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J. Dairy Sci. 87:863-867
© American Dairy Science Association, 2004.

Factors Affecting Solubility of Calcium Lactate in Aqueous Solutions

N. Kubantseva1, R. W. Hartel1 and P. A. Swearingen2

1 Department of Food Science, University of Wisconsin, 1605 Linden Drive, Madison 53706
2 Land O’Lakes, Inc., Arden Hills, MN 55112

Corresponding author: R. W. Hartel; e-mail: rwhartel{at}wisc.edu.


   ABSTRACT
TOP
 ABSTRACT
INTRODUCTION
 MATERIALS AND METHODS
 RESULTS AND DISCUSSION
 CONCLUSIONS
 ACKNOWLEDGEMENTS
 REFERENCES
 
Calcium lactate (CaL2) crystal formation on the surface of cheese continues to be a widespread problem for the cheese industry despite decades of research. To prevent those crystals from forming, it is necessary to keep the concentration of CaL2 below saturation level. The limited data available on the solubility of CaL2 at conditions appropriate for the storage of cheese are often conflicting. In this work, the solubility of L(+)-CaL2 in water was evaluated at 4, 10, and 24°C, and the effects of salt and pH at those temperatures were investigated. The effects of additional calcium and lactate ions on solubility also were studied. The results suggested that temperature and the concentration of lactate ions are the main factors influencing the solubility of CaL2, with the other parameters having limited effect.

Key Words: solubility • calcium lactate • crystallization • Cheddar cheese

Abbreviation key: CaL2 = calcium lactate, NaL = sodium lactate


   INTRODUCTION
TOP
 ABSTRACT
 INTRODUCTION
MATERIALS AND METHODS
 RESULTS AND DISCUSSION
 CONCLUSIONS
 ACKNOWLEDGEMENTS
 REFERENCES
 
For the past several decades, numerous studies have been conducted to identify why crystal appears on cheese and to find ways to prevent them. The previous studies mainly focused on the effects of milk composition and the processing conditions for Cheddar cheese manufacturing on crystal formation. The results of these studies have shown that factors such as a high lactose concentration in milk (Pearce et al., 1973), an increase in numbers of nonstarter bacteria during ripening (Johnson et al., 1990a), storing cheeses at a lower temperature that initially were held at a higher temperature (Pearce et al., 1973), and loose packaging (Johnson et al., 1990b) promote crystal formation. However, the continuing problems with calcium lactate crystals show the need for a better understanding of the process of crystal formation.

A good overview of the mechanism of CaL2 crystal formation was provided by Dybing et al. (1988) and is illustrated by the following equation:


(1)

Calcium lactate is not a native component of milk; therefore it is necessary to have sufficient amounts of both calcium and lactate ions to form crystals. Once calcium and lactate concentrations exceed the solubility limit at a certain temperature, a thermodynamic driving force for crystallization exists and crystallization might occur. The presence of nucleation sites, such as rough surfaces and cracks, will accelerate crystal formation. Additionally, differences in milk composition and the cheese manufacturing, packaging, and handling procedures might affect calcium availability and lactate synthesis and, hence, crystal formation (Dybing et al., 1988).

Very few studies have evaluated the solubility of CaL2 at different temperatures. The previous studies, reviewed and summarized by Kubantseva and Hartel (2002), show numerous discrepancies. The variability of solubility data may be due to differences in CaL2 used as a starting material (impurities, isomeric form, and the amount of water contained), differences in equilibration methods, insufficient amount of time allowed for equilibration, improper sample taking and filtering techniques, and differences in methods for calcium determination (Kubantseva and Hartel, 2002).

The present study evaluates the solubility of CaL2 in aqueous model systems with and without the addition of impurities important in cheese. By creating an aqueous solution with pH and salt content similar to cheese serum, the solubility of CaL2 in cheese serum at various temperatures can be estimated. This research provides an understanding of what parameters influence the solubility of CaL2 and provides a basis for future studies on crystallization kinetics of CaL2.


