Nitrogen and Sulfur

Nitrogen belongs to group 15 of the periodic table. In industrial processes it is typically obtained by cooling air until it becomes a liquid (liquid air is then separated by fractional distillation — nitrogen, boiling point −195.8 °C, comes off first). Liquid nitrogen is commonly used for rapid cooling purposes. In the laboratory, nitrogen can be generated by slowly heating a solution of ammonium nitrite:

NH₄NO_2(aq) —Δ→ N_2(g) + 2H₂O_(l)

Physical properties (Table 12.1): atomic number 7; relative atomic mass 14.007 u; colourless gas; electronic configuration [He] 2s² 2p³; melting point −210 °C; boiling point −195.8 °C; covalent radius 74 pm; ionic radius (3−) 171 pm; 1^st ionization energy 1402 kJ/mol; electronegativity 3.0; electron affinity −7.0 kJ/mol; density 1.25 kg/m³; principal oxidation states 3+ and 5+ (as listed in the book's table).

12.1 Reactivity of nitrogen (N₂). Nitrogen is a significant component of the air, known for its low reactivity for two reasons. First, the triple bond: with the configuration 1s² 2s² 2p³, a nitrogen atom requires three electrons to complete its octet, so it shares three electrons with another nitrogen atom, forming a triple bond (N≡N) with a bond enthalpy of +944 kJ/mol. High energy is required to break this bond to form new bonds, making N₂ very unreactive. Second, the non-polarity of its bond: both atoms are the same, having zero electronegativity difference, so the three bonded electrons are shared equally and the bond is nonpolar.

📌 Interesting Information (book's side note): the high concentration of nitrogen in the air serves to dilute oxygen, preventing every spark in our atmosphere from igniting a massive fire. In the case of large shipments of hydrocarbons or edible oils, it is crucial to utilize blankets of nitrogen or any other inert gas on ships to safeguard them from oxygen and moisture. It is also used in the laboratory to carry out reactions which require an inert atmosphere.

Ammonia (NH₃) is an important industrial compound of nitrogen, mainly used as a fertilizer. It is prepared industrially by the Haber-Bosch process:

N_2(g) + 3H_2(g) ⇌ 2NH_3(g)

12.2.1 Basicity of ammonia. Ammonia behaves as a Lowry-Brønsted base by accepting a proton (H⁺) from an acid to form ammonium:

NH_3(aq) + H⁺_(aq) ⇌ NH₄⁺_(aq)

It dissolves in water to form ammonium hydroxide (NH₄OH) and an equilibrium is established between ammonia molecules and ammonium ions in the solution:

NH_3(aq) + H₂O_(l) ⇌ NH₄⁺_(aq) + OH⁻_(aq)   K_b = 1.8 × 10⁻⁵

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

12.2.2 Structure of ammonium (NH₄⁺). The ammonia molecule has a pyramidal shape due to the lone pair of nitrogen. But when the nitrogen atom in ammonia utilizes its lone pair of electrons to form ammonium, this ion adopts a tetrahedral shape in which all the bonds are of equal length and strength.

12.2.3 Synthesis of ammonia from ammonium salts. 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):

2NH₄Cl_(s) + Ca(OH)_2(aq) → CaCl_2(aq) + 2H₂O_(l) + 2NH_3(g)

In this acid-base reaction, NH₄⁺ acts as an acid by donating H⁺ ions, while OH⁻ acts as a base by accepting H⁺ ions. This reaction displaces ammonia gas from the ammonium salt and produces salt and water. It is commonly used to identify ammonium 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 in the compound.

Quick Check 12.1 (book): (a) N₂ gas is used in food packaging because it is inert — it does not react with the food and it displaces oxygen, so oxidation/spoilage is slowed. (b) Both CO and N₂ have triple bonds, but CO is more reactive than N₂ because CO is a polar molecule (carbon and oxygen differ in electronegativity, and the bond is polarized), whereas N₂ is symmetrical and nonpolar. (c) Ammonium salts such as (NH₄)₂SO₄ or (NH₄)NO₃ are commonly used as fertilizers — but a farmer would not treat a field with an ammonium fertilizer at the same time as using lime, because lime is a base (Ca(OH)₂/CaO) and, by the reaction above, it displaces ammonia gas from the ammonium salt: the nitrogen escapes as NH₃ and the fertilizer is wasted.

