View source for Sodium carbonate

{{Short description|Chemical compound}}
{{Distinguish|text= [[Sodium bicarbonate]] (baking soda), a similar compound}}
{{chembox
| Verifiedfields = changed
| Watchedfields  = changed
| verifiedrevid  = 464362057
| Name           = Sodium carbonate
| ImageFile      = Sodium carbonate.svg
| ImageSize      = 120px
| ImageName      = Skeletal formula of sodium carbonate
| ImageFile1     = Uhličitan sodný.JPG
| ImageFile2     = Sodium-carbonate-xtal-3D-SF-C.png
| ImageName1     = Sample of sodium carbonate
| IUPACName      = Sodium carbonate
| OtherNames     = Soda ash, washing soda, soda crystals, sodium trioxocarbonate
| Section1       = {{Chembox Identifiers
|CASNo_Ref = {{cascite|correct|CAS}}
|CASNo = 497-19-8
|CASNo_Comment = (anhydrous)
|CASNo1_Ref = {{cascite|correct|CAS}}
|CASNo1 = 5968-11-6
|CASNo1_Comment = (monohydrate)
|CASNo2_Ref = {{cascite|correct|CAS}}
|CASNo2 = 6132-02-1
|CASNo2_Comment = (decahydrate)
|ChEMBL = 186314
|ChEMBL_Ref = {{ebicite|correct|EBI}}
|PubChem = 10340
|ChemSpiderID = 9916
|ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}
|RTECS = VZ4050000
|EC_number = 207-838-8
|UNII = 45P3261C7T
|UNII_Ref = {{fdacite|correct|FDA}}
|UNII1_Ref = {{fdacite|correct|FDA}}
|UNII1 = 2A1Q1Q3557
|UNII1_Comment = (monohydrate)
|UNII2_Ref = {{fdacite|correct|FDA}}
|UNII2 = LS505BG22I
|UNII2_Comment = (decahydrate)
|InChI = 1/NaHCO3.2Na/c2-1(3)4;;/h(H2,2,3,4);;/q;2*+1/p-2
|InChIKey = CDBYLPFSWZWCQE-NUQVWONBAP
|ChEBI = 29377
|ChEBI_Ref = {{ebicite|correct|EBI}}
|SMILES = [Na+].[Na+].[O-]C([O-])=O
|StdInChI_Ref = {{stdinchicite|correct|chemspider}}
|StdInChI = 1S/CH2O3.2Na/c2-1(3)4;;/h(H2,2,3,4);;/q;2*+1/p-2
|StdInChIKey_Ref = {{stdinchicite|correct|chemspider}}
|StdInChIKey = CDBYLPFSWZWCQE-UHFFFAOYSA-L
}}
| Section2       = {{Chembox Properties
|Formula = Na<sub>2</sub>CO<sub>3</sub>
|MolarMass = 105.9888{{nbsp}}g/mol (anhydrous)<br />286.1416{{nbsp}}g/mol (decahydrate)
|Appearance = White solid, [[hygroscopic]]
|Odor = Odorless
|Density = {{ubl
| 2.54{{nbsp}}g/cm<sup>3</sup> (25&nbsp;°C, anhydrous)
| 1.92{{nbsp}}g/cm<sup>3</sup> (856&nbsp;°C)
| 2.25{{nbsp}}g/cm<sup>3</sup> (monohydrate)<ref name=cod />
| 1.51{{nbsp}}g/cm<sup>3</sup> (heptahydrate)
| 1.46{{nbsp}}g/cm<sup>3</sup> (decahydrate)<ref name=crc />
}}
| Solubility = Anhydrous, g/100{{nnbsp}}mL:{{ubl
| 7 (0&nbsp;°C)
| 16.4 (15&nbsp;°C)
| 34.07 (27.8&nbsp;°C)
| 48.69 (34.8&nbsp;°C)
| 48.1 (41.9&nbsp;°C)
| 45.62 (60&nbsp;°C)
| 43.6 (100&nbsp;°C)<ref name=sioc>{{cite book|last1=Seidell |first1=Atherton |last2=Linke |first2=William F. |year=1919 |title=Solubilities of Inorganic and Organic Compounds |url=https://archive.org/details/solubilitiesino01seidgoog |publisher=D. Van Nostrand Company |place=[[New York City|New York]] |edition=2nd |page=[https://archive.org/details/solubilitiesino01seidgoog/page/n658 633]}}</ref>
}}
