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Franklin Ow

Franklin Ow

Tools in Quantitative Chemistry

Slide Duration:

Table of Contents

I. Basic Concepts & Measurement of Chemistry
Basic Concepts of Chemistry

16m 26s

Intro
0:00
Lesson Overview
0:07
Introduction
0:56
What is Chemistry?
0:57
What is Matter?
1:16
Solids
1:43
General Characteristics
1:44
Particulate-level Drawing of Solids
2:34
Liquids
3:39
General Characteristics of Liquids
3:40
Particulate-level Drawing of Liquids
3:55
Gases
4:23
General Characteristics of Gases
4:24
Particulate-level Drawing Gases
5:05
Classification of Matter
5:27
Classification of Matter
5:26
Pure Substances
5:54
Pure Substances
5:55
Mixtures
7:06
Definition of Mixtures
7:07
Homogeneous Mixtures
7:11
Heterogeneous Mixtures
7:52
Physical and Chemical Changes/Properties
8:18
Physical Changes Retain Chemical Composition
8:19
Chemical Changes Alter Chemical Composition
9:32
Physical and Chemical Changes/Properties, cont'd
10:55
Physical Properties
10:56
Chemical Properties
11:42
Sample Problem 1: Chemical & Physical Change
12:22
Sample Problem 2: Element, Compound, or Mixture?
13:52
Sample Problem 3: Classify Each of the Following Properties as chemical or Physical
15:03
Tools in Quantitative Chemistry

29m 22s

Intro
0:00
Lesson Overview
0:07
Units of Measurement
1:23
The International System of Units (SI): Mass, Length, and Volume
1:39
Percent Error
2:17
Percent Error
2:18
Example: Calculate the Percent Error
2:56
Standard Deviation
3:48
Standard Deviation Formula
3:49
Standard Deviation cont'd
4:42
Example: Calculate Your Standard Deviation
4:43
Precisions vs. Accuracy
6:25
Precision
6:26
Accuracy
7:01
Significant Figures and Uncertainty
7:50
Consider the Following (2) Rulers
7:51
Consider the Following Graduated Cylinder
11:30
Identifying Significant Figures
12:43
The Rules of Sig Figs Overview
12:44
The Rules for Sig Figs: All Nonzero Digits Are Significant
13:21
The Rules for Sig Figs: A Zero is Significant When It is In-Between Nonzero Digits
13:28
The Rules for Sig Figs: A Zero is Significant When at the End of a Decimal Number
14:02
The Rules for Sig Figs: A Zero is not significant When Starting a Decimal Number
14:27
Using Sig Figs in Calculations
15:03
Using Sig Figs for Multiplication and Division
15:04
Using Sig Figs for Addition and Subtraction
15:48
Using Sig Figs for Mixed Operations
16:11
Dimensional Analysis
16:20
Dimensional Analysis Overview
16:21
General Format for Dimensional Analysis
16:39
Example: How Many Miles are in 17 Laps?
17:17
Example: How Many Grams are in 1.22 Pounds?
18:40
Dimensional Analysis cont'd
19:43
Example: How Much is Spent on Diapers in One Week?
19:44
Dimensional Analysis cont'd
21:03
SI Prefixes
21:04
Dimensional Analysis cont'd
22:03
500 mg → ? kg
22:04
34.1 cm → ? um
24:03
Summary
25:11
Sample Problem 1: Dimensional Analysis
26:09
II. Atoms, Molecules, and Ions
Atoms, Molecules, and Ions

52m 18s

Intro
0:00
Lesson Overview
0:08
Introduction to Atomic Structure
1:03
Introduction to Atomic Structure
1:04
Plum Pudding Model
1:26
Introduction to Atomic Structure Cont'd
2:07
John Dalton's Atomic Theory: Number 1
2:22
John Dalton's Atomic Theory: Number 2
2:50
John Dalton's Atomic Theory: Number 3
3:07
John Dalton's Atomic Theory: Number 4
3:30
John Dalton's Atomic Theory: Number 5
3:58
Introduction to Atomic Structure Cont'd
5:21
Ernest Rutherford's Gold Foil Experiment
5:22
Introduction to Atomic Structure Cont'd
7:42
Implications of the Gold Foil Experiment
7:43
Relative Masses and Charges
8:18
Isotopes
9:02
Isotopes
9:03
Introduction to The Periodic Table
12:17
The Periodic Table of the Elements
12:18
Periodic Table, cont'd
13:56
Metals
13:57
Nonmetals
14:25
Semimetals
14:51
Periodic Table, cont'd
15:57
Group I: The Alkali Metals
15:58
Group II: The Alkali Earth Metals
16:25
Group VII: The Halogens
16:40
Group VIII: The Noble Gases
17:08
Ionic Compounds: Formulas, Names, Props.
17:35
Common Polyatomic Ions
17:36
Predicting Ionic Charge for Main Group Elements
18:52
Ionic Compounds: Formulas, Names, Props.
20:36
Naming Ionic Compounds: Rule 1
20:51
Naming Ionic Compounds: Rule 2
21:22
Naming Ionic Compounds: Rule 3
21:50
Naming Ionic Compounds: Rule 4
22:22
Ionic Compounds: Formulas, Names, Props.
22:50
Naming Ionic Compounds Example: Al₂O₃
22:51
Naming Ionic Compounds Example: FeCl₃
23:21
Naming Ionic Compounds Example: CuI₂ 3H₂O
24:00
Naming Ionic Compounds Example: Barium Phosphide
24:40
Naming Ionic Compounds Example: Ammonium Phosphate
25:55
Molecular Compounds: Formulas and Names
26:42
Molecular Compounds: Formulas and Names
26:43
The Mole
28:10
The Mole is 'A Chemist's Dozen'
28:11
It is a Central Unit, Connecting the Following Quantities
30:01
The Mole, cont'd
32:07
Atomic Masses
32:08
Example: How Many Moles are in 25.7 Grams of Sodium?
32:28
Example: How Many Atoms are in 1.2 Moles of Carbon?
33:17
The Mole, cont'd
34:25
Example: What is the Molar Mass of Carbon Dioxide?
34:26
Example: How Many Grams are in 1.2 Moles of Carbon Dioxide?
25:46
Percentage Composition
36:43
Example: How Many Grams of Carbon Contained in 65.1 Grams of Carbon Dioxide?
36:44
Empirical and Molecular Formulas
39:19
Empirical Formulas
39:20
Empirical Formula & Elemental Analysis
40:21
Empirical and Molecular Formulas, cont'd
41:24
Example: Determine Both the Empirical and Molecular Formulas - Step 1
41:25
Example: Determine Both the Empirical and Molecular Formulas - Step 2
43:18
Summary
46:22
Sample Problem 1: Determine the Empirical Formula of Lithium Fluoride
47:10
Sample Problem 2: How Many Atoms of Carbon are Present in 2.67 kg of C₆H₆?
49:21
III. Chemical Reactions
Chemical Reactions

43m 24s

Intro
0:00
Lesson Overview
0:06
The Law of Conservation of Mass and Balancing Chemical Reactions
1:49
The Law of Conservation of Mass
1:50
Balancing Chemical Reactions
2:50
Balancing Chemical Reactions Cont'd
3:40
Balance: N₂ + H₂ → NH₃
3:41
Balance: CH₄ + O₂ → CO₂ + H₂O
7:20
Balancing Chemical Reactions Cont'd
9:49
Balance: C₂H₆ + O₂ → CO₂ + H₂O
9:50
Intro to Chemical Equilibrium
15:32
When an Ionic Compound Full Dissociates
15:33
When an Ionic Compound Incompletely Dissociates
16:14
Dynamic Equilibrium
17:12
Electrolytes and Nonelectrolytes
18:03
Electrolytes
18:04
Strong Electrolytes and Weak Electrolytes
18:55
Nonelectrolytes
19:23
Predicting the Product(s) of an Aqueous Reaction
20:02
Single-replacement
20:03
Example: Li (s) + CuCl₂ (aq) → 2 LiCl (aq) + Cu (s)
21:03
Example: Cu (s) + LiCl (aq) → NR
21:23
Example: Zn (s) + 2HCl (aq) → ZnCl₂ (aq) + H₂ (g)
22:32
Predicting the Product(s) of an Aqueous Reaction
23:37
Double-replacement
23:38
Net-ionic Equation
25:29
Predicting the Product(s) of an Aqueous Reaction
26:12
Solubility Rules for Ionic Compounds
26:13
Predicting the Product(s) of an Aqueous Reaction
28:10
Neutralization Reactions
28:11
Example: HCl (aq) + NaOH (aq) → ?
28:37
Example: H₂SO₄ (aq) + KOH (aq) → ?
29:25
Predicting the Product(s) of an Aqueous Reaction
30:20
Certain Aqueous Reactions can Produce Unstable Compounds
30:21
Example 1
30:52
Example 2
32:16
Example 3
32:54
Summary
33:54
Sample Problem 1
34:55
ZnCO₃ (aq) + H₂SO₄ (aq) → ?
35:09
NH₄Br (aq) + Pb(C₂H₃O₂)₂ (aq) → ?
36:02
KNO₃ (aq) + CuCl₂ (aq) → ?
37:07
Li₂SO₄ (aq) + AgNO₃ (aq) → ?
37:52
Sample Problem 2
39:09
Question 1
39:10
Question 2
40:36
Question 3
41:47
Chemical Reactions II

