(8) Kolthoff, I. M., Stricks, W., hIorren, L., Analyst 78, 405 (1953). (9) Laitinen, H. A., IVoerner, D. E., AXAL.CHEY.27, 215 (1955). (10) hIeier, D. J., Myers, R. J., Swift, E. H., J . A m . Chem. SOC.71, 2340 (1949). (11) Milton, R. F., Waters, W.A., “Nethods of Quantitative bIicro Analysis,” p. 79, Edward Arnold and Co., London, 1949. (12) Ramsey, IT. J., Farrington, P. S., Swift, E. H., AXAL.CHEM. 2 2 , 3 3 2 (1950). (13) Riley, R. F., Richter, E., Rotheman, A I . , Todd, N., Myers, L. S., Jr., Nusbaum, R., J . Am. Chem. Soc. 76, 3301 (1954). (14) Willard, H. H., Cake, W.E., Ibid., 42, 264G (1920).
RECEIVED for review July 18,
1955. Accepted ,January 10, 1956. Contribution No. 2016 from the Gates and Crellin Laboratories of Chemistry.
Coulometric Titrations with Electrically Released Ethylenediaminetetraacetic Acid Titrations of Calcium, Copper, Zinc, and lead CHARLES N. REILLEY and WILLIAM W. PORTERFIELD Department o f Chemistry, University o r North Carolina, Chapel Hill, N. C.
Successful coulometric titration of calcium, copper, zinc, and lead ions was accomplished by the indirect electrical generation of ethylenediaminetetraacetic acid released upon reduction of mercuric-ethylenediamine tetraacetate chelate at a mercury pool. The pertinent equilibrium conditions are summarized in a potentialpH diagram. This diagram, in conjunction with polarograms for kinetic effects, furnished valuable information concerning the solution conditions desired. Extension of this method to the titration of many other metal ions appears feasible.
7 OULOMETRIC titrations with constant current have b een employed in various t,ypes of oxidation-reduction, precipitation, and neutralization reactions. The present investigation was undertaken to test the feasibility of extending the application of coulometric titrations into compleximetry. If a general method for the coulometric generation of ethylenedianiinetetraacetic acid [EDTA, HiY, (ethy1enedinitrilo)teti-aacetic acid] could be found, it would be highly probable that the coulometric method could be used for the successful titration of
(A
the alkaline earths, the rare earths, and such metal ions as manganese(I1 j, iron(II), copper(II), mercury(I1j, cobalt, nickel, zinc, cadmium, lead, aluminum, gallium, indium, thorium, palladium(II), and bismuth. The solution of i,his problem required the discovery of a mode of generating ethylenediaminetetraacetic acid with lOO7, current efficiency, anti the development of a compatible and sensitive end-point detection device. The properties of the end-point method selected are discussed in greater detail elsewhere (8). Mode of Generation. Several modes of coulometric titration of metal ions through the formation of stable chelates with ethylenediaminetetraacetic acid suggest themselves. The formation of the reagent, ethylenediaminetetraacetic acid, through an electrode process such as the osidation of ethylenediaminetetraacetaldehyde, was ruled out on the basis of requiring reagents which were not commonly available in pure form and because doubt existed about the tjuccessful generation a t 100% current efficiency of ethylenediaminetetraacetic acid in such a manner. It seemed that either of the following two indirect titration procedures might be employed. The first would require the coulometric generation of base in order to titrate the acid liberated when an excess of ethylene-
ANALYTICAL CHEMISTRY
444
diamine tetraacetate (at a p H of either 4.5 or 8.2) is added to the solution of the metal ion a t a corresponding pH. For example: p H = 8.2 HY-3
+
pH = 8.2 Ca++
CaI'--+H+
+
(1)
The H + would then be titrated, until the pH returns to 8.2, by the O H - formed by the following electrode reaction:
2H10
+ 2e
+ H17
+ 20H-
(2)
The simultaneous electrolytic reduction of any divalent metal ion would lead to high results and must be prevented. This procedure is also undesirable because of the necessity of having the p H of the sample solution and the ethylenediaminetetraacetic acid solution carefully controlled a t 4.5 or preferably 8.2 prior to the titration. The second indirect titration procedure would function in the following manner. A known amount of standard ethylenediaminetetraacetic acid is added in excess to the sample containing the metal ion. The excess ethylenediaminetetracetic acid is then back-titrated by a metal ion derived from the anodic dissolution of a metal.
