Nitric Acid
Nitric acid, chemically represented as HNO₃, is a highly corrosive and powerful mineral acid with significant industrial, laboratory, and environmental importance. It is a key compound in the nitrogen cycle and serves as a foundational chemical in the production of fertilisers, explosives, and a wide range of industrial materials. Often referred to historically as aqua fortis (Latin for “strong water”), nitric acid has been known to chemists since the Middle Ages and remains one of the most essential acids in modern chemical industry and research.
Chemical Composition and Physical Properties
Nitric acid is composed of hydrogen (H), nitrogen (N), and oxygen (O) atoms in a 1:1:3 ratio. Its molecular weight is 63.01 g/mol. In its pure form, it is a colourless, fuming liquid with a pungent, suffocating odour. However, due to gradual decomposition and the formation of nitrogen dioxide (NO₂), commercial nitric acid often appears yellowish or reddish-brown.
Key physical properties include:
- Boiling point: 83°C
- Freezing point: –42°C
- Density: 1.51 g/cm³ (at 20°C)
- Highly soluble in water, forming an azeotropic mixture containing approximately 68% nitric acid and 32% water (known as concentrated nitric acid).
Nitric acid is a strong oxidising agent, capable of reacting violently with many organic compounds, metals, and reducing agents. It is also corrosive to skin and tissue, producing yellow stains due to the formation of xanthoproteic compounds when it reacts with proteins.
Historical Background
The origins of nitric acid can be traced back to medieval alchemy. Around the 13th century, the German alchemist Albertus Magnus (Albert the Great) first described the preparation of nitric acid by heating saltpetre (potassium nitrate, KNO₃) with vitriol (sulphuric acid). Alchemists referred to it as aqua fortis and used it to dissolve metals such as silver and copper.
During the 16th and 17th centuries, nitric acid became essential in metallurgical and chemical experimentation. It played a critical role in developing the aqua regia mixture—one part nitric acid to three parts hydrochloric acid—capable of dissolving gold and platinum.
By the 19th century, with the advancement of industrial chemistry, nitric acid production shifted from laboratory-scale distillation to large-scale chemical manufacturing processes, especially following the development of the Ostwald process in the early 20th century.
Chemical Behaviour and Reactions
Nitric acid exhibits both acidic and oxidising properties, making it one of the most reactive mineral acids.
- Acidic Reactions: Like other strong acids, nitric acid ionises completely in water:HNO₃→H⁺+NO₃⁻\text{HNO₃} → \text{H⁺} + \text{NO₃⁻}HNO₃→H⁺+NO₃⁻It reacts with bases to form nitrates and water:HNO₃+NaOH→NaNO₃+H₂O\text{HNO₃} + \text{NaOH} → \text{NaNO₃} + \text{H₂O}HNO₃+NaOH→NaNO₃+H₂O
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Oxidising Reactions: Nitric acid acts as a powerful oxidising agent, capable of oxidising metals and non-metals alike. Its products depend on concentration and temperature:
- With copper (moderate concentration):3Cu+8HNO3→3Cu(NO3)2+2NO+4H2O3Cu + 8HNO₃ → 3Cu(NO₃)₂ + 2NO + 4H₂O3Cu+8HNO3→3Cu(NO3)2+2NO+4H2O
- With zinc or magnesium (dilute acid):Zn+2HNO3→Zn(NO3)2+H2↑Zn + 2HNO₃ → Zn(NO₃)₂ + H₂ \uparrowZn+2HNO3→Zn(NO3)2+H2↑
It also oxidises carbon, sulphur, and phosphorus to their highest oxides, producing nitrogen oxides in the process.
- Decomposition: On exposure to light or heat, nitric acid decomposes slowly:4HNO3→2H2O+4NO2+O24HNO₃ → 2H₂O + 4NO₂ + O₂4HNO3→2H2O+4NO2+O2This reaction explains the brown fumes often seen above concentrated nitric acid.
Laboratory and Industrial Preparation
1. Laboratory Preparation: Nitric acid can be produced by heating potassium nitrate (KNO₃) or sodium nitrate (NaNO₃) with concentrated sulphuric acid (H₂SO₄):
NaNO3+H2SO4→NaHSO4+HNO3NaNO₃ + H₂SO₄ → NaHSO₄ + HNO₃NaNO3+H2SO4→NaHSO4+HNO3
The vapours are condensed to obtain pure nitric acid. This method, though simple, is used only for small-scale laboratory synthesis.
2. Industrial Production – The Ostwald Process: Modern nitric acid production employs the Ostwald process, developed by Wilhelm Ostwald in 1902. It involves three main steps:
- Catalytic oxidation of ammonia (NH₃): 4NH3+5O2→4NO+6H2O4NH₃ + 5O₂ → 4NO + 6H₂O4NH3+5O2→4NO+6H2O (platinum-rhodium catalyst, 900°C)
- Oxidation of nitric oxide (NO): 2NO+O2→2NO22NO + O₂ → 2NO₂2NO+O2→2NO2
- Absorption of nitrogen dioxide in water: 3NO2+H2O→2HNO3+NO3NO₂ + H₂O → 2HNO₃ + NO3NO2+H2O→2HNO3+NO
The unreacted nitric oxide is recycled, and the resulting acid is concentrated by distillation. This process underpins most of the world’s nitric acid production today.
