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Raffi Hovasapian

Raffi Hovasapian

Resonance & Formal Charge

Slide Duration:

Table of Contents

I. Review
Naming Compounds

41m 24s

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

37m 19s

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

31m 14s

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

43m 21s

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

47m 58s

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

31m 50s

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

49m 40s

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

32m

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

31m 58s

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

25m 34s

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

37m 32s

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

32m 34s

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

23m 9s

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

39m 28s

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

36m 24s

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

30m 48s

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

32m 17s

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

26m 40s

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

40m 59s

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

29m 8s

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

46m

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

40m 53s

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

45m 53s

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

31m 51s

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

40m 52s

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

50m 11s

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

43m 52s

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

43m 4s

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

35m 34s

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

41m 14s

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

41m 58s

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

32m 24s

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

40m 6s

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

38m 43s

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

42m 42s

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

42m 3s

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

52m 3s

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

36m 25s

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

42m 6s

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

43m 9s

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

43m 38s

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

31m 30s

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

56m 28s

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

39m 55s

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

30m 10s

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

30m 7s

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

44m 56s

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

39m 23s

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

43m 9s

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

48m 41s

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

41m 23s

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

34m 19s

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

33m 21s

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

44m 45s

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

54m

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

34m 4s

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

52m 43s

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

31m 17s

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

37m 23s

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

34m 11s

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

48m 39s

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

36m 59s

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

41m 21s

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

40m 17s

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

52m 34s

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

32m 15s

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

32m 50s

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

47m 22s

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

43m 5s

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

28m 36s

Intro
0:00
Free Response
0:43
Free Response 6: Part A. i
0:44
Free Response 6: Part A. ii
3:08
Free Response 6: Part A. iii
5:02
Free Response 6: Part B. i
7:11
Free Response 6: Part B. ii
9:40
Free Response 7: Part A
11:14
Free Response 7: Part B
13:45
Free Response 7: Part C
15:43
Free Response 7: Part D
16:54
Free Response 8: Part A. i
19:15
Free Response 8: Part A. ii
21:16
Free Response 8: Part B. i
23:51
Free Response 8: Part B. ii
25:07
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Lecture Comments (15)

1 answer

Last reply by: Professor Hovasapian
Fri Dec 8, 2017 11:13 PM

Post by John Prietto on November 24, 2017

So a side question, when drawing structures like HCO2- how do we know how they are to be drawn? You put H C O with another O above C. Is there a lecture that says how you draw structures?

1 answer

Last reply by: Professor Hovasapian
Fri Apr 7, 2017 10:02 PM

Post by Magic Fu on February 26, 2017

Hi, Professor H, you wrote carbon instead of nitrogen on the example of NO3-.

1 answer

Last reply by: Professor Hovasapian
Tue Oct 7, 2014 3:13 AM

Post by Jinbin Chen on October 5, 2014

Hello, Mr. Raffi!

It's me again! First, I really want to thank you for that detailed response regarding FTC, Green's Theorem, and Stokes' Theorem in Multivariable Calculus lectures. It really clears the confusion. But I still have some questions about chemistry (and that's why I compose separate posts when actually I can just write one of them).

So now I am almost at the end of the AP Chem lecture series here, and I would like to start reviewing using a real textbook. I am actually now trying to study the content on my own since my school doesn't offer AP Chemistry at all, but here is a problem: there are so many problems at the end of each chapter of a textbook, and finishing all of them will be extremely time consuming. Do you have some advice on how to use the end-of-chapter question, or should I just use the old AP exam question to review the content?

Another question is about the new lessons on this particular course. I know that you are filming the AP Calculus lectures (I am excited about that also since I am taking this exam this year as well), and the new lessons for this course may be delayed a little bit. But may I know the approximate date and content of these new lessons on AP Chemistry?

At the end, thank you so much for these wonderful lectures and your prompt responses.

