1792
ANALYTICAL CHEMISTRY, VOL. 51, NO. 11, SEPTEMBER 1979
Chloride Interference in Cupric Ion Selective Electrode Measurements John C. Westall,'g2Francois M. M. Morel, and David
N. Hume''
Ralph M. Parsons Laboratory for Water Resources and Environmental Engineering, Massachusetts Institute of Technology, Cambridge, Massachusetts 02 139
A theory based on the establishment of a diffusion layer between the electrode surface and the bulk of solution, and the attainment of chemical equilibrium at the electrode surface is used to explain quantitatively the behavior of solid-state cupric ion-selective electrodes in chloride media: Cu( 11) from the bulk of solution is reduced at the electrode surface to Cu(I), which is stabilized by chloride complexation. The theory was experimentally verified: in solutions with different concentrations of Cuz+, CI-, and CuZf-complexing ligands; with an alternate Cu( I) stabilizing agent, CH,CN; with cupric ion selective electrodes of different standard potentials (a sulfur-rich graphite contacted electrode and a silver contacted electrode). The onset of interference of chloride ion is described as a function of concentrations of Cu2+, CI-, and Cu2+-complexing ligands, and the standard potential of the electrode. The general unsuitability of the electrode for measurement of cupric ion activity in seawater is accounted for by this theory.
Solid-state cupric ion-selective electrodes based on a mixed silver sulfide-copper sulfide phase have been used successfully for measurement of cupric ion activity in "uncomplicated" media ( I , 2 ) . These electrodes respond with a slope of nominally 29.6 mV/decade change in cupric ion activity a t 25 "C, as predicted by the Nernst equation for a reaction involving a two-electron transfer. However, difficulty is encountered when these electrodes are used in seawater or other media of high chloride ion activity. A relatively slow response and sensitivity to stirring are observed, along with a change in response slope and apparent E" of the electrode. A slope of nominally 59 mV has been observed (3-51, which would correspond to a reversible reaction involving a oneelectron transfer. This behavior is not predicted, nor can it be readily explained, by the traditional view of the mixed copper sulfide-silver sulfide electrode as an electrode of the third kind in equilibrium with Cu2+,and governed by the solubility equilibria of the metal sulfides ( I ) . In this work we shall present a theory for the behavior of the mixed sulfide cupric ion-selective electrode in chloride media, give experimental evidence in support of the theory, and discuss the analytical application of the electrode in chloride media. We show that the "chloride interference" with a cupric ion-selectiveelectrode involves the reduction of Cu2+a t the surface of the electrode with concomitant oxidation of the mixed sulfide electrode material. This redox reaction is favored by stabilization of the reaction product, cuprous ion, through complexation with chloride ion, and the "interference" may in fact be observed in other cuprous complexing media.
THEORETICAL Response of the Electrode in Uncomplicated Media. Given the construction of the cupric ion selective electrode 'Department of Chemistry, MIT, Cambridge, Mass. 02139. Present address: EAWAG, CH-8600 Duebendorf, Switzerland. 0003-2700/79/035 1-1792$01.OO/O
and thermodynamic information on the component phases, we may attempt to predict the electrode's standard potential, EoISE.This calculation has been carried out for many types of solid state ion-selective electrodes, with some success in predicting experimentally observed E0rSE(6, 7). The problem lies in the uncertainty about the exact nature and free energy of formation of the components of the solid electrode phase, as well as in the assumption that the component phases are in equilibrium. We consider the theoretical EoIsEof two commercially available cupric ion selective electrodes which were used in the experimental part of this study. The Radiometer cupric ion-selective electrode, used for the majority of the theoretical and experimental studies, consists of a sulfur-rich silver sulfide-copper sulfide mixture applied to the end of a Teflon impregnated graphite rod. As shown by Koebel (7) and Buck and Shepard (6), the standard potential of such electrodes may be related to the chemical potential, or activity, of elemental copper in the mixed sulfide; this chemical potential can, in theory, be calculated from thermodynamic information on the phases present in the active material of the electrode. As an ideal limiting case for the sulfur-rich mixed sulfide electrode, one can take the phases So,CuS, and Ag,S to be present a t equilibrium; then using the reaction describing the formation of CuS, Cuo
