NAME _________________
Chemistry 1312 General Chemistry II
Test 4: Chapter 16 Topics
Acid-Base Relationships
1. pH of solutions of weak acids and bases (review of Chapter 15)
2. Common ion effect: Use the ®ice Table and include the common ion on the
product side. Use the approximation method unless otherwise told.
3. Buffer Solutions: Weak acid or weak base and its salt: an application of the
common ion effect
a. Recognize buffer components: conjugates or the products of an acid base
reaction
b. Use the common ion calculation to find the pH of a Buffer System
1). If the concentrations are given for the components already in
solution = use Molarity in the calculations
2) If the concentrations are given for the components separately then
mixed into a solution = use moles in the calculations and figure out the
new concentrations using the sum of the volume = final volume
c. Use the Henderson-Hasselbalch Equation (the Buffer Equation) when
looking for pH and given pKa
d. Use the Henderson-Hasselbalch Equation (the Buffer Equation) when
looking for the ratio of salt to acid or vice versa to prepare a buffer of a given
pH.
d. Use the ®ice Table to analyze the effect of added acid or added base. Use
the Henderson-Hasselbalch Equation (the Buffer Equation) when looking for
the resulting pH after adding a small amount of acid or base.
4. Titration curves: apply your knowledge of calculating pH for a combination of
acid and base at different stages of the titration
a. Use stoichiometry to figure out which component remains in solution then
calculate the pH accordingly (strong or weak substance)
b. At the equivalence point, check out whether you have a weak acid or weak
base that has been neutralized. The conjugate hydrolyzes and determines the
pH of the solution
5. Buffer Capacity:
Maximum amount of acid or base that can be handled by the buffer system
6. Buffer Range:
+ or – its midpoint pH
Solubility
1. Review the solubility table for ionic compounds
2. When quantifying: use the Ksp relationship
a. Given Ksp find the solubility; molar solubility
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b. Given solubility, calculate Ksp (convert to molar solubility first)
3. Use Ksp to determine precipitation and fractional precipitation
a. Construct the equivalent of Q (the ion product). Note that ppt forms only
when Q>Ksp
b. Calculate the concentration of reactant common to both precipitates to
determine the minimum and maximum amounts required
4. Effects of pH and complex ion formation:
If one of the components reacts with the acid or forms a complex ion,
solubility of the solid (precipitate) is affected
Combine the reactions and use K x K’
Complex ion Formation
1. Write Kf reactions in the direction of forming the complex ion
2. Calculate concentrations as usual with the ®ice Table
More Practice for Chapter 16
Problem 1 (9 pts) Find the pH for the following buffers. Give a balanced reaction equation in all
cases. Use the Henderson-Hasselbalch equation to solve these 3 problems!!
a) A buffer consisting of 0.12 M sodium carbonate and 0.82 M sodium bicarbonate (Ka =
4.7 x 10-11)
b) What is the pH of 500 mL of solution containing 0.125 mol/L of propanoic acid (HPr) and
0.375 mol/L of sodium propanoate (NaPr)? (Ka for propanoic acid is 5.52 × 10-5.)
c) A frequently used buffer in biological systems, including IV fluids, is phosphate-buffered
saline (“PBS”), which consists of 1.44 g of Na2HPO4.2H2O, 0.24 g of KH2PO4, and 8.00 g
of NaCl per liter of solution. What is its pH?
(It is a very good idea to write out the reaction for this equilibrium first! See Table 15.5,
p.683, for the Ka values of the phosphates.)
Problem 2: Titration (6 pts)
You titrate 25.00 mL of 0.0500 M sulfuric acid with 0.0500 M NaOH.
a) How many mL of NaOH does it take to completely titrate the sulfuric acid?
b) Is the equivalence point at pH = 7.0, < 7.0 or > 7.0? Explain!
Repeat the titration with 0.0500 M phosphoric acid and give the answers for question a) and b) for
phosphoric acid
c) How many mL of NaOH does it take to completely titrate the phosphoric acid? (See also
Chapter 4!!!)
d) Is the equivalence point at pH = 7.0, < 7.0 or > 7.0? Explain!
Problem 3 (4 pts) What is the buffer range? What is the buffer capacity?
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Problem 4 (3 pts)
Which of the following solutions has the greatest buffer capacity?
a) 0.40 M CH3COONa/0.20 M CH3COOH
b) 0.40 M CH3COONa/0.60 M CH3COOH
c) 0.30 M CH3COONa/0.60 M CH3COOH
Problem 5 (2 pts each) Indicate whether these statements are True or False
a. Increasing the concentrations of the components of a buffer solution
True / False
will increase the buffer range.
b. Increasing the concentrations of the components of a buffer solution
True / False
will increase the buffer capacity.
c. The equivalence point in a titration is defined as the point when the
True / False
indicator changes color.
Problem 6 (4 pts)
a) Write the solubility constant Ksp for magnesium fluoride, MgF2.
b) Write the solubility constant Ksp for calcium phosphate, Ca3(PO4)2.
Problem 7 (3 pts)
The solubility of silver chloride _______________ when dilute nitric is added to it.
a) increases
b) decreases
c) does not change
d) first increases, then decreases
e) first decreases, then increases
Problem 8: Solubility Problems (2 pts each)
a) Solubility product constant for MX, Ksp = 1.8 x 10-8. What is the concentration of [M+],
and [X-]?
b) The solubility of calcium chromate is 1.56 × 10-3 g/100 mL of solution. What is the Ksp for
CaCrO4?
c) The molar solubility for compound MX equals 5 x 10-5 M. What are the values for [M+]
and [X-]. What is Ksp?
d) Calculate the solubility of barium carbonate, BaCO3, in pure water. Ksp = 2.0 × 10-9
e) The molar solubility for compound MX2 equals 3.8 x 10-12 M. What are the values for
[M+] and [X-]. What is Ksp?
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