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Franklin Ow

Franklin Ow

Acid-Base 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 (16)

0 answers

Post by Parsa Abadi on May 27, 2017

Hi there,
Quantifying acidity and basicity, what do you mean when you say at equilibrium?

1 answer

Last reply by: Professor Franklin Ow
Wed May 7, 2014 6:44 PM

Post by Edgar Suarez on May 5, 2014

I have a question, you said that milk was basic, but i thought that it was an acid, because of its lactic acid.

1 answer

Last reply by: Professor Franklin Ow
Sat Apr 26, 2014 5:17 PM

Post by Anthony Mendoza on April 25, 2014

In example 2, how did you know that HPO4 was the conjugate of weak acid H3PO4? In my head, I think I might have chosen that it was the conjugate of -PO4, or conjugate of H2PO4...

1 answer

Last reply by: Professor Franklin Ow
Sat Apr 26, 2014 5:16 PM

Post by Anthony Mendoza on April 25, 2014

Hi Professor Ow
,
Could I get some clarification on how increased electronegativity makes it a stronger acid? HF has a strong electronegativity difference, as in the F- pulls strongly on the H+, so I would imagine it would't want to let it go (and thus dissociate and thus not be that strong).

I'm trying to compare HF to the others on the list, but it's difficult because aren't the rest bases. They already have a negative charge, so obviously aren't going to want to dissociate. But how about comparing HF to HCl. HCl is a strong acid, yet has lower ElectroNegativity (since F- is supposed to be the most EN?). This part still confuses me.

Thanks for your help! Your explanations are making class a lot more manageable!

1 answer

Last reply by: Professor Franklin Ow
Sat Apr 26, 2014 5:15 PM

Post by Melyssa James on April 2, 2014

For your first example, after you determined the pH, i don't see where you calculated the percent ionization??

2 answers

Last reply by: Professor Franklin Ow
Fri Feb 28, 2014 11:40 AM

Post by Jack Miars on February 26, 2014

Can you always make the assumption that x is negligible? If not, what situations cause it to be accounted for?

1 answer

Last reply by: Professor Franklin Ow
Mon Feb 10, 2014 10:58 AM

Post by Miley Xiao on February 9, 2014

what does x at 13:38 stands for? and why are CH3CO2 and H3O the same since he wrote two xs?

0 answers

Post by Cheng Jiang on May 16, 2013

very gucci

0 answers

Post by Max Mayo on April 2, 2013

great job

Related Articles:

Acid-Base Chemistry

  • Bronsted-Lowry acid-base chemistry involves a loss/gain of a proton to/from water.
  • Conjugate pairs only differ by one proton and are inversely related in terms of acidity/basicity.
  • Acidity and basicity can be quantified primarily via pH, Ka ,pKa and percent ionization.
  • The structure of a molecule can have a significant impact on how acidic it can be.
  • ICE tables are easily applied to an acid-base equilibrium situation.

Acid-Base 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 0:06
  • Introduction 0:55
    • Bronsted-Lowry Acid & Bronsted -Lowry Base
    • Water is an Amphiprotic Molecule
    • Water Reacting With Itself
  • Introduction cont'd 4:04
    • Strong Acids
    • Strong Bases
  • Introduction cont'd 6:16
    • Weak Acids and Bases
  • Quantifying Acid-Base Strength 7:35
    • The pH Scale
  • Quantifying Acid-Base Strength cont'd 9:55
    • The Acid-ionization Constant Ka and pKa
  • Quantifying Acid-Base Strength cont'd 12:13
    • Example: Calculate the pH of a 1.2M Solution of Acetic Acid
  • Quantifying Acid-Base Strength 15:06
    • Calculating the pH of Weak Base Solutions
  • Writing Out Acid-Base Equilibria 17:45
    • Writing Out Acid-Base Equilibria
  • Writing Out Acid-Base Equilibria cont'd 19:47
    • Consider the Following Equilibrium
    • Conjugate Base and Conjugate Acid
  • Salts Solutions 22:00
    • Salts That Produce Acidic Aqueous Solutions
    • Salts That Produce Basic Aqueous Solutions
    • Neutral Salt Solutions
  • Diprotic and Polyprotic Acids 24:44
    • Example: Calculate the pH of a 1.2 M Solution of H₂SO₃
  • Diprotic and Polyprotic Acids cont'd 27:18
    • Calculate the pH of a 1.2 M Solution of Na₂SO₃
  • Lewis Acids and Bases 29:13
    • Lewis Acids
    • Lewis Bases
    • Example: Lewis Acids and Bases
  • Molecular Structure and Acidity 32:03
    • The Effect of Charge
    • Within a Period/Row
  • Molecular Structure and Acidity cont'd 34:17
    • Within a Group/Column
    • Oxoacids
  • Molecular Structure and Acidity cont'd 37:54
    • Carboxylic Acids
    • Hydrated Metal Cations
  • 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

