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Franklin Ow

Franklin Ow

Electrochemistry

Slide Duration:

Table of Contents

I. Basic Concepts & Measurement of Chemistry
Basic Concepts of Chemistry

16m 26s

Intro
0:00
Lesson Overview
0:07
Introduction
0:56
What is Chemistry?
0:57
What is Matter?
1:16
Solids
1:43
General Characteristics
1:44
Particulate-level Drawing of Solids
2:34
Liquids
3:39
General Characteristics of Liquids
3:40
Particulate-level Drawing of Liquids
3:55
Gases
4:23
General Characteristics of Gases
4:24
Particulate-level Drawing Gases
5:05
Classification of Matter
5:27
Classification of Matter
5:26
Pure Substances
5:54
Pure Substances
5:55
Mixtures
7:06
Definition of Mixtures
7:07
Homogeneous Mixtures
7:11
Heterogeneous Mixtures
7:52
Physical and Chemical Changes/Properties
8:18
Physical Changes Retain Chemical Composition
8:19
Chemical Changes Alter Chemical Composition
9:32
Physical and Chemical Changes/Properties, cont'd
10:55
Physical Properties
10:56
Chemical Properties
11:42
Sample Problem 1: Chemical & Physical Change
12:22
Sample Problem 2: Element, Compound, or Mixture?
13:52
Sample Problem 3: Classify Each of the Following Properties as chemical or Physical
15:03
Tools in Quantitative Chemistry

29m 22s

Intro
0:00
Lesson Overview
0:07
Units of Measurement
1:23
The International System of Units (SI): Mass, Length, and Volume
1:39
Percent Error
2:17
Percent Error
2:18
Example: Calculate the Percent Error
2:56
Standard Deviation
3:48
Standard Deviation Formula
3:49
Standard Deviation cont'd
4:42
Example: Calculate Your Standard Deviation
4:43
Precisions vs. Accuracy
6:25
Precision
6:26
Accuracy
7:01
Significant Figures and Uncertainty
7:50
Consider the Following (2) Rulers
7:51
Consider the Following Graduated Cylinder
11:30
Identifying Significant Figures
12:43
The Rules of Sig Figs Overview
12:44
The Rules for Sig Figs: All Nonzero Digits Are Significant
13:21
The Rules for Sig Figs: A Zero is Significant When It is In-Between Nonzero Digits
13:28
The Rules for Sig Figs: A Zero is Significant When at the End of a Decimal Number
14:02
The Rules for Sig Figs: A Zero is not significant When Starting a Decimal Number
14:27
Using Sig Figs in Calculations
15:03
Using Sig Figs for Multiplication and Division
15:04
Using Sig Figs for Addition and Subtraction
15:48
Using Sig Figs for Mixed Operations
16:11
Dimensional Analysis
16:20
Dimensional Analysis Overview
16:21
General Format for Dimensional Analysis
16:39
Example: How Many Miles are in 17 Laps?
17:17
Example: How Many Grams are in 1.22 Pounds?
18:40
Dimensional Analysis cont'd
19:43
Example: How Much is Spent on Diapers in One Week?
19:44
Dimensional Analysis cont'd
21:03
SI Prefixes
21:04
Dimensional Analysis cont'd
22:03
500 mg → ? kg
22:04
34.1 cm → ? um
24:03
Summary
25:11
Sample Problem 1: Dimensional Analysis
26:09
II. Atoms, Molecules, and Ions
Atoms, Molecules, and Ions

52m 18s

Intro
0:00
Lesson Overview
0:08
Introduction to Atomic Structure
1:03
Introduction to Atomic Structure
1:04
Plum Pudding Model
1:26
Introduction to Atomic Structure Cont'd
2:07
John Dalton's Atomic Theory: Number 1
2:22
John Dalton's Atomic Theory: Number 2
2:50
John Dalton's Atomic Theory: Number 3
3:07
John Dalton's Atomic Theory: Number 4
3:30
John Dalton's Atomic Theory: Number 5
3:58
Introduction to Atomic Structure Cont'd
5:21
Ernest Rutherford's Gold Foil Experiment
5:22
Introduction to Atomic Structure Cont'd
7:42
Implications of the Gold Foil Experiment
7:43
Relative Masses and Charges
8:18
Isotopes
9:02
Isotopes
9:03
Introduction to The Periodic Table
12:17
The Periodic Table of the Elements
12:18
Periodic Table, cont'd
13:56
Metals
13:57
Nonmetals
14:25
Semimetals
14:51
Periodic Table, cont'd
15:57
Group I: The Alkali Metals
15:58
Group II: The Alkali Earth Metals
16:25
Group VII: The Halogens
16:40
Group VIII: The Noble Gases
17:08
Ionic Compounds: Formulas, Names, Props.
17:35
Common Polyatomic Ions
17:36
Predicting Ionic Charge for Main Group Elements
18:52
Ionic Compounds: Formulas, Names, Props.
20:36
Naming Ionic Compounds: Rule 1
20:51
Naming Ionic Compounds: Rule 2
21:22
Naming Ionic Compounds: Rule 3
21:50
Naming Ionic Compounds: Rule 4
22:22
Ionic Compounds: Formulas, Names, Props.
22:50
Naming Ionic Compounds Example: Al₂O₃
22:51
Naming Ionic Compounds Example: FeCl₃
23:21
Naming Ionic Compounds Example: CuI₂ 3H₂O
24:00
Naming Ionic Compounds Example: Barium Phosphide
24:40
Naming Ionic Compounds Example: Ammonium Phosphate
25:55
Molecular Compounds: Formulas and Names
26:42
Molecular Compounds: Formulas and Names
26:43
The Mole
28:10
The Mole is 'A Chemist's Dozen'
28:11
It is a Central Unit, Connecting the Following Quantities
30:01
The Mole, cont'd
32:07
Atomic Masses
32:08
Example: How Many Moles are in 25.7 Grams of Sodium?
32:28
Example: How Many Atoms are in 1.2 Moles of Carbon?
33:17
The Mole, cont'd
34:25
Example: What is the Molar Mass of Carbon Dioxide?
34:26
Example: How Many Grams are in 1.2 Moles of Carbon Dioxide?
25:46
Percentage Composition
36:43
Example: How Many Grams of Carbon Contained in 65.1 Grams of Carbon Dioxide?
36:44
Empirical and Molecular Formulas
39:19
Empirical Formulas
39:20
Empirical Formula & Elemental Analysis
40:21
Empirical and Molecular Formulas, cont'd
41:24
Example: Determine Both the Empirical and Molecular Formulas - Step 1
41:25
Example: Determine Both the Empirical and Molecular Formulas - Step 2
43:18
Summary
46:22
Sample Problem 1: Determine the Empirical Formula of Lithium Fluoride
47:10
Sample Problem 2: How Many Atoms of Carbon are Present in 2.67 kg of C₆H₆?
49:21
III. Chemical Reactions
Chemical Reactions

43m 24s

Intro
0:00
Lesson Overview
0:06
The Law of Conservation of Mass and Balancing Chemical Reactions
1:49
The Law of Conservation of Mass
1:50
Balancing Chemical Reactions
2:50
Balancing Chemical Reactions Cont'd
3:40
Balance: N₂ + H₂ → NH₃
3:41
Balance: CH₄ + O₂ → CO₂ + H₂O
7:20
Balancing Chemical Reactions Cont'd
9:49
Balance: C₂H₆ + O₂ → CO₂ + H₂O
9:50
Intro to Chemical Equilibrium
15:32
When an Ionic Compound Full Dissociates
15:33
When an Ionic Compound Incompletely Dissociates
16:14
Dynamic Equilibrium
17:12
Electrolytes and Nonelectrolytes
18:03
Electrolytes
18:04
Strong Electrolytes and Weak Electrolytes
18:55
Nonelectrolytes
19:23
Predicting the Product(s) of an Aqueous Reaction
20:02
Single-replacement
20:03
Example: Li (s) + CuCl₂ (aq) → 2 LiCl (aq) + Cu (s)
21:03
Example: Cu (s) + LiCl (aq) → NR
21:23
Example: Zn (s) + 2HCl (aq) → ZnCl₂ (aq) + H₂ (g)
22:32
Predicting the Product(s) of an Aqueous Reaction
23:37
Double-replacement
23:38
Net-ionic Equation
25:29
Predicting the Product(s) of an Aqueous Reaction
26:12
Solubility Rules for Ionic Compounds
26:13
Predicting the Product(s) of an Aqueous Reaction
28:10
Neutralization Reactions
28:11
Example: HCl (aq) + NaOH (aq) → ?
28:37
Example: H₂SO₄ (aq) + KOH (aq) → ?
29:25
Predicting the Product(s) of an Aqueous Reaction
30:20
Certain Aqueous Reactions can Produce Unstable Compounds
30:21
Example 1
30:52
Example 2
32:16
Example 3
32:54
Summary
33:54
Sample Problem 1
34:55
ZnCO₃ (aq) + H₂SO₄ (aq) → ?
35:09
NH₄Br (aq) + Pb(C₂H₃O₂)₂ (aq) → ?
36:02
KNO₃ (aq) + CuCl₂ (aq) → ?
37:07
Li₂SO₄ (aq) + AgNO₃ (aq) → ?
37:52
Sample Problem 2
39:09
Question 1
39:10
Question 2
40:36
Question 3
41:47
Chemical Reactions II

