Enter your Sign on user name and password.

Forgot password?
Sign In | Subscribe
Start learning today, and be successful in your academic & professional career. Start Today!

Use Chrome browser to play professor video
Raffi Hovasapian

Raffi Hovasapian

Solubility Equilibria, Part II

Slide Duration:

Table of Contents

I. Review
Naming Compounds

41m 24s

Intro
0:00
Periodic Table of Elements
0:15
Naming Compounds
3:13
Definition and Examples of Ions
3:14
Ionic (Symbol to Name): NaCl
5:23
Ionic (Name to Symbol): Calcium Oxide
7:58
Ionic - Polyatoms Anions: Examples
12:45
Ionic - Polyatoms Anions (Symbol to Name): KClO
14:50
Ionic - Polyatoms Anions (Name to Symbol): Potassium Phosphate
15:49
Ionic Compounds Involving Transition Metals (Symbol to Name): Co₂(CO₃)₃
20:48
Ionic Compounds Involving Transition Metals (Name to Symbol): Palladium 2 Acetate
22:44
Naming Covalent Compounds (Symbol to Name): CO
26:21
Naming Covalent Compounds (Name to Symbol): Nitrogen Trifluoride
27:34
Naming Covalent Compounds (Name to Symbol): Dichlorine Monoxide
27:57
Naming Acids Introduction
28:11
Naming Acids (Name to Symbol): Chlorous Acid
35:08
% Composition by Mass Example
37:38
Stoichiometry

37m 19s

Intro
0:00
Stoichiometry
0:25
Introduction to Stoichiometry
0:26
Example 1
5:03
Example 2
10:17
Example 3
15:09
Example 4
24:02
Example 5: Questions
28:11
Example 5: Part A - Limiting Reactant
30:30
Example 5: Part B
32:27
Example 5: Part C
35:00
II. Aqueous Reactions & Stoichiometry
Precipitation Reactions

31m 14s

Intro
0:00
Precipitation Reactions
0:53
Dissociation of ionic Compounds
0:54
Solubility Guidelines for ionic Compounds: Soluble Ionic Compounds
8:15
Solubility Guidelines for ionic Compounds: Insoluble ionic Compounds
12:56
Precipitation Reactions
14:08
Example 1: Mixing a Solution of BaCl₂ & K₂SO₄
21:21
Example 2: Mixing a Solution of Mg(NO₃)₂ & KI
26:10
Acid-Base Reactions

43m 21s

Intro
0:00
Acid-Base Reactions
1:00
Introduction to Acid: Monoprotic Acid and Polyprotic Acid
1:01
Introduction to Base
8:28
Neutralization
11:45
Example 1
16:17
Example 2
21:55
Molarity
24:50
Example 3
26:50
Example 4
30:01
Example 4: Limiting Reactant
37:51
Example 4: Reaction Part
40:01
Oxidation Reduction Reactions

47m 58s

Intro
0:00
Oxidation Reduction Reactions
0:26
Oxidation and Reduction Overview
0:27
How Can One Tell Whether Oxidation-Reduction has Taken Place?
7:13
Rules for Assigning Oxidation State: Number 1
11:22
Rules for Assigning Oxidation State: Number 2
12:46
Rules for Assigning Oxidation State: Number 3
13:25
Rules for Assigning Oxidation State: Number 4
14:50
Rules for Assigning Oxidation State: Number 5
15:41
Rules for Assigning Oxidation State: Number 6
17:00
Example 1: Determine the Oxidation State of Sulfur in the Following Compounds
18:20
Activity Series and Reduction Properties
25:32
Activity Series and Reduction Properties
25:33
Example 2: Write the Balance Molecular, Total Ionic, and Net Ionic Equations for Al + HCl
31:37
Example 3
34:25
Example 4
37:55
Stoichiometry Examples

31m 50s

Intro
0:00
Stoichiometry Example 1
0:36
Example 1: Question and Answer
0:37
Stoichiometry Example 2
6:57
Example 2: Questions
6:58
Example 2: Part A Solution
12:16
Example 2: Part B Solution
13:05
Example 2: Part C Solution
14:00
Example 2: Part D Solution
14:38
Stoichiometry Example 3
17:56
Example 3: Questions
17:57
Example 3: Part A Solution
19:51
Example 3: Part B Solution
21:43
Example 3: Part C Solution
26:46
III. Gases
Pressure, Gas Laws, & The Ideal Gas Equation

49m 40s

Intro
0:00
Pressure
0:22
Pressure Overview
0:23
Torricelli: Barometer
4:35
Measuring Gas Pressure in a Container
7:49
Boyle's Law
12:40
Example 1
16:56
Gas Laws
21:18
Gas Laws
21:19
Avogadro's Law
26:16
Example 2
31:47
Ideal Gas Equation
38:20
Standard Temperature and Pressure (STP)
38:21
Example 3
40:43
Partial Pressure, Mol Fraction, & Vapor Pressure

32m

Intro
0:00
Gases
0:27
Gases
0:28
Mole Fractions
5:52
Vapor Pressure
8:22
Example 1
13:25
Example 2
22:45
Kinetic Molecular Theory and Real Gases

31m 58s

Intro
0:00
Kinetic Molecular Theory and Real Gases
0:45
Kinetic Molecular Theory 1
0:46
Kinetic Molecular Theory 2
4:23
Kinetic Molecular Theory 3
5:42
Kinetic Molecular Theory 4
6:27
Equations
7:52
Effusion
11:15
Diffusion
13:30
Example 1
19:54
Example 2
23:23
Example 3
26:45
AP Practice for Gases

25m 34s

Intro
0:00
Example 1
0:34
Example 1
0:35
Example 2
6:15
Example 2: Part A
6:16
Example 2: Part B
8:46
Example 2: Part C
10:30
Example 2: Part D
11:15
Example 2: Part E
12:20
Example 2: Part F
13:22
Example 3
14:45
Example 3
14:46
Example 4
18:16
Example 4
18:17
Example 5
21:04
Example 5
21:05
IV. Thermochemistry
Energy, Heat, and Work

37m 32s

Intro
0:00
Thermochemistry
0:25
Temperature and Heat
0:26
Work
3:07
System, Surroundings, Exothermic Process, and Endothermic Process
8:19
Work & Gas: Expansion and Compression
16:30
Example 1
24:41
Example 2
27:47
Example 3
31:58
Enthalpy & Hess's Law

32m 34s

Intro
0:00
Thermochemistry
1:43
Defining Enthalpy & Hess's Law
1:44
Example 1
6:48
State Function
13:11
Example 2
17:15
Example 3
24:09
Standard Enthalpies of Formation

23m 9s

Intro
0:00
Thermochemistry
1:04
Standard Enthalpy of Formation: Definition & Equation
1:05
∆H of Formation
10:00
Example 1
11:22
Example 2
19:00
Calorimetry

39m 28s

Intro
0:00
Thermochemistry
0:21
Heat Capacity
0:22
Molar Heat Capacity
4:44
Constant Pressure Calorimetry
5:50
Example 1
12:24
Constant Volume Calorimetry
21:54
Example 2
24:40
Example 3
31:03
V. Kinetics
Reaction Rates and Rate Laws

36m 24s

Intro
0:00
Kinetics
2:18
Rate: 2 NO₂ (g) → 2NO (g) + O₂ (g)
2:19
Reaction Rates Graph
7:25
Time Interval & Average Rate
13:13
Instantaneous Rate
15:13
Rate of Reaction is Proportional to Some Power of the Reactant Concentrations
23:49
Example 1
27:19
Method of Initial Rates

30m 48s

Intro
0:00
Kinetics
0:33
Rate
0:34
Idea
2:24
Example 1: NH₄⁺ + NO₂⁻ → NO₂ (g) + 2 H₂O
5:36
Example 2: BrO₃⁻ + 5 Br⁻ + 6 H⁺ → 3 Br₂ + 3 H₂O
19:29
Integrated Rate Law & Reaction Half-Life

32m 17s

Intro
0:00
Kinetics
0:52
Integrated Rate Law
0:53
Example 1
6:26
Example 2
15:19
Half-life of a Reaction
20:40
Example 3: Part A
25:41
Example 3: Part B
28:01
Second Order & Zero-Order Rate Laws

