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Ind. Eng. Chem. Res. 2006, 45, 642-647
Effect of Calcium and Magnesium Ions on Calcium Oxalate Formation in Sugar Solutions William O. S. Doherty* Sugar Research and InnoVation, Queensland UniVersity of Technology, Brisbane 4001, Australia
During the concentration of sugarcane juice, there is continuous deposition of nonsugar impurities on the surface of evaporator units. Because the scale deposit has a low thermal conductivity, its accumulation impairs heat transfer and eventually renders the process uneconomical owing to reduced throughput. Calcium oxalate is the main intractable scale formed in factories that process sugarcane. It is not removed by conventional chemical cleaning methods. This paper describes studies on the apparent solubility product of calcium oxalate in the presence of sugar and calcium and magnesium ions at different temperatures and pH’s. A solubility product model was obtained that states that the apparent solubility product of calcium oxalate decreases with increasing temperature and sugar concentration, but increases with increasing pH. The model was able to predict calcium oxalate solubility changes in raw sugarcane factory processes, indicating that the apparent solubility product of calcium oxalate declines rapidly through the evaporator set. The results also indicated that calcium oxalate solubility is increased in the presence of calcium and, in particular, magnesium ions. Introduction In sugar factories, clarified juice containing ∼11.5 mass % sucrose is passed through a number of evaporator units (effects) where most of the water is removed. Within each effect, the juice is boiled in stainless steel tubes (calandria), which are heated by circulating steam. Usually, the juice passes through five effects in series after which it contains ∼65 mass % sucrose and is then passed onto pans for crystallization. During the concentration of evaporator supply juice (ESJ) to syrup, there is continuous deposition of inorganic and organic nonsugar compounds on the surfaces of evaporator calandria tubes, resulting in the lowering of the heat-transfer coefficients (HTCs). The most important operational requirement of any evaporator set is that it can evaporate water from juice at a rate commensurate with the throughput rate of the factory. After cleaning, evaporator sets typically have a sufficiently high average HTC to meet the rate of production. However, as the heating surfaces become covered with scale, the HTC declines, and if the set becomes heavily fouled, it will not meet production rates. If the evaporator set becomes rate limiting because of the effects of the scale, techniques such as the reduction of maceration water on the milling train, the lowering of the syrup brix, and the reduction of filter wash water help to maintain the production rate. However, these techniques can only sustain the evaporator set for a limited period, and once they fail, the only option is to stop the factory so that the evaporators can be cleaned. Studies by Doherty et al.1 have shown that oxalic acid is introduced to the factory not only via the sugarcane plant, but also during processing. This explains why calcium oxalate is a major component of the scale deposits on the effect tubes.2,3 It is well-known that low-molecular-weight polymers, e.g., poly(acrylic acid) and poly(maleic acid), inhibit or retard the crystallization of calcium carbonate, calcium sulfate, and calcium oxalate. Consequently, these polymers are used in boilers, heat exchangers, cooling towers, and desalination plants to control scale formation. Although these polymers are used in the Australian sugarcane industry, their effectiveness is * Tel.: +61 7 4952 7635. Fax: +61 7 4952 7699. E-mail:
[email protected].
