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Raffi Hovasapian

Raffi Hovasapian

Precipitation Reactions

Slide Duration:

Table of Contents

I. Review
Naming Compounds

41m 24s

Intro
0:00
Periodic Table of Elements
0:15
Naming Compounds
3:13
Definition and Examples of Ions
3:14
Ionic (Symbol to Name): NaCl
5:23
Ionic (Name to Symbol): Calcium Oxide
7:58
Ionic - Polyatoms Anions: Examples
12:45
Ionic - Polyatoms Anions (Symbol to Name): KClO
14:50
Ionic - Polyatoms Anions (Name to Symbol): Potassium Phosphate
15:49
Ionic Compounds Involving Transition Metals (Symbol to Name): Co₂(CO₃)₃
20:48
Ionic Compounds Involving Transition Metals (Name to Symbol): Palladium 2 Acetate
22:44
Naming Covalent Compounds (Symbol to Name): CO
26:21
Naming Covalent Compounds (Name to Symbol): Nitrogen Trifluoride
27:34
Naming Covalent Compounds (Name to Symbol): Dichlorine Monoxide
27:57
Naming Acids Introduction
28:11
Naming Acids (Name to Symbol): Chlorous Acid
35:08
% Composition by Mass Example
37:38
Stoichiometry

37m 19s

Intro
0:00
Stoichiometry
0:25
Introduction to Stoichiometry
0:26
Example 1
5:03
Example 2
10:17
Example 3
15:09
Example 4
24:02
Example 5: Questions
28:11
Example 5: Part A - Limiting Reactant
30:30
Example 5: Part B
32:27
Example 5: Part C
35:00
II. Aqueous Reactions & Stoichiometry
Precipitation Reactions

31m 14s

Intro
0:00
Precipitation Reactions
0:53
Dissociation of ionic Compounds
0:54
Solubility Guidelines for ionic Compounds: Soluble Ionic Compounds
8:15
Solubility Guidelines for ionic Compounds: Insoluble ionic Compounds
12:56
Precipitation Reactions
14:08
Example 1: Mixing a Solution of BaCl₂ & K₂SO₄
21:21
Example 2: Mixing a Solution of Mg(NO₃)₂ & KI
26:10
Acid-Base Reactions

43m 21s

Intro
0:00
Acid-Base Reactions
1:00
Introduction to Acid: Monoprotic Acid and Polyprotic Acid
1:01
Introduction to Base
8:28
Neutralization
11:45
Example 1
16:17
Example 2
21:55
Molarity
24:50
Example 3
26:50
Example 4
30:01
Example 4: Limiting Reactant
37:51
Example 4: Reaction Part
40:01
Oxidation Reduction Reactions

47m 58s

Intro
0:00
Oxidation Reduction Reactions
0:26
Oxidation and Reduction Overview
0:27
How Can One Tell Whether Oxidation-Reduction has Taken Place?
7:13
Rules for Assigning Oxidation State: Number 1
11:22
Rules for Assigning Oxidation State: Number 2
12:46
Rules for Assigning Oxidation State: Number 3
13:25
Rules for Assigning Oxidation State: Number 4
14:50
Rules for Assigning Oxidation State: Number 5
15:41
Rules for Assigning Oxidation State: Number 6
17:00
Example 1: Determine the Oxidation State of Sulfur in the Following Compounds
18:20
Activity Series and Reduction Properties
25:32
Activity Series and Reduction Properties
25:33
Example 2: Write the Balance Molecular, Total Ionic, and Net Ionic Equations for Al + HCl
31:37
Example 3
34:25
Example 4
37:55
Stoichiometry Examples

31m 50s

Intro
0:00
Stoichiometry Example 1
0:36
Example 1: Question and Answer
0:37
Stoichiometry Example 2
6:57
Example 2: Questions
6:58
Example 2: Part A Solution
12:16
Example 2: Part B Solution
13:05
Example 2: Part C Solution
14:00
Example 2: Part D Solution
14:38
Stoichiometry Example 3
17:56
Example 3: Questions
17:57
Example 3: Part A Solution
19:51
Example 3: Part B Solution
21:43
Example 3: Part C Solution
26:46
III. Gases
Pressure, Gas Laws, & The Ideal Gas Equation

49m 40s

Intro
0:00
Pressure
0:22
Pressure Overview
0:23
Torricelli: Barometer
4:35
Measuring Gas Pressure in a Container
7:49
Boyle's Law
12:40
Example 1
16:56
Gas Laws
21:18
Gas Laws
21:19
Avogadro's Law
26:16
Example 2
31:47
Ideal Gas Equation
38:20
Standard Temperature and Pressure (STP)
38:21
Example 3
40:43
Partial Pressure, Mol Fraction, & Vapor Pressure

32m

Intro
0:00
Gases
0:27
Gases
0:28
Mole Fractions
5:52
Vapor Pressure
8:22
Example 1
13:25
Example 2
22:45
Kinetic Molecular Theory and Real Gases

31m 58s

Intro
0:00
Kinetic Molecular Theory and Real Gases
0:45
Kinetic Molecular Theory 1
0:46
Kinetic Molecular Theory 2
4:23
Kinetic Molecular Theory 3
5:42
Kinetic Molecular Theory 4
6:27
Equations
7:52
Effusion
11:15
Diffusion
13:30
Example 1
19:54
Example 2
23:23
Example 3
26:45
AP Practice for Gases

25m 34s

Intro
0:00
Example 1
0:34
Example 1
0:35
Example 2
6:15
Example 2: Part A
6:16
Example 2: Part B
8:46
Example 2: Part C
10:30
Example 2: Part D
11:15
Example 2: Part E
12:20
Example 2: Part F
13:22
Example 3
14:45
Example 3
14:46
Example 4
18:16
Example 4
18:17
Example 5
21:04
Example 5
21:05
IV. Thermochemistry
Energy, Heat, and Work

37m 32s

Intro
0:00
Thermochemistry
0:25
Temperature and Heat
0:26
Work
3:07
System, Surroundings, Exothermic Process, and Endothermic Process
8:19
Work & Gas: Expansion and Compression
16:30
Example 1
24:41
Example 2
27:47
Example 3
31:58
Enthalpy & Hess's Law

32m 34s

Intro
0:00
Thermochemistry
1:43
Defining Enthalpy & Hess's Law
1:44
Example 1
6:48
State Function
13:11
Example 2
17:15
Example 3
24:09
Standard Enthalpies of Formation

23m 9s

Intro
0:00
Thermochemistry
1:04
Standard Enthalpy of Formation: Definition & Equation
1:05
∆H of Formation
10:00
Example 1
11:22
Example 2
19:00
Calorimetry

39m 28s

Intro
0:00
Thermochemistry
0:21
Heat Capacity
0:22
Molar Heat Capacity
4:44
Constant Pressure Calorimetry
5:50
Example 1
12:24
Constant Volume Calorimetry
21:54
Example 2
24:40
Example 3
31:03
V. Kinetics
Reaction Rates and Rate Laws

36m 24s

Intro
0:00
Kinetics
2:18
Rate: 2 NO₂ (g) → 2NO (g) + O₂ (g)
2:19
Reaction Rates Graph
7:25
Time Interval & Average Rate
13:13
Instantaneous Rate
15:13
Rate of Reaction is Proportional to Some Power of the Reactant Concentrations
23:49
Example 1
27:19
Method of Initial Rates

30m 48s

Intro
0:00
Kinetics
0:33
Rate
0:34
Idea
2:24
Example 1: NH₄⁺ + NO₂⁻ → NO₂ (g) + 2 H₂O
5:36
Example 2: BrO₃⁻ + 5 Br⁻ + 6 H⁺ → 3 Br₂ + 3 H₂O
19:29
Integrated Rate Law & Reaction Half-Life