Take good care.
Jinbin

3 answers

Last reply by: Professor Hovasapian
Mon Jan 6, 2014 5:56 PM

Post by Tim Zhang on January 4, 2014

Hello Professor Hovasapian.  I have just saw two enthalpies for oxygen, "o-o"and"O2" on my book. I am so confused on which one I should use for the formation of water, could you explain it to me, thank you.

1 answer

Last reply by: Professor Hovasapian
Sun Jun 2, 2013 3:01 PM

Post by KyungYeop Kim on June 1, 2013

You said polar bonds have partially different charges, but then how are we able to calculate former charges of the atom? Don't they all depend pn whom they're bonding?

1 answer

Last reply by: Professor Hovasapian
Tue Nov 13, 2012 4:42 PM

Post by Kelly Stucker on November 13, 2012

On the NO2- example, what happened to the lone pair on nitrogen?

0 answers

Post by Ivon Nieto Ivon Nieto on May 12, 2012

Thank You So much for you resonance and formal charge structure! Really helped my understanding!

Resonance & Formal Charge

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
  • Resonance and Formal Charge 0:09
    • Resonance Structures of NO₃⁻
    • Resonance Structures of NO₂⁻
    • Resonance Structures of HCO₂⁻
    • Formal Charge
    • Formal Charge Example: SO₄²⁻
    • Formal Charge Example: CO₂
    • Formal Charge Example: HCN
    • Formal Charge Example: CN⁻
    • Formal Charge Example: 0₃

Transcription: Resonance & Formal Charge

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

Today, we are going to talk about resonance and formal charge.0004

Let's just jump right on in.0008

I am going to start off with just talking about a particular molecule, and as we go through this molecule and discuss the Lewis structure, we will talk about formal charge.0010

I think it is a little better to do it that way than to actually just discuss it and then do the example.0019

Let's go ahead and talk about NO3-.0028

This is going to be the nitrate ion: NO3-; so let's go ahead and do a Lewis structure for this particular thing.0033

Now, nitrogen is going to bring 5 valence electrons, and oxygen brings 6 valence electrons; we have 3 of them, for a total of 18.0042

We know that, when we have a negative charge on something, that is just an extra electron.0051

This is going to be +1; so it looks like we have a total of 24 electrons to distribute among this thing, and let's see what kind of a Lewis structure we can come up with for it.0055

We have the nitrogen, and we'll go ahead and put oxygens around it.0066

And again, Lewis structures don't tell us anything about geometry--that is something else; but we need the Lewis structure to tell us how things are arranged.0071

OK, so we always start off with our single bonds first; so that is 6 electrons; let's see: 6, 8, 10, 12, 14, 16, 18, 20, 22, 24.0079

This doesn't work, because all of the oxygens have 8, but the nitrogen has 6.0098

So, no, this one does not work.0103

Let's try another one; rather than erasing electrons and moving them, I'll go ahead and do them one structure at a time.0106

That way, you can see the entire process.0112

N, O, O, O; let me move this one a little bit further to the right here; OK.0115

We have 2, 4, 6; let's go ahead and put that one there, and let's put 2 there, and let's put the remaining ones there.0126

I think this will do it: 2, 4, 6, 8, 10, 12, 14, 16, 18, 20, 22, 24; 24 is taken care of; oxygen has 8; oxygen has 8; oxygen has 8; and now, nitrogen has 8.0137

One of the oxygens and nitrogen--they are actually sharing 2 pairs; so the structure for this would be: N, double bond, O, single bond, O, single bond, O; 3 electron pairs; 3 electron pairs; 2 electron pairs here.0151

That is the Lewis structure; and of course, we have to put--because this is an anion--it's a complex--we do that.0169

OK, now, as written, that is a perfectly valid Lewis structure.0176

However, we could have done this...notice that I put the double bond over on the right-hand side--these two pairs of shared electrons; well, if we don't move anything, I could just as easily have put the two electrons up here--the two pairs of electrons between this oxygen and this nitrogen.0180