55m 40s

Intro
0:00
Lesson Overview
0:10
Arrhenius Definition
1:15
Arrhenius Acids
1:16
Arrhenius Bases
3:20
The Bronsted-Lowry Definition
4:48
Acids Dissolve In Water and Donate a Proton to Water: Example 1
4:49
Acids Dissolve In Water and Donate a Proton to Water: Example 2
6:54
Monoprotic Acids & Polyprotic Acids
7:58
Strong Acids
11:30
Bases Dissolve In Water and Accept a Proton From Water
12:41
Strong Bases
16:36
The Autoionization of Water
17:42
Amphiprotic
17:43
Water Reacts With Itself
18:24
Oxides of Metals and Nonmetals
20:08
Oxides of Metals and Nonmetals Overview
20:09
Oxides of Nonmetals: Acidic Oxides
21:23
Oxides of Metals: Basic Oxides
24:08
Oxidation-Reduction (Redox) Reactions
25:34
Redox Reaction Overview
25:35
Oxidizing and Reducing Agents
27:02
Redox Reaction: Transfer of Electrons
27:54
Oxidation-Reduction Reactions Cont'd
29:55
Oxidation Number Overview
29:56
Oxidation Number of Homonuclear Species
31:17
Oxidation Number of Monatomic Ions
32:58
Oxidation Number of Fluorine
33:27
Oxidation Number of Oxygen
34:00
Oxidation Number of Chlorine, Bromine, and Iodine
35:07
Oxidation Number of Hydrogen
35:30
Net Sum of All Oxidation Numbers In a Compound
36:21
Oxidation-Reduction Reactions Cont'd
38:19
Let's Practice Assigning Oxidation Number
38:20
Now Let's Apply This to a Chemical Reaction
41:07
Summary
44:19
Sample Problems
45:29
Sample Problem 1
45:30
Sample Problem 2: Determine the Oxidizing and Reducing Agents
48:48
Sample Problem 3: Determine the Oxidizing and Reducing Agents
50:43
IV. Stoichiometry
Stoichiometry I

42m 10s

Intro
0:00
Lesson Overview
0:23
Mole to Mole Ratios
1:32
Example 1: In 1 Mole of H₂O, How Many Moles Are There of Each Element?
1:53
Example 2: In 2.6 Moles of Water, How Many Moles Are There of Each Element?
2:24
Mole to Mole Ratios Cont'd
5:13
Balanced Chemical Reaction
5:14
Mole to Mole Ratios Cont'd
7:25
Example 3: How Many Moles of Ammonia Can Form If you Have 3.1 Moles of H₂?
7:26
Example 4: How Many Moles of Hydrogen Gas Are Required to React With 6.4 Moles of Nitrogen Gas?
9:08
Mass to mass Conversion
11:06
Mass to mass Conversion
11:07
Example 5: How Many Grams of Ammonia Can Form If You Have 3.1 Grams of H₂?
12:37
Example 6: How Many Grams of Hydrogen Gas Are Required to React With 6.4 Grams of Nitrogen Gas?
15:34
Example 7: How Man Milligrams of Ammonia Can Form If You Have 1.2 kg of H₂?
17:29
Limiting Reactants, Percent Yields
20:42
Limiting Reactants, Percent Yields
20:43
Example 8: How Many Grams of Ammonia Can Form If You Have 3.1 Grams of H₂ and 3.1 Grams of N₂
22:25
Percent Yield
25:30
Example 9: How Many Grams of The Excess Reactant Remains?
26:37
Summary
29:34
Sample Problem 1: How Many Grams of Carbon Are In 2.2 Kilograms of Carbon Dioxide?
30:47
Sample Problem 2: How Many Milligrams of Carbon Dioxide Can Form From 23.1 Kg of CH₄(g)?
33:06
Sample Problem 3: Part 1
36:10
Sample Problem 3: Part 2 - What Amount Of The Excess Reactant Will Remain?
40:53
Stoichiometry II

42m 38s

Intro
0:00
Lesson Overview
0:10
Molarity
1:14
Solute and Solvent
1:15
Molarity
2:01
Molarity Cont'd
2:59
Example 1: How Many Grams of KBr are Needed to Make 350 mL of a 0.67 M KBr Solution?
3:00
Example 2: How Many Moles of KBr are in 350 mL of a 0.67 M KBr Solution?
5:44
Example 3: What Volume of a 0.67 M KBr Solution Contains 250 mg of KBr?
7:46
Dilutions
10:01
Dilution: M₁V₂=M₁V₂
10:02
Example 5: Explain How to Make 250 mL of a 0.67 M KBr Solution Starting From a 1.2M Stock Solution
12:04
Stoichiometry and Double-Displacement Precipitation Reactions
14:41
Example 6: How Many grams of PbCl₂ Can Form From 250 mL of 0.32 M NaCl?
15:38
Stoichiometry and Double-Displacement Precipitation Reactions
18:05
Example 7: How Many grams of PbCl₂ Can Form When 250 mL of 0.32 M NaCl and 150 mL of 0.45 Pb(NO₃)₂ Mix?
18:06
Stoichiometry and Neutralization Reactions
21:01
Example 8: How Many Grams of NaOh are Required to Neutralize 4.5 Grams of HCl?
21:02
Stoichiometry and Neutralization Reactions
23:03
Example 9: How Many mL of 0.45 M NaOH are Required to Neutralize 250 mL of 0.89 M HCl?
23:04
Stoichiometry and Acid-Base Standardization
25:28
Introduction to Titration & Standardization
25:30
Acid-Base Titration
26:12
The Analyte & Titrant
26:24
The Experimental Setup
26:49
The Experimental Setup
26:50
Stoichiometry and Acid-Base Standardization
28:38
Example 9: Determine the Concentration of the Analyte
28:39
Summary
32:46
Sample Problem 1: Stoichiometry & Neutralization
35:24
Sample Problem 2: Stoichiometry
37:50
V. Thermochemistry
Energy & Chemical Reactions

55m 28s

Intro
0:00
Lesson Overview
0:14
Introduction
1:22
Recall: Chemistry
1:23
Energy Can Be Expressed In Different Units
1:57
The First Law of Thermodynamics
2:43
Internal Energy
2:44
The First Law of Thermodynamics Cont'd
6:14
Ways to Transfer Internal Energy
6:15
Work Energy
8:13
Heat Energy
8:34
∆U = q + w
8:44
Calculating ∆U, Q, and W
8:58
Changes In Both Volume and Temperature of a System
8:59
Calculating ∆U, Q, and W Cont'd
11:01
The Work Equation
11:02
Example 1: Calculate ∆U For The Burning Fuel
11:45
Calculating ∆U, Q, and W Cont'd
14:09
The Heat Equation
14:10
Calculating ∆U, Q, and W Cont'd
16:03
Example 2: Calculate The Final Temperature
16:04
Constant-Volume Calorimetry
18:05
Bomb Calorimeter
18:06
The Effect of Constant Volume On The Equation For Internal Energy
22:11
Example 3: Calculate ∆U
23:12
Constant-Pressure Conditions
26:05
Constant-Pressure Conditions
26:06
Calculating Enthalpy: Phase Changes
27:29
Melting, Vaporization, and Sublimation
27:30
Freezing, Condensation and Deposition
28:25
Enthalpy Values For Phase Changes
28:40
Example 4: How Much Energy In The Form of heat is Required to Melt 1.36 Grams of Ice?
29:40
Calculating Enthalpy: Heats of Reaction
31:22
Example 5: Calculate The Heat In kJ Associated With The Complete Reaction of 155 g NH₃
31:23
Using Standard Enthalpies of Formation
33:53
Standard Enthalpies of Formation
33:54
Using Standard Enthalpies of Formation
36:12
Example 6: Calculate The Standard Enthalpies of Formation For The Following Reaction
36:13
Enthalpy From a Series of Reactions
39:58
Hess's Law
39:59
Coffee-Cup Calorimetry
42:43
Coffee-Cup Calorimetry
42:44
Example 7: Calculate ∆H° of Reaction
45:10
Summary
47:12
Sample Problem 1
48:58
Sample Problem 2
51:24
VI. Quantum Theory of Atoms
Structure of Atoms

