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{{Chembox new| Name = Calcium carbonate| ImageFile = Calcium carbonate.jpg| OtherNames =
Limestone;
calcite; aragonite;
chalk; marble| Section1 = {{Chembox Identifiers| CASNo = 471-34-1| RTECS =-->
| Section2 = {{Chembox Properties| Formula = CaCO3| MolarMass = 100.087 g/mol| Appearance = White powder.| Density = 2.83 g/cm³, solid.| Solubility = Insoluble| MeltingPt = 825 °C| BoilingPt = Decomposes| pKa =| pKb =-->
| Section3 = {{Chembox Structure| MolShape = Linear| Dipole =-->
| Section7 = {{Chembox Hazards| MainHazards = Not hazardous.| FlashPt = Non-flammable.| RPhrases = , , | SPhrases = , -->
-->
Calcium carbonate is a
chemical compound, with the
chemical formula CalciumCarbon
Oxygen3. It is a common substance found as
Rock (geology) in all parts of the world, and is the main component of seashells, snails, and eggshells. 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.
Occurrence
Calcium carbonate is found naturally as the following
minerals and rocks:
- Aragonite
- Calcite
- Vaterite or (μ-CaCO3)
- Chalk
- Limestone
- Marble
- Travertine
To test whether a mineral or rock contains calcium carbonate, strong acids, such as hydrochloric acid, can be added to it. If the sample does contain calcium carbonate, it will fizz and produce carbon dioxide and water. Weak acids such as acetic acid will react, albeit less vigorously. All of the rocks/minerals mentioned above will react with acid.
Preparation
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) or it can be prepared by passing carbon dioxide into a solution of calcium hydroxide: the calcium carbonate precipitates out, and this grade of product is referred to as a precipitate (abbreviated to PCC).
Ca(OH)2 + CO2 → CaCO3 + H2O
Chemical properties
See also: Carbonate
Calcium carbonate shares the typical properties of other carbonates. Notably:
it reacts with strong acids, releasing carbon dioxide:CaCO3 + 2HCl → CaCl2 + CO2 + H2O
it releases carbon dioxide on heating (to above 840 °C in the case of CaCO3), to form calcium oxide, commonly called quick lime:
CaCO3 → CaO + CO2
Calcium carbonate will react with water that is saturated with carbon dioxide to form the soluble calcium bicarbonate.
CaCO3 + CO2 + H2O → Ca(HCO3)2
This reaction is important in the
erosion of
carbonate rocks, forming caverns, and leads to hard water in many regions.
Uses
The main use of calcium carbonate 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 . A common contaminant is magnesium carbonate.
Calcium carbonate is widely used as an extender in paints, in particular matte emulsion paint where typically 30% by weight of the paint is either chalk or marble.
Calcium carbonate is also widely used as a filler in plastics. Some typical examples include around 15 to 20% loading of chalk in uPVC drain pipe, 5 to 15% loading of stearate coated chalk or marble in uPVC window profile. Fine ground calcium carbonate is an essential ingredient in the microporous film used in babies' diapers and some building films as the pores are nucleated around the calcium carbonate particles during the manufacture of the film by biaxial stretching.
Calcium carbonate is also used in a wide range of trade and DIY adhesives, sealants, and decorating fillers. 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.
Calcium carbonate is widely used medicinally as an inexpensive dietary calcium supplement,
antacid, and/or phosphate binder. It is also used in the pharmaceutical industry as a base material for
tablets of other pharmaceuticals.
Calcium carbonate is known as
whiting in
ceramics (art)/glazing applications, 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.
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 been a major component of blackboard chalk. Chalk may consist of either calcium carbonate or gypsum, hydrated calcium sulfate CaSO4·2H2O.
In North America, calcium carbonate has begun to replace Kaolinite 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 calcium carbonate (GCC) or precipitated calcium carbonate (PCC). The latter has a very fine and controlled particle size, on the order of 2 micron in diameter, useful in coatings for paper.
As a
food additive, it is used in some soy milk products as a source of dietary calcium.
