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J. Phys. Chem. C 2009, 113, 3666–3673
Interactions of Organic Molecules with Calcite and Magnesite Surfaces Colin L. Freeman,*,† Iosif Asteriadis,† Mingjun Yang,‡ and John H. Harding† Department of Engineering Materials, UniVersity of Sheffield, Sheffield, U.K., and Nanoscience Centre, UniVersity of Copenhagen, Copenhagen, Denmark ReceiVed: August 7, 2008; ReVised Manuscript ReceiVed: December 17, 2008
This paper presents molecular dynamics simulations of methanoic acid and methylamine molecules interacting with various calcite and magnesite surfaces in aqueous conditions. A new set of potentials for the interaction of nitrogen atoms with mineral surfaces is fitted and tested using ab initio methods. The results show that methanoic acid adsorption is stronger than methylamine adsorption onto calcite and magnesite surfaces. The competition between the surface water structure and the adsorbing molecule is shown to be a major component of the adsorption energy and explains why adsorption is stronger at particular surface indices and onto calcite surfaces rather than magnesite surfaces. 1. Introduction The dissolution and growth of crystals are complex processes. Without external interference or kinetic control, a crystal will produce a morphology with the lowest surface energy. Controlling the growth surface to produce crystals with complex shapes and surfaces with particular properties is of theoretical and technological interest. However, the control achieved in the laboratory is often poor when compared to that of biominerals in living organisms. Biominerals are frequently used by organisms for shells, skeletons, and teeth, where their structure must be carefully controlled to ensure correct material properties. Crystal growth in biominerals appears to be controlled by several different means, for example, the use of vesicles to mold and restrict concentration levels,1 the presence of foreign ions at the surface to inhibit or encourage growth,2-4 or the use of organic molecules such as polysaccharides5 or proteins6 to stabilize growing surfaces. Obviously, these methods need not be mutually exclusive. Foreign ions could be used to influence concentration levels of the solvent ions, and a large macromolecule could act as a mold or restrict the arrival of solvent ions.7 Calcite, a calcium carbonate polymorph, is the building block for many biominerals such as the shells of aquatic algae and crustaceans, which makes it an ideal system to study. When grown without external influences, calcite will produce a rhombohedral crystal with a (10.4) surface. Theoretical studies8 have demonstrated that this surface is at least 0.3 J m-2 lower in energy than any of the other surfaces. Therefore, if other surfaces are to be expressed, they must be substantially stabilized and prevented from growing out. Interactions of surfaces with different ions have been explored in several theoretical studies. Foreign ions (e.g., Mg or Sr) will frequently absorb at calcite surfaces in preference to Ca,9 which through the increased lattice mismatch (because of the change in ion size) can lead to the termination of surface growth or dislocation. These studies also demonstrate the importance of investigating the solvated surface. Water adopts a detailed layered structure at the calcite surface,10,11 and how the ions * To whom correspondence should be addressed. E-mail: c.l.freeman@ sheffield.ac.uk. Telephone: +44 (0)114 222 6021. Fax: +44 (0)114 222 5943. † University of Sheffield. ‡ University of Copenhagen.
