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Assessment of the protective effect of carbonation on portlandite crystals
Cement and Concrete Research 74 (2015) 68–77
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Assessment of the protective effect of carbonation on portlandite crystals I. Galan a,⁎, F.P. Glasser a, D. Baza b, C. Andrade b a b
Department of Chemistry, University of Aberdeen, Meston Walk, AB24 3UE Aberdeen, United Kingdom Eduardo Torroja Institute for Construction Science, CSIC, Serrano Galvache 4, 28033 Madrid, Spain
a r t i c l e
i n f o
Article history: Received 13 January 2015 Accepted 1 April 2015 Available online xxxx Keywords: B surface layer C carbonation C durability C permeability D Ca(OH)2
a b s t r a c t The kinetics of many reactions important to cement hydration and use are not well understood: this is in part due to the great complexity of many supposedly “simple” processes. One such process, carbonation of portlandite, Ca(OH)2, in moist air at ~23 °C has been investigated by microscopy and microchemical analysis. Single crystals of portlandite were grown, carbonated at relative humidities between ~ 25 and ~ 90%, and the transport properties of the self-generated calcite, CaCO3, product film were determined. The calcite films thus grown within days or weeks varied in thickness but typically were polycrystalline and epitaxial: a variety of morphologies and surface features are recorded. Permeation was measured by determining the time taken for Ca2+ ions, arising from the Ca(OH)2 substrate, to diffuse through the calcite coat into initially pure water. The spontaneous formation of self-protecting films on concrete has long been envisaged: results demonstrate that passivation can actually be achieved. © 2015 Elsevier Ltd. All rights reserved.
1. Introduction Portland cement is unstable in the thermodynamic sense in almost all service environments. Its observed deterioration often results from a combination of mechanisms which are not readily deconvoluted. One such mechanism is carbonation which results in loss of high pH and depassivation of embedded steel. Dry cement is resistant to carbonation but moisture facilitates reaction because CO2 is soluble and, in the alkaline cement environment, is converted to carbonate and bicarbonate: these latter species react readily with moist cement. The nature of the product depends on the chemical environment and temperature: while calcite is probably the stable phase, very low temperatures are also associated with formation of ikaite, CaCO3·6H2O, while the presence of Mg ions is associated with formation of monohydrocalcite, CaCO3·H2O. Mainly, however, the product is calcium carbonate which has three commonly-occurring crystalline polymorphs: vaterite, calcite and aragonite. The influence of relative humidity (RH) on the formation on CaO substrates of the different crystalline polymorphs and amorphous CaCO3 has been recently reported [1]: all three crystalline polymorphs would form at 40% RH but only calcite could be detected at RH between 60 and 80%. However, the study had as it objective hydration of CaO and the carbonation was an unwanted and uncontrolled side reaction, so its relevance is uncertain. In the title study we examine the mechanism whereby portlandite, Ca(OH)2, is converted to calcite in moist environments and how this newly grown CaCO3 layer can protect the remaining crystal. Portlandite is an abundant constituent of hydrated Portland cement, perhaps ⁎ Corresponding author. Tel.: + 44 1224274733. E-mail address: [email protected] (I. Galan).
http://dx.doi.org/10.1016/j.cemconres.2015.04.001 0008-8846/© 2015 Elsevier Ltd. All rights reserved.
comprising 20–25% of its mass, and it is known to carbonate readily. An experimental plan was prepared in which the cement was represented by isolated portlandite crystals, which were carbonated at controlled temperature and different RHs. The partially carbonated crystals were immersed in an aqueous solution in order to monitor the diffusion of Ca2+ ions through the CaCO3 cover. This process occurs by transport across two physically separated interfaces: one between air and “water” and the other between “water” and a solid, in the title paper, a single crystal of Ca(OH)2 (Fig. 1). Dissolution of carbon dioxide from the atmosphere into the pore solution has been reviewed by Cowie et al. [2]. The sorption of CO2 from a gaseous atmosphere is greatly facilitated by increasing the aqueous pH [3] and portlandite is sufficiently soluble to achieve a major rise in pH. This dissolution process is believed to be partly controlled by dislocations in solid portlandite [4]. Many other factors potentially influence transport of ions across and between the two interfaces [5–7] but in the title study, every effort has been made to reduce transport time between the two interfaces by decreasing the thickness of the aqueous layer. The kinetics of carbonation of Ca(OH)2 have been previously studied using powder [8] and single crystals [9,10]. According to [9], both calcite and vaterite were grown on Ca(OH)2 crystals exposed to humid CO2; surface nucleation of CaCO3 on portlandite crystals occurred and subsequent random growth of calcite was observed. The observations confirm that Ca(OH)2 is not topochemically converted to CaCO3, but that a dissolution–reprecipitation process occurs resulting in nucleation at many sites with eventual coalescence of carbonate crystallites. This mechanism is consistent with the observations at ~ 20 °C of a rate maximum at RH in the range of 50–80%. The dissolution–precipitation hypothesis was also confirmed by studying crystals carbonated in air
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Ca 2+ + 2OH-
AIR
VAPOUR
WATER
CALCITE
PORTLANDITE
CO2
Fig. 1. Schematic showing diffusion through calcite layer from portlandite crystal and from wet environment.
at 93% RH and in carbonated solutions [10]. According to these latter authors ‘islands’ of CaCO3 grow epitaxially on the surface of portlandite crystals in specific preferred orientations. They also claim that “no porosity is generated during the reaction, which progresses through the formation of fractures”, which “form in response to the stresses generated in the replacement rim due to the positive volume change during the transformation of portlandite into calcite” [10]. 2. Experimental First the dissolution rate of Ca(OH)2 crystals into initially pure water was measured using crystals with one exposed face. These data were used to benchmark data obtained subsequently using portlandite crystals from the same batch which had been carbonated under controlled conditions, and again subjected to dissolution. The aqueous solution was periodically analysed for Ca2+ using an ion-selective electrode. In a fixed geometry, the difference in aqueous solution compositions normalised to release per unit surface area, could then be attributed to the protective ability of the CaCO3 film. The Ca-sensitive electrode could be located at ~2 mm from the surface being measured, thereby increasing the sensitivity of measurement. The method works for two reasons: first, the solubilities of portlandite and calcite differ greatly, with portlandite being much the more soluble by a factor of ~ 100 to 1000, and second, because the time taken for Ca2 + ions to diffuse through the aqueous film and reach the sensor, the “dead time” of the experiment, as estimated from literature values for the diffusion coefficient of free calcium ions in water, is negligible. Thus the time taken for Ca2 + to diffuse to the sensor is negligible relative to the time required, often extending over weeks, for the ion constituents of Ca(OH)2 to diffuse through the CaCO3 film.
Fig. 2. SEM of a portlandite crystal showing pseudo-hexagonal prismatic morphology (bar = 1 mm).
related defects at points of contact. The purity and mineralogical nature of the crystals were confirmed by DTA/TG as well as by powder XRD. Single crystals were carbonated by exposure to CO2-containing atmosphere at controlled humidity, generally in the range of 50–90% RH, at ambient temperature. A set of samples was exposed to 100% CO2 atmospheres at 53, 75 and 90% RH. A second set of samples was exposed to air in a lab atmosphere, ~23 °C. With time, layer thicknesses of a few microns to tens of microns could be formed and the thickness was measured by SEM from cross-sections of the crystals. Fig. 3 shows the external appearance of a crystal following carbonation at 53% RH and 100% CO2 for 5 weeks. In general the morphology of the portlandite crystal, including fine features such as layer steps, is faithfully preserved. To measure carbonation kinetics and establish the protective nature of the CaCO3 layer both carbonated and uncarbonated single crystals were coated on all but one face with an impermeable epoxy coat and the unprotected surface area was measured by contrast image analysis. The uncarbonated crystals were used to benchmark the subsequent measurements. The partially coated crystals were placed in 5 ml of initially pure de-aerated water, as shown in Fig. 4, with a calcium selective electrode (with a Ag/AgCl reference electrode). In uncoated samples the portlandite crystal dissolves in the aqueous phase and the corresponding rise in calcium concentration is monitored as a function of time. The distance between the calcium electrode and the surface of the crystal was kept small, a few mm, to reduce the “dead time”. The ratio
2.1. Procedure Single crystals of portlandite were grown by mixing aqueous solutions of NaOH and CaCl2 at ~23 °C, according to the procedure described in [9]. Atmospheric carbon dioxide was excluded from the apparatus. Growth took ~ 5 weeks. The resulting product contained transparent to translucent distorted prismatic crystals or clumps of a few individual crystals ranging up to up to ~ 5–7 mm: an example is shown in Fig. 2. Individual crystals could thus be hand-picked by eye. The crystals were handled with soft-tipped tweezers to avoid inducing pressure-
Fig. 3. SEM of carbonated portlandite crystal (bar = 1 mm).
