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Franklin Ow

Franklin Ow

Chemical Reactions

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 (32)

0 answers

Post by Peter Ke on September 7, 2015

At 36:02 where did you get the subscript in Br2 from?

2 answers

Last reply by: Khalid Khan
Thu Jun 25, 2015 2:47 PM

Post by Khalid Khan on June 25, 2015

Hi Mr. Ow,

In aqueous reactions, I don't see balanced chemical equations. Is this supposed to happen?

Thanks!

0 answers

Post by Khalid Khan on June 24, 2015

Mr. Ow,

How would someone figure out if they should put a single arrow or two arrows with opposite direction?

Thanks!

1 answer

Last reply by: Professor Franklin Ow
Wed Mar 4, 2015 11:41 AM

Post by Micheal Bingham on March 3, 2015

Since an Acid and Base form water, is it possible to drink an Acid and Base reaction?

1 answer

Last reply by: Professor Franklin Ow
Wed Mar 4, 2015 11:39 AM

Post by Kate Danielle Rapinan on December 5, 2014

My textbook says that hydroxides of larger group 2 ions(Ca 2+) and down are soluble, but not Be and Mg. (In addition to group 1) What is your take on this, Mr. Ow?

1 answer

Last reply by: Professor Franklin Ow
Wed Mar 4, 2015 11:39 AM

Post by Kate Danielle Rapinan on December 5, 2014

Hello Mr. Ow,
If you are balancing equations in which all the reactants are heteronuclear, are there any rules on which one to begin balancing first?

1 answer

Last reply by: Professor Franklin Ow
Sun Oct 12, 2014 11:48 AM

Post by Saadman Elman on October 12, 2014

My Chemistry book clearly documented that All SO4 is soluble EXCEPT, Ca, Sr, Ba, Pb, Ag, Hg. With that being said, Ag SO4 can't be aqueous but it is a solid/precipitate/insoluble.

1 answer

Last reply by: Professor Franklin Ow
Sun Oct 12, 2014 11:47 AM

Post by Saadman Elman on October 12, 2014

In sample problem 1, no. 4 reaction is incorrect. The Correct formula of Silver Sulphate is Ag2 SO4 NOT Ag SO4.

Then you said Ag SO4 is aqueous.  But Ag SO4 is NOT aqueous. You can double check it. I am 100% sure that it is definitely not aqueous. It's solid, in another words insoluble.

Overall, the lecture was great as usual..

1 answer

Last reply by: Professor Franklin Ow
Fri Oct 10, 2014 9:33 AM

Post by David Gonzalez on October 9, 2014

Hi professor Franklin, hope all is well! I have a problem that's really baffling me! Basically, the problem says to balance C4H10 + O2 = CO2 + H2O. I worked out the answer C4H10 + 13 O2 = 4 CO2 + 5 H2O. What am I doing wrong???? It seems balance, right?!

Thank you in advance.

2 answers

Last reply by: david faizi
Wed Sep 24, 2014 11:07 AM

Post by david faizi on September 22, 2014

Hello,
On the sample 1 problem for KNO3 + CuCl; I don't understand why the subscript of the Cl is not on the product side with Cl as well as, why the NO3 becomes an (NO3)2.

4 answers

Last reply by: david faizi
Wed Sep 24, 2014 11:07 AM

Post by Danny Fanny on August 16, 2014

Hello Mr. Ow:

I have a issue regarding question #4 on your "Sample Problem 1" slide. The answer you give is Li2SO4 + 2AgNO3 => 2LiNO3 + AgSO4. However, if I am not mistaken,the balanced equation should be 2Li2SO4 + 2AgNO3 => 2LiNO3 + 2Ag2SO4. In addition, I believe that silver(II) sulfate, one of the products of this reaction, should be written as Ag2SO4, not AgSO4 as you had written it. If I have misunderstood this problem, please let me know!


Thanks,
   Danny

2 answers

Last reply by: Professor Franklin Ow
Tue Aug 5, 2014 7:53 PM

Post by William Kinne on August 5, 2014

On sample 1 problem three shouldn't it be 2KNO3(aq)+ CuCl2 (aq) ---> 2KCl(aq)+Cu(NO3)2 (aq)

1 answer

Last reply by: Professor Franklin Ow
Mon Jun 2, 2014 12:42 PM

Post by jared vitt on June 2, 2014

bumbleing my way through various lessons to see what interests me, thus knowing very little, dosent there need to be a catlist for chemical reactions? how it the catlist noted?

