EVANGEL
UNIVERSITY
G E N E R A L C H E M I S T R Y CHEM 111
Dr. Steve Badger
Syllabus
Fall 2007
COURSE DESCRIPTION
Fundamental principles including the laws of chemical combinations, gas laws, simpler structure of atoms,
periodic system, states of matter, chemistry of the non-metals and their important compounds, metallic
elements and their compounds, and an introduction to the chemistry of carbon compounds. Laboratory work
includes the systematic separation and identification of common anions and cations. Lab required.
COURSE OBJECTIVES
Students successfully completing General Chemistry should be able to:
1. Describe the scientific method and how chemistry
fits in the natural sciences.
2. Classify matter as element, compound, or mixture.
3. Define physical and chemical properties and
physical and chemical changes.
4. Use SI units of measurement.
5. Describe current atomic theories and some of the
experimental evidence for these theories.
6. Describe covalent and ionic bonds.
7. Use the periodic table of the elements.
8. Name a compound by looking at its formula.
9. Write the formula of a compound from its name.
10. Balance chemical equations.
11. Calculate mass relationships in chemical reactions.
12. Calculate the limiting reagent and the percent yield
of a reaction.
13. Describe the typical properties of aqueous solutions.
14. Calculate amounts of chemicals needed to make
specific molar and normal solutions.
15. Describe several reactions that occur in aqueous
solutions (acid-base, redox, etc.).
16. State and use Boyle’s Law, Charles’ Law, GayLussac Law, Avogadro’s Law, the Ideal Gas
Equation, and Dalton’s Law of Partial Pressures.
17. Describe the Kinetic Molecular Theory of Gases.
18. Outline Quantum Theory and apply it to the
electronic structure of atoms.
19. Describe the Bohr model of the atom.
20. Describe the dual nature of the electron.
21. Use the Pauli exclusion principle, Hund’s rule, and
the Aufbau principle.
22. Identify which electronic subshell is being filled on
the basis of location in the periodic table.
23. Understand and use the periodic law and the
periodic table of the elements.
24. Describe the physical and chemical periodic trends.
25. Draw Lewis dot structures using the Octet Rule and
its exceptions.
26. Describe several hybridized molecular orbitals.
27. Describe the intermolecular forces of liquids and
solids and their effects on their properties.
28. Use common units of concentration for solutions.
29. Describe the effects of temperature and pressure on
solubility.
30. Describe the colligative properties of solutions of
electrolytes and nonelectrolytes.
31. Use safe laboratory techniques to make some of the
observations supporting chemical laws and theories.
More specific objectives are listed in the course outline under each chapter title.
COURSE GOALS
Students who successfully complete the course should:
1. Be prepared for CHEM 112.
2. Understand the limitations of the scientific method.
3. Appreciate more fully the creativity of the Creator.
4. Have a clearer focus on their career goals.
REQUIRED TEXTS
Chang, Raymond. General Chemistry: The Essential Concepts, 5th edition. McGraw-Hill, 2008. (Older
editions and other texts are not acceptable. Each student must have his or her own textbook.)
Cruickshank & Chang. Problem Solving Workbook to Accompany General Chemistry: The Essential Concepts,
5th Edition, McGraw-Hill, 2008.
Slowinski, Wolsey, Masterton. Chemical Principles in the Laboratory. 8th Edition, Thomson Wadsworth
Publishing, 2005. (Used lab manuals are not acceptable! You must have a new lab book.)
DR. STEVE BADGER
GENERAL CHEMISTRY CHEM 111-2
RECOMMENDED SUPPLEMENTARY MATERIALS
Students are expected to use all of McGraw-Hill’s tutorial resources (“OnLine Learning Center,”
“ChemSkillBuilder”, and other materials) on the Internet at www.mhhe.com/physsci/chemistry/chang/.
Read the first several pages of Chang’s text to discover other helpful study aids.
SCRIPTURAL FOUNDATION
The wrath of God is being revealed from heaven against all the
godlessness and wickedness of men who suppress the truth by
their wickedness, since what may be known about God is plain
to them, because God has made it plain to them. For since the
creation of the world God's invisible qualities—His eternal
power and divine nature—have been clearly seen, being understood from what has been made, so that men are without excuse.
Rom 1:18-20, NIV.
He is the image of the invisible God, the
firstborn over all creation. For by him all things
were created: things in heaven and on earth,
visible and invisible, whether thrones or powers
or rulers or authorities; all things were created
by him and for him. He is before all things, and
in him all things hold together.
Col 1:15-17, NIV
COURSE REQUIREMENTS
1.
2.
Read this syllabus and keep it handy for easy reference. You’re expected to know what it says.
You must have your own textbook and bring it to class with you. Also, always bring paper and pencil to class to
take notes.
