Workbook/Note-Guide to accompany

Videos

Sandra Y. Etheridge, Ph.D.

Workbook/Note-Guide to accompany

College Chemistry II Videos

Sandra Y. Etheridge, Ph.D.

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Information on these discs utilized with permission of Gulf Coast Community College. The college does not provide support for this version of the software, and all questions and concerns must be directed to the Chemistry Professor. ©2007 All rights reserved

All rights reserved. No part of this workbook/note-guide may be reproduced in any form without permission in writing from the Chemistry Professor. the Chemistry Professor is a trademark of Sandra Y. Etheridge © 2007, the Chemistry Professor

Youngstown, Florida, 32466Introduction

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Introduction

The videos used for this course were made in the studios at Gulf Coast Community College by Dr. Sandra Etheridge and were designed to meet the needs of students taking chemistry by distance education. The course is referred to on the videos as CHM 1046 which is in accordance with the common course numbering system for Florida Universities and designates the course as being the second semester college or university chemistry lecture course for majors in the field. The DVDs are copies of those distance education videos and are made available courtesy of Gulf Coast Community College. The workbook/note guide written by Dr. Etheridge is designed to facilitate organized and complete note-taking to be used for study. The workbook matches the videos exactly since Dr. Etheridge wrote both the course and the workbook/note guide. If a mistake is made on the video (there are several), the error and correction are pointed out in the note guide. Further, tables and charts needed at certain points in the lecture are provided at those points in the note guide. Notations of appropriate points at which to pause the video for problem solving appear in the note guide, and utilizing those pauses as suggested facilitates learning. For those reasons it has been found that students benefit significantly from properly using the note guide. It is such an effective aid that students enrolling in the lecture course at the college frequently purchase the note guide to assist their note taking and learning. It is suggested that individuals viewing these videos have completed at least one semester of college chemistry for majors as well as have a good background in algebra. It is strongly suggested that the viewer have access to a college chemistry textbook since that reference could be used to answer questions that might arise, as well as provide additional problems for practice and the very large reference tables needed for solving problems. Hopefully you will enjoy these videos and note guide as much as Dr. Etheridge enjoyed making them for her students.

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Table of Contents Unit Page 1. Solids and Liquids 1 (Lesson 1 - 7)

Types of Intermolecular Forces 2 Phase Changes 8 The Solid State 19

2. Behavior of Solutions 25 (Lessons 8 – 11)

Properties of Solutions 26 Concentration of Solutions 28 Colligative Properties 32

Ideal vs. Real Solutions 40 3. Chemical Kinetics 41 (Lessons 12 – 15)

Factors Affecting Reaction Rates 42 Orders of Reactions 46

Reaction Mechanisms 56 4. Introduction to Chemical Equilibrium 57 (Lessons 16 – 20) Foundations 58 Le Chatelier’s Principle 64 A Survey of Problems 65 5. Acids and Bases 73 (Lessons 21 – 24) A Review of Solutions 74 The Nature of Acids and Bases 75 The Meaning of pH 83

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6. Acid-Base Equilibria 89 (Lessons 25 – 35)

Hydrolysis of Salts 90 Buffers. 93 Titration Curves (numerous) 101 The Work of Indicators 110 Polyprotic Acids 112 A Survey of Problems 117 7. More Aqueous Equilibria – Solubility Products 121 (Lessons 36 – 38)

Writing the Expression 122 Molar Solubility 123 Predicting Solubility 125 The Common Ion Effect 127 Separating Ions 128 Effects of pH

More Aqueous Equilibria – Complex Ions 133 (Lessons 39 – 41) Definitions 134 Naming Complex Ions 135 Solubility and Complex Ions 136 A Survey of Problems 138 8. Thermodynamics 141 (Lessons 42 – 45) Spontaneity 143 ∆S 145 ∆G 148 The Keq 151 A Survey Problem 156 9. Electrochemistry 157 (Lessons 46 - 49)

Introduction 158 Voltaic Cells 160 Electrolytic Cells 168 Corrosion 170 A Survey of Problems 172

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10. Nuclear Chemistry 175 (Lessons 50 - 52) Introduction 176 Nuclear Stability 177 Disintegrations 180 Fission vs. Fusion 184 Uses of Radioactivity 187 Appendix 189 Table of Common Oxidation Numbers 191 Activity Series of the Elements 193 Solubility Rules 195 Select Solubility Product Constants 195 Select Reduction Potentials 197 Equilibrium Constants for Select Weak Acids 199 Equilibrium Constants for Select Weak Bases 199 Table of Selected Indicators 201 Periodic Table of the Elements 203 (also on inside front cover)

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1

During this unit on Solids and Liquids, Dr. Etheridge will

1. discuss the several types of intermolecular forces: ion-ion, ion-dipole, and dipole- dipole.

2. describe the properties of both permanent dipoles, particularly hydrogen bonds, and temporary or instantaneous dipoles such as London Forces.

3. relate the various types of intermolecular forces to the molecules in which these forces are predominant.

4. rank the strengths of these intermolecular forces. 5. relate types of intermolecular forces to physical properties. 6. discuss the Maxwell-Boltzman Distribution Curve of Molecular Speeds and its

relationship to certain physical properties such as vapor pressure. 7. solve problems related to vapor pressure and extent of evaporation. 8. explain boiling and the factors related to boiling 9. relate ∆Hvap to strength of intermolecular forces, vapor pressure, and boiling

point. 10. introduce the Clausius-Clapeyron equation 11. discuss some of the unique properties of water, i.e. surface tension, the floating of

ice, sublimation, and the heating/cooling curve. 12. solve problems related to the heating/cooling curve. 13. explain phase diagrams using water as an example. 14. discuss the solid state with respect to structure, motion, shape and volume. 15. describe the properties of an amorphous solid. 16. explain the four types of bonds existing in a crystalline solid. 17. introduce the concepts of allotropes using those of carbon. 18. describe in detail the structure of cubic cells, including simple cubic cells, body

centered cubic cells, and face centered cubic cells. 19. solve problems dealing with unit cells. 20. look more carefully at close packing and introduce the concept of slippage in

metallic solids.

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During this unit, we will discuss the following: Types of Intermolecular Forces Phase Changes The Solid State

TYPES OF INTERMOLECULAR FORCES 1. Ion – ion Draw the structure for the NaCl lattice These attractions are called ______________________________ Describe this attraction:

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2. Ion – dipole forces Review the structure of water: It’s geometry is ____________________________ Which element is more electronegative? ________________________ Insert the dipole notations into the above structure: What kind of dipoles are these? __________________________ Describe the attractive forces that exist between the water molecule and the Na+

Explain the shielding effect of water. Describe the attraction between water and the hydrogen ion.

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The strength of interaction depends on: 1. 2. 3. 3. Dipole – dipole forces These forces are also called ______________________________________ Let’s look at two forms Permanent dipoles Describe these forces using water. What is a hydrogen bond? Hydrogen bonds bridge atoms and exist between the following:

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Hydrogen Bonds Using water and H-F, describe how the hydrogen bonds works:

Using water and ammonia, describe the hydrogen bridge. Use this to explain the phenomenal solubility of ammonia gas in water.

Temporary Dipoles Describe these dipoles and compare their force to other dipoles: Using the chlorine molecule, explain how they work What are other names for this attraction:

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Explain what is meant by this being an “induced dipole.”

These weak forces account for much of the attraction between ___________ molecules, for example.

Upon what do these forces depend?

Using the halogens, describe how size/mass of molecules impact these attractions.

Using the organic compounds from C5H12, explain the impact of shape on attraction, hence on boiling points:

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Describe the structures that have the most effective dipole force:

Describe the structures that have the least effective dipole force:

Rank the strengths of the intermolecular forces: Boiling Points and Intermolecular Forces

As intermolecular forces increase, the boiling point __________________. Rank the following in order of increasing boiling point and explain: NaCl He CO2 CH3CH2OH (ERROR: Dr. Etheridge erroneously called CH3CH2OH methyl alcohol when in fact it is ethyl alcohol. When they heard, her organic students laughed out loud!)

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Rank the following in order of increasing boiling point and explain: MgO Ne H2S CH3OH PHASE CHANGES Vapor Pressure Why do liquids evaporate?

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Explain why temperature is a factor in evaporation. Changes of State solid to liquid is ____________________ liquid to solid is ____________________ liquid to gas is _____________________ gas to liquid is _____________________ solid to gas is ______________________ gas to solid is ______________________ Give an example of an endothermic process: Give an example of an exothermic process: Describe what goes on in a closed system containing a liquid and its vapor. Distinguish between the use of “vapor” and “gas.” On what does the vapor pressure depend?

If you have a 20.0L closed container at 40.0oC, what is the minimum quantity of water needed to produce a vapor pressure of 55.3 torr?

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If you put 350.0 mL C2H5OH in an evacuated, closed container measuring 6.0 ft x 12.0 ft x 8.5 ft and the temperature is constantly maintained at 68.4oF, will all of the alcohol evaporate? The vapor pressure of the alcohol at 68.4oF is 63.7 Torr and the alcohol density is 0.7893 g/mL.

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What is boiling and when does it occur? Boiling occurs when: The normal boiling point is: What happens to boiling point as altitude changes?

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Explain how a pressure cooker works: You have the following three liquids and their vapor pressures at 25oC: A 18 Torr B 248 Torr C 547 Torr Which will boil first? Define the following: volatile flammability Are they the same? Referring again to the three liquids above: Which has the strongest intermolecular forces and how do you know? Which will evaporate the most readily?

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How does ∆Hvap relate to strength of intermolecular forces? vapor pressure? the boiling point? Vapor Pressure (continued) Describe the relationship that exists between temperature and vapor pressure (the Clausius-Clapeyron equation)

The vapor pressure of water at 44.6oC is 70.41 Torr and at 77.2oC is 316.6 Torr. Find ∆Hvap.

(Please note that Dr. Etheridge corrected her calculator input error)

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Calculate the boiling point of water in a pressure cooker at 20.0 psi if the ∆Hvap=40.7kJ/mol. (REM: one atmosphere = 14.7 psi.)

(the answer will appear on the top of the next page)

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(answer to problem on previous page: 382K)

Unique Water Surface Tension Explain what is meant by the “skin” on the surface of water and how it happens. Why can bugs walk on water? Why does ice float? (explain in detail) (Water, not ice, is at its maximum density at 3.96oC.)

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Heating/cooling curve for WATER: (fill in the information on the following curve) How much energy is needed to convert 0.500 kg ice at -20.0oC to steam At 250oC? What would be the characteristic noted on a heating/cooling curve for something like iodine or dry ice which sublimes?

100

0

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How much water at 21.4oC must evaporate to leave 4.8 liters of water at 20.0oC? An ice cube tray contains enough water at 22.0oC to make 20 ice cubes each having a mass of 25.0 g. The tray is placed in a freezer using CF2Cl2 as a refrigerant. If the ∆Hvap for CF2Cl2 is 158 J/g, what mass of CF2Cl2 must vaporize to convert all the water to ice at -5.0oC? How many grams of NH3 liquid must evaporate to freeze 200.0 g water at 40oC? The heat of vaporization of NH3 liquid is 23.35 kJ/mol.

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(the answer to the problem on the previous page is 72.9 g) Why does water in a clay jar remain cooler than water kept in a glass container? Phase Diagrams: Notes:

Diagram for Water.

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Define “triple point” and give the triple point conditions for water: THE SOLID STATE Describe the characteristics of solids with respect to: -structure -motion -shape and volume Amorphous Solid Describe an amorphous solid.

