Enter your Sign on user name and password.

Forgot password?
Sign In | Sign Up
Start learning today, and be successful in your academic & professional career. Start Today!

Use Chrome browser to play professor video
Franklin Ow

Franklin Ow

Intermolecular Forces & Liquids

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
Loading...
This is a quick preview of the lesson. For full access, please Log In or Sign up.
For more information, please see full course syllabus of General Chemistry
  • Discussion

  • Study Guides

  • Download Lecture Slides

  • Table of Contents

  • Transcription

  • Related Books & Services

Lecture Comments (8)

1 answer

Last reply by: Professor Franklin Ow
Tue Apr 7, 2015 11:57 PM

Post by Akilah Miller on April 7, 2015

Hello Professor,

I am a bit confused: on slide 5 you say that r represents the atomic size but does it not represent the distance between the the molecule and/or atoms that interact with each other?

1 answer

Last reply by: Professor Franklin Ow
Sun Feb 15, 2015 11:44 PM

Post by Anthony Linares on February 9, 2015

Greetings. am new to educator.com and I was just wondering why the video is not working.

1 answer

Last reply by: Professor Franklin Ow
Tue Oct 14, 2014 6:55 PM

Post by Luisa Gualtieri on October 13, 2014

So for long alkane chains, it is the incrase in molar mass that has the greater effect on boiling point, and not the length of the chain making an increased separation and stretching it thin in terms of dipole-dipole interaction between nuclei. Am I correct?

1 answer

Last reply by: Professor Franklin Ow
Wed May 7, 2014 6:46 PM

Post by Heather Marck on May 7, 2014

How could you tell so quickly that H2S had dipole-dipole and V.d.W?  

Intermolecular Forces & Liquids

  • The main types of IMF are ion-induced dipole, dispersion, dipole-dipole, and hydrogen bonding.
  • IMF can exert strong influence on a liquid’s physical properties, such as boiling point and vapor pressure.
  • By using IMF, we can make a strong educated guess of the relative volatility of a liquid.

Intermolecular Forces & Liquids

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 0:10
  • Introduction 0:46
    • Intermolecular Forces (IMF)
  • Intermolecular Forces of Polar Molecules 1:32
    • Ion-dipole Forces
    • Example: Salt Dissolved in Water
    • Coulomb's Law & the Force of Attraction Between Ions and/or Dipoles
  • IMF of Polar Molecules cont'd 4:36
    • Enthalpy of Solvation or Enthalpy of Hydration
  • IMF of Polar Molecules cont'd 6:01
    • Dipole-dipole Forces
  • IMF of Polar Molecules cont'd 7:22
    • Hydrogen Bonding
    • Example: Hydrogen Bonding of Water
  • IMF of Nonpolar Molecules 9:37
    • Dipole-induced Dipole Attraction
  • IMF of Nonpolar Molecules cont'd 11:34
    • Induced Dipole Attraction, London Dispersion Forces, or Vand der Waals Forces
    • Polarizability
  • IMF of Nonpolar Molecules cont'd 14:26
    • Intermolecular Forces (IMF) and Polarizability
  • Properties of Liquids 16:48
    • Standard Molar Enthalpy of Vaporization
    • Trends in Boiling Points of Representative Liquids: H₂O vs. H₂S
  • Properties of Liquids cont'd 18:36
    • Aliphatic Hydrocarbons
    • Branched Hydrocarbons
  • Properties of Liquids cont'd 22:10
    • Vapor Pressure
    • The Clausius-Clapeyron Equation
  • Properties of Liquids cont'd 25:52
    • Boiling Point
  • Properties of Liquids cont'd 27:07
    • Surface Tension
    • Viscosity
  • 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

Transcription: Intermolecular Forces & Liquids

Hi, welcome back to Educator.com.0000

Today's session in general chemistry is going to be concerning intermolecular forces and liquids.0002

As always we are going to start off with a brief introduction.0012

Then we are going to talk about the intermolecular forces first of polar compounds and finally of nonpolar compounds.0015