42m 33s

Intro
0:00
Lesson Overview
0:07
Introduction
1:01
Rutherford's Gold Foil Experiment
1:02
Electromagnetic Radiation
2:31
Radiation
2:32
Three Parameters: Energy, Frequency, and Wavelength
2:52
Electromagnetic Radiation
5:18
The Electromagnetic Spectrum
5:19
Atomic Spectroscopy and The Bohr Model
7:46
Wavelengths of Light
7:47
Atomic Spectroscopy Cont'd
9:45
The Bohr Model
9:46
Atomic Spectroscopy Cont'd
12:21
The Balmer Series
12:22
Rydberg Equation For Predicting The Wavelengths of Light
13:04
The Wave Nature of Matter
15:11
The Wave Nature of Matter
15:12
The Wave Nature of Matter
19:10
New School of Thought
19:11
Einstein: Energy
19:49
Hertz and Planck: Photoelectric Effect
20:16
de Broglie: Wavelength of a Moving Particle
21:14
Quantum Mechanics and The Atom
22:15
Heisenberg: Uncertainty Principle
22:16
Schrodinger: Wavefunctions
23:08
Quantum Mechanics and The Atom
24:02
Principle Quantum Number
24:03
Angular Momentum Quantum Number
25:06
Magnetic Quantum Number
26:27
Spin Quantum Number
28:42
The Shapes of Atomic Orbitals
29:15
Radial Wave Function
29:16
Probability Distribution Function
32:08
The Shapes of Atomic Orbitals
34:02
3-Dimensional Space of Wavefunctions
34:03
Summary
35:57
Sample Problem 1
37:07
Sample Problem 2
40:23
VII. Electron Configurations and Periodicity
Periodic Trends

38m 50s

Intro
0:00
Lesson Overview
0:09
Introduction
0:36
Electron Configuration of Atoms
1:33
Electron Configuration & Atom's Electrons
1:34
Electron Configuration Format
1:56
Electron Configuration of Atoms Cont'd
3:01
Aufbau Principle
3:02
Electron Configuration of Atoms Cont'd
6:53
Electron Configuration Format 1: Li, O, and Cl
6:56
Electron Configuration Format 2: Li, O, and Cl
9:11
Electron Configuration of Atoms Cont'd
12:48
Orbital Box Diagrams
12:49
Pauli Exclusion Principle
13:11
Hund's Rule
13:36
Electron Configuration of Atoms Cont'd
17:35
Exceptions to The Aufbau Principle: Cr
17:36
Exceptions to The Aufbau Principle: Cu
18:15
Electron Configuration of Atoms Cont'd
20:22
Electron Configuration of Monatomic Ions: Al
20:23
Electron Configuration of Monatomic Ions: Al³⁺
20:46
Electron Configuration of Monatomic Ions: Cl
21:57
Electron Configuration of Monatomic Ions: Cl¹⁻
22:09
Electron Configuration Cont'd
24:31
Paramagnetism
24:32
Diamagnetism
25:00
Atomic Radii
26:08
Atomic Radii
26:09
In a Column of the Periodic Table
26:25
In a Row of the Periodic Table
26:46
Ionic Radii
27:30
Ionic Radii
27:31
Anions
27:42
Cations
27:57
Isoelectronic Species
28:12
Ionization Energy
29:00
Ionization Energy
29:01
Electron Affinity
31:37
Electron Affinity
31:37
Summary
33:43
Sample Problem 1: Ground State Configuration and Orbital Box Diagram
34:21
Fe
34:48
P
35:32
Sample Problem 2
36:38
Which Has The Larger Ionization Energy: Na or Li?
36:39
Which Has The Larger Atomic Size: O or N ?
37:23
Which Has The Larger Atomic Size: O²⁻ or N³⁻ ?
38:00
VIII. Molecular Geometry & Bonding Theory
Bonding & Molecular Structure

52m 39s

Intro
0:00
Lesson Overview
0:08
Introduction
1:10
Types of Chemical Bonds
1:53
Ionic Bond
1:54
Molecular Bond
2:42
Electronegativity and Bond Polarity
3:26
Electronegativity (EN)
3:27
Periodic Trend
4:36
Electronegativity and Bond Polarity Cont'd
6:04
Bond Polarity: Polar Covalent Bond
6:05
Bond Polarity: Nonpolar Covalent Bond
8:53
Lewis Electron Dot Structure of Atoms
9:48
Lewis Electron Dot Structure of Atoms
9:49
Lewis Structures of Polyatomic Species
12:51
Single Bonds
12:52
Double Bonds
13:28
Nonbonding Electrons
13:59
Lewis Structures of Polyatomic Species Cont'd
14:45
Drawing Lewis Structures: Step 1
14:48
Drawing Lewis Structures: Step 2
15:16
Drawing Lewis Structures: Step 3
15:52
Drawing Lewis Structures: Step 4
17:31
Drawing Lewis Structures: Step 5
19:08
Drawing Lewis Structure Example: Carbonate
19:33
Resonance and Formal Charges (FC)
24:06
Resonance Structures
24:07
Formal Charge
25:20
Resonance and Formal Charges Cont'd
27:46
More On Formal Charge
27:47
Resonance and Formal Charges Cont'd
28:21
Good Resonance Structures
28:22
VSEPR Theory
31:08
VSEPR Theory Continue
31:09
VSEPR Theory Cont'd
32:53
VSEPR Geometries
32:54
Steric Number
33:04
Basic Geometry
33:50
Molecular Geometry
35:50
Molecular Polarity
37:51
Steps In Determining Molecular Polarity
37:52
Example 1: Polar
38:47
Example 2: Nonpolar
39:10
Example 3: Polar
39:36
Example 4: Polar
40:08
Bond Properties: Order, Length, and Energy
40:38
Bond Order
40:39
Bond Length
41:21
Bond Energy
41:55
Summary
43:09
Sample Problem 1
43:42
XeO₃
44:03
I₃⁻
47:02
SF₅
49:16
Advanced Bonding Theories

1h 11m 41s

Intro
0:00
Lesson Overview
0:09
Introduction
0:38
Valence Bond Theory
3:07
Valence Bond Theory
3:08
spᶟ Hybridized Carbon Atom
4:19
Valence Bond Theory Cont'd
6:24
spᶟ Hybridized
6:25
Hybrid Orbitals For Water
7:26
Valence Bond Theory Cont'd (spᶟ)
11:53
Example 1: NH₃
11:54
Valence Bond Theory Cont'd (sp²)
14:48
sp² Hybridization
14:49
Example 2: BF₃
16:44
Valence Bond Theory Cont'd (sp)
22:44
sp Hybridization
22:46
Example 3: HCN
23:38
Valence Bond Theory Cont'd (sp³d and sp³d²)
27:36
Valence Bond Theory: sp³d and sp³d²
27:37
Molecular Orbital Theory
29:10
Valence Bond Theory Doesn't Always Account For a Molecule's Magnetic Behavior
29:11
Molecular Orbital Theory Cont'd
30:37
Molecular Orbital Theory
30:38
Wavefunctions
31:04
How s-orbitals Can Interact
32:23
Bonding Nature of p-orbitals: Head-on
35:34
Bonding Nature of p-orbitals: Parallel
39:04
Interaction Between s and p-orbital
40:45
Molecular Orbital Diagram For Homonuclear Diatomics: H₂
42:21
Molecular Orbital Diagram For Homonuclear Diatomics: He₂
45:23
Molecular Orbital Diagram For Homonuclear Diatomic: Li₂
46:39
Molecular Orbital Diagram For Homonuclear Diatomic: Li₂⁺
47:42
Molecular Orbital Diagram For Homonuclear Diatomic: B₂
48:57
Molecular Orbital Diagram For Homonuclear Diatomic: N₂
54:04
Molecular Orbital Diagram: Molecular Oxygen
55:57
Molecular Orbital Diagram For Heteronuclear Diatomics: Hydrochloric Acid
1:02:16
Sample Problem 1: Determine the Atomic Hybridization
1:07:20
XeO₃
1:07:21
SF₆
1:07:49
I₃⁻
1:08:20
Sample Problem 2
1:09:04
IX. Gases, Solids, & Liquids
Gases

