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Franklin Ow

Franklin Ow

The Chemistry of The Transition Metals

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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The Chemistry of The Transition Metals

  • Coordination compounds contain a central metal atom/ion attached to ligands, which may be atoms and/or molecules.
  • Coordination compounds are often observed with coordination numbers of 4 or 6, giving tetrahedral, square-planar, and octahedral geometries.
  • A set of rules exists for nomenclature (naming) of coordination compounds.
  • Crystal-field theory can explain both the color and magnetism of coordination compounds.

The Chemistry of The Transition Metals

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 0:11
  • Coordination Compounds 1:20
    • Coordination Compounds
  • Nomenclature of Coordination Compounds 2:48
    • Rule 1
    • Rule 2
    • Rule 3
  • Nomenclature cont'd 4:58
    • Rule 4
    • Rule 5
    • Rule 6
    • Rule 7
    • Rule 8
  • Nomenclature cont'd 7:39
    • Rule 9
    • Rule 10
    • Rule 11
    • Nomenclature of Coordination Compounds: NH₄[PtCl₃NH₃]
    • Nomenclature of Coordination Compounds: [Cr(NH₃)₄(OH)₂]Br
  • Structures of Coordination Compounds 10:54
    • Coordination Number or Steric Number
    • Commonly Observed Coordination Numbers and Geometries: 4
    • Commonly Observed Coordination Numbers and Geometries: 6
  • Isomers of Coordination Compounds 13:13
    • Isomers of Coordination Compounds
    • Geometrical Isomers of CN = 6 Include: ML₄L₂'
    • Geometrical Isomers of CN = 6 Include: ML₃L₃'
  • Isomers cont'd 17:00
    • Structural Isomers Overview
    • Structural Isomers: Ionization
    • Structural Isomers: Hydrate
    • Structural Isomers: Linkage
    • Structural Isomers: Coordination Isomers
  • Electronic Structure 22:25
    • Crystal Field Theory
    • Octahedral and Tetrahedral Field
  • Electronic Structure cont'd 25:43
    • Vanadium (II) Ion in an Octahedral Field
    • Chromium(III) Ion in an Octahedral Field
  • Electronic Structure cont'd 28:47
    • Strong-Field Ligands and Weak-Field Ligands
  • Implications of Electronic Structure 30:08
    • Compare the Magnetic Properties of: [Fe(OH₂)₆]²⁺ vs. [Fe(CN)₆]⁴⁻
    • Discussion on Color
  • 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

Transcription: The Chemistry of The Transition Metals

Hi, welcome back to Educator.com.0000

Today's lesson from general chemistry is the chemistry of the transition metals.0002

We are first going to get introduced to what we mean by coordination compound and basically what they consist of.0010

After we get introduced to them, we are going to then turn our attention0017

to nomenclature which is learning the rules to name these coordination compounds.0021

Following this, we will then get into the geometries and structures that coordination compounds can commonly adapt.0027

Related to structures will be then something we call isomers.0036

Finally we are going to, in the last part of the lesson, go into0040

electronic structure of coordination compounds and basically what happens to the0045

valence electrons of a transitional metal when atoms or molecules become attached to them.0052

We are going to learn that that is called crystal field theory.0061

We will also get into what is called the spectral chemical series and the implications of the crystal field theory.0066

As always, we will finish up the lesson with a summary followed by some sample problems.0073

Coordination compounds is the term that refers to the following.0081

These are compounds that consist of a central metal atom or ion which have0085

attached a varying number of atoms or molecules via a coordinate covalent bond.0090

These atoms or molecules that are attached have their own names.0098

These are what we call ligands or ligands.0103

The term coordination compound can also be used interchangeably with the term coordination complex.0108

Basically a coordination compound can also have the central metal atom be electrically neutral or charged.0115

A case where we have a charged neutrality is when nickel has attached to it four carbonyl ligands.0122

We are going to learn the names of these shortly.0132

But here this is formerly nickel Ni0.0134

An example of a coordination compound where we have the transition metal to be charged, if it is FeCl2 and 2NH3s.0138

Here in this case, this is formerly Fe 2+.0149

What we are going to be focusing on in this session is coordination complexes where the central metal atom is actually a d-block element.0155

Let's go ahead and move on.0166

Now that we have been introduced to what we mean by a coordination compound, let's go ahead and move into how to name them.0168

This is what we call nomenclature of coordination compounds.0174

There is a lot of rules associated with this; let's get into it.0178

When writing formulas, you are first going to list the chemical formula or the chemical symbol of the metal first.0182