1 answer

Last reply by: Professor Franklin Ow
Fri Feb 7, 2014 10:21 AM

Post by Laura Mejia on January 31, 2014

Hi Mr. Ow,

how do you know the state of the reactants?

Chemical Reactions

  • The need to balance chemical reactions follows from the law of conservation of mass.
  • Electrolytes or electrolytic solutions contain free ions in solution.
  • Strong electrolytes dissociate completely, while weak electrolytes hardly ionize.
  • When predicting if a precipitate will form during an aqueous reaction, it is useful to refer to a table of solubility rules.
  • A net ionic equation only shows the formation of the nonaqueous species from its constituent ions.

Chemical Reactions

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

Transcription: Chemical Reactions

Hi, welcome back to Educator.com.0000

Today's lesson in general chemistry is on chemical reactions; let's go over the outline.0002

We are first going to talk about something very fundamental to all of the physical sciences.0011

That is the law of conservation of mass.0017

We are going to see that in order to obey the law of conservation of mass,0020

anytime we have a chemical equation, we have to make sure that it is completely balanced.0024

After discussing this, we will then go on to something we call chemical equilibrium.0031

Because chemical equilibrium is going to be very applicable to our discussion of reactions in solution.0037

The solutions that we are specifically talking about are electrolyte solutions, nonelectrolyte solutions, and etc.0048

After talking about electrolytes and nonelectrolytes, we are then going to dive into a couple of different reactions that can occur in aqueous solutions.0058

The first one is what we call single replacement reactions.0067

The second one is what we call double replacement reactions.0072

The third one is neutralization reactions.0076

The fourth and final one is going to be reactions that produce unstable compounds0078

that are going to further decompose and form other compounds we cannot initially predict.0084

When this section is done, we are going to be able to predict0092

the compounds that are formed from any of these types of reactions.0095

We will wrap up the session with of course our summary followed by a pair of sample problems.0101

Let's go ahead and start.0107

The law of conservation of mass is fundamental to all of the physical sciences.0110

We always will be obeying this in chemical reactions.0114

Basically the law of conservation of mass tells us the following: that matter can neither be created nor destroyed.0118

In other words, what you have, you are going to have permanently.0126

It is coming from other sources.0131

You are never going to truly eliminate it or get rid of it.0134

But matter as we know can be transformed into different states, etc.0140

If we were to translate this into a chemical reaction before and after, the law of conservation of mass basically tells us the following.0146

That the total mass before a reaction is going to be equal to the total mass after a reaction.0153

In other words, the total mass of what you start with equals to the total mass of what you end up with.0162

Let's go ahead and look at a typical way of representing a chemical reaction.0170

Basically the general format of a chemical reaction is the following.0175

We have the reactants on the left side of the arrow.0180

We have the products on the right side of the arrow.0184

We can go ahead and verbalize this the following way.0186

That A combines with B to form C and D; or A and B yields C and D.0190

Or A and B mix C and D; whatever you are more comfortable with.0197

Basically if we were to apply the law of conservation of mass, this is telling us the following.0204

That the total mass combined of A and B will be equal to the total mass combined of C and D.0209

In order to obey the law of conservation of mass, we have to do what is called balancing a chemical equation.0222

You are going to be using this from now on pretty much in every lesson.0230

This is one of the most fundamental parts of general chemistry that you want to learn to master early on.0235

When we balance a chemical reaction, let's go ahead and look at the following.0242

N2 plus H2 goes on to form NH3.0246

We want to look at the number of elements on both sides of the equation.0252

For example, here we notice that there are a total of two nitrogens on the left side.0255

But there is only one nitrogen on the right side.0262

In addition on the reactant side there are two hydrogens.0266

There are three hydrogens on the right side.0271

We say that this is an unbalanced chemical reaction.0274

We never want to leave a chemical equation unbalanced.0280

The way to go about this is the following.0285

We are going to put what are called coefficients right in front of each of the molecules to go ahead and balance this.0288

