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3.2: Basic Facts About Calcium- Its Compounds and Reactions

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    Basic Facts

    Calcium was first recognized as an element in 1808 by Humphry Davy, and the name was given after the Latin for lime: calx. Several isotopes of calcium are known. The stable isotopes are, in order of decreasing natural abundance, 40Ca (96.94%), 44Ca (2.1%), 42Ca (0.64%), and 43Ca (0.145%). 43Ca is the only isotope with a nuclear spin (I = \(\frac{7}{2}\)) different from zero, which makes it amenable to NMR studies. 45Ca is a radioactive isotope of some importance (\(\beta^{-}\) decay; 8.8 min half life).3 It has been used in studies of calcium localization and transport in biological systems.

    Calcium constitutes about 3 percent by weight of the Earth's crust, mostly in the form of sedimentary rocks of biological origin dating back some three billion years. In sea water the total concentration of calcium ranges from 5 to 50 times higher than in fresh water, which, in turn, has a calcium concentration ten times that of rain water (see Table 3.1). This explains the pleasant feeling when ordinary soaps are used in rain water. The calcium concentration in ordinary tap water varies with location; calcium is usually added to water in distributing networks in order to prevent corrosion of iron pipes. Tap water with a calcium concentration above 1.5 mM is usually classified as "hard." Interestingly, the taste of beer seems related to the calcium concentration, and it is claimed that "good" beer should have a concentration higher than that of "hard" tap water.

    In the body fluids of higher organisms the total calcium concentration is usually on the order of a few millimolar (see Table 3.1). In adult human serum, the concentration is observed to be, within narrow limits, 2.45 mM.

    Table 3.1 - Ca2+ concentrations in fluids and tissues.6-9

    Specimen Units are mM if not otherwise stated
    Sea water 10
    Fresh water 0.02 - 2
    Rain water 0.002 - 0.02
    "Hard" tap water 1.5
    "Good" beer 4
    Adult human serum 2.45 \(\pm\) 0.05
    Serum of other vertebrates 1.5 - 5
    Nematote body fluids 6
    Molluscan serum - marine 9 - 15
    Molluscan serum - fresh water 1.5 - 7.8
    Molluscan serum - land 3.3 - 12.3
    Milk 70
    Bone 0.8 - 1.0
    Mitochondria from rat liver 0.8 \(\pm\) 0.1 mmol/kg
    Endoplasmatic reticulum 8 - 10 mmol/kg
    Cytoplasm of a resting mammalian cell 0.0001
    Cytoplasm of E. coli 0.0001

    Essentials of Ca2+ Chemistry

    Since the Ca2+ ion accomplishes its biological tasks in an environment with 1 to 3 mM Mg2+, it is of particular interest to compare the properties of these two ions in order to understand how a discrimination is made in biological systems. In addition, the coordination chemistry of Ca2+ is closely related to that of Mg2+ (as well as Cd2+), though there are several obvious differences. First of all, the ionic radius of a Ca2+ ion with a given coordination number (CN) is always higher than that of an Mg2+ or Cd2+ ion with the same CN. At CN = 6, the ionic radii of Ca2+, Cd2+, and Mg2+ are 1.00, 0.95, and 0.72 Å, respectively, whereas at CN = 8 they are 1.12, 1.10, and 0.89 Å, respectively.4

    Ligand preferences of Ca2+ depend on the fact that it is a hard metal ion. Thus Ca2+ strongly prefers oxygen ligands over nitrogen or sulfur ligands; Ca2+••••N bonds are about 0.25—0.3 Å longer than Ca2+••••O bonds.5,6,10 Large differences in coordination number and geometry have been observed for Ca2+ complexes. In a study of 170 x-ray structures of Ca2+ complexes involving carboxylate groups,11 binding was found to be either (i) unidentate, in which the Ca2+ ion interacts with only one of the two carboxylate oxygens, (ii) bidentate, in which the Ca2+ ion is chelated by both carboxylate oxygens, or (iii) mixed ("\(\alpha\)-mode") in which the Ca2+ ion is chelated by one of the carboxylate oxygens and another ligand attached to the \(\alpha\)-carbon (see Figure 3.1). The Ca2+- oxygen distances span a range from 2.30 to 2.50 Å, with the average distance being 2.38 Å in the unidentate and 2.53 Å in the bidentate mode, respectively.11 Observed coordination numbers follow the order 8 > 7 > 6 > 9. By contrast, Mg2+ nearly always occupies the center of an octahedron of oxygen atoms (CN = 6) at a fixed Mg2+-oxygen distance of 2.05 ± 0.05 Å.

    clipboard_e3df9cc05febf6638080f76e002210e13
    Figure 3.1 - The three commonly observed modes of calcium carboxylate ligation. (A) The unidentate mode, in which the calcium ion interacts with only one of the two carboxylate oxygens. (B) The bidentate mode, in which the calcium ion is chelated by both oxygen atoms. (C) The \(\alpha\)-mode, in which the calcium ion is chelated by one carboxylate oxygen, and another ligand is attached to the a-carbon. Adapted from Reference II.

