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Franklin Ow

Franklin Ow

Chemical Kinetics

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 (25)

0 answers

Post by Nasser Fiture on January 12 at 01:55:24 AM

If a proposed reaction mechanism fits the experimental rate law and reaction stoichiometry, then the mechanism is known to be completely correct.
True/False

0 answers

Post by Nasser Fiture on January 12 at 01:54:20 AM

all reactions which show first-order kinetics are also known as unimolecular reactions.

0 answers

Post by Nasser Fiture on January 12 at 01:53:31 AM

the rate determining step of a chemical reaction is the step which has the transition state highest in potential energy.
True/False

0 answers

Post by Jarrah alharbi on September 18, 2016

is the ----> t1/2=In1/2 / k

or t1/2=In2/k

2 answers

Last reply by: Professor Franklin Ow
Mon Jun 20, 2016 11:35 PM

Post by Parth Shorey on June 20, 2016

I don't understand why you used the negatives in front of the recants? You said because of the slopes? I still don't understand.

1 answer

Last reply by: Professor Franklin Ow
Thu Jun 16, 2016 3:29 PM

Post by Parth Shorey on June 13, 2016

Is the Q&A active?

1 answer

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

Post by BRAD POOLE on May 7, 2015

At about the 17 min mark you said that the example was a 2nd order reaction.  How did you come up with this?  Are you just going by whatever your units are for "k"?  I always thought you used the exponents of the reactants to figure out what order it was, then again could be why I can't seem to get these kinetics problems.  

0 answers

Post by Saadman Elman on January 14, 2015

Great lecture as usual!

0 answers

Post by Saadman Elman on January 14, 2015

What you mean by we can write co-effecient as rate order ONLY for elementary steps. What you mean by elementary steps exactly?

2 answers

Last reply by: David Gonzalez
Thu Jul 31, 2014 12:25 PM

Post by David Gonzalez on July 31, 2014

Hi Professor Ow. First of all, great lecture. Although, there is one problem that I have.

In the example (around the 16-minute mark), when determining the order for S2O8, you mentioned that there was a 3x increase in initial rate from 0.015 to 0.044. I'm confused, because 0.015 x 3 is 0.045 - don't these problems need to be exact? Or can they sometimes be slightly off?

Thanks.

1 answer

Last reply by: Professor Franklin Ow
Tue Jun 24, 2014 1:29 PM

Post by brandon joyner on June 24, 2014

Also why for Iodine did you only do the first experiment for the second experiment you went from experiment 3 to 1.

4 answers

Last reply by: brandon joyner
Fri Jun 27, 2014 1:46 PM

Post by brandon joyner on June 24, 2014

For Iodine how is it a triple jump when all you had to do was go up 2?

1 answer

Last reply by: Professor Franklin Ow
Wed May 21, 2014 1:54 AM

Post by Ashley Gwemende on May 20, 2014

Where do I find a lecture of how to read a potential energy diagram ?

Chemical Kinetics

  • Kinetics studies the factors that can influence the rate of a chemical reaction.
  • The method of initial rates and the use of integrated rate laws can help solve for the rate orders and for the rate constant.
  • The slowest step of a reaction mechanism dictates the overall reaction rate.
  • The Arrhrenius equation relates the temperature of a reaction directly to its reaction rate.

Chemical Kinetics

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 0:06
  • Introduction 1:09
    • Chemical Kinetics and the Rate of a Reaction
    • Factors Influencing Rate
  • Introduction cont'd 2:27
    • How a Reaction Progresses Through Time
    • Rate of Change Equation
  • Rate Laws 7:06
    • Definition of Rate Laws
    • General Form of Rate Laws
  • Rate Laws cont'd 11:07
    • Rate Orders With Respect to Reactant and Concentration
  • Methods of Initial Rates 13:38
    • Methods of Initial Rates
  • Integrated Rate Laws 17:57
    • Integrated Rate Laws
    • Graphically Determine the Rate Constant k
  • Reaction Mechanisms 21:05
    • Step 1: Reversible
    • Step 2: Rate-limiting Step
    • Rate Law for the Reaction
  • Reaction Rates and Temperatures 26:16
    • Reaction Rates and Temperatures
    • The Arrhenius Equation
  • Catalysis 30:31
    • Catalyst
  • 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

