Sodium carbonate

Sodium carbonate
Skeletal formula of sodium carbonate
Sample of sodium carbonate
IUPAC name

Sodium carbonate
Other names

Soda ash, washing soda, soda crystals
3D model (JSmol)
ECHA InfoCard 100.007.127
EC Number 207-838-8
E number E500(i) (acidity regulators, …)
RTECS number VZ4050000
Molar mass 105.9888 g/mol (anhydrous)
286.1416 g/mol (decahydrate)
Appearance White solid, hygroscopic
Odor Odorless
  • 2.54 g/cm3 (25 °C, anhydrous)
  • 1.92 g/cm3 (856 °C)
  • 2.25 g/cm3 (monohydrate)[1]
  • 1.51 g/cm3 (heptahydrate)
  • 1.46 g/cm3 (decahydrate)[2]
Melting point 851 °C (1,564 °F; 1,124 K) (Anhydrous)
100 °C (212 °F; 373 K)
decomposes (monohydrate)
33.5 °C (92.3 °F; 306.6 K)
decomposes (heptahydrate)
34 °C (93 °F; 307 K)
Anhydrous, g/100 mL:

  • 7 (0 °C)
  • 16.4 (15 °C)
  • 34.07 (27.8 °C)
  • 48.69 (34.8 °C)
  • 48.1 (41.9 °C)
  • 45.62 (60 °C)
  • 43.6 (100 °C)[3]
Solubility Soluble in aq. alkalis,[3] glycerol
Slightly soluble in aq. alcohol
Insoluble in CS2, acetone, alkyl acetates, alcohol, benzonitrile, liquid ammonia[4]
Solubility in glycerine 98.3 g/100 g (155 °C)[4]
Solubility in ethanediol 3.46 g/100 g (20 °C)[5]
Solubility in dimethylformamide 0.5 g/kg[5]
Basicity (pKb) 3.67
−4.1·10−5 cm3/mol[2]
1.485 (anhydrous)
1.420 (monohydrate)[6]
1.405 (decahydrate)
Viscosity 3.4 cP (887 °C)[5]
Monoclinic (γ-form, β-form, δ-form, anhydrous)[7]
Orthorhombic (monohydrate, heptahydrate)[1][8]
C2/m, No. 12 (γ-form, anhydrous, 170 K)
C2/m, No. 12 (β-form, anhydrous, 628 K)
P21/n, No. 14 (δ-form, anhydrous, 110 K)[7]
Pca21, No. 29 (monohydrate)[1]
Pbca, No. 61 (heptahydrate)[8]
2/m (γ-form, β-form, δ-form, anhydrous)[7]
mm2 (monohydrate)[1]
2/m 2/m 2/m (heptahydrate)[8]
a = 8.920(7) Å, b = 5.245(5) Å, c = 6.050(5) Å (γ-form, anhydrous, 295 K)[7]
α = 90°, β = 101.35(8)°, γ = 90°
Octahedral (Na+, anhydrous)
112.3 J/mol·K[2]
135 J/mol·K[2]
−1130.7 kJ/mol[2][5]
−1044.4 kJ/mol[2]
Main hazards Irritant
Safety data sheet MSDS
GHS pictograms The exclamation-mark pictogram in the Globally Harmonized System of Classification and Labelling of Chemicals (GHS)[9]
GHS signal word Warning
NFPA 704
Flammability code 0: Will not burn. E.g., water Health code 1: Exposure would cause irritation but only minor residual injury. E.g., turpentine Reactivity code 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g., liquid nitrogen Special hazards (white): no code

NFPA 704 four-colored diamond

Lethal dose or concentration (LD, LC):
4090 mg/kg (rat, oral) [10]
Related compounds
Other anions
Sodium bicarbonate
Other cations
Lithium carbonate
Potassium carbonate
Rubidium carbonate
Caesium carbonate
Related compounds
Sodium sesquicarbonate
Sodium percarbonate
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Infobox references

Sodium carbonate, Na2CO3, (also known as washing soda, soda ash and soda crystals) is the inorganic compound with the formula Na2CO3 and its various hydrates. All forms are white, water-soluble salts. All forms have a strongly alkaline taste and give moderately alkaline solutions in water. Historically it was extracted from the ashes of plants growing in sodium-rich soils. Because the ashes of these sodium-rich plants were noticeably different from ashes of wood (once used to produce potash), sodium carbonate became known as “soda ash”.[12] It is produced in large quantities from sodium chloride and limestone by the Solvay process.


