Environ. Sci. Technol. 1988,22,768-771
Oxidation of Copper( I ) in Seawater Virender K. Sharma and Frank J. Miiiero" The Rosenstiel School of Marine and Atmospheric Science, University of Miami, 4600 Rickenbacker Causeway, Miami, Florida 33 149
The oxidation of Cu(1) in air-saturated solutions of seawater has been measured as a function of pH (5.3-8.6), temperature (5-45 "C), and salinity (5-44). The overall rate constant k (kg mol-l min-l) d[Cu(I)]/dt = -k[Cu(I)][02]
W
has been fitted to log k = 10.73 0.23pH - 2373/T - 3.33N2
+
+ 1.451
with a standard error of 0.08 in log k. The energy of activation was found to be 45.6 f 1.7 kJ mol-l. The strong chloride dependence of the rates has been attributed to the effect of temperature and ionic strength on the oxidation of Cu+ and CuClO species. At a given C1- concentration, the rates of oxidation in seawater are lower than in NaCl or NaC1-NaC104 mixtures. These differences in seawater are due to a decrease in the rates due to Mg2+ and Ca2+ and the increase in the rates due to HCOc. Possible causes of these effects are discussed. Introduction Recently, Moffett and Zika ( I ) have suggested that dynamic nonequilibria in surface seawater may lead to the formation of Cu(1). The photooxidation of dissolved organic matter may form reducing agents such as 02-and H202by the suggested sequence (2,3) org
-+ -+ hv
e202-
org+
-
02-
O2
+ 2H+
e-(aq)
H2Oz + 02
(1)
(2) (3)
The reducing agents (0; and H202)can react with Cu2+ to form Cu+: CU2f 0 2 cu+ 0 2 (4)
+
+
One might also expect that Cu(1) may occur in anoxic waters and sediments. Once reduced Cu(1) is formed in natural waters, its lifetime will be related to its rate of oxidation by O2 to Cu(I1). Moffett and Zika (I) have measured the rate in seawater (S = 35) at 25 OC as a function of pH. They also measured the rate in artificial seawater solutions (I = 0.92) where C1- was replaced by Clod-. These measurements showed a large C1- dependence and have been related ( 4 ) to the differences in the rates of oxidation of the various copper(1) chloride complexes:
+ + -
Cu+ + O2 CuClO CuC1,-
O2
O2
ko
kl
k2
products
(5)
products
(6)
products
(7)
In this paper, we will present measurements of the rate of oxidation of Cu(1) in seawater as a function of pH (5.3-8.6), temperature (5-45 "C), and salinity (5-44). These measurements have been used to derive a rate equation valid for most natural waters. The chloride dependence has been used to determine the effect of temperature on the oxidation of Cu+ and CuCl species. 768
Environ. Scl. Technol., Vol. 22, No. 7, 1988
Experimental Section The measurements were made in a 500-cm3 waterjacketed cell. The temperature of the cell was controlled to f0.02 "C with a Porma bath. The top of the cell was fitted with a glass frit used to bubble air through the solutions, and an autopipet was used to take out sample aliquots during a run. The Cu(1) concentrations were determined spectrophotometrically with bathocuproine (2,9-dimethyl-4,7-diphenyl-1,lO-phenanthroline) and the methods developed by Moffett et al. ( I , 5). A 10-cm3aliquot of solution was removed from the cell and placed into a 50-cm3 flask containing 0.05 cm3 of bathocuproine M). The solutions were stirred rapidly to quench the oxidation and diluted with water to 50 cm3, which had a final concentration of bathocuproine of 10" M. The absorbance of the solution was meagured at 484 nm with a Varian Cary 2200 UV-vis spectrophotometer. AH the measurements were made at micromolar levels of Cu(1). The rates were shown to be first order with respect to Cu(1) over the range 10-0.1 pM. Ethylenediaminetetraacetic acid (EDTA) ( M) was added ( I , 5 ) to the solutions to extend the time that the rate of disappearance of Cu(1) is pseudo first order. This may be due to the complexing of Cu2+with EDTA preventing the back-reaction with H202 ( I ) . We have found, in agreement with Moffett and Zika ( I ) , that the rates of oxidation of Cu(1) with and without EDTA are the same. The copper(1) solutions were prepared by dissolving cuprous bromide (Aldrich) in a 1 M NaCl and 0.001 M HCl solution that had been deoxygenated with nitrogen. The solutions were stirred to enhance the dissolution of CuBr. The concentration of Cu(1) in this stock solution was M and was maintained under a blanket of nitrogen. The stock solutions of bathocuproine (Aldrich) at M were freshly prepared for each experiment. The seawater used was Gulf Stream water collected off the coast of Miami. The salinity was determined with a Guildline Autosal conductance bridge using the practical salinity scale. After filtration through a 0.45-bm filter, the samples were diluted by weight with Millipore Super-Q ion-exchanged water (18
MQ).
