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Franklin Ow

Franklin Ow

Principles of Chemical Equilibrium

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
1:02:16
Sample Problem 1: Determine the Atomic Hybridization
1:07:20
XeO₃
1:07:21
SF₆
1:07:49
I₃⁻
1:08:20
Sample Problem 2
1:09:04
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 (13)

0 answers

Post by Kamal Ali on September 18 at 11:11:45 AM

Hi Pro. Ow!!!

So I am confused on the ICE problem solving for Example 1. Can you please explain to me on how to go through the process of getting 8.0 x 10^-4??

1 answer

Last reply by: Professor Franklin Ow
Thu May 28, 2015 12:35 PM

Post by Smriti Sharan on May 20, 2015

How do you know what it is not at equilibrium? And how do you know when it is?

1 answer

Last reply by: Professor Franklin Ow
Fri Apr 3, 2015 11:02 PM

Post by Sachin Ambulkar on April 3, 2015

If you are given 2SO2(g) + O2(g) (equilibrium arrow) 2SO3(g) + heat, and the system is at equilibrium, increasing the temperature increases the number of moles of which substance(s) when equilibrium is reestablished?

1 answer

Last reply by: Professor Franklin Ow
Fri Apr 3, 2015 11:01 PM

Post by Saadman Elman on January 19, 2015

I agree with Jack Miars, In sample problem 1, the concentration of SO2 at equilibrium is 1 which is already given by the way. You said, it's 0.5. (Typo). Please clarify it. Thanks.

1 answer

Last reply by: Professor Franklin Ow
Sat Aug 16, 2014 4:20 AM

Post by Neil Choudhry on August 14, 2014

How did you get 4.0M for SO3 in Sample problem 1?  It says initially it had 12.0 .  

1 answer

Last reply by: Professor Franklin Ow
Fri Feb 7, 2014 10:26 AM

Post by Jack Miars on February 5, 2014

Also, on sample problem 1, x=.5 not the concentration of SO2. It should be 1

1 answer

Last reply by: Professor Franklin Ow
Fri Feb 7, 2014 10:25 AM

Post by Jack Miars on February 5, 2014

You didn't come back to how temperature and volume affect Q and K

Related Articles:

Principles of Chemical Equilibrium

  • Chemical equilibrium is a dynamic process, having equal rates of the forward and reverse reactions.
  • The equilibrium constant describes the extent to which products are formed over reactants.
  • Le Chatelier’s Principle states that when a system at equilibrium is disturbed, it will react to restore equilibrium.
  • ICE tables can be used to calculate equilibrium concentrations or pressures.

Principles of Chemical Equilibrium

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
  • Lesson Overview 0:08
  • Introduction 1:02
  • The Equilibrium Constant 3:08
    • The Equilibrium Constant
  • The Equilibrium Constant cont'd 5:50
    • The Equilibrium Concentration and Constant for Solutions
    • The Equilibrium Partial Pressure and Constant for Gases
    • Relationship of Kc and Kp
  • Heterogeneous Equilibria 8:23
    • Heterogeneous Equilibria
  • Manipulating K 9:57
    • First Way of Manipulating K
    • Second Way of Manipulating K
  • Manipulating K cont'd 12:31
    • Third Way of Manipulating K
  • The Reaction Quotient Q 14:42
    • The Reaction Quotient Q
    • Q > K
    • Q < K
    • Q = K
  • Le Chatlier's Principle 17:32
    • Restoring Equilibrium When It is Disturbed
    • Disturbing a Chemical System at Equilibrium
  • Problem-Solving with ICE Tables 19:05
    • Determining a Reaction's Equilibrium Constant With ICE Table
  • Problem-Solving with ICE Tables cont'd 21:03
    • Example 1: Calculate O₂(g) at Equilibrium
  • Problem-Solving with ICE Tables cont'd 22:53
    • Example 2: Calculate the Equilibrium Constant
  • Summary 25:24
  • Sample Problem 1: Calculate the Equilibrium Constant 27:59
  • Sample Problem 2: Calculate The Equilibrium Concentration 30:30

