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Raffi Hovasapian

Raffi Hovasapian

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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 (11)

1 answer

Last reply by: Magic Fu
Tue Feb 14, 2017 10:43 PM

Post by Magic Fu on February 14, 2017

Hi, Professor Hovasapian.At 17:47 you said that HAr are the only source of hydrogen ions, what about the hydrogen ions of water?

0 answers

Post by Sun HaoHui on February 8, 2017

Hi, Professor Hovasapian
1. Why do you start off with F- in the beaker on the first slide you did?
2. The H+ that start off in the beaker comes from water?
3. Why would the pH drop if you add more acid to the solution?
4. (17:23)I'm really confused about how to write the major species.
Do you always break apart salt when writing major species?

3 answers

Last reply by: Professor Hovasapian
Fri Feb 26, 2016 3:48 AM

Post by chitra banarjee on April 29, 2015

How exactly do you determine what the major species are? In other words, what classifies a species are major?
Thanks!

1 answer

Last reply by: Professor Hovasapian
Sun Apr 21, 2013 8:19 PM

Post by carlos bara on April 20, 2013

Professor Hovasapian, I want to thank you very much for teaching all this material in such a wonderful way. Thank you for your time and dedication to the students. Also, I was wondering if there is any way you can upload lectures on Organic chemistry, as organic chemistry is a very complex subject and having someone like you teach it, would certainly make things less complicated. Thank you once again!

1 answer

Last reply by: Professor Hovasapian
Wed Mar 20, 2013 3:23 AM

Post by Kathryn Cosgrove on March 20, 2013

Where did you get the Ka= 1.8x10^-5 at the time 26:53 of the video?

Related Articles:

Buffer

  • When you add acid or base to a buffer, do the stoichiometry of the neutralization first ( in moles ), then do the Equilibrium problem ( in molarities ).
  • Using the Henderson Haselbalch Equation with buffer solutions can save time and effort.

Buffer

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
  • Buffers 1:20
    • Buffer Solution
    • Adding Base
    • Adding Acid
    • Example 1: Question
    • Example 1: Recall
    • Example 1: Major Species Upon Addition of NaOH
    • Example 1: Equilibrium, ICE Chart, and Final Calculation
    • Example 1: Comparison

Transcription: Buffer

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

Today, we are going to continue our discussion of buffer solutions; and after this, we are going to go even further with buffer solutions.0004

Last time, we introduced what a buffer solution was; we said that it is a solution made up of a weak acid, plus the salt of its conjugate base; or, it's a weak base plus the salt of its conjugate acid.0011

It is designed for one thing only: it is designed to resist changes in pH once you add acid or base to that buffer solution.0026

We make a buffer solution at a specified pH that we want, and we design it specifically so that, if you add acid or base to it, the pH doesn't change; or, if it changes, it changes so slightly that it is negligible; that is what a buffer is.0036

Today, we are going to talk about how buffer solutions actually work.0051

In the last lesson, you remember, we actually created a buffer solution--two of them, in fact: we did one with hydrofluoric acid and sodium fluoride, and then the last one that we did was: we did acetic acid and sodium acetate.0055

We ended up with a pH of about 4.66; so here, we are going to talk about how buffers actually do what they do--in other words, resist changes in pH.0068

Let's get started; now, let me just write down something really quickly--what it is that we actually said about buffers.0080

We said that a buffer solution is a solution at a prescribed pH that resists changes to that pH upon addition of H+ or OH-.0091

Again, when we add acid, we are adding H+; when we add base, we are adding OH-.0138

That is what we mean by adding acid or base, ultimately.0142

This is the acid; this is the base; that is what is important--the other ions don't matter.0145

OK, how does this happen?--that is the question: how--how does this happen?0154

OK, so let's take a picture of a buffer solution; we have a choice--we can either do acid, salt of conjugate base; or weak base, salt of conjugate acid.0161

Let's stick with what we have done, which is a weak acid, plus the salt of its conjugate base.0175

Hydrofluoric acid: so, we have some HF floating around in solution, some HF floating around in solution, some HF floating around in solution.0180

