Martina Vavrusova

Aqueous Solubility of Calcium l-Lactate, Calcium d-Gluconate, and Calcium d-Lactobionate: Importance of Complex Formation for Solubility Increase by Hydroxycarboxylate Mixtures

Among the calcium hydroxycarboxylates important for cheese quality, D-lactobionate [Ksp = (7.0 ± 0.3) × 10(-3) mol(3) L(-3)] and L-lactate [Ksp = (5.8 ± 0.2) × 10(-3) mol(3) L(-3)] were found more soluble than D-gluconate [Ksp = (7.1 ± 0.2) × 10(-4) mol(3) L(-3)], as indicated by the solubility products determined electrochemically for aqueous 1.0 M NaCl at 25.0 °C. Still, solubility of calcium L-lactate increases by 45% in the presence of 0.50 M sodium D-gluconate and by 37% in the presence of 0.50 M sodium D-lactobionate, while solubility of calcium D-gluconate increases by 66 and 85% in the presence of 0.50 M sodium L-lactate and 0.50 M sodium D-lactobionate, respectively, as determined by complexometric titration. Sodium L-lactate and sodium D-gluconate have only little influence on solubility of calcium D-lactobionate. The increased solubility is described quantitatively by calcium binding to D-gluconate (K1 = 14 ± 3 mol(-1) L) in 1.0 M NaCl at 25 °C, D-lactobionate (K1 = 11 ± 2 mol(-1) L), and L-lactate (K1 = 8 ± 2 mol(-1) L), as indicated by the association constants determined electrochemically. In mixed hydroxycarboxylate solutions, calcium binding is quantitatively described by the geometric mean of the individual association constants for both aqueous 1.0 and 0.20 M NaCl, indicating a 1:1 stoichiometry for complex formation.

Thermodynamics of Dissolution of Calcium Hydroxycarboxylates in Water

Aqueous solubility of calcium l-lactate, calcium d-gluconate, and calcium d-lactobionate increases with temperature (10-30 °C investigated), most significantly for the least soluble d-gluconate, while the calcium ion activity of the saturated solutions decreases with temperature, as measured electrochemically, most significantly for the most soluble d-lactobionate. This unusual behavior is discussed in relation to dairy processing and explained by endothermic binding of calcium to hydroxycarboxylate anions determined to have ΔH°ass = (31 ± 3) kJ·mol(-1) for l-lactate, (34 ± 2) kJ·mol(-1) for d-gluconate, and (29 ± 3) kJ·mol(-1) for d-lactobionate in 1:1 complexes with thermodynamic binding constants at 25 °C of Kass = 49 (l-lactate), 88 (d-gluconate), and 140 (d-lactobionate). Quantum mechanical calculations within density functional theory (DFT) confirm the ordering of strength of binding. The complex formation is entropy driven with ΔS°ass > 0, resulting in decreasing calcium ion activity in aqueous solutions for increasing temperature, even for the saturated solutions despite increasing solubility.

Calcium binding to dipeptides of aspartate and glutamate in comparison with orthophosphoserine.

Aspartate binds calcium(II) better than glutamate with Ka = 7.0 ± 0.9 L mol⁻¹ for Asp and Ka = 3.0 ± 0.8 L mol⁻¹ for Glu, respectively, as determined using calcium-selective electrodes for aqueous solutions of ionic strength 0.20 at 25 °C at pH of relevance for milk products. For the mixed peptides, the affinity seems additive with Ka = 27 ± 3 L mol⁻¹ for Asp-Glu and 22.7 ± 0.1 for Glu-Asp as compared to the expected 21 L mol⁻¹. In contrast, for Asp-Asp, the affinity is less than additive with Ka = 23 ± 5 L mol⁻¹ as compared to the expected 49 L mol⁻¹, whereas for Glu-Glu, the affinity is more than additive with Ka = 26 ± 4 L mol⁻¹ as compared to the expected 9.0 L mol⁻¹, indicating specific structural effects for Glu-Glu. Ionic strength effects, 1.0 versus 0.20 studied, are similar for Asp and Glu with decreasing affinity for higher ionic strength, whereas the dipeptides with Glu as C-terminus are more sensitive to increasing ionic strength than with Asp as C-terminus. Despite little affinity of calcium to serine with Ka = 0.9 ± 0.2 L mol⁻¹, Glu has increasing affinity for calcium in the serine dipeptide Ser-Glu with Ka = 10 ± 3 L mol⁻¹, which becomes comparable to phosphorylated serine with Ka = 22 ± 5 L mol⁻¹.

