Oxalic acid: Wikis


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Oxalic acid
CAS number 144-62-7 Yes check.svgY
ATCvet code QP53AG03
Molecular formula C2H2O4 (anhydrous)
C2H2O4·2H2O (dihydrate)
Molar mass 90.03 g/mol (anhydrous)
126.07 g/mol (dihydrate)
Appearance white crystals
Density 1.90 g/cm³ (anhydrous)
1.653 g/cm³ (dihydrate)
Melting point

101-102°C (dihydrate)

Solubility in water 9.5 g/100 mL (15 °C)
14.3 g /100 mL (25 °C?)
120 g/100 mL (100 °C)
Acidity (pKa) pKa1=1.27
MSDS External MSDS
NFPA 704
NFPA 704.svg
Flash point 166 °C
Related compounds
Related compounds oxalyl chloride
disodium oxalate
calcium oxalate
phenyl oxalate ester
 Yes check.svgY (what is this?)  (verify)
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Oxalic acid is the chemical compound with the formula that can be written in a number of equivalent ways, C2O4H2, C2O2(OH)2, and as HOOCCOOH. This colourless solid is a dicarboxylic acid. In terms of acid strength, it is about 3,000 times stronger than acetic acid. Its conjugate base, known as oxalate (C2O42-), is a reducing agent as well as a chelating agent for metal cations. Typically oxalic acid occurs as the dihydrate with the formula C2O4H2·2H2O.



Oxalic acid is mainly manufactured by the oxidation of carbohydrates or glucose using nitric acid or air in the presence of vanadium pentoxide. A variety of precursors can be used including glycolic acid and ethylene glycol.[1] A newer method entails oxidative carbonylation of alcohols to give the diesters of oxalic acid:

4 ROH + 4 CO + O2 → 2 (CO2R)2 + 2 H2O

These diesters are subsequently hydrolyzed to oxalic acid. Approximately 120M kg are produced annually.[2]


Laboratory methods

Although it can be readily purchased, oxalic acid can be prepared in the laboratory by oxidizing sucrose using nitric acid in the presence of a small amount of vanadium pentoxide as a catalyst.[3]

The hydrated solid can be dehydrated with heat or by azeotropic distillation.[4]

Of historical interest, Wöhler prepared oxalic acid by hydrolysis of cyanogen in 1824. This experiment may represent the first synthesis of a natural product.[2]


Anhydrous oxalic acid exists as two polymorphs; in one the hydrogen-bonding results in a chain-like structure whereas the hydrogen bonding pattern in the other form defines a sheet-like structure.[5] Because the anhydrous material is both acidic and hygroscopic (water seeking), it is used in esterifications.


Oxalic acid is a relatively strong acid, despite being a carboxylic acid:

C2O4H2 → C2O4H + H+; pKa = 1.27
C2O4H → C2O42− + H+; pKa = 4.28

Oxalic acid undergoes many of the reactions characteristic of other carboxylic acids. It forms esters such as dimethyl oxalate (m.p. 52.5–53.5 °C).[6] It forms an acid chloride called oxalyl chloride.

Oxalate, the conjugate base of oxalic acid, is an excellent ligand for metal ions, e.g. the drug oxaliplatin.

Oxalic acid and oxalates can be oxidized by permanganate in an autocatalytic reaction.[7]

Occurrence in nature

Oxalic acid and oxalates are abundantly present in many plants. It was first isolated from Wood-sorrel.


Oxalic acid's main applications include cleaning or bleaching. Most oxalic acid is used as a cleaning agent, especially for the removal of rust or removal of iron from minerals specimens. Many household chemical products contain oxalic acid, especially rustproofing treatments. Bar Keepers Friend is an example of a cleaner that is used in households and commercially that has oxalic acid as its active ingredient. About 25% of produced oxalic acid is used as a mordant in dyeing processes. It is used in bleaches, especially for pulpwood.[2]

Extractive metallurgy

Oxalic acid is also an important reagent in lanthanide chemistry. Hydrated lanthanide oxalates form readily in strongly acid solution in a densely crystalline easily filtered form, largely free from contamination by non-lanthanide elements. Lanthanide oxalates figure importantly in commercial processing of lanthanides, and are used to recover lanthanides from solution after separation. Upon ignition, lanthanide oxalates are converted to the oxides, which are the most common form in which the lanthanides are marketed. Upon heating, metal oxalates decompose to give the corresponding oxides.

Miscellaneous uses

Oxalic acid is used in the restoration of old wood. Its reducing properties are utilized in platinotype, the early photographic platinum/palladium printing process.

Vaporized oxalic acid, or a 6% solution of oxalic acid in sugar syrup, is used by some beekeepers as a miticide against the parasitic Varroa mite.


In humans, oxalic acid has an oral LDLo of 600 mg/kg.[8] It has an extremely irritating taste. Prolonged handling of aqueous solutions cause joint pains.


  1. ^ http://www.freepatentsonline.com/3678107.html Process for the production of oxalic acid
  2. ^ a b c Wilhelm Riemenschneider, Minoru Tanifuji "Oxalic Acid" in Ullmann's Encyclopedia of Industrial Chemistry, 2002, Wiley-VCH, Weinheim. doi: 10.1002/14356007.a18_247.
  3. ^ Practical Organic Chemistry by Julius B. Cohen, 1930 ed. preparation #42
  4. ^ Clarke H. T.;. Davis, A. W. (1941), "Oxalic Acid (Anhydrous)", Org. Synth.: 421, http://www.orgsyn.org/orgsyn/orgsyn/prepContent.asp?prep=CV1P0421 ; Coll. Vol. 1 
  5. ^ Wells, A.F. (1984) Structural Inorganic Chemistry, Oxford: Clarendon Press. ISBN 0-19-855370-6.
  6. ^ Bowden, E. (1943), "Methyl Oxalate", Org. Synth.: 414, http://www.orgsyn.org/orgsyn/orgsyn/prepContent.asp?prep=CV2P0414 ; Coll. Vol. 2 
  7. ^ Kovacs KA, Grof P, Burai L, Riedel M (2004). "Revising the Mechanism of the Permanganate/Oxalate Reaction". J. Phys. Chem. A 108: 11026. doi:10.1021/jp047061u. 
  8. ^ Safety Officer in Physical Chemistry (August 13, 2005). "Safety (MSDS) data for oxalic acid dihydrate". Oxford University. http://msds.chem.ox.ac.uk/OX/oxalic_acid_dihydrate.html. Retrieved December 30, 2009. 

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