Oxides of nitrogen are NO, 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 NO_x. Their structures, properties and uses are given in Table 12.2 of the book:

• Nitrous oxide / Nitrogen oxide (laughing gas), N₂O — formal oxidation state 1+. Structure :N≡N–Ö: . Colourless gas, water-soluble, neutral, sweet smelling, helps in combustion. Uses: dental anaesthetic, propellant for whipped ice cream, synthesis of NaN₃.
• Nitric oxide, NO (the book's table also lists N₂O₂) — formal oxidation state 2+. Structure :N=Ö with an odd electron. Colourless gas, slightly water-soluble, neutral, paramagnetic, while N₂O₂ is diamagnetic; oxidising as well as reducing in nature. Uses: biochemical messenger (lowers blood pressure and role in other body functions), synthesis of nitrosyl carbonyls.
• Nitrogen dioxide / Nitrogen peroxide / Nitrogen tetraoxide, NO₂, N₂O₄ — formal oxidation state 4+. NO₂ is a reddish-brown gas, paramagnetic, and reacts with water to form HNO₃ and HNO₂; N₂O₄ is a colourless liquid or solid. The two interconvert: 2NO₂ (brown) ⇌ N₂O₄ (colourless), cooling favouring N₂O₄ and heat favouring NO₂. Uses: rocket propellant, formation of HNO₃ by the Ostwald process, explosives.

📌 Did you know? (book's side note): the recommended name for ammonia by IUPAC is Azane and for ammonium is Azanium. These names have been derived from Azote, a Greek name for nitrogen, meaning "no life". These names are used in naming derivatives of ammonia and ammonium, e.g. sodium azide (NaN₃).

12.4 Sources of oxides of nitrogen. Here are the main categories of NO_x sources.

12.4.1 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 microorganisms using air N₂.

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

12.5 Role of NO & NO₂ in smog & PAN formation. NO_x is responsible for numerous harmful effects on living organisms.

12.5.1 Photochemical smog. Photochemical smog (Los Angeles smog) forms in the atmosphere from NO_x and volatile organic compounds (VOCs) in the sunlight. It is oxidising in nature. It contains photochemical oxidants, such as NO₂, ozone, and peroxyacyl nitrates (PANs). Photochemical smog is becoming more common than classical smog (London smog) due to increasing NO_x emissions. The formation of photochemical smog involves the following chemical reactions:

• N_2(g) + O_2(g) → 2NO_(g)
• 2NO_(g) + O_2(g) → 2NO_2(g)
• NO_2(g) —hν→ NO_(g) + O_(g)
• O_(g) + O_2(g) → O_3(g)
• RCH=RCH_(g) + O_3(g) → RCO•_(g) + RCH•_(g) →→ RC(O)OONO₂ (PAN)

📌 Did you know? (book's side note): Lahore smog consists of volatile organic compounds (VOCs), NO_x, ground level ozone (O₃), particulate matter PM2.5, CO and SO₂.

NO_x takes 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. The general structure of a PAN is:

R–C(=O)–O–O–N⁺(=O)–O⁻   i.e. RC(O)OONO₂

The R group in peroxyacetyl nitrate is CH₃, but other hydrocarbon chains may also be present. So peroxyacetyl nitrate (PAN) itself is CH₃C(O)OONO₂.

The formation of PAN is illustrated in Figure 12.5 of the book. The main component of oxidizing smog is ozone, which 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, which finally reacts with nitrogen peroxide to form peroxyacyl nitrate. In short:

Hydrocarbon —O₃→ aldehyde (RCHO) —•OH→ acyl radical (RC•=O) —O₂→ peroxyacyl radical (RC(O)OO•) —NO₂→ PAN (RC(O)OONO₂)

Quick Check 12.2 (book): (a) Structures of the oxides: N₂O is :N≡N–Ö: (a triple bond between the two nitrogens and a single bond to oxygen, the molecule being linear); NO is :N=Ö: with one unpaired electron (which is why NO is paramagnetic and reactive, both oxidising and reducing); NO₂ is an angular/bent molecule, O–N–O, again with an unpaired electron, which is why it dimerizes to N₂O₄. (b) PAN stands for peroxyacetyl nitrate; its general formula is RC(O)OONO₂ (for PAN itself, R = CH₃). (c) The formulas of the compounds responsible for PAN formation are the hydrocarbon (RCH=RCH), ozone (O₃), the aldehyde (RCHO), the acyl radical, the peroxyacyl radical and nitrogen dioxide (NO₂), as in the chain above. (d) If magnesium ribbon is ignited and placed in a jar containing N₂O, it continues to burn brightly — because N₂O supplies oxygen (Mg + N₂O → MgO + N₂). This confirms the structure :N≡N–O: in which oxygen is present as a terminal atom that can be released to support combustion, exactly as the book states for N₂O ("helps in combustion").

