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Effects of Calcium and Other Ionic Impurities on the Primary Nucleation of Burkeite Bing Shi,† W. James Frederick, Jr.,‡ and Ronald W. Rousseau* School of Chemical Engineering, Georgia Institute of Technology, Atlanta, Georgia 30332-0100
Primary nucleation of burkeite, which is a crystalline solid solution of sodium carbonate and sodium sulfate having the stoichiometry Na2CO3‚2Na2SO4, was monitored using an on-line particle-size analyzer. Observations of unusual nucleation behavior led to explorations that identified specific impurities in the system that affected nucleus formation. By varying the conditions under which nucleation occurred, calcium ions were found to inhibit nucleation, and further research showed that the substitution of calcium ions into the burkeite crystal lattice was the reason for the observed behavior. Introduction The mineral burkeite (Na2CO3‚2Na2SO4) was first prepared artificially by W. E. Burke in 1919 when he was studying the solid-liquid equilibrium of salts from Searles Lake brines.1 Its existence in nature was later confirmed in salt samples from a well during a search for potash.2 The most common industrial use of burkeite is in the detergent industry where, for ecological reasons, it has replaced phosphate carriers of detergent components. In that application, burkeite crystals usually are engineered from blocky to porous agglomerates using additives such as phosphates and ionic polymers.3-5 The present research was prompted by the presence of burkeite, a double salt of sodium sulfate and sodium carbonate, in scale formed on the heat-transfer surfaces in evaporators used to concentrate what is known as black liquor; the liquor is obtained from alkali treatment of wood pulp. (Black liquor is concentrated in multipleeffect evaporators and then burned to recover thermal energy and to regenerate pulping chemicals. During evaporation, the inorganic salts often form scale on heattransfer surfaces, which reduces evaporator throughput.) Either nucleation of burkeite on the heat-transfer surface or attachment of fine crystals to that surface is believed to be the major scaling mechanism. Accordingly, controlling the formation of fine crystals could be important in solving the scaling problem. Several properties are unusual and might have contributed to scaling. (1) Burkeite has the abnormal behavior of reduced solubility at increased temperature over the operating temperature range of black-liquor evaporators. Because the highest temperatures in the evaporator are on the heat-transfer surfaces, the solutions are more supersaturated and have the tendency to form nuclei on these surfaces. This tendency was investigated6 and found to be unimportant at relatively low heating rates. (2) Burkeite crystals tend to act as agglomerating agents.4 This tendency, when combined with the high viscosity of black liquors, might initiate * To whom correspondence should be addressed. E-mail:
[email protected]. Fax: 404-384-0185. † Current address: Pfizer Global Research and Development, 10777 Science Center Dr., San Diego, CA 92121. ‡ Current address: Department of Chemical Engineering and Environmental Science, Chalmers University of Technology, SE-412 96 Gothenburg, Sweden.
scale formation by binding burkeite crystals with other type of scale (e.g., calcium carbonate) already formed, or it might increase the growth rate of scale by binding particles of burkeite or other materials to existing scale. (3) Crystallization of burkeite has been found to be susceptible to the influence of very small amount of additives, e.g., phosphates and ionic polymers.3,5 Because the constituents of black liquor are extremely diverse, it is clear that there are several that can influence the crystallization of burkeite. The research described here is focused on primary nucleation because the conditions in solutions from which primary nuclei are generated were thought to be more likely to contribute to scaling. The nucleation mechanisms and effects of some additives on the nucleation kinetics of burkeite were investigated. Crystallization from model aqueous solutions of sodium carbonate and sodium sulfate was used to simulate blackliquor evaporation. This approach avoided difficulties in handling black liquor, which is a dark, highly viscous solution that must be processed at elevated temperatures. Also, separating crystals from black liquor is exceedingly difficult. Finally, with model solutions, the effects of individual constituents or additives can be directly determined without interference of other species, many of which are unidentified, that are normally present in black liquor. Experimental Equipment and Procedures The model solution was prepared by dissolving sodium carbonate and sodium sulfate in distilled water at a 1:2 molar ratio of sodium carbonate to sodium sulfate, which corresponds to the ratio of these species in burkeite. This solution will be referred to in the remainder of this paper as a burkeite solution. The chemicals used in the experiments were all ACS-grade reagents. The initial solution compositions were usually 30 wt % in solutes, although other compositions were used occasionally. After the solids had been completely dissolved, the solution was decanted into the crystallizer that was assembled immediately to conduct the experiment. Crystallizations were carried out in a 1-L Parr reactor (Figure 1). A stainless steel impeller driven by a motor with variable rotating speeds was used for mixing. The mixing intensity was maintained by keeping the impel-
10.1021/ie020845e CCC: $25.00 © 2003 American Chemical Society Published on Web 05/08/2003
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Figure 1. Experimental apparatus used for evaporative crystallizations
Figure 2. Typical operating profiles for evaporative crystallization experiments. P is the system pressure, T is the system temperature, and m is the mass of condensate (water) collected.
