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Ks4 chemical reactions

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Ks4 chemical reactions

  1. 1. © Boardworks Ltd 2003 oxidation and reduction neutralisation precipitation reversible reactions displacement reactions: metals exothermic and endothermic thermal decomposition displacement reactions: non-metals Types of chemical change
  2. 2. © Boardworks Ltd 2003 Thermal decomposition • A thermal decomposition is when heat causes a chemical to break down to simpler substances. • Compounds – but not elements - undergo thermal decomposition. • For compounds that contain metals we usually find: the more reactive the metal, the harder it is to decompose its compounds. For example: Potassium carbonate is not thermally decomposed. Calcium carbonate decomposes on strong heating Silver carbonate decomposes on gentle heating Getsharder
  3. 3. © Boardworks Ltd 2003 Generally, the more reactive the metal, the more difficult it is to decompose its compounds. Fill in the last column: easy, medium or hard. Potassium sodium calcium magnesium aluminium zinc iron copper mercury silver gold Increasing Compound How easy to decompose Mercury oxide Sodium oxide Iron oxide Silver oxide Zinc oxide easy hard medium easy medium Thermal decomposition
  4. 4. © Boardworks Ltd 2003 Thermal decomposition of carbonates • When carbonates are heated they release carbon dioxide. • This reaction is performed industrially to make calcium oxide (quicklime) from calcium carbonate (limestone). Quicklime is used to make concrete and to make calcium hydroxide (slaked lime). 1500°C limestone Hot air calcium oxide (lime) waste air and carbon dioxide Calcium Carbonate Calcium oxide Carbon dioxide +
  5. 5. © Boardworks Ltd 2003 Thermal decomposition of metal oxides • Most metal oxides are thermally stable (i.e. do not decompose when heated). • Oxides of the least reactive metals can be thermally decomposed more easily. • For example, silver oxide begins to break up at about 160o C and mercury oxide decomposes when heated strongly. Mercury Oxide Mercury oxygen+ Hg Hg Hg Hg O O O O Heat HgHg Hg HgOO O O HgHg Hg HgOO O O HgHg Hg HgOO O O HgHg Hg HgOO O O mercury oxide decomposes mercury metal and oxygen formed
  6. 6. © Boardworks Ltd 2003 Exothermic and endothermic reactions • Exothermic reactions give out heat (gets hot). • Endothermic reactions take in heat (gets cold). • Many chemical reactions need some energy to get them started (activation energy) but then the majority of chemical reactions are exothermic. Shuttle fuel burning- highly exothermic Ex = out (as in exit)Ex = out (as in exit) En = in (as in entrance)En = in (as in entrance)
  7. 7. © Boardworks Ltd 2003 • It is hard to think of examples of endothermic reactions but there are lots of exothermic ones that occur in the laboratory and in everyday life. • List 8 exothermic reactions. Some examples of exothermic reactions Burning wood on a fire Burning petrol in a car Burning butane in a cigarette lighter Burning gas in a gas hob Reacting an acid and alkali together Burning magnesium Rotting compost etc etc Exothermic and endothermic reactions
  8. 8. © Boardworks Ltd 2003 Displacement reactions: metals • These are reactions where two metals are competing to be combined with a non-metal. • The more reactive metal wins the competition and becomes part of a compound. • The less reactive metal is displaceddisplaced and so is present as the metal at the end of the reaction. Potassium sodium calcium magnesium aluminium zinc iron copper silver gold Increasingreactivity A more reactive metal (higher in the reactivity series) will displace a less reactive metal from its compound.