Or, I could have put it between this oxygen and this nitrogen.0203

Let me draw out those structures, also: I'm going to draw them out this way...well, yes, that is fine; I'll go ahead and do it vertically.0207

I could have done this one, and then a single bond here; so this is actually the same structure as that, but it is not the same structure as that.0215

What I mean by that is: it was just sort of random that the double bond ended up on the right-hand side; the double bond could have ended up in the middle.0231

Well, these two structures are actually called resonance structures, and a resonance structure is precisely that--it is if you can write a viable Lewis structure for a particular compound, but where the bonds are not exactly in the same place that they are for the other structure.0239

So here, the double bond is on the right, and the two singles are here; here, the double bond is up on top, and the two singles are flanking it.0259

As it turns out, these are actually not the same, but they are the same in the sense that they are both viable.0266

Because they are both viable, we have to include that.0274

Now, as you may have noticed, that is not the only viable structure; the double bond could also appear on the left; so I'm going to go ahead and write that structure also.0277

I'm going to do that one over here: single, single--now, the double bond is on the left, and it has two lone pairs, and the other oxygens--the ones with the single bonds--have three lone pairs.0286

This is yet another viable Lewis structure: I could just as easily have picked this, this, this.0298

As it turns out, all three (oops, this looks like it has 3 electrons; let me make sure...you know what, let me do this on either side of the oxygen, and make it a little bit more clear; OK)...0303

As it turns out, if you can write a viable Lewis structure where the bonds are not quite the same, but only the bonds have changed--nothing else, these are called resonance structures.0317

You actually have to notate that with a double arrow--not a double arrow like this (not a reaction double arrow), but we use a single line segment with an arrowhead on both ends.0331

We write something like this: this is our way, as chemists, of notating that the actual structure of this molecule is not this or this or this; it is actually all three simultaneously.0343

What it is--it actually ends up being an average of all of the resonance structures that you can draw for a particular molecule.0358

Some molecules--you can only draw 1 structure for it; that is the structure.0366

But, some molecules, because of the double bond and triple bond issue--they actually have more than one possible Lewis structure that is viable.0370

That means "that works"--that actually gives you an octet around each atom.0380

When that is the case, these are called resonance structures.0386

We have to represent--we have to actually list every resonance structure and connect them with this double arrow.0389

Now, here is what is interesting about this: here, it looks like this is a double bond; this is single, this is single; but, as it turns out, all the bonds--all the nitrogen-oxygen bonds in a nitrate ion--are actually equivalent.0395

They are all the same length, which means this single bond is not longer than this double bond.0408

That is what is interesting about this; so, the only reason that we draw things this way with resonance structures is that we don't have a particular notation to show that it's actually an average of these three.0416

All of these bonds are the same: imagine me taking this, putting it on top of that, putting it on top of that, and mixing it all up, and that is what I get.0429

The bond between nitrogen and oxygen is actually shorter than a single bond, but it is longer than a double bond.0437

Single bonds and double bonds and triple bonds have very specific lengths; so, as it turns out, all of these bonds are a little bit more than single, but a little bit less than double.0447

But, because we are representing these things with a notation, we don't have the notation available to actually show that; so we show it as three separate structures.0457

But it is important to remember that it is not this or this or this; it is actually an average of all three.0467

That is why all resonance structures have to be there.0474

That is all that a resonance structure is.0477

OK, so let's see: many bonds are equivalent...yes.0479

So now, I'm going to actually quickly describe how it is that resonance actually happens.0487

This is not something that you definitely need to know for general chemistry; it is something that is going to be a big deal when you go on to organic chemistry.0491

However, I want to introduce it here (it is very quick--it will only take a minute), just to show you why we can do this.0499

Now, from your perspective right now, coming up with these Lewis structures, we said, "Well, you could just as easily have picked the double bond here, the double bond here, or the double bond here."0506

As it turns out, no matter which resonance structure you start off with, there is a way to get to the other resonance structures.0516