interact with the solvent is important for understanding their surface interactions. For example, the smaller Mg ion causes more disruption to the water layers and so is more mobile and slower to reach the surface.11 The water network at the surface varies with the calcite surface; stabilizing each surface to a different degree, but crucially the (10.4) surface remains the most stable.8 Magnesite (10.4) surfaces are also more stabilized by the presence of water than their calcite counterparts because of the larger network of H-bonds that are built up by the absorbed water layer and stronger Mg-water interactions.12 Proteins and polysaccharides have been demonstrated to control calcite crystal growth,5,13,14 but understanding of their functionality is still limited. Monolayer preparation of calcite samples generally requires acidic groups16,17 in order to stimulate crystallization and direct the growth. Volkmer et al.18 used peptides partially constructed from phenylalanine (Phe) and aspartic acid (Asp) units to produce both (11.0) and (01.2) surfaces. No calcite modification was observed with only the hydrophobic Asp present. Amino acids such as glycine, lysine, and alanine can stimulate the growth of vaterite,14 as opposed to the more stable polymorph calcite, but the three different molecules caused no further modification to the morphology. The underlying mechanism of how these molecules operate is still an open-ended question. Assuming that binding is occurring through the acidic groups, the question is raised of their arrangement within the molecule. Estroff et al.19 used a tricarboxylic acid and observed that the grown surface matched the molecular structure closely, suggesting that the position of the functional groups is important. A similar result could be seen in the study by Freeman et al.20 on the effect of positional variation of carboxylic groups attached to benzene rings on the precipitation rate of calcite. Similarly, as the number of acid groups on the molecules was increased, the inhibition of calcite precipitation was also increased, although none of these scenarios produced any clear effect on the actual growth morphologies. Interestingly, they found that isophthalic acid acted as a far stronger inhibitor than the number of acid groups would suggest. This was attributed to the close matching between group separations and carbonate ions of the calcite surface. Inhibition leads to growth via particle aggregation rather than classical nucleation. Aggregation has also been seen in the
10.1021/jp807051u CCC: $40.75 2009 American Chemical Society Published on Web 02/11/2009
Interactions of Molecules with Ca and MgCO3 presence of biopolymers with carboxylic acid or sulfate groups,21 which slow crystallization. These larger multifunctional molecules produced varied morphologies of stacklike crystals. It is difficult to break down and identify the mechanisms that these large molecules use because of their multiple configurations and functional groups. Simulations can provide the option to analyze the behavior of these molecules in more detail. Yang et al.22 analyzed trisaccharide adsorption at various calcite surfaces. These saccharides possess no carboxylic groups but still demonstrate large differences in adsorption energy at calcite surfaces because of hydrogen bonding through alcohol groups. Their results for binding trends agree well with experimental analysis, particularly the differing behavior between the obtuse and acute steps at the surface. Stronger binding at the acute versus the obtuse step has also been observed with polyaspartate chains in experiments and simulations.6 Docking simulations suggest that the greater displacement of water at the obtuse step compared to that at the acute step may make binding less favorable. Alginic acid was simulated at the (10.4) calcite surface by Perry et al.,23 who noted the disruption and also the positive influence the molecule can have on the water structure. They identified ether oxygen as a key component of interfacial binding. The role of organic molecules in influencing calcium carbonate growth is still unclear. However, the possible importance of the surface water structure has emerged and, therefore, needs careful attention. In addition, understanding how particular functional groups bind to the mineral surface is vital in enabling us to break down the interactions of these large molecules into more basic units. This paper reports on a series of simulations for methanoic acid and methylamine at a range of calcite and magnesite surfaces in fully solvated conditions. 2. Methods For all of our atomistic simulations, we used the Pavese force field24,25 to model the calcite and magnesite minerals. Water is treated with the TIP3P model.26 The potentials for the organic molecules are taken from the AMBER force field set.27 Interactions among the solvent, organic molecules, and the mineral either were taken from ref 28 or, in the case of methylamine-calcite and magnesite, were fitted using the methods described in refs 28 and 34 (see section 2.1for further details). Slabs of calcite and magnesite were constructed from preoptimized surface structures using the METADISE (Minimum Energy Techniques Applied to Dislocation Interface and Surface Energy)29 code. The surfaces are minimized using the two-region approach where the atoms are allowed to relax in the region close to the surface, while atoms remain fixed in the region further from the surface to simulate the crystal bulk. The (10.4), (11.0), (10.0), (00.1)-Ca, and (00.1)-CO3 surfaces were relaxed. The last two surfaces are terminated with either Ca-Mg or carbonate ions, making them Tasker type three31 surfaces and therefore polar. Such surfaces are energetically unstable and will undergo reorganization to remove the polarity. METADISE removes the dipole by removing half of the surface ions and placing them on the bottom surface of the outer region. Solution chemistry makes an important contribution to the behavior of the molecules at the interface (see ref 30 for a detailed discussion). The pH of biological systems is approximately 6-6.5. Given the methanoic acid pKa of 3.75, we would expect the molecules to be fully ionized. However, our simulation cell contains 1200 water molecules, and therefore, the acid has a concentration of 0.05 M. If we calculate the
J. Phys. Chem. C, Vol. 113, No. 9, 2009 3667 TABLE 1: List of Potentials Derived for Simulations Described in This Papera Buckingham potentials [A exp(-F/r) - C/r6] Ca Mg Mg Mg Mg
N N Oacid (OH)b Oacid (CdO)c OW
A (eV)
F (Å)
C (eV Å-6)
6512.4 1555.3 592.2 471.6 852.2
0.253 0.253 0.297 0.297 0.297
0.0 0.0 0.0 0.0 0.0
a Further potentials can be found in ref 28. b (OH) refers to hydroxyl oxygen (i.e., with an O-H bond). c (CdO) refers to carboxyl oxygen (i.e., oxygen double bonded to carbon).
