Chemistry I Tutorials with Worksheets

Video Lectures

Displaying all 176 video lectures.
Lecture 1
Understanding Matter
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Understanding Matter
Chemistry is the study of matter and the changes it undergoes, with the atom being its basic functional unit.

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Lecture 2
Classifications of Matter
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Classifications of Matter
Matter is anything that takes up space and has mass.

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Lecture 3
Volume, compressibility and viscosity of phases of matter
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Volume, compressibility and viscosity of phases of matter
Under appropriate conditions of pressure and temperature, most substances can exist in 3 states of matter.

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Lecture 4
Homogeneous vs Heterogeneous Mixtures
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Homogeneous vs Heterogeneous Mixtures
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Lecture 5
Physical vs Chemical Changes
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Physical vs Chemical Changes
A physical change involves a change in the phases of matter, whereas a chemical change involves a change in chemical bonds.

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Lecture 6
Understanding Scientific Notation
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Understanding Scientific Notation
Scientific notation is used to turn small or large inconvenient numbers into manageable ones. A scientific notation value is comprised of 3 parts: the coefficient, base and exponent.

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Lecture 7
Precision vs Accuracy
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Precision vs Accuracy
Within every calculation made there is a level of error. Tied to this notion of error are the two major principles of precision and accuracy.

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Lecture 8
Understanding Standard Deviation
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Understanding Standard Deviation
Standard deviation measures how close our numerical data points are to one another and to the “true” value.

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Lecture 9
Understanding Metric Prefixes
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Understanding Metric Prefixes
Metric Prefixes are “labels” that can be placed in front of base units. Whenever using the metric prefix chart, remember that “1” is associated with the metric prefix.

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Lecture 10
Using the Metric Prefix Chart
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Using the Metric Prefix Chart
Whenever using the metric prefix chart, remember that “1” is associated with the metric prefix.

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Lecture 11
Determining significant figures
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Determining significant figures
Significant Figures are used to determine some level of accuracy within our recorded measurements.

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Lecture 12
Understanding calculations with Significant Figures
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Understanding calculations with Significant Figures
The number of significant figures involved in a calculation depends on whether we are adding, subtracting, multiplying or dividing.

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Lecture 13
Understanding Temperature and Heat
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Understanding Temperature and Heat
Temperature is a measure of the amount of average kinetic energy of the atoms in a compound.

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Lecture 14
Kelvins, Celsius and Fahrenheit
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Kelvins, Celsius and Fahrenheit
Although the SI unit for temperature is Kelvin you will required to know how to convert between Celsius and Fahrenheit as well.

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Lecture 15
Understanding Unit Conversions
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Understanding Unit Conversions
We use dimensional analysis as a fail proof process to convert from one unit to another.

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Lecture 16
Examining Dimensional Analysis
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Examining Dimensional Analysis
Dimensional Analysis is just simply a word problem dealing with unit conversions. It is design to begin with your given value and to end with the unknown value we need to find.

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Lecture 17
Understanding Density
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Understanding Density
Density establishes a relationship between mass and volume. Density questions can include dimensional analysis so remember to use the same approach as you would normally.

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Lecture 18
Classification of Elements in the Periodic Table
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Classification of Elements in the Periodic Table
Ever wonder where did this periodic table ever come from? At the end of the 18th century, Lavoisier compiled a list of the 23 elements known at the time. In 1869, Dmitri Mendeleev coined the term "Periodic Table".

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Lecture 19
The Groups of the Periodic Table
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The Groups of the Periodic Table
Different columns known as groups separate the elements in the periodic table. Each of these groups has a unique name used in their classification.

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Lecture 20
Phases of the Elements in the Periodic Table
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Phases of the Elements in the Periodic Table
At room temperature (between 20°C to 25°C), all elements exist as solids, liquids or gases.

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Lecture 21
The Reason for Negative and Positive Ions
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The Reason for Negative and Positive Ions
Elements gain positive and negative charge when they lose or accept electrons. They gain and lose electrons so that they have the same number of electrons as the noble gases, which are electronically perfect.

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Lecture 22
Cations as Type I Metals vs Type II Metals
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Cations as Type I Metals vs Type II Metals
Type I metals are metals that possess one charge, whereas type II metals possess more than one charge.

