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Franklin Ow

Franklin Ow

Entropy & Free Energy

Slide Duration:

Table of Contents

I. Basic Concepts & Measurement of Chemistry
Basic Concepts of Chemistry

16m 26s

Intro
0:00
Lesson Overview
0:07
Introduction
0:56
What is Chemistry?
0:57
What is Matter?
1:16
Solids
1:43
General Characteristics
1:44
Particulate-level Drawing of Solids
2:34
Liquids
3:39
General Characteristics of Liquids
3:40
Particulate-level Drawing of Liquids
3:55
Gases
4:23
General Characteristics of Gases
4:24
Particulate-level Drawing Gases
5:05
Classification of Matter
5:27
Classification of Matter
5:26
Pure Substances
5:54
Pure Substances
5:55
Mixtures
7:06
Definition of Mixtures
7:07
Homogeneous Mixtures
7:11
Heterogeneous Mixtures
7:52
Physical and Chemical Changes/Properties
8:18
Physical Changes Retain Chemical Composition
8:19
Chemical Changes Alter Chemical Composition
9:32
Physical and Chemical Changes/Properties, cont'd
10:55
Physical Properties
10:56
Chemical Properties
11:42
Sample Problem 1: Chemical & Physical Change
12:22
Sample Problem 2: Element, Compound, or Mixture?
13:52
Sample Problem 3: Classify Each of the Following Properties as chemical or Physical
15:03
Tools in Quantitative Chemistry

29m 22s

Intro
0:00
Lesson Overview
0:07
Units of Measurement
1:23
The International System of Units (SI): Mass, Length, and Volume
1:39
Percent Error
2:17
Percent Error
2:18
Example: Calculate the Percent Error
2:56
Standard Deviation
3:48
Standard Deviation Formula
3:49
Standard Deviation cont'd
4:42
Example: Calculate Your Standard Deviation
4:43
Precisions vs. Accuracy
6:25
Precision
6:26
Accuracy
7:01
Significant Figures and Uncertainty
7:50
Consider the Following (2) Rulers
7:51
Consider the Following Graduated Cylinder
11:30
Identifying Significant Figures
12:43
The Rules of Sig Figs Overview
12:44
The Rules for Sig Figs: All Nonzero Digits Are Significant
13:21
The Rules for Sig Figs: A Zero is Significant When It is In-Between Nonzero Digits
13:28
The Rules for Sig Figs: A Zero is Significant When at the End of a Decimal Number
14:02
The Rules for Sig Figs: A Zero is not significant When Starting a Decimal Number
14:27
Using Sig Figs in Calculations
15:03
Using Sig Figs for Multiplication and Division
15:04
Using Sig Figs for Addition and Subtraction
15:48
Using Sig Figs for Mixed Operations
16:11
Dimensional Analysis
16:20
Dimensional Analysis Overview
16:21
General Format for Dimensional Analysis
16:39
Example: How Many Miles are in 17 Laps?
17:17
Example: How Many Grams are in 1.22 Pounds?
18:40
Dimensional Analysis cont'd
19:43
Example: How Much is Spent on Diapers in One Week?
19:44
Dimensional Analysis cont'd
21:03
SI Prefixes
21:04
Dimensional Analysis cont'd
22:03
500 mg → ? kg
22:04
34.1 cm → ? um
24:03
Summary
25:11
Sample Problem 1: Dimensional Analysis
26:09
II. Atoms, Molecules, and Ions
Atoms, Molecules, and Ions

52m 18s

Intro
0:00
Lesson Overview
0:08
Introduction to Atomic Structure
1:03
Introduction to Atomic Structure
1:04
Plum Pudding Model
1:26
Introduction to Atomic Structure Cont'd
2:07
John Dalton's Atomic Theory: Number 1
2:22
John Dalton's Atomic Theory: Number 2
2:50
John Dalton's Atomic Theory: Number 3
3:07
John Dalton's Atomic Theory: Number 4
3:30
John Dalton's Atomic Theory: Number 5
3:58
Introduction to Atomic Structure Cont'd
5:21
Ernest Rutherford's Gold Foil Experiment
5:22
Introduction to Atomic Structure Cont'd
7:42
Implications of the Gold Foil Experiment
7:43
Relative Masses and Charges
8:18
Isotopes
9:02
Isotopes
9:03
Introduction to The Periodic Table
12:17
The Periodic Table of the Elements
12:18
Periodic Table, cont'd
13:56
Metals
13:57
Nonmetals
14:25
Semimetals
14:51
Periodic Table, cont'd
15:57
Group I: The Alkali Metals
15:58
Group II: The Alkali Earth Metals
16:25
Group VII: The Halogens
16:40
Group VIII: The Noble Gases
17:08
Ionic Compounds: Formulas, Names, Props.
17:35
Common Polyatomic Ions
17:36
Predicting Ionic Charge for Main Group Elements
18:52
Ionic Compounds: Formulas, Names, Props.
20:36
Naming Ionic Compounds: Rule 1
20:51
Naming Ionic Compounds: Rule 2
21:22
Naming Ionic Compounds: Rule 3
21:50
Naming Ionic Compounds: Rule 4
22:22
Ionic Compounds: Formulas, Names, Props.
22:50
Naming Ionic Compounds Example: Al₂O₃
22:51
Naming Ionic Compounds Example: FeCl₃
23:21
Naming Ionic Compounds Example: CuI₂ 3H₂O
24:00
Naming Ionic Compounds Example: Barium Phosphide
24:40
Naming Ionic Compounds Example: Ammonium Phosphate
25:55
Molecular Compounds: Formulas and Names
26:42
Molecular Compounds: Formulas and Names
26:43
The Mole
28:10
The Mole is 'A Chemist's Dozen'
28:11
It is a Central Unit, Connecting the Following Quantities
30:01
The Mole, cont'd
32:07
Atomic Masses
32:08
Example: How Many Moles are in 25.7 Grams of Sodium?
32:28
Example: How Many Atoms are in 1.2 Moles of Carbon?
33:17
The Mole, cont'd
34:25
Example: What is the Molar Mass of Carbon Dioxide?
34:26
Example: How Many Grams are in 1.2 Moles of Carbon Dioxide?
25:46
Percentage Composition
36:43
Example: How Many Grams of Carbon Contained in 65.1 Grams of Carbon Dioxide?
36:44
Empirical and Molecular Formulas
39:19
Empirical Formulas
39:20
Empirical Formula & Elemental Analysis
40:21
Empirical and Molecular Formulas, cont'd
41:24
Example: Determine Both the Empirical and Molecular Formulas - Step 1
41:25
Example: Determine Both the Empirical and Molecular Formulas - Step 2
43:18
Summary
46:22
Sample Problem 1: Determine the Empirical Formula of Lithium Fluoride
47:10
Sample Problem 2: How Many Atoms of Carbon are Present in 2.67 kg of C₆H₆?
49:21
III. Chemical Reactions
Chemical Reactions

43m 24s

Intro
0:00
Lesson Overview
0:06
The Law of Conservation of Mass and Balancing Chemical Reactions
1:49
The Law of Conservation of Mass
1:50
Balancing Chemical Reactions
2:50
Balancing Chemical Reactions Cont'd
3:40
Balance: N₂ + H₂ → NH₃
3:41
Balance: CH₄ + O₂ → CO₂ + H₂O
7:20
Balancing Chemical Reactions Cont'd
9:49
Balance: C₂H₆ + O₂ → CO₂ + H₂O
9:50
Intro to Chemical Equilibrium
15:32
When an Ionic Compound Full Dissociates
15:33
When an Ionic Compound Incompletely Dissociates
16:14
Dynamic Equilibrium
17:12
Electrolytes and Nonelectrolytes
18:03
Electrolytes
18:04
Strong Electrolytes and Weak Electrolytes
18:55
Nonelectrolytes
19:23
Predicting the Product(s) of an Aqueous Reaction
20:02
Single-replacement
20:03
Example: Li (s) + CuCl₂ (aq) → 2 LiCl (aq) + Cu (s)
21:03
Example: Cu (s) + LiCl (aq) → NR
21:23
Example: Zn (s) + 2HCl (aq) → ZnCl₂ (aq) + H₂ (g)
22:32
Predicting the Product(s) of an Aqueous Reaction
23:37
Double-replacement
23:38
Net-ionic Equation
25:29
Predicting the Product(s) of an Aqueous Reaction
26:12
Solubility Rules for Ionic Compounds
26:13
Predicting the Product(s) of an Aqueous Reaction
28:10
Neutralization Reactions
28:11
Example: HCl (aq) + NaOH (aq) → ?
28:37
Example: H₂SO₄ (aq) + KOH (aq) → ?
29:25
Predicting the Product(s) of an Aqueous Reaction
30:20
Certain Aqueous Reactions can Produce Unstable Compounds
30:21
Example 1
30:52
Example 2
32:16
Example 3
32:54
Summary
33:54
Sample Problem 1
34:55
ZnCO₃ (aq) + H₂SO₄ (aq) → ?
35:09
NH₄Br (aq) + Pb(C₂H₃O₂)₂ (aq) → ?
36:02
KNO₃ (aq) + CuCl₂ (aq) → ?
37:07
Li₂SO₄ (aq) + AgNO₃ (aq) → ?
37:52
Sample Problem 2
39:09
Question 1
39:10
Question 2
40:36
Question 3
41:47
Chemical Reactions II