Sodium carbonate is obtained as three different hydrates and as the anhydrous salt:

  • sodium carbonate decahydrate (natron), Na2CO3·10H2O, which readily effloresces to form the monohydrate.
  • sodium carbonate heptahydrate (not known in mineral form), Na2CO3·7H2O.
  • sodium carbonate monohydrate (thermonatrite), Na2CO3·H2O. Also known as crystal carbonate.
  • anhydrous sodium carbonate, also known as calcined soda, is formed by heating the hydrates. It is also formed when sodium hydrogen carbonate is heated (calcined) e.g. in the final step of the Solvay process.

The decahydrate is formed from water solutions crystallizing in the temperature range -2.1 to +32.0 C, the heptahydrate in the narrow range 32.0 to 35.4 C and above this temperature the monohydrate forms.[13] In dry air the decahydrate and heptahydrate lose water to give the monohydrate. Other hydrates have been reported, e.g. with 2.5 units of water per sodium carbonate unit (“pentahemihydrate”).[14]


Main applications[edit]

In terms of its largest applications, sodium carbonate is used in the manufacture of glass, paper, rayon, soaps, and detergents.[15]

Glass manufacture[edit]

Sodium carbonate serves as a flux for silica, lowering the melting point of the mixture to something achievable without special materials. This “soda glass” is mildly water-soluble, so some calcium carbonate is added to the melt mixture to make the glass produced insoluble. Bottle and window glass (Soda-lime glass) is made by melting such mixtures of sodium carbonate, calcium carbonate, and silica sand (silicon dioxide (SiO2)). When these materials are heated, the carbonates release carbon dioxide. In this way, sodium carbonate is a source of sodium oxide. Soda lime glass has been the most common form of glass for centuries.

Water softening[edit]

Sodium carbonate is used to soften water by removing Mg2+ and Ca2+. These ions form insoluble solid precipitates upon treatment with carbonate ions:

Ca2+ + CO32- → CaCO3

Sodium carbonate is an inexpensive and water-soluble source of carbonate ions.

Food additive and cooking[edit]

Sodium carbonate is a food additive (E500) used as an acidity regulator, anticaking agent, raising agent, and stabilizer. It is one of the components of kansui (かん水), a solution of alkaline salts used to give ramen noodles their characteristic flavor and texture. It is used in the production of snus to stabilize the pH of the final product. Sodium carbonate is used in the production of sherbet powder. The cooling and fizzing sensation results from the endothermic reaction between sodium carbonate and a weak acid, commonly citric acid, releasing carbon dioxide gas, which occurs when the sherbet is moistened by saliva. In China, it is used to replace lye-water in the crust of traditional Cantonese moon cakes, and in many other Chinese steamed buns and noodles. In cooking, it is sometimes used in place of sodium hydroxide for lyeing, especially with German pretzels and lye rolls. These dishes are treated with a solution of an alkaline substance to change the pH of the surface of the food and improve browning.

Inexpensive, weak base[edit]

Sodium carbonate is also used as a relatively strong base in various fields. As a common alkali, it is preferred in many chemical processes because it is cheaper than NaOH and far safer to handle. Its mildness especially recommends its use in domestic applications.

For example, it is used as a pH regulator to maintain stable alkaline conditions necessary for the action of the majority of photographic film developing agents. For example, it is a common additive in swimming pools and aquarium water to maintain a desired pH and carbonate hardness (KH). In dyeing with fiber-reactive dyes, sodium carbonate (often under a name such as soda ash fixative or soda ash activator) is used to ensure proper chemical bonding of the dye with cellulose (plant) fibers, typically before dyeing (for tie dyes), mixed with the dye (for dye painting), or after dyeing (for immersion dyeing).