ler Reynolds number above 12 000, which ensured turbulent flow throughout the crystallizer. A thermocouple and a pressure transducer were used to measure the system temperature and pressure, respectively. A PID controller adjusted the heater to provide the desired system temperature (generally 115 °C) and evaporation rate (3-5 g/min). The evaporated water was condensed and weighed so that the composition of the residual mixture in the reactor could be estimated by mass balance. Evaporated water left the crystallizer through a manually operated valve; it was condensed in a glass condenser, collected in a beaker, and weighed so that the solids content of the residual solution could be estimated by mass balance. During the experiment, the data, including system temperature, pressure, and mass of condensate, were recorded at 30-s intervals by a computer. Typical profiles of the temperature, pressure, and mass of condensate are shown in Figure 2. Crystal sampling was made possible by a vertical sampling tube whose entrance was located about 1 cm above the bottom of the crystallizer. When the sample
Figure 3. Schematic diagram of probe and particle analyzed: (a) tip of the FBRM probe (courtesy of Lasentec Inc.), (b) chords that can be scanned by the laser beam.
valve was opened, a small portion of the suspension was forced into a plastic container. Further details regarding sampling are given elsewhere.7 The thermowell inside the reactor was rearranged to accommodate the 19-mm-diameter probe of a Lasentec FBRM (focused beam reflectance measurement) D600L system. The detector, shown in Figure 3, provided online monitoring of particle dimensions (chord length). As each measured chord length is not a unique particle property, the chord length distribution (CLD) can be used only as a semiquantitative measure of particle
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Figure 4. Typical FBRM results obtained during the evaporation of sodium sulfate solutions. Note that the evaporative crystallization process has been divided into three regimes according to the trends of the curves shown in the data.
sizes. In the present experiments, the FBRM system measured the CLD at intervals of 30 s. In the experiments discussed here, the FBRM probe was configured to detect the start of nucleation by using the F (fine) electronics setting that is more sensitive to the detection of fine particles [the other setting is C (coarse) electronics]. The focal spot of the laser beam was set 100 mm beyond the probe window, which allowed for the detection of particles deeper in the solution; moving the focal plane closer to the probe window would have the opposite effect. More information of the FBRM system can be found elsewhere.8-11 Results and Discussion Interpreting the FBRM Results. Nucleation can be either primary or secondary, with the former typically showing a high-order dependence on supersaturation and the latter showing a low-order dependence. Primary nucleation often starts abruptly and results in large numbers of crystals, whereas secondary nucleation is less abrupt and results in smaller crystal numbers. Nucleation by both primary and secondary mechanisms was thought to be observed in evaporative crystallization from sodium sulfate solutions. Typical FBRM results are shown in Figure 4, which represents counts of chords in four different size ranges. The trends in the data were interpreted by dividing the evaporative crystallization process into three regimes (I, II, and III), as shown in the figure. Essentially no solute crystals were observed by the FBRM system in regime I. After about 20 min of run time, evaporation of the solution began, and the counts shown at that time can be attributed to bubbles of water vapor. No crystals were found in microscopic examination of samples withdrawn from the crystallizer during this period. Regime II was characterized by the sudden increase in counts of chords of all sizes, which was interpreted to indicate the rapid appearance of a large number of small crystals. As no crystals existed in the period immediately prior to this burst of chord counts, the sudden increase was attributed to primary nucleation. Rapid increases in chord counts (and therefore primary nucleation) lasted for only a short period of time because, with the formation of a large number of crystals, the supersaturation was lowered quickly and was no longer sufficient for primary nucleation. Regime III was characterized by a steady increase in the counts of all observed chord lengths, which was
Figure 5. Typical FBRM results showing the three operating regimes during the base-case crystallization from burkeite solutions.