The coefficients are usually going to be whole numbers.0301

I want to give you a couple tips.0304

Whenever the reactants are all homonuclear, it doesn't matter which element you try to balance first.0306

What I mean by homonuclear is that it is basically a compound composed of just one element; one element type only.0312

Examples of homonuclear would be N2, H2, O2, Cl2, etc.0326

Those again are what we call homonuclear.0335

A lot of the time, it is trial and error; so let's go ahead and start.0339

N2 plus H2 goes on to form NH3.0343

What I would like to do, I would like to maybe balance nitrogen first.0349

Again it doesn't matter what order you do it in.0352

I am going to go ahead and put a 2 in front of NH3.0356

Then I ask myself: I have now two nitrogens on the left.0361

I have now two nitrogens on the right; the last thing to fix is hydrogen.0364

Now I have two hydrogens on the left; I have six hydrogens on the right.0370

Basically we ask ourselves what multiplied by two is going to give us six.0376

The answer is going to be three.0382

My balanced chemical equation is going to be N2 plus 3H2 goes on to form 2NH3.0384

Always, this is very important, always double check your work to make sure that you have done the balancing correct.0393

As you can see, we have two nitrogens on both sides of the equation.0397

And we have a total of six hydrogens on both sides of the equation.0401

Before we move on, another thing I just want to point out is the following.0409

When you have the coefficient and a subscript, we just multiply them together to get the actual number.0413

That is how we get six hydrogens there; coefficient multiplied by the subscript.0419

Let's go ahead and now move on to another chemical reaction where the reactants are not both homonuclear.0432

For example, CH4 plus O2 goes on to form CO2 and water.0440

We have CH4 which is going to be what we call heteronuclear.0449

That is the compound is made up of more than one element.0453

My advice I usually give is the following: you always save the homonuclear reactant for the last step.0457

That is going to make your life a lot easier.0469

Let's go ahead and try to balance this.0471

Again it is CH4 plus O2 goes on to form CO2 and H2O.0474

I am going to save oxygen for last because it is part of a homonuclear compound.0481

I am left with either carbon or hydrogen to start with.0485

Again it doesn't matter for those; let's just go ahead and pick carbon.0488

Carbon, we have one on each side of the equation.0493

We have one in the CH4 and we have one in the carbon dioxide; so carbon is already balanced.0496

The only thing remaining is hydrogen now to work with before we tackle oxygen.0502

On the left side of the equation, I have a total of four hydrogens.0507

On the right side of the equation, I have a total of two hydrogens.0510

I am going to put 2 right in front of water.0515

Now my hydrogens are balanced; I have four on both sides of the equation.0518

Remember multiply the coefficient by the subscript, two times two gives us four.0522

Now we can go ahead and tackle oxygen.0527

I have a total of two oxygens on the left.0529

I have a total now of two oxygens in carbon dioxide.0534

I have a total of two oxygens in water; remember to carry through the coefficient.0538

So actually I have four oxygens total on the right side of the reaction.0545

I am going to put a 2 right in front of O2.0553

That gives me a total now of four oxygens on both sides of the reaction.0556

That is it; that is our balanced chemical reaction.0560

Another piece of advice I usually give students is the following.0565

Balancing chemical reactions can be very frustrating.0569

You don't always get it right the first time; a lot of it is trial and error.0572

If going by one element the first time doesn't work, try another element until you get it.0577

But remember it always can be balanced.0582

It is a fundamental law of nature; it always can work.0585

Right now, I want to introduce you a chemical reaction where sometimes the coefficients0593

that we are going to use initially are not going to be all whole numbers.0599

Instead the coefficients that we are going to be using are going to be fractional.0603

Remember, remember what we have to do.0609

It is going to be the coefficient times the subscript.0613

That is going to give us the total number of that element in the compound; the total number of element in compound.0620

Let's go ahead and balance this; C2H6 plus O2 goes on to CO2 and water.0633

I have two carbons on the left side; I only have one carbon on the right side.0639

I am going to go ahead and put a 2 in front of CO2; now my carbons are fixed.0645

Remember that I am leaving oxygen for last because oxygen is part of a homonuclear compound.0649

