Solvay process

The Solvay process, also referred to as the ammonia-soda process, is the major industrial process for the production of soda ash (sodium carbonate). The ammonia-soda process developed into its modern form by Ernest Solvay during the 1860s. The ingredients for this process are readily available and inexpensive: salt brine (from inland sources or from the sea) and limestone (from mines). The worldwide production of soda ash in 2005 has been estimated at 41.9 billion kilograms,^[1] which is several kilograms for each person on earth. Solvay-based chemical plants now produce roughly three-fourths of this supply, with the remainder being mined from natural deposits.

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Chemistry

  The Solvay process produces soda ash (predominantly sodium carbonate (Na₂CO₃)) from brine (as a source of sodium chloride (NaCl)) and from limestone (as a source of calcium carbonate (CaCO₃)).^[2] The overall process is:

2 NaCl + CaCO₃ → Na₂CO₃ + CaCl₂

The actual implementation of this global, overall reaction is intricate. ^[3][4][5] A simplified description can be given using the four different, interacting chemical reactions illustrated in the figure. In the first step in the process, carbon dioxide (CO₂) passes through a concentrated aqueous solution of sodium chloride (NaCl) and ammonia (NH₃).

NaCl + CO₂ + NH₃ + H₂O → NaHCO₃ + NH₄Cl (I)

In industrial practice, the reaction is carried out by passing concentrated brine through two towers. In the first, ammonia bubbles up through the brine and is absorbed by it. In the second, carbon dioxide bubbles up through the ammoniated brine, and sodium bicarbonate (NaHCO₃) precipitates out of the solution. Note that, in a basic solution, NaHCO₃ is less water-soluble than sodium chloride. The ammonia (NH₃) buffers the solution at a basic pH; without the ammonia, a hydrochloric acid byproduct would render the solution acidic, and arrest the precipitation.

The necessary ammonia "catalyst" for reaction (I) is reclaimed in a later step, and relatively little ammonia is consumed. The carbon dioxide required for reaction (I) is produced by heating ("calcination") of the limestone at 950 - 1100 C. The calcium carbonate (CaCO₃) in the limestone is partially converted to quicklime (calcium oxide (CaO)) and carbon dioxide:

CaCO₃ → CO₂ + CaO (II)

The sodium bicarbonate (NaHCO₃) that precipitates out in reaction (I) is filtered out from the hot ammonium chloride (NH₄Cl) solution, and the solution is then reacted with the quicklime (calcium oxide (CaO)) left over from heating the limestone in step (II).

2 NH₄Cl + CaO → 2 NH₃ + CaCl₂ + H₂O (III)

CaO makes a strong basic solution. The ammonia from reaction (III) is recycled back to the initial brine solution of reaction (I).

The sodium bicarbonate (NaHCO₃) precipitate from reaction (I) is then converted to the final product, sodium carbonate (Na₂CO₃), by calcination (160 - 230 C), producing water and carbon dioxide as byproducts:

2 NaHCO₃ → Na₂CO₃ + H₂O + CO₂ (IV)

The carbon dioxide from step (IV) is recovered for re-use in step (I). When properly designed and operated, a Solvay plant can reclaim almost all its ammonia, and consumes only small amounts of additional ammonia to make up for losses. The only major inputs to the Solvay process are salt, limestone and thermal energy, and its only major byproduct is calcium chloride, which is sold as road salt.

Additional details of the industrial implementation of this process are available in the report prepared for the European Soda Ash Producer's Association.^[4]

Uses of Soda Ash

Main article: Soda ash

Soda ash is used in many industrial processes, and its production is sometimes used as an indicator of economic health. The principal current uses include:^[1]

• Glass making: More than half the worldwide production of soda ash is used to make glass. Bottle and window glass (Soda-lime glass) is made by melting a mixture of sodium carbonate, calcium carbonate and silica sand (silicon dioxide (SiO₂)).
• Water treatment: Sodium carbonate is used to soften water (precipitates out Mg²⁺ and Ca²⁺ carbonates). This is used both industrially and domestically (in some washing powders).
• Making soaps and detergents: Often sodium carbonate is used as a cheaper alternative to lye (sodium hydroxide (NaOH)).
• Paper making: Sodium carbonate is used to make sodium bisulfite (NaHSO₃) for the "sulfite" method of separating lignin from cellulose.
• As a common base in many chemical factories because it is cheaper than NaOH and far safer to handle.
• Making sodium bicarbonate: NaHCO₃ is used in baking soda and fire extinguishers. Although NaHCO₃ is produced in the Solvay process, heating it to remove the ammonia it is contaminated with decomposes some NaHCO₃, so it is actually cheaper to react the finished Na₂CO₃ product with CO₂.
• Removing sulfur dioxide (SO₂) from flue gases in power stations. This is becoming more common, especially where stations have to meet stringent emission controls.