3. Use Cruickshank & Chang’s Problem Solving Workbook to Accompany General Chemistry: The Essential
Concepts, 4th Edition, McGraw-Hill, 2006.
4. Print out the PowerPoint® notes from my website (http://www.evangel.edu/Personal/badgers/Web/) and bring
them to class with you each day. (I will give you a sheet of instructions, or you can view them on my web page.)
5. Arrive punctually to the classroom, pay attention to the prof’s presentations, ask questions as necessary, and
interact with the material.
6. Read the assigned materials before the professor covers them in class.
7. Answer the questions assigned at the end of each chapter before the exam over that material.
8. Be sure you understand the “Key Equations” and “Key Words” at the end of each chapter.
9. Take the exams on the scheduled date and at the scheduled time.
10. Always bring to class a scientific calculator that you know how to use. When taking an exam, you will not be
allowed to use a calculator that can store text (e.g., TI-92). The HP30S calculator is recommended.
11. While in the classroom, turn off cell phones, pagers, and other devices that may interrupt class.
COURSE PROCEDURES
1.
Material will be presented in a lecture/discussion format. Students are encouraged to courteously interrupt the prof
with pertinent questions.
2. Powerpoint® presentations, the whiteboard, overhead transparencies, models, and handouts of labeled drawings
will be used to help illustrate concepts.
3. The professor is in the habit of asking students questions designed to provoke them to interact with the text,
scientific concepts, the author of the text, and the prof’s opinions.
4. Review sheets may be provided to help students focus on concepts they are likely to see on exams.
5. E-Mail Communication Systems: E-mail is to be the principal means of communication between faculty, staff,
administration and students. Types of communication may include assignments, registration materials,
announcements, etc. It is the responsibility of the student to check his/her Evangel University E-mail account
daily, and the student will be held accountable for any and all official communication of administrative policies,
faculty instructions and campus information sent via the Evangel University E-mail system.
ATTENDANCE POLICY
1. Regular, punctual class attendance is essential for passing chemistry.
2. Arriving more than 10 minutes late constitutes an absence. Three tardies constitute an absence.
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DR. STEVE BADGER
GENERAL CHEMISTRY CHEM 111-2
3. Leaving early without the prof’s permission constitutes an absence.
4. Being absent does not relieve you of the responsibility of any due assignments.
5. Absences are counted from the first day of classes. Absences due to late registration count as regular
absences. (If you were not present, you were absent.)
6. All absences from class (including school-sponsored activities, ministry trips, etc.) are recorded and
counted as absences. They are excused, but they still count as absences.
7. Review the detailed information at www.evangel.edu/CurrentStudents/StudentHandbook concerning
class attendance. “The all-school policy requires that students attend a minimum of 75% of the class
sessions in order to receive a passing grade.”
( See http://www.evangel.edu/CurrentStudents/StudentHandbook/Part2.asp#ClassAttendance)
8. You cannot receive credit for this class if you have more than 13 absences.
9. Nov. 2 is the last day to withdraw with a “W.” Dec. 7 is the last day to drop with either a WP or a
WF.
EXAM POLICY
The policies below have been put into effect because of the large number of frivolous excuses given for missing exams
in recent years. Please do not ask me to make an exception for you.
1. Students are expected to be present and on time for all exams.
2. No make-up exam will be given unless you have written documentation of an excusable absence signed by an
appropriate person (e.g., physician, nurse, professor, judge, coach, etc.).
3. Excusable absences are limited to “school assignment,” an illness or hospitalization, or a dire family emergency
(like a death in the immediate family).
4. All other reasons for missing a scheduled exam will not be excusable.
5. Students who are absent without an acceptable excuse will NOT be permitted to make up the missed exam and
will receive a grade of zero for that exam.
6. Here is a partial list of reasons for missing an exam that are not excusable absences:
o Overslept, alarm clock failed to work
o Had work or social conflicts
o Studied all night for the exam
o Took a friend to the hospital
o Unprepared for the exam
o Had a sick roommate
o Had another exam the same day
o Lost or stolen textbook
7. Students who miss a scheduled exam that is given on the last day of class before a holiday or on the first day of
class after a holiday will not be allowed to take that exam—unless they have written, signed documentation of a
“school assignment,” an illness or hospitalization, or a dire family emergency (like a death in the immediate
family).
8. Early exams will not be administered. If you miss a test, discuss it with the professor as soon as possible. If he
approves a make-up test, you are responsible for scheduling it as early as possible.
9. Make-up exams may not be identical to the regular exam, but they may be different (for example, it may be all
essay questions).
ASSIGNMENT POLICY
1.
2.
3.
4.
5.
6.
7.
Read the assigned chapter before the beginning of the class period in which it will be discussed.
Learn the “Key Equations” at the end of each chapter.
Read the “Summary of Facts and Concepts” at the end of each chapter.
Understand all the words in “Key Words” at the end of each chapter.