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Electrode CaCO 3
Crystal 5 ml
Fig. 6. SEM of lightly carbonated surface showing nucleation of CaCO3, with angular morphology. Fig. 4. Picture and sketch of experimental setup for measuring calcium dissolved from carbonated and uncarbonated crystals. The spacing between the electrode tip and the surface of the crystal was ~2 mm.
commercial standards and it proved stable throughout the course of experiments. 3. Results
mass of solution/mass of crystal was in the order of ~200, ensuring that the total mass of the crystal was more than sufficient to saturate the 5 ml solution. The experiment was repeated using surface-carbonated crystals and the difference in time for a rise in calcium taken as a measure of the retardation effect arising from the presence of the calcium carbonate layer. The very low solubility of CaCO3 in water (~ 3 ppm Ca2 + at 20 °C), compared to that of Ca(OH)2 (865 ppm Ca2 + at 20 °C) [11] makes this method feasible: CaCO3 does not contribute significantly to raising Ca2+ concentration: overwhelmingly, only Ca(OH)2 diffusing through the layer of CaCO3 will significantly increase calcium concentrations in solution [12]. Thus the rates at which the calcium concentration changes as a function of time is the key to kinetic analysis. Similar experiments were performed in solutions partially saturated with respect to Ca(OH)2, with initial Ca2+ concentrations ranging from 50 to 500 ppm. The 50–500 ppm concentrations were achieved by adding a weighed amount of CaO to a known volume of solution. Three replicates were used for each condition. The measurements were performed in CO2-free conditions in a glove box filled with N2: temperature was kept constant at 22–24 °C and the RH varied between 35 and 45%. Calibration of the electrode was checked daily against
Fig. 5. SEM of crystal showing dislocations, imperfections and growth steps (bar = 200 μm).
3.1. Preparation of portlandite crystals Figs. 2 and 5 show the appearance of a newly-grown portlandite crystal. The crystals have a pseudo-hexagonal prismatic morphology and do not grow as platelets, as might be expected. The longer the crystals were left ‘growing’, the more perfect the hexagonal prismatic morphology became, although some growth steps always remained. In the SEM image of Fig. 5, the surface roughness of crystals is apparent: the surface topology is dominated by growth steps, presumably resulting from the pile-up of many atomic step height layers, as well as accretion of small “seed” crystallites. The true surface area available for subsequent reaction is thus somewhat greater than that which was estimated by microscopy. 3.2. Carbonation of portlandite crystals 3.2.1. In 100% CO2 atmospheres The initial carbonation is not uniform but occurs by nucleation of CaCO3 at active sites on the portlandite surface. Figs. 6 and 7 show the first signs of nucleation in crystals exposed to CO2 atmospheres for 5 weeks at 53% RH. Multiple rhomb-shaped and dendritic crystals of CaCO3 decorate the surface, apparently at random, but probably originating at dislocations.
Fig. 7. High magnification of early stage carbonation showing dendritic growth of CaCO3. As in Fig. 6, the crystal was exposed to 100% CO2 at 53% RH for 5 weeks.
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1
2
Fig. 10. Detail of an air-carbonated surface showing different morphologies, ranging from rosette shapes (area 1) to prisms and doubly-terminated rhomboids (area 2). Fig. 8. Highly carbonated area covered with rough textured polycrystalline CaCO3 coating. Crystal was exposed to CO2 at 53% RH for 5 weeks. Note that the visual appearance of the surface is porous.
Carbonation progresses by a combination of growth from apparently randomly-nucleated sites as well as by continued nucleation at fresh sites. Eventually the surface becomes covered as the nucleated sites grow and coalesce. The carbonation mechanism shows that (i) nucleation of CaCO3 occurs at numerous sites so the resulting film is polycrystalline and (ii) from the micrographs, the film appears to become continuous or nearly so (see Fig. 8). This picture reveals no obvious evidence of preferred crystallographic orientation, in agreement with [9]. This does not mean that they are not oriented, but that the substrate is not visible so the relations cannot be determined by microscopy. The mass of CaCO3 grown on crystals carbonated at 53% RH for 5 weeks was ~ 5% of the crystal mass, as measured by TG, so layer thicknesses were on the order of tens of microns.
3.2.2. In laboratory atmosphere Figs. 9 and 10 show images of the faces of a crystal carbonated in air at room temperature and conditions for ~6 months. The layer is rough and the coalescence of crystallites makes it impossible to delineate morphological features within the layer but isolated growth also occurs at the surface as shown in Fig. 10. These surface features are believed to be forming on a substrate layer of CaCO3. Fig. 11 is a transversal cut section of a crystal following carbonation in air for ~ 6 months. The upper part of the image shows the CaCO3 ‘layer’ grown on one of the rectangular faces, the depth of which has
reached ~ 100 microns. The surface projections are conspicuous and the film is thus not homogeneously distributed across the surface. Figs. 12–14 show images of a “hexagonal face”, pseudo {0001}, of a crystal partially carbonated in laboratory air for ~6 months at progressively higher magnifications. Again different morphologies can be appreciated: ‘pyramids’, small needles and dendrites occur. Fig. 14 apparently shows ‘holes’ or ‘pores’ on the surface, which develop during the progressive formation of the carbonated layer. 3.3. Dissolution of surface carbonated and uncarbonated crystals Fig. 15 shows dissolution rates of uncarbonated and carbonated portlandite crystals when contacted with initially deionised water. Fresh, uncarbonated surfaces give a virtually instantaneous release of Ca2+ ions to solution. On the other hand, previously carbonated crystals, for example when carbonated at 75% RH, have a definite induction period before dissolution of OH− and Ca2+ ions commences. Under the conditions used, 23 °C, the solubility product constant Ksp of portlandite (5.02 × 10−6) is approximately 1000 times greater than that of calcite (3.36 × 10− 9). This means that the charge balancing ion for OH− is Ca2 + and vice versa; no other cations other than H3O+ are present and the aqueous concentration of H3O+ decreases to near zero as the pH rises. Therefore we assume, neglecting species such as Ca(OH)+, that each Ca2+ ion in solution is charge balanced by 2OH. Uncarbonated portlandite releases Ca2 + (and OH) ions almost immediately whereas release from surface-carbonated crystals is much delayed. This delay is associated with the time required for OH−
Surface
Fig. 9. SEM of a carbonated crystal exposed to air in a lab for 6 months at ~20 °C. The RH was not controlled but varied between ~25 and 70%.
Fig. 11. SEM image of transversal section showing CaCO3 layer grown on top of the crystal that was air-carbonated for 6 months indoors.
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Fig. 14. Partially carbonated hexagonal surface of crystal exposed to air for 6 months. Note the holes (black) apparently giving rise to a reticular network of CaCO3 crystals.
and Ca2+ ions to diffuse through the CaCO3 film and it is apparent that the film can be strongly protective against dissolution. Regarding the saturation degree of the solution, uncarbonated crystals eventually reach a saturation value of ~900 ppm Ca2+ after about 500 h, which accords with the known equilibrium calcium ion concentration in that solution [11]. However when CaCO3 is also present, the shape of the solubility curve in the Ca(OH)2–CaCO3–H2O system governs the equilibrium Ca2 + concentration. This curve has a degree of freedom, so its numerical value depends on the aqueous molal ratios of the two solid components. In general, the saturation concentration of Ca2+ is less than 900 ppm and actual values, as recorded in Fig. 15, lie in the range of 300–700 ppm. As expected, the dissolution rate (i) diminishes significantly with time as saturation is approached and (ii) is affected by the continuity and thickness of the CaCO3 coating which hinders the dissolution. In all surface carbonated crystals made at 75% RH, dissolution requires a significant “breakthrough” period, ranging between 50 and 100 h. The decrease both in the dissolution rate and in the saturation value is more pronounced for crystals carbonated at intermediate RH than for 90% RH (data for which are not shown) evidencing the impact of humidity on the quality of the coating layer that can be grown on single crystals. Films grown at high RH, 90%, provide little protection against dissolution of the portlandite substrate compared to films grown at
~75% RH, as evidenced by the development of a “breakthrough” period and its duration. As noted, the increase in calcium concentration, Fig. 15, is also related to the exposed surface area. To obtain quantitative measures, we normalise values: dissolved calcium was divided by the crystal exposed area. These data are shown in Figs. 16–19. The maximum values of calcium dissolved are much lower for the crystals with CaCO3 layers grown at intermediate RH than for either uncarbonated crystals or those carbonated at high (N90%) RH. As for the influence of the initial concentration in Ca2+ ions of the solution where the crystals were placed to dissolve, rates do not much vary between pure water and solutions conditioned with 50 ppm calcium as Ca(OH)2. Much higher initial concentration, 500 ppm calcium, has an important influence in the dissolution of the crystal: as expected from the law of mass action, the lower the initial concentration in the solution, the more rapidly portlandite will dissolve. This can also be appreciated by comparing the initial slope of the data, or initial dissolution rate, which is in all cases considerably lower for crystals dissolving in 500 ppm Ca solutions relative to initially pure water. Table 1 summarises average values for the induction period for the crystals in the 4 stages of carbonation and in the 3 solutions. The breakthrough period comprises the time during which the increase in the concentration of calcium in the solution is less than ~5 ppm/mm2. The uncarbonated crystals start dissolving almost immediately after being put in contact with the solution. Values for those carbonated at 90% RH do not differ very much from initially pure water, implying that
[Ca2+] / ppm
Fig. 12. Hexagonal phase of a crystal carbonated in air.
1200 1100 1000 900 800 700 600 500 400 300 200 100 0
Uncarbonated (benchmark)
Partially carbonated (75% RH)
0
100
200
300
400
500
600
Time / hours Fig. 13. Detail of the different morphologies of CaCO3 grown on the hexagonal face, {0001}, of a crystal exposed to air for ~ 6 months. The ‘pyramids’ have not been observed in n o prismatic faces 0110 of crystals carbonated under the same conditions in air (Figs. 9 and 10).