  9. 9. © Boardworks Ltd 2003 • Copper is quite low in the activity series. • Several metals will displace it from its compounds. magnesium copper sulphate solution magnesium sulphate solution copper metal Magnesium + Copper sulphate  Magnesium sulphate + Copper more reactive less reactive Magnesium wins the competition. Copper is displaced. K Na Ca Mg Al Zn Fe Cu Ag Au Displacement reactions: metals
  10. 10. © Boardworks Ltd 2003 Here are some actual photos. The colour changes from blue to red/black as copper metal is displaced. Magnesium + Copper sulphate  Magnesium sulphate + Copper more reactive less reactive Magnesium wins the competition. Copper is displaced photograph at end of reaction photograph at start of reactionK Na Ca Mg Al Zn Fe Cu Ag Au Displacement reactions: metals
  11. 11. © Boardworks Ltd 2003 The thermit reaction takes place between aluminium and iron oxide. It is so exothermic that molten iron is produced and the reaction is used to repair broken railway tracks. Aluminium + Iron Oxide  Aluminium Oxide + Iron more reactive less reactive Aluminium wins the competition. Iron is displaced and melts at the high temperatures produced. K Na Ca Mg Al Zn Fe Cu Ag Au Displacement reactions: metals iron oxide + aluminium powder magnesium fuse
  12. 12. © Boardworks Ltd 2003 Here is a photo of the thermit reaction being carried out in a laboratory. iron oxide + aluminium powder magnesium fuse Displacement reactions: metals
  13. 13. © Boardworks Ltd 2003 Predict which mixtures will result in a reaction. Metal  Solution Iron Magnesium Zinc Copper Iron chloride Magnesium nitrate Zinc nitrate Copper sulphate Yes Yes No No No No No Yes No Yes Yes Yes
  14. 14. © Boardworks Ltd 2003 L.O. Be able to describe some properties of elements in groups1, 7 and 8
  15. 15. © Boardworks Ltd 2003 TASKS •List all the elements that have 7 and 8 electrons in the outer shell •Write two halogen displacement reactions and state colour of final solution •Draw the atomic configuration of Helium, Neon and Chlorine and Lithium •State some of the properties of the halogens, noble gases and alkali metals 105
  16. 16. © Boardworks Ltd 2003 Displacement reactions: halogens • These are displacement reactions where two halogens are competing to be combined with a metal. • It is the more reactive halogen that will win and become part of a compound. • The less reactive halogen remains (or becomes) the element. Increasingreactivity Fluorine Chlorine Bromine Iodine • We can often tell which halogen is present from the colour of the solution.
  17. 17. © Boardworks Ltd 2003 For example, if chlorine solution is added to sodium bromide. sodium bromide solution sodium chloride solution bromine Chlorine + Sodium Bromide  Sodium Chloride + Bromine more reactive less reactive Chlorine wins the competition. Bromine (red) is displaced. F Cl Br I At chlorine solution Displacement reactions: halogens
  18. 18. © Boardworks Ltd 2003 The compounds of the halogens with Group 1 metals are all colourless. Halogen Halide  Chlorine solution Bromine solution Iodine Solution Potassium chloride Potassium bromide Potassium Iodide Br2 I2 Br2 I2 I2 I2 Predict what colour these will be after mixing.
  19. 19. © Boardworks Ltd 2003 Chlorine + When writing equations for halogen displacement reactions you must remember that – when in the form of the element – halogens exist in pairs. For chlorine and sodium bromide: + bromine sodium chloride Sodium bromide Cl2(aq) + 2NaBr(aq) 2NaCl(aq) + Br2(aq) F Cl Br I At Cl More reactive Br Less reactive Solution goes yellow/brown as bromine is produced. Displacement reactions: halogens
  20. 20. © Boardworks Ltd 2003 • If no reaction - not write “no reaction.” • Where there is a reaction write the names of the products and then write a chemical equation underneath. F Cl Br I At 1) iodine + sodium bromide solution 2) bromine + sodium chloride solution 3) chlorine + sodium iodide solution No reaction No reaction sodium chloride + iodine Cl2(g) + 2NaI(aq) 2NaCl(aq) + I2(aq) Predict whether or not a chemical reaction will occur.
