Log Q
=
2.14 - 2.18 E,, 0.55 2 E, 2 0
(4)
where Q is the amount of copper deposited in microcoulombs (pC). The total quantity of copper deposited at 0 V, Q*, is 540 =k 25 pC/cm2for a reduced or an oxidized electrode. Q* represents the total copper involved in the surface controlled reductions and oxidations, R1 - 0x1 and Rz - Oxz. In experiments for which E, is made more negative than 0 V, no evidence for the oxidation of bulk copper is seen (oxidation 0xa)until at least 540 pC/cm2of cathodic charge has passed at the disk electrode. All cathodic change in excess of 540 wC/cm2 is recovered as anodic charge in the potential range corresponding to the process 0x3 (oxidation of bulk copper). CONCLUSIONS
The first monolayer of copper deposited on the platinum electrode has electrochemical properties greatly different from subsequently deposited layers of copper. The reductions R1 and R2 correspond to the reduction of Cu(I1) to form zero valent copper on the platinum surface. The oxidations Oxl and Oxz correspond to the removal of the zero valent copper from the electrode surface. Vassos and Mark (1) present experimental evidence indicating that a monolayer of copper has deposited at a pyrolytic graphite surface from Cu(I1) solution after about 500 pC/cm2 of cathodic charge has passed. The value calculated from the atomic radius of copper is 490 pC/cmz. The total cathodic charge corresponding to R1 and Rz, and the total anodic charge corresponding to Oxl and Oxz, is the same, 540 i 25 MC/cm2,and within experimental error corresponds to a monolayer of zero valent copper. Deposition of a monolayer of zero valent copper is necessary before a bulk deposit forms. Reduction of Cu(I1) to form the first monolayer of zero valent copper takes place in the potential range from +0.5
to 0 V, potentials which are more anodic than the equilibrium potential for deposition of the bulk copper metal. This suggests an interaction between the copper and platinum such as intermetallic compound formation. It is known that copper and platinum form a series of very stable intermetallic compounds (7). However, oxidations quite similar to 0x1 and Ox2 are also seen on pyrolytic graphite electrodes ( I ) , and it is known that intermetallic formation between copper and carbon does not take place (7). A process similar to Rz - 0x2 is seen on gold and also on silver-plated platinum electrodes, however, no process similar to R1 - Rz is seen with these electrode materials. While intermetallic formation is possible between copper and gold, and there is some solid solution of copper in silver (7), the available evidence leads us to conclude that the processes R1 - Oxl and Rz - Ox2 at platinum are not attributable to intermetallic formation or to solid solution. Breiter (Z), who studied the behavior of Cu(I1) at an oxidized platinum electrode, concluded that Oxl is the oxidation of adsorbed copper and that Oxz is the oxidation of thin copper patches. It seems more plausible to us that Ox1 and 0 x 2 represent the oxidation of zero valent copper which has been deposited on platinum sites of varying activity by reductions R1 and Rz. The distribution of active sites is such that the empirically obtained Equation 4 is obeyed at equilibrium. The peak currents observed at $0.4 V and $0.1 V result from unknown kinetic factors and are dependent on the rate of potential scan. RECEIVED for review February 8, 1968. Accepted March 21, 1968. We are grateful to the National Aeronautics and Space Administration for the NASA Traineeship held by G.W.T. (7) M. Hansen, “Constitution of Binary Alloys,” 2nd ed., McGrawHill, New York, 1958.
Analytical Study of a Sulfide Ion-Selective Membrane Electrode in Alkaline Solution Tong-Ming Hseul and G . A. Rechnitz Department of Chemistry, State University of New York, Buffalo,N . Y. 14214
Evaluation of a silver sulfide membrane electrode has revealed sensitivity, selectivity, and other response characteristics well suited to the analytical utilization of such electrodes. Sulfide ion concentrations can be determined by direct potentiometry or potentiometric titration in alkaline solutions without interference from other common anions. The response rate of such electrodes suggests that continuous monitoring of some changing systems is feasible. The thermodynamic solubility product of silver sulfide, (1.48 k 0.01) X lo+, and the formation constant (at p = 0.1) of the SnW- complex ion, (2.062 =k 0.098) x 105, have been determined by means of the sulfide ion-sensitive membrane electrode at 25 O C .
