ABSTRACT
KULBOK, KATHERINE ELIZABETH. Elucidating the Effect of Biomolecule Structure on
Calcium Carbonate Crystal Formation. (Under the direction of Owen Duckworth).
Calcium carbonate minerals are common in soils and natural waters, and their
reactivity affects the movement of inorganic contaminants, the alkalinity and pH of marine
and freshwater ecosystems, and the global cycling of carbon. These minerals are often
biologically produced, with their formation regulated by biomolecules. This study examines
the kinetics, thermodynamics, phase, and morphology of calcium carbonate crystals
precipitated in the presence of carboxylate-containing biomolecules, including citric acid,
succinic acid, and aspartic acid. The experiments utilize a unique (NH4)2CO3 gas-diffusion
reactor, which allows in-situ measurements of chemical conditions during the precipitation
and growth of crystals.
Continuous monitoring of the in-situ conditions of pCO2, pH, and
[Ca2+] provides data on the supersaturation at which nucleation occurs and the kinetics of
mineral growth.
These data were implemented in the calculation of the kinetic rate
coefficient for each trial, which helps to quantify the rates of reaction. The use of scanning
electron microscopy and X-ray diffraction provides information on the morphology and
mineralogy of precipitates. Results show that these particular biomolecules affect the rate of
calcium carbonate formation but do not have a large effect on the polymorph of the
precipitates. The inclusion of biomolecules increased the kinetic rate coefficient by an order
of one or two magnitude compared to organic-free trials. The dominant polymorph created
in all trials was vaterite. Small amounts of calcite were precipitated only in the presence of
citric acid and at high concentration of calcium (10 mM).
Elucidating the Effect of Biomolecule Structure on Calcium Carbonate Crystal Formation
by
Katherine Kulbok
A thesis submitted to the Graduate Faculty of
North Carolina State University
in partial fulfillment of the
requirements for the degree of
Master of Science
Soil Science
Raleigh, North Carolina
2012
APPROVED BY:
_______________________________
Owen Duckworth
Committee Chair
______________________________
Jot Smyth
________________________________
Matthew Polizzotto
________________________________
Treavor Kendall
BIOGRAPHY
Katherine Kulbok was born in St. Louis, Missouri and raised in Charlottesville, Virginia.
She received a BS in Environmental Science from Virginia Tech. After graduation she
obtained a six month work visa for the UK which enabled her to live and work in London,
England. She later interned at Crater Lake National Park in Oregon where she completed a
greenhouse gas emissions survey for parks in the Klamath region. In January 2010 she began
her study in Soil Science as a master’s student at NC State University.
ii
TABLE OF CONTENTS
List of Tables ............................................................................................................... v
List of Figures ............................................................................................................. vi
Chapter 1 Introduction ............................................................................................................. 1
1.1 Introduction ............................................................................................................ 2
1.2 Global Environmental Importance of Carbonates...................................................3
1.2.1 The Global Carbon Cycle and Anthropogenic Influence ………………3
1.2.2 Ocean Acidification………………………………………………….….5
1.3 Terrestrial and Aquatic Carbonates and Their Reactivity …….……………..……7
1.3.1 Alkalinity ……………..……………………………………..………….7
1.3.2 Liming ……………………………………………………………..……7
1.3.3 Sequestration of Metals ….…..…………………………………….……8
1.4 Carbonate Crystal Chemistry …………………..…….………….…….…...……..9
1.5 Aqueous Chemistry of Carbonates ……………………………….…….…….....11
1.5.1 Dissolution ………………..………………………………………...…12
1.5.2 Growth …………….………………………………………………..…13
1.5.3 Effects of Biomolecules…………………………………...………..….14
1.6 Summary………………….………………..…….…………………………..…..15
1.7 References .............................................................................................................17
Chapter 2 Experimental Design and Kinetic Modeling for Calcium Carbonate Precipitation24
2.1 Introduction .......................................................................................................... 25
2.2 Materials and Methods ….……………………………………………………….26
2.2.1 Growth Reactor………….………………………….………….…..….26
2.2.2 Experimental Setup…………………………….…………….…..……28
2.2.3 Kinetic Modeling of Calcium Carbonate Precipitation……………..…31
2.3 Results and Discussion ........................................................................................ 34
2.3.1 pH………………….……………………………………….................. 34
2.3.2 Carbon Dioxide……….………………………………..………..……. 34
2.3.3 Calcium Data…….…………………………………..……….………. 35
2.3.4 Citric Acid Complexation…………….…………...………………..….36
2.3.5 Rates of Calcium Loss and Kinetic Modeling of Calcium Loss ……....37
2.3.6 Effects of Biomolecules on Calcium Carbonate Precipitation……...…38
2.4 Conclusions........................................................................................................... 39
2.5 References ............................................................................................................ 41
2.6 Appendix …………………………………………………………………….…..55
2.6.1 Development of Reactor……………………………………………….55
2.6.2 Limitations of the Reactor……………………………………………...57
2.6.3 References……………………………………………………………...58
Chapter 3 Determination of Phase and Morphology………………………………………..61
3.1 Introduction .......................................................................................................... 62
3.2 Materials and Methods………………………………….......................................63
3.2.1 The Chemical System............................................................................ 63
iii
3.2.2 Materials………………….................................................................... 64
3.2.3 X-ray Diffraction ……….….………………………………..……….. 64
3.2.4 Scanning Electron Microscope ………………….…………..……….. 66
3.3 Results and Discussion ........................................................................................ 66
3.3.1 XRD Results……………...................................................................... 66
3.3.2 SEM Results……………...................................................................... 67
3.3.3 Formation of CaCO3 in the Absence of Biomolecules ……………..…69
3.3.4 Effect of Biomolecules ..........................................................................70
3.4 Conclusions........................................................................................................... 71
3.5 References ............................................................................................................ 73
Chapter 4 Summary and Conclusions .................................................................................... 90
4.1 Summary of Major Findings ................................................................................ 91
4.2 Broader Implications of Major Findings .............................................................. 91
4.3 Final Thoughts .................................................................................................... 92
4.3.1 Improvements for Present Work .......................................................... 92
4.3.2 Suggestions for Future Work ............................................................... 93
4.4 References …...................................................................................................... 95
iv
LIST OF TABLES
Table 2.1 Table of kinetic rate constants for the formation of vaterite. Uncertainties are
estimated as standard deviations of replicate measurements………………………53
Table 2.2 Table of average saturation state for each set of trials. Calculations were based on
the assumption that the dissolved phase and the gas phase for CO2 are in equilibrium.
Uncertainties are estimated as standard deviations of replicate measurements.……54
Table 3.1. Intensity ratios and corresponding degrees for the determination of synthetic
vaterite (Kamhi, 1963), theoretical vaterite (Wang and Becker 2009), calcite (Graf
1961), and aragonite (de Villiers 1971)……………………………………………75
Table 3.2. Comparison of morphologies detected via x-ray diffraction and scanning electron
microscopy and the number of trials in which that polymorph was detected………83
v
LIST OF FIGURES
Figure 1.1. The CO2 concentration, in ppm, versus time, from observations for Arctic ice
floes, 1957–1958, and Point Barrow, Alaska, 1961–1968 and 1974–2003, shown as
monthly averages (dots) and a spline fit combined with seasonal harmonics
increasing in amplitude at an assumed constant rate (smooth curve) (Keeling et al.
2011)………………………………………………………………………………..21
Figure 1.2. The The main components of the natural carbon cycle. Thick lines indicate gross
primary production and respiration from land and exchange between the ocean and
atmosphere, which are the most important CO2 fluxes that take place over a shortterm timescale. Thin lines represent other natural fluxes and dashed lines indicate
carbon flux as CaCO3, which are significant over longer time scales. All units are in
Pg C yr-1 and refer to data collected during the 1980s (Prentice et al. 2001)………22
Figure 1.3. The effect of pH on the speciation of carbonates. Created with Visual MINTEQ
using 0.04 M H2CO3………………….………………………………………..…….23
Figure 1.4. Molecular structures of carboxyl-containing biomolecules including (A) aspartic
acid (B) citric acid (C) succinic acid………………………………………….…….23
Figure 2.1. A photograph of the crystal growth chamber……………………………………44
Figure 2.2. A schematic of the growth chamber, designed to allow monitoring and
manipulation of aqueous and gaseous chemical conditions during the growth of
carbonate minerals………………………………………………………………….44
Figure 2.3. A comparison of pH over time for three repetitions of a trial set containing 10
mM CaCl2. Each trial started at a pH of approximately 3 and ended at a pH of
approximately 9. This data set is a concise representation of the behavior of pH in
each experiment……………………………………………………………………..45
Figure 2.4. A comparison of CO2 concentration over time for three repetitions of a trial set
containing 10 mM CaCl2. These trials show the general increase over time of CO2
concentration……………………………………………………………………….46
Figure 2.5. Comparison of calcium concentration over time for three trials containing 10 mM
CaCl2……………………………………………………………………………….47
Figure 2.6. Comparison of calcium concentration measured by calcium ISE and atomic
adsorption spectroscopy for a trial that started with 0.01 M Ca2+ done on
09_09_2011…………………………………………………………………………48
vi
Figure 2.7. Comparison of calcium concentration measured by calcium ISE and atomic
adsorption spectroscopy for a trial that started with 0.001 M Ca2+ done on
05_26_2011. This figure shows electrode drift at the start of the experiment, and that
the electrode readings become closer to the AAS readings as time goes on…..……48
Figure 2.8. Calcium loss over time for trial 04_15_2011……………………..……………..49
Figure 2.9. Calcium ISE test done at pH approximately 3 with a solution of 1 mM CaCl2
(0.001 M Ca2+)…………………………………………………………………..…...49
Figure 2.10. Calcium ISE test done at pH approximately 10 with a solution of 1 mM CaCl2
(0.001 M Ca2+)……………………………….…………….…………………..……50
Figure 2.11. Calcium ISE test titration. A solution of 1 mM CaCl2 (0.001 M Ca2+) at
approximately pH 3 was titrated with 0.5 M NaOH at an interval of 1 drop every 30 seconds
for six minutes, until an end pH of approximately 10……………………………………….50
Figure 2.12. Calcium loss over time for a trial initially containing 1 mM CaCl2 and 0.5 mM
citric acid, completed on 06_02_2011…………………………..…………………..51
Figure 2.13. Concentration of aqueous species involving citrate formed in a solution of 1 mM
(1x10-3 M) CaCl2 with 0.5 mM (5x10-4 M) citrate at a range of pH from 3 to 10.
Model created with MINEQL+………………………………………………………51
Figure 2.14. Comparison of calcium concentration measured by calcium ISE and atomic
adsorption spectroscopy, as well as calculations for calcium in solution not
complexed by citrate. This trial that started with 0.001 M Ca2+ and contained 0.5 mM
citric acid and was done on 06_02_2011. MINEQL+ calculations were done with 1
mM Ca2+, 0.5 mM citrate, and pH and pCO2 concurrent with each time point of this
trial…………………………………………………………………………….….…52
Figure 2.15. Concentration of calcium over time for a trial containing 1 mM calcium and 0.5
mM succinic acid completed on 06_28_2011. Evidence of complexation may be seen
between hours 0 to 3……………………………..…………………………………52
Figure 2.16. SEM image at 100x of a trial containing 100 mM Ca(NO3)2 and 40 mM
NaHCO3 done with approximately 10 g of ammonium bicarbonate (NH4HCO3) on
03_17_201……………………………………………………………………………59
Figure 2.17. SEM image of a trial containing 40 mM Ca(NO3)2 and 40 mM NaHCO3 done
with 10 g of ammonium carbonate ((NH4)2CO3) with 20 mL deionized water on
06_10_2010…………………………………………………………………………59
Figure 2.18. SEM image at 1500x of a trial containing 40 mM CaCl2 and 10 g (NH4)2CO3
with 20 mL deionized water on 07_20_2010………….……………..…………….60
vii
Figure 3.1. Powder X-ray diffraction spectra of trials containing 10 mM CaCl2. Trial
04_15_2011 and 09_09_2011 were determined as vaterite, syn based on peaks at
27.05° and 32.74° (Kamhi, 1963). The trial for 04_18_2011 was determined to be
vaterite based on peaks at 26.64° and 32.47° (Wang and Becker 2009)…………….76
Figure 3.2. Powder X-ray diffraction spectra of trials containing 5 mM CaCl2. Trial
04_05_2011 was determined as vaterite, syn based on peaks at 27.05° and 32.74°
(Kamhi, 1963). The trials for 04_08_2011 and 04_11_2011 were determined to be
vaterite based on peaks at 26.64° and 32.47° (Wang and Becker 2009)……………77
Figure 3.3. Powder X-ray diffraction spectra of trials containing 1 mM CaCl2. Trial
05_26_2011 was determined as vaterite, syn based on peaks at 27.05° and 32.74°
(Kamhi, 1963). The trials for 05_12_2011 and 05_16_2011 did not detect a calcium
carbonate polymorph present, however it is obvious from the peaks present that these
forms are also vaterite……………………………………….……………………….78
Figure 3.4. Powder X-ray diffraction spectra of trials containing 1 mM CaCl2 and 0.5 mM
aspartic acid. All trials were determined to be vaterite, syn based on peaks at 27.05°
and 32.74° (Kamhi, 1963)……………………………………...…………….………79
Figure 3.5. Powder X-ray diffraction spectra of trials containing 1 mM CaCl2 and 0.5 mM
succinic acid. Trials from 06_14_2011 and 06_28_2011 both detected vaterite, syn
based on peaks at 27.05° and 32.74° (Kamhi, 1963). The trial for 06_30_2011 did
not detect a calcium carbonate polymorph present…………………………………80
Figure 3.6. Powder X-ray diffraction spectra of trials containing 1 mM CaCl2 and 0.5 mM
citric acid. No calcium carbonate polymorph was detected………………..……….81
Figure 3.7. Powder X-ray diffraction spectra of the polycarbonate material on which the
crystals were grown…………………..……..……………………………………….82
Figure 3.8. Visual representations of the three crystalline CaCO3 polymorphs (A)
rhombohedral calcite Meldrum (2010), (B) elongated aragonite Jorgensen (1976), and
(C) spheroidal vaterite Sikorski et al. (2010)………………………………………..82
Figure 3.9. Representative SEM images of common morphologies observed in crystal growth
trials using a growth solution of 10 mM CaCl2. (a) 500X and (b) 3000X Trial
04_15_2011: vaterite. (c) 500X and (d) 1000X Trial 04_18_2011: vaterite and calcite.
(e) 500X and (f) 3000X Trial 09_09_2011: vaterite and calcite. All trials used 5 g
ammonium carbonate + 20 mL deionized water…………………….………….…84
Figure 3.10. Representative SEM images of common morphologies observed in crystal
growth trials using a growth solution of 5 mM CaCl2. (a) 500X and (b) 3000X Trial
04_05_201: vaterite. (c) 500X and (d) 1200X Trial 04_08_2011: vaterite. (e) 500X
and (f) 1600X Trial 04_11_2011: vaterite. All trials used 5 g ammonium carbonate +
20 mL deionized water………………………………………………………………85
viii
Figure 3.11. Representative SEM images of common morphologies observed in crystal
growth trials using a growth solution of 1 mM CaCl2. (a) 500X and (b) 5000X Trial
05_12_201: vaterite and calcite. (c) 500X and (d) 110X Trial 05_16_2011: vaterite.
