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Understanding Chemical Equilibrium: Dynamics and Disturbances**

Chemical equilibrium occurs when the rates of the forward and reverse reactions are equal, resulting in constant concentrations of reactants and products. This dynamic state is essential for understanding reversible reactions, such as melting ice and reactions in rechargeable batteries. The equilibrium constant (K) varies with temperature and describes the ratio of concentrations at equilibrium. Disturbances, such as adding or removing reactants/products and changing temperature or pressure, can shift the equilibrium position according to Le Chatelier's Principle, impacting the system's balance.

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Understanding Chemical Equilibrium: Dynamics and Disturbances**

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  1. EQUILIBRIUM 2 • REACTION YIELDS  

  2. Equilibrium • Very few reactions proceed unhindered to completion. • Some begin reversing as soon as products are present. • Examples of reversible reactions • Melting ice block • H2O (s) H2O (l) • Ni-Cad rechargeable batteries 

  3. Equilibrium • Chemical reactions that consist of two opposing processes (forward and reverse reactions) will eventually reach an equilibrium. • The state of equilibrium is characterized by the forward and reverse reactions proceeding at the same rate • i.e. reactions do not stop ‑ we have a dynamic situation

  4. Dynamic Equilibrium • Characterized by the following criteria • amounts and concentrations of substances remain constant • total gas pressure remains constant • temperature remains constant • the reaction is incomplete (all substances involved in the reaction are present)

  5. Equilibrium N2 + 3H2 2NH3 rate 2NH3 N2 + 3H2 time Equilibrium first established Variation of the rates of the forward and reverse reactions with time

  6. The Equilibrium Law • For the general equilibriumxA + yB pC + qD • It can be stated = constant = K [C]p [D] q [A]x [B]y

  7. Equilibrium Law • K allows for the evaluation of the concentration fraction at any time. • When the system is at equilibrium the concentration fraction is constant ‑ so called the equilibrium constant (K). • For a particular reaction, K is constant for all equilibrium mixtures (provided temperature remains constant)

  8. Information From The Equilibrium Constant • If K is about 104 to 10–4 there will be significant amounts of both reactants and products present at equilibrium • If K is very large (> 104) the equilibrium mixture consists mostly of products • If K is very small (< 10–4 ) the equilibrium mixture consists mostly of reactants

  9. Le Chatelier's Principle  • Whenever a change is made to a system at equilibrium, the equilibrium position will shift to partially oppose the change 

  10. Disturbing Equilibrium • There are 4 major means of disturbing a system at equilibrium • Adding or removing a reactant or product • Changing the pressure by changing the volume (gases only) • Dilution (for solutions only) • Changing the temperature

  11. Disturbing Equilibrium • Addition of a catalyst will increase both the rate of the forward and reverse reactions equally • It will simply reduced the time taken to reach equilibrium.

  12. Effect of Temperature on Equilibria • As temperature INCREASES • For exothermic reactions, value of K decreases and amounts of products decrease • For endothermic reactions, value of K increases and amounts of products increase

  13. Effect of Temperature on Equilibria • The value of K depends on temperature • When stating a value of K, the temperature at which the constant was calculated must also be stated • Temperature is the only change that can be made to a system at equilibrium that will actually change the equilibrium constant (ie K is temperature dependant)

  14. Consider the Reaction • N2(g) + 3H2(g) 2NH3(g)

  15. Effect on Equilibrium of Adding / Removing Reactant or Product • N2(g) + 3H2(g) 2NH3(g)

  16. Effect of Adding Nitrogen • Causes the rate of the forward reaction to increase • More ammonia is formed [NH3] increases • This causes the rate of the back reaction to increase to re form more N2 and H2

  17. Effect of Adding Nitrogen [N2] concentration [H2] [NH3] time Initial equilibrium

  18. Effect of Adding Nitrogen [N2] concentration [H2] [NH3] time Initial equilibrium Nitrogen added

  19. Effect of Adding Nitrogen [N2] concentration [H2] [NH3] System returns to equilibrium time Initial equilibrium Nitrogen added

  20. Effect of Adding Nitrogen [N2] concentration [H2] [NH3] System returns to equilibrium time Initial equilibrium Nitrogen added New equilibrium established

  21. Effect of Adding Hydrogen [N2] concentration [H2] [NH3] time Hydrogen added Initial equilibrium

  22. Effect of Adding Hydrogen [N2] concentration [H2] [NH3] time Hydrogen added Initial equilibrium

  23. Effect of Adding Hydrogen [N2] concentration [H2] [NH3] System returns to equilibrium time Hydrogen added Initial equilibrium

  24. Effect of Adding Hydrogen [N2] concentration [H2] [NH3] System returns to equilibrium time Hydrogen added Initial equilibrium New equilibrium established

