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Raffi Hovasapian

Raffi Hovasapian

Partial Pressure, Mol Fraction, & Vapor Pressure

Slide Duration:

Table of Contents

I. Review
Naming Compounds

41m 24s

Intro
0:00
Periodic Table of Elements
0:15
Naming Compounds
3:13
Definition and Examples of Ions
3:14
Ionic (Symbol to Name): NaCl
5:23
Ionic (Name to Symbol): Calcium Oxide
7:58
Ionic - Polyatoms Anions: Examples
12:45
Ionic - Polyatoms Anions (Symbol to Name): KClO
14:50
Ionic - Polyatoms Anions (Name to Symbol): Potassium Phosphate
15:49
Ionic Compounds Involving Transition Metals (Symbol to Name): Co₂(CO₃)₃
20:48
Ionic Compounds Involving Transition Metals (Name to Symbol): Palladium 2 Acetate
22:44
Naming Covalent Compounds (Symbol to Name): CO
26:21
Naming Covalent Compounds (Name to Symbol): Nitrogen Trifluoride
27:34
Naming Covalent Compounds (Name to Symbol): Dichlorine Monoxide
27:57
Naming Acids Introduction
28:11
Naming Acids (Name to Symbol): Chlorous Acid
35:08
% Composition by Mass Example
37:38
Stoichiometry

37m 19s

Intro
0:00
Stoichiometry
0:25
Introduction to Stoichiometry
0:26
Example 1
5:03
Example 2
10:17
Example 3
15:09
Example 4
24:02
Example 5: Questions
28:11
Example 5: Part A - Limiting Reactant
30:30
Example 5: Part B
32:27
Example 5: Part C
35:00
II. Aqueous Reactions & Stoichiometry
Precipitation Reactions

31m 14s

Intro
0:00
Precipitation Reactions
0:53
Dissociation of ionic Compounds
0:54
Solubility Guidelines for ionic Compounds: Soluble Ionic Compounds
8:15
Solubility Guidelines for ionic Compounds: Insoluble ionic Compounds
12:56
Precipitation Reactions
14:08
Example 1: Mixing a Solution of BaCl₂ & K₂SO₄
21:21
Example 2: Mixing a Solution of Mg(NO₃)₂ & KI
26:10
Acid-Base Reactions

43m 21s

Intro
0:00
Acid-Base Reactions
1:00
Introduction to Acid: Monoprotic Acid and Polyprotic Acid
1:01
Introduction to Base
8:28
Neutralization
11:45
Example 1
16:17
Example 2
21:55
Molarity
24:50
Example 3
26:50
Example 4
30:01
Example 4: Limiting Reactant
37:51
Example 4: Reaction Part
40:01
Oxidation Reduction Reactions

47m 58s

Intro
0:00
Oxidation Reduction Reactions
0:26
Oxidation and Reduction Overview
0:27
How Can One Tell Whether Oxidation-Reduction has Taken Place?
7:13
Rules for Assigning Oxidation State: Number 1
11:22
Rules for Assigning Oxidation State: Number 2
12:46
Rules for Assigning Oxidation State: Number 3
13:25
Rules for Assigning Oxidation State: Number 4
14:50
Rules for Assigning Oxidation State: Number 5
15:41
Rules for Assigning Oxidation State: Number 6
17:00
Example 1: Determine the Oxidation State of Sulfur in the Following Compounds
18:20
Activity Series and Reduction Properties
25:32
Activity Series and Reduction Properties
25:33
Example 2: Write the Balance Molecular, Total Ionic, and Net Ionic Equations for Al + HCl
31:37
Example 3
34:25
Example 4
37:55
Stoichiometry Examples

31m 50s

Intro
0:00
Stoichiometry Example 1
0:36
Example 1: Question and Answer
0:37
Stoichiometry Example 2
6:57
Example 2: Questions
6:58
Example 2: Part A Solution
12:16
Example 2: Part B Solution
13:05
Example 2: Part C Solution
14:00
Example 2: Part D Solution
14:38
Stoichiometry Example 3
17:56
Example 3: Questions
17:57
Example 3: Part A Solution
19:51
Example 3: Part B Solution
21:43
Example 3: Part C Solution
26:46
III. Gases
Pressure, Gas Laws, & The Ideal Gas Equation

49m 40s

Intro
0:00
Pressure
0:22
Pressure Overview
0:23
Torricelli: Barometer
4:35
Measuring Gas Pressure in a Container
7:49
Boyle's Law
12:40
Example 1
16:56
Gas Laws
21:18
Gas Laws
21:19
Avogadro's Law
26:16
Example 2
31:47
Ideal Gas Equation
38:20
Standard Temperature and Pressure (STP)
38:21
Example 3
40:43
Partial Pressure, Mol Fraction, & Vapor Pressure

32m

Intro
0:00
Gases
0:27
Gases
0:28
Mole Fractions
5:52
Vapor Pressure
8:22
Example 1
13:25
Example 2
22:45
Kinetic Molecular Theory and Real Gases

31m 58s

Intro
0:00
Kinetic Molecular Theory and Real Gases
0:45
Kinetic Molecular Theory 1
0:46
Kinetic Molecular Theory 2
4:23
Kinetic Molecular Theory 3
5:42
Kinetic Molecular Theory 4
6:27
Equations
7:52
Effusion
11:15
Diffusion
13:30
Example 1
19:54
Example 2
23:23
Example 3
26:45
AP Practice for Gases

25m 34s

Intro
0:00
Example 1
0:34
Example 1
0:35
Example 2
6:15
Example 2: Part A
6:16
Example 2: Part B
8:46
Example 2: Part C
10:30
Example 2: Part D
11:15
Example 2: Part E
12:20
Example 2: Part F
13:22
Example 3
14:45
Example 3
14:46
Example 4
18:16
Example 4
18:17
Example 5
21:04
Example 5
21:05
IV. Thermochemistry
Energy, Heat, and Work

37m 32s

Intro
0:00
Thermochemistry
0:25
Temperature and Heat
0:26
Work
3:07
System, Surroundings, Exothermic Process, and Endothermic Process
8:19
Work & Gas: Expansion and Compression
16:30
Example 1
24:41
Example 2
27:47
Example 3
31:58
Enthalpy & Hess's Law

32m 34s

Intro
0:00
Thermochemistry
1:43
Defining Enthalpy & Hess's Law
1:44
Example 1
6:48
State Function
13:11
Example 2
17:15
Example 3
24:09
Standard Enthalpies of Formation

23m 9s

Intro
0:00
Thermochemistry
1:04
Standard Enthalpy of Formation: Definition & Equation
1:05
∆H of Formation
10:00
Example 1
11:22
Example 2
19:00
Calorimetry

39m 28s

Intro
0:00
Thermochemistry
0:21
Heat Capacity
0:22
Molar Heat Capacity
4:44
Constant Pressure Calorimetry
5:50
Example 1
12:24
Constant Volume Calorimetry
21:54
Example 2
24:40
Example 3
31:03
V. Kinetics
Reaction Rates and Rate Laws

36m 24s

Intro
0:00
Kinetics
2:18
Rate: 2 NO₂ (g) → 2NO (g) + O₂ (g)
2:19
Reaction Rates Graph
7:25
Time Interval & Average Rate
13:13
Instantaneous Rate
15:13
Rate of Reaction is Proportional to Some Power of the Reactant Concentrations
23:49
Example 1
27:19
Method of Initial Rates