Fig. 15. Calcium dissolved from uncarbonated (blue markers) and partially carbonated at 75% RH (red markers) crystals in water. The different markers for each colour represent different replicates. (For interpretation of the references to colour in this figure legend, the reader is referred to the web version of this article.)
Calcium in solution per unit area (ppm/mm2)
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300
73
500 ppm 50 ppm 0 ppm
250 200 150 100 50 0 0
100
200
300
400
500
600
time (hours)
Calcium in solution per unit area (ppm/mm2)
Fig. 16. Evolution of dissolved calcium (per unit of crystal area) in solutions where uncarbonated crystals were placed to dissolve. Starting calcium concentrations are 0, 50 and 500 ppm. For each concentration 3 replicates were used. Each symbol (square, circle and triangle) represents data from 3 crystals.
300
500 ppm 50 ppm 0 ppm
250 200 150 100 50 0 0
100
200
300
400
500
600
time (hours)
Calcium in solution per unit area (ppm/mm2)
Fig. 17. Evolution of dissolved calcium (per unit of crystal area) in solutions where carbonated crystals (at 53% RH) were placed to dissolve. Starting calcium concentrations are 0, 50 and 500 ppm.
300
500 ppm 50 ppm 0 ppm
250 200 150 100 50 0 0
100
200
300
400
500
600
time (hours)
Calcium in solution per unit area (ppm/mm2)
Fig. 18. Evolution of dissolved calcium (per unit of crystal area) in solutions where carbonated crystals (at 75% RH) were placed to dissolve. Starting calcium concentrations are 0, 50 and 500 ppm.
300
500 ppm 50 ppm 0 ppm
250 200 150 100 50 0 0
100
200
300
400
500
600
time (hours) Fig. 19. Evolution of dissolved calcium (per unit of crystal area) in solutions where carbonated crystals (at 90% RH) were placed to dissolve. Starting calcium concentrations are 0, 50 and 500 ppm.
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Table 1 Average breakthrough time (h) for starting dissolution of uncarbonated and carbonated crystals in 0, 50 and 500 ppm Ca solutions. Initial solution composition
Uncarbonated
Carbonated at 53% RH
Carbonated at 75% RH
Carbonated at 90% RH
0 ppm 50 ppm 500 ppm
3 3 4
50 18 98
72 56 216
5 4 34
the coating is not effective, except in the case of the 500 ppm solution where it delays the start of dissolution for 30 h. All samples carbonated at 75% RH show an induction period in excess of 50 h. This time quantitatively assesses the good protection afforded by the coating which is on the order of 10–100 microns thick (Fig. 11). The layer formed at 53% RH provides some protection but presents breakthrough times lower than the coating at 75% RH. An approximate value of the effective diffusion constant of the protective layer can be calculated (see Table 2). The difference between the numerical value of this coefficient and the diffusion coefficient of calcium at infinite dilution, 7.92 × 10−6 cm2/s [11], provides a measure of the protection afforded by the carbonate layer: see Appendix A for more details. The persistence and durability of the carbonate coating have been measured by prolonged water immersion. Figs. 20–23 show images of crystals previously carbonated at 53% RH for 5 weeks and then exposed to initially pure water (Figs. 20 and 21) and to 50 ppm Ca2+ solutions (Figs. 22 and 23) for ~30 days at ~23 °C. The white cover of CaCO3 remains largely intact while portlandite has dissolved from underneath the calcite coat (Figs. 20 and 22). Fig. 21 shows at high magnification a SEM image of the area marked A in Fig. 20. A generally dense layer of CaCO3 has become perforated allowing dissolution to concentrate at particular sites. Fig. 23 shows a detail of the area marked B in Fig. 22. Small, 2–3 micron, crystals randomly distributed and decorating a denser and more uniform substrate can be appreciated. This texture is believed to arise from partial recrystallization within the calcite protective layer in the course of the dissolution experiments. The following three SEM images show the evolution of the dissolution of a crystal partially carbonated in air. In this case, the solutions were stirred and the crystals were examined by SEM after 5, 30 and 60 min. Before the dissolution the crystals looked like the example depicted in Fig. 9. After 5 min (Fig. 24), triangular etch pitch appear and dissolution in steps starts to be visible. After 30 min (Fig. 25) pits with a defined morphology appear. The size of these pits is between 10 and 20 microns. After 60 min (Fig. 26) the pits have grown to ~ 500 microns. The layers inside the pits (up to 10) can be clearly seen. A noteworthy feature of the etch pits is their consistent orientation over distances of 1–2 mm. This orientation suggests that the calcite layer may indeed have an epitactic relation to its single crystal substrate.
A
1 mm
Fig. 20. Picture obtained by optical microscopy of a ca 1 × 3 mm carbonated crystal after exposure to dissolution in water for ~30 days. The CaCO3 cover (white) seems to remain largely intact: note especially the nearly complete margin of CaCO3 delineating the margin of the grain. The crystal had a small outgrowth of CaCO3 at point A: a higher magnification view of this area is shown in Fig. 21.
processes occurring during cement hydration, for example to control the length of the so-called dormant period of hydration. But proof of a barrier effect has often been indirect and the constitution and effectiveness of “barrier layers” remain conjectural. In the present study, it is shown that barrier layers can be formed in the course of use by spontaneous processes such as carbonation. The properties of CaCO3 formed by reaction of portlandite with water and CO2 gas have been investigated and proof obtained that in some circumstances, a layer of calcite developing on portlandite can significantly protect portlandite against subsequent carbonation and dissolution. 4.1. Influence of relative humidity Carbonation reaction proceeds very slowly in dry air but develops rapidly in the presence of water because of the need to dissolve CO2 in a water film. Once initiated, reaction liberates water which, if not evaporated, facilitates reaction. Reaction thus proceeds across two interfaces, one between portlandite and aqueous solution, the other between aqueous solution and air (see Fig. 27). At low RH, when a thin water film wets portlandite, the two interfaces may be very close, around 3 to 18 nm (10 to 60 water molecules
4. Discussion The existence of barrier layers in cement reactions has often been proposed to explain the observed kinetics of many features and
Table 2 Apparent diffusion coefficient through the layers between the surface of the crystals and the electrode (carbonate layer plus water). D (10−6 cm2/s) Uncarbonated Carbonated at 53% RH Carbonated at 75% RH Carbonated at 90% RH
1.17 0.038 0.031 0.086
Fig. 21. High magnification SEM image of CaCO3 cover: area marked as A in Fig. 20. Note the appearance of what looks to be holes.
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B
Fig. 24. SEM of an air-carbonated crystal after 5 min dissolving in water. Fig. 22. SEM image of a carbonated crystal after dissolution in 50 ppm Ca solution for ~30 days. Note that rupture of the protective CaCO3 film (B) which follows dissolution of the more soluble core of portlandite. The collapse of the calcite shell may result in part from the coating process, high vacuum and beam damage in the microscope.
thick). At the portlandite–aqueous interface, portlandite dissolves, increasing the pH of the aqueous phase, while at the air interface, atmospheric CO2 dissolves, initially as molecular CO2 but subsequently forming “carbonic acid”; the latter partially dissociates to mixtures of bicarbonate and carbonate species. Although the initial dissolution reaction is slow, the bicarbonate–carbonate equilibrium is rapidly reversible and the equilibrium speciation is pH dependent with carbonate increasingly favoured with respect to bicarbonate species as pH increases. However, the reactions at the two interfaces are only indirectly coupled: one proceeds independently of the other. The indirect link does however exist (i) because the rate of dissolution of CO2 at the air interface is much enhanced by rising pH and the source of OH− ions, essential to increase pH, is at the portlandite–solution interface and (ii), the transport of species is limited by the thickness of the liquid and its diffusive characteristics. In the experimental work reported, every effort has been made to keep the film as thin as practicable and provide direct diffusion paths. The point at which carbon containing species react with Ca(OH)2 forming CaCO3 depends on the above factors as well as temperature. But it should be recalled that CaCO3 formed is several orders of magnitude less soluble than the substrate portlandite and, as a result of this
as well as the physical conditions imposed on diffusion and mass transport by the existence of two physically separated interfaces, a layer of relatively insoluble CaCO3 often tends to develop on portlandite surfaces. The reaction is sensitive to RH. At low RH it is difficult to establish a water film on the surface and reaction is slow because the first step, dissolution of CO2 into aqueous phase, is hindered by the low and likely discontinuous surface coverage. Additionally, although water is released by the carbonation reaction, it is prone to evaporate at low humidity. On the other hand, at very high RH (N90%), the higher number of water monolayers give a continuous but thicker water coverage and the calcium carbonate may nucleate in the aqueous layer as well as at the crystal/solution interface, allowing a more complex morphology to develop in the carbonate layer; this influences product morphology and orientation according to mode of formation. Diffusion paths also lengthen as the water film thickens. A further factor is that the reaction generates more water: water balances depend on competition between the rate of generation and its evaporation or retention. At intermediate humidity, 53%, and more clearly in present results, 75%, the conditions of thickness of the water layer on the crystals and distance between the two interfaces (air/solution and solution/crystal) seem to be the optimum for development of low permeable carbonation layers which also adhere strongly to the substrate crystal. The influence of the different polymorphs formed at different RH [1] cannot be directly related to the protection offered by the film.
