Hollow Calcium Carbonate Microsphere Formation in the Presence of

Hollow Calcium Carbonate Microsphere Formation in the Presence of Biopolymers and Additives Michael F. Butler,* William J. Frith, Christopher Rawlins, Anthony C. Weaver, and Mary Heppenstall-Butler

CRYSTAL GROWTH & DESIGN 2009 VOL. 9, NO. 1 534–545

UnileVer Corporate Research, Colworth Park, Sharnbrook, Bedfordshire, MK44 1LQ, U.K. ReceiVed July 30, 2008; ReVised Manuscript ReceiVed October 1, 2008

ABSTRACT: Directed crystallization is found in many natural systems, where the inorganic material provides the organism with superior barrier and mechanical properties. Using this principle of controlled nucleation and growth on a biopolymer (pectin) substrate, hollow shells of calcium carbonate were formed in certain reaction conditions that were generally more perfect at lower reactant concentrations. The nucleation mechanism required the presence of calcium-binding acidic groups on the biopolymer. These served to initially gel the biopolymer, thus providing the template upon which crystallization could occur, and then provided the calciumrich sites at the template surface that promoted crystallization. The propensity of the system to crystallize on the gel particle surface thereby led to the formation of a complete shell of calcium carbonate, in the form of calcite, which in certain conditions was shown to effectively encapsulate small molecular weight hydrophilic species such as sodium chloride. Importantly, however, in some cases the influence of the crystallization conditions, including the nature of the molecule to be encapsulated, significantly interfered with crystallization preventing the formation of a complete shell. Observation of hollow shells in certain conditions does not therefore guarantee that a system will be usable in all cases for capsule formation. Introduction Many natural systems, such as coccoliths, diatoms, and seashells, use the principle of directed nucleation and controlled growth of inorganic materials, such as calcium carbonate, to form skeletons or shells.1-4 The template usually takes the form of a biopolymer (often a glycoprotein) rich in acidic moieties, such as aspartic and galacturonic acid. Crystallization kinetics are manipulated via active control of the supersaturation of the crystallizing species by the organism. In addition, the organism may employ additional chemical species in solution to inhibit bulk crystallization or to promote growth of the templated crystals with a particular habit. Natural inorganic composites are employed in many cases to afford superior barrier or mechanical properties to the organism, for example, to protect it from the exterior environment or predators and are normally formed at ambient temperatures and pressures from nontoxic precursor materials. All of these properties make biomimetic systems, constructed according to the principles found in nature, very attractive from an industrial perspective. The intention of the work reported in this manuscript was therefore to use natural principles to form a material that would be an effective barrier in an aqueous environment, for the encapsulation of small, readily diffusing, hydrophilic molecules that cannot be encapsulated using conventional means. Pectin was chosen because it is a readily available calcium-binding biopolymer. It is a polysaccharide used in foodstuffs as a thickener and to form gels, which it does in the presence of calcium ions.5 Most importantly for the present study, it contains the galacturonic acid monomer unit that is known to play a role in the directed growth of inorganic calcium salts in several natural systems.6 Calcium carbonate is attractive because it is nontoxic, bioavailable, and breaks down in the acidic conditions of the stomach. A complete film of calcium carbonate should therefore provide a barrier against diffusion in conditions where it is insoluble, but will rapidly * To whom correspondence should be addressed. E-mail: Michael.Butler@ Unilever.com. Telephone: +44 (0)1234 222958. Facsimile: +44 (0)1234 248010.

lose its barrier properties at low pH when release of encapsulated material is desired. Naturally present as calcite, aragonite, and vaterite forms (in order of decreasing stability), calcite is most commonly formed in the laboratory, sometimes via a metastable vaterite form. Aragonite is formed when crystallization occurs in the presence of magnesium ions. The forms are distinguished by their different crystallographic unit cells and crystal habits, with calcite usually growing as rhombohedra, aragonite as spicules, and vaterite as spherical aggregates. Several studies exist in which calcite or aragonite has been grown on two-dimensional substrates, such as self-assembled monolayers7-11 and polymer films.12-17 In these studies, relationships between the separation of functional groups, such as carboxylic acid groups, on the surface of the film and the observed crystal unit cell and habit were formulated. In many cases, the nature of the organic surface directly determined the shape of crystals nucleated on the surface. Three dimensional substrates, such as phospholipid vesicles,18 synthetic polymer dendrimers,19-21 colloidal particles,22,23 lipid bilayer stacks,24 aqueous foams,25 water-in-oil microemulsions,26 pseudovesicular double emulsions,27 fatty acid stabilized oil droplets,28 and biopolymer gel particles29 have also been studied for their effect on the growth of calcium salts. In some cases, such as the growth of calcium phosphate on liposomes18 and calcium carbonate on foams,25 oil droplets28 and biopolymer gel particles,29 thin crystalline shells were formed. Relatively little work has been performed on the investigation of the mechanism of directed crystallization in the three-dimensional case, although in the dendrimer studies the functional groups present at the surface of the molecule influenced the crystal growth,19-21 or on the possible application of application possibilities (such as encapsulation) using hollow calcium carbonate (or phosphate) shells. Hollow silica shells have been most studied as encapsulation devices,30 although other materials, such as titanium dioxide,31,32 zinc oxide,33 iron oxide,34 copper sulfide,35 nickel,36 and silver,37 have also been shown to form hollow shells.