30m 48s

Intro
0:00
Kinetics
0:33
Rate
0:34
Idea
2:24
Example 1: NH₄⁺ + NO₂⁻ → NO₂ (g) + 2 H₂O
5:36
Example 2: BrO₃⁻ + 5 Br⁻ + 6 H⁺ → 3 Br₂ + 3 H₂O
19:29
Integrated Rate Law & Reaction Half-Life

32m 17s

Intro
0:00
Kinetics
0:52
Integrated Rate Law
0:53
Example 1
6:26
Example 2
15:19
Half-life of a Reaction
20:40
Example 3: Part A
25:41
Example 3: Part B
28:01
Second Order & Zero-Order Rate Laws

26m 40s

Intro
0:00
Kinetics
0:22
Second Order
0:23
Example 1
6:08
Zero-Order
16:36
Summary for the Kinetics Associated with the Reaction
21:27
Activation Energy & Arrhenius Equation

40m 59s

Intro
0:00
Kinetics
0:53
Rate Constant
0:54
Collision Model
2:45
Activation Energy
5:11
Arrhenius Proposed
9:54
2 Requirements for a Successful Reaction
15:39
Rate Constant
17:53
Arrhenius Equation
19:51
Example 1
25:00
Activation Energy & the Values of K
32:12
Example 2
36:46
AP Practice for Kinetics

29m 8s

Intro
0:00
Kinetics
0:43
Example 1
0:44
Example 2
6:53
Example 3
8:58
Example 4
11:36
Example 5
16:36
Example 6: Part A
21:00
Example 6: Part B
25:09
VI. Equilibrium
Equilibrium, Part 1

46m

Intro
0:00
Equilibrium
1:32
Introduction to Equilibrium
1:33
Equilibrium Rules
14:00
Example 1: Part A
16:46
Example 1: Part B
18:48
Example 1: Part C
22:13
Example 1: Part D
24:55
Example 2: Part A
27:46
Example 2: Part B
31:22
Example 2: Part C
33:00
Reverse a Reaction
36:04
Example 3
37:24
Equilibrium, Part 2

40m 53s

Intro
0:00
Equilibrium
1:31
Equilibriums Involving Gases
1:32
General Equation
10:11
Example 1: Question
11:55
Example 1: Answer
13:43
Example 2: Question
19:08
Example 2: Answer
21:37
Example 3: Question
33:40
Example 3: Answer
35:24
Equilibrium: Reaction Quotient

45m 53s

Intro
0:00
Equilibrium
0:57
Reaction Quotient
0:58
If Q > K
5:37
If Q < K
6:52
If Q = K
7:45
Example 1: Part A
8:24
Example 1: Part B
13:11
Example 2: Question
20:04
Example 2: Answer
22:15
Example 3: Question
30:54
Example 3: Answer
32:52
Steps in Solving Equilibrium Problems
42:40
Equilibrium: Examples

31m 51s

Intro
0:00
Equilibrium
1:09
Example 1: Question
1:10
Example 1: Answer
4:15
Example 2: Question
13:04
Example 2: Answer
15:20
Example 3: Question
25:03
Example 3: Answer
26:32
Le Chatelier's principle & Equilibrium

40m 52s

Intro
0:00
Le Chatelier
1:05
Le Chatelier Principle
1:06
Concentration: Add 'x'
5:25
Concentration: Subtract 'x'
7:50
Example 1
9:44
Change in Pressure
12:53
Example 2
20:40
Temperature: Exothermic and Endothermic
24:33
Example 3
29:55
Example 4
35:30
VII. Acids & Bases
Acids and Bases

50m 11s

Intro
0:00
Acids and Bases
1:14
Bronsted-Lowry Acid-Base Model
1:28
Reaction of an Acid with Water
4:36
Acid Dissociation
10:51
Acid Strength
13:48
Example 1
21:22
Water as an Acid & a Base
25:25
Example 2: Part A
32:30
Example 2: Part B
34:47
Example 3: Part A
35:58
Example 3: Part B
39:33
pH Scale
41:12
Example 4
43:56
pH of Weak Acid Solutions

43m 52s

Intro
0:00
pH of Weak Acid Solutions
1:12
pH of Weak Acid Solutions
1:13
Example 1
6:26
Example 2
14:25
Example 3
24:23
Example 4
30:38
Percent Dissociation: Strong & Weak Bases

43m 4s

Intro
0:00
Bases
0:33
Percent Dissociation: Strong & Weak Bases
0:45
Example 1
6:23
Strong Base Dissociation
11:24
Example 2
13:02
Weak Acid and General Reaction
17:38
Example: NaOH → Na⁺ + OH⁻
20:30
Strong Base and Weak Base
23:49
Example 4
24:54
Example 5
33:51
Polyprotic Acids

35m 34s

Intro
0:00
Polyprotic Acids
1:04
Acids Dissociation
1:05
Example 1
4:51
Example 2
17:30
Example 3
31:11
Salts and Their Acid-Base Properties

41m 14s

Intro
0:00
Salts and Their Acid-Base Properties
0:11
Salts and Their Acid-Base Properties
0:15
Example 1
7:58
Example 2
14:00
Metal Ion and Acidic Solution
22:00
Example 3
28:35
NH₄F → NH₄⁺ + F⁻
34:05
Example 4
38:03
Common Ion Effect & Buffers

41m 58s

Intro
0:00
Common Ion Effect & Buffers
1:16
Covalent Oxides Produce Acidic Solutions in Water
1:36
Ionic Oxides Produce Basic Solutions in Water
4:15
Practice Example 1
6:10
Practice Example 2
9:00
Definition
12:27
Example 1: Part A
16:49
Example 1: Part B
19:54
Buffer Solution
25:10
Example of Some Buffers: HF and NaF
30:02
Example of Some Buffers: Acetic Acid & Potassium Acetate
31:34
Example of Some Buffers: CH₃NH₂ & CH₃NH₃Cl
33:54
Example 2: Buffer Solution
36:36
Buffer

32m 24s

Intro
0:00
Buffers
1:20
Buffer Solution
1:21
Adding Base
5:03
Adding Acid
7:14
Example 1: Question
9:48
Example 1: Recall
12:08
Example 1: Major Species Upon Addition of NaOH
16:10
Example 1: Equilibrium, ICE Chart, and Final Calculation
24:33
Example 1: Comparison
29:19
Buffers, Part II

40m 6s

Intro
0:00
Buffers
1:27
Example 1: Question
1:32
Example 1: ICE Chart
3:15
Example 1: Major Species Upon Addition of OH⁻, But Before Rxn
7:23
Example 1: Equilibrium, ICE Chart, and Final Calculation
12:51
Summary
17:21
Another Look at Buffering & the Henderson-Hasselbalch equation
19:00
Example 2
27:08
Example 3
32:01
Buffers, Part III

38m 43s

Intro
0:00
Buffers
0:25
Buffer Capacity Part 1
0:26
Example 1
4:10
Buffer Capacity Part 2
19:29
Example 2
25:12
Example 3
32:02
Titrations: Strong Acid and Strong Base

42m 42s

Intro
0:00
Titrations: Strong Acid and Strong Base
1:11
Definition of Titration
1:12
Sample Problem
3:33
Definition of Titration Curve or pH Curve
9:46
Scenario 1: Strong Acid- Strong Base Titration
11:00
Question
11:01
Part 1: No NaOH is Added
14:00
Part 2: 10.0 mL of NaOH is Added
15:50
Part 3: Another 10.0 mL of NaOH & 20.0 mL of NaOH are Added
22:19
Part 4: 50.0 mL of NaOH is Added
26:46
Part 5: 100.0 mL (Total) of NaOH is Added
27:26
Part 6: 150.0 mL (Total) of NaOH is Added
32:06
Part 7: 200.0 mL of NaOH is Added
35:07
Titrations Curve for Strong Acid and Strong Base
35:43
Titrations: Weak Acid and Strong Base

