Chelometric Indicator Titrations with the Solid-state Cupric Ion-Selective Electrode James W. Ross, Jr. and Martin S . Frant Orion Research Inc., I 1 Blackstone Street, Cambridge, Mass. 02139
SOMEYEARS AGO, Siggia, Eichlin, and Rheinhart ( I ) showed that silver and mercury metal electrodes and a small amount of the silver or mercury complex (which serves as an indicator in the usual sense) could be used for the titration of metals for which no electrode was available. The method was extended by Reilley and Schmid (2), who presented a very elegant and simple method of predicting the titration curve from pHpotential plots, and used the mercury electrode with EDTA. Because both groups were limited to mercury and silver metal electrodes, the indicator-electrode titration technique has not been as widely adapted as its usefulness would suggest. The problems have included redox interferences, sensitivity to chloride, and the need, under some conditions, to purge with nitrogen. The theory of the silver and mercury metal electrode titrations has recently been extended by Hulanicki and Trojanowicz (3). We have found that the new solid-state membrane electrodes, particularly the cupric ion electrode ( 4 ) , can be used in the same way, but with a wide variety of chelating agents, ordinary titration equipment, and without much concern about chloride or redox interference. EXPERIMENTAL All measurements were made using the Orion Cupric Ion Activity Electrode, Model 94-29, a double junction reference electrode, Model 90-02, with 10% K N 0 3 in the outer compartment, and the Orion Model 801 digital pH/mV meter. All of the chemicals used were Analytical Reagent grade, except for the l,lO-phenanthroline, which was Eastman White Label, and the tetraethylene pentamine (TEPA), Eastman Blue Label. We found it necessary to purify the TEPA for sharper end points. The following procedure was used: TEPA was added dropwise with vigorous stirring, to a 10- to 20-fold excess of concentrated HCI in an icebath. The resulting slurry of TEPA.5HCl was mixed with a n equal amount of methanol, vacuum filtered, and washed with cold methanol. The damp material was dissolved in the minimum amount of hot water,. a volume of Darco G-60 activated charcoal equal to the amount of solid material was added, and the mixture filtered after 10 minutes. The filtrate was heated to 50-55 "C, and methanol was added dropwise until the solution remained cloudy. It was then allowed t o stand at room temperature for 1-2 hours, filtered, washed with cold methanol, and air-dried overnight. This product was used for preparing standard solutions without converting to the free base. Copper indicator stock solution was prepared by titrating 0.1M Cu(NO& with the appropriate chelating agent and stopping just before the end point (90-99 %). The indicator
(1) S. Siggia, D. W. Eichlin, and R. C. Rheinhart, ANAL.CHEM.,
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Figure 1. pH-potential plot and corresponding titration curve for a metal ion to be titrated ( M ' ) using an electrode and an indicator ion (M*+)to which the electrode responds. For details, see text
is added to the solutions to be titrated a t a total copper level of I-lOz of the unknown metal ior, level. RESULTS AND DISCUSSION Because it has been our experience(5)that many present-day analytical chemists are not sufficiently familiar with the work of Reilley and Schmid, Figure 1 has been included to explain the significance and value of pH-potential plots, and to show how we determined suitable titration conditions quickly and easily. For the theoretical and mathematical background o n the use of these plots, reference (2) should be consulted. For the uppermost line in Figure 1, the response of the electrode and its metal ion-e.g., cupric ion activity electrode and cupric ion-at a fixed activity level are determined as a function of pH. In general, there is no change until the precipitation of the hydroxide occurs. In the absence of complexing agents, the precipitation of hydroxide establishes the upper limit of metal ion which can be present a t any given p H during the titration. In a similar experiment, the electrode response is determined in the presence of a 50% excess of complexing agent. This lower limit gives the level of free metal ion that will be present after the end point is reached. These two curves define the pH and metal ion levels within which useful titrations can be performed. For each nonelectrode metal ion to be titrated, concentration (as a function of pH) is calculated for a solution containing equal amounts of the free ion and the complex. Alter-
27, 1745 (1955).
(2) C. N. Reilley and R. W. Schrnid, ibid., 30, 947 (1958). (3) A. Hulanicki and M. Trojanowicz, Talaitta, 16,225 (1969). (4) J. W. Ross, Jr., National Bureau of Standards Symposium on
Ion-Selective Electrodes, Gaithersburg, Md., Jan. 30-31, 1969. 1900
s. Frant, Pittsburgh Conference on Analytical Chemistry and Applied Spectroscopy, Cleveland, Ohio, March 7, 1969. Discussion following presentation.
( 5 ) J. W. Ross, Jr. and M.