32m 17s

Intro
0:00
Kinetics
0:52
Integrated Rate Law
0:53
Example 1
6:26
Example 2
15:19
Half-life of a Reaction
20:40
Example 3: Part A
25:41
Example 3: Part B
28:01
Second Order & Zero-Order Rate Laws

26m 40s

Intro
0:00
Kinetics
0:22
Second Order
0:23
Example 1
6:08
Zero-Order
16:36
Summary for the Kinetics Associated with the Reaction
21:27
Activation Energy & Arrhenius Equation

40m 59s

Intro
0:00
Kinetics
0:53
Rate Constant
0:54
Collision Model
2:45
Activation Energy
5:11
Arrhenius Proposed
9:54
2 Requirements for a Successful Reaction
15:39
Rate Constant
17:53
Arrhenius Equation
19:51
Example 1
25:00
Activation Energy & the Values of K
32:12
Example 2
36:46
AP Practice for Kinetics

29m 8s

Intro
0:00
Kinetics
0:43
Example 1
0:44
Example 2
6:53
Example 3
8:58
Example 4
11:36
Example 5
16:36
Example 6: Part A
21:00
Example 6: Part B
25:09
VI. Equilibrium
Equilibrium, Part 1

46m

Intro
0:00
Equilibrium
1:32
Introduction to Equilibrium
1:33
Equilibrium Rules
14:00
Example 1: Part A
16:46
Example 1: Part B
18:48
Example 1: Part C
22:13
Example 1: Part D
24:55
Example 2: Part A
27:46
Example 2: Part B
31:22
Example 2: Part C
33:00
Reverse a Reaction
36:04
Example 3
37:24
Equilibrium, Part 2

40m 53s

Intro
0:00
Equilibrium
1:31
Equilibriums Involving Gases
1:32
General Equation
10:11
Example 1: Question
11:55
Example 1: Answer
13:43
Example 2: Question
19:08
Example 2: Answer
21:37
Example 3: Question
33:40
Example 3: Answer
35:24
Equilibrium: Reaction Quotient

45m 53s

Intro
0:00
Equilibrium
0:57
Reaction Quotient
0:58
If Q > K
5:37
If Q < K
6:52
If Q = K
7:45
Example 1: Part A
8:24
Example 1: Part B
13:11
Example 2: Question
20:04
Example 2: Answer
22:15
Example 3: Question
30:54
Example 3: Answer
32:52
Steps in Solving Equilibrium Problems
42:40
Equilibrium: Examples

31m 51s

Intro
0:00
Equilibrium
1:09
Example 1: Question
1:10
Example 1: Answer
4:15
Example 2: Question
13:04
Example 2: Answer
15:20
Example 3: Question
25:03
Example 3: Answer
26:32
Le Chatelier's principle & Equilibrium

40m 52s

Intro
0:00
Le Chatelier
1:05
Le Chatelier Principle
1:06
Concentration: Add 'x'
5:25
Concentration: Subtract 'x'
7:50
Example 1
9:44
Change in Pressure
12:53
Example 2
20:40
Temperature: Exothermic and Endothermic
24:33
Example 3
29:55
Example 4
35:30
VII. Acids & Bases
Acids and Bases

50m 11s

Intro
0:00
Acids and Bases
1:14
Bronsted-Lowry Acid-Base Model
1:28
Reaction of an Acid with Water
4:36
Acid Dissociation
10:51
Acid Strength
13:48
Example 1
21:22
Water as an Acid & a Base
25:25
Example 2: Part A
32:30
Example 2: Part B
34:47
Example 3: Part A
35:58
Example 3: Part B
39:33
pH Scale
41:12
Example 4
43:56
pH of Weak Acid Solutions

43m 52s

Intro
0:00
pH of Weak Acid Solutions
1:12
pH of Weak Acid Solutions
1:13
Example 1
6:26
Example 2
14:25
Example 3
24:23
Example 4
30:38
Percent Dissociation: Strong & Weak Bases

43m 4s

Intro
0:00
Bases
0:33
Percent Dissociation: Strong & Weak Bases
0:45
Example 1
6:23
Strong Base Dissociation
11:24
Example 2
13:02
Weak Acid and General Reaction
17:38
Example: NaOH → Na⁺ + OH⁻
20:30
Strong Base and Weak Base
23:49
Example 4
24:54
Example 5
33:51
Polyprotic Acids

35m 34s

Intro
0:00
Polyprotic Acids
1:04
Acids Dissociation
1:05
Example 1
4:51
Example 2
17:30
Example 3
31:11
Salts and Their Acid-Base Properties

41m 14s

Intro
0:00
Salts and Their Acid-Base Properties
0:11
Salts and Their Acid-Base Properties
0:15
Example 1
7:58
Example 2
14:00
Metal Ion and Acidic Solution
22:00
Example 3
28:35
NH₄F → NH₄⁺ + F⁻
34:05
Example 4
38:03
Common Ion Effect & Buffers

41m 58s

Intro
0:00
Common Ion Effect & Buffers
1:16
Covalent Oxides Produce Acidic Solutions in Water
1:36
Ionic Oxides Produce Basic Solutions in Water
4:15
Practice Example 1
6:10
Practice Example 2
9:00
Definition
12:27
Example 1: Part A
16:49
Example 1: Part B
19:54
Buffer Solution
25:10
Example of Some Buffers: HF and NaF
30:02
Example of Some Buffers: Acetic Acid & Potassium Acetate
31:34
Example of Some Buffers: CH₃NH₂ & CH₃NH₃Cl
33:54
Example 2: Buffer Solution
36:36
Buffer

32m 24s

Intro
0:00
Buffers
1:20
Buffer Solution
1:21
Adding Base
5:03
Adding Acid
7:14
Example 1: Question
9:48
Example 1: Recall
12:08
Example 1: Major Species Upon Addition of NaOH
16:10
Example 1: Equilibrium, ICE Chart, and Final Calculation
24:33
Example 1: Comparison
29:19
Buffers, Part II

40m 6s

Intro
0:00
Buffers
1:27
Example 1: Question
1:32
Example 1: ICE Chart
3:15
Example 1: Major Species Upon Addition of OH⁻, But Before Rxn
7:23
Example 1: Equilibrium, ICE Chart, and Final Calculation
12:51
Summary
17:21
Another Look at Buffering & the Henderson-Hasselbalch equation
19:00
Example 2
27:08
Example 3
32:01
Buffers, Part III

38m 43s

Intro
0:00
Buffers
0:25
Buffer Capacity Part 1
0:26
Example 1
4:10
Buffer Capacity Part 2
19:29
Example 2
25:12
Example 3
32:02
Titrations: Strong Acid and Strong Base

42m 42s

Intro
0:00
Titrations: Strong Acid and Strong Base
1:11
Definition of Titration
1:12
Sample Problem
3:33
Definition of Titration Curve or pH Curve
9:46
Scenario 1: Strong Acid- Strong Base Titration
11:00
Question
11:01
Part 1: No NaOH is Added
14:00
Part 2: 10.0 mL of NaOH is Added
15:50
Part 3: Another 10.0 mL of NaOH & 20.0 mL of NaOH are Added
22:19
Part 4: 50.0 mL of NaOH is Added
26:46
Part 5: 100.0 mL (Total) of NaOH is Added
27:26
Part 6: 150.0 mL (Total) of NaOH is Added
32:06
Part 7: 200.0 mL of NaOH is Added
35:07
Titrations Curve for Strong Acid and Strong Base
35:43
Titrations: Weak Acid and Strong Base

42m 3s

Intro
0:00
Titrations: Weak Acid and Strong Base
0:43
Question
0:44
Part 1: No NaOH is Added
1:54
Part 2: 10.0 mL of NaOH is Added
5:17
Part 3: 25.0 mL of NaOH is Added
14:01
Part 4: 40.0 mL of NaOH is Added
21:55
Part 5: 50.0 mL (Total) of NaOH is Added
22:25
Part 6: 60.0 mL (Total) of NaOH is Added
31:36
Part 7: 75.0 mL (Total) of NaOH is Added
35:44
Titration Curve
36:09
Titration Examples & Acid-Base Indicators