This is the real meaning of "resonance structure."0522

I'm going to start (let me see, where should I start?) with this structure, and I'm going to go ahead and make this one, and make this one, and show how resonance actually works, because I want you to understand that this is not just something that is pulled out of the sky.0526

It actually has to do with electron movement.0540

Let me draw this structure here: the nitrogen...oops, let me put the double bond over here, and let me put this single; OK, I'll go ahead and write this.0543

Now, here is the real definition of resonance structure, and it is the definition that you are going to use in organic chemistry.0567

If you can write...if you can move an electron and put it someplace, and move other electrons and put them someplace else, in such a way that you don't touch anything else (in other words, the atoms don't move; the arrangement doesn't move; only electrons move), that is a good resonance structure.0574

Here is what happens, as it turns out.0596

Now, we know that we need an octet around each atom (right?); so here, we have 2, 4, 6, 8 around the nitrogen.0599

Here is what actually happens, and here is what you are going to notate when you do get into organic chemistry.0604

This electron pair actually moves down to form a double bond with nitrogen, but in the process of moving down this way, it pushes (oops, there is no electron pair there) this electron pair on top of this.0609

What you end up getting is this structure, which is the other resonance structure.0625

Electrons have moved; now, if I want to, I can take this electron pair; it's going to move down here to form a double bond on this side of the carbon, but it is going to push this electron pair on top of the oxygen.0639

That gives me another resonance structure: now the double bond is here; now this one has two lone pairs; this one, again, has three lone pairs; and this one has three lone pairs.0652

At this point, the electrons can't move any other place to form something that is different than the others.0663

Yes, this electron pair can come down here and kick this one up here, but I end up back at this structure.0670

There is no new structure that can form.0677

That is how you actually form a resonance structure: you find one Lewis structure, and then you start moving electrons--but only electrons; you don't move anything else.0680

If anything else moves, even a little bit, it is not a resonance structure; it is a completely different species altogether.0689

Let me go ahead and erase these arrows: this goes there; this goes there; and these are the three resonance structures that we had up here.0696

As far as these double arrows are concerned, it doesn't matter whether they are horizontal or vertical or diagonal, as long as they are there.0710

They have to be there.0716

Now, could you get away with something like this?--well, you know what, if I were grading your test, yes, you could get away with it; but chances are you won't be able to get away with it anywhere else.0718

So, on the AP exam, make sure you actually write the double arrows; it's very, very important, because this is the notation for resonance structure.0727

OK, I hope that makes sense.0737

All right, so let's do an example here: let's try...you know what, let me start on a new page; OK.0741

Example: Let's try NO2-; this is the nitrite anion.0749

OK, well, so we know that nitrogen brings 5 electrons; oxygen brings 6; there are two of them, so that is 12, plus the one electron from the negative charge.0757

We have 5, plus 1 is 6; 6 plus 12...we have a total of 18 electrons.0769

OK, let's see how we can arrange this.0774

I have O over here, O over here; OK, let's see: what else do I have?--so I do 2, 4; I start with my single bonds; 6, 8, 10, 12, 14, 16, 18.0778

Well, that takes care of it, but unfortunately, I don't have 8 around nitrogen (oxygen is fine).0806

This one is not going to work; this time, I will...actually, I will leave that there; I'm sorry; but I'm going to take one of these lone pairs (it doesn't matter from which oxygen); I'm going to take this lone pair, and I'm going to bring it down here--have it share, instead of being just on the oxygen.0812

Now, I have 8, 8, and 8; yes, 2, 4, 6, 8, 10, 12, 14, 16, 18 electrons; that takes care of it.0830

That is absolutely a good Lewis structure: so we have a double bond; we have 2 lone pairs there; we have a single bond; we have 3 lone pairs here; that is one structure.0841

Well, notice: the double bond ended up on the right side; I could just as easily have picked one of these electrons to share with the nitrogen, so it could actually go on the left side, as well.0853

Another Lewis structure is this one: watch the lone pairs--it's very important.0862

