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Summer Ebs

Summer Ebs

Chemical Bonding, Part 2

Slide Duration:

Table of Contents

I. Chemistry
Properties of Matter

30m 50s

Intro
0:00
Matter
0:07
Matter
0:08
Substance, Element, and Compound
0:47
Homogeneous and Heterogeneous Mixture
1:47
Suspension, Colloid, and Solution
3:16
Physical Properties
5:25
Appearance: Color, Shape, Size, Density, and State of Matter
5:26
Behavior: Viscosity, Magnetism, Malleability, and Ductility
8:00
Physical Changes
10:29
Physical Changes
10:30
Chemical Properties
14:38
Chemical Properties
14:39
Chemical Changes
16:35
Chemical Changes
16:36
Signs of Chemical Change
16:55
Example 1: Identify the Mixtures Listed
19:21
Example 2: Physical or Chemical Change?
23:38
Example 3: How Can You Separate a Mixture of Sand, Gravel, Iron, Filings, Salt, and Water?
25:04
Example 4: Physical/Chemical Property and Change
27:27
Energy

53m 22s

Intro
0:00
Energy
0:03
Energy Overview
0:04
Potential Energy
1:48
Potential Energy
1:49
Mechanical (Elastic) Potential Energy
1:54
Chemical Potential Energy
3:15
Nuclear Energy
4:06
Gravitational Potential Energy
4:43
Kinetic Energy
7:34
Kinetic Energy
7:35
Thermal Energy
8:03
Radiant Energy
8:57
Electrical Energy
9:47
Sound
10:17
Motion
10:54
Kinetic Energy: Example
11:31
Law of Conservation of Energy
12:47
Law of Conservation of Energy
12:48
Electrical to Radiant
13:21
Chemical to Thermal
14:34
Potential to Kinetic
15:10
Friction
18:48
Energy Resources
20:06
Nonrenewable: Fossil Fuels
20:51
Nonrenewable: Nuclear
21:56
Renewable: Solar
26:50
Renewable: Wind
29:22
Renewable: Tidal
31:10
Renewable: Hydroelectric
32:30
Renewable: Geothermal
35:24
Example 1: Gravitational Potential Energy
38:40
Example 2: Kinetic Energy
42:20
Example 3: Maximum and Minimum Potential and Kinetic Energy
44:48
Example 4: Should We Use Renewable or Nonrenewable Resources to Generate Electricity?
46:31
Heat and States of Matter

48m 48s

Intro
0:00
Temperature
0:04
Temperature
0:05
Fahrenheit to Celsius
2:15
Celsius to Fahrenheit
4:29
Kelvins to Celsius and Celsius to Kelvins
5:50
Thermal Energy
8:06
Thermal Energy, Kinetic Energy, and Potential Energy
8:07
Changing Thermal Energy: Temperature
9:11
Changing Thermal Energy: State of Matter
9:37
Changing Thermal Energy: Amount of Matter
10:12
Heat
10:59
Heat
11:00
Specific Heat
12:21
Transfer of Thermal Energy
15:15
Conduction
15:16
Convection
16:43
Radiation
19:57
States of Matter
20:43
Solids: Arrangement of Atoms, Shape, Volume, and Molecular Motion
21:35
Liquids: Arrangement of Atoms, Shape, Volume, and Molecular Motion
23:49
Gases: Arrangement of Atoms, Shape, Volume, and Molecular Motion
25:33
Plasma: Arrangement of Atoms, Shape, Volume, and Molecular Motion
27:02
Changing States of Matter
27:49
Melting
27:50
Freezing
28:15
Vaporization
29:04
Boiling
29:17
Condensation
31:21
Temperature and Time Graph
32:18
Thermal Expansion
36:19
Thermal Expansion of Solids
37:16
Thermal Expansion of Liquids
38:17
Thermal Expansion of Gases
39:46
Example 1: Converting Temperatures
40:28
Example 2: Thermal Energy
43:35
Example 3: Quick Matching
44:58
Example 4: Why Does It Feel Cold When You Put Your Hand On the Table?
45:50
Example 5: Heat Transfer
46:48
Example 6: Changing States of Matter
47:29
Atoms and Elements