   MATERIALS AND METHODS
TOP
 ABSTRACT
 INTRODUCTION
 MATERIALS AND METHODS
RESULTS AND DISCUSSION
 CONCLUSIONS
 ACKNOWLEDGEMENTS
 REFERENCES
 
L(+)-CaL 2 pentahydrate (Ca (CH 3 CHOHCOO) 2•5H2O)USP grade (purity>99.0%) was purchased from PURAC America, Inc. (Lincolnshire, IL) and Sigma Chemical Co. (St. Louis, MO). Crystals of CaL2 received from Sigma Chemical Co. were needle-shaped, whereas the specimen of CaL2 purchased from PURAC America, Inc. was more powderlike. All other chemicals used were analytical grade and purchased from Sigma Chemical Co. (St. Louis, MO).

The solubility of CaL2 in aqueous model systems was evaluated in terms of temperature, presence of salt, pH, and various calcium and lactate ion concentrations. Saturated solutions of CaL2 with and without impurities were prepared and analyzed as described below.

Two methods of equilibration at constant temperature were used to obtain the saturation concentration of CaL2 in solution. The first method involved mixing an excess amount of crystalline CaL2 with deionized distilled water and allowing the crystals to dissolve until equilibrium was reached (method 1). The second method involved crystallization of CaL2 from a supersaturated solution by cooling until equilibrium was reached at the desired temperature (method 2). For the pure system (CaL2 and water), both methods were used to ensure that solutions achieved true equilibrium. Based on the results of equilibration with CaL2, only method 1 was used for experiments with added impurities since equilibration was approached more rapidly.

Prepared aqueous solutions of CaL2 were placed in flasks with a stirrer, and the flasks were set in a temperature-controlled water bath. Temperature was controlled by an Isotemp immersion circulator (Fisher Scientific, model 2150) and a refrigerated circulating bath with a digital controller (VWR Scientific, model 1157). Fluctuations of temperature were less than ± 0.1°C. After several days of mixing, samples were taken and quickly filtered (to maintain the temperature of the solution) using a syringe with an attached fiberglass filter (Millipore Corp., Bedford, MA). To determine the calcium concentration in the filtrate, the AOAC standard method for calcium determination (AOAC International, method 968.31) based on EDTA complexometric titration was used. A standard curve for calcium was established and used to determine the calcium content in samples. Titration was performed with an automatic buret (Brinkmann/Metrohm, 665 Dosimat) and an excellent reproducibility of results was observed. Samples were allowed to stir until both equilibration methods gave the same results to ensure that the system had come to equilibrium.

Temperatures of 4, 10, and 24°C were used, corresponding to typical storage temperature, typical aging temperature, and room temperature, respectively. To test the effect of salt on the solubility of CaL2, we added NaCl at 3, 4, 5, and 6 g/100 g of water prior to the addition of CaL2. The effect of pH on CaL2 solubility was tested at 4.8, 5.0, 5.2, 5.4, and 6.5, adjusted to the desired level by dropwise addition of 3.0 N HCl to the solution of CaL2 once the solution had reached equilibrium. The pH was measured using a pH meter (Accumet, model 815MP) with a relative accuracy of ± 0.01. In all cases, solubility was expressed by the amount of anhydrous salt of CaL2 dissolved in 100 g of water.

The effects of calcium and lactate ion concentrations on the solubility of CaL2 were investigated using calcium chloride (CaCl2) to adjust the calcium ion concentration and sodium lactate (NaL) to adjust the lactate ion concentration. Standard solutions of CaCl2 and NaL were prepared (100 mL each) with 5, 10, 15, or 20% excess concentration of calcium or lactate at saturation. Excess CaL2 was added to the standard solutions and the mixtures were allowed to equilibrate, as determined by the calcium ion concentration. Changes in the solubility of CaL2 relative to stoichiometric dissolution were calculated from the dissolved calcium levels. For solutions with excess calcium already added, the solubility of CaL2 was determined by subtracting the starting calcium level. For solutions with lactate already added, the calcium concentration represented the amount of CaL2.