Notice, here it has 2; now it has 3; here it has 3; now it has 2.0870

This is another viable Lewis structure, and my complete Lewis structure for this nitrite anion is this.0875

It is an average of this and this; the bond between nitrogen and oxygen is shorter than a single bond, but it is longer than a double bond (because we don't have a notation for doing a one-and-a-half bond, or a 1.2, or a 1.8; that is why we use resonance structures).0886

Now, to show you--if I wanted to use that technique of moving electrons to go to my other resonance structures, here is how I would do it.0903

Let's start with the first one: N...0911

Actually, I'm going to do it over on this side.0915

I have N, double bond, single bond, 3; I start with this structure; I move electrons to form a double bond; I move this pair here, and in the process, I kick this pair back on top of that one.0918

That gives me the next Lewis structure.0936

Now, this is single-bonded; there were two pairs to start with; this pair that was shared now is a lone pair on the oxygen.0940

This pair that was a lone pair on this left-hand oxygen is now a shared pair, and it forms the second bond.0947

There you go; that is how you get your resonance structures.0956

You move electrons; or, if you don't like the idea of moving electrons, you can just think about it as, "Well, the double bond is here; by just randomly choosing this one, I could just as easily have chosen this one."0960

So now, put the double bond here and the single bond here; that is another way of recovering the resonance structure.0973

But it would be nice if you could get used to actually moving electrons around.0978

That is a nice way of doing it, and it is the way that you are going to do it when you do resonance structures in organic chemistry, for those of you that do continue.0982

OK, let's try another one: let's try HCO2-; this is the formate anion (formic acid from which a hydrogen has been removed).0989

You have two oxygens and a hydrogen attached to the central carbon; so let's talk about the electrons.1004

We have 1 from hydrogen; we have 4 from carbon; we have 12 from oxygen, plus that 1; so we have 12, 13, 14, 15, 16, 17, 18; again, we have 18 electrons.1011

Let me make sure to put that.1024

Here is what this one looks like: we have C; we have O; we have O; and we have H.1027

We put in our single bonds; H only needs 2, so that is done; we end up with something like this.1034

2, 4, 6, 8, 10, 12, 14, 16, 18; that is one; OK, so now let's draw that out as an actual Lewis structure with lines.1049

We have carbon double-bonded to an oxygen; oxygen has 2 lone pairs; single-bonded to an oxygen--oxygen has 3 lone pairs; single-bonded to a hydrogen; that is one structure.1061

Now, does this have a resonance structure?--well, yes, it does.1074

This oxygen is double-bonded; well, this oxygen could be double-bonded, and that can be single-bonded.1077

Let's do this as far as electron movement is concerned.1082

I'm going to bring these electrons down; I'm going to push these electrons up; the other Lewis structure is...1084

And again, you are not moving anything else ("as written"--you are not lifting it, twisting it, turning it); if you were to lift this and flip it, then it actually would be the same structure as you have written here.1091

But, when you leave everything else alone and you just move electrons, things are actually different.1103

This is a C; now the double bond goes here; 2 lone pairs; this is a single bond; 3 lone pairs; H stays the same.1109

There you go: this is the complete resonance structure for the formate anion.1122

And, in fact, that is exactly what it looks like: this is what it would look like.1129

Now, just as sort of a review of the last thing that we discuss (geometry), notice: we have 3 objects around a central atom.1134

3 objects around a central atom means that it is triagonal planar (or trigonal planar, depending--I like to say triagonal).1142

There is no lone pair on the central atom, so it actually is planar; it is like this.1150

The carbon is here in the center, and you have the 1 oxygen here, 1 oxygen here, and the hydrogen here.1157

The angle between the oxygens and the hydrogens is not 90 degrees; this is just a Lewis structure; it's actually 120 degrees, because they want to be as far apart as possible.1164

Remember the valence shell-electron pair repulsion?1174

So, that is what is going on here.1177

OK, now let's talk about formal charge.1180

OK, formal charge: this is going to kind of resemble what we did with oxidation-reduction, but it's a little bit different.1185

