Exploring Nitric Acid: Properties, Uses and Synthesis


Nitric acid is a highly corrosive and powerful oxidizing agent that plays a crucial role in various industries. With its distinct properties and wide range of applications, it is important to understand its chemical and physical properties, synthesis methods, versatile uses, and safety guidelines for handling and storage. This article delves into the world of nitric acid, providing valuable insights for both professionals and enthusiasts.

Nitric acid appearance

Nitric Acid: Chemical and Physical Properties

Nitric acid, chemically represented as HNO3, is a highly corrosive and powerful oxidizing agent that exhibits a range of distinctive chemical and physical properties. As a colorless liquid with a strong and pungent odor, it is highly soluble in water, resulting in the formation of its commonly known aqueous solution, “aqua fortis.” Nitric acid is classified as a strong monoprotic inorganic acid, meaning it donates a single proton (H+) when dissolved in water.

Pure nitric acid possesses a density of approximately 1.51 g/cm³ and a boiling point of 83 °C. These properties make it a highly volatile substance, requiring careful handling and storage. Moreover, its powerful oxidizing nature enables it to readily accept electrons from other substances, allowing it to participate in a wide array of chemical reactions. Nitric acid is known for its reactivity with both organic and inorganic compounds, making it an essential reagent in numerous chemical processes.

Nitric acid resonance

Nitric acid is typically alluded to as a potent acid at room temperature. There is some discussion with respect to the worth of the acid dissociation constant, with the pKa worth being typically detailed as underneath −1. This implies that the nitric acid in weakened arrangement is completely dissociated, barring in incredibly acidic solutions. The pKa worth ascends to 1 at a temperature of 250 °C.

Nitric Acid Acid-base Properties

Nitric acid can act as a base when connected with an acid like sulfuric acid:

HNO3 + 2H2SO4 ⇌ [NO2]+ + [H3O]+ + 2HSO−4;

Equilibrium constant: K ≈ 22

The nitronium particle, [NO2]+, is the dynamic reagent in fragrant nitration responses. Since nitric acid possesses both acid and alkaline characteristics, it can go through a autoprotolysis reaction, corresponding to the self-ionization of water:

2HNO3 ⇌ [NO2]+ + NO−3 + H2O

Due to its highly corrosive nature, nitric acid can cause severe damage to various materials upon contact. It can corrode metals, organic compounds, and even some plastics. This corrosive property is attributed to its ability to donate protons and react with materials through acid-base reactions. Consequently, caution must be exercised when handling nitric acid to prevent accidents and damage to equipment.

Nitric Acid Reactions With Metals

Nitric acid is known to interact with almost all metals, the reaction rate depending on the acid concentration and the metal type. Dilute nitric acid shows the characteristic behavior of an acid when it comes into contact with the majority of metals, such as the liberation of hydrogen; e.g. magnesium, manganese and zinc:

Mg + 2HNO3 → Mg(NO3)2 + H2

Mn + 2HNO3 → Mn(NO3)2 + H2

Zn + 2HNO3 → Zn(NO3)2 + H2

Nitric acid has the capability of oxidizing non-metals that are not reactive. The products formed via reaction with these metals of lower electropositivity depend upon the temperature and the acid concentration. For instance, when exposed to a dilute nitric acid at room temperature a 3:8 stoichiometry reaction occurs:

3Cu + 8HNO3 → 3Cu(NO3)2 + 2NO + 4H2O

Additionally, the nitric oxide released may react with atmospheric oxygen, generating nitrogen dioxide. For higher concentrations of nitric acid, nitrogen dioxide is formed directly with a stoichiometry of 1:4:

Cu + 4H+ + 2NO−3 → Cu2+ + 2NO2 + 2H2O

Some elements are converted to nitrates, such as tin (Sn), arsenic (As), antimony (Sb) and titanium (Ti), which are converted to SnO2, As2O5, Sb2O5 and TiO2 respectively.

Precious metals such as gold and some platinum-group metals are nontoxic to the acid, although it may corrode silver and other less noble metals present inside some gold alloys. Nitric acid might also be employed to produce a colour change on the alloy surface, essentially detecting any low-gold concentrations.

