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Franklin Ow

Franklin Ow

Gases

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 (12)

1 answer

Last reply by: Professor Franklin Ow
Thu Aug 4, 2016 5:14 PM

Post by Parth Shorey on August 2, 2016

I still don't understand why gas density is inversely proportional to temperature?

0 answers

Post by Peter Ke on October 13, 2015

For the first example at the last lecture, you put .8205 L X ATM/ K X MOL.
Shouldn't it be .08206 L X ATM / K X MOL.

1 answer

Last reply by: Professor Franklin Ow
Mon Mar 30, 2015 11:52 AM

Post by Muhammad Ziad on March 29, 2015

Hello Professor Ow, what does it mean for some parameters to be held constant? Thanks!

1 answer

Last reply by: Okwudili Ezeh
Tue Oct 28, 2014 10:59 PM

Post by Okwudili Ezeh on October 28, 2014

Do you mind checking example 1 again. I got 13.06 liters.

0 answers

Post by Saadman Elman on June 14, 2014

It was really helpful! Thanks a lot!

1 answer

Last reply by: Professor Franklin Ow
Tue Jan 14, 2014 11:14 PM

Post by felicia ekeson on January 14, 2014

hello, I 've tried 734mmHg X 1 atm /0.388L /0.08206 X 25+273.15 but unable to get 0.015 mol

1 answer

Last reply by: Professor Franklin Ow
Thu Nov 7, 2013 5:14 PM

Post by Mark Medina on September 29, 2013

ive been having trouble determine which unit of the universal gas constant to use during a problem. During some i know im suppose to use 0.0821 and during others 8.314. how do i determine which unit of measure to use?

Related Articles:

Gases

  • The Kinetic Molecular Theory of gases are (5) postulates which describe gas behavior. Any gas that is assumed to follow this theory is called an ideal gas.
  • A series of gas laws relates the gas parameters pressure, temperature, volume and moles to each other holding all else constant.
  • The ideal gas law can be used to interpret gas density and its relationship with temperature.
  • We can apply stoichiometry techniques to reactions involving gases.

Gases

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
  • The Kinetic Molecular Theory of Gases 1:23
    • The Kinetic Molecular Theory of Gases
  • Parameters To Characterize Gases 3:35
    • Parameters To Characterize Gases: Pressure
    • Interpreting Pressure On a Particulate Level
  • Parameters Cont'd 6:08
    • Units For Expressing Pressure: Psi, Pascal
    • Units For Expressing Pressure: mm Hg
    • Units For Expressing Pressure: atm
    • Units For Expressing Pressure: torr
  • Parameters Cont'd 8:09
    • Parameters To Characterize Gases: Volume
    • Common Units of Volume
  • Parameters Cont'd 9:11
    • Parameters To Characterize Gases: Temperature
    • Particulate Level
    • Parameters To Characterize Gases: Moles
  • The Simple Gas Laws 10:43
    • Gas Laws Are Only Valid For…
    • Charles' Law
  • The Simple Gas Laws 13:13
    • Boyle's Law
  • The Simple Gas Laws 15:28
    • Gay-Lussac's Law
  • The Simple Gas Laws 17:11
    • Avogadro's Law
  • The Ideal Gas Law 18:43
    • The Ideal Gas Law: PV = nRT
  • Applications of the Ideal Gas Law 20:12
    • Standard Temperature and Pressure for Gases
  • Applications of the Ideal Gas Law 21:43
    • Ideal Gas Law & Gas Density
  • Gas Pressures and Partial Pressures 23:18
    • Dalton's Law of Partial Pressures
  • Gas Stoichiometry 24:15
    • Stoichiometry Problems Involving Gases
    • Using The Ideal Gas Law to Get to Moles
    • Using Molar Volume to Get to Moles
  • 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?
  • 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

Transcription: Gases

Hi, welcome back to Educator.com.0000

Today's lecture from general chemistry is on gases.0002

Going to start off with a brief introduction followed by the following series of topics.0009

The first is what we call the kinetic molecular theory of gases0014

which is basically a bunch of postulates that describe gas behavior.0018

We are then going to go over the parameters that are used to characterize gases--namely pressure, volume, temperature, and moles.0024

