Introduction
Vaterite is the least known polymorph of calcium carbonate and was first
described by Vater (Vater, 1893). In the 1920s and 1930s, the nature and
crystallographic structure of vaterite was still questioned and the
occurrence of a third polymorph of CaCO3 was disputed amongst various
groups (Johnston et al., 1916; Spangenberg, 1921; Heide, 1924; Rinne, 1924;
Gibson et al., 1925). In the following decades, work mainly focused on the
structure of vaterite (e.g. McConnell, 1960; Kamhi, 1963; Meyer, 1969; Mann
et al., 1991; Wang and Becker, 2009), which continues to be a source of surprises
today (Kabalah-Amitai et al., 2013). Vaterite has a major hexagonal
structure (von Olshausen, 1925) and appears in different morphologies such
as spherulitic aggregates (Han et al., 2006; Nebel and Epple, 2008; Mori et
al., 2009; Hu et al., 2012) or hexagonal plates (e.g. Johnston et al.,
1916; Kamhi, 1963; Dupont et al., 1997; Xu et al., 2006; Kawano et al.,
2009).
Vaterite occurrence in nature is more widespread than generally assumed. It
was first observed in gastropods (Mayer and Weineck, 1932), but later
studies also discovered vaterite to be related to oil field drilling (Friedman and
Schultz, 1994) and to be found in Portland cement (Friedman and Schultz, 1994) and as
stones in the urinary system (Prien and Frondel, 1947; Sutor and Wooley,
1968). Vaterite has not been found in the geologic record and is therefore
suspected to be metastable. The observation of vaterite in biogenic systems
(Mayer and Weineck, 1932; Spann et al., 2010; Nehrke et al., 2012;
Kabalah-Amitai et al., 2013) gives some constraints on its stability, which
can be on the order of years (Lowenstamm and Abbott, 1975), but not
geological ages.
The natural occurrence of vaterite and its potential economic use due to its
large specific surfaces and high porosity (Mori et al., 2009) warrants a
precise investigation of this mineral. So far, most laboratory experiments
have been designed to precipitate relatively large single crystals of vaterite for
X-ray analysis that focused on the crystal structure (e.g. Kamhi, 1963).
Vaterite precipitation experiments generally used either mixtures of several
solutions such as K2CO3, Na2CO3 and CaCl2 (Kamhi,
1963; Easton and Claugher, 1986; Han et al., 2006; Nebel and Epple, 2008) or
CaNO3 (Davies et al., 1978), sometimes with added surfactants (Mann et
al., 1991; Dupont et al., 1997; Mori et al., 2009), or additional organic
substrates (Falini et al., 1996; Xu et al., 2006; Kirboga and Oner, 2013).
These experiments provided crystals with sizes between a few hundred nanometres and a
few micrometres and were mostly restricted to the temperature range of
25–60 ∘C.
As vaterite is a relevant mineral in biogenic systems (Lowenstamm and Abbott,
1975; Pouget et al., 2009; Spann et al., 2010; Nehrke et al., 2012) it may
provide new insights into isotope fractionation during biological carbonate
formation. Vaterite has been proposed as a potential precursor phase in
biogenic CaCO3 precipitation, later transforming into aragonite or
calcite (e.g. Jacob et al., 2008). Depending on conditions prevailing during
the vaterite–aragonite or vaterite–calcite transformation, the original
isotope signature may be inherited in the final, stable CaCO3 phase. In
order to be able to use isotopic tools (for instance, to reconstruct
environmental conditions or palaeoclimate) on biogenic carbonate with
vaterite precursors or remaining vateritic components, knowledge of the
fractionation factor between vaterite and water and the clumped isotope
Δ47–T relationship of vaterite is required. However, so far
little is known about the oxygen isotope fractionation between dissolved
inorganic carbon and vaterite and, in particular, the clumped isotope
Δ47–T relationship of vaterite. Whereas vaterite was reported in
a few stable isotope studies aiming at determining the oxygen isotope
fractionation factor in the system CaCO3–H2O (e.g. Kim and O'Neil,
1997), it rarely occurred as a pure phase and thus did not allow for a
precise study focused on vaterite. Based on the limited data available,
Tarutani et al. (1969) suggested vaterite to be enriched in 18O by
+0.5 ‰ relative to calcite. Kim and O'Neil (1997) obtained a
similar value of +0.6 ‰ with respect to calcite. Both studies were
limited to either one (25 ∘C) or two temperatures (25,
40 ∘C), and a more comprehensive study is still lacking. In
addition, the clumped isotope Δ47–T relationship of vaterite has
not been assessed so far, but could give new insights into the effect of
polymorphism on isotope ratios or mineral growth-related isotope
fractionation.