+ So = CuS
AGfgs (CuS)
(1)
and taking the chemical potentials of sulfur and cupric sulfide to be their standard values a t 25 "C (pLso= 0, pcUs = AGaS8), we may calculate the chemical potential of elemental copper in the sulfur-rich mixed sulfide (pcuo = pcus - pso = JG;" (CuS)). Then using the general definition of chemical potential, p = po + RT In ( a ) , we may define an "activity" of elemental copper in the electrode material, acUo(ISE). The activity of elemental copper and the standard potential of the reaction Cu2+ + 2e- = Cuo
E " C ~ Z + , C ~ O (11)
yield the theoretical potential of the ion-selective electrode for a given activity of cupric ion in solution,
ECu~+IISE = E°CU2+,CU~ + 2F
or a t unit cupric ion activity in solution the E" of the ionselective electrode is =
E°CU2+IISE E°CU2tJCUo
+ RT 2F In
(&)
Conversely, the activity of elemental copper in the mixed sulfide phase may be found experimentally from the measured E°Cu~+,IsE. A comparison of the theoretical and experimental values for both electrodes used in this study is found in Table I. The currently marketed Orion electrode is a mixed silver sulfide-copper sulfide in contact with a silver wire (8). If the phases Ago, Ag,S, Cu2S are in equilibrium and the chemical potential of elemental silver is pAg0 = 0, and the chemical C 1979 American Chemical Society
ANALYTICAL CHEMISTRY, VOL. 51, T a b l e I. Theoreticala and E x p e r i m e n t a l Values of
phases in equilibrium soI c u SI Ag, s AgolAg2SlCu,S
' Calculated from p K ,
E o for Two Different
NO. 11, SEPTEMBER 1979
C u p r i c Ion-Selective Electrodes
theoretical E°CU2+/ISE
59 2 457
experimental
(acLl0)
(aAg0)
10-8- 6
10-3.5 1.0
10-6.3
1793
--
Eo
(%O)
1.0 10-7.0
Cu2'DSE
597 527
-
(aC1lO)
10-". 5 10- . 4
values given by Buck and Shepard (6).
potentials of Ag2S and Cu2S are their standard values a t 25 "C, we may use the energies of the reactions
2Cu0 + So = Cu2S AGfg8 (Cu2S) (IV) and the method used before, to calculate the thermodynamic activity of elemental sulfur and copper in the mixed sulfide and the theoretical E" of the silver-contacted ion selective electrode (Table I). Koebel ( 7 ) has made a more thorough analysis of the component phases of the silver contacted silver sulfide-coppen sulfide mixture, but his results are not greatly different from those presented. We see from Table I that the agreement between theoretical and experimental values of E" (and a c p ) Le., the equilibrium description of the electrodes given in Table I, is reasonable but not perfect, being better for the sulfur-rich graphite contacted electrode. We see also that the silver-contacted electrode has a much higher acuothan the graphite-contacted sulfur-rich electrode; this difference will be reflected in different susceptibilities of the electrode to the chloride interference. The use of the formalism of acuo in the mixed sulfide does not imply that the potential of the electrode is actually defined by a reversible reaction of cupric ion with elemental copper in the mixed sulfide; the use of acuo is simply a convenient way to express potentials of reactions of cupric or cuprous ions with the mixed sulfide electrode. The exact nature of the reversible reaction that defines the electrode potential is unimportant. This subject is treated elsewhere (9,101. What I S important is that cupric ion reacts reversibly with some component of the mixed sulfide in a two-electron transfer reaction, and that the cuprous ion reacts in a similar way in a one-electron transfer reaction. This is the basis of the theory for the chloride interference. Mechanism of Interference. As an explanation for the observed chloride interference with the cupric ion selective electrode, we propose a simple theory based on two fundamental ideas; (i) in unstirred solutions there is a small finite volume of solution in the immediate vicinity of the electrode that is affected by heterogeneous reactions between soluble species in solution and the mixed sulfide of the electrode; (ii) the reduction of divalent copper to monovalent copper takes place at the electrode surface. The actual mechanism of this reaction is unimportant (although we shall suggest one); it is important that Cu(1) is formed. For the sulfur rich So/Ag2S/CuS electrode, a reaction of the type cu2+
+ c u s = 2Cu+ + SO
(VI
could occur in a chloride medium when the cuprous ion is stabilized by chloride complexation: c u + + 2c1- = cuc1,-
KCuC1*-
(VI)
That is, cupric ion from the bulk of solution is reduced at the electrode surface as the mixed sulfide is oxidized. In order to simplify the expos6 of the theory, the following assumptions are made: (i) The initial bulk concentration of divalent copper is Ccucnl and of monovalent copper CcU,,,= 0.