ionization at this concentration, we find that less than 1% of the acid molecules will dissociate. Therefore, given our system size, it is appropriate to leave the molecule associated. The computational expense of modeling a box large enough to give bulk solution properties is too great. An alternative would be to dissociate the acid to form a bicarbonate ion on the surface, for which we have the necessary potential models.28 This method allows us to use a smaller box but would require the exploration of greater configurational space, given the possible positions of the ion at the mineral surface. These factors mean that a thorough consideration of pH effects would be a detailed study in itself, and therefore, we treat our molecules in associated states with neutral surfaces. These data are still invaluable to exploring the methods of binding at the interface. 2.1. Methylamine Potential. Fitting of potentials to model the interactions between the mineral and organic components of the system requires careful attention. The organic-solvent molecules are described with a largely covalent model, while the mineral is described with an ionic model. This means that there is a large disparity between the charges of the atoms in the two systems, so using standard mixing rules will not suffice. This problem can be overcome by using the methods described in refs 28 and 34. The potentials used for these “cross-term” interactions are fitted to known mineral structures. Atoms are assigned charges that maintain the ratio of charges between different atoms but also mimic the electrostatic interactions between organic and mineral components, i.e., a reduced charge model. Note that this reduced charge model is only used for the generation of the potentials and is not used subsequently in the simulations. For the purposes of our systems we fitted Nmethylamine-Cacalcite and Nmethylamine-Mgmagnesite potentials to calcium nitride, Ca3N2,35 and magnesium nitride, Mg3N2, respectively. The potential used for the system was a modified version of Wolf et al.36 for Li3N fitted, using the Ca3N2 structure in ref 37. The final parameters from the fitting are listed in Table 1. These potentials were verified by comparison of the adsorption of methylamine onto a calcite surface calculated with potential-based and ab initio methods. Five-layered slabs of CaCO3 terminated with the (10.4) surface were generated and optimized using the METADISE program. Nine methylamine molecules were placed on both surfaces of a 3 × 3 calcite slab. Molecular dynamics (MD) simulations (at 10 K), using the DL_POLY_2.16 code,32 were performed using the slab model and the same parameters as described above. Several starting configurations were used to ensure a reasonable sample of space. Local minima were identified by running the simulations until the methylamine molecules showed no further migration. The lowest-energy geometry for each surface was then used for the comparison of the adsorption energies. A further set of MD simulations were performed with a 10 ps time length in order
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Figure 1. Comparison of the adsorption energy (Eads) of a methylamine molecule onto a calcite (10.4) surface as calculated using QM methods for our new potential.