Uses and Applications
Nitric acid has a vast range of industrial, scientific, and technological applications:
1. Fertiliser Industry: The majority of nitric acid is used to manufacture ammonium nitrate (NH₄NO₃), a key nitrogen fertiliser. Ammonium nitrate supplies essential nitrogen for plant growth and is also used in controlled-release fertiliser formulations.
2. Explosives: Nitric acid is fundamental in producing nitrated organic compounds such as nitroglycerine, trinitrotoluene (TNT), and nitrocellulose, which are used in explosives, propellants, and munitions.
3. Metal Processing: It is used for pickling stainless steel, etching metals, and purifying precious metals. A mixture of nitric and hydrochloric acids, aqua regia, dissolves noble metals like gold and platinum.
4. Dyes and Pharmaceuticals: Nitric acid is employed in the nitration of aromatic compounds to produce intermediates for dyes, perfumes, and medicines such as nitrophenols and nitrobenzene.
5. Laboratory and Analytical Chemistry: Used as a reagent and oxidiser in chemical analyses, particularly for digestion of samples in elemental analysis and preparation of nitrates.
6. Rocket Propellants and Energy Materials: Concentrated nitric acid, especially when mixed with nitrogen tetroxide, serves as an oxidiser in liquid rocket fuels.
7. Environmental and Educational Uses: In environmental laboratories, nitric acid is utilised for sample digestion, metal extraction, and acidification of water samples for heavy-metal analysis.
Health and Safety Aspects
Nitric acid is highly corrosive and requires careful handling. Contact with skin causes severe burns, yellow staining, and tissue damage. Inhalation of fumes (especially nitrogen dioxide) leads to respiratory irritation, pulmonary oedema, and long-term lung damage.
Safety precautions include:
- Use of acid-resistant gloves, face shields, and fume hoods.
- Storage in amber-coloured glass or acid-resistant plastic containers to prevent light-induced decomposition.
- Neutralisation of spills with sodium bicarbonate or lime slurry.
Environmental hazards arise from nitric acid’s contribution to acid rain when released into the atmosphere, leading to soil and water acidification.
Role in the Nitrogen Cycle and Environment
Nitric acid plays a natural role in the atmospheric nitrogen cycle. It forms in the atmosphere through the oxidation of nitrogen oxides (NOₓ) generated by lightning, volcanic activity, or fossil fuel combustion:
2NO2+H2O→HNO2+HNO32NO₂ + H₂O → HNO₂ + HNO₃2NO2+H2O→HNO2+HNO3
This acid dissolves in rainwater, producing acid rain, which lowers the pH of soil and aquatic systems. Although detrimental in excess, natural nitric acid deposition also contributes trace nitrogen nutrients essential for plant growth.
Analytical and Research Applications
In research laboratories, nitric acid serves as a strong oxidising agent for:
- Digestion of organic matter in biological and environmental samples.
- Preparation of metal nitrates used in analytical standards.
- Spectroscopic analysis and atomic absorption spectroscopy (AAS) sample preparation.
- Cleaning of laboratory glassware to remove metal traces (“acid wash”).
Nitric acid is also vital in isotopic studies, particularly in the preparation of nitrate ions for nitrogen isotope ratio analysis, which aids in environmental tracing and geochemical studies.
Economic and Industrial Significance
Nitric acid is one of the top ten industrial chemicals globally in terms of production volume. Annual worldwide production exceeds 70 million tonnes, driven largely by fertiliser manufacturing. The major producers include China, India, Russia, the United States, and Western Europe.
Its economic importance extends across agriculture, defence, metallurgy, and chemical synthesis. The price of nitric acid is closely tied to ammonia and natural gas markets, given their role in feedstock supply.
Sustainability initiatives are increasingly focusing on reducing emissions of nitrogen oxides (NOₓ) from nitric acid plants. New catalytic converters and green ammonia technologies aim to lower environmental impact while maintaining production efficiency.
Environmental Challenges and Mitigation
Excessive use and release of nitrogen-based compounds, including nitric acid derivatives, contribute to environmental pollution, eutrophication, and acidification. Fertiliser runoff leads to algal blooms in water bodies, while industrial emissions contribute to air quality degradation.
To address these challenges, modern nitric acid plants employ:
- Low-NOₓ catalysts to reduce emissions.
- Carbon capture and utilisation (CCU) for integrating CO₂ and NOₓ recycling.
- Closed-loop water systems to minimise effluent discharge.
Scientific and Technological Advances
Contemporary research explores innovative uses of nitric acid in green chemistry and energy applications. Examples include:
- Electrochemical nitrate reduction for sustainable nitrogen recycling.
- Use in nano-material synthesis, where its oxidising properties help purify carbon nanotubes and graphene.
- Photochemical production studies to understand its role in atmospheric chemistry and ozone regulation.
Nitric acid remains a cornerstone reagent in materials science, catalysis, and environmental monitoring, linking industrial progress with ecological awareness.
Cultural and Historical Influence
In alchemy and early chemistry, nitric acid symbolised transformation and purification. The invention of aqua regia was legendary for its ability to dissolve gold—the “king of metals”—earning nitric acid a near-mystical reputation among medieval scholars.
During the Industrial Revolution, it became a strategic resource for explosives and fertiliser production, shaping both warfare and agriculture. Its dual nature—creative in growth yet destructive in misuse—reflects humanity’s complex relationship with chemistry itself.