Nitric acid

• 1. Introduction

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  Nitric acid, the compound with the formula HNO3, is stored with unusual care. The water-free form is kept in a shatterproof amber glass bottle with twice the volume of head space. Even then, the bottle must be vented every month. The acid slowly decomposes into nitrogen dioxide and water, and the pressure builds. Nitric acid is a highly corrosive mineral acid, and it is far more than a laboratory hazard. Where did this unstable liquid come from, and who first made it? Why do some metals dissolve in it while gold and diamond do not? And how did a chemical that destroys living tissue become a foundation of fertilizer, explosives and rocket propulsion?

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• 2. Discovering The Acid

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  The Latin text De inventione veritatis, meaning On the Discovery of Truth, is the usual starting point. It was written after about 1300 under the borrowed name pseudo-Geber. The conventional view holds that it gave the first description of nitric acid. Mineral acids in general are presumed to trace back to 13th-century European alchemy.

  More texts followed. Works falsely attributed to Albert the Great and Ramon Llull, both 13th-century figures, appeared after 1300. They describe distilling a mixture of niter and green vitriol. Their product goes by the name eau forte, or aqua fortis.

  Eric John Holmyard and Ahmad Y. al-Hassan argue for an earlier origin. They point to Arabic works that mention the acid. One is the Ṣundūq al-ḥikma, or Chest of Wisdom, attributed to Jabir ibn Hayyan in the 8th century. Another is the Taʿwīdh al-Ḥākim, attributed to the Fatimid caliph al-Hakim bi-Amr Allah (985-1021).

  The recipe in the Chest of Wisdom survives in translation. It calls for five parts of pure flowers of nitre, three parts of Cyprus vitriol and two parts of Yemen alum. Each ingredient is ground to dust separately, then loaded into a flask. The flask is plugged with palm fibre and fitted with a glass receiver. The apparatus is turned upside down, and the upper part is warmed with a gentle fire. By the translator's account, what flows down is "an oil like cow's butter."

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• 3. Chemists Close In

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  Johann Rudolf Glauber, working in the 17th century, devised a new route to nitric acid. He distilled potassium nitrate with sulfuric acid. That swap of ingredients gave later chemists a more systematic method than the old vitriol recipes.

  In 1776, Antoine Lavoisier turned to the work of Joseph Priestley. Lavoisier called nitric oxide "nitrous air." He pointed out that the acid could be made from it when "combined with an approximately equal volume of the purest part of common air, and with a considerable quantity of water."

  Henry Cavendish went further in 1785. He determined the precise composition of nitric acid. He also showed that it could be synthesized by passing a stream of electric sparks through moist air.

  Humphry Davy took the electrical approach in 1806. He ran extensive electrolysis experiments on distilled water, using a high voltage battery and electrodes and vessels that would not react. Some of his gold electrode cones doubled as vessels, bridged by damp asbestos. He concluded that nitric acid formed at the anode from nitrogen gas dissolved out of the atmosphere. Davy worked at the scale of a laboratory bench. By 1905, the same electrical idea had reached industry.

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• 4. From Air To Acid

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  In 1905, the Birkeland-Eyde process, also called the arc process, began making nitric acid industrially from atmospheric air. A very hot electric arc forced the nitrogen and oxygen in air to combine into nitric oxide. At 3,000 °C, yields reached roughly 4-5% nitric oxide, and less at lower temperatures. The gas was cooled and oxidized by leftover oxygen to nitrogen dioxide. Absorption towers then dissolved it in water to make dilute acid. The first towers bubbled the gas through water and non-reactive quartz fragments. About 20% of the nitrogen oxides stayed unreacted, so the final towers held an alkali solution to neutralize the rest.

  Around 1913, French engineer Albert Nodon tried a stranger source. Bacteria in peat bogs converted nitrogenous matter into nitrates, and Nodon electrolyzed the resulting calcium nitrate. He dug a pit into the peat and lined its sides with tarred timber stakes. Inside sat a porous earthenware vessel surrounded by crushed limestone and filled with coke around a graphite anode. A glass tube reaching nearly to the bottom pumped acid out, while a second glass pipe fed fresh water in. Cast iron cathodes were sunk into the surrounding peat. Resistance was about 3 ohms per cubic meter, with power around 10 volts. For a one hectare deposit, 6.5 feet deep, output was estimated at more than 600 tons per year.

  The Haber process for making ammonia efficiently arrived in 1913. Once ammonia was cheap, the Ostwald process overtook the energy-hungry arc process. It remains in use today.

  The Ostwald process works by creating conditions in which anhydrous ammonia burns to nitric oxide instead of to dinitrogen. The nitric oxide is then oxidized, often by atmospheric oxygen, to nitrogen dioxide. That dioxide disproportionates in water into nitric acid and more nitric oxide, which supplies the next round. Combined with the Haber process, the whole chain needs only air and natural gas as feedstocks.