to produce a configurational energy for the local minima. All the atoms, excluding organic carbon and nitrogen and those in the middle layer of the calcite slab, were free to move. The separation between the molecules and the calcite surface was systematically varied, while the orientation of the molecule with respect to the calcite surface was kept constant. The energy was averaged over 10 ps and used to calculate adsorption energies, Eads, as shown in eq 1
Eads )
(Esystem - Ecalcite - Emolecule) Nmolecule
(1)
where Esystem is the total energy of the molecules and calcite, Ecalcite is the separate energy of the calcite slab, Emolecule is the energy of the organic molecules, and Nmolecule is the number of organic molecules present within the cell. The ab initio calculations were performed using periodic plane wave density functional theory (DFT) as implemented in the CASTEP3 code.38 The generalized gradient approximation (GGA) and Perdew-Wang (PW91) exchange-correlation functional39 were used. Our simulations used ultrasoft Vanderbilt pseudopotentials40 with a cutoff at 560.0 eV. The reciprocal space integration utilized the Monkhorst-Pack sampling scheme,41 and a k-point spacing of 0.05 was applied. Simulation cells from DL_POLY were reduced to a 1 × 1 cell with a single methylamine molecule above each surface to reduce computational expense. Single-point energy calculations were performed on all final configurations produced by the 10 ps MD simulations, and adsorption energies were calculated as described in eq 1. The vacuum gap was reduced from MD simulations to approximately 15 Å to reduce computational expense. This gap was tested for energy convergence. Results of the potential tests are seen in Figure 1. The classical results agree well with the ab initio results and demonstrate the reliability of the methods used to generate the potentials. The classical results show a small shift closer to the surface but only by about 0.1 Å. 2.2. Slab Calculations. A series of simulations of hydrated calcite and magnesite slabs in contact with methanoic acid (HCOOH) and methylamine (CH3NH2) were performed using the DL_POLY 2.16 code.32 The slab model was used for the simulation cell, whereby the system was contained within a three-dimensional (3D) periodic box but separated from the other substrate layers by a substantial vacuum gap (∼40 Å). A random configuration of 600 water molecules was added to both surfaces of the slab using the Packmol program33 (see Figure 2 for an example of the simulation system). A single molecule of methanoic acid or methylamine was added to the simulation cell in close proximity to the surface. A Nose´-Hoover NVT thermostat with a 0.1 ps relaxation time was used to maintain
a constant temperature. Coulombic forces were calculated with a periodic Ewald summation, and the short-range potentials had a cutoff at 10.1 Å. Simulations were performed with 1 fs time steps. The system was warmed from 10 to 300 K over three simulations (at 10, 100, and 200 K) and then equibrilated for 100 ps. The final simulations were run for 2.0 ns, and energies were collected from the final 1 ns of these simulations. The adsorption energy of the molecule (Emolecule-mineral) was calculated at each surface with respect to the water-only solvated surface (Ewater-mineral) and the solvated molecule (Ewater-molecule). This energy represents the energy of the molecule leaving the water bulk and adsorbing onto the surface displacing some of the surface water
Emolecule-mineral ) Ewater-mineral-molecule - Ewater-mineral Ewater-molecule (2) where Ewater-mineral-molecule is the configurational energy from the simulation of the whole system. In order to calculate this energy, further simulations of the solvated molecule, bulk water, and water at the mineral surface were performed using the same methods as described above. 3. Results and Discussion 3.1. Molecules at the Aqueous Interface. The binding of the methanoic acid molecule (Figure 3a) to the surface takes place through a hydrogen bond between the carbonate oxygen and the alcohol hydrogen (H1) of the acid and the interaction between the calcium and the carbonyl oxygen (O2) of the acid group. The hydrogen bond remains at a fairly constant length of ∼1.4 Å for each surface (magnesite and calcite), as would be expected because the atoms involved are not changing. The O2 cation separation does not vary significantly between the different surfaces of the two minerals, producing values of ∼2.2 and ∼2.35 Å for magnesite and calcite, respectively. The methylamine molecule (Figure 3b) interacts with the surface through hydrogen bonds between the amine hydrogen and carbonate oxygen along with a nitrogen-cation bond. The hydrogen bonds are significantly longer (∼0.55 