HY-3
+ Ca++ AI
AI+*
-.c
+ Cay--
AT+"
+ HY-3 (excess)
+ ne +
+ H+
MYn-4
+H+
dropping mercury electrode for lead, cadmium, or zinc chelates. The large hydrogen overvoltage on mercury is also desirable. Other factors which influenced the choice of mercury-ethylenediamine tetraacetate complex were the ease of using a mercury pool as the generating electrode and the mercury drop as a corresponding end-point detection device. Back-titration was also possible by simply reversing the direction of current flow. h black substance (probably mercurous odide) was sometimes obtained near the electrode surface with anodic currents. It, however, quickly disappeared upon stirring and had no noticeable effect on the results. POTENTIAL-pII DIAGRAM
The characterization of the selected chemical system a t equilibrium is summarized by the potential-pH diagram shown in Figure 1. This diagram may be used to predict the effective~ e s ofs the coulometric titration under various p H conditions and for metal-ethylenediamine tetraacetate chelates of differing stability. I n addition the properties of the potentiometric end point under these same conditions can be elucidated,
(3) (4)
(5)
Although this method permits the use of buffered solutions of pH'a where selected reactions might occur, it does not eliminate the preparation and use of a standard ethylenediamine tetraacetate solution. The method finally evolved consisted of the liberation of ethylenediaminetetraacetic acid by the following electrode process: HgY--
+ 2e
+
Hgo
+ Y-4
(6)
The ethylenediamine tetraacetate thris liberated is used to titrate the metal ion of the sample:
Ca++
+ Y-4
-,C a y - -
(7)
In this procedure, the simultaneous reduction of divalent ions to metals does not cause any error. The selection of mercury for the metal-ethylenediamine tetraacetate complex was based on the following principles. First, it was desirable that the stability constant of the complex be large, The metal-ethylenediamine tetraacetate complex would then have sufficiently great stability to allow titrations in acid solutions. It was also desirable that the hydroxide (or oxide) of the metal selected have a sufficiently large solubility product relative to the instability constant of the metalethylenediamine tetraacetate complex, so that the complex would not decompose into insoluble hydroxides (or oxides) in alkaline solutions. ?rIercury-ethylenediamine tetraacetate complex, with an instability constant of approximately 10-z2, seemed to be a favorable choice. A second factor concerns the interference of oxygen throcigh simultaneous reduction in the electrode generation process. Polarograms show that mercury has the most positive reversible half-wave potential of the common metal ions that also possess sufficient stability with ethylenediamine tetraacetate. Even in the case of mercury it was found necessary to exclude oxygen from the electrolysis cell. Employment of metals with more negative half-wave potentials would require an even more careful elimination of oxygen. Selection of a metal with a high positive half-wave potential ia, also desirable in minimizing the extent of the undesirable side, reaction, the liberation of hydrogen. For example, it would be even more difficult to attain 100% current efficiency if zinc-. ethylenediamine tetraacetate complex were used. I n fact, Pecsok (6) reports that no reduction wave was obtained a t th(1
3
4
S
Figure 1.
7
6
8
9
10
PH Potential-pH diagram
The diagram was constructed from a combination of experimental data and reported equilibrium constants (3, 6, 9). The equilibrium constant for Equation 14 wa9 calculated from experimental data. The following equilibria are pertinent: "I+
IInYn-'
-+
H+
4
H'
+ Hn-,Y*-5
+ KH3
pK
(8)
4.7
pK1 = 10.26 pK2 = 6.16 pK8 = 2.76 (9) pK4 = 2.0 where n = 1, 2, 3, 4 respectively HgY-- -t H g + + Y-4 p K = 22.1 (10) &I7-- -L C a + + Y-4 p K = 10.7 (11) M Y - - + AI++ Y-' p K = 10, 12, 14, 15, 16,. (12) Hg(NH3)2+++ H g + + 2"~ pK1Kz 17.5 (13)
+ + +
+
..
