CHEMISTRY

Acidic Environment Assignment Part 2 1) Compare the structure and bonding of the following: (a) , (b) , and (c) sulphur dioxide

Diagram Structure Bonding

An ionic lattice. Each unit cell exhibits the anti‐ fluorite structure. Ionic bonding The anions (O2‐) are

Na2O in the face‐centred

cubic array with the cations (Na+) in all

the tetrahedral Red – O2‐ holes. Purple – Na+

Double covalent bonds join two SiO2 Covalent network oxygen atoms to each silicon atom.

Red – Silicon Black – Oxygen

Double covalent bonds join two oxygen atoms to Polar covalent a sulphur atom. SO2 molecule with bent Due to polarity, structure. there is dipole‐ dipole interaction between gaseous SO2 molecules.

2) Write equations for the reaction of the following with water:

COMPOUND REACTION WITH WATER

Carbon dioxide CO2(g) + H2O(l) H2CO3(aq) Sodium oxide Na2O(aq) + H2O(l) 2NaOH(aq) Calcium oxide CaO(aq) + H2O(l) Ca(OH)2(aq) SO2(g) + H2O(l) H2SO3(aq) SO3(g) + H2O(l) H2SO4(aq) dioxide NO2(g) + H2O(l)  HNO2(aq) + HNO3(aq)

3) Beryllium oxide is amphoteric. (a) Explain what is meant by amphoteric, and (b) Study the two equations below*. Balance them, and indicate whether BeO is acting as an or (a) ‘Amphoteric’ is a term used to describe a substance that exhibits both acidic and basic properties. Beryllium oxide is an amphoteric oxide that reacts with strong and strong bases. 2+ ‐ (b)*BeO(s) + 2HCl(aq) + 3H2O(l)  Be(H2O)4 (aq) + 2Cl (aq) BeO acting as a base 2+ + *BeO(s) + 2NaOH(aq) + H2O(l)  Be(OH)4 (aq) + 2Na (aq) BeO acting as an acid

4) Describe the origins of sulfur dioxide that are causing environmental problems

 The oxidation of hydrogen sulfide (H2S) which is a product of bacterial decomposition

 2H2S (g) + 2O2 (g)  2SO2 (g) + 2H2O (g)

 The burning of fossil fuels which usually contain sulfide minerals like FeS2. These sulfide minerals in coal are oxidised when the fuel is combusted and sulfur dioxide is released. For example,

 4FeS2 (s) + 11O2 (g)  2Fe2O3 (s) + 8SO2 (g)  Metal smelters that convert metal sulfides into metals yield vast amounts of sulfur

dioxide. For example, smelting chalcopyrite (CuFeS2) to obtain copper results in the

release of SO2

 2CuFeS2 (s) + 5O2 (g) + 2SiO2 (s)  2Cu (l) + 4SO2 (g) + 2FeSiO3 (l) “In Adelaide in 2000, petroleum refinery processing produced 40% of the sulphur dioxide emitted to the atmosphere, with motor vehicles contributing 22% and fuel combustion 15%.” 1

5) Describe the origins of oxides of nitrogen that are causing environmental problems  Nitric oxide is formed when nitrogen and oxygen react at high temperatures, for example, during lightning strikes or high temperature combustion reactions in furnaces and internal combustion engines

 N2 (g) + O2 (g)  2NO (g)  Nitric oxide, a colourless gas and neutral oxide, reacts with oxygen to form nitrogen dioxide, a brown gas and an acidic oxide

 2NO (g) + O2 (g)  2NO2 (g)

“In Adelaide in 2000, an estimated 66% of nitrogen oxides (including NO and NO2) came from motor vehicles, with a further 20% from fuel combustion.” 1

6) (a) Use www.deh.gov.au/soe/2001/atmosphere/atmosphere02‐16.html to describe the change in acidic oxides in the atmosphere using ice cores from Antarctica. Download the relevant graphs and analyse them (b) Describe the importance of data from ice cores (a) Ice core data obtained from Antarctica (refer to graph below) indicates that the atmospheric concentration of has increased since AD 1000.