42m 3s

Intro
0:00
Titrations: Weak Acid and Strong Base
0:43
Question
0:44
Part 1: No NaOH is Added
1:54
Part 2: 10.0 mL of NaOH is Added
5:17
Part 3: 25.0 mL of NaOH is Added
14:01
Part 4: 40.0 mL of NaOH is Added
21:55
Part 5: 50.0 mL (Total) of NaOH is Added
22:25
Part 6: 60.0 mL (Total) of NaOH is Added
31:36
Part 7: 75.0 mL (Total) of NaOH is Added
35:44
Titration Curve
36:09
Titration Examples & Acid-Base Indicators

52m 3s

Intro
0:00
Examples and Indicators
0:25
Example 1: Question
0:26
Example 1: Solution
2:03
Example 2: Question
12:33
Example 2: Solution
14:52
Example 3: Question
23:45
Example 3: Solution
25:09
Acid/Base Indicator Overview
34:45
Acid/Base Indicator Example
37:40
Acid/Base Indicator General Result
47:11
Choosing Acid/Base Indicator
49:12
VIII. Solubility
Solubility Equilibria

36m 25s

Intro
0:00
Solubility Equilibria
0:48
Solubility Equilibria Overview
0:49
Solubility Product Constant
4:24
Definition of Solubility
9:10
Definition of Solubility Product
11:28
Example 1
14:09
Example 2
20:19
Example 3
27:30
Relative Solubilities
31:04
Solubility Equilibria, Part II

42m 6s

Intro
0:00
Solubility Equilibria
0:46
Common Ion Effect
0:47
Example 1
3:14
pH & Solubility
13:00
Example of pH & Solubility
15:25
Example 2
23:06
Precipitation & Definition of the Ion Product
26:48
If Q > Ksp
29:31
If Q < Ksp
30:27
Example 3
32:58
Solubility Equilibria, Part III

43m 9s

Intro
0:00
Solubility Equilibria
0:55
Example 1: Question
0:56
Example 1: Step 1 - Check to See if Anything Precipitates
2:52
Example 1: Step 2 - Stoichiometry
10:47
Example 1: Step 3 - Equilibrium
16:34
Example 2: Selective Precipitation (Question)
21:02
Example 2: Solution
23:41
Classical Qualitative Analysis
29:44
Groups: 1-5
38:44
IX. Complex Ions
Complex Ion Equilibria

43m 38s

Intro
0:00
Complex Ion Equilibria
0:32
Complex Ion
0:34
Ligan Examples
1:51
Ligand Definition
3:12
Coordination
6:28
Example 1
8:08
Example 2
19:13
Complex Ions & Solubility

31m 30s

Intro
0:00
Complex Ions and Solubility
0:23
Recall: Classical Qualitative Analysis
0:24
Example 1
6:10
Example 2
16:16
Dissolving a Water-Insoluble Ionic Compound: Method 1
23:38
Dissolving a Water-Insoluble Ionic Compound: Method 2
28:13
X. Chemical Thermodynamics
Spontaneity, Entropy, & Free Energy, Part I

56m 28s

Intro
0:00
Spontaneity, Entropy, Free Energy
2:25
Energy Overview
2:26
Equation: ∆E = q + w
4:30
State Function/ State Property
8:35
Equation: w = -P∆V
12:00
Enthalpy: H = E + PV
14:50
Enthalpy is a State Property
17:33
Exothermic and Endothermic Reactions
19:20
First Law of Thermodynamic
22:28
Entropy
25:48
Spontaneous Process
33:53
Second Law of Thermodynamic
36:51
More on Entropy
42:23
Example
43:55
Spontaneity, Entropy, & Free Energy, Part II

39m 55s

Intro
0:00
Spontaneity, Entropy, Free Energy
1:30
∆S of Universe = ∆S of System + ∆S of Surrounding
1:31
Convention
3:32
Examining a System
5:36
Thermodynamic Property: Sign of ∆S
16:52
Thermodynamic Property: Magnitude of ∆S
18:45
Deriving Equation: ∆S of Surrounding = -∆H / T
20:25
Example 1
25:51
Free Energy Equations
29:22
Spontaneity, Entropy, & Free Energy, Part III

30m 10s

Intro
0:00
Spontaneity, Entropy, Free Energy
0:11
Example 1
2:38
Key Concept of Example 1
14:06
Example 2
15:56
Units for ∆H, ∆G, and S
20:56
∆S of Surrounding & ∆S of System
22:00
Reaction Example
24:17
Example 3
26:52
Spontaneity, Entropy, & Free Energy, Part IV

30m 7s

Intro
0:00
Spontaneity, Entropy, Free Energy
0:29
Standard Free Energy of Formation
0:58
Example 1
4:34
Reaction Under Non-standard Conditions
13:23
Example 2
16:26
∆G = Negative
22:12
∆G = 0
24:38
Diagram Example of ∆G
26:43
Spontaneity, Entropy, & Free Energy, Part V

44m 56s

Intro
0:00
Spontaneity, Entropy, Free Energy
0:56
Equations: ∆G of Reaction, ∆G°, and K
0:57
Example 1: Question
6:50
Example 1: Part A
9:49
Example 1: Part B
15:28
Example 2
17:33
Example 3
23:31
lnK = (- ∆H° ÷ R) ( 1 ÷ T) + ( ∆S° ÷ R)
31:36
Maximum Work
35:57
XI. Electrochemistry
Oxidation-Reduction & Balancing

39m 23s

Intro
0:00
Oxidation-Reduction and Balancing
2:06
Definition of Electrochemistry
2:07
Oxidation and Reduction Review
3:05
Example 1: Assigning Oxidation State
10:15
Example 2: Is the Following a Redox Reaction?
18:06
Example 3: Step 1 - Write the Oxidation & Reduction Half Reactions
22:46
Example 3: Step 2 - Balance the Reaction
26:44
Example 3: Step 3 - Multiply
30:11
Example 3: Step 4 - Add
32:07
Example 3: Step 5 - Check
33:29
Galvanic Cells

43m 9s

Intro
0:00
Galvanic Cells
0:39
Example 1: Balance the Following Under Basic Conditions
0:40
Example 1: Steps to Balance Reaction Under Basic Conditions
3:25
Example 1: Solution
5:23
Example 2: Balance the Following Reaction
13:56
Galvanic Cells
18:15
Example 3: Galvanic Cells
28:19
Example 4: Galvanic Cells
35:12
Cell Potential

48m 41s

Intro
0:00
Cell Potential
2:08
Definition of Cell Potential
2:17
Symbol and Unit
5:50
Standard Reduction Potential
10:16
Example Figure 1
13:08
Example Figure 2
19:00
All Reduction Potentials are Written as Reduction
23:10
Cell Potential: Important Fact 1
26:49
Cell Potential: Important Fact 2
27:32
Cell Potential: Important Fact 3
28:54
Cell Potential: Important Fact 4
30:05
Example Problem 1
32:29
Example Problem 2
38:38
Potential, Work, & Free Energy

