Barium Sulfate Crystallization from Synthetic SeawaterMatthew Boon and Franca Jones*

Department of Chemistry, Nanochemistry Research Insitute, Curtin University, GPO Box U1987, Bentley, Western Australia 6845,Australia

*S Supporting Information

ABSTRACT: Barium sulfate was crystallized in a synthetic seawater mixture thatwas chosen to better reflect ocean conditions. The synthetic seawater containedmonovalent ions, magnesium, strontium, and calcium as well as bicarbonate andboric acid. The natural pH of the synthetic seawater is 8.1, and this seawater wasused to determine the impact on morphology, nucleation rates, and incorporation offoreign ions. It was found that dendritic and diamond-shaped particles are bothformed. The main parameters influencing the formation of dendritic particles werethe saturation index and the ion ratio, but there was also a significant synergisticeffect with the other ions present. The diamond-shaped particles formed later at alower saturation index. The nucleation rate in synthetic seawater was found to behigher than expected based on an ion ratio basis. This is most probably because the divalent ions induce a higher nucleation rateby lowering the surface free energy. Strontium was found to be the dominant ion substituting for barium with some calciumsubstitution also occurring. Finally, the presence of silicate appeared to form dendritic particles with a larger aspect ratio andmore impurities being present but little other impacts. The lack of change in the homogeneous nucleation rate supports aheterogeneous nucleation hypothesis.

■ INTRODUCTION

Barium sulfate, an inorganic, sparingly soluble sulfate salt is nota common solid to be found in seawater, but it is found as abiomineral in some organisms.1,2 For example, in green alga3 ithas been found that a sulfate “trap” leads to the formation ofthe mixed Ba/Sr sulfate. There is also interest in Barite (themineral name of barium sulfate) crystallization in seawaterbecause of the “Barite paradox” whereby it appears to form inundersaturated seawater.4,5 This paradox has been linked toboth the occurrence of strontium and silicate in seawater.5−7

The presence of Barite in sediments is often taken as anindicator of significant biological productivity (i.e., as an oceanproductivity proxy) as is strontium sulfate.8,9 Thus, it isimportant that the chemistry of these proxies is is wellunderstood. Calcium carbonate and barium sulfate are alsoboth minerals of interest from a scaling perspective.10−13

Barium sulfate is most important to off-shore oil production,where its lower solubility means it is the most likely tocrystallize when aquifer water mixes with seawater.14,15 Thus,understanding the crystallization process of barium sulfate inseawater is of benefit both from a fundamental and practicalaspect.While silica and silicates are one of the most common

minerals found on Earth, making up >90% of the Earth’s crust,they can be difficult to investigate since the chemistry ofsilicates can involve inorganic polymerization of the silicatespecies, depending on the pH and composition of thesolution.16 There are, however, a few studies available oncrystallization in the presence of silicate. The Pina groupinvestigated calcium carbonate17 and strontium sulfate18,19

crystallization where they found that polysilicic acids have a

different effect to the monosilicic acids. Another study is that ofLakshminarayanan, on calcite, which showed that the silicateion was incorporated into the structure.20 Thus, the impact ofsilicate is very sensitive to pH and the mono-/polymeric form.The study involving strontium sulfate showed increasing silicateconcentration led to the formation of the hemihydrate andfinally an amorphous solid.18,19 Work by the group of Garcia-Ruiz21 has shown that barium carbonate/silicate structureshaving curved surfaces are formed from purely inorganicsources for which they coined the term “biomorphs”. From thiswork it was shown that silicate at high pH is required to formsuch structures. Barium sulfate has also been investigated in thepresence of silicate.22 Our previous work on barium sulfatecrystallization showed that at pH 10 fibrous structures areformed; however, what is lacking is similar work in seawater.This is because, as previously stated, silica has been implicatedas being an important component in causing the Bariteparadox.5−7

Some limited literature is available on the impact of thedifferent inorganic ions on barium sulfate crystallization.14,23−29

Many of these are from the perspective of off-shore oilproduction where a “seawater” is mixed with an “aquifer” water.These studies are mainly interested in the amount or rate ofscale formation or use limited ions for their “seawater”equivalent;14,24−28 of these, two were found to be mostrelevant.23,29 These results show that the morphology of theformed barium sulfate is significantly altered in the presence of

Received: May 12, 2016Revised: June 20, 2016

Article

pubs.acs.org/crystal

© XXXX American Chemical Society A DOI: 10.1021/acs.cgd.6b00729Cryst. Growth Des. XXXX, XXX, XXX−XXX

seawater. However, even literature on how different inorganicions incorporate or impact on barium sulfate is sparse. Ions thathave been investigated to date are the following:

(a) the alkali cations and halide anions,30−32 which havebeen shown to promote Barite crystallization

(b) the alkali earth metal cations such as calcium andstrontium33,34 and lead,35 which can substitute forbarium ions in the solid

(c) zinc and lanthanum,36 which both inhibit the growth ofBarite, while lanthanum ions can substitute ∼3% into theBarite lattice.

