Reaction types


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Reaction types

  1. 1. Types of Reactions Yr 10 Chemistry
  2. 2. <ul><li>There are many different types of reaction in chemistry </li></ul><ul><li>We are going to look at a selection </li></ul>Can you name any?
  3. 3. Reactions <ul><li>Dissociation </li></ul><ul><li>Ionisation </li></ul><ul><li>Combustion </li></ul><ul><li>Precipitation </li></ul><ul><li>Decomposition </li></ul><ul><li>Displacement </li></ul><ul><li>Neutralisation </li></ul><ul><li>Redox </li></ul>
  4. 4. Reaction Rates <ul><li>Every different chemical reaction occurs at a different speed, or a different rate. </li></ul><ul><li>The study of the rates of reactions is often referred to as reaction kinetics. </li></ul>
  5. 5. Rates of Reaction Chemical reactions occur when different atoms or molecules collide: <ul><li>For the reaction to happen the particles must have a certain amount of energy – this is called the ACTIVATION ENERGY. </li></ul><ul><li>The rate at which the reaction happens depends on many things: </li></ul><ul><li>The temperature of the reactants, </li></ul><ul><li>Their concentration </li></ul><ul><li>Their surface area </li></ul><ul><li>The reactivity of the reactants </li></ul><ul><li>Whether or not a catalyst is used </li></ul>
  6. 6. Rate of reaction graph Amount of product formed Time Slower reaction Fast rate of reaction here Slower rate of reaction here due to reactants being used up
  7. 7. Reaction Rates <ul><li>Reactions can occur </li></ul><ul><li>VERY fast eg explosions </li></ul><ul><li>Very slow eg. Coal converting to diamonds. </li></ul><ul><li>Or anywhere in between. </li></ul><ul><li>The rate is the speed of conversion from the reactants (the things that are combining) to the products (the things produced). </li></ul><ul><li>We will look at the factors influencing reaction rates in isolation, not as a entire system. </li></ul>
  8. 8. Reaction Rates Collision Theory <ul><li>Reactions occur when 2 reactive molecules collide with each other. </li></ul><ul><li>The rate at which these collisions occur governs the reaction rate. </li></ul><ul><li>These collisions must occur with sufficient energy and at the appropriate orientation. </li></ul>
  9. 9. Reaction Rate Effects <ul><li>Temperature </li></ul><ul><ul><li>As the temperature increases, the rate of reaction increases. </li></ul></ul><ul><ul><li>This is because the particles are moving more rapidly when they are heated, therefore they are much more likely to collide and thus react with each other. </li></ul></ul><ul><ul><li>The collisions are much more likely to have sufficient energy to allow a reaction to occur. </li></ul></ul><ul><ul><li>As a rule of thumb a reaction rate increases about 10 fold for each 10 o C rise in temperature. </li></ul></ul><ul><ul><li>Many reactions give off heat (exothermic), therefore they will cause an acceleration of the reaction as they proceed. </li></ul></ul>
  10. 10. Reaction Rate Effects <ul><li>CONCENTRATION: </li></ul><ul><ul><li>An increase in concentration will bring about an increase in the reaction rate as if there are more particles then they are more likely to collide. </li></ul></ul>
  11. 11. Concentration <ul><li>If we make one reactant more concentrated (like making a drink of orange squash more concentrated ) </li></ul><ul><li>There are more particles in the same volume to react </li></ul><ul><li>So the reaction goes faster. </li></ul>There are less red particles in the same volume so there is less chance of a collision There are more red particles in the same volume so there is more chance of a collision so the reaction goes faster
  12. 12. Reaction Rate Effects <ul><li>SURFACE AREA: </li></ul><ul><ul><li>An increase in the surface area increases the likelihood of a collision as there is a greater area on which those collisions can occur. </li></ul></ul><ul><ul><li>A reduction in particle size will bring about an increase in surface area. Therefore there is a greater area on which the reaction can occur. </li></ul></ul><ul><ul><li>Particles that have been ground up into a powder are much more reactive than when in a solid form. This is why many dusts such as wheat dust or coal dust are highly explosive. </li></ul></ul>
