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Franklin Ow

Franklin Ow

Chemical Reactions II

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 (15)

0 answers

Post by Kevin Guo on August 16 at 11:59:01 AM

Why does this lecture end at 21:56 when it's supposed to be 50 minutes long? When I click on the time stamps past 21:56, the video resets. I am currently logged in as well.

0 answers

Post by Pranav Brahmandam on March 31 at 06:25:55 PM

Hi Dr. Ow
The video is not playing after 21:23 Min.  I tried to log in second time too, but the problem persists.  


Thank You,

1 answer

Last reply by: Macy Li
Wed Jun 21, 2017 11:55 AM

Post by Mohsin Alibrahim on June 17, 2017

The second half of the video doesn't work

0 answers

Post by Bertha Paul Fils on November 21, 2016

When you explained oxidation and reduction, you mistakenly flipped the definition of each one. You associated the wrong definitions with the wrong terms. Please confirm this for me, thank you!

0 answers

Post by Parth Shorey on September 3, 2016

There is something wrong with this video, it doesn't play the other half of the lecture. At first I thought it was my computer so I played a couple of your other lectures, and all of them played just fine. However this one seemed to just pause.

1 answer

Last reply by: Professor Franklin Ow
Mon Jul 18, 2016 3:48 PM

Post by tae Sin on July 14, 2016

does every acid have a hydrogen? and if so, does this mean that the hydrogen ion is technically a proton used for donating?

2 answers

Last reply by: Professor Franklin Ow
Tue Oct 14, 2014 6:59 PM

Post by Saadman Elman on October 12, 2014

Hi, Dr. Franklin. I really like the way you teach. Your lectures are very explicit and easy to understand.

I have some questions about Strong and Weak bases. In my book it is clearly documented that NaOH, Ba OH are strong electrolytes. Even in your lecture you mentioned that Group 1 metal oxide and Some Group 2 metal oxides are Strong Bases and others are weak bases. As far as i am concerned, i think u also said it in the lecture that Weak Bases are Weak Electrolyte and Strong acids are strong electrolyte. In the beginning of this lecture when you were  explaining about Arrhenious definition of Acids and Bases, you showed us how NaOH Completely dissociates into cation and anion and thus having single arrow, but you also showed how Mg(OH)2 is NOT fully dissociating thus making it an equilibrium. So in the back of mind even though you didn't mention it i assumed that Mg (OH)2 must be a WEAK ELECTROLYTE. Why? because it is not completely dissociating. But at 17 min in this lecture when you said Mg (OH)2  are some group 2 metals which is actually Strong Base. You just said it is a strong base. Although you didn't say it is a strong electrolyte but it's self explanatory. So my questions is since Mg (OH)2 Is NOT fully dissociating than why Mg (OH)2 is strong base. It makes more sense for me to think that it is weak. Please explain it. Thanks Can't wait to read your comment.  

1 answer

Last reply by: Professor Franklin Ow
Wed Mar 5, 2014 12:36 AM

Post by SD Ryo on March 4, 2014

Hi Dr Franklin,

At 11 mins, why aren't the charges for HSO4- + H20 -><- SO4 2+ + H3O+ balanced on both the left and right side? The left side has a negative charge while the right has a 3+ charge.

1 answer

Last reply by: Professor Franklin Ow
Fri Feb 7, 2014 10:22 AM

Post by Koh Huai Ze on February 3, 2014

Hi, why is dilute H2SO4 always used to acidify potassium manganate(VII) solution and potassium dichromate(VI) solution in lab experiments? Is it because sulfuric acid is a strong dibasic acid (thus a lesser amount of acid is required), or are there other reasons? Can dilute nitric acid or dilute hydrochloric acid be used too? Thank you.

Related Articles:

Chemical Reactions II

  • Arrhenius acids ionize in water to form H+ and an anion, while Arrhenius bases ionize to form hydroxide and a cation.
  • Bronsted-Lowry acids donate a proton to water, while Bronsted-Lowry bases accept a proton from water.
  • The (7) strong acids dissociate completely in water.
  • A neutralization reaction between a Bronsted-Lowry acid and base always yields a salt and water.
  • There are several chemical aqueous reactions that form a gas byproduct.
  • Oxidation-reduction, or redox, reactions involve a transfer of electrons and a change in oxidation number.

