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Raffi Hovasapian

Raffi Hovasapian

Phase Diagrams & Solutions

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 (4)

1 answer

Last reply by: Professor Hovasapian
Fri Mar 25, 2016 10:34 PM

Post by Tania Bore on March 21, 2016

How did you get the conversion factor from IPr to liters?

1 answer

Last reply by: Professor Hovasapian
Sat Jan 11, 2014 5:32 PM

Post by Angela Patrick on January 11, 2014

Does the phase diagram for water say that at any temperature past the triple point (say 10 degrees celsius) water is incapable of becoming a solid?

Phase Diagrams & Solutions

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
  • Phase Diagrams and Solutions 0:22
    • Definition of a Phase Diagram
    • Phase Diagram Part 1: H₂O
    • Phase Diagram Part 2: CO₂
    • Solutions: Solute & Solvent
    • Ways of Discussing Solution Composition: Mass Percent or Weight Percent
    • Ways of Discussing Solution Composition: Molarity
    • Ways of Discussing Solution Composition: Mole Fraction
    • Ways of Discussing Solution Composition: Molality
    • Example 1: Question
    • Example 1: Mass Percent
    • Example 1: Molarity
    • Example 1: Mole Fraction
    • Example 1: Molality

Transcription: Phase Diagrams & Solutions

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

Last time, we talked about heating curves and changes of state.0004

Today, we are going to introduce something called a phase diagram--a very, very important diagram in chemistry.0009

And we are going to also continue our discussion of solutions.0016

Let's just jump right on in with the notion of a phase diagram: we are going to do two phase diagrams today--we are going to give you the one for H2O, and we're going to give you the one for CO2.0019

As it turns out, each one of these is actually unusual in its own way; but because carbon dioxide and water are such ubiquitous substances--things that we use all the time--these are the ones that we want to concentrate on.0030

OK, so let's go ahead and...well, let me actually write down what a phase diagram is, and then I'll start out with the one for water.0045

Let's see: A phase diagram shows which state (in other words, solid, liquid, or gas) exists (not at which--I should say "at a given"--you know what, I really think I had better stop abbreviating things and just write my words out; OK) at a given temperature and pressure.0061

So, we have temperature on the x-axis, and we have pressure on the y-axis, and the different values tell you what state the particular substance is in.0106

Here is what it looks like; so we are going to do the one for H2O now--so let's go ahead and do that.0114

Actually, you know what, I'm going to need a little bit more room on the left, so let me go over here; so that is that...and something like that.0123

Once again, we have temperature (which is in degrees Celsius), and here we have pressure (it could be in torr; it could be in atmosphere; I'm going to go ahead and do atmospheres here, if I'm not mistaken; there we go).0134

OK, so now, here is what happens--what you get is something that looks like this.0153

Put this a little bit higher here...and I'm going to drop this one just a little bit.0169

OK, this is something...I don't know...a point...0175

This is the solid phase; this is the liquid phase; and this is the gas phase; so this is temperature, so at different temperatures and different pressures, your substance (in this case, water)--sometimes it's going to be ice; sometimes it's going to be water; sometimes it's going to be steam.0180

All right, now let's actually label some of the most important points here.0207

When the...here is 1 atmosphere: so, as it turns out, at 1 atmosphere, this is going to be the boiling point.0214

So, the boiling point is going to be 100 degrees Celsius: at 100 degrees Celsius and 1 atmosphere, it is...the phase changes between liquid and gas; that is what that is.0243

OK, now, at 1 atmosphere and 0 degrees Celsius (we drop this down), we have the melting temperature, which is 0 degrees Celsius; at this point, what you end up with--at 1 atmosphere pressure and 0 degrees Celsius, the phase change is going to be solid to liquid, ice to water; that is what this is.0259

All right, so now, let's have some other interesting points here.0290

There is a point here: I'll put it right there and I'll put it right here, and I will do this--I will put Tc, and I will put Pc.0297

Pc is something called the critical pressure, and Tc is exactly what you think--it's the critical temperature.0311

Well, something very, very interesting happens.0322

Oh, wait, let me mark off one other point here: this one and this one--this is a very, very important point.0324

We signify this as T3 and P3; this is called the triple point pressure and the triple point temperature.0336

At this pressure and at this temperature, all three phases exist simultaneously.0345

You can actually have ice, water, and gas all in equilibrium with each other, in a dynamic equilibrium; all will coexist.0351

It has to be...as it turns out, this critical pressure is 0.0060 atmospheres, and the critical temperature is 0.0098 degrees Celsius.0359

That is called the triple point--let's actually write that up here: T3...I'll do a slash...P3...this is called the triple point, and the triple point is where all three phases (gas, liquid, and solid) exist simultaneously.0376