We have the salt of its conjugate base; so let's ignore the cation--let's just deal with the fluoride ion.0189

We have some F- floating around; we have F- floating around; and we have F- floating around.0196

Now, it's true: we also have a little bit of H+ floating around--I'm going to put a circle around it; H+ floating around; some H+ floating around.0204

Now, here is the idea: how does something resist a change in pH?--well, how does pH change?0213

Hydrogen ion concentration changes--it changes one of two ways: it changes by either going up or going down.0219

If I add acid to a solution, I have added more hydrogen ion to the solution, so the hydrogen ion concentration goes up; the pH drops.0226

If I add base to the solution (I mean any normal solution; I'm not talking about a buffer right now--any normal solution), well, the base ends up sort of...the hydroxide ion concentration rises; the pOH drops; the pH actually rises.0235

That is what is going on; so, what we want to do is...resisting changes upon addition of acid or base means, if I add acid, how can I eat up that acid so it doesn't float around freely in solution and change the molarity of the hydrogen ion concentration?0254

In other words, how can I sequester it--how can I bind it--how can I lock it up, so it isn't just floating around freely, changing the concentration of the hydrogen ion?0275

If I add base, how do I bind it--how do I lock it up--how do I pull it, get rid of it, so it doesn't actually end up reacting with the hydrogen ion concentration and dropping the hydrogen ion concentration, and raising the pH?0284

That is the whole idea of a buffer; how does it do that?0299

Here is how it does that: if I add base (which is OH-, right?): I add base; here is the reaction that takes place.0302

Let's say, all of a sudden, I drop in some base (which is OH-); well, the OH- is a strong base; it does one thing--it's going to seek out hydrogen ions.0316

The base that you add is going to seek out hydrogen ions; well, the only source of hydrogen ion (actually, let me erase these, because actually will just get in the way and confuse you--what is important in a buffer solution is the acid and its anion) in a buffer solution is right there--is the Hs--these Hs right here.0331

When I drop in OH here, the OH is going to pull off these Hs, and here is the reaction that is going to take place...+HF; it's going to form HOH, which is water (acid-base neutralization), plus F-.0365

What it is going to do is: it is going to...what did we say?--we said, if we add base, we need a way to bind that base so it doesn't affect the hydrogen ion concentration.0382

Well, here is your binding right here: if you add base to a buffer solution, it reacts with the hydrogen fluoride; the hydrogen fluoride is there for that reason--it is there to react with added base...production of F- ion and water.0395

Water is neutral; F- doesn't really affect anything all that much in this solution; so, what we have done is: we have sequestered the hydroxide by making it react with the hydrogen on the hydrofluoric acid, to produce water and F- ion.0409

So now, it isn't floating around freely, affecting the hydrogen ion concentration.0426

Now, if we add acid (which means adding H+), well, H+...here is what is going to happen: it is going to react with the F- now.0434

So, if I add acid, now the F- is there to react with it, to sequester it--to bind it so that it is not floating around freely.0462

Because F- is a strong conjugate base, HF is a weak acid; it is going to be mostly in that direction, not mostly in that direction; that is what the equilibrium says.0476

F- reacts with any excess acid that we add, on top of what the pH is, to produce HF; it binds the hydrogen ion, so that the hydrogen ion is not floating around freely, contributing to the concentration of free hydrogen ion.0487

Acids...the measure of the acidity of a solution is based on the concentration of hydrogen ion.0506

A buffer solution has free base floating around the F- (that is the conjugate base of the weak acid), to react with any extra acid that I might add--to bind it, so that it isn't floating around freely.0513

It also contains the weak acid itself, to react with any base that I might add, so that the base doesn't float around freely.0526

That is the whole idea; that is all that a buffer solution does--it binds any base or acid that I throw in there, and keeps it from doing any damage.0535

It keeps it from changing anything; it's as if...remember those H+s that I initially erased?--it is so that these H+s' concentration stays reasonably the same.0543

If I add H+, it would go up if these F-s were not there.0556

Because the F-s are there, they bind the added H+ so they are not floating around freely.0560