  13. 13. Surface area <ul><li>If we make the pieces of the reactants smaller we increase the number of particles on the surface which can react. </li></ul><ul><li>This makes the reaction faster. </li></ul>The particles on the surface can react When cut into smaller pieces there are more particles on the outside that can react
  14. 14. Reaction Rate Effects <ul><li>REACTIVITY: </li></ul><ul><ul><li>In general terms, excluding the noble gases, the most reactive elements are those on the sides of the periodic table. (Gp I and Gp VII) </li></ul></ul>
  15. 15. Measuring rate of reaction Two common ways: 1) Measure how fast the products are formed 2) Measure how fast the reactants are used up
  16. 16. Activation Energy <ul><li>There is a minimum amount of energy required for reaction: the activation energy . </li></ul><ul><li>Just as a ball cannot get over a hill if it does not roll up the hill with enough energy, a reaction cannot occur unless the molecules possess sufficient energy to get over the activation energy barrier. </li></ul>
  17. 17. Exothermic v Endothermic <ul><li>Energy is put into the system to get it “over the activation energy hill” </li></ul><ul><li>If there is a lot of energy out, the reaction is exothermic and the system rolls down the energy hill to a lower point. </li></ul><ul><li>If the energy out is less than the activation energy, the reaction is endothermic </li></ul>
  18. 18. Activation Energy <ul><li>This can be easily understood when the ignition of flammable liquids is considered. </li></ul><ul><li>Compare how easy it is to light a series of flammable liquids such as methanol, ethanol, hexane and ether. </li></ul><ul><li>Also compare this to your knowledge of igniting fuels such as petrol and diesel. </li></ul>
  19. 19. Catalysts <ul><li>A catalyst is something that causes a process to proceed at a faster rate without being itself consumed by the process. </li></ul><ul><li>An example of this is the platinum inside a car’s exhaust system which allows the conversion of exhaust gases to a more environmentally friendly form. The platinum facilitates the reaction, but does not get used up by the process. </li></ul>
  20. 20. Catalytic Converter
  21. 21. Catalysts <ul><li>Catalysts can be adversely affected by some substances which can inhibit their performance. </li></ul><ul><li>Such materials are referred to as poisons. An example of this is the presence of lead in petrol. A contamination of lead in fuel poisons the platinum catalytic converter and permanently prevents it from operating. </li></ul>
  22. 22. Catalysts <ul><li>Many life processes are facilitated by natural catalysts called enzymes. Without these catalysts, the activation energy is too high to allow rapid reactions, and many bodily functions would cease. </li></ul><ul><li>Without enzymes in the stomach, the process of digestion would be extremely slow, and most of the nutritional value of the foods eaten would be excreted. </li></ul>
  23. 23. Dissociation Yr 10 Chemistry
  24. 24. Dissociation reactions <ul><li>Are reactions that involve setting free ions </li></ul><ul><li>They occur when an ionic solid dissolves in water. </li></ul><ul><li>The positive and negative ions in the lattice separate in solution and are now free to move around </li></ul>
  25. 25. Example <ul><li>Sodium chloride dissociates when added to water </li></ul><ul><li>NaCl (s) + aq </li></ul><ul><li>Na + (aq) + Cl - (aq) </li></ul>
  26. 26. Solubility <ul><li>In order for dissociation to happen the ionic compound must be SOLUBLE in water. </li></ul><ul><li>The table shows soluble anions, in addition all compounds of potassium, sodium and ammonium are soluble </li></ul>Anion Formula exceptions Acetate CH 3 COO - Nitrate NO 3 - Chloride Cl - Silver, mercury, lead Sulfate SO 4 2- Barium, strontium, lead
  27. 27. Questions <ul><li>For each compound record either a null result or write the dissociation equation </li></ul><ul><li>Sodium sulfate calcium hydroxide </li></ul><ul><li>Lithium nitrate aluminium chloride </li></ul><ul><li>Potassium hydroxide barium sulfate </li></ul><ul><li>Copper sulfide ammonium carbonate </li></ul>