When we combine those, we are going to get a series of gas laws which relates all four of those parameters.0035

All of these simple gas laws are then going to culminate into what we call the ideal gas law.0042

After we study the ideal gas law, we are then going to go over some applications of the ideal gas law0048

because we can come up with several additional parameters straight from that law.0054

A unique topic then is going to be gas mixtures and what we call partial pressures.0061

Finally the last topic is going to be stoichiometry and applying it to reactions that involve a gas.0067

Then of course as always, we will finish off with a summary followed by some sample problems.0077

Basically there are five postulates to the kinetic molecular theory of gases.0085

The first one is the following.0092

It deals with gas motion and basically tells us that gases travel in straight lines obeying Newton's laws.0095

They have straight trajectories; they are in constant motion.0106

Number two, the molecules in the gas occupy no volume.0110

That is we treat them as individual points.0114

In other words, if you look at a gas sample, most of it is actually empty air.0119

Number three, when gas molecules collide, we say that they follow elastic collisions.0124

That is upon collision, there is no loss of energy.0131

There is no transfer of energy.0136

You can imagine a bunch of billiard balls colliding with each other.0138

After they collide, they bounce off each other and go their separate ways.0143

Imagine that but pretty much going on infinitely with no loss of energy.0147

Number four, there are no attractive or repulsive forces between gas molecules0153

which explains why gases are so diffuse if you will.0159

Finally when we talk about the kinetic energy of a gas,0166

the kinetic energy of a molecule is really related to its kelvin temperature.0171

Any gas that follows these five postulates of the kinetic molecular theory, we call that gas an ideal gas.0178

In reality, there is no such thing as an ideal gas.0190

But by applying this model, it allows us to make a lot of simplifications and0195

a lot of assumptions which allows us to further study the gases and0200

use models which help us to describe gas behavior pretty well actually.0206

That is the kinetic molecular theory of gases.0215

We now next turn to the parameters that are used to characterize gases.0219

The parameters are basically pressure, volume, temperature, and moles.0224

From physics, pressure is formally defined as the amount of force per unit area.0231

Pressure is equal to force per unit area.0237

If we had for example a flat surface and we put a box on it,0246

that box is applying a downward force on the surface.0252

In other words, the box is applying a certain pressure on the surface.0256

If we take the same box but we stand it upright this time, same box,0260

my area of contact is now a lot smaller than in the first picture.0266

In this case, because the area is smaller, the pressure is going to be larger.0273

That is our formal definition of pressure as defined from physics.0284

But in terms of gases, we are going to describe pressure from a particulate level diagram.0290

Imagine a container.0297

Basically we have gas particles that are moving in random directions once again following0301

the kinetic molecular theory of gases and moving in straight lines, constant motion.0307

Not only do the gas particles collide with each other but the gas particles0314

also collide with the container wall; collision with container wall.0319

When that collision occurs with the wall of the container, that itself0331

generates a force just like billiard balls colliding with each other.0334

It is this force of impact that we tend to relate to gas pressure.0339

Force of impact is proportional to gas pressure.0348

For chemistry, for gases, this is our interpretation of pressure.0362

Now that we have defined pressure, let's go ahead and talk about0370

the common units of pressure that are used to make measurements.0373

From physics, the typical units of pressure are going to be psi which is pounds per square inch and pascal.0380

All of us see psi in tires; we also see psi for water pressure too.0388

We are not going to use psi and pascal too much.0400

Really in chemistry, we are going to use these three--millimeters of mercury, atm, and torr.0402

Millimeters of mercury is what we see on a barometer for the weather report.0407

We also see that whenever you take your blood pressure reading.0414

Atm stands for atmospheres or one atmosphere of pressure.0419

Basically one atmosphere of pressure is going to be the pressure roughly at sea level.0430

Once again at sea level, we are at roughly one atm of pressure.0440

Finally torr, torr is named after Torricelli who invented the barometer.0445

Given at the bottom, we have the relationships for each of these units.0462

One atm is equal to 14.7 pounds per square inch which is equal to 760 millimeters of mercury0467

which is equal to 760 Torr which is equal to 101.325 kilopascals.0475

Of course, you see all of these equivalent statements.0481

From equivalent statements, we can then use them as conversion factors.0483

Now that we have talked about pressure, let's continue on.0490

From the kinetic molecular theory, we are told that gases travel in straight paths.0494