52m 3s

Intro
0:00
Examples and Indicators
0:25
Example 1: Question
0:26
Example 1: Solution
2:03
Example 2: Question
12:33
Example 2: Solution
14:52
Example 3: Question
23:45
Example 3: Solution
25:09
Acid/Base Indicator Overview
34:45
Acid/Base Indicator Example
37:40
Acid/Base Indicator General Result
47:11
Choosing Acid/Base Indicator
49:12
VIII. Solubility
Solubility Equilibria

36m 25s

Intro
0:00
Solubility Equilibria
0:48
Solubility Equilibria Overview
0:49
Solubility Product Constant
4:24
Definition of Solubility
9:10
Definition of Solubility Product
11:28
Example 1
14:09
Example 2
20:19
Example 3
27:30
Relative Solubilities
31:04
Solubility Equilibria, Part II

42m 6s

Intro
0:00
Solubility Equilibria
0:46
Common Ion Effect
0:47
Example 1
3:14
pH & Solubility
13:00
Example of pH & Solubility
15:25
Example 2
23:06
Precipitation & Definition of the Ion Product
26:48
If Q > Ksp
29:31
If Q < Ksp
30:27
Example 3
32:58
Solubility Equilibria, Part III

43m 9s

Intro
0:00
Solubility Equilibria
0:55
Example 1: Question
0:56
Example 1: Step 1 - Check to See if Anything Precipitates
2:52
Example 1: Step 2 - Stoichiometry
10:47
Example 1: Step 3 - Equilibrium
16:34
Example 2: Selective Precipitation (Question)
21:02
Example 2: Solution
23:41
Classical Qualitative Analysis
29:44
Groups: 1-5
38:44
IX. Complex Ions
Complex Ion Equilibria

43m 38s

Intro
0:00
Complex Ion Equilibria
0:32
Complex Ion
0:34
Ligan Examples
1:51
Ligand Definition
3:12
Coordination
6:28
Example 1
8:08
Example 2
19:13
Complex Ions & Solubility

31m 30s

Intro
0:00
Complex Ions and Solubility
0:23
Recall: Classical Qualitative Analysis
0:24
Example 1
6:10
Example 2
16:16
Dissolving a Water-Insoluble Ionic Compound: Method 1
23:38
Dissolving a Water-Insoluble Ionic Compound: Method 2
28:13
X. Chemical Thermodynamics
Spontaneity, Entropy, & Free Energy, Part I

56m 28s

Intro
0:00
Spontaneity, Entropy, Free Energy
2:25
Energy Overview
2:26
Equation: ∆E = q + w
4:30
State Function/ State Property
8:35
Equation: w = -P∆V
12:00
Enthalpy: H = E + PV
14:50
Enthalpy is a State Property
17:33
Exothermic and Endothermic Reactions
19:20
First Law of Thermodynamic
22:28
Entropy
25:48
Spontaneous Process
33:53
Second Law of Thermodynamic
36:51
More on Entropy
42:23
Example
43:55
Spontaneity, Entropy, & Free Energy, Part II

39m 55s

Intro
0:00
Spontaneity, Entropy, Free Energy
1:30
∆S of Universe = ∆S of System + ∆S of Surrounding
1:31
Convention
3:32
Examining a System
5:36
Thermodynamic Property: Sign of ∆S
16:52
Thermodynamic Property: Magnitude of ∆S
18:45
Deriving Equation: ∆S of Surrounding = -∆H / T
20:25
Example 1
25:51
Free Energy Equations
29:22
Spontaneity, Entropy, & Free Energy, Part III

30m 10s

Intro
0:00
Spontaneity, Entropy, Free Energy
0:11
Example 1
2:38
Key Concept of Example 1
14:06
Example 2
15:56
Units for ∆H, ∆G, and S
20:56
∆S of Surrounding & ∆S of System
22:00
Reaction Example
24:17
Example 3
26:52
Spontaneity, Entropy, & Free Energy, Part IV

30m 7s

Intro
0:00
Spontaneity, Entropy, Free Energy
0:29
Standard Free Energy of Formation
0:58
Example 1
4:34
Reaction Under Non-standard Conditions
13:23
Example 2
16:26
∆G = Negative
22:12
∆G = 0
24:38
Diagram Example of ∆G
26:43
Spontaneity, Entropy, & Free Energy, Part V

44m 56s

Intro
0:00
Spontaneity, Entropy, Free Energy
0:56
Equations: ∆G of Reaction, ∆G°, and K
0:57
Example 1: Question
6:50
Example 1: Part A
9:49
Example 1: Part B
15:28
Example 2
17:33
Example 3
23:31
lnK = (- ∆H° ÷ R) ( 1 ÷ T) + ( ∆S° ÷ R)
31:36
Maximum Work
35:57
XI. Electrochemistry
Oxidation-Reduction & Balancing

39m 23s

Intro
0:00
Oxidation-Reduction and Balancing
2:06
Definition of Electrochemistry
2:07
Oxidation and Reduction Review
3:05
Example 1: Assigning Oxidation State
10:15
Example 2: Is the Following a Redox Reaction?
18:06
Example 3: Step 1 - Write the Oxidation & Reduction Half Reactions
22:46
Example 3: Step 2 - Balance the Reaction
26:44
Example 3: Step 3 - Multiply
30:11
Example 3: Step 4 - Add
32:07
Example 3: Step 5 - Check
33:29
Galvanic Cells

43m 9s

Intro
0:00
Galvanic Cells
0:39
Example 1: Balance the Following Under Basic Conditions
0:40
Example 1: Steps to Balance Reaction Under Basic Conditions
3:25
Example 1: Solution
5:23
Example 2: Balance the Following Reaction
13:56
Galvanic Cells
18:15
Example 3: Galvanic Cells
28:19
Example 4: Galvanic Cells
35:12
Cell Potential

48m 41s

Intro
0:00
Cell Potential
2:08
Definition of Cell Potential
2:17
Symbol and Unit
5:50
Standard Reduction Potential
10:16
Example Figure 1
13:08
Example Figure 2
19:00
All Reduction Potentials are Written as Reduction
23:10
Cell Potential: Important Fact 1
26:49
Cell Potential: Important Fact 2
27:32
Cell Potential: Important Fact 3
28:54
Cell Potential: Important Fact 4
30:05
Example Problem 1
32:29
Example Problem 2
38:38
Potential, Work, & Free Energy

41m 23s

Intro
0:00
Potential, Work, Free Energy
0:42
Descriptions of Galvanic Cell
0:43
Line Notation
5:33
Example 1
6:26
Example 2
11:15
Example 3
15:18
Equation: Volt
22:20
Equations: Cell Potential, Work, and Charge
28:30
Maximum Cell Potential is Related to the Free Energy of the Cell Reaction
35:09
Example 4
37:42
Cell Potential & Concentration

34m 19s

Intro
0:00
Cell Potential & Concentration
0:29
Example 1: Question
0:30
Example 1: Nernst Equation
4:43
Example 1: Solution
7:01
Cell Potential & Concentration
11:27
Example 2
16:38
Manipulating the Nernst Equation
25:15
Example 3
28:43
Electrolysis

33m 21s

Intro
0:00
Electrolysis
3:16
Electrolysis: Part 1
3:17
Electrolysis: Part 2
5:25
Galvanic Cell Example
7:13
Nickel Cadmium Battery
12:18
Ampere
16:00
Example 1
20:47
Example 2
25:47
XII. Light
Light