  28. 28. Solutions <ul><li>Sodium sulfate Na + + SO 4 2- </li></ul><ul><li>calcium hydroxide null </li></ul><ul><li>Lithium nitrate Li + + NO 3 - </li></ul><ul><li>aluminium chloride Al 3+ + Cl - </li></ul><ul><li>Potassium hydroxide K + + OH - </li></ul><ul><li>barium sulfate null </li></ul><ul><li>Copper sulfide null </li></ul><ul><li>ammonium carbonate NH 4 + + CO 3 2- </li></ul>
  29. 29. Ionisation Yr 10 Chemistry
  30. 30. Ionisation reactions <ul><li>Are reactions where ions are formed </li></ul><ul><li>They occur when a molecule dissolves in water and a covalent bond is broken creating ions. </li></ul>
  31. 31. Example <ul><li>Hydrogen chloride is a covalent compound which separates into H + and Cl - dissolved in water </li></ul><ul><li>HCl (g) + H 2 O (l) H + (aq) + Cl - (aq) </li></ul><ul><li>Ammonia is another covalent compound that separates when dissolved in water </li></ul><ul><li>NH 3(g) + H 2 O (l) NH 4 + (aq) + OH - (aq) </li></ul>
  32. 32. Are these ionisation reactions? <ul><li>K 2 SO 4 K + + SO 4 2- </li></ul><ul><li>SO 2 + H 2 O H 2 SO 3 </li></ul><ul><li>PbCl 2 Pb 2+ + Cl - </li></ul><ul><li>HF+H 2 O H 3 O + + F - </li></ul><ul><li>NaOH Na + + OH - </li></ul>
  33. 33. Are these ionisation reactions? <ul><li>Nitrous oxide and water </li></ul><ul><li>Calcium carbonate and water </li></ul><ul><li>Lead oxide and water </li></ul><ul><li>Ammonium nitrate and water </li></ul><ul><li>Carbon dioxide and water </li></ul>
  34. 34. Combustion Yr 10 Chemistry
  35. 35. Combustion <ul><li>Substances can burn in oxygen to form an oxide </li></ul><ul><li>We will consider metals and hydrocarbons </li></ul>
  36. 36. Metal + Oxygen <ul><li>Many metals react with oxygen to form oxides of the metal. </li></ul><ul><li>The standard reaction is: </li></ul><ul><ul><li>metal + oxygen  metal oxide </li></ul></ul><ul><li>eg: </li></ul><ul><ul><li>Iron + oxygen  iron oxide </li></ul></ul><ul><ul><li>Magnesium + oxygen  magnesium oxide </li></ul></ul>
  37. 37. Metal + Oxygen Magnesium in Oxygen
  38. 38. Hydrocarbon + Oxygen <ul><li>Hydrocarbons are compounds made of only hydrogen and carbon </li></ul><ul><li>They burn in oxide to form the oxides of their components </li></ul><ul><ul><li>Carbon dioxide </li></ul></ul><ul><ul><li>Dihydrogen monoxide (water) </li></ul></ul>
  39. 39. Hydrocarbon + Oxygen <ul><li>Eg </li></ul><ul><li>Methane + oxygen carbon dioxide + water </li></ul><ul><li>Propane </li></ul><ul><li>Butane </li></ul><ul><li>Octane </li></ul><ul><li>Now can you write them in symbol form? </li></ul>Can you write equations for the combustion of
  40. 40. Hydrocarbon + Oxygen <ul><li>Propane + oxygen carbon dioxide + water </li></ul><ul><li>Butane + oxygen carbon dioxide + water </li></ul><ul><li>Octane + oxygen carbon dioxide + water </li></ul>
  41. 41. Precipitation Yr 10 Chemistry
  42. 42. Precipitation Reactions <ul><li>Are any reaction in aqueous form that results in an insoluble product </li></ul><ul><li>You need to remember solubilities to predict these reactions </li></ul><ul><li>Precipitation is the reverse of dissolution </li></ul>
  43. 43. Solubility <ul><li>Salts are substances composed of ions; many salts are soluble in water </li></ul><ul><li>Solubility is the amount of a substance needed to make a saturated solution at a specific temperature. </li></ul><ul><li>Solubility of solids and liquids varies widely, for example: </li></ul><ul><li>35.7g of NaCl dissolves in 100mL of H 2 O, but only 0.044g of PbI 2 dissolves in 100mL of H 2 O </li></ul>
  44. 44. Solubility anion formula Solubility Acetate C2H3O2 - All soluble Nitrate NO3 - All soluble Chloride Cl - Not Ag, Hg, Pb Sulfate SO4 2- Not Ba, Sr, Pb Carbonate CO3 2- Insoluble* Hydroxide OH - Insoluble* Sulfide S 2- Insoluble Not group I or II Na/K/NH4 salts All soluble
  45. 45. Solubility <ul><li>Solubility is temperature dependent. </li></ul><ul><li>For gases, an increase in temperature will decrease the amount of gas that can dissolve </li></ul><ul><li>For salts, the amount and rate of dissolving will increase as temperature increases : </li></ul>Temperature Amount & Rate of Dissolving of Salt
  46. 46. Precipitate Formation: Formation of Lead (II) Iodide When mixed, two clear solutions produce a dense yellow precipitate To observe the formation of a precipitate from ionic compounds in an aqueous state.