All of the measuremetlts were made in solutions saturated with air at a given temperature by bubbling. The mold concentration of O2in the solutions at a given temperature and salinity was calculated from the equation of Benson and co-workers (6, 7). The seawater solutions at pH below and above 8 were made by adding HC1 or NaOH to the solutions. Tris(hydroxymethy1)aminomethane (Tris)-seawater (0.005 rn) buffers (8,9) were used to calibrate our electrode system and to determine the pH. The calibrations were made at each salinity and temperature with the pHF on the free molal scale (8) and the equations of Miller0 (9). Results The rate constant k (kg mol-l min-l) for the oxidation of Cu(1) d [ C ~ ( I ) ] / d= t -~[C~(I)1[021 (8) has been determined as a function of pH, temperature, and
0013-936X/88/0922-0768$01.50/0
0 1988 American Chemical Society
3.2 I 1 -
Table I. Values of /3,* for Copper(1) Chloride Species at 25 Oca 1%
52
b
0
3.10 2.71
35
r 1
S
log Pz*
P1*
C
1%
P3*
b
C
b
C
6.31 5.96
5.42
5.47 5.64
4.75 4.92
5.07
Moffen o n 6 Z t k o
1
.6 -
7
2 8
9
nFrom Sharma and Miller0 (17). *Ahrland and Rawthorne (11). CFritz(12).
PH
Figure 1. Effect of pH on the rate constant for the oxidation of Cu(1) in seawater (S = 35). r. 7 ,
3 9
log k 3.2 2.5 1.8
A -
0
0.2
0.4
.Y,
0.6
0.8
1.0
(-
Figure 2. Effect of ionic strength on the rate constant for the oxidation of Cu(1) in seawater.
4.0
I
Figure 4. Speciation of Cu(1) In NaCl solutions as a fun&tion of CIusing the 0,'sfrom Fritz ( 74) (-) and Ahrland and Rawthorne ( 7 7 )
-1.
due to the difficulty of the extrapolations to zero C1-. More will be said about this in the next section.
Discussion
1
As discussed elsewhere ( 4 ) ,the C1- dependence of the oxidation of Cu(1) is given by
k 2.0
'
3 05
=
~ C & O+
acuclki + acuclzkz+ acuclsk3
(10)
where the molar fractions ai of the various species can be determined from 3.25
3 45 ($)X103
3 65
acu = (1 + &*[Cl-]
Figure 3. Effect of temperature on the rate constant for the oxidation of Cu(1) in seawater ( T , K).
