Oxalic Acid
Formula: C₂H₂O₄ — Ethanedioic acid
Appearance: White crystalline powder
Hazard: Toxic · Irritant · Kidney hazard
Properties
The simplest dicarboxylic acid — two carboxyl groups and nothing else. White crystalline solid, moderately soluble in water. Moderately strong for an organic acid (pKa₁ = 1.25, pKa₂ = 4.27 — the first dissociation approaches mineral acid strength). Powerful reducing agent; reduces permanganate and dichromate quantitatively, making it useful as a titration standard. The oxalate ion (C₂O₄²⁻) is an excellent chelating agent for calcium and iron — it forms insoluble calcium oxalate (the main component of kidney stones) and dissolves iron oxides by complexation. Found naturally in rhubarb, spinach, and wood sorrel.
Historical Context
Oxalic acid was first isolated from wood sorrel (Oxalis acetosella) by the Swedish chemist Carl Wilhelm Scheele in 1776 — the same chemist who discovered chlorine, manganese, and citric acid. Its strong affinity for iron was recognized early, and it became a standard bleaching and cleaning agent in the textile and papermaking trades.
The name “oxalic” comes from the genus Oxalis, though the acid is far more widespread in the plant world. Rhubarb leaves contain enough oxalic acid to be dangerously toxic if eaten in quantity (though rhubarb stalks are safe and the acid concentration much lower). Spinach’s oxalic acid partially blocks calcium absorption — a real but modest nutritional concern.
In analytical chemistry, oxalic acid was a cornerstone of early volumetric analysis. Its clean, quantitative reaction with potassium permanganate under acidic conditions made it the reference standard for permanganate titrations, a role it still occasionally fills today.
Experiments
Rust Removal: Make a 5–10% solution in warm water and soak rusted iron or steel objects for 30–60 minutes. The oxalate ion forms a soluble complex with Fe³⁺, dissolving rust without attacking the base metal vigorously. Objects emerge clean with minimal pitting. Rinse thoroughly and dry immediately to prevent re-rusting.
Calcium Precipitation: Add dilute oxalic acid to a calcium chloride solution — a white precipitate of calcium oxalate (CaCO₄) forms immediately, even from dilute solutions. This demonstrates the very low solubility of the salt and the selectivity of precipitation chemistry.
Permanganate Titration: React dilute acidified potassium permanganate solution with oxalic acid — the intense purple of permanganate fades to colorless as Mn⁷⁺ is reduced to Mn²⁺ and the oxalate is oxidized to CO₂. The endpoint is dramatic: the last drop of permanganate produces a persistent pink/purple tint.
Experiments using this chemical:
- Colors of Permanganate — Reduction of permanganate to Mn²⁺
Safety
Moderate hazard — toxic; harmful if ingested; kidney irritant.
Wear gloves — skin absorption is possible and the compound is irritating. Do not ingest even small amounts. The oxalate ion precipitates calcium from blood, causing hypocalcemia in large doses; kidney stones are the hazard from chronic low-level exposure. Dispose of solutions by diluting and flushing with water — the quantities in experiments are not environmentally significant.
Incompatible with: Strong oxidizers, especially permanganate and dichromate (vigorous reaction, sometimes exothermic); silver compounds (insoluble Ag₂C₂O₄); reactive metals (slow hydrogen evolution)