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Franklin Ow

Franklin Ow

Structure of Atoms

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 Mohsin Alibrahim on July 18, 2017

All due respect, but you are note are mess and you keep talking like a robot without so much elaboration. The atom is not something you see in your daily life so you can't talk like someone who is reciting bunch of common sense information

0 answers

Post by Mahsa Khallaghi zadeh on September 22, 2015

I don't know how to solve this!
A ground state H atom absorbs a photon of wavelenght 104.54nm,and its electrons attain a higher energy level. the atom then emits two photons: one photon of wavelength 235nm to reach an intermediate energy level and the second photon to return to the ground stat.what intermediate level did the electron reach?

Related Articles:

Structure of Atoms

  • Light is a form of electromagnetic radiation which can be described by its energy, wavelength, and frequency.
  • The Bohr Model, which gives a simplistic view of photoemission, incorrectly assumed that electrons traveled in fixed, circular orbits around the nucleus.
  • The advent of quantum mechanics resulted from the idea that matter has both a wave-particle nature.
  • Heisenberg’s Uncertainty Principle states that we can never know the exact location of a moving electron.
  • Schrodinger helped develop wavefunctions, giving rise to the concept of the atomic orbital.

Structure of Atoms

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 0:07
  • Introduction 1:01
    • Rutherford's Gold Foil Experiment
  • Electromagnetic Radiation 2:31
    • Radiation
    • Three Parameters: Energy, Frequency, and Wavelength
  • Electromagnetic Radiation 5:18
    • The Electromagnetic Spectrum
  • Atomic Spectroscopy and The Bohr Model 7:46
    • Wavelengths of Light
  • Atomic Spectroscopy Cont'd 9:45
    • The Bohr Model
  • Atomic Spectroscopy Cont'd 12:21
    • The Balmer Series
    • Rydberg Equation For Predicting The Wavelengths of Light
  • The Wave Nature of Matter 15:11
    • The Wave Nature of Matter
  • The Wave Nature of Matter 19:10
    • New School of Thought
    • Einstein: Energy
    • Hertz and Planck: Photoelectric Effect
    • de Broglie: Wavelength of a Moving Particle
  • Quantum Mechanics and The Atom 22:15
    • Heisenberg: Uncertainty Principle
    • Schrodinger: Wavefunctions
  • Quantum Mechanics and The Atom 24:02
    • Principle Quantum Number
    • Angular Momentum Quantum Number
    • Magnetic Quantum Number
    • Spin Quantum Number
  • The Shapes of Atomic Orbitals 29:15
    • Radial Wave Function
    • Probability Distribution Function
  • The Shapes of Atomic Orbitals 34:02
    • 3-Dimensional Space of Wavefunctions
  • Summary 35:57
  • Sample Problem 1 37:07
  • Sample Problem 2 40:23

Transcription: Structure of Atoms

Hi, welcome back to Educator.com.0000

Today's lecture in general chemistry is going to be the structure of the atom.0002

Here is the lesson overview; we are first going to do a brief introduction.0009

Followed by the introduction, we are going to talk about something we call electromagnetic radiation.0015

Then we are going to get into the early model of the atom which was called the Bohr model.0021

We are then going to discuss what is called the wave nature of matter.0027

The core of this chapter is the currently accepted theory for the structure of the atom.0034

That is going to be what is called quantum mechanics.0041

One implication of quantum mechanics is that we get to visualize what are called atomic orbitals.0045

After that, we will go ahead and summarize, followed by a few sample problems.0055

Here is our introductory slide.0063

From previous lectures, we have talked about Earnest Rutherford's gold foil experiment.0066

If you recall, Rutherford was able to find that neutrons and protons existed inside the nucleus while electrons existed outside the nucleus.0072

That was a very important finding because we were able to come up with a positively charged center that is very dense.0086

Every atom has it, which we call the nucleus.0095

But then some questions immediately arose.0098

If electrons are outside the nucleus and protons are inside the nucleus, why don't they collide with each other?0103

Why don't electrons collide with protons and crash into the nucleus because opposite charges attract?0111

If this collision occurred, then the atom should itself collapse.0119

As you see, this structure of the atom was almost paradoxical.0126

It is a very good theory; but why doesn't the atom collapse?0140

At the time, there wasn't really anything that could explain for this paradox.0144

Before we get into more detail, we now have to discuss the electron and its relationship with light.0153