interpreted to result from a combination of secondary nucleation and crystal growth. In this regime, supersaturation generated by continuing evaporation was consumed in crystal growth. Base-Case Experiments. Base-case experiments refer to evaporative crystallization of aqueous burkeite solutions (i.e., solutions having a ratio of 1 mol of sodium carbonate to 2 mol of sodium sulfate) without any additives present in the solution. They serve as control experiments throughout the investigation. Typical results obtained by the FBRM system during these runs are shown in Figure 5. (Note. The abscissa in Figure 5 is solids content, rather than time as was used in Figure 4.) Similarly to the approach with Figure 4, the data in Figure 5 also can be divided into three regimes, and initial inspection suggests that each regime is analogous to its counterpart in Figure 4. However, samples from regime I told a different story when examined under an optical microscope. It was found, rather unexpectedly, that, during the late stages of regime I, a significant number of large crystals appeared. Sample photographs taken through an optical microscope are shown in Figure 6a and b. Crystals grew very rapidly, probably because of the high supersaturation during that period, and by the end of regime I, some had grown to over 500 µm. The rapid growth rate can be explained by the high supersaturation that can be sustained through the wide metastable zone found in previous work.6 The crystals were prismatic, with some having T- or X-shaped branches, probably resulting from adherence and subsequent growth of nuclei to existing crystals. The samples showed significant agglomeration. Samples from regime II (Figure 6 c) showed several dark spots that, upon enlargement (Figure 6 d), turned out to be agglomerates of numerous small crystals. These agglomerates probably formed during sampling because, without stirring, the small crystals in suspension could quickly agglomerate and settle in the sample container. The number of small crystals in regime II was enormous in comparison to the number of large crystals in regime I. Judging from the similarities between regime II in Figure 4 and regime II in Figure 5 and the observation that so many small particles appeared in such a short period of time, regime II again was interpreted to be defined by primary nucleation. Photomicrographs of crystals from regime III are shown in Figure 6e and f. The crystals increased both in number and in size over time; as stated earlier, the
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Figure 6. Photographs of crystals sampled during different crystallization regimes: (a, b) regime I, (c, d) regime II, and (e, f) regime III.
rates at which these characteristics changed were interpreted to mean that secondary nucleation and crystal growth dominated this regime. The crystals formed during regimes I and II were examined by chemical analysis and X-ray powder diffraction. Both methods confirmed that the crystals were burkeite. However, the ratios of sodium carbonate to sodium sulfate in the crystals were different: 1:2.71 for regime I crystals and 1:2.25 for regime II crystals. This can be explained by the fact that regime I crystals were formed from solutions having compositions different from those in regime II; i.e., the carbonate-to-sulfate ratio in the solution was constantly changing during evaporation because sodium sulfate was a more preferable component in the crystal lattice. From the above observations, the regimes in the basecase experiments could be distinguished from one
another by the dominating mechanism of nucleation. In regime I, primary nucleation of burkeite produced large, but relative few, burkeite crystals; in regime II, primary nucleation of burkeite in the bulk solution produced a large number of very small crystals; as described earlier, interpretation of the data in regime III led to the conclusion that there was steady secondary nucleation and continued growth of burkeite crystals. Burkeite exhibited primary nucleation in both regimes I and II, but the crystal sizes produced in the two regimes differed greatly. In contrast, when the solute was either pure sodium carbonate or sodium sulfate, primary nucleation occurred only in regime II. This phenomenon has not been reported in the literature either because it might have escaped attention or it might be unique for burkeite. In either case, it requires further explanation.
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Figure 7. FBRM results showing that regime II nucleation consisted of two steps when 50 ppm of calcium carbonate powder was added to the solution.