(d) carbonate,37 this study showed that the growth rate isreduced in the presence of carbonate and that the 2Disland shape becomes elliptical. Eventually, witherite(BaCO3) may form.

Unfortunately, all of these were investigated separately and inpure systems. There is little to no information on what thecombination of ions does to the structure of Barite other thanits morphological changes.23,29

Seawater is itself a complex mixture of inorganic and organicspecies that changes depending on the region, depth, and timeof year. In this work, we have used a synthetic seawater recipebased on that by Berges.38 In addition, we have chosen bariumsulfate because solids such as calcium carbonate have issues ofpolymorph formation. Three common crystalline polymorphsare known for calcium carbonate (vaterite, aragonite, calcite) atroom temperature and pressure, while only Barite is known forbarium sulfate. Finally, as already alluded to, barium sulfatecrystallization in seawater is relevant for a variety of reasonsranging from off-shore oil production to ocean productivityproxies. These considerations make barium sulfate crystalliza-tion a suitable candidate as a model system, and we have used itas such to gain some fundamental insights into crystallizationgenerally.39−41

The work discussed herein investigates, as a starting point,only the inorganic ions in seawater. This combination of ions isthe most complex investigated in the literature to date, and it isimportant to determine whether the crystallization behavior ofbarium sulfate is modified from that previously observed. Inaddition, the impact of silicate ions on barium sulfatecrystallization in the synthetic seawater (SSW) will beinvestigated as it has been suggested to play a role in theBarite paradox. The impact on the morphology, nucleation rate,crystallinity, and structure (through XRD) and molecularsimulation data are presented and discussed with respect tocrystallization phenomena and to possible impacts on oceanproductivity proxies.

■ MATERIALS AND METHODSAll materials were analytical grade reagents used as received. Ultrapurewater (resistivity > 18 MΩ cm at 22 °C) was used in the preparationof all solutions.Synthetic Seawater (SSW) Solutions. The SSW solutions were

made up according to Berges38 whereby two salt solutions (saltsolution 1 and 2 listed in Table 1 below) are prepared separately andthen mixed together to form the SSW. This is done to avoidcrystallization of calcium carbonate prior to start of the crystallizationreaction. Table 1 lists the ions in the SSW and their finalconcentrations. The natural pH for the mixed system was found tobe 8.1. Ion groupings were also investigated whereby only themonovalent, divalent ions, or bicarbonate + borate were investigated.When this occurred the concentrations were equivalent to that in theSSW listed in Table 1. In these instances two separate solutions werenot necessary, and barium chloride was added to the prepared solution

containing the desired sulfate concentration to commence crystal-lization.

Once prepared, it is important to know which species are expectedto form from the SSW. The supersaturation was calculated using thePHREEQC program42 (using the phreeqc database). Each of the ionsand their concentrations were inputted and the activities werecalculated. The saturation index (SI) is then derived from thefollowing:

= KSI log[a a a .../ ]i j k sp (1)

where the ai, aj, etc. refer to the activities of the ions (and is called theion activity product), and Ksp refers to the solubility product of thesolid containing species i, j, k, etc.

Table 2 lists the species where the SI calculated by PHREEQC wasgreater than 0. For those species whose SI is <0 this implies

undersaturated conditions. As such, these species would not beexpected to crystallize.

Crystallization Morphology. The morphology of the bariumsulfate formed was determined in static, batch crystallizationexperiments. Cleaned glass vials were filled with salt solution 1 ofSSW or the required solution to the desired volume and water wasadded to achieve the desired concentration. Sodium metasilicatepentahydrate (prepared fresh at the desired pH at a stockconcentration of 1000 ppm) was then added if required to achievethe desired concentration. For the 4.6 mM silicate case a solution atthe required concentration and pH was prepared. A cleaned, round,glass coverslip was placed at the bottom of the vial and allowed toequilibrate for 1/2 h before the addition of the barium chloridecontaining salt solution or barium chloride (varying volume, 0.1 M)stock solution to commence the crystallization reaction. The totalvolume of the crystallization experiment was kept constant at 20.0 mL.After 3 days (or the desired time), the glass coverslips were removed,and the excess solution was soaked up by tissue paper. The coverslipwas then prepared for imaging.

