Nitric Acid (HNO3): Properties, Reactions, Applications

The alchemist first created nitric acid (HNO₃) by combining saltpeter (KNO₃) and concentrated sulfuric acid (H₂SO₄). The process involved the method to separate gold and silver, wherein silver demonstrated solubility while gold exhibited insolubility. The acid was referred to as aqua fortis, a term denoting its potent nature. In 1776, Lavoisier successfully demonstrated the presence of oxygen in nitric acid.

In laboratory settings, the production of nitric acid can be achieved through the process of distillation. This involves the utilization of a mixture consisting of solid potassium nitrate (KNO₃) and condensed sulfuric acid (H₂SO₄).

_distillation

KNO₃ + conc.H₂SO₄ Δ→ KHSO₄ + HNO₃

Properties of Nitric Acid

In this analysis, we will examine the distinct categories of properties associated with nitric acid, which includes both its physical and chemical attributes.

Physical Properties of HNO₃

• Pure nitric acid is a visually transparent liquid that exhibits a density of 1.5 g/cm³. The substance demonstrates a boiling point of 85 degrees Celsius when subjected to standard atmospheric pressure and experiences solidification at a temperature of -42 degrees Celsius. It is heavier than water.
• The commercially available nitric acid solution exhibits a concentration of 68%, rendering it a continuously boiling mixture.
• Undiluted nitric acid exhibits corrosive properties and possesses a high propensity for rapidly decomposing organic substances. The application of this substance results in a yellow discoloration of the skin, and even brief exposure to the acid induces the formation of yellow blisters on the skin.
• Nitric acid exhibits complete solubility in water across all ratios.
• The substance exhibits a distinctive odor characterized by its strong, sharp scent and a flavor profile that is sour in taste.
• The term “fuming nitric acid” refers to anhydrous HNO₃ that is infused with dissolved oxides of nitrogen, specifically N₂O₅, and NO₂. The substance under consideration is a volatile liquid with a yellow coloration, renowned for its significant hygroscopic properties.

Chemical Properties

Nitric Acid as an Acidic Agent

(i) Nature of Aqueous Solution

• Nitric acid is classified as a highly potent acid that undergoes complete dissociation when dissolved in an aqueous solution.

HNO₃ (aq) → H⁺ (aq) + NO₃^– (aq)

^{Hydrogen ion Nitrogen ion}

or, HNO₃ + H₂O → H₃O⁺ + NO₃^–

• Consequently, the substance exhibits the property of causing a blue litmus paper to undergo a color change to red. Additionally, it possesses the ability to neutralize basic oxides, hydroxides, carbonates, and bicarbonates, and it also engages in a reaction with ammonia.

Action With Bases

• Nitric acid (HNO₃) exhibits the ability to neutralize bases, resulting in the formation of metallic nitrates.
• The reaction between potassium hydroxide (KOH) and nitric acid (HNO₃) produces potassium nitrate (KNO₃) and water (H₂O) in an aqueous solution.

KOH (aq) + HNO₃ (aq) → KNO₃ (aq) + H₂O (l)

• Similarly, the reaction between sodium hydroxide (NaOH) and nitric acid (HNO₃) results in the formation of sodium nitrate (NaNO₃) and water (H₂O).

NaOH + HNO₃ → NaNO₃ + H₂O

• The reaction between copper(II) oxide (CuO) in a solid state and nitric acid (HNO₃) in aqueous form results in a chemical transformation.

CuO (s) + 2 HNO₃ (aq) → Cu(NO₃)₂ (aq) + H₂O (l)

^{Basic oxide}

• The reaction between Ammonia (NH₃) and nitric acid (HNO₃) results in the formation of ammonium nitrate.

NH₃ + HNO₃ → NH₄NO₃

(iii) Action With Carbonates

• The reaction between nitric acid and metallic carbonates results in the liberation of carbon dioxide.

CO₃^2- (s) + 2 H⁺ (aq) → H₂O (l) + CO₂ (g)

CaCO₃ + 2 HNO₃ → Ca(NO₃)₂ + CO₂ + H₂O

KHCO₃ + HNO₃ → KNO₃ + CO₂ + H₂O

(iv) Action on Metals (Highly Electropositive)

• The fundamental characteristic of a strong acid is that, when diluted, it should react with metals that are more electropositive, resulting in the release of hydrogen gas. However, in the case of HNO₃, this behavior is not observed in practice.

