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Raffi Hovasapian

Raffi Hovasapian

AP Practice Exam: Free Response Part III

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 (1)

0 answers

Post by Professor Hovasapian on July 18, 2012

Link to the AP Practice Exam:

http://apcentral.collegeboard.com/apc/public/repository/chemistry-released-exam-1999.pdf

Take good Care

Raffi

AP Practice Exam: Free Response Part III

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
  • Free Response 0:43
    • Free Response 6: Part A. i
    • Free Response 6: Part A. ii
    • Free Response 6: Part A. iii
    • Free Response 6: Part B. i
    • Free Response 6: Part B. ii
    • Free Response 7: Part A
    • Free Response 7: Part B
    • Free Response 7: Part C
    • Free Response 7: Part D
    • Free Response 8: Part A. i
    • Free Response 8: Part A. ii
    • Free Response 8: Part B. i
    • Free Response 8: Part B. ii

Transcription: AP Practice Exam: Free Response Part III

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

Today, we are going to close out our AP Chemistry course with the final instalment of the AP practice exam.0004

This is going to be the third part of the free response section.0010

In the last lesson, we actually did some reactions, and we did a question that involved experimental procedures.0014

This is the part of the free response section where you are actually not allowed to use calculators.0023

All you have to do is basically give good, solid explanations for why things are the way they are.0028

We are going to continue on with these types of questions.0035

The first one is going to discuss thermodynamics; so this is going to be question #6: let's just jump right on in.0037

OK, so answer the following questions in terms of thermodynamic principles and the concepts of kinetic-molecular theory.0044

Consider the reaction: CO2 gas + 2 NH3 gas goes to CO(NH2)2 solid + H2O liquid; and we have a ΔH equal to -134 kilojoules.0051

We have this reaction: we have carbon dioxide gas and ammonia gas going to this compound (which is a solid, notice), and H2O, which shows up as a liquid.0083

So now, let's go ahead and start answering some of the questions here.0094

This is A, and this is part 1 of A: so, for the reaction, indicate whether the standard entropy change is positive, negative, or 0; and they want you to justify your answer.0098

OK, so ΔS--that is the standard entropy change; well, take a look at what is happening in this reaction.0113

We have a gas, and we have another gas; we have 3 moles of gas in the reactants (which is a highly disordered system), going to a solid and a liquid (which are very highly ordered systems).0119

So, the entropy of the products is less than the entropy of the reactants; remember, entropy is disorder--it's a measure of the chaos in a system.0131

Gases are going to be a lot more disordered than a solid and a liquid; therefore, the entropy of the products is less than the entropy of the reactants.0139

Therefore, ΔS, which is products minus reactants, is actually going to end up being less than 0; it is going to end up being negative.0156

That is why: and your justification is "because reactants have a higher entropy than the products."0165

That is it; OK.0187

Part 2 says: Which factor (the change in enthalpy, ΔH, or the change in entropy, ΔS) provides the principal driving force for the reaction at 298 Kelvin?0190

Well, remember: they said that this reaction is spontaneous (right?--right up at the top, where you can read it, it said this is a spontaneous reaction).0200

Well, "spontaneous" implies that the ΔG, the free energy, is less than 0; in other words, it is negative.0210

Well, we know that ΔG is equal to ΔH minus TΔS.0218

Well, what do we know about ΔH?--it's negative.0226

Well, what do we know about ΔS?--ΔS is negative.0228

Well, if ΔS is negative, that means the -TΔS term is positive; well, if the -TΔS--if this term--is positive, and yet this is a spontaneous reaction (ΔG is negative), that means the driving force for the spontaneity of the reaction is the ΔH term.0231

That is negative; so, ΔG is negative; a negative number equals a negative number, plus something positive.0256

Because it's a spontaneous reaction, it says that that means that ΔG is negative; therefore, this negative term dominates under these conditions, so ΔH is the answer.0264

The reason is precisely this: write out the equation; you show that the -TΔS term is positive; it is positive because you know from part 1 that ΔS is negative; therefore, because it is spontaneous (negative ΔG), that means that ΔH is actually the driving force.0280

The enthalpy is the driving force for the spontaneity of this reaction at this temperature.0297

OK, part 3 says: For the reaction, how is the value of the standard free energy, ΔG, affected by an increase in the temperature?0304

Well, we use the equation again: ΔG equals ΔH minus TΔS.0315

Well, the ΔH term is negative; this term right here--the -TΔS--is positive; so, if we increase the temperature, as temperature increases, this -TΔS term increases.0324

