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Franklin Ow

Franklin Ow

Atoms, Molecules, and Ions

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 (35)

0 answers

Post by K Lee on June 14, 2017

I thought  Halogens was group 17, not 7, and Noble Gases was group 18, not 8.
Also, I thought main group was 1,2, and 13-18, not 1,2,3-6,7,8, or 1-8

0 answers

Post by K Lee on June 14, 2017

At 16:41, did you meant group 7?

0 answers

Post by Evan Wang on May 8, 2017

Can you make a lesson on balancing chemical equations? I am studying for my exams and do not completely understand the concept.

0 answers

Post by Parsa Abadi on April 23, 2017

for ex1 shouldn't it be Li2F ?

0 answers

Post by David Sondergaard on October 15, 2016

Looks like lithium mass might be wrong on periodic table

0 answers

Post by Parth Shorey on April 18, 2016

How did you get 2.02 moles for chlorine ? and 2.02 moles for C?

3 answers

Last reply by: Professor Franklin Ow
Sat Apr 9, 2016 3:35 PM

Post by Parth Shorey on April 7, 2016

Is this Q&A active?

0 answers

Post by Peter Ke on September 7, 2015

At 47:10, can you show me the process of converting 1mg and 2.73 mg into 1.44 x 10^-4?

0 answers

Post by Peter Ke on September 7, 2015

Why ammonium has a net charge of 1+ while phosphate has a net charge of 3-?

Also is there like a table where it shows all the different net charges in an ionic compound?

0 answers

Post by Akilah Futch on July 29, 2015

Professor, are atom/ion charges and oxidation number the same?

0 answers

Post by Kate Danielle Rapinan on December 5, 2014

should i use 1.01 or 1.008 for the mass of hydrogen? does it make a lot of difference in calculations?

(i have 2 periodic tables, one has mass as 1.01, the other has 1.00798)

0 answers

Post by Kate Danielle Rapinan on December 5, 2014

are there rare times when a metal would have a negative charge than a positive charge? If so, how would you denote/write it as a roman numeral?

2 answers

Last reply by: John-Paul Kliebert
Thu Dec 4, 2014 1:22 PM

Post by John-Paul Kliebert on November 7, 2014

With the Empirical Formula I don't understand when and when not to round. I looked online and saw that it said "If the number is too far to round (x.1 ~ x.9), then multiply each solution by the same factor to get the lowest whole number multiple." What counts as too far or not? Could you please help me?

1 answer

Last reply by: Professor Franklin Ow
Mon Nov 3, 2014 10:57 PM

Post by Hannah Duncan on November 2, 2014

How do you determine the empirical formula if you're not given any masses? Such as this question:
Predict the Empirical Formulas for the Ionic Compounds formed from the following pairs of elements. Name each compound.
a.) Li and N
b.) Ga and O
c.) Rb and Cl
d.) Ba and S

1 answer

Last reply by: Professor Franklin Ow
Tue Oct 14, 2014 6:54 PM

Post by Kirk Graham on October 13, 2014

Isn't the naming of iron backwards, the 2+ would be like this:  Fe2+ and the Roman Numeral is used when the name of the compound is written out, like this:  iron (II).  That's what our textbook in Gen Chem shows.

1 answer

Last reply by: Professor Franklin Ow
Wed Sep 24, 2014 2:44 AM

Post by Ahmed alkarkhi on September 2, 2014

Hey prof do you recommend using Classical names of the transition metals over the IUPAC names?

2 answers

Last reply by: Christine Park
Mon Dec 21, 2015 8:17 PM

Post by William Kinne on August 5, 2014

Where can I find chemistry practice problems?

1 answer

Last reply by: Professor Franklin Ow
Tue Jun 24, 2014 1:32 PM

Post by brandon joyner on June 21, 2014

I thought kg was the base unit for SI like sample problem1?

1 answer

Last reply by: Professor Franklin Ow
Sat Feb 15, 2014 4:06 AM

Post by Aniket D on February 14, 2014

Rule 2 of naming ionic compounds; stated that we should not use prefixes such as; mono, di, tri etc.to denote subscripts in a formula. However doesn't the ionic compound of Carbon dioxide use the prefix "di" to represent the 2 oxygen atoms in the formula?

1 answer

Last reply by: Professor Franklin Ow
Thu Nov 7, 2013 5:08 PM

Post by Nada A. on November 2, 2013

should t Halogens and Nobel gasses be group 17 and 18... why are they referred to as group 7 and 8?