In this study we present a simple method that allows vaterite precipitation
over a wide temperature range (at least between 20 and 91 ∘C) and
that provides large quantities of the mineral, enabling for example the
investigation of the oxygen isotope fractionation factor between vaterite and
H2O, and the Δ47–T relationship of vaterite.
Sketch (a) and photograph (b) of the experimental
setup used to precipitate vaterite. In the first step, CaCO3 is dissolved
in de-ionized (DI) water (outside the water bath), which is filtered after
>3 h to remove any undissolved components. The solution is then
transferred to a temperature-controlled water bath for thermal and isotopic
equilibration (flask 2 in a). NaCl is added after the equilibration
step. Mineral formation is induced by slow bubbling of N2 through the
solution. The gas stream through the solution in flask 2 is humidified by
passing it beforehand through another flask filled with de-ionized water
(flask 1).
Experimental conditions during laboratory precipitation of
CaCO3 (see Sect. 2).
Experiment
T
CaCO3,dissolved
NaCl added
Equilibration
Precipitation
no.
(∘C)
(gL-1)
(gL-1)
(h)
(h)
NA-1
23.5 ± 0.5
0.68
250
23
451
NA-3*
37.5 ± 0.5
0.70
260
21
72
NA-4
37.5 ± 0.5
0.74
244
14
341
NA-5
49.6 ± 0.5
0.70
375
16
143
NA-6
49.6 ± 0.5
0.80
262
17
573
NA-7
69.9 ± 0.5
0.70
325
3
69
NA-8
79.9 ± 0.5
0.78
280
3
47
NA-9
91.0 ± 0.5
0.70
260
3
42
* NA-2 differed in the experimental conditions and is
therefore omitted.
Experimental setup
The precipitating solution was prepared by dissolving high-purity CaCO3
(Merck Suprapur, 99.95 %) in de-ionized water. Light microscopy indicates
that this CaCO3 consists of calcite. The water was initially taken from
the local water supply, purified with a reverse osmosis technique and finally
de-ionized with an Ultrapure™ system with an
output quality of 18 MΩ cm. Trace components of the CaCO3 used
to prepare the solution are strontium (≤ 100 ppm), rubidium (≤ 20 ppm), sulfate (≤ 20 ppm), and phosphate (≤ 10 ppm).
About 370 mg of CaCO3 was dissolved in ∼ 500 mL de-ionized water at
room temperature per experiment. The water was acidified by the purging of
CO2 tank gas (normal grade, BOC UK) through the solution. The solution
was filtered after 3 h through a double layer of
Whatman® filter paper (grade 1, corresponding
to 11 µm filtration size) to remove undissolved CaCO3
crystals. Optical inspection of the filtered water via light reflectance
confirmed no large crystals to have bypassed the filtration stage. The
dissolution step appeared to be quantitative as no CaCO3 was visibly
retained on the filter paper.
Mineralogical and isotopic results of the vaterite precipitation
experiments. The mineralogy was determined by XRD analysis (Sect. 3.2). The n
values give the number of replicates measured for isotopic analysis. The
Δ47 value is given in the absolute reference frame of Dennis
et al. (2011) and is corrected for the acid digestion reaction.
Experiment
T
CaCO3,precipitated
δ18O
δ13C
Δ47
n
Mineralogy
no.
