Course title
ChemistryPre-requisite
Algebra I with a C or better for each semesterCourse description
Chemistry Curriculum
Description of Course
This course is designed to expose the student to principles and applications of chemistry. The topics covered include: measurement; matter; the atom; the periodic table; nomenclature; formulas; equations; energy; bonding; reactions; the mole; acids and bases; biochemistry; gases and solution chemistry.
Course Expectations
When the student has completed the course fulfilling all requirements; he/she will be able:
To make and record observations.
To develop a hypothesis based on their observations.
To design and conduct an experiment to test a hypothesis.
To distinguish between accuracy and precision.
To make qualitative and quantitative measurements.
To identify the number of significant figures in a measurement.
To measure mass and volume.
To calculate density of an object from experimental data.
To distinguish different Metric prefixes.
To convert from American / British standards units to Metric units and vice
versa.
To compare and contrast types of waves in terms of wavelength; frequency;
and energy.
To analyze line spectra to determine the identity of an unknown element.
To analyze covalent compounds in terms of melting point and conductivity.
To compare and contrast properties of covalent and ionic compounds.
To construct Lewis structures for covalent compounds.
To determine the average weights of each isotope of a fictitious element.
To determine the relative abundance of isotopes of a fictitious element.
To calculate from experimental data the atomic mass of a fictitious element.
Laboratory Activities
Laboratory skills are essential in any science course. The laboratory activities included this course provides students with visual and hands-on activities to help with the understanding of the concepts being learned. Students will conduct laboratory activities at home mostly using materials that should already be available; although some materials may need to be purchased. In this course; laboratory activities will account for 40% of the coursework. A total of 80% of the labs are Hands-On and account for 48 hours of the 60 hours of lab work to be completed.
Resources: Students use content powered by Florida Virtual School for the online component of the course. The Scottsdale Unified School District’s adopted materials for textbooks and other resources will be made available to students. A variety of other online websites and references will also be used as notes in the activities and labs.
Students submit the results of all labs; homework; quizzes and tests electronically to the instructor. All final exams are taken in a supervised computer lab at the end of the course. Students must pass their final exam in order to get credit for the course. Students must complete all labs prior to taking the final exam to receive lab credit for course.
Safety: As with any laboratory science; in some investigations safety issues will arise. Prior to starting any laboratory work; the student and a parent or guardian must read; sign; and submit the attached safety contract. Safety rules must be strictly adhered to and safety goggles must be worn for ANY ACTIVITY INVOLVING CHEMICALS; GLASSWARE; or A HEAT SOURCE.
Student Evaluation
Grading: A = 90%+
B = 80-89%
C = 70-79%
D = 60-69%
F = below 60% or failed the final exam
Required Materials:
High speed internet connection (unless all online work done in school eLearning lab)
Computer (see technical requirements below - unless all online work done in school eLearning lab)
Headphones (for listening to the modules)
Digital camera or phone with built in camera
Various lab materials which will be listed for each lab.
Labs:
Copies of all lab reports; pictures; graphs or other materials made will be compiled in a digital notebook. There will be at least three appropriate/relevant pictures for each lab. The individual labs are graded within each module as they are completed and turned in online. Six sample labs from the complete labs list below are included.
