Honors Chemistry
Instructor
Glenn Powell
Phone
336.625.6185
Office
Room 411
E-mail
[email protected],nc.us
Hours:
F2014 – 3rd
S2015 – 1st
Text:
Chemistry; Wilbraham, Staley, Matta, and Waterman; North Carolina ed
Description:
This is an upper division physical science elective course designed to aid the student in
developing critical thinking and study skills required for success at the college level.
Students are expected to analyze and evaluate scientific theories and data either
collected first hand or presented from external sources. It is also expected that students
will demonstrate skill in executing the various mathematical manipulations used in
chemistry as well as being able to explain the theoretical foundations for their use. This
course is generally considered to be challenging and considerable effort on the part of
the student is required to obtain satisfactory mastery.
Learning Expectations
Unit 1a –Developing an Atomic View of Matter
1. Define mass as the measure of atomic “stuff”; contrast with volume - the amount of
space an object occupies.
2. Use a multiple beam or double-pan balance to determine the mass of various
objects. Use a displacement technique to measure volume of irregularly shaped
objects.
3. Given a graph of mass vs volume of a various substances, relate the slope to the
density of the substances. (Develop proficiency in interpreting graphical
representations.)
4. Recognize that density is a intrinsic property of matter (i.e., it can be used to help
identify an unknown substance).
5. Use differences in density of solids, liquids and gases as evidence for differences in
the structure of matter in these phases.
6. Recognize that behaviors such as floating can inform us of internal details which may
not be directly observed only conceptualized.
7. Understand that internal arrangement can be changed by changing the energy
content of an object. Density also depends on temperature.
Unit 1b – States and Energy
1.
Relate observations of diffusion to particle motion and collision in the gas and liquid
phases.
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2.
Relate observations regarding the addition of energy by warming to increased
particle motion
3.
Recognize that energy is stored in an object or system in several ways; for now we
restrict our discussion to:

Kinetic – due to the motion of the particles.

Interaction (potential) – due to the arrangement of the particles in solid, liquid
and gaseous phases.
4.
Describe the ways that energy is transferred between the system and the
surroundings. These are:
5.
 Heating
 Working
 Radiating
Draw energy bar graphs to account for energy storage and transfer in all sorts of
changes. Make up a sample situation and sketch the bar graph.
6.
State the basic tenets of Kinetic Molecular Theory (KMT).
8.
Given a heating/cooling curve for a substance, identify what phase(s) is/are
present in the various portions of the curve, and what the melting and freezing
temperatures for the substance are.
9.
Given a heating/cooling curve for a substance, identify which energy storage
mode is changing for the various portions of the curve.
10.
Given a situation in which a substance at a given temperature undergoes a
change (in temperature, phase or both), sketch a heating/cooling curve that
represents the situation.
11.
State the physical meaning of the heat of fusion (Hf) and heat of vaporization (Hv)
for a given substance. Use these factors to relate the mass of a substance to the
energy absorbed or released during a phase change (at the melting or freezing
temperature).
12.
State the physical meaning of the heat capacity (c) of a substance and use this
factor to relate the mass and temperature changes to the energy absorbed or
released during a change in temperature (with no phase change).
13.
The 3 variables P, V and T are interrelated. Any factor that affects the number of
collisions has an effect on the pressure. You should be able to:

Predict the effect of changing P, V or T on any of the other variables.

P

Explain (in terms of the collisions of particles) why the change has the
effect you predicted.
Explain the basis for the Kelvin scale. Keep in mind that one must use the
absolute temperature scale to solve gas problems.
Use factors to calculate the new P, V or T. Make a decision as to how the
change affects the variable you are looking for.