1 answer

Last reply by: Professor Franklin Ow
Thu Nov 7, 2013 5:10 PM

Post by chisom anyanwu on October 23, 2013

is the molar mass same as the atomic mass?

Related Articles:

Atoms, Molecules, and Ions

  • Rutherford’s gold foil experiment suggested the presence of the nucleus.
  • Protons and neutrons reside inside the nucleus, while electrons are outside.
  • Isotopes are atoms of the same element with different number of neutrons.
  • The mole allows for conversion to/from number of units of anything.
  • Dalton’s law of multiple proportions is demonstrated when solving for empirical and molecular formulas.

Atoms, Molecules, and Ions

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.

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

            Transcription: Atoms, Molecules, and Ions

            Hi, welcome back to Educator.com.0000

            Today's lesson in general chemistry is on atoms, molecules, and ions.0003

            We are going to go ahead and look at the lesson overview right now.0009

            We are going to get a brief introduction into atomic structure which is more of a historical background.0012

            Following this, we will get into what we mean by isotopes.0019

            We will get introduced to the periodic table in this lesson followed by naming compounds.0022

            Basically ionic compounds, those that are made from metals and nonmetals.0028

            Polyatomic ions, hydrates, and also what we mean by molecular compounds.0036

            We are then going to get associated with the concept of the mole0042

            and how that can be used to determine what is called a percentage composition0047

            which is going to ultimately lead to the finale of this lesson0053

            which is on determining what is known as the empirical and molecular formulas.0057

            By the beginning of the twentieth century, physicists had already known the following.0065

            Number one, that opposite charges attracted and that like charges repelled.0071

            Number two, a subatomic particle existed that was negatively charged which had been called the electron; commonly abbreviated e-.0075

            One of the first accepted models of the atom was called the plum pudding model.0086

            What the plum pudding model was the following.0090

            It was essentially a sphere of positive charge.0092

            Embedded in the sphere of positive charge were the subatomic particles, the electrons.0099

            These were completely stationary if you will.0107

            But there was no accounting for any other type of subatomic particle.0113

            Once again, this is called the plum pudding model.0119

            A sphere of positive charge that contained electrons scattered throughout.0122

            We now get into one of the earliest theories that was accepted to explain for atomic structure.0129

            This was called John Dalton's atomic theory.0138

            John Dalton's atomic theory contained the following premises.0141

            Number one was that matter was composed of individual units called atoms.0146

            Number two, that these atoms were the smallest unit possible and that they were indivisible.0171

            That is we could not break them down any further; atoms were indivisible.0177

            Number three, the third premise of John Dalton's atomic theory was that atoms of one element were all identical to each other.0188

            Number four, atoms of one element are going to be different than atoms of another element.0213

            That is what we reference as chemical identity.0218

            Atoms of one element different than atoms of another element.0224

            Finally the fifth and final premise was that when elements combine, they can combine in small whole number ratios to form compounds.0239

            Atoms combine in small whole number ratios to form compounds.0250

            This fifth premise here is going to be looked at again.0268

            That is what we call the law of multiple proportions.0273

            I am going to go ahead and put an asterisk next to0284

            each of the premises which we actually find not be so true today.0286

            Atoms were indivisible; that is not true because we can divide an atom in a process called nuclear fission.0292

            So premise two is not completely accurate.0298

            The next one is atoms of one element are all identical to each other.0302

            We are going to talk about this--is that not all atoms of the same element are identical.0307

            Because atoms can exist in different forms which we call isotopes.0312

            That is going to be coming on later on.0317

            Following John Dalton's atomic theory was one of the first major experiments on atomic structure.0323

            This is what we call Ernest Rutherford's gold foil experiment.0331

            What Ernest Rutherford did was the following.0334

            He shot a beam of high energy particles that were positively charged.0337

            These positively charged particles were what we call alpha rays.0343

            This was shot at a very thin piece of gold foil.0348

            What was the expected result?--the expected result was the following.0355

            That all of the alpha rays should have gone straight through.0360

            Why?--because the positive sphere was mostly empty space.0366

            And that the positive ray should have been attracted to the electrons that are embedded in the sphere.0370

            This is expected result due to plum pudding.0377

            However the actual result was the following.0387

            From the gold foil, yes the majority of the rays did indeed go straight through.0390

            However what was determined was that some rays actually got deflected at various angles.0397

            You would actually have a couple rays get deflected completely backward.0405

            This was highly unusual because this was not predicted by the plum pudding model.0411