55m 40s

Intro
0:00
Lesson Overview
0:10
Arrhenius Definition
1:15
Arrhenius Acids
1:16
Arrhenius Bases
3:20
The Bronsted-Lowry Definition
4:48
Acids Dissolve In Water and Donate a Proton to Water: Example 1
4:49
Acids Dissolve In Water and Donate a Proton to Water: Example 2
6:54
Monoprotic Acids & Polyprotic Acids
7:58
Strong Acids
11:30
Bases Dissolve In Water and Accept a Proton From Water
12:41
Strong Bases
16:36
The Autoionization of Water
17:42
Amphiprotic
17:43
Water Reacts With Itself
18:24
Oxides of Metals and Nonmetals
20:08
Oxides of Metals and Nonmetals Overview
20:09
Oxides of Nonmetals: Acidic Oxides
21:23
Oxides of Metals: Basic Oxides
24:08
Oxidation-Reduction (Redox) Reactions
25:34
Redox Reaction Overview
25:35
Oxidizing and Reducing Agents
27:02
Redox Reaction: Transfer of Electrons
27:54
Oxidation-Reduction Reactions Cont'd
29:55
Oxidation Number Overview
29:56
Oxidation Number of Homonuclear Species
31:17
Oxidation Number of Monatomic Ions
32:58
Oxidation Number of Fluorine
33:27
Oxidation Number of Oxygen
34:00
Oxidation Number of Chlorine, Bromine, and Iodine
35:07
Oxidation Number of Hydrogen
35:30
Net Sum of All Oxidation Numbers In a Compound
36:21
Oxidation-Reduction Reactions Cont'd
38:19
Let's Practice Assigning Oxidation Number
38:20
Now Let's Apply This to a Chemical Reaction
41:07
Summary
44:19
Sample Problems
45:29
Sample Problem 1
45:30
Sample Problem 2: Determine the Oxidizing and Reducing Agents
48:48
Sample Problem 3: Determine the Oxidizing and Reducing Agents
50:43
IV. Stoichiometry
Stoichiometry I

42m 10s

Intro
0:00
Lesson Overview
0:23
Mole to Mole Ratios
1:32
Example 1: In 1 Mole of H₂O, How Many Moles Are There of Each Element?
1:53
Example 2: In 2.6 Moles of Water, How Many Moles Are There of Each Element?
2:24
Mole to Mole Ratios Cont'd
5:13
Balanced Chemical Reaction
5:14
Mole to Mole Ratios Cont'd
7:25
Example 3: How Many Moles of Ammonia Can Form If you Have 3.1 Moles of H₂?
7:26
Example 4: How Many Moles of Hydrogen Gas Are Required to React With 6.4 Moles of Nitrogen Gas?
9:08
Mass to mass Conversion
11:06
Mass to mass Conversion
11:07
Example 5: How Many Grams of Ammonia Can Form If You Have 3.1 Grams of H₂?
12:37
Example 6: How Many Grams of Hydrogen Gas Are Required to React With 6.4 Grams of Nitrogen Gas?
15:34
Example 7: How Man Milligrams of Ammonia Can Form If You Have 1.2 kg of H₂?
17:29
Limiting Reactants, Percent Yields
20:42
Limiting Reactants, Percent Yields
20:43
Example 8: How Many Grams of Ammonia Can Form If You Have 3.1 Grams of H₂ and 3.1 Grams of N₂
22:25
Percent Yield
25:30
Example 9: How Many Grams of The Excess Reactant Remains?
26:37
Summary
29:34
Sample Problem 1: How Many Grams of Carbon Are In 2.2 Kilograms of Carbon Dioxide?
30:47
Sample Problem 2: How Many Milligrams of Carbon Dioxide Can Form From 23.1 Kg of CH₄(g)?
33:06
Sample Problem 3: Part 1
36:10
Sample Problem 3: Part 2 - What Amount Of The Excess Reactant Will Remain?
40:53
Stoichiometry II

42m 38s

Intro
0:00
Lesson Overview
0:10
Molarity
1:14
Solute and Solvent
1:15
Molarity
2:01
Molarity Cont'd
2:59
Example 1: How Many Grams of KBr are Needed to Make 350 mL of a 0.67 M KBr Solution?
3:00
Example 2: How Many Moles of KBr are in 350 mL of a 0.67 M KBr Solution?
5:44
Example 3: What Volume of a 0.67 M KBr Solution Contains 250 mg of KBr?
7:46
Dilutions
10:01
Dilution: M₁V₂=M₁V₂
10:02
Example 5: Explain How to Make 250 mL of a 0.67 M KBr Solution Starting From a 1.2M Stock Solution
12:04
Stoichiometry and Double-Displacement Precipitation Reactions
14:41
Example 6: How Many grams of PbCl₂ Can Form From 250 mL of 0.32 M NaCl?
15:38
Stoichiometry and Double-Displacement Precipitation Reactions
18:05
Example 7: How Many grams of PbCl₂ Can Form When 250 mL of 0.32 M NaCl and 150 mL of 0.45 Pb(NO₃)₂ Mix?
18:06
Stoichiometry and Neutralization Reactions
21:01
Example 8: How Many Grams of NaOh are Required to Neutralize 4.5 Grams of HCl?
21:02
Stoichiometry and Neutralization Reactions
23:03
Example 9: How Many mL of 0.45 M NaOH are Required to Neutralize 250 mL of 0.89 M HCl?
23:04
Stoichiometry and Acid-Base Standardization
25:28
Introduction to Titration & Standardization
25:30
Acid-Base Titration
26:12
The Analyte & Titrant
26:24
The Experimental Setup
26:49
The Experimental Setup
26:50
Stoichiometry and Acid-Base Standardization
28:38
Example 9: Determine the Concentration of the Analyte
28:39
Summary
32:46
Sample Problem 1: Stoichiometry & Neutralization
35:24
Sample Problem 2: Stoichiometry
37:50
V. Thermochemistry
Energy & Chemical Reactions

55m 28s

Intro
0:00
Lesson Overview
0:14
Introduction
1:22
Recall: Chemistry
1:23
Energy Can Be Expressed In Different Units
1:57
The First Law of Thermodynamics
2:43
Internal Energy
2:44
The First Law of Thermodynamics Cont'd
6:14
Ways to Transfer Internal Energy
6:15
Work Energy
8:13
Heat Energy
8:34
∆U = q + w
8:44
Calculating ∆U, Q, and W
8:58
Changes In Both Volume and Temperature of a System
8:59
Calculating ∆U, Q, and W Cont'd
11:01
The Work Equation
11:02
Example 1: Calculate ∆U For The Burning Fuel
11:45
Calculating ∆U, Q, and W Cont'd
14:09
The Heat Equation
14:10
Calculating ∆U, Q, and W Cont'd
16:03
Example 2: Calculate The Final Temperature
16:04
Constant-Volume Calorimetry
18:05
Bomb Calorimeter
18:06
The Effect of Constant Volume On The Equation For Internal Energy
22:11
Example 3: Calculate ∆U
23:12
Constant-Pressure Conditions
26:05
Constant-Pressure Conditions
26:06
Calculating Enthalpy: Phase Changes
27:29
Melting, Vaporization, and Sublimation
27:30
Freezing, Condensation and Deposition
28:25
Enthalpy Values For Phase Changes
28:40
Example 4: How Much Energy In The Form of heat is Required to Melt 1.36 Grams of Ice?
29:40
Calculating Enthalpy: Heats of Reaction
31:22
Example 5: Calculate The Heat In kJ Associated With The Complete Reaction of 155 g NH₃
31:23
Using Standard Enthalpies of Formation
33:53
Standard Enthalpies of Formation
33:54
Using Standard Enthalpies of Formation
36:12
Example 6: Calculate The Standard Enthalpies of Formation For The Following Reaction
36:13
Enthalpy From a Series of Reactions
39:58
Hess's Law
39:59
Coffee-Cup Calorimetry
42:43
Coffee-Cup Calorimetry
42:44
Example 7: Calculate ∆H° of Reaction
45:10
Summary
47:12
Sample Problem 1
48:58
Sample Problem 2
51:24
VI. Quantum Theory of Atoms
Structure of Atoms

