Chemistry II Tutorials with Worksheets

Video Lectures

Displaying all 96 video lectures.
Lecture 1
Understanding the difference between intermolecular and intramolecular forces
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Understanding the difference between intermolecular and intramolecular forces
Intermolecular forces influence the physical properties of compounds, whereas intramolecular forces influence the chemical properties of compounds.

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Lecture 2
The First and Strongest Intermolecular Force
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The First and Strongest Intermolecular Force
This intermolecular force deals with the attraction between an ion and a polar compound.

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Lecture 3
The Second Intermolecular Force
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The Second Intermolecular Force
This intermolecular forces exists anytime hydrogen is directly connected to fluorine, oxygen or nitrogen.

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Lecture 4
The Third Intermolecular Force
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The Third Intermolecular Force
This intermolecular force deals with the attraction between two polar covalent compounds.

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Lecture 5
The Fourth Intermolecular Force
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The Fourth Intermolecular Force
This intermolecular force deals with the attraction between a polar and nonpolar covalent compound.

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Lecture 6
The Fifth Intermolecular Force
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The Fifth Intermolecular Force
This intermolecular force is found in all compounds, but it's the main intermolecular force of nonpolar covalent compounds.
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Lecture 7
Understanding the Theory of
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Understanding the Theory of "Likes" dissolve "Likes"
Solubility deals with the dissolving of a solute in a solvent in order to create a solution. In order for a solvent to dissolve a solute both components have similar polarities.

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Lecture 8
The Properties of the 3 Phases of Matter
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The Properties of the 3 Phases of Matter
Phase Diagrams show what effect temperature and pressure have on a pure substance in a closed system without any air.

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Lecture 9
Understanding the Phase Diagram
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Understanding the Phase Diagram
The phase diagram looks at the transition of matter between the phases of solid, liquid and gas as temperature and pressure are affected.

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Lecture 10
Understanding the Clasius-Clapeyron Equation
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Understanding the Clasius-Clapeyron Equation
The Clasius-Clapeyron Equation establishes a quantitative relationship between vapor pressure and temperature.

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Lecture 11
Understanding Molecular Orbital Theory
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Understanding Molecular Orbital Theory
Molecular Orbital Theory allows us to predict the distribution of electrons within a molecule. This allows us to predict properties such as bond order, magnetism and shape.

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Lecture 12
Calculating Bond Order
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Calculating Bond Order
By calculating the bond order it is possible to determine the stability of the bonds within a molecule.