So now, let's talk about Tc and Pc.0396

If I keeping raising the temperature, raising the temperature, and then increasing the pressure and increasing the pressure, something very, very interesting happens.0402

There comes a point where, at a certain pressure and at a certain temperature (in this case, it's going to be 374 degrees Celsius, and here at 218 atmospheres of pressure for water), at which point...there actually...I don't even know how to describe this.0410

There is, all of a sudden, no longer a difference: you don't have enough pressure to actually liquefy.0432

You see, if you keep raising the pressure of this, everything is just going to keep turning into a gas.0439

But, if you keep applying pressure to the system, you can keep turning that gas into a liquid, into a liquid, into a liquid.0443

At some point, you are going to reach a temperature beyond which, no matter what pressure you exert on that system, that gas will not turn into a liquid.0449

But the problem is, it is not a gas anymore, either; it is actually something called a supercritical fluid.0458

It is not a liquid, and it is not a gas; you can consider it, if you will, a fourth phase, because supercritical fluids (things at higher than critical pressures and higher than critical temperatures)--they have some very, very bizarre properties.0464

When I say bizarre, I mean very, very bizarre; you are welcome to either read about them in your books or look them up on the Internet--some really, really amazing things are happening.0480

It's a very fascinating field of chemistry and physics, working at supercritical temperatures and supercritical pressures.0490

Now, that is why this line literally stops here: it is not a gas; it is not a liquid; it is what they call a supercritical fluid.0498

It is like a totally different state of matter, and it behaves in profoundly bizarre ways.0509

That is it: that is called the critical temperature and critical pressure.0515

This is what it looks like for water: now, notice something really, really interesting here: notice that the slope here is actually negative.0520

For most substances, the slope is positive; and the reason it is positive for most substances is: most substances, as the thing becomes a solid--it actually gets more dense.0531

So, as something goes from a liquid to a solid, it becomes more dense--it becomes more tightly packed, as you would expect.0544

But water is different: when water freezes, it actually becomes less dense; it's the reason why ice floats.0550

Because of that density difference between liquid and solid, this has a negative slope.0557

It is not like that for CO2: CO2, in this case, behaves normally, because solid CO2 is more dense than liquid CO2, as you would expect.0563

But water does not behave the way we expect it to behave; that is why this is different.0574

That is it; this is a standard phase diagram; it expresses, at a particular temperature, at a particular pressure, what phases are going to exist--and, of course, the relationships.0581

We have this triple point; we have this critical temperature, critical pressure; that is it.0592

OK, so now, let's do the one for CO2.0600

CO2: OK, so let's go again here and here; so once again, we have temperature, which is going to be in degrees Celsius; we have pressure, which is going to be in atmospheres.0607

And now, we have something that looks like this.0624

This is solid; this is liquid; and this is gas.0638

OK, 1 atmosphere is actually here for CO2; so, at 1 atmosphere pressure, the temperature at which solid becomes a gas is -78 degrees Celsius.0647

So, at -78 degrees Celsius, solid carbon dioxide actually turns into a gas under normal, standard conditions of 1 atmosphere pressure.0667

It sublimes; that is why we call it "dry ice"--it doesn't actually get wet.0677

It goes directly from the solid to the gas phase; it doesn't pass through the liquid phase on its way.0680

1 atmosphere is here for solid CO2.0686

The triple point of CO2 is at...so P3 is equal to 5.1 atmospheres, and I will say the critical pressure is equal to 72.8 atmospheres--very easy to attain--72.8 atmospheres is very easy to do in the lab on a daily basis.0691

We don't really need any specialized equipment for that.0718

The critical temperature is 31 degrees Celsius--again, it's very, very easy to attain a supercritical fluid (supercritical CO2).0722

You raise the temperature to 31 degrees Celsius or a little bit higher, and then you just bring it to 72.8 atmospheres or higher, and you are in this range, where you can't tell the difference.0731

Is it a liquid? Is it a gas?--who knows: it's neither; it's both; it's a supercritical fluid.0740

Just for good measure, we will go ahead and talk about the triple temperature, the triple point, which is -56.6 degrees Celsius.0747

OK, so again, what is important to notice here is that, at 1 atmosphere pressure, which is normal conditions, CO2 sublimes.0761

"Sublime" means going straight from solid to gas phase.0776

That is what this is telling us here; and, as you see, you have a positive slope, because solid CO2 is more dense than liquid CO2, so there isn't any strange behavior.0781

But again, if you were to raise the temperature to, let's say, I don't know...I'm sorry, raise the pressure to 20 atmospheres, and then raise the temperature to...oh, I don't know, maybe 20 degrees Celsius, you will actually get liquid carbon dioxide.0793

I have never seen liquid carbon dioxide, myself; there you go--but it is possible: the phase diagram tells me so.0810

