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STANDARDS OF LEARNING CONTENT REVIEW NOTES CHEMISTRY Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin OVERVIEW Chemistry Content Review Notes are designed by the High School Science Steering Committee as a resource for students and parents. Each nine weeks’ Standards of Learning (SOLs) have been identified and a detailed explanation of the specific SOL is provided. Specific notes have also been included in this document to assist students in understanding the concepts. A “ ” section has also been developed to provide students with the opportunity to check their understanding of the content. The document is a compilation of information found in the Virginia Department of Education (VDOE) Curriculum Framework, Enhanced Scope and Sequence, and Released Test items. In addition to VDOE information, Glencoe Textbook Series and resources have been used. Finally, information from various websites is included. The websites are listed with the information as it appears in the document. The Chemistry Blueprint Summary Table is listed below as a snapshot of the reporting categories, the number of questions per reporting category, and the corresponding SOLs. It is the Chemistry Instructors’ desire that students and parents will use this document as a tool toward the students’ success on the end-of-year assessment. Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin Standards of Learning Notes & Activities Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin First Nine Weeks LAB SAFETY EQUIPMENT ATOMIC STRUCTURE SOL 1C SOL 2I SOL 2A SOL 2C SOL 2B SCIENTIFIC METHOD SOL- 1D, E REPORTING SCIENTIFIC DATA SOL- 1E, G SOL 1G SOL- 1A,E,F SOL- 1F MATTER AND ENERGY SOL 2H PERIODIC TABLE SOL 3 A SOL- 2D,E SOL 2A SOL 2F SOL 2C ELECTRON CONFIGURATION SOL 2I SOL-2 A, B,C, E,G SOL-2 A, B, C, D, G Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin TOPIC: LAB SAFETY EQUIPMENT CHEMISTRY STANDARD CH. 1C The student will investigate and understand that experiments in which variables are measured, analyzed, and evaluated produce observations and verifiable data. Key concepts include: c) proper response to emergency situations 1. What is the correct procedure to neutralize an acid? 2. When diluting an acid solution, one should always do what and why? 3. Describe the sequence of actions you would take if you spill a strong acid on your clothing in the lab. 4. Describe what you would do if your lab partner’s shirt-sleeve caught on fire. 5. When is it most important to wash your hands while in the laboratory? 6. Describe the procedure to smell a substance? Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin TOPIC: SCIENTIFIC METHOD CHEMISTRY STANDARD CH. 1D, E The student will investigate and understand that experiments in which variables are measured, analyzed, and evaluated produce observations and verifiable data. Key concepts include: d) manipulation of multiple variables, using repeated trials; and e) accurate recording, organization, and analysis of data through repeated trials. Student A Student B Student C Student D Student measurement of Temperature in ° C Reading 1 Reading 2 88.6 88.5 92.6 90.1 90.0 88.9 80.1 89.8 Reading 3 88.7 91.6 92.5 90.0 1. Use the above chart to answer the following question. Four students each took temperature readings of a sample of water. The actual temperature of the sample was 90.0 ° C. Which student’s measurements were both accurate and precise? 2. A student was asked to determine the density of a mineral. Outline the steps in which to accurately determine the density using the displacement method. 3. What is the first step that should be taken when a caustic chemical gets into a person’s eye? 4. What is the mass of a substance whose density is 0.2396 g/mL and a volume of 20.0 ml? 5. What is the minimum number of trials needed for a measurement to be reliable? Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin TOPIC: REPORTING SCIENTIFIC METHOD CHEMISTRY STANDARD CH. 1E,G The student will investigate and understand that experiments in which variables are measured, analyzed, and evaluated produce observations and verifiable data. Key concepts include: e) accurate recording, organization, and analysis of data through repeated trials. g) mathematical manipulations (SI units, scientific notation, linear equations, graphing, ratio and proportion, significant digits, dimensional analysis). 1. 2. 3. 4. 5. 6. Determine the number of significant figures in the following: a. 5006 b. 1200 c. 0.001200 d. 5 Compute the following. Record your answer to the correct number of significant figures. a. (4.33 x 103)(1.6 x 10 -2) b. 2.36 x 102 x 4.2655 x 104 c. 9.65 x 105/7.2 x 10-3 d. 2.33 x 10-3 x 1.05 x 102/1.234 The SI base unit for mass is the ______________________. The SI base unit for the volume of a liquid is the ________________. Use the factor label method or dimensional analysis to solve the following problems. a. Convert 950 m to km b. Convert 2.56 kg to g c. Convert 0.7526 mg to kg d. Convert 4.85 L to mL Convert the following temperatures. a. 543K to Celsius b. 23K to Celsius Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin CHEMISTRY STANDARD CH. 1 A, E, F The student will investigate and understand that experiments in which variables are measured, analyzed, and evaluated produce observations and verifiable data. Key concepts include: a) designated laboratory techniques e) accurate recording, organization, and analysis of data through repeated trials. f) mathematical and procedural error analysis 1. What piece of equipment would serve the purpose for each of the following? a. Adding 5 ml of 5% HCl to 5 different solutions. ______________ b. To measure 97.0 ml of water. _______________ c. To mass a solid substance. ________________ d. To dry out a precipitate for measuring. _______________ 2. Which of the following items should be used if you need approximately 100 ml of a liquid? (Choose one) a. Beaker b. Erlenmeyer flask c. Graduated cylinder 3. A student measures the mass of a piece of copper three times and records the results in the following table: Trial Mass (grams) 1 15.5 2 15.6 3 15.7 The actual mass of the metal is 17.2 grams. Are the trial measurements accurate and/or precise? Explain your answer. 