All solubility measurements were determined in triplicate using a full factorial design. Excellent reproducibility of results was observed with standard deviations in solubility values less than 0.03 g of anhydrous CaL2/100 g of water. Factors including 3 temperatures (4, 10, and 24°C), 5 salt contents (0, 3, 4, 5, and 6 g/100 g of water), and 5 pH (4.8, 5.0, 5.2, 5.4, and 6.5) were analyzed by Microsoft Excel (using single factor ANOVA, {alpha} = 0.05) to determine the statistical significance on the solubility of CaL2.


   RESULTS AND DISCUSSION
TOP
 ABSTRACT
 INTRODUCTION
 MATERIALS AND METHODS
 RESULTS AND DISCUSSION
CONCLUSIONS
 ACKNOWLEDGEMENTS
 REFERENCES
 
Temperature and Source Effects
Two different sources of CaL2 were tested in this research, and significant differences in solubility between the 2 specimens were observed. The solubility values of CaL2 as supplied by Sigma Chemical Co. were 1.88, 2.38, and 3.89 g of anhydrous CaL2/100 g of water at 4, 10, and 24°C, respectively. These values were lower than the solubility of CaL2 supplied by PURAC America, Inc., which were 3.38, 4.04, and 6.41 g of anhydrous CaL2/100 g of water at 4, 10, and 24°C, respectively. Both specimens contained the same amount of water (5 molecules), were L(+) isomers, and were produced by the sugar beet fermentation process, so there should have been no differences in production method. The only possible difference was the amount of impurities in the end product; however, both specimens of CaL2 met USP specifications. The CaL2 chosen for these experiments was purchased from PURAC America, Inc., since it was possible to obtain a large quantity of material from the same lot.

Previous research by Glass (1933) and Hill and Cocking (1912) did not find that age, acidity, or different isomeric form of the initial CaL2 affected its final solubility. Together with Macmorran (1933), Glass (1933) proposed that the differences in solubility are probably due to the method of preparation of the CaL2.

The 2 equilibration methods resulted in identical solubility concentrations, as shown in Figure 1Go, indicating that the system had reached equilibrium with both methods. Method 2 (crystallization) required a longer time to achieve equilibrium then method 1 (dissolution). Method 1 gave stable solubility results within a week, so in all subsequent experiments only method 1 was used. The pH remained unchanged (6.5 to 6.6) throughout the equilibration time.



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  		Figure 1. Solubility of calcium lactate (CaL2) in water at 4°C: {blacksquare}, method 1 (dissolution of CaL2 crystals); {square}, method 2 (crystallization of CaL2 from supersaturated solution).

 
These experiments have shown that it takes at least 1 wk at 4°C for a solution to reach equilibrium, even with method 1. At 24°C, equilibration was generally not completed until 4 d of agitation. Several hours of mixing, as reported by Cao et al. (2001) and Pateenko and Smirnov (1974), are not sufficient to ensure equilibrium. The short equilibration time could explain the lower solubility values of Pateenko and Smirnov (1974). In contrast, despite the short mixing times, Cao et al. (2001) obtained the highest (of all reported in the literature) solubility values at a given temperature. We can only speculate that this disagreement may be due to differences in starting materials.

As seen in Figure 2Go, solubility did not change during 3 wk of agitation at constant temperature, confirming that the true equilibrium was reached. The solubility of CaL2 was found to be 3.38, 4.04, and 6.41 g of anhydrous CaL2/100 g of water at 4, 10, and 24°C, respectively. These data are in general agreement with the results from previous studies (summarized by Kubantseva and Hartel, 2002), which reported solubility values of CaL2 of about 3, 4, and 6 g of anhydrous CaL2/100 g of water at 4, 10, and 24°C, respectively. As expected, solubility increased with an increase in temperature.



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  		Figure 2. Solubility of calcium lactate (CaL2) in water as a function of temperature: {blacksquare}, 4°C; •, 10°C; {blacktriangleup}, 24°C (standard deviation <= 0.01 g of anhydrous CaL2 /100 g of water).