ANALYTICAL CHEMISTRY, VOL. 41, NO. 13, NOVEMBER 1969
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nately, and as we have usually done, the potential of the equimolar mixture of metal ion and complex is measured as a function of pH with a small amount of indicator complex present. These three curves then permit a direct calculation of the titration curve for the system. In the hypothetical example in Figure 1, it was decided that a titration at pH 8 would be convenient. Point I then gives the expected electrode potential at the start of the titration. Point 2 gives the potential half-way to the end point, and Point 3 gives the potential an equal dis-
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Figure 4. Experimental titration curves using tetraethylene pentamine hydrochloride as titrant in an ammonia background. Shown are the titrations of 10-3M solutions of Cu2+, Ni2+, and Zn2+, and a 1 : 1 mixture of Cu2+and Zn2+
tance past the end point. From these three points, the titra. tion curve can be constructed. As an example, consider Figure 2, which is the actual potential-pH curve for copper and 1,IO-phenanthroline. The curves suggest that Zn2+and CdZ+can be titrated with 1,lOphenanthroline, using Cu2+ as the indicator, anywhere between pH 3 and 12, Ni2+ only between pH 3 and 8, and that the largest end point break for Ni2+ ought to be at pH 6. This has been confirmed experimentally. A lO-3M nickel perchlorate solution titrated with 10-2M 1,lO-phenanthroline gives an end point break of over 300 mV. As expected, the break between copper and nickel was small (about 60 mV). What makes this method particularly attractive is not only the ability to predict a particular titration, but also the ability to predict what happens if two or more metals are present simultaneously. Consider Figure 3, which shows the pH potential curve for the use of tetraethylene pentamine as a titrant. The data indicate that this system will work well for Pb2+,Ni2+,or Zn2+,and that large differences in end points will occur between any of these and copper. Further, the optimum separation between Ni2+ and Zn2+ should occur in the region of pH 8.5-10.5. In Figure 4, tetraethylene pentamine titrations are shown run in an ammonia background. A 10-3M Ni2+ solution gives an end point break of over 350 mV? a 10-3MZn2+ solution has a break of at least half as much, and both can probably be determined simultaneously. Also shown is the titration of an equimolar mixture of Cu2+and Zn2+. As could be qualitatively predicted, the copper break is a little smaller than the zinc break, but both are quite usable. This should make a very straight-forward method for the analysis of brass and similar alloys. EDTA is, of course, a widely used titrant. As Figure 5 shows, the copper electrode will follow nickel and cadmium titrations with EDTA in mildly acid or neutral solutions, and it will follow the calcium titrations in basic solutions, preferably a t p H 10-11.5.
ANALYTICAL CHEMISTRY, VOL. 41, NO. 13, NOVEMBER 1969
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Figure 6 shows the titration of 10-3M and 10-4M Ca2+in a background of 10-5M Cu2+ with ammonia present to adjust the pH. At this pH, the maximum CaZ+ break is obtained, but copper interferes. This is not a problem if copper is added as the indicator at relatively low levels, but if both are desired, the copper should be determined first at pH 3-4. This titration is particularly interesting, because it is superior to titrations using even the most highly selective of calcium electrodes as end point indicators (6, 7). The titration will go to lower levels, and there is no sodium interference. This same technique can be used to determine both Ca2+ and Mgz+, because the EDTA titration gives both CaZ+and MgZ+,while a titration with EGTA gives Cazf only. For (6) J. W. Ross, Jr., Science, 156, 1378 (1967). (7) J. A. King and A. K. Mukherji, Naturwissenschaften, 53, 1643
1902
ANALYTICAL CHEMISTRY, VOL. 41,
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Figure 6. Titrations of 10T3MCa2+ and 10-4M Ca2+ in a background containing ammonia and 10-6M Cuz+. At 10-4MCaZ+there is a good endpoint, but a large blank due to copper, which can be minimized by using copper-EDTA complex as the indicator example, we have found usable titration curves for CaZ+ only Mgz+ with in a mixture of 5 X 10-4M CaZf and 5 X EGTA, and determined both at the same level with EDTA. We have also gotten good end points titrating for Ca2+in a ten-fold excess of Mg2+(5 x lo-* M and 5 X lO-3M, respectively) with EGTA. ACKNOWLEDGMENT
The authors gratefully acknowledge the technical assistance of Anne Goulston.
RECEIVED for review July 7,1969. Accepted August 22, 1969. Presented in part at the Pittsburgh Conference on Analytical Chemistry and Applied Spectroscopy, Cleveland, Ohio, March 7, 1969.
NO. 13, NOVEMBER 1969