44m 45s

Intro
0:00
Light
2:14
Introduction to Light
2:15
Frequency, Speed, and Wavelength of Waves
3:58
Units and Equations
7:37
Electromagnetic Spectrum
12:13
Example 1: Calculate the Frequency
17:41
E = hν
21:30
Example 2: Increment of Energy
25:12
Photon Energy of Light
28:56
Wave and Particle
31:46
Example 3: Wavelength of an Electron
34:46
XIII. Quantum Mechanics
Quantum Mechanics & Electron Orbitals

54m

Intro
0:00
Quantum Mechanics & Electron Orbitals
0:51
Quantum Mechanics & Electron Orbitals Overview
0:52
Electron Orbital and Energy Levels for the Hydrogen Atom
8:47
Example 1
13:41
Quantum Mechanics: Schrodinger Equation
19:19
Quantum Numbers Overview
31:10
Principal Quantum Numbers
33:28
Angular Momentum Numbers
34:55
Magnetic Quantum Numbers
36:35
Spin Quantum Numbers
37:46
Primary Level, Sublevels, and Sub-Sub-Levels
39:42
Example
42:17
Orbital & Quantum Numbers
49:32
Electron Configurations & Diagrams

34m 4s

Intro
0:00
Electron Configurations & Diagrams
1:08
Electronic Structure of Ground State Atom
1:09
Order of Electron Filling
3:50
Electron Configurations & Diagrams: H
8:41
Electron Configurations & Diagrams: He
9:12
Electron Configurations & Diagrams: Li
9:47
Electron Configurations & Diagrams: Be
11:17
Electron Configurations & Diagrams: B
12:05
Electron Configurations & Diagrams: C
13:03
Electron Configurations & Diagrams: N
14:55
Electron Configurations & Diagrams: O
15:24
Electron Configurations & Diagrams: F
16:25
Electron Configurations & Diagrams: Ne
17:00
Electron Configurations & Diagrams: S
18:08
Electron Configurations & Diagrams: Fe
20:08
Introduction to Valence Electrons
23:04
Valence Electrons of Oxygen
23:44
Valence Electrons of Iron
24:02
Valence Electrons of Arsenic
24:30
Valence Electrons: Exceptions
25:36
The Periodic Table
27:52
XIV. Intermolecular Forces
Vapor Pressure & Changes of State

52m 43s

Intro
0:00
Vapor Pressure and Changes of State
2:26
Intermolecular Forces Overview
2:27
Hydrogen Bonding
5:23
Heat of Vaporization
9:58
Vapor Pressure: Definition and Example
11:04
Vapor Pressures is Mostly a Function of Intermolecular Forces
17:41
Vapor Pressure Increases with Temperature
20:52
Vapor Pressure vs. Temperature: Graph and Equation
22:55
Clausius-Clapeyron Equation
31:55
Example 1
32:13
Heating Curve
35:40
Heat of Fusion
41:31
Example 2
43:45
Phase Diagrams & Solutions

31m 17s

Intro
0:00
Phase Diagrams and Solutions
0:22
Definition of a Phase Diagram
0:50
Phase Diagram Part 1: H₂O
1:54
Phase Diagram Part 2: CO₂
9:59
Solutions: Solute & Solvent
16:12
Ways of Discussing Solution Composition: Mass Percent or Weight Percent
18:46
Ways of Discussing Solution Composition: Molarity
20:07
Ways of Discussing Solution Composition: Mole Fraction
20:48
Ways of Discussing Solution Composition: Molality
21:41
Example 1: Question
22:06
Example 1: Mass Percent
24:32
Example 1: Molarity
25:53
Example 1: Mole Fraction
28:09
Example 1: Molality
29:36
Vapor Pressure of Solutions

37m 23s

Intro
0:00
Vapor Pressure of Solutions
2:07
Vapor Pressure & Raoult's Law
2:08
Example 1
5:21
When Ionic Compounds Dissolve
10:51
Example 2
12:38
Non-Ideal Solutions
17:42
Negative Deviation
24:23
Positive Deviation
29:19
Example 3
31:40
Colligatives Properties

34m 11s

Intro
0:00
Colligative Properties
1:07
Boiling Point Elevation
1:08
Example 1: Question
5:19
Example 1: Solution
6:52
Freezing Point Depression
12:01
Example 2: Question
14:46
Example 2: Solution
16:34
Osmotic Pressure
20:20
Example 3: Question
28:00
Example 3: Solution
30:16
XV. Bonding
Bonding & Lewis Structure

48m 39s

Intro
0:00
Bonding & Lewis Structure
2:23
Covalent Bond
2:24
Single Bond, Double Bond, and Triple Bond
4:11
Bond Length & Intermolecular Distance
5:51
Definition of Electronegativity
8:42
Bond Polarity
11:48
Bond Energy
20:04
Example 1
24:31
Definition of Lewis Structure
31:54
Steps in Forming a Lewis Structure
33:26
Lewis Structure Example: H₂
36:53
Lewis Structure Example: CH₄
37:33
Lewis Structure Example: NO⁺
38:43
Lewis Structure Example: PCl₅
41:12
Lewis Structure Example: ICl₄⁻
43:05
Lewis Structure Example: BeCl₂
45:07
Resonance & Formal Charge

36m 59s

Intro
0:00
Resonance and Formal Charge
0:09
Resonance Structures of NO₃⁻
0:25
Resonance Structures of NO₂⁻
12:28
Resonance Structures of HCO₂⁻
16:28
Formal Charge
19:40
Formal Charge Example: SO₄²⁻
21:32
Formal Charge Example: CO₂
31:33
Formal Charge Example: HCN
32:44
Formal Charge Example: CN⁻
33:34
Formal Charge Example: 0₃
34:43
Shapes of Molecules

41m 21s

Intro
0:00
Shapes of Molecules
0:35
VSEPR
0:36
Steps in Determining Shapes of Molecules
6:18
Linear
11:38
Trigonal Planar
11:55
Tetrahedral
12:45
Trigonal Bipyramidal
13:23
Octahedral
14:29
Table: Shapes of Molecules
15:40
Example: CO₂
21:11
Example: NO₃⁻
24:01
Example: H₂O
27:00
Example: NH₃
29:48
Example: PCl₃⁻
32:18
Example: IF₄⁺
34:38
Example: KrF₄
37:57
Hybrid Orbitals

40m 17s

Intro
0:00
Hybrid Orbitals
0:13
Introduction to Hybrid Orbitals
0:14
Electron Orbitals for CH₄
5:02
sp³ Hybridization
10:52
Example: sp³ Hybridization
12:06
sp² Hybridization
14:21
Example: sp² Hybridization
16:11
σ Bond
19:10
π Bond
20:07
sp Hybridization & Example
22:00
dsp³ Hybridization & Example
27:36
d²sp³ Hybridization & Example
30:36
Example: Predict the Hybridization and Describe the Molecular Geometry of CO
32:31
Example: Predict the Hybridization and Describe the Molecular Geometry of BF₄⁻
35:17
Example: Predict the Hybridization and Describe the Molecular Geometry of XeF₂
37:09
XVI. AP Practice Exam
AP Practice Exam: Multiple Choice, Part I

52m 34s

Intro
0:00
Multiple Choice
1:21
Multiple Choice 1
1:22
Multiple Choice 2
2:23
Multiple Choice 3
3:38
Multiple Choice 4
4:34
Multiple Choice 5
5:16
Multiple Choice 6
5:41
Multiple Choice 7
6:20
Multiple Choice 8
7:03
Multiple Choice 9
7:31
Multiple Choice 10
9:03
Multiple Choice 11
11:52
Multiple Choice 12
13:16
Multiple Choice 13
13:56
Multiple Choice 14
14:52
Multiple Choice 15
15:43
Multiple Choice 16
16:20
Multiple Choice 17
16:55
Multiple Choice 18
17:22
Multiple Choice 19
18:59
Multiple Choice 20
20:24
Multiple Choice 21
22:20
Multiple Choice 22
23:29
Multiple Choice 23
24:30
Multiple Choice 24
25:24
Multiple Choice 25
26:21
Multiple Choice 26
29:06
Multiple Choice 27
30:42
Multiple Choice 28
33:28
Multiple Choice 29
34:38
Multiple Choice 30
35:37
Multiple Choice 31
37:31
Multiple Choice 32
38:28
Multiple Choice 33
39:50
Multiple Choice 34
42:57
Multiple Choice 35
44:18
Multiple Choice 36
45:52
Multiple Choice 37
48:02
Multiple Choice 38
49:25
Multiple Choice 39
49:43
Multiple Choice 40
50:16
Multiple Choice 41
50:49
AP Practice Exam: Multiple Choice, Part II

