Nitrogen and sulfur
Contents: 8 sections
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.
Check you have it
Question 1
What is the maximum volume of sulfur dioxide gas measured at room conditions produced from burning 100 dm³ of diesel fuel containing 0.8346 g of sulfur?
Answer: D.
S + O₂ → SO₂
n(S) = 0.8346 / 32.1 = 0.0260 mol
so n(SO₂) = 0.0260 mol as well.
At room conditions the molar gas volume is 24.0 dm³ mol⁻¹:
V = 0.0260 × 24.0 = 0.624 dm³ = 624 cm³, which is D.
The 100 dm³ of diesel is not used in the calculation at all. It is there to set the scene, and the mass of sulfur it contains has already been given. Reaching for it is the first trap.
B, 312 cm³, is half the answer, from using 64.1 for the molar mass, which is SO₂ rather than S. The mass given is the mass of sulfur, so it is sulfur's 32.1 that converts it to moles.
The 1 : 1 ratio is worth stating because it makes the rest trivial: every sulfur atom in the fuel ends up in one SO₂ molecule, which is why the sulfur content of a fuel is such a direct measure of the acid rain it causes.
Question 2
The structure of the sulfur dioxide molecule is shown. What is the shape of the sulfur dioxide molecule?
Answer: B.
Three regions repel into a trigonal planar arrangement at 120°, but shape is named from the atoms only, never from the lone pairs. Ignore the lone pair when you describe the shape and what is left is bent, which Cambridge calls non-linear. The bond angle is squeezed slightly below 120°, to about 119°, because a lone pair repels more strongly than a bonding pair.
A is what you get by counting only the two bonds and forgetting the lone pair, which is CO2 with its 180° angle. The difference between CO2 and SO2 is that one lone pair and nothing else.
C, pyramidal, needs four regions with one lone pair, as in NH3. D, tetrahedral, needs four regions and no lone pairs, as in CH4. Both would require sulfur to have more around it than it has.
Count regions first, then subtract the lone pairs to name the shape. Doing it in that order gets every one of these right.
Question 3
Carbon monoxide, CO, nitrogen dioxide, NO2, and sulfur dioxide, SO2, are all atmospheric pollutants.
Which reaction occurs in the atmosphere?
Answer: B.
NO₂ + SO₂ → NO + SO₃
The nitrogen goes from +4 down to +2, so NO₂ is reduced to NO, and the sulfur goes from +4 up to +6. The NO is then reoxidised to NO₂ by oxygen in the air, so the nitrogen dioxide is regenerated and acts as a catalyst for the whole process. The SO₃ dissolves in rain to give sulfuric acid.
C describes the catalytic converter, not the atmosphere. Carbon monoxide does reduce nitrogen oxides, but only on a hot platinum surface. In open air the two gases are far too dilute and too cool to react.
A is false. Carbon monoxide is not spontaneously oxidised in the atmosphere. It is stable enough to persist for weeks, which is exactly why it is a dangerous pollutant.
D has the wrong oxidising agent. Carbon dioxide is already fully oxidised and cannot oxidise anything.
The nitrogen dioxide cycle is the mechanism behind acid rain from sulfur dioxide, and the catalytic part is what makes it so effective: a small amount of NO₂ can convert a great deal of SO₂.
What the syllabus asks for on this topicSyllabus points
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.
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