Nitric Acid – Preparation, Properties and Uses

Thirty vocabulary flashcards summarising essential terms, reactions, apparatus, and industrial aspects associated with nitric acid from the lecture notes.

Nitric Acid (HNO₃)

A strong, monobasic, highly corrosive mineral acid that acts as a powerful oxidising agent.

Aqua Fortis

The historical name for nitric acid, meaning “strong water,” used by alchemists because of its corrosive action on metals.

Laboratory Preparation of Nitric Acid

Distillation of KNO₃ or NaNO₃ with concentrated H₂SO₄ in an all-glass apparatus below 200 °C to collect HNO₃ vapours.

Potassium Nitrate (KNO₃)

Also called Bengal saltpetre or nitre; a solid reactant heated with conc. H₂SO₄ to produce laboratory nitric acid.

Sodium Nitrate (NaNO₃)

Also called Chile saltpetre; an alternative nitrate used with conc. H₂SO₄ to prepare HNO₃ in the lab.

Concentrated Sulphuric Acid (H₂SO₄)

A strong, non-volatile acid that displaces the more volatile nitric acid from nitrates during laboratory preparation.

Potassium Bisulphate (KHSO₄)

The acid salt by-product formed when KNO₃ is distilled with concentrated H₂SO₄ below 200 °C.

Sodium Bisulphate (NaHSO₄)

The acid salt by-product obtained when NaNO₃ is treated with concentrated H₂SO₄ below 200 °C.

Temperature Limit (< 200 °C)

Critical control in the lab setup to avoid glass damage, HNO₃ decomposition, and formation of hard Na₂SO₄/K₂SO₄ crusts.

Ostwald’s Process

Industrial manufacture of HNO₃ via catalytic oxidation of NH₃ to NO, oxidation to NO₂, and absorption in water.

Platinum Gauze Catalyst

The catalyst used at 700-800 °C in the catalytic chamber of Ostwald’s process to oxidise NH₃ to NO.

Catalytic Chamber

First stage of Ostwald’s process where NH₃ and air (1 : 10) contact Pt gauze, forming NO and steam exothermically.

Oxidation Chamber

Second stage of Ostwald’s process (≈50 °C) where NO is further oxidised by excess O₂ to NO₂.

Absorption Tower

Packed with acid-resistant quartz; absorbs NO₂ in water and excess air to yield ~50 % nitric acid.

Constant Boiling Mixture (68 % HNO₃)

Azeotrope of HNO₃ and water that boils at 121 °C and prevents concentration by simple boiling.

Fuming Nitric Acid (98 % HNO₃)

Highly concentrated nitric acid obtained by distilling the 68 % azeotrope with conc. H₂SO₄ under reduced pressure.

Decomposition of Nitric Acid

Self-oxidation producing NO₂, H₂O and O₂; responsible for yellow-brown colour in stored acid.

Nitrogen Dioxide (NO₂)

Reddish-brown gas formed by HNO₃ decomposition and released during its oxidising reactions with metals and non-metals.

Nascent Oxygen [O]

Reactive oxygen atom liberated during HNO₃ reduction, responsible for the acid’s strong oxidising power.

Oxidising Agent Nature of HNO₃

Ability to oxidise metals, non-metals and compounds while itself being reduced to NO, NO₂, N₂O, etc.

Passivity of Iron and Aluminium

Formation of a protective oxide film by conc. or fuming HNO₃ that renders these metals temporarily inert.

Aqua Regia

Mixture of conc. HNO₃ : HCl (1 : 3); generates nascent Cl₂ able to dissolve noble metals like Au and Pt.

Brown Ring Test

Analytical test for nitrates/HNO₃ using freshly prepared FeSO₄ and conc. H₂SO₄ to form a brown FeSO₄·NO ring.

Nitroso Ferrous Sulphate (FeSO₄·NO)

Unstable brown complex formed at the interface during the brown ring test, confirming nitrate presence.

Fixation of Atmospheric Nitrogen

Natural conversion of N₂ to nitrates via lightning-produced NO and NO₂ that dissolve to form HNO₃ in rain.

Acid Rain

Rainwater containing dissolved nitric (and sulphuric) acids formed from atmospheric NO₂ and other oxides.

Xanthoproteic Reaction

Yellow staining of skin proteins by HNO₃ due to formation of xanthoproteic acid; illustrates the acid’s corrosive action.

Explosive Manufacture

Use of HNO₃ to nitrate toluene, glycerol and cellulose producing TNT, nitroglycerine and guncotton respectively.