In other words, it becomes more positive.0342

ΔH is fixed--it is fixed at -134; if I keep increasing the temperature (ΔS is also fixed, whatever it happens to be, but), this temperature, as it goes up--this whole term, -TΔS, goes up.0346

But -TΔS is positive, so it becomes more positive; or, ΔG becomes less negative and, if you want to say so, eventually positive.0361

Again, you can raise the temperature, eventually, so high that it is going to overrun the ΔH term, and it is going to make the reaction all of a sudden become non-spontaneous.0384

That is it--basic thermodynamic principles.0394

ΔG=ΔH-TΔS; ΔG is negative--that means spontaneous; ΔH negative means exothermic; ΔH positive means endothermic; if ΔS is positive, that means that entropy has increased from reactants to products; if ΔS is negative, that means entropy has decreased from reactants to products.0398

ΔG positive means non-spontaneous (or spontaneous in the reverse direction); that is what all of these mean.0423

All of these are accessible via this one equation.0428

OK, so now, let's talk about part B.0432

Let's see, what does part B say?--Some reactions that are predicted by their ΔG to be spontaneous at room temperature do not proceed at a measurable rate at room temperature.0438

#1: Account for this apparent contradiction.0454

Well, as it turns out, there is no contradiction here.0458

ΔG is a thermodynamic property; rate is a kinetic property--that is all you have to write.0461

ΔG is thermodynamics; rate is kinetics; so thermodynamics don't tell me how fast a reaction is going to go--they are telling me that, if I can get the reaction to go somehow, that it will go.0469

In other words, once I get it started, I don't have to do anything to it; rate is kinetics.0491

If you want to draw a little energy diagram, you have this; you have this; you have something like this; OK.0495

Now, ΔG--this part right here--that is thermodynamics; this part right here--the activation energy (remember, the energy of activation--it is the energy in order to get over this initial hump, in order for the reaction to proceed forward)--this is what controls the rate.0506

This is thermodynamics; this part here is what controls the rate.0526

The lower this is, the faster the reaction; in other words, more of the reactants have enough energy to get over this hill.0529

But, if this is really, really high, there aren't going to be a lot of particles that have enough energy to get over that hill, so the reaction rate is going to be slow.0536

ΔG is thermodynamics; rate is kinetics; a slow reaction (this is one thing that you definitely want to say) means a high activation energy; you definitely want to mention activation energy here.0545

You don't just want to leave it as "ΔG is thermodynamics, and rate is kinetics"; you want to say which energy actually controls the rate: it's the activation energy that controls the rate.0564

ΔG tells you about thermodynamics.0575

OK, so that is question #1; now, #2 says: A suitable catalyst increases the rate of such a reaction; what effect does the catalyst have on ΔG for the reaction?0578

OK, well, catalysts affect rates, not ΔG (and if you want to mention it) or Keq.0592

A catalyst does not change the equilibrium constant; a catalyst does not change the thermodynamic property of a given reaction under a given set of conditions.0611

The only thing that a catalyst does is reduce the activation energy and/or provide an alternative pathway, which is tantamount to reducing the activation energy.0621

A catalyst makes reactions faster; it does not actually change the nature of the reaction.0632

That is thermodynamics; and again, a catalyst affects kinetics; ΔG is thermodynamics; the two are completely different domains.0639

Both are important, but they are actually handled separately; that is why we had different chapters--we had a kinetics chapter and a thermodynamics chapter.0648

Both are important, but a catalyst has no effect on the rate.0655

And again, if you want to use this same thing right here, you can do that; this energy difference is thermodynamics, ΔG; this energy difference is kinetics.0660

That is that; OK, now, let's go to question #7.0672

OK, question #7: Answer the following questions, which refer to 100-milliliter samples of aqueous solutions at 25 degrees Celsius in the stoppered flasks shown above.0682

In front of you, as you are looking at these questions, you should see some flasks (4 of them); they are all .1 Molar; one of them is sodium fluoride; one of them is magnesium chloride; one of them is C2H5OH (which is ethanol); and one of them is CH3COOH, which is acetic acid, or vinegar.0697

So, we have sodium fluoride, magnesium chloride, ethanol, and acetic acid.0717

OK, so now, we are going to answer some questions about it: Which solution has the lowest electrical conductivity?0722

Well, the lowest electrical conductivity means the least dissociated into free ions.0728

Therefore, C2H5OH has the lowest conductivity.0756

The lowest conductivity means the least dissociated into free ions; of these things (the sodium fluoride, the magnesium chloride, the ethanol, or the acetic acid), the one that is dissociated the least is going to be the ethanol.0779