What we have to do now is fix hydrogens.0655

I have six hydrogens on the left; I have two on the right side.0658

Then we ask ourselves: what coefficent, what number times two is going to give me six?0663

That is going to be a three; so now my hydrogens are fixed; six on each side.0669

Now we have the following C2H6 plus O2 goes on to form 2Cl2 and 3H2O.0675

Let's go ahead and look at the oxygen count right now.0685

I have two oxygens on the left side.0688

Here on the right side, I have a total of four oxygens in CO2.0691

I have a total of three oxygens in water.0695

That means I have a grand total of seven oxygens on the right side and only have two on the left side.0700

This is when we have a case of an even-odd number.0710

If you look at the previous examples, we have always had an even number on one side and an even number on the other side.0713

In that case, it is always convenient just to use a whole number.0720

But when you have an even-odd pair like this, there is going to be a little trick we are going to do.0724

Basically you ask yourself the following.0729

You see that I am shy of oxygens on the left side.0732

We have to ask ourselves: what coefficient times two is going to give me seven?0736

Basically 2 goes down stairs; 7 goes upstairs; it is very simple.0743

When you have an even-odd pair, you will always use a fractional coefficient.0750

Basically the bottom number is going to be the subscript from the homonuclear compound.0757

The top number is always going to be the greater number; what you are trying to get the homonuclear compound up to.0771

Once again when you have an even-odd pair, use fractional coefficients.0785

The fractional coefficient I am going to use is 7/2.0789

Remember a few sessions back, we talked about Dalton's atomic theory.0796

Part of his theory was that elements combine in small whole number ratios to form compounds.0802

What that means is the following: we never want to leave a chemical equation in a fractional form.0808

In other words, what we are going to do right now, we are going to eliminate the fraction.0816

We are going to be left with nothing but whole numbers.0819

The fraction we have 7/2; if we want to get rid of a fraction, we simply divide by its denominator.0825

But I just can't multiply oxygen by 2; if I change one element, I have to do it proportionally to everything.0833

My grand and final answer is going to be the following: it is going to be 2C2H6 plus 7O2.0842

That is going to go on and form four carbon dioxides and six waters.0858

This gives us our balanced chemical equation with no fractions.0877

Let's just go ahead and do one final check, make sure we have done it right.0880

Let's go ahead and see if we have the same number of elements on each side.0883

I have 4 carbons total here; 4 carbons total here; carbons are good to go.0889

I have 12 hydrogens here; I have 12 hydrogens here; hydrogens are good to go.0895

Finally I have 14 oxygens here.0903

I have 8 oxygens in CO2 and 6 oxygens in the water; giving me a grand total of 14 oxygens.0907

As you can see, this method does work.0920

Again the final thought is to never leave a coefficient as a fraction.0923

Now that we have learned how to balance chemical equations, let's move on to the next topic of the session.0932

This is what we call chemical equilibrium.0939

When an ionic compound fully dissociates... what I mean by fully dissociates is that it completely breaks up.0944

It forms individual cations and anions.0951

Whenever this happens, we use a single arrow in the forward direction.0954

For example, hydrochloric acid is going to be fully dissociating when dissolved in water.0959

That is going to break up into its respective cation and anion.0966

But not all ionic compounds do this.0972

A lot of ionic compounds really incompletely dissociate which means they don't fully break up into cation and anion.0976

When this occurs, we use a new type of arrow.0985

It looks like a pair of opposite arrows.0991

This is what we call an equilibrium arrow.0995

What the equilibrium tells us is that the reaction is reversible.1004

In other words... let's look at HF, as hydrofluoric acid breaks apart.1019

In other words, as we move in the forward direction, H+ and F- instantaneously form.1024

However look at the reverse arrow; let's go ahead and translate that.1032

What that is telling us is that H+ and F- are going to somehow find each other again.1037

They are going to recombine and reform the original hydrofluoric acid molecule.1042

This is what we call dynamic equilibrium.1048

That is it is a constant process; it is not static whatsoever.1052

Basically both the forward and the reverse reactions happen simultaneously.1057

In other words, as soon as a compound dissociates, the ions are going to recombine into the original compound.1063

