Calcium Carbonate - Structure, Formula, Preparation, Properties, Uses, FAQs

Structure of Calcium Carbonate (CaCO₃)

Under normal conditions, CaCO₃ (the mineral calcite) has a thermodynamically stable hexagonal shape. Other forms, such as denser 2.83g/cm³ orthorhombic CaCO₃ as well as hexagonal CaCO₃, are possible. It is feasible to prepare aragonite at temperatures exceeding 85°C along with vaterite at 60°C, for example. Calcite has calcium atoms coordinated with six oxygen atoms, while aragonite has nine oxygen atoms coordinated. The vaterite's chemical composition is unknown.

Magnesium carbonate (MgCO₃) has a calcite structure, but strontium carbonate (SrCO₃) and barium carbonate (BaCO₃) have an aragonite structure, which reflects their greater ionic radii.

Calcium carbonate is an ionic compound made up of calcium ions ($\mathrm{Ca}^{2+}$) and carbonate ions ($\mathrm{CO}_3{ }^{2-}$).

Carbonate Ion ($\mathrm{CO}_3{ }^{2-}$) Structure

• The carbonate ion has a trigonal planar geometry.

• One carbon atom lies at the center and is bonded to three oxygen atoms.

• All C–O bonds are equivalent due to resonance, giving partial double-bond character.

• The bond angle between O–C–O is 120°.

• The ion is planar and carries an overall –2 charge.

Arrangement in Calcium Carbonate

• $\mathrm{Ca}^{2+}$ ions and $\mathrm{CO}_3{ }^{2-}$ ions are held together by strong electrostatic (ionic) forces.

• Each calcium ion is surrounded by carbonate ions, forming a stable crystal lattice.

Crystal Forms

Calcium carbonate exists mainly in two crystalline forms:

1. Calcite (most stable form)

   • Each $\mathrm{Ca}^{2+}$ ion is coordinated by six oxygen atoms from different carbonate ions.

2. Aragonite (less stable)

   • Nine oxygen atoms coordinate each $\mathrm{Ca}^{2+}$ ion.

Occurrence of Calcium Carbonate

1. Geological Sources: Calcite, aragonite, and vaterite are pure calcium carbonate minerals found in nature. Limestones, chalks, travertines, and marbles are examples of industrially significant calcium carbonate source rocks. Corals thrive in warm, clear tropical waters rather than the frigid waters of the polar regions. Since the ample light and filterable food, calcium carbonate sources such as planktons (coccoliths and planktic foraminifera), sponges, echinoderms, coral algae and mollusks are typically found in shallow water locations.

2. Biological Sources: Calcium carbonate is found in eggshells, snail shells, and most seashells, and it is also a commercial source of this chemical. Oyster shells have recently been recognized as a dietary calcium source as well as a useful industrial source. On a diet, dark green vegetables like broccoli and kale contain a lot of calcium carbonate, but they aren't useful as an industrial supply.

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Preparation of CaCO₃

1. Carbon dioxide and slaked lime are used as raw materials to make CaCO₃. Calcite is formed when carbon dioxide is transported through slaked lime.

Adding sodium carbonate to calcium chloride is another way to produce calcite.

$\begin{aligned} & \mathrm{Ca}(\mathrm{OH})_2+\mathrm{CO}_2 \rightarrow \mathrm{CaCO}_3+\mathrm{H}_2 \mathrm{O} \\ & \mathrm{CaCl}_2+\mathrm{Na}_2 \mathrm{CO}_3 \rightarrow \mathrm{CaCO}_3+2 \mathrm{NaCl}\end{aligned}$

2. Calcium hydrogen-carbonate is formed when there is an excess of carbon dioxide in the atmosphere. It's made by passing carbon dioxide gas through calcium hydroxide on a big scale (slaked lime). Carbon dioxide, on the other hand, creates the soluble calcium hydrogen-carbonate when it is released in excess.

$\mathrm{Ca}(\mathrm{OH})_2+\mathrm{CO}_2 \rightarrow \mathrm{CaCO}_3+\mathrm{H}_2 \mathrm{O}$

Commercial Production of Calcium Carbonate

Commercially available calcium carbonate comes in two grades. In the industrial world, both classes compete primarily on particle size and product attributes.

1. Ground Calcium carbonate: This is made by extracting and processing naturally occurring calcium carbonate deposits. The form of GCC crystals is irregularly rhombohedral, with a wider size distribution.

$\mathrm{CaCO}_3+\mathrm{H}_2 \mathrm{SO}_4 \rightarrow \mathrm{CaSO}_4+\mathrm{H}_2 \mathrm{O}+\mathrm{CO}_2$

2. Precipitated Calcium carbonate: It is produced by a by-product of some bulk chemical operations. The shape of PCC crystals varies depending on the product, along with particles are more homogenous as well as regular, with restricted size distribution.

