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Franklin Ow

Franklin Ow

Periodic Trends

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 (9)

0 answers

Post by Parth Shorey on April 25, 2016

I still don't understand why paramagnetic is attracted to external magnetic field?

1 answer

Last reply by: Professor Franklin Ow
Thu May 28, 2015 12:36 PM

Post by Iman Abbas on May 27, 2015

Why is that when we talk about electron affinity and ionization, we refer to just the gas phase only?

0 answers

Post by Saadman Elman on November 27, 2014

Thanks for the video. The explanation was very explicit.

2 answers

Last reply by: Shawn Freeman
Sun May 24, 2015 6:33 AM

Post by John Wadsworth on December 7, 2013

In the two exceptions, the expected configuration of Cr would be [Ar]4s^2 3d^4 and not [Ar]4s^2 3d^6.

1 answer

Last reply by: Professor Franklin Ow
Thu Nov 7, 2013 5:12 PM

Post by Baoer Ye on October 8, 2013

why is Cl in the 7th column number?And O in the  6th ?

Related Articles:

Periodic Trends

  • A consequence of quantum mechanics was the establishment of certain periodic trends, including ionization energy, electron affinity, atomic size and electronegativity.
  • The periodic table is very organized and informative concerning an element’s physical and chemical properties.
  • An atom’s electron configuration can be determined following the Aufbau Principle.
  • Orbital box diagrams are drawn following the Pauli Exclusion Principle and Hund’s Rule.

Periodic Trends

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 0:09
  • Introduction 0:36
  • Electron Configuration of Atoms 1:33
    • Electron Configuration & Atom's Electrons
    • Electron Configuration Format
  • Electron Configuration of Atoms Cont'd 3:01
    • Aufbau Principle
  • Electron Configuration of Atoms Cont'd 6:53
    • Electron Configuration Format 1: Li, O, and Cl
    • Electron Configuration Format 2: Li, O, and Cl
  • Electron Configuration of Atoms Cont'd 12:48
    • Orbital Box Diagrams
    • Pauli Exclusion Principle
    • Hund's Rule
  • Electron Configuration of Atoms Cont'd 17:35
    • Exceptions to The Aufbau Principle: Cr
    • Exceptions to The Aufbau Principle: Cu
  • Electron Configuration of Atoms Cont'd 20:22
    • Electron Configuration of Monatomic Ions: Al
    • Electron Configuration of Monatomic Ions: Al³⁺
    • Electron Configuration of Monatomic Ions: Cl
    • Electron Configuration of Monatomic Ions: Cl¹⁻
  • Electron Configuration Cont'd 24:31
    • Paramagnetism
    • Diamagnetism
  • Atomic Radii 26:08
    • Atomic Radii
    • In a Column of the Periodic Table
    • In a Row of the Periodic Table
  • Ionic Radii 27:30
    • Ionic Radii
    • Anions
    • Cations
    • Isoelectronic Species
  • Ionization Energy 29:00
    • Ionization Energy
  • Electron Affinity 31:37
    • Electron Affinity
  • Summary 33:43
  • Sample Problem 1: Ground State Configuration and Orbital Box Diagram 34:21
    • Fe
    • P
  • Sample Problem 2 36:38
    • Which Has The Larger Ionization Energy: Na or Li?
    • Which Has The Larger Atomic Size: O or N ?
    • Which Has The Larger Atomic Size: O²⁻ or N³⁻ ?

Transcription: Periodic Trends

Hi, welcome back to Educator.com.0000

Today's presentation from general chemistry is going to be on periodic trends.0003

We are going to go ahead and as always start off with a brief introduction.0009

Then we are going to get into something we call electron configuration of atoms.0014

After electron configuration of atoms, we are then going to see how we can use the periodic table0018

to help us determine electron configurations--that is what we call the periodic trend.0023

After that, we are going to go ahead and look at additional periodic trends.0028

Followed by summary and finally a few sample problems.0032

We in the previous presentation discussed quantum mechanics and how it gave rise to what was called an atomic orbital.0038

One of the grand outcomes of quantum mechanics was the establishment of periodicity.0048

What periodicity refers to is the following.0055

That patterns of elements and how they are organized throughout the periodic table can be quite predictable0058

just off of organization in their location on the periodic table.0069

The patterns that we are dealing with have to do with both the physical and the chemical properties of atoms.0075

For today this presentation, we are just going to look at a series of these periodic trends.0082