42m 33s

Intro
0:00
Lesson Overview
0:07
Introduction
1:01
Rutherford's Gold Foil Experiment
1:02
Electromagnetic Radiation
2:31
Radiation
2:32
Three Parameters: Energy, Frequency, and Wavelength
2:52
Electromagnetic Radiation
5:18
The Electromagnetic Spectrum
5:19
Atomic Spectroscopy and The Bohr Model
7:46
Wavelengths of Light
7:47
Atomic Spectroscopy Cont'd
9:45
The Bohr Model
9:46
Atomic Spectroscopy Cont'd
12:21
The Balmer Series
12:22
Rydberg Equation For Predicting The Wavelengths of Light
13:04
The Wave Nature of Matter
15:11
The Wave Nature of Matter
15:12
The Wave Nature of Matter
19:10
New School of Thought
19:11
Einstein: Energy
19:49
Hertz and Planck: Photoelectric Effect
20:16
de Broglie: Wavelength of a Moving Particle
21:14
Quantum Mechanics and The Atom
22:15
Heisenberg: Uncertainty Principle
22:16
Schrodinger: Wavefunctions
23:08
Quantum Mechanics and The Atom
24:02
Principle Quantum Number
24:03
Angular Momentum Quantum Number
25:06
Magnetic Quantum Number
26:27
Spin Quantum Number
28:42
The Shapes of Atomic Orbitals
29:15
Radial Wave Function
29:16
Probability Distribution Function
32:08
The Shapes of Atomic Orbitals
34:02
3-Dimensional Space of Wavefunctions
34:03
Summary
35:57
Sample Problem 1
37:07
Sample Problem 2
40:23
VII. Electron Configurations and Periodicity
Periodic Trends

38m 50s

Intro
0:00
Lesson Overview
0:09
Introduction
0:36
Electron Configuration of Atoms
1:33
Electron Configuration & Atom's Electrons
1:34
Electron Configuration Format
1:56
Electron Configuration of Atoms Cont'd
3:01
Aufbau Principle
3:02
Electron Configuration of Atoms Cont'd
6:53
Electron Configuration Format 1: Li, O, and Cl
6:56
Electron Configuration Format 2: Li, O, and Cl
9:11
Electron Configuration of Atoms Cont'd
12:48
Orbital Box Diagrams
12:49
Pauli Exclusion Principle
13:11
Hund's Rule
13:36
Electron Configuration of Atoms Cont'd
17:35
Exceptions to The Aufbau Principle: Cr
17:36
Exceptions to The Aufbau Principle: Cu
18:15
Electron Configuration of Atoms Cont'd
20:22
Electron Configuration of Monatomic Ions: Al
20:23
Electron Configuration of Monatomic Ions: Al³⁺
20:46
Electron Configuration of Monatomic Ions: Cl
21:57
Electron Configuration of Monatomic Ions: Cl¹⁻
22:09
Electron Configuration Cont'd
24:31
Paramagnetism
24:32
Diamagnetism
25:00
Atomic Radii
26:08
Atomic Radii
26:09
In a Column of the Periodic Table
26:25
In a Row of the Periodic Table
26:46
Ionic Radii
27:30
Ionic Radii
27:31
Anions
27:42
Cations
27:57
Isoelectronic Species
28:12
Ionization Energy
29:00
Ionization Energy
29:01
Electron Affinity
31:37
Electron Affinity
31:37
Summary
33:43
Sample Problem 1: Ground State Configuration and Orbital Box Diagram
34:21
Fe
34:48
P
35:32
Sample Problem 2
36:38
Which Has The Larger Ionization Energy: Na or Li?
36:39
Which Has The Larger Atomic Size: O or N ?
37:23
Which Has The Larger Atomic Size: O²⁻ or N³⁻ ?
38:00
VIII. Molecular Geometry & Bonding Theory
Bonding & Molecular Structure

52m 39s

Intro
0:00
Lesson Overview
0:08
Introduction
1:10
Types of Chemical Bonds
1:53
Ionic Bond
1:54
Molecular Bond
2:42
Electronegativity and Bond Polarity
3:26
Electronegativity (EN)
3:27
Periodic Trend
4:36
Electronegativity and Bond Polarity Cont'd
6:04
Bond Polarity: Polar Covalent Bond
6:05
Bond Polarity: Nonpolar Covalent Bond
8:53
Lewis Electron Dot Structure of Atoms
9:48
Lewis Electron Dot Structure of Atoms
9:49
Lewis Structures of Polyatomic Species
12:51
Single Bonds
12:52
Double Bonds
13:28
Nonbonding Electrons
13:59
Lewis Structures of Polyatomic Species Cont'd
14:45
Drawing Lewis Structures: Step 1
14:48
Drawing Lewis Structures: Step 2
15:16
Drawing Lewis Structures: Step 3
15:52
Drawing Lewis Structures: Step 4
17:31
Drawing Lewis Structures: Step 5
19:08
Drawing Lewis Structure Example: Carbonate
19:33
Resonance and Formal Charges (FC)
24:06
Resonance Structures
24:07
Formal Charge
25:20
Resonance and Formal Charges Cont'd
27:46
More On Formal Charge
27:47
Resonance and Formal Charges Cont'd
28:21
Good Resonance Structures
28:22
VSEPR Theory
31:08
VSEPR Theory Continue
31:09
VSEPR Theory Cont'd
32:53
VSEPR Geometries
32:54
Steric Number
33:04
Basic Geometry
33:50
Molecular Geometry
35:50
Molecular Polarity
37:51
Steps In Determining Molecular Polarity
37:52
Example 1: Polar
38:47
Example 2: Nonpolar
39:10
Example 3: Polar
39:36
Example 4: Polar
40:08
Bond Properties: Order, Length, and Energy
40:38
Bond Order
40:39
Bond Length
41:21
Bond Energy
41:55
Summary
43:09
Sample Problem 1
43:42
XeO₃
44:03
I₃⁻
47:02
SF₅
49:16
Advanced Bonding Theories

1h 11m 41s

Intro
0:00
Lesson Overview
0:09
Introduction
0:38
Valence Bond Theory
3:07
Valence Bond Theory
3:08
spᶟ Hybridized Carbon Atom
4:19
Valence Bond Theory Cont'd
6:24
spᶟ Hybridized
6:25
Hybrid Orbitals For Water
7:26
Valence Bond Theory Cont'd (spᶟ)
11:53
Example 1: NH₃
11:54
Valence Bond Theory Cont'd (sp²)
14:48
sp² Hybridization
14:49
Example 2: BF₃
16:44
Valence Bond Theory Cont'd (sp)
22:44
sp Hybridization
22:46
Example 3: HCN
23:38
Valence Bond Theory Cont'd (sp³d and sp³d²)
27:36
Valence Bond Theory: sp³d and sp³d²
27:37
Molecular Orbital Theory
29:10
Valence Bond Theory Doesn't Always Account For a Molecule's Magnetic Behavior
29:11
Molecular Orbital Theory Cont'd
30:37
Molecular Orbital Theory
30:38
Wavefunctions
31:04
How s-orbitals Can Interact
32:23
Bonding Nature of p-orbitals: Head-on
35:34
Bonding Nature of p-orbitals: Parallel
39:04
Interaction Between s and p-orbital
40:45
Molecular Orbital Diagram For Homonuclear Diatomics: H₂
42:21
Molecular Orbital Diagram For Homonuclear Diatomics: He₂
45:23
Molecular Orbital Diagram For Homonuclear Diatomic: Li₂
46:39
Molecular Orbital Diagram For Homonuclear Diatomic: Li₂⁺
47:42
Molecular Orbital Diagram For Homonuclear Diatomic: B₂
48:57
Molecular Orbital Diagram For Homonuclear Diatomic: N₂
54:04
Molecular Orbital Diagram: Molecular Oxygen
55:57
Molecular Orbital Diagram For Heteronuclear Diatomics: Hydrochloric Acid
02:16
Sample Problem 1: Determine the Atomic Hybridization
07:20
XeO₃
07:21
SF₆
07:49
I₃⁻
08:20
Sample Problem 2
09:04
IX. Gases, Solids, & Liquids
Gases

35m 6s

Intro
0:00
Lesson Overview
0:07
The Kinetic Molecular Theory of Gases
1:23
The Kinetic Molecular Theory of Gases
1:24
Parameters To Characterize Gases
3:35
Parameters To Characterize Gases: Pressure
3:37
Interpreting Pressure On a Particulate Level
4:43
Parameters Cont'd
6:08
Units For Expressing Pressure: Psi, Pascal
6:19
Units For Expressing Pressure: mm Hg
6:42
Units For Expressing Pressure: atm
6:58
Units For Expressing Pressure: torr
7:24
Parameters Cont'd
8:09
Parameters To Characterize Gases: Volume
8:10
Common Units of Volume
9:00
Parameters Cont'd
9:11
Parameters To Characterize Gases: Temperature
9:12
Particulate Level
9:36
Parameters To Characterize Gases: Moles
10:24
The Simple Gas Laws
10:43
Gas Laws Are Only Valid For…
10:44
Charles' Law
11:24
The Simple Gas Laws
13:13
Boyle's Law
13:14
The Simple Gas Laws
15:28
Gay-Lussac's Law
15:29
The Simple Gas Laws
17:11
Avogadro's Law
17:12
The Ideal Gas Law
18:43
The Ideal Gas Law: PV = nRT
18:44
Applications of the Ideal Gas Law
20:12
Standard Temperature and Pressure for Gases
20:13
Applications of the Ideal Gas Law
21:43
Ideal Gas Law & Gas Density
21:44
Gas Pressures and Partial Pressures
23:18
Dalton's Law of Partial Pressures
23:19
Gas Stoichiometry
24:15
Stoichiometry Problems Involving Gases
24:16
Using The Ideal Gas Law to Get to Moles
25:16
Using Molar Volume to Get to Moles
25:39
Gas Stoichiometry Cont'd
26:03
Example 1: How Many Liters of O₂ at STP are Needed to Form 10.5 g of Water Vapor?
26:04
Summary
28:33
Sample Problem 1: Calculate the Molar Mass of the Gas
29:28
Sample Problem 2: What Mass of Ag₂O is Required to Form 3888 mL of O₂ Gas When Measured at 734 mm Hg and 25°C?
31:59
Intermolecular Forces & Liquids