(e) 500X and (f) 100X Trial 05_26_2011: vaterite. All trials used 5 g ammonium
carbonate + 20 mL deionized water…………………………………………………86
Figure 3.12. Representative SEM images of common morphologies observed in crystal
growth trials using a growth solution of 1 mM CaCl2 + 0.5 mM citric acid. (a) 500X
and (b) 3500X Trial 06_02_201: vaterite and calcite. (c) 500X and (d) 100X Trial
06_07_2011: vaterite and calcite. (e) 500X and (f) 1000X Trial 06_09_2011: vaterite
and calcite. All trials used 5 g ammonium carbonate + 20 mL deionized water…..87
Figure 3.13. Representative SEM images of common morphologies observed in crystal
growth trials using a growth solution of 1 mM CaCl2 + 0.5 mM succinic acid. (a)
500X and (b) 4000X Trial 06_14_201: vaterite. (c) 500X and (d) 1500X Trial
06_28_2011: vaterite. (e) 500X and (f) 1600X Trial 06_30_2011: vaterite. All trials
used 5 g ammonium carbonate + 20 mL deionized water…………………………...88
Figure 3.14. Representative SEM images of common morphologies observed in crystal
growth trials using a growth solution of 1 mM CaCl2 + 0.5 mM aspartic acid. (a)
500X and (b) 1000X Trial 07_18_201: vaterite. (c) 500X and (d) 1500X Trial
07_19_2011: vaterite. (e) 500X and (f) 2000X Trial 07_21_2011: vaterite. All trials
used 5 g ammonium carbonate + 20 mL deionized water…………………………89
ix
Chapter 1
Introduction
1
1.1 Introduction
Carbonate minerals are common, highly reactive components of the Earth’s surface.
Among the most abundant mineral families are the calcium carbonates, which form when
dissolved carbonate reacts with calcium ions to precipitate crystalline solids (Morse 1983).
Calcium carbonates occur in many environments, and thus they play important roles in the
regulation of natural systems. In soils and natural waters, calcium carbonates buffer pH
(Pankow 1991) and can immobilize toxic metals (Cai et al. 2010). On a larger scale, reactions
involving calcium carbonate in both terrestrial and marine systems partially regulate
atmospheric CO2 concentration (Dickinson et al. 2002).
However, the increasing
concentration of atmospheric carbon dioxide has resulted in increased acidity in natural
waters, and thus the inhibition of calcium carbonate formation and promotion of calcium
carbonate dissolution.
As the most abundant biogenic mineral, calcium carbonate precipitation is key to the
ecology of many ecosystems.
Its precipitation is frequently templated by molecules
produced by microorganisms. The diversity of microorganisms leads to an extensive number
of biomolecules present in the Earth’s systems, and thus an extensive range of possible
effects on CaCO3 formation.
Low molecular mass carboxylic acids are very common
biomolecules and have a relatively simple structure, thus making them an ideal family of
model compounds for experiments that probe the mechanisms of biomineralization. The
connection between biomolecule structure and the resulting morphology of calcium
carbonate may provide insight into CaCO3 precipitation. This topic is becoming increasingly
important as natural systems which rely on CaCO3 precipitation are put under stress through
increased acidity.
2
The work presented in this thesis utilizes a novel chemical reactor to quantitatively
and qualitatively examine the effect of calcium concentration and biomolecule structure on
calcium carbonate precipitation reactions. In this chapter, we discuss the environmental
importance of carbonates and provide an introduction to the chemistry of the carbonate
system. In chapter two, we describe the utilization of this reactor, and the changes in aqueous
chemistry that can be measured in it; additionally, the results are applied to a kinetic model to
relate precipitation rate to thermodynamic driving force. In chapter three, we discuss the
phase and morphology of precipitates from the reactor. Chapter four provides concluding
remarks and the outlook for future applications of this work.
1.2 Global Environmental Importance of Carbonates
1.2.1 The Global Carbon Cycle and Anthropogenic Influence
Anthropogenic emissions have greatly influenced the chemistry of the atmosphere.
The concentration of carbon dioxide in the atmosphere has increased from approximately 280
ppmv in 1850 to nearly 390 ppmv today (Figure 1.1), with a current rate of increase of 1.7
ppmv per year. Emissions of carbon dioxide from fuel combustion increased from 2.3 Pg C
yr-1 in the 1950s to 7.3 Pg C yr-1 in 2003. Atmospheric carbon dioxide concentration has
risen approximately in proportion to the increasing rate of carbon dioxide emitted from fossil
fuel combustion (Keeling et al. 2011).
Due to its high reactivity and abundance at the Earth’s surface calcium carbonate is a
major part of the carbon cycle, both on land and in marine settings. The global pool of carbon
is 1023 g, and 6.5x1022 g C are contained in carbonate rocks (Schlesinger 1997). The three
3
reservoirs of carbon are atmospheric, oceanic, and terrestrial. The terrestrial pool includes
soil, plants, organisms, and geological reservoirs such as rock carbonates and fossil organic
carbon. Each system has a complex carbon cycle of its own, but overall carbon exchanges
occur between the atmosphere and the land, and the atmosphere and the ocean. Marine
sediments are gradually incorporated into the geological reservoir, and land carbon is quickly
transported to the oceans via river transport and weathering. An overview of the carbon
cycle can be seen in Figure 1.2 (Prentice et al. 2001). The formation of the geological
reservoir occurs over the course of several millennia, although these reservoirs are quickly
transformed to atmospheric carbon via fossil fuel burning. The formation and dissolution of
CaCO3 happens on a shorter time interval, as biogenic production of calcium carbonate by
marine organisms causes nearly 50% of dissolved inorganic carbon from rivers to be returned
to the atmospheric pool, while the other half is buried as deep-sea carbonate sediments
(Prentice et al. 2001).
Atmospheric carbon dioxide is considered a major factor in driving global climate
change. The changes in the composition of the troposphere are connected to changes in
radiative forcing by the so called Greenhouse Effect, by which gases such as CO2 transmit
short wave radiation reflected from the surface of the planet into space but do not allow
longer wavelength infrared radiation to escape. The capture of long wave radiation traps heat
that otherwise would be lost. The putative result of this process is an increasing trend in the
global mean temperature (Berkeley Earth Team, 2011), and several models predict a
continued increase of 2–4°C by the end of the next century, as well as a wide variety of
concomitant changes to the climate and the Earth system (IPCC, 2007).
4
1.2.2 Ocean Acidification
One of the major motivations for this work is developing a better understanding of the
impact of ocean acidification on carbonate precipitation and dissolution. Carbon dioxide
reacts with water to form carbonic acid, which upon dissociation produces hydrogen ions and
thus an increase in acidity.
Approximately 30 to 40% of atmospheric carbon dioxide
becomes dissolved in the world’s oceans. Subsequently, ocean pH has decreased by 0.1 pH
units since prior to the Industrial Revolution, resulting in a phenomenon known as ocean
acidification (Zalasiewicz et al. 2008; Fabricius et al. 2011). This downward trend in pH has
resulted in a decrease in the saturation state of carbonate minerals in seawater and an increase
in rates of calcium carbonate dissolution, which acts as a mechanism to buffer increased
[H+]. In order for the carbonate species to be present at any meaningful concentration, the
pH of the system must be greater than 8, as can be seen in Figure 1.3. Decreased rates of
calcification have already been seen in organisms that build skeletons with aragonite (Morse
et al. 2007). Additionally, terrestrial ecosystems take up approximately 20% of
anthropogenic carbon dioxide emissions, and are also subject to the same acidity increasing
reactions that occur when carbon dioxide reacts with water (Mitchell et al. 2010).
The following equation shows that decreasing pH comes from the release of H+ ions
during the dissociation of carbonic acid created by the reaction of carbon dioxide and water:
CO2 + H2O ↔ H2CO30 ↔ HCO3- + H +↔ CO32- + H +
(1.1)
The increase of H+ in turn shifts equation 1.1 towards the left, increasing more acidic species,
H2CO2 and HCO3-, and decreasing the concentration of carbonate ions (Andersson et al.
2007).
Calcium carbonate minerals are formed when aqueous species of Ca2+ and CO32-
react to form solid phase CaCO3:
5
Ca2+(aq) + CO32-(aq) ↔ CaCO3(s)
Log KSP = -3.22
(1.2)
where KSP indicates the solubility product for the given calcium carbonate polymorph
(Busenberg and Plummer 1986). Aragonite and calcite are both present in marine systems,
although more often aragonite is the primary polymorph (De Choudens-Sanchez and
Gonzalez 2009). The saturation state (Ω) for both aragonite and calcite are both dependent on
the concentration of CO32- ions available in seawater (Feely et al. 2004):
 arg 
cal 
[Ca 2+ ][CO32 ]
KSP arg
Log K*SParg = -8.3 (1.3)
[Ca 2+ ][CO32 ]
KSPcal
Log K*SPcal = -8.5 (1.4)
Precipitation only occurs at a saturation state greater than one (supersaturation), so if the
solution is less than one (undersaturation) organisms will not be able to induce precipitation
of either aragonite or calcite (Orr et al. 2005). Additionally, a saturation state of less than one
indicates that net dissolution may occur, indicating that existing calcium carbonate minerals
may be subject to dissociation (Tyrrell 2007).
Increased pCO2 and decreasing ocean pH will affect the ability of the numerous
organisms that utilize or produce CaCO3 to survive and compete with other species. These
changes in the prevalent species will have profound effects on the ecology of many marine
systems (Doney 2009). Many of these organisms incorporate biomolecules during the
biomineralization of calcium carbonate (Hollingsworth 2009). The results from the proposed
6
research will lead to a greater understanding of the role of biomolecule structure on the
morphology and formation kinetics of calcium carbonates.
1.3 Terrestrial and Aquatic Carbonates and Their Reactivity
1.3.1 Alkalinity
Dissolved carbonates buffer pH in most natural waters and physiological systems.
Alkalinity refers to the acid neutralization capacity of a system, whereas acidity refers to the
base neutralization capacity. Therefore, alkalinity can be defined as the amount of acid
required to bring the pH to a specific end point. The equation for carbonate alkalinity is:
Alkalinity = CB – CA = [HCO3-] + 2[CO32-] + [OH-] – [H+]
(1.5)
Where CB = concentration of the base and CA = concentration of the acid (Pankow 1991).
Carbonate concentration is the major source of alkalinity in most natural waters and,
due to its relative abundance and reactivity, and calcium carbonate is therefore a major factor
in controlling alkalinity in natural waters and engineered systems. Additionally, the
carbonate supplied by the dissolution of CaCO3 can be used to raise pH.
1.3.2 Liming
Calcium carbonate is often used to neutralize soil acidity in a practice known as
liming. Calcium carbonate reacts with water to produce a hydroxide ion which in turn can
bind to H+ or Al3+, the cations responsible for exchangeable acidity in soils. The general
equation for liming is:
7
CaCO3 + H2O↔ Ca2+ + HCO3- + OH-
(1.6)
Typically, North Carolina soils are naturally acidic and thus require liming to produce
productive crops. In acidic soils that are high in aluminum, as is common in North Carolina,
liming typically results in the overall reaction:
Al-soil + 3CaCO3 + 3H2O ↔ 3Ca-soil + 2Al(OH)3 + 3CO2
(1.7)
The resulting form of aluminum, Al(OH)3, is highly insoluble and generally does not
contribute to phytotoxicity.
Ground dolomitc rock is the most commonly used liming
material in North Carolina, but pure calcium carbonate is considered the standard for liming
materials and is used in calculations for determining how much liming agent is needed
(Crozier and Hardy 2003).
1.3.3 Sequestration of Metals
Calcium carbonate has been shown to be a useful tool in the remediation of
potentially toxic metals, either by binding to and thus immobilizing the metal or by raising
pH and changing the speciation to a non-toxic form. Potentially toxic metals are a concern
when they are in a mobile and soluble state, and can therefore be taken up by plants or
transported to drinking water. The remediation technique of chemical stabilization uses a
benign material to reduce the solubility of the chemical of concern, and thus render it nontoxic. Other soil remediation techniques include physically removing the contaminated soil,
mixing the soil with clean soil or covering the soil with clean soil, using acids to leach the
8
contaminants out of the soil, and using plants to extract the chemical of concern. Chemical
stabilization is deemed to be the most cost-effective of these techniques.
Chemical
stabilization can occur via multiple pathways including precipitation, adsorption, and
complexation (Chen et al. 2000).
Calcite and aragonite have been utilized as a means of removing lead from aqueous
solutions for remediation purposes (Godelitas et al., 2003). Results show that both
polymorphs are successful at removing Pb, and that the dissolution of calcium carbonate
occurs simultaneously with the heterogeneous nucleation of lead carbonates.
These
precipitants have been identified as cerussite (PbCO3) and hydrocerussite (Pb3(CO3)2(OH)2)
crystals. Additionally, Pb was found to both adsorb to CaCO3 crystal surfaces and absorb
into CaCO3 surface layers. Calcium carbonate was also found to reduce the extractability of
Cd and Pb in soils and to reduce the uptake of both metals by wheat shoots (Chen et al.
2000).
In addition to lead and cadmium, calcium carbonate has been proven to be a
successful tool for remediation of soils containing copper (Achal et al. 2011) as well as
chromium and nickel (Cai et al. 2010). Carbonate minerals can raise pH and thus mitigate
aluminum and manganese toxicity (Geebelen et al. 2002), as discussed in section 1.3.2.
1.4 Carbonate Crystal Chemistry
Carbonates are highly abundant and reactive (Reeder 1983). One of the reasons for
their abundance is due to the nature of carbon dioxide. Atmospheric carbon dioxide may
diffuse into aqueous systems and form minerals with a wide variety of divalent metal ions
that may be present in the system (Pankow 1991). Calcium carbonate in particular has a
9
diverse array of polymorphs which are present in nature (Reeder 1983). The reactivity of this
particular carbonate and the transition between polymorphs makes calcium carbonate a very
interesting chemical system.
There are eight known polymorphs of calcium carbonate. Calcite, aragonite, and
vaterite are the three crystalline forms comprised of pure calcium carbonate (Weiner and
Dove 2003). Calcite is the thermodynamically stable CaCO3 structure in laboratory
conditions; aragonite and vaterite are metastable polymorphs that also commonly form
(Wray and Daniels 1957). Calcite and aragonite are the most common forms of calcium
carbonate and can be coprecipitated in the same solution. In natural systems, calcite is the
more stable form under surface conditions, although aragonite is often the dominant
polymorph in marine settings (De Choudens-Sanchez and Gonzalez 2009).
Both aragonite and calcite are significant in seawater precipitation and dissolution
reactions. Calcite has hexagonal structure and a unit cell comprised of two CaCO3 units
arranged in an acute rhombohedron. This morphology is common in calcite crystals.
Aragonite has orthorhombic structure that has a higher coordination number for Ca2+ ions
(Morse 1983). Consequently, aragonite is more likely to isomorphically substitute with
cations that have larger ionic radii than Ca2+ , such as Sr2+ and Ba2+, while calcite often
contain cations with smaller ionic radii, such as Mg2+ (Morse et al. 2007). Aragonite
frequently adopts a needle-like morphology in its particles.
