Calcium carbonate is a chemical compound with the formula CaCO3. It is
a common substance found in rocks as the minerals calcite and
aragonite (most notably as limestone, which contains both of those
minerals) and is the main component of pearls and the shells of marine
organisms, snails, and eggs.
Calcium carbonate is the active
ingredient in agricultural lime and is created when calcium ions in
hard water react with carbonate ions to create limescale. It is
medicinally used as a calcium supplement or as an antacid, but
excessive consumption can be hazardous.
4.1 Geological sources
4.2 Biological sources
5.1 Carbonate compensation depth
5.2 Role in taphonomy
6.1 Industrial applications
6.2 Health and dietary applications
6.3 Agricultural use
6.4 Environmental applications
8.1 With varying CO2 pressure
8.2 With varying pH, temperature and salinity: CaCO3 scaling in
Solubility in a strong or weak acid solution
9 See also
11 External links
Calcium carbonate shares the typical properties of other carbonates.
it reacts with acids, releasing carbon dioxide:
CaCO3(s) + 2H+(aq) → Ca2+(aq) + CO2(g) + H2O (l)
it releases carbon dioxide upon heating, called a thermal
decomposition reaction, or calcination (to above 840 °C in the
case of CaCO3), to form calcium oxide, commonly called quicklime, with
reaction enthalpy 178 kJ/mole:
CaCO3 (s) → CaO (s) + CO2 (g)
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 rock, forming
caverns, and leads to hard water in many regions.
An unusual form of calcium carbonate is the hexahydrate, ikaite,
Ikaite is stable only below 6 °C.
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
Alternatively, calcium carbonate is prepared from calcium oxide. Water
is added to give calcium hydroxide then carbon dioxide is passed
through this solution to precipitate the desired calcium carbonate,
referred to in the industry as precipitated calcium carbonate
CaO + H2O → Ca(OH)2
displaystyle ce Ca(OH)2 + CO2 -> CaCO3(v) + H2O
The thermodynamically stable form of CaCO3 under normal conditions is
hexagonal β-CaCO3, (the mineral calcite). Other forms can be
prepared, the denser,(2.83 g/cc) orthorhombic λ-CaCO3 ( the mineral
aragonite) and μ-CaCO3, occurring as the mineral vaterite. The
aragonite form can be prepared by precipitation at temperatures above
85 °C, the vaterite form can be prepared by precipitation at
Calcite contains calcium atoms coordinated by 6 oxygen
atoms, in aragonite they are coordinated by 9 oxygen atoms. The
vaterite structure is not fully understood.
MgCO3 has the calcite structure, whereas strontium and barium
carbonate (SrCO3 and BaCO3) adopt the aragonite structure, reflecting
their larger ionic radii.
Calcium carbonate chunks from clamshell
Calcite, aragonite and vaterite are pure calcium carbonate minerals.
Industrially important source rocks which are predominantly calcium
carbonate include limestone, chalk, marble and travertine.
Calcite is the most stable polymorph of calcium carbonate. It is
transparent to opaque. A transparent variety called Iceland spar
(shown here) is used for optical purposes.[clarification needed]
Eggshells, snail shells and most seashells are predominantly calcium
carbonate and can be used as industrial sources of that chemical.
Oyster shells have enjoyed recent recognition as a source of dietary
calcium, but are also a practical industrial source. Dark
green vegetables such as broccoli and kale contain dietarily
significant amounts of calcium carbonate, however, they are not
practical as an industrial source.
Beyond Earth, strong evidence suggests the presence of calcium
carbonate on Mars. Signs of calcium carbonate have been detected at
more than one location (notably at Gusev and Huygens craters). This
provides some evidence for the past presence of liquid water.
Carbonate is found frequently in geologic settings and constitutes an
enormous carbon reservoir.
Calcium carbonate occurs as aragonite,
calcite and dolomite. The carbonate minerals form 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 as 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 exist at higher latitudes but have a very slow growth rate. The
calcification processes are changed by ocean acidification.