+ +
NH3HgY-- -+ HgY-iYH3 p K = 6.6 HgO(r) HzO 20HH g + + p K = 25.52 Hg,ClZ + 2'21Hgz++ pK 17.96 0.612 us. S.C.E. Hg -+ Hg++ 2e EO = 2Hg -+ Hg,++ 2e Eo = 0.547 vs. S.C.E.
+
-+
+ +
+
+ +
(14) (15) (16) (17) (18)
44s
V O L U M E 28, N O . 4, A P R I L 1 9 5 6 For the purpose of constructing the diagram for use in establishing optimum titration conditions, it is desirable to obtain a line representing the system a t a point prior to the end pointhalf way to the end point, for example-and another line a t a point after the titration-as far past the end point, for example, as the other line was prior to the end point. These data were obtained experimentally by mixing the proper chemicals in the correct proportions to represent these two situations and then measuring the potential of the system a t various p H values. I n this manner lines I, 11, 111, and V were obtained. The remaining lines were calculated on the basis of the data given in Equations 8 to 18. The exact concentrations employed are listed beneath each species in the reactions I to V, each reaction corresponding to the line in Figure 1 bearing the same designation. Prior to End Point. The most positive potential that may be obtained is ultiniately limited by the formation of mercuric oxide. This region is plotted in Figure 1 from the data given in Equations 15 and 17. In practice a complexing agent-e.g., aninionia -is added to prevent precipitation of mercuric oxide and the potential limit is consequently even more negative. In the titration of metal ions which have stability constants with ethylenediaminetetraacetic acid greater than that of mercuric cthylenediamine tetraacetate, the following displacement reaction occurs prior to titration [assuming no “complex effect” ( 9 ) ] : SPI”4f
R”I+
+ HgI‘-- + M++ M Y - - + Hg(NH3),+++ 2 H + (below p H 4) + NH3HgY-- + If++ -, illy-- + Hg(hTHa)z+’ + H + (above pH 4) --+.
-
(19) (20)
The equilibrium electrode process is then given by the reaction: Hg(NH3),++ (0.001M )
+ 2 H + + 2e
Hg
+ (o.oon1) 2NH4+
(1)
I n agreement with Equation I, the slope of curve I in Figure 1 corresponds t o a two-electron, two-proton reaction. The metal ions, whose stability constants with ethylenediarninetetraacetic acid are less than that of mercuric ion, do not undergo reactions 19 or 20 over the entire p H range. The weaker the chelating power of the metal, the higher the p H must be for extensive conversion. Thus, for example, at p H 9 all metal ions of p K (with ethylenediaminetetraacetic acid) greater than 11 are “leveled” to the species Hg(NH3)2++which in turn would be the species actually titrated. Consequently, these nietal ions separately titrated would yield identical potentiometric titration curves. M7hen mixed, no separate breaks would result but only a final break representing the titration of the sum of these ions. On the other hand, a t p H 4, only metal ions of pk’ (with ethylenediaminetetraacetic acid) greater than 16 are “leveled” to Hg(NHz)z++, whereas metal ions of pK (with ethylenediaminetetraacetic acid) weaker than 11 are so unstable that they m-ould not titrate a t all. Calcium, with an instability constant (pK) of 10.70 (ethylenediaminetetraacetic acid), serves as a typical example. Between p H values 4 and 9, the potential, a t the point where half of the calcium has been titrated, is given by the following electrode reaction:
H + + Ca++ (0.00111)
+ NHaHgY-+ 2e (0.0 1111)
-
+ +
Cay-Hg “4’ (11) (o.Ooln1) (O.Od1)
A3 seen in Figure 1, the slope of the experimental potential-pH curve I1 docs correspond to this two-electron, one-proton reaction. The intersection of curves I and I1 occurs a t a pH value (approximately 9) where Reaction 20 is essentially complete towards the right. The intersection of curves I1 and I11 occurs at a p H value (approximately 4) where Reaction 20 is essentially complete towards the left. At pH 9, calcium will titrate as effectively as any metal ion of greater stability with ethylenediaminetetraacetic acid. At p H 4, calcium will not titrate.