The strong oxidizing properties of nitric acid allow it to react with most non-metallic compounds, and such reactions may yield explosive end products. On rare occasions, noble metals, certain alloys and reduced metals could be immune to the effects of nitric acid. Normally, the high concentration of nitric acid produces nitrogen dioxide via oxidation, and the presence of small quantities of nitrous acid (HNO2) accelerates this reaction.

Although iron, chromium and aluminium can dissolve in dilute nitric acid, concentrated nitrogen acid will lead to the formation of a protective passivating layer between the metal and the acid, which is termed passivation. The passivation of metals with nitric acid is observed in elements such as iron, cobalt, chromium, nickel and aluminium, at concentrations which tend to be in the range of 20% to 50% by volume.

Reactions with non-metals

Nitric acid is a formidable oxidizing agent, interacting with many organic materials and capable of causing explosive reactions. Hydroxyl groups commonly eliminate hydrogen from the organic molecule, forming water while the nitro group supplants the hydrogen. The nitration of organic substances with nitric acid is the main mode of production for various explosives like nitroglycerin and trinitrotoluene (TNT). As an array of incredibly instable byproducts may be created, these chemical processes necessarily require temperature regulation and the discarding of byproducts to extract the desired product.

When the acid is not in contact with metallic elements apart from nitrogen, oxygen, inert gases, silicon, and halogens not including iodine, the components generally oxide to their highest oxidization state, forming acid salts plus nitrogen dioxide in concentrated acid and nitric oxide in weak acidity.

C (graphite) + 4HNO3 → CO2 + 4NO2 + 2H2O;

3C (graphite) + 4HNO3 → 3CO2 + 4NO + 2H2O.

Concentrated nitric acid oxidizes I2, P4, and S8, respectively into HIO3, H3PO4, and H2SO4. It will cause oxidization in graphite and noncrystalline carbon, however it fails to react with diamond; it can separate diamond from the graphite which it oxidizes.

Nitric Acid Synthesis

Nitric acid is created through a synthesis of nitrogen dioxide (NO2) and water.

4NO2 + 2H2O → 2HNO3 + NO + NO2 + H2O

The general reaction is:

3NO2 + H2O → 2HNO3 + NO

Typically, the nitric oxide created by the reaction is restored to its original form via oxidation by the oxygen in the atmosphere, resulting in the production of supplementary nitrogen dioxide.

Introducing nitrogen dioxide through hydrogen peroxide can assist with increasing the quantity of acid produced.

2NO2 + H2O2 → 2HNO3

Commonly used commercial nitric acid concentrations typically range from 52% to 68%. Wilhelm Ostwald, a German chemist, coined the name Ostwald process for the manufacturing of nitric acid. This approach involves the oxidation of anhydrous ammonia to obtain nitric oxide by using a platinum or rhodium gauze catalyst at temperatures of around 227 °C and pressure of 9 atmospheres (910 kPa).

4NH3(g) + 5O2(g) → 4NO(g) + 6H2O(g) (ΔH = −905.2 kJ/mol)

Nitric oxide then interacts with oxygen from the atmosphere to result in nitrogen dioxide.

2NO(g) + O2(g) → 2NO2(g) (ΔH = −114 kJ/mol)

This afterwards is taken in by water to generate nitric acid and nitric oxide.

3NO2(g) + H2O(l) → 2HNO3(aq) + NO(g) (ΔH = −135.74 kJ/mol)

The nitric oxide can be changed back into an oxidized form, or if the final step happens in oxygen-containing atmosphere:

4NO2(g) + O2(g) + 2H2O(l) → 4HNO3(aq)

The HNO3 in liquid form that is obtained can be increased in concentration up to approximately 68% by mass through distillation. Dehydration with a strong H2SO4 can elevate the concentration even further, up to 98%. At this point, the ultimate product can be synthesized from nitrogen, hydrogen, and oxygen, which all originate from the air and natural gas. The process of conversion relies on ammonia which can be produced from the Haber method.

Laboratory synthesis

In the lab, nitric acid can be manufactured by breaking down copper(II) nitrate through the application of heat, creating nitrogen dioxide and oxygen gases which are subsequently passed through water to form nitric acid.