It can be seen that prior to the 1800s, CO2 (ppm) showed relatively minor fluctuation ranging from approximately 275ppm‐285ppm. This range may represent the normal

concentration of CO2 that works in conjunction with other atmospheric constituents such as methane, nitrous oxide and water vapour to contribute to a natural greenhouse effect. During and after the 1800s, the data from the ice core samples show a very steep

increase in CO2 (ppm), from approximately 285ppm in 1800 to 330ppm before the year 2000. This dramatic increase is accounted for by the occurrence of the Industrial Revolution in the 1800s. The burning of fossil fuels to provide power released vast amounts of carbon dioxide and other combustion products into the atmosphere. Since the 1800s, technological advances have placed greater demands on industry and transport in particular, in terms of the generation of power by burning coal. Consequently, in the recent years, there has been a rapid increase in the concentration of atmospheric carbon dioxide.

Carbon dioxide concentration from ice core and air samples since AD 1000.Source: CSIRO Atmospheric Research.

(b) Ice core data can be used to

“reconstruct past climate conditions and climate changes through time...an ice core paleoclimate record can be compared to and combined with other paleoclimate records, for example, from ocean sediment cores, tree rings, coral records and spleotherm (stalactites and stalagmites) records to establish the Earth’s climate history. This climate history can then be used to establish the limits of natural climate variability which can be used to constrain climate models and to help us reduce uncertainty in future climate predictions.” (Secrets from Antarctic Ice, Mark Curran)

In other words, data from ice cores is important as it allows us to not only unlock the composition of the ancient atmosphere, but also to understand how our past actions have affected the climate. By understanding causal relationships, such as increased global temperature due to an increase in the concentration of atmospheric carbon dioxide, which in turn is largely due to the burning of fossil fuels, we can forecast climate change based on models of the climate constructed using evidence collected

from ice core analysis, such as gas composition of trapped air to determine CO2 levels. Strategies can then be imposed to rectify dim prospects if our present actions are predicted to have an adverse impact on the atmosphere. For example, as ice core research has revealed a link between greenhouse gases and climate change in the past, we need to reduce greenhouse gas emissions, in particular, carbon dioxide.

7) Look up http://www.rta.nsw.gov.au/constructionmaintenance/completedprojects/m5east/m5east currentairqualitydata/currentdata.html (a) Identify the gases that are monitored in the M5 tunnel, (b) Explain why these gases are monitored, and (c) What will happen if the gases in the tunnel exceed the ‘safe’ limit? (a) Carbon monoxide and nitrogen dioxide (b) “As a condition of approving the M5 East Freeway project, the Department of Planning placed strict air quality limits on the operation of the tunnel. As a result, air quality is closely monitored and must continually conform within the limits set.” Air quality limits must not be exceeded as accumulation of the monitored gases will have adverse health effects on users of the tunnel. Both carbon monoxide and nitrogen dioxide are toxic to humans. When inhaled, carbon monoxide binds with haemoglobin to form carboxyhaemoglobin. This reduces an individual’s oxygen‐carrying capacity. At low concentrations, this causes mild headaches and fatigue. These symptoms intensify as concentration increases and respiratory complications become evident.

At high concentrations, there is the risk of unconsciousness, convulsions, coma and death. Exposure to nitrogen dioxide at low concentrations results in irritation of the eyes, nose, lungs and throat causing coughing, shortness of breath, fatigue and nausea. Existing respiratory complications such as asthma and emphysema are also aggravated. Particularly susceptible to the health effects of the two gases are young children, the elderly and pregnant women. The build‐up of nitrogen dioxide, a red‐brown gas, also contributes to the formation of haze in the tunnel and reduces visibility. (c) Vehicle emission levels are under constant monitoring in the M5 tunnel. If the concentrations of the gases in the tunnel exceed the ‘safe’ limit, tunnel operators ventilate the tunnel, venting gases from the portals (entrances and exits) to reinstate air quality standards.

