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Oxalic acid

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Oxalic acid is an organic acid with the systematic name ethanedioic acid and chemical formula HO−C(=O)−C(=O)−OH, also written as (COOH)2 or (CO2H)2 or H2C2O4. It is the simplest dicarboxylic acid. It is a white crystalline solid that forms a colorless solution in water. Its name comes from the fact that early investigators isolated oxalic acid from flowering plants of the genus Oxalis, commonly known as wood-sorrels. It occurs naturally in many foods. Excessive ingestion of oxalic acid or prolonged skin contact can be dangerous.

Oxalic acid
Structural formula of oxalic acid
Skeletal formula of oxalic acid
Space-filling model of oxalic acid
Oxalic acid dihydrate
Oxalic acid dihydrate
Names
IUPAC name
1,2-ethanedioic acid
Preferred IUPAC name
Oxalic acid[1]
Systematic IUPAC name
Ethanedioic acid[1]
Other names
Wood bleach
(Carboxyl)carboxylic acid
Carboxylformic acid
Dicarboxylic acid
Diformic acid
Identifiers
3D model (JSmol)
3DMet
385686
ChEBI
ChEMBL
ChemSpider
DrugBank
ECHA InfoCard 100.005.123 Edit this at Wikidata
EC Number
  • 205-634-3
2208
KEGG
MeSH Oxalic+acid
RTECS number
  • RO2450000
UNII
UN number 3261
  • InChI=1S/C6H6O6/c3-1(4)2(5)6/h(H,3,4)(H,5,6) checkY
    Key: MUBZPKHOEPUJKR-UHFFFAOYSA-N checkY
  • OC(=O)C(=O)O
Properties
C2H2O4
Molar mass 90.034 g·mol−1 (anhydrous)
126.065 g·mol−1 (dihydrate)
Appearance White crystals
Odor Odorless
Density 1.90 g/cm3 (anhydrous, at 17 °C)[2]
1.653 g/cm3 (dihydrate)
Melting point 189 to 191 °C (372 to 376 °F; 462 to 464 K)
101.5 °C (214.7 °F; 374.6 K) dihydrate
  • In g/L:
  • 46.9 (5 °C)
  • 57.2 (10 °C)
  • 75.5 (15 °C)
  • 95.5 (20 °C)
  • 118 (25 °C)
  • 139 (30 °C)
  • 178 (35 °C)
  • 217 (40 °C)
  • 261 (45 °C)
  • 315 (50 °C)
  • 376 (55 °C)
  • 426 (60 °C)
  • 548 (65 °C)
[3]
Solubility 237 g/L (15 °C) in ethanol
14 g/L (15 °C) in diethyl ether[4]
Vapor pressure <0.001 mmHg (20 °C)[5]
Acidity (pKa) pKa1 = 1.25
pKa2 = 4.14[6]
Conjugate base Hydrogenoxalate
−60.05·10−6 cm3/mol
Thermochemistry[7]
91.0 J/(mol·K)
109.8 J/(mol·K)
−829.9 kJ/mol
Pharmacology
QP53AG03 (WHO)
Hazards
Occupational safety and health (OHS/OSH):
Main hazards
 Corrosive
GHS labelling:
GHS05: CorrosiveGHS07: Exclamation markGHS08: Health hazard
Danger
H302+H312, H318, H402
P264, P270, P273, P280, P301+P312+P330, P302+P352+P312, P305+P351+P338+P310, P362+P364, P501
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasFlammability 1: Must be pre-heated before ignition can occur. Flash point over 93 °C (200 °F). E.g. canola oilInstability 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazard ACID: Acid
3
1
0
ACID
Flash point 166 °C (331 °F; 439 K)
Lethal dose or concentration (LD, LC):
1000 mg/kg (dog, oral)
1400 mg/kg (rat)
7500 mg/kg (rat, oral)[8]
NIOSH (US health exposure limits):
PEL (Permissible)
 TWA 1 mg/m3[5]
REL (Recommended)
 TWA 1 mg/m3 ST 2 mg/m3[5]
IDLH (Immediate danger)
 500 mg/m3[5]
Safety data sheet (SDS) External MSDS
Related compounds
Related compounds
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
☒N verify (what is checkY☒N ?)

Oxalic acid has much greater acid strength than acetic acid. It is a reducing agent[9] and its conjugate bases hydrogen oxalate (HC2O−4) and oxalate (C2O2−4) are chelating agents for metal cations. It is used as a cleaning agent, especially for the removal of rust, because it forms a water-soluble ferric iron complex, the ferrioxalate ion. Oxalic acid typically occurs as the dihydrate with the formula H2C2O4·2H2O.

