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Sodium carbonate

Soda ash

Washing soda

Soda crystals

Identifiers

CAS number 497-19-8 ,

5968-11-6 (monohydrate)

6132-02-1 (decahydrate)

PubChem 10340

ChemSpider 9916

UNII 45P3261C7T

EC number 207-838-8

ChEBI CHEBI:29377

ChEMBL CHEMBL186314

RTECS

number

VZ4050000

Jmol-3D

images

Image 1 (http://chemapps.stolaf.edu/jmol/jmol.php?

model=%5BNa%2B%5D.%5BNa%2B%5D.%5BO-%5DC%28%5BO-

%5D%29%3DO)

Sodium carbonateFrom Wikipedia, the free encyclopedia

Sodium

carbonate (also

known as

washing soda

orsoda ash),

Na2CO3 is a

sodium salt of

carbonic acid. It

most commonly

occurs as a

crystalline

heptahydrate,

which readily

effloresces toform a white

powder, the

monohydrate.

Sodium

carbonate is

domestically well

known for its

everyday use as

a water softener.

It can beextracted from

the ashes of

many plants. It is

synthetically

produced in

large quantities

from salt

(sodium

chloride) and

limestone by amethod known

as the Solvay

process.

Other names

SMILES

InChI
http://chemapps.stolaf.edu/jmol/jmol.php?model=%5BNa%2B%5D.%5BNa%2B%5D.%5BO-%5DC%28%5BO-%5D%29%3DOhttp://en.wikipedia.org/wiki/File:Sodium-carbonate-xtal-3D-SF-C.pnghttp://en.wikipedia.org/wiki/File:Uhli%C4%8Ditan_sodn%C3%BD.JPGhttp://en.wikipedia.org/wiki/File:Uhli%C4%8Ditan_sodn%C3%BD.JPGhttp://en.wikipedia.org/wiki/File:Uhli%C4%8Ditan_sodn%C3%BD.JPGhttp://en.wikipedia.org/wiki/File:Sodium_carbonate.pnghttp://en.wikipedia.org/wiki/File:Sodium_carbonate.pnghttp://en.wikipedia.org/wiki/Solvay_processhttp://en.wikipedia.org/wiki/Limestonehttp://en.wikipedia.org/wiki/Sodium_chloridehttp://en.wikipedia.org/wiki/Sodium_chloridehttp://en.wikipedia.org/wiki/Water_softenerhttp://en.wikipedia.org/wiki/Efflorescencehttp://en.wikipedia.org/wiki/Heptahydratehttp://en.wikipedia.org/wiki/Carbonic_acidhttp://en.wikipedia.org/wiki/Sodiumhttp://en.wikipedia.org/wiki/File:Sodium-carbonate-xtal-3D-SF-C.pnghttp://en.wikipedia.org/wiki/International_Chemical_Identifierhttp://en.wikipedia.org/wiki/Simplified_molecular-input_line-entry_systemhttp://en.wikipedia.org/wiki/Solvay_processhttp://en.wikipedia.org/wiki/Limestonehttp://en.wikipedia.org/wiki/Sodium_chloridehttp://en.wikipedia.org/wiki/Water_softenerhttp://en.wikipedia.org/wiki/Efflorescencehttp://en.wikipedia.org/wiki/Heptahydratehttp://en.wikipedia.org/wiki/Crystalhttp://en.wikipedia.org/wiki/Carbonic_acidhttp://en.wikipedia.org/wiki/Salt_(chemistry)http://en.wikipedia.org/wiki/Sodiumhttp://chemapps.stolaf.edu/jmol/jmol.php?model=%5BNa%2B%5D.%5BNa%2B%5D.%5BO-%5DC%28%5BO-%5D%29%3DOhttp://en.wikipedia.org/wiki/Jmolhttp://en.wikipedia.org/wiki/RTECShttps://www.ebi.ac.uk/chembldb/index.php/compound/inspect/CHEMBL186314http://en.wikipedia.org/wiki/ChEMBLhttps://www.ebi.ac.uk/chebi/searchId.do?chebiId=29377http://en.wikipedia.org/wiki/ChEBIhttp://esis.jrc.ec.europa.eu/lib/einecs_IS_reponse.php?genre=ECNO&entree=207-838-8http://en.wikipedia.org/wiki/EC_number_(chemistry)http://fdasis.nlm.nih.gov/srs/srsdirect.jsp?regno=45P3261C7Thttp://en.wikipedia.org/wiki/Unique_Ingredient_Identifierhttp://www.chemspider.com/9916http://en.wikipedia.org/wiki/ChemSpiderhttp://pubchem.ncbi.nlm.nih.gov/summary/summary.cgi?cid=10340http://en.wikipedia.org/wiki/PubChemhttp://www.commonchemistry.org/ChemicalDetail.aspx?ref=497-19-8http://en.wikipedia.org/wiki/CAS_registry_numberhttp://en.wikipedia.org/wiki/File:Sodium-carbonate-xtal-3D-SF-C.pnghttp://en.wikipedia.org/wiki/File:Uhli%C4%8Ditan_sodn%C3%BD.JPGhttp://en.wikipedia.org/wiki/File:Sodium_carbonate.png