8) Explain what is being done by governments and industry to reduce the release of oxides of sulphur and nitrogen In order to reduce emissions of oxides of sulphur and nitrogen, the Australian Government has developed a balanced mix of policy responses by targeting industry, households, governments and communities. Australia is also currently developing a climate change forward agenda to cover the next 20‐30 years. A few examples of the strategies undertaken by the government and industry to reduce emissions include:  The Greenhouse Gas Abatement Program (GGAP) ‐ assists Australia in meeting its Kyoto Protocol target. The objective is to reduce Australia's net greenhouse gas emissions by supporting activities that are likely to result in substantial emission reductions or substantial sink enhancement $400 million has been allocated to the Program.  A commitment to increase the use of clean renewable energy in Australia. For example, legislation that requires the generation of 9,500 gigawatt hours of extra renewable electricity per year by 2010, enough power to meet the residential electricity needs of four million people. This initiative is being achieved by establishing an innovative market in renewable energy certificates and is expected to deliver in excess of $2 billion of investment in renewable energy in Australia;  Working to increase the use of alternative fuels such as compressed natural gas (CNG) and liquefied petroleum gas (LPG) and to improve consumer awareness of the fuel efficiency of their motor vehicles.  Regulating the energy efficiency of equipment and many appliances used by Australian households and businesses. This reduces the amount of coal being burnt.  Advocating that power companies burn low‐sulfur coal rather than high‐sulfur coal – they can also switch to natural gas which produces very little sulfur dioxide on combustion.  Reducing reliance on fossil fuels by seeking alternate forms of power production such as hydroelectricity, solar power and wind power.  Collecting the sulfur dioxide produced by smelting metal sulfides and using it to make sulphuric acid

 Reducing acidic emissions from smoke stacks by a process called “scrubbing” – the acidic gases are passed through a slurry of a base such as CaO; the sulfur dioxide reacts with the CaO to form solid calcium sulphite

 Ensuring that exhausts from motor vehicles pass through a catalytic converter – NOx

converted back into nitrogen gas – 2NO (g) + 2CO (g)  N2 (g) + 2CO2 (g)

9) Assess (make a judgement of value, quality, outcomes, results or size) the evidence which indicates increases in atmospheric concentration of oxides of sulfur and nitrogen. Include arguments for and against the statement; has human activity increased the levels rather than natural sources? Have some areas shown decrease in concentration of these oxides? Why? Make a clear judgement

10) Identify the causes of . Include equations Acid rain is caused by pollution of the atmosphere by acidic oxides such as sulphur dioxide and oxides of nitrogen which dissolve in rainwater to produce solutions of various acids. Sources of these pollutant emissions multiplied during and after the Industrial Revolution of the 1800s and early 1900s. The following are examples of how acidic oxides are released into the atmosphere, where they contribute to the formation of acid rain:

 When coal is burnt, sulphide minerals are oxidised during combustion, releasing SO2

 4FeS2 (s) + 11O2 (g)  2Fe2O3 (s) + 8SO2 (g)

 When metal smelters convert metal sulfides into metals yield vast amounts of SO2, e.g.

smelting chalcopyrite (CuFeS2) to obtain Cu results in the release of SO2

 2CuFeS2 (s) + 5O2 (g) + 2SiO2 (s)  2Cu (l) + 4SO2 (g) + 2FeSiO3 (l)  Nitric oxide is formed when nitrogen and oxygen react at high temperatures, for example, during lightning strikes or high temperature combustion reactions in furnaces and internal combustion engines

 N2 (g) + O2 (g)  2NO (g)  Nitric oxide, a colourless gas and neutral oxide, reacts with oxygen to form nitrogen dioxide, a brown gas and an acidic oxide

 2NO (g) + O2 (g)  2NO2 (g)

 CO2 is the product of complete combustion of fossil fuels such as petrol, kerosene and diesel oil

 2C8H18 (l) + 25O2 (g)  16CO2 (g) + 18H2O (g)