(∘C)
(mg)
(‰)
(‰)
(‰)
(–)
NA-1
23.5 ± 0.5
120
-8.57 ± 0.16
-18.21 ± 0.06
0.689 ± 0.003
3
vaterite
NA-3*
37.5 ± 0.5
5
-11.29 ± 0.20
-20.39 ± 0.10
0.639 ± 0.020
1
vaterite (95 %),calcite (5 %)
NA-4
37.5 ± 0.5
50
-13.30 ± 0.37
-26.06 ± 0.18
0.672 ± 0.027
3
vaterite (>95 %),rest: calcite
NA-5
49.6 ± 0.5
15
-13.85 ± 0.26
-21.39 ± 0.03
0.605 ± 0.005
2
vaterite
NA-6
49.6 ± 0.5
235
-15.06 ± 0.22
-25.26 ± 0.17
0.634 ± 0.008
3
vaterite
NA-7
69.9 ± 0.5
15
-16.92 ± 0.15
-21.71 ± 0.03
0.577 ± 0.010
3
vaterite
NA-8
79.9 ± 0.5
80
-17.54 ± 0.03
-25.86 ± 0.10
0.553 ± 0.018
3
calcite (49 %),aragonite (24 %),vaterite (27 %)
NA-9
91.0 ± 0.5
90
-19.21 ± 0.15
-25.00 ± 0.16
0.545 ± 0.005
5
vaterite (94 %),aragonite (6 %),calcite (<1 %)
* NA-2 differed in the experimental conditions and is
therefore omitted.
The filtered CaCO3 solution was then thermally and isotopically
equilibrated at a set temperature in a temperature-controlled water bath
(a sketch of the experimental setup is depicted in Fig. 1). The experimental
temperatures ranged from 23 to 91 ∘C. The solution was enclosed in a
500 mL Erlenmeyer flask with a rubber stopper. The rubber stopper contained
two feed-throughs for tubes that were used to maintain a constant gas flow
through the solution. Humidified and thermally equilibrated CO2 gas
(same temperature as the solution) was passed through the CaCO3 solution
(Table 1) at a rate of ∼ 0.03–0.1 mL s-1 to prevent carbonate
precipitation before complete isotopic equilibrium was achieved. The CO2
gas was humidified and adjusted to the experimental temperature by bubbling
it slowly through an Erlenmeyer flask filled with de-ionized water and
contained in the temperature-controlled water bath. The equilibration period
varied between 3 h at 91 ∘C and 23 h at 23 ∘C. pH values
during equilibration are below pH 6 due to the continuous CO2 flux.
Photomicrographs of CaCO3 minerals precipitated in the
laboratory experiment. Scale bar is 100 µm in (a)-(c) and
200 µm in (d). Vaterite crystals formed at 50 ∘C
in experiment NA-6 (a), at 70 ∘C (b) and at 91 ∘C
(d). At 80 ∘C a mixture of aragonite, calcite and vaterite
was precipitated (c).
After equilibration, NaCl was added, reaching concentrations between 4.2 and
6.4 mol L-1 (Table 1). The added NaCl (Sigma
Aldrich®) has a purity of ≥ 99 %
and contains minor traces of sulfates (≤ 200 ppm), alkaline earth
metals (≤ 100 ppm) and bromides (≤ 100 ppm). Carbonate
precipitation was induced by slowly bubbling N2 tank gas (BOC UK, normal
grade) through the solution. The N2 gas was humidified and adjusted to
the experiment temperature using the same procedure as for the CO2 gas.
The bubbling rate was set to about 1 bubble per second
(∼ 0.03 mL s-1). The gas stream was humidified to prevent water
evaporation and a potential change in the solution δ18O value over
time. Minerals always formed on the bottom or the side wall of the Erlenmeyer
flask. No crystals were observed on the surface of the solution. After 2–19
days, the solution was passed through a double layer of
Whatman® filter paper (grade 1 with a pore
size of 11 µm). Crystals on the glass walls were loosened by a thin
PVC plastic tube and flushed out with de-ionized water. The precipitated
minerals were air-dried at room temperature before microscopic and XRD (X-ray
diffraction) analysis was conducted.
Samples
Mineral description, microscopy and SEM
Depending on the experiment, temperature and duration, between 5 and 235 mg
CaCO3 was precipitated (Tables 1, 2). Low carbonate recovery < 20 mg
is linked to short experiment duration (3–6 days) at lower temperatures
(< 70 ∘C, Table 2). In contrast, experiments with a longer
duration of 14–24 days at temperatures ≤ 50 ∘C yielded on
average 135 mg. At 80–91 ∘C it was sufficient to allow 2 days
for mineral precipitation to obtain 80–90 mg of calcite. Note that in all
experiments the initially dissolved amount of calcite was similar at about
360 (±20) mg in 500 mL of de-ionized water (Table 1).
