Energy Changes | Jamb Chemistry

Energy Changes in Physical and Chemical Changes

1. Enthalpy change ($\Delta H$) represents heat absorbed or released in a process.
2. A positive $\Delta H$ indicates energy absorption (endothermic).
3. A negative $\Delta H$ indicates energy release (exothermic).
4. Phase changes involve $\Delta H$, e.g., melting ($+\Delta H$) and freezing ($-\Delta H$).
5. Chemical reactions may release energy (e.g., combustion) or absorb energy (e.g., decomposition).
6. Bond breaking absorbs energy, contributing to $+\Delta H$.
7. Bond formation releases energy, contributing to $-\Delta H$.
8. Heat of reaction depends on the difference between reactants' and products' enthalpies.
9. The reaction enthalpy is calculated as:
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   $\Delta H = \Sigma H_{products} - \Sigma H_{reactants}$
10. Energy changes are essential for understanding reaction mechanisms.
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Dissolution of Substances in/with Water

11. Dissolution involves interactions between solute particles and water molecules.
12. Sodium ($Na$) reacts exothermically with water ($-\Delta H$):
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    $2Na + 2H_2O \rightarrow 2NaOH + H_2$
13. Potassium ($K$) reacts more vigorously than sodium, releasing more energy.
14. Sodium hydroxide ($NaOH$) dissolves exothermically:
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    $NaOH + H_2O \rightarrow Na^+ + OH^- + heat$
15. Ammonium chloride ($NH_4Cl$) dissolves endothermically ($+\Delta H$):
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    $NH_4Cl \rightarrow NH_4^+ + Cl^- (absorbing heat).$
16. The enthalpy of solution depends on solute-solvent interactions.
17. Exothermic dissolution warms the solution (e.g., NaOH in water).
18. Endothermic dissolution cools the solution (e.g., $NH_4Cl$ in water).
19. Dissolution may also involve entropy changes, driving the process.
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Endothermic and Exothermic Reactions

20. Endothermic reactions absorb heat from the surroundings ($+\Delta H$).
21. Examples of endothermic processes: photosynthesis, melting, evaporation.
22. Exothermic reactions release heat to the surroundings ($-\Delta H$).
23. Examples of exothermic processes: combustion, condensation, freezing.
24. In endothermic reactions, products have higher enthalpy than reactants.
25. In exothermic reactions, products have lower enthalpy than reactants.
26. Calorimetry measures the heat changes of reactions.
27. The sign of $\Delta H$ determines heat flow direction.
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Entropy as an Order-Disorder Phenomenon

28. Entropy ($S$) measures the disorder or randomness of a system.
29. Systems tend to move toward higher entropy (greater disorder).
30. Mixing gases increases entropy as molecules spread randomly.
31. Dissolution of salts increases entropy due to ion dispersion.
32. Phase transitions affect entropy (e.g., melting increases $\Delta S$, freezing decreases $\Delta S$).
33. Entropy of gases is higher than liquids, which is higher than solids.
34. Chemical reactions with more moles of products than reactants increase entropy.
35. Decreasing temperature reduces entropy, while increasing temperature raises it.
36. Spontaneous processes often involve entropy increases.
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Spontaneity of Reactions

37. A reaction is spontaneous if it occurs without external intervention.
38. Spontaneity depends on changes in enthalpy ($\Delta H$) and entropy ($\Delta S$).
39. Gibbs free energy ($\Delta G$) determines spontaneity:
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    $\Delta G = \Delta H - T\Delta S$
40. A negative $\Delta G$ indicates a spontaneous reaction.
41. A positive $\Delta G$ indicates a non-spontaneous reaction.
42. Spontaneous reactions release free energy into the surroundings.
43. Temperature can influence spontaneity by affecting $T\Delta S$.
44. Exothermic reactions with increased entropy ($-\Delta H, +\Delta S$) are always spontaneous.
45. Endothermic reactions with decreased entropy ($+\Delta H, -\Delta S$) are non-spontaneous.
46. Reactions with opposing $\Delta H$ and $\Delta S$ depend on temperature.
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Gibbs Free Energy and Equilibrium

