of 6
EARTH, ATMOSPHERIC,
AND PLANETARY SCIENCES
Catalysis and chemical mechanisms of calcite
dissolution in seawater
Adam V. Subhas
a,1
, Jess F. Adkins
a
, Nick E. Rollins
b
, John Naviaux
a
, Jonathan Erez
c
, and William M. Berelson
b
a
Division of Geological and Planetary Sciences, California Institute of Technology, Pasadena, CA 91125;
b
Department of Earth Sciences, University of
Southern California, Los Angeles, CA 90089; and
c
Institute of Earth Sciences, Hebrew University of Jerusalem, Jerusalem 9190401, Israel
Edited by Mark H. Thiemens, University of California, San Diego, La Jolla, CA, and approved June 26, 2017 (received for review March 6, 2017)
Near-equilibrium calcite dissolution in seawater contributes sig-
nificantly to the regulation of atmospheric CO
2
on 1,000-y
timescales. Despite many studies on far-from-equilibrium disso-
lution, little is known about the detailed mechanisms responsi-
ble for calcite dissolution in seawater. In this paper, we dissolve
13
C-labeled calcites in natural seawater. We show that the time-
evolving enrichment of
δ
13
C in solution is a direct measure of
both dissolution and precipitation reactions across a large range
of saturation states. Secondary Ion Mass Spectrometer profiles
into the
13
C-labeled solids confirm the presence of precipitated
material even in undersaturated conditions. The close balance
of precipitation and dissolution near equilibrium can alter the
chemical composition of calcite deeper than one monolayer into
the crystal. This balance of dissolution–precipitation shifts signifi-
cantly toward a dissolution-dominated mechanism below about
Ω =
0
.
7. Finally, we show that the enzyme carbonic anhydrase
(CA) increases the dissolution rate across all saturation states, and
the effect is most pronounced close to equilibrium. This finding
suggests that the rate of hydration of CO
2
is a rate-limiting step
for calcite dissolution in seawater. We then interpret our dissolu-
tion data in a framework that incorporates both solution chem-
istry and geometric constraints on the calcite solid. Near equilib-
rium, this framework demonstrates a lowered free energy barrier
at the solid–solution interface in the presence of CA. This frame-
work also indicates a significant change in dissolution mechanism
at
Ω =
0
.
7, which we interpret as the onset of homogeneous etch
pit nucleation.
mineral dissolution
|
isotope geochemistry
|
oceanography
|
catalysis
T
he production and dissolution of calcium carbonate miner-
als provide a crucial link between the marine carbon and
alkalinity cycles. The ocean has absorbed about 25 to 30
%
of
anthropogenic
CO
2
emissions, dropping mean surface ocean
pH since the industrial era (1). As ocean pH decreases, sed-
imentary carbonate minerals will dissolve to compensate for
the loss of buffering capacity, eventually restoring atmospheric
p
CO
2
to about its preindustrial level (2, 3). This reaction will
mostly take place in the deep ocean, where the calcite satura-
tion state
Ω = [Ca
2+
][CO
2
3
]
/
K
sp
<
1
. Most of the deep ocean
is only mildly undersaturated such that pelagic dissolution is pri-
marily a near-equilibrium phenomenon. However, attempts to
quantify the relationship between calcite dissolution rate and
are highly variable between different studies, both in functional
form and absolute value (4–9).
With the exception of very early work by Berner and Morse
(7), few studies have attempted to unpack the chemical species
responsible for calcite dissolution in seawater. Many studies
choose instead to derive empirical relationships between sat-
uration state and dissolution rate (5, 6, 10, 11). In contrast,
freshwater and dilute solution dissolution studies have made
large advances in identifying key chemical species responsible for
observed dissolution rates, starting from early work (12–14) and
culminating in a dissolution model that incorporates both aque-
ous species and the distribution of ion complexes on the calcite
surface (15, 16). These groups recover an essentially linear rela-
tionship between dissolution rate and saturation state.
Calcite dissolution rates in freshwater under circumneutral pH
conditions appear largely independent of solution pH (10), but
instead respond nonlinearly to mineral surface processes (17–
20). For example, faster dissolution rates are typically associ-
ated with a large increase in etch pit nucleation below some
critical saturation state,
critical
(21, 22). Even in freshwater,
however, a link between the generation of etch pits and other
surface features, and the overlying solution chemistry, is poorly
understood (20).