Under normal conditions, vaterite is the least stable calcium carbonate polymorph.
However, the presence of organic molecules can help to stabilize vaterite. Recent work has
10
discovered that inorganic molecules can stabilize and even promote vaterite precipitation.
The diffusion of ammonia gas into aqueous calcium solutions can result in the stabilization of
the vaterite polymorph (Becker and Hu 2009). Vaterite, which commonly forms spheroidal
particles, may also be a precursor to a more stable form, much like amorphous calcium
carbonate.
Amorphous calcium carbonate is the least thermodynamically stable form of CaCO3,
and is therefore the most soluble, approximately 120 times more soluble than calcite.
According to Ostwald’s step rule, amorphous calcium carbonate will be the first form to
precipitate from a supersaturated solution and will rapidly crystallize into one of the more
stable polymophs, if it is not stabilized by another element. Amorphous calcium carbonate is
also the smallest of the polymorphs, averaging spherules of 40 to 120 nm compared to a
normal CaCO3 crystal of 1 to 10 µm. Only recently has the synthesis and stabilization of
amorphous calcium carbonate been possible, but usually these experiments require the use of
toxic chemicals or only last a few days (Meiron et al. 2011). Certain organisms have been
found to produce molecules, such as chitin, that temporarily stabilize amorphous calcium
carbonate. The amorphous form is either used as an antecedent for a more stable form, or
can be used to help transport calcium into the bloodstream (Meiron et al., 2011).
1.5 Aqueous Chemistry of Carbonates
Calcium carbonate crystals are formed when aqueous species of Ca2+ and CO32- react to
form solid phase CaCO3 (equation 1.2). This reaction is highly pH dependent, given the
11
distribution of carbonate species over pH range. The deprotonation of carbonic acid occurs at
log K0 =- 6.36 and the deprotonation of bicarbonate occurs at log K0 = -10.33 (Figure 1.3).
Therefore, carbonate species and thus calcium carbonate crystals will start to be present at
approximately pH 8.5 and will become the dominant species at pH 10.33.
CO2(g) + H2O ↔ H2CO30
log K0 = -1.46 (1.8)
H2CO30 ↔ H + + HCO3-
log K0 = -6.36 (1.9)
HCO3- ↔ H + + CO32-
log K0 = -10.33
(1.10)
As shown in equations 1.2 – 1.4, precipitation of CaCO3 is driven by dissolved carbonate,
which is a significant species only at relatively high pH (Figure 1.3).
Lindsay (1979) used equations 1.8-1.10 to derive an equation relating the activity of
CO32- to pH and the activity of CO2(g) :
log CO32- = -18.15 + 2pH + log CO2(g)
(1.11)
This equation shows that carbonate activity increases with increasing pH and increasing
activity of CO2(g). Single-ion activity is the product of the ion concentration multiplied by
the ion activity coefficient. When ionic strength is assumed to be near zero, then the activity
coefficient is assumed to be one and the activity will be equal to the concentration.
1.5.1 Dissolution
Dissolution of calcium carbonate can be a complicated process and cannot be
12
described as simply crystal growth equations with a negative sign, especially in regards to the
ocean where the chemical system is quite complicated.
Generally, in undersaturated
solution, dissolution may proceed by proton or water promoted pathways:
MCO3(s) + H+  M2+ + HCO3-
(1.12)
MCO3(s) + H2O  M2+ + HCO3- + OH-
(1.3)
Equation 1.12 is the important reaction pathway at acidic pH values while the pHindependent water-promoted pathway (equation 1.12) dominates for circumneutral to
alkaline conditions (Busenberg and Plummer, 1986). Dissolution can be transport-controlled
or surface-controlled, and the determination of which will control dissolution is dependent on
factors such as pH and saturation state (Rickard 1983; Morse et al. 2007).
1.5.2 Growth
Saturation state depends on the concentration of aqueous calcium and carbonate in the
system (equations 1.3 and 1.4). Undersaturation indicates a low product for [Ca2+][CO32-].
Therefore, in a state of undersaturation any solids present will be likely to dissolve and no
solids will form (equation 1.2 moves to the left). Conversely, supersaturated solutions tend to
cause crystal growth and, if solids are absent, they may nucleate from solution (equation 1.2
moves to the right) (Pankow 1991). If the rate of crystallization is slow, then the atoms will
form a few crystals, generally with the same orientation.
13
However, if the rate of
crystallization is fast, then many crystals will form and the resulting crystals will be small
and randomly oriented.
Research using atomic force microscopy has found that calcite growth occurs by
either step flow (Teng et al. 2000), by growth at areas of crystal imperfections (Hillner et al.
1992), or by two-dimensional surface nucleation (Dove and Hochella 1993). Studies at the
macroscopic scale have determined that the kinetics of calcite formation depend on
parameters such as supersaturation state, pH, pCO2, ionic strength, and temperature.
However, research has been unable to determine a conclusive relationship between growth
mechanisms and supersaturation (Teng et al. 2000).
1.5.3 Effects of biomolecules
Calcium carbonate phases are the most abundant biominerals (Weiner and Dove
2003). Biominerals are inorganic solids formed by living organisms and are created for a
specialized purpose, such as providing a shell or skeleton for an organism. These minerals
must precipitate to a certain size and shape, and organisms use biomolecules to ensure that
the mineral will grow to the correct specifications. Proteins and polysaccharides may serve as
templating agents during the formation of biominerals, affecting phase or morphology
(Mukkamala et al. 2006). Additives involved in CaCO3 precipitation may work in many
different ways to affect reaction kinetics and ultimately nucleation and growth (Henderson et
al. 2008).
Biomolecules can act as a template through stereochemical recognition and
14
electrostatic matching (Sugawara et al. 2003). These molecules can also inhibit surface-step
growth by binding to surface-step edges and changing the free energies of the step-edge, and
thus modifying crystal shape. The modifiers will bind to the site that provides the best
geometric and chemical fit (Orme et al. 2001). Although often times the exact role of the
biomolecule is unknown, it is accepted that the anionic groups of the biomolecules interact
with the metal ions (Sugawara et al. 2003).
Biomolecules are often complex structures, such as proteins or polysaccharides, and
can be difficult to isolate or to define chemically. The proposed experiment aims to
determine the influence of biomolecule structure on calcium carbonate formation, and
therefore the biomolecule involved must be small enough to be easily manipulated, and must
have a known chemical structure. Carboxyls are acidic functional groups that have been
shown to be structurally important in the active regions of many biomolecules. Carboxylic
acids follow the chemical formula R-C=O(-OH), where R denotes an aliphatic compound
(Sparks 2003), and the deprotonated forms of these acids are known as carboxylates. A list of
biomolecules utilized to this experiment includes aspartic acid, citric acid, and succinic acid
(Figure 1.4). These molecules were chosen because of the biological prevalence and
structural relationships to one another.
1.6 Motivation and Goals
Despite the fact that calcium carbonate precipitation and the importance of certain
biomolecules pertaining to formation processes has been extensively studied, critical
15
uncertainties exist about how the structure of small biomoleclues affects the formation of
carbonates. A better understanding of these processes may lead to improved predictive
models of aqueous carbonate chemistry; which are imperative at a time when global carbon
cycling faces drastic changes due to increasing atmospheric carbon dioxide concentrations.
The overall goal of this study is to better understand the effect of biomolecule
structure on the critical supersaturation necessary for and the rate of the precipitation of
calcium carbonate crystals, as well as the resulting phase and morphology of products.
Specific objectives are: (1) to measure the critical supersaturation required for nucleation as
well as the growth rate of calcium carbonate crystals, and to relate these parameters to
solution condition by a chemical model and (2) to determine the effects of biomolecule
structure on the mechanisms and products of calcium carbonate growth.
16
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Kocuria flava CR1, based on microbially induced calcite precipitation. Ecological
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under rising pCO2 and ocean acidification: Observations from Devil’s Hole, Bermuda.
Aquatic Geochemistry. 13, 237-264.
Becker, U. and Hu, Q. (2009). Controlled growth of different calcium carbonate polymorphs
as induced by the presence of dissolved molecules and mineral surfaces. Goldschmidt
Conference. June 21 -26, Davos, Switzerland.
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growth kinetics of calcite and aragonite. U.S. Geol. Surv. Bull. 1578, 139-168.
Cai, G.B., Zhao, G.H., Wang, X.K., Yu, S.H. (2010). Synthesis of polyacrylic acid stabilized
amorphous calcium carbonate nanoparticles and their application for removal of toxic heavy
metal ions in water. Journal of Physical Chemistry. 114, 12948-12954.
Chen, Z.S., Lee, G.J., and Liu, J.C. (2000). The effects of chemical remediation treatments of
the extractability and speciation of cadmium and lead in contaminated soils. Chemosphere.
41, 235-242.
Crozier, C. and Hardy, D.H. (2003). Soil facts: Soil acidity and liming for agricultural soils.
NC State University and North Carolina A&T State University Cooperative Extension.
De Choudens-Sanchez, V. and Gonzalez, L. A. (2009). Calcite and aragonite precipitation
undercontrolled instantaneous supersaturation: Elucidating the role of CaCO3 saturation state
and Mg/Ca ratio on calcium carbonate polymorphism. Journal of Sedimentary Research. 79,
363-376.
Dickinson, S.R., Henderson, G.E., McGrath, K.M. (2002). Controlling the kinetic versus
thermodynamic crystallization of calcium carbonate. Journal of Crystal Growth. 244, 369378.
Doney, S.C., W.M. Balch, V.J. Fabry, and R.A. Feely. 2009. Ocean acidification: A critical
emerging problem for the ocean sciences. Oceanogr. 22:16-25.
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Dove P.M. and Hochella M.F., Jr. (1993) Calcite precipitation mechanisms and inhibition by
orthophosphate: In situ observations by Scanning Force Microscopy. Geochimica et
Cosmochimica Acta. 57, 705–714.
Essington, Michael E. (2004). Soil and water chemistry: an integrative approach.
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CRC
Fabricius, K.E., Langdon, C., Uthicke, S., Humphrey, C., Noonan, S., De’ath, G., Okazaki,
R., Muehllehner, N., Glas, M.S., Lough, J.M. (2011). Losers and winners in coral reefs
acclimatized to elevated carbon dioxide concentrations. Nature Climate Change. 1, 165-169.
Feely, R.A., Sabine, C.L., Lee, K., Berelson, W., Kleypas, J., Fabry, V.J., Millero, F.J.
(2004). Impact of anthropogenic CO2 on the CaCO3 system in the oceans. Science. 305, 362366.
Geebelen, W., Vangronsveld, J., Adriano, D.C., Carleer, R., Clijsters, H. (2002).
Amendment-induced immobilization of lead in a lead-spiked soil: Evidence from
phytotoxicity studies. Water, Air, and Soil Pollution. 140, 261-277.
Godelitas, A., Astilleros, J.M., Hallam, K., Harissopoulos, S., and Putnis, A. (2003).
Interaction of calcium carbonates with lead in aqueous solutions. Environmental Science &
Technology. 37, 3351-3360.
Henderson, G.E., Murray, B.J., and McGrath, K.M. (2008). Controlled variation of calcite
morphology using simple carboxylic acids. Journal of Crystal Growth. 310, 4190-4198.
Hillner P.E., Manne S., Gratz A.J., and Hansma P.K. (1992) AFM images of dissolution and
growth on a calcite crystal. Geology. 20, 359-362.
Hollingsworth, M.D. (2009). Calcite biocomposites up close. Science. 326, 1194-1195.
IPCC. 2007. Summary for Policymakers, In S. Solomon, et al., eds. Climate Change 2007:
The Physical Science Basis. Contribution of Working Group I to the Fourth Assessment
Report of the Intergovernmental Panel on Climate Change. Cambridge University Press,
Cambridge, United Kingdom and New York, NY, USA.
Keeling, C.D., Piper, S.C., Whorf, T.P., Keeling, R.F. (2011). Evolution of natural and
anthropogenic fluxes of atmospheric CO2 from 1957 to 2003. Tellus. 63B, 1-22.
Klein, C., and Hurlbut, C.S. (1993). Manual of Mineralogy. John Wiley & Sons, Inc., New
York, NY.
Meiron, O.E., Bar-David, E., Aflalo, E.D., Shechter, A., Stepensky, D.,Berman, A., and Sagi,
A. (2011). Solubility and bioavailability of stabilized amorphous calcium carbonate. Journal
of Bone and Mineral Research. 26, 364-372.
18
Mitchell, M.J., Jensen, O.E., Cliffe, K.A., and Maroto-Valer, M.M. (2010). A model of
carbon dioxide dissolution and mineral carbonation kinetics. Proceedings of the Royal
Society. 466, 1265-1290.
Morse, J. W. (1983). The kinetics of calcium carbonate dissolution and precipitation. In
Carbonates: Mineralogy and Chemistry, Reeder, R. J., Ed. Mineralogical Society of
America: Washington D.C.
Morse, J.W., Arvidson, R.S., and Luttge, A. (2007). Calcium carbonate formation and
dissolution. Chemical Reviews. 107, 342-381.
Mukkamala, S.B., Anson, C.E., and Powell, A.K. (2006). Modelling calcium carbonate
biomineralisation processes. Journal of Inorganic Biochemistry. 100, 1128-1138.
Oh, N.H. and Richter, D.D. (2004). Soil acidification induced by elevated atmospheric CO2.
Global Change Biology. 10, 1936-1946.
Orme, C.A., Noy, A., Wierzbicki, A., McBride, M.T., Grantham, M., Teng, H.H., Dove,
P.M., and DeYoreo, J.J. (2001). Formation of chiral morphologies through selective binding
of amino acids to calcite surface steps. Nature. 411, 775-779.
Orr, J.C., et al. (2005). Anthropogenic ocean acidification over the twenty-first
century and its impact on calcifying organisms. Nature. 437, 681-686.
Pankow, J.F. (1991). Aquatic Chemistry Concepts. Lewis Publishers, Boca Raton, FL.
Prentice, I.C., Farquhar, G.D., Fasham, M.J.R., Goulden, M.L., Heimann, M., Jaramillo, V.J.,
Kheshgi, H.S., Le Quere, C., Scholes, R.J., and Wallace, D.W.R. (2001). Climate Change
2001: The Scientific Basis. Climate Change 2001, IPPC Thrid Assessment Report.
Plummer, L.N. and Busenberg, E. (1982). The solubilities of calcite, aragonite and vaterite in
CO2-H2O solutions between 0 and 90°C, and an evaluation of the aqueous model for the
system CaCO3-CO2-H2O. Geochimica et Cosmochimica Acta. 46, 1011-1040.
Quere, C.L., Aumont, O., Bopp, L., Bousquet, P., Ciais, P., Francey, R., Heimann, M.,
Keeling, C.D., Keeling, R.F., Kheshgi, H., Peylin, P., Piper, S.C., Prentice, I.C., Rayner, P.J.
(2003). Two decades of ocean CO2 sink and variability. Tellus. 55B, 649-656.
Reeder, R.J. (1983). Crystal chemistry of the rhombohedral calcites. In Carbonates:
Mineralogy and Chemistry, Reeder, R. J., Ed. Mineralogical Society of America:
Washington D.C.
