1. pets
2. aquariums

Notes
iment redox and solute dynamics. Ecology 67: 875882.
KRABBENHOFT, D. P. 1984. Hydrologic and geochemical controls
of freshwater
ferromanganese
deposit formation at Trout Lake, Vilas County,
Wisconsin. M.S. thesis, Univ. Wisconsin-Madison. 137 p.
-.
1988. Hydrologic and geochemical investigations of aquifer-lake interactions at Sparkling
Lake, Wisconsin. Ph.D. thesis, Univ. WisconsinMadison. 213 p.
LEE, D. R. 1977. A device for measuring seepage flux
in lakes and estuaries. Limnol. Oceanogr. 22: 140147.
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J. A. CHERRY, AND J. S. PICKENS. 1980.
Groundwater transport of a salt tracer through a
sandy lake bed. Limnol. Oceanogr. 25: 45-61.
LODGE, D. M., T. K. KRA’IZ, AND G. M. CAPELLI. 1986.
Long-term dynamics of three crayfish species in
Trout Lake, Wisconsin. Can. J. Fish. Aquat. Sci.
43: 993-998.
-,
AND J. G. LDRMAN. 1987. Reductions in submersed macrophyte biomass and species richness
by the crayfish Orconectes rusticus. Can. J. Fish.
Aquat. Sci. 44: 591-597.
LOEB, S. L., AND C. R. GOLDMAN. 1979. Water and
nutrient transport via groundwater
from Ward
Valley into Lake Tahoe. Limnol. Oceangr. 24:
1146-l 154
-,
AND S. H. HACKLEY. 1988. The distribution
of submerged macrophytes in Lake Tahoe, California and Nevada, and the possible influence of
groundwater
seepage. Int. Ver. Theor. Angew.
Limnol. Verh. 23: In press.
Limnol.
Oceanogr., 34(l), 1989, 239-244
0 1989, by the American
Society of Limnology
and Oceanography,
239
MCCREARY, N. J., S. R. CARPENTER, AND J. E. CHANEY.
1983. Coexistence and interference in two submersed freshwater perennial plants. Oecologia 59:
393-396.
N~NBERG,
G., AND R. H. PETERS.
1984. Biological
availability
of soluble reactive phosphorus in anoxic and oxic freshwaters. Can. J. Fish. Aquat. Sci.
41: 757-765.
OKW~EZE, E. E. 1983. Geophysical investigations of
the bedrock and groundwater-lake
flow systems in
the Trout Lake region of Vilas County, northern
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108:
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Submitted: 29 March 1987
Accepted: 12 August 1988
Revised: 28 September 1988
Inc
Temperature dependence of Cu(I1) carbonate complexation in
natural seawater
Abstract-Cu(I1)
carbonate complexation was
examined in natural seawater between 5” and
35°C. The reaction enthalpy for the equilibrium
Cl?’ + co,*- = CuCO,O, expressed in terms of
total carbonate ion concentration,
is -2.5 kcal
mol-l. The formation constant sw@l= [CUCO,~]
[cu*+]-1[co,*-],-‘,
expressed in terms of total
carbonate ion concentration in seawater, thus decreases by about 25% between 25” and 5°C.
Acknowledgments
This work was supported in part by NSF grants OCE
83-08890 and OCE 84-00548. We thank C. H. Culberson and A. J. Paulson for constructive criticism.
It is generally appreciated that chemical
complexation is a dominant factor affecting
the behavior of trace metals in seawater.
Solution complexation exerts important
controls on metal bioavailability, toxicity,
solubility, and adsorptive behavior. In view
of the general importance of complexation
on metal behavior in seawater, it is, at first
sight, surprising that there is very little direct evidence that would allow a description
of chemical complexation at temperatures
appropriate to normal oceanic conditions.
The vast majority of complexation inves-
Notes
240
tigations performed in the laboratory are
conducted at temperatures 1 20°C, although -90% of ocean waters (Montgomery 1958) have temperatures < 8°C.