4. A student measured the temperature of a solution and found it to be 46.0◦C at standard pressure. The theoretical temperature of that boiling solution is 45.0◦C. What is the percent error according to the student’s measurement? Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin CHEMISTRY STANDARD CH. 1G The student will investigate and understand that experiments in which variables are measured, analyzed, and evaluated produce observations and verifiable data. Key concepts include: g) mathematical manipulations including SI units, scientific notation, linear equations, graphing, ratio and proportions, significant digits and dimensional analysis. DIRECTIONS: Show all work using dimensional analysis 1. Convert 3.5 x 10-3 L to mL. 2. How many kilograms are there in 4223 g? 3. Using the rules for significant figures, what is the correct answer for 2.5 m times 44.8 m? 4. (3.5 x 106)(5.3 x 10-5) = Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin CHEMISTRY STANDARD CH. 1 F The student will investigate and understand that experiments in which variables are measured, analyzed, and evaluated produce observations and verifiable data. Key concepts include f) mathematical and procedural error analysis DIRECTIONS: Determine the percent error in the following problems. 1. Experimental Value = 1.44 g Accepted Value = 1.54 g 2. Experimental Value = 22.2 L Accepted Value = 22.4 L 3. Experimental Value = 128.6 mg Accepted Value =128.3 mg 4. Experimental Value = 9.98 x 10-3 g Accepted Value =1.03 x 10-2 g Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin TOPIC: MATTER AND ENERGY CHEMISTRY STANDARD CH. 2H The student will investigate and understand that the placement of elements on the periodic table is a function of their atomic structure. The periodic table is a tool used for the investigations of: h) chemical and physical properties. Chemical and Physical Properties The student will be given an object or substance and must identify the physical and chemical properties of that object. (This assignment is to be given after lecture on chemical and physical properties and how to correctly identify each.) 1. Describe the chemical properties of this object below, citing any evidence you have collected: 2. Describe the physical properties of this object below, citing any evidence you have collected: Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin TOPIC: ATOMIC STRUCTURE CHEMISTRY STANDARD CH. 2I The student will investigate and understand that the placement of elements on the periodic table is a function of their atomic structure. The periodic table is a tool used for the investigations of: i) historical and quantum models. Students will create a two page foldable for the models of an atom. Each outside flap will be the names of the scientists (Democritus, Dalton, Rutherford, Bohr), and the inside picture will be each scientist’s contribution to the historical model. A one page foldable will be created for the quantum models with Heisenberg and Planck. Students will be creative using mnemonics and/or graphics. Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin CHEMISTRY STANDARD CH. 2A The student will investigate and understand that the placement of elements on the periodic table is a function of their atomic structure. The periodic table is a tool used for the investigations of a) average atomic mass, mass number, and atomic number Element Ne Al Ge P Br Ca K I Sr Sn Average Atomic mass Mass number Atomic number Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin CHEMISTRY STANDARD CH. 2C The student will investigate and understand that the placement of elements on the periodic table is a function of their atomic structure. . The periodic table is a tool used for the investigations of: c) mass and charge characteristics of subatomic particles Warm-up: Complete the following chart: Subatomic Particle Symbol Charge Relative mass Location in atom protons neutrons electrons Atomic Structure Element/Ion 3 1 Mass Number Protons Neutrons Electrons H+ C Li+ 35 Cl- 39 24 Atomic Mass H 12 7 Atomic Number K Mg2+ 77 As3- 108 Ag 30 -2 S 238 U Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin CHEMISTRY STANDARD CH. 2B The student will investigate and understand that the placement of elements on the periodic table is a function of their atomic structure. The periodic table is a tool used for the investigations of: b) isotopes, half lives, and radioactive decay DIRECTIONS: Answer the following questions Average Atomic Mass 1. Rubidium has two common isotopes, Rb-85 and Rb-87. If the abundance of Rb-85 is 72.2% and the abundance of Rb-87 is 27.8%, what is the average atomic mass of rubidium? 2. Uranium has 3 common isotopes. If the abundance of U-234 is 0.01%, the abundance of U-235 is 0.71% and the abundance of U-238 is 99.28%, what is the average atomic mass? 3. Titanium has five common isotopes: Ti-46 (8.0%), Ti-47 (7.8%), Ti-48 (73.4%), Ti-49 (5.5%), and Ti-50 (5.3%). What is the average atomic mass of titanium? 4. Explain why atoms have different isotopes. In other words, how is it that helium can have three different sized atoms and they all are still the element Helium. 5. Draw and label 3 possible isotopes of hydrogen. Which isotope can you predict to be the most abundant? Why? Correctly name the 3 isotopes using the symbol “H”? 6. Select an element on the Period Table. Create 4 isotopes for your element and form an average atomic mass calculation to support the given atomic mass. Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin TOPIC: PERIODIC TABLE CHEMISTRY STANDARD CH. 3A The student will investigate and understand how conservation of energy and matter is expressed in chemical formulas and balanced equations. Key concepts include a) nomenclature Writing Binary Formulas Cation Anion Na+ Formula Cation Anion Cl- Fe+2 O-2 Ba+2 F- Fe+2 O-2 K+ S-2 Cr+2 S-2 Li+ Br- Cr+3 S-2 Al+3 I- Cu+ Cl- Zn+2 S-2 Cu+2 Cl- Ag+ O-2 Pb+2 O-2 Mg+2 P-3 Pb+4 O-2 Ni+2 O-2 Mn+2 Br- Ni+3 O-2 Mn+4 Br- Formula Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin CHEMISTRY STANDARD CH. 2 D, E The student will investigate and understand that the placement of elements on the periodic table is a function of their atomic structure. The periodic table is a tool used for the investigations of: d) families or groups; e) periods Counting Valence Electrons Element Group Number Number of Valence Electrons Ne Al Ge P Br Ca K I Sr Sn Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin CHEMISTRY STANDARD CH. 2 F The student will investigate and understand that the placement of elements on the periodic table is a function of their atomic structure. The periodic table is a tool used for the investigations of: f) trends including atomic radii, electronegativity, shielding effect, and ionization energy. DIRECTIONS: State the periodic trend and explain. PERIODIC TRENDS LEFT TO RIGHT EXPLAINATION TOP TO BOTTOM EXPLAINATION ATOMIC RADII ELECTRONEGATIVITY SHIELDING EFFECT IONIZATION ENERGY Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin CHEMISTRY STANDARD CH. 2 C The student will investigate and understand that the placement of elements on the periodic table is a function of their atomic structure. The periodic table is a tool used for the investigations of: c) mass and charge characteristics of subatomic particles. DIRECTIONS: Complete the chart below. Ions and Subatomic Particles Ion Symbol S Protons Electrons Charge 2- K1+ Ba2+ Fe3+ Fe2+ F1O2P3Sn4+ Sn2+ N3Br1Mg2+ Cu1+ Cu2+ U6+ Mn5+ Cl1Se2- Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin Second Nine Weeks IONIC AND COVALENT BONDING AND COMPOUNDS SOL- 3 A, C, D MOLE/EMPIRICAL AND MOLECULAR FORMULA SOL- 4A, B Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin TOPIC: IONIC COMPPOUNDS CHEMISTRY STANDARD CH. 3A,C, D The student will investigate and understand how conservation of energy and matter is expressed in chemical formulas and balanced equations. Key concepts include a) nomenclature c) writing chemical formulas d) bonding types Multiple Choice 1. What is the ionic charge on the copper ion in the ionic compound that has the formula CuSO4 ? a. b. c. d. +1 -2 +3 +2 2. Which element when combined with Bromine would most likely form an ionic compound? a. b. c. d. potassium carbon phosphorus iodine 3. An –ite or –ate ending on the name of a compound indicates that the compound _________________. a. b. c. d. is a binary ionic compound is a binary molecular compound contains a polyatomic ion contains a polyatomic cation 4. In general, ionic compounds _____________. a. b. c. d. are amorphous solids at room temperature conduct electricity when in the solid state conduct electricity when they are dissolved in water all of these Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin 5. Which of the following is the correct name for the compound MnF3? a. b. c. d. Manganese fluoride (III) Manganese (III) fluoride Manganese (I) fluoride (III) Manganese (III) fluoride (III) 6. According to the periodic table, Mg will most likely react with elements in which of these groups? a. b. c. d. 1 3 17 18 7. Which of these compounds is most likely to contain an ionic bond? a. b. c. d. H2 SO2 CH4 CaCl2 8. What is the name of NH4OH? a. b. c. d. Ammonium hydroxide Nitrogen oxygen hydride Nitrogen hydroxide Ammonium oxygen hydride 9. All of the following formulas are correct except: a. b. c. d. NaOH CaF2 H3PO4 AlSO4 10. In forming NaCl, energy is required to a. b. c, d. change chlorine to a gas add an electron to a chlorine atom remove an electron from a sodium atom bring together the sodium ions and chloride ions Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin Critical Thinking: 1. Why are there no rules for naming Group 18 ions? 2. Determine the ratios of cations to anions that are most likely in the formulas for ionic substances of the following elements: a. An alkali metal and a halogen b. An alkaline earth metal and a halogen c. An alkali metal and member of Group 16 d. An alkaline earth metal and a member of Group 16 3. Compound B has lower melting and boiling points than compound A does. At the same temperature, compound B vaporizes faster and to a greater extent than compound A. If only one of these compounds is ionic, which one would you expect it to be? Why? TOPIC: COVALENT COMPOUNDS AND BONDING CHEMISTRY STANDARD CH. 3 A, C, D The student will investigate and understand how conservation of energy and matter is expressed in chemical formulas and balanced equations. Key concepts include a) nomenclature c) writing chemical formulas d) bonding types Multiple Choice 1. Which of the following is the name of the molecule PCl3? a. b. c. d. Phosphorus trichloride Phosphorus chloride Potassium trichloride Potassium chloride 2. The type of formula that shows the arrangements of atoms and bonds is calleda. Empirical b. Chemical c. Molecular d. Structural Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin 3. What shape does the molecule BF3 have? a. Bent b. Linear c. Tetrahedral d. Trigonal planar 4. H H :Z–Z: H H The figure above shows a compound containing hydrogen (H) and an unknown element Z. To which group on the periodic table does element Z belong? 5. Bonding between two elements of equal electronegativity would bea. 100% covalent b. Primarily ionic c. 50% ionic d. Metallic in character 6. The small attractive force between molecules of nitrogen gas, N2, is due toa. b. c. d. Hydrogen bonds Van der Waals forces Dipole-dipole attraction Magnetism 7. Which of the following is a polar covalent bond? a. b. c. d. NaCl O2 Al SO2 8. In chemical compounds, covalent bonds form whena. b. c. d. The electronegativity difference between two atoms is very large Electrons are completely transferred between two metals Pairs of electrons are shared between two nonmetal atoms Two nonmetal ions are attracted to each other by opposite charges Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin 9. The Lewis electron dot system represents electrons in thea. b. c. d. Outer energy level Inner level Middle level Core level 10. An alien landed on Earth and created the periodic table shown. The astronaut was trying to determine what type of bond would be present in several compounds. The type of bond in a compound containing D and B would bea. b. c. d. A metallic bond A nonmetallic bond A covalent bond An ionic bond SOL Release Chemistry Test 2005 Critical Thinking 1. Ionic compounds tend to have higher boiling points than covalent substances do. Both ammonia, NH3, and methane, CH4, are covalent compounds, yet the boiling point of ammonia is 130 C higher than that of methane. What might account for the large difference? 