Raffi Hovasapian

Raffi Hovasapian

Acids and Bases

Slide Duration:

Table of Contents

Section 1: 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
Section 2: 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
Section 3: 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
Section 4: 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
Section 5: 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
Section 6: 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
Section 7: 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
Section 8: 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
Section 9: 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
Section 10: 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
Section 11: 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
Section 12: 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
Section 13: 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
Section 14: 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
Section 15: 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
Section 16: 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 (22)

2 answers

Last reply by: David Gonzalez
Mon Jul 28, 2014 5:45 PM

Post by David Gonzalez on July 1, 2014

Hi, great lesson! You mentioned that hydrogen cations are dangerous (acidic) - but how come they don't cause damage in the body when released into the medium? Thank you!

2 answers

Last reply by: Tim Zhang
Sun Apr 6, 2014 6:19 PM

Post by Tim Zhang on April 3, 2014

I am a little confuesed about the equailibrium of the acid and base reation. For example, if the reaction has a stronger base on the left side, will the equilibirum favor the products? And, could you expain what will be produced when Oxygen ion mixed with water?

1 answer

Last reply by: Professor Hovasapian
Fri Apr 4, 2014 7:08 PM

Post by antoni szeglowski on April 1, 2014

Thank you so much Professor Hovasapian.  Last exam I took in chemistry I scored a 90...the class average of about 1500 students taking the exam was around a 55.  

This is all thanks to you Professor Hovasapian.  I look forward to your physical chemistry lectures!

If only you taught all the courses on educator!

1 answer

Last reply by: Professor Hovasapian
Thu Dec 26, 2013 3:08 PM

Post by Burhan Akram on December 25, 2013

Hi Prof. Raffi,

Just wanted to say that you explain things VERY WELL and I deeply understand the material. Also, after watching your lessons, I started to imagin things in real world how they relate to Chemical reactions in chemistry; Very helpful. I might end up doing a minor in Chem along Math now :D

Thank You Again

Burhan

1 answer

Last reply by: Professor Hovasapian
Wed Sep 11, 2013 4:35 PM

Post by Stephanie Dahlström on September 11, 2013

I'm a little confused about the H3O+ - molecule. O is -2 and with two H+ attached to it, the shell should be completely full, which means that it has achieved noble gas structure, right? So why would H2O want to take another H+ to make it an ion?

1 answer

Last reply by: Professor Hovasapian
Sat Jul 6, 2013 6:41 PM

Post by KyungYeop Kim on July 5, 2013

Question: When you look at different compounds, how do you know which one is stronger or weaker Lewis acid or base?
For instance, BF3 and BCL3, how do I know which one's stronger? Also is it any different with Lewis bases?

1 answer

Last reply by: Professor Hovasapian
Wed Jun 19, 2013 10:03 PM

Post by Jeff Q on June 19, 2013

Hi Raffi. In example 4, when you gave the [OH-] in solution, you assumed that the solution was water. Is that an assumption we should usually make? That is, when a problem just says "solution" should we presume it is an aqueous solution?

1 answer

Last reply by: Professor Hovasapian
Wed Nov 28, 2012 1:43 PM

Post by Dustin Voelzke on November 27, 2012

I sit in my class like a deer in headlights thinking that i'll never understand chemistry. My professor is a smart man but cannot articulate this stuff as beautifully as you can. My grades have picked up ever since I started this and I have a new found interest in the subject matter. Sorry if this comment sounds like an as seen on TV ad. Just wanted to say thank you.

1 answer

Last reply by: Professor Hovasapian
Tue Nov 27, 2012 3:41 PM

Post by Kamiko Darrow on November 27, 2012

Hi Raffi,

Your Chemistry lectures are so helpful and easy to follow. I am enjoying your lessons a lot, and I wish you were teaching at Boston University!

1 answer

Last reply by: Professor Hovasapian
Sun Jul 29, 2012 8:08 PM

Post by kwasi agyeman on July 28, 2012

There is a white block in the middle of the lesson that will not go away. I cannot see.

Related Articles:

Acids and Bases

  • The only thing that moves in an A/B reaction is Hydrogen Ion (H+).
  • Diferent acids have different strengths: the extent to which they dissociate to release Hydrogen Ion -- more dissociation means stronger acid.
  • pH is just another way of measuring the amount of Hydrogen Ion floating around in solution.