30m 12s

Intro
0:00
Atoms
0:05
Atoms
0:06
Atomic Structure
1:01
Electron Cloud
1:02
Nucleus, Protons, and Neutrons
1:43
Quarks
2:07
Protons, Neutrons, Electrons
2:40
Protons, Neutrons, Electrons: Location
2:42
Protons, Neutrons, Electrons: Electric Charge
3:05
Examples
4:10
Electron Configuration
5:32
Electron Configuration
5:33
Elements
12:22
Atomic Number
13:05
Carbon
13:15
Oxygen
14:49
Important Elements for Living Things
16:25
Isotopes
17:04
Isotopes
17:05
Example 1: Atomic Structure and Electrical Charge
21:16
Example 2: Electron Configuration
23:13
Example 3: Electron Configuration
24:57
Example 4: Use the Periodic Table to Complete the Table Below
26:08
Periodic Table

47m 23s

Intro
0:00
Periodic Table
0:06
Atomic Number, Chemical Symbol, and Atomic Mass
0:07
Groups and Periods
4:14
Groups and Periods
4:15
Electron Dot Diagrams
10:05
Electron Dot Diagrams
10:06
Ion Formation
19:09
An Ion Forms When an Atom Gains or Loses Electrons
19:10
A Positive Ion Forms When an Atom Loses and Electron
20:25
A Negative Ion Forms When an Atom Gains an Electron
26:49
Oxidation Numbers
28:51
Oxidation Numbers
28:52
Metals, Nonmetals, Metalloids
34:52
Metals, Nonmetals, Metalloids
34:53
Example 1: Group and Period
37:39
Example 2: Electron Dot Diagrams
39:50
Example 3: How do Fluorine and Calcium Become Ions?
42:10
Example 4: What Are 2 Ways to Find the Oxidation Number of Sodium?
44:58
Chemical Bonding, Part I

51m 6s

Intro
0:00
Chemical Bonds Form Compounds
0:17
Atoms and Electrons
0:18
H2O
2:14
HCl
3:36
C6H12O6
4:16
Ca(NO3)2
5:06
Review: Dot Diagrams
7:10
Review: Ion Formation
8:30
Ionic Bond
9:57
Ionic Bond
9:58
Sodium and Fluorine
10:41
Magnesium and Chlorine
16:30
Covalent Bond
22:19
Covalent Bond
22:20
Hydrogen and Carbon
23:58
Hydrogen and Oxygen
27:28
Multiple Covalent Bonds
29:03
Single Covalent Bond
29:04
Double Covalent Bond
29:40
Triple Covalent Bond
31:50
Polar and Nonpolar Molecules
33:33
Polar Molecules
33:34
Unequal sharing of Electrons and Electronegativities
35:02
Nonpolar Molecules
37:46
Example 1: Elements and Atoms
38:42
Example 2: Dot Diagram of the Bond That Forms Between Magnesium and Oxygen
41:17
Example 3: Dot Diagram of the Bond That Forms Between Nitrogen and Oxygen
45:24
Example 4: Polar or Nonpolar?
47:22
Chemical Bonding, Part 2