Transcription: Principles of Chemical Equilibrium

Hi, welcome back to Educator.com.0000

Today's lecture from general chemistry is on the principles of chemical equilibrium.0003

Let's go ahead and take a look at the lesson overview.0010

We will first start off with a brief introduction and then get in right to0012

the core of everything which is basically what we call the equilibrium constant.0017

We are going to first define this equilibrium constant and then go into different0021

types of equilibria including what we call homogeneous equilibria and heterogeneous equilibria.0026

We will then go and see how we can change the value of Keq.0033

We will also introduce something called the reaction quotient.0039

We then jump into a very fundamental principle in all of general chemistry.0042

That is called Le Chatelier's principle.0046

After that we will get into the quantitative part of the chapter which involves what we call ICE tables.0050

We will wrap everything up with a very brief summary followed by a pair of sample problems.0056

Chemical equilibrium, exactly what this is.0065

This refers to the simultaneous occurrence of a forward and a reverse reaction at the same rate.0068

We see that equilibrium is a dynamic process, not a static process.0075

For example, let's say you had two beakers here; these beakers are closed.0079

Let's say that this one beaker had H2O in it.0086

Let's just say this is time zero.0091

At time zero, I have nothing but pure liquid water.0094

But we know this from everyday experience inside a water bottle.0097

If you let some time progress, we notice that the water level is going to drop a little.0101

That is because some of the liquid water has entered the gas phase.0108

But there is also going to be a point in time where this water level is not going to drop forever.0116

It is going to reach a minimum.0125

It reaches a minimum because as soon as the vaporization process occurs, the condensation process also occurs.0129

We say for this case that the rate of evaporation is equal to the rate of condensation.0142

If we were to write this out in a chemical reaction, this would be H2O liquid.0158

Now we introduce a new type of reaction arrow which is this.0164

That goes to H2O gas; this is our equilibrium arrow.0169

It basically shows that the forward direction is happening simultaneously with the reverse direction.0180

Let's now get into how we can represent equilibrium numerically.0190

Consider the following reaction.0198

Small a big A plus small b big B, equilibrium sign and then small c big C small d big D.0199

In this reaction, let the lowercase letters represent the stoichiometric coefficients.0207

You know the moles after we balance the chemical equation.0222

It turns out that after a chemical reaction has reached equilibrium, it is experimentally determined that the ratio of0226

product to reactant concentration raised to the stoichiometric coefficients is actually constant at any given temperature.0235

Basically the concentration of C raised to the c power times the0243

concentration of D raised to the d power over the concentration of A0249

raised to the a power times the concentration of B raised to the b power.0256

This whole ratio of products to reactants raised to the coefficients is equal to some constant that we call K.0262

Sometimes you are going to see this called Keq.0272

This is formally what we call the equilibrium constant.0275

The only thing that can change the value of Keq is temperature.0285

The equilibrium constant is temperature dependent.0290

We can typically represent K in two different ways.0302

Kc is when molarities are used; Kp is when partial pressures are used.0307

Every textbook is a little different.0326

But for the partial pressures, the typical units are going to be atmosphere or bar.0328

Once again Kc and Kp are just the equilibrium expressions when0338

molarities and partial pressures are used for Kc and Kp respectively.0343

For solutions, the equilibrium constant can be expressed in units of molarity just like we discussed.0353

However it turns out that the equilibrium constant is unitless.0361

Kc, how do we go ahead and do that?0364

Mole over liter raised to some power divided by mole over liter raised to some power.0368

It turns out that the equilibrium expression is always referenced to 1 molarity.0378

This is really moles over liter per 1 molar.0385

This is moles over liter per one molar raised to the y power and raised to the x power.0394

