Unit 12: Unit 12: Nitrogen and Sulfur — Long Questions

Nitrogen
Introduction
• Nitrogen belongs to group 15 of the periodic table.
• It's atomic number is 7 amu and atomic mass 14 amu.
• It is a non-metal element belongs to p-block of periodic table.
Preparation of Nitrogen in laboratory
In laboratory settings, nitrogen can be generated by slowly heating a solution of ammonium nitrite.
NH₄NO₂(aq) → N₂(g) + 2H₂O(g)
Uses
• In industrial processes, nitrogen is typically obtained by cooling air until it becomes a liquid.
• Liquid nitrogen is commonly used for rapid cooling purposes.

Reactivity of Nitrogen (N₂)
Nitrogen is a significant component of the air, known for its low reactivity due to its small size, symmetrical electronic cloud, and nonpolar triple bond. With an electronic configuration of 1s² 2s² 2p³, nitrogen requires three electrons to complete its octet, forming a triple bond by sharing three electrons with another nitrogen atom as shown in Figure. This bond has a bond enthalpy of +944 kJ/mol⁻¹. High energy is required to break this bond to form new bonds, making N₂ very unreactive. The second reason for its lack of reactivity is the non-polarity of its bond. Both the atoms are the same having zero electronegativity difference. This causes equal sharing of the three bonded electrons between the two atoms making the bond nonpolar.

Ammonia (NH₃)
Industrial production of Ammonia: Ammonia (NH₃) is an important industrial compound of nitrogen, which is mainly used as a fertilizer. It is prepared industrially by Haber-Bosch process.
N₂(g) + 3H₂(g) ⇌ 2NH₃(g)

Basicity of Ammonia
Ammonia behaves as a Lowry-Bronsted base by accepting a proton (H⁺) from an acid to form ammonium ion:
NH₃(aq) + H⁺(aq) ⇌ NH₄⁺(aq)
It dissolves in water to form ammonium hydroxide (NH₄OH) and equilibrium is established between ammonia molecules and ammonium ions in the solution.
NH₃(aq) + H₂O(ℓ) ⇌ NH₄⁺(aq) + OH⁻(aq) Kₑ=1.8×10⁻⁵

Ammonia solution is a weak base due to the low basicity constant (K) and the equilibrium position being towards the far left side.

Structure of Ammonium (NH₄⁺) ion
• Ammonia molecule has pyramidal shape due to lone pair of nitrogen.
• When nitrogen atom in ammonia utilizes its lone pair of electrons to form ammonium, this ion adopts a tetrahedral shape.
• All the bonds are of equal length and strength in ammonium ion as depicted in figure.

Synthesis of Ammonia from Ammonium salts
Laboratory preparation: In the laboratory, ammonia gas can be synthesized by heating an ammonium salt such as ammonium chloride (NH₄Cl) with a base like calcium hydroxide Ca(OH)₂ as shown in figure.
2NH₄Cl(s) + Ca(OH)₂(aq) → CaCl₂(aq) + 2H₂O(ℓ) + 2NH₃(g)

In this acid-base reaction,
Acid: NH₄⁺ acts as an acid by donating H⁺ ions,
Base: OH⁻ acts as a base by accepting H⁺ ions.
This reaction displaces ammonia gas from the ammonium salt and produces salt and water.

Uses
• It is commonly used to identify ammonium ion in salt analysis.
• If a gas with a pungent smell is released and turns moistened red litmus paper blue, it indicates the presence of ammonium ion in the compound.

Oxides of Nitrogen • Oxides of nitrogen are NO, NO₂, N₂O, NO₂, N₂O₄ and N₂O₅ in which oxidation states range from I to V
• N₂O₄ and N₂O₅ decay quickly to other oxides.
• NO and NO₂ are collectively called as NOₓ.
• The structures, properties and uses of these oxides are given in Table 12.2.