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Properties

Molecular

formula

Na2CO3

Molar mass 105.9885 g/mol (anhydrous)

124.00 g/mol (monohydrate)

286.14 g/mol (decahydrate)

Appearance White solid, hygroscopicOdor Odorless

Density 2.54 g/cm3 (anhydrous)

2.25 g/cm3 (monohydrate)

1.51 g/cm3 (heptahydrate)

1.46 g/cm3 (decahydrate)

Melting

point851 C (anhydrous)[1]

100 C (decomp, monohydrate)

33.5 C (decomp, decahydrate)32 C (heptahydrate)

Boiling point1633 C (anhydrous)

Solubility in

water

71 g/L (0 C)

215 g/L (20 C)

455 g/L (100 C)[1]

Solubility insoluble in ethanol, acetone

Basicity

(pKb)

4.67

Refractive

index (nD)

1.485 (anhydrous)

1.420 (monohydrate)

1.405 (decahydrate)

Structure

Coordination

geometry

trigonal planar

Thermochemistry

Std enthalpy

of

formation

fHo298

1131 kJmol1[2]

Standard

molar

entropy

So298

136 Jmol1K1[2]

Hazards

MSDS MSDS

(http://www.chem.tamu.edu/class/majors/msdsfiles/msdssodiumcarb.htm#Material)
http://www.chem.tamu.edu/class/majors/msdsfiles/msdssodiumcarb.htm#Materialhttp://en.wikipedia.org/wiki/MSDShttp://en.wikipedia.org/wiki/Sodium_carbonate#cite_note-b1-2http://en.wikipedia.org/wiki/Standard_molar_entropyhttp://en.wikipedia.org/wiki/Sodium_carbonate#cite_note-b1-2http://en.wikipedia.org/wiki/Standard_enthalpy_change_of_formationhttp://en.wikipedia.org/wiki/Coordination_geometryhttp://en.wikipedia.org/wiki/Refractive_indexhttp://en.wikipedia.org/wiki/Acid_dissociation_constanthttp://en.wikipedia.org/wiki/Acetonehttp://en.wikipedia.org/wiki/Ethanolhttp://en.wikipedia.org/wiki/Solubilityhttp://en.wikipedia.org/wiki/Sodium_carbonate#cite_note-UNEP-1http://en.wikipedia.org/wiki/Waterhttp://en.wikipedia.org/wiki/Solubilityhttp://en.wikipedia.org/wiki/Boiling_pointhttp://en.wikipedia.org/wiki/Sodium_carbonate#cite_note-UNEP-1http://en.wikipedia.org/wiki/Melting_pointhttp://en.wikipedia.org/wiki/Densityhttp://en.wikipedia.org/wiki/Odorhttp://en.wikipedia.org/wiki/Hygroscopichttp://en.wikipedia.org/wiki/Molar_masshttp://en.wikipedia.org/wiki/Chemical_formula


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The manufacture of glass is one of the most important uses of sodium carbonate. Sodium carbonate acts 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 pre-melt mixture to make the

glass produced insoluble. This type of glass is known as soda lime glass: "soda" for the sodium carbonate and

"lime" for the calcium carbonate. Soda lime glass has been the most common form of glass for centuries.

