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Raffi Hovasapian

Raffi Hovasapian

Solubility Equilibria

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
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Lecture Comments (9)

1 answer

Last reply by: Professor Hovasapian
Wed Dec 20, 2017 7:28 AM

Post by Vishal Raman on December 20, 2017

Just for a clarification, you can still compare the solubilities of the ions with not 1:1 ratios, right?  You just need to calculate the solubility of each one individually.

1 answer

Last reply by: Professor Hovasapian
Fri Mar 4, 2016 9:46 PM

Post by Nadan Cha on March 4, 2016

Hello Professor,

Firstly, thank you for your lecture -- it is fantastic as always =).

For relative solubilities you said we can decide the order of solubilities only for salts that produce the same number of ions when they dissociate, right?

Do you mean the TOTAL number of ions they produce? So, if a salt produces 1 mol of ion A and 3 mol of ion B, can I compare solubility with a salt that produces 2 mol of ion C and 2 mol of ion D (because they both produce 4 mol of ions in total)?

Thank you so much!

1 answer

Last reply by: Professor Hovasapian
Tue Dec 30, 2014 11:48 PM

Post by Rafael Mojica on December 30, 2014

Please do and Organic Chemistry course, you will save thousands of premed lives!

0 answers

Post by Mohamed Yassin on November 30, 2013

Hi there, why is the connection to the site (Educator.com) is so slow even though i have good connection with other site. is your site under maintenance? why is continuously buffing and reversing as you watch to to the start of the video it is so annoying! please advise what to do.

Thank you

1 answer

Last reply by: Professor Hovasapian
Wed May 15, 2013 1:37 AM

Post by Kendrick Miyano on May 14, 2013

What do you mean by "The higher the Ksp, the farther the equilibrium is to the right,"?

Related Articles:

Solubility Equilibria

  • Insoluble salts do dissolve, but very, very little. Their Equilibrium constants are called Ksp.
  • ICE charts are still used, as these are Equilibrium problems.

Solubility Equilibria

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 0:48
    • Solubility Equilibria Overview
    • Solubility Product Constant
    • Definition of Solubility
    • Definition of Solubility Product
    • Example 1
    • Example 2
    • Example 3
    • Relative Solubilities

Transcription: Solubility Equilibria

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

Today, we are going to continue our discussion of aqueous equilibria, and we are going to discuss solubility equilibria.0004

We are going to discuss the equilibria involved in salts that actually don't dissolve very much, unlike the salts that you are accustomed to, like table salt, for example.0010

When you drop it into water, it completely dissociates into free sodium and free chloride; there are some salts that--they just basically--you drop them in water, and they just sink to the bottom as solids.0020

Now, they do dissolve a little bit, and we are actually able to measure how much they do dissolve; but for all practical purposes, they dissolve so little that, to the naked eye, it looks like they just sink to the bottom like sand.0031

OK, let's jump in and start with some definitions and see what we can do.0043

OK, so some salts (and again, "salt" is just a generic term for an ionic compound--metal-nonmetal) dissolve completely (oh, I really need to learn to spell here); those are the ones we know from our solubility chart way back when; those are soluble.0048

For example, sodium chloride (like we mentioned)--when you drop it into water (I'll just put water on top of the arrow here), you end up with (so this is a solid)--when you drop it into water--you get aqueous sodium ion and chloride ion; the aq means that they are floating around in solution.0083

Or, let's say, potassium nitrate (so KNO3, solid): drop it in water; you end up with free potassium ion, plus free nitrate ion.0101

I'm not going to put the subscripts; again, I think they're sort of superfluous at this point--I mean, we know what we are talking about; we are talking about soluble salts; we know that it's floating around in solution; any time you have an ion like this, we are talking about aqueous equilibria, so we know that it's in water.0115

OK, now, others dissolve only very little, where mostly the solid stays intact.0130

OK, so notice: these arrows go one way; that means that, when you drop this in water, all of it dissolves; when you drop this in water, all of it dissolves.0158