            Instead what came about was the following implication; an implication was the following.0416

            That matter and atoms contained a dense center of positive charge.0424

            Because it is positive charge that was making the alpha rays deflect.0441

            This positive charge is what we today refer to as the nucleus.0449

            The nucleus is at the center of all atoms and is positively charged.0456

            To resummarize the implications of the gold foil experiment was that all atoms0464

            have a dense center that is positively charged which we call the nucleus.0469

            Inside the nucleus, two subatomic particles exist.0474

            The proton is positively charged; the neutron has zero charge.0477

            Number three, the nucleus is responsible for the mass of the atom; that is what gives an atom weight.0482

            Finally electrons, which were already known to have existed, exist outside the nucleus and have negligible weight.0489

            Let's go ahead and summarize these subatomic particles--neutron, proton, and electron.0499

            Please make a note that these relative masses and the relative charges are indeed in relation to each other.0510

            A neutron is going to be assigned a mass of 1 and a charge of 0.0520

            A proton is assigned a mass of 1 and a charge of +1.0524

            An electron has a negligible mass, a mass of 0, and a charge of -1.0529

            Of course, we symbolize the neutron lower-case n, proton p, and the electron is going to be e.0535

            One fallacy in John Dalton's atomic theory was that atoms of the same element are all identical to each other.0547

            But we are going to learn that atoms of the same element can exist in different forms; these are called isotopes.0555

            Isotopes are atoms of the same element; they differ only in their number of neutrons.0563

            Number of protons is going to be equal to the number of electrons; this is because of charged neutrality.0570

            But the only difference is the number of neutrons; number of n differs.0577

            We are going to introduce something called isotope format.0586

            For example, carbon has two isotopes, carbon-12 and carbon-13.0589

            Anytime you see the number immediately following the hyphen, this is what we call the mass number of the isotope.0597

            The mass number of the isotope is simply equal to the number of neutrons plus the number of protons.0607

            When we write an isotope in a specific format, it is always going to have the following design.0618

            Element symbol, X; A on top and Z on the bottom as a subscript.0628

            X is simply your element symbol; A is your mass number; Z is simply your atomic number.0634

            The atomic number is strictly equal to the number of protons of the element.0648

            A mathematical relation follows that if you take the difference of A minus Z, you are going to get the following.0659

            The number of neutrons plus the number of protons and then minus the number of protons; as you can see protons cancels.0668

            When you take the difference of A minus Z, you simply get the number of neutrons.0677

            For example, if you take 12, 6 and carbon.0683

            We have 12 is equal to the number of neutrons plus the number of protons.0689

            6 is on the bottom; that is the number of protons which is also equal to the number of electrons.0698

            Therefore A minus Z is equal to 6; that is equal to the number of neutrons.0704

            So carbon-12 has six protons, six neutrons, and six electrons.0712

            However if you do the same analysis for carbon-13, we are going to get six protons, seven neutrons, and six electrons.0721

            As you can see, isotopes strictly differ by the number of neutrons present inside the nucleus.0730

            Now we get introduced to the periodic table.0739

            The periodic table is essentially organized into different rows and columns.0742

            Each column is also referred to as a group; each row is referred to as a period.0749

            Typically there is always a staircase that is embedded in all of our periodic tables.0766

            Anything to the left side of the staircase is what we call a metal.0776

            Anything to the right side of the staircase is what we call a nonmetal.0785

            When we look at the periodic table, you see always the following.0797

            You see a number on top; the number of top is what we call the element number.0802

            The element number is actually the atomic number, the number of protons.0810

            The number on the bottom is the atomic mass of the element.0816

            Finally right in the middle is your element symbol.0823

            We are going to talk now more about different areas of the periodic table.0830

            We talked about metals; metals have the following characteristics.0839

            They appear lustrous which is shiny; they are good conductors of heat and of electricity.0843

            They are malleable which means we can form dinged sheets out of them.0850

            And they are ductile which means we can pull them into fine wires.0856

            Nonmetals are going to be the opposite.0866

            They are going to be electrically and thermally insulated which is what we mean by poor conductivity.0869

            In addition, they are brittle which means that if we take a hammer and hit them, they are going to crumble.0878

            Instead are not too malleable; so will crumble.0885

            But right in the middle of the border between metals and nonmetals, that is what we find semimetals.0891

            Semimetals are also known as metalloids.0900

            They border that staircase that we drew previously; border the staircase.0904

            Because they are called semimetals or metalloids, they are going to have properties of both.0915