Where the oceanic crust is subducted under a continental plate
sediments will be carried down to warmer zones in the asthenosphere
and lithosphere. Under these conditions calcium carbonate decomposes
to produce carbon dioxide which, along with other gases, give rise to
explosive volcanic eruptions.
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 carbonate compensation depth can range from
4–6 km below sea level.
Role in taphonomy
Calcium carbonate can preserve fossils through permineralization. Most
of the vertebrate fossils of the Two Medicine Formation—a geologic
formation known for its duck-billed dinosaur eggs—are preserved by
CaCO3 permineralization. This type of preservation conserves high
levels of detail, even down to the microscopic level. However, it also
leaves specimens vulnerable to weathering when exposed to the
Trilobite populations were once thought to have composed 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, which had purely chitinous shells.
The main use of calcium carbonate is in the construction industry,
either as a building material or limestone aggregate for road building
or as an ingredient of cement or as the starting material for the
preparation of builder's lime by burning in a kiln. However, because
of weathering mainly caused by acid rain, calcium carbonate (in
limestone form) is no longer used for building purposes on its own,
but only as a raw/primary substance for building materials.
Calcium carbonate is also 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 various impurities present, and
separates from the purified iron.
In the oil industry, calcium carbonate is added to drilling fluids as
a formation-bridging and filtercake-sealing agent; it is also a
weighting material which increases the density of drilling fluids to
control the downhole pressure.
Calcium carbonate is added to swimming
pools, as a pH corrector for maintaining alkalinity and offsetting the
acidic properties of the disinfectant agent.
It is also used as a raw material in the refining of sugar from sugar
beet; It is calcined in a kiln with anthracite to produce calcium
oxide and carbon dioxide. This burnt lime is then slaked in sweet
water to produce a calcium hydroxide suspension for the precipitation
of impurities in raw juice during carbonatation.
Calcium carbonate has traditionally been a major component of
blackboard chalk. However, modern manufactured chalk is mostly gypsum,
hydrated calcium sulfate CaSO4·2H2O.
Calcium carbonate is a main
source for growing Seacrete, or Biorock. Precipitated calcium
carbonate (PCC), pre-dispersed in slurry form, is a common filler
material for latex gloves with the aim of achieving maximum saving in
material and production costs.
Fine ground calcium carbonate (GCC) is an essential ingredient in the
microporous film used in diapers and some building films as the pores
are nucleated around the calcium carbonate particles during the
manufacture of the film by biaxial stretching. GCC or PCC is used as a
filler in paper because they are cheaper than wood fiber. In terms of
market volume, GCC are the most important types of fillers currently
used. Printing and writing paper can contain 10–20% calcium
carbonate. In North America, calcium carbonate has begun to replace
kaolin in the production of glossy paper. Europe has been practicing
this as alkaline papermaking or acid-free 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 micrometres.
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. It is also a popular filler in
plastics. Some typical examples include around 15 to 20% loading
of chalk in unplasticized polyvinyl chloride (uPVC) drain pipe, 5 to
15% loading of stearate coated chalk or marble in uPVC window profile.
PVC cables can use calcium carbonate 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
Polypropylene compounds are often
filled with calcium carbonate to increase rigidity, a requirement that
becomes important at high use temperatures. Here the percentage is
often 20–40%. It also routinely used as a filler in thermosetting
resins (sheet and bulk molding compounds) and has also been mixed
with ABS, and other ingredients, to form some types of compression
molded "clay" poker chips. Precipitated calcium
carbonate, made by dropping calcium oxide into water, is used by
itself or with additives as a white paint, known as
Calcium carbonate is added to a wide range of trade and do it yourself
adhesives, sealants, and decorating fillers.
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.