|SolubleOther = Soluble in aq. [[alkali]]s,<ref name=sioc /> [[glycerol]]<br /> Slightly soluble in aq. [[ethanol|alcohol]]<br /> Insoluble in [[carbon disulfide|CS<sub>2</sub>]], [[acetone]], alkyl [[acetate]]s, alcohol, [[benzonitrile]], liquid [[ammonia]]<ref name=doc00>{{cite book|title = A Dictionary of Chemical Solubilities: Inorganic|url = https://archive.org/details/in.ernet.dli.2015.163725|edition = 2nd|first1 = Arthur Messinger|last1 = Comey|first2 = Dorothy A.|last2 = Hahn|place = New York|publisher = The MacMillan Company|date = February 1921|pages = 208–209}}</ref>
|Solubility1 = 98.3{{nbsp}}g/100{{nnbsp}}g (155&nbsp;°C)<ref name=doc00 />
|Solvent1 = glycerine
|Solubility2 = 3.46{{nbsp}}g/100{{nnbsp}}g (20&nbsp;°C)<ref name=chemister />
|Solvent2 = ethanediol
|Solubility3 = 0.5{{nbsp}}g/kg<ref name=chemister />
|Solvent3 = dimethylformamide
|MeltingPtC = 851
|MeltingPt_notes = (Anhydrous)<br /> {{convert|100|C|F K}}<br /> decomposes (monohydrate)<br /> {{convert|33.5|C|F K}}<br /> decomposes (heptahydrate)<br /> {{convert|34|C|F K}}<br /> (decahydrate)<ref name=crc>{{CRC90}}</ref><ref name=pphoic>{{cite book|last = Pradyot|first = Patnaik|year = 2003|title = Handbook of Inorganic Chemicals|publisher = McGraw-Hill |isbn = 978-0-07-049439-8|page = 861}}</ref>
|RefractIndex = 1.485 (anhydrous)<br /> 1.420 (monohydrate)<ref name=pphoic /><br /> 1.405 (decahydrate)
|MagSus = −4.1·10<sup>−5</sup> cm<sup>3</sup>/mol<ref name=crc />
|Viscosity = 3.4 cP (887&nbsp;°C)<ref name=chemister />
}}
| Section3       = {{Chembox Structure
| CrystalStruct = [[Monoclinic]] (γ-form, β-form, δ-form, anhydrous)<ref name=scr>{{cite journal|title = Sodium carbonate revisited|first1 = Michal|last1 = Dusek|first2 = Gervais|last2 = Chapuis|first3 = Mathias|last3 = Meyer|first4 = Vaclav|last4 = Petricek|journal = [[Acta Crystallographica Section B]]|url = http://infoscience.epfl.ch/record/82110/files/publ_03_dusek_a.pdf|issn = 0108-7681|year = 2003|volume = 59|issue = 3|access-date = 2014-07-25|pages = 337–352|doi = 10.1107/S0108768103009017|pmid = 12761404| bibcode=2003AcCrB..59..337D }}</ref><br /> [[Orthorhombic]] (monohydrate, heptahydrate)<ref name=cod>{{cite journal|title = Crystal Structure of Sodium Carbonate Monohydrate, Na<sub>2</sub>CO<sub>3</sub>. H<sub>2</sub>O|first = J. P.|last = Harper|journal = Zeitschrift für Kristallographie - Crystalline Materials|url = http://www.crystallography.net/1011295.html|issn = 2196-7105|access-date = 2014-07-25|pages = 266–273|year = 1936|volume = 95|issue = 1|doi = 10.1524/zkri.1936.95.1.266|editor-last1 = Antipov|editor-first1 = Evgeny|editor-last2 = Bismayer|editor-first2 = Ulrich|editor-last3 = Huppertz|editor-first3 = Hubert|editor-last4 = Petrícek|editor-first4 = Václav|editor-last5 = Pöttgen|editor-first5 = Rainer|editor-last6 = Schmahl|editor-first6 = Wolfgang|editor-last7 = Tiekink|editor-first7 = E. R. T. |editor-last8 = Zou|editor-first8 = Xiaodong|url-access = subscription}}</ref><ref name=7h2o>{{cite journal|title = Sodium Carbonate Heptahydrate|first1 = C.|last1 = Betzel|first2 = W.|last2 = Saenger|first3 = D.|last3 = Loewus|journal = Acta Crystallographica Section B|pages = 2802–2804|year = 1982|volume = 38|issue = 11|doi = 10.1107/S0567740882009996| bibcode=1982AcCrB..38.2802B }}</ref>
