Enter your Sign on user name and password.

Forgot password?
Sign In | Subscribe
Start learning today, and be successful in your academic & professional career. Start Today!
Use Chrome browser to play professor video
Raffi Hovasapian

Raffi Hovasapian

Oxidation-Reduction & Balancing

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
Loading...
This is a quick preview of the lesson. For full access, please Log In or Sign up.
For more information, please see full course syllabus of AP Chemistry
  • Discussion

  • Study Guides

  • Download Lecture Slides

  • Table of Contents

  • Transcription

  • Related Books & Services

Lecture Comments (9)

1 answer

Last reply by: Professor Hovasapian
Tue Dec 2, 2014 2:43 AM

Post by Long Tran on November 30, 2014

hello Professor
on the example 2 i understand that Oxidation state that Fe2+ to Fe3+. My question is what substance is oxidized? is it Fe2+ or Fe3+
Thank you for lecture

1 answer

Last reply by: Professor Hovasapian
Mon May 13, 2013 5:18 PM

Post by Nawaphan Jedjomnongkit on May 13, 2013

Thank you for the lecture but I just don't understand why Fe will be the one that want to lose e- not Mn, i.e. why the reaction will not happen in the opposite direction , while in the table Fe is on the right of Mn that mean it's EN will be more than Mn from the periodic trend.

1 answer

Last reply by: Professor Hovasapian
Thu May 9, 2013 3:21 AM

Post by Fiona McLagan on May 8, 2013

Thank you, Dr Hovasapian, for all you do.

1 answer

Last reply by: Professor Hovasapian
Wed Aug 22, 2012 8:05 PM

Post by Etienne Carrier on August 22, 2012

I can't thank you enough about these lessons. I'm doing one physics and two chemistry subjects at university this semester, and I would be completely lost without this site and all of you awesome teachers. Thanks a million!

0 answers

Post by Linnea Huson on May 21, 2012

I am so grateful for these videos! They are clear and concise. I especially like how the teacher goes step by step through a process. If I didn't have this video I wouldn't be making it through my current chemistry class. Thanks again.

Oxidation-Reduction & Balancing

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
  • Oxidation-Reduction and Balancing 2:06
    • Definition of Electrochemistry
    • Oxidation and Reduction Review
    • Example 1: Assigning Oxidation State
    • Example 2: Is the Following a Redox Reaction?
    • Example 3: Step 1 - Write the Oxidation & Reduction Half Reactions
    • Example 3: Step 2 - Balance the Reaction
    • Example 3: Step 3 - Multiply
    • Example 3: Step 4 - Add
    • Example 3: Step 5 - Check

Transcription: Oxidation-Reduction & Balancing

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

Today, we are going to begin our discussion of electrochemistry.0004

I personally love electrochemistry.0008

Electrochemistry is absolutely everywhere: the life that you enjoy in the modern world that we live in is all about electrochemistry.0012

Basically, just take a look at your cell phone; batteries--batteries make everything possible in life--computers, cell phones, any number of things.0024

Batteries are just a simple application of actually very, very simple electrochemistry.0034

Today, we are going to start our discussion of it; and in today's lesson, we are actually going to be talking about--doing a little review on--oxidation-reduction, because that is what electrochemistry is.0040

It is just about oxidation-reduction chemistry.0051

We'll define what we mean; again, I know we talked about it a while back, but now we are really going to get into what oxidation-reduction means.0055

We are going to talk in great detail about where the electrons go, how they move, things like that.0063

Today, after a brief discussion of oxidation-reduction and a review, we are going to talk about balancing oxidation-reduction reactions.0069

It's not like balancing other reactions, where you just have to sort of fiddle with coefficients: that is part of it, but because oxidation-reduction, under many circumstances, involves some really, really complex changes, in terms of atoms shifting around and electrons moving around, things don't look the same as they did maybe before and after.0078

Oftentimes, you look at a reaction that is asked to be balanced--the oxidation-reduction methods--and you think, "Oh, wow, what is this?"0104

So again, a slightly different procedure--but again, very, very systematic; so today, we are just going to develop some of the tools of oxidation-reduction.0113