Acids and Bases

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
  • Acids and Bases 1:14
    • Bronsted-Lowry Acid-Base Model
    • Reaction of an Acid with Water
    • Acid Dissociation
    • Acid Strength
    • Example 1
    • Water as an Acid & a Base
    • Example 2: Part A
    • Example 2: Part B
    • Example 3: Part A
    • Example 3: Part B
    • pH Scale
    • Example 4

Transcription: Acids and Bases

Welcome back to Educator.com; welcome back to AP Chemistry.0000

Last time we finished off our discussion of equilibrium, a very, very important topic, and now we are going to go on to probably the second central topic of chemistry, which is acids and bases.0005

They're very, very important in all aspects of chemistry, particularly in biochemistry, because much of the reaction that takes place in your body is, of course, acid-base reactions.0016

We are going to spend some time discussing acids and bases and writing reactions and talk about what is going on qualitatively, and then, we will also deal with it quantitatively, and we are going to introduce a new equilibrium constant, called the Ka or the Kb.0026

Now, fortunately, it isn't really anything new; it isn't new in the sense that it is a different way of writing the equilibrium constant.0043

It is just a different symbol for it; it is still exactly what you learned with equilibrium, and everything that we learned with equilibrium will again play the same part.0049

We will be doing ICE charts, so we will get plenty of practice in this.0058

If you feel like you don't have a solid grasp of equilibrium, acids and bases are a wonderful place to continue to practice that until you get a solid grasp of it.0062

Let's go ahead and get started with some definitions.0071

OK, what we are going to be using: there are a couple of models for acid-base behavior, and the primary model--at least, the one that is used most often--is something called the Bronsted-Lowry acid-base model.0075

Let's just talk about what that is; we'll just define it.0087

Bronsted-Lowry...the name itself doesn't really matter all that much--for historical reasons, I suppose it's nice to give credit to the people who deserve credit--but it is the model that is important: acid-base model.0093

Now, an acid is a hydrogen ion donor, and a base is a hydrogen ion acceptor--a very, very simple definition, and let's talk a little bit about what this means.0109

So, an acid is a compound that has an H+ that can come off--that can potentially come off.0133

I'll write "potentially come off"--in other words, it has a hydrogen ion to donate, if it needs to donate it.0161

More often than not, it will actually donate it--it will give up the hydrogen ion.0170

A base is something (I shouldn't say something--it's a compound--yeah, right, it's just "something")...is a compound that can potentially take that (not that; it will be that, but let's be as general as possible)...that can take an H+ ion.0176

That is it: so, a base is a compound that has a hydrogen attached to it, but the hydrogen can potentially separate from that compound, and just float around freely as hydrogen ion.0221

That is, in fact, the acid: when we speak about an acid, that is what an acid is: it is just free hydrogen ions floating around in solution--that is what does the damage.0233

A base is something that has the capacity to either take a hydrogen ion from something (an acid) that has it to give--in other words, rip it off--or, if there are free hydrogen ions floating around, a base can actually grab onto it.0241

That is it: so an acid is a hydrogen ion donor; a base is a hydrogen ion acceptor.0260

Acid-base chemistry--they come in pairs; when there is some acid, there is some base, usually; so we will see what that means in just a minute.0266

OK, there are two ways to write the reaction of an acid in water--or with water, I should say: to write the reaction of an acid with water.0277

And again, an acid is something that you are actually dropping in water, and once it is in water, it actually tends to come apart.0304

Here is what actually happens: we will just write H, and we will write A; so H is the H that actually comes off; this A can be anything--it's just a generic other part of the molecule.0312

...Plus H2O, goes to H3O, plus A-.0326

What has happened is that the H has gone from this HA over, and attached itself to the H2O; it has actually attached itself to the oxygen, and it is H3O+, now.0335

H2O is neutral; H+ is plus charge--that is why now, this has a plus charge; and because this was HA, it's a neutral ionic compound, but this H left its electron when it left (it left as H+), it has an A-.0346

There is another way to write this reaction, which is the way that I am certainly accustomed to doing it, and in certain circumstances, it helps to write it this way as opposed to this way.0365

It is why we actually have two ways of doing it, because it depends on the problem and how convenient it is to write one way or the other.0375

HA...it's the actual dissociation reaction: when you put this in water, it becomes free ion, H+ ion, and free A- ion.0382