Followed by the ligand symbols in alphabetical order.0188

After this, if the complex is neutral overall, brackets are not needed.0194

However if the complex is an ion, you are going to enclose the entire complex in brackets followed by the charge as a subscript.0200

An example would be the following, Fe(CN)64-.0214

This tells me that the entire complex ion here has an overall 4- charge.0225

But something like this, FeCl2 followed by (NH3)2,0231

this tells me that the overall complex ion is neutral overall because you don't see brackets.0241

When the ligand is a molecule instead of an atom, sometimes the point of attachment0249

can also be emphasized in the order in which you list the letters of the molecule.0253

For example, we can say FeCl2 and (OH2)2; this is basically water.0262

We are listing the oxygen first because it is the oxygen that is0271

attached to the iron center, not the hydrogens; because point of attachment.0279

There is more rules; when naming the complex... this is always harder.0298

You are going to list the ligand names first followed by the metal name.0302

A Roman numeral in parentheses indicates the oxidation number of the central metal ion.0306

If you have a neutral ligand, neutral ligands are always going to retain their full name with the exception of the following four.0314

H2O is going to be now formally be called aqua.0324

NH3 is called ammine; CO is called carbonyl; NO is called nitrosyl.0328

Anionic ligands will also change their suffixes.0336

If your anion such as chloride ends in ?ide, it is going to become with an ?ido suffix.0340

For example, chloride is going to become chlorido.0347

If something ends in ?ate, it becomes with an ?ato suffix.0354

For example, sulfate is going to become sulfato.0359

Finally if you have something ending in ?ite, it becomes ?ito.0365

An example of this will be nitrite anion becoming nitrito.0372

Furthermore if you have more than one of the same ligand, we are going to0380

use Greek prefixes to indicate the number of each ligand, omitting mono.0384

We are going to use di- for two, tri- for three, tetra- for four, etc.0390

Next sometimes the ligand name already contains a Greek prefix.0401

If that happens, we are going to use the following system.0408

Bis- is going to be replacing di-; tris- is going to replace tri-.0412

Tetrakis- is going to replace tetra-.0418

The other scenario where you have to use these unique prefixes is when the ligand is polydentate.0421

What a polydentate ligand is that it has more than one point of attachment to the metal.0429

Once again a polydentate ligand means it has more than one point of attachment to the metal.0446

We are going to go ahead and look at more examples of those in greater detail later on in this lecture.0451

Finally the final four rows, we are going to list the ligand names in alphabetical order ignoring any prefixes.0459

If the complex ion is anionic, the suffix ?ate is added to the metal ion name.0466

An exception is iron; you don't call it ironate.0472

But instead we are going to use I believe the Latin phrase, ferrate.0475

If the complex ion is anionic, you are going to list the cation name first.0481

If the complex ion is cationic, you are going to show the anion name last.0486

Let's go ahead and look at an example.0491

Let's take the following, NH4Pt in brackets Cl3NH3.0495

In this case, my cation is ammonium.0504

My anion is the entire complex which means we don't call it platinum.0509

But we call it platinate.0517

We know that the ammonium is a 1+ overall charge.0520

We know that each chlorine is 1- each, giving me 3- overall.0527

We know that ammonia is neutral ligand overall.0535

This tells us that the platinum must have an oxidation state of 2+.0538

Now that we have established that, we can go ahead and name the compound.0546

This is going to be ammonium.0549

It is the name of the cation followed by the following.0554

We list the ligands in alphabetical order--ammonium ammine and then trichlorito-platinate(II).0556

Let's go ahead and do one more example for nomenclature.0572

This one is going to have in brackets chromium with four ammines attached and with two OH groups attached and then bromine.0579

In this case, bromine is our anion; this entire complex ion is our cation.0594

We know that amine is 0 overall; we know that hydroxo is 1- overall.0602

We know that bromine is 1- overall, telling us that chromium must be 3+ here.0608

Let's go ahead and name it; list the ligands in alphabetical order.0617

This is going to be tetra-amine followed by di...0621

Now hydroxide becomes renamed as hydroxito; then chromium(III) bromide which is the complete name.0633

That is nomenclature for coordination compounds.0648

Let's now move on to the structures of coordination compounds--what do they really look like.0653

We define the coordination number, which is also known as steric number, to be0658

the total number of points of attachment to the central metal atom or ion.0663

When I mean total points of attachment, I mean from all of the ligands.0671

The most commonly observed coordination numbers for these transition metal complexes is the coordination number 4.0675

When we have a coordination number of 4, we can get two common geometries.0684

Number one--tetrahedral; or number two--square planar.0688

In tetrahedron case, you have your central metal atom.0692

Then with the four ligands attached like so in a traditional tetrahedral format.0695