Photomicrographs of CaCO3 minerals precipitated in the
laboratory experiments. Scale bar is 100 µm. Vaterite crystals
formed at 23 ∘C (a), at 37 ∘C in experiment NA-3
(b) and NA-4 (c), and at 50 ∘C in NA-5
(d).
Vaterite can be distinguished from other CaCO3 polymorphs by its morphology. Calcite rhombohedra and aragonite needles can be recognized by
light microscopy (e.g. Fig. 2c). Vaterite crystals can be similar in size
but are more irregular and show a spherulitic shape (Figs. 3, 4). Inspection
of large vaterite crystals under normal and polarized light reveals a complex
growth history. Various globular segments of 50–100 µm with an
internal spherulitic growth pattern coalesce into one larger crystal
(Fig. 4). Vaterite crystals showed a typical size of 50 µm
(Figs. 2, 3, 4), whereas in a few experiments, crystals of up to
500 µm were observed. Experiments at 70 and
91 ∘C also resulted in vaterite crystals in the 50 µm size
range; however, these are composed of many small (∼ 10 µm)
globular sub-segments. A peculiarity of vaterite crystals precipitated at
23 ∘C is the combination of rounded, spherical shapes with sharp
angular forms (Fig. 3a). Together with the larger crystals sizes observed at
this temperature; this points towards slower mineral growth. All minerals of
this study were investigated with XRD in addition to optical microscopy.
Note that optical microscopy alone may be ambiguous and should be
complemented by additional methods (e.g. XRD).
Scanning electron microscope (SEM) images were made at the Institute of Earth
Sciences at Heidelberg University to investigate the morphology in more
detail. The scanning electron microscope LEO 440 was used for imaging. It has
a tungsten cathode, was operated at an accelerating voltage of 20kV, and enables a minimal resolution of ca. 5 nm. Samples were sputtered with a thin
gold layer for imaging and with carbon for elemental analysis. A summary with
characteristic vaterite aggregates is shown in Fig. 5. The size of individual
grains that make up the vaterite aggregates was investigated using SEM images
(Figs. 5, 6). The grain size decreases with increasing temperature, from
about 100 µm at 23 ∘C, 10 to 20 µm at
50 ∘C, to < 10 µm at 91 ∘C. The relatively
large grains detected at 23 ∘C are internally composed of
∼ 2 µm long elongated fibrous crystallites (Fig. 6a). The
smallest grains observed at 91 ∘C are of a similar size of a few
micrometres. Sub-micrometre-sized crystallites and framboidal aggregates
(e.g. Nehrke and Van Cappellen, 2006) were not observed. The minerals show a
radial growth pattern from a central nucleus leading to spherical
conglomerate particles (Fig. 5a). The spherical shape is still dominant at
50 ∘C (Fig. 5b); however, it changes to the growth of flat platelets at
91 ∘C (Fig. 5c). The radiating growth pattern at 91 ∘C is
restricted to two dimensions with a tree-like branching structure
characteristic for diffusion-controlled dentritic crystallization (Fig. 5d).
Close-up photomicrographs of vaterite minerals. Scale bar is
50 µm in (a) and (c), and 200 µm in
(b). Panel (b) shows a vaterite crystal using polarized light.
Additional elemental analyses on carbon-sputtered vaterite grains using
energy-dispersive X-ray spectroscopy (EDS) at the SEM system at Heidelberg
University revealed minor traces of sodium and chloride to be occasionally
incorporated in the vaterite mineral. Quantitative analysis was not
attempted due to the very sporadic and dispersed occurrence of the sodium
chloride crystals.
XRD analysis
The carbonate samples were analysed at the National History Museum, London,
using an Enraf Nonius FR 590 powder diffractometer with Cu-Kα
radiation (40 kV, 35 mA). In brief, the sample powder was placed in a thin
layer on a sapphire substrate and measured by fixed beam-sample-detector
geometry. Analysis times were adjusted to the counting statistics and varied
between 10 and 90 min. Signals and phase fractions were evaluated by
comparing measured spectra with a mineral database using the program X'Pert
Highscore (PANalytical B.V., 2009). Peak positions were calibrated with two
standards (silver behenate and quartz). For phase quantification, a pure
calcite standard and an aragonite standard were additionally measured.
