Abstract

Nitric acid (HNO₃) is a highly corrosive mineral acid and strong oxidizing agent that represents one of the most important industrial chemicals worldwide. This inorganic compound exists as a colorless liquid at room temperature but often develops a yellow tint due to decomposition into nitrogen oxides. The commercial form typically consists of a 68% aqueous solution that forms an azeotrope with water. Nitric acid demonstrates complete dissociation in aqueous solutions with a pKa of -1.4, classifying it as a strong acid. Its molecular structure exhibits resonance stabilization with bond lengths of 1.20-1.21 Å for the terminal N-O bonds and 1.41 Å for the N-OH bond. Primary applications include fertilizer production, explosives manufacturing, metal processing, and organic synthesis through nitration reactions. Annual global production exceeds 60 million metric tons, with the Ostwald process serving as the dominant industrial manufacturing method.

Introduction

Nitric acid stands as a fundamental chemical in modern industrial chemistry, with applications spanning fertilizer production, explosives manufacturing, metallurgy, and organic synthesis. This mineral acid belongs to the class of strong inorganic acids and serves as a primary reagent for introducing nitro functional groups into organic molecules. Historical records indicate its discovery in medieval alchemy, with pseudo-Geber's "De inventione veritatis" (circa 1300) containing the earliest known descriptions of its preparation through distillation of nitrate salts with vitriol. The compound's systematic name according to IUPAC nomenclature is nitric acid, with the molecular formula HNO₃ and a molar mass of 63.01 g/mol. Its significance in industrial chemistry stems from its dual functionality as both a strong acid and powerful oxidizing agent, properties that enable diverse chemical transformations across multiple sectors.

Molecular Structure and Bonding

Molecular Geometry and Electronic Structure

The nitric acid molecule exhibits a planar structure with Cₛ symmetry, though slight deviations from perfect planarity occur due to the orientation of the hydroxyl group. According to VSEPR theory, the nitrogen atom adopts sp² hybridization with bond angles of approximately 130° between oxygen atoms. The molecular geometry features two nearly equivalent terminal N-O bonds measuring 1.20-1.21 Å, which display partial double bond character due to resonance stabilization. The N-OH bond measures 1.41 Å in the gas phase, significantly longer than the terminal bonds. Resonance structures show delocalization of electrons across the nitrate group, with the major canonical forms contributing to the molecule's stability. The nitrogen atom carries a formal positive charge while the oxygen atoms bear partial negative charges, creating a molecular dipole moment of 2.17 ± 0.02 D. Restricted rotation occurs about the N-OH single bond with a rotational barrier of approximately 9 kJ/mol.

Chemical Bonding and Intermolecular Forces

Covalent bonding in nitric acid involves polar bonds with significant ionic character due to the electronegativity difference between nitrogen (3.04) and oxygen (3.44). The O-N-O bond angle of approximately 130° results from the repulsion between oxygen lone pairs. Intermolecular forces include strong hydrogen bonding between the acidic proton and oxygen atoms, with an average hydrogen bond energy of 25-30 kJ/mol. The extensive hydrogen bonding network in concentrated aqueous solutions contributes to the compound's high boiling point relative to its molecular weight. The molecule's polarity, with an estimated dielectric constant of approximately 50 for pure acid, facilitates its solvation of ionic compounds. Comparative analysis with related acids shows shorter N-O bond lengths than in nitrous acid (HNO₂) but longer than those in nitrate ions, reflecting the intermediate bonding character in nitric acid.

