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Table of Contents

I. Basic Concepts & Measurement of Chemistry

Basic Concepts of Chemistry

16m 26s

- Intro0:00
- Lesson Overview0:07
- Introduction0:56
- What is Chemistry?0:57
- What is Matter?1:16
- Solids1:43
- General Characteristics1:44
- Particulate-level Drawing of Solids2:34
- Liquids3:39
- General Characteristics of Liquids3:40
- Particulate-level Drawing of Liquids3:55
- Gases4:23
- General Characteristics of Gases4:24
- Particulate-level Drawing Gases5:05
- Classification of Matter5:27
- Classification of Matter5:26
- Pure Substances5:54
- Pure Substances5:55
- Mixtures7:06
- Definition of Mixtures7:07
- Homogeneous Mixtures7:11
- Heterogeneous Mixtures7:52
- Physical and Chemical Changes/Properties8:18
- Physical Changes Retain Chemical Composition8:19
- Chemical Changes Alter Chemical Composition9:32
- Physical and Chemical Changes/Properties, cont'd10:55
- Physical Properties10:56
- Chemical Properties11:42
- Sample Problem 1: Chemical & Physical Change12:22
- Sample Problem 2: Element, Compound, or Mixture?13:52
- Sample Problem 3: Classify Each of the Following Properties as chemical or Physical15:03

Tools in Quantitative Chemistry

29m 22s

- Intro0:00
- Lesson Overview0:07
- Units of Measurement1:23
- The International System of Units (SI): Mass, Length, and Volume1:39
- Percent Error2:17
- Percent Error2:18
- Example: Calculate the Percent Error2:56
- Standard Deviation3:48
- Standard Deviation Formula3:49
- Standard Deviation cont'd4:42
- Example: Calculate Your Standard Deviation4:43
- Precisions vs. Accuracy6:25
- Precision6:26
- Accuracy7:01
- Significant Figures and Uncertainty7:50
- Consider the Following (2) Rulers7:51
- Consider the Following Graduated Cylinder11:30
- Identifying Significant Figures12:43
- The Rules of Sig Figs Overview12:44
- The Rules for Sig Figs: All Nonzero Digits Are Significant13:21
- The Rules for Sig Figs: A Zero is Significant When It is In-Between Nonzero Digits13:28
- The Rules for Sig Figs: A Zero is Significant When at the End of a Decimal Number14:02
- The Rules for Sig Figs: A Zero is not significant When Starting a Decimal Number14:27
- Using Sig Figs in Calculations15:03
- Using Sig Figs for Multiplication and Division15:04
- Using Sig Figs for Addition and Subtraction15:48
- Using Sig Figs for Mixed Operations16:11
- Dimensional Analysis16:20
- Dimensional Analysis Overview16:21
- General Format for Dimensional Analysis16:39
- Example: How Many Miles are in 17 Laps?17:17
- Example: How Many Grams are in 1.22 Pounds?18:40
- Dimensional Analysis cont'd19:43
- Example: How Much is Spent on Diapers in One Week?19:44
- Dimensional Analysis cont'd21:03
- SI Prefixes21:04
- Dimensional Analysis cont'd22:03
- 500 mg → ? kg22:04
- 34.1 cm → ? um24:03
- Summary25:11
- Sample Problem 1: Dimensional Analysis26:09

II. Atoms, Molecules, and Ions

Atoms, Molecules, and Ions

52m 18s

- Intro0:00
- Lesson Overview0:08
- Introduction to Atomic Structure1:03
- Introduction to Atomic Structure1:04
- Plum Pudding Model1:26
- Introduction to Atomic Structure Cont'd2:07
- John Dalton's Atomic Theory: Number 12:22
- John Dalton's Atomic Theory: Number 22:50
- John Dalton's Atomic Theory: Number 33:07
- John Dalton's Atomic Theory: Number 43:30
- John Dalton's Atomic Theory: Number 53:58
- Introduction to Atomic Structure Cont'd5:21
- Ernest Rutherford's Gold Foil Experiment5:22
- Introduction to Atomic Structure Cont'd7:42
- Implications of the Gold Foil Experiment7:43
- Relative Masses and Charges8:18
- Isotopes9:02
- Isotopes9:03
- Introduction to The Periodic Table12:17
- The Periodic Table of the Elements12:18
- Periodic Table, cont'd13:56
- Metals13:57
- Nonmetals14:25
- Semimetals14:51
- Periodic Table, cont'd15:57
- Group I: The Alkali Metals15:58
- Group II: The Alkali Earth Metals16:25
- Group VII: The Halogens16:40
- Group VIII: The Noble Gases17:08
- Ionic Compounds: Formulas, Names, Props.17:35
- Common Polyatomic Ions17:36
- Predicting Ionic Charge for Main Group Elements18:52
- Ionic Compounds: Formulas, Names, Props.20:36
- Naming Ionic Compounds: Rule 120:51
- Naming Ionic Compounds: Rule 221:22
- Naming Ionic Compounds: Rule 321:50
- Naming Ionic Compounds: Rule 422:22
- Ionic Compounds: Formulas, Names, Props.22:50
- Naming Ionic Compounds Example: Al₂O₃22:51
- Naming Ionic Compounds Example: FeCl₃23:21
- Naming Ionic Compounds Example: CuI₂ 3H₂O24:00
- Naming Ionic Compounds Example: Barium Phosphide24:40
- Naming Ionic Compounds Example: Ammonium Phosphate25:55
- Molecular Compounds: Formulas and Names26:42
- Molecular Compounds: Formulas and Names26:43
- The Mole28:10
- The Mole is 'A Chemist's Dozen'28:11
- It is a Central Unit, Connecting the Following Quantities30:01
- The Mole, cont'd32:07
- Atomic Masses32:08
- 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'd34:25
- 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 Composition36:43
- Example: How Many Grams of Carbon Contained in 65.1 Grams of Carbon Dioxide?36:44
- Empirical and Molecular Formulas39:19
- Empirical Formulas39:20
- Empirical Formula & Elemental Analysis40:21
- Empirical and Molecular Formulas, cont'd41:24
- Example: Determine Both the Empirical and Molecular Formulas - Step 141:25
- Example: Determine Both the Empirical and Molecular Formulas - Step 243:18
- Summary46:22
- Sample Problem 1: Determine the Empirical Formula of Lithium Fluoride47:10
- 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