41m 23s

Intro
0:00
Potential, Work, Free Energy
0:42
Descriptions of Galvanic Cell
0:43
Line Notation
5:33
Example 1
6:26
Example 2
11:15
Example 3
15:18
Equation: Volt
22:20
Equations: Cell Potential, Work, and Charge
28:30
Maximum Cell Potential is Related to the Free Energy of the Cell Reaction
35:09
Example 4
37:42
Cell Potential & Concentration

34m 19s

Intro
0:00
Cell Potential & Concentration
0:29
Example 1: Question
0:30
Example 1: Nernst Equation
4:43
Example 1: Solution
7:01
Cell Potential & Concentration
11:27
Example 2
16:38
Manipulating the Nernst Equation
25:15
Example 3
28:43
Electrolysis

33m 21s

Intro
0:00
Electrolysis
3:16
Electrolysis: Part 1
3:17
Electrolysis: Part 2
5:25
Galvanic Cell Example
7:13
Nickel Cadmium Battery
12:18
Ampere
16:00
Example 1
20:47
Example 2
25:47
XII. Light
Light

44m 45s

Intro
0:00
Light
2:14
Introduction to Light
2:15
Frequency, Speed, and Wavelength of Waves
3:58
Units and Equations
7:37
Electromagnetic Spectrum
12:13
Example 1: Calculate the Frequency
17:41
E = hν
21:30
Example 2: Increment of Energy
25:12
Photon Energy of Light
28:56
Wave and Particle
31:46
Example 3: Wavelength of an Electron
34:46
XIII. Quantum Mechanics
Quantum Mechanics & Electron Orbitals

54m

Intro
0:00
Quantum Mechanics & Electron Orbitals
0:51
Quantum Mechanics & Electron Orbitals Overview
0:52
Electron Orbital and Energy Levels for the Hydrogen Atom
8:47
Example 1
13:41
Quantum Mechanics: Schrodinger Equation
19:19
Quantum Numbers Overview
31:10
Principal Quantum Numbers
33:28
Angular Momentum Numbers
34:55
Magnetic Quantum Numbers
36:35
Spin Quantum Numbers
37:46
Primary Level, Sublevels, and Sub-Sub-Levels
39:42
Example
42:17
Orbital & Quantum Numbers
49:32
Electron Configurations & Diagrams

34m 4s

Intro
0:00
Electron Configurations & Diagrams
1:08
Electronic Structure of Ground State Atom
1:09
Order of Electron Filling
3:50
Electron Configurations & Diagrams: H
8:41
Electron Configurations & Diagrams: He
9:12
Electron Configurations & Diagrams: Li
9:47
Electron Configurations & Diagrams: Be
11:17
Electron Configurations & Diagrams: B
12:05
Electron Configurations & Diagrams: C
13:03
Electron Configurations & Diagrams: N
14:55
Electron Configurations & Diagrams: O
15:24
Electron Configurations & Diagrams: F
16:25
Electron Configurations & Diagrams: Ne
17:00
Electron Configurations & Diagrams: S
18:08
Electron Configurations & Diagrams: Fe
20:08
Introduction to Valence Electrons
23:04
Valence Electrons of Oxygen
23:44
Valence Electrons of Iron
24:02
Valence Electrons of Arsenic
24:30
Valence Electrons: Exceptions
25:36
The Periodic Table
27:52
XIV. Intermolecular Forces
Vapor Pressure & Changes of State

52m 43s

Intro
0:00
Vapor Pressure and Changes of State
2:26
Intermolecular Forces Overview
2:27
Hydrogen Bonding
5:23
Heat of Vaporization
9:58
Vapor Pressure: Definition and Example
11:04
Vapor Pressures is Mostly a Function of Intermolecular Forces
17:41
Vapor Pressure Increases with Temperature
20:52
Vapor Pressure vs. Temperature: Graph and Equation
22:55
Clausius-Clapeyron Equation
31:55
Example 1
32:13
Heating Curve
35:40
Heat of Fusion
41:31
Example 2
43:45
Phase Diagrams & Solutions

31m 17s

Intro
0:00
Phase Diagrams and Solutions
0:22
Definition of a Phase Diagram
0:50
Phase Diagram Part 1: H₂O
1:54
Phase Diagram Part 2: CO₂
9:59
Solutions: Solute & Solvent
16:12
Ways of Discussing Solution Composition: Mass Percent or Weight Percent
18:46
Ways of Discussing Solution Composition: Molarity
20:07
Ways of Discussing Solution Composition: Mole Fraction
20:48
Ways of Discussing Solution Composition: Molality
21:41
Example 1: Question
22:06
Example 1: Mass Percent
24:32
Example 1: Molarity
25:53
Example 1: Mole Fraction
28:09
Example 1: Molality
29:36
Vapor Pressure of Solutions

37m 23s

Intro
0:00
Vapor Pressure of Solutions
2:07
Vapor Pressure & Raoult's Law
2:08
Example 1
5:21
When Ionic Compounds Dissolve
10:51
Example 2
12:38
Non-Ideal Solutions
17:42
Negative Deviation
24:23
Positive Deviation
29:19
Example 3
31:40
Colligatives Properties

34m 11s

Intro
0:00
Colligative Properties
1:07
Boiling Point Elevation
1:08
Example 1: Question
5:19
Example 1: Solution
6:52
Freezing Point Depression
12:01
Example 2: Question
14:46
Example 2: Solution
16:34
Osmotic Pressure
20:20
Example 3: Question
28:00
Example 3: Solution
30:16
XV. Bonding
Bonding & Lewis Structure

48m 39s

Intro
0:00
Bonding & Lewis Structure
2:23
Covalent Bond
2:24
Single Bond, Double Bond, and Triple Bond
4:11
Bond Length & Intermolecular Distance
5:51
Definition of Electronegativity
8:42
Bond Polarity
11:48
Bond Energy
20:04
Example 1
24:31
Definition of Lewis Structure
31:54
Steps in Forming a Lewis Structure
33:26
Lewis Structure Example: H₂
36:53
Lewis Structure Example: CH₄
37:33
Lewis Structure Example: NO⁺
38:43
Lewis Structure Example: PCl₅
41:12
Lewis Structure Example: ICl₄⁻
43:05
Lewis Structure Example: BeCl₂
45:07
Resonance & Formal Charge

36m 59s

Intro
0:00
Resonance and Formal Charge
0:09
Resonance Structures of NO₃⁻
0:25
Resonance Structures of NO₂⁻
12:28
Resonance Structures of HCO₂⁻
16:28
Formal Charge
19:40
Formal Charge Example: SO₄²⁻
21:32
Formal Charge Example: CO₂
31:33
Formal Charge Example: HCN
32:44
Formal Charge Example: CN⁻
33:34
Formal Charge Example: 0₃
34:43
Shapes of Molecules

41m 21s

Intro
0:00
Shapes of Molecules
0:35
VSEPR
0:36
Steps in Determining Shapes of Molecules
6:18
Linear
11:38
Trigonal Planar
11:55
Tetrahedral
12:45
Trigonal Bipyramidal
13:23
Octahedral
14:29
Table: Shapes of Molecules
15:40
Example: CO₂
21:11
Example: NO₃⁻
24:01
Example: H₂O
27:00
Example: NH₃
29:48
Example: PCl₃⁻
32:18
Example: IF₄⁺
34:38
Example: KrF₄
37:57
Hybrid Orbitals