Physical Properties

Phase Behavior and Thermodynamic Properties

Anhydrous nitric acid exists as a colorless, mobile liquid at room temperature with a density of 1.512 g/cm³ at 20 °C. The compound freezes at -41.6 °C to form white crystalline solid and boils at 83 °C under standard atmospheric pressure. Commercial concentrated nitric acid typically constitutes a 68% aqueous solution that forms an azeotrope with water, boiling at 120.5 °C with a density of 1.41 g/cm³. The heat of formation measures -207 kJ/mol with an entropy of 146 J/(mol·K). Specific heat capacity ranges from 2.1 J/(g·K) for the pure acid to 3.5 J/(g·K) for dilute solutions. The dynamic viscosity of anhydrous nitric acid measures 0.76 mPa·s at 25 °C, increasing significantly with water addition. Vapor pressure reaches 48 mmHg at 20 °C for the concentrated solution, with the pure acid exhibiting higher volatility. Two stable hydrates form under appropriate conditions: the monohydrate (HNO₃·H₂O) melting at -37.7 °C and the trihydrate (HNO₃·3H₂O) melting at -18.5 °C.

Spectroscopic Characteristics

Infrared spectroscopy of nitric acid reveals characteristic vibrational frequencies at 3550 cm⁻¹ (O-H stretch), 1710 cm⁻¹ (N=O stretch), 1320 cm⁻¹ (N-O stretch), and 950 cm⁻¹ (O-H bend). Proton NMR spectroscopy shows a singlet at approximately 11 ppm in dimethyl sulfoxide solution, consistent with strongly acidic protons. Nitrogen-14 NMR exhibits a resonance at -20 ppm relative to nitromethane. UV-Vis spectroscopy demonstrates weak absorption in the visible region with λmax at 302 nm (ε = 15 M⁻¹cm⁻¹) for the pure acid, while decomposed samples show additional absorption at 400 nm due to nitrogen dioxide formation. Mass spectrometric analysis reveals a molecular ion peak at m/z 63 with major fragmentation patterns including loss of OH (m/z 46), O (m/z 47), and formation of NO₂⁺ (m/z 46). Raman spectroscopy shows strong bands at 1300 cm⁻¹ and 1040 cm⁻¹ corresponding to symmetric and asymmetric stretching vibrations of the nitrate group.

Chemical Properties and Reactivity

Reaction Mechanisms and Kinetics

Nitric acid functions as a strong oxidizing agent with standard reduction potential of 0.96 V for the NO₃⁻/NO couple. Oxidation reactions typically proceed through initial formation of nitrous acid intermediates, which catalyze further reaction through autocatalytic mechanisms. The decomposition kinetics follow second-order behavior with an activation energy of 80 kJ/mol for the reaction 4HNO₃ → 2H₂O + 4NO₂ + O₂. Thermal stability decreases with increasing temperature and concentration, with half-life of approximately 10 hours at 100 °C for concentrated solutions. Reaction with metals exhibits diverse behavior depending on concentration and metal reactivity. Dilute acid (2-10%) reacts with active metals such as zinc and magnesium to produce hydrogen gas: Zn + 2HNO₃ → Zn(NO₃)₂ + H₂. Concentrated acid oxidizes copper through the reaction Cu + 4H⁺ + 2NO₃⁻ → Cu²⁺ + 2NO₂ + 2H₂O with a rate constant of 0.15 M⁻¹s⁻¹ at 25 °C. Passivation occurs with iron, chromium, and aluminium due to formation of protective oxide layers that inhibit further reaction.

Acid-Base and Redox Properties

Nitric acid demonstrates complete dissociation in aqueous solution with pKa = -1.4, classifying it among the strongest mineral acids. The acid dissociation constant increases with temperature, reaching pKa = 1 at 250 °C. In concentrated sulfuric acid, nitric acid acts as a base, protonating to form the nitronium ion (NO₂⁺) with equilibrium constant K ≈ 22. This property enables its use in electrophilic aromatic substitution reactions. Redox properties include multiple reduction pathways depending on concentration and reducing agent strength. Dilute acid typically reduces to nitric oxide (NO) with standard reduction potential E° = 0.96 V, while concentrated acid reduces to nitrogen dioxide (NO₂) with E° = 0.80 V. Further reduction to nitrous oxide (N₂O) occurs under strong reducing conditions with E° = 1.11 V. The compound exhibits stability in acidic conditions but decomposes rapidly in alkaline media through nitrate formation. Autoprotolysis occurs according to the equilibrium 2HNO₃ ⇌ NO₂⁺ + NO₃⁻ + H₂O with equilibrium constant K = 2.4 × 10⁻² at 25 °C.