 
Salt Effect
The effect of salt (in the range between 0 and 6 g of NaCl/100 g water) on the solubility of CaL2 in water is shown in Figure 3Go. Statistical analysis indicated significant differences (P < 0.02) between the solubilities of CaL2 with various salt concentrations added (see Table 1Go). However, all differences are quite small and for practical purposes can be considered insignificant. These results are consistent with those reported by Pearce et al. (1973), who also noted only a small effect of salt on the solubility of CaL2. Again, no change in pH values (6.5 to 6.6) was observed for any NaCl level.



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  		Figure 3. Solubility of calcium lactate (CaL2) in water as a function of NaCl at various temperatures: {blacksquare}, 4°C; •, 10°C; {blacktriangleup}, 24°C (standard deviation <= 0.01 g of anhydrous CaL2/100 g of water).

 

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  		Table 1. Solubility of calcium lactate (CaL2) at different pH and temperature in the presence of added NaCl.
 
pH Effect
The effects of pH on CaL2 solubility are presented in Table 1Go. A change in pH had only a minor effect on the solubility of CaL2 within the studied range; however, due to the small standard error, the effect of pH was statistically significant at most pH levels (see Table 1Go). A slight increase in solubility with a decrease in pH was observed. From the NaCl experiments, it was concluded that chloride ions have almost no effect on the solubility of CaL2, therefore it was possible to use HCl to adjust pH and identify the pH effect as a single factor. Roth-Bassell and Clydesdale (1992) compared the solubility of CaL2 in water at pH values of 2.0 and 7.0 by using HCl to adjust the pH. Their results showed a 20% decrease in the solubility of CaL2 when the pH was raised from 2.0 to 7.0, a trend generally consistent with our observations.

Calcium and Lactate Ions Effects
The solubility of CaL2 in water in the presence of additional calcium and lactate ions is shown in Figures 4Go and 5Go, respectively. Additional calcium and lactate ions are defined as the amounts (in %) added above the equilibrium level. At all temperatures, the addition of calcium ions (5 to 20% calcium above the saturation level) to the CaL2 solution had virtually no effect on the solubility of CaL2 in water and did not change the pH (between 6.5 and 6.6). Statistically, however, the effect was significant (P < 0.001) at some levels of addition due to a small standard error. Dybing et al. (1988) also showed that the addition of calcium ions in the form of CaCl2 to cheese serum had no effect on crystal formation.



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  		Figure 4. Effect of an addition of calcium ions on the solubility of CaL2: {blacksquare}, 4°C; •, 10°C; {blacktriangleup}, 24°C (standard deviation = 0.03 g of anhydrous CaL2/100 g of water).

 


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  		Figure 5. Effect of an addition of lactate ions on the solubility of CaL2: {blacksquare}, 4°C; •, 10°C; {blacktriangleup}, 24°C (standard deviation <= 0.03 g of anhydrous CaL2/100 g of water).

 
On the other hand, when calcium ions were added to the CaL2 solution at much higher concentrations (200% or more calcium above the saturation level), a decrease (P < 0.001) in the solubility of CaL2 was observed at all 3 temperatures. A decrease in pH to 5.7 to 5.8 was also observed.

At all temperatures, the addition of excess lactate ions (5 to 20% lactate above saturation level) caused a decrease (P < 0.001) in the solubility of CaL2 in water, with the effect as large as 14 to 16% depending on the temperature (see Figure 5Go). The addition of NaL had no significant effect on the pH (between 6.5 and 6.8). When the lactate ion concentration was increased to 100% or more above saturation, the solubility of CaL2 significantly (P < 0.001) decreased (by 33 to 43%), while the pH increased to 6.9 to 7.2. With added lactate ions, the equilibrium shifted towards the formation of solid CaL2, according to Le Châtelier’s principle, so less CaL2 was in the soluble form at equilibrium.