limited.4 Thus, it has become necessary to obtain a better understanding of factors that control calcium oxalate formation so that effective ways of inhibiting calcium formation can be found. Recently, there has been much activity in studying the kinetics, thermodynamics, and mechanisms of composite fouling of calcium oxalate and amorphous silica.5-7 The results for the kinetics of calcium oxalate monohydrate (COM) in water showed that increasing the initial COM concentration at 60 °C led to a corresponding increase in the rate constant (k) of calcium oxalate formation and a decrease in the order of the reaction n. The n values were found to be in the range of 2.4-4.2, which is similar to those reported by Nielsen8 but higher than those of Nancollas.9 Yu et al.6 showed that, as the initial COM concentration increased, the solubility concentration of COM exhibited a maximum before leveling off. Further work carried out by Yu et al.7 showed that higher COM precipitation rates occurred in sucrose solutions than in water, with higher k and lower n values. The extent of the increase was dependent on the sucrose concentration. To obtain complimentary information on calcium oxalate formation in sugar mill evaporators, this work investigated calcium oxalate formation in sugar solutions in the presence and absence of calcium and magnesium ions. Calcium and magnesium ions are present at concentrations of up to 400 and up to 300 mg‚kg-1, respectively, in clarified cane juice prior to evaporation in the evaporator units. Experimental Section Materials. Concentrated stock and dilute solutions of calcium chloride, potassium oxalate, sodium hydroxide, and hydrochloric acid were prepared with analytical reagent chemicals and Millipore deionized water. The sucrose used in the study was analytical grade and did not contain any measurable quantities of oxalic acid or other organic acids. The concentrations of calcium and magnesium in the analytical sucrose were 2 and 1 mg‚kg-1, respectively. Precipitation of Calcium Oxalate. Methodology. For each precipitation study, potassium oxalate solution (1.63 mM‚L-1) was added to calcium chloride solution (1.40 mM‚L-1) or to a
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solution containing calcium chloride and sugar. Prior to the addition of potassium oxalate, the pH of the calcium chloride/ calcium chloride-sugar solution was adjusted to the desired pH by the addition of sodium hydroxide or hydrochloric acid, and the solution was maintained at the appropriate temperature in a water bath. The mixture was stirred for 1 min and then left in a water bath for an additional 30 min in order for equilibrium to be established. Recent work by Yu et al.5 has established that equilibrium was attained within 30 min, as similar results were obtained even after 4 h of standing time. The mixture was then filtered through a filtration apparatus that was thermostated at the same temperature as the mixture in order to prevent precipitation due to a decrease in temperature. To ensure that no calcium oxalate precipitated in the filtrate, 0.5 mL of concentrated hydrochloric acid was added to the filtration apparatus’s receiver vessel prior to filtration. The residue on the membrane filter was later recovered after thorough washing with water to remove sugar. Six groups of tests were conducted. Each test group used varying sugar concentrations in nine steps from 0 to 40.7 mass % for a constant temperature and pH value. Temperatures of 60 and 80 °C and pH values of 5.0, 7.5, and 8.0 were used in the six test groups. Analysis of Filtrate. The oxalate ion concentration of the filtrate was determined using the complexometric method of Burriel-Marti et al.10 There is a decrease in optical density when a ferric salicylate complex is treated with oxalate ions. The decrease in optical density with increasing oxalate ion concentration obeys Beer’s law. The calcium concentration of the filtrate was determined using an atomic absorption spectrophotometer (AAS) coupled to a graphite furnace detector. It was necessary to use a very sensitive AAS system as samples to be processed were diluted to sugar concentrations less than 1 mass %. This was to reduce the interference of the sugar organic matrix on the calcium determination. The experimental error involved in the precipitation studies was less than 6.5%. Analysis of Calcium Oxalate Precipitate. The crystallographic forms of the precipitates were determined by X-ray powder diffraction (XRD) using a Rigaku diffraction camera with an X-ray