Advanced Chemistry + Advanced Chemistry Lab

Course title

Pre-requisite

Chemistry

Course description

https://www.davidsononline.org/academics/courses/science/advanced-chemistry/

Students participate in:

Semester 1 - Two 1.5 hour Live Sessions/ week that also include labs less frequently than semester 2

Semester 2 - Two 1.5 hour Live Sessions/ week + One 2 hour Lab. Live Session/ Week

Advanced ChemistryCourse Overview

This Advanced Chemistry course is designed to be the equivalent of the general chemistry course taken during the first year of college. Students successfully completing this course will be endowed with an exceptional understanding of the fundamentals of chemistry and achieve proficiency in solving chemical problems. This course will contribute to the development of each student’s ability to think critically and to express their ideas, in both oral and written fashion, with clarity and logic.

Laboratory investigation is a central pillar of the Advanced Chemistry course. Labs will include an emphasis on experimental procedures. Each week during the first semester students will participate in a Friday lab elective and should expect to spend one to two hours each week on lab work outside of regular class time.

Course ContentUnit 1: Lab Safety

The main idea of this unit is to have students work safely when handling chemicals. Students will learn about various types of safety equipment present in a chemistry lab. They will also learn how to safely handle chemicals and equipment in a chemistry lab.

Unit 2: First Year Chemistry Review

Students will get a brief overview on concepts they learned in their previous chemistry course. This refresher unit quickly covers:

classifying matter according to a scheme,
understanding the difference between measured numbers and exact numbers,
solving problems using various units of measurement,
relating atomic theory with atomic structure based on indirect evidence,
describing atomic structure and the properties of atoms, molecules and matter,
describing key terms,
using the periodic table,
comparing empirical formulas from molecular formulas,
calculating empirical and molecular formulas from experimental data,
discussing differences between ionic and molecular compounds,
naming inorganic compounds,
writing balanced chemical equations to describe a chemical reaction for synthesis, decomposition, single replacement, metathesis, redox, combustion, and acid-base reactions,
calculating the molar mass of a substance,
using the molar mass and Avogadro’s number to interconvert among mass, moles and number of particles of a substance, work problems involving mole concepts, molarity, percent composition, empirical formulas, and molecular formulas,
and solving stoichiometric problems.

Unit 3: Aqueous Reactions & Solution Stoichiometry

Students study chemical reactions more deeply. Students will describe the nature of aqueous solutions through water as a solvent and strong and weak electrolytes as solutes, identify common strong and weak acids, determine the solubility of ionic compounds from general solubility rules, write molecular, ionic, and net ionic equations, and identify metathesis reactions that go to completion (formation of a gas, precipitate or molecular product). Students will also predict the products for reactions that are redox, neutralization, and precipitation reactions, and perform stoichiometric calculations on acid-base volumetric (titrations), precipitation, and redox reactions.

Unit 4: Thermochemistry

Students will investigate enthalpy, calorimetry, and Hess’s law. Students will describe the energy flow between a system and its surroundings, explain the significance of the first law of thermodynamics and use the law to calculate ∆E, q, and w, define and distinguish among heat, temperature, work, energy, kinetic and potential energy, calculate the enthalpy change associated with phase changes, determine the enthalpy change or stoichiometric quantities for thermochemical equations, describe a state function, use standard enthalpies of formation to calculate ∆H for a reaction, solve calorimetry problems using q = mc∆T, use Hess’s Law to calculate the enthalpy change for a reaction, use standard enthalpies of formation to calculate ∆H for a reaction, interconvert among calories and Joules.

Unit 5: Electronic Structure of Atoms

Students will investigate Bohr theory, the wave mechanical behavior of an atom, and quantum mechanics. Students will determine various aspects of the electromagnetic spectrum including relative frequencies, wavelengths and energies, quantitatively and qualitatively relate frequency, wavelength and speed of a wave, describe Planck’s concept of quantized energy and calculate the energy of a photon using the relationship λ = hν, relate Bohr’s model of the atom to the quantum theory, calculate the energy difference resulting from the change in energy levels of an electron, state the meaning and possible values of the quantum numbers and assign the quantum numbers to a given sublevel or orbital, and use the quantum numbers, Aufbau Principle, and Hund’s Rule to assign an electron configuration for a given element or ion.

