Over a limited concentration range, C1- interferes with the determination of phosphate by forming CuC1. When I, is 500 pA cm-2, a CuCl film is formed a t C1- concentrations greater than 0.5 ppm. A typical example is shown in Figure 4. The reduction of CuCl occurs a t a potential coincident with 71 for the phosphate stripping so the determination must be based upon 72; however, 72 is decreased from 15 to 13 sec in this mixture which contains 10 ppm P043-and 1.5 ppm C1-. The decrease in 72 is significant when the C1- concentration is greater than 0.5 ppm and exceeds 10% of the phosphate concentration. Apparently the presence of CuCl interferes with the nucleation and/or growth of the phosphate film (11). The stripping method was compared to the spectrophotometric procedure for the tietermination of phosphate in a set of five identical grab samples of lake water. The
samples were stabilized by addition of 1% (volume) chloroform and were filtered through 0.45-p Millipore membrane prior to use. The mean and standard deviation were 52 f 20 ppb by the electrochemical method and 58 f 15 ppb by spectrophotometry. The relatively high standard deviation of the former is probably the result of surface active agents in the samples. In the present state, the method is primarily useful for determinations in turbid or colored samples where the spectrophotometric procedure is subject to interference. Received for review August 22, 1973. Accepted November 5, 1973. This research was supported by the United States Department of the Interior Office of Water Resources Research Allotment Grant A-041-ILL.
Potentiometric Measurement of Copper in Seawater with Ion Selective Electrodes Raymond Jasinski, Isaac Trachtenberg, and Dmetro Andrychuk Texas Instruments Incorporated, P. 0. Box 5936, Dallas, Texas 75222
The data reported confirm the thesis that, with proper precautions in sample handling and measurement, the ionic cupric copper concentration of seawater can be monitored with ion-selective electrodes. The measured cupric ion content is identical to the total soluble cupric copper content: when the electrode response is calibrated directly in seawater, to account for the inorganic complex ion distribution; when organic chelating agents with high stability constants relative to cupric ion are absent; and when the proper Nernstian slope is used in calculation. At and below 1 ppb copper, account must be taken of the injection of soluble copper (or silver) into solution by the sensor electrode. Seawater samples from three different sources were studied. The analysis of an open ocean water was straightforward. However, analysis of the near shore waters showed abnormalities which can be explained by the presence of small quantities of chelating agents.
Recent advances in analytical electrochemistry have resulted in commercially available ion selective electrodes which are reasonably specific to some of the trace metal cations of interest in the characterization of seawaters ( I ) , and which have sensitivities comparable to the concentrations of these cations in seawater (nominally 10-aM) (2). Heretofore, it has not been considered possible to measure copper directly in seawater by this potentiometric technique (3, 4 ) . Considerable instability in the measured potentials as well as a general lack of sensitivity a t the parts-per-billion concentration level have been reported ( 4 ) . Recently, the use of such electrodes for the measure( 1 ) J. Ross, in "Ion Selective Electrodes," R. Durst, Ed., Chapter 2, Nat. Bur. Stand. ( U . S . ) Spec. Pubi. 314 (1969). (2) J. Riley and G. Skirrow, "Chemical Oceanography," Vol. 1 , Academic Press, New York. N . Y . , 1965, p 164. (3) T. Warner, Mar. Tech. SOC.,-6thAnn. Preprints, 2, 1495 (1970). (4) R. Durst, Paper No. 330, 142nd National Meeting, Electrochemical
Society, Miami Beach, Fla., October 1972.