Sodium bicarbonate (NaHCO3) or baking soda, also a component in fire extinguishers, is often generated from sodium carbonate. Although NaHCO3 is itself an intermediate product of the Solvay process, the heating needed to remove the ammonia that contaminates it decomposes some NaHCO3, making it more economic to react finished Na2CO3 with CO2:

Na2CO3 + CO2 + H2O → 2NaHCO3

In a related reaction, sodium carbonate is used to make sodium bisulphite (NaHSO3), which is used for the “sulfite” method of separating lignin from cellulose. This reaction is exploited for removing sulphur dioxide from flue gases in power stations:

Na2CO3 + SO2 + H2O → NaHCO3 + NaHSO3

This application has become more common, especially where stations have to meet stringent emission controls.

Sodium carbonate is used by the cotton industry to neutralize the sulfuric acid needed for acid delinting of fuzzy cottonseed.


Sodium carbonate is used by the brick industry as a wetting agent to reduce the amount of water needed to extrude the clay. In casting, it is referred to as “bonding agent” and is used to allow wet alginate to adhere to gelled alginate.Sodium carbonate is used in toothpastes, where it acts as a foaming agent and an abrasive, and to temporarily increase mouth pH.

Physical properties[edit]

The integral enthalpy of solution of sodium carbonate is −28.1 kJ/mol for a 10% w/w aqueous solution.[16] The Mohs hardness of sodium carbonate monohydrate is 1.3.[6]

Occurrence as natural mineral[edit]

Structure of monohydrate at 346 K.

Sodium carbonate is soluble in water, and can occur naturally in arid regions, especially in mineral deposits (evaporites) formed when seasonal lakes evaporate. Deposits of the mineral natron 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.

The anhydrous mineral form of sodium carbonate is quite rare and called natrite. Sodium carbonate also erupts from Ol Doinyo Lengai, Tanzania’s unique volcano, and it is presumed to have erupted from other volcanoes in the past, but due to these minerals’ instability at the earth’s surface, are likely to be eroded. All three mineralogical forms of sodium carbonate, as well as trona, trisodium hydrogendicarbonate dihydrate, are also known from ultra-alkaline pegmatitic rocks, that occur for example in the Kola Peninsula in Russia.

Extraterrestrially, known sodium carbonate is rare. Deposits have been identified as the source of bright spots on Ceres, interior material that has been brought to the surface.[17] While there are carbonates on Mars, and these are expected to include sodium carbonate,[18] deposits have yet to be confirmed, this absence is explained by some as being due to a global dominance of low pH in previously aqueous Martian soil.[19]



Trona, trisodium hydrogendicarbonate dihydrate (Na3HCO3CO3·2H2O), is mined in several areas of the US and provides nearly all the domestic consumption of sodium carbonate. Large natural deposits found in 1938, such as the one near Green River, Wyoming, have made mining more economical than industrial production in North America.
There are important reserves of trona in Turkey; two million tons of soda ash have been extracted from the reserves near Ankara.
It is also mined from some alkaline lakes such as Lake Magadi in Kenya by dredging. Hot saline springs continuously replenish salt in the lake so that, provided the rate of dredging is no greater than the replenishment rate, the source is fully sustainable.[citation needed]

Barilla and kelp[edit]

Several “halophyte” (salt-tolerant) plant species and seaweed species can be processed to yield an impure form of sodium carbonate, and these sources predominated in Europe and elsewhere until the early 19th century. The land plants (typically glassworts or saltworts) or the seaweed (typically Fucus species) were harvested, dried, and burned. The ashes were then “lixiviated” (washed with water) to form an alkali solution. This solution was boiled dry to create the final product, which was termed “soda ash”; this very old name refers to the archetypal plant source for soda ash, which was the small annual shrub Salsola soda (“barilla plant”).

