Chemical Equilibra | Jamb Chemistry

Reversible Reactions

1. Reversible reactions proceed in both forward and backward directions.
2. They are represented by a double arrow ($\leftrightarrow$).
3. An example is the decomposition of ammonium chloride:
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   $NH_4Cl \leftrightarrow NH_3 + HCl$
4. Reversible reactions do not go to completion under normal conditions.
5. They reach a state of balance known as dynamic equilibrium.
6. Reactions with comparable rates in both directions are typically reversible.
7. The reaction between steam and iron:
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   $3Fe + 4H_2O \leftrightarrow Fe_3O_4 + 4H_2$
8. Another example is the dissociation of dinitrogen tetroxide:
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   $N_2O_4 \leftrightarrow 2NO_2$
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Dynamic Equilibrium

9. Dynamic equilibrium occurs when the rate of the forward reaction equals the rate of the reverse reaction.
10. At equilibrium, the concentrations of reactants and products remain constant.
11. Equilibrium does not mean the reactants and products are in equal concentrations.
12. Dynamic equilibrium can occur in both physical and chemical processes.
13. A common example is the evaporation and condensation of water in a closed system.
14. The position of equilibrium depends on initial conditions such as concentration, temperature, and pressure.
15. At equilibrium, macroscopic properties (e.g., color, pressure) remain constant.
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Factors Governing the Equilibrium Position

16. The position of equilibrium indicates the relative concentrations of reactants and products.
17. Factors affecting equilibrium include temperature, pressure, and concentration.
18. A change in temperature shifts the equilibrium position depending on the reaction's enthalpy ($\Delta H$).
19. Increasing reactant concentration shifts equilibrium to the right (toward products).
20. Increasing product concentration shifts equilibrium to the left (toward reactants).
21. In gaseous systems, increasing pressure shifts equilibrium toward the side with fewer gas molecules.
22. Decreasing pressure shifts equilibrium toward the side with more gas molecules.
23. Adding a catalyst does not affect the equilibrium position but speeds up attainment of equilibrium.
24. Removing products as they form drives the equilibrium forward.
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Le Chatelier’s Principle

25. Le Chatelier’s principle states that if a system at equilibrium is disturbed, it shifts to counteract the disturbance.
26. Increasing temperature shifts equilibrium in the endothermic direction.
27. Decreasing temperature shifts equilibrium in the exothermic direction.
28. Increasing concentration of reactants shifts equilibrium toward products.
29. Decreasing reactant concentration shifts equilibrium toward reactants.
30. Increasing pressure shifts equilibrium toward the side with fewer gas molecules.
31. Decreasing pressure shifts equilibrium toward the side with more gas molecules.
32. Adding inert gases at constant volume does not affect equilibrium position.
33. Changing the volume of the system alters pressure, affecting equilibrium.
34. In the steam and iron reaction:
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    $3Fe + 4H_2O \leftrightarrow Fe_3O_4 + 4H_2$ Increasing $H_2O$ shifts equilibrium to the right.
35. For the reaction $N_2O_4 \leftrightarrow 2NO_2$, increasing pressure shifts equilibrium to the left (fewer gas molecules).
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Equilibrium Constant K

36. The equilibrium constant ($K$) quantifies the position of equilibrium.
37. It is expressed as:
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    $K = \frac{{[products]}^{stoichiometric coefficients}}{[{reactants}]^{stoichiometric coefficients}}$
38. For $N_2O_4 \leftrightarrow 2NO_2$:
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    $K = \frac{[NO_2]^2}{[N_2O_4]}$
39. A large $K$ value indicates products are favored at equilibrium.
40. A small $K$ value indicates reactants are favored.
41. $K$ depends on temperature but not on concentrations or pressures of individual species.
42. Changing temperature alters $K$ based on whether the reaction is endothermic or exothermic.
43. Increasing temperature increases $K$ for endothermic reactions.
44. Increasing temperature decreases $K$ for exothermic reactions.
45. $K$ remains unchanged if concentration or pressure is altered without changing temperature.
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Effects of Factors on Equilibrium Constant

46. Temperature is the only factor that changes $K$.
47. For the steam and iron reaction:
    • If the reaction is exothermic, increasing temperature decreases $K$.
    • If the reaction is endothermic, increasing temperature increases $K$.
48. For $N_2O_4 \leftrightarrow 2NO_2$:
    • Increasing temperature increases $K$ as the reaction is endothermic.
49. Changes in pressure or volume do not alter $K$, but they shift the equilibrium position.
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Examples of Equilibrium