32m 15s

Intro
0:00
Multiple Choice
0:12
Multiple Choice 42
0:13
Multiple Choice 43
0:33
Multiple Choice 44
1:16
Multiple Choice 45
2:36
Multiple Choice 46
5:22
Multiple Choice 47
6:35
Multiple Choice 48
8:02
Multiple Choice 49
10:05
Multiple Choice 50
10:26
Multiple Choice 51
11:07
Multiple Choice 52
12:01
Multiple Choice 53
12:55
Multiple Choice 54
16:12
Multiple Choice 55
18:11
Multiple Choice 56
19:45
Multiple Choice 57
20:15
Multiple Choice 58
23:28
Multiple Choice 59
24:27
Multiple Choice 60
26:45
Multiple Choice 61
29:15
AP Practice Exam: Multiple Choice, Part III

32m 50s

Intro
0:00
Multiple Choice
0:16
Multiple Choice 62
0:17
Multiple Choice 63
1:57
Multiple Choice 64
6:16
Multiple Choice 65
8:05
Multiple Choice 66
9:18
Multiple Choice 67
10:38
Multiple Choice 68
12:51
Multiple Choice 69
14:32
Multiple Choice 70
17:35
Multiple Choice 71
22:44
Multiple Choice 72
24:27
Multiple Choice 73
27:46
Multiple Choice 74
29:39
Multiple Choice 75
30:23
AP Practice Exam: Free response Part I

47m 22s

Intro
0:00
Free Response
0:15
Free Response 1: Part A
0:16
Free Response 1: Part B
4:15
Free Response 1: Part C
5:47
Free Response 1: Part D
9:20
Free Response 1: Part E. i
10:58
Free Response 1: Part E. ii
16:45
Free Response 1: Part E. iii
26:03
Free Response 2: Part A. i
31:01
Free Response 2: Part A. ii
33:38
Free Response 2: Part A. iii
35:20
Free Response 2: Part B. i
37:38
Free Response 2: Part B. ii
39:30
Free Response 2: Part B. iii
44:44
AP Practice Exam: Free Response Part II

43m 5s

Intro
0:00
Free Response
0:12
Free Response 3: Part A
0:13
Free Response 3: Part B
6:25
Free Response 3: Part C. i
11:33
Free Response 3: Part C. ii
12:02
Free Response 3: Part D
14:30
Free Response 4: Part A
21:03
Free Response 4: Part B
22:59
Free Response 4: Part C
24:33
Free Response 4: Part D
27:22
Free Response 4: Part E
28:43
Free Response 4: Part F
29:35
Free Response 4: Part G
30:15
Free Response 4: Part H
30:48
Free Response 5: Diagram
32:00
Free Response 5: Part A
34:14
Free Response 5: Part B
36:07
Free Response 5: Part C
37:45
Free Response 5: Part D
39:00
Free Response 5: Part E
40:26
AP Practice Exam: Free Response Part III

28m 36s

Intro
0:00
Free Response
0:43
Free Response 6: Part A. i
0:44
Free Response 6: Part A. ii
3:08
Free Response 6: Part A. iii
5:02
Free Response 6: Part B. i
7:11
Free Response 6: Part B. ii
9:40
Free Response 7: Part A
11:14
Free Response 7: Part B
13:45
Free Response 7: Part C
15:43
Free Response 7: Part D
16:54
Free Response 8: Part A. i
19:15
Free Response 8: Part A. ii
21:16
Free Response 8: Part B. i
23:51
Free Response 8: Part B. ii
25:07
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Lecture Comments (26)

1 answer

Last reply by: Professor Hovasapian
Fri Jan 13, 2017 7:11 PM

Post by Areez Khaki on December 9, 2016

Hi professor, i dont understand why you would need to balance the reaction solution with the barium chloride example, you added a 2 to KCl to "balence the equation"  what did you mean by that

1 answer

Last reply by: Professor Hovasapian
Sun Feb 22, 2015 7:45 PM

Post by Jason Smith on February 21, 2015

Is it possible for two soluble ionic compounds to mix together and create two non-soluable compounds?

1 answer

Last reply by: Professor Hovasapian
Sun Nov 2, 2014 2:48 AM

Post by David Gonzalez on November 1, 2014

Hi professor, great lecture! Always so thorough and informative in your explanations. I have one question: why does aqueous chemistry always take place in water? Can't the solvent be anything else? Thank you!

1 answer

Last reply by: Professor Hovasapian
Fri Jan 17, 2014 9:24 PM

Post by Jérémie Lessard on January 15, 2014

Hi Professor Hovasapian !

My question refers to the second example (part 26:10). If the ions do not react with each others, can we say that they are in an dynamic equilibrium ? (since there is no precipating species)

Thank you !

Jérémie

0 answers

Post by Professor Hovasapian on December 22, 2013

Hi Burhan.

I hope all is well with you.

The topics you mention run in a straight sequence one after the other for my course, so start with the Kinetics lesson and just move forward. The redox stuff is under Electrochemistry.

I hope that helps.

Best wishes.

Raffi

1 answer

Last reply by: Professor Hovasapian
Sun Dec 22, 2013 3:51 PM

Post by Burhan Akram on December 22, 2013

Hello Professor Raffi,

I live in B.C, Canada and the grade 12 Chemistry I am studying has the following,

1) Reaction Rates
2) Equilibrium
3) Solubility
4) Acids & Bases
5) Redox Reactions

Could you please tell me which lessons I should watch to help me study efficiently rather than watching lessons which I don't require for this course?

I would really appreciate that,

Thank You

Burhan

3 answers

Last reply by: Professor Hovasapian
Tue Dec 3, 2013 12:48 AM

Post by Alexis Yates on October 15, 2013

Is there a way to measure the level of dissociation of a compound mathematically of is it something that I should just try to memorize?

1 answer

Last reply by: Professor Hovasapian
Sun Sep 1, 2013 10:51 PM

Post by Stephanie Dahlström on September 1, 2013

You're a great teacher! I've had so much trouble understanding chemistry until I started watching your lectures. You make it really easy to understand. Thank you so much!

5 answers

Last reply by: Professor Hovasapian
Fri Feb 7, 2014 6:50 PM

Post by Antie Chen on April 14, 2013

What's the formula of acetic acid? I consider it should be CH3COOH, and one molecule contains 4 Hydrogen atoms.
What's the formula H2C2H3O2?

2 answers

Last reply by: Thomas Dykstra
Wed Jul 11, 2012 2:43 PM

Post by Pierre-Alexandre Leblanc on April 21, 2012

why cant we fast forward to examples on any videos

Precipitation Reactions

  • When Ionic Compound dissolve, they actually dissociate – break up into free-floating separate ions. We call such compounds soluble.
  • There are guidelines for which ions form compounds that are soluble, and which are not – this is readily available as a chart in your book or on the web.
  • When one solution of dissolved Ionic Compound is mixed with another solution of dissolved Ionic Compound, sometimes a Cation-Anion pair will be such that their combination is not soluble in water.
  • This Insoluble compound forms a solid in solution and, generally, falls to the bottom of the container – it is called a Precipitate.
  • When mixing two soluble ionic compounds, switch Cation-Anion pairs and check the solubility guidelines to decide if a precipitate forms.