  21. 21. © Boardworks Ltd 2003 Reversible and irreversible reaction • Most chemical reactions are considered irreversible in that the new products are not readily changed back into reactants. For example, once you have reacted magnesium with hydrochloric acid it is very hard to get the magnesium back. • In the equations for irreversible reactions reactants and products are joined by a “one-way arrow.” magnesium + hydrochloric magnesium + hydrogen acid chloride
  22. 22. © Boardworks Ltd 2003 • Although most chemical reactions are difficult to reverse it is possible to find reactions ranging from irreversible through to the fully reversible. • One of the best known reversible processes is heating copper sulphate. Note the double arrow symbol in the chemical equation hydrated copper sulphate Heat anhydrous copper sulphate steam CuSO4.5H20 CuSO4 + 5H2O these decompose these combine Reversible reactions
  23. 23. © Boardworks Ltd 2003 Equilibrium reactions • There are some reactions in which both the “forward and backward” reactions occur to a substantial extent under the same conditions. • These lead to equilibrium mixtures of reactants and products. • One of the most important of these reactions occurs in the Haber Process. N2(g) + 3H2(g) 2 NH3(g) However long you leave the reaction going you still get a mixture of nitrogen, hydrogen and ammonia.
  24. 24. © Boardworks Ltd 2003 Getting more product at equilibrium • There are some simple rules that can be used to move the position of an equilibrium towards reactants or products: 1. Exothermic reactions give more product at lower temperatures. (Endothermic – the opposite) 2. Increasing the pressure in gas reactions favours whichever side of the chemical equation has least gas molecules. What conditions will favour formation of more ammonia? 3H2(g) + N2 (g)  2NH3 (g) (exothermic) Low temperature High pressure
  25. 25. © Boardworks Ltd 2003 Precipitation reactions • A precipitation reaction is any reaction that produces an insoluble compound when two aqueous solutions are mixed. • It is impossible to predict whether or not we will get precipitation reactions unless we know something about the physical states (especially solubility) of the various reactants and products. Here are the symbols that we use in chemical equations to say what the physical state is: –(s) solid –(l) liquid –(g) gas –(aq) aqueous (dissolved in water)
  26. 26. © Boardworks Ltd 2003 A precipitation reaction that is often used to measure reaction rates occurs between sodium thiosulphate and hydrochloric acid. Sodium + hydrochloric sodium + sulphur + water + sulphur thiosulphate acid chloride dioxide Both reactants are colourless and dissolved (aq) Sulphur is insoluble and precipitates. This makes the solution go cloudy. aqueous aqueous aqueous solid liquid gassolid Precipitation reactions – first example
  27. 27. © Boardworks Ltd 2003 Most metal hydroxides (except sodium, potassium and calcium) are insoluble. Reactions leading to their formation give precipitates. Copper + ammonium copper + ammonium sulphate hydroxide hydroxide sulphate aqueous aqueous solid aqueoussolid Copper hydroxide is insoluble and precipitates. A pale blue solid settles at the bottom of the test tube. Both reactants are dissolved (aq). Copper sulphate is blue. Precipitation reactions – second example
  28. 28. © Boardworks Ltd 2003 Another metal hydroxide that precipitates is iron(III) hydroxide. Like many transition metals its compounds are coloured. Iron + sodium iron + sodium chloride hydroxide hydroxide chloride aqueous aqueous solid aqueoussolid Iron hydroxide is insoluble and precipitates. A deep brown solid settles at the bottom of the test tube. Both reactants are dissolved (aq) (iron chloride is yellow). Precipitation reactions – third example
