NitrogenThe lack of reactivityWe know that nitrogen is an unreactive gas because it is mixed with reactive oxygenin Earths atmosphere and reacts very little. The reason for this lack of reactivityis the very strong NN bond in the molecule. The two nitrogen atoms share threeelectron pairs, which form a triple covalent bond, and each atom retains a lone pairof electrons. This is shown in Figure 12.1. The bond energy for nitrogen is 946 kJ mol1compared with the bond energy in fluorine, F2, which is 158 kJ mol1.Hodder CIE revision guide 2010Chemistry g 12.126 Aug 2010Eleanor JonesN NFigure 12.1 Structure of nitrogenAlthough nitrogen is unreactive, it does react under the right conditions. For example,burning magnesium reacts with the nitrogen in air to form magnesium nitride:3Mg(s) N2(g) Mg3N2(s)At high temperatures nitrogen will react with oxygen to form oxides of nitrogen, forexample:N2(g) O2(g) 2NO(g)In car engines (and in thunderstorms) nitrogen combines with oxygen to produce amixture of oxides of nitrogen often referred to as NOx.Nitrogen also reacts with hydrogen to form ammonia:N2(g) 3H2(g) 2NH3(g)This is the basic reaction in the Haber process for the manufacture of ammonia. Thisis covered in Chapter 7 and in more detail later in this section.Finally, the roots of some plants of the pea and bean family have nodules thatcontain bacteria able to fix nitrogen chemically. The bacteria convert the nitrogeninto ammonium ions, which the plants can use to make proteins.11412 Nitrogen and sulfurAmmoniaAmmonia, NH3, is an alkaline gas. You might expect it to have a trigonal planarshape, but it is pyramidal with a lone pair of electrons occupying the apex, as isshown in Figure 12.2.Hodder CIE revision guide 2010Chemistry g. 12.226 August 2010Eleanor JonesHNHH 107Figure 12.2 Structure of ammoniaRemember that the lone pair takes up more space than a pair of bonding electrons.In the case of ammonia this reduces the H N H bond angles to about 107. Thehigh electronegativity (3.0) of nitrogen means that ammonia can form hydrogenbonds, and as a result it is very soluble in water.Ammonia is also a base, accepting a proton to form the ammonium ion, NH4:NH3(g) H(aq) NH4(aq)This ion is tetrahedral, the proton forming a coordinate (dative) bond with thenitrogen atom using the lone pair (Figure 12.3).Hodder CIE revision guide 2010Chemistry g 12.326 August 2010Eleanor JonesHHNHHFigure 12.3 Structure of the ammonium ionIn the laboratory, ammonia is easily displaced from ammonium compounds bywarming with a strong base such as sodium hydroxide:NH4(aq) OH(aq) NH3(g) H2O(l)The Haber processWe have already looked at some aspects of the Haber process in Chapters 7 and 8.You may wonder why you have to study this manufacturing process. Ammonia isone of the most important bulk chemicals manufactured, mainly because of its usein the production of fertilisers. In 2008, more than 150 million tonnes of ammoniawere produced worldwide.The original Haber process was developed almost 100 years ago and still forms thebasis of ammonia production, although the pressures and temperatures have beenrefined to give maximum production at an optimum price. Most recent researchhas focused on improving the iron catalyst. Both osmium and ruthenium are moreeffective as catalysts, but both are more expensive than iron. Therefore, the catalystused depends on the market price of ammonia.Cambridge International AS and A Level Chemistry Revision Guidecontent guidance115For the examination you need to be able to state the essential operating conditionsfor an ammonia plant (Figure 12.4), and to interpret these in terms of reactionkinetics and the position of equilibrium.Hodder CIE revision guide 2010Chemistry g 12.426 August 2010Eleanor JonesCompressorN2(g) 3H2(g)CondenserCatalystReactor(400 C, 200 atm)NH3(l)storageUnreactedN2(g) and H2(g)Figure 12.4 The Haber processThe set of conditions chosen are a compromise: pressure around 20 000 kPa (or 200 atm) temperature around 670 K catalyst iron (which has a lifetime of around 5 years)As well as its use in fertiliser manufacture (mainly as the sulfate or nitrate), ammoniacan be oxidised using a platinum catalyst to form nitrogen monoxide:4NH3(g) 5O2(g) 4NO(g) 6H2O(l) Ho 909 kJ mol1This exothermic reaction is the starting point for the manufacture of nitric acid.The hot gas is cooled, reacted with more oxygen to form nitrogen dioxide, and thendissolved in water to form nitric acid:2NO(g) O2(g) 2NO2(g)2NO2(g) N2O4(g)3N2O4(g) 2H2O(l) 4HNO3(aq)Nitrogen compounds and pollutionNitratesNitrogen compounds, particularly nitrates, are used to make fertilisers. These havehad significant benefits in increasing crop yields around the world. However, ifexcessive amounts of fertilisers are used they can get washed into streams