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How can an amorphous solid be identified? In contrast, the melting point of a crystalline solid can be described as: Types of Bonding:

1. Ionic (describe)

2. Molecular (describe)

3. Metallic (describe)

4. Network (describe)

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Allotropes: What are allotropes? Consider carbon: How many allotopes does it have? ______________ List them:

Describe Graphite: Describe Diamond: Describe Fullerenes Buckminsterfullerene

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Crystalline Structures: There are actually _______ types of cells. For the cubic cell, there are ______types? Simple Cubic Cell: In the space to the right, draw a cubic cell

and note the square that defines the cubic cell. The volume of this simple cubic cell is defined as ________ or _____________. How many atoms are actually in a unit cell? ________ Body Centered Cubic Cell: Try to draw a bcc and note the

square that defines it. The volume of this body centered cube is defined as _______ or _______________. How many atoms are in a unit bcc cell? __________ Face Centered Cubic Cell: The volume of this face centered cube is defined as _______ or ________________. How many atoms are in a unit fcc cell? ______.

The face centered cube is characterized by cubic close packing.

Explain how this structure allows for slippage among metal atoms.

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Brief Review and a small Addition: What condition or conditions can produce an amorphous solid? What characteristics are often noted in these amorphous solids? List three allotropic forms of oxygen: Solving problems dealing with unit cells:` To determine density of a solid, use the formula d = Where: Z = M = NA= e3 = Ag crystallizes in a FCC arrangement with an atomic radius of 144 pm. Find the density of silver. (When Dr. Etheridge substituted in the formula, she did not write in the “cube;

however, she did use it in her calculations.)

Z .MNA

.e3

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BaO crystallizes in a FCC arrangement with O= defining the lattice structure and Ba2+ filling the octahedral holes. Find density of BaO if the radius (Ba+2) = 149 pm and the radius (O=) = 126 pm. A Closer look at Close Packing Note the two different kinds of holes that are formed. The hole that is formed over a sphere is a/an ____________________ hole.

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During this unit on Behavior of Solutions, Dr. Etheridge will

1. discuss factors affecting both quantity of solute dissolved and rate at which dissolving occurs.

2. describe the less quantitative methods of expressing concentration. 3. list the identities and concentrations of “standard desk reagents.” 4. review molarity. 5. note the three types of percent solutions and calculations involving each. 6. cover ppm and ppb solutions. 7. introduce mole fraction and mole percent with regard to solutions. 8. discuss molality of solutions. 9. introduce the concept of colligative properties, including vapor pressure, boiling

point elevation, and freezing point depression 10. use colligative properties to introduce the Rast method of determining molecular

weights. 11. introduce the factors used to address degree of ionization when dealing with

colligative properties. 12. discuss osmosis and the methods of determining osmotic pressure. 13. close with a brief discussion of the differences between ideal and real solutions.

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During this unit, we will discuss the following: Properties of Solutions

Concentration of Solutions

Colligative Properties

Ideal vs. Real Solutions

PROPERTIES OF SOLUTIONS Factors affecting Dissoving:

List the factors that affect the quantity of solute dissolved:

List the factors that affect the rate of dissolving:

List the factors that affect the rate of dissolving gases:

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A closer look at the temperature factor.

As the temperature increases, the solubility of gases __________________.

A closer look at the pressure factor.

As you increase the pressure, the solubility of gases ___________________

proportionally.

State Henry’s law:

Expressions of Concentration: Describe and define:

Saturated solutions

Unsaturated solutions

Supersaturated solutions

Concentrated vs. dilute:

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Standard Desk Reagents (fill this in)

Solution concentrated dilute

HCl

HNO3

H2SO4

HC2H3O2

NH4OH

NaOH

CONCENTRATION OF SOLUTIONS

Molarity What is molarity (give the formula for finding it):

How would you make 250 mL of a 0.30 M Ni(NO3)2 solution?

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Percent Solutions

The three types of percent solutions are

1.

2.

3.

Calculate the weight of lye needed to make 200. g of a 15.0% w/w solution.

Calculate the volume of ethyl alcohol required to make 1.00 L of a 45% v/v solution.

How could you make 2.00 L of a 7.5% w/v solution of whatever?

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Ppm, Ppb What types of solutions are these designations used for?

How many ions of magnesium can be found in 3.0 liters of salt water containing 2.5 ppb Mg. Use 1.00 g/mL for the density of salt water.

Mole Fraction What is mole fraction?

What is the mole fraction of acetic acid in 250 mL of a 0.44 M solution having a specific gravity of 1.01?

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What is the relationship between mole fraction and mole percent?

Molality

Molality allows us to deal with ___________________ properties.

The designation for molality is __________.

By definition, molality is:

Determine the molality of a solution made by dissolving 10.0 g C2H5OH in 100.0 g water.

(Watch out! Dr. Etheridge erred when she said, “moles of solute per tenth kg of solvent is molality. She should have said “moles of solute per kg of solvent is molality.)

For that 10.90 g C2H5OH in 100.0 g water, find molarity, mole fraction, and mole percent of alcohol if the solution has a density of 0.98 g/mL.

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COLLIGATIVE PROPERTIES What are “colligative properties”?

The three colligative properties with which we will deal are:

Vapor Pressure As temperature increases, vapor pressure _________________. As temperature

decreases, the vapor pressure __________________.

When dealing with solutions,

VPsolvent = xsolvent x Ppure solvent

This is called __________________________ Law.

VPsolute = xsolute x Ppure solute

Therefore, VPsolution = xsolvent x Ppure solvent + xsolute x Ppure solute

It would be a good idea to return to your College Chemistry I notes on gas laws and review Dalton’s Law of Partial Pressures.

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The brief review is an opportunity for you to fill out your notes.

Calculate the vapor pressure of a sugar solution made by dissolving 10.0 g C6H12O6 in 50.0 g water at 100oC

Boiling Point Elevation What happened to the boiling point of the water when a solute (sugar) was added?

Why didn’t the solution boil at 100oC?

What must happen for the solution to boil?

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Boiling Point Curve for Water:

State the rule for change in boiling point of water upon addition of a solute:

The boiling point constant for water is ______________

The formula describing this is:

Calculate the molality of an aqueous solution that boils at 100.722oC.

VP

temp. oC 100

760

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Other Common Solvents: You may refer to your textbook or other suitable reference such as a CRC Handbook of Chemistry and Physics for the physical properties of additional solvents.

Mixtures of volatile liquids:

Rem: Dalton’s Law of Partial Pressures Pt = Pa + Pb

Calculate the composition of a benzene in toluene solution that will boil at 90oC and 1.00 atm. At 90oC, the VPbenzene is 1022 Torr and the VPtoluene is 406 Torr.

(It really doesn’t matter which you use as solvent and which as solute. As a matter of fact you can certainly use Pt = xaPa + xbPb and that will eliminate any question about solute and solvent.)

Freezing Point Depression What happened to the freezing point of the water when a non-volatile solute (sugar) was added?

Why didn’t the solution freeze at 0oC?

What must happen for the solution to freeze?

What is the freezing point depression constant for water? ______________

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Explain how addition of a solute lowers the freezing point of water. To put into a few words that which Dr. Etheridge was saying: there are more solvent molecules in the ice at the surface than there are molecules of solvent in the liquid form at the surface of the ice.

Can you explain how antifreeze works?

Can you do calculations regarding molecular weight determination by freezing point depression?

Find the gmw of an unknown compound if 0.300 g dissolved in 60.00 g camphor has a freezing point of 177.9oC

You may wish to try the problem here, on your own. Dr. Etheridge will work this problem in the next lesson:

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From the previous lesson: Find the gmw of an unknown compound if 0.300 g dissolved in 60.00 g camphor has a freezing point of 177.9oC

This method of determining the molecular weight is called the ___________ method.

Why must these solutes be “non-volatile, non-electrolyte”?

Ions in Solution It is important that we know how many “particles” are present. What happens when we add a substance that is ionic?

The factor that compensates for extent of ionization is called the ____________

factor and is represented by the letter ______. We may not be able to assume 100% ionization.

Give the formula for finding the van’t Hoff factor:

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Consider the salt, ZnCl2. We would expect a 1.0 m solution to freeze at (3)(-1.86oC) or -5.58oC. However, in reality, the solution freezes at -5.21oC. What is the van’t Hoff factor for ZnCl2 under these circumstances?

What is the extent of ionization of ZnCl2 in the above situation?

Osmotic Pressure What is osmosis?

Consider a cell placed in water: What happens to the cell?

Explain:

More molecules pass through the membrane from the side of _______________

numbers to side of _____________ lesser numbers. Another way of saying this is:

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Explain the concept of reverse osmosis:

How can we prevent osmosis?

We use _________ to represent osmotic pressure.

Give the formula for calculating osmotic pressure:

How can you allow for ionization in these problems?

Calculate the osmotic pressure of a 3.6 M NaCl solution having a van’t Hoff factor of 2.6.

This method of solving the problem is most effective for dilute solutions.

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IDEAL vs. REAL SOLUTIONS Designate each of the following as “ideal” or “real.”

When substances are mixed, the volume of the solution is the sum of the component volumes.

Mixing neither absorbs nor evolves heat.

Each substance in solution obeys Raoult’s law.

The vapor pressure above an ideal solution obeys Dalton’s Law of Partial Pressures.

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During this unit on Chemical Kinetics, Dr. Etheridge will

1. discuss the impact of changing temperature, concentration, types of reactants, and catalysts on rates of reaction.

2. pay particular attention to the way in which catalysts function and the types of catalysts commonly used.

3. introduce the concept of “Energy of Activation” and diagrams that relate potential energy to progress of reactions.

4. demonstrate the difference in diagrams for spontaneous, endothermic, and exothermic reactions.

5. introduce the concept of orders of reactions. 6. discuss ways of describing the rate of a reaction and introduce the rate or

proportionality constant. 7. demonstrate how to determine a rate expression for several types of reactions. 8. introduce half life, its meaning and expression. 9. go through a series reaction orders: 0 order, first order, etc., describing the plot of

the reaction types. 10. lead the student through a series of problems dealing with first order reactions and

half life. 11. will explain how the carbon dating process works. 12. address second order reactions, the ways to express them, the half life expression,

and the plot which describes them. 13. discuss overall order of reaction. 14. introduce the Arrhenius Equation. 15. conclude with reaction mechanisms

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During this unit, we will discuss the following: Factors Affecting Reaction Rates Orders of Reactions Reaction Mechanisms

FACTORS AFFECTING REACTION RATES Temperature: Why does heating a reaction cause it to occur faster? Heating a reaction increases both the ________________________________ of collisions. Do all collisions result in a reaction? Why or why not? Specific effects of temperature.

For every 10oC the temperature of a reaction is increased, what will happen to the rate?

The formula to describe this is:

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Concentration Does changing the concentration change the rate of the reaction? Elaborate: How can we find out whether changing the concentration impacts the rate. Nature of the Reactants Describe this: How do we know what to expect? Catalysts What are catalysts? What is the difference between positive catalysts and retardants? What are homogeneous catalysts? It is important to note that homogeneous catalysts actually take part in the reaction mechanism. We’ll discuss this more later. Recall the decomposition of KClO3, a common way of producing oxygen in the lab. Write the equation for this decomposition. What is the purpose of the MnO2?

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Heterogeneous Catalysts Describe the heterogeneous catalyst and how it works. Using the example of hydrogen and oxygen gases with platinum as a catalyst, explain how a heterogeneous catalyst might be expected to work. Energy of Activation

Exothermic Reaction

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Endothermic Reaction Draw a curve for a spontaneous reaction How would you draw and describe a reaction that is just the opposite of a spontaneous reaction?

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ORDERS OF REACTIONS Expressing Rates How can we describe the rate of a reaction? For the equation: 2 N2O5 → 4 NO2 + O2, the rate of appearance of oxygen can be expressed as: This is the same as describing it relative to NO2 in which case the change is expressed as: This is the same as describing it relative to the disappearance of N2O5 which would be: Consider the reaction: A + B → C Write the rate expression: What is “k”? That rate constant is relative to what?