Silver Nitrate (AgNO₃)

A light-sensitive nitrate prepared with HNO₃, widely used in photography and analytical chemistry.

Nitrous Oxide (N₂O)

‘Laughing gas’ generated by thermal decomposition of NH₄NO₃, one of the thermal reactions of nitrates.

Explain the comprehensive properties of Nitric Acid (HNO₃) as a strong mineral acid.

Nitric Acid is a strong, monobasic, highly corrosive mineral acid. Its key property is its powerful oxidizing nature, primarily due to the formation of nascent oxygen when it reacts. This oxidizing power enables it to react with and oxidize numerous metals and non-metals, often producing different reduction products of nitrogen depending on the concentration of the acid and the reactivity of the substance. It is also a significant component in various industrial processes and is known for its ability to nitrate organic compounds.

Describe the historical significance and origin of the name “Aqua Fortis” for nitric acid.

"Aqua Fortis," meaning “strong water,” was the historical name given to nitric acid by alchemists. This name reflects their early observations of its potent corrosive action, particularly on numerous metals, highlighting its distinction from more benign liquids and its chemical strength. It underscores its long history of recognition as a powerful chemical agent.

Outline the detailed laboratory preparation process for Nitric Acid, including the specific reagents, apparatus, and control measures.

Nitric acid is prepared in the laboratory by the distillation of a nitrate salt (like KNO3 or NaNO3) with concentrated H2SO4. The reaction occurs in an all-glass apparatus (retort) to prevent corrosion. The process involves heating the mixture below 200 \degree C to ensure that the more volatile nitric acid vaporizes and is then condensed and collected. The chemical reaction is typically: KNO3(s) + H2SO4(l) \xrightarrow{

Explain the crucial reasons for maintaining the temperature below 200 \degree C during the laboratory preparation of nitric acid.

Maintaining the temperature below 200 \degree C is critical for several reasons: Firstly, higher temperatures can cause the decomposition of nitric acid, leading to a poorer yield and the formation of nitrogen dioxide (NO2). Secondly, at higher temperatures, the glass apparatus (retort) can be severely corroded by nitric acid vapors, potentially leading to breakage. Lastly, if sodium nitrate is used, heating above 200 \degree C would lead to the formation of hard, crusty sodium sulfate (Na2SO4), which is difficult to remove from the apparatus, whereas sodium bisulphate (NaHSO4) formed below 200 \degree C is easily removable.

Elaborate on the role of concentrated Sulphuric Acid (H2SO4) in the laboratory synthesis of nitric acid.

Concentrated Sulphuric Acid (H2SO4) is used in the laboratory preparation of nitric acid because it is a strong, non-volatile acid. Its higher boiling point allows it to displace the more volatile nitric acid from nitrate salts (KNO3 or NaNO3) when heated. This displacement is a key principle of preparing a more volatile acid from its salt using a less volatile, stronger acid.

Describe the full Ostwald’s Process for the industrial manufacture of nitric acid, detailing the principal reactions and conditions in each of its three main stages.

Ostwald’s Process is the industrial method for synthesizing nitric acid.

1. Catalytic Oxidation Chamber: Ammonia (NH3) and air (in a 1:10 ratio) are passed over a platinum gauze catalyst heated to 700-800 \degree C. This exothermic reaction oxidizes ammonia to nitric oxide (NO) and steam: 4NH3(g) + 5O2(g) \xrightarrow{Pt, 700-800 \degree C} 4NO(g) + 6H2O(g).
2. Oxidation Chamber: The nitric oxide (NO) formed is cooled to approximately 50 \degree C and further oxidized by excess oxygen from the air to nitrogen dioxide (NO2): 2NO(g) + O2(g) \xrightarrow{} 2NO_2(g).
3. Absorption Tower: Nitrogen dioxide (NO2) is then absorbed in water in an absorption tower packed with acid-resistant quartz. Excess air is supplied to ensure complete oxidation of any unreacted NO to NO2, which then reacts with water to form nitric acid and nitric oxide, the latter being recycled: 3NO2(g) + H2O(l) \xrightarrow{} 2HNO_3(aq) + NO(g). The resulting nitric acid is about 50 \%.

Explain the significance and function of the Platinum Gauze Catalyst in the Ostwald’s process.

The platinum gauze catalyst is crucial in the first stage (catalytic oxidation chamber) of Ostwald's process. It facilitates the catalytic oxidation of ammonia (NH_3) to nitric oxide (NO) at high temperatures (700-800 \degree C). Without the catalyst, this reaction would not proceed efficiently or selectively at the desired rate, enabling the primary conversion step for nitric acid production.

