The calcium bicarbonate is an inorganic salt with chemical formula Ca (HCO 3 ) 2 . It originates in nature from the calcium carbonate present in limestone stones and minerals such as calcite.
Calcium bicarbonate is more soluble in water than calcium carbonate. This characteristic has allowed the formation of karst systems in limestone rocks and in the structuring of caves.
The groundwater that passes through the cracks becomes saturated in its displacement of carbon dioxide (CO 2 ). These waters erode limestone rocks releasing calcium carbonate (CaCO 3 ) that will form calcium bicarbonate, according to the following reaction:
CaCO 3 (s) + CO 2 (g) + H 2 O (l) => Ca (HCO 3 ) 2 (aq)
This reaction occurs in caves where very hard waters originate. Calcium bicarbonate is not in solid state but in an aqueous solution, together with Ca 2+ , bicarbonate (HCO 3 – ) and carbonate ion (CO 3 2- ).
Subsequently, by decreasing the saturation of carbon dioxide in the water, the reverse reaction occurs, that is, the transformation of calcium bicarbonate into calcium carbonate:
Ca (HCO 3 ) 2 (aq) => CO 2 (g) + H 2 O (l) + CaCO 3 (s)
Calcium carbonate is poorly soluble in water, this causes its precipitation to occur as a solid. The above reaction is very important in the formation of stalactites, stalagmites and other speleothems in the caves.
These rocky structures are formed from the drops of water that fall from the ceiling of the caves (upper image). The CaCO 3 present in the water droplets crystallizes to form the mentioned structures.
The fact that calcium bicarbonate is not found in a solid state has made its use difficult, with few examples being found. Likewise, it is difficult to find information on its toxic effects. There is a report of a set of side effects from its use as a treatment to prevent osteoporosis.
Two HCO 3 – anions and a Ca 2+ cation interacting electrostatically in the upper image . According to the image, the Ca 2+ should be located in the middle, since this way the HCO 3 – would not repel each other due to their negative charges.
The negative charge on HCO 3 – is delocalized between two oxygen atoms, through resonance between the carbonyl group C = O and the bond C – O – ; while in CO 3 2– , it is delocalized between the three oxygen atoms, since the C – OH bond is deprotonated and can therefore receive a negative charge by resonance.
The geometries of these ions can be considered as spheres of calcium surrounded by flat triangles of carbonates with a hydrogenated end. In terms of size ratio, calcium is notably smaller than HCO 3 – ions .
Ca (HCO 3 ) 2 cannot form crystalline solids, and actually consists of aqueous solutions of this salt. In them, the ions are not alone, as in the image, but surrounded by H 2 O molecules .
How do they interact? Each ion is surrounded by a hydration sphere, which will depend on the metal, the polarity and the structure of the dissolved species.
Ca 2+ coordinates with the oxygen atoms in water to form an aqueous complex, Ca (OH 2 ) n 2+ , where n is generally considered to be six; that is, an “aqueous octahedron” around calcium.
While the anions HCO 3 – interact either with hydrogen bonds (O 2 CO – H — OH 2 ) or with the hydrogen atoms of water in the direction of the negative charge delocalizes (HOCO 2 – H – OH, dipole interaction- ion).
These interactions between Ca 2+ , HCO 3 – and water are so efficient that they make calcium bicarbonate very soluble in that solvent; Unlike CaCO 3 , in which the electrostatic attractions between Ca 2+ and CO 3 2– are very strong, precipitating from the aqueous solution.
Besides water, there are CO 2 molecules around, which react slowly to supply more HCO 3 – (depending on the pH values).
So far, the sizes and charges of the ions in Ca (HCO 3 ) 2 , nor the presence of water, explain why the solid compound does not exist; that is to say, pure crystals that can be characterized by X-ray crystallography. Ca (HCO 3 ) 2 is nothing more than ions present in the water from which the cavernous formations continue to grow.
If Ca 2+ and HCO 3 – could be isolated from the water avoiding the following chemical reaction:
Ca (HCO 3 ) 2 (aq) → CaCO 3 (s) + CO 2 (g) + H 2 O (l)
Then these could be grouped into a white crystalline solid with stoichiometric ratios 2: 1 (2HCO 3 / 1Ca). There are no studies about its structure, but it could be compared with that of NaHCO 3 (since magnesium bicarbonate, Mg (HCO 3 ) 2 , does not exist as a solid either), or with that of CaCO 3 .
Stability: NaHCO 3 vs Ca (HCO 3 ) 2
NaHCO 3 crystallizes in the monoclinic system, and CaCO 3 in the trigonal (calcite) and orthorhombic (aragonite) systems. If Na + were replaced by Ca 2+ , the crystal lattice would be destabilized by the greater difference in sizes; In other words, Na +, because it is smaller, forms a more stable crystal with HCO 3 – compared to Ca 2+ .