After we talk about the intermolecular forces in these classes of compounds, we will then0024

get into properties of liquids and how intermolecular forces affect these properties.0030

As always we will go ahead and do a brief summary of the presentation followed by some sample problems.0038

For the introduction, what we are looking at is basically the attractive forces that hold molecules together.0048

If you have a molecule and an identical molecule comes along right beside it,0056

what is responsible for those two molecules interacting with each other or attracting each other?0063

We call these forces of attraction intermolecular forces.0067

For my presentation, I am going to be abbreviating intermolecular forces as IMF.0073

What we are going to see is a good application of IMF.0080

We will see that the strength of a liquid's IMF can easily influence how it behaves physically.0085

Let's now get into the intermolecular forces for polar molecules.0092

The first type of IMF for polar molecules is what we call an ion-dipole force.0098

An ion-dipole force is going to occur between a polar molecule and an ionic compound.0104

For the example that I want to look at on the particulate level, this is going to be table salt dissolved in water.0111

Here I have water and table salt.0118

As soon as I dissolve table salt in water, the sodium chloride is going to break up into ions.0126

We have Na+ and Cl1-.0131

We have already discussed polarity previously.0136

We know that water is partial positive here, partial positive here, and partial negative there.0139

Something interesting is going to happen.0146

The partial negative end of oxygen is going to get attracted to the full positive ion that is sodium.0148

In addition, the partial positive end of water is going to get attracted to the full ion that is chloride.0160

This is what we call an ion-dipole force.0171

It is basically the attraction between an actual ion that has a full charge and a compound that has a dipole moment.0175

Coulomb's law describes the force of attraction between ions and/or dipoles.0188

Basically Coulomb's law tells us that this attraction is going to0194

be proportional to the following: q1q2 over r20200

where q1 and q2 are the relative charges of the cation and anion; charges of cation and anion.0210

r is just the distance between the two nuclei; ionic distance.0227

We see a very important proportionality.0237

That the force of attraction is going to be directly proportional to the magnitude of the charges.0242

It is going to be inversely proportional to 1 over r2 or inversely proportional to the atomic size.0254

Once again this is what we call Coulomb's law.0264

It tells us, it really quantifies the degree of attraction between charges and/or dipoles.0267

There is one important physical parameter that describes what we just illustrated.0280

That is an ion interacting with a solvent.0285

This is what we call the enthalpy of solvation which is also known as the enthalpy of hydration.0290

Basically what it is is the following.0299

It is another type of ΔH like we learned before.0302

It is also a ΔH of solvation.0307

It describes the energy change when an ion becomes solvated.0311

In other words, when an ion becomes surrounded by solvent molecules.0323

What we want to know, remember mother nature favors low energy.0336

The more negative this energy, the more likely an ion will be dissolved or solvated by water.0339

An ion is more likely to be solvated by water the larger it is and the more positive the charge.0346

Let's go ahead and move on then.0358

Ion-dipole was the first type of IMF for a polar compound.0364

Let's now move on to the second type.0369

The second type of IMF that can occur in liquids is what we call dipole-dipole forces.0372

Dipole-dipole forces occur between molecules that have a permanent dipole moment.0379

Then the strength of the dipole-dipole force tends to increase with overall polarity.0386

For example, if we take a carbon-hydrogen bond versus an OH bond,0391

the OH bond is the more polar of the two and therefore will have a stronger IMF.0397

How does this all relate to boiling point?0411

Basically the stronger the IMF, the more energy required to overcome the attraction.0415

When we think of energy, we can think of boiling point.0421

Boiling point is the temperature at which the boiling process is going to occur.0425

Remember temperature is the measure of average kinetic energy.0428

Really when we say the more energy required to overcome the attraction, we really mean a higher boiling point.0432

This was again dipole-dipole forces.0442

The next type of intermolecular force is what we call hydrogen bonding.0445

Hydrogen bonding you can think of as a very strong type of dipole-dipole.0450

This is going to occur in a molecule that has hydrogen that is directly bonded to nitrogen, oxygen, or fluorine.0454