Å) than those of the acid molecule and show a greater variation in length during the simulation. This would be expected because of the weaker nature of the amine hydrogen bonds. The N cation separation does not vary significantly between the different surfaces of the two minerals, producing values of ∼2.55 and ∼1.95 Å for calcite and magnesite, respectively. Adsorption energies are listed in Table 2. Several interesting features immediately emerge. First, the methylamine molecule adsorption energies are all less binding than those of methanoic acid for the same surfaces, with the exception of the (10.0) magnesite surface. This may be expected because of the smaller charges of the methyl group versus those of the acidic oxygens, and therefore, the total interaction between a methylamine molecule and a surface compared to that of a methanoic acid molecule and a surface is weaker. Second, the pattern of the adsorption energies for the five surfaces is generally the same for magnesite and calcite with the two different molecules. For example, at the magnesite surface, the adsorption energy has a significantly stronger binding for the (10.4) and (00.1)-Mg surfaces compared to that of the (11.0) and (00.1)-CO3 surfaces with both methanoic acid and methylamine. While at the calcite surface, the (10.4), (10.0), and (00.1)-CO3 surfaces produce stronger binding adsorption energies with both methanoic acid and methylamine compared to those of the (11.0) and (00.1)-Ca surfaces. This suggests that
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Figure 2. Example system for the simulations showing the magnesite mineral slab and water. In this and subsequent figures (unless otherwise noted), the Mg ions are depicted in green, carbons in gray, oxygens in red, and hydrogens in light blue.
Figure 3. Example positions of (a) methanoic acid and (b) methylamine at the (a) magnesite (10.0) and (b) calcite (11.0) surfaces. Cations (Mg and Ca) are depicted in green, and N atoms are dark blue. Water molecules are faded to aid visualization of the adsorbate molecule.
TABLE 2: Adsorption Energy of Methanoic Acid and Methylamine onto Calcite and Magnesite Surfaces with Respect to Watera adsorption energy (kJ mol-1) methylamine surface index 10.4 11.0 10.0 001-Ca-Mg 001-CO3 mean
calcite
magnesite
-39.3 ( 13.7 9.2 ( 20.4 13.7 ( 10.4 63.8 ( 15.6 -24.6 ( 14.4 12.5 ( 18.2 n/ab 33.6 ( 15.0 -0.2 ( 13.6 76.0 ( 10.0 -10.1 ( 13.1 34.02 ( 15.9
methanoic acid calcite
magnesite
-64.1 ( 10.0 -8.6 ( 19.8 4.9 ( 4.0 39.9 ( 14.0 -57.5 ( 13.4 48.0 ( 19.5 64.1 ( 15.6 4.5 ( 16.7 -44.5 ( 17.1 37.2 ( 12.4 -19.4 ( 10.4 24.2 ( 16.5
a Values are listed with error bars for maximum and minimum values based on fluctuations during simulations. Note the more negative the value, the stronger the binding of the molecule. b The methylamine molecule failed to remain in contact with the surface during simulations, so no adsorption energy could be calculated.
the structure of the surface water-molecule interface is similar for each particular molecule at the calcite or magnesite surface. Third, adsorption of both molecules onto the magnesite surface is weaker than onto the calcite surface for all surfaces. This is surprising, given that the smaller size of the Mg ion
compared to that of the Ca ion means the magnesite-molecule interactions are more attractive than the comparable calcite interactions. However, we must remember that our energy is calculated with respect to the fully solvated mineral surface. Therefore, the presence of methanoic acid at the surface means some water has been displaced. It has previously been noted for other molecular adsorptions that displacement of water may be a factor in determining the adsorption energy.6 Adsorption of water monolayers onto calcite and magnesite surfaces has shown that water adsorption is stronger at the magnesite surface than at calcite surfaces because of the strength of the Mg-Owater interaction and the additional intermolecular interactions that form between the water molecules.12 Therefore, the adsorption energies reported are a feature of the loss of water from the surface, rather than only the strength of the molecule-surface interaction. 3.1.1. Magnesite Surfaces. For magnesite surfaces, methylamine adsorption is less favorable by at least 30 kJ mol-1 for the (11.0) and (00.1)-CO3 surfaces compared to those of the (10.4), (10.0), and (00.1)-Mg surfaces. Examination of the z-density of the water and methylamine molecules at the surface shows different behaviors (Figure 4). For the (10.4) and (00.1)-
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Figure 4. z-Density of water oxygen and methylamine nitrogen at the (a) (00.1)-Mg, (b) (00.1)-CO3, (c) (10.0), (d) (11.0), and (e) (10.4) magnesite surfaces.