Sodium carbonate is also used as a relatively strong base in various settings. For example, sodium carbonate is

used as a pH regulator to maintain stable alkaline conditions necessary for the action of the majority of

photographic film developing agents.

It is a common additive in municipal pools used to neutralize the corrosive effects of chlorine and raise pH.

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.

In taxidermy, sodium carbonate added to boiling water will remove flesh from the skull or bones of trophies to

create the "European skull mount" or for educational display in biological and historical studies.

In chemistry, it is often used as an electrolyte. This is because electrolytes are usually salt-based, and sodium

carbonate acts as a very good conductor in the process of electrolysis. In addition, unlike chloride ions, which

form chlorine gas, carbonate ions are not corrosive to the anodes. It is also used as a primary standard for acid-

base titrations because it is solid and air-stable, making it easy to weigh accurately.

Domestic use

In domestic use, it is used as a water softener in laundering. It competes with the magnesium and calcium ions in

hard water and prevents them from bonding with the detergent being used. Sodium carbonate can be used to

remove grease, oil and wine stains. It is sold as washing soda, soda crystals, or sal soda. Sodium carbonate is

also used as a descaling agent in boilers such as those found in coffee pots and espresso machine.

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 carbonate test

The sodium carbonate test (not to be confused with sodium carbonate extract test) is used to distinguishbetween some common metal ions, which are precipitated as their respective carbonates. The test can

distinguish between Cu, Fe and Ca/Zn/Pb. Sodium carbonate solution is added to the salt of the metal. A blue

precipitate indicates Cu2+ ion. A dirty green precipitate indicates Fe2+ ion. A yellow-brown precipitate indicates

Fe3+ ion. A white precipitate indicates Ca2+, Zn2+ or Pb2+ ion. The compounds formed are, respectively,

copper(II) carbonate, iron(II) carbonate, iron(III) oxide, calcium carbonate, zinc carbonate and lead(II)

carbonate. This test is used to precipitate the ion present as almost all carbonates are insoluble. While this test is

useful for telling these cations apart, it fails if other ions are present, because most metal carbonates are insoluble

and will precipitate. In addition, calcium, zinc and lead ions all produce white precipitates with carbonate,

making it difficult to distinguish between them. Instead of sodium carbonate, sodium hydroxide may be added,

this gives nearly the same colours, except that lead and zinc hydroxides are soluble in excess alkali, and canhence be distinguished from calcium. For the complete sequence of tests used for qualitative cation analysis, see

qualitative inorganic analysis.
http://en.wikipedia.org/wiki/Qualitative_inorganic_analysishttp://en.wikipedia.org/wiki/Sodium_hydroxidehttp://en.wikipedia.org/wiki/Lead(II)_carbonatehttp://en.wikipedia.org/wiki/Zinc_carbonatehttp://en.wikipedia.org/wiki/Calcium_carbonatehttp://en.wikipedia.org/wiki/Iron(III)_oxidehttp://en.wikipedia.org/wiki/Iron(II)_carbonatehttp://en.wikipedia.org/wiki/Copper(II)_carbonatehttp://en.wikipedia.org/wiki/Dyeinghttp://en.wikipedia.org/wiki/Espresso_machinehttp://en.wikipedia.org/wiki/Descalinghttp://en.wikipedia.org/wiki/Titrationhttp://en.wikipedia.org/wiki/Electrolytehttp://en.wikipedia.org/wiki/Taxidermyhttp://en.wikipedia.org/wiki/German_cuisinehttp://en.wikipedia.org/wiki/Lyehttp://en.wikipedia.org/wiki/Sodium_hydroxidehttp://en.wikipedia.org/wiki/Chlorinehttp://en.wikipedia.org/wiki/Developer_(photography)http://en.wikipedia.org/wiki/Base_(chemistry)http://en.wikipedia.org/wiki/Soda_lime_glasshttp://en.wikipedia.org/wiki/Calcium_carbonatehttp://en.wikipedia.org/wiki/Silicahttp://en.wikipedia.org/wiki/Flux_(metallurgy)http://en.wikipedia.org/wiki/Glass