You will recognize these three elements being the most electronegative.0462

Therefore their bond with hydrogen will be the most polar.0467

Hydrogen bonding is important of course biologically because we are going to find out that liquids that0474

have hydrogen bonding tend to be number one, nonvolatile, and number two, have relatively higher boiling points.0480

You can think of water as our typical example.0487

Hydrogen here, partial positive, partial positive, partial negative.0493

We can have another water molecule come into play just right there.0497

We can have an interaction between the partial negative oxygen with the partial positive oxygen.0502

You can imagine another water molecule just like that; we can have even more attraction.0508

What we see is basically that we get a network of hydrogen bonding; network of strong hydrogen bonding.0520

Why is this biologically important?0537

We know water because of its strong hydrogen bonding is going to be nonvolatile.0539

Imagine if all the world's oceans all of a sudden evaporated.0546

Life would be unbearable; our planet would be too hot.0551

Because of hydrogen bonding, water does not evaporate.0554

That allows us to have a good cooling effect by the ocean.0557

Again that is also going to lead to a high boiling point.0563

Once again it is in our best interest to have water to be a very stable compound.0568

Again it is due to hydrogen bonding.0575

What we are going to go into now, now that we are done with polar molecules,0579

let's get into the intermolecular forces for nonpolar molecules.0583

We are first going to look at the solubility of a nonpolar gas in a polar solvent.0588

Wait a second, I thought polar and nonpolar are not supposed to mix.0596

But it turns out that for gases, there can be a slight degree of mixing.0599

The typical example is going to be carbon dioxide in water which is basically your average carbonated beverage like soda.0605

We all know that CO2 is nonpolar.0614

Let's go ahead and look at its Lewis structure.0619

Each oxygen is partial negative; the carbon is partial positive.0623

What can happen is the following.0629

A water molecule can come into play right here just like that.0632

Partial positive, partial positive, and partial negative.0641

Even though carbon dioxide is nonpolar, at any one point in time, at any point in time,0646

an electron cloud can get slightly distorted; electron density can distort.0660

As soon as we get a distortion, that instant we form a temporary dipole moment.0672

It is that temporary dipole moment which can then interact with the polar water molecule.0685

Let's go ahead and take a look at the interaction between two nonpolar molecules.0698

The interaction between two nonpolar molecules is what we call a London dispersion force or a van der Waals force.0703

Another name for it is induced dipole attraction.0712

Let's go ahead and take a look at molecular fluorine as an example.0716

Here is our typical Lewis structure for a molecular fluorine.0722

Because it is homonuclear diatomic, it is nonpolar overall.0726

However, at any point in time, we can have distortion of the electron cloud once again; distortion of electron density.0730

You can imagine then that at any one point in time, if what I am drawing represents the F2 molecule,0746

we can get one side to be partial positive and we can get one side to be partial negative.0754

This is not going to last forever.0760

It is a temporary dipole moment just like it was for carbon dioxide.0761

Basically if I have this temporary dipole moment, another F2 molecule can come into play, can come right alongside it.0770

It too can also have a temporary dipole moment.0781

We can get an interaction between the partial positive end of one F20784

molecule with the partial negative end of the second F2 molecule.0789

Because of the temporary nature though of this interaction, we can also0794

infer that London dispersion forces are going to be incredibly weak.0799

In fact, dispersion forces are the weakest type of intermolecular force due to temporary nature of attraction.0806

This brings into play a very important concept.0829

You are going to see this concept in this chapter.0832

You are going to see this concept in organic chemistry quite a bit actually when you get to that level.0835

To what extent does an electron cloud or electron density get distorted?0841

The extents to which this happen is known as polarizability.0848

Basically the more polarizable the molecule, the easier it is to0854

have an induced dipole moment and therefore the stronger this attractive force.0860

