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20: Entropy and Free Energy

Last updated

Mar 21, 2025
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- 19.15: Common Ion Effect
- 20.1: Entropy

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Topic hierarchy

20.1: Entropy
This page explains entropy as a measure of disorder in systems, illustrating that natural processes tend to increase disorder over time. It discusses how chemical reactions often lead to increased entropy, especially when transitioning to gas or increasing temperature. The content provides examples of reactions that demonstrate changes in entropy, highlighting methods to predict whether entropy increases or decreases based on molecular states and quantities.
20.2: Standard Entropy
This page discusses geothermal energy sources that utilize steam from geysers to fulfill energy needs. It also explains the concept of entropy, which quantifies molecular motion and has a standard value at absolute zero. As temperature increases, entropy rises, and the page details how to calculate entropy changes for reactions like water vaporization and liquid formation. It emphasizes that reactions can still be favorable despite a decrease in entropy when they are exothermic.
20.3: Spontaneous and Nonspontaneous Reactions
This page discusses nitroglycerin as an unstable explosive, spontaneous reactions that favor product formation with decreased enthalpy and increased entropy, and contrasts them with nonspontaneous, endothermic reactions. It notes that spontaneity does not guarantee speed, as seen in slow combustion, and highlights reversible reactions, such as carbonic acid decomposition, to further illustrate these principles.
20.4: Free Energy
This page discusses the steam engine's significance in modern railroads and introduces free energy ( $G$ ) in chemical reactions. It explains how free energy is influenced by enthalpy and entropy, detailing the Gibbs free energy equation that evaluates reaction spontaneity.
20.5: Calculating Free Energy Change $\left( \Delta G^\text{o} \right)$
This page explains the process of baking, emphasizing the importance of heating ingredients to specific temperatures for chemical reactions. It discusses the Gibbs free energy change ( $\Delta G^\text{o}$ ), illustrated with an example of methane and water where the calculated $\Delta G^\text{o}$ is +142.0 kJ/mol at 25°C, indicating a non-spontaneous reaction at that temperature. It also warns of the need to carefully apply these thermodynamic values across different temperatures.
20.6: Temperature and Free Energy
This page discusses the reactions of iron ore and coke producing iron and carbon dioxide at high temperatures. It also explains the decomposition of calcium carbonate into calcium oxide and carbon dioxide, which requires temperatures exceeding approximately 835°C. Below this temperature, the products are not detectable, but measurable carbon dioxide is produced above 700°C. The quicklime production process leverages the removal of carbon dioxide to favor product formation.
20.7: Changes of State and Free Energy
This page discusses energy changes in water during phase transitions, particularly melting and vaporization. It explains that heating water leads to steam formation and increased molecular disorder. At the melting point (0°C), Gibbs free energy ( $\Delta G$ ) reaches zero, enabling calculations of entropy change ( $\Delta S$ ).
20.8: Calculations of Free Energy and $K_\text{eq}$
This page explains the formation of stalactites and stalagmites in caves through mineral solutions reacting with carbon dioxide to produce calcium carbonate deposits. It covers equilibrium concepts, stating that at equilibrium, the free energy change ( $\Delta G$ ) is zero, and the equilibrium constant ( $K_{eq}$ ) indicates product or reactant favorability. The page includes examples of calculating $K_{eq}$ from $\Delta G^\circ$ and vice versa using thermodynamic equations.

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