Curve
I V represents the equilibrium reaction:
At p H below 4, reaction IV exists to the left; above 4, to the right. The dotted line (curve IV) a t pH 4 corresponds to the condition where the concentration of HgY -- and NH,HgI’-are equal (0.005.M) in a 0.05X ammonium ion solution. After the End Point. Curve I11 in Figure 1 corresponds to the situation past the end point by a definite amount and a t a pH greater than 4. NH,HgY-(n 1)H+ 2 e - XHI’ H,I’”-‘ Hg (O.OlA1) (0.05J1) (0.00131) (111)
+ +
+
+
+
Below pH 4, the reaction is illustrated by curve V. HgY-nH+ 2e + H,Yn-‘ Hg (0.0 lllI) (0.00 1111)
+
+
+
End-Point Break. The break in the end point (as well as the extent of the titration reaction) is given by the difference between the upper curves-e.g., I, 11-and lower curves (I11or V). This end-point region is illustrated for the case of calcium ion titration by the shaded portion of Figure 1. For example, a t pH 8.5, the potential will be around 0.1 volt us. S.C.E. a t a point in the titration corresponding to half-titration (CaY-- = C a t + = O . O O l J 1 ) . The potential a t an equal point past the end point (HY--- = O.OOlil1) would be -0.035 volt us. S.C.E Metal iom which form more stable ethylenediamine tetraacetate complexes than calcium give approximately the same results a t this pH. Using Figure 1, one can also conclude that calcium could not be titrated in solutions as acid as p H 4. Reasonable end-point breaks and the extent of reaction increases a t higher pH values and a pH of 7 or greater is necessary for effective titrations. This is based on the arbitrary assumption of 0.1 volt difference between upper and lower lines. A t p H of 4 to 4.5 metal ions with stability constants greater than approximately 10’5 [assuming no [‘coniplex effect” (9)] can be titrated selec1,ively in the presence of calcium. Thus, for example, a two-component mixture-e.g., lead and calcium-could be analyzed in the same solution by titration a t p H 4.5 (for lead) and then changing pH to 8.5 for the titration of calcium. Interference of Chloride. On the edge of Figure 1 are illustrated the interference levels of chloride ion caused by Reactions 16 and 18. For example, in the presence of l O - 3 . V chloride ion, the most positive potential obtainable m-ould be approsiniately 0.2 volt. Chloride ion a t this concentration, therefore, interferes with titrations carried out a t p H 7 or less. To carry out a titration a t pH 3, the chloride ion should be less than 10-6-11 according to Figure 1. Actually and fortunately, this is not necessary in practice. The potential that actually occurs will be a “mixed potential,” with Reactions 16 and 18 competing with Reaction 19. Because the smallest cboncentration of any species involved in Reaction 19 is lO-3Af (the latent hydrogen ion concentration would be even greater when solution is properly buffered), Reaction 19 mill greatly predominate over Reactions 16 and 18 (if chloride concentration were 1 0 - 6 X ) in determining the net potential. Only when the titration is 99.9% complete will the opposing concentrations (chloride us. metal ion) be of the same magnitude. If sufficient time were allowed a t the beginning of the titration for equilibrium, the chloride ion concentration would be reduced from, say approximately 10-6 to l O - T X , a t the expense of generating an equivalent amount of mercuric diamine. The error introduced upon subsequent titration of the mercuric diamine would correspond to approximately 5 X 10-7.11 metal ion a t most-a negligible error. K~~ETICS
The potential-pH diagram is a useful graphic summary of the pertinent thermodynamics of the system. ‘From such a diagram valuable predictions can be postulated. For example, in the
ANALYTICAL CHEMISTRY
446
titration of calcium a t pH 8, the potential of the mercury pool
at the start of the titration will be in the neighborhood of 0.1 volt us. S.C.E. If we forcibly changed this potential to -0.05 volt, current would flow and the situation in the titration vessel would change until a new state of equilibrium mas reached which would correspond to the exact titration of the calcium From the number of coulombs required, the quantity of calcium could be calculated. The diagram does not, however, tell the rate of flow of current and thus the length of time required for the analysis. The diagram predicts equilibrium, not rates.