With certain salts that don't dissolve completely, an equilibrium is established; here there is no equilibrium--I mean, there is, but basically, there is no solid sodium chloride left anywhere (until you get to a saturate solution, which we will talk about in a subsequent lesson); but really, you are talking about complete dissociation.0169

Well, let's take an example like magnesium fluoride: it turns out, magnesium fluoride doesn't dissolve very much.0187

As a solid, when you drop it in water, it is true that some magnesium ion and fluoride ion do break free; but it's actually very, very little that does so.0194

So, there is this equilibrium that is established, and most of the equilibrium is, in fact, over here; which is why, when you drop this in the water, and you stir it around, it just sort of sinks to the bottom; it looks like nothing has dissolved.0210

Well, we know that at the molecular level, some has; and we are able to measure it, but most of it just sort of stays intact; but there is an equilibrium that is established.0222

So now, because of this equilibrium situation--because we have some that is dissolved, but not a lot of it--we can write an equilibrium constant for it.0232

Well, the equilibrium constant is just Mg2+ (the concentration--remember, it's just reactants over the products), times F-, raised to its coefficient, 2 (because we know that, when this dissociates, it produces 1 Mg ion and 2 fluoride ions; that is what happens here.0243

Well, we call this Ksp; instead of Keq, when we are dealing with solubility problems where we have a solid that is in equilibrium, in solution with its free ions, we call this the Ksp.0265

It is the called the solubility product constant.0289

The best part is: you are just talking about equilibrium: there is no new math that is going on here--there is no new concept.0298

All of the equilibrium that you have learned before applies here; it's all we have done--just sort of changed the label; instead of Keq, we call it Ksp.0303

It is called the solubility product constant, or just the solubility product; you will see both of those terms used.0312

Now, I just want to describe what is happening physically, because I want you to actually be able to think about this when you are doing your problems.0323

If I take some magnesium fluoride, solid, and if I drop it in solution, it basically piles up at the bottom, because it's not very soluble.0333

However, some of it does dissolve: you have a little bit of magnesium ion and a little bit of fluoride ion actually floating around in solution.0341

There is this equilibrium between the solid magnesium fluoride and its ions.0351

This is called a saturated solution--you will often hear it referred to as that: a saturated solution is one that is in equilibrium with the solid that is present in solution, at the bottom of the beaker.0359

That is what they mean by "saturated": I can take sodium chloride, and I can drop it into a beaker; I can dissolve it and dissolve it and dissolve it; at some point, no matter how much more salt I add, it's just going to sink to the bottom, because it's completely saturated.0371

There is going to be so much--there are going to be so many sodium ions and chloride ions floating around in solution that the solution just doesn't have enough capacity to carry any more ions.0384

The salt, the sodium chloride, which is normally soluble, will actually sit at the bottom of the beaker as a solid; this is called a saturated solution.0399

So, for these problems, you will often see--it will often start by saying "a saturated solution of sodium chloride," "a saturated solution of sodium phosphate"; that is what this means.0408

It means that the ions are in equilibrium with their solid; that is what this equation says.0419

I just wanted you to know that you will often hear it as the "saturated solution."0427

That means no more ions are going into solution; it's staying as a solid.0432

OK, now, as it turns out (let me rewrite this Mg): magnesium fluoride is in equilibrium with magnesium ion...solid...plus fluoride ion; as it turns out, the Ksp for this is 6.4x10-9.0436

This is a small number: that means that...so we said that the Ksp is equal to the magnesium ion concentration in moles per liter, times the fluoride ion concentration squared, is equal to 6.4x10-9; that is a very tiny number.0462

You know what that means--that means the concentration of free magnesium and free fluoride in solution is very, very tiny; that is the whole idea.0481

OK, let's do a barium hydroxide: BaOH2, solid--it is going to be in equilibrium with a little bit of barium ion, plus 2 moles of hydroxide ion; and the Ksp for this equals 5.0x10-3.0491

Again, it's a tiny number, which means that there is very little barium and very little hydroxide floating around.0512

