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Dynamic Equilibrium

16. Dynamic Equilibrium. 16.1 Irreversible and Reversible Reactions 16.2 Dynamic Nature of Chemical Equilibrium 16.3 Examples of Chemical Equilibrium 16.4 Equilibrium Law 16.5 Determination of Equilibrium Constants 16.6 Equilibrium Constant in Terms of Partial Pressures

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Dynamic Equilibrium

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  1. 16 Dynamic Equilibrium 16.1 Irreversible and Reversible Reactions 16.2 Dynamic Nature of Chemical Equilibrium 16.3 Examples of Chemical Equilibrium 16.4 Equilibrium Law 16.5 Determination of Equilibrium Constants 16.6 Equilibrium Constant in Terms of Partial Pressures 16.7 Equilibrium Position 16.8 Partition Equilibrium of a Solute Between Two Immiscible Solvents 16.9 Significances of Equilibrium Constants 16.10 Factors Affecting Equilibrium

  2. 16.1 Irreversible and Reversible Reactions

  3. 16.1 Irreversible and Reversible Reactions (SB p.88) Irreversible Reactions • Chemical reactions that take place in one direction only • It goes on until at least one of the reactants is used up • e.g. Reaction between Na and H2O to form NaOH and H2 • 2Na(s) + 2H2O(l)  2NaOH(aq) + H2(g)

  4. N2(g) + H2(g) 2NH3(g) 16.1 Irreversible and Reversible Reactions (SB p.88) Reversible Reactions • Chemical reactions that can go in either way • Do not proceed to completion • e.g. Reaction of N2 and H2 to form NH3

  5. N2(g) + H2(g) 2NH3(g) Check Point 16-1 16.1 Irreversible and Reversible Reactions (SB p.88) Reversible Reactions • Reaction going from left to right is called forward reaction • Reaction going from right to left is called backward reaction

  6. CaCO3(s) CaO(s) + CO2(g) 16.1 Irreversible and Reversible Reactions (SB p.89) Thermal Decomposition of Calcium Carbonate • CaCO3 is heated to 800 oC, it decomposes to form CaO and CO2 • CaO and CO2react to form CaCO3

  7. CH3COOH(l) + CH3CH2OH(l) CH3COOCH2CH3(l) + H2O(l) 16.1 Irreversible and Reversible Reactions (SB p.89) Esterification • Does not go to completion no matter how long the reaction mixture is heated • Ethyl ethanoate reacts with water to give ethanoic acid and ethanol

  8. 16.1 Irreversible and Reversible Reactions (SB p.89) Esterification The laboratory set-up for esterification under reflux

  9. 16.2 Dynamic Nature of Chemical Equilibrium

  10. 16.2 Dynamic Nature of Chemical Equilibrium (SB p.90) Equilibrium State • Both the forward and backward reactions proceed simultaneously • After a certain period of time, a state is reached when there is no observable changes of the reaction mixture • This state is known as equilibrium state

  11. 16.2 Dynamic Nature of Chemical Equilibrium (SB p.90) Chemical Equilibrium • At the equilibrium state: • Rate of forward reaction • = Rate of backward reaction • Concentrations of reactants and products remain constant

  12. 16.2 Dynamic Nature of Chemical Equilibrium (SB p.90) Dynamic Equilibrium An example of dynamic equilibrium

  13. H2(g) + I2(g) 2HI(g) 16.2 Dynamic Nature of Chemical Equilibrium (SB p.90) Dynamic Equilibrium Changes in concentrations of hydrogen, iodine and hydrogen iodide with respect to time

  14. H2(g) + I2(g) 2HI(g) Let's Think 1 16.2 Dynamic Nature of Chemical Equilibrium (SB p.90) Dynamic Equilibrium Changes in the rates of forward and backward reactions with respect to time

  15. 16.3 Examples of Chemical Equilibrium

  16. Br2(aq) + H2O(l) H+(aq) + Br-(aq) + HOBr(aq) yellowish brown (all products are colourless) 16.3 Examples of Chemical Equilibrium (SB p.92) Reaction of Bromine with Water • Colour of solution is related to the amount of Br2(aq) which is yellowish brown

  17. Br2(aq) + H2O(l) H+(aq) + Br-(aq) + HOBr(aq) yellowish brown (all products are colourless) 16.3 Examples of Chemical Equilibrium (SB p.92) Reaction of Bromine with Water • When an alkali is added, a colourless solution is observed • Amount of Br2(aq)  • Equilibrium position shifts to the right( H+(aq) is removed)

  18. Br2(aq) + H2O(l) H+(aq) + Br-(aq) + HOBr(aq) yellowish brown (all products are colourless) 16.3 Examples of Chemical Equilibrium (SB p.92) Reaction of Bromine with Water • When an acid is added, a yellowish brown solution is observed • Amount of Br2(aq)  • Equilibrium position shifts to the left

  19. Cr2O72-(aq) + H2O(l) 2CrO42-(aq) + 2H+(aq) yellow orange 16.3 Examples of Chemical Equilibrium (SB p.92) Interconversion of Dichromate(VI) and Chromate(VI) Ions • When an alkali is added, the solution changes from orange to yellow • Amount of Cr2O72-(aq)  • Amount of CrO42-(aq)  • Equilibrium position shifts to the right

  20. Cr2O72-(aq) + H2O(l) 2CrO42-(aq) + 2H+(aq) yellow orange Check Point 16-3 16.3 Examples of Chemical Equilibrium (SB p.92) Interconversion of Dichromate(VI) and Chromate(VI) Ions • When an acid is added, the solution changes back to orange • Amount of Cr2O72-(aq)  • Amount of CrO42-(aq)  • Equilibrium position shifts to the left

  21. 16.4 Equilibrium Law

  22. aA + bB cC + dD Equilibrium constant Equilibrium Law 16.4 Equilibrium Law (SB p.93) (Constant at a fixed temperature)

  23. H2(g) + I2(g) 2HI(g) Equilibrium Law 16.4 Equilibrium Law (SB p.93)

  24. 2H2(g) + 2I2(g) 4HI(g) Equilibrium Law 16.4 Equilibrium Law (SB p.93) • If the coefficients of the equation are multiplied by 2, Kc2 = (Kc1)2

  25. Equilibrium Law 16.4 Equilibrium Law (SB p.93) If the coefficients in a balanced chemical equation are multiplied by a common factor n, the new equilibrium constant will be the original one to the power n.