Sources of oxides of nitrogen
The main categories of NOₓ sources.

(i) Natural Sources
Natural sources include lightning, volcanic eruption, forest fires, and denitrifying bacteria in soil. NO is produced when N₂ and O₂ in the air react during lightning. N₂O is produced by microorganism using air N₂.

(ii) Anthropogenic (Man-made) Sources
The main anthropogenic sources of NOₓ are the combustion of fossil fuels in vehicles and power plants. Other sources include chemical plants, biomass burning, welding etc.

Role of NO and NO₂ in smog and PAN formation
NOₓ is responsible for numerous harmful effects on living organisms.

Photochemical Smog
Formation: Photochemical smog (Los Angeles smog) forms in the atmosphere from NO₂ and volatile organic compounds (VOC's) in the sunlight.
• It is oxidizing in nature.
• Photochemical oxidants, such as NO₂, ozone, and peroxyacyl nitrates (PANs) can react and oxidize specific compounds in the atmosphere.
• Photochemical smog is becoming more common than classical smog (London smog) due to increasing NOₓ emissions.

Reactions The formation of photochemical smog involves the following chemical reactions.
N₂(g) + O₂(g) → 2NO(g)
2NO(g) + O₂(g) → 2NO₂(g)
NO₂(g) --hv--> NO(g) + O(g)
O(g) + O₂(g) → O₃(g)
RCH = RCH(g) + O₃(g) → RCO₃(g) + RCH'₂(g) → RC(O)OONO₂

Formation of Peroxyacyl Nitrates (PANs)
Starting material: NOₓ is starting material of PAN formation.
Process: NOₓ take part in a series of reactions leading to the formation of ozone (O₃), aldehydes, peroxyacyl nitrates (PANs) and peroxybenzoyl nitrate (PBN).
PAN is one of the members of peroxyacyl nitrates.

Steps of PAN production
• In peroxyacetyl nitrate (PAN) the alkyl (R−) group is methyl (CH₃−). Other hydrocarbon chains may also be present.
• The main component of oxidizing smog is ozone.
• Ozone oxidizes hydrocarbon to produce aldehyde.
• The aldehyde then reacts with hydroxyl radical to produce acyl radical.
• The acyl radical reacts with oxygen to produce peroxyacyl radical.
• Peroxyacyl radical finally reacts with nitrogen peroxide to form peroxyacyl nitrate.

CATALYTIC CONVERTER
Construction
• A catalytic converter is a ceramic or metallic monolith with a honeycomb-like structure.
• Its inner channels have a layer of alumina to provide a high surface area.
• Noble expensive metals such as Pt (Platinum), Pd (Palladium), and Rh (Rhodium) are dispersed on the alumina.

Working
These metals catalyze three redox reactions to remove the half harmful exhaust gases. The three-way converter converts harmful CO, NO, and hydrocarbons into CO₂, N₂, and water. These precious metals can also be recycled.

Reduction
2NO(g) + 2CO(g) --Pt/Rh--> N₂(g) + 2CO₂(g)

Oxidation
2CO(g) + O₂(g) --Pt/Pd--> 2CO₂(g)
2C₂H₄ + 6O₂(g) --Pt/Pd--> 4CO₂(g) + 4H₂O(g)

NITRIFICATION AND DENITRIFICATION
Nitrification and denitrification
Nitrogen cycle: Nitrification and denitrification are the two phases of the nitrogen cycle.

Nitrification It involves the conversion of ammonium (NH₄⁺) to nitrite (NO₂⁻) and nitrate (NO₃⁻)

Denitrification It involves the conversion of nitrate (NO₃⁻) to Nitrogen (N₂).
These two processes are also involved in the wastewater treatment to remove nitrogen. Some differences between nitrification and denitrification are given in the table 12.3.