2. Create a concept map using the following terms: valence electrons, non-polar, covalent compounds, polar, dipoles, and Lewis structures. 3. Create a Venn Diagram to compare and contrast ionic and covalent bonding. Make sure you have 5 similarities and 5 differences. Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin TOPIC: MOLE/EMPIRICAL AND MOLECULAR FORMULA CHEMISTRY STANDARD CH. 4 A, B The student will investigate and understand that quantities in a chemical reaction are based on molar relationships. Key concepts include a) Avogadro’s principle and molar volume b) Stoichiometric relationships Multiple Choice: 1. The molar mass (gram formula mass) for the compound sodium thiosulfate, Na2S2O3 is a. 71 grams b. 153 grams c. 158 grams d. 254 grams 2. The empirical formula for C6H12 is ---a. C3H6 b. C2H4 c. CH3 d. CH2 3. How many moles of copper are equivalent to 3.44 x 1023 atoms of copper? a. 0.571 moles b. 1.75 moles c. 5.41 x 1021 moles d. 5.71 x 1022 moles 4. What is the mass of one mole of CO2? a. b. c. d. 24 g 28 g 44 g 56 g Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin 5. Which of the following shows the correct number of atoms of each element in the formula Mg(NO3)2? a. b. c. d. 1 magnesium atom, 2 nitrogen atoms, and 6 oxygen atoms 1 magnesium atom, 2 nitrogen atoms, and 5 oxygen atoms 1 magnesium atom, 1 nitrogen atom, and 6 oxygen atoms 1 magnesium atom, 1 nitrogen atom, and 5 oxygen atoms Practice: 1. Formula Mass: Determine the molar mass of the following. Show all work and units. a. CuSO4 b. CaCO3 c. H3PO4 d. Al2(SO4)3 2. Determine the empirical formula for the following: a. 80.0 g of carbon and 20.0 g of oxygen b. 71.5% Ca and 28.5% O 3. Percent Composition. Determine the percent composition of the following compounds. Show all work. a. NO2 b. Al2(SO4)3 Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin THIRD Nine Weeks CHEMICAL REACTIONS SOL 3E STOICHIOMETRY SOL- 1G, 3B, 4B THERMOCHEMISTRY SOL- 3E, 4B, 5D, 5E, 5F CHEMICAL KINETICS SOL- 3F GAS LAWS SOL- 4A, 5A, 5B, 5C TOPIC: CHEMICAL REACTIONS Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin CHEMISTRY STANDARD CH. 3e The student will investigate and understand how conservation of energy and matter is expressed in chemical formulas and balanced equations. Key concepts include: e) reaction types Directions: Determine the reaction type of each of the following reactions. CO 1. Fe2O3 + 2. Cr + 3. Eu + HF 4. NH4Cl HCl + NH3 5. C12H22O11 + O2 6. Zn + 7. SiO2 + 8. Pb(NO3)2 + 9. KClO3 10. C6H6 + S8 FeO + CO2 Cr2S3 HCl C EuF3 + CO2 + ZnCl2 + Si + KCl + H2O H2 CO H3AsO4 O2 H2 PbHAsO4 + HNO3 O2 CO2 + H2O Directions: Complete the following table: Reaction Energy Description Exothermic Endothermic Sign of ΔH, Positive or Negative? Temperature, Increase or Decrease? Energy, Product or Reactant? TOPIC: STOICHIOMETRY Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin CHEMISTRY STANDARD CH. 1g The student will investigate and understand that experiments in which variables are measured, analyzed and evaluated produce observations and verifiable data. Key concepts include: g) mathematic manipulations (SI units, scientific notation, linear equations, graphing, ratio and proportion, significant digits, dimensional analysis) 1. Determine the number of significant figures in the following: a. 602 b. 1200 c. 0.00345 d. 0.1040 e. 34.08 f. 3 g. 0.970 2. Record your answer to the correct number of sig. figs. a. (7.502 x 102 )(5.43 x 104) b. 9.01 x 106 / 1.22 x 105 3. Write the following numbers in scientific notation, or translate the numbers to expanded notation. c. 9.32 x 10-5 d. 7.68 x 107 a. 556,000,000,000 b. 0.00000751 4. Use conversion factors to solve the following: SHOW ALL WORK!! a. Convert 250 m to km. b. Convert 18.50 kg to g. c. Convert 3.54 L to kL. d. Convert 0.5420 kg to mg 5. Convert the following temperatures to Kelvin or to degree’s Celsius. a. 100˚C b. 45˚C c. 273 K d. 400 K Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin CHEMISTRY STANDARD CH. 3B The student will investigate and understand how conservation of energy and matter is expressed in chemical formulas and balanced equations. Key concepts include b) balancing chemical equations. DIRECTIONS: Balance the following equations: CO 1. Fe2O3 + 2. Cr + S8 3. Eu + HF FeO + Cr2S3 EuF3 + 4. C12H22O11 + O2 5. Zn + 6. SiO2 + HCl C 7. Pb(NO3)2 + 8. NaCl + 9. KClO3 10.C6H6 + CO2 H2 CO2 + ZnCl2 + Si + Cl2 + KCl + O2 H2 CO H3AsO4 H2O H2O PbHAsO4 + H2 + HNO3 NaOH O2 CO2 + H2O TOPIC: STOICHIOMETRY Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin CHEMISTRY STANDARD CH. 4B The student will investigate and understand that quantities in a chemical reaction are based on molar relationships. Key concepts include: b) stoichiometric relationships. 1. Fill in the blanks. a. The amount of a substance equal to 6.02 x 1023 particles is called a ________________________. b. 6.02 x 1023 is also known as __________________________________. c. If you have 1 mole of copper, you have 6.02 x 1023 ________________________. 2. Molar (Formula) Mass. Determine the molar mass of the following and show all work and units. a. CuSO4 b. CaCO3 c. H3PO4 d. Al2(SO4)3 3. Mole Conversions. Solve the following problems. Show all work a. How many grams are there in 1.55 x 1023 molecules of CO2? b. How many atoms are there in 2.18 moles of nitrogen gas? c. What is the volume, at STP, occupied by 4.30 moles of oxygen gas? d. How many grams are there in 56.32 liters of carbon monoxide, at STP? e. How many moles are in 8.67 x 1025 atoms of sulfur? 4. Percent Composition. Determine the percent composition of the following compounds. Show all work. a. NO2 b. Al2(SO4)3 c. What is the percent of Mg in MgCl2? Solve the following problems and show all work. 5. Determine the empirical formula for 71.5% Ca and 28.5% O. 6. Determine the molecular formula for the empirical formula of C3H5O2, with a molar mass of 146 g/mol. 7. 