History

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The preparation of salts of oxalic acid from plants had been known, at least since 1745, when the Dutch botanist and physician Herman Boerhaave isolated a salt from wood sorrel, akin to kraft process.[10] By 1773, François Pierre Savary of Fribourg, Switzerland had isolated oxalic acid from its salt in sorrel.[11]

In 1776, Swedish chemists Carl Wilhelm Scheele and Torbern Olof Bergman[12] produced oxalic acid by reacting sugar with concentrated nitric acid; Scheele called the acid that resulted socker-syra or såcker-syra (sugar acid). By 1784, Scheele had shown that "sugar acid" and oxalic acid from natural sources were identical.[13] The modern name was introduced along with many other acid names by de Morveau, Lavoisier and coauthors in 1787.[14]

In 1824, the German chemist Friedrich Wöhler obtained oxalic acid by reacting cyanogen with ammonia in aqueous solution.[15] This experiment may represent the first synthesis of a natural product.[16]

Production

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Industrial

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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.[17] 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 120,000 tonnes are produced annually.[16]

Historically oxalic acid was obtained exclusively by using caustics, such as sodium or potassium hydroxide, on sawdust, followed by acidification of the oxalate by mineral acids, such as sulfuric acid.[18] Oxalic acid can also be formed by the heating of sodium formate in the presence of an alkaline catalyst.[19]

Laboratory

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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.[20]

The hydrated solid can be dehydrated with heat or by azeotropic distillation.[21]

Structure

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Anhydrous

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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.[22] Because the anhydrous material is both acidic and hydrophilic (water seeking), it is used in esterifications.

Dihydrate

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The dihydrate H
2C
2O
4·2H
2O has space group C52h–P21/n, with lattice parameters a = 611.9 pm, b = 360.7 pm, c = 1205.7 pm, β = 106°19′, Z = 2.[23] The main inter-atomic distances are: C−C 153 pm, C−O1 129 pm, C−O2 119 pm.[24]

Reactions

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Acid–base properties

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Oxalic acid's pKa values vary in the literature from 1.25 to 1.46 and from 3.81 to 4.40.[25][26][27] The 100th ed of the CRC, released in 2019, has values of 1.25 and 3.81.[28] Oxalic acid is relatively strong compared to other carboxylic acids:

H2C2O4 ⇌ HC2O−4 + H+            pKa1 = 1.27
HC2O−4 ⇌ C2O2−4 + H+            pKa2 = 4.27

Oxalic acid undergoes many of the reactions characteristic for other carboxylic acids. It forms esters such as dimethyl oxalate (m.p. 52.5 to 53.5 °C, 126.5 to 128.3 °F).[29] It forms an acid chloride called oxalyl chloride.

Metal-binding properties

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Transition metal oxalate complexes are numerous, e.g. the drug oxaliplatin. Oxalic acid has been shown to reduce manganese dioxide MnO2 in manganese ores to allow the leaching of the metal by sulfuric acid.[30]

Oxalic acid is an important reagent in lanthanide chemistry. Hydrated lanthanide oxalates form readily in very strongly acidic solutions as a densely crystalline, easily filtered form, largely free of contamination by nonlanthanide elements:

2 Ln3+ + 3 H2C2O4 → Ln2(C2O4)3 + 6 H+

Thermal decomposition of these oxalates gives the oxides, which is the most commonly marketed form of these elements.[31]

Other

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Oxalic acid and oxalates can be oxidized by permanganate in an autocatalytic reaction.[32]

Oxalic acid vapor decomposes at 125–175 °C into carbon dioxide CO
2 and formic acid HCOOH. Photolysis with 237–313 nm UV light also produces carbon monoxide CO and water.[33]

Evaporation of a solution of urea and oxalic acid in 2:1 molar ratio yields a solid crystalline compound H2C2O4·2CO(NH2)2, consisting of stacked two-dimensional networks of the neutral molecules held together by hydrogen bonds with the oxygen atoms.[34]

Occurrence

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Biosynthesis

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At least two pathways exist for the enzyme-mediated formation of oxalate. In one pathway, oxaloacetate, a component of the Krebs citric acid cycle, is hydrolyzed to oxalate and acetic acid by the enzyme oxaloacetase:[35]

[O2CC(O)CH2CO2]2− + H2O → C2O2−4 + CH3CO−2 + H+

It also arises from the dehydrogenation of glycolic acid, which is produced by the metabolism of ethylene glycol.