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19
Rumpel, C., Kogel-Knabner, I. (2011). Deep soil organic matter—a key but poorly
understood component of terrestrial C cycle. Plant Soil. 338, 143-158.
Schlesinger, W.H. (1997). Biogeochemistry: An analysis of global change. Academic Press,
San Diego, CA.
Sparks, D.L. (2003). Environmental Soil Chemistry. Academic Press: San Diego, CA.
Sugawara, T., Suwa, Y., Ohkawa, K., and Yamamoto, H. (2003). Growth of calcite with
chiral phosphoserine copolypeptides. Macromolecular Rapid Communications. 24, 847-851.
Teng, H.H., Dove, P.M., and De Yoreo, J.J. (2000). Kinetics of calcite growth: Surface
processes and relationships to macroscopic rate laws. Geochimica et Cosmochimica Acta. 64,
2255-2266.
Tyrrell, T. (2007). Calcium carbonate cycling in future oceans and its influence on future
climates. Journal of Plankton Research. 30, 141-156.
Weiner, S. and P.M. Dove (2003) An Overview of Biomineralization and the Problem of the
Vital Effect. In Biomineralization. Eds. P.M. Dove, S. Weiner and J.J. De Yoreo.
Mineralogical Society of America, Washington, D.C., v. 54, 1-31.
Wray, J.L. and Daniels, F. (1957). Precipitation of calcite and aragonite. Journal of the
American Chemical Society. 79, 2031-29034.
Zalasiewicz, J., Williams, M., Smith, A., Barry, T.L., Coe, A.L., Bown, P.R., Brenchley, P.,
Cantrill, D., Gale, A., Gibbard, P., Gregory, F.J., Hounslow, M.W., Kerr, A.C., Pearson, P.,
Knox, R., Powell, J., Waters, J.M., Oates, M., Rawson, P., and Stone, P. (2008). Are we now
living in the Anthropocene? Geological Society of America Today. 2, 4-8.
20
Figure 1.1. The CO2 concentration, in ppm, versus time, from observations for Arctic ice
floes, 1957–1958, and Point Barrow, Alaska, 1961–1968 and 1974–2003, shown as monthly
averages (dots) and a spline fit combined with seasonal harmonics increasing in amplitude at
an assumed constant rate (smooth curve) (Keeling et al. 2011).
21
Figure 1.2. The main components of the natural carbon cycle. Thick lines indicate gross
primary production and respiration from land and exchange between the ocean and
atmosphere, which are the most important CO2 fluxes that take place over a short-term
timescale. Thin lines represent other natural fluxes and dashed lines indicate carbon flux as
CaCO3, which are significant over longer time scales. All units are in Pg C yr-1 and refer to
data collected during the 1980s (Prentice et al. 2001).
22
Figure 1.3. The effect of pH on the speciation of carbonates. Created with Visual MINTEQ
using 0.04 M H2CO3.
A
B
C
Figure 1.4. Molecular structures of carboxyl-containing biomolecules including (A) aspartic
acid (B) citric acid (C) succinic acid.
23
Chapter 2
Experimental Design and Kinetic Modeling for Calcium Carbonate
Precipitation
24
2.1 Introduction
Biomineralization is defined as the process by which organisms form minerals from
dissolved species (Weiner and Dove 2003). Biominerals are a combination of organic and
inorganic molecules, and this combination provides these minerals with strength-to-weight
properties that allow them to be more versatile than minerals comprised solely of inorganic
materials. These properties can allow microbes to use the minerals for structural support,
filtration, or light collection (Dove 2010).
Calcium carbonate biomineralization is an
important part of many geological processes and has implications for engineering and
environmental remediation (Li et al. 2010). Biogenic calcium carbonate can have a diverse
array of structures, illustrating the complexity of biominerization processes (Teng et al.
1999).
The same organism can even produce different morphological structures.
For
example, the marine organism Globigerina pachyderma has a shell made of calcium
carbonate that coils in a counter-clockwise direction in the Arctic and Antarctic, but coils in
the clockwise direction in temperate and tropical areas (Addadi and Weiner 2001). Often,
these differences are caused by biomolecules present during growth. Biomineral morphology
and polymorphism is influenced by inorganic and organic compounds that can act as either
impurities or growth templates (Teng et al. 1999). For example, when non-chiral calcite was
grown in the presence of the chiral aspartic acid, the resulting crystals take on the chirality of
the aspartic acid (Orme et al. 2001). Additionally, the surface of a crystal can be chiral, even
if the bulk of the crystal structure is not (Addadi and Weiner 2001). In addition to phase and
morphology, biomolecules can increase the rate of formation by lowering the energy barrier
of the crystallization (Elhadj et al. 2006). The terrace-ledge-kink model holds that crystals
25
grow via step propagation that mostly occurs at kink sites, where the step edge is not
complete (De Yoreo et al. 2009). The more recently developed stereochemical recognition
model suggests that peptides and proteins can bind to otherwise unstable faces by lowering
the surface energy of the crystal lattice, and therefore biomolecules can use this as a means of
altering crystal shape (De Yoreo and Dove 2004).
Many molecules used by organisms to facilitate or control biomineralization contain
complex polymeric materials that are difficult to study in the laboratory.
However,
researchers have found that simple amino acids and polypeptides can cause the same
morphological changes as complex modifiers (De Yoreo and Dove 2004).
Many
biomolecules are thought to contain carboxylate moieties in the precipitation active regions
of the molecule (Addadi and Weiner 1985). Carboxylates may change calcium carbonate
precipitation by blocking carbonate sites (Mukkamala et al. 2006) or acting as growth
templates (Teng et al. 1999). In this chapter, we describe the effect of small carboxylate
bearing biomolecules on the rate of calcium carbonate precipitation. We also describe the
performance of a novel reactor designed to allow for in-situ monitoring of chemical
conditions within the reactor during mineral formation.
2.2 Materials and Methods
2.2.1 Growth Reactor
Much research has been conducted to investigate calcium carbonate precipitation and
dissolution reactions. However, previous research on CaCO3 biominerization often has
26
focused solely on reactants at the start of the experiment and the products at the conclusion of
the experiment (Dickinson et al. 2002, Ruiz-Agudo et al. 2011) or has focused on a singular
aspect such as thermodynamics (Wolf et al. 1996).
These studies have thus neglected
quantifying the kinetics of the reaction. In order to better understand this chemical system
and the reactions taking place within it, more data must be collected on solution parameters
(such as [Ca2+], pH, and pCO2) throughout the experiment. These data can then be used in
kinetic modeling to determine rate coefficients. In order to make such measurements, a new
approach must be used for the completion of calcium carbonate precipitation experiments.
The development of a growth chamber which would allow for access to the solution during
the experiment was an integral part of examining this problem.
A novel reactor (Figure 2.1, 2.2) has been utilized to study CaCO3 formation at the
bench scale. The reactor was made of polycarbonate to minimize the interaction between the
reactor material and the chemical solutions enclosed within. The reactor contains an upper
chamber to house a volatile solid or solution, which serves as the carbonate source and pH
driver. The top of the chamber has a hole through which a pH electrode, a calcium electrode,
and a UV-visible spectrophotometer are allowed access to the experimental system. Once the
electrodes are in place the hole is filled with a rubber stopper. The outside of the rubber
stopper is covered in caulk to ensure that this point does not allow access to the atmosphere.
There is also a port that enables carbon dioxide samples to be continuously drawn out and a
similar port that allows the samples that have been analyzed (but not chemically changed) to
be replaced back into the experiment system.
27
2.2.2 Experimental Setup
All experiments were carried out inside of the reactor, with solution samples and solid
precipitates were analyzed after completion. Calcium carbonate crystals have been
precipitated by the sublimation of CO2(g) into CaCl2(aq). Solutions of calcium chloride (Fisher
Scientific, dihydrate) at concentrations of 1 mM, 5 mM, and 10 mM CaCl2 were placed in the
bottom chamber of the reactor. Experiments involving biomolecules used 1 mM CaCl2 and
0.5 mM of aspartic acid (Acros Organics), citric acid (Acros Organics), and succinic acid
(Acros Organics). The pH of each solution was titrated to pH = 3 via the addition of 1 M
HCl (Acros), which ensured that no calcium carbonate solids formed before the start of the
experiment. A sample of 5 g ammonium carbonate (Fisher Scientific) plus 20 mL deionized
water was placed in the upper section of the chamber. As NH3(g) and CO2(g) volatize, a fan
circulates gas from the upper chamber to the aqueous solution below, and a stir bar present in
the lower chamber was used to ensure the solution was constantly and uniformly mixed. The
carbon dioxide gas served as a source of carbonate and the ammonia gas increased the pH of
the solution, thus promoting calcium carbonate precipitation. By varying the initial [Ca2+]
concentration, the ratio of [Ca2+]/[CO32-] was altered and thus the kinetics of the reaction
were shifted.
The chamber is equipped with a pH meter (Fisher Scientific Accumet Excel XL25), a
calcium ion selective electrode (ISE; Vernier), a CO2 gas analyzer (Licor LI-840A
CO2/H2O), and a UV-visible spectrophotometer (Ocean Optics USB4000) with a
transmission dip probe (Ocean Optics TP300-UV-VIS) and halogen light source (Ocean
Optics LS-1-LL). Visible spectroscopy has been used to determine the time at which
28
nucleation occurs by measuring light absorbance and turbidity within the solution. The
spectrophotometer (Ocean Optics Jaz Spectrophotometer) is set up with a probe inside the
growth chamber to take data measurements in-situ. However, this method was found to be
unreliable and is not discussed further. The calcium ISE has a gel-filled membrane with a
quantification range of 0.2 to 40,000 mg Ca2+ L-1. To protect the reactor and facilitate
collection of precipitates, the bottom piece of the chamber is covered with a circular
polycarbonate film that has been cut to match the area. The polycarbonate film is supplied
by Grainger and is clear with a thickness of 0.0127 cm. A scanning electron microscope was
used to image the crystals collected on the polycarbonate film at the end of the experiment
and X-ray diffraction was used to determine the phase of the solid precipitates, as described
in Chapter 3.
Gas samples were continuously analyzed from a gas port located at the top of the
chamber and atmospheric CO2 concentrations (PCO2) are measured using a carbon dioxide
analyzer (Licor). Gas samples were moved by a diaphragm pump (Hargraves) to the CO2
analyzer, ensuring that sample flow remained constant. A calcium ISE was present in the
growth chamber during experiments and used to measure the concentration of Ca2+ in
solution. Data on Ca2+ loss from solution has been used to help determine the rate of
crystallization. The concentration of calcium ions in solution has also been determined
through atomic absorption spectroscopy analysis of solution samples taken at intervals
throughout the experiment to confirm in-situ measurements. The chemical equilibrium
modeling software MINEQL+ has been used to calculate CaCO3 speciation in selected
experiments. The concentration of biomolecule was not measured during the experiment.
29
Initial experiments included pure calcium chloride solutions, with calcium
concentrations of 1 mM, 5 mM, and 10 mM Ca2+. Experiments with biomolecules were
completed using 1 mM Ca2+ solutions and 0.5 mM biomolecule concentrations, resulting in
an initial molar ratio of 2 for [Ca2+]/[biomolecule]. Experiments for each solution condition
were conducted in repetitions of three. Experiments were carried out at ambient laboratory
temperature (25 ± 2 °C). The 1 mM calcium concentration trial completed on 05_12_2011
was excluded from data calculations. This trial resulted in a linear calcium loss curve that
did not fit with any of the other data. The succinic acid trial completed on 06_30_2011 was
also excluded from data calculations because it was highly abnormal when compared to the
other trials.
The experiments were terminated when the calcium concentration read by the
calcium ISE neared zero and precipitates were visible by eye in the reactor. Experiment
duration varied from four to eight hours. Average experiment lengths for trials without
biomolecules were 3.9 hours for 1 mM CaCl2, 5 hours for 5 mM CaCl2, and 7.5 hours for 10
mM CaCl2. Experiments (1 mM CaCl2) with aspartic acid averaged 6.2 hours, citric acid
averaged 5.5 hours, and succinic acid averaged 6.6 hours.
Atomic absorption spectroscopy (Thermo Scientific iE 3000 Series) was also used to
verify Ca2+ concentrations. Solution samples were taken from a port at the bottom of the
reactor throughout the experiment, in order to be analyzed via AAS.
Sample of
approximately 10 mL were taken at varying time intervals. These solutions were filtered in
order to remove and solid precipitates. After filtration 3 drops of 1 M HCl was added to each
sample in order to bring the pH to approximately 3, and thus prohibit the growth of crystals
after sampling.
30
2.2.3 Kinetic Modeling of Calcium Carbonate Precipitation
Kinetic modeling was conducted to determine the kinetic rate coefficient (k) for each
experiment.
Modeling was done via a Microsoft Excel spreadsheet that allowed for
comparison of experimental parameters. The kinetic rate constant describes a loss of calcium
over time (mol Ca2+ L-1 hr-1) and was determined for each experiment.
The rate of calcium carbonate precipitation can be related to saturation state using the
equation:
R  k    1
(2.1)
Where R is the rate of precipitation, k is the rate constant, and Ω is the saturation state of the
solution (Romanek et al. 2011).
To utilize equation 2.1, several simplifying assumptions must be made. Firstly, we
assume that the rate of precipitation formation can also be related to the loss of calcium from
dCa 2
solution over time,
:
dt
R
dCa 2
dt
(2.2)
dCa 2
is calculated from the two-side limit at a given point on a graph of Ca2+ vs. time.
dt
Therefore, the equation used for kinetic mode becomes:
31
dCa 2+
=-k(Ω-1)
dt
(2.3)
Secondly, data collected via the calcium ISE is assumed to represent the
concentration of free Ca2+ in the system. Thirdly, it is assumed that atmospheric CO2 is in
equilibrium with the aqueous phases. Even if this assumption is not strictly true, it is
reasonable to assume that the equilibrium value will be indicative of relative concentration
across experiments. However, the assumption is supported by the relative stability of pH
during the part of the experimental course used for this analysis (section 2.3.1). In regards to
ionic strength, the assumption was made that the effect on the system is minimal and thus it
was not included. The solutions used to precipitate calcium carbonate have low salinity and
are relatively dilute.
Given these parameters the activity coefficients are assumed to
approach unity and the ionic strength is negligible.
The concentration of aqueous carbon dioxide is thus equal to the concentration of
H2CO3*, and H2CO3* is defined as the product of Henry’s law constant and the partial
pressure of carbon dioxide:
(CO2)aq = (H2CO3*)
(2.4)
(H2CO3*) = KH pCO2
(2.5)
Where KH is Henry’s law constant, KH=10-1.15. The concentration of aqueous carbonate can
now be determined, given the equation:
(CO32- ) 
K 2 K 1( H 2CO3*)
( H  )2
32
(2.6)
Where K1 and K2 are dissociation constants for carbonic acid; K1 = 10-6.3 and K2= 10-10.3.
Finally, it is assumed that the precipitation product is vaterite, as discussed in chapter three.
The saturation state of vaterite is defined as:
vaterite 
[Ca 2 ][CO32- ]
KSPvaterite
(2.7)
Where KSPvaterite is the solubility product for vaterite, log KSPvaterite = -7.9 (Plummer and
Busenberg 1982).