    In Table 3.2, stability constants for the binding of Ca2+ and Mg2+ to various ligands are collected. We may note that selectivity of Ca 2+ over Mg2+ is not very great for simple carboxylate ligands, but that it tends to increase for large multidentate ligands, such as EDTA and in particular EGTA. The Ca2+ sites in many intracellular proteins with "EF-hand" binding sites (see Section V. C) bind Ca2+ about 104 times more strongly than Mg2+.

    Table 3.2 - Ca2+ and Mg2+ (where available) stability constants (log K) for different organic and biochemical ligands. Most values are at ionic strength 0.1 and 25 °C.5,6,12-15

    a) EGTA: ethylenebis(oxyethylenenitrilo)tetraacetate

    b) EDTA: ethylenedinitrilotetraacetate

    Ligand Ca2+ Mg2+
    Acetate 0.5 0.5
    Lactate 1.1 0.9
    Malonate 1.5 2.1
    Aspartate 1.6 2.4
    Citrate 3.5 3.4
    Nitrilotriacetate 6.4 5.5
    EGTAa 10.9 5.3
    EDTAb 10.6 8.8
    Glycine (Gly) 1.4 3.4
    \(\gamma\)-Carboxyglutamic acid (Gla) 1.3
    Gly-Gly dipeptide 1.2
    Gla-Gla dipeptide 3.2
    Macrobicyclic amino cryptate [2.2.2] 4.5
    Fluo-3 6.2 2.0
    Fura-2 6.9 2.0
    BAPTA 7.0 1.8
    Quin-2 7.1 2.7
    Phospholipase A2 3.6
    Thrombin fragment 1 3.7 3.0
    Trypsinogen 3.8
    Chymotrypsinogen 3.9
    Chymotrypsin 4.1
    Calmodulin, N-terminal 4.5 3.3
    Trypsin 4.6
    Calmodulin, C-terminal 5.3
    Protein kinase C \(\sim\)7
    \(\alpha\)-Lactabumin \(\sim\)7
    Rabbit skeletal muscle
    Troponin C, Ca2+/Mg2+ sites 7.3 3.6
    Carp parvalbumin \(\sim\)8.5 4.2
    Bovine calbindin D9K 8.8 \(\sim\)4.3

    Another difference in ligand-binding properties of Mg2+ and Ca2+ can be seen by comparing the rates of substitution of water molecules in the inner hydration sphere by simple ligands, according to

    \[M(H_{2}O)_{n}^{2+} + L \xrightarrow{k} ML(H_{2}O)_{n-1}^{2+} + H_{2}O\]

    This rate (log k, with k in s-1) has been determined to be 8.4 for Ca2+ and 5.2 for Mg2+.16

    The formation of biominerals is a complex phenomenon. In order to obtain a feeling for the conditions under which inorganic solid phases in biological systems are stable, it is of some interest to look at solubility products. Solubility products, Kspo, have a meaning only if the composition of the solid phase is specified. For a solid compound with the general composition (A)k(B)l(C)m the solubility product is defined as

    \[K_{sp}^{o} = [A]^{k} [B]^{l} [C]^{m} \tag{3.1}\]

    where [A], [B], etc., denote activities of the respective species, usually ionic, in equilibrium with the solid. Activities are concentrations multiplied by an activity coefficient, \(\gamma\), nearly always less than unity. Activity coefficients for ions in real solutions can be estimated from Debye-Hückel theory17 if the ionic strength of the solution is known. In human blood plasma, the ionic strength, I, is about 0.16, and the activity coefficient for Ca2+ at 37 °C is 0.34. In many discussions it may be sufficient to equate concentrations with activities.

    The solid phase involved is essentially assumed to be an infinitely large, defect- and impurity-free crystal with a well-defined structure. Microscopic crystals have higher solubilities than large crystals, a well-known phenomenon that leads to "aging" of precipitates, in which larger crystals grow at the expense of smaller ones.

    Many anionic species appearing in the solubility products may also be involved in protonation equilibria in solution, such as those of phosphoric acid: H2PO4- \(\rightleftharpoons\) H+ + HPO42-; HPO42- \(\rightleftharpoons\) PO43- + H+; etc. When the prospects for the formation of a solid phase under certain solution conditions are investigated, the activity, or concentration, of the particular anionic species specified in the solubility product must be known, not only "total phosphate" or "total calcium," etc. The data in Table 3.3 show that, at pH > 5, the most stable (i.e., insoluble) solid calcium phosphate is hydroxyapatite.

    Table 3.3 - Solubility products, at pH 5 and 25 °C, for solid calcium phosphates

    Solid Phase -log Kspo -log Kspo of corresponding Mg2+ compound where applicable
    CaSO4 • 2H2O (sulfate, "gypsum") 5.1 < 1.0
    Ca(OH)2 (hydroxide) 5.3 10.7
    CaHPO4 • 2H2O (hydrogen phosphate) 6.6
    CaCO3 (carbonate, "calcite," "aragonite") 8.5 7.5
    CaC2O4 • H2O (oxalate, "whewellite") 10.5 5.0
    \(\beta\)-Ca3(PO4)2 (\(\beta\)-phosphate) 29
    Ca5(PO4)3OH (hydroxyapatite) 58

    3.2: Basic Facts About Calcium- Its Compounds and Reactions is shared under a not declared license and was authored, remixed, and/or curated by LibreTexts.

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