|SpaceGroup = C2/m, No. 12 (γ-form, anhydrous, 170&nbsp;K)<br /> C2/m, No. 12 (β-form, anhydrous, 628&nbsp;K)<br /> P2<sub>1</sub>/n, No. 14 (δ-form, anhydrous, 110&nbsp;K)<ref name=scr /><br /> Pca2<sub>1</sub>, No. 29 (monohydrate)<ref name=cod /><br /> Pbca, No. 61 (heptahydrate)<ref name=7h2o />
|PointGroup = 2/m (γ-form, β-form, δ-form, anhydrous)<ref name=scr /><br /> mm2 (monohydrate)<ref name=cod /><br /> 2/m 2/m 2/m (heptahydrate)<ref name=7h2o />
|LattConst_a = 8.920(7)&nbsp;Å
|LattConst_b = 5.245(5)&nbsp;Å
|LattConst_c = 6.050(5)&nbsp;Å (γ-form, anhydrous, 295&nbsp;K)<ref name=scr />
|LattConst_beta = 101.35(8)
|Coordination = Octahedral (Na<sup>+</sup>, anhydrous)
}}
| Section4       = {{Chembox Thermochemistry
|DeltaHf = −1130.7{{nbsp}}kJ/mol<ref name=crc /><ref name=chemister>{{cite web|last = Anatolievich|first = Kiper Ruslan|website =chemister.ru|url = http://chemister.ru/Database/properties-en.php?dbid=1&id=66|title = sodium carbonate|access-date = 2014-07-25}}</ref>
|Entropy = 135{{nbsp}}J/mol·K<ref name=crc />
|DeltaGf = −1044.4{{nbsp}}kJ/mol<ref name=crc />
|HeatCapacity = 112.3{{nbsp}}J/mol·K<ref name=crc />
}}
| Section5       = {{Chembox Hazards
|MainHazards = Irritant
|ExternalSDS = [https://web.archive.org/web/20080521003259/http://www.chem.tamu.edu/class/majors/msdsfiles/msdssodiumcarb.htm#Material MSDS]
|GHSPictograms = {{GHS07}}<ref name="sigma">{{Sigma-Aldrich|id=451614|name=Sodium carbonate|accessdate=2014-05-06}}</ref>
|GHSSignalWord = Warning
|HPhrases = {{H-phrases|319|313+333}}<ref name="sigma" />
|PPhrases = {{P-phrases|305+351+338}}<ref name="sigma" />
|NFPA-H = 2
|NFPA-F = 0
|NFPA-R = 0
|NFPA_ref = <ref name=css>{{cite web|title = Material Safety Data Sheet – Sodium Carbonate, Anhydrous|url = http://www.conservationsupportsystems.com/system/assets/msds/sodium_carbonate_msds.pdf|website =conservationsupportsystems.com|publisher = ConservationSupportSystems|access-date = 2014-07-25}}</ref>
|LD50 = 4090 mg/kg (rat, oral)<ref>{{cite web|url=https://pubchem.ncbi.nlm.nih.gov/#tab/sidsrcname=ChemIDplus&query=497-19-8&input_type=text|title=ChemIDplus - 497-19-8 - CDBYLPFSWZWCQE-UHFFFAOYSA-L - Sodium carbonate [NF] - Similar structures search, synonyms, formulas, resource links, and other chemical information|first=Michael|last=Chambers}}</ref>
}}
| Section6       = {{Chembox Related
|OtherAnions = [[Sodium bicarbonate]]
|OtherCations = [[Lithium carbonate]]<br /> [[Potassium carbonate]]<br />[[Rubidium carbonate]]<br /> [[Cesium carbonate]]
|OtherCompounds = [[Sodium sesquicarbonate]]<br /> [[Sodium percarbonate]]
}}
}}