Toward the goal of obtaining metal speciation models appropriate to the normal
range of oceanic conditions, we here describe investigations of the influence of temperature on Cu(I1) complexation in natural
seawater. Copper was selected as an element
of particular interest because of the important role of chemical complexation in influencing copper toxicity in natural aquatic
systems (Spenser 1957; Johnson 1964; Steemann Nielsen and Wium-Anderson 1970;
Sunda and Guillard 1976; Anderson and
Morel 1978).
Previous work on Cu(I1) solution chemistry has shown that dissolved copper in
seawater is partitioned principally between
organic complexes, CuL,,, and the inorganic species CUCO,~ (Bately and Florence
1976; Van den Berg 1982, 1984a; Zuehlke
andKester 1983a,b; ByrneandMiller 1985;
Sunda and Hanson 1987):
+ t-2Lorg9- = CU(Lorg)n(2-nq)
+ co32-.
(1)
In an attempt to produce quantitative copper complexation models appropriate to
natural systems, it is useful to separate Eq.
1 into organic and inorganic components:
CuCO3O
Cu2+ + nL 0%q- * Cu(L,,),(Z-“4)
(2)
and
cu2+ +
O*
(3)
Quantitative characterization of Eq. 2 for a
wide variety of organic ligands (Martell and
Smith 1974, 1977, 1982; Smith and Martell
1975) indicates that the strength of Cu(II)organic ligand associations is remarkable for
a divalent metal, generally surpassed only
by Hg(I1) and Pd(I1).
Investigations of the temperature dependence of reaction 3 appropriate to seawater
are complicated by the interactions of C032with the major seawater constituents Na+,
Mg2+, and Ca2+. In this work we have obviated the necessity of constructing a temperature-dependent major-ion seawater
speciation scheme by conducting our investigations directly in natural seawater.
co32-
= cuco
3
According to such a procedure, complexation results are expressed in terms of total,
rather than free, carbonate ion concentration. In previous CUCO,~ complexation
work at 25”C, good agreement was observed
between results obtained in natural seawater and simple synthetic media (Byrne and
Miller 1985). With the same ultraviolet
spectroscopic techniques used in previous
trace metal-carbonate investigations (Byrne
198 1; Byrne and Miller 1985), we have conducted studies which apparently constitute
the first direct examination of a trace metal’s
temperature-dependent complexation behavior in natural seawater.
The procedures used here closely follow
the methods used by Byrne and Miller ( 1985)
in UV spectroscopic examinations of Cu(II)carbonate complexation in seawater at 25°C:
The absorbance of copper-enriched seawater, 5 x 10-6M Cu(II), vs. a copper-free seawater blank was monitored at 280 nm. Measurements were performed at 5”, 15”, 25”,
and 3 5°C. Use of an open-topped 1O-cm cell
housed in the temperature-controlled well
of a spectrophotometer (Cary 17D) permitted simultaneous measurements of absorbance and pH. Our titrations were conducted between pH 8 and 5.6. About 30
titration points were obtained in each of our
four experiments. Absorbances ranged between -0.125 and 0.005. Carbonate ion
concentrations in our experiments ranged
between -1 x lop4 M and 1 x lop7 M.
Seawater solutions were titrated with HCl
using calibrated Gilmont microburets.
Measurements of pH were obtained on the
NBS scale with a Ross combination electrode. Seawater was obtained from Gulf of
Mexico surface water and diluted to S = 35
(see list of symbols). The initial total alkalinity (TA) of our seawater was obtained
through potentiometric analysis (Culberson
et al. 1970; Johnson et al. 1977). The results
of a typical titration are shown in Fig. 1. In
contrast to the work of Byrne and Miller
(1985), the present study did not include
seawater samples having greater than normal total CO2 concentrations. As a consequence of the much lower total CO2 levels
and somewhat more acidic conditions (pH
5 8.0) in our study, our data analyses were
conducted with a truncated version of the
Notes
241
Symbols used in this work.