Figure 5. Time-averaged image of methylamine at the (a) (00.1)-Mg and (b) (00.1)-CO3 magnesite surfaces over 1 ns. The positions of the oxygen water (red) and methylamine nitrogen (blue) are plotted at regular intervals to demonstrate their average positions. The magnesite surface is also shown at a single time step for reference. The nitrogen atom is seen to reside mostly in the first water layer at the (00.1)-CO3 surface compared to that of the (00.1)-Mg surface, where it is between the first and second layers.
Mg surfaces, the nitrogen of the methylamine resides in the second water layer above the surface (note the first water layer on the (10.4) surface has a significantly smaller density than the second). However, for the (11.0) and (00.1)-CO3 surfaces, the nitrogen is within the first water layer. At the (10.0) surface, the nitrogen is generally closer to the surface than in any of the water structures and is therefore only partially within the water layer. This suggests that in the (10.4) and (00.1)-CO3 cases the presence of the methylamine molecule is causing a greater disruption to the water structure at the surface, and therefore, a weaker binding adsorption energy is recorded. Viewing two images of the system in Figure 5, one can see this occurring. The nitrogen atom at the (00.1)-Mg surface is sitting between the water layers, while the acid molecule at the (00.1)-CO3 surface is completely within the water layers. Methanoic acid adsorption is energetically less favorable by at least 30 kJ mol-1 for the (11.0), (10.0), and (00.1)-CO3 surfaces compared to that of the (10.4) and (00.1)-Mg surfaces. Examination of the z-density of the water and methanoic acid molecules suggests a pattern similar to that seen for methylamine (Figure 6). At the (00.1)-Mg surface, the O2 of the acid resides in the second water layer on the surface, while at the (10.4) surface, the O2 resides only partially within the first water layer
(and on the bulk water side of this layer). This compares with the (11.0), (10.0), and (00.1)-CO3 surfaces, where O2 is fully within the first water layer. This suggests that in the three latter cases the presence of the acid molecule is causing a greater disruption to the water structure at the surface, and therefore, a weaker binding adsorption energy is recorded. Interactions at the (11.0) surface behave differently between the two molecules. In the case of methylamine, the molecule was able to reside closer to the surface than much of the water structure, while with the acid, the molecule is completely encompassed within the water layer, and the relative adsorption energy is less binding. It is interesting to note the differences in the water structure at the different surfaces. For example, the water orders differently at the two polar surfaces. At the magnesiumterminated surface (Figure 5), the layers are relatively merged together with less distinction between them when compared to that of the carbonate-terminated surface. These differing water structures are obviously a result of the mineral surface structure. At the (00.1)-CO3 surface, the carbonate ions lie flat above the magnesium ions, the water molecules are very close, and the oxygen atoms are attempting to close the gap with the cations. In contrast, at the (00.1)-Mg surface, the
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Figure 6. z-Density of water oxygen and acid oxygen (O2) at the (a) (00.1)-Mg, (b) (00.1)-CO3, (c) (10.0), (d) (11.0), and (e) (10.4) magnesite surfaces.