364
ANALYTICAL CHEMISTRY, VOL. 46, NO. 3, MARCH 1974
ment of copper in natural waters has been described ( 5 ) . However the buffer system employed is not compatible with seawater, resulting in the loss of soluble cupric ion from the sample, presumably by coprecipitation with the alkaline earth fluorides. The data presented in the present paper imply that many of the problems reported for measurements of copper in seawater have been due to alteration of the cupric ion content of the sample by the containers (6) and, in some cases, by the electrodes themselves. This paper will show that, with the proper precautions, direct measurements of copper in seawater can be made with ion selective electrodes. In fact, it is possible thus to study the solution contaminating processes with these electrodes. An evaluation was then made of the response of the commercial cupric ion selective electrode in three different seawaters, in terms of intrinsic electrode sensitivity (limit-ofdetection); electrode dynamics ( e . g . , response time, stirring dependencies. electrode stability); naturally occurring chelating agents and possible interfering metal cations.
EXPERIMENTAL Orion Model 94-29 cupric ion selective electrodes were used in this study. Most as-received electrodes had to be equilibrated in seawater before potentials representative of the soluble cupric ion concentration could be obtained. The equilibration time period did vary with electrode. and in some cases was of the order of 24 hours. The Orion double junction Ag/AgCl electrode was used as reference with 0.1N K N 0 3 as the outer filling solution. Glass barreled. single junction electrodes gave rise to a slow, continuous increase in the potential measured by the copper electrode, presumably due to the slow contamination of the test solution by soluble silver from the reference electrode. Potentials were readout to within 0.1 mV on a Corning Model 101 digital voltmeter. All measurements were made in polyethylene and Teflon beakers and bottles to minimize contamination of the seawater sample. As a precaution, all containers were first acid washed; on oc(5) M . Smith and S. Manahan. Ana/. Chem., 45,836 (1973). (6) J. Alexander and E, Corcoran, Limnoi. Oceanogr.. 12, 236 (1967).
casion it was found that a new polyethylene bottle would otherwise inject an impurity into the test solution which would generate an increase in the potential of the ion selective electrode. In the p H studies reported, the quantities of acid required to reach the p H values listed were determined with a glass p H electrode on auxiliary samples and the same quantities of acid added to the test samples. It was not possible, otherwise, to avoid contamination of the test solution by the particular glass p H electrode used. The analytical method was developed with samples of seawater taken 5 miles off Port Isabel, Texas; these waters were assumed to be representative of relatively uncontaminated open-ocean water. The technique was then applied to nearshore seawaters taken 15 miles south of and 1 mile off Freeport, Texas, and in outer San Diego harbor. The inorganic carbon content of all three seawater samples was 21 f 1 ppm. Seawater was also simulated with a 0.5M NaCl solution, p H 8, t o obtain solutions having lower residual copper ion contents. The electrochemistry of cupric ion in this solution relative to that in seawater will be discussed. An atomic absorption spectrometric method similar to those described in References 7, 8, and 9 was used as an independent method of analyzing for the copper content of the water samples. In this technique, copper was extracted from water a t a 40:l ratio into a solution of 2% oxine in ethyl propionate, which was then aspirated into an acetylene/air flame. Calibration was achieved by spiking, both pre-extracted and as-received, seawater with known amounts of copper. Identical results were obtained when both glass and Teflon bottles were used as the extraction vessels even though the absorption of copper onto the glass was indicated by the ion selective electrode. Evidently the extractant will remove such absorbed copper. The standard deviation of this method was established t o be zk0.06 ppb between 0.2 and 10 ppb copper. It is to be expected that the AA method will measure all the inorganic copper and will extract at least some of the organic chelate copper, so that copper analyses so obtained represent the maximum copper concentration accessible to the ion selective electrode.