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Use the following table to record the data from that hypothetical reaction:

RUN [A] [B] Rate (hrs-1)

notes: Now, calculate k It is very important to note that the rate, R, will change with time! The proportionality constant, k, will be the same, regardless of the time or the concentration. However, k will change with temperature.

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Consider the following reaction: 2A + 3B + C → D + 2E and fill in the following table

RUN [A] [B] [C] Rate (day-1)

Determine the rate expression: Half-Life What is meant by half-life? The expression for half-life is _________________

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0o (Zero) Order Reactions What is a zero order reaction? Write the rate expression for this type reaction: Write the half-life expression for this reaction: If [A]t is plotted vs t, the plot is ___________________________ with a slope = _______ 1o Reactions (aka 1st order reactions) What type reactions are generally first order reactions? Give the rate expression for these reactions: When integrated, the equation becomes: (watch the sign) The half-life is expressed as: Describe the plot when ln[A] is plotted against time:

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The conversion of cyclopropane into cyclopropene is a first order reaction with k = 5.5 x 10-2 hr-1. If the initial [cyclopropane] is 0.55 M, how long will it take for the concentration to become 0.15 M? Consider the decomposition of H2O2, a first order reaction with k = 1.10 x 10-3 min-1. The initial [H2O2] = 0.025 M. What is the concentration after 2.5 hours? A patient is given a medication, 75% of which decomposes in 60 min. What is the half-life of the medication?

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4 PH3(g) → P4(g) + 6 H2(g) If the t1/2 is 35.0 sec at 690oC, how long will it take for 85% of the PH3 to decompose? CH3N=NCH3(s) → N2(g) + C2H6(g) R = k[CH3N=NCH3] t1/2 = 2.0 min @ 200.oC. If you begin with 1.0 g solid in a 500 mL flask, what is the pressure of N2 after 10 min.?

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Carbon Dating Explain how the carbon dating process works: The half-life of C-14 is __________________ Since so much of what we attempt to carbon date (artwork, bones, clay pottery, to name a few) is derived from living, carbon-based things, it makes sense to use carbon as a significant material for this process. However, there are severe limitations to the carbon dating process, namely that its accuracy has time limitations. An artifact has a decay rate of 0.186 disintegrations/sec for each gram of sample. If a young, living tree decays at 0.260 disintegrations/sec for each gram, how old is the artifact? (Check Dr. Etheridge’s math. The answer may be closer to 2768 yrs.)

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2o Reactions (second order reactions) There are several ways to express second order reactions. List the types we will use: Give the basic expression for a 2o reaction. Upon integration, this expression becomes: Give the half-life expression for this second order reaction. Describe the plot when is plotted vs time. When Iodine atoms combine to form I2 at 30oC, the reaction is second order

and has a rate constant of 6.2 x 109 M-1s-1.. If the initial concentration of I atoms is 0.120 M, calculate the concentration after 45 sec. What is the half life of this reaction? What would the half-life be if the initial concentration were 0.800 M?

1[A]t

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Overall Order of Reaction Considering the rate expression: R = k[A]2[B]o[C] This reaction is _________ order with respect to A, _________ order with respect to B, and ___________ order with respect to C. The reaction is ____________ order, overall. How is this found? **Rate expressions are always given in terms of _______________________________. The Arrhenius Equation Label as needed Write the Arrhenius Equation and identify each variable in the equation.

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Upon integration, the Arrhenius Equation becomes: Since R ∞ k, give another way of expressing the Arrhenius Equation: If raw mile sours in about 4 hours at 28oC and in 48 hours when refrigerated at 5oC, calculate the Ea for the souring of milk. Assume this is a 1st degree reaction. Now, let’s go to Reaction Mechanisms

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REACTION MECHANISMS Consider the reaction: H2O2 + 3 I- + 2 H3O+ → 4 H2O + I3

- for which the rate expression is R = k[H2O2]x [I-]y [H3O+]z The elementary steps are as follows: H3O+ + H2O2 → H3O2

+ + H2O fast H3O2

+ + I- → H2O + HOI slow HOI + H3O+ + I- → 2 H2O + I2 fast I2 + I- → I3

- fast Therefore, the rate determining step is: Thus, the rate expression should be: How do we deal with a factor involved in controlling the reaction rate if the factor is not a reagent used? Thus, the rate expression must be modified to become:

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During this unit on Introduction to Chemical Equilibrium, Dr. Etheridge will

1. develop the concept of equilibrium from rate expressions. 2. address writing the equilibrium expression. 3. explain the use of Kc as opposed to Keq. 4. discuss the states of matter that should and should not be used in these

expressions. 5. develop the concept of equilibrium when pressures of gasses is involved. 6. discuss the numerical value of the equilibrium expression relative to the quantities

of products and reactants involved. 7. introduce Le Chatelier’s Principle and the concept of stress. 8. show impact of various factors on the equilibrium. 9. work a series of problems dealing with molecular equilibria.

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During this unit, we will discuss the following: Foundations Le Chatelier’s Principle A Survey of Problems

FOUNDATIONS The Meaning of Equilibrium Describe/define equilibrium in terms of the freezing of water and the melting of ice. Rem: The key word in this is RATES. The equilibrium is DYNAMIC – ongoing. Consider: A + 2B D + 3C Plot concentrations vs. time for this reaction.

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Note: The concentrations reach a point at which they do not change. At that point, Rf = Rr Now, plot the rates of forward and reverse reactions vs time: Mark the point at which equilibrium exists. Writing the Expression Again, consider: A + 2B D + 3 C If the expression for the rate of the forward reaction is Rf = kf[A][B]2 and the rate of the reverse reaction is Rr = kr[D][C]3, then at equilibrium the rates are equivalent. Now, derive the expression for Q, the Law of Mass Action or Keq. Again, consider the reaction A + 2B D + 3C using a different set of reagents such that the Rf = kf[A][B], then the rate of the reverse reaction will be:

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Therefore, if you know the balanced equation, you can write the expression for Mass Action. If you also know the rate expression for either the forward or reverse reaction, then the other one may be determined. 2A + B 2 C + D + E If we determine that Rf = [A]2, then the rate of the reverse reaction MUST be: If a reagent and its numerical coefficient doesn’t appear in the rate expression for the forward reaction, then it must appear where? Now, consider 2A + 3B 2 C + D if Rf = kf[A]2 [B]3, then the rate for the reverse reaction is: why? Recall the reaction: 2 A + 3 B 2 C + D by experimentation: Rf = kf [A]2 [B]3 and by elimination Rr = kr [C2] [D] at equilibrium, Rf = Rr

then: Thus, the equilibrium is equal to the product of the [______________________] divided by the product of the [_____________________]. Don’t forget that!

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There are many equilibrium constants, but the one we use with molarity is Kc, instead of Keq. Should solids and liquids (not solutions) be included in these expressions? Why or why not? Therefore, only solutions and gases will appear in the Kc expressions. Consider: H2(g) + Br2 2 HBr(g) Write the equilibrium expression: Oops, Dr. Etheridge didn’t mean to say [H2] squared, she intended to say [H2]. Try this one: 2 H2(g) + O2(g) 2 H2O(g) Try this one: H2(g) + ½ O2(g) H2O(g) Try this one: BaCO3(s) BaO(s) + CO2(g) And this one: BaO(s) + CO2(g) BaCO3(s)

62

The Pressure Factor: How how P = MRT can be derived from PV = nRT Consider: PCl5(g) PCl3(g) + Cl2(g) Write the equilibrium expression, Kc, for this reaction: Derive the Kp for this reaction: Write the Kp for N2(g) + 3 H2(g) 2 NH3(g) Why must the Keq be temperature dependent? This is an important question and you should take the time to derive the proof.

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k2 Now, write the equation for ln (save this) Assume 6.00 moles PCl5 was placed in a 2.00 L container and allowed to decompose. If after equilibrium was established, 2.50 moles PCl5 remained, determine the Kc

PCl5(g) PCl3(g) + Cl2(g) Initially 5.00 moles NOBr were placed in a 1.00 L flask. At equilibrium, it was found that 12.0% of the NoBr has dissociated. Determine Kc. Then determine Kp at 25.0oC. 2 NOBr(g) 2 NO(g) + Br2(g) If the equilibrium constant is small, as in Kc for the above problem, what does the equilibrium favor? A large equilibrium constant favors the ______________________

k1

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LE CHATELIER’S PRINCIPLE What does this principle address: The principle of stress When stress is placed on a system at equilibrium, what must happen? What constitutes stress? When stress occurs, what does the reaction do? Stress and the direction of shift 2 NO(g) + O2(g) 2 NO2(g) ∆H = -114.1 k What happens when we add O2 add NO add NO2 remove O2 increase the temperature add a catalyst increase the pressure increase volume of container add Argon gas

65

A SURVEY OF PROBLEMS Consider the reaction: COBr2(g) CO(g) + Br2(g) [COBr2] = 2.0 M, [CO] = 3.0 M, [Br2] = 4.0 M at 150oC. Calculate a) Kc and b) Kp

Using information from the above equation: Empty the container. Now add COBr2 until equilibrium is reached and the COBr2 concentration is 6.0 M. Find the equilibrium concentration of CO and Br2.

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Using the same equation: Again, empty the container. Now add 5.00 M COBr2. Find the equilibrium concentration of each substance. Again, use the same equation and again, empty the container. Now add CO until its concentration is 2.0 M and add Br2 to 3.00 M. Calculate the new equilibrium concentration of each substance.

67

Consider the reaction (again): COBr2(g) CO(g) + Br2(g) [COBr2] = 2.0 M, [CO] = 3.0 M, [Br2] = 4.0 M at 150oC. Now, add 1.5 M COBr2. Calculate the new equilibrium concentration of each

substance in equilibrium. Recall the Kc was 6. Some things to think about regarding COBr2(g) CO(g) + Br2(g)

If more COBr2 is added to system when it is at equilibrium, does the equilibrium still exist? How can you prove a system is or is not at equilibrium? If you have a system at equilibrium and you add stress to the system, how can you prove the direction of shift?

68

Consider the reaction (again) COBr2(g) CO(g) + Br2(g) [COBr2] = 2.0 M, [CO] = 3.0 M, [Br2] = 4.0 M at 150oC. Add 2.0 M CO. Calculate the new equilibrium concentration of each substance. This is the problem Dr. Etheridge added on the graphics page.

Using the same equation. [COBr2] = 2.0 M, [CO] = 3.0 M, [Br2] = 4.0 M at 150oC. We want to increase the [COBr2] to 4.0 M at equilibrium. How much Br2 must be added to accomplish the increase in concentration of the [COBr2] to 4.0M, without adding any more carbon monoxide?

69

Use the same equation. [COBr2] = 2.0 M, [CO] = 3.0 M, [Br2] = 4.0 M at 150oC. How much COBr2 must be removed to have an equilibrium concentration of CO become 1.0 M?

Consider the reaction (again) COBr2(g) CO(g) + Br2(g) [COBr2] = 2.0 M, [CO] = 3.0 M, [Br2] = 4.0 M at 150oC. If the volume of the container is instantaneously halved, what is the new equilibrium concentration of each substance?

70

What’s this? A new equation!! 2 NO2(g) 2 NO(g) + O2(g) [NO2] = 3.0 M, [NO] = 6.0 M, [O2] = 5.0 M a) Calculate Kc b) How much O2 must be removed to increase the [NO] at equilibrium to 7.0 M? Consider the equation: Ni(s) + 4 CO(g) Ni(CO)4(g) You have CO(g) at 3.0 atm and 200oC. Ni(s) is added until the PCO is reduced to 1.0 atm at equilibrium. Find Kp.

71

Consider: C(s) + CO2(g) 2 CO(g) You have CO2(g) at 4.00 atm and 150oC. C(s) is added until P = 6.00 atm at equilibrium. Calculate Kp

Consider: CO(g) + Cl2(g) COCl2(g) P = 1.5 atm 2.0 atm 5.0 atm at equilibrium

CO(g) is added until PT = 10.0 atm. Calculate PT after equilibrium has been reestablished.

