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9.3 The acidic environment: 2. Acid oxides in the atmosphere

Extract from Chemistry Stage 6 Syllabus (Amended October 2002). © Board of Studies, NSW.
[Edit: 11 Jun 10]

Prior learning: Preliminary modules 8.3.3.

Background: Just as elements show a pattern in properties in the Periodic Table so metal oxides and non-metal oxides show a pattern in properties. Metal oxides are usually basic and non-metal oxides usually acidic. The extent of the acidity or basicity of an oxide can often be predicted from the element's position in the Periodic Table.

identify oxides of non-metals which act as acids and describe the conditions under which they act as acids

• Carbon dioxide (CO₂ ), sulfur dioxide (SO₂) and nitrogen dioxide (NO₂) all dissolve in water forming acid solutions. Most non-metal oxides (except for CO, NO and N₂O which are neutral) are said to be acidic.
  
  
• To detect that a non-metal oxide gas is acidic with indicator paper, the paper must be moist. Moisture enables the gas to dissolve and form the acid that produces hydrogen ions. Reaction of a hydrogen ion with an indicator causes the colour change.

analyse the position of these non-metals in the Periodic Table and outline the relationship between position of elements in the Periodic Table and acidity/basicity of oxides

• When the Periodic Table outline for oxides below is compared with the information in a full Periodic Table of the type used for HSC exams, it can be seen that:
  • metal oxides are mostly basic
  • non-metal oxides are mostly acidic
  • oxides of the five elements close to the borderline between metals and non-metals are amphoteric, that is, they show both acidic and basic properties.

    

define Le Chatelier's principle

• In 1885, the French chemist, Le Chatelier, put forward a principle for predicting the effect of change on reversible reactions:

  The concentrations of reactants and products in a mixture at equilibrium will alter so as to counteract any change in concentration, temperature or gas pressure.

identify factors which can affect the equilibrium in a reversible reaction

• The factors that can affect the equilibrium in a reversible reaction are:
  • change in concentration
  • change in temperature
  • change in gas pressure.

identify data, plan and perform a first-hand investigation to decarbonate soft drink and gather data to measure the mass changes involved and calculate the volume of gas released at 25^oC and 100kPa

• In order to carry out a valid investigation, you will need to plan the experimental design to remove the carbon dioxide from a quantity of soft drink in a way that allows you to make accurate and reliable measurements. Variations on two basic methods could be used successfully in a school or domestic situation: a warming method, that uses the drop in water solubility of a gas on heating, and a salting-out method, that uses the stronger attraction soluble salt ions have for water than dissolved gas molecules have for water. These are described generally below.
  
  
• In planning your investigation, consider the ideas presented below and, from them, identify the types of data you will need to gather during the investigation and be able to explain the quantitative analysis that will be required for this data to be useful.
  
  
• Perform the investigation, making modifications as needed and analysing the effect of these adjustments.
  
  
• Measure and record the quantitative data required for the calculations to determine the volume of gas released at 25^oC and 100kPa. Carry out repeat trials to ensure reliable data is collected.

describe the solubility of carbon dioxide in water under various conditions as an equilibrium process and explain in terms of Le Chatelier's principle

• The solubility of carbon dioxide gas in water can be fully described using four equilibrium equations:
  
  
  1. CO_2(g) CO_2(aq)
     
     
  2. H₂O_(l) + CO_2(aq) H₂CO_3(aq)
     
     
  3. H₂CO_3(aq) H⁺_(aq) + HCO₃^-_(aq)
     
     
  4. HCO₃^-_(aq) H⁺_(aq) + CO₃^2-_(aq)
     
     
• An equilibrium shift to the left or right in any of the equations above leads to a change in the concentration of the reactants and products in that reaction. This influences the concentrations in the other three reactions causing the equilibrium in those reactions to also shift to the left or right respectively. For example, a decrease in the pressure above the solution causes reaction 1 to shift to the left and results in a decrease in the pH of the solution as the other equlibrium reactiuons also shift to the left. Likewise, an equilibrium shift to the right in equations 2, 3 or 4 results in more carbon dioxide gas being dissolved in equation 1.


• In accordance with Le Chatelier's principle:
  
  
  • addition of acid (increased concentration of H⁺) shifts equilibrium to the left in equations 3 and 4 above and therefore results in a decrease in dissolved carbon dioxide.
    