(b) Catalytic converter works on the principle of harmful gases into less toxic gases by redox reaction. The metal catalysts catalyzes these redox reactions to remove the half harmful exhaust gases. The three-way converter converts harmful CO, NO and hydrocarbons into CO₂, N₂ and water.

(c) Hybrid cars have catalytic converters to reduce the emission of harmful gases in atmosphere. Electric cars do not need catalytic converters these cars do not emit harmful gases.

Sulphur

Sulphur is a member of group 16 which is also called the Chalcogen family. Some physical properties of sulphur are given in table 12.4.

Reactivity of sulphur

• Sulphur usually forms single bonds with other sulphur atoms instead of double bonds due to poor overlapping of the orbitals.
• Sulphur forms larger molecules and structures through a process called catenation.
• S₈ is a crown-like molecule

Oxidation states of sulphur

• Sulphur exhibits oxidation states of −2, 0, +2, +4, and +6.
• The oxidation state is determined by the number of unpaired electrons.

• Under standard conditions, sulphur and oxygen react to produce sulphur dioxide (SO₂) in which sulphur has an oxidation state of +4.
• In formation of sulphur trioxide (SO₃), sulphur with an oxidation state of +6. It requires high energy.

Electronic configurations of sulphur for attaining different oxidation states is shown in figure.

a) Oxygen atom in ground state (2s² 2p⁴) - There are no d-orbitals in 2ⁿᵈ shell therefore excitation is not possible.

b) Sulphur atom in ground state (3s² 3p⁴) - 2+ oxidation state due to two unpaired electrons

c) Sulphur atom in 1st excited state (3s² 3p³ 3d¹) - 4+ oxidation state due to four unpaired electrons

d) Sulphur atom in second excited state (3s¹ 3p³ 3d²) - 6+ oxidation state due to six unpaired electrons

STABILITY OF OXIDATION STATES OF SULPHUR

Sulphur displays a range of oxidation states, principally from -2 to +6.

Factor effects stability of oxidation states of sulphur

The stability of these states is influenced by the factors like pH, temperature, nature of the compound and the chemical environment.

• pH: In general, under acidic conditions, reduced forms (like H₂S, -2 oxidation state) are more stable, while under basic or neutral conditions, oxidized forms (like SO₄²⁻, +6 oxidation state) become more stable in water.

• Thermodynamics and kinetics: Although sulphur (+6), as in SO₃, is thermodynamically the most stable kinetic limitations can prevent it from forming readily at standard temperatures, making sulphur (+4), as in SO₂, the more frequent form.

• Nature of the compound: The stability of sulphur's oxidation state can be affected by the elements it is bonded to, and the overall chemical environment. In acidic environment SO₄²⁻ is kinetically stable due to strong O-S bonds.

• Catalyst: A catalyst can also enhance the rate of formation of a specific oxidation state, for example vanadium in the contract process increases the rate of formation of SO₃.

Reactions of Sulphur

Sulphur can combine with many elements to form a wide variety of inorganic and organic compounds.

Reactivity of sulphur

It is unreactive to water under normal conditions, dilute non-oxidising acids, and noble gases. Its ability to catenate allows it to form ring structures and linear chains.

Some important reactions of sulphur

(i) Burning of sulphur

Sulphur burns in the air to form SO₂ with a blue colour flame. The other main sulphur oxide is SO₃ which requires higher temperature and a catalyst for its formation.

S₍ₛ₎ + O₂₍ₘ₎ ⟶ SO₂₍ₘ₎

2SO₂₍ₘ₎ + O₂₍ₘ₎ →(Pt) 2SO₃₍ₘ₎

Sulphur can be oxidized by nitric acid to produce SO₂ and H₂SO₄

S₍ₛ₎ + 6HNO₃₍₁₎ ⟶ H₂SO₄₍₁₎ + 6NO₂₍ₘ₎ + 2H₂O₍₁₎

(ii) Sulphur as oxidizing agent

When sulphur reacts with elements that have lower electronegativity, it acts as an oxidizing agent and forms their sulfides. It tarnishes Ag, Cu, and Zn by forming a coating of metal sulfide.