PLAT IN U M
c I
FOIL
ANODE
:S
6 '
-MERCURY POOL CATHODE
Figure 2.
Coulometric titration cell
In the constant-potential coulometric titration the rate of current flow generally would be limited in the most favorable case by the rate of diffusion of the species being titrated. The over-all reaction in the titration of calcium (Equation 11) would be controlled by the rate of diffusion of the species on the left of the equation. I n practice the solution is well buffered and contains a large excess of NH,HgY-- [ammine(ethylenediaminetetraacetato)mercurate(II)] and, therefore, the diffusion of calcium ions to the surface of the mercury pool might impose the upper limit on the rate of reaction. The limiting factor actually was found to be controlled by a kinetic process in the diffusion
Table I. Cation
Amount Taken, Rig.
Aniount Found, hfg.
cu
3.52
3.48 3.54 3.59 Av. 3 .54 7.10 7.00 7.06 Av. 7.06 14.01 14.10 14.20 .4v. 14.10
7.05
14.10
Zn
0.120
Av. 0.300
Av. 3.27
Av. 6.54
Av.
layer and, consequently, the over-all rate was much lower. This over-all rate process is studied simply by obtaining a polarogram for this solution. Figure 1 shows that a reduction wave should appear near +0.1 volt us. S.C.E. corresponding to Equation 11. The height should be proportional to the concentration of calcium ion if no kinetic effect is operative. A second reduction wave should appear a t -0.02 volt us. S.C.E. corresponding to Reaction 111, and its height would be proportional to the NH3HgY-- concentration. In practice only the second wave appears. If the same procedure is carried out a t pH 9.3, more favorable results are obtained. The polarogram now shows two steps, the first wave ( N +0.075 volt us. S.C.E.) corresponding to Reaction I in place of Reaction 11, and the second wave (- -0.050 volt us. S.C.E.) corresponding again t o Reaction 111. Thus by changing pH conditions slightly, a different set of species [Hg(NH3)2+' and Cay-- in place of NH,HgY-- and Ca++] occurs whose electrochemical kinetics is more favorable to rapid reduction. Thus, a combination of potential-pH diagrams and the polarograph as complementary tools is seen to be very useful for establishing best conditions for coulometric analysis a t controlled potential. In the constant current method, the over-all electrode reaction is forced to take place a t a given rate. If the desired reaction is too slom, another reaction will occur simultaneously and the entire process may not allow stoichiometric calculations through Faraday's law. I n practice the electrode process is never controlled by diffusion of the species being titrated, since this rate would be less than the current flow in the neighborhood of the end point. Therefore, the electrode reaction is arranged such that the product(s) of the electrode reaction will react in solution with the species being titrated. In this way the same over-all reaction occurs, but in roughly two stages. One occurs a t the electrode surface and the other in the bulk of the solution. The electrode reaction must be rapid, but the solution reaction can be as slow as can be tolerated in an analysis. The major electrode reaction in the coulometric titration of calcium a t constant current is given by Equation 111. Because the rate of this electrode reaction is controlled by the diffusion of NH,HgY-- to the electrode surface, the cell was constructed to enhance diffusion by employing a large electrode surface to solution volume ratio
Coulometric Titration of Metal Ions with EDTA
I
Error, hfg.