Again you are going to call this an equilibrium arrow.1072

We are going to be using this quite heavily in the latter half of this general chemistry class.1075

Now we are going to talk about specific compounds when they dissolve in water.1085

They are called electrolytes and nonelectrolytes.1091

When you hear the word electrolytes, it obviously sounds like electricity or electron.1094

Basically electrolytes are going to be ionic compounds that form ions when dissolved in water.1101

For example, we can go ahead and take a salt, table salt, sodium chloride.1107

Sodium chloride is going to dissociate and form the respective ions; sodium cation and chloride anion.1112

The following is going to be a fundamental requirement for a solution to be conductive.1122

Basically a solution has to have free ions in order to conduct an electric current.1128

That is why they call them electrolytes.1133

Strong electrolytes completely dissociate.1138

For strong electrolytes which completely dissociate, we are going to use a regular single arrow in the forward direction.1143

Weak electrolytes incompletely dissociate.1151

Remember going back to the previous slide, for ionic compounds that incompletely dissociate, we use an equilibrium arrow.1155

Finally there are compounds that when dissolved in water, they do not form ions whatsoever.1162

These are going to be nonelectrolytes; these are going to be typically molecular compounds.1171

For example, if you take ordinary glucose for example, C6H12O6; you go ahead and dissolve it in water.1176

You don't form individual ions; it just becomes aqueous.1190

You see that there are no ions whatsoever.1197

Glucose is going to be a nonelectrolyte solution.1200

Now that we have gotten those vocabulary terms out of the way,1205

what we are going to do right now, we are going to study specific types of aqueous reactions.1208

We are going to learn how to predict the products.1215

The first type of aqueous reaction is what we call a single replacement reaction; also known as single displacement.1218

A single replacement reaction is always going to have the following general format.1228

A... this is usually going to be a metal... plus B and C.1233

Remember this is going to be the metal; this is going to be the nonmetal; usually.1240

Or the cation and the anion goes on to form AC and B.1247

You notice that A has come in and essentially displaced B from its bond with C.1252

We go on to form AC and B; this is called a single replacement.1259

Let's go ahead and look at a couple of examples.1263

You take lithium metal and you react it with copper(II) chloride.1266

Lithium is going to come in and knock out copper; we form lithium chloride instead.1273

Copper is also going to come out as copper solid; however let's reverse the process.1279

If we take copper solid and go on and react it with lithium chloride, we get no reaction at all.1286

Let's go ahead and see why this is going to occur and how you are going to be able to predict this.1295

All of the experiments have been done already.1301

The results are summarized in what is called an activity series.1303

Basically an activity series is the following.1307

You see that lithium is right at the first element.1310

We call lithium a very active metal.1314

It can come in and knock out essentially any other element.1318

Let's look at sodium; the way to use this activity series is the following.1325

Sodium can knock out any element below it, but it cannot knock out any element above it.1329

For example, sodium will not be able to knock out calcium.1335

But it will be able to displace magnesium in a single replacement reaction.1340

Depending on your instructor, you may or may not have to memorize this activity series; definitely be on the lookout for that.1346

Another very common type of single replacement reaction occurs between the following: a metal and acid.1353

For example, when zinc solid is going to get dissolved in hydrochloric acid, we get the formation zinc chloride and hydrogen gas.1363

Essentially zinc has come in and knocked out hydrogen.1373

Let's go ahead and see if we can find hydrogen on this activity series.1377

We know that zinc is going to be right here and right below it is going to be H2.1381

Yes, this activity series does predict that zinc is going to come in and displace hydrogen in a single replacement reaction.1389

Here is the thing; when hydrogen gets displaced, it always forms hydrogen gas.1398

When displaced in a single replacement reaction, it is going to form H2 gas.1405

This is how we use an activity series again for single replacement reactions.1417

The next type of reaction we want to talk about is going to be called a double replacement reaction.1422

It is also known as double displacement.1428

Basically instead of one quote and quote knockout, two occur.1432

We are going to have two swaps; I always like to show brackets with this.1437

Basically A is now going to hook up with D and now C is going to hook up with B to form AD and CB.1443

Please note how the cation always hooks up with the other anion, etc.1452

An example of this is going to be the reaction between sodium chloride and lead(II) nitrate.1459