Transcription: Chemical Kinetics

Hi, welcome back to Educator.com.0000

Today's lecture from general chemistry is going to be on chemical kinetics.0002

Here is our brief overview of the lesson today.0008

As always we will go ahead and start off with our introduction.0010

Then we are going to get into what is really at the heart of chemical kinetics.0013

That is the rate of chemical reactions; that is how fast does a reaction occur?0017

We are then going to see what are the mathematical equations which actually attempt to quantify the rate of an equation.0021

That is what we call a rate law.0029

Our objective is to then try to derive these equations for any given reaction.0031

There is two sets of experiments that you can do to perform this.0037

Number one is what we call the method of initial rates.0040

Number two is what we call the integrated rate laws.0043

We will then jump into how a reaction proceeds, the step by step process which is0045

called a reaction mechanism, followed by the effect that temperature has on reaction rates.0051

We will go ahead and wrap up the lesson with a brief overview of0059

catalysis followed by our summary and a pair of sample problems.0062

Chemical kinetics basically is the area of chemistry that examines the factors which can influence the rate of a reaction.0070

What do we mean by factors that can influence the rate?0081

There are several--temperature, pressure, the reactant concentration, the addition of a catalyst, and mechanical force.0084

We are going to see that basically as temperature goes up, the rate goes up.0095

We are going to see that in general as pressure goes up, the rate goes up.0104

We are going to see that as the concentration of reactant goes up, the rate goes up.0112

When we add a catalyst, we are going to define what that exactly means.0120

But this also increases the rate.0123

Finally mechanical force, this is what you probably do in lab.0126

All we mean by mechanical force is something as simple as stirring or shaking.0129

Of course this will increase the rate of the reaction.0137

This is what we mean by factors which can influence the rate of a reaction.0142

Let's take a look though how a typical reaction progresses through time.0150

Consider the reaction A plus 2B goes to C.0154

This is telling me that for every one of A, two Bs are required.0158

One C is going to be produced.0165

Remember we are only looking here at the forward direction only.0167

When we go in the forward direction, we say that reactants are consumed which means concentration of reactant go down.0176

The products are made which means that the concentration of product goes up.0190

It looks like because twice the amount of B is going to be consumed as A,0201

the rate of consumption of B is going to be double that of A.0206

It is very typical for us in chemical kinetics to graph this.0221

We can go ahead and graph the concentration of the reactant with respect to time and the concentration of product too.0225

At time zero, I have no product; I only have reactants.0237

Down here right at the origin, this is the concentration of C initial.0240

The way that concentrations change with respect to time is not linear.0247

But instead it is going to be characterized by a simple curve just like that.0251

You notice that at some time t, the concentration of C stops changing.0260

We are going to plateau; it is not going to increase forever.0264

How about the concentration of A and B?0269

The concentration of A let's say is here and the concentration of B.0272

Let's say we had equal amounts.0278

Let me go ahead and draw now a curve that represents the concentration of A changing with respect to time.0280

It is going to be a mirror image of the concentration of C changing because they are both 1:1 ratio.0287

But how is the concentration of B changing?0293

The concentration of B, it is going to drop at double the rate.0296

The concentration of B is going to drop much much much... quicker than the concentration of A.0302

You notice that for all of the reactions, there reaches a point where the curves plateau0315

which means that the concentration of reactants and the concentration of products no longer changes.0325

We are going to talk about this region in another lecture.0337

This is what we call the equilibrium region.0340

In chemical kinetics, what we are interested in is really the start of the reaction, the early part0346

of the reaction where we actually can monitor change in the concentration of reactants and products.0351

That is the representation graphically; how about mathematically?0360

It turns out that the rate of change is going to be equal to the following.0364

The change in the concentration of A, Δ[A]Δt, this is going to be equal to0368

1/2 Δ[B]Δt which is equal to the change in concentration of C Δ[C]Δt.0382

We have to put a negative sign in front of the reactants such that the overall number is0392

going to be positive because Δ[A]Δt is going to be slope of the curve.0401

This is going to be slope of the curve you see.0408

This is going to be the slope of that curve too.0411

This gives us a general equation for relating the rates of change of each reactant and product concentration with respect to time.0416

