Calcium hydrogen carbonate

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Calcium ions are dissolved in (tap) water in the form of calcium hydrogen carbonate (also: calcium bicarbonate, empirical formula: Ca (HCO 3 ) 2 , theoretical molar mass : 162.11 g · mol −1 ). The substance itself cannot be represented as a pure substance under normal conditions. It is therefore not possible to determine material properties such as melting point and the like. Ä. to specify. Its solubility at 20 ° C and 1 atm is 1.66 g per 100 g H 2 O, which is significantly higher than that of calcium carbonate (1.4 mg per 100 g H 2 O at 20 ° C).

The balance with carbon dioxide

Calcium hydrogen carbonate is formed when limestone , which essentially consists of calcium carbonate, is weathered through the action of water and carbon dioxide . The carbon dioxide partially dissolves with water as carbonic acid , which further dissociates to hydrogen carbonate and thereby protonssupplies. At the same time, calcium carbonate dissolves in small quantities in water, releasing carbonate ions. These take over the hydrogen ions provided by the carbon dioxide and also become hydrogen carbonate ions. During the carbonic acid dissolution of limestone, one equivalent proportion of hydrogen carbonate ions comes from the carbonic acid, the other from the stone.

In order to keep calcium hydrogen carbonate in solution , a certain concentration of so-called "associated carbonic acid" is necessary. Chemically speaking, this is no different from any other carbonic acid; it's all about the proportion. In the dissociation equilibrium with the hydrogen carbonate ions present, this associated carbonic acid sets the pH value of the water just so low that the proportion of carbonate ions, which in turn depends on this pH value, together with the calcium concentration present, is the solubility product of Calcium carbonate does not yet exceed.

If there is more than the corresponding amount of free carbon dioxide in the solution, this amount of carbon dioxide is called “excess”. It can react with more limestone and loosen it. The proportion of it that dissolves further lime and goes into the additional calcium hydrogen carbonate is referred to as "lime-aggressive carbonic acid". The rest of the excess carbon dioxide increases the associated carbon dioxide to the new higher level.

Calcium hydrogen carbonate only exists in aqueous solution in the coexistence of equivalent amounts of calcium and hydrogen carbonate ions. When the water evaporates or when it is heated, the carbon dioxide escapes from the solution; it can also be diminished by photosynthesis. As a result, the dissociation equilibrium of the carbonic acid shifts again in the direction of the carbonate ions, i.e. to the left side of the above reaction equation. Thus the solubility product of the calcium carbonate is exceeded again and insoluble limestone is formed again. This process is the basis for the formation of lime sinter , travertine or tufa , but also the formation of lime sediments ( sea ​​chalk ) in lakes and oceans. Also, the occurring in the northern Alpenvorland Nagelfluh said conglomerate (sediments of molasses or glacial gravels) can thus occur by gravel or crushed by the binder lime is baked to a natural concrete.

Waterworks adapt the carbon dioxide content to the lime content of the drinking water in such a way that a thin layer of lime forms in iron pipes, which prevents rusting. In order not to reduce the line cross-section too much, the carbon dioxide content must be constantly adjusted. The speed with which the above equilibrium is established also plays a role in this constant adaptation: With pure calcium carbonate this takes an extraordinarily long time; in the presence of foreign ions (e.g. Mg 2+ or SO 4 2− ), on the other hand, the equilibrium is established very quickly; therefore dolomite is used to bind excess carbonic acid.

The calcium hydrogen carbonate content of the tap water makes the main contribution to the hardness of the water in the so-called carbonate hardness .

Individual evidence

  1. Entry on calcium carbonate. In: Römpp Online . Georg Thieme Verlag, accessed on October 22, 2013.
  2. ^ E. Schweda: Jander / Blasius: Inorganic Chemistry I - Introduction & Qualitative Analysis . 17th edition. Hirzel, 2012, ISBN 978-3-7776-2134-0 , pp. 253 .
  3. E. Riedel, Christoph Janiak: Inorganic Chemistry . 8th edition. de Gruyter, 2011, ISBN 3-11-022566-2 , p. 532 .
  4. E. Riedel, Christoph Janiak: Inorganic Chemistry . 8th edition. de Gruyter, 2011, ISBN 3-11-022566-2 , p. 610 .