Also notice: for every molecule of barium hydroxide that dissociates, 1 ion of barium and 2 ions of hydroxide...the stoichiometry is very, very important; and you will see that in a minute.0518

OK, now we want to talk about the two things that are most important when discussing solubility equilibria.0530

These are the things that we do not want to confuse; and it is the single biggest problem for students when they read their problems; they are in such a hurry, or they are so stressed out, that they actually end up mixing up these two things that I'm going to write down.0537

Solubility: solubility is the amount, in either moles or grams--the amount of solid that dissolves under a given set of conditions.0551

For example, if I said 10.0 grams per 100 milliliters of water, this is an expression of solubility; it is telling me that, if I have 100 milliliters of water, I can dissolve 10 grams of solid; that means 10 grams of a salt will completely dissociate in that solution.0582

Not more than 10 grams: if I put 10.2 grams, 10 grams will dissolve; .2 grams will sit at the bottom as a saturated solution.0607

This is an expression of saturation: it says that I can't squeeze, at a given temperature, more than 10 grams of a solid into solution--into 100 milliliters of solution.0616

It might say it this way: 1.2x10-3 moles per liter of, let's say...let's just pick something at random: potassium nitrate.0627

Now, obviously, potassium nitrate is going to be more soluble than this; but I'm just using this as an example.0640

This means that, if I have solid potassium nitrate, and if I have 1 liter of water, that means I can only dissolve 1.2x10-3 moles of this potassium...0646

You know what, I'm not going to use potassium nitrate; I know it's going to end up sticking in somebody's mind, and it's going to end up confusing you kids, so we don't want to do that--just some random salt.0657

That means 1.2x10-3 moles of this salt will dissolve in 1 liter of water.0667

If I have 2 liters of water, well, it's going to be 2.4x10-3 moles.0672

But, this is an expression of solubility; it tells me an amount of the solid salt that dissolves under a given set of conditions.0677

This can change, depending on the conditions.0686

Now, solubility product is what we just defined; it is a constant--a constant that expresses the relationship between the concentrations of free ion in solution.0689

It is a constant; it doesn't change; that is the whole idea behind Ksp.0749

A certain amount of something might dissolve in solution; a certain different amount might dissolve under a different set of circumstances; however, what doesn't change...0755

Remember when we talked about an equilibrium position, versus the equilibrium constant?0764

Equilibrium positions can change, but the overall equilibrium constant does not; so, the solubility can change--can be different--but the solubility product is a constant; the Ksp expresses the relationship in a given aqueous solution.0768

MgF2, Mg2+ + 2 F-; we said that the Ksp was, what, 6.4x10-9.0786

Different amounts of Mg and F- might be in solution, depending on how the solution was made; but the product of the magnesium ion concentration in solution, times the fluoride ion concentration, squared, always equals the same constant.0802

That is the difference: do not mistake solubility for solubility product, or solubility product constant; that is the single biggest mistake that kids make.0821

Other than that, the problems to deal with this are actually quite simple.0830

We have been doing a ton of these equilibrium problems; we should be well-versed by now.0833

We do Before, Change, After; we do Initial, Change, Equilibrium; nothing is different, except now we just have a different constant to work with, Ksp instead of Keq or Ka or Kb.0838

OK, let's just jump into some examples, and I think it will start to bring everything to light.0849

So, Example 1: Copper (1) bromide has a solubility of 1.9x10-4 moles per liter at 25 degrees Celsius.0856

So, if I have one liter of water at 25 degrees Celsius, I can only dissolve 1.9x10-4 moles of copper bromide in that solution; no more will go in there.0886

Anything more than that, and it will have an equilibrium--it will be saturated.0899

OK, what is its Ksp?0903

Well, let's write out the reaction first; let's write out the equilibrium: Copper bromide (it's copper (1) bromide, so it's CuBr)--when it dissociates, it dissociates into copper +1, plus a bromide -.0911

Let's write its Ksp expression: chemistry--start with a reaction--write the equilibrium expression; it will get you started, and it will keep you on track.0929