'''Sodium carbonate''' (also known as '''washing soda''', '''soda ash''', '''sal soda''', and '''soda crystals''') is the [[inorganic compound]] with the formula {{chem2|Na2CO3|}} and its various [[hydrate]]s. All forms are white, odorless, water-soluble salts that yield alkaline solutions in water. Historically, it was extracted from the ashes of plants grown in sodium-rich soils, and 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".<ref>{{cite web|url=https://www.usgs.gov/centers/national-minerals-information-center/soda-ash-statistics-and-information|title=Soda Ash Statistics and Information|publisher=United States Geographical Survey|access-date=2024-03-03}}</ref> It is produced in large quantities from [[sodium chloride]] and [[limestone]] by the [[Solvay process]], as well as by carbonating sodium hydroxide which is made using the [[chloralkali process]].

==Hydrates==
Sodium carbonate is obtained as three [[hydrate]]s and as the anhydrous salt:
* sodium carbonate decahydrate ([[natron]]), Na<sub>2</sub>CO<sub>3</sub>·10H<sub>2</sub>O, which readily [[efflorescence|effloresces]] to form the monohydrate.
* sodium carbonate heptahydrate (not known in mineral form), Na<sub>2</sub>CO<sub>3</sub>·7H<sub>2</sub>O.
* sodium carbonate monohydrate ([[thermonatrite]]), Na<sub>2</sub>CO<sub>3</sub>·H<sub>2</sub>O. Also known as '''crystal carbonate'''.
* anhydrous sodium carbonate (natrite), also known as calcined soda, is formed by heating the hydrates. It is also formed when sodium hydrogencarbonate 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&nbsp;°C, the heptahydrate in the narrow range 32.0 to 35.4&nbsp;°C and above this temperature the monohydrate forms.<ref>{{cite journal|title=On the transition temperatures of the transition temperatures of the hydrates of sodium carbonate as fix points in thermometry |journal=Journal of the American Chemical Society |volume=36 |issue=3 |pages=485–490 |author= T.W.Richards and A.H. Fiske|doi=10.1021/ja02180a003 |year=1914 |url=https://zenodo.org/record/1428987}}</ref> 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 ("Penta hemihydrate").<ref>{{cite web |url=http://www.minsocam.org/ammin/am15/am15_69.pdf |author=A. Pabst |title=On the hydrates of sodium carbonate }}</ref>

===Washing soda===
Sodium carbonate decahydrate (Na<sub>2</sub>CO<sub>3</sub>·10H<sub>2</sub>O), also known as washing soda, is the most common hydrate of sodium carbonate containing 10 molecules of [[water of crystallization]]. Soda ash is dissolved in water and crystallized to get washing soda.

:{{chem2 | Na2CO3 + 10 H2O -> Na2CO3*10H2O }}

It is one of the few metal [[carbonate]]s that is soluble in water.

==Applications==
Some common applications of sodium carbonate include:
* As a cleansing agent for domestic purposes like washing clothes. Sodium carbonate is a component of many dry soap powders. It has [[detergent]] properties through the process of [[saponification]], which converts fats and grease to water-soluble [[Salt (chemistry)|salt]]s (specifically, soaps).<ref name=Ullmann/>
* It is used for lowering the [[hardness of water]]<ref name="Cornell" /> (see {{section link|#Water softening}}).
* It is used in the manufacture of [[glass]],<ref name="Himmel">{{Cite book |last1=Himmelblau |first1=David M. |title=Basic principles and calculations in chemical engineering |last2=Riggs |first2=James B. |date=2022 |publisher=Pearson |isbn=978-0-13-732717-1 |edition=Ninth |series=International series in the physical and chemical engineering sciences |location=Boston}}</ref> [[soap]],<ref name="Himmel" /> and [[paper]] {{xref|(see {{slink||Glass manufacture}})}}.
* It is used in the manufacture of sodium compounds like [[borax]] (sodium borate).{{np}}{{pb}}

