Influencing Crystallization Damage in Porous Materials through the

Properties, characterization, and decay of sticky rice–lime mortars from the Wugang Ming dynasty city wall (China). Ya Xiao , Xuan Fu , Haibing Gu ,...
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Langmuir 2000, 16, 947-954

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Influencing Crystallization Damage in Porous Materials through the Use of Surfactants: Experimental Results Using Sodium Dodecyl Sulfate and Cetyldimethylbenzylammonium Chloride Carlos Rodriguez-Navarro,*,† Eric Doehne,† and Eduardo Sebastian‡ The Getty Conservation Institute, 1200 Getty Center Drive, Suite 700, Los Angeles, California 90049, and Departamento de Mineralogı´a y Petrologı´a, Universidad de Granada, Fuente Nueva s/n, 18002 Granada, Spain Received May 12, 1999. In Final Form: September 20, 1999 The interactions of two ionic surfactants on the decay of ornamental stone (porous limestone) by salt crystallization, a common and damaging weathering process, were studied. Conductivity and/or surface tension measurements allowed calculations of the critical micellar concentration (cmc) of sodium dodecyl sulfate (SDS) and cetyldimethylbenzylammonium chloride (CDBAC) in distilled water and saturated sodium sulfate solution (both with and without the addition of calcite), total surfactant adsorption onto calcite (Γtot), and the area (As) which a surfactant molecule occupies at the liquid-air interface. In aqueous calcite suspensions SDS shows a strong cmc reduction due to Ca2+ binding to the micelles that undergo sphereto-rod transition at SDS concentrations > cmc, while calcium dodecyl sulfate (Ca(DS)2) precipitation contributes to a reduction of DS- concentration in the bulk solution. Adsorption of DS- on calcite is promoted in the saturated saline solution where Γtot reaches values of 1 × 10-2 mmol m-2. CDBA+ is preferentially adsorbed onto calcite in water (Γtot ) 1.89 × 10-3 mmol m-2) while adsorption on calcite in saturated sodium sulfate solution is limited (Γtot ) 2.18 × 10-4 mmol m-2) due to competition with Na+ for calcite adsorption sites. Limited CDBA+ adsorption onto calcite, and significant As reduction in the saline solution, results in sphere-to-rod (or disk) micelle shape transition. The previous results together with in situ, high magnification environmental scanning electron microscopy (ESEM) studies and macroscale salt crystallization tests revealed that: (a) the adsorption behavior of DS- induces crystallization of mirabilite at high supersaturation, resulting in nonequilibrium crystal shapes that promote significant damage to the stone; (b) rodlike and/or disc-shaped CDBA+ micelles enhance solute solubilization and transport to mirabilite nuclei growth sites, inducing crystallization of euhedral crystals formed at low supersaturation and distributed homogeneously throughout the stone pore system. While CDBAC initially reduces stone damage by salt crystallization, it can ultimately result in enhanced damage when in contact with water due to mechanical weakening of the stone and rehydration of previously dehydrated mirabilite crystals within the stone pores. The implications of these results in the conservation of ornamental stone are discussed.

1. Introduction In the last few decades, the increasing awareness of the need for preserving our cultural heritage, particularly architectural and sculptural stone, has led to the examination of surfactants as new conservation materials.1 Surfactants in conservation are used as biocides (alkylammonium salts2-4), cleaning agents (ionic and nonionic surfactants2,5,6), surface-protection agents (cationic surfactants7,8), clay-stabilizing agents (alkylammonium salts9) and stone desalination agents (cationic and anionic surfactants10). Additionally, biosurfactants, naturally * Corresponding author. Presently at the Universidad de Granada. Phone: 34 958 243340. Fax: 34 958 243368. E-mail: carlosrn@ goliat.ugr.es. † The Getty Conservation Institute. ‡ Universidad de Granada. (1) Snethlage, R.; Wendler, E. Mater. Res. Symp. Proceedings 1991, 185, 193. (2) Lazzarini, L.; Tabasso, M. L. Il Restauro della Pietra; CEDAM: Padova, Italy, 1986; Chapter 4. (3) Monte, M.; Nich, D. Sci. Technol. Cultural Heritage 1997, 6, 209. (4) Tiano, P.; Tosini, I.; Rizzi, M. Sci. Technol. Cultural Heritage 1997, 6, 129. (5) Buys, S.; Oakley, V. Conservation and Restoration of Ceramics; Butterworth and Heinemann, Oxford, U.K., 1993; p 243. (6) Warren, J. Conservation of Brick; Butterworth and Heinemann: Oxford, U.K., 1999; Chapter 10. (7) Wendler, E.; Klemm, D. D.; Snethlage, R. In Proceedings of the 5th International Conference on the Durability of Building Materials and Components; Clarendon Press: Brighton, U.K., 1991; p 203.

