AS 12 Nitrogen and sulfur | A Level Chemistry 9701

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🎯What you need to be able to do

Explain why nitrogen is unreactive, using the strength of its triple bond and its lack of polarity.
Explain the basicity of ammonia in Brønsted–Lowry terms, the structure and formation of the ammonium ion, and how ammonia is displaced from ammonium salts.
State the natural and man-made sources of nitrogen oxides, and explain how they are removed from vehicle exhausts by catalytic converters.
Understand that NO and NO₂ react with unburned hydrocarbons to form PAN, part of photochemical smog.
Describe the role of NO and NO₂ in acid rain, directly and as catalysts for the oxidation of sulfur dioxide.

📚The chemistry

This is the shortest inorganic topic, and the most environmental one. It uses ideas from earlier pages — bond energy (topic 3), Brønsted–Lowry acids and bases (topic 7), redox and homogeneous catalysis (topic 8) — and applies them to air pollution.

Why nitrogen is unreactive

Nitrogen makes up 78% of the air and does almost nothing. Two reasons:

The N≡N triple bond is very strong, 944 kJ mol⁻¹. Any reaction must begin by breaking or weakening it, so the activation energy of nitrogen’s reactions is very high.
The molecule is non-polar. Both atoms are identical, so there is no δ+ or δ− end to attract a nucleophile or an electrophile, and nothing for a polar reagent to attack.

That is why ammonia synthesis needs 450 °C, 200 atmospheres and an iron catalyst.

Ammonia and the ammonium ion

Ammonia is a Brønsted–Lowry base: the lone pair on the nitrogen atom accepts a proton. In water it is a weak base, because the equilibrium lies to the left:

\[ \mathrm{NH_3(aq) + H_2O(l) \rightleftharpoons NH_4^{+}(aq) + OH^{-}(aq)} \]

With an acid, the ammonium ion forms in an acid–base reaction — with hydrogen chloride gas, as white smoke of ammonium chloride:

\[ \mathrm{NH_3(g) + HCl(g) \rightarrow NH_4Cl(s)} \] \[ \mathrm{NH_3 + H^{+} \rightarrow NH_4^{+}} \]

In NH₄⁺ the new N–H bond is a coordinate (dative covalent) bond: both electrons come from the nitrogen’s lone pair. Once it has formed, all four N–H bonds are identical. With four bonding pairs and no lone pair, the ion is tetrahedral, bond angle 109.5° (compared with 107° in pyramidal NH₃).

Displacing ammonia from ammonium salts

Warm any ammonium salt with a base such as sodium hydroxide, and ammonia is released:

\[ \mathrm{NH_4Cl(s) + NaOH(aq) \rightarrow NaCl(aq) + NH_3(g) + H_2O(l)} \] \[ \mathrm{NH_4^{+}(aq) + OH^{-}(aq) \rightarrow NH_3(g) + H_2O(l)} \]

This is an acid–base reaction: NH₄⁺ is the acid, donating a proton to OH⁻, the stronger base. It is also the test for the ammonium ion: the gas turns damp red litmus paper blue. Farmers need to know it too — spreading lime (a base) on soil at the same time as an ammonium fertiliser releases the nitrogen as ammonia gas.

Oxides of nitrogen

Nitrogen and oxygen do not react at normal temperatures, but they do in the extreme heat of a lightning strike (natural) and inside a car or aircraft engine or a power station furnace (man-made):

\[ \mathrm{N_2(g) + O_2(g) \rightarrow 2NO(g)} \]

In the air, nitrogen monoxide is then oxidised to brown nitrogen dioxide:

\[ \mathrm{2NO(g) + O_2(g) \rightarrow 2NO_2(g)} \]

Catalytic converters

A catalytic converter in a car exhaust removes nitrogen oxides together with carbon monoxide. Exhaust gases pass over a honeycomb coated with platinum, palladium and rhodium — a heterogeneous catalyst with a very large surface area. On the metal surface, the oxides of nitrogen are reduced to nitrogen while carbon monoxide is oxidised to carbon dioxide:

\[ \mathrm{2CO(g) + 2NO(g) \rightarrow 2CO_2(g) + N_2(g)} \]

Both products are already present in air. Unburned hydrocarbons are oxidised to CO₂ and H₂O at the same time.

Photochemical smog and PAN

In sunlight, nitrogen oxides react with unburned hydrocarbons from vehicle exhausts to form peroxyacetyl nitrate, PAN. PAN is a component of photochemical smog, the brown haze over busy cities on sunny, still days. It irritates the eyes and lungs and damages plants. You do not need its structure or the reaction steps — just that NO and NO₂ plus unburned hydrocarbons give PAN.

Acid rain

Rain is naturally slightly acidic (about pH 5.6) from dissolved carbon dioxide. Acid rain is more acidic than that, and nitrogen oxides contribute in two ways.

Directly: nitrogen dioxide reacts with water and oxygen to form nitric acid:

\[ \mathrm{4NO_2(g) + O_2(g) + 2H_2O(l) \rightarrow 4HNO_3(aq)} \]

As a catalyst: nitrogen dioxide speeds up the oxidation of sulfur dioxide (from burning sulfur-containing fuels) to sulfur trioxide, and is regenerated:

\[ \mathrm{SO_2(g) + NO_2(g) \rightarrow SO_3(g) + NO(g)} \] \[ \mathrm{NO(g) + \tfrac{1}{2}O_2(g) \rightarrow NO_2(g)} \]

Adding these two steps gives the overall reaction, SO₂ + ½O₂ → SO₃: NO₂ is used in the first step and re-formed in the second, so it is a homogeneous catalyst (all species are gases). The sulfur trioxide then dissolves in rain: SO₃ + H₂O → H₂SO₄.