Fig. 23. Detail of SEM picture of CaCO3 cover of partially dissolved crystals (area marked as B in Fig. 22).
Fig. 25. SEM of an air-carbonated crystal after 30 min dissolving in water.
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ages. Crystal morphologies become more regular and suggestive of near-equilibrium conditions of growth. Thus the film properties are best determined by actual measurement of their diffusive properties. It is also noteworthy that the conditions of formation and measurement are such that they could readily be achieved in practice. Formation occurred in limited volume of water, but this ideal thickness is spontaneously developed by exposure to air and CO2. Optimum conditions for growth are close to self-stabilising as water lost by evaporation is replaced by water liberated at the interface with substrate portlandite.
4.3. Film properties by electron microscopy
Fig. 26. SEM of an air-carbonated crystal after 60 min dissolving in water.
4.2. Influence of the neoformed carbonation coverage on further carbonation Results show that the quality of the film is very dependent on the conditions prevailing in the course of its initial growth. In the most favourable conditions the film had sufficiently low permeability to diffusion of Ca(OH)2 as to require 10 to 100 h for breakthrough of ions from substrate portlandite to affect the aqueous phase. And, once breakthrough is achieved, the rate of diffusion through the protective calcite film is slowed relative to unprotected crystals. Experience showed that it was difficult to judge the performance of coatings by microscopy alone: at higher magnifications necessary to detect pores or imperfections, it was only possible to examine a small area. The granularity of the film, as observed by electron microscopy, was not necessarily an indication of its permeation properties. Nor was the early-formed partial coverage an indication of the quality of the fully developed cover. Thus the early stages of surface cover gave a wide range of morphologies: initially, crystals were sparse and had acicular, rosette, dendritic or other morphologies typical of growth from supersaturated solutions. However many of the more successful films were grown slowly over weeks or months and had more regular prismatic morphologies which appeared to consist of an interlocking network of prismatic grains. It is believed that films, once formed, may not be morphologically stable but in favourable cases may recrystallise with ageing and growth, thereby gaining in strength and coherence and with infilling of pores. Thus while the early stages of growth of the calcite layer are driven by morphological features and defects on the substrate, the coalescence of growth and the internal recrystallization and accretion of more material govern the film properties at longer
AIR : CO2
(Ca+2 OH-) (CO3= H+) Pore solution
Although electron microscopy has been immensely valuable to characterise the surface state of portlandite crystals and visualise features such as the microstructure and thickness of calcite films, it has limitations. The principle limitation is the inability of electron microscopy to estimate the permeation properties of the protective layers. None of the layers look particularly impermeable even when actual measurements disclose extremely low film permeabilities, as evidenced from breakthrough times. In this respect only actual measurements of film properties can be treated as a reliable indicator of permeation properties.
4.4. Overall perspective We have shown that in certain regimes, calcium carbonate films can form spontaneously which are remarkably effective in limiting diffusion of calcium hydroxide from cement and in protecting further carbonation of substrate portlandite. The question immediately arises: to what extent can data derived from portlandite single crystals be used to develop carbonation resistant concrete? Can spontaneous carbonation processes be used as part of a protective system for durable concrete? We have not done the determination of whether the protective film can be established on commercial portlandite and Portland cements, the latter containing only 20–25% polycrystalline portlandite. But we predict that the self-protection mode is most likely to be demonstrated by the behaviour of cements based on portlandite where solubility is most effectively suppressed by formation of optimised calcite coatings. Portland cements will however require a great deal of further investigation: the paste is only 20–25% portlandite although carbonation of other phases, e.g., CSH, may contribute to establishment of protective surface layers. Many unknowns remain but it is at least possible to envisage a new generation of cementitious materials which will have selfgenerated and self-protective surface layers which ensure longer service lifes in conditions where leaching, carbonation and resistance to penetration by adverse chemicals in the service environment occur.
AIR : CO2
(Ca+2 OH-) (CO3= H+) Pore solution CaCO3
Ca(OH)2
Ca(OH)2
AIR : CO2
Pore solution CaCO3
Ca(OH)2
Fig. 27. Schematic showing the initial establishment and evolution (from left to right in the figure) of a relatively impermeable layer of calcite on a portlandite substrate in a moist atmosphere. The water film is coloured blue. The thickness of the stable water film relates to the relative humidity with RH values in the range of 53–75% giving the best coverage. (For interpretation of the references to colour in this figure legend, the reader is referred to the web version of this article.)
I. Galan et al. / Cement and Concrete Research 74 (2015) 68–77
5. Conclusions The results obtained have enabled to draw up the following conclusions: 1. Portlandite crystals develop a layer of calcium carbonate whose properties of continuity and thickness depend on the external RH. 2. The protective ability of the carbonated film has been for the first time assessed by: a. Measurement of the calcium leached from uncarbonated and carbonated crystals. This protective ability afforded by optimum carbonation has been measured in different conditions and optimum conditions established. b. Calculation of an apparent diffusion coefficient of the calcium through the layers between the surface of the crystals and the electrode (carbonate layer plus water). The so calculated coefficients are 1-2 orders of magnitude lower for carbonated than for uncarbonated crystals. 3. Despite the polycrystalline nature of the CaCO3 layer it provides significant protection against diffusion of Ca2+ and OH− ions. 4. The foundations are laid whereby spontaneous processes can be used to develop the self-protecting ability of lime and possibly Portland cement against a range of degradative processes. Appendix A
h
A one-dimensional solution of Fick's second law, C ðx; t Þ ¼ C s i 1−erf 2pxffiffiffiffiffi was applied where x is the distance between the surface Dt
of the crystal and the position of the electrode, t is the time taken to reach the concentration C at the electrode, D is the apparent diffusion coefficient through the layer of carbonate plus the water and Cs is the concentration at the surface of the crystal or the saturation value. This equation was chosen assuming that the dissolution of Ca(OH)2 is much faster than the diffusion of the ions. Table 2 includes D values for pure water. For these calculations, the saturation value was taken as 900 ppm, a time of 170 h was chosen (as the time required to reach steady state) and the distance between surface of the crystal and electrode was 2 mm.
77
The coefficient for the uncarbonated case is not equal to the coefficient in water, but is around seven times smaller and additionally, the diffusion coefficients for the carbonated crystals are significantly lower, 1–2 orders of magnitude, the lowest value being obtained for the crystal carbonated at 75% RH. Two explanations can be given for these differences: - The own kinetics of dissolution which can be not fully instantaneous, being slower in the case of the presence of the carbonated layer. - Being the dissolution instantaneous, there is a time to build up the saturation degree at just the crystal solution interface.
Even though a perfect calculation cannot be achieved with present results, and the values in Table 2 also include the diffusion through the 2 mm of water, it can be clearly appreciated that the coating slows down the process considerably: for the layer grown at 75% RH, D is 250 times less than the corresponding value for initially pure water. References [1] E. Dubina, L. Korat, L. Black, J. Strupi-Šuput, J. Plank, Influence of water vapour and carbon dioxide on free lime during storage at 80 C, studied by Raman spectroscopy, Spectrochim. Acta A Mol. Biomol. Spectrosc. 111 (2013) 299–303. [2] J. Cowie, F.P. Glasser, The reaction between cement and natural waters containing dissolved carbon dioxide, Adv. Cem. Res. 4 (1991) 119–134. [3] W. Stumm, J.J. Morgan, Aquatic Chemistry, 3rd ed. Wiley, NY. USA, 1996. [4] I. Galan, F. Glasser, C. Andrade, D. Baza, Dissolution of portlandite, 13th International Congress on the Chemistry of Cement, 2011, pp. 1–7. [5] D.E. Giles, I.M. Ritchie, B.- Xu, The kinetics of dissolution of slaked lime, Hydrometallurgy 32 (1993) 119–128. [6] J. Wang, T.C. Keener, G. Li, S.- Khang, The dissolution rate of Ca(OH)2 in aqueous solutions, Chem. Eng. Commun. 169 (1998) 167–184. [7] K. Johannsen, S. Rademacher, Modelling the kinetics of calcium hydroxide dissolution in water, Acta Hydrochim. Hydrobiol. 27 (1999) 72–78. [8] S.- Shih, C.- Ho, Y.- Song, J.- Lin, Kinetics of the reaction of Ca(OH)2 with CO2 at low temperature, Ind. Eng. Chem. Res. 38 (1999) 1316–1322. [9] J.R. Johnstone, F.P. Glasser, Carbonation of Single Crystals of Portlandite in Cement Paste, 9th International Congress on the Chemistry of Cement (ICCC) 1992, pp. 370–376. [10] E. Ruiz-Agudo, K. Kudlacz, C.V. Putnis, A. Putnis, C. Rodriguez-Navarro, Dissolution and carbonation of portlandite [Ca(OH)2] single crystals, Environ. Sci. Technol. 47 (2013) 11342–11349. [11] D.R. Lide, CRC Handbook of Chemistry and Physics, 2005. [12] I. Galan, F.P. Glasser, D. Baza, C. Andrade, Calcium carbonate coating on portlandite crystals, 33rd Cement and Concrete Science Conference, 2013.
