Limnol. Oceanogr., 41(3), 1996, 396-407
0 1996, by the American Society of Limnology and Oceanography, Inc.
Controls on iron(II1) hydroxide solubility in seawater:
The influence of pH and natural organic chelators
Kenshi Kuma, Jun Nishioka,’ and Katsuhiko Matsunaga
Department of Fisheries Oceanography and Marine Science, Faculty of Fisheries, Hokkaido University,
041 Japan
Hakodate,
Abstract
Iron solubilities of Fe(III) hydroxide in coastal and oceanic waters and in the ultraviolet (UV)-irradiated
seawaters over a pH range of 5.7-8.2 at 20°C were determined by a sim:Dle filtration (0.025 pm) involving
y-activity measurement of 59Fe.At pH ranges of 5.7-7.2 (coastal water) and of 5.7-7.6 (oceanic water), only
the Fe(OH),+ species is significant. The calculated solubility products, log*&,, for coastal and oceanic waters
were 4.8-5.0 and 4.4-4.6, respectively. The solubilities within the pH range of 7.8-8.2 are relatively independent of pH and aging time. Solubility in the oceanic water was about one order of magnitude lower than
that in the coastal water, and UV irradiation reduced solubility to 40.1 r-&I. The vertical profiles of ambient
Fe(III) solubility (pH 8.0-8.2) in oceanic waters have the following features in common: solubility in the
surface mixed layer is high and variable (0.3-0.6 nM), generally corresponding with the depth of high
chlorophyll a concentrations; solubility minima (0.15-0.2 nM) occur at a depth of 50-200 m. These results
suggestthat natural organic Fe(III) chelators exist in significant concentrations and control the dissolved iron
concentration in seawaters.
Although iron is an element of great biological and
geochemical importance, its oceanic chemistry, such as
inorganic speciation and organic complexes, is very complex and not yet fully understood. The dissolved inorganic
species of Fe(III) in seawater are predominantly the hydrolysis products Fe(OH),+, Fe(OH),O, and Fe(OH),(Byrne and Kester 1976; Zafiriou and True 1980; Motekaitis and Martell 1987). However, there are discrepancies between the estimated concentrations of dominant
hydroxo-complex species of Fe(III) and values (0. l-l 0
nM) for the thermodynamic solubility of Fe(III) hydroxide in seawater (e.g. Byrne and Kester 1976; Zafiriou and
True 1980; Zhuang et al. 1990). These discrepancies are
probably due to the existence of natural organic Fe(III)
chelators in seawater. Recent reports (Gledhill and van
den Berg 1994; Wells et al. 1995; Rue and Bruland 1995)
pointed out that iron complexation with organic ligands
is possible in oceanic waters, although most researchers
have argued that organic iron complexes do not seem to
be significant in the open ocean.
The oceanic distributions and biogeochemical behavior
of dissolved and particulate iron are controlled by comL Present address: Biology Department, Central Research Institute of Electric Power Industry, 1646 Abiko, Chiba, 270-l 1
Japan.
Acknowledgments
We thank K. Toya, K. Suzuki, T. Takabayashi, H. Kawakami,
and A. Katsumoto for help with the field and technical assistance. We also thank the crews of the Oshoro Maru and the
Hokusei Maru of Hokkaido University for help in sampling.
We are grateful to Eden Rue and Kenneth Bruland for offering
forthcoming data and to anonymous reviewers for comments
that helped improve the manuscript.
This research was partially supported by a Grant-in-Aid
(06640626) for Scientific Research from the Ministry of Education, Science and Culture of Japan.
396
plex interactions among input, internal cycling, and the
removal process. Over the last 10 yr, reliable iron concentrations in s’eawatershave been determined by the use
of clean sampling and more sensitive analytical methods
(Bruland et al. :I979; Landing and Bruland 1987; Wu and
Luther 1994). Martin and Gordon (1988) and Martin et
al. (1989) reported that dissolved iron (<0.4-pm fraction)
in the northeast subarctic Pacific exhibits a nutrient-type
profile, with depletion (0.02-o. 1 nM) in the surface water
and relatively constant higher concentrations (0.6-0.8 nM)
in deep waters. Active removal of iron by phytoplankton
in surface waters results in surface depletion of dissolved
iron (Martin and Gordon 1988). The oxidative decomposition of organic matter sinking into deeper waters
should regener,ate both iron and the major nutrients. In
addition, both surface-enriched (0.35-0.6 nM) and subsurface-depleted (0.02-0.2 nM) dissolved iron (<0.4-,
0.3-, or 0.2-pm fraction) profiles were observed by Bruland et al. (199 1, 1994) in the North Pacific central gyre
and by Wu and Luther (1994) in the western North Atlantic Ocean. Their dissolved iron concentrations in deep
waters are relatively constant (0.3-0.7 nM). Bruland et
al. (1991) suggested that eolian input caused the surface
dissolved iron maximum and that phytoplankton uptake
and particle scavenging might contribute to the subsurface dissolved iron minimum.
We used solubility measurements of Fe(III) hydroxide
in coastal and open-ocean surface waters, and in UVirradiated seawater over the pH range of 5.7-8.2 at 20°C
to study the do minant hydroxo-complex species of Fe(III)
and the existence of natural organic Fe(III) chelators in
seawater. The Fe(III) hydroxide solubilities were experimentally dete::mined by a simple filtration technique involving y-activity measurement of 5gFe.Furthermore, the
vertical distributions of Fe(III) hydroxide solubility in
oceanic waters, (O-1,500 m) were measured by the same
filtration method. The method provides indirect evidence
Fe(III) solubility in seawater
that natural organic Fe(III) chelators exist in significant
concentrations and control the dissolved iron concentration in seawater.
Methods
Sample collection and treatment-Seawater samples
were collected in the eastern Indian Ocean and the western North Pacific Ocean (Fig. l), and at a coastal region
near Hokkaido, Japan, in the northern Japan Sea with
ultraclean trace-element sampling techniques (5-liter GoFlo bottles on Kevlar line). Samples were drawn immediately from Go-Flo bottles into acid-cleaned polyethylene bottles. The open-ocean water samples were passed
through acid-cleaned 0.45-pm Millipore filters in a class
100 clean air bench on board as soon as the samples were
collected. The coastal seawater sample was filtered in a
clean room at the laboratory within 6 h of collection. The
filtrates in precleaned bottles were immediately frozen to
- 20°C until analysis in the laboratory. The freezing treatment was used to prevent possible microbial degradation
of natural organic Fe(III) chelators in the filtered seawater
until analysis of solubility.