Table 1. Synthetic Seawater Stock Solutions and FinalConcentrationsa

stock solution final concentration

g L−1 mM ppm

Salt solution 1NaCl 21.19 363 21,213KCl 0.599 8.04 599KBr 0.0863 0.725 86.2NaF 0.0028 0.0657 2.76H3BO3 0.0230 0.372 23.2NaHCO3 0.174 2.07 173Na2SO4 3.55 varied (0.25) 0−3551 (35.5)Salt solution 2MgCl2·6H2O 9.592 41.2 8,375CaCl2·2H2O 1.344 9.14 3,307SrCl2·6H2O 0.0218 0.082 79.5Na2SiO3·9H2O 1000 varied 0−813BaCl2 20.8236 varied (0.06) 0−5205 (12.5)

aValues in parentheses are the concentrations used to obtain an SI of2.5.

Table 2. Saturation Index (SI) from PHREEQC forSupersaturated Species and the Database Solubility Product

SI log Ksp

aragonite 0.38 −8.34barite 2.52 −9.97calcite 0.52 −8.48dolomite 1.87 −17.09

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Scanning Electron Microscopy (SEM). The coverslips harvestedfrom the batch experiments were placed on carbon coated SEM stubs,and carbon paint was applied to the circumference to help avoidcharging effects. The stubs were then dried in a desiccator prior tobeing sputter-coated with gold and viewing on an Evo Zeiss SEMinstrument. Energy dispersive X-ray spectroscopy (EDS) analysis ofthe solids was conducted on the gold sputtered samples using anaccelerating voltage of 15 keV.Powder X-ray Diffraction (XRD). In order to obtain sufficient

solids for XRD, the crystallization experiment was scaled up from20 mL to 4 L. The proportions were equivalent to the small-scaleexperiments, conducted at room temperature, ∼22 °C, and the solidswere magnetically stirred for 3 h before being left for 3 days to stand.After this time, the supernatant was decanted, and the solids obtainedby filtration through a 0.2 μm membrane. The solids were washedthree times with ultrapure water. The solids were then dried in a

desiccator, and the XRD patterns were collected on a Bruker D8Advance instrument using Cu Kα radiation. A low background holderwas used and spun at 30 rpm. The 2 theta range was 15−50 deg with astep size of 0.001° and a divergence slit of 0.3°.

Thermal Analysis. Approximately 15 mg of sample was heated in aplatinum pan using a TA Instruments SDT 2960 simultaneous DSC-TGA instrument. The temperature ranged from ambient to 800 °C at5 °C per min in air at a flow rate of 40 mL/min. The temperature ofthe instrument was calibrated against the melting points of indium,zinc, tin, silver, and gold. In addition, the balance was calibrated overthe temperature range with standard alumina weights as provided bythe vendor.

DLS. The nucleation behavior of the system was investigated usinga Malvern Nano ZS instrument to determine the derived counts versustime. The intensity of the scattered light, adjusted for optical filters, isused by the instrument to calculate the derived counts. These counts

Figure 1. SEM images of barium sulfate formed at SI = 2.5, pH 8 (a) in SSW and (b) in pure water, (Ba2+/SO42− 1:1).

Figure 2. SEM image of barium sulfate formed at pH 8 in pure water (Ba2+/SO42− 1:1) and SI = 2.5 in the presence of K+, F−, Na+, Cl−, Br−. Circled

particles show a lengthened c axis.

