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Franklin Ow

Franklin Ow

Bonding & Molecular Structure

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 (2)

0 answers

Post by xin jin on October 23, 2016

Is the definition of bond energy correct?
Shouldn't it be the energy required to break part the bonds?

0 answers

Post by Saadman Elman on November 28, 2014

You didn't talk about the molecular geometry of the very last sample problem which is Trigonal Bipyramid because it has 5 electronic group with no lone pair. Please let me know after you are done verifying it. It is the very last sample problem ----> PF5.

As usual the lecture was awesome. Especially you made the lewis structure extremely easy. I feel like you didn't put enough time in the concept of molecular geometry. So i was little rusty on that particular concept only. So i went to youtube for molecular polarity. But over all once again, very good lecture. You made this chapter easy and it was fun as well.

Related Articles:

Bonding & Molecular Structure

  • Lewis structures are depictions of covalent bonding in molecules and ions.
  • Formal charge can be used to help construct the best Lewis structures for any molecule/ion.
  • Using VSEPR theory, a Lewis structure can be used to predict molecular geometry and molecular polarity.

Bonding & Molecular Structure

Lecture Slides are screen-captured images of important points in the lecture. Students can download and print out these lecture slide images to do practice problems as well as take notes while watching the lecture.

  • Intro 0:00
  • Lesson Overview 0:08
  • Introduction 1:10
  • Types of Chemical Bonds 1:53
    • Ionic Bond
    • Molecular Bond
  • Electronegativity and Bond Polarity 3:26
    • Electronegativity (EN)
    • Periodic Trend
  • Electronegativity and Bond Polarity Cont'd 6:04
    • Bond Polarity: Polar Covalent Bond
    • Bond Polarity: Nonpolar Covalent Bond
  • Lewis Electron Dot Structure of Atoms 9:48
    • Lewis Electron Dot Structure of Atoms
  • Lewis Structures of Polyatomic Species 12:51
    • Single Bonds
    • Double Bonds
    • Nonbonding Electrons
  • Lewis Structures of Polyatomic Species Cont'd 14:45
    • Drawing Lewis Structures: Step 1
    • Drawing Lewis Structures: Step 2
    • Drawing Lewis Structures: Step 3
    • Drawing Lewis Structures: Step 4
    • Drawing Lewis Structures: Step 5
    • Drawing Lewis Structure Example: Carbonate
  • Resonance and Formal Charges (FC) 24:06
    • Resonance Structures
    • Formal Charge
  • Resonance and Formal Charges Cont'd 27:46
    • More On Formal Charge
  • Resonance and Formal Charges Cont'd 28:21
    • Good Resonance Structures
  • VSEPR Theory 31:08
    • VSEPR Theory Continue
  • VSEPR Theory Cont'd 32:53
    • VSEPR Geometries
    • Steric Number
    • Basic Geometry
    • Molecular Geometry
  • Molecular Polarity 37:51
    • Steps In Determining Molecular Polarity
    • Example 1: Polar
    • Example 2: Nonpolar
    • Example 3: Polar
    • Example 4: Polar
  • Bond Properties: Order, Length, and Energy 40:38
    • Bond Order
    • Bond Length
    • Bond Energy
  • Summary 43:09
  • Sample Problem 1 43:42
    • XeO₃
    • I₃⁻
    • SF₅

Transcription: Bonding & Molecular Structure

Hi, welcome back to Educator.com.0000

Today's lecture from general chemistry is going to be on bonding and molecular structure.0002

As always, we will start off with a brief introduction.0010

After that, we are going to then go into the different types of chemical bonds.0014

After types of different chemical bonds, we are then going to discuss yet another periodic trend0019

which is called electronegativity and how it relates to something we call bond polarity.0024

The next two bullets deal with what we call Lewis electron dot structures.0030

These are going to become quite prominent in our future discussions in general chemistry.0035

After discussing Lewis structures, we will go into something we call resonance and formal charges.0042

After that, we then jump into a type of bonding theory called VSEPR theory or valence shell electron pair repulsion.0049

Followed by molecular polarity and bond properties.0058

As usual, we will finish it off with a brief summary followed by some sample problems.0064

In the previous lecture on molecules and atomic structure and periodic trends, we introduced the concept of valence electrons.0072

Remember valence electrons are an atom's or ion's outermost electrons.0082

That is they are the farthest away from the nucleus and the least tightly held.0088

It turns out that we have highlighted them for a reason.0094

It is because they are the primary electrons involved in chemical reactions and in chemical bonding.0097

In this lecture, we are going on to focus our attention on valence electrons and their qualitative representations.0107