===Glass manufacture===
Sodium carbonate serves as a [[Flux (metallurgy)|flux]] for [[silica]] (SiO<sub>2</sub>, melting point 1,713&nbsp;°C), 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 insoluble. Bottle and window glass ("[[soda–lime glass]]" with transition temperature ~570&nbsp;°C) is made by melting such mixtures of sodium carbonate, calcium carbonate, and silica sand ([[silicon dioxide]] (SiO<sub>2</sub>)). 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. It is also a key input for tableware glass manufacturing.<ref name=Ullmann/>

===Water softening===
{{See also|Hard water|Water softening}}
Hard water usually contains calcium or magnesium ions. Sodium carbonate is used for removing these ions and replacing them with sodium ions.<ref name="Cornell">{{cite web |url=https://www.ccmr.cornell.edu/wp-content/uploads/sites/2/2015/11/Water-Hardness-Reading.pdf |title=Water Hardness Reading |website=Cornell Center for Materials Research}}</ref>

Sodium carbonate is a water-soluble source of carbonate. The calcium and magnesium ions form insoluble solid precipitates upon treatment with [[carbonate]] ions:
{{block indent|{{chem2|Ca(2+)  +  CO3(2-) -> CaCO3 (s)}}}}
The water is softened because it no longer contains dissolved calcium ions and magnesium ions.<ref name="Cornell" />

===Food additive and cooking===
Sodium carbonate has several uses in cuisine, largely because it is a stronger base than baking soda ([[sodium bicarbonate]]) but weaker than [[lye]] (which may refer to [[sodium hydroxide]] or, less commonly, [[potassium hydroxide]]). Alkalinity affects [[gluten]] production in kneaded doughs, and also improves browning by reducing the temperature at which the [[Maillard reaction]] occurs. To take advantage of the former effect, sodium carbonate is therefore one of the components of {{nihongo3||かん水|kansui}}, a solution of alkaline salts used to give [[Japanese cuisine|Japanese]] [[ramen]] noodles their characteristic flavour and chewy texture; a similar solution is used in [[Chinese cuisine]] to make [[lamian]], for similar reasons. [[Cantonese cuisine|Cantonese]] bakers similarly use sodium carbonate as a substitute for lye-water to give [[moon cake]]s their characteristic texture and improve browning. In [[German cuisine]] (and Central European cuisine more broadly), breads such as [[pretzel]]s and [[lye roll]]s traditionally treated with lye to improve browning can be treated instead with sodium carbonate; sodium carbonate does not produce quite as strong a browning as lye, but is much safer and easier to work with.<ref name="McGee">{{cite news |last1=McGee |first1=Harold |author-link=Harold McGee |title=For Old-Fashioned Flavor, Bake the Baking Soda |url=https://www.nytimes.com/2010/09/15/dining/15curious.html |access-date=25 April 2019 |work=[[The New York Times]] |date=24 September 2010}}</ref>

Sodium carbonate is used in the production of [[sherbet (powder)|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.

Sodium carbonate also finds use in the [[food industry]] as a [[food additive]] ([[E number|European Food Safety Authority number]] E500) as an acidity regulator, [[anticaking agent]], [[leavening agent|raising agent]], and stabilizer. It is also used in the production of {{lang|no|[[snus]]}} to stabilize the pH of the final product.

===Other applications===
Sodium carbonate is also used as a relatively strong [[Base (chemistry)|base]] in various fields. As a common alkali, it is preferred in many chemical processes because it is cheaper than [[sodium hydroxide]] 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 [[Developer (photography)|film developing]] agents. It is also a common additive in [[swimming pool]]s 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 as [[mordant]] to ensure proper chemical bonding of the dye with cellulose (plant) fiber. It is also used in the [[froth flotation process]] to maintain a favourable [[pH]] as a float conditioner besides [[CaO]] and other mildly basic compounds.
=== Porphyrins,Fractionated. ===
In Clinical Chemistry, Sodium Carbonate is the preferred preservative used in the collection of urine for the quantitative analysis of Porphyrins,Fractionated. Random or 24-hour urine sample types are collected with 5 grams of Sodium Carbonate, protected from light and refrigerated.
==== Method: ====
# Measure out 5 grams Sodium Carbonate, preferrably using a scale.
# Put the measured Sodium Carbonate from step 1 above into a sterile urine collection container protected from light for a Random urine sample or a 24-hour urine sample.
# Collect urine sample into the container as directed by a Physician or laboratory personnel.
# Store the collected urine refrigerated and protected from light until testing is complete.