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CHEMISTRY – Semester 1 (5510eL)
LAB SEQUENCE
Table of Contents:
1. The Cat’s Meow (Hands-on) – 0.5 hour
2. Raisin’ It Up (Hands-on) – 1.5 hours
3. Density of Solids; Liquids; & Gases (Hands-on) – 2 hours
4. Quantitative and Qualitative Measurements (Hands-on) – 2 hours
5. Hit the Bulls eye (Hands-on) – 0.5 hour
6. Accuracy and Precision: Balance Lab (Virtual) – 1 hour
7. Isotopes & Atomic Mass (Hands-on) – 1.5 hours
8. Calculating Molar Mass (Hands-on) – 1 hour
9. Line Spectra (Virtual) – 1 hour
10. Electron Probability (Hands-on) – 1.5 hours
11. Graphing Periodic Trends (Virtual) – 2 hours
12. Properties of Metals and Non-metals (Hands-on) – 1 hour
13. Formula Writing: Ionic Compounds (Hands-on) – 2 hours
14. Covalent Bonding (Virtual) – 1 hour
15. Chemical Names & Formulas (Virtual) – 1.5 hours
16. Determining Empirical Formulas (Virtual) – 1 hours
17. Chemical and Physical Properties (Hands-on) – 2 hours
18. Conservation of Mass (Hands-on) – 1.5 hours
19. Decomposition of Baking Soda (Hands-on) – 1.5 hours
20. Chemical Reactions (Virtual) – 1 hour
21. S’Mores Stoichiometry (Hands-on) – 2 hours
22. Percent Yield (Virtual) – 1 hour
TOTAL: 30 hours (Hands-on = 21 hours; Virtual = 9 hours)
CHEMISTRY – Semester 2 (5511eL)
LAB SEQUENCE
1. Properties of Water (Hands-on) – 2.5 hours
2. Phase Changes (Hands-on) – 2 hours
3. Gas Laws (Virtual) – 1 hour
4. Paper Chromatography (Hands-on) – 2 hours
5. Solubility of a Salt (Hands-on) – 2 hours
6. Molarity of a Solution (Hands-on) – 2 hours
7. Molar Mass Determination: Freezing Point Depression Method (Hands-on) – 1 hour
8. Calculating Calories (Hands-on) – 1 hour
9. Heat of Fusion (Hands-on) – 2 hours
10. Calorimetry (Virtual) – 1 hour
11. Entropy and Enthalpy (Hands-on) – 1.5 hours
12. Acids and Bases (Hands-on) – 3 hours
13. Neutralization of Antacids (Hands-on) – 1 hour
14. Acid – Base Titration (Virtual) – 1 hour
15. Equilibrium (Hands-on) – 1 hour
16. Oxidation-Reduction (Hands-on) – 1 hour
17. Radioactive Decay (Hands-on) – 1 hour
18. Constructing a Solar Still (Hands-on) – 2 hours
19. Polymers (Hands-on) – 2 hours
20. The Egg: A Biochemical Storehouse (Hands-on) – 1 hour
TOTAL: 30 hours (Hands-on = 27 hours; Virtual = 3 hours)
CHEMISTRY – Semester 1 (5510eL)
LAB DETAILS
Module 1 – Chemistry and You
3. Lesson 1.04 Matter and Measurements
Lab: Density of Solids; Liquids; and Gases
Type: Hands-on.
Materials: scale; Metric ruler; graduated cylinder; wood cube; wood block; 10 pennies; iron nail; small piece of copper; water; 20-30 mL vegetable oil; 5 g salt; red; blue & green food coloring; zip-lock baggie.
Time needed: 2 hours.
Summary: Density is an important physical property of matter. The relationship between an object’s mass and its volume is called density. Density is calculated using the equation:
density = mass/volume. The common unit of density is g/cm3; but since 1 cm3 = 1 mL; we can also use g/mL. In this laboratory activity; students will take measurements of mass and volume for different objects and use their measurements to calculate the object’s density.
Objectives: 1. to measure mass and volume.
2. to define density.
3. to calculate density of an object from experimental data.
AZ State Standards: Strand 1; Concept 2; #1 & 5; Strand 5; Concept 1 #1.
Source: adapted from S. Panzilius: Maine West High School.
4. Lesson 1.06 Unit Conversions
Lab: Qualitative and Quantitative Measurements
Type: Hands-on.
Materials: scale; Metric ruler; graduated cylinder; wood block; cardboard box; small metal pieces; 10 pennies; water.
Time needed: 2 hours.
Summary: Everyone makes and uses measurements. Measurements are fundamental to science as well. It is important to be able to make measurements and to be able to convert between American / British Standard units of measurement and Metric units of measurement.
Qualitative measurements are results given in a descriptive form without using numbers or measuring devices. They are based on observation. Quantitative measurements are results given in a numerical form usually using some type of measuring device. In this investigation; students will make both qualitative measurements and quantitative measurements. They will also practice converting from American / British Standard units of measurement to Metric units of measurement and vice versa.
Objectives: 1. To make qualitative and quantitative measurements.
2. To describe Metric units of length; mass; and volume.
3. To distinguish different Metric prefixes.
4. To convert from American / British standards units to Metric units and vice versa.
AZ State Standards: Strand 1; Concept 2; #1 & 5; Strand 5; Concept 1 #1.
Source: adapted from S. Panzilius: Maine West High School.
Module 4
17. Lesson 4.01 Chemical vs. Physical
Lab: Chemical and Physical Changes
Type: Hands-on.
Materials: apple; beaker; graduated cylinder; test tubes; sodium chloride; sodium bicarbonate; acetic acid; piece of paper; matches; candle; cabbage juice; calcium chloride; ice; watch glass; stove burner; beaker; food coloring; safety goggles.
Time needed: 2 hours.
