Calcium is an important part of human bones and teeth. It plays a vital role in every stage of human growth. The lack of calcium can cause health problems such as convulsions, rickets in children, and osteoporosis in older people. Milk has always been highlighted by researchers as a good source of calcium. In milk, about one-third of the calcium is soluble in the water phase (serum phase), mainly as free calcium ion (Ca2+), calcium citrate complex (CaCitr−), calcium hydrogen phosphate complex (CaHPO40), and a few other calcium complexes, while the other about two-thirds of the calcium is part of the casein protein in the form of micellar calcium. Compared with micellar calcium, soluble calcium is easier to be absorbed in the intestine, and CaCitr− accounts for 70% of the total soluble calcium. Therefore, to improve the nutritive value of milk and the technological properties of milk, it is very important to understand the properties of CaCitr−complexes and calcium citrate.

In the work of Paper I, we identified four solid forms of calcium citrate, including hexahydrate (CCH), tetrahydrate (CCT), dihydrate (CCD), and anhydrous calcium citrate (CCA), through thermogravimetric experiments. The Differential Scanning Calorimeter (DSC) was used to characterize the transformation of CCH as a gradual transformation along the series CCH  CCT  CCD  CCA. The enthalpy change of dehydration (ΔH0dehydr) for each step was determined experimentally. By investigating the solubility of the various crystals at different temperatures, it was found that CCH is the most thermodynamically stable form among the four different forms of calcium citrates at 25 ℃. The dissolution enthalpy (ΔH0dissol) and solubility product (Ksp) of each form were determined by EDTA titration at varying temperatures. We also explored the solubility of a commercial form of CCT at different pH, and the results showed that pH significantly affects the ion distribution in CCT solutions. At the pH of the small intestine around 8, the concentration of Ca2+ in a saturated solution of CCT is still higher than 1 mM, and this concentration is a necessary condition for absorption of Ca2+ in the small intestine.

In Paper II, strontium and calcium as elements of the same main group are compared for their binding to citrate. Addition of strontium chloride to a calcium citrate tetrahydrate (CCT) suspension significantly increased the activity and the concentration of Ca2+, as strontium was found to replace calcium in calcium citrate complexes and to form strontium citrate precipitates with lower solubility than calcium citrate and release more Ca2+ into solution. The association constants of CaCitr− and SrCitr− were estimated by measuring the conductivity of strontium citrate and calcium citrate in dilute solutions. The enthalpy change (ΔH0) of the coordination reaction was obtained from the temperature variation of the association constants. We have also measured the solubility of the strontium citrate pentahydrate at different pH by EDTA titration, and the results showed that pH significantly affects solubility and the ion distribution in aqueous solution.

According to the study reported in Paper III, CaCitr− accounts for 70% of the soluble calcium in milk because of the strong coordination bond formed between calcium ions and citrate ions. In agreement with this, a spontaneous supersaturation phenomenon during codissolution of calcium hydroxycarboxylates and citrate salts in water was identified. In this part of the work, we investigated the effect of temperature on the kinetics of the supersaturation formation in calcium lactate and sodium citrate aqueous solutions (CaLact2 + Na3Citr). By combining small-angle X-ray scattering and calcium ion activity measurement, we followed and analyzed the entire supersaturation process, which includes three steps: step I dissolution of calcium lactate assisted by citrate, step II a metastable homogeneous state, and step III formation of precipitate. In step III, the activation enthalpy of the decrease of calcium ion concentration is significantly lower than that for precipitate formation, a difference which should be the main reason for the occurrence of supersaturation and for the long lag phase of step II prior to step III.