Chemical Reactions II

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 0:10
  • Arrhenius Definition 1:15
    • Arrhenius Acids
    • Arrhenius Bases
  • The Bronsted-Lowry Definition 4:48
    • Acids Dissolve In Water and Donate a Proton to Water: Example 1
    • Acids Dissolve In Water and Donate a Proton to Water: Example 2
    • Monoprotic Acids & Polyprotic Acids
    • Strong Acids
    • Bases Dissolve In Water and Accept a Proton From Water
    • Strong Bases
  • The Autoionization of Water 17:42
    • Amphiprotic
    • Water Reacts With Itself
  • Oxides of Metals and Nonmetals 20:08
    • Oxides of Metals and Nonmetals Overview
    • Oxides of Nonmetals: Acidic Oxides
    • Oxides of Metals: Basic Oxides
  • Oxidation-Reduction (Redox) Reactions 25:34
    • Redox Reaction Overview
    • Oxidizing and Reducing Agents
    • Redox Reaction: Transfer of Electrons
  • Oxidation-Reduction Reactions Cont'd 29:55
    • Oxidation Number Overview
    • Oxidation Number of Homonuclear Species
    • Oxidation Number of Monatomic Ions
    • Oxidation Number of Fluorine
    • Oxidation Number of Oxygen
    • Oxidation Number of Chlorine, Bromine, and Iodine
    • Oxidation Number of Hydrogen
    • Net Sum of All Oxidation Numbers In a Compound
  • Oxidation-Reduction Reactions Cont'd 38:19
    • Let's Practice Assigning Oxidation Number
    • Now Let's Apply This to a Chemical Reaction
  • Summary 44:19
  • Sample Problems 45:29
    • Sample Problem 1
    • Sample Problem 2: Determine the Oxidizing and Reducing Agents
    • Sample Problem 3: Determine the Oxidizing and Reducing Agents

Transcription: Chemical Reactions II

Hi, welcome back to Educator.com.0000

Today we are going to be talking about additional chemical reactions; especially those that happen in aqueous solution.0003

Let's go ahead and go over the lesson overview; we are going to cover a few more areas.0013

We are going to build upon what we discussed last time on chemical reactions.0019

The very first type of reactions that we are going talk about are going to be involving what we call acids and bases.0024

That is going to be our first really general category.0035

We are going to be then jumping into another category of acids and bases.0040

These are going to be nontraditional ones; something that we couldn't look at just by formula; I will build on that.0046

After that, we are going to then go on to what we call redox reactions.0058

We are going to learn how to identify when redox reactions occur and exactly what is occurring during the reaction.0064

Finally we go on to our summary and sample problems.0071

We are first going to tackle acids.0078

The first interpretation of what an acid and a base is is called the Arrhenius definition.0081

The Arrhenius definition tells us the following.0090

That acids are going to dissolve in water and dissociate into a proton and anion.0093

For example, we can take hydrochloric acid; this is going to dissociate in water.0100

This is going to form H+ which we call proton and its anion, Cl1-.0109

This is very familiar to what we talked about in the previous session on the first part of chemical reactions.0119

If you recall, we talked about what was labeled an electrolyte, strong and weak.0127

This is pretty much an electrolyte; the Arrhenius acids essentially behave as the electrolytes.0135

We can go ahead and look at another example.0144

HNO3 aqueous can also break up into H+ aqueous and NO31- aqueous.0147

Just want to do a reminder for strong electrolytes.0162

Remember we are going to use a single arrow pointing in the forward direction0165

to show we get full dissociation--that we completely break up into ions.0168

However there are also weak electrolytes.0175

There is going to be Arrhenius acids that are going to be weak electrolytes.0179

For example, hydrofluoric acid is the typical example.0183

If everybody tried to recall, remember this was called an equilibrium arrow.0187

This is going to form H+ aqueous and F- aqueous.0193

Now that we have covered the Arrhenius definition of an acid,0201

let's go ahead and cover the Arrhenius definition of a base.0204

According to Arrhenius, bases dissolve in water and dissociate into a cation and hydroxide.0208

In other words, Arrhenius bases are also electrolytes.0217

The typical example will be sodium hydroxide.0221

Sodium hydroxide is going to be a strong electrolyte; we use a single arrow.0228

That is going to dissociate and break up into its cation, Na1+, and hydroxide, OH1-.0233

There is also going to be other types of Arrhenius bases that don't completely dissociate.0243

Again we are going to use a double arrow for them.0249

We can imagine maybe the following; magnesium hydroxide aqueous.0253

This one again is not going to completely dissociate; we use an equilibrium arrow.0259

That is going to give us Mg2+ aqueous and two hydroxides aqueous.0264

The Arrhenius interpretation of an acid and a base is relatively straightforward.0273

Once again they simply are ionic compounds that break up into their respective cations and anions when in aqueous solution.0278

The Arrhenius definition is fine for most purposes.0289

However what we are going to cover now is called the Bronsted-Lowry definition.0292

The Bronsted-Lowry definition, the main difference is the following.0297

It always takes to in count the rule of water.0301

Let's first start off with the definition of acids.0306

Acids dissolve in water; what they do is the following.0309

They are going to donate a proton to water.0313

The first example that we talked about was hydrochloric acid.0318

Let's go ahead and talk about that once again.0321

This time interpreting it from the Bronsted-Lowry perspective; HCl aqueous.0324

This time we are going to follow the definition.0330

Instead of just dissociating into cation and anion, the acid is going to react with water.0333