This implies the following.0500

That gases are going to travel in straight paths until they collide with something,0502

either with each other or the wall of the container.0505

What that says is that gases are going to expand to fill their entire container.0508

Hence the volume of a gas is strictly determined by the container that it is placed into.0514

The reason why we can do this is because gases are compressible.0522

Remember that most of a gas is empty air, that they have negligible volume.0526

I could take the same amount of gas in a ten liter bottle0533

and compress it easily to a smaller bottle, no problem.0536

The common units of volume are going to be liters of course and milliliters.0541

Sometimes milliliters, you also see cc or cubic centimeters.0545

Another postulate from the kinetic molecular theory is the kelvin temperature.0553

We say that the kinetic energy of a gas is proportional to its kelvin temperature.0559

In other words, temperature of a gas is directly related to how fast these particles are moving.0563

Because it is directly related, this says the following.0570

The faster a gas is, the hotter its kelvin temperature.0573

We can also interpret this at the particulate level.0577

We are saying that kinetic energy is going to be proportional to the temperature in kelvin.0583

If you consider a gas sample, if we apply heat to this, of course that is going to result in0593

much faster motion because the gases will get more kinetic energy; faster motion, higher kinetic energy.0606

Finally the last parameter is the mole amount of a gas.0625

For gases, the amount of a gas is going to be related to moles of course.0632

Those are the four parameters--pressure, volume, temperature in kelvin, and moles.0637

Now that we have introduced the four parameters, we now get into what is called the simple gas laws.0645

The simple gas laws basically do the following.0650

They are a series of equations that relate the four parameters of a gas that we just covered.0653

There are restrictions for these gas laws to work.0660

First we assume ideal behavior.0663

That is the gas is going to follow all five postulates of the kinetic molecular theory.0665

All other parameters are held constant.0671

If I compare pressure and volume for example, that means temperature and moles are being held constant.0675

All else is held constant.0680

Let's go ahead and tackle each of these gas laws now.0682

The first gas law is called Charles's law.0685

Charles's law states that volume and temperature are directly related, holding pressure and moles constant.0688

Volume is directly related to temperature; let's take a look at this.0696

If I have this container and these gas particles are moving in random directions,0702

if I heat it, the gas particles are going to have more kinetic energy.0712

We are going to result in an expansion of the container if the container is flexible.0719

This is going to be larger volume.0728

That is why if you try heating a balloon up, you see the balloon expanding.0733

It is because the temperature is going to increase the kinetic energy of the molecules.0739

They are going to push outward on the container, on the wall of the balloon.0746

The equation to do this, to quantify Charles's law, is V1 over T1 is equal to V2 over T2.0751

That is if we know the initial volume and initial temperature, we can get0762

either the final volume or the final temperature, whichever is not given.0765

The restriction for this equation is that this must be in kelvin.0771

Temperature must be in kelvin.0775

Volume, it can be in any units as long as they are identical units.0777

That is a pretty straightforward equation to use.0787

Volume and temperature are directly proportional to each other.0790

The next gas law is what we call Boyle's law.0794

Boyle's law states the following.0796

That pressure and volume of a gas are inversely related when temperature and moles are held constant.0799

That is pressure is inversely related to volume.0805

Let's think about this; let's say I had two gas particles in a small container.0810

All of a sudden, let's say the size of the container has increased.0819

I am holding everything else constant.0824

When this happens, I am going to have a smaller rate of collision with the container wall.0828

Because my rate of collisions with the container wall is going to be smaller, my force drops off.0844

Therefore my pressure is going to drop off.0853

Once again pressure is inversely related to volume.0859

If you ever go to the higher elevation, you notice that a potato chip bag or a snack bag is always larger.0862

It is because at the higher elevation, the outside pressure is much smaller.0870

To compensate, the air inside the bag is going to expand.0875

This is why balloons also tend to pop the higher they go.0880

Because as the elevation increases, the air pressure gets lower.0884

The air molecules inside the balloon expand against the walls of the balloon.0887