O
-80
‘
0.4
I
I
I
I
I
I
1.0
/
2.0
I
1
I
4.0
,
I
,
_I
,
10.0
20.0
p p b Cut* ADDED TO SEA WATER
Figure 1. Shown is the ’response of a cupric ion selective electrode to the additions of cupric chloride to the Port Isabel seawater. Also plotted are the electrode potentials vs. the total copper present in the sample-Le., the added copper plus an initial copper Concentration of 1.6 ppb
RESULTS AND DISCUSSION Plotted in Figure 1 are the potentials of one Orion electrode, immersed in a sample of the Port Isabel seawater, to which were added increments of cupric chloride-KC1 solution. Also plotted are the same potentials but against the concentration of copper added plus a residual (initial) copper concentration of 1.6 ppb. In this second plot, a slope of 52 mV/decade change in copper concentration is obtained; replicate experiments with the same electrode on the same batch of seawater yielded slopes identical to within f 2 mV per decade copper. The validity of this value for the slope at the lower copper concentration will be discussed. It is difficult to reconcile such reproducible measurements with the slope of 29 mV/decade anticipated from the Nernst equation and a divalent ion [ e . g . (IO)]. Until the sensing mechanism is clarified for the Orion type electrode in seawater, the measured slope can be regarded as an empirical parameter relating measured potential to the logarithm of the cupric ion concentration. Since this slope is reproducible, it is nevertheless of practical utility in studying the soluble copper concentration in seawater. For brevity, this parameter will be designated in this paper as the “Nernstian slope.” Most of the inorganic anions normally found in seawater a t pH 8, particularly OH-, C032-, C1-, and S 0 4 2 - , form complex ions with cupric ion. However, the formation constants are not sufficiently high to remove completely all cupric ion from solution ( 1 1 , 12), no strong polynuclear (7) J. Culp, R. Windham, and R . Whealy, Anal. Chem., 43, 1321 (1971). (8) S. Sachdev and P. West, Anal. Chim. Acta., 44, 301 (1969). (9) J. Jones and R. E d d y , Anal. Chim. Acta., 43, 165 (1968). (10) P. Ruetschi and R . Amlie, J. Electrochem. SOC.,112,665 (1965). (11) A. Zirino and S. Yamamoto, Limnol. Oceanogr., 17, 661 (1972).
r
complexes exist, and the total concentration of these ions is relatively constant and large compared to the copper concentration. Hence, a proportionality should exist between the measured free cupric ion concentration and the total copper concentration actually present in solution, and this was demonstrated by the response curve shown. This copper is designated as the inorganic soluble cupric copper and is the quantity measured. There are at least four general reactions whereby substances found in seawater itself can react with such copper and render it electrochemically inactive: 1) Reduction of C U + ~to Cu+, 2) precipitation of insoluble cupric species, 3) absorption of C U + ~ onto particulate matter, and 4) complex ion formation with organic chelating agents. Process 1 is unlikely in oxygen-containing seawater since dissolved air is sufficient to oxidize cuprous ion to cupric. Also, unlikely is process 2 since most common cupric salts are soluble a t the ppb concentrations of copper normally found in seawater. The presence of sulfide ion, which would precipitate copper in this concentration range, is unlikely in oxygen-containing solutions although it is likely in anaerobic waters. Reaction 3 is unlikely in the open ocean where the concentration of particulates is low. Organic chelating agents can be found in some seawaters (2, 6, 13) and, if the formation constant with cupric ions is sufficiently high, and their concentration is sufficiently high, there would be complete removal of cupric ion from solution. The electrode behavior signaling such a condition was determined by the following experiment. An ion selective electrode was immersed in 1M KC1, pH 5.5, containing approximately 0.8 ppb Cu2+, as determined by AA analysis. EDTA (ethylenediaminetetraacetic acid disodium salt) was added to make the solution 10F7M, sufficient to complex completely 6.3 ppb cupric copper. After 16-minute exposure to this solution, the electrode drift had decreased to less than 0.2 mV/min, and CuClz was added to make the solution 1.4 ppb total copper; after 34 minutes, CuC12 was again added to the solution in sufficient amount t o make it 2.0 ppb total copper. Table I shows the potential-time dependence observed. Thus, the transfer of an electrode from a solution of ppb Sillen and A. Martell, “Stability Constants of Metal Ion Complexes,” 2nd ed., Chemical Society London, Spec. Publ., 1964. (13) P. West and R. Bustin, Environ. Lett., 3, 247 (1972). (12) L.
