+ MODEL
ARTICLE IN PRESS
Earth and Planetary Science Letters xx (2006) xxx – xxx
www.elsevier.com/locate/epsl
Experimental measurement of boron isotope fractionation in seawater
Kateryna Klochko a,⁎, Alan J. Kaufman a , Wengsheng Yao b ,
Robert H. Byrne b , John A. Tossell c
a
Department of Geology, University of Maryland, College Park, MD 20742, USA
College of Marine Sciences, University of South Florida, St. Petersburg, FL 33701, USA
Department of Chemistry and Biochemistry, University of Maryland, College Park, MD 20742, USA
b
c
Received 17 November 2005; received in revised form 24 May 2006; accepted 28 May 2006
Editor: M.L. Delaney
Abstract
The boron isotopic composition of marine carbonates is considered to be a tracer of seawater pH. Use of this proxy benefits
from an intimate understanding of chemical kinetics and thermodynamic isotope exchange reactions between the two dominant
boron-bearing species in seawater: boric acid B(OH)3 and borate ion B(OH)−4 . However, because of our inability to quantitatively
separate these species in solution, the degree of boron isotope exchange has only been known through theoretical estimates. In this
study, we present results of a spectrophotometric procedure wherein the boron isotope equilibrium constant (11–10KB) is determined
empirically from the difference in the dissociation constants of 11B(OH)3 and 10B(OH)3 in pure water, 0.6 mol kg− 1 H2O KCl and
artificial seawater. Within experimental uncertainty, our results show no dependence of 11–10KB on temperature, but 11–10KB at
25 °C in pure water was statistically different than results obtained in solutions at high ionic strength. 11–10KB of the seawater
(S = 35, BT = 0.01 mol kg− 1 H2O) at 25 °C is 1.0272 ± 0.0006. This result is significantly larger than the theoretical value used in
numerous paleo-pH studies (11–10KB = 1.0194).
© 2006 Elsevier B.V. All rights reserved.
Keywords: boron isotope equilibrium constant; marine carbonates; seawater pH proxy
1. Introduction
In aqueous solution, the equilibrium distribution of
B(OH)3 and B(OH)4− is known to be strongly pH dependent, such that at values higher than 8.6 borate ion
dominates while at lower pH boric acid is the dominant
species [1]:
BðOHÞ3 þ H2 O⇔BðOHÞ−4 þ Hþ
ð1Þ
⁎ Corresponding author. Tel.: +1 301 405 4457; fax: +1 301 314
9661.
E-mail address: [email protected] (K. Klochko).
In modern seawater (pH = 8.2) borate ion comprises
∼ 28.5% of boron species (assuming the dissociation
constant of boric acid pKB = 8.597 (at 25 °C); [2]), representing ca. 6% of seawater alkalinity, a measure of the
capacity of the ocean to neutralize anthropogenic CO2
and thus resist changes in pH. Increases in pCO2 over the
next 50 years are predicted to overwhelm the ocean's
neutralizing capacity [3], resulting in declining pH and
potentially devastating effects on marine ecosystems,
especially corals and other calcifying organisms in
shallow oceanic environments. Boron isotope measurements of coral and foraminiferal carbonate over the past
century and through the recent geological past [4–15]
hold the promise of charting temporal pH changes, but this
0012-821X/$ - see front matter © 2006 Elsevier B.V. All rights reserved.
doi:10.1016/j.epsl.2006.05.034
EPSL-08220; No of Pages 10
ARTICLE IN PRESS
2
K. Klochko et al. / Earth and Planetary Science Letters xx (2006) xxx–xxx
requires quantitative estimates of isotope exchange reactions between aqueous species, as well as evaluation of
possible microenvironmental effects at the sites of
calcification within organisms [16,17].
Differences in molecular coordination and vibrational frequencies between boron species in solution control
the magnitude of isotope fractionation (α = 1 / 11–10KB)
[18]. The isotope exchange of 10B and 11B between the
two species is described by the reaction:
10
BðOHÞ3 þ11 BðOHÞ−4 ⇔11 BðOHÞ3 þ10 BðOHÞ−4
ð2Þ
with an isotopic equilibrium constant defined by:
11−10
KB¼
½11 B ðOHÞ03 d ½10 B ðOHÞ−4 ½10 B ðOHÞ03 d ½11 B ðOHÞ−4 ð3Þ
In 1977, Kakihana et al. [19,20] provided the first
theoretical estimate of the magnitude of isotope exchange between borate and boric acid (11–10KB = 1.0194
at 25 °C), based on reduced partition function ratio
calculations from spectroscopic data on molecular vibrations (Table 1). In spite of the fact that in a later
publication with new collaborators [21] it was suggested
that the boron isotope equilibrium constant could be
much larger than that predicted in 1977, it is the earlier
value [19] that has been consistently used in paleo-pH
reconstructions. Supporting the lower value, a recent
analysis based on experimental vibrational frequencies
from various sources using force field modeling reported
11–10
KB = 1.0176 ± 0.0002 [22]. In contrast, more rigorous theoretical treatments of the spectroscopic data [23–
25] using ab-initio molecular orbital theory have
indicated significantly higher 11–10KB values (ranging
from 1.0260-1.0267 at 25 °C) for these species.