This is constant.0565

OH-, if I add that...these are there to react with the OH- to bind it so that it doesn't pull these out of solution and reduce the hydrogen ion concentration.0567

That is what a buffer solution does, and this is how it does it.0578

OK, let's do an example.0583

OK, now, I have to tell you something about these examples that I am going to do: I'm going to do 2 examples; they are going to be reasonably detailed, and they are going to be very, very, very important.0588

If you understand (both of the examples are the same--just different species; I just thought I would do it twice, so that you get a sense of what is going on) one or both of these examples, make sure you understand them completely.0598

Take your time with these: you will understand the nature of buffer solutions completely.0612

I promise you, everything that you need to know about buffers is contained in these examples--both technically and conceptually.0616

OK, Example 1: We want you to calculate the change in pH when 0.0100 moles of solid NaOH is added to 1.0 liters of the buffer solution from the last example of the last lesson (and don't worry, I am going to actually review what the last lesson was).0623

Then, we want you to compare this to the change in pH that occurs when 0.0100 moles of NaOH is added to 1.0 liters of pure water.0690

We want you to calculate the change that occurs--the change in pH--when .0100 moles of solid sodium hydroxide is added to 1 liter of the buffer solution from the last example of the last lesson.0729

Let's review what that was: Recall, we had 0.55 molarity acetic acid and 0.45 Molar sodium acetate.0741

This was our buffer solution.0758

Remember that?--we ended up with a hydrogen ion concentration that was 2.2x10-5, and the pH was 4.66.0767

That is our prescribed pH; we have a buffer solution; our pH is 4.66; now, to this buffered solution, we are going to add .0100 mol of sodium hydroxide, and we want to see what the final pH is.0787

That is what is going on here; OK.0799

Now, let's just take a look at where we are starting.0802

Recall: well, let's see--so we need...our HAc concentration, if you remember right, is going to be the initial 0.55; that was minus x; remember, we started with .55 Molar HAc; some of that dissolved--that was the x.0814

Well, remember, x was really, really tiny; so, for all practical purposes, this is x right here.0838

0.55, minus the 2.2x10-5: well, that equals 0.55 Molar; this is so small that it actually doesn't have an effect on the initial concentration.0846

The concentration of HAc floating around in solution is that.0863

The concentration of Ac-...same thing: it was 0.45+x=0.45+2.2x10-5, and again, it's so small that this is essentially 0.45.0870

This is the concentration of acetic acid and acetate ion floating around in the buffer solution, before we add the sodium hydroxide.0891

Now, this is where we stand as a buffer before the addition (so let me write that down--this is where we stand as a buffer...)--before OH- is added.0900

And again, we are adding sodium hydroxide, but the sodium can be ignored; it's the OH- that you are actually adding; you are adding base.0922

OK, buffer problems consist of two parts: the first part--you have to do the stoichiometry problem; and the second--you do the equilibrium problem.0928

In solution, these happen simultaneously; but in order to do the mathematics, we have to treat them as if they happen separately.0954

OK, so let's see what is going on.0964

What is the first thing that we do?--major species.0971

Major species (let me actually...OK) upon addition of sodium hydroxide: OK, so now that I have added the sodium hydroxide, here is what I have floating around in solution.0979

I have my acetic acid, HAc; I have my acetate ion, Ac-; I have water; I have put in sodium hydroxide, so I have (sodium hydroxide is a strong base--fully dissociates into sodium ion and hydroxide ion, so) sodium ion floating around; and I have hydroxide ion floating around.1013

OK, let's take a look at what happens here.1038

I have HAc, Ac-, H2O, Na, and OH-.1041

What did we say about adding hydroxide?--when you add a base, it is going to react; the first thing that it is going to do is react completely with any source of hydrogen ions that are available to it.1045

OK, the only source of hydrogen ion available to it is that, so these two will dominate what happens immediately in the reaction, before the system actually has a chance to come to equilibrium.1058

This is the stoichiometry part of the problem.1070

It will always be like this: when you add the base or the acid, you do the stoichiometry (this reaction); then, you do the equilibrium.1073

Let's go ahead: so, OH- is a very strong base and will react with the only source (not with the only source; there are a couple of sources, but) the primary source--with the best source of H+, which is HAc.1081

