Required Materials:
General Chemistry: Whitten, Davis, Peck, and Stanley Thomson Brooks/Cole Pub. 10th Ed., 2014
General Chemistry II: Course Manual Alfare, Carlo MCCC, 13th Edition 2019
General Chemistry II: Laboratory Manual Alfare, Carlo MCCC, 14th Edition 2019
Catalog Description:
Theoretical and practical aspects of kinetics, simple and ionic chemical equilibria, thermodynamics, spectrophotometry,electrochemistry, nuclear chemistry, and the major families of chemical elements, with emphasis on the coordination chemistry of the transition elements. The laboratory work includes qualitative cation and anion analysis.
Corequisite: MAT 146 or higher
Carlo Alfare Professor of Chemistry & Course Coordinator
COURSE OUTLINE TABLE
OF CONTENTS
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Grading will be based on the point system as indicated below.
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Minimum Course Grade Assignment:*
A 605 Points (93%) B - 520 Points (80%)
A - 585 Points (90%) C + 500 Points (77%)
B + 565 Points (87%) C 440 Points (68%)
B 540 Points (83%) D 380 Points (58%)
*Acceptable laboratory and recitation participation and performance along with a passing grade on the final examination are required to pass the course. See the Course Objectives for more details.
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I
Chemical
Thermodynamics |
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0. Math Preparation on Logarithms 1. Thermochemistry |
Chap. 13: Liquids and Solids Chap. 15: Chemical Thermodynamics Whitten: Chapter 13 |
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II Chemical Kinetics |
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1. Thermochemistry (complete) | Chap. 16: Chemical Kinetics |
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III Chemical Equilibrium |
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2.Spectrophotometry | Chap. 17: Chemical Equilibrium |
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Spectrophotometry (Lab Topic Only) |
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2.Spectrophotometry
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Chap. 4: The Structure of Atoms Whitten: Chapter 4 |
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EXAMINATION I | WEEKS 1 - 4 | OBJECTIVES I - III | |
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IV Electrochemistry |
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3 Equilibrium | Chap. 21: Electrochemistry Whitten: Chapter 11 |
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V Acids and Bases |
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3 Equilibrium | Chap. 10: Reactions in Aqueous Whitten: Chapter 10 |
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VI Ionic Equilibria |
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4. Electrochemistry & Electrochemical Cells |
Chap. 18, 19: Ionic Equilibria: Acids, Bases and Buffers CHE 101 Exp. 10 & 11 |
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VI Ionic Equilibria |
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5. pH Meter and pH Measurements |
Chap. 20: Ionic Equilibria: Solubility Product |
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EXAMINATION II | WEEKS 5 - 8 | OBJECTIVES IV - VI | |
9 | VII Chemistry of the Representative Elements I: The Metals | 27 | 6. pH Titration | Chap. 27: Metals II: Properties and Reactions Whitten: Chapter 5 & 7 |
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VIII Chemistry of the Rep. Elements II: The Metalloids and Nonmetals |
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7. Solubility Product | Chap. 28: Nonmetals and Metalloids Whitten: Chapter 8 |
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IX The Transition Elements |
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8 Qualitative Analysis | Chap. 26: Metals I: Metalurgy |
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IX Coordination Chemistry |
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8 Qualitative Analysis | Chap. 25: Coordination Compounds Whitten: Appendix B |
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X Nuclear Chemistry |
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8 Qualitative Analysis | Chap. 22: Nuclear Chemistry Whitten: Chapter 6 |
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XI Organic Chemistry |
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13. Organic Model Building |
Chap. 23:Organic Chemistry I:Formulas, Names and Properties Chap. 24: Organic Chemistry II: Shapes and Reactions |
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Homework assignments are on a weekly basis to help
you learn the course material according to the performance objectives and
to help you to test your mastery of the material considered. They are not
to be considered "exclusive" but representative of the material. If you
feel the need to do additional reading or problems, you can ask your instructor
for guidance. All but the last column will be found in Whitten, Davis,Peck, and Stanley.
