Kinetics & equilibria CHEM Unit 2, Pt 5

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Kinetics & equilibria CHEM Unit 2, Pt 5
2012-04-04 11:48:05
chemistry kinetics unit equilibria

Starting with reaction rates!
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  1. Even if 2 particles collide, a reaction would not take place unless: (and what is this theory called?)
    • They collide in the right orientation
    • They collide with at least a certain minimum amount of kinetic energy.
    • (Called the Collision Theory)
  2. What is this diagram called?
    An Enthalpy profile diagram.
  3. Imagine a Maxwell-Boltzmann Distribution.
  4. Which coloured line represents the greatest temperature?
    • The Blue line.
    • A greater proportion of molecules will have energies greater than the activation energy. Because an increase in temperature means on average, molecules have more kinetic energy.
  5. Why does increase in temperature increase the rate of reaction?
    • More collisions - more chance of reaction
    • More particles with a greater energy than the activation energy.
  6. What does increase in concentration/pressure have on rate of reactions? Why?
    • Increase the rate of reaction.
    • Particles collide more.
  7. Explain what would happen to rate of reaction if surface area is increased?
    • A greater rate of reaction.
    • Because more particles will come into contact with the reactants, thus more collisions.
  8. How do catalysts increase the rate of reaction? (Simply)
    • They lower the activation energy.
    • Because they provide a different way for bonds to be broken and remade.
    • More particles would have enough energy to react.
  9. Give a definition of a catalyst.
    A substance that increases the rate of a reaction by providing an alternative reaction pathway/route with a lower activation energy. The catalyst is chemically unchanged at the end of the reaction.
  10. What is a homogenous catalyst?
    A catalyst that is in the same state as the reactants.
  11. How does a catalyst work? (Sophisticated)
    • Increases the rate of reaction by forming an activated complex - a lot energy intermediate. Provide alternative reaction route/pathway.
    • The activation energy required to form the intermediate is lower than that needed to make the products directly from the reactants.
    • Therefore lowers the activation energy and thus more collisions would result in a successful reaction.
  12. Imagine the enthalpy/reaction profile of a catalysed reaction.
  13. List some of the methods with which rates of reaction can be measured.
    • Measuring the volume of gas at intervals to create graph. The steeper the graph, the faster the rate. (eg. decomposition of hydrogen peroxide (produce oxygen gas and water) investigating effect of catalyst Mg(IV)O)
    • Measure change in mass of the reaction mixture. (as gas escapes from mixture)
    • Monitor colour change
    • Monitor time taken for precipitate that clouds a solution to make something behind it invisible. (Problem with this one is that result is subjective - hard to know exact point when mark "disappears". Also only works if reactants at beginning were transparent.)
  14. What other 3 methods can be used to finding the rate of reaction? (Fancy names)
    • Titrimetric analysis: removing small portions at a time, adding reagent to immediately stop reaction, titrating to see concentration.
    • Colorimetric analysis: use photoelectric colorimeter (can be used to calculate changes of conc)
    • Conductimetric analysis: involves measuring conductivity changes in mixture over time - reflect changes in ions present in solution. (so can be used to measure changes in conc of different components in mixture)
  15. What is a dynamic equilibrium?
    • Involves two opposing processes that occur at equal rates.
    • Constant macroscopic properties (properties external observer can see), while the microscopic processes (in a molecular scale) continue to occur.
  16. A dynamic equilibrium can only happen in a _______ system.
    closed system
  17. What can you change to alter the position of equilibrium? What does this mean? And what has NO EFFECT on position of equilibrium?
    • Concentration, pressure of temperature.
    • A point of equilibrium will be reached with different amounts of reactants and products.
    • A catalyst has NO EFFECT on position of equilibrium (it would still make equilibrium be reached faster)
  18. If there is a change in _______, _______ or ________, the equilibrium will move to help ________ the change.
    • If there is a change in concentration, pressure, or temperature, the equilibrium will move to help counteract the change.
    • eg. raise in temp - position of equilibrium will shift to try and cool down the system.
  19. What is the effect of increasing the concentration of the reactant have on the position of the equilibrium?
    • More product is formed - equilibrium shifts to the right (to counteract the increased reactants).
    • Vice versa for increasing conc of product or decreasing conc.
  20. How does change in pressure affect the position of the equilibrium?
    • Increasing pressure shifts equilibrium to the side with fewer gas molecules. This reduces the pressure.
    • Decreasing the pressure shift equilibrium to side with more gas molecules. Raises pressure.
    • Look at total number of moles on each side!
    • (Only affects equilibria involving gases)
  21. How does change in temperature affect the position of the equilibrium?
    • Increasing temp shifts the equilibrium in the endothermic direction to absorb this heat.
    • Decreasing temp shifts the equilibrium in the exothermic direction to replace the heat.
  22. What 2 equations do I need to know and to be able to test positions of equilibrium on? Colours?
    • Iodine (I) chloride (brown liquid) + chlorine gas ---- Iodine (III) chloride (yellow solid).
    • ICl(l) (brown liquid) + Cl2(g) ICl3 (yellow solid)
    • Dinitrogen tetroxide (colourless gas) ----- nitrogen dioxide (brown gas)
    • N2O4(g) (colourless gas) 2NO2(g) (brown gas)
  23. In industrial equilibrium reactions, what factors are often compromised and why?
    • Temperature and pressure.
    • Temp: for an exothermic forward reaction, lower temps mean higher yield of product. However, low temp mean slow rate of reaction. So, a compromise between maximum yield and a faster reaction is made.
    • Pressure: Increasing pressure often mean higher yield (if products have less molecules), and also mean faster rate of reaction. However, high pressures are expensive. Compromise between maximum yield and expense.