Acid rain damages limestone buildings, leaches aluminium and nutrients from soils, harms trees, and lowers the pH of lakes enough to kill fish.

✏️Worked example

2.14 g of ammonium chloride is warmed with excess aqueous sodium hydroxide. (a) Write an ionic equation and identify the Brønsted–Lowry acid and base. (b) Calculate the maximum volume of ammonia produced at room conditions. (c) Describe a test to confirm the gas is ammonia. (d) Explain why the bond angle changes when ammonia forms the ammonium ion. [A_r: N 14.0, H 1.0, Cl 35.5; molar volume 24.0 dm³ mol⁻¹]

(a)

\[ \mathrm{NH_4^{+}(aq) + OH^{-}(aq) \rightarrow NH_3(g) + H_2O(l)} \]

NH₄⁺ donates a proton, so it is the acid; OH⁻ accepts it, so it is the base.

(b) M_r(NH₄Cl) = 14.0 + 4(1.0) + 35.5 = 53.5.

\[ n = \frac{2.14}{53.5} = 0.0400\ \mathrm{mol} \qquad V = 0.0400 \times 24.0 = 0.960\ \mathrm{dm^3} = 960\ \mathrm{cm^3} \]

(c) Hold damp red litmus paper in the gas: it turns blue. Ammonia is the only common alkaline gas. (A glass rod dipped in concentrated hydrochloric acid also gives white smoke of NH₄Cl.)

(d) NH₃ has three bonding pairs and one lone pair. The lone pair repels more strongly than bonding pairs, closing the H–N–H angle to 107°. In NH₄⁺ the lone pair has become the fourth (dative) bonding pair, so there are four equal bonding pairs and no lone pairs: the ion is tetrahedral, 109.5°.

Check it. The ratio is 1 : 1 — one ammonium ion gives one ammonia molecule — so moles of gas equal moles of salt. 960 cm³ from about 2 g of salt is reasonable: 1 mol of any gas is 24 dm³, and 0.04 mol is 4% of that.

Calling this a displacement redox reaction, or using dry litmus. Nothing changes oxidation number (N stays −3 throughout); it is a proton transfer. And the litmus must be damp: ammonia only produces OH⁻ ions, which change the indicator, when it dissolves in water.

📝Practise

Work through these, then reveal the answer. Each question targets a different objective from the list above.

1. Give two reasons why nitrogen gas is unreactive.

(1) The N≡N triple bond is very strong (944 kJ mol⁻¹), so a great deal of energy is needed to break it and reactions of nitrogen have a high activation energy. (2) The N₂ molecule is non-polar: the two atoms are identical, so there are no partial charges to attract electrophiles or nucleophiles.

2. Explain, using the Brønsted–Lowry theory, why an aqueous solution of ammonia is alkaline.

Ammonia acts as a base: the lone pair on nitrogen accepts a proton from a water molecule, which acts as an acid. NH₃ + H₂O ⇌ NH₄⁺ + OH⁻. This produces hydroxide ions, so [OH⁻] > [H⁺] and the solution is alkaline. The equilibrium lies to the left — only a small fraction of the ammonia is protonated — so ammonia is a weak base, with a pH of about 11 at 0.1 mol dm⁻³ rather than 13.

3. Explain how nitrogen monoxide forms in a car engine, and why it does not form in the air around the car.

Air, containing N₂ and O₂, is drawn into the engine. The combustion of the fuel produces very high temperatures, which supply enough energy to overcome the very high activation energy of breaking the N≡N bond, so N₂ + O₂ → 2NO. At normal air temperatures, almost no molecules have energy equal to that activation energy, so the reaction does not occur at a measurable rate.

4. Write an equation for the removal of nitrogen monoxide and carbon monoxide in a catalytic converter, and state the type of catalysis.

2CO(g) + 2NO(g) → 2CO₂(g) + N₂(g). NO is reduced (N +2 → 0) and CO is oxidised (C +2 → +4). The catalyst (platinum, palladium, rhodium) is a solid and the reactants are gases, so this is heterogeneous catalysis; the reaction takes place on the metal surface, which is why the catalyst is spread thinly over a honeycomb.

5. Show, with equations, how nitrogen dioxide acts as a catalyst in the formation of acid rain.

Step 1: SO₂ + NO₂ → SO₃ + NO (NO₂ oxidises SO₂). Step 2: NO + ½O₂ → NO₂ (NO is reoxidised by air). Overall: SO₂ + ½O₂ → SO₃. NO₂ is consumed in the first step and regenerated in the second, so it is a catalyst; all species are gases, so it is homogeneous. The SO₃ then forms sulfuric acid in rain: SO₃ + H₂O → H₂SO₄.

6. What is PAN, and what conditions lead to its formation?

PAN is peroxyacetyl nitrate, a component of photochemical smog that irritates eyes and lungs. It forms when oxides of nitrogen (NO and NO₂) react with unburned hydrocarbons, both from vehicle exhausts, in the presence of sunlight. Its formation is favoured in busy cities on sunny, still days, when the pollutants are not dispersed.

🔗Go deeper — other people’s work

These are external resources, not mine. If one stops working, tell me and everything above it on this page still stands.

Chemguide (Jim Clark) — the pages on ammonia and ammonium ions, and on catalytic converters as heterogeneous catalysis
Royal Society of Chemistry — resources on acid rain and air quality