In this paper, we use a method that unpacks the relation-
ship between calcite dissolution kinetics and seawater chemistry
(
Methods
and ref. 9). In short,
13
C
-labeled calcites are placed in
undersaturated seawater in a closed system with no headspace.
The evolving seawater
δ
13
C
traces mass transfer from solid to
solution. Dissolution will add
13
C
to solution; precipitation will
add seawater carbon (
99%
12
C
) to the solid surface. We first
demonstrate precipitation in undersaturated solutions using Sec-
ondary Ion Mass Spectrometry (SIMS) analysis of the calcite
solid. Next, we show that a box model of calcite dissolution and
precipitation provides information about gross precipitation and
dissolution fluxes at the calcite surface near equilibrium. The
model predicts a decrease in relative importance of precipita-
tion below
Ω = 0
.
7
. Finally, we show that the enzyme carbonic
anhydrase (CA) catalyzes the dissolution of calcite in seawater.
CA is a cosmopolitan enzyme known for its rapid equilibration
of carbonic acid and aqueous
CO
2
. A mechanism of dissolu-
tion through defect-assisted etch pit nucleation (21) is proposed,
which also demonstrates a distinct change in reaction energetics
at
Ω = 0
.
7
. Near equilibrium, CA reduces the free energy barrier
to dissolution. At high [CA], this energy barrier is similar to that
recovered from freshwater dissolution experiments.
Significance
The experimental system described here provides constraints
on the relative balance of gross dissolution and precipitation
fluxes contributing to the observed net dissolution rate of cal-
cite in seawater. We show that our dissolution rates fit well
within a framework that accounts for the geometry of the dis-
solving mineral surface. We further show that carbonic anhy-
drase (CA) catalyzes calcite dissolution, which implicates the
hydration of aqueous CO
2
as a rate-limiting step for calcite
dissolution in seawater. The presence of carbonic anhydrase
in carbonate-rich environments such as coral reefs or sinking
marine particles is poorly understood. However, our findings
suggest that CA activity would significantly enhance the rate
at which alkalinity is cycled between solids and seawater in
these environments.
Author contributions: A.V.S., J.F.A., N.E.R., J.E., and W.M.B. designed research; A.V.S.,
N.E.R., and J.N. performed research; A.V.S., J.F.A., N.E.R., J.N., and W.M.B. analyzed data;
and A.V.S., J.F.A., J.N., and W.M.B. wrote the paper.
The authors declare no conflict of interest.
This article is a PNAS Direct Submission.
1
To whom correspondence should be addressed. Email: asubhas@gps.caltech.edu.
This article contains supporting information online at
www.pnas.org/lookup/suppl/doi:10.
1073/pnas.1703604114/-/DCSupplemental
.
www.pnas.org/cgi/doi/10.1073/pnas.1703604114
PNAS
|
August 1, 2017
|
vol. 114
|
no. 31
|
8175–8180
Measuring Dissolution
Precipitation in the Calcite Solid
The net rate of
CaCO
3
dissolution is the result of a balance
between dissolution and precipitation reactions near equilib-
rium. Our experimental system has a strong isotopic gradient
between seawater and the
13
C
-labeled mineral, such that the
processes of precipitation and dissolution will each leave distinct
isotopic signatures on the solid. Seawater
12
C
will precipitate and
increase the
12
C
/
13
C
ratio of the reactive calcite surface. Disso-
lution will expose pure
13
C
carbonate such that the surface com-
position approaches 100%
13
C
as dissolution outpaces precipi-
tation at lower saturation state. A change in Mg/Ca should also
be measurable because the initial calcite mineral contains only
trace
Mg
2+
.
We measured the near-equilibrium balance of dissolution and
precipitation by reacting several
13
C
-labeled, Mg-free calcites for
48 h in saturated and slightly undersaturated seawater. Using
SIMS, we then measured carbon isotope and Mg/Ca profiles
through this reacted surface and compared them to an unreacted
control (Fig. 1 and
SI Appendix
, Fig. S1
). The unreacted con-
trol experiments (dashed yellow) show some enrichment of
12
C
due to surface contamination. The supersaturated experiments
(dotted blue), which will have precipitated calcite, show persis-
tent
12
C
and Mg/Ca enrichment above the unreacted control. The
undersaturated experiments (solid red) start at a surface compo-
sition similar to precipitated calcite, indicating that
12
C
has been
incorporated even in undersaturated conditions. Undersaturated
profiles then transition down to the composition of the unreacted
control, matching the unreacted pure
13
C
calcite composition at
a depth of about 4 nm. Fig. 1,
Inset
shows that all
12
C
/
13
C
com-
positions converge within error at about 130 nm deep, indicating
that underlying
13
C
calcium carbonate has been reached.
