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Raffi Hovasapian

Raffi Hovasapian

Cell Potential

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 (5)

0 answers

Post by Christian Fischer on May 17, 2014

Hi Raffi . I have a question wioth respect to electrochemistry which seems like a paradox for me, and I thought you - with your big brain  - might be able to guide me in the right direction of understanding it. Here it comes:

The symbol of charge is Q but the SI unit of charge is coulumb which is the Charge of approximately 6.241×1018 electrons. But charge is not itself defined, only in terms of Coulomb, and coulomb is defined in termes of Charge. Its SI definition of Coulomb is the charge transported by a constant current of one ampere in one second:

1C=1A*1s = (q/s)*s=q= charge,

Here is the question
It seems to me that Coulumbs are defined in termes Amps which are defined in terms of charge but charge itself is not a unit of measurement, so how is it possible to define coulumbs and amps in terms of charge when charge is a property and not something we can measure? Charge is part of the equation for Amps A=q/s and Coulumbs=1A*1s = (q/s)*s=q= charge, How does it make sense?

3 answers

Last reply by: Professor Hovasapian
Wed May 14, 2014 1:35 AM

Post by Rafael Mojica on May 2, 2014

Hello Raffi Hovasapian,

I carefully looked into my chart and the only rxn for Mn that i was able to find was Mn --  Mn2+ +2e. How can i manipulate that equation? because i was not able to find the specific one with the hydrogen and the oxide.

Cell Potential

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
  • Cell Potential 2:08
    • Definition of Cell Potential
    • Symbol and Unit
    • Standard Reduction Potential
    • Example Figure 1
    • Example Figure 2
    • All Reduction Potentials are Written as Reduction
    • Cell Potential: Important Fact 1
    • Cell Potential: Important Fact 2
    • Cell Potential: Important Fact 3
    • Cell Potential: Important Fact 4
    • Example Problem 1
    • Example Problem 2

Transcription: Cell Potential

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

Today, we are going to introduce a very, very, very important notion--probably the central notion, obviously, of electrochemistry; and it's called cell potential.0004

Last time, we talked about galvanic cells and how, if you mix two species and those species have a tendency to...if there is going to be some sort of an oxidation-reduction reaction, then electrons are going to spontaneously flow from one of those species to the next.0012

The galvanic cell exploits that tendency by separating the species, connecting the species with a wire or with a circuit, and then having the electrons flow through that wire.0033

Well, what you end up doing is (what we have done there is) create a battery.0045

Now, if we cut that wire and put something in between that wire (a heater, a cell phone, a computer, whatever), we can actually use that spontaneous flow of electrons to do work for us.0049

Again, that is all a battery is: it's a galvanic cell where the oxidizing agent and the reducing agent are separated, and the minute you actually put that into some device, you have closed the circuit.0061

When you flip that thing on, that is when you have actually close the circuit, and now electrons can flow, and you can operate your phone, operate your computer...whatever it is that you need to do.0073

OK, well, we want to be able to assign some numerical values to this; like, for example, if I put this species and this species together, are the electrons going to flow quickly? Are they going to flow slowly?0083

How strong is the tendency of electrons to flow?--we want to be able to control this, because if electrons are just going to sort of trickle through, that is not really going to be much use to us.0097

And if they are going to fly through at high speeds, well, they're going to end up doing damage to the material.0108

So, we need to know how to control this; that is the whole idea behind science--science is about understanding nature so that we can exercise control over nature, or at least control over the things that make our lives better; that is the whole idea.0113

OK, so let's start with some definitions, and we'll hopefully get a better sense of what this thing called a galvanic cell does and how it works.0128

We are going to define something called a cell potential.0137

So, definition, and cell potential (or it's also called electromotive force)...now, before I actually write this definition, let me tell you what we mean by the word "potential," when we talk about science.0142

It's exactly what you think it is: when we say something has potential, it hasn't happened yet, but it can happen.0168

So, for example, if you are on top of a mountain, skiing, and you are right there at the edge, you have the potential to go very, very, very fast.0176

But, you haven't dropped off the edge yet and started your skiing; so it's a measure of what could happen--that is what it is.0188

We can actually assign numerical values to what could happen, once we open the circuit, in this case--once we drop onto the mountain, once we open the faucet.0196

That is actually a good way to think about this cell potential; so I'm going to write the definition, and then I'm going to go back to this faucet idea and talk about it; and I think it's a good way to think about it.0207