In 1989, a researcher introduced CaCO3 into the
Whetstone Brook in Massachusetts . His hope was that the calcium carbonate 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 aluminum 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. Nowadays, calcium carbonate is used to neutralise acidic conditions in both soil and water.
Calcination Equilibrium
{| border="1" cellspacing="0" cellpadding="2" style="margin: 0 0 0 0.5em; background: white; border-collapse: collapse; border-color: #C0C090;" align="right"! | Equilibrium Pressure of CO2 over CaCO3
CRC Handbook of Chemistry and Physics 44th ed., p2292|-| 550 °C| 0.055 kPascal (unit)|-| 587 °C| 0.13 kPascal (unit)|-| 605 °C| 0.31 kPascal (unit)|-| 680 °C| 1.80 kPascal (unit)|-| 727 °C| 5.9 kPascal (unit)|-| 748 °C| 9.3 kPascal (unit)|-| 777 °C| 14 k
Pascal (unit)|-| 800 °C| 24 kPascal (unit)|-| 830 °C| 34 k
Pascal (unit)|-| 852 °C| 51 kPascal (unit)|-| 871 °C| 72 kPascal (unit)|-| 881 °C| 80 k
Pascal (unit)|-| 891 °C| 91 kPascal (unit)|-| 898 °C| 101 kPascal (unit)|-| 937 °C| 179 kPascal (unit)|-| 1082 °C| 901 kPascal (unit)|-| 1241 °C| 3961 k
Pascal (unit)|-|}Calcination of limestone using charcoal fires to produce
calcium oxide has been practiced since antiquity by cultures all over the world. The answer to the question, "how hot does the fire have to be?" is usually given as 825 °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 CO2 pressure is only a tiny fraction of the partial CO2 pressure in air, which is about 0.035 k
Pascal (unit). At temperatures above 550 °C the equilibrium CO2 pressure begins to exceed the CO2 pressure in air. So above 550 °C, calcium carbonate begins to outgas CO2 into air. But 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, then the partial pressure of CO2 in the kiln can be as high as 20 k
Pascal (unit). The table shows that this equilibrium 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 pressure must significantly exceed the ambient pressure of CO2. And for it to happen rapidly, the equilibrium pressure must exceed total atmospheric pressure of 101 kPascal (unit), which happens at 898 °C.
Solubility of calcium carbonate in water
Solubility in pure water with varying CO2 pressure
Calcium carbonate is poorly soluble in pure water. The equilibrium of its solution is given by the equation (with dissolved calcium carbonate on the right):
:{| width="450"
| width="50%" height="30"| CaCO3 Ca2+ + CO32–|
Ksp = 3.7×10–9 to 8.7×10–9 at 25 °C|}
where the
solubility product for is given as anywhere from
Ksp = 3.7×10–9 to
Ksp = 8.7×10–9 at 25 °C, depending upon the data source. CSUDHCRC Handbook of Chemistry and Physics, 44th ed. What the equation means is that the product of molar concentration of calcium ions (mole (unit) of dissolved Ca2+ per liter of solution) with the molar concentration of dissolved CO32– cannot exceed the value of
Ksp. 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 CO32– combines with H+ in the solution according to:
:{| width="450"
| width="50%" height="25"| HCO3– H+ + CO32– |
Ka2 = 5.61×10–11 at 25 °C|}
HCO3– is known as the
bicarbonate ion.
Calcium bicarbonate is many times more soluble in water than calcium carbonate -- indeed it exists
only in solution.
