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Raffi Hovasapian

Raffi Hovasapian

Spontaneity, Entropy, & Free Energy, Part IV

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 (7)

0 answers

Post by bob singh on April 13, 2015

Professor,

For some reason, around 8:29, I am not getting -1378 kJ for my delta G. My calculations are:
(2(-394)+4(-229))-(-163). I get -1704 kJ.

3 answers

Last reply by: Professor Hovasapian
Sun Apr 21, 2013 8:33 PM

Post by Antie Chen on April 21, 2013

Excuse me, what's the whole name of Q? and in the calculating of Q, liquid and solid won't involve because they don't have pressure?

1 answer

Last reply by: Professor Hovasapian
Wed Sep 26, 2012 11:00 PM

Post by nguyen yen on September 26, 2012

Dear Dr Hovasapian,
I'm confused with the value of R when you talked about RTlnQ and the pressure is expressed in atm
We use 8.31 and in other case it's 0.082
Please explain.

Spontaneity, Entropy, & Free Energy, Part IV

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
  • Spontaneity, Entropy, Free Energy 0:29
    • Standard Free Energy of Formation
    • Example 1
    • Reaction Under Non-standard Conditions
    • Example 2
    • ∆G = Negative
    • ∆G = 0
    • Diagram Example of ∆G

Transcription: Spontaneity, Entropy, & Free Energy, Part IV

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

Today, we are going to continue our discussion of free energy, and we are going to introduce the notion of reactions that take place under non-standard conditions.0004

How do we account for the free energy change that takes place for something that is not one atmosphere, or not one Molar concentration?0014

Well, this is what we are going to do today; and we are going to introduce a very, very, very important equation.0022

So, let us go ahead and get started.0028

In the last example of the last lesson, we used that ΔG=ΔH-TΔS to calculate the free energy of a given reaction.0031

Well, before we actually introduce this notion of non-standard conditions, let's just talk a little bit about the other way that we can actually get ΔG, instead of using the ΔH-TΔS.0043

Let's say--well, let's just start with a definition.0058

OK, definition: The standard free energy of formation, which is symbolized as ΔG with a little f down at the bottom, is the change in free energy that accompanies the formation of one mole of a substance from its constituent elements with all reactants and products in their standard states.0063

There you go: so, for example, the ΔG of formation of, let's say, water: H2 gas + O2 gas goes to 1 mole (right, 1 mole of a substance, so) of H2O liquid.0162

If we want to calculate the standard entropy change for that, we would do whatever it is that chemists and physicists do, and this would actually end up being one-half here (right? no, the other way around--what am I doing?--this is H2...oh, that's nice; I love that...plus 1/2 O2, goes to H2O).0180

There is some ΔG of formation associated with that, some number.0206

This number, the ΔG of formation--this is the number that you find at the end of thermodynamic tables.0212

This definition is actually exactly like the definition for enthalpy, if you remember, from way back then when we did thermochemistry.0222

There was something called the standard heat of formation (the standard enthalpy of formation), and it was the change in enthalpy that accompanies the formation of one mole of that particular substance (that we want the enthalpy of formation for), under standard conditions, with everything being in their standard states.0229

This definition is completely analogous: the idea here is that, instead of using the ΔH=ΔG-TΔS, we can actually use the equation as written, and just take the ΔG of the products, minus the ΔG of the reactants, and get the ΔG of the reaction.0249

We actually end up getting the same number; so we have tables, or we have the equation ΔH-TΔS.0267

So, let's just do a quick example here.0274

Example: What is the ΔG for 2 CH3OH + 3 O2 (in other words, the combustion of methanol gas); this is gas; this is gas; 2 CO2 gas plus 4 H2O.0279

We want to know what the standard free energy change is for this.0310

OK, now the entries...I'm going to actually write out the entries that you would see in a thermodynamic table, just so that you actually see it; I know that you can flip to the back of your book, but I would like you to see the numbers anyway.0315