A-l
P
AUSTRALIA
A
100"E
1lOOE
130°E
120'E
-T-I-
14OOE
50°N
.B-1
Dissolved iron, nutrient, and chlorophyll a concentrations-The iron concentrations in the defrosted seawater
samples collected from the coastal region and at stations
A-3 and B-2 in the open ocean (Fig. 1) were determined
with a graphite furnace-atom absorption spectrophotometer after Preconcentration by APDC/DDDC-chloroform
organic extraction (Bruland et al. 1979; Landing and Bruland 1987). The overall extraction efficiency, determined
by a 5gFeradiotracer, was 99 + 1% (n = 5). The precision
for iron analysis was 5.7% (n = 8) at the 1 nM level. The
iron concentrations were 1.5 nM in the coastal water and
5 1 nM in the open-ocean water. Major nutrient concentrations were determined by a Technicon autoanalyzer.
Chl a concentrations were determined by the fluorometric
method of Parsons et al. (1984).
Fe(III) hydroxide solubility measurements-The solubilities of Fe(III) hydroxide in the 0.45~pm-filtered seawater and in the UV-irradiated seawater were measured
over a pH range of 5.7-8.2 at 20°C by a simple filtration
technique involving y-activity measurement of 5gFe. A
UV-irradiation treatment was used to decompose the natural organic Fe(III) chelators in the filtered seawater by
photo-oxidative degradation of organic matter. A portion
of each sample was placed in an acid-cleaned quartz tube
and irradiated for 3 h by a 450-W high-pressure Hg-vapor
UV lamp.
Acid-cleaned 0.025~pm Millipore filters were used to
distinguish between dissolved Fe(II1) solubility and particulate Fe(II1) in seawater. Although the 0.025~pm-filterable fraction in the experiment of Fe(II1) hydroxide
solubility may contain smaller iron colloids, the filtration
technique is among the best approaches we presently have.
Colloids are traditionally defined as particles in the l1,OOO-nm size range (Vold and Vold 1966). However,
@B-2
40°N
l B-3
PACIFIC
NORTH
30°N
OCEAN
B
4
,
160°E
170°E
150°E
140"E
Fig. 1. Sampling locations in the eastern Indian Ocean, November 1993 (A- 1, 20”15’S, 106”OO’E;A-2, 20”15’S, 107’30’E;
A-3,20”15’S, 109”00’E), and in the western North Pacific Ocean,
June 1994 (B- 1, 44”00’N, 155”OO’E; B-2, 4 l”OO’N, 155”OO’E;
B-3, 35”00’N, 155”OO’E).
130'E
filters could retain particles that are smaller than the pore
size of the filter, depending on the characteristics of the
particles and the volume of water passing through the
filter. It is important to know what percentage of the iron
concentration in the filterable fraction is made up of small
iron colloids and soluble organic iron complexes, in addition to dissolved inorganic species of Fe(II1). Our seawater samples were coastal waters collected at a coastal
region in the northern Japan Sea and open-ocean surface
waters from the eastern Indian Ocean.
To examine the effect of aging time on Fe(II1) hydroxide
solubility in seawater, we added a small amount of radioactive ferric 5gFe (New England Nuclear Corp., No.
398
Kuma et al.
NEZ-037) solution, previously spiked with a small known
amount of stable ferric Fe, to 100 ml of 0.45-pm-filtered
seawater or filtered UV-irradiated seawater (20°C) in acidcleaned 125-ml polypropylene bottles. The final iron concentration was - 100 nM. In general, the addition of dissolved Fe(III) into seawater results in rapid hydrolytic
precipitation of metastable amorphous hydrous ferric oxide - Fe(III) hydroxide-which
slowly converts to more
stable solid phases with aging time (Crosby et al. 1983;
Wells and Mayer 199 1; Kuma and Matsunaga 1995).
The bottles containing the seawater solution and radiolabeled Fe(III) hydroxide were immersed in a constant-temperature water bath at 20&0.2”C. After standing
in the dark for 3 d and 1, 3, and 5 weeks at 2O”C, and
then for 4 d at 10°C each 7.5-ml sample aliquot was
filtered through a 0.025-pm Millipore filter and acidified
by addition of 10 ~1 of concentrated HCl to prevent adsorption of filtered Fe(III) on the wall of the collecting
vial. The y-activity of the 2.5-ml acidified sample filtrates
was measured in counting vials by means of a gamma
counter (Aloka ARC-30 1 B). The 0.025-pm-filtered Fe(III)
concentrations-Fe(III)
hydroxide solubility- were calculated from the counts (corrected with an average counting efficiency for 5gFe of the scintillation counter), the
volume of solution in the vial, and the amount of Fe per
count. The relative standard deviation of five filtration
replicates passing through each filter was 2.7% at an average of 122 cpm obtained by counting for 30 min. After
the solubility measurement, the pH of the seawater solutions was measured on the NBS scale with a combination pH electrode with a silver-silver chloride reference
electrode. The pH values of seawater solutions ranged
from 8.0 to 8.2.
To examine the effect of pH on the Fe(III) hydroxide
solubility in seawater, we added a small amount of radioactive ferric 5gFesolution to 60 ml of pretreated seawater (20°C) in precleaned bottles, in a quantity (- 100
nM for pH > 6 and - 200 nM for the pH range of 5.76.0) that exceeded the solubility level of Fe(III) hydroxide.
After the seawater aged in the dark for 1 d at 2O”C, small
volumes of diluted, subboiled HCl solution were added
a few times per day for a week at 20°C to obtain a desired
pH (5.7-6.5) or pH range (6.5-8.0). Milli-Q SP (low TOC)
water was used to dilute reagents and wash apparatus.
For pH of < 6.5, the desired pH was attained within a
week as a result of buffering systems in the seawater. After
standing in the dark for 3 and 5 weeks at 2O”C, the 7.5ml-sample aliquots were filtered through 0.025~pm filters
and then were subjected to the same experimental procedures described above. The pH of the seawater solutions was measured during and after the solubility experiment.
To determine the influence of preserving the seawater
samples by freezing at -20°C on the solubility measurement, we compared the solubility (3.2 nM) of Fe(III)
hydroxide aged for 3 weeks (20°C) in a defrosted coastal
water sample after freezing for 1 week to that (3.4 nM)
in the water analyzed without freezing just after collection. We found no significant difference in the solubilities
(aged for 3 weeks at 20°C) of a defrosted oceanic surface-
water sample that had been frozen for 2 months and a
sample that had been frozen for 1.5 yr (solubilities of 0.40
and 0.43 nM, respectively).
