"කැල්සියම් කාබනේට්" හි සංශෝධන අතර වෙනස්කම්

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[[ගොනුව:Calcite.GIF|thumb|right|Crystal structure of calcite]]
 
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'''කැල්සියම් කාබනේට්''' is a [[chemical compound]] with the [[chemical formula]] [[කැල්සියම්|Ca]][[Carbon|C]][[Oxygen|O]]<sub>3</sub>. It is a common substance found in [[Rock (geology)|rock]] in all parts of the world, and is the main component of [[seashells|shells of marine organisms]], [[snail]]s, [[pearls]], and [[eggshell]]s. Calcium carbonate is the active ingredient in [[agricultural lime]], and is usually the principal cause of [[hard water]]. It is commonly used medicinally as a [[කැල්සියම්]] supplement or as an [[antacid]], but excessive consumption can be hazardous.
 
'''කැල්සියම් කාබනේට්'''.
== රසායනික ගුණ ==
{{seealso|කාබනේට}}
කැල්සියම් කාබනේට් shares the typical properties of other carbonates. Notably:
* it reacts with strong acids, releasing carbon dioxide:
:CaCO<sub>3(s)</sub> + 2 HCl<sub>(aq)</sub> → CaCl<sub>2(aq)</sub> + CO<sub>2(g)</sub> + H<sub>2</sub>O<sub>(l)</sub>
* it releases carbon dioxide on heating (to above 840&nbsp;°C in the case of CaCO<sub>3</sub>), to form [[calcium oxide]], commonly called [[quicklime]], with reaction [[enthalpy]] 178 kJ / mole:
:CaCO<sub>3</sub> → CaO + CO<sub>2</sub>
 
කැල්සියම් කාබනේට් will react with water that is saturated with carbon dioxide to form the soluble [[calcium bicarbonate]].
:CaCO<sub>3</sub> + CO<sub>2</sub> + H<sub>2</sub>O → Ca(HCO<sub>3</sub>)<sub>2</sub>
 
This reaction is important in the [[erosion]] of [[carbonate rock]]s, forming [[cavern]]s, and leads to hard water in many regions.
 
== නිපැයීම ==
The vast majority of calcium carbonate used in industry is extracted by mining or quarrying. Pure calcium carbonate (e.g. for food or pharmaceutical use), can be produced from a pure quarried source (usually marble).
 
Alternatively, calcium carbonate is prepared by [[calcining]] crude [[calcium oxide]]. Water is added to give [[කැල්සියම් හයිඩ්‍රොක්සයිඩ්]], and [[carbon dioxide]] is passed through this solution to precipitate the desired calcium carbonate, referred to in the industry as precipitated calcium carbonate (PCC):<ref>{{cite web|title = Solvay Precipitated Calcium Carbonate: Production|publisher = Solvay S. A. |date=2007-03-09|accessdate = 2007-12-30|url = http://www.solvaypcc.com/safety_environment/0,0,1000044-_EN,00.html}}</ref>
 
: CaCO<sub>3</sub> → CaO + CO<sub>2</sub>
: CaO + H<sub>2</sub>O → Ca(OH)<sub>2</sub>
:Ca(OH)<sub>2</sub> + CO<sub>2</sub> → CaCO<sub>3</sub> + H<sub>2</sub>O
 
== Occurrence ==
Calcium carbonate is found naturally as the following [[mineral]]s in the form of [[Polymorphism (materials science)|polymorphs]]:
* [[Aragonite]]
* [[Calcite]]
* [[Vaterite]] or (μ-CaCO<sub>3</sub>)
 
The [[trigonal]] crystal structure of calcite is most common.
 
The calcium carbonate minerals occur in the following [[Rock (geology)|rocks]]:
* [[හුණුකූරු]]
* [[Limestone]]
* [[Marble]]
* [[Travertine]]
 
{{gallery
|Image:Calcite-HUGE.jpg|කැල්සයිට්
|Image:Aragonite.jpg|Aragonite
|Image:MarbleUSGOV.jpg|Marble
|Image:Kalktuff-Block Schloss-Tuebingen 2.jpg|Travertine
}}
 
== Geology ==
කාබනේට් is found frequently in geologic settings. It is found as a polymorph. A polymorph is a [[mineral]] with the same chemical formula but different chemical structure. [[Aragonite]], [[calcite]], [[limestone]], [[chalk]], [[marble]], [[travertine]], [[tufa]], and others all have CaCO<sub>3</sub> as their formula but each has a slightly different chemical structure. Calcite, as calcium carbonate is commonly referred to in geology is commonly talked about in marine settings. Calcite is typically found around the warm tropic environments. This is due to its chemistry and properties. Calcite is able to precipitate in warmer shallow environments than it does under colder environments because warmer environments do not favour the dissolution of CO<sub>2</sub>. This is analogous to CO<sub>2</sub> being dissolved in soda. When you take the cap off of a soda bottle, the CO<sub>2</sub> rushes out. As the soda warms up, [[carbon dioxide]] is released. This same principle can be applied to calcite in the ocean. Cold water carbonates do exist at higher latitudes but have a very slow growth rate.
 