This is a little bit more descriptive of what is really going on, in the sense that HA dissolves, like sodium chloride dissolves, and it breaks up into a hydrogen ion and the other ion (whatever the rest of the molecule is).0393

What is really happening is: there is an actual transfer.0409

This is the acid; it has the hydrogen to give up--it has the hydrogen to donate.0413

In this case, water is acting as the base; it has the capacity to take that hydrogen, or to accept that hydrogen.0418

So, when we write it this way, we actually talk about an acid conjugate base pair, which I will actually get to in just a minute; but I just want to sort of talk about these two equations, and how to represent them.0426

When we talk about an acid dissociating, or an acid reacting with water, we can either write it this way, or we can write it this way.0443

This H3O+ and this H+ are the same thing.0452

This H+ is just an H+; when it is attached to an H2O, we write it as H3O+, but the chemistry is the same.0456

Now, let's go ahead and write the equilibrium constant for this.0466

The Keq--well, we said the Keq is the concentration of product, concentration of product, divided by concentration of reactant, concentration of reactant.0470

This is liquid water; it doesn't show up--remember, liquids and solids don't show up in the equilibrium constant.0483

This is going to be: the concentration of H3O+ in moles per liter, raised to its coefficient (which is 1), the concentration of A-, raised to its coefficient (which is 1), divided by the concentration of HA; this is aqueous.0488

OK, water doesn't show up; when we have an acid written like this, or (let me actually write this version of it: H+, A-, over HA) when we are dealing with an acid, we don't call it Keq; we call it Ka.0507

This is the acid dissociation constant; it is the equilibrium constant for the reaction of an acid with water, or for the reaction of an acid dissolved in water.0530

When you dissolve it in water, we think about it as just a dissociation, like any ionic compound.0543

But what is really happening--it is actually reacting with water.0547

The acid is giving up the H; the water is acting as the base, accepting the H.0550

In some sense, when you look at this, this equation, this equilibrium--what it expresses, what it represents...it represents a competition between water and the conjugate base for this H.0555

In other words, does the H stay with the A, or does it go with the water?0570

If it goes with the water, that means the reaction is over on this side: H3O+ is formed.0574

If the H doesn't want to leave the A--if it wants to stay with the A--that means most of the reaction is on this side.0580

That is what the Ka measures.0586

So, the equilibrium constant is called the acid dissociation constant.0588

It is the same as any other equilibrium, except it is representing what is happening when the acid reacts with water, when the acid is dropped into water.0595

That is it; so, from now on, when we are discussing acid-base chemistry, it is going to be Ka.0605

Later on, when we discuss bases, we are going to have an analogous equation, and it is going to be Kb.0610

But it is just an equilibrium constant.0615

And what does an equilibrium constant represent? It represents the extent to which a reaction has moved forward.0619

Once it reaches equilibrium, how much of this, how much of this, how much of this?--it's a measure of the extent of reaction.0626

How far forward has it gone? How far has it not gone?--that is all it is a measure of.0634

The higher the Ka, that means the dissociation; that means it's to the right, because the numerator is a bigger number.0639

OK, now, let's talk about polyprotic acids really quickly.0648

For acid dissociation, 1 H+ leaves at a time.0656

For example, if I have the acid H3PO4, which we know of as phosphoric acid, it doesn't do this: H3PO4--it doesn't just dissociate into 3 H+s plus PO43-.0683

It doesn't do that; that is not how acids behave.0699

This doesn't happen; what does happen is that 1 H+ leaves at a time, so the dissociation of phosphoric acid is three steps, and it looks like this.0704

H3PO4 + H2O is going to be in equilibrium with H3O+ + H2PO4-.0722

That is the first dissociation; there is some Ka associated with that first equilibrium.0733

And then, H2PO4 has two more Hs to give up; so, H2PO4 (actually, you know what, I think I'm going to write it as just a regular dissociation--I'm sorry; I think it will be a little more clear) now gives up another one.0739

It gives up 1 at a time: H+, and it becomes HPO42-; there is an equilibrium associated with that.0762

OK, and then we have HPO4; it has another H+ to give up, and now it's PO43-; there is a third acid dissociation constant.0773

And each one of these is actually different; these phosphoric acid dissociates to differing degrees--these are not the same number.0787

In other words, this equilibrium might be really far forward; this equilibrium may be really far this way; it just depends.0796

The key idea is: for a polyprotic acid (polyprotic just means having more than one hydrogen to give up), it gives them up one at a time.0805