• The nascent hydrogen, which is produced when metal reacts with HNO₃, is immediately oxidized by additional acid to form water. The reduction products of the acid can include nitrogen dioxide, nitric oxide, nitrous oxide, or even ammonia in the form of ammonium nitrate.

• If a very dilute solution of HNO₃ (about 1%) is used with magnesium or manganese, some hydrogen will be produced. This hydrogen will escape further oxidation due to the highly diluted nature of the acid. In this case, the hydrogen may not be pure, as it is likely accompanied to some extent by the gaseous reduction products of the acid.

Mg + very dil.(1%) 2 HNO₃ → Mg(NO₃)₂ + H₂↑

Mn + very dil.(1%) 2 HNO₃ → Mn(NO₃)₂ + H₂↑

Nitric Acid as an Oxidizing Agent

Nitric acid is an excellent oxidizing agent due to its ability to decompose and release nascent oxygen.

• Concentrated nitric acid (HNO₃) decomposes to produce nitrogen dioxide (NO₂), water (H₂O), and oxygen (O). The reaction can be represented as follows:

2 HNO₃ → 2 NO₂ + H₂O + [O]

• When moderately concentrated (1:1) HNO₃ is used, it produces nitric oxide along with nascent oxygen. The balanced chemical equation for the reaction is:

(1:1) 2 HNO₃ → 2 NO + H₂O + 3[O]

• Even more dilute acid yields nitrous oxide

dil. 2 HNO₃ → N₂O + H₂O + 4[O]

Nitric acid is distinguished by a notable concentration of oxygen and demonstrates a tendency for decomposition, leading to the creation of nitrogen and oxygen oxides.

• Similar to other oxidizing agents, nitric acid functions as an electron acceptor. There are multiple methods by which this might be achieved. Two of the more significant ones are:

4 HNO₃ (aq) + 2e^– → 2 NO₃^– (aq) + 2 H₂O (l) + 2 NO₂ (g)

8 HNO₃ (aq) + 6e^– → 6 NO₃^– (aq) + 4 H₂O (l) + 2 NO (g)

giving nitrogen dioxide and nitrogen oxide respectively.

• The supply of electrons is facilitated by the reducing agent, which actively participates in the chemical reaction. Frequently, this substance is a metal, such as copper. This process results in the formation of ions.

Cu → Cu²⁺ + 2e^–

• By combining this ionization with the two equations given above, the reaction shown between copper and nitric acid is obtained.
• Similar behavior can be observed among various metals, although the specifics may differ depending on the metal, the acid’s condition, and the temperature used.
• In addition to its well-known corrosive properties, nitric acid exhibits the ability to induce oxidation reactions on specific non-metallic elements and compounds that lack oxygen, such as hydrogen sulfide (H₂S) and potassium iodide (KI), as well as those with a low oxygen content, such as hydrogen peroxide (H₂O₂) and sulfur dioxide (SO₂).

(i) Action of HNO₃ on Metals

Nitric acid has the ability to dissolve various metals, with the exception of gold and platinum. The effect of HNO₃ on metals is reliant upon:

(a) Concentration of the acid

(b) Temperature at which the reaction is carried out

(c) Nature of the metal used

(d) Presence of other impurities

In each reaction, the metal undergoes conversion into its corresponding nitrate, while simultaneously generating nascent hydrogen. The hydrogen atom undergoes a reaction with nitric acid, resulting in the formation of various compounds such as NO₂, NO, N₂O, and NH₃, among others. Furthermore, the hydrogen atom is oxidized to produce H₂O.

Armstong’s View of the Action

(a) Metal (M) + 2 HNO₃ → M(NO₃)₂ + 2[H] Primary reaction

(b) 2 HNO₃ + 2 H → 2 H₂O + 2 NO₂

2 HNO₃ + 4 H → 3 H₂O + N₂O₃

2 HNO₃ + 6 H → 4 H₂O + 2 NO { Secondary reaction

2 HNO₃ + 8 H → 5 H₂O + N₂O

2 HNO₃ +10 H → 6 H₂O + N₂

2 HNO₃ + 16 H → 6 H₂O + 2 NH₃

The reaction between nitric acid and various metals, including those positioned below hydrogen in the activity series, is observed. This observation appears to contradict the established principle that only metals positioned above hydrogen in the reactivity series can react with acids. However, the reaction between nitric acid and metals positioned below hydrogen is not primarily due to the acidic nature of nitric acid but rather stems from its oxidizing properties.