In fact, it is probably going to be not dissociated at all; it is a molecular compound--it isn't ionic at all.0793

Acetic acid will dissociate a little bit into H+ and acetate; magnesium chloride and sodium fluoride will dissociate completely, because they are ionic compounds.0798

This is a molecular compound: there is no dissociation (or very, very little).0809

It is true that this H is slightly acidic, but it is acidic on the order of water; so you are talking about 1x10-14; it is not going to dissociate--there is no conductivity.0814

OK, now B says: Which has the lowest freezing point?--explain.0825

OK, colligative properties: when you add a solute to a solvent, you create a solution.0836

You lower the vapor pressure; you lower the freezing point; you elevate the boiling point; and you activate the osmolarity of that solution.0843

OK, lowest freezing point means highest number of solute particles.0854

The highest number of solute particles between sodium fluoride, magnesium chloride, ethanol, and acetate--well, sodium fluoride will dissolve into 1 mole of sodium ions and 1 mole of fluoride ion.0877

Magnesium chloride will dissociate and dissolve into 1 mole of magnesium ions and 2 moles of chloride ions, for a total of 3 moles of particles.0889

C2H5OH will not dissolve at all, or dissociate at all.0900

And the CH3COOH will be partial dissociation into H+ and acetate ion.0905

Of those, magnesium chloride actually produces the most solute particles--free particles floating around in the solvent, which is water.0910

Again, the identity of the solute particles does not matter; all that matters is the number of particles floating around, which is why they are called colligative properties.0922

Identity is irrelevant; it just matters how many particles are there.0931

In other words, how many particles are interfering with the solvent molecules doing what they normally do?0935

That is it.0940

OK, C: Above which solution is the pressure of the water vapor greatest?--explain.0943

OK, here we go with our implications: more solute particles means a lower vapor pressure; therefore, a lower vapor pressure implies that C2H5OH is our particular answer to C, because C2H5OH has the fewest solute particles.0957

It has the fewest solute particles because it doesn't dissociate at all; it is just the amount of ethanol that you actually drop into water is that much ethanol; it doesn't break up into more particles like sodium fluoride, magnesium chloride, and acetic acid (to some extent).1000

OK, let's see what we have for part D--D says: Which solution has the highest pH?--explain.1016

OK, well, highest pH means the most basic; it doesn't necessarily mean that it is going to be basic--it just means you might have a solution of 6 (which is acidic), and another solution which is 4 (which is also acidic), but the 6 has a higher pH.1025

Let's take a look at acetic acid: CH3COOH--that is going to be reasonably acidic; C2H5OH--well, it is not really going to dissociate much, in terms of hydrogen ion, so probably the pH is going to be reasonably neutral.1046

So yes, compared to those two, it is going to have a higher pH than the acetic acid.1061

The magnesium chloride--neither magnesium nor chloride is such that it is going to cause any sort of hydrolysis of the water; it is not going to do anything to the water molecules to release a hydrogen ion or hydroxide ion.1065

So, you are looking at a neutral pH there.1079

However, sodium fluoride--when sodium fluoride dissociates, what you are going to end up with is this: fluoride is the anion of a weak acid, HF (hydrofluoric acid).1081

So, when F- is released into solution, what is going to happen: it is going to react with the water (HOH) to produce HF (hydrofluoric acid or hydrogen fluoride, really), plus OH-.1105

This free OH- in solution--that is going to create a basic solution; you are going to have a pH which is higher than 7.1119

So, your answer is sodium fluoride; and this is the reason why: because F- is the anion of a weak acid; therefore, it will cause hydroxide ion to form.1125

If you write this, and just write this equation, that is all you need to write--very simple; very straightforward.1136

This is just recognition of chemistry; this part of the AP exam really tests the extent to which you understand the chemistry (regardless of the math).1141

OK, so we are down to our last problem in the AP chemistry exam; so this is kind of exciting--it has been a long road, but it has absolutely been worth it, I hope.1153

Question #8: Consider the carbon dioxide molecule and the carbonate ion (so we have CO2 and CO32-).1166

The first question says: Draw the complete Lewis electron dot structure for each species.1181

OK, so we have done this many times--these are the last lessons that we did, in fact, before we started the AP practice exam.1186

The Lewis dot structure for CO2 is this (and remember, you have to include the lone pairs of electrons); that is CO2.1194

