Is Caco3 A Binary Compound

Chemical chemical compound

Calcium carbonate

Names
IUPAC name Calcium carbonate
Other names Aragonite; calcite; chalk; lime; limestone; marble; oyster; pearl
Identifiers
CAS Number	471-34-1 Y
3D model (JSmol)	Interactive image Interactive image
ChEBI	CHEBI:3311 Y
ChEMBL	ChEMBL1200539 N
ChemSpider	9708 Y
DrugBank	DB06724
ECHA InfoCard	100.006.765
EC Number	207-439-nine
Due east number	E170 (colours)
KEGG	D00932 Y
PubChem CID	10112
RTECS number	FF9335000
UNII	H0G9379FGK Y
CompTox Dashboard (EPA)	DTXSID3036238
InChI InChI=1S/CH2O3.Ca/c2-1(3)iv;/h(H2,2,3,four);/q;+2/p-two Y Key: VTYYLEPIZMXCLO-UHFFFAOYSA-L Y InChI=1/CH2O3.Ca/c2-1(3)four;/h(H2,2,3,4);/q;+2/p-2 Key: VTYYLEPIZMXCLO-NUQVWONBAS
SMILES [Ca+two].[O-]C([O-])=O C(=O)([O-])[O-].[Ca+2]
Properties
Chemical formula	CaCO3
Molar mass	100.0869 g/mol
Appearance	Fine white powder; chalky taste
Aroma	odorless
Density	ii.711 thousand/cm3 (calcite) 2.83 g/cm3 (aragonite)
Melting betoken	1,339 °C (two,442 °F; 1,612 G) (calcite) 825 °C (1,517 °F; i,098 M) (aragonite)[4] [v]
Boiling point	decomposes
Solubility in water	0.013 g/50 (25 °C)[1] [2]
Solubility product (K sp)	3.iii×10−nine [3]
Solubility in dilute acids	soluble
Acidity (pK a)	9.0
Magnetic susceptibility (χ)	−iii.82×10−5 cmiii/mol
Refractive alphabetize (n D)	1.59
Structure
Crystal structure	Trigonal
Space group	32/m
Thermochemistry
Std molar entropy (S ⦵ 298)	93 J·mol−i·Thousand−1 [vi]
Std enthalpy of formation (Δf H ⦵ 298)	−1207 kJ·mol−1 [six]
Pharmacology
ATC code	A02AC01 (WHO) A12AA04 (WHO)
Hazards
NFPA 704 (burn diamond)	0 0 0
Lethal dose or concentration (LD, LC):
LD50 (median dose)	6450 mg/kg (oral, rat)
NIOSH (United states health exposure limits):
PEL (Permissible)	TWA 15 mg/kthree (total) TWA 5 mg/yardthree (resp)[vii]
Safety data sheet (SDS)	ICSC 1193
Related compounds
Other anions	Calcium bicarbonate
Other cations	Beryllium carbonate Magnesium carbonate Strontium carbonate Barium carbonate Radium carbonate
Related compounds	Calcium sulfate
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa). Nverify (what is Y N  ?) Infobox references

Chemical compound

Crystal structure of calcite

Calcium carbonate is a chemical compound with the formula CaCO₃ . Information technology is a common substance plant in rocks every bit the minerals calcite and aragonite (well-nigh notably as limestone, which is a type of sedimentary stone consisting mainly of calcite) and is the main component of eggshells, gastropod shells, shellfish skeletons and pearls. Calcium carbonate is the active ingredient in agricultural lime and is created when calcium ions in hard h2o react with carbonate ions to create limescale. Information technology has medical employ as a calcium supplement or as an antacid, but excessive consumption tin exist chancy and cause hypercalcemia and digestive issues.^[8]

Chemistry [edit]

Calcium carbonate shares the typical properties of other carbonates. Notably it

• reacts with acids, releasing carbon dioxide (technically speaking, carbonic acid, but that disintegrates quickly to CO₂ and H₂O):

CaCO_iii(s) + ii H⁺(aq) → Ca²⁺(aq) + CO₂(g) + H₂O(l)

• releases carbon dioxide upon heating, called a thermal decomposition reaction, or calcination (to above 840 °C in the case of CaCO_three ), to grade calcium oxide, CaO, commonly called quicklime, with reaction enthalpy 178 kJ/mol:

${\displaystyle {\ce {CaCO3(s)->[\Delta ]CaO(s){+}CO2\uparrow }}}$

Calcium carbonate reacts with water that is saturated with carbon dioxide to course the soluble calcium bicarbonate.