In the aforementioned reactions, the oxidizing agent is the nitrate ion as opposed to the nitric acid. Platinum (Pt) and gold (Au) exhibit resistance to oxidation. The production of water is a consistent outcome of the oxidizing reaction involving HNO₃. The formation of free hydrogen is not observed, except in the case of active metals such as Mg and Mn, which release hydrogen gas when reacting with cold and highly diluted HNO₃.

So, Armstrong’s view is that metals above hydrogen in the electrochemical series react with HNO₃ by making new oxygen.

Action With Metals

(a) Reaction With Metals More Electropositive Then Hydrogen

Action With Zinc

Variations in the concentration of acid lead to the formation of distinct products.

With highly diluted HNO₃

Zn + 2 HNO₃ → Zn(NO₃)₂ + 2[H]4

HNO₃ + 8[H] → NH₃ + 3 H₂O

HNO₃ + NH₃ → NH₄NO₃

4 Zn + 10 HNO₃ → 4 Zn(NO₃)₂ + NH₄NO₃ + 3 H₂O

With diluted HNO₃

Zn + 2 HNO₃ → Zn(NO₃)₂ + 2[H]4

2 HNO₃ + 8[H] → N₂O + 5 H₂O

4 Zn + 10 HNO₃ → 4 Zn(NO₃)₂ + 5 H₂O + N₂O

With Moderately Concentrated HNO₃

Zn + 2 HNO₃ → Zn(NO₃)₂ + 2[H]3

HNO₃ + 3[H] → NO + 2 H₂O]2

3 Zn + 8 HNO₃ → 3 Zn(NO₃)₂ + 2 NO + 4 H₂O

With Concentrated HNO₃

Zn + 2 HNO₃ → Zn(NO₃)₂ + 2[H]

HNO₃ + [H] → H₂O + NO₂]2

Zn + 4 HNO₃ → Zn(NO₃)₂ + 2 H₂O + 2 NO₂

Action With Magnesium

Magnesium with highly diluted nitric acid (HNO₃)

Mg + highly dil. 2 HNO₃ → Mg(NO₃)₂ + H₂ ↑

Magnesium also reacts similarly to the aforementioned reactions.

The reaction of magnesium with moderately concentrated and concentrated nitric acid is zinc.

Action With Iron (Fe)

With Highly Concentrated HNO₃

When iron is subjected to concentrated nitric acid (HNO₃), a layer of ferrosoferric oxide (Fe₃O₄) is formed on the iron surface, creating an impermeable coating. Iron does not undergo any additional chemical reactions; instead, it assumes a passive state. Analogous patterns of behavior are observed in the cases of aluminum, chromium, cobalt, and nickel.

With moderately conc. HNO₃

Fe + 3 HNO₃ → Fe(NO₃)₃ + 3[H]

HNO₃ + [H] → NO + H₂O]3

Fe + 6 HNO₃ → Fe(NO₃)₂ + 3 NO + 3 H₂O

^{(Ferric nitrate)}

With diluted HNO₃

Fe + 2 HNO₃ → Fe(NO₃)₂ + 2[H]4

2 HNO₃ + 8[H] → N₂O + 5 H₂O

4 Fe + 10 HNO₃ → 4 Fe(NO₃)₂ + NO₂ + 5 H₂O

With highly diluted HNO₃

Fe + 2 HNO₃ → Fe(NO₃)₂ + 2[H]4

HNO₃ + 8[H] → NH₃ + 3 H₂O

HNO₃ + NH₃ → NH₄NO₃

4 Fe + 10 HNO₃ → 4 Fe(NO₃)₂ + NH₄NO₃ + 3 H₂O

Action With Tin (Sn)

