Nitrogen is a diatomic molecule in its elemental form, N2, and it makes up 78% of the atmosphere on Earth. It can be separated from air by cryogenic fractional distillation or pressure swing adsorption.
Nitrogen is very unreactive because of the strong triple bond between the atoms in the molecule. In the atmosphere it reacts with oxygen in the presence of lightning to form nitrogen(II) oxide.
N2(g) + O2(g) ⇾ 2NO(g)
Nitrogen(II) oxide is oxidised forming nitrogen(IV) oxide which is soluble in rain water.
2NO(g) + O2(g) ⇾ NO2(g)
NO2(g) + H2O(l) + ½O2(g) ⇾ 2HNO3(aq)
This is one way by which atmospheric nitrogen is converted into a soluble form (nitrate ions, NO3–(aq)) that can be taken up by plants and then converted into amino acids and proteins.
Nitrogen(II) oxide is also formed in car engines, a consequence of the high temperatures and pressures needed to ignite the fuel. Nitrogen oxides are pollutants contributing to the build up of photochemical smog in cities and nitrogen(IV) oxide plays a role in the formation of acid rain, catalysing the oxidation of sulfur dioxide to sulfur trioxide which dissolves in water forming sulfuric acid, H2SO4.
SO2(g) + NO2(g) ⇾ SO3(g) + NO(g)
NO(g) + ½O2(g) ⇾ NO2(g)
The nitrogen cycle
The nitrogen cycle charts the processes that convert nitrogen from one form to another by plants, bacteria and increasingly by industry – fixation, nitrification, ammonification and denitrification.
Atmospheric nitrogen must be converted into a usable form to be taken up by plants (known as nitrogen fixation). Some nitrogen is fixed in the presence of lightning as discussed above but most is fixed by bacteria that combine atmospheric nitrogen with hydrogen to make ammonia, NH3. These nitrogen-fixing bacteria usually live in the root nodules of leguminous plants providing ammonia in exchange for carbohydrates. Different soil-based bacteria convert the ammonia firstly into nitrate(III) ions, NO2–, and then into nitrate(V) ions, NO3–, which are absorbed by plants (nitrification).
Nitrogen is also fixed industrially in the Haber process which uses high temperatures and pressures to react nitrogen with hydrogen (from the gasification of methane) in the presence of an iron-based catalyst.
N2(g) + 3H2(g) ⇌ 2NH3(g) ΔrH⦵ = -92.3 kJ mol-1
The reaction is exothermic but unfavourable in terms of the overall change in entropy. In addition, the high activation energy required to break the N⫴N bond means that the rate of reaction is exceedingly slow. Le Chatelier’s principle would suggest that increasing the pressure and lowering the temperature would both shift the equilibrium position to the right. High pressures are used in the Haber process but the temperature is maintained at ∼400°C in order to increase the rate of reaction and for the catalyst to be effective.
Even so, the conversion rate is only 15%. The equilibrium mixture of gases is cooled under high pressure and the ammonia condenses. Unreacted hydrogen and nitrogen gases are then returned to the reaction vessel. The Haber process is a very high energy process overall.
Experiments show that the reaction mechanism is made up of a number of steps:
- N2(g) is adsorbed onto the surface of the catalyst
- the N⫴N is broken – this is the rate determining step
- H2(g) is adsorbed onto the surface of the catalyst
- the H-H bond is broken
- nitrogen and hydrogen atoms bond to form a molecule of ammonia which is adsorbed
- NH3(g) is released from the catalyst surface
Denitrification comprises a series of reactions undertaken by denitrifying bacteria in which nitrate(V) ions are reduced to nitrate(III) ions, then ammonium ions and finally nitrogen gas.
Manufacturing fertilisers
Around 85% of the ammonia made in the Haber process is needed for the manufacture of nitrogen based fertilisers e.g. ammonium nitrate, NH4NO3. Fertilisers are applied to soils to treat mineral deficiencies – this is the result of crops being removed before natural decay processes (ammonification) are able to return nutrients to the soil.
Ammonium nitrate is the product of the acid-base reaction between ammonia and nitric acid.
NH3(aq) + HNO3(aq) ⇾ NH4NO3(aq)
Water is evaporated from the ammonium nitrate solution and the solid compound converted into pellets.
Nitric acid is made industrially in the Ostwald process. Ammonia reacts with oxygen at a temperature of 900°C and a pressure of 810kPa in the presence of a platinum-rhodium catalyst.
4NH3(g) + 5O2(g) ⇌ 4NO(g) + 6H2O(g) ΔrH⦵ = −905 kJ mol-1
The nitrogen(II) oxide is cooled to ∼250˚C before it is reacted with more oxygen forming nitrogen(IV) oxide, which dimerises to form dintrogen tetroxide.
2NO(g) + O2(g) ⇌ 2NO2(g) ΔrH⦵ = -114 kJ mol-1
2NO2(g) ⇌ N2O4(g) ΔrH⦵ = -57 kJ mol-1
The nitrogen oxide gases are then passed through a tower to meet a stream of water which readily absorbs them. Both of the following reactions are disproportionation reactions (the oxidation state of nitrogen increases and decreases from reactants to products).