Some of the HCO3– combines with H+ in solution according to:
:{| width="450"
| width="50%" height="25"|H2CO3 H+ + HCO3– |
Ka1 = 2.5×10–4 at 25 °C|}
Some of the H2CO3 breaks up into water and dissolved carbon dioxide according to:
:{| width="450"
| width="50%" height="25"| H2O + CO2(dissolved) H2CO3 |
Kh = 1.70×10–3 at 25 °C|}
And dissolved carbon dioxide is in equilibrium with atmospheric carbon dioxide according to:
:{| width="450"
| width="50%" |\frac{P_{\mathrm{CO}_2-->{}\ =\ k_\mathrm{H}| where
kH = 29.76 atm/(mol/L) at 25°C (
Henry's law), \scriptstyle P_{\mathrm{CO}_2} being the CO2 partial pressure.|}
{| border="1" cellspacing="0" cellpadding="2" style="margin: 0 0 0 0.5em; background: white; border-collapse: collapse; border-color: #C0C090;" align="right"! colspan="3" | Calcium
Ion Solubility as a function of carbon dioxide
partial pressure at 25 °C]| width="120" align="center" bgcolor="E0E0E0"| (mol/L)|-|10−12 ||align="right"|12.0|| align="right"|5.19 × 10−3|-|10−10 ||align="right"|11.3|| align="right"|1.12 × 10−3|-|10−8 ||align="right"|10.7|| align="right"|2.55 × 10−4|-|10−6 ||align="right"|9.83|| align="right"|1.20 × 10−4|-|10−4 ||align="right"|8.62|| align="right"|3.16 × 10−4|-|
3.5 × 10−4 ||align="right"|
8.27|| align="right"|
4.70 × 10−4|-|10−3 ||align="right"|7.96|| align="right"|6.62 × 10−4|-|10−2 ||align="right"|7.30|| align="right"|1.42 × 10−3|-|10−1 ||align="right"|6.63|| align="right"|3.05 × 10−3|-|1 ||align="right"|5.96|| align="right"|6.58 × 10−3|-|10 ||align="right"|5.30|| align="right"|1.42 × 10−2|}
For ambient air, \scriptstyle P_{\mathrm{CO}_2} is around 3.5×10–4 atmospheres (or equivalently 35
Pascal (unit)). The last equation above fixes the concentration of dissolved CO2 as a function of \scriptstyle P_{\mathrm{CO}_2}, independent of the concentration of dissolved CaCO3. At atmospheric partial pressure of CO2, dissolved CO2 concentration is 1.2×10–5 moles/liter. The equation before that fixes the concentration of H2CO3 as a function of . For =1.2×10–5, it results in =2.0×10–8 moles per liter. When is known, the remaining three equations together with
:{| width="450"
| width="50%" height="25"| H2O H+ + OH–|
K = 10–14 at 25 °C|}
(which is true for all aqueous solutions), and the fact that the solution must be electrically neutral,
:2 + = + 2 +
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 calcium carbonate 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 and (in the form of p
H) as a function of ambient partial pressure of CO2 (
Ksp = 4.47×10−9 has been taken for the calculation). At atmospheric levels of ambient CO2 the table indicates the solution will be slightly alkaline. The trends the table shows are
1) As ambient CO2 partial pressure is reduced below atmospheric levels, the solution becomes more and more alkaline. At extremely low \scriptstyle P_{\mathrm{CO}_2}, 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.
2) As ambient CO2 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 Ca2+.
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 calcium carbonate 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 calcium carbonate becomes less soluble as a result and the excess precipitates as lime scale. This same process is responsible for the formation of
stalactites and stalagmites in limestone caves.
Two hydrated phases of calcium carbonate, monohydrocalcite and
Ikaite, may precipitate from water at ambient conditions and persist as metastable phases.