            They tend to be good conductors of electricity but poor conductors of heat; they tend to be malleable.0923

            The typical examples are of course silicon and germanium.0928

            We see these two being used in a lot of electrical devices and what we call semiconductors.0938

            What we are going to analyze next are some of the major columns of the periodic table.0960

            Group 1 is what we call the alkali metals such as lithium and sodium and potassium, etc.0966

            Group 1 metals tend to have the following characteristics.0974

            That is they are moisture sensitive which means they have a tendency to react with water.0977

            They can be explosive when they react with water.0982

            Group 2 are what we call the alkaline earth metals such as calcium and magnesium and barium.0986

            These are harder than group 1 and less moisture sensitive.0995

            Group 6 are the halogens which are fluorine, chlorine, bromine, and iodine.1001

            All of these occur naturally in nature diatomically which means they occur in pairs.1007

            When someone says fluorine, it is really F2, Cl2, Br2, and I2; the halogens are incredibly reactive.1015

            Finally group number 8, the rightmost column of the periodic table are what we call the noble gases such as helium, argon, and neon.1028

            Unlike the group 7s, these are going to be monoatomic; that is they occur just as individual units.1039

            They are called noble gas species because they are relatively inert and have a smaller tendency to react.1045

            Now that we have been introduced to the periodic tables and the elements that comprise them, let's now talk about different types of compounds.1057

            The first type of compound we are going to study are called ionic compounds.1066

            Ionic compounds are formed between a metal and a nonmetal or they contain what are known as polyatomic ions.1069

            Because of the strong electrostatic attractions that exist between opposite charges, ionic compounds tend to have very high melting points.1078

            Below are a couple of these polyatomic ions; polyatomic could be the following.1087

            Something like CO32- which is called carbonate or SO42- which is called sulfate.1094

            NO31- which is called nitrate; you see why these are called polyatomic ions.1107

            Because it is several atoms coming together to form one entity; those are what we call polyatomic.1111

            That is different than monoatomic ions of course which are just formed from one unit.1120

            Something like Cl-, O2-, Al3+, etc.; polyatomic versus monoatomic.1124

            It turns out that we can actually use the periodic table to predict the ionic charge for main group elements.1135

            Main group elements are going to be group 1 and group 2, then group 3 to 6, and then 7 and 8.1141

            It turns out that all of the group 1's tend to form a 1+ charge.1160

            All of the group 2's tend to form a 2+ charge; group 3's tend to form a 3+ charge.1167

            We are going to skip 4 and 5 because they tend not to form ionic compounds to such a large degree.1174

            Column 6, we are going to get a 2- charge; column 7 is going to be a 1- charge.1183

            Group 8, we say no charge because group 8 are the noble gases and they tend not to react with other species.1188

            You noticed also very importantly that the positive charges, which are what are called cations, these are going to be formed by metals.1196

            The negative charges are what are called anions; these are all nonmetals.1208

            I am sorry; I am leaving out group 5.1219

            Group 5 is of course... this is going to be a 3- charge.1222

            So you have your cations and then your anions.1226

            Once again cations are going to be formed by metals and anions are going to formed by nonmetals.1229

            Now we have to learn how to name ionic compounds, such as NaCl, such as BaF2.1237

            We need to now go over the systematic procedure for doing so.1247

            You are going to name the cation first followed by the anion and the anion will end in ?ide.1253

            Something like NaCl; the cation is simply going to be sodium.1259

            Cl is usually chlorine; but now we are going to become chloride.1267

            So the correct name of NaCl is simply sodium chloride; very simple.1275

            You do not use prefixes mono, di, tri, etc, to denote subscripts in a formula.1284

            Something like for example BaF2; let's go ahead and name it.1289

            That is going to be barium followed by fluoride; you do not say difluoride.1294

            No, no difluoride, for example; again we ignore the prefixes.1302

            For certain transition metals, you are going to use a Roman numeral in parentheses to indicate the metal charge.1310

            If we want Fe2+, we are going to denote it as iron(II); Fe3+ cation is going to be Fe(III).1317

            Again you should definitely ask your instructor as to which transition metals he or she1331

            would like you to know for sure where you have to use a parentheses to indicate charge.1336

            Finally you can also have a hydrate; a hydrate is when you have water attached.1343

            For example, CoCl2·5H2O, we are going to use the prefixes followed by the word hydrate.1350

            Five using the Greek prefix would be penta; we would just say that this is the pentahydrate compound of CoCl2.1360

