15.1 - Solutions of Acids or Bases Containing Containing a Common Ion
The common ion effect is an application of Le Châtelier’s principle
This effect makes a solution of NaF and HF less acidic than a solution of HF alone
The common ion effect is also important in solutions of polyprotic acids.
It is more important than ever to be systematic and to focus on the chemistry occurring in the solution before thinking about mathematical procedures
The way to do this is always to write the major species first and consider the chemical properties of each one
15.2 - Buffered Solutions
A buffered solution is one that resists a change in its pH when either hydroxide ions or protons are added.
The most important practical example of a buffered solution is our blood, which can absorb the acids and bases produced in biological reactions without changing its pH.
Contains a weak acid (HA) and its salt (NaA) or a weak base (B) and its salt (BHCl)
The pH in the buffered solution is determined by the ratio of the concentrations of the weak acid and weak base.
As long as this ratio remains virtually constant, the pH will remain virtually constant
15.3 - Buffering Capacity
A buffer with a large capacity contains large concentrations of the buffering components
The capacity of a buffered solution is determined by the magnitudes
The pKa of the weak acid to be used in the buffer should be as close as possible to the desired pH
15.4 - Titrations and pH Curves
The progress of a titration is represented by plotting the pH of the solution versus the volume of added titrant; the resulting graph is called a pH curve or titration curve
Strong acid: A strong base titrations show a sharp change in pH near the equivalence point
The shape of the pH curve for a strong base–strong acid titration is quite different before the equivalence point from the shape of the pH curve for a strong base–weak acid titration
The strong base–weak acid pH curve shows the effects of buffering before the equivalence point
For a strong base–weak acid titration, the pH is greater than 7 at the equivalence point because of the basic properties
Indicators are sometimes used to mark the equivalence point of an acid-base titration
The endpoint is where the indicator changes color
The goal is to have the endpoint and the equivalence point be as close as possible
Equivalence (stoichiometric) point: The point in the titration where an amount of base has been added to exactly react with all the acid originally present
The equivalence point is defined by the stoichiometry, not by the pH
The amount of acid present, not its strength, determines the equivalence point
15.5 - Acid-Base Indicators
The most common acid-base indicators are complex molecules that are themselves weak acids
The endpoint is defined by the change in color of the indicator. The equivalence point is defined by the reaction stoichiometry
Choosing an indicator is easier if there is a large change in pH near the equivalence point of the titration
The dramatic change in pH near the equivalence point in a strong acid–strong base titration produces a sharp endpoint; that is, the complete color change usually occurs over one drop of added titrant
There is a wide choice of suitable indicators. The results will agree within one drop of titrant, using indicators with endpoints as far apart as pH 5 and pH 9
15.6 - Solubility Equilibria and the Solubility Product
An important consequence of solubility involves the use of a suspension of barium sulfate to improve the clarity of X rays of the gastrointestinal tract
Pure liquids and pure solids are never included in an equilibrium expression
The solubility product is an equilibrium constant and has only one value for a given solid at a given temperature
On the other hand, if a common ion is present in the solution, the solubility varies according to the concentration of the common ion
Ksp is an equilibrium constant; solubility is an equilibrium position
Carbon dioxide dissolved in groundwater makes it acidic, increasing the solubility of calcium carbonate and eventually producing huge caverns.
As the carbon dioxide escapes to the air, the pH of the dripping water goes up and the calcium carbonate precipitates, forming stalactites and stalagmites
15.7 - Precipitation and Qualitative Analysis
We will use the ion product, which is defined just like the expression for Ksp for a given solid except that initial concentrations are used instead of equilibrium concentrations.
Mixtures of metal ions in an aqueous solution are often separated by selective precipitation by using a reagent whose anion forms a precipitate with only one or a few of the metal ions in the mixture
The classic scheme for qualitative analysis of a mixture containing all the common cations involves first separating them into five major groups based on solubilities
Each group is then treated further to separate and identify the individual ions
Group I—Insoluble chlorides
Group II—Sulfides insoluble in acid solution
Group III—Sulfides insoluble in basic solution
Group IV—Insoluble carbonates
Group V—Alkali metal and ammonium ions
15.8 - Equilibria Involving Complex Ions
A complex ion is a charged species consisting of a metal ion surrounded by ligands.
Metal ions and ligands one at a time in steps characterized by equilibrium constants called formation constants for stability constants.
When reactions are added, the equilibrium constant for the overall process is the product of the constants for the individual reactions
If the anion of the solid is a good base, the solubility is greatly increased by acidifying the solution.
In cases where the anion is not sufficiently basic, the ionic solid often can be dissolved in a solution containing a ligand that forms stable complex ions with its cation