Actually, you can go to any fire extinguisher and just shake it up; that is liquid carbon dioxide.0819

OK, in fact, I will discuss that right now.0825

At 25 degrees Celsius and high pressures (pressures in the range of about 5.1 to 72.8), CO2 is a liquid and used in fire extinguishers.0830

When you open up the valve of a fire extinguisher, the liquid CO2 is all of a sudden exposed to 1-atmosphere conditions.0863

It immediately vaporizes--it immediately turns into a gas.0874

CO2 gas is heavier than air--it actually sinks; when it sinks, what it does is: it actually ends up smothering the fire.0878

It covers it up like a blanket; it literally just smothers it--sits on top of the fire.0887

And, because it is sitting on top of the fire, oxygen gas can't get to the fire; that is how it puts a fire out.0892

Not only that--in the process of going from liquid to vapor, that is actually an endothermic process; well, the energy for that endothermic process is coming from the surrounding air, so it's literally sucking heat out of the air.0900

So, what we feel is cold: because it is cold, that is even...that actually helps to retard the fire even more.0916

You have cooler temperatures, and you have no oxygen getting to the flame.0924

OK, as far as seeing fire extinguishers and this...it looks like you are actually blowing this white smoke at people (or at the fire); that white smoke is not carbon dioxide; that is the (because of the cold temperatures) water in the air, in the atmosphere, that is actually spontaneously condensing.0930

In other words, what you are doing is: you are forming a cloud, right here on the ground; that is it--that is what is going on there.0954

OK, so this was phase diagrams; I just wanted you to sort of see what they are.0960

You will see them on the AP exam; you will see them in your future work; so it's just good to know what they are.0965

All right, now let's go ahead and start talking about solutions.0972

Today, we are just going to talk about...I'll give you a brief introduction to solutions; we are going to talk about the different ways to actually represent solutions.0978

We are going to talk about molarity, mass percent, molality, mole fraction...things like that: the different ways of representing concentration.0985

Up until now, when we say "concentration," we are talking about moles per liter.0992

Well, it's true--the primary unit of concentration is moles per liter--but there are other ways to represent concentration, or the mixture of a solution--how much of a solute is in a solvent.0997

OK, so let us define what we mean by a solution: it is just a homogeneous mixture--that is it.1009

You take some salt; you drop it into water; you stir it up; and all of a sudden, you have a salt solution.1021

Salt is your solute--it is the thing that dissolves; water is your solvent--it's the thing that does the dissolving.1027

That is it.1033

So, a solute is the thing being dissolved, and the solvent is the medium doing the dissolving.1035

The medium...we are mostly going to be concerned with liquid solvents; it doesn't have to be a liquid...the medium doing the dissolving.1054

For example, the soda that you drink--the Coke, the diet Pepsi, the Sprite, things like that: actually, the solvent is water; the solute is actually a gas--it's carbon dioxide gas under high pressure.1071

When you put that under pressure, the carbon dioxide actually dissolves in the water; so you have a gas which is a solute, not a solid like salt.1083

Air--the air that you breathe--it is basically just oxygen gas that is mixed in with nitrogen gas: it is a solution--that is what it is.1094

OK, so let's say a sugar solution (just as a quick example): the sugar is the solute, and the water is the solvent.1106

OK, now, ways of discussing solution composition: how can we numerically represent how much of a solute is in our solution?1127

Well, here is how we do it--composition: OK, the first way is by something called mass percent.1151

Or, you will often see it as weight percent.1165

It is (so the mass percent)...now, these definitions are very, very precise.1172

It is the mass of the solute, divided by the mass of the solution; OK, the mass of the solution is the mass of the solute, plus the mass of the solvent.1181

The total mass of the solution, times 100: that is mass percent.1200

The second one (you already know) is molarity: it is the most common unit used to talk about concentration; it is moles of solute per liters of total solution (liters of solution--total volume--the volume of the solute and the volume of the solvent).1210

Because, again, even if it is a solid solute, that stuff has volume, it actually makes the volume bigger.1243

OK, the third one is something called mole fraction, and there is actually a symbol for it: it is the Greek letter chi.1251

The mole fraction of A is equal to the moles of A, divided by total moles.1260

So, if we happen to have 2 things...well, let's say we want to do the mole fraction of sodium chloride solution...or sugar solution (let's just take sugar; let's stick with sugar): I take the moles of sugar, divided by the total moles (the moles of sugar, plus the moles of the water).1277

It's a fraction: a fraction is always a part over the whole.1295

That is all this is--nothing that you don't already know.1299

And last, something called molality (not molarity, but molality): it is actually equal to the moles of solute (this is an interesting one) over kilograms of solvent.1303

Notice how these two are actually separate: solute and solvent.1323

OK, well, let's just do an example.1328

Example: A solution is made by mixing 2.5 grams of isopropanol with 100 grams of H2O.1332