The sodium carbonate concentration in soda ash varied very widely, from 2–3 percent for the seaweed-derived form (“kelp“), to 30 percent for the best barilla produced from saltwort plants in Spain. Plant and seaweed sources for soda ash, and also for the related alkalipotash“, became increasingly inadequate by the end of the 18th century, and the search for commercially viable routes to synthesizing soda ash from salt and other chemicals intensified.[20]

Leblanc process[edit]

In 1792, the French chemist Nicolas Leblanc patented a process for producing sodium carbonate from salt, sulphuric acid, limestone, and coal. In the first step, sodium chloride is treated with sulfuric acid in the Mannheim process. This reaction produces sodium sulfate (salt cake) and hydrogen chloride:

2NaCl + H2SO4 → Na2SO4 + 2HCl

The salt cake and crushed limestone (calcium carbonate) was reduced by heating with coal.[15] This conversion entails two parts. First is the carbothermic reaction whereby the coal, a source of carbon, reduces the sulfate to sulfide:

Na2SO4 + 2C → Na2S + 2CO2

The second stage is the reaction to produce sodium carbonate and calcium sulfide:

Na2S + CaCO3 → Na2CO3 + CaS

This mixture is called black ash. The soda ash is extracted from the black ash with water. Evaporation of this extract yields solid sodium carbonate. This extraction process was termed lixiviation.

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.[20][21]

Solvay process[edit]

In 1861, the Belgian industrial chemist Ernest Solvay developed a method to convert sodium chloride to sodium carbonate using ammonia and carbon dioxide:[15]

NaCl + NH3 + CO2 + H2O → NaHCO3 + NH4Cl

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

2NaHCO3 → Na2CO3 + H2O + CO2

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

2NH4Cl + CaO → 2NH3 + CaCl2 + H2O

The Solvay process recycles its ammonia. It consumes only brine and limestone and calcium chloride is its only waste product. The process is substantially more economical than the Leblanc process, which generates two waste products, calcium sulfide and hydrogen chloride. The Solvay process quickly came to dominate sodium carbonate production worldwide. By 1900, 90% of sodium carbonate was produced by the Solvay process, and the last Leblanc process plant closed in the early 1920s.[15]

Hou’s process[edit]

This process was developed by Chinese chemist Hou Debang in the 1930s. The earlier steam reforming byproduct carbon dioxide was pumped through a saturated solution of sodium chloride and ammonia to produce sodium bicarbonate by these reactions:

CH4 + 2H2OCO2 + 4H2
3H2 + N2 → 2NH3
NH3 + CO2 + H2ONH4HCO3
NH4HCO3 + NaClNH4Cl + NaHCO3

The sodium bicarbonate was collected as a precipitate due to its low solubility and then heated up to approximately 80 °C (176 °F) or 95 °C (203 °F) to yield pure sodium carbonate similar to last step of the Solvay process. More sodium chloride is added to the remaining solution of ammonium and sodium chlorides; also, more ammonia is pumped at 30-40 °C to this solution. The solution temperature is then lowered to below 10 °C. Solubility of ammonium chloride is higher than that of sodium chloride at 30 °C and lower at 10 °C. Due to this temperature-dependent solubility difference and the common-ion effect, ammonium chloride is precipitated in a sodium chloride solution.

The Chinese name of Hou’s process, lianhe zhijian fa (联合制碱法), means “coupled manufacturing alkali method”: Hou’s process is coupled to the Haber process and offers better atom economy by eliminating the production of calcium chloride, since ammonia no longer needs to be regenerated. The byproduct ammonium chloride can be sold as a fertilizer.

See also[edit]