Answer the assigned “Questions and Problems” at the end of each chapter. Write both the question and the
answer in a notebook that you bring to class each day. Write nothing else in this notebook.
Homework assignments should be completed before the prof finishes covering that chapter.
If a student has a hardship, he/she should discuss it with the professor, in advance, if possible.
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DR. STEVE BADGER
GENERAL CHEMISTRY CHEM 111-2
EVALUATION PROCEDURES
1.
Anyone caught cheating on an examination will receive a zero on that exam with no opportunity to take it again.
This includes looking at notes or the text or talking with someone.
2. If you look at someone else’s exam during a test, you will receive a zero on that exam with no opportunity to take
it again.
3. Exam questions are typically multiple choice, true-false, matching, and a few questions that require writing longer
answers. Bring a No. 2 pencil with you to class when exams are administered.
4. You will not be allowed to use a graphing calculator (that can store text) on exams. The HPS30S calculator is
recommended.
5. All grading will be done on a percentage basis.
6. The simple average of the exam grades and your lab grade will be the final grade.
7. The professor may use the completion of assignments, participation in class discussions, and your demeanor in
class to “weight” a student’s grade in the course.
8. The professor will occasionally give unannounced tests (“pop quizzes”) over current homework material. Grades
from pop quizzes will be averaged and this average used as a single test grade and averaged in with the others.
9. Grading scale:
92-100
A
82-87
B
72-77
C
62-67
D
90-91
A80-81
B70-71
C60-61
D88-89
B+
78-79
C+
68-69
D+
59 and less F
10. Class grades will not be placed on a “curve”; however, grades for some exams may be skewed.
11. Evangel University is committed to the provision of reasonable accommodations for students with disabilities, as
defined in Section 504 of the Rehabilitation Act of 1973. If you think you qualify for accommodations, notify me
or talk with people in the Academic and Career Development office (AB-1, room 218) as soon as possible.
COURSE OUTLINE
Chapter 1 Introduction
Upon completion of this chapter, the student should be able to:
1. Offer examples of how chemistry is used in everyday life.
2. Set forth what the scientific method is and how it is used.
3. Classify materials in terms of homogeneous and heterogeneous mixtures.
4. Distinguish between compounds and elements.
5. Compare physical versus chemical properties.
6. Recall from memory commonly used prefixes used with SI units.
7. Solve problems involving density, volume, and mass.
8. Convert between the Kelvin, Celsius, and Fahrenheit temperature scales.
9. Apply scientific notation and correct number of significant figures in problem solving.
10. Discuss the difference between accuracy and precision.
11. Utilize factor-label method of problem solving.
12. Recall from memory common conversion factors for the metric system to English system. Examples of such
conversion factors include grams to pounds, centimeters to inches and liters to gallons.
Chapter 2 Atoms, Molecules, and Ions
Upon completion of this chapter, the student should be able to:
1. Restate the points of Dalton’s atomic theory.
2. Distinguish between the law of definite proportions and the law of multiple proportions.
3. Argue how Dalton’s third hypothesis is another way of stating the law of conservation of mass.
4. Explain how electrons were discovered and how Millikan’s oil drop experiment determined the charge of the electron.
5. Predict the path of alpha particles, beta particles, and gamma rays as they pass between two oppositely charged
electrical plates.
6. Set forth how Rutherford’s experiment concluded that atoms are mostly empty space with very small central cores,
which are known as nuclei.
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DR. STEVE BADGER
GENERAL CHEMISTRY CHEM 111-2
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Predict the path of protons, electrons, and neutrons as they pass between oppositely charged electrical plates.
Compute the number of electrons, protons, and neutrons in atoms and ions.
Give examples of isotopes.
Recall from memory the names for the three isotopes of hydrogen.
Predict if an element is a metal, nonmetal, or metalloid.
Classify elements as alkali metals, alkaline earth metals, or noble gases.
List several examples of diatomic molecules.
Classify ions in terms of monatomic ions, polyatomic ions, cations, and anions.
Distinguish between molecular and empirical formulas.
Predict correct formulas for ionic compounds.
Name common ionic compounds, molecular compounds, binary acids, oxoacids, bases, and hydrates given their
respective chemical formulas.
18. Predict the chemical formulas of common ionic compounds, molecular compounds, binary acids, oxoacids, bases, and
hydrates given their respective names.
Chapter 3 Stoichiometry
Upon completion of this chapter, the student should be able to:
1. Convert between grams and atomic mass units (AMUs).
2. Calculate average atomic mass given the mass and natural abundance of each isotope.
3. Use Avogadro’s number to determine the number of units present in a given number of moles.
4. Convert between mass, number of moles, and number of atoms (molecules) of an element (compound).
5. Establish the molecular mass and molar mass given the molecular formula.
6. Compute the percent composition (mass percent) given the chemical formula for an ionic or molecular compound.
7. Describe the experimental procedure used to determine empirical formulas.
8. Establish the molecular formula given the mass of each element present (or mass percent of each element) and the
compound’s molar mass.