Contents lists available at ScienceDirect
Cement and Concrete Research journal homepage: http://ees.elsevier.com/CEMCON/default.asp
Assessment of the protective effect of carbonation on portlandite crystals I. Galan a,⁎, F.P. Glasser a, D. Baza b, C. Andrade b a b
Department of Chemistry, University of Aberdeen, Meston Walk, AB24 3UE Aberdeen, United Kingdom Eduardo Torroja Institute for Construction Science, CSIC, Serrano Galvache 4, 28033 Madrid, Spain
a r t i c l e
i n f o
Article history: Received 13 January 2015 Accepted 1 April 2015 Available online xxxx Keywords: B surface layer C carbonation C durability C permeability D Ca(OH)2
a b s t r a c t The kinetics of many reactions important to cement hydration and use are not well understood: this is in part due to the great complexity of many supposedly “simple” processes. One such process, carbonation of portlandite, Ca(OH)2, in moist air at ~23 °C has been investigated by microscopy and microchemical analysis. Single crystals of portlandite were grown, carbonated at relative humidities between ~ 25 and ~ 90%, and the transport properties of the self-generated calcite, CaCO3, product film were determined. The calcite films thus grown within days or weeks varied in thickness but typically were polycrystalline and epitaxial: a variety of morphologies and surface features are recorded. Permeation was measured by determining the time taken for Ca2+ ions, arising from the Ca(OH)2 substrate, to diffuse through the calcite coat into initially pure water. The spontaneous formation of self-protecting films on concrete has long been envisaged: results demonstrate that passivation can actually be achieved. © 2015 Elsevier Ltd. All rights reserved.
1. Introduction Portland cement is unstable in the thermodynamic sense in almost all service environments. Its observed deterioration often results from a combination of mechanisms which are not readily deconvoluted. One such mechanism is carbonation which results in loss of high pH and depassivation of embedded steel. Dry cement is resistant to carbonation but moisture facilitates reaction because CO2 is soluble and, in the alkaline cement environment, is converted to carbonate and bicarbonate: these latter species react readily with moist cement. The nature of the product depends on the chemical environment and temperature: while calcite is probably the stable phase, very low temperatures are also associated with formation of ikaite, CaCO3·6H2O, while the presence of Mg ions is associated with formation of monohydrocalcite, CaCO3·H2O. Mainly, however, the product is calcium carbonate which has three commonly-occurring crystalline polymorphs: vaterite, calcite and aragonite. The influence of relative humidity (RH) on the formation on CaO substrates of the different crystalline polymorphs and amorphous CaCO3 has been recently reported [1]: all three crystalline polymorphs would form at 40% RH but only calcite could be detected at RH between 60 and 80%. However, the study had as it objective hydration of CaO and the carbonation was an unwanted and uncontrolled side reaction, so its relevance is uncertain. In the title study we examine the mechanism whereby portlandite, Ca(OH)2, is converted to calcite in moist environments and how this newly grown CaCO3 layer can protect the remaining crystal. Portlandite is an abundant constituent of hydrated Portland cement, perhaps ⁎ Corresponding author. Tel.: + 44 1224274733. E-mail address: [email protected] (I. Galan).
http://dx.doi.org/10.1016/j.cemconres.2015.04.001 0008-8846/© 2015 Elsevier Ltd. All rights reserved.
comprising 20–25% of its mass, and it is known to carbonate readily. An experimental plan was prepared in which the cement was represented by isolated portlandite crystals, which were carbonated at controlled temperature and different RHs. The partially carbonated crystals were immersed in an aqueous solution in order to monitor the diffusion of Ca2+ ions through the CaCO3 cover. This process occurs by transport across two physically separated interfaces: one between air and “water” and the other between “water” and a solid, in the title paper, a single crystal of Ca(OH)2 (Fig. 1). Dissolution of carbon dioxide from the atmosphere into the pore solution has been reviewed by Cowie et al. [2]. The sorption of CO2 from a gaseous atmosphere is greatly facilitated by increasing the aqueous pH [3] and portlandite is sufficiently soluble to achieve a major rise in pH. This dissolution process is believed to be partly controlled by dislocations in solid portlandite [4]. Many other factors potentially influence transport of ions across and between the two interfaces [5–7] but in the title study, every effort has been made to reduce transport time between the two interfaces by decreasing the thickness of the aqueous layer. The kinetics of carbonation of Ca(OH)2 have been previously studied using powder [8] and single crystals [9,10]. According to [9], both calcite and vaterite were grown on Ca(OH)2 crystals exposed to humid CO2; surface nucleation of CaCO3 on portlandite crystals occurred and subsequent random growth of calcite was observed. The observations confirm that Ca(OH)2 is not topochemically converted to CaCO3, but that a dissolution–reprecipitation process occurs resulting in nucleation at many sites with eventual coalescence of carbonate crystallites. This mechanism is consistent with the observations at ~ 20 °C of a rate maximum at RH in the range of 50–80%. The dissolution–precipitation hypothesis was also confirmed by studying crystals carbonated in air
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Ca 2+ + 2OH-
AIR
VAPOUR
WATER
CALCITE
PORTLANDITE
CO2
Fig. 1. Schematic showing diffusion through calcite layer from portlandite crystal and from wet environment.
at 93% RH and in carbonated solutions [10]. According to these latter authors ‘islands’ of CaCO3 grow epitaxially on the surface of portlandite crystals in specific preferred orientations. They also claim that “no porosity is generated during the reaction, which progresses through the formation of fractures”, which “form in response to the stresses generated in the replacement rim due to the positive volume change during the transformation of portlandite into calcite” [10]. 2. Experimental First the dissolution rate of Ca(OH)2 crystals into initially pure water was measured using crystals with one exposed face. These data were used to benchmark data obtained subsequently using portlandite crystals from the same batch which had been carbonated under controlled conditions, and again subjected to dissolution. The aqueous solution was periodically analysed for Ca2+ using an ion-selective electrode. In a fixed geometry, the difference in aqueous solution compositions normalised to release per unit surface area, could then be attributed to the protective ability of the CaCO3 film. The Ca-sensitive electrode could be located at ~2 mm from the surface being measured, thereby increasing the sensitivity of measurement. The method works for two reasons: first, the solubilities of portlandite and calcite differ greatly, with portlandite being much the more soluble by a factor of ~ 100 to 1000, and second, because the time taken for Ca2 + ions to diffuse through the aqueous film and reach the sensor, the “dead time” of the experiment, as estimated from literature values for the diffusion coefficient of free calcium ions in water, is negligible. Thus the time taken for Ca2 + to diffuse to the sensor is negligible relative to the time required, often extending over weeks, for the ion constituents of Ca(OH)2 to diffuse through the CaCO3 film.
Fig. 2. SEM of a portlandite crystal showing pseudo-hexagonal prismatic morphology (bar = 1 mm).
related defects at points of contact. The purity and mineralogical nature of the crystals were confirmed by DTA/TG as well as by powder XRD. Single crystals were carbonated by exposure to CO2-containing atmosphere at controlled humidity, generally in the range of 50–90% RH, at ambient temperature. A set of samples was exposed to 100% CO2 atmospheres at 53, 75 and 90% RH. A second set of samples was exposed to air in a lab atmosphere, ~23 °C. With time, layer thicknesses of a few microns to tens of microns could be formed and the thickness was measured by SEM from cross-sections of the crystals. Fig. 3 shows the external appearance of a crystal following carbonation at 53% RH and 100% CO2 for 5 weeks. In general the morphology of the portlandite crystal, including fine features such as layer steps, is faithfully preserved. To measure carbonation kinetics and establish the protective nature of the CaCO3 layer both carbonated and uncarbonated single crystals were coated on all but one face with an impermeable epoxy coat and the unprotected surface area was measured by contrast image analysis. The uncarbonated crystals were used to benchmark the subsequent measurements. The partially coated crystals were placed in 5 ml of initially pure de-aerated water, as shown in Fig. 4, with a calcium selective electrode (with a Ag/AgCl reference electrode). In uncoated samples the portlandite crystal dissolves in the aqueous phase and the corresponding rise in calcium concentration is monitored as a function of time. The distance between the calcium electrode and the surface of the crystal was kept small, a few mm, to reduce the “dead time”. The ratio
2.1. Procedure Single crystals of portlandite were grown by mixing aqueous solutions of NaOH and CaCl2 at ~23 °C, according to the procedure described in [9]. Atmospheric carbon dioxide was excluded from the apparatus. Growth took ~ 5 weeks. The resulting product contained transparent to translucent distorted prismatic crystals or clumps of a few individual crystals ranging up to up to ~ 5–7 mm: an example is shown in Fig. 2. Individual crystals could thus be hand-picked by eye. The crystals were handled with soft-tipped tweezers to avoid inducing pressure-
Fig. 3. SEM of carbonated portlandite crystal (bar = 1 mm).