${\displaystyle {\ce {CaCO3(s){+}CO2(g){+}H2o(l)-> Ca(HCO3)2(aq)}}}$

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

An unusual form of calcium carbonate is the hexahydrate ikaite, CaCO₃·6H₂O. Ikaite is stable only below 8 °C.

Grooming [edit]

The vast majority of calcium carbonate used in industry is extracted by mining or quarrying. Pure calcium carbonate (such as for nutrient or pharmaceutical use), can exist produced from a pure quarried source (ordinarily marble).

Alternatively, calcium carbonate is prepared from calcium oxide. H2o is added to requite calcium hydroxide and so carbon dioxide is passed through this solution to precipitate the desired calcium carbonate, referred to in the industry as precipitated calcium carbonate (PCC) This procedure is called carbonatation:^[9]

${\displaystyle {\ce {CaO{+}H2O->Ca(OH)2}}}$

${\displaystyle {\ce {Ca(OH)2{+}CO2->CaCO3{+}H_2O}}}$

Structure [edit]

The thermodynamically stable form of CaCO₃ nether normal conditions is hexagonal β-CaCO_three (the mineral calcite).^[10] Other forms tin can be prepared, the denser (ii.83 g/cmⁱⁱⁱ) orthorhombic λ-CaCO_three (the mineral aragonite) and hexagonal μ-CaCO_iii , occurring every bit the mineral vaterite.^[x] The aragonite form tin be prepared by precipitation at temperatures above 85 °C; the vaterite class tin exist prepared by precipitation at 60 °C.^[10] Calcite contains calcium atoms coordinated by six oxygen atoms; in aragonite they are coordinated by nine oxygen atoms.^[10] The vaterite construction is non fully understood.^[11] Magnesium carbonate (MgCO₃ ) has the calcite structure, whereas strontium carbonate (SrCO₃ ) and barium carbonate (BaCO₃ ) adopt the aragonite structure, reflecting their larger ionic radii.^[10]

Occurrence [edit]

Geological sources [edit]

Calcite, aragonite and vaterite are pure calcium carbonate minerals. Industrially of import source rocks which are predominantly calcium carbonate include limestone, chalk, marble and travertine.

Biological sources [edit]

Eggshells, snail shells and most seashells are predominantly calcium carbonate and tin can be used as industrial sources of that chemical.^[13] Oyster shells accept enjoyed recent recognition as a source of dietary calcium, but are also a practical industrial source.^{[fourteen] [15]} Dark greenish vegetables such as broccoli and kale contain dietarily meaning amounts of calcium carbonate, but they are non practical equally an industrial source.^[16]

[edit]

Beyond Globe, strong testify suggests the presence of calcium carbonate on Mars. Signs of calcium carbonate take been detected at more than than one location (notably at Gusev and Huygens craters). This provides some testify for the past presence of liquid water.^{[17] [eighteen]}

Geology [edit]

Surface precipitation of

CaCO₃ equally tufa in Rubaksa, Federal democratic republic of ethiopia

Carbonate is found frequently in geologic settings and constitutes an enormous carbon reservoir. Calcium carbonate occurs as aragonite, calcite and dolomite as meaning constituents of the calcium wheel. The carbonate minerals course the rock types: limestone, chalk, marble, travertine, tufa, and others.

In warm, clear tropical waters corals are more abundant than towards the poles where the waters are cold. Calcium carbonate contributors, including plankton (such equally coccoliths and planktic foraminifera), coralline algae, sponges, brachiopods, echinoderms, bryozoa and mollusks, are typically found in shallow water environments where sunlight and filterable food are more abundant. Cold-water carbonates do be at higher latitudes but accept a very tiresome growth rate. The calcification processes are changed by ocean acidification.