In tropic settings, the waters are warm and clear. Consequently, you will see many more coral in this environment than you would towards the poles where the waters are cold. Calcium carbonate contributors such as corals, algae, and microorganisms are typically found in shallow water environments because as filter feeders they require sunlight to produce calcium carbonate.
 
=== Carbonate compensation depth ===
The [[carbonate compensation depth]] (CCD) is the point in the ocean where the rate of precipitation of calcium carbonate is balanced by the rate of dissolution due to the conditions present. Deep in the ocean, the temperature drops and pressure increases. Calcium carbonate is unusual in that its solubility increases with decreasing temperature. Increasing pressure also increases the solubility of calcium carbonate. The CCD can range from 4–6&nbsp;km below sea level.
 
== භාවිත ==
=== කාර්මික යෙදවුම් ===
The main use of කැල්සියම් කාබනේට් is in the construction industry, either as a building material in its own right (e.g. marble) or limestone aggregate for roadbuilding or as an ingredient of [[cement]] or as the starting material for the preparation of builder's lime by burning in a kiln.
 
කැල්සියම් කාබනේට් is also used in the purification of [[iron]] from [[iron ore]] in a [[blast furnace]]. කැල්සියම් කාබනේට් is calcined ''in situ'' to give calcium oxide, which forms a slag with various impurities present, and separates from the purified iron.<ref>{{cite web|title = Blast Furnace|publisher = Science Aid|accessdate = 2007-12-30|url = http://www.scienceaid.co.uk/chemistry/industrial/blastfurnace.html}}</ref>
 
කැල්සියම් කාබනේට් is also used in the [[oil industry]] in [[drilling fluid]]s as a formation bridging and filtercake sealing agent and may also be used as a weighting material to increase the density of [[drilling fluid]]s to control downhole pressures.
 
කැල්සියම් කාබනේට් is also one of the main sources used in growing [[Seacrete]], or [[Biorock]].
 
Precipitated කැල්සියම් කාබනේට්, pre-dispersed in slurry form, is also now widely used as filler material for latex gloves with the aim of achieving maximum saving in material and production costs.<ref name = precaco3>{{cite web|title = Precipitated කැල්සියම් කාබනේට් uses|url = http://www.aristocratholding.com/calris-5.html}}</ref>
 
කැල්සියම් කාබනේට් is widely used as an extender in paints,<ref name = reade>{{cite web|title = Calcium carbonate Powder|publisher = Reade Advanced Materials |date=2006-02-04|accessdate = 2007-12-30|url = http://www.reade.com/Products/Minerals_and_Ores/calcium_carbonate.html}}</ref> in particular matte emulsion paint where typically 30% by weight of the paint is either chalk or marble.
 
කැල්සියම් කාබනේට් is also widely used as a filler in plastics.<ref name = reade/> Some typical examples include around 15 to 20% loading of chalk in [[Polyvinyl chloride|unplasticized polyvinyl chloride]] (uPVC) drain pipe, 5 to 15% loading of stearate coated chalk or marble in uPVC window profile. [[Polyvinyl chloride|PVC]] cables can use කැල්සියම් කාබනේට් at loadings of up to 70 phr (parts per hundred parts of resin) to improve mechanical properties (tensile strength and elongation) and electrical properties (volume resistivity). [[Polypropylene]] compounds are often filled with කැල්සියම් කාබනේට් to increase rigidity, a requirement that becomes important at high use temperatures.<ref name= Imerys>{{cite web|url=http://www.imerys-perfmins.com/calcium-carbonate/eu/calcium-carbonate-plastic.htm |title=කැල්සියම් කාබනේට් in plastic applications |accessdate=2008-08-01 |publisher=Imerys Performance Minerals}}</ref> It also routinely used as a filler in [[Thermosetting plastic|thermosetting resins]] (Sheet and Bulk moulding compounds)<ref name = Imerys/> and has also been mixed with [[acrylonitrile butadiene styrene|ABS]], and other ingredients, to form some types of compression molded "clay" Poker chips.
 
Fine ground කැල්සියම් කාබනේට් is an essential ingredient in the microporous film used in babies' [[diapers]] and some building films as the pores are nucleated around the කැල්සියම් කාබනේට් particles during the manufacture of the film by biaxial stretching.
 
කැල්සියම් කාබනේට් is also used in a wide range of trade and [[do it yourself]] adhesives, sealants, and decorating fillers.<ref name = reade/> Ceramic tile adhesives typically contain 70 to 80% limestone. Decorating crack fillers contain similar levels of marble or dolomite. It is also mixed with putty in setting [[stained glass]] windows, and as a resist to prevent glass from sticking to kiln shelves when firing glazes and paints at high temperature.
 
කැල්සියම් කාබනේට් is known as ''whiting'' in [[ceramics (art)|ceramics]]/glazing applications,<ref name = reade/> where it is used as a common ingredient for many glazes in its white powdered form. When a glaze containing this material is fired in a kiln, the whiting acts as a [[flux]] material in the glaze.
 