Fe(III) hydroxide solubility measurement in open-ocean
waters-Ambient Fe(III) hydroxide solubility in the water columns of the eastern Indian Ocean (O-800 m) and
western North Pacific Ocean (O-l ,500 m) were measured
by the same filtration technique and experimental procedures described above. After adding radiolabeled
5gFe(III) solution to 60 ml of the pretreated seawater to
make a final iron concentration of - 100 nM the samples
were stored in the dark for 3 weeks at 20°C. The sample
aliquots were then passed through 0.025~pm filters. In a
previous study (Kuma and Matsunaga 1995), the solubility of Fe(III) hydroxide in seawater (20°C) - pH 8 was
nearly constant with aging time after 1 week, presumably
because it was a’; saturation equilibrium. The pH values
of seawater solulions measured after the filtration experiment were within a range of 8.0-8.2.
Theory
Solubility-Tk.e
solubility product of the ferric solid
phase can be conveniently formulated as
*KS0 = [Fe3+][H+]-3.
(1)
The total concentration of dissolved Fe(III) in a solubility
equilibrium also depends on the equilibria among the
dissolved forms of Fe(III) in seawater. The following
equation, formu’.ated by Byrne and Kester (1976), can be
written to approximate the solubility behavior of ferric
oxide in seawater for pH > 5.0:
T[Fe(III)d] = *Ks,(*P,[H+]2 + *P2[H+] + *p3
+ *&[H-‘I-‘).
(2)
and
*p4
are
the
formation
constants
for
*Pl, v29 v3,
Fe(OH)2+, Fe(CbH)2+, Fe(OH),O, and Fe(OH),- , represented by *PI = [Fe(OH)2+][H+]/[Fe3+],
*f12 =
[Fe(OH)2+][H+]2/[Fe3+], *p3 = [Fe(OH)30][H+]3/[Fe3+],
and *p4 = [Fe(OH),- J[H+ J4/[Fe3+], and T[Fe(III)d] is the
concentration of total dissolved Fe(III). The stoichiometric solubility product, *Z&o, is calculated from the
equation T[Fe(III)d] = *&0*/32[H+] by using the results
of the filtration experiments in this study and the value
oflog*P,(-6.3 to -6.5)givenbyBaesandMesmer(1976),
Byrne et al. (1988), and Hudson et al. (1992). If a dissolved ferric hydroxo-complex species becomes dominant in a certain pH range, the slope of logarithmic
T[Fe(III)d] vs. pH in the pH range should approach a
theoretical slope. Fe(OH)2+, Fe(OH),+, Fe(OH),O, and
Fe(OH),- species have the theoretical slopes of - 2, - 1,
0, and 1.
Conditional ::tability constant -The solution equilibrium between f;:rric ion and organic ligand is further affected by specific inorganic complexation of the ferric ion
with seawater a.nions (such as OH-, Cl-, and C032-) and
Fe(III) solubility in seawater
399
by organic ligand association with H+ and major cations
(such as Ca2+ and Mg2+ ). These side reactions can dramatically reduce the free ferric ion and free organic ligand
concentrations and must be considered when defining a
conditional stability constant, XFeL, for a specific set of
solution conditions (Ringbom 1963). The conditional stability constant (assuming 1 : 1 complexation) with respect
to Fe3+ can be formulated as
K’ FeL= [FeL]/[Fe3+][L’].
(3)
[FeL] is the concentration of iron complexed by natural
organic ligand and [L’] is the concentration of ligand not
complexed with iron, free [L-l, and complexes of L with
the major cations, protons, and possibly other trace metals in seawater. The value of KfFeLis conditional on the
seawater composition (salinity, pH, competing trace metals). In this study, we estimated how strong natural organic ligands must be to be measured by the Fe(III) hydroxide solubility experiment.
Results and discussion
Fe(IZI) hydroxide solubility-The
solubilities of Fe(III)
hydroxide in surface seawaters and in the UV-irradiated
seawaters from the coastal region and from the open ocean
seem to decrease rapidly with aging time before day 7
and then to be nearly constant, presumably because they
are at saturation equilibrium (Fig. 2). The internal structure of Fe(III) hydroxide formed in seawater at 20°C
changes to more highly ordered phases through polymerization, dehydration, and formation of crystalline
products with aging time, resulting in decreasing solubility and dissolution rate of Fe(III) hydroxide with time
(Crosby et al. 1983; Wells et al. 1983; Kuma and Matsunaga 1995).
In Fig. 2, the Fe(III) solubility (2.8-2.9 nM) in coastal
surface water (non-UV irradiated) is about one order of
magnitude higher than that (0.24-0.27 nM) in oceanic
surface water (non-UV irradiated), probably due to the
coastal water having higher concentrations of organic ligands, which were possibly released by marine organisms
and riverine and atmospheric inputs. In addition, no detectable change of the Fe(III) solubility was observed when
water temperature was reduced from 20 to 10°C (Fig. 2).
The solubility (0.15-0.20 nM) in the UV-irradiated coastal water is about twice as high as that (0.07-0.09 nM) in
the UV-irradiated oceanic water. A possible explanation
for this difference is more incomplete decomposition of
organic ligands in the coastal water; it has been reported
that dissolved organic C (DOC) in seawater is not completely destroyed by UV irradiation (Kerr and Quinn 1980;
Mills et al. 1982), but the efficiency depends on exposure
time and the components of the organic substances.
The solubility results (Table 1) for the effect of pH are
presented graphically in Fig. 3 for coastal water (Fig. 3A)
and open-ocean surface water (Fig. 3B) as a plot of
-logT[Fe(III)d]
(mol liter-l) vs. pH. The line labeled
Fe(OH)2+ was determined with data having pH ranges of
5.7-7.23 (coastal water) and 5.7-7.63 (open-ocean water)
E
=
0.2
i
Aging time
(d)
Fig. 2. Effect of aging time (5 weeks at 20°C and then 4 d
at 10°C) on the Fe(III) hydroxide solubility (<0.025-pm fraction) in the 0.45~pm-filtered coastal water and open-ocean surface water (eastern Indian Ocean) and in UV-irradiated (3 h)
seawaters. Coastal water-o; open-ocean water-Cl; UV-irradiated coastal water-e; UV-irradiated open-ocean water-m.
Error bars on the solubility represent + 1 SD, estimated from
counting errors, where errors are larger than the symbol.