Now, CO3- looks something like this.1202

OK, now, this is actually part of a resonance structure, so I am going to go ahead and actually draw all three resonance structures; it's not necessary to draw all three, but you know what, it is not a bad idea.1217

That way, they won't penalize you, because they might actually ask you a question regarding resonance.1228

This is going to be CO; I'm going to put the double bond on the left; I'm going to do 2-, 2, 4, 6, 2, 4, 6; and I'll go ahead and draw this other resonance structure over here--it doesn't matter--vertical or horizontal.1233

2-, CO; I'm going to draw the double bond on the left this time; there is 6; there is 2; there you go--that is CO2, and this resonance structure; so this structure is actually an average of all three.1253

This is the Lewis structure; however, you would be OK if you just drew one of them--that is not a problem.1270

That is part 1; now, part 2 says: Account for the fact that the carbon-oxygen bond length in CO32- is greater than the carbon-oxygen bond length in CO2.1276

Here is the reason why: notice the Lewis structure here--these are fixed double bonds; CO2 has fixed double bonds, which have a specific length.1293

Now, CO32- has resonance structures, and each bond has both single-bond character and double-bond character.1318

Remember, when we have a resonance structure, it isn't that it is this or this or this; it is all three simultaneously.1338

It has both single-bond character and double-bond character: that is the reason why.1345

CO2 has fixed double bonds, which have a specific length; CO32- has resonance structures, and each bond has both single-bond character and double-bond character.1366

Therefore, because (let me see, how shall I write this) the bond is between single and double, it is longer than a straight double bond; there you go--something like that.1375

As long as you mention resonance structure, and you say that it has single-bond and double-bond character, the rest of it is actually pretty implicit; there you go.1418

OK, now part B--it says: Consider the molecules CF4 and SF4; so, we have CF4, and we have SF4: carbon tetrafluoride; sulfur tetrafluoride.1433

Part 1 says: Draw the complete Lewis structure for each molecule.1450

OK, so when we do a Lewis structure for CF4, we are going to end up with something like this; and here we go with the lone pairs; that is CF4.1455

Now, SF4 is going to look almost the same, except we are going to have a lone pair on sulfur, also.1473

There you go: those are the two Lewis structures, CF4 and SF4.1490

The only difference: S has those extra two electrons; and those two electrons are actually on the sulfur itself.1494

Sulfur can accommodate more than an octet, because it is in the third row; it has d orbitals available to it.1501

OK, now, part 2 says: In terms of molecular geometry, account for the fact that the CF4 molecule is nonpolar, whereas the SF4 molecule is polar.1508

OK, molecular geometry: we have 4 objects around a central atom, and the central atom does not have any lone pairs, so remember that table that we were looking at.1519

CF4 has a tetrahedral geometry; that geometry looks like this: C, this is F (I'm not going to draw the lone pairs); there is an F (no, I'm not going to have that; I can't have random lines on a Lewis structure--that is not going to work); F; F; F.1532

Here, as it turns out, because fluorine is more electronegative than carbon, you are going to have a dipole that way; you are going to have a dipole to the back and to the left; you are going to have a dipole to the back and to the right; and you are going to have a dipole pulling forward.1563

All of these dipoles are going to cancel (the bond dipoles--the individual bond dipoles); so all right, bond dipoles cancel, which makes it nonpolar.1578

Now, CF4 looks something like...I'm sorry, SF4 looks like this: SF4...well, you have 5 objects around this central atom: you have 1, 2, 3, 4; your fifth object--remember, when we talk about an object, it includes lone pairs.1598

Now, because you have 5 objects around a central atom, it is actually going to be triagonal bipyramidal.1617

But because it has one lone pair, what we are going to have is a seesaw arrangement of the atoms.1624

SF4 looks like this: there is an F here; there is an F there; there is an F going backward; there is an F coming forward; and there is a lone pair there.1631

The dipole bond, dipole bond...those two end up canceling.1645

However, now we have a dipole pulling to the left and back, and it is pulling to the left and front; the net dipole is going to be in that direction (to the left).1650

The vertical components of this dipole and this dipole cancel, but the horizontal ones--they are both to the left, so this is going to act like a bar magnet that has a negative side and a positive side.1663

Therefore, this is polar.1673

So here, the bond dipoles do not cancel; therefore, you have a polar molecule.1676

OK, and that, my friends, takes care of the AP practice exam.1699

It has been my pleasure presenting AP Chemistry and going through this AP Chemistry exam with you.1703

I wish you the best of luck in all that you do.1709

Take good care, and we will see you next time; thank you for joining us here at Educator.com; goodbye.1711

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