55m 40s

Intro
0:00
Lesson Overview
0:10
Arrhenius Definition
1:15
Arrhenius Acids
1:16
Arrhenius Bases
3:20
The Bronsted-Lowry Definition
4:48
Acids Dissolve In Water and Donate a Proton to Water: Example 1
4:49
Acids Dissolve In Water and Donate a Proton to Water: Example 2
6:54
Monoprotic Acids & Polyprotic Acids
7:58
Strong Acids
11:30
Bases Dissolve In Water and Accept a Proton From Water
12:41
Strong Bases
16:36
The Autoionization of Water
17:42
Amphiprotic
17:43
Water Reacts With Itself
18:24
Oxides of Metals and Nonmetals
20:08
Oxides of Metals and Nonmetals Overview
20:09
Oxides of Nonmetals: Acidic Oxides
21:23
Oxides of Metals: Basic Oxides
24:08
Oxidation-Reduction (Redox) Reactions
25:34
Redox Reaction Overview
25:35
Oxidizing and Reducing Agents
27:02
Redox Reaction: Transfer of Electrons
27:54
Oxidation-Reduction Reactions Cont'd
29:55
Oxidation Number Overview
29:56
Oxidation Number of Homonuclear Species
31:17
Oxidation Number of Monatomic Ions
32:58
Oxidation Number of Fluorine
33:27
Oxidation Number of Oxygen
34:00
Oxidation Number of Chlorine, Bromine, and Iodine
35:07
Oxidation Number of Hydrogen
35:30
Net Sum of All Oxidation Numbers In a Compound
36:21
Oxidation-Reduction Reactions Cont'd
38:19
Let's Practice Assigning Oxidation Number
38:20
Now Let's Apply This to a Chemical Reaction
41:07
Summary
44:19
Sample Problems
45:29
Sample Problem 1
45:30
Sample Problem 2: Determine the Oxidizing and Reducing Agents
48:48
Sample Problem 3: Determine the Oxidizing and Reducing Agents
50:43
IV. Stoichiometry
Stoichiometry I

42m 10s

Intro
0:00
Lesson Overview
0:23
Mole to Mole Ratios
1:32
Example 1: In 1 Mole of H₂O, How Many Moles Are There of Each Element?
1:53
Example 2: In 2.6 Moles of Water, How Many Moles Are There of Each Element?
2:24
Mole to Mole Ratios Cont'd
5:13
Balanced Chemical Reaction
5:14
Mole to Mole Ratios Cont'd
7:25
Example 3: How Many Moles of Ammonia Can Form If you Have 3.1 Moles of H₂?
7:26
Example 4: How Many Moles of Hydrogen Gas Are Required to React With 6.4 Moles of Nitrogen Gas?
9:08
Mass to mass Conversion
11:06
Mass to mass Conversion
11:07
Example 5: How Many Grams of Ammonia Can Form If You Have 3.1 Grams of H₂?
12:37
Example 6: How Many Grams of Hydrogen Gas Are Required to React With 6.4 Grams of Nitrogen Gas?
15:34
Example 7: How Man Milligrams of Ammonia Can Form If You Have 1.2 kg of H₂?
17:29
Limiting Reactants, Percent Yields
20:42
Limiting Reactants, Percent Yields
20:43
Example 8: How Many Grams of Ammonia Can Form If You Have 3.1 Grams of H₂ and 3.1 Grams of N₂
22:25
Percent Yield
25:30
Example 9: How Many Grams of The Excess Reactant Remains?
26:37
Summary
29:34
Sample Problem 1: How Many Grams of Carbon Are In 2.2 Kilograms of Carbon Dioxide?
30:47
Sample Problem 2: How Many Milligrams of Carbon Dioxide Can Form From 23.1 Kg of CH₄(g)?
33:06
Sample Problem 3: Part 1
36:10
Sample Problem 3: Part 2 - What Amount Of The Excess Reactant Will Remain?
40:53
Stoichiometry II

42m 38s

Intro
0:00
Lesson Overview
0:10
Molarity
1:14
Solute and Solvent
1:15
Molarity
2:01
Molarity Cont'd
2:59
Example 1: How Many Grams of KBr are Needed to Make 350 mL of a 0.67 M KBr Solution?
3:00
Example 2: How Many Moles of KBr are in 350 mL of a 0.67 M KBr Solution?
5:44
Example 3: What Volume of a 0.67 M KBr Solution Contains 250 mg of KBr?
7:46
Dilutions
10:01
Dilution: M₁V₂=M₁V₂
10:02
Example 5: Explain How to Make 250 mL of a 0.67 M KBr Solution Starting From a 1.2M Stock Solution
12:04
Stoichiometry and Double-Displacement Precipitation Reactions
14:41
Example 6: How Many grams of PbCl₂ Can Form From 250 mL of 0.32 M NaCl?
15:38
Stoichiometry and Double-Displacement Precipitation Reactions
18:05
Example 7: How Many grams of PbCl₂ Can Form When 250 mL of 0.32 M NaCl and 150 mL of 0.45 Pb(NO₃)₂ Mix?
18:06
Stoichiometry and Neutralization Reactions
21:01
Example 8: How Many Grams of NaOh are Required to Neutralize 4.5 Grams of HCl?
21:02
Stoichiometry and Neutralization Reactions
23:03
Example 9: How Many mL of 0.45 M NaOH are Required to Neutralize 250 mL of 0.89 M HCl?
23:04
Stoichiometry and Acid-Base Standardization
25:28
Introduction to Titration & Standardization
25:30
Acid-Base Titration
26:12
The Analyte & Titrant
26:24
The Experimental Setup
26:49
The Experimental Setup
26:50
Stoichiometry and Acid-Base Standardization
28:38
Example 9: Determine the Concentration of the Analyte
28:39
Summary
32:46
Sample Problem 1: Stoichiometry & Neutralization
35:24
Sample Problem 2: Stoichiometry
37:50
V. Thermochemistry
Energy & Chemical Reactions

55m 28s

Intro
0:00
Lesson Overview
0:14
Introduction
1:22
Recall: Chemistry
1:23
Energy Can Be Expressed In Different Units
1:57
The First Law of Thermodynamics
2:43
Internal Energy
2:44
The First Law of Thermodynamics Cont'd
6:14
Ways to Transfer Internal Energy
6:15
Work Energy
8:13
Heat Energy
8:34
∆U = q + w
8:44
Calculating ∆U, Q, and W
8:58
Changes In Both Volume and Temperature of a System
8:59
Calculating ∆U, Q, and W Cont'd
11:01
The Work Equation
11:02
Example 1: Calculate ∆U For The Burning Fuel
11:45
Calculating ∆U, Q, and W Cont'd
14:09
The Heat Equation
14:10
Calculating ∆U, Q, and W Cont'd
16:03
Example 2: Calculate The Final Temperature
16:04
Constant-Volume Calorimetry
18:05
Bomb Calorimeter
18:06
The Effect of Constant Volume On The Equation For Internal Energy
22:11
Example 3: Calculate ∆U
23:12
Constant-Pressure Conditions
26:05
Constant-Pressure Conditions
26:06
Calculating Enthalpy: Phase Changes
27:29
Melting, Vaporization, and Sublimation
27:30
Freezing, Condensation and Deposition
28:25
Enthalpy Values For Phase Changes
28:40
Example 4: How Much Energy In The Form of heat is Required to Melt 1.36 Grams of Ice?
29:40
Calculating Enthalpy: Heats of Reaction
31:22
Example 5: Calculate The Heat In kJ Associated With The Complete Reaction of 155 g NH₃
31:23
Using Standard Enthalpies of Formation
33:53
Standard Enthalpies of Formation
33:54
Using Standard Enthalpies of Formation
36:12
Example 6: Calculate The Standard Enthalpies of Formation For The Following Reaction
36:13
Enthalpy From a Series of Reactions
39:58
Hess's Law
39:59
Coffee-Cup Calorimetry
42:43
Coffee-Cup Calorimetry
42:44
Example 7: Calculate ∆H° of Reaction
45:10
Summary
47:12
Sample Problem 1
48:58
Sample Problem 2
51:24
VI. Quantum Theory of Atoms
Structure of Atoms