(Table 2) and is representative of the equation T[Fe(III)d]
= *K,,*p,[H+], assuming that within each pH range only
the species Fe(OH)2+ is significant. Published equilibrium
constants for ferric hydroxo-complexes indicate that only
the monomeric species, Fe(OH)2+, can exist in significant
concentrations in the pH range of 5.50-6.50 (Fox 1988).
The constant term in this equation, *Kso*p2, was calculated from an average of the values of T[Fe(III)d]/[H+]
within each pH range. Least-squares linear regression of
-logT[Fe(III)d]
vs. pH gives slopes that are closely consistent with the theoretical slope of - 1.O (Table 2).
The solubility product, *Kso, (Table 2) was calculated
with the value range of log*& (-6.3 to -6.5). For coastal
water (non-UV irradiated) aged for 3 and 5 weeks and
UV-irradiated water aged for 5 weeks (Fig. 3A), the calculated log*K,, was 4.8-5.0 in good agreement with values calculated from the results of simple filtration experiments (Byrne and Kester 1976) and dialysis experiments
(Kuma et al. 1992) for Fe(III) hydroxide in seawater.
However, the log*K,, (4.5-4.7) in UV-irradiated coastal
water aged for 3 weeks was smaller than the value for
water aged for 5 weeks. This result suggests that within
this acidified pH range, the dissolved iron concentration
in samples aged for 3 weeks was not at saturation equilibrium because of the slower proton-promoted dissolution rate of the amorphous phase formed in the UVirradiated coastal water. In addition, the log*Kso for both
the non-UV- and UV-irradiated open-ocean surface waters was 4.4-4.6 and did not change significantly by an
aging time of between 3 and 5 weeks (Fig. 3B).
These results suggest that a more active amorphous
phase or smaller particles with higher solubility and faster
dissolution rate would be formed in seawater containing
400
Kuma et al.
Table 1. Fe(III) hydroxide solubility in seawater at 20°C determined by simple filtration
with 0.025~pm filter.
Natural water
Aged 3 weeks
Aged 5 weeks
T[Fe(III)d]
PH
(nM)
T[Fe(III)d]
PH
UV-irradiated water
Aged 3 weeks
Aged 5 weeks
T[Fe(III)tl]
T[Fe(III)d]
PH
(nM)
b-W
PI-I
(nM)
5.70
6.00
6.26
6.52
6.80
7.17
7.53
7.77
7.87
8.19
47.74
42.4 1
14.95
8.27
3.27
1.43
1.27
0.89
0.52
0.61
5.76
6.00
6.24
6.58
6.89
7.29
7.53
7.70
7.87
8.06
34.39
21.29
4.20
1.69
1.10
0.30
0.32
0.25
0.19
0.19
Coastal water
5.72
6.06
6.35
6.62
6.78
7.19
7.57
7.69
7.91
8.16
73.02
45.30
13.40
6.23
4.00
3.23
2.53
1.94
1.46
1.85
5.71
6.04
6.32
6.59
6.75
7.22
7.58
7.75
7.83
8.17
5.72
5.98
6.21
6.60
6.99
7.30
7.63
7.75
7.92
8.14
28.24
15.69
5.36
1.91
0.85
0.55
0.37
0.27
0.24
0.25
5.77
5.99
6.21
6.60
6.99
7.32
7.63
7.79
7.97
8.14
88.36
5.72
25.6 1
54.42
6.01
23.88
20.24
6.29
6.76
7.46
6.54
3.63
5.00
6.78
1.97
2.89
7.23
0.69
7.49
1.93
0.74
2.33
7.65
0.49
7.87
1.28
0.52
8.18
1.66
0.42
Open-ocean surface water
34.42
5.76
16.93
18.25
5.97
11.22
6.21
6.20
4.68
2.92
6.56
2.12
6.90
0.94
0.95
7.31
0.55
0.19
0.36
7.53
0.17
7.75
0.3 1
0.15
0.30
7.92
0.20
8.11
0.28
0.13
a larger amount of natural organic Fe(III) chelators. Such
an active form, usually a very fine crystalline or amorphous solid phase with disordered lattice, may persist in
metastable equilibrium with the solution; this form is
more soluble than the stable solid phase (Stumm and
Morgan 198 1). The surface properties and colloidal stability of particles in natural waters are affected by naturally occurring dissolved organic substances (O’Melia
1987), probably resulting in the slowing of polymerization
and coagulation rates by dissolved organic ligands.
The Fe(III) solubilities in the non-UV- and UV-irradiated seawaters in the pH range of 7.8-8.2 are relatively
independent of pH and aging time (3-5 weeks) (Fig. 3).
The UV-irradiation treatment reduced the solubility (average value between pH 7.8 and 8.2) from 1.56 to 0.52
nM for coastal water and from 0.27 to 0.18 nM for oceanic
water. Particularly in the pH 7-8 range, the solubility
values in the oceanic water are similar to the Byrne and
Kester (1976) filtration values (0.05~pm filter and ultrafilter). In addition, the relative standard deviation of five
filtration replicates repeatedly passing through a 0.025pm filter was 1.3% at an average of 85 cpm (30 min of
counting). These results suggest that the filters in this
method are not sorbing significant amounts of the monomeric Fe(III) hydroxide species and organic Fe(III) chelator, which may be likely to sorb on filters made of
organic materials. If adsorption on filters were playing a
major role, the high precision for filtration replicates
through one filter would not be obtained.
The higher solubility in the coastal water is probably
due to the higher concentration of organic ligands or to
the presence of specific organic ligands with higher affinity
with ferric ion. The solubility in the UV-irradiated coastal
water is about three times that of the UV-irradiated oceanic water. The concentrations of Fe(OH),+ in equilibrium with Fe(III) hydroxide in coastal and oceanic waters
at pH 8.1 were found to be 0.25 and 0.1 nM from the
equations log[Fc(OH),+] = -pH - 1.5 and -pH - 1.9
(Table 2, Fig. 3), respectively. We assumed that the relatively constant solubility in a pH range of - 7.8-8.1 for
the UV-irradiat’ed oceanic water (aged for 5 weeks) is due
to Fe(OH),- for mation (Stumm and Morgan 198 1; Morel
and Hering 1993) and calculated a maximum value of
the formation constant for Fe(OH),-, log*@,. The value
of log *p4 was --22.6 to -22.8 from the equation
*P4= [Fe(oH)4-I[H+12*P21[Fe(OH)2+l,
using the solubility value {[Fe(OH),+] = 0.11 nM and
[Fe(OH), + ]+[ Fe(OH),-] = 0.19 nM at pH 8.06) and
The value of log*@, constructs an equation
log*B2.
log[Fe(OH),-] =pH - 18.2 (Table 2) as presented in Fig.