Read
Week Chap. Questions and Problems Course Manual
1a 13.9-11 48, 50, 52 Unit I: 1 - 5
1b 15 1-6, 11 - 14, 16, 20, *Calc. DeltaGo &
24 - 26,28, 34, 38b*, 59, So(Fe2O3)
60, 72 - 74, 76, 78, 83 – 85,
1c 15 92 - 96, 102,104a, 105, 106, 107a Unit I: 6 - 8
111a, 112a, 114, 115, 117, 121, 122
2a 16 1 - 6, 8b, 10, 11 Unit II: 1 – 3
2b 16 13 - 17, 19, 20, 30, 49 - 52, Unit II: 4 - 5
65, 66, 69, 70, 72
2c 16 22, 28, 18(optional) Unit II: 6 - 7
3a 17 1 - 8, 14, 17, 19, 22, 24, 28, 32, Unit III: 1 - 2
37 – 42, 53 - 56, 58 - 61, 69,
70, 76, 78, 81, 82, 84, 86
3b 17 44 - 46, 65 Unit III: 3 - 5
4 21 1, 2, 4, 5, 8, 10, 12, 16 - 19, Unit IV: 1 - 3
21, 24, 26, 32, 36, 38, 44
5 21 47 - 50, 54, 55, 58, 62, *Calc. Delta Go & Delta G
78 - 81, 87a,b*, 98, 100a,b
6a 10 1 - 4, 6 - 13, 17 - 20, 24 – 27, Unit V: 1 - 4
31 - 36, 39 - 41, 43 – 47, 50,
57, 61, 63, 67 - 69
6b 5.9 67 - 71 Unit V: 5 – 7
7a 18 1 - 3, 4a, 6c, 8c, 10 - 12, 15, 17, Unit VI: 1
22a,d, 24, 26, 30, 31, 36, 38,
40, 42, 47, 48, 54 - 57
7b 19 4,7,10,12,19,22,32,40,48,53 Unit VI: 2
7c 18 64, 68 – 72, 75, 79 – 82, 86 Unit VI: 3
8 20 1-6, 8b, 18, 20*, 28 Unit VI: 4 - 7
* & in 0.12 M AgNO3
& in pure H2O
9a 27 1- 4, 6 – 14,16, 18 – 20, 22a, Unit VII: 1 – 11
24a, 25 - 31
9b 28 1-5, 7, 9, 12, 24, 33 Unit VIII: 1 - 17
10 28 8, 11, 13, 15, 16, 19 - 21,
25 - 32, 34, 35, 38 - 40,
45-55, 66,73
11a 27 33 – 38, 41 – 43, 47 Unit IX: 1 - 20
11b 26 1 – 5, 7, 9 - 13, 15, 22, 26a, 33 Unit IX: 21 - 25
12a 25 1 - 8, 15, 24, 26, 28 Unit IX: 26,27
12b 25 32a, 36a,e, 38, 39, 45a,c, Unit IX: 28 - 31
46a,b,c, 47*, 48, 49a,b *Also analyze with the
VB Theory
13 22 1, 3, 4, 6, 9, 20 - 22, 25-28, Unit X: 1 - 15
31a,d, 34,38, 40, 50 - 52,
54, 63 - 65, 67, 68, 70 - 72
14a 23 1 - 15, 18 - 30, 32, 33, 36, Unit XI: 1 - 5
40c-d, 41 ,44, 45, 48 - 52,
56, 57, 60, 65, 68, 69, 71
14b 24 1 – 5, 8, 10, 11 Unit XI: 6
23 79, 80, 85 - 87, 91, 93 Unit XI: 7 – 8
NOTE: Anyone taking Organic Chemistry before General Chemistry II will not be graded on Organic Quizzes and will receive a prorated grade for Experiment 9.
Textbook OWLv2 On-Line Mastery and Homework Assignments: This should be your
major source of extra help in this course.
Tutoring sites on my home page: www.mccc.edu/~calfare
Science Learning Center (MS 211) for help with labs and homework
Chemistry Survival Skills by Margaret Brault and Margaret MacDevitt. This book will help the chemistry student be more successful in the course.
Chemistry Survival Skills computer disk by Stanley and Ruth Chabrey. See Appendix E of this manual for a description.
Problem Solving for General Chemistry by Leslie Kinsland. Contains more practice problems with answers to all exercises.
Student Study Guide by Raymond Davis. Chapter summaries, study goals, 80 drill and concept questions per chapter with answers.