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are quoted in kilocounts per second (kcps) and should relate to thetotal number of particles within the measurement volume. The particlecounts in this work are <100 kcps when no nucleation has occurredand increase once homogeneous nucleation has occurred; the higherthe derived counts the greater the nucleation rate must be.Molecular Modeling. The molecular modeling of barium sulfate

was performed using static empirical potentials, and the ability forstrontium to substitute was determined from the following reactionscheme:

+ → ++−

+n nBa SO Sr (aq) Ba Sr SO Ba (aq)x x n n42

42

(2)

The methodology is discussed in detail in ref 36. The simulation wasperformed on a 3 × 3 × 2 supercell to make a relatively cubic cell andSr was substituted for Ba. To check for consistency and possiblesources of error, a BaSO4 supercell had Ba

2+ substituted with Sr2+ up tohalf the occupancy and then a SrSO4 supercell was simulated with Sr2+

substituted with Ba2+ up to half the occupancy. The results from thesesimulations were combined to give information on the movements of

the lattice parameters, the cell volume, and the replacement energydefined as

= + − +E E nE E nE( ) ( )repl substituted hydr,Ba initial hydr,Sr (3)

where Erepl is the replacement energy, Esubstituted is the final energy withthe cation substituted, Ehydr,Ba is the hydration energy of the bariumion, Ehydr,Sr is the hydration energy of the strontium ion, and Einitial isthe initial energy of the pure barium sulfate.

■ RESULTS AND DISCUSSION

Barite in SSW. Despite the system being supersaturatedwith respect to four different solids (see Table 2), only Baritewas seen to crystallize. Dolomite, aragonite, or calcite werenever observed. The lack of calcite and aragonite is not thatsurprising given the low SI values. Obviously, for these solidsthe driving force was not even sufficient for heterogeneousnucleation. The lack of dolomite is perhaps surprising given an

Figure 3. SEM image of barium sulfate formed at pH 8 in pure water (Ba2+/SO42− 1:1) and SI = 2.5 in the presence of (a) bicarbonate and borate

ions and (b) the divalent cations.

Figure 4. SEM images of barium sulfate formed at pH 8 in pure water (Ba2+/SO42− 1:1) and SI = (a) 2.0, (b) 2.5, (c) 3.5, (d) 4.0.

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SI of 1.87, but dolomite is known to be difficult to form atambient temperatures.43 As shown in Figure 1a, the

crystallization of barium sulfate in SSW resulted in the presenceof two particle populations. The first population is the large

Figure 5. SEM images of barium sulfate formed at different ion ratios; those on the left have a fixed concentration of sulfate ions and those on theright have a fixed concentration of barium ions. Solids formed at pH = 8, SI = (a, b) 2, (c, d) 2.5, (e, f) 3.0.

Figure 6. SEM images of barium sulfate formed in SSW at pH 8.1, at different times; (a) 15 min, (b) 30 min.

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dendritic particles that have many intergrowths on theirsurfaces. The second population is that highlighted in Figure1a by the circle; these are small diamond-shaped particles seenpreviously when calcium is present.33,44 The dendritic particleshapes are consistent with that found by Benton et al.;23

however, an interesting point can be raised. The particlesformed in the SSW are at an SI of 2.5, yet when compared toBarite crystallized at pH 8 (the pH of the SSW is 8.1) and thesame SI with a 1:1 Ba2+/SO4

2− ion ratio, the particles look quitedifferent (see Figure 1b). Further work was conducted todetermine the root cause of this difference.The possibility of the ions determining the dendritic shape

was investigated by separating them into different groupings.

Thus, these groupings were the monovalent ions, the divalentions, and borate + carbonate.Monovalent ions are thought to promote crystal

growth;30−32 thus, their presence could influence the formationof the dendritic particles. The presence of these ions at pH 8 ata cation to anion ratio of 1:1 and equivalent SI show a variety ofparticle morphologies (Figure 2), none of which are dendriticin nature. It is clear that while increased ionic strength mayaccount for a lengthening of the c-axis (see circled particles inFigure 2), it does not account for the dendritic growth.When crystallization of barium sulfate occurs in the presence

of borate + bicarbonate at pH 8, the morphology is not thatdissimilar to the pH 8, SI = 2.5 particles formed (compare

Figure 7. SEM image of barium sulfate formed at 90 min in SSW and EDS spectra at the different spots shown (A, B, C) in the SEM image

Figure 8. Derived counts (kilocounts per second) versus time for the SSW system (natural pH 8, Ba2+/SO42− = 0.012), pH 8 in pure water with a

ratio of Ba2+/SO42− = 1, and pH 8 in pure water with a ratio of Ba2+/SO4

2− = 0.169.