In ceramics/glazing applications, calcium carbonate is known as
whiting, 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. Ground calcium
carbonate is 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, 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 water can be used to clean tarnish on
Health and dietary applications
500-milligram calcium supplements made from calcium carbonate
Calcium carbonate is widely used medicinally as an inexpensive dietary
calcium supplement for gastric antacid (e.g., Tums). 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 filler for tablets and other
Calcium carbonate is used in the production of calcium oxide as well
as toothpaste and has seen a resurgence as a food preservative and
color retainer, when used in or with products such as organic
Excess calcium from supplements, fortified food 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 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 hypercalcaemia, 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 arose. During the past 15 years, it has been
reported in women taking calcium supplements above the recommended
range of 1.2 to 1.5 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
As a food additive it is designated E170, and it has an INS number
of 170. Used as an acidity regulator, anticaking agent, stabiliser or
colour it is approved for usage in the EU, USA and Australia
and New Zealand. It is used in some soy milk and almond milk
products as a source of dietary calcium; one study suggests that
calcium carbonate might be as bioavailable as the calcium in cow's
Calcium carbonate is also used as a firming agent in many
canned or bottled vegetable products.
Agricultural lime, powdered chalk or limestone, is used as a cheap
method for neutralising acidic soil, making it suitable for
In 1989, a researcher, Ken Simmons, 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 amount 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 calcium carbonate is used to neutralize acidic
conditions in both soil and water. 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
Calcium carbonate is also 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.
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
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 kPa.
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. However, 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 kPa.
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 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 kPa, which happens at
Equilibrium pressure of CO2 over CaCO3 (P) vs. temperature (T).
With varying CO2 pressure
Travertine calcium carbonate deposits from a hot spring
Calcium carbonate is poorly soluble in pure water (47 mg/L at
normal atmospheric CO2 partial pressure as shown below).
The equilibrium of its solution is given by the equation (with
dissolved calcium carbonate on the right):
CaCO3 ⇌ Ca2+ + CO32−
Ksp = 3.7×10−9 to 8.7×10−9 at 25 °C
where the solubility product for [Ca2+] [CO32−] is given as anywhere
from Ksp = 3.7×10−9 to Ksp = 8.7×10−9 at 25 °C, depending
upon the data source. What the equation means is that the
product of molar concentration of calcium ions (moles 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
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:
H2CO3 ⇌ H+ + HCO3−
Ka1 = 2.5×10−4 at 25 °C
Some of the H2CO3 breaks up into water and dissolved carbon dioxide
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:
displaystyle frac P_ ce CO2 [ ce CO2 ] = k_ ce H
where kH = 29.76 atm/(mol/L) at 25 °C (Henry constant),
displaystyle P_ ce CO2
being the CO2 partial pressure.
For ambient air,
displaystyle P_ ce CO2
is around 3.5×10−4 atmospheres (or equivalently 35 Pa). The last
equation above fixes the concentration of dissolved CO2 as a function
displaystyle P_ ce CO2
, 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 [CO2]. For [CO2]=1.2×10−5, it results in
[H2CO3]=2.0×10−8 moles per liter. When [H2CO3] is known, the
remaining three equations together with
Calcium ion solubility as a function of CO2 partial pressure at
25 °C (Ksp = 4.47×10−9)
displaystyle scriptstyle P_ ce CO2
5.19 × 10−3
1.12 × 10−3
2.55 × 10−4
1.20 × 10−4
3.16 × 10−4
3.5 × 10−4
4.70 × 10−4
6.62 × 10−4
1.42 × 10−3
3.05 × 10−3
6.58 × 10−3
1.42 × 10−2
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[Ca2+] + [H+] = [HCO3−] + 2[CO32−] + [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 calcium carbonate has been put in contact
with pure water or with a neutral pH solution; in the case where the
initial 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 (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 with a maximum CaCO3 solubility of
As ambient CO2 partial pressure is reduced below atmospheric levels,
the solution becomes more and more alkaline. At extremely low
displaystyle P_ ce 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
displaystyle P_ ce CO2 =10^ -12 mathrm atm
, the [Ca2+] [OH−]2 product is still below the solubility product of
Ca(OH)2 (8×10−6). For still lower CO2 pressure, Ca(OH)2
precipitation will occur before CaCO3 precipitation.
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, CaCO3·H2O
and ikaite, CaCO3·6H2O, may precipitate from water at ambient
conditions and persist as metastable phases.