Solubility at atmospheric CO2 pressure with varying p
H
We now consider the problem of the maximum solubility of calcium carbonate in normal atmospheric conditions (\scriptstyle P_{\mathrm{CO}_2} = 3.5 × 10−4 atm) when the p
H of the solution is adjusted. This is for example the case in a swimming pool where the p
H is maintained between 7 and 8 (by addition of NaHSO4 to decrease the p
H or of NaHCO3 to increase it). From the above equations for the solubility product, the hydratation reaction and the two acid reactions, the following expression for the maximum can be easily deduced:
_\mathrm{max} = \frac{K_\mathrm{sp}k_\mathrm{H--> {K_\mathrm{h}K_\mathrm{a1}K_\mathrm{a2--> \frac{^2}{P_{\mathrm{CO}_2-->
showing a quadratic dependence in . The numerical application with the above values of the constants gives
{| border="1" cellspacing="0" cellpadding="4" style="margin: 0 0 0 0.5em; background: }; border-collapse: collapse; border-color: };"|-| width="170" align="center" |
pH|width="40"|7.0|width="40"|7.2|width="40"|7.4|width="40"|7.6|width="40"|7.8|width="40"|8.0|width="40"|8.2|width="40"|8.27|width="40"|8.4|-| width="170" align="center" |max (10-4mol/L or °F)'|width="40"|1590|width="40"|635|width="40"|253|width="40"|101|width="40"|40.0|width="40"|15.9|width="40"|6.35|width="40"|4.70|width="40"|2.53|-| width="170" align="center"|max (mg/L)|width="40"|6390|width="40"|2540|width="40"|1010|width="40"|403|width="40"|160|width="40"|63.9|width="40"|25.4|width="40"|18.9|width="40"|10.1|}Comments:
- decreasing the pH from 8 to 7 increases the maximum Ca2+ concentration by a factor 100
- note that the Ca2+ 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/OCl- hypochlorous acid in the case of "chlorine" maintenance) results in a maximum Ca2+ concentration of 1010 mg/L. This means that successive cycles of water evaporation and partial renewing may result in a very hard water before CaCO3 precipitates. Addition of a calcium sequestrant or complete renewing of the water will solve the problem.
References
See also
External links
{{Chembox new| Name = Calcium carbonate| ImageFile = Calcium carbonate.jpg| OtherNames = Limestone; calcite; aragonite; chalk; marble| Section1 = {{Chembox Identifiers| CASNo = 471-34-1| RTECS =-->
| Section2 = {{Chembox Properties| Formula = CaCO3| MolarMass = 100.087 g/mol| Appearance = White powder.| Density = 2.83 g/cm³, solid.| Solubility = Insoluble| MeltingPt = 825 °C| BoilingPt = Decomposes| pKa =| pKb =-->
| Section3 = {{Chembox Structure| MolShape = Linear| Dipole =-->
| Section7 = {{Chembox Hazards| MainHazards = Not hazardous.| FlashPt = Non-flammable.| RPhrases = , , | SPhrases = , -->
-->
Calcium carbonate is a
chemical compound, with the chemical formula
CalciumCarbonOxygen3. It is a common substance found as Rock (geology) in all parts of the world, and is the main component of
seashells, snails, and
eggshells. 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.
Occurrence
Calcium carbonate is found naturally as the following minerals and rocks:
To test whether a mineral or rock contains calcium carbonate, strong acids, such as hydrochloric acid, can be added to it. If the sample does contain calcium carbonate, it will fizz and produce carbon dioxide and water. Weak acids such as acetic acid will react, albeit less vigorously. All of the rocks/minerals mentioned above will react with acid.
Preparation
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) or it can be prepared by passing carbon dioxide into a solution of calcium hydroxide: the calcium carbonate precipitates out, and this grade of product is referred to as a precipitate (abbreviated to PCC).
Ca(OH)2 + CO2 → CaCO3 + H2O
Chemical properties
See also: Carbonate
Calcium carbonate shares the typical properties of other carbonates. Notably:
it reacts with strong acids, releasing carbon dioxide:CaCO3 + 2HCl → CaCl2 + CO2 + H2O
it releases carbon dioxide on heating (to above 840 °C in the case of CaCO3), to form calcium oxide, commonly called quick lime:
CaCO3 → CaO + CO2
Calcium carbonate will react with water that is saturated with carbon dioxide to form the soluble calcium bicarbonate.
CaCO3 + CO2 + H2O → Ca(HCO3)2
This reaction is important in the
erosion of
carbonate rocks, forming
caverns, and leads to
hard water in many regions.
Uses
The main use of calcium carbonate 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 . A common contaminant is magnesium carbonate.
Calcium carbonate is widely used as an extender in paints, in particular matte emulsion paint where typically 30% by weight of the paint is either chalk or marble.
Calcium carbonate is also widely used as a filler in plastics. Some typical examples include around 15 to 20% loading of chalk in uPVC drain pipe, 5 to 15% loading of stearate coated chalk or marble in uPVC window profile. Fine ground calcium carbonate is an essential ingredient in the microporous film used in babies'
diapers and some building films as the pores are nucleated around the calcium carbonate particles during the manufacture of the film by biaxial stretching.