   CONCLUSIONS
TOP
 ABSTRACT
 INTRODUCTION
 MATERIALS AND METHODS
 RESULTS AND DISCUSSION
 CONCLUSIONS
ACKNOWLEDGEMENTS
 REFERENCES
 
The observed decrease in CaL2 solubility suggests that an excess level of lactate ions may be a primary reason for CaL2 crystal appearance on cheese. Therefore, lactose content should be controlled in the cheese milk to prevent CaL2 crystals from forming. Temperature fluctuations (from higher to lower) during storage also may be a reason for crystal appearance as the solubility of CaL2 decreases with a decrease in temperature. Parameters such as pH, NaCl content, and excess calcium ions did not affect CaL2 solubility and, if they are responsible for CaL2 crystal formation, it is not due to a decrease in CaL2 solubility.


   ACKNOWLEDGEMENTS
TOP
 ABSTRACT
 INTRODUCTION
 MATERIALS AND METHODS
 RESULTS AND DISCUSSION
 CONCLUSIONS
 ACKNOWLEDGEMENTS
REFERENCES
 
This project was funded by Land O’Lakes, Inc. (Arden Hills, MN). The authors would like to express their appreciation to Baomin Liang, John Lucey, and Mark Johnson of the Wisconsin Center for Dairy Research for assistance and discussion.

Received for publication September 12, 2003. Accepted for publication November 21, 2003.


   REFERENCES
TOP
 ABSTRACT
 INTRODUCTION
 MATERIALS AND METHODS
 RESULTS AND DISCUSSION
 CONCLUSIONS
 ACKNOWLEDGEMENTS
 REFERENCES
 


Association of Official Analytical Chemists. 1990. Official Methods of Analysis. Vol. II. 15th ed. Arlington, VA.

Cao, X., H-J. Lee, H. S. Yun, and Y-M. Koo. 2001. Solubilities of calcium and zinc lactate in water and water-ethanol mixture. Korean J. Chem. Eng. 18:133–135.

Chan, Y-W. 2001. Crystallization kinetics of calcium lactate. Independent Study Report, University of Wisconsin, Madison.

Dybing, S. T., J. A. Wiegand, S. A. Brudvig, E. A. Huang, and R. C. Chandan. 1988. Effect of processing variables on the formation of calcium lactate crystals on Cheddar cheese. J. Dairy Sci. 71:1701–1710.[Abstract/Free Full Text]

Glass, N. 1933. The variation in solubility of calcium lactate. Q. J. Pharm. Pharmacol. 6:522–530.

Hill, C. A., and T. T. Cocking. 1912. Note on calcium lactate. Yearb. Pharm. 481–486.

Johnson, M. E., B. A. Riesterer, and N. F. Olson. 1990a. Influence of nonstarter bacteria on calcium lactate crystallization on the surface of Cheddar cheese. J. Dairy Sci. 73:1145–1149.[Abstract/Free Full Text]

Johnson, M. E., B. A. Riesterer, C. Chen, B. Tricomi, and N. F. Olson. 1990b. Effect of packaging and storage conditions on calcium lactate crystallization on the surface of Cheddar cheese. J. Dairy Sci. 73:3033–3041.[Abstract]

Kubantseva, N. and R. W. Hartel. 2002. Solubility of calcium lactate in aqueous solutions. Food Rev. Intl. 18:135–149.

Macmorran, G. H. 1933. The solubility of calcium lactate. Pharm. J. 130:245.

Pateenko, S. K., and V. A. Smirnov. 1974. Solubility of calcium lactate. Khlebopek. Konditer. Prom. (Bread and Candy Manufacturing) 1:27–29.

Pearce, K. N., L. K. Creamer, and J. Gilles. 1973. Calcium lactate deposits on rindless Cheddar cheese. N.Z. J. Dairy Sci. Technol. 8:3–7.

Roth-Bassell, H. A., and F. M. Clydesdale. 1992. In vitro solubility characteristics of six calcium salts. J. Food Prot. 55:1003–1005.


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