26m 40s

Intro
0:00
Kinetics
0:22
Second Order
0:23
Example 1
6:08
Zero-Order
16:36
Summary for the Kinetics Associated with the Reaction
21:27
Activation Energy & Arrhenius Equation

40m 59s

Intro
0:00
Kinetics
0:53
Rate Constant
0:54
Collision Model
2:45
Activation Energy
5:11
Arrhenius Proposed
9:54
2 Requirements for a Successful Reaction
15:39
Rate Constant
17:53
Arrhenius Equation
19:51
Example 1
25:00
Activation Energy & the Values of K
32:12
Example 2
36:46
AP Practice for Kinetics

29m 8s

Intro
0:00
Kinetics
0:43
Example 1
0:44
Example 2
6:53
Example 3
8:58
Example 4
11:36
Example 5
16:36
Example 6: Part A
21:00
Example 6: Part B
25:09
VI. Equilibrium
Equilibrium, Part 1

46m

Intro
0:00
Equilibrium
1:32
Introduction to Equilibrium
1:33
Equilibrium Rules
14:00
Example 1: Part A
16:46
Example 1: Part B
18:48
Example 1: Part C
22:13
Example 1: Part D
24:55
Example 2: Part A
27:46
Example 2: Part B
31:22
Example 2: Part C
33:00
Reverse a Reaction
36:04
Example 3
37:24
Equilibrium, Part 2

40m 53s

Intro
0:00
Equilibrium
1:31
Equilibriums Involving Gases
1:32
General Equation
10:11
Example 1: Question
11:55
Example 1: Answer
13:43
Example 2: Question
19:08
Example 2: Answer
21:37
Example 3: Question
33:40
Example 3: Answer
35:24
Equilibrium: Reaction Quotient

45m 53s

Intro
0:00
Equilibrium
0:57
Reaction Quotient
0:58
If Q > K
5:37
If Q < K
6:52
If Q = K
7:45
Example 1: Part A
8:24
Example 1: Part B
13:11
Example 2: Question
20:04
Example 2: Answer
22:15
Example 3: Question
30:54
Example 3: Answer
32:52
Steps in Solving Equilibrium Problems
42:40
Equilibrium: Examples

31m 51s

Intro
0:00
Equilibrium
1:09
Example 1: Question
1:10
Example 1: Answer
4:15
Example 2: Question
13:04
Example 2: Answer
15:20
Example 3: Question
25:03
Example 3: Answer
26:32
Le Chatelier's principle & Equilibrium

40m 52s

Intro
0:00
Le Chatelier
1:05
Le Chatelier Principle
1:06
Concentration: Add 'x'
5:25
Concentration: Subtract 'x'
7:50
Example 1
9:44
Change in Pressure
12:53
Example 2
20:40
Temperature: Exothermic and Endothermic
24:33
Example 3
29:55
Example 4
35:30
VII. Acids & Bases
Acids and Bases

50m 11s

Intro
0:00
Acids and Bases
1:14
Bronsted-Lowry Acid-Base Model
1:28
Reaction of an Acid with Water
4:36
Acid Dissociation
10:51
Acid Strength
13:48
Example 1
21:22
Water as an Acid & a Base
25:25
Example 2: Part A
32:30
Example 2: Part B
34:47
Example 3: Part A
35:58
Example 3: Part B
39:33
pH Scale
41:12
Example 4
43:56
pH of Weak Acid Solutions

43m 52s

Intro
0:00
pH of Weak Acid Solutions
1:12
pH of Weak Acid Solutions
1:13
Example 1
6:26
Example 2
14:25
Example 3
24:23
Example 4
30:38
Percent Dissociation: Strong & Weak Bases

43m 4s

Intro
0:00
Bases
0:33
Percent Dissociation: Strong & Weak Bases
0:45
Example 1
6:23
Strong Base Dissociation
11:24
Example 2
13:02
Weak Acid and General Reaction
17:38
Example: NaOH → Na⁺ + OH⁻
20:30
Strong Base and Weak Base
23:49
Example 4
24:54
Example 5
33:51
Polyprotic Acids

35m 34s

Intro
0:00
Polyprotic Acids
1:04
Acids Dissociation
1:05
Example 1
4:51
Example 2
17:30
Example 3
31:11
Salts and Their Acid-Base Properties

41m 14s

Intro
0:00
Salts and Their Acid-Base Properties
0:11
Salts and Their Acid-Base Properties
0:15
Example 1
7:58
Example 2
14:00
Metal Ion and Acidic Solution
22:00
Example 3
28:35
NH₄F → NH₄⁺ + F⁻
34:05
Example 4
38:03
Common Ion Effect & Buffers

41m 58s

Intro
0:00
Common Ion Effect & Buffers
1:16
Covalent Oxides Produce Acidic Solutions in Water
1:36
Ionic Oxides Produce Basic Solutions in Water
4:15
Practice Example 1
6:10
Practice Example 2
9:00
Definition
12:27
Example 1: Part A
16:49
Example 1: Part B
19:54
Buffer Solution
25:10
Example of Some Buffers: HF and NaF
30:02
Example of Some Buffers: Acetic Acid & Potassium Acetate
31:34
Example of Some Buffers: CH₃NH₂ & CH₃NH₃Cl
33:54
Example 2: Buffer Solution
36:36
Buffer

32m 24s

Intro
0:00
Buffers
1:20
Buffer Solution
1:21
Adding Base
5:03
Adding Acid
7:14
Example 1: Question
9:48
Example 1: Recall
12:08
Example 1: Major Species Upon Addition of NaOH
16:10
Example 1: Equilibrium, ICE Chart, and Final Calculation
24:33
Example 1: Comparison
29:19
Buffers, Part II

40m 6s

Intro
0:00
Buffers
1:27
Example 1: Question
1:32
Example 1: ICE Chart
3:15
Example 1: Major Species Upon Addition of OH⁻, But Before Rxn
7:23
Example 1: Equilibrium, ICE Chart, and Final Calculation
12:51
Summary
17:21
Another Look at Buffering & the Henderson-Hasselbalch equation
19:00
Example 2
27:08
Example 3
32:01
Buffers, Part III

38m 43s

Intro
0:00
Buffers
0:25
Buffer Capacity Part 1
0:26
Example 1
4:10
Buffer Capacity Part 2
19:29
Example 2
25:12
Example 3
32:02
Titrations: Strong Acid and Strong Base

42m 42s

Intro
0:00
Titrations: Strong Acid and Strong Base
1:11
Definition of Titration
1:12
Sample Problem
3:33
Definition of Titration Curve or pH Curve
9:46
Scenario 1: Strong Acid- Strong Base Titration
11:00
Question
11:01
Part 1: No NaOH is Added
14:00
Part 2: 10.0 mL of NaOH is Added
15:50
Part 3: Another 10.0 mL of NaOH & 20.0 mL of NaOH are Added
22:19
Part 4: 50.0 mL of NaOH is Added
26:46
Part 5: 100.0 mL (Total) of NaOH is Added
27:26
Part 6: 150.0 mL (Total) of NaOH is Added
32:06
Part 7: 200.0 mL of NaOH is Added
35:07
Titrations Curve for Strong Acid and Strong Base
35:43
Titrations: Weak Acid and Strong Base

42m 3s

Intro
0:00
Titrations: Weak Acid and Strong Base
0:43
Question
0:44
Part 1: No NaOH is Added
1:54
Part 2: 10.0 mL of NaOH is Added
5:17
Part 3: 25.0 mL of NaOH is Added
14:01
Part 4: 40.0 mL of NaOH is Added
21:55
Part 5: 50.0 mL (Total) of NaOH is Added
22:25
Part 6: 60.0 mL (Total) of NaOH is Added
31:36
Part 7: 75.0 mL (Total) of NaOH is Added
35:44
Titration Curve
36:09
Titration Examples & Acid-Base Indicators

52m 3s

Intro
0:00
Examples and Indicators
0:25
Example 1: Question
0:26
Example 1: Solution
2:03
Example 2: Question
12:33
Example 2: Solution
14:52
Example 3: Question
23:45
Example 3: Solution
25:09
Acid/Base Indicator Overview
34:45
Acid/Base Indicator Example
37:40
Acid/Base Indicator General Result
47:11
Choosing Acid/Base Indicator
49:12
VIII. Solubility
Solubility Equilibria