  47. 47. Demonstration <ul><li>The formation of PbI 2(s) from Pb(NO 3 ) 2 and 2KI solutions. </li></ul>Pb 2- NO 3 - K + I - + PbI 2(s) K + NO 3 -
  48. 48. Double Replacement Reactions <ul><li>Occurs when the elements in a solution of reacting compounds exchange places, or replace one another. </li></ul><ul><li>AB + CD  AD + CB </li></ul><ul><li>For this demonstration: </li></ul><ul><li>Pb(NO 3 ) 2(aq) + 2KI (aq)  PbI 2(s) + 2KNO 3(aq) </li></ul>
  49. 49. Precipitates <ul><li>A precipitate is an insoluble solid formed when two ionic solutions react. </li></ul><ul><li>Pb(NO 3 ) 2 and KI are ionic compounds that dissociate into ions as they dissolve in water, but when mixed together, undergo a double replacement reaction to form an insoluble yellow precipitate, PbI 2 . </li></ul>
  50. 50. Write equations for … <ul><li>Sodium chloride + silver nitrate </li></ul><ul><li>Lead sulfide + copper chloride </li></ul><ul><li>Barium chloride + potassium sulfate </li></ul><ul><li>Lead nitrate + sodium carbonate </li></ul><ul><li>Magnesium chloride + potassium carbonate </li></ul><ul><li>Barium nitrate + potassium sulfate </li></ul>
  51. 51. Practical <ul><li>Carry out precipitation practical </li></ul>
  52. 52. Decomposition Yr 10 Chemistry
  53. 53. Decomposition reactions <ul><li>Are those where a compound breaks down when heated </li></ul><ul><li>Calcium carbonate decomposes with heat into calcium oxide and carbon dioxide </li></ul><ul><li>CaCO 3 CaO + CO 2 </li></ul><ul><li>Remember carbon dioxide can be tested for using limewater </li></ul>
  54. 54. Practical <ul><li>You will decompose calcium carbonate (white solid) to form calcium oxide (white solid) and carbon dioxide (colourless gas) </li></ul><ul><li>How will you know the reaction has occurred? </li></ul><ul><li>What test could you do? </li></ul>
  55. 55. Which decomposes quickest? <ul><li>The following carbonate solutions are provided </li></ul><ul><li>Sodium, calcium, magnesium, copper, zinc </li></ul><ul><li>Your challenge is to design a practical that will allow you to find out which decomposes quickest on heating </li></ul>
  56. 56. Displacement Reactions Year 10 Chemistry
  57. 57. The Activity Series of Metals <ul><li>Remember that metals vary in their activity. </li></ul><ul><li>Chemists have made a list and placed the metals in order of their activity. </li></ul>
  58. 58. The Activity Series of Metals most reactive least reactive K Ca Na Mg Al Zn Fe Cu Ag Au
  59. 59. Displacement Reactions <ul><li>Because some metals are more active than others, they want to react and form compounds more than others. </li></ul><ul><li>If a more reactive metal is combined with a less reactive metal salt, the more reactive metal will displace the less reactive metal from its salt. </li></ul><ul><li>The less reactive metal will precipitate out of the solution. </li></ul>
  60. 60. Displacement Reactions <ul><li>For example consider what happens if Mg is mixed with a solution of CuSO 4 . </li></ul><ul><li>Mg + CuSO 4  </li></ul><ul><li>Magnesium is more active than copper. </li></ul><ul><li>The magnesium pushes the copper out of solution. </li></ul><ul><li>The magnesium reacts with the SO 4 ion forming magnesium sulfate. </li></ul><ul><li>Copper metal forms. </li></ul>Mg SO 4 + Cu (s)
  61. 61. Displacement Reactions <ul><li>Consider : </li></ul><ul><li>Fe + Al 2 (SO 4 ) 3  </li></ul><ul><li>In this case, according to the activity series, aluminium is the more active metal. </li></ul><ul><li>It has already reacted with the sulfate group. </li></ul><ul><li>The iron is not active enough to push the aluminium out of solution. </li></ul><ul><li>No visible reaction is observed. </li></ul>NVR