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Lecture 23
Nonmetals and anion formation
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Nonmetals and anion formation
Unlike metals, nonmetals will under normal conditions gain electrons to become more like their closest noble gas.

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Lecture 24
Understanding and Recognizing Element Symbols
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Understanding and Recognizing Element Symbols
Many of the symbols used for elements are easy to recognize. However, some elements in the periodic table have Latin, Greek and other origins and so their symbol may not match their name as clearly.

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Lecture 25
Discussing the Monoatomic Elements
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Discussing the Monoatomic Elements
Many elements in the periodic table exist naturally by themselves, such as sodium metal, Na (s).

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Lecture 26
Discussing the Diatomic and Polyatomic Elements
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Discussing the Diatomic and Polyatomic Elements
Some elements in the periodic table exist in pairs called diatomic elements. Polyatomic elements exist as more than a pair of elements together, such as phosphorus which exists as P4.

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Lecture 27
Understanding Atomic Mass & Isotopes
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Understanding Atomic Mass & Isotopes
Under mass conversions, connections between grams, moles and molecules are established.

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Lecture 28
Calculating the average mass of an element
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Calculating the average mass of an element
Isotopes are elements with the same number of protons, but different number of neutrons.

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Lecture 29
Avogadro’s Number, Moles, Molecules & Atoms
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Avogadro’s Number, Moles, Molecules & Atoms
The moles of a compound or element serve as the bridge between grams, molecules and atoms.

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Lecture 30
Atoms to Moles to Grams
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Atoms to Moles to Grams
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Lecture 31
Molecules to Moles to Grams
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Molecules to Moles to Grams
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Lecture 32
Discussing Atomic Structure
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Discussing Atomic Structure
The atomic number represents the identity of the element and reveals the number of protons, whereas the atomic mass reveals the number of protons and neutrons. The atom is the fundamental unit of an element and it’s composed of the 3 subatomic particles.

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Lecture 33
Discussing Atomic Number & Atomic Mass
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Discussing Atomic Number & Atomic Mass
The number of protons gives the identity of an element. Ions arise when there is an unequal number of protons and electrons.

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Lecture 34
The Conservation of Matter
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The Conservation of Matter
Modern Atomic Theory states that matter is composed of small, indivisible particles called atoms. The First Law of Thermodynamics states that matter can neither be created nor destroyed, all that happens is that it changes forms.

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Lecture 35
Law of Definite Proportions
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Law of Definite Proportions
No matter where you obtain a compound, whether it’s from a lab experiment or from collection out in the field, the ratio of elements in it will remain constant.

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Lecture 36
Law of Multiple Proportions
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Law of Multiple Proportions
When element A and element B combine they can form different compounds in different ratios to one another. Dividing these different ratios should generate whole number answers.

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Lecture 37
Explanation for Bond Formation
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Explanation for Bond Formation
A chemical bond is the connection of different elements through the sharing, transferring or pooling of electrons.

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Lecture 38
Understanding an Ionic Bond
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Understanding an Ionic Bond
An ionic bond involves the transferring of an electron from one element to another element.

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Lecture 39
Understanding a Covalent Bond
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Understanding a Covalent Bond
A covalent bond involves the sharing of electron(s) between elements.
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Lecture 40
Understanding a Metallic Bond
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Understanding a Metallic Bond
A metallic "bond" isn't really a bond. Think of electrons as being loosely connected or "pooling" on the surface of many metals.

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Lecture 41
Understanding the Ionic Bonding Model
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Understanding the Ionic Bonding Model
An ionic bond forms when a cation and anion combine because of their opposite charges.

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Lecture 42
Understanding the Polyatomic Ions
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Understanding the Polyatomic Ions
Polyatomic Ions are compounds composed of different elements. All polyatomic ions have a positive or negative charge.

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Lecture 43
The 4 Characteristics of Polyatomic Ions with Halogens
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The 4 Characteristics of Polyatomic Ions with Halogens
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Lecture 44
Naming System for Polyatomic Ions with Halogens
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Naming System for Polyatomic Ions with Halogens
The naming system for polyatomic ions with halogens is based on the number of oxygens present.
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Lecture 45
Understanding how Ionic Compounds are formed
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Understanding how Ionic Compounds are formed
Nomenclature is the system used in chemistry for the naming of different compounds and molecules. Metals lose electrons to be positive ions, whereas nonmetals gain electrons to become negative ions.