It is the driving force or pull or push (depending on how you want to think about it) which causes electron flow in a galvanic cell.0217

You have a faucet at home: well, the water in the line is under pressure; you know this, because when you turn it on, water actually comes out.0246

When the faucet is closed, no water is coming out; however, the water company is sending water to your home; they are pressurizing the water, so there is a certain push against the valve in your faucet.0256

That is the whole idea: when you open it, you actually unleash that pressure.0271

Well, the pressure is the potential; the water has the potential to come out with a certain amount of force when you open the valve.0275

Here, the electrons have the potential to flow across that wire when you close that circuit.0284

It is actually a measure of the push or the pull, depending on which direction you want to see it; it's what drives the electrons forward.0291

Sometimes the drive is high; sometimes the drive is low; we are going to be assigning numerical values.0299

But don't let this word fool you: it will often talk about the potential, the cell potential; when we have a galvanic cell, what is the numerical value that we can assign to how badly the electrons want to go from the reducing agent to the oxidizing agent?0305

That is what this is: electromotive force.0324

That is a little bit more descriptive: "electromotive" means it's the force that is moving the electrons.0326

How powerful is the force that is moving the electrons--is it pushing the electrons through that wire really fast, or is it just sort of nudging them through?0332

That is the whole idea; OK.0341

The unit...well, actually, let's do the symbol first: the symbol for...let me see; maybe I'll do...yes, that's fine; I'll do the symbol; it doesn't matter which order.0343

The symbol is this: it's just E, and the cell--that is the potential of the cell.0360

The unit is the volt.0370

It is equivalent to Joules per coulomb.0375

We won't worry about...we have seen Joules before; that's a unit of energy; coulomb is a unit of charge.0382

Don't worry about this unit...really, you just need to concern yourself with this V, volt, for the time being.0389

We will actually get into...when we discuss electrolysis, we will actually discuss what the volt really is, and what coulomb is, and things like that.0396

But, for now, just know that it's a unit, and it's a measure of some ability for something to push or pull those electrons, to move them.0404

When we talk about 22 volts, that is what we are talking about: we are talking about the potential to actually move electrons.0413

OK, so now, let's take a look at a galvanic cell, and let's see how we are actually going to measure the cell potential.0423

We are going to take one of the cells that we have already looked at before, and we are going to create...we're going to put a little something in between here--something called a voltmeter or a potentiometer (a digital voltmeter, actually).0432

It is going to give us some number, and it is going to tell us--give us a numerical value for the potential, for the strength of that push or pull of the electrons.0445

This is connected this way; we have our two electrodes (oops, let me do this; OK)--so we have our electrodes, and we have a zinc solution here, and we have a copper solution here.0455

This is copper metal (because they are both metals, we can go ahead and use them as the electrodes); this is zinc metal.0474

Now, what happens here is: electrons are going to flow this way.0485

In other words, zinc metal is going to dissolve; it's going to turn into zinc ion and go into solution.0492

Copper ion is going to meet at the interface where the electrons are coming, and it's going to start to form copper.0497

So, this is going to be the oxidation; this is going to be the anode; this is going to be the cathode.0505

Now, when we actually run this--when we put some digital voltmeter in between here and everything is connected, but we don't let any current actually flow (in other words, we don't let any electrons do any flowing), what we want to measure is the potential for flow; what is the pressure?0509

These electrons are sort of building up; we don't want to open up the circuit just yet--what we want to measure is the pressure at that point.0532

Well, here is what happens: if you were to take this measurement, you would get a reading of 1.10.0539

So, the cell potential for this particular galvanic cell, made up of zinc solution, zinc metal electrode and copper solution, copper metal electrode, is 1.10 volts.0545

That is it: 1.10 volts...for right now, it's just a numerical value.0563

So again, what we do is: we put this digital voltmeter in there, and we want to measure the tendency--how badly these electrons want to get over here.0569

The higher this number, the higher the pressure; the higher this number, the more badly the electrons want to flow.0579

That is the idea.0586

It's a measure of how strong the push or the pull is; we don't let anything flow just yet.0590

If we were to open the circuit, yes, the electrons would start flowing; zinc would start to become oxidized; copper ion would start to be reduced, and the circuit would be closed; everything would be fine.0595

But right now, we are just concerned with the potential of this cell to do work.0605