A catalytic converter is a ceramic or metallic material 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.

These metals catalyze three redox reactions to remove the harmful exhaust gases. The three-way converter converts harmful CO, NO_x and hydrocarbons into CO₂, N₂ and water. These precious metals can also be recycled.

Reduction (NO_x reduced to nitrogen):

2NO_(g) + 2CO_(g) —Pt/Rh→ N_2(g) + 2CO_2(g)

Oxidation (CO and hydrocarbons oxidized):

2CO_(g) + O_2(g) —Pt/Pd→ 2CO_2(g)

2C₂H_4(g) + 6O_2(g) —Pt/Pd→ 4CO_2(g) + 4H₂O_(g)

Quick Check 12.3 (book): (a) The reduction and oxidation reactions in the vehicle exhaust are exactly the three equations above: NO + CO reduced to N₂ + CO₂ over Pt/Rh, and CO and hydrocarbons oxidized to CO₂ + H₂O over Pt/Pd. (b) The basic principle of a catalytic converter is that a catalyst (Pt, Pd, Rh on a high-surface-area alumina washcoat) speeds up redox reactions without itself being consumed, so harmful gases are converted into harmless N₂, CO₂ and water at exhaust temperatures instead of being released unchanged; the honeycomb gives a huge surface for the gases to meet the catalyst. (c) Hybrid cars have a petrol engine as well as an electric motor, so they produce exhaust and do have catalytic converters; a purely electric car has no exhaust at all, so it has nothing to convert and needs no catalytic converter. (This follows directly from the converter's purpose — cleaning exhaust gases — stated in the chapter.)

Nitrification and denitrification are the two phases of the nitrogen cycle. Nitrification involves the conversion of ammonium (NH₄⁺) to nitrite (NO₂⁻) and nitrate (NO₃⁻), while denitrification involves the conversion of nitrate (NO₃⁻) to nitrogen (N₂). Some differences between nitrification and denitrification are given in Table 12.3 of the book:

• Nitrification: ammonia NH₃/NH₄⁺ is converted into nitrite (NO₂⁻) and nitrate (NO₃⁻). Nitrifying bacteria, aerobic conditions, pH 6.5–8.0, optimum temperature 20 °C–30 °C. Plants absorb these nitrites and nitrates for their nutrition as they cannot assimilate nitrogen directly from the atmosphere.
• Denitrification: nitrite (NO₂⁻) and nitrate (NO₃⁻) are converted back to N₂ that is released into the atmosphere. Denitrifying bacteria, anaerobic conditions, pH 7.0–9.0, optimum temperature 26 °C–38 °C. It is important for wastewater treatment and useful for aquatic life supported with Anammox that oxidizes NH₄⁺ with NO₂⁻ to form N₂ gas.

Oxidation of nitrogen (nitrification):

NH₄⁺ → NO₂⁻ → NO₃⁻

2NH₄⁺_(aq) + 3O_2(g) → 2NO₂⁻_(aq) + 4H⁺_(aq) + 2H₂O_(l)

2NO₂⁻_(aq) + O_2(aq) → 2NO₃⁻_(aq)

Reduction of nitrogen (denitrification):

NO₃⁻ → NO₂⁻ → NO → N₂O → N₂

2NO₃⁻_(aq) + 10e⁻ + 12H⁺_(aq) → N_2(g) + 6H₂O_(l)

The key memory hook: nitrification is oxidation (nitrogen's oxidation state rises, −3 in NH₄⁺ → +3 in NO₂⁻ → +5 in NO₃⁻) and needs oxygen; denitrification is reduction (oxidation state falls back to 0 in N₂) and happens without oxygen.