ANALYTICAL CHEMISTRY, VOL. 46, NO. 3, MARCH 1974
365
Table I. Effect of EDTA on Potential Time, minutes
Potential, mV
10
-107.4 - 113.9 -117.0 -119.5
16
-120.5
0
r
4 16 added copper 17
19 22
-118.1 - 118.8 -121.1
34 34 added copper
-123.4
35
36
-121.8 - 121.9
38
-123.2
cupric ion to a solution of a strong chelating agent results in a slow drift of the electrode potential and to highly negative potentials-not a t all representative of the total copper present in the solution. EDTA reacts with the copper originally present in the solution, which may be a slow process at these concentration levels, as well as reacting with the CuS dissolving off the electrode surface (14). Both processes would result in a decreasing electrode potential. The addition of cupric ion results in only a temporary shift to more positive potentials, which is consistent with a slow reaction of EDTA and parts-per-billion concentrations of copper. Calculation of the residual copper concentration present from the maximum shift in potential on addition of copper indicates abnormally high residual copper contents-in this case: 7 ppb rather than the nominal 0.8 ppb actually present. That this insensitivity to copper was not due to a spurious interaction of the spike with the electrodes or the container wall was shown by spiking the same solution contained in a Teflon beaker without electrodes, mixing by agitating the beaker, and transferring the electrodes from another portion of the spiked solution. There was no difference in potential between the solutions in the two beakers and the original solution, indicating no increase in the free copper concentration from the spike. It is also to be expected that. in such solutions, a plot of potential change us. added copper would be sigmoidal in form, as in a titration curve, until the concentration of chelate is exceeded (which could be a method of estimating the concentration of such species). Examples of these four effects in actual seawater solutions will be given below. Those metal cations (other than copper) which could generate a response from the ion selective electrode are ions which form sulfides more insoluble than CuS-e.g., Fe3+ and Agl*. The soluble ferric iron content of seawater has been estimated to be no more than 0.05 ppb ( 1 5 ) , and, as such, should not represent a significant interference to the measurement of cupric ion. To test this conclusion, sufficient ferric nitrate, pH 1.7, was added to Port Isabel seawater to make the concentration of 1 ppb total iron, assuming complete dissolution. An increase of 1.4 mV was observed for the potential of the cupric ion selective electrode. The addition of a more concentrated ferric nitrate solution, sufficient to make the solution 100 ppb iron, resulted in the same shift in potential. It is thus to be concluded that the interference by ferric iron, if any, is minor. The small shift in potential observed could be explained in terms of slow nucleation of ferric hydroxide (or phosphate ( I s ) ,or a small change in the pH of the test (14) G. Johansson and K. Edstrom, Talanfa, 19, 1623 (1972). (15) G . Lewis and E. Goldberg, J . Mar. Res., 13, 183 (1954).