            Let's go ahead and put everything together and do a couple of examples.1372

            Al2O3, FeCl3, CuI2·3H2O; we are going to name the cation first.1378

            Al2O3, that is going to be aluminum followed by oxygen becoming oxide.1394

            FeCl3, Fe is one of those we need to know.1402

            Iron, something in parentheses, and then chlorine becomes chloride.1406

            Now we have to come up with the charge; each chlorine is a 1- charge.1411

            There is three of them, giving me a 3- overall.1416

            For iron, iron is going to be 3+ here, giving us iron(III) chloride in the end.1422

            Don't forget for the ionic compounds, the net charge should equal zero to maintain charge neutrality.1431

            Right here we are going to have copper.1442

            Copper is one of those where you need a Roman numeral for, followed by iodide.1444

            Then we have three waters which becomes trihydrate.1451

            Here iodine is a 1- charge; there is two of them giving me a -2 overall.1457

            Copper must be a +2 to balance it; I am going to indicate that in the parentheses as Roman numeral.1464

            Now we have to do the other way.1474

            Before I gave you the formula and you were asked to provide the name.1476

            Now let's go ahead and write the formula given the name.1480

            Let's go ahead and look at the following examples; barium phosphide; ammonium phosphate.1485

            Barium phosphide; barium is Ba; phosphide is phosphorus which is P.1511

            What I would like to do is I would like to put the charges up in the top just for my own purposes.1518

            Barium is 2+; phosphorus is 3-; what I always tell my students to do is then I use the crossing rule.1524

            Basically if you cross the charges, they become the subscripts for each of the other elements.1530

            Ba is going to get subscript of 3; phosphorus is going to get a subscript of 2.1537

            You can convince yourself to make sure that the charges cancel.1543

            Three of the bariums times 2+ plus two of the phosphorus times 3-, yes they do cancel to zero.1547

            Next one is ammonium phosphate; ammonium is NH4; phosphate is PO4.1556

            Ammonium is a 1+ charge; phosphate is a 3- charge; let's go ahead and cross them again.1564

            When we cross them, we are going to get (NH4)3PO4 with a 1.1572

            Let's go ahead and make sure that they all work out.1582

            Three of the ammoniums times a +1 charge plus one of the phosphates and a 3- charge; yes it does cancel to zero.1585

            That is some nice sample problems on naming ionic compounds.1596

            The other type of compound we are going to learn about now is called a molecular compound.1603

            Molecular compounds are not a metal and nonmetal but instead they are composed from two nonmetals.1608

            Something like CO2; something like CO; and of course H2O itself.1613

            To name a molecular compound, the rules are going to be very similar with one main exception.1620

            We are going to name the first element first by its entire name; then the second element has a ?ide suffix.1627

            This time we do pay attention to the subscripts.1636

            We are going to use the Greek prefixes, mono, di, tri, tetra, etc, to determine what the subscript is.1639

            However you omit mono for the first nonmetal.1647

            Something like CO2, name the first element first which is carbon.1651

            The second element is oxygen becoming oxide.1656

            There is two of them which is how we come up with carbon dioxide.1660

            Something like N2O5; there is two nitrogens which is di.1665

            The first element gets the full name; so dinitrogen; then five oxides or just pentoxide.1670

            So actually naming a molecular compound is always easier than naming an ionic compound.1679

            Why?--because you don't have to worry about charge.1685

            Now onto what we call the concept of the mole.1691

            When you go to the supermarket and you purchase eggs, they are always in either a half dozen or a dozen.1697

            It is very hard and very seldom that you are going to purchase just an individual egg.1704

            It is convenience; it is convenience for us to refer to eggs as a dozen.1708

            For chemistry, we have the equivalent of our dozen.1716

            The mole is our chemists' dozen for convenience.1721

            When you go into lab, you don't measure just the weight of for example one atom.1730

            You don't deal with an individual atom; it is not practical.1735

            What the mole is all about is that it goes from the microscale to a macroscale; it is more practical.1739

            Basically our dozen which is what we call the mole... the mole is abbreviated simply m-o-l.1755

            It is equal to the following: 6.022 times 1023 units of whatever you are measuring.1763

            This is equivalent of saying one dozen is equal to twelve units of whatever you are measuring.1774

            Again our dozen, the chemists' dozen is what we call the mole.1782

            This mole is probably... this number you have to memorize I am pretty sure.1786

            6.022 times 1023 is what we call Avogadro's number; Avogadro's number.1792

            The mole we are going to find is a central unit; it can connect the following quantities.1803