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Electrode CaCO 3
Crystal 5 ml
Fig. 6. SEM of lightly carbonated surface showing nucleation of CaCO3, with angular morphology. Fig. 4. Picture and sketch of experimental setup for measuring calcium dissolved from carbonated and uncarbonated crystals. The spacing between the electrode tip and the surface of the crystal was ~2 mm.
commercial standards and it proved stable throughout the course of experiments. 3. Results
mass of solution/mass of crystal was in the order of ~200, ensuring that the total mass of the crystal was more than sufficient to saturate the 5 ml solution. The experiment was repeated using surface-carbonated crystals and the difference in time for a rise in calcium taken as a measure of the retardation effect arising from the presence of the calcium carbonate layer. The very low solubility of CaCO3 in water (~ 3 ppm Ca2 + at 20 °C), compared to that of Ca(OH)2 (865 ppm Ca2 + at 20 °C) [11] makes this method feasible: CaCO3 does not contribute significantly to raising Ca2+ concentration: overwhelmingly, only Ca(OH)2 diffusing through the layer of CaCO3 will significantly increase calcium concentrations in solution [12]. Thus the rates at which the calcium concentration changes as a function of time is the key to kinetic analysis. Similar experiments were performed in solutions partially saturated with respect to Ca(OH)2, with initial Ca2+ concentrations ranging from 50 to 500 ppm. The 50–500 ppm concentrations were achieved by adding a weighed amount of CaO to a known volume of solution. Three replicates were used for each condition. The measurements were performed in CO2-free conditions in a glove box filled with N2: temperature was kept constant at 22–24 °C and the RH varied between 35 and 45%. Calibration of the electrode was checked daily against
Fig. 5. SEM of crystal showing dislocations, imperfections and growth steps (bar = 200 μm).
3.1. Preparation of portlandite crystals Figs. 2 and 5 show the appearance of a newly-grown portlandite crystal. The crystals have a pseudo-hexagonal prismatic morphology and do not grow as platelets, as might be expected. The longer the crystals were left ‘growing’, the more perfect the hexagonal prismatic morphology became, although some growth steps always remained. In the SEM image of Fig. 5, the surface roughness of crystals is apparent: the surface topology is dominated by growth steps, presumably resulting from the pile-up of many atomic step height layers, as well as accretion of small “seed” crystallites. The true surface area available for subsequent reaction is thus somewhat greater than that which was estimated by microscopy. 3.2. Carbonation of portlandite crystals 3.2.1. In 100% CO2 atmospheres The initial carbonation is not uniform but occurs by nucleation of CaCO3 at active sites on the portlandite surface. Figs. 6 and 7 show the first signs of nucleation in crystals exposed to CO2 atmospheres for 5 weeks at 53% RH. Multiple rhomb-shaped and dendritic crystals of CaCO3 decorate the surface, apparently at random, but probably originating at dislocations.
Fig. 7. High magnification of early stage carbonation showing dendritic growth of CaCO3. As in Fig. 6, the crystal was exposed to 100% CO2 at 53% RH for 5 weeks.
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1
2
Fig. 10. Detail of an air-carbonated surface showing different morphologies, ranging from rosette shapes (area 1) to prisms and doubly-terminated rhomboids (area 2). Fig. 8. Highly carbonated area covered with rough textured polycrystalline CaCO3 coating. Crystal was exposed to CO2 at 53% RH for 5 weeks. Note that the visual appearance of the surface is porous.
Carbonation progresses by a combination of growth from apparently randomly-nucleated sites as well as by continued nucleation at fresh sites. Eventually the surface becomes covered as the nucleated sites grow and coalesce. The carbonation mechanism shows that (i) nucleation of CaCO3 occurs at numerous sites so the resulting film is polycrystalline and (ii) from the micrographs, the film appears to become continuous or nearly so (see Fig. 8). This picture reveals no obvious evidence of preferred crystallographic orientation, in agreement with [9]. This does not mean that they are not oriented, but that the substrate is not visible so the relations cannot be determined by microscopy. The mass of CaCO3 grown on crystals carbonated at 53% RH for 5 weeks was ~ 5% of the crystal mass, as measured by TG, so layer thicknesses were on the order of tens of microns.
3.2.2. In laboratory atmosphere Figs. 9 and 10 show images of the faces of a crystal carbonated in air at room temperature and conditions for ~6 months. The layer is rough and the coalescence of crystallites makes it impossible to delineate morphological features within the layer but isolated growth also occurs at the surface as shown in Fig. 10. These surface features are believed to be forming on a substrate layer of CaCO3. Fig. 11 is a transversal cut section of a crystal following carbonation in air for ~ 6 months. The upper part of the image shows the CaCO3 ‘layer’ grown on one of the rectangular faces, the depth of which has
reached ~ 100 microns. The surface projections are conspicuous and the film is thus not homogeneously distributed across the surface. Figs. 12–14 show images of a “hexagonal face”, pseudo {0001}, of a crystal partially carbonated in laboratory air for ~6 months at progressively higher magnifications. Again different morphologies can be appreciated: ‘pyramids’, small needles and dendrites occur. Fig. 14 apparently shows ‘holes’ or ‘pores’ on the surface, which develop during the progressive formation of the carbonated layer. 3.3. Dissolution of surface carbonated and uncarbonated crystals Fig. 15 shows dissolution rates of uncarbonated and carbonated portlandite crystals when contacted with initially deionised water. Fresh, uncarbonated surfaces give a virtually instantaneous release of Ca2+ ions to solution. On the other hand, previously carbonated crystals, for example when carbonated at 75% RH, have a definite induction period before dissolution of OH− and Ca2+ ions commences. Under the conditions used, 23 °C, the solubility product constant Ksp of portlandite (5.02 × 10−6) is approximately 1000 times greater than that of calcite (3.36 × 10− 9). This means that the charge balancing ion for OH− is Ca2 + and vice versa; no other cations other than H3O+ are present and the aqueous concentration of H3O+ decreases to near zero as the pH rises. Therefore we assume, neglecting species such as Ca(OH)+, that each Ca2+ ion in solution is charge balanced by 2OH. Uncarbonated portlandite releases Ca2 + (and OH) ions almost immediately whereas release from surface-carbonated crystals is much delayed. This delay is associated with the time required for OH−
Surface
Fig. 9. SEM of a carbonated crystal exposed to air in a lab for 6 months at ~20 °C. The RH was not controlled but varied between ~25 and 70%.
Fig. 11. SEM image of transversal section showing CaCO3 layer grown on top of the crystal that was air-carbonated for 6 months indoors.
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Fig. 14. Partially carbonated hexagonal surface of crystal exposed to air for 6 months. Note the holes (black) apparently giving rise to a reticular network of CaCO3 crystals.
and Ca2+ ions to diffuse through the CaCO3 film and it is apparent that the film can be strongly protective against dissolution. Regarding the saturation degree of the solution, uncarbonated crystals eventually reach a saturation value of ~900 ppm Ca2+ after about 500 h, which accords with the known equilibrium calcium ion concentration in that solution [11]. However when CaCO3 is also present, the shape of the solubility curve in the Ca(OH)2–CaCO3–H2O system governs the equilibrium Ca2 + concentration. This curve has a degree of freedom, so its numerical value depends on the aqueous molal ratios of the two solid components. In general, the saturation concentration of Ca2+ is less than 900 ppm and actual values, as recorded in Fig. 15, lie in the range of 300–700 ppm. As expected, the dissolution rate (i) diminishes significantly with time as saturation is approached and (ii) is affected by the continuity and thickness of the CaCO3 coating which hinders the dissolution. In all surface carbonated crystals made at 75% RH, dissolution requires a significant “breakthrough” period, ranging between 50 and 100 h. The decrease both in the dissolution rate and in the saturation value is more pronounced for crystals carbonated at intermediate RH than for 90% RH (data for which are not shown) evidencing the impact of humidity on the quality of the coating layer that can be grown on single crystals. Films grown at high RH, 90%, provide little protection against dissolution of the portlandite substrate compared to films grown at
~75% RH, as evidenced by the development of a “breakthrough” period and its duration. As noted, the increase in calcium concentration, Fig. 15, is also related to the exposed surface area. To obtain quantitative measures, we normalise values: dissolved calcium was divided by the crystal exposed area. These data are shown in Figs. 16–19. The maximum values of calcium dissolved are much lower for the crystals with CaCO3 layers grown at intermediate RH than for either uncarbonated crystals or those carbonated at high (N90%) RH. As for the influence of the initial concentration in Ca2+ ions of the solution where the crystals were placed to dissolve, rates do not much vary between pure water and solutions conditioned with 50 ppm calcium as Ca(OH)2. Much higher initial concentration, 500 ppm calcium, has an important influence in the dissolution of the crystal: as expected from the law of mass action, the lower the initial concentration in the solution, the more rapidly portlandite will dissolve. This can also be appreciated by comparing the initial slope of the data, or initial dissolution rate, which is in all cases considerably lower for crystals dissolving in 500 ppm Ca solutions relative to initially pure water. Table 1 summarises average values for the induction period for the crystals in the 4 stages of carbonation and in the 3 solutions. The breakthrough period comprises the time during which the increase in the concentration of calcium in the solution is less than ~5 ppm/mm2. The uncarbonated crystals start dissolving almost immediately after being put in contact with the solution. Values for those carbonated at 90% RH do not differ very much from initially pure water, implying that
[Ca2+] / ppm
Fig. 12. Hexagonal phase of a crystal carbonated in air.
1200 1100 1000 900 800 700 600 500 400 300 200 100 0
Uncarbonated (benchmark)
Partially carbonated (75% RH)
0
100
200
300
400
500
600
Time / hours Fig. 13. Detail of the different morphologies of CaCO3 grown on the hexagonal face, {0001}, of a crystal exposed to air for ~ 6 months. The ‘pyramids’ have not been observed in n o prismatic faces 0110 of crystals carbonated under the same conditions in air (Figs. 9 and 10).