3 s+cn:e
',om
E1e:Ircde
5brf3ce
Figure 1. Schematic concentration profiles of Cu(I1) as a function of distance from the electrode surface at different times in an unstirred solution, expressed as a fraction of the bulk concentration. As time progresses, the diffusion layer becomes thicker, the concentration gradient becomes flatter, and the rate of material transport to the electrode surface becomes slower
(ii) Bulk concentrations are not significantly affected by reactions a t the electrode surface. (iii) Chloride ion is in great excess over copper; the concentration of chloride ion is not affected by reactions a t the electrode surface. (iv) The interference may be described by Reactions V and VI. Reaction VI is fast compared to Reaction V. (v) Diffusion coefficients of all species are equal. Effect of Diffusive Transport. Since a net chemical reaction does occur a t the electrode surface. it is necessary to consider diffusion and reaction kinetics to establish at what point equilibrium conditions are approached a t the electrode surface. The change in concentration of cupric ion a t the electrode surface can be expressed by an equation of the form:
(3) D is the diffusion flux of Cu2+from the bulk to the electrode surface, rf and rb are the forward and backward rates, respectively, of Reaction V. For equilibrium conditions to exist at the electrode surface, not only must the system be a t steady state ((dC',p)/dt) = 0, but also the diffusion transport term must be negligibly small with respect to rf and rb. The sensitivity to stirring observed for the electrode in chloride media suggests that rt and rb are not extremely fast compared t o diffusive transport. and that near-equilibrium conditions will be reached only in unstirred sol &ions after time has been allowed for a broad diffusion layer to be built up at the electrode surface. A physical picture of this condition is a layer immediately a t the electrode surface where quantitative conversion of Cu(I1) to Cu(1) can take place. This surface layer is separated from the bulk of solution by a broad diffusion layer through which transport is very slow (Figure 1). The theory presented in this paper is dependent on attainment of near-equilibrium a t the electrode surface. Hulanicki and Lewenstam ( I I ) have employed a similar diffusion layer model to describe the interference of cupric ion in measurements of cuprous ion concentrations with a chalcocite membrane electrode. Near-Equilibrium at the Electrode Surface. For the given initial bulk concentrations of divalent and monovalent copper, and under the assumption that the diffusion coefficients of all species are equal, the stoichiometry 3f Reaction
1794
ANALYTICAL CHEMISTRY, VOL. 51, NO. 11, SEPTEMBER 1979 E (Cu2’
530
I
c c u z IO-^
200-
M
1 -2
-3
-1
log
cC,.
I 0 (moles
I
1tter”i
Flgure 3. Calculated potentials of the cupric ion-selective electrode as a function of chloride concentration (Equations 6 and 9) for three different bulk concentrations of Cu2+. In the domain of the horizontal line, Cu2+ is the main species of copper at the electrode surface and the interference does not occur; in the domain of the sloping line, CuCI; is the main species of copper at the electrode surface and the interference does occur. The intersection of the lines is defined as CHa. “A” curves: ac,o = 10-8.5. ”B” curves: acuo = 10-6.5
I
W
/ I / ) /
-3
RT E = E O ’ C ~ Z + / C ~ O+ 2 F In
(G) CCU(III
(6)
where Eo’CUz+ICUo is E0CUz+!CUo corrected with the activity coefficient of Cu2+ in the given ionic medium. Case II. Interference Fully Developed. At very high chloride concentration, Reactions V and VI proceed quantitatively to the right (at the electrode surface) and the concentration of copper at the electrode surface can be approximated by C’cUci2-= ~ C W I I )
(7)
Using the concentration of C ’ C ~given ~ ~ above ~ - and equilibrium for Reaction VI, we find the concentration of Cu+ at the electrode surface,
Then we may use the concentration of Cu+ at the electrode surface to calculate the expected electrode potential
Equation 9 predicts an electrode response slope of 59 mV per
-2 -I 0 tog cC[- (moles , l i t e r - ’ )
Figure 4. Calculated potentials of the cupric ion-selective electrode as a function of chloride concentration (Equations 9 and 12) for three values of competing ligand concentration; KCuL = lo4, Ccu(II,= M, acuo =
decade change in bulk copper Concentration. Case III. Onset of Interference. Considering the transition from a medium of low chloride ion concentration to a medium of high chloride ion concentration, we may designate the chloride concentration at which the potential predicted by Equation 6 (no chloride) is equal to the potential predicted by Equation 9 (high chloride) as the critical chloride concentration, or the chloride concentration of onset of interference. Setting the right hand side of Equation 6 equal to the right hand side of Equation 9 and solving for Ccritc1-yields p i t