  Commercial solutions usually fall between 52% and 68% nitric acid by mass, the maximum that distillation can reach. Concentrated sulfuric acid can dehydrate the mixture further, to 98%. Historically, higher concentrations were made by dissolving extra nitrogen dioxide in the acid. The last plant in the United States to use that process stopped in 2012.

  Laboratory methods borrow from industry. Sodium nitrate reacts with sulfuric acid, and distillation separates the acid from the solid residue. The result is the 68.5% azeotrope, which sulfuric acid or magnesium nitrate can concentrate further. Another route heats copper(II) nitrate to release nitrogen dioxide and oxygen, then passes both through water or hydrogen peroxide. More recently, electrochemical methods have been developed to produce the anhydrous acid directly from concentrated nitric acid.

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• 5. Concentrations And Colors

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  At 68% HNO3 in water, nitric acid forms an azeotrope that boils at 120.5 C at 1 atm. This is the solution known as concentrated nitric acid. It is a colorless liquid at room temperature. An older density scale occasionally appears in the literature, in which the concentrated acid is specified as 42 Baumé. Two solid hydrates are also known, a monohydrate called oxonium nitrate and a trihydrate.

  Above 86% HNO3, the solution is called fuming nitric acid. The pure acid gives off white fumes in air. Acid carrying dissolved nitrogen dioxide gives off reddish-brown vapors instead, which explains the names white fuming and red fuming nitric acid. Dissolved nitrogen dioxide can tint the liquid yellow, and at higher temperatures even red. Light and heat both drive the decomposition that produces it. The dissolved gas also raises the density. Anhydrous acid with 25% nitrogen dioxide is around 1.60 g/cm3, and the maximum density comes at 40% nitrogen dioxide.

  Commercial-grade fuming nitric acid contains 98% HNO3, with a density of 1.50 g/cm3. It is often used in the explosives industry. It is less volatile and less corrosive than the anhydrous acid, at an approximate concentration of 21.4 M. Adding 0.6 to 0.7% hydrogen fluoride produces an inhibited grade, either white or red. The fluoride builds a protective metal fluoride layer, which guards metal tanks against corrosion.

  White fuming nitric acid is very close to anhydrous. It is sold as 99.9% nitric acid by assay, about 24 molar. One specification allows at most 2% water and 0.5% dissolved nitrogen dioxide. The anhydrous liquid has a low viscosity of 0.76 mPa·s. It solidifies at -42 C into white crystals and boils at 83 C.

  The molecule itself has an uneven shape. Two terminal N-O bonds are nearly equivalent and short, at 1.20 and 1.21 Å, because resonance gives each some double bond character. The third N-O bond is longer, at 1.41 Å in the gas phase, since its oxygen carries a hydrogen atom. The NO2 and NOH planes tilt away from each other by 2 degrees, and rotation about the N-OH bond is restricted.

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• 6. Reactions With Metals

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  The pKa of nitric acid is usually reported as less than -1, though there is some disagreement over the exact value. In dilute solution, that means the acid is fully dissociated except in extremely acidic conditions. The pKa rises to 1 at 250 °C. Nitric acid can even act as a base toward sulfuric acid, and that reaction yields the nitronium ion, NO2+. The nitronium ion is the active reagent in aromatic nitration.

  Magnesium, manganese and zinc behave simply, because dilute nitric acid liberates hydrogen from them. Copper and silver are non-active metals, and nitric acid oxidizes them anyway. With dilute acid at ambient temperature, copper reacts in a 3:8 stoichiometry and releases nitric oxide. That gas can meet atmospheric oxygen and become nitrogen dioxide. With more concentrated acid, copper reacts in a 1:4 stoichiometry and gives nitrogen dioxide directly. Most metals end up as nitrates. Tin, arsenic, antimony and titanium end up as the oxides SnO2, As2O5, Sb2O5 and TiO2.

  Chromium, iron and aluminium dissolve readily in dilute nitric acid. Concentrated acid does the opposite. It forms a metal-oxide layer that shields the rest of the metal, a process called passivation. Typical passivation concentrations run from 20% to 50% by volume. Iron, cobalt, chromium, nickel and aluminium are all passivated this way.

  Pure gold and the platinum-group metals do not react with nitric acid at all. Pure gold does dissolve in aqua regia, a blend of concentrated nitric acid and hydrochloric acid. Colored gold is different. Alloys with less gold contain metals such as silver and copper that nitric acid dissolves, which changes the color of the surface. Jewelry shops exploit this as a cheap, quick test for alloys under 14 karats and for rapid checks of gold purity.

  Among the non-metals, concentrated nitric acid oxidizes iodine, white phosphorus and sulfur into HIO3, H3PO4 and H2SO4. It attacks graphite and amorphous carbon but leaves diamond alone, so it can separate diamond from graphite by oxidizing the graphite. Many of these reactions are kinetically sluggish. A small amount of nitrous acid greatly speeds them up.