Precipitation Reactions

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
  • Precipitation Reactions 0:53
    • Dissociation of ionic Compounds
    • Solubility Guidelines for ionic Compounds: Soluble Ionic Compounds
    • Solubility Guidelines for ionic Compounds: Insoluble ionic Compounds
    • Precipitation Reactions
    • Example 1: Mixing a Solution of BaCl₂ & K₂SO₄
    • Example 2: Mixing a Solution of Mg(NO₃)₂ & KI

Transcription: Precipitation Reactions

Hello, and welcome back to Educator.com.0000

Welcome back to AP Chemistry.0002

Today, we're going to be talking about precipitation reactions; more generally, over the next couple of lessons, we're going to be talking about aqueous solutions and solution stoichiometry.0005

We'll be talking, particularly, about precipitation reactions today, and then in the next couple of lessons, we'll be talking about acid, base, and oxidation-reduction reactions.0015

Those are the three primary reactions in chemistry; it's either going to be some kind of precipitation, some acid-base, or some oxidation-reduction.0024

These are primarily aqueous because most of the chemistry that you really are going to end up doing is going to be in some sort of a water medium.0032

That is what aqueous means--just chemistry that takes place in the water--wet chemistry, in other words.0041

Let's go ahead and get started and see what we can do.0049

The first thing I want to talk about--we talked about naming ionic compounds; ionic compounds, remember, are compounds that have a cation and an anion, and they come together in the appropriate proportions to form a neutral compound.0054

Well, something very, very interesting happens to a fair number of ionic compounds--in fact, most ionic compounds to some extent--but some more than others.0070

They actually dissolve; and they do more than just dissolve; they actually dissociate, and we're going to discuss that right now, before we get into the precipitation reactions formally.0078

OK, so let's talk about the dissociation of ionic compounds.0088

Dissociation is exactly what the word looks like: it's dis-association.0099

Let's take something like sodium chloride: when they come together, you have this extended crystal of sodium chloride: sodium chloride, sodium chloride, sodium chloride, one after the other, in this long chain--a three-dimensional crystal structure, which is what you have when you look at salt.0103

When you drop it in water, the individual molecules of sodium chloride don't just separate--the ions just completely come apart from each other, so you have chloride ions and sodium ions floating around freely in the solution, each one of those surrounded by water molecules.0117

That is what dissociation means: it means it actually completely comes apart into its constituent elements.0133

Now that I have said that, let's talk about how it actually looks in...how we write it out formally.0141

Let's take NaCl, and I'll put S: S stands for solid--just regular salt; now, you're going to see this written a couple of ways: sometimes (which is often the fashion nowadays), they'll go ahead and just write H2O, which means they are taking solid salt and dropping it in some water; and then you see an arrow and what it is that happens.0148

I do it a little bit differently, and there is a reason why I do it differently, and the reason is I like the actual ionic compound to be what you focus on, not necessarily the water.0171

So, what I do is: I do that, and I put a little H2O on top of the arrow, which means I have dissolved it in water.0184

It could be any kind of amount; so water is irrelevant--it just happens to be the medium in which we're doing our chemistry: aqueous chemistry.0190

Solid salt dissolved in the water: what you end up getting is Na+ + Cl-.0197

Now, sodium chloride happens to be something called a strong electrolyte, and what that means is: it comes apart completely.0202

This NaCl--a crystal of NaCl--doesn't stick together, doesn't clump down at the bottom of the beaker.0210

Yes, when you first drop it in there--yes, you see it at the bottom, but when you stir it around, it dissolves: that means that the sodium and the chloride ions are leaving from the surface of the crystal, and now they're just floating around freely.0216

Once it's fully dissolved, when you look at the solution, you can't actually see the salt; it just looks like a slightly cloudy solution, but it's homogeneous.0227

That is what dissociation means: it comes apart into its constituent compounds.0237

Another example would be something like iron (2) sulfate.0243

Remember iron (2) sulfate?--two oxidation states, so we have to call it that.0249

This is a solid; when we drop it in water, we're going to get a free iron ion floating around in solution, plus SO42-; and notice this polyatomic ion--it stays together.0253

Dissociation of ionic compounds means it dissociates into its individual ions, not into its individual elements.0266

So, a polyatomic ion stays together as a group: carbonate, dichromate--these things stay together.0273

There is also something else that we write--a little (aq) that you see; this (aq) means that it is dissolved; it is in water; it's surrounded by water--solvated; (aq), (aq), to differentiate it from the solid.0278

Again, iron sulfate--iron (2) sulfate--is also a strong electrolyte, and we'll talk about solubility in a minute--the extent to which these things actually do come apart completely.0296

But, it dissociates completely; you are not going to find any iron, any sulfate, in solid form.0307

OK, now let's talk about something...an acid; and we're going to use acetic acid, so I'm going to write H2C2H3O2 (acetic acid is just vinegar).0315

When I drop it in water, something interesting happens with acetic acid; now, acid--for those of you that are familiar with acids (if you are not, not a problem--we're going to be discussing it in the next lesson)--an acid is something that has hydrogens that it gives up.0327

So, when you drop an acid in water, the actual compound in water--it actually gives up one of its hydrogen ions, and it gives it up one at a time.0342

But what happens first is: when you drop acetic acid in water, acetic acid itself stays together as a molecule, and it actually dissolves in water.0351

So, we actually write it like this: H2C2H3O2, and we put a little (aq) here.0361

Well, this (aq) means that the whole molecule is surrounded by water, so you can't actually see the individual molecules of acetic acid anymore.0368

Then, what we do--we have a double arrow here; now that it is an aqueous solution, now it gives up one of its hydrogen ions--it dissociates into its ionic constituents: hydrogen ion plus (if you don't mind, instead of writing C2H3O2, I'm just going to write Ac--that is a common abbreviation for the acetate ion).0377

So, notice, in this particular case, what we have done here is just write the particular ionic compound dissociates into its constituents.0401

This ionic compound dissociates into its constituents; acetic acid happens to be a weak electrolyte; in other words, it doesn't dissociate completely into its hydrogen ion and acetate ion form.0409

In fact, most of it stays in this form--the acetic acid form.0421

But, it's dissolved in water; I take pure acetic acid, and drop it in water, and you can't tell the difference: when you look, it just looks like a solution of water.0425

Well, there is acetic acid in there; it is surrounded by water molecules, but it hasn't dissociated, which is why you see this double arrow.0433

Notice, up here, you see single arrows; double arrow means that it goes both ways.0441

In fact, later in the course, we'll find a way to actually quantify the extent to which this dissociates--meaning we'll attach some math to it and measure how much of this dissociates.0445

As it turns out, acetic acid is about 1 to 2% dissociated, under normal conditions.0456

Just to let you know, you will often see it like this: you will see it from here--this is usually how it's written, but I wanted you to see what it is that actually happens, chemically.0461

Pure acetic acid, dropped in water, is solvated by the water, and then it will dissociate into its constituent ions partially; weak electrolyte/strong electrolyte.0472

Strong electrolyte means completely dissociated in aqueous solution.0482

OK, now for the most part we're going to be concerned with electrolytes that dissociate completely, but every once in a while we will discuss things like acetic acid.0486

All right, having said this, now let's talk about precipitation reactions.0497

Precipitation reactions (rxn is just a shorthand for reaction): If I have a solution of some salt (and again, salt is just a fancy word for ionic compound)--sodium chloride, and I mix it with a solution of iron sulfate, something might happen.0504

As it turns out, there are certain salts that are so tightly bound together that, even when I drop them in water, they don't dissociate.0526

Or, if they do, they dissociate so little that, for all practical purposes, they don't dissociate.0537

We can measure very, very tiny amounts, but really, what happens--it just sinks to the bottom.0544