The equation for this, for Boyle's law is P1V1 is equal to P2V2.0894

Once again for this equation to work, P1 and P2 must be identical units.0902

V1 and V2 also must be identical units.0913

Once again this is a rather straightforward equation to use.0917

We can calculate any final pressure or volume given the other three parameters.0919

That is Boyle's law.0926

The next gas law is what we call Gay Lussac's law.0929

Gay Lussac's law is pressure and temperature.0933

It tells us that pressure and temperature of a gas are directly related when volume and moles are held constant.0936

That is pressure is directly related to temperature; that just makes sense.0943

When we have gas particles just like this, let's say this is colder.0950

That is going to generate some pressure.0958

But if we take the same volume, the same box, the same amount, and I heat this sample up, that is going to0960

result in obviously a higher rate of collision with the walls of the container; more collisions with container wall.0969

That is going to increase my force.0982

Because my force goes up, my pressure goes up also.0985

On extreme temperature differences, a car tire is always going to be at a lower pressure when it is colder.0992

Later in the day when it gets much hotter, the pressure will slightly increase because of this difference.1002

The equation for Gay Lussac's law is the following.1009

P1 over T1 is equal to P2 over T2.1011

Once again the units of pressure must be identical.1016

But remember that our temperature is always related in kelvin whenever discussing a gas.1021

That is Gay Lussac's law.1029

The last simple gas law is what we call Avogadro's law.1033

Avogadro's law is the following.1037

That the volume and moles of a gas are directly related when temperature and pressure are held constant.1039

V is proportional to n.1046

Just think about maybe a car tire.1049

You put more air into it; you increase the amount of air.1054

What happens?--the volume increases.1059

The equation for Avogadro's law is V1 over n1 is equal to V2 over n2.1064

Once again the volume units must be identical.1073

n1 and n2 will always be in moles.1080

That is Avogadro's law.1084

If you look at the four gas laws, Charles's law, Boyle's law, Gay Lussac's law,1086

and Avogadro's law, really Boyle's law is the only one that stands out.1091

It is the only where we have something times something is equal to the product of something else.1096

Every other gas law is division on the left side of the equation and on the right side of the equation.1102

Please make a note of that, Boyle's law is definitely the one that stands out.1110

In case your instructor ever requires you to memorize these gas laws.1119

When we put all of these simple gas laws together, they culminate into one equation.1124

This grand equation is what we call the ideal gas law where PV is equal to nRT.1132

When we do this, there are a couple of restrictions.1138

That pressure must be in atm; volume must be in liters; n is simply moles.1140

The temperature must be in units of kelvin as we always have said.1147

There is something we haven't introduced yet; that is what R is.1152

R is what we call the universal gas constant.1155

It is equal to 0.08206 liters atmosphere K mol.1158

Once again you may or may not have to memorize this.1164

Definitely refer to your instructor for that.1166

That is a relatively straightforward equation to use.1169

Pretty much for an ideal gas, I can determine the pressure, volume, moles,1171

or kelvin temperature given any of the other three parameters.1177

Again this is the ideal gas law.1183

Probably something you want to be comfortable with is to solve for a single variable.1187

Pressure here is going to be equal to nRT over V.1193

Volume is equal to nRT over P; n is equal to PV over RT.1198

Temperature is going to be equal to PV over nR.1206

That is again the ideal gas law.1211

The ideal gas law is relatively straightforward to use.1214

But another important aspect of it is that we can derive and come to many conclusions using this law.1217

The first relationship that we are going to get from the ideal gas law is what is called standard temperature and pressure.1226

It becomes very difficult to compare gases because there is so many parameters--pressure, volume, temperature, and moles.1235

A set of universal conditions has been defined.1242

This set of universal conditions is called standard temperature and pressure or STP for short.1246

Standard temperature is 273.15 kelvin; standard temperature is 1.0 atm.1252

When these values are plugged into the ideal gas law, we can go ahead and solve for the ratio of volume to moles.1258

When we solve for this ratio of volume to moles, we get exactly 22.4 liters per mole.1265

This is what we call molar volume; its significance is the following.1271

That at STP, one mole of any ideal gas regardless of its identity occupies a volume exactly 22.4 liters.1277