  25. Effect of Adding Product • Leads to Formation of more Reactants • A nett back reaction occurs

  26. Effect of Adding Ammonia [N2] concentration [H2] [NH3] time Initial equilibrium

  27. Effect of Adding Ammonia [N2] concentration [H2] [NH3] time Ammonia added Initial equilibrium

  28. Effect of Adding Ammonia [N2] concentration [H2] [NH3] System returns to equilibrium time Ammonia added Initial equilibrium

  29. Effect of Adding Ammonia [N2] concentration [H2] [NH3] System returns to equilibrium time Ammonia added Initial equilibrium New equilibrium established

  30. Effect of Changing Reactant / Product • Addition of Reactant leads to more Products being formed (Nett Forward Reaction) • Addition of Product leads to more Reactants being formed (Nett Back Reaction) • Removal of Reactant leads to less Products being formed (Nett Back Reaction) • Removal of Product leads to less Reactants being formed (Nett Forward Reaction)

  31. Changing Pressure • Pressure can be changed by increasing or decreasing the volume of the container while keeping the temperature constant. • Need to examine 2 examples

  32. Changing Pressure • 2SO2(g) + O2(g) 2SO3(g) • 3 gas particles 2 gas particles • A nett forward reaction • involves a reduction in the number of gas particles, • so a reduction in pressure • A nett back reaction • Involves an increase in the number of gas particles • So an increase in pressure

  33. Changing Pressure • 2SO2(g) + O2(g) 2SO3(g) • 3 gas particles 2 gas particles • Using Le Chatelier’s Principle • An increase in pressure will lead to • Be adjusted by a reduction in pressure • A nett forward reaction will occur increasing the amount of sulphur trioxide present at equilibrium

  34. Changing Pressure • 2SO2(g) + O2(g) 2SO3(g) SO2 5 O2 3 SO3 1 TOTAL 9

  35. Changing Pressure • 2SO2(g) + O2(g) 2SO3(g) Nett forward reaction Increased pressure SO2 1 O2 1 SO3 5 TOTAL 7

  36. Changing Pressure • N2O4(g) 2NO2(g) • 1 gas particles 2 gas particles • Colourless Brown • A nett forward reaction • involves an increase in the number of gas particles, • so an increase in pressure • A nett back reaction • Involves a decrease in the number of gas particles • So a decrease in pressure

  37. Changing Pressure • N2O4(g) 2NO2(g) • An equilibrium mixture of the gases was compressed • Initially darkened - [NO2] increases • Then colour of gas mixture fades • Nett backward reaction

  38. Changing Pressure • N2O4(g) 2NO2(g) [N2O4] concentration [NO2] time Initial equilibrium

  39. Changing Pressure • N2O4(g) 2NO2(g) [N2O4] concentration [NO2] time Initial equilibrium Increase of pressure

  40. Changing Pressure • N2O4(g) 2NO2(g) [N2O4] concentration [NO2] time Initial equilibrium Increase of pressure System returns to equilibrium

  41. Changing Pressure • N2O4(g) 2NO2(g) [N2O4] concentration [NO2] time Initial equilibrium Increase of pressure System returns to equilibrium New equilibrium established

  42. Adding an inert gas • Total pressure of equilibrium system can be changed without changing the volume of the container by adding an inert gas • There is no increase in concentrations of reactants or products • No change in equilibrium

  43. Dilution • When dilution occurs, a net reaction results which produces the greater number of particles • The effect of diluting the solution by adding water is • A net reaction in the direction that produces more particles

  44. Dilution • Fe3+(aq) + SCN–(aq) Fe(SCN)2+(aq) • 2 particles in soln 1 particle in soln • Dilution of this equilibrium will result in a nett back reaction • Results in an increase of [Fe3+] and [SCN–]

  45. Change in Temperature • Using Le Chatelier’s Principle • Exothermic reaction can be written as • Reactants Products + energy • Heating increases the energy of the substances • Principle says the reaction will oppose an increase in energy by removing energy • A nett back reaction occurs • Less product and more reactants now present

  46. Change in TemperatureExothermic A + B C + D [D] concentration [C] [B] [A] time Initial equilibrium

  47. Change in TemperatureExothermic A + B C + D [D] Temperature increases concentration [C] [B] [A] time System returns to equilibrium Initial equilibrium Temperature increases

  48. Change in TemperatureExothermic A + B C + D [D] Temperature increases concentration [C] [B] [A] time System returns to equilibrium Initial equilibrium Temperature increases New equilibrium established

  49. Change in TemperatureEndothermic A + B C + D [D] concentration [C] [B] [A] time Initial equilibrium

  50. Change in TemperatureEndothermic A + B C + D [D] Temperature increases concentration [C] [B] [A] time System returns to equilibrium Initial equilibrium Temperature increases

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