            The mole, we can equal and calculate individual units; the mole we can also get grams.1809

            When we go from mole to units, we are going to multiply this way.1823

            By Avogadro's number, 6.022 times 1023 units for every one mole.1830

            But when we go the opposite direction, we are going to divide.1838

            We are going to divide by Avogadro's number.1842

            Mole can also be used to go to mass.1854

            When we go to mass, we can actually multiply the mole by a rate conversion factor of gram for every one mole.1858

            When we go the opposite direction, we are going to divide by one mole.1871

            We are going to divide by grams; so we can see that mole is truly a central unit.1879

            The conversion factor of gram per mole, that is what we call the molar mass.1892

            This is equal to the grams that one mole of a substance weighs.1903

            We are going to come back to this quite heavily.1920

            We are going to do a lot of example problems utilizing the mole.1924

            Where do we get this molar mass from then?--on the periodic table.1929

            On the periodic table, atomic masses... remember it is the number below the element symbol... are also called molar masses.1935

            When you look at C-6, 12.101 on the periodic table, this bottom number here it the molar mass of carbon.1943

            One mole of carbon weighs exactly 12.01 grams; that is how you write it out.1953

            That is going to be a conversion factor.1960

            How many moles are in 25.7 grams of sodium?1962

            You say 25.7 grams of sodium times something over something.1967

            That is going to give us moles of sodium.1973

            Remembering our dimensional analysis, now grams of sodium is going to go downstairs.1978

            Moles of sodium is going to go upstairs.1984

            You just look up the value; that is going to be 22.99 grams of sodium for every one mole of sodium.1988

            Giving you your answer in units of moles of sodium.1994

            Let's now go ahead and utilize Avogadro's number.1998

            How many atoms are in 1.2 moles of carbon?2002

            1.2 moles of carbon times something over something is going to give us our answer in units of atoms of carbon.2008

            To get cancelled, moles of carbons goes downstairs; then atoms of carbon goes upstairs.2021

            We know from our flow chart with the mole that this is involving Avogadro's number.2029

            One mole will go downstairs and 6.022 times 1023 goes upstairs to give our answer in units of atoms of carbon.2034

            Whenever you see the word atoms or molecules, that is pretty much a good give away2046

            that you are going to be using Avogadro's number as your conversion factor.2053

            Again always keep in mind that; especially if the word moles is mentioned too.2060

            Now that we have been introduced to what the mole can help us get,2068

            we are going to find something else that the mole can be used to calculate,2073

            what is called molar mass of not only an atom but also of a compound.2079

            Let me say that again; the mole can be used to calculate the molar mass of a compound.2084

            The molar mass of an element is already provided to us on the periodic table.2088

            Basically the molar mass is simply the net sum of all molar masses of each element.2093

            The molar mass of carbon dioxide is simply equal to2111

            the molar mass of carbon plus the molar mass of oxygen2114

            plus the molar mass of oxygen because we have two oxygens.2122

            That is going to be equal to 12.01 grams plus 16.00 grams plus 16.00 grams.2126

            Giving us a grand total of 44.01 grams; one mole of CO2 weighs exactly 44.01 grams.2135

            Having said that, we can also use molar mass as a conversion factor for a compound.2149

            How many grams are in 1.2 moles of carbon dioxide?2154

            1.2 moles of CO2 times something over something is going to give us our answer in units of grams of CO2.2158

            The mole is going to go downstairs; the grams is going to go upstairs.2169

            Right there in the previous problem, we calculated this conversion factor or molar mass of CO2.2175

            Anytime you want to relate mass and moles, it is always molar mass to be your conversion factor.2183

            It is going to be 1 on the bottom; it is going to be 44.01 on top.2190

            Giving us our answer in units of grams of CO2.2195

            The next application of the mole is what we call percentage composition.2206

            Percentage composition is given to us to be the following.2211

            Percent composition of an element tells us basically the parts of the element divided by total parts times 100.2216

            It basically tells us your relative amount that a specific element makes up of the compound.2236

            This equation that we actually use is going to then be the total mass of the element2244

            divided by the molar mass of the compound; all of that times 100.2256

            Let's go ahead and answer this question.2269

            How many grams of carbon dioxide are contained in 65.1 grams of CO2?2270

            This problem is strictly wanting to know how much carbon is going to be contained in so much CO2.2276

            Let's go ahead and first calculate the percentage composition of carbon in CO2.2283