Formal charge: OK, I'm going to actually start off by drawing three resonance structures for a particular molecule, and then, from there, discussing why we have this idea of formal charge.1196

OK, let me define what it is, really quickly.1211

Formal charge, by definition: it is the valence of the neutral atom...so when we talk about formal charge, we are talking about the formal charge on each atom in a molecule or in an ion.1214

OK, each atom has to have a formal charge.1232

It's the valence of the neutral atom, minus the number of assigned electrons when the atom is bonded.1236

OK, and as far as the number of assigned electrons when the atom is bonded...let me see, Rule 1 is: Lone pairs count as two electrons; each bond counts as 1 electron.1260

We will see what that means in just a minute.1291

Let's do SO42-.1293

OK, so let me just draw a little line here: so we're going to do the sulfate anion, SO42-.1297

I'm going to draw a couple of resonance structures here.1305

I'm just going to go ahead and draw them out; I'm just not going to go through the process--I'll let you confirm that these are viable.1308

There is that; there is that; there is that; there is that; and this is two pairs; this is the most frustrating part for me--writing in all of these lone-pair electrons.1315

2-; there is another resonance structure that has the double bonds horizontal and the single bonds vertical; 3, 4, 5, 6; 1, 2, 3, 4; 1, 2, 3, 4, 5, 6; 1, 2, 3...oh, let me put them down below here.1330

OK, so this is another resonance structure; however, something really, really interesting happens.1350

There is another resonance structure, believe it or not; there are quite a few, in fact.1355

I'm going to write it this way: All single-bonded oxygens (1, 2, 3, 4, 5, 6; 1, 2, 3, 4, 5, 6; 1, 2, 3, 4, 5, 6; 1, 2, 3, 4, 5, 6; 1, 2, 3, 4, 5, 6)...1361

OK, now, notice: this is a perfectly viable resonance structure; this is viable, and this is viable.1375

The reason this is viable is because, remember, sulfur is third-row; anything after the third row has d orbitals actually available to it.1383

So, it can actually carry greater than its octet; it can carry an octet, but if it needs to, it can carry greater than an octet--either as bonds or as lone pairs.1391

There is no rule being violated here.1401

Oxygens all have 8; sulfur has...more than 8: 2, 4, 6, 8, 10, 12; 2, 4, 6, 8...yes, that is correct; this one has...but notice here, this has 2, 4, 6, 8; so sulfur here has 8.1404

It is a perfectly good resonance structure, but it looks different than these.1419

I can...if I want to, I can go ahead and put that; I won't for the time being.1425

Let's just call these two structures equivalent; and the reason that they are equivalent is because they both have a pair of double bonds, two double bonds, two single bonds; two double bonds, two single bonds.1429

The only thing different with these two is the arrangement of the bonds; everything else is the same.1442

Here, you have four single bonds; so it is a resonance structure, because it is perfectly good--that is all a resonance structure has to be; it has to be viable.1449

It is non-equivalent.1458

Now, the question is: when we actually come to deciding what is the actual, real bonding taking place here, is it going to be 4 single bonds, or 2 double bonds and 2 single bonds?1461

This is where the idea of formal charge comes from.1476

What we are going to do is: we are going to actually calculate the formal charge on those three structures, to see if we can discern some sort of a pattern.1482

Well, as it turns out, we will be able to.1490

Let's calculate...well, you know what, I'm not going to write that out; I'll just tell you.1492

We are going to calculate the formal charge on those.1497

Let's do this one: this one...1500

And again, formal charge is on each individual atom.1507

So, the formal charge on sulfur: well, the valence of sulfur is 6 (right?--because it brings 6 electrons--a normal, neutral atom, right?); the formal charge is the valence minus the number of electrons assigned to it when it is bonded.1511

In a normal, neutral, free atom, its valence is 6; we subtract the electrons that are assigned to it in the bonding; well, since it has 4 single bonds, and we said that single bonds count as 1 electron, that is 4.1527