An important drawback of both ab-initio calculations and calculations that rely on experimental
frequencies is that their outcome is extremely sensitive
to both the choice of vibrational frequencies and the
detailed theoretical methods applied to calculate
molecular forces. The most detailed overview and
comparison of various quantum methods of computation was presented in [26], which demonstrated that
11–10
KB could range between ∼ 1.020 and ∼ 1.050
depending on the data used and the theoretical
methods applied. The best way to avoid errors is to
follow one of the main rules of thumb in computational science, which requires the use of strictly
comparable methods for obtaining and treatment of
the data (e.g. experimental vibrational frequencies).
Several studies have been conducted to estimate
boron isotope fractionation that accompanies adsorption
of dissolved boron species on various substances such as
marine clay [27], a N-methyl-D-glucamine resin phase
[28], and boron bearing minerals [21] (Table 1). Others
[29] estimated boron isotope fractionation by fitting
through the inorganic calcite precipitation data [30].
However, to the best of our knowledge, although 11–10KB
has been determined in 0.6 mol kg− 1 H2O KCl [31], no
direct experimental 11–10KB value for seawater has yet
been reported.
The method applied here and in [31] is purely experimental and based upon fundamental observations of
chemical equilibrium. Experiments in this study were designed to directly determine the boron isotope equilibrium
constant over a range of solution conditions. These were
specifically considered in order to evaluate the effect of
different temperatures, medium composition and total
boron on the isotopic equilibrium constant. Thousands of
at-sea measurements of seawater pH [32–34] demonstrate
Table 1
Published estimates of 11–10KB in seawater
Estimates based on spectroscopic data on molecular vibrations
Treatment (T°C)
11–10
Empirical spectra and Force Field modeling (26.8 °C) [19,20]
Empirical spectra and Force Field modeling (26.8 °C) [22]
Ab-initio molecular orbital theory (25 °C) [23,24]
Ab-initio molecular orbital theory (25 °C) [25]
Evaluation of all of the above treatments [26]
1.0194
1.0176(2)
1.0260
1.0267
≥1.0300
KB
Estimates from adsorption and precipitation experiments
Method (T°C)
11–10
Boron adsorption on marine clay (25 °C) [27]
Boron mineral precipitation (25 °C) [21]
Best fit value for boron adsorption on resin (10 and 30 °C) [28]
Best fit value [29] for inorganic calcite precipitation [30]
1.0330(20)
1.0395(190)
1.0290
1.0260
KB
ARTICLE IN PRESS
K. Klochko et al. / Earth and Planetary Science Letters xx (2006) xxx–xxx
that the precision of spectrophotometric pH measurements
obtained via absorbance ratios is on the order of ±0.0004
units or better. In this work we have used precise spectrophotometric pH measurement procedures to accurately
determine 11–10KB from differences in the dissociation
constants of 11B(OH)3 and 10B(OH)3 [31].
2. Methods
2.1. Theory
Measurements of 11–10KB can be obtained through
pH observations in borate/boric acid buffers composed
exclusively of 11B on one hand and exclusively 10B on
the other. The pH of solutions containing B(OH)30 and
B(OH)4− can be written as:
!
½BðOHÞ−4 ð4Þ
pH ¼ pKB þ log
½BðOHÞ03 where pH = − log [H+], pKB = − log KB, KB is an
equilibrium constant appropriate to Eq. (1) and square
brackets denote total concentrations of B(OH)30 and
B(OH)4− in all forms (e.g., free B(OH)4− plus ion pairs).
The difference between pH measured in solutions of
11
B (11pH) and the pH of solutions containing solely
10
B (10pH) can be written as:
11
pH −10 pH ¼ p11 KB −p10 KB
þ log
½11 B ðOHÞ03 d ½10 B ðOHÞ−4 !
ð5Þ
½10 B ðOHÞ03 d ½11 B ðOHÞ−4 where
p11 KB −p10 KB ¼ logð11−10 KB Þ
ð6Þ
In the case that the 11B and 10B buffers are
prepared using identical procedures, and buffering by
H+ exchange couples other than B(OH)30 and B(OH)4−
(Eq. (1)) are negligible, the [11B(OH)30]/[11B(OH)4−]
and [10B(OH)30]/[10B(OH)4−] ratios in (Eq. (5)) are
identical. In this case, the final term in Eq. (5) is log
(1) = 0 and the equation can be written as:
DpH ¼ 11 pH −10 pH ¼ logð11−10 KB Þ
ð7Þ
Eq. (7) shows that 11–10KB is measured as a small
difference in pH between two solutions, one composed
using 11B, and the other composed using 10B. As such,
highly precise pH measurements are required. Very
precise measurements of pH can be obtained via spec-
3
trophotometric analysis of indicator absorbance ratios R
in the following form [31,35,36]:
R−e1
ð8Þ
pH ¼ pKI þ log
e2 −Rde3
where pKI quantifies the equilibrium characteristics of H+
exchange between protonated and unprotonated forms of
thymol blue, and e1, e2 and e3 are constants that are
dependent on the molar absorbance vs. wavelength characteristics of fully protonated (intense yellow color) and
fully unprotonated (intense blue color) forms of thymol
blue. The physical/chemical characteristics of the various
terms in Eq. (8) are presented in [36].