occurring surfactants resulting from biological activity,11,12 are found in natural environments13,14 and deteriorating ornamental stone.15,16 To date, some research has been dedicated to the use of surfactants in the conservation of sculptural-architectural stone,10,17 however, fundamental knowledge with respect to the interaction of surfactants with the stone surface and its implications on the decay processes is lacking. (8) Simon, S.; Boehm, H. P.; Snethlage, R. In Proceedings of the VIIth International Congress on Deterioration and Conservation of Stone; Delgado-Rodrigues, J., Henriquez, F., Jeremias, J. T., Eds.; LNEC: Lisbon, Portugal, 1992; p 851. (9) Wendler, E.; Charola, A. E.; Fitzner, B. In Proceedings of the 8th International Congress on Deterioration and Conservation of Stone; Riederer, J., Ed.; Berlin, Germany, 1996; p 1159. (10) Puehringer J.; Engstro¨m L. In Proceedings of the Vth International Congress on Deterioration and Conservation of Stone; Fe´lix, G., Ed.; E Ä cole Polytechnique Fe´de´rale de Lausanne: Lausanne, Switzerland, 1985; p 241. (11) Prescott, L. M.; Harley, J. P.; Klein, D. A. Microbiology, 3rd ed.; Brown Publishers: Dubuque, IA, 1996; p 925. (12) Okubo, T.; Kobayashi, K. J. Colloid Interface Sci. 1998, 205, 433. (13) West, C. C.; Harwell, J. H. Environmental Sci. Tech. 1992, 26, 2324. (14) Folk, R. L. J. Sedimentary Petrology 1993, 63, 990. (15) De la Torre, M. A.; Gomez Alarcon, G.; Vizcaino, C.; Garcia, M. T. Biogeochemistry 1993, 19, 129. (16) Wakefield, R. D.; Jones, M. S. Q. J. Eng. Geo. 1998, 31, 301. (17) Price, C. A. Stone Conservation: An Overview of Current Research; The Getty Conservation Institute: Los Angeles, CA, 1996.

10.1021/la990580h CCC: $19.00 © 2000 American Chemical Society Published on Web 12/04/1999

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Weathering of ornamental stone occurs through the singular or combined action of different decay phenomena: e.g., dissolution, hydrolysis, oxidation, freeze-thaw, physical erosion, thermal expansion, biological activity, and salt crystallization.18 In particular, pressure generation in confined spaces (i.e., pores) due to crystallization of soluble salts (so-called salt weathering19) has been found responsible for significant damage to porous stones in buildings,18-22 as well as in geologic settings both on Earth22 and possibly on other planets (i.e., Mars23). Recently, applications of surfactant solutions to salt contaminated stones have been proposed as a potential treatment to mitigate the effects of salt weathering.10,17,24 Surfactants can induce significant kinetic and morphological changes in crystallizing salts, mostly due to preferential adsorption at particular crystal faces.25 Additionally, adsorption of surfactants at the liquid-solid interface induces wetability changes on various minerals (e.g., calcite26-28) therefore changing the transport/flow of saline solution through stone pores.29 Surfactants may reduce the solution-stone contact angle resulting in faster capillary transport, inducing salt crystallization on the stone surface10 and promoting the formation of less harmful efflorescence as opposed to harmful subflorescence.20 Surfactants may also hinder (or promote) salt crystallization (e.g., crystallization inhibitors30). However, little experimental data are available to explain the action of surfactants in stone desalination, as well as possible side effects connected to their use in common conservation treatments (i.e., cleaning). Salts, on the other hand, cause changes in the behavior and performance of the surfactant system (e.g., reduction of cmc,31 and changes in micelle shape32) a fact that deserves more fundamental research as well. The aim of this work is to study the effects of two types of ionic surfactants (anionic and cationic) on the process of salt crystallization in a porous calcareous stone, with a three-fold goal: (a) a better understanding of surfactant chemistry at the liquid-solid interface (calcite-solution) particularly in high ionic strength solution (i.e., saturated salt solutions); (b) determination of the role of surfactants in salt crystallization processes; and (c) evaluation of their use in desalination or mitigation treatments for porous ornamental stones affected by salt weathering. (18) Sebastian, E.; Rodriguez-Navarro, C. Ingenierı´a Civil 1995, 96, 167. (19) Evans, I. S. Revue Ge´ omorphologie Dynamique 1970, 19, 155. (20) Rodriguez-Navarro, C.; Doehne, E.; Ginell, W. S.; Sebastian, E. In Proceedings of the 3rd International Congress on Restoration of Architectural Heritage and Building; Sebastian, E., Valverde, I.; Zezza, U., Eds.; Universidad de Granada: Granada, Spain, 1996; p 509. (21) Rodriguez-Navarro, C.; Doehne, E. Earth Surf. Processes Landforms 1999, 24, 191. (22) Goudie, A. S.; Viles, H. Salt Weathering Hazards; John Wiley and Sons: Chichester, U.K., 1997. (23) Rodriguez-Navarro, C. Geophysical Res. Lett. 1998, 25, 3249. (24) Puehringer, J. Salt Migration and Degradation by SaltsA Hypothesis; Swedish Council for Building Research: Stockholm, Sweden, 1983. (25) Adamson, A. W.; Gast, A. P. Physical Chemistry of Surfaces, 6th ed.; John Wiley and Sons: New York, 1997; Chapter 8. (26) Goujon, G.; Cases, J. M.; Mustaftschiev, B. J. Colloid Interface Sci. 1976, 56, 587. (27) Somasundaran, P.; Goddard, E. D. Modern Aspects Electrochem. 1979, 13, 207. (28) Thomas, M. M.; Clouse, J. A.; Longo, J. M. Chem. Geology 1993, 109, 201. (29) Tadros, T. F. Surfactants; Academic Press: London, 1984. (30) Cody, R. D. J. Sedimentary Petrology 1991, 61, 704. (31) Armstrong, J. K.; Chowdhry, B. Z.; Snowden, M. J.; Leharne, S. A. Langmuir 1998, 14, 2004. (32) Dutt, G. B.; van Stam, J.; De Schryver, F. C. Langmuir 1997, 13, 1957.