40m 17s

Intro
0:00
Hybrid Orbitals
0:13
Introduction to Hybrid Orbitals
0:14
Electron Orbitals for CH₄
5:02
sp³ Hybridization
10:52
Example: sp³ Hybridization
12:06
sp² Hybridization
14:21
Example: sp² Hybridization
16:11
σ Bond
19:10
π Bond
20:07
sp Hybridization & Example
22:00
dsp³ Hybridization & Example
27:36
d²sp³ Hybridization & Example
30:36
Example: Predict the Hybridization and Describe the Molecular Geometry of CO
32:31
Example: Predict the Hybridization and Describe the Molecular Geometry of BF₄⁻
35:17
Example: Predict the Hybridization and Describe the Molecular Geometry of XeF₂
37:09
XVI. AP Practice Exam
AP Practice Exam: Multiple Choice, Part I

52m 34s

Intro
0:00
Multiple Choice
1:21
Multiple Choice 1
1:22
Multiple Choice 2
2:23
Multiple Choice 3
3:38
Multiple Choice 4
4:34
Multiple Choice 5
5:16
Multiple Choice 6
5:41
Multiple Choice 7
6:20
Multiple Choice 8
7:03
Multiple Choice 9
7:31
Multiple Choice 10
9:03
Multiple Choice 11
11:52
Multiple Choice 12
13:16
Multiple Choice 13
13:56
Multiple Choice 14
14:52
Multiple Choice 15
15:43
Multiple Choice 16
16:20
Multiple Choice 17
16:55
Multiple Choice 18
17:22
Multiple Choice 19
18:59
Multiple Choice 20
20:24
Multiple Choice 21
22:20
Multiple Choice 22
23:29
Multiple Choice 23
24:30
Multiple Choice 24
25:24
Multiple Choice 25
26:21
Multiple Choice 26
29:06
Multiple Choice 27
30:42
Multiple Choice 28
33:28
Multiple Choice 29
34:38
Multiple Choice 30
35:37
Multiple Choice 31
37:31
Multiple Choice 32
38:28
Multiple Choice 33
39:50
Multiple Choice 34
42:57
Multiple Choice 35
44:18
Multiple Choice 36
45:52
Multiple Choice 37
48:02
Multiple Choice 38
49:25
Multiple Choice 39
49:43
Multiple Choice 40
50:16
Multiple Choice 41
50:49
AP Practice Exam: Multiple Choice, Part II

32m 15s

Intro
0:00
Multiple Choice
0:12
Multiple Choice 42
0:13
Multiple Choice 43
0:33
Multiple Choice 44
1:16
Multiple Choice 45
2:36
Multiple Choice 46
5:22
Multiple Choice 47
6:35
Multiple Choice 48
8:02
Multiple Choice 49
10:05
Multiple Choice 50
10:26
Multiple Choice 51
11:07
Multiple Choice 52
12:01
Multiple Choice 53
12:55
Multiple Choice 54
16:12
Multiple Choice 55
18:11
Multiple Choice 56
19:45
Multiple Choice 57
20:15
Multiple Choice 58
23:28
Multiple Choice 59
24:27
Multiple Choice 60
26:45
Multiple Choice 61
29:15
AP Practice Exam: Multiple Choice, Part III

32m 50s

Intro
0:00
Multiple Choice
0:16
Multiple Choice 62
0:17
Multiple Choice 63
1:57
Multiple Choice 64
6:16
Multiple Choice 65
8:05
Multiple Choice 66
9:18
Multiple Choice 67
10:38
Multiple Choice 68
12:51
Multiple Choice 69
14:32
Multiple Choice 70
17:35
Multiple Choice 71
22:44
Multiple Choice 72
24:27
Multiple Choice 73
27:46
Multiple Choice 74
29:39
Multiple Choice 75
30:23
AP Practice Exam: Free response Part I

47m 22s

Intro
0:00
Free Response
0:15
Free Response 1: Part A
0:16
Free Response 1: Part B
4:15
Free Response 1: Part C
5:47
Free Response 1: Part D
9:20
Free Response 1: Part E. i
10:58
Free Response 1: Part E. ii
16:45
Free Response 1: Part E. iii
26:03
Free Response 2: Part A. i
31:01
Free Response 2: Part A. ii
33:38
Free Response 2: Part A. iii
35:20
Free Response 2: Part B. i
37:38
Free Response 2: Part B. ii
39:30
Free Response 2: Part B. iii
44:44
AP Practice Exam: Free Response Part II

43m 5s

Intro
0:00
Free Response
0:12
Free Response 3: Part A
0:13
Free Response 3: Part B
6:25
Free Response 3: Part C. i
11:33
Free Response 3: Part C. ii
12:02
Free Response 3: Part D
14:30
Free Response 4: Part A
21:03
Free Response 4: Part B
22:59
Free Response 4: Part C
24:33
Free Response 4: Part D
27:22
Free Response 4: Part E
28:43
Free Response 4: Part F
29:35
Free Response 4: Part G
30:15
Free Response 4: Part H
30:48
Free Response 5: Diagram
32:00
Free Response 5: Part A
34:14
Free Response 5: Part B
36:07
Free Response 5: Part C
37:45
Free Response 5: Part D
39:00
Free Response 5: Part E
40:26
AP Practice Exam: Free Response Part III

28m 36s

Intro
0:00
Free Response
0:43
Free Response 6: Part A. i
0:44
Free Response 6: Part A. ii
3:08
Free Response 6: Part A. iii
5:02
Free Response 6: Part B. i
7:11
Free Response 6: Part B. ii
9:40
Free Response 7: Part A
11:14
Free Response 7: Part B
13:45
Free Response 7: Part C
15:43
Free Response 7: Part D
16:54
Free Response 8: Part A. i
19:15
Free Response 8: Part A. ii
21:16
Free Response 8: Part B. i
23:51
Free Response 8: Part B. ii
25:07
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Lecture Comments (12)

1 answer

Last reply by: Professor Hovasapian
Tue Aug 2, 2016 2:53 AM

Post by tae Sin on July 29, 2016

Hi, Mr. Hovasapian.

At the part about the vapor pressure, you were talking about vapor pressure with test tubes and the water. Should the gas react to the air when you insert the tube into the tube underwater? And does the air causes the gas to change the pressure of the test tube underwater?

I like the lesson. Thanks!

1 answer

Last reply by: Professor Hovasapian
Sat Nov 15, 2014 10:38 PM

Post by Datevig Daghlian on November 13, 2014

Dear Professor Hovasapian,

   Thank you very much for your lecture! I had never seen the Ideal Gas law from that perspective. Thank you and God bless!

Thank You,
George Daghlian

3 answers

Last reply by: Professor Hovasapian
Tue Jan 28, 2014 3:04 AM

Post by Deborah Lee on January 27, 2014

Hi Professor Hovasapian,

I hope you are well.

We refer to Nitrogen as the gas (N_2) when calculating partial pressure, mol and mass. Is there a reason why you refer to Nitrogen as N when calculating Mass % ?

Thank you for your lectures, I find them extremely helpful!

1 answer

Last reply by: Professor Hovasapian
Wed Aug 14, 2013 2:44 PM

Post by Mark Andrews on August 14, 2013

Back when you introduced the Mol Fractions and said you'd skip over the maths, I was wondering if you would be able to do a video on that because while I was able to replicate your example with the numbers you provided, I kind of got lost with how you got to the 616.2.

I realise that 640-23.8 = 616.2

My thought was to approach the problem by converting 640 Torr as the P total
Then converting 23.8 Torr and subtracting that from the "total pressure" then using pV=nRT

When I followed that procedure I had P Total = 0.842
H2O Pressure worked out at 0.0313 which gave me the 0.8107 = 616.2

It has taken me a while to work out I ended up with the same answer but was that the mathematical process you didn't go into?