And then, next lesson, we will start with some definitions on what electrochemistry really is.0121

Let's just go ahead and get started.0127

Let's do a quick definition of electrochemistry, to begin with: So, electrochemistry--it is using oxidation-reduction chemistry to generate an electric current.0130

That is it--an electric current means we want to get some electrons moving in one direction or another.0161

That is it; that is all that is happening--an electric current is just moving electrons--that is it.0169

When we move electrons through a wire, we can make things run: a light bulb, a heater, a computer...whatever it is.0175

That is it; it is just moving electrons.0182

OK, so let us recall: oxidation is just a fancy word for a process that loses electrons--so it's a loss of electrons.0185

That is it--nothing more, nothing less; it's a fancy word--it's going to look like a lot of things are going on--but it just means something has lost electrons.0200

When sodium goes from sodium metal to sodium +1, it has lost an electron; it has been oxidized; it has undergone oxidation--loss of electrons.0208

Reduction is the opposite; it is the gain of electrons (I'll put plural, because it doesn't have to be just one).0220

When oxygen goes from oxygen to...let's use sulfur instead.0232

When sulfur goes from sulfur to sulfur 2-, well, it has gained 2 electrons--it has been reduced--it has undergone a reduction.0237

Oxidation and reduction--they go together: when something is reduced, something has been oxidized; when something is oxidized, something has been reduced.0248

The analogy is like with acids and bases: when you have acid, if something gives a proton, generally something is taking the proton; if something is taking a proton, it is generally taking it from something that is giving the proton.0257

That is the whole idea: in acid-base chemistry, it was the proton that bounced back and forth between two species; here, in oxidation-reduction, it is the electron that is bouncing back and forth between two species.0269

You are going to discover, as you go on--and hopefully through our discussion--that there are a lot of similarities between acid-base chemistry and oxidation-reduction chemistry, when you look at it globally.0282

OK, so let's just do a quick reaction: 2 sodium molecules, plus a chlorine molecule (chlorine is a diatomic gas) goes and forms 2 units of sodium chloride.0293

Notice, I didn't say "molecule"; an ionic compound is not really a molecule--it's an extended crystal lattice--crystal network.0311

OK, so here is what is actually happening: sodium has an electron; sodium has an electron; and chlorine is bonded with chlorine in a single bond.0326

Now, in this particular...for those of you who are sort of going along this course systematically, I haven't discussed the notion of bonding yet.0339

Those were several topics in sort of the middle of the chemistry that I decided to skip, so that I could go to kinetics and equilibrium and things like that.0349

Just know that this single bond just represents 2 electrons: and we will discuss it later on--don't worry--it's not going to interfere with anything here.0356

This is just notation.0365

So, what happens is the following: well, sodium gives up its electron to one of the chlorine atoms, and this sodium gives up its electron to one of the chlorine atoms.0367

Or, another way of saying it: one of the chlorine atoms rips that electron away from sodium, and this chlorine rips it away from this sodium.0378

We have a convention: whenever we talk about the movement of electrons, we always go from the electron to the direction--to where it's going--not the other way around.0387

And also, when we are talking about a single electron (this is more of a convention in organic chemistry, but I think it is appropriate to sort of be introduced to it here)--something called a fishhook--only one thing like that; we use an arrow for two electrons--the movement of two electrons.0396

Here, because one electron is going that way, we use a fishhook.0410

What you end up with is the following: well, when sodium loses its electron, it becomes a sodium +; this becomes a sodium +; when chlorine atom takes one of the electrons, this bond actually breaks, and what you get is a chloride ion.0414

You get a chloride ion; now, they slam together, because positive and negative charges are attracted to each other, and what you get is salt.0429

Salt is only possible with an oxidation-reduction reaction: sodium has been oxidized--it lost electrons; chlorine has been reduced--it gained electrons: oxidation-reduction.0440

The shorthand is redox: when we talk about oxidation-reduction processes, we talk about a redox process.0454

When we talk about balancing an oxidation-reduction reaction, let's balance a redox reaction.0460

So now, we are going to be using "redox."0464

So now, what we want to do is do something that we did earlier, but I think it's good to review: we want to assign oxidation states to certain compounds and elements, to decide what has been oxidized and what has been reduced, because it's not quite so easy to figure it out sometimes.0468