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Lecture 13
Solutions and Molarity
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Solutions and Molarity
Solutions are homogeneous mixtures that result from the dissolving of a solute by a solvent. Molarity represents the moles of solute dissolved per liters of solution.
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Lecture 14
The Theory of “Likes” dissolve “Likes”
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The Theory of “Likes” dissolve “Likes”
A pure solvent with similar polarity can successfully dissolve a solute to create a solution.
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Lecture 15
Molarity vs Molality
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Molarity vs Molality
As we already know MOLARITY represents the moles of solute per liters of solution. Now, we will deal with the new term of MOLALITY, which represents the moles of solute per kilograms of solvent.
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Lecture 16
Mass Percent & Mole Fraction
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Mass Percent & Mole Fraction
Molarity serves as the bridge between moles to liters and can also incorporate the concept of mass percent and mole fraction. Mass Percent is represented by the mass of the solute per mass of solution, which is the combined mass of the solute and solvent.
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Lecture 17
Saturated, Unsaturated and Supersaturated Solutions
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Saturated, Unsaturated and Supersaturated Solutions
When an ionic solid dissolves, ions leave the solid and become dispersed in the solvent. This solution can be classified as saturated, unsaturated and supersaturated.
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Lecture 18
Henry’s Law
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Henry’s Law
Henry’s Law relates the solubility of a gas over a solution based on the factors of temperature, volume and pressure.
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Lecture 19
The Colligative Properties
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The Colligative Properties
The 4 Colligative Properties help to explain what happens to a pure solvent as solute is added to it. The four colligative properties of Boiling Point, Freezing Point, Vapor Pressure and Osmotic Pressure will either increase or decrease the addition of solute to a pure solvent.
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Lecture 20
Defining Vapor Pressure
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Defining Vapor Pressure
Vapor Pressure is achieved by the equilibrium rate of vaporization equaling the rate of condensation.
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Lecture 21
Understanding Chemical Kinetics
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Understanding Chemical Kinetics
Chemical Kinetics is the study of reaction rates, which is the speed in which reactants are consumed to make products.
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Lecture 22
Kinetics in a Chemical Reaction
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Kinetics in a Chemical Reaction
The word “kinetics” is derived from the Greek word “kinesis”, which means motion. So Chemical Kinetics deals with the speed of motion experienced by reactants as they are allowed to react.
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Lecture 23
Slow vs Fast Reactions
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Slow vs Fast Reactions
Depending on certain conditions, a chemical reaction can either happen very quickly in seconds or take place over several years.
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Lecture 24
Kinetics & the Concentration of Compounds
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Kinetics & the Concentration of Compounds
In order for a chemical reaction to occur two molecules must collide. The more concentrated a solution then the greater the chance of them colliding.
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Lecture 25
Kinetics & the Surface Area of Compounds
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Kinetics & the Surface Area of Compounds
For a collision between molecules to be successful, molecules join at their active sites. The larger their surface area then the more places the molecules can successful join.
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Lecture 26
Kinetics & Temperature
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Kinetics & Temperature
The General Rule is increasing the reaction temperature by 10°C will cause the rate to double.
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Lecture 27
Kinetics & the Catalyst
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Kinetics & the Catalyst
A catalyst helps to speed up the rate of a reaction by lowering the energy of activation (Ea).
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Lecture 28
Understanding Average or General Rate
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Understanding Average or General Rate
General or Average Rate is the change in the concentration of a compound over a period of time.
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Lecture 29
Understanding the Rate Law
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Understanding the Rate Law
When it comes to the rate of a reaction the Rate Law focuses on the reactant concentrations, while ignoring the product concentrations.
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Lecture 30
The Rate Determining Step
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The Rate Determining Step
If the reaction says it’s a SLOW step or a ONE-STEP MECHANISM then we can simply look at the coefficients of the reactants to determine the reaction orders.
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Lecture 31
The Arrhenius Equation
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The Arrhenius Equation
Temperature is one of the four factors that directly affect the rate of a reaction. It affects the rate of a reaction by affecting the rate constant k. The Arrhenius Equation tries to explain how temperature plays a central role in affecting the rate constant k.
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Lecture 32
The other forms of the Arrhenius Equation
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The other forms of the Arrhenius Equation
When dealing with TWO rate constants or TWO temperatures then we use the Two–Point Form of the Arrhenius Equation. In order to relate the plot of a graph to the Arrhenius equation then we manipulate it into the plot­­–wise approach of the equation.
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Lecture 33
The Integrated Rate Laws
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The Integrated Rate Laws
When we include the variable of time to our Rate Law then we obtain the Integrated Rate Laws.