SEM images of vaterite aggregates in the order of increasing
precipitation temperatures. Scale bars are 100 µm in (a)
and (b), 200 µm in (c) and 10 µm in
(d). Samples NA-1 (a), NA-4 (b), and NA-9
(c and d) were used for imaging.
For most samples, the dominant XRD peaks were found at 20.98(±0.04),
24.86(±0.02), 27.03(±0.03), 32.74(±0.03), 43.79(±0.09), and
50.0(±0.04)∘ (2θ, Fig. 7). In contrast, the
characteristic and dominant calcite peak of the calcite standard is observed
at 29.46∘; those of the aragonite calibration standard are at 26.36,
27.35, 33.25 and 46.01∘ (2θ, Fig. 8). The aragonite peaks at
26.36 and 27.35∘ were resolved for the pure aragonite standard, but
the peak at 27.35∘ overlapped with the vaterite peak at
∼ 27.03∘ for aragonite–vaterite mixtures. Our laboratory
CaCO3 samples are clearly different from aragonite and calcite, but
coincide with the XRD data and d-spacing of vaterite. Kabalah-Amitai et
al. (2013) measured vaterite d-spacing of 2.07 and 3.63Å, corresponding
to 43.69 and 24.50∘ (2θ at Cu-Kα radiation). Earlier
work of Dupont et al. (1997) determined similar d-spacing values of 4.254,
3.591, 3.307, 2.741, 2.07 and 1.826Å, corresponding to 20.86, 24.77,
26.94, 32.64, 43.69 and 49.90∘ (2θ at Cu-Kα
radiation). Our own results are close to these values confirming the
precipitates to be composed of vaterite. In the case of the 37 ∘C
experiment, calcite is additionally present as a minor phase (about 5 %,
visible in the peak at 29.46∘, Fig. 7), whereas only traces of
calcite are found in the 91 ∘C experiment. An exact quantification
of this small calcite fraction was not possible due to a general uncertainty
of about 3 % in the phase quantification. Half of the mineral phase at
80 ∘C consists of calcite, whereas the other half is made up of
equal proportions of aragonite and vaterite (Table 2, Fig. 8). Aragonite is
also a minor phase in the 91 ∘C experiment (6 %, Fig. 7).
High-resolution SEM images of vaterite aggregates showing the
internal structure in the order of increasing precipitation temperatures
(a, 23; b, 50; c and d, 91 ∘C). Two
different growth structures are observed at 91 ∘C: aggregation of
µm-sized flat crystallites (c) and tree-like branching
growth resulting in flat platelets (d).
Isotope analysis
Oxygen, carbon and clumped isotopes were analysed at the Qatar Stable Isotope
Laboratory at Imperial College. Details of the sample preparation and mass
spectrometric procedures are given in Kluge et al. (2015). In brief, per
analysis ∼ 5 mg sample was dissolved in orthophosphoric acid at
70 ∘C to produce CO2 for the mass spectrometric measurement.
The CO2 is cleaned manually, comprising a step for cryogenic water
separation and one for contaminant removal via porous polymers
(Porapak™ Q). Analyses were done using two dual-inlet isotope ratio mass spectrometers (Thermo Scientific MAT 253) that
measure sample and reference gas alternately. Individual analyses have a
precision of 0.2 ‰ for δ18O, 0.1 ‰ for δ13C, and
0.03 ‰ for Δ47, based on replicate analyses of standards.
Samples were measured repeatedly (typically three times) to reduce the
uncertainty.
Carbonate δ18O values follow the trend determined by the
temperature-dependent isotopic fractionation between water and calcite
(Fig. 9). For calculation of the expected carbonate δ18O values, we
used the fractionation factors of Kim and O'Neil (1997) and a water value of
-6.4 ± 0.7 ‰ (-7.8 ± 0.5 ‰ for the later
repeat experiments NA-4 and NA-6). The water δ18O value was not
directly measured but is based on back-calculation of aragonite and calcite
samples that were precipitated in close temporal connection (within 3–12
days) of the vaterite experiments using the same water source and the same
precipitation technique (see Kluge et al., 2015). The back-calculated water
values from the pure Ca(HCO3)2 solution, used as a reference, correspond
to the observed range of surface and ground water values of the London
metropolitan area (Darling, 2003). For all experiments, the water was taken
from the local water supply. Although we have a defined water reference value
based on the pure Ca(HCO3)2 solution, we cannot exclude short-term
fluctuations of the tap water δ18O. However, it is unlikely that it
exceeds the observed long-term variability of ±0.7 ‰
(June–October 2012) and ±0.5 ‰ (March–April 2013) for the
back-calculated solution δ18O value.
