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Raffi Hovasapian

Raffi Hovasapian

Complex Ions & Solubility

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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Complex Ions & Solubility

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
  • Complex Ions and Solubility 0:23
    • Recall: Classical Qualitative Analysis
    • Example 1
    • Example 2
    • Dissolving a Water-Insoluble Ionic Compound: Method 1
    • Dissolving a Water-Insoluble Ionic Compound: Method 2

Transcription: Complex Ions & Solubility

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

Today, we are going to round out our discussion of aqueous equilibria, and we are going to finish off by discussing complex ions and solubility.0004

We talked about solubility, and we talked about complex ions; now, we are going to talk about how we can mix the two.0013

Let's go ahead and move forward.0021

OK, so let's recall, from a couple of lessons ago: we did something called this classical qualitative analysis.0029

Recall: the classical qualitative analysis--remember, we had this mixture of ions, all of these metal ions, and we added different reagents, and we used this process of selective precipitation to precipitate out certain groups.0042

Group 1: the insoluble chlorides; Group 2 was those that are insoluble in acidic solution; those that are insoluble in basic solution; and then the carbonates; and then the alkali metals.0070

We are sort of taking this mixture, and we are sort of separating it out, based on solubility.0083

OK, well, now we are going to combine this: so we had this...I won't write out all of the ions again...mixture of ions.0088

The first thing that we did was: we added dilute hydrochloric acid; and when we did that, we precipitated out the insoluble chlorides.0107

There were three of them, if you remember: there was lead (2) chloride; there was silver chloride; and there was mercury (1) chloride, which is Hg2Cl2 (and the reason is because mercury (1) tends to aggregate as 2 mercury atoms together...but it's still mercury 1+).0123

And then, of course, we had the others; OK.0142

Well, here is the question: now I have these things sort of sitting at the bottom of the solution; so I filter it out, and now I have this solid clump of lead chloride, silver chloride, and mercury chloride.0147

Well, now, the next step: How do I separate these into their individual ions (lead 2, silver 1, mercury 1)--how do I do that?0163

OK, well, from a practical standpoint, we need to basically put all of these three back into solution again; so we throw away these others; we take this; and we have to find a way to dissolve these back again, so that we can run a different set of reactions to separate these individually.0175

We separated them from the group, but now we need to separate them; how do we do that?0196

OK, so again, we need to find a way to re-dissolve these so that we can do some chemistry to separate these three out.0201

Well, as it turns out (let me write this down...well, let me write out the whole thing): OK, now the question becomes: How do we separate these three--how do we separate these three?0213

OK, the answer is: we have to re-dissolve them and run some other chemistry.0233

OK, well, the "run some other chemistry" part is probably going to be taken care of when you actually do your lab work for this particular General Chemistry course, or your AP Chemistry course.0254

What we are going to concentrate on is: How do we re-dissolve them--how do we take these and put them back into solution (or at least one of these species back into solution) to separate them out?)0267

OK, as it turns out, forming complex ions with insoluble salts (that is what these are--they are insoluble salts; their Ksps are really tiny--silver chloride, lead chloride...they don't dissolve very much; they just sit there at the bottom of the beaker) makes them soluble.0279

We will concentrate on AgCl; so I'm just going to pick one of these to show you how to re-dissolve that.0329

If you have this precipitate, that is what is going to happen; anytime you put the chloride--when you have silver and chloride together in solution, they are going to bind, and they are going to drop to the bottom of the flask as a precipitate.0346

Well, in order to do some more chemistry with that, I need to find a way to dissolve silver chloride.0357

Well, silver chloride doesn't dissolve in water; how can I, therefore, dissolve it?0363

Well, here is how you do it: let's start with the actual silver chloride equilibrium.0368

AgCl, solid, is in equilibrium with Ag+ + Cl-; OK, if you remember (well, you don't have to remember), the Ksp for this is 1.6x10-10; so clearly, not very soluble at all.0377

All right, well, watch what happens: so I have this solution; I have a whole bunch of silver chloride, solid, down at the bottom; so this is AgCl, solid; and there are a few ions of silver and a few ions of chloride; there you go.0396

Now, here is what we do to actually dissolve this stuff that is sitting at the bottom.0417

We need to find a way to dissolve it to run more chemistry on it.0423

We add NH3; well, in the last lesson, you saw what happens when you have silver solution reacting with NH3; here is what happens.0428

The free silver, the little bit of free silver that is actually floating around in solution, the few ions--well, they actually react with the ammonia that you drop into the solution according to the complex ion equilibria.0438