Unit 6: Chemical Bonding

Students will investigate periodic trends by interpreting trends within the periodic table in terms of: atomic radii, ionization energy, electron affinity, and ionic radii, distinguishing between metals and nonmetals and semimetals, describe how effective nuclear charge varies with position on the periodic table, and compare the relative energies of atomic energy levels and of sublevels. Students will also use periodic trends and electronegativity to predict bond types, compare and contrast different types of bonding, compare bond strength with ionic sizes of elements on the periodic table, relate the enthalpy dissociation of ionic bonding to bond strength, draw Lewis structures for various atoms, ions, and molecules, draw resonance structures for various molecules, use formal charges to determine the most likely resonance structure, compare oxidation numbers and formal charges for atoms in a molecule, relate electronegativity values to bond polarity, and compare and contrast bond distance and bond energy for single and multiple bonds. While investigating molecular shapes, students will use the VSEPR model to predict molecular geometry, determine molecular polarity using dipole moments of individual bonds, compare VSEPR structures to the hybridization of orbitals, compare and contrast sigma and pi bonds, predict the number of sigma and pi bonds in a structure, compare and contrast valence bond theory with molecular orbital theory, and contrast molecular orbitals (delocalized) and orbitals derived from the valence-bond theory.

Unit 7: Gases

Students will examine the relationship between pressure, volume and temperature of ideal gases, apply Charles’, Boyle’s, Gay-Lusaac’s, Dalton’s, and the ideal gas laws quantitatively and qualitatively, analyze the kinetic molecular theory, use Graham’s Law to relate the molar masses of gases to their rates or times of effusion, describe how real gases deviate from ideal behavior, show how Van Der Waals’ equation allows for real conditions, use the ideal gas law equation to calculate the density or molar mass of a gas and solve stoichiometric calculations at standard and non-standard conditions, and use the molar volume at STP conditions in calculations.

Unit 8: Intermolecular Forces

In this unit, students will describe the intermolecular forces such as dipole-dipole, hydrogen bonding, and London dispersion forces, describe the effects that IM forces have on the properties of liquids and solids such as melting point, boiling point, vapor pressure, viscosity, state of matter, phase changes and solubility, characterize the processes of evaporation, condensation, sublimation, fusion at the particle level, distinguish among ionic, molecular, network covalent and metallic solids with regard to particle structure, physical properties, and inter- and intramolecular forces, apply the concepts of unit cells and crystal lattices for solids to calculations involving atomic radii, volume, density or identity, explain the relationship of boiling point to vapor pressure, using phase diagrams, and be able to calculate energy of various phase changes for water and carbon dioxide.

Unit 9: Solution Chemistry

Students will define and describe solution formation, energy changes, salvation and hydration as they relate to solutions, describe the unique characteristics of water due to its extensive hydrogen bonding, compare and contrast saturated, unsaturated and supersaturated solutions; and be able to interpret graphs and charts of solubility, make calculations involving molarity, molality, mass percent and mole fractions as a means of expressing concentration, analyze the effects of colligative characteristics on the properties of solutions such as electrolytes vs. non-electrolytes, solve problems involving freezing point depression, boiling point elevation, vapor pressure lowering, and increase in osmotic pressure, use Raoult’s Law to relate vapor pressure lowering to solute mole fraction, explain different properties of colloidal systems such as size of particles, Tyndall’s effect, and Brownian’s motion.