By the way it was after watching your FREE lectures that I decided to purchase a prescription.  I think you're a very good teacher and far easier to understand than my current lecturer at Uni who has lost me in the subject about 4 weeks ago. Thank you for way you conduct your lessons.



1 answer

Last reply by: Professor Hovasapian
Wed Aug 7, 2013 1:26 AM

Post by Charles Zhou on August 6, 2013

For example 1, when I was dealing with the moles of oxygen, I use n=PV/RT. But I think the volume of 0.550L is the total volume, we need to subtract the Vapor volume by using the molar Fraction: n()2)/n(vapor)=P(O2)/P(vapor). So I got the n(O2) equals to 0.0181. It's really close to 0.0182. And the mass of KClO3 is 1.49. I don't know if I'm right. Thanks for your help and I really enjoy your class.

Related Articles:

Partial Pressure, Mol Fraction, & Vapor Pressure

  • The total pressure in a sample containing a mixture of gases is the sum of the individual pressures of each gas.
  • Mol Fraction is like any other fraction: part over the whole. Mol Fraction of A is the mols of A divided by the total number of mols in the sample.
  • Given any liquid, some molecules always escape into the gas phase. In a closed container, these gas-phase particles exert a pressure. This is called the Vapor Pressure of the liquid at a given temperature.
  • Nitrogen content of an Organic compound can be determined by the Dumas Method. The compound is passed over hot Copper (2) Oxide and reacts.
  • The product gases are passed through a solution of Potassium Hydroxide to remove the Carbon Dioxide. The remaining gas is Nitrogen saturated with water vapor. In a given experiment 0.225 g of the unknown compound produces 27.8 mL of Nitrogen saturated with water vapor at 25 degrees Celsius and 730 torr. What is the Mass % Nitrogen in the compound? (Vapor pressure of water at this temp is 23.8 torr.

Partial Pressure, Mol Fraction, & Vapor Pressure

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
  • Gases 0:27
    • Gases
    • Mole Fractions
    • Vapor Pressure
    • Example 1
    • Example 2

Transcription: Partial Pressure, Mol Fraction, & Vapor Pressure

Hello, and welcome back to Educator.com; welcome back to AP Chemistry.0000

In the last lesson, we introduced gas laws; we introduced the ideal gas law; and we talked a fair amount about pressure.0004

Today, we're going to continue on, and we're going to talk about mixtures of gases and something called partial pressure.0011

It's going to use exactly the ideal gas law--again, very, very simple mathematically--so let's just go ahead and jump in and see what we can do!0017

OK, it's very, very simple: let's just say that, if I have a mixture of gases--let's just say 3 gases--each one has a certain pressure.0028

Let's say, if I keep them in three separate containers--there is a pressure; there is a pressure; there is a pressure.0045

If I dump them all into one container, each one is contributing to the total pressure.0048

Well, it's exactly what you think: the total pressure of the system is equal to the partial pressure of 1, plus the partial pressure of 2, plus the partial pressure of the third.0053

Now, when we say partial, that means that it is part of a mixture--that is where the "partial" comes from.0065

It's not as if you are taking part of the pressure of the gas.0071

So, gas 1 contributes to the pressure; gas 2 contributes to the pressure; gas 3 contributes to the pressure; because now, you have a mixture of them; there is a total pressure that we can measure, but we can actually divide it up.0074

Each one actually contributes some kind of pressure to it.0089

That is all that is going on here; so we speak of these individual pressures of the individual gases as the partial pressure of that gas.0091

So, if I have a mixture of...well, for example, air: air is a mixture of nitrogen and oxygen gas; if I take a sample of it in a certain volume, the partial pressure of O2 is going to be a certain something, and the partial pressure of N2 is going to be something else.0100

That is all that is going on here--nothing strange.0115

Let's go ahead and just deal with some of the mathematics, because I think if you see the mathematics, where it comes from, it will make a little bit more sense as to what it is that is going on.0119

We, again, start off with our ideal gas law; so let me put that over here: PV=nRT: pressure times the volume equals the number of moles times the gas constant times the temperature in Kelvin (volume in liters, pressure in atmospheres).0129

For each gas in the container, for a given volume and a given temperature, we have the following.0144

The partial pressure of 1 is equal to the number of moles of 1, times RT, over V.0173

R is constant: given volume, given temperature--those are constant, so when we rearrange this equation, for each gas we can use that equation.0179

So, the number of moles of 1, times R, times T, divided by the volume, gives me the pressure of the one gas.0187

The partial pressure of 2 is equal to the number of moles of 2, times RT, over V: just a rearrangement of the ideal gas law.0197

And partial pressure of 3: exactly what you think: it's n3RT over V.0205

Well, we have our basic relationship; if we just add the individual pressures, it is going to give us our total pressure.0211

So, total pressure is equal to P1 + P2 + P3; now, I'll just substitute these values in for this.0217

I end up with n1RT/V +n2RT/V, + n3RT/V.0228

Well, if you notice--RT/V--you can factor it out; so let me pull that out; RT/V times n1 +n2 + n3.0241

So, the total pressure is equal to RT/V times the total number of moles.0255

The total number of moles comes from the moles of 1, plus the moles of 2, plus the moles of the other.0269

As you see, the number of moles and pressure are related.0274

Pressure is just a different expression--a different way of representing the number of moles of something.0279

That is what is nice about pressure; it is just the other side of the coin, if you will, of an amount.0285

So, we speak of the partial pressure; we speak of the total number of moles; the total number of moles is related to the partial pressure through this proportionality constant.0291

That is all that is going on.0303

That is what this right here is essentially saying: it's saying that the number of moles is related to the pressure through a proportionality constant.0305

Number of moles and pressure are just two sides of the same coin--amount; that is what it is representing--amount.0313

Moles is the standard unit by which we measure things in chemistry; well, when you are dealing with gases, pressure is more convenient--to speak about pressure and moles.0324

Here is our conversion; that is all that is going on here.0333

OK, so the total number of moles of a gas is what is important, not the identity; that is another thing that this is telling me.0337

It is the total number of moles that matter of the gas; the identity of the gas is completely irrelevant.0345

Now, let's introduce something called a mole fraction: for our purposes, we're probably not going to be using it all that much, as far as the problems that we do for gases, but mole fraction--I want to introduce it here.0353

It is something that is very important; it will show up again when we discuss solutions, not too long from now.0365

So, it's exactly what you would expect it to be when we talk about a fraction: a fraction is just a part over the whole.0370

Well, a mole fraction, which is symbolized by the Greek letter chi...the mole fraction of, let's say, a certain...something, is the number of moles of that something over the total number of moles.0378

It's just the part over the whole; that is it.0394

If I have 10 moles total--let's say I have 8 moles of nitrogen mixed with two moles of oxygen--well, the mole fraction of oxygen is 2/10, and the mole fraction is 8/10; it's just the number of moles of the part, over the total number of moles floating around.0399

That is it; that is all mole fraction is.0418

So, as it turns out, I'll go ahead and skip the math, but in terms of pressure...which--you can do this yourself if you want to...0420

Again, you have n1; just expand this out with PV=nRT; just put the n1=PV/RT, and you will see that the PV's over the RT's--they all cancel, and what you end up with is...0431

You can also find the mole fraction of a gas by measuring the pressure of that gas, individual gas, over the total pressure of the system; that is it.0448