Very, very important: for acids, it releases one at a time; it doesn't happen all at once.0813

So, H2PO4: it doesn't just release both hydrogens--it releases one completely, and the second one comes off very little, in fact--so, just something that you should know.0819

OK, so let's talk about acid strength.0830

Acid strength: Acid strength is defined by the position of the equilibrium constant--or by the position of equilibrium; in other words, which is measured by the equilibrium constant.0837

It's measured, it's defined, by the position of the equilibrium; and the equilibrium, we said, was some acid plus water, going to hydronium ion...or just H3O+...0856

We call it hydronium, by the way--let me just write that name down: H3O+ is called a hydronium.0874

Anything with a positive charge tends to have an -onium ending; it's the same as H+--you can treat them the same.0882

So, the acid strength is defined by the position of the equilibrium; if the equilibrium lies really far to the right, that means, if this acid has completely given up all of its Hs over to water, it is a strong acid.0891

So, the farther to the right the equilibrium is, that means it is a stronger acid.0903

We talk about a strong acid being something that actually does not want to keep its H's; it wants to give them up completely.0914

It wants to be A-, and it wants its H to be H+.0922

It does not want to be together.0928

OK, now, let's talk about...let's see; so, actually, let me write that again.0930

Let me say: farther to the right, stronger acid--so, when we talk about acid strength, we are talking about relative strength; it's always...we are comparing it to something.0942

Stronger acid--well, that means a higher Ka, right?--because the Ka is products over reactants; well, farther to the right means there is more product and very little reactant--a really big numerator, a really small denominator, and a huge number--stronger acid.0958

OK, now let's go back to blue here.0976

A strong acid implies a weak conjugate base.0983

Acids and bases come in conjugate pairs--"conjugate base"; we get the conjugate base by just pulling off the hydrogen from the acid.0988

So, let's write our equation again: HA + H2O is in equilibrium with H3O+ + A-.0999

This is the acid; this is its conjugate base.1010

I pulled off the H; that leaves my conjugate base.1021

If I take a base (in this particular case, again, this is the acid--this has the H to give up; water is acting as the base--it is going to take that H, or accept it, if it wants to), when it accepts the H and it becomes H3O+, this is called the conjugate acid.1026

That is it; if you have an acid, you pull of the H; you are left with a conjugate base.1048

If you have a base, and if you add a hydrogen ion to it, you have its conjugate acid.1053

That is it--nothing more than that; no more, no less: don't read any more into that.1060

OK, so a strong acid means a weak conjugate base, which is exactly what you would expect.1064

A strong acid does not want to hold onto its H's; it will give them up completely; that means it has a weak conjugate base.1070

A weak conjugate base means that it doesn't want this H; that is the whole idea.1077

A strong base wants the H--will take the H; it's a competition.1084

A weak acid has a strong conjugate base.1091

A weak acid is one that does not want to give up its H's; and the reason it doesn't want to give up its H's is because its conjugate base is so strong that it literally holds onto its H's without letting go.1104

That is the idea; so, weak acid means conjugate base.1116

When you are thinking about this, you have to pick a perspective; you have to pick a point of reference; that is the whole idea.1121

We decide to take the reference point as the acid itself; we call it strong; we call its conjugate base weak.1127

That means the conjugate base doesn't want the H; it gives it up.1134

A weak acid is one that does not want to give up its H; that means that its conjugate base is very, very strong--it will not release its H.1139

So, it's really a question of perspective; you need a point of reference; and the point of reference that we have taken is the acid.1149

We talked about a strong acid and a weak acid: a strong acid has a weak conjugate base; a weak acid has a strong conjugate base.1157

OK, for strong acids, no Ka exists...no Ka is listed, we should say.1172

So, for example, hydrochloric acid is a very strong acid; its dissociation (it's not even in equilibrium)--it completely breaks up into free H+ and free Cl-.1184

You won't find any of this at all.1195

Well, the Ka for this is H+(Cl-)/HCl.1198

Well, since there is a whole bunch of this and a whole bunch of that, and virtually none of this, this denominator is really tiny; the numerator is really big.1207

So, you have a Ka which is huge; in fact, the Ka is so huge that we can't even measure this.1216

Because we can't measure the Cl concentration accurately, we don't even list a Ka for it; it's just off the charts.1222