4Al + 3O2 → 2Al2O3 , How many moles of oxygen are required to react completely with 0.84 mol of Al? How many moles of aluminum are needed to form 2.3 mol of Al2O3? 8. The combustion of acetylene gas is represented by this equation: 2C2H2(g) + 5O2(g) → 4CO2(g) + 2H2O(g) Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin a. How many grams of oxygen are required to burn 13.0g of C2H2 ? b. How many moles of C2H2 are needed to react completely with 98.0 g of water? 9. Calcium carbonate can be decomposed by heating. CaCO3(s) → CaO) (s) + CO2(g) What is the percent yield of this reaction if 24.8 g of CaCO3 is heated to give 13.1 g of CaO? 10. Copper reacts with sulfur to form copper (I) sulfide. 2Cu(s) + S(s) → Cu2S(s) What is the limiting reactant when 80.0 g of Cu reacts with 25.0 g of S? 11. How many moles of CS2 form when 2.7 mol of C reacts? 5C + 2SO2 → CS2 +4CO TOPIC: THERMOCHEMISTRY Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin CHEMISTRY STANDARD CH. 3E The student will investigate and understand how conservation of energy and matter is expressed in chemical formulas and balanced equations. Key concepts include: e) reaction types See Previous “Try It” in Chemical Reactions for CH.3E CHEMISTRY STANDARD CH. 4B The student will investigate and understand that quantities in a chemical reaction are based on molar relationships. Key concepts include: b) stoichiometric relationships See previous “Try It” in Stoichiometry for CH.4b Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin CHEMISTRY STANDARD CH. 5E The student will investigate and understand that the phases of matter are explained by kinetic theory and forces of attraction between particles. Key concepts include: e) molar heats of fusion and vaporization Solve the following problems. Show all work. 1. How much heat is released as a 75.0 g sample of ethanol gas at the boiling point condenses to a liquid? Heat of vaporization is 879 J/g. 2. The heat of vaporization of water is 540. cal/g. How many calories would be needed to convert 3 moles of water to vapor? Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin CHEMISTRY STANDARD CH. 5F The student will investigate and understand that the phases of matter are explained by kinetic theory and forces of attraction between particles. Key concepts include: f) specific heat capacity Solve the following problems. Show all work. 1. How much heat is lost as a 500. g cube of aluminum is cooled from 200°C to 25.0°C? The specific heat for aluminum is 0.897 J/g°C 2. How much heat is required to melt 550.0 g of Cu that has already been heated to its melting point? Heat of fusion is 205 J/g. 3. How many joules of heat are required to raise the temperature of 1.00 kg of water from 10.2˚C to 26.8˚C? 4. How much heat is released when 274 g of water cools from 85.2˚C to 38.4˚C? TOPIC: CHEMICAL KINETICS Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin CHEMISTRY STANDARD CH. 3F The student will investigate and understand how conservation of energy and matter is expressed in chemical formulas and balanced equations. Key concepts include: f) reaction rates, kinetics and equilibrium 1. Define the following terms: catalyst, activated complex, activation energy, heating curve, phase diagram, and heat capacity. 2. Draw labeled (reactants, products, activation energy, Δ H) reaction progress diagrams for endothermic and exothermic reactions. TOPIC: GAS LAWS Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin CHEMISTRY STANDARD CH. 4A The student will investigate and understand that quantities in a chemical reaction are based on molar relationships. Key concepts include: a) Avogadro’s principle and molar volume. Solve the following problems. Show all work. 1. What volume of oxygen gas at STP can be produced when 17.5 grams of potassium chlorate is decomposed? 2KClO3 → 2KCl + 3O2 2. Chlorine gas will react with 3.45 L of hydrogen gas to yield what mass of hydrogen chloride gas at STP? H2 + Cl2 → HCl Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin CHEMISTRY STANDARD CH. 5B The student will investigate and understand that quantities in a chemical reaction are based on molar relationships. Key concepts include: b) partial pressure and gas laws. Solve the following problems. Show all work. 1. A mixture of He, Ne, Ar and Xe has a total pressure of 2.5 atmospheres. Calculate the partial pressure of the He if the pressure due to Ne, Ar and Xe, respectively, is 0.3, 0.8 and 1.0 atmospheres. 2. A sample of carbon dioxide gas was collected over water at 24 °C. If the pressure of the sample was 758 mm Hg and the vapor pressure of water is 20. mm Hg, what is the pressure due to the carbon dioxide. Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin CHEMISTRY STANDARD CH. 5B The student will investigate and understand that quantities in a chemical reaction are based on molar relationships. Key concepts include: b) partial pressure and gas laws. 1. If 4.50 g of methane gas (CH4) is introduced into an evacuated 2.00 L container at 35°C, what is the pressure in the container, in atmospheres? 2. How many grams of iron would be produced from 52.5g of iron (III) oxide reacting with aluminum? Al + Fe2O3 → Al2O3 + Fe 3. In the reaction between calcium hydride and water, the theoretical yield of hydrogen gas from 75.0 g of calcium hydride is 7.18 g. In running this reaction the actual amount of hydrogen produced was 6.94 g. What is the percent yield of this reaction? CaH2 + H2O → Ca(OH)2 + H2 4. Hydrogen gas can be produced in the laboratory by the reaction of magnesium metal with hydrochloric acid. Mg + HCl → MgCl2 + H2 What is the limiting reactant, when 6.00 g of HCl is added to 5.00 g of Mg to product hydrogen gas? Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin CHEMISTRY STANDARD CH. 5B The student will investigate and understand that the phases of matter are explained by kinetic theory and forces of attraction between particles. Key concepts include: b) partial pressure and gas laws 1. A gas with a volume of 300 mL at 150°C is heated until its volume is 600 mL. What is the new temperature of the gas if the pressure is constant? 