c1- =
Equation 10 shows that the insensitivity of the electrode to interference by chloride is improved by (i) higher bulk divalent copper concentration, (ii) lower activity of elemental copper in the mixed sulfide phase. Predicted Electrode Behavior. On the basis of Equations 6, 9, and 10, we can make predictions of electrode behavior
ANALYTICAL CHEMISTRY, VOL. 51, NO. 11, SEPTEMBER 1979
valid for unstirred solutions. These are shown in Figures 2-4. In Figure 2, electrode potential as a function of bulk copper concentration in media of low and high chloride ion concentration is shown. The response in the medium of high chloride concentration is 59 mV/decade change in bulk Cu(I1) concentration. In Figure 3 is shown the effect of chloride concentration on electrode reponse for several values of bulk Cu(I1) concentration. As is clear from Equation 10, CcritCl-is slightly greater for greater bulk Cu(I1) concentrations. In Figure 4 is shown the effect of added cupric i o n complexing ligand on the electrode potential in chloride media; cupric ion and the ligand react according to the reaction: cu2+
+ L2- = CUL
Table 11. Equilibrium Constantsa (log,, K ) Adapted from Reference 12 Ag+
+ c1-
c u + + c1Ag’ + CH3CN Cu’ t CH,CN
Go
PI
Pz
P3
0 4
-9.7 -6.7
3.4
5.3 5.5
53 51
5.1
0.7
0.8
., .
3.9
4.1
... .. .
...
...
... ... .. .
E“ (Cu2+/Cuo) = +377 mV. E” (Cu+/Cuo) = 521 mV. E’ (Ag+/Ago)= 799 mV. a
-1.0
(VI11
KCuL
1795
i
When the ligand is in great excess over copper, we may write the free concentration of cupric ion a t the electrode surface as a fraction of the total divalent copper a t the electrode surface
t
Then for low chloride concentration (Case I) the electrode potential is (cf. Equation 6)
RT E = E 0 ’ C U o / C U ~ ++ - In 2F
(:+)
-
-50
+
= Eo‘CU~+ICU~
-3.0
-2.G log
-I
G
cC I - ( m o l e s
0
G
Iiter-li
Figure 5. Computed regime of solid AgCl formation at the mixed sulfide electrode as a function of Ccuz+and Cc,-; acuO = and a,,o = 10-35
and the critical chloride concentration is (cf. Equation 10)
As is intuitively obvious, addition of a cupric ion complexing ligand stabilizes the divalent form of copper and extends the domain of the cupric ion response of the electrode to higher values of chloride ion concentration. In Figure 3 is also shown the predicted behavior of electrodes of different constructions, e.g., the sulfur rich graphite contacted electrode and the silver contacted electrode. As is shown by Equation 10, the electrode with a higher activity of elemental copper is more sensitive to the chloride interference. General Theory. In the preceding discussion only one chemical species, CuCL-, was considered as the end product of the interference reaction. However, since the actual mixed sulfide electrode does contain silver and sulfur, it is necessary to discuss the possible participation of these components in the behavior of the electrode. I t is possible that the reduction of aqueous Cu(I1) a t the electrode surface is coupled to the oxidation of silver sulfide, and that the liberated Ag(1) is stabilized by chloride complexation (cf. Reactions V and VI). Also, the formation of a new solid phase, AgC1, could occur a t the electrode surface. The importance of these reactions of silver compared to the corresponding reactions of copper depends on the relative activities ncuoand aAg0in the mixed sulfide. Soluble sulfide species could form from simple dissolution of the electrode (the dissolution process is different from the redox process described for Cu(1) formation); but unless the total copper Concentration were exceedingly low, the concentration of sulfide species would be negligible. The domain of Ccu(II)in
which sulfide species become important is the theoretical limit of detection of the electrode. In view of all these possible complications, it was deemed worthwhile to consider the hypothetical case in which every one of the possible side reactions has an effect, Le., the case in which all of these reactions at the electrode surface were at equilibrium with one another. With the approximation that the diffusion coefficients for all species are equal and the assumption that the bulk concentration of divalent copper is Ccu(II)and the bulk concentrations of other species are Ccu(Ij - C A ~ (= I ) Cs(-IIj = 0, the material balance equation a t the electrode surface becomes The chemical equilibrium computer program MINEQL (13) was used to compute equilibrium among all of the relevant reactions in Table I1 for several sets of values of q u o , a@, Ca-, and CCuz+.Although there is considerable uncertainty in