56m 22s

Intro
0:00
Bonding Atoms Make Compounds
0:05
Binary Compounds
0:06
Reviwew: Oxidation Number
1:14
Naming Ionic Compounds
1:45
Naming Ionic Compounds
1:46
NaCl
2:26
MgCl2
5:04
Al2S3
6:52
Writing Formulas of Ionic Compounds
10:03
Writing Formulas of Ionic Compounds
10:04
Beryllium Fluoride
10:17
Lithium Nitride
12:24
Calcium Bromide
13:53
Polyatomic Ions
15:31
Polyatomic Ions
15:32
Ammonium Phosphate
17:21
Aluminum Hydroxide
19:37
Magnesium Chlorate
20:54
NaOH
21:47
(NH4)2O
22:17
Mg(NO3)2
22:56
Special Ions
23:28
Iron (III) Iodide
24:28
Lead (IV) Chloride
26:30
Chromium (III) Oxide
27:31
Fe3P2
29:18
CuI2
31:51
PbBr2
33:04
Naming Covalent Compounds
33:57
Naming Covalent Compounds
33:58
Examples
35:03
Ionic or Covalent?
39:50
Ionic vs. Covalent: Electron
39:51
Ionic vs. Covalent: State At Room Temperature
10:23
Ionic vs. Covalent: Metal, Nonmetal, Metalloids
41:02
Ionic vs. Covalent: Naming
41:35
Example 1: Write the Names or Formulas for Each Ionic Compound
42:50
Example 2: Write the Names or Formulas for Each Covalent Compound
46:13
Example 3: Name the Following Ionic Compounds
49:44
Example 4: Provide the Formulas for the Following Ionic Compounds
52:19
Example 5: Ionic or Covalent?
54:21
Chemical Reactions

49m 13s

Intro
0:00
Chemical Reactions
0:05
Chemical Reactions
0:06
Chemical Formula Example
0:54
Reactants and Products
3:50
Conservation of Mass
4:58
The Total Mass of the Reactant Must Equal the Total Mass of the Products
4:59
Balancing Chemical Equations
6:42
Balancing Equations
11:12
Example 1: Balancing Equations
11:27
Example 2: Balancing Equations
14:15
Example 3: Balancing Equations
16:28
Types of Reactions
19:17
Synthesis
19:18
Decomposition
20:09
Single-Displacement
20:54
Double-Displacement
22:12
Combustion
23:34
Energy in Chemical Reactions
24:41
Chemical Reactions and Activation Energy
24:42
Endergonic Reactions
25:55
Exergonic Reactions
27:51
Rate of Chemical Reactions
29:42
Rate of Chemical Reactions Overview
29:43
Temperature
30:51
Concentration
31:26
Agitation
32:08
Surface Area
32:29
Pressure
33:06
Catalysts and Inhibitors
33:18
Example 1: Translate Into Chemical Equations
34:32
Example 2: Law of Conservation of Mass
37:35
Example 3: Balance the Following Equations
40:33
Example 4: Math Each Equation With the Correct Type of Reaction
44:58
Example 5: Exothermic or Endothermic Reaction?
48:21
Solutions, Acids, and Bases

29m

Intro
0:00
Solutions
0:06
Definition of Solution
0:07
Solute and Solvent
0:26
Example: Salt Water
0:35
Example: Carbonated Water
1:03
Dissolving
1:49
Dissolving
1:50
Example: Liquid Dissolves a Solid at the Surface of the Solid
3:54
Aqueous Solutions: Water as Solvent
4:42
Increasing the Rate of Dissolving
5:33
Stir
5:34
Crush
6:37
Heat
7:36
Solubility
8:31
Definition of Solubility
8:32
Compare the Solubility of Sugar in Water vs. Salt in Water
8:44
Factors that Affect Solubility
11:45
Concentration
12:45
Concentration
12:46
pH Scale
15:21
pH Scale: Acids, Neutral, and Bases
15:22
Acids and Bases
18:01
Chemical Properties
18:02
Physical Properties
18:43
pH Scale
19:31
Examples of Acids and Bases
19:36
Acids and Bases React Together to Form Salt and Water
20:09
Example 1: Identify the Solutes and Solvents for the Following Solutions
21:26
Example 2: Temperature and the Rate of Dissolving/Solubility of a Solid
23:57
Example 3: How Can You Make a Solution Have a Higher Concentration?
25:44
Example 4: Acids and Bases
27:57
II. Physics
Waves