Transcription: Acid-Base Chemistry

Hi, welcome back to Educator.com.0000

Today's lesson from general chemistry is on acid base chemistry.0002

We are going to first start off as usual with our introduction followed by how to quantify acid and base strength.0007

After we learn how to quantify how acidic or how basic something is, we will then go on to0015

writing out aqueous acid base equilibria reactions which is going to become very fundamental for this section.0022

After that, we will learn something we call acidic and basic salts0028

followed by multiple aqueous equilibria which is for acids that have more than one hydrogen.0031

We will then get into some specific topics, namely Lewis acids and bases, followed by molecular structure and acidity.0038

We will wrap up the session with our summary followed by a pair of sample problems.0047

Let's go ahead and begin.0054

We are going to look at two different types of molecules in this chapter--what we call a Bronsted-Lowry acid and a Bronsted-Lowry base.0057

A Bronsted-Lowry acid basically donates a proton to water.0065

A Bronsted-Lowry base is going to accept a proton from water.0076

You see that water is involved in both of these reactions.0088

For a Bronsted-Lowry acid, you can say HA aqueous plus water goes on to form A- aqueous because it loses its proton.0092

It is going to give it to water; H2O then becomes H3O1+ aqueous.0105

This special ion is what we call the hydronium ion.0113

Anytime you see the hydronium cation, it is always indicative of acid.0119

Let's go ahead and look at the reaction for a base.0125

We can have a generic base B aqueous plus H2O liquid goes on to form...0127

The base is going to accept the proton; it becomes BH1+ aqueous.0136

H2O loses a proton; it is going to become OH1- aqueous.0141

This special guy here of course is what we call the hydroxide ion.0145

This is always indicative of aqueous base.0150

Once again hydronium cation is indicative of aqueous acid.0153

Hydroxide anion is indicative of aqueous base.0157

You see that water is involved in both of the reactions.0161

In the acid reaction, water is the base.0166

In the base reaction, water is the acid.0169

We say that water is an amphiprotic molecule meaning it can function both as an acid and as a base.0172

When we take a look at water reacting with itself, we see the dual role of water.0179

For example, one of these water molecules can lose a proton giving hydroxide.0186

Therefore the other water molecule is going to gain a proton forming hydronium.0191

What is the extent of this reaction?0198

It turns out that if we write out the equilibrium expression for this reaction, it is going to be0202

equal to the hydroxide ion concentration at equilibrium times the hydronium ion concentration at equilibrium.0207

It turns out that this value, this product is only 1.0 times 10-14 at 25 degrees Celsius.0216

We give this K a very special name; this is what we call Kw.0225

This is called the auto-dissociation constant for water.0230

It is a very handy one because if you know the hydroxide ion concentration, you can calculate hydronium and vice versa.0236

When we talk about acids and bases, there are going to be what we call strong and weak acids and bases.0247

Basically strong acids are going to donate a proton to water and fully dissociate into a proton and an anion.0253