For solutions at constant temperature, ionic strength
and composition, pKI is constant and Eqs. (7) and (8)
can be combined in the following form:
11
R −e1
11−10
DpH ¼ logð
KB Þ ¼ log
ð9Þ
e2 −11 R d e3
10
R −e1
−log
10
e2 − R d e3
Eq. (9) shows that 11–10KB can be measured solely
through observations of thymol blue absorbance ratios:
A
435 A
R ¼ 596
ð10Þ
where 596A and 435A are absorbances at 596 and 435 nm
measured in buffered solutions that contain the thymol
blue indicator.
A rigorous treatment of 11–10KB determinations in
solutions buffered not only by B(OH)30/B(OH)4− but also
by hydrolysis (H2O ↔ H+OH−) and H+ exchange by the
spectrophotometric indicator (I2− + H+ ↔ HI−) is given
in [31]. Including the influence of hydrolysis and several
micromolar concentrations of indicator, it can be shown
[31] that
!
!
11
10
R −e1
R −e1
11−10
logð
KB Þ ¼ log
−log
e2 −11 Re3
e2 −10 Re3
ð11Þ
11 x
−log 10
x
where
x¼
½Naþ þ ½Hþ −Kw ½Hþ −1 −½I2− −IT
BT −½Naþ −½Hþ þ Kw ½Hþ −1 þ ½I2− þ IT
ð12Þ
Within Eq. (12), BT is the total boron concentration;
[Na+] is the solution concentration of Na+ that is added
as both NaOH and the sodium form of thymol blue; IT is
a total indicator concentration, and the concentration of
ARTICLE IN PRESS
4
K. Klochko et al. / Earth and Planetary Science Letters xx (2006) xxx–xxx
unprotonated indicator ([I2−]) is calculated using the
thymol blue equilibrium characteristics determined by
[36].
The final term in Eq. (11) is insignificant if [Na+] ≫
([H + ] − K w [H + ] − 1 − [I 2− ] − I T ) and (B T − [Na + ]) ≫
(− [H+] + Kw[H+]− 1 + [I2− ] + IT). Under our experimental
conditions, wherein [Na+] ranges between 0.004 and
0.02 mol kg− 1 H2O, and (BT − [Na+]) ranges between
0.006 and 0.03 mol kg− 1 H2O, pH ≈ 8.6, IT ≈ 3 ×
10− 6 mol kg− 1 H2O and ΔpH ≤ 0.03, the term log
(11x/10x) is smaller than 0.0001 and can be neglected in
calculations of 11–10KB.
2.2. Sample preparation and materials
2.2.1. Sample solutions
Boron is one of the major constituents of natural seawater and has a concentration of 4.4 ⁎ 10− 4 mol kg− 1 H2O
at a ∼ 34.8 salinity [37]. The spectrophotometric measurements in [31] were obtained at a total boron concentration
of 0.05 mol kg− 1 H2O, which is significantly higher than
the concentration of boron in natural seawater. To investigate the potential significance of polyborate formation
on the 11–10KB determinations, experiments were conducted at a range of total boron concentrations, including:
0.01, 0.025 and 0.05 mol kg− 1 H2O.
Insofar as the activity quotient for the reaction (2) is
very close to unity, the value for the isotope exchange in
this reaction (11–10KB) should be applicable over a wide
range of solution compositions and ionic strengths [31].
To evaluate this contention, experiments were conducted
in three different synthetic solutions: “pure water”,
0.6 mol kg− 1 H2O KCl and synthetic seawater. To assess
the effects of temperature on 11–10KB, experiments in
“pure water”, 0.6 mol kg− 1 H2O KCl and seawater were
performed at 25 and 40 °C.
Equimolal 10B(OH)3 and 11B(OH)3 in “pure water”
were prepared by weight in 15 MΩ Milli-Q water. Approximately 25 g (± 0.0005 1 s.d.) of each solution were
then weighed into 10 cm spectrophotometric cells. Equal
amounts of 1.0 M NaOH approximately 0.5 g (± 0.0001 1
s.d.), were then added (by weight) to each cell. These
procedures resulted in two solutions that were equimolal
in Na+, and whose concentrations of 10B and 11B were
identical. For all “pure water” solutions, BT ≈ 0.05 mol
kg− 1 H2O, [Na+] ≈ 0.02 mol kg− 1 H2O, and pH ≈ 8.6. In
each experiment usually four to six sample cells (n = 4–
6) were prepared for both isotopically labeled solutions.
Upon preparation each sample cell was promptly
transferred from the scale into a custom designed
thermostated cell holder, where the cells were left to
thermally equilibrate for at least 45 min.
A second set of solutions was prepared in 0.6 mol kg− 1
H2O KCl medium, resulting in an ionic strength close of
that of natural seawater. A single batch of 0.6 mol kg− 1
H2O KCl was prepared by weight in 15 MΩ Milli-Q
water. Similar to the procedure using “pure water”,
equimolal 10B(OH)3 and 11B(OH)3 solutions were prepared by gravimetric addition of boric acid and 1.0 M
NaOH to the 0.6 mol kg− 1 H2O KCl solution. For the
0.6 mol kg− 1 H2O KCl solutions, BT ≈ 0.01, 0.025 and
0.05 mol kg− 1 H2O.