PCC has smaller particles, is purer, is less abrasive, and has a brighter appearance than GCC.

$\mathrm{CaCO}_3(\mathrm{~s}) \xrightarrow{\Delta} \mathrm{CaO}(\mathrm{s})+\mathrm{CO}_2(\mathrm{~g})$

Calcium carbonate produces carbon dioxide when it reacts with dilute acids.

$\mathrm{CaCO}_3+2 \mathrm{HCl} \rightarrow \mathrm{CaCl}_2+\mathrm{H}_2 \mathrm{O}+\mathrm{CO}_2$

Properties of Calcium Carbonate (CaCO₃)

1. Calcium carbonate is a fluffy calcium powder that decomposes into carbon dioxide when heated to 1200 K.
2. When it reacts with dilute acid, it produces carbon dioxide as a by-product.
3. At a temperature of 1200 K, calcium carbonate decomposes into calcium oxide and carbon dioxide.
4. When carbon dioxide reacts with dilute acids, calcium carbonate is generated.
5. At a temperature of 1200 K, calcium carbonate decomposes into calcium oxide and carbon dioxide.
6. When carbon dioxide reacts with dilute acids, calcium carbonate is generated.
7. Calcium carbonate (CaCO₃) has a molecular weight of 100.0869 g/mol.
8. CaCO₃ has a molar mass of 100.0869 g/mol.

Application of Calcium Carbonate

1. The pulp and paper industry uses a lot of calcium carbonate. It can be used as a filter and a pigment, resulting in a whiter, higher-quality pigment than other minerals.
2. It is used in the construction industry as filler in concrete to improve its durability and aesthetics, as well as to cleanse metals for use in construction applications.
3. Calcium carbonate is also used in fertilizers to deliver calcium to plants and to keep the pH of the soil stable.
4. Calcium carbonate can be used as a supplement in vitamins and as a food ingredient for livestock and people.
5. Calcium carbonate is used in water and wastewater treatment plants to remove acidity and pollutants.

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Some Solved Examples

Question 1: 4 gm of impure $\mathrm{CaCO}_3$ on treatment with excess HCl produces $0.88 \mathrm{gm} \mathrm{CO}_2$. What is the percent purity of the $\mathrm{CaCO}_3$ sample?
1) 75
2) 25
3) 90
4) (correct) 50

Solution:

Ratio of elements in $\mathrm{CaCO}_3$ :
Ca: 1, C: 1, $0: 3$

$\text { Molar mass }=1 \times 40+1 \times 12+3 \times 16=100$
$\mathrm{CaCO}_3 \xrightarrow{\Delta} \mathrm{CaO}+\mathrm{CO}_2$
$\begin{aligned}
& \% \text { purity }=\frac{\text { mass of pure product }}{\text { mass of impure product obtained }} \times 100 \% \\
& \therefore \mathrm{x}=\frac{0.88}{44} \mathrm{~mol}=0.02 \mathrm{~mol}=2 \mathrm{gm} \\
& \therefore \% \text { purity }=\frac{2}{4} \times 100=50
\end{aligned}$
Hence, the answer is option (4).

Question 2: In curing cement plaster, water is sprinkled from time to time. This helps in:

1) keeping it cool

2) (correct) developing interlocking needle­-like crystals of hydrated silicates

3) Hydrating sand and gravel mixed with cement

4) converting sand into silicic acid.

Solution:

Water helps in the hydration of calcium aluminates and calcium silicates which change into their colloidal gels by developing interlocking needle-like crystals of hydrated silicates.

The hydrolysis leads to the formation of calcium hydroxide which is useful in binding the particles of calcium silicate together.

Hence, the answer is option (2).

Question 3: Which one of the following alkaline earth metal sulfates has its hydration enthalpy greater than its lattice enthalpy?
1) $\mathrm{CaSO}_4$
2) (correct) $\mathrm{BeSO}_4$
3) $\mathrm{BaSO}_4$
4) $\mathrm{SrSO}_4$

Solution:

We know that
Solubility of sulfates of alkaline metals -
Solubility in water decreases from Be to Ba
Solubility order

$\Rightarrow \mathrm{BeSO}_4>\mathrm{MgSO}_4>\mathrm{CaSO}_4>\mathrm{SrSO}_4>\mathrm{BaSO}_4$
Hence, the answer is option (2).

Question 4: The solubilities of carbonates decrease down the magnesium group due to a decrease in

1) lattice energies of solids

2) (correct) hydration energies of cations

3) inter­-ionic attraction

4) entropy of solution formation.

Solution:

In the alkaline earth metal group, the hydration energy of metal carbonates decreases down the group, which means that carbonates become more stable down the group.

Hence, the answer is option (2).

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