10.1021/cg8008333 CCC: $40.75  2009 American Chemical Society Published on Web 12/03/2008

Calcium Carbonate Crystallization

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In the studies reported herein, morphologies of calcium carbonate crystals grown in the presence of an anionic biopolymers, low methoxy (LM) pectin, will be described as an extension of our previous report on calcium carbonate crystallization in the presence of biopolymers, in which pectin was found to be particularly effective at causing the formation of hollow shells of calcium carbonate crystals.29 Electron microscopy was used to provide high resolution of the particles, whereas turbidity, pH, and rheological measurements were used to provide information on the crystallization kinetics in order to obtain an idea of the limitations of reaction/processing conditions on the formation of hollow particles, to give an idea of their usefulness as capsules. Encapsulation studies, in which crystallization was performed in the presence of small ionic (sodium chloride) and molecular (ascorbic acid and tryptophan) species, were used to test whether the hollow shells could indeed be used as capsules. Ascorbic acid was chosen for both its relevance (vitamin C) and its known ability to influence calcium carbonate crystallization. Tryptophan was chosen for its relevance and the fact that it contains amine groups that are likely to interact with the carboxylic acid groups possessed by pectin as well as possibly influence calcium carbonate crystallization. Many studies demonstrating the formation of inorganic hollow shells claim their potential use as encapsulation systems, although there are not many where a complete study, including encapsulation and controlled release measurements, has been performed. The current manuscript aims to highlight both the possibilities and limitations of calcium carbonate shells grown under the conditions reported. Experimental Section Materials and Sample Preparation. The biopolymer used in this study was LM pectin (LM12, supplied by Kelco, with a 35% degree of esterification (DE)). For confocal laser scanning microscopy experiments to determine the location of pectin in the crystallized system, the pectin was labeled with fluorescein isothiocyanate by dissolving fluorescein isothiacynate in sodium bicarbonate solution and reacting with unlabeled pectin. The reagents used for forming calcium carbonate were calcium chloride, sodium bicarbonate, and potassium hydroxide (all reagent grade supplied by Sigma Chemical Company). Other materials used in the study were sodium chloride, ascorbic acid, and L-tryptophan (reagent grade, supplied by Sigma Chemical Company). For the optical and electron microscopy studies, control samples of calcium carbonate were made by mixing 10 mL of a 0.01 M calcium chloride solution with 10 mL of a 0.02 M sodium bicarbonate solution, in the absence of any other additivess. Precipitation of calcium carbonate crystals was triggered by the dropwise addition of 5 M potassium hydroxide to the mixture to raise the pH to 10.5 ((0.1). To prepare calcium carbonate crystals in the presence of pectin or sodium alginate, 10 mL of x M aqueous calcium chloride was added to 10 mL of a solution containing 2x M aqueous sodium bicarbonate, where x was varied between 0.005 and 0.4 M and y% (w/v) biopolymer, where y was varied between 0.4 and 3.2 in a 30 mL bottle, while continuously stirring. Potassium hydroxide was added dropwise with a micropipette until the pH reached 10.5 ((0.1). Calcium carbonate is produced from the following reaction:

|

H2CO3

+ CaCl2 + 2NaHCO3 w CaCO3 + 2NaCl + HCO3 + H + CO23 + 2H

where the pH of the solution determines the configuration of the carbonate ions in solution. Above pH 10.5 carbonate ions become the dominant species. After a minimum time of 48 h of continuous stirring, the solution was centrifuged in 40 mL tubes for 10 min at 3300 rpm. The supernatant was removed and the crystals were dispersed in 20 mL of deionized water. More detailed investigations of the crystallization process were made using rheology, turbidity, and titration

