1. 17.3 Acid-BaseTitrations
A known concentration of base (or acid) is slowly added to a solution of acid (or base).
For example, a standard solution of NaOH (the titrant) is added to an acid of to determine its concentration

2. A pH meter or acid-base indicators are used to determine when the solution has reached the equivalence point, at which the stoichiometric amount of acid equals that of base.

3. Strong Acid—Strong Base Titrations
From the start of the titration to near the equivalence point, the pH goes up slowly.

4. Just before and after the equivalence point, the pH increases rapidly.

5. At the equivalence point, moles acid = moles base, and the solution contains only water and the salt from the cation of the base and the anion of the acid.

6. As more base is added, the increase in pH again levels off.

7. The shape of a pH curve for the titration of a strong base with a strong acid is seen in Fig The pH starts out high and decreases as the acid is added, rapidly dropping to the equivalence point.
See sample exercise 17.6

8. Weak Acid—Strong Base Titrations
Unlike in the previous case, the conjugate base of the acid affects the pH when it is formed.
The pH at the equivalence point will be >7.
Phenolphthalein is commonly used as an indicator in these titrations.

9. At each point below the equivalence point, the pH of the solution during titration is determined from the amounts of the acid and its conjugate base present at that particular time.

10. See sample exercise 17.7

11. With weaker acids, the initial pH is higher and pH changes near the equivalence point are more subtle.

12. Weak Base—Strong Acid Titrations
The pH at the equivalence point in these titrations is < 7.
Methyl red is the indicator of choice.

13. Titrations of Polyprotic Acids
In these cases there is an equivalence point for each dissociation.

14. 17.4 Solubility Equilibria
Equilibria considered so far include acids and bases.
These equilibria are homogeneous
All species are in the same phase
Other equilibria involve the dissolution or precipitation of ionic compounds.
These reactions are heterogeneous since all components are not in the same phase.
By considering solubility equilibria (rather than just rules for solubility), quantitative predictions can be made about the amount of a given compound that will dissolve.

15. The Solubility-Product Constant, Ksp
Consider the equilibrium that exists in a saturated solution of BaSO4 in water:
The equilibrium constant expression for this equilibrium is
Ksp = [Ba2+] [SO42−]
where the equilibrium constant, Ksp, is called the solubility-product constant (or, simply, solubility product).
BaSO4(s)
Ba2+(aq) + SO42−(aq)

16. Even though [BaSO4] is excluded from the equilibrium-constant expression, some undissolved BaSO4(s) must be present in order for the system to be at equilibrium.
The Ksp is the equilibrium constant for the equilibrium that exists between a solid ionic solute and its ions in a saturated aqueous solution.
See Appendix D for Ksp values for many ionic solids.
The Ksp for BaSO4 is 1.1 x 10-10, indicating that only a very small amount of the solid will dissolve in water.

17. Solubility and Ksp Ksp is not the same as solubility.
The solubility of a substance is the quantity that dissolves to form a saturated solution.
Solubility is generally expressed as the mass of solute dissolved in 1 L (g/L) or 100 mL (g/mL) of solution, or in mol/L (M).
Ksp (a unitless number) is the equilibrium constant for the equilibrium between an ionic solid and its saturated solution.
The magnitude of Ksp is a measure of how much of the solid dissolves to form a saturated solution.

18. Relationship between solubility and Ksp:
Calculating Ksp from Solubility: See sample exercise
Calculating Solubility from Ksp: See sample exercise