  29. 29. © Boardworks Ltd 2003 Precipitation and solubility To work out whether a precipitate will be formed we need to know the solubility of the compounds that may be formed. Here are a few general guidelines: Soluble Insoluble All sodium, potassium and ammonium salts All nitratesnitrates Most chlorides, bromides and iodides. (halides) Silver and lead halides Most sulphatessulphates Lead, barium and calcium sulphates Sodium, potassium and ammonium carbonatescarbonates Most carbonates Sodium, potassium, ammonium and calcium hydroxidehydroxide Most hydroxides
  30. 30. © Boardworks Ltd 2003 To work out whether a precipitate will be formed when many ionic compounds react there are four stages: 1 Write down the names of the reactants. Sodium chloride & lead nitrate 2 Write down the ions in the reactants. (Ignore numbers) 3 Swap over the + and – ions. 4 Are the products going to be soluble or insoluble? Na+ Cl- Pb2+ NO3 - Pb2+ Cl- Na+ NO3 - Lead chloride is insoluble so there will be a precipitate Sodium + lead lead + sodium chloride nitrate chloride nitrate aqueous aqueous solid aqueoussolid Precipitation and solubility
  31. 31. © Boardworks Ltd 2003 Will there be a precipitate if I mix sodium sulphate and magnesium nitrate? Sodium nitrate & Magnesium sulphate 1 Write down the names of the reactants. 2 Write down the ions in the reactants. 3 Swap over the + and – ions. 4 Are the products going to be soluble or insoluble? Na+ SO4 2- Mg2+ NO3 - Mg2+ SO4 2- Na+ NO3 - Both the products are soluble there will be no precipitate. Sodium + magnesium magnesium + sodium sulphate nitrate sulphate nitrate aqueous aqueous aqueous aqueous
  32. 32. © Boardworks Ltd 2003 Will there be a precipitate if I mix sodium sulphate and barium nitrate? Sodium sulphate & barium nitrate1 Write down the names of the reactants. 2 Write down the ions in the reactants. 3 Swap over the + and – ions. 4 Are the products going to be soluble or insoluble? Na+ SO4 2- Ba2+ NO3 - Ba2+ SO4 2- Na+ NO3 - Barium sulphate is insoluble so there will be a precipitate. Sodium + barium barium + sodium sulphate nitrate sulphate nitrate aqueous aqueous solid aqueoussolid
  33. 33. © Boardworks Ltd 2003 Separating Precipitates – reminder!
  34. 34. © Boardworks Ltd 2003 Neutralisation reactions • AcidsAcids are substances that: • Turn litmus red. • Turn universal indicator yellow, orange or red. • Have a pH below 7. • Form solutions containing H+ ions. • BasesBases are substances that: • Turn litmus blue. • Turn universal indicator dark green, blue or purple. • React with the H+ ions in acids. • Are called alkalis if they dissolve in water. 1 2 14131211109876543 Increasingly acid Increasingly alkali
  35. 35. © Boardworks Ltd 2003 Neutralisation reactions: acids • Common AcidsCommon Acids are Name of acid Formula Strong or Weak? Sulphuric Hydrochloric Nitric Ethanoic (vinegar) H2SO4 HCl HNO3 CH3COOH strong strong strong weak • SaltsSalts Sulphuric acid Sulphates Nitric acid Nitrates Chlorides Hydrochloric acid
  36. 36. © Boardworks Ltd 2003 Neutralisation reactions: bases • Common alkalisCommon alkalis are Name of alkali Formula Strong or Weak? Sodium Hydroxide Potassium Hydroxide Calcium Hydroxide Ammonium Hydroxide NaOH KOH Ca(OH)2 NH4OH strong strong strong weak • Common basesCommon bases (neutralise acids but don’t dissolve) are Type of compound Contain React with acids to give Metal Hydroxides Metal Oxides Metal Carbonates OH- O2- CO3 2- water + a salt water + a salt water + a salt + CO2
  37. 37. © Boardworks Ltd 2003 Neutralisation reactions: acid + base A neutralisation reaction is where an acidacid reacts with a basebase to produce a neutral solution of aa saltsalt and waterwater. 1 2 14131211109876543 Increasingly acid Increasingly alkali sodium hydroxide pH 14 hydrochloric acid pH 1 neutralisation sodium chloride pH 7