measurements: sample preparation for these studies is given in more detail in the appropriate section. To test for incorporation and retention of ionic species in the hollow shells, sodium chloride was dissolved in both 10 mL of 0.01 M calcium chloride and 10 mL of 0.02 M sodium bicarbonate, to a concentration of 1% by weight. LM pectin was added to the latter solution and dissolved, as before. The pH was adjusted to 10.5 ((0.1) by adding 5 M potassium hydroxide dropwise with a micropipette. In order to study the interactions of sodium chloride with calcium carbonate without biopolymer present the above method was perfomed in the presence of sodium chloride but the absence of biopolymer. To test for incorporation of molecular species, different concentrations of ascorbic acid and L-tryptophan were included in the pectin, calcium chloride, and sodium bicarbonate solutions as described above for sodium chloride incorporation. Transmission Optical Microscopy. Transmission optical microscopy (Leitz Diaplan, set up for Kohler illumination) was used to determine the crystal habit of the calcium carbonate crystals, that had precipitated and aged for three days, in the control samples and in the presence of LM pectin and sodium alginate. The crystals were obtained by centrifuging the suspension in which they formed and washing with 0.1 M sodium hydroxide solution (to prevent dissolution during washing) three times. A drop of the final, cleaned, suspension was placed on a microscope slide beneath a standard glass coverslip and observed under bright-field conditions with and without crossed polarizers. Confocal Laser Scanning Microscopy (CLSM). CLSM images of the crystallized system containing 0.4% (w/v) fluorescently labeled pectin were obtained using a Leica TCS SP1 spectral scanning confocal microscope equipped with an oil immersion objective (magnification × 60, numerical aperture 1.4). Images were acquired simultaneously in transmitted light brightfield and fluorescence modes. Laser excitation was with an argon ion laser operating at 488 nm. Emission was collected in the range from 500-550 nm. The samples were either imaged in aqueous solution under the crystallization conditions or imaged after being washed and dried on the glass slide. Electron Microscopy. Scanning transmission electron microscopy (STEM) was used to study the calcium carbonate crystal morphology over a range of pectin, calcium chloride, and sodium biocarbonate concentrations. One drop of a suspension of crystals obtained after centrifugation was placed onto a carbon coated copper grid. The drop was allowed to air-dry and then sputter coated with a 10 nm layer of gold/palladium metal. The treated crystals were examined in a JEOL1200EX electron microscope with an ASID10 scanning attachment at a range of accelerating voltages between 20 and 120 kV. In addition, a back-scattered X-ray detector was used to perform a compositional analysis of the crystals, both with and without potentially encapsulated substances. In order to improve the spectrum, sample holders with low and high copper content were experimented with. The images and data were recorded using Oxford Instruments INCA software. Scanning electron microscopy was used to to study the calcium carbonate crystal morphology over a range of pectin, calcium chloride, and sodium biocarbonate concentrations and to examine calcium carbonate crystals after microtoming. Suspensions of crystals were first washed with water to remove any extraneous material and resuspended in water. For the microtoming experiments, a drop of the sample was placed in a small cryo “top hat” holder and then plunge frozen in liquid nitrogen. The frozen sample was then transferred to a Cressington CFE50 freeze facture machine and the top of the frozen sample microtomed using the swinging knife. After several passes of the knife, the sample was removed from the machine, allowed to thaw and the drop of sample spread onto a plasma glowed 25 mm Thermanox coverslip. The sample was allowed to dry and the coverslip was mounted on a 25 mm SEM stub. The sample was then sputter coated with gold/palladium (∼25 nm thickness) and examined using a JEOL JSM 6060 scanning electron microscope operated at 25-35 kV and the sample stage tilted to 45°. pH Measurement. The crystallization process was followed by measuring the pH as the reaction proceeded in a series of controlled titrations for samples with defined CaCl2, NaHCO3, and pectin concentrations. In the first series of titrations, drops of CaCl2 solution were titrated into 25 mL of NaHCO3 solution, with a CaCl2/NaHCO3 concentration ratio of 1:2. Concentrations of CaCl2 used (in M) were 0.0050, 0.0070, 0.0085. 0.0086, 0.0087, 0.0088, 0.0089, 0.0090, 0.0100,

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Figure 1. Scanning tranmisssion electron micrographs of calcium carbonate formed (a) in the presence of 0.4% (w/v) LM pectin, (b) in the absence of LM pectin, showing “rosette”-like crystals in the former case and rhombohedral crystals in the latter case. and 0.0200. The pH of the stirred mixture was continuously monitored, and the final, stabilized pH was noted. No more than 2 min elapsed between adding the drops. The second series of titrations were similar to the first, but with the addition of pectin to the NaHCO3 solution to a concentration of 0.2% (w/v). The titration commenced immediately after the addition of pectin to the NaHCO3 solution. The concentrations of CaCl2 used were 0.010, 0.020, 0.025, 0.030, 0.035, 0.040, 0.045, 0.050, and 0.075. The CaCl2/ NaHCO3 concentration ratio was 1:2, as before. The third series of titrations were similar to the second, but in this series the time before the titration commenced after the addition of pectin to the NaHCO3 solution was varied. In order to test for the influence of atmospheric CO2 on the reaction equilibria in the crystallization conditions used, a solution was made up to 0.2 M NaHCO3 and 0.2% w/v LM pectin. The pH was measured continuously from the time of addition. Three mixes were made, one container was kept open, one container had a lid put on it to reduce the headspace to a relatively small volume and the last was purged with nitrogen gas to remove carbon dioxide from the solution. Finally, to test for the influence of stirring time of pectin and NaHCO3, solutions of 0.2% w/w pectin and 0.06 M NaHCO3 were made and left stirring for three different amounts of time: none, 1 h, and overnight. 0.03 M CaCl2 was then titrated into 25 mL of the mix and the pH was recorded. Turbidity. In situ crystallization experiments were performed using an ultraviolet/visible wavelength (UV/vis) spectrophotometer (PerkinElmer Lambda 40) to measure the absorption of 300 nm wavelength light at 25 °C as a function of time. Two sample preparation conditions were used. In the first method, a volume of 0.4 M NaHCO3 was added to a 5 mL of 0.02 M CaCl2. The mixture was shaken and about 3 mL was transferred into a PMMA cuvette, placed in the spectrophotometer and the turbidity measurement was started. The time between solution addition and starting the measurement was noted so that the data could be plotted using the time of mixing as the zero time. In the second method, 5 mL CaCl2 was added to a 5 mL solution of NaHCO3 of known concentration and 0.2% w/w pectin. The concentration of CaCl2 was half-that of the NaHCO3. The solution was shaken and about 3 mL was transferred to a PMMA cuvette, placed in the spectrophotometer and the turbidity measurement was started. The time between solution addition and starting the measurement was noted so that the data could be plotted using time of mixing as zero time. In both methods, each volume ratio was repeated three times. Viscosity. Viscosity measurements were made at six different stages of the crystallization process using a Haake stress controlled viscometer equipped with a cone and plate geometry (60 mm diameter, cone angle 5°) in order to follow the structural changes to the biopolymer upon addition of calcium ions and the formation of crystalline calcium carbonate particles. In the first stage, a 0.8% (w/v) pectin solution was diluted to 0.05, 0.1, 0.2, and 0.4% (w/v) samples respectively using water. In the second stage, a 0.8% (w/v) pectin solution was diluted to 0.05, 0.1, 0.2,