Sulfur is a member of group 16 which is also called the Chalcogen family. Some physical properties of sulfur are listed in Table 12.4 of the book: atomic number 16; relative atomic mass 32.06 a.m.u.; physical appearance solid yellow; electronic configuration [Ne] 3s² 3p⁴; melting point 113 °C (honey-yellow); boiling point 445 °C (dark brown); covalent radius 104 pm; ionic radius 184 pm; 1^st ionization energy 1000 kJ/mol; electronegativity 2.5; electron affinity −200 kJ/mol; density 2.07 g/cm³; common oxidation states +4 and +6; common bonding: covalent bond.

12.8.1 Reactivity of sulfur. Sulfur usually forms single bonds with other sulfur atoms instead of double bonds due to poor overlapping of the orbitals. As a result, it forms larger molecules and structures through a process called catenation. S₈ is a crown-like molecule, as shown in Figure 12.7 of the book. This stable ring structure is one reason sulfur is so unreactive at room temperature (see Section 12.10).

12.8.2 Oxidation states of sulfur. Sulfur exhibits oxidation states of −2, 0, +2, +4, and +6. The electronic configuration of sulfur is depicted in Figure 12.8 of the book, where the oxidation state is determined by the number of unpaired electrons:

• Oxygen atom in ground state (2s² 2p⁴): there are no d-orbitals in the 2^nd shell, therefore excitation is not possible — oxygen cannot expand its octet like sulfur.
• Sulfur atom in ground state (3s² 3p⁴): 2 unpaired electrons → 2+ oxidation state (in the book's counting).
• Sulfur atom in 1^st excited state (3s² 3p³ 3d¹): 4 unpaired electrons → 4+ oxidation state.
• Sulfur atom in 2^nd excited state (3s¹ 3p³ 3d²): 6 unpaired electrons → 6+ oxidation state.

Under standard conditions, sulfur and oxygen react to produce sulfur dioxide (SO₂) in which sulfur has an oxidation state of +4. However, to form sulfur trioxide (SO₃) with an oxidation state of +6, high energy is required — which is why SO₂ forms easily but SO₃ needs a catalyst and high temperature (see the contact process, Section 12.12).

Sulfur displays a range of oxidation states, principally from −2 to +6. The stability of these states is influenced by 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 sulfur (+6), as in SO₃, is thermodynamically the most stable, kinetic limitations can prevent it from forming readily at standard temperatures, making sulfur (+4), as in SO₂, the more frequent form.

Nature of the compound. The stability of sulfur'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 contact process increases the rate of formation of SO₃.

Exam shortcut: acidic → reduced form wins (H₂S, −2); basic → oxidized form wins (SO₄²⁻, +6); and "most stable in theory" is not always "most common in practice" — kinetics decides what you actually get, which is why SO₂ (+4) is the everyday form and SO₃ (+6) needs a catalyst.

Sulfur can combine with many elements to form a wide variety of inorganic and organic compounds. 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. Here are some important reactions of sulfur:

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

S_(s) + O_2(g) → SO_2(g)

2SO_2(g) + O_2(g) —Δ, Pt→ 2SO_3(g)

Sulfur can be oxidised by nitric acid to produce NO₂ and H₂SO₄:

S_(s) + 6HNO_3(aq) → H₂SO_4(aq) + 6NO_2(g) + 2H₂O_(l)

ii. With less electronegative elements — sulfur as oxidising agent. When sulfur 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_(l) + S_(s) → HgS_(s)  •  2Cu_(s) + S_(s) → Cu₂S_(s)

iii. With cyanide. Sulfur converts cyanide into thiocyanate which is also known as pseudohalide:

KCN_(s) + S_(s) —Δ→ KSCN_(s)

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

S_(s) + 3F_2(g) → SF_6(g)  •  2S_(l) + Cl_2(g) → S₂Cl_2(l)  •  S₂Cl₂ + Cl₂ → 2SCl₂

Quick Check 12.4 (book): (a) Oxidation states of S: in SO₃, +6 (3 × −2 = −6, so S = +6); in H₂SO₄, +6 (2 × +1 + S + 4 × −2 = 0, so S = +6); in SO₃²⁻, +4 (S + 3 × −2 = −2, so S = +4); in S₂O₃²⁻, +2 on average (2S + 3 × −2 = −2, so 2S = +4). (b) The most common oxidation states are +4 and +6 — exactly the "common oxidation states" in Table 12.4 — because SO₂ (+4) forms readily from sulfur and oxygen, and +6 is the thermodynamically most stable state (SO₃, SO₄²⁻). (c) Variable oxidation states come from sulfur's 3d orbitals: unlike oxygen (no d-orbitals in shell 2), sulfur can promote electrons into vacant 3d orbitals, giving 2, 4 or 6 unpaired electrons and hence the +2, +4 and +6 states (Section 12.8.2).