Sodium is the cation; it is going to hook up now with the other anion which is nitrate.1465

Lead, Pb2+, is going to hook up with the other anion chloride.1471

We get sodium nitrate and lead(II) chloride.1474

Please note... this is very important when you do this... that the charges never change, especially for the transition metals.1479

Remember when we went back to nomenclature and you learned that transition metals can have different oxidation numbers.1488

Here this is going to be Pb2+; there should be really a 2 there; sorry about that.1496

Here in the product side, it is going to remain Pb2+.1505

Please do not ever change a metal's charge in a double displacement reaction.1509

Pb2+ is going to remain Pb2+; that is why it is PbCl2 and not just PbCl for example.1523

The next type of equation though, something very related to this, is the following.1532

That is called a net ionic equation.1537

A net ionic equation is usually written to show the formation of the precipitate.1539

A precipitate is just a fancy name for the solid that is formed in a reaction.1544

Basically a net ionic equation shows the formation of the precipitate, in this case, lead(II) chloride, from its constituent ions.1550

That is basically going to be Pb2+ combining with two chlorides.1558

It is always aqueous plus aqueous goes on to form the solid.1563

Now that we have covered net ionic equations, how do you predict that a compound is going to be solid?1570

Or how do you predict if it is going to remain aqueous in these reactions?1579

This is something you may or may not have to memorize also.1584

So please be on the lookout for that from your instructor.1586

This is what we call the solubility rules for ionic compounds.1591

Basically this table is divided into two areas; first the compounds that are usually soluble.1595

What soluble means is that the compound is going to remain broken up into ions and dissolved.1602

Something that is soluble is going to be essentially aq or aqueous in the chemical reaction.1608

You are going to find ammonium salts, group 1 salts, nitrates, etc.1616

What is important in this solubility table is to pay attention to the exceptions.1623

For example, acetates are usually soluble; for example, sodium acetate will be aqueous in water.1627

However there are exceptions such as silver, Ag1+, and aluminum, Al3+.1635

While sodium acetate is going to be aqueous, silver acetate and aluminum acetate are exceptions.1640

They are going to be solid; that is they are going to be insoluble.1647

So you always have to pay attention to the exceptions.1654

The same goes for the other side of this table; compounds are usually insoluble.1657

They are going to be remaining intact and not break up relatively speaking.1662

They are going to remain as solids; so carbonates, phosphates, etc, are going to be solids.1667

However there are exceptions.1673

Something like magnesium carbonate is going to be mostly a solid in water.1676

But an exception is ammonium; ammonium carbonate is going to remain aqueous in water.1681

Let's go ahead and apply another type of double displacement reaction.1688

That is what we call a neutralization reaction.1697

A neutralization reaction always has the generic formula.1700

An acid is going to react with a base.1703

It is always going to form a salt, water, and it is always going to give off heat.1706

Anytime you mix an acid and a base, it is going to feel warm to the touch.1713

A typical reaction would be between hydrochloric acid and sodium hydroxide.1720

Neutralization reactions are going to be typically double displacement reactions.1725

Hydrogen is going to hook up with hydroxide.1729

Chloride is going to hook up with sodium.1733

Let's do the salt; the salt is going to be sodium chloride.1736

From the table of solubility rules which we just covered, this type of chloride is going to be soluble which is aqueous.1742

You see that when hydrogen combines with hydroxide, that is essentially going to form water.1750

That is going to be a liquid.1757

There we have it; a double displacement reaction in the form of what we call a neutralization reaction.1759

Let's do one last example; that is going to be between sulfuric acid and potassium hydroxide.1765

Sulfate is going to hook up with potassium.1772

Again hydrogen is going to hook up with hydroxide.1775

This time the salt is going to be potassium sulfate.1778

I need two potassiums to balance sulfate's 2- charge.1786

Look at the table of solubility rules; that is going to be aqueous.1790

Again I am going to always form water; let's go ahead and balance it.1794

I am going to need two potassium hydroxides to go ahead and react with the sulfuric acid.1799

I am also going to form two waters as a result.1808

Remember always balance your chemical reactions.1811

We saw that a neutralization reaction is a type of double displacement reaction.1817