Where the oceanic crust is subducted nether a continental plate sediments will be carried downward to warmer zones in the asthenosphere and lithosphere. Under these weather calcium carbonate decomposes to produce carbon dioxide which, along with other gases, give rise to explosive volcanic eruptions.

Carbonate bounty depth [edit]

The carbonate bounty depth (CCD) is the point in the bounding main where the rate of precipitation of calcium carbonate is balanced by the rate of dissolution due to the weather condition present. Deep in the ocean, the temperature drops and pressure increases. Calcium carbonate is unusual in that its solubility increases with decreasing temperature.^[nineteen] Increasing force per unit area also increases the solubility of calcium carbonate. The carbonate compensation depth can range from 4,000 to 6,000 meters beneath sea level.

Office in taphonomy [edit]

Calcium carbonate can preserve fossils through permineralization. Most of the vertebrate fossils of the Ii Medicine Germination—a geologic formation known for its duck-billed dinosaur eggs—are preserved by CaCO_iii permineralization.^[20] This type of preservation conserves high levels of particular, even down to the microscopic level. However, information technology also leaves specimens vulnerable to weathering when exposed to the surface.^[20]

Trilobite populations were once thought to accept equanimous the majority of aquatic life during the Cambrian, due to the fact that their calcium carbonate-rich shells were more easily preserved than those of other species,^[21] which had purely chitinous shells.

Uses [edit]

Construction [edit]

The master use of calcium carbonate is in the construction industry, either equally a building cloth, or limestone amass for road building, as an ingredient of cement, or as the starting cloth for the training of builders' lime by burning in a kiln. Nevertheless, because of weathering mainly caused by acid rain,^[22] calcium carbonate (in limestone form) is no longer used for building purposes on its own, just only as a raw primary substance for building materials.

Calcium carbonate is besides used in the purification of iron from iron ore in a blast furnace. The carbonate is calcined in situ to give calcium oxide, which forms a slag with diverse impurities present, and separates from the purified atomic number 26.^[23]

In the oil industry, calcium carbonate is added to drilling fluids equally a germination-bridging and filtercake-sealing amanuensis; it is too a weighting textile which increases the density of drilling fluids to control the downhole pressure. Calcium carbonate is added to swimming pools, equally a pH corrector for maintaining alkalinity and offsetting the acidic backdrop of the disinfectant amanuensis.^[24]

It is also used as a raw fabric in the refining of sugar from sugar beet; information technology is calcined in a kiln with anthracite to produce calcium oxide and carbon dioxide. This burnt lime is then slaked in fresh water to produce a calcium hydroxide suspension for the precipitation of impurities in raw juice during carbonatation.^[25]

Calcium carbonate in the form of chalk has traditionally been a major component of blackboard chalk. However, modern manufactured chalk is mostly gypsum, hydrated calcium sulfate CaSO_iv·2H_iiO. Calcium carbonate is a main source for growing biorock. Precipitated calcium carbonate (PCC), pre-dispersed in slurry form, is a mutual filler material for latex gloves with the aim of achieving maximum saving in material and production costs.^[26]

Fine basis calcium carbonate (GCC) is an essential ingredient in the microporous film used in diapers and some edifice films, as the pores are nucleated effectually the calcium carbonate particles during the manufacture of the film past biaxial stretching. GCC and PCC are used as a filler in newspaper because they are cheaper than wood cobweb. In terms of market place volume, GCC are the almost of import types of fillers currently used.^[27] Printing and writing paper can comprise 10–20% calcium carbonate. In North America, calcium carbonate has begun to replace kaolin in the product of glossy paper. Europe has been practicing this as alkali metal papermaking or acid-costless papermaking for some decades. PCC used for paper filling and paper coatings is precipitated and prepared in a variety of shapes and sizes having characteristic narrow particle size distributions and equivalent spherical diameters of 0.4 to 3 micrometers.^{[ citation needed ]}