- Intro0:00
- Lesson Overview0:06
- The Law of Conservation of Mass and Balancing Chemical Reactions1:49
- The Law of Conservation of Mass1:50
- Balancing Chemical Reactions2:50
- Balancing Chemical Reactions Cont'd3:40
- Balance: N₂ + H₂ → NH₃3:41
- Balance: CH₄ + O₂ → CO₂ + H₂O7:20
- Balancing Chemical Reactions Cont'd9:49
- Balance: C₂H₆ + O₂ → CO₂ + H₂O9:50
- Intro to Chemical Equilibrium15:32
- When an Ionic Compound Full Dissociates15:33
- When an Ionic Compound Incompletely Dissociates16:14
- Dynamic Equilibrium17:12
- Electrolytes and Nonelectrolytes18:03
- Electrolytes18:04
- Strong Electrolytes and Weak Electrolytes18:55
- Nonelectrolytes19:23
- Predicting the Product(s) of an Aqueous Reaction20:02
- Single-replacement20:03
- Example: Li (s) + CuCl₂ (aq) → 2 LiCl (aq) + Cu (s)21:03
- Example: Cu (s) + LiCl (aq) → NR21:23
- Example: Zn (s) + 2HCl (aq) → ZnCl₂ (aq) + H₂ (g)22:32
- Predicting the Product(s) of an Aqueous Reaction23:37
- Double-replacement23:38
- Net-ionic Equation25:29
- Predicting the Product(s) of an Aqueous Reaction26:12
- Solubility Rules for Ionic Compounds26:13
- Predicting the Product(s) of an Aqueous Reaction28:10
- Neutralization Reactions28:11
- Example: HCl (aq) + NaOH (aq) → ?28:37
- Example: H₂SO₄ (aq) + KOH (aq) → ?29:25
- Predicting the Product(s) of an Aqueous Reaction30:20
- Certain Aqueous Reactions can Produce Unstable Compounds30:21
- Example 130:52
- Example 232:16
- Example 332:54
- Summary33:54
- Sample Problem 134:55
- 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 239:09
- Question 139:10
- Question 240:36
- Question 341:47

Chemical Reactions II

55m 40s

- Intro0:00
- Lesson Overview0:10
- Arrhenius Definition1:15
- Arrhenius Acids1:16
- Arrhenius Bases3:20
- The Bronsted-Lowry Definition4:48
- Acids Dissolve In Water and Donate a Proton to Water: Example 14:49
- Acids Dissolve In Water and Donate a Proton to Water: Example 26:54
- Monoprotic Acids & Polyprotic Acids7:58
- Strong Acids11:30
- Bases Dissolve In Water and Accept a Proton From Water12:41
- Strong Bases16:36
- The Autoionization of Water17:42
- Amphiprotic17:43
- Water Reacts With Itself18:24
- Oxides of Metals and Nonmetals20:08
- Oxides of Metals and Nonmetals Overview20:09
- Oxides of Nonmetals: Acidic Oxides21:23
- Oxides of Metals: Basic Oxides24:08
- Oxidation-Reduction (Redox) Reactions25:34
- Redox Reaction Overview25:35
- Oxidizing and Reducing Agents27:02
- Redox Reaction: Transfer of Electrons27:54
- Oxidation-Reduction Reactions Cont'd29:55
- Oxidation Number Overview29:56
- Oxidation Number of Homonuclear Species31:17
- Oxidation Number of Monatomic Ions32:58
- Oxidation Number of Fluorine33:27
- Oxidation Number of Oxygen34:00
- Oxidation Number of Chlorine, Bromine, and Iodine35:07
- Oxidation Number of Hydrogen35:30
- Net Sum of All Oxidation Numbers In a Compound36:21
- Oxidation-Reduction Reactions Cont'd38:19
- Let's Practice Assigning Oxidation Number38:20
- Now Let's Apply This to a Chemical Reaction41:07
- Summary44:19
- Sample Problems45:29
- Sample Problem 145:30
- Sample Problem 2: Determine the Oxidizing and Reducing Agents48:48
- Sample Problem 3: Determine the Oxidizing and Reducing Agents50:43

IV. Stoichiometry

Stoichiometry I

42m 10s

- Intro0:00
- Lesson Overview0:23
- Mole to Mole Ratios1:32
- 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'd5:13
- Balanced Chemical Reaction5:14
- Mole to Mole Ratios Cont'd7:25
- 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 Conversion11:06
- Mass to mass Conversion11:07
- 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 Yields20:42
- Limiting Reactants, Percent Yields20:43
- Example 8: How Many Grams of Ammonia Can Form If You Have 3.1 Grams of H₂ and 3.1 Grams of N₂22:25
- Percent Yield25:30
- Example 9: How Many Grams of The Excess Reactant Remains?26:37
- Summary29:34
- 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 136:10
- Sample Problem 3: Part 2 - What Amount Of The Excess Reactant Will Remain?40:53