The shape of these SIMS profiles is influenced by the incor-
poration of
12
C
from precipitation, and also by the mixing of
surface signals down into the calcite interior during secondary
ion sputtering and excavation. To more quantitatively measure
the total number of moles of seawater C and
Mg
2+
incorporated
into our calcites, we integrated each of the curves in Fig. 1 and
SI Appendix
, Fig. S1
over the 140-nm SIMS profile (
SI Appendix
,
Eqs.
S1
and
S2
and Table S2
). This integral can be converted
to a “reactive thickness” and compared with previous estimates.
Briefly, the measured
12
C
/
13
C
mole fraction measured at each
SIMS analysis cycle was converted to a number of moles of
12
C
.
All cycles were then summed over the entire profile to estimate
the total number of moles of
12
C
incorporated into our calcites
over the 48-h experimental period. To remove the influence of
surface contamination, the control was subtracted from both the
supersaturated and undersaturated experiments. The total num-
ber of moles was then converted into an effective thickness of
12
C
present in the supersaturated and undersaturated experi-
ments (
SI Appendix
, Table S2
). This calculation estimates that
1.0 nm to 3.3 nm of
12
C
calcite was incorporated in undersat-
urated conditions. Mg/Ca measurements also confirm the pres-
ence of new
Mg
2+
in the solid (
SI Appendix
, Table S2
). We also
found that 3.1 nm to 9.3 nm of
12
C
calcite was added in super-
saturated experiments, in agreement with previous precipitation
experiments (ref. 23 and
SI Appendix
, Fig. S3
). One calcite mono-
layer is about 0.5 nm, suggesting that, even under conditions of
net dissolution, seawater can react with two to six monolayers of
calcite via gross dissolution and precipitation reactions.
Most calcite dissolution in the deep sea happens near equilib-
rium, where we have documented a large influence of dissolu-
tion and precipitation reactions on the composition of the calcite
surface. Calcites in sediments, which contain primary environ-
mental information in their oxygen isotope and/or Mg/Ca ratios,
will experience similar dissolution–precipitation reactions in the
deep ocean (24–26). These reactions will impart secondary pore-
water isotopic and chemical information, and our experiments
suggest they can potentially do so deeply into the calcite surface.
Fig. 1.
Vertical logarithmic
12
C
/
13
C SIMS profiles of reacted calcite grains
under three different experimental conditions. Solid lines are the mean
isotope ratio of all profiles collected under each experimental condition.
Shaded areas are the SD of all profiles collected under each experimental
condition. In the first 15 nm of these profiles,
= 0.95 profiles (solid red)
transition from a supersaturated (dotted blue) composition to an unreacted
(dashed yellow) composition. This
12
C enrichment demonstrates that sea-
water carbon has incorporated into the calcite solid in undersaturated con-
ditions. (
Inset
) SIMS entire profile, with experiment ratios converging with
the unreacted control run at depth. All curves are depth-corrected for the
thickness of gold coating. Note that the
x
axis is flipped relative to the con-
vention of plotting a
13
C
/
12
C ratio.
Future work will target understanding the extent to which these
processes modify deep-sea sedimentary calcites, and should
prove useful in unmixing the secondary and primary environmen-
tal signals that these calcites record.
Measuring and Modeling Dissolution
Precipitation in the
Solution
In addition to measuring dissolution–precipitation reactions in
the solid, we measured the appearance of
13
C
in seawater dis-
solved inorganic carbon (DIC) using a Picarro cavity ring-down
spectrometer (9). The dissolution of
13
C
calcite produces a lin-
ear increase of seawater
δ
13
C
over time. The slope of this line is
a direct measure of the net dissolution rate (e.g., rates in ref. 9).