As you can see that after this is done, we see that all units cancel.0402

Keq is actually one of the few unitless values in all of your general chemistry studies.0411

That is Kc.0420

It turns out that if we choose to do our problem with Kp, maybe this is going to be in atmospheres.0422

This is also relative to 1 atm, raised to some power divided by atm raised to some power.0429

We see it again that the units cancel; Keq again is unitless.0438

What is the relationship between Kc and Kp?0451

Kc and Kp are directly proportional to each other.0455

Depending on the reaction, sometimes they are nearly identical.0461

But the exact relationship between the two is the following where0465

Kp is equal to Kc times RT over Δn where Δn is0469

equal to the moles of gaseous product minus the moles of gaseous reactant.0479

Once again Kp and Kc can be very similar.0489

However they are not quite the same thing.0495

You should really check with your instructor to see what he or she prefers.0499

For chemical equilibria, we are usually dealing with very small concentrations.0507

Because of this, we assume that pure solids and pure liquids are going to remain relatively unchanged.0512

If you have the following reaction, A solid plus B aqueous goes on to form C aqueous plus D solid,0520

it is only the aqueous species that are going to affect the equilibrium, affect Keq.0532

Hence pure solids and pure liquids are going to be assigned an arbitrary value of 1 when incorporated into the expression for K.0546

That is they have no effect.0555

Kc for this reaction here would be simply the concentration of C times 1 divided by the0557

concentration of B times 1 which is just equal to the concentration of C over concentration of B.0566

Once again pure liquids and pure solids do not appear in the expression for K, in the Keq expression.0573

Again this is going to be specifically for heterogeneous equilibria.0593

What are some ways where we can manipulate K?0599

The first way of manipulating K is by multiplying an entire chemical equation by a factor.0602

For example, let's take A aqueous plus 2B aqueous goes on to form 3C aqueous.0608

In this case, K as you see is equal to the concentration of C cubed0621

divided by the concentration of A times the concentration of B squared.0627

Let's go ahead and take this chemical reaction and multiply everything through by 2.0633

2A aqueous plus 4B aqueous goes on to form 6C aqueous.0638

It turns out therefore that this Knew is going to be equal to concentration of C to the sixth0647

divided by the concentration of A squared times the concentration of B raised to the fourth power.0654

We see very nicely that Knew is simply equal to the original Kc squared.0661

The rule of thumb is the following.0670

That when multiplying a reaction by a factor, K is going to be raised to that factor.0672

K is raised to this factor.0695

That is the first way of algebraically manipulating K.0699

Once again this is by multiplying through a chemical reaction by a factor.0702

The second way is by taking the reverse reaction.0709

For example, let's take not A plus 2B going to 3C0712

but 3C aqueous goes on to form A aqueous plus 2B aqueous.0717

In this case, K is equal to concentration of A times the concentration of B squared divided by the concentration of C cubed.0726

We see that this is going to be 1 over kforward.0739

kreverse is simply the reciprocal of kforward.0745

The third way of changing K is by adding a series of individual chemical reactions together to form a net balanced chemical equation.0753

Let's go ahead and see.0760

For example, A aqueous plus B aqueous goes on to form C aqueous.0765

C aqueous plus D aqueous goes on to form E aqueous.0777

Let's go ahead and add these two together.0788

When we add these two together, we are going to get A aqueous plus B aqueous plus D aqueous goes on to form E aqueous.0790

You see that C is going to be cancelled out because it is going to be formed and consumed simultaneously.0803

What we want to look at now are the expressions for k for each of these.0815

Here k, I will call this k1 for reaction one, is equal to the0819

concentration of C divided by the concentration of A times the concentration of B.0823

Here k2 is equal to the concentration of E divided by the concentration of C times the concentration of D.0830

Here I will call this third one knet.0841

That is equal to the concentration of E divided by the concentration of A times the concentration of B times the concentration of D.0845

We see that when we add individual reactions together to get a net reaction, the equilibrium constant0855

of the net reaction is simply equal to the product of each individual equilibrium expression constant.0864