We are not concerned with the actual work yet; we will be.0610

OK, now let's define something called a standard reduction potential.0615

Let me write the definition, and then we will explain what is going on.0634

It is the potential of a given species to become reduced by oxidizing hydrogen gas to hydrogen ion, when all species are in their standard states.0641

"Standard states" means 1 Molar solute concentrations, and 1 atmosphere for gases.0689

Now, let me tell you what this exactly means: in order to be able to actually measure something in science, we have to have a standard by which to measure it.0705

For example, in order to know that something weighs 27 grams, we have to know what 1 gram is.0719

Well, 1 gram is not some objective measure; it isn't just one gram that just fell from the sky; we actually have to decide what we mean--as a scientific community, what we mean by one gram.0724

So, in some sense, it's arbitrary; it's not arbitrary--there is a reason for choosing a particular measure--but we actually have to choose a point of reference, a standard against which to measure.0736

That is the whole idea: if I said, "This is how much an inch is," once you know what an inch is, you can go ahead and measure this based on the standard.0746

Well, the scientific community has decided that this reaction, the reaction of...well, they have decided that the hydrogen atom being oxidized to hydrogen ion, or hydrogen ion being reduced the hydrogen atom--they have assigned that a reduction potential of 0 volts.0753

Now, let me draw what I just said, and it will make sense what I actually just said.0782

OK, if I take a cell like this--a standard cell--this is going to look slightly different, because now we're going to be dealing with hydrogen gas.0788

Hydrogen gas--I'm going to have to bubble it in, so the arrangement is going to look different; but it's exactly the same.0798

We have our porous disk; we have our little digital voltmeter; we have our electrode (and here I'm going to go ahead and use copper).0805

OK, so if I put some copper here, and if I put some copper ion over there (the anion doesn't matter)...and now, on this side, watch this little arrangement.0817

Here is a little tube, and hydrogen gas is being pumped in this way; so the hydrogen gas is going this way, and it's actually bubbling out.0831

Well, the wire--there is a wire that goes down through this tube, and at the end of that wire is a little platinum electrode.0841

OK, it's a little platinum electrode; so the little platinum electrode--and here, we have a bunch of hydrogen ion--basically just an acidic solution.0849

This hydrogen ion is in contact with this electrode.0860

Well, hydrogen gas is being bubbled on top--so basically, we are flooding this electrode with hydrogen gas from above.0863

There is no liquid in this tube: we are just bubbling in--we are pushing in hydrogen gas, and it is bubbling out from underneath.0871

That is what is happening: so what is happening is that this electrode is being saturated with hydrogen gas.0879

Now, when I do this, and I measure this cell potential, I'm going to get a number, 0.34.0888

Well, again, we need to be able to assign certain numbers to certain cell potentials.0898

If I take this one side to always be hydrogen gas/hydrogen ion solution, and if I just change this species (zinc, permanganate, copper, manganese, cobalt, iron, whatever), I actually, by using hydrogen as my standard that I set at 0--as it turns out, I can actually write a cell potential for this reaction.0907

Here, what is happening is: electrons are flowing this way.0933

So, what is happening in this is: copper 2+ ion is gaining 2 electrons to become copper metal.0940

OK, let me write it a little smaller, because to the right, I want to write its standard reduction potential.0950

Copper ion, plus 2 electrons, becomes copper metal.0956

The standard reduction potential, which is just E (not cell)--the standard reduction potential for copper 2+...0963

You know what, I need more room; this is not going to work; OK, I'm going to write it right below: copper 2+ plus 2 electrons goes to copper solid, copper metal.0975

The cell potential for copper 2+ reducing to copper metal is equal to 0.34 volts.0990

By choosing hydrogen as 0, by international agreement, it gives us the standard reduction potential for a given species.0999

So, mind you, there are two things going on: we defined something called a cell potential--that is the potential for the whole cell; we also defined something called a standard reduction potential--this is the potential for a given species to become reduced, relative to the hydrogen electrode.1010

That is the whole idea: we have chosen hydrogen as 0 volts; therefore, we can assign a value, based on this cell--that means copper, in going from copper 2+ to copper--it has a standard reduction potential (this is a positive number) of .34 volts.1032

It is a measure of how badly copper wants to become reduced.1051

That is all that is going on here: we have chosen hydrogen as our standard; we set it equal to 0; because copper is the one being reduced, in this case, the oxidation that is taking place is the following.1057