generator with Cu KR radiation of wavelength 1.5418 Å. The diffraction patterns were indexed according to parameters obtained from the XRD file. The crystal morphologies were examined using a scanning electron microscope (SEM) operating at a gun voltage of 20 kV. Precipitation of Calcium Oxalate in the Presence of Excess Calcium Ions. The formation of calcium oxalate at 60 and 80 °C and at pH 5 and 7.5 in the presence of calcium ions (150625 mg‚kg-1 based on dry solids) was studied using the procedure outlined in the section on the precipitation of calcium oxalate. The experiments were conducted at two sugar concentrations, namely, 13.9 and 22.2 mass %. The filtrates were analyzed as described in the section on the precipitation of calcium oxalate. Precipitation of Calcium Oxalate in the Presence of Calcium and Magnesium Ions. Calcium oxalate formation was evaluated in two sets of synthetic juice solutions at 60 and 80 °C and at pH of 5 and 7.5 by the procedure described in the section on the precipitation of calcium oxalate. Both sets of synthetic juice solutions contained calcium ions (125-400 mg‚kg-1 based on dry solids) and magnesium ions (62.5-320
mg‚kg-1 based on dry solids). One set of solution contained 11.9 mass % sugar, and the second set contained 15.0 mass % sugar. Results and Discussion Calcium Oxalate Formation in Simple Sugar Systems. (i) Solubility Product. One of the fundamental relationships determining the precipitation of slightly soluble compounds is the solubility product. CaC2O4 dissociates in water to according to eq 1
CaC2O4 T Ca2+ + C2O42-
(1)
Its solubility product is defined as
Ks ) (aCa) × (aC2O4)
(2)
where the solubility product Ks is a constant, and aCa and aC2O4 are the activities of calcium and oxalate ions, respectively. In very dilute solutions, the activity coefficient tends to unity, and the solubility product becomes a product of the concentration of the ions, thus
Ks ) (cCa) × (cC2O4)
(3)
where cCa and cC2O4 are the molar concentrations of calcium and oxalate ions, respectively. Precipitation of calcium oxalate can commence once the product of the actual concentrations of calcium and oxalate ions exceeds Ks, and precipitation can continue until that product is equal to Ks. In this work, because the experiments were conducted in nonideal conditions (i.e., at high concentrations of the latticeforming ions and in the presence of sucrose), the concentrations of calcium and oxalate ions were higher than their corresponding equilibrium activities due to electrical interactions. It was not possible to calculate the activity coefficients of the ions and to identify and quantify the main complexes in the sugar solutions as the effect of sugar on activity coefficient is not known. As shown by X-ray powder diffraction studies in this study, two of the three major phases of calcium oxalate, i.e., the mono-, di-, and trihydrates, are in some cases formed at any one time. As such, the values reported in the text as the solubility products of calcium oxalate should be regarded as “apparent solubility products of calcium oxalate monohydrate”, Ksa. Ksa is not a constant and exceeds the activity solubility product Ks to an increasing extent as the activity coefficients of the ions become smaller with increasing concentration. (ii) Calcium Oxalate Solubility. The Ksa values of calcium oxalate in water and in sugar solution were determined at pH values of 5, 7, and 8.5 and at 60 and 80 °C using eq 3. Individual plots of each of the six groups of tests conducted at 60 °C are shown in Figure 1a-c. The plots for the tests carried out at 80 °C were similar. In each plot, the apparent solubility product is plotted against the sugar concentration, and an exponential trend curve is added. There does appear to be a log-linear trend, with a sharp peak at ∼7.3 mass %. Multivariable regression analysis procedures were carried out on the full set of data obtained and again on a reduced data set including only those data for sugar concentrations above 9.6 mass %. The latter are considered more relevant to the raw sugar factory conditions, where the sugar concentration of the main
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Figure 2. Fit of the log-linear regression of the actual data; all data included.
Figure 3. Fit of the log-linear regression to the actual data; only data with a sucrose concentration of 9.6 mass % and above are included. Table 1. Conditions for Juice Processing in a Raw Sugar Factory
Figure 1. (a-c) Plots of apparent solubility product for test groups.