However, we also observe curvature in plots of
δ
13
C
versus time
(Fig. 2). Because an experiment is conducted at a fixed saturation
state and mineral surface area, curvature in Fig. 2
B
cannot rep-
resent a change in bulk solution chemistry, but instead reflects
the instantaneous balance of fixed dissolution and precipitation
rates at the mineral surface. Precipitation adds
12
C
to the solid
surface, which changes the isotopic composition of the dissolv-
ing solid. The calcite surface will continue to change its
13
C
/
12
C
ratio until it comes into steady state with respect to gross dis-
solution and precipitation reactions. Curves of
δ
13
C
versus time
thus straighten out when the calcite surface comes into steady
8176
|
www.pnas.org/cgi/doi/10.1073/pnas.1703604114
Subhas et al.
EARTH, ATMOSPHERIC,
AND PLANETARY SCIENCES
Decreasing
precipitation
rate
A
B
Experiment Data
Best Model Fit
Top-Binned Model Fits
r
fb
= 5.0
3.4
2.2
1.5
1.0
13
C (‰)
0
5
10
15
01234
0.5
1.5
2.5
3.5
Time (days)
Time (days)
0
0.5
1
1.5
2
2.5
3
3.5
10
-8
mol dissolved
0
1
2
3
4
5
6
7
Fig. 2.
Results of the dissolution–precipitation model (
SI Appendix
, Fig. S3
and Eqs.
S2
S4
). (
A
) Model output of bulk solution
δ
13
C under our experi-
mental conditions, assuming a reactive calcite layer thickness of five mono-
layers. Decreasing the precipitation rate (increasing
r
bf
; see
Measuring and
Modeling Dissolution–Precipitation in the Solution
for details) increases the
net dissolution rate and decreases curvature. (
B
) Model–data comparison for
a dissolution experiment conducted at
Ω =
0
.
87. The red curve is the abso-
lute best-fit
k
diss
and
r
fb
over the entire parameter search space; the gray
curves are the next 15 best fits.
state with respect to these dissolution and precipitation fluxes
(Fig. 2
A
). Using the observation of precipitated calcite even in
undersaturated conditions, we developed a box model to quantify
the impact of dissolution and precipitation reactions on solution
δ
13
C
versus time (see
SI Appendix
, Fig. S4
for model description).
All experimental data with sufficient data density from ref. 9
were fit with output from the box model. We modeled a suite
of dissolution rate curves over a range of gross dissolution (
R
f
)
and precipitation rates (
R
b
). The magnitude of these rates, and
their ratio (
r
fb
=
R
f
/
R
b
), sets the amount of initial curvature.
For example, when
r
fb
= 1, the precipitation rate is equal to the
dissolution rate, and calcite is in equilibrium with solution (the
blue line in Fig. 2
A
). When
r
fb
= 5, the dissolution rate is 5 times
faster than the precipitation rate (the green line in Fig. 2
A
). The
size and shape of curvature observed in Fig. 2 is also set by the
volume of calcite that is allowed to react with seawater. Consis-
tent with our SIMS results, the model gives a better fit to the
data if multiple monolayers of calcite are reacting with seawa-
ter, and gives a best fit with five monolayers (
SI Appendix
, Fig.
S7
). An example data fit using model output is shown in Fig.
2
B
and
SI Appendix
, Fig. S6
; details of the fitting routine are in
SI Appendix
. Ranges of acceptable dissolution and precipitation
rates were calculated from the acceptable model-fit
R
f
and
r
fb
values, shown as the spread of gray curves that adequately fit the
dataset in Fig. 2
B
.
Dissolution studies in seawater have typically related net dis-
solution rate with undersaturation [1 –
(4, 6–8, 27, 28)]. This
framework is historically linked to a derivation of net dissolution
rate,
R
diss
=
R
f
R
b
=
k
f
{
CaCO
3
}−
k
b
[Ca
2+
][CO
2
3
]
,
[1]
where
R
is a rate and
k
f
and
k
b
are the specific dissolution and
precipitation rate constants, respectively. This expression, along
with the assumption that activity of the solid is 1, defines a linear
relationship between net dissolution and
1
(see
SI Appendix
for derivation). Furthermore, the precipitation term in Eq.
1
is
the only term that explicitly includes information about the con-
centrations of calcium and carbonate ion in solution. In other
words, gross dissolution is independent of solution composition,
and saturation state only affects the rate of gross precipitation.
Such dissolution behavior has been demonstrated in freshwater
(e.g., ref. 16), but has never been investigated directly in sea-
water. Our modeling results here are the complement to this
freshwater approach. Model-fit gross dissolution and precipita-
tion rates, shown in Fig. 3, cannot be fit using a constant disso-
lution rate and a decreasing precipitation rate, as predicted by
the linear model of Eq.