OF THE SEVERAL TYPES of ion-selective electrodes now available (]), the solid-state membrane electrodes appear to hold the most promise for the measurement of anions. Recently Frant and Ross (2) reported on a single crystal lan(1) G. A. Rechnitz, Chem. Eng. News, 45 (25), 146 (1967). (2) M. S. Frant and J. W. Ross, Jr., Science, 154, 1553 (1966).
1054
ANALYTICAL CHEMISTRY
thanum fluoride membrane electrode which is so fluoridespecific that fluoride ion activity can be measured as easily as hydrogen ion activity with a glass electrode. The electrode has been employed for the detection of the end point in titrations (3) of fluoride with thorium, lanthanum, and calcium, for the direct determination of fluoride in tungsten (4), and for the formation constants evaluation (5) of HF and HF2species in acidic fluoride media. The present study was undertaken to provide some quantitative information regarding the analytical usefulness of an analogous solid-state membrane electrode, containing a crystal of silver sulfide as the membrane, to measure sulfide ion concentrations in alkaline solutions and to use such elecOn leave from National Taiwan University, Taiwan. (3) J. J. Lingane, ANAL.CHEM., 39, 881 (1967). (4) B. A. Raby and W. E. Sunderland, ANAL.CHEM.,39, 1304 (1967). ( 5 ) K. Srinivasan and G. A. Rechnitz, ANAL.CHEM., 40,509 (1968).
trodes to determine the formation constant of the tri-thiostannate(1V) ion (SnS3*-). EXPERIMENTAL
Chemicals. Stock solutions of sodium sulfide were prepared by dissolving Fisher reagent grade NazS.9H20 in water. The concentration of sodium sulfide was determined in alkaline media using potassium iodate as oxidant (6). Sulfide was oxidized to sulfate, and the excess of potassium iodate was back-titrated with standard thiosulfate solution by iodometry. The sodium sulfide solution was stored in polyethylene reagent bottles to avoid contact with atmospheric oxygen and carbon dioxide. The solution was restandardized every week; no significant change in sulfide concentration was observed during the period. A series of standard solutions of sodium sulfide was prepared by suitable dilution of the stock solution, keeping the ionic strength of 0.1M by addition of appropriate volumes of 1M sodium nitrate solution. All other chemicals were of Fisher reagent grade and were used without further purification. Stock solutions 0.1M in silver(1) were prepared by direct weight from AgN03. Stannic chloride, SnC14 5H20, of Baker analytical reagent grade was dissolved in excess hydrochloric acid and treated with an excess of sodium sulfide solution until a clear yellow solution was obtained. To this solution an excess of hydrochloric acid was added to decompose the thiosalt and to precipitate the yellow stannic sulfide completely. The precipitated sulfide was repeatedly washed with water until free from chloride ions. Borate and sodium carbonate buffer solutions were used to adjust the pH and the ionic strength if necessary. All solutions were prepared from water which had been both deionized and distilled. Apparatus. A Beckman research model pH meter was used for all potentiometric measurements. E m f . us. time curves for the dynamic response measurements were obtained by displaying the output signal of the pH meter on a Photovolt Model 43 potentiometric recorder at variable chart speeds. The indicator electrode employed in this study was a Model 94-16 sulfide ion-sensitive membrane electrode obtained from Orion Research, Inc. This electrode is similar in principle to that of a conventional glass pH electrode, except that the membrane material is a disk-shaped section of crystalline silver sulfide. To prevent the contamination of the reference electrode, a Fisher 13-639-57cracked bed type saturated calomel electrode in conjunction with a Fisher 13-639-55 remote reference junction was used as the reference electrode system. All measurements were carried out in a thermostated cell at 25” i= 0.02 OC. Procedure. Potentiometric measurements for the construction of the calibration curves were carried out in the conventional manner. The electrochemical cell (120-cc capacity) consisted of a sulfide ion-sensitive membrane electrode, a pH type glass electrode, and a reference electrode. To avoid the absorption of atmospheric oxygen and carbon dioxide during the measurement, purified nitrogen gas was bubbled through the test solution kept at a constant volume of 50 cc. Magnetic stirring (Teflon bar) was employed. Electrode potentials were attained rapidly and in every case the equilibrium potentials were read within 2 minutes after immersion of the electrodes into the test solution. The reproducibility for different solutions of the same sulfide ion concentration was within =tl mV if the electrode surface was kept clean, but the electrode potential tended to drift gradually to more positive values over an extended period. Electrode response curves were obtained after rapid concentration of the test solution by the addition of more concentrated sodium sulfide solution. In all cases, the resulting e.m.f. us. time curves were smooth and reproducible ( 6 ) P. 0. Bethge, Anal. Chim. Acta, 10,310 (1954).