3NO2(g) + H2O(l) → 2HNO3 + NO(g) ΔrH⦵ = -117 kJ mol-1
3N2O4(g) + 2H2O(l) → 4HNO3(aq) + 2NO(g) ΔrH⦵ = -103 kJ mol-1
The NO and any unreacted NO2 and N2O4 are recycled. All of these reactions are exothermic which makes the overall process a useful heat source.
Nitric acid is also used in the manufacture of explosives, detergents, paints, dyes and nylons (polyamide polymers).
The environmental problems of nitrate fertilisers are well known. Nitrates are soluble in water and leach from soils into groundwater. This promotes the overgrowth of algae in lakes and rivers causing eutrophication.
Bonding in nitrogen oxides
Nitrogen bonds with oxygen to from a number of oxides and oxyanions, the most common of which are listed below – examiners often ask students to draw the dot/cross structure for these molecules.
- dinitrogen oxide, N2O (nitrogen has an oxidation state of +1)
- nitrogen(II) oxide, NO (nitrogen has an oxidation state of +2)
- nitrogen(IV) oxide, NO2, and dinitrogen tetroxide, N2O4 (nitrogen has an oxidation state of +4)
- nitrate(V) ion, NO3–
- nitrate(III) ion, NO2– (nitrite ion)
Practice questions
- The exhaust systems of some diesel engines use ammonia, stored in the form of urea, (NH2)2CO, to remove nitrogen(IV) oxide from the exhaust emissions.
(a) In the first stage of the reaction, urea thermally decomposes to form isocyanic acid, HNCO(g), and ammonia. Write a balanced equation for the reaction.
(b) In the second stage of the reaction, isocyanic acid reacts with water vapour to form more ammonia and carbon dioixde. Write a balanced equation for the reaction.
(c) In the third stage of the reaction, the ammonia reacts with nitrogen(IV) oxide as shown below.
8NH3(g) + 6H2O(g) ⇾ 7N2(g) + 12H2O(g)
(i) Use your answers to parts (a) and (b) in addition to the equation given above to construct the overall equation for the reduction of nitrogen(IV) oxide by urea.
(ii) Explain why this is a reduction reaction.
- Sodium azide, NaN3, is used as a propellant in vehicle airbags because it decomposes on impact producing nitrogen gas which rapidly fills the airbag.
(a) Write a balanced equation for the decomposition of sodium azide to form nitrogen and sodium.
(b) The azide ion, N3–, is isoelectronic with carbon dioxide, as shown in the resonance structure below.
(i) What does ‘isoelectronic’ mean?
(ii) Draw a dot and cross diagram for the azide ion.
- Write a balanced half equation for the conversion of nitrogen(II) oxide to dinitrogen oxide, N2O, by denitrifying bacteria under acidic conditions.
- The reaction between nitrogen and hydrogen to from ammonia is a key step in the industrial manufacture of nitric acid.
N2(g) + 3H2(g) ⇌ 2NH3(g) ΔrH⦵ = -92 kJ mol-1
(a) Suggest why this reaction is described as a homogeneous equilibrium.
(b) le Chatelier’s principle can be used to predict how different conditions affect the equilibrium position in the reaction. Explain how changing pressure, temperature and using a catalyst affect the equilibrium yield of NH3 and the overall rate of reaction.
Answers
- (a) (NH2)2CO(g) ⇾ HNCO(g) + NH3(g)
(b) HNCO(g) + H2O(g) ⇾ NH3(g) + CO2(g)
(c) (i) 4(NH2)2CO(g) + 6NO2(g) ⇾ 7N2(g) + 8H2O(g) + 4CO2(g)
(ii) The oxidation state of nitrogen is reduced from +4 in NO2 to 0 in N2.
2. (a) 2NaN3(s) ⇾ 2Na(s) + 3N2(g)
(b) (i) Isoelectronic means that N3– has the same number of electrons as CO2.
(ii)
3. 2NO(g) + 2H+(aq) + 2e– ⇾ N2O(g) + H2O(l)
4. (a) In a homogenous equilibrium the reactants and products are in the same physical state.
(b) le Chatelier’s principle states that when an external change is made to a system in dynamic equilibrium, the system responds to minimise the effect of the change.
There are more moles of gases on the left than on the right side of the equation so increasing the pressure will cause the equilibrium position to shift to the right to the side with fewer moles of gas, increasing the yield of ammonia (increasing the pressure will also increase the rate of the reaction as there will be more collisions per unit time).
The forward reaction is exothermic so decreasing the temperature would cause the equilibrium position to shift to the right, in the exothermic direction, increasing the yield of ammonia (however decreasing the temperature would decrease the rate of the reaction as fewer collisions would have the necessary activation energy).
Using a catalyst has no effect on the equilibrium position but equilibrium is achieved faster.