and 0.4% (w/v) samples respectively using 0.02 M NaHCO3. In the third stage, a 0.5% (w/v) pectin solution was diluted to 0.05, 0.1, 0.2, and 0.4% (w/v) samples respectively using 0.02 M NaHCO3 and 0.01 M CaCl2. In the fourth stage, the pH of the mixture from stage 3 was adjusted to pH (10.5 ( 0.1) by the addition of 1.0 and 0.1 M NaOH. The sample was left stirring gently for at least 48 h to complete the crystallization process. In the fifth stage, the mixture from stage 4 was centrifuged and the supernatant and sediment were separated. In the sixth stage, the mixture from stage 4 was acidified to pH (2.0 ( 0.1) using 2 M HCl. Aliquots of the sample were taken at the various stages of the preparation of the product and placed in the viscometer. Excess sample was removed and the sample was covered with a vapor retention hood to minimize evaporation. Controlled stresses, with values between 0.01 and 10 Pa, were applied to the sample, and the viscosity was measured as a function of applied stress. In addition, the time dependence of the sample rheology was monitored in one of three ways: (1) aliquots were taken from a reaction mixture and placed in the rheometer at intervals, whereupon the flow curve was measured; (2) samples were prepared and placed in the rheometer and the flow curve of that sample was then recorded over a range of times; (3) samples were prepared and placed in the rheometer, a steady shear stress was applied and the shear rate was recorded as a function of time.

Results Electron Microscopy. Figure 1 shows scanning transmission electron microscope (STEM) images of calcium carbonate crystals formed with and without added LM pectin. With no additive present, conventional calcite rhombohedra formed. When biopolymer was present in the mixture, however, radially arranged “rosette-like” spherical or dumb-bell-shaped aggregates of rhombohedra formed, whose level of organization depended on the biopolymer present. Previous work has shown that these crystals are comprised of the calcite polymorph that formed at the onset of crystallization. The degree of radial ordering was shown previously in this type of morphology by the presence of Maltese crosses present in optical micrographs of samples containing crystalline aggregates, placed between crossed polarizers.29 The crystallization kinetics were also different for crystals formed in the presence of LM pectin compared to the control samples. Whereas crystallization was immediate in the control samples, there was an induction time before the crystals grew in the presence of the biopolymers. The induction time was dependent on the concentration of the LM pectin. For solutions containing 0.2% (w/v) LM pectin, crystallization was observed, by turbidity measurement, to begin after 46.0 ( 6.9 min. At an



Figure 2. Electron micrographs showing evidence for the prevalence of hollow calcium carbonate shells formed in the presence of 0.4% (w/v) LM pectin: (a-c) different shells from a mechanically disrupted sample, (d) scanning electron micrograph of a calcium carbonate “rosette-like” aggregate formed in the presence of 0.4% (w/v) LM pectin, then washed in water, frozen, and microtomed.

LM pectin concentration of 0.5% (w/v) the induction time was measured to be 30.0 ( 12.5 min. Some of the electron micrographs suggested that at least some of the as-formed crystalline rosette-like aggregates were in the form of hollow shells of calcite formed from closely arranged individual crystals. After subjecting the sample to crushing in a press, many more broken particles were obtained, revealing them to consist of hollow shells of calcium carbonate with a thickness of approximately 0.5 µm, shown in Figure 2. In some cases, also shown in Figure 2, the hollow shells contained an internal crystal growth. Figure 2d shows a scanning electron microscope (SEM) image of a rosette-like aggregate that had been washed in water, frozen and then microtomed, also revealing it to be a hollow shell rather than a solid object. To test whether all of the rosette-like crystal aggregates were hollow shells of calcium carbonate, compositional maps were made of many particles, chosen at random, from the measurement of the energy spectrum of back-scattered X-rays detected in the STEM. A typical calcium map of a particle, using the size of the calcium peak in the X-ray spectrum to identify the local concentration of calcium, is shown in Figure 3b along with the corresponding image of the particle (Figure 3a). An excess of calcium at the edge of the particle is clearly visible, demonstrating that the particle was indeed hollow and consisted of a shell of calcium carbonate approximately 0.5-1 µm thick.