35m 6s

Intro
0:00
Lesson Overview
0:07
The Kinetic Molecular Theory of Gases
1:23
The Kinetic Molecular Theory of Gases
1:24
Parameters To Characterize Gases
3:35
Parameters To Characterize Gases: Pressure
3:37
Interpreting Pressure On a Particulate Level
4:43
Parameters Cont'd
6:08
Units For Expressing Pressure: Psi, Pascal
6:19
Units For Expressing Pressure: mm Hg
6:42
Units For Expressing Pressure: atm
6:58
Units For Expressing Pressure: torr
7:24
Parameters Cont'd
8:09
Parameters To Characterize Gases: Volume
8:10
Common Units of Volume
9:00
Parameters Cont'd
9:11
Parameters To Characterize Gases: Temperature
9:12
Particulate Level
9:36
Parameters To Characterize Gases: Moles
10:24
The Simple Gas Laws
10:43
Gas Laws Are Only Valid For…
10:44
Charles' Law
11:24
The Simple Gas Laws
13:13
Boyle's Law
13:14
The Simple Gas Laws
15:28
Gay-Lussac's Law
15:29
The Simple Gas Laws
17:11
Avogadro's Law
17:12
The Ideal Gas Law
18:43
The Ideal Gas Law: PV = nRT
18:44
Applications of the Ideal Gas Law
20:12
Standard Temperature and Pressure for Gases
20:13
Applications of the Ideal Gas Law
21:43
Ideal Gas Law & Gas Density
21:44
Gas Pressures and Partial Pressures
23:18
Dalton's Law of Partial Pressures
23:19
Gas Stoichiometry
24:15
Stoichiometry Problems Involving Gases
24:16
Using The Ideal Gas Law to Get to Moles
25:16
Using Molar Volume to Get to Moles
25:39
Gas Stoichiometry Cont'd
26:03
Example 1: How Many Liters of O₂ at STP are Needed to Form 10.5 g of Water Vapor?
26:04
Summary
28:33
Sample Problem 1: Calculate the Molar Mass of the Gas
29:28
Sample Problem 2: What Mass of Ag₂O is Required to Form 3888 mL of O₂ Gas When Measured at 734 mm Hg and 25°C?
31:59
Intermolecular Forces & Liquids

33m 47s

Intro
0:00
Lesson Overview
0:10
Introduction
0:46
Intermolecular Forces (IMF)
0:47
Intermolecular Forces of Polar Molecules
1:32
Ion-dipole Forces
1:33
Example: Salt Dissolved in Water
1:50
Coulomb's Law & the Force of Attraction Between Ions and/or Dipoles
3:06
IMF of Polar Molecules cont'd
4:36
Enthalpy of Solvation or Enthalpy of Hydration
4:37
IMF of Polar Molecules cont'd
6:01
Dipole-dipole Forces
6:02
IMF of Polar Molecules cont'd
7:22
Hydrogen Bonding
7:23
Example: Hydrogen Bonding of Water
8:06
IMF of Nonpolar Molecules
9:37
Dipole-induced Dipole Attraction
9:38
IMF of Nonpolar Molecules cont'd
11:34
Induced Dipole Attraction, London Dispersion Forces, or Vand der Waals Forces
11:35
Polarizability
13:46
IMF of Nonpolar Molecules cont'd
14:26
Intermolecular Forces (IMF) and Polarizability
14:31
Properties of Liquids
16:48
Standard Molar Enthalpy of Vaporization
16:49
Trends in Boiling Points of Representative Liquids: H₂O vs. H₂S
17:43
Properties of Liquids cont'd
18:36
Aliphatic Hydrocarbons
18:37
Branched Hydrocarbons
20:52
Properties of Liquids cont'd
22:10
Vapor Pressure
22:11
The Clausius-Clapeyron Equation
24:30
Properties of Liquids cont'd
25:52
Boiling Point
25:53
Properties of Liquids cont'd
27:07
Surface Tension
27:08
Viscosity
28:06
Summary
29:04
Sample Problem 1: Determine Which of the Following Liquids Will Have the Lower Vapor Pressure
30:21
Sample Problem 2: Determine Which of the Following Liquids Will Have the Largest Standard Molar Enthalpy of Vaporization
31:37
The Chemistry of Solids

25m 13s

Intro
0:00
Lesson Overview
0:07
Introduction
0:46
General Characteristics
0:47
Particulate-level Drawing
1:09
The Basic Structure of Solids: Crystal Lattices
1:37
The Unit Cell Defined
1:38
Primitive Cubic
2:50
Crystal Lattices cont'd
3:58
Body-centered Cubic
3:59
Face-centered Cubic
5:02
Lattice Enthalpy and Trends
6:27
Introduction to Lattice Enthalpy
6:28
Equation to Calculate Lattice Enthalpy
7:21
Different Types of Crystalline Solids
9:35
Molecular Solids
9:36
Network Solids
10:25
Phase Changes Involving Solids
11:03
Melting & Thermodynamic Value
11:04
Freezing & Thermodynamic Value
11:49
Phase Changes cont'd
12:40
Sublimation & Thermodynamic Value
12:41
Depositions & Thermodynamic Value
13:13
Phase Diagrams
13:40
Introduction to Phase Diagrams
13:41
Phase Diagram of H₂O: Melting Point
14:12
Phase Diagram of H₂O: Normal Boiling Point
14:50
Phase Diagram of H₂O: Sublimation Point
15:02
Phase Diagram of H₂O: Point C ( Supercritical Point)
15:32
Phase Diagrams cont'd
16:31
Phase Diagram of Dry Ice
16:32
Summary
18:15
Sample Problem 1, Part A: Of the Group I Fluorides, Which Should Have the Highest Lattice Enthalpy?
19:01
Sample Problem 1, Part B: Of the Lithium Halides, Which Should Have the Lowest Lattice Enthalpy?
19:54
Sample Problem 2: How Many Joules of Energy is Required to Melt 546 mg of Ice at Standard Pressure?
20:55
Sample Problem 3: Phase Diagram of Helium
22:42
X. Solutions, Rates of Reaction, & Equilibrium
Solutions & Their Behavior

38m 6s

Intro
0:00
Lesson Overview
0:10
Units of Concentration
1:40
Molarity
1:41
Molality
3:30
Weight Percent
4:26
ppm
5:16
Like Dissolves Like
6:28
Like Dissolves Like
6:29
Factors Affecting Solubility
9:35
The Effect of Pressure: Henry's Law
9:36
The Effect of Temperature on Gas Solubility
12:16
The Effect of Temperature on Solid Solubility
14:28
Colligative Properties
16:48
Colligative Properties
16:49
Changes in Vapor Pressure: Raoult's Law
17:19
Colligative Properties cont'd
19:53
Boiling Point Elevation and Freezing Point Depression
19:54
Colligative Properties cont'd
26:13
Definition of Osmosis
26:14
Osmotic Pressure Example
27:11
Summary
31:11
Sample Problem 1: Calculating Vapor Pressure
32:53
Sample Problem 2: Calculating Molality
36:29
Chemical Kinetics

37m 45s

Intro
0:00
Lesson Overview
0:06
Introduction
1:09
Chemical Kinetics and the Rate of a Reaction
1:10
Factors Influencing Rate
1:19
Introduction cont'd
2:27
How a Reaction Progresses Through Time
2:28
Rate of Change Equation
6:02
Rate Laws
7:06
Definition of Rate Laws
7:07
General Form of Rate Laws
7:37
Rate Laws cont'd
11:07
Rate Orders With Respect to Reactant and Concentration
11:08
Methods of Initial Rates
13:38
Methods of Initial Rates
13:39
Integrated Rate Laws
17:57
Integrated Rate Laws
17:58
Graphically Determine the Rate Constant k
18:52
Reaction Mechanisms
21:05
Step 1: Reversible
21:18
Step 2: Rate-limiting Step
21:44
Rate Law for the Reaction
23:28
Reaction Rates and Temperatures
26:16
Reaction Rates and Temperatures
26:17
The Arrhenius Equation
29:06
Catalysis
30:31
Catalyst
30:32
Summary
32:02
Sample Problem 1: Calculate the Rate Constant and the Time Required for the Reaction to be Completed
32:54
Sample Problem 2: Calculate the Energy of Activation and the Order of the Reaction
35:24
Principles of Chemical Equilibrium

34m 9s

Intro
0:00
Lesson Overview
0:08
Introduction
1:02
The Equilibrium Constant
3:08
The Equilibrium Constant
3:09
The Equilibrium Constant cont'd
5:50
The Equilibrium Concentration and Constant for Solutions
5:51
The Equilibrium Partial Pressure and Constant for Gases
7:01
Relationship of Kc and Kp
7:30
Heterogeneous Equilibria
8:23
Heterogeneous Equilibria
8:24
Manipulating K
9:57
First Way of Manipulating K
9:58
Second Way of Manipulating K
11:48
Manipulating K cont'd
12:31
Third Way of Manipulating K
12:32
The Reaction Quotient Q
14:42
The Reaction Quotient Q
14:43
Q > K
16:16
Q < K
16:30
Q = K
16:43
Le Chatlier's Principle
17:32
Restoring Equilibrium When It is Disturbed
17:33
Disturbing a Chemical System at Equilibrium
18:35
Problem-Solving with ICE Tables
19:05
Determining a Reaction's Equilibrium Constant With ICE Table
19:06
Problem-Solving with ICE Tables cont'd
21:03
Example 1: Calculate O₂(g) at Equilibrium
21:04
Problem-Solving with ICE Tables cont'd
22:53
Example 2: Calculate the Equilibrium Constant
22:54
Summary
25:24
Sample Problem 1: Calculate the Equilibrium Constant
27:59
Sample Problem 2: Calculate The Equilibrium Concentration
30:30
XI. Acids & Bases Chemistry
Acid-Base Chemistry