Because they fully dissociate, we do not use an equilibrium arrow.0262

We only use an arrow pointing in the forward direction.0267

We can take a molecule HX aqueous; this is going to react with water.0273

We only use a single arrow to show full dissociation.0278

That is going to form H3O1+ aqueous and X1- aqueous.0281

It is very imperative that you know what are the strong acids.0287

For our purposes right now... you should always confirm this with your instructor.0292

But there are going to be seven strong acids that you should know.0295

HClO4, HClO3, HNO3, H2SO4, HCl, HBr, and HI.0299

Again these are the seven strong acids that fully dissociate when dissolved in water.0313

The strong bases also are going to become fully protonated.0319

They are going to fully dissociate when dissolved in water.0324

They are going to form a cation and a hydroxide.0328

The typical ones are going to be MOH or MOH2.0332

These are basically group 1 and 2 hydroxides.0340

We can take the typical one, sodium hydroxide.0349

You don't even have to show water because all it does is that0353

it is going to fully break apart into sodium ions and hydroxide anion.0356

Once again these are the strong bases and the strong acids that you should know.0362

Again check with your instructor because every instructor is going to be0368

slightly different on the ones he or she requires you to memorize.0372

We have talked about strong species, what are the weak species.0378

Basically weak species, weak acid and bases do not completely dissociate at all, do not completely dissociate at all.0381

For example, HF aqueous plus water is a weak acid.0389

We are going to use of course an equilibrium arrow just like we do for any type of weak electrolyte.0396

That is going to give us hydronium cation plus F1- aqueous.0403

That is the typical weak acid.0410

The typical weak base is going to going to be ammonia NH3 aqueous plus H2O liquid.0412

Once again equilibrium arrow.0420

That goes on to form NH41+ aqueous and hydroxide aqueous.0421

Anything related to ammonia which is what we call... this is ammonia/amines.0429

These are all going to be our typical weak bases.0439

Again these do not completely dissociate.0449

Let's go ahead and now see how to quantify acidity and basicity.0452

To quantify acid and base strength, we talk about what is known as the pH scale.0457

When we talk about the pH scale, this is on a scale of 0 to 14 where 7 is the neutral area.0464

Any pH less than 7 is acidic.0476

Any pH greater than 7 is going to be basic.0480

The equations for pH is the following.0483

pH is equal to ?log of concentration of hydronium at equilibrium.0485

pOH is equal to ?log of the concentration of hydroxide at equilibrium.0492

There is some common things that you can know that are acidic--citric juice, sodas, stomach acid.0501

Then there is some things that are basic that you should know also.0513

Any household detergents tend to be basic.0516

Milk and human blood also is slightly basic.0519

You have to also make sure that you be able to go both ways.0528

Sometimes you are asked to calculate the hydronium ion concentration at equilibrium.0532

That is going to be anti-log.0537

That is going to be log inverse of ?pH which is 10 raised to the ?pH.0538

The hydroxide ion concentration at equilibrium is equal to log inverse of ?pOH which is equal to 10 raised to the ?pOH.0548

There is going to be a relationship between pH and the pOH.0563

Basically the pH plus the pOH is going to be equal to 14.00 at 25 degrees Celsius.0566

Again you should always consult with your instructor whether or not you have to commit these to memory.0577

This is the pH scale that is typically used for quantifying acid and base strength.0589

Let's now move on to another way of quantifying acid base strength.0596

This is going to be something we are going to encounter quite a bit.0599

This is what we call the acid ionization constant Ka and pKa.0603

For any of the acids where HA reacts with water, you are going to form hydronium and A1-.0608

There is always going to be a certain extent to which this acid is going to dissociate.0624

We are going to call this a Ka.0633

This is just the same equilibrium expression that we have been working with so much already.0636

Remember it is going to be products, hydronium ion, at equilibrium times0640

A1- at equilibrium divided by the concentration of HA at equilibrium.0645

You can see that for stronger acids, the concentration of hydronium at equilibrium is going to be high.0652