Student Solutions Manual by Yi-Noo Tang and Wendy Keeney-Kennicutt. Answers and solutions to all even-numbered end-of-chapter exercises.
Schaum's Outline of College Chemistry by Jerome
Rosenberg. Theory and problems (778) with complete solutions.
Goggles and gloves for the laboratory.
Mercer's Academic Integrity Policy:
Mercer County Coummunity College is committed to academic integrity – the honest, fair and continued pursuit of knowledge, free from fraud or deception. Read the booklet on Academic Integrity. Violations will result in failure in the test, lab or , if serious or repeated, the course.
Academic integrity is violated when a student:
Violators will be penalized in accordance with college policy.
General Chemistry II is intended
to extend your initial exposure to a broad realm of fundamental concepts
in chemistry. It will assist you in attaining a basic understanding of
these concepts, and it will help you to develop essential skills in these
areas. The lectures, recitation discussions, laboratory sessions, homework
assignments, quizzes, and examinations provide an integrated selection
of activities which can lead you to success, provided that you are conscientious.
The average student should spend as many hours outside of class, studying
and working on the course, as they spend in class. In order to receive
credit for the course, you must at least meet the minimum requirements
described below. Additional effort and achievement will be especially rewarding,
however.
Participation in Chemistry laboratory courses is permitted provided the student has completed the required prerequisites, is a minimum of 16 years of age, or by permission of the instructor and the Dean of the division.
It is the college policy that a student taking the class as an Audit must declare this at the time of registration, and may not attend the laboratory, may not take exams, and may not have quizzes graded.
If you need an accommodation, you must bring the form at least 2 weeks before it will be used.
Students behavior deemed unsafe by the laboratory instructor will be grounds for removing the student from the course with a grade of WI or F.
1. You must satisfactorily complete the assigned laboratory experiments. (Missing 3 or more will constitute an F or W for the course)
2. You must participate in weekly recitations (missing 3 or more may constitute an F for the course)
3. You must complete the weekly quizzes and hour tests, as assigned.
4. You must achieve a passing grade on a comprehensive final examination.
5. You must complete a minimum of six hours of work on chemistry at home each week.
6. You must demonstrate your level of performance (see page 3 for "grading") by mastering a large part of the material covered by lectures, films, homework, laboratory work and the textbooks as detailed in the specific course objectives that follow.
7. With a C or better in CHE 101 you will be expected to know the material in that course and you can expect test questions from that course.
Unit VI: Ionic Equilibria: 1. Define the following terms, using examples where appropriate: a. Ionization constant i. Common ion effectb. pH and pOH j. Complex ionsc. Weak electrolyte k. Instability constantd. Dissociation l. Formation constante. Polyprotic acid m. Hydrolysisf. Indicator n. Solubility product constantg. Buffer o. Equivalence pointh. Hydrolysis constant p. Endpoint 2. Describe the ionization of water and its ionization constant. 3. Calculate the hydrogen ion concentration and the hydroxide ion concentration of pure water. 4. Given one of the following: [H+], [OH-], pH, or pOH of a solution, calculate the other three. 5. Given the concentration of a strong acid or strong base, calculate the pH and pOH of the solution.6. Calculate the pH of weak acids or bases, given their equilibrium concentrations, and vice versa. 7. Given the ionization constant and the initial (total) concentration of a monoprotic weak acid, or a monohydroxy weak base, calculate the hydrogen ion concentration and the hydroxide ion concentration and the concentration of all other species of the solution at equilibrium (you should make the appropriate assumptions). 8. From the information given in (7), calculate the percent ionization in a monoprotic weak acid solution or in a monohydroxy weak base solution. 9. Calculate the equilibrium concentrations of all species present when a polyprotic weak acid dissociates. 10. Given the initial (or total) concentration of a monoprotic weak acid or a monohydroxy weak base and given the pH or the solution at equilibrium, calculate the ionization constant of the weak acid or the weak base, and vice versa. 11. Explain the nature, preparation and use of buffer solutions. 12. Given the initial concentrations of a weak acid and its salt or that of a weak base and its salt (buffers) and the ionization constant(s), deduce the equili-brium conditions with proper assumptions and calculate the resulting pH of the mixture at equilibrium, and after dilution or small additions of acids or bases. 