Calcium carbonate is widely used equally an extender in paints,^[28] in particular matte emulsion paint where typically thirty% past weight of the pigment is either chalk or marble. Information technology is also a popular filler in plastics.^[28] Some typical examples include around 15 to twenty% loading of chalk in unplasticized polyvinyl chloride (uPVC) drainpipes, 5% to xv% loading of stearate-coated chalk or marble in uPVC window profile. PVC cables tin use calcium carbonate at loadings of upward to seventy phr (parts per hundred parts of resin) to improve mechanical properties (tensile forcefulness and elongation) and electrical properties (volume resistivity).^{[ commendation needed ]} Polypropylene compounds are often filled with calcium carbonate to increase rigidity, a requirement that becomes important at high usage temperatures.^[29] Hither the pct is often 20–forty%. It also routinely used equally a filler in thermosetting resins (canvass and bulk molding compounds)^[29] and has likewise been mixed with ABS, and other ingredients, to form some types of compression molded "clay" poker chips.^[xxx] Precipitated calcium carbonate, made by dropping calcium oxide into h2o, is used past itself or with additives as a white paint, known as whitewashing.^{[31] [32]}

Calcium carbonate is added to a wide range of trade and exercise it yourself adhesives, sealants, and decorating fillers.^[28] Ceramic tile adhesives typically incorporate seventy% to lxxx% limestone. Decorating cleft fillers incorporate similar levels of marble or dolomite. It is also mixed with putty in setting stained glass windows, and every bit a resist to foreclose glass from sticking to kiln shelves when firing glazes and paints at loftier temperature.^{[33] [34] [35] [36]}

In ceramic glaze applications, calcium carbonate is known as whiting,^[28] and is 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. Basis calcium carbonate is an abrasive (both as scouring powder and equally an ingredient of household scouring creams), in particular in its calcite form, which has the relatively depression hardness level of three on the Mohs scale, and will therefore not scratch glass and most other ceramics, enamel, bronze, iron, and steel, and have a moderate effect on softer metals like aluminium and copper. A paste made from calcium carbonate and deionized h2o can be used to clean tarnish on silvery.^[37]

Health and diet [edit]

500-milligram calcium supplements made from calcium carbonate

Calcium carbonate is widely used medicinally every bit an cheap dietary calcium supplement for gastric antacid^[38] (such as Tums and Eno). It may be used as a phosphate folder for the treatment of hyperphosphatemia (primarily in patients with chronic kidney failure). It is used in the pharmaceutical industry as an inert filler for tablets and other pharmaceuticals.^[39]

Calcium carbonate is used in the production of calcium oxide as well as toothpaste and has seen a resurgence every bit a food preservative and color retainer, when used in or with products such as organic apples.^[xl]

Calcium carbonate is used therapeutically as phosphate binder in patients on maintenance haemodialysis. It is the nigh common form of phosphate binder prescribed, specially in non-dialysis chronic kidney disease. Calcium carbonate is the most commonly used phosphate binder, just clinicians are increasingly prescribing the more than expensive, non-calcium-based phosphate binders, particularly sevelamer.

Excess calcium from supplements, fortified nutrient, and high-calcium diets can cause 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 x days, combined with alkaline powders, which provided symptomatic relief for peptic ulcer disease. Over the next several decades, the Sippy regimen resulted in kidney failure, alkalosis, and hypercalcaemia, mostly in men with peptic ulcer affliction. These agin furnishings were reversed when the regimen stopped, just information technology was fatal in some patients with protracted vomiting. Milk-alkali syndrome declined in men after effective treatments for peptic ulcer affliction arose. Since the 1990s information technology has been nigh oft reported in women taking calcium supplements above the recommended range of 1.two to 1.5 grams daily, for prevention and treatment of osteoporosis,^{[41] [42]} and is exacerbated past dehydration. Calcium has been added to over-the-counter products, which contributes to inadvertent excessive intake. Excessive calcium intake tin can pb to hypercalcemia, complications of which include vomiting, abdominal hurting and altered mental condition.^[43]

Equally a food additive information technology is designated E170,^[44] and it has an INS number of 170. Used as an acidity regulator, anticaking agent, stabilizer or color it is approved for usage in the EU,^[45] USA^[46] and Australia and New Zealand.^[47] It is "added by law to all United kingdom of great britain and northern ireland milled breadstuff flour except wholemeal".^{[48] [49]} It is used in some soy milk and almond milk products as a source of dietary calcium; at least one written report suggests that calcium carbonate might be as bioavailable equally the calcium in cow'due south milk.^[fifty] Calcium carbonate is also used equally a firming agent in many canned and bottled vegetable products.