Before going into that, we need to first discuss what types of bonds occur.0116

Really for our purposes for general chemistry, there are two types of bonds.0124

The first type is an ionic bond.0127

This exists between a metal and a nonmetal or between a cation and an anion.0131

We actually touched upon this when we first learned about polyatomic ions and naming ionic compounds.0138

Remember that in an ionic bond, it is really an electrostatic attraction occurring,0145

resulting in complete transfer of electrons always from the metal to the nonmetal.0150

That is not the type of bond we are going to be focusing on in this lecture.0157

We are going to be focusing on what is called a molecular bond.0161

A molecular bond exists between nonmetals.0166

Remember we learned how to name molecular compounds earlier on.0168

In a molecular bond, electrons are not being transferred, instead they are shared.0174

This sharing of electrons occurs mutually or not.0180

Because of this sharing, a molecular bond is also referred to as a covalent bond0184

which is how I am going to start to refer to it from now on, a covalent bond.0189

Next we are going to examine this sharing of electrons and0194

see the extent of how mutual or not the covalent bond is.0197

In order to address the question whether or not the sharing of electrons in0208

a covalent bond is mutual or not, we need to now introduce another periodic trend.0213

This periodic trend is what we call electronegativity.0218

I am going to abbreviate it as EN for short.0222

Electronegativity, it is literally a quantity, it is a value that describes the ability of0227

an atom to attract electrons to itself from a chemical bond in a gas phase.0232

Again it is the ability of an atom to attract electrons to itself in a chemical bond.0240

This is usually determined in the gas phase.0246

Electronegativity is more or less like your GPA which is on a scale of 0 to 4 with 4 being the highest.0250

High electronegativity values tend to be those elements that have very strong nuclear charges.0260

They attract electrons quite readily to themselves.0266

Those atoms that have very strong nuclear charges also tend to be very small.0272

What we are going to do right now, we are going to reexamine our periodic table.0277

Let's go ahead and repeat the periodic trends that we have already been through.0283

Going this way remember is atomic size increasing.0288

Going this way it was increasing ionization energy.0295

If elements that have high electronegativity values have strong nuclear charges and tend to be small,0305

that means the highest electronegativity values are going to be right here.0311

In other words, electronegativity is going to parallel ionization energy.0315

Again electronegativity is going to be parallel to ionization energy and opposite to atomic size.0323

There is a little acronym we can remember to help us with this.0330

F-o-n-c-l-b-r-i-s-c-h, fonclbrisch is the pronunciation of this little silly acronym.0334

But it is a nice crude way of remembering the order of electronegativity0345

in case you don't readily have a periodic table in front of you.0350

Fluorine is going to be the highest electronegative atom.0354

Carbon and hydrogen are going to be the lowest; fonclbrisch.0358

We are going to classify a bond with specific words depending on the0366

difference in the electronegativities of the two atoms that share the bond.0373

If the difference in electronegativity is great, then the sharing of electrons we say is not mutual.0379

That means the atom with the higher electronegativity gets a greater share of the electron density within the covalent bond.0387

Such a bond is what we call a polar covalent bond.0395

We schematically indicate polarity using partial charges and a dipole arrow.0400

Let's go ahead and take a look at an example.0405

We look at the hydrochloric acid and the bond between hydrogen and chlorine.0409

If I look at fonclbrisch, chlorine is going to be more electronegative than hydrogen.0414

That means chlorine is pulling all of the electron density to it.0425

In other words, all of the electron density accumulates around chlorine.0432

It becomes partially negative due to all that electron density.0437

To represent partial, we use the Greek symbol, δ.0442

If chlorine is δ-, that means hydrogen is going to be δ+0447

because all of its electron density is essentially being removed away from it and toward the chlorine.0451

This produces what we call a dipole moment in the molecule.0460

A dipole moment is going to be represented using a dipole arrow.0469

The dipole arrow is always going to point toward the more electronegative atom.0473

Again a dipole arrow points toward electronegative atom; toward the more electronegative atom.0482

Basically what a dipole arrow is, a dipole arrow or dipole moment, it basically shows where the electron density has accumulated.0491

It points toward the center of higher electron density.0503

Shows area / points toward area of accumulated electron density.0510

Let's go ahead and look at the opposite situation.0531

If the difference in electronegativity is negligible, then we have a pretty much equal sharing of electron density.0535

It is more or less mutual.0543

Such a bond is called a nonpolar covalent bond; nonpolar covalent bond.0545

For example, if you just look at typical hydrogen H2, we have a difference in electronegativity between the two atoms to be zero.0551

They are identical; this is a nonpolar covalent bond.0559

You may be wondering what exactly is considered negligible and what is considered great.0563

I am going to leave that up to your instructor and up to your textbook that you are using because0570

everybody can have a slightly different gauge of the difference in electronegativity and0575

what they consider to be negligible and what they consider to be rather significant.0581