With varying pH, temperature and salinity: CaCO3 scaling in swimming
In contrast to the open equilibrium scenario above, many swimming
pools are managed by addition of sodium bicarbonate (NaHCO3) to about
2 mM as a buffer, then control of pH through use of HCl, NaHSO4,
Na2CO3, NaOH or chlorine formulations that are acidic or basic. In
this situation, dissolved inorganic carbon (total inorganic carbon) is
far from equilibrium with atmospheric CO2. Progress towards
equilibrium through outgassing of CO2 is slowed by (i) the slow
reaction H2CO3 ⇌ CO2(aq) + H2O; (ii) limited aeration in a deep
water column and (iii) periodic replenishment of bicarbonate to
maintain buffer capacity (often estimated through measurement of
In this situation, the dissociation constants for the much faster
reactions H2CO3 ⇌ H+ + HCO3‾ ⇌ 2 H+ + CO32− allow the
prediction of concentrations of each dissolved inorganic carbon
species in solution, from the added concentration of HCO3− (which
constitutes more than 90% of
Bjerrum plot species from pH 7 to pH 8 at
25 °C in fresh water). Addition of HCO3− will increase
CO32− concentration at any pH. Rearranging the equations given
above, we can see that [Ca2+] = Ksp / [CO32−], and [CO32−] = Ka2
× [HCO3−] / [H+]. Therefore, when HCO3− concentration is known,
the maximum concentration of Ca2+ ions before scaling through CaCO3
precipitation can be predicted from the formula:
Ca2+max = (Ksp / Ka2) × ([H+] / [HCO3−])
The solubility product for CaCO3 (Ksp) and the dissociation constants
for the dissolved inorganic carbon species (including Ka2) are all
substantially affected by temperature and salinity, with the
overall effect that Ca2+max increases from fresh to salt water, and
decreases with rising temperature, pH, or added bicarbonate level, as
illustrated in the accompanying graphs.
The trends are illustrative for pool management, but whether scaling
occurs also depends on other factors including interactions with Mg2+,
B(OH)4− and other ions in the pool, as well as supersaturation
effects. Scaling is commonly observed in electrolytic chlorine
generators, where there is a high pH near the cathode surface and
scale deposition further increases temperature. This is one reason
that some pool operators prefer borate over bicarbonate as the primary
pH buffer, and avoid the use of pool chemicals containing calcium.
Solubility in a strong or weak acid solution
Solutions of strong (HCl), moderately strong (sulfamic) or weak
(acetic, citric, sorbic, lactic, phosphoric) acids are commercially
available. They are commonly used as descaling agents 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.
In the case of a strong monoacid with decreasing acid concentration
[A] = [A−], we obtain (with CaCO3 molar mass = 100 g):
(g/L of acid)
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
displaystyle scriptstyle [mathrm Ca ^ 2+ ]simeq frac [mathrm
A ^ - ] 2
. When the concentration decreases, [HCO3−] becomes non-negligible
so that the preceding expression is no longer valid. For vanishing
acid concentrations, one can recover the final pH and the solubility
of CaCO3 in pure water.
In the case of a weak monoacid (here we take acetic acid with pKA =
4.76) with decreasing total acid concentration [A] = [A−]+[AH], we
(g/L of acid)
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 pKA, so that
the weak acid is almost completely dissociated, yielding in the end as
many H+ ions as the strong acid to "dissolve" the calcium carbonate.
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 [HCO3−], [CO32−], [Ca2+],
[H+] and [OH−]. The system may be reduced to a seventh degree
equation for [H+] the numerical solution of which gives
(g/L of acid)
where [A] = [H3PO4] + [H2PO4−] + [HPO42−] + [PO43−] is the total
acid concentration. Thus phosphoric acid is more efficient than a
monoacid since at the final almost neutral pH, the second dissociated
state concentration [HPO42−] is not negligible (see phosphoric
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Magnesium (increases motility)
Aluminium (decreases motility)
Combinations and complexes
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Drugs for treatment of hyperkalemia and hyperphosphatemia (V03AE)
Sodium zirconium cyclosilicate
Calcium acetate/magnesium carbonate