9. Balance chemical equations.
10. Interpret the meaning of chemical equations in terms of molecules, moles, and masses.
11. Distinguish between products and reactants in a chemical equation.
12. Predict the products formed by combustion reactions.
13. Use stoichiometric methods to predict the mass (number of moles) of the products formed given the mass of each
reactant (number of moles of each reactant).
14. Use stoichiometric methods to deduce the limiting reagent, excess reagent, the amount of expected products produced,
and the amount of excess reagent left over upon completion of the reaction given the mass (number of moles) of each
reactant in the chemical equation.
15. Use stoichiometric methods to predict the theoretical yield and percent yield given the mass (number of moles) of each
reactant and the actual yield of a reaction.
16. Calculate the mass (number of moles) of each reactant required given the percent yield and the mass (number of
moles) of products desired.
___________________________________________ 0 SEP 14, EXAM 1
Chapter 4 Reactions in Aqueous Solution
Upon completion of this chapter, the student should be able to:
1. Distinguish between solute, solvent, and solution.
2. Classify common compounds as strong electrolytes, weak electrolytes, or nonelectrolytes (strong or weak acids or
strong or weak bases).
3. Suggest why water is often called a universal solvent utilizing the terms polar solvent, dissociation, ionization, and
hydration.
4. Describe precipitation reactions using the terms solubility and precipitate.
5. Classify common ionic compounds as soluble or insoluble.
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DR. STEVE BADGER
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GENERAL CHEMISTRY CHEM 111-2
Given the reactants of a chemical reaction, predict the resulting products and write the molecular equation, ionic
equation, and net ionic equation and identify spectator ions.
Distinguish between Arrhenius acids and bases and Brφnsted acids and bases.
Compare and contrast the properties of acids and bases.
List common examples of monoprotic, diprotic and triprotic acids.
Justify how some ions can act as an acid or as a base (amphoteric).
Explain, by using a chemical equation, how ammonia (NH3) is classified as a Brφnsted base.
Predict the products formed by acid-base neutralization reactions.
Discuss what factor results in an oxidation-reduction reaction.
Identify oxidation half-reactions, reduction half-reactions, oxidizing agents and reducing agents.
Assign oxidation numbers to elements in compounds and ions.
Categorize redox reactions in terms of combination reactions, decomposition reactions, displacement reactions, and
disproportionation reactions.
Predict the results of a chemical reaction involving metals given the activity series (electrochemical series).
Predict the results of a chemical reaction involving halogens given the halogen activity series.
Compute the molarity of a solution given the mass (number of moles) of solute and the volume of solution.
Describe the method for preparing a specific molar solution given the volume of solution required and the solute to be
used.
Relate in detail how to prepare a specific dilute solution given a known stock solution using dilution techniques.
Predict the mass of a precipitate formed using gravimetric analysis methods.
Deduce the mass percent of specific ions present in an original solution given the results of a gravimetric analysis.
Use the terms titration, standard solution, equivalence point, and indicator to describe quantitative studies of acid-base
neutralization reactions.
Determine the concentration of an unknown acid (base) given the results of an acid-base titration.
Predict the amount (mass, moles, or volume of solution) of an acid (base) required to neutralize a base (acid).
Predict the volume of an oxidizing (or reducing) agent solution required to oxidize (or reduce) a specific volume of
reducing (or oxidizing) agent solution provided that the net ionic equation is given.
Chapter 5 Gases
Upon completion of this chapter, the student should be able to:
2. Recall from memory at least ten common substances that are gases at one atmosphere and 25o C.
3. Distinguish between the terms gas and vapor.
4. List four physical characteristics of all gases.
5. Define the terms velocity, acceleration, force, newton, energy, joule, kinetic energy, pressure and pascal.
6. Describe how a simple barometer is constructed and how it functions.
7. Convert among these pressure units: torr, mmHg, atmospheres, and pascals.
8. State the difference between open-tube manometers and closed-tube manometers and indicate how each is used.
9. Write, explain, and apply each of the following:
ƒ Avogadro’s law (V ∝ n).
ƒ Boyle’s law (P ∝ 1/ V and P1V1 = P2V2).
ƒ Ideal gas law (PV = nRT).
ƒ Charles’ law (P ∝ T and V1/ T1 = V2/ T2).
10. Describe the Kelvin temperature scale.
11. Recall from memory the value and units of the gas constant, R.
12. State what standard temperature and pressure are and demonstrate that at STP one mole of gas occupies 22.4 liters.
13. Perform calculations involving density, the Ideal gas equation and molar mass.
14. Use the Ideal gas equation to determine the moles of a gas and use the number of moles in stoichiometric-based
problems.
