Oxides of Nitrogen, Sulfur Dioxide and Acid Rain

AS · 12 min

Although nitrogen is very unreactive, the extreme temperatures of lightning and of car engines are enough to make it combine with oxygen. The resulting oxides of nitrogen, NO and $\ce{NO2}$ , are pollutants with three effects you must know: they are removed by catalytic converters, they react with unburned hydrocarbons to form the photochemical smog component PAN, and they cause acid rain, both directly and by catalysing the oxidation of sulfur dioxide. This note covers each of those learning outcomes, with the equations and oxidation-number reasoning that Paper 2 questions ask for.

Oxides of nitrogen

The two oxides you need are:

oxide	name	oxidation number of N	appearance
$\ce{NO}$	nitrogen monoxide (nitric oxide)	$+2$	colourless gas
$\ce{NO2}$	nitrogen dioxide	$+4$	brown gas, acidic, toxic

Together they are often written $\ce{NO_x}$ .

Natural occurrence

Lightning. The enormous temperature in a lightning strike provides enough energy to break the $\ce{N#N}$ and $\ce{O=O}$ bonds: $\ce{N2(g) + O2(g) -> 2NO(g)}$
The nitrogen monoxide is then oxidised by oxygen in the air at ordinary temperatures: $\ce{2NO(g) + O2(g) -> 2NO2(g)}$
Some soil bacteria also release oxides of nitrogen.

Man-made occurrence

Internal combustion engines. In the cylinder of a petrol or diesel engine, air (mostly $\ce{N2}$ and $\ce{O2}$ ) is compressed and the fuel burns, reaching temperatures well above 1000 °C. At these temperatures nitrogen and oxygen from the air combine: $\ce{N2 + O2 -> 2NO}$ . The NO is exhausted and oxidised to $\ce{NO2}$ in the atmosphere.
Power stations and furnaces that burn fuels at high temperatures produce oxides of nitrogen in the same way.

Watch out

The nitrogen in the oxides of nitrogen from car engines comes from the air, not from the fuel. Petrol is a mixture of hydrocarbons and contains essentially no nitrogen. The reaction happens because of the high temperature in the engine.

The reaction $\ce{N2 + O2 -> 2NO}$ is endothermic and has a very high activation energy (the $\ce{N#N}$ bond must be broken), which is why it only happens at very high temperatures.

Catalytic removal in catalytic converters

A catalytic converter is fitted to the exhaust system of a car. It contains a ceramic honeycomb coated with a thin layer of the metals platinum, palladium and rhodium, giving a very large surface area. These are heterogeneous catalysts: they are in a different phase (solid) from the reacting gases.

The main reactions remove NO together with the other two pollutants, carbon monoxide and unburned hydrocarbons:

\ce{2CO(g) + 2NO(g) -> 2CO2(g) + N2(g)}

Unburned hydrocarbons are also oxidised by NO (or by oxygen), for example with octane:

\ce{C8H18(g) + 25NO(g) -> 8CO2(g) + 9H2O(g) + $\frac{25}{2}$N2(g)}

and CO and hydrocarbons are oxidised by oxygen: $\ce{2CO + O2 -> 2CO2}$ .

In $\ce{2CO + 2NO -> 2CO2 + N2}$ , nitrogen is reduced from $+2$ to 0 and carbon is oxidised from $+2$ to $+4$ . The harmful NO becomes harmless nitrogen.

Method

How a heterogeneous catalyst works (catalytic converter)

Adsorption: the gas molecules (CO, NO) form weak bonds to the surface of the metal at active sites.
Reaction: bonds within the adsorbed molecules are weakened, so the activation energy is lowered and the molecules react on the surface.
Desorption: the products ( $\ce{CO2}$ , $\ce{N2}$ ) leave the surface, freeing the active sites for more molecules.

Tip

Catalytic converters work only once they are hot (about 300 °C and above), which is why short journeys cause more pollution. They are "poisoned" by lead, so cars with catalytic converters must use unleaded petrol: lead compounds adsorb strongly onto the catalyst surface and block the active sites.