Rodriguez-Navarro et al.

2. Materials and Methods Two common ionic surfactants were selected for this study: (a) reagent grade anionic sodium dodecyl sulfate (SDS, CH3(CH2)11OSO3Na, m.w. 288.4 Sigma); and (b) cationic cetyldimethylbenzylammonium chloride (CDBAC, CH3(CH2)15NCl(CH3)2C6H5, m.w. 413, Pfaltz & Bauer). They were used as provided by the manufacturer. A stock saturated (1.37 M) sodium sulfate (reagent grade Na2SO4, Sigma) solution was prepared. Calcite (reagent grade CaCO3, Baker) suspensions were prepared by adding 0.5 g of calcite, with 2.75 m2g-1 BET surface area, to either 100 cm3 distilled water, or 100 cm3 saturated sodium sulfate solution. Critical micellar concentrations (cmc’s) in distilled water and saturated sodium sulfate solution, with and without the addition of calcite, were determined by conductimetry33,34 (Orion Model 160 Conductivity Meter) and/or surface tension measurements25,35 (pullout ring method: DuNouy Tensiometer, model CSC No. 70535). All measurements were performed at 20 (1 °C and 50 ( 5% relative humidity. The pH was measured using a Hanna pH meter with ( 0.01 pH units accuracy. Advancing contact angles were obtained by taking still images with a video camera (coupled with 50× lenses) of sessile drops deposited using a syringe on a polished calcitic marble surface placed on an optical bench. The height h and the contact diameter d of the drop were measured, and the contact angle θ was calculated as follows

(2θ) ) 2hd

tan

(1)

Viscosities of sodium sulfate solutions with and without surfactant added (micellar concentration) were determined using an Ubbelohde-type viscometer (Koehler Instrument Company) kept in a water bath at 20 ( 0.2 °C, and calibrated with respect to distilled water. Evaporation rates of solutions were determined by weight loss measurements following water evaporation at constant temperature (20 ( 1 °C) and relative humidity (50 ( 5%) in open beakers of equal diameter. Maximum surfactant adsorption onto calcite Γtot, both in the absence (distilled water) and in the presence of electrolyte (1.37 M Na2SO4 solution), was calculated from cmc data as follows

Γtot )

[cmc]Cc - [cmc]sol [Cc]‚S

(2)

where [cmc]Cc and [cmc]sol are the surfactant cmc’s (mmol dm-3) in calcite suspension and in the solution (distilled water or 1.37 M Na2SO4 solution), respectively; [Cc] is the calcite concentration (g dm-3) and S is the calcite specific surface area (m2 g-1). The area, As, which a surfactant molecule occupies at the liquid-air interface at saturation adsorption was calculated from surface tension measurements using the Gibbs adsorption isotherm36

AS ) -

RT 1 NA dγ/d lnc

(3)

where γ is the surface tension, c is the surfactant concentration, NA is the Avogadro’s number, and R and T are the gas constant and the temperature, respectively. Phases formed in the calcite suspensions in the presence of surfactant were analyzed by X-ray diffraction (XRD, Philips model PW 1710) and scanning electron microscope (SEM, Zeiss model DMS 950) coupled with a microanalysis EDX (Link QK 2000). Sodium sulfate crystallization in a porous calcareous stone (ornamental mid-Jurassic oolitic limestone from Monks Park, U.K., which has been shown to be very susceptible to salt weathering21) as well as the changes induced by the presence of each surfactant were observed in situ, at high magnification, (33) Dominguez, A.; Fernandez, A.; Gonzalez, N.; Iglesias, E.; Montenegro, L. J. Chem. Education 1997, 74, 1227. (34) Perez-Rodriguez, M.; Prieto, G.; Rega, C.; Varela, L. M.; Sarmiento, F.; Mosquera, V. Langmuir 1998, 14, 4422. (35) Stumm, W.; Morgan, J. J. Aquatic Chemistry: Chemical Equilibria and Rates in Natural Waters, 3rd ed.; John Wiley and Sons: New York, 1996; Chapter 9. (36) Schott, H. J. Colloid Interface Sci. 1998, 205, 496.

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Table 1. pH and cmc Values (mmol dm-3) of SDS and CDBAC Surfactants in the Various Solutions and Calcite (Cc) Suspensions Tested cmc

a

pH

solution/suspension

SDS

CDBAC

control

SDSa

CDBACa

water Na2SO4 saturated solution Cc (in water) Cc (in Na2SO4 saturated solution)

8.00 0.52 1.50 1.90

0.170 0.019 0.430 0.049

5.66 7.20 8.30 9.12

5.93 7.02 9.60 9.19

7.07 7.35 7.75 8.94

Concentration of surfactant ) cmc.