3B. However, we still do not know whether the constant
solubility withj n a pH range of - 7.8-8.1 is due to
Fe(OH),- formation or iron complex formation with natural organic ligands that remain in the UV-irradiated
oceanic water because of incomplete UV-photo-oxidation of organic ligands.
In this study, the minimum iron solubility in the UVq
Fe(III) solubility in seawater
irradiated open-ocean waters (pH 8.0-8.2) was 0.07-0.09
nM (Fig. 2), nearly consistent with the concentrations
(0.08-O. 13 nM) of Fe(OH),+ obtained by extrapolation
of the Fe(OH)2+ line to pH 8.0-8.2 (Fig. 3B, Table 2).
This result suggests that Fe(OH),+ is the dominant dissolved inorganic species of Fe(III) in seawater at the normal pH of seawater. The previous works claiming to demonstrate the existence of the Fe(OH),O species, which has
been suggested to be the dominant dissolved species in
seawater of -pH 8, are flawed because they do not account for the existence of natural organic Fe(III) chelators
that result in elevating the iron solubility. Zhuang et al:
(1990) found that the saturated concentration of dissolved
atmospheric iron in surface seawater collected from the
North Pacific Ocean and SargassoSeaand passed through
a 0.05-pm filter was 5-8 nM. In previous studies (Kuma
et al. 1992; Kuma and Matsunaga 1995), we reported that
the Fe(III) hydroxide solubility in seawater collected from
Funka Bay, Japan, determined by dialysis ( 1,000 Da) and
filtration (0.025 pm) experiments, is - 10 nM, possibly
because of the existence of natural organic Fe(III) chelators in coastal seawater. Sunda (1988-l 989) gave a value for the solubility of truly dissolved Fe(III) of 1.5 nM
at pH 8.2, whereas Wells (1988-l 989), in a review of iron
chemistry, argued for a maximum solubility of 0.1 nM.
Hudson et al. (1992) also concluded that their experimental results are consistent with a solubility of Fe(III)
that is < 1 nM at pH 8. The discrepancy in the different
values is probably due to the existence of natural organic
Fe(III) chelators in seawater.
A
Coastal
z
s:
5
water
6
7
8
I
I
I
P
Fe(OH),+
10.0
5
Ambient Fe(III) hydroxide solubility in open-ocean waters-In the eastern Indian Ocean, the oceanographic regime did not change over a longitudinal transect (Table
3). Fe(III) hydroxide solubility, Chl a concentration, nitrate plus nitrite concentration, and temperature for each
station are shown in Fig. 4. The vertical profile of nutrient
concentration and temperature showed oligotrophic subtropical waters over this transect. The Fe(III) solubility
profiles (Fig: 4A) indicated higher and variable solubility
(0.3-0.6 nM) in the surface mixed layer (O-50 m) with
high Chl a concentration (Fig. 4B), minima (-0.2 nM)
at depths of 100-200 m, and gradually increasing solubility with depth relatively in association with the increase in nutrient concentration (Fig. 4C). The vertical
profile of Fe(III) solubility in the UV-irradiated seawater
(Sta. A-3, Fig. 4A) is similar to that of the non-UVirradiated seawater, although the solubility decreased to
-0.1 nM because of incomplete UV-photo-oxidation of
organic ligands as described above.
A three-station transect in the western North Pacific
Ocean comprised the extent of the spatial study. The
oceanographic regime changed dramatically over this
transect, from oligotrophic water at lower latitude to eutrophic water at higher latitude (Table 4). A vertical crosssection of salinity at 155”OO’E (Fig. 5) indicated a subtropical water mass (North Pacific Current front) with
higher salinity at lower latitude (35-40°N), a subarctic
water mass with lower salinity at higher latitude (4244”N), and the boundary zone at midlatitude (40-42”N).
401
Oceanic
6
7
9
water
8
9
PH
Fig. 3. Effects of pH on the Fe(III) hydroxide solubilities
(<O.O25+m fraction) at 20°C in the 0.45-pm-filtered coastal
water and in UV-irradiated coastal water and in the 0.45+mfiltered open-ocean surface water and the UV-irradiated openocean water. Natural water aged for 3 weeks-O; natural water
aged for 5 weeks-e; UV-irradiated water aged for 3 wecksCl; UV-irradiated water aged for 5 weeks-m. Numbered lines
represent equations (Table 2) between the concentration of
Fe(OH),+ and pH for coastal and oceanic waters determined
from the Fe(III) solubility data vs. pH (Table 1). (1: -pH 1.5; 2: -pH - 1.8; 3: -pH - 1.9). Line Fe(OH),- represents
pH - 18.2 (Table 2).
The vertical profiles of Fe(III) solubility (Fig. 6A) indicate
higher solubility (0.3-0.5 nM) in the surface mixed layer
(O-50 m), generally corresponding with the depth of high
Chl a concentrations (Fig. 6B), and minima (-0.15 nM)
at depths of 50-200 m, just below the mixed layer. However, the minimum value of solubility had a slightly
northward increase (0.14-o. 17 nM) in subsurface water
and a northwardly shallower depth (5-50 m) of high solubility in surface water. The solubility levels in middepth
waters tended to increase with depth at lower latitude and
to be nearly constant within a range of 0.25-0.4 nM at
high latitude, probably because of a hydrographic change
over this transect. A vertical mixing or upwelling was
Kuma et al.
402
Table 2. Thermodynamic solubility values of Fe(III) hydroxide at 20°C (solubility) in the
coastal and oceanic seawaters and their UV-irradiated seawaters determined from the Fe(III)
solubility vs. pH data (Table 1). 1-For both the natural coastal water aged for 3 and 5 weeks
in the pH range of 5.7-6.78 and the UV-irradiated coastal water aged for 5 weeks in the pH
range of 5.7-7.17; 2-for the UV-irradiated coastal water aged Ibr 3 weeks in the pH range
of 5.7-7.23; 3 -for both the natural oceanic water aged for 3 and 5 weeksin the pH range of
5.7-7.63 and the UV-irradiated oceanic water aged for 3 and 5 weeks in the pH range of 5.7-
7.29.
Seawater(agingtin: e, pH range)
Solubility
2
1
Observed slope of
logT[Fe(III)d] vs. pH
1og*JL3
WFdOW2+ 1
(mol liter- l)
log*P,
3
- 1.2OkO.27
4.8-5.0
-1.11+0.1r3
4.5-4.7
- l.lOkO.26
4.4-4.6
-pH - 1.5
-
-pH - 1.8.