42m 33s

Intro
0:00
Lesson Overview
0:07
Introduction
1:01
Rutherford's Gold Foil Experiment
1:02
Electromagnetic Radiation
2:31
Radiation
2:32
Three Parameters: Energy, Frequency, and Wavelength
2:52
Electromagnetic Radiation
5:18
The Electromagnetic Spectrum
5:19
Atomic Spectroscopy and The Bohr Model
7:46
Wavelengths of Light
7:47
Atomic Spectroscopy Cont'd
9:45
The Bohr Model
9:46
Atomic Spectroscopy Cont'd
12:21
The Balmer Series
12:22
Rydberg Equation For Predicting The Wavelengths of Light
13:04
The Wave Nature of Matter
15:11
The Wave Nature of Matter
15:12
The Wave Nature of Matter
19:10
New School of Thought
19:11
Einstein: Energy
19:49
Hertz and Planck: Photoelectric Effect
20:16
de Broglie: Wavelength of a Moving Particle
21:14
Quantum Mechanics and The Atom
22:15
Heisenberg: Uncertainty Principle
22:16
Schrodinger: Wavefunctions
23:08
Quantum Mechanics and The Atom
24:02
Principle Quantum Number
24:03
Angular Momentum Quantum Number
25:06
Magnetic Quantum Number
26:27
Spin Quantum Number
28:42
The Shapes of Atomic Orbitals
29:15
Radial Wave Function
29:16
Probability Distribution Function
32:08
The Shapes of Atomic Orbitals
34:02
3-Dimensional Space of Wavefunctions
34:03
Summary
35:57
Sample Problem 1
37:07
Sample Problem 2
40:23
VII. Electron Configurations and Periodicity
Periodic Trends

38m 50s

Intro
0:00
Lesson Overview
0:09
Introduction
0:36
Electron Configuration of Atoms
1:33
Electron Configuration & Atom's Electrons
1:34
Electron Configuration Format
1:56
Electron Configuration of Atoms Cont'd
3:01
Aufbau Principle
3:02
Electron Configuration of Atoms Cont'd
6:53
Electron Configuration Format 1: Li, O, and Cl
6:56
Electron Configuration Format 2: Li, O, and Cl
9:11
Electron Configuration of Atoms Cont'd
12:48
Orbital Box Diagrams
12:49
Pauli Exclusion Principle
13:11
Hund's Rule
13:36
Electron Configuration of Atoms Cont'd
17:35
Exceptions to The Aufbau Principle: Cr
17:36
Exceptions to The Aufbau Principle: Cu
18:15
Electron Configuration of Atoms Cont'd
20:22
Electron Configuration of Monatomic Ions: Al
20:23
Electron Configuration of Monatomic Ions: Al³⁺
20:46
Electron Configuration of Monatomic Ions: Cl
21:57
Electron Configuration of Monatomic Ions: Cl¹⁻
22:09
Electron Configuration Cont'd
24:31
Paramagnetism
24:32
Diamagnetism
25:00
Atomic Radii
26:08
Atomic Radii
26:09
In a Column of the Periodic Table
26:25
In a Row of the Periodic Table
26:46
Ionic Radii
27:30
Ionic Radii
27:31
Anions
27:42
Cations
27:57
Isoelectronic Species
28:12
Ionization Energy
29:00
Ionization Energy
29:01
Electron Affinity
31:37
Electron Affinity
31:37
Summary
33:43
Sample Problem 1: Ground State Configuration and Orbital Box Diagram
34:21
Fe
34:48
P
35:32
Sample Problem 2
36:38
Which Has The Larger Ionization Energy: Na or Li?
36:39
Which Has The Larger Atomic Size: O or N ?
37:23
Which Has The Larger Atomic Size: O²⁻ or N³⁻ ?
38:00
VIII. Molecular Geometry & Bonding Theory
Bonding & Molecular Structure

52m 39s

Intro
0:00
Lesson Overview
0:08
Introduction
1:10
Types of Chemical Bonds
1:53
Ionic Bond
1:54
Molecular Bond
2:42
Electronegativity and Bond Polarity
3:26
Electronegativity (EN)
3:27
Periodic Trend
4:36
Electronegativity and Bond Polarity Cont'd
6:04
Bond Polarity: Polar Covalent Bond
6:05
Bond Polarity: Nonpolar Covalent Bond
8:53
Lewis Electron Dot Structure of Atoms
9:48
Lewis Electron Dot Structure of Atoms
9:49
Lewis Structures of Polyatomic Species
12:51
Single Bonds
12:52
Double Bonds
13:28
Nonbonding Electrons
13:59
Lewis Structures of Polyatomic Species Cont'd
14:45
Drawing Lewis Structures: Step 1
14:48
Drawing Lewis Structures: Step 2
15:16
Drawing Lewis Structures: Step 3
15:52
Drawing Lewis Structures: Step 4
17:31
Drawing Lewis Structures: Step 5
19:08
Drawing Lewis Structure Example: Carbonate
19:33
Resonance and Formal Charges (FC)
24:06
Resonance Structures
24:07
Formal Charge
25:20
Resonance and Formal Charges Cont'd
27:46
More On Formal Charge
27:47
Resonance and Formal Charges Cont'd
28:21
Good Resonance Structures
28:22
VSEPR Theory
31:08
VSEPR Theory Continue
31:09
VSEPR Theory Cont'd
32:53
VSEPR Geometries
32:54
Steric Number
33:04
Basic Geometry
33:50
Molecular Geometry
35:50
Molecular Polarity
37:51
Steps In Determining Molecular Polarity
37:52
Example 1: Polar
38:47
Example 2: Nonpolar
39:10
Example 3: Polar
39:36
Example 4: Polar
40:08
Bond Properties: Order, Length, and Energy
40:38
Bond Order
40:39
Bond Length
41:21
Bond Energy
41:55
Summary
43:09
Sample Problem 1
43:42
XeO₃
44:03
I₃⁻
47:02
SF₅
49:16
Advanced Bonding Theories

1h 11m 41s

Intro
0:00
Lesson Overview
0:09
Introduction
0:38
Valence Bond Theory
3:07
Valence Bond Theory
3:08
spᶟ Hybridized Carbon Atom
4:19
Valence Bond Theory Cont'd
6:24
spᶟ Hybridized
6:25
Hybrid Orbitals For Water
7:26
Valence Bond Theory Cont'd (spᶟ)
11:53
Example 1: NH₃
11:54
Valence Bond Theory Cont'd (sp²)
14:48
sp² Hybridization
14:49
Example 2: BF₃
16:44
Valence Bond Theory Cont'd (sp)
22:44
sp Hybridization
22:46
Example 3: HCN
23:38
Valence Bond Theory Cont'd (sp³d and sp³d²)
27:36
Valence Bond Theory: sp³d and sp³d²
27:37
Molecular Orbital Theory
29:10
Valence Bond Theory Doesn't Always Account For a Molecule's Magnetic Behavior
29:11
Molecular Orbital Theory Cont'd
30:37
Molecular Orbital Theory
30:38
Wavefunctions
31:04
How s-orbitals Can Interact
32:23
Bonding Nature of p-orbitals: Head-on
35:34
Bonding Nature of p-orbitals: Parallel
39:04
Interaction Between s and p-orbital
40:45
Molecular Orbital Diagram For Homonuclear Diatomics: H₂
42:21
Molecular Orbital Diagram For Homonuclear Diatomics: He₂
45:23
Molecular Orbital Diagram For Homonuclear Diatomic: Li₂
46:39
Molecular Orbital Diagram For Homonuclear Diatomic: Li₂⁺
47:42
Molecular Orbital Diagram For Homonuclear Diatomic: B₂
48:57
Molecular Orbital Diagram For Homonuclear Diatomic: N₂
54:04
Molecular Orbital Diagram: Molecular Oxygen
55:57
Molecular Orbital Diagram For Heteronuclear Diatomics: Hydrochloric Acid
1:02:16
Sample Problem 1: Determine the Atomic Hybridization
1:07:20
XeO₃
1:07:21
SF₆
1:07:49
I₃⁻
1:08:20
Sample Problem 2
1:09:04
IX. Gases, Solids, & Liquids
Gases

35m 6s

Intro
0:00
Lesson Overview
0:07
The Kinetic Molecular Theory of Gases
1:23
The Kinetic Molecular Theory of Gases
1:24
Parameters To Characterize Gases
3:35
Parameters To Characterize Gases: Pressure
3:37
Interpreting Pressure On a Particulate Level
4:43
Parameters Cont'd
6:08
Units For Expressing Pressure: Psi, Pascal
6:19
Units For Expressing Pressure: mm Hg
6:42
Units For Expressing Pressure: atm
6:58
Units For Expressing Pressure: torr
7:24
Parameters Cont'd
8:09
Parameters To Characterize Gases: Volume
8:10
Common Units of Volume
9:00
Parameters Cont'd
9:11
Parameters To Characterize Gases: Temperature
9:12
Particulate Level
9:36
Parameters To Characterize Gases: Moles
10:24
The Simple Gas Laws
10:43
Gas Laws Are Only Valid For…
10:44
Charles' Law
11:24
The Simple Gas Laws
13:13
Boyle's Law
13:14
The Simple Gas Laws
15:28
Gay-Lussac's Law
15:29
The Simple Gas Laws
17:11
Avogadro's Law
17:12
The Ideal Gas Law
18:43
The Ideal Gas Law: PV = nRT
18:44
Applications of the Ideal Gas Law
20:12
Standard Temperature and Pressure for Gases
20:13
Applications of the Ideal Gas Law
21:43
Ideal Gas Law & Gas Density
21:44
Gas Pressures and Partial Pressures
23:18
Dalton's Law of Partial Pressures
23:19
Gas Stoichiometry
24:15
Stoichiometry Problems Involving Gases
24:16
Using The Ideal Gas Law to Get to Moles
25:16
Using Molar Volume to Get to Moles
25:39
Gas Stoichiometry Cont'd
26:03
Example 1: How Many Liters of O₂ at STP are Needed to Form 10.5 g of Water Vapor?
26:04
Summary
28:33
Sample Problem 1: Calculate the Molar Mass of the Gas
29:28
Sample Problem 2: What Mass of Ag₂O is Required to Form 3888 mL of O₂ Gas When Measured at 734 mm Hg and 25°C?
31:59
Intermolecular Forces & Liquids