That is the third way of algebraically manipulating Keq.0874

Let's now go on to another aspect of the equilibrium constant.0884

It is what we call the reaction quotient.0887

The equilibrium constant is good when we actually have the values at equilibrium.0890

But what happens if we use values not at equilibrium?0894

When concentrations or pressures are inserted into the expression for K that are not at equilibrium,0899

the ratio of products to reactants is now what we call the reaction quotient symbolized Q.0906

Q is going to be equal to products raised to some power divided by reactants raised to0911

some power except that these molarities and partial pressures are not at equilibrium.0916

The significance is the following.0942

Q can be used to predict which way a reaction will shift to reobtain a state of equilibrium.0944

Basically we are going to have the following general rules.0976

If Q is greater than K, that means we have too much product, not enough reactant.0980

We are going to shift left.0987

If Q is less than K, we have too much reactant relative to product.0992

We are going to shift right.1001

Finally if Q is identical to K, we are at a equilibrium state.1005

Shifting left is the same as making more reactant.1015

Shifting right is the same as making more product.1028

At equilibrium, there is no net change; neither direction is favored over the other.1036

Once again Q can be very useful for determining what direction a reaction will shift if any.1044

Now we have come into one of the most fundamental principles from all of general chemistry.1055

This is called Le Chatelier's principle.1060

Le Chatelier's principle tells us that when a system at equilibrium is disturbed,1063

it will react to counteract the disturbance in an attempt to restore equilibrium.1068

One of the best examples we can think of is from us.1078

When we bleed, when we lose blood, what is the natural thing that our body does in order to counteract this?1082

Your body is going to try to make more blood.1091

When we lose blood, our bodies try to make more in order to counteract the lost.1092

This is a nice example of Le Chatelier's principle.1112

But now we are going to apply this to chemical reactions.1114

In Le Chatelier's principle, you notice that there is the word disturb.1119

There are several ways of disturbing a chemical system that is at equilibrium already.1124

Number one is a change in reactant or product concentration.1127

Number two if we are dealing with gases, it is going to be a change in reactant or product partial pressure.1132

The third one is the change in reaction volume.1137

The fourth one is going to be a change in temperature.1140

Let's now study each of these.1144

Basically if the concentration of reactant goes up, then we are going to shift away from it.1150

We are going to shift right.1163

If the concentration of product goes down, we don't have enough of it.1165

We are going to shift right; that is one situation.1170

If the partial pressure of a reactant goes up, partial pressure is just the same as concentration.1178

We know from ideal gas law that pressure is proportional to amount.1189

If partial pressure of the reactant goes up, we are going to shift right.1194

If the partial pressure of the product goes down, we are going to shift right.1204

For temperature, I am going to do this on the last slide because I am running out of room right now.1215

I will come back to changes in temperature and also to change in vessel volume.1228

But now let's get into some problem solving with some ICE tables.1235

We will now approach chemical equilibrium from a quantitative view.1238

An ICE table allows for one to study a chemical system at three points in time.1241

Basically what are all reactant and product initial amounts?1246

That is what the I stands for.1250

During the reaction on the way to equilibrium, what are their changes in concentration/pressure?1252

That is what the C stands for.1256

Finally what are their final volumes once equilibrium has been achieved?1258

That is what the E stands for.1262

We are going to look at this right now.1263

Consider 2 water gas goes to 2H2 gas plus O2 gas.1266

K is equal to 2.4 times 10-5 at some temperature.1271

At equilibrium, the concentration of H2O gas is 0.11 molar.1275

The concentration of H2 gas is 0.019 molar; calculate O2 at equilibrium.1279

The very first thing we want to do is set up our ICE table.1285

2H2O gas goes on to form 2H2 gas plus O2 gas.1289

What I like to do, I just like to set up the letters I-C-E right underneath it.1296