2. Calculate the volume of a gas in liters at 1 atm if its volume at 900 mmHg is 1500 mL. 3. A 3.50 L gas sample at 20°C and a pressure of 650 mmHg is allowed to expand to a volume of 8.00 L. What is the final temperature in degrees Celsius if the final pressure of the gas is 425 mm Hg? 4. A gas cylinder contains nitrogen gas at 10 atm and temperature of 20°C. The cylinder is left in the sun and the temperature of the gas increases to 50°C. What is the pressure in the cylinder? Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin CHEMISTRY STANDARD CH. 5C The student will investigate and understand that the phases of matter are explained by kinetic theory and forces of attraction between particles. Key concepts include: c) vapor pressure. 1. How is the normal boiling point of a substance related to its vapor pressure? 2. The atmospheric pressure of a mountain in South America is 400 mm Hg. Using the vapor pressure curve of water from your textbook (page 401), determine the boiling point of water at this altitude. 3. Sketch a graph of the vapor pressure of a liquid. Show the dependency of pressure on temperature. Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin CHEMISTRY STANDARD CH. 5D The student will investigate and understand that the phases of matter are explained by kinetic theory and forces of attraction between particles. Key concepts include: d) phase changes. Directions: Fill in the blanks with the correct response. The particles in a ______________ phase are very close together in an orderly, fixed and usually crystalline arrangement. _______________ is an endothermic change of state in which a solid becomes a liquid. The temperature and pressure at which a solid becomes a liquid is its ____________________. Because particles in the ____________ phase have enough kinetic energy to be able to move past each other easily, they take the shape of their container. While many liquids flow readily, many are resistant to flow, or are ________. Because they are held close together, liquid particles are more affected by forces between particles. They have attraction for each other, or ____________, as well as attraction for particles of solid surfaces, called _______________. Liquids tend to form spherical drops because of ____________, or the tendency to decrease their surface area to the smallest possible size. Particles in a liquid can gain enough kinetic energy to leave the surface and become a gas in a process called ______________. Attractive forces between ____________ particles do not have a great effect, which makes the particles independent of each other. The temperature and pressure at which the number of liquid particles become gas particles is the same as the number of gas particles returning to the liquid phase is called a substance’s _______________. Gas particles lose energy and become liquid during ________________. The process during which a liquid loses energy and becomes a sold is called ____________. The temperature at which this change occurs is the _______________ of the substance. When particles of solid become gas particles without first melting, the substance undergoes a process called _____________. The reverse of this process, when a gas becomes a solid directly, is called ________________. Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin TOPIC: ORGANIC CHEMISTRY AND BIOCHEMISTRY CHEMISTRY STANDARD CH. 6A, B The student will investigate and understand how basic chemical properties relate to organic chemistry and biochemistry. Key concepts include: a) Unique properties of carbon that allow multi-carbon compounds b) Uses in pharmaceuticals and genetics, petrochemicals, plastics and food DIRECTIONS: Answer the questions below. 1. List the 3 factors that make the bonding of carbon atoms unique. 2. Draw Lewis structures, identify shape, and describe polarities of the following molecules: CH4, C2H6, CH3CH2OH, C6H6 3. List 6 natural, biological polymers. 4. List 6 common pharmaceuticals that are organic compounds. 5. List 6 man-made, synthetic polymers. Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin FOURTH Nine Weeks SOLUTIONS SOL- 3F, 4D, 5G ACIDS AND BASES SOL- 4D Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin TOPIC: SOLUTIONS CHEMISTRY STANDARD CH. 4C The student will investigate and understand that quantities in a chemical reaction are based on molar relationships. Key concepts include: c) solution concentrations. DIRECTIONS: Solve the following problems. Show all work. 1. What is the molality of a solution in which 3.0 moles of NaCl is dissolved in 1.5 kg of water? 2. How many grams of I2 should be added to 750 g of CCl4 to prepare a 0.020 m solution? 3. Calculate the molarity of 5.85 g of NaCl in 2.00 L of solution. 4. How many moles are needed to make 2.0 L of 0.30M solution of Na2SO4? 5. How would you correctly prepare 125 mL of a 0.30M solution of copper(II) sulfate (CuSO4) from a 2.00M solution of CuSO4? 6. What is the boiling point of a solution that contains 1.25 mol of CaCl2 in 1400 g of water? Kb for water = 0.512 °C/m 7. What is the freezing point of Na2SO4 in 1750 g of water? Kf of water = 1.86 °C/m Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin CHEMISTRY STANDARD CH. 3F The student will investigate and understand that quantities in a chemical reaction are based on molar relationships. Key concepts include: f) reaction rates, kinetics and equilibrium DIRECTIONS: Answer the questions below. 1. What is a reversible chemical reaction? How does the format of the arrow differ from a reaction that goes to completion? 2. What is Le Châtelier’s principle? How would the equilibrium in the following reaction be impacted by: CO (g) + O2 (g) ↔ CO2 (g) a. increasing [CO2]? b. decreasing pressure? c. decreasing [CO]? Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin TOPIC: ACIDS, BASES, and SOLUTIONS CHEMISTRY STANDARD CH. 4D The student will investigate and understand that quantities in a chemical reaction are based on molar relationships. Key concepts include: d) acid/base theory: strong electrolytes, weak electrolytes, and nonelectrolytes; dissociation and ionization, pH and pOH; and the titration process. DIRECTIONS: Answer the questions below. 