Now that we know the different types of bonds and what we are exactly dealing with,0591

we want to now introduce a schematic way to represent valence electrons.0596

We first are going to talk about atoms.0601

Basically we represent valence electrons schematically as dots.0603

If you recall, you can just go by the periodic table to determine how many valence electrons exist for main group elements.0611

For main group elements, the number of valence electrons equals to the column number that the element is in.0621

Let's go ahead and look at sodium.0642

Sodium is in column one of the periodic table; it has one valence electron.0647

All I am going to do, I am going to represent that as one dot.0651

That is it; that is the electron dot structure for sodium.0654

Let's go ahead and look at magnesium.0659

Magnesium is in column two of the periodic table; it has two valence electrons.0661

I am going to put one of those dots there.0666

I am just going to put another dot maybe on another side of the element if you will.0668

Boron is an element in column number three.0675

It is going to have three valence electrons so I am going to show that to have three dots.0681

Carbon, column number four, it is going to have one dot here, here, here, and then there.0685

On to nitrogen, nitrogen is column five; it has five valence electrons.0694

One dot, two dot, three dot, fourth dot; where do I go put the fifth dot?0701

Really there is not too much of a physical significance here, but I am going to start pairing up.0705

Remember when we talked about atomic orbitals and Hund's rule.0711

Remember we singly fill orbitals before pairing up; I am just going to follow that.0715

Let's go ahead and apply that to oxygen.0721

Oxygen has six valence electrons because it is in column six.0722

One, two, three, four, five, and six.0726

On to fluorine now is in column seven.0732

One, two, three, four, five, six, and seven.0738

Finally we can do maybe krypton.0743

Krypton is going to be a noble gas in column eight.0748

One, two, three, four, five, six, seven, and eight.0751

These are the representative elements from their respective column numbers and how we represent them in terms of Lewis dot structures.0757

That was pretty simple, representing valence electrons of atoms, but now let's go bigger.0774

Let's go ahead and look at electron dot structures for polyatomic species.0780

Basically single bonds are going to be shown as a single line each containing two valence electrons.0785

Something like this; this is what we call a single bond.0792

There is going to be exactly two valence electrons in there.0799

I am going to represent valence electrons as VE for short now.0804

Single bonds are not the only type of covalent bonding that we can have.0809

We can also have double bonds.0812

For example, in oxygen, you can have a double bond like that.0815

In a double bond, there is going to be exactly four of the valence electrons.0822

You can also have triple bonds; for example, nitrogen N2 contains a triple bond.0828

In the triple bond, there is just going to be six valence electrons.0835

Not all electrons are participating in the bonding.0841

There are what we call nonbonding electrons.0846

Basically for general chemistry purposes, we are going to represent nonbonding electrons as lone pairs.0848

For example, when I go back to the Lewis structure for oxygen each oxygen contains two lone pairs like that.0855

That is what we call nonbonding electrons.0864

Also for nitrogen just to finish it up each nitrogen also contains one lone pair each.0868

How exactly am I coming up with these Lewis structures?0877

There happens to be a very systematic procedure for us to follow.0881

Let's go ahead and take a look at them right now.0885

Step number one, we are going to determine the central atom.0889

There is a couple of guidelines we can use to help us do this.0892

The first point is the following.0896

It is often the element that appears once in the formula.0897

It is never hydrogen; it is often carbon.0901

If there is more than one element that appears by itself, then you choose the least electronegative element.0907

That is number one is to determine the central atom.0915

Number two, we are going to arrange all other atoms around the center and0917

connect it to the central atom using only single bonds first; using only single bonds first.0921

Let's go ahead and look at water; water is H2O.0928

Here oxygen is the only element appearing by itself.0933

I am going to put it right to the center; step one is done.0938

Now I move on to step two; I surround the center with every other atom.0942

I connect it to the center using a single bond only.0946

Step two is done; let's go ahead and move on.0950

Step three, we are going to complete what is called the octet.0953

Remember we looked at this earlier on that why do atoms gain or lose electrons?0957

It is because they want to achieve eight electrons total, a noble gas configuration.0965

This magic number of eight is what we call an octet.0971

Every atom needs to have an octet in a Lewis structure in order for it to be quote and quote acceptable.0975

There are a couple of exception though--elements that don't need an octet.0982

Hydrogen, hydrogen is completely happy with just two electrons.0987

That is what we call the duet rule for that purpose.0992

Beryllium, beryllium is completely happy with just the four electrons.1004

Boron and aluminum, these elements are fine with just six electrons.1005

Phosphorus and sulfur, phosphorus and sulfur and really elements in period three or more, these actually can have more than an octet.1006