33m 47s

Intro
0:00
Lesson Overview
0:10
Introduction
0:46
Intermolecular Forces (IMF)
0:47
Intermolecular Forces of Polar Molecules
1:32
Ion-dipole Forces
1:33
Example: Salt Dissolved in Water
1:50
Coulomb's Law & the Force of Attraction Between Ions and/or Dipoles
3:06
IMF of Polar Molecules cont'd
4:36
Enthalpy of Solvation or Enthalpy of Hydration
4:37
IMF of Polar Molecules cont'd
6:01
Dipole-dipole Forces
6:02
IMF of Polar Molecules cont'd
7:22
Hydrogen Bonding
7:23
Example: Hydrogen Bonding of Water
8:06
IMF of Nonpolar Molecules
9:37
Dipole-induced Dipole Attraction
9:38
IMF of Nonpolar Molecules cont'd
11:34
Induced Dipole Attraction, London Dispersion Forces, or Vand der Waals Forces
11:35
Polarizability
13:46
IMF of Nonpolar Molecules cont'd
14:26
Intermolecular Forces (IMF) and Polarizability
14:31
Properties of Liquids
16:48
Standard Molar Enthalpy of Vaporization
16:49
Trends in Boiling Points of Representative Liquids: H₂O vs. H₂S
17:43
Properties of Liquids cont'd
18:36
Aliphatic Hydrocarbons
18:37
Branched Hydrocarbons
20:52
Properties of Liquids cont'd
22:10
Vapor Pressure
22:11
The Clausius-Clapeyron Equation
24:30
Properties of Liquids cont'd
25:52
Boiling Point
25:53
Properties of Liquids cont'd
27:07
Surface Tension
27:08
Viscosity
28:06
Summary
29:04
Sample Problem 1: Determine Which of the Following Liquids Will Have the Lower Vapor Pressure
30:21
Sample Problem 2: Determine Which of the Following Liquids Will Have the Largest Standard Molar Enthalpy of Vaporization
31:37
The Chemistry of Solids

25m 13s

Intro
0:00
Lesson Overview
0:07
Introduction
0:46
General Characteristics
0:47
Particulate-level Drawing
1:09
The Basic Structure of Solids: Crystal Lattices
1:37
The Unit Cell Defined
1:38
Primitive Cubic
2:50
Crystal Lattices cont'd
3:58
Body-centered Cubic
3:59
Face-centered Cubic
5:02
Lattice Enthalpy and Trends
6:27
Introduction to Lattice Enthalpy
6:28
Equation to Calculate Lattice Enthalpy
7:21
Different Types of Crystalline Solids
9:35
Molecular Solids
9:36
Network Solids
10:25
Phase Changes Involving Solids
11:03
Melting & Thermodynamic Value
11:04
Freezing & Thermodynamic Value
11:49
Phase Changes cont'd
12:40
Sublimation & Thermodynamic Value
12:41
Depositions & Thermodynamic Value
13:13
Phase Diagrams
13:40
Introduction to Phase Diagrams
13:41
Phase Diagram of H₂O: Melting Point
14:12
Phase Diagram of H₂O: Normal Boiling Point
14:50
Phase Diagram of H₂O: Sublimation Point
15:02
Phase Diagram of H₂O: Point C ( Supercritical Point)
15:32
Phase Diagrams cont'd
16:31
Phase Diagram of Dry Ice
16:32
Summary
18:15
Sample Problem 1, Part A: Of the Group I Fluorides, Which Should Have the Highest Lattice Enthalpy?
19:01
Sample Problem 1, Part B: Of the Lithium Halides, Which Should Have the Lowest Lattice Enthalpy?
19:54
Sample Problem 2: How Many Joules of Energy is Required to Melt 546 mg of Ice at Standard Pressure?
20:55
Sample Problem 3: Phase Diagram of Helium
22:42
X. Solutions, Rates of Reaction, & Equilibrium
Solutions & Their Behavior

38m 6s

Intro
0:00
Lesson Overview
0:10
Units of Concentration
1:40
Molarity
1:41
Molality
3:30
Weight Percent
4:26
ppm
5:16
Like Dissolves Like
6:28
Like Dissolves Like
6:29
Factors Affecting Solubility
9:35
The Effect of Pressure: Henry's Law
9:36
The Effect of Temperature on Gas Solubility
12:16
The Effect of Temperature on Solid Solubility
14:28
Colligative Properties
16:48
Colligative Properties
16:49
Changes in Vapor Pressure: Raoult's Law
17:19
Colligative Properties cont'd
19:53
Boiling Point Elevation and Freezing Point Depression
19:54
Colligative Properties cont'd
26:13
Definition of Osmosis
26:14
Osmotic Pressure Example
27:11
Summary
31:11
Sample Problem 1: Calculating Vapor Pressure
32:53
Sample Problem 2: Calculating Molality
36:29
Chemical Kinetics

37m 45s

Intro
0:00
Lesson Overview
0:06
Introduction
1:09
Chemical Kinetics and the Rate of a Reaction
1:10
Factors Influencing Rate
1:19
Introduction cont'd
2:27
How a Reaction Progresses Through Time
2:28
Rate of Change Equation
6:02
Rate Laws
7:06
Definition of Rate Laws
7:07
General Form of Rate Laws
7:37
Rate Laws cont'd
11:07
Rate Orders With Respect to Reactant and Concentration
11:08
Methods of Initial Rates
13:38
Methods of Initial Rates
13:39
Integrated Rate Laws
17:57
Integrated Rate Laws
17:58
Graphically Determine the Rate Constant k
18:52
Reaction Mechanisms
21:05
Step 1: Reversible
21:18
Step 2: Rate-limiting Step
21:44
Rate Law for the Reaction
23:28
Reaction Rates and Temperatures
26:16
Reaction Rates and Temperatures
26:17
The Arrhenius Equation
29:06
Catalysis
30:31
Catalyst
30:32
Summary
32:02
Sample Problem 1: Calculate the Rate Constant and the Time Required for the Reaction to be Completed
32:54
Sample Problem 2: Calculate the Energy of Activation and the Order of the Reaction
35:24
Principles of Chemical Equilibrium

34m 9s

Intro
0:00
Lesson Overview
0:08
Introduction
1:02
The Equilibrium Constant
3:08
The Equilibrium Constant
3:09
The Equilibrium Constant cont'd
5:50
The Equilibrium Concentration and Constant for Solutions
5:51
The Equilibrium Partial Pressure and Constant for Gases
7:01
Relationship of Kc and Kp
7:30
Heterogeneous Equilibria
8:23
Heterogeneous Equilibria
8:24
Manipulating K
9:57
First Way of Manipulating K
9:58
Second Way of Manipulating K
11:48
Manipulating K cont'd
12:31
Third Way of Manipulating K
12:32
The Reaction Quotient Q
14:42
The Reaction Quotient Q
14:43
Q > K
16:16
Q < K
16:30
Q = K
16:43
Le Chatlier's Principle
17:32
Restoring Equilibrium When It is Disturbed
17:33
Disturbing a Chemical System at Equilibrium
18:35
Problem-Solving with ICE Tables
19:05
Determining a Reaction's Equilibrium Constant With ICE Table
19:06
Problem-Solving with ICE Tables cont'd
21:03
Example 1: Calculate O₂(g) at Equilibrium
21:04
Problem-Solving with ICE Tables cont'd
22:53
Example 2: Calculate the Equilibrium Constant
22:54
Summary
25:24
Sample Problem 1: Calculate the Equilibrium Constant
27:59
Sample Problem 2: Calculate The Equilibrium Concentration
30:30
XI. Acids & Bases Chemistry
Acid-Base Chemistry