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Figure 3a to Figure 1b). This no doubt reflects the lowconcentration of these ions in SSW thus having minimalimpact. When divalent ions are present, roughened particles areformed similar to the particles formed in the presence ofbicarbonate + borate ions but with a more 1:1 aspect ratio(Figure 3b). These particles show the roughening observed inthe dendritic particles but not the gross shape. Energydispersive X-ray spectroscopy (EDS) analysis of the bariumsulfate particles formed in SSW showed that Sr2+ ions wereindeed present (Supporting Information, SFigure 1). Thissuggests that the formation of the dendritic particles is not duesolely to the presence of these ions. Benton et al.23 proposedthat the dendritic shape of the particles was due to highsupersaturation and a low barium/sulfate ion ratio. Themorphology at pH 8 versus SI was, therefore, investigated(see Figure 4).It is seen that the dendritic barium sulfate formed in seawater

is similar in morphology to particles formed at higher SI. Thus,SI does appear to be an important parameter in the formationof dendritic particles.The ion ratio was also investigated. This involved performing

the experiments at pH 8 and variable SI (= 2, 2.5, and 3) witheither varying barium or sulfate ion concentrations (high or lowbarium/sulfate ratios). As Figure 5 shows, dendritic-likeparticles are formed when the supersaturation is high enough.Mavredeki et al.29 found that at a lower SI and low ion ratio,more diamond-shaped particles will form, supporting theresults found here (see Figure 5a). This shape is also observedat the molecular level.45

The main conclusion from this is that dendritic particles canform whenever there is a high enough SI. When sulfate is inexcess the particles are more X shaped, while when barium ionsare in excess the particles appear more square like andintergrown. Even when the ion ratio is 1:1 dendritic particlesare possible provided the SI is sufficiently high.It can be seen that the particles formed in SSW have some

characteristics similar to those formed in the presence of excesssulfate (Figure 5e) and at high supersaturation ratio (Figure4d), but the particles formed in the presence of SSW are muchmore rounded (see Supporting Information, SFigure 2) and

neither the “pure” (changing SI or ion ratio) systems nor thoseof the ion groupings can explain the morphology adequately.Thus, the formation of the dendritic particles is duepredominantly by the SI first and foremost and to some extentthe ion ratio (as proposed by Benton et al.23 and found byKowacz et al.46) but also to a synergistic effect with the ionspresent.Timed experiments were undertaken to determine how the

two populations of Barite particles in SSW formed. At 15 min(Figure 6a) very few diamond-shaped particles are formed(shown by arrows), and the vast majority of particles aredendritic in nature suggesting the dendritic particles are formedfirst when the supersaturation is high. By 30 min, many morediamond-shaped particles are clearly seen, suggesting theseparticles nucleate later than the dendritic particles at lowersupersaturation. Perhaps more interestingly, the impurityassociated with these particles shows an interesting trend asdetermined by EDS (Figure 7).At 90 min, the particles formed are shown in Figure 7. These

particles span the range from purely diamond-shaped todendritic and to shapes in between. EDS on the dendriticparticles shows Sr present (Spot and EDS spectrum A). TheEDS spectrum from an almost diamond-shaped particle showsthat calcium and not strontium is present (Spot and EDSspectrum B). This is expected from the previous ion ratioresults (Figure 5e) and because calcium ions are known toproduce a diamond-shaped morphology at high enoughconcentrations.33,44 Finally, a particle in between these twoextremes shows both calcium and strontium are present (Spotand EDS spectrum C). This suggests that strontiumincorporation/substitution occurs in dendritic particles firstand as time goes on, smaller diamond-shaped particles nucleate,which initially have calcium ion incorporation/substitution.Over time, as these small diamond-shaped particles grow,strontium ion incorporation/substitution is able to occur inthese particles also. This suggests that strontium ionincorporation/substitution is preferred on the fast growingarms of the dendrites (which appears to be the b axis of Barite).Particles of pure strontium sulfate can also show morphologies

Figure 9. Derived counts versus time for barium sulfate formed at pH 8 (Ba2+/SO42− 1:1) with different ions present.

