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Loading home…N₂ has a very strong triple bond (N≡N) between the two atoms.
A large amount of energy is needed to break this bond.
Therefore N₂ is not highly reactive under ordinary conditions.
It makes up about 78% of dry air by volume.
High temperature or catalysts are needed for many nitrogen reactions (e.g. Haber process).
Inertness makes N₂ useful as a protective atmosphere.
N₂ + 3H₂ ⇌ 2NH₃ is the Haber synthesis of ammonia.
A typical industrial compromise uses moderate temperature (about 400–450 °C).
High pressure (about 150–300 atm) favours the forward reaction (fewer gas moles).
An iron catalyst (with promoters) speeds attainment of equilibrium.
Ammonia is removed by liquefaction to shift equilibrium forward.
Le Chatelier’s principle guides these choices for yield and rate.
Ammonia is a Brønsted base and accepts H⁺ from acids.
NH₃ + HCl → NH₄Cl (ammonium chloride).
NH₃ + HNO₃ → NH₄NO₃ (ammonium nitrate).
2NH₃ + H₂SO₄ → (NH₄)₂SO₄ (ammonium sulphate).
These salts are important fertilizers.
The reactions show the basic character of NH₃.
Sulphur is burned: S + O₂ → SO₂.
SO₂ is oxidized: 2SO₂ + O₂ ⇌ 2SO₃ (V₂O₅ catalyst, ~450 °C, 1–2 atm).
SO₃ is absorbed in concentrated H₂SO₄ to form oleum (H₂S₂O₇).
Oleum is diluted carefully with water to give H₂SO₄.
Direct absorption of SO₃ in water is avoided (forms mist).
The process is a classic application of equilibrium and catalysis.
Important oxides include N₂O, NO, NO₂ and N₂O₅ (among others).
NO and NO₂ are often grouped as NOx from combustion.
NOx contribute to photochemical smog and acid rain.
They irritate the respiratory system and damage materials.
Lightning and engines can form NO from N₂ and O₂ at high T.
Control of vehicle and industrial emissions reduces NOx pollution.
SOx (mainly SO₂) come from burning sulphur-containing fuels.
NOx form in high-temperature combustion of air–fuel mixtures.
Both dissolve/react in air and water to give acidic solutions.
This produces acid rain that harms lakes, soil, forests and buildings.
They also contribute to smog and respiratory illness.
Flue-gas treatment and cleaner fuels help reduce emissions.
Nitrogen is the most abundant element in the atmosphere (~78% by volume as N₂).
It occurs in combined form in proteins, nitrates and ammonium salts.
N₂ is a colourless, odourless diatomic gas, slightly soluble in water.
The N≡N triple bond makes it kinetically inert at room temperature.
At high temperature it can react with O₂ to form NO (e.g. in engines/lightning).
It reacts with H₂ under Haber conditions to form NH₃.
Active metals such as Li and Mg can form nitrides on heating.
Nitrogen is used to provide an inert atmosphere and in ammonia manufacture.
Liquid nitrogen is a cryogenic coolant.
Oxidation states of nitrogen range widely (−3 in NH₃ to +5 in HNO₃/NO₃⁻).
In NH₃, nitrogen has oxidation state −3.
N₂ fixation (natural and industrial) converts inert N₂ into bioavailable forms.
The nitrogen cycle links atmosphere, soil and living organisms.
Exam answers should stress abundance, inertness reason, and key industrial reaction.
Give balanced equations for Haber synthesis and nitride formation where asked.
Contrast physical inertness with the importance of nitrogen compounds.
Mention that “inert” here means not highly reactive, not noble-gas complete inactivity.
Link properties to the strong triple bond and high bond dissociation energy.
State safety: confined N₂ can cause asphyxiation by displacing O₂.
Conclude with uses: fertilizer chain via NH₃, inert blanketing, cryogenics.
The Haber process synthesizes ammonia: N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH negative (exothermic).
Fresh gases are compressed, heated, and passed over an iron catalyst.
High pressure favours the forward reaction because 4 gas moles become 2.
Low temperature favours yield (exothermic forward) but slows the rate.
A moderate temperature (~400–450 °C) is a compromise for useful rate and yield.
The catalyst increases rate without changing the equilibrium constant.
Ammonia is condensed and removed; unreacted gases are recycled.
Removing NH₃ shifts equilibrium toward products (Le Chatelier).
Hydrogen often comes from natural gas reforming; nitrogen from air.
Plant design balances cost of compression with improved yield at high P.
Impurities that poison the catalyst must be removed from feed gases.
Ammonia is used for fertilizers, nitric acid (Ostwald), and many chemicals.
Exam answers should write the equation, state conditions, and justify each with Le Chatelier + kinetics.
Distinguish equilibrium position (K) from rate (catalyst/T).
Note that increasing T decreases K for this exothermic reaction.
Describe the recycle loop as essential for economy.