In fact, Ca (HCO 3 ) 2 (aq) needs water to evaporate so that its ions can group together in a crystal; but its crystal lattice is not strong enough to do so at room temperature. As the water is heated, the decomposition reaction occurs (equation above).
With the Na + ion in solution, it would form the crystal with the HCO 3 – before its thermal decomposition.
The reason then why Ca (HCO 3 ) 2 does not crystallize (theoretically) is due to the difference in ionic radii or sizes of its ions, which cannot form a stable crystal before decomposition.
Ca (HCO 3 ) 2 vs CaCO 3
If, on the other hand, H + were added to the CaCO 3 crystalline structures , their physical properties would drastically change. Perhaps, their melting points drop significantly, and even the morphologies of the crystals end up modified.
Would it be worth trying the synthesis of solid Ca (HCO 3 ) 2 ? Difficulties could exceed expectations, and a salt with low structural stability may not provide significant additional benefits in any application where other salts are already used.
Physical and chemical properties
Ca (HCO 3 ) 2
162.11 g / mol
It does not appear in solid state. It is found in aqueous solution and attempts to turn it into a solid by evaporation of water, have not been successful as it becomes calcium carbonate.
16.1 g / 100 ml at 0 ° C; 16.6 g / 100 ml at 20º C and 18.4 g / 100 ml at 100º C. These values are indicative of a high affinity of water molecules for Ca (HCO 3 ) 2 ions , as explained in the previous section. Meanwhile, only 15 mg of CaCO 3 dissolve in one liter of water, which reflects its strong electrostatic interactions.
Because Ca (HCO 3 ) 2 cannot form a solid, its solubility cannot be determined experimentally. However, given the conditions created by the CO 2 dissolved in the water surrounding the limestone, the mass of calcium dissolved at a temperature T could be calculated; mass, which would be equal to the concentration of Ca (HCO 3 ) 2 .
At different temperatures, the dissolved mass increases as shown by the values at 0, 20 and 100 ° C. Then, according to these experiments, it is determined how much of the Ca (HCO 3 ) 2 dissolves in the vicinity of CaCO 3 in an aqueous medium gasified with CO 2 . Once the gaseous CO 2 escapes , the CaCO 3 will precipitate, but not the Ca (HCO 3 ) 2 .
Melting and boiling points
The crystal lattice of Ca (HCO 3 ) 2 is much weaker than that of CaCO 3 . If it can be obtained in a solid state, and the temperature at which it melts is measured in a fusiometer, a value would surely be obtained well below 899ºC. Similarly, the same would be expected in determining the boiling point .
It is not combustible.
Since this compound does not exist in solid form, it is unlikely that it represents a risk to handle its aqueous solutions, since both Ca 2+ and HCO 3 ions – are not harmful at low concentrations; and therefore, the greater risk that would be to ingest these solutions, could only be due to a dangerous dose of calcium ingested.
If the compound were to form a solid, even though it may be physically different from CaCO 3 , its toxic effects may not go beyond simple discomfort and dryness after physical contact or by inhalation.
-Calcium bicarbonate solutions have long been used to wash old papers, especially works of art or historically important documents.
-The use of bicarbonate solutions is useful, not only because they neutralize the acids in the paper, but they also provide an alkaline reserve of calcium carbonate. The latter compound provides protection for future damage to the paper.
-Like other bicarbonates, it is used in chemical yeasts and in effervescent tablet or powder formulations. In addition, calcium bicarbonate is used as a food additive (aqueous solutions of this salt).
-Bicarbonate solutions have been used in the prevention of osteoporosis. However, side effects such as hypercalcemia, metabolic alkalosis, and kidney failure have been observed in one case.
-Calcium bicarbonate is occasionally administered intravenously to correct the depressive effect of hypokalemia on cardiac function.
-And finally, it provides calcium to the body, which is a mediator of muscle contraction, at the same time that it corrects acidosis that can occur in a hypokalemic condition.
- Wikipedia. (2018). Calcium bicarbonate. Taken from: en.wikipedia.org
- Sirah Dubois. (October 03, 2017). What Is Calcium Bicarbonate? Recovered from: livestrong.com
- Science Learning Hub. (2018). Carbonate chemistry. Recovered from: sciencelearn.org.nz
- PubChem. (2018). Calcium Bicarbonate. Recovered from: pubchem.ncbi.nlm.nih.gov
- Amy E. Gerbracht & Irene Brückle. (1997). The Use of Calcium Bicarbonate and Magnesium Bicarbonate Solutions in Small Conservation Workshops: Survey Results. Recovered from: cool.conservation-us.org