It is going to donate a proton to water.0339

So HCl aqueous plus H2O liquid goes to the following.0342

If HCl loses its hydrogen, all that is left is going to be Cl1- aqueous.0347

If H2O is accepting the hydrogen from HCl, H2O becomes the following; H3O1+ aqueous.0356

There is a couple of things I want to point out here.0368

Number one, we give H3O1+ a unique name.0370

This is what we call the hydronium cation.0376

The hydronium cation, anytime you see it, it is always indicative that an acid is present in water; acid in water.0382

Basically a Bronsted-Lowry acid will always form hydronium in water; Bronsted-Lowry acids form H3O1+.0395

This is the Bronsted-Lowry interpretation of what an acid is.0408

Let's go ahead and look at another example.0414

Let's go ahead and look at HClO4; HClO4 aqueous.0421

Once again this is going to go ahead and react with water liquid.0428

That goes on to form ClO41- aqueous plus H3O1+ aqueous.0433

Always remember everyone that you always want to pay attention to balancing; don't forget.0444

You always make sure your equation is balanced; not only the coefficients but also the charges.0449

My overall charge on the reactant side is 0 and 0 here.0454

Overall here it is -1 and +1 on the right side giving me a net sum of 0 too.0461

By breaking up into hydronium and the anion, we are going to get the balanced chemical equation also.0471

There is a couple of different types of acids that we want to talk about right now.0480

Right now when you just see one hydrogen in the formula of the acid, that is what we call monoprotic acids.0485

But of course there are going to be many other types of acids that contain more than one hydrogen.0497

These are what we call polyprotic acids.0504

The typical example of a polyprotic acid would be sulfuric acid, H2SO4 aqueous.0510

H2SO4 specifically, this is called a diprotic acid; the prefix di meaning there is two of them.0520

Let's go ahead and react this with water; here is the rule of thumb.0527

For polyprotic acids, anytime you write them reacting with water, you never take off both of the hydrogens at once.0535

You always do one deprotonation at a time; for polyprotic acids, again one deprotonation at a time.0544

That means the following; I am not going to get just sulfate.0563

I am going to get just one proton is going to be taken off.0568

That is going to be HSO41- aqueous; again we always form H3O1+ aqueous.0571

The rule of thumb again, don't forget, it is going to be one deprotonation at a time.0586

Let's go ahead and take HSO4- and deprotonate it one more time.0597

That is going to be HSO41- aqueous plus H2O liquid.0603

We have another rule of thumb that we also have to go over.0610

For polyprotic acids, only the first step is going to be significantly strong.0614

Everything else is going to be considered a weak step.0621

In other words, HSO41- will be functioning as a weak electrolyte.0624

We are going to use therefore an equilibrium arrow.0635

We are going to get now sulfate plus H3O1+ aqueous.0639

This is the basic definition of Bronsted-Lowry acids.0648

Once again a couple rules to remember.0654

That Bronsted-Lowry acids are going to react with water and are going to donate a proton to water.0657

If we are dealing with polyprotic acids, you only write their deprotonations one step at a time.0662

Finally for the polyprotic acids, only the first step is considered a strong electrolyte step.0671

Everything else after that is going to be considered weak.0679

You may be wondering: how do I know what are the strong electrolytes for acids?0683

How do I know which ones are the weak electrolytes for acids?0688

We are going to go over that right now.0691

Basically the strong electrolytes are what we call strong acids.0693

Traditionally there are seven strong acids that are normally taught to a freshman level chemistry course.0700

Your instructor may give you more or less; so definitely pay attention to him or her.0706

But the basic seven that are always strong are the following.0712

Hydrochloric acid, HBr, HI, HNO3, H2SO4, HClO4, and HClO3.0716

Pretty much it is pretty safe to assume that any acid that is not on this list you can assume to be weak.0734

Again these are the seven strong acids that I would recommend definitely committing to memory.0742

You will always use a single arrow for them.0747

But remember, make a note, for sulfuric acid which is what we just looked at.0750

That is the polyprotic acid; only the first step is strong.0756

Now that we have talked about the Bronsted-Lowry definition of an acid,0762

let's go ahead and talk about how Bronsted-Lowry defines a base.0766

Bases dissolve in water and what they do, instead of donating a proton, they accept a proton from water.0773

The typical example of a Bronsted-Lowry base is going to be ammonia.0779

Ammonia... this is definitely something you should just have it memorized... is NH3.0785

Ammonia is going to react with water, H2O liquid.0790

It turns out that ammonia is going to be what we call a weak base.0794

Weak bases are going to be weak electrolytes; we are going to use an equilibrium arrow for them.0800

NH3 aqueous plus H2O liquid, it is going to accept the proton from water.0806

NH3 is going to become... remember your polyatomic ions... NH41+ aqueous.0812

H2O is going to lose a hydrogen this time forming hydroxide.0818

We see it pretty clear cut; Bronsted-Lowry acids always generate hydronium.0826

And Bronsted-Lowry bases are always going to generate hydroxide when reacting with water.0831

