more active, mole for mole, than osmium as a catalyst in the cerium(1V)-arsenic(II1) reaction. Ruthenium could be added a t any time during the titration without causing an error. Although Ru(II1) also catalyzes the arsenic(II1)-silver(I1) reaction in perchloric acid, it is not as effective as Os(VII1).The results for coulometric titrations at relatively low current densities are shown in Table V, using 5 or 10 ml of 2.004 x M arsenic(II1) with os04 as catalyst in HC104 and RuC13 as catalyst in H N 0 3 media. It is evident that titration efficiencies of 99.9% or better are readily attainable. Using the Sargent current source, the lowest current setting gave a titration time of 400 s for 5 ml of solution containing 10 pmol of arsenic(II1). Lower concentrations could readily be determined by using smaller currents or shorter titration times. Even with the Sargent source, 0.25 pmol could be titrated in 10 s. Similar recovery rates were observed with silver(1) concentrations as low as 0.01 M and acid concentrations of 1M and 2 M for nitric and perchloric acid, respectively. At relatively high current densities, the current efficiencies were calculated using the equation % current efficienty =
(Et - E x ) / E , X 100
where Et = total equivalents of arsenic(III),E , = equivalents generated at the current density under study, and E x = equivalent of Ag(I1) generated at 100%current density. The results thus obtained at various concentrations of silver(1) in HN03 and HC104 media are presented in Table VI. Using 0.1 M Ag NO3 at current densities below 6 mA cm-2, the current efficiencies were better than 99.5% both for 4-6 M HC104 and 3-5 M “03. At lower concentrations of acid and silver(I), the current efficiency dropped off. Nitric acid is superior to perchloric acid under these conditions, as expected from the lower formal potential of the Ag(II)/Ag(I) couple. It is concluded that tin oxide electrodes can be used for the quantitative coulometric generation of silver(I1) as a titrant in perchloric or nitric acid solutions at room temperature. The
limits of permissible current density depend upon the nature and concentration of acid, the concentration of silver(I), and the rate of the titration reaction. Kinetic information on the uncatalyzed silver(I1) arsenic(II1) reaction as determined from constant rate titrimetry is given elsewhere (13).
LITERATURE CITED (1) D. G. Davis and J. J. Lingane, Anal. Chim. Acta, 18, 245 (1958). (2) H. A. Laitinen and Noel H. Watkins, Anal. Chem., 47, 1352 (1975). (3) K. Kirnura and Y. Murakima, Mikrochim. Mikrochim. Acta, 36/37, 727 (1951). (4) A. A. Noyes, J. L. Hoard, and K. S. Pitzer, J. Am. Chem. SOC., 57, 1221 (1935). (5) A. A. Noyes, C. D. Coryell, F. Stitt, and A. Kossiakoff. J. Am. Chem. SOC., 59, 1316 (1937). (6) H. N. Po, J. H. Seinehard, and T. L. Allen, lnorg. Chem., 7, 244 (1968). (7) M. Fleischman,D. Pletcher. and A. Rafinski, J. Elecfraanal. Chem., 38, 329 (1972). (8) J. B. Kirwin, F. D. Peat, P. J. Proll, and L. H. Sutcliffe, J. Phys. Chem., 67, 1617 (1963). (9) J. B. Kirwin, F. D. Peat, P. J. Proll, and L. H. Sutcliffe, J. Phys. Chem., 67, 2288 (1963). (10) G. A. Rechnitzand S. B. Samochnick, Talanta, 11, 713 (1964). (1 1) G. A. Rechnitz and S. B. Sarnochnick, Talanta, 12, 479 (1965). (12) G. W. Harrington,H. A. Laitinen, and V. Trendafilov, Anal. Chem., 45, 433 (1973). (13) J. M. Conley, Ph.D. Thesis, University of Illinois, Urbana, Ill., 1975. (14) B. F. Rider and M. G. Mellon, lnd. Eng. Chem., Anal. .Ed., 18, 96 (1946). (15) E. A. Klobbie, Chem. Zenfralbl., 69, 65 (1898). (36) V. G. Levich, ActaPhysicochim. URSS, 17, 257 (1942). (17) D. P. Gregory and A. C. Riddiford, J. Chem. SOC.,3756 (1956). (18) P. W. Carrand J. Jordan, Anal. Chem.. 45, 634 (1973). (19) “International Critical Tables of Numerical Data, Physics, Chemistry and Technology,” Vol. 111, McGraw-Hill, New York, N.Y., 1928, pp 54-59. (20) R. N. Adarns, “Electrochemistry at Solid Electrodes,” Marcel Dekker, New York, N.Y., 1969, p 85. (21) J. Heyrovsky and J. Kuta, “Principles of Polarography,” Academic Press, New York, N.Y., 1966, pp 104-108. (22) H. Kim and H. A. Laitinen, J. Electrochem. SOC., 122, 53 (1975). (23) M. Fleischman, D. Pletcher, and A. Rafinski, J. Appl. Electrochem., 1, 1 (1971). (24) D. Elliott, D. L. Zellmer, and H. A. Laitinen, J. Necfrochem. Soc., 117, 1343 (1970). (25) H. Pardue, Anal. Chem., 39, 600 (1967). (26) C. Surasiti and E. B. Sandell, Anal. Chim. Acta, 22, 261 (1960).
RECEIVEDfor review January 22,1976. Accepted March 15, 1976. This work was supported by the National Science Foundation.
Potentiometric Response of Silver(I) Sulf ide/Copper(II) Sulfide Membranes to Chelons and Applications for End-Point Detection in Chelometric Titrations Virgil K. Olson,” James D. Carr, Robert D. Hargens,’ and R. Ken Force2 Department of Chemistry, University of Nebraska-Lincoln,
Lincoln, Neb. 68588
In the absence of Cu(ll) ion, the Ag2S/CuS membrane electrode responds to chelons in solution via the cupric ion-chelon complexation equilibria at the membrane surface. At pH N 12 in 3 M ammonla buffered solution, potentiometric titrations in the absence of Cu(ll) yield sharp peak-shaped end points for a number of metal chelate systems with relatively large conditional stability constants. Metals having smaller bonditional stabllity yield S-shaped curves. Precision approaches the limit characteristic of the volumetric equipment used. A small titration blank occurs attributable to competitlve equilibria between ammonia and the chelon employed. Present address, Dorsey Laboratories, P.O. Box 83288, Lincoln, N e b . 68501. Present address, Department of Chemistry, University of Rhode Island, Kingston, R.I.
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ANALYTICAL CHEMISTRY, VOL. 48, NO. 8 , JULY 1976
Other investigators have studied and described methods for applying Ag2S/CuS membranes as indicating electrodes in chelometric titrations of metal ions (1,2).These procedures are based on the addition of the metal complex C U L ~ + ~ (where L is the chelon used in the titration) to the solution containing the metal ion Mn+ to be titrated. The equilibrium Mn+ + CuLm+2 + C$+
+ MLn+”
established in conformity with the relative stability of the metal complexes establishes a finite concentration of Cu(I1) to poise the electrode. When the solution is titrated, the effect of additional chelating agent on the equilibrium and the corresponding response of the Ag2S/CuS membrane to Cu(I1) over the course of the titration provides the means of potentiometric end-point detection.
Recently, we observed that a AgZS/CuS membrane electrode produced sharp peak-shaped end points in the absence of Cu(I1) when used for Ni(I1)-EDTA titrations in ammonia buffered solutions (Figure 1).Further investigation revealed similar titration curves could be obtained with other metal ions or complexing agents. We report here results of studies on the nature of this electrode response and its general applicability to complexometric titrations.