47. Gibbs free energy ($\Delta G^\circ$) combines enthalpy and entropy effects.
48. At equilibrium, $\Delta G^\circ = 0$, indicating no net reaction.
49. $\Delta G < 0$: Reaction is spontaneous in the forward direction.
50. $\Delta G > 0$: Reaction is non-spontaneous, favoring the reverse reaction.
51. The equilibrium constant ($K$) relates to $\Delta G^\circ$:
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    $\Delta G^\circ = -RT \ln K$
52. Spontaneity aligns with the sign of $\Delta G^\circ$.
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Relationship

53. $\Delta G^\circ = \Delta H^\circ - T\Delta S^\circ$ links energy, entropy, and spontaneity.
54. A negative $\Delta G^\circ$ reflects favorable enthalpy and entropy conditions.
55. Endothermic reactions can be spontaneous if $T\Delta S > \Delta H$.
56. Exothermic reactions are generally spontaneous if entropy increases.
57. At high temperatures, $T\Delta S$ dominates $\Delta H$.
58. At low temperatures, $\Delta H$ dominates $T\Delta S$.
59. Reaction feasibility depends on the interplay of $\Delta H$ and $\Delta S$.
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Simple Illustrations of Entropy

60. Mixing gases demonstrates entropy increase as molecules spread.
61. Dissolution of salts shows entropy increase due to ion dispersion.
62. Crystallization decreases entropy as ions organize into a lattice.
63. Combustion increases entropy as solid/liquid fuels convert to gases.
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Solve Simple Problems Using Gibbs energy change

64. Gibbs energy change calculates reaction spontaneity:
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    $\Delta G = \Delta H - T\Delta S$
65. Example: Given $\Delta H = -50kJ/mol$, $\Delta S = 100J/mol·K$, $T = 298K$:
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    $\Delta G = -50 - (298 \times 0.1) = -79.8kJ/mol.$ Reaction is spontaneous.
66. $\Delta G$ changes with temperature, affecting spontaneity.
67. Positive $\Delta H$ and $\Delta S$ require high $T$ for spontaneity.
68. Negative $\Delta H$ and $\Delta S$ require low $T$ for spontaneity.
69. At equilibrium:
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    $\Delta G = 0 \Rightarrow \Delta H = T\Delta S$
70. Reactions with $\Delta G > 0$ are non-spontaneous under given conditions.
71. Solving problems involves converting $\Delta S$ to consistent units (e.g., $kJ/mol·K$).
72. The direction of reaction shifts if $T$ alters $T\Delta S$.
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Applications and Importance

73. Understanding $\Delta H$ helps in energy-efficient reactions.
74. $\Delta S$ predicts molecular organization and disorder.
75. $\Delta G$ determines biochemical reaction feasibility.
76. Combustion engines rely on $\Delta G$ for fuel efficiency.
77. Industrial processes optimize conditions using $\Delta H, \Delta S, \Delta G$.
78. Spontaneity informs storage stability of chemicals.
79. Phase diagrams derive from thermodynamic properties.
80. Thermodynamics explains natural processes like weather patterns.
81. Reaction spontaneity predicts environmental degradation rates.
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Miscellaneous

82. $\Delta G$ links chemistry, biology, and physics through energy concepts.
83. Exothermic spontaneous reactions are favorable for energy harnessing.
84. Endothermic spontaneous reactions (e.g., ice melting) rely on entropy increases.
85. Entropy underpins statistical mechanics in molecular dynamics.
86. Understanding entropy helps design refrigeration systems.
87. Equilibrium constants depend on $\Delta G^\circ$ values.
88. Thermodynamics unites macroscopic and microscopic perspectives.
89. Reaction coupling (e.g., in cells) uses $\Delta G$ relationships for efficiency.
90. Spontaneity principles drive innovation in sustainable energy.
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Summary

91. $\Delta H$ measures heat changes during reactions.
92. Dissolution can be exothermic or endothermic depending on interactions.
93. $\Delta S$ explains disorder in chemical systems.
94. $\Delta G$ governs reaction spontaneity.
95. Temperature affects the balance between $\Delta H$ and $\Delta S$.
96. Thermodynamic principles optimize industrial reactions.
97. Entropy and enthalpy influence natural phenomena.
98. Spontaneity aligns with energy conservation laws.
99. Equilibrium conditions balance enthalpy and entropy contributions.
100. $\Delta G = \Delta H - T\Delta S$ is the cornerstone of reaction thermodynamics.

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