Meaning
Units
A
AO
I
0
,
20
40
[CO;-],
60
1
80
MOLES/LITER
1
120
I
100
1
140
x106
Fig. 1. An example of our absorbance vs. carbonate
ion concentration data is shown for a titration at 15°C.
The pH in this experiment ranged between 8.01 and
Solution absorbance
Sum absorbance of Cu2+,
CuCl+, and CUSO,~
liters2 mo1-2 cm-l
Product of c, and J,
A,
BA
eq liter-’
Borate alkalinity
BO
Sum concentrations of Cu2+,
CuCl+, and CuSO,O divided
by Cu2+ concentration
CA
Carbonate alkalinity
eq liter-’
Reaction enthalpy
kcal mol-I
Bicarbonate dissociation constant
Boric acid dissociation constant
Total alkalinity
eq liter-l
Gas constant
cal deg-’ mol-I
Cu2+ hydrolysis constant
mol liter-’
CuC030 formation constant liters rnol-’
CuCI+ formation constant
liters mol-’
s0,2-p, CuSO,O formation constant liters mol-*
CUCO,~ molar absorptivity
liters mol-’ cm-*
>
Salinity
5.62.
analytical equation used by Byrne and Miller (1985).
Spectrophotometric data obtained in our
seawater titrations were analyzed with the
equation
A
Cu, x 10
an enthalpy for CuCl+ formation (Smith and
Martell 1974) equal to 1.6 kcal mol-‘. In
the absence of enthalpy data appropriate to
CuSO,Oformation in seawater, and in view
of the minor influence of copper sulfate
complexation on our analyses, the term
so,2&w42-l was assumed constant at all
temperatures in this work.
Total carbonate concentrations in our exA, + AJC032-lT
periments
were calculated with the rela= B, + &*[H+]-’ + ,Wp1[C032-]T (4)
tionships
where A is absorbance, Cu, is total copper
CA=TA-BA
(7)
concentration, [C032-]T is the total molar
concentration of CO 32-,A, is a constant in- and
corporating the sum absorbance contribuCA = [C032-]T{2+(K;)-1[H+]}.
(8)
tions of Cu 2+, CuCl+, and CuSO,O (Byrne
and Miller 1985), A, is equal to cl x sWfll TA was calculated as the initial seawater TA
where tl is the molar absorbance of CuCO,O, minus the concentration of added HCl. BA
and Jl is defined as
was determined from pH and the total boron concentration (BT) in S = 35 seawater:
[CuCO,O]
& = [cu2+][co32-]T*
(5)
The term B, in the present work was approximated as
I?. = 1 + cl&[C1-] + ,,,2-f?,[S0,2-].
(6)
At 25°C we set &, = 0.57 and so,~-pl[S0,2-]
= 0.10. At other temperatures cl@lwas calculated with the van’t Hoff relationship and
BA = K/B x B&K/B + [H+])
(9)
where B, = 4.2 x lop4 moles kg-l seawater.
Bicarbonate (KG) and boric acid (KL) dissociation constants compatible with the NBS
pH scale were obtained with Eq. 6 and table
3 in the work of Miller0 (1979), and [H+] is
calculated as 10--pH.
The term pl* in Eq. 4 presented some
potential difficulties in our analysis since
previous careful work at 25°C has produced
242
Notes
Table 1. CuCO,O complexation results. The parameters log &3, and log A, determined in each of our leastsquares analyses are shown together with 95% CL. reflecting the precision of each estimate.