That means its formal charge is +2.1541

Now, oxygen: well, let me draw in the lone pairs on at least one of them (all of the others have the lone pairs, but let me just choose one of them, because they are all equivalent here).1544

The normal valence of oxygen is 6, minus...how many electrons are assigned to it?--well, lone pairs count as 2; bonds count as 1; so we have 2, 4, 6, 7 electrons.1555

In this resonance structure, the formal charge on sulfur is +2; the formal charge on oxygen is -1.1571

OK, now, let's do the formal charge on the other resonance structure.1577

These are equivalent, so I don't have to do both.1584

2-...I forgot the 2- there...1590

Let me go ahead and do it this way.1599

OK, on the sulfur, the valence is 6 (right?--normal valence is 6); we subtract the assigned electrons; well, we said bonds count as 1 electron, so 1, 2, 3, 4, 5, 6.1601

The formal charge on sulfur in this resonance structure is 0.1622

Oxygen: we have two types of oxygen--we have one with a single bond; well, it is 6, minus 1, 2, 3, 4, 5, 6, 7: -1.1627

Here, we have 6, minus 1, 2, 3, 4, 5, 6, equals 0.1642

Notice: in this case, sulfur has a formal charge of +2 and -1; here, you have a formal charge of 0 on sulfur; on two of the oxygens, you have a formal charge of 0, and on two of the other oxygens, you have a formal charge of -1.1649

Here is the rule: OK, atoms in molecules try to achieve formal charges as close to 0 as possible.1664

The second rule is: Any negative formal charge resides on the most electronegative atom.1699

And again, this has to do with minimizing energy--that is the whole idea.1730

All systems, throughout all of physics and all of science--biology and chemistry...all of it is about one thing, ultimately: it is about minimizing energy.1735

That is it: the basic law of science is that any system will seek the lowest energy point; that is it.1744

Here, a formal charge of +2 and -1--well, it's certainly viable, but here, notice, you have: the sulfur has a formal charge of 0; 2 of the oxygens have a formal charge of 0.1753

Now, it's true that the other oxygens have a charge of -1, but oxygen is more electronegative than sulfur, so it can actually carry that charge.1765

Between these two distributions of electrons, this is the one that dominates, because the formal charges on more of the atoms are as close to 0 as possible.1773

Here, it's +2 and -1; here, it's 0, 0, -1.1786

So, as it turns out, even though this resonance structure is a viable structure, and probably contributes a little bit to the overall average, the real bonding of this molecule is actually this.1789

It is those two resonance structures that actually contribute 99.99% of the actual structure of the sulfate anion.1802

In other words, it looks more like this than it does like that.1811

Now again, it is a viable structure, and it probably contributes a little bit, but virtually nothing.1816

So, when we write the resonance structures for SO42-, you really only have to include those two resonance structures; you can actually eliminate that one.1821

You can eliminate it, precisely because formal charge tells you that it is going to end up like this.1829

This is going to be the minimum energy; you are not going to find any molecules in this particular state.1835

You might; I mean, it is quantum mechanics, after all, so you might find some; but more than likely not.1840

That is it--that is what formal charge is: it is a way of...1847

And again, we know what charges are; charges mean the extent to which something has lost an electron or gained an electron; well, these are covalent bonds, so we're talking about polarity.1851

Even though oxygen is highly electronegative, the idea of sulfur actually giving up two electrons is not really likely, because again, sulfur has a reasonably high electronegativity itself.1862

But here, it doesn't have to give up anything; these oxygens don't have to give up anything; there is no output of energy--0.1874

0 is always a number that we seek.1882

Here, yes, it's carrying a negative charge, each of the oxygens; but it is electronegative--it can handle it.1885

This is what we are going to go with.1891

Now, let's just do some examples to round out this lesson of formal charge.1893

And then, because a lot of the problems that you actually get on the AP exam, calculating formal charge...you won't necessarily have to decide between resonance structures, like we did here.1899