1
. Instead, gross dissolution rates show
a strong, nonlinear dependence on saturation state, changing by
almost four orders of magnitude over the entire range of satura-
tion states measured here. This nonlinearity implies that either
the dissolution rate constant or the activity of the solid in Eq.
1
changes as a function of saturation state.
As undersaturation increases, the range of acceptable precipi-
tation rates grows much larger than the range of acceptable dis-
solution rates, as shown by the size of the boxes in Fig. 3
A
. The
ratio of gross dissolution to precipitation (
r
fb
=
R
f
/
R
b
) also sig-
nificantly increases below about
Ω = 0
.
7
, suggesting a decreased
contribution of precipitation to the net dissolution rate farther
from equilibrium (Fig. 3
B
). The linear dissolution framework
(Eq.
1
) predicts a continuous decline in precipitation as a func-
tion of
1
. Our model results instead show an abrupt change
in the balance of precipitation to dissolution at
Ω = 0
.
7
. Once
r
fb
becomes larger than 10, precipitation accounts for only 10% or
less of the net dissolution rate. Changes in
r
fb
therefore have a
relatively small effect on the goodness of fit to the experimental
data, and model fits with high
r
fb
accommodate a larger accept-
able range in
r
fb
. Such an abrupt shift in the balance of precipita-
tion to dissolution helps to explain the strongly nonlinear disso-
lution rates observed in many seawater calcite dissolution studies
(e.g., refs. 6 and 7). It also implies a distinct change in dissolution
mechanism.
Catalysis via Carbonic Anhydrase and a Link Between
Solution Chemistry and Surface Features
The nonlinearity of a gross dissolution flux brings into question
the chemical reactions responsible for calcite dissolution in sea-
water. In dilute solution, three calcite dissolution mechanisms
are proposed to operate at the calcite surface (12, 16). Water-
catalyzed dissolution is predicted to be independent of solution
chemistry (the activity of the solid is always assumed to be 1).
At low pH (
<
5), calcite dissolution depends almost exclusively
on the transport of hydrogen ion to the mineral surface (14,
16, 29).
CO
2
itself is thought to be relatively unreactive with
CaCO
3
, but
H
2
CO
3
, produced through the hydration of
CO
2
,
is an uncharged species that acts as a proton donor to promote
dissolution (16, 29, 30).
There is increasing evidence that CA, an enzyme that catalyzes
the equilibration of
CO
2
and
H
2
CO
3
, can enhance the dissolu-
tion rate of calcium carbonate in karst systems and their analogs
Subhas et al.
PNAS
|
August 1, 2017
|
vol. 114
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no. 31
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8177
1
10
100
0
0.2
0.4
0.6
0.8
1.0
r
fb
= R
f
/R
b
1-
Ω
Gross precipitation (
R
b
)
tGross dissolution (
R
f
)
0
0.2
0.4
0.6
0.8
1.0
1-
Ω
-9
-10
-11
-12
-13
-14
-15
Log Rate (mol s
-1
)
A
B
Fig. 3.
Results of the dissolution–precipitation model. (
A
) Each net disso-
lution rate is represented by a pair of blue (dissolution) and yellow (precip-
itation) gross rates. Lines in the boxes are the median of the best fits of
R
f
;
box boundaries are the 25th and 75th percentile values for
R
f
that best fit
the experimental moles dissolved versus time data. Gross precipitation rates
R
b
are the median
R
f
divided by the median
r
fb
(
=
R
f
/
R
b
) of the best fits
to the experimental data; box boundaries are the 25th and 75th percentile
values for
R
b
. Overall, dissolution and precipitation rates are very close to
each other, leading to a net dissolution rate that is the difference between
two large gross fluxes. The precipitation rate variance increases (larger yel-
low box size) after 1
0
.
3. (
B
) Box plot of the best-fitting
r
fb
values
for the all experiments without CA. A significant jump in
r
fb
is evident after
1
0
.
3 in both the absolute value of
r
fb
and the range of acceptable
values.
(31–35). These studies used an open-system dissolution reactor,
in which
CO
2
gas was bubbled into an experimental chamber.
There are two effects of CA in such a system. The first is rapid
equilibration of the bubbled gas and the solution
p
CO
2
, resulting
in an experimental system limited by the kinetics of gas exchange
and with a poorly constrained saturation state, as documented
by ref. 16. Second, CA could enhance chemical mechanisms act-
ing at the calcite–water interface. Our results are from closed-
system dissolution experiments that have no headspace and thus
isolate this second chemical mechanism of CA on calcite disso-
lution (ref. 9 and
Methods
).
