Nitrogen and sulfur

Why the nitrogen molecule is so unreactive

Nitrogen makes up about 78% of the air by volume and is obtained in bulk by the fractional distillation of liquid air, yet at room temperature it does almost nothing. It does not burn, it is not attacked by dilute acids or by alkalis, and it is not oxidised by the oxygen it is mixed with. That is why nitrogen is used as an inert atmosphere for storing reactive chemicals and for filling crisp packets. You have to explain this with two features of the N2 molecule, not with the properties of the nitrogen atom.

The first feature is the strength of the triple bond. The two nitrogen atoms are joined by one sigma bond and two pi bonds, formed by the overlap of p orbitals on each atom. The bond enthalpy of N2 is about 945 kJ mol-1, which makes it one of the strongest bonds you meet. Any reaction of nitrogen has to break that bond, so the activation energy is very high and the reaction is extremely slow unless the temperature is very high or a catalyst is present.

The second feature is that the molecule is non-polar. Nitrogen and oxygen differ in electronegativity, but the two atoms in N2 are identical, so the bonding electrons are shared equally and there is no permanent dipole. The molecule has no partially positive end for a negative ion or a lone pair to attack, and no partially negative end for a positive ion to attack. Polar reagents such as H+, OH- and water, which attack polar or easily polarised bonds, find nothing to attack in N2 at room temperature.

Put the two ideas together. Each nitrogen atom in N2 has a lone pair, so the molecule can in principle act as a Lewis base, but the lone pair is held in a very small, tightly bound orbital and the molecule is not polarisable, so N2 is a very weak ligand and also has a very high first ionisation energy of about 1400 kJ mol-1. Nitrogen only reacts when something supplies enough energy to break the triple bond or a catalyst provides a lower-energy route: the Haber process does it with an iron catalyst at about 450 °C and 200 atm, nitrogen-fixing bacteria do it with the enzyme nitrogenase, and a lightning strike does it with a sudden burst of electrical energy.

• Bond enthalpy of N2 is about 945 kJ mol-1, compared with about 158 kJ mol-1 for an N-N single bond, so breaking the triple bond costs roughly six times as much energy.
• N2 is non-polar, with no permanent dipole, so polar reagents have no partially charged end to attack.
• Nitrogen is fixed naturally by lightning and by nitrogen-fixing bacteria, and industrially in the Haber process: N2(g) + 3H2(g) <-> 2NH3(g).

Ammonia as a Bronsted-Lowry base, and the ammonium ion

Ammonia is a pyramidal molecule: the nitrogen atom carries one lone pair and three N-H bonds pushed down below it, with an H-N-H bond angle of about 107 °. That lone pair is what makes ammonia a base. A Bronsted-Lowry base is a proton (H+) acceptor, and in water ammonia accepts a proton from a water molecule, which acts as the acid: NH3(aq) + H2O(l) <-> NH4+(aq) + OH-(aq).

The equilibrium lies well to the left, so a solution of ammonia is only weakly alkaline and is described as a weak base: only a small fraction of the molecules has taken a proton, so the concentration of hydroxide ions is small. A 0.1 mol dm-3 solution of ammonia has a pH of about 11, whereas 0.1 mol dm-3 sodium hydroxide, which is fully ionised, has a pH of about 13. The pair NH3 and NH4+ is a conjugate acid-base pair, related by the gain or loss of one proton.

Because it accepts protons, ammonia neutralises acids to give ammonium salts: NH3(aq) + HCl(aq) -> NH4Cl(aq), 2NH3(aq) + H2SO4(aq) -> (NH4)2SO4(aq), and NH3(aq) + HNO3(aq) -> NH4NO3(aq). Ammonium nitrate and ammonium sulfate are made on a huge scale as nitrogen fertilisers, and ammonium chloride is used as a soldering flux and in zinc-carbon dry cells. All these salts are ionic, made of NH4+ ions and the anion of the acid, and most are very soluble in water. Their solutions are slightly acidic, because the ammonium ion itself can donate a proton to water: NH4+(aq) + H2O(l) <-> NH3(aq) + H3O+(aq).

The same lone pair explains the shape of the product. When the lone pair on the nitrogen of ammonia attacks the proton of an acid, a dative (coordinate) covalent bond forms: both electrons in the new N-H bond come from the nitrogen atom. The product is the ammonium ion, NH4+, which is tetrahedral with bond angles of about 109.5 °. Once the bond has formed there is nothing to distinguish it from the other three: all four N-H bonds are the same length and the same strength. There is no lone pair left on the nitrogen, so the ammonium ion cannot accept another proton.

