Industrial synthesis and environmental consequences
Modern production of nitric acid is a complex multi-stage process. It begins with the synthesis of ammonia from hydrogen and atmospheric nitrogen. The final and slowest stage is the interaction of nitrogen dioxide (IV) with water in the presence of oxygen.
It is this final stage that is associated with serious environmental problems. Due to process imperfections, some of the gases are not absorbed, leading to the inevitable emission of nitrogen oxides into the atmosphere. Nitrogen dioxide colors the smoke from factory chimneys a brownish-orange hue, which environmentalists call the “fox tail”. These emissions are highly toxic and can cause massive diebacks of coniferous forests even at a considerable distance from the chemical plant.
Historically, another method was used—metathetical decomposition of nitrates. For this, a mixture of sulfuric acid and potassium nitrate was heated to high temperatures, resulting in the release of the target product's vapors. Today, this method is of historical significance only.
Oxidizing properties and chemical activity
Nitric acid acts as a very strong oxidizer, however, the mechanism of its action strictly depends on the concentration of the solution:
- In concentrated solutions, the Acidum nitricum molecules themselves act as the oxidizing agent. When reacting with non-metals (carbon, sulfur), the corresponding higher oxides or acids are formed, and nitrogen is reduced to monoxide.
- In dilute solutions, the nitrate ion takes over the main oxidizing function, but only with the support of an acidic environment (hydrogen ions). It is important to understand that in an acidic environment, the oxidizing capacity of the nitrate anion is significantly higher than that of the hydrogen cation.
The key rule of nitric acid chemistry: gaseous hydrogen is never released upon interaction with metals. Instead, various nitrogen reduction products are formed. In reactions with active metals and using very dilute solutions (e.g., with aluminum), reduction goes as deep as possible—down to the ammonium cation with the formation of ammonium nitrate.
Metal passivation and applications
Concentrated acid has a specific effect on certain metals. Upon contact with iron, chromium, and aluminum, it instantly forms a thin but incredibly durable oxide film on their surface. This process is called passivation. The film mechanically blocks the reagent's access to the metal, completely stopping further dissolution. Thanks to passivation, concentrated acid is safely transported in standard steel railway tank cars.
The majority of nitric acid produced globally is directed toward creating agricultural fertilizers. In addition, it is widely used in metallurgy for dissolving metals and in organic chemistry as a nitrating agent for introducing a nitro group into molecules.
Nitrates: salts of nitric acid
Nitrates have been successfully synthesized for all known metals. Their main common feature is excellent solubility in water, making them ideal reagents for ion-exchange reactions.
In neutral solutions, the nitrate ion acts as an extremely weak oxidizing agent. But if a strong acid is added to the solution, its oxidizing properties sharply increase, becoming comparable to nitric acid itself.
Upon thermal exposure (heating), all solid nitrates decompose. The nature of the products depends on the chemical activity of the metal:
- Nitrates of active metals lose part of their oxygen upon heating, turning into nitrites.
- Nitrates of less active metals (e.g., copper) break down much more deeply. As a result of the reaction, a metal oxide is formed, free oxygen is actively released, and brown gas—nitrogen dioxide—is produced.
Briefly: answers to topical questions
How does the depth of nitrogen reduction depend on acid concentration and metal activity?
The depth of nitrogen reduction reaches its maximum upon interaction with active metals and the use of very dilute acid.
During the reaction of nitric acid (Acidum nitricum) with metals, hydrogen is not released. Instead, various nitrogen oxides ($NO_2$, $NO$, etc.) are formed. When reacting with active metals, nitrogen can be reduced as deeply as possible—down to the ammonium ion, forming ammonium nitrate ($NH_4NO_3$).
| Metal activity | Acid concentration | Reduction product |
|---|---|---|
| Less active (Cu) | Dilute | Nitrogen oxide ($NO$) |
| Active (Al) | Very dilute | Ammonium ion ($NH_4^+$) |
To remember the metals passivated by concentrated nitric acid, use the mnemonic "AH, Iron!" — Aluminum, Chromium, Iron.
Class and exam questions
Why is hydrogen not evolved during the reaction of nitric acid with metals?
In an acidic medium, the nitrate ion is a much stronger oxidizing agent than the hydrogen ion. Therefore, electrons from the metal are accepted by nitrogen, forming various oxides or ammonium salts, while hydrogen is bound into water.
What is the "fox tail" at a chemical plant?
This is a visual marker of pollution—a brownish-orange smoke consisting of nitrogen dioxide, which is uncontrollably emitted into the atmosphere during the final stages of industrial synthesis.
Does the medium affect the oxidizing capacity of nitrate salts?
Yes, tremendously. In a neutral environment, nitrates are weak oxidizers, but upon acidification of the solution with a strong acid, their activity sharply increases.
What else to study on this topic
- Reaction equations for the oxidation of copper with concentrated and dilute nitric acid.
- In-depth analysis of the nitration mechanism of organic compounds.
- Ammonia synthesis as the initial stage of the industrial chain.
- Oxidation of carbon and sulfur by concentrated solutions.
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