The entry for...well, actually, I'm going to do the entry for all of them: so let's go to CH3OH as a gas; O2 as a gas; CO2 as a gas, and (oops, here we go with the stray lines again; we don't want that) H2O gas.0325

Now, we have entries that look like this: there is going to be something that says ΔHformation, and it's going to say kilojoules per mole.0358

And then, there is going to be an entry that says ΔG of formation, and that is in kilojoules per mole.0367

And then, they are going to have S, which is going to be in Joules per mole-Kelvin.0373

You know what, I really, really, really need to make this a little bit more legible; I understand--my apologies.0381

S--and notice, there is no ΔS; this is just S--this is going to be in Joules per mole-Kelvin.0388

We have: -201, 0 (remember, 0--elements, the enthalpy is 0), and we have -394, -242; the ΔG is -163, and the ΔG for elements is also 0; this happens to also be 394, which is really interesting, in and of itself; -229.0398

And then, we have 240, 205, 214, and 189.0426

The entropy of an element is not 0: oxygen gas is a disordered system; there are a lot of particles of gas: 205--that is a lot of entropy.0435

It's not 0; it's not the same as these.0443

When I run this--when I take 4 times the ΔG of H2O (so I'm working in this column right here--I just wanted you to see what the other entries are), plus 2 times the ΔG of CO2, minus 2 times the ΔG of CH3OH gas, minus 3 times the ΔG of O2...when I actually run the reaction, I end up with (I'll run the calculation) a ΔG.0447

Notice, the f is gone; those were ΔG's of formation of the individual elements; that is what is in the table; the final outcome is just my ΔG, my free energy.0477

And it's not in moles; it's not in kilojoules per mole; it's just in kilojoules.0488

And the reason is because I have accounted for the moles from the balanced reaction: I am taking 2 times something, 3 times something--I am using the stoichiometric coefficients.0492

I end up with -1378 kilojoules--137,800 Joules.0502

That is a lot of energy; that reaction--the burning of methanol--is highly spontaneous under standard conditions.0515

So, that is what this means.0526

Let's see: OK, so now, we have a good sense of dealing with free energy; we treat it the same way we treat ΔH, or H, or S--we just take products minus reactants.0530

We also have the equation ΔH-TΔS, to calculate it that way.0542

Now, let's think about what this actually means.0547

This number here tells me that it's spontaneous as written: that means it will happen (eventually) if I don't do anything.0551

That doesn't mean it's going to happen quickly; it just means that it will happen, and I don't have to do anything about it.0559

Thermodynamically, it's spontaneous.0565

Spontaneity is a thermodynamic statement, not a kinetic statement; it just says that the energy (from your perspective) of the CO2 and the water is lower than the energy of the methanol and the oxygen gas.0567

So, it's a downhill thing.0587

We just needed to get it over that hump; that is the kinetics part.0590

OK, so now let's think about what this means: this ΔG is a measure of the tendency, the potential, for a reaction to move forward.0592

Well, let's say it starts moving forward: well, as it starts to move forward, more CO2 and water are going to form; methanol is going to be used up; oxygen is going to be used up.0606

Let's say that half the reaction has gone forward: does that mean, at that point, that the free energy is the same?0619

As it turns out, no: the idea is that free energy wants to get down to 0--or, in this case, wants to get up to 0.0626

0 free energy means the system is at equilibrium; so, you can have a negative free energy; that means that the reaction, as written, wants to go to the right.0635

It's spontaneous to the right.0648

If it's positive, that means the reaction wants to go to the left.0650

The idea is: the reaction wants to get to a point where the ΔG is 0.0655

So, as a reaction proceeds forward, the ΔG actually changes until it gets to a point where it's 0; that is equilibrium--that is what all reactions do--they move toward equilibrium.0660