Figure 7. z-Density of water oxygen and acid oxygen (O2) at the (a) (00.1)-Ca, (b) (00.1)-CO3, (c) (10.0), (d) (11.0), and (e) (10.4) calcite surfaces.
magnesium ions are now in the outermost position, and the carbonate ions in the layer below rotate to project one of their oxygen atoms into the magnesium layer. Therefore, the water is now close to both the carbonate and magnesium ions. Clearly, adsorption at different surfaces varies because of the surface structure and its effect on the fluid structure above. The z-density plots suggest that the more disorganized water structures, where the first layer is less clearly defined, e.g., (00.1)-Mg and (10.4), have the lowest adsorption energies. If the water structure is less clearly defined, it implies that disruption of this water may be energetically easier. 3.1.2. Calcite Surfaces. Calcite surfaces also demonstrate a trend similar to those of the magnesite surfaces. The adsorption of methanoic acid is significantly stronger (greater than 50 kJ mol- 1) at the (10.4), (10.0), and (00.1)-CO3 surfaces than at the (11.0) and (00.1)-Ca surfaces. The oxygen atom generally resides in a similar position relative to the water structure at
the calcite surface, as it did at the magnesite surface. Differences in adsorption between the calcite and magnesite surfaces are therefore hard to identify. The water structure at the (00.1)-Ca surface still shows two peaks close together, and the acid oxygen sits in the second peak. However, the first water peak is far smaller than the second at the calcite interface compared to that of magnesite (where they are approximately equal). This suggests that the second water layer at the calcite surface is more significant, and the disruption caused by the presence of the acid molecule makes the adsorption less favorable than at the magnesite surface. In contrast, the acid molecule at the (00.1)-CO3 calcite surface now resides within the first peak and the gap between the layers. Therefore, the molecule spends approximately one-half of its time between the layers causing less disruption to the water structure, leading to a stronger binding than at the magnesite surface. The water structure at the (10.0) surface is more sharply defined at the calcite surface
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Freeman et al. The adsorption energies we have calculated consider the adsorption of a single molecule onto the surface. One molecule adsorbing within a layer of water will cause a large disruption to the water structure, but if multiple molecules were adsorbed, then these molecules would be able to stabilize each other in much the same way the water molecules do, setting up potential domains of molecules. Eventually, we can imagine the adsorption of a monolayer of the molecules. As the adsorption energies per molecule and the molecular density at the surface increase, adsorption is likely to change significantly. Molecular monolayers have been considered by Cooper and de Leeuw,42 who performed energy minimizations of methanoic acid monolayers onto calcite surfaces. They report organized monolayers of the acid molecules, which imply intermolecular interactions that help stabilize the adsorption. Because their simulations are carried out with respect to a vacuum, direct comparison of the energies is not possible, but it is interesting to note the small differences observed between their two polar surfaces. Simulations performed in relation to water are needed to elucidate this further.
Figure 8. z-Density of water oxygen and methylamine nitrogen at the (a) (00.1)-CO3, (b) (10.0), (c) (11.0), and (d) (10.4) calcite surfaces. Note that no z-density plot is shown for the (00.1)-Ca surface, as the molecule always entered the bulk in the simulation.