Calcium carbonate is also used in a wide range of trade and DIY adhesives, sealants, and decorating fillers. 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.
Calcium carbonate is widely used medicinally as an inexpensive dietary calcium supplement, antacid, and/or phosphate binder. It is also used in the pharmaceutical industry as a base material for tablets of other pharmaceuticals.
Calcium carbonate is known as
whiting in
ceramics (art)/glazing applications, 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.
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 been a major component of blackboard chalk. Chalk may consist of either calcium carbonate or
gypsum, hydrated calcium sulfate CaSO4·2H2O.
In North America, calcium carbonate has begun to replace Kaolinite 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 calcium carbonate (GCC) or precipitated calcium carbonate (PCC). The latter has a very fine and controlled particle size, on the order of 2 micron in diameter, useful in coatings for paper.
As a
food additive, it is used in some soy milk products as a source of dietary calcium.
In 1989, a researcher introduced CaCO3 into the Whetstone Brook in Massachusetts . His hope was that the calcium carbonate 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 aluminum 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. Nowadays, calcium carbonate is used to neutralise acidic conditions in both soil and water.
Calcination Equilibrium
{| border="1" cellspacing="0" cellpadding="2" style="margin: 0 0 0 0.5em; background: white; border-collapse: collapse; border-color: #C0C090;" align="right"! | Equilibrium Pressure of CO2 over CaCO3
CRC Handbook of Chemistry and Physics 44th ed., p2292|-| 550 °C| 0.055 k
Pascal (unit)|-| 587 °C| 0.13 kPascal (unit)|-| 605 °C| 0.31 kPascal (unit)|-| 680 °C| 1.80 k
Pascal (unit)|-| 727 °C| 5.9 kPascal (unit)|-| 748 °C| 9.3 k
Pascal (unit)|-| 777 °C| 14 kPascal (unit)|-| 800 °C| 24 k
Pascal (unit)|-| 830 °C| 34 k
Pascal (unit)|-| 852 °C| 51 k
Pascal (unit)|-| 871 °C| 72 kPascal (unit)|-| 881 °C| 80 k
Pascal (unit)|-| 891 °C| 91 kPascal (unit)|-| 898 °C| 101 kPascal (unit)|-| 937 °C| 179 kPascal (unit)|-| 1082 °C| 901 kPascal (unit)|-| 1241 °C| 3961 kPascal (unit)|-|}Calcination of limestone using charcoal fires to produce
calcium oxide has been practiced since antiquity by cultures all over the world. The answer to the question, "how hot does the fire have to be?" is usually given as 825 °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 CO2 pressure is only a tiny fraction of the partial CO2 pressure in air, which is about 0.035 kPascal (unit). At temperatures above 550 °C the equilibrium CO2 pressure begins to exceed the CO2 pressure in air. So above 550 °C, calcium carbonate begins to outgas CO2 into air. But 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, then the partial pressure of CO2 in the kiln can be as high as 20 kPascal (unit). The table shows that this equilibrium 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 pressure must significantly exceed the ambient pressure of CO2. And for it to happen rapidly, the equilibrium pressure must exceed total atmospheric pressure of 101 kPascal (unit), which happens at 898 °C.
Solubility of calcium carbonate in water
Solubility in pure water with varying CO2 pressure
Calcium carbonate is poorly soluble in pure water. The equilibrium of its solution is given by the equation (with dissolved calcium carbonate on the right):
:{| width="450"
| width="50%" height="30"| CaCO3 Ca2+ + CO32–|
Ksp = 3.7×10–9 to 8.7×10–9 at 25 °C|}
where the solubility product for is given as anywhere from
Ksp = 3.7×10–9 to
Ksp = 8.7×10–9 at 25 °C, depending upon the data source. CSUDHCRC Handbook of Chemistry and Physics, 44th ed. What the equation means is that the product of molar concentration of calcium ions (
mole (unit) of dissolved Ca2+ per liter of solution) with the molar concentration of dissolved CO32– cannot exceed the value of
Ksp. 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 CO32– combines with H+ in the solution according to:
:{| width="450"
| width="50%" height="25"| HCO3– H+ + CO32– |
Ka2 = 5.61×10–11 at 25 °C|}
HCO3– is known as the bicarbonate ion.