The density of C3H8O is 0.786 (the C3H8O is the isopropanol, which, by the way, looks like this, in case you want the structure--there is another H here).1366

It is a 3-carbon chain, and in the middle carbon, it has a hydrogen attached to it; it also has a hydroxide attached to it; so that is isopropanol--that is basically rubbing alcohol, is what that is.1390

OK, grams per milliliter...your task is to express the concentration in the four different ways described above.1402

Now, I have to say: some people are kind of sticklers about this; when they say "concentration," they are speaking strictly about moles per liter.1430

But concentration, in its generic term, means how much of a solute is in a particular solution.1437

So, any of these could actually be used for concentration; so it just depends on what we are talking about.1443

If somebody talks about concentration, it is always good to ask, "What unit are we using? Are we using molality? Are we using molarity? mole fraction? mass percent? Something else called normality?" (which we won't discuss)--what?1448

OK, so once again, we have 2.5 grams of isopropanol, which is a liquid, mixed with 100 grams of H2O, also a liquid.1462

All right, so let's go ahead and do our first one, which is mass percent.1472

#1, which is mass percent: well, we said that mass percent is the mass of (let me just write it again) the isopropanol, over total mass, times 100.1476

Well, what is the mass of the isopropanol?--it is 2.5 grams.1496

What is the total mass?--well, the total mass is 100 grams, plus the 2.5 grams, times 100; so it's 2.5/102.5; you end up with 2.44% by mass.1499

2.44% by mass, 2.44% by weight; that means that, if I have 100 grams of that substance, 2.44 grams of it is going to be isopropanol.1518

If I have 50 grams of that substance, 1.22 grams is going to be isopropanol.1534

2.44% of any particular amount is going to be isopropanol; the rest of that solution is going to be water.1541

It is an expression of how much solute there is in our solution.1549

OK, #2, molarity: Well, molarity is the moles of isopropanol, over liters of solution.1553

So, here we have some conversions that we need to make.1570

Let's talk about moles of solute first: moles of isopropanol--well, we have 2.5 grams, and 1 mole of isopropanol is 60 grams; so we end up with 0.0417 mol; so we have the numerator.1573

Now, H2O: well, 100 grams is 100 milliliters, equals 0.100 liter (because water is 1 gram per milliliter in general--it's at 4 degrees Celsius, but not a problem).1593

Now, isopropanol (here we are doing volume): isopropanol--we said we have 2.5 grams, and we said that 1 milliliter of that is 0.786 grams.1611

So, what we have here is 3.18 milliliters, which is equivalent to 0.00318 liters.1631

Our molarity is equal to the moles, 0.0417 mol, divided by this plus this; 0.10318 liters, and we end up with 0.404 molarity.1644

That means, for every liter of solution, it contains .404 moles of that isopropanol.1673

OK, let's see...mole fraction: so, the chi of the isopropanol is equal to, again, the moles of the isopropanol (which we got from the previous one), over the total moles.1685

OK, so we said that our isopropanol contained 0.0417 mol.1711

Let me write these a little bit better here: 0.0417 mol.1720

H2O, on the other hand: we have 100 grams of it, and 1 mole of H2O is 18 grams; we end up with 5.56 mol.1728

Therefore, our chi of isopropanol is equal to 0.0417 mol, divided by 5.56, plus 0.0417; you end up with 0.0074 (I hope you'll forgive me if I don't write 7.4x10-3--I am actually not a big fan of scientific notation, myself--I prefer decimals).1743

That is the mole fraction.1773

And last but not least, molality (which will show up again when we discuss the colligative properties--boiling point, elevation, freezing point, depression, osmotic pressure, vapor pressure of solution, things like that): it is the moles of solute (oh, actually, let's just go ahead and...because we are dealing with isopropanol, let's just say "moles of isopropanol"), divided by the kilograms of water.1777

This is the one that is actually different; you are not actually combining things in the denominator--separate solute and solvent.1804

It equals 0.0417 mol, divided by 0.100 kilogram (right?--100 grams, .1 kilograms), equals 0.417 molal.1810

We say, for molarity, Molar; for molality, we say molal.1832

There we go.1836

OK, so these are the four ways that you are going to see concentration talked about.1840

In this particular chapter, we are going to talk about molarity mostly, but we are also going to be talking about molality, because again, when we discuss vapor pressure and the colligative properties in the next lesson, as it turns out, the mathematical expression requires that the concentration be expressed in molality.1845

So, I will go ahead and stop it there, and next time, we will discuss vapor pressure of a solution, and begin discussing colligative properties; and we will finish off our discussion of solutions.1863

Until then, thank you for joining us here at Educator.com.1873

We'll see you next time; goodbye.1875

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