Calcium carbonate

What is it?	Calcium carbonate is a common substance found in rocks, shells, and bones.
Where is it found?	It's in limestone, chalk, marble, and the shells of marine animals.
Uses	It is used in making chalk, antacids, and cement.
Fun fact	Calcium carbonate helps keep our bones strong!

Calcium carbonate is a chemical compound with the chemical formula CaCO₃. It is found in rocks as the minerals calcite and aragonite, most notably in chalk and limestone. You can also find it in eggshells, gastropod shells, shellfish skeletons, and pearls. Materials that contain much calcium carbonate or look like it are called calcareous.

Crystal structure of calcite

Calcium carbonate is used in agricultural lime and can also be found as limescale, which forms when calcium ions in hard water react with carbonate ions. It has medical uses as a calcium supplement or an antacid, but eating too much can be harmful and may cause health problems like hypercalcemia and digestive issues.

Chemistry

Calcium carbonate behaves like other carbonates. It reacts with acids, producing carbonic acid that breaks down into carbon dioxide and water. When heated, it releases carbon dioxide and forms calcium oxide, also known as quicklime.

It also reacts with water that has carbon dioxide in it, turning into soluble calcium bicarbonate. This process helps shape caves and can cause hard water in some places. There is a special form of calcium carbonate called ikaite, which stays stable only in very cold temperatures below 8 °C.

Preparation

Most calcium carbonate used in factories comes from digging up rocks or mining. To make very pure calcium carbonate for food or medicine, people can use marble or make it from calcium oxide. They mix calcium oxide with water to get calcium hydroxide, and then they add carbon dioxide. This makes the calcium carbonate form, which is called precipitated calcium carbonate.

There is also a way to make calcium carbonate by mixing calcium chloride with ammonium carbonate in a special container called a desiccator. The ammonium carbonate breaks down into ammonia, carbon dioxide, and water. The carbon dioxide moves into the calcium chloride solution and helps create calcium carbonate.

Structure

Calcium carbonate, written as CaCO₃, usually appears in nature as a mineral called calcite. Another form, called aragonite, can be made when special conditions are used, like higher temperatures. Scientists are still learning more about a third form called vaterite. Different metals, like magnesium, strontium, and barium, can also form similar structures with carbonate, but they look a bit different because of their size.

Polymorphs

Calcium carbonate can form in three different ways, called polymorphs. The most common one is called calcite, followed by aragonite, and then vaterite, which is the least common.

These polymorphs have different shapes and structures. Calcite has a special shape called trigonal, aragonite has an orthorhombic shape, and vaterite can have more than one structure at the same time. All three can form together in water under normal conditions, but calcite is usually the main one that forms.

Crystal structure of calcite and aragonite

Some living things, like molluscs and arthropods, can make these different forms of calcium carbonate, often to build shells or for other protective purposes. They can choose which form to make, often using special molecules to help them do this.

Occurrence

Calcium carbonate is found in many places on Earth and beyond. It makes up important rocks like limestone, chalk, marble, and travertine. It is also the main part of eggshells, snail shells, and most seashells. Even some green vegetables like broccoli and kale have calcium carbonate in them.

Calcite is the most stable polymorph of calcium carbonate. It is transparent to opaque. A transparent variety called Iceland spar (shown here) was used to create polarized light in the 19th century.

Scientists have also found signs of calcium carbonate on Mars, suggesting there may have once been liquid water there. This helps us learn more about the planet and its history.

Geology

Calcium carbonate is found in many places in nature and stores a lot of carbon. It appears as minerals like aragonite, calcite, and dolomite, and it makes up rocks such as limestone, chalk, marble, travertine, and tufa.

Surface precipitation of CaCO3 as tufa in Rubaksa, Ethiopia

In warm, clear tropical waters, there are many creatures like corals, plankton, algae, sponges, and mollusks that help create calcium carbonate. These animals usually live in shallow waters where they can get sunlight and food. In colder waters, calcium carbonate can still form, but it grows very slowly. When ocean floor moves under continents, calcium carbonate can break down and release carbon dioxide, which can lead to volcanic eruptions.

Calcium carbonate helps preserve fossils by filling in tiny spaces, keeping even very small details safe. However, when these fossils are brought to the surface, they can be damaged by weather.

Uses

Calcium carbonate is commonly used in construction. It can be a building material, part of road-building mixtures, an ingredient in cement, or a starting material for making lime used in buildings. However, because of damage from acid rain, it is no longer used alone for buildings but as a base for other materials.

500-milligram calcium supplements made from calcium carbonate

It is also used to help clean iron from its raw form in factories and to keep swimming pools at the right balance. In factories that make sugar, it helps remove unwanted materials. Chalk, made from calcium carbonate, was once a main part of blackboards but now most chalk is made from a different material. Calcium carbonate is also used to help grow special structures in water and as a filler in products like rubber gloves, diapers, and paper. It can make paints, plastics, and many other everyday items cheaper and better. It is even used in adhesives, tiles, and cleaning products.