Photochemical smog and PAN

In sunny, still weather in busy cities, oxides of nitrogen and unburned hydrocarbons from vehicle exhausts react together, driven by sunlight, to form a brownish haze called photochemical smog.

A key component is peroxyacetyl nitrate (PAN), $\ce{CH3CO-O-O-NO2}$ (molecular formula $\ce{C2H3NO5}$ ). It forms when NO and $\ce{NO2}$ react with the products of oxidation of hydrocarbons in the air.

Key result

Atmospheric oxides of nitrogen (NO and $\ce{NO2}$ ) react with unburned hydrocarbons to form peroxyacetyl nitrate, PAN, which is a component of photochemical smog. PAN irritates the eyes and lungs (causing breathing problems) and damages plants.

The syllabus does not require the mechanism of PAN formation: you need to know what reacts (NO $_x$ and unburned hydrocarbons) and what forms (PAN, in photochemical smog). Catalytic converters reduce PAN formation by removing both of its sources.

Sulfur dioxide

Sulfur dioxide, $\ce{SO2}$ , is a colourless, toxic gas with a choking smell. Most atmospheric sulfur dioxide from human activity comes from burning fossil fuels that contain sulfur compounds (especially coal, and some oil), for example in power stations:

\ce{S(s) + O2(g) -> SO2(g)}

Volcanoes are a natural source.

Sulfur dioxide is an acidic oxide (see Period 3 oxides, hydroxides and chlorides): it dissolves in water to give sulfurous acid, $\ce{SO2 + H2O <=> H2SO3}$ . In the atmosphere it is also oxidised to sulfur trioxide, which forms sulfuric acid, a much stronger acid:

\ce{SO3(g) + H2O(l) -> H2SO4(aq)}

The oxidation of $\ce{SO2}$ to $\ce{SO3}$ by oxygen alone is very slow. In the atmosphere, nitrogen dioxide speeds it up.

Acid rain

Rain is naturally slightly acidic (pH about 5.6) because it dissolves carbon dioxide. Acid rain has a pH below this, typically 4–4.5 or lower, because of dissolved acids made from oxides of nitrogen and sulfur.

The direct role of nitrogen dioxide

Nitrogen dioxide dissolves in rain water and is oxidised to nitric acid:

\ce{4NO2(g) + O2(g) + 2H2O(l) -> 4HNO3(aq)}

(With water alone, $\ce{NO2}$ disproportionates: $\ce{2NO2 + H2O -> HNO3 + HNO2}$ , giving nitric acid and nitrous acid. Nitrogen goes from $+4$ to $+5$ and to $+3$ .)

The catalytic role of NO and $\ce{NO2}$

Nitrogen dioxide oxidises sulfur dioxide to sulfur trioxide and is itself reduced to nitrogen monoxide. The nitrogen monoxide is then re-oxidised by oxygen in the air, regenerating nitrogen dioxide:

\ce{SO2(g) + NO2(g) -> SO3(g) + NO(g)}

\ce{NO(g) + 1/2O2(g) -> NO2(g)}

Adding the two steps, NO and $\ce{NO2}$ cancel:

\ce{SO2(g) + 1/2O2(g) -> SO3(g)}

Key result

$\ce{NO2}$ is a catalyst for the oxidation of atmospheric $\ce{SO2}$ : it takes part in the reaction but is regenerated at the end, and it speeds up a reaction that would otherwise be very slow. The $\ce{SO3}$ formed dissolves in rain to give sulfuric acid: $\ce{SO3 + H2O -> H2SO4}$ .

Oxidation numbers: S goes from $+4$ ( $\ce{SO2}$ ) to $+6$ ( $\ce{SO3}$ ), oxidised; N goes from $+4$ ( $\ce{NO2}$ ) to $+2$ (NO), reduced, and then back to $+4$ .

This is an example of homogeneous catalysis: the catalyst ( $\ce{NO2}$ ) is in the same phase (gas) as the reactants, and it works by an alternative route involving an intermediate (NO) with a lower activation energy.