Unit 10: Kinetics

Students will describe the collision theory and the requirements for an effective collision, list factors that affect the rate of a reaction, use experimental data to determine the rate law and rate order of a reaction and to predict a reaction mechanism, interpret graphs of endothermic and exothermic reactions identifying the activation energy, enthalpies, and the reaction course with and without a catalyst, determine a zero, first or second order reaction from graphical analysis of concentration vs. time plots, explain the role of a catalyst in a reaction and distinguish between homogeneous and heterogeneous catalysts, predict how temperature and concentration affect the rate of a reaction over time, use data to calculate the half life of a reaction, and generally describe the meaning and use of the Arrhenius equation and be able to solve problems involving activation energy and the Arrhenius equation.

Unit 11: Equilibrium & Acid/Base Theory

Students will discuss the concept of equilibrium, write the equilibrium expression for a given equilibrium system in terms of concentrations or pressures, calculate values for any of the equilibrium constants, given Kc or Kp, calculate the equilibrium concentrations for the species in the system, predict the changes in equilibrium that will occur when various stresses are placed on the system (Le Chatelier’s Principle): concentration change, temperature change, pressure change, and addition of a catalyst, calculate pH, pOH, pK, Ka, Kb, ionization constant, percent ionization, Ksp, use the reaction quotient, Q, to determine the initial direction of a reaction needed to establish equilibrium, write the Kw expression for water, explain the common ion effect, identify strong and weak acids and bases and write dissociation equations for each, predict the direction of equilibrium from knowledge of the strength of the acid-base conjugate pair in water, solve problems involving concentrations of substances necessary to produce a precipitate, and concentrations of ions involved in simultaneous equilibrium, graphically determine pKa for a weak acid from a titration curve, given the composition of a buffer system, determine its pH before and after the addition of known amounts of strong acid or base, determine the proportions in which a weak acid and its conjugate base should be mixed to give a buffer of a specified pH, and use the Henderson-Hasselbach equation in equilibrium (buffered) reactions.

Unit 12: Thermodynamics

Students will discuss the laws of thermodynamics, define entropy, second law of thermodynamics, PV work, enthalpy and free energy, use Hess’s Law to solve problems of energy, entropy and free energy, relate the signs of ∆H and ∆S to determine the direction of a reaction/determine the spontaneity of a reaction, predict the sign of entropy change for a given reaction, apply the relationship between ΔS, surroundings, ΔH, and Temperature, describe how the signs of ΔH, ΔS, and ΔG relate to the spontaneity of a reaction, calculate the free energy change for a given reaction, and calculate ΔS for reactions or phase changes from Absolute Entropy values.

Unit 13: Electrochemistry

Students will assign oxidation states to various elements in compounds and molecules, recognize reactions that undergo redox reactions by comparing their oxidation states, recognize oxidation/reducing agents in various reactions, balance redox reactions using the half reaction method, balance redox reactions in acidic or basic solutions, draw and label parts of an e-cell showing electron flow, differentiate between galvanic and electrolytic cells, use the Nernst equation to calculate the EMF at non-standard conditions, apply Faraday’s Law to electrolytic cells in calculating amount of products formed, time required, or current required, use the table of standard reduction potentials to determine cell voltages, explain the electrochemical nature of lead storage batteries, corrosion (anode and cathode protection), and fuel cells, define terms such as redox, anode, cathode, oxidizing agent, reducing agent, emf, electrode, Faraday’s Law, voltaic cells, and galvanic cells.

Unit 14: Nuclear Chemistry

Students will research alpha, beta, and gamma radiation by describing each type of radiation in aspect of mass, charge, penetrating power, and symbol. Students will also write balanced nuclear equations for radioactive decay and nuclear transformations, given a table of nuclear masses, calculate Δ mass for a nuclear reaction and relate it to the energy change, ΔE (binding energy), generally describe the functioning, reactions, and positive and negative aspects of fission and fusion reactors, and generally discuss the rates of radioactive decay.