The morphology observed was dependent on pectin and reactant concentrations. For a given pectin concentration, the size of the aggregates was approximately the same up to a sodium bicarbonate concentration around 0.10 M. As the pectin concentration increased the average size of the aggregates decreased (although a propensity for a bimodal distribution of particle size appeared to occur at and above 0.4% w/v pectin, particularly at the higher reactant concentrations). Rosette-like aggregates with greater perfection were generally observed at lower reactant and pectin concentrations. At pectin concentrations greater than 0.4% w/v and sodium bicarbonate concentrations greater than 0.1 M the rosette morphologies deteriorated to the extent that non-rosette-like aggregates of rhombohedra formed. Confocal Laser Scanning Microscopy. Figure 4a,b shows the transmission optical microscopy image and associated fluorescence image obtained from a suspension of rosette-like aggregates formed in the presence of fluorescently labeled pectin at a concentration of 0.4% (w/v). While there is a fluorescent background, indicating that some of the pectin was uniformly dispersed throughout the sample, there are clear concentrations of fluorescence that coincide with the position of the crystalline aggregates. A higher magnification image of washed crystals formed in the presence of 0.4% (w/v) pectin, to highlight the


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Figure 3. (a) Image of “rosette-like” calcium carbonate formed in the presence of 0.4% (w/v) LM pectin, (b) calcium density map of the same particle, showing that the particle was formed from a hollow shell of calcium carbonate.

Figure 4. Confocal laser scanning microscope images of “rosette-like” aggregates formed in the presence of 0.4% (w/v) pectin. (a) Bright-field transmission image, (b) fluorescence image, showing the presence of pectin in solution throughout the image as well as concentrated regions of pectin coincident with the positions of the crystalline aggregates (the image width is 794 µm), (c) higher magnification bright field transmission image of washed “rosette-like” aggregates formed in the presence of 0.4% (w/v) pectin, (d) corresponding fluoresence image. A dumb-bell-shaped aggregate, showing two regions of pectin at either end of the dumb-bell, is arrowed. (The image width is 62.5 µm.)

concentrated fluoresent regions is shown in Figure 4c,d. It is clear that the fluorescently labeled pectin is concentrated in the centers of the aggregates (in some cases the aggregates formed

as dumb-bells, and the two concentrated regions of pectin can be seen in each end of the dumb-bell), although some is dispersed in the shell.



Figure 5. Titration of CaCl2 into 25 mL of NaHCO3 ([NaHCO3] ) 2x [CaCl2]). (a) In the absence of pectin, (b) in the presence of 0.2% (w/v) pectin (sample made immediately before titration).

Figure 6. Variation in pH with time after addition of NaHCO3 to a solution containing 0.2% (w/v) LM pectin.

pH Measurement. Figure 5a shows the pH measured during titrations of CaCl2 into 25 mL of NaHCO3 in the absence of pectin at a CaCl2/NaHCO3 molar ratio of 1:2, for a range of CaCl2 concentrations. Above CaCl2 concentrations of 0.0087 M the pH gradually decreased, with a change in the rate of decrease at volumes of CaCl2 between 10 and 20 mL, corresponding to a pH around 8.25. The point at which the rate of change of pH with CaCl2 concentration changed coincided with the visual observation of precipitation. Below 0.0087 M no change in the gradient and no precipitation was observed. When pectin was present in the system, shown in Figure 5b for the titration of CaCl2 into a fresh mixture of NaHCO3 and 0.2% pectin, no precipitation or change in gradient of the pH-CaCl2 curve was observed until the CaCl2 concentration was 0.035 M. Precipitation occurred around pH values of about 7.7. During the course of the investigations, the sample conditions (in particular the headspace above the solution and the duration of stirring of the NaHCO3/pectin mixture prior to CaCl2 addition) were shown to be important. As shown in Figure 6, after NaHCO3 was added to a solution containing 0.2% (w/v) LM pectin the pH increased with time during stirring at a rate, and to a final value, dependent on headspace conditions. Low partial pressures of CO2 in the headspace in the nitrogen-purged sample resulted in a much higher solution pH. The sample with the smallest headspace reduced the amount of CO2 that could be removed from solution compared to an open sample, resulting in the smallest change in pH. Figure 7 shows the titration curve for 0.03 M CaCl2 added to a 0.06 M NaHCO3/0.2% pectin mixture that had been left for different amounts of time prior to the titration. The amount of stirring time prior to titration affected the starting and final pH and also whether precipitation occurred on addition of CaCl2. In this case, no precipitation occurred in the freshly prepared NaHCO3/pectin mixture that had the lowest initial pH, whereas precipitation and the change

Figure 7. Titration of CaCl2 into 25 mL NaHCO3 containing 0.2% (w/v) LM pectin ([NaHCO3] ) 2x [CaCl2]), where the NaHCO3/LM pectin mixture had been left for different periods of time prior to titration.

Figure 8. The variation in absorbance of light, at a wavelength of 300 nm, with time for equal volumes of CaCl2 mixed with (NaHCO3 + 0.2% LM pectin) ([NaHCO3] ) 2x [CaCl2]).

in gradient in the titration curve did occur in the samples left for 1 h and overnight (at a pH of about 8.4). Turbidity. Figure 8 shows the variation in absorbance at 300 nm with time for equal volumes of CaCl2 and a fresh NaHCO3/ 0.2% (w/v) pectin mixture with a 1:2 CaCl2/NaHCO3 molar ratio. Upon addition of CaCl2 the absorbance increased to a constant value for CaCl2 concentrations up to 0.035 M. Above 0.035 M CaCl2 the absorbance increased, reached a maximum, and then decreased again to a constant value that was higher than the starting value. The maximum in absorbance was reached at earlier times as the CaCl2 concentration increased. Visual inspection of the latter samples after the experiment was over showed that crystallization and sedimentation had occurred. The former samples contained no evidence of crystallization, and the gradual turbidity increase was not related to crystallization.