Calcium d-Saccharate: Aqueous Solubility, Complex Formation, and Stabilization of Supersaturation

Molar conductivity of saturated aqueous solutions of calcium d-saccharate, used as a stabilizer of beverages fortified with calcium d-gluconate, increases strongly upon dilution, indicating complex formation between calcium and d-saccharate ions, for which, at 25 °C, Kassoc = 1032 ± 80, ΔHassoc° = −34 ± 6 kJ mol–1, and ΔSassoc° = −55 ± 9 J mol–1 K–1, were determined electrochemically. Calcium d-saccharate is sparingly soluble, with a solubility product, Ksp, of (6.17 ± 0.32) × 10–7 at 25 °C, only moderately increasing with the temperature: ΔHsol° = 48 ± 2 kJ mol–1, and ΔSassoc° = 42 ± 7 J mol–1 K–1. Equilibria in supersaturated solutions of calcium d-saccharate seem only to adjust slowly, as seen from calcium activity measurements in calcium d-saccharate solutions made supersaturated by cooling. Solutions formed by isothermal dissolution of calcium d-gluconate in aqueous potassium d-saccharate becomes spontaneously supersaturated with both d-gluconate and d-saccharate calcium salts, from which only calcium d-saccharate slowly precipitates. Calcium d-saccharate is suggested to act as a stabilizer of supersaturated solutions of other calcium hydroxycarboxylates with endothermic complex formation through a heat-induced shift in calcium complex distribution with slow equilibration upon cooling.

Spontaneous supersaturation of calcium citrate from simultaneous isothermal dissolution of sodium citrate and sparingly soluble calcium hydroxycarboxylates in water

Strongly supersaturated homogeneous calcium citrate solutions are formed spontaneously when solid sodium citrate and solid calcium hydroxycarboxylates are dissolved simultaneously in water or when solid sodium citrate is dissolved in an already saturated aqueous solution of the calcium hydroxycarboxylate at ambient conditions. Maximal supersaturation of calcium citrate was found to decrease for an increasing value of the stability constant for calcium binding: L-lactate < D-gluconate < citrate, indicating citrate assisted dissolution through competitive complex formation as a thermodynamic factor controlling spontaneous supersaturation for up to a factor of more than twenty. Time elapsing prior to initiation of precipitation of calcium citrate was found to be shorter for a higher degree of supersaturation and lasted between hours and days. During subsequent precipitation equilibrium solubility of calcium citrate was approached with a simultaneous increase in water activity. Both thermodynamic and kinetic factors are suggested to be important for the spontaneous supersaturation, which seems to explain the paradoxal but well-stablished high bioavailability of calcium from the sparingly soluble calcium citrate and the high mobility of calcium in the presence of citrate during biomineralization.

Supersaturation of calcium citrate as a mechanism behind enhanced availability of calcium phosphates by presence of citrate

Dissolution of amorphous calcium phosphate (ACP) in aqueous citrate at varying pH has been studied with perspective of increasing availability of calcium from sidestreams of whey protein, lactose and/or cheese production or on development of new functional foods. ACP formed as an initial precipitate in 0.10 mol L-1 equimolar aqueous calcium chloride, sodium citrate, and sodium hydrogenphosphate was used as model for mineral residues formed during milk processing. Upon acidification of the ACP suspension by hydrochloric acid decreasing pH from 6.5 to 4.5, the transformations of ACP occurred through an 8 h period of supersaturation prior to a slow precipitation of calcium citrate tetrahydrate. This robust supersaturation, which may explain increased availability of calcium phosphates in presence of citrate, presented a degree of supersaturation of 7.1 and was characterized by precipitation rates for 0.10 mol L-1 equimolar aqueous calcium chloride, sodium hydrogencitrate, and sodium hydrogenphosphate with pH 5.5, and for 0.10 mol L-1 equimolar aqueous calcium chloride, sodium hydrogencitrate, and sodium dihydrogenphosphate with pH 4.1, with a degree of supersaturation of 2.7. The crystallization processes were similar according to Avramis model with a half-life for precipitation of approximately 5 h independent of the degree of supersaturation. Ion speciation based on measurement of pH, and total concentrations of calcium, phosphate and citrate, and of conductivity and calcium ion activity during precipitation indicates a low driving force for precipitation with calcium citrate complex dominating at pH 5.5 and calcium hydrogencitrate complex dominating at pH 4.1. Calcium hydrogencitrate is suggested to be the species involved in the crystal growth followed by solid state transformation to calcium citrate tetrahydrate.