Several calcium supplement formulations have been documented to comprise the element atomic number 82,^[51] posing a public health business.^[52] Lead is ordinarily found in natural sources of calcium.^[51]

Agriculture and aquaculture [edit]

Agricultural lime, powdered chalk or limestone, is used every bit a inexpensive method for neutralising acidic soil, making it suitable for planting, besides used in aquaculture industry for pH regulation of pond soil before initiating culture.^[53]

Household cleaning [edit]

Calcium carbonate is a key ingredient in many household cleaning powders like Comet and is used every bit a scrubbing agent.

Pollution mitigation [edit]

In 1989, a researcher, Ken Simmons, introduced CaCO₃ into the Whetstone Brook in Massachusetts.^[54] His hope was that the calcium carbonate would counter the acid in the stream from acrid rain and relieve the trout that had ceased to spawn. Although his experiment was a success, it did increase the amount of aluminium ions in the surface area of the brook that was non treated with the limestone. This shows that CaCO₃ can be added to neutralize the effects of acid pelting in river ecosystems. Currently calcium carbonate is used to neutralize acidic conditions in both soil and water.^{[55] [56] [57]} Since the 1970s, such liming has been proficient on a large scale in Sweden to mitigate acidification and several g lakes and streams are limed repeatedly.^[58]

Calcium carbonate is as well used in flue gas desulfurisation applications eliminating harmful So₂ and NO₂ emissions from coal and other fossil fuels burnt in large fossil fuel power stations.^[55]

Calcination equilibrium [edit]

Calcination of limestone using charcoal fires to produce quicklime has been practiced since antiquity by cultures all over the world. The temperature at which limestone yields calcium oxide is usually given every bit 825 °C, but stating an absolute threshold is misleading. Calcium carbonate exists in equilibrium with calcium oxide and carbon dioxide at whatsoever temperature. At each temperature at that place is a partial force per unit area of carbon dioxide that is in equilibrium with calcium carbonate. At room temperature the equilibrium overwhelmingly favors calcium carbonate, because the equilibrium CO₂ force per unit area is just a tiny fraction of the partial CO₂ pressure level in air, which is about 0.035 kPa.

At temperatures above 550 °C the equilibrium CO₂ pressure level begins to exceed the CO_ii pressure in air. So above 550 °C, calcium carbonate begins to outgas CO₂ into air. Nevertheless, in a charcoal fired kiln, the concentration of CO₂ will be much higher than it is in air. Indeed, if all the oxygen in the kiln is consumed in the fire, so the partial force per unit area of CO₂ in the kiln can be as loftier as 20 kPa.^[59]

The table shows that this partial pressure is not achieved until the temperature is nearly 800 °C. For the outgassing of CO₂ from calcium carbonate to happen at an economically useful rate, the equilibrium force per unit area must significantly exceed the ambience pressure of CO_two . And for it to happen rapidly, the equilibrium pressure must exceed total atmospheric pressure level of 101 kPa, which happens at 898 °C.

Equilibrium pressure of CO_two over CaCO₃ (P) versus temperature (T).^[60]

P (kPa)	0.055	0.thirteen	0.31	1.80	five.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

Solubility [edit]

With varying CO_ii force per unit area [edit]

Calcium carbonate is poorly soluble in pure water (47 mg/L at normal atmospheric CO₂ partial pressure equally shown below).

The equilibrium of its solution is given past the equation (with dissolved calcium carbonate on the right):

CaCO3 ⇌ Ca2+ + CO 2− 3

K sp = 3.7×x−9 to 8.7×10−ix at 25 °C

where the solubility product for [Ca^two+][CO 2− 3 ] is given as anywhere from One thousand _sp = 3.seven×10⁻⁹ to 1000 _sp = 8.vii×10^−nine at 25 °C, depending upon the data source.^{[60] [61]} What the equation means is that the product of molar concentration of calcium ions (moles of dissolved Ca^two+ per liter of solution) with the tooth concentration of dissolved CO ii− 3 cannot exceed the value of K _sp. This seemingly uncomplicated solubility equation, withal, must be taken along with the more complicated equilibrium of carbon dioxide with water (run into carbonic acrid). Some of the CO ii− iii combines with H⁺ in the solution according to

HCO − 3 ⇌ H+ + CO ii− three

One thousand a2 = 5.61×x−11 at 25 °C

HCO − iii is known equally the bicarbonate ion. Calcium bicarbonate is many times more than soluble in h2o than calcium carbonate—indeed it exists only in solution.