36m 25s

Intro
0:00
Solubility Equilibria
0:48
Solubility Equilibria Overview
0:49
Solubility Product Constant
4:24
Definition of Solubility
9:10
Definition of Solubility Product
11:28
Example 1
14:09
Example 2
20:19
Example 3
27:30
Relative Solubilities
31:04
Solubility Equilibria, Part II

42m 6s

Intro
0:00
Solubility Equilibria
0:46
Common Ion Effect
0:47
Example 1
3:14
pH & Solubility
13:00
Example of pH & Solubility
15:25
Example 2
23:06
Precipitation & Definition of the Ion Product
26:48
If Q > Ksp
29:31
If Q < Ksp
30:27
Example 3
32:58
Solubility Equilibria, Part III

43m 9s

Intro
0:00
Solubility Equilibria
0:55
Example 1: Question
0:56
Example 1: Step 1 - Check to See if Anything Precipitates
2:52
Example 1: Step 2 - Stoichiometry
10:47
Example 1: Step 3 - Equilibrium
16:34
Example 2: Selective Precipitation (Question)
21:02
Example 2: Solution
23:41
Classical Qualitative Analysis
29:44
Groups: 1-5
38:44
IX. Complex Ions
Complex Ion Equilibria

43m 38s

Intro
0:00
Complex Ion Equilibria
0:32
Complex Ion
0:34
Ligan Examples
1:51
Ligand Definition
3:12
Coordination
6:28
Example 1
8:08
Example 2
19:13
Complex Ions & Solubility

31m 30s

Intro
0:00
Complex Ions and Solubility
0:23
Recall: Classical Qualitative Analysis
0:24
Example 1
6:10
Example 2
16:16
Dissolving a Water-Insoluble Ionic Compound: Method 1
23:38
Dissolving a Water-Insoluble Ionic Compound: Method 2
28:13
X. Chemical Thermodynamics
Spontaneity, Entropy, & Free Energy, Part I

56m 28s

Intro
0:00
Spontaneity, Entropy, Free Energy
2:25
Energy Overview
2:26
Equation: ∆E = q + w
4:30
State Function/ State Property
8:35
Equation: w = -P∆V
12:00
Enthalpy: H = E + PV
14:50
Enthalpy is a State Property
17:33
Exothermic and Endothermic Reactions
19:20
First Law of Thermodynamic
22:28
Entropy
25:48
Spontaneous Process
33:53
Second Law of Thermodynamic
36:51
More on Entropy
42:23
Example
43:55
Spontaneity, Entropy, & Free Energy, Part II

39m 55s

Intro
0:00
Spontaneity, Entropy, Free Energy
1:30
∆S of Universe = ∆S of System + ∆S of Surrounding
1:31
Convention
3:32
Examining a System
5:36
Thermodynamic Property: Sign of ∆S
16:52
Thermodynamic Property: Magnitude of ∆S
18:45
Deriving Equation: ∆S of Surrounding = -∆H / T
20:25
Example 1
25:51
Free Energy Equations
29:22
Spontaneity, Entropy, & Free Energy, Part III

30m 10s

Intro
0:00
Spontaneity, Entropy, Free Energy
0:11
Example 1
2:38
Key Concept of Example 1
14:06
Example 2
15:56
Units for ∆H, ∆G, and S
20:56
∆S of Surrounding & ∆S of System
22:00
Reaction Example
24:17
Example 3
26:52
Spontaneity, Entropy, & Free Energy, Part IV

30m 7s

Intro
0:00
Spontaneity, Entropy, Free Energy
0:29
Standard Free Energy of Formation
0:58
Example 1
4:34
Reaction Under Non-standard Conditions
13:23
Example 2
16:26
∆G = Negative
22:12
∆G = 0
24:38
Diagram Example of ∆G
26:43
Spontaneity, Entropy, & Free Energy, Part V

44m 56s

Intro
0:00
Spontaneity, Entropy, Free Energy
0:56
Equations: ∆G of Reaction, ∆G°, and K
0:57
Example 1: Question
6:50
Example 1: Part A
9:49
Example 1: Part B
15:28
Example 2
17:33
Example 3
23:31
lnK = (- ∆H° ÷ R) ( 1 ÷ T) + ( ∆S° ÷ R)
31:36
Maximum Work
35:57
XI. Electrochemistry
Oxidation-Reduction & Balancing

39m 23s

Intro
0:00
Oxidation-Reduction and Balancing
2:06
Definition of Electrochemistry
2:07
Oxidation and Reduction Review
3:05
Example 1: Assigning Oxidation State
10:15
Example 2: Is the Following a Redox Reaction?
18:06
Example 3: Step 1 - Write the Oxidation & Reduction Half Reactions
22:46
Example 3: Step 2 - Balance the Reaction
26:44
Example 3: Step 3 - Multiply
30:11
Example 3: Step 4 - Add
32:07
Example 3: Step 5 - Check
33:29
Galvanic Cells

43m 9s

Intro
0:00
Galvanic Cells
0:39
Example 1: Balance the Following Under Basic Conditions
0:40
Example 1: Steps to Balance Reaction Under Basic Conditions
3:25
Example 1: Solution
5:23
Example 2: Balance the Following Reaction
13:56
Galvanic Cells
18:15
Example 3: Galvanic Cells
28:19
Example 4: Galvanic Cells
35:12
Cell Potential

48m 41s

Intro
0:00
Cell Potential
2:08
Definition of Cell Potential
2:17
Symbol and Unit
5:50
Standard Reduction Potential
10:16
Example Figure 1
13:08
Example Figure 2
19:00
All Reduction Potentials are Written as Reduction
23:10
Cell Potential: Important Fact 1
26:49
Cell Potential: Important Fact 2
27:32
Cell Potential: Important Fact 3
28:54
Cell Potential: Important Fact 4
30:05
Example Problem 1
32:29
Example Problem 2
38:38
Potential, Work, & Free Energy

41m 23s

Intro
0:00
Potential, Work, Free Energy
0:42
Descriptions of Galvanic Cell
0:43
Line Notation
5:33
Example 1
6:26
Example 2
11:15
Example 3
15:18
Equation: Volt
22:20
Equations: Cell Potential, Work, and Charge
28:30
Maximum Cell Potential is Related to the Free Energy of the Cell Reaction
35:09
Example 4
37:42
Cell Potential & Concentration

34m 19s

Intro
0:00
Cell Potential & Concentration
0:29
Example 1: Question
0:30
Example 1: Nernst Equation
4:43
Example 1: Solution
7:01
Cell Potential & Concentration
11:27
Example 2
16:38
Manipulating the Nernst Equation
25:15
Example 3
28:43
Electrolysis

33m 21s

Intro
0:00
Electrolysis
3:16
Electrolysis: Part 1
3:17
Electrolysis: Part 2
5:25
Galvanic Cell Example
7:13
Nickel Cadmium Battery
12:18
Ampere
16:00
Example 1
20:47
Example 2
25:47
XII. Light
Light

44m 45s

Intro
0:00
Light
2:14
Introduction to Light
2:15
Frequency, Speed, and Wavelength of Waves
3:58
Units and Equations
7:37
Electromagnetic Spectrum
12:13
Example 1: Calculate the Frequency
17:41
E = hν
21:30
Example 2: Increment of Energy
25:12
Photon Energy of Light
28:56
Wave and Particle
31:46
Example 3: Wavelength of an Electron
34:46
XIII. Quantum Mechanics
Quantum Mechanics & Electron Orbitals

54m

Intro
0:00
Quantum Mechanics & Electron Orbitals
0:51
Quantum Mechanics & Electron Orbitals Overview
0:52
Electron Orbital and Energy Levels for the Hydrogen Atom
8:47
Example 1
13:41
Quantum Mechanics: Schrodinger Equation
19:19
Quantum Numbers Overview
31:10
Principal Quantum Numbers
33:28
Angular Momentum Numbers
34:55
Magnetic Quantum Numbers
36:35
Spin Quantum Numbers
37:46
Primary Level, Sublevels, and Sub-Sub-Levels
39:42
Example
42:17
Orbital & Quantum Numbers
49:32
Electron Configurations & Diagrams