using an environmental scanning electron microscope (ESEM, Electroscan model E-III). The ESEM is especially useful in the study of surfactant systems because samples can be studied in their natural state, even when they contain water.37 Furthermore, the possibility of controlling pressure (between ∼ 2 up to 20 Torr) and temperature (using a Peltier stage) in the ESEM chamber, and in particular in the sample surroundings, allows condensation/evaporation cycles to be performed. Thus, in situ dynamic dissolution/precipitation studies of saline systems can be performed, and salts with hydrates (i.e., sodium sulfate decahydrate) can be directly observed.38 Recent papers describe the fundamentals of this technique and its applications for in situ studies of various systems and processes (e.g., emulsions39 and salt crystallization40). Three saline solutions were prepared: (a) blank, pure saturated sodium sulfate solution; (b) 2 mmol dm-3 SDS in saturated sodium sulfate solution; and (c) 0.8 mmol dm-3 CDBAC in saturated sodium sulfate solution. Monks Park oolitic limestone samples (2 × 2 × 2 mm) were immersed in each saline solution before being placed on a steel sample holder, and this was immediately introduced into the ESEM chamber and placed on a Peltier stage. Full dissolution of the salt was maintained in the ESEM chamber at a pressure of 6 Torr, keeping the sample holder at a constant temperature of 8 °C (using the Peltier stage). Evaporation and salt crystallization were promoted by reducing the pressure to 3 Torr. Additional dissolution/crystallization cycles were performed by raising the pressure up to 6 Torr and reducing it down to 3 Torr. The whole process was recorded using a VCR. Still images were also taken. The pressure reduction rate, which controls the evaporation rate and, therefore, the supersaturation ratio and crystallization rate,20-21 was a low 0.2 Torr/min. Macroscale salt crystallization tests using Monks Park limestone were also performed. Limestone blocks (3 × 3 × 28 cm) were placed vertically in 300 cm3 glass beakers filled with the above-mentioned saline solutions (pure and with surfactants added). The solution surface was covered with melted paraffin to avoid excessive evaporation and/or creeping of salts. Details on the experimental setup for the macroscale salt crystallization test can be found elsewhere.21,40 Following the capillary rise of the solution through the stone pore system, evaporation and crystallization occurred. The resulting damage to the stone (i.e., cracks, flakes, and granular disintegration) was recorded photographically. After total loss of the liquid from the beakers, the stone blocks were removed and immersed in distilled water to eliminate the remaining salts. The total weight loss due to granular disintegration and scaling following salt crystallization was determined. Small pieces (0.5 × 0.5 × 2 cm) were cut from the decayed stone blocks (before salt extraction) and their porosity and pore size distribution were evaluated using a mercury intrusion porosimeter (Micromeritic model 3300). The evaporation rate of each solution was evaluated by weighing the beakers (with the solution, stone and paraffin) at constant intervals.

3. Results and Discussion 3.1. Surface Tension and Conductivity Measurements. 3.1.1. The H2O-Surfactant-CaCO3 System. Fig(37) Danilatos, G. Microsc. Res. Tech. 1993, 25, 354. (38) Doehne, E. 1994 In Proceedings of the 3rd International Symposium on the Conservation of Monuments in the Mediterranean Basin; Fassina, V., Ott, H., Zezza, F., Eds.; Graffo: Venice, Italy, 1994; p 143-150. (39) Stokes, D. J.; Thiel, B. I.; Donald, A. M. Langmuir 1998, 14, 4402. (40) Rodrigue-Navarro, C.; Doehne, E. Am. Laboratory 1999, 31, 28.

Figure 1. Conductivity and surface tension vs concentration of SDS (a) and CDBAC (b) in water. cmc’s are indicated with arrows. The dotted arrow corresponds to (secondary) SDS cmc following micelle shape transition. Open squares (surface tension) and open dots (conductivity) correspond to water without calcite added. Full squares (surface tension) and full dots (conductivity) correspond to water with calcite added.

ure 1 shows solution conductivity and surface tension of both distilled water and calcite suspensions versus concentration of SDS (Figure 1a) and CDBAC (Figure 1b). Table 1 shows surfactant cmc values obtained from both conductivity and surface tension measurements (pH values are also indicated). There is a good agreement between cmc values calculated from surface tension and conductivity data. The cmc of SDS in water is in good agreement with published data.34,41 We observed a minimum of SDS surface tension at values below cmc, which can be attributed to impurities.42 Surface tension measurements from Avranas et al.42 show a slightly higher cmc for CDBAC in water (4.9 ( 0.2 mmol dm-3). Figure 1b shows that this latter value fits well with our findings of the surface tension of CDBAC. However the conductivity slope change is more evident, and therefore in this case we used this technique for cmc determination. In the presence of calcite, SDS solution shows a conductivity curve slope change first at 1.5 mmol dm-3, (41) Esposito, C.; Colicchio, P.; Facchiano, A.; Ragone, R. J. Colloid Interface Sci. 1998, 200, 310. (42) Avranas, A.; Malasidou, E.; Mandrazidou, I. J. Colloid Interface Sci. 1998, 207, 363.

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Rodriguez-Navarro et al. Table 2. Calculated Total Adsorbed Surfactant, ΓTot (mmol m-2) onto Calcite

surfactant

Γtot Cc aqueous suspension

Γtot Cc suspension in 1.37 M Na2SO4 solution

SDS CDBAC

a 1.89 × 10-3

1.00 × 10-2 2.18 × 10-4

a cmc reduction in calcite suspension (if compared with SDS cmc in water) makes this value negative, according to eq 2.