Ground කැල්සියම් කාබනේට් (GCC) or Precipitated කැල්සියම් කාබනේට් (PCC) is used as a filler in paper. GCC and PCC are cheaper than wood fiber, so adding it to paper is cost efficient for the paper industry. Printing and writing paper can be made of 10 - 20% calcium cabonate.
 
In North America, කැල්සියම් කාබනේට් has begun to replace [[Kaolinite|kaolin]] in the production of glossy paper. Europe has been practicing this as alkaline [[papermaking]] or acid-free papermaking for some decades. Carbonates are available in forms: ground කැල්සියම් කාබනේට් (GCC) or precipitated කැල්සියම් කාබනේට් (PCC). The latter has a very fine and controlled particle size, on the order of 2 micrometres in diameter, useful in [[Coated paper|coatings for paper]].
 
It is used in swimming pools as a [[pH]] corrector for maintaining [[alkalinity]] "buffer" to offset the acidic properties of the disinfectant agent.
 
It is commonly called [[chalk]] as it has traditionally been a major component of blackboard chalk. Modern manufactured chalk is now mostly [[gypsum]], hydrated [[calcium sulfate]] CaSO<sub>4</sub>·2H<sub>2</sub>O.
 
Ground කැල්සියම් කාබනේට් is further used as an [[abrasive]] (both as scouring powder and as an ingredient of household scouring creams), in particular in its calcite form, which has the relatively low hardness level of 3 on the [[Mohs scale of mineral hardness]], and will therefore not scratch [[glass]] and most other [[ceramics]], [[Vitreous enamel|enamel]], [[bronze]], [[iron]], and [[steel]], and have a moderate effect on softer metals like [[aluminium]] and [[තඹ]].
 
=== Health and dietary applications ===
[[ගොනුව:500 mg calcium supplements with vitamin D.jpg|thumb|500 milligram calcium supplements made from කැල්සියම් කාබනේට්]]
කැල්සියම් කාබනේට් is widely used medicinally as an inexpensive dietary calcium supplement or [[antacid|gastric antacid]].<ref name = medline>{{cite web|publisher = [[National Institutes of Health]]|work = Medline Plus|title = කැල්සියම් කාබනේට් |date=2005-10-01|accessdate = 2007-12-30|url = http://www.nlm.nih.gov/medlineplus/druginfo/medmaster/a601032.html}}</ref> It may be used as a [[phosphate binder]] for the treatment of [[hyperphosphatemia]] (primarily in patients with [[chronic renal failure]]). It is also used in the pharmaceutical industry as an inert [[Excipients#Fillers and diluents|filler]] for [[tablet]]s and other pharmaceuticals.<ref>{{cite book|author = Herbert A. Lieberman, Leon Lachman, Joseph B. Schwartz|title = Pharmaceutical Dosage Forms: Tablets|year = 1990|isbn = 0824780442|page=153|publisher = Dekker|location = New York}}</ref>
 
කැල්සියම් කාබනේට් is used in the production of toothpaste and is also used in homeopathy as one of the constitutional remedies. Also, it has seen a resurgence as a food preservative and color retainer, when used in or with products such as organic apples or food<ref>http://chemistry.about.com/od/foodcookingchemistry/a/cadditives.htm</ref>.
Excess calcium from supplements, fortified food and high-calcium diets, can cause the "milk alkali syndrome," which has serious toxicity and can be fatal. In 1915, Bertram Sippy introduced the "Sippy regimen" of hourly ingestion of milk and cream, and the gradual addition of eggs and cooked cereal, for 10 days, combined with alkaline powders, which provided symptomatic relief for peptic ulcer disease. Over the next several decades, the Sippy regimen resulted in renal failure, alkalosis, and hypercalcemia, mostly in men with peptic ulcer disease. These adverse effects were reversed when the regimen stopped, but it was fatal in some patients with protracted vomiting. Milk alkali syndrome declined in men after effective treatments for peptic ulcer disease. During the past 15 years, it has been reported in women taking calcium supplements above the recommended range of 1.2 to 1.5&nbsp;g daily, for prevention and treatment of osteoporosis, and is exacerbated by dehydration. Calcium has been added to over-the-counter products, which contributes to inadvertent excessive intake. Excessive calcium intake can lead to [[hypercalcemia]], complications of which include vomiting, abdominal pain and altered mental status.<ref>{{cite journal|url=http://content.nejm.org/cgi/content/short/358/18/1952|title=Clinical problem-solving, back to basics|author=Ilan Gabriely et al.|journal=New England Journal of Medicine|year=2008|volume=358|page=1952|pmid=18450607|doi=10.1056/NEJMcps0706188|issue=18}}</ref>
 
As a [[food additive]] it is designated E170.<ref>{{cite web|title=Food-Info.net : E-numbers : E170 කැල්සියම් කාබනේට්|url=http://www.food-info.net/uk/e/e170.htm}} 080419 food-info.net</ref> It is used in some [[soy milk]] products as a source of dietary calcium; one study suggests that කැල්සියම් කාබනේට් might be as [[bioavailable]] as the calcium in cow's milk.<ref>{{cite journal|author = Y. Zhao, B. R. Martin and C. M. Weaver|title = Calcium Bioavailability of කැල්සියම් කාබනේට් Fortified Soymilk Is Equivalent to Cow's Milk in Young Women|year = 2005|journal = J. Nutr.|volume = 135|issue = 10|pages = 2379–2382|pmid = 16177199}}</ref> කැල්සියම් කාබනේට් is also used as a [[firming agent]] in many canned or bottled vegetable products.
 