Let's go ahead and look at the molecular halogens as our example.0866

F2, CL2, Br2, and I2.0873

Under standard conditions, fluorine is a gas; chlorine is a gas.0880

Bromine is a liquid; I2 is a solid.0885

Why is that though?--why are we going from gas to liquid to solid?0890

In other words, as we go down this column, why does the intermolecular forces increase?0894

Because that is what explains for going to a more condensed state from gas to liquid to solid.0906

Why does the IMF increase?--the answer is because of polarizability.0911

Iodine is the largest out of these molecules.0921

Because it is the largest, the electrons are not held as tightly.0927

Because the electrons are not held as tightly, we get easier distortion.0939

More likely for electron cloud distortion to occur.0945

Because it is more likely for the electron cloud distortion to occur, we can form stronger dispersion forces.0955

In other words, polarizability increases with molar mass.0976

Once again polarizability tends to increase with molar mass; that is it.0991

That pretty much explains why we change physical state as you go down the column of0997

the halogens from fluorine to iodine, from a gas all the way to a solid.1003

We finished our discussion on the attractive forces that can occur for liquids.1012

Let's go ahead and apply them now.1019

We are going to apply these to the physical properties of liquids.1021

There are several of them.1026

The first physical property that we are going to be talking about is what we call standard molar enthalpy of vaporization.1027

It is yet another energy; it is another type of enthalpy.1034

The definition of ΔH of vaporization is the energy required to vaporize one mole of a liquid under standard conditions.1038

But in order to vaporize a liquid, you have to overcome the intermolecular forces,1046

the IMF whether it be dipole-dipole, hydrogen bonding, or dispersion.1051

In other words, the stronger a liquid's IMF, the more energy required and the larger the ΔH of vaporization.1056

Let's go ahead and look at the trends and boiling points of representative liquids.1064

Here we will examine two inorganic liquids, water and hydrogen sulfide.1069

For water, let's look at the IMF; we have hydrogen bonding here.1075

For hydrogen sulfide, there is no hydrogen bonding; only dipole-dipole and van der Waals.1080

Because hydrogen bonding is going to be the strongest type of intermolecular force here,1089

we conclude that water has the stronger IMF and therefore the higher ΔH of vaporization.1094

That is how we apply intermolecular forces to solving these types of problems.1110

Let's go ahead and take a look at some representative organic liquids.1118

The first type of organic liquids I want to talk about are what we call aliphatic hydrocarbons.1122

Aliphatic hydrocarbons are basically straight-chain alkanes.1126

Here we have CH3CH2CH3 and CH3CH2CH2CH3.1131

First of all, if you recall the electronegativities of carbon and hydrogen are very similar.1137

In fact, they are nonpolar bonds; both of these molecules are nonpolar.1142

Because both of these are nonpolar, the only type of intermolecular force are van der Waals in each of them.1149

What is the only difference then?--the only difference is chain length or molar mass.1156

We see here that this has the longer chain which means higher molar mass.1162

Just like we said previously, as the molar mass increases, the polarizability1171

also tends to increase, therefore the strength of the intermolecular force.1176

Therefore this molecule here, CH3CH2CH2CH3, has the stronger IMF1180

because of intermolecular forces... excuse me... of the higher molar mass and stronger IMF.1186

The next two, CH3CH2OH versus CH3OCH3.1197

In this left molecule, we have hydrogen bonding and dipole-dipole.1202

In the right molecule, there is no hydrogen bonding; only dipole-dipole and van der Waals.1210

We are going to obviously pick the one with the hydrogen bonding.1221

That is going to give us the stronger IMF and therefore the higher ΔH of vaporization.1224

These are again just examples of what we call aliphatic hydrocarbons, straight-chain alkanes.1247

We now move on to branched out hydrocarbons.1252

For branched hydrocarbons, it is important to go ahead and draw the Lewis structures out.1257

This one we did already, CH3CH2CH2CH3.1263

Nothing is going on here; this is aliphatic.1268

But when we move to its counterpart on the right side, we are1271

going to get a new type of structure that looks like this.1275

This is what we call a branched alkane; this is no longer aliphatic.1283

What we have to take away from this is the following.1290

That any type of branching is going to lower the surface area of your compound.1292