• Bronsted-Lowry base = proton acceptor; in water, NH3 + H2O <-> NH4+ + OH-.
• Ammonium ion = NH4+, tetrahedral, four identical N-H bonds, one of them formed by a dative bond from the nitrogen lone pair.
• Ammonium salts are ionic and soluble, and their solutions are slightly acidic because NH4+ donates a proton to water.

Displacing ammonia from ammonium salts

Adding a strong alkali to an ammonium salt displaces the ammonia. Warm the salt with sodium hydroxide solution and ammonia gas comes off: NH4+(aq) + OH-(aq) -> NH3(g) + H2O(l), or in full, NH4Cl(aq) + NaOH(aq) -> NaCl(aq) + NH3(g) + H2O(l) and (NH4)2SO4(aq) + 2NaOH(aq) -> Na2SO4(aq) + 2NH3(g) + 2H2O(l). This is again an acid-base reaction, but now the two bases are competing for the same proton.

The rule is that the proton ends up on the stronger base. The hydroxide ion is a stronger base than ammonia, so it takes the proton from the ammonium ion and leaves neutral ammonia behind. The reaction is also helped along by heating, because ammonia is a gas and it bubbles out of the solution instead of staying to react back again, so the removal of the product keeps driving the reaction to the right.

This displacement is the standard test for an ammonium ion in an unknown salt. Warm the substance with sodium hydroxide solution and hold a piece of damp red litmus paper in the gas: it turns blue, because ammonia dissolves in the water on the paper to give hydroxide ions. Holding a glass rod dipped in concentrated hydrochloric acid near the mouth of the tube is a confirming test: a white smoke of fine ammonium chloride particles forms, NH3(g) + HCl(g) -> NH4Cl(s). Ammonia also has a sharp, pungent smell, but you must never put your nose to the tube and sniff - waft the vapour towards you with your hand.

The same reaction is the basis of the laboratory determination of the nitrogen content of a fertiliser: the ammonium salt is decomposed with alkali and the ammonia produced is distilled into a measured volume of standard acid, which is then titrated. It is also why ammonium salts must not be stored next to lime, and why ammonium sulfate fertiliser mixed with slaked lime on a farm loses its nitrogen as ammonia gas.

Oxides of nitrogen: where they come from

Nitrogen and oxygen sit next to each other in the air without reacting at room temperature, but heat them enough and they combine: N2(g) + O2(g) -> 2NO(g). The reaction is endothermic and needs the very high activation energy required to break the N-N triple bond, so it happens only where the temperature is around 2000 °C. That is exactly what a lightning strike provides, and it is also what happens inside the cylinder of a petrol or diesel engine when the fuel-air mixture burns. The nitric acid plant does it deliberately, oxidising ammonia over a platinum-rhodium gauze at about 900 °C: 4NH3(g) + 5O2(g) -> 4NO(g) + 6H2O(g).

The natural sources are therefore lightning, forest fires and volcanoes, together with the work of soil bacteria. Nitrifying bacteria oxidise ammonium compounds in the soil to nitrite and then to nitrate, and denitrifying bacteria break nitrate down and release nitrogen monoxide and nitrous oxide, N2O, back into the air. The man-made sources are any high-temperature combustion of air: road vehicles, aircraft, ships, power stations, gas boilers and furnaces. N2O is also released from nitrogen fertilisers spread on fields, and it matters because it is a powerful greenhouse gas as well as an ozone-depleting gas in the stratosphere.

Once nitrogen monoxide is released it does not stay as NO for long. It is slowly oxidised by oxygen in the air: 2NO(g) + O2(g) -> 2NO2(g). The reaction is slow at the low concentrations found in the atmosphere, so NO can travel a long way before it becomes NO2, which is one reason the effects of traffic emissions are felt far from the road.

You should be able to work out the oxidation state of nitrogen in each oxide and to link that to its behaviour. The lower oxides are neutral, and the higher ones are acidic and dissolve in water to give nitrous and nitric acids.

Removing nitrogen oxides from exhaust gases

A modern petrol car carries a three-way catalytic converter in the exhaust pipe. The gases pass through a ceramic honeycomb coated with a washcoat of aluminium oxide, on which platinum, palladium and rhodium are dispersed as very fine particles, so that a huge surface area is in contact with the gas stream. Two different jobs are done at once: the oxidation of carbon monoxide and unburned hydrocarbons, 2CO(g) + O2(g) -> 2CO2(g) and 2C8H18(l) + 25O2(g) -> 16CO2(g) + 18H2O(l), and the reduction of the nitrogen oxides to nitrogen, 2NO(g) + 2CO(g) -> N2(g) + 2CO2(g) and 2NO2(g) + 4CO(g) -> N2(g) + 4CO2(g). The nitrogen formed is harmless, because it is the gas the air is mostly made of.