after the initial mixing period (pH and ionic strength were kept constant). Potentiometric titrations were carried out using the incremental addition of standard hydrochloric acid solutions (or appropriate metal salt solutions) to the sample solutions containing known amounts of sodium sulfide. The formation of the thiostannate complex ion was studied for several series of solutions at constant concentrations of sodium sulfide and different pH values. The freshly prepared stannic sulfide was suspended in 200 ml of buffer solution and kept in an airtight Erlenmeyer flask. After shaking, known amounts of the suspension were vacuum-transferred to a 100-ml Erlenmeyer flask and constant amounts of sodium sulfide standard solution and 1M NaN03 solution were added ( p = 0.1). The reaction mixture was shaken vigorously for 3 hours to complete the reaction. The resulting suspension was allowed to stand for an hour or more in a thermostat before the potential and pH measurements were made. pH readings were converted to hydrogen ion concentrations, [H+], using the activity coefficient of 0.76 for univalent ions (7) at ionic strength of 0.10M. RESULTS AND DISCUSSION
Response of Sulfide Ion-Selective Membrane Electrode to Silver(1). The sulfide electrode contains a solid silver sulfide membrane separating an internal reference solution from the test solution. The membrane is an ionic conductor which allows silver ions to pass between the test solution and the reference solution, which is kept at a fixed silver ion concentration. The distribution of silver ions between the two solutions develops a voltage which depends on the silver ion activity in the sample solution. The experimental cell can be represented as Sulfide electrode
I
test solution / I 0.1M 0.1 (NaN03) NaN03
1
SCE
p =
The potential of this cell is given by E
=
constant
+ 2.303 RT - log F
CZA~+
where the “constant” is the sum of the potentials at the internal electrode (silver-silver chloride electrode), the saturated calomel electrode (SCE), the liquid-junction potential between the test solution, and the reference electrode, and the potential across the membrane when the silver ion activity in the test solution is unity. Since the measurements were carried out at a constant ionic strength of 0.1M, the junction potential was kept approximately constant. If one then neglects the junction potential, and introduces EOAg+,Ag = 0.7991 volt, EsCE= 0.2415 volt at 25 O C , and 0.76 as the activity coefficient of the Ag+ ion, Equation 1 becomes
E
=
0.5506
+ 0.05916 log [AS+]
(2)
and the “constant” term in Equation 1 is seen to equal 0.5576 volt using the activity scale. The actual potential of the silver sulfide membrane electrode us. the saturated calomel electrode as a function of silver ion concentration is demonstrated in Figure 1. The straight line has been drawn, over a range of Ag+ ion concentrations from 10-1 to lO-*M, with the slope of 59.0 mV per log [Ag+] unit and the intercept of 0.5535 volt. These values are in agreement with those predicted by the Nernst relationship, The experimental results suggest that the solid silver sulfide (7) H. A. Laitinen, “Chemical Analysis,” McGraw-Hill, New York, 1960. VOL 40, NO. 7, JUNE 1968
1055
6oo
t ei W I
300
t
600
t
0
I
I
I
I
I
2
3
4
-log
2
1
3
hi]
