A. Reactivity of metals The reactivity series, metal oxides and extractions 1. Three metals, X, Y and Z were put into water. The reactions are shown below: a) Use the diagrams to put metals X, Y and Z in order of reactivity, starting with the most reactive. (1) Z X Y (1) b) When a metal reacts with water, it produces hydrogen gas and a metal hydroxide. Describe how you can test for the products. (2) Hydrogen gives a squeak pop with a lit splint (1) Metal hydroxide turns blue with litmus/purple with universal indicator (1) c) Give two variables that should be controlled in this investigation. (2) any two from: same temperature of the water same mass / number of moles of the metal same surface area of the metal 2. A piece of magnesium ribbon was added to dilute hydrochloric acid. a) Give two observations that are evidence for a chemical reaction taking place. (2) Any two from: gas given off / fizzing, magnesium gets smaller then disappears / thermal energy released / surroundings get warmer/ temperature increases (2) b) Write the word and balanced symbol equation, including state symbols, for the reaction. (4) magnesium + sulfuric acid magnesium sulfate + hydrogen [1] Mg(s) + H 2 SO 4 (aq) MgSO 4 (aq) + H 2 (g) [1 formula equation/1 balancing/1 state symbols] 3. The reaction between aluminium powder and iron(iii) oxide (Fe 2 O 3 ) is used in the rail industry. a) Write a word equation and balanced symbol equation for the reaction that takes place. (3) Aluminium + iron(iii) oxide aluminium oxide + iron (1) 2Al + Fe 2 O 3 Al 2 O 3 + 2Fe (2) b) Compare the reaction above to the reaction with powdered aluminium and copper(ii) oxide and explain why there is a difference. (1) Aluminium reacts more vigorously with copper(ii) oxide than with iron(iii) oxide. (1) Copper is lower than iron in the reactivity series, so the reaction is more violent. (1)
4. A student carried out some displacement reactions using three metals and three sulfate solutions. The results are shown in the table below: Iron sulfate (FeSO 4 ) Copper sulfate (CuSO 4 ) Magnesium sulfate (MgSO 4 ) Iron (Fe) Copper (Cu) Magnesium (Mg) a) i) Explain what is observed when iron reacts with copper sulfate. (2) iron nail turns (from grey to) brown (1) solution turns (from blue to) pale green (1) ii) HT: Write an ionic equation for the reaction between iron and copper sulfate solution. (2) Fe + Cu 2+ Fe 2+ + Cu (1 mark for reactants and 1 mark for products) b) Explain why there is no observation between copper and iron sulfate. (2) Copper does not react with iron sulfate (1) (because) copper is less reactive than iron (1) c) i) Explain what is observed when magnesium reacts with iron sulfate. (2) magnesium ribbon turns (from silver to) grey (1) solution turns (from colourless to) pale green (1) ii) HT: Write a half equation to show the reduction of iron ions (Fe 2+ ) when magnesium reacts with iron sulfate. Use the half equation to explain why Fe 2+ ions are reduced. Fe 2+ + 2e- Fe (1) (Fe 2+ ions are reduced because they have) gained electrons (1)) B. Reactions of metals part 1 Metals & acids and strong & weak acids (HT) 1. Zinc reacts with hydrochloric acid. a) Write a word and a balanced symbol equation, with state symbols, to show this reaction. (3) zinc + hydrochloric acid zinc chloride + hydrogen (1) Zn(s) + 2HCl(aq) ZnCl 2 (aq) + H 2 (g) [2] b) HT: write an ionic equation for the reaction. (2) Zn(s) + 2H + (aq) Zn 2+ (aq) + H 2 (g) [2] c) HT: Give both half equations to show the electron transfers taking place. (2) Zn Zn 2+ + 2e [1] 2H + + 2e H 2 [1] d) HT: Explain why this reaction is a redox reaction. (4) Each zinc atom loses two electrons, [1] to two hydrogen ions, [1] zinc atoms oxidised as lose electrons, [1] hydrogen ions reduced as gain electrons [1] 2. a) HT: Explain why ethanoic acid (found in vinegar) is described as a weak acid, whereas nitric acid is a strong acid. (4) Any 4 from (accept dissociate for ionise) Ethanoic acid does not ionise completely when added to water, [1] reaction is reversible, [1] majority of molecules remain intact, [1] only a small fraction form H + (aq) ions [1]. Therefore,