History

See also: alkali, potash, soda ash, Leblanc process, and trona

The name "soda ash" is based on the principal historical method of obtaining alkali, which was by using water to extract it from ashes. Wood fires yielded potash and the active ingredient potassium carbonate. The word "soda" (from the Middle Latin) originally referred to certain plants that grow in salt marshes; it was discovered that the ashes of these plants yielded the useful alkali "soda ash." The cultivation of such plants for production of soda ash reached a particularly high state of development in the 18th Century in Spain, where the plants are named barrilla; the English word is "barilla."^[6][7][8] The ashes of kelp also yield soda ash, and were the basis of an enormous 18th Century industry in Scotland. Alkali was also mined from dry lakebeds in Egypt.

By the late 1700s, however, these sources were insufficient to meet Europe's burgeoning demand for alkali for soap, textile, and glass industries.^[9] In 1791, the French physician Nicolas Leblanc developed a method to manufacture soda ash using salt, limestone, sulfuric acid, and coal. Although the Leblanc process came to dominate alkali production in the early 1800s, the expense of its inputs and its polluting byproducts (including hydrochloric acid gas) made it apparent that it was far from an ideal solution.^[2][9]

It has been reported that, in 1811, the renowned French physicist Augustin Jean Fresnel discovered that sodium bicarbonate precipitates when carbon dioxide is bubbled through ammonia-containing brine— which is the chemical reaction central to the Solvay process. The discovery wasn't published. As has been noted by Desmond Reilly, "The story of the evolution of the ammonium-soda process is an interesting example of the way in which a discovery can be made and then laid aside and not applied for a considerable time afterwards."^[10] Serious consideration of this reaction as the basis of an industrial process dates from the British patent issued in 1834 to H. G. Dyan and J. Henning. There were several attempts to reduce this reaction to industrial practice, with varying success.

In 1861, the Belgian industrial chemist Ernest Solvay turned his attention to the problem; he was apparently largely unaware of the extensive earlier work.^[2] His solution, an 80-foot-tall gas absorption tower in which carbon dioxide bubbled up through a descending flow of brine, together with efficient recovery and recycling of the ammonia, proved effective, and by 1864, Solvay and his brother Alfred had acquired good financial backing and constructed a plant in the Belgian town of Charleroi. The new process proved more economical and less polluting than the Leblanc method, and its use spread. In 1874, the Solvays expanded their facilities with a new, larger plant at Nancy, France.

In the same year, Ludwig Mond visited Solvay in Belgium and acquired rights to use the new technology. He and John Brunner formed the firm of Brunner, Mond & Co., and built a Solvay plant at Winnington, near Northwich, Cheshire, England. The facility started up in 1874. Mond was instrumental in making the Solvay process a commercial success; he made several refinements between 1873 and 1880 that removed byproducts that could slow or halt the mass production of sodium carbonate through use of the process.

In 1884, the Solvay brothers licensed Americans William B. Cogswell and Rowland Hazard to produce soda ash in the United States, and formed a joint venture (Solvay Process Company) to build and operate a plant in Solvay, New York.

By the 1890s, Solvay process plants produced the majority of the world's soda ash.

In 1938, large natural deposits of the mineral Trona were discovered near the Green River in Wyoming. Sodium carbonate can be mined from this source less expensively than it can be produced by the Solvay process, and with the closing of the original Solvay, New York plant in 1986, there have been no Solvay-based plants operating in North America. Throughout the rest of the world, however, the Solvay process remains the major source of soda ash.

Byproducts and Wastes

The principal byproduct of the Solvay process is calcium chloride (CaCl₂) in aqueous solution. The process has other waste and byproducts as well.^[4] Not all of the limestone that is calcined is converted to quicklime and carbon dioxide (in reaction II); the residual calcium carbonate and other components of the limestone becomes wastes. In addition, the salt brine used by the process is usually purified to remove magnesium and calcium ions, typically to form carbonates; otherwise, these impurities would lead to scale in the various reaction vessels and towers. These carbonates are additional waste products.