4 -log
Figure 1. Cell e m f . as a function of Ag+ concentration p =
0
A. B.
(3)
6
7
[s“]
Figure 2. S2- calibration curves as a function of ionic strength
0.1M (adjusted with NaNOa)
membrane material used in the indicator electrode is a stoichiometric silver sulfide, containing no excess sulfur in the crystal lattice. Response of Solid Silver Membrane Electrode to Sulfide Ions in Alkaline Media. If the original sample contains no silver ions, a small number of ions are produced by the very low, but finite, solubility of the silver sulfide membrane. The resulting silver ion activity depends on the sulfide ion activity in the sample solution and can be calculated from the solubility product (KeP)of silver sulfide:
5
HS-
p = p =
0.1M 0.3M
C. u
= 0.5M
+ HLOe H2S + OH-;
Kh? =
Klo Ki
=
0.628 X 10-7 (/J = 0.1)
only the first needed to be considered, since the second hydrolysis step has a negligible effect on the sulfide ion concentration in alkaline solutions of sodium sulfide. The thermodynamic ionization constant of HS- ion, K’z = 1.2 X 10-15 or pK‘2 = 14.92 (7-9), can be converted to an ionization constant in terms of concentrations at any given ionic strength by using the extended Debye-Huckel equation (7) (EDHE)2 d i
When this expression for U A ~ +is substituted in Equation 1, Equation 4 relating the electrode potential to the sulfide ion activity in the sample solution results:
E
=
constant
- 2.303 RT - log as%2F
(4)
Le., pK2 = pKI2 - __---. The calculated values of 1 d p pK2 are 14.44 (Kz = 3.635 X lO-l5) at ionic strength of O.lM, 14.21 (K2 = 6.134 X lO-’5) at 0.3M, and 14.09 (Kz = 8.10 X 10-15) at OSM, respectively. From these values one can calculate the sulfide ion concentration, [S2-], via
+
or, in practicality, at constant ionic strength E
=
constant - 0.0296 log [S2-]
(5)
The “constant” term in Equation 5 is the same as in Equation 2, except that it now contains an additional factor involving the solubility product of silver sulfide. The polarity of the potential also changes, so that the cell e.m.f. becomes increasingly negative in more concentrated sulfide ion solutions. Sodium sulfide solutions were used as the source of sulfide ions in experiments carried out to test these equations, and the concentration of the sulfide ion was calculated from the known hydrolysis equilibria. Of the two steps in the hydrolysis Sz-
+ H20 e HS- + OH-;
Kh,
and 1056
ANALYTICAL CHEMISTRY
=
KtG
-
K2
=
2.773 (p = 0.1)
where [S2-It is the total concentration of sulfide present ([Sz-] [HS-I). Equation 6 indicates that at constant sodium sulfide concentration and constant ionic strength, [Sz-] will increase with increasing pH values. Substituting the value of
+
(E)>> 1 , one obtains
[Sz-] in Equation 5 and simplifying [H+l
E
=
constant - 0.0296 pH
(7)
Equation 7 represents the relationship between the potentials and the pH values at a given concentration of sodium sulfide (p = constant). The potential of the silver sulfide membrane electrode us. SCE as a function of sulfide ion concentration is shown in (8) L. Bruner and J. Zawadzki, Z . Anorg. Chem., 63, 136 (1909); 67, 454 (1910). (9) I. M. Kolthoff, J . Phys. Chem., 35,2711 (1931).
1
800
0 4
I
3
I
w’
0 v)
Slopes 8 29.5 mV
d
w
6001
,
I
I
10.0
9.0
11.0
12.0
PH
Figure 3. Calibration curves as a function of pH and sodium sulfide concentration at constant ionic strength 0.1M D. 5 X 10-aM Na2S E. 1 X lO-*MNanS F. 2 X lO-’MNa& p =
Figure 2. These measurements were made with solutions containing sodium sulfide, 0.01 M sodium hydroxide, and sodium nitrate at ionic strengths of 0.1M, 0.3M, and OSM, respectively. The sodium sulfide concentrations were varied between 10-1 and 3 X 10-5M. Small amounts of NaOH were added to increase the buffer action of the test solutions and the pH values were varied from 11.4 to 11.8. The sulfide ion concentrations were calculated by Equation 6 using [H+] calculated from the readings of the pH meter and the activity coefficient at a given ionic strength as described above. The plotted data, over a range of [S2-] from 10-z.2 to lO-’M, follow straight lines. The response curves are Nernstian with slopes of 30.0 mV per concentration decade for an ionic strength of 0.1M (line A ) and 29.7 mV for both 0.3M (line B) and 0.5M (line C) and are within experimental error of the theoretical values given by Equation 5 . The intercepts obtained from these lines on the potential axis were -0.911 volt for line A , -0.883 volt for line B, and -0.864 volt for line C, respectively. A plot of these interceptional potentials against -v$can be drawn as a straight line with an intercept of -0.946 volt, which is equal to “0.5576
+ 0*05916 -log K~;’as repre2
sented by the “constant” term of Equation 4, and of Equation The thermodynamic solubility product constant of silver sulfide thus calculated is (1.48 f 0.1) X 10-b1 at 25 “C,which can be compared with the literature values of 6.2 x 10-s2, by e.m.f. measurement techniques (IO), and 7 X 10-50 by calculation from free energy data (11). In the present study, the constancy of the “constant” term in Equation 5 is dependent not only on the electrode characteristics but also on the stability of the sodium sulfide solutions with respect to the oxidation of sulfide ion by atmospheric oxygen and the absorption of COz during the measurements. For the electrode used in this study, the value of the 5 when the activity coefficient of S2- ion is unity.