Stoichiometry II

42m 38s

- Intro0:00
- Lesson Overview0:10
- Molarity1:14
- Solute and Solvent1:15
- Molarity2:01
- Molarity Cont'd2:59
- 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
- Dilutions10:01
- Dilution: M₁V₂=M₁V₂10:02
- Example 5: Explain How to Make 250 mL of a 0.67 M KBr Solution Starting From a 1.2M Stock Solution12:04
- Stoichiometry and Double-Displacement Precipitation Reactions14:41
- Example 6: How Many grams of PbCl₂ Can Form From 250 mL of 0.32 M NaCl?15:38
- Stoichiometry and Double-Displacement Precipitation Reactions18:05
- 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 Reactions21:01
- Example 8: How Many Grams of NaOh are Required to Neutralize 4.5 Grams of HCl?21:02
- Stoichiometry and Neutralization Reactions23:03
- 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 Standardization25:28
- Introduction to Titration & Standardization25:30
- Acid-Base Titration26:12
- The Analyte & Titrant26:24
- The Experimental Setup26:49
- The Experimental Setup26:50
- Stoichiometry and Acid-Base Standardization28:38
- Example 9: Determine the Concentration of the Analyte28:39
- Summary32:46
- Sample Problem 1: Stoichiometry & Neutralization35:24
- Sample Problem 2: Stoichiometry37:50

V. Thermochemistry

Energy & Chemical Reactions

55m 28s

- Intro0:00
- Lesson Overview0:14
- Introduction1:22
- Recall: Chemistry1:23
- Energy Can Be Expressed In Different Units1:57
- The First Law of Thermodynamics2:43
- Internal Energy2:44
- The First Law of Thermodynamics Cont'd6:14
- Ways to Transfer Internal Energy6:15
- Work Energy8:13
- Heat Energy8:34
- ∆U = q + w8:44
- Calculating ∆U, Q, and W8:58
- Changes In Both Volume and Temperature of a System8:59
- Calculating ∆U, Q, and W Cont'd11:01
- The Work Equation11:02
- Example 1: Calculate ∆U For The Burning Fuel11:45
- Calculating ∆U, Q, and W Cont'd14:09
- The Heat Equation14:10
- Calculating ∆U, Q, and W Cont'd16:03
- Example 2: Calculate The Final Temperature16:04
- Constant-Volume Calorimetry18:05
- Bomb Calorimeter18:06
- The Effect of Constant Volume On The Equation For Internal Energy22:11
- Example 3: Calculate ∆U23:12
- Constant-Pressure Conditions26:05
- Constant-Pressure Conditions26:06
- Calculating Enthalpy: Phase Changes27:29
- Melting, Vaporization, and Sublimation27:30
- Freezing, Condensation and Deposition28:25
- Enthalpy Values For Phase Changes28:40
- 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 Reaction31:22
- Example 5: Calculate The Heat In kJ Associated With The Complete Reaction of 155 g NH₃31:23
- Using Standard Enthalpies of Formation33:53
- Standard Enthalpies of Formation33:54
- Using Standard Enthalpies of Formation36:12
- Example 6: Calculate The Standard Enthalpies of Formation For The Following Reaction36:13
- Enthalpy From a Series of Reactions39:58
- Hess's Law39:59
- Coffee-Cup Calorimetry42:43
- Coffee-Cup Calorimetry42:44
- Example 7: Calculate ∆H° of Reaction45:10
- Summary47:12
- Sample Problem 148:58
- Sample Problem 251:24

VI. Quantum Theory of Atoms

Structure of Atoms

42m 33s

- Intro0:00
- Lesson Overview0:07
- Introduction1:01
- Rutherford's Gold Foil Experiment1:02
- Electromagnetic Radiation2:31
- Radiation2:32
- Three Parameters: Energy, Frequency, and Wavelength2:52
- Electromagnetic Radiation5:18
- The Electromagnetic Spectrum5:19
- Atomic Spectroscopy and The Bohr Model7:46
- Wavelengths of Light7:47
- Atomic Spectroscopy Cont'd9:45
- The Bohr Model9:46
- Atomic Spectroscopy Cont'd12:21
- The Balmer Series12:22
- Rydberg Equation For Predicting The Wavelengths of Light13:04
- The Wave Nature of Matter15:11
- The Wave Nature of Matter15:12
- The Wave Nature of Matter19:10
- New School of Thought19:11
- Einstein: Energy19:49
- Hertz and Planck: Photoelectric Effect20:16
- de Broglie: Wavelength of a Moving Particle21:14
- Quantum Mechanics and The Atom22:15
- Heisenberg: Uncertainty Principle22:16
- Schrodinger: Wavefunctions23:08
- Quantum Mechanics and The Atom24:02
- Principle Quantum Number24:03
- Angular Momentum Quantum Number25:06
- Magnetic Quantum Number26:27
- Spin Quantum Number28:42
- The Shapes of Atomic Orbitals29:15
- Radial Wave Function29:16
- Probability Distribution Function32:08
- The Shapes of Atomic Orbitals34:02
- 3-Dimensional Space of Wavefunctions34:03
- Summary35:57
- Sample Problem 137:07
- Sample Problem 240:23

VII. Electron Configurations and Periodicity

Periodic Trends

38m 50s

- Intro0:00
- Lesson Overview0:09
- Introduction0:36
- Electron Configuration of Atoms1:33
- Electron Configuration & Atom's Electrons1:34
- Electron Configuration Format1:56
- Electron Configuration of Atoms Cont'd3:01
- Aufbau Principle3:02
- Electron Configuration of Atoms Cont'd6:53
- Electron Configuration Format 1: Li, O, and Cl6:56
- Electron Configuration Format 2: Li, O, and Cl9:11
- Electron Configuration of Atoms Cont'd12:48
- Orbital Box Diagrams12:49
- Pauli Exclusion Principle13:11
- Hund's Rule13:36
- Electron Configuration of Atoms Cont'd17:35
- Exceptions to The Aufbau Principle: Cr17:36
- Exceptions to The Aufbau Principle: Cu18:15
- Electron Configuration of Atoms Cont'd20:22
- Electron Configuration of Monatomic Ions: Al20:23
- Electron Configuration of Monatomic Ions: Al³⁺20:46
- Electron Configuration of Monatomic Ions: Cl21:57
- Electron Configuration of Monatomic Ions: Cl¹⁻22:09
- Electron Configuration Cont'd24:31
- Paramagnetism24:32
- Diamagnetism25:00
- Atomic Radii26:08
- Atomic Radii26:09
- In a Column of the Periodic Table26:25
- In a Row of the Periodic Table26:46
- Ionic Radii27:30
- Ionic Radii27:31
- Anions27:42
- Cations27:57
- Isoelectronic Species28:12
- Ionization Energy29:00
- Ionization Energy29:01
- Electron Affinity31:37
- Electron Affinity31:37
- Summary33:43
- Sample Problem 1: Ground State Configuration and Orbital Box Diagram34:21
- Fe34:48
- P35:32
- Sample Problem 236:38
- 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