X-ray diffraction pattern of crystals from the laboratory
experiments. The minerals that grew at 23 ∘C show a pure vaterite
signal. Similarly, minerals formed at 50 and 70 ∘C yield an almost
pure vaterite signal with a non-quantifiable fraction of calcite (traces,
≤ few %). The samples at 37 and 91 ∘C contain a minor
fraction of calcite and aragonite (≤ 6 % in total).
The δ13C values vary between -18 and -26 ‰ and reflect the
negative signature of the CO2 tank gas used during the equilibration
phase (Table 2). The δ13C values do not show a temperature dependence.
Clumped isotope Δ47 values of vaterite samples decrease with
increasing temperature and are similar to calcite or aragonite–calcite
mixtures precipitated at the same temperature (Fig. 10). Δ47 values
of calcite and calcite–aragonite mixtures were taken from Kluge et al. (2015)
for comparison.
X-ray diffraction pattern of crystals from laboratory experiments.
The experiment at 80 ∘C produced a mixture of calcite, aragonite and
vaterite (lower panel). For comparison, the XRD pattern of pure calcite is
shown (upper panel). This example shows calcite that precipitated at
25 ∘C from a pure CaCO3 supersaturated solution without NaCl
addition. The peaks in the lower panel are labelled according to the related
mineral structure (A: aragonite; C: calcite; V: vaterite).
Discussion
Vaterite was obtained over the entire experimental temperature range of
23–91 ∘C. It is detected either as the only phase (23, 50,
70 ∘C) or as the major phase (≥ 94 %) with minor
contributions from calcite or aragonite (37, 91 ∘C). An exception is
the experiment at 80 ∘C where all anhydrous CaCO3 polymorphs
were precipitated simultaneously. On average 80 mg of vaterite was formed per
experiment. This amount may be increased by longer experiment runs or by
up-scaling of the setup using larger beakers with the same solution
concentrations. A longer experiment duration appears to be the most effective
approach. Considering the experiments from 23 to 70 ∘C only, the yield
increases exponentially with the duration, reaching a recovery rate (relative
to the initially dissolved CaCO3) of 70 % after ∼ 570 h
(Fig. 11). An extrapolation of this trend is not easily possible as the precipitation rate will eventually decrease due to the continuously decreasing supersaturation. Another option
of increasing the CaCO3 supersaturation in the initial solution was not
tested, but has to be treated carefully. A higher initial chemical potential
may produce a higher yield but also lead to the precipitation of other forms
of CaCO3 such as ikaite (calcium carbonate hexahydrate,
CaCO3(H20)6) or amorphous calcium carbonate (Kawano et al.,
2009).
This study shows that vaterite precipitation is not limited to a certain
temperature range, e.g. to room temperature or from 10 to 48 ∘C
(Gussone et al., 2011), but can be performed at least up to 91 ∘C. A
pressurized reaction vessel that prevents boiling of the solution could be
used to extend vaterite mineral formation to much higher temperatures (e.g.
Kluge et al., 2015). A thermally and isotopically equilibrated CaCO3
supersaturated solution could be injected into the thermally equilibrated
and saturated NaCl solution of a pressurized reaction vessel.
Oxygen isotope fractionation factor 1000 lnα (CaCO3-H2O) of vaterite (circles) relative to expected values for
calcite (Kim and O'Neil, 1997; solid line). The sample at 80 ∘C
consists of a mixture of calcite, aragonite and vaterite. At 37 and
50 ∘C repeat experiments were performed yielding fractionation
factors in agreement with the previous experiment at the same temperature.