They form the monoamine complex ion; that K1, remember, was 2.1x103.0456

Now, this one goes and reacts with another molecule of ammonia, and it forms the diamine complex ion.0465

So, when this happens (and let me just go ahead and write the K), let's see what is going on here.0480

I have some solid sodium chloride; it is in equilibrium with some silver and some chloride, based on the Ksp, so this is what is happening.0489

Mostly it's this, but a little bit of this; but the moment I drop ammonia into this solution, the ammonia starts binding to the silver--the free silver that is floating around--and it binds it up as this diamine complex ion.0497

Well, when that happens, that is what is happening; you are basically tying up this silver concentration as something else.0511

Well, when you do that, you have reduced the silver ion--the free silver ion concentration.0520

Well, if you have reduced the free silver ion concentration, Le Chatelier's Principle says this reaction will adjust itself to re-increase the silver ion concentration.0525

The only way the silver ion concentration is going to increase again is if this dissociates to release more free silver ion.0535

Yes, it's going to also release chloride, but chloride doesn't matter in this case.0543

That is what happens; when it dissociates, that is dissolving; so, every time it dissolves to produce more silver ion, up to the value of the Ksp, and then we add ammonia--that ammonia is going to react with the silver ion and pull it out of solution again.0547

If that pulls it out of solution, it is going to create an empty space for silver ion; more of the AgCl will dissolve; more of it will dissolve; more of it will dissolve, until all of it dissolves.0563

Because, what you have done is: you have locked up the silver as diamine silver complex ion.0574

It is kind of like the same thing that we did when we added acid to an insoluble hydroxide.0580

The acid reacts with the hydroxide, forming water, pulling the reaction forward, causing the insoluble salt to dissolve.0587

It is the same thing here: silver chloride is insoluble in water; therefore, to that water, I add ammonia.0594

The ammonia binds with the silver, pulling this reaction forward, dissolving the silver ion.0600

Now, I have silver ion in solution; yes, it's bound up, but it is still silver ion.0606

And now, I can do other chemistry to it or separate it out from other solutions; that is the whole idea.0612

OK, so let's actually write these reactions in order...oh, let me write down what is going on here; it's very, very important.0618

Ag+ reacts with NH3 to form a complex ion (and as we know, ions in solution are soluble); the Ag+ concentration is reduced (that's nice; it would be good if I could spell); as compensation for that, AgCl responds (oh, I'm having a hard time spelling tonight) by dissociating to release more silver ion (in other words, it dissolves), until the Ksp value of 1.6x10-10 is reached again.0628

Because that is what the Ksp is; the Ksp is the concentration of that, times the concentration of that; once that is reached, no more of this will dissolve.0705

But, if I add ammonia to it, it reacts with the silver, depleting the silver ion concentration; now, the Ksp doesn't match--there is chloride ion, but there is virtually no silver ion.0714

More of this dissolves to form silver ion until it reaches the Ksp value.0724

That is what is going on here.0729

OK, now let's do the chemistry; so we have AgCl in equilibrium with Ag+ + Cl-.0730

The Ksp for that, we said, is 1.6x10-10: very, very insoluble.0744

We have Ag+ + NH3 going to form the monoamine complex ion.0751

Oh, in case you are wondering, ammonia, when it acts as a ligand, is called "amine"; so that is why I'm saying "amine" instead of "ammonia"; I should have mentioned that earlier.0765

K1 equals 2.1x103, and then the monoamine reacts with another molecule of ammonia to form the diamine complex ion.0779

Symbols, symbols, symbols...I tell you...2...no, this is 8.2, times...oh, -3; it's positive 3; I'm telling you, there are so many places to make mistakes when you do this chemistry, because there is so much symbolism.0800

That is the thing: concepts are not difficult--it's the symbolism that gets in the way, and that is what is often frustrating about science.0819

It's not too difficult, but just the symbolism makes you crazy!0825

OK, well, we can form...so this is what is going on: there is equilibrium among all of this stuff.0829

Well, we have some things on the left-hand side and the right-hand side that are actually the same; so let's add these equations for a net reaction--what is happening in solution when you take silver chloride, solid, and drop in ammonia?0839

All of these reactions are taking place; what is the net reaction?0852

Well (let's do this in blue), that is on the right; that is on the left; they cancel; that cancels with that; and we are left with the following.0856

Solid silver chloride will react with 2 moles of ammonia to form the diamine complex ion, plus chloride ion.0872

The Knet (remember what we said: we multiply these) is equal to Ksp times K1 times K2, and it ends up being 2.8x10-3; so now, let's look at this.0893

