Contents: 8 sections
Syllabus points
- Explain the lack of reactivity of nitrogen.
- Describe the structure of ammonia, its basicity and its reaction with acids.
- Describe the industrial production of ammonia and the conditions used.
- Explain the environmental consequences of the oxides of nitrogen and sulfur.
- Describe the industrial production of sulfuric acid by the Contact process.
Nitrogen and why it is unreactive
Nitrogen makes up about 78% of the atmosphere and takes part in almost nothing at room temperature.
The reason is the triple bond in N₂. Its bond energy is around 994 kJ mol⁻¹, which is very high, and the molecule is also non-polar, so there is no site for a nucleophile or an electrophile to attack. Breaking that bond requires either extreme conditions or a catalyst, which is why fixing nitrogen industrially is such a demanding process.
That unreactivity is useful in itself: nitrogen is used as an inert atmosphere in food packaging and in chemical plant.
Ammonia
Ammonia has a pyramidal shape with a bond angle of about 107°. Nitrogen has three bonding pairs and one lone pair, and the lone pair repels more strongly than the bonding pairs, squeezing the angle down from the tetrahedral 109.5°.
That lone pair explains ammonia's chemistry. It can be donated to a proton, so ammonia is a base:
NH₃ + H⁺ → NH₄⁺
The bond formed is dative covalent, because both electrons come from the nitrogen. Once formed it is identical to the other three, so the ammonium ion is a regular tetrahedron with bond angles of 109.5°.
Ammonia is very soluble in water, partly because it forms hydrogen bonds with it, and the solution is weakly alkaline:
NH₃ + H₂O ⇌ NH₄⁺ + OH⁻
With acids it gives ammonium salts. With hydrogen chloride gas the two produce dense white fumes of solid ammonium chloride, which is a useful observation.
Warming an ammonium salt with sodium hydroxide releases ammonia, identified by its sharp smell and by turning damp red litmus blue. That is the standard test for the ammonium ion.
Making ammonia: the Haber process
N₂ + 3H₂ ⇌ 2NH₃, and the forward reaction is exothermic.
| Condition | Typical value | Why |
|---|---|---|
| Temperature | About 450 °C | A compromise: low favours yield, high favours rate |
| Pressure | About 200 atmospheres | Fewer gas molecules on the right, so high pressure favours ammonia |
| Catalyst | Iron | Speeds the reaction without shifting the equilibrium |
The temperature is the point worth understanding rather than memorising. The reaction is exothermic, so by Le Chatelier a low temperature would give a higher yield, but at low temperature the rate is uneconomically slow. The figure chosen is a compromise between yield and rate.
The catalyst does not increase the yield. It lets equilibrium be reached quickly enough to be worth doing.
Ammonia matters because it is the route to nitrogenous fertilisers, and so to a large share of the world's food supply.
Oxides of nitrogen and the environment
Nitrogen and oxygen do not react at ordinary temperatures, but in the very high temperatures inside a car engine they do:
N₂ + O₂ → 2NO
Once in the air, NO is oxidised further to NO₂. The consequences are:
- Acid rain. NO₂ dissolves and is oxidised to nitric acid, lowering the pH of rain, damaging trees and aquatic life and eroding limestone buildings.
- Photochemical smog. Nitrogen oxides react with unburnt hydrocarbons in sunlight to form ozone and other irritants at ground level.
- Catalysing the oxidation of sulfur dioxide, which makes the acid rain problem worse still.
A catalytic converter removes them. Over a platinum and rhodium surface:
2CO + 2NO → 2CO₂ + N₂
This is heterogeneous catalysis: the gases react on the solid surface, which is why the catalyst is spread as a thin coating over a honeycomb to maximise the area exposed.
Sulfur dioxide and acid rain
Sulfur dioxide comes mainly from burning fossil fuels that contain sulfur compounds. In the atmosphere:
SO₂ + H₂O → H₂SO₃
and further oxidation gives sulfuric acid. The effects are the same family as those of the nitrogen oxides: acidified lakes, damaged forests, and eroded limestone and marble.
Flue gas desulfurisation removes it at the power station, typically by reaction with calcium oxide or calcium carbonate:
CaO + SO₂ → CaSO₃
The calcium sulfite can be oxidised to calcium sulfate and sold as gypsum for plasterboard, which turns a pollutant into a product.
Sulfuric acid: the Contact process
Three stages:
1. Burn sulfur.
S + O₂ → SO₂
2. Oxidise the sulfur dioxide, which is the equilibrium step:
2SO₂ + O₂ ⇌ 2SO₃
The conditions are about 450 °C, 1 to 2 atmospheres, and a vanadium(V) oxide catalyst. The forward reaction is exothermic, so again the temperature is a compromise. The pressure is low because the yield is already around 96%, and higher pressure would cost more than the small extra yield is worth. That is a genuinely useful contrast with the Haber process, where high pressure is worth paying for.
3. Absorb the sulfur trioxide. It is dissolved in concentrated sulfuric acid to give oleum, and the oleum is then diluted:
SO₃ + H₂SO₄ → H₂S₂O₇
H₂S₂O₇ + H₂O → 2H₂SO₄
Sulfur trioxide is not added straight to water, because the reaction is violently exothermic and produces a dense mist of acid droplets that is dangerous and difficult to condense.
Common mistakes
- Explaining nitrogen's unreactivity by saying it is a gas. The reason is the strong triple bond and the non-polar molecule.
- Giving the ammonia bond angle as 109.5°. The lone pair reduces it to 107°.
- Saying the Haber catalyst increases the yield. It increases the rate only.
- Saying a high temperature is used in the Haber process for a better yield. It lowers the yield and is used for rate.
- Adding SO₃ directly to water in the Contact process. It is absorbed in concentrated sulfuric acid first.
- Confusing the pressures: high in the Haber process, low in the Contact process, and knowing why the answer differs.