Here, this is easy: you see what 0 charge went to +1; this 0 charge went to -1; you can see what has been oxidized and reduced; but it isn't always that clear.0486

So, assigning...let's not say assigning; let's say: Let's assign oxidation states to decide what has been oxidized, and what, reduced.0496

OK, now, here I'm going to say: Look up the rules for assigning oxidation states in your book.0534

We actually have it in a previous lesson, in the earlier part of the year--actually, I think it's maybe Lesson 6, Lesson 7, somewhere around there--where I actually wrote out the rules; but you will find them in your book, of course.0550

Or, you can find them online; in your book, just look in the index, under "Oxidation states" or "Rules for assigning oxidation states," and it will direct you to the page.0563

I'm not going to rewrite them here; I just want to be able to use those rules, so look up the rules for assigning oxidation states.0572

Now, if you are still confused by what we mean by an oxidation state, an oxidation state refers to the extent to which a particular element has lost or gained an electron or electrons.0580

It will make more sense when we actually do the problems; it's sort of difficult to define.0593

So, let's see: so, when we assign oxidation states, we are assigning them for each element, OK?0600

An element might be in a compound, but we are assigning it to each individual element--to each single atom.0608

OK, so we'll just do some examples.0615

Example 1: Assign oxidation states to each atom in the following: Well, how about oxygen?0621

Oxygen: well, we know it's a gas, so not a problem: so, we are assigning oxidation states to each individual atom.0646

Oxygen gas is a diatomic molecule; in this particular case, oxygen and oxygen, it looks like this.0654

OK, there actually isn't a double bond here, but don't worry about that; I'll actually discuss why that is later.0667

Those of you who actually learned that oxygen is bonded as a molecule--you learn it as a double bond to make sure that the whole bonding scheme that you learn--that whole octet stuff--is satisfied.0675

Oxygen is actually a diradical, and I actually teach it as, an oxygen is a diradical.0686

For our purposes, that is not a problem; what is important here is that these two--since they are equal, neither one of them has actually taken electrons from the other.0691

The oxidation state is 0; in fact, for any element--for any element--the oxidation state is 0.0702

That is one of the rules.0711

B: How about chloride ion--what is the oxidation state of the chlorine atom?0713

Well, it is very simple: the oxidation state of the chlorine atom is the charge on that ion: it's -1; that is it.0718

How about magnesium 2+, the magnesium 2+ ion--what is the oxidation state of the magnesium?0728

Well, the oxidation state of any ion is the charge on that ion: +2.0734

OK, now we get into some interesting territory: NH3: what is the oxidation state of H, and what is the oxidation state of N?0741

All right, when hydrogen is bonded to a nonmetal, it carries an oxidation state of +1, always.0752

It's one of the rules; the other rule is, when it is bonded to a metal, it carries an oxidation state of negative 1--always.0765

Well, OK: here is where we do something interesting.0771

The sum of the oxidation states of the individual elements--the total sum--has to equal the charge on the whole species.0776

This is NH3; NH3 is a neutral molecule: there is no charge on it--it is just a 0 charge.0786

Well, if hydrogen carries a +1 charge when it's bonded to a nonmetal, and there are three of them, here is what I do: I'm actually going to do it over here...NH3...here is how I do it.0793

Each hydrogen has a +1 charge; there are three of them, for a total of a +3 charge.0810

Well, +3, and the total charge is 0 on this species; so nitrogen has to be a -3, because there is one nitrogen.0815

The oxidation state on nitrogen is -3.0826

Basically, what that means is that nitrogen, based on the difference in something called electronegativity (which, again, shouldn't concern us right now), has taken the electron from one hydrogen, an electron from another, an electron from another.0829

If I were to separate these, I would find that nitrogen is now carrying a -3 charge, and each hydrogen is carrying a +1 charge.0842

That is what oxidation state is actually saying: it's saying where the electrons are--where did they go?0850

In this case, the electrons are with nitrogen, even though this whole thing is bonded; it's not an ionic compound--we are just assigning oxidation states.0856