stream is not below 11.5 mass %. The second analysis avoided the complications of an apparent peak in the solubility product at around 2.5-7.3 mass % sugar concentration. Combinations of linear-linear, log-log, and log-linear regressions were tried. The better regressions, with the smallest standard errors and highest residual R2 values, were obtained with the log-linear form. The regression expression for the full set of data was
log Ksa ) -8.143 - 0.003T - 0.011Suc + 0.060pH
(4)
where T is temperature in °C and Suc is sugar concentration (mass %). The standard error for Ksa is 0.137 × 10-8; R2 ) 0.73; degree of freedom, F ) 49. The coefficients of all terms were significant. The fit of all measured data to the regression predictions is shown in Figure 2. The regression expression for the partial set data is given by
log Ksa ) -8.107 - 0.002T - 0.010Suc + 0.049pH
(5)
where the standard error of Ksa is 0.115 × 10-8, R2 ) 0.74, F ) 31. The coefficients of all terms were significant. Put simply, eq 5 states that the Ksa value of calcium oxalate decreases with increasing temperature and sugar concentration, but increases with increasing pH. It also indicates that small changes in pH have a greater effect on Ksa than small changes in temperature and sugar concentration. The second degree of
processing stage
temperature (°C)
sugar (mass %)
pH
predicted Ksa value
warm juice clarified juice effect 1 effect 2 effect 3 effect 4 final effect
75 96 115 109 104 87 57
12 12 17 25 30 46 65
7.8 7.4 7.3 7.3 7.2 7.2 7.1
8.59 × 10-9 7.51 × 10-9 6.06 × 10-9 5.15 × 10-9 4.64 × 10-9 3.43 × 10-9 2.49 × 10-9
dissociation of oxalic acid, which lies in the pH range 4-7, might explain the influence of solution pH. The regression expression of eq 5 can be compared to the van’t Hoff equation where there exists a linear relationship between Ks and the inverse temperature (in kelvin). The fit of the measured data to the regression predictions for the partial set is shown in Figure 3. The predictive capabilities of the regression are good at lower solubility product values but poor at high values. This suggests that another factor might be influencing solubility factors at higher values (i.e., lower sugar concentration). It is of interest to examine how the regression expression for Ksa of calcium oxalate predicts the change in Ksa during the processing of limed raw sugar juice to evaporator syrup. To do this, a typical set of information on raw juice conditions is required for use with the chosen regression expression. A set of juice conditions from an Australian sugar factory is provided in Table 1. The final column in Table 1 shows Ksa of calcium oxalate as determined by the log-linear regression expression for the partial data set. The results indicate that Ksa of calcium oxalate declines rapidly through the evaporator set. The solubility of calcium oxalate in warm juice is well over 3 times that in the final effect.
Ind. Eng. Chem. Res., Vol. 45, No. 2, 2006 645 Table 2. X-ray Powder Diffraction Data (d Values) of Calcium Oxalate under Different Conditions 7.3 mass % sugar, pH 8.5, temp 60 °C
7.3 mass % sugar, pH 5, temp 60 °C
34.41 mass % sugar, pH 5, temp 60 °C
5.89 5.76 5.53a 5.41 5.29a 5.01a 4.73 4.00 3.93 3.88 3.63 2.96 2.35
6.10 5.92 5.80 5.42 5.29a 5.21 3.93 3.89 3.64 2.87a 3.63 2.96 2.36
6.27b 6.12 5.98 5.85 5.36 3.92b 3.67b 2.98 2.62 2.36 2.27 1.89b
a
Calcium oxalate trihydrate. b Calcium oxalate dihydrate.
The model (eq 5) compares well with the actual situation in the factory examined, in which increasing amounts of calcium oxalate were progressively formed through the evaporator set.2 This is in contrast to the model developed by Walford and Walthew,11 which predicts that the lowest solubility condition appears in stages 1-3. There is a suspicion that the techniques (e.g., mixing sugar solutions with excess of calcium oxalate crystals) used by Walford and Walthew11 to establish the solubilities might not be very relevant to the actual process conditions, in which there is precipitation of more than one form of calcium oxalate hydrate. The degree of hydration in the crystals might be dependent on the water activity in the solution and, hence, on the concentration of sugar present at the time. Analysis of Calcium Oxalate Precipitates. (i) X-ray Powder Diffraction Analysis. The X-ray powder diffraction data (as d values) of some of the precipitates are presented in Table 2. Interpretation of the d values indicated that some of the precipitates were made up of mixtures of calcium oxalate trihydrate, calcium oxalate monohydrate, and calcium oxalate dihydrate. Higher proportions of calcium oxalate dihydrate were formed for precipitates obtained at 34 and 40.7 mass % sugar concentrations. The higher solubility product values obtained in the sugar range of 2.5-7.3 mass % (i.e., the peak) is due to a higher proportion of calcium oxalate trihydrate formed. Calcium oxalate trihydrate is more soluble than the two other calcium oxalate phases. (ii) Scanning Electron Microscopy. Figure 4a-c shows scanning electron micrographs of calcium oxalate crystals prepared under various conditions. In each of the micrographs, the crystals are not well-developed, and there are no welldeveloped faces. This confirms the X-ray powder diffraction spectra, which consisted of broad and diffused peaks. As shown in Figure 4a, the calcium oxalate precipitated in 7.3 mass % sugar at pH 5 and 60 °C consists of monoclinic (typical of calcium oxalate monohydrate) and truncated crystals and tiny crystal clusters. The average crystal length varies from less than 3 to 12 µm. When calcium oxalate was prepared under similar conditions but at 80 °C, the crystallites were no longer present, and there appeared to be an increase in the average crystal length (Figure 4b). At the higher temperature, aggregation and coagulation of the crystals occurred. Figure 4c is the micrograph of calcium oxalate prepared in the presence of 34 mass % sugar. Here, serrated crystals are the main morphology with the sizes typically smaller than at lower sucrose concentrations. The sizes vary from about 3 to 10 µm.
