As was discussed in Chapter 6, the natural tendency of chemical systems is to seek a state of minimum Gibbs function. Once the minimum is achieved, movement in any chemical direction will not be spontaneous. It is at this point that the system achieves a state of equilibrium.
- 9.1: Prelude to Chemical Equilibria
- The small is great, the great is small; all is in equilibrium in necessity... - Victor Hugo in “Les Miserables”
- 9.2: Chemical Potential
- Equilibrium can be understood as accruing at the composition of a reaction mixture at which the aggregate chemical potential of the products is equal to that of the reactants.
- 9.3: Activities and Fugacities
- To this point, we have mostly ignored deviations from ideal behavior. But it should be noted that thermodynamic equilibrium constants are not expressed in terms of concentrations or pressures, but rather in terms of activities and fugacities .
- 9.4: Pressure Dependence of Kp - Le Châtelier's Principle
- Since the equilibrium constant is a function of change in Gibbs energy, which is defined for a specific composition (all reactants in their standard states and at unit pressure (or fugacity), changes in pressure have no effect on equilibrium constants for a fixed temperature. However, changes in pressure can have profound effects on the compositions of equilibrium mixtures.
- 9.5: Degree of Dissociation
- Reactions such as the one in the previous example involve the dissociation of a molecule. Such reactions can be easily described in terms of the fraction of reactant molecules that actually dissociate to achieve equilibrium in a sample. This fraction is called the degree of dissociation.
- 9.6: Temperature Dependence of Equilibrium Constants - the van ’t Hoff Equation
- The value of Kp is independent of pressure, although the composition of a system at equilibrium may be very much dependent on pressure. Temperature dependence is another matter. Because the value of is dependent on temperature, the value of Kp is as well. The form of the temperature dependence can be taken from the definition of the Gibbs function.
- 9.7: The Dumas Bulb Method for Measuring Decomposition Equilibrium
- A classic example of an experiment that is employed in many physical chemistry laboratory courses uses a Dumas Bulb method to measure the dissociation of N2O4(g) as a function of temperature. In this experiment, a glass bulb is used to create a constant volume container in which a volatile substance can evaporate, or achieve equilibrium with other gases present.
- 9.8: Acid-Base Equilibria
- A great many processes involve proton transfer, or acid-base types of reactions. As many biological systems depend on carefully controlled pH, these types of processes are extremely important.
- 9.9: Buffers
- Buffer solutions, which are of enormous importance in controlling pH in various processes, can be understood in terms of acid/base equilibrium. A buffer is created in a solution which contains both a weak acid and its conjugate base. This creates to absorb excess H+ or supply H+ to replace what is lost due to neutralization. The calculation of the pH of a buffer is straightforward using an ICE table approach.
- 9.10: Solubility of Ionic Compounds
- The solubility of ionic compounds in water can also be described using the concepts of equilibrium. Ksp is the solubility product and is the equilibrium constant that describes the solubility of an electrolyte.
- 9.E: Chemical Equilibria (Exercises)
- Exercises for Chapter 9 "Chemical Equilibria" in Fleming's Physical Chemistry Textmap.
- 9.S: Chemical Equilibria (Summary)
- Summary for Chapter 9 "Chemical Equilibria" in Fleming's Physical Chemistry Textmap.
Contributors and Attributions
Patrick E. Fleming (Department of Chemistry and Biochemistry; California State University, East Bay)