  62. 62. Is it displacement? <ul><li>If it is displacement then write out the equation </li></ul><ul><li>Na + CuS </li></ul><ul><li>Fe + Ag 2 O </li></ul><ul><li>CuSO 4 + Mg </li></ul><ul><li>MgO + K </li></ul><ul><li>Na 2 O + Al </li></ul><ul><li>Al 2 O 3 + Ca </li></ul>
  63. 63. Neutralisation – Reactions with Acids Year 10 Chemistry
  64. 64. Reactions of Acids <ul><li>There are some standard reactions of acids: </li></ul><ul><ul><li>acid + alkali </li></ul></ul><ul><ul><li>acid + metal </li></ul></ul><ul><ul><li>acid + metal oxide </li></ul></ul><ul><ul><li>acid + carbonate </li></ul></ul><ul><li>These reactions are called Neutralisation Reactions because they neutralise an acid. </li></ul>
  65. 65. PROPERTIES OF ACIDS & BASES <ul><li>Acids </li></ul><ul><ul><li>Produce hydrogen ions (H + ) in H 2 O </li></ul></ul><ul><ul><li>Taste sour </li></ul></ul><ul><ul><li>Turn blue litmus (lichen)  red </li></ul></ul><ul><ul><li>Act as electrolytes in solution </li></ul></ul><ul><ul><li>Neutralise solutions containing hydroxide ions (OH - ) </li></ul></ul><ul><ul><li>React with several metals releasing H 2(g)  corrosion </li></ul></ul><ul><ul><li>React with carbonates releasing CO 2(g) </li></ul></ul><ul><ul><li>Destroy body tissue </li></ul></ul>
  66. 66. <ul><li>Bases </li></ul><ul><ul><li>Produce or cause an increase in hydroxide ions (OH - ) in H 2 O </li></ul></ul><ul><ul><li>Taste bitter </li></ul></ul><ul><ul><li>Turn red litmus  blue </li></ul></ul><ul><ul><li>Act as electrolytes in solution </li></ul></ul><ul><ul><li>Neutralise solutions containing hydrogen ions (H + ) </li></ul></ul><ul><ul><li>Have a slippery, ‘soapy’ feel </li></ul></ul><ul><ul><li>Destroy body tissue/ dissolve fatty (lipid) material </li></ul></ul>
  67. 67. Measuring Acids and Bases <ul><li>Acids and bases strengths are measured using the pH scale. </li></ul><ul><li>(pH stands for pondus hydrogenii , or potential of hydrogen.) </li></ul><ul><li>Acids have a pH between 0 and 7 </li></ul><ul><li>Bases have a pH between 7 and 14 </li></ul><ul><li>pH is a logarithmic scale like the Richter scale. Therefore 2 is 10 times more acidic than 3 and 10 is ten times more basic than 9. </li></ul><ul><li>7 is pH Neutral. </li></ul>
  68. 68.
  69. 69. Acid + Alkali <ul><li>An acid will react with an alkali to form a salt and water. </li></ul><ul><li>Neutralization reactions are a special type of double replacement reactions </li></ul><ul><li>The standard reaction is: </li></ul><ul><ul><li>acid + alkali  salt + water </li></ul></ul><ul><li>eg: </li></ul><ul><li>hydrochloric acid + sodium hydroxide  sodium chloride + water </li></ul><ul><li>s ulfuric acid + potassium hydroxide  potassium sulfate + water </li></ul><ul><li>Write out the balanced chemical equations for each reaction </li></ul>
  70. 70. <ul><li>Write the equations for the following neutralization reactions </li></ul><ul><li>Phosphoric acid and Calcium hydroxide </li></ul><ul><li>H 3 PO 4 + Ca(OH) 2  Ca 3 (PO 4 ) 2 + H 2 O </li></ul><ul><li>Nitric acid and cesium hydroxide </li></ul><ul><li>HNO 3 + CsOH  H 2 O + CsNO 3 </li></ul>Neutralization Reactions
  71. 71. Acid + Metal <ul><li>An acid will react with a metal to form a salt and hydrogen gas. </li></ul><ul><li>The standard reaction is: </li></ul><ul><ul><li>acid + metal  salt + hydrogen gas </li></ul></ul><ul><li>eg: </li></ul><ul><ul><li>hydrochloric acid + zinc  zinc chloride + hydrogen gas </li></ul></ul><ul><ul><li>s ulfuric acid + sodium  sodium sulfate + hydrogen gas </li></ul></ul><ul><li>Write out the balanced chemical equations for each reaction </li></ul>
  72. 72. Metal + Acid <ul><li>Reactions of metals with different acids form different salts: </li></ul><ul><ul><li>Hydrochloric acid  chlorides </li></ul></ul><ul><ul><li>Sulfuric acid  sulfates </li></ul></ul><ul><ul><li>Nitric acid  nitrates </li></ul></ul><ul><ul><li>Phosphoric acid  phosphates </li></ul></ul>
  73. 73.