Quick Check 12.5 (book): (a) SO₂ and sulfite (SO₃²⁻) salts preserve food because they act as antimicrobial and antioxidant preservatives — they check the growth of microbes and slow oxidation that spoils food. ⚠️ The chapter asks this question but gives no dedicated passage on food preservation; this answer states the standard function of SO₂/sulfite as a preservative and is flagged as such. (b) Sulfur with HCl: sulfur is unreactive towards dilute non-oxidising acids (Section 12.10), so no reaction with HCl. Sulfur with NaOH: ⚠️ the chapter gives no NaOH reaction for sulfur, so no book equation is claimed here — the standard school reaction is disproportionation in hot concentrated alkali, 3S + 6NaOH → 2Na₂S + Na₂SO₃ + 3H₂O, and it is flagged as outside this chapter's text.

12.11 Uses of sulfur and its compounds.

12.11.1 Vulcanisation. Sulfur is used as a cross linker for the rubber molecular chains through di-sulfide (S–S) bridges. This is called vulcanization of rubber. It improves the strength of rubber.

12.11.2 Fertilizer. Sulfur is an essential nutrient for plant growth. When soils become depleted in sulfate, sulfur can be restored in soil by ammonium sulfate, elemental sulfur and sulfur-coated fertilizers such as sulfur-coated urea. Soil components and microbes convert elemental sulfur into soluble forms for the use of plants. Gypsum is also used as a fertilizer. ⚠️ The book prints gypsum as CaSO₄·5H₂O; standard gypsum is the dihydrate, CaSO₄·2H₂O — the book's "5" is flagged here as an apparent slip, and the standard formula is the one to use in exams.

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

10KNO_3(s) + 3S_(s) + 8C_(s) → 2K₂CO_3(s) + 3K₂SO_4(s) + 6CO_2(g) + 5N_2(g)

12.12 Role of sulfur in organic synthesis. Carbon–sulfur 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, thioethers, sulfones, etc.

12.12.1 Drugs. Sulfa drugs are the antibacterial sulfonamides, such as penicillins and cephalosporins contain sulfur. ⚠️ The book's sentence here is compressed and confusing — it runs "the common drug β-lactam such as omeprazole" together. Read separately, as the chemistry requires: penicillins and cephalosporins are β-lactam antibacterials that contain sulfur, while omeprazole, used in GERD (Gastroesophageal reflux disease), contains a sulfoxide group. Omeprazole is not itself a β-lactam.

12.12.2 Dyes. Sulfur dyes are synthesized by the process of thionation or sulfurization of organic compounds that contain nitro or amino groups. These compounds contain sulfur linkages. They generally give black, brown, khaki, blue, and green colours. Some examples are sulfur black 1, sulfur blue and sulfur brilliant green.

12.12.3 Odorants / Fragrances. Mercaptans are used to give odour to natural gas. Some thiols have pleasant odours on high dilution. For example, thioterpineol is the key ingredient in the aroma of grapefruit. cis-Galbanum 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.

The major portion of sulfur, around 85%, is used for the production of sulfuric acid (H₂SO₄). It has a tetrahedral structure with two S–O and two S=O bonds. Bond lengths from Figure 12.10 of the book: S=O 142.2 pm, S–O 157.4 pm, O–H 97 pm.

12.13.1 Contact process. Sulfuric acid is produced by the contact process. The process can be divided into the following stages.

Sulfur / pyrite burners. The process starts with the combustion of molten sulfur or by heating pyrites such as iron pyrite in excess of air to produce sulfur dioxide SO₂.

S_(s) + O_2(g) → SO_2(g)   ΔH = −297.03 kJ/mol

4FeS_(s) + 6O_2(g) → 2Fe₂O_3(s) + 4SO_2(g)   (the book writes the pyrite as FeS in this equation)

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.

As₂O_3(g) + 2Fe(OH)_3(s) → 2FeAsO_3(s) + 3H₂O_(l)

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.