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Consider: NH4HS(s) NH3(g) + H2S(g) 6.1589 g Solid NH4HS is placed in a 4.00 L container at 24.0o. After equilibrium has been established, PT = 0.70 atm with some solid remaining. a) Determine Kp b) Determine the percent decomposition c) If the volume of the container is increased, what happens to the amount of solid?

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During this unit on Acids and Bases, Dr. Etheridge will

1. review concentration of solutions, dilution of solutions, and stoichiometry regarding solutions.

2. review the solubility rules that the student should memorize. 3. discuss and review molecular, ionic, and net ionic equations. 4. discuss the first of three acid-base theories, the Arrhenius Theory and its

limitations. 5. introduce the Lowry-Brönsted theory and show how acids and bases are identified

under that theory. 6. show how conjugate acid-base pairs are related 7. introduce the Lewis Theory of acids and bases and explain how this theory is

more expansive than either the Arrhenius or Lowry-Brönsted. 8. discuss the relative strengths of acids and bases with respect to magnitude of

anionic charge and oxygen content. 9. compare acid strength using reaction information. 10. demonstrate how to develop a sequence of acid and base strengths. 11. discuss the meaning of pH and the role of water. 12. develop the Ka, Kb, and Kw concepts. 13. work several pH problems.

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During this unit, we will discuss the following: A Review of Solutions The Nature of Acids and Bases The Meaning of pH

A REVIEW OF SOLUTIONS Molarity: What is molarity? This can be used for both ______________ and ________________. Dilution of Solutions: Give the formula relationship for dilution of solutions: Solution Stoichiometry: Calculate the concentration of all ions remaining in solution when 20.0 mL of 0.20 M Ba(NO3)2 and 30.0 mL of 0.40 M H2SO4 are mixed.

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Solubility Rules: 1. 2. 3. 4. Molecular, Ionic, and Net Ionic Equations: For molecular equations, show everything as ___________________________. For ionic equations, show all strong acids, strong bases, and soluble salts as __________. Show all else as molecules For net ionic equations, show only those things that are __________________________, following the rules of ionic equations, above. Ex: Ba(NO3)2(aq) + H2SO4(aq) Total ionic equation: Net ionic equation: THE NATURE OF ACIDS AND BASES List the strong acids: List the strong bases:

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Recall the characteristics of acids: Recall the characteristics of bases: Acid-Base Theories We will study the Arrhenius Theory the Lowry-Brönsted Theory the Lewis Theory The Arrhenius Theory How does this theory define acids? bases? This is a very severely limited theory of acids and bases. The Lowry-Brönsted Theory Acids are _________________________ Bases are _________________________ Consider the equation: HCl + H2O → H3O+ + Cl- identify the acid and base Consider the equation: H-Cl + :NH3 → NH4

+ + Cl- identify the acid and base Why did this NOT fit the Arrhenius definition?

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The Lowry-Brönsted Theory (cont.) Write the equation that explains the intense solubility of ammonia gas in water and use it to explain that solubility. Write the conjugate acid-base pair for ammonia as a base and as an acid. How are the components of a conjugate acid-base pairs related? Write the conjugate acid-base pair for water as a base and as an acid. Consider the reaction: Identify the base, acid, conjugate acid, and conjugate base. HS- + HNO2 → H2S + NO2- Consider this one, complete it, and label the parts like the equation above. HCN + HSO3

- →

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The Lewis Theory Acids are _____________________________________. Why will they accept electrons? Bases are ____________________________________. Why do they have this characteristic? Use BF3 and F- as an example: Use the hydronium ion and the hydroxide ion as another example: Use the Cu2+ ion and ammonia molecules as yet another example: Relative Strengths of Acids and Bases Which of the following is more likely to lose a proton? NH3 or NH4

+ Which of the following is more likely to gain a proton? HSO4

- or SO4= and why?

Which is more likely to lose a proton? HSO4

- or H2SO4 and why?

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Complete the following to make a summary statement: For molecules/ions that differ only in the number of H+, When H3BO3, H2BO3

-, and HBO3= are compared, what is the order of decreasing acid

strength and why? Complete the statement: For molecules/ion combinations that differ only in the number of H+, again, the Strong acids are strong because the conjugate base is ________________. (If the conjugate base were strong, it would hold those hydrogen ions tightly.) Consider BO3

3-, HBO32-, and H2BO3

- and indicate the order of base strengths. How would you compare the SO4

= and HSO4- and how did you arrive at your

conclusion? The greater the negative charge, the _________________ the base. Explain:

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Comparing Acids with different anions Explain how do we do this: Compare these acids: HClO3, HIO3, and HBrO3 What can be said about the size of the central atom and strength of the acid? Comparing acids with different oxygen content What can be said about the quantity of oxygen present and the acid strength? Explain. Which acid in each of these pairs is stronger? H2SO4 or H2SO3 H3AsO4 or H3AsO3 H3AsO4 or H3PO4 Comparing acids using reaction information Which is the stronger acid, CH4 or H2S? WHY? CH3

- + H2S → CH4 + HS-

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Using the following equation, determine which is the stronger acid, CH4 or H2S? CH3

- + H2S → CH4 + HS- Using this equation, determine which is the stronger acid, H2S or HNO2 HS- + HNO2 → H2S + NO2

- Consider the reverse of the reactions: Which is the stronger base, CH3

- or HS- CH4 + HS- → no reaction Which is the stronger base, NO2

- or HS-? H2S + NO2

- → no reaction Now, based on the above reactions, order the bases from strongest to weakest: Based on the above reactions, order the acids from strongest to weakest: On basis of that, predict the results of the following: CH3

- + HNO2 → Add these two reactions to the sequence: CH4 + OH- → no reaction H2S + OH- → HS- + H2O Rewrite the sequence of base strengths including the OH- and try these reactions: HNO2 + OH- → H2O + NO2

- →

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Given the following series of reactions, develop a sequence of base strengths: HOAc + C2H5O- → OAc- + C2H5OH H2S + C2H5O- → HS- + C2H5OH H2O + C2H5O- → OH- + C2H5OH H2O + HS- → no reaction H2S + OAc- → no reaction Write the bases from strongest to weakest (you may want to pause the video for a moment and work this out.) Consider the situation: NH4OH + NH4Cl + Mg2+ → no reaction Why? Consider this: Ca(OH)2 + Mg2+ → Mg(OH)2 ↓ + Ca2+

Which is the stronger base, Ca(OH)2 or Mg(OH)2 and how do you know? Why are the strong acids strong? Why are they all about the same strength? When do their strengths differ? List them in order of greater to lesser strengths: What is this effect called? And why does it work?

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THE MEANING OF pH Finding the [H+] Consider the strong acid, HCl In a 0.12 M HCl solution, the [H+] (or [H3O+]) is __________ For a strong acid, the [H+] is the same as the acid concentration. For a weak acid, however, we must use an equilibrium expression: Give the equilibrium expression for the ionization of acetic acid: H2O + HOAc H3O+ + OAc- Does the concentration of water change significantly? Rewrite the expression allowing for water constancy. REM: It’s a Ka Calculate the percent ionization of an aqueous 0.30 M HOAc solution. For acetic acid, the Ka = 1.76 x 10-5 Learn that. When can you drop x in these calculations? Identify x. What does x represent? ______________________________

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What does pH mean? What is the pH of a 0.20 M HCl solution? What is the pH of a 0.20 M HOAc? You know the Ka for HOAc. Just in case you don’t, the Ka for acetic acid is 1.76 x 10-5.

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The pH Scale (and the Role of Water) In a perfectly neutral sample of water, the [H+] = [OH-] = ____________________ Calculate the molarity of water: Does the molarity of water change in the process of ionization? Write the Keq for water: Rewrite as the Kw, the equilibrium expression for water: (Well, not ALWAYS. If you increase the temperature significantly, the Kw will increase.) Derive the expression for the sum of pH and pOH What, then, is the pH scale range? REM: The pH scale is designed for dilute solutions. We generally consider the solutions being no more concentrated than 1.0 M, although upon rare occasion we may use them.

86

If [H+] > 1 x 10-7, the solution is _____________ and the pH is ______________. If [H+] < 1 x 10-7, the solution is _____________ and the pH is ______________. If [H+] = [OH-] = 1 x 10-7, the solution is ____________ and the pH is ___________. On this scale, mark the acid and basic areas 14 7 0

0 7 14 Calculate the pH of a 0.20 M NaOH. Calculate the pH of the weak base, NH4OH, if the concentration is 0.20 M Pay attention to the Kb concept.

87

A table of Ka and Kb is found in the appendix of this book, as well as in your textbook. Calculate the pH of a 0.15 M HCN solution. Calculate the pH of a 0.15 M HNO3 solution. Calculate the pH of a 0.10 M Ba(OH)2 solution. Is it strong or weak? Calculate the pH of a solution made by diluting 15.0 mL of 6.0 M HOAc to 800 mL

88

What is the pH of a 0.022M aniline, C6H5NH2, solution? Treat this compound like HNH2, but do not allow the C6H5 portion to dissociate from the nitrogen.

89

During this unit on Acid-Base Equilibria, Dr. Etheridge will

1. introduce the concept of salts as an acid-base pair.

2. discuss the hydrolysis of salts. 3. examine the relationship between Ka and Kb. 4. introduce the common ion effect. 5. develop the buffer concept and the common kinds of buffers. 6. describe the impact on equilibrium and pH when a strong base is added to a

solution of a weak acid and its salt. 7. derive the Hendenson-Hasselbalch equations weak acids with the salt of the weak

acid and weak bases with the salt of the weak base. 8. show how to develop the pH range of buffers. 9. develop the concept that a titration curve has four areas and introduce each area. 10. introduce the titration curve for a weak acid with a strong base and show its

pattern. 11. introduce the titration curve for a strong acid with a weak base and show its

pattern 12. introduce the titration curve for a weak acid with a weak base and show its

pattern. 13. introduce indicators, how they work, and how to select the proper one. 14. discuss dissociation patterns of polyprotic acids. 15. show how the dissociation patterns of polyhydroxy bases are different from

polyprotic acids. 16. work several problems involving first and second dissociations of polyprotic

acids. 17. develop the combined K1K2K3 expression. 18. work problems related to these dissociations.

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During this unit, we will discuss the following: Hydrolysis of Salts Buffers Titration Curves The Work of Indicators Polyprotic Acids A Survey of Problems

HYDROLYSIS OF SALTS Classification of Salts: Consider the salt, NaCl. What acid ____________ and what base ____________ could have been used to produce this salt? Are these acid and base strong or weak? Consider the salt, NaOAc. What acid/base pair could have been used to produce this? Write the equation. This NaOAc is the salt of a _________acid and a __________ base. It is the weak one that reacts with water (undergoes hydrolysis). Write the reaction: Therefore, the solution produced is ____________ with a pH ________ 7. Give the net ionic equation. The salt of a weak acid and a strong base will hydrolyze to produce a/an ____________. The salt of a strong acid and a weak base will hydrolyze to produce a/an ____________.

91

Using this principle, explain why NaHCO3 produces a basic solution: Using this principle, explain what NH4Cl will do in water: REM: The salt of a weak base and strong acid produces a/an __________ solution. The salt of a strong base and weak acid produces a/an __________ solution. The salt of a strong base and strong acid produces a/an __________ solution. The salt of a weak base and a weak acid produces a/an __________ solution. Reaction with Water Examine the relationship between Ka and Kb by writing the equilibrium expressions for the following

HA H+ + A- and A- + H2O HA + OH- Ka x Kb = Therefore, Ka x Kb = And … This is an important relationship and you would do well to remember how to find Ka from Kb and vice versa.

92

pH of Salt Solutions Calculate the pH of a 0.14 M NaOAc solution Calculate the pH of a 0.080 M NaOBr. The Ka for HOBr is 6.2 x 10-9. Calculate the pH of a 0.10 M (CH3)3NHCl. The Kb for trimethylammonium chloride is 6.3 x 10-5.

