Franklin Ow

Energy & Chemical Reactions

Slide Duration:Table of Contents

Section 1: 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

Section 2: 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

Section 3: 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

Section 4: 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

Section 5: 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

Section 6: 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

Section 7: 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

Section 8: 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

Section 9: 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

Section 10: 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

Section 11: 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

Section 12: 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

Section 13: 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

Section 14: 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

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

# General Chemistry Energy & Chemical Reactions

Section 5: Thermochemistry: Lecture 1 | 55:28 min

Lecture Description

In this lesson our instructor talks about energy and chemical reactions. First, he talks about the first law of thermodynamics, calculating ∆U, Q, and W, and constant-volume calorimetry. Then he discusses constant-pressure conditions, calculating enthalpy, phase changes, heats of reaction, and using standard enthalpies of formation. He also lectures on enthalpy from a series of reactions and coffee-cup calorimetry.

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1 answer

Last reply by:

Wed Aug 26, 2020 4:51 PM

Post by Abdullah Alqahtani on June 24, 2017

YOU are very bad lecturer

1 answer

Last reply by: Professor Franklin Ow

Wed Jul 20, 2016 11:02 PM

Post by Magic Fu on July 20, 2016

In 15:44. Why Styrofoam has smaller heat capacities, and still need lot of heat? Shouldn't it be the opposite instead?

0 answers

Post by Peter Ke on September 16, 2015

At 45:00, why did you multiply 1.00 g/mL x 100mL because 1 Liter = 1000 mL so I was thinking that you should multiply it by 1000 and get 1000 grams of solution?

Also, when you got -339 kJ shouldn't it be 339 J without the kilo? Because if you were to set up the equation -100g(4.18 J/gC)(24.21 C - 23.40 C), don't the grams and celsius cancel each other out and you would be left with J?

Please correct me if I'm wrong.

0 answers

Post by Peter Ke on September 15, 2015

At 23:12. I understand how you got Qreaction = - Qbomb, because you subtract Qbomb on both side but I don't understand how you got -Cbomb(DeltaT). Please explain. Thx!

0 answers

Post by Peter Ke on September 15, 2015

At 16:04, can you show me how you got 30.7, cause the answer I got was like 495.1 which totally made no sense. So, please show me how you got 30.7. Thx

0 answers

Post by Peter Ke on September 15, 2015

At 8:59, why 488 is negative? Please explain it in a easier way. Thx

1 answer

Last reply by: Professor Franklin Ow

Mon Jun 1, 2015 5:41 PM

Post by Ivan de La Grange on May 31, 2015

For Sample problem 1, shouldn't it be .0327 grams?

2 answers

Last reply by: Denise Bermudez

Sun Mar 8, 2015 10:04 AM

Post by Denise Bermudez on March 2, 2015

hi professor

can you please explain or re-do how you got the answer of 30.7 degrees celsius. I set the equation correctly but i cant figure out how you got your answer with it. The part that confused me was the *4.18 J/ g celsius.

thanks!

0 answers

Post by Saadman Elman on January 31, 2015

Please look at 13:00 - 14:00 min, You did first part of the problem but you didn't do the second part of the problem. The question was asked Calculate delta U for burning of fuel if the system gives off 875 joule. When all is said and done, the answer that you will end up is -998.6 Joule which is the answer for delta you. You skipped that part. Please verify my answer.

0 answers

Post by Saadman Elman on January 9, 2015

Correct me if i am wrong. You said Insulator such as Styrofoam has small heat capacities. Since they retain the heat so i think they have higher heat capacities.

1 answer

Last reply by: Ryan Lozon

Sun Dec 21, 2014 11:34 AM

Post by Ryan Lozon on December 21, 2014

You mean insulators such as styrofoam having larger heat capacities than metals, right? Because they require more energy to raise one gram by 1 Celsius?

3 answers

Last reply by: Professor Franklin Ow

Mon Aug 4, 2014 11:31 PM

Post by brandon joyner on July 1, 2014

Why does 725 have a negative sign?

2 answers

Last reply by: Peidong, He

Sun Mar 16, 2014 11:53 AM

Post by Peidong, He on March 15, 2014

should insulators, such as Styrofoam, have a higher specific heat capacity?

0 answers

Post by KyungYeop Kim on July 9, 2013

Regarding a specific situation in which a nonspontaneous reaction under standard conditions becomes spontaneous at lower temperature, how can I describe this phenomenon in relation to enthalpy, entropy, and free energy? and how can I explain it in terms of temperature change? I've succeeded so far in determining that Î”G>0 since it's nonspontaneous under standard conditions, but what about Î”H and Î”S?

Given the equation = Î”G = Î”H*TÎ”S; I think the fact that temperature(T) can go either from positive to negative or negative to negative seems to confuse me. Are we assuming, in saying lowering temperature, that the T goes from negative(-) to negative(-)?

I know it's a complex problem, and I apologize if I'm asking too much, but I would like to know what the answer is and why. Thank you always!

0 answers

Post by KyungYeop Kim on July 9, 2013

I have a question about effective nuclear charge. As you go down a group in the periodic table, why is it that the effective nuclear charge decreases? from what I know, is it true that as the attraction decreases down the group, it somehow counterbalances the increase in nuclear charge? I'm confused.