Now that we have gone through a brief introduction, let's go now into0429

what is at the core of chemical kinetics--that is the rate law.0434

A rate law is going to be a mathematical equation which relates the rate of reaction to the concentration of the reactants.0439

Basically the rate law tells us that the rate of a reaction is proportional to the concentration of reactants typically.0447

When we write out the rate law, all rate laws have the following equation0458

Of the form rate which is going to be usually in molarity per second...0461

is equal to some constant k times the concentration of A raised to0467

some power times the concentration of the B raised to some power, etc.0472

Let's go ahead and define what each of these mean.0478

k is what we call the rate constant.0484

The rate constant is going to be unique to a reaction at a certain temperature.0491

In addition, we are going to see that the rate constant, the units of k vary.0508

We are going to see that very soon.0519

x and y are what we call rate orders.0522

Once again x and y are what we call rate orders.0529

If x is equal to 1, we say reaction is first order with respect to the concentration of A.0537

If x is equal to 2, we say the reaction is second order with respect to the concentration of A.0554

In addition, if the sum of the rate orders x plus y is equal to 0, we say zero order overall.0564

If the sum of the orders is equal to 1, we say first order overall.0577

If x plus y is equal to 2, we say second order overall.0585

This is just some terminology that we want to introduce and clarify.0592

Let's go ahead and get back to the units of k.0598

The units of k will vary.0601

For zero order, the units of k is just going to be reciprocal seconds.0604

For first order, the units of k is going to... excuse me.0619

For zero order.. my apologies.0627

For zero order, the units of k is going to be molarity per second.0629

For first order, the units of k is just reciprocal seconds.0636

For second order overall, the units of k is going to be inverse molarity inverse second.0640

You can easily plug that back into the equation for k and see that the units will cancel.0649

This is the general rate law.0657

Rate is equal to some constant k times the concentration of A raised to0659

the x power times the concentration of B raised to the y power.0662

But let's go ahead now and see what the significance of the rate orders are.0666

Consider the following reaction: A plus 2B goes to C.0673

The rate law is given to us to be k times A squared times the concentration of B.0678

Let's go ahead and study what happens if we change the concentration of one of these reactants.0686

If the concentration of A doubles holding B constant, then we see that the rate is going to increase by 4.0692

It is going to quadruple.0710

If the concentration of A triples holding B constant, we see that the rate is going to increase by a factor of 9.0714

What if the concentration of B doubles holding concentration of A constant?0728

If that happens, the rate is just going to increase by 2.0739

If the concentration of B triples holding A constant, we see that the rate is going to increase by a factor of 3.0746

You see what the significance of the rate order is.0756

The rate order is really proportional to the sensitivity of a reaction rate on0760

the concentration of a specific reactant molecule, on the concentration of a specific reactant.0778

Pretty much we see that the larger the value of these rate orders, the more sensitive0788

your reaction is to a change in concentration of a specific reactant.0795

Basically what the kinetics deals with is focusing on this rate law and solving for x, y, and k.0802

In other words, our goal is to solve for the rate constant and rate orders.0811

We can do this in one of two ways.0820

The first experimental way to derive a rate law is what we call the method of initial rates.0822

In the method of initial rates, you basically do what we just did.0828

You change one reactant concentration holding all else constant and seeing how the rate varies.0834

Let's go ahead and take a look at the following.0857

You are usually given some table of data.0859

The table of data is going to list different concentrations, different molarities of each reactant and the rate that was measured.0865

For example, in experiment number one, this concentration of iodide was found.0873

This concentration of thiosulfate was found; this is the initial rate.0881

In experiment number two, we see that the concentration of thiosulfate was fixed.0886

The concentration of iodide was tripled; we see that the rate also tripled.0894

What does that mean?--basically as the concentration of iodide tripled, the rate tripled.0901

This is a 1:1 correspondence.0913

Anytime you have this 1:1 correspondence, it is a rate order of 1; rate order of 1.0917