OK, so go ahead and do this here...zero...0867

E: How about NO2?0873

Well, what is the oxidation state on oxygen; what is the oxidation state on nitrogen?0876

OK, one of the rules is that oxygen always has an oxidation state of -2.0881

It is very electronegative: it will take electrons.0885

The oxidation state is -2 on each individual oxygen atom; there are 2 oxygen atoms, so the total negative charge is -4.0888

Well, -4...the charge on the species...it's a neutral species, so the nitrogen has to be a +4.0897

I'm going to erase this; I'm going to need the space here.0911

Once again, the rules for assigning oxidation states: oxygen always has a -2 (except in some circumstances, but we won't deal with those); there are two oxygen atoms--the total charge is -4.0917

Well, -4 plus what is going to give me a 0 charge?--negative 4, plus 4.0930

There is only one nitrogen, so that entire nitrogen is carrying a +4 charge--oxidation state.0936

OK, let's try F: let's try NO3-, nitrate.0942

What is the oxidation state on that; what is the oxidation state on this?0949

Oxygen has a -2 charge; there are three of them, for a total of -6.0954

-6, plus what, will give me a -1 charge for the total species?0959

+5: nitrogen is a +5; oxygen is a -2.0964

I hope this is starting to make sense--it's very, very important to be able to assign oxidation states to any compound.0969

G: let's try CrO42- (chromate ion).0976

Well, we want to know the oxidation state on oxygen and the oxidation state on chromium.0984

Oxygen has a -2 charge; there are four of them; that means a total charge of -8.0990

Well, the charge on the whole species is -2: -8, plus what, is equal to -2?0997

+6: since +6, there is only one chromium: that means chromium has a +6 oxidation state; that means oxygen has a -2.1003

Again, the oxidation state is listed for each atom, but then we take the number of atoms that are in there and multiply it by the oxidation states to work with the total charge.1012

HMnO4-: this is the permanganate ion.1026

What is the oxidation state on manganese? What is the oxidation state on oxygen?1032

Well, we know what the oxidation state on oxygen is--it's -2, always; -2: there are 4 of them, for a total of -8.1036

-8, plus what, will give me a -1?1044

+7: the oxidation state on manganese is +7.1047

In this compound, manganese has lost 7 electrons to oxygen; in this compound, chromium has lost 6 electrons to oxygen; in this compound, nitrogen has lost 5 electrons to oxygen; in this compound, nitrogen has lost 4 electrons to oxygen.1052

In this compound, nitrogen has gained three electrons from hydrogen; that is what oxidation state is saying--how many electrons have been lost or gained in the formation of that compound.1070

That is all that is going on here.1083

OK, now, let's take the next step.1085

Is the following a redox reaction?1091

Example 2: Is the following a redox reaction?1094

If some of this is strange, I would encourage you to go back and take a look at the previous lesson, where we actually talk about assigning oxidation states.1107

I write out the rules; I talk a little bit more in-depth about oxidation-reduction--this is just sort of a review to get things going again.1113

OK, is the following a redox reaction (in other words, is it an oxidation-reduction reaction)?--some reactions are not; some reactions are acid-base--there is no real oxidation-reduction going on.1123

Some reactions--nothing changes; so a redox reaction is when there is a change in oxidation state from reactants to products.1132

You will see what we mean in just a minute.1143

In other words, do oxidation states change in going from reactant to product?1146

That is it: reactant to product--this is how you decide if a particular reaction is an oxidation-reduction reaction: you assign oxidation states to each thing in that equation, and then you check to see, on both sides...the reactant--like the iron on the reactant side and the iron on the product side...has the oxidation state changed?1164

If it has, it's an oxidation-reduction reaction; electrons have been shifted around.1185

OK, so let's do 8 H+ (don't worry about how this looks; it's a perfectly good equation)...MnO4-...(this is what I was talking about; some of these can look really weird--don't let the symbolism scare you)... [8 H+ + MnO4- + 5 Fe2+ → Mn2+ + 5 Fe3+ + 4 H2O]1190

You are presented with this equation, and you are asked, "Is the following an oxidation-reduction reaction?"1219

Well, I have to assign oxidation states: let's do it, just like we did in the previous example.1224