Of course for square planar, your central metal atom.0702

Now the ligands are arranged 90 degrees to each other, giving you square planar overall.0706

The other common coordination number is coordination number of 6.0716

For coordination number of 6, we are going to basically get the main geometry of octahedral geometry.0721

Here you have your central metal atom.0730

You have the four ligands forming a square base.0733

Then you have one ligand perpendicular to that; another ligand perpendicular right below.0739

We call the four ligand positions that are horizontal, that is formed as square plane, we call them equatorial positions.0746

Of course then for the ligands that are perpendicular to that plane, these are called the axial positions.0763

Again tetrahedral and square planar are for a coordination number of 4.0777

Octahedral is for a coordination number of 6.0780

Now related to structures of coordination compounds, we are now going to discuss isomers.0784

Recall that isomers are basically the following.0792

These are compounds that have the exact same formula but have different chemical and/or physical properties.0795

Specifically for the coordination number of 6, we can have a large number of different isomers.0802

The first type of isomers we are going to talk about is for coordination number 6.0810

These are what are called geometrical isomers.0815

When we have a geometrical isomer, these are two formulas that you want to look for.0818

ML4L'2 and ML3L'3 where L and L' are different ligands.0823

When we look at ML4L'2, we can have one possible configuration0842

where the four Ls that are identical to each other,0847

they form the square plane like so--L' here and L' here.0851

In the other possibility, let's switch an L' with an L; good.0860

When the L's are completely opposite to each other, 180 degrees apart, this is what we call the trans isomer.0877

However when the two Ls here and here are not 180 degrees apart, when they are0888

perpendicular to each other like this, that is what we call the cis isomer.0897

Again this is what we call trans and cis isomerism.0904

The other type of geometrical isomer that exists for a coordination number of 6 is ML3L'3.0908

In this case, we could have the following two possibilities.0916

M and then L', L', L'... L, L, and L.0922

The other possibility, let's substitute an L' for an L.0941

In the left scenario, you see that the L's actually form a plane with each other like so that bisects the octahedron.0958

This is what we call the mer isomer.0973

Mer stands for meridian such that the three identical ligands share a plane that forms a meridian through the compound.0977

The other possibility on the right is when the three ligands are actually sharing the same face of the octahedron.0990

That is what we call the fac isomer, fac being short for facial.1000

That is geometrical isomers.1009

Let's now take a look at the category of what we call structural isomers.1012

When we talk about the word structural isomers, this is a very general category.1020

Basically structural isomers are going to retain a ligand position.1032

Excuse me... they are going to change ligand position for structural isomers; change ligand position.1049

There are basically four types of structural isomers that we want to deal with.1057

One is what we call ionization isomers; two is what we call hydrate isomers.1062

Three is what we call linkage isomers; finally four is what we call coordination isomers.1072

Ionization isomers literally have a cation or an anion switch places in or out of the brackets of the coordination sphere.1087

That is, in one isomer, a certain ligand is attached to the metal.1120

In the second isomer, it is not.1125

Let's go ahead and look at an example.1127

Imagine here in brackets, cobalt, bromine, and then five amines attached, closed brackets, and then sulfate.1130

We can imagine a situation where the sulfate just substitutes for the bromine.1140

We get CO(NH3)5SO4 and then bromine.1146

My compound is still neutral overall because this is a salt after all.1155

The only difference is the location of sulfate with bromine.1162

The second type of structural isomer is what we call a hydrate isomer.1167

This simply involves ligand exchange with water.1171

Imagine chromium with three chlorines attached, coordinated to six waters.1184

We can imagine a situation where we have the waters now directly bonded to the chromium metal.1193

The three chlorines serving as counterions.1204

Linkage isomers, linkage isomers involve what are called ambidentate ligands.1210

They involve ambidentate ligands.1220

Basically the word ambidentate means the ligand can attach in more than one location to the metal.1225

Again the ligand can attach in different locations to the metal.1239

You could imagine a situation where you have transition metal; let's take a NO2 group.1247

But you could also imagine the ligand being directly attached to the oxygen instead of the nitrogen.1257

In this case, we get a pair of linkage isomers.1263

Finally we can have also a fourth type of structural isomer called coordination isomers.1267

Basically you have two metals switching location with each other.1275

An example would be a chromium with the six amines attached followed by iron with the six cyano groups attached closed brackets.1292