The above calculations resulted in an experimental value for
dCa 2
and for (  1) ,
dt
which allowed for determination of k using Microsoft Excel. In order to ensure that the
kinetic rate constant was being calculated during calcium carbonate precipitation,
experimental data was chosen from the area of the calcium concentration graph that seemed
to display a linear regression for calcium loss. For each point on the curve, the difference
between the right and left hand side of equation 2.3 was calculated as the residual. The Goal
Seek function in Excel was then unitized to minimize the sum of the square of residuals by
varying the value of k.
33
2.3 Results and Discussion
2.3.1 pH
As seen in Figure 2.3, pH follows a general trend upwards over time. All solutions
were started at a pH of approximately 3, in order to inhibit any solid precipitation from
occurring before the start of the experiment.
Once the solution was exposed to the
ammonium carbonate in the reactor, a rapid increase of pH took place. The solutions then
stabilized at a pH of approximately 9. The 10 mM trial set is a suitable representation for
what occurred in all other trials. The trend for pH is similar in the presence of biomolecules
and also without biomolecules.
2.3.2 Carbon dioxide
Figure 2.4 shows the general trend for CO2 concentration increase over the duration
of the experiments. All trials should have started with a CO2 concentration slightly below
1000 ppm, normal for indoor carbon dioxide levels. Although there were discrepancies
between trials, the general trend was an increase in CO2 until eventual peaking at
approximately 7000 ppm.
Carbon dioxide concentrations followed similar patterns
regardless of whether or not the trial included biomolecules.
34
2.3.3 Calcium Data
Figure 2.5 shows the course of ISE data over time. The ISE reports concentrations
close to the initial concentration of 0.01 M. The concentration remains roughly constant until
it steadily drops, corresponding to the onset of precipitation. In addition, some samples were
collected, filtered, acidified, and analyzed for total Ca by atomic absorption spectroscopy
(AAS). In most cases, there was reasonable agreement between AAS and the ISE (Figure
2.6).
In some trials, excessive drift occurred in the [Ca2+] measured by ISE, which then
seemed to correct itself as time increased (Figure 2.7). In Figure 2.8 the solution initially
contained 10 mM CaCl2 at pH 3, which means that the initial concentration of calcium in
solution should be 0.01 M Ca2+. However, this figure shows an initial calcium concentration
of 0.014 M Ca2+, outside the range of error (± 20%). This figure also clearly shows an
increase following the initial decrease, which was not expected given the chemistry of the
system. It should be noted however, that the AAS and ISE measurements track closely
during calcium concentration decrease.
The most practical way to measure Ca in the reactor solution is an ISE. However,
these electrodes are widely known to perform inconsistently. In order to better understand the
limitations of this instrument, a series of tests were completed. Each test was done with 1
mM CaCl2 (0.001 M Ca2+), exactly as was done in the previous experiments. The first test
was run at a pH of approximately 3 (Figure 2.9), the second experiment was run at a pH of
approximately 10 (Figure 2.10), and the third experiment involved titration with 0.5 M
35
NaOH from an initial pH of approximately 3 to an end pH of approximately 10 (Figure
2.11). Each test shows drift at the beginning as the sensors become equilibrated with the
solution. The results show that the calcium ISE works fairly accurately at constant pH
(Figure 2.9, 2.10), but during titration a noticeable drift upward is observed as the pH of the
solution increases. This may account for drift in ISE seen in our experiments as pH changes.
The results of this testing, as well as comparison to AAS measurements, suggest that
ISE is susceptible to drift beyond normal error as pH changes. This manifests itself as a
physically unrealistic increase in concentration as pH increases in the reactor. However, the
electrode is reasonably reliable at constant pH, as is the case during crystallization. This
assertion is supported by the close agreement between AAS and ISE measurements as
calcium concentration decreases during precipitation. The calcium concentration from ISE is
thus used in our kinetic analyses.
2.3.4 Citric Acid Complexation
Experiments with citrate acid yielded graphs of calcium concentration over time with
a tiered shaped that did not match the shape of the graphs in all solution trials. All other
experiments, both those with biomolecules and without, shared a similar downward slope. A
representative graph for citric acid calcium loss can be seen in Figure 2.12. Additionally,
there is a significant discrepancy between AAS measurements of total calcium and ISE
measurements of Ca. Chemical speciation calculations were conducted using MINEQL+ at
concentrations of 1 mM and 0.5 mM total Ca and citrate, respectively, over a pH range of 3
36
to 10, also similar to actual experimental conditions. The results showed that at pH 5-8,
approximately 20-40% of calcium in the system was complexed by citrate. These results can
be seen in Figure 2.13. Further investigation using modeling-based experimental values of
pH, CO2, and total calcium concentration (as measured by AAS) indicates that calcium not
complexed by citrate is closer to the calcium ISE measurements rather than the AAS
measurements (Figure 2.14). The discrepancy between the calculated free calcium and ISE
measurement may be due to uncertainty in thermodynamic data for Ca-citrate complexes.
Based on its chemistry and the observable patterns of Ca decrease, complexation reactions
may also occur with succinic acid (Figure 2.15). However, a literature search yielded no
stability constants for succinic acid- calcium complexes, and thus we are unable to model
chemical speciation in this system.
2.3.5 Rates of Calcium Loss and Kinetic Modeling of Calcium Loss
Data calculated from kinetic modeling (2.2.3) is shown in Table 2.1. Trials with 10
mM Ca2+ had an average of 2.7 ± 1.8 × 10-6 mol Ca2+ L-1 hr-1 5 mM Ca2+ had an average of 2.3
7
± 1.3 × 10-7 mol Ca2+ L-1 hr-1 , 1 mM Ca
2+
had an average of 1.6 ± 1.1 × 10-6. In general, these
coefficients would be predicted to vary only weakly with Ca concentration. However, results
indicate that at 5 mM Ca2+ the kinetic rate coefficient is one order of magnitude smaller than
at 10 or 1 mM Ca2+. This discrepancy could be due to a difference in pH, pCO2, or [Ca2+]
between trials, as all parameters greatly impact the calculation of k. Even so, the kinetic rate
coefficients are consistently smaller than those in organic-free trials.
37
Modeling indicates that trials containing biomolecules produce larger kinetic rate
coefficients.
All biomolecule trials resulted in average kinetic rate coefficients with a
magnitude of 10-5 mol Ca2+ L-1 hr-1 with respect to vaterite. Aspartic acid had an average of 1.7
± 1.7 × 10-5 mol Ca2+ L-1 hr-1, succinic acid had an average of 3.9 ± 3.5 × 10-5 mol Ca2+ L-1 hr-1, and
citric acid had an average of 6.0 ± 7.8 × 10-5 mol Ca2+ L-1 hr-1. Trials with citrate were
specifically difficult to model due to a reduced data range caused by metal complexation.
Figure 2.12 shows the effects of citrate complexation on calcium loss over time. The first
rapid decrease (t ~ 0.5 hr) is believed to be calcium loss due to complexation with citrate.
The second area of rapid decrease (t ~ 4) is most likely calcium loss due to CaCO3 formation.
The smaller data range of CaCO3 formation, starting at t ~ 4, for citric acid trials is a source
of error throughout these experiments. In all measurements, standard deviations are large,
highlighting difficulty in using the reactor and the limitations of the ISE (as discussed in the
appendix on reactor development). Many factors go into the calculation of the kinetic rate
coefficient, and slight differences in pH, pCO2, or [Ca2+] will cause noticeable difference
between coefficients. However, trials without biomolecules had average kinetic coefficient
rates one or two magnitude smaller than those with biomolecules, indicating a substantial
effect on precipitation rate.
2.3.6 Effects of Biomolecules on Calcium Carbonate Precipitation
Biomolecules may affect the preicpitaion of calcium carbonate in several different
ways. Classical nucleation theory describes how crystal formation starts with ions that form
38
into initially unstable aggregates. These particles are unstable due to surface free energy
occurring at the crystal-solvent interface. Crystals increase in stability as they grow in size,
since the bulk lattice energy begins to overcome the surface free energy (Raiteri and Gale
2010). The Ostwald-Lussac rule states that the lowest energy phase will be precipitated first,
and biomineralization follows this rule as well (Weiner and Dove 2003). As noted earlier, a
study by Elhadj et al. (Elhadj et al. 2006) found that biomolecules can increase the rate of
crystal formation by lowering the energy barrier for solutes to bind to solids. This increase in
rate comes from an increase in the kinetic coefficient, which is due to the reduction of the
diffusive barrier. The data on kinetic rate constants indicates that this is the case for citric,
aspartic, and succinic acid.
Supersaturation is the first step towards nucleation (Weiner and Dove 2003). The
addition of biomolecules can increase the rate of formation by creating nucleation sites that
allow precipitation to occur at a smaller saturation state (Lasaga 1998). This fits with the
calculations of saturation state for each trial, in which biomolecule trials produced a lower Ω
with regards to vaterite.
Stereochemical recognition, or geometric matching, provides surface charge potential
that can increase the formation of clusters at the solute-solvent interface (Mann et al. 1990).
The biomolecules used in these trials contain carboxylates, which bind strongly to and may
closely resemble the structure of CaCO3. Therefore, stereochemical recognition may provide
a plausible explanation for the increased rate of formation for trials including biomolecules.
39
2.4 Conclusions
Carboxylate-containing biomolecules affect the rate of calcium carbonate
precipitation. The addition of citric, aspartic, and succinic acid results in a larger kinetic rate
constant and thus a faster rate of formation.
The classical nucleation theory and the
stereochemical recognition model provide rationales for how these biomolecules that contain
carboxylic acid, which is structurally similar to carbonate, increase the rate of formation by
lowering surface energy barriers. Additionally, these biomolecules can produce nucleation
sites that promote precipitation at lower saturation states and lower the energy barrier,
allowing for faster rates of formation.
The presence of citric acid will change the kinetics of precipitation reactions, due to
the formation of calcium citrate complexes which remove aqueous calcium ions from
solution until the system reaches a more basic pH and calcium carbonate formation can
commence. This also has implications for reducing crystal size (Chapter 3). Precipitate
phase is another important aspect of calcium carbonate precipitation and will be discussed in
Chapter 3.
40
2.5 References
Addadi, L. and Weiner, S. (1985). Interactions between acidic proteins and crystals:
Stereochemical requirements in biomineralization. Proceedings of the National Academy of
Sciences. 82, 4110-4114.
Addadi, L. and Weiner, S. (2001). Crystals, asymmetry and life. Nature. 411, 753-755.
De Yoreo, J.J. and Dove, P.M. (2004). Shaping crystals with biomolecules. Science. 306,
1301-1302.
De Yoreo, J.J., Zepeda-Ruiz, L.A., Friddle, R.W., Qiu, S.R., Wasylenki, L.E., Chernov,
A.A., Gilmer, G.H., and Dove, P.M. (2009). Rethinking classical crystal growth models
through molecular scale insights: Consequences of kink-limited kinetics. Crystal Growth and
Design. 9, 5135-5144.
Dickinson, S.R., Henderson, G.E., McGrath, K.M. (2002). Controlling the kinetic versus
thermodynamic crystallization of calcium carbonate. Journal of Crystal Growth. 244, 369378.
Dove, P.M. (2010). The rise of skeletal biominerals. Elements. 6, 37-42.
Elhadj, S., De Yoreo, J.J., Hoyer, J.R., and Dove, P.M. (2006). Role of molecular charge and
hydrophilicity in regulating the kinetics of crystal growth. Proceedings of the National
Academy of Sciences. 103, 19237-19242.
Lasaga, A.C. (1998). Kinetic theory in the earth sciences. Princeton University Press,
Princeton, N.J.
41
Li, W., Liu, L., Chen, W., Yu, L., Li, W., and Yu, H. (2010). Calcium carbonate precipitation
and crystal morphology induced by microbial carbonic anhydrase and other biological
factors.
Mann, S., Heywood, B.R., Rajam, S., and Walker, J.B.A. (1990). Structural and
stereochemical relationships between Langmuir monolayers and calcium carbonate
nucleation. Journal of Applied Physics. 24, 154-164.
Mukkamala, S.B., Anson, C.E., and Powell, A.K. (2006). Modelling calcium carbonate
biomineralisation processes. Journal of Inorganic Biochemistry. 100, 1128-1138.
Orme, C.A., Noy, A., Wierzbicki, A., McBride, M.T., Grantham, M., Teng, H.H., Dove,
P.M., DeYoreo, J.J. (2001). Formation of chiral morphologies through selective binding of
amino acids to calcite surface steps. Nature. 411, 775-779.
Pankow, J.F. (1991).
Chapter 1: Overview.
In Aquatic Chemistry Concepts.
Lewis
Publishers. Boca Raton, FL.
Plummer, L.N., Busenberg, E. (1982). The solubilities of calcite, aragonite and vaterite in
CO2-H2O solutions between 0 and 90°C, and an evaluating of the aqueous model for the
system CaCO3-CO2-H2O. Geochmica et Cosmochimica Acta. 46, 1011-1040.
Raiteri, P. and Gale, J.D. (2010). Water is the Key to Nonclassical Nucleation of Amorphous
Calcium Carbonate. Journal American Chemical Society. 132, 17623-17634.
Romanek, C.S., Morse, J.W., and Grossman, E.L. (2011). Aragonite kinetics in dilute
solutions. Aquatic Geochemistry. 17, 339-356.
42
Ruiz-Agudo, E., Putnis, C.V., Rodriquez-Navarro, C., Putnis, A. (2011). Effect of pH on
calcite growth at constant αCa2+/αCO32- ratio and supersaturation.
Geochimica et
Cosmochimica Acta. 75, 284-296.
Teng, H.H., Dove, P.M., and DeYoreo, J.J. (1999). Reversed calcite morphologies induced
by
microscopic
growth
kinetics:
Insight
into
biomineralization.
Geochimica
et
Cosmochimica Acta. 63, 2507-2512.
Weiner, S. and Dove, P.M. (2003) An Overview of Biomineralization and the problem of the
vital effect. In Biomineralization. Eds. P.M. Dove, S. Weiner and J.J. De Yoreo.
Mineralogical Society of America, Washington, D.C., v. 54, 1-31.
Wolf, G., Lerchner, J., Schmidt, H., Gamsjager, H., Konigsberger, E., Schmidt, P. (1996).
Journal of Thermal Analysis. 46, 353-359.
43
Figure 2.1. A photograph of the crystal growth chamber.
Figure 2.2. A schematic of the growth chamber, designed to allow monitoring and
manipulation of aqueous and gaseous chemical conditions during the growth of carbonate
minerals.
44
Figure 2.3. A comparison of pH over time for three repetitions of a trial set containing 10
mM CaCl2. Each trial started at a pH of approximately 3 and ended at a pH of approximately
9. This data set is a concise representation of the behavior of pH in each experiment.
45
Figure 2.4. A comparison of CO2 concentration over time for three repetitions of a trial set
containing 10 mM CaCl2. These trials show the general increase over time of CO2
concentration.
46
Figure 2.5. Comparison of calcium concentration over time for three trials containing 10
mM CaCl2.
47
Figure 2.6. Comparison of calcium concentration measured by calcium ISE and atomic
adsorption spectroscopy for a trial that started with 0.01 M Ca2+ done on 09_09_2011.