H2 gas is losing...and it's becoming 2 H+ + 2 electrons.1068

Here is what is happening: hydrogen gas is being bubbled in here; when hydrogen gas hits this electrode, this electrode--because electrons are being pulled this way by .34 volts--every hydrogen molecule that passes over this electrode is split in half.1076

Each one of those electrons from each hydrogen atom is ripped off; those two electrons travel through the wire; they come down here; the copper 2+ ion takes those two electrons and becomes copper metal.1095

That is what is happening.1109

Hydrogen gas is turning into hydrogen ion; now you have two more hydrogen ions going into solution; that is what happens when we run this cell.1112

You get a positive value: spontaneously, between copper and hydrogen, copper will take the electrons; hydrogen will give up the electrons spontaneously.1119

If you put copper in the presence of hydrogen gas, that is what will happen.1130

Now, let's run another...and the cell potential for copper is .34 volts.1134

Now, let's do another one: let's do the same thing--we are going to have the same setup on the right, because that is our standard: a hydrogen electrode is our standard; we are going to pump in hydrogen gas.1142

We are going to have a platinum electrode down at the bottom; it's going to be saturated with this hydrogen gas as the hydrogen gas bubbles all over it.1156

We have a porous disk; we have some hydrogen ion, except this time, I have zinc in here.1166

I have this; I have my digital voltmeter; I have my zinc metal--this is zinc metal; well, something very interesting happens in this case.1176

Now, electrons flow this way, as it turns out, spontaneously.1187

Electrons flow this way: zinc metal gives up 2 electrons; they travel through the wire.1192

It comes over here; one H+ grabs an electron to become a hydrogen atom; another H+ grabs an electron to become a hydrogen atom; it turns into hydrogen gas; it bubbles off as hydrogen gas.1200

Now, you are forming hydrogen gas; the zinc melts; when I measure this potential, before I actually close the circuit, I end up with this: -0.76.1214

So here, the reaction that takes place is: Zn2+...let me see...+ 2 electrons equals -0.76 volts.1225

Now, watch: see how I have written this.1250

I have written this as a reduction, but what is happening to zinc is not a reduction.1254

Zinc is being oxidized; what is actually happening in this cell is this thing.1259

Zinc is becoming zinc ion, plus 2 electrons; however, when we said a standard reduction potential--all standard reduction potentials are written as reductions.1264

So, when I flip this equation around, that is why this actually has a negative value--because, relative to hydrogen, in this case, it isn't hydrogen that is oxidized; it is the hydrogen ion that is reduced.1278

Zinc gets oxidized; so here, the reaction is this...let me correct this...2 electrons...goes to zinc metal...1293

So, this reaction is what actually takes place: and because we have automatically assigned this 0 volts, this -.76 is the standard reduction potential for zinc ion.1318

That is the whole idea: all standard reduction potentials are written as reductions--that is why they are called reduction potentials.1334

It is the potential for this species to reduce.1341

But, because we have set some species (in this case, hydrogen) as our standard...as it turns out, when hydrogen is in contact with certain species, it will end up being oxidized.1344

When it's in contact with other species, it will actually be reduced.1354

If it's oxidized, your reduction potential for that species is going to be positive; if hydrogen is what is reduced, and the other thing is oxidized, the standard reduction potential is going to be negative.1359

Standard states: I forgot that little 0 on top: that means standard states.1374

That means 1 Molar solution, 1 atmosphere pressure.1378

-0.76...that is what is going on here.1382

Now, let us write what it is that we just did here: All reduction potentials are written as reductions; that is the whole idea.1387

We need a standard by which to measure them, which is why we write them as reductions.1409

We have the following: there is a table of reduction potentials; it's in your book; it's in the back of your book.1419

It is going to be on the AP exam: these are not things you have to memorize, but you have to understand what they mean when you look at them.1429

It is just like any thermodynamic table data or Ksp data: it gives you a numerical value for how strong a tendency any given species has to be reduced, relative to the hydrogen electrode, which we have set at 0 volts.1435

A partial view of a standard reduction potential looks like this: you will see copper 2+, plus 2 electrons, goes to copper; you will see this: equals 0.34 volts.1453

That is what it says; you will see: 2 H+ + 2 electrons goes to H2 gas.1469

You will see 0.00 volts--that is our standard; there are going to be certain numbers that are going to be higher; there are going to be certain numbers that are going to be lower.1478