33m 47s

Intro
0:00
Lesson Overview
0:10
Introduction
0:46
Intermolecular Forces (IMF)
0:47
Intermolecular Forces of Polar Molecules
1:32
Ion-dipole Forces
1:33
Example: Salt Dissolved in Water
1:50
Coulomb's Law & the Force of Attraction Between Ions and/or Dipoles
3:06
IMF of Polar Molecules cont'd
4:36
Enthalpy of Solvation or Enthalpy of Hydration
4:37
IMF of Polar Molecules cont'd
6:01
Dipole-dipole Forces
6:02
IMF of Polar Molecules cont'd
7:22
Hydrogen Bonding
7:23
Example: Hydrogen Bonding of Water
8:06
IMF of Nonpolar Molecules
9:37
Dipole-induced Dipole Attraction
9:38
IMF of Nonpolar Molecules cont'd
11:34
Induced Dipole Attraction, London Dispersion Forces, or Vand der Waals Forces
11:35
Polarizability
13:46
IMF of Nonpolar Molecules cont'd
14:26
Intermolecular Forces (IMF) and Polarizability
14:31
Properties of Liquids
16:48
Standard Molar Enthalpy of Vaporization
16:49
Trends in Boiling Points of Representative Liquids: H₂O vs. H₂S
17:43
Properties of Liquids cont'd
18:36
Aliphatic Hydrocarbons
18:37
Branched Hydrocarbons
20:52
Properties of Liquids cont'd
22:10
Vapor Pressure
22:11
The Clausius-Clapeyron Equation
24:30
Properties of Liquids cont'd
25:52
Boiling Point
25:53
Properties of Liquids cont'd
27:07
Surface Tension
27:08
Viscosity
28:06
Summary
29:04
Sample Problem 1: Determine Which of the Following Liquids Will Have the Lower Vapor Pressure
30:21
Sample Problem 2: Determine Which of the Following Liquids Will Have the Largest Standard Molar Enthalpy of Vaporization
31:37
The Chemistry of Solids

25m 13s

Intro
0:00
Lesson Overview
0:07
Introduction
0:46
General Characteristics
0:47
Particulate-level Drawing
1:09
The Basic Structure of Solids: Crystal Lattices
1:37
The Unit Cell Defined
1:38
Primitive Cubic
2:50
Crystal Lattices cont'd
3:58
Body-centered Cubic
3:59
Face-centered Cubic
5:02
Lattice Enthalpy and Trends
6:27
Introduction to Lattice Enthalpy
6:28
Equation to Calculate Lattice Enthalpy
7:21
Different Types of Crystalline Solids
9:35
Molecular Solids
9:36
Network Solids
10:25
Phase Changes Involving Solids
11:03
Melting & Thermodynamic Value
11:04
Freezing & Thermodynamic Value
11:49
Phase Changes cont'd
12:40
Sublimation & Thermodynamic Value
12:41
Depositions & Thermodynamic Value
13:13
Phase Diagrams
13:40
Introduction to Phase Diagrams
13:41
Phase Diagram of H₂O: Melting Point
14:12
Phase Diagram of H₂O: Normal Boiling Point
14:50
Phase Diagram of H₂O: Sublimation Point
15:02
Phase Diagram of H₂O: Point C ( Supercritical Point)
15:32
Phase Diagrams cont'd
16:31
Phase Diagram of Dry Ice
16:32
Summary
18:15
Sample Problem 1, Part A: Of the Group I Fluorides, Which Should Have the Highest Lattice Enthalpy?
19:01
Sample Problem 1, Part B: Of the Lithium Halides, Which Should Have the Lowest Lattice Enthalpy?
19:54
Sample Problem 2: How Many Joules of Energy is Required to Melt 546 mg of Ice at Standard Pressure?
20:55
Sample Problem 3: Phase Diagram of Helium
22:42
X. Solutions, Rates of Reaction, & Equilibrium
Solutions & Their Behavior

38m 6s

Intro
0:00
Lesson Overview
0:10
Units of Concentration
1:40
Molarity
1:41
Molality
3:30
Weight Percent
4:26
ppm
5:16
Like Dissolves Like
6:28
Like Dissolves Like
6:29
Factors Affecting Solubility
9:35
The Effect of Pressure: Henry's Law
9:36
The Effect of Temperature on Gas Solubility
12:16
The Effect of Temperature on Solid Solubility
14:28
Colligative Properties
16:48
Colligative Properties
16:49
Changes in Vapor Pressure: Raoult's Law
17:19
Colligative Properties cont'd
19:53
Boiling Point Elevation and Freezing Point Depression
19:54
Colligative Properties cont'd
26:13
Definition of Osmosis
26:14
Osmotic Pressure Example
27:11
Summary
31:11
Sample Problem 1: Calculating Vapor Pressure
32:53
Sample Problem 2: Calculating Molality
36:29
Chemical Kinetics

37m 45s

Intro
0:00
Lesson Overview
0:06
Introduction
1:09
Chemical Kinetics and the Rate of a Reaction
1:10
Factors Influencing Rate
1:19
Introduction cont'd
2:27
How a Reaction Progresses Through Time
2:28
Rate of Change Equation
6:02
Rate Laws
7:06
Definition of Rate Laws
7:07
General Form of Rate Laws
7:37
Rate Laws cont'd
11:07
Rate Orders With Respect to Reactant and Concentration
11:08
Methods of Initial Rates
13:38
Methods of Initial Rates
13:39
Integrated Rate Laws
17:57
Integrated Rate Laws
17:58
Graphically Determine the Rate Constant k
18:52
Reaction Mechanisms
21:05
Step 1: Reversible
21:18
Step 2: Rate-limiting Step
21:44
Rate Law for the Reaction
23:28
Reaction Rates and Temperatures
26:16
Reaction Rates and Temperatures
26:17
The Arrhenius Equation
29:06
Catalysis
30:31
Catalyst
30:32
Summary
32:02
Sample Problem 1: Calculate the Rate Constant and the Time Required for the Reaction to be Completed
32:54
Sample Problem 2: Calculate the Energy of Activation and the Order of the Reaction
35:24
Principles of Chemical Equilibrium

34m 9s

Intro
0:00
Lesson Overview
0:08
Introduction
1:02
The Equilibrium Constant
3:08
The Equilibrium Constant
3:09
The Equilibrium Constant cont'd
5:50
The Equilibrium Concentration and Constant for Solutions
5:51
The Equilibrium Partial Pressure and Constant for Gases
7:01
Relationship of Kc and Kp
7:30
Heterogeneous Equilibria
8:23
Heterogeneous Equilibria
8:24
Manipulating K
9:57
First Way of Manipulating K
9:58
Second Way of Manipulating K
11:48
Manipulating K cont'd
12:31
Third Way of Manipulating K
12:32
The Reaction Quotient Q
14:42
The Reaction Quotient Q
14:43
Q > K
16:16
Q < K
16:30
Q = K
16:43
Le Chatlier's Principle
17:32
Restoring Equilibrium When It is Disturbed
17:33
Disturbing a Chemical System at Equilibrium
18:35
Problem-Solving with ICE Tables
19:05
Determining a Reaction's Equilibrium Constant With ICE Table
19:06
Problem-Solving with ICE Tables cont'd
21:03
Example 1: Calculate O₂(g) at Equilibrium
21:04
Problem-Solving with ICE Tables cont'd
22:53
Example 2: Calculate the Equilibrium Constant
22:54
Summary
25:24
Sample Problem 1: Calculate the Equilibrium Constant
27:59
Sample Problem 2: Calculate The Equilibrium Concentration
30:30
XI. Acids & Bases Chemistry
Acid-Base Chemistry