What are some other typical bases, Bronsted-Lowry bases, that you want to be on the lookout for?0840

Pretty much anything that is related to ammonia is going to be a weak base.0847

Specifically the molecules that I am talking about that are related to ammonia, these are going to be called amines.0861

Amines contain nitrogen and usually with carbon and hydrogen.0871

NH2R would be an example where R is a molecule that contains hydrogen and carbon; what we call a hydrocarbon.0890

Not only NH2R but NHR2; not only NHR2 but NR3.0904

These are the again generic formulas for amines; also Bronsted-Lowry bases.0913

Just by looking at the name, you can usually tell if it is going to be a base or not.0922

A lot of the organic bases that are also biomolecules, they all end in the same suffix.0926

What you want to look for is a ?ine suffix.0935

You actually know a lot already of examples that are weak Bronsted-Lowry bases that are also bimolecules.0942

For example, here is one of my favorite.0951

Of course it is going to be caffeine; you see this caffeine ends in ?ine.0956

Adrenaline is another one that ends in ?ine.0961

A lot of the drugs also end in ?ine; they are also bases and amines.0968

For example, morphine, cocaine, and heroine; the list goes on and on.0975

I am sure you will encounter these in some of your biology classes.0983

But just by looking at the name again, the ?ine suffix usually indicates if something is an amine.0986

What are the strong bases and what are the weak bases?0998

Pretty much the strong bases are going to be group 1 metal hydroxides and some group 2.1002

Sodium hydroxide, potassium hydroxide, lithium hydroxide, etc; those are the group 1 metal hydroxides.1023

Some of the group 2 metal hydroxides that are also considered strong are going to be calcium hydroxide and magnesium hydroxide.1032

Pretty much it is safe to assume... again you should check with your instructor about this.1045

It is safe to assume that anything that is not on this list of strong bases can be considered to be weak.1050

Especially ammonia and its derivatives which we call amines.1057

Now that we have talked about the Bronsted-Lowry interpretation of acids and bases...1065

We saw that they have one thing in common; that they all react with water.1071

But in one situation, that is with acids, water is accepting the hydrogen.1076

In the second situation, that is with bases, water is actually donating the hydrogen.1083

Let's take a look more closely at that.1089

Water has the unique ability of functioning as both a Bronsted-Lowry acid and also as a Bronsted-Lowry base.1093

This is what we call the amphiprotic nature of water.1099

Because of this, let's go ahead and further scrutinize water reacting with itself.1105

H2O liquid plus H2O liquid; this is going to be a weak process.1111

If we focus just on maybe the left water molecule... let's call that the acid.1124

Let's call the other water molecule the base.1135

If the acid is going to function as an acid, it is going to donate its hydrogen.1143

But if it donates a hydrogen, it becomes hydroxide.1149

The second water molecule which we are labeling as the base is going to accept the hydrogen.1157

All of a sudden it becomes H3O1+.1163

This process is known as the autoionization of water.1168

It is physically representative of what happens in normal water.1172

The water that you get out of your tap, it is not just H2O.1178

There is a little bit of hydroxide and there is a little bit of hydronium.1184

But the majority is clearly going to be water.1187

Again this is what we call the autoionization of water.1192

When we go ahead and look at quantitative acid-base chemistry,1200

down the line, this is going to come into play a lot more.1204

Now that we have covered the traditional acids and bases, Arrhenius and Bronsted-Lowry,1210

we are now going to look at some more acids and bases1218

that are going break the trend from Bronsted-Lowry and Arrhenius.1224

So far we learned that acids can be identified as basically having one or more hydrogens.1229

Just by looking at the formula; the hydrogens are usually at the beginning.1235

Bases were pretty easy to spot; they either contained hydroxide or they were amines related to ammonia.1240

There are molecules that fit neither profile.1249

Specifically we are going to now look at oxides.1254

Basically oxides are anything that contain oxygen.1257

Certain oxides of metals and nonmetals, they contain neither hydrogen or hydroxide.1260

However they can still react with water to produce hydronium or hydroxide.1266

What we are going to do now, we are going to spend some time on going over1275

what are called the acidic oxides and what are called the basic oxides.1279

Certain oxides of nonmetals... remember certain oxides of nonmetals are actually acidic.1286

That is they will dissolve in water to form hydronium.1294

Here I present to you three examples.1301

They are going to be the oxides of carbon, of sulfur, and of nitrogen.1305

Let's first look at the oxide of carbon, carbon dioxide.1315

Carbon dioxide can go ahead and react with water to form H2CO3; this is called carbonic acid.1318

Carbonic acid remember is a compound we learned before; it is a relatively unstable product.1329

It can go on to react with water to form bicarbonate and hydronium.1337

Essentially carbonic acid is going to function as a Bronsted-Lowry acid.1343

This process is very interesting; this is actually what is happening in your body.1351

