Ap Chem Chapter 17 Outline
AP Chemistry
Chapter 17 Additional Aspects of Aqueous Equilibria
Chapter 17. Additional Aspects of Equilibrium
Common Student Misconceptions
•
•
•
•
•
Students often believe that the pH at the equivalence point for any titration is 7.00.
In terms of problem-solving skills, this is probably the most difficult chapter for most students.
Students tend to find buffers particularly difficult to understand.
Students often forget to consider volume changes that occur when two solutions are mixed (this will have an effect on the concentration of the species present).
Students tend to confuse Ksp and solubility.
17.1 The Common Ion Effect
•
The dissociation of a weak electrolyte is decreased by the addition of a strong electrolyte that has an ion in common with the weak electrolyte.
•
For example, consider the ionization of a weak acid, acetic acid.
HC2H3O2(aq) ⇋ H+(aq) + C2H3O2–(aq)
•
•
This causes a reduction in the [H+] and a decrease in the percent ionization of the acetic acid.
•
By adding sodium acetate, we have disturbed the acetic acid equilibrium.
•
•
If we add additional C2H3O2– ions by the addition of a strong electrolyte, (e.g., NaC2H3O2) the equilibrium is shifted to the left.
In effect, we have added a product of this equilibrium (i.e., the acetate ion).
• This phenomenon is called the common-ion effect.
Common ion equilibrium problems are solved following the same pattern as other equilibrium problems.
• However, the initial concentration of the common ion (from the salt) must be considered.
Sample Exercise 17.1 (p. 662)
What is the pH of a solution made by adding 0.30 mol of acetic acid (HC2H3O2) and 0.30 mol of sodium acetate
(NaC2H3O2) to enough water to make 1.0 L of solution?
(4.74)
-1-
AP Chemistry
Chapter 17 Additional Aspects of Aqueous Equilibria
Practice Exercise 17.1
Calculate the pH of a solution containing 0.085 M nitrous acid (HNO2, Ka = 4.5 x 10-4) and 0.10 M
Chapter 17 Additional Aspects of Aqueous Equilibria
Chapter 17. Additional Aspects of Equilibrium
Common Student Misconceptions
•
•
•
•
•
Students often believe that the pH at the equivalence point for any titration is 7.00.
In terms of problem-solving skills, this is probably the most difficult chapter for most students.
Students tend to find buffers particularly difficult to understand.
Students often forget to consider volume changes that occur when two solutions are mixed (this will have an effect on the concentration of the species present).
Students tend to confuse Ksp and solubility.
17.1 The Common Ion Effect
•
The dissociation of a weak electrolyte is decreased by the addition of a strong electrolyte that has an ion in common with the weak electrolyte.
•
For example, consider the ionization of a weak acid, acetic acid.
HC2H3O2(aq) ⇋ H+(aq) + C2H3O2–(aq)
•
•
This causes a reduction in the [H+] and a decrease in the percent ionization of the acetic acid.
•
By adding sodium acetate, we have disturbed the acetic acid equilibrium.
•
•
If we add additional C2H3O2– ions by the addition of a strong electrolyte, (e.g., NaC2H3O2) the equilibrium is shifted to the left.
In effect, we have added a product of this equilibrium (i.e., the acetate ion).
• This phenomenon is called the common-ion effect.
Common ion equilibrium problems are solved following the same pattern as other equilibrium problems.
• However, the initial concentration of the common ion (from the salt) must be considered.
Sample Exercise 17.1 (p. 662)
What is the pH of a solution made by adding 0.30 mol of acetic acid (HC2H3O2) and 0.30 mol of sodium acetate
(NaC2H3O2) to enough water to make 1.0 L of solution?
(4.74)
-1-
AP Chemistry
Chapter 17 Additional Aspects of Aqueous Equilibria
Practice Exercise 17.1
Calculate the pH of a solution containing 0.085 M nitrous acid (HNO2, Ka = 4.5 x 10-4) and 0.10 M