EXPERIMENTAL Membrane Electrodes. The AgZS/CuS membrane electrodes used in the study were constructed in our laboratory. The membranes were prepared using the procedure of Czaban and Rechnitz ( 3 ) and mounted using procedures peviously described ( 4 ) . The response of each membrane was tested using solutions prepared from cupric nitrate with the ionic strength held constant a t 0.1 with potassium nitrate. All of the membranes used showed Nernstian response (slope M. N 30 mV) and detection limits below Apparatus. Potential measurements were made using a Corning Model 12 expanded scale pH meter or a Heathkit DVM. Automated potentiometric titrations were performed using a Sargent-Welch Model DG recording titrator. Saturated calomel and double junction reference electrodes were employed. Reagents. All solutions were prepared with deionized and distilled water and stored in polyethylene bottles. Solutions of copper and nickel prepared from high purity wire (Ventron Alfa Inorganics) by dissolution in nitric acid served as primary standards. Other metal cation solutions were prepared from reagent grade nitrate salts except for those prepared from CaC03 (Mallinckrodt, primary standard grade), MgO (Fisher, USP), HgClz (Fisher, reagent grade), cadmium wire (Alfa Inorganics), and Mn(CH&00)~-4H20(Matheson, Coleman and Bell, practical grade). Complexone solutions were prepared from Matheson, Coleman and Bell NaZHz-EDTA (disodium salt of ethylenediaminetetraacetic acid, reagent grade), H4-CyDTA (cyclohexanedinminetetraacetic acid, practical grade), and “trien” (triethylenetetramine, technical grade). mol Procedure. Solutions containing between 3 and 300 X of metal ion were titrated with 0.002 or 0.1 M titrant. Initial sample volumes were adjusted to approximately 50 ml after adding 10 ml of concentrated ammonia. When an automated titrator was employed, the rate of titrant addition was manually set to the lowest available setting (0.04 ml/min) prior to the end point.
RESULTS AND DISCUSSION pH Effects. Figure 2 shows the AgZS/CuS membrane response in the presence of various complexing agents as a function of pH. Differences in the NH3 and EDTA NH3 profiles correspond roughly to the overall change in potential observed during EDTA titrations as a function of p H in the absence of Cu(I1). A reduction in ammonia concentration raised the profiles to more positive potentials and reduced the magnitude of the potential difference (Le., the NH3 profile approached the K N 0 3 profile and the EDTA NH3 profile approached the EDTA profile). The NH3 and trien profiles suggested a larger overall potential change was to be expected for trien titrations with a reduced p H dependence of the magnitude of the change between pH 10 and 12. This proved to be the case although trien titrations yielded round S-shaped end points instead of the sharp peak-shaped end points observed with EDTA titrations. The addition of hydroxide ion or a complexone to the solution affects the membrane potential via the complexation equilibria at the sample/membrane interface. In the alkaline pH region, hydroxide ion, ammonia, aminocarboxylates, or polyamines act to effectively “buffer” Cu(I1) membrane surface activity to very low levels (see pCu scale in Figure 2). In solutions containing no Cu(I1) ion, relatively stable and reproducible potentials were obtained in high pH ammonia buffered solutions in the presence of the various complexing agents. With decreasing pH, longer equilibrium times were required to achieve stable and reproducible potentials and hysteresis effects reflecting the membranes previous environment were evident. The pH profiles shown in Figures 2 and
+
+
-4000
10
20 30 40 rnl 0.05 @ Na3H-EDTA
Figure 1. Experimental titration curves using 0.0538 M EDTA as titrant in a 3 M ammonia background Shown are the titrations of 1.02 mmol of Ni(l1) alone and in combination with 0.996 mmol of Cu(ll) when conducted by hand (0) or with an automated titrator (-)