            We are going to say 12.01 grams because there is only one carbon.2294

            Divided by the entire molar mass of CO2 which is 44.01 grams; times 100.2299

            That is going to give us 27.3 percent.2306

            Carbon dioxide is only 27.3 percent carbon; the majority is oxygen.2309

            That tells me that any sample I hold of CO2 in my hand, 27.3 percent of that is carbon.2317

            We are going to use percentage as our conversion factor.2324

            We are going to say .273 times the 65.1 grams of CO2 is going to give me my grams of carbon.2328

            We are going to get 17.8 grams as our answer.2339

            Once again anytime you see a problem that relates the amount of an element2345

            contained in a certain amount of compound, you want to use percentage composition as your tool.2351

            Finally we are now going to see how we can determine what is known as an empirical and molecular formula.2360

            An empirical formula is the simplest whole number ratio of a formula.2367

            For example, CO2, this is an exact 1:2 ratio of carbon to oxygen.2371

            That is my smallest ratio; I cannot go smaller than that.2378

            However let's go ahead and look at one that is maybe C2H6.2382

            In this example, I have a ratio of 2:6 of carbon to hydrogen.2388

            But you notice that I can divide; I can divide by a lowest common multiple which is 2.2395

            I am going to get 1:3 ratio instead; therefore 1:3 is actually my smallest ratio.2402

            That is going to be CH3 which is my empirical formula.2410

            What we are going to do is we are going to determine how to calculate the empirical formula.2415

            The empirical formula is determined from elemental analysis.2422

            Basically you have an unknown sample and you put it through...2428

            You perform what is called an elemental analysis test on it.2434

            After you perform this analysis, you are going to get a printout2440

            of the percentage that a certain element makes up in your unknown.2445

            You are going to get x percent of for example carbon.2450

            You are going to get y percent of for example oxygen.2454

            You are going to get z percent for example of hydrogen.2457

            These percentages again are only to be used for empirical formulas.2461

            They may or may not necessarily give you the actual formula; only the relative amounts are reported.2466

            If you want to get the actual formula, which is known as the molecular formula, you can determine it.2473

            But only if the molar mass is already provided for you.2478

            Let's go ahead and look at a typical example.2486

            A compound used as an additive for gasoline to help prevent engine knock has the following percent composition by mass.2490

            71.65 percent chlorine; 24.27 percent carbon; 4.07 percent hydrogen.2497

            The molar mass is known to be 98.96 grams per mole.2505

            Determine both the empirical and molecular formulas.2509

            The very first step is to get your percentages into grams; get percentage to grams.2516

            After we get everything into grams, then we can go to our central unit from the chapter which is moles.2527

            Again whenever you are in doubt, get to moles.2533

            The nice thing about percentage is that percentage is always made out of 100.2536

            If we assume 100 gram of compound, the percentages automatically become grams.2541

            We have 71.65 grams of chlorine, 24.27 grams of carbon, and 4.07 of hydrogen.2548

            Again assuming 100 gram of compound.2560

            Now we are going to go ahead and get to moles.2572

            This is something we know how to do by now.2575

            When we go ahead and get to moles, we are going to get 2.02 moles of chlorine.2577

            We are going to get 2.02 moles of carbon.2584

            We are going to get 4.04 moles of hydrogen.2587

            Step one is done; now step two.2596

            Step two, we are going to take our moles and we are going to divide.2601

            We are going to divide by the smallest moles present; divide by smallest moles present; everything.2604

            We are going to get a nice whole number that is going to become our subscripts in our empirical formula.2617

            The result is the empirical formula subscript.2626

            When we take chlorine and divide it by the smallest number which is 2.02, we are going to get 1.2642

            Carbon divided by 2.02, get 1.2648

            For hydrogen, 4.04 divided by 2.02, we are going to get 2.2651

            Giving us our empirical formula of CH2Cl.2655

            What happens if you don't get a perfect whole number?--what if you get 1.5, 1.5, 2?2665

            Dalton's law of multiple proportions tells us that we cannot combine elements in other than a small whole number ratio.2684

            Basically any time you have a decimal number, we are just going to then2694

            multiply everything by the same factor to get everything in whole numbers.2697

            When we do that, we can multiply 1.5 by 2 and everything by 2.2702

            Giving us nothing but whole numbers; 3, 3, and 4.2707

            If that happened, then those numbers would become the empirical; we would get C3H4Cl3.2711

            To get the molecular formula, you simply do the following.2722

            You take the molar mass of the molecular formula which is given to you.2726

            You divide it by the molar mass of your empirical formula.2733

            That is always going to give you a nice whole number.2738

            When we do this, we are going to get 98.96 divided by our molar mass of our empirical formula which is CH2Cl.2741