    
  • addition of base (reacts with and reduces the concentration of H⁺) shifts equilibrium to the right in equations 3 and 4 above and therefore leads to an increase in the concentraction of dissolved carbon dioxide.
    
    
  • addition of a soluble carbonate (increased concentration of CO₃^2-) shifts equilibrium in equation 4 above to the left and therefore results in a decrease in dissolved carbon dioxide.

calculate volumes of gases given masses of some substances in reactions, and calculate masses of substances given gaseous volumes, in reactions involving gases at 0^oC and 100kPa or 25^oC and 100kPa

• Calculations require the writing of balanced equations and use of mole relationships:

  Example 1: Calculate the volume of carbon dioxide released at 100kPa and 25^oC by the reaction of 10.0 g of calcium carbonate with excess acid.

  From the equation above, it can be seen that one mole of CaCO₃ releases one mole of CO₂.

  10.0 g CaCO₃   =    10.0 g/100 g mol^-1     =    0.100 mol

  Thus 0.100 mol of CO₂ gas is released = 0.100 x 24.8 L = 2.48 L

  Example 2: If 1.00 L of hydrogen sulfide gas was collected at 101.3 kPa and 0^oC from the reaction of excess acid with iron(II) sulfide, calculate the mass of FeS.

  From the equation, it can be seen that one mole of FeS produces one mole of H₂S.

  1.0 L gas at 100kPa and 0^oC is 1.00/22.7 mol = 0.0441 mol

  Thus, 0.0441 mol of FeS reacted = 0.0441 mol x 87.9 g mol^-1 = 3.87 g

identify natural and industrial sources of sulfur dioxide and oxides of nitrogen

• Some coal or oil reserves contain considerable quantities of sulfur compounds. Most sulfur dioxide (SO₂) released into the atmosphere comes from the burning of such coal or oil in electric power stations. SO₂ is the major contributor to acid rain that can affect places thousands of kilometres from the source. The low sulfur content of Australian coal is one reason why Australia is the world's largest exporter of coal.
  
  
• Smelters where metal sulfides are heated in air to remove the sulfur as SO₂ are becoming a smaller source of atmospheric SO₂. Most of the SO₂ produced is now used to make sulfuric acid. The huge amount of SO₂ that used to be released from Mt Isa mine's smelter in Queensland is now changed to sulfuric acid and transported by rail 150 km away to make ammonium sulfate and superphosphate fertilisers at Phosphate Hill.
  
  
• Volcanoes are an unpredictable source of SO₂. After Mt Pinatubo erupted in the Philippines in 1991, oxidation of the emitted SO₂ and reaction with water formed an aerosol of sulfuric acid droplets. This aerosol reduced global temperatures by about 0.5^oC.

National Pollutant Inventory Database for Australia 2008/2009 data within Australia - Sulfur dioxide from All Sources, Department of Environment, Water, Heritage and the Arts, Australia.

The figures for Oxides of Nitrogen Department of Environment, Water, Heritage and the Arts, Australia.

describe using equations, examples of chemical reactions which release sulfur dioxide and chemical reactions which release oxides of nitrogen

• In a metal sulfide smelter, the ore is heated in air and converts to a metal oxide, releasing sulfur dioxide.

  e.g. 2ZnS + 3O₂ 2ZnO + 2SO₂

• When coal, petroleum or natural gas are burnt, sulfur in sulfur compounds is converted to sulfur dioxide.

  S + O₂ SO₂

• Lightning strikes cause reaction between the two most common gases in the atmosphere.

  N₂ + O₂ 2NO

• High temperature combustion reactions in furnaces and internal combustion engines produce significant amounts of NO [called nitrogen monoxide, nitric oxide or nitrogen(II) oxide] above 1300^oC.

  N₂ + O₂ 2NO

• Colourless, neutral nitrogen monoxide reacts with oxygen in the air to form brown, acidic nitrogen dioxide [nitrogen(IV) oxide].

  2NO + O₂ 2NO₂

analyse information from secondary sources to summarise the industrial origins of sulfur dioxide and oxides of nitrogen and evaluate reasons for concern about their release into the environment

• Information about sources of acid rain, mostly produced by SO₂ emissions and to a lesser extent by NOx emissions, can be found by using Internet search engines or by going to sites for scientific magazines and journals and using site search engines. In choosing equipment, determine the suitability and effectiveness of your browser search facilities to locate the relevant information.
  