For both jobs to work at the same time the engine must run with the air-to-fuel ratio held very close to its ideal value, so that there is enough reducing gas (CO and hydrocarbons) to reduce the NOx but also enough oxygen to oxidise the CO and the hydrocarbons. A lambda sensor in the exhaust monitors the oxygen content and the engine management system adjusts the fuelling many times a second. Catalysts only work above a light-off temperature of roughly 250 °C, so most of the NOx from a short journey is produced in the first couple of minutes, before the converter has warmed up.

The catalyst is poisoned by substances that stick to its active sites. Lead, which used to be added to petrol as tetraethyllead, coats the platinum and stops it working permanently, and sulfur compounds in the fuel also block the active sites and can be oxidised to sulfur dioxide and sulfate. That is why leaded petrol has been phased out wherever catalytic converters are used, and why low-sulfur petrol is now the standard.

Vehicles are not the only source. Power stations and industrial furnaces use selective catalytic reduction: ammonia or urea solution is injected into the flue gas, which then passes over a catalyst of vanadium(V) oxide on titanium dioxide, where 4NO(g) + 4NH3(g) + O2(g) -> 4N2(g) + 6H2O(g). The ammonia is the reducing agent, and the catalyst makes it react with NO rather than simply being carried out of the stack. You cannot prevent the formation of NOx completely, because the air in the furnace contains both nitrogen and oxygen; the only options are to lower the flame temperature, to recirculate some exhaust gas, and then to clean up what is left.

PAN and photochemical smog

Sunlight does chemistry that darkness cannot. Nitrogen dioxide absorbs ultraviolet light and splits apart: NO2 + ultraviolet light -> NO + O. The oxygen atom produced is extremely reactive and combines at once with an oxygen molecule: O + O2 -> O3. That is how ground-level ozone, a harmful pollutant, is made in the lower atmosphere. It has nothing to do with the ozone layer high in the stratosphere, which is beneficial because it absorbs ultraviolet radiation before it reaches us.

If only NOx were present, ozone would reach a steady level, because NO destroys ozone again. Unburned hydrocarbons from vehicle exhausts break that balance. They are attacked by hydroxyl radicals and converted into peroxy radicals, and the oxidation of acetaldehyde gives in particular the peroxyacetyl radical, CH3C(O)OO. These radicals oxidise NO to NO2 without consuming ozone, so ozone accumulates instead of being removed. The peroxyacetyl radical then reacts with NO2 to form peroxyacetyl nitrate, PAN, whose formula is CH3C(O)OONO2.

PAN is a lachrymator: it irritates the eyes and the lining of the lungs, and it damages crops and trees. It is one of the mixture of pollutants - ozone, aldehydes, NO2, PAN and unburned hydrocarbons - known as photochemical smog, the brown haze that hangs over large cities. Because NO2 is what gives the haze its colour and because the chemistry needs ultraviolet light, photochemical smog builds up during the morning and is worst in the early afternoon on a hot, still, sunny day, when there is little wind to disperse it. Los Angeles, Mexico City, Athens and Beijing are the classic examples.

PAN is a stable enough molecule to survive being carried by the wind for days, and it is thermally unstable, so it slowly decomposes and releases NO2 far from the traffic that made it. In that way PAN acts as a reservoir that transports nitrogen oxides into the countryside, where the NO2 can then take part in further ozone formation and in acid rain.

• Photodissociation: NO2 + ultraviolet light -> NO + O, followed by O + O2 -> O3.
• Unburned hydrocarbons give peroxy radicals, including the peroxyacetyl radical CH3C(O)OO.
• PAN = CH3C(O)OONO2, formed from the peroxyacetyl radical and NO2; an eye and lung irritant.
• Photochemical smog needs three things together: sunlight, oxides of nitrogen and unburned hydrocarbons.

Sulfur dioxide and the catalytic role of nitrogen dioxide

Sulfur dioxide comes mostly from burning fossil fuels that contain sulfur. Coal and crude oil contain sulfur as sulfide and as organic sulfur compounds, and burning them gives SO2: S(s) + O2(g) -> SO2(g). Metal extraction is the second source: roasting sulfide ores in an excess of air gives the metal oxide and sulfur dioxide, for example 2ZnS(s) + 3O2(g) -> 2ZnO(s) + 2SO2(g). Volcanoes release it naturally. Industrially, sulfur dioxide is also made deliberately, by burning sulfur, because it is the starting material for the Contact process: 2SO2(g) + O2(g) <-> 2SO3(g), carried out at about 450 °C and a modest pressure over a vanadium(V) oxide catalyst.