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Lecture 46
Rules for Naming a Binary Ionic Compound
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Rules for Naming a Binary Ionic Compound
An ionic compound involves the bond between a cation (positive ion) and an anion (negative ion).

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Lecture 47
Understanding the Base Name of the nonmetal
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Understanding the Base Name of the nonmetal
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Lecture 48
Rules for naming Ionic Compounds with polyatomics
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Rules for naming Ionic Compounds with polyatomics
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Lecture 49
Rules for Naming Ionic Hydrates
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Rules for Naming Ionic Hydrates
An ionic hydrate is when an ionic compound is bonded to one or more molecules of water.
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Lecture 50
Rules for Naming Molecular Compounds
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Rules for Naming Molecular Compounds
A molecular compound represents a covalent compound, in which nonmetals are connected together.

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Lecture 51
Rules for Naming Binary Acids
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Rules for Naming Binary Acids
A binary acid is a covalent compound that contains an H+ ion connected to negative ion.

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Lecture 52
Rules for Naming Oxyacids
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Rules for Naming Oxyacids
An oxyacid is a covalent compound that contains an H+ ion connected to a polyatomic ion containing oxygen.
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Lecture 53
Empirical Formula vs Molecular Formula
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Empirical Formula vs Molecular Formula
The empirical formula represents the simplest formula for a compound.
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Lecture 54
Understanding Combustion Analysis
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Understanding Combustion Analysis
Combustion analysis is a technique where an organic compound is combusted (exploded) and the products are quantitatively measured. In a combustion reaction the organic compound is combusted to form carbon dioxide and water as products.
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Lecture 55
Understanding the difference between Atomic Mass & Molecular Mass
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Understanding the difference between Atomic Mass & Molecular Mass
Mass Percent is the percentage of a particular element within a compound. In order to find the atomic mass of an element you must look it up on your periodic table.
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Lecture 56
Understanding Mass Percent
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Understanding Mass Percent
In order to calculate the molecular mass or weight of a compound you must look up the atomic mass or weight of each element.
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Lecture 57
Balancing Chemical Reactions
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Balancing Chemical Reactions
When balancing an equation always make sure the types and number of atoms on both sides of the arrow are equal. Balancing a chemical reaction involves following the Law of Conservation of Mass and Energy. The total number of elements must be the same on both sides.

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Lecture 58
Understanding the Types of Functional Groups
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Understanding the Types of Functional Groups
The reactivity of an organic compound is based on the presence of its Functional Group.

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Lecture 59
Defining Stoichiometry
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Defining Stoichiometry
We use stoichiometry when given the KNOWN quantity of a compound or element and asked to find the UNKNOWN quantity of another compound or element in a balanced chemical equation.

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Lecture 60
The Stoichiometric Chart
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The Stoichiometric Chart
Remember whenever you have to find the unknown quantity of a compound or element just start with your given quantity and use the stoichiometric chart for guidance.

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Lecture 61
Understanding the limiting reactant or reagent
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Understanding the limiting reactant or reagent
The limiting reactant or limiting reagent is the compound that determines the maximum amount of product obtained in a chemical reaction.

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Lecture 62
Understanding Percent Yield
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Understanding Percent Yield
The percent yield determines how efficiently your reaction was in producing products. The higher the percent yield then the more successful the chemical reaction.

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Lecture 63
Definition of Molarity
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Definition of Molarity
Ever wonder how a mass amount such as moles can be converted into the volume amount of liters? Well, molarity serves as the bridge between moles and liters.

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Lecture 64
The Difference Between a Solute and Solvent
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The Difference Between a Solute and Solvent
In a homogeneous mixture, the smaller amount is the solute, and the larger amount is the solvent. When dissolving a solute into a solvent they make a solution.

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Lecture 65
Molarity and Stoichiometry
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Molarity and Stoichiometry
Stoichiometry is used when given the known quantity of one compound and asked the find the unknown quantity of another compound or element. Now we throw molarity into the mix.

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Lecture 66
Stoichiometric Chart and Molarity
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Stoichiometric Chart and Molarity
This new Stoichiometric Chart is used anytime we are given the known quantity of a compound in units such as mL, L or M and asked to find the unknown quantity of another compound.