Figure 4. Scanning electron micrographs of calcium oxalate (a) formed in 7.3 mass % sucrose at pH 5 at 60 °C, (b) formed in 7.3 mass % sucrose at pH 5 at 80 °C, and (c) formed in 34 mass % sucrose at pH 5 at 60 °C. Table 3. Solubility of Calcium Oxalate in Sugar Solutions Containing Calcium Ions temp (°C) 60 60 60 60 60 60 60 60 80 80 80 80 80 80 80 80
pH
sugar (mass %)
concentration of calcium ions (mg‚kg-1 based on dry solid)
Ksa
5.0 5.0 5.0 5.0 7.5 7.5 7.5 7.5 5.0 5.0 5.0 5.0 7.5 7.5 7.5 7.5
13.9 22.2 13.9 22.2 13.9 22.2 13.9 22.2 13.9 22.2 13.9 22.2 13.9 22.2 13.9 22.2
150 250 375 625 150 250 375 625 150 250 375 625 150 250 375 625
2.03 × 10-8 2.10 × 10-8 3.35 × 10-8 1.28 × 10-7 1.65 × 10-8 3.11 × 10-8 3.65 × 10-8 4.80 × 10-8 9.16 × 10-9 2.29 × 10-8 1.85 × 10-8 1.44 × 10-7 5.00 × 10-9 1.88 × 10-8 1.75 × 10-8 6.90 × 10-8
Effect of Excess Calcium Ions on Calcium Oxalate Formation. The results of the effect of calcium ions on the solubility of calcium oxalate are presented in Table 3. The results are reported as apparent solubility products for convenience. What is actually being carried out in the investigation is an examination of the solubility of calcium oxalate in excess calcium.
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Table 4. Concentrations of Oxalate Ions in Synthetic Juice Solutions concentration temp (°C)
pH
sugar (mass %)
Ca (mg‚kg-1)
Mg (mg‚kg-1)
oxalate (mol‚L-1)
60 60 60 60 60 60 60 60 60 60 60 60 60 60 60 60 80 80 80 80 80 80 80 80 80 80 80 80 80 80 80 80
5.0 5.0 5.0 5.0 5.0 5.0 5.0 5.0 7.5 7.5 7.5 7.5 7.5 7.5 7.5 7.5 5.0 5.0 5.0 5.0 5.0 5.0 5.0 5.0 7.5 7.5 7.5 7.5 7.5 7.5 7.5 7.5
11.9 11.9 11.9 11.9 15.0 15.0 15.0 15.0 11.9 11.9 11.9 11.9 11.9 15.0 15.0 15.0 11.9 11.9 11.9 11.9 15.0 15.0 15.0 15.0 11.9 11.9 11.9 11.9 15.0 15.0 15.0 15.0
125 312.5 125 312.5 160 400 160 400 125 312.5 125 312.5 160 400 160 400 125 312.5 125 312.5 160 400 160 400 125 312.5 125 312.5 160 400 160 400
62.5 62.5 250 250 80 80 320 320 62.5 62.5 250 250 80 80 320 320 62.5 62.5 250 250 80 80 320 320 62.5 62.5 250 250 80 80 320 320
5.14 × 10-5 2.69 × 10-5 2.21 × 10-4 8.63 × 10-5 8.11 × 10-5 3.22 × 10-5 2.31 × 10-4 1.12 × 10-4 2.17 × 10-5 3.91 × 10-5 1.95 × 10-4 1.18 × 10-4 9.06 × 10-5 3.74 × 10-5 2.91 × 10-4 1.99 × 10-4 7.41 × 10-5 4.61 × 10-5 1.81 × 10-4 1.41 × 10-4 5.66 × 10-5 4.44 × 10-5 1.74 × 10-4 9.68 × 10-5 1.28 × 10-4 8.80 × 10-5 2.35 × 10-4 1.09 × 10-4 1.63 × 10-4 5.49 × 10-5 2.29 × 10-4 1.58 × 10-4
The regression expression for the data of Table 3 (all variables are significant as shown by the t-statistic values) is
log(Ksa × 109) ) 1.507 - 0.009T - 0.050pH + 0.016Suc + 0.002cCa (6) R ) 0.93, F ) 17.0 2
t-statistics on variables: T (°C) ) -2.1, pH ) -1.5, Suc ) 1.66, cCa ) 5.83 where cCa is the calcium ion concentration (mg‚kg-1 based on dry solids). The results of these regressions indicate that the solubility of calcium oxalate increases with increasing calcium ion concentration. If these results are related to factory conditions, then the current practice of using lime/lime saccharate for juice clarification should reduce the precipitation of calcium oxalate in the clarifier through an increased calcium level. However, this would inevitably result in scaling of the evaporators by calcium oxalate. Effect of Calcium and Magnesium Ions on