42m 35s

Intro
0:00
Waves
0:05
Introduction to Waves
0:06
Mechanical Waves
1:24
Electromagnetic Waves
1:50
Mechanical Waves
2:13
Transverse
2:14
Longitudinal (Compressional Waves)
4:00
Properties of Waves
7:26
Transverse and Compressional Waves: Wavelength
7:27
Transverse and Compressional Waves: Frequency (Hz)
9:32
Transverse and Compressional Waves: Amplitude
11:30
Wavelength and Frequency are Related
13:40
Wave Speeds
15:01
Wave Speeds
15:02
Behavior of Waves
18:06
Reflection
18:33
Refraction
22:42
Diffraction
24:25
Electromagnetic Waves
26:00
Electromagnetic Waves
26:01
Visible Light
30:49
Visible Light
30:50
Opaque
34:25
Translucent
34:54
Transparent
35:41
Example 1: Label the Transverse Wave
36:59
Example 2: Label the Compressional Wave
38:13
Example 3: What Happens to the Frequency of a Wave as the Wavelength Increases?
39:12
Example 4: Law of Reflection and Light Wave
40:48
Motion

37m 21s

Intro
0:00
Distance vs. Displacement
0:04
Distance
0:05
Displacement
0:49
Speed
4:47
Speed
4:48
Instantaneous Speed
6:14
Average Speed
6:40
Velocity
7:25
Distance-Time Graphs
8:21
Distance-Time Graphs
8:22
Acceleration
13:38
Acceleration Definition
13:39
Acceleration Equation
15:23
Positive Acceleration
18:43
Negative Acceleration
18:52
Speed-Time Graphs
20:56
Speed-Time Graphs
20:57
Example 1: Displacement, Distance, and Average Speed
25:15
Example 2: Velocities
28:02
Example 3: Acceleration
28:59
Example 4: Distance and Time
30:19
Example 5: Speed and Time
34:08
Forces

35m 3s

Intro
0:00
Force
0:04
Force Definition
0:05
Net Force
1:44
Balanced Forces
3:06
Unbalanced Forces
4:23
Forces Examples
5:09
Friction
7:53
Friction Definition
7:54
Static Friction
8:23
Sliding Friction
9:35
Rolling Friction
10:11
Fluid Friction
11:13
Air Resistance
12:10
Newton's Laws of Motion
14:06
First Law of Motion
14:07
Inertia
15:56
Newton's Laws Continued
17:13
Second Law of Motion
17:14
Third Law of Motion
18:35
Gravitational Force
24:17
Gravity and Gravitational Force
24:18
Example 1: Horizontal Force, Frictional Force, and Net Force
28:36
Example 2: Net Force and Acceleration
29:38
Example 3: Gravitational Force
30:35
Example 4: Force of Air Resistance and Net Force
32:32
Density & Buoyancy

23m 43s

Intro
0:00
Density
0:05
Definition of Density
0:06
Density = Mass / Volume
1:01
Density of Irregular Objects
3:58
Density of Irregular Objects
3:59
Buoyant Force
7:46
Buoyancy
7:47
Archimedes' Principle
9:23
Floating and Sinking
12:47
Floating and Sinking: Looking at Density
12:48
Example 1: Density of an Object
16:15
Example 2: Density of Yourself
17:28
Example 3: Using Archimedes' Principle to Predict If an Object Will Sink or Float in Water
19:38
Example 4: Will Aluminum, Gold, and Oil Float or Sink When Placed Into Water?
22:06
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Lecture Comments (21)

0 answers

Post by Amanda Rimeikis on February 20, 2017

Great job!!

0 answers

Post by aimun amatul-hayee on March 2, 2015

what is the charge of the transition metals?

0 answers

Post by sadia sarwar on November 15, 2014

why is N2 alone nitrogen when S4N2 is tetrasulfur dinitride?