andrivers. This can have two effects:11612 Nitrogen and sulfur The nitrates can get into drinking water supplies, from which they are difficultto remove. They can affect the ability of babies under 6 months to carry oxygenin the bloodstream. They can increase the growth of aquatic vegetation, which then decays reducingoxygen levels in streams and rivers, and affecting other forms of aquatic life.Nitrogen oxidesThe combustion of motor fuels produces temperatures high enough to form oxidesof nitrogen, NOx, in exhaust gases. These gases in air at street level are known tocontribute to respiratory problems. They also have a pollution effect in the upperatmosphere where they catalyse the oxidation of sulfur dioxide. The main sourceof sulfur dioxide in the atmosphere is the combustion of fuels (mainly coal and oil)that contain sulfur or its compounds. Some of the sulfur dioxide is removed from thewaste gases emitted by major users of these fuels, such as power stations.The exact mechanism is uncertain, but the following is a possibility:2NO2(g) H2O(g) HNO2(l) HNO3(l)SO2(aq) HNO3(l) NOHSO4(l)NOHSO4(l) HNO2(l) H2SO4(l) NO2(g) NO(g)SO2(aq) 2HNO2(l) H2SO4(l) 2NO(g)The other form of atmospheric pollution caused by nitrogen oxides is the formationof ozone, O3, in the presence of unburnt hydrocarbons from vehicles. Ozone formedat street level can cause respiratory problems. However, once nitrogen oxides escapeinto the upper atmosphere they cause the reverse problem, removing ozone andreducing the ability of ozone to protect the Earths surface from harmful ultravioletradiation.The emission of nitrogen oxides from vehicles has been greatly reduced since theintroduction of catalytic converters. These also reduce the emission of carbonmonoxide and unburnt hydrocarbons (see also Chapter 8).2NOx(g) xO2(g) N2(g)2CO(g) O2(g) 2CO2(g)CnH2n2(g) [(3n 1)/2]O2(g) nCO2(g) (n 1)H2O(g)SulfurThe Contact processSulfuric acid is an extremely important industrial chemical that is used in theproduction of a huge range of manufactured goods. Figure 12.5 shows some of themajor uses of sulfuric acid. In 2008, world production of sulfuric acid was more than200 million tonnes.Cambridge International AS and A Level Chemistry Revision Guidecontent guidance117Hodder CIE revision guide 2010Chemistry g 12.526 August 2010Eleanor JonesUses of sulfuric acidPetroleumrefining20%Fertilisermanufacture30%Chemicalsincluding(NH4)2SO416%Picklingsteel 10.6%Storagebatteries10.2%Pigments, textiles 4.6%Explosives 2.7%Other uses 5.9%Figure 12.5 Uses of sulfuric acidFor the examination you need to be able to state the essential operating conditionsfor the manufacture of sulfuric acid by the Contact process (Figure 12.6), and tointerpret these in terms of reaction kinetics and the position of equilibrium.SulfurburnerSO2 SO3converterAbsorptiontowerOleumstorage fordilution toH2SO4S(l) ExhauststackAirSO2 inlet Air inletConcentratedsulfuric acidCooling SO2 recycledCoolingCoolingFigure 12.6 The Contact processThe conditions chosen are a compromise: pressure around 100200 kPa (or 12 atm) temperature around 700 K catalyst vanadium(V) oxide (although some plants now use ruthenium whichis more expensive)11813 Introduction to organic chemistrySulfur dioxideAcid rainWe have already seen that the combustion of sulfur-containing fuels releases sulfurdioxide into the atmosphere, and that in the presence of oxides of nitrogen, this canbe converted into sulfuric acid. This decreases the pH of rain. Acid rain can harmplants and animals directly, and indirectly by making lakes acidic. It also releasestoxic metals such as aluminium from soils, and below pH 4.5 no fish are likely tosurvive. This has further effects up the food chain. Acid rain also increases theerosion of limestone-based buildings and statues. The production and some effectsof acid rain are shown in Figure 12.7.Hodder CIE revision guide 2010Chemistry g 12.76 September 2010Eleanor JonesAcid rainPower stations,industries and vehiclesrelease gases intothe atmosphereSulfur dioxide and nitrogen oxidesOxidation in clouds producessulfuric acid and nitric acidAcid polluted airblown by the windWater dropletsbecome pollutedWet deposition acid rainTrees andvegetation affectedSoils polluted and acid rainleaches aluminium from soilsRivers and lakes pollutedaffecting water lifeFigure 12.7 Acid rainFood preservationAlthough sulfur dioxide is a pollutant that causes widespread damage to the environment, it also has beneficial uses, particularly as a preservative for foodstuffs. Ithas the advantage of being a reducing agent (preventing spoilage by oxidation) andhas anti-microbial properties. It is used in the preservation of fruit products (suchas dried fruits and juices) and in winemaking. It can be a problem for people whosuffer from asthma.