2Cu(NO3)2 → 2CuO + 4NO2 + O2

Then, following the Ostwald process:

2NO2 + H2O → HNO2 + HNO3

Alternatively, an equal molar amount of any nitrate salt such as sodium nitrate can be combined with sulfuric acid (H2SO4), and then those ingredients can be boiled off at nitric acid’s boiling temperature of 83°C. A rugged residue of metal hydrogen sulfate stays put in the distillation container. The burgundy fuming nitric acid created can be modified to colorless nitric acid.

NaNO3 + H2SO4 → HNO3 + NaHSO4

The diluted NOx can be simply eliminated by reduced pressure at room temperature (approx. 10-30 min at 200 mmHg or 27 kPa) resulting in white fumes of nitric acid. Alternatively, by reducing the temperature and pressure, this process may be done in a single step avoiding the production of nitrogen dioxide.

Making Concentrated (68%) Nitric Acid

To then raise the concentration of nitric acid up to 68% in the laboratory, distillation is carried out in combination with either sulphuric acid or magnesium nitrate functioning as drying agents. As a measure of precaution, the distillation should be done at low pressure to stop decomposition of the acid. Incidentally, in industrial production, highly concentrated nitric acid is attained by combining nitrogen dioxide with 68% nitric acid in a absorption tower. Whether nitrogen oxides remain combined in solution or vaporise during the procedure decides if one is left with white or red fumes of nitric acid. In a recent development, electrochemical means have been put forth for attain anhydrous acid from a concentrated nitric acid material.

Versatile Applications of Nitric Acid in Various Industries

Nitric acid finds extensive use in numerous industries due to its versatile properties. Some of the key applications include:

  • Fertilizer Production: Nitric acid is a vital component in the manufacturing of ammonium nitrate, a commonly used fertilizer. It provides essential nitrogen for plant growth, enhancing crop yields and overall agricultural productivity.
  • Explosives and Propellants: Nitric acid is utilized in the production of explosives and propellants. It is a key ingredient in the manufacturing of dynamite, TNT, and other powerful explosives. In the aerospace industry, it is used in the production of rocket propellants.
TNT synthesis
  • Metallurgy: Nitric acid plays a significant role in metal refining and etching processes. It is used for dissolving and purifying metals like gold and silver, as well as for etching designs on metal surfaces.
  • Dyes and Pigments: Nitric acid is involved in the production of dyes, pigments, and colorants. It is utilized in the synthesis of various organic compounds, including aniline dyes, which find applications in the textile industry.
  • Pharmaceutical and Laboratory Applications: Nitric acid is used in the pharmaceutical industry for the synthesis of various drugs and pharmaceutical intermediates. It is also an important reagent in laboratory settings for chemical analysis and research purposes.

Ensuring Safety with Nitric Acid: Handling, Storage, and Protective Measures

Due to its corrosive and hazardous nature, handling and storing nitric acid require adherence to strict safety guidelines. Here are some essential measures to ensure safety:

  • Personal Protective Equipment (PPE): When working with nitric acid, proper PPE should be worn, including safety goggles, gloves, lab coats, and closed-toe shoes. This protects the eyes, skin, and respiratory system from potential contact or inhalation.
  • Ventilation: Nitric acid should be handled in well-ventilated areas or under fume hoods to prevent the accumulation of hazardous vapors. Adequate ventilation helps maintain a safe working environment.
  • Storage: Nitric acid should be stored in tightly sealed containers made of compatible materials, such as glass or high-density polyethylene (HDPE). Storage areas should be cool, dry, and away from direct sunlight. It is important to keep nitric acid separate from incompatible substances to prevent reactions and potential hazards.
  • Handling and Spillage: Nitric acid should be handled with care, avoiding contact with skin and eyes. In case of spillage, appropriate spill kits should be readily available, and spills should be cleaned up promptly using neutralizing agents and proper disposal methods.


In conclusion, a thorough grasp of the chemical and physical properties of nitric acid is crucial for safely harnessing its versatility and power in various industries. By adhering to proper handling, storage, and protective measures, professionals can create a secure working environment while maximizing the benefits of nitric acid. By respecting its corrosiveness and reactivity and implementing safety protocols, individuals can effectively utilize nitric acid to drive advancements in fields such as metallurgy, explosives production, and pharmaceutical development. Through ongoing research and a commitment to safety, nitric acid’s potential can be harnessed to contribute to advancements and innovations across diverse sectors.


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