- Intro0:00
- Lesson Overview0:08
- Introduction1:10
- Types of Chemical Bonds1:53
- Ionic Bond1:54
- Molecular Bond2:42
- Electronegativity and Bond Polarity3:26
- Electronegativity (EN)3:27
- Periodic Trend4:36
- Electronegativity and Bond Polarity Cont'd6:04
- Bond Polarity: Polar Covalent Bond6:05
- Bond Polarity: Nonpolar Covalent Bond8:53
- Lewis Electron Dot Structure of Atoms9:48
- Lewis Electron Dot Structure of Atoms9:49
- Lewis Structures of Polyatomic Species12:51
- Single Bonds12:52
- Double Bonds13:28
- Nonbonding Electrons13:59
- Lewis Structures of Polyatomic Species Cont'd14:45
- Drawing Lewis Structures: Step 114:48
- Drawing Lewis Structures: Step 215:16
- Drawing Lewis Structures: Step 315:52
- Drawing Lewis Structures: Step 417:31
- Drawing Lewis Structures: Step 519:08
- Drawing Lewis Structure Example: Carbonate19:33
- Resonance and Formal Charges (FC)24:06
- Resonance Structures24:07
- Formal Charge25:20
- Resonance and Formal Charges Cont'd27:46
- More On Formal Charge27:47
- Resonance and Formal Charges Cont'd28:21
- Good Resonance Structures28:22
- VSEPR Theory31:08
- VSEPR Theory Continue31:09
- VSEPR Theory Cont'd32:53
- VSEPR Geometries32:54
- Steric Number33:04
- Basic Geometry33:50
- Molecular Geometry35:50
- Molecular Polarity37:51
- Steps In Determining Molecular Polarity37:52
- Example 1: Polar38:47
- Example 2: Nonpolar39:10
- Example 3: Polar39:36
- Example 4: Polar40:08
- Bond Properties: Order, Length, and Energy40:38
- Bond Order40:39
- Bond Length41:21
- Bond Energy41:55
- Summary43:09
- Sample Problem 143:42
- XeO₃44:03
- I₃⁻47:02
- SF₅49:16

Advanced Bonding Theories

1h 11m 41s

- Intro0:00
- Lesson Overview0:09
- Introduction0:38
- Valence Bond Theory3:07
- Valence Bond Theory3:08
- spᶟ Hybridized Carbon Atom4:19
- Valence Bond Theory Cont'd6:24
- spᶟ Hybridized6:25
- Hybrid Orbitals For Water7:26
- Valence Bond Theory Cont'd (spᶟ)11:53
- Example 1: NH₃11:54
- Valence Bond Theory Cont'd (sp²)14:48
- sp² Hybridization14:49
- Example 2: BF₃16:44
- Valence Bond Theory Cont'd (sp)22:44
- sp Hybridization22:46
- Example 3: HCN23:38
- Valence Bond Theory Cont'd (sp³d and sp³d²)27:36
- Valence Bond Theory: sp³d and sp³d²27:37
- Molecular Orbital Theory29:10
- Valence Bond Theory Doesn't Always Account For a Molecule's Magnetic Behavior29:11
- Molecular Orbital Theory Cont'd30:37
- Molecular Orbital Theory30:38
- Wavefunctions31:04
- How s-orbitals Can Interact32:23
- Bonding Nature of p-orbitals: Head-on35:34
- Bonding Nature of p-orbitals: Parallel39:04
- Interaction Between s and p-orbital40:45
- Molecular Orbital Diagram For Homonuclear Diatomics: H₂42:21
- Molecular Orbital Diagram For Homonuclear Diatomics: He₂45:23
- Molecular Orbital Diagram For Homonuclear Diatomic: Li₂46:39
- Molecular Orbital Diagram For Homonuclear Diatomic: Li₂⁺47:42
- Molecular Orbital Diagram For Homonuclear Diatomic: B₂48:57
- Molecular Orbital Diagram For Homonuclear Diatomic: N₂54:04
- Molecular Orbital Diagram: Molecular Oxygen55:57
- Molecular Orbital Diagram For Heteronuclear Diatomics: Hydrochloric Acid1:02:16
- Sample Problem 1: Determine the Atomic Hybridization1:07:20
- XeO₃1:07:21
- SF₆1:07:49
- I₃⁻1:08:20
- Sample Problem 21:09:04