salinity, The results for the effect of pH on the rate constant for seawater (S = 35) at 5 and 25 "C are shown in Figure 1. Our results at 25 "C are in good agreement with the results of Moffett and Zika (I)-see Figure 1. The slope of log k versus pH is 0.23 f 0.04 and independent of temperature. This value is in excellent agreement with the slope of 0.22 found by Moffett and Zika (I) at 25 "C. The small variation of the rate with pH indicates that acidic and basic species (H', OH-, CuOH, and CuCIOH-) are not strongly involved in the oxidation. The effect of ionic strength, I = O.O199S/(l- 10-3S) (IO), on the rates is shown in Figure 2. The values of log k are nearly linear with 11/2 at low temperature but show a second-degree functionality at higher temperatures. The decrease in the rates at higher ionic strengths (or salinity) is due to the formation of copper(1) chloride complexes (I, 4 ) . The effect of temperature on the values of log h is shown at fixed salinities in Figure 3. Within the experimental error of the measurements, the slopes of log k versus 1/T (K) are linear and independent of salinity. All of the measurements have been fitted to log k = 10.73 + 0.23pH - 2373/T - 3.33I1I2+ 1.451 (9) which has a standard error of 0.08 in log k . This equation gives an energy of activation for the oxidation of Cu(l) of 45.6 f 1.7 kJ mol-l. This equation should be valid for estuarine and seawater solutions. The results for pure water may be unreliable
+ pz*[Cl-]Z + p3*[C1-]3)-l (11)
acuc1 = P1*[C1-I~cu
(12)
= Pz* [C1-I2acu
(13)
"CUCI*
acuc1, = P3* [C1-I3acu The stepwise association coqstants Pi are given by PI
= ( [ C ~ C 1 I / [ C ~[c1-l)(YcucdYcuYcl) +l
(15)
P2
= ~ ~ ~ ~ ~ ~ z - l / ~ ~ ~ + l ~ ~ ~ - 1 2 ~ (16) ~ Y c u c ~ , / Y c
P3
= ( [ C U C ~ /~ [Cu'l ~ - I ~ ~ ~ ~ 1 3 ~ ~ ~(17) c u c ~ 3 /
where [i] and y1are respectively the concentrations and activity coefficients of species i and P,* = [CUC~;~"]/ [Cu'] [Cl-ILis the stoichiometric constant. In the earlier analysis ( 4 ) of the experimental measurements of Moffett and Zika (I) for artificial seawateri the values of &* were taken from the work of Ahrland and Rawthorne (11) at I = 5.0. Recently, Fritz (12-14) has reanalyzed these results as well as literature data. He has cqlculated values of &* and P3* in various ionic media at various temperatures and determined Pitzer (15) activity coefficient parameters for the various ion pairs. Fritz (12) was not able to determine a reliable value of &* from solubility data. We have extrapolated the values of Ahrland and Rawthorne (11)at I = 5.0 to infinite dilution using Pitzer's (15) equations and assuming ycu = yNaand In ycucl = 0.1321 (16). The various values of &* at I = 0.7 in NaCl are compared in Table I. The differences are quite significant and lead to differences in the speciation of Cu' and CuClO (see Figure 4). This difference in speciation can lead to Envlron. Sci. Technol., Vol. 22, No. 7, 1988 769
0.25
"0
0.50
0.75
[cl-l
Figure 5. Values of kla,, from Moffett and Zika ( 7 ) plotted versus [CI-].
A 1
A
Ahrlond a n d R a w t h o r n s
0
Fritz
Table 11. Effect of Mg2' and HCOc on the Oxidation of Cu(1) a t t h e Same Cl- Concentration
I
\
Figure 8. Values of log k versus [CI]"* for the Oxidation Cu(1) in seawater.
mNa
0.71 0.76
mMg
mCa
mHC03
0.71 0.70
mso,
0.70 0.0023 0.70 0.54 0.08 0.54 0.08 0.70 0.54 0.08 0.70 0.0023 seawater (s= 43.6, [CI-] = 0.71 m )
3.75 f 0.02 3.76 f 0.01 4.21 f 0.02 2.35 f 0.04 2.35 f 0.04 2.82 f 0.02 2.81 f 0.02
0.50 0.50 0.50 0.0016 0.38 0.06 0.50 0.38 0.06 0.50 0.0016 seawater (8 = 30, [Cl-] = 0.49 rn)
4.05 f 0.02 4.41 f 0.02 2.61 f 0.02 2.97 k 0.04 2.94 f 0.04
0.03
0.70
0I - _ i _ _ _ _ _ L 0 025 0 50
[Cll
075
Figure 6. Values of klaCuCI, from Moffett and Zika ( 7 ) plotted versus [Ci-1.
kiio'i
%"
Figure 7. Values of kla,, for seawater at various temperatures.