You can work in moles, or you can work in pressure, because again, pressure and moles are just different ways of representing the same thing: an amount.0457

Notice, mole fraction doesn't have any units.0469

OK, we can also rearrange this; we can rearrange this to: the partial pressure of some individual gas is equal to the mole fraction of the gas, times the total pressure.0473

That is just expressing a fraction; that is all that is going on here--straight math--nothing you haven't seen before.0491

I did want to introduce it to you, because we will be talking about it more and more as we go on.0498

Now, I want to talk about something called vapor pressure; I'm going to use vapor pressure to segue into our first problem.0505

Let's talk--I'm going to draw a little bit of a thing here; if I were to take, let's say, a test tube, and I have a little trough of, let's say, water, and if I invert the test tube this way--and let's say I have some sample of a gas, and put a little fire under there, and I have a little tube that is going into the water and up, up...0516

So now, some gas is escaping; it is traveling through the tube; it is going underneath the water (and this tube is actually filled with water, to the top, just like we did with the mercury--except, in this case, we are dealing with just H2O).0555

Well, the gas is going start bubbling up, right?--it's going to start bubbling, and it's going to make its way to the top.0570

When it makes its way to the top, the gas is actually going to start collecting at the top here.0578

As it does so, it's going to push the water level down; it's going to push--the gas is expanding; it's creating a volume up in the top container.0582

So, what happens is: the water level in there is going to drop as more gas collects.0592

Now, let's see: we have collected some gas, and let's say now the water level is there.0600

What is up here is the gas that we have collected.0605

When we collect a gas, using this method, over whatever liquid this happens to be (more often than not, it will be water; and for our purposes it will be water), something very interesting happens.0612

There is a certain amount of gas in here that is going to exert a pressure, because gas pressure--that is what is pushing the water level down, back into the trough of the water.0624

Well, something else interesting happens: any time I have some liquid in a closed container, some of the liquid particles at the surface escape and become gas particles--at any temperature (most standard temperatures).0635

So, in a closed container, if I have some water, it isn't just water-and-there-is-nothing-floating-around-on-top; as it turns out, water itself--some of the water molecules--will escape into the vapor phase, and so now, you have some water vapor on top of the liquid water.0654

That water vapor is a gas, and a gas exerts a pressure.0670

We call the gas that has escaped from the liquid phase of something, that is hovering above it in a container, the vapor pressure.0675

Let me give you a reasonable definition here.0685

We talk about the vapor pressure, and that is the pressure exerted by the gaseous molecules of a given liquid (which, in our case, is water) that have escaped the liquid phase.0690

That is it: if I have a container of a liquid, it isn't just an empty vacuum; as it turns out, some of the molecules actually escape, become gaseous, and--because they become gaseous--they actually exert a certain pressure on the walls of the container.0729

That pressure is what we call the vapor pressure; all liquids do this.0744

That is it; what this means for our purposes here is: we have collected a gas, but we have collected it over water.0748

Well, water, at different temperatures, has different vapor pressures, certainly; so here, we not only have gas particles, but we also have particles of water vapor.0757

This becomes a mixture of gases; and, because it is a mixture of gases, our Dalton's law of partial pressure applies.0769

The total pressure in here is going to be the pressure of the gas, plus the pressure of the water at that particular temperature.0776

There is a table of vapor pressure of water at different temperatures, and that is what we are going to use in our next example.0782

I just wanted to introduce this idea of vapor pressure; again, it will show up again and again, particularly when we talk about solutions.0789

We will discuss mole fraction in more detail, but I wanted you to see what was going on for the sake of this example.0800

Now, let's go ahead and do our Example #1.0806

This is a standard problem in chemistry, using gases, using stoichiometry, using...this is a typical problem that you will actually see on the AP exam in the free response section.0812

Solid KClO3, which is potassium chlorate, was heated and decomposed to form solid potassium chloride and oxygen gas.0827

Well, the O2 was collected over water with that thing that we just designed; so we heated up the potassium chlorate in the test tube; the oxygen gas was collected over water--pushed the water level down and collected in that little space, up above, at 25 degrees Celsius and a total pressure of 640 torr.0851

That means the total pressure in that space above is 640 torr; it consists of oxygen pressure and the water vapor pressure, which we're going to talk about next.0889

Now, the volume collected was 0.550 liters.0898

The vapor pressure of water at this temperature, at 25 degrees Celsius, is 23.8 torr; that means there is enough water vapor, floating around above the water, to create a pressure of 23.8 torricelli.0914

Our question is this: What is the partial pressure of O2 collected, and the mass of KClO3 decomposed?0937

Let's see what is going on here: if you refer back to the diagram, we burned the KClO3; we produced oxygen gas.0970

Oxygen gas traveled through the tube, went up into the tube, and pushed the water level down; now, oxygen gas is collecting in that space above in the vertical tube.0978

It is also mixed with some water vapor, because again, any time you collect something over a liquid, some of that liquid turns into vapor.0987

So, it is saying that the total pressure is 640 torr, while the vapor pressure at this particular temperature that we ran this experiment is 23.8 torr.0994

We're asking for "What is the partial pressure of the oxygen?" and "How much KClO3 was decomposed, based on the fact that I collected .550 liters?"1004

Wow, I collected .550 liters; how am I going to find out how much KClO3 was actually decomposed?1016

Well, let's start with an equation; again, we are doing chemistry.1022

We do KClO3; we heat it up; we end up with KCl + O2.1027

I see a three--no, that is ozone; no, it's oxygen gas that we collected, not ozone.1038

When I see a three and I see a two, I immediately do this: just 3, 6, 3, 2, 6, and then just put a 2 here to balance out the Cl.1043

So, 2 moles of potassium chlorate produce 2 moles of potassium chlorite, solid, plus 3 moles of oxygen gas.1056

We know that the total pressure is equal to the partial pressure of the two gases: water vapor and oxygen: the partial pressure of O2 plus the partial pressure of H2O.1066

Well, we know the total pressure: that is 640 torr; we're looking for the partial pressure of O2, and we are given the vapor pressure of water--it's 23.8 torr.1078

So, the first part of the problem is easy: the partial pressure of O2 collected is just the 640 minus the 23.8, which gives us 616.2 torricelli.1094

That's it: I just use the standard: the total pressure is equal to the sum of the partial pressures.1110

I have 2 gases; I have the partial pressure of water vapor at 25 degrees Celsius; so, my partial pressure of O2 is 616.2 torr.1115

Now, let's see if we can--the next thing we want to do is: we want to basically find the number of moles of oxygen.1125

From the moles of oxygen, we're going to use the reaction stoichiometry to find the number of moles of potassium chlorate.1133

From the moles of potassium chlorate, we're going to deduce the number of grams of potassium chlorate.1138

That is what we are doing here.1143

Now, let's go ahead and--again, we want to work in atmospheres, so let's convert this 616.2 torr; it's going to be times 1 atm=760 torricelli; so that gives us 0.811 atmospheres.1145

So, the partial pressure in atmospheres: the partial pressure of O2...0.811 atmospheres is the partial pressure of O2.1167

Now, we have PV=nRT; in this case, partial pressure of O2, times V, equals the number of moles of O2, times RT, because we're dealing only with the oxygen.1182

But, the equation applies.1202

We rearrange, because we want to find the number of moles, so the number of moles of O2 is equal to the partial pressure of O2, times the volume, divided by RT.1204