34m 4s

Intro
0:00
Electron Configurations & Diagrams
1:08
Electronic Structure of Ground State Atom
1:09
Order of Electron Filling
3:50
Electron Configurations & Diagrams: H
8:41
Electron Configurations & Diagrams: He
9:12
Electron Configurations & Diagrams: Li
9:47
Electron Configurations & Diagrams: Be
11:17
Electron Configurations & Diagrams: B
12:05
Electron Configurations & Diagrams: C
13:03
Electron Configurations & Diagrams: N
14:55
Electron Configurations & Diagrams: O
15:24
Electron Configurations & Diagrams: F
16:25
Electron Configurations & Diagrams: Ne
17:00
Electron Configurations & Diagrams: S
18:08
Electron Configurations & Diagrams: Fe
20:08
Introduction to Valence Electrons
23:04
Valence Electrons of Oxygen
23:44
Valence Electrons of Iron
24:02
Valence Electrons of Arsenic
24:30
Valence Electrons: Exceptions
25:36
The Periodic Table
27:52
XIV. Intermolecular Forces
Vapor Pressure & Changes of State

52m 43s

Intro
0:00
Vapor Pressure and Changes of State
2:26
Intermolecular Forces Overview
2:27
Hydrogen Bonding
5:23
Heat of Vaporization
9:58
Vapor Pressure: Definition and Example
11:04
Vapor Pressures is Mostly a Function of Intermolecular Forces
17:41
Vapor Pressure Increases with Temperature
20:52
Vapor Pressure vs. Temperature: Graph and Equation
22:55
Clausius-Clapeyron Equation
31:55
Example 1
32:13
Heating Curve
35:40
Heat of Fusion
41:31
Example 2
43:45
Phase Diagrams & Solutions

31m 17s

Intro
0:00
Phase Diagrams and Solutions
0:22
Definition of a Phase Diagram
0:50
Phase Diagram Part 1: H₂O
1:54
Phase Diagram Part 2: CO₂
9:59
Solutions: Solute & Solvent
16:12
Ways of Discussing Solution Composition: Mass Percent or Weight Percent
18:46
Ways of Discussing Solution Composition: Molarity
20:07
Ways of Discussing Solution Composition: Mole Fraction
20:48
Ways of Discussing Solution Composition: Molality
21:41
Example 1: Question
22:06
Example 1: Mass Percent
24:32
Example 1: Molarity
25:53
Example 1: Mole Fraction
28:09
Example 1: Molality
29:36
Vapor Pressure of Solutions

37m 23s

Intro
0:00
Vapor Pressure of Solutions
2:07
Vapor Pressure & Raoult's Law
2:08
Example 1
5:21
When Ionic Compounds Dissolve
10:51
Example 2
12:38
Non-Ideal Solutions
17:42
Negative Deviation
24:23
Positive Deviation
29:19
Example 3
31:40
Colligatives Properties

34m 11s

Intro
0:00
Colligative Properties
1:07
Boiling Point Elevation
1:08
Example 1: Question
5:19
Example 1: Solution
6:52
Freezing Point Depression
12:01
Example 2: Question
14:46
Example 2: Solution
16:34
Osmotic Pressure
20:20
Example 3: Question
28:00
Example 3: Solution
30:16
XV. Bonding
Bonding & Lewis Structure

48m 39s

Intro
0:00
Bonding & Lewis Structure
2:23
Covalent Bond
2:24
Single Bond, Double Bond, and Triple Bond
4:11
Bond Length & Intermolecular Distance
5:51
Definition of Electronegativity
8:42
Bond Polarity
11:48
Bond Energy
20:04
Example 1
24:31
Definition of Lewis Structure
31:54
Steps in Forming a Lewis Structure
33:26
Lewis Structure Example: H₂
36:53
Lewis Structure Example: CH₄
37:33
Lewis Structure Example: NO⁺
38:43
Lewis Structure Example: PCl₅
41:12
Lewis Structure Example: ICl₄⁻
43:05
Lewis Structure Example: BeCl₂
45:07
Resonance & Formal Charge

36m 59s

Intro
0:00
Resonance and Formal Charge
0:09
Resonance Structures of NO₃⁻
0:25
Resonance Structures of NO₂⁻
12:28
Resonance Structures of HCO₂⁻
16:28
Formal Charge
19:40
Formal Charge Example: SO₄²⁻
21:32
Formal Charge Example: CO₂
31:33
Formal Charge Example: HCN
32:44
Formal Charge Example: CN⁻
33:34
Formal Charge Example: 0₃
34:43
Shapes of Molecules

41m 21s

Intro
0:00
Shapes of Molecules
0:35
VSEPR
0:36
Steps in Determining Shapes of Molecules
6:18
Linear
11:38
Trigonal Planar
11:55
Tetrahedral
12:45
Trigonal Bipyramidal
13:23
Octahedral
14:29
Table: Shapes of Molecules
15:40
Example: CO₂
21:11
Example: NO₃⁻
24:01
Example: H₂O
27:00
Example: NH₃
29:48
Example: PCl₃⁻
32:18
Example: IF₄⁺
34:38
Example: KrF₄
37:57
Hybrid Orbitals

40m 17s

Intro
0:00
Hybrid Orbitals
0:13
Introduction to Hybrid Orbitals
0:14
Electron Orbitals for CH₄
5:02
sp³ Hybridization
10:52
Example: sp³ Hybridization
12:06
sp² Hybridization
14:21
Example: sp² Hybridization
16:11
σ Bond
19:10
π Bond
20:07
sp Hybridization & Example
22:00
dsp³ Hybridization & Example
27:36
d²sp³ Hybridization & Example
30:36
Example: Predict the Hybridization and Describe the Molecular Geometry of CO
32:31
Example: Predict the Hybridization and Describe the Molecular Geometry of BF₄⁻
35:17
Example: Predict the Hybridization and Describe the Molecular Geometry of XeF₂
37:09
XVI. AP Practice Exam
AP Practice Exam: Multiple Choice, Part I

52m 34s

Intro
0:00
Multiple Choice
1:21
Multiple Choice 1
1:22
Multiple Choice 2
2:23
Multiple Choice 3
3:38
Multiple Choice 4
4:34
Multiple Choice 5
5:16
Multiple Choice 6
5:41
Multiple Choice 7
6:20
Multiple Choice 8
7:03
Multiple Choice 9
7:31
Multiple Choice 10
9:03
Multiple Choice 11
11:52
Multiple Choice 12
13:16
Multiple Choice 13
13:56
Multiple Choice 14
14:52
Multiple Choice 15
15:43
Multiple Choice 16
16:20
Multiple Choice 17
16:55
Multiple Choice 18
17:22
Multiple Choice 19
18:59
Multiple Choice 20
20:24
Multiple Choice 21
22:20
Multiple Choice 22
23:29
Multiple Choice 23
24:30
Multiple Choice 24
25:24
Multiple Choice 25
26:21
Multiple Choice 26
29:06
Multiple Choice 27
30:42
Multiple Choice 28
33:28
Multiple Choice 29
34:38
Multiple Choice 30
35:37
Multiple Choice 31
37:31
Multiple Choice 32
38:28
Multiple Choice 33
39:50
Multiple Choice 34
42:57
Multiple Choice 35
44:18
Multiple Choice 36
45:52
Multiple Choice 37
48:02
Multiple Choice 38
49:25
Multiple Choice 39
49:43
Multiple Choice 40
50:16
Multiple Choice 41
50:49
AP Practice Exam: Multiple Choice, Part II

32m 15s

Intro
0:00
Multiple Choice
0:12
Multiple Choice 42
0:13
Multiple Choice 43
0:33
Multiple Choice 44
1:16
Multiple Choice 45
2:36
Multiple Choice 46
5:22
Multiple Choice 47
6:35
Multiple Choice 48
8:02
Multiple Choice 49
10:05
Multiple Choice 50
10:26
Multiple Choice 51
11:07
Multiple Choice 52
12:01
Multiple Choice 53
12:55
Multiple Choice 54
16:12
Multiple Choice 55
18:11
Multiple Choice 56
19:45
Multiple Choice 57
20:15
Multiple Choice 58
23:28
Multiple Choice 59
24:27
Multiple Choice 60
26:45
Multiple Choice 61
29:15
AP Practice Exam: Multiple Choice, Part III

32m 50s

Intro
0:00
Multiple Choice
0:16
Multiple Choice 62
0:17
Multiple Choice 63
1:57
Multiple Choice 64
6:16
Multiple Choice 65
8:05
Multiple Choice 66
9:18
Multiple Choice 67
10:38
Multiple Choice 68
12:51
Multiple Choice 69
14:32
Multiple Choice 70
17:35
Multiple Choice 71
22:44
Multiple Choice 72
24:27
Multiple Choice 73
27:46
Multiple Choice 74
29:39
Multiple Choice 75
30:23
AP Practice Exam: Free response Part I