We just fill in this table right now.1300

You are told that at equilibrium, the water is 0.11 molar and that the H2 is 0.019 molar.1303

We don't know what O2 is.1313

We can just call that the concentration of O2 at equilibrium.1315

We know the expression for K is equal to 2.4 times 10-5.1322

That is going to be equal to the concentration of H2 squared times1328

the concentration of O2 divided by the concentration of H2O squared.1332

That is going to be equal to 0.019 squared times the concentration of O2 at equilibrium divided by 0.11 squared.1338

When all is said and done, the concentration of O2 at equilibrium is going to be 8.0 times 10-4 molar.1349

Don't forget the units.1359

You notice that we didn't have to fill in the rest of the table1363

because we were already at halfway there to the equilibrium values.1366

This is a nice usage of the ICE table.1371

Let's go ahead and do another example though.1374

A 1 liter flask was filled with so much of SO2 and so much NO2 at some temperature.1377

After equilibrium was reached, 1.3 moles of NO was present.1382

The reaction is the following; calculate the value of Kc at this temperature.1385

The problem wants us to calculate Kc at this temperature.1390

We are going to go ahead and set up the ICE table--I, C, and E.1396

SO2 here is going to be 2.00 molar initially.1406

O2 is going to be 2.00 molar also.1411

After equilibrium was reached, 1.30 moles of NO gas was present.1415

That is what goes right here, 1.30 molar.1419

Let's fill in the rest of the table.1426

If no initial values are mentioned of SO3 and NO, it is safe to assume that they are zero.1427

Let's go ahead and do the change.1434

The SO2 is going to go down by some amount x.1438

O2 is going to go down by some amount x.1441

SO3 goes up by some amount x.1443

O is going to go up by some amount x.1452

We notice because these are all 1:1 mole ratio.1457

At equilibrium, it is just going to be the addition of the two rows added together; sum.1465

That is going to be 2.00 minus x, 2.00 minus x, x, and x.1472

But guess what? x is 1.30 molar because we were told that in the beginning; that is very nice.1479

Therefore this is 1.30 molar; this is 0.70 molar; this is 0.70 molar.1485

We know enough to calculate the value of Kc.1494

Kc is going to be 1.30 squared divided by 0.70 squared.1498

That is going to give us an answer of 3.5 for our answer.1504

This is how we use the ICE tables to help us come up with the equilibrium constant.1517

I want to quickly summarize this and then jump into sample problems.1526

Chemical equilibrium is a dynamic process where the rate of forward and reverse reactions is equal.1531

Equilibrium constant K quantifies how reactant or product favored a reaction is once it has reached equilibrium.1536

If K is between 0 and 1, we say that it is reactant favored.1543

If K is greater than 1, we say that it is product favored.1552

Le Chatelier's principle states that a system at equilibrium when disturbed will react to counteract the disturbance.1559

We already saw the concentration and partial pressures.1565

But let's go ahead and see what I left out before; this is now temperature.1573

For an exothermic reaction, if the temperature increases, it turns out we are going to shift left.1579

If the temperature decreases, we are going to shift right.1593

Basically you treat the word heat as a molecule.1599

Again if you treat the word heat as a molecule on the product side, it becomes1609

a lot more intuitive as you will see what happens when we result in the change.1616

Another thing we can disturb an equilibrium with is with the volume.1622

If the volume of the reaction vessel goes up, that means the pressure is lower.1629

That means we are going to shift to make more gases.1636

Shift toward side of reaction with more gas molecules.1640

Once again if the volume goes up, we are going to shift toward the side of the reaction with more gas molecules.1654

Those are the ways we can disrupt a system that is at equilibrium and see how it is going to counteract the stress.1662

Finally ICE tables allow for determination of the reactant and product concentrations/pressures given the equilibrium constant K.1668

That is our summary.1676

Now let's go ahead and jump into a pair of sample problems.1678

Consider 2SO3 gas goes to 2SO2 plus O2.1681

Initially 12 moles of SO3 is placed into a 3 liter flask at some temperature.1685

