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ANALYTICAL CHEMISTRY, VOL. 51, NO. 13, NOVEMBER 1979
Copper Ion-Selective Electrode for Determination of Inorganic Copper Species in Fresh Waters Renato Stella and M. T. Ganrerli-Valentini Laboratorio di Radiochimica e Centro di Radiochimica ed Analisi per Attivazione del C.N.R., Istituto di Chimica Generale ed Inorganica, Universita' di Pavia, Wale Taramelli 12, 27100 Pavia, Italy
A method is described for determination of soluble copper inorganic species in fresh waters. Assuming that OH- and C032- are the most important ligands, the presence and the distribution of CuOH', Cu2(OH)$+, CuC03(aq), Cu(C03)$- are deduced by measuring Cu2+ concentration as a function of pH in controlled media. The cupric Ion-specific electrode is used. The procedure devised for copper speciation in natural waters requires that, beside free cupric ion concentration [Cu"], only total carbonate carbon C, and pH be known. Some examples of application to actual samples are given.
Copper toxicity toward aquatic life is dependent on chemical forms and hence the knowledge of chemical species distribution is of primary interest (1)in the studies concerning copper transport and biological interaction in natural waters. T h e copper ion-selective electrode has commonly been applied to distinguish between free and bound copper ion (2, 3). In the present work this technique was exploited to study copper behavior in weakly basic and hydrocarbonate solutions, following the line traced by Stiff ( 4 ) . The investigated concentration range simulated natural systems. This research was part of an interdisciplinary investigation on the ecosystems of the Po and Ticino rivers in Northern Italy. T o minimize any possible interference or change in chemical equilibria, the standard addition method was discarded and the direct use of the electrode was extended down to W 0M Cu2+. As the hydrogen ion, H30+, is competitive with copper toward the ligands OH- and C032-, accurate p H measurements were simultaneously made with Cu2+activity determination, thus allowing a more precise identification of all complexation and precipitation equilibria involved. EXPERIMENTAL Electrode Calibration. Before examining simulated Cu(II)/OH- and CU(II)/HCO,--CO,~-systems, the Orion 94-20 A copper ion-selective electrode was carefully calibrated (5-7). The following set of operational conditions was selected: (1)Ionic strength. In all tested solutions, this was fixed at 0.05 M by adding KNO,. (2) Dissolved oxygen interference. Before running any measurement, formaldehyde up t o lo-, M was added to prevent the CuS membrane oxidation. Bubbling a N2 gaseous stream through the solution was avoided as C02also was carried off, thus altering carbonate-hydrocarbonate equilibria. (3) Stirring speed and electrode distance. These were both kept constant through all the measurements. (4) Operational details. A total volume of 100 cm3 was used for all solutions undergoing copper measurement, all the equipment being screened from incident light. (5) TPX plastic containers were always used as this plastic showed little or no adsorption and release of trace metals. Glassor polytheneware were prevented from coming in contact with diluted copper solutions. A 9.95 X M copper solution was prepared by dissolving the appropriate "Suprapure" C U ( C ~ O ~ ) ~ amount . ~ H ~ Oin tridistilled water and the titer was made by EDTA complexometric 0003-2700/79/0351-2148$01.OO/O
titration, both visual and potentiometric. From this solution more M, were prepared by stepwise dilution. diluted ones, up t o Still more diluted ones were prepared according to Hansen (8) as copper ion buffers with nitrilotriacetic acid (NTA) at buffered pHs. Calibration graphs resulted as straight lines down to lo-'' M Cu2+. The cupric ion-selective electrode response was measured with a high impedance electrometer using a single junction Orion 90-01 reference electrode. A tenfold change in activity yielded an electrode response change of about 29.5 f 0.5 mV as expected. Concentration instead of activity was plotted vs. the mV electrode response as the concentration is directly known. Procedure. The effect of copper complexation by the OHligand (or cupric ion hydrolysis)was studied as a function of ligand concentration in a selected pH range (pH 6 to 9). Measured amounts of copper standard solution and "Suprapure" 40% carbonate free NaOH, properly diluted, were added to a stock solution containing formaldehyde and KNO,, lo-, M and 5 X