47m 22s

Intro
0:00
Free Response
0:15
Free Response 1: Part A
0:16
Free Response 1: Part B
4:15
Free Response 1: Part C
5:47
Free Response 1: Part D
9:20
Free Response 1: Part E. i
10:58
Free Response 1: Part E. ii
16:45
Free Response 1: Part E. iii
26:03
Free Response 2: Part A. i
31:01
Free Response 2: Part A. ii
33:38
Free Response 2: Part A. iii
35:20
Free Response 2: Part B. i
37:38
Free Response 2: Part B. ii
39:30
Free Response 2: Part B. iii
44:44
AP Practice Exam: Free Response Part II

43m 5s

Intro
0:00
Free Response
0:12
Free Response 3: Part A
0:13
Free Response 3: Part B
6:25
Free Response 3: Part C. i
11:33
Free Response 3: Part C. ii
12:02
Free Response 3: Part D
14:30
Free Response 4: Part A
21:03
Free Response 4: Part B
22:59
Free Response 4: Part C
24:33
Free Response 4: Part D
27:22
Free Response 4: Part E
28:43
Free Response 4: Part F
29:35
Free Response 4: Part G
30:15
Free Response 4: Part H
30:48
Free Response 5: Diagram
32:00
Free Response 5: Part A
34:14
Free Response 5: Part B
36:07
Free Response 5: Part C
37:45
Free Response 5: Part D
39:00
Free Response 5: Part E
40:26
AP Practice Exam: Free Response Part III

28m 36s

Intro
0:00
Free Response
0:43
Free Response 6: Part A. i
0:44
Free Response 6: Part A. ii
3:08
Free Response 6: Part A. iii
5:02
Free Response 6: Part B. i
7:11
Free Response 6: Part B. ii
9:40
Free Response 7: Part A
11:14
Free Response 7: Part B
13:45
Free Response 7: Part C
15:43
Free Response 7: Part D
16:54
Free Response 8: Part A. i
19:15
Free Response 8: Part A. ii
21:16
Free Response 8: Part B. i
23:51
Free Response 8: Part B. ii
25:07
Loading...
This is a quick preview of the lesson. For full access, please Log In or Sign up.
For more information, please see full course syllabus of AP Chemistry
  • Discussion

  • Study Guides

  • Download Lecture Slides

  • Table of Contents

  • Transcription

  • Related Books & Services

Lecture Comments (14)

1 answer

Last reply by: Professor Hovasapian
Wed Jan 18, 2017 8:03 PM

Post by Vivian Wang on January 8, 2017

Hi Professor Raffi,

I still don't really understand why in the common ion effect example when you removed the 2x and said that it was negligible compared to the 0.03. May you please explain again why? and does this apply to all ksp problems? like we always regard it as negligible?

Thanks!
Vivian

1 answer

Last reply by: Professor Hovasapian
Thu Apr 7, 2016 2:04 AM

Post by john lee on April 2, 2016

What if I add NaF and CaF2 at the same time and speed? Can we know which one will dissolve more by K?

1 answer

Last reply by: Professor Hovasapian
Sat Aug 29, 2015 8:37 PM

Post by Derek Marshall on August 28, 2015

Hey Professor,

    I had a question about the Ca(OH)2 dissociation.  I was wondering would it be faster/more efficient/more beneficial to add acid or just to add more water to the solution if you wanted to dissolve all the solid?

Thanks!
Derek Marshall

1 answer

Last reply by: Professor Hovasapian
Tue Apr 7, 2015 10:59 PM

Post by Lyngage Tan on April 4, 2015

in example 2. Ba3(PO4)2 -> 3Ba+ + 2PO4^2- i think the charges should be 3Ba^2+ + 2PO4 ^ 3- so we end up with 6+ and 6- charges to make the molecule neutral. pls correct me if i made a mistake on this analysis. thanks.

1 answer

Last reply by: Marcus Kirkegaard
Thu Dec 12, 2013 7:02 AM

Post by Marcus Kirkegaard on December 12, 2013

I dont get how in forholde the common ion effect two Q sow solve disse kinds of problems? Zn 2 + ions can be precipitated as insoluble Zn (OH) 2 at sufficiently high [OH-]. Indicate
from Zn (OH) 2 solubility product if Zn 2 + ion remains in solution by dissolving the
0.1 M ZnCl 2 in water. Ksp of Zn (OH) 2 is given as 1.8 * 10 ^ -14
??

3 answers

Last reply by: Professor Hovasapian
Thu Aug 1, 2013 7:17 PM

Post by Christian Fischer on July 31, 2013

Hi Raffi.

Can you help me clarify if the following is a correct understanding of the relationship between  Solubility and Solubility Product? I really hope this is as clear and simple as i think it sounds :)

If the Ksp of CaF2=4,0*10^(-11)  Then this means that at equilibrium the product [Ca][F]^2 will ALWAYS be 4,0*10^(-11). So if we add 0.001 moles of NaF, the reaction CaF2 = Ca + 2F will be pushed left due to excess of Fluoride ( because of common ion effect) until the product [Ca][F]^2 equals the value 4,0*10^(-11), and therefore less CaF2 dissociates at equilibrium so the solubility has changed at equilibrium due to common ion effect but NOT (and NEVER) the ksp?

Second: If we add 1000moles of Ca(+2) and we only have 0.01moles of fluoride then the calculated ksp product {Ca][F]^2 would obviously be off since Calcium is a limiting reagent and the concentration of Fluoride  is too large.. Does that mean that when we have a limiting reagent or too much of another ion, we can't calculate the equilibrium constant Ksp?

Thanks again for your great videos! Have a great day,
Christian

Related Articles:

Solubility Equilibria, Part II

  • The Common Ion effect is very pronounced for insoluble salts.
  • Certain salts become more soluble as the pH of the solution drops. This is just an application of Le Chatelier’s Principle.
  • Just as for regular Equilibrium, where we defined the Reaction Quotient Q, we have the analogous Qsp.
  • Qsp can be calculated with concentrations at any time, then compared to Ksp to decide if a precipitate will form.

Solubility Equilibria, Part II

Lecture Slides are screen-captured images of important points in the lecture. Students can download and print out these lecture slide images to do practice problems as well as take notes while watching the lecture.

  • Intro 0:00
  • Solubility Equilibria 0:46
    • Common Ion Effect
    • Example 1
    • pH & Solubility
    • Example of pH & Solubility
    • Example 2
    • Precipitation & Definition of the Ion Product
    • If Q > Ksp
    • If Q < Ksp
    • Example 3

Transcription: Solubility Equilibria, Part II

Hello, and welcome back to Educator.com, and welcome back to AP Chemistry.0000

Today, we are going to continue our discussion of solubility; it's going to be the second part of solubility, and there is going to be one more part after this.0004

This is a very, very important application of the equilibrium concept--of the aqueous equilibrium concept.0011

So, we want to make sure (the problems are sufficiently different among themselves that we want to make sure) that we have a really solid grasp of this, because again, if this acid-base equilibria, solubility equilibria, and subsequent lessons (complex ion equilibria)--if we get a good grasp of what is going on, and if we are able to handle the chemistry, then virtually everything in chemistry and biology will be an absolute walk in the park, I promise you.0017

It really just does get easier when you understand this.0042

OK, let's get started.0044

Today, we are going to talk about something called the common ion effect for solubility.0047

Now, we discussed the common ion effect when we talked about acid-base equilibria earlier; this is pretty much the same thing--it's just an application of Le Chatelier's Principle.0052

I'll go ahead and define it and discuss it really briefly, and we'll just launch into some examples.0062

Common ion effect: the solubility of an ionic solid (salt, in other words) decreases if the solution into which you drop the solid already contains an ion (or ions, actually--I should say "already contains ions," because you might have one or both) common to the solid.0067

OK, so again, the solubility of an ionic solid decreases if the solution into which you drop the solid already contains ions common to the solid.0140

For example, let's say, if you have a beaker, and in there you decide to throw in some sodium chloride (salt), which is completely soluble; so you mix it up, and now you have some free chloride ion and some free sodium ion floating around.0148