Figure 2.7. Comparison of calcium concentration measured by calcium ISE and atomic
adsorption spectroscopy for a trial that started with 0.001 M Ca2+ done on 05_26_2011. This
figure shows electrode drift at the start of the experiment, and that the electrode readings
become closer to the AAS readings as time goes on.
48
Figure 2.8. Calcium loss over time for trial 04_15_2011.
Figure 2.9. Calcium ISE test done at pH approximately 3 with a solution of 1 mM CaCl2
(0.001 M Ca2+).
49
Figure 2.10. Calcium ISE test done at pH approximately 10 with a solution of 1 mM CaCl2
(0.001 M Ca2+).
Figure 2.11. Calcium ISE test titration. A solution of 1 mM CaCl2 (0.001 M Ca2+) at
approximately pH 3 was titrated with 0.5 M NaOH at an interval of 1 drop every 30 seconds
for six minutes, until an end pH of approximately 10.
50
Figure 2.12. Calcium loss over time for a trial initially containing 1 mM CaCl2 and 0.5 mM
citric acid, completed on 06_02_2011.
Figure 2.13. Concentration of aqueous species involving citrate formed in a solution of 1
mM (1x10-3 M) CaCl2 with 0.5 mM (5x10-4 M) citrate at a range of pH from 3 to 10. Model
created with MINEQL+.
51
Figure 2.14. Comparison of calcium concentration measured by calcium ISE and atomic
adsorption spectroscopy, as well as calculations for calcium in solution not complexed by
citrate. This trial that started with 0.001 M Ca2+ and contained 0.5 mM citric acid and was
done on 06_02_2011. MINEQL+ calculations were done with 1 mM Ca2+, 0.5 mM citrate,
and pH and pCO2 concurrent with each time point of this trial.
Figure 2.15. Concentration of calcium over time for a trial containing 1 mM calcium and 0.5
mM succinic acid completed on 06_28_2011. Evidence of complexation may be seen
between hours 0 to 3.
52
Table 2.1 Table of kinetic rate constants for the formation of vaterite. Uncertainties are
estimated as standard deviations of replicate measurements.
Date
04_15_2011
04_18_2011
09_09_2011
Solution (mM Ca2+)
10
10
10
10 mM Average
k (mol Ca2+ L-1 hr-1)
1.4 ×10-6
1.9 ×10-6
4.7 ×10-6
2.7 ± 1.8 ×10-6
04_05_2011
04_08_2011
04_11_2011
5
5
5
5 mM Average
1.2 ×10-7
2.1 ×10-7
3.7 ×10-7
2.3 ± 1.3 ×10-7
05_16_2011
05_26_2011
1
1
1 mM Average
8.0 ×10-7
2.3 ×10-6
1.6 ± 1.1×10-6
06_02_2011
06_07_2011
06_09_2011
1 + 0.5 mM Citric Acid
1 + 0.5 mM Citric Acid
1 + 0.5 mM Citric Acid
Citric Acid Average
1.8 ×10-5
1.3x10-5
1.5x10-4
6.0 ± 7.8 ×10-5
06_14_2011
06_28_2011
1 + 0.5 mM Succinic Acid
1 + 0.5 mM Succinic Acid
Succinic Acid Average
6.3 ×10-5
1.4 ×10-5
3.9 ± 3.5 × 10-5
07_18_2011
07_19_2011
07_21_2011
1 + 0.5 mM Aspartic Acid
1 + 0.5 mM Aspartic Acid
1 + 0.5 mM Aspartic Acid
Aspartic Acid Average
1.2 ×10-5
3.5 ×10-5
2.7 ×10-6
1.7 ± 1.7 ×10-5
53
Table 2.2 Table of average saturation state for each set of trials. Calculations were based on
the assumption that the dissolved phase and the gas phase for CO2 are in equilibrium.
Uncertainties are estimated as standard deviations of replicate measurements. High
supersaturation may reflect the limitations of assumptions in the kinetic model.
Solution (mM Ca2+)
Avg. Ω
10
3427 ± 2651
5
53389 ± 48673
1
3066 ± 2155
1 + 0.5 mM Citric Acid
214 ± 149
1 + 0.5 mM Succinic Acid
245 ± 191
1 + 0.5 mM Aspartic Acid
807 ± 806
54
2.6 Appendix
2.6.1 Development of Reactor
A set of experiments were completed utilizing a trial-and-error approach, in order to
determine the best course of action for the experimental methodology of this research. The
chemicals and concentrations used were directly correlated to those in the literature from
related research projects.
Three of these trials were analyzed using Scanning Electron
Microscopy. Improper storage of the crystals most likely led to the transformation from
vaterite to calcite, which can be seen in Figure 2.16 – 2.18.
Initial experiments were carried using calcium nitrate (Ca(NO3)2) (Fisher Scientific)
as calcium source and sodium bicarbonate (NaHCO3) (Fisher Scientific) as a carbonate
source.
Ammonium bicarbonate (NH4HCO3) (Fisher Scientific) acted as an additional
carbonate source and also a means of driving pH upward so that the precipitation reaction
could take place. Aqueous solutions of calcium nitrate and sodium bicarbonate were placed
in the bottom of the growth chamber in the same experimental design described above
(section 2.2.2 Experimental Setup). Experiments were carried out at concentrations of 100
mM Ca(NO3)2 and 40 mM NaHCO3, 60 mM Ca(NO3)2 and 40 mM NaHCO3, and then at
concentrations of 40 mM Ca(NO3)2 and 40 mM NaHCO3. The solutions were brought to pH
3 with a 1 M solution of HCl. Ammonium bicarbonate was weighed out to approximately 10
g and was placed in the top of the chamber, and in these experiments no water was added to
the NH4HCO3. A trial using 100 mM Ca(NO3)2, 40 mM NaHCO3, and approximately 10 g
of ammonium bicarbonate resulted in solids that were analyzed with scanning electron
microscopy (Figure 2.16).
55
Additional literature search determined that ammonium carbonate was often used in
such experiments in the place of ammonium bicarbonate, since the increased ammonium
levels lead to a more rapid pH increase (Naka et al. 2002). Additionally, a small amount of
deionized water had been used in other studies to help speed up the volatilization process
(Dickinson et al. 2002). The next experiment used approximately 10 g of ammonium
carbonate ((NH4)2CO3) with 20 mL deionized water in the top of the chamber, with an
aqueous solution of 40 mM Ca(NO3)2 and 40 mM NaHCO3 in the bottom of the chamber.
Solutions were again brought to approximately pH 3. The solids formed in this trial were
analyzed with scanning electron microscopy (Figure 2.17).
The literature search also indicated that other experiments used calcium chloride
(CaCl2) instead of calcium nitrate and sodium bicarbonate (Dickinson et al. 2002). Since
chloride is also a common ion present in seawater, the next set of experiments replaced the
calcium nitrate and sodium bicarbonate with calcium chloride. These experiments were done
with concentrations of 5, 10, 20, 30, and 40 mM CaCl2. Approximately 10 g of ammonium
carbonate was used, submerged in 20 mL of deionized water. As always, solutions were
brought to a pH of approximately 3. A trial with 40 mM CaCl2 was analyzed with scanning
electron microscopy (Figure 2.18).
These previous trials helped created the experimental method used for this thesis.
Calcium chloride was chosen as the calcium source due to its prevalence in the literature as
well as its relation to seawater chemistry.
Ammonium carbonate was chosen as the
carbonate source and as a means of increasing pH. This source was chosen for its increased
ammonium concentration; however the mass was decreased from 10 g to 5 g to help slow
56
reaction rates.
Deionized water was added to the ammonium carbonate to help with
volatilization. Calcium chloride concentrations of 1, 5, and 10 mM were chosen so that the
calcium ISE would last longer and work more effectively. As crystals grow on the electrode
they compromise the effectiveness of the equipment.
The lower calcium chloride
concentration allowed for the experiments to be completed more rapidly and thus the calcium
ISE would spend less time submerged in the calcium solution, allowing it a longer shelf life.
Carbon dioxide samples were originally taken at intervals and injected into a Licor
carbon dioxide analyzer. This technique was then replaced with the Licor LI-840A which
allowed for continuous CO2 analysis.
2.6.2 Limitations of the Reactor
Experiments utilizing the reactor were often unwieldy and challenging. Calcium data
was complicated by included complexation with biomolecules, as well as ISE drift. Issues
with the carbon dioxide analyzer happened on more than one occasion, and differences in
CO2 concentrations within trial sets can be seen in the appendix. The Licor analyzer is fed a
constant stream of gas from the top of the reactor with the aid of a pump. The pump failed
on multiple occasions, and this may have led to discrepancies in CO2 data. It should be noted
that the original intent was to directly measure dissolved CO2 using optical fluoresce
methods. However, production hold up prevented the manufacture of this device. Also, the
volatile nature of ammonium carbonate may have lead to different concentrations of carbon
dioxide being produced.
57
2.6.3 References
Dickinson, S.R., Henderson, G.E., McGrath, K.M. (2002). Controlling the kinetic versus
thermodynamic crystallization of calcium carbonate. Journal of Crystal Growth. 244, 369378.
Naka, K., Tanaka, Y., Chujo, Y. (2002). Effect of anionic starburst dendrimers on the
crystallization of CaCO3 in
58
Figure 2.16. SEM image at 100x of a trial containing 100 mM Ca(NO3)2 and 40 mM
NaHCO3 done with approximately 10 g of ammonium bicarbonate (NH4HCO3) on
03_17_2010.
Figure 2.17. SEM image of a trial containing 40 mM Ca(NO3)2 and 40 mM NaHCO3 done
with 10 g of ammonium carbonate ((NH4)2CO3) with 20 mL deionized water on 06_10_2010.
59
Figure 2.18. SEM image at 1500x of a trial containing 40 mM CaCl2 and 10 g (NH4)2CO3
with 20 mL deionized water on 07_20_2010.
60
Chapter 3
Determination of Phase and Morphology
61
3.1 Introduction
There are three common crystalline polymorphs of CaCO3 in nature: calcite,
aragonite, and vaterite. Calcite has a rhombohedral structure, aragonite is orthorhombic, and
vaterite is hexagonal. Under ambient conditions, calcite is the thermodynamically stable
phase, whereas aragonite is typically metastable, and vaterite is unstable (Xu et al., 2010a).
Because of these trends in stability, calcite and aragonite are the polymorphs most often
found in nature (Spann et al., 2010). However, each polymorph may play significant roles in
geology, biology, or industrial applications.
Calcite and dolomite (CaMg(CO3)2) are the most abundant rock-forming carbonate
minerals.
carbonates.
Together these two minerals account for over 90% of naturally existing
Calcium carbonate has more natural polymorphic variation than any other
mineral (Reeder 1983). The main reason for this phenomenon is that the metastable form
aragonite was likely abundant throughout the ancient ocean (just as it still is in the modern
ocean), but this polymorph may transform into calcite during diagenesis.
Aragonite
formation is favored in the presence of magnesium and or high turbidity. Under ambient
conditions, aragonite will transform to calcite (Reeder 1983); however, transformations from
one polymorph to another do not frequently occur in seawater. Proposed reasons for this are
inhibition of aragonite dissolution as well as inhibition of calcite nucleation and growth,
mostly likely caused by the presence of magnesium (Carlson, 1983).
Vaterite has high specific surface area, high solubility, and lower density (ρ = 2.65 g
cm-3) than either calcite or aragonite (Carlson 1983; Borah et al., 2011). In nature, vaterite is
often used in shell regeneration, shell formation, and pearl production (Spann et al., 2010).
62
Due to its high solubility and low stability, vaterite is often a precursor to other polymorphs
of calcium carbonate (Becker and Hu 2009). Although vaterite is the least stable polymorph,
it can potentially be stabilized by the addition of organic or inorganic molecules. Research
has shown that positively charged ammonium ions can adsorb to polar surfaces of vaterite,
thus inhibiting growth and promoting the hexagonal morphology indicative of vaterite
(Becker and Hu, 2009).
A goal of this study was to understand the phase and morphology of calcium
carbonate precipitates formed in the presence of carboxyl-containing biomolecules. Products
of crystal growth experiments (chapter 2) were examined by X-ray diffraction and scanning
electron microscopy. In order to aptly observe the morphological effects of biomolecules,
these crystals were to be compared to calcium carbonate solids grown in an organic-free
solution.
3.2 Materials and Methods
3.2.1 The Chemical System
Crystal growth results from the combination of aqueous calcium with aqueous
carbonate, as described in detail in chapter 2. The following reaction sequence relates
atmospheric CO2 gas to CaCO3 precipitates.
CO2(g) ↔ CO2(aq)
(3.1)
CO2(aq) + H2O ↔ H2CO3
(3.2)
H2CO3 ↔ H+ + HCO3-
(3.3)
63
HCO3- ↔ H+ + CO32-
(3.4)
Ca2+ + CO32- ↔ CaCO3(s)
(3.5)
Gaseous carbon dioxide is dissolved in water to from an aqueous species (3.1), which then
combines with water to form carbonic acid (3.2). Carbonic acid dissociates, producing a
hydrogen ion and bicarbonate (3.3), which in turn dissociates to form a hydrogen ion and
carbonate (3.4). Calcium ions combine with the carbonate species to solid calcium carbonate
(3.5) (Mitchell et al. 2011). The calcium carbonate formed can conceptually be any of the
poymorphs of CaCO3, depending on the conditions of formation.
3.2.2 Materials
Solid calcium carbonate crystals were collected from the reactor in which they were
grown (chapter 2). The crystals were grown on a polycarbonate sheet at the bottom of the
reactor. At the end of the experiment, the solution was drained from the reactor by a port at
the bottom of the chamber. The crystals were then allowed to air dry briefly and were then
placed in plastic bags which were sealed and then placed in a freezer. Samples are identified
by an eight digit code (XX_XX_XXXX) that corresponds to the date of the growth
experiment that the produced the sample.
3.2.3 X-ray Diffraction
X-ray diffraction (XRD) analyzes the scattering of X-rays by a crystal lattice via
Bragg’s Law:
64
n λ =2 d sin θ
(3.6)
where λ is the wavelength of X-rays, d is the spacing between atomic planes in the mineral, θ
is the incident angle of the incoming light, and n is the order of reflection. Because most
minerals contain a unique combination of atomic plane spacings, diffraction patterns are
typically diagnostic of specific phases. By comparison to known standards, minerals can thus
be identified.
A study by Dickinson and McGrath (2001) found that not only is XRD satisfactory
for determining the present polymorphs of calcium carbonate, but that XRD is superior to
both Raman and infrared spectroscopies. After the completion of the crystal growth trials, a
Rigaku SmartLab XRD was used to analyze the phase of solid crystal precipitates. Due to
small sample quantities (often times less than 50 mg), samples were analyzed on the
polycarbonate sheet on which they precipitated during the experiment. The XRD analysis
was conducted using parallel beam setting with angle of 1°. This glancing angle helped to
ensure that the X-ray beam made contact with the crystals on the surface of the
polycarbonate sheet. A diffractogram was also collected on a clean piece of polycarbonate as
a control, and this diffractogram was compared with the experimental diffractograms to
ensure that the polycarbonate sheet was not interfering with the diffraction analysis.
Fingerprinting of the phases present was completed by the computer program PDXL
(Rigaku). The diffractograms produced by this software were then confirmed with numbers
from the literature (Wang and Becker, 2009 and Kamhi, 1963).