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Lecture 34
Zeroth, First and Second Order Reactions
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Zeroth, First and Second Order Reactions
Each ordered reaction has its own unique integrated rate law equation as well as half-life equation that you will be expected to know.
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Lecture 35
Understanding Chemical Equilibrium
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Understanding Chemical Equilibrium
A reaction reaches equilibrium once the rate of the forward reaction equals the rate of the reverse reaction. Once at equilibrium there is no net change in the concentration of reactants or products.
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Lecture 36
Understanding the Equilibrium Constant
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Understanding the Equilibrium Constant
Associated with any reaction at equilibrium is the equilibrium constant K. Its numerical value determines if reactants or products are more greatly favored within a reaction. The equilibrium constant K is a ratio of products to reactants.
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Lecture 37
The Different Types of Equilibrium Constants
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The Different Types of Equilibrium Constants
When dealing with equilibrium units in terms of pressure the equilibrium constant of Kp must be used, but if the equilibrium units are in molarity then we must use the equilibrium constant of Kc.
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Lecture 38
Understanding an ICE Chart
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Understanding an ICE Chart
In order to calculate the equilibrium concentrations of compounds we need to use an ICE Chart. An ICE Chart should be used when we are missing more than one equilibrium amount for compounds in our balanced equation.
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Lecture 39
Understanding Le Chatelier’s Principle
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Understanding Le Chatelier’s Principle
According to Le Chatelier’s Principle if a system (chemical reaction) is at equilibrium and we disturb it then the system will readjust to maintain it’s equilibrium state.
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Lecture 40
Identifying Binary Acids
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Identifying Binary Acids
Most acids have one common feature and that is the presence of the hydronium ion, which is represented by H+ or H3O+. Binary acids can be identified by the fact that they all possess an H+ ion attached to an electronegative element.
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Lecture 41
Identifying Oxyacids
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Identifying Oxyacids
They are created by the hydration of nonmetal oxides. Conversely, metal oxides create bases when hydrated by water.
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Lecture 42
Acids and Electrolytes
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Acids and Electrolytes
Strong Acids are strong electrolytes that completely dissociate into ions when dissolved in water. Weak acids are weak electrolytes that don’t completely ionize, but instead reach a state of equilibrium.
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Lecture 43
Strength of Binary Acids
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Strength of Binary Acids
When looking at a binary acid we look at both electronegativity and size to determine their strength.
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Lecture 44
The Strength of Oxyacids
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The Strength of Oxyacids
The strength of an oxyacid depends on the number of oxygen atoms and the electronegativity of the nonmetal. If your oxyacid has 2 or more oxygens than hydrogens then your oxyacid is a strong acid.
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Lecture 45
Comparing the Strength of Oxyacids
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Comparing the Strength of Oxyacids
When comparing the strengths of different oxyacids remember: If they have different number of oxygens then the more oxygens the more acidic. If they have the same number of oxygens then the more electronegative the nonmetal the more acidic.
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Lecture 46
Exceptions to the Strength of Oxyacids
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Exceptions to the Strength of Oxyacids
With some rules there are exceptions. Oxalic acid and Iodic acid have two more oxygen’s than hydrogens but are weak acids because carbon are iodine have low electronegativity.
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Lecture 47
Bases and Electrolytes
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Bases and Electrolytes
Strong Bases are strong electrolytes that completely dissociate into ions when dissolved in water. Weak bases are weak electrolytes that don’t completely ionize, but instead reach a state of equilibrium.
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Lecture 48
Bases and Group 1A
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Bases and Group 1A
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Lecture 49
Bases and Group 2A
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Bases and Group 2A
Group 2A metals, from Ca2+ to Ba2+, when combined with OH –, H –, O2– or NH2 – makes a STRONG BASE.
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Lecture 50
Amines
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Amines
Amines are compounds with only nitrogen or hydrogen (i,e. NH3) or with carbon, nitrogen and hydrogen (i,e. CH3NH2). Neutral Amines are considered weak bases. Positive Amines are considered weak acids.
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Lecture 51
Discussing Arrhenius Acids and Bases
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Discussing Arrhenius Acids and Bases
According to Svante Arrhenius the two most important ions fundamental to the concept of acids and bases were H3O+ , the hydronium ion, and OH –, the hydroxide ion. An Arrhenius acid increases the H+ ion when dissolved in a solvent.
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Lecture 52
Discussing Brønsted – Lowry Acids and Bases
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Discussing Brønsted – Lowry Acids and Bases
In 1923, Johannes Brønsted and Thomas Lowry developed a new set of definitions for acids and bases. According to Brønsted and Lowry, an acid was classified as a proton acceptor while a base is a proton acceptor.
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Lecture 53
Understanding Amphoteric Species
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Understanding Amphoteric Species
An amphoteric species is one that can act as an acid or base depending on the reactants involved.
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Lecture 54