Syllabus for Advanced Chemistry 2023-2024

Course Description: This Advanced Chemistry course is designed to be the equivalent of the general chemistry course taken during the first year of college. Students successfully completing this course will be endowed with an exceptional understanding of the fundamentals of chemistry and will achieve proficiency in solving chemical problems. This course will contribute to the development of each student’s ability to think critically and to express their ideas, in both oral and written fashion, with clarity and logic. Students must be disciplined, self- motivated and industrious.

Course Objectives:

Quantitatively and qualitatively describe matter and its changes by applying concepts of liquids, solids, gases, solutions, chemical reactions, atomic theory, chemical bonding, stoichiometry, equilibrium, kinetics, and thermodynamics.
Apply and analyze chemical concepts through chemical calculations such as percent composition, molar masses, empirical formulas, gas laws, mole fractions, chemical kinetics, and standard electrode potentials and their use.
Create, conduct, and analyze the laboratory experiments to engage and reinforce concepts taught throughout the course.
Demonstrate critical and independent thinking, collaboration, and an appreciation for the natural world.

Materials:

OPTIONAL Text: Chemistry: The Central Science, 15th edition, Brown, LeMay, Bursten.
MyLab & Mastering access: Students will purchase MyLab & Mastering through our Canvas course to access homework problems and the online textbook.
A scientific calculator is required. A graphing calculator will be extremely useful, but it is not required.
Students will purchase lab kits from Science Interactive and will need to provide common household supplies (see list).

Course Requirements:

Laboratory: Hands-on laboratory investigation is a central pillar of the Advanced Chemistry course. Labs will include an emphasis on experimental procedures. Each week during the second semester, students will participate in a Friday lab elective and should expect to spend one to two hours each week on lab work outside of regular class time.

Homework: Homework will be assigned with every unit. Regardless of whether homework is assigned or not, time should be given to this course on a regular basis through re-reading the text, watching videos, and working on extra practice problems from the textbook. It is imperative that students show all work with units, use significant figures (where applicable), and place a box around your answer for full credit.

Tests and Quizzes: There will be short quizzes on topics covered in class. There will also be unit

exams that go along with each unit covered in class. Honorlock will be used on all exams.

Grading:

Assignments are divided into three categories: Formative assignments (homework and quizzes), summative assignments (major tests and projects), and the semester exam. The semester exam score may replace one unit exam score if the result is a higher overall semester grade. No other opportunities (extra credit, corrections, etc.) will be offered to improve low scores. In exceptional cases, and at the teacher's discretion, replacement of scores will not be allowed. Grades are NOT weighted, but below is the approximate percentage that each category covers:

Assignment type

% of overall grade

Homework, Quizzes (formative)

25-30%

Tests (summative)

50-55%

Semester Exam

20-25%

Laboratory Grading:

Students will be graded based on the quality of their lab reports and demonstrated understanding of the topics covered in each lab. Individual labs are associated with a total possible number of points that is reflected in the complexity of the lab. All lab work is graded within the Science Interactive cloud. Lab reports will collectively count for approximately 90% of the total grade and approximately 10% will come from lab etiquette (preparation, engagement, safety, etc.).

Laboratory Guidelines: The laboratory component is an essential part of the Advanced Chemistry course, because it is the time when we practice chemistry and apply what we have learned in class to real situations. You will learn techniques and develop skills that will be essential as you transition to advanced science courses in college.

Through our carefully designed lab kits, we bring an authentic lab experience into your homes. Below are additional details to set you up for success in the lab:

Expectation

Purpose

Steps to take

You will perform individual experiments while collaborating with peers.

We learn more when we work with each other. In a traditional laboratory course, each lab group performs the same experiment at the same time. You will be able to ask and answer questions, as well as share data when appropriate.

Have an appropriate work space reserved for lab days. The space should allow you to:

safely mix chemicals.
use an open flame.
have your camera on so that others can see your experiment.

You should have access to a sink for most labs.

Have all supplies ready to go.

During lab, you need to keep up with the class to ensure meaningful collaboration and that we finish the lab in a reasonable amount of time.