If you take in too much carbon dioxide, the carbon dioxide is going to react with the water in our bodies.1356

It is going to form H2CO3; H2CO3 in turn is going to form hydroxide.1363

If we get too much carbon dioxide in our body, that is actually make our blood and etc more acidic.1367

That is going to be very very dangerous.1378

The next oxide is that of sulfur; that is sulfur dioxide.1381

Sulfur dioxide is a slightly different reaction.1387

Sulfur dioxide is going to react with molecular oxygen, O2, to form sulfur trioxide.1390

Sulfur trioxide is going to go ahead and react with water to form sulfuric acid.1397

You know sulfuric acid is one of the seven strong acids that we covered.1402

That is definitely going to go on to form hydronium.1408

Finally the last oxide is that of nitrogen; that is going to be nitrogen dioxide.1415

Nitrogen dioxide is going to go ahead and react with water.1421

We are going to go ahead and form a strong acid, nitric acid, and a weak acid, nitrous acid.1424

Again both of them are going to go on and form hydronium.1431

Again it is the oxides of carbon, sulfur, and nitrogen that are going to react with water somehow to give us hydronium.1436

That is what we call acidic oxides.1447

Now that we have covered acidic oxides, let's go on to basic oxides.1451

Basic oxides are going to be oxides of certain metals1455

because they are going to go on and react with water and to form hydroxide.1459

The typical examples are going to be group 1 and group 2 oxides; group 1 and group 2 oxides.1465

Here I want to present to you calcium oxide.1477

Calcium oxide is going to react with water to go ahead and form calcium hydroxide.1480

Remember calcium hydroxide was one of those relatively strong bases that we talked about.1486

That is going to go ahead and form Ca2+ and two hydroxides.1494

The other example is the group 1 oxides.1499

Sodium oxide is going to go ahead and react with water to form sodium hydroxide.1502

We saw that sodium hydroxide was one of those strong bases also.1506

It is going to go ahead and go on to form Na1+ and hydroxide.1511

We have learned the Arrhenius interpretation of an acid and base.1517

We have learned the Bronsted-Lowry interpretation.1520

Now we have learned that acids and bases will also include oxides of metals and nonmetals.1523

We now move on to the second half of this presentation on chemical reactions.1535

This type of chemical reaction is what we call oxidation-reduction reactions; which are also known as redox reactions.1542

In an oxidation-reduction reaction, basically one reactant is going to get oxidized and a second reactant is going to get reduced.1551

The most basic, the simplest interpretation of an oxidation-reduction reaction is one that involves the loss or gain of oxygen atoms.1564

That is pretty easy to remember; when you hear the word oxidation, it pretty much sounds like oxygen.1576

Let's go ahead and look at the following example.1583

We are going to take iron(III) oxide; we are going to react it with carbon monoxide.1585

We are going to form iron solid and carbon dioxide gas.1592

You see that one compound has gained oxygens and one compound has completely lost them.1597

Fe2O3 has lost oxygens; CO has gained oxygens.1605

Now to build upon our terminology, if a reactant gains oxygens, we say that it has become oxidized.1619

Therefore carbon monoxide became carbon dioxide; it was oxidized.1632

If a reactant is going to lose oxygen, we say that it has become reduced.1642

In this example, it is the iron(III) oxide that has become reduced.1648

Two more vocabulary terms; the oxidized reactant is what we call the reducing agent.1656

The reduced reactant is what we call the oxidizing agent.1663

Let's go ahead and move on then.1672

Again the simplest interpretation of a redox reaction is involving a loss or gain of oxygens.1677

Another interpretation involves a transfer of electrons.1683

The reducing agent is going to be the reactant that loses electrons while the oxidizing agent gains electrons.1691

This is very important that electron transfer is always from reduced agent to oxidizing agent in a redox reaction.1699

Let's go ahead and take a look at the following example.1714

Zinc solid plus Pb2+ is going to form is going to form Zn2+ and lead solid.1717

Look what happened here; zinc solid became Zn2+; Pb2+ became lead solid.1726

In other words, zinc gave up electrons and became a cation; Pb2+ gained electrons to get reduced.1737

In other words, look at the charge.1752

Zinc started off with a charge of 0 and it became 2+.1753

Lead 2+ started out here on the reactant side and it became 0.1760

We can easily tie this in to the activity series that we talked about in chemical reactions from the previous section.1769

Basically good reducing agents include those at the top of the activity series that are above hydrogen.1779

Basically the higher on the list, the better its reducing ability.1788

Now that we saw that we can interpret a redox reaction as a transfer of electrons, let's go ahead and talk about another one.1796

This is what we call a change in oxidation number.1806

This is going to be the less obvious one.1811

It is easy to see the change in number of oxygen atoms.1815

It is easy to see a change in ionic charge.1819

So this one we are going to spend a little more time on, takes a little getting used to.1824