Understanding Lewis Acids
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Understanding Lewis Acids
In the 1920s, Gilbert Lewis proposed a new set of definitions for acids and bases. According to the Lewis definition an acid is an electron pair acceptor.
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Lecture 55
Understanding Lewis Bases
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Understanding Lewis Bases
Compounds that have lone pair(s) on a central element can act as Lewis bases. Compounds that possess a negative charge can act as Lewis Bases.
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Lecture 56
Understanding pH vs pOH
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Understanding pH vs pOH
The pH and pOH of a compound helps to determine its acidity and basicity. pH stands for the negative logarithmic function of your hydronium concentration, whereas pOH stands for the negative logarithmic function of hydroxide concentration.
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Lecture 57
The connection to Hydronium ion and Hydroxide ion concentrations
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The connection to Hydronium ion and Hydroxide ion concentrations
If we know the pH or pOH concentration then we can determine the concentration of hydronium ions or hydroxide ions.
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Lecture 58
The pH Scale
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The pH Scale
If the pH is less than 7 then the solution is acidic, if the pH is equal to 7 then the solution is neutral and if the pH is greater than 7 then the solution is basic.
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Lecture 59
Understanding the Auto Ionization Reaction
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Understanding the Auto Ionization Reaction
In a self-ionization reaction two water molecules react with one another, where one acts an acid and the other as a base.
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Lecture 60
Calculating the pH of a Strong Acid or Strong Base
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Calculating the pH of a Strong Acid or Strong Base
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Lecture 61
Calculating the pH of a Weak Acid or Weak Base
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Calculating the pH of a Weak Acid or Weak Base
If we have a WEAK ACID or WEAK BASE then we DO use an ICE chart to find the pH of solution because the yield of H+ and OH – ions are unknown.
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Lecture 62
Understanding Ka and Kb Expressions
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Understanding Ka and Kb Expressions
Associated with any weak acid or weak base is a Ka or Kb value respectively. Weak acids possess a Ka value less than 1, while weak bases possess a Kb value less than 1. The equilibrium expressions of Ka and Kb are the same as other equilibrium constants we’ve seen.
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Lecture 63
Understanding Ka and Kb
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Understanding Ka and Kb
The greater the Ka value then the stronger the acid, while the greater the Kb the stronger the base. Ka and Kb are connected by the following equation: Kw = Ka * Kb
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Lecture 64
Transition Metal Cations and Acidity
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Transition Metal Cations and Acidity
When an acid and base undergo a neutralization reaction they form ions, which can create acidic, basic or neutral solutions. Cations (positive ions) can create either acidic or neutral solutions.
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Lecture 65
Main Group Metal Cations and Acidity
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Main Group Metal Cations and Acidity
If the main group metal within an ionic compound has a charge of +3 or higher then it creates an acidic solution. If it’s less than +3 then it creates a neutral solution and is classified as a non-significant ion.
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Lecture 66
Positive Amines and Acidity
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Positive Amines and Acidity
Amines are compounds with only nitrogen or hydrogen (i,e. NH3) or with carbon, nitrogen and hydrogen (i,e. CH3NH2). Positive amines within compounds create acidic solutions.
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Lecture 67
Understanding a Buffer
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Understanding a Buffer
A buffer is a solution composed of a weak acid with its conjugate base. Buffers resist drastic changes to the pH if a strong acid or strong base is added.
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Lecture 68
Method 1 for a Creating a Buffer
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Method 1 for a Creating a Buffer
As we stated earlier a buffer is composed of a weak acid and its conjugate base, but there are actually three ways to create a buffer.
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Lecture 69
Method 2 for a Creating a Buffer
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Method 2 for a Creating a Buffer
The second method in creating a buffer is mixing a strong acid with a weak base. In this case since we have a strong species mixing with a weak species then we must make sure the weak species is higher in amount.
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Lecture 70
The Henderson – Hasselbalch Equation
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The Henderson – Hasselbalch Equation
Whenever we have a buffer we can skip the ICE Chart and use the Henderson Hasselbalch.
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Lecture 71
Understanding an Acid–Base Titration Curve
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Understanding an Acid–Base Titration Curve
Titration curves help us to understand the neutralization occurring between an acid and base in order to identify the equivalence point. The equivalence point is when the moles of acid and base titrating are equal in amount.
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Lecture 72
The Equivalence Point
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The Equivalence Point
Depending on the types of acids and bases mixing the equivalence point can be less than, greater than or equal to 7.
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Lecture 73
Understanding Weak Acid–Base Titrations
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Understanding Weak Acid–Base Titrations
Whenever we react a WEAK ACID with a STRONG BASE or a STRONG ACID with a WEAK BASE we use an ICF CHART. In an ICF Chart the units will be in moles instead of Molarity.
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Lecture 74
Understanding Strong Acid–Base Titrations
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Understanding Strong Acid–Base Titrations