Let's now see what the rate order is for S2O82-, thiosulfate.0927

For thiosulfate, we are going to look at experiments two and three.0933

You see that as the concentration of thiosulfate tripled... I'm sorry.0941

We have to do experiments one and three, not two and three.0953

Here as the concentration of thiosulfate tripled, iodide was held constant.0956

What happened to the rate?--the rate tripled.0966

As S2O82- tripled, the rate tripled holding everything else constant.0968

This was also a 1:1 correspondence; the rate order is also 1.0975

Therefore we have the following rate law.0986

The rate of this reaction is equal to some constant k times the concentration of I-0989

raised to the first power and the concentration of S2O82- raised to the first power.0994

The next thing now that we have the rate orders, we can go ahead and now solve for k.1002

We can solve for k by simply plugging in any experiment--one, two, or three--directly into the equation.1011

Solve for k using any experiment number.1019

For example, let's use experiment one; the rate was 0.044 molarity per second.1029

That is going to be equal to the rate constant k times the concentration of I- which is1037

0.125 molar times the concentration of thiosulfate, 0.150 molar, all raised to the first power.1041

Then you can just use your algebra to go ahead and solve for the value of k for this second order overall reaction.1050

Here we are going to get units of inverse molarity inverse seconds.1063

That is how we use the method of initial rates--very straightforward experiment to determine the rate law.1070

The second way of determining a rate law is to use what we call integrated rate laws.1078

Integrated rate laws, they quantify the relationship between the reactant concentration and time; and time.1085

Basically these are all derived mathematically.1095

You should always ask your instructor if you need to know how to derive it or not.1100

For zero order rate law, the rate is equal to the rate constant k.1104

The integrated rate law is the following.1108

The concentration of A at any given time is equal to the initial concentration of A minus kt.1111

For the first order overall, the integrated rate law is natural log of A0 minus natural log of A equal to kt.1116

Finally the second order integrated rate law is 1 over A minus 1 over A0 is equal to kt.1125

The nice thing about these equations is that they are all linear.1133

If several concentrations are determined at different times, you can get the rate constant graphically.1139

For a zero order overall reaction, we are basically going to graph the concentration of A as a function of t.1146

That is going to give us a nice straight line with slope equal to ?k.1156

For first order overall, we can go ahead and graph the natural log of the concentration of A versus t.1162

We are also going to get a straight line whose slope is equal to ?k.1174

For second order overall, we are going to plot 1 over the concentration of A which is equal to time.1179

Here we are going to get a nice straight line with a positive slope equal to k.1189

Once again this is a graphical determination of the rate constant anytime1194

you have data from several different time intervals and measured reactant concentration.1198

Another thing we like to talk about too is the half-life.1210

For zero order, the half-life is equal to the initial concentration of A over 2k.1214

For first order, the half-life is equal to the natural log of 2 over k.1221

For second order, the half-life is equal to 1 over k times the concentration of A0.1228

Once again you should ask your instructor to make sure if you have to know the derivation or not.1235

Basically the half-life is very important because it tells us the time required to reach 1/2 of the initial concentration of A.1241

Again that is what we call integrated rate laws.1261

The next thing we are going to look at is what we call a reaction mechanism.1267

A reaction mechanism basically represents the step by step reactions which when combined give you the overall net reaction.1271

For example, let's say we had the following given, step one.1278

Step one was 2NO gas going on to form N2O2 gas.1285

This tells us that the reaction is occurring both in the forward and reverse directions.1293

This is what we call the reversible reaction; this is usually very very fast.1299

Step number two is going to be O2 gas plus N2O2 gas going on to form 2NO2 gas.1305

You are told that this reaction is very very slow.1318

The overall reaction is going to be 2NO gas plus O2 gas going on to form 2NO2 gas.1323

You notice that N2O2 gets cancelled out.1336

Any item in a chemical reaction mechanism that gets cancelled is what we call an intermediate.1342

That is it is both formed and consumed during the course of a chemical reaction; formed and consumed during a reaction.1351

The reason why we care about what the slow step is is because of the following.1370

The slow step, which is in this case step two, is like the weakest link in your chain.1376