the values of the stability constants involved, the chloro complexes of silver are approximately of the same stability tis the chloro complexes of copper(1). The calculated equilibrium distribution of copper and silver species for acuo = 10-s.5and U A ~ O = 10-3.5(values corresponding to a sulfur-rich electrode such as was used here) although much more complicated than that found when only the CuCL- ion was considered as a product, led to very nearly the same electrode potential values as a function of chloride concentration. Other equilibrium computations predict that the formation of the solid phase AgCl should occur when the product Ccl-. CcUz+is greater than the limit shown in Figure 5 . The precipitation of CuCl is predicted not to occur. The computed limits of precipitation are dependent on the values used for aAp and ac,o (hence, the electrode type). It is worth noting one other effect which could occur in the mixed sulfide electrode, but is not treated by this theory: the concentration polarization of the mixed sulfide phase when
1796
ANALYTICAL CHEMISTRY, VOL. 51, NO. 11, SEPTEMBER 1979
Table 111. Descriptions of Experiments” 250
I. Effect of Chloride Alone
B.
. +’ +/
ia-1
[CU2+] A
S I
0 , l o o. 3
10-4.0-10-6.0 10-3.4, 10-4.4 10-5.4
i
10-3.0-100.3
3
11. Effect of Cupric Ion Complexing Agents
[CU+21
[Cl-l 10-3.0-100.3
10-4.4
10-3.0-100.5
/
+/+’ KNOj
0 * O O P + ’
[Ligand]
10-4.4
1
/+
i
/
10-3.1 1 0 - 3 . 4 1 0 - 3 . 7 , 10-2.3’ (tartrate)
111. Effect of Cuprous Ion Complexing Agent
(Acetonitrile)
/+
[CU , + I
[CH, CNI
10-5.4
10-1 .0-100.6
/
IV. Seawater with Cupric Complexing Agent (NTA) [Cu2+1 50-150%of NTA a
1a-1
w-41
lo-’.’,
50 -6
10-4.4,
All concentrations are in moles per liter.
I
-3
-4
-5 log Cr-2.
(moles M e r - 1
Figure 6. Measured potential of graphffe contacted cupric ion-selective electrode in 2 M KNO, and 2 M KCI
a significant fraction of the electrode material is oxidized. Polarization is most likely t o occur in stirred solutions or in solutions of high cupric ion and chloride ion activity when the redox processes at the electrode proceed most rapidly. EXPERIMENTAL Reagents. The standard 0.1 M Cu2+solution was made by dissolution of copper shot in nitric acid; the concentrated KCl and KNOBsolutions were filtered through a 0.22-fim Nucleopore filter prior to use: the other chemicals were used without prior treatment. The acetonitrile was Spectro from Eastman Organics, and the ethanol was 95%. All solutions were made with water which had been distilled, deionized, and redistilled. Apparatus. All measurements were made a t 25.0 f 0.1 “C in a thermostated cell isolated from the atmosphere. The cell was shielded from light because of observed photosensitivity in cupric ion-selective electrodes. Solutions were purged of oxygen with N2 (or 1%COSin N,) prior to introduction of the cupric ion selective electrode, after which nitrogen (or 1%COPin NZ)was passed over the solution a t a very low flow rate to minimize gas flow induced circulation of the sample solution. Potential measurements were made with the Orion 801A or Orion 701 Digital pH meters against an Orion Model 90-01 Double Junction Electrode. Potential readings were taken in unstirred solutions when the rate of potential drift became less than 1 mV/h (as measured on a Valtec Model 1024 X-Y-t recorder); equilibration times greater than 2 h were not uncommon. Two cupric ion-selective electrodes were used: the Radiometer Selectrode Model F-3000 and the Orion Model 94-29A cupric ion-selective electrode. The Radiometer electrode was used in all of the experiments except one because of the ease with which the chloride-poisoned electroactive element can be restored. Procedure. Four experiments were carried out to measure the response of the Radiometer cupric ion-selective electrode to various total concentrations of cupric ion, cuprous (and argentous) complexing agents, and cupric complexing agents. The experiments are described in Table 111. An additional experiment similar to I was carried out with the Orion electrode. In each set the solutions were brought to a constant ionic strength and buffered: in Experiments I and 11, an appropriate amount of KNO, was added to give a total ionic strength of 2.0 equiv L-’. In Experiment 111, appropriate amounts of 95% ethanol were added to prevent changes in the dielectric constant of the medium. Experiments 1-111 were buffered at pH 4.7 by a M acetic acid/W3 M sodium acetate buffer. The medium for Experiment IV was water collected from Buzzards Bay, Mass., filtered successively through a 0.45-fim and a 0.22-fim Nucleopore filter, refrigerated, and used within 48 h of collection. The medium was buffered at pH 7.0 by the natural alkalinity and 170C 0 2 in N, passed over the top of the solution.