H+: the oxidation state on it is +1; permanganate--we said that the oxidation state was +7, and for oxygen it's -2 (remember, -2 is always the same for oxygen).1230

There are 4 of them, for a total of -8; -8 plus 7 gives me a -1, so the manganese has an oxidation state of +7.1245

Iron: well, the oxidation state is +2; here, manganese--the oxidation state is +2.1253

There you go: manganese has gone from a +7 oxidation state to a +2; the oxidation state on manganese has changed.1261

It doesn't matter that here it is with a compound, here it's alone; that doesn't matter--that is just chemistry.1274

Iron is +3 (I'll go back to blue); this is +3; sure enough, iron has gone from a +2 oxidation state to a +3; oxidation states have changed--yes.1280

Oxygen -2; hydrogen +1--this is a nonmetal; hydrogen hasn't changed; oxygen hasn't changed; manganese has changed--the oxidation state has gone down.1294

Iron has changed--the oxidation state has gone up.1309

This is where the word reduction comes from: reduction means the number...from +7 to +2, the number is reduced from +7 to +2; it has gained electrons.1313

Iron: +2 to +3--it has been oxidized; oxidation is actually an older term than reduction was--that is why we call it oxidation and reduction--that's why it doesn't seem to really match all that well.1329

+2, +3: this has been oxidized.1342

Oxidation and reduction come in pairs: so you see, this is an oxidation-reduction reaction; and the reason we know that--because oxidation states have been changed.1345

One of them has gone down; one of them has gone up; it will always be that way--if one of them has gone down, I promise you, somewhere, something has gone up, because oxidation-reduction comes in pairs.1355

OK, now Example 3: OK, so the Example 3...Example 3 and Example 4 are going to be very, very, very, very important examples.1366

I am going to outline a method for actually balancing an oxidation-reduction reaction--two methods: under acidic conditions and under basic conditions.1379

The reason this is the case is because often, oxidation-reduction reactions happen under acidic or basic conditions.1391

We will run through the process; we will be very, very careful about this; this is very, very important to be able to do.1401

OK, so make sure you understand it completely.1407

Balance the following reaction under acidic conditions...so the first one is going to be under acidic conditions.1410

Balance the following under acidic conditions...and we know what acidic conditions means--it means H+, hydrogen ion.1417

That is what acid is; acid is just free hydrogen ion; OK.1432

MnO4- + Fe2+ goes to Mn2+ + Fe3+.1439

We are presented with this equation, and we need to balance it.1453

Now notice: this doesn't look like the other equations that we are given--it's not just about putting a coefficient here and a coefficient there.1456

Fe, Fe, Mn, Mn...all of a sudden, there is an oxygen here; there is no oxygen on this side; how are we going to introduce an oxygen?1462

And how are we going to balance charge?--because balancing doesn't just mean balancing the individual molecules or atoms; it means balancing charge.1470

Here we have 3+ and 2+; this is 5+ over here; this is 2+ and 1-; this is 1+ over here.1480

This is all kinds of crazy.1487

So, an oxidation-reduction reaction...now, we don't know if it's an oxidation-reduction reaction yet.1489

This is what we are going to check; so we are going to do the same thing that we did.1493

When we discover that it is an oxidation-reduction, confirm it; then, we're going to go ahead and take the next step and balance it.1496

So, let's do that: well, here we go--the oxidation state of iron is 2+; the oxidation state of this is 3+; the oxidation state of this was 7+ (that we said from previous work); and it's 2+; so yes.1502

So, the first thing we need to do is...this is the first step: Write the oxidation and reduction half-reactions.1516

Like we said, oxidation and reduction come in pairs: we can break them up into an oxidation part and a reduction part.1539

That is what we are doing: we are going to write two half-reactions.1545

OK, well, the oxidation--we know what the oxidation is: it is the iron going from this to this, so we write that.1551

We write: Fe2+ goes to Fe3+ (that is the oxidation part).1559

And the reduction part--the reduction: well, manganese is going from +7 to +2, so we write the whole species: MnO4- becoming Mn2+.1574

That is the first step: identify the oxidation half-reaction; identify the reduction half-reaction.1592