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BUFFERS Le Chatelier’s Principle (again) Consider the equation: HOAc H+ + OAc- What happens when acetate ion is added? What happens to [H+]? What happens to the pH? What happens to the molarity of the undissociated HOAc? Calculate the pH of a 0.080 M HOAc. Then calculate the pH after the solution is made 0.10 M in NaOAc? How did the added OAc- affect the pH?

94

The Common Ion Affect What was the “common ion” in the problem on the previous page? _________________ Again, consider HOAc H+ + OAc- Give two ways in which we could increase the [OAc-], thus producing a common ion effect. Find the pH of the solution made by adding 10.0 mL of 0.20 M NaOH to

40.0 mL of 0.10 M HOAc.

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Now we are ready for the definition of a buffer. What is a buffer? The common kinds of buffers are: 1. 2. 3. What about the salt of a strong acid/strong base? Why?

1. Use HOAc + OAc-

What if you add extra H+

Here is a problem for you to contemplate. Dr. Etheridge will show you how to work it in the first part of the next lesson.

Calculate the pH of a 0.10 M HOAc that is 0.12 M in NaOAc. Then calculate the pH when 1.0 mL of 1.0 M HCl is added to 30 mL of the above.

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Now, lets work the problem without the addition of the NaOAc. Calculate the pH of a 0.10 M HOAc that is 0.12 M in NaOAc. Then, calculate the pH when 1.0 mL of 1.0 M HCl is added to 30 mL of the above. Now, compare the results of the pH with the buffering NaOAc and without the buffering NaOAc. Consider the equilibrium: HOAc H+ + OAc-

Describe what happens to the equilibrium and the pH when a strong base, NaOH is added to this solution.

97

Recall this problem from the previous lesson?

Calculate the pH of a 0.10 M HOAc that is 0.12 M in NaOAc. HOAc H+ + OAc- You may wish to quickly work it here, again. Now, modify it to read like this: Calculate the pH of a 0.10 M HOAc that is 0.12 M in NaOAc. Add

1.0 mL of 1.0 M NaOH to 30 mL of the buffer. Find the pH.

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The Henderson-Hasselbalch Equation Derive the Henderson-Hasselbalch Equation: This equation can only be used when x can be dropped. Calculate the pH of a solution that is 0.130 M HOAc and 0.080M in NaOAc. What weight of NaOAc must be added to 100 mL of a 0.14 M HOAc to produce a buffer with a pH of 4.9?

99

Rewrite the Henderson-Hasselbalch equation for a weak base with the salt of the weak base: Calculate the pH of a solution that is 0.10 M NH4OH and 0.12 M NH4OH. Buffer Range: List the factors that control the range over which a buffer can work: Consider 100 mL of a buffer containing 15 mmoles HOAc and 10 mmoles NaOAc. What will happen if we add more than 15 mmoles of strong base or 10 mmoles of strong to this buffer and WHY?

100

Some Important Buffers: List three buffers important to the human body and the physiology they impact. In this next section, we will cover Titration Curves which involves a combination of all acid-base equilibria we have discussed to date. Each titration curve involves a series of problems in order to plot the curve and Dr. Etheridge works a very large number in this series to produce the plot. If you understand exactly what she is doing at each point, you may wish to fast forward your video to the final series of points to check your work. However, if you have any doubts or errors in your work, be sure to view each section.

101

TITRATION CURVES Weak Acid vs. Strong Base We will plot volume of strong base against pH. 4 Areas of a titration curve: I. II. III. IV.

Plot a titration curve when 20.0 mL of 0.10 M HOAC is titrated with 0.10 M NaOH

v. NaOH pH 0 mL

5 mL

10 mL

15 mL

19 mL

19.5 mL

20 mL

20.5 mL

21 mL

25 mL

30 mL

Go to the next page for some space to work these sections:

102

Area I: Find the pH of the acid solution before any base is added Area II: Find the pH of the buffer

begins in this section.

103

Area III. Find the pH at the equivalence point. This is the point at which equal numbers of millimoles of H+ and OH- are present.

begins in this section Area IV. Find the pH when excess base is added

104

Now, plot each of the pH values on the vertical axis vs volume on the horizontal. Draw a smooth curve connecting the points. Mark area I, the area in which no base has been added Mark area II, the area in which the buffer exists. Mark area III, the equivalence point Mark area IV, the area in which excess base is used. Describe the SHAPE of each area, paying particular attention to the gradual incline, the transition state, and the sharper curve.

105

Plot a titration curve when 20 mL of 0.10 M HOAc is titrated with 0.08 M NaOH. Here is some working space. Make your own notes.

v. NaOH pH

These calculations are continued in the next lesson:

106

Notes and calculation space: Information from the plot:

107

Weak Base vs. Strong Acid: Plot a titration curve when 20.0 mL of 0.10 M NH4OH is titrated with 0.10 M HCl. The four areas are very similar to that seen previously Area I – NH4OH, only Area II – buffer Area III – equivalence point Area IV – excess HCl Calculations:

v. HCl pH 0 mL

5 mL

10 mL

15 mL

19 mL

19.5 mL

20 mL

20.5 mL

21 mL

25 mL

30 mL

108

Calculations, continued:

109

Plot and describe the curve:

Calculate the equivalence pH when 0.12 M HOAc is titrated with 0.08 M NaOH.

110

THE WORK OF INDICATORS How does an indicator work, colorwise? When are indicators most effective? Refer to the table of indicators below and in the Appendix. What characteristics must an effective indicator have?

Be sure you can look at a titration curve and make a reasoned deduction about whether the acid and base being titrated are strong or weak. For this next section, let’s refer to this table of indicators: Please note that this table is not quite the same as the table appearing in your video.

Indicator pH Range Methyl Orange 3.2 – 4.4 Bromocresol Green 3.8 – 5.4 Litmus 4.5 – 8.3 Bromocresol Purple 5.2 – 6.8 Bromothymol Blue 6.0 – 7.6 Thymol Blue 8.0 – 9.6 Phenolphthalein 8.3 – 10.0

111

How Indicators Work: Describe the chemistry of indicators, using bromocresol green If the ratio of HIn to In- must be ________________ to view the color as HIn.

Suppose our indicator, bromocresol green, has a pKa of 4.8. Let’s determine the pH range of this indicator.

It is interesting to note that the generally accepted, experimental range for bromocresol green is 3.8 – 5.4, which is not quite as we calculated. But, it isn’t bad.

112

POLYPROTIC ACIDS Dissociation Patterns Describe the dissociation pattern of the weak acid, H2SO3: Consider the influence of the first dissociation on the second dissociation: Write the K1 for the acid: Write the K2 for the acid: Write the dissociations series for H3PO4

Bases are different. Write the dissociation for Mg(OH)2 Thus, there is no serial dissociation of most bases.

113

The pH Picture: Calculate the pH from the first dissociation of 0.10 M H2SO3. The K1 = 1.71 x 10-2.

Now, calculate the pH when the second dissociation occurs. The K2 = 6.0 x 10-8. Because the [H+] is virtually unchanged by the second dissociation , the pH of the solution is the result of the first dissociation. To find the pH of the polyprotic acid solution, look at the _______________ dissociation.

114

Calculate the pH of a 0.14 M H2CO3. K1 = 4.30 x 10-7 and K2 = 5.61 x 10-11 Calculate the [HCO3

-], [H+], and [CO3=] in 0.14 M H2CO3.

K1 = 4.30 x 10-7 and K2 = 5.61 x 10-11. In summary: [HCO3

-] comes from the first dissociation and no significant amount dissociates in the second step. [H+] comes from the first dissociation, also with no significant amount produced in the second step.

[CO3=] comes from the second dissociation, only, and is quite small.

115

Write the K1 and K2 expressions for H2SO3 and then multiply them: Calculate the [SO3

=] in a solution that is 0.10 M in H2SO3 with a pH of 3.2. The K1 = 1.71 x 10-2 and the K2 = 6.0 x 10-8. Why was it reasonable to assume that the [H2SO3] is 0.10 M?

116

An important observation: Compare the first and second ionizations for acids. Consider the step-wise dissociation of H3PO4 and the equilibrium expressions. Write the combined K1K2K3 expression:

117

A Survey of Problems: Calculate the percent ionization and the pH when a regular aspirin tablet (325 mg) is dissolved in 200 mL water. Ka = 3.3 x 10-4, and the formula for aspirin is HC9H7O4 Place these 0.1 M solutions in order from lowest to highest pH: KF, CaCl2, and NH4Br. Explain your logic.

118

What volume of 1.00 M NaOH must be added to 400 mL of 0.16 M nicotinic acid, HC6H4NO2, to produce a buffer with a pH of 5.2? The Ka = 1.4 x 10-5. What is the pH of a 0.15 M solution of CH3H3NO3? The Kb for CH3NH2 is 3.7 x 10-4.

119

White vinegar is 5.0% by mass acetic acid in water with the solution having a specific gravity of 1.007. What is the pH? Calculate the pH at equivalence point when 0.12 M hydrocyanic acid is titrated with 0.16 M KOH. The Ka for hydrocyanic acid is 4 x 10-10.

120

Notes

121

During this unit on More Aqueous Equilibria, Dr. Etheridge will

1. describe the dynamic equilibrium occurring when a virtually insoluble substance is placed in water.

2. explain and differentiate between same-source and separate-source situations when dealing with Ksp.

3. calculate the molar solubility of several compounds. 4. describe the situations in which Q > Ksp, Q < Ksp, and Q = Ksp. 5. introduce the common ion effect and the impact on solubility. 6. discuss separating ions such as two cations existing in the same solution. 7. explain the impact of pH on solubilities to cause precipitation, prevent

precipitation, and dissolve types of compounds. 8. introduced the complex ion concept. 9. define key terms such as complex, central metal ion, ligand, donor atom,

coordination sphere, coordination number, monodentate ligand, polydentate ligand, and metal chelate.

10. list, explain, and give examples of the seven rules involved in naming complex ions.

11. discuss the relationship between dissolving and complex ions. 12. briefly note the relationship between the Kd and the Kf. 13. Review the concepts via a survey of problems.

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This unit will be broken into two parts: A: Solubility Product Constants

and B: Complex Ion Formation

During A: Solubility Product Constants, we will discuss the following:

Writing the Expression Molar Solubility Predicting Solubility The Common Ion Effect Separating Ions Effects of pH

WRITING THE EXPRESSION We will be dealing with substances that are very nearly insoluble. Using AgCl, describe the dynamic equilibrium existing between the solid and dissolved states: Write the equilibrium expression for the solubility product of AgCl. Why isn’t AgCl(s) involved in the equilibrium expression?

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Write the equilibrium expression for the Ksp of Ca3(PO4)2: Same Source vs. Separate Source Equilibria What is a “same source” situation? What is a “separate source” situation? MOLAR SOLUBILITY Same Source Situations Using AgCl(s) and its dissolving, write the Ksp expression. Note that this produces a saturated solution. What represents the concentration of the saturated solution? Is this a same source situation? ________________ Using PbCl2 and its dissolving, write the Ksp expression. Is this a same source situation? _______________ Using Ca3(PO4)2, write the Ksp expression: Is this a same source situation? ______________

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Using this, let’s calculate the molar solubility of some compounds. Calculate the molar solubility of AgCl. Ksp = 1.8 x 10-10. What volume of water is needed to dissolve 1.0 g AgCl? Calculate the pH of a saturated Co(OH)2. Ksp = 2.5 x 10-16

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PREDICTING SOLUBILITY The Solubility Rules are included in the Appendix in case you want to look at them. Separate Source Situations This often occurs when a compound is to be dissolved in a solution already containing a dissolved substance. Will a precipitate form when 1.00 mg AgNO3 is added to 1.00 L of a 0.0004 M NaCl? Ksp = 1.8 x 10-10 Describe the situation if Q > Ksp: Describe the situation if Q = Ksp: Describe the situation if Q < Ksp:

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REM: Molar solubility deals with __________________ solutions! When Q < K, the solution is unsaturated and there is no equilibrium When Q = K, the reaction is in equilibrium When Q > K, the reaction will move to remove ions by forming precipitates. What is the maximum mass of Sr(NO3)2 that will dissolve in a 0.0030 M Na2SO4? Ksp = 2.8 x 10-7

How many grams of Hg2(NO3)2 will dissolve in a 0.00020 M NaCl? Ksp for Hg2Cl2 = 1.1 x 10-18

How many grams of NaCl will dissolve in a 0.00020 M Hg2(NO3)2? Ksp for Hg2Cl2 = 1.1 x 10-18.