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Lecture 67
The Rules for the Solubility of a compound
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The Rules for the Solubility of a compound
Whether you use the term molarity, concentration or solubility they all represent the maximum amount of solute that can successfully be dissolved in a solvent.

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Lecture 68
The meaning of the terms soluble and aqueous
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The meaning of the terms soluble and aqueous
Saying a compound is soluble means that the compound will dissolve into aqueous ions in a solvent. These ions can conduct electricity and are called electrolytes.

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Lecture 69
The meaning of the terms insoluble and precipitate
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The meaning of the terms insoluble and precipitate
Saying a compound is insoluble means that the compound will not dissolve into aqueous ions and forms a solid called a precipitate, which will not conduct electricity.

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Lecture 70
The Classification of Electrolytes
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The Classification of Electrolytes
Whenever you add a solute to a solvent there are three possible outcomes: the solute dissolves completely, dissolves partially or doesn't dissolve at all into ions.
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Lecture 71
The Strong Electrolytes
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The Strong Electrolytes
Strong Electrolytes break up into ions completely and are broken down into 3 categories: Strong Acids, Strong Bases and Soluble Ionic Compounds.

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Lecture 72
The Weak Electrolytes
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The Weak Electrolytes
Weak Electrolytes partially break up into ions and are broken down into 2 categories: Weak Acids and Weak Bases.

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Lecture 73
The Non Electrolytes
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The Non Electrolytes
Non-Electrolytes do NOT break up into ions at all and are referred to as molecular compounds. Now that we have a more firm grasp on the types of electrolytes, we can now apply this knowledge to electrolyte based questions.

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Lecture 74
Writing Chemical Reactions
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Writing Chemical Reactions
A balanced molecular equation is an equation showing the neutral forms of compounds. All that happens is that soluble ionic compounds separate and become ions in the chemical equation. This represents the total or full ionic equation.

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Lecture 75
Balanced Molecular Equation
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Balanced Molecular Equation
In a balanced molecular equation reactant compounds exist as ions that recombine to form products.

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Lecture 76
Total Ionic Equation
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Total Ionic Equation
In a total ionic equation only the compounds that are aqueous can break up into ions.

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Lecture 77
Net Ionic Equation
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Net Ionic Equation
In a net ionic equation the spectator ions are removed and what remains is the net ionic equation. Remember when asked to write a complete and balanced molecular equation that a reaction only occurs if a solid, liquid or gas is formed as a product.

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Lecture 78
Understanding Oxidation versus Reduction
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Understanding Oxidation versus Reduction
Oxidation and reduction deals with the transferring of electrons between reactants. The reactant that loses electrons is oxidized, while the reactant that gains electrons is reduced.

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Lecture 79
Rules for Oxidation Numbers
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Rules for Oxidation Numbers
We use these oxidation number rules whenever we are asked to determine the oxidation number of an entire compound or the individual elements within a compound.

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Lecture 80
Determining Oxidation Numbers
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Determining Oxidation Numbers
Using your newly acquired rules for oxidation numbers you can now determine the oxidation number of any element.
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Lecture 81
Pressure vs Force
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Pressure vs Force
Pressure is defined as the force exerted per unit of surface area.

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Lecture 82
The different units for Pressure
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The different units for Pressure
Although the SI unit for pressure is the Pascal, most professors use atm, mmHg or torrs as the everyday units for pressure. The conversion of units of pressure may sometimes be required in order to answer a quantitative question.

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Lecture 83
Understanding Dalton's Law
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Understanding Dalton's Law
According to Dalton’s Law, the total pressure inside a container is obtained by adding all the partial pressures of each non-reacting gas.

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Lecture 84
Understanding mole fraction and partial pressure
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Understanding mole fraction and partial pressure
The partial pressure of a gas is based on the moles of that gas. The relationship of the different moles of gases helps us determine the mole fraction of each gas.

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Lecture 85
Boyle's Law
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Boyle's Law
Boyle’s Law states that pressure and volume are inversely proportionally, which basically means they are opposites, at constant moles and temperature.

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Lecture 86
Charles Law
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Charles Law
Charles Law states that volume and temperature are directly proportionally if pressure and moles are constant. So if one is high then the other is high and vice versa.