-
-pH - 1.9
-22.6 to -22.8
-
-
hWKW4-l
(mol liter - l)
observed at higher latitude, as shown in the vertical profiles of salinity (Fig. 5), nitrate plus nitrite concentration
(Fig. 6C), and temperature (Fig. 6D).
Another feature associated with the iron profiles was
pH - 18.2
the extreme diirerence in iron concentration between the
Fe(III) hydroxide solubility (<0.025-pm fraction) and the
dissolved iron concentration (~0.45~pm fraction) at stations A-3 and B-2 (Fig. 7). The dissolved iron concen-
Table 3. Hydrographic, chlorophyll a, nutrient, and iron data collected at stations in the
eastern Indian Ocean, November 1993. Fe(III) hydroxide solubility (<0.025-pm fraction)Fe(III); “dissolved” iron concentration (<0.45-pm fraction)-[FeId. (Not analyzed-N.)
Sta.
A-l
A-2
A-3
Depth
(m)
0
5
10
50
100
200
500
800
1,000
0
5
10
50
100
200
500
800
1,000
0
5
10
50
100
200
500
800
1,000
22.8 1
22.84
22.83
21.84
20.46
17.88
9.39
5.75
5.01
23.85
24.00
24.00
22.99
21.85
18.44
9.61
5.74
5.01
22.95
22.99
22.98
20.67
19.76
17.20
9.22
5.68
4.97
35.09
35.11
35.11
35.30
35.55
35.77
34.66
34.54
34.60
34.87
34.87
34.88
35.08
35.22
35.58
34.69
34.56
34.59
35.18
35.43
35.43
35.44
35.65
35.78
34.63
34.55
34.59
Chl a
(a liter-I)
N03
+N02
PO,
SiO,
Fe(III)
0.030
0.022
0.015
0.047
0.000
0.000
0.000
0.000
N
0.024
0.035
0.035
0.043
0.007
0.000
0.000
0.000
N
0.03 1
0.035
0.02 1
0.024
0.002
0.000
0.000
0.000
N
1.36
1.34
1.27
1.22
2.19
3.87
13.63
28.46
N
1.30
1.36
1.22
1.63
5.44
4.48
14.67
26.46
N
1.79
1.42
1.19
1.00
1.53
2.07
13.34
30.27
N
(PM)
0.17
0.17
0.16
0.20
0.28
0.36
1.22
2.61
N
0.15
0.16
0.16
0.18
0.54
0.47
1.32
2.61
N
0.16
0.15
0.17
0.19
0.23
0.31
1.15
2.55
N
2.94
2.94
3.01
4.56
4.75
4.82
6.50
66.98
N
2.92
2.88
3.01
3.72
7.80
5.14
7.08
59.69
N
3.19
2.81
3.72
3.01
3.65
3.52
4.75
61.47
N
0.38
0.40
0.50
0.39
0.29
0.20
0.25
0.51
N
0.50
0.59
0.44
0.38
0.21
0.20
0.24
0.36
N
0.37
0.32
0.29
0.38
0.21
0.31
0.31
0.53
N
[Fe]d
(nM)
N
N
N
N
N
N
N
N
N
N
N
N
N
N
N
N
N
N
1.07
0.14
0.07
0.80
0.43
1.08
0.73
0.64
N
403
Fe(III) solubility in seawater
trations were higher (0.3-1.1 nM) than the Fe(III) solubility at depths of 0 and 50-800 m and were lower (0.070.16 nM) than the Fe(III) solubility at depths of 5-10 m.
The higher dissolved iron concentration at O-m depth
may result from the presence of dissolved Fe(III), dissolved Fe(II), and colloidal iron in the <0.45-pm size
fraction through the atmospheric input of soluble iron in
open-ocean surface water (Zhuang et al. 1992). In addition, the colloidal iron phases at depths of 50-800 m may
be present in the 0.45~pm fraction, probably because of
the formation of colloidal-size material containing iron
through the microbiological decomposition of organic
matter. Wu and Luther (1994) reported that the colloidal
iron concentration in the 0.2-0.4-pm size fraction in the
water column of the western North Atlantic Ocean was
relatively high (0.2-0.3 nM) at depths of 50-750 m and
decreased to below the detection limit at depths <750
m. At depths of 5-10 m, depletion of dissolved iron,
despite higher Fe(III) solubility, may result from active
removal of dissolved iron by phytoplankton and high
adsorption of dissolved and colloidal iron on the surface
of marine particles such as clay minerals and organic
materials. The dissolution kinetics of colloidal iron
(~0.45~pm fraction) may be important in estimating iron
availability to phytoplankton.
Martin and Gordon (1988) and Martin et al. (1989)
reported that dissolved iron (<0.4-pm fraction) in the
northeast subarctic Pacific exhibits a nutrient-type profile,
with depletion (0.02-0.1 nM) in the surface water and
relatively constant higher concentrations (0.6-0.8 nM) in
deep waters. In addition, surface-enriched and subsurface-depleted dissolved iron profiles were observed by
Bruland et al. (199 1, 1994) in the North Pacific central
gyre and by Wu and Luther (1994) in the western North
Atlantic Ocean. The dissolved iron (<0.3-pm fraction)
profile from the North Pacific showed that surface mixedlayer concentrations are as high as 0.37 nM, but decrease
sharply through the upper seasonal thermocline to a subsurface minimum of only 0.02-0.05 nM at depths of 75100 m. Dissolved iron at depths < 100 m exhibits a nutrient-type profile with a relatively constant value (0.30.45 nM) in deep waters. In the western North Atlantic,
dissolved iron ( < 0.2~pm fraction) concentrations were
much higher in surface waters (-0.6 nM) and decreased
sharply through the euphotic zone to a subsurface minimum of 0.2-0.3 nM at depths of 30-90 m. Below 100
m, dissolved iron gradually increased to 0.7-0.8 nM at
1,000-m depth.
The surface-enriched and subsurface-depleted dissolved iron profiles are remarkably similar to our ambient
Fe(III) hydroxide solubility profiles. This result suggests
that a substantial fraction of the dissolved iron coming
from atmospheric sources could be retained in the oligotrophic surface mixed layer despite the biological demand for iron. Active biological iron removal in the surface layer could result in the nutrient-type profile. The
vertical distribution of natural organic Fe(III) chelators
is one of the most important factors controlling dissolved
iron in oceanic waters.