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Here’s a situation that is very common in chemistry: If the [Hg22+]

is 0.00010 M, what is the minimum [Cl-] needed to just begin precipitation of Hg2Cl2? Ksp for Hg2Cl2 is 1.1 x 10-18

If the [Hg2

2+] is 0.0010 M, what is the percent of Hg22+ remaining if the

[Cl-] is made 0.0040 M? See above problem for Ksp THE COMMON ION EFFECT Determine the quantity of AgCl that will dissolve in one liter of a solution that is 0.10 M NaCl. Ksp AgCl is 1.8 x 10-10

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Determine the quantity of Ag2CO3 that can dissolve in 500 mL of a 0.0050 M Na2CO3 solution. Ksp for Ag2CO3 is 8.1 x 10-12. SEPARATING IONS Separation of ions is the objective in qualitative analytical chemistry. After the mixture of ions is separated, each ion can be tested. The key, however, is to separate the ions in the first place. GIVEN: A solution is 0.15M in both Ni2+ and Cd2+. How could the two ions be separated from each other? How can this be done? If we use S= to effect the separation, it is necessary to be VERY careful to add enough to precipitate one of the ions without precipitating the other ion. A question must be asked: How particular must we be in the separation? If 70% of one ion is precipitated, is that enough? If 90% of one ion is precipitated, is that good enough? Will it take 99% of one ion being precipitated? Let’s have a look:

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It is necessary to separate the Ni2+ and Cd2+ from a solution that is 0.15 M in each ion. If we use S= for this separation:

a) which cation will precipitate first?

b) what percent of the first ion to precipitate will have precipitated when the second cation is ready to precipitate? Ksp for NiS = 3.0 x 10-21 Ksp for CdS = 3.6 x 10-29

130

A solution is 0.010 M in both AgNO3 and Pb(NO3)2. Which ion will precipitate first when Cl is added? What is the concentration of the ion which precipitates first when the second ion is just ready to precipitate? Ksp (AgCl) is 1.8 x 10-10 Ksp for PbCl2 is 1.7 x 10-5 EFFECTS OF pH Review of Diprotic Acids Recall the stepwise dissociation and the K1K2 using the example of H2A.

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The H2S Factor Why is/was H2S used extensively? What is the K1? ____________ K2? ___________________ What is the concentration of a saturated H2S solution? _________________ Therefore, the combined K1K2 expression for H2S is: Find the pH needed to prevent precipitation of ZnS in a saturated H2S solution that is 0.010 M in Zn(NO3)2. The Ksp for ZnS is 1.1 X 10-21.

132

Calculate the maximum [Cd3+] possible in a 0.30 M HCl saturated with H2S. The Ksp for CdS is 3.6 x 10-29. Dissolving Precipitates CaCO3, limestone, is insoluble in water, but is quite soluble in HCl(aq). Write the equation for the reaction of limestone with HCl(aq). FeS is insoluble in water but is soluble in HCl(aq). Write the equation for this dissolving. Thus, insoluble salts of weak acids may be dissolved in _________________________.

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Explain how a strong acid dissolves CaCO3. In this manner the concentration of one ion may be sufficiently reduced so that the product of the concentrations of the ions is less than the Ksp. How could you dissolve the insoluble Al(OH)3? How could you dissolve AgCN? What is the danger, here? There are other ways to dissolve these, but that is in COMPLEX IONS

COMPLEX ION FORMATION

During B: Complex Ion Formation, we will discuss the following: Definitions Naming Complex Ions Solubility and Complex Ions A Survey of Problems

134

DEFINITIONS Give definitions for the following: Complex: Coordinate covalent bonds: Examples: Central metal ion: Ligand: Donor atom: Coordination sphere: Coordination number: The most common coordination numbers are ___________________ Monodentate ligand Polydentate Metal chelate

135

Naming Complex Ions: Rule 1: Rule 2: Rule 3: Rule 4: Rule 5: Rule 6: Rule 7: Four ligands with special names: H2O NH3 Name these: [Co(NH3)6]Cl3 _______________________________________________ [Ag(NH3)2]+ _________________________________________________ [Ag(NH3)2]Br _______________________________________________ [Cr(H2O]4Cl2]Cl _____________________________________________ Na2[PtCl6] __________________________________________________

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SOLUBILITY AND COMPLEX IONS Write the equation for the reaction between Cu2+ and note the colors: Write the Kf What is the Kf? Write the equation for the formation of Fe(SCN)2+ and note the colors. Give the Kf expression The larger the Kf, the _____________________ the complex. Historically, these have been written as Kd instead of the Kf, but the more modern approach is to write them ad Kf What is the mathematical relationship between the Kd and Kf.

137

Consider the formation of the Ag(NH3)2+

Write the equation for both steps: Why is it necessary to add an excess of the complexing agent, NH3? Calculate the number of moles of NH3 required to dissolve 1.0 g of AgCl(s) into 1.0 L of water. How does this problem differ from other we have been working? In determining the quantity of Ag+ to be complexed, the expression does not show up well on the video. The expression was 0.0070 M – 0.000000026 M ≅ 0.0070 M Solve the problem:

138

Rework the above problem using the combined expression: A SURVEY OF PROBLEMS: A saturated solution of Co(OH)2 has a pH of 9.17. Calculate the Ksp of Co(OH)2

139

Calculate the molar solubility of Al(OH)3 in 0.10 M NaOH. Kf Al(OH)3) = 7.7 x 1033, Ksp Al(OH)3 = 3 x 10-34. First, however, Dr. Etheridge is going to have you find the molar solubility of Al(OH)3 in water. Her point will be made after the problem is worked. Calculate the molar solubility of AgCl in 6.00 M NH3. Kf [Ag(NH3)2]+ = 1.6 x 107, Ksp = 1.8 x 10-10. Make sure you put the plus sign on the diammine silver complex.

140

Calculate the [Cu2+] at equilibrium in a solution made by dissolving 0.20 mol CuSO4 in 1.00 L of 6.0 M NH3. Kf Cu(NH3)4

2+ = 1.1 x 1013. Calculate the molar solubility of AgCN with and without consideration

of hydrolysis. Ksp = 5.97 x 10-17, Ka = 4.0 x 10-10. (Oops, Dr. Etheridge left out Ag+, but she caught it. However, she may have a math error in this problem. x3 = 1.49 x 10-21, thus x = 1.1 x 10-7)

141

During this unit on Thermodynamics, Dr. Etheridge will

1. review thermodynamic terms studied in the earlier course, i.e. system, surroundings, exothermic, endothermic, enthalpy change, enthalpy of formation, state function, and the First Law of Thermodynamics.

2. discuss what spontaneity is and isn’t. 3. relate energy and spontaneity 4. explain why certain circumstances are spontaneous. 5. describe the properties of entropy. 6. introduce the Third Law of Thermodynamics. 7. discuss perfect order. 8. compare enthalpy and entropy changes and use this information to estimate select

physical properties, such as boiling point. 9. introduce the Second Law of Thermodynamics. 10. derive the ∆G. 11. work a series of problems involving ∆G. 12. use the relationships among the three state functions studied and draw conclusions

regarding certain reactions. 13. discuss the relationship between the ∆G and the Keq. 14. utilize the van’t Hoff equation to calculate the Keq at other than standard state

conditions. 15. relate time and spontaneity.

You will need to locate a table of thermodynamic properties of substances at standard state conditions. Such a table is too extensive for this noteguide.

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During this unit, we will discuss the following: Spontaneity ∆S ∆G The Keq A Survey of Problems

Review of Terms: Give definitions for System Surroundings Exothermic Endothermic Enthalphy Change Enthalpy of Formation

143

State Function 1st Law of Thermodynamics SPONTANEITY What it is and isn’t What is a spontaneous reaction? What are non-spontaneous reactions? Energy and spontaneity Give examples of spontaneous reactions/actions. Give examples of spontaneous reactions that gain energy. Entropy, S The designation for entropy is ______. Is it a state function? __________ Systems spontaneously move to _______________________________________. WHY? Putting things into order requires the input of ____________________.

144

Disorder is the result of _________________________________. Does whether or not the change is endothermic or exothermic matter? _______________ Another name for disorder is called _______________________. Give the characteristics of entropy: What are the units of entropy? _____________________ Can you make a general statement regarding the state of matter of a substance and its entropy? REM: The entropy values of the elements are NOT zero. 3rd Law of Thermodynamics How can you have a situation in which there is perfect order? Why must it be at absolute zero? State the 3rd Law of Thermodynamics: Why is it important?

145

∆S Phase Changes Phase changes are _________________________________ Compare the enthalpy change in this process: H2O(l) H2O(g) -285.830 kJ/mol -241.818 kJ/mol Calculate ∆H for the above process: The process is _________________________. What does that say about the entropy or randomness? Repeating the above for entropy H2O(l) H2O(g) 69.91 J/mol.K 188.7 j/mol.K Calculate the ∆S for the process? What is this value? As heat is absorbed, randomness ____________ Therefore, T∆S = and ∆S is proportional to ______________

146

What happens to the process if we apply heat? ∆S = Which will have the greater impact on randomness? Heating a solid or heating a gas? WHY? State ∆S for phase changes:

∆S = ≅ for phase changes (the “proportional” sign didn’t show well on the TV screen)

Again, look at the boiling of water: ∆S = ________________________ ∆H = _________________________ Estimate the temperature at which boiling of water should occur:

147

Changes in Reactions Consider: CaCO3(s) CaO(s) + CO2(g) 92.9 J/mol.K 39.75 j/mol.K 213.74 J/mol.K Calculate the ∆S for the reaction. Should we expect the reaction to be spontaneous? Why? If we were to reverse the reaction, how would the ∆S compare to the ∆S above? Predicting Spontaneity In summary 1. 2.

3. 4.

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2nd Law of Thermodynamics State the law (in more than one way): When ∆Suniv is positive (+), the reaction is _________________________ When ∆Suniv is negative (-), the reaction is __________________________________ ∆G Give the term that ∆G represents: Deriving ∆G: What does the Gibbs Free Energy combine? What does it allow us to do? Derive the ∆G: (be careful, some of the minus (-) signs on the video look like small =)

149

What does this mean? ∆Gsystem =

and finally, ∆Gsystem = More simply stated: ∆G = Consider the relationship at constant temperature If ∆G is < 0, the process is _____________________ If ∆G is > 0, the process is _______________________________ If ∆G is = 0, the process is _______________________________ Calculating the ∆G: (refer to thermodynamics table in the back of a textbook or CRC Handbook) Calculate the ∆G for the formation of CO2 using ∆Ho and So The ∆G for an element at standard state conditions is __________________ Again, consider: ∆G = ∆H - T∆S If ∆G is negative, the reaction is ___________________ If ∆H is positive, can ∆G be negative? If T∆S is positive, can the reaction be spontaneous?

150

Consider dissolving sufficient NH4NO3 to produce a 1.0 M solution. Write the equation and include the thermodynamic values needed: Using the relationships among the three state functions, what conclusions can you draw?