            We actually get 48.468; when we do this, we get approximately 2.2752

            You take that whole number and you multiply the subscripts of the empirical formula by this.2758

            I get a molecular formula of C2H4Cl2 which is the formula of the actual compound.2764

            You can always double check your work.2773

            You can add up the molar mass of C2H4Cl2.2774

            We are going to get very close to 98.96.2778

            Let's go ahead and summarize.2783

            Atoms are composed of a central nucleus that is positively charged.2786

            Protons and neutrons reside within the nucleus; electrons are outside.2790

            We saw that isotopes are the same element; differ only by their number of neutrons.2796

            We have went through some specific rules for naming ionic and molecular compounds.2801

            We also saw that the mole is a central unit that allows for conversion between number of atoms and molecules and for mass.2806

            Finally the mole is a central unit that is required for several types of different problems2814

            Including percentage composition and the empirical and molecular formula problems.2822

            Let's go ahead and look at some sample problems.2831

            You have 1 milligram of lithium metal reacting with molecular fluorine gas.2834

            The resulting fluoride salt has a mass of 7.3 mg.2838

            Determine the empirical formula of lithium fluoride; you have 1.00 mg of lithium.2842

            After it reacts with the molecular fluorine gas which is F2, you get 3.73 mg of your lithium fluoride.2850

            If I start with only 1 mg of lithium and I wind up with 3.73 mg of compound, why did my mass get heavier?2865

            The mass got heavier because it reacted with fluorine; fluorine came on board.2874

            We can actually determine the mass of fluorine.2879

            It is just the difference, 3.73 mg minus 1.00 mg.2882

            That is going to give us 2.73 mg which is the mass of fluorine reacted.2888

            We got all of our masses now; now let's get everything into moles.2902

            2.73 mg of fluorine is going to become 1.44 times 10-4 moles of fluorine.2908

            1.00 mg of lithium is going to become 1.44 times 10-4 moles of lithium.2919

            Now that we have all of our moles, we can divide by the smallest number present.2930

            For this problem, they are identical.2934

            When we divide everything, we are going to get Li1 and F1 for our empirical formula; or just LiF.2937

            This answer actually makes sense because we know that lithium is a 1+ charge and that fluorine is a 1- charge.2947

            Indeed the charges do balance each other out.2954

            Now moving onto sample problem two.2962

            How many atoms of carbon are present in 2.67 kg of C6H6?2964

            In this case, we are asked about an element within a certain compound.2970

            That sounds a lot like percentage composition.2976

            Let's go ahead and calculate the percentage composition of carbon in C6H6.2978

            The percentage composition is equal to the total mass of carbon...2984

            There is six of them by the molar mass 12.01.2988

            Divided by the total mass of C6H6 which is 6 by 12.01 plus the 6 hydrogens by 1.008.2991

            All of that times 100; when we do this, we get 92.3 percent.3001

            So C6H6 is mostly carbon; 92 percent carbon.3007

            Remember what we did last time.3013

            We are going to take our percentage composition 0.923.3015

            We are going to multiply it by the total mass 2.67 kg.3019

            That is going to give us the 2.46 kg of carbon in this specific sample.3023

            The question is asking for atoms of carbon though.3031

            Somehow we have to go from kilograms to atoms.3035

            Remember anytime you see the word atoms or molecules, it is going to be via an Avogadro number.3040

            Our first step is to go from kg just to regular g.3047

            Then from g to the central unit which is moles.3052

            Then of course moles onto atoms; we can go ahead and do this.3056

            We are going to say 2.46 kg of carbon times something over something.3063

            G goes on top; kg on the bottom; that is going to be 103 kg over 1 kg.3071

            Now onto moles; times something over something; g goes downstairs to get cancelled.3078

            Moles goes upstairs to get carried through the final answer; that is molar mass.3084

            You look up the molar mass of carbon which is 12.01 grams for every one mole.3090

            Then finally times something over something, giving us our answer in atoms of carbon.3095

            That is Avogadro's number; mole on the bottom; Avogadro's on top which is atoms here.3102

            That is 1 mole on the bottom and 6.022 times 1023 atoms on top.3110

            You should get an answer of 1.23 times 1026 atoms of carbon.3116

            I want to thank you for your attention.3128

            This concludes our lecture on atoms, molecules, and ions.3130

            Thank you for using Educator.com.3135

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