=== Environmental applications ===
In 1989, a researcher, Ken Simmons, introduced CaCO<sub>3</sub> into the [[Whetstone Brook]] in Massachusetts.<ref>{{cite news|author = [[Associated Press]]|title =
Limestone Dispenser Fights Acid Rain in Stream |date=1989-06-13|url = http://query.nytimes.com/gst/fullpage.html?res=950DEFD9173FF930A25755C0A96F948260|publisher = [[New York Times]]}}</ref> His hope was that the කැල්සියම් කාබනේට් would counter the acid in the stream from acid rain and save the trout that had ceased to spawn. Although his experiment was a success, it did increase the amounts of aluminium ions in the area of the brook that was not treated with the limestone. This shows that CaCO<sub>3</sub> can be added to neutralize the effects of acid rain in [[river]] ecosystems. Currently කැල්සියම් කාබනේට් is used to neutralize acidic conditions in both soil and water.<ref>{{cite journal|author = R. K. Schreiber|title = Cooperative federal-state liming research on surface waters impacted by acidic deposition|year = 1988|journal =Water, Air, & Soil Pollution|volume = 41|issue = 1|pages = 53–73|url=http://www.springerlink.com/content/x71m4r306055p651/}}</ref><ref>{{cite web|title = Effects of low pH and high aluminum on Atlantic salmon smolts in Eastern Maine and liming project feasibility analysis|year = 2006|author = Dan Kircheis; Richard Dill|publisher = National Marine Fisheries Service and Maine Atlantic Salmon Commission|url = http://www.mainesalmonrivers.org/pages/Liming%20Project%20Rpt.pdf|format = reprinted at Downeast Salmon Federation}}</ref>. Since the 1970s, such ''liming'' has been practiced on a large scale in Sweden to mitigate acidification and several thousand lakes and streams are limed repeatedly.<ref>{{cite journal|author = M. Guhren, C. Bigler and I. Renberg|title = Liming placed in a long-term perspective: A paleolimnological study of 12 lakes in the Swedish liming program|year = 2007|journal =Journal of Paleolimnology|volume = 37|pages = 247–258|doi=10.1007/s10933-006-9014-9}}</ref>
 
== Calcination equilibrium ==
[[Calcination]] of limestone using charcoal fires to produce [[calcium oxide|quicklime]] has been practiced since antiquity by cultures all over the world. The temperature at which limestone yields calcium oxide is usually given as 825&nbsp;°C, but stating an absolute threshold is misleading. Calcium carbonate exists in equilibrium with calcium oxide and carbon dioxide at any temperature. At each temperature there is a [[partial pressure]] of carbon dioxide that is in equilibrium with calcium carbonate. At room temperature the equilibrium overwhelmingly favors calcium carbonate, because the equilibrium CO<sub>2</sub> pressure is only a tiny fraction of the partial CO<sub>2</sub> pressure in air, which is about 0.035 kPa.
 
At temperatures above 550&nbsp;°C the equilibrium CO<sub>2</sub> pressure begins to exceed the CO<sub>2</sub> pressure in air. So above 550&nbsp;°C, calcium carbonate begins to outgas CO<sub>2</sub> into air. But in a charcoal fired kiln, the concentration of CO<sub>2</sub> will be much higher than it is in air. Indeed if all the oxygen in the kiln is consumed in the fire, then the partial pressure of CO<sub>2</sub> in the kiln can be as high as 20 kPa.
 
The table shows that this equilibrium pressure is not achieved until the temperature is nearly 800&nbsp;°C. For the outgassing of CO<sub>2</sub> from calcium carbonate to happen at an economically useful rate, the equilibrium pressure must significantly exceed the ambient pressure of CO<sub>2</sub>. And for it to happen rapidly, the equilibrium pressure must exceed total atmospheric pressure of 101 kPa, which happens at 898&nbsp;°C.<br clear="right">
 
{| class=wikitable border=1
! {{chembox header}} colspan=18|Equilibrium pressure of CO<sub>2</sub> over CaCO<sub>3</sub> (P) vs temperature (T).<ref name=crc>{{RubberBible86th}}</ref>
|-
|'''P (kPa)'''||0.055||0.13||0.31||1.80||5.9||9.3||14||24||34||51||72 ||80||91||101||179||901||3961
|-
|'''T (°C)'''||550||587||605||680||727||748||777||800||830||852||871||881||891||898||937||1082||1241
|}
 