This was just for our motivation; you will just have to find the formal charge.1909

Examples: so we'll do CO2 as our first example.1915

OK, and again, we are going to need a Lewis structure for this; so C, O, O; there we go.1920

The normal valence of carbon is 4; how many electrons are assigned to it in CO2?--well, 1, 2, 3, 4 (right? bonds count for one electron); it equals 0.1930

Oxygen: the valence is 6; 1, 2, 3, 4, 5, 6; 1, 2, 3, 4, 5, 6; both oxygens are equivalent; 6 minus 6 equals 0.1942

0 and 0--a very, very stable molecule, by the way--you know that already; CO2.1952

That is it: the formal charge on carbon is 0; the formal charge on oxygen is 0.1957

That is all you are doing--very, very simple.1962

OK, let's do HCn, which is hydrogen cyanide, or hydrocyanic acid if you happen to drop it into water, depending on what you want to call it.1965

The Lewis structure is triple bond, single bond--like that.1973

So now, what is the valence on carbon?--it is 4; how many bonds are assigned to it?--1, 2, 3, 4; so on carbon, the formal charge is 0.1980

On hydrogen--well, hydrogen has a normal valence of 1; it is assigned 1; that is 0.1992

And now, nitrogen: the normal valence of nitrogen is 5, and it is assigned 1, 2, 3, 4, 5 (right?--lone pairs--each electron counts).2001

5 minus 5 equals 0; very nice.2010

How about Cn---how about the cyanide anion?2015

Well, the cyanide anion has a Lewis structure that looks like this (and you are more than welcome to confirm it, if you want to; hopefully you do, because it will give you some practice in drawing Lewis structures); something like that.2018

Carbon has a valence of 4; 1, 2, 3, 4, 5--minus 5 electrons that are assigned to it, for a total formal charge of -1 on carbon.2031

Nitrogen: normal valence is 5; it has 1, 2, 3, 4, 5: minus 5 equals 0.2044

There you go: notice, the sum of the formal charges equals the charge on the species.2052

Let me say that again: The sum of the formal charges of each atom in a molecule equals the total charge on the species.2058

The formal charge on carbon is -1; the formal charge on nitrogen is 0; -1 plus 0 is -1--that happens to be the charge on the species, cyanide.2070

That will always be the case; it was like that with oxidation numbers, too.2079

OK, now, let us try (oh, should I do it on this page; yes)...let's do ozone.2083

Ozone is O3; OK, so we have 18 electrons; as it turns out, ozone looks like this.2093

OK, and we'll do lone pair, lone pair, lone pair, lone pair...and you are absolutely welcome to confirm this if you want to.2106

Well, actually, you know what, this is not a charged species; therefore, I will not put brackets around it; however, it does have another resonance structure.2116

The double bond can be on the other side; this can be single; this can be double; so we get three lone pairs here; one lone pair there; two lone pairs here.2126

Now, we have three different types of oxygen: we have this oxygen, this oxygen, this oxygen.2137

O1: the valence is 6; let's look at this one and call it O1.2142

1, 2, 3, 4, 5, 6, 7; 7 are assigned; it has a formal charge of -1.2150

O2: well...2158

Let me make these electrons a little bit clearer; that looks like 3 electrons there--that is not going to work.2160

OK, now, this one is...normal oxygen is 6; the valence of this is 1, 2, 3, 4, 5: this one is +1.2165

The third oxygen is: normal valence of 6; 1, 2, 3, 4, 5, 6: it is 0.2179

-1, +1, +0, is 0; notice, ozone carries a 0 charge; however, one of these oxygens is carrying a formal charge of +1 (this one), and the other is carrying a formal charge of -1 (this one).2187

This is the complete resonance structure, along with formal charge.2202

OK, so that is it; that is our discussion of resonance and formal charge.2208

Thank you for joining us for AP Chemistry, and thank you for joining us here at Educator.com.2213

We'll see you next time; goodbye.2218

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