366
sample due to the use of pH 1.7 supporting electrolyte for the iron spike. Additions of silver ion were made to 0.5M NaC1, pH 8, over the range 10-8 to 10-5M-i.e., the total solubility of silver species in the electrolyte (16). Electrode responses were similar to those observed for additions of cupric ion, although the “Nernstian slope” did appear to be higher for silver-i.e., about 54 mV/decade rather than the 50 mV/d.ecade slope obtained for Cu2+ additions. Above 10-5M Agl+, there was no change in potential on increasing additions, Additions of cupric ion to solutions spiked with silver generated the expected response. Thus, silver can constitute a possible interference for the measurement of copper in seawater. However, the normal Concentration of silver in seawater is an order of magnitude less than that of copper ( 2 ) , and thus, silver interference should not be a practical problem. When this situation is not the case, correction would have to be made. The next factor considered in developing the potentiometric method for measuring Cu2+ in seawater was the possible dependency of the “Nernstian slope” on copper ion concentrations. Figure 2 is a composite plot of the responses of one ion selective electrode in three different chloride solutions, as a function of total copper content over the concentration range of 2 to 100,000 ppb. The points designated as 0 were measured in 1M KCl, pH 2; the points designated A were taken in 0.5M NaCl, pH 8; and the points designated as 0 were taken in the Port Isabel seawater. All copper concentrations were corrected for the residual copper concentrations in the solutions, as determined by AA analysis. Consider first the performance of the electrode a t high copper concentrations (>60 ppm). This lack of sensitivity to changes in the copper concentration (the plateau extends at least another order of magnitude higher than shown) was most likely due to the reaction: AgzS CuClz z CuS 2 AgCl, which becomes important a t high copper and chloride concentrations; the resulting electrode is then sensitive to chloride ion uia AgCl (1, 17) rather than cupric ion. Below 1000 ppb copper, the electrode behavior exhibited a response proportional to log [Cu2i-]; between 100 and 10 ppb, the “Nernstian slope” was 50 mV/decade of change in Cu2+ concentration, rather than 29 mV/decade observed in nitrate and sulfate media. The apparent independence of the actual potentials to the composition of the supporting electrolyte was most likely fortuitous. Also, shown in Figure 2 is the calibration curve taken for the same electrodes in the same seawater sample-but three weeks later. The prime conclusions to be drawn from Figure 2 are the following: the same “Nernstian slope” is found in all three solutions; and, in particular, the slope is independent of cupric ion concentration above 2 ppb copper. This identity of the “Nernstian slope” for the Port Isabel water, the O.5N NaCl solution and the KC1 solution implies that the abnormal 50 mV/decade slope is a function of the electrode material and the high chloride ion concentration and not due to the presence of small quantities of chelating agents in the solution. The effect of such chelating agents on the slope will be discussed below in terms of the near shore seawaters. Extrapolating the curves of Figure 2 to copper concentrations below 1 and 2 ppb copper in the salt and in the seawater solutions gave rise to a number of inconsistencies which could be
+
+
(16) J . N. Butler, D. Cogley. and W. Zurosky, J . Electrochem. SOC., 115,445 (1968). (17) J. Koryta, Anal. Chim. Acta., 61, 329 (1972).
ANALYTICAL CHEMISTRY, VOL. 46, NO. 3, MARCH 1974
@r
b
STIRRING RATE 720 r ~ m
y1,
101 = I m pH2 l4l=OSmNoCI,pHB ( 0 )=
SEAWATER, pHE -75
I
10
100
1000
10,000
100,000
Cu“Ippb1
Figure 2. Plotted are the response of the cupric ion selective electrode vs. the total copper content in seawater, in 0.5M NaCI, and in 1M KCI, pH 2. The upper curve represents the same measurement on the same seawater taken three weeks later
explained in terms of a dependence of the “Nernstian slope” on cupric ion concentrations. For example, calculating the residual copper concentration from the potential change resulting from the first addition of 0.63 ppb copper to the 0.5M NaCl solution and a slope of 50 mV/ decade gave a value of 0.4 ppb; calculation from the second addition gave 0.3 ppb; and calculation from the third and subsequent additions gave 0.24 ppb with a spread of k0.04 ppb. Atomic absorption spectrometric analysis of the stock 0.5M NaCl solution indicated a residual copper concentration of 0.22 & 0.12 ppb. The possibility of interference from trace quantities of residual resin in the deionized water is unlikely in view of the results whereby additions of water gave a small increase in potential. These inconsistencies were somewhat more pronounced with the samples of a second batch of Port Isabel seawater analyzed in the same manner. The potential change on the first addition of 0.6 ppb copper indicated a residual copper concentration of 1.8 ppb; the fourth, fifth, and sixth additions indicated 1.3 ppb. Atomic absorption spectrometric analysis indicated 0.7 f 0.15 ppb and definitely not 1.8 ppb. Assuming the AA results to be correct and plotting the potentials against the residual concentrations plus the concentrations of added copper resulted in a concentration-dependent slope below 1 ppb total copper. The form of this curve resembled that observed when the concentration of the ion being sensed approaches the solubility product concentration of the electrode material (18). Although, potentiometric analysis is not prohibited in such concentration ranges, the correct slope obviously must be used. If electrode solubility is indeed involved in determining this change in slope with copper concentration, the potential readings and the slope should be stirring-dependent in this concentration range. As the stirring rate in a solution of fixed copper concentration is increased, the potential should decrease and the “Nernstian slope” should increase for the following reasons: the electrode senses only that copper (or silver) in the vicinity of the electrode surface, regardless of its source; and, with stirring, the small quantities of copper injected into the solution by the electrode should be swept away from the electrode surface and the potential should decrease to a value more representative of the lower bulk copper concentration. Listed in Table I1 are the potentials of three ISE’s in unspiked samples of Port Isabel seawater a t three stirring rates. Shown in Figure 3 are the effects of stirring rate on the curves re(18) J. Ross, in “Ion Selective Electrodes,” R. Durst, Ed., Mat. Bur. Stand. ( U . S . ) Spec. P u b / . . 314, 67 (1969).