Sulfur dioxide cannot wait for the Contact process in the atmosphere. The direct gas-phase oxidation SO2 + 0.5O2 -> SO3 has a high activation energy and is far too slow at the temperatures and concentrations found in the air, so without help the sulfur dioxide would simply be dispersed. Nitrogen dioxide supplies that help. The reaction NO2(g) + SO2(g) -> SO3(g) + NO(g) takes place readily, because NO2 can give up one of its oxygen atoms to the sulfur dioxide, and the nitrogen monoxide produced is oxidised straight back to NO2 by the air: 2NO(g) + O2(g) -> 2NO2(g).

Add the two steps together and the overall change is 2SO2(g) + O2(g) -> 2SO3(g). The nitrogen dioxide has been used in the first step and regenerated in the second, so it comes out of the reaction unchanged: it is a homogeneous catalyst for the oxidation of sulfur dioxide. It is worth noticing the contrast with the Contact process, where the same oxidation is catalysed by a heterogeneous solid catalyst, vanadium(V) oxide, in a different phase from the gases.

The sulfur trioxide formed is the acid anhydride of sulfuric acid: SO3(g) + H2O(l) -> H2SO4(aq). In cloud droplets sulfur dioxide is also oxidised directly by hydrogen peroxide and by ozone to sulfate. Because the oxidation happens in the gas phase and inside droplets, the sulfuric acid is often carried far from the power station or smelter that emitted the SO2 and deposited hundreds of kilometres downwind, which is why acid deposition is a problem between countries and not just next to the chimney.

Acid rain: the direct role of NOx and the damage it does

Unpolluted rain is not neutral. Carbon dioxide from the air dissolves in it and forms carbonic acid, CO2(g) + H2O(l) <-> H2CO3(aq), which gives clean rain a pH of about 5.6. Rain is called acid rain when its pH falls well below that, usually below about 5, and the extra acidity comes mainly from sulfuric acid and nitric acid.

Nitrogen oxides act directly by making nitric acid. Nitrogen dioxide dissolves in water and disproportionates: 2NO2(g) + H2O(l) -> HNO3(aq) + HNO2(aq), and the nitrous acid is then oxidised: 2HNO2(aq) + O2(g) -> 2HNO3(aq). The overall change is 4NO2(g) + 2H2O(l) + O2(g) -> 4HNO3(aq). Whichever route you write, the nitrogen monoxide that is released in the air is first oxidised to NO2 by 2NO + O2 -> 2NO2. The same reaction is used deliberately in industry, where NO2 is absorbed in water to make nitric acid: 3NO2(g) + H2O(l) -> 2HNO3(aq) + NO(g), with the NO recycled. Nitrogen oxides act indirectly, as a catalyst, by oxidising sulfur dioxide to sulfur trioxide, which then forms sulfuric acid - the route you have just met.

The damage caused by acid rain is the standard exam material, and you should be able to attach an equation or a mechanism to each effect. In lakes and rivers, acid water releases aluminium ions, Al3+, from the clay minerals in the soil and the sediment; the Al3+ damages the gills of fish and the water becomes unable to support fish and invertebrates, so the lake becomes clear but dead. In soils, acid rain displaces and washes out nutrient ions such as Ca2+ and K+, and it slows the bacterial decomposition on which the soil depends, so trees on thin soils above granite, which cannot neutralise the acid, lose condition.

Acid rain attacks buildings and statues made of limestone and marble, because both are calcium carbonate: CaCO3(s) + H2SO4(aq) -> CaSO4(s) + CO2(g) + H2O(l). The calcium sulfate either dissolves away or forms a crust that flakes off, which is why the detail on old stonework is lost. Metals are corroded too, especially iron and zinc, because both dissolve in dilute sulfuric and nitric acids. Acid rain also carries nitrate into the soil as a nutrient, which can disturb the balance of species in a habitat, and nitric acid contributes to the formation of ground-level ozone.

Control is tackled at the source. Sulfur dioxide is reduced by burning low-sulfur fuel and by flue-gas desulfurisation, in which the flue gas is scrubbed with a slurry of powdered limestone or lime so that CaCO3(s) + SO2(g) -> CaSO3(s) + CO2(g); the calcium sulfite is oxidised to calcium sulfate and sold as gypsum for plasterboard. Nitrogen oxides are reduced by catalytic converters on vehicles, by selective catalytic reduction in power stations, and by lowering the flame temperature in the furnace so that less NO is formed in the first place. International agreements limit how much sulfur and nitrogen each country may emit, because the pollution crosses borders.