  
• Some suitable starting points from which to gather information include:
  
  
• search engines, such as Google
  
  
• popular scientific journals, such as Nature New Scientist
  
  
• science specific Internet sites, such as Nova and Science Daily

Specific information on industrial origins of pollutants in Australia can be found in the National Pollutant Inventory Database. Reasons for concern about release of these gases into the environment can be found by:

• searching the Internet using a search engine
  
  
• accessing material safety data sheets (MSDS) for these gases using CD-ROMs, such as the one supplied to schools with the Chemical Safety In Schools (CSIS) package.
  
  
• analyse the information you access to evaluate the reasons for concern about the release of sulfur dioxide and the oxides of nitrogen into the environment. Your evaluation might describe cause and effect relationships and could be based on criteria such as economic, social cost, health effects, environmental, political and energy benefit. List the criteria or values you use in your evaluation.

assess the evidence which indicates increases in atmospheric concentration of oxides of sulfur and nitrogen

• There is extensive evidence for an increase of over 25% in atmospheric carbon dioxide levels over the last two hundred years. The evidence comes from quantitative analysis of trapped air bubbles in Antarctic ice and measurement of carbon isotopes in old trees, grass seeds in museum collections and calcium carbonate in coral.
  
  
• Finding evidence for increases in atmospheric sulfur oxides and nitrogen oxides is more difficult for the following reasons:
  1. Whereas atmospheric CO₂ concentrations are about 360 parts per million (ppm), the levels for SO₂ and NO_x are only about 0.001 ppm in populated parts of the Earth.
  2. The chemical instruments able to measure very low concentrations, like those for SO₂, have only been commercially available since the 1970s.
  3. CO₂ changes to carbonate ions when it dissolves in water and most carbonates are insoluble. Seashells and coral are made up of carbonates that came from atmospheric CO₂. Isotope ratio measurements using mass spectrometers on shells and corals of different ages give clues as to past atmospheric CO₂ concentrations.

     On the other hand, SO₂ eventually forms sulfate ions and NO₂ forms nitrate ions. Most sulfates and all nitrates are water-soluble. Soluble sulfates and nitrates circulate in the hydrosphere and biosphere and are chemically changed while insoluble carbonates tend to stay in inert forms such as shells or coral.

• Some evidence that has been observed that indicate the increasing atmospheric concentrations of oxides of sulfur (SO_x) and nitrogen (NO_x) are:
  1. Higher atmospheric concentrations of SO_x and NO_x have been detected in industrial areas, particularly in Europe. These correspond to an increase in acidity in lakes and rivers in these areas.
  2. Damage to forests, buildings and aquatic life has also been observed in industrialised areas. This is thought to be due to acid rain formed by elevated levels of atmospheric SO_x and NO_x.


• Information can be found at http://www.apis.ac.uk/ that indicates the atmospheric concentrations of oxides and nitrogen have been decreasing in recent decades. Acces this and other souces of information to make your assessment. Make sure you assess the validity and reliability of the secondhand sources you access and make your final assessment based on the most valid and reliable information.

explain

• Distilled water in contact with the atmosphere is not neutral. It has a pH of 5.5 to 6, due to absorption of the acidic gas CO₂ from the atmosphere.
  
  
• In Australia, unpolluted rainwater has a pH between 5 and 6. If the pH is below 5, an acidic substance, such as SO₂ or NO₂, has dissolved in the water, which is sometimes called acid rain. In the Northern Hemisphere, pHs as low as 2 have been recorded in acid rain. The source of the SO₂ or NO₂ could be hundreds or thousands of kilometres from where the acid rain falls.
  
  
• SO₂ sources, such as fossil fuel burning power stations and metal sulfide smelters, are larger but fewer in number than NO₂ sources, like internal combustion engines in vehicles.
  
  
• If the quantity of acid rain is greater than the capacity of an environment to neutralise it then the following can occur:
  1. Soil pH can drop, making it difficult for plants to absorb sufficient calcium or potassium.
  2. Soil chemistry can change, leading to the death of important micro-organisms and release of normally insoluble aluminium and mercury into soil water.
  3. Protective waxes can be lost from leaves, causing leaf damage.
  4. Buildings made of carbonates, such as concrete, mortar, limestone and marble, can be gradually dissolved away.
  5. Aquatic animals can die as water acidity drops below pH 5.
  6. Smog and acid rain can combine to form killer fog, as happened after the Second World War in London, when many homes burnt sulfur dioxide-releasing coal.

 Acid rain effects described United States Environmental Protection Agency, US

(These web sites last checked 27 June 2008)

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