1
V
P T
V T
Unit 2a – Describing Substance
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1.
Recognize that a pure substance is composed of one kind of particles and that
mixtures contain more than one kind of particles.
2.
Recognize that a pure substance has one set of characteristic properties (density,
mp, bp), whereas a mixture exhibits properties that are a mixture of the properties
of the substance they contain.
3.
Recognize that some pure substances are “compounded” of simpler particles in a
definite ratio.
4.
Describe the evidence that supports the idea that the simple particles have a
property we call charge
5.
Describe a model of the atom that accounts for the fact that neutral particles can
become either positively or negatively charged
6.
Describe the evidence that distinguishes ionic from molecular or atomic solids.
7.
Given the formula of an ionic or molecular substance, state its name.
8.
Given the name of ionic or molecular substance, write its formula.
9.
From the name or formula of a substance determine whether that substance is
ionic or molecular.
Unit 2b -- Atoms
1. Draw the models of the atom proposed by Dalton, Thomson and Rutherford and
state the key features of each.
2. State and explain the evidence that enabled Thomson to claim that cathode rays
were really a stream of negatively charged particles.
3. Describe Rutherford's experiment and why the results he obtained could NOT be
explained by the Thomson model.
4. Name the three basic subatomic particles; state their
a. location in the atom
b. charge
c. mass relative to a proton
5. Distinguish between the atomic number and mass number for an element.
6. Write the symbol of the isotope of an element given the number of protons and
neutrons. Given the symbol of an isotope, state the number of protons, neutrons and
electrons found in an atom or ion of that isotope.
7. Explain how different isotopes of an element are alike; explain how they differ.
8. Determine the average molar mass of an element, given data regarding isotopic
abundance.
Unit 2c -- Electrons
1. Describe the characteristics of waves.
2. Relate the wavelength and frequency of electromagnetic radiation. Solve problems
using c=λf.
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3. Relate the frequency and energy of electromagnetic radiation. Solve problems using
E=hf.
4. Relate energy, frequency, and wavelength of the major regions of the
electromagnetic spectrum.
5. Describe a photon in terms of the energy it carries.
6. Relate the lines in an emission spectrum to the jumps made by electrons from higher
to lower levels in atoms.
7. State the differences between classical mechanics and quantum mechanics in terms
of models of the atom. Contrast orbit and orbital.
8. List the different types of atomic orbitals, and the number and shape of each, for
energy levels 1 through 4.
9. Write an orbital diagram for any element up to atomic number 36.
10. Write the electron configuration for any element in the Periodic Table. Given the
electron configuration, identify the element.
11. Define valence electrons; relate the number of valence electrons of an element to
the group number.
Unit 2d -- Periodicity
1. Define ionization energy; indicate the general trend in ionization energies with
increasing atomic number:
a)
within a given row.
b)
within a given family of elements.
2. Demonstrate your understanding of the concept of grouping of elements in terms of
similar chemical properties. This involves:
a)
correctly grouping a list of elements.
b)
stating to which family an element belongs.
c)
identifying an "unknown" element or at least identifying the group to
which it
belongs.
3. Explain the behavior of elements in a given family in terms of their tendency to
achieve a noble gas electron arrangement. Extend this to groups II and VI as well as
I, VII and VIII.
4. State the general characteristics of three major families: the noble gases, the alkali
metals, and the halogens. This includes:
a)
tendency to gain or lose electrons.
b)
relative reactivity within the group.
c)
equations for reactions with H2, H2O or each other.
d)
trends in m.p. and b.p.
5. State the trends in physical and chemical properties across a given row of the
Periodic Table. This includes:
a)
metallic vs non-metallic character.
b)
atomic size and density.
c)
attraction for electrons.
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6. Using the Periodic Table, predict the formula of the compound formed by a given
pair of elements.
Unit 2e – Bonding
1. Describe the conditions that make a chemical bond primarily covalent or ionic. Use
electronegativity values to predict bond types between atoms that are covalently
bonded.
2. Draw electron-dot diagrams of for ionic and molecular compounds.
3. State the central idea behind VSEPR; for representative elements state
a. the number of electron pairs on the central atom
b. the spatial arrangement of these electron pairs
c. the overall shape of the resulting hydride or halide
4. When carbon atoms form multiple bonds, state the number of electron pairs used for
single bonds by central carbon atoms of that compound and how this affects the
shape of the molecule.
5. Apply your understanding of molecular geometry and bond polarity by predicting
whether a molecule will be polar or non-polar.
6. Predict the solubility of various molecules in solvents based on the polarity of solvent
and solute molecules.
7. Use the concept of hydrogen bonding to explain some of the properties of water
and the structure of complex organic molecules, such as proteins and DNA.
8. Recognize that the attractions between atoms leading to a chemical bond lower
the energy of the bonded atoms. Separating bonded atoms requires an input of
energy.
9. Sketch a reasonable energy bar chart diagram for both an endothermic and an
exothermic chemical reaction. From such a diagram, determine on which side of
the equation to place the chemical potential energy term.
Unit 3a – Counting Particles
1.
Use experimental data to determine the number of items in a sample without
actually counting them.
2.
Given the chemical formula of a substance, determine the molar mass.
3.
Given the mass of a substance, determine the number of moles of the sample, the
number of atoms, or molecules in the sample
4.
Given the number of moles of a substance, find the mass of the sample, the
number of atoms or molecules in the sample
5.
Given the formula of a compound, determine its % composition.
6.
Given data about the % composition of a sample, determine the empirical formula