Describe the characteristics of the “Constant Boiling Mixture” of nitric acid and water, and explain how “Fuming Nitric Acid” is obtained.

A solution of 68 \% nitric acid (HNO3) and 32 \% water forms a constant boiling mixture or azeotrope, which boils at 121 \degree C at standard atmospheric pressure. This means its composition does not change upon boiling, preventing further concentration by simple distillation. To obtain "fuming nitric acid" (approximately 98 \% HNO3), the 68 \% azeotrope is distilled with concentrated Sulphuric Acid (H2SO4) under reduced pressure. The H2SO4 acts as a dehydrating agent, absorbing water and allowing the more concentrated HNO_3 to distill off.

Explain the phenomenon of "Decomposition of Nitric Acid" and its observable consequence.

Nitric acid is inherently unstable and undergoes self-oxidation decomposition, especially when exposed to light or heat. This leads to the formation of nitrogen dioxide (NO2), water, and oxygen: 4HNO3(aq) \xrightarrow{} 4NO2(g) + 2H2O(l) + O_2(g). The accumulation of reddish-brown nitrogen dioxide gas dissolved in the acid is responsible for the characteristic yellow-brown color observed in stored nitric acid.

Detail how "Nascent Oxygen" is formed from nitric acid and its role in the acid's strong oxidizing power.

Nascent oxygen (symbolized as [O]) is a highly reactive, newly formed oxygen atom. It is liberated when nitric acid acts as an oxidizing agent and is itself reduced. For example, in dilute nitric acid's reaction, the reduction product often includes nascent oxygen: 2HNO3 \xrightarrow{} 2NO + H2O + 3[O]. This nascent oxygen is extremely reactive and is primarily responsible for the powerful oxidizing ability of nitric acid, enabling it to oxidize various substances that molecular oxygen (O_2) would not.

Explain the diverse oxidizing agent nature of Nitric Acid when reacting with metals and non-metals, illustrating by varying reduction products.

Nitric acid is a strong oxidizing agent, and its reduction products vary depending on its concentration, the temperature, and the nature of the substance it oxidizes. Generally, concentrated HNO3 is reduced to nitrogen dioxide (NO2), while dilute HNO3 can be reduced to nitric oxide (NO), nitrous oxide (N2O), nitrogen (N2), or even ammonium ions (NH4^+, which can form ammonium nitrate). For example, with active metals, very dilute HNO3 may produce N2O or NH4NO3, whereas with less active metals or concentrated acid, NO_2 or NO are typical products. This variability highlights its flexible oxidizing power.

Describe the phenomenon of "Passivity" observed for iron and aluminium when exposed to concentrated or fuming nitric acid.

Passivity is the state where a metal, despite being thermodynamically reactive, becomes chemically inert due to the formation of a thin, invisible, and protective oxide layer on its surface. When iron or aluminium are treated with concentrated or fuming nitric acid, these metals form a dense, adherent, and non-porous layer of their respective oxides (Fe3O4 or Al2O3). This oxide film prevents further reaction between the metal and the acid, rendering the metals "passive" or temporarily unreactive.

Provide a detailed explanation of "Aqua Regia," its composition, and its unique ability to dissolve noble metals.

Aqua Regia, meaning "royal water," is a highly corrosive, fuming yellow or reddish liquid. It is a fresh mixture of concentrated nitric acid (HNO3) and concentrated hydrochloric acid (HCl) in a molar ratio of 1:3. Its unique ability to dissolve noble metals like gold (Au) and platinum (Pt), which are inert to individual acids, stems from a synergistic effect. Nitric acid oxidizes the metal to its ions (e.g., Au^{3+}), while hydrochloric acid reacts with these metal ions to form stable chloro-complexes (e.g., [AuCl4]^{−}). This formation of stable complexes effectively removes the metal ions from solution, shifting the equilibrium of the oxidation reaction forward, thus allowing continuous dissolution of the noble metal.

Elaborate on the "Brown Ring Test" as an analytical method, detailing the procedure, required reagents, and the chemical species responsible for the brown ring.

The Brown Ring Test is a classic qualitative analytical test used to detect the presence of nitrate ions (NO3^−) or nitric acid. The procedure involves taking a solution containing nitrate, adding freshly prepared ferrous sulfate (FeSO4) solution, and then carefully adding concentrated Sulphuric Acid (H2SO4) along the side of the test tube. The H2SO4, being denser, settles at the bottom, forming a distinct layer. At the interface of the two layers, a characteristic brown ring forms. This brown ring is due to the formation of an unstable nitroso ferrous sulfate complex, [Fe(H2O)5NO]^{2+} or FeSO_4·NO. The initial reactions involve nitrate being reduced to nitric oxide (NO) by Fe^{2+}, and then NO reacting with unreacted Fe^{2+} to form the brown complex.