Figure 2. Surface tension versus log of concentration of SDS (a) and CDBAC (b) in water and sodium sulfate saturated solution.

Figure 3. X-ray diffraction patterns of SDS (Sigma) and Ca(DS)2 (precipitated in the calcite suspension). Cu KR X-ray radiation, λ ) 1.5469 Å.

and a second at 10 mmol dm-3 (Figure 1a). Surface tension measurements confirm these results. A precipitate in the water-calcite-SDS system was formed at SDS concentration above 1.5 mmol dm-3. XRD (Figure 3) and SEMEDX analyses show this precipitate to be calcium dodecyl sulfate Ca(DS)243,44. The existence of two conductivity and surface tension changes in our experiments is consistent with a sphere-to-rod transition of micelle shape induced by the presence of Ca2+ ions at DS- concentrations > 1.5 mM and < 10 mM. Other studies confirm that multivalent cations such as Ca2+ or Al3+ strongly enhance the formation of rodlike micelles in anionic surfactant solutions due to counterion adsorption at the micelle Stern layer.43-47 In agreement with our results are the findings of Algorva et (43) Kallay, N.; Pastuovic, M.; Matijevic, E. J. Colloid Interface Sci. 1985, 106, 452.

al.,47 showing the transition from sphere to rodlike micelles by a drop in the solution surface tension at concentrations above the cmc. It can be concluded that conductivity, K, measurements are also sensitive to sphere-to-rod transition, and consequently, K can be used as an indicator of such change in micelle shape. Overall, dodecyl sulfate underwent a significant cmc reduction in calcite suspensions. On the other hand, CDBAC shows a significant cmc increase in the aqueous calcite suspension (Table 1), due to enhanced CDBA- adsorption onto calcite. Table 2 shows the total adsorbed surfactant Γtot, calculated according to eq 2. The slightly basic pH (7.7-9.6) of the CaCO3-H2Osurfactant system (Table 1), close to the point of zero charge of calcite (pzc, pH 8.2-9.548), reduces the electrostatic adsorption of both ionic surfactants. However, limited adsorption of DS- onto calcite in water may take place via replacement of any potential determining negative ions in the Stern layer (i.e., HCO3- and CO32-) followed by Ca2+-DS- bonding (through the DS- functional charged -OSO3- groups). However, precipitation of Ca(DS)2 will significantly reduce the concentration of DS- in the bulk solution, therefore limiting DS- adsorption onto calcite. The cationic surfactant can be adsorbed onto calcite via replacement of any potential determining positive ions (i.e., Ca2+, CaHCO3+, and CaOH+) adsorbed on calcite surfaces, followed by bonding of CDBA+ to carbonate ions. 3.1.2. The Na2SO4-H2O-Surfactant-CaCO3 System. cmc’s of both surfactants are reduced to very low values in the saline solution (Figure 2 and Table 1). Reduction of cmc values,31,49,50 changes in micelle aggregation number,32 and transitions from sphere to rodlike46 or disk shaped51 micelle structures, have been observed in the presence of electrolytes. The cmc reduction is due to the binding of the counterion into the micelle, which thereby reduces electrostatic repulsion among charged ionic surfactant heads46 and allows micelles to form at lower surfactant concentrations. Table 3 shows As values of the various surfactant solutions. As values are reduced in the saline solutions, this reduction being more significant in the case of CDBAC. According to the theory of Israelachvili et al.,52 such As reduction in a saline solution will promote sphere-to-rod (or disk) shape transition of micelles. (44) Baviere, M.; Bazin, B.; Aude, R. J. Colloid Interface Sci. 1983, 92, 1983. (45) Algorva, A. W.; Petrov, J.; Petsev, D.; Ivanov, I. B.; Broze, G.; Mehreteab, A. Langmuir 1995, 11, 1530. (46) Algorva, A. W.; Danov, K. D.; Kralchevsky, P. A.; Broze, G.; Mehreteab, A. Langmuir 1998, 14, 4036. (47) Algorva, A. W.; Danov, K. D.; Petkov, P. A.; Kralchevsky, P. A.; Broze, G.; Mehreteab, A. Langmuir 1997, 13, 5544. (48) Somasundaran, P.; Agar, G. E. J. Colloid Interface Sci. 1967, 24, 433. (49) Rupprecht, H.; Gu, T. Colloid Polymer Sci. 1991, 269, 506. (50) Nevskaia, D. M.; Guerrero-Ruiz, A.; Lopez-Gonzales, J. D. J. Colloid Interface Sci. 1998, 205, 97. (51) Swanson-Vethamuthu, M.; Feitosa, E.; Brown, W. Langmuir 1998, 14, 1590. (52) Israelachvili, J. N.; Mitchell, D. J.; Ninham, B. W. J. Chem. Soc., Faraday Trans. 2 1976, 72, 1525.