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Figure 9. State diagram in terms of (a) CaCl2 and LM pectin concentrations, for CaCl2 titrated into a fresh mix of NaHCO3 and LM pectin, (b) CaCl2 concentration and time before titration into the NaHCO3/LM pectin mixture for a fixed 0.2% (w/v) LM pectin concentration, showing where calcium carbonate crystallization occurred.

Figure 10. (a) Flow curve for pectin solutions of different concentration with added CaCl2, showing the existence of a yield stress. (b) Evolution of the flow curve with time for a mixture of 0.2% (w/v) LM pectin, CaCl2, and NaHCO3 pH adjusted to 10.5.

Figure 11. (a) Electron micrographs of calcium carbonate crystals grown in the presence of 0.4% (w/v) pectin and 1% (w/v) sodium chloride, (b) in the presence of 1% (w/v) sodium chloride.

The titration and turbidity data were used to construct a state diagram shown in Figure 9a, in terms of the CaCl2 and LM pectin concentrations, for the conditions at which crystallization occurred, for the situation when the NaHCO3/LM pectin mixture was made immediately before addition of CaCl2. Higher concentrations of CaCl2 were required to cause calcium carbonate crystallization as the concentration of pectin in the solution increased. Figure 9b shows the state diagram in terms of the CaCl2 concentration and the stirring time of the NaHCO3/LM pectin mixture before titration, for a fixed LM pectin concentration of 0.2% (w/v). With increasing time before titration a greater amount of CaCl2 was needed to cause calcium carbonate

crystallization after titration of CaCl2 into the NaHCO3/LM pectin mixture. Viscosity. Solutions of pectin in water and NaHCO3 displayed a Newtonian response, shown by the generally linear dependence of viscosity on pectin concentration. Figure 10a shows the flow curves (plotted as viscosity versus shear stress) for a range of pectin concentrations upon adding CaCl2. The flow curves showed the existence of a yield stress for all pectin concentrations used (that was most apparent at the higher concentrations), indicating that the presence of calcium ions caused gelation of the system. Bulk gelation was confirmed by visual inspection


Figure 12. (a) Electron micrograph of “rosette”-like calcium carbonate grown in the presence of 0.4% (w/v) LM pectin and 1% (w/v) sodium chloride, with (b) calcium and (c) chloride density maps of the same particle, showing that the particle was a hollow shell of calcium carbonate containing encapsulated sodium chloride.

of the sample upon removal from the rheometer at the end of the experiment. With time a decrease in viscosity and yield stress was measured. When NaOH was added to induce crystallization there was a significant drop in the viscosity. The dynamics of this lowering are shown in the flow curves in Figure 10b for a 0.2% (w/v) pectin sample, which shows that the change in rheology of the sample occurred within the first 15 min after addition of NaOH. Subsequent reacidification of the final reaction mixture to dissolve the calcium carbonate caused the viscosity of the solution to rise once again, indicating that the amount of calcium in solution was the primary determinant of the presence of gelled pectin.


Encapsulation. When calcium carbonate was formed in the presence of 1% (w/v) sodium chloride as well as LM pectin, rosette-like crystalline aggregates, shown in Figure 11a, were obtained that were similar in appearance to those formed in the presence of pectin only. X-ray spectra from points within the “rosettes” formed in the presence of pectin clearly showed the presence of peaks from sodium and chlorine as well as a calcium peak from calcium carbonate. In the absence of pectin, but presence of sodium chloride, modified rhombohedra, but no rosettes, were obtained, shown in Figure 11b. Calcium maps of calcium carbonate aggregates formed in the presence of LM pectin and sodium chloride, shown in Figure 12 along with the corresponding image of the particle, demonstrated that, as for the samples containing biopolymer only, they were hollow shells of calcium carbonate. Furthermore, sodium and chlorine maps, also shown in Figure 12, detected the presence of sodium chloride within these hollow shells, even after storage of the particles in pure water for 1 week. That is, they demonstrated that these hollow calcium carbonate shells had encapsulated and retained ionic species. To confirm that the sodium and chlorine peaks were indeed from sodium and chlorine located inside the hollow calcium carbonate particles, measurement of the intensity of the sodium and chlorine peaks was performed with increasing accelerating voltage, which corresponds to increasing penetration into the sample. The sodium and chlorine signals increased with increasing accelerating voltage, thereby demonstrating that the sodium chloride was located within the calcium carbonate shell. The sodium and chlorine maps did show, however, that in some cases the sodium chloride was unevenly distributed within the hollow shells of calcium carbonate. Results with the molecular species were very different. Figures 13a shows a control sample of calcium carbonate grown in the presence of 20% w/v ascorbic acid and Figure 13b shows the results of calcium carbonate crystallization on a gelled LM pectin template in the presence of 20% w/v ascorbic acid. In this case, no hollow rosette-like aggregates formed and the system was clearly unsuitable for encapsulation. In the case of L-tryptophan, that contains basic groups capable of interacting with the carboxylic acid groups present on the biopolymer and forming a complex, the formation of rosette-like aggregates was completely prevented unless very low concentrations of added L-tryptophan were present. Figure 14 shows electron micrographs of calcium carbonate crystals grown in the presence of three concentrations of L-tryptophan: 1.0, 0.1, and 0.01% (w/ v). Rosette-like aggregates of calcium carbonate crystals were only obtained at the lowest concentration of L-tryptophan. At the higher concentrations, “stack-like” aggregates, reminiscent of those obtained in the presence of nongelling biopolymers containing carboxylic acid groups such as xanthan and gellan29 were obtained. The possibility that L-tryptophan affected calcium carbonate growth was discounted, since the calcium carbonate crystals that formed in the presence of L-tryptophan were not significantly different from the control samples containing no additive. These observations suggested that the LM pectin still affected the calcium carbonate crystallization even when it was unable to form a gelled particle. Discussion The combined results from electron microscopy and confocal laser scanning microscopy clearly showed that the rosette-like aggregates were, in fact, hollow shells of calcium carbonate formed around a pectin core. The electron micrographs of both mechanically damaged and microtomed whole aggregates and