Let's take a look at it and think about it.0910

We have this reaction taking place: solid silver chloride; drop in some ammonia; the ammonia reacts with silver and forms the monoamine; the monoamine reacts with another ammonia and forms the diamine.0915

Mostly this is going to be there; I can add these equations together for a total reaction, a net reaction; that means one mole of silver chloride, solid, will react with 2 moles of ammonia to form 1 mole of the diamine complex ion and a free chloride ion.0926

The total equilibrium constant for this reaction--these over this; this doesn't count--is 2.8x10-3.0948

2.8x10-3 is small, but it's not super small; so it's large enough, so this definitely is a classical equilibrium situation; there is going to be a fair amount of this, this, and this floating around in solution (and perhaps even some of this, depending on how much was actually there).0957

So, given that, now we can do our example.0976

Example: Calculate the solubility of AgCl in 11 Molar ammonia solution.0986

There you go: remember, solubility is the amount of a solid that dissolves.1005

We calculated the solubility of silver chloride earlier; it just has to do with the Ksp value--writing out the Ksp, doing your ICE chart.1009

Well now, we are not saying "Calculate the solubility of AgCl in pure water"; we are saying, "Calculate the solubility of AgCl in 11 Molar NH3."1016

When that happens, we know what happens when it reacts with NH3.1026

This happens; we add the equations--this is our net equation; this is our net equilibrium constant; that is what we use in our ICE chart.1031

That is the whole idea: the chemistry first, then the math.1040

OK, so now, let me go ahead and move to another page here, and write...yes, that is fine: so, we have AgCl + 2 NH3 in equilibrium with Ag(NH3)2+ + chloride ion.1045

Our initial concentration of AgCl: well, this is a solid, so we don't care how much is actually there--it doesn't affect the equilibrium, so we don't really care about it.1081

How much ammonia do we have--what is the concentration of ammonia?--well, the ammonia concentration, before anything happens, is 11 Molar; that was the thing that was given.1090

There is none of this formed yet, and there is none of this formed yet.1100

This is before anything takes place.1105

The change: what kind of a change takes place?--well, a certain amount of the silver chloride is going to dissolve; this is what we are looking for--this is the x, solubility, how much dissolves.1107

Well, for every mole of this that dissolves, 2 moles of NH3 is used up.1123

So, it's -2x; that is it; this 2 and this 2--that is where they come from.1131

Well, for every mole of this used up, one mole of the complex ion shows up, and one mole of that shows up.1137

Our equilibrium is that; that doesn't matter, because it's a solid; 11.0-2x; x; and x.1146

The Knet is the diamine concentration, times the chloride ion concentration, divided by the NH3 concentration squared (law of mass action, right?).1164

This is not a liquid; this is solution--this is aqueous; OK--this is a mixture of water--all right.1185

Now, well, this is x; this is x; this is 11.0-2x; we know what the Knet is--we had it already; that is 2.8x10-3, is equal to x, times x, divided by 11.0-2x squared, equals x squared over 11.0-2x squared.1192

Now, I know what you are thinking: can we actually go ahead and ignore this and use the 5% rule to decide that it is valid--so, can we make the math easier?1227

Yes, you can if you want to; but notice something here: you actually don't have to simplify here, because this is squared and this is squared.1238

This is equal to x over 11.0-2x squared; so it's the whole thing squared; so because the top is squared and the bottom is squared, I can just take the square root of both sides.1245

When I do that, I get (the square root of this, the square root of that)...I end up with 0.05291=x/11.0-2x (and let's not have any stray lines here).1266

When we multiply through, we get 0.5821-0.10582x=x; and we end up with 0.5821=1.10582x; and we get x is equal to 0.53 moles per liter.1292

There we go: we just calculated the solubility.1324

In 11 Molar ammonia solution, silver chloride: .53 moles of silver chloride will dissolve in a liter of that solution.1329

Compare this with the solubility of AgCl in pure H2O (in pure water): 1.3x10-5 moles per liter--very big difference.1343

1.3x10-5 is virtually nothing; .53--that is very significant solubility--that is very, very soluble.1372

There we go: by simply taking an insoluble salt (like silver chloride or any other insoluble salt), if that insoluble salt--if the cation actually ends up forming a complex ion, we can use the fact that it forms a complex ion to dissolve that insoluble salt.1382

If we need to run some chemistry on it, in this particular case, we just need to add some ammonia to the solution, and sure enough, what is insoluble in water will dissolve in the ammonia, by an application of Le Chatelier's Principle, by the formation of a complex ion--very, very important.1402