The light that we see as humans is a form of energy, what we call electromagnetic radiation.0166

Radiation can be pretty much characterized by three parameters.0174

Energy which we are going to symbolize with capital E.0178

Frequency which is going to be symbolized with the Greek symbol nu(ν).0181

Finally wave length which is going to be symbolized with the Greek symbol lambda(λ).0186

Energy, we have talked about energy before.0191

This is going to be usually in units of joules.0194

Frequency, it is like the frequency on your radio.0197

That is going to be in the units of hertz.0200

Hertz is the same thing as reciprocal seconds.0203

Finally wavelength, this is going to be usually in meters or nanometers.0207

Basically if we have a travelling wave, you can think of a sine or a cosine function.0218

We have a traveling wave.0224

The wavelength is just the distance between two adjacent peaks or two adjacent troughs.0226

Frequency is the rate at which a wave passes a specific point in time.0236

The higher the frequency, really you can think of it as the faster it is passing that same point in time.0244

What is the mathematical relationship between wavelength and light?--it is the following.0252

c is equal to λ times ν; it is equal to a constant.0258

This constant is something you will recognize; it is the speed of light.0263

The speed of light is approximately 3.0 times 108 meters per second.0271

The important thing from this equation to remember is the proportionality relationship between wavelength and frequency.0279

As we see, they are inversely related; inversely proportional.0287

In other words, the longer the wavelength, the shorter or smaller the frequency.0293

Now that we have been introduced to the three parameters used to characterize electromagnetic radiation, let's now look at the different types of radiation.0302

The different types of radiation collectively is what we call the electromagnetic spectrum.0312

You will find this in any beginning physics course and general chemistry course.0320

Probably you have encountered this some time or another.0326

On the left side of the spectrum is what we have, the long wavelength value; this is long λ.0333

Remember because λ is inversely related to frequency, this means that this is small frequencies.0343

Another important relationship is the following; that energy is directly related to frequency.0353

Long λ, small frequency, also means small energy.0361

Right off the left end here is going to be the radio waves,0370

the same radio waves that we listen, that we use to listen to music, etc.0374

After radio waves is going to be microwaves; after microwaves is going to be infrared.0380

What infrared is, it is abbreviated IR; it is pretty much...0387

You can think of infrared as any type of heat that you feel coming from a warm object.0390

Then we get into the visible region.0397

This is the visible region that you and I as humans see.0400

As you can see, this is in the order of R-O-Y-G-B-I-V.0403

The colors of the rainbow are always in that specific order.0410

They are literally ordered by their energies.0413

In this case, red is going to be the weakest colored light in terms of energy.0417

Violet is going to be the greatest colored, the greatest energetic light in the visible region.0425

After visible comes ultraviolet; the UV light basically that comes from the sun.0439

After UV, we are getting into x-rays.0446

You can already tell, these are terms you recognize already from everyday language.0449

We are getting indeed more and more energetic.0454

Finally we finish it off with gamma rays.0456

Gamma rays is the type of energy that can be emitted from a nuclear power plant.0459

Now that we have studied the different types of electromagnetic radiation, we want to now try to relate it to the electron.0468

In order to do that, I need to now go over with you what happens when light passes through a prism.0479

When white light passes through a prism, we have all seen this.0486

It is separated into the different colors of the visible spectrum, and we essentially get the rainbow.0490

But that is white light; what happens when we do the following?0497

When we excite an element vapor, maybe with electricity, it turns out that the light that we see0501

as humans is not a rainbow but instead it is going to be specific colors of light.0510

In other words, it is different wavelengths unique to the element are emitted.0515

Sodium vapor, when it is excited, we get an orange light.0520

This is the same light that is used in residential street lamps.0525

It is the same color light that you see when hot water touches the blue flame on your stove.0529

It is because of sodium ions get excited in the water, giving off the same colored light, orange-yellow.0535

Xenon vapor, xenon vapor, sometimes you will see xenon being used in very nice expensive headlights on your car.0543

You can recognize that as a nice blue color, purple color.0555

The emission of light can be used as a fingerprint identification for elements.0560

In other words, different elements give off different colors or emit different colors.0568

Different elements give off different colors.0583

Why is this so?--why do elements give off different colors?0587

In order to do that, we are now going to introduce what was commonly accepted at the time for atom structure.0591

This is called the Bohr model of the atom.0597

Niels Bohr came up with a model as to why different elements give off different wavelengths of light.0600