So, let me write this down: As a reaction proceeds toward equilibrium, the ΔG changes until it hits equilibrium--until it reaches 0, which is the equilibrium point.0675

OK, now, we have been calculating ΔG with that little 0 on top--standard ΔG's under standard conditions (1 atmosphere pressure; 25 degrees Celsius; 1 Molar concentration for aqueous species--standard conditions).0723

Well, what happens when we run a reaction not under standard conditions--what if it's at 200 degrees Celsius; what if it's 3 atmospheres of one gas and 17 atmospheres of another gas?0746

What happens then--how does it affect the ΔG?0761

Is it affected at all?0764

As a matter of fact, it is; so, when we change the conditions, when we change the pressures, the temperatures, the ΔG of the reaction as written changes.0766

Rather than going through a sort of discussion of a derivation of this equation that I'm going to write down, I'm just going to write down the equation, because we just want you to be able to use it and understand that this ΔG is under standard conditions, but we're not always running reactions under standard conditions.0777

We often...most of the time, we're not doing it under standard conditions.0799

So, for reactions run under non-standard conditions, the ΔG of the reaction equals the standard ΔG, plus a certain term--a correction factor, if you will, for the non-standard part.0803

That is the whole point; at standard conditions, it's this; but if it's non-standard, then we're going to have to add a little term--either plus or minus, one way or the other--to adjust for the fact that it's non-standard.0838

That is RT ln(Q), and Q is exactly what you think it is: it's the reaction quotient.0849

If I have (let me do this one in red...I thought I said red...let's see) aA + bB going to cC + dD, the reaction quotient is equal to the concentration of C, or the pressure of C (depending on if it's aqueous or gas), raised to the power of c, times the concentration of D raised to the power of d, over A raised to the power of a, B raised to the power of b.0856

At any given moment, under those circumstances, we use Q for at that moment; we use K for equilibrium concentrations.0896

So, that is what this is; this is a very, very important equation.0905

It is telling me that, if I am going to run a reaction under non-standard conditions (let's say higher pressure--and we will deal with pressure mostly with our discussions), it is equal to...well, if I want to know what the free energy change of a reaction is, at a non-standard...I go ahead and I calculate the free energy change under standard conditions, and then I add to it R (which is the gas constant--I'll write that down: R=8.3145--that is Joules per mole-Kelvin), times the temperature (which is the absolute temperature in Kelvin), times the logarithm of the reaction quotient.0909

The reaction quotient is this thing, as stated in the particular problem.0950

Now, let's just go ahead and do an example.0959

Oh, and by the way, notice: ΔG and R--R is Joules; ΔG...when you calculate ΔG, more often than not, you will calculate it in kilojoules, depending on if you are going to use the thermodynamic data in the back or the equation.0962

Make sure your units match, OK?--so Joules, Joules or kilojoules, kilojoules: so just make sure that the units match.0977

OK, so let's go ahead and do our example.0986

Example: Calculate (and notice, there is no 0--there is no little degree sign on top, because it is not standard anymore; this is the ΔG of a reaction) ΔG of the reaction at 25 degrees Celsius (so, in this case, at least the temperature will be) for the reaction CO (carbon monoxide gas), plus 2 H2 gas, going to CH3OH liquid.0990

We want to calculate the ΔG of the reaction, at 25 degrees Celsius, for the formation of methanol from carbon monoxide gas and hydrogen gas, when the partial pressure of the CO gas is equal to 6 atmospheres (so notice, it's not 1 atmosphere anymore), and the partial pressure of the H2 gas equals 2.0 atmospheres (not one atmosphere anymore).1035

So now, the pressures are higher.1059

This is a gas; this is a gas; this is a liquid; so--this is a liquid--it doesn't show up in the expression for the reaction quotient.1062

Let's go ahead and write this out (let me actually write the equation again; let me do it in blue).1071

We have: carbon monoxide gas, plus two moles of hydrogen gas, is going to form one mole of methanol liquid.1079