1 answer

Last reply by: Professor Ebs
Wed Jan 8, 2014 12:06 AM

Post by Yisrael Harris on January 2, 2014

Questions about polyatomic ions:

1. Is there a name for the ionic compounds formed from polyatomic ions?
2. Such compounds are not binary compounds, are they?
3. I notice that most of the polyatomic ions on your list contain O. Do most polyatomic ions contain O?
4. Are there many other polyatomic ions besides the ones in your list?
5. Are there properties which make certain covalent compounds more likely than others to become polyatomic ions?
6. I notice that all the polyatomic ions on your list except ammonium are -ve. Are most polyatomic ions -ve?
7. I believe that the ionic compounds formed from polyatomic ions are not named exactly like other ionic compounds: the rule of changing the ending of negative ion to "-ide" doesn't seem to apply when the negative ion is a polyatomic ion.

1 answer

Last reply by: Professor Ebs
Wed Apr 24, 2013 11:25 AM

Post by Yisrael Harris on April 23, 2013

When ionic bonds form, what is it that causes the metal to lose the electron(s)? Is it the non-metal pulling that electron away? In other words, is the non-metal the direct cause of the metal losing its electron, or does the metal lose its electron independently of the non-metal, and then the non-metal takes that free electron?

1 answer

Last reply by: Professor Ebs
Wed Apr 24, 2013 11:22 AM

Post by CeCe Tang on March 24, 2013

On your periodic table, why is Group 13,2A? Shouldn't it be 3A?

1 answer

Last reply by: Professor Ebs
Fri Feb 22, 2013 4:49 PM

Post by Doreen Wagner on February 22, 2013

5. Example IV Lead (IV) oxide:
she wrote PbO4, but should it be Pb2O4, because Pb has a 4-/+and O is 2- charged.


1 answer

Last reply by: Professor Ebs
Fri Feb 22, 2013 4:44 PM

Post by Doreen Wagner on February 22, 2013

4. On Example III polyatomic compounds:
she wrote calcium phosphate as Ca(PO4)2, but wouldn't it actually be Ca3(PO4)2, because PO4 has a -3 charge and Ca has a 2+charge.

1 answer

Last reply by: Professor Ebs
Fri Feb 22, 2013 4:00 PM

Post by Doreen Wagner on February 22, 2013

3.same example: Disulfur dichloride... if you have S, your dot diagram would have 6 electrons in the dot diagram (so 2 empty spaces), then Cl would have 7 electrons in the dot diagram, and they would each have 1 electron to share. If S has 2 empty spaces for electrons, I can see that 2 Cl could have room to bond, but I don't see that you need another Sulfur, so how does it fit in to become Disulfur? Also, if S has a 2- charge and Cl has 1- charge, wouldn't the final chemical formula be SCl2?

1 answer

Last reply by: Professor Ebs
Fri Feb 22, 2013 5:02 PM

Post by Doreen Wagner on February 22, 2013

2. same example 2 (covalent compounds): Hexaboron monosilicide: B6Si... so if I do dot diagrams for Si and attach B so they share an electron, I can only attach 4 B, not 6. How would 6 Boron bond when Si has only 4 electrons to share?

1 answer

Last reply by: Professor Ebs
Fri Feb 22, 2013 4:30 PM

Post by Doreen Wagner on February 22, 2013

on Example 2, Covalent Compounds:
on the Iodine pentafluoride question, why isn't it Pentafluorine iodide because the rule says: "the element with the lower group # goes 1st in the name. If they are in the same group, the element with the higher period # goes 1st." (so, it's like reading... Left to right, top to bottom. or am I wrong on this...

because Fluorine is in period 2 and iodine is in period 5, would we say fluorine first, then iodine?

Also, since they are in the same group, wouldn't you have an overall charge of 6-, since you have 5 of the F (with 1- each) and 1 of I (with 1- each) so an overall charge of 6- ?
I'm confused on that...