Y4 10-6
loot
r J
‘8 12-10
0-
-1277
-
0
n p Q14-14
z-too_I
C
- 16
-200 -
t
EDTA
16-18
NHf
+
- 20
18-22
-300
2
4
6
6
1012
PH
Figure 2. Ag&CuS membrane response as a function of pH in 0.1 M KN03, M EDTA, 1 M NH3, M EDTA 1 M NH3, and M trien
+
I
I
I
\\
1001
i
-J 100
\\
I
TRIEN’
-501
4
,
I
2
4
I
15
\ I
I 49
8 1 0 1 2
6
PH Figure 3. Ag2S membrane response as a function of pH in 0.1 M KN03, 4X M EDTA, and 5 X M trien
3, while qualitatively correct, do not necessarily represent readily reproducible potentials, particularly in the moderately alkaline and acidic p H regions. Behavior of AgzS Membranes. Since the membrane potential is ultimately determined by the membrane silver ion activity (5), a study was made to determine what significance could be attributed to interactions between Ag(1) ion and the complexones a t the membrane surface. On the basis of known stability constants, Ag(1)-complexone interactions were expected to be insignificant relative to Cu(11). T o confirm this, p H profiles for a AgzS membrane were obtained in solutions containing KN03, EDTA, and trien (Figure 3). The silver ion activity as a function of p H for the AgZS membrane was then compared to the silver ion activity calculated for a Ag2S/CuS membrane (see pAg axes in Figures 2 and 3). The Ag(1) ion activity for the Ag2S/CuS membrane was calculated by means of the equation UAg+
=
[(K~s/Kf~s)0,Cu2+]1’2
ANALYTICAL CHEMISTRY, VOL. 48, NO. 8, JULY 1976
(1) 1229
I
Figure 4. Experimental titration curves illustrating typical responses
in the region of the end point for the various metal determinations summarized in Table I The Hg(II)-EDTA and Hg(ll)-CyDTA titration curves were actually twice the size illustrated
be equivalent to -3.5 X low5M complexone in 3 M ammonia solution. Experimental Titrations. Clearly a number of parameters can have marked effects on the integrity of the end point since anything which influences the Cu(I1) activity a t the membrane/solution interface can be expected to influence the observed potential. In practice, no problems arose when normal titration procedures were used. Potential readings were steady in stirred solution and showed little drift with time, although long term changes to more positive potentials could be observed with decreasing p H because of loss of ammonia. In general, the shapes of the end points correlate well with the conditional stabilities of the ML complex or can be interpreted in terms of the expected chemical kinetics influencing the complexation equilibria. A large conditional stability constant combined with rapid complexation kinetics preventing the buildup of free ligand in solution until excess titrant is introduced, can be expected to facilitate a sharp end point. The Ni(I1)-EDTA complex, having a relatively large conditional stability gives a sharp end point in 3 M ammonia while Ba(I1)-EDTA titrations under the same conditions give S-shaped breaks. Changing the titrant to CyDTA for Ba(I1) titrations results in sharper end points as a result of the increased stability. While the conditional stability for Ni(I1)trien is sufficiently large, the exchange reaction Ni(NHs),Z+
as the soluand K g S = 6.3 X using K f F S = 6.3 X bility constants for AgzS and CuS, respectively (6). The much lower silver ion activity observed for AgZS/CuS membranes relative to AgzS membranes in the presence of EDTA or trien indicates the Cu(I1)-complexone equilibria is of far greater significance. Thus, membrane behavior can be interpreted directly in terms of Cu(I1)-complexone interactions and the Ag(1) ion membrane activity can be assumed to remain dependent on the Cu2+/S2-/Ag+ equilibria. Stirring Effects. In the presence of small amounts of complexing agent (solution also about 3 M in NH3), the membrane potential was found to reflect lower Cu(I1) ion activity in stirred than in quiet solutions. Stopping the stirring