The Temperature Factor Consider the process: 2 Fe2O3(s) + 3 C(s) → 4 Fe(s) + 3 CO2(g) ∆H (2)(-824.2 kJ/mol) (3)(0) (4)(0) (3)(-393.5 kJ/mol)

∆H = +467.9kJ

∆G (2)(-742.2k/mol) (3)(0) (4)(0) (3)(-394.359 kJ/mol) ∆G = +301.323 kJ S (2)87.40J/molK (3)5.74J/molK (4)27.28J/molK (3)213.74 J/molK ∆S = +560.32 J/K The reaction is endothermic/exothermic? ________________________ The reaction is/is not spontaneous? ____________________________

151

Can we alter the temperature to make this reaction feasible and spontaneous? Calculate the minimum temperature at which the ∆G is no longer positive, ∆G = 0

Now, is the reaction feasible?

THE Keq The Q/Keq Relationship Give the “formula” for Q Using A + 2B → C + D Give Q Give Keq recall: If Q > Keq, the reaction _________________________________ If Q > Keq, the reaction _________________________________ If Q = Keq, the reaction _________________________________

152

Reaction ∆G Q/Keq ln(Q/Keq)

spontaneous not spontaneous at equilibrium

Notes: Calculating the Keq Calculate the ∆G for the ionization of acetic acid. What does that tell us?

153

Limitation of the Gibbs Free Energy Expression: How can we figure the equilibrium constant at some other temperature? Let’s look at the van’t Hoff equation: Write this equation: Calculate the Keq for the decomposition of CaCO3 at 25oC and 100oC. Refer to thermodynamics tables in a textbook or CRC Handbook Write the equation: Calculate the ∆G Calculate the Keq at 25oC. Using ∆H and the van’t Hoff equation, calculate the Keq at 100oC. Compare the Keq at 25oC and at 100oC.

154

∆G and the Keq

Consider the equation for the evaporation of ethanol, a liquid alcohol: C2H5OH(l) C2H5OH(g) ∆G = 7.00 kJ Calculate the Keq for the evaporation of alcohol. In the next lesson we will discuss how to make the ∆G have a wider range of application.

155

How can the ∆G be made useful at other temperatures? Give the formula: Consider the reaction: Br2(l) + Cl2 2 BrCl(g) at 25oC. P(Cl2) = 0.010 atm P(BrCl) = 0.22 atm and we do not know if the reaction is in equilibrium. In which direction will the reaction proceed? Time Be a philosopher and relate time and spontaneity.

156

A SURVEY OF PROBLEMS For the reaction: CuS(s) + H2(g) → Cu(s) + H2S(g)

(a) calculate K at 298.15K and 1.00 atm. (b) estimate the K at 510oC and 1.00 atm (c) calculate the ∆S at 298.15K and 1.00 atm (d) estimate ∆G at 510oC and 1.00 atm (e) estimate the temperature above which the reaction is expected to be

spontaneous

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During this unit on Electrochemistry, Dr. Etheridge will

1. give a brief review of oxidation and reduction. 2. review the activity series and develop the concept of why it works that way. 3. discuss standard electrode potentials from the standpoint of reduction. 4. introduce the voltaic cell (galvanic cell), specifically the Daniell cell, and address

the types of reaction occurring at each electrode. 5. expand the galvanic cell concept, including batteries, and add standard cell

notation. 6. explain the SHE and why it is important. 7. discuss concentration, electrode potential, and ∆G. 8. describe the chemistry of the dry cell battery and the lead storage battery. 9. introduce electrolysis in general and several important electrolytic cells. 10. develop the concept of equivalent weight using Faraday’s Law. 11. describe the chemistry of corrosion and corrosion protection. 12. conclude with a brief survey of problems.

158

During this unit, we will discuss the following: Introduction Voltaic Cells Electrolytic Cells Corrosion

INTRODUCTION Review of Oxidation-Reduction: Oxidation: Reduction: Oxidizing agent: Reducing agent: Cu + HNO3 → Cu(NO3)2 + NO + H2O What’s oxidized? ____________ What’s reduced? ___________________ What’s the reducing agent? __________________ oxidizing agent? ______________

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Single Replacement Series Remember how it works? Zn + HCl → Cu + ZnCl2 → Why does the activity series work that way? Standard Electrode Potentials Now let’s look from the tendency to GAIN electrons: Zn + H+ → WHY? Cu + Zn2+ → WHY?

160

Putting It All Together: Stop the video and locate a table of standard reduction potentials at 25oC. There is a brief table on page 197 in this book. However you can find one at the web site: chemistrygeek.com/reductiontable.pdf or a more extensive one is found at en.wikipedia.org/wiki/Table_of_standard_electrode_potentials. Using this table, give a couple of examples to show how it works. VOLTAIC CELLS Electrochemical Cells: Explain this: Zn2+ + Cu → no reaction Notes:

161

Label this Daniell cell as shown: Explain how this works: Specific Galvanic Cells: Define a galvanic cell: Another name for a galvanic cell is a ________________________. REM: Anode: Cathode: Daniell Cell:

162

Cell Notation Show the method of using cell notation to describe a galvanic cell: Give the cell notation for a Daniell cell Give the cell notation for Ni as the anode and Br2 as the cathode: (Do put in a salt bridge)

It would be possible to use a non-reactive solid material such as Pt or C for the electrode in the reaction involving Br2 and the Br-.

Calculate the voltage of this cell: Ni(s) | Ni2+ || Br1- (1.0 M)| Br2(l)

163

Here are two half-cells. Write the cell notation and calculate the voltage when they are arranged in a galvanic cell. Fe3+ + 1 e- → Fe2+ 0.769 v ClO3

1- + H2O + 2 e- → ClO2- + 2 OH- 0.271 v

Dr. Etheridge doesn’t mean it is in your textbook, rather she is referring to this noteguide. Consider the reaction of Cu and HNO3. Write the half-cells as they should go. Pay very close attention to what the copper actually reacts with. Explain how you arrive at the half cells to be used and their direction: The Standard Hydrogen Electrode (SHE) What is the SHE and why is it important? Describe the set-up of the SHE

164

More about galvanic cells: On what does the voltage depend? On what does the voltage NOT depend? What is the difference among the various batteries, AAA, AA, C, D, etc. Concentration and Electrode Potential Recall: ∆G = -RT lnK When an oxidation-reduction reaction occurs producing an electrical current, is the reaction spontaneous? The electrical current produced is the _______________________________________ What is the relationship between the ∆G and the standard cell potential? What is a Faraday? What is a coulomb?

165

Examine the relationship ∆G=-nFE using the Daniell Cell: Is the reaction spontaneous? How do you know? For the reaction: 2 Al(s) + 3 Cu2+ → 3 Cu(s) + 2 Al3+ the ∆G = -1165 kJ. What would be the maximum voltage at standard state conditions?

166

Concentration and Electrode Potential (cont.) Does the voltage change as the concentration changes? Answer the question by deriving the expression relating moles, Faradays, etc, i.e. the Nernst equation. What happens at standard state conditions and equilibrium? Thus, there is a relationship between the voltage and the equilibrium constant. Calculate the voltage for the Daniell cell if the [Cu2+] = 0.02 M and the [Zn2+] is 1.4 M.

167

Primary and Secondary Cells What is the difference between primary and secondary cells? Give examples of primary cells: Give examples of secondary cells: Describe the chemistry of the “Dry Cell Battery” Why can’t it be recharged? Describe the chemistry of the lead storage battery:

168

ELECTROLYTIC CELLS Electrolysis Describe an electrolytic cell: Electrolysis of Molten (fused) NaCl Draw and describe the reactions that occur during the electrolysis of molten NaCl: Note the anode and cathode, their charges, and the specific reactions occurring there. Electrolysis of aqueous NaCl Draw and describe the reactions that occur during the electrolysis of NaCl(aq): Note the anode and cathode, their charges, and the specific reactions occurring there.

169

Electrolysis of water: Draw and describe the reactions that occur during the electrolysis of water: Note the anode and cathode, their charges, and the specific reactions occurring there. REMEMBER Oxidation ALWAYS occurs at the ____________________

In a galvanic cell, the anode is _______________

In an electrolytic cell, the anode is ______________________ Faraday’s Law What is the equivalent weight of a metal? What is Faraday’s Law? Get these definitions: 1 faraday 1 coulomb

170

Electroplating Calculate the mass of copper plated out when 1.60A of current is applied to a CuSO4 solution for 1.0 hour.

CORROSION The Chemistry of Corrosion What do we generally mean by “corrosion”? Describe the circumstances where corrosion is most likely to occur. Describe the process for iron that explains the ready corrosion of iron. Note the anodic and cathodic reactions:

171

Complete the explanation of iron corrosion by showing how “rust” is formed. Corrosion Protection How can we protect against corrosion? Explain “passivity.” Why do we coat iron with zinc and how does it work? (Use the term “sacrificial anode.” Explain the chemistry of “tarnish.” (Error: silver sulfide is Ag2S)

172

A Survey of Problems How many grams of cobalt will be deposited from a solution of CoCl2 using a 20.0 amp current for 54.5 minutes? (Watch out! Dr. Etheridge wrote down the wrong number of coulombs. It should

have calculated to be 65,400 C. BUT, she carried the correct value in her calculator and came out with the correct answer in the end.)

To which electrode should a spoon be attached in order to silver plate a spoon from a Ag+ solution? Explain. If you decided to plate a spoon from a block of silver, to which electrode should the block be attached?

173

A current of 9.0 amps flowed for 45 minutes through water containing a small quantity of NaOH. How many liters of gas were formed at the anode at 27.0oC and 750 Torr?

174

Notes:

175

During this unit on Nuclear Chemistry, Dr. Etheridge will

1. note the differences between nuclear and chemical reactions. 2. review isotopic notation and isotopes. 3. introduct mass defect and the binding energy it involves. 4. describe the characteristics of stable and unstable nuclei, including the band of

stability. 5. give a series of common emissions and their impact on the nucleus. 6. explain K capture. 7. discuss Radioactive Decay Law. 8. review carbon dating and how it works. 9. discuss fission vs. fusion 10. introduce nuclear reactors, parts, and common fuels. 11. describe Breeder Reactors and the problems inherent in them. 12. briefly describe several important uses of radioactivity.

176

During this unit, we will discuss the following: Introduction Nuclear Stability Disintegrations Fission vs. Fusion Uses of Radioactivity

INTRODUCTION Briefly note the work of

Henri Becquerel

Marie Curie

Pierre and Marie Curie Note the differences between nuclear reactions and chemical reactions:

177

NUCLEAR STABILITY The Nucleus Recall isotopic notation and write the notation for U-238. Describe the characteristics of the nucleus: Isotopes: What are they? Give names for and compare occurrences of hydrogen isotopes Where are they found/used? Are they all radioactive?

178

Mass Defect What holds the nucleus together? Where does the energy come from? Explain mass defect and the energy it involves (include “binding energy”) Neutron-Proton Ratio: Describe most naturally occurring nuclei: Describe the characteristics of exceptionally stable nuclei (include the magic numbers)

179

What do magic numbers suggest? Explain and draw the “band of stability.”

One theory about how the nucleus is held together via electron exchange. Make notes if you choose.

180

DISINTEGRATIONS Common Emissions

Type Symbol Identity Charge penetration

beta

positron

alpha particle

proton

neutron

gamma

How does a nucleus reduce the number of neutrons? Explain. Write the equation

Give an equation showing the loss of a neutron: How can a nucleus increase the neutron to proton ratio? Another emission to accomplish this is….

181

What happens in positron emission? Explain K capture: Explain alpha (α) emission: Give a series of balanced equation for the disintegration of U-238 by α, α, β, β, followed by K-capture. (A periodic table appears on the next page)

182

Half-Life Radioactive decay follows _____________________________________ Give the formulas for R and –kt and half life: State Radioactive Decay Law:

183

Dr. Etheridge gives the derivation for several equations from the kinetics studied earlier. Give the derivations if you see fit or at least give the newly stated equations for use in radioactivity: Co-60, used in the treatment of cancer, has a half life of 5,271 years.