-
:
I
+\
I
\
Flgure 7. Measured potential of cupric ion-selective electrode as a function of chloride concentration at various bulk cupric ion concentrations. “A” curves: graphite contacted electrode. “B” curve: silver contacted electrode RESULTS AND DISCUSSION T h e response of the electrode in chloride media was investigated in Experiment I. In Figure 6 are seen the 29 mV/decade electrode response slope in a chloride free medium and the -60 mvldecade slope in 2 M KC1. In Figure 7 is seen the effect on electrode response of increasing chloride concentration for three values of total copper. The values of CcritCi-found from Figure 7 agree reasonably well with the predicted values (Table IV), and show that the electrode is less sensitive t o interference a t higher bulk copper concentrations. The effect on electrode response of a cupric ion complexing agent added to the chloride medium was seen in Experiment I1 (Figure 8). Stabilization of cupric ion by complexation favors the cupric ion response over the cuprous response, and greater chloride concentrations are necessary to cause interference. In order to confirm the hypothesis that the interference is the result of cuprous or argentous complexation, another cuprous/argentous complexing agent acetonitrile, was employed in Experiment 111. The response of t h e electrode in
ANALYTICAL CHEMISTRY, VOL. 51, NO. 11, SEPTEMBER 1979
250
1797
t I
c 0 *
a 3
C’
+
N o Ligand
0
Tartrate
10-23M
o
Oxalate
I O - ~ ~ M
A
oxalate Oxalate
M
IO-^,^
M
I Io-‘
10-2
I
IO0 -I
C c l - ( m o l e s ‘liter 1 Figure 8. Measured potential of graphite contacted cupric ion selective electrode as M ligands. CCu2+=
Table IV. Comparison of Experimental and Theoretical Values of CCdtcl- and C C n t C H , c ~
a function of chloride concentration in the presence of competing 250
A. Experiment I
experimental calculateda computedb
Ccu2
ccntcl-
ccnt
10’0.2
10-0.1
10-4.4
100.0 10-0.3
10-0.1
10-0.3
10-5.4
10-0.6
10.0.3
10-0.5
10-3.4
c1-
ccr’tc1-
B. Experiment I1 Ccu2
+
10-4.4
exFrimenta1 calculatedc computedd ccrit CCntCH,CN centCH ,CN c1100.3
loo.*
a Calculated from the intersection of the lines given by Equations 6 and 9 ; aCuo = 1 0 - 8 . 5 . The electrode potential was computed as a function of chloride from all of the pertinent reactions in Table I; the potential VS. chloride concentration curve was plotted and the value of cCritclobtained by extrapolation;^^,^ = 1 0 - 8 . 5aAg0 , = 10-3.5. Calculated from the intersection of the lines by Equation 6, and Equation 9 written for the reaction Cu2++ Cuo + 6CH,CN = 2 Cu(CH,CN):; aCuO = 1 0 - 8 . 5 . Computed as in b , except the reactions involving CH,CN were considered.