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such as those in Figure 7, spot C (see Supporting Information,SFigure 3).DLS. In addition to morphological investigation of the

system, we undertook to investigate the impact SSW and theion groupings had on nucleation of Barite. The derived countsfrom the DLS investigation versus time can be used as anindicator of nucleation. That is, in the ideal system (noaggregation or settling) the derived counts will increase due tonucleation events until supersaturation is depleted and growthtakes over, and the derived counts then should remain steady.In classical nucleation theory, collisions are required to achievethe metastable “clusters” at their critical size; thus it is expectedthat the nucleation rate will be related to the hydrodynamicconditions (and therefore the stirring rate) of the system underinvestigation. Reproducibility of the data, therefore, wasensured by using the same stirrer as well as the same stirringspeed (see Supporting Information, SFigure 4). The obtainedderived counts data (Figure 8) suggest that aggregation or

settling does occur, particularly as time progresses (counts donot remain steady but decrease over time). However, there isalso a significant difference between the pH 8 system at 1:1 andthe SSW system. The results suggest that the nucleation rate ishigher for the 1:1 system despite the similar SI. The nucleationrate is often given as (in classical nucleation theory):

β γ= Ω − ΦJ v kTexp( / )2 3 2(4)

where Ω is the pre-exponential factor, β and v are the shapefactors and the molecular volume, respectively, k is theBoltzmann constant, T is temperature, γ is the surface freeenergy and Φ is the supersaturation. Assuming the pre-exponential term does not change and since the initialsupersaturation (and T) is constant for these experiments,this implies J is changing due to changes in the surface freeenergy. The cause of the increase in the surface free energycould be due to a larger critical nucleus for nucleation toproceed (assuming classical nucleation theory). Reducednucleation rates are also observed at a low barium to sulfateion ratio as found by Kowacz et al.46 Having said this, there isno consistent decrease in the nucleation rate as the barium ionratio decreases with respect to sulfate. The nucleation rate forthe pure 0.169 ion ratio experiment was similar but slightly lessthan that for the SSW case where the ion ratio is lower at 0.012.This may be because the nucleation rate plateaus on furtherdecreases in the ion ratio, or it may be that in SSW other ionsincrease nucleation.The presence of divalent ions (Figure 9) shows the

nucleation rate is higher than that observed for the 1:1 systemat pH 8. When divalent ions are present the counts perhaps stillincrease at >90 min. Thus, nucleation may still be occurringthroughout the 120 min of investigation. With reference to thenucleation rate equation (eq 4), the presence of ions such asSr2+ and Ca2+ may form a more soluble solid with a lowersurface free energy. Certainly both SrSO4 and CaSO4.2H2O aremore soluble sulfates than Barite, and their substitution into thelattice could potentially form a more soluble Barite leading to ahigher nucleation rate as suggested by Monnin.4 This wouldsupport the hypothesis that in SSW the presence of the divalentions will increase nucleation relative to the pure system havingthe same ion ratio.The presence of borate + carbonate shows a similar

nucleation rate to the 1:1 pH 8 system. This supports themorphology experiments, whereby due to the low concen-trations of these impurities, little impact on nucleation rate wasobserved relative to the control system. For the monovalentions it can be seen that their presence results in a lowernucleation rate. This is an unusual result given that previouswork has suggested an increased 2D nucleation rate whenmonovalent ions are present.30−32 This may be due toaggregation of the particles, leading to significant settling inthe system and an undercounting of particles. This resultrequires further investigation.

Barite in SSW+Silicate. In the presence of silicate (2.82mM) the morphology of the barium sulfate formed in SSW isnot vastly altered (see Figure 10). The main difference appearsto be the length/width ratio of the dendritic particles (thedendritic particles look longer generally), the larger number ofdiamond-shaped Barite, and the presence of an amorphoussilicate layer, seen more clearly in the backscatter image (Figure10b). The EDS clearly shows the presence of Si, while Sr maybe missed due to the closeness of the Sr/Si peaks. In addition

Figure 10. (a) SEM image of barium sulfate formed in the presence ofSSW and silicate (2.82 mM), (b) the backscatter image emphasizingthe amorphous silicate component, and (c) the EDS spectrum fromthis region.

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when silicate is present, more impurity ions appear such as Na,Ca, and Mg.The presence of silicate in solution did not alter the

nucleation behavior of the SSW system significantly (Figure11). The hypothesis put forward by Pina and Tamayo19 is thatsilicate polymerizes at lower pHs and promotes heterogeneousnucleation. This is supported by these results because if theimpact of silicate is to increase heterogeneous nucleation onpre-existing nuclei this will not impact on particle counts to asignificant degree.The solids formed in SSW at different silicate concentrations