  1. ^ a b c d Harper, J.P (1936). Antipov, Evgeny; Bismayer, Ulrich; Huppertz, Hubert; Petrícek, Václav; Pöttgen, Rainer; Schmahl, Wolfgang; Tiekink, E.R.T.; Zou, Xiaodong, eds. “Crystal Structure of Sodium Carbonate Monohydrate, Na2CO3. H2O”. Zeitschrift für Kristallographie – Crystalline Materials. 95 (1): 266–273. doi:10.1524/zkri.1936.95.1.266. ISSN 2196-7105. Retrieved 2014-07-25.
  2. ^ a b c d e f g Lide, David R., ed. (2009). CRC Handbook of Chemistry and Physics (90th ed.). Boca Raton, Florida: CRC Press. ISBN 978-1-4200-9084-0.
  3. ^ a b Seidell, Atherton; Linke, William F. (1919). Solubilities of Inorganic and Organic Compounds (2nd ed.). New York: D. Van Nostrand Company. p. 633.
  4. ^ a b Comey, Arthur Messinger; Hahn, Dorothy A. (February 1921). A Dictionary of Chemical Solubilities: Inorganic (2nd ed.). New York: The MacMillan Company. pp. 208–209.
  5. ^ a b c d Anatolievich, Kiper Ruslan. “sodium carbonate”. Retrieved 2014-07-25.
  6. ^ a b c Pradyot, Patnaik (2003). Handbook of Inorganic Chemicals. The McGraw-Hill Companies, Inc. p. 861. ISBN 978-0-07-049439-8.
  7. ^ a b c d Dusek, Michal; Chapuis, Gervais; Meyer, Mathias; Petricek, Vaclav (2003). “Sodium carbonate revisited” (PDF). Acta Crystallographica Section B. 59 (3): 337–352. doi:10.1107/S0108768103009017. ISSN 0108-7681. Retrieved 2014-07-25.
  8. ^ a b c Betzel, C.; Saenger, W.; Loewus, D. (1982). “Sodium Carbonate Heptahydrate”. Acta Crystallographica Section B. 38 (11): 2802–2804. doi:10.1107/S0567740882009996.
  9. ^ a b c Sigma-Aldrich Co., Sodium carbonate. Retrieved on 2014-05-06.
  10. ^ Chambers, Michael. “ChemIDplus – 497-19-8 – CDBYLPFSWZWCQE-UHFFFAOYSA-L – Sodium carbonate [NF] – Similar structures search, synonyms, formulas, resource links, and other chemical information”.
  11. ^ “Material Safety Data Sheet – Sodium Carbonate, Anhydrous” (PDF). ConservationSupportSystems. Retrieved 2014-07-25.
  12. ^ “” (PDF).
  13. ^ T.W.Richards and A.H. Fiske (1914). “On the transition temperatures of the transition temperatures of the hydrates of sodium carbonate as fix points in thermometry”. Journal of the American Chemical Society. 36 (3): 485–490. doi:10.1021/ja02180a003.
  14. ^ A. Pabst. “On the hydrates of sodium carbonate”.[permanent dead link]
  15. ^ a b c d Christian Thieme (2000). “Sodium Carbonates”. Ullmann’s Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a24_299. ISBN 978-3527306732.
  16. ^ “”.
  17. ^ De Sanctis, M. C.; et al. (29 June 2016). “Bright carbonate deposits as evidence of aqueous alteration on (1) Ceres”. Nature. 536 (7614): 54–57. Bibcode:2016Natur.536…54D. doi:10.1038/nature18290. PMID 27362221. Retrieved 2016-06-30.
  18. ^ Jeffrey S. Kargel (23 July 2004). Mars – A Warmer, Wetter Planet. Springer Science & Business Media. pp. 399–. ISBN 978-1-85233-568-7.
  19. ^ Grotzinger, J. and R. Milliken (eds.) 2012. Sedimentary Geology of Mars. SEPM
  20. ^ a b
    Clow, Archibald and Clow, Nan L. (1952). Chemical Revolution, (Ayer Co Pub, June 1952), pp. 65–90. ISBN 0-8369-1909-2.
  21. ^ Kiefer, David M. (January 2002). “It was all about alkali”. Today’s Chemist at Work. 11 (1): 45–6.

Further reading[edit]

External links[edit]

H2CO3 He
BeCO3 B C (NH4)2CO3,
O F Ne
Al2(CO3)3 Si P S Cl Ar
Sc Ti V Cr MnCO3 FeCO3 CoCO3 NiCO3 CuCO3 ZnCO3 Ga Ge As Se Br Kr
Rb2CO3 SrCO3 Y Zr Nb Mo Tc Ru Rh Pd Ag2CO3 CdCO3 In Sn Sb Te I Xe
BaCO3   Hf Ta W Re Os Ir Pt Au Hg Tl2CO3 PbCO3 (BiO)2CO3 Po At Rn
Fr Ra   Rf Db Sg Bh Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Og
La2(CO3)3 Ce2(CO3)3 Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
Ac Th Pa UO2CO3 Np Pu Am Cm Bk Cf Es Fm Md No Lr