IX. Gases, Solids, & Liquids

Gases

35m 6s

- Intro0:00
- Lesson Overview0:07
- The Kinetic Molecular Theory of Gases1:23
- The Kinetic Molecular Theory of Gases1:24
- Parameters To Characterize Gases3:35
- Parameters To Characterize Gases: Pressure3:37
- Interpreting Pressure On a Particulate Level4:43
- Parameters Cont'd6:08
- Units For Expressing Pressure: Psi, Pascal6:19
- Units For Expressing Pressure: mm Hg6:42
- Units For Expressing Pressure: atm6:58
- Units For Expressing Pressure: torr7:24
- Parameters Cont'd8:09
- Parameters To Characterize Gases: Volume8:10
- Common Units of Volume9:00
- Parameters Cont'd9:11
- Parameters To Characterize Gases: Temperature9:12
- Particulate Level9:36
- Parameters To Characterize Gases: Moles10:24
- The Simple Gas Laws10:43
- Gas Laws Are Only Valid For…10:44
- Charles' Law11:24
- The Simple Gas Laws13:13
- Boyle's Law13:14
- The Simple Gas Laws15:28
- Gay-Lussac's Law15:29
- The Simple Gas Laws17:11
- Avogadro's Law17:12
- The Ideal Gas Law18:43
- The Ideal Gas Law: PV = nRT18:44
- Applications of the Ideal Gas Law20:12
- Standard Temperature and Pressure for Gases20:13
- Applications of the Ideal Gas Law21:43
- Ideal Gas Law & Gas Density21:44
- Gas Pressures and Partial Pressures23:18
- Dalton's Law of Partial Pressures23:19
- Gas Stoichiometry24:15
- Stoichiometry Problems Involving Gases24:16
- Using The Ideal Gas Law to Get to Moles25:16
- Using Molar Volume to Get to Moles25:39
- Gas Stoichiometry Cont'd26:03
- Example 1: How Many Liters of O₂ at STP are Needed to Form 10.5 g of Water Vapor?26:04
- Summary28:33
- Sample Problem 1: Calculate the Molar Mass of the Gas29:28
- 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

- Intro0:00
- Lesson Overview0:10
- Introduction0:46
- Intermolecular Forces (IMF)0:47
- Intermolecular Forces of Polar Molecules1:32
- Ion-dipole Forces1:33
- Example: Salt Dissolved in Water1:50
- Coulomb's Law & the Force of Attraction Between Ions and/or Dipoles3:06
- IMF of Polar Molecules cont'd4:36
- Enthalpy of Solvation or Enthalpy of Hydration4:37
- IMF of Polar Molecules cont'd6:01
- Dipole-dipole Forces6:02
- IMF of Polar Molecules cont'd7:22
- Hydrogen Bonding7:23
- Example: Hydrogen Bonding of Water8:06
- IMF of Nonpolar Molecules9:37
- Dipole-induced Dipole Attraction9:38
- IMF of Nonpolar Molecules cont'd11:34
- Induced Dipole Attraction, London Dispersion Forces, or Vand der Waals Forces11:35
- Polarizability13:46
- IMF of Nonpolar Molecules cont'd14:26
- Intermolecular Forces (IMF) and Polarizability14:31
- Properties of Liquids16:48
- Standard Molar Enthalpy of Vaporization16:49
- Trends in Boiling Points of Representative Liquids: H₂O vs. H₂S17:43
- Properties of Liquids cont'd18:36
- Aliphatic Hydrocarbons18:37
- Branched Hydrocarbons20:52
- Properties of Liquids cont'd22:10
- Vapor Pressure22:11
- The Clausius-Clapeyron Equation24:30
- Properties of Liquids cont'd25:52
- Boiling Point25:53
- Properties of Liquids cont'd27:07
- Surface Tension27:08
- Viscosity28:06
- Summary29:04
- Sample Problem 1: Determine Which of the Following Liquids Will Have the Lower Vapor Pressure30:21
- Sample Problem 2: Determine Which of the Following Liquids Will Have the Largest Standard Molar Enthalpy of Vaporization31:37

The Chemistry of Solids

25m 13s

- Intro0:00
- Lesson Overview0:07
- Introduction0:46
- General Characteristics0:47
- Particulate-level Drawing1:09
- The Basic Structure of Solids: Crystal Lattices1:37
- The Unit Cell Defined1:38
- Primitive Cubic2:50
- Crystal Lattices cont'd3:58
- Body-centered Cubic3:59
- Face-centered Cubic5:02
- Lattice Enthalpy and Trends6:27
- Introduction to Lattice Enthalpy6:28
- Equation to Calculate Lattice Enthalpy7:21
- Different Types of Crystalline Solids9:35
- Molecular Solids9:36
- Network Solids10:25
- Phase Changes Involving Solids11:03
- Melting & Thermodynamic Value11:04
- Freezing & Thermodynamic Value11:49
- Phase Changes cont'd12:40
- Sublimation & Thermodynamic Value12:41
- Depositions & Thermodynamic Value13:13
- Phase Diagrams13:40
- Introduction to Phase Diagrams13:41
- Phase Diagram of H₂O: Melting Point14:12
- Phase Diagram of H₂O: Normal Boiling Point14:50
- Phase Diagram of H₂O: Sublimation Point15:02
- Phase Diagram of H₂O: Point C ( Supercritical Point)15:32
- Phase Diagrams cont'd16:31
- Phase Diagram of Dry Ice16:32
- Summary18:15
- 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 Helium22:42

X. Solutions, Rates of Reaction, & Equilibrium

Solutions & Their Behavior

38m 6s

- Intro0:00
- Lesson Overview0:10
- Units of Concentration1:40
- Molarity1:41
- Molality3:30
- Weight Percent4:26
- ppm5:16
- Like Dissolves Like6:28
- Like Dissolves Like6:29
- Factors Affecting Solubility9:35
- The Effect of Pressure: Henry's Law9:36
- The Effect of Temperature on Gas Solubility12:16
- The Effect of Temperature on Solid Solubility14:28
- Colligative Properties16:48
- Colligative Properties16:49
- Changes in Vapor Pressure: Raoult's Law17:19
- Colligative Properties cont'd19:53
- Boiling Point Elevation and Freezing Point Depression19:54
- Colligative Properties cont'd26:13
- Definition of Osmosis26:14
- Osmotic Pressure Example27:11
- Summary31:11
- Sample Problem 1: Calculating Vapor Pressure32:53
- Sample Problem 2: Calculating Molality36:29

