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12. Nitrogen and sulfur
A section of Chemistry, 9701
Listing 10 of 350 questions
Release of sulfur dioxide, SO2, into the atmosphere causes acid rain. Explain why high levels of SO2 may be found in the atmosphere near power stations that burn fossil fuels. The SO2 released can be converted in the atmosphere into sulfuric acid, H2SO4, by reaction with nitrogen dioxide gas, NO2, and water, H2O. SO2 + NO2 NO + SO3 SO3 + H2O H2SO4 During one year, 1590 tonnes of SO2 was released into the atmosphere by a fossil-fuel burning power station. Use the equations to calculate how many tonnes of H2SO4 were formed in the atmosphere. Assume that all of the SO2 released was converted into H2SO4.  mass of H2SO4 = tonnes Describe how NO2 is also produced by these power stations. State one natural cause of NO2 being formed in the atmosphere. Explain why NO2 can be described as a catalyst in the oxidation of atmospheric SO2. 
9701_s19_qp_23
THEORY
2019
Paper 2, Variant 3
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Nitric acid, HNO3, can be made by reacting nitrogendioxide with water. The enthalpy change for the reaction can be measured indirectly using a Hess’ cycle. ∆Hr 3NO2+ H2O2HNO3+ NOExplain what is meant by the term enthalpy change of formation. Complete the Hess’ cycle using the values given in the table and hence calculate the enthalpy change, ∆Hr, for this reaction. Show your working. substance ∆Hf / kJ mol–1 NO234.0 H2O–286 HNO3–173 NO91.1 ∆Hr 3NO2+ H2O2HNO3+ NO∆Hr = kJ mol–1  Nitrogen and oxygen do not react at normal atmospheric temperatures. Explain why. Nitrogen oxides can be formed naturally in the Earth’s atmosphere from nitrogen and oxygen in the air. State one way that nitrogen oxides are produced naturally. Nitrogen dioxide, NO2, acts as a homogeneous catalyst in the oxidation of atmospheric sulfurdioxide. Explain why NO2 is described as a homogeneous catalyst. Write equations which describe the two reactions occurring when NO2 acts as a catalyst in the formation of sulfurtrioxide from sulfurdioxide. 
9701_s20_qp_21
THEORY
2020
Paper 2, Variant 1
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Calciumnitrate, Ca(NO3)2, reacts with ammonia, carbondioxide and water to form a mixture of ammoniumnitrate and calciumcarbonate. Ca(NO3)2 + 2NH3 + CO2 + H2O → 2NH4NO3 + CaCO3 Explain why ammonia is described as a Brønsted‑Lowry base in this reaction. The product mixture can then be added to soil. State two reasons why this mixture of products is added to some soils.  Complete the table to name the shape and give the bond angle of each species. name of shape bond angle / ° CO2 NH3 H2O  
9701_s20_qp_21
THEORY
2020
Paper 2, Variant 1
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Nitric acid can be made in a 3-stage process. Stage 1 Ammonia is oxidised by oxygen from the air, to form nitrogenmonoxide and water. This reaction is carried out at 10–13 atmospheres pressure and 900 °C in the presence of a platinum catalyst. Stage 2 Nitrogenmonoxide reacts with more oxygen to form nitrogendioxide. 2NO + O2 2NO2 ∆H = –114 kJ mol–1 Stage 3 Nitrogendioxide reacts with water to make nitricacid and nitrogenmonoxide. 3NO2 + H2O → 2HNO3 + NO Write an equation to show the reaction occurring in stage 1. Draw a ‘dot-and-cross’ diagram to show the arrangement of outer electrons in a molecule of ammonia.  In the boxes, give the oxidation numbers of nitrogen in the nitrogen-containing species for the reaction in stage3. 3NO2 + H2O → 2HNO3 + NO  Explain why the reaction in stage3 is described as a disproportionation reaction. Include reference to transfer of electrons in your answer. The release of nitrogenmonoxide into the atmosphere causes atmospheric pollution. State and explain the effect of nitrogenmonoxide gas in contact with moist air. The nitricacid made in stage3 can then be reacted with ammonia to form ammoniumnitrate. Stage 3 3NO2 + H2O → 2HNO3 + NO NH3 + HNO3 → NH4NO3 Calculate the volume of nitrogen dioxide, measured at room temperature and pressure, required to make 40tonnes of ammoniumnitrate. [1 tonne = 1000 kg] Show your working.  volume of nitrogen dioxide = State one use of ammonium nitrate. 
9701_s20_qp_22
THEORY
2020
Paper 2, Variant 2
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Questions Discovered
350