This is what we call an expanded octet.1016

Don't worry, we are going to look at a lot of examples where this occurs.1021

We are going to complete the octet when we read this using lone pairs of electrons.1026

When I go ahead and look at H2O or water, there is two electrons here.1031

There is two electrons here which means there is four electrons.1039

Now I want to give oxygen eight electrons.1042

That means I need four more electrons to get to eight which is two lone pair.1046

I am just going to do it like that.1050

What we need to do now in step four is to perform a valence electron count.1053

Mother nature has given us these elements.1058

We have to limit ourselves to the number of valence electrons provided by each element.1061

We can only use the total number of valence electrons, not any less, not any more.1068

We have to also account for the charge if it is a polyatomic ion.1074

For each positive charge, we are going to delete one valence electron.1079

For each negative charge, we are going to add one valence electron.1083

Don't worry, we are going to do a lot more examples.1087

But let's go ahead and check out water.1089

For hydrogen in water, there is two hydrogens.1092

Each hydrogen is in column one of the periodic table which means I am going to get two electrons just from hydrogen.1096

There is one oxygen.1103

Oxygen is in column six giving me six electrons just from oxygen.1105

The grand total of valence electrons in water is therefore eight.1111

Let's see if we have used eight in this molecule here.1117

I have two from the single bond here, two here, giving me a grand total of four, six, and eight.1123

Yes, this is a valid Lewis structure for water because everybody's octet rule is good to go.1131

Hydrogen's duet rule is satisfied.1139

I have used exactly the allotted indicated number of valence electrons.1142

If too many electrons are used, we need to then resort to multiple bonding.1153

We are going to try double bonds first, followed by triple bonds.1158

But here is the warning.1162

For each multiple bond added, we need to delete a lone pair from each atom sharing the bond.1165

Let's go ahead and look at an example; that is carbonate.1171

Carbonate is CO32-; this is a nice polyatomic ion we can look at.1176

Step one is to determine the central atom.1182

That is going to be the element I am pairing by itself.1184

In this case, it is only carbon.1187

Step two, we are going to surround the central atom with all other atoms.1190

We are going to connect it to the center using a single bond just like that.1195

Step three, we are going to complete everybody's octet using a lone pair of electrons.1201

Oxygen here needs six more electrons because it already has two nearby.1209

Same thing for this oxygen; same thing for this oxygen.1215

The central carbon has six nearby electrons already.1221

It only need two more electrons or one lone pair.1224

Let's go ahead and do a valence electron count now.1231

Don't forget this is a polyatomic ion.1234

I am going to put that in brackets and show the charge.1236

I have three oxygens each contributing six electrons giving me a grand total of eighteen valence electrons just from oxygen alone.1241

I have one carbon contributing four electron.1253

That is going to be four valence electrons just from the carbon.1256

Remember we must account for the charge; a 2- charge means we add two electrons.1259

Remember for negative charges, we add one electron per negative charge.1267

For cations, we take away one electron per positive charge.1273

My grand total of valence electrons is going to be twenty-four valence electrons that I must use in this compound.1278

When we go ahead and look at carbonate now over here, do we have exactly twenty-four?1287

The answer is no; it turns out that we have twenty-six electrons total.1292

We are over our limit.1298

What we are going to do right now is we are going to resort to multiple bonding.1300

We are going to try double bonds first followed by triple bonds.1303

What I am going to do, let's do this in red.1307

I am going to try one double bond first just like that; what else?1310

We have to delete a lone pair from each atom that shares the bond.1316

Let's go ahead and do that right now.1322

There is one lone pair gone from oxygen.1325

The one lone pair gone from the carbon.1327

Let's go ahead and recount.1332

Yes, this Lewis structure, the way we have it here, has exactly twenty-four valence electrons.1334

Every oxygen has an octet; so does the carbon.1340

This is a valid Lewis structure for carbonate.1344

If you follow these steps every time, you will be fine.1349

But let's go ahead and look at one more example.1355

Let's go ahead and look at nitrogen, N2.1359

Nitrogen, let's go ahead and connect it with a single bond just like that.1363

Now I am going to complete everybody's octet using lone pair of electrons.1368

Each nitrogen here needs six more electrons to get eight.1373

Remember they share those two electrons in the bond.1377

When we go ahead and do the count, we have two nitrogen each contributing1380

five valence electrons which means ten valence electrons are allotted to us.1384

When I look at this molecule and when I count the number of valence electrons, we are too many.1390

We have used twelve; we are going to resort to multiple bonding right now.1395

Sorry, we used fourteen; we used fourteen electrons; we only can use ten.1402

I am going to try a double bond.1408

When I do that, I have to delete a lone pair from every atom involved.1410

Let's go ahead and recount; when we recount, this is twelve electrons.1415

I am still way over my limit; double bonds have failed us.1420

Now we resort to triple bond.1424

I have to remove another lone pair of electrons from each atom sharing the bond.1428