Copper +, times Br-; this is solid; these are aqueous; we know that--OK.0943

The solubility of a salt means how much has dissolved in a given volume.0951

Now, how much of a given salt has dissolved in a given volume?0958

This is telling me that 1.9x10-4 moles per liter of copper (1) bromide dissociates; well, for every 1 mole of copper bromide that dissociates, 1 mole of copper ion is produced; 1 mole of bromide ion is produced.0963

What does that mean?--that means that 1.9x10-4 moles per liter, at 25 degrees Celsius--that means, when it dissolves, it produces 1.9x10-4 Molar copper ion and 1.9x10-4 bromide ion.0984

So, the solubility of a salt means how much has dissolved; we know that--in this case, the copper and a bromide--they are the solubility, because for every amount that dissolves, it produces (because the ratio is 1:1, 1:1) that much of the free ion.1004

Again, if 1.9x10-4 moles dissolves in 1 liter, well, that produces 1.9x10-4 moles of free copper ion in that liter.1055

It produces 1.9x10-4 moles of free bromide ion in that liter.1065

Those are our solubilities, moles per liter; we have our concentrations, so we're done.1070

Let's just see what this looks like in an ICE chart.1076

CuBr in equilibrium with Cu+ + Br-; OK, our Initial, our Change, our Equilibrium...1079

So, this is a solid, so it doesn't matter; it doesn't actually show up in the Ksp expression; and because it doesn't show up in the Ksp expression, we don't care about it.1092

Before anything happens, we drop this into water; before it dissociates, there is none of this, there is none of this, and there is a whole bunch of this.1103

How much dissociates?--well (let's do this in red), that much dissociates, so -1.9x10-4.1110

Now, we said it doesn't matter, but it is good to put it here to remind us what is dissolving; and the negative sign means it's dissolving, it's breaking up; that much of this copper bromide is disappearing.1122

+1.9x10-4, +1.9x10-4; this does not matter; we have 1.9x10-4 moles per liter of copper ion floating around; we have 1.9x10-4 moles per liter bromide ion floating around.1135

So now, our Ksp is equal to our copper ion times our bromide ion; it equals 1.9x10-4, times 1.9x10-4, and we get a Ksp equal to 3.61x10-8.1162

We have calculated a Ksp from a solubility; the solubility is the amount of solid that dissolves; well, for every amount of solid that dissolves, the reaction creates a certain amount of one ion and a certain amount of the other ion, depending upon the stoichiometric coefficients.1184

Let's do another example.1202

Notice: we omitted the units; this is moles per liter, moles per liter; we don't do moles squared per liter squared--we just, for the Ksp, ignore the units; we don't write it down.1210

OK, Example 2: all right, the solubility of bismuth (3) iodide is 7.78x10-3 grams per liter.1220

So notice, this time they gave it to us in grams per liter.1252

Calculate its Ksp.1255

Well, let's start with our reaction: bismuth (3) iodide (that is a solid; that is going to be...you know what, let me go ahead and do this in another color here--let me go back to black)...so, we have bismuth iodide, solid; it is in equilibrium with a bismuth 3+ ion; but notice, for every "molecule"--for every unit--of bismuth iodide that dissolved, it produces one mole of bismuth ion, and it produces 3 moles of iodide.1263

Now, our stoichiometric coefficients are going to be different.1306

That is our reaction; let's write our Ksp expression.1310

Our Ksp is going to be bismuth 3+, times iodide -, cubed; yes, that's exactly right; OK.1313

Now, one of the things that you have probably noticed with these equilibrium problems, acid-base problems, is the sheer notational busy-ness--not complexity, but you have charges; you have exponents; you have coefficients; just write slowly and keep track of them all.1326

Otherwise, it will drive you crazy, and it will throw you off.1347

If you are going to make a mistake, you want to make a conceptual mistake; you don't want to make a notational or arithmetic mistake--that is what is important.1351