13. Use the common ion effect to calculate equilibrium concentrations or the equilibrium constant for weak acids or weak bases, given initial concentrations of all species. 14. Illustrate the three kinds of hydrolysis: a. Salt of a strong base and a weak acidb. Salt of a weak base and a strong acidc. Salt of a weak base and a weak acid and apply the ideas of hydrolysis to calculate the concentration of all ions and the pH at equilibrium, given initial conditions, or predict the acidic, basic, or neutral nature of salts in water. 15. Standardize a pH meter and correctly measure the pH of a solution with a pH meter. 16. Given the ionization constant of a weak acid or a weak base, and given the weak acid or the weak base and a salt (strong electrolyte) of the acid or of the base and necessary apparatus, prepare a required volume of a buffer solution of a desired pH.17. Given the initial concentration of the titrants and the ionization constants where applicable, predict the end point and the shape of a titration curve of pH against volume of acid or base added to a base or to an acid, respectively, for the following cases: a. Strong acid titrated with a strong base (or the reverse)b. Weak acid titrated with a strong basec. Strong acid titrated with a weak base. 18. Calculate the pH at any point of the addition in (17). 19. Perform any of the titrations in (17) in the laboratory, properly using burettes. 20. Explain the nature of the curve in the titrations of (17) in terms of the vertical rise and the two plateaus and why they are so. 21. Explain an alternate means of preparing the solution that exists at the end point in (17). 22. Select the right indicator according to the range of pH in which the end point of an acid-base titration lies. 23. Calculate the solubility (or concentration of the ions) given the solubility product constant, and vice versa. 24. Use the common ion effect to calculate equilibrium concentrations or the equilibrium constant for slightly soluble salts, given initial concentrations of all species. 25. Given the concentration of a solution of a cation and the concentration of a separate solution of an anion, of a slightly soluble salt, and given its Ksp, mathematically determine if a precipitate will form if given volumes of the two solutions are mixed. 26. Given the information in (25), for the case where a precipitate forms, calculate the number of moles (and grams) of the solid formed, the percent precipitation, and the final concentration of each of the ions remaining in solution. 27. Predict, mathematically, which ion will precipitate when a precipitating agent is added to a solution of two or more ions. 28. Determine the molar solubility of salts in solvents that form complex ions with the solute added. 29. Write instability constant and formation constant expressions from the chemical equation. 30. Relate the instability constant to the formation constant for complex ion formation.Unit VII: Chemistry of the Representative Elements I: The Metals: 1. Distinguish among metals, nonmetals, and metalloids (semi-metals) with respect to chemical properties, physical properties, and positions in the Periodic Table. 2. Write the outer shell electron configuration of any of the representative elements. 3. From the electron configuration of any element, determine which family or group it belongs to and vice versa. 4. Describe the reactions of the representative metals, their oxides, and their hydroxides with water, acids, or bases. 5. Describe the trends in metallic behavior , electronegativity, ionization energy, electron affinity, and atomic radii throughout the periodic table. 6. Deduce, using simple thermodynamics, what type of chemical reaction can be used to produce free metals from their compounds. 7. Illustrate some similarities in chemical behavior of the Group IA, IIA and IIIA elements, especially diagonal relationships, and the relative reactivities within each group. 8. Interpret diagonal relationships in terms of ionic potential. 9. Predict and explain the values of the stable oxidation states for the representative metals, and which will be more stable. 10. Interpret the trends in oxidation states exhibited by the atoms within a group in terms of the relative stabilities of high and low oxidation states. 11. Describe trends in any row or column of the periodic table with respect to: a. Atomic radius f. Metallic propertiesb. Ionic radius properties g. Oxidizing/reducing propertiesc. Ionization potential h. Ionic potentiald. Electron affinity i. Polarization of ionse. Electronegativity j. Hydrolysis 12. Use ionic potential to compare the relative degree of ionic-covalent bonding and physical properties of compounds composed of the representative elements. 13. Discuss the Solvay Process14. Relate equivalents, equivalent weight, weight, normality, and volume in an acid base titration, and use it in calculations.
