"කැල්සියම් කාබනේට්" හි සංශෝධන අතර වෙනස්කම්

|ExactMass = 100.0869 g/mol

|Appearance = Fine white powder

|SolubilityProduct = 4.8{{e|-9}}<ref>{{cite book|last =Patnaik|first=Pradyot|year=2003|title=Handbook of Inorganic Chemical Compounds |publisher=McGraw-Hill|pages=|isbn =0070494398|url=http://books.google.com/?id=Xqj-TTzkvTEC|accessdate=2009-06-06}}</ref>

|Solvent = dilute acids

|SolubleOther = soluble

|BoilingPt = decomposes

|RefractIndex = 1.59

|Section8 = {{Chembox Related

|OtherAnions = [[Calcium bicarbonate]]

|OtherCpds = [[Calcium sulfate]]

}}

}}

 

'''කැල්සියම් කාබනේට්''' is a [[chemical compound]] with the [[chemical formula]] [[Calcium|Ca]][[Carbon|C]][[Oxygen|O]]₃. 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 [[calcium]] supplement or as an [[antacid]], but excessive consumption can be hazardous.

 

{{seealso|කාබනේට}}

කැල්සියම් කාබනේට් shares the typical properties of other carbonates. Notably:

:CaCO_3(s) + 2 HCl_(aq) → CaCl_2(aq) + CO_2(g) + H₂O_(l)

:CaCO₃ → CaO + CO₂

 

This reaction is important in the [[erosion]] of [[carbonate rock]]s, forming [[cavern]]s, and leads to hard water in many regions.

 

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).

 

:Ca(OH)₂ + CO₂ → CaCO₃ + H₂O

 

 

The [[trigonal]] crystal structure of calcite is most common.

 

The calcium carbonate minerals occur in the following [[Rock (geology)|rocks]]:

 

{{gallery

}}

 

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 CaCO₃ 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 CO₂. This is analogous to CO₂ being dissolved in soda. When you take the cap off of a soda bottle, the CO₂ 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.

 

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.

 

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.

 

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 [[copper]].

 

 

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>.

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.

 

In 1989, a researcher, Ken Simmons, introduced CaCO₃ 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 CaCO₃ 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>

 

 

 

 

{| class=wikitable border=1

|}

 

{| class=wikitable border=1 align="right"

|-

!<math>\scriptstyle P_{\text{CO}_2}</math> (atm)

:{| width="450"

| width="50%" height="30"| CaCO₃ {{eqm}} Ca²⁺ + CO₃^2–

|}

 

 

:{| width="450"

| width="50%" height="25"| HCO₃^– {{eqm}} H⁺ + CO₃^2– &nbsp;&nbsp;

|}

 

:{| width="450"

| width="50%" height="25"|H₂CO₃ {{eqm}} H⁺ + HCO₃^– &nbsp;&nbsp;

|}

 

:{| width="450"

| width="50%" height="25"| H₂O + CO₂(dissolved) {{eqm}} H₂CO₃ &nbsp;&nbsp;

|}

 

:{| width="500"

| width="45%" |<math>\frac{P_{\text{CO}_2}}{[\text{CO}_2]}\ =\ k_\text{H}</math>

|}

 

:{| width="450"

| width="50%" height="25"| H₂O {{eqm}} H⁺ + OH^–

|}

 

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).

 

 

The table on the right shows the result for [Ca²⁺] and [H⁺] (in the form of pH) as a function of ambient partial pressure of CO₂ (''K''_sp = 4.47×10⁻⁹ has been taken for the calculation).

 

 

 

 

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 CO₂ much higher than atmospheric. As such water percolates through කැල්සියම් කාබනේට් rock, the CaCO₃ dissolves according to the second trend. When that same water then emerges from the tap, in time it comes into equilibrium with CO₂ levels in the air by outgassing its excess CO₂. 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]], CaCO₃·H₂O and [[ikaite]], CaCO₃·6H₂O, may [[precipitate]] from water at ambient conditions and persist as metastable phases.

 

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 × 10⁻⁴ 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]] NaHSO₄ to decrease the pH or of [[sodium bicarbonate]] NaHCO₃ 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 [Ca²⁺] 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>

|}

Comments:

 

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 CaCO₃ that can be "dissolved" by one liter of an acid solution can be calculated using the above equilibrium equations.

 

{| border="1" cellspacing="0" cellpadding="4" style="margin: 0 0 0 0.5em; background: {{{bgc|white}}}; border-collapse: collapse; border-color: {{{bc|#C0C090}}};"

where the initial state is the acid solution with no Ca²⁺ (not taking into account possible CO₂ dissolution) and the final state is the solution with saturated Ca²⁺. For strong acid concentrations, all species have a negligible concentration in the final state with respect to Ca²⁺ and A^- so that the neutrality equation reduces approximately to 2[Ca²⁺] = [A^-] yielding <math>\scriptstyle[\mathrm{Ca}^{2+}] \simeq \frac{[\mathrm{A}^-]}{2}</math>. When the concentration decreases, [HCO₃^-] 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 CaCO₃ in pure water.

 

 

{| border="1" cellspacing="0" cellpadding="4" style="margin: 0 0 0 0.5em; background: {{{bgc|white}}}; border-collapse: collapse; border-color: {{{bc|#C0C090}}};"

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 CaCO₃ which can be dissolved is approximately the same. This is because in the final state, the pH is larger than the p''K''_A, so that the weak acid is almost completely dissociated, yielding in the end as many H⁺ ions as the strong acid to "dissolve" the කැල්සියම් කාබනේට්.

 

 

{| border="1" cellspacing="0" cellpadding="4" style="margin: 0 0 0 0.5em; background: {{{bgc|white}}}; border-collapse: collapse; border-color: {{{bc|#C0C090}}};"

where [A] = [H₃PO₄] + [H₂PO₄^-] + [HPO₄^2-] + [PO₄^3-] 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 [HPO₄^2-] is not negligible (see [[phosphoric acid#pH and composition of a phosphoric acid aqueous solution|phosphoric acid]]).

 

{{colbegin|3}}

{{colend}}

 

{{reflist|2}}

 

 

{{Calcium compounds}}

 

{{DEFAULTSORT:Calcium Carbonate}}

 

[[ar:كربونات الكالسيوم]]

[[bg:Калциев карбонат]]

[[bn:চুনাপাথর]]

[[ca:Carbonat de calci]]

[[cs:Uhličitan vápenatý]]