It is equal to 0.811; that is what we just found--that is the partial pressure--times 0.550 liters--that is the volume that it is actually encased in, the volume at the top of the tube; R is 0.08206, and we are at 298 Kelvin.1215

Well, when I do that, I end up with 0.0182 moles of oxygen; that is how many moles of oxygen I have.1241

Well, now, 0.0182 moles of O2 (the mole relationship is 2 moles of potassium chlorate produce 3 moles of O2).1250

Therefore, I know that this came from 0.0126 moles of KClO3.1266

Now, I take 0.01216--I'm sorry, I forgot a 1 there--moles of KClO3, times the molar mass of KClO3, which happens to be 122.6 grams per mole, for a grand total of 1.49 grams KClO3 decomposed.1276

This is no more than a standard stoichiometry problem, using mole ratios, except it's under gaseous conditions; therefore, I have to use partial pressures and ideal gas laws; that's all.1303

Moles--that is the nice thing about the ideal gas law: it contains moles, so we can use it in our stoichiometry.1317

Don't be fooled; this is a gas problem; it's not really a gas problem--this is a stoichiometry problem.1323

It's a standard chemistry problem--dealing with gases, therefore--for gases, we have one technique in our toolbox: PV=nRT.1329

Well, actually, we have two techniques; the other is total pressure is equal to the sum of the partial pressures of however many gases; that is it--those are the two things that we are using when we are dealing with gases.1338

Ideal gas law, Dalton's law of partial pressures: everything else is just classical stoichiometry.1351

OK, let's see what else we can do...let's move forward here, and let's deal with another example.1358

This example is going to be typical of the kind of example that you see in the AP exam: it's going to have a lot of wording; a lot is going on.1372

Don't that let that intimidate you.1380

A lot of times, half the problem will be just the description of what is happening, so that you have a sense of what is going on, so that you can deduce the chemistry, or deduce the math, from that.1383

Again, don't get intimidated.1394

Let's just start and see where we go.1395

The nitrogen content of an organic compound can be determined by something called the Dumas method.1398

The compound is passed over hot copper oxide, and reacts as follows (so I'm going to go ahead and write the reaction).1403

It is: you take the compound, you pass it over hot copper oxide, and you end up producing nitrogen gas; you end up producing CO2 gas, and you end up producing water vapor ("water gas").1410

Let's see: total pressure, CO2...OK.1430

Now, they say the product gases are passed through a solution of potassium hydroxide to remove the carbon dioxide.1438

You produce three gases, and then you pass it through a solution of potassium hydroxide; that binds the carbon dioxide and gets it out of the way.1444

The remaining gas is nitrogen, saturated with water vapor.1455

Nitrogen saturated with water vapor: that just means nitrogen and water vapor, in the same gas mixture; that is all that this means.1459

In a given experiment, 0.225 grams of the unknown compound produces 27.8 milliliters of nitrogen saturated with water vapor, at 25 degrees Celsius and 730 torr.1468

What is the mass percent of nitrogen in the compound?1483

The vapor pressure of water at this temperature is 23.8 torr.1486

What are they asking for?--they want to know what is the mass percent of nitrogen in this compound.1490

Well, we know mass percent...so let's just write out the equation so we don't get lost.1496

Mass percent equals the mass of nitrogen over the total mass, times 100; so that is all we want.1501

We have the total mass already: our compound is .225 grams--we already have one of the numbers.1516

All we need to do is find the mass of nitrogen.1522

Let's take a look and see how we are going to do this.1526

Again, mass of nitrogen--chances are, we are going to have to find the number of moles of nitrogen, and the number of moles is going to come from the ideal gas equation, so let's just work forward.1528

We are dealing with a mixture here: we have a mixture of water vapor and nitrogen, because the carbon dioxide has been removed by the potassium hydroxide solution.1539

Therefore, our total pressure of the system is equal to the partial pressure of the nitrogen gas, plus the partial pressure of the water vapor (I just wrote the nitrogen).1550

OK, well, let's see what we have.1564

In a given experiment, unknown compound, nitrogen saturated with water, 730 torr--that is our total pressure.1566

We have 730 torr; we need the partial pressure of nitrogen in order to find the number of moles of nitrogen, and they gave us the vapor pressure of water, which is 23.8 torr.1574

So again, we can find the partial pressure of nitrogen with a simple arithmetic problem: 730-23.8, and we end up with 702 (is that correct?--23, 24, 731)--yes, 702 torr.1588

Wait, is that correct?--730, minus...no, that is not correct; that is not 702 torr; let me see, let's do some quick arithmetic; arithmetic has never been my strong suit.1617

730, 23.8, 24, 25, 26, 27, 28, 29...it's going to be 706 (I think), point 2, torr--excellent.1631

Then, we're going to multiply that by...we're going to convert that to atmospheres, so 1 atmosphere over 760 torr...1649

OK, so I don't have a calculator at my disposal; my number was originally incorrect when I did this, but I'm going to use my original number.1664

I'm going to do it as if it were 702, but again, the division and the number...it's the process that is important, not the actual number here.1670

I'm going to use the original number that I got, which was 0.924 atmospheres.1680

That is the partial pressure of the nitrogen gas.1688

Well, PV=nRT; the partial pressure of the nitrogen gas, times the volume, equals the number of moles of nitrogen gas, times RT.1694

Rearrange: the number of moles of nitrogen gas equals the P of N2, times the V, over RT.1707

The pressure is 0.924 atmospheres, times...and we collected 27.8 milliliters, so we want to convert that to liters: .0278 liters; the gas constant is .08206, and our temperature is 25 degrees Celsius; which is 298 Kelvin.1716

Again, liter, Kelvin, atmosphere.1744

We end up with a number of moles that is equal to 0.00105 moles of nitrogen gas.1748

Well, 0.00105 moles of nitrogen gas, times 28 grams per mole (it's 28 because this is N2: nitrogen is 14, and 2 is 28)--we end up with 0.0294 grams of N2.1759

OK, mass percent equals 0.0294 grams, divided by the total number of grams, which was 0.225, in our compound; when we multiply by 100%: 13% nitrogen by mass.1785

There we go.1810

We converted the nitrogen in a compound to a gas--to nitrogen gas, to carbon dioxide gas, to water gas.1816

We took care of the carbon dioxide gas; therefore, now we have a mixture of nitrogen and water vapor.1827

We know, at a certain temperature, that the vapor pressure of water is a certain number; we can look that up in a chart, or we are given it in the problem.1834

We can find the partial pressure of nitrogen.1841

From the partial pressure of nitrogen, we use our one basic technique, which is the ideal gas law, to find the number of moles of nitrogen.1843

From the number of moles of nitrogen, we find the number of grams of nitrogen.1851

Take the number of grams of nitrogen; divide by the number of grams of the total compound; and that gives us a mass percent.1855

So, again, the problem is routine, in the sense that it is nothing you haven't done before.1861

You are doing the same problems over and over again.1867

But now, we are doing it in the context of gases, so the technique that we use is Dalton's law of partial pressures and the ideal gas law.1869

Everything is going to be a variation of that: if you sort of look at it that way globally, I think a lot of this will start to make more sense.1878

And again, we start with the chemistry; start with an equation; balance the equation; see if you can understand what is going on--let the problem wash over you.1886

There is nothing strange that is happening here; it is basic math.1896

It just needs to be arranged in a certain way.1899

OK, so that is Dalton's law of partial pressures, and a little bit of mole fraction, and things like that.1902

With that, I will say "Thank you" for joining us here at Educator.com, and we'll see you next time--goodbye!1908

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