Synthetic seawater was prepared using a standard
recipe [38] and had an ionic strength of 0.70 mol kg− 1
H2O and salinity of ∼ 35. The only boron in this
solution, as well as the solutions prepared using “pure
water” and 0.6 mol kg− 1 H2O KCl, came from the
isotopically labeled boric acid. A single batch of
synthetic seawater was prepared by adding seawater
constituent salts (NaCl, KCl, Na2SO4, CaCl2, MgCl2)
by weight to 15 MΩ Milli-Q water. The preparation of
equimolal 10 B(OH) 3 and 11 B(OH)3 solutions in
seawater was identical to the procedures used for
“pure water” and 0.6 mol kg− 1 H2O KCl. For the
synthetic seawater, BT ≈ 0.01 and 0.05 mol kg− 1 H2O.
2.2.2. Materials
Experimental reagents including boric acid as 10B
(OH)3 and 11B(OH)3 (99+ atom % pure 10B and 11B
respectively, 99.95% by weight boric acid), 1.0 M
NaOH, and the pH indicator thymol blue, were
obtained from Sigma-Aldrich. The boric acid and
KCl were stored in a desiccator containing P2O5 for
3 d prior to use. Salts used in this work (NaCl, KCl,
Na2SO4, CaCl2, MgCl2) were analytical grade reagents
also obtained from Sigma-Aldrich. We calculate that a
1% isotopic impurity of the 10B and 11B labeled boric
acids would result in a pH uncertainty of around
0.0001 units. Given 0.05% neutral impurity of the
labeled boric acids we calculate an additional
uncertainty of 0.00036 pH units. Either of these
uncertainties are less than our analytical precision.
2.3. Instrumental procedures
Absorbance measurements for determination of pH
in 10B(OH)30/10B(OH)4− and 11B(OH)30/11B(OH)4− buffers prepared as described above were obtained using
an HP 8453 and CARY 400 Bio UV–Vis spectrophotometers. Both instruments are equipped with
water-jacketed sample cell holders, which were thermostated at 25 or 40 °C within ± 0.1 °C and maintained
with a Neslab refrigerating circulator. First, the sample
was transferred to the instrument cell holder for a blank
ARTICLE IN PRESS
K. Klochko et al. / Earth and Planetary Science Letters xx (2006) xxx–xxx
5
absorbance measurement. Then, indicator was added to
the sample for simultaneous measurements of absorbances (λA) at wavelength λ = 730,596 and 435 nm.
When calculating the absorbance ratio R (Eq. (10)) the
absorbance measurement at 730 nm (730 A) was
subtracted from 596A and 435A as means of compensating for possible baseline shifts:
A− A
R ¼ 596 730
ð13Þ
435 A−730 A
the grand mean ΔpH was chosen as an upper bound
uncertainty estimate. The standard error of the mean
with 95% confidence interval (2σm) was determined as:
2d1s:d:
2rm ¼ pffiffiffi
ð14Þ
n
2.4. Data analysis
2rm V¼ 2:302dð11−10 KÞd 2rm
For each of the experimental conditions described
above (various BT, temperature and medium composition) four to six samples were prepared for each isotope.
For example, in seawater experiments (BT = 0.01 mol
kg− 1 H2O, at 25 °C), each sample was analyzed at least
five times, producing a maximal instrumental error of
± 0.0004 (1 s.d.). For n = 5 (number of samples of each
isotope) there were five values of ΔpH calculated using
Eq. (9) by pairing the mixtures. The grand mean ΔpH of
these five values was 0.0117 ± 0.0003 (1 s.d.). Instrumental contributions to uncertainties were only ± 0.0001
(1 s.d.). Therefore, the larger uncertainty associated with
The results from multiple experiments conducted
under identical experimental conditions were combined
in a weighted mean with the individual uncertainties
incorporated into a weighted error as [39]:
X x̄
X 1
weighted mean ¼
=
ð16Þ
2
ð2rm VÞ
ð2rm VÞ2
sffiffiffiffiffiffiffiffiffiffiffiffiffiffiffiffi
1
variation on mean ¼ P 1
ð17Þ
Table 2
KB determined for three different media at 25 and 40 °C
11–10
⁎ CARY 400 Bio UV–Vis double-beam spectrophotometer.
where n = number of samples for each isotope or number of pairs.
11–10
KB was calculated using Eq. (9) and the uncertainty (2σm) associated with the grand mean ΔpH was
propagated as:
ð15Þ
ð2rm ÞV 2
where x̄ are 11–10KB values obtained in each of the
multiple experiments under identical experimental
ARTICLE IN PRESS
6
K. Klochko et al. / Earth and Planetary Science Letters xx (2006) xxx–xxx
Fig. 1. Boron isotope equilibrium constant (11–10KB) determined in “pure water”, 0.6 mol kg− 1 H2O KCl and seawater (S = 35) over a range of
total boron concentrations (BT ≈ 0.01, 0.025 and 0.05 mol kg− 1 H2O) and two temperatures (25, and 40 °C). The error bars represent 95%
confidence interval (±2rm′ — standard error of a mean) of the analyses within each experiment. The results of multiple experiments (N > 1)
conducted under the same experimental conditions are presented as weighted mean of these results with the associated weighted uncertainty of
the mean (±2σm).
conditions and 2σm′ — are uncertainties associated with
each 11–10KB.