of the compound.
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7.
Given the empirical formula and information about the molar mass of the
compound, determine the molecular formula.
Unit 3b – Representing Chemical Change
1. Describe chemical changes in terms of rearranging atoms to form new substances.
2. Recognize that the total number of atoms does not change during a reaction because every reactant
atom must be included in a product molecule.
3. Recognize that the total number of particles (sum of the coefficients) can change during a reaction
because of differences in the bonding ratios of each substance
4. Learn to describe reactions in terms of macroscopic observations
5. Learn to describe reactions in terms of microscopic behavior of atoms
6. Learn to write balanced equations to represent these changes symbolically.
7. Explain that the coefficients in a chemical equation describe the quantities of
a. the individual atoms or molecules involved
b. the moles of the substances involved.
8. Observe basic patterns in the way substances react and learn to generalize them to
other reactions students encounter.
a. Synthesis reactions
b. Decomposition reactions
c. Combustion reactions
d. Single replacement reaction
8.
e. Double replacement (ionic) reactions
Describe endo- and exothermic reactions in terms of storage or release of chemical potential energy
Unit 3c – Stoichiometry
1.
Recognize that while mass and the number of atoms are conserved, the number of
molecules or moles in a reaction is not conserved.
2.
3.
Given a chemical reaction stated in words, write a balanced chemical equation.
Starting with a balanced chemical equation, and the number of moles of a reactant or product,
determine the number of moles of any other reactant or product involved.
4.
Starting with a balanced chemical equation, and the mass of a reactant or
product, determine the mass of any other reactant or product involved.
5.
Starting with a balanced chemical equation, the mass of a reactant and mass of
product actually produced, calculate the percent yield.
6.
Starting with a balanced chemical equation, and the mass of the reactants,
determine which reactant is limiting, and what the theoretical yield of a product is.
7.
Given a balanced chemical equation and the amounts of reactants, sketch
diagrams to represent the reaction mixture before and after the reaction.
8.
Use the concept that
P  n to determine the partial pressure of a particular gas in
a mixture.
9.
Determine by experiment the volume of a mole of gas at STP.
10. Use the ideal 
gas law equation to determine the number of moles in a sample of gas not at standard
conditions.
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11.
Relate the molar concentration (molarity) of a solution to the number of moles and
volume of the solution.
12.
Use information regarding molarity and volume to determine the number of moles
of a reactant or product in a given volume of solution.
13.
Sketch energy bar graph diagrams to represent energy storage and eventual
transfer in exothermic and endothermic reactions.
14.
Extend the use of the BCA table, first used in situations where the mass of reactant
or product was known, to cases involving volume of a gas or solution or energy of
reaction.
Units 4 – Reaction Rates and Equilibrium
1. Use Collision Theory to explain how the following factors affect the rate of a reaction.
a. concentration of reactants
b. temperature of the system
c. presence of a catalyst
d. surface area
2. Sketch a curve for a chemical reaction that shows the chemical potential energy of
the reactants, activated complex and products. Sketch energy bar graphs showing
both kinetic and chemical potential energy for these three points in a reaction. On
either type of graph indicate the activation energy and the energy of reaction.
3. Show how the addition of catalyst changes the shape of the potential energy curve
or the energy bar graph for a reaction.
4
Explain, in terms of the distribution of molecular kinetic energy, how a small increase
in the temperature can produce a large increase in the rate of a reaction.
5. List the conditions necessary for a state of equilibrium to exist.
6. State Le Chatelier’s principle in your own words.
7. Given a chemical equation representing a system at equilibrium, predict qualitatively
the effect of the following changes:
a.
b.
c.
d.
change the concentration of a reactant or product
change in temperature
change in pressure
addition of a catalyst
8. Given a chemical equation representing a system at equilibrium, write the expression
for the equilibrium constant.
9. Perform calculations involving the equilibrium constant and the concentrations of the
reactants and products.
Unit 5 – Acids and Bases
1. Identify the properties common to the class of compounds known as acids and
bases.
2. Indicate your understanding of the following terms: strong, weak, concentrated,
dilute as they apply to acids by using them properly in describing solutions.
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3. Use the Bronsted-Lowry definition of acids and bases to identify the proton donor,
proton acceptor, conjugate acid and conjugate base in a given equation.
4. Given the mass (or number of moles) of a known strong acid or strong base and the
total volume of solution, calculate the [H+] and [OH–] in the solution.
5. Given the [H+] or [OH–] of a solution, find the [OH–], the [H+], and the pH.
6. Given the pH of a solution:
a) state whether the solution is acidic or basic.
b) calculate the [H+] and [OH–] of the solution.
7.
Given the volume and concentration of known acid (or base) used to titrate a
base (or acid), calculate the concentration of the unknown solution.
Unit 6 -- Nucleus
1. Distinguish between a chemical change and a nuclear change.
2. Distinguish between ionizing and non-ionizing radiation.
3. Describe for each of the three types of nuclear radiation :
a. source
b. characteristics
c. penetrating power
d. biological hazard
e. potential uses
4. Given an equation for a nuclear decay reaction in which one reactant or product is
missing, supply the missing term.
5. Define half-life; given problems involving the fraction of original material remaining,
elapsed time, or half-life, solve for the missing quantity.
6. Distinguish between nuclear decay and nuclear fission.
7. Identify fissionable isotopes and balance equations for nuclear fission reactions.
8. Distinguish between nuclear fission and nuclear fusion.
Requirements:
Black ink pens
Scientific calculator
Resources:



http://www.chemtutor.com/
http://dbhs.wvusd.k12.ca.us/webdocs/ChemTeamIndex.html
http://antoine.frostburg.edu/chem/senese/101/index.shtml
Evaluation:
The principle means of assessing student performance in this course will be by tests and quizzes
which together will be weighted as 60% of the final score. Most tests will be in two sections. One
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section will be mostly objective type questions intended to evaluate progress toward basic
proficiency in the goals as listed above. Students will be expected to work and re-test as needed
to attain proficiency in these goals. The other section will be composed of open ended questions
intended to assess the student’s ability to engage in higher level thinking. Attaining basic
proficiency will assure the student of receiving credit for this course while demonstrated ability in
higher order thinking will be necessary to attain an A grade.
This course includes a laboratory component which will contribute 30% to the final score.
Documentation of lab work will be by written report. There will be three types of report used during
the semester; a formal report in the style described at Labwrite.edu, informal reports which are
usually very short, and lab practicum reports which are usually very short and submitted as a
group. A rubric will be provided for the formal and informal reports.
The remaining 10% is based on class assignments. This category may also include some homework.
In general homework is done to improve the performance on tests and is not scored except as it
applies to goal 0. It is in the students best interest to make an honest effort to complete all
assignments whether a formal grade will be assigned or not.
Course Schedule:
It is difficult to predict with any certainty exactly where in the course we will be on any
given future date. The SCOS encompasses 20 of the 25 chapters in the text and a
semester is 18 weeks in length with at least a week of this devoted to final exams and
other institutional events. You may see from the table below that we will not be following
the text in exactly the order provided by the authors. The chapters omitted will be 1, 21,
22, 23, and 24. The instructor posts the weeks topics on the board in the classroom. The
student is expected to stay abreast of the current topic and timeline in order to have
read the related text material and to be adequately prepared to participate in class
activities. Notification of tests and other important events are also noted on the weekly
calendar on the board.
Unit
Approx.
SCOS Goals
Text Chapter
ES 2.1.5, 3.2.4, and extensions
2, 3, 13
Time (wk)
1a. Atomic View of
2
Matter
beyond ES
1b. States and Energy
2
ES: 2.1.1, 2.1.2, 2.1.3, 2.1.4, and
13, 14, 17,
ext. multi-step change and nonwater systems.
2a. Describing
<1
Substance
ES:1.2.1, 1.2.3, 1.2.4, 1.3.1, 1.3.2,
2, 9
3.2.1
2b. Atoms
<2
ES: 1.1.1, 1.1.3,
4
2c. Electrons
1
ES: 1.1.2, 1.1.3, 1.2.2, and ext.
4, 25
with non-conforming configs.
Page 9
2d. Periodicity
<2
ES: 1.2.1, 1.3.1, 1.3.2, and ext. by
4, 5
density and reactivity.
2e. Bonding
1
ES: 1.2.1, 1.2.2, 1.2.3, 1.2.5, 1.3.2,
10
3.2.4
3a. Counting Particles
<1
ES: 2.2.5
11
3b. Representing
1
ES: 2.2.2, 2.2.3
12
3c. Stoichiometry
1
ES: 2.2.1, 2.2.2, 2.2.4,
5, 6
4. Reaction Rates and
<1
ES: 3.1.1, 3.1.2, 3.1.3
7, 8
5. Acids and Bases
<1
ES: 2.2.3, 3.2.1, 3.2.2, 3.2.3
15, 16, 19
6. Nucleus
<1
ES: 1.1.1, 1.1.4
20
Chemical Change
Equilibrium
Examinations:
The final exam for this class will contribute to the final score in the same proportion as the three sixweeks scores. The date and time of the exam will be announced.
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