Fig. 15. Calcium dissolved from uncarbonated (blue markers) and partially carbonated at 75% RH (red markers) crystals in water. The different markers for each colour represent different replicates. (For interpretation of the references to colour in this figure legend, the reader is referred to the web version of this article.)
Calcium in solution per unit area (ppm/mm2)
I. Galan et al. / Cement and Concrete Research 74 (2015) 68–77
300
73
500 ppm 50 ppm 0 ppm
250 200 150 100 50 0 0
100
200
300
400
500
600
time (hours)
Calcium in solution per unit area (ppm/mm2)
Fig. 16. Evolution of dissolved calcium (per unit of crystal area) in solutions where uncarbonated crystals were placed to dissolve. Starting calcium concentrations are 0, 50 and 500 ppm. For each concentration 3 replicates were used. Each symbol (square, circle and triangle) represents data from 3 crystals.
300
500 ppm 50 ppm 0 ppm
250 200 150 100 50 0 0
100
200
300
400
500
600
time (hours)
Calcium in solution per unit area (ppm/mm2)
Fig. 17. Evolution of dissolved calcium (per unit of crystal area) in solutions where carbonated crystals (at 53% RH) were placed to dissolve. Starting calcium concentrations are 0, 50 and 500 ppm.
300
500 ppm 50 ppm 0 ppm
250 200 150 100 50 0 0
100
200
300
400
500
600
time (hours)
Calcium in solution per unit area (ppm/mm2)
Fig. 18. Evolution of dissolved calcium (per unit of crystal area) in solutions where carbonated crystals (at 75% RH) were placed to dissolve. Starting calcium concentrations are 0, 50 and 500 ppm.
300
500 ppm 50 ppm 0 ppm
250 200 150 100 50 0 0
100
200
300
400
500
600
time (hours) Fig. 19. Evolution of dissolved calcium (per unit of crystal area) in solutions where carbonated crystals (at 90% RH) were placed to dissolve. Starting calcium concentrations are 0, 50 and 500 ppm.
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Table 1 Average breakthrough time (h) for starting dissolution of uncarbonated and carbonated crystals in 0, 50 and 500 ppm Ca solutions. Initial solution composition
Uncarbonated
Carbonated at 53% RH
Carbonated at 75% RH
Carbonated at 90% RH
0 ppm 50 ppm 500 ppm
3 3 4
50 18 98
72 56 216
5 4 34
the coating is not effective, except in the case of the 500 ppm solution where it delays the start of dissolution for 30 h. All samples carbonated at 75% RH show an induction period in excess of 50 h. This time quantitatively assesses the good protection afforded by the coating which is on the order of 10–100 microns thick (Fig. 11). The layer formed at 53% RH provides some protection but presents breakthrough times lower than the coating at 75% RH. An approximate value of the effective diffusion constant of the protective layer can be calculated (see Table 2). The difference between the numerical value of this coefficient and the diffusion coefficient of calcium at infinite dilution, 7.92 × 10−6 cm2/s [11], provides a measure of the protection afforded by the carbonate layer: see Appendix A for more details. The persistence and durability of the carbonate coating have been measured by prolonged water immersion. Figs. 20–23 show images of crystals previously carbonated at 53% RH for 5 weeks and then exposed to initially pure water (Figs. 20 and 21) and to 50 ppm Ca2+ solutions (Figs. 22 and 23) for ~30 days at ~23 °C. The white cover of CaCO3 remains largely intact while portlandite has dissolved from underneath the calcite coat (Figs. 20 and 22). Fig. 21 shows at high magnification a SEM image of the area marked A in Fig. 20. A generally dense layer of CaCO3 has become perforated allowing dissolution to concentrate at particular sites. Fig. 23 shows a detail of the area marked B in Fig. 22. Small, 2–3 micron, crystals randomly distributed and decorating a denser and more uniform substrate can be appreciated. This texture is believed to arise from partial recrystallization within the calcite protective layer in the course of the dissolution experiments. The following three SEM images show the evolution of the dissolution of a crystal partially carbonated in air. In this case, the solutions were stirred and the crystals were examined by SEM after 5, 30 and 60 min. Before the dissolution the crystals looked like the example depicted in Fig. 9. After 5 min (Fig. 24), triangular etch pitch appear and dissolution in steps starts to be visible. After 30 min (Fig. 25) pits with a defined morphology appear. The size of these pits is between 10 and 20 microns. After 60 min (Fig. 26) the pits have grown to ~ 500 microns. The layers inside the pits (up to 10) can be clearly seen. A noteworthy feature of the etch pits is their consistent orientation over distances of 1–2 mm. This orientation suggests that the calcite layer may indeed have an epitactic relation to its single crystal substrate.
A
1 mm
Fig. 20. Picture obtained by optical microscopy of a ca 1 × 3 mm carbonated crystal after exposure to dissolution in water for ~30 days. The CaCO3 cover (white) seems to remain largely intact: note especially the nearly complete margin of CaCO3 delineating the margin of the grain. The crystal had a small outgrowth of CaCO3 at point A: a higher magnification view of this area is shown in Fig. 21.
processes occurring during cement hydration, for example to control the length of the so-called dormant period of hydration. But proof of a barrier effect has often been indirect and the constitution and effectiveness of “barrier layers” remain conjectural. In the present study, it is shown that barrier layers can be formed in the course of use by spontaneous processes such as carbonation. The properties of CaCO3 formed by reaction of portlandite with water and CO2 gas have been investigated and proof obtained that in some circumstances, a layer of calcite developing on portlandite can significantly protect portlandite against subsequent carbonation and dissolution. 4.1. Influence of relative humidity Carbonation reaction proceeds very slowly in dry air but develops rapidly in the presence of water because of the need to dissolve CO2 in a water film. Once initiated, reaction liberates water which, if not evaporated, facilitates reaction. Reaction thus proceeds across two interfaces, one between portlandite and aqueous solution, the other between aqueous solution and air (see Fig. 27). At low RH, when a thin water film wets portlandite, the two interfaces may be very close, around 3 to 18 nm (10 to 60 water molecules
4. Discussion The existence of barrier layers in cement reactions has often been proposed to explain the observed kinetics of many features and
Table 2 Apparent diffusion coefficient through the layers between the surface of the crystals and the electrode (carbonate layer plus water). D (10−6 cm2/s) Uncarbonated Carbonated at 53% RH Carbonated at 75% RH Carbonated at 90% RH
1.17 0.038 0.031 0.086
Fig. 21. High magnification SEM image of CaCO3 cover: area marked as A in Fig. 20. Note the appearance of what looks to be holes.
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B
Fig. 24. SEM of an air-carbonated crystal after 5 min dissolving in water. Fig. 22. SEM image of a carbonated crystal after dissolution in 50 ppm Ca solution for ~30 days. Note that rupture of the protective CaCO3 film (B) which follows dissolution of the more soluble core of portlandite. The collapse of the calcite shell may result in part from the coating process, high vacuum and beam damage in the microscope.
thick). At the portlandite–aqueous interface, portlandite dissolves, increasing the pH of the aqueous phase, while at the air interface, atmospheric CO2 dissolves, initially as molecular CO2 but subsequently forming “carbonic acid”; the latter partially dissociates to mixtures of bicarbonate and carbonate species. Although the initial dissolution reaction is slow, the bicarbonate–carbonate equilibrium is rapidly reversible and the equilibrium speciation is pH dependent with carbonate increasingly favoured with respect to bicarbonate species as pH increases. However, the reactions at the two interfaces are only indirectly coupled: one proceeds independently of the other. The indirect link does however exist (i) because the rate of dissolution of CO2 at the air interface is much enhanced by rising pH and the source of OH− ions, essential to increase pH, is at the portlandite–solution interface and (ii), the transport of species is limited by the thickness of the liquid and its diffusive characteristics. In the experimental work reported, every effort has been made to keep the film as thin as practicable and provide direct diffusion paths. The point at which carbon containing species react with Ca(OH)2 forming CaCO3 depends on the above factors as well as temperature. But it should be recalled that CaCO3 formed is several orders of magnitude less soluble than the substrate portlandite and, as a result of this
as well as the physical conditions imposed on diffusion and mass transport by the existence of two physically separated interfaces, a layer of relatively insoluble CaCO3 often tends to develop on portlandite surfaces. The reaction is sensitive to RH. At low RH it is difficult to establish a water film on the surface and reaction is slow because the first step, dissolution of CO2 into aqueous phase, is hindered by the low and likely discontinuous surface coverage. Additionally, although water is released by the carbonation reaction, it is prone to evaporate at low humidity. On the other hand, at very high RH (N90%), the higher number of water monolayers give a continuous but thicker water coverage and the calcium carbonate may nucleate in the aqueous layer as well as at the crystal/solution interface, allowing a more complex morphology to develop in the carbonate layer; this influences product morphology and orientation according to mode of formation. Diffusion paths also lengthen as the water film thickens. A further factor is that the reaction generates more water: water balances depend on competition between the rate of generation and its evaporation or retention. At intermediate humidity, 53%, and more clearly in present results, 75%, the conditions of thickness of the water layer on the crystals and distance between the two interfaces (air/solution and solution/crystal) seem to be the optimum for development of low permeable carbonation layers which also adhere strongly to the substrate crystal. The influence of the different polymorphs formed at different RH [1] cannot be directly related to the protection offered by the film.