were fully characterized using powder XRD and TGA. TheTGA data shows that as more silicate is present, more waterappears to be associated with the solids (see SupportingInformation, SFigure 5). This was also observed at pH 1022 andwould explain the increased impurity ions associated with thesolids when silicate is present. In terms of the XRD data(Figure 12), the peaks were matched to Barite except for thoseat ∼20° 2θ where a doublet was expected and a triplet isobserved (see Figure 12b). The best match for this solid was amixed Ba1−xSrxSO4 solid where x = 0.25 (powder diffraction filenumber: 04-007-7651), although we suspect calcium is alsosubstituted to some extent. As silicate is introduced, the peaksshift, and at 4.6 mM sodium metasilicate addition, a broad XRDpattern is observed. Many peaks still correspond to aBa1−xSrxSO4 solid in this broad pattern but two peaks (at 2θ= 28 and 46°) are unaccounted for.Rietveld refinement of the XRD patterns was undertaken and

the lattice parameters calculated (calculated and differencepatterns can be seen for 0 mM sodium metasilicate in theSupporting Information, SFigure 6). The lattice parameters

derived from the Rietveld refinement were plotted against thesilicate concentration and are shown in the SupportingInformation (SFigure 7). For the a and c lattice parameter,there is no distinct trend with silicate concentration, though ageneral decreasing trend is observed. However, for the b latticeparameter there appears to be some concentration dependentbehavior. The impact on the b lattice parameter appears to besmall until high silicate concentrations are reached. It should bestressed that although these reactions were scaled up to 4 L, thesolids formed are in the milligram region and as such had to beperformed on low background holders; this is not ideal forRietveld refinement and could be the source of some error.Thus, these results require further validation.Since the lattice parameters change with silicate concen-

tration, this suggests that either more Sr (and/or Ca33) will beincorporated into the Barite structure or that the silicate is alsoincorporating as seen at higher pHs.22 If this is the strontiumion further incorporating it is possible that this is a kineticphenomenon meaning that the presence of silicate promotes ametastable phase (Sr1−xBaxSO4) before converting to thethermodynamic product, Ba1−xSrxSO4. The presence of silicateon the surface of the metastable phase may also stop it fromdissolving and reforming Barite as is found when silicate stopsferrihydrite from converting to goethite.47−49 However, thesystem is not saturated with respect to celestite (SrSO4). Inaddition, significant silicate substitution into Barite was notobserved at a pH of 7;22,19 thus, it is unlikely that this isoccurring either.Another possibility is that the presence of Si may increase the

local Sr concentration as the silicate levels rise, and this theninteracts with sulfate such that the formation of barium sulfate

Figure 11. DLS derived counts versus time for barium sulfate formed in SSW and different silicate concentrations.

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with significant strontium ion incorporation occurs. Somesupport for this hypothesis is that Sr-carbonate complexation isstronger than Ba-carbonate complexation.50 If the same is truefor Sr-silicate complexation compared to Ba-silicate complex-ation, this could be a possible mechanism for increasedstrontium substitution. A similar behavior was also observedwith barium carbonate in the presence of silicate when calciumions were present.21,51 As calcium ion concentration wasincreased, calcium carbonate rather than barium carbonateformed despite the lower solubility of barium carbonate. Assilicate concentrations increase, it is possible that strontiumsilicate formation becomes dominant due to the local highconcentration of strontium, and this in turn could lead to morestrontium incorporation into the Barite structure.To gain further insight into the strontium ion substitution

into Barite, empirical molecular modeling was undertaken andcompared to the XRD results. The results from the simulationshow that the lattice parameters change as expected with the

degree of Sr2+ substitution (see Supporting Information,SFigure 8). Because of the smaller size of Sr2+ compared toBa2+, the lattice parameters and unit cell volume all decreasewith the degree of substitution. The b-axis unlike the otherparameters shows a nonlinear trend with Sr2+ incorporation.Since the a and c lattice parameters should decrease withincreasing Sr incorporation, comparison with the XRD datashows that the degree of substitution is somewhat variable fromsample to sample, but that the general decreasing trend isconfirmed. The nonlinear trend with the b lattice parameter isalso observed from the Rietveld analysis of the XRD data.The replacement energy is found to be more and more

positive (i.e., less energetically favorable, see SupportingInformation, SFigure 9) the more Sr2+ is incorporated intothe structure. Given the thermodynamic stability of Baritecompared to celestite, this appears reasonable. Calculating adefect energy according to

Figure 12. XRD patterns of (a) barium sulfate formed in SSW at different sodium metasilicate concentrations and (b) close up of 19−22° 2θ region(patterns have been displaced in the y direction for easier comparison).