Well, now if you end up dropping, let's say, lead chloride or silver chloride (which are not soluble salts--there is a little bit of solubility)--if you drop them in there, because the chloride ion is already in there, it's going to actually suppress the dissociation of the silver chloride.0160

Let's just go ahead and do a problem, and I think it will make more sense, as opposed to doing sort of a run-down explanation, and then doing the problem.0180

As we do the problem, I'll talk about what it is that is actually going on, and I'll draw you a picture of the solution.0188

Example 1: Calculate the solubility (in other words, how much dissolves--an amount per volume) of calcium fluoride (which is CaF2) in a 0.030 Molar sodium fluoride solution.0194

OK, now: up here in the definition, I said that you have this solution that has a common ion, and then you drop some of the salt in it; it could be the other way around--you could drop the silver chloride in there, and then add, let's say, the sodium chloride to that.0224

It doesn't matter the order in which you do it, because again, once the system comes to equilibrium, that Ksp expression--it is a measure of equilibrium concentration.0237

It really doesn't matter which order you do it; just for the purposes of definition, I decided to choose the fact that the solution already has the chloride ion, and into it I'm dropping the silver chloride.0248

It doesn't matter the order in which you do that.0262

Calculate the solubility of a calcium fluoride in a .030 Molar sodium fluoride solution.0265

OK, so the Ksp of calcium fluoride, which we do need, is equal to 4.0x10-11; so just by looking at the Ksp, we know that the calcium fluoride is very, very insoluble--very little of it actually dissociates.0272

Well, when we are doing solubility product, as we do any kind of chemistry, we write a reaction, just to get a sense of it, so we know what is going on, so we can follow it.0292

We don't want to do things in our head.0301

Oftentimes, just by writing things down, pieces come together.0303

CaF2 is going to be in equilibrium with one calcium ion (which is a 2+ charge), and it's going to be in equilibrium with two...I'm sorry, it dissociates into one calcium and two fluoride ion.0306

Now, the Ksp expression for this is equal to Ca2+, times F- squared (right?--law of mass action, equilibrium constant).0322

Here is what is happening: we have this solution, which is .030 Molar sodium fluoride; so I have some sodium ion floating around and fluoride ion floating around, because sodium fluoride is fully soluble salt, like sodium chloride.0334

OK, so there is my .03 Molar; now, I'm going to drop some calcium fluoride into here: that is what I'm doing--I'm taking this calcium fluoride, and I'm dropping it into this solution, and I want to see--I want to calculate how much of it actually dissolves in this solution.0350

Well, we did a solubility problem earlier, in the last lesson, but that was in pure water; now, there is this common ion, fluoride.0373

You see, that is what we mean by common ion: the ion that is common to the salt is already in solution.0383

So, what is going to happen, by Le Chatelier's Principle: since there is fluoride ion already in there, it is as if you take this equilibrium, and it's as if you have increased the concentration of this fluoride ion.0390

What it is actually going to do is: it's going to push the equilibrium that way.0402

What that means is that it's going to suppress the dissociation; another way of thinking about it is: because of the presence of the fluoride ion already in solution, this calcium fluoride, solid--when you drop it in solution, it's not going to dissociate very much--even less that what it normally does (based on this), simply because there is fluoride ion already in solution.0406

If there is fluoride ion already in solution, it's not going to allow it to dissolve.0428

That is the whole idea behind the common ion effect: it tends to push a reaction back, or it tends to pull it forward, depending on what is happening, chemically.0432

Let's go ahead and just run this problem; this is a simple ICE chart, except now, the ICE chart looks slightly different.0442

Again, just let the chemistry--the situation--decide what the ICE chart looks like.0449

Let me write my dissociation again: CaF2 dissociates into calcium 2+, plus 2 F-; I have an Initial; I have a Change; I have an Equilibrium.0454

Well, this--it doesn't matter what it is; it's a solid, so we don't really care what happens to it; there is no calcium to begin with in the solution; but the fluoride ion--before (again, Initial--this is before) anything happens, before...when you drop it in, before the system comes to equilibrium, before it dissociates, the fluoride ion concentration is the concentration of fluoride ion already in there.0470

Well, we know that that is 0.030 Molar.0497

Well, now the Change: it starts to dissociate; a certain amount is going to dissolve; a certain amount is going to show up, 1:1; and this is 1:2--2x is going to show up in terms of the fluoride ion (right?--this is 1:1; this is 1:2).0503

For every mole of this that dissociates, for every unit of this, it creates 2 units of fluoride ion; that is what this x and 2x are.0521

Our equilibrium concentrations are: this one doesn't matter; this one is x; and this one is 0.030+2x.0529

These are our equilibrium concentrations of calcium and fluoride respectively; so now, we go ahead, and we put them into the Ksp expression.0539

I don't know if I should do this on the same page or not; let me see...I don't know; that is OK--I'll go ahead and do it on the next page--it's not a problem.0549

OK, so now the Ksp (let me write the expression again)--it equals the calcium 2+ concentration (that is calcium, not copper), and the F- concentration squared.0563

Well, the calcium concentration is x; the fluoride ion concentration is 0.030+2x (let me make that 2x a little more clear), squared; OK.0576

Now, we are going to have a "squared" here; we are going to end up multiplying it, x; we are going to have a cubic equation; again, for those of you that are dealing with graphing calculators, a cubic equation is not a problem.0593

Or, if you are dealing with mathematical software, a cubic equation is definitely not a problem--you will get your answer back in a microsecond.0603

You know what, when you are sort of in a fast test, and you don't...you sort of want some approximation techniques.0610

We do the same approximation that we did before: we know that calcium fluoride is already not very soluble (we have a Ksp which is 10-11); we know that the common ion, fluoride, is going to suppress that dissociation even more.0616

So, this x--this x right here--compared to the .030, it's probably going to be really, really small; so for all practical purposes, we can just ignore it.0630

We are going to rewrite this as...put it approximately, and you will see it is a fantastic, fantastic approximation...x times 0.030 squared.0640

And again, it is this x that you drop, not this x.0652

Well, when we do this calculation, we get x equal to 4.4x10-8; that is the solubility--remember, x--that was the amount that dissolved; that was the amount of calcium that showed up; twice that, is the fluoride that showed up.0656

So, this is in moles per liter--molarity--so 4.4x10-8 moles of calcium fluoride dissolve if you happen to drop it in a solution which is already .030 Molar sodium fluoride.0680

Again, the presence of the fluoride suppresses the dissociation--doesn't allow the formation--because when it dissociates, it forms fluoride ion.0696

There is already fluoride ion in solution, so it's going to keep it from dissociating.0704

Now, let's go ahead and check the validity of this approximation.0708

We go by the 5% rule; so we'll take 4.4x10-8, divided by the 0.030 (that is what we compared it to), times 100.0712

Well, get a look at this: this is 0.000148%.0726

This is a percentage; there are actually 2 more decimal places, if you want to just use the decimal.0733

So, .000148% is a lot less than 5%; it's not even measurable, for all practical purposes; so we are absolutely, absolutely justified, in almost every single solubility problem, for reducing this--eliminating it--and just taking the approximation the way that we have done before.0738

We are not doing anything new here.0758

There you go: common ion effect--the only thing that is different is the ICE chart--in the line that says Initial, instead of having just one thing and then 00 for the ions, one of those ions is going to have a number, because there is going to be a certain amount of it in solution when you drop the salt in there.0759

OK, now let's move on to the second aspect of the common ion effect: pH and solubility.0779

As it turns out, the pH of a solution that you drop something in...depending on what you are dropping in, it has an effect on the extent to which it is actually soluble.0791

For example, something called milk of magnesia--maybe your parents might remember it--it is a suspension of magnesium hydroxide.0803

Magnesium hydroxide is not a very soluble salt; however, in acid solution, it becomes almost completely soluble.0811

The idea is: you drink it as this sort of suspension of solid particles in a liquid--you can taste it; it's kind of chalky and stuff--but when it comes into the acid, the hydrochloric acid is acidic solution; it actually dissolves it according to what is necessary.0819

Now, I'm going to show you what the chemistry actually looks like; but let me go ahead and define what is going on first.0835

The pH of a solution can (that doesn't mean "will"; it can) affect the solubility of a salt, because the H+ or OH- ions in solution (right?--because a less-than-7 pH--that means there is a high concentration of hydrogen ion; a pH of greater than 7--there is a high concentration of hydroxide ion--these are free ions floating around) are going to react with the free ions of the slightly dissociated salt, and it's going to change the chemistry.0842