There are a few exceptions where using the table of solubility rules1823

and simply going by a generic format for double replacement does not work.1829

These three reactions here are going to form compounds initially that are unstable.1835

These unstable compounds are going to in turn produce additional products which we just can't predict.1842

We are going to now introduce these to you.1850

The first type, number one, is going to be the reaction between a metal carbonate...1852

This is usually going to be a group 1 metal... reacting with an acid, HA.1858

Usually A is going to be the following: Cl, Br, iodine, nitrate, and also can be sulfate.1868

Sulfuric acid, nitric acid, hydrochloric acid, hydrobromic acid, hydroiodic acid, etc.1885

These are what we call the strong acids.1894

Again depending on your instructor, you may have to know a number of different ones.1897

So please definitely be on the lookout for that.1902

When we do the double displacement format, we are going to get H2CO3, carbonic acid, in combination with an HA aqueous.1905

It doesn't stop there; what happens is the following; that carbonic acid is actually unstable in water.1915

It is going to decompose to form CO2 and liquid water; that is one example.1920

Again a metal carbonate reacting with a strong acid goes on to form carbon dioxide gas and liquid water.1928

Similar to metal carbonate reacting with acid, we are also going to see metal sulfites reacting with acid.1937

Very similar, we are going to form instead of H2CO3, we are going to get H2SO3, sulfuric acid, in combination with the salt.1945

Sulfuric acid like carbonic acid is unstable in water.1955

It is going to form not CO2 but SO2 in combination with water.1959

Once again it is going to be a group 1 metal sulfite1965

reacting with strong acid to give you sulfur dioxide gas and liquid water.1968

The third and final example that is an exception is going to be the following.1975

It is going to be an ammonium salt where X is usually a halide.1980

It is going to be the ammonium salt reacting with a metal hydroxide.1988

Again this is usually going to be a group 1 hydroxide.1994

When you do the double displacement, the M is going to hook up with X.1998

Ammonium is going to hook up with hydroxide.2003

It turns out that ammonium hydroxide is unstable in water.2006

It is going to decompose and form aqueous ammonia and liquid water.2011

Once again the third exception, something we just can't predict, is the following.2017

It is going to be an ammonium halide salt in combination with a group 1 metal hydroxide.2021

You are going to form ammonia, NH3, in combination with liquid water.2029

That pretty much covers solutions that are occurring in aqueous environments.2036

Let's go ahead and summarize what we learned.2043

We started off talking about the law of conservation of mass.2046

Remember a big implication of the law of conservation of mass is that we always have to balance any chemical equation.2049

We went on to discuss electrolytes.2059

We learned that strong and weak electrolytes are going to each dissociate to different extents when dissolved in water.2061

We learned several rules; we saw an activity series; we learned the table of solubility rules.2071

We learned the three exceptions that we can use to help us predict the product2079

of any reaction that occurs in aqueous solutions, especially single replacement, double replacement, and neutralization reactions.2085

Let's go ahead and tackle some sample problems.2097

Here I have presented to you the reactants.2100

I want you to come up with the products and of course always balance the chemical equation.2104

Let's go ahead and look at this.2110

We have a zinc carbonate, that is going to be a metal carbonate, reacting with a strong acid, H2SO4.2113

That is one of the exceptions; what is going to happen is the following.2123

We are going to get zinc sulfate.2127

That is going to be aqueous; use your table of solubility rules.2134

This is going to form not H2CO3 but H2O liquid and CO2 gas.2138

Remember H2CO3 is going to be one of those unstable compounds in water.2145

H2CO3 we saw breaks up into liquid water and CO2 gas; let's go on.2153

Next one, ammonium bromide reacting with lead(II) acetate.2163

In this case, lead is going to now hook up with bromine.2172

Ammonium is now going to hook up with acetate.2178

This is yet again another double displacement reaction.2182

When we go ahead and draw the products, let's go ahead and do ammonium now with acetate, NH4C2H3O2, and lead(II) bromide.2186

Remember Pb2+ remains Pb2+; let's not forget the physical states.2201

Ammonium acetate is going to be aqueous.2208

Lead(II) bromide is going to be the solid.2211

Most importantly always balance your chemical equation.2214

I am going to need two of these and I am going to need two of these for my balanced chemical equation.2217