Error,
-0.04 0.02 0.07 0.02 0.05 -0.05 0.01 0.01 -0.09 0.00 0.10 0.00
70
-1.2 0.6 a 2.1 0.6 0.7 -0.7 0.1 0.1 -0.6 0.0 0.6 0.0
0.121 0.121 0.121 0.121
0.001 0.001 0.001 0.001
0.8 0.8 0.8 0.8
0.309 0.309 0.301 0.306 3.28 3.30 3.28 3.28 6.58 6.56 6.49 6.55
0,009 0.009 0.001 0 . 006 0.01 0.03 0.01 0.02 0.04 0.02 -0.05 ' 0.01
3.0 3.0 0.8 2.3 0.3 0.9 0.3 0.6 0.6 0.3 -0.7 0.1
'
Cation
Amount Taken, hfg.
Zn (Conld.)
13 .08
Pb
10 36
Amount Found, Rlg. 13.17 13.08 13.12 Av. 13.12 10.36 10.36 10.31 Av. 10.34 20.80 20,90 20.44 20.72 41.73 41.60 41.48 41.60 2.02 2.03 2.00 2.01 4.96 4.97 4.96 4.96 9.89 9.95 9.96 9.93
20.72
Bv. 41.44
Av. Ca
2.01
Av. 4.93
Av. 9.86
Av.
Error, hIg.
Error, %
0.09 0.00 0.04 0.04
0.6 0.0 0.3 0.3
0.00 0.00 -0.05 -0.02
0.0 0.0 -0.5 -0.2
0.08 0.18 -0.28 0.00 0.29 0.16 0.04 0.16 0.01 0.02 -0.01 0.00 0.03 0.04 0.03 0.03 0.03 0.09 0.10 0.07
0.4 0.9 -1.4 0.0 0.7 0.4 0.1 0.4 0.5 1.0 -0.5 0.0 0.6 0.8 0.6 0.6 0.3 0.9 1.0 0.7
447
V O L U M E 28, N O . 4, A P R I L 1 9 5 6 and providing efficient stirring action. With these two factors fixed, it was then desirable to find the critical lower concentration necessary to yield 100% current efficiency-that is, to prevent simultaneous generation of hydrogen. This concentration for the current employed was found by following the potential of the mercury pool as measured amounts of the stock mercuricethylenediamine tetraacetate solution was added. At the critical concentration the potential shifts suddenly towards a much more po;itive potential. A more detailed discussion of this procedure is given elsewhere (1, 2, 6). The concentration necessary was 7 X 10-3JI. In practice a threefold excess was added to ensure against any transient effects such as lower stirring action.
The standard calcium solution was prepared by dissolving approximately 0.9 gram of calcium nitrate monohydrate in 200 ml. of water and standardizing this solution against the standard ethylenediamine tetraacetate solution a t a pH of 9, using Eriochrome Schwarz T indicator. The mercuric ethylenediamine tetraacetate solution was prepared by dissolving approximately 16.S grams of mercuric nitrate and 18.6 grams of disodium dihydrogen ethylenediamine tetraacetate in 500 ml. of water. This solution (approximately 0.lM) proved to have a slight excess of mercuric ion, which was taken into account by the pretitration procedure. An approximate 0.LU ammonium nitrate solution acted as a buffer, as aqueous ammonia was used in bringing each sample to the proper pH. A saturated potassium sulfate solution was used as an electrolyte in the anode compartment.
APPARATUS
A specially constructed titration cell was used (see Figure 2) with a mercury pool cathode and a platinum foil anode. A Leeds 6: Northrup p H meter (Model 76G4) was employed to measure potential reading from the 700-mv. scale. The endpoint system consisted of a saturated calomel reference electrode and a mercury J-tube as an indicator electrode in which a drop of mercury is held in a cup connected through an enclosed platinum wire to a tube of mercury. The constant current supply and associated timer are described in detail elsewhere ( 7 ) . The current magnitude was 43.49 == ! 0.04 ma. REAGENTS AND SOLUTIONS
All chemicals were reagent grade. The standard copper solution was prepared by dissolving an accurately weighed quantity of electrolytic sheet copper in 5 ml. of concentrated nitric acid and diluting to 250 ml., producing a 0.02111 standard solution.