Burkeite Nucleation in the Presence of Calcium Ions. The presence of calcium is ubiquitous in black liquor; its concentration ranges from 118 to 1050 mg/ kg of dry solids.12 Calcium ions in black liquor have been reported to form chelates with lignin fragments,13 and these chelates tend to dissociate at higher temperatures, releasing calcium ions to form calcium scale on heattransfer surfaces. Grace14 noted that the addition of calcium carbonate likely caused the black liquor to form sodium scale. Calcium carbonate has been observed to be present in sodium scale,15 but essentially no research has been conducted on the relationship between these two most common (calcium and sodium) scales. In the present work, calcium carbonate powder was added in experiments that were otherwise the same as the basecase experiments so as to determine the effects on burkeite nucleation. These experiments were found to be very important in disclosing the nucleation mechanisms of burkeite in the base-case experiments. In the new experiments, calcium carbonate (calcite) powder was added during the preparation of the solution of sodium carbonate and sodium sulfate. The powder particles had a size of about 5 µm, which was estimated using an optical microscope. The solubility products of sodium carbonate and sodium sulfate are about 4.95 × 10-9 at 25 °C, and they increase slightly as temperature increases. In an aqueous solution of sodium carbonate and sodium sulfate, calcium carbonate is almost insoluble because of the abundance of carbonate ions. [Note that calcium carbonate has two other polymorphs: aragonite and vaterite. Vaterite is unstable at the conditions of the present experiments and so was not considered here. Aragonite can be produced in stable form following procedures outlined by Tai and Chen.16 Subsequent experiments with aragonite showed that it and calcite influenced burkeite crystallization in the same manner.] Figure 7 shows the FBRM results obtained during evaporative crystallization when 50 ppm calcium carbonate powder was added to the initial solution. Two abrupt increases in chord counts occurred in regime II, and four such increases were observed when 100 ppm was added to the initial solution. There were two runs in which 50 ppm CaCO3 was added and four with the addition of 100 ppm. Results from all six experiments reflect intermittent primary nucleation and negligible secondary nucleation until regime II is ended. Photographs taken through an optical microscope clearly revealed the tiny particles of calcium carbonate,
which could be easily contrasted with the larger burkeite crystals formed during regime I. The large crystals formed in regime I were much more numerous than those from the base-case experiments, but they were of similar morphology. The agglomerates from regime II resembled those seen during the base-case experiments, suggesting that they formed by the same mechanism as in the base-case experiments. Calcium carbonate crystals in the samples were obviously diminishing as the runs progressed until they totally disappeared after the intermittent primary nucleation period. By using potentiometric titration against EDTA solutions, the calcium concentrations in the large crystals were determined to be in the range of 300-500 ppm in both cases (with 50 and 100 ppm of calcium carbonate powder added), which is an order of magnitude higher than the calcium concentration would have been in the solutions (20 and 40 ppm, respectively) if all of the added calcium carbonate had dissolved. In other words, the burkeite crystals appeared to be functioning as calcium scavengers during the experiments. The change in nucleation mechanism and the unusually high concentration of calcium in the large crystals imply a close tie between the presence of calcium and formation of the large burkeite crystals. The impact of calcium in these experiments also provided a clue to the appearance of large crystals in the base-case experiments. Calcium can come from various sources; for example, calcium was listed as one of the impurities in the sodium carbonate reagents. The actual amounts of calcium were analyzed by titration with EDTA solutions at pH ) 10 and were found to range as high as 40-50 ppm in two commercial sources (Fisher Scientific and JT Baker) of sodium carbonate used in the current work. Because of these observations, we determined the calcium content of crystals from the base-case experiments and found the large crystals to contain calcium concentrations in the range of 150-300 ppm. Such values are about half those found for crystals produced in experiments involving the addition of calcium carbonate powder, and they are significantly higher than the initial concentration of calcium in the solutions. These results suggest that the base-case experiments were also influenced by calcium. To test this possibility, experiments were performed in which the calcium concentration was controlled, in contrast to the basecase experiments. Two-Stage Crystallization. To remove the insoluble calcium carbonate from the solution prepared in the base-case experiments, the solutions were passed through a filter with a pore size of 0.45 µm. There was no change in results when the base-case experiments were repeated using the filtered solutions; that is, formation of the large crystals in regime I continued to occur. These results suggest that the size of the insoluble calcium carbonate might be smaller than the pore size of the filter or that the solubility of calcium carbonate is higher in burkeite solutions than water. The latter could result from salting-in by the sodium carbonate and sodium sulfate in the solution, but no data were found to support this possibility. Another method of reducing calcium concentration in the crystallizing solutions is to use the results from the previous experiments showing that calcium was incorporated into the large crystals, thereby making the solution concentration of calcium at the end of regime II very low (less than the detection limit of the calcium-
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as many transition-metal ions. It has been applied in heat exchangers to control fouling caused by calcium and magnesium compounds. When the pH is over 11, EDTA exists mostly in its fully dissociated form (designated by Y4-), and this happens to be the case in the solutions of sodium carbonate and sodium sulfate studied here. At 20 °C, the equilibrium constant of the reaction
Ca2+ + Y4- T CaY2-
Figure 8. Typical photograph taken through an optical microscope of a sample from the second-stage crystallization. Note the absence of large crystals that were observed in the base-case experiments.