l
5
7
I
i
2
-
3 4 5
7
--
l
5
IO
7
1
A l --Lr
2 3 4 5 7 1 0
Cu** I P P ~ I
Figure 3. Cupric ion selective electrodevresponse in seawater as
a function of added and total copper at two stirring rates T a b l e 11. P o t e n t i a l us. Stirring R a t e Potential for electrodes, mV
Stirring rate, rpm
1
2
3
275 600 720
-69.0 -72 6 -75.0
.. -120.6 -123.7
-119 7 -123.7
...
Table 111. Comparison of Electrodes Copper added, ppb
0 .o
0.6 1.3 1.9 3.2 4.4 5.7
Potential of electrodes, mV 1
2
-59.8 -55.9 -51.8 -48.4 -43.2 -38.4 -36.3
-112.0 -108.7 -105.2 -102.3 -96.4 -91.8 -90.0
3
-108.3 -104.7 - 100.9 -97.8 -92.6 -88.5 -86.4
lating potential and copper concentration. The stock solution was the Port Isabel seawater; the stirring rates were 275 rpm (curve a ) and 720 rpm (curve b ) . Potentials are plotted us. added copper and LIS. added copper plus an 0.7 ppb residual. In the case of the data taken at 275 rpm. the slope of the linear portion of the curve was 37 mV/decade in this concentration range; a t 720 rpm, the slope was 50 mV/ decade. From these data then (Table I1 and Figure 3), it is quite likely that these sensor electrodes have a finite, though small, solubility or corrosion rate in these solutions, resulting in a changing effective “Kernstian slope” below 2 ppb total copper. The remaining factors considered were potential response as a function of the particular electrode used. response times, and electrode stability. The reproducibility in potential response from electrode to electrode for a fixed amount of copper in solution is illustrated by the data in Table 111. Shown are the measured potentials for three of the Orion copper ion selective electrodes US. the same double-junction reference electrode, in the Port Isabel seawater to which were added increasing amounts of cupric ion. As might be expected, there were significant differences between electrodes in the absolute values of potentials for a given amount of copper in solution-Le., differences in Eo. There were also differences in the total potential changes for a given change in the copper ion concentration-ie., different “Nernstian slopes.” Seven electrodes in all were evaluated in this study. The steady state “Nernstian slopes” ranged from 44 to 5 2 mV/decade; none showed slopes approximately 30 mV/decade. A N A L Y T I C A L C H E M I S T R Y , VOL. 46, NO. 3, M A R C H 1974
367
L
0
-
-10
-i -
-30 -40
-60
2
_^_NIo
A-
/
-70 -80
0850
1000
1200 0800
1600
1400
1200
0800
1200
1600
TIME Of DAY
Figure 4. Steady state potentials of three cupric ion selective electrodes against the same double junction reference electrode, as a function of the time of day. Also plotted are the potential differences ( € 3 - E l ) between two of the cupric ion selective electrodes -170
-50
I -150
r
1 -130
1 -120
I
I
-90
-70
POTENTIAL lmv)
Figure 6. Potential response as a function of pH for the Freeport seawater
Table IV. Response Times for C u ISE Sensors, Adding 0.6 ppb to 6.0 ppb Copper
-75
-E
-100
1
i
4:
!