Essentially you have the nucleus right at the middle.0608

You have what are called circular orbits around the nucleus.0613

Immediately you will recognize this as analogous to a solar system.0621

Not only is this called the Bohr model, this is also called the planetary model of the atom.0630

Basically there is a couple of things we want to remember.0643

Electrons which we know from Rutherford's experiment exist outside the nucleus.0646

They travel in these circular orbits around the nucleus.0651

The larger the orbit, the higher the energies.0657

Basically the farther out you go, the higher the energy.0659

Electrons can instantaneously transfer between orbits, but they cannot exist between.0664

They can exist in n equal to 1, n equal to 2, n equal to 3, but not in between.0672

Finally how did Bohr explain for light emission?0679

Basically energy is always going to be released in the form of light when an electron does the following.0683

As it falls from a higher energy to a lower energy level.0689

If we look right here, we have an electron going back to the n equal 1 level from equal 2.0694

Energy is going to be emitted in the form of light that you and I see.0704

We will write this as n equal 2 to a n equal 1 transition.0711

Once again energy is released in the form of light as you go from a higher energy to a lower energy level.0718

Or as you go from an outer orbit to an inner orbit; outer orbit to an inner orbit.0724

Again this is what we call the planetary model of the atom or Bohr's model.0737

The first element that we will study is of course going to be the simplest one.0745

This was hydrogen.0748

When hydrogen in the gas phase is excited, we can go ahead and take a look at the colors of light that are given off.0750

The Balmer series is commonly known; here I show it to you.0756

It shows four lines, each of a different color, in the visible region.0763

The technical term for these fine separate lines is what we call discrete.0769

This is what we call the Balmer series or hydrogen's photoemission spectrum.0776

Photo, it means light; emission means given off.0781

It turns out that we can come up with an equation that actually predicts the wavelength of light that is given off.0786

For hydrogen, this equation is what we call the Rydberg equation.0796

It is basically 1 over λ is equal to a constant R times the following.0802

1 over n1 squared minus 1 over n2 squared.0808

R is what we call the Rydberg constant.0813

It is equal to roughly 1.1 times 107 reciprocal meters.0819

You should probably ask your instructor if you need to memorize this or not.0823

λ, this is going to be the wavelength in meters of the light that you and I detect with our eyes.0827

This is going to be in meters.0838

Again this is the wavelength of light that you and I detect that is given off.0840

1 is going to be equal to the initial energy level or orbit.0846

n2 is going to be equal to the final energy level or the orbit.0859

This equation works remarkably well for hydrogen; the interesting to note is the following.0866

For this equation to work, the energy level of an orbit can never be equal to zero.0877

We cannot take 1 over 0 for example; the important implication is the following.0884

Because n can never be 0, electrons cannot exist inside the nucleus.0890

We answer the question, why doesn't the atom collapse?0898

Because using this mathematical equation, we cannot have an electron existing inside the nucleus.0902

Now that we have talked about atomic spectroscopy and photoemission,0914

we want to now get into a little more philosophical arguments0919

and the early work that has been done in quantum mechanics.0924

This is what we call the wave nature of matter.0929

Matter which of course includes electrons which traditionally viewed as behaving as particles.0933

This is coming from classical physics.0942

What we mean by particles, that means we can plot a trajectory.0949

We can come up with an equation for it basically.0954

The Bohr model came from classical physics.0957

It thought of electrons as travelling in these fixed circular trajectories around a nucleus.0961

However an interesting finding had occurred.0969

When you take a wave and you pass it through a very small slit,0974

you can capture the resulting image on a piece of photographic film.0988

What happens is you get an alternating pattern.0996

The beam is essentially going to split into such a pattern.1000

The dark area represents what we call destructive interference.1012

The other areas represent constructive interference.1023

Again this is all coming from physics.1028

Some of you may have had this already in your physics course.1030

Basically destructive interference is when two waves essentially cancel each other out.1035

Constructive interference is when the waves, their amplitudes are combined.1046

This is highly characteristic of any wave.1057

You get this alternating pattern between destructive and constructive interference when a wave has passed through a very narrow slit.1062

This experiment was then repeated with a beam of electrons.1073

When the beam of electrons passes through this very narrow slit, it turns out that you get the same pattern.1079

This is a highly interesting because we thought of electrons and matter as only behaving as particles.1091