Some of the HCO − 3 combines with H⁺ in solution according to

H2CO3 ⇌ H+ + HCO − three

Chiliad a1 = 2.v×10−four at 25 °C

Some of the H₂CO₃ breaks upward into water and dissolved carbon dioxide according to

HtwoO + CO2(aq) ⇌ HiiCOiii

K h = 1.70×ten−3 at 25 °C

And dissolved carbon dioxide is in equilibrium with atmospheric carbon dioxide according to

P COtwo / [COtwo] = k H

where k H = 29.76 atm/(mol/L) at 25 °C (Henry constant), P COii being the COtwo partial pressure.

For ambient air, P _CO2 is around 3.5×ten⁻⁴ atmospheres (or equivalently 35 Pa). The terminal equation above fixes the concentration of dissolved CO_ii as a function of P _CO2 , independent of the concentration of dissolved CaCO_three . At atmospheric fractional pressure of CO₂ , dissolved CO₂ concentration is ane.2×10⁻⁵ moles per liter. The equation before that fixes the concentration of H₂CO₃ equally a role of CO₂ concentration. For [CO₂ ] = ane.ii×ten⁻⁵ , it results in [H₂CO_iii] = 2.0×10⁻⁸ moles per liter. When [H_twoCO_three] is known, the remaining three equations together with

Calcium ion solubility every bit a function of CO₂ partial force per unit area at 25 °C (K _sp = 4.47×10⁻⁹ )

P COii (atm)	pH	[Ca2+] (mol/50)
10−12	12.0	v.19×10−3
10−10	11.3	1.12×10−three
10−eight	10.7	2.55×x−4
x−6	nine.83	1.twenty×10−4
10−4	eight.62	three.16×10−four
3.5×x−4	8.27	4.70×x−iv
10−3	7.96	6.62×10−4
x−two	7.30	1.42×ten−three
ten−1	half dozen.63	3.05×10−3
1	five.96	vi.58×ten−iii
10	5.30	ane.42×ten−2

HiiO ⇌ H+ + OH−

K = ten−xiv at 25 °C

(which is true for all aqueous solutions), and the fact that the solution must be electrically neutral, i.eastward., the overall charge of dissolved positive ions [Ca^two+] + 2 [H⁺] must be cancelled out past the overall charge of dissolved negative ions [HCO − 3 ] + [CO 2− 3 ] + [OH⁻], get in possible to solve simultaneously for the remaining five unknown concentrations (note that the previously mentioned class of the neutrality is valid only if calcium carbonate has been put in contact with pure h2o or with a neutral pH solution; in the example where the initial water solvent pH is non neutral, the balance is not neutral).

The side by side table shows the issue for [Ca²⁺] and [H⁺] (in the course of pH) as a function of ambient partial pressure of CO₂ (One thousand _sp = 4.47×ten⁻⁹ has been taken for the calculation).

• At atmospheric levels of ambient CO₂ the table indicates that the solution volition be slightly alkaline with a maximum CaCO₃ solubility of 47 mg/L.
• As ambient CO₂ partial pressure is reduced below atmospheric levels, the solution becomes more and more alkali metal. At extremely low P _CO2 , dissolved CO₂ , bicarbonate ion, and carbonate ion largely evaporate from the solution, leaving a highly alkaline solution of calcium hydroxide, which is more soluble than CaCO₃ . Note that for P _CO2 = x⁻¹² atm, the [Ca²⁺][OH⁻]₂ production is still below the solubility product of Ca(OH)₂ ( viii×10^{−half dozen} ). For nevertheless lower CO₂ pressure, Ca(OH)₂ precipitation will occur before CaCO₃ precipitation.
• Every bit ambient CO_two partial force per unit area increases to levels higher up atmospheric, pH drops, and much of the carbonate ion is converted to bicarbonate ion, which results in higher solubility of Ca^two+ .