acetonitrile was similar to the response in chloride solutions, and the Ccrit~H3CN occurred approximately where predicted by theory (Table IV). Experiment I11 is interesting from another point of view. In the previous experiments in chloride media, the relative stabilities of the cuprous and argentous complexes were approximately equal, and it was impossible to tell whether aqueous copper, or aqueous silver, or both, were released from the electrode as cupric ion was reduced. However, since the cuprous acetonitrile complexes are several orders of magnitude more stable than the argentous aceto-
~
3:
075
~
0
25
153
F r 3 c . c r o( LTA T t i o t e d
Figure 9. Measured potential of graphite contacted cupric ion-selective electrode in seawater spiked with NTA and titrated with Cu2+
nitrile complexes, (12),and the CCntCHaCN occurs approximately where predicted for the formation of cuprous acetonitrile complex, we may conclude that in the acetonitrile medium Cu(1) is released. Experiment IV was performed to evaluate the usefulness of cupric compleximetric titrations performed in seawater for the purpose of determining “natural chelating capacity”. Samples of seawater spiked with different concentrations of NTA (a relatively strong chelator, probably stronger than most natural chelators) were titrated with Cu(1I) from 50% to 150% M NTA the effect of the total NTA (Figure 9). In 2 X of chloride is negligible; before the equivalence point, copper speciation is controlled by NTA (see Equation 13 or Figure 4 for the effect of KCuLon the CCntCl ), and after the equivalence point the concentration of cupric ion is high enough that
1798
ANALYTICAL CHEMISTRY, VOL. 51, NO. 11, SEPTEMBER 1979
chloride does not interfere (see Equation 10 or Figure 3 for effect of Ccu(II)on the Ccr'tcI-). However, in 8 X 10 M NTA the electrode response is completely dominated by the chloride interference and no titration breakpoint is seen. Actually this result could have been predicted from Equation 13 with Ccr't~l set to the chloride concentration of seawater, 0.5 M. One experiment was carried out to compare the behavior in chloride media of a cupric ion selective electrode of a different construction-hence a different acuo and aha. Since the acuo and aAgoof the silver contacted electrode are quite different from those of the sulfur-rich graphite contacted electrode (Table I), the theoretical speciation, electrode response, and precipitation regime calculated previously for the latter electrode do not apply to the former. However, one may still predict that, because the activity of elemental copper and silver are much higher in the silver contacted electrode. the interference begins a t lower chloride concentrations. This expectation is borne out in Figure 9; the experimental value of CCntC1= M at M total copper is almost an order of magnitude lower than for the other electrode. Finally, mention should be made of those experiments to which the theoretical predictions of Figure 2-5 were not seen, Le., experiments in which formation of solid AgCl at the electrode or polarization of the electrode did occur. At concentrations of copper and chloride greater than the limit of the precipitation region in Figure 7 , electrode behavior was observed which was not continuous with behavior outside of this region. Whether this behavior should be attributed to electrode polarization of to AgCl (s) formation or to both, is uncertain.
CONCLUSION The use of the cupric ion-selective electrode in chloride media (seawater) has been suggested for the monitoring of ionic cupric copper concentration ( 4 ) and the assessment of cupric ion activity with respect to organic complexation ( 3 , 5). These practices will be interpreted in light of the theory presented in this paper. The cupric ion-selective electrode has formally two modes of response: 29 mV/decade in the absence of chloride, and 59 mV/decade in media of high chloride concentration. If the electrode is in the 29 mV/decade response domain, cupric ion activity may be measured normally as in a nonchloride medium. If the electrode is in the 59 mV/decade response domain, it is impossible to measure cupric ion activity; only the total copper concentration, as cuprous ion, can be measured and this requires a recalibration of the electrode in the chloride medium. If the electrode response is in the transition zone between the two response domains, it is very difficult to interpret electrode potential measurements a t all. In seawater (Ca- = 0.5 M), in the absence of cupric ion complexation, the transition from the 59 mV/decade response to the 29 mV/decade response occurs at approximately 4 X M total cupric ion for the