Occurrence in foods and plants

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Stems of Oxalis triangularis contain oxalic acid.

Early investigators isolated oxalic acid from wood-sorrel (Oxalis). Members of the spinach family and the brassicas (cabbage, broccoli, brussels sprouts) are high in oxalates, as are sorrel and umbellifers like parsley.[36] The leaves and stems of all species of the genus Chenopodium and related genera of the family Amaranthaceae, which includes quinoa, contain high levels of oxalic acid.[37] Rhubarb leaves contain about 0.5% oxalic acid, and jack-in-the-pulpit (Arisaema triphyllum) contains calcium oxalate crystals. Similarly, the Virginia creeper, a common decorative vine, produces oxalic acid in its berries as well as oxalate crystals in the sap, in the form of raphides. Bacteria produce oxalates from oxidation of carbohydrates.[16]

Plants of the genus Fenestraria produce optical fibers made from crystalline oxalic acid to transmit light to subterranean photosynthetic sites.[38]

Carambola, also known as starfruit, also contains oxalic acid along with caramboxin. Citrus juice contains small amounts of oxalic acid.

The formation of naturally occurring calcium oxalate patinas on certain limestone and marble statues and monuments has been proposed to be caused by the chemical reaction of the carbonate stone with oxalic acid secreted by lichen or other microorganisms.[39][40]

Production by fungi

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Many soil fungus species secrete oxalic acid, which results in greater solubility of metal cations and increased availability of certain soil nutrients, and can lead to the formation of calcium oxalate crystals.[41][42] Some fungi such as Aspergillus niger have been extensively studied for the industrial production of oxalic acid;[43] however, those processes are not yet economically competitive with production from oil and gas.[44] Cryphonectria parasitica may excrete oxalic acid containing solutions at the advancing edge of its chestnut cambium infection. The lower pH (<2.5) of more concentrated oxalic acid excretions may degrade cambium cell walls and have a toxic effect on chestnut cambium cells. Cambium cells that burst provide nutrients for a blight infection advance. [45] [46]

Biochemistry

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The conjugate base of oxalic acid is the hydrogenoxalate anion, and its conjugate base (oxalate) is a competitive inhibitor of the lactate dehydrogenase (LDH) enzyme.[47] LDH catalyses the conversion of pyruvate to lactic acid (end product of the fermentation (anaerobic) process) oxidising the coenzyme NADH to NAD+ and H+ concurrently. Restoring NAD+ levels is essential to the continuation of anaerobic energy metabolism through glycolysis. As cancer cells preferentially use anaerobic metabolism (see Warburg effect) inhibition of LDH has been shown to inhibit tumor formation and growth,[48] thus is an interesting potential course of cancer treatment.

Oxalic acid plays a key role in the interaction between pathogenic fungi and plants. Small amounts of oxalic acid enhances plant resistance to fungi, but higher amounts cause widespread programmed cell death of the plant and help with fungi infection. Plants normally produce it in small amounts, but some pathogenic fungi such as Sclerotinia sclerotiorum cause a toxic accumulation.[49]

Oxalate, besides being biosynthesised, may also be biodegraded. Oxalobacter formigenes is an important gut bacterium that helps animals (including humans) degrade oxalate.[50]

Applications

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Oxalic acid's main applications include cleaning or bleaching, especially for the removal of rust (iron complexing agent). Its utility in rust removal agents is due to its forming a stable, water-soluble salt with ferric iron, ferrioxalate ion. Oxalic acid is an ingredient in some tooth whitening products. About 25% of produced oxalic acid is used as a mordant in dyeing processes. It is also used in bleaches, especially for pulpwood, cork, straw, cane, feathers, and for rust removal and other cleaning, in baking powder, and as a third reagent in silica analysis instruments.

Niche uses

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Honeybee coated with oxalate crystals

Oxalic acid is used by some beekeepers as a miticide against the parasitic varroa mite.[51]

Dilute solutions (0.05–0.15 M) of oxalic acid can be used to remove iron from clays such as kaolinite to produce light-colored ceramics.[52]

Oxalic acid can be used to clean minerals like many other acids. Two such examples are quartz crystals and pyrite.[53][54][55]

Oxalic acid is sometimes used in the aluminum anodizing process, with or without sulfuric acid.[56] Compared to sulfuric-acid anodizing, the coatings obtained are thinner and exhibit lower surface roughness.