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Lecture 87
Avogadro's Law
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Avogadro's Law
Avogadro’s Law states that volume and moles are directly proportionally at constant pressure and temperature. So if one is high then the other is high and vice versa.

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Lecture 88
Understanding PV = nRT, the Ideal Gas Law
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Understanding PV = nRT, the Ideal Gas Law
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Lecture 89
Gay Lussac's Law
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Gay Lussac's Law
Gay-Lussac’s Law states that pressure and temperature are directly proportionally at constant volume and moles. So if one is high then the other is high and vice versa.

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Lecture 90
Ideal Gas Stoichiometry
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Ideal Gas Stoichiometry
We’ve dealt with gases on previous topics such as stoichiometry, where the gas was either a reactant or product. Now we can relate Stoichiometry to the Ideal Gas Law.

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Lecture 91
Understanding Standard State Conditions
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Understanding Standard State Conditions
In order to accurately study the effect that changes in pressure, temperature and moles have on volume, chemists will often run their experiments under Standard State conditions.

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Lecture 92
Standard Molar Volume
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Standard Molar Volume
At STP, 1 mole of any ideal gas will occupy exactly 22.4 Liters of volume inside a container. Knowing the meaning of STP allows you to now further manipulate the Ideal Gas Law.

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Lecture 93
Combined Gas Law
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Combined Gas Law
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Lecture 94
Using the Ideal Gas Law to find density or molar mass
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Using the Ideal Gas Law to find density or molar mass
The Ideal Gas Law can be further manipulated into new equations. These new equations can help us find the density or molar mass of a gas.

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Lecture 95
Understanding kinetic energy & Root Mean Square Speed
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Understanding kinetic energy & Root Mean Square Speed
Increasing the temperature allows a gas to absorb thermal energy and convert into kinetic energy. Kinetic energy allows the gas to move and the speed at which it moves gives us the root mean square speed.

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Lecture 96
Understanding Graham's Law of Effusion
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Understanding Graham's Law of Effusion
Effusion is the rate at which a gas escapes a container with an opening in it. When comparing the rate or speed of two gases then we must use Graham’s Law of Effusion.

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Lecture 97
Understanding the Kinetic Molecular Theory
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Understanding the Kinetic Molecular Theory
A gas is seen as a collection of molecules or individual atoms that are in constant motion. The Kinetic-Molecular Theory tries to use data of real gases to predict how an ideal gas would behave if they existed.

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Lecture 98
The Characteristics of an Ideal Gas
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The Characteristics of an Ideal Gas
Ideal gases are imaginary and don’t actually exist. However, if they did exist then there would be certain characteristics that a gas would need to possess in order to be an ideal gas.

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Lecture 99
Thermodynamics vs Thermochemistry
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Thermodynamics vs Thermochemistry
Our system is simply our chemical reaction. Anything outside of that is considered our surroundings.

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Lecture 100
System vs Surroundings
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System vs Surroundings
Thermodynamics deals with the conversion of energy from one form to another. Thermochemistry is the branch of chemistry dealing with thermal (heat) energy.

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Lecture 101
Understanding the internal energy of the system
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Understanding the internal energy of the system
The internal energy (ΔE or ΔU) of the system can be calculated from the heat and work of the system.

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Lecture 102
Heat vs. Work
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Heat vs. Work
The signs of heat (q) and work (w) of the system can be either negative or positive depending on the key words stated.

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Lecture 103
Calculating work
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Calculating work
Work is one key variable to find the internal energy of the system. It’s equals to – PΔV.

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Lecture 104
Calculating heat and the internal energy of the system
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Calculating heat and the internal energy of the system
Once we’ve calculated work we can calculate the internal energy of the system once we also calculate the heat released or absorbed. Under certain conditions either q (heat) or w (work) can be equal to zero. This makes it easier to calculate the internal energy of the system.

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Lecture 105
Calculating the internal energy of the system in a vacuum
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Calculating the internal energy of the system in a vacuum
When work is done against a vacuum the pressure is equal to 0 atm. Since ΔE = q + w, the equation becomes only ΔE = q.