Let's go ahead and identify and define what we mean by an oxidation number first.1829

Basically an oxidation number is going to be a pseudo charge.1833

It has no physical meaning and don't think of it as an ionic charge.1839

But it is a convenient way for us to keep track of electrons.1842

What you want to remember is the following.1849

During the redox reaction, an oxidation number is going to change for the compound that is involved.1851

If the oxidation number has increased, we say that the compound the element was a part of was oxidized.1859

If the oxidation number has gone down, then reduction has occurred.1868

There is a couple rules that we have to know in order to do this.1879

This is in order of priority; please remember that.1885

Rule number one: for all homonuclear species including monatomic, diatomic, and polyatomic, the oxidation number is 0 for each element.1891

For example, if you just see sodium solid, that is going to be a neutral atom sodium; its oxidation number is 0.1906

Let's look at some diatomic homonuclear species.1919

That would be Cl2 gas, Br2 liquid, I2 solid, N2 gas, H2 gas, and F2 gas.1923

These are what we call homonuclear diatomic compounds.1941

Each element in the compound has an oxidation number of 0.1946

There are some homonuclear polyatomic species; for example, sulfur.1953

The natural form of sulfur is actually S8 solid.1960

In S8 solid, all eight sulfurs, each have an oxidation number of 0.1966

That is the first rule; not too bad to remember.1975

The second rule is actually very straightforward too.1979

It tells us that for monatomic ions, the oxidation number is identical to the charge.1982

When we look at Na1+ aqueous, N3- aqueous, and O2- aqueous, these are what we mean by monatomic ions.1989

The ionic charge is equal to the oxidation state.2004

Next rule, this deals with fluorine.2010

Fluorine is the first element that gets very high priority.2013

Fluorine is going to be having an oxidation number of -1 when part of a heteronuclear compound.2017

For example, if you look at the molecule HF, this is a heteronuclear compound.2025

The fluorine here is going to have an oxidation number of -1; pretty straightforward.2034

The next element which gets a pretty big priority is going to be oxygen.2040

Oxygen is going to have an oxidation number of -2 which is very very common.2045

For example, in water, oxygen has an oxidation number of -2.2051

However there is going to be one main exception for oxygen; this is in what we call peroxides.2063

Remember peroxides we had you memorize for polyatomic ions; an example would be Na2O2.2069

In this case, for sodium peroxide, each oxygen is -1 oxidation number.2090

Once again peroxides is a big exception for oxygen.2103

Going on, after fluorine and oxygen comes the remainder of the halogens.2110

The halogens, chlorine, bromine, and iodine, are also pretty high on the list of rules here.2117

They are going to get oxidation numbers of -1.2127

Coming up is hydrogen; hydrogen is almost always a +1 oxidation number when part of a heteronuclear compound.2133

For example, if you look at HF which we just dealt with.2143

Heteronuclear compound so hydrogen is going to get a +1 oxidation number.2147

There is an exception however; this is when hydrogen is going to be in a compound that contains a metal.2151

When it is directly bonded to a metal; for example, NaH.2159

In this case, hydrogen is going to get an oxidation number of -1.2165

These types of compounds are what we call metal hydrides.2171

When hydrogen is part of a compound with a metal, what we call metal hydrides.2176

The final rule is going to be something very familiar to us.2183

Remember when we learned how to name ionic compounds.2189

We said that the overall charge, the net charge, must be equal to the overall charge of the compound itself.2193

Along a parallel track, the net sum of all oxidation numbers in the compound must add up to the overall charge.2200

This is very familiar; let's look at those examples we just talked about; HF.2212

We said that hydrogen is going to get +1 and fluorine is going to get -1.2220

HF is neutral overall so we get +1 and -1 adding up to 0.2227

If you look at NaH, Na is going to get a +1 charge.2232

Remember this is an ionic compound so we go by group 1; we just go by the charge.2241

Hydrogen we said is going to get -1; once again +1 and -1 giving us 0.2246

Let's do one final look; that is going to be H2O.2253

What get priority?--is it hydrogen or oxygen?2256

Going back to your rules, oxygen is going to get the priority.2260

That is going to have a -2 oxidation state.2263

Hydrogen here we said is going to get a +1 oxidation state when part of a heteronuclear compound.2266

However how many hydrogens are there?--there is two of them.2274

When we do the arithmetic, we are going to get +2 plus -2 adding up to 0 again.2278

This is again very familiar territory where the net sum of the oxidation numbers must equal to the overall charge of the compound.2290

Now that we have talked about the rules...2302

Unfortunately, these are rules you are just going to have to commit time to memorizing.2306

Let's go ahead and apply them to do some problems.2310

We are going to now assign oxidation numbers to all elements in each of these compounds.2315

Ammonia, nitrogen and three hydrogens; if you look at your list, hydrogen is going to get the priority.2324