43m 44s

Intro
0:00
Lesson Overview
0:06
Introduction
0:55
Bronsted-Lowry Acid & Bronsted -Lowry Base
0:56
Water is an Amphiprotic Molecule
2:40
Water Reacting With Itself
2:58
Introduction cont'd
4:04
Strong Acids
4:05
Strong Bases
5:18
Introduction cont'd
6:16
Weak Acids and Bases
6:17
Quantifying Acid-Base Strength
7:35
The pH Scale
7:36
Quantifying Acid-Base Strength cont'd
9:55
The Acid-ionization Constant Ka and pKa
9:56
Quantifying Acid-Base Strength cont'd
12:13
Example: Calculate the pH of a 1.2M Solution of Acetic Acid
12:14
Quantifying Acid-Base Strength
15:06
Calculating the pH of Weak Base Solutions
15:07
Writing Out Acid-Base Equilibria
17:45
Writing Out Acid-Base Equilibria
17:46
Writing Out Acid-Base Equilibria cont'd
19:47
Consider the Following Equilibrium
19:48
Conjugate Base and Conjugate Acid
21:18
Salts Solutions
22:00
Salts That Produce Acidic Aqueous Solutions
22:01
Salts That Produce Basic Aqueous Solutions
23:15
Neutral Salt Solutions
24:05
Diprotic and Polyprotic Acids
24:44
Example: Calculate the pH of a 1.2 M Solution of H₂SO₃
24:43
Diprotic and Polyprotic Acids cont'd
27:18
Calculate the pH of a 1.2 M Solution of Na₂SO₃
27:19
Lewis Acids and Bases
29:13
Lewis Acids
29:14
Lewis Bases
30:10
Example: Lewis Acids and Bases
31:04
Molecular Structure and Acidity
32:03
The Effect of Charge
32:04
Within a Period/Row
33:07
Molecular Structure and Acidity cont'd
34:17
Within a Group/Column
34:18
Oxoacids
35:58
Molecular Structure and Acidity cont'd
37:54
Carboxylic Acids
37:55
Hydrated Metal Cations
39:23
Summary
40:39
Sample Problem 1: Calculate the pH of a 1.2 M Solution of NH₃
41:20
Sample Problem 2: Predict If The Following Slat Solutions are Acidic, Basic, or Neutral
42:37
Applications of Aqueous Equilibria

55m 26s

Intro
0:00
Lesson Overview
0:07
Calculating pH of an Acid-Base Mixture
0:53
Equilibria Involving Direct Reaction With Water
0:54
When a Bronsted-Lowry Acid and Base React
1:12
After Neutralization Occurs
2:05
Calculating pH of an Acid-Base Mixture cont'd
2:51
Example: Calculating pH of an Acid-Base Mixture, Step 1 - Neutralization
2:52
Example: Calculating pH of an Acid-Base Mixture, Step 2 - React With H₂O
5:24
Buffers
7:45
Introduction to Buffers
7:46
When Acid is Added to a Buffer
8:50
When Base is Added to a Buffer
9:54
Buffers cont'd
10:41
Calculating the pH
10:42
Calculating the pH When 0.010 mol NaOH is Added to 1.0 L of the Buffer
14:03
Buffers cont'd
14:10
Calculating the pH When 0.010 mol NaOH is Added to 1.0 L of the Buffer: Step 1 -Neutralization
14:11
Calculating the pH When 0.010 mol NaOH is Added to 1.0 L of the Buffer: Step 2- ICE Table
15:22
Buffer Preparation and Capacity
16:38
Example: Calculating the pH of a Buffer Solution
16:42
Effective Buffer
18:40
Acid-Base Titrations
19:33
Acid-Base Titrations: Basic Setup
19:34
Acid-Base Titrations cont'd
22:12
Example: Calculate the pH at the Equivalence Point When 0.250 L of 0.0350 M HClO is Titrated With 1.00 M KOH
22:13
Acid-Base Titrations cont'd
25:38
Titration Curve
25:39
Solubility Equilibria
33:07
Solubility of Salts
33:08
Solubility Product Constant: Ksp
34:14
Solubility Equilibria cont'd
34:58
Q < Ksp
34:59
Q > Ksp
35:34
Solubility Equilibria cont'd
36:03
Common-ion Effect
36:04
Example: Calculate the Solubility of PbCl₂ in 0.55 M NaCl
36:30
Solubility Equilibria cont'd
39:02
When a Solid Salt Contains the Conjugate of a Weak Acid
39:03
Temperature and Solubility
40:41
Complexation Equilibria
41:10
Complex Ion
41:11
Complex Ion Formation Constant: Kf
42:26
Summary
43:35
Sample Problem 1: Question
44:23
Sample Problem 1: Part a) Calculate the pH at the Beginning of the Titration
45:48
Sample Problem 1: Part b) Calculate the pH at the Midpoint or Half-way Point
48:04
Sample Problem 1: Part c) Calculate the pH at the Equivalence Point
48:32
Sample Problem 1: Part d) Calculate the pH After 27.50 mL of the Acid was Added
53:00
XII. Thermodynamics & Electrochemistry
Entropy & Free Energy

36m 13s

Intro
0:00
Lesson Overview
0:08
Introduction
0:53
Introduction to Entropy
1:37
Introduction to Entropy
1:38
Entropy and Heat Flow
6:31
Recall Thermodynamics
6:32
Entropy is a State Function
6:54
∆S and Heat Flow
7:28
Entropy and Heat Flow cont'd
8:18
Entropy and Heat Flow: Equations
8:19
Endothermic Processes: ∆S > 0
8:44
The Second Law of Thermodynamics
10:04
Total ∆S = ∆S of System + ∆S of Surrounding
10:05
Nature Favors Processes Where The Amount of Entropy Increases
10:22
The Third Law of Thermodynamics
11:55
The Third Law of Thermodynamics & Zero Entropy
11:56
Problem-Solving involving Entropy
12:36
Endothermic Process and ∆S
12:37
Exothermic Process and ∆S
13:19
Problem-Solving cont'd
13:46
Change in Physical States: From Solid to Liquid to Gas
13:47
Change in Physical States: All Gases
15:02
Problem-Solving cont'd
15:56
Calculating the ∆S for the System, Surrounding, and Total
15:57
Example: Calculating the Total ∆S
16:17
Problem-Solving cont'd
18:36
Problems Involving Standard Molar Entropies of Formation
18:37
Introduction to Gibb's Free Energy
20:09
Definition of Free Energy ∆G
20:10
Spontaneous Process and ∆G
20:19
Gibb's Free Energy cont'd
22:28
Standard Molar Free Energies of Formation
22:29
The Free Energies of Formation are Zero for All Compounds in the Standard State
22:42
Gibb's Free Energy cont'd
23:31
∆G° of the System = ∆H° of the System - T∆S° of the System
23:32
Predicting Spontaneous Reaction Based on the Sign of ∆G° of the System
24:24
Gibb's Free Energy cont'd
26:32
Effect of reactant and Product Concentration on the Sign of Free Energy
26:33
∆G° of Reaction = -RT ln K
27:18
Summary
28:12
Sample Problem 1: Calculate ∆S° of Reaction
28:48
Sample Problem 2: Calculate the Temperature at Which the Reaction Becomes Spontaneous
31:18
Sample Problem 3: Calculate Kp
33:47
Electrochemistry