Oxalic acid is also widely used as a wood bleach, most often in its crystalline form to be mixed with water to its proper dilution for use.[citation needed]

Semiconductor industry

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Oxalic acid is also used in electronic and semiconductor industries. In 2006 it was reported being used in electrochemical–mechanical planarization of copper layers in the semiconductor devices fabrication process.[57]

Proposed uses

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Reduction of carbon dioxide to oxalic acid by various methods, such as electrocatalysis using a copper complex,[58] is under study as a proposed chemical intermediate for carbon capture and utilization.[59]

Content in food items

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[60][clarification needed]

Vegetable Content of oxalic acid
(%)a
Amaranth 1.09
Asparagus 0.13
Beans, snap 0.36
Beet leaves 0.61
Beetroot 0.06[61]
Broccoli 0.19
Brussels sprouts 0.02[61]
Cabbage 0.10
Carrot 0.50
Cassava 1.26
Cauliflower 0.15
Celery 0.19
Chicory 0.2
Chives 1.48
Collards 0.45
Coriander 0.01
Corn, sweet 0.01
Cucumber 0.02
Eggplant 0.19
Endive 0.11
Garlic 0.36
Kale 0.02
Lettuce 0.33
Okra 0.05
Onion 0.05
Parsley 1.70
Parsnip 0.04
Pea 0.05
Bell pepper 0.04
Potato 0.05
Purslane 1.31
Radish 0.48
Rhubarb leaves 0.52[62]
Rutabaga 0.03
Spinach 0.97 (ranges from 0.65% to 1.3%
on fresh weight basis)[63]
Squash 0.02
Sweet potato 0.24
Swiss chard, green 0.96 [61]
Tomato 0.05
Turnip 0.21
Turnip greens 0.05
Watercress 0.31

Toxicity

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Oxalic acid has an oral LDLo (lowest published lethal dose) of 600 mg/kg.[64] It has been reported that the lethal oral dose is 15 to 30 grams.[65] The toxicity of oxalic acid is due to kidney failure caused by precipitation of solid calcium oxalate.[66]

Oxalate is known to cause mitochondrial dysfunction.[67]

Ingestion of ethylene glycol results in oxalic acid as a metabolite which can also cause acute kidney failure.

Kidney stones

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Most kidney stones, 76%, are composed of calcium oxalate.[68]

Notes

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^a Unless otherwise cited, all measurements are based on raw vegetable weights with original moisture content.