We said that is how a buffer works; if it is base you are adding, it's going to react with this; if it's acid you are adding, it's going to react with that; that is the whole idea behind a buffer.1119

Here is the reaction that takes place: write the reaction--this is chemistry; you need the reactions in order to see what is going on.1127

OH- will react with HAc to produce Ac- + HOH.1134

OK, OH- takes this H, releasing Ac-; HOH--that is water.1148

This is the reaction that takes place; now, we have to do the stoichiometry--this reaction takes place, and then what happens after the reaction takes place?1156

Well, we are going to have some more Ac- floating around; we are going to have a little less HAc floating around; then, we can do the equilibrium of Ac and HAc.1168

So again, hydroxide...this is going to react with that before anything else happens.1179

So, let's do this: I'm going to rewrite it again, a little bit bigger: I'm going to write OH- (a little more room), HAc (and you will see why in just a minute), Ac-, + HOH.1185

OK, so we are going to do something similar to an ICE chart, except I like to use Before, Change, and After.1203

The reason I do this is: I like to do Initial, Change, and Equilibrium when I'm dealing with equilibrium situations.1211

Before, Change, and After is used (I like to use it) for the stoichiometry part of the problem, before anything happens, after this reaction has taken place, but before any equilibrium has been established.1219

That is why I don't like to use ICE, because this isn't really an equilibrium situation.1236

This is a before-and-after situation.1240

So, before anything happens, I have dropped in 0.0100 moles of OH-.1242

Now notice, this is a stoichiometry problem (let me actually write this here: this is stoich...)--this is the stoichiometry part of the buffer problem.1252

The stoichiometry: in stoichiometry, we work in moles; in equilibrium, we work in concentrations (moles per liter)--it's very, very important that you recognize that, and you will see why in a minute.1264

In stoichiometry, we are working in moles; notice: we are not working in concentrations just yet--that is for the equilibrium part.1278

So, HAc; well, we said that we have 1 liter of this solution, this buffer solution; and that contains .55 moles per liter of the HAc.1285

Well, 1.0 liter times 0.55 moles per liter gives me 0.55 moles; it's one of the conveniences of using one liter of a buffered solution--because the molarity actually becomes the moles.1303

But, it is important--again, with stoichiometry, we work in moles; with equilibrium, we work in molarity.1320

The Ac- concentration: well, we just said in the previous slide that the Ac concentration is 0.45 moles per liter.1328

We have one liter of this solution; therefore, its concentration is 1.0 liters, times 0.45 moles per liter; that means there are 0.45 moles of Ac floating around.1338

Water doesn't matter.1354

The change: all of the OH- that I have added is going to react with any HAc that is available; .010 moles reacts with .55 moles of this; this is going to vanish.1357

That is what this reaction says: this OH- is being sequestered--it's being bound up in the form of water.1378

It is going to vanish: 0.0100; this is also going to disappear--this is 1:1.1385

For every one mole of hydroxide, one mole of HAc reacts; therefore, it's going to be -0.0100.1396

This--for every one mole of hydroxide added, one mole of Ac- shows up; therefore, this is going to be +0.0100; water doesn't matter.1405

After this reaction has taken place, we have no hydroxide left; it has been sequestered--it has been bound by this.1420

We have 0.55 minus 0.01; 0.54 moles of that left over; water doesn't matter.1431

0.45 moles; we have produced 0.01 moles; we have 0.46 moles.1443

Now, there is .54 moles of HAc floating around; there is .46 moles of Ac floating around; there is no more OH floating around--that has been bound up; it has formed water (it doesn't matter; water is water).1456

Now, we can do the equilibrium part.1470

Again, major species: now what do we have floating around in solution?1480

Well, we have HAc, a little bit less than what we had before, because it reacted with the OH-; we have Ac-, a little bit more than we had before, because it was produced upon reaction with the OH-.1486

We have H2O, and we still have the sodium ion floating around: irrelevant, irrelevant.1501