Summary: Matter has both physical properties and chemical properties. Likewise; changes that occur to matter can be either physical or chemical. Physical changes are changes that do not alter the chemical composition of the substance. In chemical changes; also called chemical reactions; one or more substances change into different substances. A chemical change always results in a change in chemical composition of the substances involved. In this laboratory; students will conduct a series of mini-experiments in order to examine physical changes and chemical changes.
Objectives: 1. To differentiate between physical and chemical changes in matter.
2. To observe evidence for chemical changes (reactions).
AZ State Standards: Strand 5; Concept 1; #1 & 2; Strand 5; Concept 4; #2.
Source: adapted from S. Panzilius: Maine West High School.
CHEMISTRY – Semester 2 (5511eL)
LAB SEQUENCE
5. Lesson 5.06 – Solubility and Concentration
Lab: Solubility Curve of a Salt
Type: Hands-on.
Materials: ammonium chloride crystals; large test tube; 400 mL beaker; thermometer; scale;
scoop; graduated cylinder; safety goggles; oven mitt; stove top burner.
Time needed: 2 hours.
Summary: The solubility of a solute is the amount of solute dissolved in a given amount of a certain solvent at equilibrium; under specified conditions (the ability to dissolve). Increasing the temperature usually increases the solubility of solids in liquids (endothermic changes only); and decreasing the temperature has the reverse effect (exception; gaseous solutions). In this activity; students will construct a solubility curve representing data collected experimentally. Masses of salt will be varied and temperatures required to dissolve it will be recorded.
Objective: 1. To produce a solubility curve.
2. To apply knowledge of solubility curves to predict solubility at different
temperatures.
AZ State Standards: Strand 5; Concept 1; #1.
Source: unknown.
12. Lesson 7.01 – Acids & Bases
Lab: Acids & Bases
Type: Hands-on.
Materials: red cabbage; paring knife; large pot; stove top burner; small Dixie cups or 10 test tubes & test tube rack; vinegar; lime or lemon juice; orange juice; soda water; pickle juice; windex; tomato juice; shampoo; liquid soap; coke; milk; vitamin C; apple slice; baking soda; borax; aspirin; washing soap; table salt; coffee; sugar; mortar and pestle; spoon or small scoop; dropper; pH paper.
Time needed: 3 hours.
Summary: Acids and bases are present in many household products. Acid compounds give foods a sour taste. In solution; acids are electrolytes and they conduct electricity. Bases are compounds that react with acids to form water and a salt. Bases taste bitter and will have a slippery feel. Both acids and bases will cause indicators to change color. In this investigation; students will observe the properties of acids and bases and will determine which household substances are acids and which are bases
Objectives: 1. To describe properties of acids and bases.
2. To produce an acid-base indicator.
3. To identify common products as acids or bases..
AZ State Standards: Strand 5; Concept 1 #3; Strand 5; Concept 4; #11.
Source: unknown.
16. Lesson 7.06 – Oxidation and Reduction
Lab – Oxidation-Reduction
Type: Hands-on.
Materials: 20 dull pennies; ¼ cup vinegar (acetic acid); 1 tsp sodium chloride; shallow glass or clear plastic dish; plastic spoon or fork; 2 clean steel screws or nails (not galvanized) or paper clips; water; measuring spoons; paper towels.
Time needed: 1 hour.
Summary: Oxidation and reduction always occur simultaneously. The substance gaining oxygen is oxidized and the substance losing oxygen is reduced. No oxidation occurs without reduction; and no reduction occurs without oxidation. Reactions that involve these processes are thus called oxidation-reduction reactions or redox reactions for short. In this laboratory; students will explore oxidation-reduction reactions.
Objectives: 1. To define oxidation and reduction.
2. To assign oxidation numbers to elements in a reaction.
3. To describe and explain what occurs during a oxidation-reduction reaction.
AZ State Standards: Strand 5; Concept 4; #12.
Source: Florida Virtual Labs.
Bibliography
Florida Virtual Labs (2011) brainhoney.com
Glencoe - Earth Science: Geology; the Environment; and the Universe.
Harris; Hal (2007) “Decomposition of Baking Soda”. http://www.umsl.edu/~chickosj/chem11/Labmat/DecompositionOfBakingSoda.pdf
http://www.hschem.org/Laboratory/labs.htm#LAB ACTIVITIES (2011) “S’mores Stoichiometry”.
Jaeger; Dave and Suzanne Weisker (1996) Chemistry: Visualizing Matter. Laboratory Experiments. Holt; Rinehart; and Winston: Austin; TX.
Panzilius; Stefan: Maine West High School.
