Sodium carbonate
| Sodium carbonate | |||||
|---|---|---|---|---|---|
| |||||
| General | |||||
| Other names | Soda ash Washing soda | ||||
| Molecular formula | Na2CO3 | ||||
| Molar mass | 106.0 g/mol | ||||
| Appearance | Dark White solid | ||||
| CAS number | [497-19-8] | ||||
| Properties | |||||
| Density and phase | 2.5 g/cm3, solid | ||||
| Solubility in water | 30 g/100 ml (20 °C) | ||||
| Melting point | 851 °C | ||||
| Boiling point | decomposes | ||||
| Basicity (pKb) | ? | ||||
| Structure | |||||
| Coordination geometry |
? | ||||
| Crystal structure | ? | ||||
| Hazards | |||||
| MSDS | External MSDS | ||||
| EU classification | Irritant (Xi) | ||||
| NFPA 704 |
| ||||
| R-phrases | Template:R36 | ||||
| S-phrases | Template:S2, Template:S22, Template:S26 | ||||
| Flash point | non flammable | ||||
| RTECS number | VZ4050000 | ||||
| Supplementary data page | |||||
| Structure and properties |
n, εr, etc. | ||||
| Thermodynamic data |
Phase behaviour Solid, liquid, gas | ||||
| Spectral data | UV, IR, NMR, MS | ||||
| Related compounds | |||||
| Other anions | Sodium bicarbonate | ||||
| Other cations | Lithium carbonate Potassium carbonate | ||||
| Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa) Infobox disclaimer and references | |||||
Sodium carbonate (also known as washing soda or soda ash), Template:SodiumTemplate:Carbonate, is a sodium salt of carbonic acid. It most commonly occurs as a crystaline heptahydrate which readily effloresces to form a white powder, the monohydrate. It has a cooling alkaline taste, and can be extracted from the ashes of many plants. It is produced artificially in large quantities from common salt.
Uses
Sodium carbonate is used in the manufacture of glass, pulp and paper, detergents, and chemicals such as sodium silicates and sodium phosphates. It is also used as an alkaline agent in many chemical industries.
Domestically it is used as a water softener during laundry. It competes with the ions magnesium and calcium in hard water and prevents them from bonding with the detergent being used. Without using washing soda, additional detergent is needed to soak up the magnesium and calcium ions. Called washing soda in the detergent section of stores, it effectively removes oil, grease, and alcohol stains.
Sodium carbonate is widly used in photographic processes as a pH regulator to maintain stable alkaline conditions necessary for the action of the majority of developing agents.
Sodium carbonate is also used by the brick industry as a wetting agent to reduce the amount of water needed to extrude the clay.
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 and in the early manufacture of glass. Sodium carbonate has three known forms of hydrates: 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, sulphuric acid, limestone, and coal. First, sea salt (sodium chloride) was boiled in sulphuric acid to yield sodium sulphate and hydrochloric acid gas, according to the chemical equation
Next, the sodium sulphate 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 sulphide 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.