That will be one coordination compound.1308

But you can also have another one where you just switch places.1310

Now iron is going to get the six amines.1313

Chromium is going to get the cyanos just like that.1318

Again these are what are known as structural isomers where we actually involve a change in attachment to the metal.1326

Now that we have covered isomers, let's go ahead and take a look at how the electronic structure is in coordination compounds.1336

To examine how the d-block electrons of the central metal atom or ion are1345

arranged once ligands are attached, we use something called crystal field theory.1349

Crystal field theory treats the lone pairs on a ligand as point negative charges.1356

These negative charges influence of course how the metal d-orbitals are arranged.1362

Let's go ahead and examine the most common type of geometry which is an octahedral field and a tetrahedral field.1367

For an octahedral field, you can imagine the following.1375

Your transition metal in the middle; then the ligands are on the outside.1380

These are all treated again as point negative charges like that.1384

Here is the octahedral framework.1392

It is determined that of the five metal d-orbitals, two get raised in energy and three go down in energy.1400

It turns out that dxy, the dxz, and the dyz orbitals lie in between the ligands.1413

Because they lie in between the ligands, there is no electron-electron repulsion.1430

Their energies go down.1440

The two remaining d-orbitals, dx2-y2 and dz2, are found to lie directly pointing toward the ligand positions.1445

Because they lie directly toward the ligand positions, there is electron-electron repulsion occurring.1462

Because of this, their energies go up.1470

You can imagine the five d metal orbitals.1479

All of a sudden, once the ligands attach, you get two going up and two going downward.1482

The two metal orbitals that are together, the dz2 and the dx2-y2, these are what we call the eg orbitals.1490

The three on the bottom, the dxy, the dxz, and the dyz, these are what we call the t2g orbitals.1499

The difference in energy between the eg and the t2g levels are what we call Δo.1506

Δo is what is known as the CFSE.1512

The CFSE is what we call the crystal field splitting energy.1517

Now that we have look at an octahedral field and we know how the d orbitals are what we call split,1532

let's now take a look at electron configuration using crystal field theory.1538

Consider a vanadium(II) ion in an octahedral field.1543

Vanadium has the configuration... vanadium 2+ is going to have the configuration argon followed by 3d3.1550

We are only dealing with 3d electrons here.1560

Because this is an octahedral field, you get the following 2-3 split.1563

Just as we have learned how to fill orbitals using Aufbau principle, we are still going to follow that.1570

d3, with the first d electron, second, and the third.1577

For d3, this is the only possible ground state configuration.1582

Once again this is the only possible ground state configuration.1590

However something different happens when we consider a chromium(III) ion in an octahedral field.1598

The configuration for chromium 3+ is going to be argon 3d4.1608

Now we actually have more than one possible ground state configuration because I can do one, two, three, four.1617

But also I can have one, two, three, and four.1629

Basically is the configuration going to be t2g4?1637

Or is the configuration going to be t2g3 eg1.1644

What that depends on is the magnitude of the crystal field splitting energy.1651

It depends on Δo.1659

Basically if Δo is too large, then we are going to get the t2g4.1662

If Δo is small enough, that is negligible enough, then yes1674

it will be relatively easy to promote that last electron to the upper levels.1681

We will get the t2g3 eg1 configuration.1686

The situation where Δo is too large, this is what we call the low spin scenario.1693

The situation where Δo is very small is what we call the high spin scenario.1702

How do you determine if Δo is small or large?1713

That is strictly going to be dependent on the ligands.1717

Let's go ahead and take a look at a series of common ligands next.1722

Here we have what is known as the spectral chemical series.1727

The spectral chemical series lists commonly encountered ligands in coordination chemistry starting with the carbonyl and cyano ligands.1736

Wrapping it up on the other end of the spectrum are the halides.1750

The way to read this is the following.1754

Ligands at the leftmost end are called strong field ligands because they result in a large crystal field splitting energy.1756

As a result, with these ligands, we get low spin configurations.1764

Ligands at the rightmost end are called weak field ligands because they result in a small crystal field splitting energy.1769

As a result, we often get high spin configurations with these ligands.1778

What you have to be on the lookout for are both possible high and low spin scenarios.1783

These will arise for configurations d4, d5, d6, and d7.1792

What is the big deal?--you may be wondering; why bother with this?1800

Now let's go ahead and examine the implications.1803

This is where it gets very interesting; the implications are the following.1805

We can use crystal field splitting theory to help us predict both the color and the magnetic properties of coordination compounds.1810

Let's first compare the magnetic properties of the following.1820

Iron with six aqua ligands attached making this 2+ versus now iron with six cyano ligands attached making that 4-.1824