c
z
w
c
0
a
1
Time, min
Total potential change, R of electrodes 1
2
3
1 4 6 8
41 76 85
34
90
10
95 100
86 97
45 74 79 87 92
100
100
12
65 77
-125 tI
i -150
1 1
04
I
l
!
07
l
I
I
I
I 3
1
I
4
l
l
l
l
7
l
l
10
Cut* ADDED (ppb)
Figure 5. Potential response on the addition of cupric ion to the Freeport seawater, at pH 3 and at pH 8
Shown in Table IV are the percentages of total potential change for three ion selective electrodes as a function of time after the addition of a 0.63 ppb copper spike to seawater previously spiked with 6.0 ppb copper. The use of this higher ambient copper concentration avoids complications from changing “Nernstian slopes,” possible contamination from organic chelating agents, and absorption effects. Obviously, changes in copper concentration were detected almost immediately after addition of the spike; however, a t least 10 minutes were required to reach steady-state. Also, among the electrodes, there were some differences, in response times. There was indication also that the response times depended on the chemistry of the particular salt solution used. After the addition of 0.6 ppb copper to a 0.5M NaCl solution containing 1 ppb copper, 368
for example, 60% of the final total signal was achieved in 1 minute, 94% in 5 minutes, and 99% in 8 minutes. More extreme examples of this effect of solution composition on response time were found in waters containing significant amounts of organic chelating agents. Data on the long-term stability of the electrode system are shown in Figure 4. Plotted as a function of time of day in the same seawater solution are the potential readings of three ion selective electrodes us. the same double-junction reference electrode (points 0 , A, and A ) and the differences between the potentiais of electrode 3 and electrode 1. All three electrodes showed a steady drift to higher potentials with time, while the potential differences between electrodes 2 and 3 were relatively constant. These effects can be explained in terms of the slow injection of copper (or silver) into the sample by the three ion selective electrodes, which is consistent with the stirring dependence data previously discussed. The abnormally high values in potential a t 0800 hr were the readings taken before the stirrer was turned on, which is also consistent with the explanation of slow dissolution of copper (or silver) from the sensor electrodes. This experiment was repeated but with a conventional glass-barreled single-junction Ag/AgCl reference electrode. The measured potentials were much less stable, showing variations of &20 mV during an 8-hour time period, which is further indication that this type of reference electrode is not suitable for monitoring copper in seawater.
ANALYTICAL CHEMISTRY, VOL. 46, NO. 3, MARCH 1974
creases as pH decreases, which, in general, is what would be expected from the pH dependence of copper-organic chelates. To explain these data in that context, however, the interfering species would have to be particularly basic so that a small decrease in p H a t pH 8 would lead to a large shift in copper concentration. The same experiment carried out on the San Diego water resulted in a small shift in potential on changing pH from 8 to 5 (20 mV), but a large decrease in potential on changing pH from 5 to 3 (50 mV). The addition of copper spikes to the San Diego seawaters developed a curve similar to that shown in Figure 5 for the Freeport water but with a more reasonable “Nernstian slope” a t high copper concentrations ( > 2 ppb)-ie., 60 mV/decade rather than 125 mV/decade. The actual potentials observed for the San Diego samples were also displaced to more negative values for a given amount of copper relative to those observed for the Port Isabel samples but not as negative as observed from the Freeport samples. For example, the San Diego samples containing 3.8 ppb copper generated a potential of -75.8 mV; at the same concentration, the Port Isabel sample generated a potential of -55 mV. Atomic absorption spectrometric analysis indicated the presence of 0.3 i~ 0.12 ppb copper in the San Diego water, which was sufficient to have generated potentials of the order of -110 mV for the unspiked solution rather than the -140 mV actually observed. Rather than simply having an inherently low copper content, it would appear that the San Diego sample also contained interfering species, presumably organic chelating agents, but to a lesser extent than the Freeport sample.