Here we go; each nitrogen has an octet.1433

I have used exactly ten valence electrons.1437

Again follow this basic five-step procedure; you will get it right every time.1441

You notice that for carbonate we put the carbon-oxygen, the double bond on the twelve o'clock position if you will.1451

I could have put the double bond along any other one with oxygen.1458

That is what we call resonance structures.1466

Resonance structures are basically different Lewis structure of a same molecule.1468

For carbonate, this is what we drew originally.1472

Again I could have easily put it down here.1480

I am just putting my lone pairs in right now.1489

Or I could have easily put it right here.1492

These are what are called resonance structures.1504

These are Lewis structures of the same molecule.1506

In chemistry, we like to indicate the presence of resonance structures with a double ended arrow like that.1509

We like to keep track of electrons around an atom relative to its nuclear charge.1521

This is what we call the formal charge.1526

The formal charge is equal to the following.1530

It is equal to the number of valence electrons minus the number of lines and dots attached.1532

Let's go ahead and look at this structure for the carbonate ion.1538

Let's go ahead and look at this oxygen here.1547

This oxygen right there, let's go ahead and use this equation.1552

Formal charge here is equal to number of valence electrons.1556

Six minus the lines and dots attached, the number of lines and dots attached.1560

I have six dots and one line attached giving me a grand total of seven.1565

That formal charge is -1.1571

We usually like to indicate the formal charge just as a tiny superscript around the atom.1573

I personally like to circle them like that.1579

This oxygen here is identical; it is three lone pairs and one bond.1583

This is also -1.1588

Let's look at the carbon.1593

The carbon is going to be four valence electrons minus...1594

I have how many lines attached total?--four.1601

This carbon has a 0 formal charge.1606

This oxygen here, that is going to be six minus the number of lines and dots attached.1609

I have two lines attached and four dots giving me a grand total of six.1615

That one is also a 0 formal charge.1620

Notice what do the sum of my formal charges equal to?1626

The sum of my formal charges is equal to -2 which is the overall charge of the ion.1630

The nice thing about formal charges is that when I add up all formal charges,1637

it must equal to the overall charge of the molecule or ion.1641

This is yet another way to determine if the Lewis structure that you have drawn is valid or not.1650

Remember it must equal to the overall charge of the molecule or ion, the sum of the formal charges that is.1658

We have already assigned the formal charge to each element in carbonate.1670

Again the sum of the formal charges equals to the overall charge of the molecule or ion.1677

It turns out that not all resonance structures for a molecule or ion are made equal.1684

In terms of carbonate, all three were identical.1691

We always had one carbon-oxygen double bond and two carbon-oxygen single bonds.1694

All three were identical.1699

But now let's go ahead and look at sulfuric acid.1701

Sulfuric acid is an example of a molecule that can have resonance.1705

But we are going to find that the resonance structures are not identical.1709

For example, when you look at sulfuric acid, one way to draw its Lewis structure is right here with all single bonds.1713

I am just going to go ahead and put in the lone pairs to complete everybody's octet just like that.1721

Yet another way to draw sulfuric acid is with two double bonds like this.1729

Remember that sulfur is one of those atoms that can have more than eight electrons.1739

In this case, it has ten electrons.1744

Everybody else has an octet and hydrogen's duet rule is good.1747

Is it the resonance structure on the left or the resonance structure on the right?1752

Which one best portrays sulfuric acid?1756

We have to go through what makes up a good resonance structure.1761

It is all about formal charges.1765

Good resonance structures are going to maximize the number of covalent bonds.1767

They are going to minimize separation of charge.1773

That means formal charges of zero are ideal.1776

They are going to obey the rule, the octet rule as much as possible obviously.1780

Finally if there is a negative formal charge, we are going to try as1786

best as we can to place it on the most electronegative atoms.1789

If we go ahead and look at it, here we have -1 formal charge, -1 formal charge.1798

Sulfur is actually having a +2 formal charge right here.1804

That is equal to six minus four.1809

But when we look at the Lewis structure on the right, this is 0.1813

This is 0, 0, 0 and now sulfur actually has a formal charge of 0.1818

You cannot beat that; you cannot get better than all 0 formal charges.1826

The Lewis structure on the right, the resonance structure on the right, is what we call the major contributor.1831

The resonance structure on the left is what we call the minor contributor.1842

We have seen how we can use formal charge to determine what is a quote and quote good Lewis structure.1849

Now we are going to use Lewis structures to help us determine how a molecule physically looks like.1856

That is what is its shape?--what is its geometry?1863

The theory that helps us to determine this, that is going to give us some guidelines is what we call VSEPR theory.1870