15. State Dalton’s law of partial pressures and utilize it in problems involving mixtures of gases including the collection of
gases over water.
16. Define mole fraction and verify that Pa = XaPtotal.
17. Discuss the four assumptions upon which the kinetic molecular theory of gases is based.
18. Suggest how the kinetic molecular theory of gases qualitatively explains the following:
ƒ The compressibility of gases.
ƒ Avogadro’s law.
ƒ Boyle’s law.
ƒ Dalton’s law of partial pressure.
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DR. STEVE BADGER
GENERAL CHEMISTRY CHEM 111-2
ƒ Charles’ law.
22. Perform calculations using root-mean-square speeds.
23. Describe the process of gaseous diffusion.
24. Argue how a real gas, behaving non-ideally, differs from an ideal gas as described by the four assumptions in the
kinetic molecular theory of gases.
25. Conclude under what conditions a real gas will approximate an ideal gas.
26. Apply van der Waal’s equation to real gases.
Chapter 6 Energy Relationships in Chemical Reactions
Upon completion of this chapter, the student should be able to:
1. Classify common processes as endothermic or exothermic.
2. Use thermochemical equations and stoichiometry to determine amount of heat lost or gained in a chemical reaction.
3. Perform calculations involving specific heat, mass and temperature change.
4. Sketch the main components of a constant-volume bomb calorimeter.
5. Determine heats of reactions given experimental data collected in a calorimetry experiment.
6. Calculate standard enthalpy of reactions given the standard enthalpy of formations for products and reactants.
7. Apply Hess’s law to a multi-step process to determine standard enthalpy of reaction.
8. Describe heat of solution, lattice energy, heat of hydration, heat of dilution, system, surrounding, and internal energy.
9. Classify properties of materials as state functions or non-state functions.
10. Restate the First Law of Thermodynamics.
11. Recall the sign conventions for work and heat used in the textbook.
12. Apply heat and work relationships to gas-phase problems.
13. Define H in term of E, P, and V.
14. Calculate change in internal energy (∆E) given thermochemical equations.
15. Define and explain the following terms:
• Exothermic
• Thermochemistry
• Energy
• Enthalpy (∆H)
• Open system
• Radiant energy
• Calorimetry
• Closed system
• Thermal energy
• Heat capacity
• Isolated system
• Chemical energy
• Specific heat
• Endothermic
• Potential energy
_____________________________________ 000 OCT 12, EXAM 2
Chapter 7 The Electronic Structure of Atoms
Upon completion of this chapter, the student should be able to:
1. Explain how Planck’s theory challenged classical physics.
2. Define wavelength, frequency, and amplitude of waves.
3. Utilize the relationship between speed, wavelength, and frequency (hertz).
4. Describe Maxwell’s theory of electromagnetic radiation.
5. Recall from memory the speed of light (3.00 x 108 m/s).
6. Apply the metric prefix nano in calculations involving wavelength of light.
7. Classify various regions of the electromagnetic spectrum in terms of energy, frequency and wavelength.
8. Use Planck’s equation to determine energy, frequency, or wavelength of electromagnetic radiation.
9. Describe the photoelectric effect as explained by Einstein using such terms as threshold frequency, photons, kinetic
energy, binding energy, light intensity and number of electrons emitted.
10. Show how Bohr’s model of the atom explains emission, absorption and line spectra for the hydrogen atom.
11. Compare Bohr’s model of the atom and that of the sun and surrounding planets.
12. Predict the wavelength (frequency) of EMR emitted (absorbed) for electronic transitions in a hydrogen atom.
13. Use the terms ground state and excited state to describe electronic transitions.
14. Describe De Broglie’s relationship involving the wavelength of particles.
15. Explain the major components of a laser and list three properties that are characteristic of a laser.
16. Describe Heisenberg’s uncertainty principle.
17. Contrast orbits (shells) in Bohr’s theory with orbitals in quantum theory.
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DR. STEVE BADGER
GENERAL CHEMISTRY CHEM 111-2
18. Discuss the concept of electron density.
19. Recall from memory the four quantum numbers (n, ℓ, mℓ, ms) and their relationships.
20. Relate the values of the angular momentum quantum number, ℓ, to common names for each orbital (s, p, d, f) and
describe their shapes.
21. Account for the number of orbitals and number of electrons associated with each value of ℓ, the angular momentum
quantum number.
22. Categorize orbital energy levels in many-electron atoms in order of increasing energy.
23. Write the four quantum numbers for all electrons in multi-electron atoms.
24. Use the Pauli exclusion principle and Hund’s rule to predict the electron configuration for multielectron atoms
25. Deduce orbital diagrams from diamagnetic and paramagnetic data.
26. Derive the ground state electron configuration of multi-electron atoms using the Aufbau principle.
27. List exceptions to the expected electron configuration for common metals (e.g., Cr, Mo, Cu, Ag, Au).
Chapter 8 The Periodic Table
Upon completion of this chapter, the student should be able to:
1. Explain the periodic table as described by Mendeleev & Meyer and indicate their shortcomings.
2. Explain the basis of the periodic table as described by Moseley and how it predicted properties of “missing” elements.