3. Results
11–10
KB results obtained over a wide range of
medium composition, total boron concentrations
(BT = 0.01, 0.025, 0.05 mol kg− 1 H2O) and 25–40 °C
temperature range were quite coherent (Table 2 and
Fig. 1). Spectroscopic measurements in “pure water”
media (BT = 0.05 mol kg− 1 H2O) yielded 11–10KB
values of 1.0308 ± 0.0023 and 1.0289 ± 0.0048 at 25
and 40 °C, respectively. 11–10KB values in 0.6 mol kg− 1
H2O KCl with identical BT at these two temperatures
were 1.0250 ± 0.0005 and 1.0262 ± 0.0035. Finally, 11–
10
KB values in synthetic seawater (BT = 0.05 mol kg− 1
H2O) were 1.0272 ± 0.0003 and 1.0269 ± 0.0027 at 25
and 40 °C, respectively. Higher errors were observed
for measurements at 40 °C than for the measurements at
25 °C, especially for the “pure” water solutions. Within
the experimental uncertainties, our 11–10 KB results
show no temperature dependence within the range
25–40 °C. However, considering only the 25 °C
analyses, there is a statistical difference between 11–
10
KB values obtained in “pure water” solution and at
higher ionic strength. It should also be noted that the
measurement uncertainties obtained using CARY 400
double-beam instrument were much smaller than those
obtained using HP 8453.
11–10
KB results obtained in 0.6 mol kg− 1 H2O KCl
and artificial seawater are very similar (Table 2 and
Fig. 1). Since approximately 44% of the total borate in
seawater is ion paired with Na+, Ca2+ and Mg2+ [37],
it can be concluded that the relatively weak ion paring
of borate with major seawater cations has a negligible
influence on 11–10 KB. Furthermore, the excellent
agreement in 11–10KB results obtained using total
boron concentration between 0.01 and 0.05 mol kg− 1
H2O indicates that polyborate formation has a
negligible influence on the equilibrium results
obtained in this work. In a recent publication [31] it
was pointed out that experiments conducted over a
range of boron concentrations would allow extrapolation of 11–10KB results to values appropriate at the
natural boron levels of seawater (416 μmol/kg
seawater). This appears to be unnecessary because
there is no statistically discernable trend in 11–10KB
results over a substantial range of boron concentrations. The likely mechanism behind this constancy in
11–10
KB is seen in the observation [40] that boron
fractionations into the four coordinate sites of B(OH)4−
and B3O3(OH)4− are very similar [31].
ARTICLE IN PRESS
K. Klochko et al. / Earth and Planetary Science Letters xx (2006) xxx–xxx
4. Discussion
Of greatest interest to oceanic pH reconstructions, the
KB result for artificial seawater [11–10KB = 1.0272
± 0.0006 (2σm′) at 25 °C (BT = 0.01 mol kg− 1 H2O)] is
significantly larger than the 1977 estimate of 1.0194
[19,20] and the spectral and the force field modeling
estimate of 1.0176 [22]. On the other hand, our empirical result is in agreement with the recent ab-initio
calculations [23–25], which indicate 11–10KB values
between 1.0260 to 1.0267.
The relationship between pH and boron isotopic
composition of borate in seawater (δ11Bborate), given
pKB = 8.597 at 25 °C [2], 11–10KB as either 1.0272
± 0.0006 (this study) or 1.0194 [19,20] and boron
isotopic composition of seawater (δ11Bsw = 39.5‰) can
be summarized as:
11–10
y11 Bsw −y11 Bborate
pH ¼ pKB −log 11
11−10
y Bsw −
K B y11 Bborate −1000 ð11−10 K B −1Þ
!
ð18Þ
The experimental 11–10KB value obtained in this work
results in a calculated relationship between δ11Bborate and
pH, that lies below the curve based on the previous
theoretical estimate (Fig. 2). It is important to recognize
that both curves intend to display the boron isotopic
7
composition of borate species (B(OH)4−) in seawater, not
the carbonate minerals that serve as a proxy for seawater
isotopic compositions.
Calibrations of the δ11B of carbonate materials vs.
pH have been conducted using controlled experiments
with inorganic calcite precipitates and both coral and
foraminifera species grown under constant pH conditions in modified seawater solutions [30,41–43].
While the coral analyses appear to conform to the
Kakihana et al. reference curve — calculated using the
most widely accepted pKB value of 8.597 [2] — it is
noteworthy that both foraminiferal and inorganic
carbonates had δ11B values that fell significantly
below the theoretical constraint for B(OH)4−. It has
been argued [44] that because these calibration
measurements broadly mirror the “shape” of the
Kakihana et al. curve, a constant offset at different
pH may be used to correct for each of the studied
species. In contrast, the borate–pH curve produced
from 11–10KB obtained in this study (1.0272 ± 0.0006)
fits best with the data from the controlled O. universa
measurements [41], but the slope of this curve is
steeper over the studied pH range, and hence does not
as faithfully mirror the control data. Slight differences
in pKB values are shown to have a profound effect on
the slope of these curves [25] with higher values
reversing the orientation of the two calculated slopes,
Fig. 2. δ11B of B(OH)−4 in seawater based on theoretical 11–10KB = 1.0194 [19, 20] and empirical 11–10KB = 1.0272± 0.0006 (2σm′ , n = 5) obtained in this
study at 25 °C (BT ≈ 0.01 mol kg− 1 H2O); and the results of the inorganic calcite precipitation experiments [30], cultured O. universa and G. sacculifer
foraminifera species [41,42] and cultured scleractinian corals A. nobilis and P. cylindrica [43]. The pHNBS values from [30,41,42] were recalculated to fit
the seawater pH scale (pHSWS = pHNBS − 0.14).