65
3.2.4 Scanning Electron Microscope
A JEOL JSM-5900LV Scanning Electron Microscope (NCSU Center for Electron
Microscopy) was used to analyze the morphology of solid samples after the completion of
the experiments. Experiments that produced samples with a larger quantity (viz. 5 mM and
10 mM CaCl2 solutions) were scraped off of the polycarbonate sheet they formed on, and
then mounted directly on a SEM stub (Ladd-Research Industries). Examples that resulted in
a smaller quantity of crystals (viz. 1 mM CaCl2 solutions) were mounted on the stub while
still attached to the polycarbonate sheet they had formed on.
The crystals and the
polycarbonate film were both attached to the stubs with double sided tape, than the edges of
the sample and the stub were coated with conductive silver paint (Ladd-Research Industries),
in order to ensure ground contact would be achieved within the microscope. The samples
were coated to a thickness of 0.0125 nm with a Au/Pd mixture using a Hummer 6.2 sputter
coater.
3.3 Results and Discussion
3.3.1 XRD Results
XRD difractograms were compared to known CaCO3 polymorph data (Table 3.1).
An intensity ratio of 100 indicates the most intense peak, and is the best reference for
determining which phase is present in the diffractogram. Other peaks are referenced to the
most intense peak. The parameter “2θ” refers to the angle of the diffracted X-ray. For crystal
growth trials beginning from a solution of 10 mM and 5 CaCl2, only vaterite was detected
(Figures 3.1 and 3.2). Two trials (05_12_2011 and 05_16_2011), for crystals grown from 1
66
mM CaCl2 solution showed only small peaks that were barely above background, possibly
due to a small mass of CaCO3. However, there did appear to be a peak at approximately 32°
for both trials, which is indicative of the polymorph vaterite; thus, these trials are assigned
vaterite as the dominant mineral phase. The third trial (05_26_2011) from 1 mM CaCl2
resulted in an improved diffractogram, which indicated the presence of vaterite (Figure 3.3).
Thus, X-ray diffraction analysis of crystals grown in organic-free solution detected only the
presence of vaterite.
Similar results were obtained for crystal grown in the presence of aspartic acid and
succinic acid. All three trials for 1 mM CaCl2 + 0.5 mM aspartic acid resulted in the vaterite
(07_18_2011, 07_19_2011, and 07_21_2011) (Figure 3.4). Similarly, trials for the 1 mM
CaCl2 + 0.5 mM succinic acid experiments resulted diffraction patterns consistent with
vaterite (Figure 3.5).
Trials including 1 mM CaCl2 + 0.5 mM citric acid produced
diffractograms that were highly similar to that of the polycarbonate blank (06_02_2011,
06_07_2011, and 06_09_2011) (Figure 3.6). The lack of peaks suggests either the presence
of amorphous phases or an insufficient mass of material for X-ray diffraction analysis. A
sample of clean polycarbonate film was analyzed by XRD as a blank for comparison with the
background of the diffractogram (Figure 3.7)
3.3.2 SEM Results
Microscopy can be used to view morphologies that are diagnostic of the different
phases of CaCO3. Visual representations of these polymorphs can be seen in Figure 3.8. In
the absence of growth modifiers, calcite tends to form rhombohedral crystals. Aragonite
67
tends to from needle-shaped laths. Vaterite normally forms spheriodal particles (Xu et al.,
2010a). The results of microscope analysis can be compared to results obtained from the
XRD (Table 3.2). The results achieved from SEM varied slightly from those obtained by
XRD, suggesting that SEM may have better sensitivity than XRD for the detection of minor
phases.
Trials grown from 10 mM CaCl2 solution were identified as vaterite (04_15_2011)
and as vaterite and calcite (04_18_2011 and 09_09_2011).
Vaterite crystals were
approximately 20 to 30 µm. Calcite crystals were approximately 10 µm. One trial resulted
in much larger crystals (04_18_2011) and these were approximately 40 µm for vaterite and
approximately 20 µm for calcite (Figure 3.9).
All three trials from 5 mM CaCl2 were identified as vaterite (04_05_2011,
04_08_2011, and 04_11_2011). The vaterite crystals were approximately 20 to 30 µm
(Figure 3.10). The first trial from 1 mM CaCl2 was identified as vaterite and calcite
(05_12_2011) while the other two trials only showed vaterite (05_16_2011 and
05_26_2011). Vaterite crystals were approximately 20 µm and the calcite crystals were also
approximately 20 µm in diameter (Figure 3.11).
All three trials grown from 1 mM CaCl2 + 0.5 mM citric acid contained vaterite and
calcite polymorphs (06_02_2011, 06_07_2011, and 06_09_2011). Citric acid resulted in the
smallest crystals.
Vaterite crystals were approximately 10 µm;
calcite crystals were
observed to be approximately 20 µm in diameter (Figure 3.12). This was also the only
solution that resulted in calcite crystals larger than the vaterite crystals. Additionally, in
68
Figure 3.12B an intergrowth of calcite and vaterite can be seen. It is possible that this image
documents the transformation of vaterite into calcite, but this is difficult to ascertain from a
single image.
All three trials grown from 1 mM CaCl2 + 0.5 mM succinic acid showed only the
vaterite polymorph (06_14_2011, 06_28_2011, and 06_30_2011). Vaterite crystals were
approximately 10 to 20 µm in diameter (Figure 3.13). Similarly, all three trials grown from
1 mM CaCl2 + 0.5 mM aspartic acid were identified as only vaterite (07_18_2011,
07_19_2011, and 07_21_2011).
Vaterite crystals were approximately 20 to 30 µm in
diameter (Figure 3.14).
Overall, scanning electron microscopy confirmed XRD measurements. Vaterite was
identified in all experiments, with a less common presence of calcite. In no trial was
aragonite identified. A summary of the polymorph identified and the frequency can be seen
in Table 3.2.
3.3.3 Formation of CaCO3 in the Absence of Biomolecules
A study by Dickinson et al. examined the morphologies of calcium carbonate grown
in solutions with varying pCO2 and [Ca2+]. This study found that the growth of solid calcium
carbonate is primarily influenced by the amount of calcium ions present in the aqueous
solution. This is a reasonable conclusion, since supersaturation (which leads to nucleation) is
highly dependent on calcium concentration. Results from this study showed that low pCO2
and low [Ca2+] promote growth of the thermodynamically stable polymorph calcite. Also,
lower concentrations lead to multinucleated calcite crystals. The number of nucleation sites
69
increased with higher concentration of calcium in the initial solution.
The perfect
rhombohedral calcite was seen at approximately 6 M calcium solution, but was highly
dependent on pCO2. Vaterite was found to be dependent on both pCO2 and [Ca2+], and the
quantity of vaterite increased with the increase of each parameter.
High calcium
concentrations occasionally resulted in aragonite, although this was a small component
(Dickinson et al. 2002).
As mentioned earlier, vaterite is an unstable polymorph of CaCO3 under ambient
conditions, yet ammonium has been shown to promote and even stabilize vaterite (Becker
and Hu 2009). The presence of ammonia gas could therefore be a cause for the prevalence of
vaterite in these experiments. Another possible cause is the elevated concentration of pCO2
throughout the experiment. As Dickinson et al. (2000) noted, an increase in either pCO2 or
[Ca2+] often lead to the presence of vaterite. That same study also mentioned that at lower
[Ca2+] and higher pCO2 vaterite was present in aggregated plates rather than individual platelike crystals. The vaterite produced by the experiments discussed within this thesis was also
the aggregated plate form.
3.3.4 Effect of Biomolecules
Organisms produce biominerals to perform a specific function, and they create these
biogenic materials with properties that best suit that function (Sugawara et al. 2003). With
the addition of three different biomolecules, we expected to observe a variety of
morphologies and sizes.
70
Orme et al. (2001) found that the addition of chiral amino acids will produce chiral
calcite crystals. Our experiments included the addition of aspartic acid, yet those trials did
not produce calcite. One major difference between our trials and the trials done by Orme et
al. (2001) was the chemical system itself, namely the addition of ammonium carbonate in our
system. Orme et al. (2001) used NaCl, NaHCO3, and CaCl2 × 2H2O in their growth solution,
and did not use any ammonium-containing compounds. Becker and Hu (2009) found that
ammonium has the ability to produce and stabilize the normally metastable form of vaterite.
Xu et al. (2010b) examined the polymorphs of CaCO3 created in the presence of
trisodium citrate at high temperatures. This study produced rhombohedral calcite as well as
what the authors refer to as “dumbbell-like crystals”. In the absence of trisodium citrate, the
rhombohedral form of calcite was produced. XRD confirmed that all crystals produced were
calcite.
Again, these experiments were done without ammonium compounds, further
supporting the assertion that the high concentration of ammonium in our system is a likely
reason for the abundance of vaterite produced.
3.4 Conclusions
In our experiments, biomolecules do not have a pronounced effect of the
polymorphism of calcium carbonate precipitation. All trials, with and without biomolecules,
result in the formation of vaterite. All vaterite crystals appear to be in aggregated plate form.
However, citric acid seems to promote calcite production, possibly by transformation through
vaterite. Trials without biomolecules create calcite, though not consistently.
71
Citric acid produces the smallest vaterite crystals and comparatively large calcite
crystals. Succinic acid and aspartic acid produced vaterite crystals of relatively the same
size.
Neither succinic acid nor aspartic acid promoted calcite formation.
It has been
proposed that high pCO2 will result in system that is kinetically controlled and will produce
vaterite (Dickinson et al. 2002). Since all biomolecule trials were under high pCO2 and
produced vaterite, citric acid must be producing calcite by thermodynamic control, possibly
by lowering the energy barrier required for calcite production. Greater concentrations of
calcium favored larger vaterite crystals.
72
3.5 References
Borah, B.M., Dey, S.K., and Das, G. (2011). Crystal to calcite: Fabrication of pure calcium
carbonate polymorph in the solid state. American Chemical Society: Crystal Growth and
Design. 11, 2773-2779.
Becker, U. and Hu, Q. (2009). Controlled growth of different calcium carbonate polymorphs
as induced by the presence of dissolved molecules and mineral surfaces. Goldschmidt
Conference. June 21 -26, Davos, Switzerland.
Carlson, W. (1983). The polymorphs of CaCO3 and the aragonite-calcite transformation. In
Carbonates: Mineralogy and Chemistry, Reeder, R. J., Ed. Mineralogical Society of
America: Washington D.C.
De Villiers, J.P.R. (1971). Crystal structures of aragonite, strontianite, and witherite.
American Mineralogist. 56, 758-767,
Dickinson, S.R., and McGrath, K.M. (2001). Quantitative determination of binary and
tertiary calcium carbonate mixtures using powder X-ray diffraction. The Royal Society of
Chemistry. 126, 1118-1121.
Dickinson, S.R., Henderson, G.E., McGrath, K.M. (2002). Controlling the kinetic versus
thermodynamic crystallization of calcium carbonate. Journal of Crystal Growth. 244, 369378.
Graf, D.L. (1961). Crystallographic tables for the rhombohedral carbonates. American
Mineralogist. 46, 1283-1316
Kamhi, S. R. (1963). On the structure of vaterite, CaCO3. Acta Crystallographica. 16, 770772.
Orme, C.A., Noy, A., Wierzbicki, A., McBride, M.T., Grantham, M., Teng, H.H., Dove,
P.M., and DeYoreo, J.J. (2001). Formation of chiral morphologies through selective binding
of amino acids to calcite surface steps. Nature. 411, 775-779.
Reeder, R.J. (1983). Crystal chemistry of the rhombohedral calcites. In Carbonates:
Mineralogy and Chemistry, Reeder, R. J., Ed. Mineralogical Society of America:
Washington D.C.
Spann, N., Harper, E.M., Aldridge, D.C. (2010). The unusual mineral vaterite in shells of the
freshwater bivalve Corbicula fluminea from the UK. Naturwissenschaften. 97, 743-751.
Sugawara, T., Suwa, Y., Ohkawa, K., and Yamamoto, H. (2003). Growth of calcite with
chiral phosphoserine copolypeptides. Macromolecular Rapid Communications. 24, 847-851.
73
Wang J., Becker U. (2009). Structure and carbonate orientation of vaterite (CaCO3).
American Mineralogist. 94, 380-386.
Xu, X., Zhao, Y., Lai, Q., Hao, Y. (2010a). Effect of polyethylene glycol on phase and
morphology of calcium carbonate. Journal of Applied Polymer Science. 119, 319-324.
Xu, X, Lai, Q., Zhao, Y., Hao, Y., Zeng, H., and Wang, L. (2010b). Preparation of calcium
carbonate crystals in the presence of trisodium citrate. Crystal Research and Technology. 45,
712-716.
74
Table 3.1. Intensity ratios and corresponding degrees for the determination of synthetic
vaterite (Kamhi, 1963), theoretical vaterite (Wang and Becker 2009), calcite (Graf 1961), and
aragonite (de Villiers 1971).
CaCO3 Polymorph
Vaterite - synthetic
Vaterite - theoretical
Calcite
Aragonite
Intensity Ratio
100
79.68
62.02
100
79.47
61.25
100
20.16
13.71
100
58.62
29.52
75
2 θ (degrees)
27.05
32.74
24.89
26.64
32.47
42.98
29.42
39.44
36.00
26.24
27.25
36.15
800
0
20
Integrated Intensity (cps deg)
30
Ca C O3
Ca C O3
Ca C O3
Ca C O3
Ca C O3
200
04_15_2011
Ca C O3
400
Ca C O3
Intensity (cps)
600
Ca C O3
Ca C O3
Meas. data:sample_04_15_11/Data 1
Ca C O3
40
50
40
50
Ca C O3
500
400
300
200
100
0
20
30
2-theta (deg)
0
20
50
40
50
400
0
20
30
2-theta (deg)
0
20
Ca C O3
Ca C O3
50
30
Ca C O3
09_09_2011
40
50
40
50
Ca C O3
100
40
0
20
30
2-theta (deg)
Figure 3.1. Powder X-ray diffraction spectra of trials containing 10 mM CaCl2. Trial
04_15_2011 and 09_09_2011 were determined as vaterite, syn based on peaks at 27.05° and
32.74° (Kamhi, 1963). The trial for 04_18_2011 was determined to be vaterite based on
peaks at 26.64° and 32.47° (Wang and Becker 2009).
76
Ca C O3
Ca C O3
Ca C O3
Ca C O3
Ca C O3
40
800
Ca C O3
Intensity (cps)
30
Ca C O3
100
Integrated Intensity (cps deg)
04_18_2011
Ca C O3
200
Ca C O3
400
Ca C O3
600
Ca C O3
Integrated Intensity (cps deg)
Intensity (cps)
800
Ca C O3
1000
Integrated Intensity (cps deg)
0
20
30
40
50
40
50
Ca C O3
Ca C O3
Ca C O3
100
Ca C O3
Ca C O3
04_05_2011
Ca C O3
200
Ca C O3
300
Ca C O3
Intensity (cps)
400
Ca C O3
400
200
0
20
30
2-theta (deg)
30
50
40
50
Ca C O3
Ca C O3
Ca C O3
Ca C O3
40
Ca C O3
400
0
20
30
0
20
Ca 800
C O3
30
40
50
40
50
400
0
20
30
2-theta (deg)
Figure 3.2. Powder X-ray diffraction spectra of trials containing 5 mM CaCl2. Trial
04_05_2011 was determined as vaterite, syn based on peaks at 27.05° and 32.74° (Kamhi,
1963). The trials for 04_08_2011 and 04_11_2011 were determined to be vaterite based on
peaks at 26.64° and 32.47° (Wang and Becker 2009).