The first one we are going to look at is what is called electron configuration.0089

Electron configuration of an atom is going to answer the following questions about an atom's electrons.0095

Number one, what type of atomic orbitals are involved--s, p, d, or f?0102

Number two, how many electrons are in each of these orbitals?0111

Electron configuration typically has the following format.0117

It is related to quantum numbers actually.0122

It is always the principal quantum number n, followed by the letter of the atomic orbital.0126

That is going to be s, p, d, or f.0139

There is always going to be a superscript.0143

This superscript is going to be also positive whole numbers--1, 2, 3, etc.0147

Basically this is equal to the number of electrons in that orbital; number of electrons in specific orbital.0156

What we are going to do now, we are going to use the periodic table to help us to determine an atom's electron configuration.0175

Basically when you look at the periodic table, at first it looks like one big mess, but it is incredibly organized.0182

There is so much information we can gather from this very important tool.0189

This first two columns, group 1 and group 2, this is what we call the s block; s block.0196

After the s block comes this whole area in green.0209

It is basically what we call the p block.0212

All of the transition metals are in yellow.0216

That is what we call the d block.0220

The lanthanides and actinides are at the bottom.0224

That is what we call the f block.0229

To read this, the n level, remember n comes first.0233

Basically n is basically just the row that the element is in.0237

n is equal to row number.0241

For example, when you look at hydrogen, you see 1s.0245

When you look at lithium, 2s.0249

When you look at sodium, it is 3s for example.0251

Another part, that s is actually helium.0255

You want to count helium as being right adjacent to hydrogen, just for electron configuration purposes.0259

The d block, this is where n changes.0266

Even though the d block doesn't start until the fourth row of the periodic table,0270

the first n value for the d block is going to be always one less than the row.0274

n is equal to one less than actual row.0279

In the fourth row, the first n value is actually 3.0285

Same thing for the f block; same thing for the f block.0292

It is going to be a little different.0295

You notice here that the f block, we don't start until n equal to 4.0296

But the f block actually starts in the sixth row.0302

The n is going to be two less than the actual row.0308

Basically now what you want to do is you want to count blocks you are in.0320

For example, we can go ahead and let's just start off with hydrogen.0326

Hydrogen has just the one electron; hydrogen is going to be 1s1.0332

When we move over to helium right here, that is going to be a total of two electrons.0343

Helium is going to be 1s2.0350

Lithium, lithium has a total of three electrons.0356

But remember we always start from the beginning; we always start from hydrogen.0360

Starting at the beginning, you get hydrogen 1s1; helium was 1s2.0366

Now lithium coming down through the next row is going to be 2s and just the one block.0372

Remember one block is equal to one electron.0378

When you see something like this, 1s1 1s2, you only count the maximum.0382

You only count the last one when they both have the same n value and the same letter.0387

Therefore I only count the last one; it is just 1s2 2s1 for lithium.0393

I know this may seem a little weird at first.0401

But let's go ahead and look at it more in detail now.0405

How do we fill orbitals?--the filling orbitals is...0407

We are going to base it off of what we call the Aufbau principle; the Aufbau principle.0411

Lithium is what we just looked at; lithium was again 1s2 2s1.0417

Let's go ahead and look at oxygen; oxygen is going to be 1s2 2s2.0427

Now oxygen is in the p block; it is four blocks over.0436

That is going to be 2p4; there is something you always want to check.0441

The superscripts remember tell me the number of electrons.0447

The sum of the superscripts should equal to the number of electrons in that atom.0453

That is how you should always check.0464

When we look at lithium, lithium is element number 3.0466

Indeed, 2 plus 1 gives us 3; that is good to go.0469

Remember how we only count the last one, using Aufbau principle.0473

1s1 1s2 was what we initially had; and then 2s1.0479

If I add those up, I get 4; we know that is not right.0487

That is why you only count the last one; that is oxygen.0491

Let's go ahead and do another example as chlorine.0498

You should follow along on the periodic table.0500

Chlorine, this is going to be... again you always start from the beginning.0502

Again only count the last one; 1s2 completes the first row.0506

2s2 is going to be lithium and beryllium.0511

2p6 completes the p block in the second row.0516

We go down to the next line.0520

It is just like reading sentences in a paragraph.0521

You go from top to bottom, left to right; top to bottom, left to right.0525

After 2p6 comes 3s2; finally now we are into the block for chlorine.0530

Chlorine is in the 3p block five blocks over, 3p5.0536

Again double check, do all of the superscripts add up to the atomic number of the element?0542

