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The Acidic Environment

The Acidic Environment. YEAR 12 CHEMISTRY. ACIDS AND BASES. Acids are compounds that: have a low pH (below 7) taste sour turn blue litmus paper red react with bases to neutralize them and produce salts release H 2 gas in reactions with active metals aqueous solutions conduct electricity

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The Acidic Environment

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  1. The Acidic Environment YEAR 12 CHEMISTRY

  2. ACIDS AND BASES Acids are compounds that: • have a low pH (below 7) • taste sour • turn blue litmus paper red • react with bases to neutralize them and produce salts • release H2 gas in reactions with active metals • aqueous solutions conduct electricity • furnish H+ • burn the skin if strong • dissolve carbonates These are some common acids. What others can you think of? Make a list in your notes.

  3. ACIDS AND BASES Bases are compounds that: • have a high pH (above 7) • taste bitter • turn red litmus paper blue • react with acids to neutralize them and produce salts • are slippery feeling • aqueous solutions conduct electricity • burn skin if strong • react with fats to form soap • furnish OH- What other bases/alkalis do you know?

  4. ACIDS AND BASES Some common laboratory acids HCl hydrochloric acid HClO4perchloric acid HBrhydrobromic acid HNO3 nitric acid HI hydroiodic acid H2SO4 sulfuric acid Five of these acids are classified as monoprotic acids. They only have one hydrogen which they are able to donate. Sulfuric acid is classified as a diprotic acid because it has two acidic hydrogens that it can donate. Similarly, an acid which has three donatable hydrogens would be classified as triprotic.

  5. ACIDS AND BASES Some common laboratory bases NaOH sodium hydroxide Ca(OH)2 calcium hydroxide KOH potassium hydroxide Mg(OH)2 magnesium hydroxide NH3 ammonia Na2CO3 sodium carbonate Bases are often found in everyday products such as many cleaning products (sodium hydroxide), antacid products (magnesium hydroxide )and fertilisers (ammonia). It is a common misconception that bases are not as dangerous as acids. In fact, many bases can be as much or more corrosive than many acids.

  6. pH Indicators • A simple explanation of pH is that it is a measure of acidity/basicity (more in-depth explanation to follow) • Many substances change colour as they are exposed to different pH levels. These can be used to “indicate” the pH of substances when the colour ranges are known • Some natural products such as litmus, cabbage, grapes and tea are natural indicators while others such as phenolphthalein and methyl orange are synthetic • pH indicators are themselves acids or bases as they donate or accept protons (more on this later) Notice the variety of ranges where different indicators change colours. Some have more than one change.

  7. pH of some common substances

  8. pH Indicators Some specific examples: • Litmus red5 8blue • Phenolphthalein colourless8.3 10red • Methyl orange red3.1 4.4yellow • Bromothymol blue yellow6 7.6blue Note: colour changes within these ranges are gradual

  9. A pH problem An unknown solution produces the colours above. What is the pH range of this solution? 8-8.3

  10. Uses of Indicators Soil testing • Most plants cannot survive outside pH 5.5-7.5 • Different plants require different pH level • 2 methods for testing soil pH • Mix small sample with universal indicator and sprinkle BaSO4 powder on top – read colour • Mix soil with water in a test tube and add indicator – read colour • Changing soil pH • Too acidic – add NH3, CaO (lime), CaCO3 • Too basic – add manure, pine bark, peat

  11. Uses of Indicators Pool acidity • Pool should be pH of 7.4 • Tested with a meter or an indicator such as phenol red • Changing the pH • Too acidic – add CaOCl2 (pool chlorine), Na2CO3 • Too basic – add HClsoln, NaHSO4

  12. Non-metal Oxides Sources • Atmospheric O2 is very reactive and reacts with many substances to form oxides • Natural formation • CO2 – from respiration (“burning” sugars for energy) • NO2 – from lightning strikes (N2 + 2O2 in the air  2NO2) • SO2 – released from volcanoes or H2S + O2 SO2 + H2O (H2S produced by bacterial decomposition of organic matter) • Human causes (bushfires and burning fossil fuels) • CO2 – fossil fuel combustion product • NO – high temperature combustion product • NO2 – NO is easily oxidised in the air (NO + O2  NO2) • SO2 – burning coal that contains S as an impurity • SO3 – SO2 is easily oxidised in the air

  13. Acidic Non-metal Oxides Many non-metal oxides react with water in the atmosphere to produce acids; CO2 + H2O  H2CO3 (carbonic acid) SO2 + H2O  H2SO3 (sulfurous acid) SO3 + H2O  H2SO4 (sulfuric acid) 2NO2 + H2O  HNO3 + HNO2 (nitric and nitrous acid) These non-metal oxides are all gases Their acidic products all contribute to the acidity of rain

  14. Oxide Trends in the Periodic Table Oxides tend to increase in acidity from left to right In general: • Metal oxides are basic (left side) • Non-metal oxides are acidic (right side) Exceptions: • Amphoteric oxides (i.e. Al, Be, Ga, Sn, Pb) Why this trend? This is due to electronegativity increasing from left to right (see following slides for more details)

  15. Oxides on the left side of the PT • Electrons are transferred to the O2- • This is due to the ionic nature of these bonds because of a large difference in electronegativities, therefore ions are formed in solution For example: elementelectronegativity Na 0.93 O 3.44 This means: Na2O(aq) Na+ + O2- and O2- + H+ OH-(readily) This overall consumption of H ions leads to an increase of pH (i.e. Basic)

  16. Oxides on the right side of the PT • Electrons are shared with the O2- • This is due to the covalent nature of these bonds because of a small difference in electronegativities, thus no ions are formed For example: elementelectronegativity S 2.58 O 3.44 This means: Due to a partially positive S central atom, SO3(aq) + H2O  H+ + HSO4- This overall production of H ions leads to an decrease of pH (i.e. acidic). (See following slide for details of this reaction)

  17. Sulfur trioxide forms an acid - O O OH- S S H+ H+ O O O OH O Sulfur trioxide water Sulfuric acid solution

  18. Industrial Pollution Research The Industrial Revolution brought about many changes in the 19th Century. One of these changes involves an increase in the release of sulfur dioxide SO2 and oxides of nitrogen such as NO2. • Summarise the industrial origins of SO2 and NO2 and evaluate reasons for concern about their release into the environment • Consider health issues, visibility issues, effects on the environment (rivers, soil, plants), and effects on buildings

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