===Precursor to other compounds===
Sodium {{em|bicarbonate}} (NaHCO<sub>3</sub>) or baking soda, also a component in fire extinguishers, is often generated from sodium carbonate. Although NaHCO<sub>3</sub> is itself an intermediate product of the Solvay process, the heating needed to remove the ammonia that contaminates it decomposes some NaHCO<sub>3</sub>, making it more economical to react finished Na<sub>2</sub>CO<sub>3</sub> with CO<sub>2</sub>:
{{block indent|Na<sub>2</sub>CO<sub>3</sub> + CO<sub>2</sub> + H<sub>2</sub>O → 2NaHCO<sub>3</sub>}}

In a related reaction, sodium carbonate is used to make [[sodium bisulfite]] (NaHSO<sub>3</sub>), which is used for the "sulfite" method of separating [[lignin]] from cellulose. This reaction is exploited for removing [[sulfur dioxide]] from flue gases in power stations:
{{block indent|Na<sub>2</sub>CO<sub>3</sub> + SO<sub>2</sub> + H<sub>2</sub>O → NaHCO<sub>3</sub> + NaHSO<sub>3</sub>}}
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.

It is also used to form carbonates of other metals by ion exchange, often with the other metals' sulphates.

===Miscellaneous===
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.

Sodium carbonate is also used in the processing and tanning of animal hides.<ref>{{cite web|title=Home Tanning Hides and Furs|url= https://shareok.org/bitstream/handle/11244/331379/oksa_ANSI-3998_2007-06.pdf?sequence=1&isAllowed=y|access-date=16 April 2024}}</ref>

==Physical properties==
The integral [[enthalpy of solution]] of sodium carbonate is −26.7&nbsp;kJ/mol .<ref>{{cite web|url=https://chem.libretexts.org/Courses/Howard_University/General_Chemistry%3A_An_Atoms_First_Approach/Unit_4%3A__Thermochemistry/09%3A_Thermochemistry/Chapter_9.05%3A_Enthalpies_of_Solution |title=Joshua Halpern |date=22 November 2014 }}</ref> The [[Mohs scale of mineral hardness|Mohs hardness]] of sodium carbonate monohydrate is 1.3.<ref name=pphoic />

==Occurrence as natural mineral==

[[File:Na2CO3.H2O-bas.png|160px|thumbnail|left|Structure of monohydrate at 346&nbsp;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 [[mummy|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 hydrogendi carbonate dihydrate, are also known from ultra-alkaline [[pegmatite|pegmatitic rocks]], that occur for example in the [[Kola Peninsula]] in Russia.

Extra terrestrially, 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.<ref>{{cite journal |title=Bright carbonate deposits as evidence of aqueous alteration on (1) Ceres |journal=Nature |date= 29 June 2016 |last=De Sanctis |first=M. C. |display-authors=etal |volume=536 |issue= 7614|doi=10.1038/nature18290 |pages=54–57 |pmid=27362221|bibcode=2016Natur.536...54D |s2cid=4465999 }}</ref> While there are [[carbonates on Mars]], and these are expected to include sodium carbonate,<ref name="Kargel2004">{{cite book|author=Jeffrey S. Kargel|title=Mars - A Warmer, Wetter Planet|url=https://books.google.com/books?id=0QY0U6qJKFUC&pg=PA399|date=23 July 2004|publisher=Springer Science & Business Media|isbn=978-1-85233-568-7|pages=399–}}</ref> 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]].<ref>Grotzinger, J. and R. Milliken (eds.) 2012. Sedimentary Geology of Mars. SEPM</ref>