Calcium Oxalate Formation. Whereas the previous section deals with simple sugar systems containing calcium ions, this section deals with mixtures of calcium and magnesium ions in sugar systems. Two sets of solutions were studied at 60 and 80 °C and at pH of 5 and 7.5. X-ray powder diffraction identified calcium oxalate monohydrate as the predominant phase obtained from these solutions. Traces of calcium oxalate trihydrate crystals were also identified. The results for the concentration of oxalate ions that remained after precipitation of calcium oxalate in the various synthetic juice solutions are presented in Table 4. The results are reported
in terms of the oxalate concentration instead of the solubility products because of the difficulty in calculating the solubility products from such complex systems. Regression was carried out on the data of Table 4. The variable sugar concentration was found not to be significant in its influence on the measured oxalate concentration. The regression expression for the other variables, which are significant (as shown in the t-statistics), was found to be
log(oxalate × 106) ) 1.209 + 0.0053T + 0.043pH 0.001cCa + 0.002cMg (7) R2 ) 0.89, F )25.8 t-statistics on variables: T (°C) ) 1.99, pH ) 2.00, cCa ) -4.37, cMg ) 8.95 where cMg is the concentration of magnesium ions (mg‚kg-1 based on dry solids). It is seen that magnesium ions inhibited the precipitation of calcium oxalate even in the presence of high concentrations of calcium ions. The inhibitory effect of magnesium ions on calcium phosphate formation has been reported by Hidi12 and Feestra et al.13 In the regression of the data in Table 4, the concentration of soluble oxalate doubled (on the average from 118 to 235 mM‚L-1) in the presence of 138 mg‚kg-1 magnesium ions based on dry solids. As a consequence of this finding, the addition of soluble magnesium ions to clarified juice (perhaps in the form of magnesium chloride) might well be a way of inhibiting the precipitation of calcium oxalate in the evaporator set. However, it might be more effective to add a formulation consisting of a mixture of magnesium ions and a polymeric scale inhibitor. Conclusions The formation of calcium oxalate in various sugar solutions was extensively studied. Using simple sugar systems, it was found that, in dilute solutions, the apparent solubility product increases slightly in sucrose concentration to 7.3 mass % before markedly decreasing. A regression expression for the apparent solubility product of calcium oxalate with a sugar concentration of 9.6 mass % and above was obtained because this is the concentration range employed in sugar factories. The expression states that the apparent solubility product of calcium oxalate decreases with increasing temperature and sugar concentration, but increases with increasing pH. The expression fitted well to the actual observed situation in a raw sugar factory. The solubility studies indicated that the deposition of calcium oxalate in evaporators is due to solubility effects. The residual oxalate ions not removed during juice clarification will precipitate later in the evaporators as a result of solubility effects. In addition, the oxalate ions formed during processing will also form in the evaporators. In the case of solubility studies