ethanoic acid does not produce as high a concentration of H + (aq) ions as a strong acid of equal concentration. [1] Nitric acid is a strong acid, because its molecules ionise completely in water [1] b) HT: Magnesium reacts with ethanoic acid and nitric acid. What difference would you see if magnesium carbonate was reacted with ethanoic acid of the same concentration as nitric acid? (2) Effervescence/ fizzing with ethanoic acid (1) would be slower (than with nitric acid) (1) C. Reactions of metals part 2 ph scale, neutralisation, salt and titration (chem) 1. Magnesium carbonate reacts with nitric acid. The equation is shown below: MgCO 3 (aq) + 2 HNO 3 (aq) ---> Mg(NO 3 ) 2 (aq) + H 2 O(l) + CO 2 (g) Extended writing: Plan a method to produce dry crystals of magnesium nitrate. (6) 0 marks Level 1 (1 2 marks) Level 2 (3 4 marks) Level 3 (5 6 marks) No relevant content Magnesium carbonate is added to nitric acid. Excess magnesium carbonate is added to dilute nitric acid. The excess magnesium carbonate is filtered off. Excess magnesium carbonate is added to dilute nitric acid. The excess magnesium carbonate is filtered off. The saturated solution is heated for a few minutes to evaporate some of the water and left in a warm place until dry crystals are formed or the crystals are dried on filter paper. 2. i) HT: You are given a 0.50mol/dm 3 solution of nitric acid (strong) and ethanoic acid (weak). Calculate the concentration of each acid, giving your answer in g/dm 3 to 3 significant figures. (2) 31.5 g/dm 3 nitric acid (1) 30.0 g/dm 3 ethanoic acid (1) ii) The solution of ethanoic acid has a ph of 4 and the solution of nitric acid a ph of 1. How many times greater is the concentration of H + ions in the nitric acid compared to the concentration in the ethanoic acid? 1,000 times greater (1)
3. CHEMISTRY ONLY: a) A titration is carried out between hydrochloric acid and sodium hydroxide. The following results show the volumes of acid added to neutralize the sodium hydroxide. Volume of acid added (cm 3 ) Rough Trial 1 Trial 2 Trial 3 15.70 15.30 15.25 15.30 Calculate the mean volume of solution added and explain your answer. (3) (15.3 + 15.25 + 15.3) 3 = 15.28 cm 3 (2 1 mark for method and 1 mark for answer with unit) Volume of rough titration has been disregarded) (1) b) HT: In another investigation, it takes 27.00cm 3 of hydrochloric acid to neutralise 25.00cm 3 of sodium hydroxide at a concentration of 1.0 mol/dm 3. Calculate the concentration of hydrochloric acid in g/cm 3. (4) Number of moles of sodium hydroxide = concentration volume = 1 mol/dm 3 (25 1000) dm 3 = 0.025 mol (1) The equation for the reaction shows that 1 mole of sodium hydroxide reacts with 1 mole of hydrochloric acid. So there is 0.025 mol of HCl in 27 cm 3 of solution. So the concentration of HCl in mol/dm 3 = number of moles volume = 0.025 mol (27 1000) dm 3 = 0.925 mol/dm 3 (1) The mass of 1 mole of HCl is (1 + 35.5) = 36.5 g (1) So the concentration in g/dm 3 = 36.5 g/mol 0.925 mol/dm 3 = 33.8 g/dm 3 (1) D. Electrolysis part 1 Electrolysis of a molten and solution 1. The diagram shows how molten lead bromide is electrolysed. Lead bromide contains Pb 2+ and Br - ions. a) Explain why molten lead bromide conducts electricity. (1) The ions are free to move/carry the electrical charge (1)
b) HT: Write the half equations, including the state symbols for the changes at the anode and cathode. (4) 2Br - Br 2 + 2e - (2) Pb 2+ + 2e - Pb (2) 2. The diagram shows how sodium chloride is electrolysed in the laboratory: a) Name the products A and B? (2) A: chlorine (1) B: hydrogen (1) b) Give one use of substance A. (1) To make bleach/plastics/used for sterlising water (any one for 1 mark) c) A few drops of universal indicator was added to the solution after the reaction and it turned blue. Explain why. (2) The solution is alkaline (1) because sodium hydroxide (ions) is produced (1) d) HT: Write the half equations, including the state symbols for the changes as the anode and cathode. (4) anode: 2Cl (aq) Cl 2 (g) + 2e or 2Cl (aq) 2e Cl 2 (g) (2) cathode: 2H + (aq) + 2e H 2 (g) (2) E. Electrolysis part 2 Using electrolysis to extract metals 1. Aluminium is extracted from Aluminium oxide (Al 2 O 3 ) by electrolysis. Aluminium contains Al 3+ and O 2- ions. a) Suggest why aluminium was only discovered in the 1800s, despite it being a common element in the Earth s crust. (3) Aluminium in its ore is bonded to other elements in compounds difficult to break down. [1] Aluminium could only be extracted once electrical cells used to pass electricity were made [1] Aluminium could then be made from one of its molten compounds by electrolysis. [1]
The following diagram shows how aluminium is extracted from aluminium oxide by electrolysis: b) Why is molten aluminium oxide dissolved in molten cryolite? (2) To melt at a lower temperature (1) and therefore save on costs (1) c) Why are the carbon anodes replaced regularly in the industrial electrolysis of aluminium oxide? (2) Any two from Oxygen produced at the hot carbon anodes [1] reacts with the carbon to produce carbon dioxide, burning away anodes. [1] d) HT: Write half equations for the changes at each electrode and explain which of the ions are oxidised and reduced. (4) cathode: Al 3+ + 3e Al [1] Al 3+ ions gain electrons - reduced [1] anode: 2O 2 O 2 + 4e [1] O 2 ions lose electrons - oxidised [1]