Unit VIII: Chemistry of the Representative Elements II: The Metalloids and Nonmetals:1. Define the following terms, including examples where appropriate: a. Allotropism e. Disproportionationb. Catenation f. Polymersc. Three center bonds g. Oxoaniond. Amorphous h. Hydride 2. Compare metalloids and nonmetals in terms of the oxidation states displayed and the processes employed in their production. 3. Contrast the methods of preparation of the metalloids with those for the production of the nonmetals. 4. Describe the molecular structure, bonding, geometry and name of the allotropic forms of the pure metalloids and nonmetals. 5. Predict the important oxidation states of the nonmetals and metalloids. 6. Determine the oxidation state of the nonmetals and metalloids in ions and in neutral compounds. 7. Illustrate examples of catenation among nonmetals and metalloids by drawing structural formulas, and which element does it best. 8. Describe the two general methods for the preparation of nonmetals and metalloid hydrides. 9. Relate the ease of preparation and stability of nonmetal and metalloid hydrides to their standard enthalpies and free energies of formation. 10. Compare the relative acidic strength of the hydrides for the elements in both vertical columns and horizontal rows. 11. Write equations for the hydrolysis of nonmetal anions such as sulfide, nitride, phosphide and carbide. 12. Draw the structure of diborane and describe the bonding in this substance and why BH3 is not the simplest stable boron hydride. 13. Write the equation for the reaction of nonmetal oxides with water. 14. Describe three methods of preparation of nonmetal oxides, writing chemical equations for each. 15. Give the formulas of the important nonmetal oxides. 16. Give the structures, hybridization, and resonance forms of NO, NO2, CO, CO2, SO2, SO3 and show the valence bonds ( d and p ) that form. 17. Write equations for the preparation of the compounds in (16). 18. Compare the structures of nonmetal oxides on the bases of bonding preferences exhibited by the non-metals. 19. Compare the structures of P4, P4O6 and P4O10, and give reaction for the preparation of the oxides from phosphorous. 20. Discuss the molecular structure of quartz. 21. Give examples of 3 methods of preparing oxoacids and their anions. 22. Given the structure or the name or the formula for the following oxoacids (and the oxoanions), given one of them: a. HClO f. H2SO4 k. H3PO3 b. HClO2 g. H2S2O3 l. H3PO4 c. HClO3 h. HNO2 m. H2CO3 d. HClO4 i. HNO3 n. H3BO3 e. H2SO3 j. H3PO2 o. H2C2O4 and extend these structures to other members of the same families where appropriate. 23. List the oxidation state for the central element in the oxoacids and oxoanions in (22). 24. Give the bonding, geometries, resonance forms, and hybridization (where appropriate) for the oxoacids and oxoanions in (22). 25. Give the names and structures for the salts which form from the oxoacids in (22). 26. Give the equation for the formation of the oxoacids from the anhydrides, where appropriate (from CO2, N2O3, N2O5, P4O6, P4O10, SO2, SO3). 27. Compare the acidic strengths and oxidizing abilities of the oxoacids of the nonmetals. 28. Predict formulas for the halogen compounds of the nonmetals. Give geometries for these compounds based on the electron repulsion theory and list hybridizations where appropriate. 29. Use electronic structure and relative size of the atoms to determine possibility for existing and relative stability of the nonmetal halogen compounds. 30. Compare and explain the relative reactivities among the halogens and among the noble gases. 31. Explain the valence bond formation in the N2 molecule. 32. Explain why nitrogen is relatively unreactive, when compared to other nonmetals. 33. Define and explain nitrogen fixation and the nitrogen cycle in nature. 34. Discuss the preparation of and bonding in noble gas compounds. Include geometries and hybridization. 35. Indicate the composition of the two most abundant components of the atmosphere. 36. Indicate the six major pollutants of the air. 37. Indicate five sources for these pollutants. Unit IX: The Transition Elements: 1. Describe similarities and differences between A and B groups of the periodic table. 2. Distinguish between representative, transition, and inner transition elements. 3. Explain why Group IIB is sometimes considered a representative group. 4. Compare the properties among the transition elements horizontally as well as vertically including atomic radii, ionic radii, important oxidation states, ionization energy, hardness, melting points, and density. 5. List at least five characteristics that the transition elements have in common with each other.6. Write the electronic configurations of the first row transition elements, noting the anomalies and the reason for them. 7. Define "lanthanide contraction" and predict its effect on the properties of the transition elements in period 6. 8. Predict the possible oxidation states of the transition metals and give the more important oxidation states of the first row transition metals. 