Strong acids--full dissociation; that is the idea.1228

Once again, strong acid means full dissociation.1233

That means this H and Cl completely come apart; in solution, you will never find any Cl attached to an H.1240

It will free Cl- and free H+.1249

That is what makes it a strong acid; because there is so much H+ floating around freely, that is what does the damage.1252

So, when we say "strong acid," we are talking about a chemical property.1259

Strong acids do damage because there is so much H+; there is so much H+ because the H and Cl do not want to stay together--they want to be separate.1263

Strong acid--full dissociation: that's very important.1274

OK, let's do an example to get a sense, or a feel for what is going on, for some numbers here.1279

So, Example 1: HF, which is hydrofluoric acid, has a Ka of 7.2x10-4.1290

It's a small number; it's a weak acid--it's not very far forward.1301

Most of it is in this form; OK, so let me write the dissociation: HF goes to H+ + F-.1305

When I measure this, this, and this, and I put these on top of that--the Ka--I get this number.1314

It's pretty tiny; a tiny Ka means that most of this equilibrium will be found on this side, the reactant side.1320

In other words, there is very little dissociation; most of it just stays HF.1327

Hydrocyanic acid, HCn--it has a Ka of 6.2x10-10--wow, really, really small.1333

So, when I see HCn--when I drop some of that in water (that is Cn-), this would be the dissociation.1343

This is so tiny that this is telling me that most of this reaction is over on this side.1351

I won't find a lot of free H+ and Cn-; I'll find some--I was able to measure it--but most of it is this.1356

That is what the small Ka means; a small Ka means it hasn't dissociated very much.1362

A big Ka means it has dissociated a lot more.1368

Between these two, this is bigger; this is smaller; this is a stronger acid, because its Ka is bigger--it has dissociated more.1371

It is still a weak acid, but compared to hydrocyanic, it's stronger than hydrocyanic.1383

That is the difference; when we speak about the strength of acids, we are often going to be comparing; it's relative.1390

We are going to be talking about 2 or 3 or 4 species.1396

Now, the question is: which one has the stronger conjugate base?--that is the question.1402

Which acid has the stronger conjugate base?1409

Well, for HF, the conjugate base is--just take off the H; this is the conjugate base.1421

For HCn, it is the acid minus the H; that is the conjugate base.1427

Which has the stronger conjugate base? Well, the stronger conjugate base comes from the weaker acid, right?1433

Weak acid=strong conjugate base; OK, between these two, the weaker acid is the hydrocyanic acid.1445

In other words, it's a weak acid because it doesn't dissociate very much; that is confirmed with this really, really small Ka.1456

It doesn't dissociate very much; it's a weaker acid; it has a stronger conjugate base.1464

In other words, the base is so strong that it doesn't want to release its hydrogens; it wants to hold on to it.1470

That is what a strong base does: a strong base grabs onto its hydrogens and holds onto it.1476

A strong acid wants to give up its hydrogens; that is why we have strong acid, weak base--yes, stronger acid, weaker base; weaker acid, stronger base.1482

So, in this case, F- is a weaker conjugate base than Cn-.1495

That is it; we are using Ka values to decide, depending on what species we are talking about.1509

This Ka means HF is a stronger acid than HCn.1515

It means its conjugate base is weaker than Cn-; that is what is going on here.1520

OK, so now, let's discuss water as an acid and a base.1526

Let's define something called amphoteric: an amphoteric substance is one that can behave as an acid or a base.1538

It should be "and a base," because it actually does both; "as an acid or a base"--in other words, something that can go both ways.1561

Well, let's look at the dissociation of water.1571

I'm going to write it slightly differently.1575

I'm going to write the first water: so now, we are talking about water as a species dissolved in water.1577

It's plain old water; I know it's a little weird, but think about it this way.1583

I'm going to write it as HOH + H2O.1587

These are both water, but I wanted to emphasize what is happening.1591

This HOH, believe it or not, actually dissociates just a little bit; it breaks up into H+ and OH-.1596

Well, this H goes over here to become H3O+ + OH-.1602

Another way of writing this is HOH, as free dissociation, without involving the water species, breaks up into H+ + OH-1611

Well, let me rewrite it: HOH is in equilibrium with a little bit of H+ and a little bit of OH-.1622

The Keq for this is equal to...well, this is liquid water; liquid doesn't show up in the equilibrium concentration; this is aqueous; this is aqueous; this is floating around, in other words, and this is floating around in the water, to a certain degree.1634