The detection of vaterite minerals over the large temperature interval of
this study and its predominating character is surprising, given that many
other studies emphasized the low stability of vaterite (e.g. McConnel,
1959). McConnel (1959) states that vaterite dissolves at room temperature at
contact with water. However, our precipitates were air-dried at room
temperature on Whatman® filter paper and
stayed wet for a few hours, but did not transform into calcite. Furthermore,
vaterite minerals were stored between several weeks and a year before being
analysed by XRD and SEM. Despite long storage periods vaterite did not
transform into other CaCO3 polymorphs; this implies that vaterite can be
precipitated and stored for periods that are long enough to enable precise
and detailed experimental analyses. Independent evidence for the stability of
vaterite over years comes from biogenic samples such as bivalves, mollusks
and other marine organisms (Lowenstamm and Abbott, 1975; Spann et al., 2010;
Nehrke et al., 2012).
Isotopic analysis of vaterite
Our study presents a well-defined method for the isotopic study of vaterite.
The technique for vaterite precipitation was modified from the procedures of
McCrea (1950), O'Neil et al. (1969) and Kim and O'Neil (1997) that were
developed for isotopic studies on calcite and aragonite minerals. The values
of Kim and O'Neil (1997) are commonly used as a reference for the oxygen
isotope fractionation between calcite and water. Furthermore, for the
temperature calibration of clumped isotopes, an analogous method was applied
to precipitate calcite (Ghosh et al., 2006; Zaarur et al., 2013). The
closeness of the methodology for synthetic vaterite and calcite precipitation
ensures good comparability of the obtained results with widely used
calibration and fractionation data.
The long equilibration procedure used in our experimental approach (Table 1)
enables isotopic equilibration between the dissolved inorganic carbon (DIC)
and water and among the DIC species. A 99 % equilibrium between oxygen
isotopes in water and DIC takes about 9 h at 25 ∘C and a pH of
∼ 8, whereas it is less than 2 h at temperatures above 40 ∘C
(Beck et al., 2005). For comparison,
the equilibration duration was 23 h at 23 ∘C, 14–21 h at
37.5 ∘C, and 3–17 h above 40 ∘C (Table 1). This provides
the necessary basis for a meaningful isotopic analysis of the precipitated
vaterite which has not been attempted in a systematic manner so far.
Measured Δ47 values of vaterite (circles) relative to
expected values following the calibration line of Kluge et al. (2015). The
sample at 80 ∘C (marked by an asterisk) consists of a mixture of
calcite, aragonite and vaterite. The calibration line of Kluge et al. (2015)
was mainly determined on calcite.
CaCO3 formed per experiment vs. duration. The initially
dissolved CaCO3 amount was identical in all experiments (about 370 mg).
The yield of experiments conducted below 80 ∘C follows an
exponential relationship with duration.
Beyond the isotopic equilibration of the DIC with water, the precipitation
rate and the ionic concentration of the solutions can affect isotope values.
The growth rate has to be considered in the interpretation of isotope values
as differences can be substantial between slow and rapidly grown minerals (up
to 2 ‰ for δ18O; Coplen, 2007; Dietzel et al., 2009;
Gabitov et al., 2012). The growth rate in our study was not precisely
monitored, but can be estimated from the derivative of the general
relationship between experiment duration and mineral yield (Fig. 11). The
mineral growth for temperatures below 80 ∘C started very slowly with
a value of ∼ 1.4 × 10-10 mol s-1 and increased
exponentially (e.g. to 1.1 × 10-9 mol s-1 after
350 h). Converting it relative to the growth surface (using the walls of the
Erlenmeyer flask as a first-order estimate), the growth rate was about
3.7 × 10-9 mol (m2 s)-1 at the beginning of the mineral
formation and increased to 2.9 × 10-8 mol
(m2 s)-1 at 350 h. These values only
give an impression of the order of magnitude due to the limitation of the
assumptions, but already demonstrate that the mineral formation was unlikely
affected by rapid growth disequilibrium that is typically encountered at
values above 10-7–10-8 mol
(m2 s)-1 (compare with, for example, Watkins et al., 2013). In contrast, in other
vaterite precipitation techniques, two solutions were mixed, leading to almost
instantaneous precipitation (e.g. Nebel and Epple, 2008). As rapid mineral
growth may induce disequilibrium fractionation related to a mineral surface
effect (Watson, 2004; Dietzel et al., 2009; Watson and Müller, 2009;
DePaolo, 2011; Reynard et al., 2011; Gabitov et al., 2012; Gabitov, 2013),
experiments with quasi-instantaneous mineral growth are not suitable for
isotope studies. Consequently, we neither rapidly grew vaterite nor conducted
isotope measurements on potentially fast growth phases.