Review the materials list that was emailed to you. The list is also available in Canvas.
Check the announcements to see if any labs are scheduled for the upcoming week.
Gather lab materials at least 1-2 days prior to the lab.
Ensure all supplies are out of their packaging and neatly organized in your lab space.

Keep your camera on.

Safety is always our main priority in lab. While you may have to complete some lab portions outside of class, steps involving strong chemicals or use of a flame will usually be done in class. As the instructor, I need to be able to watch students to ensure they wear goggles and abide by the guidelines discussed in our safety lab. I also need to observe them using correct laboratory techniques and procedures.

Enter class with your camera on and keep it on throughout the entire lab.

Complete all pre-labs (Explorations) ahead of time.

You will have a better understanding of what the lab entails after completing the pre-lab. When all students come to class prepared, the entire class benefits.

Diligently check the announcements for upcoming assignments and labs.
Remember that you are part of a team and your preparation impacts your peers.
Don’t wait until the last minute to complete the Exploration section for an upcoming lab.

Classroom Policies:

Assignments that are original writing will be run through CopyLeaks, which is a plagiarism checker.
All tests and quizzes will be taken with Honorlock turned on to monitor web activity during your quiz or test.
If you are having tech problems with assignment or exam submissions, you need to contact Ben Brown immediately so that your work is not late.
You may be required to work with a STEM Mentor to support you with the course. You will be notified if that is the case.

For other class policies, please see the DAO handbook and the “Syllabus Addendum for Core DAO Classes”.

Course Topics:

I. Matter and Measurement

a. Matter, its classification and properties

b. Units of measurement

c. Uncertainty in measurement

d. Dimensional Analysis

The student will:

1. classify matter according to a classification scheme among the properties of homogeneous, heterogeneous, pure substance, mixture, solution, element and compound

2. understand the difference between measured numbers and exact numbers and that

uncertainties always exist in measured numbers

3. solve problems using various units of measurement including those with length, mass, temperature, volume, and density

4. distinguish between accuracy and precision in measurements

5. solve problems using the metric system and dimensional analysis and proper significant figures

6. name common polyatomic ions given the formulas and vice versa

II. Elements, Molecules, and Ions

a. Atomic structure, isotopes, atomic numbers, mass numbers

b. The periodic table

c. Molecules and molecular compounds

d. Ions and ionic compounds

e. Naming inorganic compounds

The student will:

1. relate atomic theory with atomic structure based on indirect evidence

2. describe atomic structure and the properties of atoms, molecules and matter

3. define and describe key terms such as isotopes, atomic number, mass number, chemical and empirical formulas

4. use the periodic table to accurately predict trends within the families and periods

5. distinguish among metals, nonmetals, and metalloids on the periodic table

6. compare empirical formulas from molecular formulas and be able to calculate empirical and molecular formulas from experimental data

7. discuss differences between ionic and molecular compounds

8. name inorganic compounds, including acid using a set of systematic rules

III. Stoichiometry: Calculations with Chemical Formulas

a. Chemical equations

b. Atomic and molecular weights; the mole

c. Masses of atoms and molecules

d. Empirical formulas from chemical analysis

e. Quantitative information from balanced equations

f. Limiting reagents

g. % composition

The student will:

1. write balanced chemical equations to describe a chemical reaction for synthesis,

decomposition, single replacement, metathesis, redox, combustion, and acid-base reactions 2. calculate the molar mass of a substance, use the molar mass and Avogadro’s number to interconvert among mass, moles and number of particles of a substance

3. work problems involving mole concepts, molarity, percent composition, empirical formulas, and molecular formulas

4. solve stoichiometric problems involving percent yield, and limiting reagents

IV. Aqueous Reactions and Solution Stoichiometry

a. Solution composition and concentration

b. Properties of solutes in solution

c. Solutions of acids, bases, and salts; neutralization

d. Ionic equations

e. Metathesis reactions

f. Solution stoichiometry and chemical analysis

The student will:

1. describe the nature of aqueous solutions through water as a solvent and strong and weak electrolytes as solutes

2. identify common strong and weak acids

3. determine the solubility of ionic compounds from general solubility rules

4. write molecular, ionic, and net ionic equations

5. identify metathesis reactions that go to completion (formation of a gas, precipitate or molecular product)

6. predict the products for reactions that are redox, neutralization, and precipitation reactions 7. perform stoichiometric calculations on acid-base volumetric (titrations), precipitation, and redox reactions

V. Thermochemistry

a. The first law of thermodynamics

b. State Functions

c. Enthalpy

d. Enthalpy changes

e. Calorimetry including working problems with calories and specific heat

f. Hess’s Law

g. Enthalpy of formation

The student will:

1. describe the energy flow between a system and its surroundings

2. Explain the significance of the first law of thermodynamics and use the law to calculate ∆E, q and w

3. define and distinguish among heat, temperature, work, energy, kinetic and potential energy 4. calculate the enthalpy change associated with phase changes

5. determine the enthalpy change or stoichiometric quantities for thermochemical equations 6. use Hess’s Law to calculate the enthalpy change for a reaction

7. describe a state function

8. use standard enthalpies of formation to calculate ∆H for a reaction

9. solve calorimetry problems using q = mc∆T

10. interconvert among calories, Calories, and Joules

VI. Electronic Structure of Atoms

a. Light and quanta

b. The Bohr model of the atom and electron energies

c. Wave behavior of matter

d. The quantum mechanical model of the atom

e. Atomic orbitals

f. Electron configurations

The student will:

1. determine from the electromagnetic spectrum: relative frequencies, wavelengths and energies 2. quantitatively and Qualitatively relate frequency, wavelength and speed of a wave

3. describe Planck’s concept of quantized energy and calculate the energy of a photon using the relationship λ = hν

4. relate Bohr’s model of the atom to the quantum theory

5. calculate the energy difference resulting from the change in energy levels of an electron 6. state the meaning and possible values of the quantum numbers and assign the quantum numbers to a given sublevel or orbital

7. use the quantum numbers, Aufbau Principle, and Hund’s Rule to assign an electron configuration for a given element or ion

VII. Periodic Properties of the Elements

a. Atomic sizes

b. Ionization energies

c. Electron affinities

d. Metals, nonmetals, and semimetals

e. Group trends for Groups 1, 2, 16, 17, and 18

The student will:

1. interpret trends within the periodic table in terms of: atomic radii, ionization energy, electron affinity, and ionic radii

2. distinguish between meals and nonmetals and semimetals

3. describe how effective nuclear charge varies with position on the periodic table

4. compare the relative energies of atomic energy levels and of sublevels

VIII. Basic Concepts of Chemical Bonding

a. Ionic bonding and energetics of ionic bonding

b. Ionic sizes

c. Covalent Bonding

d. Lewis Structures

e. Bond polarity and Electronegativity

f. Covalent bond strength

g. Oxidation numbers and formal charge

The student will:

1. use periodic trends and electronegativity to predict bond types

2. compare and contrast different types of bonding

3. compare bond strength with ionic sizes of elements on the periodic table

4. relate the enthalpy dissociation of ionic bonding to bond strength

5. draw Lewis structures for various atoms, ions, and molecules

6. draw resonance structures for various molecules

7. use formal charges to determine the most likely resonance structure

8. compare oxidation numbers and formal charges for atoms in a molecule

School country

United States

School state

Nevada

School city

Reno

High school

Davidson Academy of Nevada

School / district Address

Davidson Academy Online, 9665 Gateway Dr., Suite A, Reno, NV

School zip code

89521

Requested competency code

Lab Science

Date submitted

10/23/2023

Approved

Yes

Approved competency code

LADV
Advanced science

LCHM
Chemistry

Approved date

11/14/2023

Online / Virtual

Yes