References

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  1. ^ a b "Front Matter". Nomenclature of Organic Chemistry : IUPAC Recommendations and Preferred Names 2013 (Blue Book). Cambridge: The Royal Society of Chemistry. 2014. pp. P001–P004. doi:10.1039/9781849733069-FP001. ISBN 978-0-85404-182-4.
  2. ^ Record in the GESTIS Substance Database of the Institute for Occupational Safety and Health
  3. ^ Apelblat, Alexander; Manzurola, Emanuel (1987). "Solubility of oxalic, malonic, succinic, adipic, maleic, malic, citric, and tartaric acids in water from 278.15 to 338.15 K". The Journal of Chemical Thermodynamics. 19 (3): 317–320. doi:10.1016/0021-9614(87)90139-X.
  4. ^ Radiant Agro Chem. "Oxalic Acid MSDS". Archived from the original on 2011-07-15. Retrieved 2012-02-02.
  5. ^ a b c d NIOSH Pocket Guide to Chemical Hazards. "#0474". National Institute for Occupational Safety and Health (NIOSH).
  6. ^ Bjerrum, Jannik; Sillén, Lars Gunnar; Schwarzenbach, Gerold Karl; Anderegg, Giorgio (1958). Stability constants of metal-ion complexes, with solubility products of inorganic substances. London: Chemical Society.
  7. ^ CRC handbook of chemistry and physics : a ready-reference book of chemical and physical data. William M. Haynes, David R. Lide, Thomas J. Bruno (2016-2017, 97th ed.). Boca Raton, Florida. 2016. ISBN 978-1-4987-5428-6. OCLC 930681942.{{cite book}}: CS1 maint: location missing publisher (link) CS1 maint: others (link)
  8. ^ "Oxalic acid". Immediately Dangerous to Life or Health Concentrations (IDLH). National Institute for Occupational Safety and Health (NIOSH).
  9. ^ Ullmann's Encyclopedia of Industrial Chemistry. Wiley. 2005. pp. 17624/28029. doi:10.1002/14356007. ISBN 9783527306732.
  10. ^ See:
    • Herman Boerhaave, Elementa Chemiae (Basil, Switzerland: Johann Rudolph Im-hoff, 1745), volume 2, pp. 35-38. (in Latin) From p. 35: "Processus VII. Sal nativum plantarum paratus de succo illarum recens presso. Hic Acetosae." (Procedure 7. A natural salt of plants prepared from their freshly pressed juice. This [salt obtained] from sorrel.)
    • Henry Enfield Roscoe and Carl Schorlemmer, ed.s, A Treatise on Chemistry (New York, New York: D. Appleton and Co., 1890), volume 3, part 2, p. 105.
    • See also Wikipedia's articles "Oxalis acetosella" and "Potassium hydrogen oxalate".
  11. ^ See:
    • François Pierre Savary, Dissertatio Inauguralis De Sale Essentiali Acetosellæ [Inaugural dissertation on the essential salt of wood sorrel] (Jean François Le Roux, 1773). (in Latin) Savary noticed that when he distilled sorrel salt (potassium hydrogen oxalate), crystals would sublimate onto the receiver. From p. 17: "Unum adhuc circa liquorem acidum, quem sal acetosellae tam sincerissimum a nobis paratum quam venale destillatione fundit phoenomenon erit notandum, nimirum quod aliquid ejus sub forma sicca crystallina lateribus excipuli accrescat, ..." (One more [thing] will be noted regarding the acid liquid, which furnished for us sorrel salt as pure as commercial distillations, [it] produces a phenomenon, that evidently something in dry, crystalline form grows on the sides of the receiver, ...) These were crystals of oxalic acid.
    • Leopold Gmelin with Henry Watts, trans., Hand-book of Chemistry (London, England: Cavendish Society, 1855), volume 9, p. 111.
  12. ^ See:
    • Torbern Bergman with Johan Afzelius (1776) Dissertatio chemica de acido sacchari [Chemical dissertation on sugar acid] (Uppsala, Sweden: Edman, 1776).
    • Torbern Bergman, Opuscula Physica et Chemica, (Leipzig (Lipsia), (Germany): I.G. Müller, 1776), volume 1, "VIII. De acido Sacchari," pp. 238–263.
  13. ^ Carl Wilhelm Scheele (1784) "Om Rhabarber-jordens bestånds-delar, samt sått at tilreda Acetosell-syran" (On rhubarb-earth's constituents, as well as ways of preparing sorrel-acid), Kungliga Vetenskapsakademiens Nya Handlingar [New Proceedings of the Royal Academy of Science], 2nd series, 5 : 183-187. (in Swedish) From p. 187: "Således finnes just samma syra som vi genom konst af socker med tilhjelp af salpeter-syra tilreda, redan förut af naturen tilredd uti o̊rten Acetosella." (Thus it is concluded [that] the very same acid as we prepare artificially by means of sugar with the help of nitric acid, [was] previously prepared naturally in the herb acetosella [i.e., sorrel].)
  14. ^ "OXALIQUE : Définition de OXALIQUE". www.cnrtl.fr. Retrieved 2024-09-27.
  15. ^ See:
    • F. Wöhler (1824) "Om några föreningar af Cyan" (On some compounds of cyanide), Kungliga Vetenskapsakademiens Handlingar [Proceedings of the Royal Academy of Science], pp. 328–333. (in Swedish)
    • Reprinted in German as: F. Wöhler (1825) "Ueber Cyan-Verbindungen" (On cyanide compounds), Annalen der Physik und Chemie, 2nd series, 3 : 177-182.
  16. ^ a b c Riemenschneider, Wilhelm; Tanifuji, Minoru (2000). "Oxalic Acid". Ullmann's Encyclopedia of Industrial Chemistry. doi:10.1002/14356007.a18_247. ISBN 3527306730.
  17. ^ Eiichi, Yonemitsu; Tomiya, Isshiki; Tsuyoshi, Suzuki and Yukio, Yashima "Process for the production of oxalic acid", U.S. patent 3,678,107, priority date March 15, 1969
  18. ^ Von Wagner, Rudolf (1897). Manual of chemical technology. New York: D. Appleton & Co. p. 499.
  19. ^ "Oxalic acid | Formula, Uses, & Facts | Britannica". 29 August 2024.
  20. ^ Practical Organic Chemistry by Julius B. Cohen, 1930 ed. preparation #42
  21. ^ 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.
  22. ^ Wells, A.F. (1984) Structural Inorganic Chemistry, Oxford: Clarendon Press. ISBN 0-19-855370-6.
  23. ^ Sabine, T. M.; Cox, G. W.; Craven, B. M. (1969). "A neutron diffraction study of α-oxalic acid dihydrate". Acta Crystallographica Section B. 25 (12): 2437–2441. doi:10.1107/S0567740869005905.
  24. ^ Ahmed, F. R.; Cruickshank, D. W. J. (1953). "A refinement of the crystal structure analyses of oxalic acid dihydrate". Acta Crystallographica. 6 (5): 385–392. doi:10.1107/S0365110X53001083.
  25. ^ Bjerrum, J., et al. (1958) Stability Constants, Chemical Society, London.
  26. ^ Haynes, W. M. (ed.). (2014). CRC Handbook of Chemistry and Physics, 95th ed., Boca Raton; London; New York: CRC Press.
  27. ^ Clayton, G. D. and Clayton, F. E. (eds.). Patty's Industrial Hygiene and Toxicology, Volume 2A, 2B, 2C: Toxicology, 3rd ed., New York: John Wiley Sons, 1981–1982, p. 4936.
  28. ^ Rumble, J. (ed.). (2019). CRC Handbook of Chemistry and Physics, 100th ed., CRC Press.
  29. ^ Bowden, E. (1943). "Methyl oxalate". Organic Syntheses: 414; Collected Volumes, vol. 2.
  30. ^ Sahoo, R. N.; Naik, P. K.; Das, S. C. (December 2001). "Leaching of manganese from low-grade manganese ore using oxalic acid as reductant in sulphuric acid solution". Hydrometallurgy. 62 (3): 157–163. Bibcode:2001HydMe..62..157S. doi:10.1016/S0304-386X(01)00196-7. Retrieved 4 December 2021.
  31. ^ DezhiQi (2018). "Extraction of Rare Earths From RE Concentrates". Hydrometallurgy of Rare Earths Separation and Extraction. pp. 1–185. doi:10.1016/B978-0-12-813920-2.00001-5. ISBN 9780128139202.
  32. ^ Kovacs K. A.; Grof P.; Burai L.; Riedel M. (2004). "Revising the mechanism of the permanganate/oxalate reaction". Journal of Physical Chemistry A. 108 (50): 11026–11031. Bibcode:2004JPCA..10811026K. doi:10.1021/jp047061u.
  33. ^ Higgins, James; Zhou, Xuefeng; Liu, Ruifeng; Huang, Thomas T.-S. (1997). "Theoretical Study of Thermal Decomposition Mechanism of Oxalic Acid". The Journal of Physical Chemistry A. 101 (14): 2702–2708. Bibcode:1997JPCA..101.2702H. doi:10.1021/jp9638191.
  34. ^ Harkema, S.; Bats, J. W.; Weyenberg, A. M.; Feil, D. (1972). "The crystal structure of urea oxalic acid (2:1)". Acta Crystallographica Section B. 28 (5): 1646–1648. doi:10.1107/S0567740872004789.
  35. ^ Dutton, M. V.; Evans, C. S. (1996). "Oxalate production by fungi: Its role in pathogenicity and ecology in the soil environment". Canadian Journal of Microbiology. 42 (9): 881–895. doi:10.1139/m96-114..
  36. ^ Rombauer, Rombauer Becker, and Becker (1931/1997). Joy of Cooking, p.415. ISBN 0-684-81870-1.
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  38. ^ Attenborough, David. "Surviving." The Private Life of Plants: A Natural History of Plant Behaviour. Princeton, NJ: Princeton UP, 1995. 265+. "OpenLibrary.org: The Private Life of Plants" Print.
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  40. ^ Frank-Kamemetskaya, Olga; Rusakov, Alexey; Barinova, Ekaterina; Zelenskaya, Marina; Vlasov, Dmitrij (2012). "The Formation of Oxalate Patina on the Surface of Carbonate Rocks Under the Influence of Microorganisms". Proceedings of the 10th International Congress for Applied Mineralogy (ICAM). pp. 213–220. doi:10.1007/978-3-642-27682-8_27. ISBN 978-3-642-27681-1.
  41. ^ Dutton, Martin V.; Evans, Christine S. (1 September 1996). "Oxalate production by fungi: its role in pathogenicity and ecology in the soil environment". Canadian Journal of Microbiology. 42 (9): 881–895. doi:10.1139/m96-114.
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