VSEPR theory stands for valence shell electron pair repulsion.1878

Basically the premise behind this theory is the following.1885

That a molecule is going to adopt a configuration, basically a shape or geometry, which minimizes electron-electron repulsion.1888

Remember electrons are the same charges; they don't want to be close together.1898

In other words, the conformation is going to maximize the distance between thereby electrons.1903

There is different types of electron-electron interactions.1910

Lone pair-lone pair interactions are going to be the worst because it turns out that lone pairs occupy the greatest volume.1913

We want to avoid lone pairs being too close together.1927

After lone pair-lone pair interactions is going to be lone pair-bonding pair interactions and finally bonding pair-bonding pair interactions.1933

Right now in the next slide, we are going to introduce molecular geometries, their names, and their respective bond angles.1946

What I mean by bond angle is if you have a central atom A surrounded by two atoms like this,1955

basically the bond angle is going to be the angle of that.1962

Bond angle just like in a typical geometry class.1967

Let's go ahead; now this looks like a lot of material.1973

This looks very sometimes intimidating; but again it is not too bad.1978

Let's go ahead and just take each molecule one by one.1984

Number one, the steric number.1987

The steric number is just a fancy name for the number of electron groups around the central atom.1990

What I mean by an electron group is basically any type of bond.2002

A single bond counts as one group; a double bond counts as one group.2011

A triple bond counts as one group; or a lone pair.2015

Basically one lone pair counts as one group.2024

The basic geometry, the basic geometry is also known as the electron geometry.2031

Basically the basic geometry or electron geometry is everything in this first column.2043

What they have in common is that the central atom E contains zero lone pairs.2051

When we have a situation where there is going to be only two electron groups around a central atom with no lone pairs,2059

that results in a linear geometry or a bond angle of 180 degrees.2068

When we have three bonding groups and no lone pairs around a central atom,2073

we get what is called a trigonal planar geometry and a bond angle of 120.2077

When we come up to four electron groups around a center and no lone pairs, this is what we call tetrahedral.2083

You are going to see that we have what is called a broken wedge and a solid wedge notation.2092

That implies three-dimensional nature of the molecule.2098

Broken wedge means it is pointing away from you into the paper.2101

Solid wedge means it is pointing toward out of the paper.2107

If you look at this molecular here and imagine its three-dimensional nature, we do get a pyramid.2112

A tetrahedron basically is what it is called; let's go on.2118

Five bonding groups around the center atom and no lone pairs, that is what we call trigonal bipyramid.2124

We have two bond angles going on, 120 and 90 degrees.2130

Finally six bonding groups around the central atom, no lone pairs,2135

we get an octahedron or an octahedral geometry with a 90 degree bond angle.2138

What happens as soon as we place lone pairs on a central atom?2145

As soon as we place lone pairs on the central atom, we get everything here.2150

We are going to use a new name to describe their geometries.2157

That is what we call the molecular geometry or the molecular shape.2161

Please make note, there is a difference between electron geometry and molecular geometry.2166

Sometimes they are identical.2171

When they have zero lone pairs like we just covered, electron geometry is the same as molecular geometry.2174

For steric number three, we saw that the electron geometry was trigonal planar.2182

But as soon as we put a lone pair from one of the Xs,2189

that lone pair is a great area of electron density.2195

It is going to push the Xs away from it.2199

We don't have trigonal planar anymore.2203

Instead the geometry name is called bent or angular.2204

We have a bond angle of less than 120.2208

Steric number four, with one lone pair, we get what is called a trigonal pyramid with a bond angle less than 109.2212

If we have two lone pairs, we have a bond angle of much less than 109.2222

Steric number five, one lone pair is what we call a seesaw geometry.2228

If we have two lone pairs, it is called a t-shape geometry.2236

If we have three lone pairs, it is called a linear geometry with a bond angle of 180.2241

Finally steric number six, one lone pair is called square pyramid.2246

Two lone pairs is going to be called square planar.2252

Three lone pairs is t-shape; four lone pairs is going to be linear.2255

You should definitely consult with your instructor and determine which ones he or she exactly would like you to know by heart.2261

Now that we have gone over the names of the geometries and how molecules can look like,2274

we are going to use a molecule shape to determine if it itself is polar.2280

We have already talked about bond polarity, but what about an entire molecule?2286

A step by step procedure is the following.2290

Step one, we are going to draw the Lewis structure.2292

Step two, from the Lewis structure, we are going to determine the molecular shape.2296

Then we are going to assign all bond dipoles.2300

Remember that is the arrow that points toward the more electronegative atom.2302

Then we are going to do the following analysis.2307

If the dipoles cancelled each other, that means there is no net direction of electron density and the molecule is nonpolar.2310