  74. 74. Metal Oxides <ul><li>We have already seen how m any metals will react with oxygen to form a metal oxide: </li></ul><ul><ul><li>eg: sodium + oxygen  sodium oxide </li></ul></ul><ul><li>Non-metals also react with oxygen to form oxides: </li></ul><ul><ul><li>eg: carbon + oxygen  carbon dioxide </li></ul></ul>
  75. 75. Forming Basic Oxides <ul><li>Metal oxides (that are soluble) react with water to form alkaline solutions. </li></ul><ul><li>Hence these oxides are called Basic Oxides . </li></ul><ul><ul><li>eg: magnesium oxide + water  magnesium hydroxide </li></ul></ul><ul><ul><li> MgO + H 2 O  Mg (OH) 2 </li></ul></ul><ul><li>Only the very active metals form metal oxides which are soluble </li></ul>
  76. 76. Basic Oxides <ul><li>Try these: </li></ul><ul><ul><li>Sodium oxide + water  </li></ul></ul><ul><ul><li>Potassium oxide + water  </li></ul></ul><ul><ul><li>Calcium oxide + water  </li></ul></ul>Sodium hydroxide Na 2 O + H 2 O  2NaOH Potassium hydroxide K 2 O + H 2 O  2KOH Calcium hydroxide CaO + H 2 O  Ca(OH) 2
  77. 77. Forming Acidic Oxides <ul><li>Non-metal oxides react with water to form acidic solutions. </li></ul><ul><li>Hence these oxides are called Acidic Oxides . </li></ul><ul><ul><li>Eg: sulfur dioxide + water  sulfurous acid </li></ul></ul><ul><ul><li> SO 2 + H 2 O  H 2 SO 3 </li></ul></ul>
  78. 78. Acidic Oxides <ul><li>Try these: </li></ul><ul><ul><li>sulfur tri oxide + water  </li></ul></ul><ul><ul><li>carbon dioxide + water  </li></ul></ul><ul><ul><li>nitrogen di oxide + water  </li></ul></ul>sulfuric acid carbonic acid Nitric acid SO 3 + H 2 O  H 2 SO 4 CO 2 + H 2 O  H 2 CO 3 NO 2 + H 2 O  HNO 3
  79. 79. Acid Rain <ul><li>Many of these non-metal oxides are being released into the atmosphere. </li></ul><ul><li>Here they are reacting with atmospheric water to form acids. </li></ul><ul><li>They then fall as Acid Rain : </li></ul><ul><ul><li>sulfur oxides, nitrogen oxides, carbon oxides </li></ul></ul>
  80. 80. Metal Oxides + Acid <ul><li>Metal oxides will dissolve in acid to form a salt and water. </li></ul><ul><ul><li>Eg: metal oxide + acid  salt + water </li></ul></ul>copper oxide + sulfuric acid Copper sulfate + water CuO + H 2 SO 4 CuSO 4 + H2O
  81. 81. Metal Oxides in Acid <ul><li>Try these: </li></ul><ul><ul><li>zinc oxide + hydrochloric acid  </li></ul></ul><ul><ul><li>aluminium oxide + sulfuric acid  </li></ul></ul><ul><ul><li>lead oxide + sulfuric acid  </li></ul></ul>zinc chloride + water ZnO + 2HCl  ZnCl 2 + H 2 O Al 2 O 3 + H 2 SO 4  Al 2 SO 4 + H 2 0 aluminium sulfate + water lead sulfate + water PbO + H 2 SO4  PbSO 4 + H 2 O
  82. 82. Acid + Carbonate <ul><li>An acid will react with a carbonate to form a salt and carbon dioxide gas and water. </li></ul><ul><li>The standard reaction is: </li></ul><ul><ul><li>acid + carbonate  salt + carbon dioxide + water </li></ul></ul><ul><li>eg: </li></ul><ul><li>nitric acid + calcium carbonate  calcium nitrate + carbon dioxide + water </li></ul><ul><li>sulfuric acid + lead carbonate  lead sulfate + carbon dioxide + water </li></ul><ul><li>Write out the balanced chemical equations for each reaction </li></ul>
