Oxalic acid

Oxalic acid, also known systematically as ethanedioic acid, is the simplest dicarboxylic acid. Its chemical formula is HO−C(=O)−C(=O)−OH, often written as (COOH)₂ or H₂C₂O₄. It appears as a white crystalline solid that dissolves in water to give a colorless solution. The name comes from early researchers who first isolated it from flowering plants in the genus *Oxalis*, commonly called wood-sorrels. The substance occurs naturally in many foods, but ingesting large amounts can be toxic, and concentrated forms can cause chemical burns.

It is a much stronger acid than acetic acid and acts as a reducing agent. Its conjugate bases—hydrogenoxalate (HC₂O₄⁻) and oxalate (C₂O₄²⁻)—are chelating agents for metal cations. Oxalic acid is used as a cleaning agent, particularly for rust removal, because it forms a water-soluble ferric iron complex called ferrioxalate. It typically occurs as a dihydrate, with the formula H₂C₂O₄·2H₂O.

The preparation of oxalic acid salts from plants was known by 1745, when Dutch botanist and physician Herman Boerhaave isolated a salt from wood sorrel. In 1773, François Pierre Savary of Fribourg, Switzerland, isolated oxalic acid itself from its salt in sorrel. In 1776, Swedish chemists Carl Wilhelm Scheele and Torbern Olof Bergman produced oxalic acid by reacting sugar with concentrated nitric acid; Scheele called it *socker-syra* (sugar acid). By 1784, Scheele had shown that this "sugar acid" was identical to oxalic acid from natural sources. The modern name was introduced in 1787 by de Morveau, Lavoisier, and co-authors. In 1824, German chemist Friedrich Wöhler obtained oxalic acid by reacting cyanogen with ammonia in aqueous solution, an experiment that may represent the first synthesis of a natural product.

Industrially, oxalic acid is mainly manufactured by oxidizing carbohydrates like glucose with nitric acid or air in the presence of vanadium pentoxide. Another process uses oxygen to regenerate nitric acid, with precursors including glycolic acid and ethylene glycol; as of 2011, only Mitsubishi in Japan used this method. A newer approach involves oxidative carbonylation of alcohols to form diesters of oxalic acid, which are then hydrolyzed. About 120,000 tonnes are produced annually. Historically, oxalic acid was obtained exclusively by treating sawdust with caustics like sodium or potassium hydroxide, then acidifying the resulting oxalate with mineral acids such as sulfuric acid. It can also be formed by heating sodium formate in the presence of an alkaline catalyst.

In the laboratory, oxalic acid can be prepared by oxidizing sucrose with nitric acid and a small amount of vanadium pentoxide as a catalyst. The dihydrate can be converted to the anhydrous form by heating or azeotropic distillation.

Anhydrous oxalic acid exists as two polymorphs: in one, hydrogen bonding creates a chain-like structure; in the other, it forms a sheet-like structure. Because it is both acidic and hydrophilic, the anhydrous material is used in esterifications. The dihydrate has space group C₅²h–P2₁/n, with lattice parameters a = 611.9 pm, b = 360.7 pm, c = 1205.7 pm, β = 106°19′, and Z = 2. Key inter-atomic distances are C−C 153 pm, C−O1 129 pm, and C−O2 119 pm.

Oxalic acid’s pKa values vary in the literature, but the 2019 100th edition of the CRC lists them as 1.25 and 3.81. It is relatively strong compared to other carboxylic acids. It undergoes typical carboxylic acid reactions, forming esters like dimethyl oxalate (melting point 52.5–53.5 °C) and an acid chloride called oxalyl chloride.

Transition metal oxalate complexes are numerous, including the drug oxaliplatin. Oxalic acid reduces manganese dioxide in manganese ores, allowing leaching of the metal by sulfuric acid. It is an important reagent in lanthanide chemistry, where hydrated lanthanide oxalates form readily in strongly acidic solutions as dense, easily filtered crystals largely free of nonlanthanide contamination. Thermal decomposition of these oxalates yields the oxides, the most commonly marketed form of these elements.

Oxalic acid and oxalates can be oxidized by permanganate in an autocatalytic reaction. Oxalic acid vapor decomposes at 125–175 °C to carbon dioxide and formic acid. Photolysis with 237–313 nm UV light produces carbon monoxide and water. Evaporation of a 2:1 molar ratio solution of urea and oxalic acid yields a solid crystalline compound, H₂C₂O₄·2CO(NH₂)₂, consisting of stacked two-dimensional networks held together by hydrogen bonds.

At least two pathways exist for enzyme-mediated oxalate formation. In one, oxaloacetate from the Krebs cycle is hydrolyzed to oxalate and acetic acid by oxaloacetase. It also arises from dehydrogenation of glycolic acid, produced by ethylene glycol metabolism.

Early investigators isolated oxalic acid from wood-sorrel (*Oxalis*). Members of the spinach family, brassicas (cabbage, broccoli, brussels sprouts), sorrel, and umbellifers are high in oxalates.