It is like the slowest person on your track and field relay team.1382

The slow step determines, it limits how fast a reaction can go.1387

We call this the rate limiting step.1392

The rate limiting step is equal to the rate of the overall reaction.1399

If we were to write out the rate law for this, we would get the following.1408

The rate is equal to the rate constant k2...1412

This is k2, of the second step.1417

Times the concentration of O2 raised to the first power1419

times the concentration of N2O2 raised to the first power.1422

We can always use the coefficients as the orders if the reaction you are1427

looking at is a part of the mechanism, what we call an elementary reaction.1438

Coefficient is equal to the rate orders for elementary steps only.1444

Otherwise you would have to go through integrated rate laws or method of initial rates1451

to go through the whole process again to find what x and y are.1454

When we look at this, we have a problem.1458

We have N2O2 appearing in this rate law.1460

This is the intermediate.1463

You can never have an intermediate appearing in the rate law.1465

We have to do something about this; what we do is the following.1468

We use what is called a steady state approximation.1473

You utilize the fast equilibrium step where the k1 times the concentration of NO squared1481

equals to k-1 times the concentration of N2O2 where k1 represents the forward.1492

k-1 represents the reverse.1502

What we do then is we solve for the intermediate.1505

Concentration of N2O2 is equal to k1 over k-1 times the concentration of NO2 squared.1507

We then plug this back into our experimental rate law.1518

Rate is equal to k2 times the concentration of O21522

times k1 over k-1 times the concentration of NO squared.1526

You can collect all of the constants together and just call that what we call kobserved.1534

We get left with O2 times the concentration of NO squared.1539

Here we have our final rate law that is going to be the rate law for the overall reaction.1546

It is just by coincidence here that the rate orders are the coefficients.1555

That is not usually the case.1559

But that is how we solve for it where kobserved is equal to k2k1 over k-1.1560

Once again this is how you deal with reaction mechanism problems.1571

Let's now move on to the next topic.1577

This is the relationship between temperature and reaction rate.1579

Basically in general, as temperature increases, so does the reaction rate; just think about this.1583

You know the tea bag is going to brew faster in warm water than cold water.1588

You can see that visually happening right before you.1594

In order for a reaction to occur, reactant molecules must do two things.1597

They must collide; they must collide with sufficient energy.1602

They must collide in the proper orientation; sufficient energy and proper orientation.1606

Basically we can look at a sample here.1613

Let the y-axis be fraction of sample; let the x-axis be temperature.1616

At any given temperature, I am going to have a bell curve distribution of molecules just like that.1628

Let's call this T1.1639

What happens to T2?--what happens when we have a hotter temperature?1641

When I have a hotter temperature, my bell curve is going to shift just like that.1649

I call this T2; T2 is greater than T1.1652

Let's say that in order for the sufficient energy, in order for the1658

reaction to occur, let's go ahead and make that as a dotted line.1663

I am going to call this dotted line EA; EA is equal to activation energy.1669

What this is, it is the minimum energy required for the reaction to proceed, for collisions to occur.1677

Basically anything below EA, anything less than EA, you get zero collisions and no reaction.1691

Anything greater than or equal to EA, you get collisions; therefore a reaction will occur.1704

Basically you see that as you go from T1 to T2,1712

the fractional molecules with an energy greater than EA significantly increases.1717

At T2, larger percent of molecules with an energy greater than or equal to EA.1723

This is why as temperature goes up, so does the rate of a reaction in general.1739

We have a nice equation which can actually quantify this.1749

This is called the Arrhenius equation.1754

The natural log of k1 over k2 equals to EA over R times 1 over T2 minus1756

1 over T1 where R is our universal gas constant in terms of energy, 8.314 joules per K times mole.1766

Temperatures T1 and T2 are kelvin temperatures; what this basically says is the following.1779

If I do a series of reactions at different temperatures and I calculate the1789

rate constant, I can then do graphical determination of the activation energy.1794

If I graph natural log of k1 over k2 as a function of 1 over temperature,1800

I am going to get a nice straight line with slope equal to ?EA over R.1809

Again the Arrhenius equation is very useful because it gives us graphical approximation of the activation energy for a reaction.1815

Again this is the relationship between rate and temperature.1829

Finally the last factor we are going to study is a catalyst.1833

Basically a catalyst's job is to do the following.1837

A catalyst assists reactant molecules to be in the proper orientation for proper collision to occur.1840

What that does is that the activation energy is lower.1847

If this reaction represents without a catalyst, the activation energy is going be basically right here.1853