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Lecture 106
Heat capacity vs Specific Heat Capacity
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Heat capacity vs Specific Heat Capacity
By using calorimetry we can measure the thermal energy of a reaction (system) by measuring the change in heat of the surroundings.

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Lecture 107
Bomb Calorimeter
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Bomb Calorimeter
A bomb calorimeter measures the amount of calories within a substance through combustion. In other words, we “blow it up” and measure the amount of heat it releases.

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Lecture 108
Coffee Cup Calorimeter
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Coffee Cup Calorimeter
The enthalpy or heat of a reaction can be calculated through the use of a coffee cup calorimeter.

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Lecture 109
Calculating the Heat of the Solution
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Calculating the Heat of the Solution
The heat of the solution can be determined by first determining the heat of water.

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Lecture 110
Calculating the enthalpy of the reaction
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Calculating the enthalpy of the reaction
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Lecture 111
Understanding Thermochemical Equations
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Understanding Thermochemical Equations
A thermochemical equation is a stoichiometric question that now involves the variable of ΔHReaction. Instead of doing a mole to mole comparison we will now do a mole to ΔH comparison.

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Lecture 112
Understanding Hess's Law
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Understanding Hess's Law
A chemical reaction usually doesn’t happen in one step, but through a series of steps. To determine the overall energy or heat involved we use Hess’s Law.

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Lecture 113
Using Hess's Law
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Using Hess's Law
To determine the overall energy or heat of a reaction look to isolate each compound given within that chemical reaction.

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Lecture 114
Understanding a Formation Equation
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Understanding a Formation Equation
In a formation equation one mole of product is formed from the standard or elemental forms of each element.

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Lecture 115
Understanding Heats of Formation
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Understanding Heats of Formation
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Lecture 116
The Electromagnetic Spectrum
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The Electromagnetic Spectrum
Wavelength is the distance from one crest of a wave to another, whereas frequency is the number of waves within a second. The electromagnetic spectrum consists of varying forms of energy within oscillating electric and magnetic fields.

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Lecture 117
Wavelength & Frequency
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Wavelength & Frequency
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Lecture 118
The relationship between Wavelength & Frequency
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The relationship between Wavelength & Frequency
Wavelength and Frequency are inversely proportionally, meaning that if one increases then the other one must be decrease.

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Lecture 119
Converting between Wavelength & Frequency
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Converting between Wavelength & Frequency
To convert between wavelength & frequency we use the equation c = ν λ, where c equals the speed of light.

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Lecture 120
Understanding Refraction
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Understanding Refraction
Both energy & matter have different properties, including the way they travel and interact with their surroundings. In refraction, a particle changes direction as it moves from one medium to another. Energy (light) and matter refract differently.

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Lecture 121
Understanding Diffraction
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Understanding Diffraction
In diffraction, a particle bends around an object it encounters. Energy (light) and matter will experience this phenomenon differently.

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Lecture 122
Understanding the difference between Constructive and Destructive Interference
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Understanding the difference between Constructive and Destructive Interference
In constructive interference, waves that are in phase will combine together and increase the amplitude. In destructive interference, waves are out of phase and cancel one another out.

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Lecture 123
Understanding Planck's constant
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Understanding Planck's constant
Energy can be thought of as small "bundles or packets" of electromagnetic radiation. To determine this energy we use Planck's constant (h), 6.626 x 10^–34 J * s.
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Lecture 124
Understanding the Photoelectric Effect
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Understanding the Photoelectric Effect
The Photoelectric Effect was theorized by Albert Einstein to help explain what would happen to the electrons on the surface of a metal when a photon (light particle) with enough energy struck.

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Lecture 125
Illustrating what happens when a photon strikes the surface of a metal
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Illustrating what happens when a photon strikes the surface of a metal
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Lecture 126
The De Broglie Wavelength Equation
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The De Broglie Wavelength Equation
The De Broglie Wavelength equation relates wavelength to velocity or speed.

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Lecture 127
Understanding the Atom
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Understanding the Atom
Electrons can move up and down to different orbitals or shells based on absorbing or releasing of energy. Bohr's Model tries to explain what happens to the electron when it absorbs or emits energy.

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Lecture 128
Absorption vs Emission
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Absorption vs Emission
In absorption, an electron gains energy and becomes excited. In this excited state, the electron moves to higher energy level. In emission, the electron releases its excess energy to go down to a lower energy level.