If the surface area is lower, that means there is less area of the molecule to be exposed to an adjacent molecule.1304

In other words, the IMF is going to be lower also.1314

This molecule here that is branched has the weaker IMF and therefore the lower ΔH of vaporization.1320

Next another physical property is what we call vapor pressure.1331

Vapor pressure is going to be commonly measured in our traditional units1337

for a gas pressure which is atm, torr, and even kilopascal.1341

What vapor pressure is is the following.1347

It is going to be the pressure of a vapor that is directly above the surface of its liquid in dynamic equilibrium.1349

If you take a small beaker and you cover the beaker and you close it, let's say you have a liquid in here.1358

That liquid is going to go into a gas phase naturally.1368

But not all of it because as soon as the gas forms, it is going to go back and condense back down to a liquid.1375

This is in constant motion and not a static condition.1381

This is what we call dynamic equilibrium; it is a dynamic process.1387

Basically whenever we have this type of closed system, if we measure the pressure of1393

the vapor right above the liquid, that is what we call vapor pressure.1399

Now that we have a little drawing to help support the idea of vapor pressure, we can now put it into play.1408

Vapor pressure is actually going to parallel what you and I know as smell.1415

Basically a liquid with a strong odor, we are detecting its gas phase with our nose.1421

A liquid with strong odor means it vaporizes easily.1427

Its vapor pressure is going to be relatively high.1430

If you compare gasoline versus water, gasoline we know smells very strongly.1433

Water has no odor; we can conclude that gasoline has a higher vapor pressure.1445

The technical term for this with a higher vapor pressure is what we call volatile; volatile.1456

We say water is therefore is nonvolatile.1463

We do have an equation that allows us to quantify the relationship between vapor pressure and temperature.1473

This is what we call the Clausius-Clapeyron equation.1484

The Clausius-Clapeyron equation is basically the natural log of P2 over P1 which is equal to...1488

ΔH of vaporization over R times 1 over T1 minus 1 over T2.1507

All this is saying is the following.1515

If T2 is greater than T1, that automatically implies that P2 is greater than P1.1519

In other words, vapor pressure increases with temperature; P increases with temperature.1527

Just think about that.1535

A gas station is going to smell a lot worse on a warmer day.1537

It is because the vapor pressure of any gasoline that is dropped is going to be much higher than on a cooler day.1541

That is again the Clausius-Clapeyron equation.1550

The next physical property of liquids is what we call normal boiling point.1553

Normal boiling, we all know what boiling point is.1557

It is the temperature at which a liquid is going to boil.1560

But normal boiling point is specifically the process that occurs at atmospheric pressure.1562

Since boiling requires energy, a high boiling point implies that lots of energy is needed to boil a liquid.1569

A high boiling point means very strong IMF which also means1577

higher ΔH of vaporization which also means lower vapor pressure.1588

Let me go ahead and summarize that.1601

Strong IMF, high boiling point, high ΔH of vaporization, lower vapor pressure.1603

Simply knowing the intermolecular force, we can take an educated guess at a lot of the physical properties of a liquid.1618

Now onto the last physical properties we will cover.1630

This one is what we call surface tension.1634

Surface tension is something measurable.1636

It basically quantifies the tendency of a liquid to minimize its surface area.1639

To do this, it is going to form spherical drops.1644

A liquid with a relatively high surface tension easily forms spherical droplets1649

such as water which has very strong IMF due to hydrogen bonding.1655

Liquids with low surface tension, they tend not to form spherical droplets; nonspherical droplets.1663

Finally we reached the last physical parameter of liquids.1686

It is what we call viscosity; viscosity is also measurable.1693

It basically quantifies a liquid's resistance to flow.1698

What that really translates to is how thick a liquid is.1702

If you think of something like honey or like syrup, we say those are very viscous liquids.1706