With highly diluted HNO₃

Sn + 2 HNO₃ → Sn(NO₃)₂ + 2[H]4

HNO₃ + 8[H] → NH₃ + 3 H₂O

HNO₃ + NH₃ → NH₄NO₃

4 Sn + 10 HNO₃ → 4 Sn(NO₃)₂ + 3 H₂O + NH₄NO₃

With concentrated HNO₃

2 HNO₃ → H₂O + 2 NO₂ + [O]2

Sn + 2 [O] → SnO₂

SnO₂ + H₂O → H₂SnO₃

Sn + 4 HNO₃ → H₂SnO₃ + 4 NO₂ + H₂O

^{(Meta stannic acid)}

(b) Reaction With Metals Less Electropositive Than Hydrogen

Metals such as copper, silver, and mercury do not release any nascent hydrogen. However, it is certain that HNO₃ reacts by producing nascent oxygen and metallic oxides. These oxides then undergo further reaction with the acid, resulting in the formation of metal nitrates.

Action With Copper

With cold concentrated HNO₃

2 HNO₃ → 2 NO₂ + H₂O + [O]

Cu + [O] → CuO

CuO + 2 HNO₃ → Cu(NO₃)₂ + H₂O

Cu + 4 HNO₃ → Cu(NO₃)₂ + 2 H₂O + 2 NO₂

With acid of moderate strength (cold)

2 HNO₃ → 2 NO + H₂O + 3[O]

Cu + [O] → CuO]3

CuO + 2 HNO₃ → Cu(NO₃)₂ + H₂O]3

3 Cu + 8 HNO₃ → 3 Cu(NO₃)₂ + 4 H₂O + 2 NO

HNO₃ Vapor when passed over strongly heated copper gives nitrogen.

5 Cu + 2 HNO₃ → 5 CuO + H₂O + N₂ ↑

^{Heated Vapor}

(c) Action On Noble Metals

Noble metals, such as gold and platinum, exhibit a lack of reactivity when exposed to concentrated nitric acid (HNO₃). On the contrary, they exhibit solubility in aqua regia, a solution composed of three parts concentrated nitric acid (HNO₃).

3 HCl + HNO₃ → NOCl + 2 H₂O + 2[Cl]

^{Nitrosyl chloride Nascent chlorine}

The newly formed chlorine undergoes a chemical reaction with the metals as,

Pt + 4[Cl] → PtCl₄

Au + 3[Cl] → AuCl₃

The compounds PtCl₄ and AuCl₃ have the potential to undergo a reaction with an excess of HCl, resulting in the formation of chloroplatinic acid and chloroauric acid, respectively.

PtCl₄ + 2 HCl → H₂PtCl₆

AuCl₃ + HCl → HAuCl₄

The action of Nitric Acid on Non-metals

Nitric acid can oxidize non-metallic elements like sulfur, carbon, phosphorous, and iodine to make their own oxyacids. Concurrently, the process of reduction occurs in nitric acid, resulting in the generation of nitrogen dioxide (NO₂).

(i) Sulfur is Oxidized to Sulfuric Acid

2 HNO₃ → H₂O + 2 NO₂ + [O]]3

S + 3[O] → SO₃

SO₃ + H₂O → H₂SO₄

S + 6 HNO₃ → H₂SO₄ + 6 NO₂ + 2 H₂O

or, 2 HNO₃ → 2 NO₂ + H₂O + [O]]3

S + H₂O + 3[O] → H₂SO₄

S + 6 HNO₃ → H₂SO₄ + 6 NO₂ + 2 H₂O

(ii) Carbon is Oxidized to Carbon dioxide

2 HNO₃ → 2 NO₂ + H₂O + [O]]2

C + 2[O] → CO₂

C + 4 HNO₃ → CO₂ + 4 NO₂ + 2 H₂O

or, 2 HNO₃ → 2 NO₂ + H₂O + [O]]2

C + [O] → CO₂

CO₂ + H₂O → H₂CO₃

4 HNO₃ + C → 4 NO₂ + H₂O + H₂CO₃

(iii) Phosphorous is Oxidized to Phosphoric Acid

2 HNO₃ → 2 NO₂ + H₂O + [O]]5

2 P + 5[O] → P₂O₅

P₂O₅ + 3 H₂O → 2 H₃PO₄

2 P + 10 HNO₃ → 2 H₃PO₄ + 10 NO₂ + 2 H₂O

(iv) Iodine is Oxidized to Iodic Acid

2 HNO3 → H₂O + 2 NO₂ + [O]]5

I₂ + 5[O] → I₂O₅

I₂O₅ + H₂O → 2 HIO₃

I₂ + 10 HNO₃ → 2 HIO₃ + 10 NO₂ + 4 H₂O

^{Ionic acid}

The Action of Nitric Acid with Compounds

(i) Action of Nitric Acid with Sulfur dioxide (SO₂)