The higher Fe(III) solubility in the surface mixed layer
Fe(III) hydroxide solubility (Fe, nM)
0.0
0.2
0.4
0.2
0.0
0.4
0.0
0.2
0.4
0.6
A-2
I
a
I
I
I
C
.
NO,+NO,(pM)
600
900
t
1000 -
B
1
.
Fig. 4. Vertical profiles of Fe(III) hydroxide solubility, chlorophyll a concentration, nitrate+nitrite concentration, and temperature at stations A- 1 (0), A-2 (A), and A-3 (Cl) (UV-irradiated
seawaterat each depth-m in the easternIndian Ocean. Error
bars on the solubility represent + 1 SD, estimated from counting
errors, where errors are larger than the symbol.
is probably due to higher concentrations or stronger affinity of natural organic Fe(III) chelators that were possibly released metabolically by phytoplankton and bacteria. In general, DOC concentration in the open ocean
is at a maximum at the surface and decreasessharply with
depth (Tanoue 1992; Tupas et al. 1994). In addition,
Coale and Bruland ( 1988,199O) reported that the stronger
Cu-complexing organic ligand concentration is generally
high and variable (l-3 nM) at the depth of high primary
production (~200 m). They observed a concomitant
maximum in the Cu-complexing ligand and primary production and interpreted it as evidence of a phytoplankton
source.
Laboratory culture experiments also have indicated that
marine phytoplankton and bacteria can produce extracellular substances, such as siderophores, that have a strong
Kuma et al.
404
Table 4. As Table 3, but in the western North Pacific Ocean, June 1994.
Chl a
Sta.
B-l
B-2
B-3
Depth
(m)
0
5
10
50
100
200
500
800
1,200
1,500
0
5
10
50
100
200
500
800
1,200
1,500
0
5
10
50
100
200
500
800
1,200
1,500
T
(“0
8.35
8.35
8.35
4.3 1
1.90
2.61
3.14
2.76
2.33
N
13.02
13.02
13.02
10.18
7.36
5.75
3.81
3.01
2.48
N
21.46
21.46
21.46
18.42
16.71
12.54
4.92
3.70
2.7 1
2.26
s
(73
32.96
32.95
32.96
33.07
33.19
33.63
34.12
34.3 1
34.44
N
33.93
33.92
33.92
33.79
33.83
33.75
34.06
34.32
34.43
N
34.55
34.55
34.54
34.60
34.70
34.44
34.0 1
34.26
34.43
34.5 1
kg li-
ter-I)
0.477
0.345
0.37 1
0.138
0.043
0.016
0.018
0.025
N
0.02 1
0.949
0.90 1
0.932
0.148
0.436
0.013
N
0.008
0.00 1
0.003
0.216
0.188
0.183
0.388
0.047
0.007
N
0.005
N
N
chelating affinity for copper, iron, and other metals (e.g.
Trick et al. 1983; Trick 1989; Reid and Butler 199 1).
However, these extracellular substances have not been
found in many phytoplankton and bacteria. In this study,
the Fe(III) hydroxide solubility in the surface mixed layer
had no significant correlation with Chl a concentration.
The solubility in the surface mixed layer in the eastern
Indian Ocean and western North Pacific Ocean was nearly
the same at all stations, although the Chl a concentrations
in the mixed layer in the North Pacific were variable (0.21.O pg liter- ‘; Fig. 6B) and about one order of magnitude
higher than those in the Indian Ocean (0.02-0.05 pg liter-l, Fig. 4B). If the strong ligand is produced by phytoplankton, it seemsunlikely that all phytoplankton would
produce similar amounts of the same ligand. The distribution of organic ligand probably reflects distributions of
particular source species because many phytoplankton
species occupy relatively distinct vertical and regional
ranges in the surface water (Taylor and Waters 1982).
Coale and Bruland (1990) suggested that the distribution
of organic ligand would not be expected to track trends
in total productivity
or Chl a if such a scenario is operative.
No3
+NOz
PO,
14.92
14.41
12.39
23.82
29.80
26.76
44.66
39.80
N
31.31
4.85
4.20
4.83
15.13
11.28
21.44
N
41.63
46.33
37.42
0.47
0.38
1.70
3.79
6.24
13.63
N
33.05
N
N
(CLW
1.31
1.30
1.23
2.04
2.16
2.18
3.12
2.91
N
2.37
0.48
0.53
0.50
1.02
0.93
1.64
N
2.81
3.07
2.66
0.12
0.13
0.16
0.33
0.54
1.07
N
2.68
N
N
SiO,
Fe(III)
14.83
15.42
13.84
43.87
46.60
43.70
63.90
69.83
0.34
0.42
0.24
0.17
0.17
0.36
0.29
0.23
82?3
10.40
10.65
9.39
20.60
19.05
37.8 1
N
52.96
80.17
80.40
4.68
4.56
5.26
7.27
9.40
21.35
N
77.02
N
N
0.32
0.29
0.34
0.37
0.15
0.20
0.21
[Fe]d
0-W
N
N
0.42
0.32
0.27
0.36
0.40
0.40
0.52
0.14
0.16
ON43
k
N
N
N
N
N
N
N
N
N
N
N
0.39
0.16
0.14
0.30
0.90
0.83
N
1.oo
N
N
N
N
N
N
N
N
N
N
N
N
Odate et al. (1990) reported the distribution of cyanobacteria and other picophytoplankton (0.2-2.0-pm size
fraction) in the western North Pacific Ocean (36. 5-44.0°N,
155.O”E) in June 1989, which is the same transect and
season we used. Cyanobacteria were most abundant in
the surface of subtropical water (36.5-38.O”N) and less
abundant in subarctic water (39.5-44.O”N). However, the
cell density of other picophytoplankton was low in the
subtropical area and high in the subarctic area, and the
Chl a concentration accurately represented the abundance
other than cyanobacteria. In adof picophytoplankton
dition, the Chl a contribution by the different size fractions (<2 pm , Z-1 0 pm, > 10 pm) of the phytoplankton
communities in the western North Pacific showed regional and tem;?oral variations (Odate and Maita 19881989). Therefore, it is difficult to speculate as to the nature
of organic Fe(III) chelators in the surface layer until the
specific ligand or class of ligands and its source have been
identified.