If the dipoles seem to point in a general direction, that is they don't cancel,2319

then we say the molecule is expected to be polar.2325

Let's go ahead and look at a few general examples.2327

For example, if we have EXX and we have one dipole going this way and one dipole going this way,2331

this is a 180 degree bond angle, those dipoles completely cancel each other.2340

This would be an example of a nonpolar scenario.2346

We can do a couple of more examples.2351

EXXX and we can have one dipole, for example, let's go here, one going here, one going here.2353

In this case, all three dipoles, they point in completely opposite directions at a 120 degree bond angle.2366

This would be also an example of a nonpolar molecule.2373

Let's go ahead and look at another one.2378

Now EXX and X; this would be trigonal pyramid geometry--one dipole here, here, and here.2381

In this case, because this is a three-dimensional pyramid, these dipoles don't cancel.2392

They all point away from the central atom.2397

This would be an example of a polar molecule.2405

Finally one last example, EXX; let's have two lone pairs here.2409

In this case, we can have one dipole going that way; one dipole that way.2418

Remember this is a bent geometry; my net dipole is going downward.2422

This would be also an example of a polar molecule.2427

These types of problems take a little getting used to.2431

But with more and more practice, I am sure...2433

I know you will get better and better and more confident in it.2435

The final item that we want to look at for covalent bonding is bond properties.2441

There is really three properties that we want to look at now.2447

What is called bond order, bond length, and bond energy; these are the three parameters.2450

The first one is called bond order; basically bond order is a fancy name.2457

A bond order of one is going to be indicative of a single bond.2462

Bond order of two is going to be indicative of a double bond.2466

Bond order of three is going to be indicative of a triple bond.2469

We are going to come across bond order again in a future chapter.2472

I am just going to leave it at that for now.2479

Bond length is literally the length of the bond.2483

It is the distance between two nuclei sharing a bond.2487

The shorter bond, the greater the interaction between bond nuclei.2492

We have greater ability for electrons to be shared.2497

What that means is that we get a stronger bond.2501

The shorter the bond, the stronger it is; shorter means stronger.2504

That is easy to remember because both of the words here start with an s.2511

Bond energy, bond energy is the third parameter.2516

Bond energy is also known as bond enthalpy; that is right.2519

It is going to be yet another type of ΔH.2524

Remember ΔH from thermodynamics; the formal definition is the following.2526

Bond enthalpy refers to the amount of energy required to form a bond between two nuclei in the gas phase.2532

That is it is the amount of energy required; what does that mean?2541

If it is required that means it is going to be an endothermic process which means2545

bond enthalpies are typically going to be reported as positive values and in units of kilojoules per mole.2553

If for example a carbon-carbon single bond has a bond enthalpy of 347 kilojoules per mole.2562

As we go to a carbon-carbon double bond which is stronger, we expect more energy.2567

A stronger bond is going to require more energy to break.2572

A carbon-carbon double bond is 630.2576

Carbon-carbon triple bond is going to be the strongest and have the highest bond enthalpy of around 800 kilojoules per mole.2578

Notice that they are all positive values.2586

Let's summarize what we went over today.2591

First point is that Lewis structures are basic visual depictions of covalent bonding in molecules and ions.2593

A formal charge is a nice tool to help us determine a valid Lewis structure.2600

Using VSEPR theory, we were able to predict the electron and molecular geometries for any given molecule.2606

That is how is it going to look three-dimensionally?2616

Let's now go over a sample problem.2623

This sample problem is basically the following; here we are given three molecules.2627

For each molecule or ion, we want to determine a) the best Lewis structure,2631

b) we want to determine the electron and molecular geometries and determine if it is polar or nonpolar overall.2637

Let's go ahead and do this one.2644

Xenon trioxide, my central atom is going to be xenon.2646

I am going to surround the xenon with the oxygens like that.2652

I am going to connect everything to the center like that.2655

I am then going to proceed and complete everybody's octet.2659

Each oxygen is going to need three lone pairs.2664

The xenon is going to need one lone pair itself.2669

Let's go ahead and see if we have used the exact number of valence electrons allotted to us.2672

Three oxygens contribute six valence electrons; that is eighteen from oxygen.2681

The one xenon contributes eight valence electrons.2691

I have a grand total of twenty-six valence electrons that I can use.2696

Let's go ahead and see if we have used twenty-six.2700

It is two, four, six... two, four, six, eight, ten, twelve, fourteen, sixteen, eighteen, twenty, twenty-two, twenty-four, and twenty-six.2704

