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Chemistry in Action

Chemistry in Action. W Richards Worthing High School. Balancing equations. Sodium + water sodium hydroxide + hydrogen Na + H 2 O NaOH + H 2. O. Na. Na. H. H. H. H. H. O. Consider the following reaction:. +. +.

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Chemistry in Action

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  1. Chemistry in Action W Richards Worthing High School

  2. Balancing equations Sodium + water sodium hydroxide + hydrogen Na + H2O NaOH + H2 O Na Na H H H H H O Consider the following reaction: + + This equation doesn’t balance – there are 2 hydrogen atoms on the left hand side (the “reactants” and 3 on the right hand side (the “products”)

  3. Balancing equations Sodium + water sodium hydroxide + hydrogen O O Na Na Na Na H H H H H H H H O O 2Na(s) + 2H2O(l) 2NaOH(aq) + H2(g) We need to balance the equation: + + Now the equation is balanced, and we can write it as:

  4. Some examples Mg + O2 Zn + HCl Fe + Cl2 NaOH + HCl CH4 + O2 Ca + H2O NaOH + H2SO4 CH3OH + O2 MgO ZnCl2 + H2 FeCl3 NaCl + H2O CO2 + H2O Ca(OH)2 + H2 Na2SO4 + H2O CO2 + H2O 2 2 2 3 2 2 2 2 3 2 2 2 2 2 4

  5. Testing for carbon dioxide Gas Limewater Limewater turns milky/cloudy

  6. Adding acid to carbonates Calcium carbonate + hydrochloric acid calcium chloride + carbon dioxide + water CaCO3(s) + HCl(aq) CaCl2(aq) + CO2(g) + H2O(l) Carbonates are compounds containing carbon and oxygen. When an acid is added to a carbonate the carbonate starts to fizz. A gas called _________ _______ is produced. 2

  7. Flame tests

  8. Flame tests Compounds containing lithium, sodium, potassium, calcium and barium can be recognised by burning the compound and observing the colours produced: Lithium Red Sodium Yellow Potassium Lilac Calcium Brick red Barium Green

  9. Metal ions Calcium is in group 2 and has two electrons in its outer shell, so it will form a Ca2+ ion. Chlorine is in group 7 so a chloride ion will be Cl- Metal compounds in a solution contain metal ions. For example, consider calcium chloride: Calcium chloride has the formula CaCl2

  10. Metal ions and precipitates Ca2+(aq) + OH- Ca(OH)2 (s) Some metal ions form precipitates, i.e. an insoluble solid that is formed when sodium hydroxide is added to them. Consider calcium chloride: 2

  11. Metal ions and precipitates Ca2+(aq) + OH- Ca(OH)2 (s) Some metal ions form precipitates, i.e. an insoluble solid that is formed when sodium hydroxide is added to them. Consider calcium chloride: 2

  12. Testing for chloride and sulphate ions For each test state: 1) The colour of the precipitate 2) What compound it is Test 1: Chloride ions Add a few drops of dilute nitric acid to the chloride ion solution followed by a few drops of silver nitrate. Precipitate formed = silver chloride (white) Test 2: Sulphate ions Add a few drops of dilute hydrochloric acid to the sulphate ion solution followed by a few drops of barium chloride. Precipitate formed = barium sulphate (white again)

  13. Ammonium, nitrate, bromide and iodide ions Ammonium ions: Add sodium hydroxide and test the gas using damp litmus paper – ammonia gas turns damp litmus paper blue. Nitrate ions: Add sodium hydroxide followed by aluminium powder and test using damp litmus paper. Bromide and iodide ions: Add a few drops of dilute nitric acid followed by a few drops of silver nitrate solution. A pale yellow precipitate should be formed for bromide ions and a darker yellow precipitate for iodide ions.

  14. Thermal decomposition Copper carbonate: CuCO3 (s) CuO(s) + CO2 (g) (Green – Black) Zinc carbonate: ZnCO3 (s) ZnO(s) + CO2 (g) (White – Yellow) A “thermal decomposition” reaction occurs when a compound breaks down (“decomposition”) through the action of heat. • Practical work: • Perform a thermal decomposition reaction on each of these compounds and state: • The colour changes you observed • The reaction that happened

  15. Sulphuric acid Sulphur + oxygen sulphur dioxide Sulphur dioxide + oxygen sulphur trioxide Sulphur trioxide + conc. sulphuric acid oleum Oleum + water sulphuric acid Sulphuric acid has many important uses – car batteries, detergents, fertilisers etc. How sulphuric acid is made: Step 1: Burn sulphur in air: Step 2: Pass the sulphur dioxide over a vanadium oxide catalyst at 450OC: Step 3: Dissolve the sulphur trioxide in sulphuric acid: Step 4: Add water to the oleum:

  16. Reversible Reactions A + B C + D Endothermic reactions Exothermic reactions Increased temperature: Increased temperature: A + B C + D A + B C + D Decreased temperature: Decreased temperature: A + B C + D A + B C + D When a reversible reaction occurs in a CLOSED SYSTEM (i.e. no reactants are added or taken away) an EQUILIBRIUM is achieved – in other words, the reaction goes at the same rate in both directions: More products Less products Less products More products

  17. Sulphuric acid Endothermic 2SO2 + O2 2SO3 Exothermic Step 2 in the manufacture of sulphuric acid is an example of a reversible reaction: What would happen if the temperature was decreased? The reaction would favour the production of sulphur trioxide BUT the reaction would happen at a slower rate. Solution – use 450OC as a compromise

  18. Iron and Steel The iron contains roughly 5% carbon and different metals and is very ________. In order to reduce these impurities and convert the iron into _________ the molten iron is transferred into another furnace where it is mixed with recycled scrap iron and pure ___________. The oxygen reacts with the metal impurities to form ________ oxides. Calcium carbonate is also added to remove some of the acidic oxides as _______ when the furnace is tilted. Words – slag, brittle, steel, oxygen, acidic In previous work we considered the role of the blast furnace in extracting iron from its ore.

  19. Making steel – the reactions Silicon + oxygen Silicon oxide Calcium carbonate calcium oxide + carbon dioxide Silicon oxide + calcium oxide calcium silicate Steel with chromium and nickel is called stainless steel Steel with a high carbon content is strong but brittle Steel with a low carbon content is easily shaped 1) Mixing oxygen with silicon impurities: 2) Decomposition of limestone: 3) Adding these products together:

  20. Titanium Titanium chloride + sodium titanium + sodium chloride In this reaction the titanium is displaced my a more reactive metal. This reaction is done in an argon atmosphere to avoid any further reactions. Titanium ions have a charge of 4+ and gain four electrons to become titanium atoms. This is a reduction reaction. Titanium is a strong metal used in planes, replacement hip joints, bikes etc. Two steps are used in its manufacture: Step 1: Convert titanium dioxide (ore) to titanium chloride Step 2: Displace the titanium using sodium or magnesium:

  21. Aluminium ++++ ---- Al H2SO4 Aluminium is a fairly reactive metal that doesn’t corrode due to forming a thin layer of aluminium oxide. This explains why greenhouses don’t rust and don’t need to be painted. A thicker layer of aluminium oxide can be made artificially. There are two stages: 1) Remove the natural layer by placing the aluminium in sodium hydroxide. 2) Use electrolysis on sulphuric acid with the aluminium as the positive electrode. This is called anodising.

  22. Electroplating ++++ ---- Silver electrode Object to be plated Solution containing silver ions

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