(10) J. R. Goates, A. G. Cole, E. L. Gray, and N. D. Faux, J . Am. Chem. SOC.,73,707(1951). (11) J. R. Goates, M. B. Gordon, and N. D. Faux, J. Am. Chem. SOC.,74, 835 (1952).
0.500
M HCI, cc
Figure 4. Titration of 50 ml of 0.055M Na2S with 0.500M hydrochloric acid -Sulfide ion-selective indicator electrode - - - pH-type glass indicator electrode “constant” term remained constant within i:1 mV for freshly prepared sodium sulfide solutions. The slope of the calibration curves was not found to change. The dependence of the electrode potentials (us. SCE) on the pH and the sodium sulfide concentration is shown in Figure 3. The concentrations of sodium sulfide used were 5 x 10-3 M (line D),1 X lO-3M (line E), and 2 X lO-‘M (line F), respectively. The pH values were varied from 9 to 12 by use of borate or sodium carbonate buffer solutions, while the ionic strength was kept constant at 0.1M. The results show that the electrode, at constant sulfide concentration, obeys the theoretically expected pH relationship. Apparently, the electrode responds to the Hf activity in a reversible manner as well as to the Sz- activity in sulfide electrolytes, and the potential us. pH plots of Figure 3 have a uniform slope of 29.5 mV per pH unit, in agreement with that predicted by Equation 7. The electrode obeys the theoretical potential function for sulfide concentrations at a given pH. The experimental results obtained suggest that the sulfide ion-sensitive membrane electrode can be used to determine either the concentration of free sulfide ion or the total sulfide concentration if the pH value of the sample solution is determined simultaneously in aqueous solution. Effect of Anions on Electrode Response. Tests of the effect of several anions-e.g, C1-, Br-, I-, SCN-, NOa-, s04’-, C2042-,and Cr042--on the sulfide ion-sensitive membrane electrode were made over the concentration range of sodium sulfide from 2 x 10-4 to 2 x 10-2 M . The concentration ratios of univalent anions to total sulfide ion were varied from 1 to 480, while those of divalent anions were varied from 0.3 to 160. The pH values of the solutions unavoidably varied from 11.8 to 12.3 during this procedure but were taken into account in calculating the free S2-concentrations. The ionic strength was held constant at 0.1M. For all of the ions tested, the resulting calibration curves agreed with the calibration curve due to SZ-, alone, within the experimental error both in respect to the slope of the curve and the actual cell e.m.f. Thus, the numerical selectivity ratios of the electrode for sulfide ion with respect to these anions VOL. 40, NO. 7, JUNE 1968
0
1057
Fraction Titrated
0.4
-1000
0
-1000 1
1.2
08
Fraction 0.0
0.4
,
,
I
Titrated 1.2 ,
I
1.6
I
1
- 800 -600
-600
-
I
I
2
-400
-400
-
- 200
-200
-
00
-
+
v) 0
I I I
00
I 01
I
W
I 1 I
I
I I
-
I
I
t 200
t200
+ 400
+40 1
+ 600 0
8
12
AgN03,
cc
16
20
At u = 0.1M G. 1.473 X 10-2MNa2S,1.012 X 10-1MAgN03 H. 1.473 X 10-3MNaS, 1.012 X 10-2M AgN03
cannot be evaluated and are, for practical purposes, infinite. In view of this finding, it is clear that this electrode is well suited for the selective analytical measurement of sulfide activity in a wide variety of practical systems, Potentiometric Titrations. In view of the excellent response of the sulfide ion-sensitive membrane electrode to changes in pH at a constant concentration of sodium sulfide solution and in sulfide ion concentration (Figures 2 and 3), one would expect the electrode to be suitable for indicating the course of a potentiometric titration of sulfide. This possibility was tested with the titration of 50 cc of 0.055M sodium sulfide solution with 0.500M hydrochloric acid, as shown by the solid curve in Figure 4. For comparison, the pH titration curve of the same system, using the glass electrode as an indicator electrode, is shown by the dashed curve. The titration curve determined with the sulfide ion-sensitive membrane electrode os. SCE indicates that The two titration end points observed coincide with the end points in the pH titration. The amount of titrant consumed before the first end point is larger than that consumed from the first to the second end point; the amount of titrant consumed for the second end point corresponds to the theoretical value. The potential break at the second end point is about twice as large as that at the first end point and is equal to the breaks of the pH titration curve for the first and second end points. Since considerable amounts of hydrogen sulfide gas are lost (through hydrolysis) when sodium sulfide is first dissolved in water, the concentrations of hydroxyl ion and bisulfide ions are not equal and are represented by the relationship [OH-lpresent
I
I
4
Figure 5. Titration of 50 ml of Na2S with A g N 0 3
=
[OH-ltote.~ - [HS-I - 2[S2-l
and the unequal amounts of titrant used for the titration of sodium sulfide to the first end point and from the first to 1058
I
I
w” >
I
ANALYTICAL CHEMISTRY
I
+600
0
2
I
I
4
6
0.07367 M
,
.