(a) What is the decay constant for a 3.00 mg sample of Co-60? (b) What is the activity of this sample? and (c) What will the rate of decay become after 10 years?

184

Carbon Dating (Review) Explain how it works: C-14 decays by ____________ and has a half-life of ______________ years Important Series Heavy elements tend to undergo _____________________emissions Disintegration of heavy elements eventually wind up producing ___________________. When Dr. Etheridge referred to the “activity series” as she discussed heavy elements, she was actually referring to the periodic table. No elements above (past) ___________ have any stable isotope FISSION VS. FUSION Fission: Give the equations for the bombardment of U-235 with a neutron: Explain how this sets off a chain reaction: How can a chain reaction be prevented? What is meant by a “critical mass”?

185

How can these reactions be controlled? Explain. Fusion: What is fusion? When the two scientists announced their ability to perform “cold nuclear fusion” it was very exciting. If, indeed, it had happened, why would it be so important? Nuclear Reactors; What are the common fuels? Draw the fundamentals of the reactor and label the parts:

186

Explain how the system produces electricity: How is the reaction controlled? What are Breeder Reactors? How do they work? Explain how it “produces more fuel than it consumes.” Why don’t we use breeder reactors?

187

USES OF RADIOACTIVITY You should list and know these uses of radioactivity: Begin with U-235. Write a series of balanced equations showing an α, β, 1n emission, followed by a K-capture. A wooden artifact is shown to contain 6.8 x 10-6g C-14. A living piece of wood contains 5.8 x 10-5g of C-14. How old is the artifact? t1/2 for C-14 is 5730 years.

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NOTES:

189

APPENDIX

TABLE OF CONTENTS

Table of Common Oxidation Numbers 191 Activity Series of the Elements 193 Solubility Rules 195 Select Solubility Product Constants 195 Select Reduction Potentials 197 Equilibrium Constants for Weak Acids 199 Equilibrium Constants for Weak Bases 199 Table of Selected Indicators 201 Periodic Table of the Elements 203

190

191

Table of Common Oxidation Numbers

1+ elements 2+ elements 3+ elements 4+ elements H+ Mg++ Al+++ Pb++++ (plumbic) Li+ Ca++ Bi+++ Sn++++ (stannic) Na+ Sr++ Cr+++ (chromic) K+ Ba++ Fe+++ (ferric) Cs+ Zn++ Mn+++ (manganic) 5+ elements Ag+ Cd++ Co+++ (cobaltic) Cu+ (cuprous) Cr++ (chromous) Ni+++ (nickelic) As+++++ (arsenic) Hg2

++ (mercurous) Mg++ (manganous) As+++ (arsenous) Fe++ (ferrous) Co++ (cobaltous) Ni++ (nickelous) 1+ radicals* Sn++ (stannous) Pb++ (plumbous) H3O+

(hydronium) Cu++ (cupric) NH4

+ (ammonium) Hg++ (mercuric) 1- elements 2- elements 3- elements H- (hydride) O= N≡

F- S= P≡

Cl- Br- I- 1- radicals* 1- radicals (cont) 2- radicals* 3- radicals* C2H3O2

- (acetate) HSO4- (bisulfate) SO4

= (sulfate) BO3≡ (borate)

CN- (cyanide) HSO3- (bisulfite) SO3

= (sulfite) PO4≡ (phosphate)

NO3- (nitrate) HS- (bisulfide) S2O3

= (thiosulfate) PO3≡ (phosphite)

NO2- (nitrite) HCO3

- (bicarbonate) CO3= (carbonate) AsO4

≡ (arsenate) OH- (hydroxide) HC2O4

- (binoxolate) C2O4= (oxalate) AsO3

≡ (arsenite) SCN- (thiocyanate) O2

= (peroxide) ClO- (hypochlorite) SiO3

= (silicate) ClO2

- (chlorite) CrO4= (chromate)

ClO3- (chlorate) Cr2O7

= (dichromate) ClO4

- (perchlorate) BrO3

- (bromate) BrO4

- (perbromate) MnO4

- (permanganate) * The term “radicals” is an older term for polyatomic ions.

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193

ACTIVITY SERIES OF THE ELEMENTS

K Ba Sr Ca Na Mg Al Mn Zn Cr Fe Cd Co Ni Sn Pb H Sb As Bi Cu Ag Pd Hg Pt Au

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195

Solubility Rules

1. All nitrates, acetates, compounds of sodium and potassium, and all ammonium compounds are soluble.

2. All chlorides are soluble except those chlorides of silver, mercurous, and lead, although PbCl2 is moderately soluble in hot water.

3. All sulfates are soluble except those of barium and lead. Calcium sulfate, mercurous sulfate, and silver sulfate are only slightly soluble.

4. All carbonates and phosphates are insoluble except those of sodium, potassium, and ammonium.

5. All hydroxides are insoluble except sodium, potassium, ammonium, and barium.

6. All sulfides are insoluble except those of sodium, potassium, and ammonium and the alkaline earth metals.

7. There are a few uncommon compounds of sodium, potassium, and ammonium that are not soluble.

8. All silver compounds are insoluble except silver nitrate and silver perchlorate. Two important moderately soluble silver salts are silver acetate and silver sulfate.

Select Solubility Product Constants Formula Ksp

AgCl 1.8 x 10-10

AgBr 5 x 10-13

AgI 8.5 x 10-17

Hg2Cl2 1.1 x 10-18

PbCl2 1.7 x 10-5

PbI2 8.3 x 10-9 AgCO3 8.1 x 10-12 BaCO3 1.6 x 10-9 CaCO3 4.8 x 10-9 CuCO3 2.5 x 10-10

Ca(OH)2 1.3 x 10-6 Ag2SO4 1.7 x 10-5 BaSO4 1.5 x 10-9 SrSO4 2.8 x 10-7 Ag2S 1 x 10-50 (?) ZnS 1.1 x 10-21

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Select Reduction Potentials

K+ + 1 e- → Ko -2.924

Ca2+ + 2 e- → Cao -2.840

Na+ + 1 e- → Nao -2.713

Mg2+ + 2e- → Mgo -2.356

Zn2+ + 2e- → Zno -0.763

Fe2+ + 2e- → Feo -0.440

Ni2+ + 2e- → Nio -0.257

Pb2+ + 2e- → Pbo -0.125

2 H+ + 2e- → H2o 0.00

Cu2+ + 1e- → Cu+ +0.520

Cu2+ + 2e- → Cuo

+ 0.34

Fe3+ + 1e- → Fe2+ +0.771 Br2(l) + 2e- → 2 Br- +1.065

Ag+ + e- → Ago + 1.98

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EQUILIBRIUM CONSTANTS FOR WEAK ACIDS

ACID Ka Acetic HC2H3O2 or HOAc 1.76 x 10-5

K1 4.3 x 10-7 Carbonic H2CO3 K2 5.61 x 10-11

Formic HCOOH 2.1 x 10-4 Hydrocyanic HCN 4.0 x 10-10

Hydrofluoric HF 6.9 x 10-4

K1 1.0 x 10-7 Hydrosulfuric H2S K2 1.3 x 10-13

Nitrous HNO2 4.5 x 10-4 K 1 3.8 x 10-2

Oxalic H2C2O4 K2 5.0 x 10-5 K1 7.1 x 10-3 K2 6.3 x 10-8 Phosphoric H3PO4 K3 4.4 x 10-13

K1 1.71 x 10-2 Sulfurous K2 6.0 x 10-8

EQUILIBRIUM CONSTANTS FOR WEAK BASES

BASE Kb Ammonium hydroxide NH4OH 1.8 x 10-5 Aniline C6H5OH 7.4 x 10-10

Ethylamine C2H5OH 4.3 x 10-4 Methylamine CH3OH 4.2 x 10-4

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TABLE OF SELECTED INDICATORS

NAME pH RANGE COLORS pKa

Methyl violet 0.0 – 1.6 Yellow to blue Thymol blue 1.2 - 2.8 Red to yellow 1.65 Methyl orange 3.2 - 4.4 Red to yellow 3.46 Bromophenol blue 3.0 - 4.6 Yellow to blue 4.10 Congo red 3.0 -5.0 Blue to red Bromocresol green 3.8 - 5.4 Yellow to blue 4.90 Methyl red 4.4 - 6.2 Red to yellow 5.00 Bromocresol purple 5.2 - 6.8 Yellow to purple 6.40 Litmus 4.5 - 8.3 Red to blue Bromothymol blue 6.0 -7.6 Yellow to blue 7.30 Phenol red 6.8 - 8.2 Yellow to red 8.00 Thymol blue 8.0 - 9.6 Yellow to blue 9.20 Phenolphthalein 8.3 - 10.0 Colorless to red 9.5 Alizarin yellow R 10.0-12.0 Yellow to red Indigo carmine 11.4-13.0 Blue to yellow

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PERIODIC TABLE OF THE ELEMENTS 1A 8A 1

1 H 1.0079

2A

3A

4A

5A

6A

7A

2 He 4.0026

2

3 Li 6.941

4 Be 9.012

5 B 10.811

6 C 12.011

7 N 14.007

8 O 15.999

9 F 18.998

10 Ne 20.180

3

11 Na 22.990

12 Mg 24.305

3B

4B

5B

6B

7B

8B

1B

2B

13 Al 26.982

14 Si 28.086

15 P 30.974

16 S 32.066

17 Cl 35.453

18 Ar 39.948

4

19 K 39.098

20 Ca 40.078

21 Sc 44.956

22 Ti 47.88

23 V 50.942

24 Cr 51.996

25 Mn 54.938

26 Fe 55.847

27 Co 58.933

28 Ni 58.693

29 Cu 63.546

30 Zn 65.39

31 Ga 69.723

32 Ge 72.61

33 As 74.922

34 Se 78.96

35 Br 79.904

36 Kr 83.80

5

37 Rb 85.468

38 Sr 87.62

39 Y 88.906

40 Zr 91.224

41 Nb 92.906

42 Mo 95.94

43 Tc (98)

44 Ru 101.07

45 Rh 102.91

46 Pd 106.42

47 Ag 107.87

48 Cd 112.41

49 In 114.82

50 Sn 118.71

51 Sb 121.76

52 Te 127.60

53 I 126.90

54 Xe 131.29

6

55 Cs 132.90

56 Ba 137.33

57 La 138.91

72 Hf 178.49

73 Ta 180.95

74 W 183.84

75 Re 186.21

76 Os 190.23

77 Ir 192.22

78 Pt 195.08

79 Au 196.97

80 Hg 200.59

81 Tl 204.38

82 Pb 207.2

83 Bi 208.98

84 Po (209)

85 At (210)

86 Rn (222)

7

87 Fr (223)

88 Ra 226.03

89 Ac 227.03

104 Rf (261)

105 Db (262)

106 Sg (263)

107 Bh (262)

108 Hs (265)

109 Mt (266)

110 (269)

111 (272)

112 (277)

unk

114

unk

116 unk

118

LANTHANIDES

58 Ce 140.12

59 Pr 140.91

60 Nd 144.24

61 Pm (145)

62 Sm 150.36

63 Eu 151.96

64 Gd 157.25

65 Tb 158.93

66 Dy 162.50

67 Ho 164.93

68 Er 167.26

69 Tm 168.93

70 Yb 173.04

71 Lu 174.97

ACTINIDES

90 Th 232.04

91 Pa 231.04

92 U 238.03

93 Np 237.05

94 Pu (244)

95 Am (243)

96 Cm (247)

97 Bk (247)

98 Cf (251)

99 Es (252)

100 Fm (257)

101 Md (258)

102 No (259)

103 Lr (260)

Elements 110, 111, 112, and 114 have not been named. Proposed elements 116 and 118 are in question.