SO_2(g) + ½O_2(g) ⇌ SO_3(g)   ΔH = −98.98 kJ/mol

V₂O₅ as catalyst works in the following two steps:

Oxidation of SO₂ into SO₃ by V⁵⁺:   SO_2(g) + V₂O_5(s) → SO_3(g) + V₂O_4(s)

Oxidation of V⁴⁺ back into V⁵⁺ by oxygen (catalyst regeneration):   V₂O_4(s) + ½O_2(g) → V₂O_5(s)

Absorption tower. Sulfur trioxide is cooled and can be converted to sulfuric acid by reacting with water.

H₂O_(l) + SO_3(g) → H₂SO_4(l)   ΔH = −176.6 kJ/mol

Mixing SO₃ with water is not feasible because the reaction is extremely exothermic and acidic vapour or mist is produced rather than a liquid solution. Mainly, sulfur trioxide is dissolved in recirculating hot 98.5% sulfuric acid to form oleum (fuming sulfuric acid). The term fuming sulfuric acid or oleum is used for the mixture of sulfur trioxide with 100 percent sulfuric acid to form oleum. Oleum undergoes a reaction with water to make a highly concentrated solution of H₂SO₄ whose concentration can be adjusted.

H₂SO_4(l) + SO_3(g) → H₂S₂O_7(l)  •  H₂S₂O_7(l) + H₂O_(l) → 2H₂SO_4(aq)

Quick Check 12.6 (book): (a) High temperature and a catalyst are needed to form SO₃ because the oxidation of SO₂ is slow at ordinary temperatures even though it is exothermic (ΔH = −98.98 kJ/mol) — the V₂O₅ catalyst provides the two-step V⁵⁺/V⁴⁺ path above, and 420–450 °C gives a practical rate without pushing the exothermic equilibrium too far back. (b) Dehydration reactions of conc. H₂SO₄ with starch and oxalic acid: ⚠️ the chapter asks this in Quick Check 12.6 but gives no starch or oxalic-acid equation in its text, so no book equation is claimed. The pattern follows the book's stated dehydrating action (it removes water from starch, wood and paper to produce carbon, steam and heat): starch chars to carbon as water is removed, and oxalic acid is dehydrated by conc. H₂SO₄ to carbon monoxide and carbon dioxide (H₂C₂O₄ → CO + CO₂ + H₂O), the standard school equation, flagged as outside this chapter's text. (c) How V₂O₅ works: by the two-step cycle above — V⁵⁺ oxidizes SO₂ to SO₃ and is itself reduced to V⁴⁺, then oxygen re-oxidizes V⁴⁺ back to V⁵⁺, so the catalyst is regenerated and not consumed.

12.13.2 Physical properties. Sulfuric acid is soluble in water and hygroscopic in nature. It readily absorbs water vapour from the air. Anhydrous H₂SO₄ is a very polar liquid. It is highly corrosive to various materials. On contacting the skin, it causes chemical burns. Some physical properties are given in Table 12.5 of the book: molar mass 98.08 g/mol; physical appearance colourless viscous liquid; odour odourless; melting point 10 °C; boiling point 290 °C; specific gravity (15 °C) 1.83 g/cm³; viscosity 25.24 centipoise; vapour pressure (25 °C) 0.001 torr. Sulfuric acid is a highly corrosive substance. It can badly burn clothes, plastic, rubber and injured the human skin, eyes, etc. Handle it very carefully.

12.13.3 Chemical properties.

i. Self-ionization (autoprotolysis) and ionization in water. It self-ionizes or undergoes autoprotolysis as follows:

2H₂SO_4(l) ⇌ H₃SO₄⁺_(aq) + HSO₄⁻_(aq)   K = 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. Sulfuric acid is a strong acid as shown by its pK_a1 value:

H₂SO_4(aq) ⇌ H₃O⁺_(aq) + HSO₄⁻_(aq)   (pK_a1 = −2)

But hydrogensulfate (HSO₄⁻) is a far weaker acid due to a positive pK_a2 value:

HSO₄⁻_(aq) + H₂O_(l) ⇌ H₃O⁺_(aq) + SO₄²⁻_(aq)   (pK_a2 = 1.92)

ii. Dehydrating agent. Concentrated sulfuric 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.