Fig. 23. Detail of SEM picture of CaCO3 cover of partially dissolved crystals (area marked as B in Fig. 22).
Fig. 25. SEM of an air-carbonated crystal after 30 min dissolving in water.
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ages. Crystal morphologies become more regular and suggestive of near-equilibrium conditions of growth. Thus the film properties are best determined by actual measurement of their diffusive properties. It is also noteworthy that the conditions of formation and measurement are such that they could readily be achieved in practice. Formation occurred in limited volume of water, but this ideal thickness is spontaneously developed by exposure to air and CO2. Optimum conditions for growth are close to self-stabilising as water lost by evaporation is replaced by water liberated at the interface with substrate portlandite.
4.3. Film properties by electron microscopy
Fig. 26. SEM of an air-carbonated crystal after 60 min dissolving in water.
4.2. Influence of the neoformed carbonation coverage on further carbonation Results show that the quality of the film is very dependent on the conditions prevailing in the course of its initial growth. In the most favourable conditions the film had sufficiently low permeability to diffusion of Ca(OH)2 as to require 10 to 100 h for breakthrough of ions from substrate portlandite to affect the aqueous phase. And, once breakthrough is achieved, the rate of diffusion through the protective calcite film is slowed relative to unprotected crystals. Experience showed that it was difficult to judge the performance of coatings by microscopy alone: at higher magnifications necessary to detect pores or imperfections, it was only possible to examine a small area. The granularity of the film, as observed by electron microscopy, was not necessarily an indication of its permeation properties. Nor was the early-formed partial coverage an indication of the quality of the fully developed cover. Thus the early stages of surface cover gave a wide range of morphologies: initially, crystals were sparse and had acicular, rosette, dendritic or other morphologies typical of growth from supersaturated solutions. However many of the more successful films were grown slowly over weeks or months and had more regular prismatic morphologies which appeared to consist of an interlocking network of prismatic grains. It is believed that films, once formed, may not be morphologically stable but in favourable cases may recrystallise with ageing and growth, thereby gaining in strength and coherence and with infilling of pores. Thus while the early stages of growth of the calcite layer are driven by morphological features and defects on the substrate, the coalescence of growth and the internal recrystallization and accretion of more material govern the film properties at longer
AIR : CO2
(Ca+2 OH-) (CO3= H+) Pore solution
Although electron microscopy has been immensely valuable to characterise the surface state of portlandite crystals and visualise features such as the microstructure and thickness of calcite films, it has limitations. The principle limitation is the inability of electron microscopy to estimate the permeation properties of the protective layers. None of the layers look particularly impermeable even when actual measurements disclose extremely low film permeabilities, as evidenced from breakthrough times. In this respect only actual measurements of film properties can be treated as a reliable indicator of permeation properties.
4.4. Overall perspective We have shown that in certain regimes, calcium carbonate films can form spontaneously which are remarkably effective in limiting diffusion of calcium hydroxide from cement and in protecting further carbonation of substrate portlandite. The question immediately arises: to what extent can data derived from portlandite single crystals be used to develop carbonation resistant concrete? Can spontaneous carbonation processes be used as part of a protective system for durable concrete? We have not done the determination of whether the protective film can be established on commercial portlandite and Portland cements, the latter containing only 20–25% polycrystalline portlandite. But we predict that the self-protection mode is most likely to be demonstrated by the behaviour of cements based on portlandite where solubility is most effectively suppressed by formation of optimised calcite coatings. Portland cements will however require a great deal of further investigation: the paste is only 20–25% portlandite although carbonation of other phases, e.g., CSH, may contribute to establishment of protective surface layers. Many unknowns remain but it is at least possible to envisage a new generation of cementitious materials which will have selfgenerated and self-protective surface layers which ensure longer service lifes in conditions where leaching, carbonation and resistance to penetration by adverse chemicals in the service environment occur.
AIR : CO2
(Ca+2 OH-) (CO3= H+) Pore solution CaCO3
Ca(OH)2
Ca(OH)2
AIR : CO2
Pore solution CaCO3
Ca(OH)2
Fig. 27. Schematic showing the initial establishment and evolution (from left to right in the figure) of a relatively impermeable layer of calcite on a portlandite substrate in a moist atmosphere. The water film is coloured blue. The thickness of the stable water film relates to the relative humidity with RH values in the range of 53–75% giving the best coverage. (For interpretation of the references to colour in this figure legend, the reader is referred to the web version of this article.)
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5. Conclusions The results obtained have enabled to draw up the following conclusions: 1. Portlandite crystals develop a layer of calcium carbonate whose properties of continuity and thickness depend on the external RH. 2. The protective ability of the carbonated film has been for the first time assessed by: a. Measurement of the calcium leached from uncarbonated and carbonated crystals. This protective ability afforded by optimum carbonation has been measured in different conditions and optimum conditions established. b. Calculation of an apparent diffusion coefficient of the calcium through the layers between the surface of the crystals and the electrode (carbonate layer plus water). The so calculated coefficients are 1-2 orders of magnitude lower for carbonated than for uncarbonated crystals. 3. Despite the polycrystalline nature of the CaCO3 layer it provides significant protection against diffusion of Ca2+ and OH− ions. 4. The foundations are laid whereby spontaneous processes can be used to develop the self-protecting ability of lime and possibly Portland cement against a range of degradative processes. Appendix A
h
A one-dimensional solution of Fick's second law, C ðx; t Þ ¼ C s i 1−erf 2pxffiffiffiffiffi was applied where x is the distance between the surface Dt
of the crystal and the position of the electrode, t is the time taken to reach the concentration C at the electrode, D is the apparent diffusion coefficient through the layer of carbonate plus the water and Cs is the concentration at the surface of the crystal or the saturation value. This equation was chosen assuming that the dissolution of Ca(OH)2 is much faster than the diffusion of the ions. Table 2 includes D values for pure water. For these calculations, the saturation value was taken as 900 ppm, a time of 170 h was chosen (as the time required to reach steady state) and the distance between surface of the crystal and electrode was 2 mm.
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The coefficient for the uncarbonated case is not equal to the coefficient in water, but is around seven times smaller and additionally, the diffusion coefficients for the carbonated crystals are significantly lower, 1–2 orders of magnitude, the lowest value being obtained for the crystal carbonated at 75% RH. Two explanations can be given for these differences: - The own kinetics of dissolution which can be not fully instantaneous, being slower in the case of the presence of the carbonated layer. - Being the dissolution instantaneous, there is a time to build up the saturation degree at just the crystal solution interface.
Even though a perfect calculation cannot be achieved with present results, and the values in Table 2 also include the diffusion through the 2 mm of water, it can be clearly appreciated that the coating slows down the process considerably: for the layer grown at 75% RH, D is 250 times less than the corresponding value for initially pure water. References [1] E. Dubina, L. Korat, L. Black, J. Strupi-Šuput, J. Plank, Influence of water vapour and carbon dioxide on free lime during storage at 80 C, studied by Raman spectroscopy, Spectrochim. Acta A Mol. Biomol. Spectrosc. 111 (2013) 299–303. [2] J. Cowie, F.P. Glasser, The reaction between cement and natural waters containing dissolved carbon dioxide, Adv. Cem. Res. 4 (1991) 119–134. [3] W. Stumm, J.J. Morgan, Aquatic Chemistry, 3rd ed. Wiley, NY. USA, 1996. [4] I. Galan, F. Glasser, C. Andrade, D. Baza, Dissolution of portlandite, 13th International Congress on the Chemistry of Cement, 2011, pp. 1–7. [5] D.E. Giles, I.M. Ritchie, B.- Xu, The kinetics of dissolution of slaked lime, Hydrometallurgy 32 (1993) 119–128. [6] J. Wang, T.C. Keener, G. Li, S.- Khang, The dissolution rate of Ca(OH)2 in aqueous solutions, Chem. Eng. Commun. 169 (1998) 167–184. [7] K. Johannsen, S. Rademacher, Modelling the kinetics of calcium hydroxide dissolution in water, Acta Hydrochim. Hydrobiol. 27 (1999) 72–78. [8] S.- Shih, C.- Ho, Y.- Song, J.- Lin, Kinetics of the reaction of Ca(OH)2 with CO2 at low temperature, Ind. Eng. Chem. Res. 38 (1999) 1316–1322. [9] J.R. Johnstone, F.P. Glasser, Carbonation of Single Crystals of Portlandite in Cement Paste, 9th International Congress on the Chemistry of Cement (ICCC) 1992, pp. 370–376. [10] E. Ruiz-Agudo, K. Kudlacz, C.V. Putnis, A. Putnis, C. Rodriguez-Navarro, Dissolution and carbonation of portlandite [Ca(OH)2] single crystals, Environ. Sci. Technol. 47 (2013) 11342–11349. [11] D.R. Lide, CRC Handbook of Chemistry and Physics, 2005. [12] I. Galan, F.P. Glasser, D. Baza, C. Andrade, Calcium carbonate coating on portlandite crystals, 33rd Cement and Concrete Science Conference, 2013.