than at the magnesite surface, and now the acid oxygen is not fully encompassed within the first water layer at the calcite surface. This may explain the stronger binding at the calcite surface. The z-density profiles for the (11.0) and (10.4) surfaces are very similar, and it is surprising to see adsorption at the (10.4) surface significantly stronger than at the (11.0) surface. Adsorption of methylamine also produces the weakest binding for the (11.0) and (00.1)-CO3 surfaces compared to that of the (10.4) and (10.0) surfaces. Adsorption at the (00.1)-Ca surface was not possible as the molecule moved into the bulk in all simulations. This implies that adsorption is unfavorable at the (00.1)-Ca surface, as was seen for methanoic acid. In general, the energy differences for the methylamine molecule adsorption are smaller than those for the methanoic acid molecule. Comparing the z-density plots for the calcite and magnesite surfaces again shows some differences. Nitrogen at the carbonate-terminated surface is mostly between the two water layers leading to a lower adsorption energy than at the magnesite surface. The (10.0), (10.4), and (11.0) z-density plots are similar for both interfaces. They provide no immediate clues as to why the adsorption of methylamine at the (10.0) surface is much more favorable at the calcite surface than at the magnesite surface. It is somewhat surprising that the two minerals produce different trends in adsorption energy when adsorbing the same molecule. However, the water structure does vary significantly between the minerals with the same surface as seen in ref 12. For example, at the (10.4) magnesite surface, the water molecules are close enough together to develop an intermolecular hydrogen bond arrangement that is not seen at the (10.4) calcite surface. Because the water structure can vary, the adsorption energies also vary. Adsorption at the (00.1) surface is particularly interesting. This is a polar surface that can be terminated with either the cation or the carbonate ion. For magnesite, the adsorption is stronger at the cation (Mg)terminated surface than at the carbonate-terminated surface for both molecules. The opposite result is true for calcite, where adsorption is stronger at the carbonate-terminated surface. Calculations of polysaccharide adsorption onto calcite surfaces22 observed the same preference as our simulations. This demonstrates a significant difference between the behavior of molecular adsorption for the two minerals.
4. Conclusions In this paper, we have developed and tested new potentials for the interactions of calcite and magnesite with nitrogen-based functional groups. The addition of these potentials to the library we have previously developed28 allows us to model the interactions of peptides and proteins with mineral surfaces. These molecular types are often involved in biomineralization processes, so a reliable potential is a necessity. Our simulations of small molecules at the magnesite and calcite surfaces demonstrate several interesting features about their adsorption. First, the behavior of the water layer may be an important component of adsorption at all surfaces in aqueous conditions. Less disruption to the surface water structure during adsorption generally equates with lower adsorption energy. Second, adsorption at magnesite surfaces is generally weaker than at calcite surfaces because of the stronger binding of the water to the magnesite surface. It would be interesting to explore this further with molecular monolayers and see if this trend continues at all surface concentrations. Third, the methylamine molecule has a significantly weaker bond to the mineral surface than that of the methanoic acid molecule. It would be expected that this trend for functional groups would scale to larger molecules and may therefore be very important for the binding of peptides and proteins, as their backbone and side chains frequently exhibit acid- or nitrogen-based groups. References and Notes (1) Hetherington, N. B. J.; Kulak, A. N.; Sheard, K.; Meldrum, F. C. Langmuir 2006, 22, 1955–1958. (2) Paquette, J.; Reeder, R. J. Geochim. Cosmochim. Acta 1995, 59, 735–749. (3) Reeder, R. J. Geochim. Cosmochim. Acta 1996, 60, 1543–1552. (4) Lam, R. S. K.; Charnock, J. M.; Lennie, A.; Meldrum, F. C. CrystEngComm 2007, 9, 1226–1236. (5) Henriksen, K.; Stipp, S. L. S.; Young, J. R.; Marsh, M. E. Am. Mineral. 2004, 89, 1709–1716. (6) Elhadj, S.; Salter, E. A.; Wierzbicki, A.; De Yoreo, J. J.; Han, N.; Dove, P. M. Cryst. Growth Des. 2006, 6, 6197–201. (7) DiMasi, E.; Kwak, S.-Y.; Amos, F. F.; Olszta, M. J.; Lush, D.; Gower, L. B. Phys. ReV. Lett. 2006, 97, 045503. (8) de Leeuw, N. H.; Parker, S. C.; Rao, K. H. Langmuir 1998, 14, 5900–5906. (9) de Leeuw, N. H. J. Phys. Chem. B 2002, 106, 5241–5249. (10) Fenter, P.; Geissbu¨hler, P.; DiMasi, E.; Srajer, G.; Sorensen, L. B.; Sturchio, N. C. Geochim. Cosmochim. Acta 2000, 64, 1221–1228. (11) Kerisit, S.; Parker, S. C. J. Am. Chem. Soc. 2004, 126, 10152– 10161. (12) de Leeuw, N. H. Am. Mineral. 2002, 87, 679–689.
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