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Other applications

Sodium carbonate is a food additive (E500) used as an acidity regulator, anti-caking agent, raising agent, and

stabilizer. It is one of the components ofkansui, a solution of alkaline salts used to give ramen noodles their

characteristic flavor and texture. It is also used in the production ofsnus (Swedish-style snuff) to stabilize the pH

of the final product. In Sweden,snus is regulated as a food product because it is put into the mouth, requires

pasteurization, and contains only ingredients that are approved as food additives.

Sodium carbonate is also 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.

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, in a solution with common salt, may be used for cleaning silver. In a non-reactive container

(glass, plastic or ceramic) aluminium foil and the silver object are immersed in the hot salt solution. The elevated

pH dissolves the aluminium oxide layer on the foil and enables an electrolytic cell to be established. Hydrogen

ions produced by this reaction reduce the sulphide ions on the silver restoring silver metal. The sulphide can be

released as small amounts of hydrogen sulphide. Rinsing and gently polishing the silver restores a highly polished

condition. [3]

Occurrence

Sodium carbonate crystallizes from water to form three different hydrates:

1. sodium carbonate decahydrate (natron)

2. sodium carbonate heptahydrate (not known in mineral form)

3. sodium carbonate monohydrate (mineral thermonatrite).

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.

Production
http://en.wikipedia.org/wiki/Kola_Peninsulahttp://en.wikipedia.org/wiki/Pegmatitehttp://en.wikipedia.org/wiki/Tronahttp://en.wikipedia.org/wiki/Ol_Doinyo_Lengaihttp://en.wikipedia.org/wiki/Mummyhttp://en.wikipedia.org/wiki/Natronhttp://en.wikipedia.org/wiki/Natronhttp://en.wikipedia.org/wiki/Sodium_carbonate#cite_note-3http://en.wikipedia.org/wiki/Alginatehttp://en.wikipedia.org/wiki/Moon_cakehttp://en.wikipedia.org/wiki/Citric_acidhttp://en.wikipedia.org/wiki/Sherbet_(powder)http://en.wikipedia.org/wiki/Snuffhttp://en.wikipedia.org/wiki/Snushttp://en.wikipedia.org/wiki/Ramen


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Mining

Trona, trisodium hydrogendicarbonate dihydrate (Na3HCO3CO32H2O), 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 ofdredging is no greater than the replenishment rate, the source is fully sustainable.

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 glassworts or saltworts) or the seaweed (typicallyFucus 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 23 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.[4]

Leblanc process

Main article: Leblanc process

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 hydrogen chloride gas, according to the chemical equation

2 NaCl + H2SO4 Na2SO4 + 2 HCl

Next, the sodium sulfate was blended with crushed limestone (calcium carbonate) and coal, and the mixture was

burnt, producing calcium sulfide.

Na2SO4 + CaCO3 + 2 C Na2CO3 + 2 CO2 + CaS

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.[4][5]