ion-selective electrode). This led to the implementation of two-stage crystallization in which the first stage was expected to deplete calcium brought in as a contaminant by the reagents so that the second-stage crystallization could start without any calcium present in the solution. The experiments were carried out as follows: The first stage was conducted in the same manner as the basecase experiments, except that the ratio of sodium carbonate and sodium sulfate was slightly reduced. This was done so that the solutions entering the second stage had approximately the same ratio of carbonate to sulfate as in the base-case experiments. (Recall that burkeite crystallization takes more sulfate than carbonate from solution.) The evaporation was stopped just after regime II occurred, and the suspension was cooled to 90-100 °C and filtered to remove crystals. The suspension should not be cooled further lest burkeite crystals dissolve. After the first stage, the residual crystals left in the crystallizer were washed away, and the crystallizer was rinsed thoroughly with distilled water. During the experiments, all contacting surfaces were rinsed with distilled water to avoid contamination. The filtered solution was reintroduced into the crystallizer to conduct the second-stage crystallization in which the operating conditions were the same as in the base-case experiments. As expected, photographs of crystal samples (Figure 8) showed that the second-stage crystallization produced only agglomerates of the tiny crystals that are typical of regime II nucleation. The complete absence of large crystals found in the base-case experiments and in experiments involving the addition of calcium carbonate confirmed the close relationship between the calcium in the solution and the formation of large burkeite crystals. Effects of EDTA. Another way to reduce the concentration of calcium ions in the solution is by the addition of EDTA (ethylenediamine tetraacetic sodium salt, dihydrate crystals, molecular weight ) 372.24). EDTA is a common chelating agent because it forms metal-ligand complexes with Ca2+ and Mg2+, as well
(1)
is 5.0 × 1010. When an excess amount of EDTA is added, the calcium ion concentration becomes negligibly small, and even calcium carbonate solids will dissolve, as has been observed experimentally. When 400 ppm EDTA was added to crystallizing burkeite solutions, no large crystals formed in regime I. Upon addition of 100 ppm EDTA, a small number of large burkeite crystals formed, but the number was much lower than in the base-case experiments. The threshold to inhibit fully the formation of large burkeite crystals in regime I seemed to be between 200 and 400 ppm. Considering that other transition-metal ions could exist in the solution and that EDTA could attack stainless steel at the operating temperature, the amount of EDTA required to sequester calcium was judged to be reasonable. This set of experiments further confirmed the close tie between calcium and the large crystals of burkeite. Effects of Other Inorganic Compounds. Inorganic compounds such as CaC2O4 (calcium oxalate), Ca3(PO4)2, MgCO3, BaCO3, PbCO3, K2SO4, CuSO4, and TiO2 powder were added individually at 100 ppm levels to determine their effects on burkeite nucleation or, to be more specific, on the formation of large crystals. Some of these compounds are natural contaminants in the reagents used, and others were selected because they were readily available. At the concentrations investigated, only the calcium compounds had any effect on the crystallization process. When added to the crystallizing solutions, calcium oxalate and calcium phosphate had effects similar to those of calcium carbonate. In these cases, the large crystals were much more numerous than in the basecase experiments. Because carbonate ions were abundant in the solution and because the solubility limits of these calcium salts are all very low, it was unclear whether calcium oxalate and calcium phosphate had been partially converted to calcium carbonate and whether it was the calcium carbonate thus produced that was responsible for the appearance of large crystals. Calcium has the potential to form numerous compounds with sodium, carbonate, and sulfate ions and water molecules present in this system. These compounds include polymorphs of calcium carbonate, calcium sulfate, the double salts of sodium carbonate and calcium carbonate, and the double salts of sodium sulfate and calcium sulfate, plus the possible hydrates of these salts. It is still unclear whether the solid forms of calcium compounds or calcium ions in the solution were responsible for the observed behavior. If it was the solid form of a calcium compound that caused the formation of large burkeite crystals, the mechanism could be burkeite nucleation on a solid template. This mechanism is based on the premise that lattice matching can reduce the free energy change in the formation of a stable nucleus. Several solid compounds of calcium