2Ag₍ₛ₎ + S₍ₛ₎ ⟶ Ag₂S₍ₛ₎

Hg₍₁₎ + S₍ₛ₎ ⟶ HgS₍ₛ₎

2Cu₍ₛ₎ + S₍ₛ₎ ⟶ Cu₂S₍ₛ₎

Sulphur does not react with Au and Pt

(iii) Reaction with cyanate

Sulphur converts cyanide into thiocyanate which is also known as pseudohalide

Application of reaction

It is used to analyze Fe³⁺.

KCN₍ₛ₎ + S₍ₛ₎ →(Δ) KSCN₍ₛ₎

(iv) Reaction with fluorine

Sulphur reacts directly with F₂ to form SF₄ and SF₆. Sulphur hexafluoride (SF₆) is a gas and is very unreactive. It is used as an insulator gas in electric devices. Sulphur readily reacts with Cl₂ to form S₂Cl₂ (yellow liquid) which further reacts with Cl₂ to form SCl₂ (red liquid).

S₍ₛ₎ + 3F₂₍ₘ₎ →(Δ) SF₆₍ₘ₎

2S₍ₛ₎ + Cl₂₍ₘ₎ →(Δ) S₂Cl₂₍₁₎

Cl₂₍ₘ₎ + S₂Cl₂₍₁₎ ⟶ 2SCl₂₍₁₎

Uses of sulphur and its compounds

Vulcanization Sulphur is used as a cross linker for the rubber molecular chains. This is called vulcanization it improves the strength of rubber.

Fertilizer

Sulphur is an essential nutrient for plant growth. When soils become depleted in sulfate, sulphur can be restored in soil by applying sulphur containing N/P fertilizers, or sulfur-coated fertilizers such as sulfur-coated urea. Soil components and microbes convert elemental sulphur into soluble forms for the use of plants.

Gypsum (CaSO₄.5H₂O) is also used as a fertilizer.

Gun powder

Gun powder is a coarse blend of 75% potassium nitrate (KNO₃), 15% wood charcoal, and 10% sulphur. Charcoal carbon is the main fuel, nitrate is the oxidizer and sulphur is the additional fuel that burns the powder faster. The following reaction in burning takes place:

10KNO₃₍ₛ₎ + 3S₍ₛ₎ + 8C₍ₛ₎ →(Δ) 2K₂CO₃₍ₛ₎ + 3K₂SO₄₍ₛ₎ + 6CO₂₍ₘ₎ + 5N₂₍ₘ₎

Role of Sulphur in organic synthesis

Carbon-Sulphur bonds are prevalent in a wide range of compounds with biological, pharmaceutical, and material properties. These bonds form a large number of organic compounds containing a variety of functional groups such as thiols or mercaptans, thioeters, sulfoxides, sulfones, etc.

Drugs

Sulfa drugs are the antibacterial sulfonamides such as penicillins and cephalosporins contain sulphur. The common drug omeprazole, used in GERD (Gastroesophageal reflux disease) contains sulfoxide group

Dyes

Sulfur dyes are synthesized by the process of thionation or sulfurization of organic compounds that contain nitro or amino groups. These compounds contain sulphur linkages. They generally give black, brown, blue, and green colours.

Some examples of sulfur colors are:

• Black
• Blue
• Brilliant green

Odorants/Fragrances

Mercaptans are used to give odour to natural gas. Some thiols have pleasant odours on high dilution.

Example Thioterpineol is the key ingredient in the aroma of grapefruit. cis-glabanum oxathiane is a fragrant compound. It is used in fine fragrances, soaps, shampoos and shower gels. Many naturally occurring odorants are produced synthetically and also applied as flavouring agents.