Effects of acid rain

Buildings and statues made of limestone or marble ( $\ce{CaCO3}$ ) are eroded: $\ce{CaCO3(s) + H2SO4(aq) -> CaSO4(s) + H2O(l) + CO2(g)}$ .
Metal structures corrode faster: $\ce{Fe(s) + 2H+(aq) -> Fe^2+(aq) + H2(g)}$ .
Lakes and rivers become acidic, killing fish and other aquatic life; acid also releases toxic aluminium ions from soils into the water.
Forests are damaged: leaves are harmed and nutrients such as $\ce{Mg^2+}$ and $\ce{Ca^2+}$ are leached from the soil.

Reducing acid rain (background)

Catalytic converters cut NO $_x$ from vehicles.
Flue-gas desulfurisation removes $\ce{SO2}$ from power-station chimneys using powdered calcium carbonate or calcium oxide (Group 2 chemistry): $\ce{CaCO3(s) + SO2(g) -> CaSO3(s) + CO2(g)}$ , and the calcium sulfite can be oxidised to calcium sulfate (gypsum).
Removing sulfur from fuels before they are burned, and switching to low-sulfur fuels or renewable energy.
Adding lime to acidified lakes neutralises them.

Worked examples

How oxides of nitrogen form in an engine

Explain how nitrogen monoxide is formed in a car engine, and write equations for its formation and for its conversion to nitrogen dioxide in the atmosphere.

Solution

Air drawn into the engine contains nitrogen and oxygen. Burning the fuel produces a very high temperature, which provides enough energy to break the strong $\ce{N#N}$ triple bond (overcome the high activation energy), so nitrogen and oxygen from the air combine:

$\ce{N2(g) + O2(g) -> 2NO(g)}$

In the atmosphere, NO is oxidised by oxygen:

$\ce{2NO(g) + O2(g) -> 2NO2(g)}$

Catalytic converter: oxidation numbers

Write an equation for the removal of nitrogen monoxide and carbon monoxide in a catalytic converter. Use oxidation numbers to identify the oxidising agent, and name a metal used as the catalyst.

Solution

$\ce{2CO(g) + 2NO(g) -> 2CO2(g) + N2(g)}$

N: $+2$ in NO to 0 in $\ce{N2}$ , reduced. C: $+2$ in CO to $+4$ in $\ce{CO2}$ , oxidised. NO is the oxidising agent (it oxidises CO and is itself reduced).

Catalyst: platinum (or palladium or rhodium).

The catalytic role of nitrogen dioxide

Write equations to show how nitrogen dioxide catalyses the oxidation of sulfur dioxide in the atmosphere, and explain why $\ce{NO2}$ is described as a catalyst.

Solution

$\ce{SO2(g) + NO2(g) -> SO3(g) + NO(g)}$

$\ce{NO(g) + 1/2O2(g) -> NO2(g)}$

Overall: $\ce{SO2(g) + 1/2O2(g) -> SO3(g)}$

$\ce{NO2}$ is used in the first step but regenerated in the second, so it is not used up overall and does not appear in the overall equation, yet it increases the rate of oxidation of $\ce{SO2}$ . That is the definition of a catalyst. The $\ce{SO3}$ then forms sulfuric acid in rain: $\ce{SO3 + H2O -> H2SO4}$ .

Exam-style: balancing a hydrocarbon reaction in a converter

Unburned heptane, $\ce{C7H16}$ , reacts with nitrogen monoxide in a catalytic converter to form carbon dioxide, water and nitrogen. Write a balanced equation.

Solution

Carbon: 7 $\ce{CO2}$ . Hydrogen: 16 H gives 8 $\ce{H2O}$ .

Oxygen needed on the right: $7 \times 2 + 8 = 22$ , all from NO, so 22 NO.

Nitrogen: 22 N gives 11 $\ce{N2}$ .

\ce{C7H16(g) + 22NO(g) -> 7CO2(g) + 8H2O(g) + 11N2(g)}

Check: C 7, H 16, O 22, N 22 on each side.