I
,
8
I
io
,
2
Na2S, cc
Figure 6. Titration of 50 ml of 0.0202M A g N 0 8 with 0.07367M NazS At
p =
0.1M
the second end point result from the fact that the sample solution initially contains more hydroxyl ions than bisulfide ions. Because the hydroxyl ion can also be titrated with hydrochloric acid, the positions of the two end points observed in the pH titration coincide with those of the potentiometric titration using the sulfide ion-sensitive electrode. The f a t that the potential break observed at the second end point is twice that observed at the first end point is due to the formation of undissociated hydrogen sulfide in such acidic solutions in the vicinity of the second end point. The sulfide ion concentration thus decreases more markedly than at the first end point where HS- is produced. At the second end point, where the main species in solution is H2S,the potentialpH response is 59.16 mV per pH unit instead of the 29.6 mV at the first end point (Equation 7) because [S2-] = 0.1
*
KI
*
K2/[Hf]2
and, substituting in Equation 5 , E
=
constant - 0.05916 pH
(8)
where 0.1 is the molar solubility of HzS in water at 25 “C, and K, and Kz are the first and second dissociation constants of hydrogen sulfide, respectively. The “constant” term in Equation 8 is not the same as that of Equation 5 , but contains the additional factors K I and K2. The potentiometric titration of sodium sulfide with silver nitrate was also followed by means of the silver sulfide membrane electrode. In Figure 5 the titration curves of both high and low concentrations of sodium sulfide are compared in terms of the fraction titrated. Curve G , obtained for the titration of 1.43 X M sodium sulfide ( p = 0.1) with 0.1012M silver nitrate, shows the large potential change at the Ag2Sequivalent point. The point of maximal rate of change of potential corresponds to 99.6z of the sulfide ion titrated and coincides with the theoretical equivalence point. The
-
second small potential inflection observed between E = 175 10-6.2 M and 420 mV, corresponding to [Ag+] = 10-2.2 (Figure l ) , indicates the formation of silver hydroxide (KBp= 2 X lop8). Since sodium sulfide solutions contain a small amount of free OH- ion initially, some Ag+ will be consumed to form insoluble AgOH after all of the sulfide ion has been precipitated as Ag2S. Curve H shows the titration of 1.473 X M sodium sulfide with 1.012 x lov2 M silver nitrate. The point of maximal rate of change of potential coincides with that of the more concentrated sample, but the small second potential break has disappeared because the OH- concentration is negligibly small. Figure 6 shows the titration curve for the reverse titration of 2.02 X 10-2 M silver nitrate (p = 0.1) with 7.367 X M sodium sulfide using a 0.1M NaNOa agar-agar bridge as a junction between the sample and the SCE. The large potential change occurs at 98% of the theoretical volume and reflects the presence of OH-, produced by hydrolysis of sulfide, in the titrant. It is clear that the sulfide electrode responds faithfully to sulfide activity and is well suited as an indicator electrode in potentiometric titrations, provided side reactions and other chemical complications are taken into account. Dynamic Response Characteristics. The response characteristics of the sulfide ion-sensitive membrane electrode were evaluated by exposing the electrode to a rapid change in sulfide concentration (approximately twofold) and recording the resulting e.m.f. us. time function. The concentration experiment was carried out over a tenfold range of initial sulfide ion concentration (p = 0.1M). Some of the resulting response curves, reproduced in Figure 7, are typical of the results obtained. All curves were smooth and of identical shape; the expected (on the basis of calibration curves) e.m.f. value was attained in all cases. The shapes of the e.m.f. us. time curves are qualitatively similar to those obtained with silver halide precipitate membrane electrodes (12, I S ) , but response times are considerably shorter. The response half-time, tIl2,of the electrode used depends on the stirring method, efficiency of solution mixing, and the cleanliness of the electrode surface, but is independent of the initial sulfide concentrations over the entire range studied. The response half-times obtained are of the order of 1 second for the present experimental conditions-Le., one eighth of those for the AgI precipitate mem(12) G. A. Rechnitz, M. R. Kresz, and S. B. Zamochnick, ANAL. CHEM.,38, 973 (1966). (13) G. A. Rechnitz and M. R. Kresz, ANAL. CHEM.,38, 1786 (1966).