  83. 83. Ions in Solution <ul><li>Acidic solutions – contain more H + than OH - </li></ul><ul><li>Basic solutions – contain more OH - than H + </li></ul><ul><li>Neutral solutions – contain equal amounts of H + and OH - </li></ul>
  84. 84. Autoionization of Water <ul><li>H 2 O + H 2 O  H 3 O + + OH - </li></ul><ul><li>Water is the usual solvent for acids and bases </li></ul><ul><li>It produces equal numbers of H 3 O + and OH - </li></ul>
  85. 85. Redox Yr 10 Chemistry
  86. 86. Redox <ul><li>Are reactions that involve the transfer of electrons. One species is oxidised the other is reduced </li></ul><ul><li>Oxidation – loss of electrons/addition of oxygen/ increase of ox no. </li></ul><ul><li>Reduction – gain of electrons/ loss of oxygen/ decrease of ox no. </li></ul>
  87. 87. Oxidisers <ul><li>The following are oxidisers and the product that usually results </li></ul><ul><li>Oxygen water </li></ul><ul><li>Fluorine fluoride ion </li></ul><ul><li>Hydrogen ion hydrogen </li></ul><ul><li>Hydrogen peroxide water </li></ul>
  88. 88. Reducers <ul><li>The following are reducers and the product that usually results </li></ul><ul><li>carbon carbon dioxide </li></ul><ul><li>Zinc Zinc ions </li></ul><ul><li>Hydrogen water </li></ul><ul><li>Sulfur dioxide sulfate ions </li></ul>
  89. 89. Metal + Water is redox <ul><li>Only very active metals react with water to form an alkali and hydrogen gas. </li></ul><ul><li>The standard reaction is: </li></ul><ul><ul><li>Metal + water  metal alkali + hydrogen gas </li></ul></ul><ul><li>eg: </li></ul><ul><ul><li>Sodium + water  sodium hydroxide + hydrogen gas </li></ul></ul><ul><ul><li>Caesium + water  caesium hydroxide + hydrogen gas </li></ul></ul>oxidation reduction
  90. 90. Metal + Water Caesium in water
  91. 91. Complete the equation and Identify the oxidiser and reducer <ul><li>Calcium oxide + hydrogen </li></ul><ul><li>Fluorine + hydrogen </li></ul><ul><li>Sodium hydroxide + carbon </li></ul><ul><li>Carbon monoxide + oxygen </li></ul><ul><li>Chlorine + magnesium </li></ul><ul><li>Aluminium + oxygen </li></ul>
  92. 92. Oxidation <ul><li>Is the Loss of Electrons </li></ul><ul><li>is Oxidation </li></ul><ul><li>LEO </li></ul><ul><li>eg </li></ul><ul><li>Aluminium aluminium ions </li></ul>
  93. 93. Reduction <ul><li>Gain of Electrons </li></ul><ul><li>Is Reduction </li></ul><ul><li>GER </li></ul><ul><li>eg </li></ul><ul><li>Iron ions iron </li></ul>
  94. 94. Redox <ul><li>If both reduction and oxidation take place we can remember </li></ul><ul><ul><ul><ul><li>O xidation </li></ul></ul></ul></ul><ul><ul><ul><ul><li>I s </li></ul></ul></ul></ul><ul><ul><ul><ul><li>L oss of electrons </li></ul></ul></ul></ul><ul><ul><ul><ul><li>R eduction </li></ul></ul></ul></ul><ul><ul><ul><ul><li>I s </li></ul></ul></ul></ul><ul><ul><ul><ul><li>G ain in electrons </li></ul></ul></ul></ul>