# 17: Additional Aspects of Aqueous Equilibria

This chapter address equilibria with solutions containing more than one solute.

• 17.1: The Common-Ion Effect
The common-ion effect is used to describe the effect on an equilibrium involving a substance that adds an ion that is a part of the equilibrium.
• 17.2: Buffered Solutions
Buffers are solutions that resist a change in pH after adding an acid or a base. Buffers contain a weak acid ($$HA$$) and its conjugate weak base ($$A^−$$). Adding a strong electrolyte that contains one ion in common with a reaction system that is at equilibrium shifts the equilibrium in such a way as to reduce the concentration of the common ion. Buffers are characterized by their pH range and buffer capacity.
• 17.3: Acid-Base Titrations
The shape of a titration curve, a plot of pH versus the amount of acid or base added, provides important information about what is occurring in solution during a titration. The shapes of titration curves for weak acids and bases depend dramatically on the identity of the compound. The equivalence point of an acid–base titration is the point at which exactly enough acid or base has been added to react completely with the other component.
• 17.4: Solubility Equilibria
The solubility product (Ksp) is used to calculate equilibrium concentrations of the ions in solution, whereas the ion product (Q) describes concentrations that are not necessarily at equilibrium. The equilibrium constant for a dissolution reaction, called the solubility product (Ksp), is a measure of the solubility of a compound. Whereas solubility is usually expressed in terms of mass of solute per 100 mL of solvent, Ksp is defined in terms of the molar concentrations of the component ions.
• 17.5: Factors that Affect Solubility
Ion-pair formation, the incomplete dissociation of molecular solutes, the formation of complex ions, and changes in pH all affect solubility. There are four explanations why the solubility of a compound can differ from the solubility indicated by the concentrations of ions: (1) ion pair formation, in which an anion and a cation are in intimate contact in solution and not separated by solvent, (2) the incomplete dissociation of molecular solutes, (3) the formation of complex ions, and (4) changes
• 17.6: Precipitation and Separation of Ions
A mixture of metal ions in a solution can be separated by precipitation with anions such as $$\ce{Cl-}$$, $$\ce{Br-}$$, $$\ce{SO4^2-}$$, $$\ce{CO3^2-}$$, $$\ce{S^2-}$$, $$\ce{Cr2O4^2-}$$, $$\ce{PO4^2-}$$, $$\ce{OH-}$$ etc. When a metal ion or a group of metal ions form insoluble salts with a particular anion, they can be separated from others by precipitation. We can also separate the anions by precipitating them with appropriate metal ions.
• 17.7: Qualitative Analysis for Metallic Elements
In qualitative analysis, the identity, not the amount, of metal ions present in a mixture is determined. The technique consists of selectively precipitating only a few kinds of metal ions at a time under given sets of conditions. Consecutive precipitation steps become progressively less selective until almost all the metal ions are precipitated. Other additional steps are needed to separate metal ions that precipitate together.
• 17.8: Additional Aspects of Aqueous Equilibria (Exercises)
These are homework exercises for Chapter 17 "Additional Aspects of Aqueous Equilibria" of the Brown et al. Textmap, which addresses more complex equilibria with solutions containing more than one solute.
• 17.9: Additional Aspects of Aqueous Equilibria (Summary)
This is the summary for Chapter 17 "Additional Aspects of Aqueous Equilibria" of Brown et al. Textmap, which addresses more complex equilibria with solutions containing more than one solute.