Decide what is being oxidized and what is reduced.1598

Iron is oxidized; manganese, reduced; OK.1600

Step 2 (let's move to the next page): Step 2 comes in 4 parts: Balance the elements like usual, like you always do; balance the elements except oxygen and hydrogen.1603

B: Balance oxygen, using H2O--I don't want to say "with"--balance the oxygen using H2O (oh, I had better write it out).1628

Balance hydrogen using H+.1655

And then, finally, balance the charge.1661

You are going to run through this, this, this, and this for each half-reaction; OK.1667

So, let's take the oxidation first.1673

Oxidation: Let's write our oxidation again; and we said that the oxidation reaction is Fe2+, going to Fe3+.1677

OK, let's run through the list: let's balance the elements, except oxygen and hydrogen--well, there is no hydrogen...Fe, Fe; there is one Fe...that is balanced.1688

There is no oxygen; there is no hydrogen; let's balance the charge: I have a +2 on this side; I have a +3 on this side--the only way I can balance the charge is by adding electrons to one side or the other until the charge balances.1696

Fe2+; this is 3+; I'm going to add an electron over here; 3+, minus 1, is 2+; 2+...I'm done.1711

This half-reaction is balanced.1720

Now, we'll do the reduction: the reduction reaction is...we said it is: MnO4- goes to Mn2+.1724

OK, let's balance the elements except hydrogen and oxygen: 1 Mn, 1Mn; we're good.1736

OK, that is done; balance the oxygen using H2O; there are 4 oxygens over here, so over here, I write 4 H2O--four oxygens.1744

Well, in the process of writing four oxygens by using water, I have introduced 8 hydrogens; so now, I balance the hydrogens over here.1755

I write 8 H+; now the hydrogens are balanced.1762

The last thing: balance the charge: I have 2+ on this side; I have 8+, minus 1 is 7+...so I have a total of 7+ on this side, and I have 2+ on that side; how can I balance these?1767

I can't add positive charge; I can only add negative charge, so I'm going to add five electrons to that side.1791

7, minus 5, is 2+; 2+; I'm done.1799

Now, that reaction...and now, the reduction half-reaction is balanced.1806

Part 3 (I'm going to do this in blue): Third step--multiply one or both of the half-reactions by an integer to equalize the number of electrons on each side.1811

There we go: so, multiply one or both of the half-reactions by an integer to equalize the number of electrons on each side.1856

Well, I have one electron on that side; I have 5 electrons on here; so I'm going to multiply this whole equation by 5.1865

I end up with the following (I'm going to do this in red): 5 Fe2+ goes to 5 Fe3+ + 5 electrons; this one I leave alone; I don't have to...because now, the electrons are balanced.1877

So, I have: 5 electrons, plus 8 protons (proton=hydrogen ion), plus permanganate ion, goes to manganese 2+, plus 4 waters.1897

Now, I add.1916

Oh, OK, so I have done that; now (let me write down...), the fourth part is: Add and cancel the identical species on each side.1926

Add the two half-reactions, and cancel identical species, just like you would an algebraic equation.1936

So, I have 5 Fe2+ goes to 5 Fe3+, plus 5 electrons; and I have 5 electrons, plus 8 hydrogen ion, plus a permanganate ion, going to a manganese 2+, plus 4 H2O.1950

I add, and I cancel species: 5 electron cancels 5 electron; that is why we did that--we want to cancel the electrons; we don't want electrons in our final answer.1980

We get: 8 H+ + 5 Fe2+ + MnO4- becomes 5 Fe3+ + Mn2+ + 4 H2O.1989

The last part, which...if you want to do it or not, it's up to you...check!2010

In other words, go back and make sure the charge balances; make sure the elements balance and the charges balance.2017

Trust me, this is correct.2025

OK, so here is the real question: what does the above equation say?2028

We were given an original equation, and we were asked to balance it; well, we assigned oxidation states, and we discovered that we had an oxidation-reduction process going on.2040

We separated it into an oxidation half-reaction and a reduction half-reaction; we went through the process of balancing each half-reaction; we made sure that the electron numbers were equal.2049

We added the equations, and then, when we added and canceled everything, we came up with this thing.2060