OK, so now, this is our Ksp expression; we have our reaction; 1 mole of this produces 1 mole of bismuth ion; it produces 3 moles of that ion; but notice, this is in grams per liter.1359

These have to be in moles per liter, so we need to change this solubility to molarity--not a problem; we'll just divide it by the molar mass of the bismuth iodide.1374

Let's go ahead and do that first; so let's do 7.78x10-3 (because again, we are looking for Ksp, so we need to find this, and we need to find that; but it's in moles per liter, but the solubility is in grams).1388

This is in grams per liter, times...well, let's see: we want: 1 mole of the bismuth iodide happens to be 589.7 grams; and when we do this, we get 1.32x10-5 moles per liter; so this is the solubility.1403

Once again, that means 1.32x10-5 moles of bismuth iodide will dissolve in 1 liter of water.1439

OK, well, let's do our chart: BiI3 is in equilibrium with Bi3+ + 3 I-; our Initial; our Change; and our Equilibrium (this Equilibrium one is the one that we are going to put back into our Ksp expression).1446

OK, we don't care about this (no, I'm not going to have any stray lines here, because these problems are confusing enough without these stray lines), so nothing there.1468

Before anything happens, 00: this is: we just drop it into water, before it comes to equilibrium.1479

Well, a certain amount of the bismuth iodide is going to dissolve; that is the solubility; that is this number, right here; so it is -1.32x10-5.1485

Well, for every 1 mole of this, 1 mole of that is created; so it's going to be +1.32x10-5.1500

But, for every mole of this that dissolves, 3 moles of iodide are produced; so, if 1.32x10-5 moles per liter dissolve, this is going to be 3 times 1.32x10-5; what we end up with (this doesn't matter)--we end up with 1.32x10-5 (that's 0 plus that); 0 plus this equals 3.96x10-5--that is the concentration of our bismuth; this is the concentration of our iodide.1510

And now, we go ahead and we put these numbers into this expression, and we get (let me write it again on this page) Ksp is equal to bismuth 3+ ion concentration, times the iodide concentration cubed, equal to 1.32x10-5, times 3.96x10-5, cubed.1553

We end up with 8.20x10-19; that is our Ksp.1587

Now, notice what we did: this 3 here shows up in two places--it shows up in the Ksp as the exponent; it shows up for this value, based on the stoichiometry.1594

For every mole of bismuth iodide that dissolves, 3 moles of iodide go into solution.1610

Therefore, 1.32x10-5 moles per liter of bismuth iodide that dissolve release 3.96x10-5 moles per liter of iodine.1618

When you put this value into this, this "cubed" shows up, now, differently.1629

The stoichiometry gives you this number, but you still have to use this number, because it's part of the Ksp expression; don't think that it is one or the other--it is both.1635

Be very, very careful when you are doing these solubility products.1644

OK, let's do Example #3: OK, so in this one, we're going to go in reverse--we gave you the solubility and you found the Ksp; now we're going to give you the Ksp; let's see what the solubility is.1647

The Ksp of magnesium hydroxide is 8.9x10-12 at 25 degrees Celsius; calculate its solubility.1671

OK, so "calculate its solubility"--they are saying, at 25 degrees Celsius, given a certain moles per liter, given 1 liter of water, how much magnesium hydroxide will actually dissolve?1696

That is what they are asking; OK.1710

Simple enough: well, let's write the equation.1712

Magnesium hydroxide (this is chemistry; you have to have an equation) is in equilibrium, when it dissociates, with one magnesium ion, plus two hydroxide ions.1716

OK, well, we have the Ksp; we want the solubility.1729

Well, the solubility is the amount of this that actually dissolves; OK, so let's do our ICE chart--here is what an ICE chart looks like now.1734

It doesn't matter how much we start with: 0, 0--this is before anything happens.1742

As it comes to equilibrium, as some of this dissolves, a certain amount is going to dissolve; that is -x.1750

Well, for every x that dissolves, x of this shows up; so this is going to be +x.1756