ARTICLE IN PRESS
8
K. Klochko et al. / Earth and Planetary Science Letters xx (2006) xxx–xxx
such that our empirical 11–10KB constraint would better
mirror the control data.
It is evident that the controls on the boron isotopic
composition of marine carbonates are more complex
than suggested by simple isotope exchange equilibrium
reactions. In this regard both vital effects and the
chemistry of seawater at the site of carbonate formation
warrant further investigation.
Boron isotope redistribution during biosynthesis of
carbonate is possible given the requirement that marine
calcifiers modify seawater in order to concentrate carbonate ion and maintain saturation at the site of calcification
[45,46]. Saturation is achieved by seawater vacuolization
and modification within the cytoplasm, which elevates
both pH and alkalinity. In foraminifera, photosynthesis
plays a critical role in elevating the pH through the uptake
of vacuole CO2 [45], while in corals endergonic enzymatic reactions that exchange protons for Ca2+ result in
higher pH at the site of calcification [47,48]. Micro-sensor
studies indicate that the pH of the calcifying fluid in the
foraminifera G. sacculifer rises to as high as 8.6 in daylight [17] relative to ambient seawater (pH = 8.2). Similarly, pH in the symbiotic coral Galaxea rises from 8.2 to
8.5 at the polyp surface and further to 9.3 in the calcifying
fluid [16]. It is not known whether boron is delivered to
the calcifying site in the form of neutral B(OH)30 or
charged B(OH)4−. If the neutral species are more likely to
pass through to the calcifying site, then 11B enriched
borate would form from the heavier boric acid via reequilibration due to higher pH of the site. Thus, elevated
pH in these biogenically modified solutions would drive
both chemical and isotopic equilibrium resulting in 11B
enrichment of borate ion by as much as 4–5‰.
Mineralogic differences between marine calcifiers may
also play an important role in boron isotope distributions.
While it has been concluded that the main boron-bearing
species incorporated into the carbonate lattice of aragonite
is the borate ion, with no observed fractionation [49,50],
in calcite up to 90% of the incorporated species may be
boric acid [50]. This could result in heavier boron isotopic
composition of calcite relative to aragonite provided no
other processes are involved (e.g. biological).
Both biological and mineralogical effects would
appear to result in enrichments of the heavy isotope in
carbonates, which is consistent with the larger value of the
boron isotope equilibrium constant empirically determined in this study (Fig. 2). Thus we presently find no
logical basis for offset corrections to lower measured δ11B
values in marine carbonates, based on the use of the 1977
equilibrium constant.
For the time being, while time series changes in pH are
valid [14,15], obtaining absolute values for paleo-pH
using the carbonate proxy remains a challenge. Existing
calibrations [30,41–43] are useful but limited to only two
species of foraminifera (G. sacculifer and O. universa)
and two species of relatively fast-growing branching
corals (Acropora nobilis and Porites cylindrica). Notably,
massive corals (e.g. P. lutea) with much slower growth
rate are more widely used in paleo-reconstructions and
should be the focus of new culture experiments. In our
view, progress towards the use of carbonates as absolute
pH proxies will require additional controlled culture and
micro-sensor studies of pertinent species and better understanding of boron speciation at the site of calcification
with respect to boron transport and incorporation into the
crystal lattice.
Empirical and culture experiments are the key to a
better understanding of past changes in oceanic pH, but
they will also be critical in charting the course of
anthropogenic effects on oceanic pH and marine
ecosystems through the next century.
Acknowledgments
This work was supported in part by NOAA research grant NAO40AR4310096 to R.H. Byrne, and
by NSF grant EAR 0001031 and DOE grant DE-FG
02-94ER14467 to J.A. Tossell. Additional support for
K. Klochko came from NSF grant EAR 040709-8545
to A.J. Kaufman. We would like to thank the editor
Dr. Peggy Delaney and three anonymous reviewers
for their helpful comments and suggestions. We are
also grateful to Dr. John Hayes for pointing us
towards this line of research, and to Dr. George Helz
and Dr. William McDonough for their guidance and
discussion on earlier versions of this manuscript.
References
[1] J.P. Hershey, M. Fernandez, P.J. Milne, F.J. Millero, The
ionization of boric acid in NaCl, Na–Ca–Cl and Na–Mg–Cl
solutions at 25 °C, Geochim. Cosmochim. Acta 50 (1986)
143–148.
[2] A.G. Dickson, Thermodynamics of the dissociation of boric acid
in synthetic seawater from 273.15 to 318.15 K, Deep-Sea Res. 37
(1990) 755–766.
[3] R.A. Feely, C.L. Sabine, K. Lee, W. Berelson, J. Kleypas, V.J.
Fabry, F.J. Millero, Impact of anthropogenic CO2 on the CaCO3
system in the oceans, Science 305 (2004) 362–366.
[4] J. Gillardet, C.J. Allegre, Boron isotopic compositions of corals:
seawater or diagenesis record? Earth Planet. Sci. Lett. 136 (1995)
665–676.
[5] N.G. Hemming, G.N. Hanson, Boron isotopic composition and
concentration in modern marine carbonates, Geochim. Cosmochim. Acta 56 (1992) 537–543.