This looks alright; everybody's octet is complete.2717

I have used exactly twenty-six valence electrons.2720

Let's look at the formal charges though.2724

The formal charge of the xenon is going to be eight minus five which is +3.2726

Each oxygen is -1.2735

Remember formal charges of 0 are best; +3 is pretty far from 0.2739

There is another Lewis structure that we can draw; it is basically a resonance structure.2745

Remember that xenon is an element in period three or higher.2750

It can actually come up with more than an octet.2756

It turns out that we can come up with a Lewis structure for xenon trioxide like that.2764

When I do this, look what happens to all my formal charges.2771

Oxygens are now 0; xenon has a formal charge of now 0.2775

That is going to be better than our left Lewis structure.2780

Our best answer here is going to be that one.2784

Polarity, oxygen is more electronegative than xenon.2789

I have one dipole that way, one that way, and I have one that way.2793

This is a trigonal pyramid geometry with a bond angle less than 109.5.2798

When I think about this three-dimensionally, these three arrows that we have just drawn do not cancel.2810

This molecule is expected to be polar; the dipole arrows do not cancel.2815

I31-, this is going to be triiodide.2825

In this case, it is going to be iodine, iodine, iodine.2829

Let's go ahead and connect everything to the center of a single bond.2833

Let's go ahead and assign lone pairs to complete everybody's octet now.2836

This iodine on the left needs three; this iodine on the right needs three.2841

The iodine in the middle is going to need two lone pairs.2845

Again this is all going to be a -1 anion.2849

Let's go ahead and do the valence electron count.2853

I have three iodines each contributing seven valence electrons giving me a grand total of twenty-one.2855

I have one additional valence electron from the -1 charge which means I can only use a grand total of twenty-two valence electrons.2861

That is eleven pair--one, two, three, four, five, six, seven, eight, nine, ten.2871

We are too few now; we are two electrons shy.2880

We are not going to try multiple bonding.2886

But remember iodine is in period three or higher.2889

I can actually have another lone pair on the central atom to have more than eight electrons.2893

Let's see if the formal charges though are happy.2901

This iodine here on the left has a formal charge of 0.2903

This one on the right has formal charge of 0.2907

This one in the center has a formal charge of -1.2910

Does everything add up to the overall charge?2916

The answer is yes; this is a valid Lewis structure.2919

Remember iodine is somewhat electronegative in fonclbrisch.2923

It is going to be perfectly defined with a negative formal charge.2927

That is the Lewis structure for triiodide.2932

What is the geometry then?--the geometry is what we call linear.2937

That is going to be a 180 degree bond angle.2941

This molecule is going to also be nonpolar because this molecule is homonuclear which means we have no difference in electronegativity among the atoms.2945

Last example is going to be SF5 for sulfur pentafluoride.2958

Central atom in the middle is sulfur.2964

I am going to surround the center with all fluorines like that.2966

I am going to connect everything to the center using a single bond like that.2972

The next step is to complete everybody's octet using lone pair of electrons.2978

Let's go ahead and do that.2984

You see sulfur has already an octet; in fact it has more than an octet.2993

But that is okay because it is in period three or higher.2997

Everybody seems to be good to go.3002

Let's go ahead and do a valence electron count.3003

I have five sulfurs each contributing six valence electrons giving me thirty... I am sorry, my apologies.3007

I have one sulfur contributing six valence electrons giving me six.3016

I have five fluorines each contributing seven valence electrons giving me a grand total of thirty-five valence electrons.3024

When I add this up, I get a grand total of forty-one valence electrons; forty-one valence electrons.3042

Let's go ahead and see if this is right or wrong.3054

There is a little typo here; I am very sorry about that.3062

This shouldn't be SF5; this should be PF5.3065

I am going to go ahead and change that; my apologies.3072

I am going to go ahead and make that PF5; now it is phosphorus pentafluoride.3075

Phosphorus is going to be in the middle now.3080

I am going to change sulfur to phosphorus.3084

Phosphorus is going to be in column five.3087

That is going to contribute five valence electrons.3092

My total is going to be forty valence electrons.3095

Let's see if we used forty valence electrons or not which is twenty pair.3098

Two, four, six, eight, ten, twelve, fourteen, sixteen, eighteen... excuse me.3104

Two, four, six, eight, ten, twelve, fourteen, sixteen, eighteen, twenty, twenty-two, twenty-four, twenty-six, twenty-eight, thirty.3112

Thirty-two, thirty-four, thirty-six, thirty-eight, and a grand total of forty electrons or twenty pair.3122

Yes, this is the valid Lewis structure for phosphorus pentafluoride.3128

This was molecular shape and geometry using VSEPR theory.3134

This molecule is going to be nonpolar because everything is cancelled with all the dipoles.3141

Everything is going to be cancelled just like that.3148

Thanks everyone for your attention; I will see next time on Educator.com.3154

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