M, respectively. To avoid COz absorption a Nz atmosphere was constantly kept over the solution up to the end of the experiment. "Suprapure" diluted HC104was added stepwise to lower pH from initial 9.7 to 6 and electrode responses were correspondingly recorded. After each addition at least 10 min was required to reach constancy of the electrode response. Interactions between copper and CO': anions were also studied using a copper solution at a total concentration lower than 4.97 X lo4 M to which a NaHC0, or Ca(HC03)2solution was added. Concentrations of the ligands, C032- and OH-, are obviously related to the pH of the system. Therefore pH was continuously changed from 9 to 6 by stepwise addition of "Suprapure" HC104 and the electrode responses were correspondingly recorded. In the first case, a 0.105 M Na2C03solution in known aliquots was added to the copper solution containing formaldehyde and KNOB. In the second case, a standard Ca(HC03)2solution was prepared by bubbling a COPstream through a CaC03 suspension in water for 1h. After the excess CaC03 was filtered off, pH was measured and the filtrate potentiometrically titrated to evaluate the alkalinity which corresponds to [HCOJ at pHs 8. The total carbonate, CT,was deduced from the [HCOC] using the diagram reported by Stumm (9) or the following formula:
cT=
[
["I2
+ [H+]K, + KiKZ [H+]K,
X
[HCOS-]
The value of CT determined this way was 1.63 X lo-' M, which was only 5% higher than [HC03-]. RESULTS AND DISCUSSION Cu(II)/OH- System. Results are shown in Figure 1 where a -2 slope of the log [Cu2+]vs. pH curve indicates the presence of the C U ( O H )precipitate. ~ Over the straight portion (up to the lowest p H value allowing precipitation) log *K,, for the following equilibrium may be calculated:
Cu(OH),(s)
+ 2H+
Cu2+ + 2H20
(1)
A value of log *K, = 9.58 f 0.12 (I = 0.05) is obtained, which is very close to the 9.60 value that is calculated, for both CuO and C U ( O H ) ~from , Schindler's formula (10) when the influence of the molar surface, and hence the precipitate particle size as a function of pH, is taken into account. At lower p H 1979 American Chemical Society
ANALYTICAL CHEMISTRY, VOL. 51, NO. 13,NOVEMBER 1979 c
2149
+
+= 0
3
0
v
a
,O? IN
3-
4-
3-1
5
\
9
'9
s'
6
-1
7
8
9
10 p H -
Flgure 1. Cupric ion concentration (log) as a pH function (carbonate M; (2)Cu, = 3.88 X M; free solutions). (1)Cu, = 1.55 X
(3)CU, = 9.95 x 10-4 M
lo-'
values, where no precipitate is formed, the curves fit the predicted Cu2+ concentrations relative to the following acid-base (or hydrolysis) equilibrium: Cu2+
+ HzO F! CuOH+ + H+
(2)
8
7
10
9
PH-
Figwe 2. Cupric ion concernation (log) as a pH function in the pesence of NaHCO,. (1)CU, = 2.49 X lo-' M, CT = 2.63 X M; (2)CUT = 1.94x 1 0 - ~ M, C, = 2.63 x io-, M; (3)CU, = 3.88 x M, c, = 5.25 x io-, M; (4)CU, = 3.88 x 1 0 - ~ M, cT= 2.60 x io-, M; (5) CUT = 3.88 X M, CT = 1.10 X M; (6)CUT = 4.98 x 10-4 M, C, = 2.60 x 10-3 M
7
-d + a1
0
A
I
log *K1 = -7.52 ( I = 0.05) (9)
3-
The influence of the latter equilibrium is strongly evidenced at the lowest total copper concentrations. Figure 1 also shows that the higher the total copper concentration, the lower the pH a t which the precipitate disappears. A complete description of the Cu(II)/OH- system should include also the second hydrolysis product Cu(OH)2(aq),but a lack of convincing data on the *& constant of the following equilibrium:
Cu2++ 2 H 2 0 F! Cu(OH)z(aq) + 2H+
(3)
makes the problem still open. In a recent paper Vuceta and Morgan (11) suggested, on the basis of potentiometric titrations, a value of log *Pz = -13.7 in agreement with those previously reported by Quintin (12) and by Spivakovskii and Makovskaya (13)and discarded the hypothesis suggested by Mesmer and Baes (14) of log *p2 = -17.3. The measured [Cu2+]as a function of pH suggests a log *pz value of about -16. This may easily be derived from the displacement of the reported curves in a position intermediate between those theoretically calculated by Vuceta and Morgan (11). A more recent work by W. G. Sunda and P. J. Hanson (15) suggests a value very close to -16. As the problem of *p2 value is not solved, the second hydrolysis product Cu(OH)2(aq)was not taken into an account. Anyway it seems to add a contribution much lower than could be predicted on the basis of Vuceta and Morgan results. Cu(1I)/CO3*-, HC03- System. T h e variation of -log [Cu2+]vs. pH is reported in Figures 2 and 3: they show that no relevant difference exists when calcium instead of sodium hydrocarbonate is used.