PROCEDURE
A 20-ml. portion of mercuric-ethylenediamine tetraacetate stock solution and 55 ml. of ammonium nitrate solution were mixed and placed in the titration cell, which had a mercury pool covering the entire bottom of the flask, and brought to a p H of 8.5 with concentrated ammonia solution. This was done conveniently by adding the base until the potential of the indicator read approximately + O . l to +0.12 volt os. 8.C.E. (see Figure 1). Potassium sulfate electrolyte was placed in the side arms of the cell so that the center arm had a higher liquid level than the outer arm, and both were higher than the solution level in the flask. The mercuric-ethylenediamine tetraacetate solution was then bubbled out with nitrogen for about 10 minutes to remove dissolved oxygen. Failure to remove oxygen was found to give high results. The excess mercuric ion \vas then pretitrated, using 100-second intervals in the beginning and 10-second intervals when the change in voltage revealed an approaching end point. A sample of metal ion mas than pipetted into the titration cell, and the pH was again adjusted by adding concentrated aqueous ammonia t o the solution, while in the titration cell, until a potential of approximately +lo0 mv. was reached (see Figure 1). The solution was bubbled out again with nitrogen for about 5 minutes, and the sample was then titrated using the same time intervals as in the pretitration. When an end point was reached, the pretitration and the sample titration were plotted and the time interval between the two end points was found as illustrated in Figure 3. The current was checked every 50 to 100 seconds to ensure its remaining constant throughout the titration. The results of the titration of four typical cations are summarized in Table I. ACKNOW LEDGM EKT
This research mas supported by the United States Air Forcc through the Office of Scientific Research of the Air Research and Development Command. -40
n
‘
1
I
300
I
1
I
TIME
Figure 3.
1
600
IN
I
t
900
8
I
I
1
17.00
1
1500
SECONDS
Typical titration curve for calcium
The standard zinc solution was prepared by dissolving 0.3271 gram of 30-mesh granular zinc metal in 5 ml. of concentrated nitric acid and diluting to 250 ml., producing a 0.02001dI solution. The standard lead solution mas prepared by dissolving l.G56 grams of lead nitrate in 250 ml. of mater, producing a 0.02000ilI solution. The standard ethylenediamine tetraacetate solution was prepared by dissolving approximately 18.6 grams of disodium dihydrogen ethylenediamine tetraacetate (Bersmorth Chemical Co.) in 1 liter of water and standardizing this solution against the standard lead solution a t a pH of 9, using Eriochrome Schwarz T indicator and a tartrate buffer.
LITERATURE CITED
(1) Adams, R. N., Reilley, C. N., Furman, N. H., ANAL.CHEV.25, 1160 (1953). (2) Badoz-Lamhling, J., Anal. Chim. Acta 7, ii85 (1952). (3) Bjerrum, Jannik, ‘‘hieta1 Ammine Formation in Aqueous Solution,” dissertation, Copenhagen, 1941. (4) Furman, N. H., Reilley, C. N., Cooke, W.D., ANAL.CHEW23, 1665 (1951). ( 5 ) Latimer, W.AI., “Oxidation Potentials,” 2nd ed., Prentice Hall, New York, 1952 (6) Pecsok, R. L., J . Chenz. Eclztc. 29, 597 (1952). (7) Reilley, C. K.,unpublished manuscript. (8) Reilley, C. K., J . Chem. E d z ~ c 31, . 543 (1951). (9) Schwarzenbach, Gerald, “Die komplexometrische Titration,” Ferdinand Enke Verlag, Stuttgart, 1955. RECEIVED for review September 30, 195.5. Accepted January 17, 1956.