Ka is going to be high also.0665

This is what we call the acid ionization constant Ka.0673

Ka is usually written in scientific notation.0678

But there is another way we can go around that, something more convenient system to use.0682

This is called pKa; pKa is just like the pH.0687

pKa is equal to ?log of Ka.0691

You see that for large Ka values, pKa is actually small.0694

Again this is going to be true for stronger acids.0703

The nice thing about pKa is that pKa, the advantage is that we don't have to use scientific notation when we compute it.0707

It is going to be a nice small number.0716

pKa advantage avoids use of scientific notation.0718

That is the acid ionization constant Ka and pKa.0730

Let's go ahead and do a sample problem now.0734

Calculate the pH of a 1.2 molar solution of acetic acid where Ka is 1.8 times 10-5.0736

Also calculate the percent ionization.0741

Anytime you have a calculation like this, this always involves the use...0746

This is always going to involve the use of ICE tables just like we learned last time.0756

Let's go ahead and write out the equation.0766

Acetic acid CH3CO2H plus H2O liquid goes on to form CH3CO21- aqueous plus hydronium aqueous.0768

When we are reading the statement, calculate the pH of a 1.2 molar solution of acetic acid,0786

that 1.2 molar, remember the problem doesn't mention equilibrium here for 1.2 molar.0790

We assume this is initial and then zero and zero.0796

This is going to be ?x, +x, and +x.0800

This is then 0.2 minus x, x, and x.0804

We are going to set it up then to the expression.0808

Ka is equal to 1.8 times 10-5 which is equal to x squared over 1.2 minus x.0810

You remember from our last session that we talked about the assumption that x is negligible.0819

We can do this for any weak acid.0827

Assume x is negligible for weak acids and bases.0829

This equation then is approximately x squared over 1.2.0840

When we solve for x, that is going to give us the hydronium ion concentration at equilibrium which is going to be 0.0046 molar.0846

When we go ahead and solve for pH, as expected, we should get an acidic pH because this is an acid after all.0857

This is going to be pH is equal to 2.33.0864

The rules for sig figs when you do the logarithm function, the number of sig figs0867

in the concentration is equal to the number of digits after the decimal in pH.0877

As you can here for 0.0046, there is two significant figures.0894

I am going to report my pH value to two digits after the decimal.0899

That is how we quantify acid base strength.0905

Another way we can quantify acid base strength is to do the other one besides acids--that is bases.0908

Instead of using Ka, we use Kb which is the base ionization constant.0917

For Kb, we are going to say B aqueous plus H2O liquid goes on to form BH+1 aqueous and hydroxide aqueous.0922

Basically Kb is going to be equal to the concentration of BH1+ at equilibrium0936

times hydroxide at equilibrium divided by the concentration of B at equilibrium.0944

You can see it is analogous to Ka.0951

Stronger bases are going to have lots of OH- at equilibrium which means Kb is going to be large.0955

Just like pKa, pKb is going to be equal to the negative log of Kb.0969

If Kb is large, this means pKb is small.0976

Once again for stronger bases, pKb is small.0984

That is how we do the base ionization constant.0998

Finally the other ways to quantify acid base strength is what we call percent ionization.1003

For acids, this is going to be equal to the hydronium ion concentration1014

at equilibrium divided by the initial concentration of the acid times 100.1022

Then for bases, this is equal to the concentration of BH1+ at equilibrium divided by the initial concentration of base times 100.1027

Basically strong acids are going to have relatively large percent ionizations near the 99 percent.1039

For weak acids and bases, it is going to be the extreme.1046

The percent ionization is actually going to be single digits, sometimes less than1050

one percent just to show you how drastic the differences can be.1054

Now that we have gone through quantifying acid base strength,1059

let's go ahead and turn our attention to writing out acid base equilibria.1062

As you can see, in order to set up the ICE table, it is very imperative that you be1066

able to write out the correct equilibria, or else the entire problem will be messed up.1070