Indeed it does.0548

We are going to do electron configuration, two formats; the first way we just did.0551

The second way is to do what we call a noble gas notation.0557

This is going to save us time.0561

Basically what you want to look at is the preceding noble gas; look at preceding noble gas of the element.0563

If I look at lithium, the preceding noble gas for lithium is going to be helium.0577

Helium's configuration is 1s2.0584

If I look at oxygen, the preceding noble gas is also helium.0588

Helium is also represented by 1s2.0595

Finally for chlorine, chlorine's noble gas is going to be argon.0598

Argon is going to be represented... excuse me, not argon... neon.0605

Chlorine's preceding noble gas is neon; that is represented by this entire thing, 1s2 2s2 2p6.0613

The shortcut notation is just to put the noble gas in brackets like this.0623

Followed by all the other electrons to the immediate right; helium in brackets, 2s1.0629

Oxygen is going to be helium in brackets, 2s2 2p4.0636

Finally chlorine, chlorine is going to be neon in brackets and then 3s2 3p5.0644

We want to make a note then; what do these electrons mean?0655

Everything in brackets that is represented by the noble gas, think about this.0662

These are the lowest n values.0666

That means these are the innermost electrons, closest to the nucleus.0672

This is what we call the core electrons.0678

Everything else that is to the right of the bracket has the highest n value, which we just underlined right now.0681

We are going to give these electrons a very important name which is going to come into play later on.0690

These are what we call valence electrons.0695

Valence electrons are the electrons farthest from the nucleus.0699

They are the outermost electrons; outermost electrons.0702

A very important finding is that you can always tell the number of0710

valence electrons for main group elements just by looking at the periodic table.0715

For main group elements, the number of valence electrons equals simply to the column number.0721

Let's go ahead and think about this.0736

For lithium, the valence electron is 2s1.0738

There is only one of them; lithium is in column 1.0742

For oxygen, you have valence electrons of 2s22p4.0746

That is a grand total of six valence electrons; oxygen is in column 6.0751

Finally chlorine has valence electrons of 3s2 3p5; that is a grand total of seven.0756

Look where chlorine is--in column 7 of the periodic table.0763

There is another way of representing the electrons in an atom.0770

This is what we call orbital box diagrams.0779

Basically orbital box diagrams are just schematic depiction to show filling of atomic orbitals.0783

Two rules are essentially followed; very important rules.0792

Pauli exclusion principle, the Pauli exclusion principle states that an atomic orbital can hold a maximum of two electrons which must be opposite spin.0796

Once again the Pauli exclusion principles states that an atomic orbital can hold a maximum of two electrons which must be opposite spin.0807

Hund's rule states that a degenerate set of orbitals are to be filled singly with parallel spin.0818

What degenerate means is that they are going to have the same n value, the same principal energy level.0825

After each orbital is singly filled, you can then go ahead and insert a second electron but with opposite spin.0831

That is a lot to take in; let's go ahead and look at chlorine.0840

Chlorine again is 1s2 2s2 2p6 3s2 3p5.0843

What we are going to do is we are going to use a line or a box to represent each of these orbitals.0855

Remember an s orbital from the quantum mechanics presentation, there is only one per energy level.0860

A p orbital, there are three per energy level.0871

A d orbital means five per energy level.0877

Finally f was seven per energy; again this is only when applicable.0882

For example, a d orbital is not going to appear until n equal to 3.0890

F orbitals won't appear until n equal to 4.0898

That is what I mean by when applicable.0900

Again the periodic table is your best friend in this chapter, in this presentation.0902

You have to use the periodic table and you can see this easily.0906

We are going to go ahead and draw the orbital box diagram for chlorine.0910

That is it; one line represents the 1s orbital; 2s, also represented by one line.0915

2p, remember three per n; one, two, three; this is 1s2s; this is 2p.0923

3s is one line; finally 3p, three lines.0932

Again depending on what textbook you use, sometimes the lines will be instead drawn as boxes.0938

They are the same.0946

What we are going to do right now, we are going to fill electrons using Pauli exclusion principle and Hund's rule.0947

1s2, if I pair electrons, two electrons per orbital, they have to be of opposite spin.0954