Structure of sulfuric acid

It has tetrahedral structure with two S−O and two S=O bonds

Contact process

The major portion of sulfur, around 85% is used for the production of sulphuric acid (H₂SO₄). Contact process for the industrial production of sulphuric acid. The Contact process can be divided into the following stages.

Sulfur/pyrite Burners

The process starts with the combustion of molten sulphur or by heating pyrites such as iron pyrite (FeS) in excess of air to produce sulphur dioxide SO₂.

2S₍ₛ₎ + O₂₍ₘ₎ ⟶ 2SO₂₍ₘ₎ ΔH = −297.03 kJ/mol

4FeS₍ₛ₎ + 6O₂₍ₘ₎ ⟶ 2Fe₂O₃₍ₛ₎ + 4SO₂₍ₘ₎

Purification Unit

If pyrite ore is used as a sulfur source, the SO₂ gas formed may contain contaminants like dust particles, vapours, and arsenic oxide. These contaminants affect the efficiency of the catalyst. Hence, the gas needs to pass through the purification unit.

In an arsenic purifier, gelatinous ferric hydroxide Fe(OH)₃ present in horizontal shelves, absorbs arsenic oxide As₂O₃.

As₂O₃₍ₛ₎ + 2Fe(OH)₃₍₁₎ ⟶ 2FeAsO₍₁₎ + 3H₂O₍₁₎

Contact Tower and Heat Exchangers

Purified SO₂ and air, preheated at 420°C−450°C, are fed to the first converter stage of the contact tower at 1−2 atm pressure. Here, these gases come in contact with vanadium pentoxide (V₂O₅) catalyst.

SO₂₍ₘ₎ + ½O₂₍ₘ₎ ⇌ SO₃₍ₘ₎ ΔH = −98.98 kJ mol⁻¹

The catalyst works in two steps

Oxidation of SO₂ (into SO₃ by V⁵⁺)

SO₂₍ₘ₎ + V₂O₅₍ₛ₎ ⟶ SO₃₍ₘ₎ + V₂O₄₍ₛ₎

Oxidation of V⁴⁺ back into V⁵⁺ by oxygen (catalyst regeneration)

V₂O₄₍ₛ₎ + ½O₂₍ₘ₎ ⟶ V₂O₅₍ₛ₎

Physical properties of Sulphuric acid

• Solubility Sulphuric acid is soluble in water and hygroscopic in nature.
• Absorbs water vapours: It readily absorbs water vapour from the air.
• Polarity: Anhydrous H₂SO₄ is a very polar liquid.
• Corrosive: It is highly corrosive to various materials. On contacting the skin, it causes chemical burns.

Chemical properties

(i) Self Ionization

It self-ionizes or undergoes autoprotonolysis as follows.

H₂SO₄₍₁₎ + H₂SO₄₍₁₎ ⟶ H₃SO₄⁺₍ₐᵩ₎ + HSO₄⁻₍ₐᵩ₎ Ka = 2.7 × 10⁻⁴

The equilibrium constant value is greater than that of water which makes it to be used as a non-aqueous protic solvent.

Sulphuric acid is a strong acid as shown by its pKₐ value:

H₂O₍₁₎ + H₂SO₍ₐᵩ₎ ⇌ H₃O⁺₍ₐᵩ₎ + HSO₄⁻₍ₐᵩ₎ (pKₐ₁ = −2)

But hydrogen sulfate (HSO₄⁻) is a far weaker acid due to a positive pKₐ₂ value:

HSO₄⁻₍ₐᵩ₎ + H₂O₍₁₎ ⇌ H₃O⁺₍ₐᵩ₎ + SO₄²⁻₍ₐᵩ₎ (pKₐ₂ = 1.92)

(ii) Dehydrating agent

Concentrated sulphuric acid is a powerful dehydrating agent that removes water from many substances such as sucrose, starch, wood, and paper to produce carbon, steam, and heat.