In this situation, this is Fe2+.1841

In this situation on the right, this is also Fe2+.1844

Both of these complexes are d6.1848

The only difference obviously then is the identity of the ligand.1853

Because here H2O, octahedral d6; here we are going to have cyano d6.1858

Let's start with the extreme.1867

Cyano is one of the first ligands you encounter.1868

It is one of the strongest field ligands.1871

Because it is strong field, that means it is going to be low spin.1874

One, two, three, four, five, six.1878

You notice that all of my valence electrons are paired.1882

For the cyano compound, we expect diamagnetism.1886

However for the aqua compound, it is determined that it is actually high spin.1893

One, two, three, four, five, six.1898

You see that we get plenty of unpaired electrons.1901

For the aqua complex, we expect paramagnetism.1904

You see how useful crystal field theory can be in explaining for the magnetic behavior of certain coordination complexes.1909

Now a brief discussion on color.1919

We want to review what is known as complimentary colors.1922

Complimentary colors is coming really from physics.1930

This tells us that the color of light that an object absorbs will be complimentary1934

to the color of light that is transmitted to our eyes, basically what you and I see.1940

You have tried to construct something called a color wheel like so.1947

Basically complimentary colors are going to be opposite to each other.1957

For example, if something is going to absorb orange color light, something is going to look blue to us.1961

How does this relate to crystal field theory?1972

Basically that crystal field splitting energy Δo is going to be proportional1976

to the energy of the photon absorbed which means that it is1984

complimentary to the energy of the photon transmitted or given off1995

or the photon that travels to our eyes, what you and I see.2010

If Δo is large, a weaker color of light is going to be given off.2016

If Δo is small, a stronger color of light is going to be given off.2034

If Δo is very large, you basically have a good bet of getting something like red, orange, or yellow.2047

If Δo is small enough, the object will most likely look violet blue or green to us or any combination thereof.2055

Once again using crystal field theory, we are able to explain for both the color and magnetic properties of coordination compounds.2067

This is very fascinating material actually.2074

To summarize, coordination compounds contain a central metal atom or ion attached to ligands which may be atoms or molecules themselves.2079

Coordination compounds are often observed with coordination numbers of 4 or 6, giving tetrahedral, square planar, and octahedral geometries.2089

Finally as we just demonstrated, crystal field theory can explain for both the color and magnetism of coordination compounds.2097

Now let's go ahead and tackle some sample problems; name the following compounds.2108

Here we see that the coordination complex is a cation and that the chlorides are the anions.2115

We know hydroxide is 1-; we know that water is neutral.2129

We know that each chloride here is 1-.2135

This tells us that in order to cancel the overall charge of the chlorides, that iron here must be 3+.2139

Let's go ahead and name them; alphabetical order.2151

We are going to get aqua first; there is five of them.2153

That is going to be penta-aqua followed by hydroxide; but hydroxide becomes hydroxido.2156

Then followed by iron; we don't use ferrate because this is a cation coordination sphere.2166

Then chloride; so pentaaqua hydroxido-iron(III) chloride.2175

Now onto the bottom one; here we have also another one.2184

Your bracket is your cation; your sulfate of course is going to be your anion.2189

We know that ammine is 0; we know that water is 0.2196

Which means that in order to balance the charges here, we have three sulfates.2200

Actually this should only be one sulfate here... my apologies.2214

We know that sulfate here is 2-; the cobalt here must be 2+ overall.2222

Let's go ahead and name this compound.2228

This is going to be triamine triaqua cobalt(II) and then sulfate.2230

That is sample problem one, focusing on nomenclature of coordination compounds.2244

The last sample problem deals with magnetism.2252

What change of magnetic properties if any can be expected when NO21- ligands2255

in an octahedral field are replaced by chloride ligands in a d6 complex?2261

When we look up the spectral chemical series, we see that the NO21- ligand is going to be a stronger field ligand than Cl1-.2266

This is going to give us a larger Δo.2279

This is going to give us the smaller Δo.2283

Remember large Δo means it is going to be low spin.2287

Small Δo means it is going to be high spin.2294

Let's go ahead and draw the possibilities out.2298

d6 low spin means that all of my electrons are paired.2301

That is going to be diamagnetic.2308

Here we have d6 high spin.2310

All of a sudden, we get at least one unpaired electron.2315

That is going to be paramagnetic.2319

To answer this question, when you replace the NO21- ligands with the chloride ligands,2321

you go from a compound that is diamagnetic to one that exhibits paramagnetism.2327

That is our lecture today on the transition elements.2335

I want to thank you for your time.2339

I will see you next time on Educator.com.2340

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