Oxalic acid

Oxalic acid
Oxalic acid ball-and-stick model	Oxalic acid space-filling model
Names
	IUPAC name ethanedioic acid
Identifiers
CAS Number	144-62-7 ;
3D model (JSmol)	Interactive image;
ECHA InfoCard	100.005.123
CompTox Dashboard (EPA)	DTXSID0025816 ;
	SMILES OC(=O)C(O)=O;
Properties
Chemical 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; pKa2=4.28
Hazards
NFPA 704 (fire diamond)	3 1
Flash point	166 °C
Related compounds
	Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa). verify (what is ?) 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 (C₂O₄^2-), 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.

Preparation

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]

Structure

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.

Reactions

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

C₂O₄H₂ → C₂O₄H⁻ + H⁺; pK_a = 1.27

C₂O₄H⁻ → C₂O₄²⁻ + H⁺; pK_a = 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 (Oxalis). Its presence makes it dangerous to eat unripe carambola or monstera fruits. It is also suggested to consume members of the spinach family in moderation as they are high in oxalates.

Rhubarb leaves contain about 0.5% oxalic acid (which can convert into even more toxic oxalates after cooking), as well as other glycoside toxin; consequently only the stalks are used for food. ^[8] Arisaema triphyllum (Jack-in-the-Pulpit), a north american plant, contains calcium oxalate crystals in all its parts, and was used by early natives as both a poison (for their enemies) and food (after special preparation including drying and cooking).

Applications

Oxalic acid's main applications include cleaning or bleaching, especially for the removal of rust, e.g. Bar Keepers Friend is an example of a household cleaner containing oxalic acid. 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 of 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 convert to the oxides, which are the most common form in which the lanthanides are marketed.

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.

Safety

In humans, oxalic acid has an oral LD_Lo of 600 mg/kg.^[9]

References

^ http://www.freepatentsonline.com/3678107.html Process for the production of oxalic acid
^ ^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.
^ Practical Organic Chemistry by Julius B. Cohen, 1930 ed. preparation #42
^ Clarke H. T.;. Davis, A. W. (1941). "Oxalic Acid (Anhydrous)". Organic Syntheses: 421{{cite journal}}: CS1 maint: multiple names: authors list (link); Collected Volumes, vol. 1.
^ Wells, A.F. (1984) Structural Inorganic Chemistry, Oxford: Clarendon Press. ISBN 0-19-855370-6.
^ Bowden, E. (1943). "Methyl Oxalate". Organic Syntheses: 414; Collected Volumes, vol. 2.
^ 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.{{cite journal}}: CS1 maint: multiple names: authors list (link)
^ US National Library of Medicine and National Institutes of Health. Rhubarb leaves poisoning.
^ Safety Officer in Physical Chemistry (August 13, 2005). "Safety (MSDS) data for oxalic acid dihydrate". Oxford University. Retrieved December 30, 2009.

External links

[1] ttp://www.freepatentsonline.com/3678107.html Process for the production of oxalic acid

[Ullmann-2] Wilhelm Riemenschneider, Minoru Tanifuji "Oxalic Acid" in Ullmann's Encyclopedia of Industrial Chemistry, 2002, Wiley-VCH, Weinheim. doi:10.1002/14356007.a18_247.

[cohen-3] Practical Organic Chemistry by Julius B. Cohen, 1930 ed. preparation #42

[4] Clarke H. T.;. Davis, A. W. (1941). "Oxalic Acid (Anhydrous)". Organic Syntheses: 421{{cite journal}}: CS1 maint: multiple names: authors list (link); Collected Volumes, vol. 1.

[5] Wells, A.F. (1984) Structural Inorganic Chemistry, Oxford: Clarendon Press. ISBN 0-19-855370-6.

[6] Bowden, E. (1943). "Methyl Oxalate". Organic Syntheses: 414; Collected Volumes, 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.{{cite journal}}: CS1 maint: multiple names: authors list (link)

[8] US National Library of Medicine and National Institutes of Health. Rhubarb leaves poisoning.

[9] Safety Officer in Physical Chemistry (August 13, 2005). "Safety (MSDS) data for oxalic acid dihydrate". Oxford University. Retrieved December 30, 2009.

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