==Production==
The initial large-scale chemical procedure was established in England in 1823 to manufacture soda ash.<ref name="Himmel" />

===Mining===
[[Trona]], also known as [[sodium sesquicarbonate|trisodium hydrogendicarbonate dihydrate]] (Na<sub>3</sub>HCO<sub>3</sub>CO<sub>3</sub>·2H<sub>2</sub>O), is mined in several areas of the US and provides nearly all the US 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;<ref>{{cite news |date=2021-08-09 |title=Ciner Weighs Sale of Stake in $5 Billion Soda Ash Unit |language=en |work=Bloomberg.com |url=https://www.bloomberg.com/news/articles/2021-08-09/ciner-said-to-weigh-sale-of-stake-in-5-billion-soda-ash-unit |access-date=2023-12-04}}</ref> two million tons of soda ash have been extracted from the reserves near Ankara.

===Barilla and kelp===
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 [[glasswort]]s or [[saltwort]]s) or the seaweed (typically ''[[Fucus]]'' species) were harvested, dried, and burned. The ashes were then "[[Leaching (chemistry)|lixivated]]" (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 derives from the Arabic word ''soda'', in turn applied to ''[[Salsola soda]]'', one of the many species of seashore plants harvested for production. "Barilla" is a commercial term applied to an impure form of [[Pearlash|potash]] obtained from coastal plants or [[kelp]].<ref>{{cite book |last1=Hooper |first1=Robert |author-link1=Robert Hooper (physician) |title=Lexicon Medicum |date=1802 |publisher=Longman |location=London |pages=1198–9 |edition=1848|oclc= 27671024}}</ref>

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 [[alkali]] "[[potash]]", 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.<ref name="Clow52">
Clow, Archibald and Clow, Nan L. (June 1952). ''Chemical Revolution''. Ayer. pp. 65–90. {{ISBN|0-8369-1909-2}}.</ref>

===Leblanc process===
{{Main|Leblanc process}}
In 1792, the French chemist [[Nicolas Leblanc]] patented a process for producing sodium carbonate from salt, [[sulfuric 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]]:
{{block indent|2NaCl + H<sub>2</sub>SO<sub>4</sub> → Na<sub>2</sub>SO<sub>4</sub> + 2HCl}}

The salt cake and crushed [[limestone]] ([[calcium carbonate]]) was reduced by heating with [[coal]].<ref name=Ullmann>{{cite encyclopedia|author=Christian Thieme|encyclopedia=Ullmann's Encyclopedia of Industrial Chemistry|publisher=Wiley-VCH|location=Weinheim|year=2000|doi=10.1002/14356007.a24_299|isbn = 978-3527306732|chapter = Sodium Carbonates}}</ref> This conversion entails two parts. First is the [[carbothermic reaction]] whereby the coal, a source of [[carbon]], [[Redox|reduces]] the [[sulfate]] to [[sulfide]]:
{{block indent|Na<sub>2</sub>SO<sub>4</sub> + 2C → Na<sub>2</sub>S + 2CO<sub>2</sub>}}

The second stage is the reaction to produce sodium carbonate and [[calcium sulfide]]:
{{block indent|Na<sub>2</sub>S + CaCO<sub>3</sub> → Na<sub>2</sub>CO<sub>3</sub> + 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 [[Leaching (chemistry)|lixiviating]].

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.<ref name="Clow52"/><ref name="Kiefer">{{cite journal |last1=Kiefer |first1=David M. |date=January 2002 |url=http://pubs.acs.org/subscribe/journals/tcaw/11/i01/html/01chemchron.html |title=It was all about alkali |journal=Today's Chemist at Work |volume=11 |issue=1 |pages=45–6}}</ref>

===Solvay process===
{{Main|Solvay process}}

In 1861, the [[Belgium|Belgian]] industrial chemist [[Ernest Solvay]] developed a method for making sodium carbonate by first reacting [[sodium chloride]], [[ammonia]], water, and carbon dioxide to generate [[sodium bicarbonate]] and [[ammonium chloride]]:<ref name=Ullmann/>