11) Explain the formation of acid rain and its effects The term ‘acid rain’ is used to describe rainwater that has a higher hydrogen ion concentration than normal, that is, higher than about 10‐5 mol/L (< pH 5). When the atmosphere is polluted with acidic oxides such as sulfur dioxide and nitrogen dioxide, rainwater can become quite acidic due to the high solubility of these gases in water which dissolve to produce solutions of various acids:

 Sulfur dioxide forms weak sulphurous acid

 SO2(g) + H2O(l) H2SO3(aq)  Sulfur trioxide produces strong sulphuric acid

 SO3(g) + H2O(l) H2SO4(aq)  Sulphurous acid can be catalytically oxidised to produce sulphuric acid

 2H2SO3(aq) + O2(g)  2H2SO4(aq)  Nitrogen dioxide produces weak nitrous acid and strong

 NO2(g) + H2O(l)  HNO2(aq) + HNO3(aq)  In the presence of water and oxygen, nitrous acid is catalytically oxidised to nitric acid

 2HNO2 (aq) + O2  2HNO3(aq)  Although pure water has a pH of 7.0 (neutral), natural water contains dissolved gases including carbon dioxide, which makes the water weakly acidic due to the presence of carbonic acid.

 CO2(g) + H2O(l) H2CO3(aq)

Effects of acid rain include:  Corrosion of marble statues and building facades. The calcium carbonate of the marble is attacked by the sulphuric acid in acid rain and the surface of the marble in converted into insoluble calcium sulphate

 CaCO3(s) + H2SO4(aq)  CaSO4(s) + H2O(l) + CO2(g)  Corrosion of metallic structures composed of iron and steel. The iron is oxidised by the hydrogen ions in the acid and becomes chemically weathered

 Fe(s) + H2SO4(aq)  FeSO4(aq) + H2(g)  Acidification of soils which inhibits the growth of plant seedlings. Basic minerals in the soil (such as dolomite and limestone) are attacked and dissolved by acidic water. Many types of sandstone have grains that are cemented together with calcite (calcium carbonate). The acid rain dissolves this cement and thus causes significant chemical weathering and erosion. Increasing acidification also affects nitrogen fixing bacteria that are vital to soil health  Acidification of lakes which places stress on populations of aquatic organisms. The presence of hydrogen ions interferes with the carbon dioxide/carbonate equilibrium in the water. The amount of dissolved carbon dioxide in the water drops as carbonate ions are removed. This has adverse effects on photosynthetic organisms 2‐ +  CO3 (aq) + 2H (aq) H2CO3(aq)  H2O(l) + CO2(g)

CO2(aq) CO2(g) Many aquatic invertebrates cannot reproduce in an acidic environment. Most fish eggs will not survive in water if the pH drops below 5.5. Below a pH of 5, adult fish will die as they cannot extract sufficient calcium from the water to maintain their skeletons  Chemical leaching of toxic heavy metals from bedrock and soil into lakes, which contributes to death of marine life due to heavy metal poisoning. Reductions in the numbers of any organism affects the entire food chain

 Devastating effects on plant life. For example, in the Black Forest in Germany, pine needles lose their waxy coating and turn brown. The trees become denuded of foliage, growth is slowed or inhibited, and in extreme cases, death of the forest occurs. Also, mineral nutrients such as potassium, calcium and magnesium that are required for plant growth can be removed when acid rain soaks into the ground. Some insoluble minerals can also be dissolved by the acidified water and cause a release of toxic levels of metal ions. High levels of aluminium ions (formed when clay minerals are attacked by acid) in the soil interfere with normal mineral uptake by plant roots

12) Evaluate (make a judgement based on criteria; determine the value of) reasons for concern about the release of oxides of sulfur and nitrogen in the environment. Concern for environment; evidence for increase of oxide levels and resulting environmental problems; include equations; what is being done? Judgement 13) Normal rainwater has a pH of 5.6. (a) How would you use any of the following indicators to determine the smallest pH range of normal rainwater? (b) What would the pH of acid rain be?