Synthesis and Preparation Methods

Laboratory Synthesis Routes

Laboratory-scale synthesis typically employs the reaction of nitrate salts with concentrated sulfuric acid. Sodium nitrate reacts according to the equation NaNO₃ + H₂SO₄ → HNO₃ + NaHSO₄, with distillation at 83 °C separating the nitric acid from the sodium bisulfate residue. This method produces the 68% azeotrope, which can be dehydrated using concentrated sulfuric acid or magnesium nitrate to achieve higher concentrations. Alternative routes include thermal decomposition of copper(II) nitrate at 170-200 °C: 2Cu(NO₃)₂ → 2CuO + 4NO₂ + O₂, followed by absorption of the nitrogen dioxide in water or hydrogen peroxide. Electrochemical methods have been developed for producing anhydrous acid through oxidation of nitrogen species at platinum electrodes. Small quantities of high-purity nitric acid can be prepared by direct reaction of nitrogen dioxide with water: 3NO₂ + H₂O → 2HNO₃ + NO, though this method requires recycling of the nitric oxide byproduct.

Industrial Production Methods

Industrial production primarily utilizes the Ostwald process, which involves catalytic oxidation of ammonia. The process begins with ammonia combustion at 850-900 °C over platinum-rhodium catalysts: 4NH₃ + 5O₂ → 4NO + 6H₂O. The nitric oxide rapidly oxidizes to nitrogen dioxide: 2NO + O₂ → 2NO₂, which subsequently disproportionates in water: 3NO₂ + H₂O → 2HNO₃ + NO. The nitric oxide is recycled to maximize yield, with overall stoichiometry NH₃ + 2O₂ → HNO₃ + H₂O. Modern plants achieve conversion efficiencies exceeding 95% with production capacities up to 1,000 metric tons per day. Concentration from the 60-70% product to 98-99% fuming nitric acid employs extractive distillation with concentrated sulfuric acid or magnesium nitrate solutions. Environmental considerations include catalytic reduction of nitrogen oxide emissions and energy optimization through heat recovery systems. The global production capacity exceeds 60 million metric tons annually, with China, the United States, and Europe representing the largest producing regions.

Analytical Methods and Characterization

Identification and Quantification

Qualitative identification employs the brown ring test, where addition of iron(II) sulfate to a nitrate solution followed by careful addition of concentrated sulfuric acid produces a brown ring of FeNO²⁺ at the interface. Quantitative analysis typically utilizes ion chromatography with conductivity detection, achieving detection limits of 0.1 mg/L for nitrate ions. Spectrophotometric methods employ reactions with brucine or chromotropic acid, producing colored complexes measurable at 410 nm and 430 nm respectively. Titrimetric methods include alkalimetric titration with standardized sodium hydroxide solution using methyl orange indicator, though potentiometric endpoint detection provides greater accuracy for concentrated solutions. Gasometric methods measure nitrogen oxide evolution upon reduction with devarda's alloy or other reducing agents. Inductively coupled plasma optical emission spectroscopy enables simultaneous determination of nitric acid and metal impurities with detection limits below 1 ppb for most elements.

Purity Assessment and Quality Control

Commercial nitric acid grades are specified according to concentration and impurity profiles. Reagent grade (68-70%) typically contains less than 0.00001% heavy metals, 0.0001% chloride, and 0.0005% sulfate. Trace metal grade acid undergoes additional purification through sub-boiling distillation to achieve parts-per-trillion impurity levels. Quality control parameters include specific gravity (1.413-1.416 at 20 °C for 68% acid), residual nitrogen dioxide content (max 0.2% for white fuming grade), and non-volatile matter (max 0.005%). Stability testing monitors nitrogen dioxide formation through UV-Vis spectroscopy at 400 nm. Industrial grades are classified according to concentration ranges: weak acid (30-55%), concentrated acid (55-70%), and fuming acid (85-100%). Specifications for rocket propellant grades include maximum water content (0.5% for WFNA), hydrogen fluoride content (0.6-0.7% for inhibited grades), and aluminum corrosion rates (max 0.0005 inches per month).