This is the energy that you must overcome for the reaction to go form products.1867

I can then proceed and draw another curve where this is a catalyst now.1873

With the catalyst, you see that the activation energy EA is much lower.1878

Activation energy catalyst is going to be always much less than the activation energy no catalyst.1884

Because of that, with the lower activation energy, that means a faster reaction.1892

A catalyst again speeds up a reaction by lowering the activation energy.1900

We see that of course the nice about catalysts is that they can be reused over and1907

over again because during the course of a reaction, they are recovered; they are recovered.1913

That is catalysis.1921

Let's now get into our summary before we jump into our sample problems.1923

Kinetic studies the factors that can influence the rate of a chemical reaction.1927

We saw that we can have two main experiments to help us determine1932

the rate law--the method of initial rates and the integrated rate law.1936

We found that in a reaction mechanism, that the slowest step dictates the overall reaction.1940

That is what we call the rate limiting step.1946

Finally we introduced the Arrhenius equation which gives a mathematical relationship between the reaction rate and the temperature.1952

The nice thing about this equation again, this gave us graphical estimate of EA.1963

Let's now get into a pair of sample problems.1975

A certain first order reaction has a half-life of twenty minutes.1977

Calculate the rate constant; that is part A.1980

Part B, how much time is required for this reaction to be 75 percent complete?1984

Let's go ahead; you are told that the reaction is first order.1989

For a first order reaction, T1/2 is equal to the natural log of 2 over the rate constant k.1992

You are told that the rate constant is 20.0 minutes.2001

Here the rate constant is simply going to be equal to the natural log of 2 divided by 20.0 minutes.2007

That gives us our answer in units of reciprocal minutes; that is part A.2014

Let's go ahead and do part B now.2024

How much time is required for this reaction to be 75 percent complete?2026

As soon as you see the word time, you should immediately, immediately, think integrated rate law because2029

method of initial rates does not have time in it; only integrated rate laws does.2036

For the first order integrated rate law, it is the natural log of the initial concentration2041

of A minus the natural log of the concentration of A is equal to kt.2047

We know what k already is because we already solved for that in part A.2056

We are good to go on that.2060

The question is asking for how much time is required.2062

This is what we are trying to solve for.2065

All that matters is is what is the identity of A0 and what is the identity of A?2068

You are told that how much time is required for this reaction to be 75 percent complete?2074

Let's say A0 is going to be 100.2079

If the reaction is 75 percent complete, that means only 25 percent of A is remaining.2083

A is going to be 25.2090

Again we now have enough information to solve for t.2092

We are going to get our final answer of t in units of minutes.2097

How do you know if you have done something wrong?2105

You should always check your answer because if you get a ?t, again that just doesn't make physical sense.2107

You know you have done something mathematically wrong.2119

Always check your answer for a negative time.2121

That is sample problem number one.2125

Let's go ahead and move on to sample problem number two.2127

Consider the following reaction.2130

2N2O5 gas goes on to form 4NO2 gas plus O2 gas.2132

Here we are given several rate constants that were found at several temperatures.2137

The only equation that we know that deals with this is the Arrhenius equation.2145

The natural log of k1 over k2 is equal to2150

EA over R times 1 over T2 minus 1 over T 1.2154

We know that we are going to be using this equation quite easily.2161

Here we can calculate EA from there.2165

The nice thing about this is because this is going to give us a nice straight line,2170

we can use any pair of k and T data points to go ahead and solve for EA.2174

Once again this is going to be using the Arrhenius equation to solve for EA.2193

Our units of EA is going to be in kilojoules per mole.2199

Next one is what is the order of the reaction.2206

This is kind of a trick question; this is something I have asked students before.2208

This is something I have seen asked by other instructors before.2213

The order of the reaction, you don't have to do any work for that.2217

The reason is because they already give you the units for k.2219

The units of the rate constant tell us the rate order.2226

It is only first order where the units of k is reciprocal time; first order overall.2234

Again just watch out for that when you do problems.2248

Again the units of k tell us a great deal of information without doing any work.2251

That is our lecture from general chemistry concerning kinetics.2257

I want to thank you for your time.2262

I will see you next time on Educator.com.2263

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