Well, what is kind of interesting is: when you take, let's say, a mixture of a salt that is fully dissociated, and you pour in another solution of a salt that is also dissociated, now ions are sort of slamming into each other randomly.0549

Well, some of those ions, when they stick together--they stick together so tightly that they literally fall out of solution as a solid--they just...literally, a solid shows up in the middle of solution and drops to the bottom.0562

We call that a precipitate.0573

Precipitation reactions are reactions where, when you mix things, a solid all of a sudden shows up in the reaction mixture.0575

Sometimes, it will be a liquid, but more often than not, it will be a solid.0585

The way we decide what is actually going to stick together: we use something called the Solubility Guidelines.0588

The Solubility Guidelines--I'm going to refer you to your text, because I think it's important that you find some of these important tables in your text, and be familiar with them--at least where they are.0597

Basically, they are just guidelines to tell you which ionic compounds tend to dissolve in water, and which don't dissolve in water--which just fall to the bottom without dissolving at all.0610

You're going to be referring to this over and over again.0620

Your teacher may ask you to memorize it; maybe not; for the most part, it's just a reference to decide, when you're faced with a reaction of two ionic compounds, is there going to be a precipitate?--and you decide based on solubility.0624

You'll see what I mean in just a minute.0637

So, a Solubility Guideline for ionic compounds looks like the following.0639

I just want you to see what you can actually expect to see in your book as a table.0650

You're going to see something that looks like this: it says (you know what, let me do this in red--it'll be a lot better): Soluble Ionic Compounds, and then it has a column called Important Exceptions.0655

You will see something that says like this: Compounds Containing...and it will list a bunch of ions: nitrate, carbonate, iodide, bromide, things like that.0698

Over here, it will say None, because any ionic compound that is compounded with a nitrate ion--they are all going to be soluble--completely soluble.0714

Then, let's say there is an entry for iodide: this means that all compounds that have an iodide in them (like magnesium iodide, lead iodide, something like that)--in general, they tend to be soluble.0728

Then, the exceptions are: in the case of things that are mixed with iodide, lead2+ ion, mercury22+ ion, and silver.0741

So, you're going to see something like this--a bunch of charts; you're going to see ions, and then you're going to see the exceptions.0753

Then, there is going to be another set, and it's going to say (oops, I have all kinds of stray lines showing up here--let's see...)--then you're going to see another that says Insoluble Ionic Compounds.0759

It's also going to have a column of important exceptions, and it's going to list a bunch of ions that...those ions that have those particular ion it are generally insoluble--they don't dissolve in water.0784

Then, there is going to be some Exceptions.0798

For example, you will see hydroxide; generally, anything that is bound to a hydroxide or multiple hydroxides--for the most part, most of the metals--most of the ionic compounds--that have hydroxide in them are insoluble in water, the exceptions being the alkali metals.0802

So, over in this column, it will say "alkali metal," and you will know that, when you see a hydroxide with an alkali metal, like lithium hydroxide or potassium hydroxide, you will know that it is an exception to this rule and they are completely soluble.0819

Now that we know what this looks like, again, I encourage you to find it in your book; it will usually be in one of the chapters, and it will say Solubility Guidelines for Ionic Compounds, or something like that.0834

Get to know that very, very well.0845

OK, so let's do a problem, and I think everything will come together nicely.0848

Let's look at magnesium nitrate, Mg(NO3)2 (right?--magnesium is 2+; NO3 is 2-; we need two of them to balance the charge) (aq) + sodium hydroxide--again, (aq), and Mg(OH)2; this one...we'll say S, and we have NaNO3, sodium nitrate.0852

OK, when you see an equation like this, the question is, "What is going on here?"0889

This is a shorthand notation for what is going on chemically.0894

Now, we're going to go ahead and describe what is going on chemically, and how you decide what is going on chemically.0898

The first thing we want to do is to make sure that this is actually balanced, so let's see: I have two nitrates over here; I have one magnesium, one magnesium; two hydroxides, one hydroxide--so I'm going to stick a 2 in front of that; now, that gives me 2 sodiums--I'm going to stick a 2 there--2 sodiums; 2 nitrates; 2 nitrates; I'm balanced--2 in front of the sodium hydroxide, 2 in front of the sodium nitrate.0904

Now, notice this aqueous, aqueous, aqueous; there is an (aq) over here, too--my apologies (and let me go back to blue).0931

This aqueous--what this means is that, even though we write it as a molecule, well...if I look under the Solubility Guidelines for magnesium nitrate, anything that has a nitrate is soluble.0938

It means that a solution of this (which--aqueous means that you have dropped some magnesium nitrate in water)--it's completely dissolved; it's a strong electrolyte--it dissolves completely; it's soluble.0950

So, it looks like this: what is floating around in solution is a magnesium 2+ ion and two free nitrate ions.0961

You also have two sodium ions, because sodium hydroxide is also soluble.0972

Hydroxide is generally insoluble, but the exception is the alkali metals; sodium is an alkali metal; the hydroxide of an alkali metal is completely soluble, so you have two sodium ions floating around, and you have two hydroxide ions floating around.0980

That is where the 2 comes from, here.0996

Now, here is what is going on: everything on the left side is what happens before a reaction takes place.0999

This arrow--everything here is after the reaction has taken place; now, let's see what has really happened.1005

Well, what is going to happen here is: this is going to slam into nitrate, but when it slams into nitrate, it's just going to come apart again, because it's soluble.1011

It's going to slam into sodium, but when it slams into sodium, they're both positively charged--they're going to bounce off each other.1020

But, when it slams into a hydroxide, because a hydroxide, under the Solubility Guidelines, is insoluble, it won't bounce off, and it won't dissolve again; the binding strength is so strong that they literally stick together, and they stick together so completely that it literally turns into a solid again.1026

So, this is...essentially, what happens is: you are switching partners.1046

The cation of one goes with the anion of the other, and this cation goes with this anion.1051

So, Mg and OH--so we have to make sure that's appropriate, so it's going to be Mg(OH)2, because hydroxide is -1 and magnesium is +2.1058

Because it is insoluble, that is why we write it as a solid.1069

It literally falls out of solution, down to the bottom of the flask.1075

You will also see this written with an arrow pointing down; that is an old symbol for precipitate.1078

I tend to prefer this symbol, myself.1085

The symbol for a gas is an arrow pointing up, meaning the gas bubbles off.1087

Now, let's look at the other combination--sodium and nitrate: well, sodium nitrate--we know that anything that has a nitrate in it is going to be soluble, so they stay as free ions; they might slam together, but they come apart again, because sodium nitrate is soluble.1092

So, you have two sodium ions and two nitrate ions that are, again, floating around in solution.1107

Now, watch what happens: you know that, in algebra, whenever you have something on the left side of the equality sign which is the same as the right side of the equality sign, they cancel out.1115

The same thing here: I have two nitrates on the left, two nitrates on the right; I can cancel those--they don't really participate in any chemistry.1125

I have two sodiums on the left, two sodiums on the right; now, I'm going to write a third equation, which is basically just what is left over.1133

Mg2+ + 2OH- → Mg(OH)2 as a precipitate.1143

These three equations--the first one is called a molecular equation; you're just giving molecular formulas--you're not really talking about what is going on.1154

The second equation tells you everything that is in solution; everything is right there--that is called the total ionic equation.1162

The total ionic is my personal favorite, because you see everything that is involved--there is nothing hidden.1172

This is it; this is the chemistry; the total ionic--that is where the chemistry takes place.1178

The third equation is called the net ionic; it's the actual chemical reaction that takes place.1183

We cancel the nitrate and the sodium because they don't do anything--they just float around; they are called spectator ions.1189

I'm not sure if the names are altogether that important; what is important is their function, which in this case is nothing--they just sit there and do nothing.1195