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Lecture 129
Understanding Absorption & Emission
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Understanding Absorption & Emission
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Lecture 130
Different Emission Series
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Different Emission Series
We know that emission is the releasing of energy, but different energies are released depending on which shell the electron falls. These different energies appear in different places on the electromagnetic spectrum.

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Lecture 131
The Quantum Mechanical Picture of the Atom
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The Quantum Mechanical Picture of the Atom
The 4 Quantum Numbers provide us the coordinates to find the theoretical location of an electron.

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Lecture 132
The Quantum Sublevels
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The Quantum Sublevels
The main atomic sub-levels are the s, p, d and f. The sublevels have a set number of electron orbitals, each which can hold two electrons.

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Lecture 133
The Principal Quantum Number (n)
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The Principal Quantum Number (n)
The principal quantum number tells us the size and energy of an electron orbital.

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Lecture 134
Electron Capacity
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Electron Capacity
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Lecture 135
The Angular Momentum Quantum Number (l)
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The Angular Momentum Quantum Number (l)
The angular momentum quantum number, also known as the azimuthal quantum number, tells us the shape of the electron orbitals.

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Lecture 136
The Magnetic Quantum Number (mL)
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The Magnetic Quantum Number (mL)
The magnetic quantum number deals with the orientation of the orbital in the space around the nucleus.

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Lecture 137
Understanding Periodic Law
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Understanding Periodic Law
The electron configuration of an element is the distribution of its electrons within atomic orbitals. The arrangement of elements within the periodic table is based on increasing atomic mass.

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Lecture 138
The Auf Bau Principle
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The Auf Bau Principle
According to the Auf Bau Principle, electrons fill in lower energy orbitals before moving to higher energy orbitals.

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Lecture 139
The s,p, d, f sublevels and the Periodic Table
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The s,p, d, f sublevels and the Periodic Table
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Lecture 140
The Electron Configuration of Fluorine
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The Electron Configuration of Fluorine
By using the Periodic Table, you can determine the electron configuration of fluorine.

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Lecture 141
Hund's Rule
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Hund's Rule
According to Hund's Rule, electrons of similar energy will first half fill their orbitals before completely filling them.

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Lecture 142
Understanding Paramagnetism & Diamagnetism
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Understanding Paramagnetism & Diamagnetism
Whether an element or ion is paramagnetic or diamagnetic depends if there are unpaired electrons.

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Lecture 143
Inner Core vs Valence Electrons
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Inner Core vs Valence Electrons
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Lecture 144
The Spin Quantum Number
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The Spin Quantum Number
The electron configuration gives the distribution of electrons within an element, while the quantum number gives the probable location of those electrons.
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Lecture 145
The Pauli Exclusion Principle
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The Pauli Exclusion Principle
Within an electron orbital there are two electrons. One spins up and the other spins down.
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Lecture 146
Revisiting the Quantum Mechanical Picture of the Atom
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Revisiting the Quantum Mechanical Picture of the Atom
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Lecture 147
Quantum Numbers and Atomic Orbital Shape
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Quantum Numbers and Atomic Orbital Shape
Each atomic orbital has designated shapes that are determined by the angular momentum quantum number (l).

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Lecture 148
Understanding Atomic Radius
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Understanding Atomic Radius
Periodic Trends are patterns present in the periodic table based on an element's size and electron distribution. Atomic radius is defined as half the distance between the nuclei in a molecule of two identical elements.

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Lecture 149
Understanding Ionization Energy
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Understanding Ionization Energy
Ionization energy is the energy required to remove an electron from a gaseous element.

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Lecture 150
Exceptions to Ionization Energy in Groups 5A & 6A
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Exceptions to Ionization Energy in Groups 5A & 6A
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Lecture 151
Electron Affinity
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Electron Affinity
Electron affinity is the energy required to add 1 mole of electrons to a gaseous element.

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Lecture 152
Ionic Radius
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Ionic Radius
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Lecture 153
Exceptions to Ionization Energy in Groups 2A & 3A
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Exceptions to Ionization Energy in Groups 2A & 3A
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Lecture 154
Understanding Exergonic & Endergonic Reactions
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Understanding Exergonic & Endergonic Reactions
Lattice Energy is the energy required to combine gaseous ions into a solid, crystalline structure.