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Figure 4. ESEM photomicrographs showing bulky aggregates of hollow faces and prismatic (a) mirabilite crystals in the blank. Prismatic (b) and rombohedral/hopper (c) mirabilite crystals in the sample with SDS. Euhedral mirabilite crystals in the sample with CDBAC (d). Table 3. Calculated Area per Surfactant Molecule, As (Å2) for the Various Solutions and Calcite (Cc) Suspensions

SDS CDBAC

water

Na2SO4 saturated solution

Cc (in water)

Cc (in Na2SO4 saturated solution)

59 132

42 23

61 158

44 24

Therefore, one can expect CDBAC to form a significant number of rod or disk shaped micelles in the saline solution, while this behavior is less evident in the SDS case. cmc values increase when calcite is added (Figure 2 and Table 1). This shift is interpreted as a result of surfactant adsorption onto calcite. While in water the surfactant preferentially adsorbed onto calcite is the cationic, in the saturated sodium sulfate solution it is the anionic (a value 2 orders of magnitude higher) (Table 2). At high pH values the positive counterion, i.e., Na+, will be adsorbed onto the calcite surface negative sites, contributing to a charge reduction of the double layer and decreasing the calcite negative ζ potential.27 Na+ ions will compete with the cationic surfactant for calcite adsorption sites, while the anionic surfactant may easily replace the weak adsorbed SO42- ions, thus resulting in stronger DSadsorption (even though the system pH is 9.19). However, limited CDBA+ adsorption onto calcite in the saline solution can occur through binding to SO42- counterions located in the calcite Stern layer.53 It must be pointed out that even though the overall surface charge of calcite above and below the pzc is negative or positive, respectively, different faces of this mineral can have variations on surface energy and charge distribution.54 Therefore some (53) Saleeb, F. Z.; Hanna, H. S. J. Chem. United Arab Republic 1969, 12, 229.

faces may adsorb a particular ionic surfactant even though the overall charge of the calcite may be of opposite sign. The implications of the above-discussed surfactant behavior on salt crystallization in a porous calcareous stone will be discussed in the following section. 3.2. Salt Crystallization in the Presence of Surfactants. In situ ESEM observations of sodium sulfate crystallization in the porous limestone showed: (a) in the blank stone, crystallization of mirabilite occurred either as bulky aggregates of hollow faced or prismatic crystals (Figure 4a) corresponding to crystallization at high supersaturation (i.e., crystal shapes are far from the equilibrium shape21,55); (b) in the presence of SDS, mirabilite growth occurred as rhombohedral or prismatic (Figure 4b) or isolated hopper crystals (Figure 4c), also indicative of crystallization at high supersaturation ratios; (c) in the presence of CDBAC, sodium sulfate decahydrate crystallized as euhedral, bulky, rhombohedral mirabilite crystals formed at low supersaturation ratios,55 filling the pores of the stone (Figure 4d); (d) dehydration/hydration cycles result in significant damage in the sample with salt and cationic surfactant. Figure 5 shows photographs of the process. Crystallization of mirabilite in large pieces of stone (Figure 6) resulted in scale formation and significant stone loss in the blank (30 wt % loss) and the slab submitted to crystallization in the presence of SDS (32 wt % loss), while reduced damage occurred in the slab submitted to salt crystallization in the presence of CDBAC (28 wt % loss). It is well known that impurities, including surfactants, induce morphology and kinetic changes on the crystallization and growth of a variety of crystals.29,30,56,57 It seems (54) Parker, S. C.; Titiloye, J. O.; Watson, G. W. Philos. Trans. Royal Soc. London A 1993, 344, 37. (55) Sunagawa, I. Bull. Mineralogie 1981, 104, 81. (56) Kubota, N.; Yokota, M.; Mullin, J. W. J. Crystal Growth 1997, 182, 86.

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Figure 6. State of the limestone slabs following 30 days of salt crystallization through capillary suction in the macroscale experiment: (a) blank, (b) sample with SDS added, (c) sample with CDBAC added.

Figure 5. ESEM snapshots of the damage created in the oolitic limestone by the hydration of thenardite to form mirabilite in the presence of CDBAC: (a) before, (b) during, and (c) after hydration.

that the cationic surfactant, being concentrated in the bulk of the saline solution, due to its low adsorption onto calcite, forms a large amount of micelles and may induce nucleation of sodium sulfate at low supersaturation ratios. Tadros29 reports that solute solubilization by micelles results in enhanced transport from the bulk solution toward growing crystal nuclei. Measured As values of CDBAC are consistent with the existence of rodlike or disk shaped micelles in the saline solution, rodlike and disk shaped micelles having a higher solubilization capacity than spherical ones.46 Therefore, a significant enhancement of solute transport toward growing mirabilite nuclei in the presence of CDBAC will occur, resulting in mirabilite crystallization at low supersaturation. On the other hand, DS- ions will be less concentrated in the bulk saline solution due to their preferential adsorption onto calcite, in addition to being precipitated as sparingly soluble Ca(DS)2. Additionally, the limited As reduction in the saline solution with SDS is consistent with limited sphere-to-rod (or disk) transition of micelle shape. Hence the amount of dodecyl sulfate micelles (if any) will be low, and with limited solubilization capacity. In the presence (57) Black, S. N.; Bromley, L. A.; Cottier, D.; Davey, R. J.; Dobbs, B.; Rout, J. E. J. Chem. Soc., Faraday Trans. 1991. 87, 3409.