If we react a STRONG ACID with a STRONG BASE then we DO NOT use an ICF Chart or an ICE Chart.
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Lecture 75
Understanding the Solubility Product Constant
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Understanding the Solubility Product Constant
Associated with any ionic compound is the solubility product constant, which measures how soluble the compound will be in a solvent.
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Lecture 76
Understanding Spontaneity
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Understanding Spontaneity
Thermodynamics is the branch of physical science concerned with heat and its transformations to and from other forms of energy. A reaction that requires no outside energy source is classified as a natural process and is spontaneous.
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Lecture 77
The 2nd and 3rd Law of Thermodynamics
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The 2nd and 3rd Law of Thermodynamics
Entropy is the disorder or chaos associated with a system’s inability to convert thermal energy into mechanical work. The 2nd Law of Thermodynamics states systems (chemical reactions) spontaneously move to state of disorder.
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Lecture 78
Increasing or Decreasing Entropy
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Increasing or Decreasing Entropy
During a phase change as our molecules grow farther apart then entropy will increase.
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Lecture 79
Calculating Entropy of the Universe
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Calculating Entropy of the Universe
The 2nd Law of Thermodynamics states that the entropy of the universe is always increasing and so it must always be greater than zero.
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Lecture 80
Total Entropy and Spontaneity
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Total Entropy and Spontaneity
If the entropy is greater than zero then we classify the process as spontaneous.
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Lecture 81
Understanding Gibbs Free Energy
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Understanding Gibbs Free Energy
If our Gibbs Free Energy is less than zero then a reaction is spontaneous. If a reaction is spontaneous in one direction then it will be non-spontaneous in the opposite direction.
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Lecture 82
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 83
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 84
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 85
Cell Potential and Spontaneity
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Cell Potential and Spontaneity
There is a relationship between Gibbs Free Energy, the equilibrium constant K and the cell potential and spontaneity.
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Lecture 86
Understanding Nuclear Reactions
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Understanding Nuclear Reactions
Radioactive reactions deal with the chemical instability of the nucleus in an atom. Heavy (large atomic mass) elements undergo radioactive reactions in order to increase the stability of their nuclei.
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Lecture 87
Types of Nuclear Reactions
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Types of Nuclear Reactions
In a radioactive decay or emission reaction the radioactive particle is ejected from the nucleus and forms a product. In a radioactive capture or absorption reaction the radioactive particle is taken into the element and so is seen as a reactant.
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Lecture 88
Understanding Alpha Decay or Emission
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Understanding Alpha Decay or Emission
An alpha particle is comprised of 2 protons and 2 neutrons. An alpha decay or alpha emission occurs when an unstable nucleus ejects an alpha particle to create a new element.
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Lecture 89
The Alpha Particle
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The Alpha Particle
The alpha particle is one of the largest radioactive particles with the highest ionizing power, but lowest penetrating power.
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Lecture 90
Understanding Beta Decay or Emission
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Understanding Beta Decay or Emission
A beta particle has no atomic mass and is represented by an electron. A beta decay or beta emission occurs when an unstable nucleus ejects a beta particle to create a new element.
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Lecture 91
The Beta Particle
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The Beta Particle
Beta particles are smaller in size, and therefore have more penetrating power. Luckily, they are less radioactively damaging because of their lower ionizing power.
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Lecture 92
Understanding Gamma Radiation
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Understanding Gamma Radiation
A gamma particle has no atomic mass and no atomic number and is represented by the sign gamma. Gamma radiation is involved in the electromagnetic spectrum. Gamma rays possess the highest energy, while radio waves have lowest energy in terms of the spectrum.
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Lecture 93
The Gamma Particle
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The Gamma Particle
The gamma particle does not create a new element like the other radioactive particles, but instead causes the excitation of electrons within an element. Gamma Particles have lowest ionizing power, but are the most dangerous because of their highest penetrating power.
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Lecture 94
Understanding Electron Capture
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Understanding Electron Capture
In an electron capture or electron absorption reaction our electron particle is a reactant and not a product.
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Lecture 95
Understanding Positron Emission
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Understanding Positron Emission
A positron particle is referred to as the anti-electron particle because it looks like a positively charged electron. A positron decay or positron emission occurs when an unstable nucleus ejects an alpha particle to create a new element.
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Lecture 96
Radioactive Decay Rates
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Radioactive Decay Rates
Radioactive reactions follow principles dealing with Chemical Kinetics. Radioactive decay follows a 1st order mechanism and so has a similar integrated rate law equation from Chemical Kinetics.
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