Let's go ahead and take some practice in writing out equilibria reactions involving acids and bases.1075

For example, HNO2 aqueous, this is going to react with water liquid.1080

Because this is weak, we use an equilibrium arrow.1088

HNO2 is going to lose a proton giving nothing but NO21- aqueous.1093

H2O is going to gain it; that becomes H3O1+ aqueous.1097

Let's go ahead and look at a base, CH3NH2 aqueous plus H2O liquid.1103

This is going to be weak base; we use an equilibrium arrow.1111

The amine is going to gain a proton, CH3NH31+ aqueous.1114

H2O loses the hydrogen giving us hydroxide.1121

A couple of notes anytime you are dealing with an amine.1127

Four amines, you always attach the H to the nitrogen atom.1132

That is note number one.1141

Note number two, you see that when an acid loses a proton, you see that the charge always decreases.1143

HNO2 becomes NO21-; for each hydrogen loss, charge decreases by one.1150

You see for the base of course, for each hydrogen gained, the charge is going to increase by one.1162

That is why we go from CH3NH20 to CH3NH31+.1172

This leads us into what we call conjugate pairs.1181

This is going to become very important conceptually for the rest of our discussion.1184

Consider the following equilibrium.1188

HF aqueous plus H2O liquid goes on to form H3O1+ aqueous and F- aqueous.1192

The only difference between HF and F- is a proton.1207

The only difference between H2O and H3O1+ is a proton.1212

When you add two molecules on opposite sides of the equilibrium that1216

differ by only one hydrogen, these are what we call conjugate pairs.1222

HF and F- are conjugates.1226

Water and hydronium are also a conjugate pair.1229

HF is the Bronsted-Lowry acid, making F- its counterpart the conjugate base.1234

The reason why it is called a conjugate base is because there is a tendency for the reverse reaction to happen1244

where F- aqueous can react with water going on to form HF aqueous and hydroxide.1252

This equation is likely to happen because HF is only a weak acid.1260

What that means is that it doesn't fully dissociate.1264

That means that F- will react with water to reform at least some of the initial HF as we can see from this reaction.1267

In general, the stronger the acid, the weaker its conjugate base.1279

Something like HCl which fully dissociates means that Cl1- will not react with water1284

to reform HCl because HCl has a great tendency to remain dissociated in water.1290

In general, the stronger an acid, the weaker its conjugate base.1297

In general, the stronger a base, the weaker its conjugate acid.1300

What we see is that acid base conjugate strengths are going to be inversely related.1304

Now let's go ahead and examine salt solutions and predict if they are acidic, basic, or neutral.1314

Basically you want to go by the following three rules to help us predict1322

if a salt is acidic, basic, or neutral just going off of its formula.1326

Salts said to produce acidic aqueous solutions contain the following combinations.1330

A conjugate acid of a weak base and a conjugate base of a strong acid--something like NH4Cl aqueous.1335

Cl1- is not going to react with water to give us anything because Cl- is the conjugate of HCl.1349

We are not going to reform HCl.1357

However NH41+ will have a tendency to react with water because it is the conjugate of only a weak base.1359

Here we are going to NH3 aqueous and hydroxide aqueous.1368

Excuse me... NH3 aqueous plus hydronium aqueous; there we go.1375

You see that we form hydronium which makes sense that we expect this solution1381

therefore to be acidic when the salt is dissolved in water.1386

Let's look at basic aqueous solutions.1393

Salts that produce aqueous solutions that are basic contain a group 1 and group 21395

metal cation like Na in combination with the conjugate base of a weak acid like F.1400

Na+ is not going to react with water to form anything.1409

It is going to remain solvated; we don't have to worry about that.1417

Also F-, F- is the conjugate of HF, a weak acid, which means there will be1422

a tendency for the following reaction to occur, formation of hydroxide and HF aqueous.1429

