Sodium carbonate

From Academic Kids

</table> Sodium carbonate or soda ash, Na2CO3, is a sodium salt of carbonic acid. It is a white crystalline compound with a cooling alkaline taste, and found in the ashes of many plants. It is produced artificially in large quantities from common salt. It is used in the manufacture of: glass, chemicals such as sodium silicates and sodium phosphates, the pulp and paper industries, the manufacture of detergents and for the treatment of water. It is also used as an alkaline agent in many chemical industries.

Occurrence

Sodium carbonate is soluble in water, but can occur naturally in arid regions, especially in the mineral deposits (evaporites) formed when seasonal lakes evaporate. Deposits of the mineral natron, a combination of sodium carbonate and sodium bicarbonate, have been mined from dry lake bottoms in Egypt since ancient times, when natron was used in the preparation of mummies.

Production

In 1791, the French chemist Nicolas Leblanc patented a process for producing sodium carbonate from salt, sulfuric acid, limestone, and coal. First, sea salt (sodium chloride) was boiled in sulfuric acid to yield sodium sulfate and hydrochloric acid gas, according to the chemical equation

2 NaCl + H2SO4Na2SO4 + 2 HCl

Next, the sodium sulfate was blended with crushed limestone (calcium carbonate) and coal, and the mixture was burnt, producing sodium carbonate along with carbon dioxide and calcium sulfide.

Na2SO4 + CaCO3 + 2 C → Na2CO3 + 2 CO2 + CaS

The sodium carbonate was extracted from the ashes with water, and then collected by allowing the water to evaporate.

The hydrochloric acid produced by the Leblanc process was a major source of air pollution, and the calcium sulfide byproduct also presented waste disposal issues. However, it remained the major production method for sodium carbonate until the late 1880s.

In 1861, the Belgian industrial chemist Ernest Solvay developed a method to convert sodium chloride to sodium carbonate using ammonia. The Solvay process centered around a large hollow tower. At the bottom, calcium carbonate (limestone) was heated to release carbon dioxide:

CaCO3CaO + CO2

At the top, a concentrated solution of sodium chloride and ammonia entered the tower. As the carbon dioxide bubbled up through it, sodium bicarbonate precipitated:

NaCl + NH3 + CO2 + H2ONaHCO3 + NH4Cl

The sodium bicarbonate was then converted to sodium carbonate by heating it, releasing water and carbon dioxide:

2 NaHCO3 → Na2CO3 + H2O + CO2

Meanwhile, the ammonia was regenerated from the ammonium chloride byproduct by treating it with the lime (calcium hydroxide) left over from carbon dioxide generation:

CaO + H2OCa(OH)2
Ca(OH)2 + 2 NH4ClCaCl2 + 2 NH3 + 2 H2O

Because the Solvay process recycled its ammonia, it consumed only brine and limestone, and had calcium chloride as its only waste product. This made it substantially more economical than the Leblanc process, and it soon came to dominate world sodium carbonate production. By 1900, 90% of sodium carbonate was produced by the Solvay process, and the last Leblanc process plant closed in the early 1920s.

Sodium carbonate is still produced by the Solvay process in much of the world today. However, large natural deposits found in 1938 near the Green River in Wyoming, have made its industrial production in North America uneconomical.de:Natriumcarbonat es:Carbonato sdico nl:Natriumcarbonaat ja:炭酸ナトリウム pl:Węglan sodu ru:Карбонат натрия

sv:Natriumkarbonat

Properties

General

Name Sodium carbonate
Chemical formula Na2CO3
Appearance White solid

Physical

Formula weight 106.0 amu
Melting point 1124 K (851 °C)
Boiling point Decomposes at ?
Density 2.5 ×103 kg/m3
Crystal structure ?
Solubility 10.9 g in 100g water

Thermochemistry

ΔfH0liquid -1102 kJ/mol
ΔfH0solid -1131 kJ/mol
S0liquid, 1 bar ? J/mol·K
S0solid 135 J/mol·K

Safety

Ingestion May cause irritation.
Inhalation Harmful, should be avoided especially in the long term.
Skin Irritation and possible burns.
Eyes Severe irritation, injury may result.
More info Hazardous Chemical Database (http://ull.chemistry.uakron.edu/erd/chemicals/8/7741.html)

SI units were used where possible. Unless otherwise stated, standard conditions were used.

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