  26. 2HI(g) H2(g) + I2(g) Equilibrium Law 16.4 Equilibrium Law (SB p.94) • If the chemical equation is reversed,

  27. S8(s) + 8O2(g) 8SO2(g) Equilibrium Law 16.4 Equilibrium Law (SB p.94) • For chemical reactions involving reactants or products that are in solid or liquid states, their equilibrium concentrations are not included in the expression of Kc

  28. 16.4 Equilibrium Law (SB p.94) Equilibrium Law • Unit of Kc depends on the powers of the concentration of substances in the equilibrium law • Related to the stoichiometric coefficients in the balanced equation

  29. Let's Think 2 Check Point 16-4 16.4 Equilibrium Law (SB p.94) Equilibrium Law = no unit

  30. 16.5 Determination of Equilibrium Constants

  31. Equilibrium System: CH3COOH(l) + CH3CH2OH(l) CH3COOCH2CH3(l) + H2O(l) 16.5 Determination of Equilibrium Constants (SB p.95)

  32. Equilibrium System: CH3COOH(l) + CH3CH2OH(l) CH3COOCH2CH3(l) + H2O(l) 16.5 Determination of Equilibrium Constants (SB p.95)

  33. Equilibrium System: CH3COOH(l) + CH3CH2OH(l) CH3COOCH2CH3(l) + H2O(l) 16.5 Determination of Equilibrium Constants (SB p.96) For experiment 1:

  34. Equilibrium System: CH3COOH(l) + CH3CH2OH(l) CH3COOCH2CH3(l) + H2O(l) 16.5 Determination of Equilibrium Constants (SB p.96) For experiment 2:

  35. Equilibrium System: CH3COOH(l) + CH3CH2OH(l) CH3COOCH2CH3(l) + H2O(l) = 4.065 (no unit) 16.5 Determination of Equilibrium Constants (SB p.96)

  36. Equilibrium System: CH3COOH(l) + CH3CH2OH(l) CH3COOCH2CH3(l) + H2O(l) 16.5 Determination of Equilibrium Constants (SB p.96) • Conc. H2SO4 acts as a positive catalyst • The time required to reach the equilibrium becomes shorter

  37. 16.5 Determination of Equilibrium Constants (SB p.96) Effect of a positive catalyst on the time required to reach the equilibrium

  38. Equilibrium System: Fe3+(aq) + NCS-(aq) [Fe(NCS)]2+(aq) 16.5 Determination of Equilibrium Constants (SB p.97) • [Fe(NCS)]2+(aq) is blood-red in colour • Its concentration can be determined by measuring the absorbance of a particular wavelength using a colorimeter

  39. Equilibrium System: Fe3+(aq) + NCS-(aq) [Fe(NCS)]2+(aq) 16.5 Determination of Equilibrium Constants (SB p.97) • Known amounts of Fe3+(aq) and NCS-(aq) are mixed together • Concentrations of Fe3+(aq) (or NCS-(aq)) at equilibrium • = Initial conc. of Fe3+(aq) (or NCS-(aq)) – Conc. of Fe(NCS)2+(aq) formed

  40. Equilibrium System: Fe3+(aq) + NCS-(aq) [Fe(NCS)]2+(aq) Check Point 16-5A 16.5 Determination of Equilibrium Constants (SB p.97)

  41. Example 16-5A Example 16-5B Let's Think 3 Check Point 16-5B Example 16-5C

  42. 16.6 Equilibrium Constant in Terms of Partial Pressures

  43. aA(g) + bB(g) cC(g) + dD(g) 16.6 Equilibrium Constant in Terms of Partial Pressures (SB p.101) Equilibrium Constant in Terms of Partial Pressures For the following reaction: The equilibrium constant in terms of partial pressures (Kp) is expressed as:

  44. Example 16-6A Example 16-6B Check Point 16-6

  45. 16.7 Equilibrium Position

  46. N2(g) + 3H2(g) 2NH3(g) 16.7 Equilibrium Position (SB p.102) Equilibrium Position • e.g. The Haber process has a Kc value of 0.062 dm6 mol-2 at 500 oC • Value of Kc is a function of temperature • Not affected by the amount of gases that are mixed together at the start of reaction

  47. 16.7 Equilibrium Position (SB p.103) Equilibrium Position

  48. Example 16-7 Check Point 16-7 16.7 Equilibrium Position (SB p.103) Equilibrium Position • Each specific set of equilibrium concentrations of the reaction mixture is known as equilibrium position

  49. 16.8 Partition Equilibrium of a Solute Between Two Immiscible Solvents

  50. 16.8 Partition Equilibrium of a Solute Between Two Immiscible Solvents (SB p.104) Partition Coefficient • I2 is added to a beaker containing water and 1,1,1-trichloroethane which are immiscible • I2 will dissolve in both layers to different extent

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