41m 16s

Intro
0:00
Lesson Overview
0:08
Introduction
0:53
Redox Reactions
1:42
Oxidation-Reduction Reaction Overview
1:43
Redox Reactions cont'd
2:37
Which Reactant is Being Oxidized and Which is Being Reduced?
2:38
Redox Reactions cont'd
6:34
Balance Redox Reaction In Neutral Solutions
6:35
Redox Reactions cont'd
10:37
Balance Redox Reaction In Acidic and Basic Solutions: Step 1
10:38
Balance Redox Reaction In Acidic and Basic Solutions: Step 2 - Balance Each Half-Reaction
11:22
Redox Reactions cont'd
12:19
Balance Redox Reaction In Acidic and Basic Solutions: Step 2 - Balance Hydrogen
12:20
Redox Reactions cont'd
14:30
Balance Redox Reaction In Acidic and Basic Solutions: Step 3
14:34
Balance Redox Reaction In Acidic and Basic Solutions: Step 4
15:38
Voltaic Cells
17:01
Voltaic Cell or Galvanic Cell
17:02
Cell Notation
22:03
Electrochemical Potentials
25:22
Electrochemical Potentials
25:23
Electrochemical Potentials cont'd
26:07
Table of Standard Reduction Potentials
26:08
The Nernst Equation
30:41
The Nernst Equation
30:42
It Can Be Shown That At Equilibrium E =0.00
32:15
Gibb's Free Energy and Electrochemistry
32:46
Gibbs Free Energy is Relatively Small if the Potential is Relatively High
32:47
When E° is Very Large
33:39
Charge, Current and Time
33:56
A Battery Has Three Main Parameters
33:57
A Simple Equation Relates All of These Parameters
34:09
Summary
34:50
Sample Problem 1: Redox Reaction
35:26
Sample Problem 2: Battery
38:00
XIII. Transition Elements & Coordination Compounds
The Chemistry of The Transition Metals

39m 3s

Intro
0:00
Lesson Overview
0:11
Coordination Compounds
1:20
Coordination Compounds
1:21
Nomenclature of Coordination Compounds
2:48
Rule 1
3:01
Rule 2
3:12
Rule 3
4:07
Nomenclature cont'd
4:58
Rule 4
4:59
Rule 5
5:13
Rule 6
5:35
Rule 7
6:19
Rule 8
6:46
Nomenclature cont'd
7:39
Rule 9
7:40
Rule 10
7:45
Rule 11
8:00
Nomenclature of Coordination Compounds: NH₄[PtCl₃NH₃]
8:11
Nomenclature of Coordination Compounds: [Cr(NH₃)₄(OH)₂]Br
9:31
Structures of Coordination Compounds
10:54
Coordination Number or Steric Number
10:55
Commonly Observed Coordination Numbers and Geometries: 4
11:14
Commonly Observed Coordination Numbers and Geometries: 6
12:00
Isomers of Coordination Compounds
13:13
Isomers of Coordination Compounds
13:14
Geometrical Isomers of CN = 6 Include: ML₄L₂'
13:30
Geometrical Isomers of CN = 6 Include: ML₃L₃'
15:07
Isomers cont'd
17:00
Structural Isomers Overview
17:01
Structural Isomers: Ionization
18:06
Structural Isomers: Hydrate
19:25
Structural Isomers: Linkage
20:11
Structural Isomers: Coordination Isomers
21:05
Electronic Structure
22:25
Crystal Field Theory
22:26
Octahedral and Tetrahedral Field
22:54
Electronic Structure cont'd
25:43
Vanadium (II) Ion in an Octahedral Field
25:44
Chromium(III) Ion in an Octahedral Field
26:37
Electronic Structure cont'd
28:47
Strong-Field Ligands and Weak-Field Ligands
28:48
Implications of Electronic Structure
30:08
Compare the Magnetic Properties of: [Fe(OH₂)₆]²⁺ vs. [Fe(CN)₆]⁴⁻
30:09
Discussion on Color
31:57
Summary
34:41
Sample Problem 1: Name the Following Compound [Fe(OH)(OH₂)₅]Cl₂
35:08
Sample Problem 1: Name the Following Compound [Co(NH₃)₃(OH₂)₃]₂(SO₄)₃
36:24
Sample Problem 2: Change in Magnetic Properties
37:30
XIV. Nuclear Chemistry
Nuclear Chemistry

16m 39s

Intro
0:00
Lesson Overview
0:06
Introduction
0:40
Introduction to Nuclear Reactions
0:41
Types of Radioactive Decay
2:10
Alpha Decay
2:11
Beta Decay
3:27
Gamma Decay
4:40
Other Types of Particles of Varying Energy
5:40
Nuclear Equations
6:47
Nuclear Equations
6:48
Nuclear Decay
9:28
Nuclear Decay and the First-Order Kinetics
9:29
Summary
11:31
Sample Problem 1: Complete the Following Nuclear Equations
12:13
Sample Problem 2: How Old is the Rock?
14:21
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Lecture Comments (12)

1 answer

Last reply by: Peter Ke
Mon Sep 7, 2015 2:17 PM

Post by Peter Ke on September 7, 2015

At 4:43, why X= 32.6 and not 31.8 and 34.1?

2 answers

Last reply by: Debora Oppong
Fri Jan 29, 2016 1:19 AM

Post by Davon Shackleford on December 20, 2014

Would 5.0008 be four sig figs because the three zeros demonstrate the zero sandwiched between non-zero digit numbers?

5 answers

Last reply by: Denny Yang
Thu Jun 4, 2015 4:41 PM

Post by brandon oneal on June 19, 2014

You said the rule of thumb is to put the prefix multipliers with the base unit. I thought kg was the base unit for mass?

0 answers

Post by James Pelezo on March 24, 2013

Suggestions: a polite FYI...
Accuracy is the variation of data about an 'accepted' value and is quantified by %Error as indicated in your lecture. However, some scientist use the %Error relationship = [(Experimental - Accepted)/Accepted]100%. For a set of results, the %Error may calculate to be a 'negative' number indicating that the average is 'below' the accepted value, or if 'positive' the variation of data is 'above' the accepted value.

Tying Accuracy into Precision:
Precision is the variation of data about the 'average' of a set of measurements all obtained in the same way and is, as indicated, quantified by Standard Deviation. However, I would add that the 'utility' of STDV can be more than just trying to obtain a 'small' range value.

If one is trying to 'reproduce' a set of data points defined in a research paper or published academic experiment, the credibility of ones experimental methods can be confirmed if a given 'Accepted Value' (assuming there is one) is found to be within the bounds of the calculated STDV. Such results indicate that the experimental procedures were 'correctly' followed as defined by the author of the experiment.

If the Accepted Value is 'outside' of the bounds of the STDV, then the experimental results of the one reproducing the experiment would be questionable, or the research paper is flawed. Most academic experiments are well developed and typically have reliable 'Accepted Values'. Having the Accepted Value within the bounds of the STDV is a credible method of verifying ones experimental methods. Great lecture on this topic. Keep up the good work. :-)

Tools in Quantitative Chemistry

  • Dimensional analysis is the use of conversion factors to convert between different units, including the standard SI units of measurement.
  • Percent error is related to accuracy, while standard deviation is related to precision.
  • Significant figures are related to precision, and must be considered when performing calculations.
  • Significant figures are a relate to a measurement’s uncertainty.

Tools in Quantitative Chemistry

Lecture Slides are screen-captured images of important points in the lecture. Students can download and print out these lecture slide images to do practice problems as well as take notes while watching the lecture.

  • Intro 0:00
  • Lesson Overview 0:07
  • Units of Measurement 1:23
    • The International System of Units (SI): Mass, Length, and Volume
  • Percent Error 2:17
    • Percent Error
    • Example: Calculate the Percent Error
  • Standard Deviation 3:48
    • Standard Deviation Formula
  • Standard Deviation cont'd 4:42
    • Example: Calculate Your Standard Deviation
  • Precisions vs. Accuracy 6:25
    • Precision
    • Accuracy
  • Significant Figures and Uncertainty 7:50
    • Consider the Following (2) Rulers
    • Consider the Following Graduated Cylinder
  • Identifying Significant Figures 12:43
    • The Rules of Sig Figs Overview
    • The Rules for Sig Figs: All Nonzero Digits Are Significant
    • The Rules for Sig Figs: A Zero is Significant When It is In-Between Nonzero Digits
    • The Rules for Sig Figs: A Zero is Significant When at the End of a Decimal Number
    • The Rules for Sig Figs: A Zero is not significant When Starting a Decimal Number
  • Using Sig Figs in Calculations 15:03
    • Using Sig Figs for Multiplication and Division
    • Using Sig Figs for Addition and Subtraction
    • Using Sig Figs for Mixed Operations
  • Dimensional Analysis 16:20
    • Dimensional Analysis Overview
    • General Format for Dimensional Analysis
    • Example: How Many Miles are in 17 Laps?
    • Example: How Many Grams are in 1.22 Pounds?
  • Dimensional Analysis cont'd 19:43
    • Example: How Much is Spent on Diapers in One Week?
  • Dimensional Analysis cont'd 21:03
    • SI Prefixes
  • Dimensional Analysis cont'd 22:03
    • 500 mg → ? kg
    • 34.1 cm → ? um
  • Summary 25:11
  • Sample Problem 1: Dimensional Analysis 26:09

Transcription: Tools in Quantitative Chemistry

Hi, welcome back to Educator.com.0000

Today's lesson in general chemistry is on the tools of quantitative chemistry.0002