43m 44s

Intro
0:00
Lesson Overview
0:06
Introduction
0:55
Bronsted-Lowry Acid & Bronsted -Lowry Base
0:56
Water is an Amphiprotic Molecule
2:40
Water Reacting With Itself
2:58
Introduction cont'd
4:04
Strong Acids
4:05
Strong Bases
5:18
Introduction cont'd
6:16
Weak Acids and Bases
6:17
Quantifying Acid-Base Strength
7:35
The pH Scale
7:36
Quantifying Acid-Base Strength cont'd
9:55
The Acid-ionization Constant Ka and pKa
9:56
Quantifying Acid-Base Strength cont'd
12:13
Example: Calculate the pH of a 1.2M Solution of Acetic Acid
12:14
Quantifying Acid-Base Strength
15:06
Calculating the pH of Weak Base Solutions
15:07
Writing Out Acid-Base Equilibria
17:45
Writing Out Acid-Base Equilibria
17:46
Writing Out Acid-Base Equilibria cont'd
19:47
Consider the Following Equilibrium
19:48
Conjugate Base and Conjugate Acid
21:18
Salts Solutions
22:00
Salts That Produce Acidic Aqueous Solutions
22:01
Salts That Produce Basic Aqueous Solutions
23:15
Neutral Salt Solutions
24:05
Diprotic and Polyprotic Acids
24:44
Example: Calculate the pH of a 1.2 M Solution of H₂SO₃
24:43
Diprotic and Polyprotic Acids cont'd
27:18
Calculate the pH of a 1.2 M Solution of Na₂SO₃
27:19
Lewis Acids and Bases
29:13
Lewis Acids
29:14
Lewis Bases
30:10
Example: Lewis Acids and Bases
31:04
Molecular Structure and Acidity
32:03
The Effect of Charge
32:04
Within a Period/Row
33:07
Molecular Structure and Acidity cont'd
34:17
Within a Group/Column
34:18
Oxoacids
35:58
Molecular Structure and Acidity cont'd
37:54
Carboxylic Acids
37:55
Hydrated Metal Cations
39:23
Summary
40:39
Sample Problem 1: Calculate the pH of a 1.2 M Solution of NH₃
41:20
Sample Problem 2: Predict If The Following Slat Solutions are Acidic, Basic, or Neutral
42:37
Applications of Aqueous Equilibria

55m 26s

Intro
0:00
Lesson Overview
0:07
Calculating pH of an Acid-Base Mixture
0:53
Equilibria Involving Direct Reaction With Water
0:54
When a Bronsted-Lowry Acid and Base React
1:12
After Neutralization Occurs
2:05
Calculating pH of an Acid-Base Mixture cont'd
2:51
Example: Calculating pH of an Acid-Base Mixture, Step 1 - Neutralization
2:52
Example: Calculating pH of an Acid-Base Mixture, Step 2 - React With H₂O
5:24
Buffers
7:45
Introduction to Buffers
7:46
When Acid is Added to a Buffer
8:50
When Base is Added to a Buffer
9:54
Buffers cont'd
10:41
Calculating the pH
10:42
Calculating the pH When 0.010 mol NaOH is Added to 1.0 L of the Buffer
14:03
Buffers cont'd
14:10
Calculating the pH When 0.010 mol NaOH is Added to 1.0 L of the Buffer: Step 1 -Neutralization
14:11
Calculating the pH When 0.010 mol NaOH is Added to 1.0 L of the Buffer: Step 2- ICE Table
15:22
Buffer Preparation and Capacity
16:38
Example: Calculating the pH of a Buffer Solution
16:42
Effective Buffer
18:40
Acid-Base Titrations
19:33
Acid-Base Titrations: Basic Setup
19:34
Acid-Base Titrations cont'd
22:12
Example: Calculate the pH at the Equivalence Point When 0.250 L of 0.0350 M HClO is Titrated With 1.00 M KOH
22:13
Acid-Base Titrations cont'd
25:38
Titration Curve
25:39
Solubility Equilibria
33:07
Solubility of Salts
33:08
Solubility Product Constant: Ksp
34:14
Solubility Equilibria cont'd
34:58
Q < Ksp
34:59
Q > Ksp
35:34
Solubility Equilibria cont'd
36:03
Common-ion Effect
36:04
Example: Calculate the Solubility of PbCl₂ in 0.55 M NaCl
36:30
Solubility Equilibria cont'd
39:02
When a Solid Salt Contains the Conjugate of a Weak Acid
39:03
Temperature and Solubility
40:41
Complexation Equilibria
41:10
Complex Ion
41:11
Complex Ion Formation Constant: Kf
42:26
Summary
43:35
Sample Problem 1: Question
44:23
Sample Problem 1: Part a) Calculate the pH at the Beginning of the Titration
45:48
Sample Problem 1: Part b) Calculate the pH at the Midpoint or Half-way Point
48:04
Sample Problem 1: Part c) Calculate the pH at the Equivalence Point
48:32
Sample Problem 1: Part d) Calculate the pH After 27.50 mL of the Acid was Added
53:00
XII. Thermodynamics & Electrochemistry
Entropy & Free Energy

36m 13s

Intro
0:00
Lesson Overview
0:08
Introduction
0:53
Introduction to Entropy
1:37
Introduction to Entropy
1:38
Entropy and Heat Flow
6:31
Recall Thermodynamics
6:32
Entropy is a State Function
6:54
∆S and Heat Flow
7:28
Entropy and Heat Flow cont'd
8:18
Entropy and Heat Flow: Equations
8:19
Endothermic Processes: ∆S > 0
8:44
The Second Law of Thermodynamics
10:04
Total ∆S = ∆S of System + ∆S of Surrounding
10:05
Nature Favors Processes Where The Amount of Entropy Increases
10:22
The Third Law of Thermodynamics
11:55
The Third Law of Thermodynamics & Zero Entropy
11:56
Problem-Solving involving Entropy
12:36
Endothermic Process and ∆S
12:37
Exothermic Process and ∆S
13:19
Problem-Solving cont'd
13:46
Change in Physical States: From Solid to Liquid to Gas
13:47
Change in Physical States: All Gases
15:02
Problem-Solving cont'd
15:56
Calculating the ∆S for the System, Surrounding, and Total
15:57
Example: Calculating the Total ∆S
16:17
Problem-Solving cont'd
18:36
Problems Involving Standard Molar Entropies of Formation
18:37
Introduction to Gibb's Free Energy
20:09
Definition of Free Energy ∆G
20:10
Spontaneous Process and ∆G
20:19
Gibb's Free Energy cont'd
22:28
Standard Molar Free Energies of Formation
22:29
The Free Energies of Formation are Zero for All Compounds in the Standard State
22:42
Gibb's Free Energy cont'd
23:31
∆G° of the System = ∆H° of the System - T∆S° of the System
23:32
Predicting Spontaneous Reaction Based on the Sign of ∆G° of the System
24:24
Gibb's Free Energy cont'd
26:32
Effect of reactant and Product Concentration on the Sign of Free Energy
26:33
∆G° of Reaction = -RT ln K
27:18
Summary
28:12
Sample Problem 1: Calculate ∆S° of Reaction
28:48
Sample Problem 2: Calculate the Temperature at Which the Reaction Becomes Spontaneous
31:18
Sample Problem 3: Calculate Kp
33:47
Electrochemistry

41m 16s

Intro
0:00
Lesson Overview
0:08
Introduction
0:53
Redox Reactions
1:42
Oxidation-Reduction Reaction Overview
1:43
Redox Reactions cont'd
2:37
Which Reactant is Being Oxidized and Which is Being Reduced?
2:38
Redox Reactions cont'd
6:34
Balance Redox Reaction In Neutral Solutions
6:35
Redox Reactions cont'd
10:37
Balance Redox Reaction In Acidic and Basic Solutions: Step 1
10:38
Balance Redox Reaction In Acidic and Basic Solutions: Step 2 - Balance Each Half-Reaction
11:22
Redox Reactions cont'd
12:19
Balance Redox Reaction In Acidic and Basic Solutions: Step 2 - Balance Hydrogen
12:20
Redox Reactions cont'd
14:30
Balance Redox Reaction In Acidic and Basic Solutions: Step 3
14:34
Balance Redox Reaction In Acidic and Basic Solutions: Step 4
15:38
Voltaic Cells
17:01
Voltaic Cell or Galvanic Cell
17:02
Cell Notation
22:03
Electrochemical Potentials
25:22
Electrochemical Potentials
25:23
Electrochemical Potentials cont'd
26:07
Table of Standard Reduction Potentials
26:08
The Nernst Equation
30:41
The Nernst Equation
30:42
It Can Be Shown That At Equilibrium E =0.00
32:15
Gibb's Free Energy and Electrochemistry
32:46
Gibbs Free Energy is Relatively Small if the Potential is Relatively High
32:47
When E° is Very Large
33:39
Charge, Current and Time
33:56
A Battery Has Three Main Parameters
33:57
A Simple Equation Relates All of These Parameters
34:09
Summary
34:50
Sample Problem 1: Redox Reaction
35:26
Sample Problem 2: Battery
38:00
XIII. Transition Elements & Coordination Compounds
The Chemistry of The Transition Metals