resulted in slow drifting to more positive potentials, while stirring a quiet solution produced a rapid change to a relatively stable and more negative potential. With EDTA, a difference of -100 mV was observed. The effect could be readily observed just beyond the end point of a titration by stopping and starting the stirring action. This behavior is consistent with slow dissolution of the membrane and the effect of stirring on the Cu(I1) activity gradient in the vicinity of the membrane/ solution interface. Increasing the complexone concentration resulted in reduced potential differences in stirred and quiet solutions. The actual dissolution of the AgZS/CuS was found to be negligible from an analytical applications standpoint. Polished membranes showed only traces of pitting or tarnish under magnification after several weeks of titrations with strong complexing agents in 3 M "3. No deterioration in electrode response was observed with extended use. Ammonia Blank. The presence of ammonia introduces a small titration blank. While the magnitude of the blank increases with ammonia concentration, the relationship was not found to be linear. Nor did the blank increase with extended time of membrane contact with the solution prior to titration. These observations indicate the blank is due neither to introduction of contaminant metal ions in the reagent ammonia nor from dissolution of the membrane. Instead, the blank can be attributed to the solution equilibria involved, it being necessary to achieve a small but finite complexone concentration before competitive equilibria between the titrant and ammonia begins to influence the membrane surface Cu(I1) activity. For EDTA, CyDTA, or trien, the blank was found to 1230
ANALYTICAL CHEMISTRY, VOL. 48, NO. 8, JULY 1976
+ trien
-
Ni(trien)2+
+ nNH3
(2)
seems to be kinetically slow in 3 M NH3 at pH 12 relative to the normal titration time scale. Lowering the pH to 10 facilitates the exchange and a more rapid titration rate without reducing the magnitude of the overall titration curve. As shown by Figure 1,some titration curves move gradually to more positive potentials prior to the end point. When titrations were conducted by hand, the change to more positive potentials was somewhat erratic while relatively smooth peak-shaped curves could be obtained using the automated titrator (see Figure 1).The sharpness of the peak was observed to depend on the rate of titrant addition. Too rapid a titration rate resulted in rounding of the peak. While the evidence is not conclusive in pinpointing the mechanistic nature of the peak phenomena, a t least two possibilities provide plausible explanations. An increase in Cu(I1) activity at the membranelsolution interface prior to the end point could be caused by a gradual decrease in pH (resulting from loss of ammonia or from protons contributed by the EDTA titrant) or adsorptioddesorption equilibria of the analyte metal ion on the electrode membrane surface and corresponding effects on the Cu(I1) activity. Figure 4 shows the titration curves for several metal ions while Table I summarizes results and special titration conditions. Precision was found to be limited by the volumetric apparatus. Accuracy was improved by making a pretitration prior to the addition of the sample aliquot. The magnitude of the potentiometric break was considerably larger for Hg(I1) than for other metal ions, supporting the observation by other investigators ( 2 ) that the membrane responds directly to Hg(I1). Such behavior is believed to be due to the lower solubility of HgS and displacement of Cu(I1) by Hg(I1) to produce a mercury sensitive membrane/solution interface. High pH conditions of 3 M NH3 solutions preclude the titration of many metals. However, in some cases, the titration can be made feasible by a reverse titration or by the addition of a secondary complexing agent to suppress hydrolysis of the metal ion. Lead, for example, can be titrated in the ammonia buffered solution if a small amount of tartrate is added. Hydrolysis was a problem in Mg(I1)-EDTA titrations. The use of sodium citrate to suppress the hydrolysis and CyDTA as titrant significantly improved the accuracy of results for