When concentrated nitric acid reacts with sulfur dioxide, it is oxidized to sulfuric acid.

conc. 2 HNO₃ → 2 NO₂ + H₂O + [O]

SO₂ + [O] → SO₃

SO₃ + H₂O → H₂SO₄

SO₂ + 2 HNO₃ → 2 NO₂ + H₂SO₄

dil. 2 HNO₃ → 2 NO + H₂O + 3[O]

SO₂ + [O] → SO₃]3

SO₂ + H₂O → H₂SO₄]3

3 SO₂ + 2 HNO₃ + 2 H₂O → 2 NO + 3 H₂SO₄

(ii) Action of Nitric Acid with Hydrogen Sulfide

When concentrated nitric acid reacts with hydrogen sulfide it oxidizes to sulfur.

conc. 2 HNO₃ → 2 NO₂ + H₂O + [O]

H₂S + [O] → H₂O + S

H₂S + 2 HNO₃ → 2 NO₂ + 2 H₂O + S

mod. dil. 2 HNO₃ → 2 NO + H₂O + 3[O]

H₂S + [O] → H₂O + S ]3

3 H₂S + 2 HNO₃ → 2 NO + 4 H₂O + 3 S

(iii) Action of Nitric Acid with Ferrous Sulfate

The oxidation of ferrous sulfate takes place in the presence of sulfuric acid, leading to the production of ferric sulfate.

conc. 2 HNO₃ → 2 NO₂ + H₂O + [O]

2 FeSO₄ + H₂SO₄ + [O] → Fe₂(SO₄)₃ + H₂O

2 FeSO4 + H₂SO4 + 2 HNO₃ → Fe₂(SO₄)₃ + 2 H₂O + 2 NO₂

(1:1) 2 HNO₃ → 2 NO + H₂O + 3[O]

2 FeSO₄ + H₂SO₄ + [O] → Fe₂(SO₄)₃ + H₂O]3

6 FeSO₄ + 3 H₂SO₄ + 2 HNO₃ → 3 Fe₂(SO₄)₃ + 2 NO + 4 H₂O

(iv) Action of Nitric Acid with Potassium Iodide

The reaction between concentrated nitric acid (HNO₃) and potassium iodide (KI) results in the oxidation of potassium iodide to iodine (I₂), while the nitric acid itself undergoes reduction to nitrogen dioxide (NO₂).

conc. 2 HNO₃ → 2 NO₂ + H₂O + [O]

2 KI + [O] → K₂O + I₂

K₂O + 2 HNO₃ → 2 KNO₃ + H₂O

2 KI + 4 HNO₃ → 2 KNO₃ + 2 H₂O + 2 NO₂ + I₂

Applications of Nitric Acid (HNO₃)

• Nitric acid has been employed in the production of various explosives, such as TNT (trinitrotoluene), dynamite, nitroglycerine, and guncotton, through its utilization in conjunction with concentrated sulfuric acid (H₂SO₄).
• Nitric acid holds significant importance as a reagent in laboratory settings.
• Nitric acid has been employed in the synthesis of dyes, pharmaceuticals, and fragrances.
• Additionally, it finds application in the refinement of precious metals such as silver and gold.

Video on Nitric Acid

References

• https://byjus.com/chemistry/nitric-acid/
• https://thechemco.com/chemical/nitric-acid/
• https://pubchem.ncbi.nlm.nih.gov/compound/Nitric-Acid#section=Structures
• https://www.dcceew.gov.au/environment/protection/npi/substances/fact-sheets/nitric-acid
• https://www.products.pcc.eu/en/blog/nitric-acid-v-characteristics-uses-and-hazards/
• https://www.nagwa.com/en/explainers/342152103047/