Temp (“K)
BO
1% PI*
1% SWBI
278.15kO.l
288.15kO.l
298.15kO.l
308.15 aO.1
1.36
1.39
1.42
1.45
-8.74
-8.42
-8.11
-7.82
4.79,+0.01
4.85,+0.01
4.91,+0.01
4.98,kO.Ol
I%*estimates at variance by nearly a factor
of three. Our studies (unpubl.) using the NBS
pH scale and the Tris (free H+) scale over a
range of temperatures have demonstrated
that offsets between the NBS scale and free
H+ scale in seawater are small compared to
the observed uncertainties in PI*. In order
to assessthe role of p,* uncertainties in our
analysis, we examined the &3i results provided by a range of p,* estimates. The pi*
estimates used in our study follow from the
assessments of Van den Berg (19843) and
Paulson and Kester (1980) at 25°C plus the
estimate AH = 12.0 kcal mol-l (Baes and
Mesmer 198 1) for the first Cu2+ hydrolysis
step:
Cu2+ + H20 = CuOH+ + H+;
&* = [CuOH+][H+]/[Cu2+].
(10)
With the procedures outlined above, we used
Eq. 4 in nonlinear least-squares analyses
of our spectrophotometric titration data (A,
[C032-]T, [H+]). Best-fit estimates for A,, Al,
and Jl were obtained as in previous work
(Byrne and Miller 1985).
Our results (e.g. Fig. 1) indicated that bestfits (Table 1) of our data were obtained
through calculations consistent with the pi*
(25°C) work of Paulson and Kester (1980).
A linear least-squares analysis of our swpl
data with the van? Hoff relationship (Stumm
and Morgan 198 1) provides a concise summary of our formation constant results:
bA,
8.269+0.006
8.32940.008
8.39OkO.007
8.457kO.004
Our enthalpy results indicate that &r decreasesby about 25% between 25” and 5°C.
Alternative choices of PI* and AH (hydrolysis) have a minor influence on this conclusion. Using hydrolysis enthalpies equal to
9, 12, and 15 kcal mol-1 and the pl* (25°C)
value of Paulson and Kester (1980), we obtained best-fit nH estimates within the range
2.37 I AH I 2.63 kcal mol-l. Formation
constants obtained with the value PI* (25°C
0.7 M) = 10-‘.(j6 (Van den Berg 1984b) and
a hydrolysis enthalpy equal to 12.0 kcal
mall’ provided log sw@lvalues -0.04 units
smaller than our Table 1 results and a calculated AH value only 0.3 3 kcal smaller than
our best-fit (Eq. 12) estimate.
In addition to our direct examination of
the temperature dependence of &i, it is
possible to use our A, data (Table 1) to assessnH. For comparative purposes we have
also shown our best-fit sw@lresults (Fig. 2).
Although assumed pi* values consistent with
I
’
I
’
I
1
1
’
I
- 8.5
- 8.4
4.9 -
log &?, = 6.740 - 542.1(1/T).
- 8.2
(11)
4.8 -
The slope of Eq. 11 is equal to nH/2.303R,
where R is the gas constant. Consequently,
our enthalpy estimate for CuC030 formation appropriate to S = 35 seawater is calculated as
AH = (2.48k0.29)kcal mol-l
(12)
with indicated 95% C.L.
3.2
3.3
3.4
3.5
3.6
+x103
Fig. 2. Least-squares fits are shown for our formation constant data (,&I,), and for the product A, =
x &?, where E, is the CuCO,O molar absorptivity.
;he slope of log ,,$I1 vs. T-l is expected to be slightly
more negative than for log A, vs. T-l because t, is
expected to increase with decreasing temperature.