39m 3s

Intro
0:00
Lesson Overview
0:11
Coordination Compounds
1:20
Coordination Compounds
1:21
Nomenclature of Coordination Compounds
2:48
Rule 1
3:01
Rule 2
3:12
Rule 3
4:07
Nomenclature cont'd
4:58
Rule 4
4:59
Rule 5
5:13
Rule 6
5:35
Rule 7
6:19
Rule 8
6:46
Nomenclature cont'd
7:39
Rule 9
7:40
Rule 10
7:45
Rule 11
8:00
Nomenclature of Coordination Compounds: NH₄[PtCl₃NH₃]
8:11
Nomenclature of Coordination Compounds: [Cr(NH₃)₄(OH)₂]Br
9:31
Structures of Coordination Compounds
10:54
Coordination Number or Steric Number
10:55
Commonly Observed Coordination Numbers and Geometries: 4
11:14
Commonly Observed Coordination Numbers and Geometries: 6
12:00
Isomers of Coordination Compounds
13:13
Isomers of Coordination Compounds
13:14
Geometrical Isomers of CN = 6 Include: ML₄L₂'
13:30
Geometrical Isomers of CN = 6 Include: ML₃L₃'
15:07
Isomers cont'd
17:00
Structural Isomers Overview
17:01
Structural Isomers: Ionization
18:06
Structural Isomers: Hydrate
19:25
Structural Isomers: Linkage
20:11
Structural Isomers: Coordination Isomers
21:05
Electronic Structure
22:25
Crystal Field Theory
22:26
Octahedral and Tetrahedral Field
22:54
Electronic Structure cont'd
25:43
Vanadium (II) Ion in an Octahedral Field
25:44
Chromium(III) Ion in an Octahedral Field
26:37
Electronic Structure cont'd
28:47
Strong-Field Ligands and Weak-Field Ligands
28:48
Implications of Electronic Structure
30:08
Compare the Magnetic Properties of: [Fe(OH₂)₆]²⁺ vs. [Fe(CN)₆]⁴⁻
30:09
Discussion on Color
31:57
Summary
34:41
Sample Problem 1: Name the Following Compound [Fe(OH)(OH₂)₅]Cl₂
35:08
Sample Problem 1: Name the Following Compound [Co(NH₃)₃(OH₂)₃]₂(SO₄)₃
36:24
Sample Problem 2: Change in Magnetic Properties
37:30
XIV. Nuclear Chemistry
Nuclear Chemistry

16m 39s

Intro
0:00
Lesson Overview
0:06
Introduction
0:40
Introduction to Nuclear Reactions
0:41
Types of Radioactive Decay
2:10
Alpha Decay
2:11
Beta Decay
3:27
Gamma Decay
4:40
Other Types of Particles of Varying Energy
5:40
Nuclear Equations
6:47
Nuclear Equations
6:48
Nuclear Decay
9:28
Nuclear Decay and the First-Order Kinetics
9:29
Summary
11:31
Sample Problem 1: Complete the Following Nuclear Equations
12:13
Sample Problem 2: How Old is the Rock?
14:21
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Lecture Comments (4)

0 answers

Post by terry terrison on October 11, 2017

is this forum really moderated as the site claims? I'd like to know the solution to the problem with the numbers written in. I did not get -0.248. Please explain. I don't see the point of subscribing to this site if I cannot get simple problems explained clearly (what is the point of this video?)

0 answers

Post by Maria Mazos on January 24, 2017

18:24-can y please explain why -0.248?

1 answer

Last reply by: Maria Mazos
Tue Jan 24, 2017 8:18 PM

Post by Okwudili Ezeh on October 30, 2014

Please could you clearly state how you got the answer to the problem you solved at 18:24. You did not plug in any numbers and so how are we expected to understand what you did?

Entropy & Free Energy

  • Nature favors states of low energy and high entropy.
  • Entropy is related to disorder or chaos, and is the focus of the second and third laws of thermodynamics.
  • There are a series of equations that allow us to calculate both entropy andGibb’s free energy under a given set of conditions.

Entropy & Free Energy

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.

  1. Intro
    • Lesson Overview
      • Introduction
        • Introduction to Entropy
        • Entropy and Heat Flow
        • Entropy and Heat Flow cont'd
        • The Second Law of Thermodynamics
        • The Third Law of Thermodynamics
        • Problem-Solving involving Entropy
        • Problem-Solving cont'd
        • Problem-Solving cont'd
        • Problem-Solving cont'd
        • Introduction to Gibb's Free Energy
        • Gibb's Free Energy cont'd
        • Gibb's Free Energy cont'd
        • Gibb's Free Energy cont'd
        • Summary
          • Sample Problem 1: Calculate ∆S° of Reaction
            • Sample Problem 2: Calculate the Temperature at Which the Reaction Becomes Spontaneous
              • Sample Problem 3: Calculate Kp
                • Intro 0:00
                • Lesson Overview 0:08
                • Introduction 0:53
                • Introduction to Entropy 1:37
                  • Introduction to Entropy
                • Entropy and Heat Flow 6:31
                  • Recall Thermodynamics
                  • Entropy is a State Function
                  • ∆S and Heat Flow
                • Entropy and Heat Flow cont'd 8:18
                  • Entropy and Heat Flow: Equations
                  • Endothermic Processes: ∆S > 0
                • The Second Law of Thermodynamics 10:04
                  • Total ∆S = ∆S of System + ∆S of Surrounding
                  • Nature Favors Processes Where The Amount of Entropy Increases
                • The Third Law of Thermodynamics 11:55
                  • The Third Law of Thermodynamics & Zero Entropy
                • Problem-Solving involving Entropy 12:36
                  • Endothermic Process and ∆S
                  • Exothermic Process and ∆S
                • Problem-Solving cont'd 13:46
                  • Change in Physical States: From Solid to Liquid to Gas
                  • Change in Physical States: All Gases
                • Problem-Solving cont'd 15:56
                  • Calculating the ∆S for the System, Surrounding, and Total
                  • Example: Calculating the Total ∆S
                • Problem-Solving cont'd 18:36
                  • Problems Involving Standard Molar Entropies of Formation
                • Introduction to Gibb's Free Energy 20:09
                  • Definition of Free Energy ∆G
                  • Spontaneous Process and ∆G
                • Gibb's Free Energy cont'd 22:28
                  • Standard Molar Free Energies of Formation
                  • The Free Energies of Formation are Zero for All Compounds in the Standard State
                • Gibb's Free Energy cont'd 23:31
                  • ∆G° of the System = ∆H° of the System - T∆S° of the System
                  • Predicting Spontaneous Reaction Based on the Sign of ∆G° of the System
                • Gibb's Free Energy cont'd 26:32
                  • Effect of reactant and Product Concentration on the Sign of Free Energy
                  • ∆G° of Reaction = -RT ln K
                • Summary 28:12
                • Sample Problem 1: Calculate ∆S° of Reaction 28:48
                • Sample Problem 2: Calculate the Temperature at Which the Reaction Becomes Spontaneous 31:18
                • Sample Problem 3: Calculate Kp 33:47

                Transcription: Entropy & Free Energy

                Hi, welcome back to Educator.com.0000

                Today's lesson from general chemistry is on entropy and free energy.0003

                This is basically going to be our second and final lecture on thermodynamics.0009

                As usual we are going to start off with a brief introduction.0013

                We are going to get into something very important in thermodynamics which is what we call entropy.0017

                It is going to be related to spontaneous processes.0023

                After we get into entropy, we are going to go through the second and third laws of thermodynamics followed by problem solving.0027

                After entropy, we are going to then get into what is called Gibbs free energy followed by definition and problem solving.0038

                Then we are going to go ahead and conclude with a summary and a pair of sample problems.0047

                We have consistently said, you have heard me say over and over again that nature favors states of low energy.0055

                In other words, reactions occur to obtain a lower energy state in the end.0063

                This is also known as a downhill reaction.0069

                In other words, nature tends to favor reactions that start in a high energy state and go to a low energy state.0072

                That is going to be the natural tendency.0079

                In this lecture, we are going to complete our discussion from thermodynamics and quantify the above statement.0081

                In the end, we are going to see that nature not only favors0089

                states of low energy but also states of what we call high entropy.0091

                Consider the following closed reaction vessels that contain gas molecules.0099

                Let me go ahead and draw a closed reaction vessel here.0104

                We can make this one gas particle; the red one can be another gas particle.0109

                The black one can be another gas particle; we are now going to...0115

                Because these are gases, gases naturally expand; what can happen is the following.0120

                What are some possible different configurations for these three gas particles if they naturally expand?0128

                Situation number one is when all three gas particles are on the same side.0135

                We can have a second situation when only the blue and the black are together and the red is different.0140

                We can have another situation where the blue is by itself and the red and black are together.0148

                You notice that these two situations here are actually the same0162

                where we have two gas particles on one side versus one on the other.0169

                It turns out that configuration number two actually has more what we call microstates.0177

                Once again configuration number two is going to have what we call more microstates.0196

                We can even do one more to show you how this is possible.0204

                Now red and blue can be together with black by itself.0209

                This is also considered a configuration number two.0215

                Configuration number one, there is only one way to put three particles on the same side.0220

                Configuration number one has one microstate.0227

                Because gases expand, we are going to see that configuration number two is going to be the more probable configuration.0237

                Why?--because there are a higher number of ways to do it.0246

                There are a higher number of microstates to achieve, configuration number two, where one gas particle is on one side.0250

                Two gas particles are on the other side.0257

                We symbolize the term microstate with capital W.0260

                For configuration number one, W is 1; for configuration number two, W is simply 3.0264

                We can quantify that this is going to be more probable.0273

                In other words, this is more likely to occur.0281

                The equation we use to quantify this is following.0286

                S is equal to kB times the natural log of W where S is what we call entropy.0289

                I want you to think of entropy as basically disorder or chaos.0297

                We see that S is directly proportional to W.0308

                In other words, as W goes up, so does S.0313

                What does that mean though?0318

                We already said that a high value of W means it is going to be more likely to occur0321

                which says that very high entropy values are going to relate to more likely processes.0326