This is the chemistry.1202

So, again, the first equation: it is called the molecular, because we give the molecular formulas--the balanced molecular formulas.1203

The second equation is called total ionic (your book might call it complete ionic), because you are listing every ionic species.1217

The third, which is the actual chemistry taking place, is the net ionic.1227

It's called ionic simply because, on the left-hand side, you have ions, and here you have a precipitate--a solid that sinks to the bottom.1234

Notice, it's neutral; charges are balanced; neutral--zero charge on the right side; 2+, 2-, zero charge on the left side; zero net charge.1242

That is what is happening with precipitation reactions.1252

You're going to really mix two soluble compounds, and sometimes the ions--one cation and one anion--will be such that they are actually insoluble in solution; they'll bind really tightly, and they'll drop to the bottom of the flask as a solid.1255

You pour out the water; you filter it, and you collect the solid; you weigh it, and you do all of the other things that you're going to end up doing in a lab.1269

Let's do an examples.1276

Let's move forward: let's call this one Example 1 (the one previous was just a discussion): this time, we're not going to write the right half; we're going to actually figure out what is going on as a product.1282

The question is: what happens when you mix a solution (sol'n is a shorthand for solution) of barium chloride and potassium sulfate--BaCl2 and K2SO4?1298

Well, let's write out the molecular formula and see what happens; we'll switch partners and see what we get.1326

So, we have BaCl2; that is aqueous (well, actually, let's not write--let's just write BaCl2 first); and then we'll write K2SO4 and draw a little arrow; now, let's decide what is happening here.1332

Barium chloride: we look under the solubility guidelines, and we discover that it is soluble.1349

So, this is going to be a Ba2+, and again, full dissociation; that is what solubility means--full dissociation.1354

This Cl2 breaks up into 2 chloride ions, 2Cl- + K2SO4: sulfate may or may not be soluble, but in the case of a sulfate that is put together with an alkali metal, which potassium is, they are completely soluble.1362

So, again, they dissolve completely; that means we have 2 potassium ions in the solution, plus an SO42-.1381

Now, we switch partners; the barium goes with the sulfate.1392

Now, when we put barium and sulfate together, as it turns out, barium sulfate is not soluble; therefore, it has a little arrow pointing down (or you can put an S for solid); that is a precipitate.1397

The other combination is, as far as the molecular formula is concerned, just potassium and chloride, which is KCl.1412

Again, it isn't K2Cl2; we have to put them together the way we did when we named compounds.1421

Potassium is a +1; chlorine is a -1; it's just KCl.1427

Now, we can balance the molecular formula.1431

I actually should have done that first, instead of starting to write the total ionic; forgive me.1433

All you have do is switch partners; so, there are 2 potassiums here; I'm going to put a 2 there; here it's 2 chloride--I'm going to come over here--that is 2 chloride; 1 barium, 1 barium; 1 sulfate, 1 sulfate; good.1438

All I needed was this 2, and the equation is balanced.1449

Now, we already took care of the reactants: barium, chloride, and potassium sulfate; we broke them up into ions.1453

We said that barium sulfate is a precipitate, so it stays together; it does not separate into ions; they actually bind together.1460

Potassium chloride is soluble; therefore, it stays as 2 free potassium ions plus 2 free chloride ions.1469

Now, I take a look at what is on the left and what is on the right; the 2 chloride and the 2 potassium cancel with the 2 chloride and the 2 potassium, and I end up with Ba2+ ion + a sulfate ion; it gives me barium sulfate as a solid that falls to the bottom of the flask.1478

The first is my molecular equation that I got just by switching partners and combining them in the appropriate stoichiometric quantities.1500

The second is a total ionic equation, meaning everything that is happening, based on the Solubility Guidelines.1508

So, these two are soluble; therefore, they are floating around as free solutions; the barium sulfate is not; therefore, it sticks together.1514

The potassium chloride is soluble; therefore, it is floating around as free ions.1520

We cancel what is on the left and the right, and we are left with our net ionic.1527

Molecular, total ionic, net ionic: this is a complete description of the chemistry that is going on here.1530

I hope that this is making sense; it's reasonable straightforward--a lot of symbolism, but there is nothing strange going on here.1536

Your intuition should be able to guide you through, as long as you realize that this molecular--when you're dealing with ionic compounds, they may or may not dissolve.1543

What you are checking is solubility--if they are soluble, you write them out as free ions, and then just mix and match.1552

Really, though, you're not even mixing and matching; you're just switching partners.1559

OK, let's see; let's do one and see what happens here.1563

Our second example is going to be: let's mix a solution (well, let's actually spell out the problem); what happens when a solution of magnesium nitrate (let me write it over here--I don't have enough room), which is Mg(NO3)2, and potassium iodide, are mixed?1577

OK, well, let's write out the molecular formula and see what we get.1620

That is Mg(NO3)2 + KI (we want to write the molecular formula, so all we do is switch partners) magnesium and iodide; they come together, and it's going to be MgI2, and then potassium and nitrate are going to come together, so it's going to be KNO3.1626

Now, I want to balance this; that is the first thing I do, once I come up with a molecular by switching partners.1650

There is 1 magnesium, 1 magnesium; there are 2 nitrates--there is 1 nitrate here, so I put a 2, which gives me 2 potassiums; I stick a 2 in front of the K on the left; that gives me two potassium.1656

That gives me 2 iodide, but there is 2 iodide on the right; I'm fully balanced--all I needed was a 2 here and a 2 there.1668

Oops, some random lines are showing up again, which we do not want.1675

Let me put that Mg there; so now, let's check for solubility, because that is what we do next.1683

Well, magnesium nitrate is fully soluble; therefore, we have Mg2+ + 2NO3-.1690

Potassium iodide: fully soluble; therefore, we have 2 potassium ions and 2 iodide ions.1703

Now, let's look over here--magnesium iodide; it turns out magnesium iodide is also soluble; therefore, you have magnesium 2+, plus 2I-.1711

And potassium nitrate: definitely soluble; so 2K+ + 2NO3-.1723

So, notice, in this case: you ended up with--you mixed two solutions, where both things are fully soluble; and when you switched partners, you ended up with, well, something where both of the products are actually also soluble.1732

So, what is going on here--even though you can write a molecular formula and switch partners, there is no chemistry going on here.1747

All that is really happening is a bunch of free ions floating around, slamming into each other, and then coming apart again, because everything is soluble; every combination is soluble.1753

Really, there is no reaction here, even though we can write a molecular formula.1762

That is what is really important to remember: just because you can write a chemical equation doesn't mean there is any chemistry going on here.1767

Chemistry means something changes; magnesium, nitrate, potassium, iodide on the left, floating around, free in solution; magnesium, nitrate, potassium, and iodide, floating around in solution on the right; there is no difference--everything cancels; nothing happens here.1772

There is no net ionic, because nothing actually happens.1787

That is the important lesson.1791

So, take a look at your two ionic compounds; write a molecular formula on the left; switch partners, and write the appropriate chemical formula on the right.1795

Then, go back to the reactants and do the total ionic; see which one of those four compounds that you have is soluble, and which is not.1805

That will be your total ionic equation.1814

Then, you cancel similar species in the total ionic that are on the left and on the right, like we did for both of these examples; and what you are left with is your net ionic equation.1818

Your net ionic equation is the chemistry.1830

Here, in the last example, there is no net ionic; there is no chemistry; so, just because we can write it doesn't mean that something is actually happening.1833

These things don't really exist as compounds; they are soluble; they're just floating around, free, in solution.1841

This, the total ionic, is what is telling you everything that is going on; nothing is hidden--this tells you that nothing is going on.1848

Again, your intuition should be just as good; you know that free ions floating around--they slam into each other; there is nothing here that is mysterious--you just have to follow the logic.1857

OK, so thank you for joining us here at Educator.com for the discussion of precipitation reactions.1867

We'll see you next time; take care.1872

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