Notes
243
of Cu(I1) is
Table 2. The inorganic complexation
the PI* (25°C) estimate of Van den Berg dicalculated for S = 35 seawater at 25” and 5°C. For all
minished our calculated log swfll values by conditions
the carbonate alkalinity has been set equal
N 0.04 units, concomitant changes in log A 1 to 2.1 x 10m3eq liter-l. The carbonate ion concentrawere of the order of only 0.004. Since A, is tion is calculated with Eq. 8 with -log K; (25°C) =
identified as the product cl x sw@1,
a fit of our 9.109 and - log Ki (S’C) = 9.391. Copper carbonate
formation constants are calculated with Eq. 11. The
log A I data against 1/T should provide an CU(CO,),~formation constant used in this calculation
appropriate assessment of nH if c1is nearly at 25°C was taken from the work of Byrne and Miller
constant. It is generally observed that molar (1985): log s~2 = 7.58. At 5°C we assumed log Jz =
absorptivities increase slightly with de- 7.45. This assumption implies that m = 0 for the
creasing temperature. Consequently, the reaction CUCO,~ + COj2- = CU(CO,),~- in seawater.
best-fit to our log Al data shown in Fig. 2
Fraction of total inorganic Cu(U)
produces the estimate AH >, 2.5kO.4 kcal
25°C
25°C
mol-1 with 95% C.L. Although previous
(pH 8.2)
(pH 7.6)
(pAT.2)
(pAT.6)
Species
work provides no basis for direct compar- cuc0,0
78.3
58.2
80.2
73.2
isons with our seawater nH results, pre- CU2f
10.5
28.7
4.6
14.3
10.3
1.9
6.0
3.8
vious carbonate complexation studies pro- CuCl’ + cuso,o
5.7
4.4
3.1
2.0
duced the following estimates for Mg2+ and CuOH’
7.6
2.0
4.3
0.9
cu(co,),zCa2+ complexation in pure water (Martell
and Smith 1982):
A survey of enthalpy data for organic li~WWOsO) = 3 kcal mol-I, and
gands
indicates that temperature-induced
nH(CaCO,O) = 3 kcal mol-l.
changes in reaction 2 are of a similar magThe Fig. 2 fit of our formation constant nitude. For many organic complexation reresults indicates that the log Jl value at actions (Martell and Smith 1977, 1982) the
25°C consistent with our entire Jl data set enthalpy change associated with Eq. 2 is of
is log Jl = 4.92. This result is in good the order of - 5 kcal mol-* 5 AH I 0. This
agreement with the results of Byrne and modest increase in Cu2+-organic ligand afMiller (1985) at 25°C: log swpl = 4.88. Both finities with decreasing temperature is charof these estimates in turn are in good accord acteristically offset by concomitant enthalpy
with formation constants derived from work changes in organic ligand protonation conat 25°C in 0.7 M NaClO,: at 25°C in 0.7 M stants, -10 kcal mol-l I AH I -5 kcal
NaClO,, Byrne and Miller (1985) obtained mol-l, so that in most cases Cu2+--organic
the result ,& = [CUCO,~]/([CU~+]M,,,~- = ligand complexation should also exhibit
2.39 x 105, where M,,+
= [COJ2-] + modest decreases between 25” and 5°C.
[NaCO,-1. This result can be combined with Consequently, as a first-order assessment,
the assessment log J,/p’i
= -0.447 (Can- we conclude that reaction 1 will exhibit only
trell and Byrne 1987), yielding the result a small temperature dependence in seawalog swfll = 4.93. Alternatively, the estimate ter.
,G,= [CUCO,~]/([CU~+][CO,~-]) = 5.32 x lo5
Alan L. Soli
(Byrne and Miller 1985) plus the estimate
[co32-]/[co32-]T = 0.15 40.0 1 in seawater Collegium of Natural Sciences
(Miller0 and Schreiber 1982; Whitfield
Eckerd College
1979) produces the result log &3i = 4.90.
St. Petersburg, Florida 33733
Although the change in sw@lwith temRobert H. Byrne
perature in seawater is small, it should be
recognized that the temperature depen- Marine Science Department
dence of Kt2, the bicarbonate ion dissocia- University of South Florida
tion constant, is such that at constant pH St. Petersburg 3370 1
[C032-]T decreases by approximately a factor of two between 25” and 5°C. Conse- References
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Notes
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Submitted: 6 April 1988
Accepted: 30 June 1988
Revised: 7 October 1988