Waterman; Edward L. and Stephen Thompson (2005) Chemistry. Small-Scale Chemistry Laboratory Manual. Prentice Hall: Needham; Massachusetts.
Wilbrahim; Antony C.; Staley; Dennis; D.; Matta; Michael S.; and Edward L. Waterman (2005) Chemistry. Prentice Hall: Needham; Massachusetts.
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Name: _________________________________________
Date: ______________________________
Formula Writing Lab for Ionic Compounds
Introduction:
An enormous amount of ionic compounds are known to exist today; each having its own unique chemical formula and name. There is no way chemists could possibly test each compound for the weight ratios of the atoms in the compound in order to determine its chemical formula. A systematic method of writing formulas was needed to deal with the vast amount of compounds; without having to do an experiment every time a new compound was discovered. In this activity you will try to discover the underlying principle used to write formulas for all ionic compounds.
Procedure:
1. Color all the positive ions (cations) red.
2. Color all the negative ions (anions) green.
3. Cut out all the positive and negative ion squares carefully (don't round the comers or cut
the squares larger or smaller than they are).
4. Locate the data table on page two and the list of names of ions on the bottom of this page.
5. Put the two combining substances together; so that the heights of the squares match.
Keep adding more of the same squares until the heights match! Then; fill in the table and
record the formula of the compound.
6. Proceed on to the next formula on the data table. Once the data table is complete; answer
the questions in the analysis.
Reference:
Cations Name Anions Name
K+ Potassium cation Cl- Choride anion
Na+ Sodium cation Br- Bromide anion
(NH4)+ Ammonium Cation (NO3)- Nitrate anion
Fe2+ Iron (II) cation O2- Oxide anion
Mg2+ Magnesium Cation S2- Sulfide anion
Fe3+ Iron (III) Cation (SO4)2- Sulfate anion
Al3+ Aluminum Cation (PO4)3- Phosphate anion
Data Table
Combining Substances # of Cation
(red) squares
used Number of Anion (green) squares used The charge on each Cation The charge on each Anion The total positive charge The total negative charge Formula
Aluminum &
Bromine 1 3 +3 -1 +3 -3 AlBr3
Sodium & Oxygen
Iron (II) &
Sulfide
Aluminum & Nitrate
Potassium & Sulfate
Sodium & Nitrate
Ammonium & Phosphate
Iron (III) & Chlorine
Iron (II) & Chlorine
Ammonium & Sulfide
Aluminum & Sulfide
Aluminum & Oxygen
Iron (III) & Sulfate
Magnesium & Phosphate
Iron (III) & Nitrate
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Analysis:
1) Looking at the data table on the previous page; what do you think the rule is for making formulas for ionic compounds?
______________________________________________________________________________
______________________________________________________________________________
______________________________________________________________________________
2) Using your rule from number one; write the ionic formula for each of the following combining substances:
a) Lithium and Chromate
Li+ (CrO4)2-
b) Calcium and Hydroxide
Ca2+ (OH)-
c) Lead (IV) and Oxide
Pb4+ O2
3) Using your ion sheet; create three new ionic compounds that do not contain any of the ions used thus far in this activity. List the combining substances; and write the formulas.
Combine 2 Substances Formula
a)
b)
c)
4) In ions like Iron (III) or Lead (IV); what do you think the Roman numerals in parentheses represent?
Cut outs for Formula Writing Lab
Cations Anions
Mg2+
Mg2+ Mg2+ O2- O2- O2-
A13+
A13+
A13+
(PO4)3-
(PO4)3-
(PO4)3-
Fe 2+
Fe 2+
Fe 2+
S2-
S2-
S2-
Na+ Na+ Na+ Cl- Cl- Cl-
K+ K+ K+ Br- Br- Br-
(NH4)+ (NH4)+ (NH4)+ (NO3)- (NO3)- (NO3)-
Fe3+
Fe3+
Fe3+
(SO4)2-
(SO4)2-
(SO4)2-
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Trends in the Periodic Table
Background Information
The arrangement of elements in the periodic table reveals many important trends; or patterns; in the properties of the elements. Many of the trends illustrated by the table are periodic. That is; they repeat within the table. For example; a periodic trend may involve low values at the top of a group; high values at the bottom; and low values once again at the top of the next group. Recall that the atomic number of an element is equal to the number of protons in the nucleus. The atomic mass is the average number of protons and neutrons in the nucleus of an element. Atomic radius is the distance from the center of the atom to the outermost electrons in the atom. The first ionization energy of an element is the energy needed to remove the most loosely held electron of an atom of the element.
Problem: How are elements arranged in the periodic table?