  Proteins also react. In the xanthoproteic test, concentrated nitric acid is added to a sample and the mixture is heated. If proteins containing amino acids with aromatic rings are present, the mixture turns yellow. Adding a base such as ammonia shifts the color to orange. The same yellow-to-orange shift turns up when the acid touches skin.

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• 7. Fertilizer To Rockets

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  Of the 26 million tonnes of nitric acid produced annually as of 1987-75-80% went into fertilizer. Neutralizing the acid with ammonia gives ammonium nitrate. The other main applications are explosives, nylon precursors and specialty organic compounds.

  The nitro group explains the explosives. When nitric acid reacts with organic material, a hydroxyl group typically strips a hydrogen from the molecule to form water, and a nitro group takes the hydrogen's place. A mixture of nitric and sulfuric acids installs nitro groups on aromatic compounds by electrophilic aromatic substitution. Concentrated sulfuric acid or oleum absorbs the excess water. Nitroglycerin and trinitrotoluene, or TNT, are made this way. Many possible byproducts form, so these reactions need careful thermal control, and the byproducts must be removed.

  Not every nitro compound is dangerous. Some are shock- and thermally-sensitive explosives, but a few are stable enough for munitions and demolition. Others are more stable still and serve as synthetic dyes and medicines, among them metronidazole. Chemists can also reduce a nitro group to an amine, which turns nitrobenzenes into aniline compounds.

  Oxidation is the second skill. Adipic acid, the precursor to nylon, is produced on a large scale by oxidizing "KA oil" with nitric acid. KA oil is a mixture of cyclohexanone and cyclohexanol.

  Rockets use the same oxidizing power. Red fuming nitric acid, white fuming nitric acid, mixtures with sulfuric acid, and versions with hydrogen fluoride inhibitor have all served as oxidizers in liquid-fueled rockets. Inhibited red fuming nitric acid was one of three liquid fuel components for the BOMARC missile.

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• 8. Odd Jobs

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  Gold refiners combine nitric acid with hydrochloric acid as aqua regia to dissolve noble metals, with gold forming chloroauric acid. Nitric acid alone can turn copper into cupric nitrate. Because it dissolves certain metals selectively and dissolves many metal salts, it is useful in gold parting. Recrystallization and selective precipitation of the resulting salts can purify gold and other metals beyond 99.9%.

  Analytical chemists use very dilute nitric acid, between 0.5% and 5.0%, as the matrix for finding trace metals by ICP-MS, ICP-AES, GFAA and flame atomic absorption. The acid must be ultrapure trace metal grade, because tiny amounts of metal ions could skew the results. For turbid water, sludge and solid samples, digestion typically uses a 50% solution mixed with Type 1 deionized water. In electrochemistry, the acid serves as a doping agent for organic semiconductors and helps purify raw carbon nanotubes.

  Woodworkers use roughly 10% nitric acid to age pine and maple artificially. The color it produces is a grey-gold that closely resembles very old wax- or oil-finished wood.

  Corrosion is the point in etching and cleaning. Printmakers etch with it, and it pickles stainless steel and cleans silicon wafers in electronics. Nital, a solution of nitric acid, water and alcohol, etches metals to reveal their microstructure, and ISO 14104 details that well known procedure. Used alone or with hydrochloric acid, nitric acid cleans glass cover slips and slides for high-end microscopy. It also cleans glass before silvering when silver mirrors are made.

  Commercial blends of 5-30% nitric acid and 15-40% phosphoric acid clean food and dairy equipment. They mainly remove precipitated calcium and magnesium compounds, whether deposited from the process stream or left by hard water. The phosphoric acid helps passivate ferrous alloys against corrosion. Nitric acid also works as a spot test for alkaloids like LSD, producing different colors depending on the alkaloid.

  At the far end of the list sits nuclear fuel reprocessing. In PUREX and related methods, nitric acid dissolves many different actinides. The resulting nitrates are converted into complexes that can be extracted selectively, which separates the metals from each other.

  For an acid spill on skin, the standard first aid is irrigation with large quantities of water. Washing continues for at least 10-15 minutes, which cools the tissue around the burn and prevents secondary damage. Contaminated clothing is removed immediately, and the skin underneath is washed thoroughly.

  The major hazard is chemical burns. Nitric acid hydrolyzes proteins and fats, which breaks down living tissue such as skin and flesh. Concentrated acid stains human skin yellow through its reaction with keratin, and the stains turn orange when neutralized. Xanthoproteic acid forms where the acid meets epithelial cells. Local skin color changes therefore signal inadequate safety precautions.

  Systemic effects are unlikely, and the substance is not considered a carcinogen or mutagen. The danger is local and immediate. As a strong oxidizing agent, nitric acid can also react violently with many compounds. It is one of the most common types of acid used in acid attacks, which makes careful handling and the quick rinse described above matter far beyond the laboratory.

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