Calcium bicarbonate is many times more soluble in water than calcium carbonate -- indeed it exists
only in solution.
Some of the HCO3– combines with H+ in solution according to:
:{| width="450"
| width="50%" height="25"|H2CO3 H+ + HCO3– |
Ka1 = 2.5×10–4 at 25 °C|}
Some of the H2CO3 breaks up into water and dissolved carbon dioxide according to:
:{| width="450"
| width="50%" height="25"| H2O + CO2(dissolved) H2CO3 |
Kh = 1.70×10–3 at 25 °C|}
And dissolved carbon dioxide is in equilibrium with atmospheric carbon dioxide according to:
:{| width="450"
| width="50%" |\frac{P_{\mathrm{CO}_2-->{}\ =\ k_\mathrm{H}| where
kH = 29.76 atm/(mol/L) at 25°C (
Henry's law), \scriptstyle P_{\mathrm{CO}_2} being the CO2 partial pressure.|}
{| border="1" cellspacing="0" cellpadding="2" style="margin: 0 0 0 0.5em; background: white; border-collapse: collapse; border-color: #C0C090;" align="right"! colspan="3" |
Calcium Ion Solubility
as a function of
carbon dioxide partial pressure at 25 °C]| width="120" align="center" bgcolor="E0E0E0"| (mol/L)|-|10−12 ||align="right"|12.0|| align="right"|5.19 × 10−3|-|10−10 ||align="right"|11.3|| align="right"|1.12 × 10−3|-|10−8 ||align="right"|10.7|| align="right"|2.55 × 10−4|-|10−6 ||align="right"|9.83|| align="right"|1.20 × 10−4|-|10−4 ||align="right"|8.62|| align="right"|3.16 × 10−4|-|
3.5 × 10−4 ||align="right"|
8.27|| align="right"|
4.70 × 10−4|-|10−3 ||align="right"|7.96|| align="right"|6.62 × 10−4|-|10−2 ||align="right"|7.30|| align="right"|1.42 × 10−3|-|10−1 ||align="right"|6.63|| align="right"|3.05 × 10−3|-|1 ||align="right"|5.96|| align="right"|6.58 × 10−3|-|10 ||align="right"|5.30|| align="right"|1.42 × 10−2|}
For ambient air, \scriptstyle P_{\mathrm{CO}_2} is around 3.5×10–4 atmospheres (or equivalently 35
Pascal (unit)). The last equation above fixes the concentration of dissolved CO2 as a function of \scriptstyle P_{\mathrm{CO}_2}, independent of the concentration of dissolved CaCO3. At atmospheric partial pressure of CO2, dissolved CO2 concentration is 1.2×10–5 moles/liter. The equation before that fixes the concentration of H2CO3 as a function of . For =1.2×10–5, it results in =2.0×10–8 moles per liter. When is known, the remaining three equations together with
:{| width="450"
| width="50%" height="25"| H2O H+ + OH–|
K = 10–14 at 25 °C|}
(which is true for all aqueous solutions), and the fact that the solution must be electrically neutral,
:2 + = + 2 +
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 calcium carbonate 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 and (in the form of p
H) as a function of ambient partial pressure of CO2 (
Ksp = 4.47×10−9 has been taken for the calculation). At atmospheric levels of ambient CO2 the table indicates the solution will be slightly alkaline. The trends the table shows are
1) As ambient CO2 partial pressure is reduced below atmospheric levels, the solution becomes more and more alkaline. At extremely low \scriptstyle P_{\mathrm{CO}_2}, 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.
2) As ambient CO2 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 Ca2+.
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 calcium carbonate 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 calcium carbonate becomes less soluble as a result and the excess precipitates as lime scale. This same process is responsible for the formation of
stalactites and
stalagmites in limestone caves.
Two hydrated phases of calcium carbonate,
monohydrocalcite and Ikaite, may precipitate from water at ambient conditions and persist as metastable phases.