Next one, potassium nitrate aqueous reacting with copper(II) chloride.2226

This is going to be another example of a double displacement reaction.2234

Nitrate is going to hook up with copper.2241

Potassium is going to hook up with chlorine.2244

We are going to get KCl; when you look up the table of solubility rules, that is going to be aqueous.2249

Forming copper(II) nitrate; when you look up the table of solubility rules, that is also going to be aqueous.2258

This is an example of a double displacement reaction that forms no precipitate which is not too uncommon.2267

Final example of sample problem one, lithium sulfate reacting with silver nitrate.2275

Let's go ahead and do the double displacement.2282

Lithium is now going to hook up with nitrate.2285

Sulfate is going to now hook up with silver.2288

When we go ahead and do this, we are going to get lithium nitrate.2294

That is going to be aqueous when you use the table of solubility rules; and silver sulfate.2300

Silver sulfate is going to also be aqueous when you look up the table of solubility rules.2309

We have to balance this; this is going to be two of these and two of those.2318

I forgot, for the third example, the potassium nitrate reacting with copper chloride, I left it unbalanced.2325

Shame on me; again let's go ahead and fix that.2330

We are going to need two of these and two of those.2333

I hope this sample problem one was really testing your ability to not only predict the products2339

but also the physical states in addition of course to balancing.2345

The final set of sample problems involves the following.2351

I have seen a lot of tests.2355

A lot of the times, you have to translate from a sentence or a phrase into a chemical equation.2359

This is really testing your ability of nomenclature also.2365

So this is a nice cumulative type of problem.2368

Solid zinc reacts with aqueous lead(II) nitrate to form aqueous zinc(II) nitrate and solid lead.2373

Let's go ahead and translate it.2382

Zinc solid reacts with lead(II) nitrate going on to form zinc(II) nitrate aqueous and solid lead.2385

Let's go ahead and balance this; we are going to need two zinc nitrates.2407

No, it is going to be zinc(II) nitrate; I'm sorry; this is going to be two of these.2413

I am going to erase that coefficient; very sorry about that.2417

My balanced chemical equation is good to go.2423

This is an example of a single replacement reaction.2426

If you go ahead and look up the activity series, zinc will come in and knock out lead.2430

Second example, solid iron reacts with molecular chlorine gas to form solid iron(III) chloride.2437

Solid iron reacts with molecular chlorine; remember what we mean by molecular chlorine?2446

That is going to be Cl2 gas not just Cl.2453

That is going to go ahead and form solid iron(III) chloride, FeCl3 solid.2457

Let's go ahead and balance this.2465

Iron is good to go already; I have one on each side.2466

But I have two chlorines on the left and I have three chlorines on the right side.2469

We can go ahead and simply put a 3 in front of the Cl2 and a 2 in front of the FeCl3.2474

Finally I am going to need 2 irons.2482

This is again an example of an even-odd pair where you could have used fractional coefficients; that would have worked too.2485

This is an example of what we call a combination reaction where we have the form A plus B goes on to from AB.2493

Once again a combination reaction.2505

Last example, we have solid barium peroxide forming solid barium oxide and molecular oxygen gas.2508

Let's go ahead and write this and balance it; we have solid barium peroxide.2516

You are told that this forms solid barium oxide in combination with molecular oxygen gas.2529

Remember molecular oxygen; very important; that is going to be O2.2536

My bariums are good to go; the only thing we have to balance are the oxygens.2542

I have two oxygens on the left; I have three on the right.2548

Let's go ahead and put a 2 in front of BaO.2551

That is going to give me four oxygens on the right side now and two bariums.2557

All I have to do is put a 2 in front of there.2561

So I have two bariums on both sides of the equation.2564

I have four oxygens on both sides of the equation.2567

This is what we call a decomposition reaction where we have one compound AB essentially dissociating into A and B.2570

One compound forming multiple products.2585

This is a nice cumulative exercise where it really tests your ability,2589

your knowledge of the nomenclature, and of course to balance a chemical equation.2594

Thank you very much for your time.2600

I will see you next time on Educator.com.2601

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