Each hydrogen is +1; there is three of them giving me +3 overall.2334

Ammonia is neutral; that means nitrogen has to be high or low enough of an oxidation number to balance the +3 from hydrogen.2341

Therefore nitrogen is going to be -3 oxidation number.2352

Next one is a barium chloride.2356

We recognize this as a metal and a nonmetal; this is an ionic compound.2359

Barium is column 2; that is going to be +2.2364

Each chlorine, group 7, is -1; there is two of them.2368

2 plus -2 is going to give us 0; so we are good to go.2373

C6H12O6, this is a molecular compound, nothing but nonmetals.2379

This is actually the formula for glucose; let's take it one by one.2384

Oxygen gets priority; each oxygen is -2; there is six of them giving me -12 overall.2391

After oxygen, hydrogen is going to get the next priority.2401

There is +1 oxidation number times 12 hydrogens giving me +12 overall.2406

What does that mean for carbon then?2418

If this whole molecule is neutral, that means carbon has an oxidation number of 0 in this compound.2423

When I add all the numbers together, 0 times 6 added to +12 added to -12, I add up to 0.2431

Finally, oxygen difluoride, this is a molecular compound; oxygen difluoride.2440

Which one is going to get priority?--fluorine is going to get priority.2449

Each fluorine is -1; there is two of them giving me -2 overall.2452

That means the oxygen has to be +2 to balance the -2 oxidation state.2459

Now that we know how to assign oxidation numbers to elements within a compound, let's apply this now to the following chemical reaction.2469

We have CH4 plus 2O2 goes on to form carbon dioxide and water.2482

Let's go ahead and assign oxidation numbers to each element.2489

And see if there is any changes from the reactant side to the product side.2492

Carbon is in CH4; hydrogen is going to get priority.2498

The hydrogen is going to be +1; there is four of them which means carbon has to be -4.2504

Molecular oxygen is a homonuclear diatomic so each oxygen is going to be 0.2517

Carbon dioxide; in carbon dioxide, oxygen gets priority; each oxygen is -2.2525

There is two of them giving me -4 overall which means carbon has to be +4.2533

Finally water; oxygen gets priority; it is going to be -2.2540

Therefore each hydrogen in water must be +1 oxidation number to balance the -2 overall.2550

Have there been any elements in this problem that have undergone changes from reactant side to product side?2560

Let's go ahead and see; carbon changed from a -4 to +4.2565

In other words, carbon's oxidation number increased.2571

In addition, oxygen's oxidation number decreased from 0 to -2.2579

Carbon went up from -4 to 4.2592

Therefore, remember what we said for compounds that have an element whose oxidation number has gone up.2598

That compound we say has been oxidized.2606

So CH4 has been oxidized; it is the reducing agent.2610

Molecular oxygen has been reduced; it is the oxidizing agent.2626

This is how we, as you can see, use the concept of oxidation numbers2643

to help us determine exactly what is going on in a redox reaction,2650

what compound is being oxidized, and what compound is being reduced.2654

Let me go ahead and summarize this lesson on chemical reactions.2660

We covered many interpretations of what we mean by an acid and base.2665

The first one was called an Arrhenius interpretation.2672

The second one was a Bronsted-Lowry interpretation.2676

Right after Bronsted-Lowry, we also covered the common strong acids and bases that you will just have to memorize.2682

There is no other way of going about it.2690

In addition, we saw that acids and bases include oxides.2693

What is interesting about these oxides is that they contain neither hydrogen nor hydroxide in their formulas.2698

Yet they still react with water to form hydronium or hydroxide.2705

Finally the last type of chemical reaction we covered were called redox reactions.2711

Redox reactions basically involve a transfer of electrons.2716

We saw that we can use the concept of oxidation numbers2720

to help us keep track of the oxidizing agent and of the reducing agent.2725

Now that we have covered all this material, let's go ahead and tackle some sample problems.2734

This is again a very traditional type of problem.2739

Write out the reaction for nitric acid reacting with water.2742

Let's go ahead and do that; HNO3 aqueous plus H2O liquid.2748

You have to try to think back: is nitric acid one of the seven strong acids?2760

The answer is absolutely yes.2765

Because of that, we are not going to use an equilibrium arrow.2767

But we are going to use a single arrow in the forward direction.2770

That is going to go ahead and form hydronium of course and nitrate aqueous.2774

The second part of the question is the following: is this redox?2782

If this is a redox reaction, we have to have a transfer of electrons involved.2787

In order to determine that, we have to see if any element's oxidation number2794

has changed from the reactant side going to the product side.2801

Let's go ahead and assign oxidation numbers to all elements in this chemical reaction.2806

In nitric acid, oxygen is going to get priority.2814

That is going to be -2 by three.2817

Hydrogen is going to get the next one which +1.2822

Therefore in order to have everything add up to the overall charge which is 0, nitrogen has to be +5.2826