[6] M.R. Palmer, P.N. Pearson, S.J. Cobb, Reconstructing past ocean
pH-depth profiles, Science 282 (1998) 1468–1471.
ARTICLE IN PRESS
K. Klochko et al. / Earth and Planetary Science Letters xx (2006) xxx–xxx
[7] M.R. Palmer, P.N. Pearson, A 23,000-year record of surface
water pH and pCO2 in the western equatorial Pacific Ocean,
Science 300 (2003) 480–482.
[8] P.N. Pearson, M.R. Palmer, Middle Eocene seawater pH and
atmospheric carbon dioxide concentrations, Science 284 (1999)
1824–1826.
[9] P.N. Pearson, M.R. Palmer, Atmospheric carbon dioxide over the
past 60 million years, Nature 406 (2000) 695–699.
[10] A. Sanyal, N.G. Hemming, G.N. Hanson, W.S. Broecker,
The pH of the glacial ocean as reconstructed from boron
isotope measurements on foraminifera, Nature 373 (1995)
234–236.
[11] A. Sanyal, N.G. Hemming, W.S. Broecker, Changes in pH in the
eastern equatorial Pacific across stage 5–6 boundary based on
boron isotopes in foraminifera, Glob. Biogeochem. Cycles 11
(1997) 125–133.
[12] A.J. Spivack, C.F. You, H.J. Smith, Foraminiferal boron isotope
ratios as a proxy for surface ocean pH over the past 21 Myr,
Nature 363 (1993) 149–151.
[13] A. Vengosh, Y. Kolodny, A. Starinsky, A.R. Chivas, M.T.
McCulloch, Coprecipitation and isotopic fractionation of boron
in modern biogenic carbonates, Geochim. Cosmochim. Acta 55
(1991) 2901–2910.
[14] B. Hönisch, N.G. Hemming, Surface ocean pH response to
variations in pCO2 through two full glacial cycles, Earth Planet.
Sci. Lett. 236 (2005) 305–314.
[15] C. Pelejero, E. Calvo, M.T. McCulloch, J.F. Marshall, M.K.
Cagan, J.M. Lough, B.N. Opdyke, Preindustrial to modern
interdecadal variability in coral reef pH, Science 309 (2005)
2204–2207.
[16] F.A. Al-Horani, S.M. Al-Moghrabi, D. de Beer, The mechanism
of calcification and its relation to photosynthesis and respiration
in the scleractinian coral Galaxea fascicularis, Mar. Biol. 142
(2003) 419–426.
[17] B.B. Jorgensen, J. Erez, N.P. Revsbech, Y. Cohen, Symbiotic
photosynthesis in a planktonic foraminifera Globigerinoides
sacculifer (Brady), studied with microelectrodes, Limnol.
Oceanogr. 30 (1985) 1253–1267.
[18] H.C. Urey, The thermodynamic properties of isotopic substances,
J. Chem. Soc. 57 (1947) 562–581.
[19] H. Kakihana, M. Kotaka, S. Satoh, M. Nomura, M. Okamoto,
Fundamental studies on the ion-exchange separation of boron
isotopes, Bull. Chem. Soc. Jpn. 50 (1977) 158–163.
[20] H. Kakihana, M. Kotaka, Equilibrium constants for boron
isotope-exchange reactions, Bull. Res. Lab. Nucl. React. 2
(1977) 1–12.
[21] T. Oi, J. Kato, T. Ossaka, H. Kakihana, Boron isotope
fractionation accompanying boron mineral formation from
aqueous boric acid–sodium hydroxide solutions at 25 °C,
Geochem. J. 25 (1991) 377–385.
[22] C. Sanchez-Valle, B. Reynard, I. Daniel, Ch. Lecuyer, I.
Martinez, J.-C. Chervin, Boron isotopic fractionation between
minerals and fluids: new insights from in situ high pressure-high
temperature vibrational spectroscopic data, Geochim. Cosmochim. Acta 69 (2005) 4301–4313.
[23] T. Oi, Calculations of reduced partition function ratios of
monomeric and dimeric boric acids and borates by the ab initio
molecular orbital theory, J. Nucl. Sci. Technol. 37 (2000)
166–172.
[24] T. Oi, S. Yanase, Calculations of reduced partition function ratios
of hydrated monoborate anion by the ab initio molecular orbital
theory, J. Nucl. Sci. Technol. 38 (2001) 429–432.
9
[25] Y. Liu, J.A. Tossell, Ab initio molecular orbital calculations for
boron isotope fractionations on boric acids and borates,
Geochim. Cosmochim. Acta 69 (2005) 3995–4006.
[26] R. Zeebe, Stable boron isotope fractionation between dissolved B
(OH)3 and B(OH)−4 , Geochim. Cosmochim. Acta 69 (2005)
2753–2766.
[27] M.R. Palmer, A.J. Spivack, J.M. Edmond, Temperature and pH
controls over isotopic fractionation during adsorption of boron
on marine clay, Geochim. Cosmochim. Acta 51 (1987)
2319–2323.
[28] A. Sonoda, Y. Makita, K. Ooi, N. Takagi, T. Hirotsu, pH-dependence of the fractionation of boron isotopes with N-Methyl-DGlucamine resin in aqueous solution systems, Bull. Chem. Soc.