The numbers that are higher--they will reduce this; the numbers that are lower--they will be oxidized by hydrogen.1487

Now, if I don't use hydrogen--if I just do something here and something here--the numbers that are higher will reduce; the numbers that are below--the species--those will be oxidized.1498

We'll explain in just a minute what we mean.1509

We also have: Zn2+ + 2 electrons goes to Zn = -0.76 volts.1512

OK, positive reduction potential; 0 reduction potential; negative reduction potential.1523

Between this and this--because this is positive, this will happen spontaneously; between this and this--because this is negative, what is spontaneous is this one; that means this one has to be reversed.1534

However, in a table that we look at, all of them are written as reductions; notice, all of the electrons are on the left.1542

It shows the ion species gaining electrons to become another species--reduction potentials.1549

OK, so let's see--a couple of things we should know about these: as we said before, this 0 little superscript--that means standard reduction potential...standard states.1555

All solutes (in other words, all ionic compounds) are at 1 Molar concentrations, and all gases are at 1 atmosphere.1573

We are bubbling in hydrogen gas at 1 atmosphere pressure--not 5 atmospheres; not .6 atmospheres; 1 atmosphere pressure.1592

The concentration of hydrogen ion in that solution: 1 molarity.1600

The concentration of zinc ion: 1 molarity; that is how we take this measurement.1604

OK, here are the important things: the higher the reduction potential, the greater the tendency to be reduced.1609

Copper has a greater tendency to be reduced than hydrogen ion; hydrogen ion has a greater tendency to be reduced than zinc.1641

Copper has a greater tendency to be reduced than zinc.1648

Between two species, the one with the higher standard reduction potential will become reduced, and the other will be oxidized.1658

Therefore, it must be flipped.1697

Copper and zinc: if I create a galvanic cell with copper and zinc...copper: .34; zinc is -.76; this has a higher reduction potential than this, so this will be reduced; this stays as written.1707

Because this is reduced, this is going to be oxidized; I have to flip this equation, and upon flipping this equation, I reverse the sign of this.1721

That is one of the problems that we are going to do right now.1729

OK, a couple of more things before we get to the example: the standard potential for a cell--the standard cell potential (this is the potential of the whole cell), which is symbolized E0cell, comes from adding the standard reduction potentials for each half-reaction.1733

Oxidation half, reduction half: remember, you break up an oxidation-reduction into two; each one of those has a standard reduction potential.1783

We add the equations (oxidation and reduction); we add the cell potentials; that gives us the...we add the reduction potentials for each species; that gives us the total cell potential.1791

OK, one thing you have to keep in mind, though, when doing this--very, very important: When multiplying a half-reaction by an integer to equalize electron number (remember when we were equalizing electron numbers so that we can actually cancel them?), do not multiply the standard reduction potential by that number--by that integer.1805

Remember when we did enthalpy?--if we multiply an equation by an integer, we have to multiply the enthalpy by that integer.1859

It is because enthalpy is an extensive property: it depends on the amount of material that we are dealing with.1867

Standard reduction potential is an intensive property: it doesn't matter how much--it doesn't change.1880

For example, mass is an intensive property: the more of something, the greater the mass.1887

The density...1892

Mass is not an intensive property; mass is an extensive property.1893

Density is intensive--it doesn't matter how much of something I have--whether it's this much gold or that much gold--gold has one density.1898

It is a property of the material: it doesn't depend on how much of that material is present.1906

Standard reduction potential is intensive: I don't multiply it by anything, just because I do five of those reactions as opposed to one of those reactions.1911

In other words, that is an intensive property--one that does not depend on quantity.1920

One that does is called extensive.1943

OK, so let's do an example.1949

OK, so let's see: Given the following cell and data, give the balanced cell reaction; state the anode; state the cathode; and calculate the cell potential; OK.1955

For the following cell and data, give the cathode, the anode, the balanced reaction, and the standard cell potential.1972

OK, so we have some cell, and we have some electrodes connected; this one is going to be aluminum; this one is going to be magnesium; and here is our data.2006

We have some magnesium ion here; we have some aluminum ion here; we have the following data.2025

Aluminum 3+, plus 3 electrons, goes to aluminum, and the standard reduction potential for that is -1.66 volts (straight out of a reduction potential table that you will be using all of the time with electrochemistry problems).2032