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Figure 13. Electron micrographs of (a) calcium carbonate crystals grown in the presence of 20% (w/v) ascorbic acid, (b) calcium carbonate crystals grown in the presence of 0.4% (w/v) LM pectin and 20% (w/v) ascorbic acid.

the elemental maps of whole aggregates revealed shells of calcium carbonate, whereas the fluorescence images revealed a concentrated region of LM pectin at the center of each aggregate around which the calcium carbonate had grown. Furthermore, the elemental maps of the aggregates showed that, for LM pectin at least, hollow shells were a common morphology rather than being a rarity. The mechanism of formation of these hollow shells is proposed to be a consequence of the interplay between the biopolymer gelation and calcium carbonate crystallization kinetics, and is shown schematically in Figure 15. It is believed that the pectin initially forms physically cross-linked gel particles upon mixing the polymer-containing solutions with the calcium chloride solution, with sizes on the order of a few micrometers. Pectin and alginate are natural block copolymers, with one block that has a propensity to bind to calcium ions. In the presence of calcium ions, calcium-binding blocks on different biopolymer molecules are linked in short, relatively ordered, stretches to form physical cross-links and a gel is formed within minutes (as shown by the viscosity results). At low pectin and calcium ion concentrations the gels that form were initially relatively weak and therefore easily broken into small particulate gel pieces under shear. The calcium ions that are bound in the gel are then proposed to act as nuclei, or regions of high local supersaturation, that favor calcium carbonate nucleation and therefore direct growth when the bulk supersaturation conditions are conducive for crystallization (i.e., when the pH is raised to 10.5). This latter process occurs over longer time-scales than the formation of the gel templates, shown by the polymer concentrationdependent induction times. The induction time, τ, can be mathematically expressed as38

log τ ∝

(

βυ2γ2s

)

1 (2.303kBT) (log Ω)2 3

where β is a shape factor for the calcite nuclei () 16π/3 for spherical shapes), ν is the molar volume of calcite ()1.89 × 10-5 m-3), γs is the surface energy of the calcite nuclei, kB is Boltzmann’s constant, Ω is the solution saturation and T is the temperature. As the polymer concentration changes the two factors that may be affected are the solution supersaturation, as the polymer binds more calcium from solution, and the nucleus surface energy, as the polymer may become associated with the growing crystal. Since an increase in pectin concentration

reduces the supersaturation of calcium ions in solution by binding an increasing quantity of it, the decrease in induction time with increasing pectin concentration may be explained by a decrease in surface energy that is increasingly favorable for nucleation. Although not tested in the currrent study, one possible explanation for this would be if the separation between the templating carboxylate groups became increasingly matched to the lattice spacing between calcium ions on the templated crystal plane. Previous studies on the growth of calcium carbonate on two-dimensional chitosan biopolymer films have observed such a lattice matching effect that led to templated crystal growth.12,13 Lattice-matching of calcium carbonate single crystals to underlying self-assembled monolayer templates has also been previously reported,7,8,10 although some workers have shown for vaterite, at least,39 that direct epitaxy of calcium carbonate does not necessarily occur at such organic interfaces. The formation of calcium carbonate shells causes the observed large decrease in viscosity of the particle suspensions after the pH was raised to 10.5, since hard crystalline aggregates interact less than the soft biopolymer particles that were initially formed. It is worth noting that in some studies of calcium carbonate shell formation a mechanism occurred that involved the initial crystallization of a vaterite precursor that subsequently dissolved and recrystallized as a calcite shell on the vaterite surface.41-45 If this mechanism had occurred in the current case, the significant lowering of the viscosity would have only been measured if the entire biopolymer bead had become filled with hard, inorganic, material. No evidence was found for this having occurred, and samples observed at all stages after the increase in turbidity that heralded calcium carbonate crystallization appeared as calcite shells. The viscosification of the sample when reacidifed to dissolve the calcium carbonate confirmed that the interaction of the soft gel particles was mainly responsible for the measured viscosity. Importantly, for crystallization to occur on the particle surface, bulk crystallization must be suppressed. Studies on twodimensional calcium carbonate crytallisation on polymer films have shown that the presence of a calcium-sequestering polyelectrolyte in bulk solution is necessary for this effect to occur.12,13,40 In the present study, the CLSM images showed that pectin existed in bulk solution as well as in the particle gel cores. The pectin in bulk solution therefore sequestrated calcium from the bulk solution and yielded favorable supersaturation conditions at the gel particle surfaces for crystallization to occur,