sulfur-rich graphite contacted electrode. Thus for bulk copper concentrations somewhat less M only total copper concentration (as cuprous than 4 X ion a t the electrode surface) can be measured, but for bulk copper concentrations somewhat greater than 4 X 10-5 M, cupric ion activity can be measured. On the other hand. the silver contacted electrode remains in the 59 mV/decade response range for all total bulk concentrations. The presence of a cupric ion complexing ligand will affect the results given above only if the ligand is able to out-compete chloride ion for control of copper at the electrode surface. As shown in Figure 8, this does not occur with 5 X 10% M NTA a t pH 7; in general it is unlikely that anything but an exceptionally strong cupric complexing ligand will bring the
electrode out of the 59 mV/decade response domain. If the electrode remains in the 59 mV/decade response domain, it is impossible to measure cupric ion activity directly. The remarks above pertain to unstirred solutions which have been allowed to approach equilibrium in the region immediately at the electrode surface. In stirred solutions the situation is much more complicated; the electrode potential is dependent on the rate of material transport to the electrode surface, and thus highly dependent on stirring rate. Also, the rate of reduction of divalent copper a t the electrode surface is much higher than in stirred solutions, and the mixed sulfide is much more likely to become polarized and exhibit an apparent shift in E". Oglesby, Duer, and Millero (14) in a study of the effect of chloride ion and ionic strength on the cupric ion selective electrode observed that in stirred solutions of different NaCl concentrations, the response slope of the ion-selectiveelectrode (rnV/decade CcucII,)was greater a t higher chloride concentrations. This result can readily be interpreted in terms of the current model. Since the solution was being stirred, the system never reached chemical equilibrium, but only steady state. The term D in Equation 3 was not small with respect to the reaction rates rf and rh, and the concentration of cupric ion at the electrode surface was being held artifically high by diffusive transport from the bulk of solution. Hence the response slope was closer to the "no interference" value of 29 mV/decade. As the concentration of chloride ion was increased, the reduction of cupric ion at the electrode surface became more rapid, and the diffusive transport of Cu2+to the electrode surface became relatively less significant, and the response slope approached the "fully developed chloride interference" value of 59 mV/decade Cu(1I). In addition, Oglesby et al. noticed a slow increase in E" of the electrode after use in chloride; this increase in E" is consistent with the depletion of silver and/or copper from the mixed sulfide. This work exemplifies the difficulties of working in stirred solutions of high chloride concentration with the cupric ion selective electrode.
ACKNOWLEDGMENT The authors thank John G. Rueter, Jr., for performing several of the experiments. LITERATURE CITED (1) J. W. Ross, in "Ion Selective Electrodes", R. Durst, Ed., Chapter 2, Nat. Bur. Stand., (U.S.A.) Spec. Pub/. 314 (1969). 12) E. H. Hansen, C. G. Lamm, and J. Ruzicka, Anal. Chim. Acta, 59, 403 (1972). (3) W. Sunda, FhD. Thesis, Massachusetts Instlute of Technology and Woods Hole Oceanographic Institution, 1975. (4) R . Jasinski, I.Trachtenberg. and D. Andrychuk, Anal. Chem., 46, 364 (1374). (5) P . M. Williams and R . J. Baldwin, Mar. Sci. Comm., 2 , 161 (1976). (6) R. P. Buck and V. R. Shepard, Anal. Chem., 46, 2097 (1974). (7) M. Koebel, Anal. Chem., 46, 1559 (1974). (8) Orion Research Incorporated, Cambridge, Mass., "Analytical Methods Guide", 1972. (9) J. Westail, W.D. Thesis, Departmentof Chemisby, Massachusetts Institute of Technology, Cambridge, Mass., 1977. (10) J. Westall, F. Morel, D. N. Hume, manuscript in preparation. (11) A . Hulanicki and A. Lewenstam, Talanta, 23, 661 (1976). (12) L. G. Sillen and A. E. Martell, "Stability Constants", Special Publications No. 17 and No. 25, The Chemical Society, London, 1964 and 1971. (13) J. C. Westall, J. L. Zachary, and F. M. M. Morel, "MINEQL, A Computer Program fOc the Caicuhtbn of Chemical EquilibriumSpeciation of Aqueous Systems". Technical Note No. 18, Ralph M. Parsons Laboratory, Massachusetts Institute of Technology, Cambridge. Mass., 1975. (14) G. B. Oglesby, W. C. Duer, and F. J. Millero, Anal. Chem., 49, 877 (1977).
REL'EI~XD for review March 28, 1979. Accepted June 25, 1979. This work supported by National Science Foundation Grant No. DES 75-10523 and Environmental Protection Agency Grant No. R803738.