So, because the hydrogen ion or the hydroxide ions in solution will suppress or induce--so it could push it either way--dissociation (dissociation is just dissolving--dissociation is the fancy term for it)...0900

OK, so here is the example: we are going to use calcium hydroxide, simply because magnesium hydroxide is always used; it's the same chemistry.0925

Calcium hydroxide is Ca(OH)2; it has a Ksp...I'll actually write it in just a minute; but it dissociates into Ca2+ + 2 OH-.0935

It has a Ksp equal to 1.3x10-6--not very soluble.0948

Now, watch what happens: if we take this and just sort of drop it into normal solution, we can change the pH a couple of ways: we can add hydroxide solution to it and raise the pH, making it more basic; or we can drop acid into the solution, raising the hydrogen ion concentration and dropping the pH below 7.0956

Watch what happens: now, if we add OH (so adding OH---in other words, raising the pH), we'll suppress dissociation (I'll call it...yes, I'll call it dissociation) by the common ion effect that we just discussed.0976

See: one of the things that is produced upon dissociation is sodium hydroxide; well, if I add hydroxide to the solution, now I'm increasing this hydroxide ion concentration; I'm raising it.1006

By Le Chatelier's Principle, the system is going to adjust in order to lower this hydroxide back down.1018

The only way to lower the hydroxide is to push the equilibrium this way by forming calcium hydroxide, solid--taking hydroxides out of free solution and sequestering them as the solid.1024

It's going to push the equilibrium this way, so it's going to suppress the dissociation.1036

The dissociation is to the right; it will suppress the dissociation by the common ion effect, right?1041

One of the ions produced is the hydroxide; if you add hydroxide, which is a common, or if there happens to be hydroxide ion in solution (like the previous problem, which was already in solution when we dropped it in there), it's going to suppress the dissociation.1054

However, adding H+ will increase the solubility, and here is why (you can almost imagine why).1067

It will increase the solubility by reacting with the OH- formed from dissociation to form water.1077

Well, if you drop this in solution in a certain amount of hydroxide, and now it's floating around in solution a little bit; well, now, if you add hydrogen ion to it (an acid, lowering the pH), that acid that you drop is going to react with the OH.1116

Remember, H+ and OH, whenever they are in proximity--they form water.1131

Well, when they form water, this hydroxide is going to be pulled--is going to be eaten up; as this is eaten up, it's going to leave an emptiness here.1134

In order for this Ksp to be viable, this is going to dissociate more and create more hydroxide.1142

As you add more hydrogen ion, it's going to react with the hydroxide; it's going to form water, leaving a dearth of hydroxide; therefore, some more is going to dissolve.1151

Some more is going to dissolve; so the more acid I actually add...some more is going to dissolve.1161

The acid is not actually dissolving the solid; the acid is reacting with the hydroxide, which creates a depletion of hydroxide ion; therefore, by Le Chatelier's Principle, this reaction will move in a direction to produce more hydroxide ion that has been depleted.1166

There is a very big difference: the acid is not eating up the solid when you drop acid into it: it's actually an application of Le Chatelier's Principle that is taking effect.1184

OK, to form water...let me finish this...removing OH (it might be nice if I actually wrote the word so that you could read it)...removing OH-...pulling the reaction forward.1195

OK, let's see what this actually looks like.1222

OK, so here is what is actually happening: let's see, let's just use one unit of calcium hydroxide, OH-, OH-; OK.1228

There is, of course, a bunch of calcium hydroxide, solid, at the bottom; so this is calcium hydroxide, solid.1240

They are in equilibrium with each other; this is a saturated solution--remember, the saturated solution means that the solid is in equilibrium with the free ions according to the dissociation reaction, and there is a solubility product associated with that.1249

The concentration of calcium, times the concentration of free hydroxide squared (because two units are produced upon dissociation)--those two concentrations, multiplied, equals the Ksp for this particular salt at that particular temperature.1261

That is what the Ksp measures.1275

OK, well, here is what happens: when I take hydrogen ion (so in this case, let's add hydrogen ion)--when I add hydrogen ion--when I make the solution acidic, or when it's already acidic and I drop the salt in there, here is what happens.1277

I'll do 2 H+s; one of these H+s reacts with one of these OH-: H+ + OH- forms neutral water.1293

That went away; the second H+, plus the second OH-, forms another molecule of water; that went away.1305

It is reacting, and it's eating up the hydroxide; now, what you end up with is a solution that just has calcium in it, and some--again, this is the calcium hydroxide, solid; that is that; and now, basically, because you have this calcium left over, now some more of this is going to dissolve according to: calcium hydroxide goes to calcium 2+, plus OH-, 2.1315

Because now there is no hydroxide floating around, more of this will dissolve in order to release OH-, OH-; and it will release another calcium--the calcium itself doesn't really matter all that much.1357

But, because there is no OH- in here, it wants to move in this direction; in order to move in this direction to produce hydroxide ion, this has to dissolve--that is the whole idea.1371

OK, let's see: let's do another example here, Example 2, for this particular thing; let's do barium phosphate.1384

Barium phosphate is Bo3(PO4)2; it is in equilibrium with 3 barium ion, plus 2 phosphate ion.1398

So again, the pH is going to affect the solubility of a salt if either the OH- that you add or the H+ that you add (or if those OH- or H+ are already in solution), if they react with one of these ions.1410

That is the whole idea: when they react with one of these ions, they are going to...well, they are either going to suppress it and push the equilibrium that way, or they are going to react with it and take it out of solution by reacting with it, pulling the reaction that way--in other words, making this more soluble or less soluble.1428

Well, if we have barium phosphate in acidic solution, or upon addition (again, the order doesn't matter) of H+, the following reaction takes place.1447

You remember phosphate: hydrogen phosphate is a weak acid, which means the phosphate is a reasonably strong conjugate base.1478

If it comes in contact with H+, it's going to react with the H+ according to: PO43- + H+ is going to almost completely form hydrogen phosphate 2-.1487

Again, because this is this is the conjugate...one of the products of the dissociation of the salt happens to be something that will react with hydrogen ion to produce hydrogen phosphate.1504

When it reacts with hydrogen ion, it pulls this out of solution--it sequesters it; it lowers the concentration of free phosphate ion.1514

When you lower the concentration of free phosphate ion, this equilibrium is going to shift in the direction which increases the phosphate ion.1523

The only way to do that is by creating more phosphate ion; the only way to do that is by dissociating, by dissolving.1531

It is going to pull the reaction that way; that is what is going on with pH.1538

pH is about hydrogen ion-hydroxide ion concentration; if one of the ions from the dissolving salt reacts with the hydrogen ion or the hydroxide, the equilibrium is going to shift one way or the other.1546

It is going to suppress dissociation by common ion effect; it is going to pull the reaction forward to your right--it's going to pull the reaction forward by reacting with one of the species that dissolves.1559

OK, now let's go ahead and talk about something called precipitation, itself.1570

We have talked about solubility; we have talked about Ksp; now what happens when we actually mix solutions that have certain salts that, when they come together, they precipitate?1579

You remember way back, in the early part of the course, we talked about precipitation reactions, using the solubility chart to decide which salts will precipitate and which will not?1590

Well, now we are to discuss it quantitatively; we are going to use the molarities and volumes to decide whether a precipitate will form or will not form.1598

Let's talk about precipitation, and we'll do a problem, and we'll finish up this lesson.1609

Precipitation (oh, I'm going to need to spell it right; let's try this one more time): OK, we're going to define something called the ion product.1615

Define the ion product: well, it's very, very easy--it's the same as the Ksp expression, but it represents initial concentrations--initial concentrations, not equilibrium concentrations; initial--very, very important--initial concentrations, not equilibrium concentrations.1634

Do you remember back when we did equilibrium--we were just discussing equilibrium in general?1681

We had this thing called Q; we said it is the concentration of species in that particular flask (or whatever) at any given time (not necessarily at equilibrium).1685

This is analogous; and, in fact, we use Q again.1697

So, let's do...for example, we have been working with calcium fluoride, so let's go ahead and stick with calcium fluoride.1700

2 F-: let's write the Ksp expression; this is at equilibrium--equals calcium 2+ concentration, times F- squared (that is the law of mass action).1709