Mass- and surface area-normalized dissolution rates from many
13
C
dissolution experiments are plotted as a function of under-
saturation (1 –
) and CA concentration in Fig. 4. Calcite dis-
solution rates in the presence of CA are always enhanced over
the uncatalyzed rates. An experiment performed in the presence
of BSA showed no significant enhancement of dissolution rate
(
SI Appendix
, Fig. S9
), demonstrating that proteinaceous dis-
solved organic matter has no significant effect on dissolution rate.
Increasing [CA] enhances dissolution at all saturation states,
and CA has the largest effect close to equilibrium. Far-from-
equilibrium enhancement of dissolution rate in Fig. 4,
Inset
shows similar rate enhancements to values obtained in fresh-
water experiments (31). This result is surprising, because many
studies in freshwater have either dismissed carbonic acid as a
major proton donor (7, 29) or have not found catalysis via CA in
freshwater at low
p
CO
2
(31). Our documented rate increase of
2.5 orders of magnitude at [CA] = 0.04 mg/mL and
0
.
85
suggests that, in contrast to freshwater, carbonic acid is a major
proton donor close to equilibrium in seawater.
The strong nonlinearities in our dissolution rate data cannot
be explained using the simple dissolution framework of Eq.
1
.
Furthermore, the addition of CA, although it increases the dis-
solution rate of calcite, does not produce a linear response of
dissolution rate to saturation state. We were thus compelled to
apply a model of dissolution to our rate data that incorporates
features of the calcite solid into the control of dissolution rates
(17, 21, 22, 36). For many solid–solution systems, a description
of solution chemistry alone is indeed insufficient to predict crys-
tal growth or dissolution kinetics. In addition to saturation state
and its contribution to free energy, the crystal growth theory of
ref. 37 incorporates energetics associated with the crystal itself,
such as the free energy of crystal edges, faces, and lattice defects
in contact with the solution. Recently, several studies have suc-
cessfully mapped this theory of crystal growth and precipitation
onto the dissolution of quartz, feldspar, diopside, and calcite (21,
36, 38). This model relates net dissolution rate (
R
diss
) to several
physical–chemical parameters (
h
,
ω
,
C
e
, and
a
; see
SI Appendix
),
as well as the dissolution velocity
β
at defects, surface defect
density
n
s
, and the interfacial energy barrier at nucleation sites
α
(21, 39),
ln
[
|
R
diss
|
(1
Ω)
2
/
3
|
σ
|
1
/
6
]
= ln[
h
β
C
e
(
ω
2
hn
s
a
)
1
/
3
]
πα
2
ω
h
3(
k
B
T
)
2
1
σ
.
[2]
Saturation state control is found in the
and
σ
=
ln
(Ω)
terms;
k
B
T
is the Boltzmann constant multiplied by temperature in
kelvins, i.e., the system’s thermal energy.
Freshwater (Cubillas et al., 2005)
1-
Ω
Log
R
d
(g cm
-2
day
-1
)
0.1
0.2
0.3
0.4
0.5
0.6
0.7
0.8
0.9
1.0
-2
-3
-4
-5
-6
-7
-8
Seawater Uncatalyzed
0.01 mg/mL CA
0.02
0.04
1
3
5
7
0.6
0.7
0.8
0.9
1
x10
-3
g cm
-2
day
-1
Fig. 4.
The relationship between saturation state, carbonic anhydrase con-
centration, and calcite dissolution rate in seawater. Semilog plot of disso-
lution rate versus undersaturation (1 –
), including freshwater data from
ref. 10. The linear–linear
Inset
at bottom right shows the far-from-equilibrium
dissolution rate increase as a function of carbonic anhydrase. The
x
axis
(1 –
) is the same as in the main figure; the
y
axis (dissolution rate) is in units
of 10
3
g
·
cm
2
·
d
1
. For clarity,
Inset
does not show freshwater data.
8178
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www.pnas.org/cgi/doi/10.1073/pnas.1703604114
Subhas et al.