of SDS, high supersaturation ratios will be reached before crystallization of hollow-shaped or prismatic mirabilite crystals occurs. DS- may have a secondary role as a crystallization inhibitor of mirabilite by being preferentially adsorbed on the initial unstable mirabilite nuclei. For a stable nucleus to form, very high supersaturation ratios must be reached. A similar crystallization inhibition process has been reported for calcite, calcium oxalates, and Ba and Ca sulfates in the presence of various additives.57-62 It seems that the substrate alone (i.e., calcite) does not induce heterogeneous crystallization of mirabilite at low supersaturation ratios, since crystals formed onto calcite in the blank limestone are hollowfaced (i.e., formed at a high supersaturation). This is striking, because it should be expected that a porous, defect-rich, high surface area substrate would not be compatible with a highly supersaturated solution (i.e., heterogeneous nucleation should easily occur). However, high supersaturation ratios may be sustained within small pores63 due to the effect of pressure buildup into capillaries (due to the Laplace effect of curvature of a solution64). The smaller the pore, the higher the supersaturation at which crystallization occurs, when the molar volume of the condensed phase (Vc) is larger than the molar volume of the solute (Vs).64 In the case of sodium sulfate Vc > Vs, therefore, sodium sulfate in solution will be more stable than in a condensed phase at high pressures. When the solution arrives to fractures or the surface, the resulting pressure reduction will induce rapid supersaturation followed by salt crystallization (and damage to the stone). Strong evaporation at the stone surface will also contribute to the rapid increase of the supersaturation ratio.21 The presence of a cationic surfactant dramatically changes this process by inducing early crystallization of equilibrium-shape mirabilite crystals within the pores at low supersaturation values. (58) Tadros, M. E.; Mayes, I. J. Colloid Interface Sci. 1979, 72, 245. (59) Singh, R. P.; Gaur, S. S.; White, D. J.; Nancollas, G. H. J. Colloid Interface Sci. 1987, 118, 379. (60) Fu¨rendi-Milhofer, H.; Tunik, L.; Filipovic-Vincelovic, N.; Skrtic, D.; Garti, N. Scan. Microsc. 1995, 9, 1061. (61) Klepetsanis, P. V.; Koutsoukos, P. G. J. Cryst. Growth 1998, 193, 156. (62) He, S.; Kan, A. T.; Tomson, M. B. Appl. Geochem. 1999, 14, 17. (63) Putnis, A.; Prieto, M.; Fernandez-Diaz, L. Mineralogical Magazine 1995, 132, 1. (64) Kashchiev, D.; van Rosmalen, G. M. J. Colloid Interface Sci. 1995, 169, 214.

Surfactant Effect on Crystallization Damage

Langmuir, Vol. 16, No. 3, 2000 953

Table 4. Solution Viscosity (η), Evaporation Rate (Er), and Calcite-Solution Contact Angle (θ) in the Presence of Ionic Surfactants (Surfactant Concentration ) cmc)a Er

η

θ

solution

control

SDS

CDBAC

control

SDS

CDBAC

control

SDS

CDBAC

water Na2SO4

1.003 1.594

nd 1.625

nd 1.637

8.0 × 10-3 7.6 × 10-3

8.4 × 10-3 8.6 × 10-3

8.2 × 10-3 7.2 × 10-3

42° 37°

nd 23°

nd 20°

a

η viscosity in cStokes; Er in g cm-2 h-1; nd ) not determined.

Figure 7. Evaporation through the stone porous system vs time for the various saline solutions (with and without surfactant added) in the macroscale experiments.

The supersaturation ratio reached when a salt crystallizes is proportional to the crystallization pressure, P (atm), defined as65

P)

C RT ln Vs Cs

(4)

where Vs is the molar volume (cm3 mol-1), C is the actual solution concentration and Cs is the saturation concentration (i.e., C/Cs is the supersaturation ratio), R is the gas constant (cm3 atm K-1 g mol-1) and T is the temperature (K). It can be expected that in the cases of the blank and the saline solution with SDS added, the high supersaturation ratio reached before crystallization start will result in a high crystallization pressure,66 and therefore, it will cause significant damage to the porous stone.21,67 A damage reduction should take place in the case of the saline solution with CDBAC added since the crystal morphologies correspond to lower supersaturation ratios, i.e., lower crystallization pressure. This is consistent with the damage observed following the macroscale salt crystallization experiment. The macroscale salt crystallization experiments showed that the evaporation rate of the saline solution with SDS added was higher than in the blank sample (Figure 7). The stone with the cationic surfactant experienced a significant evaporation rate reduction. Mercury intrusion porosimetry (Figure 8) revealed that both anionic and cationic surfactants induce salt crystallization within the stone pore system (i.e., salts fill or plug pores), the porosity reduction being more significant in the case of the cationic surfactant. No porosity changes were detected in the blank following salt crystallization, meaning that salts concentrate in the scales and fell off with the stone debris, but they did not fill the pores of the remaining stone block. The differences in solution evaporation behavior may be explained by the combined action of two phenomena: (a) salts precipitating at low supersaturation ratios in the (65) Correns, C. W. Disc. Faraday Soc. 1949, 5, 267. (66) Winkler, E. M.; Singer, P. C. Geo. Soc. Am. Bull. 1972, 83, 3509. (67) Rodriguez-Navarro, C.; Doehne, E.; Sebastian, E. Geo. Soc. Am. Bull. 1999, 111, 1250.