That is what Hund's rule tells me.0965

If I put one as spin up, the second electron is going to be spin down.0968

We are representing electrons basically as arrows.0971

2s2, same thing; one spin up, one spin down.0975

The 2p orbital, you notice that there are three of them.0981

That is what we call a degenerate set; they have the same n value.0986

Hund's rule tells me I am going to fill these p orbitals singly first before doubling up.0991

I am going to do this with parallel spin.0996

For example, this is one electron of the six.0999

When I go to the next one and to the next one, that is 2p3 so far.1006

Again I have filled the atomic orbitals singly; now I can double up.1011

They are going to be all of opposite spin--2p4 2p5 and finally 2p6 in that order.1016

Yes it does matter how you write it; very important.1025

Go ahead now onto 3s2; spin up, spin down.1028

Now onto 3p5, once again one electron at a time per orbital before doubling up with opposite spin.1033

One, two, three, now four and five.1041

Once again we have applied Pauli exclusion principle and Hund's rule to what are called orbital box diagrams.1047

The Aufbau principle generally works.1059

There are of course a couple of exceptions that we need to account for.1062

Again you should ask your instructor which one you need to know.1066

But the two most common exemptions are in the first d row.1069

That is going to be copper and chromium.1072

Chromium we would expect to be [argon] 4s2 3d6; expected.1075

But actual configuration is going to be [argon] 4s1 3d5.1086

Copper, copper we would expect [argon] 4s2 3d9; expected.1096

But actual is going to be [argon] 4s1 3d10.1106

We are going to save the reasoning for perhaps a higher level course.1117

But for now, what you want to remember is these exceptions arise from experimental data.1126

The simple explanation is that the configuration is more stable due to what is called a half-filled or a completely filled orbital.1133

You see that this s orbital can hold two electrons.1144

It is holding only one right now; that is half-filled.1149

Here you have five of the d orbitals.1152

Each have two electrons, giving you a maximum of ten.1154

That is going to be a completely filled orbital.1157

This is half-filled; this is completely filled; half-filled and completely filled.1159

Something else to point out, this is going to be true not just for copper and chromium, but for transition metals in general.1172

It is to count the valence electrons.1179

Remember for main group elements, we go by the column number; that is it.1182

But for transition metals, they are not main group elements.1186

What rule do we go by for counting valence electrons?1189

Basically it is a simple rule.1192

You count all of the electrons that come after the noble gas for transition metals.1195

Chromium is going to have a total of six valence electrons.1201

Copper is going to have a total of eleven valence electrons.1207

Once again for transition metals, all electrons coming after the noble gas in brackets will be counted as valence.1212

Let's now take a look at the electron configuration for monatomic ions.1224

We are going to examine the configurations in orbital box diagrams.1230

Aluminum, aluminum is 1s2 2s2 2p6 3s2 and 3p1.1235

What does it mean to be aluminum 3+?1248

3+, that means that we remove three electrons.1252

Is it easier to remove electrons closest or farthest from the nucleus?1257

Electrons that are close to the nucleus, they are being drawn toward the nucleus because of that positive charge.1263

Electrons farther out from the nucleus are not held as tightly so they are easier to remove.1271

When you form cations, you remove valence electrons first.1279

In other words, remove the outermost electrons; 1s2, for aluminum 3+, 2s2 2p6.1288

The first electron I am going to remove is the 3p1 because that is farthest out, followed by 3s2.1304

Aluminum 3+'s configuration is 1s2 2s2 2p6.1312

Let's go ahead and repeat this now for chlorine.1317

Chlorine is 1s2 2s2 2p6 3s2 3p5.1321

Chlorine 1-, that 1- means we have added an additional electron.1330

Once again we are going to add it to the outermost orbital because that is most easily accessible.1335

You add electrons also to valence orbitals.1340

Chlorine 1- is 1s2 2s2 2p6 3s2; then 3p5, you add one more, becomes 3p6.1350

The reason why I wanted to bring this up is because there is something in common that these electron configurations of the ions share.1363

They all end in 2p6 which means they have a completely filled p block.1375

What elements are in that six p block column?--it is the noble gases.1381

This is all noble gas configuration; this is what we call noble gas configuration; noble gas configuration.1388

You can go ahead and check with the periodic table that all of them end in p6; all of the noble gases.1401

The valence electrons for a noble gas, two and six, two and six, which means eight valence electrons now.1409

This number 8 is going to come into play in the next presentation.1423

It is going to become very important.1427

But what I want you to take away from this section is the following.1429

That elements react to form ions and covalent bonds in order to achieve eight electrons.1437

