Sodium carbonate: Difference between revisions
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== Occurrence == |
== Occurrence == |
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Sodium carbonate is soluble in [[water]], but can occur naturally in arid regions, especially in the [[mineral]] deposits (''evaporites'') formed when seasonal [[lake]]s 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 [[mummy|mummies]]. |
Sodium carbonate is soluble in [[water]], but can occur naturally in arid regions, especially in the [[mineral]] deposits (''evaporites'') formed when seasonal [[lake]]s 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 [[mummy|mummies]]. Sodium carbonate is three known forms of hydrates, they are sodium carbonate decahydrate, sodium carbonate heptahydrate and sodium carbonate monohydrate |
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== Production == |
== Production == |
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Revision as of 12:57, 3 September 2005
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General |
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|---|---|
| Name | Sodium carbonate |
| Chemical formula | Na2CO3 |
| Appearance | White solid |
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Physical |
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| 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 |
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Thermochemistry |
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| ΔfH0liquid | -1102 kJ/mol |
| ΔfH0solid | -1131 kJ/mol |
| S0liquid, 1 bar | ? J/mol·K |
| S0solid | 135 J/mol·K |
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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 |
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SI units were used where possible. Unless otherwise stated, standard conditions were used. | |
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. Sodium carbonate is three known forms of hydrates, they are sodium carbonate decahydrate, sodium carbonate heptahydrate and sodium carbonate monohydrate
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
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.
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:
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:
The sodium bicarbonate was then converted to sodium carbonate by heating it, releasing water and carbon dioxide:
Meanwhile, the ammonia was regenerated from the ammonium chloride byproduct by treating it with the lime (calcium hydroxide) left over from carbon dioxide generation:
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.