OK, so we have that the ΔG of the reaction is equal to the standard free energy change, plus RT ln(Q).1091

Well, let's see what Q is, first: Q is equal to the concentration of the products, divided by the concentration of the reactants.1103

In this case, these are gases, so we are going to use pressures; this is a liquid, so it doesn't show up in the numerator (it's just 1 up there).1111

The partial pressure of CO2, times the partial pressure of H2, squared (the stoichiometric coefficient: you know this already--that is what this is--just the reaction quotient under gaseous conditions).1121

OK, so now, let's go ahead and calculate: we have taken care of this; we'll plug it in in just a minute.1134

We know what T is: it's going to be 298; we know what R is--that is a constant; logarithm is just a mathematical operation; we need to calculate what ΔG is.1142

ΔG, standard free energy change--when I look at a thermodynamic table, I'm going to end up with...I'm going to do this one, minus 166 kilojoules (I'm going to...actually, let me do it this way); it's 1, times 1, times -166, minus 1 times -137, plus 2 times 0 (right?--we use the standard--something that we have done all along).1152

You end up with a ΔG of -29 kilojoules, which is equal to 29,000 Joules.1193

Good; so now, we have: our ΔG of reaction is equal to -29,000, plus 8.3145 Joules per Kelvin, times 298 Kelvin, times the logarithm of 1, over...well, what was the partial pressure of CO?--it was 6.0 atmospheres; the partial pressure of the hydrogen was 2.0 atmospheres squared.1203

When we do all of the math, we end up with -29,000, plus a -7,874; I wanted you to see it in both of its forms.1245

Under standard conditions, we just calculated that ΔG is 29,000; now, to that, based on the fact that the pressures are actually higher, this equation tells me that it's even more spontaneous, because this number is negative.1263

The total is going to be -36,874 Joules; so the reaction as written is even more spontaneous under conditions of higher pressure.1278

That is all this is: we use this equation to find the free energy change of a reaction under conditions that are non-standard.1294

That is it; that is all that this equation allows us to do.1304

Based on the reaction quotient, that means as written, I just throw into a flask CO gas at 6 atmospheres and hydrogen gas at 2 atmospheres; the free energy change for this reaction is going to be -36,874.1307

This is highly spontaneous: it has a tendency--it's going to start to move toward equilibrium; that is the whole idea.1322

OK, so let's see what else: we want to notice a couple of things about this.1330

Notice: ΔG is more negative, so it's more spontaneous, under conditions of higher pressure (or, I should say, under conditions of pressures higher than standard--pressures higher than 1 atmosphere).1337

But we could have predicted this from Le Chatelier's Principle: watch.1379

CO (now again, prediction is one thing: we are actually able to get a number, so--qualitative: we could predict it qualitatively, but we want to be able to do the quantitative; we want a number for it) + 2 H2 goes to CH3OH liquid.1383

Here, we have three particles of gas; here, we have no particles of gas.1405

Well, under 1 atmosphere of pressure, it's at ΔG; now, I have 6 atmospheres of CO; I have 2 atmospheres of the H2; I have increased the pressure of the system.1412

If I increase the pressure of the system, well, the system is going to want to offset that pressure, right?1424

That is the whole idea: it's Le Chatelier's Principle--it's going to do the opposite of--it's going to try to get itself back to where it was, to offset the effect.1431

Increasing the pressure, the system is going to want to decrease the pressure; how does it decrease pressure?--well, it decreases pressure by lowering the number of particles that are bouncing around against the walls to create that pressure, which means it's going to move in that direction.1439

It is going to be more spontaneous; it is going to push the reaction forward.1455

You see, all of these things are coming together; that is the whole idea behind chemistry--we want to get you to see the big picture: spontaneity, equilibrium, Le Chatelier's Principle, quantitative aspects, free energy, entropy--all of this is tied together.1459