In inland plants, such as that in Solvay, New York, the byproducts have been deposited in "waste beds;" the weight of material deposited in these waste beds exceeded that of the soda ash produced by about 50%. These waste beds have led to water pollution, principally by calcium and chlorine ions. The waste beds in Solvay, New York substantially increased the salinity in nearby Onondaga Lake, which is among the most polluted lakes in the U.S. ^[11] and is a superfund pollution site.^[12] As such waste beds age, they do begin to support plant communities which have been the subject of several scientific studies. ^[13][14]

At seaside locations, such as those at Saurashtra, Gujarat, India,^[15] and at Osborne, South Australia,^[16] the CaCl₂ solution may be discharged directly into the sea, apparently without substantial environmental harm. In the "modified" Solvay process, the CaCl₂ is supplanted by ammonium chloride (NH₄Cl). NH₄Cl, which can be used in fertilizer, may have greater commercial value than CaCl₂, thus reducing the extent of waste beds.

Carbon Sequestration and the Solvay Process

Main article: carbon sequestration

Variations in the Solvay process have been proposed for carbon sequestration. One idea is to react carbon dioxide, produced perhaps by the combustion of coal, to form solid carbonates (such as sodium bicarbonate) that could be permanently stored, thus avoiding carbon dioxide emission into the atmosphere. ^{[17] [18]} Variations in the Solvay process have been proposed to convert carbon dioxide emissions into sodium carbonates, but carbon sequestration by calcium or magnesium carbonates appears more promising.

References

1. ^ ^{a b} Kostick, Dennis (2006). "Soda Ash", chapter in 2005 Minerals Yearbook, United States Geological Survey. See Table I.
2. ^ ^{a b c} Kiefer, David M. (February 2002). "Soda Ash, Solvay Style". Today's Chemist at Work 11 (2): 87-88, 90.
3. ^ Speight, James (2001). Chemical Process and Design Handbook. McGraw Hill. DOI:10.1036/0071374337. ISBN 0071374337. 
4. ^ ^{a b c} "Process Best Practices Reference Document (BREF) for Soda Ash," report produced by the European Soda Ash Producer's Association, March 2004. Retrieved 2006-10-14.
5. ^ Moore, John T. Edd (2005). Chemistry Made Simple. Broadway Books, 190. ISBN 0767917022. 
6. ^ The barilla used for soda ash production refers to any of several bushy plants that are well adapted to grow in salt marshes, and that are common in Spain and Italy. The ashes of these plants can contain as much as 30% sodium carbonate. The principal species for soda ash production were the "saltworts" Salsola soda or Salsola kali, but several other species could also be used.
7. ^ Pérez, Joaquín Fernández (1998). "From the barrilla to the Solvay factory in Torrelavega: The Manufacture of Saltwort in Spain," Antilia: The Spanish Journal of History of Natural Sciences and Technology, Vol. IV, Art. 1. ISSN 1136-2049. Retrieved Oct. 21, 2006.
8. ^ Grieve, M. (1931). A Modern Herbal, ISBN 0-486-22798-7 & 0486227995. See section on glasswort. Retrieved October 21, 2005.
9. ^ ^{a b} Kiefer, David M.. "It was all about alkali". Today's Chemist at Work 11 (1): 45-6.
10. ^ Reilly, Desmond (December 1951). "Salts, Acids & Alkalis in the 19th Century. A Comparison between Advances in France, England & Germany". Isis 42 (4): 287-296.
11. ^ Onondaga Lake Partnership. Retrieved 2006-10-14.
12. ^ U.S. Environmental Protection Agency, superfund ID NYD986913580. Retrieved 2006-10-14.
13. ^ Cohn, E.V.J., Rostanski, A., Tokarska-Guzik, B., Trueman, I.C., and Wozniak, G. (2001). "The flora and vegetation of an old Solvay process tip in Jaworzno (Upper Silesia, Poland)." Acta Soc. Bot. Pol. 70(1):47-60.
14. ^ Michalenko, Edward M. (1991). "Pedogenesis and invertebrate microcommunity succession in immature soils originating from chlor-alkali wastes," doctoral dissertation, College of Environmental Science and Forestry, State University of New York.
15. ^ "Technology in the Indian Soda Ash Industry", Technology Status Report #148 (October, 1995), Department of Scientific and Industrial Research, Ministry of Science & Technology, India. Retrieved 2006-10-14.
16. ^ Penrice Soda Holdings Limited. Retrieved 2006-10-14.
17. ^ Huijgen, W.J.J. and Comans, R.N.J. (February, 2003). "Carbon dioxide sequestration by mineral carbonation: Literature Review," Report ECN C-03-016, Energy Research Centre of the Netherlands. Retrieved 2006-10-14.
18. ^ Lackner, Klaus S (2002). "Carbonate Chemistry for Sequestering Fossil Carbon". Annual Review of Energy and the Environment 27 (1): 193-232. doi:doi:10.1146/annurev.energy.27.122001.083433.