But what this shows, this shows that electrons and therefore matter also has characteristics of a wave.1101

Matter has also wave-like characteristics; wave-like character in addition to having particle character too.1110

This term was what we call the wave particle duality of nature.1145

This early experiment pretty much helped to give rise to a new1154

school of thought early in the twentieth century and late nineteenth century.1160

Basically the new thought viewed matter as having both a wave and particle duality.1168

There were many famous well-known scientists whose names you are going to recognize1178

that helped to contribute to this new movement known as quantum mechanics.1183

Einstein described energy in what he called photons which were essentially discrete quantized packets of energy.1189

In addition to being described as a moving wave, photons he thought were described as tiny packets of energy, something like that.1207

Two more scientists were Hertz and Planck.1217

Hertz and Planck helped to describe what was called the photoelectric effect.1220

Basically when a metal absorbs energy, light can be emitted in the form of photons1225

having an energy E is equal to hν minus φ1232

where h is equal to Planck's constant and phi(φ) is what we call the threshold energy.1237

Basically it is literally the energy required to remove an electron from a metal surface.1249

Energy needed for electron removal from a metal surface.1261

Another scientist that came along was de Broglie.1276

He helped to develop an equation which gave the wave length of a moving particle.1279

Here we are finally being able to quantify the wave-like nature of matter.1283

Basically the wavelength of a moving particle is equal to λ which is equal to h over mv.1290

h is Planck's constant again.1298

m is going to be the mass of the particle but in units of kilograms.1300

Finally v is just your ordinary velocity in units of meters per second.1306

In this case, λ will be in units of meters then.1314

Again we were able to early on from these scientists' contributions,1320

quantify both the wave character and the particle character of matter.1327

Moving on, another important fundamental contribution to quantum mechanics was Heisenberg.1338

Heisenberg came along and developed his very groundbreaking what we call uncertainty principle.1346

He basically said the following.1354

That for a moving electron, we cannot simultaneously ever know both the position and the momentum.1356

Both the position and momentum of a moving electron can never be simultaneously known.1364

Instead only a probabilistic determination of a moving electron's position can be formulated.1370

What does that mean?--that means we can give a pretty good guess or1376

estimate of where an electron may be found, where it may occur around a nucleus.1382

Finally the next scientist was Schrödinger.1390

Schrödinger helped to develop the concept of what was called a wave function.1394

What a wave function is, it is a mathematical function that describes the energy of a moving particle with respect to time.1398

A wave function then, this is what we also call an atomic orbital.1407

An atomic orbital is basically a probability map of where an electron may occur around the nucleus.1408

A probability map of electron location around a nucleus.1427

Moving on, if we took a look at these wave functions, we will see that they will be somewhat complicated.1444

We are going to leave that for a upper division quantum mechanics course.1452

But what I want you to take away is the following.1458

If we were to solve the wave function and we were to look at these atomic orbitals1462

in more detail, the solutions contain what are called quantum numbers.1469

Basically there are four quantum numbers.1474

Number one is what we call the principal quantum number, n.1476

n describes the overall energy of an electron.1479

I want you to think back to the Bohr model of the atom.1484

We use lowercase n to describe each ring around a nucleus, each orbit.1486

It is pretty much the same thing.1491

As the energy level increases, so does the n value.1493

Please note, n cannot be 0.1498

n is only going to be a positive whole number, that is not zero.1500

The second quantum number is what we call the angular momentum quantum number.1506

This is going to be symbolized with lowercase l.1513

This is going to describe the shape of the atomic orbital.1516

Where in other words, if we take the wave function, this mathematical equation,1521

and we plot it in three-dimensional space, we get an image basically.1525

These images... we are going to take a look at it on the next slide... all have very characteristic shapes.1531

l, possible values are going to be 0, 1, 2, all the way up to and including n-1.1539

If l is equal to 0, the orbital designation is what we call an s orbital.1550

Again this is going to make more sense in a slide or two.1562

If l is equal to 1, we call the characteristic shape, a p orbital.1565

If l is equal to 2, we call the characteristic shape, a d orbital.1571

Finally if l is equal to 3, we call the characteristic shape of the wave function in three-dimensional space, an f orbital.1577

Moving on, the third quantum number is what we call the magnetic quantum number.1588