== ද්‍රාව්‍යතාව ==
=== With varying CO<sub>2</sub> pressure ===
{| class=wikitable border=1 align="right"
! {{chembox header}} colspan="3"|Calcium ion solubility as a function of<br /> [[carbon dioxide|CO<sub>2</sub>]] [[partial pressure]] at 25&nbsp;°C (''K''<sub>sp</sub> = 4.47×10<sup>−9</sup>)
|-
!<math>\scriptstyle P_{\text{CO}_2}</math> (atm)
![[pH]]
![Ca<sup>2+</sup>] (mol/L)
|-
|10<sup>−12</sup> ||12.0||5.19 × 10<sup>−3</sup>
|-
|10<sup>−10</sup> ||11.3||1.12 × 10<sup>−3</sup>
|-
|10<sup>−8</sup> ||10.7||2.55 × 10<sup>−4</sup>
|-
|10<sup>−6</sup> ||9.83||1.20 × 10<sup>−4</sup>
|-
|10<sup>−4</sup> ||8.62||3.16 × 10<sup>−4</sup>
|-
!3.5 × 10<sup>−4</sup>!!8.27!!4.70 × 10<sup>−4</sup>
|-
|10<sup>−3</sup> ||7.96||6.62 × 10<sup>−4</sup>
|-
|10<sup>−2</sup> ||7.30||1.42 × 10<sup>−3</sup>
|-
|10<sup>−1</sup> ||6.63||3.05 × 10<sup>−3</sup>
|-
|1 ||5.96||6.58 × 10<sup>−3</sup>
|-
|10 ||5.30||1.42 × 10<sup>−2</sup>
|}
 
කැල්සියම් කාබනේට් is poorly soluble in pure water (47&nbsp;mg/L at normal atmospheric CO<sub>2</sub> partial pressure as shown below).
 
The equilibrium of its solution is given by the equation (with dissolved කැල්සියම් කාබනේට් on the right):
:{| width="450"
| width="50%" height="30"| CaCO<sub>3</sub> {{eqm}} Ca<sup>2+</sup> + CO<sub>3</sub><sup>2–</sup>
| ''K''<sub>sp</sub> = 3.7×10<sup>−9</sup> to 8.7×10<sup>−9</sup> at 25&nbsp;°C
|}
 
where the [[solubility product]] for [Ca<sup>2+</sup>][CO<sub>3</sub><sup>2–</sup>] is given as anywhere from ''K''<sub>sp</sub> = 3.7×10<sup>−9</sup> to ''K''<sub>sp</sub> = 8.7×10<sup>−9</sup> at 25&nbsp;°C, depending upon the data source.<ref name = crc/><ref>{{cite web|title = Selected Solubility Products and Formation Constants at 25&nbsp;°C|publisher = [[California State University, Dominguez Hills]]|url = http://www.csudh.edu/oliver/chemdata/data-ksp.htm}}</ref> What the equation means is that the product of molar concentration of calcium ions ([[mole (unit)|moles]] of dissolved Ca<sup>2+</sup> per liter of solution) with the molar concentration of dissolved CO<sub>3</sub><sup>2–</sup> cannot exceed the value of ''K''<sub>sp</sub>. This seemingly simple solubility equation, however, must be taken along with the more complicated equilibrium of [[carbon dioxide]] with [[water]] (see [[carbonic acid]]). Some of the CO<sub>3</sub><sup>2–</sup> combines with H<sup>+</sup> in the solution according to:
 
:{| width="450"
| width="50%" height="25"| HCO<sub>3</sub><sup>–</sup> {{eqm}} H<sup>+</sup> + CO<sub>3</sub><sup>2–</sup> &nbsp;&nbsp;
| ''K''<sub>a2</sub> = 5.61×10<sup>−11</sup> at 25&nbsp;°C
|}
 
HCO<sub>3</sub><sup>–</sup> is known as the [[bicarbonate]] ion. [[Calcium bicarbonate]] is many times more soluble in water than කැල්සියම් කාබනේට්—indeed it exists ''only'' in solution.
 
Some of the HCO<sub>3</sub><sup>–</sup> combines with H<sup>+</sup> in solution according to:
 
:{| width="450"
| width="50%" height="25"|H<sub>2</sub>CO<sub>3</sub> {{eqm}} H<sup>+</sup> + HCO<sub>3</sub><sup>–</sup> &nbsp;&nbsp;
| ''K''<sub>a1</sub> = 2.5×10<sup>−4</sup> at 25&nbsp;°C
|}
 
Some of the H<sub>2</sub>CO<sub>3</sub> breaks up into water and dissolved carbon dioxide according to:
 
:{| width="450"
| width="50%" height="25"| H<sub>2</sub>O + CO<sub>2</sub>(dissolved) {{eqm}} H<sub>2</sub>CO<sub>3</sub> &nbsp;&nbsp;
| ''K''<sub>h</sub> = 1.70×10<sup>−3</sup> at 25&nbsp;°C
|}
 
And dissolved carbon dioxide is in equilibrium with atmospheric carbon dioxide according to:
 
:{| width="500"
| width="45%" |<math>\frac{P_{\text{CO}_2}}{[\text{CO}_2]}\ =\ k_\text{H}</math>
| where ''k''<sub>H</sub> = 29.76 atm/(mol/L) at 25&nbsp;°C ([[Henry's law|Henry constant]]), <math>\scriptstyle P_{\text{CO}_2}</math> being the CO<sub>2</sub> partial pressure.
|}
 