  38. 38. © Boardworks Ltd 2003 Neutralisation - naming salts To name the salt formed in a neutralisation: 1 The first part of the name of the salt comes from the first name of the base So Ammonium hydroxide gives ammonium ………… Magnesium oxide gives magnesium …………... 2 The acid gives the last part of the name of the salt. So Sulphuric acid make sulphatessulphates Nitric acid makes nitratesnitrates Hydrochloric acid makes chlorideschlorides Eg. Sodium hydroxide + nitric acid forms: Calcium carbonate + sulphuric acid forms: Sodium nitrate calcium sulphate
  39. 39. © Boardworks Ltd 2003 Name the salt formed in these neutralisations: Base Acid Salt? Calcium hydroxide Hydrochloric acid Magnesium oxide Nitric acid Calcium carbonate Sulphuric acid Aluminium hydroxide Nitric acid Potassium hydroxide Sulphuric acid Calcium chloride Magnesium nitrate Calcium sulphate Aluminium nitrate Potassium sulphate + 
  40. 40. © Boardworks Ltd 2003 Neutralisation reactions: hydroxides Each OH- ion reacts with one H+ ion. Reaction with hydroxides: H+ + OH-  H2O Eg. Potassium +hydrochloric  water + potassium hydroxide acid chloride KOHOH + HHCl  HH22OO + KCl Eg. Calcium + sulphuric  water + calcium hydroxide acid sulphate Ca(OHOH)22 + HH22SO4  2H2H22OO + CaSO4
  41. 41. © Boardworks Ltd 2003 Neutralisation Reactions: oxides Neutralisation reactions usually lead to water being formed. Reaction with oxides: 2H+ + O2-  H2O Eg. Calcium + hydrochloric  water + calcium oxide acid chloride CaOO + 2H2HCl  HH22OO + CaCl2 Eg. Sodium + sulphuric  water + sodium oxide acid sulphate Na2OO + HH22SO4  HH22OO + Na2SO4
  42. 42. © Boardworks Ltd 2003 Neutralisation Reactions: carbonates Each carbonate ion provides one oxygen to join with two H+ ions. At the same time carbon dioxide is released. Carbonates: 2H+ + CO3 2-  H2O + CO2 Eg. Potassium + hydrochloric  water + carbon + potassium carbonate acid dioxide chloride K2COCO33 + 2H2HCl  HH22OO + COCO22 + 2KCl Eg. calcium + nitric  water + carbon + calcium carbonate acid dioxide nitrate CaCOCO33 + 2H2HNO3  HH22OO + CO2 +Ca(NO3)2
  43. 43. © Boardworks Ltd 2003 Neutralisation equations Eg. Potassium + hydrochloric  + hydroxide acid Complete the word equation Eg. KOH + HHCl  + water Potassium chloride HH22OO KCl Replace the words with the correct formula Check that it balancesbalances (same number of each type of atom each side). Eg. KOH + HHCl  HH22OO + KCl Reactants 1*K 1*O 2*H 1*Cl Products 2*H 1*O 1*K 1*Cl 
  44. 44. © Boardworks Ltd 2003 Eg. Magnesium + nitric  + oxide acid Complete the word equation Eg. MgO + HHNO3  + water Magnesium nitrate HH22OO Mg(NO3)2 Replace the words with the correct formula Check that it balancesbalances (Same number of each type of atom each side. Reactants 1*Mg 1*O 1*H1*H 1*NO1*NO33 Products 2*H2*H 1*O 1*Mg 2*NO2*NO33 Eg. MgO + HHNO3  H2O + Mg(NO3)2 2 2  Neutralisation equations
  45. 45. © Boardworks Ltd 2003 Write balanced equations going through the same stages as the previous examples. 1. word equation 2. formulae 3. balance a) sodium hydroxide + hydrochloric acid b) magnesium oxide + hydrochloric acid c) sodium hydroxide + sulphuric acid d) ammonium hydroxide + hydrochloric acid e) calcium hydroxide + nitric acid
  46. 46. © Boardworks Ltd 2003 • Insoluble salts can be separated by filtering. • Soluble salts are obtained by evaporating. bunsen burner evaporating basingauze tripod heat-proof mat vapour Put these in the correct order. A. Check the pH frequently by testing drops of the solution. B. Add the acid slowly to the alkali. C. When neutral pour into the evaporating basin. D. Put on safety specs. E. Allow to cool F. Heat. D B A C F E
  47. 47. © Boardworks Ltd 2003 Redox Reactions • Redox is a short way of saying: Reduction and oxidation Oxidation meant addingOxidation meant adding oxygen to a substance.oxygen to a substance. Rusting (iron becoming iron oxide) is an example of oxidation. Reduction meant takingReduction meant taking oxygen away.oxygen away. Extracting iron from iron oxide in the blast furnace is reduction. • Early on in chemistry these words had very straightforward meanings.