Chemical Kinetics

37m 45s

- Intro0:00
- Lesson Overview0:06
- Introduction1:09
- Chemical Kinetics and the Rate of a Reaction1:10
- Factors Influencing Rate1:19
- Introduction cont'd2:27
- How a Reaction Progresses Through Time2:28
- Rate of Change Equation6:02
- Rate Laws7:06
- Definition of Rate Laws7:07
- General Form of Rate Laws7:37
- Rate Laws cont'd11:07
- Rate Orders With Respect to Reactant and Concentration11:08
- Methods of Initial Rates13:38
- Methods of Initial Rates13:39
- Integrated Rate Laws17:57
- Integrated Rate Laws17:58
- Graphically Determine the Rate Constant k18:52
- Reaction Mechanisms21:05
- Step 1: Reversible21:18
- Step 2: Rate-limiting Step21:44
- Rate Law for the Reaction23:28
- Reaction Rates and Temperatures26:16
- Reaction Rates and Temperatures26:17
- The Arrhenius Equation29:06
- Catalysis30:31
- Catalyst30:32
- Summary32:02
- Sample Problem 1: Calculate the Rate Constant and the Time Required for the Reaction to be Completed32:54
- Sample Problem 2: Calculate the Energy of Activation and the Order of the Reaction35:24

Principles of Chemical Equilibrium

34m 9s

- Intro0:00
- Lesson Overview0:08
- Introduction1:02
- The Equilibrium Constant3:08
- The Equilibrium Constant3:09
- The Equilibrium Constant cont'd5:50
- The Equilibrium Concentration and Constant for Solutions5:51
- The Equilibrium Partial Pressure and Constant for Gases7:01
- Relationship of Kc and Kp7:30
- Heterogeneous Equilibria8:23
- Heterogeneous Equilibria8:24
- Manipulating K9:57
- First Way of Manipulating K9:58
- Second Way of Manipulating K11:48
- Manipulating K cont'd12:31
- Third Way of Manipulating K12:32
- The Reaction Quotient Q14:42
- The Reaction Quotient Q14:43
- Q > K16:16
- Q < K16:30
- Q = K16:43
- Le Chatlier's Principle17:32
- Restoring Equilibrium When It is Disturbed17:33
- Disturbing a Chemical System at Equilibrium18:35
- Problem-Solving with ICE Tables19:05
- Determining a Reaction's Equilibrium Constant With ICE Table19:06
- Problem-Solving with ICE Tables cont'd21:03
- Example 1: Calculate O₂(g) at Equilibrium21:04
- Problem-Solving with ICE Tables cont'd22:53
- Example 2: Calculate the Equilibrium Constant22:54
- Summary25:24
- Sample Problem 1: Calculate the Equilibrium Constant27:59
- Sample Problem 2: Calculate The Equilibrium Concentration30:30

XI. Acids & Bases Chemistry

Acid-Base Chemistry

43m 44s

- Intro0:00
- Lesson Overview0:06
- Introduction0:55
- Bronsted-Lowry Acid & Bronsted -Lowry Base0:56
- Water is an Amphiprotic Molecule2:40
- Water Reacting With Itself2:58
- Introduction cont'd4:04
- Strong Acids4:05
- Strong Bases5:18
- Introduction cont'd6:16
- Weak Acids and Bases6:17
- Quantifying Acid-Base Strength7:35
- The pH Scale7:36
- Quantifying Acid-Base Strength cont'd9:55
- The Acid-ionization Constant Ka and pKa9:56
- Quantifying Acid-Base Strength cont'd12:13
- Example: Calculate the pH of a 1.2M Solution of Acetic Acid12:14
- Quantifying Acid-Base Strength15:06
- Calculating the pH of Weak Base Solutions15:07
- Writing Out Acid-Base Equilibria17:45
- Writing Out Acid-Base Equilibria17:46
- Writing Out Acid-Base Equilibria cont'd19:47
- Consider the Following Equilibrium19:48
- Conjugate Base and Conjugate Acid21:18
- Salts Solutions22:00
- Salts That Produce Acidic Aqueous Solutions22:01
- Salts That Produce Basic Aqueous Solutions23:15
- Neutral Salt Solutions24:05
- Diprotic and Polyprotic Acids24:44
- Example: Calculate the pH of a 1.2 M Solution of H₂SO₃24:43
- Diprotic and Polyprotic Acids cont'd27:18
- Calculate the pH of a 1.2 M Solution of Na₂SO₃27:19
- Lewis Acids and Bases29:13
- Lewis Acids29:14
- Lewis Bases30:10
- Example: Lewis Acids and Bases31:04
- Molecular Structure and Acidity32:03
- The Effect of Charge32:04
- Within a Period/Row33:07
- Molecular Structure and Acidity cont'd34:17
- Within a Group/Column34:18
- Oxoacids35:58
- Molecular Structure and Acidity cont'd37:54
- Carboxylic Acids37:55
- Hydrated Metal Cations39:23
- Summary40:39
- Sample Problem 1: Calculate the pH of a 1.2 M Solution of NH₃41:20
- Sample Problem 2: Predict If The Following Slat Solutions are Acidic, Basic, or Neutral42:37