For ambient air, <math>\scriptstyle P_{\text{CO}_2}</math> is around 3.5×10<sup>−4</sup> atmospheres (or equivalently 35 [[Pascal (unit)|Pa]]). The last equation above fixes the concentration of dissolved CO<sub>2</sub> as a function of <math>\scriptstyle P_{\text{CO}_2}</math>, independent of the concentration of dissolved CaCO<sub>3</sub>. At atmospheric partial pressure of CO<sub>2</sub>, dissolved CO<sub>2</sub> concentration is 1.2×10<sup>–5</sub> moles/liter. The equation before that fixes the concentration of H<sub>2</sub>CO<sub>3</sub> as a function of [CO<sub>2</sub>]. For [CO<sub>2</sub>]=1.2×10<sup>–5</sub>, it results in [H<sub>2</sub>CO<sub>3</sub>]=2.0×10<sup>−8</sup> moles per liter. When [H<sub>2</sub>CO<sub>3</sub>] is known, the remaining three equations together with
 
:{| width="450"
| width="50%" height="25"| H<sub>2</sub>O {{eqm}} H<sup>+</sup> + OH<sup>–</sup>
| ''K'' = 10<sup>−14</sup> at 25&nbsp;°C
|}
 
(which is true for all aqueous solutions), and the fact that the solution must be electrically neutral,
 
:2[Ca<sup>2+</sup>] + [H<sup>+</sup>] = [HCO<sub>3</sub><sup>–</sup>] + 2[CO<sub>3</sub><sup>2–</sup>] + [OH<sup>–</sup>]
 
make it possible to solve simultaneously for the remaining five unknown concentrations (note that the above form of the neutrality equation is valid only if කැල්සියම් කාබනේට් has been put in contact with pure water or with a neutral pH solution; in the case where the origin water solvent pH is not neutral, the equation is modified).
 
[[ගොනුව:CanarySpring.jpg|thumb|right|Travertine කැල්සියම් කාබනේට් deposits from a [[hot spring]]]]
 
The table on the right shows the result for [Ca<sup>2+</sup>] and [H<sup>+</sup>] (in the form of pH) as a function of ambient partial pressure of CO<sub>2</sub> (''K''<sub>sp</sub> = 4.47×10<sup>−9</sup> has been taken for the calculation).
 
* At atmospheric levels of ambient CO<sub>2</sub> the table indicates the solution will be slightly alkaline with a maximum CaCO<sub>3</sub> solubility of 47&nbsp;mg/L.
 
* As ambient CO<sub>2</sub> partial pressure is reduced below atmospheric levels, the solution becomes more and more alkaline. At extremely low <math>\scriptstyle P_{\text{CO}_2}</math>, dissolved CO<sub>2</sub>, bicarbonate ion, and carbonate ion largely evaporate from the solution, leaving a highly alkaline solution of [[කැල්සියම් හයිඩ්‍රොක්සයිඩ්]], which is more soluble than CaCO<sub>3</sub>. Note that for <math>\scriptstyle P_{\text{CO}_2}</math> = 10<sup>−12</sup> atm, the [Ca<sup>2+</sup>][OH<sup>-</sup>]<sup>2</sup> product is still below the solubility product of Ca(OH)<sub>2</sub> (8×10<sup>−6</sup>). For still lower CO<sub>2</sub> pressure, Ca(OH)<sub>2</sub> precipitation will occur before CaCO<sub>3</sub> precipitation.
 
* As ambient CO<sub>2</sub> partial pressure increases to levels above atmospheric, pH drops, and much of the carbonate ion is converted to bicarbonate ion, which results in higher solubility of Ca<sup>2+</sup>.
 
The effect of the latter is especially evident in day to day life of people who have hard water. Water in aquifers underground can be exposed to levels of CO<sub>2</sub> much higher than atmospheric. As such water percolates through කැල්සියම් කාබනේට් rock, the CaCO<sub>3</sub> dissolves according to the second trend. When that same water then emerges from the tap, in time it comes into equilibrium with CO<sub>2</sub> levels in the air by outgassing its excess CO<sub>2</sub>. The කැල්සියම් කාබනේට් becomes less soluble as a result and the excess precipitates as lime scale. This same process is responsible for the formation of [[stalactites]] and [[stalagmite]]s in limestone caves.
 
Two hydrated phases of කැල්සියම් කාබනේට්, [[monohydrocalcite]], CaCO<sub>3</sub>·H<sub>2</sub>O and [[ikaite]], CaCO<sub>3</sub>·6H<sub>2</sub>O, may [[precipitate]] from water at ambient conditions and persist as metastable phases.
 