Magnetic quantum number is going to be symbolized with lowercase ml.1593

It describes the spatial orientation of the wave function in three-dimensional space.1599

ml will be equal to ?l all the way to 0 and then all the way to +l.1604

There is something important to make a note of.1614

If l is equal to 0, then ml is just equal to 0.1618

l equal to 0 is an s orbital.1625

ml is just equal to 0; that is only one value.1627

If l is equal to 1, then ml can be equal to -1, 0, +1.1633

If l is equal to 2, ml can be equal to -2, -1, 0, +1, and +2.1641

Finally if l is equal to 3, we get to -3, -2, -1, 0, +1, +2, and +3.1650

The reason why I am bringing this up is going to play a very big role in the next presentation.1660

But please make a note.1665

s orbitals, if you look at it, there is only one value of ml.1667

That means there is only one s orbital per energy level.1671

What we are looking at is not the specific values of ml, but the numbers, how many ml values do we have.1682

For the p orbitals, -1, 0, 1, which means we are going to have three p orbitals per energy level.1691

For the d orbitals, there are five ml values which means we are going to have five d orbitals per energy level.1702

Finally f orbitals, there are seven ml values which means we are going to have seven f orbitals per energy level.1711

Finally the last quantum number is what we call the spin quantum number.1723

The spin quantum number is going to be symbolized with lowercase ms.1728

Basically there is only two values, + or ? 1-1/2.1733

It describes the relative spin of an electron.1737

You can have either spin up or spin down.1741

We are going to be using that terminology a lot in the next presentation.1743

Those are the four quantum numbers that are used to characterize and describe an atomic orbital for an atom.1749

What do these atomic orbitals look like?1757

What do the wave functions when they are plotted in three-dimensional space look like?1760

There is two ways to go about this.1767

The first one is what we call the radio wave function.1769

When we do a radio wave function, it is basically describing electron density at different distances from the nucleus.1776

When we plot for example what we call a 3s orbital, this is going to be our zero line.1787

We are going to get graph that looks like that.1798

This zero line means zero electron density.1803

When we go ahead and look at a 3p orbital, we are going to get something like this.1811

Finally let's go ahead and look at 3d.1821

3d is going to go ahead and look like this.1826

We need a point of reference in order to compare it.1830

Because this is hard for us to look at this overlapped.1832

The point of reference is going to be the same on each of these graphs--here, here, and here.1835

What the x-axis is is the distance from the nucleus.1842

Basically let's examine the s orbital.1853

The s orbital tells us that there is quite a high density of electrons very close to the nucleus.1857

The p orbital tells us that we also have a good density of electrons close to the nucleus, but it is less than 3s.1866

Finally for 3d orbitals, the electron density is farther out from the nucleus.1875

This is very important.1882

The value of 3 is basically the value of n.1886

As you can see, as you go from s to p to d in the same n value,1891

basically the electron density is going to be very high farther and farther out.1903

In other words, the atomic orbital size is increasing; atomic orbital increases in size.1908

That is what we can take away from these radio wave functions.1922

Another way of plotting this is what we call the probability distribution function.1930

This is basically the probability, not electron density, but probability of finding an electron at a certain distance from the nucleus.1935

When we do it for 3s, we are now going to get a graph that looks like this.1944

We do it for 3p, we are going to get a graph that looks like that.1951

Finally for 3d, we are going to get something that looks like that.1957

Once again the y-axis is going to be probability.1963

The x-axis again is going to be distance from the nucleus.1968

What we want to point out here is the intercepts.1982

We want to point out the intercepts here.1988

These intercepts literally mean I have gotten zero probability of finding an electron that specific distance from the nucleus.1991

These are what we call radial nodes.2000

What a node is when we have zero probability at that point from the nucleus.2005

As you can see, this 3s orbital has two nodes, 3p has one node, 3d has zero nodes.2010

The simple equation predicts the number of radial nodes for the atomic orbital.2018

The number of radial nodes equals to n minus l minus 1.2023

These are the different ways of plotting wave functions.2039

But that is two-dimensional graphs.2043

What happens now when we plot it in a three-dimensional space?2046

This is what we typically show at this level of general chemistry.2050

You may have had this even in high school.2057

An s orbital, as you can see, how do you tell?2062

There is only m value; ml value.2067

An s orbital is basically going to be a sphere.2070

A sphere therefore is going to have greatest density right at the center or close to the nucleus.2078