6
7I
8
9
7
-
Flgure 3. Cupric ion concentration (log) as a pH function in the presence of Ca(HCO,),. (1)CUT = 2.62 X lo-' M; (2)CUT = 4.98 X M; M; (4)CuT = 4.98 X lo-' M. CT = 1.23 X (3)CuT = 1.94 X
lo-, M in all samples
Part of the total inorganic carbon is subtracted by the formation of the following calcium species (1):
+ C032- * CaC03(aq) Ca2+ + H+ + C0:CaHC03+
Ca2+
log Ke, 3.2
11.6
(4) (5)
The amounts of CaC03(aq)and CaHC03+thus formed may be calculated; they were not higher than 1% of the total
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ANALYTICAL CHEMISTRY, VOL. 51,NO. 13, NOVEMBER 1979
- (CU++I
*CUT
y/
,
roJ
m-7
10-6
10-6
,
+
1
C'T
Figure 4. Cupric ion concentration as a function of total added copper concentration-malachite precipitation. (1)CT = 2.65 X lo-, M, pH 7.88; (2)CT = 2.49 X lo-, M, pH 7.70
calcium and for this reason both equilibria were neglected in this work. The upper curves, for which CuT 1 1.94 X 10" M in Figure 2 and CuT = 4.98 X lo4 M in Figure 3, show, for the straight portion, a -1.5 slope which may be related to the presence of precipitated malachite C U ~ ( O H ) ~ C undergoing O~ the following reaction: 1/2Cuz(OH),C03 + 3 / 2 H+ $ Cu2+ + 1/2HC03+ H20 (6) log *K,, = 3.54 ( I = 0.05) (16) For lower total-copper concentration, the curve slope is close to -1, thus indicating that the ion pair CuC03(aq) is the predominant species; C032- concentration shows in fact a similar pH dependence in the range between the two pK values of the carbonic acid, 6.3 and 10.25. The following equilibrium takes place: Cu2+ C032- CuCO,(aq) (7)
*
+
cu++
log K1 = 6.04 ( I = 0.05) (9) Equilibrium 7 obviously does not prevent the simultaneous occurrence of equilibrium 2 but the former is certainly predominant owing to the higher equilibrium constant value, not considering any effect due to a favorable ligand concentration; this view is supported by the fact that curve bending is displaced toward lower pH values. For the malachite precipitation equilibrium C U O H ( C O ~ ) ~ ,Cu2+ ~ + OH+ 1/2 C032- (8)
*
the calculated -log K , = 14.42 f 0.09 (I = 0.05) is surprisingly lower than 15.94, the corresponding value obtained from Schindler's data (10). The measured value was confirmed by checking [Cu2+]and pH on the same malachite saturated solution for 7 days; ageing did not result in any change. The [Cu2+]measurements were repeated in a modified precipitation experiment where copper was added and all other parameters were kept constant: results reported in Figure 4 were consistent with pK,, = 14.42. Such a remarkable discrepancy cannot be ascribed to a lack of precision in the experimental data, but it may be due to a particle size effect similar to the one already reported for C U ( O H )or ~ CuO. According to Hepburn (17) the formation of basic cupric carbonate (malachite) is due to an absorption process; to an extent which is a continuous function of both particle size and concentration of C 0 2 in solution, the absorbent and the
I
4
7
6
9
8
PH
Figure 5. Copper carbonate and hydroxo complex distribution as a pH function. Cu, = 2.50 X lo-' M; C, = 2.62 X lo-, M (NaHCO,); L = OH- or C o t - ; x = 1 or 2;y = 1 or 2 X 1
8r
A
I1
41
71 /
81
/
I 91 I
I
1
I I
101
1
6
7
8
9 pH
.
Flgure 6. Copper carbonate and hydroxo complex distribution as a M; C, = 2.62 X lo-, M (NaHCO,); pH function. C y = 3.88 X L = OH- or COS2-;x = 1 or 2; y = 1 or 2
absorbed materials being, respectively, cupric oxide and un-ionized carbonic acid. To calculate the distribution of the most probable copper species in fresh waters at carbonate-hydrocarbonate natural levels, the following equilibria, besides equilibria 2 and 7, were considered:
+
2Cu2+ + 2 H z 0 F! Cuz(OH)Z+ 2H+ log Keq ( I = 0.05) (9,10, 16) - 10.54 (9) cu2+ + 2C0s2- $ Cu(C03)22-
9.28
(10)
To check the reliability of the assumed set of equilibria, the
ANALYTICAL CHEMISTRY, VOL. 51, NO. 13, NOVEMBER 1979
Table I. Copper Hydroxide and Carbonate Species Distribution in Natural Watersa natural samples pH CT CUTb [Cu'+] [CuOH'] [Cu,(OH),"] Po A1 Po C1 P o F1
7.63 7.59 7.86 7.62
1.71 X 1.85 X 2.74 X 2.08 x
4.72 X 3.79 X 1.39 X 4.25 x
Ticino 2 a All concentrations in moles per liter.