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Lecture 155
Understanding Lattice Energy
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Understanding Lattice Energy
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Lecture 156
Understanding Dipole Arrows
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Understanding Dipole Arrows
Dipole arrows are used anytime a molecule possesses a dipole moment, which happens when a molecule is polar. If a polar bond exists then dipole arrows must be used.

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Lecture 157
Dipole Arrows and Bond Classifications
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Dipole Arrows and Bond Classifications
Depending on the difference in electronegativity, a chemical bond fits within a certain category.

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Lecture 158
Understanding Electronegativity
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Understanding Electronegativity
Electronegativity measures how likely an element within a bond will attract electrons more or less to itself.

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Lecture 159
Understanding Lewis Dot Symbols
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Understanding Lewis Dot Symbols
Lewis Dot Structures or Electron Dot Structures are diagrams that show how elements in a molecule use their valence electrons to form bonds. For Main Group Elements, the number of valence electrons is based on the group number of that element.

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Lecture 160
Rules for Drawing Molecular Compounds
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Rules for Drawing Molecular Compounds
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Lecture 161
Understanding Incomplete Octets
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Understanding Incomplete Octets
An incomplete octet occurs when the central element in a Lewis Dot Structure is surrounded by less than 8 valence electrons.

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Lecture 162
Understanding the Expanded Octet
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Understanding the Expanded Octet
An expanded octet occurs when the central element in a Lewis Dot Structure is surrounded by more than 8 valence electrons.

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Lecture 163
Shortcut to Drawing Polyatomic Ions
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Shortcut to Drawing Polyatomic Ions
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Lecture 164
Why is formal charge important
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Why is formal charge important
Formal Charge is the charge given to an element within a molecule. It assumes that electrons are shared equally between elements, regardless of electronegativity. Many Lewis Dot Structures break the octet rule. Formal Charge is used to predict the most correct structure.

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Lecture 165
The Formal Charge Formula
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The Formal Charge Formula
Whenever using the formal charge formula the three acceptable answers are -1, 0 and +1.
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Lecture 166
Understanding Electronic Geometry
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Understanding Electronic Geometry
Whenever we draw a molecule we must take into account two different systems of geometrical shape. The electronic geometry of a compound treats surrounding elements and lone pairs on the central element as the same.

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Lecture 167
Understanding Molecular Geometry
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Understanding Molecular Geometry
The molecular geometry of a compound treats surrounding elements and lone pairs on the central as different.

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Lecture 168
VSEPR Model for Bonding
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VSEPR Model for Bonding
According to the VSEPR Model, bond angles result from surrounding elements and lone pairs around the central element positioning themselves at an optimal distance. The VSEPR theory gives bond angles based on the number of groups around the central element.

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Lecture 169
Ideal Bond Angles
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Ideal Bond Angles
The more lone pairs on the central element then the more compressed the bond angle, and the greater the deviation from an ideal bond angle.

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Lecture 170
Lone Pairs & Bond Angles
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Lone Pairs & Bond Angles
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Lecture 171
Understanding Molecular Polarity
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Understanding Molecular Polarity
Dipole moments or polarity happen in molecules when there is an unequal sharing of electrons.

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Lecture 172
The First Rule for Molecular Polarity
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The First Rule for Molecular Polarity
When the central element has no lone pairs and has the same surrounding elements then usually the molecule is nonpolar.

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Lecture 173
The Second Rule for Molecular Polarity
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The Second Rule for Molecular Polarity
When the central element has lone pairs then we must use element dipole arrows and lone pair dipole arrows to determine polarity.

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Lecture 174
Understanding Overlapping Atomic Orbitals
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Understanding Overlapping Atomic Orbitals
How would an s orbital combine with a p orbital to form a covalent bond? The answer is through hybridization where a hybrid oribital is formed. A sigma bond is formed by the direct end to end overlapping of atomic orbitals.

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Lecture 175
Understanding Hybrid Orbitals
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Understanding Hybrid Orbitals
For an s orbital to bond with a p orbital the central element must first undergo hybridization.

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Lecture 176
Simplifying Hybridization
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Simplifying Hybridization
To determine the hybridization of a central element we look at the number of groups around it.

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