OK, so now let's take sort of a global view of what it is that we have done.1419

Let me go back to blue.1424

We have now seen two ways to dissolve a water-insoluble ionic compound.1426

One, which we discussed a few lessons ago, is to acidify the solution.1458

If we acidify the solution (in other words, if we add hydrogen ion to the solution), the anion of the solid reacts with the added H+ (oops, not going to have that), thus pulling the dissociation forward (forward--to the right--dissolving).1468

I'll just write "forward."1518

We have seen two ways to dissolve a water-insoluble ionic compound; so when we have something that is actually not soluble in water--it doesn't dissolve--we need to find a way to dissolve it.1523

Well, one of the ways that we have is to add acid to that solution; when we add acid to the solution, the acid (in other words, the H+) reacts with the anion of the solid; and, by Le Chatelier's Principle, it pulls the dissociation forward until everything is dissolved.1533

That is the whole idea.1549

A couple of examples: the examples that we did of that were like magnesium hydroxide--remember, we had magnesium hydroxide, solid.1552

Well, the equilibrium, the Ksp for that...it's true that it is in equilibrium with magnesium and 2 hydroxide ions, but take a look at the Ksp for this: it is (oh, wait, do I even have the Ksp for this?)...oh, it looks like I forgot to write the Ksp, but if I'm not mistaken, it was somewhere on the order of 10-12.1561

It's very, very insoluble; however, when we added H+ to this solution, here is what happens.1584

That OH- ends up reacting with that H+ to form water.1594

When the water is formed, this is depleted; this is now reacting with that; well, since this is depleted, Le Chatelier's Principle wants to produce more hydroxide, which means...the only way to produce more hydroxide is: this has to dissolve to release free hydroxide.1602

More will dissolve; more will dissolve; more will dissolve until the whole thing is dissolved.1619

We have bypassed that, the fact that it is insoluble, by adding acid to the solution.1624

Another example would be calcium carbonate: calcium carbonate...1631

I could be wrong about the 10-12; that could be 10-16...but it doesn't matter; it's insoluble--that is the whole idea.1637

...goes to calcium 2+, plus carbonate 2-; well, the anion of the salt (this is solid)--small Ksp--very small--not very soluble.1645

Well, this reacts with the acid; when it reacts with the acid, it forms the bicarbonate ion; that means it is now sequestered as bicarbonate ion.1664

Well, now there is an empty spot there; in order to fill up that empty spot and reach the Ksp value again, more of this has to dissolve.1679

More dissolves; more dissolves until it is all dissolved.1688

OK, #2: The other way that we have to dissolve an insoluble ionic solid is to add a solution of a ligand that forms a stable complex ion with (I'm always putting an 'e' at the end of my "with") the cation (this time, with the positive--with the cation) of the ionic solid.1693

Our example (well, almost our example; we used silver chloride, but I'm going to write it out as silver bromide): AgBr, solid, is in equilibrium with Ag+ + Br- (do I have another page?--yes, I do; good).1741

Well, the Ag+ reacts with 2 molecules of ammonia to form that diamine complex ion.1761

In the process of forming that diamine complex ion, now the silver is bound up as this complex ion.1776

Well, because now there is no silver ion in solution anymore, the Ksp value is not satisfied for this.1783

So, Le Chatelier's Principle: when you deplete this, the system will react by trying to raise that concentration again.1791

In order to raise the concentration of free silver ion, this has to dissolve; it has to dissociate to produce, so more dissolves.1798

You add more ammonia--it binds the silver and pulls the silver out of solution as a complex ion; more dissolves until eventually, you have dissolved all of your silver bromide.1806

And that is the whole idea; so again, when we are dealing with insoluble salts, yes, it is true--the problems that we did--we dealt with them directly; but in certain circumstances, we do have to find a way to dissolve them.1817

And now, we have two ways of dissolving them: we can acidify the solution by adding hydrogen ion (in which case, the hydrogen ion--if the hydrogen ion reacts with the anion of the insoluble salt, that acidified solution will actually dissolve that salt, where it is not otherwise soluble in water).1832

Or, if the cation of the insoluble salt forms a stable complex ion with a given ligand (it could be any ligand), if we add a solution of that ligand, by doing that, we can make that insoluble salt all of a sudden soluble, so that we can do further chemistry with it.1854

OK, that brings us to the end, actually, of our discussion of aqueous equilibria.1876

With that, I will thank you again for joining us here at Educator.com.1882

We will see you next time for a further discussion of chemistry; take good care; goodbye.1887

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