43m 44s

Intro
0:00
Lesson Overview
0:06
Introduction
0:55
Bronsted-Lowry Acid & Bronsted -Lowry Base
0:56
Water is an Amphiprotic Molecule
2:40
Water Reacting With Itself
2:58
Introduction cont'd
4:04
Strong Acids
4:05
Strong Bases
5:18
Introduction cont'd
6:16
Weak Acids and Bases
6:17
Quantifying Acid-Base Strength
7:35
The pH Scale
7:36
Quantifying Acid-Base Strength cont'd
9:55
The Acid-ionization Constant Ka and pKa
9:56
Quantifying Acid-Base Strength cont'd
12:13
Example: Calculate the pH of a 1.2M Solution of Acetic Acid
12:14
Quantifying Acid-Base Strength
15:06
Calculating the pH of Weak Base Solutions
15:07
Writing Out Acid-Base Equilibria
17:45
Writing Out Acid-Base Equilibria
17:46
Writing Out Acid-Base Equilibria cont'd
19:47
Consider the Following Equilibrium
19:48
Conjugate Base and Conjugate Acid
21:18
Salts Solutions
22:00
Salts That Produce Acidic Aqueous Solutions
22:01
Salts That Produce Basic Aqueous Solutions
23:15
Neutral Salt Solutions
24:05
Diprotic and Polyprotic Acids
24:44
Example: Calculate the pH of a 1.2 M Solution of H₂SO₃
24:43
Diprotic and Polyprotic Acids cont'd
27:18
Calculate the pH of a 1.2 M Solution of Na₂SO₃
27:19
Lewis Acids and Bases
29:13
Lewis Acids
29:14
Lewis Bases
30:10
Example: Lewis Acids and Bases
31:04
Molecular Structure and Acidity
32:03
The Effect of Charge
32:04
Within a Period/Row
33:07
Molecular Structure and Acidity cont'd
34:17
Within a Group/Column
34:18
Oxoacids
35:58
Molecular Structure and Acidity cont'd
37:54
Carboxylic Acids
37:55
Hydrated Metal Cations
39:23
Summary
40:39
Sample Problem 1: Calculate the pH of a 1.2 M Solution of NH₃
41:20
Sample Problem 2: Predict If The Following Slat Solutions are Acidic, Basic, or Neutral
42:37
Applications of Aqueous Equilibria

55m 26s

Intro
0:00
Lesson Overview
0:07
Calculating pH of an Acid-Base Mixture
0:53
Equilibria Involving Direct Reaction With Water
0:54
When a Bronsted-Lowry Acid and Base React
1:12
After Neutralization Occurs
2:05
Calculating pH of an Acid-Base Mixture cont'd
2:51
Example: Calculating pH of an Acid-Base Mixture, Step 1 - Neutralization
2:52
Example: Calculating pH of an Acid-Base Mixture, Step 2 - React With H₂O
5:24
Buffers
7:45
Introduction to Buffers
7:46
When Acid is Added to a Buffer
8:50
When Base is Added to a Buffer
9:54
Buffers cont'd
10:41
Calculating the pH
10:42
Calculating the pH When 0.010 mol NaOH is Added to 1.0 L of the Buffer
14:03
Buffers cont'd
14:10
Calculating the pH When 0.010 mol NaOH is Added to 1.0 L of the Buffer: Step 1 -Neutralization
14:11
Calculating the pH When 0.010 mol NaOH is Added to 1.0 L of the Buffer: Step 2- ICE Table
15:22
Buffer Preparation and Capacity
16:38
Example: Calculating the pH of a Buffer Solution
16:42
Effective Buffer
18:40
Acid-Base Titrations
19:33
Acid-Base Titrations: Basic Setup
19:34
Acid-Base Titrations cont'd
22:12
Example: Calculate the pH at the Equivalence Point When 0.250 L of 0.0350 M HClO is Titrated With 1.00 M KOH
22:13
Acid-Base Titrations cont'd
25:38
Titration Curve
25:39
Solubility Equilibria
33:07
Solubility of Salts
33:08
Solubility Product Constant: Ksp
34:14
Solubility Equilibria cont'd
34:58
Q < Ksp
34:59
Q > Ksp
35:34
Solubility Equilibria cont'd
36:03
Common-ion Effect
36:04
Example: Calculate the Solubility of PbCl₂ in 0.55 M NaCl
36:30
Solubility Equilibria cont'd
39:02
When a Solid Salt Contains the Conjugate of a Weak Acid
39:03
Temperature and Solubility
40:41
Complexation Equilibria
41:10
Complex Ion
41:11
Complex Ion Formation Constant: Kf
42:26
Summary
43:35
Sample Problem 1: Question
44:23
Sample Problem 1: Part a) Calculate the pH at the Beginning of the Titration
45:48
Sample Problem 1: Part b) Calculate the pH at the Midpoint or Half-way Point
48:04
Sample Problem 1: Part c) Calculate the pH at the Equivalence Point
48:32
Sample Problem 1: Part d) Calculate the pH After 27.50 mL of the Acid was Added
53:00
XII. Thermodynamics & Electrochemistry
Entropy & Free Energy

36m 13s

Intro
0:00
Lesson Overview
0:08
Introduction
0:53
Introduction to Entropy
1:37
Introduction to Entropy
1:38
Entropy and Heat Flow
6:31
Recall Thermodynamics
6:32
Entropy is a State Function
6:54
∆S and Heat Flow
7:28
Entropy and Heat Flow cont'd
8:18
Entropy and Heat Flow: Equations
8:19
Endothermic Processes: ∆S > 0
8:44
The Second Law of Thermodynamics
10:04
Total ∆S = ∆S of System + ∆S of Surrounding
10:05
Nature Favors Processes Where The Amount of Entropy Increases
10:22
The Third Law of Thermodynamics
11:55
The Third Law of Thermodynamics & Zero Entropy
11:56
Problem-Solving involving Entropy
12:36
Endothermic Process and ∆S
12:37
Exothermic Process and ∆S
13:19
Problem-Solving cont'd
13:46
Change in Physical States: From Solid to Liquid to Gas
13:47
Change in Physical States: All Gases
15:02
Problem-Solving cont'd
15:56
Calculating the ∆S for the System, Surrounding, and Total
15:57
Example: Calculating the Total ∆S
16:17
Problem-Solving cont'd
18:36
Problems Involving Standard Molar Entropies of Formation
18:37
Introduction to Gibb's Free Energy
20:09
Definition of Free Energy ∆G
20:10
Spontaneous Process and ∆G
20:19
Gibb's Free Energy cont'd
22:28
Standard Molar Free Energies of Formation
22:29
The Free Energies of Formation are Zero for All Compounds in the Standard State
22:42
Gibb's Free Energy cont'd
23:31
∆G° of the System = ∆H° of the System - T∆S° of the System
23:32
Predicting Spontaneous Reaction Based on the Sign of ∆G° of the System
24:24
Gibb's Free Energy cont'd
26:32
Effect of reactant and Product Concentration on the Sign of Free Energy
26:33
∆G° of Reaction = -RT ln K
27:18
Summary
28:12
Sample Problem 1: Calculate ∆S° of Reaction
28:48
Sample Problem 2: Calculate the Temperature at Which the Reaction Becomes Spontaneous
31:18
Sample Problem 3: Calculate Kp
33:47
Electrochemistry

41m 16s

Intro
0:00
Lesson Overview
0:08
Introduction
0:53
Redox Reactions
1:42
Oxidation-Reduction Reaction Overview
1:43
Redox Reactions cont'd
2:37
Which Reactant is Being Oxidized and Which is Being Reduced?
2:38
Redox Reactions cont'd
6:34
Balance Redox Reaction In Neutral Solutions
6:35
Redox Reactions cont'd
10:37
Balance Redox Reaction In Acidic and Basic Solutions: Step 1
10:38
Balance Redox Reaction In Acidic and Basic Solutions: Step 2 - Balance Each Half-Reaction
11:22
Redox Reactions cont'd
12:19
Balance Redox Reaction In Acidic and Basic Solutions: Step 2 - Balance Hydrogen
12:20
Redox Reactions cont'd
14:30
Balance Redox Reaction In Acidic and Basic Solutions: Step 3
14:34
Balance Redox Reaction In Acidic and Basic Solutions: Step 4
15:38
Voltaic Cells
17:01
Voltaic Cell or Galvanic Cell
17:02
Cell Notation
22:03
Electrochemical Potentials
25:22
Electrochemical Potentials
25:23
Electrochemical Potentials cont'd
26:07
Table of Standard Reduction Potentials
26:08
The Nernst Equation
30:41
The Nernst Equation
30:42
It Can Be Shown That At Equilibrium E =0.00
32:15
Gibb's Free Energy and Electrochemistry
32:46
Gibbs Free Energy is Relatively Small if the Potential is Relatively High
32:47
When E° is Very Large
33:39
Charge, Current and Time
33:56
A Battery Has Three Main Parameters
33:57
A Simple Equation Relates All of These Parameters
34:09
Summary
34:50
Sample Problem 1: Redox Reaction
35:26
Sample Problem 2: Battery
38:00
XIII. Transition Elements & Coordination Compounds
The Chemistry of The Transition Metals