=== With varying pH ===
We now consider the problem of the maximum solubility of කැල්සියම් කාබනේට් in normal atmospheric conditions (<math>\scriptstyle P_{\mathrm{CO}_2}</math> = 3.5 × 10<sup>−4</sup> atm) when the pH of the solution is adjusted. This is for example the case in a swimming pool where the pH is maintained between 7 and 8 (by addition of [[sodium bisulfate]] NaHSO<sub>4</sub> to decrease the pH or of [[sodium bicarbonate]] NaHCO<sub>3</sub> to increase it). From the above equations for the solubility product, the hydration reaction and the two acid reactions, the following expression for the maximum [Ca<sup>2+</sup>] can be easily deduced:
:<math>[\text{Ca}^{2+}]_\text{max} = \frac{K_\text{sp}} {K_\text{h}K_\text{a1}K_\text{a2}k_\text{H}} \frac{[\text{H}^+]^2}{P_{\text{CO}_2}}</math>
showing a quadratic dependence in [H<sup>+</sup>]. The numerical application with the above values of the constants gives{{Citation needed|date=September 2009}}
 
{| border="1" cellspacing="0" cellpadding="4" style="margin: 0 0 0 0.5em; background: {{{bgc|white}}}; border-collapse: collapse; border-color: {{{bc|#C0C090}}};"
|-
|{{chembox header}} width="170" align="center" |'''pH'''
|7.0
|7.2
|7.4
|7.6
|7.8
|8.0
|8.2
|8.27
|8.4
|-
|{{chembox header}} width="170" align="center" |'''[Ca<sup>2+</sup>]<sub>max</sub> (10<sup>−6</sup>mol/L)'''
|180
|71.7
|28.5
|11.4
|4.52
|1.80
|0.717
|0.519
|0.285
|-
|{{chembox header}} width="170" align="center"|'''[Ca<sup>2+</sup>]<sub>max</sub> (mg/L)'''
|7.21
|2.87
|1.14
|0.455
|0.181
|0.0721
|0.0287
|0.0208
|0.0114
|}
Comments:
* decreasing the pH from 8 to 7 increases the maximum Ca<sup>2+</sup> concentration by a factor 100. Water with a pH maintained to 7 can dissolve up to 15.9 g/L of CaCO<sub>3</sub>. This explains the high Ca<sup>2+</sup> concentration in some mineral waters with pH close to 7.
* note that the Ca<sup>2+</sup> concentration of the previous table is recovered for pH = 8.27
* keeping the pH to 7.4 in a swimming pool (which gives optimum HClO/ClO<sup>−</sup> [[hypochlorous acid|ratio]] in the case of "chlorine" maintenance) results in a maximum Ca<sup>2+</sup> concentration of 1010&nbsp;mg/L. This means that successive cycles of water evaporation and partial renewing may result in a very [[hard water]] before CaCO<sub>3</sub> precipitates (water with a Ca<sup>2+</sup> concentration above 120&nbsp;mg/L is considered very hard). Addition of a calcium [[chelation|sequestring agent]] or complete renewing of the water will solve the problem.
 
=== Solubility in a strong or weak acid solution ===
Solutions of [[strong acid|strong]] ([[hydrochloric acid|HCl]]), moderately strong ([[sulfamic acid|sulfamic]]) or [[weak acid|weak]] ([[acetic acid|acetic]], [[citric acid|citric]], [[sorbic acid|sorbic]], [[lactic acid|lactic]], [[phosphoric acid|phosphoric]]) acids are commercially available. They are commonly used as [[descaling agent]]s to remove [[limescale]] deposits. The maximum amount of CaCO<sub>3</sub> that can be "dissolved" by one liter of an acid solution can be calculated using the above equilibrium equations.
* In the case of a strong monoacid with decreasing acid concentration [A] = [A<sup>-</sup>], we obtain (with CaCO<sub>3</sub> molar mass = 100 g):
 
{| border="1" cellspacing="0" cellpadding="4" style="margin: 0 0 0 0.5em; background: {{{bgc|white}}}; border-collapse: collapse; border-color: {{{bc|#C0C090}}};"
|-
|width="160" {{chembox header}} |'''[A] (mol/L)'''
|1
|10<sup>−1</sup>
|10<sup>−2</sup>
|10<sup>−3</sup>
|10<sup>−4</sup>
|10<sup>−5</sup>
|10<sup>−6</sup>
|10<sup>−7</sup>
|10<sup>−10</sup>
|-
|width="160" {{chembox header}} |'''Initial pH'''
|0.00||1.00||2.00||3.00||4.00||5.00||6.00||6.79||7.00
|-
|width="160" {{chembox header}} |'''Final pH'''
|6.75||7.25||7.75||8.14||8.25||8.26||8.26||8.26||8.27
|-
|width="160" {{chembox header}} |'''Dissolved CaCO<sub>3</sub> (g per liter of acid)'''
|50.0||5.00||0.514||0.0849||0.0504||0.0474||0.0471||0.0470||0.0470
|}
 
where the initial state is the acid solution with no Ca<sup>2+</sup> (not taking into account possible CO<sub>2</sub> dissolution) and the final state is the solution with saturated Ca<sup>2+</sup>. For strong acid concentrations, all species have a negligible concentration in the final state with respect to Ca<sup>2+</sup> and A<sup>-</sup> so that the neutrality equation reduces approximately to 2[Ca<sup>2+</sup>] = [A<sup>-</sup>] yielding <math>\scriptstyle[\mathrm{Ca}^{2+}] \simeq \frac{[\mathrm{A}^-]}{2}</math>. When the concentration decreases, [HCO<sub>3</sub><sup>-</sup>] becomes non negligible so that the preceding expression is no longer valid. For vanishing acid concentrations, we recover the final pH and the solubility of CaCO<sub>3</sub> in pure water.
 