INDICATOR COLOUR IN LOW pH pH CHANGE COLOUR IN HIGH pH Phenolphthalein Colourless 8.3 – 10.0 Magenta (pink) Methyl orange Red 3.1 – 4.4 Yellow Bromothymol blue Yellow 6.0 – 7.6 Blue (a) Rinse two small beakers with distilled water. Fill each beaker with 25mL of normal rainwater. Place 3 drops of the bromothymol blue indicator into one of the beakers. Agitate the solution by swirling the contents of the beaker until a uniform colour is observed. This colour should be yellow, indicating a pH of 6.0 or lower, which corresponds with the given pH of normal rainwater. To obtain the smallest pH range of the sample requires the use of another indicator from the table, methyl orange. Hence, place 3 drops of methyl orange into the remaining beaker. Swirl the beaker to disperse the indicator throughout the sample and observe the colour of the solution. According to the table, methyl orange is yellow for a pH of 4.4 or higher, so the colour of the rainwater should be yellow. Thus, the smallest pH range of normal rainwater that can be obtained by using any of the three indicators in the table is pH 4.4‐6.0. (b) The pH of acid rain would be less than pH 5.0. Many readings show the pH of acid rain to be in the range pH 4.0 – 5.0, though industrial regions areas which are severely polluted have shown readings as low as pH 3.6. Such readings have been recorded in the United States and Europe. (Thickett, 2006).

14) The uptake of elements by plants from the soil depends on the pH. Study the chart below. The wider the band, the greater the uptake of the element. (a) Identify the elements that are taken up at pH 8.0, (b) Azaleas grow best in a pH range of 5.0 – 6.0. Predict which elements azaleas are likely to take up, and (c) Describe a method to test the pH of soil (a) At pH 8.0, the greatest uptake is of calcium, magnesium, nitrogen, phosphorus and potassium. There is very minimal uptake of boron, iron and manganese.

(b) The elements that azaleas are most likely to take up include boron, iron, manganese, nitrogen, and potassium. Azaleas also take up calcium, magnesium and phosphorous, but in lesser amounts than the aforementioned elements. (c) A test tube is one‐third filled with soil. Distilled water is added to within 2 cm of the top. The tube is stoppered and shaken, and the soil is allowed to settle for 3 minutes. A small amount of white barium sulphate suspension can be added to aid the settling process. A Pasteur pipette is then used to withdraw a sample of the supernatant water into two clean test tubes. Universal indicator is added to the first tube and the colour compared with a pH chart. A narrow range indicator (e.g. methyl orange, bromocresol green, chlorophenol red, Bromothymol blue, or phenol red – these are common indicators used by horticulturalists) is then selected to test the other water sample to determine a more accurate soil pH. The experiment is repeated 2 more times.

References

 http://www.epa.sa.gov.au/pdfs/info_acidrain.pdf  http://www.parliament.nsw.gov.au/Prod/Parlment/HansArt.nsf/V3Key/LC200603280 40  http://connector.clearblueday.net/files/RTA_Tunnel_Safety_Brochure.pdf  http://books.google.com/books?id=qciCdSFpFPkC&pg=PA93&lpg=PA93&dq=interpene trating+face+centred+cubic&source=web&ots=vK8YqrGXkR&sig=41SItttLYMBhzAaLpL dpY4oGWSU#PPA94,M1  http://en.wikipedia.org/wiki/Sulfur_dioxide  http://www.chemguide.co.uk/atoms/structures/giantcov.html  http://www.rta.nsw.gov.au/constructionmaintenance/completedprojects/m5east/m5e astcurrentairqualitydata/currentdata.html  www.deh.gov.au/soe/2001/atmosphere/atmosphere02‐16.html  http://www.worldofenergy.com.au/factsheet_coal/07_fact_coal_adv_disadv.html  http://www.bom.gov.au/inside/eiab/reports/caa03/chapter4/emissions_reduction_str ategies.shtml  Stewart, C et al. (ed) (2007) ecoscience: The 34th Harry Messel International Science School 2007, The University of Sydney: Sydney, p 40  Thickett, G (2006) Chemistry 2: HSC Course, Jacaranda: Milton, Qld