                If W large, S is large; process more likely to occur.0337

                In this equation, we see that there is a proportionality constant called kB.0351

                kB is what we call the Boltzmann constant.0356

                That is equal to 1.38 times 10-23 joules per kelvin.0363

                Because natural log of W is going to be unitless, that means the units of entropy are also joules per kelvin.0369

                This entire equation is known as the Boltzmann equation.0377

                S is equal to kB natural log of W.0383

                The units of entropy are going to be joules per kelvin.0386

                Let's go ahead and relate to entropy to what we have already discussed0393

                in our first lecture from thermodynamics which involved heat flow and energy.0397

                If you recall from our first thermodynamic lecture, if heat leaves the system, Q is going to be a negative value.0402

                If heat enters the system, Q is going to be a positive value.0410

                Entropy is going to be a state function and can be related to Q where ΔS is equal to Q over T.0415

                Temperature, because you see capital T, has to be in kelvin here.0426

                Again because this is a state function, what we are interested in is0433

                a change in entropy just like we talked about for ΔE or ΔU.0437

                If heat enters the system, Q is going to be positive which means ΔS is going to be positive.0448

                Let's see if that makes sense.0460

                Remember if heat enters the system, we are supplying thermal energy for molecules to absorb.0462

                If they are going to absorb this thermal energy, that gives them more molecular motion such as vibrations.0470

                That increases the number of states an electron can occupy.0477

                Basically we are increasing the amount of disorder.0483

                We are increasing the amount of chaos.0485

                Yes, it does make sense that ΔS is a positive value.0488

                ΔS is also related to heat flow during phase changes.0494

                It is found that Q is equal to n times ΔH of the phase change.0500

                This is from our first thermodynamics; from thermo I.0506

                Therefore it follows that ΔS is equal to Q over T which equals to nΔH of the phase change over the temperature.0514

                What does that mean then?0525

                That means for endothermic processes such as melting and vaporization, endothermic means Q is going to be a positive value.0527

                ΔS is positive.0540

                This agrees with the previous slide where when we put energy in, everything gets a little more chaotic and a little more disorderly.0542

                For example, as you go for melting, you go from a very orderly state, a solid, to a less orderly state, to a liquid.0553

                When go from a liquid to a gas, a liquid is less orderly.0563

                A gas is going to be the most disordered state.0567

                Everything is a lot more chaotic there.0572

                ΔS is going to be positive when heat enters the system.0576

                Here we have the phase change backing up this point.0580

                In other words then, we can summarize this very nicely.0586

                If a process is likely to occur, then ΔS is going to be a positive value.0590

                We say that the process is expected to be spontaneous if ΔS is positive.0596

                This brings us into our second law of thermodynamics.0605

                The second law of thermodynamics states that a process will be spontaneous if ΔS total is greater0609

                than 0 where ΔS total is equal to ΔS of the system plus ΔS of the surroundings.0616

                In other words, nature is going to favor processes that contribute to the overall entropy of the entire universe.0623

                In other words, nature favors processes where the amount of entropy or disorder increases as a result of the process.0631

                We see this occurring in nature quite a bit.0640

                If you think of a virus, what does a virus naturally do?0643

                A virus is going to naturally expand and go from cell to cell consuming the resources there and moving on and on.0648

                If viruses expand, they get more disorderly, more chaotic.0658

                We see this in herd migration.0662

                Herd migration, flocks of animals tend to naturally expand from area to area.0667

                They tend not to stay just in a single location.0678

                Of course the universe itself, the universe theoretically started off as a single point.0682

                It is just constantly expanding and expanding.0689

                Again we see that entropy is seen in several naturally occurring examples.0694

                Then you may be wondering, if ΔS positive is going to be for a spontaneous process, what does it mean when entropy is zero?0704

                That brings us into the third law of thermodynamics.0715

                Zero entropy really does not exist because the conditions to obtain zero entropy do not exist experimentally.0718

                The third law of thermodynamics states that entropy is zero only for a perfect flawless crystal at absolute zero.0728

                But absolute zero has not been obtained experimentally.0736

                Only in the most extreme conditions where we have zero thermo motion, that is molecules cease to move, entropy is completely zero.0741

                In other words, entropy is always going to be a nonzero quantity in practice.0750

                Let's now focus on problem solving involving entropy.0758

                The following are typical types of questions involving entropy.0762

                Sometimes a question can ask you simply to predict the sign of ΔS given a chemical reaction.0765

                This is going to be an example of a qualitative process, of a qualitative problem.0770

                Basically when the process is a physical change, ΔS is going to be positive for0777

                any endothermic process just like we talked about, including melting, vaporization, and sublimation.0781

                When you go from a solid to a liquid, from a liquid to a gas, or from a0788

                solid to a gas, ΔS of these processes are going to be expected to be positive.0793

                They are opposite processes, are all exothermic which means ΔS is expected to be negative.0801

                Liquid to solid, gas to liquid, and gas to solid.0807

                Freezing, condensation, and deposition; again ΔS is expected to be negative.0812

                Let's go on to another type of problem.0822

                Another type of problem involves a quantitative problem.0828

                What you want to do here is pay close attention to the physical states.0834

                From a solid to a liquid to a gas, entropy increases.0838

                We have already seen that.0841

                What you want to keep track of is the change in physical states from reactant to product side.0843

                Let's go ahead and look at an example.0850

                Fe solid plus O2 gas goes on to form Fe2O3 solid.0856

                We are going to need two of these, three of these, and two of these.0864

                Excuse me... we will need four irons.0869

                You notice that we start off with a solid and a gas.0873

                But on the product side, we have no gases here; no gases on product side.0878

                If we are eliminating the highest disorderly state which is a gas, we are starting off with it.0887

                We are not winding with it; we are actually getting less disorderly.0894

                ΔS for this process expected to be a negative value.0898

                Let's go on to another type of problem.0903

                If the reaction involves all gases, what you want to keep track of0906

                is the net change in gas molecules from reactant to product side.0909

                For example, N2 gas plus 3H2 gas goes on to form 2NH3 gas.0913

                On the left side, we start off with four gas molecules.0923

                On the right side, we wind up with only two gas molecules.0930

                For all intents and purposes, gas molecule A is going to be the same as gas molecule B for entropy counting purposes.0935

                Because we start off with four gas molecules and we only wind up with two,0945

                we are actually decreasing the amount of disorder in this process.0949

                ΔS is expected to be negative for this example.0953

                Sometimes the problem is going to ask you to calculate ΔS0958

                for what is called the system, surroundings, and total where0960

                ΔS total is equal to ΔS system plus ΔS of the surroundings.0963

                We have talked about system and surroundings in thermodynamics I.0967

                If you recall, Q of the system equals to ?Q of the surroundings.0972

                This is going to come into play.0978

                If a refrigerator coolant absorbs heat from stored food items, the heat then vaporizes the coolant which boils at -27 degrees Celsius.0979

                ΔH of vaporization is equal to -22.0 kilojoules per mole.0988

                Calculate ΔS total when 1.471 moles of the coolant vaporizes exchanging heat with the food items that are stored at 4 degrees Celsius.0993

                What we have to do is the following.1002

                We know we are using this equation for sure.1005

                ΔS total is equal to ΔS of the system plus ΔS of the surroundings.1007

                Let's go ahead and define the system; it doesn't matter which one as the coolant.1018

                Let's go ahead and define the surroundings as the food items inside the refrigerator.1024

                This is then going to be equal to...1031

                We know that ΔS is equal to Q of the system over T.1035

                The ΔS of the surroundings is going to be equal to Q of the surroundings over T.1043

                But we also know that this relationship here where Q of surroundings is equal to ?Q of the system.1050

                We can then go ahead and plug everything in.1060

                That is going to be equal to nΔH of vaporization over the temperature of the coolant.1062

                That is going to be equal to ?nΔH of vaporization but this time over the temperature of the food.1072

                Really the only difference is the temperature that the system and surroundings are stored at1080

                if we assume complete transfer of heat from the system to the surroundings.1085

                When all is said and done, we are going to get -0.248... don't forget the units.1092

                Here this is going to be kilojoules per kelvin or -248 joules per kelvin.1098

                Again you must be careful on how you define your system and surroundings.1108

                Be consistent throughout the entire equation.1113

                Finally another type of problem involves standard molar entropies of formations which is the1118

                change in entropy when a compound is formed from its constituent elements under standard conditions.1123

                Recall that we covered standard molar enthalpies of formation and used the summation equation.1131

                Similarly ΔS of the reaction is going to be equal to summation nS of the products minus summation nS of the reactants.1137

                These values here again just like for the enthalpy, you are going to look up in the appendix.1148

                Don't forget that n is going to be the stoichiometric coefficient.1157

                Unlike standard molar enthalpies of formation, known compound is going to have a standard molar enthalpy of formation.1169

                Why?--because the zero is only obtained under the conditions of the third law of thermodynamics.1175

                Let's go ahead and examine our... excuse me... this is very important.1182

                Again you never can assume that the molar entropy of formation is a zero value ever.1191

                It is only true for ΔH and, we are going to learn, for one more--what we call ΔG later.1200