Exam-hard: removing sulfur dioxide from a power station

A power station burns $1.00$ tonne ( $1.00 \times 10^{6}\ \text{g}$ ) of coal containing 2.00% sulfur by mass.

(a) Calculate the volume of $\ce{SO2}$ , at room conditions, formed if all the sulfur is converted to $\ce{SO2}$ .

(b) Calculate the minimum mass of calcium carbonate needed to remove this $\ce{SO2}$ , using $\ce{CaCO3 + SO2 -> CaSO3 + CO2}$ .

(c) Explain, with equations, how any $\ce{SO2}$ that escapes is converted into sulfuric acid in the atmosphere, and the role of nitrogen oxides.

( $A_r$ : S 32.1; $M(\ce{CaCO3}) = 100.1$ )

Solution

(a) Mass of S $= 0.0200 \times 1.00 \times 10^{6} = 2.00 \times 10^{4}\ \text{g}$

$n(\ce{S}) = \dfrac{2.00 \times 10^{4}}{32.1} = 623\ \text{mol} = n(\ce{SO2})$

$V = 623 \times 24.0 = 1.50 \times 10^{4}\ \text{dm}^3$ ( $15.0\ \text{m}^3$ )

(b) $n(\ce{CaCO3}) = 623\ \text{mol}$ ; mass $= 623 \times 100.1 = 6.24 \times 10^{4}\ \text{g} = 62.4\ \text{kg}$

(c) $\ce{NO2}$ oxidises $\ce{SO2}$ : $\ce{SO2 + NO2 -> SO3 + NO}$ ; the NO is re-oxidised by air: $\ce{NO + 1/2O2 -> NO2}$ , so $\ce{NO2}$ acts as a catalyst. Sulfur trioxide dissolves in rain: $\ce{SO3 + H2O -> H2SO4}$ , giving acid rain.

Watch out

NO is colourless; $\ce{NO2}$ is brown. Do not mix them up in observations.
In the catalytic cycle, NO is an intermediate and $\ce{NO2}$ is the catalyst. Both appear in the steps but neither appears in the overall equation.
Catalytic converters remove NO by reducing it to $\ce{N2}$ ; they do not "filter" or "absorb" the gases.
PAN is formed from NO $_x$ and hydrocarbons, not from $\ce{SO2}$ . Acid rain involves $\ce{SO2}$ and NO $_x$ ; photochemical smog involves NO $_x$ and hydrocarbons.

Exam tip

"State and explain the occurrence" of NO $_x$ : give both a natural source (lightning) and a man-made source (internal combustion engines), and explain that high temperature provides the energy to break the $\ce{N#N}$ bond.
"Catalytic removal" questions usually want an equation (often $\ce{2CO + 2NO -> 2CO2 + N2}$ ) and the name of a catalyst metal; sometimes the three steps of heterogeneous catalysis.
For the catalytic role of $\ce{NO2}$ , examiners want both equations and the statement that $\ce{NO2}$ is regenerated.
Balancing practice is common: hydrocarbons with NO in a converter, and oxidation of NO. Always check O and N separately.

Practical skills

Sulfur dioxide can be detected in the laboratory because it is a reducing agent: it decolourises acidified aqueous potassium manganate(VII) (purple to colourless). The qualitative analysis notes use the same reaction to identify sulfite ions, which release $\ce{SO2}$ with acid. Nitrogen dioxide is recognisable as a brown gas, for example from the thermal decomposition of Group 2 nitrates. Both gases are toxic, so any experiment producing them is carried out in a fume cupboard. The acidity of rain water samples can be compared using a pH meter, calibrated with buffer solutions before use.