Elements react to achieve eight valence electrons; again just take my word for now.1452

This number eight is going to become hugely important later on.1466

One application of knowing the electron configuration for an atom is we are going to be able to predict an atom's magnetic behavior.1475

Paramagnetism arises when the atom has at least one unpaired electron; one unpaired electron.1483

Why that is important is because paramagnetic species are going to be attracted to an external magnetic field.1491

In other words, they will stick to a magnet if you will.1497

Diamagnetism, the prefix di, you could probably tell what this means already.1501

Diamagnetism arises when an atom or ion contains no unpaired electrons.1505

In other words, all electrons are paired; all electrons are paired.1510

Diamagnetic species are going to be repelled by an external magnetic field.1517

An example of a paramagnetic species, we have a lot of them, maybe like 1s2 2s1.1524

Here we have the one unpaired electron; this will be for lithium.1535

Then let's go ahead and look at an example of a diamagnetic species.1540

Diamagnetic species will be just 1s and then 2s.1544

1s2 2s2, that is going to be for beryllium.1549

Zero unpaired electrons, that is what we mean.1556

Electron configuration was the first big periodic trend.1560

We saw how useful the periodic table came into play to help us predict configuration.1563

Let's now take a look at other periodic trends.1568

Atomic radius, atomic radius literally is the size of the atom.1573

It is the distance an atom's valence electrons are from its nucleus.1577

It is also known as atomic radius or atomic size; it turns out the following.1582

That in a column of the periodic table, atomic size increases top to bottom.1588

The reason is because of the following.1593

As you go down a column, you are increasing the n value.1595

As n increases, so does the size of the atomic orbital, remember the higher the energy.1600

In a row of the periodic table, atomic size decreases left to right.1606

Let's go over the reasoning for that.1613

Within the same row of the periodic table, you are pretty much keeping the same n value.1615

n does not change at all; what does change is the following.1620

That is going to be the number of protons inside the nucleus.1625

As the element gains more protons, its nuclear charge is going to increase1629

which essentially pulls the electrons in to a greater degree.1636

This results in contraction; once again atomic size is going to decrease left to right.1642

Ionic radii, when we look at ionic radii, remember that electrons are added to or taken away from the valence shell.1651

The outermost shell that is; that is very important.1659

For anions, because we are adding electrons to the outermost shell, anions typically are larger than the parent anion.1663

For example, O2- will be greater than just oxygen by itself.1672

Cations, we are removing electrons from the outermost shell.1677

Cations are typically going to be smaller than the parent ion.1682

For example, aluminum 3+ is going to be smaller than aluminum 0.1686

How about for what we call isoelectronic species?1694

Isoelectronic species means we have the same number of electrons.1697

For example, N3- versus O2-, you can look at their electron configurations.1700

They are going to have the same electron configurations, the same number of electrons.1709

The rule for isoelectronic species is just to go by nuclear charge.1712

The element basically with the greater nuclear charge, the greater number of protons,1718

is going to have a greater pull on the electrons, decreasing the size.1724

Therefore N3- which has fewer protons than oxygen is going to be greater in size than O2-.1730

Another periodic trend is what we call ionization energy; let's start off with the definition.1743

Ionization energy is the energy required to remove a valence electron in the gas phase.1750

Again you have to put in energy to overcome the attraction the electron has with the nucleus.1756

Removal of an electron is an endothermic process.1765

Typically then ionization energies are going to be positive values.1770

Think about ΔH in units of kilojoules per mole.1775

Ionization energy is completely opposite to atomic size.1782

Therefore smaller atoms have very large ionization energies; just think about that.1785

When we have smaller atoms, we have the electrons much closer, the valence electrons much much closer to the nucleus.1791

That attraction with the nucleus is incredibly strong.1799

You have to put in more energy to overcome that attraction.1802

For our purposes, fluorine is going to be the smallest of our atoms and therefore the largest ionization energy.1807

There is some additional terminology to get across.1825

An atom's first ionization energy refers to removing a valence electron from the ground state.1828

For example, oxygen goes on to form... excuse me.1834

For example, calcium can go on to form Ca1+ plus an electron when energy is applied.1849

We are just removing the first valence electron.1862

This is what is called the first ionization energy.1865

The second, third, etc, ionization energies then refer to subsequent removal of the electrons from a cation.1868