At equilibrium, 3 moles of SO2 is present.1689

Calculate the value of Kc at this temperature.1692

Let's go ahead and rewrite this.1696

2SO3 gas goes on to form 2SO2 gas plus O2 gas.1698

Let's go ahead and see what we can fill in, I-C-E.1707

Initially 12 moles is placed in a 3 liter flask of SO3.1712

That is going to be 4.0 molar.1716

At equilibrium, SO2 is present, 3 moles; that is going to be 1.0 molar.1721

Usually you can fill in two-thirds of the table with just one or two sentences.1728

SO2 initial is not mentioned; that is zero.1733

O2 initial is not mentioned; that is zero also.1735

We have to pay attention to the coefficents.1740

The coefficients tells us the relative amount of reactant and product that is going to be involved.1742

Really this is going to be -2x; SO2 is going to be +2x.1747

O2 is just going to be +x.1752

At equilibrium, we get 4.0 minus 2x, 1 molar, and x.1756

The nice thing about this is that because we know the 2x and the 1 molar, we can conclude that x is 0.5 molar.1763

When we do that, we are going to get here 0.5 molar.1776

Here for SO3, we are going to get 3.0 molar.1780

Great, we now have all of our molarities at equilibrium.1786

Kc is equal to the concentration of SO2 squared times the concentration1791

of O2 divided by concentration of SO3 squared, all equilibrium values.1797

This is going to be 0.5 squared times 0.5 divided by 3.0.1805

When all is said and done, this is going to give us a Kc of 0.056 for this reaction.1815

That is another usage of the ICE table.1827

Finally the last sample problem is sample problem number two.1831

At a particular temperature, K is equal to 3.75.1834

If all four gases had initial concentrations of 0.800, calculate their equilibrium concentrations.1837

Let's go ahead and set it up--I, C, and E.1843

0.800 molar initial, 0.800 molar, 0.800 molar, and 0.800 molar.1848

Because we have nonzero amounts of all reactants and products, we don't know which way we are going to shift.1858

This is a twist on things; now we have to solve for Q.1868

Q is going to tell us which way we are going to shift.1872

This is going to be equal to concentration of SO3 initial times the concentration of1875

NO initial divided by concentration of SO2 initial times the concentration of NO2 initial.1882

That is just going to be equal to 1, 1.00 which is going to be less than K.1891

Because Q is less than K, that means we are going to shift to the right.1900

Therefore the sign that goes here is going to be ?x.1907

This is ?x; this is +x; this is +x.1911

We are going to get 0.800 minus x; 0.800 minus x.1916

This is going to be 0.800 plus x.1923

This is going to be 0.800 plus x; K is equal to 3.75.1926

That is equal to 0.800 plus x squared divided by 0.800 minus x squared.1934

We can use what we call a perfect square to solve for this.1946

The square root of both sides is what we are going to do next.1951

When we go ahead and do that, we are going to get 1.93 is equal to 0.800 plus x divided by 0.800 minus x.1956

When all is said and done, x is going to be equal to 0.25 molar.1966

All we do next is simply plug it back into all of the expressions at equilibrium.1972

The concentration of SO3 at equilibrium is equal to the concentration of NO at equilibrium.1978

That is going to be 0.800 plus 0.25.1986

I will let you guys do the sum; that is going to be 1.050 molar.1992

The concentration of SO2 at equilibrium is equal to the concentration of NO2 at equilibrium.2004

That is going to be equal to 0.800 minus 0.25.2012

We are going to get 0.55 molar for these guys.2020

That is use of the ICE table when we have all four initial values present.2026

Once again you have to find which way it is going to shift.2034

You do that by evaluating Q.2036

That is our lecture from general chemistry on the principles of chemical equilibria.2039

I want to thank you for your time.2045

I will see you next time on Educator.com.2046

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