Solubility at atmospheric CO2 pressure with varying p
H
We now consider the problem of the maximum solubility of calcium carbonate in normal atmospheric conditions (\scriptstyle P_{\mathrm{CO}_2} = 3.5 × 10−4 atm) when the p
H of the solution is adjusted. This is for example the case in a swimming pool where the p
H is maintained between 7 and 8 (by addition of NaHSO4 to decrease the p
H or of NaHCO3 to increase it). From the above equations for the solubility product, the hydratation reaction and the two acid reactions, the following expression for the maximum can be easily deduced:
_\mathrm{max} = \frac{K_\mathrm{sp}k_\mathrm{H--> {K_\mathrm{h}K_\mathrm{a1}K_\mathrm{a2--> \frac{^2}{P_{\mathrm{CO}_2-->
showing a quadratic dependence in . The numerical application with the above values of the constants gives
{| border="1" cellspacing="0" cellpadding="4" style="margin: 0 0 0 0.5em; background: }; border-collapse: collapse; border-color: };"|-| width="170" align="center" |
pH|width="40"|7.0|width="40"|7.2|width="40"|7.4|width="40"|7.6|width="40"|7.8|width="40"|8.0|width="40"|8.2|width="40"|8.27|width="40"|8.4|-| width="170" align="center" |max (10-4mol/L or °F)'|width="40"|1590|width="40"|635|width="40"|253|width="40"|101|width="40"|40.0|width="40"|15.9|width="40"|6.35|width="40"|4.70|width="40"|2.53|-| width="170" align="center"|max (mg/L)|width="40"|6390|width="40"|2540|width="40"|1010|width="40"|403|width="40"|160|width="40"|63.9|width="40"|25.4|width="40"|18.9|width="40"|10.1|}Comments:
- decreasing the pH from 8 to 7 increases the maximum Ca2+ concentration by a factor 100
- note that the Ca2+ 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/OCl- hypochlorous acid in the case of "chlorine" maintenance) results in a maximum Ca2+ concentration of 1010 mg/L. This means that successive cycles of water evaporation and partial renewing may result in a very hard water before CaCO3 precipitates. Addition of a calcium sequestrant or complete renewing of the water will solve the problem.
References
See also
External links
Calcium carbonate - Wikipedia, the free encyclopedia
Calcium carbonate is a chemical compound with the chemical formula Ca C O 3. It is a common substance found as rock in all parts of the world, and is the main component of shells ...
Calcium Carbonate An
What is Calcium Carbonate? Calcium Carbonate is an exceptional mineral. The chemical formula CaCO 3 covers a raw material, which is widespread throughout nature, whether dissolved ...
Calcium Carbonate
CALCIUM CARBONATE Calcium is a mineral used to treat osteoporosis (bone loss), kidney stones, and menstrual cramps. Link to Shop http://my-word-here.info/1/calcium_c...
MedlinePlus Drug Information: Calcium Carbonate
Calcium Carbonate ... Why is this medication prescribed? Return to top. Calcium carbonate is a dietary supplement used when the amount of calcium taken in the diet is not ...
calcium carbonate
Environmental Fate - Ecotoxicology - Human Health - A to Z Index - Home . GENERAL INFORMATION . Description: An inorganic substance with multiaction ...
IMA Europe - Industrial Minerals Association
Calcium Carbonate. Whole mountain ranges are made from Calcium Carbonate in the form of chalk, limestone, marble and dolomite. It constitutes more than 4% of the earth’s crust.
calcium carbonate - Hutchinson encyclopedia article about calcium ...
calcium carbonate. White solid, found in nature as limestone, marble, and chalk. It is a valuable resource, used in the making of iron, steel, cement, glass, slaked lime, bleaching ...
Longcliffe Ltd | Calcium carbonate of the highest purity
Longcliffe is one of the UK's leading producers of high quality calcium carbonates extracting, processing and delivering over 4000 tonnes of limestone products every working day.
calcium-carbonate
Calcium Carbonate Supplier, Product, Exporter, Factory, China - Talc ...
We is an international white minerals company supplying high quality calcium carbonate and talc. ... Management Ideals. About Shengtai HuanQiu . We are pursuing to do the best.