Calcium carbonate is also used in health and food. It can be a cheap way to add calcium to diets or to help with stomach issues. It is found in medicines, toothpaste, and food products like apple preserves. It also helps balance the pH in soil and water to protect plants and fish. In factories, it is added to plastics to make them stronger and more stable.

Calcination equilibrium

Calcination of limestone using charcoal fires to make quicklime has been done for a very long time by many different cultures around the world. Limestone turns into calcium oxide when heated to about 825 °C, but this temperature can change depending on other conditions.

Calcium carbonate can change into calcium oxide and carbon dioxide depending on the temperature. At room temperature, calcium carbonate stays as it is because the air has very little carbon dioxide. But when the temperature goes above 550 °C, calcium carbonate starts to release carbon dioxide into the air. In a special kiln with charcoal fires, there can be much more carbon dioxide than in normal air. This makes the process happen at lower temperatures than expected.

For this process to work well and quickly, the temperature needs to be around 898 °C.

Equilibrium pressure of CO₂ over CaCO₃ (P) versus temperature (T).
P (kPa)	0.055	0.13	0.31	1.80	5.9	9.3	14	24	34	51	72	80	91	101	179	901	3961
T (°C)	550	587	605	680	727	748	777	800	830	852	871	881	891	898	937	1082	1241

Solubility

Travertine calcium carbonate deposits from a hot spring

Calcium carbonate does not dissolve easily in plain water. It can mix better when the air around has less of a certain gas, or when the water has special added ingredients.

When acids, like those found in some cleaning products, are used, calcium carbonate can dissolve more. This helps get rid of hard spots in pipes or on surfaces caused by calcium carbonate.

CaCO₃ ⇌ Ca²⁺ + CO2−3

K_sp = 3.7×10⁻⁹ to 8.7×10⁻⁹ at 25 °C

HCO−3 ⇌ H⁺ + CO2−3

K_a2 = 5.61×10⁻¹¹ at 25 °C

H₂CO₃ ⇌ H⁺ + HCO−3

K_a1 = 2.5×10⁻⁴ at 25 °C

H₂O + CO₂(aq) ⇌ H₂CO₃

K_h = 1.70×10⁻³ at 25 °C

⁠P_CO₂/[CO₂]⁠ = H v {\displaystyle H_{\rm {v}}}

where H v {\displaystyle H_{\rm {v}}} = 29.76 atm/(mol/L) at 25 °C (Henry volatility), and P_CO₂ is the CO₂ partial pressure.

Calcium ion solubility as a function of CO₂ partial pressure at 25 °C (K_sp = 4.47×10⁻⁹)
P_CO₂ (atm)	pH	[Ca²⁺] (mol/L)
10⁻¹²	12.0	5.19×10⁻³
10⁻¹⁰	11.3	1.12×10⁻³
10⁻⁸	10.7	2.55×10⁻⁴
10⁻⁶	9.83	1.20×10⁻⁴
10⁻⁴	8.62	3.16×10⁻⁴
3.5×10⁻⁴	8.27	4.70×10⁻⁴
10⁻³	7.96	6.62×10⁻⁴
10⁻²	7.30	1.42×10⁻³
10⁻¹	6.63	3.05×10⁻³
1	5.96	6.58×10⁻³
10	5.30	1.42×10⁻²

H₂O ⇌ H⁺ + OH⁻

K = 10⁻¹⁴ at 25 °C

[A] (mol/L)	1	10⁻¹	10⁻²	10⁻³	10⁻⁴	10⁻⁵	10⁻⁶	10⁻⁷	10⁻¹⁰
Initial pH	0.00	1.00	2.00	3.00	4.00	5.00	6.00	6.79	7.00
Final pH	6.75	7.25	7.75	8.14	8.25	8.26	8.26	8.26	8.27
Dissolved CaCO₃ (g/L of acid)	50.0	5.00	0.514	0.0849	0.0504	0.0474	0.0471	0.0470	0.0470

[A] (mol/L)	[Ca²⁺] ≈ 0.5 [A⁻]	10⁻¹	10⁻²	10⁻³	10⁻⁴	10⁻⁵	10⁻⁶	10⁻⁷	10⁻¹⁰
Initial pH	2.38	2.88	3.39	3.91	4.47	5.15	6.02	6.79	7.00
Final pH	6.75	7.25	7.75	8.14	8.25	8.26	8.26	8.26	8.27
Dissolved CaCO₃ (g/L of acid)	49.5	4.99	0.513	0.0848	0.0504	0.0474	0.0471	0.0470	0.0470

[A] (mol/L)	1	10⁻¹	10⁻²	10⁻³	10⁻⁴	10⁻⁵	10⁻⁶	10⁻⁷	10⁻¹⁰
Initial pH	1.08	1.62	2.25	3.05	4.01	5.00	5.97	6.74	7.00
Final pH	6.71	7.17	7.63	8.06	8.24	8.26	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