50. Dissociation of $N_2O_4$ into $NO_2$ is temperature-dependent, with more $NO_2$ at higher temperatures.
51. In the steam and iron reaction, reducing $H_2O$ shifts equilibrium to the left.
52. For $CO + H_2O \leftrightarrow CO_2 + H_2$, increasing $CO_2$ shifts equilibrium to the left.
53. The Haber process ($N_2 + 3H_2 \leftrightarrow 2NH_3$) demonstrates equilibrium manipulation using pressure and temperature.
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Significance of Dynamic Equilibrium

54. Dynamic equilibrium allows reversible reactions to stabilize under specific conditions.
55. It explains chemical processes in biological systems, like oxygen binding to hemoglobin.
56. Equilibrium principles are essential for industrial processes (e.g., Haber process).
57. Understanding equilibrium helps in optimizing chemical yields.
58. Dynamic equilibrium governs natural systems like carbon dioxide dissolution in oceans.
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Applications of Le Chatelier’s Principle

59. Le Chatelier’s principle is used in industrial synthesis to maximize product yield.
60. In the Haber process, high pressure favors $NH_3$ formation.
61. For sulfuric acid production, low temperatures favor $SO_3$ formation.
62. Removing ammonia from the system in the Haber process drives the reaction forward.
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Determining Equilibrium Constants

63. Experimental measurements of reactant and product concentrations determine $K$.
64. Temperature-dependent changes in $K$ are used to study reaction thermodynamics.
65. For $N_2O_4 \leftrightarrow 2NO_2$, $K$ increases with temperature as dissociation is endothermic.
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Solving Problems with Equilibrium Constants

66. Given $K$ and initial concentrations, equilibrium concentrations can be calculated.
67. For $N_2O_4 \leftrightarrow 2NO_2$, if $[N_2O_4] = 0.1M$ and $K = 0.36$, $[NO_2]$ can be found.
68. The expression:
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    $K = \frac{[NO_2]^2}{[N_2O_4]}$ allows determination of unknown concentrations.
69. Changes in temperature require recalculating $K$ for equilibrium adjustments.
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Significance of Equilibrium Constants

70. $K$ predicts the extent of a reaction under specific conditions.
71. A very large $K$ value indicates almost complete reaction to form products.
72. $K$ helps determine whether to favor reactants or products in industrial reactions.
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Experimental Observations

73. Equilibrium shifts can be observed using color changes (e.g., $N_2O_4 \leftrightarrow 2NO_2$).
74. Temperature effects on equilibrium can be studied in closed systems.
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Dynamic Nature of Equilibrium

75. Even at equilibrium, particles continue to react in both directions.
76. The system appears static macroscopically but is dynamic at the molecular level.
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Limitations of Le Chatelier’s Principle

77. Le Chatelier’s principle qualitatively predicts shifts but does not quantify changes.
78. It does not apply to changes in the rate of reaction caused by catalysts.
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Summary of Key Examples

79. For $N_2O_4 \leftrightarrow 2NO_2$, increasing pressure favors $N_2O_4$.
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Miscellaneous

81. Dynamic equilibrium concepts apply to environmental cycles, like the carbon cycle.
82. Equilibrium constants are used in drug formulation to ensure stability.
83. Le Chatelier’s principle guides waste reduction in industrial reactions.
84. Reversible reactions ensure chemical equilibrium in living organisms.
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Advanced Applications

85. In the contact process, $SO_2 + O_2 \leftrightarrow SO_3$, high pressure favors $SO_3$ production.
86. Biological equilibrium maintains pH balance in blood using carbonic acid and bicarbonate ions.
87. Industrial processes optimize equilibrium conditions for maximum efficiency.
88. Understanding equilibrium constants aids in designing better chemical reactors.
89. Equilibrium principles support advancements in renewable energy.
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Calculations

90. Calculating $K$ provides insight into reaction conditions.
91. For endothermic reactions, equilibrium constants increase with temperature.
92. For exothermic reactions, equilibrium constants decrease with temperature.
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Final Considerations

93. Dynamic equilibrium reflects a balance of molecular activity.
94. Le Chatelier’s principle applies universally to all equilibrium systems.
95. The equilibrium constant is temperature-dependent but not influenced by catalysts.
96. Equilibrium principles are foundational to industrial and environmental chemistry.
97. Understanding shifts in equilibrium enhances control over chemical processes.
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Conclusion

98. Reversible reactions highlight the complexity of chemical processes.
99. Dynamic equilibrium balances forward and reverse reaction rates.
100. Le Chatelier’s principle and equilibrium constants are essential for optimizing reactions.

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