1 answer

Last reply by: Professor Ebs
Thu Nov 29, 2012 3:47 PM

Post by Nathanael Shim on November 28, 2012

So gas is not a compound right?

Related Articles:

Chemical Bonding, Part 2

  • A binary compound is a compound made up of 2 elements.
  • Binary compounds can be ionic or covalent.
  • To name ionic compounds, write the name of the positive ion first; change the ending of the negative ion to –ide.
  • Ionic compounds must be neutral. The number of atoms in the formula (written as subscript) must be written so that the charges are balanced.
  • A covalently bonded group of atoms can lose or gain electrons to become a polyatomic ion.
  • Polyatomic ions have special names, but the ionic compounds that they form are named and written just like other ionic compounds.
  • You can use a numerical prefix for each type of element in a covalent compound, but you do not need to use mono for one. If there is no prefix, that means there is only one atom of that type of element.
  • Some transition metals have more than one oxidation number. When you write these names, you must use roman numerals to express which of the two oxidation numbers are being used.

Chemical Bonding, Part 2

When naming ionic compounds, the name of which ion goes first?
The positive ion is first
What is the name of the ionic compound that forms when Hydrogen (with an oxidation number of 1+) and Chlorine (with an oxidation number of 1-) bond?
Hydrogen chloride
What is the chemical formula of Magnesium Bromide? (Mg2+ and Br1−)
MgBr2
What is the name of an ion that is made up of more than one atom?
A polyatomic ion
What is the chemical formula for Magnesium Hydroxide (Mg2+ and OH1−)?
Mg(OH)2
What type of compound is diarsenic pentoxide? How do you know?
A covalent compound this is a bond between a metalloid and a nonmetal so it is covalent and the numerical prefixes are included in the names, which is another indication that it is a covalent compound.
What type of compound is copper (II) carbonate? How do you know?
An ionic compound Copper is a metal and carbon is a nonmetal, copper (II) has an oxidation number of 2+ and carbonate is a polyatomic ion with an oxidation number of 2-
What is the chemical formula for phosphorus trihydride?
PH3
What type of compounds are solids at room temperature and form when electrons are gained or lost?
Ionic compounds
What is the name for a compound that is made up of only two elements?
Binary compound

*These practice questions are only helpful when you work on them offline on a piece of paper and then use the solution steps function to check your answer.

Answer

Chemical Bonding, Part 2

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
  • Bonding Atoms Make Compounds 0:05
    • Binary Compounds
    • Reviwew: Oxidation Number
  • Naming Ionic Compounds 1:45
    • Naming Ionic Compounds
    • NaCl
    • MgCl2
    • Al2S3
  • Writing Formulas of Ionic Compounds 10:03
    • Writing Formulas of Ionic Compounds
    • Beryllium Fluoride
    • Lithium Nitride
    • Calcium Bromide
  • Polyatomic Ions 15:31
    • Polyatomic Ions
    • Ammonium Phosphate
    • Aluminum Hydroxide
    • Magnesium Chlorate
    • NaOH
    • (NH4)2O
    • Mg(NO3)2
  • Special Ions 23:28
    • Iron (III) Iodide
    • Lead (IV) Chloride
    • Chromium (III) Oxide
    • Fe3P2
    • CuI2
    • PbBr2
  • Naming Covalent Compounds 33:57
    • Naming Covalent Compounds
    • Examples
  • Ionic or Covalent? 39:50
    • Ionic vs. Covalent: Electron
    • Ionic vs. Covalent: State At Room Temperature
    • Ionic vs. Covalent: Metal, Nonmetal, Metalloids
    • Ionic vs. Covalent: Naming
  • Example 1: Write the Names or Formulas for Each Ionic Compound 42:50
  • Example 2: Write the Names or Formulas for Each Covalent Compound 46:13
  • Example 3: Name the Following Ionic Compounds 49:44
  • Example 4: Provide the Formulas for the Following Ionic Compounds 52:19
  • Example 5: Ionic or Covalent? 54:21
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