A. Energy Changes part 1 Exothermic and Endothermic Reactions 1. How do we describe energy conservation in reactions? (2) Amount of energy in the Universe at the end of a reaction [1] Is the same as before the reaction takes place [1] 2. What is an exothermic reaction? (2) One that transfers energy to the surroundings [1] The temperature of the surroundings increases [1] 3. Give two examples of exothermic reactions (2) Any two from combustion/oxidation/neutralisation (accept examples e.g. hydrochloric acid and sodium hydroxide) [2] 4. What is an endothermic reaction? (2) One that takes in energy from the surroundings [1] The temperature of the surroundings decreases [1] 5. Give one example of an endothermic reaction. (1) Any one from thermal decompositions/sports injury packs (accept examples e.g. citric acid and sodium hydrogen carbonate) [1] 6. Extended response question: A student wishes to investigate which of three metals will give the largest exothermic reaction when they react with hydrochloric acid. Describe how she would carry out the experiment making sure it was a fair test. (6) Level 3 (5-6 marks) Complete description of experiment and fair test Level 2 (3-4 marks) Partial description of experiment and fair test Level 1 (1-2 marks) Limited description of experiment and fair test The reaction will be exothermic therefore the rise in temperature will be measured [1] the greatest rise in temperature signifies the most exothermic [1] Mass of metal/surface area must be the same [1] Temperature of acid must be the same [1] Concentration of acid must be the same [1] Mention of agitation or stirring [1]
B. Energy Changes part 2 Reaction profiles 7. What is activation energy? (1) The energy needed for a reaction to occur [1] 8. Extended response question: The following reaction is exothermic: H 2 (g) + F 2 (g) 2HF (g) Draw a reaction profile to show this reaction, including the relative energies of the reactants and products, the activation energy and the overall energy change (6) 9. (HT only) Below is a balanced chemical reaction between nitrogen gas and oxygen gas to produce nitrogen monoxide. The bond energy between two nitrogen atoms is 942kJ/mol, between two oxygen atoms is 494 Kj/mol and between a nitrogen atom and an oxygen atom is 607kJ/mol. Calculate the overall energy change for this reaction. (4) N 2 + O 2 2NO Bond energies for the reactants 1,436 kj/mol [1] Bond energies for the products 1,214 kj/mol [1] Answer = + [1] 222kJ/mol [1] 10. (HT only) Draw a reaction profile for the above reaction. (4) Correctly labelled axes [1] energy of products above reactants [1] both reactants and products labelled on correct lines [1] curved line showing activation energy [1] 11. Describe the energy release in an exothermic reaction in terms of bond energies. (2) Energy released from forming bonds is greater [1] than energy needed to break bonds [1]
C. Energy Changes part 3 Cells, Batteries and Fuel Cells 12. A student has a number of cells all with a potential difference of 1.5V. Explain how they could construct a battery with a potential difference of 12V (2) 8 cells [1] All cells in series [1] 13. Below is a table of standard electrode potentials Metal electrode Standard electrode potential E O in volts Calcium -2.76 Magnesium -2.38 Aluminium -1.66 Zinc -0.76 Iron -0.41 Lead -0.13 When a cell contains an aluminium electrode and a zinc electrode the potential difference of the cell is 0.90 Volts. Calculate the potential difference of a cell with magnesium and iron electrodes. (2) Idea of 2.38V 0.41V [1] 1.97 V [2] 14. How would you increase the potential difference of this cell? (1) Change the magnesium electrode for a calcium electrode/change the iron electrode for a lead electrode [1] 15. (HT) Write the half equations for the electrode reactions in a hydrogen fuel cell Cathode (2) 2H 2 (g) 4H + (aq) + 4e - correct chemicals including electrons [1] balanced [1] Anode (2) 4H + (aq) + O 2 (g) + 4e - 2H 2 O (g) correct chemicals including electrons [1]balanced [1] 16. Hydrogen fuel cells are used on space craft to produce electricity. Give an advantage and a disadvantage of using hydrogen fuel cells on a space craft. (2) One advantage no pollution/ water can be drunk by astronauts [1] One disadvantage hydrogen fuel takes a lot of space/explosive [1]