📌 Interesting Information (book's side note): a common laboratory demonstration is the dehydration of table sugar, where a black porous carbon snake protrudes out of the apparatus: C₁₂H₂₂O_11(s) + H₂SO_4(l) → 12C_(s) + 11H₂O_(g) + H₂SO_4(aq). It also dehydrates ethyl alcohol to ethene or to ethoxyethane depending upon the conditions: C₂H₅OH_(l) + H₂SO_4(l) → C₂H_4(g) + H₂O_(l) + H₂SO_4(aq).

iii. With sodium chloride. Hydrochloric acid (HCl) gas is formed when sulfuric acid reacts with sodium chloride: NaCl_(s) + H₂SO_4(l) → NaHSO_4(s) + HCl_(g).

iv. With metals. Reactions of sulfuric acid with metals depend upon the metal, concentration of the acid, and temperature. Metals that are above hydrogen in the electrochemical series such as Fe, Al, Zn, Mn, Ni, and Mg react directly with dilute sulfuric acid to produce hydrogen gas and metal sulfates. But with cold conc. H₂SO₄, they liberate SO₂ and metal sulfates.

Zn_(s) + dil. H₂SO_4(aq) → H_2(g) + Zn²⁺_(aq) + SO₄²⁻_(aq)

Zn_(s) + H₂SO_4(l) → SO_2(g) + Zn²⁺_(aq) + SO₄²⁻_(aq) + 2H₂O_(l)

v. As oxidising agent. Metals like Cu, Ag, and Hg react with hot conc. H₂SO₄ to form metal sulfates. Sulfuric acid is not regarded as a typical oxidising agent due to the stability of SO₄²⁻ anion. This anion is weakly oxidising. However hot concentrated sulfuric acid is a moderately strong oxidising agent due to high temperature, high concentration of protons (H⁺), and formation of nascent oxygen. Hot concentrated sulfuric acid oxidises Cu, as given below:

Cu_(s) + 2H₂SO_4(aq) → SO_2(g) + 2H₂O_(l) + SO₄²⁻_(aq) + Cu²⁺_(aq)

Quick Check 12.7 (book): (a) Dehydration reactions of conc. H₂SO₄ with starch and oxalic acid — see Quick Check 12.6(b) above: the chapter asks this but gives no starch or oxalic-acid equation, so the answers there are flagged as the standard school pattern, not book equations. (b) Conc. H₂SO₄ reacts with NaCl to give NaHSO₄ + HCl gas (equation iii above); this reaction is important because it is the laboratory preparation of hydrogen chloride gas. (c) Difference between the oxidising power of cold and hot sulfuric acid: cold dilute acid is not an oxidising agent in these reactions — metals above hydrogen simply displace H₂; cold concentrated acid already liberates SO₂ instead of H₂, and hot concentrated acid is a moderately strong oxidising agent (high temperature + high H⁺ concentration + nascent oxygen) that oxidises even metals below hydrogen such as Cu, Ag and Hg to their sulfates with SO₂ evolved.

12.13.4 Uses and industrial applications. Sulfuric acid is considered a king of chemicals and its consumption is an indicator of the industrial progress of a country.

1. A major portion of the acid is used in making fertilizers, normally three fourth is used to digest the phosphate rock. Sulfuric acid also reacts with ammonia to make ammonium sulfate fertilizer.
2. It is used in the extraction of metals from ores such as Cu, Ni, steel etc.
3. It is utilized as a catalyst in oil and coal refining, polymers, synthetic rubber, and plastic industries.
4. It is used in the pulp and paper industry and involved in the production of pesticides, insecticides, herbicides, varnishes, dyes, pharmaceuticals, soaps and detergents.
5. It is used for nitration in making explosives such as trinitrotoluene (TNT), nitroglycerine, picric acid, nitrocellulose, etc.
6. It is involved in the food industry for making sugar, starch, and corn syrup.
7. It is used in the paint industry for making titanium dioxide (TiO₂) pigment.
8. It is used to dry gases in industrial processes.
9. 35.67% acid is used in lead storage batteries.
10. It is used as a laboratory reagent.

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

📌 Did You Know? (book's side note): calcium phosphate (Ca₃(PO₄)₂) and fluorapatite (Ca₅(PO₄)₃F) are used to make single superphosphate Ca(H₂PO₄)₂ and phosphoric acid (H₃PO₄). H₃PO₄ is further used to make double and triple superphosphates and ammonium phosphate.