All of them are written as reductions; remember, all standard reduction potentials--that is why they are called "reduction potentials."2052

The equations are written as reductions; we decide, depending on which is higher and which is lower, what stays reduction and what becomes oxidized.2058

Magnesium 2+, plus 2 electrons, goes to magnesium: the reduction potential of that is (ooh, look at these crazy lines; OK) -2.37 volts.2069

All right, so we have -1.66 volts, and we have -2.37 volts.2086

All right, they are both negative, but the aluminum reaction has a higher reduction potential than this.2092

-1.66 is a higher number than -2.37.2103

So, the aluminum will stay as written--it will be reduced; the magnesium is going to end up being oxidized; so we have to flip the magnesium equation.2107

That is how we decide: we look at the reduction potentials: the one that is higher stays as a reduction; the other one gets flipped as reduced.2117

Let's write that: so we are going to write the reduction--the reduction is: Aluminum 3+, plus 3 electrons, goes to aluminum; and our standard reduction potential is -1.66 volts.2125

Our oxidation (which we had to flip) is going to be magnesium going to magnesium ion, plus 2 electrons; and because we actually flipped the equation, we reverse the sign of the reduction potential: it becomes 2.37 volts.2146

OK, I tend to put brackets around those.2165

Now, I need to balance the reaction; well, I have an oxidation-reduction--a reduction reaction, an oxidation reaction--I have the standard reduction potentials (oh, this vocabulary!); now I need to equalize the electrons.2169

I multiply this equation by 2 and this equation by 3, and I end up with 3 Al3+...no, 2 (I can't even do basic arithmetic!)...2 Al3+ + 6 electrons goes to 2 Al.2183

And again, remember: we don't change anything: -1.66 volts--that is an intensive property.2204

3 Mg goes to 3 Mg2+ + 6 electrons; this is 2.37 volts (I actually like to put a positive sign there).2211

And now, we add: we add the equation to balance; we add the standard reduction potentials to get our cell.2226

6 electrons goes with 6 electrons; 2 Al3+ + 3 magnesium goes to 2 Al + 3 magnesium 2+; the E of the cell is equal to...well, when I add those two, I get 0.71 volts.2233

0.71 volts: that is what happens.2255

When I put aluminum and magnesium in the cell, the way I described a little bit earlier, aluminum will pull 6 electrons from 3 magnesium atoms.2262

Aluminum will turn into solid aluminum; magnesium, upon losing electrons, will turn into magnesium ion; and the driving force, the pressure behind this process, is .71 volts.2272

Again, don't worry if you have a sense of what "volt" means; we will get to that a little bit later; but that is it--we can assign a numerical value to how strong this process is once we close the circuit.2286

Anode--oxidation: oxidation takes place in the magnesium compartment; reduction--cathode: cathode-reduction takes place in the aluminum compartment.2301

The balanced reaction; the cell potential; good.2315

Let's do the other one in blue--Example 2: OK, a galvanic cell is based on the following reaction.2320

MnO4- + H+ + ClO3- becomes ClO4- + Mn2+ + H2O.2347

OK, a galvanic cell is based on the following reaction; OK, our problem is to calculate the standard cell potential for this reaction.2370

Well, OK: let's take a look at what we have.2393

This is balanced as written; we can double-check that--it's not a problem--but it is balanced as written, because you see the H+; you see the H2O; everything looks like it is done.2399

Permanganate--manganese is being reduced; it's going from positive 7 to positive 2.2408

Chlorine is being oxidized.2417

Let me actually do this one: 2x3 is 6; this is going to be +5; this is going to be +7; this is going from a positive 5 state to a positive 7, so it's being oxidized.2421

As written, our permanganate is being reduced, and our chloride is being oxidized; so we can either read it off, or...let's go ahead and take a look at the half-reactions, the way we have been doing so far.2430

When we look up the half-reactions in a reduction potential chart (a table of reduction potentials), we get the following.2444

We get: MnO4- + 5 electrons + 8 H+ (this is exactly what it says in the chart--this is what you are looking for) goes to Mn2+ + 4 H2O.2452

It says that the reduction potential--standard reduction potential--is 1.51 volts.2471

OK, now, the other species that we notice in there is ClO4.2476

Well, this one, plus 2 H+, plus 2 electrons, goes to ClO3- (remember, everything is written as a reduction; electrons are always on the left-hand side; everything is written as a reduction in a standard reduction potentials chart), plus H2O.2485