Q (or Qsp--you can write Q; I tend to write Qsp, simply just to sort of remind myself that I'm dealing with solubilities--either one of those is fine): it is the initial calcium ion concentration (so I'll put a little subscript 0, which means "initial")--times the initial fluoride ion concentration, squared.1721

Initial means...when you first drop certain amounts of these ions into solution, when you mix solutions, before any reaction takes place, before anything comes together to form solid calcium fluoride (if it does)...that is the initial concentration.1751

So, here is the mathematics of it--well, not the mathematics--here is how you decide.1771

If Q is greater than the Ksp (because you know that the Ksp is defined for every single salt at a given...), then a precipitate will form, and continue to form until the ion concentrations satisfy the Ksp.1779

So again...well, actually, let me write the other one, and then we'll talk about what this really means.1824

If Q is less than Ksp, then no precipitate forms.1833

OK, now recall: Ksp is a measure of the ion concentration in equilibrium with the solid.1849

I want to go ahead and write that down: recall, the Ksp is a measure of free ion concentrations in equilibrium with the solid from which they come.1859

OK, so, for example, let's go ahead and use...well, you know what, let's just go ahead and actually do a problem; again, I think if we do the problem--and we will discuss what is actually going on as we do the problem--I think it will make more sense.1901

OK, so...but remember, Ksp is a measure of the free ion concentrations in equilibrium with the solid from which it comes.1921

So, in the case of calcium fluoride (which is a solid), calcium 2+, plus 2 F-, this is what we call a saturated solution.1928

A saturated solution is when you have free ions that make up the particular solid, and you have solid that is visible in the solution.1940

Like we said, if you keep dropping a bunch of salt into a certain fixed amount of water, at some point the salt is not going to dissolve anymore; you are just going to stir it around, and the salt is going to settle at the bottom.1951

Well, that means there is some chloride ion; there is some sodium ion; and there is a bunch of solid sodium chloride; there is an equilibrium that is established there.1961

There is a Ksp associated with that.1970

When you...well, let's just actually do the problem; it will make more sense.1972

OK, so let me do this in blue.1977

Example: 750 milliliters of a 3.0x10-3 molarity silver nitrate solution is mixed with 300.0 milliliters of a 2.5x10-2 Molar potassium chloride solution.1984

Will a precipitate form?2034

OK, so notice: earlier, we were talking about just taking a certain solution and dropping some solid into it and seeing how much dissolves.2039

Now, what we are doing is: we are taking some solutions of free ion; we are mixing them together; and, if a couple of those ions form in soluble salts, we want to find out, based on how much of each free ion there is in there, if an actual solid will drop to the bottom--if a precipitate will suddenly appear.2052

Let's draw this out: we have a solution; we have 750 milliliters of a silver nitrate.2074

This is a silver nitrate solution, AgNO3; well, we know that things that have nitrate in them are completely soluble, so what you have is free silver ion and free nitrate ion floating around; it just looks like some water, because it's fully soluble.2081

Now, we also have another solution: we have 300 milliliters of this potassium chloride solution.2097

Well, potassium chloride is also a completely soluble salt, so we write KCl as a solution; but really, what you have is free potassium and free chloride--it's fully dissolved.2103

And again, these are not saturated solutions; there is no solid anywhere--these are just solutions of potassium chloride and silver nitrate.2114

Now, what we do is: we mix them together.2121

OK, when we mix them together, now what we have is, well, 750 milliliters, plus 300; now we have 1050 milliliters of the KCl and the AgNO3 solution, mixed.2124

Before any reaction takes place, what you have is free silver ion, free nitrate ion, free potassium ion, and free chloride ion.2141

These are the major species in solution; and again, major species--let's go take a look at these major species, and let's see if a reaction is going to take place.2152

Well, silver nitrate--if they slam back into each other, they are just going to dissolve again--they are fully soluble.2161

If potassium and chloride ion slam into each other, they are just going to dissolve again; that doesn't matter.2168

Positive and positive: they are not going to...when they slam into each other, they are just going to bounce right off.2173

So, the only two possibilities are potassium and nitrate coming together, and silver and chloride coming together.2178

Well, potassium nitrate is also soluble; so, if they slam together and they stick, it's just going to dissolve again; but, silver and chloride, when they slam into each other--the Ksp...let me see...the Ksp happens to be (let me actually write it here): the Ksp of AgCl equals 1.6x10-10.2187

Well, we know from our solubility chart that it is insoluble, so when these slam together, chances are they will precipitate.2213

However, we need to find out if they will precipitate, based on these amounts.2220

These concentrations of free silver and free chloride--they have to equal (the product has to equal) the Ksp before something will precipitate.2225

Remember, that is the whole idea: if it's equal to the Ksp or bigger, it will precipitate; if it's less than that, it hasn't reached the Ksp value; therefore, it is not a saturated solution.2234

The ions are free; there is no precipitate.2245

In this case, if we didn't know from our solubility chart that this is insoluble, we can just take a look at the Ksp and know that this is one of the possibilities, that AgCl may precipitate here, because we are looking at a small Ksp.2247

OK, so let's actually do the math.2262

So now, let's go ahead and write out the possible reaction: AgCl is going to be in equilibrium with Ag+ + Cl-, 1:1.2265

Our Ksp expression equals Ag+, Cl-; our Qsp is equal to the initial concentration of Ag, times the initial concentration (it's the same as that, except initial concentrations)...now, let's calculate this.2280

We'll calculate this; we'll multiply them together; we'll see what we get.2299

Let's rewrite the Ksp on this page: Ksp equals 1.6x10-10, for our reference.2303

OK, go ahead and move on to blue.2311

The silver ion concentration is going to equal...well, concentration is moles per liter; the total number of moles of silver, and the total volume.2316

Well, the total volume has changed; I added 750 milliliters to 300 milliliters; so now, that silver ion is not just floating around in 750 milliliters--it's floating around in 1050 milliliters.2325

I need to know how many moles that was.2336

Well, we had 750.0 milliliters of a 2.0 (I'm sorry, this says 3.0, not 2.0)...volume times molarity: 3.0x10-3 millimoles per milliliter; that is going to give us moles (millimoles), over the total number of liters/total number of milliliters/total volume.2338

Plus 300 milliliters (right?): the moles of silver, divided by the total volume that it is floating around in--you end up with 2.14x10-3 molarity.2372

Now, let's calculate the initial chloride ion concentration.2388

Well, it's equal to the total chloride ion and total volume; well, the initial chloride solution--we had 300 milliliters of it, and the molarity was 2.5x10-2 millimoles per milliliter.2392

When you multiply those two, that will give us the moles of chloride ion that we started off with in the second beaker.2412

And now we have mixed them, and again, our total volume is 750 + 300 milliliters, and we get 7.14x10-3 Molar.2418

Well, Qsp is equal to that times that (right?--this expression), and there are no exponents, so we're OK.2433

It is going to be 2.14x10-3, times 7.14x10-3; that equals 1.52x10-5, which is definitely bigger than the Ksp, which is 1.6x10-10; so yes, AgCl (a little arrow pointing down is an old symbol for a precipitate) will precipitate.2447

That is it: so this is a nice, quantitative method of letting us know whether something actually will precipitate, based on the product of the ion concentrations.2482

If it is bigger than the Ksp, then yes, a precipitate will show up.2495

If it is less than the Ksp...well, Ksp measures equilibrium; something that is in equilibrium with the solid from which those ions come--if these numbers, multiplied together, are less than the Ksp, well, that means there is no solid for there to be any equilibrium reached.2499

OK, thank you for joining us here at Educator.com to discuss solubility equilibria, and we have one more lesson to discuss that same subject.2517

See you next time; goodbye.2525

Educator®

Please sign in for full access to this lesson.

Sign-InORCreate Account

Enter your Sign-on user name and password.

Forgot password?

Start Learning Now

Our free lessons will get you started (Adobe Flash® required).
Get immediate access to our entire library.

Sign up for Educator.com

Membership Overview

  • Unlimited access to our entire library of courses.
  • Search and jump to exactly what you want to learn.
  • *Ask questions and get answers from the community and our teachers!
  • Practice questions with step-by-step solutions.
  • Download lesson files for programming and software training practice.
  • Track your course viewing progress.
  • Download lecture slides for taking notes.

Use this form or mail us to .

For support articles click here.