Table I.
[NazSI
x
103, M 8.000 12.00 7.969 12.00 3.993 9.963 6.000 8.000 3.020 9.963
-E,voit 0.6673 0.6762 0.6747 0.6836 0.6732 0.6850 0.6806 0.6871 0.6776 0.6924
4 x
IO-^ M
I x
IO-^
-
8.10 x
IO-^
2.36 x
IO-^
% M-
P
7601
I 0
.
0
IO
.
*
I
.
. # .
20
30 t,
*
I
.
40'
. , . 50
sec,
Figure 7. Dynamic response characteristics of sulfide ionselective electrode AI
=
0.1M
brane electrode (1.3)- and are probably determined mainly by the mixing efficiency. Since the electrode used is an extremely low-solubility solid-state silver sulfide ionic conductor (electrical resistance is less than 1 megohm), the rapid response to changes in sulfide ion concentration is not unexpected. The high response rate of this electrode suggests that continuous monitoring of sulfide ion activities in aqueous solutions would be feasible. Complex Formation Measurements. Because of the excellent sensitivity, selectivity, and response characteristics of the sulfide electrode, it should be very well suited for the study of ion association and complex formation phenomena. The tin(IV)-suKde system was chosen to test this possibility. Stannic sulfide is soluble in high concentrations of sulfide ion to form the thiostannate complex ion. SnSz (S)
+ S2- e SnSaz-
Formation Constants of Thiostannate Complex Ion at 25 OC and = 0.1M [SZ-I [SnSs2-] x 108 ts2-1~ x 103, x 103, g-ion11 PH m+],g-ion/l g-ion11 g-ion11 0.6310 8.298 3.827 x lV9 6.643 1.358 1.259 8.450 2.697 x 10-9 9.341 2.659 1.122 8.610 1.866 X W9 5,758 2.211 2.239 8.760 1.321 X lWg 8.136 3.865 1 .000 8.998 7.811 X 2.149 1.844 2.512 9.104 5.980 X 4.133 5.831 1.778 9.450 2.697 x 1b1O 1.319 4.681 2.951 9.728 1.422 X 1.154 6.847 1.413 10.00 7.600 x lo-" 0.2954 2.724 4.467 10.30 3.810 X 10-11 0.4682 7.532
KI X 10-6 2.152 2.112 1.970 1.726 1.844 2.321 2.633 2.320 1.928 1.609 Av. 2.062& 0.098
VOL 40, NO. 7, JUNE 1968
1059
Babko and Lisetskaya (14) in 1956 determined the solubility of stannic sulfide in hydrogen sulfide solution at different pH values and confirmed that the gram-molecular ratio of the components SnSz and NazSis 1 to 1. According to their work, the equilibrium constant of the formation reaction is 1.1 x lo5 in alkaline sulfide solution as measured by a colorimetric and light-scattering method. However, the direct measurement of the sulfide ion concentration in such equilibrium solution by means of potentiometric techniques has not previously been studied. We therefore carried out complex formation studies by measuring the equilibrium potentials with the sulfide ion-sensitive membrane electrode (us. SCE) and the pH values of several series of equilibrium suspensions containing tin(1V) and sulfide at a constant ionic strength of 0.10 M. Table I gives the results of experiments carried out at different total sodium sulfide concentrations and varying pH values. The first column represents the original concentrations of sodium sulfide added to the suspension of stannic sulfide. The second and third columns give the equilibrium potentials and the corresponding free sulfide ion concentrations derived from the potential ES. -log [S2-] calibration curve ( p = 0.1M). The fourth and fifth columns give the pH and the corresponding hydrogen ion concentrations calculated by using the activity coefficient of hydrogen ion, 0.76, at an ionic strength of 0.1M. Substituting these experimental values of the free (14) A. K. Babko and G. S. Lisetskaya, Zh. Neorg. Khim., 1, 969 (1956).