For every x of this that dissolves, it produces 2 hydroxides; so this is 2x.1762

Well, this doesn't matter; this is x; this is 2x; these are the two things that show up in the Ksp expression.1771

So, let's go ahead and write: Ksp is equal to...well, we said it was 8.9x10-12; that is equal to the magnesium ion concentration, times the hydroxide ion concentration squared; it's equal to x, times 2x squared (be very careful how you do this), equals x times 4x squared, equals 4x cubed.1780

We have 4x3=8.9x10-12; when we divide by 4 and take the cube root, we end up with 1.3x10-4.1818

This is our solubility, right here.1834

You remember: it's the -x; 1.3x10-4: solubility--we include the unit, so this is moles per liter--which means that, in 1 liter of water, 1.3x10-4 moles of magnesium hydroxide will dissolve.1839

There you go; OK, now let's talk about something called relative solubilities.1863

This is going to be more of a qualitative thing, so we're not really going to be doing any math here.1872

Relative solubilities: if you have a group of salts, you can use Ksp values, of course, to decide the order of solubilities--only for salts that produce the same number of ions when they dissociate.1878

OK, so if you have a group of salts, you can use the Ksp values; you can compare them to decide the order of solubilities, only for salts that produce the same number of ions upon dissociation.1941

OK, so let's just do a quick example here.1953

If I have silver iodide, well, the Ksp of silver iodide is 1.5x10-16; if I have barium sulfate, the Ksp is equal to 1.5x10-9.1956

And, if I have nickel (2) carbonate (or nickel carbonate--I guess nickel is always +2 for the most part), the Ksp is equal to 1.4x10-7.1973

Well, each of these produces...when Ag dissociates, it's 1 mole of Ag ion and 1 mole of iodide ion; here, barium sulfate--1 mole of barium ion and 1 mole of sulfate ion (sulfate ion--the S, the O4--they don't dissociate; SO4 is an ion itself); nickel carbonate--when it dissociates...1 mole of nickel; 1 mole of carbonate.1986

So, in each case, each is producing 2 moles of ion total.2011

Because we can do that, we can compare these Ksps.2016

Well, the higher the Ksp...that means the equilibrium is farther to the right; if it's farther to the right, that means the concentration of the individual ions is higher.2019

Well, a higher concentration of free ion means that more of a salt has dissolved.2029

Therefore, that has a bigger solubility than this; this has a bigger solubility than that.2035

The order of solubilities, in terms of greatest to least, is: nickel carbonate is more soluble than barium sulfate, is more soluble than silver iodide.2041

You can do this, because they produce the same number of ions.2053

Because they produce the same number of ions, you can just compare Ksps; the one with the biggest Ksp is the most soluble--it dissolves the most.2056

Now, let's try...what about copper (2) sulfide (which has a Ksp equal to 8.5x10-45); how about barium phosphate (it has a Ksp of 6x10-39--very, very little dissolution); and iron (3) hydroxide (notice, it's FeOH3, not FeOH2; OK, Ksp is equal to 4.0x10-38).2066

When this dissolves, it produces 1 mole of copper, 1 mole of sulfide; this produces 3 moles of barium, 2 moles of phosphate; this produces 1 mole of iron, 3 moles of hydroxide.2110

You can't do it--you cannot compare; you cannot say that this -38 is bigger than this -39, is bigger than this 10-45.2125

Numerically, it's bigger; but, because these produce different numbers of ions upon dissolution, you cannot do it.2135

If you have a group of ions that produce the same number of ions upon dissolution, then you can use your Ksp values to qualitatively decide which is the most soluble.2143

You will see questions like this, absolutely, on the AP exam.2155

OK, so this was our general introduction to solubility equilibria; we are going to spend, actually, 2 or 3 more lessons on solubility equilibria, because they do tend to get a little bit more involved.2160

Again, it's a lot like buffers and titrations; I definitely want you to understand this, because, if you can understand these equilibria, then most of the rest of chemistry is actually an absolute breeze.2171

So, we'll see you next time; thank you for joining us here at Educator.com; goodbye.2182

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