The effect of the latter is especially evident in twenty-four hours-to-mean solar day life of people who accept hard water. H2o in aquifers underground tin exist exposed to levels of CO₂ much higher than atmospheric. Every bit such h2o percolates through calcium carbonate rock, the CaCO₃ dissolves co-ordinate to the 2d trend. When that same water then emerges from the tap, in time it comes into equilibrium with CO₂ levels in the air past outgassing its excess CO₂ . The calcium carbonate becomes less soluble as a consequence, and the backlog precipitates as lime calibration. This same process is responsible for the formation of stalactites and stalagmites in limestone caves.

Two hydrated phases of calcium carbonate, monohydrocalcite CaCO₃·H_twoO and ikaite CaCO₃·6H₂O, may precipitate from water at ambient conditions and persist as metastable phases.

With varying pH, temperature and salinity: CaCO_iii scaling in swimming pools [edit]

In contrast to the open equilibrium scenario above, many swimming pools are managed by addition of sodium bicarbonate (NaHCO₃ ) to about 2 mM equally a buffer, and then command of pH through use of HCl, NaHSO₄ , Na₂CO_three , NaOH or chlorine formulations that are acidic or bones. In this situation, dissolved inorganic carbon (total inorganic carbon) is far from equilibrium with atmospheric CO₂ . Progress towards equilibrium through outgassing of CO_ii is slowed by

1. the tiresome reaction

   H_twoCO_iii ⇌ CO_ii(aq) + H₂O;^[62]

2. limited aeration in a deep water column; and
3. periodic replenishment of bicarbonate to maintain buffer capacity (often estimated through measurement of full alkalinity).

In this situation, the dissociation constants for the much faster reactions

H_twoCO_iii ⇌ H⁺ + HCO − 3 ⇌ ii H⁺ + CO 2− 3

permit the prediction of concentrations of each dissolved inorganic carbon species in solution, from the added concentration of HCO − 3 (which constitutes more than than 90% of Bjerrum plot species from pH 7 to pH 8 at 25 °C in fresh water).^[63] Add-on of HCO − three will increment CO ii− iii concentration at any pH. Rearranging the equations given above, we can see that [Caⁱⁱ⁺] = Thousand _sp / [CO 2− iii ] , and [CO 2− 3 ] = Thousand _a2 [HCO − 3 ] / [H⁺ ] . Therefore, when HCO − 3 concentration is known, the maximum concentration of Caⁱⁱ⁺ ions before scaling through CaCO₃ precipitation can exist predicted from the formula:

[Caⁱⁱ⁺ ]_max = Chiliad _sp / K _a2 × [H⁺ ] / [HCO − three ]

The solubility product for CaCO₃ (K _sp) and the dissociation constants for the dissolved inorganic carbon species (including Thousand _a2) are all essentially affected by temperature and salinity,^[63] with the overall outcome that [Ca^two+ ]_max increases from freshwater to saltwater, and decreases with ascent temperature, pH, or added bicarbonate level, every bit illustrated in the accompanying graphs.

The trends are illustrative for pool management, merely whether scaling occurs also depends on other factors including interactions with Mg²⁺ , [B(OH)₄]⁻ and other ions in the puddle, too every bit supersaturation effects.^{[64] [65]} Scaling is commonly observed in electrolytic chlorine generators, where there is a high pH most the cathode surface and scale deposition further increases temperature. This is one reason that some pool operators prefer borate over bicarbonate as the principal pH buffer, and avoid the use of pool chemicals containing calcium.^[66]

Solubility in a strong or weak acid solution [edit]

Solutions of strong (HCl), moderately strong (sulfamic) or weak (acetic, citric, sorbic, lactic, phosphoric) acids are commercially available. They are ordinarily used as descaling agents to remove limescale deposits. The maximum amount of CaCO₃ that tin can exist "dissolved" by one liter of an acid solution tin can be calculated using the above equilibrium equations.