Solvay process

Main article: Solvay process
http://en.wikipedia.org/wiki/Solvay_processhttp://en.wikipedia.org/wiki/Sodium_carbonate#cite_note-Kiefer-5http://en.wikipedia.org/wiki/Sodium_carbonate#cite_note-Clow52-4http://en.wikipedia.org/wiki/Calcium_sulfidehttp://en.wikipedia.org/wiki/Leblanc_processhttp://en.wikipedia.org/wiki/Solvent_extractionhttp://en.wikipedia.org/wiki/Calcium_sulfidehttp://en.wikipedia.org/wiki/Carbon_dioxidehttp://en.wikipedia.org/wiki/Carbonhttp://en.wikipedia.org/wiki/Calcium_carbonatehttp://en.wikipedia.org/wiki/Sodium_sulfatehttp://en.wikipedia.org/wiki/Calcium_sulfidehttp://en.wikipedia.org/wiki/Calcium_carbonatehttp://en.wikipedia.org/wiki/Limestonehttp://en.wikipedia.org/wiki/Hydrochloric_acidhttp://en.wikipedia.org/wiki/Sodium_sulfatehttp://en.wikipedia.org/wiki/Sulfuric_acidhttp://en.wikipedia.org/wiki/Sodium_chloridehttp://en.wikipedia.org/wiki/Chemical_equationhttp://en.wikipedia.org/wiki/Hydrogen_chloridehttp://en.wikipedia.org/wiki/Sodium_sulfatehttp://en.wikipedia.org/wiki/Sodium_chloridehttp://en.wikipedia.org/wiki/Limestonehttp://en.wikipedia.org/wiki/Sulfuric_acidhttp://en.wikipedia.org/wiki/Nicolas_Leblanchttp://en.wikipedia.org/wiki/Leblanc_processhttp://en.wikipedia.org/wiki/Sodium_carbonate#cite_note-Clow52-4http://en.wikipedia.org/wiki/Potashhttp://en.wikipedia.org/wiki/Alkalihttp://en.wikipedia.org/wiki/Saltworthttp://en.wikipedia.org/wiki/Barillahttp://en.wikipedia.org/wiki/Kelphttp://en.wikipedia.org/wiki/Salsola_sodahttp://en.wikipedia.org/wiki/Fucushttp://en.wikipedia.org/wiki/Saltworthttp://en.wikipedia.org/wiki/Glassworthttp://en.wikipedia.org/wiki/Halophytehttp://en.wikipedia.org/wiki/Lake_Magadihttp://en.wikipedia.org/wiki/Green_River,_Wyominghttp://en.wikipedia.org/wiki/Sodium_sesquicarbonatehttp://en.wikipedia.org/wiki/Trona


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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:

CaCO3 CaO + CO2

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:

NaCl + NH3 + CO2 + H2O NaHCO3 + NH4Cl

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

dioxide:

2 NaHCO3 Na2CO3 + H2O + CO2

Meanwhile, the ammonia was regenerated from the ammonium chloride byproduct by treating it with the lime

(calcium hydroxide) left over from carbon dioxide generation:

CaO + H2O Ca(OH)2Ca(OH)2 + 2 NH4Cl CaCl2 + 2 NH3 + 2 H2O

Because the Solvay process recycles its ammonia, it consumes only brine and limestone, and has 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.

Hou's process

Developed by Chinese chemist Hou Debang in 1930s. The earlier steam reforming byproduct carbon dioxide

was pumped through a saturated solution of sodium chloride and ammonia to produce sodium bicarbonate via

the following reactions:

NH3 + CO2 + H2O NH4HCO3NH4HCO3 + NaCl NH4Cl + NaHCO3

The sodium bicarbonate was collected as a precipitate due to its low solubility and then heated to yield pure

sodium carbonate similar to last step of the Solvay process. More sodium chloride is added to the remainingsolution of ammonium and sodium chlorides; also more ammonia is pumped at 30-40C to this solution. The

solution temperature is then lowered to below 10C. Solubility of ammonium chloride is higher than that of

sodium chloride at 30C and lower at 10C. 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 () 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.