  48. 48. © Boardworks Ltd 2003 Redox reactions: oxidation and ions • Many redox reactions involve metals and their oxides. • Whenever metals react with oxygen they form ionic compounds and the metal loses electrons to form positively charged ions. • Eg. When magnesium burns to form magnesium oxide magnesium atoms (no charge) become magnesium ions (2+ charge) by losing 2 electrons to oxygen atoms. Mg O2 e- to give Mg2+ O2- Oxidation involves loss of electrons.
  49. 49. © Boardworks Ltd 2003 Redox reactions: electron loss • Think about what has happened to the magnesium when it reacts with oxygen. – It has been oxidised. – It has lost electrons by changing from Mg Mg2+ • Magnesium can also lose electrons to things other than oxygen (e.g. to chlorine or sulphur) and since these also involve Mg Mg2+ these too must be oxidation. Mg Oxidation is the loss of electrons. Mg2+ S2- S Mg2+ O2- O Mg2+ Cl-Cl- Cl Mg Mg2+ S2- S Mg2+ O2- O Mg2+ Cl-Cl- Cl
  50. 50. © Boardworks Ltd 2003 Redox reactions: electron gain • Exactly the same reasoning applies to reduction. • Reduction can be the removal of oxygen (e.g. from iron oxide to form iron or from aluminium oxide in the electrolysis to extract aluminium.) • When this happens the metal gets back its electrons. – Aluminium has been reduced. – Aluminium has gained electrons Al3+ O2- O2- O2- Al3+ Oxygen removed Reduction is the gain of electrons. Al Al O O 1 ½
  51. 51. © Boardworks Ltd 2003 Redox Reactions: oil rig An easy way of remembering this is “Oil RigOil Rig”! O oxidationO oxidation II isis L lossL loss R reductionR reduction II isis G gainG gain of electrons
  52. 52. © Boardworks Ltd 2003 Redox Reactions:Two for one! • Whenever something is oxidised, something else is reduced. • This should be obvious if we use the oil rig definition. • If something loses electrons – then something else must have gained them. • For example, when burning magnesium: – Magnesium loses electrons (Mg Mg2+ …..oxidation) – Oxygen gains electrons (O O2- …….reduction) The overall reaction is both RedReduction and OxOxidation = RedoxRedox
  53. 53. © Boardworks Ltd 2003 Say whether the substance in red type is oxidised or reduced. CalciumCalcium + oxygen calcium oxide ZincZinc oxide + hydrogen zinc + water CopperCopper chloride copper + chlorine AluminiumAluminium + iron oxide iron + aluminium oxide oxidised reduced reduced oxidised
  54. 54. © Boardworks Ltd 2003 If the first substance is oxidised, what has been reduced or vice versa (use whichever definition of oxidation and reduction seems easier to apply). CalciumCalcium + oxygen calcium oxide ZincZinc oxide + hydrogen zinc + water CopperCopper chloride copper + chlorine AluminiumAluminium + iron oxide iron + aluminium oxide oxidised reduced reduced oxidised Oxygen is reduced. Each oxygen atom gains 2 e-. Hydrogen is oxidised. It gains oxygen. Chlorine is oxidised. It gains an electron Cl- ½Cl2 Iron is reduced. It loses oxygen.