Goals: In this investigation; you will use a modified periodic table to examine data on the properties of elements. You will graph the data and look for trends in the arrangement of elements in the table.
Materials
graph paper 3 pencils of different colors
Procedure
1. Look at the modified periodic table on page 68. Each box contains data about a particular element. Study the key to discover what each number reveals about the element.
2. Study the modified periodic table on page 68 and try to find trends; or patterns; in atomic mass; atomic radius; and first ionization energy. Look at what happens to the values of each with increasing atomic number—either going down a group or across a period.
3. On a sheet of graph paper; make a graph of atomic mass versus atomic number for elements with atomic numbers 3 to 20 in the modified periodic table on page 68. Plot atomic mass on the y-axis and atomic number on the x-axis. Connect the data points using a different colored pencil to connect the points of elements that are in different periods.
4. On another sheet of graph paper; make a graph of atomic radius versus atomic number for elements with atomic numbers 3 to 20. Plot atomic radius on the y-axis and atomic number on the x-axis. Connect the data points with a different colored pencil for each period.
5. Graph on another sheet of graph paper; make a graph of first ionization energy versus atomic number for elements with atomic numbers 3 to 20. Plot first ionization energy on the y-axis and atomic number on the x-axis. Connect the data points with a different colored pencil for each period.
Analysis
1. What happens to atomic mass when going down each group of the periodic table? When going across each period?
2. What happens to atomic radius when going down each group? When going across each period?
3. What happens to first ionization energy when going down each group? When going across each period?
4. Describe the trend you see in atomic mass as the atomic number increases. Does this trend continue beyond the elements you graphed? Explain.
5. Interpret data Describe the trends you see in atomic radius as atomic number increases. Does the pattern repeat itself? If so; at which groups?
6. Describe the trends you see in first ionization energy as atomic number increases. Look at your graph and the periodic table included with this activity. At which element is the first ionization energy going to reach another low point?
Conclusions
1. What properties of the elements show periodic; or repeating; trends when plotted versus atomic number? Explain your answer.
2. Why is the arrangement of elements in the periodic table a good one for displaying trends and making them easily recognizable and understandable?
NAME _________________________________________
COUNTING BY MEASURING MASS
Purpose: To determine the mass of several samples of chemical compounds and use the data to count atoms.
Materials: safety goggles; plastic spoon; weighing paper; small beaker; scale; water; aluminum; NaCl; CaCO3.
Procedure: 1. Measure the mass of about 10 mL of water.
2. Measure the mass of the aluminum bar.
3. Measure the mass of one teaspoon of NaCl.
4. Measure the mass of one teaspoon of CaCO3.
5. Complete the table below showing all calculations.
Data Table (Show your work on a separate sheet of paper):
H2O (l) Al (s) NaCl (s) CaCO3 (s)
Mass (grams)
Molar mass (g/mol)
Moles of each compound
Moles of each element
Atoms of each element
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Analysis:
1. Calculate the molar mass for each substance (for example: CO2 would have a molar mass of 44 g/mol - each of the two oxygen atoms has a mass of 16 g/mol (32 g/mol total) and the carbon has a mass of 12 g/mol).
2. Calculate the moles of each compound. Moles = mass x (1 mole of substance/molar mass of substance). If we had 132 g of CO2 there would be 132 g / 88 g/mol = 2.5 mol CO2).
3. Calculate the moles of each element. Ex: moles of carbon = 2.5 mol CO2 x 1 mol carbon / 1 mol CO2 = 2.5 moles carbon. Moles of oxygen = 2.5 mol CO2 x 2 mol oxygen / 1 mol CO2 = 5 moles oxygen.
4. Calculate the atoms of each element. Multiply the moles of the element by Avogadro’s number.
2.5 moles carbon x 6.02 x 1023 atoms of carbon / 1 mole of carbon = 1.525 x 1024 mol carbon.
5. Which of the four samples contain the most moles?
6. Which of the three compounds contains the most atoms?
Problems (SHOW ALL WORK! USE A SEPARATE SHEET IF NECESSARY!):
1. Determine the number of moles and particles in 84 g of:
a. water
b. aluminum
c. NaCl
d. CaCO3
2. Determine the percent composition for:
a. H2O H = O =
b. CaCO3 Ca = C = O =
School Country
United StatesSchool state
ArizonaSchool city
ScottsdaleHigh school
Arcadia High SchoolSchool Address
8500 E. Jackrabbit RoadSchool zip code
85018Requested competency code
Lab ScienceDate submitted
Approved
YesApproved competency code
- LCHM
- Chemistry