{{block indent|NaCl + NH<sub>3</sub> + CO<sub>2</sub> + H<sub>2</sub>O → NaHCO<sub>3</sub> + NH<sub>4</sub>Cl}}

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

{{block indent|2NaHCO<sub>3</sub> → Na<sub>2</sub>CO<sub>3</sub> + H<sub>2</sub>O + CO<sub>2</sub>}}

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

{{block indent|2NH<sub>4</sub>Cl + CaO → 2NH<sub>3</sub> + CaCl<sub>2</sub> + H<sub>2</sub>O}}

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.<ref name=Ullmann/>

The second step of the Solvay process, heating sodium bicarbonate, is used on a small scale by home cooks and in restaurants to make sodium carbonate for culinary purposes (including [[pretzels]] and [[alkali noodles]]). The method is appealing to such users because sodium bicarbonate is widely sold as baking soda, and the temperatures needed ({{convert|250|F|C}} to {{convert|300|F|C}}) to convert baking soda to sodium carbonate are readily achieved in conventional kitchen [[oven]]s.<ref name="McGee"/>

===Hou's process===
{{unsourced section|date=October 2025}}

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

{{block indent|{{chem2|[[methane|CH4]] + 2 [[water|H2O]] → [[carbon dioxide|CO2]] + 4 [[hydrogen|H2]]}}}}
{{block indent|{{chem2|3 [[hydrogen|H2]] + [[nitrogen|N2]] → 2 [[ammonia|NH3]]}}}}
{{block indent|{{chem2|[[ammonia|NH3]] + [[carbon dioxide|CO2]] + [[water|H2O]] → [[ammonium bicarbonate|NH4HCO3]]}}}}
{{block indent|{{chem2|[[ammonium bicarbonate|NH4HCO3]] + [[sodium chloride|NaCl]] → [[ammonium chloride|NH4Cl]] + [[sodium bicarbonate|NaHCO3]]}}}}

The sodium bicarbonate was collected as a precipitate due to its low solubility and then heated up to approximately {{Convert|80|C|}} or {{Convert|95|C|}} 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&nbsp;°C to this solution. The solution temperature is then lowered to below 10&nbsp;°C. Solubility of ammonium chloride is higher than that of sodium chloride at 30&nbsp;°C and lower at 10&nbsp;°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'' ({{zh|c=联合制碱法|labels=no}}), 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 any ammonia generated gets used by the reaction. The by-product ammonium chloride can be sold as a fertilizer.

==See also==
* [[Residual sodium carbonate index]]

==References==
{{reflist}}

==Further reading==
* {{cite book | last1 = Eggeman | first1 = T. | chapter = Sodium Carbonate | doi = 10.1002/0471238961.1915040918012108.a01.pub3 | title = Kirk-Othmer Encyclopedia of Chemical Technology | year = 2011 | pages = 1–11 | isbn = 978-0471238966 }}
* {{cite book | last1 = Thieme | first1 = C. | chapter = Sodium Carbonates | doi = 10.1002/14356007.a24_299 | title = Ullmann's Encyclopedia of Industrial Chemistry | year = 2000 | isbn = 978-3527306732 }}

==External links==
{{Commons category|Sodium carbonate}}
* [http://www.ansac.com American Natural Soda Ash Company]
* [http://www.inchem.org/documents/icsc/icsc/eics1135.htm International Chemical Safety Card 1135]
* [https://archive.today/20061015212751/http://www.fmcchemicals.com/Products/SodaAsh/tabid/1471/Default.aspx FMC Wyoming Corporation]
* [http://www.pburch.net/dyeing/FAQ/sodaash.shtml Use of sodium carbonate in dyeing]
* [https://web.archive.org/web/20181116131601/http://www.inclusive-science-engineering.com/sodium-carbonate-manufacturing-synthetic-processes-chlor-alkali-industry/ Sodium carbonate manufacturing] by synthetic processes

{{Carbonates}}
{{Sodium compounds}}
{{Authority control}}

[[Category:Carbonates]]
[[Category:E-number additives]]
[[Category:Household chemicals]]
[[Category:Photographic chemicals]]
[[Category:Sodium compounds]]
[[Category:Types of ash|soda ash]]