Applications and Uses

Industrial and Commercial Applications

Fertilizer production consumes approximately 80% of global nitric acid output, primarily for ammonium nitrate manufacture through neutralization with ammonia: NH₃ + HNO₃ → NH₄NO₃. The explosives industry utilizes nitric acid for nitration of organic compounds including toluene to trinitrotoluene (TNT), glycerol to nitroglycerin, and cellulose to nitrocellulose. Adipic acid production for nylon synthesis employs nitric acid oxidation of cyclohexanol/cyclohexanone mixtures. Metal processing applications include pickling of stainless steel, etching of copper printed circuits, and purification of precious metals through aqua regia (3HCl:HNO₃). The nuclear industry uses nitric acid for fuel reprocessing through the PUREX process, where it dissolves uranium and plutonium from spent nuclear fuel. Smaller-scale applications include wood aging through controlled oxidation of lignin, semiconductor cleaning, and analytical chemistry as a digestion acid for trace metal analysis.

Research Applications and Emerging Uses

Research applications focus on nitric acid's role in atmospheric chemistry, particularly in stratospheric ozone depletion through formation of polar stratospheric clouds. Materials science employs nitric acid for surface functionalization of carbon nanotubes and graphene through oxidation reactions. Electrochemical applications include doping of organic semiconductors and synthesis of metal-organic frameworks. Emerging technologies utilize nitric acid in energy storage systems, particularly for flow batteries using nitrate redox couples. Catalysis research investigates nitric acid as a nitrating agent for sustainable synthesis of pharmaceutical intermediates and fine chemicals. Environmental applications involve nitric acid in flue gas treatment for simultaneous NOx removal and ammonia neutralization. The compound's role in nitrogen cycle chemistry continues to be investigated through isotopic tracing studies using ¹⁵N-labeled nitric acid. Recent patent activity focuses on improved production methods, recycling technologies, and specialized formulations for electronics manufacturing.

Historical Development and Discovery

The discovery of nitric acid dates to medieval alchemy, with the earliest reliable descriptions appearing in the 13th century. Pseudo-Geber's "De inventione veritatis" (circa 1300) described distillation of potassium nitrate with iron sulfate, producing what was termed "aqua fortis" (strong water). Arabic alchemical works including Jabir ibn Hayyan's "Ṣundūq al-ḥikma" contained similar recipes using Cyprus vitriol (copper sulfate) and Yemen alum. The 17th century brought systematic investigation, with Johann Rudolf Glauber developing improved production methods using potassium nitrate and sulfuric acid. Antoine Lavoisier recognized nitric acid's composition from nitrogen and oxygen in 1776, while Henry Cavendish established its precise elemental composition in 1785 through electric spark experiments in air. The 19th century saw Humphry Davy's electrolytic experiments and the development of the first industrial production methods. The Birkeland-Eyde process (1905) utilized arc discharge for atmospheric nitrogen fixation, quickly superseded by the Ostwald process (1913) following the invention of the Haber process for ammonia synthesis. Technological improvements throughout the 20th century focused on catalyst development, energy efficiency, and environmental protection.

Conclusion

Nitric acid represents a cornerstone of industrial chemistry with unparalleled importance in fertilizer production, explosives manufacturing, and chemical synthesis. Its unique combination of strong acidity and powerful oxidizing capability enables diverse chemical transformations across multiple sectors. The compound's molecular structure exhibits interesting bonding characteristics with resonance stabilization and restricted rotation. Industrial production through the Ostwald process demonstrates remarkable efficiency and scalability, though environmental considerations continue to drive process improvements. Future research directions include development of sustainable production methods, novel applications in energy storage, and advanced materials processing. The compound's fundamental chemistry continues to present challenges in understanding reaction mechanisms, decomposition pathways, and atmospheric implications. As global demand for nitrogen-based products increases, nitric acid will maintain its position as an essential chemical building block for modern society.