lo-' lo-'
3.70 X 10"O 1.00 X lo-'' 3.23 X lo-'' 1.80 x
lo-'' lo-'' lo-'
4.50 X 2.76 X 9.95 X 1.02 X
lo-'' lo-'' lo-''
[CuCO,]
[Cu(CO,),"-]
2.43 X 6.38 X 10." 5.95 X 1.40 X
2.53 X 10." 6.46 X 1.73 X lo-'' 1.72 X 10"
By atomic absorption including organically bound copper. carbonate species may however be calculated. It is sufficient to measure the natural pH value, the corresponding free cupric ion concentration, and to calculate the total carbonate carbon concentration CT from the measured alkalinity. Several examples of these calculations are reported in Table I.
cu, LLy I CU" .. '/CUT-
3.79 X 9.33 X 5.62 X 1.80 X
2151
(CU++)
CONCLUSIONS Though natural surface waters are very complicated systems which cannot be easily simulated, the simplified speciation model described here allows one to obtain useful informations a t least for the most commonly soluble inorganic copper species. Other inorganic ligands such as phosphate and chloride may, in fact, interfere, as suggested by Stiff ( 4 ) ,but because of the low formation constants of the complexes and the relative low concentration of the ligands their contribution may be neglected. ACKNOWLEDGMENT The authors are grateful to S. Meloni for valuable discussion concerning this work.
91
10
LITERATURE CITED G. K. Pagenkopf, R. C. Russo, and R. V. Thurston, J . Fish. Res. Board
6
7
8
9 PH-
Flgure 7. Copper carbonate and hydroxo complex distribution as a pH function. Cu, = 2.62 lo-' M; CT = 2.47 X M [Ca(HCO,),]; L = OH- or C03*-; x = 1 or 2; y = 1 or 2
sum of all calculated concentrations is also reported in Figures 5,6, and 7 for a comparison with the amount directly obtained as CUT - [Cu2+]a t different pH values. In Figure 6 this comparison is limited to pH 7.2 as malachite precipitation occurs beyond this point. The close similarity of the two curves means that the whole system is sufficiently described by the proposed equilibria. In Figure 5 the discrepancy observed for pH values higher than 7.8 may be due to the fact that the Cu(OH),(aq) reaches the highest concentration in this region; even if its contribution could be exactly calculated, this should make available a reliable value for its equilibrium constant *&. Application to River Samples. Water samples of the Ticino and Po Rivers were collected and immediately filtered through a Millipore HA fiiter (0.45-pm pore size). The filtered water still contains all colloids holding trace metals and may be fractionated using the ultrafiitration technique. The soluble copper distribution among the most probable hydroxy and
Can., 31, 462 (1974). S.Ramamrty and D. J. Kushner, J . Fish. Res. B a r d . Can., 32, 1755 (1975). Y. K. Chau and K. Lum Shue Chan, Water Res., 8 , 383 (1974). M. J. Stiff, Wafer Res., 5 , 585 (1971). D. Midgley, Anal. Chim. Acta, 87, 7 (1976). M. J. Smith and S. E. Manahan, Anal. Chem., 45, 836 (1973). W. J. Blaedel and D. E. Dinwddie, Anal. Chem., 46, 873 (1974). E. H. Hansen, C. G. Lamm, and J. Ruzicka, Anal. Chim. Acta, 59, 403 (1972). W. Stumm and J. J. Morgan, "Aquatic Chemistry, an Introduction Emphasizing Chemical Equilibria in Natural Waters", Wiley-Interscience, New York, 1970, Chap. 4, 5, and 6, pp 118-299. P. Schindler, M. Reinert, and H. Gamsjager, Helv. Chlm. Acta, 51, 1845 (1968). J. Vuceta and J. J. Morgan, Limnol. Oceanogr., 22, 742 (1977). M. Quintin, C . R . Hebd. Seances Acad. Sci., 204, 968 (1937). V. B. Spivakovskiiand G. V. Makovskaya, Russ. J . Inorg. Chem., 13, 815 (1968). R. E. Mesmer and C. F. Baes Jr., ORNL-NSF-EATC-3 (1974). W. G. Surida and P. J. Hanson, "Chemical Modeling in Aqueous Systems: Speciation, S a p h , Solubility, and Kinetics", E. A. Jenne, Ed.; ACS Symp. Ser., 93, 179. L. G.Sillen and A. E. Marten, "Stability Constants of MetaCIon Complexes", Chem. SOC. Spec. Pub/. No. 17 (1964); No. 25 (1971). J. R. Irons Hepburn, J. Chem. SOC.,2883 (1927).
RECEIVED for review March 12,1979. Accepted July 24,1979. This work was supported by the ENEL (Italian Electrical Energy Agency) which also gave permission for publication.