Figure 14. Electron micrograph of calcium carbonate crystals grown in the presence of 0.4% (w/v) LM pectin and L-tryptophan (left column) and L-tryptophan only (right column) at L-tryptophan concentrations of (a) 1%, (b) 0.1%, (c) 0.01% (w/v).

when it did. Sequestration of calcium provides a simple explanation for the difference in titration results between systems with and without biopolymer and the resulting state diagram. The more biopolymer that is present in the system, the more free calcium ions are required to achieve favorable local supersaturations for crystallization at the gel particle surface. At low calcium ion concentrations (for a given biopolymer concentration), there is an insufficient supersaturation of calcium to result in crystallization, as shown by the turbidity results. Finally, it should be noted that ample evidence was gathered for the operation of a third kinetic effect, the equilibration of the carbonate ion concentration in solution. In most of the present study this was discounted by performing experiments immediately after mixing the biopolymer and sodium bicarbonate. However, it was shown that the complex set of dissociation reactions between hydrogen carbonate, carbonate, and carbon dioxide present in the atmosphere above the reaction vessel had

an important influence on the solution pH which, in turn, influenced the supersaturation conditions governing the crystallization conditions for calcium carbonate. Likewise, the concentration of reactants and templating pectin biopolymer had a large effect on the quality of the hollow rosette-like aggregates. Hollow particles only formed under conditions of lower supersaturation and low pectin concentrations where more individual polymer molecules would have been in solution and able to suppress bulk crystallization, making surface-induced crystallization more likely. At the higher reactant concentrations the rates of bulk crystallization will be higher, and at higher pectin concentrations more of the pectin is likely to be gelled, meaning that there is less free polyelectrolyte in solution to suppress bulk crystallization. Although not investigated in detail in the present study, it is sufficient to say that a complete understanding of the process conditions necessary for producing calcium carbonate shells by the current proposed mechanism


Butler et al.

Figure 15. Proposed templating mechanism for the growth of hollow shells of calcium carbonate in the presence of LM pectin.

will be necessary if these systems are to be used for any applications. The results of this study, however, indicate that there will be large challenges for commercial scale-up since the best quality shells were formed at low reactant concentrations. Nevertheless, the ability to form hollow shells of calcium carbonate leads to the possibility that this system could be used for the encapsulation of functional molecules, particularly those that are hydrophilic. The effective encapsulation of such readily diffusing species as ions led to the conclusion that complete shells of calcium carbonate provided highly effective barriers against diffusion, even in an aqueous environment. Furthermore, the ability to readily form hollow shells, with overall diameters in the range 1-10 µm, via a simple mixing process, leads to the possibility that they may be manufactured at large enough scales to be used as a practical encapsulation system in a wide range of applications where encapsulation is necessary. Examples of these applications involve the pharmaceutical and foods industries, where many examples exist of materials that must be stored or used in an aqueous environment but require encapsulation owing to problems with bitterness or stability. The results for encapsulation of ascorbic acid and L-tryptophan illustrate some potential pitfalls in the use of templated inorganic materials for encapsulation, however. The results obtained for calcium carbonate crystallization in the presence of both L-tryptophan and ascorbic acid with pectin or alginate demonstrated the possibility of a competition between the interactions of the biopolymer, added molecules and calcium carbonate preventing hollow shells from forming. For ascorbic acid, the acidic groups in ascorbic acid and the biopolymers will both compete for interaction with calcium carbonate. For L-tryptophan, the basic amine groups in L-tryptophan interact with the acid groups in pectin or alginate and prevent a gel template from forming in the first place, so that hollow calcium carbonate capsules cannot be grown unless the tryptophan is present at very low (and impractical) concentrations. Of course, the list of molecules studied was by no means exhaustive or representative of all the species that could potentially require encapsulation. The study does show, however, that the formation of hollow shells in the absence of any additive does not necessarily mean that a viable encapsulation system has been found and, by extension, that investigation of the encapsulation ability of

these types of inorganic hollow sphere-forming systems must be studied on a case-by-case basis. Conclusions We have shown that it is possible to use the principles of directed crystallization of inorganic materials on acidic biopolymers, which are employed by nature to form such elaborate structures as coccoliths, diatoms, and sea-shells, to successfully form complete, hollow, shells of radially arranged calcite crystals. The calcite crystals form an arranged assembly on a gelled pectin biopolymer substrate that directs the crystallization process via interaction of acidic functional groups with calcium ions. Furthermore, direct evidence was provided that these shells can successfully encapsulate rapidly diffusing species (using the ionic species sodium chloride as an example) providing the possibility for the effective encapsulation of a range of functional hydrophilic molecules of interest to the pharmaceutical and foods industries. There is a caVeat, however, since it was shown that it was not possible to form the hollow inorganic shell in the presence of molecules that either disrupted the formation of the gel template or interacted preferentially with the calcium ions and prevented them from crystallizing into a shell. The formation of hollow shells in the absence of any additive does not, therefore, necessarily mean that a viable encapsulation system has been found. Acknowledgment. The authors thank Unilever for permission to publish this paper. Karolina Barck is acknowledged for assistance in preparing samples for electron microscopy analysis.

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CG8008333

Hollow Calcium Carbonate Microsphere Formation in the Presence of

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