EARTH, ATMOSPHERIC,
AND PLANETARY SCIENCES
10
-22
-21
-20
-19
-18
-17
-16
-15
-14
-13
-12
-11
ln
|
R
diss
|
(1-
Ω
)
2/3
|
σ
|
1/6
2 4
6
8
0
|1/
σ
|
Freshwater
Seawater Uncatalyzed
0.01 mg/mL CA
0.02 mg/mL CA
0.04 mg/mL CA
Ω
calcite
0.9
0.8
0.7
0.5
0.1
Homogeneous
2D-Dissolution
Defect-Driven
Dissolution
Fig. 5.
Dissolution rate data from Fig. 4 plotted in the framework of Eq.
2
. The
y
axis dissolution rate (
R
diss
) is in moles per square meter per second.
A top axis of corresponding
values is included, and regimes of dissolu-
tion with their hypothesized mechanisms are shown. To the right of the
kink (closer to equilibrium), note the decreasing slope with increasing [CA].
Freshwater data are included for comparison. Linear fits to the data in this
framework are presented in Table 1, along with an estimate of the interfa-
cial surface energy
α
. The kink in these data represents a change in disso-
lution mechanism from defect-only nucleated dissolution near equilibrium
to homogeneous nucleation far from equilibrium at a kink point around
Ω =
0
.
7.
Catalyzed and uncatalyzed dissolution rate data are plotted as
the left-hand side of Eq.
2
versus
|
1
|
in Fig. 5, where
|
1
|
= 0
is complete undersaturation. The slope of a straight line in this
space gives the interfacial energy barrier
α
; the intercept gives
information about the kinetic rate constant
β
and density of etch
pit nucleation sites
n
s
. As seen in several other mineral dis-
solution studies (21, 22, 40), our uncatalyzed data plot as two
straight lines in this space with a “kink” at
|
1
|
= 3
, which cor-
responds to
= 0.71. This transition in slope denotes two dis-
tinct regimes of dissolution, which has been interpreted previ-
ously as a transition from defect-assisted nucleation of etch pits
near equilibrium to homogeneous etch pit nucleation farther
from equilibrium (21). It could also be interpreted as a transi-
tion from the opening of hollow cores to the propagation of step
waves in the framework of Lasaga and Luttge (36). This transi-
tion to homogeneous nucleation of etch pits is also concurrent
with the saturation state at which precipitation becomes unim-
portant to the overall dissolution rate in our
δ
13
C
tracer data
(Fig. 3).
Table 1. Fits for dissolution rate data shown in Fig. 5 in the framework of Eq. 2
Near equilibrium
Far from equilibrium
Experiment
Intercept
Slope
α
, mJ
·
m
2
Intercept
Slope
α
, mJ
·
m
2
SW Uncat.
17.2
±
0.5
0.69
±
0.11
15
±
1.5
12.3
±
0.1
2.4
±
0.1
29
±
1
0.01 mg/mL CA
16.0
±
0.1
0.42
±
0.02
12
±
0.5
12.1
±
0.2
1.9
±
0.1
26
±
2
0.02 mg/mL CA
15.1
±
0.2
0.35
±
0.05
11
±
1.5
11.5
±
0.1
2.6
±
0.2
30
±
2
0.04 mg/mL CA
15.5
0.03
3
11.3
±
0.1
2.5
±
0.1
29
±
2
Freshwater (10)
13.7
±
0.3
0.05
±
0.15
4
±
13
Near-equilibrium fit data from 1
/σ >
3; far-from-equilibrium fits are from 0
<
1
/σ <
3. There are no errors
on the 0.04 mg/mL data near equilibrium because the fit was made with two points. Freshwater data were
taken from ref. 10. Only a single fit was performed on freshwater data, due to the lack of kink.
Near equilibrium (to the right in Fig. 5), slopes decrease with
increasing [CA] (Table 1). The slope of this line is diagnostic of
the rate-limiting step in calcite dissolution near equilibrium. A
decrease in the free energy barrier as a function of [CA] suggests
that a greater availability of carbonic acid effectively decreases
the energetic barrier to etch pit nucleation, by increasing the con-
centration of carbonic acid at defects on the calcite surface. The
addition of CA also increases the intercept (Table 1), changing
either the density of nucleation sites (
n
s
) or the rate of step retreat
(
β
; Eq.
2
). Adding CA does not seem to significantly change the
transition between defect-assisted and homogeneous nucleation:
The kink point in Fig. 5 does not move significantly given the den-
sity of our data. The transition between these two regimes may
be controlled instead by calcite saturation state (i.e., [
CO
2
3
]).