One mole equals 22.4 liters; that is an equivalence statement.1288

From that, we can use that as a typical conversion factor.1294

Once again molar volume at STP only, 22.4 liters per mole.1298

Another application we can derive from the ideal gas law is gas density.1304

Gas density is going to be measured in grams per liter.1312

We are not going to be in its derivation.1315

But the density of a gas in grams per liter is equal to the following.1317

It is equal to the molar mass of the gas in grams per mole times1322

the pressure in atm divided by the universal gas constant times the kelvin temperature.1330

You can convince yourself that when all the units cancel, we are left with grams per liter.1337

This equation once again is relatively straightforward to use.1342

However there is an important thing that cannot be overlooked.1345

We now have a relationship between density and temperature for gases.1348

You see here that density is going to be inversely related to the kelvin temperature.1353

That means the following.1359

That as temperature of a gas goes up, the gas density decreases.1360

As temperature goes up, gases tend to become lighter.1366

Therefore they tend to rise; this explains why hot air balloons rise.1373

As you heat the gas within the walls of the balloon, the gas becomes less dense than air.1379

It results in a lower density and results in rising of the hot air, bringing the balloon upwards.1385

Once again density is inversely related to the kelvin temperature of a gas.1394

A final gas law that focuses on pressure, this is called Dalton's law of partial pressures.1401

Dalton's law of partial pressures refers to gas mixtures.1409

It tells us the following.1412

Pretty much that the whole is equal to the sum of the parts.1413

The sum of the individual pressures of each gas component is equal to the total pressure of the gas mixture.1419

These individual pressures, the technical term is called partial pressures.1426

Basically very simple--the total gas pressure of a mixture is equal to the partial pressure of1431

the first gas plus the partial pressure of the second gas, etc.1443

Once again the total pressure is equal simply to the sum of all individual pressures.1448

We now come back to stoichiometry.1457

Stoichiometry is something that we spend a great deal of time on.1463

At the basis of stoichiometry was the following.1468

We want to go from moles of A to moles of B.1472

To do this, we use the conversion factor, the mole to mole ratio.1475

From moles of B, you can go to grams using molar mass.1481

You can go to atoms and molecules using Avogadro.1489

You can go to liters if it is a solution using molarity.1501

The same thing applies on the other side to go to moles of A for example.1507

We spent a deal of time doing mole to mole conversion and also mass to mass conversions.1512

All we are going to do now, we are going to apply our knowledge of stoichiometry to gases.1517

If everything is about pretty much getting to moles first, we have an ideal gas law that helps us do that.1524

Typically for gas stoichiometry problems, we are going to use ideal gas law where n is equal to PV over RT.1532

If we are at standard temperature and pressure, we could take a shortcut.1540

We can just use the molar volume to get to moles because we know that one mole equals 22.4 liters.1544

In this case, the ideal gas law not needed.1552

But again this is only at STP.1557

Let's go ahead and do a sample problem then; the question is the following.1561

How many liters of oxygen gas at standard temperature and pressure are needed to form 10.5 grams of water vapor?1566

As soon as I see the letters STP, I know that I am dealing with 1 atm pressure and 273.15 kelvin.1573

I also know that one mole of a gas is going to be equal to exactly 22.4 liters.1582

The first thing you always do in stoichiometry is to make sure the chemical equation is balanced like we have always done.1590

Here we are going to need two hydrogens and two waters.1598

What do we have here?--we have the mass of the water vapor.1605

Somehow we want to go from mass of water vapor all the way to liters of O2 gas.1613

Because I am at STP, the liters of O2 gas is going to1622

come from molar volume which is one mole is equal to 22.4 liters.1627

But in order to get the moles of O2, I first need the moles of H2O.1641

Before moles of H2O, we then have our mass of H2O which is given.1651

There is our basic flow chart; it is pretty much three main steps.1656

Let's go ahead and do this.1660

10.5 grams of water vapor times 1 mole of water divided by its molar mass of 18.016 grams of water.1664

That gives me moles of A.1679

Now from moles of A to moles of B using the mole to mole ratio1681

which is 1 mole of O2 over 2 moles of H2O.1684

Finally now that I am at moles of O2, I can go ahead and1690

use molar volume as a conversion factor to go and get volume.1693

22.4 liters for every one mole of O2.1697

When all is said and done, we get a volume of 6.5 liters that are1703

required at STP for this reaction to make 10.5 grams of water vapor.1707

Let's now go ahead and summarize this lecture.1715

We started off today's lecture with the kinetic molecular theory of gases.1719

It is basically five postulates which describe ideal gas behavior.1723

We then proceeded to tackle the simple gas laws which basically relates the1728

four parameters used to characterize gases--pressure, volume, kelvin temperature, and moles.1732