Applications of Aqueous Equilibria

55m 26s

- Intro0:00
- Lesson Overview0:07
- Calculating pH of an Acid-Base Mixture0:53
- Equilibria Involving Direct Reaction With Water0:54
- When a Bronsted-Lowry Acid and Base React1:12
- After Neutralization Occurs2:05
- Calculating pH of an Acid-Base Mixture cont'd2:51
- Example: Calculating pH of an Acid-Base Mixture, Step 1 - Neutralization2:52
- Example: Calculating pH of an Acid-Base Mixture, Step 2 - React With H₂O5:24
- Buffers7:45
- Introduction to Buffers7:46
- When Acid is Added to a Buffer8:50
- When Base is Added to a Buffer9:54
- Buffers cont'd10:41
- Calculating the pH10:42
- Calculating the pH When 0.010 mol NaOH is Added to 1.0 L of the Buffer14:03
- Buffers cont'd14:10
- Calculating the pH When 0.010 mol NaOH is Added to 1.0 L of the Buffer: Step 1 -Neutralization14:11
- Calculating the pH When 0.010 mol NaOH is Added to 1.0 L of the Buffer: Step 2- ICE Table15:22
- Buffer Preparation and Capacity16:38
- Example: Calculating the pH of a Buffer Solution16:42
- Effective Buffer18:40
- Acid-Base Titrations19:33
- Acid-Base Titrations: Basic Setup19:34
- Acid-Base Titrations cont'd22:12
- Example: Calculate the pH at the Equivalence Point When 0.250 L of 0.0350 M HClO is Titrated With 1.00 M KOH22:13
- Acid-Base Titrations cont'd25:38
- Titration Curve25:39
- Solubility Equilibria33:07
- Solubility of Salts33:08
- Solubility Product Constant: Ksp34:14
- Solubility Equilibria cont'd34:58
- Q < Ksp34:59
- Q > Ksp35:34
- Solubility Equilibria cont'd36:03
- Common-ion Effect36:04
- Example: Calculate the Solubility of PbCl₂ in 0.55 M NaCl36:30
- Solubility Equilibria cont'd39:02
- When a Solid Salt Contains the Conjugate of a Weak Acid39:03
- Temperature and Solubility40:41
- Complexation Equilibria41:10
- Complex Ion41:11
- Complex Ion Formation Constant: Kf42:26
- Summary43:35
- Sample Problem 1: Question44:23
- Sample Problem 1: Part a) Calculate the pH at the Beginning of the Titration45:48
- Sample Problem 1: Part b) Calculate the pH at the Midpoint or Half-way Point48:04
- Sample Problem 1: Part c) Calculate the pH at the Equivalence Point48:32
- Sample Problem 1: Part d) Calculate the pH After 27.50 mL of the Acid was Added53:00

XII. Thermodynamics & Electrochemistry

Entropy & Free Energy

36m 13s

- Intro0:00
- Lesson Overview0:08
- Introduction0:53
- Introduction to Entropy1:37
- Introduction to Entropy1:38
- Entropy and Heat Flow6:31
- Recall Thermodynamics6:32
- Entropy is a State Function6:54
- ∆S and Heat Flow7:28
- Entropy and Heat Flow cont'd8:18
- Entropy and Heat Flow: Equations8:19
- Endothermic Processes: ∆S > 08:44
- The Second Law of Thermodynamics10:04
- Total ∆S = ∆S of System + ∆S of Surrounding10:05
- Nature Favors Processes Where The Amount of Entropy Increases10:22
- The Third Law of Thermodynamics11:55
- The Third Law of Thermodynamics & Zero Entropy11:56
- Problem-Solving involving Entropy12:36
- Endothermic Process and ∆S12:37
- Exothermic Process and ∆S13:19
- Problem-Solving cont'd13:46
- Change in Physical States: From Solid to Liquid to Gas13:47
- Change in Physical States: All Gases15:02
- Problem-Solving cont'd15:56
- Calculating the ∆S for the System, Surrounding, and Total15:57
- Example: Calculating the Total ∆S16:17
- Problem-Solving cont'd18:36
- Problems Involving Standard Molar Entropies of Formation18:37
- Introduction to Gibb's Free Energy20:09
- Definition of Free Energy ∆G20:10
- Spontaneous Process and ∆G20:19
- Gibb's Free Energy cont'd22:28
- Standard Molar Free Energies of Formation22:29
- The Free Energies of Formation are Zero for All Compounds in the Standard State22:42
- Gibb's Free Energy cont'd23:31
- ∆G° of the System = ∆H° of the System - T∆S° of the System23:32
- Predicting Spontaneous Reaction Based on the Sign of ∆G° of the System24:24
- Gibb's Free Energy cont'd26:32
- Effect of reactant and Product Concentration on the Sign of Free Energy26:33
- ∆G° of Reaction = -RT ln K27:18
- Summary28:12
- Sample Problem 1: Calculate ∆S° of Reaction28:48
- Sample Problem 2: Calculate the Temperature at Which the Reaction Becomes Spontaneous31:18
- Sample Problem 3: Calculate Kp33:47

Electrochemistry

41m 16s

- Intro0:00
- Lesson Overview0:08
- Introduction0:53
- Redox Reactions1:42
- Oxidation-Reduction Reaction Overview1:43
- Redox Reactions cont'd2:37
- Which Reactant is Being Oxidized and Which is Being Reduced?2:38
- Redox Reactions cont'd6:34
- Balance Redox Reaction In Neutral Solutions6:35
- Redox Reactions cont'd10:37
- Balance Redox Reaction In Acidic and Basic Solutions: Step 110:38
- Balance Redox Reaction In Acidic and Basic Solutions: Step 2 - Balance Each Half-Reaction11:22
- Redox Reactions cont'd12:19
- Balance Redox Reaction In Acidic and Basic Solutions: Step 2 - Balance Hydrogen12:20
- Redox Reactions cont'd14:30
- Balance Redox Reaction In Acidic and Basic Solutions: Step 314:34
- Balance Redox Reaction In Acidic and Basic Solutions: Step 415:38
- Voltaic Cells17:01
- Voltaic Cell or Galvanic Cell17:02
- Cell Notation22:03
- Electrochemical Potentials25:22
- Electrochemical Potentials25:23
- Electrochemical Potentials cont'd26:07
- Table of Standard Reduction Potentials26:08
- The Nernst Equation30:41
- The Nernst Equation30:42
- It Can Be Shown That At Equilibrium E =0.0032:15
- Gibb's Free Energy and Electrochemistry32:46
- Gibbs Free Energy is Relatively Small if the Potential is Relatively High32:47
- When E° is Very Large33:39
- Charge, Current and Time33:56
- A Battery Has Three Main Parameters33:57
- A Simple Equation Relates All of These Parameters34:09
- Summary34:50
- Sample Problem 1: Redox Reaction35:26
- Sample Problem 2: Battery38:00