You see that because we form hydroxide, these types of salts are expected to be basic when dissolved in water.1437

Let's now move onto neutral salt solutions containing a group 1 and 2 metal cation and the conjugate base of a strong acid.1445

For example, NaCl, we said that Na1+, this is not going to reform hydroxide when reacting with water.1454

Cl1- is not going to react with water to reform HCl.1464

In either case, both the cation and anion do not react with water to form hydroxide or hydronium.1469

Neither is going to contribute to pH.1476

This is what we call a neutral salt solution.1479

We have so far talked about monoprotic acids, those acids that only contain one hydrogen atom.1485

But what happened for acids like sulfuric acid or phosphoric acid?1491

These are what we call polyprotic acids.1496

To calculate the pH of a 1.2 molar solution of sulfuric acid which is diprotic, we want to remember one thing.1499

It is that the first dissociation is the only one contributing to pH; first dissociation only contributes to pH.1507

Once again the first dissociation only contributes to pH1521

meaning that Ka1 is going to be relatively much greater than Ka21525

which is going to be relatively much greater than Ka3.1530

For H2SO3 aqueous plus water, we are going to write out the dissociation stepwise.1534

That is very important; write out dissociation stepwise; write out steps one at a time.1545

H2SO3 aqueous plus H2O liquid goes on to form HSO31- aqueous and H3O1+ aqueous.1558

This is what we call Ka1.1569

The next one is HSO31- aqueous is going to take its turn.1572

Plus H2O liquid goes on to form H3O1+ aqueous and SO32- aqueous.1577

For this, we are going to see Ka2.1587

We are basically saying Ka1 is much greater than Ka2.1590

To complete calculate the pH, Ka1 is going to be equal to 1.2 times 10-2.1594

You can look that up which is approximately x squared over 1.2 minus x.1604

You can go ahead and use this to solve for the hydronium ion concentration which gets us pH.1610

If you look up Ka2, Ka2 is 6.6 times 10-8.1617

That is more than six orders of magnitude less than Ka1.1622

This shows how drastic a difference the first deprotonation is versus successive deprotonations.1629

Now let's take a look at the salt of a diprotic acid.1638

Calculate the pH of a 1.2 molar solution of Na2SO3.1642

Here Na2SO3, which one of these cation or anion is going to react with water?1647

We know that Na1+ is not going to react with water.1653

There is going to be no reaction; the only thing that is left is sulfite.1659

SO32- aqueous is going to react with water.1664

We just saw that SO32- is the conjugate of a weak acid.1668

Then there is going to be a tendency for this reaction to occur where we reform HSO31- aqueous and generate hydroxide.1674

We don't use Ka for this one; we use Kb.1684

But we don't have Kb; we were only given Ka.1690

But we do know that Kw is equal to Ka times Kb.1693

This is a very important equation which is equal to 1.0 times 10-14.1697

It turns out that Kb for this reaction which we call Kb1, Kb1 represents the first protonation.1704

This is going to be equal to Kw over not Ka1 but Ka2 only.1711

The reason why we use Ka2 is because we don't generate SO32-.1717

It is not formed until both protons have been removed.1726

SO32- not formed until second deprotonation which on the previous slide if you go back is associated with Ka2.1730

Again this is how you deal with a salt solution of a polyprotic acid.1745

Now let's go ahead and take a look at acids from a different perspective.1755

So far we have talked about Bronsted-Lowry acids and bases which involve a transfer of protons.1759

However we can also look at it from a more general approach.1764

This is what we call the Lewis theory of acids and bases.1768

Lewis acids accept a lone pair of electrons to form a new covalent bond.1772

Lewis bases donate a lone pair; Lewis acids accept a lone pair.1778

You want to look for metal cations that tend to be Lewis acids.1785

You want to look for elements that do not have a complete octet,1792

that do not require a complete octet which is boron, aluminum, and beryllium.1796