The standard reduction potential for this one is 1.19 volts.2507

OK, so just by looking at these: this has a higher reduction potential than that; that means this stays as is; this reaction gets reversed.2511

That means this is going to be oxidized to that.2519

So, when we do that, we reverse that; so let's do it.2522

We are going to write: MnO4- + 5 electrons + 8 hydrogen ions goes to Mn2+ + 4 H2O; and we leave the reduction potential as is, equal to 1.51 volts.2527

This one we reverse--we write: ClO3- + H2O goes to ClO4- + 2 H+ + 2 electrons; and the reduction potential becomes -1.19 volts.2546

Now, I need to make sure that the electrons balance; so I'm going to multiply this equation by 2, and I'm going to multiply this equation by 5; I'm going to rewrite what I have.2563

I have: 2 MnO4- + 10 electrons + 16 hydrogen ion → 2 Mn2+ + 8 H2O; and still, nothing changes as far as the reduction potential (1.51 volts).2576

Here, I have 5 ClO3- + 5 H2O goes to 5 ClO4- + 10 H+ + 10 electrons.2595

Reduction potential: -1.19 volts; remember, we reversed that.2608

Now, let's add this: 10 electrons cancels 10 electrons; 10 H+ leaves 6 H+; 5 H2O leaves 3 H2O; and I end up with the following.2613

2 permanganate ions, plus 6 hydrogen ions, plus 5 chlorate ions, produce (if I were to close the circuit) 2 manganese, plus 5 perchlorate, plus 3 H2O.2629

The standard reduction for that cell is equal to 0.32 volts.2659

There we go: if I create a cell based on permanganate and chlorate, the potential for that cell is .32.2666

That means that is a measure of the tendency for this reaction to happen, if I were to open the circuit.2676

That is what is going on.2684

Now, notice all of these species: that is an ion; that is an ion; that is an ion; and that is an ion.2686

When I actually draw out the physical arrangement for this thing, here is what it looks like.2696

The cell itself is going to look like this: I'm going to have my digital voltmeter...and because they are both ions, I'm going to have ClO3- (Cl...well, I'm going to leave this as ClO3- for right now), and I'm going to put the MnO4- in there.2705

This is platinum; none of these species that has been oxidized or reduced actually becomes a metal, so I can't really use that as an electrode.2737

Therefore, I have to provide some surface for this chemistry to take place.2746

Again, the MnO4 will go to the surface; the electrons will come, and they will join together on that surface.2750

It provides a platform, a meeting place, for the two species.2756

This is also going to be a platinum electrode here.2760

This is going to measure 0.32.2765

OK, when we open the circuit (so again, this is measuring the pressure that the electrons--the extent to which the electrons want to go this way; when I open the faucet here, that is when the electrons are going to start to flow), here is what happens.2768

The ClO3 turns into ClO4-, so ClO4- starts to show up over here.2793

This starts to go away.2802

And, MnO4- starts to turn into Mn2+.2804

This starts to go away; Mn2+ starts to show up in solution.2811

That is what is going on.2816

We can set up a standard reduction potential for any species that we want, relative to the hydrogen electrode.2821

Some are going to be positive numbers; some are going to be negative numbers, because hydrogen, we set at 0; that was our choice (international agreement).2829

Because we have, now, a table of all of these reduction potentials, well, I can create any galvanic cell I want, just by mixing and matching species.2838

All I have to look at is which one has a higher reduction potential.2846

The one that has the higher reduction potential, between any two species that I choose--that is going to be reduced.2849

The other one--the equation has to be flipped, because that is going to be oxidized (right?--oxidation-reduction: they come in pairs).2855

I balance those half-reactions; I add them the way that I did for the acid-base section earlier, last lesson; when I add those, I add the reduction potentials, and that gives me the standard cell potential for that galvanic cell.2861

It gives me a measure of just how badly that cell wants to start pulling electrons.2881

The higher that number (the higher the cell potential), the more work I'm going to get out of that particular process.2888

That is what is going on.2897

That is what is important: it is really, really important that you actually understand what is happening, physically.2899

If you don't get this, none of the math will make sense.2905

Absolutely none of the math will make sense.2909

So, hopefully, think about this; think about what is going on; we will do some more problems later on.2912

Until then, thank you for joining us here at Educator.com.2917

We'll see you next time; goodbye.2920

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