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HSC Chemistry

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  1. Module 1: Properties and Structure of Matter
    1.1 Properties of Matter
  2. 1.2 Atomic Structure and Atomic Mass
  3. 1.3 Periodicity
  4. 1.4 Bonding
  5. Module 2: Introduction to Quantitative Chemistry
    2.1 Chemical Reactions and Stoichiometry
  6. 2.2 Mole Concept
  7. 2.3 Concentration and Molarity
  8. 2.4 Gas Laws
  9. Module 3: Reactive Chemistry
    3.1 Chemical Reactions
  10. 3.2 Predicting Reactions of Metals
  11. 3.3 Rates of Reactions
  12. Module 4: Drivers of Reactions
    4.1 Energy Changes in Chemical Reactions
  13. 4.2 Enthalpy and Hess's Law
  14. 4.3 Entropy and Gibbs Free Energy
  15. Module 5: Equilibrium and Acid Reactions
    5.1 Static and Dynamic Equilibrium
    5 Topics
  16. 5.2 Factors that Affect Equilibrium
    2 Topics
  17. 5.3 Calculating the Equilibrium Constant
    2 Topics
  18. 5.4 Solution Equilibria
  19. Module 6: Acid/Base Reactions
    6.1 Properties of Acids and Bases
    7 Topics
  20. 6.2 Using Brønsted–Lowry Theory
    2 Topics
  21. 6.3 Quantitative Analysis
    1 Topic
  22. Module 7: Organic Chemistry
    7.1 Nomenclature
    2 Topics
  23. 7.2 Hydrocarbons
    2 Topics
  24. 7.3 Products of Reactions Involving Hydrocarbons
  25. 7.4 Alcohols
    1 Topic
  26. 7.5 Reactions of Organic Acids and Bases
  27. 7.6 Polymers
    2 Topics
  28. Module 8: Applying Chemical Ideas
    8.1 Analysis of Inorganic Substances
    3 Topics
  29. 8.2 Analysis of Organic Substances
  30. 8.3 Chemical Synthesis and Design
  31. Working Scientifically
    Working Scientifically Overview
    1 Topic
Lesson 15, Topic 5
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Entropy and Enthalpy in Non-equilibrium Systems

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  • Enthalpy (H) is the heat content of a system.
  • Entropy (S) is a measure of the distribution of available energy, but is also a measure of disorder.
  • These determine whether a reaction is:
    • Spontaneous – Proceeds without needing a continual input of energy from an external source.
    • Non-spontaneous – Requires a continual input of energy from an external source.

The Gibbs free energy formula predicts the spontaneity of a reaction.

\Delta G\degree = \Delta H\degree - T\Delta S\degree
  • If \Delta G\degree > 0, reaction is non-spontaneous.
  • If \Delta G\degree = 0, system is at equilibrium.
  • If \Delta G\degree < 0, reaction is spontaneous.

Combustion

Consider the complete combustion of octane.

C_8H_{18 (l)} + \frac{25}2 O_{2 (g)} \rightarrow 8CO_{2 (g)} + 9H_2O_{(l)}
  • Combustion is exothermic as it always releases heat, i.e. \Delta H\degree < 0
  • Entropy can either be positive or negative, depending on the number of moles of gas on each side of the reaction
  • For octane, entropy decreases as there are fewer moles of gas resulting from the reaction, i.e. \Delta S \degree < 0

However, \Delta G\degree < 0 for all combustion reactions, regardless of temperature of molar ratios, as enthalpy is always extremely negative. Therefore, combustion is a spontaneous reaction.

Photosynthesis

6CO_{2 (g)} + 6H_2O_{(l)} \rightarrow C_6H_{12}O_{6 (aq)} + 6O_{2 (g)}
  • Photosynthesis is endothermic, absorbing energy from the sun, i.e. \Delta H\degree > 0
  • Entropy decreases as there are 6 moles of gas on each side but more overall particles on the reactants size, i.e. \Delta S\degree < 0

For photosynthesis, \Delta G\degree > 0 so it is a non-spontaneous reaction which requires a constant input of energy to proceed.