Table I. Summary of Metal Ion Determinations Metalhitrant= Ni(I1)-EDTA Ca(I1)-EDTAC Ca(I1)-EDTA Cd(I1)-EDTA Zn(I1)-EDTA Ba(I1)-EDTA Hg(II)-EDTAd Pb(I1)-EDTAd Mn(I1)-EDTAe Pb(I1)-EDTAf Sr(I1)-EDTA Ca(I1)-CyDTA Mg(I1)-CyDTAg Zn(I1)-CyDTA Ba(I1)-CyDTA Cd(I1)-CyDTA Pb (11)-CyDTAf Sr(I1)-CyDTA Hg(I1)-CyDTA Ni(I1)-trien Cd(I1)-trien Zn(I1)-trien
Amount expected, gmoles
Amount found, gmoles
(primary standard)
Re1 std dev, %b 0.15 0.38 0.34 0.21 0.05 0.00 0.19 0.04 0.05 0.26 0.36
5.513 137.8 58.10 256.0 3.296 282.5 274.5 263.5 219.6 103.6
5.524 138.6 57.55 255.2 3.279 284.0 272.2 263.4 219.7 103.9
137.8 200.1 204.8 211.2 58.1 54.9 51.8 226.0
139.2 200.9 203.5 209.6 55.3 51.3 225.5
0.11 0.40 0.14 0.07 0.43 0.30 0.62 0.00
255.0 58.1 51.2
254.5 58.4 51.8
0.35 0.75 0.20
57.5
Representative titration curve Figure 1 Figure 4 Figure 4 Ni(I1)-EDTA Figure 4 Zn(I1)-EDTA Figure 4 Ni(I1)-EDTA Ni(I1)-EDTA Ni(I1)-EDTA Figure 4 Figure 4 Figure 4 Ca(I1)-CyDTA Ca(I1)-CyDTA Ca(I1)-CyDTA Figure 4 Ca(I1)-CyDTA Figure 4 Figure 4 Figure 4 Cd(I1)-trien
a A minimum of three and maximum of five replicate titrations were used for each metal ion determination. Normal titrant concentration was 0.1 M. Mean deviation 0.25%. Titrant 0.002 M for this determination. Excess EDTA added prior to the concentrated ammonia and a known amount of excess Ca2+.The titration was then continued with EDTA. e Ascorbic acid (0.1 g) added to each sample. f Sodium tartrate (0.5 mmol) added to each sample. g Sodium citrate added to each sample.
Mg(11)determinations. Proper choice and concentration levels of secondary complexing agents is necessary as the presence of other complexones in the sample solution may reduce the conditional stability of the desired complexation sufficiently to adversely affect the break. While rather elaborate schemes can be applied to facilitate the titrations of various metal ions, i t is our feeling that if hydrolysis or kinetics interferes, other available titration methods ( I , 2) should normally be considered first. Although membranes other than Ag2S/CuS were not tested, it would not be surprising to find other Ag2S/MS membranes showing similar behavior with overall behavioral differences being consistent with differences in the conditional stability of the membrane metal ion-solution complexone. In summary, the response of the AgZS/CuS membrane to complexones such as EDTA in ammonia buffered solutions offers a simple method for end-point detection in complexometric titrations for a number of commonly titrated metal ions. The method eliminates the need for addition of an indicating metal ion or preparation of specific pH buffers. For
some metal ions, end points are peak-shaped making it possible to read end points directly without determining the point of maximum slope.
ACKNOWLEDGMENT Portions of this work were completed using instrumentation made available by the Colorado State University Department of Chemistry, Fort Collins, Colo., and Dorsey Laboratories, Lincoln, Neb. LITERATURE CITED (1) J. W. Ross and M. S. Frant, Anal. Chem., 41, 1900 (1969). (2) E. W. Baumann and R. M. Wallace, Anal. Chem., 41, 2072 (1969). (3)J. D. Czaban and G. A. Rechnitz. Anal. Chem., 45, 471 (1973). (4) V. K. Olson, J. D. Carr, R. C. Hargens, and R. C. Larson, J. Chem. Educ., 51, 791 (1974). (5) R. A. Durst, Ed., "Ion Selective Electrodes", Nat. Bur. Stand. (US.), Spec. Publ. 314. Washinaton. D.C.. 1969. DD 79-81. (6) J. A. Dean, Ed., "L&?ge'sHandbook Of'Chemistry", McGraw-Hill, New York, N.Y. 1973, Chap. 5, p 7 ff.
RECEIVEDfor review February 6,1976. Accepted March 15, 1976.
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