# Chapter 16 Lecture Outline

## 16.2 Buffers: Solutions That Resist pH Change

Introduction to Buffers

## 16.3 Buffer Effectiveness: Buffer Range and Buffer Capacity

1. Calculate for a 0.1 M acetic acid/0.1 M sodium acetate buffer. Solution

2. Effect of adding a base to a generic weak acid. (mathcad)

3. Add 0.05 moles of sodium hydroxide to 1 liter of acetic acid buffer. (Solution)

4. Add 0.05 moles of sodium hydroxide to 1 liter of water.

5. Formic acid buffer
1. 10.0 mL of 1.0 M formic acid is mixed with 1.00 g of sodium formate and diluted to 250.0 mL.
2. Add 5.0 mL of 1.0 M sodium hydroxide to the formic acid/sodium formate solution.
3. Add 5.0 mL of 2.0 M hydrochloric acid to the formic acid/sodium formate solution.
4. Solutions

6. Graph results of adding acid or base to a buffer with spreadsheet

## 16.4 Titrations and pH Curves

### Strong Acid-Strong Base Titration

1. Strong Acid Titration Problem
1. Start with 50.0 mL of 0.100 M HCl, Calculate pH (on board)
2. Add 1 mL 0.200 M NaOH, Calculate pH (on board)
3. Add 5 mL 0.200 M NaOH, Calculate pH (Group 1)
4. Add 10 mL 0.200 M NaOH, Calculate pH (Group 2)
5. Add 15 mL 0.200 M NaOH, Calculate pH (Group 3)
6. Add 20 mL 0.200 M NaOH, Calculate pH (Group 4)
7. Add 25 mL 0.200 M NaOH, Calculate pH (Class first, then on board)
8. Add 26 mL 0.200 M NaOH, Calculate pH (Class first, then on board)
9. Add 30 mL 0.200 M NaOH, Calculate pH (Group 1)
10. Add 50 mL 0.200 M NaOH, Calculate pH (Group 2)

2. Summerize results

 mL 0.100 M HCl mL 0.200 M NaOH pH 50 0 1.00 50 1 1.03 50 5 1.14 50 10 1.30 50 15 1.51 50 20 1.84 50 25 7.00 50 26 11.42 50 30 12.10 50 50 12.70

3. Show graph in spreadsheet.

4. Solutions

### Titration of a weak acid

1. Calculate for 100.0 mL of 0.1 M acetic acid solution.

1. Add 1 mL of 0.1 M sodium hydroxide.

2. Add 10 mL of 0.1 M sodium hydroxide.

3. Add 50 mL of 0.1 M sodium hydroxide.

4. Add 75 mL of 0.1 M sodium hydroxide.

5. Add 100 mL of 0.1 M sodium hydroxide.

6. Add 110 mL of 0.1 M sodium hydroxide.

2. Graph results with Spreadsheet

3. Solutions

### Graphical Review of titrations

1. Effect of concentration and Ka on Titration curves
1. 0.1 M Strong Acid, 0.1 M Weak acid (Ka = 1.8x10-5, 0.1 M NaOH)
2. 0.1 M Strong Acid, 0.1 M Weak acid (Ka = 1.8x10-7, 0.1 M NaOH)
3. 0.1 M Strong Acid, 0.1 M Weak acid (Ka = 1.8x10-9, 0.1 M NaOH)
4. 0.01 M Strong Acid, 0.01 M Weak acid (Ka = 1.8x10-5, 0.1 M NaOH)

## 16.5 Solubility Equilibria and the Solubility Product Constant

1. Introductiion to Solubility
1. Soluble Salts: NaCl (s) --> Na1+ (aq) + Cl1- (aq) ( internet © Saunders, 1997)

2. Insoluble Salts: Pb2+(aq) + 2 I1-(aq) --> PbI2(s). ( internet © Saunders, 1997)

3. Equlibrium: PbCl2(s) <--> Pb2+(aq) + 2 Cl1-(aq). ( internet © Saunders, 1997)

2. Measuring Pb2+ ions by AA ( internet © Saunders, 1997)

3. Shifting Equlibrium
1. Dilute Ag1+ and Cl1-( internet © Saunders, 1997)

2. Concentrated Ag1+ and Cl1-( internet © Saunders, 1997)

3. Adding Cl1- stepwise to Pb2+ ( internet © Saunders, 1997

4. Qualitative Analysis: internet © Saunders, 1997

5. Conversion of PbCl2 to PbCrO4 ( internet © Saunders, 1997 Animation internet © Saunders, 1997)

### Solubility Equlibrium Problems

1. Molar Solubility. 9.3x10-3 g of CaCO3 dissolves in 1 L of water. What is the "molar solubility" Solutions

2. What is Ksp for CaCO3? Solutions

3. How much AgCl will dissolve in 500 mL of water? Ksp = 1.56 x 10-10 Solutions

4. How much PbCl2 will dissolve in 500 mL of water? Ksp = 1.6 x 10-5. Solutions

5. Given Ksp = 1.5 x 10-11 for Mg(OH)2
1. Calculate the mass of Mg(OH)2 that will dissolve in 250 mL of water.
2. Calculate the mass of Mg(OH)2 that will dissolve in 250 mL of water with the pH buffered to 13.
3. Calculate the mass of Mg(OH)2 that will dissolve in 250 mL of water with the pH buffered to 4.
4. Solutions

6. For AgCl, Ksp = 1.56 x 10-10
1. What is the solubility in water
2. What is the solubility in 0.10 M NaCl
3. What is the solubility in 1.0 M NaCl
4. What happens when 50 mL of 1.0 x10-4 M NaCl is added to 50 mL of 1.0x10-6 M AgNO3.
5. Increase the concentration of NaCl to 0.50 M and repeat.
6. Solutions

7. For PbCl2, Ksp = 1.6 x 10-5
1. What is the solubility in water
2. What is the solubility in 0.10 M NaCl
3. What is the solubility in 1.0 M NaCl
4. What happens when 50 mL of 0.20M NaCl is added to 50 mL of 1.0x10-4 M Pb(NO3)2.
5. Increase the concentration of NaCl to 2.0 M and repeat.
6. Solutions

8. Given a solution containing Ba2+ and Fe2+, Start at pH 1 and adjust the pH so that:
1. The first ion begins to precipitate
2. The second ion begins to precipitate
3. How much of the first has precipitated at this point?
4. How much of the second ion has precipitated at pH 14?
5. Solutions