This is our balanced equation, under acidic conditions, for permanganate and iron reacting.2064

What does the above equation say? Here is what it says.2073

If I have a permanganate solution, and let's say that that permanganate solution--to it, I add a solution that contains iron 2+ ion (like, for example, FeCl2); well, if I mix those two solutions together, something is going to happen.2077

Well, something won't happen until I add some acid to make it an acidic reaction--until I drop the pH below 7.2096

Actually, well...yes, fairly below 7; when that happens--when I have a solution where I have hydrogen ion mixed with iron ion and permanganate ion, all of a sudden, electrons start to shuffle around and start to shift.2104

Bonds are broken; bonds are formed; here is what happens.2123

Each iron atom loses one electron; each of those five electrons are gained by manganese; manganese, in the process of gaining those electrons, doesn't want to be attached to oxygen anymore.2127

Well, oxygen wants to be attached to something, so it starts to grab some of the hydrogen ion.2142

What you end up with is 5 atoms of iron, each of which has now lost an electron; you end up with manganese that started with 7 electrons missing--now it only has 2 electrons missing, but is not attached to oxygen.2147

And then, oxygen went to bind with H to form 4 molecules of water.2163

That is what the above equation says.2169

You were given a process, and you had to balance it, in this case, under acidic conditions.2171

Manganese oxidized iron ion; iron ion reduced manganese.2180

Permanganate ion is called an oxidizing agent, because it oxidizes something--in this case, it oxidized the iron.2194

We can speak of permanganate oxidizing iron ion, or we can say that manganese oxidized iron ion, because it's manganese that is being reduced.2203

Oxygen isn't changing oxidation state--it's still -2; but, in the chemical reaction, the species that actually was dropped in there and was floating around was the permanganate.2212

That is fine, as long as you understand what is going on.2224

Something is oxidized; something is reduced; it often happens in these oxidation-reduction processes that things get shifted around; bonds are broken; bonds are formed...which is why things look so complicated.2229

But don't let that fool you; it is only a transfer of electrons.2241

5 iron ions, each of which lost one electron--those 5 electrons all went to manganese; manganese dropped from a +7 to a +2.2247

In the process of dropping from a +7 to a +2, it released its oxygen; the bond with oxygen broke.2258

Well, fine; so that stayed manganese +2; it's very, very stable.2264

In the process of losing the electrons, those 5 iron ions ended up going to become iron (3) ions; now, they are stable--they are fine.2269

In the process of breaking the bond with manganese, those oxygens needed something to attach to; they ended up using the acid, the free hydrogen proton--the hydrogen ion floating around--to form water.2280

This process will not take place until acid is added--that is the thing.2293

That is why this oxidation-reduction reaction takes place under acidic conditions.2298

What we are left with is this beautiful, balanced reaction to let us know what is going on.2303

When you look at this reaction, what is happening is exactly what it says is happening.2309

Bonds are being broken; electrons are being transferred; and you end up with this state right here.2315

I hope that made sense.2326

OK, I'm actually going to stop this lesson here; next lesson, I will actually do another balancing reaction, except this time, instead of acidic conditions, I want to do basic conditions.2327

I would recommend that you sort of go through this process again; make sure that you understand what is happening, but really, more than anything else, make sure you understand what is happening--the chemistry--what is going on.2338

It isn't just symbols that are being dropped out of the sky; there is a very, very real occurrence here.2350

It is just a transfer of electrons and the breaking and reshuffling of bonds.2355

Until next time, thank you for joining us here at Educator.com.2359

Take care; goodbye.2362

Educator®

Please sign in for full access to this lesson.

Sign-InORCreate Account

Enter your Sign-on user name and password.

Forgot password?

Start Learning Now

Our free lessons will get you started (Adobe Flash® required).
Get immediate access to our entire library.

Sign up for Educator.com

Membership Overview

  • Unlimited access to our entire library of courses.
  • Search and jump to exactly what you want to learn.
  • *Ask questions and get answers from the community and our teachers!
  • Practice questions with step-by-step solutions.
  • Download lesson files for programming and software training practice.
  • Track your course viewing progress.
  • Download lecture slides for taking notes.