These do not need a full octet.1802

They have room to accommodate an extra lone pair.1806

Lewis bases however donate a lone pair.1810

You want to look for anions for visual clues.1813

Of course you have to have lone pairs to start with; look for lone pairs in the Lewis structure.1817

How do they relate to Bronsted-Lowry acids and bases?1831

You want to remember that all Bronsted-Lowry acids/bases are also Lewis acids and bases.1834

However the reverse statement is not always true; the converse is not always true.1849

Again these are what we call Lewis acids and Lewis bases.1861

We can take a look at a very interesting example.1865

Let's look at ammonia then, NH3 aqueous.1869

I am going to put its lone pair right there.1872

It is going to react with H2O liquid going on to form NH41+ aqueous and hydroxide aqueous.1874

Let me put the lone pairs on the oxygens just to highlight the point.1885

You see that the nitrogen atom has lost a lone pair making ammonia a Lewis base.1889

You see that the oxygen atom has gained a lone pair making water the Lewis acid.1899

Just like a Bronsted-Lowry acid reacts with a Bronsted-Lowry base,1906

a Lewis acid is also going to react with a Lewis base.1910

Now that we have interpreted acids and bases differently, let's go ahead and look at factors that influence acidity strictly off of molecular structure.1916

Factor number one, the effect of charge, basically acidity is higher with increasing charge.1925

For example, H2SO4 versus HSO41-, this is going to be the stronger acid.1939

Remember we said that for polyprotic acids that successive deprotonations don't really contribute to pH.1950

Successive steps are negligible to pH; look at this.1957

The reason is the following; a hydronium cation is positively charged.1966

It is easier for something with a high charge to give up something positively charged.1971

HSO4 has its -1 charge.1975

It is going to be a relatively harder time to give up that positively charged cation.1977

Once again acidity is higher with increasing charge.1983

Now, within a period/row--CH31- aqueous, NH21- aqueous, OH1- aqueous, and HF aqueous.1988

It turns out that acidity is going to increase left to right.2003

Acidity goes up left to right in a row; the reason is the following.2008

What else increases left to right in a row?--this parallels electronegativity.2018

Remember we say that Bronsted-Lowry acids, which are also Lewis acids, they donate a proton and accept a lone pair.2028

The more electronegative the atom, the more easier it is to accept a lone pair.2039

Higher EN is easier to accept a lone pair.2045

Again acidity increases left to right in a row.2055

Within a group or column is the next one we will look at--HF, HCl, HBr, and HI.2058

It is experimentally determined that acidity increases down a column.2065

This doesn't parallel electronegativity; however this does parallel atomic size of the anion.2073

Parallels atomic size of atom directly bonded to H.2081

Again that is key; it has to be directly bonded to the H.2088

What happens is when HA become A1-, A1- is gaining a lone pair.2095

Let's look at this; it starts off with three lone pairs in HA.2103

It comes with four lone pairs in A-.2107

We are going to be better off adding electron density to a larger volume.2113

Easier to add more electrons to a larger volume.2118

In other words, I- is more stable than F- which makes HI a stronger acid2129

because the dissociation of HI is much more likely to occur and to a greater extent2144

than the dissociation of HF where F- is not as stable because of its smaller size.2151

Oxo acids, for example, HNO3 versus HNO2.2158

In oxo acids, acidity increases with the number of oxygen atoms per hydrogen.2167

To answer the question why, let's go ahead and look at the Lewis structures.2184

HNO3 here; HNO2 is going to be right there.2190

Basically what is happening is we know that HNO3 is a stronger acid.2201

It is one of the seven you have memorized; HNO2 is not.2207

What makes the difference?--the only difference is one oxygen atom.2211

That explains for why HNO3 is more acidic.2214

What happens is the oxygen atoms are highly electronegative.2217

They are going to withdraw electron density away from the bond with hydrogen.2220