3. Identify elements that correspond to each of the following groups:
ƒ representative elements
ƒ lanthanides
ƒ transition metals
ƒ noble gases
ƒ actinides
4. Describe the electron configuration of cations and anions and identity ions and atoms that are isoelectronic.
5. Apply the concept of effective nuclear charge and shielding constants (screening constants) to justify why the first
ionization energy is always smaller than the second ionization energy of a given atom.
6. Predict the trends from left to right and top to bottom of the periodic table for each of the following:
ƒ atomic radius
ƒ electron affinity
ƒ ionic radius
ƒ metallic character
ƒ ionization energy
7. Relate why hydrogen could be placed in a class by itself when reviewing its chemical properties.
8. Provide examples of Group 1A elements reacting with oxygen to form oxides, peroxides, and superoxides.
9. Predict the reaction of alkali metals with water.
10. Describe the reactivity of alkaline earth metals with water.
11. Relate how strontium-90 could lead to human illness.
12. Compare the reactivity of boron, a metalloid, to aluminum.
13. Identify the metals, nonmetal, and metalloids of Group 4A.
14. Recall the reactions that form nitric acid, phosphoric acid and sulfuric acid.
15. List the halides (halogens).
16. Indicate the three hydrohalic acids that are strong acids and the one hydrohalic acid that is a weak acid.
17. Explain why the name for Group 8A has changed from inert gases to noble gases.
18. List the three “coinage” metals and explain their relative inertness.
19. Rationalize the characteristics of the properties of oxides of the third period elements.
20. Classify oxides as acidic, basic, or amphoteric.
21. Explain why concentrated bases such as NaOH should not be stored in glass containers.
______________________________________ 000 NOV 2, EXAM 3
Chapter 9 Chemical Bonding I: The Covalent Bond
Upon completion of this chapter, the student should be able to:
1. Identify the valence electrons for all representative elements.
2. Rationalize why alkali metals and alkaline earth metals usually form cations and oxygen and the halogens usually form
anions using Lewis dot symbols in the discussion.
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DR. STEVE BADGER
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GENERAL CHEMISTRY CHEM 111-2
Use Lewis dot symbols to show the formation of both ionic and molecular compounds.
Define lattice energy, Coulomb’s law, and the Born-Haber cycle.
Demonstrate how the Born-Haber cycle is an application of Hess’s law and use the Born-Haber cycle to determine
lattice energy for an ionic solid.
Identify covalent compounds, the type of covalent bonds present, and the number of lone pairs of electrons using
Lewis structures.
Relate types of bonds to bond length and bond strength.
Compare and contrast various properties expected for ionic compounds versus covalent compounds.
Identify ionic, polar covalent and (nonpolar) covalent bonds using the concepts of electronegativity.
Predict the relative changes in electronegativity with respect to position on the periodic table.
Use the concept of electronegativity to rationalize oxidation numbers.
Use Lewis dot and the octet rule to write Lewis structures of compounds and ions.
Apply the concept of formal charge to predict the most likely Lewis structure of a compound.
Explain how Lewis structures are inadequate to explain observed bond length (bond types) in some compounds and
how the concept of resonance must be invoked.
Recall several common examples in which the octet rule fails.
Demonstrate, using Lewis structures, the formation of a coordinate covalent (dative) bond.
Use Lewis structures and bond energies to predict heats of reaction.
Rationalize why ¥H for breaking chemical bonds is positive and the formation of chemical bonds is negative.
Chapter 10 Chemical Bonding II: Molecular Geometry and Hybridization of Atomic Orbitals
Upon completion of this chapter, the student should be able to:
1. Identify using the VSEPR model, what category (and thus the corresponding molecular geometry, angle(s) and sketch)
a molecular or ion belongs given its formula.
2. Rationalize the observed decrease in angles for AB4E0, AB3E1, and AB2E2 and for AB3E0 and AB2E1.
3. Apply the VSEPR model to compounds with more than one central atom.
4. Use the concepts of electronegativity, dipole moments, and VSEPR geometries to identify polar and nonpolar
molecules.
5. Use dipole moment concepts to predict properties of cis and trans isomers.
6. Relate how a microwave oven functions and how the type of bonds present effects the amount of energy absorbed.
7. Sketch and justify how potential energy changes versus the interatomic distance for a diatomic molecule.
8. Use Valence Bond theory, hybrid orbitals, and hybridization to explain the geometries predicted by VSEPR model.
9. Identify what type of hybrid orbitals are in common compounds and ions.
10. Apply the concepts of sigma and pi bonds and Valence Bond theory to explain properties of double and triple bonds
and the concept of resonance.