Let's go ahead and look at water; water is something we have done before.2836

Here oxygen is -2; here each hydrogen is +1.2840

That adds up to 0; so we are good to go on that.2847

Let's go ahead and do hydronium; oxygen is -2.2850

Hydrogen, when it is part of a heteronuclear compound, it is going to be +1; there is three of them.2856

Let's see if this adds up to the overall charge of hydronium which is just +1.2862

I have +3 and -2; I do get a sum of +1; it definitely works out.2867

The last one is nitrate; nitrate, we have -2 for oxygen times three of them.2875

Each nitrogen therefore has to be what?2882

You see that nitrate is -1 overall charge.2886

If I already have -6 from oxygen, nitrogen has to be +5.2891

Now that we have assigned all oxidation numbers, let's see if anything has changed.2898

Nitrogen remains +5, oxygen remains a -1, and all hydrogens remain +1.2904

Not a single oxidation number has changed; therefore this reaction is not redox.2912

Again this reaction is not redox.2923

The final two questions are focusing on redox reactions.2929

It is basically having you identify the oxidizing agent and the reducing agent.2936

That involves first assigning oxidation numbers once again; let's go ahead and do that.2941

We have silicon solid; that is a monoatomic homonuclear compound; so that is neutral, 0.2946

Two Cl2s, molecular chlorine is a homonuclear diatomic; both of them are 0s.2957

Now we reached silicon tetrachloride.2964

Once again chloride is going to be getting the priority here, -1; there is four of them.2968

The overall molecule is neutral; therefore silicon must be +4.2973

We have seen that silicon goes from 0 to +4 and that chlorine goes from 0 to -1.2980

This is also how you tell if you are doing it right.2992

Because when you do a redox reaction, you have to have the following combination always.2995

Exactly one element is going to have an increase in oxidation number and one is going to go down.3000

In this case, that is what we exactly have.3006

We say therefore that silicon, because it experienced an increase in oxidation number, was oxidized.3009

Molecular chlorine, each chlorine experienced a decrease so molecular chlorine was reduced.3019

That makes silicon the reducing agent and that makes molecular chlorine the oxidizing agent.3029

The last problem is a little longer chemical equation.3043

But again I want to present this to you because it doesn't make a difference how long the chemical reaction is.3047

You are always going to have one being oxidized and one being reduced.3052

Just follow the rules, assign the oxidation numbers, and you will be fine all the time.3056

We have sodium iodide; we recognize that immediately as metal and nonmetal; that is going to be ionic.3062

Therefore iodine is going to be -1; sodium is going to be +1.3068

Let's go ahead and look at sulfuric acid.3073

Each hydrogen is +1; there is two of them; the overall molecule is neutral.3083

So far I have 2 and -8 giving me -6 overall.3090

Which means sulfur must be +6 to help me balance out everything to 0.3095

Next one is manganese(IV) oxide.3102

Manganese(IV) oxide we recognize as being ionic also; metal and nonmetal.3105

Therefore each oxygen is -2; so manganese here is going to be +4 overall.3109

Moving on, sodium sulfate.3118

We recognize that as also being ionic; it is a metal and nonmetal.3121

Each sodium is going to be +1; there is two of them.3127

Each oxygen is -2; there is four of them.3132

Everything has to equal up to the overall charge which 0.3138

I have +2 and -8 overall which is going to be -6; that means sulfur has to be +6.3141

Next compound, manganese(II) sulfate.3152

Once again this is cation and anion so we just go by ionic charges.3157

Each oxygen is -2 overall; there is four of them.3164

Sulfur is going to be +6; how did I get that?3172

We know from our polyatomic table of ions that you memorized, sulfate is -2.3177

If oxygen is already -8, that means sulfur has to be +6 to give me a -2 overall charge.3183

Therefore manganese... we know that sulfate is -2 so manganese has to be +2.3194

Two more compounds; each iodine is going to be 0; that is homonuclear diatomic.3203

Finally in water, each oxygen is -2 and here each hydrogen is +1; there is two of them.3209

It looks like a lot of material, but let's go ahead and see if anything has changed.3218

Sodium, has sodium changed?--the answer is no.3225

Sodium started off +1; in the product side, it has remained +1; so sodium is not involved here.3230

The next one is iodine; iodine started off -1 and iodine became 0.3242

You see the oxidation number of iodine has changed.3250

In fact we say that I- has become oxidized; therefore it is the reducing agent.3255

We have one more left; we have to now find which element or compound became reduced.3273

That is going to be, looking here, that is going to be manganese.3281

Manganese started off +4 and manganese became +2; look, it went down.3287

We would be in trouble if we had two elements go up or if we had two elements go down.3296

Remember you always have one of each for a redox reaction to occur.3301

In this case, manganese Mn4+ was reduced, making it the oxidizing agent.3307

We are going to talk about redox reactions quantitatively down the line later on in this class.3323

We will come back to this before you know it.3331

Thank you for using Educator.com again; I will see you all next time.3337

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