Summary

NO ( $+2$ , colourless) forms from $\ce{N2 + O2 -> 2NO}$ at very high temperatures: lightning (natural) and internal combustion engines (man-made); then $\ce{2NO + O2 -> 2NO2}$ (brown).
Catalytic converters (Pt, Pd, Rh) remove NO: $\ce{2CO + 2NO -> 2CO2 + N2}$ ; hydrocarbons are also oxidised. Adsorption, reaction, desorption.
NO $_x$ + unburned hydrocarbons form PAN, a component of photochemical smog.
Acid rain directly: $\ce{4NO2 + O2 + 2H2O -> 4HNO3}$ .
Catalytic role: $\ce{SO2 + NO2 -> SO3 + NO}$ , then $\ce{NO + 1/2O2 -> NO2}$ ; then $\ce{SO3 + H2O -> H2SO4}$ .
$\ce{SO2}$ comes from burning sulfur-containing fossil fuels; acid rain damages limestone, metals, lakes and trees.

Practice

Question

State the oxidation number of nitrogen in NO, $\ce{NO2}$ and $\ce{HNO3}$ .
Write the equation for the formation of nitrogen monoxide during a lightning strike, and explain why this reaction does not happen at room temperature.
Name the three metals used in catalytic converters and state why they are coated on a honeycomb.
Write an equation for the reaction between nitrogen monoxide and carbon monoxide in a catalytic converter.
State what PAN is a component of, and name the two types of pollutant that react to form it.
Write equations to show the catalytic role of nitrogen dioxide in the oxidation of sulfur dioxide.
Write an equation for the formation of nitric acid from nitrogen dioxide, oxygen and water.
Balance the equation for the reaction of octane with nitrogen monoxide in a catalytic converter, using whole numbers: $\ce{C8H18 + NO -> CO2 + H2O + N2}$ .
Explain, with an equation, why limestone statues are damaged by acid rain.
A car engine produces $0.300\ \text{g}$ of NO per kilometre. Calculate the volume of NO (at room conditions) produced on a $50.0\ \text{km}$ journey, and the minimum mass of CO that would be needed in the catalytic converter to remove it all by $\ce{2CO + 2NO -> 2CO2 + N2}$ . ( $A_r$ : N 14.0, O 16.0, C 12.0)

Answers

NO $+2$ ; $\ce{NO2}$ $+4$ ; $\ce{HNO3}$ $+5$ .
$\ce{N2(g) + O2(g) -> 2NO(g)}$ . The $\ce{N#N}$ triple bond is very strong, so the activation energy is very high; at room temperature almost no molecules have enough energy to react. Lightning supplies a very high temperature.
Platinum, palladium and rhodium. The honeycomb gives a very large surface area for the gases to adsorb onto, while using only a small mass of the expensive metals.
$\ce{2CO(g) + 2NO(g) -> 2CO2(g) + N2(g)}$ .
Photochemical smog. Oxides of nitrogen (NO and $\ce{NO2}$ ) and unburned hydrocarbons.
$\ce{SO2 + NO2 -> SO3 + NO}$ ; $\ce{NO + 1/2O2 -> NO2}$ (overall $\ce{SO2 + 1/2O2 -> SO3}$ , with $\ce{NO2}$ regenerated).
$\ce{4NO2(g) + O2(g) + 2H2O(l) -> 4HNO3(aq)}$ .
$\ce{2C8H18 + 50NO -> 16CO2 + 18H2O + 25N2}$ . Check: C 16; H 36; O 50 = 32 + 18; N 50.
Limestone and marble are calcium carbonate, which reacts with the acids in acid rain (and the soluble/flaking product is washed away): $\ce{CaCO3(s) + H2SO4(aq) -> CaSO4(s) + H2O(l) + CO2(g)}$ (or with $\ce{2HNO3}$ to give $\ce{Ca(NO3)2}$ ).
Mass of NO $= 0.300 \times 50.0 = 15.0\ \text{g}$ . $n(\ce{NO}) = 15.0 / 30.0 = 0.500\ \text{mol}$ ; $V = 0.500 \times 24.0 = 12.0\ \text{dm}^3$ . CO : NO $= 1 : 1$ , so $n(\ce{CO}) = 0.500\ \text{mol}$ ; mass $= 0.500 \times 28.0 = 14.0\ \text{g}$ .