These two are going to dominate.1508

Now, the system can come to equilibrium according to: HAc, H+, plus Ac-; now we can do our ICE chart; Initial, Change, Concentration--I'm sorry, Initial, Change, Equilibrium.1511

OK, ICE charts--we have to deal in molarity.1527

We said we have (I'm going to make this a little bit bigger--I'm sorry--I need some room to actually do the chemistry here, so let me go...): HAc is in equilibrium with Ac-, plus H+.1530

Let me keep it consistent: I usually always write the H+ first.1553

H+ + Ac-: Initial, Change, Equilibrium.1558

We said we had 0.54 moles of HAc left; it is floating around in 1.0 liter; so the molarity is 0.54 Molar.1564

H+ hasn't come to equilibrium yet; this is initial--we said we had 0.46 moles of Ac- floating around in the 1 liter, so the concentration is 0.46 molarity.1578

This will diminish; this will augment; this will augment; at equilibrium, we have 0.54-x; x; and we have 0.46+x.1592

Now, we have our Ka, 1.8x10-5, equals the hydrogen ion concentration x, times the Ac- concentration, 0.46+x, divided by the HAc concentration of 0.54-x.1608

OK, and this is approximate, because x is going to be very, very small in this case.1630

We have x times 0.46, divided by 0.54.1635

When we do this, we get x is equal to 2.11x10-5, which gives us a pH equal to 4.68.1642

There we go: now, recall: before the addition, we had a hydrogen ion concentration of 2.2x10-5; that was a pH equal to 4.66.1658

After we added the sodium hydroxide, now our concentration is 2.11x10-5; now, the pH is equal to 4.68.1675

The ΔpH is 0.02; it is virtually nothing.1694

This is the power of a buffer solution.1700

Notice what happened: I had a buffer solution at a prescribed pH of 4.66; I added .010 moles (it doesn't seem like a lot, but it's a fair amount) of sodium hydroxide; the pH only changed by .02.1711

In fact, the hydrogen ion concentration (if you were to take the difference of the 2.2x10-5 and the 2.11)--the Δ of the hydrogen ion concentration is 9x10-7--it's virtually nothing.1725

The hydrogen ion concentration wasn't affected at all.1738

That is the power of a buffer solution.1741

OK, now we are going to actually do the comparison that we said we were going to do.1744

We want to now find what the pH is going to be if I take neutral water (not a buffer solution, just regular water) and drop in .010 moles of sodium hydroxide.1749

OK, so now, let's do the comparison.1759

Now, our major species upon addition of the hydroxide (sodium hydroxide) to water: we have--well, we have water; sodium ion; hydroxide ion; and we have water; that is it.1769

Sodium hydroxide is a strong base--completely dissociated--that is floating around; that is floating around; that is floating around.1801

Something is going to dominate the concentration here; well, water can behave like a base; OH can behave like a base; the difference between these two...it's that that is going to...1807

So basically, what we are going to do is take the H+ concentration (remember, we said that is equal to 10-14), over the OH- concentration: 10-14 over 0.0100 equals 1.0x10-12.1820

Our pH is going to be 12; well, neutral water is pH=7; the ΔpH (it might be nice if I actually wrote properly here) is equal to 12.0-7.0; look at the difference.1842

If I add .01 moles of sodium hydroxide to pure water: .01--that is not very much!--but apparently, it is a lot--the pH changes by 5 units: it goes from pH 7 to pH 12.1878

But, in a buffer solution, I add the same amount of sodium hydroxide, and the pH only goes up by .02--virtually nothing; almost not even noticeable.1896

This is the power of a buffer solution.1906

OK, now, I know that I had a second example, but I think I'm going to hold off on that second example, and I'm going to actually start the next lesson with that second example, just to give you a chance to sort of work on this one a little bit.1909

And then, it will be something nice to start the other one with, as a review of the previous lesson; because this is very, very, very important, and I don't want to end up concentrating all of it into this one lesson.1921

So, again, go through this very, very carefully--it's profoundly important that you understand what is happening in a buffer solution.1931

If you understand the chemistry, the math should make complete sense.1937

Thank you for joining us here at Educator.com.1941

We'll see you next time; goodbye.1943

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