We are going to get introduced into how chemists go ahead and make measurements0010

and the basic concepts and methods of performing everyday calculations.0015

Our first point in the lesson is going to be the SI units of measurement.0023

Once we become familiar with the units of measurement,0027

we will then get into explaining how exactly good or bad your measurement was.0030

This is done by something we call percent error and also standard deviation.0039

We will then get into a discussion of what we mean by precision versus accuracy.0044

We will then get into something a little more quantifiable which is called significant figures and uncertainty.0048

Followed by the following: identifying sig figs; then using these significant figures in calculations.0054

We will then get into a very important concept which is fundamental to all of the physical sciences.0063

This is called dimensional analysis which is your use of what we call conversion factors for problem solving.0069

Finally we will wrap up the lesson with a brief summary followed by some sample problems.0078

The SI units of measurement are what we call the international system of units.0085

They are included as part of the measuring system.0090

It is used in all parts of the world except the United States.0093

The basic SI unit for mass of course is going to be the kilogram.0098

For length, it is going to be the meter.0106

For volume, it is going to be the liter.0109

Kg again stands for kilogram; lowercase m is going to be abbreviating meter.0114

Finally capital L is going to be symbolized for liter.0122

Again mass is going to be kilogram; length is going to be meter; volume is going to be liter.0128

When you make a measurement, how good or poor is it?0139

How close was a measurement to an actual or already known value?0143

This is what we call percentage error.0148

Percentage error is equal to the actual value minus the experimental value divided by the actual value.0150

Actual value is the known value.0158

The experimental value is what you actually measure in an experiment; what you measured.0163

Of course we always like to get as close as possible to the actual.0172

To illustrate this very simple equation to use, suppose you weighed an object and recorded a mass of 86.2 grams.0177

The known value is 97.9 grams; calculate the percentage error.0185

Known value is going to be actual.0189

What you recorded, this is our experimental value.0194

The percent error... let's just plug everything into the equation.0202

It is going to be equal to the actual value 97.9 grams minus the experimental value of 86.2 grams.0206

Divided by the actual value which was 97.9 grams.0216

All of this is going be multiplied by 100 to get you your percentage error.0222

The next statistic that we can also calculate is something called standard deviation.0229

Sometimes when you repeat an experiment several times, it is useful to know how successful you were in reproducing your results.0235

If you did for example ten trials of the same measurement,0243

how well were you able to get the same measurement in each of the ten trials?0248

For our purposes we want of course the standard deviation to be as low as possible.0255

Standard deviation has the following equations.0261

It is equal to the square root of all of these terms where x is equal to basically each individual score.0263

X-bar is simply the average of all the x's.0271

N is the number of values or the number of trials.0274

Summation of course means we are going to add up across all the values.0278

Let's go ahead and illustrate the use of this equation.0283

It looks like a lot but it is really not too bad.0287

You weighed an object three separate times with masses of 32.6, 31.8, 34.1 milligrams.0291

Calculate your standard deviation.0297

All I am going to do is set it up for us right now.0300

The standard deviation is going to be equal the square root of everything underneath it; summation x minus x-bar.0302

X is each score; 32 minus 6, minus the average, squared.0311

I am going to put all of this in brackets showing that we are going to add what is next.0321

Plus, 31.8 minus x-bar, squared; then plus 34.1 minus x-bar, squared; then closed brackets.0326

That is what is meant by summation x minus x-bar, squared.0349

Then divided by n minus 1 where n is going to be equal to 3 trials.0353

All of that divided by 3 minus 1.0358

All we need then to finish this problem is x-bar.0362

X-bar is just your average of the three values.0368

That is just going to be 32.6 plus 31.8 plus 34.1.0371

All of that divided by 3, giving you your average score; that is standard deviation.0377

The next item is what we call precision and accuracy.0385

Precision simply represents how well you were able to reproduce your results over several trials.0388

High levels of precision tend to support accountability and reliability.0394

Precision is going to be quantifiable by standard deviation; measured by standard deviation.0399

Again you want as low a deviation as possible; low standard deviation desirable.0410

The word accuracy however represents how well you were able to measure a value0422

in relation to an already known or published value from the literature.0427

Of course this is going to be measured by percent error.0431

Once again we want as low a percentage error as possible; low percent error is also desirable.0443

A lot of people use the terms precision and accuracy interchangeably.0454

But of course as you have just seen they are not the same.0457

One represents your ability to reproduce a certain measurement.0461

The second represents how well you were able to get to a known value.0466

The next item is very important whenever you make measurements.0473

This is called significant figures and uncertainty.0476

Consider the following two rulers; let me go ahead and draw one ruler.0480

I am going to go ahead and draw a second ruler on the bottom.0486

Now the difference is going to be their tick marks; one, two, three.0489

Then the next one is going to be the following; one and two.0495

Let's go ahead and measure a piece of wood.0510

That is this length right here measured by green.0514

I am going to do the same thing with the other ruler; just like that.0517

How would you measure the green stick with the above ruler?0529

One person could say maybe 1.4 cm; or another person could say 1.5 cm.0533

It turns out that neither person is incorrect.0541

Why?--because the 1 is what we know for sure.0545

We know that green object is longer than 1 for sure; 1 cm.0549

Really this last digit here, the .4 and the .5, is really uncertain.0554

It is the digit of uncertainty.0560

The digit of uncertainty is strictly up to you the user; up to person making the measurement.0566

Because the ruler is so poor, the tick marks are so large,0578

there is no way we can determine if it is .4 or .5.0584

It is not provided by the tick marks on the ruler at all.0587

However when we go ahead and look at this ruler down here,0593

we finally get the tick marks that is representative of the appropriate digit.0596

So you hear, one person could say 1.4.0601

Because it looks like the green mark is right on the .4, one person could even say 1.40 cm.0606

The other person could say 1.41 cm; one could say 1.42cm.0613

Someone could even say it is a little less; 1.39 cm.0619

It turns out that because of the tick marks, we can go one digit more.0623

We can provide a better measurement; better measurement due to what we call a higher level of precision.0629

Again it is a better measurement due to higher level of precision.0651

The rule of thumb is the following.0656

How far or how many digits do you know how to record a measurement?0658

You are always going to go one digit past whatever is given to you or whatever the limit is on the ruler.0663

Go one digit past what is given by the ruler or instrument.0671

For example when we go ahead and read a graduated cylinder, you always want to look at the bottom of the curve.0690

You see how this water level is slightly curved.0698

That is what you call the meniscus.0701

You always want to look at the bottom of it.0705

What is provided to us here on this graduated cylinder?0707

50 milliliters is right here; 55 milliliters is right there.0712

We know that the meniscus is approximately 53 milliliters.0726

The graduated cylinder can tell us if it is 51, 52, or 53.0735

What the rule is is we are going to go one digit past this.0741

We can say something like 53.0; we can say something like 53.1.0745

We could even go under and say 52.9 milliliters or even 52.8 milliliters.0750

Once again you are going to one digit past the last digit of certainty provided by the instrument.0757

When you make a measurement, it is important to always write it down with the correct number of these significant figures.0765

What are their significant figures?--what are the rules for identifying them?0774

The rules for sig figs are the following; all nonzero digits are significant.0779

A zero is significant when it is in between nonzero digits.0786

A zero is significant when at the end of a decimal number.0791

A zero is not significant when starting a decimal number.0795

Let's go take a look at a couple examples; all nonzero digits are significant.0799

For example 562, we have a grand total of three sig figs because none of them are zeros.0804

A zero is significant when it is in between nonzero digits.0810

Something like 501; this is three sig figs.0814

The zero counts because it is in between nonzero digits.0818

5001; both zeros count because they are in between nonzero digits.0822

How about 50010?--only the two zeros count here because they are in between nonzero digits.0829

This last zero here does not count; we only have four sig figs here.0837

A zero is significant when at the end of a decimal number; for example 500. and 500.0.0843

It turns out that in 500. all of these are significant.0853

The zeros come at the end of a decimal number.0858

Here, 500.0, these are all significant because they come at the end of a decimal number.0860

A zero is not significant when starting a decimal number.0870

For example 0.00321, the two zeros do not count here because they start a decimal.0873

We only have three significant figures here.0882

However 0.003210, these do not count at all.0884

However you see the zero here, that comes at the end of a decimal number.0891

This definitely does count; we have a grand total of four significant figures.0895

When we use significant figures in calculations, we have to learn how to incorporate the rules now.0906

For multiplication and division, the answer will have the same number of sig figs as the fewest number of sig figs present.0913

For addition and subtraction, the answer will have the same number of decimal places as the fewest number of decimal places present.0921

For example. 0.321 times 0.57, we are going to get an answer that is only two sig figs.0929