Let's go ahead and now take Ca1+.1876

Ca1+ now has its turn of being ionized to form Ca2+ plus an electron.1878

Once again this is what is called the first ionization energy.1887

This will be calcium's second ionization energy; very good.1891

The next periodic trend is what we call electron affinity.1899

Electron affinity refers to the energy that is released when an electron is added to an atom in the gas phase.1902

Which means addition of electrons is an exothermic process.1910

We will represent that as for example oxygen plus an electron goes on to form O1-.1915

We can also build on that; O1- plus an electron goes on to form O2-, etc.1922

Again this is what we call electron affinity.1929

Again it is going to be typically an exothermic process which means these are1932

going to be reported in less than 0, negative kilojoules per mole.1937

Electron affinity therefore parallels ionization energy.1945

We are going to find that nonmetals are going to be more apt to1951

gain electrons to form anions which will have the largest electron affinities.1955

What we want to do, we want summarize all of the periodic trends looking at the periodic table.1961

As you go left to right upward, this is a direction of increasing ionization energy and electron affinity.1969

However both of these are opposite to atomic size which increases this way.1982

This is increasing atomic size.1990

Really we use our start point and end point, fluorine being at one corner and then perhaps francium being at another corner.1998

Again fluorine being one of the smallest atoms is going to have the largest ionization energies, a very large electron affinity, etc.2010

These are our periodic trends; of course they come from consequences of quantum mechanics.2024

The periodic table we saw is very organized and informative concerning an element's physical and chemical properties.2033

An atom's electron configuration we learned, we do the filling of the orbitals based off of the Aufbau principle.2040

Finally orbital box diagrams are schematic representations of electrons following two important rules--the Pauli exclusion principle and Hund's rule.2049

Let's go ahead and tackle some sample problems.2063

For each of the following, give the ground state configuration and orbital box diagram.2066

Ground state simply refers to the neutral atom.2071

It is called ground state because it refers to the configuration that represents the lowest energy for the atom; of lowest energy.2078

Iron, the preceding noble gas for iron is going to be argon.2090

Then that is going to be 4s2 3d6.2098

Let's go ahead and draw it out.2102

I am just going to do the valence, the valence electrons; that is 4s.2105

Remember for these, how many d orbitals are there per energy level?2109

There is five of them; you have to draw them always as a degenerate set; degenerate set.2112

Let's go ahead and fill them; 4s2 first; now 3d6.2118

Remember one at a time with parallel spin before pairing up.2122

When you pair up, opposite spin.2127

Finally a phosphorus, phosphorus, you can go ahead and look at the preceding noble gas.2132

Preceding noble gas for phosphorus is going to be neon.2140

After neon is going to be 3s2... it is going to be 3s2 3p3.2143

Let's go ahead and draw that out, 3s.2154

Remember there is three degenerate orbitals for p's; let's go ahead and fill the electrons.2157

3s2 first; now 3p3; one, two, and three; let's just go ahead and complete them.2163

Iron is going to have a grand total of eight valence electrons.2174

Phosphorus is going to have a grand total of five valence electrons then.2179

We expect both of these, because they do have unpaired electrons, we expect both iron and phosphorus to be paramagnetic.2186

Let's go on to now sample problem number two.2200

For each of the following, circle the element and ion that has the larger indicated trait.2203

Basically we are going to use the periodic table all the way.2207

Ionization energy, when we look at sodium and lithium, they are in the same column.2210

Lithium is above sodium; sodium is going to have the smaller ionization energy.2215

Lithium should have the larger one; again the reason is because lithium is smaller.2222

It is going to require energy to remove its valence electron than in sodium.2233

Atomic size for oxygen or nitrogen, oxygen and nitrogen are in the same row.2245

Atomic size decreases left to right; oxygen is going to have the smaller size.2251

Nitrogen is going to have the larger size; remember why?2256

Because nitrogen is going to have fewer protons in its nucleus; so a smaller nuclear charge.2260

Because it is a smaller nuclear charge, electrons are going to be farther out from the nucleus.2270

Last one is atomic size, this time for O2- and N3-.2281

Again anytime you see ions, try to think isoelectronic species.2286

Again for isoelectronic species, you want to just go by the atomic number again; the number of protons.2292

Once again oxygen is going to have more protons.2300

Nitrogen is going to have fewer protons.2306

For the same reasoning as the previous question, N3- should be larger because of the fewer protons.2310

The attraction is less with the electrons; and they are farther out.2319

It was good to see everyone on Educator.com again; thank you for your attention.2325

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