| 1. | B | |
| 2. | A | |
| 3. | C | |
| 4. | C | |
| 5. | C | |
| 6. | C | |
| 7. | ||
| a) i) | ||
 |
||
| ii) | ||
| Filtration | ||
| iii) | ||
| Neutralisation | ||
| b)i) | ||
| hydrogen chloride | ||
| ii) | ||
| Hydrogen chloride shifts the equilibrium position to the left. | ||
| 8. | ||
| a)i) | ||
| It is endothermic | ||
| ii) | ||
| A decrease in pressure will cause an increase in the proportion of synthesis gas [or vice versa] (1). According to Le Chatelier’s Principle, when at equilibrium, a closed system will shift the position of the equilibrium to undo any change imposed. As there are more moles of gas on the right hand side, decreasing pressure will shift the equilibrium to the right. (1) | ||
| iii) | ||
| It will be increased | ||
| iv) | ||
| No change | ||
| 9. | ||
| a) | ||
| increased | ||
| b) | ||
| no change | ||
| c) | ||
| decreased | ||
| 10. | ||
| a) | ||
| Natural gas | ||
| b) | ||
| Sulfur dioxide may be produced causing acid rain pollution. | ||
| d) | ||
| According to Le Chatelier’s, the increase in pressure will cause the equilibrium to establish to the right (1), increasing the yield of methanol(1) |
| 11. | |
 |
|
| 12. | |
 |
|
| 13. | |
| a) | |
| Circle OXYGEN and NITROGEN | |
| b) | |
| Both processes are endothermic (1), According to Le Chatelier’s principle, the higher temperature will favour the endothermic/ reverse reaction (1). | |
| c) | |
| Contact Process: | |
| There are 3 molecules of gas on the left-hand side of the equation, but only 2 on the right.According to Le Chatelier’s Principle, if you increase the pressure the system will respond by favouring the reaction which produces fewer molecules. [That will cause the pressure to fall again.]
In order to get as much sulphur trioxide as possible in the equilibrium mixture, you need as high a pressure as possible. High pressures also increase the rate of the reaction [gas reactants closer together]. However, the reaction is done at pressures close to atmospheric pressure! BUT At atmospheric pressure, there is already a 99.5% conversion of sulphur dioxide into sulphur trioxide. The reason for atmospheric pressure is a cost benefit consideration. The very small improvement that you could achieve by increasing the pressure isn’t worth the expense of producing those high pressures. |
|
| Haber Process: | |
| According to Le Chatelier’s Principle, if you increase the pressure the system will respond by favouring the reaction which produces fewer molecules. [That will cause the pressure to fall again.]
In order to get as much ammonia as possible in the equilibrium mixture, you need as high a pressure as possible. In addition, the higher the pressure the better in terms of the rate of a gas reaction [gas reactants closer together]. BUT Strong pipes and vessels to withstand the very high pressure are expensive to build. Energy used in pressuring and maintaining the pressure increase the running costs of the plant. A compromise between these conflicting consideration is found 400 atmospheres (200 -450) is a compromise pressure chosen on economic grounds. If the pressure used is too high, the cost of generating it exceeds the price you can get for the extra ammonia produced. |
|
| 14. | |
 |
|
| 15. | |
 |
|
| Answer: The equilibrium is shifted to the right, increasing the concentraion of the hypochlorite ion.
Explanation: [The greater the concentration of the hypoclorite ion, the greater the bleaching efficiency. Adding NaOH cause a neutralisation to occur, removing hydrogen ions from the equilibrium reaction. This will cause, by Le Chatelier’s principle, the equilibrium to shift to the right to undo the change, increasing the concentration of the hypochlorite ion. ] |
|
| 16. | |
 |
|
| 17. | |
 |
|
| i) | EXOTHERMIC |
| ii) | No Change |
| 18. | |
 |
|
| The forward reaction is exothermic. If cooled, the equilibrium position will shift to undo the change (Le Chatelier’s Principle), so generates heat, by increasing the rate of the forward reaction (exothermic), shifting the equilibrium to the right, and increasing the production of methane | |
| On cooling, to below 100C, the water condenses, changing from gas to liquid, and so creates fewer moles of gas on the product side. The equilibrium position will shift to undo that change (Le Chatelier’s Principle), by increasing the moles of gas on the product side, increasing the production of methane | |
| 19. | |
| a) | Rates of forward and reverse reactions are equal. |
| b) | Iodine will move from the chloroform partition into the KI partition. |
| c) | 30gl-1 |
| 20. | |
| i) | |
 |
|
| ii) | Remain unchanged |