Jpn. 73 (2000) 1131–1133.
[29] M. Pagani, D. Lemarchand, A. Spivack, J. Gaillardet, A critical
evaluation of the boron isotope-pH proxy: the accuracy of ancient
ocean pH estimates, Geochim. Cosmochim. Acta 69 (2005)
953–961.
[30] A. Sanyal, M. Nugent, R.J. Reeder, J. Bijma, Seawater pH
control on the boron isotopic composition of calcite: evidence
from inorganic calcite precipitation experiments, Geochim.
Cosmochim. Acta 64 (2000) 1551–1555.
[31] R.H. Byrne, W. Yao, K. Klochko, J.A. Tossell, A.J. Kaufman,
Experimental evaluation of the isotopic exchange equilibrium
10
B(OH)3 + 11B(OH)−4 = 11B(OH)3 + 10B(OH)−4 in aqueous solution, Deep-Sea Res., Part 1, Oceanogr. Res. Pap. 53 (2006)
684–688.
[32] R.H. Byrne, S. McElligott, R.A. Feely, F.J. Millero, The role of
pHT measurements in marine CO2-system characterizations,
Deep-Sea Res. 46 (1999) 1985–1997.
[33] T.D. Clayton, R.H. Byrne, Spectrophotometric pH measurements: total hydrogen ion concentration scale measurements and
at-sea results, Deep-Sea Res. 40 (1993) 2115–2129.
[34] T.D. Clayton, R.H. Byrne, J.A. Breland, R.A. Feely, F.J. Millero,
D.M. Campbell, P.P. Murphy, M.F. Lamb, The role of pH
measurements in modern oceanic CO2-system characterizations:
precision and thermodynamic consistency, Deep-Sea Res. 42
(1995) 411–429.
[35] R.H. Byrne, Standardization of standard buffers by visible
spectrometry, Analytical Chemistry 59 (1987) 1479–1481.
[36] H. Zhang, R.H. Byrne, Spectophotometric pH measurements of
surface seawater at in-situ conditions: absorbance and protonation
behavior of thymol blue, Marine Chemistry 52 (1996) 17–25.
[37] R.H. Byrne, D.R. Kester, Inorganic speciation of boron in
seawater, Journal of Marine Research 32 (1974) 119–127.
[38] A.G. Dickson, C. Goyet (Eds.), Handbook of Methods for the
Analysis of the Various Parameters of the Carbon Dioxide
System in Sea Water, U.S. Department of Energy, Oak Ridge,
TN, 1994.
[39] P.R. Bevington, D.K. Robinson, Data Reduction and Error
Analysis for the Physical Sciences, McGraw-Hill, 2003.
[40] T. Oi, Ab initio molecular orbital calculations of reduced partition
function ratios of polyboric acids and polyborate anions, Z.
Naturforsch. 55a (2000) 623–628.
[41] A. Sanyal, N.G. Hemming, W.S. Broecker, David W. Lea,
Howard J. Spero, G.N. Hanson, Oceanic pH control on the boron
isotopic composition of foraminifera: evidence from culture
experiments, Paleoceanography 11 (1996) 513–517.
[42] A. Sanyal, J. Bijma, H.J. Spero, D.W. Lea, Empirical relationship
between pH and the boron isotopic composition of G. sacculifer:
implications for the boron isotope paleo-pH proxy, Paleoceanography 16 (2001) 515–519.
ARTICLE IN PRESS
10
K. Klochko et al. / Earth and Planetary Science Letters xx (2006) xxx–xxx
[43] B. Hönisch, N.G. Hemming, A.G. Grottoli, A. Amat, G.N.
Hanson, J. Bijma, Assessing scleractinian corals as recorders for
paleo-pH: empirical calibration and vital effects, Geochim.
Cosmochim. Acta 68 (2004) 3675–3685.
[44] B. Hönisch, N.G. Hemming, Ground-truthing the boron isotopepaleo-pH proxy in planktonic foraminifera shells: partial
dissolution and shell size effects, Paleoceanography 19 (2004),
doi:10.1029/2004PA001026.
[45] J. Erez, The source of ions for biomineralization in foraminifera
and their implications for paleoceanographic proxies, Rev.
Mineral. Geochem. 54 (2003) 115–149.
[46] S. Weiner, P.M. Dove, An overview of biomineralization
processes and the problem vital effect, Rev. Mineral. Geochem.
54 (2003) 1–29.
[47] D. Allemand, E. Tambutte, J.-P. Girard, J. Jaubert., Organic
matrix synthesis in the scleractinian coral Stylophora pistillata:
role in biomineralization and potential target of the organotin
tributyltin, J. Exp. Biol. 201 (1998) 2001–2009.
[48] A. Cohen, T.A. McConnaughey, Geochemical perspectives on
coral mineralization, Rev. Mineral. Geochem. 54 (2003)
151–187.
[49] N.G. Hemming, R.J. Reeder, G.N. Hanson, Mineral-fluid
partitioning and isotopic fractionation of boron in synthetic
calcium carbonate, Geochim. Cosmochim. Acta 59 (1995)
371–379.
[50] S. Sen, J.F. Stebbins, N.G. Hemming, B. Ghosh, Coordination
environments of boron impurities in calcite and aragonite
polymorphs: an 11B MAS NMR study, Am. Mineral. 79 (1994)
818–825.