differences in rate constants assigned to these species. This is shown in Figure 5 for the data of Moffett and Zika ( I ) . The results indicate that Cu+ and CuClO are the most reactive species in agreement with earlier analysis, but the values of ko and k, are a function of the stability constants used. We will use the more reliable values of pi* from Fritz (12) in our analysis of our seawater results. It should be pointed out that changes in the fraction of CuC1,- and CuC1,2- cannot account for the changes in the rate of oxidation of Cu(1) in dilute C1- solutions. This is demonstrated in Figure 6 where the nonlinear behavior of k/acuc12 as a function of C1- is shown. A plot of k/cucu for seawater as a function of the molal concentration of C1- ( I ) is shown in Figure 7. The results converge to a value of log ko = 6.1 f 0.2 independent of temperature. The nonlinear dependence of k / a c u can be attributed to an ionic strength dependence of k , log kl = 11.51 - 2024/T + 0.461 (18) It is also possible to account for the curvature in the plots of k/crcu by assuming that the CuC1,- species is being oxidized. Since this is not the case in NaC1-NaC1O4 mixtures at a constant ionic strength we do not feel this explanation is reasonable. Since the other components of seawater are also changing when seawater is diluted, other factors could contribute to the nonlinear dependence of k l a c u . These effects can be elucidated by comparing our seawater results to the artificial seawater results of Moffett and Zika (1) and our results for NaCl and NaCl-NaC10,
(In,
770
Environ. Sci. Technol., Vol. 22, No. 7, 1988
1% k
mixtures (Iq-see Figure 8. Our seawater results are in reasonable agreement at a given C1- concentration with the results of Moffett and Zika (1). These comparisons indicate that over a narrow ionic strength (0.1-1.0 m) the C1- concentration controls the rate of Cu(1) oxidation. The small differences a t low and high C1- concentrations in seawater can be attributed to changes in ionic strength (low Cl-> and changes in composition (high C1-). The results in NaCl (0.5-1 m) and NaC1-NaC10, mixtures ( I = 1.0) also show that over a narrow ionic strength range C1controls the rate. The seawater results, however, are much lower than those of the NaCl or NaCI-NaC104 mixtures at the same C1- concentration. To elucidate these d'ifferences, we have measured the rate of oxidation in NaC1, NaMgC1, NaCaC1, NaCIHC03, NaClSO,, and NaMgClHCO, mixtures at a C1- concentration of 0.7 and 0.5 m. The results are given in Table 11. The addition of Mg2+or ea2+to NaCl at a constant C1- concentration causes a decrease in the rate, while the addition of NaHCO, causes the rate to increase. The SO-: ion has no effect within our experimental error. The rate of oxidation of Cu(1) for the artificial seawater solution made up of Na-Mg-C1-HCO3 is in good agreement with the measured value in seawater at the same C1- concentration. These results demonstrate that Mg2+and Ca2+cause a decrease and HCO, causes an increase in the rate of oxidation of Cu(1) at a constant Cl-. The decrease in the rate by Mg2+and Ca2+in seawater or artificial seawater can be attributed to a number of factors. In our initial thoughts we felt that the addition of Mg2+or Ca2+could slow down the rate of oxidation by forming ion pairs with the charged copper(1) chloride complexes: Mg2++ CuC1,Mg2++ CuC1:-
-
MgCuC12+
(19)
MgCuCl,'
(20)
-+
The formation of these complexes would decrease the more
reactive Cu+ and CuClO species at a fixed Cl- concentrations. Preliminary spectroscopic measurements of the CuC12- and CuC12- species in NaC1-MgC12 solutions indicate that Mg2+does not interact with these species. We, thus, do not feel that this is a likely cause of the observed effects. We feel that a more likely cause of the decrease in the rates is due to the slow exchange of Mg2+(or Ca2+)complexes with Cu2+: MgL Cu2+ CuL + Mg2+ (21)
-
+