39m 3s

Intro
0:00
Lesson Overview
0:11
Coordination Compounds
1:20
Coordination Compounds
1:21
Nomenclature of Coordination Compounds
2:48
Rule 1
3:01
Rule 2
3:12
Rule 3
4:07
Nomenclature cont'd
4:58
Rule 4
4:59
Rule 5
5:13
Rule 6
5:35
Rule 7
6:19
Rule 8
6:46
Nomenclature cont'd
7:39
Rule 9
7:40
Rule 10
7:45
Rule 11
8:00
Nomenclature of Coordination Compounds: NH₄[PtCl₃NH₃]
8:11
Nomenclature of Coordination Compounds: [Cr(NH₃)₄(OH)₂]Br
9:31
Structures of Coordination Compounds
10:54
Coordination Number or Steric Number
10:55
Commonly Observed Coordination Numbers and Geometries: 4
11:14
Commonly Observed Coordination Numbers and Geometries: 6
12:00
Isomers of Coordination Compounds
13:13
Isomers of Coordination Compounds
13:14
Geometrical Isomers of CN = 6 Include: ML₄L₂'
13:30
Geometrical Isomers of CN = 6 Include: ML₃L₃'
15:07
Isomers cont'd
17:00
Structural Isomers Overview
17:01
Structural Isomers: Ionization
18:06
Structural Isomers: Hydrate
19:25
Structural Isomers: Linkage
20:11
Structural Isomers: Coordination Isomers
21:05
Electronic Structure
22:25
Crystal Field Theory
22:26
Octahedral and Tetrahedral Field
22:54
Electronic Structure cont'd
25:43
Vanadium (II) Ion in an Octahedral Field
25:44
Chromium(III) Ion in an Octahedral Field
26:37
Electronic Structure cont'd
28:47
Strong-Field Ligands and Weak-Field Ligands
28:48
Implications of Electronic Structure
30:08
Compare the Magnetic Properties of: [Fe(OH₂)₆]²⁺ vs. [Fe(CN)₆]⁴⁻
30:09
Discussion on Color
31:57
Summary
34:41
Sample Problem 1: Name the Following Compound [Fe(OH)(OH₂)₅]Cl₂
35:08
Sample Problem 1: Name the Following Compound [Co(NH₃)₃(OH₂)₃]₂(SO₄)₃
36:24
Sample Problem 2: Change in Magnetic Properties
37:30
XIV. Nuclear Chemistry
Nuclear Chemistry

16m 39s

Intro
0:00
Lesson Overview
0:06
Introduction
0:40
Introduction to Nuclear Reactions
0:41
Types of Radioactive Decay
2:10
Alpha Decay
2:11
Beta Decay
3:27
Gamma Decay
4:40
Other Types of Particles of Varying Energy
5:40
Nuclear Equations
6:47
Nuclear Equations
6:48
Nuclear Decay
9:28
Nuclear Decay and the First-Order Kinetics
9:29
Summary
11:31
Sample Problem 1: Complete the Following Nuclear Equations
12:13
Sample Problem 2: How Old is the Rock?
14:21
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Lecture Comments (8)

1 answer

Last reply by: Professor Franklin Ow
Tue Jul 15, 2014 1:36 AM

Post by Meredith Roach on July 14, 2014

Needs some better explanations to build understanding

1 answer

Last reply by: Professor Franklin Ow
Wed May 21, 2014 2:02 AM

Post by Anthony Mendoza on May 19, 2014

In sample Question 1, Part d, how did you choose which form of the element you put in the answer? Why Cl2(g) versus Cl-(aq)?

1 answer

Last reply by: Professor Franklin Ow
Wed May 21, 2014 2:00 AM

Post by Anthony Mendoza on May 19, 2014

Hi,
I'm still getting confused on what's going on in the picture at 21:30.
My instructor put the same picture on the board, but I got lost on it, too.

What is making the electrons go to the right?
Why are there cations on the right side of the salt bridge? (why not the opposite)?
Why is the right side the + cathod? (Isn't it getting a flow of electrons to it, so would be more on the negative side than the neutral side?)
Where are the electrons (that flow from left to right) coming from? From the solid A? If so, why is the aqueous portion needed on the left? What role does it play?

I feel totally lost with the picture. It's like I need a play by play on what's going on and why...I'm good with redox as far as balancing reactions and all that, but having trouble applying it to the picture.

I appreciate the help

1 answer

Last reply by: Professor Franklin Ow
Mon Feb 24, 2014 11:35 AM

Post by Ruth Arthur on February 23, 2014

At 14:17, you said the charges on the left side of the equation has an net charge of +1. Won't the + & - charge on the 2H3O and NO3 cancel out? I just got a little confused there.

Related Articles:

Electrochemistry

  • Electrochemistry studies how chemical energy can be produced and converted into mechanical energy, which in turn can perform work such as powering an electrical device.
  • A potential is produced when there is a transfer of electrons from the reducing agent to the oxidizing agent which can happen in neutral, acidic, or basic media.
  • The Nernst equation allows for determination of the cell potential when not at standard conditions.

Electrochemistry

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
  • Lesson Overview 0:08
  • Introduction 0:53
  • Redox Reactions 1:42
    • Oxidation-Reduction Reaction Overview
  • Redox Reactions cont'd 2:37
    • Which Reactant is Being Oxidized and Which is Being Reduced?
  • Redox Reactions cont'd 6:34
    • Balance Redox Reaction In Neutral Solutions
  • Redox Reactions cont'd 10:37
    • Balance Redox Reaction In Acidic and Basic Solutions: Step 1
    • Balance Redox Reaction In Acidic and Basic Solutions: Step 2 - Balance Each Half-Reaction
  • Redox Reactions cont'd 12:19
    • Balance Redox Reaction In Acidic and Basic Solutions: Step 2 - Balance Hydrogen
  • Redox Reactions cont'd 14:30
    • Balance Redox Reaction In Acidic and Basic Solutions: Step 3
    • Balance Redox Reaction In Acidic and Basic Solutions: Step 4
  • Voltaic Cells 17:01
    • Voltaic Cell or Galvanic Cell
    • Cell Notation
  • Electrochemical Potentials 25:22
    • Electrochemical Potentials
  • Electrochemical Potentials cont'd 26:07
    • Table of Standard Reduction Potentials
  • The Nernst Equation 30:41
    • The Nernst Equation
    • It Can Be Shown That At Equilibrium E =0.00
  • Gibb's Free Energy and Electrochemistry 32:46
    • Gibbs Free Energy is Relatively Small if the Potential is Relatively High
    • When E° is Very Large
  • Charge, Current and Time 33:56
    • A Battery Has Three Main Parameters
    • A Simple Equation Relates All of These Parameters
  • Summary 34:50
  • Sample Problem 1: Redox Reaction 35:26
  • Sample Problem 2: Battery 38:00

Transcription: Electrochemistry

Hi, welcome back to Educator.com.0000

Today's lesson in general chemistry is electrochemistry.0003

We are going to start off with a brief introduction in electrochemistry--what it is about.0010

We are going to go ahead and go into the main reactions that are0015

at the core of electrochemistry which are what we call redox reactions.0020

A big application of electrochemistry is what we call a voltaic cell.0025

We are going to take a look at that followed by quantifying aspects in electrochemistry which deal with what we call electrochemical potentials.0028

After we introduce the potentials, we will then go ahead and get into quantitative problems0040

that include current, charge, and time, followed by a brief summary and sample problems.0046

Electrochemistry is the branch of chemistry which deals with harnessing chemical energy and converting it into mechanical energy.0054

Basically it studies the processes in which energy is released during a chemical reaction0064

so that it can be used to perform some type of work.0071

The perfect example of this is the battery which powers everyday machines including vehicles, computers, and phones.0076

Basically a battery is a chemical process that produces energy that is then converted into electricity.0084

Basically in this lecture, we are going to examine both the concepts and equations that are at the core of electrochemistry.0095

An electrochemical reaction is one that involves a complete transfer of electrons from one reactant to another.0104

The reactant which loses the electrons is said to be oxidized or undergo oxidation.0111

The nice little acronym that allows us to remember this is OIL or oxidation-is-loss.0119

The reactant that gains electrons is said to be reduced or undergo reduction.0125

The acronym for that is RIG or reduction-is-gain.0135

OILRIG, oxidation is loss; reduction is gain.0138

An oxidation reduction reaction is also known as a redox reaction for short.0143

Let's now look at how to identify each type of a reactant.0151

To identify which reactant is being oxidized and which is being reduced, it helps to0158

keep track of oxidation states which we learned in the first half of general chemistry.0164

Let's go ahead and consider the following reaction.0169

CH4 gas plus O2 gas goes on to form CO2 gas and H2O gas.0175

Let's go ahead and balance this guy.0191

We are going to need two of these and two of those.0194

This is the reaction that represents the combustion of methane.0201

We are going to now assign oxidation states to each of the elements.0207

Hydrogen is +1 of course which means overall this carbon is going to be -4.0214

Here in O2, this is a homonuclear diatomic; each oxygen is 0 here.0222

Here in CO2, oxygen is -2 each which means carbon is going to be +4.0230

Here oxygen is -2; each hydrogen is +1.0238

We see that there are two elements that experience a change in oxidation state.0243

Carbon starts off as -4; it becomes +4 in the product side.0249

Oxygen starts off as 0; it becomes -2 on the product side.0261

We see that carbon's oxidation state has increased.0269

If the oxidation number goes up, that means it has lost electrons.0275

Oxygen's oxidation state has gone down which means it has gained electrons.0284

If carbon has lost electrons, we say that it has been oxidized.0293

If oxygen gains electrons, we say that it has been reduced.0303

But it is not just the element that undergoes the process.0308

It is that entire compound.0311

To be more accurate, we say that CH4 gas was oxidized.0313

We say that O2 gas was reduced.0323

We are going to introduce two more terms here.0330

CH4 gas was oxidized; if it was oxidized, it performed the reduction.0333

Therefore CH4 gas is also known as the reducing agent.0341

O2 gas was reduced.0353

If it was reduced, it performed the oxidation which means that O2 gas is what we call the oxidizing agent.0357