1. Define the following terms: conjugate acid, acid, base, conjugate base, Bronsted-Lowry acid, Bronsted-Lowry base, indicator, neutralization and pH scale. 2. Solve the following problems. a. Calculate the pH and pOH for each solution: [H+] = 5.0 x 10-6 and [OH-] = 4.5 x 10-11 b. Calculate the [H+] for each solution: pH = 5.0 and pH = 12.20 3. List the properties of an acid and a base. 4. Identify the conjugate acid and conjugate base in the following equation. NH3 + H2O ↔ NH4+ + OH- Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin CHEMISTRY STANDARD CH. 5G The student will investigate and understand that the phases of matter are explained by kinetic theory and forces of attraction between particles. Key concepts include: g) colligative properties. DIRECTIONS: Answer the questions below. 1. What is the freezing point of a solution of 200 grams of ethylene glycol (HOCH2CH2OH) mixed with 400 grams of water? 2. What is the boiling point elevation of a solution of 26.8 grams of NaCl dissolved in 500 grams of water? 3. Calculate the boiling point and freezing point of a solution of 220 g of sucrose (C12H22O11) dissolved in 1.0 kg of water. How does this compare to the boiling point and freezing point of pure water? Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin ANSWERS FIRST NINE WEEKS STANDARD CH. 1 F 1. Experimental Value = 1.44 g Accepted Value = 1.54 g 6.5% 2. Experimental Value = 22.2 L Accepted Value = 22.4 L 0.89% 3. Experimental Value = 128.6 mg Accepted Value =128.3 mg 0.23% 4. Experimental Value = 9.98 x 10-3 g Accepted Value =1.03 x 10-2 g 3.11% STANDARD CH. 2H Chemical and Physical Properties The student will be given an object or substance and must identify the physical and chemical properties of that object. (This assignment is to be given after lecture on chemical and physical properties and how to correctly identify each.) 1. Describe the chemical properties of this object below, citing any evidence you have collected: Answers will vary. 2. Describe the physical properties of this object below, citing any evidence you have Answers will vary. collected: STANDARD CH. 2I .Students will create a two page foldable for the models of an atom. Each outside flap will be the names of the scientists (Democritus, Dalton, Rutherford, Bohr), and the inside picture will be each scientist’s contribution to the historical model. A one page foldable will be created for the quantum models with Heisenberg and Planck. Students will be creative using mnemonics and/or graphics. Answers will vary. Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin STANDARD CH. 2A Element Ne Al Ge P Br Ca K I Sr Sn Average Atomic mass 20.179 26.9812 72.59 30.97376 79.904 40.08 39.0983 126.904 87.62 118.69 Mass number 20 27 72 31 80 40 39 127 88 119 Atomic number 10 13 32 15 35 20 19 53 38 50 Atomic mass answers may slightly vary depending on the Periodic Table used. STANDARD CH. 2C Warm-up: Complete the following chart: Subatomic Particle Symbol Charge Relative mass Location in atom protons P + 1 In nucleus neutrons N 0 1 In nucleus electrons E - 0 Outside nucleus or electron cloud Atomic Structure Element/Ion Atomic Number Atomic Mass Mass Number Protons Neutrons Electrons 3 H 1 1.00797 3 1 2 1 1 H+ 1 1.00797 1 1 0 0 12 C 6 12.0115 12 6 6 6 7 Li+ 3 6.941 7 3 4 2 17 35.453 35 17 18 18 39 K 19 39.0983 39 19 20 19 Mg2+ 12 24.305 24 12 12 10 33 74.922 77 33 44 36 47 107.868 108 47 61 47 35 Cl- 24 77 As3- 108 Ag Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin 30 -2 16 32.064 30 16 14 18 238 92 238.03 238 92 146 92 S U STANDARD CH. 2B DIRECTIONS: Answer the following questions Average Atomic Mass 1. Rubidium has two common isotopes, Rb-85 and Rb-87. If the abundance of Rb-85 is and the abundance of Rb-87 is 27.8%, what is the average atomic mass of rubidium? 85.556 Uranium has 3 common isotopes. If the abundance of U-234 is 0.01%, the abundance of U-235 is 0.71% and the abundance of U-238 is the average atomic mass? 237.9783 72.2% 2. 3. Titanium has five common isotopes: Ti-46 (8.0%), Ti-47 (7.8%), Ti-48 (73.4%), Ti-49 and Ti-50 (5.3%). What is the average atomic mass of titanium? 47.923 99.28%, what is (5.5%), 4. Explain why atoms have different isotopes. In other words, how is it that helium can have three different sized atoms and they all are still the element Helium. Atoms of the same element will have the same number of protons but different numbers of neutrons. 5. Draw and label 3 possible isotopes of hydrogen. Which isotope can you predict to be most abundant? Why? Correctly name the 3 isotopes using the symbol “H”? Answers will vary. the 6. Select an element on the PTOE. Create 4 isotopes for your element and form an average atomic mass calculation to support the given atomic mass. Answers will vary. STANDARD CH. 3A Writing Binary Formulas Cation Anion Formula Cation Anion Formula Na+ Cl- NaCl Fe+2 O-2 FeO Ba+2 F- BaF2 Fe+3 O-2 Fe2O3 K+ S-2 K2S Cr+2 S-2 CrS Li+ Br- LiBr Cr+3 S-2 Cr2S3 Al+3 I- AlI3 Cu+ Cl- CuCl Zn+2 S-2 ZnS Cu+2 Cl- CuCl2 Ag+ O-2 Ag2O Pb+2 O-2 PbO Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin Mg+2 P-3 Mg3P2 Pb+4 O-2 PbO2 Ni+2 O-2 NiO Mn+2 Br- MnBr2 Ni+3 O-2 Ni2O3 Mn+4 Br- MnBr4 STANDARD CH. 2 D, E Counting Valence Electrons Element Ne Al Ge P Br Ca K I Sr Sn Group Number 18 13 14 15 17 2 1 17 2 14 Number of Valence Electrons 8 3 4 5 7 2 1 7 2 4 STANDARD CH. 2 F DIRECTIONS: State the periodic trend and explain. ATOMIC RADII LEFT TO RIGHT DECREASE ELECTRONEGATIVITY INCREASE SHIELDING EFFECT SAME IONIZATION ENERGY INCREASE PERIODIC TRENDS EXPLAINATION STRONGER NUCLEAR ATTRACTION MORE VALNCE ELECTRONS SAME ENERGY LEVEL MORE VALENCE ELECTRONS TOP TO BOTTOM INCREASE EXPLAINATION DECREASE LARGER RADIUS INCREASE MORE ENERGY LEVELS MORE SHIELDING DECREASE MORE ENERGY LEVELS STANDARD CH. 2 C Ions and Subatomic Particles Ion Symbol Protons Electrons Charge 2- S 16 18 -2 K1+ 19 18 +1 Ba2+ 56 54 +2 Fe3+ 26 23 +3 Fe2+ 26 24 +2 Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin F1- 9 10 -1 O2- 8 10 -2 P3- 15 18 -3 Sn4+ 50 46 +4 Sn2+ 50 48 +2 N3- 7 10 -3 Br1- 35 36 -1 Mg2+ 12 10 +2 Cu1+ 29 28 +1 Cu2+ 