Chapter 7: Thermodynamics
Introduction
Thermodynamics is the study of energy transformations and the principles governing these changes, providing a framework to understand why certain processes occur and others do not. At its core are three key concepts: enthalpy (H), entropy (S), and Gibbs free energy (G), which together determine whether a reaction is energetically favorable . Recall that enthalpy describes the heat exchange in a system at constant pressure, indicating whether a reaction absorbs or releases thermal energy. In this Chapter, other thermodynamic entities such as Entropy and Gibbs free energy will be explored to explain spontaneous process as governed by the Second Law of Thermodynamics.
Thermodynamics explains everyday phenomena, such as why ice melts at room temperature. At temperatures above 0°C, increasingly disordered water molecules as ice transitions to liquid outweighs the heat required to break hydrogen bonds, resulting in a negative ΔG or exergonic process.
a. b.
Figure 7.1: a. When water is frozen, the molecules are, "locked" into place due to the strong hydrogen bonding that occurs between molecules (left, "1" = Hydrogen Bond). (CC BY-SA 4.0, 2020: Haritz Perez via WikiMedia Commons ) As we know, heat is required to melt ice, meaning it is an endo thermic change. b. However, the requisite heat is compensated for by the increasing randomness in liquid water, which disperses energy (right). Since all systems, unconstrained, must tend towards disorder, ice spontaneously melts despite melting being an endothermic phase change. (CC BY-SA 3.0, 2007: Thomas Splettstoesser via WikiMedia Commons )
This principle also distinguishes between exergonic reactions, which release energy (ΔG<0), and endergonic reactions, which require energy input (ΔG>0). These ideas are central to understanding biological systems, such as ATP hydrolysis, where the release of free energy drives essential cellular processes.
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- 7.1: Spontaneous Change
- Chemical and physical processes have a natural tendency to occur in one direction under certain conditions. A spontaneous process occurs without the need for a continual input of energy from some external source, while a nonspontaneous process requires such. Systems undergoing a spontaneous process may or may not experience a gain or loss of energy, but they will experience a change in the way matter and/or energy is distributed within the system.
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- 7.2: Entropy and The Second Law of Thermodynamics
- Entropy (S) is a state function whose value increases with an increase in the number of available microstates.For a given system, the greater the number of microstates, the higher the entropy. During a spontaneous process, the entropy of the universe increases.
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- 7.3: Gibbs Free Energy - Relating H and S
- We can predict whether a reaction will occur spontaneously by combining the entropy, enthalpy, and temperature of a system in a new state function called Gibbs free energy (G). The change in free energy (ΔG) is the difference between the heat released during a process and the heat released for the same process occurring in a reversible manner. If a system is at equilibrium, ΔG = 0. If the process is spontaneous, ΔG < 0. If the process is not spontaneous as written.
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- 7.4: Gibbs Free Energy for Non-Standard
- For a reversible process (with no external work), the change in free energy can be expressed in terms of volume, pressure, entropy, and temperature. If ΔG° < 0, then K > 1, and products are favored over reactants. If ΔG° > 0, then K < 1, and reactants are favored over products. If ΔG° = 0, then K = 1, and the system is at equilibrium. We can use the measured equilibrium constant K at one temperature and ΔH° to estimate the equilibrium constant for a reaction at any other temperature.
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- 7.5: Gibbs Energy and Equilibrium
- For a reversible process (with no external work), the change in free energy can be expressed in terms of volume, pressure, entropy, and temperature. If ΔG° < 0, then K > 1, and products are favored over reactants. If ΔG° > 0, then K < 1, and reactants are favored over products. If ΔG° = 0, then K = 1, and the system is at equilibrium. We can use the measured equilibrium constant K at one temperature and ΔH° to estimate the equilibrium constant for a reaction at any other temperature.