* In the case of a weak monoacid (here we take acetic acid with p''K''<sub>A</sub> = 4.76) with decreasing total acid concentration [A] = [A<sup>-</sup>]+[AH], we obtain:
 
{| border="1" cellspacing="0" cellpadding="4" style="margin: 0 0 0 0.5em; background: {{{bgc|white}}}; border-collapse: collapse; border-color: {{{bc|#C0C090}}};"
|-
|width="160" {{chembox header}} |'''[A] (mol/L)'''
|1
|10<sup>−1</sup>
|10<sup>−2</sup>
|10<sup>−3</sup>
|10<sup>−4</sup>
|10<sup>−5</sup>
|10<sup>−6</sup>
|10<sup>−7</sup>
|10<sup>−10</sup>
|-
|width="160" {{chembox header}} |'''Initial pH'''
|2.38||2.88||3.39||3.91||4.47||5.15||6.02||6.79||7.00
|-
|width="160" {{chembox header}} |'''Final pH'''
|6.75||7.25||7.75||8.14||8.25||8.26||8.26||8.26||8.27
|-
|width="160" {{chembox header}} |'''Dissolved CaCO<sub>3</sub> (g per liter of acid)'''
|49.5||4.99||0.513||0.0848||0.0504||0.0474||0.0471||0.0470||0.0470
|}
We see that for the same total acid concentration, the initial pH of the weak acid is less acid than the one of the strong acid; however, the maximum amount of CaCO<sub>3</sub> which can be dissolved is approximately the same. This is because in the final state, the pH is larger than the p''K''<sub>A</sub>, so that the weak acid is almost completely dissociated, yielding in the end as many H<sup>+</sup> ions as the strong acid to "dissolve" the කැල්සියම් කාබනේට්.
 
* The calculation in the case of [[phosphoric acid]] (which is the most widely used for domestic applications) is more complicated since the concentrations of the four dissociation states corresponding to this acid must be calculated together with [HCO<sub>3</sub><sup>-</sup>], [CO<sub>3</sub><sup>2-</sup>], [Ca<sup>2+</sup>], [H<sup>+</sup>] and [OH<sup>-</sup>]. The system may be reduced to a seventh degree equation for [H<sup>+</sup>] the numerical solution of which gives
 
{| border="1" cellspacing="0" cellpadding="4" style="margin: 0 0 0 0.5em; background: {{{bgc|white}}}; border-collapse: collapse; border-color: {{{bc|#C0C090}}};"
|-
|width="160" {{chembox header}} |'''[A] (mol/L)'''
|1
|10<sup>−1</sup>
|10<sup>−2</sup>
|10<sup>−3</sup>
|10<sup>−4</sup>
|10<sup>−5</sup>
|10<sup>−6</sup>
|10<sup>−7</sup>
|10<sup>−10</sup>
|-
|width="160" {{chembox header}} |'''Initial pH'''
|1.08||1.62||2.25||3.05||4.01||5.00||5.97||6.74||7.00
|-
|width="160" {{chembox header}} |'''Final pH'''
|6.71||7.17||7.63||8.06||8.24||8.26||8.26||8.26||8.27
|-
|width="160" {{chembox header}} |'''Dissolved CaCO<sub>3</sub> (g per liter of acid)'''
|62.0||7.39||0.874||0.123||0.0536||0.0477||0.0471||0.0471||0.0470
|}
 
where [A] = [H<sub>3</sub>PO<sub>4</sub>] + [H<sub>2</sub>PO<sub>4</sub><sup>-</sup>] + [HPO<sub>4</sub><sup>2-</sup>] + [PO<sub>4</sub><sup>3-</sup>] is the total acid concentration. We see that phosphoric acid is more efficient than a monoacid since at the final almost neutral pH, the second dissociated state concentration [HPO<sub>4</sub><sup>2-</sup>] is not negligible (see [[phosphoric acid#pH and composition of a phosphoric acid aqueous solution|phosphoric acid]]).
 
== මේවාත් බලන්න ==
{{colbegin|3}}
* [[Cuttlebone]]
* [[Cuttlefish]]
* [[Gesso]]
* [[Limescale]]
* [[Marble]]
* [[Ocean acidification]]
{{colend}}
 
== References ==
{{reflist|2}}
 
== බාහිර සබැඳුම් ==
* {{ICSC|1193|11}}
* {{PubChemLink|516889}}
* [[ATC codes]]: {{ATC|A02|AC01}} and {{ATC|A12|AA04}}
* [http://calcium-carbonate.org.uk/calcium-carbonate.asp The British Calcium Carbonate Association – What is calcium carbonate]
 
{{Calcium compounds}}
{{Antacids}}
 
{{DEFAULTSORT:Calcium Carbonate}}
"https://si.wikipedia.org/wiki/කැල්සියම්_කාබනේට්" වෙතින් සම්ප්‍රවේශනය කෙරිණි