77
Ca C O3
Ca C O3
200
Ca C O3
Ca C O3
04_11_2011
Ca C O3
400
Ca C O3
600
Ca C O3
2-theta (deg)
Ca C O3
Intensity (cps)
Ca C O3
Ca C O3
0
20
800
Integrated Intensity (cps deg)
04_08_2011
Ca C O3
200
Ca C O3
400
Ca C O3
Integrated Intensity (cps deg)
Intensity (cps)
600
Ca C O3
800
1200
Meas. data:sample_05_12_11/Data 1
1000
05_12_2011
Intensity (cps)
800
600
400
200
Integrated Intensity (cps deg)
0
20
30
40
50
2.0e+004
1.5e+004
1.0e+004
5.0e+003
0.0e+000
20
30
40
50
2-theta (deg)
1000
Meas. data:sample_05_16_11/Data 1
800
05_16_2011
Intensity (cps)
600
400
200
Integrated Intensity (cps deg)
0
20
30
40
50
40
50
600
400
200
0
20
30
Integrated Intensity (cps deg)
Meas. data:05_26_11/Data 1
Ca C O3
05_26_2011
500
0
20
Ca
C O3
1200
25
Ca C O3
Ca C O3
Ca C O3
Intensity (cps)
1000
Ca C O3
2-theta (deg)
30
35
30
35
600
0
20
25
2-theta (deg)
Figure 3.3. Powder X-ray diffraction spectra of trials containing 1 mM CaCl2. Trial
05_26_2011 was determined as vaterite, syn based on peaks at 27.05° and 32.74° (Kamhi,
1963). Trials 05_12_2011 and 05_16_2011 did not detect a calcium carbonate polymorph
present, however it is obvious from the peaks present that these forms are also vaterite.
78
0
20
30
Ca C O3
Ca C O3
Ca C O3
Ca C O3
Ca C O3
07_18_2011
40
50
40
50
Ca C O3
2000
1000
0
20
30
500
0
20
30
Ca C O3
Ca C O3
07_19_2011
1000
Ca C O3
Ca C O3
Ca C O3
Ca C O3
Meas. data:07_19_11/Data 1
Ca C O3
Ca C O3
2-theta (deg)
1500
40
50
40
50
Ca C O3
1000
0
20
30
2-theta (deg)
500
30
Ca C O3
07_21_2011
Ca C O3
Meas. data:07_21_11a/Data 1
Ca C O3
Ca C O3
Ca C O3
Ca C O3
0
20
Ca C O3
Intensity (cps)
1000
Integrated Intensity (cps deg)
1500
Ca C O3
Intensity (cps)
Ca C O3
Ca C O3
Ca C O3
1000
2000
Integrated Intensity (cps deg)
Meas. data:07_18_11/Data 1
Ca C O3
Ca C O3
Integrated Intensity (cps deg)
Intensity (cps)
2000
40
50
40
50
Ca C O3
1000
0
20
30
2-theta (deg)
Figure 3.4. Powder X-ray diffraction spectra of trials containing 1 mM CaCl2 and 0.5 mM
aspartic acid. All trials were determined to be vaterite, syn based on peaks at 27.05° and
32.74° (Kamhi, 1963).
79
30
Ca C O3
40
50
40
50
Ca C O3
2000
1000
0
20
30
06_28_2011
1000
30
Ca C O3
Ca C O3
Meas. data:06_28_11/Data 1
Ca C O3
Ca C O3
Ca C O3
Ca C O3
Ca C O3
Intensity (cps)
Integrated Intensity (cps deg)
0
20
Ca C O3
2-theta (deg)
2000
40
50
40
50
Ca C O3
2000
0
20
30
2-theta (deg)
3000
Intensity (cps)
Ca C O3
500
3000
Integrated Intensity (cps deg)
06_14_2011
Ca C O3
Ca C O3
1000
Ca C O3
Meas. data:06_14_11_x2/Data 1
Ca C O3
Ca C O3
0
20
Ca C O3
Intensity (cps)
1500
Integrated Intensity (cps deg)
2000
Ca C O3
2500
Meas. data:06_30_11/Data 1
06_30_2011
2000
1000
0
20
2.0e+004
30
40
50
40
50
1.0e+004
0.0e+000
20
30
2-theta (deg)
Figure 3.5. Powder X-ray diffraction spectra of trials containing 1 mM CaCl2 and 0.5 mM
succinic acid. Trials from 06_14_2011 and 06_28_2011 both detected vaterite, syn based on
peaks at 27.05° and 32.74° (Kamhi, 1963). The trial for 06_30_2011 did not detect a
calcium carbonate polymorph present.
80
3000
Integrated Intensity (cps deg)
Intensity (cps)
Meas. data:06_02_11/Data 1
2000
06_02_2011
1000
0
20
6.0e+004
30
40
50
40
50
3.0e+004
0.0e+000
20
30
2-theta (deg)
2500
Meas. data:06_07_11/Data 1
Integrated Intensity (cps deg)
Intensity (cps)
2000
06_07_2011
1500
1000
500
0
20
25
30
35
30
35
2000
0
20
25
2-theta (deg)
3000
Integrated Intensity (cps deg)
Intensity (cps)
Meas. data:06_09_11/Data 1
2000
06_09_2011
1000
0
20
25
30
35
30
35
2000
0
20
25
2-theta (deg)
Figure 3.6. Powder X-ray diffraction spectra of trials containing 1 mM CaCl2 and 0.5 mM
citric acid. No calcium carbonate polymorph was detected.
81
Meas. data:blank/Data 1
( C13 H8 O S )n
( C13 H8 O S )n
( C13 H8 O S )n
0.0e+000
20
)n
H8OOSS)n
C13H8
((C13
( C13 H8 O S )n
( C13 H8 O S )n
Intensity (cps)
4.0e+003
Integrated Intensity (cps deg)
6.0e+003
2.0e+003
blank
30
40
50
40
50
( C13 H8 O S )n
1.0e+005
5.0e+004
0.0e+000
20
30
2-theta (deg)
Figure 3.7. Powder X-ray diffraction spectra of the polycarbonate material on which the
crystals were grown.
A
B
C
Figure 3.8. Visual representations of the three crystalline CaCO3 polymorphs (A)
rhombohedral calcite Meldrum (2010), (B) elongated aragonite Jorgensen (1976), and (C)
spheroidal vaterite Sikorski et al. (2010).
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Table 3.2. Comparison of morphologies detected via x-ray diffraction and scanning electron
microscopy and the number of trials (in parentheses) in which that polymorph was detected.
Solution
10 mM CaCl2
XRD
Vaterite (3)
SEM
Vaterite & Calcite (2)
Vaterite (1)
5 mM CaCl2
Vaterite (3)
Vaterite (3)
1 mM CaCl2
Vaterite (3)
Vaterite (2)
Vaterite & Calcite (1)
1 mM CaCl2 + 0.5 Citric
None identified (3)
Vaterite & Calcite (3)
Vaterite (3)
Acid
Vaterite (2)
None identified (1)
1 mM CaCl2 + 0.5 Aspartic
Vaterite (3)
Vaterite (3)
Acid
1 mM CaCl2 + 0.5 Succinic
Acid
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a
b
c
d
e
f
Figure 3.9. Representative SEM images of common morphologies observed in crystal
growth trials using a growth solution of 10 mM CaCl2. (a) 500X and (b) 3000X Trial
04_15_2011: vaterite. (c) 500X and (d) 1000X Trial 04_18_2011: vaterite and calcite. (e)
500X and (f) 3000X Trial 09_09_2011: vaterite and calcite. All trials used 5 g ammonium
carbonate + 20 mL deionized water.
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a
b
c
d
e
f
Figure 3.10. Representative SEM images of common morphologies observed in crystal
growth trials using a growth solution of 5 mM CaCl2. (a) 500X and (b) 3000X Trial
04_05_201: vaterite. (c) 500X and (d) 1200X Trial 04_08_2011: vaterite. (e) 500X and (f)
1600X Trial 04_11_2011: vaterite. All trials used 5 g ammonium carbonate + 20 mL
deionized water.
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a
b
c
d
e
f
Figure 3.11. Representative SEM images of common morphologies observed in crystal
growth trials using a growth solution of 1 mM CaCl2. (a) 500X and (b) 5000X Trial
05_12_201: vaterite and calcite. (c) 500X and (d) 110X Trial 05_16_2011: vaterite. (e)
500X and (f) 100X Trial 05_26_2011: vaterite. All trials used 5 g ammonium carbonate + 20
mL deionized water.
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a
b
c
d
e
f
Figure 3.12. Representative SEM images of common morphologies observed in crystal
growth trials using a growth solution of 1 mM CaCl2 + 0.5 mM citric acid. (a) 500X and (b)
3500X Trial 06_02_201: vaterite and calcite. (c) 500X and (d) 100X Trial 06_07_2011:
vaterite and calcite. (e) 500X and (f) 1000X Trial 06_09_2011: vaterite and calcite. All
trials used 5 g ammonium carbonate + 20 mL deionized water.
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a
b
c
d
e
f
Figure 3.13. Representative SEM images of common morphologies observed in crystal
growth trials using a growth solution of 1 mM CaCl2 + 0.5 mM succinic acid. (a) 500X and
(b) 4000X Trial 06_14_201: vaterite. (c) 500X and (d) 1500X Trial 06_28_2011: vaterite.
(e) 500X and (f) 1600X Trial 06_30_2011: vaterite. All trials used 5 g ammonium carbonate
+ 20 mL deionized water.
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a
b
c
d
e
f
Figure 3.14. Representative SEM images of common morphologies observed in crystal
growth trials using a growth solution of 1 mM CaCl2 + 0.5 mM aspartic acid. (a) 500X and
(b) 1000X Trial 07_18_201: vaterite. (c) 500X and (d) 1500X Trial 07_19_2011: vaterite.
(e) 500X and (f) 2000X Trial 07_21_2011: vaterite. All trials used 5 g ammonium carbonate
+ 20 mL deionized water.
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Chapter 4
Summary and Conclusions
90
4.1 Summary of Major Findings
This thesis investigated the effect of carboxylate containing biomolecules on the
kinetic rate coefficients and polymorphism of calcium carbonate precipitation. The aim of
this study was to elucidate how these biomolecules affect the kinetics and subsequent
structure of calcium carbonate crystals.
Constant monitoring of [Ca2+], pH, and pCO2
throughout the reaction process resulted in data that allowed for application of a kinetic
model. Post experiment analysis of solid precipitates provided insight into which polymorph
of CaCO3 had been formed.
A major conclusion of this work is that the presence of the carboxylate-containing
biomolecules used in this study increases the kinetic rate coefficient, and thus increases the
rate of calcium carbonate formation. Biomolecules did not have an effect on the phase of the
products, although citric acid consistently produced a fraction of calcite crystals. All other
crystals produced were that of the metastable form vaterite, possibly due to the presence of
ammonium ions in solution during formation.
4.2 Broader Implications of Major Findings
Vaterite is a metastable polymorph and has the ability to transform into other CaCO3
polymorphs (Becker and Hu 2009). This study has implications for carbon sequestration
techniques. Vaterite is not an ideal polymorph for sequestration, given its transient nature.
However, it may be possible to dictate which polymorph forms next by manipulating solution
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parameters. The high concentration of ammonium present in these experiments is most
likely the factor that caused the formation of vaterite. However, more work would need to be
done to prove this hypothesis.
4.3 Final Thoughts
4.3.1 Improvements for Present Work
The high concentration of ammonium in experimental conditions led to the accidental
corroboration of another hypothesis. It has been suggested that ammonium ions can promote
and even stabilize vaterite production (Becker and Hu 2009). To fully understand the effect
of biomolecules on this system it would be necessary to complete these experiments at a
much lower concentration of ammonium. The use of ammonium bicarbonate or a large
reduction in the mass of ammonium carbonate used would decrease the concentration of
ammonium in solution but still allow for the formation of aqueous carbonate and an increase
in pH. Sodium carbonate (Na2CO3) would work as a source of carbonate and as a means of
increasing pH, however these experiments utilize a gas diffusion method, and Na2CO3 would
not work in this capacity. Another option that has been utilized in peer reviewed literature is
to bubble CO2/N2 mixed gas into the aqueous CaCl2 solution.
However, vaterite is a metstable CaCO3 polymorph and it is possible that these
crystals could have been precursors to other polymorphs that may have formed had time
continued. The sensitive nature of the calcium ISE led to the termination of experiments
92
once aqueous calcium had been depleted. An interesting twist on this experiment would
have been to complete additional trials without the ISE and let them go for a set amount of
time (approximately 48 hours, similar to many other trials) and then examine the phase of the
crystals produced with SEM.
4.3.2 Suggestions for Future Work
Ocean acidification is a major concern for the future health of our world’s oceans
(Fabricius et al. 2011). Given the many important roles of calcium carbonate in marine
systems, as well as the current downward pH trends of these same systems, it would be
beneficial to examine calcium carbonate formation kinetics under marine conditions.
However, this is a difficult task since these conditions vary greatly. It would be possible to
manipulate the temperature of the system during formation, as well as to include salts which
would impact the ion activity of the system (and thus the kinetics of formation). Examples of
ions to add in order to mimic ocean conditions would be chloride, sodium, magnesium,
sulfur, and potassium. Additionally, experiments that vary the concentration of CO2 in the
system would investigate the effect of carbon dioxide on calcium carbonate formation.
The use of atomic force microscopy would be an insightful tool for any similar future
projects. AFM is a type of scanning probe microscopy that can achieve resolution at the
nanometer scale and is often used to image growth/dissolution reactions. The microscope
consists of a cantilever with a sharp tip that runs over a material and creates a topographical
93
image. The work done in this thesis has shown that carboxylate-containing biomolecules do
affect calcium carbonate precipitation by increasing the rate of formation. The use of AFM
would allow for an examination of precipitation reactions on the nanometer scale.
Additional computer modeling may also shed light on calcium carbonate growth
reactions. The use of atomistic simulation can predict and model how molecules will act
based on structural geometries and surface energies. Computer programs such as Materials
Studio allow for computation of surface free energies as well as visualizations of how and
where biomolecules will bind to crystal faces. This tool will allow further insight of the
mechanisms of biomolecule-calcium carbonate reactions.
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4.4 References
Becker, U. and Hu, Q. (2009). Controlled growth of different calcium carbonate polymorphs
as induced by the presence of dissolved molecules and mineral surfaces. Goldschmidt
Conference. June 21 -26, Davos, Switzerland.
Fabricius, K.E., Langdon, C., Uthicke, S., Humphrey, C., Noonan, S., De’ath, G., Okazaki,
R., Muehllehner, N., Glas, M.S., Lough, J.M. (2011). Losers and winners in coral reefs
acclimatized to elevated carbon dioxide concentrations. Nature Climate Change. 1, 165-169.
95