They are very resistant to flow.1712

Basically if an IMF is very large, the viscosity tends to also be large.1714

That makes sense that the intermolecular force of attraction is large.1720

Liquid is not going to want to separate from each other.1724

It is going to want to stay intact and stay together.1726

In other words, it is going to be somewhat thick if you will or not as flow-like as water or something like that.1728

Again this is what we call viscosity.1741

We now reach our summary slide.1746

Basically in this chapter, we introduced the types of intermolecular forces that are relevant to liquids.1750

Number one is what we call ion-induced dipole; number two is what we call dispersion.1757

Number three is what we call dipole-dipole; number four is what we call hydrogen bonding.1762

Remember ion-induced dipole is going to be between a free ion and a polar molecule.1768

Dispersion is going to occur between nonpolar compounds.1780

Dipole-dipole and hydrogen bonding are going to be for polar compounds.1784

After we went over all of the different types of intermolecular forces, we then applied them1791

and used them to make educated guesses on how a molecule is going to behave physically.1796

We saw that the IMF can exert strong influence on a liquid's physical properties such as boiling point and vapor pressure.1803

In other words, by using IMFs, we can make a strong educated guess on the relative volatility of a liquid.1813

For sample problem one, let's go ahead and determine which of the following liquids will have the lower vapor pressure.1823

As soon as I see this molecule here which is OH, I see hydrogen bonding immediately.1830

That tells me this molecule is going to have relatively strong intermolecular forces.1836

Here CH3OCH3, that is going to look a lot like this.1841

I don't have any hydrogen bonding at all.1847

But I do have dipole-dipole and of course van der Waals or dispersion.1849

For here, CH3CH2CH3, there is none of that.1858

There is no dipole-dipole; there is no hydrogen bonding.1864

This molecule is completely nonpolar.1867

I am only left with van der Waals or dispersion.1871

What low vapor pressure means is that we are going to have strong intermolecular forces.1878

Therefore the molecule that is expected to have the lower vapor pressure is going to have the strongest intermolecular forces.1886

That is the CH3CH2OH right there.1893

Finally sample problem two, we are going to take the same compounds.1899

But now we are going to determine which of the following will have the largest ΔH of vaporization.1903

As soon as I see largest ΔH of vaporization, that means a lot of1912

energy is required to vaporize one mole of a liquid under standard conditions.1916

That tells me I am dealing with a liquid with relatively strong IMF.1920

Using the same rationale, it is going to be the CH3CH2OH again expected1927

to have the largest ΔH of vaporization because this has hydrogen bonding.1931

This one in the middle has only really dipole-dipole and dispersion.1938

This one here has only dispersion.1946

What I encourage you to do is when you tackle these types of problems,1951

you should always try to from the formula get the structure, the Lewis structure that is,1955

because that will really help you see any of types of bonds.1963

Step two is then go ahead and identify the types of IMF present.1966

Hopefully that is enough to distinguish the liquids from each other.1979

But if IMF identical, then what you need to go by is really1984

the size of the molecule or its molar mass; go by molar mass.1993

Again what was the word that describes the effect of molar mass and size on distortion of electron cloud?1999

This is all related to polarizability; good.2006

That was our general chemistry lecture on intermolecular forces and liquids.2015

I want to thank you guys for your attention.2021

I will see you next time on Educator.com.2024

Educator®

Please sign in for full access to this lesson.

Sign-InORCreate Account

Enter your Sign-on user name and password.

Forgot password?

Start Learning Now

Our free lessons will get you started (Adobe Flash® required).
Get immediate access to our entire library.

Sign up for Educator.com

Membership Overview

  • Unlimited access to our entire library of courses.
  • Search and jump to exactly what you want to learn.
  • *Ask questions and get answers from the community and our teachers!
  • Practice questions with step-by-step solutions.
  • Download lesson files for programming and software training practice.
  • Track your course viewing progress.
  • Download lecture slides for taking notes.

Use this form or mail us to .

For support articles click here.