It is equal to the H+ concentration times the OH- concentration--the products.1650

Well, as it turns out, for water, we actually call it Kw.1656

We have done the experiment; and, as it turns out experimentally, we have actually measured, at 25 degrees Celsius...experimentally, at 25 degrees Celsius, the hydrogen ion concentration is equal to the hydroxide ion concentration; that is what this equation says.1662

For each H that is produced, one OH is produced.1685

They have the same concentration; it equals 1.0x10-7 moles per liter.1688

The concentration of H+ and OH- in standard, neutral, run-of-the-mill water, is 1.0x10-7.1696

Well, what does that mean?--Kw is H+ times OH-, so Kw equals 1.0x10-7 squared, equals 1.0x10-14.1706

Water has a Ka of 1.0x10-14.1723

We give it a special name; we call it Kw for water--it's very, very important that you know this.1728

Kw is 1.0x10-14 at 25 degrees Celsius--always.1734

OK, what does this mean?--it does have a meaning.1743

This means (and here is where it gets really, really important): in any, in any aqueous solution, at 25 degrees Celsius, no matter what is in it--no matter what--no matter what is in it (in the solution), the product of H+ and OH- must always equal 1.0x10-14.1753

OK, we need to stop and think about this, because this is really important.1811

Any aqueous solution, no matter what is in it--sodium chloride, potassium phosphate, the complex ion, magnesium hydroxide, phosphoric acid, permangenic acid--anything--if I were to measure the hydrogen ion concentration and the hydroxide ion concentration, they have to--the products have to equal--have to equal 1.0x10-14.1816

That is what this says; at any given moment, an aqueous solution (in other words, water--something that is in water)--the product of the hydrogen ion and the OH- ion concentration always 1x10-14.1842

That doesn't mean that each one is 1x10-7; that is just for neutral water.1858

It just says that, if...so this is equal to the H+ concentration times the OH- concentration--that means, if the H+ concentration rises, the OH- concentration has to drop in order to retain the constant value of 1.0x10-14.1863

So again, they don't have to equal each other; their product has to be a constant.1883

If one goes up, the other goes down; that is the relationship--that is the mathematical relationship.1888

OK, their product has to be 10-14; and it is going to be very, very important in just a minute.1894

So, let's just...if the concentrations do equal each other, we call it neutral.1899

If the hydrogen ion concentration is bigger than the OH- concentration, we call it acidic.1910

That is where we get the idea of "acidic"; so, when you say that orange juice is acidic, that means that, if you take some orange juice, and if you measure the concentration of H+ versus the concentration of OH-, the H+ concentration is going to be higher.1919

But, the product of the two concentrations is still going to be 10-14.1934

If the H+ concentration is less than the OH- concentration, we call that solution basic.1940

That's it; OK, so let's do an example.1948

Nice, simple math: Calculate H+ and OH- (I'm really, really sorry about my handwriting; I know) concentration for the following solutions.1956

OK, let's see: Oh, "at 25 degrees Celsius."1982

A: We have: 1.5x10-5 molarity OH-.1991

Well, we know what the concentration of OH- is; it's right there--it's 1.5x10-5.2000

Well, we know that the H+ concentration, times the OH- concentration, at 25 degrees Celsius, equals 1.0x10-14, so we just plug it in to find the H+ concentration.2010

H+ concentration equals 1.0x10-14, divided by the OH- concentration; I have just done a simple division.2026

1.0x10-14 (ten to the negative fourteen, not fifteen), divided by OH-, which is 1.5x10-5, and I get a concentration of 6.7x10-10.2038

Which one is a higher concentration? OH- is 10-5; H+ is 10-10; the OH- is a bigger concentration--the solution is basic.2058

It will still hurt you, so don't...it doesn't mean that a base won't hurt you; an acidic solution will hurt you--a basic solution will also hurt you.2072

OK, the idea is: we want a neutral solution.2080

By the way, blood pH is just about...we will get to that in a minute; sorry about that.2083

OK, so let's do another one; let's do a 2 Molar solution--2.0 Molar solution of H+.2091

Well, again, we have the H+ concentration: in this case, a 2.0 Molar solution, a solution that contains 2 molarity H+; that is the hydrogen ion concentration.2101

Well, we know that the H+ concentration, times the OH- concentration, equals 1.0x10-14; therefore, the OH- concentration equals 1.0x10-14, divided by 2.0, which was that.2112