  55. 55. © Boardworks Ltd 2003 • Across: 5 tells us whether acid or alkali 11 reaction of an acid with a base • Down 1 a solid forms in a solution 2 loss of electrons 3 competition reaction 4 gives solutions containing H+ ions 6 to break down into smaller particles 7 removal of oxygen 8 state of balance 9 soluble base 10 ionic compound formed in neutralisations
  56. 56. © Boardworks Ltd 2003 Match them up Thermal decomposition Dehydrating copper sulphate Endothermic A solid forms within a solution Metal displacement A salt and water is formed Reversible reaction Alkali Precipitation Reaction in a state of balance Neutralisation Thermit reaction Oxidation Removal of oxygen Reduction Breaking up with heat Soluble base Takes in energy – gets cold Equilibrium Loss of electrons
  57. 57. © Boardworks Ltd 2003 When heated the orange powder erupted like a volcano producing a huge pile of green powder that had less mass than the orange material. What type of reaction is this? 1. Neutralisation 2. Thermal decomposition 3. Displacement 4. Precipitation
  58. 58. © Boardworks Ltd 2003 When the two colourless solutions mixed a yellow solid formed which sank to the bottom of the test tube. What type of reaction is this? 1. Neutralisation 2. Thermal decomposition 3. Displacement 4. Precipitation
  59. 59. © Boardworks Ltd 2003 When the copper was placed in the silver nitrate solution snow-like crystals of silver seemed to grow out from the copper. What type of reaction is this? 1. Equilibrium 2. Thermal decomposition 3. Displacement 4. Precipitation
  60. 60. © Boardworks Ltd 2003 When the washing soda was added to the lemon juice it fizzed and the pH rose towards 7. What type of reaction is this? 1. Neutralisation 2. Thermal decomposition 3. Displacement 4. Oxidation
  61. 61. © Boardworks Ltd 2003 Which of the oxides shown will thermally decompose most easily? 1. Mercury oxide 2. Potassium oxide 3. Iron oxide 4. Silver oxide
  62. 62. © Boardworks Ltd 2003 Which of the salts below might be formed when nitric acid neutralises a metal hydroxide? 1. Potassium hydroxide 2. Potassium nitrate 3. Ammonium nitrate 4. Calcium sulphate
  63. 63. © Boardworks Ltd 2003 Which of the mixtures below will result in a metal displacement reaction? 1.Potassium oxide and gold 2.Magnesium and sodium nitrate 3.Copper and silver nitrate 4.Aluminium and calcium sulphate
  64. 64. © Boardworks Ltd 2003 Which of the mixtures below will result in a non-metal displacement reaction? 1.Potassium chloride and iodine 2.Potassium bromide and iodine 3.Potassium fluoride and chlorine 4.Potassium iodide and chlorine
  65. 65. © Boardworks Ltd 2003 Which of the elements in red (below) is oxidised in the reaction? (Oil Rig!) 1.Ca + CuCuO  CaO + Cu 2.2Li + 2HHCl  2LiCl + H2 3.2AlAl + Fe2O3  Al2O3 + 2Fe 4.HNO3 + CuCuO  CuNO3 + H2O
  66. 66. © Boardworks Ltd 2003 Which compound can you be sure is soluble in water? 1. Manganese nitrate 2. Osmium iodide 3. Thallium chloride 4. Palladium sulphate

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