11. State what physical property of oxygen is not accounted for by Valence Bond theory but is in Molecular Orbital
theory.
12. Explain the difference between bonding and anti-bonding orbitals using the concepts of constructive and destructive
interference of waves.
____________________________________ 0000 NOV 16, EXAM 4
Chapter 12 Intermolecular Forces and Liquids and Solids
Upon completion of this chapter, the student should be able to:
1. Characterize the properties of gases, liquids and solids in terms of density, compressibility and motion of molecules.
2. Distinguish between intermolecular and intramolecular forces.
3. Identify and give examples of the following forces:
ƒ ion–ion
ƒ induced dipole–induced dipole
ƒ ion–dipole
ƒ van der Waals
ƒ ion–induced dipole
ƒ dispersion
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DR. STEVE BADGER
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GENERAL CHEMISTRY CHEM 111-2
ƒ dipole–dipole
ƒ hydrogen bonding
ƒ dipole–induced dipole
Relate polarizability and intermolecular forces.
Suggest why H2O, HF and NH3 do not follow the expected trend as shown by the plot in Figure 11.7
Use the concepts of intermolecular forces to explain surface tension, capillary action, cohesion, adhesion and viscosity.
Describe the structure of water and relate it to why ice is less dense then water.
Sketch density vs. temperature curve for water and relate how this plot has major significance for aquatic life.
Distinguish between crystalline and amorphous solids and give examples of each.
Identify coordination numbers for atoms in simple cubic, face-centered cubic and body-centered cubic structures.
Determine the number of atoms contained in a unit cell for typical crystal structures.
Use geometric concepts to relate a, the edge length of a unit cell, to r, the radius of atoms, in simple cubic, facecentered cubic and body-centered cubic structures.
Calculate what fraction of a unit cell is occupied by atoms, and show that the face-centered cubic structure has the
closest packing for simple cubic, face-centered cubic and body-centered cubic structures.
Perform calculations involving crystal structure, density, atomic radius and the number of atoms per unit cell.
Describe how x-ray diffraction is used to determine geometric parameters of solids.
Characterize ionic, covalent, molecular and metallic crystals including general properties and examples of each.
Use Clausius–Clapeyron equation to solve for vapor pressure at a specific temperature given associated data.
Sketch a typical heating curve and identify various aspects of it.
Use phase diagrams to identify what phase(s) is/are present given specific conditions.
Discuss the following:
ƒ evaporation
ƒ boiling point
ƒ deposition
ƒ condensation
ƒ melting point
ƒ molar heat of sublimation
ƒ molar heat of vaporization
ƒ supercooling
ƒ phase changes
ƒ molar hear of fusion
ƒ sublimation
ƒ critical temperature and pressure
Chapter 13 Physical Properties of Solutions
Upon completion of this chapter, the student should be able to:
1. Use the terms saturated, unsaturated and supersaturated to describe solutions.
2. Distinguish between crystallization and precipitation.
3. Predict relative solubilities given the next dipole moment of solute and solvent.
4. Define, determine and inter-convert between each of the following:
molarity
mole fraction
molality
percent by mass
5. Suggest a shortcoming of molarity and explain why molality is a preferred concentration unit under certain conditions.
6. Use the concept of fractional crystallization to show how dissolved solids can be separated.
7. Describe how thermal pollution may effect the oxygen content in lakes or streams.
8. State Henry’s law and use it to determine the solubility of gases in liquids.
9. Rationalize why two common materials (NH3 or CO2) when dissolved in water do not follow Henry’s law.
10. Define colligative properties and give four examples (vapor–pressure lowering, freezing–point lowering, boiling–point
elevation, and osmotic pressure).
11. Use Raoult’s law to find vapor pressures or concentrations of solutions.
12. Describe the apparatus used in fractional distillation.
13. Predict the plot of pressure versus mole fraction for an ideal solution.
14. Rationalize the possibility of either positive or negative deviations from Raoult’s law by non-ideal solutions.
15. Perform calculations involving boiling-point elevation, freezing-point depression, Kf, Kb and molality.
16. Use the concepts of osmotic pressure to describe the processes of osmosis and reverse osmosis.
17. Describe the following terms: semi-permeable membrane, isotonic, hypertonic, hypotonic, crenation
18. Use the concepts of colligative properties to determine molar mass.
19. Define the van’t Hoff factor and demonstrate how it is incorporated into the colligative property equations.
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DR. STEVE BADGER
GENERAL CHEMISTRY CHEM 111-2
20. Give examples of common types of colloids and describe the dispersing medium and dispersed phase for each.
21. Describe the Tyndall effect.
22. Identify hydrophilic and hydrophobic colloids and describe the cleansing action of soap.
____________________________________ 0000 DEC 12, EXAM 5
_____________________ 0000 DEC 14, 7:30 A.M., FINAL EXAM
HOMEWORK ASSIGNMENTS
Answer these representative items in the “Questions and Problems” at the end of each chapter. Working these
problems should help prepare you for the exams. Write both the question and the answer in a notebook that
you bring to class each day. Write nothing else in this notebook.