And we end up with an OH- concentration of 5.0x10-15--a very, very small number; in this case, the hydrogen ion concentration hugely, hugely out-values the OH- concentration.2132

This is acidic--I would say highly acidic; that is it; very, very dangerous.2149

OK, let's do another one here.2158

Example 3 (or, actually, I'm not sure if it's Example 3, but...I'll just do an example): At 100 degrees Celsius, the Kw equals 5.13x10-13.2163

So remember: we said that, at different temperatures, the equilibrium constant is different.2183

At 25 degrees Celsius, the Kw is 1.0x10-14; but if I change the temperature to 100 degrees Celsius, now it has gone up to 5.13x10-13.2187

It has gone up.2199

The question to you is: Is this reaction, a dissociation of water into H+ plus OH---is it endo- or is it exo-thermic?2202

OK, well, let's see what happens here: at 100 degrees Celsius, they are telling me that the Kw equals 5.13x10-13.2227

They want to know if this reaction is endo- or exothermic.2240

Well, let's see what we did: we increased the temperature, and by increasing the temperature, we actually pushed the reaction forward, because that is what this Kw is telling us.2243

This Kw of 5.13x10-13 is bigger than the 1.0x10-14, right?--negative 13 is bigger than negative 14.2256

That means that the equilibrium has actually shifted to the right; that means more product has been produced.2267

That means that the reaction shifted to the right when I increased the temperature.2273

Well, if it increases to the right when I increase the temperature, that means that temperature, or heat, must have been one of the reactants, because, if it's one of the reactants, and if I increase this temperature, the system is going to want to offset what I do to it (which is an increase) by decreasing the temperature.2283

Well, in order to decrease the temperature, it has to use up this extra heat that I have pumped into it.2311

In order to use up the heat, it has to shift the reaction that way; it has to produce more H and more OH-, which is exactly what it did, because now the equilibrium constant has gone up.2317

When an equilibrium constant goes up, that means the reaction has moved to the right.2329

The numerator of the equilibrium constant has increased; so, in this case, because heat is on the left side of this equilibrium arrow, this is an endothermic reaction.2333

So notice, we used our discussion of equilibrium and Le Chatelier's Principle to actually find out, qualitatively, that this is an endothermic reaction.2346

In order to push this reaction forward--in order to produce more H+ and more OH-, I actually have to increase the temperature, which means that it is an endothermic reaction.2356

I don't know what the ΔH is, but I do know that it is going to be positive for this reaction.2368

OK, B: what is the H+ and OH- concentration in a neutral solution at 100 degrees Celsius?2375

Well, we just said that Kw is equal to the concentration of OH- times H+ at 100 degrees Celsius.2398

We gave it already; it's 5.13x10-13.2407

Well, in a neutral solution--go back a slide or two--neutral means that the OH- concentration is equal to the H+ concentration.2412

So, if I call this X and call that X, I end up with: X squared is equal to 5.23x10-13.2425

I get: X is equal to 7.16x10-7 molarity.2435

At 100 degrees Celsius, there is 7.16x10-7 moles of H+ per liter of water, and 7.16x10-7 moles of OH- per liter of water.2444

That is all that is going on here; molarity is moles per liter.2461

OK, let's see what else we can do.2466

Now, we are going to close this discussion off by introducing the pH scale.2474

Now, I have to (I probably shouldn't say this, but I'm going to go ahead and do it, because I tell all my students this)...working with concentrations is perfectly valid--concentration is something we always work with; concentration of H+, concentration of OH-, concentration of Cl-...2479

With acids and bases, because these concentrations tend to be really, really small, as you have noticed (you know, 10-7, 10-6...), chemists have come up with a way to have numbers that actually are a little more easy to deal with--normal numbers like 1, 2, 3, 4, 5, 6, 7, 8, 9, 10.2499

So, they have created this thing called a pH scale.2516

Here is what it means: the p of anything is equal to negative log of the concentration of that anything.2519

So, in other words, if I talk about pH, that is equal to the negative log of the hydrogen ion concentration.2532

If I talk about pOH, that is equal to the negative log of the hydroxide ion concentration.2540

If I say pKa (so this p function means whatever it is...p of whatever--just take the negative log of that whatever), pKa is the negative log of Ka.2546