sulfide ion concentration, [S2-], the hydrogen ion concentration, [H+],and Kz = 3.635 x 10-15 ( p = 0.1M) in Equation 6, one obtains the unreacted total sulfide ion concentration, [S2-],, given in the sixth column. The concentrations of the thiostannate complex ion formed can then be calculated by subtracting [Sz-]i from the concentrations of sodium sulfide initially added ([SnS32-] = [NazS] - [S2-],. The values of the formation constant, K,, calculated via Equation 9 are listed in the last column in Table I--e.g.,
K,
=
[SnS32-]/[S2-]
(9)
The calculated values are fairly constant, with an average value of (2.062 f 0.098) X lo5at 25 "C and p = O.lM, which agrees well with a previously reported value of 1.1 X lo6 at 20 "C(14). Although stannic sulfide also dissolves in alkaline solution to form the hydroxystannate complex ion, SnSzOH-, the present method does not suffer from this complication because the electrode measures only the activity of the free sulfide ion. Thus, the use of sulfide ion-sensitive membrane electrode for the determination of equilibrium constants in complex formation systems offers considerable advantages over conventional methods and could be profitably applied to a variety of practical problems.
RECEIVED for review Feburary 8, 1968. Accepted March 18, 1968. The financial support of NSF Grant GP-6485, NIH Grant GM-14544, and the Alfred P. Sloan Foundation is gratefully acknowledged. One of the authors (T.M.H.) also acknowledges a Fulbright travel grant.
Capacitance Measurements on Platinum Electrodes for the Estimation of Organic Impurities in Water Leonard0 Formaro and Sergio Trasatti Laboratory of Electrochemistry and Metallurgy, University of Milan, Via Venezian 21, 20133 Milan, Italy
The measurement of the double-layer capacitance on platinum electrodes can be used as an experimental method for the determination of last traces of single organic substances in aqueous solutions as well as for the evaluation of industrial water quality. Principles of the method and experimental techniques needed to ensure a high sensitivity are described. Possibilities and limits of the method are discussed taking into account the nature of the organic impurities. Results obtained in the analysis of solutions containing traces of styrene, toluene, acrylonitrile, and amyl alcohol, and in the test of some samples of industrial water are shown. I n the latter case, although the sensitivity is lower than in the former, it was possible to detect carbon contents down to 0.03 & 0.005 ppm.
THERE ARE only a few experimental methods ensuring sufficient accuracy and satisfactory sensitivity for the control of organic impurity level in water. Titration with permanganate allows water quality evaluation down to 10 f 2 ppm (I). The determination by an infrared analyzer of the C02 formed in the oxidation with 0 2 of the carbonaceous material present in (1) R. B. Schaffer, C. E. Van Hall, G. N. McDerrnott, D. Barth, V. A. Stenger, S. J. Sebesta, and S. H. Griggs, J. Water Pollution ControI Federation, 37, 1545 (1965).
1060
ANALYTICAL CHEMISTRY
water shows a sensitivity of 1-2 ppm. The precision of the method is 2 0 z at these water pollution levels ( I ) . It provides a measure of all the carbonaceous material in a water sample, both organic and inorganic. However, if a measure of organic carbon alone is desired, the inorganic carbon content of the sample can be removed during sample preparation ( 2 ) . If no volatile carbonaceous materials are present in the sample, the solution can be concentrated and the practical sensitivity goes down to 0.1 pprn. However, this possibility is rather rare. These methods never provide the sensitivity required in the control of organic impurities in laboratory water in which the pollution level must be low enough for the most careful experimental work. In many cases, the sensitivity of these methods is poor, even for industrial purposes. An electrochemical method ensuring an accuracy of about 15% at carbon contents of 0.03 ppm will be described in this paper. The test of a water sample can be carried out in a few minutes and no concentration of the sample is required to increase the sensitivity. This method is unaffected by low contents in inorganic materials. (2) C. E. Van Hall and V. A. Stenger, ANAL.CHEM., 39,503 (1967).