The presence of solubility minima at narrow depth
ranges just below the surface mixed layer indicates that
the organic ligands produced are consumed or degraded
in the surface layer. The possible destruction mechanisms
Fe(III) solubility in seawater
Latitude
44
43
42
41
40
405
(ON)
39
38
Fe(III) hydroxide solubility (Fe, nM)
37
36
35
34
;;o
,
Oi2s
,
of
0.2
0.a
0.4
T
400
3
z
600
3
800
1000
1200
1400
1600
I
,
,
,
Chl a (,ug liter-‘)
0.0 0.2 0.4 0.6 0.8
I
,
I
NO,+NO,
0
10
20
B-3
.
(PM)
30
_
40
0
5
I
* 1 ,
Temp. (“C)
10 15 20
25
P
200 I400
,
8
600
800 ,
t
B-l
B!2
BT3
Fig. 5. Vertical cross-section of salinity at 155”OO’Ein the
western North Pacific Ocean, indicating a subtropical water mass
(North Pacific Current front) with higher salinity at lower latitude (35-4O”N), subarctic water mass with lower salinity at higher latitude (42-44”N), and the boundary zone at midlatitude
(40-42”N).
include photo-oxidation-reduction reactions in sunlit surface waters and heterotrophic absorption and subsequent
biological oxidation. The existence of organic Fe(III) chelators in the surface waters may be strongly related to the
photoreduction of Fe(III), probably through a photoinduced ligand-to-metal charge transfer reaction under sunlight in marine systems (e.g. O’Sullivan et al. 199 1; Miller
and Kester 1994; Kuma et al. 1995). Furthermore, the
photoproduction of Fe(II) in seawater may play a role in
making iron bioavailable (Johnson et al. 1994; Miller and
1400
I
1600
B
1.1
.I,,,
c
,,,,,,,,,
66
Fig. 6. Vertical profiles of Fe(III) hydroxide solubility, chlorophyll a concentration, nitrate+ nitrite concentration, and temperature at stations B-l (0), B-2 (A), and B-3 (0) in the western
North Pacific Ocean.
Fe(III)
0.1
aI-
hydroxide solubility
and
dissolved iron concentration
2
(Fe. nM)
2.cisaI-
IE 2oa
5
8
Q 400
Fig. 7. Vertical profiles of Fe(III) hydroxide solubility
(<0.025-pm fraction) (lJ,A) and dissolved iron concentration
(<0.45-pm fraction) (0) at station A-3 in the eastern Indian
Ocean and at station B-2 in the western North Pacific Ocean.
Error bars on the dissolved iron concentration represent + 1 SD,
estimated from the precision for iron analysis, where errors are
larger than the symbol.
600
800
1000
1600’
’
’ ’
’ ’ ’
’
’
a 1
406
Kuma et al.
Kester 1994). The subsequent increasing solubility in
middepth waters may be due to the organic Fe(III) chelators released through decomposition of organic matter.
The identity of organic ligands in middepth waters may
be different from those in surface waters. Coale and Bruland ( 1988, 1990) observed at least two Cu-complexing
ligand classes that have distinctly different vertical distributions and binding strengths in vertical profiles from
the northeast Pacific. The stronger ligands class (1.5-3
nM) is present in the surface mixed layer, generally corresponding with the depth of high primary production,
and the weaker class is present at higher concentrations
(8-10 nM) throughout the water column and has no apparent structure to its vertical distribution.
Fe(III) chelators exist in significant concentrations and
control the dissolved iron concentration in oceanic waters, the chemical composition and vertical distribution
of the organic ligands are still unknown. A future challenge will be to search oceanic regions that may have
surface waters with higher Fe(III) hydroxide solubility
that is strongly related to the biological activity and to
qualitatively and quantitatively measure the organic ligands in seawater. These natural organic ligands may play
an important role in the biological availability of iron in
oceanic waters.
References
BAES,C. F., AND R. E. MESMER. 1976. The hydrolysis of catConditional stability constant-In our study, we estiions. A critical review of hydrolytic species and their stamate how strong ligands must be to be measured by Fe(III)
bility consta:lts in aqueous solution. Wiley.
hydroxide solubility. For at least half of a ligand to be
BRULAND, K. W.,J. R. DONAT, AND D. A. HUTCHINS. 1991.
detected, [FeL]/[L’] 2 1 is necessary. Thus, K’&Fe3+]
Interactive influences of bioactive trace metals on biological
2 1 is, by definition, [FeL]/[L’] = KFeL[Fe3+] (Eq. 3). At
production in oceanic waters. Limnol. Oceanogr. 36: 15551577.
pH 8.15 and with log *KS, = 4.4-4.6 (for oceanic water,
-,
R.P.FRANKs,G.A. KNAUER,ANDJ.H. MARTIN. 1979.
Table 2), [Fe3+] is calculated to be - 10A20M from Eq.
Sampling and analytical methods for the determination of
1. Therefore, ligands with KFeL 2 l/l Om20= 1020 M-l
copper, cadmium, zinc, and nickel at the nanogram per liter
would be detected. At KFeL = 1020M- l, half of the ligands
level in seawater. Anal. Chim. Acta 105: 223-245.
would be detected, and >90% would be detected at KtFeL
-,
K. J. ORIANS,AND J. P. COWEN. 1994. Reactive trace
2 102’ M-l .
metals in the stratified central North Pacific. Geochim. CosRecently, Rue and Bruland (1995) reported that two
mochim. Acta 58: 3 17 l-3 182.
ligands that add up to - 2 nM are present in surface waters
BYRNE,R. H., AI\D D. R. KESTER. 1976. Solubility of hydrous
in the central North Pacific, with the weaker of the two
ferric oxide and iron speciation in seawater. Mar. Chem.
ligands having KFeL = 1021.5M-l. This result suggests
4: 255-274.
-,
L. R. Ku MP,AND K. J. CANTRELL. 1988. The influence
that essentially all of the ligand should be detected with
of temperature and pH on trace metal speciation in seathe Fe(III) hydroxide solubility method we used. Howwater. Mar. Chem. 25: 163-18 1.
ever, there are differences between these Fe(III) hydroxide
COALE,
K. H., AND K. W. BRULAND. 1988. Copper complexsolubility data (Tables 3, 4) at the surface waters in our
ation in the northeast Pacific. Limnol. Oceanogr. 33: 1084study and the concentrations of Fe(III)-complexing or1101.
ganic ligand observed at other sites (Gledhill and van den
-,AND--.
1990. Spatial and temporal variability in
Berg 1994; Rue and Bruland 1995). A possible explacopper complexation in the North Pacific. Deep-Sea Res.
nation for this difference is that the solubility data may
37: 3 17-336.
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Submitted: 22 March 1995
Accepted: I7 October 1995
Amended: 27 November 1995