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HSC CHEMISTRY CORE TOPIC 2

HSC CHEMISTRY CORE TOPIC 2. THE ACIDIC ENVIRONMENT. INDICATORS. LE CHATELIER’S PRINCIPLE. TITRATION. pH. VOLUMETRIC ANALYSIS. NEUTRALISATION. CHEMICAL EQUILIBRIUM. THE ACIDIC ENVIRONMENT. HISTORY Lavoisier Davy Arrhenius. CARBOXYLIC ACIDS. ACIDIC OXIDES. BRONSTED

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HSC CHEMISTRY CORE TOPIC 2

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  1. HSC CHEMISTRY CORE TOPIC 2 THE ACIDIC ENVIRONMENT

  2. INDICATORS LE CHATELIER’S PRINCIPLE TITRATION pH VOLUMETRIC ANALYSIS NEUTRALISATION CHEMICAL EQUILIBRIUM THE ACIDIC ENVIRONMENT HISTORY Lavoisier Davy Arrhenius CARBOXYLIC ACIDS ACIDIC OXIDES BRONSTED LOWRY THEORY ACID RAIN ESTERIFICATION

  3. Subsection 1 Indicators were identified with the observation that the colour of some flowers depends on soil composition

  4. Indicators • substances that have distinctive colours in different types of chemical environments • natural acid-base indicators are vegetable dyes that provide the colour of flowers and vegetables • LITMUS is a pink mixture of compounds extracted from lichens grown mainly in the Netherlands • today many indicators used are manufactured dyes

  5. INDICATORS

  6. INDICATORS

  7. INDICATORS Universal Indicator is a mixture of • thymol blue • methyl red • bromothymol blue • phenolphthalein • dissolved in methanol, propan-1-ol and water

  8. INDICATORS • A solution gives the following colours in each indicator. Deduce the approximate pH Indicator Colour methyl orange yellow bromothymol blue yellow phenolphthalein colourless >4.4 <6.0 <8.0 pH of solution is 4.4 – 6.0

  9. INDICATORS • A solution gives the following colours in each indicator. Deduce the approximate pH Indicator Colour bromothymol blue blue thymolphthalein colourless phenol red red > 7.6 < 9.5 > 8.4 pH of solution is 8.4 – 9.5

  10. INDICATORS USES • Universal indicator used to test soil acidity/alkalinity (pH) • plants have preference for alkaline/neutral/acid soils – choice of crop • diseases that affect plants thrive in soils with a particular pH range • pH affects availability of nutrients • Phenol red used to test acidity/alkalinity of a swimming pool – level of disinfection

  11. Clay Minerals in Soil Metal ions bind to negative surface H+ are able to displace the surface cations from the clay. If aluminium ions are displaced by H+ the soil becomes toxic to crops growing in the soil.

  12. Subsection 2 While we usually think of the air around us as neutral, the atmosphere naturally contains acidic oxides of carbon, nitrogen and sulfur. The concentrations of these acidic oxides have been increasing since the Industrial Revolution

  13. Oxides of Period 3 NaOH(s)  MgO(s) + H2O(aq)g S(s) + O2(g)g SO2(g) + H2O(aq)g P4(s) + O2(g)gP2O5(s) P2O5(s) + H2O(aq)g

  14. Oxides of Period 3 NaOH(s)g Na+(aq) + OH-(aq) MgO + H2O g Mg(OH)2 S + O2g SO2 SO2 + H2O g H2SO3 (sulfurous acid) 2H2SO3 + O2g 2H2SO4 (sulfuric acid) P4 + 5O2g2P2O5 P2O5 + 3H2O g 2H3PO4 (phosphoric acid) Cl2O7 + H2O g 2HClO4 (perchloric acid)

  15. Oxides of Period 3

  16. Acid/Base Properties of Oxides BASIC OXIDES • metal oxides and hydroxides (ionic compounds) • soluble oxides react with water to form alkaline/basic solutions Na2O(s) + H2O(l) 2NaOH(aq) • react with acids/acidic oxides to form salts CuO(s) + H2SO4(aq) CuSO4(aq) + H2O(l) CaO(s) + SO2(g) CaSO3(s)

  17. Acid/Base Properties of Oxides ACIDIC OXIDES • generally oxides of non-metals (covalent compounds • called acid anhydrides • react with water to produce an acidic solution CO2(g) + H2O(l) H2CO3(aq) • react with bases to form salts CO2(g) + CaO(s) CaCO3(s) 2NaOH(l) + SiO2(s)  Na2SiO3(s) + H2O(l)

  18. Acid/Base Properties of Oxides AMPHOTERIC OXIDES • react with acids and bases • Al2O3, BeO, ZnO, PbO, SnO Al2O3(s) + 6HCl(aq)g 2AlCl3(aq) + 3H2O(l) Al2O3(s) + 2NaOH(aq) g2NaAlO2(aq) + H2O(l) NEUTRAL OXIDES • do not react with acids or bases • CO, N2O, NO

  19. 105 Db 107 Bh Group A Oxides basic amphoteric acidic 8A 1A 3A 4A 5A 6A 7A Group B

  20. Salts • in general acids and bases form salts • the type of salt is determined by the acid • HCl  chloride salts • HNO3  nitrate salts • H2SO4  sulfate salts • H3PO4  phosphate salts

  21. Solution with Na+ and Cl- ions Na+ Cl- Solid NaCl A Saturated Solution of NaCl Dissolution NaCl(s)g Na+(aq) + Cl-(aq) NaCl System at 25oC

  22. Solution with Na+ and Cl- ions Na+ Cl- Radioactive 24NaCl Solid NaCl A Saturated Solution of NaCl Radioactive 24NaCl is introduced into the beaker

