Join Dr. Franklin Ow’s General Chemistry online course where he explains difficult concepts in easy-to-understand language and follows with plenty of step-by-step examples you will likely see on homework and exams.

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I. Basic Concepts & Measurement of Chemistry

  Basic Concepts of Chemistry 16:26
   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 29:22
   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 52:18
   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 43:24
   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 55:40
   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 42:10
   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 42:38
   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 55:28
   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 42:33
   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 38:50
   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 52:39
   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 1:11:41
   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 62:16 
   Sample Problem 1: Determine the Atomic Hybridization 67:20 
    XeO₃ 67:21 
    SF₆ 67:49 
    I₃⁻ 68:20 
   Sample Problem 2 69:04 

IX. Gases, Solids, & Liquids

  Gases 35:06
   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 33:47
   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 25:13
   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 38:06
   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 37:45
   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 34:09
   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 43:44
   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 55:26
   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 36:13
   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 41:16
   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 39:03
   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 16:39
   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 

Duration: 16 hours, 21 minutes

Number of Lessons: 24

This course is ideal for both high school and college students taking their first Chemistry course. Balancing an entire sheet of chemistry equations, memorizing the Periodic Table, and answering tough vocabulary questions is much easier when you have a strong foundation in General Chemistry.

Additional Features:

  • Free Sample Lessons
  • Closed Captioning (CC)
  • Downloadable Lecture Slides
  • Study Guides
  • Instructor Comments

Topics Include:

  • Atoms, Molecules, & Ions
  • Chemical Reactions
  • Stoichiometry
  • Thermodynamics
  • Structure of Atoms
  • Periodic Trends
  • Gases, Solids, & Liquids
  • Chemical Equilibrium
  • Acids & Bases
  • Electrochemistry
  • Transition Elements
  • Nuclear Chemistry

Dr. Franklin Ow obtained both his B.S. and Ph.D. in chemistry from the University of California, Los Angeles. He currently teaches chemistry classes at three universities and has won multiple awards in his academic field.

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