This slow exchange would cause the overall oxidation rates of Cu(1) to be slower due to the back-reaction of Cu(I1) with reductants such as 0; or Hz02( I ) . Since the addition of NTA, CO2- and B(OH)4-gives similar rates at levels sufficient to complex Cu2+in Na-Mg-C1 solutions, this exchange reaction may be common for other ligands able to complex Mg2+and Cu2+in natural waters such as humic material. Further measurements are needed to examine the rates of the back-reaction in the presence of various ligands. The increase in the rate of oxidation due to an increase in HCO, and C032-in NaCl was attributed by Moffett and Zika ( I ) to the slowing down of the back-reaction of Cu(I1). Since the level of EDTA in our experiments was more than sufficient to theoretically complex the Cu(II), we feel that the increase in the rate may be related to the formation of copper(1) carbonate complexes: Cu+ + HC03- CuHC030 (22) cu+
-
-
+ c0:-
cuco3-
(23)
These copper(1) carbonate complexes may be more reactive than CuClO. Another possibility could be due to the different rates of reduction of CuEDTA and CuC03 complexes. For example
--
+ 02CuC03 + 02-
CuEDTA
+ O2 + EDTA Cu+ + O2 + C032Cu+
(24)
(25)
or similar reaction with HzOzcould occur. If the rates of reaction 24 are greater than those of reaction 25, the rates of Cu(1) oxidation will be faster in carbonate solutions. Further kinetic and thermodynamic data are needed in NaC1-HC03 solutions over a wide range of C1- and ionic strength to elucidate these effects. The nonlinear behavior of k/acu in seawater can be compared to the linear behavior for artificial seawater. The results shown in Figure 9 show good agreement at low levels of C1-. The deviations at higher C1- are largely due to HC03- in seawater that was not in the artificial solutions used by Moffett and Zika (1). They also used a higher Mg/C1 ratio in their artificial solutions. As the seawater solutions are diluted, the effect of HC03- becomes less important. If corrections arb made for these effects, the seawater results fall on the same linear curve as the artificial seawater results of Moffett and Zika (1). In summary, our results for the oxidation of Cu(1) in seawater indicated that the rates are controlled by the concentration of Mg2+, Ca2+,C1-, and HC03-. Further
Figure 9. Values of kla,, for seawater ( I = 0.1-0.91) and artificial seawater ( I = 0.92) plotted versus [CI-1.
kinetic and thermodynamic measurements are needed to completely characterize the causes of changes due to Mg2+, Ca2+,and HC03- on the rates of oxidation.
Supplementary Material Available Two tables detailing the effects of pH, salinity, and temperature on the oxidation of Cu(1) in seawater (2 pages) will appear following these pages in the microfilm edition of this volume of the journal. Photocopies of the supplementary material from this paper or microfiche (105 X 148 mm, 24X reduction, negatives) may be obtained from Microforms Office, American Chemical Society, 1155 16th St., N.W., Washington, DC 20036. Full bibliographic citation (journal, title of article, authors' names, inclusive pagination, volume number, and issue number) and prepayment, check or money order for $10.00 for photocopy ($12.00 foreign) or $10.00 for microfiche ($11.00 foreign), are required. Registry No. Cu, 7440-50-8; Mg, 7439-95-4; Ca, 7440-70-2.
Literature Cited (1) Moffett, J. W.; Zika, R. G. Mar. Chem. 1983, 13, 239. (2) Zafiriou, 0. G. Mar. Chem. 1977, 5, 497. (3) Zika, R. G. In Marine Organic Chemistry;Duursma, E. K., Dawson, R., Eds.; Elsevier, Amsterdam, 1981; pp 299-325. (4) Millero, F. J. Geochim. Cosmochim. Acta 1985, 49, 547.
(5) Moffett, J. W.; Zika, R. G.; Petasne, R. G. Anal. Chim. Acta 1985, 175, 171. (6) Benson, B. B.; Krause, D., Jr.; Peterson, M. A. J. Solution Chem. 1979,8, 655. (7) Benson, B. B.; Krause, D., Jr. Limnol. Oceanogr. 1984,29, 620.
( 8 ) Ramette, R. W.; Culberson, C. H.; Bates, R. G. Anal. Chem. 1977, 49,867. (9) Millero, F. J. Limnol. Oceanogr. 1986, 31, 839. (10) Millero, F. J. Ocean Sci. Eng. 1982, 7, 403. (11) Ahrland, S.; Rawthorne, J. Acta Chem. Scand. 1970, 24, 157. (12) (13) (14) (15) (16) (17)
Fritz, J. J. J. Phys. Chem. 1980, 84, 2241. Fritz, J. J. J. Phys. Chem. 1981, 85, 890. Fritz, J. J. J. Solution Chem. 1984, 13, 369. Pitzer, K. S.; Mayorga, G. J. Phys. Chem. 1973, 77, 2300. Millero, F. J.; Schreiber, D. R. Am. J. Sci. 1982,282,1508. Sharma, V. K.; Millero, F. J. J. Solution Chem. (in press).
Received for review May 5, 1987. Revised manuscript received December 23,1987. Accepted January 25,1988. The support of the Office of Naval Research (N00014-87-G-0116)and the Oceanographic Section of the National Science Foundation (OCE-8600284) for this study is acknowledged.
Environ. Sci. Technol., Vol. 22, No. 7, 1988 771