When we culminated all of these simple gas laws, we arrived at the ideal gas law.1740

The ideal gas law allows us to come up with many applications including density and its relationship with temperature.1745

Finally all of our stoichiometry skills that we established previously can easily apply to gas problems.1754

That is our summary; let's go ahead and do a series of sample problems.1764

Here is sample problem one; you have a 827 milligram sample of a gas.1769

It occupies 0.270 liters when measured at a temperature of 88 degrees Celsius and a pressure of 975 millimeters of mercury.1775

Calculate the molar mass of the gas; let's take it one by one.1784

Here we have mass; here we have volume, temperature, and pressure.1788

The question is asking for molar mass.1797

Molar mass, we all know to be in units of grams per mole.1800

We have the grams already; that is the 827 milligrams or the 0.827 grams.1805

All we have to get then is the moles.1812

Once we have that, we can divide the two numbers to give us the molar mass.1815

We need to get the moles of this gas which is n.1821

We are given pressure, volume, and temperature; that is three out of the four parameters.1826

We can go ahead and use the ideal gas law to help us do this.1834

Moles is equal to PV over RT.1838

Pressure is 975 millimeters of mercury.1845

We have to then go ahead and convert this to atm remember.1851

That is our restriction.1853

We are going to multiply this by the volume in liters.1857

It is already in liters.1860

We are going to divide this by the universal gas constant.1863

We are going to then multiply this by the kelvin temperature.1871

88 plus 273.15; this gets us 0.012 moles.1876

Now we can go ahead and proceed to solve for the molar mass.1889

0.827 grams over 0.012 moles.1892

That is going to be equal to roughly 69 grams per mole for molar mass.1899

This is another nice application of the ideal gas law.1907

It can be used to determine the molar mass of a gas that follows ideal behavior.1911

Let's go ahead and now proceed on to sample problem two.1919

What mass of silver(I) oxide is required to form 388 milliliters of O2 gas1923

when measured at 734 millimeters of mercury and 25 degrees Celsius?1930

Mass is what we want to get; we are given volume.1937

We are given pressure; we are given temperature; guess what?1940

You have chemical equation here which pretty much means you have a stoichiometry problem.1945

Always the first step is to balance.1951

When we go ahead and balance this, we are going to need 2 silver oxides and 4 silvers.1955

We want to go from basically the following.1962

We are given the pressure, the temperature, and the volume of O2.1967

That is three out of four parameters.1971

We can go ahead and get the moles of O2.1973

n of O2 is equal to PV over RT.1976

That is going to be equal to 734 millimeters of mercury times 1 atm divided 760 millimeters of mercury.1981

Going to multiply that by the volume in liters which is 0.388 liters,1994

divided by the universal gas constant, 0.08206 liters atmosphere K mol,2000

times the temperature in kelvin, 25 plus 273.15.2009

Then the moles of O2, we get 0.015 moles of oxygen gas.2015

We want to go from moles of O2 which is what we have.2023

Somehow we want to go all the way to the mass of silver(I) oxide.2027

We know how to do that.2032

This is really now just a matter of doing something we have already learned.2033

We are going to go from the moles of O2 to the moles of silver(I) oxide using the mole to mole ratio.2038

Then on from there is to the mass of silver(I) oxide using the molar mass.2048

Let's go ahead and finish this up.2055

You have 0.015 moles of O2.2058

The mole to mole ratio is going to be 2 moles of silver(I) oxide for every 1 mole of oxygen gas.2062

Then we are going to go ahead and multiply this by the molar mass of silver(I) oxide to get to grams.2072

Its molar mass is 231.74 grams for every 1 mole of silver(I) oxide.2078

We get roughly 7.0 grams of silver(I) oxide that are required.2089

That is another stoichiometry problem that involves gases.2097

Thank you all for your attention; I will see you all next time on Educator.com.2102

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