Because the kinetic interconversion of
CO
2
and
H
2
CO
3
does
not change the thermodynamic saturation state
, one might not
expect this kink point to change significantly due to enhanced
hydration kinetics. Far from equilibrium (to the left in Fig. 5),
slopes are insensitive to [CA], indicating that interface energies
at etch pit nucleation sites are insensitive to the concentration of
H
2
CO
3
. Instead, transport of
H
2
CO
3
simply limits the delivery
of protons to the mineral surface, limiting the overall propagation
of etch pits once formed (i.e., modulating
β
or
n
s
in Eq.
2
).
A transition to homogeneous etch pit nucleation is also con-
sistent with the mechanism shift diagnosed using relative disso-
lution and precipitation fluxes above. It is possible that precipi-
tation limits net dissolution near equilibrium by occupying sites
that would otherwise dissolve. This constraint is freed when etch
pits begin to form everywhere on the solid surface. Instead of
being limited to defects, dissolution is now allowed to proceed
everywhere on the mineral surface, which erases the influence of
precipitation reactions on the isotopic composition of the solid,
and on the time-evolving solution
δ
13
C
.
Conclusion
In general, calcites react much more slowly and with less pre-
dictability in seawater than in freshwater, a problem that has
plagued marine chemists for decades. Furthermore, calcite dis-
solution repeatedly has been shown to respond nonlinearly to
saturation state, implying the presence of multiple dissolution
mechanisms. We show here that this strongly nonlinear disso-
lution behavior in seawater is due to the combined effects of
solution chemistry and geometric constraints on the propagation
of dissolution features on the solid surface. Gross precipitation
and dissolution fluxes influence the incorporation of both cation
and anion tracers into the calcite solid. Treatment of bulk rate
data using our box model demonstrates a change in the balance
of dissolution and precipitation at
Ω = 0
.
7
. Despite the chemical
complexities that arise in seawater, our results suggest that car-
bonic acid availability is key, such that increasing its formation
using CA drastically increases calcite dissolution near equilib-
rium. This reaction pathway appears to be rate-limiting in sea-
water near equilibrium. Incorporation of geometric constraints
Subhas et al.
PNAS
|
August 1, 2017
|
vol. 114
|
no. 31
|
8179
on dissolution rates suggests that the balance of precipitation and
dissolution reactions changes fundamentally once etch pits freely
nucleate and propagate across the calcite surface. These findings
have implications not only for the reactivity differences of calcite
between freshwater and seawater but also for how calcite dis-
solves in natural environments in the presence of increased
CO
2
and carbonic anhydrase.
Methods
For dissolution rate measurements, we followed the methods detailed in
ref. 9. More details can be found in
SI Appendix
. Briefly,
13
C-labeled car-
bonates were placed in undersaturated seawater in gas-impermeable bags
with no headspace. The seawater was sampled over time for its DIC and
δ
13
C, measured on a modified Picarro cavity ringdown spectrometer (CRDS)
with Liason autosampler. The
δ
13
C signals were then converted into the
number of moles dissolved versus time. Saturation state was measured
by measuring DIC and total alkalinity on seawater before, during, and
after every experiment. Solutions of lyophilized carbonic anhydrase from
bovine erythrocites purchased from Sigma Aldrich (C2624) were made up
in either deionized water or seawater. These solutions were added to
the dissolution experiments to bring the total [CA] in each experiment to
0, 0.01, 0.02, or 0.04 mg/mL. Final experimental alkalinities were always
checked; in several experiments, the alkalinity of initial experimental sea-
water was checked after CA addition to confirm the magnitude of alkalin-
ity change due to CA. Dissolution rates calculated in this way were plot-
ted against calculated saturation state and carbonic anhydrase concen-
tration. Details on the box model and SIMS analysis can be found in
SI
Appendix
.
ACKNOWLEDGMENTS.
We acknowledge Alex Gagnon for helpful discus-
sions on the formulation of the box model, and Yunbin Guan for help
with SIMS analysis. We also thank Sijia Dong for discussions in general
about carbonate dissolution in seawater. We thank Mathis Hain and one
anonymous reviewer, whose careful reading and detailed comments greatly
improved this manuscript. Thanks go to National Science Foundation Grad-
uate Research Fellowship Program and the Resnick Institute Graduate Fel-
lowships for supporting A.V.S. and J.N. We acknowledge support from NSF
Grants OCE1220600 and OCE1220302.
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