involving inorganic ions, the effects of calcium and magnesium ions on calcium oxalate formation were demonstrated. The results showed the detrimental effect of using lime/lime saccharate for juice clarification as calcium ions increase the solubility of calcium oxalate. The results also indicated that magnesium ions enhance calcium oxalate solubility more so than calcium ions. The addition of soluble magnesium salts or the addition of a mixture of magnesium ions and a scale inhibitor to clarified juice could be a feasible option for minimizing calcium oxalate precipitation
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in the evaporators. It is considered that the latter method should be examined to determine its beneficial effects on scale reduction. Acknowledgment The financial support of Sugar Research Development Corporation and Australian sugar mills are gratefully appreciated. Literature Cited (1) Doherty, W. O. S.; Galea, C. F.; Falzon, K. L. Oxalic Acid LeVels in First Expressed Juice through to Syrup; SRI Technical Report 2295; Sugar Research Institute: MacKay, Australia, 1999; p 5. (2) Crees, O. L.; Cuff, C.; Doherty, W. O. S.; Senogles, E. Examination of evaporator scales from the far Northern regions of the sugar industry. Proc. Aus. Soc. Sugar Cane Technol. 1992, 14, 238-245. (3) Doherty, W. O. S.; Senogles, E.; Crees, O. L. The preparation of calcium oxalate dihydrate crystals. Cryst. Res. Technol. 1994, 29 (4), 517524. (4) Doherty, W. O. S.; Senogles, E.; Crees, O. L. Polymeric additives: Effects on crystallization of calcium oxalate scales. Cryst. Res. Technol. 1995, 31 (3) 281-286. (5) Yu, H.; Sheikholeslami, R.; Doherty, W. O. S. Mechanisms, thermodynamics and kinetics of composite fouling of calcium oxalate and amorphous silica in sugar mill evaporators. Chem. Eng. Sci. 2002, 57 (11), 1969-1978.
(6) Yu, H.; Sheikholeslami, R.; Doherty, W. O. S. Composite fouling characteristics of calcium oxalate monohydrate and amorphous silica by a novel approach simulating successive effects of a sugar mill evaporator. Ind. Chem. Eng. Sci. 2002, 41, 3379-3388. (7) Yu, H.; Sheikholeslami, R.; Doherty, W. O. S. Composite fouling of calcium oxalate and amorphous silica in sugar solutions. Ind. Chem. Eng. Sci. 2003, 42, 904-910. (8) Nielsen, A. E. The kinetics of calcium oxalate precipitation. Acta Chem. Scand. 1960, 14, 1654-1659. (9) Nancollas, G. H. The mechanisms of precipitation of biological minerals. The phosphates, oxalates and carbonates of calcium. Croatica Chem. Acta 1983, 56, 741-752. (10) Burriel-Marti, F.; Ramirez-Munoz, J.; Fernandez-Caldas, E. Determination of oxalate ion and calcium ion by indirect colorimetry. Anal. Chem. 1953, 25, 573-585. (11) Walford, S. N.; Walthew, D. C. Preliminary model of oxalate formation in evaporator scale. Proc. S. Afr. Sug. Technol. Assoc. 1996, 70, 231. (12) Hidi, P. The effect of magnesium on clarification and the phosphate content of raw sugars. Aust. J. Appl. Sci. 1964, 15, 35-40. (13) Feenstra, K.; De-Bruyn, P. L. The influence of small amounts of magnesium ions on the formation of calcium phosphate in moderately supersaturated solutions. J. Colloid Interface Sci. 1981, 83, 583-588.
ReceiVed for reView August 3, 2005 ReVised manuscript receiVed November 3, 2005 Accepted November 4, 2005 IE0509037