That makes this partial positive.2225

That is going to make the rest of the molecule partial negative.2228

This weakens the bond with the hydrogen.2233

It is going to allow for hydrogen to leave the acid much more easily.2236

Here in HNO2, the withdrawing effect is not as great.2242

The electron density is not so lopsided.2250

It is going to be a harder time for hydrogen to fall off here with fewer oxygen atoms.2255

This is what we call an electron withdrawing inductive effect.2260

Once again this is what we call an electron withdrawing inductive effect.2267

That is oxo acids.2273

Carboxylic acids, when we go ahead and look at carboxylic acids, you also go by inductive effect.2277

When you go by the inductive effect, just look for the presence of electronegative groups.2291

For example, we can take this carboxylic acid that has two fluorines versus a carboxylic acid that has no fluorines.2298

They are structurally similar.2307

You see here that with the two fluorines, we are going to get a greater withdrawing effect.2310

Again the stronger the withdrawing effect, the more easily the hydrogen is going to quote and quote fall off.2320

H falls off more easily.2329

Here in acetic acid, there is going to be a minimum withdrawing effect from just the oxygens.2336

Once again for carboxylic acids, you want to go by the inductive effect.2356

You want to look for nearby electronegative atoms.2360

Hydrated metals cations also can be acidic.2364

If we look at Fe(H2O)63+, this is aqueous.2367

This can go ahead and react with water.2375

It can actually function as a Bronsted-Lowry acid.2379

We are going to get Fe Fe(H2O)5OH plus H3O1+ aqueous.2382

How likely is this reaction to occur?2394

This reaction is more likely to occur when this charge is very high.2396

Again for hydrated metal cations, acidity goes up with metal charge.2402

Acidity goes up with metal charge.2412

For example, the Ka of Fe(H2O)63+ is going to be greater2415

than the Ka of Fe(H2O)62+ aqueous just strictly because of charge.2426

That is molecular structure and acidity.2438

Let's go ahead and summarize the section.2441

Bronsted-Lowry acid base chemistry involves a loss or gain of a proton to or from water.2443

Conjugate pairs only differ by one proton and are inversely related in terms of acidity and basicity.2449

We learned many ways of quantifying acid base strength, namely pH, pKa, and Ka and then percent ionization.2456

Finally we saw qualitatively how the structure of a molecule can have a significant impact on how acidic it can be.2466

That is our summary of the lesson.2476

Let's now jump into a pair of sample problems.2478

Calculate the pH of a 1.2 molar solution of NH3 where Kb is 1.8 times 10-5.2481

Just like the previous lecture, the first step for any equilibrium problem is to write out the actual equilibria.2489

NH3 aqueous plus H2O liquid goes on to form NH41+ aqueous and hydroxide aqueous.2496

Let's set up the problem; 1.2 molar is given to us.2510

This is going to be 0 and 0; this is ?x, +x, and +x.2512

This thing goes to 1.2 minus x at equilibrium, x, and x.2519

Kb is 1.8 times 10-5; this is going to be approximately x squared over 1.2.2524

When we go ahead and solve for x, we get the hydroxide ion concentration at equilibrium which is going to be 0.0046 molar.2533

When we solve for pH, we had better get a pH that is basic because this is ammonia after all.2545

We get 11.66; this is sample problem one.2550

Let's now move on to sample problem two--predicting if the following salt solutions are acidic, basic, or neutral.2558

Here potassium bromide, we have a group 1 cation.2563

Br1-, this is the conjugate of HBr which is a strong acid.2569

When we have this combination, we expect this salt solution to be neutral.2576

Sodium, group 1, HPO42-.2581

This is going to be the conjugate of phosphoric acid which is going to be a weak acid.2587

We expect this compound to be basic.2593

Finally lithium cyanide, this is going to be group 1 here.2597

CN is going to be the conjugate of HCN which is considered to be a weak acid.2602

We expect this compound here to also be basic when dissolved in water.2608

That is our lesson on acid base chemistry.2617

I want to thank you for your time.2620

I will see you next time on Educator.com.2621

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