You definitely want to keep track of these terms because they are somewhat confusing.0369

A reducing agent performs the reduction; but it itself is oxidized.0373

A reducing agent performs the reduction... excuse me.0379

An oxidizing agent performs the oxidation; it itself is reduced.0385

Let's now learn how to balance redox reactions.0389

Redox reactions can occur in neutral solutions and in acidic solutions and basic solutions.0397

We are going to learn how to balance redox reactions in all three of these types of medium.0407

In neutral solutions, let's go ahead and start off.0413

Cl2 gas can react with Zn2+ to go and form Cl1- aqueous plus zinc solid like that.0422

What we are going to do is we are going to balance these guys.0447

When we balance in neutral solutions, you must make sure that the0452

stoichiometry is balanced and the net charge on both sides of the equation.0456

When we go ahead and look at this, we see that we have one zinc on each side.0467

That is good to go.0472

We have two chlorines on the left and only one on the right.0473

Let's go ahead and put the two chlorines there.0476

My net balanced charge here is 2+.0478

My net charge here is going to be 2-.0482

Let's go ahead and see what the half reactions are.0492

For the half reactions, we have Cl2 gas going to become 2Cl1- aqueous.0501

That means that each chlorine is gaining an electron.0516

I am going to need two electrons on the left side.0520

Zn2+ is becoming zinc solid.0524

That means I am going to need two electrons on this side.0530

When we look at the half reactions, these are both reduction.0550

But you can't have two reduction half reactions.0558

You always have to have one oxidation; you always have to have one reduction.0561

It turns out that the chlorine one is going to be the one that reacts.0566

What is going to happen is we are going to flip this one.0577

That is going to be zinc solid going to form Zn2+ plus two electrons.0589

This immediately told me that we have done something wrong because my net charge on each side is not the same.0599

My overall reaction is going to be Cl2 gas plus zinc solid going on to form 2Cl1- aqueous and Zn2+ aqueous.0608

As you can see, now my overall charge is fixed.0624

This is 0 on the left.0628

This is 0 on the right side for our balanced redox reaction in neutral solutions.0629

Let's now jump into solutions where we have acidic or basic conditions.0637

For example, in acid, silver solid plus nitrate aqueous goes on to form NO2 gas and Ag+1 aqueous.0644

Step one is to break the unbalanced equation into your half reactions.0652

My half reaction here is going to be silver solid going to Ag1+ aqueous.0658

NO31- aqueous going on to form NO2 gas.0671

Those are our half reactions.0681

Step two, we are going to balance each of the half reactions by being careful to follow each of these following steps.0683

Balance all elements except oxygen and hydrogen by stoichiometry.0692

For the silver half reaction, we have one of each on side.0699

We are good to go on that.0703

For nitrate going to nitrogen dioxide, we have one nitrogen on each side.0705

We are good on that; part A is done.0711

Part B, we are going to balance oxygen by adding water to the side that is deficient in oxygen.0715

Here for nitrate, I have three oxygens; for nitrogen dioxide, I only have two oxygens.0721

I am going to add H2O to the side that is deficient in oxygen.0727

Let's go ahead and move on now to step number three.0737

If in acidic solution, we are going to add hydronium for0742

every missing hydrogen to the side that is deficient in H0745

and the same number of water molecules to the other side.0749

Our half reaction again was NO31- aqueous going on to form NO2 gas.0752

We added water last time; we are going to follow part C.0762

We are going to add hydronium for every missing hydrogen to the side that is deficient in hydrogen.0767

On the right side of this equation, I have the two hydrogens here.0773

I am going to add two hydroniums on the left side.0776

Remember you are adding one hydronium for every hydrogen.0780

At the same time, you are going to add water, the same number of water molecules to the other side.0784

Let me go ahead and do that in blue; plus two more H2Os.0789

If this was basic solution, we would add one water for0796

every missing hydrogen to the side that is deficient in H0799

and the same number of hydroxide to the other side.0803

What we are going to do right now is we are going to go on to part E.0807

We are going to balance the net charge by adding electrons to the side that is deficient in negative charge.0812

Let's go ahead and look at this.0818

Right now we have two hydroniums aqueous plus NO31- aqueous.0820

That goes on to form NO2 gas plus 3H2O liquid.0829

We have a net charge of +1 on the left side.0842

We have a net charge of 0 on the right side which means that0851

we have to add an electron to the left side here to get us to 0.0856

We will then multiply the half reactions by integers that will lead to cancellation of electrons.0875

In other words, each half reaction should have an identical number of electrons.0880

Ag solid going to Ag1+ aqueous.0887

It was now NO31- aqueous going to NO2 gas.0895

I believe this was three waters; this was two hydroniums aqueous.0906

Let's go ahead and fix the electrons; I am going to need one electron here.0916

I am only going to need one electron right there.0922

Because the number of electrons is the same in each reaction, we are good to go.0930

We are going to now recombine the half reactions to get the net balanced equation.0939

Be sure to cancel like terms.0943

When I add these two reactions together, the electrons cancel out.0946

We get silver solid plus 2H3O1+ aqueous plus NO31- aqueous0950

going on to form Ag1+ aqueous plus NO2 gas and 3H2O liquid.0961

This should be liquid here.0971

Let's go ahead and double check if everything is balanced.0973

On the left side, I have one silver.0977

On the right side, I have one silver; that is good to go.0979

On the left side, I have six hydrogens.0982

On the right side, I have six hydrogens.0984

On the left side, I have one nitrogen.0987

On the right side, I have one nitrogen.0989

On the left side, I have a total of five oxygens.0992

On the right side, I have a total of five oxygens.0996

Stoichiometry is good to go; let's do a final check on the charge.0999

I have a total charge of 1+ on the left side.1003

I have a total charge of 1+ on the right side.1010

Yes, this is our overall balanced redox reaction.1012

Let's go ahead now and look at how a simple battery works.1017

The first battery was referred to as a voltaic cell, also known as a galvanic cell shown below.1023

Let's go ahead and look at the redox reaction for one we already looked at.1031

This was going to be then Cl2 gas plus zinc solid going on to form 2Cl1- aqueous and Zn2+ aqueous.1040

Basically the redox reaction was the following; we can have a metal electrode here.1070

The metal electrode is going to be dipped in a salt solution that contains its ion.1091

If let's say the metal electrode was represented by A solid, the salt solution1099

could be for example A+ and maybe X1- for our salt solution.1107

Again I have another metal electrode on the right side of a different identity.1121

This could be B solid for example; this would be B+ and X1-.1125

These two electrodes are going to be connected by a conductive wire here.1133

We get a flow of electrons from A to B.1143

The flow of electrons is going this way.1147

Anytime you have a flow of electrons, that generates an electric current which can then be captured to power some type of device.1151

What we see here is that A solid is going to be losing electrons to form A1+ aqueous like so.1162

That is its half reaction.1175

B solid is going to be gaining the electrons to form B1- aqueous.1177

We need something called a salt bridge to help us complete the circuit and to balance the charge.1195

Here because we are forming a lot of A+, some type of ion has to come in and balance the positive charge.1205

Maybe some more X- comes in from the salt bridge.1218

The salt bridge can be for example A+ and X1-.1222

Because in the B side we are forming a lot of negative charge,1232

some positive cation is going to come in from the salt bridge also.1240

We see that here that this is going to be representing the oxidation half reaction.1249

Here on the right side, this is going to be the reduction half reaction.1255

We can now also look at the charges.1272

If the flow of electrons is from A to B, that means that...1275

Remember that electrons are going to go away from a negative side.1280

This is the negative charge.1284

They are going to go toward the positive side which is going to be where the electrons gravitate toward.1287

If you ever look at a battery, a battery always has a negative sign and a positive sign representing these two electrode terminals.1294

The negative electrode is what we call the anode.1303

The positive terminal is what we call the cathode.1308

A nice little trick to remember this is when you cross the t in the word cathode, it looks like a positive sign.1315

We can also have something we call cell notation.1324

Cell notation typically looks like this.1328

On the left side goes the anode; on the right side goes the cathode.1344

The single lines represent changes in physical state.1357

The double line just simply represents a salt bridge.1367

Here we can easily write in A solid going to A1+ aqueous.1372

Here we can write in B1+ aqueous going to B solid.1382

Let me go ahead and correct this then.1399

This is really B1+ aqueous plus an electron going to be solid.1402

Again this is what we call cell notations.1409

The anode half reaction is written on the left.1411

The cathode half reaction is written on the right side.1415

Another thing about the voltaic cell is that we see that the anode here is getting consumed.1421

The solid is becoming ionized.1431

We see here that the cathode, we are actually depositing more solid onto the electrode.1434

This is going to become heavier, become larger in mass.1442

Why does a battery die?1453

It is essentially because one of the terminals, the anode gets consumed to the point where no more reduction or oxidation can occur.1456

If you ever have a rechargeable battery, you know that a rechargeable battery you have to plug in.1472

What does that do?1479

When we put a rechargeable battery into the wall outlet and charge it overnight, we are basically driving the reverse reaction.1480

The reverse reaction is then replenishing the anode while consuming the cathode.1488