  23. Solution with Na+ and Cl- ions Na+ 24Na+ Cl- Cl- Radioactive 24NaCl Solid NaCl 24NaCl 24NaCl A Saturated Solution of NaCl CHEMICAL EQUILIBRIUM Dissolution NaCl(s)g Na+(aq) + Cl_(aq) Recrystallisation NaCl(s)f Na+(aq) + Cl_(aq)

  24. Chemical Equilibrium

  25. CHEMICAL EQUILIBRIUM • only occurs in a closed system • no interchange of matter between system and surroundings • occurs in physical and chemical systems • temperature of the system remains constant • the rate of the forward reaction equals the rate of the reverse reaction • refers to reversible reactionswhere the forward reaction occurs simultaneouslywith reverse reaction

  26. CHEMICAL EQUILIBRIUM

  27. Le CHATELIER’S PRINCIPLE • French chemist Henri Le Chatelier (1850-1936) studied changes in systems that were in a state of equilibrium • If a stress is applied to a system in a state of chemical equilibrium, the system changes to relieve the stress • may be changes in • concentration • volume and pressure • temperature

  28. CONCENTRATION • increasing the concentration of a reactant or product will cause the system to favour the direction which will decrease the concentration of that substance • decreasingthe concentration of a reactant or product means the rate of the reaction using up that substance will decreasein rate • the rate of the other reaction to produce that substance will now have the faster rate

  29. CONCENTRATION • equilibria with solids and pure liquids e.g. C(s) + H2O(g)n CO(g) + H2(g)

  30. PRESSURE • changes in pressure have little effect on solids and liquids as they are only very slightly compressible • changes in pressure have significant effects on the concentration of GASES • gas pressure is proportional to the number of molecules • changing the partial pressureof a gas changes its concentration

  31. Partial Pressure of a Gas

  32. PRESSURE Adding an inert gas • does not change the partial pressures of any of the other gases • no effect on equilibrium

  33. VOLUME CHANGES • reducing the volume of the gas by half doublesthe pressure of the gas

  34. VOLUME CHANGES • decreasing volume - reaction rate increases in the direction that produces the smaller number of molecules in the reaction above there will be a shift to the right as the forward reaction increasesin rate to minimise the pressure changes • e.g. 2SO2(g) + O2(g)n 2SO3(g) • producing less molecules reduces the pressure • increasing the volume decreases the pressure • causes a shift to the left in the above reaction as this produces the larger number of molecules

  35. CHEMICAL EQUILIBRIUM 2NO2 n N2O4 • analyse the changes made to the equilibrium system shown in the diagram at the left

  36. TEMPERATURE • increasing the temperature increases the reaction rate • this is due to an increase in the fraction of collisions in which the total kinetic energy of reacting particles is at least equal to the activation energy • increasing the temperature favours the endothermic reaction • decreasingthe temperature favours the exothermic reaction

  37. TEMPERATURE

  38. Mass-Gas Volume • Calculate the mass of 18.25L of ammonia gas at 25.0oC and 100.0kPa. • At 100.0 kPa and 25.0oC, how many litres of carbon dioxide gas will be produced when 75.0g of calcium carbonate is decomposed into calcium oxide?

  39. Mass-Gas Volume • Solid lithium hydroxide has been used in space craft to remove carbon dioxide from air. Lithium carbonate and water are formed. What mass of lithium hydroxide would be needed to remove 250.0 L of carbon dioxide at 100.0 kPa and 25.0oC?

  40. Acid Rain Sulfur dioxide • natural and man-made sources • reactions to produce it • sulfur compounds in coal S(s) + O2(g) SO2(g) • smelting of sulfide ores ZnS(s) + O2(g) Zn(s) + SO2(g) • oxidation of H2S – decay and industrial 2H2S + 3O2(g)  2H2O(l) + 2SO2(g) • reactions to produce acidic solutions • effects • living things and environment • corrosion metals, limestone buildings (CaCO3)

  41. Acid Rain Nitrogen oxides NOx (NO, NO2) • natural and man-made sources • reactions to produce it • high temperature engines N2 + O2g2NO (neutral oxide) 2NO + O2g 2NO2 (acidic oxide) • reactions to produce acidic solutions • 2NO2 + H2O g HNO2 + HNO3 • effects • living things and environment

  42. Production of Ozone photodecomposition NO2g NO + O ozone formation O + O2g O3 regeneration of nitrogen dioxide O3 + NO g O2 + NO2 Ozone is a secondary pollutant in the troposphere

  43. Photochemical Smog

  44. What is acid rain? More appropriate term is “acidic deposition” -snow, fog, sleet, haze, dry deposition What is Acid Rain? Pure water: pH 7 Natural rain: pH 5-6 Acid rain pH < 5

  45. Acid Rain 1730s – originated at height of Industrial Revolution 1872 – Robert Smith, an English chemist, coined the phrase “acid rain” 1950s – lake acidification first described 1960s – became more noticeable and subsequently became worse in rural areas  tall chimneys on factories allow wind to transport pollutants far away from sources of production

  46. Acid Rain Out west, in the Rocky Mountains scientists are finding that power plant emissions are saturating high-elevation watersheds in Colorado with acid-causing nitrogen. Evergreen forests are losing their needles and tree health is declining throughout the forest range

  47. Acid Rain Acid rain damage Blue Ridge Mountains North Carolina

  48. Acid Rain Acid rain damage on monument CaCO3 (s) + H2SO4 (aq)  CaSO4 (aq) + CO2 (g) + H2O (l)

  49. Acid Rain Tasmania - Queenstown emerged as a boomtown of the 1890s when gold and minerals were discovered at Mount Lyell. The strange but arresting 'moonscape' that surrounds the town was caused by acid-rain during the mining era.

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