OK, now, the last thing that we are going to talk about is the following.1476

ΔG=0: we said that, when the ΔG of a reaction (either at standard or non-standard conditions)--when it equals 0, that means equilibrium.1482

It does not mean completion; and here is what we mean by that.1502

If I start with some CO gas and some H2 gas, and here I have the CH3OH, energetically, the free energy change is negative, right?1509

So, the ΔG is less than 0; it's going from a higher energy to a lower energy--that is what makes this spontaneous--but here is what happens.1524

As the CO and the H2 are used up (let me do this in red), they are going to diminish; OK, there is going to be less of this, more of this.1533

I'm going to come up; it's going to come to a point where there is going to be a mixture of the CO, the H2, and the CH3OH, when there is an equality.1547

It doesn't mean ("spontaneous")--this negative 36,000 that we got doesn't say that, if I put this in a flask and somehow the reaction takes place, that all of a sudden all of the carbon monoxide and the H2 are going to vanish, and the only thing left in the flask is going to be the methanol.1563

That is not what this is saying.1582

"Spontaneous" is a measure of the tendency of a system to reach equilibrium, not completion.1584

When we say "completion," that means there are no more reactants left.1591

When we say "equilibrium," that means it has reached a point where there is a little bit of everything left: that is what ΔG means.1595

So, pictorially, it looks like this.1603

In some sense (again, I'm not sure if this is the best way to represent it, but it's not a bad way of representing it): ΔG=0, equilibrium--it represents a low point in energy.1612

That means...over here, the ΔG is negative; it's spontaneous in this direction; it is going to seek out this--in other words, it's going to form more.1634

But, if it went to completion...as it turns out, at that point, as written, it is going to end up with a positive ΔG.1646

Well, a positive ΔG means it's spontaneous in the reverse direction; so now, the reaction is going to go this way.1658

It is going to go this way and this way until it reaches this point, where the ΔG of the reaction is equal to 0.1664

It is going to seek out equilibrium; it's not going to seek out completion.1673

Let me say that again: Reactions...ΔG going to 0...reactions will seek equilibrium; they will not seek out completion.1677

That doesn't mean that there are going to be no reactants left, and that it's going to be all products.1687

That is not what "spontaneous" talks about; "spontaneous" is a measure of the tendency of a reaction to seek out equilibrium.1691

It wants to get to a point where ΔG equals 0; it's like a ball rolling down a hill--it will roll down a hill, it will roll up the other side, and it will roll back down, roll up this side, until it finds a point of lowest energy.1699

Reactions move toward the valley of lowest free energy (zero).1714

I shouldn't say "lowest free energy"--zero, because negative free energy implies that it is actually moving toward that 0.1720

This is what you want to remember: Completion is not the same as equilibrium.1727

Yes, there are some reactions...like, for example, if you put hydrogen and oxygen in a flask and you ignite it--yes, it is going to be all water; there is going to be no hydrogen and no oxygen in there--at least, none that is measurable.1733

But, believe it or not...there is none that is practically measurable, but believe it or not, it actually is in equilibrium.1744

It has gone so far--there is so little hydrogen and oxygen gas left--that, for all practical purposes, yes, that reaction has come to completion.1753

In that case, completion and equilibrium are almost the same.1759

But, that doesn't mean that they are; the system is actually at equilibrium--there is still just a little bit of hydrogen and oxygen left, because the free energy for that reaction is here.1764

It is not going to end up going to form all water or all something else.1775

Think about that for a little bit.1780

OK, next time we get together, we are actually going to do a series of problems, because I know we have done only a handful of problems for the thermodynamic section.1782

We spent a lot of time discussing it, but we definitely need to wrap our minds around it by doing more practice.1790

So, the next time I see you, we are going to introduce one more concept of equilibrium, and then we are going to spend our time on some problems.1797

Thank you for joining us here at Educator.com.1803

See you next time; goodbye.1806

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