But that doesn't always go on forever; we cannot use a rechargeable battery infinitely.1494

It also too is going to eventually die.1500

That is because no process is 100 percent efficient.1504

You will never ever regain all of your anode.1508

It is going to become less and less efficient as time goes on.1514

Let's now go ahead and look at the quantitative part of electrochemistry.1518

All half reactions have a potential measured in volts, when you purchase for example a 1.5 volt battery.1524

This is just like a chemical reaction having a change in energy associated with it.1531

These potentials are all relative to the hydrogen reduction half reaction which is arbitrarily assigned 0 volts.1536

That is going to be 2H+ aqueous plus two electrons going on to form H2 gas.1544

This potential is what we symbolize as E standard.1555

That is going to be 0.00 volts.1559

It is very important to be able to use the following table of standard reduction potentials.1561

Here I have listed for you five half reactions from this table.1570

You see that these are all reduction half reactions.1575

On the right side are the reduction potentials; E standard values.1584

It is important that we understand how to decipher this.1590

Point number one is a large E0 means that the reduction is more likely to occur.1595

Again a large E0 is reduction more likely to occur.1606

For example, the reduction of Cl2 has a potential of +1.36 volts.1611

The reduction of Zn2+ has a potential of -0.76 volts which means that Cl2 is more likely to be reduced than Zn2+.1622

In fact if Cl2 is going to be...1638

If we combine the two, Cl2 gas plus two electrons going to 2Cl1-1644

and Zn2+ aqueous plus two electrons going on to zinc solid,1654

which one is going to be the reduction half reaction?1664

Which one is going to be the oxidation half reaction?1667

Because this half reaction for chlorine is more positive, that is going to remain reduction.1672

Zinc solid is going to become oxidized.1677

The combination of Cl2 gas plus two electrons going on to form 2Cl- and1682

zinc solid going on to form Zn2+ plus two electrons for a balanced redox reaction of1690

Cl2 gas plus zinc solid going on to form Zn2+ aqueous and 2Cl1- aqueous.1697

In other words, Cl2 gets reduced; zinc is going to be oxidized.1708

This is very important because when we read this table, we can always draw an arrow in the following direction.1720

Anything that follows this direction, you are going to have a successful redox reaction.1729

Basically we can summarize it the following way, what this line means.1738

Any element or ion can oxidize any other species...1743

Any element or ion can oxidize any other species to the bottom and right of it, to the bottom right of it.1765

For example, Cl2 can oxidize silver; it can oxidize H2.1780

It can oxidize nickel; it can oxidize zinc.1786

Cl2 itself will be reduced because its reduction potential is so positive.1790

We see therefore that when we go this direction, these are going to be very strong oxidizing agents.1798

When we go this direction on the right side, these are going to be stronger reducing agents.1812

In this table, we see that Cl2 is the strongest oxidizing agent.1822

It itself is most likely to be reduced.1827

Zinc solid is the strongest reducing agent.1830

It itself is going to be more likely to be oxidized.1832

Let's now look at calculating the potential when not at standard conditions.1837

To calculate the potential for a voltaic cell when not at standard conditions, we use what is called the Nerst equation.1843

This is just like ΔG is equal to ΔG0 minus RT natural log of Q; very very similar.1851

Once again Q is going to be the reaction quotient.1857

The temperature is in kelvin of course.1864

R is going to be 8.314 joules per k mole.1868

n is going to be the moles of electrons.1875

You get that from the balanced redox reaction; moles of electrons from balanced redox reaction.1878

F is what we call Faraday's constant.1890

Faraday's constant, you should always ask your instructor if you have to know it or not.1898

It is basically 9.65 times 104.1904

The units, I am going to draw in red, are coulombs for every mole of electron.1909

Once again coulombs per moles of electron or C over mole.1918

Again you are going to use this equation anytime you are dealing with nonstandard conditions1924

which is basically not 1 atm of pressure and not 1 molarity of concentration.1929

It can also be shown that at equilibrium, E is going to be 0.00 volts1937

which means that E0 is equal to RT over nF times the natural log of K.1943

Again this is any of the K values--Ka, Kb, Kf, Ksp, etc.1948

Again these are the two equations that are heavily used in electrochemistry.1960

Finally it can be shown that we can relate Gibbs free energy to electrochemistry.1968

The equation is ΔG0 is equal to ?nFE0.1977

This equation tells us that if E0 is large...1982

If E0 is large, that means the potential is very high which means the redox reaction is likely to occur.1988

Remember what we said if something is likely to occur>2000

If something is likely to occur, ΔG0 is negative.2003

Look, that makes perfect sense because if E is positive... we know F is positive.2008

We know n is positive making that whole side of the equation negative giving us a negative Gibbs free energy.2013

Basically if E0 is very large, that means the reaction is in the forward direction.2020

It is highly product favored which as we have seen translates to a lower energy state.2026

Let's now see how to factor in the lifetime of a battery.2032

A battery basically has three main parameters--the voltage of the cell,2038

the current that is drawn from it by the device,2043

and the total amount of time the battery can run at the specified current.2045

A simple equation relates all of these parameters.2050

n is equal to It over F where I is your current in amps.2053

Time, t is your time in seconds.2063

Of course F is your Faraday's constant, 9.65 times 104 coulombs per mole of electrons.2069

In this equation, n is not going to be a whole number.2079

n is going to be the mole of electrons that is determined stoichiometrically.2082

Let's go ahead and summarize this before we jump into our sample problems.2091

Electrochemistry we saw studies how chemical energy can be produced and converted into2095

mechanical energy which in turn can perform work such as powering an electrical device.2100

A potential is produced when there is a transfer of electrons2107

from the reducing agent to the oxidizing agent which as2110

we have seen can happen in neutral, acidic, or basic media.2114

Finally the Nerst equation allows for determination of the cell potential when not at standard conditions.2118

Sample problem number one, basically here...2128

This is a very common type of problem by the way.2132

You are given a list of half reactions here with their reduction potentials.2134

E0 is right here.2140

You are just asked to answer the following questions, all qualitative problems.2146

What is the strongest oxidizing agent?--what is the strongest reducing agent?2150

Remember that the stronger oxidizing agents are the ones that are going to be2154

most likely to be reduced which corresponds to a very high reduction potential.2158

Here the strongest oxidizing agent is Cl2 gas.2163

The strongest reducing agent therefore is going to be zinc solid.2168

Part B, list the species that silver solid can reduce if any.2173

Silver solid is not going to reduce species.2181

But instead it is going to oxidize species.2187

Silver solid is going to only... excuse me.2189

Silver solid can reduce anything to the top and left of it.2198

That is only going to be Cl2 gas; again reduction occurs to top and left.2205

List the species that zinc solid can reduce if any.2216

Zinc solid because it is such a strong oxidizing agent, it can reduce any species that is to the top and left of it.2220

That is going to be Ni2+, H+, Ag1+, and Cl2 gas.2231

Finally part D, what species is most easily oxidized?2240

The species that is most easily oxidized is going to be the strongest reducing agent.2244

Therefore this is going to be zinc solid.2254

What species is most easily reduced?2257

That is going to be the strongest oxidizing agent which is going to be Cl2 gas.2259

You see, I just want to point out a couple things.2264

That part A and part D, these are really the same question that are asked differently.2266

Again there is going to be multiple ways of phrasing, of posing the same question.2272

Sample problem number two, if a battery contains 0.355 grams of cadmium and excess2280

nickel hydroxide, how long can it operate if it draws 1.25 milliamps of current?2288

You are given the balanced redox reaction below.2294

Because this question mentions current and how long, there is only one equation that deals with that.2299

That is n is equal to It over F.2305

We know that I is amps; that is going to be 0.00125 amps.2308

We know Faraday's constant already, 96500 coulombs per mole of electrons.2316

The question is asking us to solve for time.2324

t is going to be equal to n times F over I.2327

All we need is n.2332

Remember that n is the mole of electrons that is determined stoichiometrically.2334

Let's go ahead and get it.2339

You are told that you have an excess amount of the nickel precursor.2340

The 0.355 grams of cadmium is going to be our limiting reactant here.2345

0.355 grams of cadmium times 1 mole of cadmium divided by 112.41 grams.2350

That is going to be equal to 0.00316 moles of cadmium solid.2365

The redox reaction is right here; that is cadmium zero going to cadmium hydroxide.2374

In this reaction, cadmium is 2+.2388

For cadmium to go to cadmium 2+, it has to give up two electrons.2395

My mole to mole ratio of the electrons to cadmium is a 2:1 ratio.2401

0.00316 moles of cadmium times 2 moles of electrons for every 1 mole of cadmium.2408

That is going to give me the value of n which is 0.00632 moles.2419

When all is said and done, you should get a time of 487904 seconds which is approximately 136 hours.2428

That is an awfully nice lifetime for a battery to operate, 136 hours.2441

If you look at these elements, it is nickel cadmium.2450

I am sure you have heard of batteries that are called nickel cadmium or cadmium nickel batteries2453

which are very nice batteries that help to power devices such as your smartphones and cameras.2459

That is our lecture from electrochemistry; I want to thank you for your time.2469

I will see you next time on Educator.com.2474

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