Why?--because here in 0.321, you have three sig figs.0941

Here in 0.57, you have two sig figs.0946

When we do the addition and subtraction with the same numbers, 0.321 minus 0.57,0949

you see now we go by digits after the decimal places.0956

Here there is three digits after the decimal; here there is two.0960

Our answer is going to have two digits after the decimal.0963

Finally when you have mixed operations, you never want to round until the end.0971

You want to carry all digits through.0976

Now that we have talked about significant figures, let's go ahead and discuss what we mean by dimensional analysis.0982

Dimensional analysis utilizes the following.0989

It utilizes ratios of different units that we call conversion factors to convert from one unit to another.0991

If you want to go for example from unit A to unit B, how do we go ahead and do that?0999

The general format is the following; we are going to take unit A.1007

We are going to multiply by this ratio, something over something.1011

That is going to go ahead and give me unit B.1015

Unit A goes downstairs to get cancelled.1018

Unit B goes upstairs to get carried through to the final answer.1022

This ratio right here of unit B to unit A, that is what we call your conversion factor.1027

Let's go ahead and look at a couple of examples.1038

There are 4 laps in 1 mile; how many miles are in 17 laps?1040

We are going to say 17 laps times something over something.1045

That is going to give us our answer in units of miles which is what the question is asking for.1050

I start with laps; but I want to cancel it.1056

It is going to go downstairs to get cancelled.1059

Miles goes upstairs to get carried through to the final answer.1061

The conversion factor is actually given to us in the problem because they state that 4 laps is equal to 1 mile.1065

This is then 1 mile on top and 4 laps on the bottom.1071

We are going to get 4.25 miles for your answer.1075

This is going to round to 4.3 miles.1081

Let me tell you why: 17 laps, you have two sig figs.1085

However for conversion factors, we are assuming that a conversion factor are what we call exact numbers.1090

That is they have infinite precision.1097

They have infinite significant figures; there is no uncertainty.1100

You are going to ignore conversion factors for sig fig purposes.1104

Once again you are going to ignore conversion factors for sig fig purposes.1113

Let's go ahead and look at one last example here on dimensional analysis.1121

One pound is 454 grams; how many grams are in 1.22 pounds?1125

1.22 pounds times something over something is going to give us our answer in grams.1131

I want to cancel the pounds; that goes downstairs.1141

I want to keep grams; that is going to go upstairs.1144

You are told that 1 pound is 454 grams; 1 pound on the bottom and 454 grams on top.1147

That is going to give us an answer of 553.88 grams.1155

Here we have three sig figs; our answer is going to round to 554 grams.1162

That again, that is what we call dimensional analysis.1170

It is important to really master this because we are going to be using this incredibly heavily throughout all our lessons in general chemistry.1173

One last example then; suppose a diaper cost us 35 cents.1185

You have a newborn who goes through about 14 diapers a day.1189

How much is spent on diapers in one week?1192

Suppose you have a diaper costing 25 cents; that is actually a statement already.1196

We are told that 25 cents costs us each diaper; 25 cents per diaper.1203

We are trying to multiply through and cancel units.1211

The other item we see here that has diapers is 14 diapers a day; that is another ratio.1214

14 diapers goes upstairs to get cancelled; then 1 day on the bottom.1220

Finally I want to get my answer into week or dollars per week.1226

This is going to now be 7 days on top to get cancelled, divided by 1 week on the bottom.1233

When all is said and done, you are going to get an answer of1241

24 dollars and 50 cents per week to be spent on diapers.1243

As you can see, the reason why we are doing these examples1250

that are not chemistry yet is to show you that dimensional analysis1253

which we use in chemistry and the physical sciences can actually be very easily applied to everyday life.1256

The next type of dimensional analysis deals with unit conversion.1266

You have heard of terms like centi and milli and kilo before.1271

But how do you convert between the three?1276

We are going to learn that dimensional analysis is all behind this; converting between units.1278

You should ask your chemistry instructor which units you actually have to know.1284

But let me go ahead and just point out a few.1288

Mega is 106; kilo is 103; deci is 10-1; centi is 10-2.1291

Milli means 10-3; micro is 10-6; nano is 10-9.1301

Once again please ask your instructor if you have to memorize any of these1309

if at all or if they are going to be given to you.1313

Now that we have been introduced to these prefixes and what they mean,1317

let's go ahead and see how we can use them in calculations.1322

For example, 500 milligrams is equal to how many kilograms?1326

What I always like to say is that sometimes it helps to convert to the base unit first.1330

What I mean by the base unit is that it has no prefixes.1335

In other words, let's get to the unit with no prefixes first.1341

First step is to go from milligrams to regular grams then onto kilograms.1347

This is going to be a two step process.1355

500 mg, just going to set it up, times something over something.1358

The first step is to get g; g goes upstairs; mg goes downstairs to get cancelled.1364

Once I am in g, now I can go to kg; times something over something.1369

That is going to give us our answer in units of kg.1375

You see that g is upstairs here; it is going to go downstairs to get cancelled.1378

Kg goes upstairs to get carried through to the final answer.1383

What numbers do we put in and where?--the rule is the following.1388

For the prefix multipliers like kilo and milli, you put the multiplier with the base unit; put multiplier with the base unit.1393

Once again you should put the multiplier with the base unit.1409

For milli, milli stands for 10-3.1414

That is going to go with the base unit here; 10-3 on top; 1 on the bottom.1419

Kilo stands for 103; that is going to go with the base unit.1424

1 goes to kg; 103 is going to go the g.1428

Once again you always put the multiplier with the base unit.1433

Then you get your answer in kilograms.1441

Let's go ahead and do a last one; this is centimeters to micrometers.1443

The first step is to go from centimeters to regular meters; then from regular meters on to micrometers.1447

34.1 centimeters times something over something is going to give me my answer in meters.1457

Cm goes downstairs to get cancelled.1465

M goes on top to get carried through to the final answer.1467

Once I am in meters, I can then go on and get micrometers; times something over something.1471

That is going to give me my answer in micrometers.1476

You see that m is on top; it has to go downstairs to get cancelled.1479

Then micrometers goes upstairs to get carried through to the final answer.1483

When we look up the prefix for centi, centi stands for 10-2.1488

That goes with the prefix-less unit; 10-2 on top; 1 on the bottom.1492

When we look up the multiplier for micro, it is 10-6.1497

That goes downstairs with meters; 1 on top.1503

That is going to get you your answer in units of micrometers.1507

To summarize, when you perform measurements, you want to gauge how well you are doing.1514

We can do this by percentage error calculation which again is going to tell us a little about your accuracy.1519

We can also calculate what is called the standard deviation.1527

That is going to tell us of how precise you were.1531

We also learned the concept of significant figures; significant figures is related to precision.1535

Finally we learned a very fundamental concept in all of the physical sciences which is dimensional analysis1545

which follows the same basic pattern where we can go from unit A to unit B using a conversion factor.1551

There is our summary; now let's go ahead and tackle some sample problems.1567

An intramuscular medication is given at a dosage of 5.00 milligrams per kilogram of body weight.1571

If you give 0.425 grams of medication to a patient, what is the patient's weight in pounds?1584

0.425 grams of medication is here.1592

Somehow we want to go from grams of medication to whatever the question is asking for which is pounds of body weight.1597

In addition we see in the first sentence that we have a statement here1611

that 5 milligrams of medication are given per kilogram of body weight.1615

That represents a very nice ratio.1619

5.00 mg of medication for every kilogram kg of body weight; that is our conversion factor.1622

This unit is 5 milligrams of medication; we are given 2.45 grams of medication.1637

We have to get the 0.425 grams of medication into milligrams first.1643

0.425 grams of medication times something over something is going to give us our answer in units of milligrams of medication.1648

Milligrams goes upstairs; g goes downstairs.1659

When we look up the prefix for milli, it is 10-3.1662

That goes with the prefix-less unit on the bottom; then 1 on top.1666

That is going to give us milligrams of medication which is going to be 425.1670

We can then take the 425 milligrams of medication, multiply it by something over something.1677

That is going to give us our answer in kilograms of body weight.1684

Mg is going to go on the bottom; kg is going to go on top.1698

We know that from the sentence, it is 1 kg for every 5.00 mg.1701

Finally we can then take our kilogram of body weight and we can go to pounds of body weight.1708

We do that from the conversion factor where 1 kilogram is equal to approximately 2.20 pounds.1717

We are going to take kilograms of weight, multiply it by something over something.1726

That is going to give us our answer in pounds.1732

Pounds goes upstairs; kg goes downstairs; just 2.20 divided by 1.1736

When all is said and done, we should get an answer of 187 pounds using the correct number of sig figs.1743

That is our lesson from general chemistry on quantitative tools.1752

I want to thank you for your attention.1758

I will see you next time on Educator.com.1760

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