For these, we have three ml values; remember we call that p orbitals.2086

p orbitals are also known as a dumbbell shape.2092

Basically those nodes are right at the middle.2099

We have electron density farther from the nucleus, not at the center.2105

Next one, here five of the ml values.2111

These are there for the d orbitals; again you have nodes at the center.2117

Finally you have seven here ml values; these are going to be the f orbitals.2128

What I want you to take away from this is that again as you go2135

from s to p to d to f, the orbital size increases which means you are going to2139

have a higher probability of finding an electron farther away from the nucleus.2147

Again these are the characteristic shapes of atomic orbitals.2154

Let's go ahead and summarize this presentation on the structure of the atom and quantum mechanics.2159

We first started off the presentation with a look at light.2166

We saw that light is a form of what we call electromagnetic radiation.2172

It can be described by energy, wavelength, and frequency.2176

The early model of the atom was the Bohr model or also known as the planetary model.2181

It gave a very simplistic view of photoemission which incorrectly assumed that electrons travelled in fixed circular orbits around the nucleus.2187

That is where quantum mechanics came in.2196

Quantum mechanics resulted from the idea that matter has both a wave particle nature.2198

Electrons don't travel in fixed orbits.2206

Instead we can only give a probable location of the electron.2208

This is coming from Heisenberg's uncertainty principle.2215

Finally Schrödinger helped to develop what were called wave functions which gave rise to concept of the atomic orbital.2219

Let's go ahead and tackle some sample problems right now.2227

The energy required to dislodge electrons from sodium metal via the photoelectric effect is 275 kilojoules per mole.2231

What wavelength in nanometers of light has sufficient energy per photon to dislodge an electron from sodium?2238

This is 275 kilojoules per mole; that is our energy.2246

We need to go from here to wavelength in energy per photon; maybe in kJ per photon.2258

Basically we are in kJ per mole right now.2269

We want to go to the wavelength of light that has enough energy to dislodge an electron from sodium.2272

This wavelength is going to be in meters.2288

When we look at 275 kilojoules per mole, we want to go to per photon.2294

That is going to be individual particles.2302

We are just going to use our Avogadro's number; divided by 6.022 times 1023 photons.2304

That is going to give us 4.57 times 10-22 kilojoules per photon.2315

We want to go from energy to wavelength.2325

This is going to be energy is equal to hν.2330

ν, this is going to be c over λ.2337

This is our Planck's constant equation basically.2345

That allows us to go from energy to frequency to wavelength.2350

But remember Planck's constant is 6.626 times 10-34 joules times second.2354

Here we are in kilojoules; let's go ahead and get that to regular joules.2366

When we do that, we get 4.57 times 10-19 joules per photon.2371

We are in good shape; we can plug everything directly into the equation.2379

When we do this first part, we can solve for frequency.2384

When we do that, frequency is going to be equal to the energy divided by h.2391

That gives us 6.9 times 1014 reciprocal seconds.2397

When we plug that into ν is equal to c over λ, we can go ahead and solve for λ.2403

We are going to get a wavelength of 435 nanometers.2412

435 nanometers is roughly violet blue light; that is one common sample problem.2418

Another sample problem, let's go ahead and look at it.2427

An electron in the n equals 6 level of the hydrogen atom relaxes to a lower energy level, emitting 93.8 nanometers of light.2430

What is the principal level to which the electron relaxed?2437

Relaxed is just a technical word for fall to.2440

This is involving the hydrogen atom and light that is emitted.2448

We are going to use the Rydberg equation.2452

1 over λ is equal to R times 1 over n1 squared minus 1 over n2 squared.2455

The wavelength of light that emitted is 93.8 nanometers, but we need to get that into meters.2465

That is going 1 over 9.38 times 10-8 meters; that equals to R.2470

1.097 times 107 reciprocal meters times 1 over n1 squared.2478

The n equal to 6 is n1; that is 6 squared.2488

What we are trying to solve for is where the electron fell back to--that is the final energy state.2493

When all is said and done, n2 is equal to 1.2499

How do you know if you have done some miscalculation?2504

Remember n can only be a positive whole number.2506

If you get a really off number, has a lot of decimal places, you probably did something incorrectly.2510

This transition therefore is n equals 6 to n equal to 1.2521

It is going to give off a photon with a wavelength of 93.8 nanometers in hydrogen.2526

That is the structure of the atom and introduction to quantum mechanics.2535

Thank you for using Educator.com; I will see you later.2539

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