What is the chemical composition and significance of "Nitroso ferrous sulphate" in the context of the Brown Ring Test?

Nitroso ferrous sulphate (FeSO4·NO or more accurately, the pentaaquanitrosyliron(II) cation, [Fe(H2O)_5NO]^{2+}) is the unstable brown complex responsible for the characteristic "brown ring" in the Brown Ring Test. Its formation is the definitive indication of the presence of nitrate ions. In this complex, nitric oxide (NO), formed from the reduction of nitrate by Fe^{2+}, coordinates with an Fe^{2+} ion, leading to the brown color. The complex is unstable and decomposes upon heating or agitation, causing the brown ring to disappear.

Describe the natural process of "Fixation of Atmospheric Nitrogen" and the role of nitric acid in it.

Fixation of atmospheric nitrogen refers to the natural processes by which inert atmospheric nitrogen gas (N2) is converted into more reactive nitrogen compounds, such as nitrates, which are usable by plants. One significant natural process involves lightning. During lightning discharges, the extremely high temperatures promote the direct reaction between atmospheric nitrogen and oxygen to form nitric oxide (NO): N2(g) + O2(g) \xrightarrow{lightning} 2NO(g). This NO quickly oxidizes further to nitrogen dioxide (NO2): 2NO(g) + O2(g) \xrightarrow{} 2NO2(g). The NO2 then dissolves in rainwater to form nitric acid (HNO3) and nitrous acid (HNO2): 2NO2(g) + H2O(l) \xrightarrow{} HNO3(aq) + HNO_2(aq). The nitric acid so formed falls to the earth as acid rain (though beneficial for nitrogen supply when in normal quantities) and reacts with alkaline substances in the soil to form nitrates, vital plant nutrients.

Explain the formation of "Acid Rain" and nitric acid's contribution to it.

Acid rain is precipitation (rain, snow, fog, etc.) that contains elevated levels of nitric and sulphuric acids, making it significantly more acidic than normal rain. Nitric acid contributes substantially to acid rain. It is primarily formed when nitrogen oxides (NOx), released from sources like vehicle exhausts, industrial combustion processes, and natural lightning, undergo reactions in the atmosphere. Specifically, nitrogen dioxide (NO2) reacts with water (H2O) and oxygen (O2) in the air to form nitric acid (HNO3): 4NO2(g) + O2(g) + 2H2O(l) \xrightarrow{} 4HNO_3(aq). This dissolved nitric acid then falls back to the Earth, causing environmental damage.

Describe the "Xanthoproteic Reaction" and its implications when nitric acid comes into contact with skin.

The Xanthoproteic reaction is a chemical reaction that occurs when nitric acid (HNO_3) comes into contact with proteins containing aromatic amino acids (e.g., tyrosine, tryptophan, phenylalanine). The nitric acid nitrates the benzene rings in these amino acids, forming yellow-colored nitro compounds. This is why skin (which contains proteins) turns yellow upon contact with nitric acid. This reaction serves as an illustration of the strong corrosive and nitrating action of nitric acid on organic matter, particularly proteins, and highlights the need for caution when handling the acid.

Discuss the significant role of nitric acid in the "Explosive Manufacture" industry.

Nitric acid is a cornerstone chemical in the manufacture of various explosives, as it is a powerful nitrating agent. It is used to introduce nitro (-NO_2) groups into organic compounds, forming highly unstable and energetic nitro compounds. Key examples include:

• TNT (Trinitrotoluene): Produced by nitrating toluene with concentrated nitric acid (often with sulphuric acid).
• Nitroglycerine: Formed by the nitration of glycerol.
• Guncotton (Nitrocellulose): A highly flammable and explosive material formed by nitrating cellulose.
  The introduction of multiple nitro groups imparts high energy and instability, making these compounds suitable for explosive applications.

Besides its use in explosives, nitric acid has several other crucial applications. It is used in the preparation of various nitrate salts, such as Silver Nitrate (AgNO3), which is light-sensitive and widely employed in photography (as silver halides are formed from AgNO3) and in analytical chemistry. A vastly important application is in the production of fertilizers. Nitric acid reacts with ammonia to produce ammonium nitrate ($$NH_