Figure 8. Mercury intrusion porosimetry pore-size distribution plots of limestone slabs following tests of salt crystallization in the presence of surfactants.

presence of cationic surfactant may block the stone pores early in the process, thus reducing solution flow-rate. This is consistent with porosity data showing pore filling by the salts; (b) the low CDBAC As values (Table 3) and the reduced adsorption of the cationic surfactant if compared to the anionic, increases the surfactant concentration in the saline solution-air interface, thus resulting in slower evaporation. This is consistent with the evaporation rate experiments (Table 4) showing that the sodium sulfate solution with cationic surfactant has the lowest evaporation rate. The CDBAC-saturated saline solution evaporation rate in the macroscale experiment (Figure 7) is slightly smaller than the blank or the SDS-saturated saline solution in the first week. However, later on, the evaporation rate was steadily reduced, when salt was crystallizing in the stone pores. Therefore the crystallization of the salts plugging the pores seems to be the main phenomenon responsible for the observed solution transport and evaporation behavior. Contact angle, surface tension, and viscosity of the saline solutions in the presence of surfactants (Table 4) will not account for the observed differences in capillary suction and the resulting evaporation at the stone surface. These parameters have very similar values in both the cationic and the anionic saline solution systems. The crystallization of mirabilite in a homogeneous fashion along the pores of the stone, while reducing stone damage at first, may result in greater damage later. Immersion in water resulted in complete disintegration of the stone block with cationic surfactant (61% weight loss). Less damage was observed in the case of the block with anionic surfactant (52% weight loss), while no additional damage occurred in the blank sample. Mirabilite can undergo rapid dehydration, transforming into thenardite (Na2SO4) at room temperature.38 In contact with water, hydration of thenardite to form mirabilite results in significant hydration pressure generation.68 If thenardite salts are evenly distributed throughout the stone pores, complete stone disintegration may take place following water exposure. This seems to occur in the case of the stone with surfactants; particularly in the CDBAC (68) Winkler, E. M.; Wilhelm, E. J. Geo. Soc. Am. Bull. 1970, 81, 567.

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Langmuir, Vol. 16, No. 3, 2000

case. This damaging process may be enhanced by the reported role of cationic surfactants in weakening stone by means of enhancing crack propagation and reducing mechanical strength due to adsorption at developing crack tips, resulting in interfacial energy reduction between contacting minerals (cement).69,70 4. Conclusions - and CDBA+ adsorption behavior and

Differences in DS micelle shape evolution promote changes in sodium sulfate crystallization as well as in the resulting damage associated with the crystallization of this salt in a porous calcareous stone. SDS induces crystallization and growth of nonequilibrium mirabilite crystals formed in high supersaturation conditions, resulting in high crystallization pressures, due to: (i) preferential adsorption of DS- onto calcite, together with Ca(DS)2 precipitation, that result in a reduction of free DS- in the solution as monomer and, particularly, as aggregates (micelles). A secondary role of SDS as crystallization inhibitor is not discarded. However, this latter role is not sufficient to avoid crystallization of the salt as damaging subflorescence, instead of harmless efflorescence; (ii) the greater evaporation rate of sodium sulfate solution with SDS, if compared with the blank sample and the solution with CDBAC, results in faster transport toward the evaporation front located a few millimeters below the limestone block surface, followed by crystallization of mirabilite as subflorescence at high supersaturation ratios. CDBAC promotes crystallization of mirabilite in the stone pores at low supersaturation, resulting in salt crystals with equilibrium morphologies, lower crystallization pressures (if compared with the blank sample and the solution with SDS), and lesser initial damage to (69) Dunning, J. D.; Lewis, W. L.; Dunn, D. E. J. Geophysical Res. 1980, 85, 5344. (70) Dunning, J. D.; Huf, W. L. J. Geophysical Res. 1983, 88, 6491.

Rodriguez-Navarro et al.

the stone blocks. Reduced adsorption of CDBA+ onto calcite in the saline solution enhances micellation. Micelles may promote crystallization and growth of mirabilite at low supersaturation ratios by solubilizing the solutes and promoting their transport toward mirabilite nuclei growth sites. This effect can be promoted by a transition from sphere to rodlike or disk shaped micelles, with higher solubilization capacity. However, the even distribution of mirabilite crystals throughout the stone pore system, may, eventually, result in significant damage if dehydrated mirabilite rehydrates. This damaging phenomenon may be enhanced by the mechanical weakening effect of cationic surfactants. In conclusion, care must be taken when using ionic surfactants in treatments for the conservation of porous calcareous materials. This study demonstrates that an impregnation with the ionic surfactants tested (SDS and CDBAC) is not an effective method for stone desalination when salts with hydrates are present. In fact, the side effects of their application can be serious (e.g., cationic surfactant). However, additional research is necessary to evaluate the effects of other surfactants or surfactant mixtures on similar decay processes. The testing of salt crystallization inhibitors, such as phosphates, could be a potential path to follow. The implication for understanding of decay induced by surface-active biological byproducts (i.e., biosurfactants) is also a potentially fruitful new area of research. Acknowledgment. This study has been financially supported by the Getty Conservation Institute (Project name: Preservation of Porous Calcareous Stone), by the Research Group RNM-0179 of the Junta de Andalucı´a and by the Research Project PB96-1445 (DGICYT). We are grateful to K. Elert and C. Selwitz for lively discussions and excellent suggestions, and to W.S. Ginell and A. Tagle for their help and continuous support during the development of this research. LA990580H