Traditionally, vaterite was synthesized from mixtures of CaCl2,
K2CO3 (Kamhi, 1963) and admixtures of calgon (McConnell, 1960), or
included other surfactants (Mori et al., 2009). In other experiments, a
CaCO3 supersaturated solution was treated with surfactants (Dupont et
al., 1997) or polymeric substances (Kirboga and Oner, 2013). In a few
experiments, Na2CO3 replaced K2CO3 as the solution containing
the carbonate ion (Nebel and Epple, 2008). The use of CaCl2 and
especially K2CO3 could impact on the isotopic values of the
forming minerals via preferential fractionation related to the hydration
sphere of the Ca2+ and K+ ions (Taube, 1954; Sofer and Gat, 1972;
O'Neil and Truesdell, 1991) and thus should either be restricted to low
concentrations or avoided. Our method uses only NaCl as an additive that has
been confirmed not to affect the isotope values of the DIC (e.g. O'Neil and
Truesdell, 1991).
Before discussing the measured vaterite δ18O values and its
implication for the oxygen isotope fractionation factor αCaCO3-H2O, we note that we did not analyse the oxygen
isotope composition of the solution per se. However, as the solution water
is ultimately taken from the local water supply that reflects the London
Metropolitan ground- and surface water δ18O of -6 to
-8 ‰ (Darling, 2003) and which is cross-examined using values from
independent NaCl-free experiments carried out in parallel, we have a defined
reference water δ18O value (-6.4 ± 0.7,
-7.8 ± 0.5 ‰ for the later repeat experiments NA-4 and NA-6;
see Sect. 3.3). Thus, our results give a first-order guideline with respect
to the temperature dependence of αCaCO3-H2O and
can provide an upper limit for the deviation of the fractionation factor of
vaterite compared to aragonite and calcite. We compare our values to data of
Kim and O'Neil (1997) due to similar experimental procedures and their
universal use. Note that the exact value of the equilibrium oxygen isotope
fractionation factor for calcite is under debate (Coplen, 2007; Dietzel et
al., 2009; Gabitov et al., 2012; Kluge et al., 2014).
The fractionation 1000 lnα(CaCO3-H2O) of vaterite
closely follows the values of Kim and O'Neil (1997) for calcite (Fig. 9).
Over the entire experimental temperature range, vaterite values agree with the calcite values within
uncertainty. Repeat experiments at 37 and
50 ∘C confirm the initial results by providing almost identical
fractionation factors. The average deviation is 0.0 ± 0.4 ‰
and thus the vaterite oxygen isotope fractionation factor
αCaCO3-H2O cannot be distinguished from that of
calcite. Tarutani et al. (1969) observed vaterite to be enriched in 18O
by 0.5 ‰ relative to calcite at 25 ∘C. Kim and
O'Neil (1997) detected a similar difference of 0.6 ‰ at 25 and
40 ∘C. Taking into account the measurement uncertainties, our experiments are consistent with both studies, which indicate a small or negligible difference of the fractionation
αCaCO3-H2O between vaterite and calcite. Our study
additionally constrains this value over the larger temperature range from
23 to 91 ∘C.
Carbonate clumped isotope Δ47 values are only determined by the
mineral formation temperature at equilibrium conditions and are independent
of the solution δ18O and δ13C values (for reviews see Eiler, 2007, 2011). We use the Δ47–T calibration of Kluge
et al. (2015) as a reference relationship as it was determined in the same
laboratory, using the same preparation and measurement techniques (T in K,
Δ47 in ‰):
Δ47(T)=0.98×(-3.407× 109/T4+2.365× 107/T3-2.607× 103/T2-5.880/T)+0.293.
Equation (1) is given in the absolute reference
frame of Dennis et al. (2011). Vaterite Δ47 values scatter around
the Δ47–T line of Eq. (1) with an average difference of
-0.003 ± 0.013 ‰ and are thus indistinguishable from the
calibration line (Fig. 10). Subtle differences in the mineral structure of
the CaCO3 polymorph vaterite appear to be irrelevant for the
13C–18O clumping.