It is a way of taking numbers that are really odd, like 6.6x10-3, and turning them into actual numbers that you can deal with: 3.16--something like that.2562

Personally, I don't care for it myself; I like dealing strictly with what we are dealing with.2571

We are dealing with concentrations; the equilibrium constant is expressed in concentrations; so, I think it is perfectly valid to deal with concentrations.2578

But again, this is so deeply entrenched in chemistry that literally nobody talks about concentration anymore when they speak about acid-base.2588

They speak about pH and pOH and pKa.2596

That is the only reason; but they are actually the same thing.2599

It is not like we are introducing a new concept; we are just introducing a way of changing the numbers of concentrations that tend to be really small, and turning them into numbers that make more sense, that look better--not make more sense, just look better, aesthetically.2605

OK, so one more example: if I said, "What is the pCl-?" well, that is going to be the -log of the Cl- concentration, if we happen to be talking about Cl-.2618

It is just a p function, and it is the negative log of the concentration we are talking about.2630

OK, let's do an example.2636

What is the pH and pOH of a solution which is 1.4x10-3 Molar hydroxide?2642

OK, well, they want the pH, and they want the pOH.2668

Well, I know the pH is the negative log of the hydrogen ion concentration; pOH is the negative log of the hydroxide ion concentration.2672

Well, they give me the hydroxide ion concentration, so why don't I just deal with that one first.2681

pOH is equal to negative log of the OH- concentration, which is negative log of 1.4x10-3.2685

I just stick in this in the calculator; I take the logarithm of it; I change the sign; and I end up with 2.85.2699

So notice, 2.85 is a much more attractive number than 1.4x10-3; at least, that is what many people think.2709

I, personally, don't think so, but that is fine; I'm a chemist; we deal with pHs; we'll deal with pHs.2716

But, don't think it is something different; it is not.2722

If you want to, you are more than welcome to deal strictly in concentrations, and if you ever have to find a pH, just at the last minute take the negative log of it: no harm, no foul.2724

Now, we want the pH; well, the pH--we know that the H concentration, times the OH- concentration, is equal to 1.0x10-14.2735

The H+ concentration equals 1.00x10-14 (well, let me just stick with 1.0; let's just do 2 significant figures here), divided by 1.4x10-3.2748

That is going to equal 7.2x10-12, and then, when I take the pH of that (that is going to be the negative log of the H concentration), -log of 7.2x10-12 equals 11.14.2767

11.14 pH; 2.85 pOH: notice something--what is 11.14 plus 2.85? Yes, you guessed it--it's about 14.2791

Now, let's explain why.2806

We said that Kw is equal to the hydrogen ion concentration, times the hydroxide ion concentration.2812

Let's take the negative log of both sides.2822

-log of Kw is equal to -log of this whole thing: H+, OH-...2824

Well, -log of Kw is pKw; that is it--equals -log of H+, right?--plus the negative log of OH-.2836

You end up with pH plus pOH; well, pKw is the negative log of 1.0x10-14, equals pH plus pOH.2860

Because Kw is 1.0x10-14, the negative log of that is 14; that is our final relationship.2884

So, for any aqueous solution at 25 degrees Celsius, pH plus the pOH of the solution equals 14.2895

This is just a restatement of the fact that the H+ concentration times the OH- concentration equals 1.0x10-14.2914

These are the same thing, except one deals with these numbers (these scientific notation numbers, the small ones); one deals with numbers that are a little bit more tractable.2928

In any aqueous solution, the product--the product--of the hydrogen ion and hydroxide concentrations is 10-14.2937

The pH plus the pOH of that solution is equal to 14.2948

OK, so we have introduced acid and base; actually, we have introduced acid mostly; we have talked a little bit about base, mostly in the context of conjugate base.2955

We have talked about the dissociation of acids and what happens in water.2964

We have introduced the idea of an equilibrium constant for that acid dissociation.2970

We have introduced the idea of water acting as both acid and base, an amphoteric substance.2976

And we have introduced the very, very important property that the product of the hydrogen ion and the hydroxide ion concentrations is 10-14; or, if we want to use the p scale, we can talk about the pH and the pOH equaling 14 for any aqueous solution at 25 degrees Celsius.2981

Next time, we will continue our discussion of acids and bases, and we will actually get into some of the more equilibrium-type problems that are involved.3000

Thank you for joining us here at Educator.com and AP Chemistry.3008

We'll see you next time; goodbye.3010

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