                Let's examine our final state function.1207

                This is an introduction to Gibbs free energy; free energy is ΔG.1210

                It is defined as the amount of energy available to do work.1215

                We said that mother nature favors high entropy already and low energy.1220

                It is that energy that we are really talking about.1225

                We will see that the spontaneous processes will all have ΔG less than zero.1227

                To go ahead and do this, we can do ΔG is equal to ΔH minus TΔS.1241

                This is the equation that relates all three thermodynamic properties together.1258

                We already said that a positive ΔS is going to be a spontaneous process.1266

                We want ΔG to be negative for a spontaneous process.1279

                What conditions give you ΔG negative?1287

                ΔG negative is going to be spontaneous under the following conditions--for very large ΔS values and at high temperatures.1290

                As you can see, if ΔS is large and temperature is large,1312

                this term is going to be larger than this term giving us1318

                an overall negative ΔG value; very large ΔS at high temperatures.1322

                Basically ΔH is going to be a relatively negative value; ΔH, small as possible.1329

                Again this is what we call Gibbs free energy.1345

                Just like our previous state functions, we have standard molar free energies of formation1350

                where ΔG of the reaction is equal to summation nΔG of the products minus summation nΔG of the reactants.1355

                Like enthalpy, the free energies of formation are zero for all compounds in the standard state.1363

                Something like O2 gas, H2 gas, N2 gas, Cl2 gas, Br2 liquid, and I2 solid just to name a couple of examples.1369

                These are all going to have a ΔG of formation which is equal to ΔH of formation equal to zero.1385

                But again S, the molar entropy of formation, is not going to be zero whatsoever.1397

                Let's now study the temperature dependence of ΔG.1409

                It is found that we already said that ΔG of the system is equal to ΔH of the system minus TΔS of the system.1414

                The reason why this is important is because of the following--this term right here.1424

                ΔS of the system is in this equation.1429

                But the second law tells us that ΔS of the universe has to be a positive value.1433

                If we go strictly by entropy, we need to know the ΔS of the total, not just of the system or the surroundings.1446

                However this equation here tells us that all I have to know is ΔG of the system; that is all.1454

                Let's go ahead and see what conditions are going to be right for this.1466

                We can go ahead and make a chart here; ΔH, T, and ΔS.1474

                Then this is going to be sign of ΔG.1481

                Let's go ahead and delete temperature here--sign of ΔG, and then ΔH, ΔS.1488

                Let's say ΔH is a negative value--if this is less than zero and ΔS is a negative value.1495

                If ΔH is negative and ΔS is negative, it will be very difficult to get ΔG to be negative.1505

                The only time ΔG can be negative is going to be at a low temperature here.1516

                What if ΔH is negative and ΔS is positive?1523

                If ΔH is negative and ΔS is positive, then this sign of ΔG is going to be negative at high temperature.1531

                What if ΔH is positive and ΔS is positive?1547

                If ΔH is positive and ΔS is positive, ΔG is going to be negative at high temperatures.1551

                Finally what if ΔH is positive and ΔS is negative?1560

                If ΔH is positive and ΔS is negative, ΔG will never be negative according to the equation at any temperature.1565

                This is going to be greater than zero at all temperatures.1574

                Again as you can see, ΔG of the system is greatly influenced by1579

                not only the signs of ΔH and ΔS but also the kelvin temperature.1585

                Finally let's go ahead and examine the effect of reactant and product concentration on the sign of free energy.1595

                It is found that ΔG of the reaction is equal to ΔG0 plus RT natural log of Q.1602

                Remember what Q is?--Q is what we call the reaction quotient.1609

                This equation relates reactant and product concentration to the free energy not at1618

                standard conditions which means the concentrations are not going to be 1 molar.1623

                Or the pressures are not going to be 1 atm; or partial pressures not equal to 1 atm.1630

                Finally it can also be shown at equilibrium that ΔG not of the reaction is equal to ?RT natural log of K1639

                where this is our equilibrium constant which we have discussed so much in the previous sections.1647

                This is any equilibrium constant.1654

                This can be Ka, Kb, Kp, Ksp, Kf, etc; any of them.1656

                Basically this equation tells us that if K is large, the reaction is product favored.1667

                ΔG of the reaction is a negative value; that makes sense.1672

                If K is very large, that means the reaction is highly product favored.1677

                In other words, the reaction as written is going to be likely to occur.1682

                If the reaction is likely to occur, ΔG of that reaction should be negative.1687

                Let's go ahead and summarize our thermodynamics lecture before getting into our problems.1694

                We see that nature not only favors states of low energy but also now high entropy.1701

                Entropy is basically disorder or chaos.1707

                It is the focus of the second and third laws of thermodynamics.1710

                Basically what we have seen in this lecture is that there is a series of equations that govern entropy and free energy.1713

                They allow us to calculate both of them under a given set of conditions.1722

                Let's go ahead and do sample problem number one.1729

                Calculate ΔS of the reaction under standard conditions for the following.1732

                We have aluminum oxide reacting with 3H2 going on to form two aluminums plus three waters.1738

                Anytime you see a balanced chemical equation and it asks you simply to calculate1746

                ΔS not of the reaction, you are going to use the summation equation.1751

                Use ΔS of the reaction is equal to summation n S molality of all products minus summation n S molality of all reactants.1755

                I am just going to setup the problem for you.1772

                Again you are going to look up these values in the appendix of your textbook.1775

                Let's go ahead and do the products first.1779

                This is going to be 2 moles of aluminum times the standard molar entropy of formation of aluminum solid.1781

                Because it is summation, I am going to add... plus.1799

                Let me put that in parentheses.1803

                Plus 3 moles of the water vapor times its standard molar entropy.1804

                This is all in brackets; that is going to be subtracted from the reactant part.1813

                Here this is going to be 1 mole of Al2O3 solid times1819

                its molar entropy plus 3 moles of H2 gas times its molar entropy.1830

                Again you must pay attention to the physical states because the molar entropy of formation for say water vapor1848

                is going to be different than from liquid water which is going to be different than from solid water.1856

                Once again pay attention to the physical states.1862

                When this is all said and done, you are going to get your answer1865

                in of course units of joules per kelvin or kilojoules per kelvin.1868

                It depends what the question specifies of course.1874

                Let's go ahead and do sample problem number two.1880

                The reaction N2O4 going to 2NO2 is not spontaneous under standard conditions.1883

                Calculate the temperature at which the reaction becomes spontaneous.1889

                What that translates to is, in other words, at what temperature does ΔG become negative?1893

                Anytime you are asked to calculate a temperature at which something becomes spontaneous,1902

                pretty much you always use this equation: ΔG is equal to ΔH minus TΔS.1908

                They are asking you when does this change sign?1916

                In other words, when does ΔG go from positive to negative?1920

                What you simply is you set this whole thing equal to zero.1926

                Zero is equal to ΔH minus TΔS; you solve for temperature.1931

                Temperature is going to be equal to ΔH over ΔS.1936

                This is all going to be standard conditions of course.1941

                To get ΔH, you are going to use the summation equation.1946

                To get ΔS, you are going to use the summation equation1950

                which means you have to look up the formation values in the appendix.1954

                Look up formation values in appendix; appendices.1959

                When you do this, you are going to get a temperature value.1969

                At temperatures lower than T, ΔG is going to be positive.1976

                At temperatures higher than what you calculate for T, ΔG is going to be a negative value.1987

                Once again anytime you are asked to calculate the temperature at which something becomes spontaneous,2006

                almost always you are going to use this equation: ΔG is equal to ΔH minus TΔS.2015

                Set ΔG equal to zero and solve for the temperature.2021

                Finally our third sample problem, calculate ΔG of the reaction at 25 degrees Celsius2029

                for the formation of 1 mole of C2H5OH gas from C2H4 gas and H2O gas.2036

                The first thing we probably want to do is translate this into a balanced chemical equation.2047

                The reactants are C2H4 gas; this is going to react with H2O gas.2054

                That is going to form 1 mole of C2H5OH gas.2063

                Let's see if everything is balanced.2069

                Two carbons, six hydrogens, one oxygen; yes, this is nicely balanced.2071

                You are asked to calculate ΔG of the reaction at 25 degrees Celsius.2078

                Again ΔG of the reaction can simply be calculated from ΔH of the reaction minus TΔS of the reaction.2083

                Again to get the ΔH, you are going to use the summation equation.2094

                To get the ΔS, you are going to use the summation equation.2099

                Again the kelvin temperature is already specified for you right there.2102

                That gives you ΔG of the reaction; they then ask you to calculate Kp.2108

                The only equation in this lecture that involves the equilibrium constant is ΔG is equal to ?RT natural log of K.2114

                We can then insert this ΔG of a reaction in there.2126

                K therefore is equal to 10 raised to the -ΔG0 over RT2131

                where of course R is equal to 8.314 joules per k mole.2140

                Always double check your work.2151

                Make sure that K is a positive value because again2152

                we saw in previous lectures that K can never ever be negative.2158

                That is our lecture from general chemistry on thermodynamics.2165

                I want to thank you for your time and attention.2169

                I will see you next time on Educator.com.2171

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