• In the example of a strong monoacid with decreasing acrid concentration [A] = [A⁻ ], we obtain (with CaCO₃ tooth mass = 100 m/mol):

[A] (mol/L)	ane	10−1	10−2	x−3	10−4	ten−five	x−6	ten−7	ten−10
Initial pH	0.00	1.00	2.00	iii.00	4.00	5.00	6.00	half-dozen.79	7.00
Final pH	6.75	7.25	7.75	viii.fourteen	8.25	8.26	8.26	8.26	eight.27
Dissolved CaCO3 (g/50 of acrid)	fifty.0	five.00	0.514	0.0849	0.0504	0.0474	0.0471	0.0470	0.0470

where the initial state is the acrid solution with no Ca²⁺ (not taking into account possible CO₂ dissolution) and the final country is the solution with saturated Caⁱⁱ⁺ . For stiff acrid concentrations, all species have a negligible concentration in the terminal state with respect to Ca²⁺ and A⁻ so that the neutrality equation reduces approximately to 2[Ca²⁺ ] = [A⁻ ] yielding [Ca^two+ ] ≈ 0.5 [A⁻ ]. When the concentration decreases, [HCO − iii ] becomes non-negligible and so that the preceding expression is no longer valid. For vanishing acid concentrations, ane can recover the final pH and the solubility of CaCO_three in pure water.

• In the case of a weak monoacid (here we take acetic acid with pK _a = 4.76) with decreasing full acid concentration [A] = [A⁻ ] + [AH], nosotros obtain:

[A] (mol/L)	[Catwo+ ] ≈ 0.5 [A− ]	10−1	10−2	10−three	10−4	10−5	10−6	10−7	10−x
Initial pH	2.38	two.88	3.39	3.91	4.47	5.fifteen	6.02	6.79	seven.00
Final pH	6.75	7.25	vii.75	8.14	viii.25	8.26	eight.26	8.26	viii.27
Dissolved CaCO3 (yard/L of acid)	49.five	4.99	0.513	0.0848	0.0504	0.0474	0.0471	0.0470	0.0470

For the aforementioned total acid concentration, the initial pH of the weak acid is less acrid than the one of the stiff acid; nonetheless, the maximum amount of CaCO₃ which tin exist dissolved is approximately the aforementioned. This is because in the final state, the pH is larger than the pOne thousand _a, so that the weak acid is almost completely dissociated, yielding in the end as many H⁺ ions as the potent acid to "dissolve" the calcium carbonate.

• The adding in the example of phosphoric acrid (which is the most widely used for domestic applications) is more complicated since the concentrations of the iv dissociation states corresponding to this acrid must be calculated together with [HCO − three ], [CO 2− 3 ], [Ca^two+ ], [H⁺ ] and [OH⁻ ]. The organization may be reduced to a 7th degree equation for [H⁺ ] the numerical solution of which gives

[A] (mol/50)	ane	10−ane	10−2	10−3	10−4	10−5	x−6	10−vii	10−x
Initial pH	1.08	1.62	2.25	three.05	iv.01	5.00	five.97	6.74	seven.00
Final pH	6.71	7.17	vii.63	8.06	8.24	viii.26	8.26	8.26	8.27
Dissolved CaCO3 (g/L of acid)	62.0	7.39	0.874	0.123	0.0536	0.0477	0.0471	0.0471	0.0470

where [A] = [H₃PO₄] + [H₂PO − four ] + [HPO ii− 4 ] + [PO 3− four ] is the full acid concentration. Thus phosphoric acid is more than efficient than a monoacid since at the final almost neutral pH, the second dissociated state concentration [HPO two− 4 ] is not negligible (run into phosphoric acid).

See also [edit]

Electron micrograph of needle-like calcium carbonate crystals formed equally limescale in a kettle

Around 2 k of calcium-48 carbonate

• Cuttlebone
• Cuttlefish
• Gesso
• Limescale
• Marble
• Ocean acidification
• Whiting issue
• List of climate technology topics
• Lysocline

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External links [edit]

• International Chemical Safety Carte 1193
• CID 516889 from PubChem
• ATC codes: A02AC01 (WHO) and A12AA04 (WHO)
• The British Calcium Carbonate Clan – What is calcium carbonate Archived 24 May 2008 at the Wayback Auto
• CDC – NIOSH Pocket Guide to Chemical Hazards – Calcium Carbonate

Is Caco3 A Binary Compound,

Source: https://en.wikipedia.org/wiki/Calcium_carbonate

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