References

1. ^ ab "Sodium Carbonate" (http://www.chem.unep.ch/irptc/sids/oecdsids/Naco.pdf). UNEP Publications.
http://www.chem.unep.ch/irptc/sids/oecdsids/Naco.pdfhttp://en.wikipedia.org/wiki/Sodium_carbonate#cite_ref-UNEP_1-1http://en.wikipedia.org/wiki/Sodium_carbonate#cite_ref-UNEP_1-0http://en.wikipedia.org/wiki/Atom_economyhttp://en.wikipedia.org/wiki/Haber_processhttp://en.wikipedia.org/wiki/Common-ion_effecthttp://en.wikipedia.org/wiki/Sodium_bicarbonatehttp://en.wikipedia.org/wiki/Ammonium_chloridehttp://en.wikipedia.org/wiki/Sodium_chloridehttp://en.wikipedia.org/wiki/Ammonium_bicarbonatehttp://en.wikipedia.org/wiki/Ammonium_bicarbonatehttp://en.wikipedia.org/wiki/Waterhttp://en.wikipedia.org/wiki/Carbon_dioxidehttp://en.wikipedia.org/wiki/Ammoniahttp://en.wikipedia.org/wiki/Brinehttp://en.wikipedia.org/wiki/Steam_reforminghttp://en.wikipedia.org/wiki/Hou_Debanghttp://en.wikipedia.org/wiki/Calcium_chloridehttp://en.wikipedia.org/wiki/Waterhttp://en.wikipedia.org/wiki/Ammoniahttp://en.wikipedia.org/wiki/Calcium_chloridehttp://en.wikipedia.org/wiki/Ammonium_chloridehttp://en.wikipedia.org/wiki/Calcium_hydroxidehttp://en.wikipedia.org/wiki/Calcium_hydroxidehttp://en.wikipedia.org/wiki/Waterhttp://en.wikipedia.org/wiki/Calcium_oxidehttp://en.wikipedia.org/wiki/Calcium_hydroxidehttp://en.wikipedia.org/wiki/Carbon_dioxidehttp://en.wikipedia.org/wiki/Waterhttp://en.wikipedia.org/wiki/Sodium_bicarbonatehttp://en.wikipedia.org/wiki/Ammonium_chloridehttp://en.wikipedia.org/wiki/Sodium_bicarbonatehttp://en.wikipedia.org/wiki/Waterhttp://en.wikipedia.org/wiki/Carbon_dioxidehttp://en.wikipedia.org/wiki/Ammoniahttp://en.wikipedia.org/wiki/Sodium_chloridehttp://en.wikipedia.org/wiki/Carbon_dioxidehttp://en.wikipedia.org/wiki/Calcium_oxidehttp://en.wikipedia.org/wiki/Calcium_carbonatehttp://en.wikipedia.org/wiki/Solvay_processhttp://en.wikipedia.org/wiki/Ammoniahttp://en.wikipedia.org/wiki/Ernest_Solvayhttp://en.wikipedia.org/wiki/Belgium


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en.wikipedia.org/wiki/Sodium carbonate 8/8

2. ^ ab Zumdahl, Steven S. (2009). Chemical Principles 6th Ed. Houghton Mifflin Company. p. A23. ISBN 0-

618-94690-X.

3. ^ Finishing techniques in Metalwork - Philadelphia Museum of Art

(http://www.philamuseum.org/booklets/7_44_85_1.html?page=2)

4. ^ ab Clow, Archibald and Clow, Nan L. (1952). Chemical Revolution, (Ayer Co Pub, June 1952), pp. 6590.

ISBN 0-8369-1909-2.

5. ^ Kiefer, David M. (January 2002). "It was all about alkali"

(http://pubs.acs.org/subscribe/journals/tcaw/11/i01/html/01chemchron.html). Today's Chemist at Work11 (1):

456.

External links

American Natural Soda Ash Company (http://www.ansac.com)

International Chemical Safety Card 1135 (http://www.inchem.org/documents/icsc/icsc/eics1135.htm)

FMC Wyoming Corporation

(http://www.fmcchemicals.com/Products/SodaAsh/tabid/1471/Default.aspx)

Use of sodium carbonate in dyeing (http://www.pburch.net/dyeing/FAQ/sodaash.shtml)

See also

Residual Sodium Carbonate Index

Retrieved from "http://en.wikipedia.org/w/index.php?title=Sodium_carbonate&oldid=560564109"

Categories: Carbonates Household chemicals Photographic chemicals Sodium compounds

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