XIII. Transition Elements & Coordination Compounds

The Chemistry of The Transition Metals

39m 3s

- Intro0:00
- Lesson Overview0:11
- Coordination Compounds1:20
- Coordination Compounds1:21
- Nomenclature of Coordination Compounds2:48
- Rule 13:01
- Rule 23:12
- Rule 34:07
- Nomenclature cont'd4:58
- Rule 44:59
- Rule 55:13
- Rule 65:35
- Rule 76:19
- Rule 86:46
- Nomenclature cont'd7:39
- Rule 97:40
- Rule 107:45
- Rule 118:00
- Nomenclature of Coordination Compounds: NH₄[PtCl₃NH₃]8:11
- Nomenclature of Coordination Compounds: [Cr(NH₃)₄(OH)₂]Br9:31
- Structures of Coordination Compounds10:54
- Coordination Number or Steric Number10:55
- Commonly Observed Coordination Numbers and Geometries: 411:14
- Commonly Observed Coordination Numbers and Geometries: 612:00
- Isomers of Coordination Compounds13:13
- Isomers of Coordination Compounds13:14
- Geometrical Isomers of CN = 6 Include: ML₄L₂'13:30
- Geometrical Isomers of CN = 6 Include: ML₃L₃'15:07
- Isomers cont'd17:00
- Structural Isomers Overview17:01
- Structural Isomers: Ionization18:06
- Structural Isomers: Hydrate19:25
- Structural Isomers: Linkage20:11
- Structural Isomers: Coordination Isomers21:05
- Electronic Structure22:25
- Crystal Field Theory22:26
- Octahedral and Tetrahedral Field22:54
- Electronic Structure cont'd25:43
- Vanadium (II) Ion in an Octahedral Field25:44
- Chromium(III) Ion in an Octahedral Field26:37
- Electronic Structure cont'd28:47
- Strong-Field Ligands and Weak-Field Ligands28:48
- Implications of Electronic Structure30:08
- Compare the Magnetic Properties of: [Fe(OH₂)₆]²⁺ vs. [Fe(CN)₆]⁴⁻30:09
- Discussion on Color31:57
- Summary34:41
- Sample Problem 1: Name the Following Compound [Fe(OH)(OH₂)₅]Cl₂35:08
- Sample Problem 1: Name the Following Compound [Co(NH₃)₃(OH₂)₃]₂(SO₄)₃36:24
- Sample Problem 2: Change in Magnetic Properties37:30

XIV. Nuclear Chemistry

Nuclear Chemistry

16m 39s

- Intro0:00
- Lesson Overview0:06
- Introduction0:40
- Introduction to Nuclear Reactions0:41
- Types of Radioactive Decay2:10
- Alpha Decay2:11
- Beta Decay3:27
- Gamma Decay4:40
- Other Types of Particles of Varying Energy5:40
- Nuclear Equations6:47
- Nuclear Equations6:48
- Nuclear Decay9:28
- Nuclear Decay and the First-Order Kinetics9:29
- Summary11:31
- Sample Problem 1: Complete the Following Nuclear Equations12:13
- Sample Problem 2: How Old is the Rock?14:21

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For more information, please see full course syllabus of General Chemistry

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0 answers

Post by Nasser Fiture on January 12 at 01:55:24 AM

If a proposed reaction mechanism fits the experimental rate law and reaction stoichiometry, then the mechanism is known to be completely correct.

True/False

0 answers

Post by Nasser Fiture on January 12 at 01:54:20 AM

all reactions which show first-order kinetics are also known as unimolecular reactions.

0 answers

Post by Nasser Fiture on January 12 at 01:53:31 AM

the rate determining step of a chemical reaction is the step which has the transition state highest in potential energy.

True/False

0 answers

Post by Jarrah alharbi on September 18, 2016

is the ----> t1/2=In1/2 / k

or t1/2=In2/k

2 answers

Last reply by: Professor Franklin Ow

Mon Jun 20, 2016 11:35 PM

Post by Parth Shorey on June 20, 2016

I don't understand why you used the negatives in front of the recants? You said because of the slopes? I still don't understand.

1 answer

Last reply by: Professor Franklin Ow

Thu Jun 16, 2016 3:29 PM

Post by Parth Shorey on June 13, 2016

Is the Q&A active?

1 answer

Last reply by: Professor Franklin Ow

Thu May 28, 2015 12:28 PM

Post by BRAD POOLE on May 7, 2015

At about the 17 min mark you said that the example was a 2nd order reaction. How did you come up with this? Are you just going by whatever your units are for "k"? I always thought you used the exponents of the reactants to figure out what order it was, then again could be why I can't seem to get these kinetics problems.

0 answers

Post by Saadman Elman on January 14, 2015

Great lecture as usual!

0 answers

Post by Saadman Elman on January 14, 2015

What you mean by we can write co-effecient as rate order ONLY for elementary steps. What you mean by elementary steps exactly?

2 answers

Last reply by: David Gonzalez

Thu Jul 31, 2014 12:25 PM

Post by David Gonzalez on July 31, 2014

Hi Professor Ow. First of all, great lecture. Although, there is one problem that I have.

In the example (around the 16-minute mark), when determining the order for S2O8, you mentioned that there was a 3x increase in initial rate from 0.015 to 0.044. I'm confused, because 0.015 x 3 is 0.045 - don't these problems need to be exact? Or can they sometimes be slightly off?

Thanks.

1 answer

Last reply by: Professor Franklin Ow

Tue Jun 24, 2014 1:29 PM

Post by brandon joyner on June 24, 2014

Also why for Iodine did you only do the first experiment for the second experiment you went from experiment 3 to 1.

4 answers

Last reply by: brandon joyner

Fri Jun 27, 2014 1:46 PM

Post by brandon joyner on June 24, 2014

For Iodine how is it a triple jump when all you had to do was go up 2?

1 answer

Last reply by: Professor Franklin Ow

Wed May 21, 2014 1:54 AM

Post by Ashley Gwemende on May 20, 2014

Where do I find a lecture of how to read a potential energy diagram ?