9. Indicate the relative importance of higher and lower oxidation states as one moves horizontally or vertically among the transition metals. 10. Use the relative stabilities of oxidation states to determine which of 2 compounds will be more easily oxidized (or reduced) or which will be the better oxidizing agent (or reducing agent). 11. Give formulas and names to the more important oxides and hydroxides (and their anions) of the first row transition metals and compare their relative oxidizing abilities. 12. Compare the relative acidity of the oxides and of the hydroxides of each transition element that has more than one important oxide of hydroxide. 13. Explain the use of silver in the black and white photographic process. 14. Explain the physiological action of mercury. 15. Discuss the coinage metals and why they are called that. 16. Discuss the two oxidation states of mercury and the unique structure and bonding they produce. 17. List the iron, palladium, and platinum triads, and why Group VIII is structured that way. 18. List the platinum metals and some of their important properties. 19. Define or describe the following terms relating to metallurgy, giving examples where appropriate: a. Ore i. Blast furnaceb. Amalgam j. Cast ironc. Flotation process k. Pig irond. Gangue l. Steele. Slag m. Calcinationf. Flux n. Bessemer converterg. Roasting o. Open hearth furnaceh. Smelting p. Mond process20. Identify and describe the three main steps involved in extracting a metal from its ore, and give examples of each. 21. Define an alloy. 22. Name at least two important alloys and describe their composition and applications. 23. List at least two properties in which an alloy differs from its components. 24. Define and compare the terms paramagnetism, ferromagnetism and diamagnetism, and give examples of elements exhibiting each type. 25. Define the term domain and relate it to the degree of magnetism which a substance exhibits. 26. Define the following terms relating to coordination chemistry, giving examples where appropriate: a. Complex compound l. Enantiomersb. Coordinate covalent bond m. Racemicc. Ligand n. Inner orbital complexd. Coordination sphere o. Outer orbital complexe. Chelating group p. High spin complexf. Monodentate ligand q. Low spin complexg. Polydentate ligand r. Degenerateh. Coordination number s. Crystal Field Theoryi. Stereoisomerism t. Crystal field splittingj. Geometrical isomerism u. Valence Bond Theoryk. Optical isomerism v. Donor atom 27. Given the formula or structure of a transition metal complex, identify the ligands, chelating groups, coordination sphere, coordination number, and donor atom. 28. Name transition metal complexes (using the rules of nomenclature) given the formula, and vice versa. 29. Identify and draw isomers of some transition metal complexes, identifying cis, trans, and optical isomers, or given the structure, identify which isomer is present. 30. State what nonsuperimposable mirror images means and how this relates to coordination compounds and optically active coordination compounds. 31. Define polarized light and explain what happens to it when it is passed through a solution of each of a pair (or a mixture) of optical isomers.32. Use the valence bond theory to explain the nature of the bonding in coordination complexes. 33. For transition metal complexes, use the valence bond theory to explain: a. The nature of the coordinate covalent bond.b. Their electron configuration (before and after complexing).c. Their structure, geometry, and hybridization.d. Whether an inner or outer orbital complex will form (be more stable).e. The number of unpaired electrons that results.f. Their magnetic properties.g. Their color.h. The faults in the theory. 34. For transition metal complexes, use the crystal field theory (ligand field theory) to explain: a. The nature of the bond formed and compare it to the coordinate
2. Demonstrate the ability to correctly and effectively manipulate chemicals and glassware by working alone.
4. Demonstrate the ability to correctly and effectively use laboratory balances.3. Demonstrate the ability to correctly and effectively collect and analyze data from an experiment by working alone.
5. Demonstrate the ability to correctly and accurately do quantitative analysis such as titrations, pipetting and preparation of solutions by working alone.
6. Demonstrate the ability to correctly and effectively collect and treat data on the computer.
7. Demonstrate the ability to correctly and effectively use instruments like Spectrophotometers, voltmeters and pH meters.
8. Utilize critical thinking and quantitative reasoning skills in observing, organizing and analyzing data, synthesizing information, interpreting results, and communicating the results of the analyses and laboratory investigations.
9. Perform chemical experimentation in a safe and scientific manner, using proper scientific and laboratory safety procedures.
10. Students must show work, thought process and/or justification for answers when
necessary on laboratory reports. They should also be clear and legible.