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were tested to determine whether burkeite could grow on the solid surfaces. The possibility that calcium carbonate could act as a nucleation template is reasonable because more large crystals were formed when calcium carbonate powder was added to the solutions. The calcium carbonate powder used was calcite, which has a crystal lattice that does not match that of burkeite. Even so, an experiment was performed in which a large calcite crystal (approximately 4 cm × 3 cm × 2 cm) was suspended in the solution during evaporation. A thorough examination of the calcite crystal at the end of the run failed to find any burkeite crystals attached to the calcite. Therefore, the possibility that calcite served as the nucleation template was rejected. The FBRM results obtained when this large calcite crystal was suspended in the solution resembled the results in the base-case experiments rather than those when calcium carbonate powder was added. Apparently, the relatively small surface area of the mineral made a difference between this experiment and the ones when calcium carbonate (calcite) powder was added. This experiment suggests, but cannot prove, that dissolved calcium is responsible for the appearance of large burkeite crystals. Metastable Limits. The effects of calcium compounds have been recognized as the principal factor in the formation of large burkeite crystals. Aside from the hypothesis of nucleation on templates proposed above, another possible mechanism is that calcium ions in solution are responsible for the appearance of large burkeite crystals. According to the published work on nucleation kinetics,17-19 it is known that impurity ions can either inhibit or accelerate nucleation, although there are more examples of the former. However, if solids affect nucleation (as might be the case with calcite addition), the effects are usually accelerative; otherwise, if the effect on nucleation is inhibitive, only a small portion of solution will be affected because the interfacial area between the solid particles and the solution is generally small, and the vast majority of the solution will not be affected by the solid particles. The only exception is solid particles having very large interfacial areas such as colloidal dispersions (e.g., gelatin), in which the colloidal particles can inhibit nucleation in aqueous solutions. Whether impurities accelerate or inhibit nucleation can usually be determined from their effect on the metastable limit. If the comparison between the metastable limits for contaminated and uncontaminated solutions shows an inhibitive effect by calcium, it could support the hypothesis that calcium ions inhibit nucleation of burkeite. A series of experiments was performed in which the polythermal method was used to determine the metastable limits of these solutions. In these experiments, the valve controlling the vapor flow was kept closed and the system temperature was increased at a constant rate until nucleation was detected using the FBRM probe. Supersaturation was generated by increasing the system temperature (0.02-6 °C/min) rather than by evaporation. The former provides an easy way to control the rates at which supersaturation is generated. Two types of solutions were used in these experiments, the base-case solution (contaminated by calcium) and solutions remaining after a first-stage crystallization (uncontaminated). The latter was considered to show intrinsic nucleation kinetics because the calcium
Figure 9. FBRM results (count rates of chords in the 1-10 µm range) showing the nucleation starting temperature Tonset during measurements of the metastable limits. Table 1. Metastable Limits [Nucleation Temperature (°C)] with and without Calciuma solution content (%)
base caseb
100 ppm EDTAc
400 ppm EDTAc
second stagec
30 28.3
109 ( 1 136 ( 1.5
96.7 ( 1 109 ( 1
96.5 ( 1 108.7 ( 1
109.8 ( 1.5
a
Heating rate ≈ 0.5 °C/min. b 3-5 ppm Ca2+. c