Answer these “Questions & Problems” at the end of each chapter
Chapter 1
Introduction
Chapter 2
Atoms, Molecules, & Ions
Chapter 3
Stoichiometry
Chapter 4
Reactions in Aqueous Solution
Chapter 5
Gases
Chapter 6
Energy Relationships in Chemical Reactions
1, 4, 5, 8, 10, 12, 13, 14, 15, 18, 22, 28, 32, 34, 40
1, 9, 10, 12, 14, 16, 18, 25, 26, 38, 40, 42, 44, 46
1, 2, 5, 6, 13, 16, 21, 22, 26, 48, 55, 57, 59, 60
1, 2, 3, 4, 5, 6, 8, 10, 14, 16, 18, 19, 20, 30, 32, 34, 36, 38,
40, 42, 54, 60, 65, 78
1, 2, 3, 4, 5, 14, 18, 20, 22, 27, 32, 36, 40, 58, 60, 65, 82
1, 2, 7, 11, 12, 16, 18, 19, 24, 27, 32, 37, 40, 44, 46, 48, 60
Chapter 7
The Electronic Structure of Atoms
1, 2, 3, 5, 8, 10, 12, 18, 26, 28, 32, 34, 44, 48, 54, 58, 60, 62,
69, 72, 78
Chapter 8
The Periodic Table
4, 5, 6, 7, 8, 16, 18, 20, 22, 23, 24, 28, 30, 32, 38, 44, 49, 52,
57, 60, 62, 63, 72, 84
Chapter 9
Chemical Bonding I: The Covalent Bond
Chapter 10 Chemical Bonding II:
Molecular Geometry & Hybridization of Atomic Orbitals
Chapter 12
Intermolecular Forces & Liquids & Solids
Chapter 13
Physical Properties of Solutions
1, 7, 11, 14, 16, 18, 22, 24, 25, 30, 34, 38, 42, 45, 48
1, 2, 5, 8, 10, 12, 13, 18, 23, 29, 32, 36, 38, 43
1, 3, 8, 10, 12, 14, 16, 22, 25, 27, 32, 36, 38, 53, 82, 83
1, 2, 8, 10, 11, 14, 16, 20, 22, 23, 26, 27, 29, 36, 37, 39, 43,
52, 67, 74, 78
EXAM DATES
Chaps Date
Exam 1
1-3
Sep 14
Exam 2
4-6
Oct 12
Exam 3
7-8
Nov 2
Last day to withdraw W—Nov 2
Chaps
Date
Exam 4
9-10
Nov 16
Exam 5
12-13
Dec 12
Comprehensive final
Dec 14
Last day to withdraw WPWF—Dec 7
Lab grade and pop quiz average count as an exam too!
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DR. STEVE BADGER
GENERAL CHEMISTRY CHEM 111-2
ADDITIONAL INFORMATION
If you need to meet with the professor, please schedule an appointment at the end of class period. My office is
Z-312G and my phone number is 865-2815 ext 8327. You can email me at [email protected]. At times I
use AOL’s Instant Messenger with the screen name DocBadger.
LABORATORY
Every student must have his or her own lab textbook. (Slowinski, Wolsey, Masterton. Chemical Principles in
the Laboratory. 8th Edition, Thomson Wadsworth Publishing, 2005.)
A used book with missing pages is not acceptable, so be careful not to purchase a used book.
You must hand in a completed pre-lab assignment as you enter the lab.
DATE
EXP’T
Sep 3-5
1
Check in, The Densities of Liquids and Solids
Sep 10-12
2
Paper Chromatography
Sep 17-19
3
Fractional Crystallization
Sep 24-26
4
Determination of a Chemical Formula
Oct 1-3
5
Identification of a Compound by Mass Relationships
Oct 8-10
14
Heat Effects & Calorimetry
Oct 15-17
9
Molar Mass of a Volatile Liquid
Oct 22-24
TITLE
TBA
Oct 29-31
10
Analysis of an Aluminum-Zinc Alloy
Nov 5-7
13
Lewis Structures
Nov 12-14
7
Analysis of an Unknown Chloride
Nov 26-28
15
Vapor Pressure and Heat of Vaporization of Liquids
Dec 3-5
20
Rates of Chemical Reaction I: Iodination of Acetone
Dec 10-12
21
Rates of Chemical Reaction II: A Clock Reaction
No make up labs
Cell phones must be turned off and kept out of sight during class & lab. Thanks!
(Revised 11:19 AM August 28, 2007)
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