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= −E E E n( )/defect substituted initial (5)

where n is the number of barium ions substituted shows thatthe defect energy is actually more negative the more strontiumions are substituted into the structure, to a limiting value of−1.188 eV. Thus, provided there is a driving force fordehydrating the strontium ion, incorporation into the Baritestructure should be facile. Finally, the simulations that gave theclosest XRD pattern to that found to the experimental SSWpattern were superimposed (see Supporting Information,SFigure 10). They ranged from ∼20−30% Sr substitution,and this agreed reasonably well with the database match ofBa75Sr25SO4.Consequences for Marine Environment Proxies.

Strontium isotopes are often used in seawater systems as areference system to measure other isotopic fractionation buthas been found to be fractionated toward the lighter isotopesdue to kinetic effects occurring during crystallization.52 It is alsoimportant to remember that Barite is often used as a proxy forocean productivity.53 As first noted by Dymond et al.53 andlater by others,54,55 the export of barium from the upper oceanhighly correlated with organic carbon. The results presentedhere show that the presence of silicate may increase the Sr2+ ionsubstitution into Barite. Incorporation of foreign ions into solidlattices changes their solubility53 perhaps leading to anunderestimation if they dissolve away. For example, the Kspof pure Barite is ∼9.9, while Barite formed from seawater isfound to have a Ksp of ∼8.1, at least an order of magnitudemore soluble.52 Another possibility is that if silicate is able toprotect the solids from the solution as mentioned above foriron (hydrox)oxides,47−49 the solids will be less soluble andremain even in undersaturated solutions (perhaps explainingthe Barite paradox4,5). Finally, the presence of silicate may alsocause kinetic fractionation of the Sr2+ isotopes if thesubstitution into Barite is kinetically controlled. The solubilityand isotopic fractionation aspects are important missing piecesof information, which are currently being investigated.

■ CONCLUSIONSBarium sulfate formation in seawater has a dendriticmorphology at the supersaturation investigated here (SI =2.5). This morphology is determined by the SI and the cation/anion ratio but is also due to synergistic effects with the otherions present. It appears that, in this batch system, the dendriticparticles form first with a significant degree of strontiumsubstitution followed by nucleation of the diamond-shapedparticles with significant calcium substitution. Over time, thediamond-shaped particles grow and also incorporate strontiumions. In terms of the nucleation behavior, the presence ofdivalent cations appears to lower the surface free energy, andthis may explain the nucleation of the diamond shaped particleseven after the supersaturation has decreased on formation ofthe dendritic particles.When solution silicate is present in addition to the SSW, the

impact on Barite morphology is minimal. Longer particlesappear to form with a greater variety of impurity ions present.No significant increase in the nucleation rate is evident whensilicate is present, but this may be due to heterogeneousnucleation.Finally, as would be expected, strontium ions are

incorporated to a significant degree into the lattice. However,what is perhaps more significant is that the presence of solublesilicate appears to promote further incorporation. This could be

due to Sr-silicate interactions. If silicate does induce furtherstrontium incorporation into the Barite lattice, this would haveconsequences both for the solubility of the resulting solidformed and possibly also for the isotopic abundance ofstrontium present. These two aspects are the focus of furtherstudies.

■ ASSOCIATED CONTENT*S Supporting InformationThe Supporting Information is available free of charge on theACS Publications website at DOI: 10.1021/acs.cgd.6b00729.

SEM images of dendritic Barite particles formed in SSW,EDS spectrum of barium sulfate formed in SSW, SEM ofstrontium sulfate particles, DLS results versus stirringrate, TGA results for barium sulfate formed in SSW andsilicate concentrations, XRD pattern with Rietveldrefined and difference pattern, change in the latticeparameters determined from Rietveld refinement versussilicate concentration and different silicate concentra-tions, simulation results for barium sulfate withincorporation of Sr, simulated XRD (Ba33Sr15SO4), andexperimental XRD (for solids obtained from SSW)superimposed on each other (PDF)

■ AUTHOR INFORMATIONCorresponding Author*E-mail: F.Jones@curtin.edu.au. Phone: (618) 9266 7677. Fax:(618) 9266 2300.

NotesThe authors declare no competing financial interest.

■ ACKNOWLEDGMENTSWe thank Peter Chapman for doing the thermal analysis andWannapa Boonwannich for use of the SrSO4 image, and wewish to acknowledge the Curtin Centre for Materials Researchfor use of the SEM and XRD facilities.

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