Mention typical pressures in the hundreds of atmospheres in classical plants.
Link inertness of N₂ to the need for catalyst and elevated temperature.
End with the industrial importance of NH₃ in food production.
Always quote balanced equations and oxidation numbers when discussing N and S compounds in exams.
Sulphuric acid manufacture by the Contact process has three chemical stages.
First: S + O₂ → SO₂ (or roasting of sulphide ores).
Second: 2SO₂ + O₂ ⇌ 2SO₃; this step is exothermic and uses V₂O₅ catalyst.
Moderate temperature (~450 °C) balances rate and equilibrium yield.
Only slight pressure above atmospheric is used; the reaction decreases gas moles, so high P helps a little but is costly.
Excess air/O₂ favours SO₃ formation.
Third: SO₃ is dissolved in concentrated H₂SO₄ to form oleum, then diluted to H₂SO₄.
Direct SO₃ + water forms a corrosive mist, so oleum route is preferred.
In H₂SO₄, sulphur has oxidation state +6.
H₂SO₄ is a strong acid, dehydrating agent, and sulfonating agent.
Le Chatelier explains why temperature must not be too high despite faster rate.
Catalyst choice allows a workable rate at the compromise temperature.
Uses include fertilizers (superphosphate), detergents, dyes, and batteries.
Exam answers need equations, conditions, and clear Le Chatelier reasoning.
State why SO₃ is not absorbed directly in water.
Mention purification of SO₂ gas before the converter.
Note that Contact process replaced older lead-chamber methods for concentrated acid.
Link oxidation of SO₂ to environmental control of SO₂ emissions.
Conclude with the economic importance of H₂SO₄ as a key industrial chemical.
Always quote balanced equations and oxidation numbers when discussing N and S compounds in exams.
NOx (NO, NO₂) form when N₂ and O₂ react at high temperatures in engines and furnaces.
SO₂ forms when sulphur in coal/oil is burned; some SO₃ may also form.
In air, NO is oxidized to NO₂; SO₂ can be oxidized to SO₃.
These oxides dissolve and react with moisture to give HNO₃ and H₂SO₄.
Rain with lower than normal pH is called acid rain.
Acid rain damages limestone buildings, metal structures, soils and aquatic life.
NOx also help form photochemical smog with unburnt hydrocarbons and sunlight.
Health effects include asthma and other respiratory problems.
Control methods: low-sulphur fuels, flue-gas desulphurization, catalytic converters.
Catalytic converters reduce NOx, CO and unburnt hydrocarbons in exhaust.
Tall chimneys dilute pollutants locally but do not destroy them.
Monitoring ambient SO₂/NOx guides environmental policy.
SO₂ can be recognised by its sharp smell and reducing properties in lab tests.
SO₂ + water gives sulphurous acid (H₂SO₃); further oxidation yields sulphuric acid.
Exam answers should name sources, chemical path to acid rain, effects, and controls.
Write simplified equations: 2SO₂ + O₂ → 2SO₃; SO₃ + H₂O → H₂SO₄.
For NOx: 2NO + O₂ → 2NO₂; related nitric acid formation in droplets.
Distinguish sulphurous vs sulphuric acid where relevant.
Emphasise prevention at source as better than treating damaged ecosystems.
Always quote balanced equations and oxidation numbers when discussing N and S compounds in exams.
The Ostwald process oxidizes ammonia to nitric acid on an industrial scale.
Step 1: 4NH₃ + 5O₂ → 4NO + 6H₂O over a Pt–Rh catalyst at high temperature.
Step 2: 2NO + O₂ → 2NO₂.
Step 3: 4NO₂ + 2H₂O + O₂ → 4HNO₃ (overall absorption/oxidation in water).
Ammonia feedstock comes from the Haber process, linking N₂ → NH₃ → HNO₃.
Nitric acid is used for fertilizers (NH₄NO₃), explosives, and nitration reactions.
In NH₃, N is −3; in HNO₃, N is +5 — a large oxidation-state change.
The first step is highly exothermic and carefully temperature-controlled.
Excess air ensures oxidation of NO to NO₂.
Absorption towers produce aqueous HNO₃ of useful concentration.
Side reactions and catalyst poisoning must be managed for yield.
Lab preparation of NH₃ often uses NH₄Cl + NaOH (heat): NH₄Cl + NaOH → NH₃ + NaCl + H₂O.
That lab gas can be tested with HCl fumes (white NH₄Cl smoke).
Exam answers should give the three Ostwald stages with equations.
State catalyst and conditions for NH₃ oxidation.
Link Haber and Ostwald as a nitrogen-industry chain.
Mention concentrated HNO₃ as an oxidizing acid.
Note environmental care: NOx leaks are pollutants.
Conclude with major uses of HNO₃ in agriculture and industry.
Always quote balanced equations and oxidation numbers when discussing N and S compounds in exams.