As discussed, three different groups of seawater samples were studied. The most representative of the open ocean was the Port Isabel sample taken 5 miles offshore from a nonindustrialized area; the Freeport samples taken 1 mile offshore were most likely representative of nearshore water, the composition of which may have been influenced by industrial activity; and the San Diego sample was also of a nearshore water, the composition of which may have been influenced by industrial activity and especially heavy shipping. The analytical method described for the measurement of ionic copper was developed primarily with one large sample of the seawater taken off Port Isabel. Application of the method to the other two seawater samples indicated a much more complex chemistry. Shown in Figure 5 are the measured potentials after adding copper spikes to the Freeport sample, as received and a t p H 3. These potential readings were taken a t least 15 minutes after addition; in all cases, there was still a slow drift to more positive potentials. The “Nernstian slope” between the last two points was 125 mV/decade (copper additions >5 ppb were not made to avoid possible complications from pH changes induced by the p H 2 CuClz spikes); the initial potential was abnormally low (-171.0 mV) compared to what was anticipated from the AA analysis (0.4 ppb) and compared to the potentials observed for 0.2 ppb copper in 0.5M NaCl (-116.5 mV); and the residual copper concentration calculated from the potential change on adding the first copper spike of 0.63 ppb (and the 50 mV/decade slope already established for this electrode in the Port Isabel water) was 1.1 ppb. Further additions gave calculated residual copper values decreasing to 0.4 ppb with the final spike. All effects are essentially what would be expected from the presence of small quantities (10-7M) of chelating agents, as discussed; indeed, such materials are produced, used, and discharged into the ocean from industrial sites a t Freeport, Texas. The addition of copper spikes to this solution a t pH 3 resulted in a monotonic increase in potential rather than the titration effect observed at pH 8 (Figure 5). Figure 6 shows the steady-state response of the sensor in the Freeport sample as a function of pH. The implication of these data is that the free copper concentration in-
ACKNOWLEDGMENT The authors wish to acknowledge the assistance of Roy Deviney in performing the experimental measurements and the assistance of Gary K. Rice in evaluating the data. Received for review June 7 , 1973. Accepted October 9, 1973. This work was sponsored by Advanced Research Proiects Aeencv on Order No. 2199 and administrated bv the“ Office-of Naval Research under Contract NO. OOIC 72-C-0368.
Spectrochemical Method for the Determination of 36 Elements in Industrial Effluent V. M. LeRoy and A. J. Lincoln lnstrumental Analysis Laboratory, Engelhard Minerals & Chemicals Carp., P.O. Box 2307, Newark, N.J. 071 1 4
A quantitative spectrochemical method is described for the determination of 36 elements in industrial effluent with limits of element sensitivities for the analytical curves in the 0.0005-0.1 mg/100 ml range. One Set Of analytical curves is used for all samples. The samples are weighed salt residues obtained through evaporation to dryness at 120 “C. Any sample volume can be used. Matrix effects are significantly reduced by combining total energy burns in an argon-oxygen atmosphere with 1:6 salt to high purity graphite powder dilutions. Gerrnaniurn dioxide, added to the graphite powder, is the internal
standard. Conditions of analysis, salt weight ratios, preparation of standards, precision, and accuracy are discussed.
Industrial waste and inadequate domestic sewage treatment plants have become the primary targets for initiating programs to achieve responsible control of the two major sources of water pollution. In the precious metals industry, a continuous study of plant waste streams has always been a natural sequence in the refining and manufacturing processes. In addition to the analyses for pre-
ANALYTICAL CHEMISTRY, VOL. 46, NO. 3, MARCH 1974
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