Dehydration of table sugar

C₁₂H₂₂O₁₁₍ₛ₎ + H₂SO₄₍₁₎ ⟶ 12C₍ₛ₎ + 11H₂O₍ₘ₎ + H₂SO₄₍ₐᵩ₎

A black porous carbon mass called carbon snake protrudes out of the apparatus

Dehydration of alcohol

It also dehydrates ethyl alcohol to ethene or ethoxyethane depending upon the reaction conditions

C₂H₅OH₍₁₎ + H₂SO₄₍₁₎ ⟶ C₂H₄₍ₘ₎ + H₂O₍₁₎ + H₂SO₄₍ₐᵩ₎

(iii) Reaction with sodium chloride

Hydrochloric acid (HCl) gas, is formed when sulphuric acid reacts with sodium chloride

NaCl₍ₛ₎ + H₂SO₄₍₁₎ ⟶ NaHSO₄₍ₛ₎ + HCl₍ₘ₎

(iv) Reaction with metals

Reactions of sulphuric acid with metals depend upon the metal, concentration of the acid, and temperature.

Metals that are above hydrogen in electrochemical series such as Fe, Al, Zn, Mn, Ni, and Mg react directly with dilute sulphuric acid to produce hydrogen gas and metal sulfates.

Zn₍ₛ₎ + H₂SO₄₍ₐᵩ₎ ⟶ H₂₍ₘ₎ + Zn²⁺₍ₐᵩ₎ + SO₄²⁻₍ₐᵩ₎

But with cold conc. H₂SO₄, they liberate SO₂ and form sulfates.

Zn₍ₛ₎ + 2H₂SO₄₍₁₎ ⟶ SO₂₍ₘ₎ + Zn²⁺₍ₐᵩ₎ + SO₄²⁻₍ₐᵩ₎ + 2H₂O₍₁₎

Metals like Cu, Ag, and Hg react with hot conc. H₂SO₄ to form metal sulfates

(v) Oxidizing power

• Sulphuric acid is not regarded as a typical oxidizing agent due to the stability of SO₄²⁻ anion. This anion is weakly oxidizing.

• Hot concentrated sulphuric acid is a moderately strong oxidizing agent due to high temperature, high concentration of protons (H⁺), and formation of nascent oxygen

• Hot concentrated sulphuric acid oxidizes Cu, as given below:

Cu₍ₛ₎ + 2H₂SO₄₍₁₎ ⟶ SO₂₍ₘ₎ + 2H₂O₍₁₎ + SO₄²⁻₍ₐᵩ₎ + Cu²⁺₍ₐᵩ₎

Uses and Industrial Applications

Sulphuric acid is considered a king of chemicals and its consumption is an indicator of the industrial progress of a country.

• Fertilizers: A major portion of the acid is used in making fertilizers, normally 77.67 % is used to digest the phosphate rock containing calcium phosphate Ca₃(PO₄)₂. Sulphuric acid also reacts with ammonia to make ammonium sulfate fertilizer

• Extraction of metals: It is used in the extraction of metals from ores such as Cu, Ni, steel etc.

• As a catalyst: It is utilized as a catalyst in oil and coal refining, polymers, synthetic rubber, and plastic industries.

• Industrial usage: It is used in the pulp and paper industry and involved in the production of pesticides, insecticides, herbicides, varnishes, dyes, pharmaceuticals, soaps and detergents.

• Explosives: It is used for nitration in making explosives such as trinitrotoluene (TNT), nitroglycerin, picric acid, nitrocellulose, etc.

• Food industry: It is involved in the food industry for making sugar, starch, and corn syrup.

• Paint industry: It is used in the paint industry for making titanium dioxide (TiO₂) pigment.

• Drying agent: It is used to dry gases in industrial processes.

• Batteries: 35.67% acid is used in lead storage batteries.

• Reagent: It is used as a laboratory reagent.

• Even though it is used in various industries, it is rarely contained in the final product.