29 27 +2 U6+ 92 86 +6 Mn5+ 25 20 +5 Cl1- 17 18 -1 Se2- 34 36 -2 SECOND NINE WEEKS SOL 3C 1. D 2.A 3.C 4. C 5. B 6. C 7. D 8. A 9. D 10. D SOL 3 A,C,D 1. A 2. D 3. B 4. Group 15 5. A 6. B 7. D 8. C 9. A SOL 4A 1. C 2. C 3. A 4. C 5. A THIRD NINE WEEKS CHEMISTRY STANDARD CH. 3e 11. CO Fe2O3 + S8 12. Cr + 13. Eu + HF 14. NH4Cl 15. C12H22O11 + 16. Zn + 17. SiO2 + FeO + CO2 double replacement Cr2S3 synthesis EuF3 + HCl + NH3 decomposition O2 HCl C H2 single replacement CO2 + ZnCl2 + Si + H2O combustion H2 single replacement CO single replacement Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin 18. Pb(NO3)2 + 19. KClO3 20. C6H6 + H3AsO4 KCl + O2 PbHAsO4 + O2 CO2 + HNO3 double replacement decomposition H2O combustion Directions: Complete the following table: Reaction Energy Description Sign of ΔH, Positive or Negative? Exothermic Endothermic + Temperature, Increase or Decrease? increase decrease Energy, Product or Reactant? product reactant STANDARD CH. 1g 6. Determine the number of significant figures in the following: c. d. 7. 602 1200 3 2 c. 0.00345 3 d. 0.1040 4 e. 34.08 f. 3 4 1 g. 0.970 3 Record your answer to the correct number of sig. figs. c. d. 8. (7.502 x 102 )(5.43 x 104) 9.01 x 106 / 1.22 x 105 4.07 x 107 7.39 x 101 Write the following numbers in scientific notation, or translate the numbers to regular notation. c. d. 9. 556,000,000,000 5.56 x 1011 0.00000751 7.51 x 10-6 c. 9.32 x 10-5 0.0000932 d. 7.68 x 107 76,800,000 Use conversion factors to solve the following: SHOW ALL WORK!! c. d. Convert 250 m to km. 0.25 km Convert 18.50 kg to g. 1.85 x 104 c. Convert 3.54 L to kL. 3.54 x 10 -3 d. Convert 0.5420 kg to mg 5.42 x 10 5 10. Convert the following temperatures to Kelvin or to degree’s Celsius. b. 100˚C 373K b. 45˚C 318K c. 273 K 0°C d. 400 K 127°C CH. 3B 21. Fe2O3 + CO 2FeO + 22. 16Cr + 3S8 8Cr2S3 23. 2Eu + 6HF 2EuF3 + CO2 3H2 Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin 24. C12H22O11 + 12 O2 25. Zn + 26. SiO2 + 27. Pb(NO3)2 + 28. 2NaCl + 2 H2O 29. 2 KClO3 30. 2C6H6 + 2 HCl 2C 12 CO2 + 11H2O ZnCl2 + Si + 2CO H3AsO4 2KCl + 15O2 H2 PbHAsO4 + 2 HNO3 Cl2 + H2 + 2NaOH 3O2 12CO2 + 6 H2O STANDARD CH. 4B 12. Fill in the blanks. d. The amount of a substance equal to 6.02 x 1023 particles is called a _mole_______________________. e. 6.02 x 1023 is also known as ____Avogadro’s number_____________. f. If you have 1 mole of copper, you have 6.02 x 10 23 __atoms__________. 13. Molar (Formula) Mass. Determine the molar mass of the following and show all work and units. b. CuSO4 159.58g/mol b. CaCO3 100.06g/mol c. H3PO4 97.96g/mol d. Al2(SO4)3 342.05g/mol 14. Mole Conversions. Solve the following problems. Show all work f. How many grams are there in 1.55 x 1023 molecules of CO2? 11.3g g. How many atoms are there in 2.18 moles of nitrogen gas? 1.31 x 1024atoms h. What is the volume, at STP, occupied by 4.30 moles of oxygen gas?96.3L i. j. How many grams are there in 56.32 liters of carbon monoxide, at STP?70.40g How many moles are in 8.67 x 1025 atoms of sulfur?144 moles 15. Percent Composition. Determine the percent composition of the following compounds. Show all work. b. NO2 30.5%N, 69.5%O b. Al2(SO4)3 15.8%Al, 28.1%S, 56.1%O c. What is the percent of Mg in MgCl2?25.5%Mg Solve the following problems and show all work. 16. Determine the empirical formula for 71.5% Ca and 28.5% O. CaO 17. Determine the molecular formula for the empirical formula of C3H5O2, with a molar mass of 146 g/mol. C6H10O4 18. 4Al + 3O2 → 2Al2O3 , How many moles of oxygen are required to react completely with 0.84 mol of Al? .63 mol O2 How many moles of aluminum are needed to form 2.3 mol of Al 2O3?4.6 mol Al 19. The combustion of acetylene gas is represented by this equation: 2C2H2(g) + 5O2(g) → 4CO2(g) + 2H2O(g) a. How many grams of oxygen are required to burn 13.0g of C 2H2 ? 39.9 g O2 b. How many moles of C2H2 are needed to react completely with 98.0 g of water? 5.44 mol C2H2 20. Calcium carbonate can be decomposed by heating. CaCO3(s) → CaO(s) + CO2(g) What is the percent yield of this reaction if 24.8 g of CaCO 3 is heated to give 13.1 g of CaO?94.2% 21. Copper reacts with sulfur to form copper (I) sulfide. 2Cu(s) + S(s) → Cu2S(s) Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin What is the limiting reactant when 80.0 g of Cu reacts with 25.0 g of S?LR=Cu 22. How many moles of CS2 form when 2.7 mol of C reacts? 5C + 2SO2 → CS2 +4CO 0.54 mol CS2 CHEMISTRY STANDARD CH. 3E See Previous “Try It” in Chemical Reactions for CH.3E STANDARD CH. 4B See previous “Try It” in Stoichiometry for CH.4b STANDARD CH. 5D Solve the following problems. Show all work. 3. How much heat is released as a 75.0 g sample of ethanol gas at the boiling point condenses to a liquid? Heat of vaporization is 879 J/g.6.59 x 104J 4. The heat of vaporization of water is 540. cal/g. How many calories would be needed to convert 3 moles of water to vapor?3 x 104 cal STANDARD CH. 5E Solve the following problems. Show all work. 5. How much heat is lost as a 500. g cube of aluminum is cooled from 200°C to 25.0°C? The specific heat for aluminum is 0.897 J/g°C 8 x 104 J 6. How much heat is required to melt 550.0 g of Cu that has already been heated to its melting point? Heat of fusion is 205 J/g. 1.128 x 105 J 7. How many joules of heat are required to raise the temperature of 1.00 kg of water from 10.2˚C to 26.8˚C? 6.95 x 104 J 8. How much heat is released when 274 g of water cools from 85.2˚C to 38.4˚C? -5.37 x 104 J STANDARD CH. 3F 1. Define the following terms: catalyst, activated complex, activation energy, heating curve, phase diagram, and heat capacity. a. Catalyst – a substance that changes the rate of a chemical reaction without being consumed. Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin b. c. d. e. f. 2. Activated complex – molecule in an unstable state intermolecular to the reactant and products in the chemical reaction. Activation energy – the minimum amount of energy required to start a chemical reaction. Heating curve – shows the change in temperature of a substance as heat is added and the substance undergoes phase changes. Phase diagram – graph of the relationship between the physical state of a substance, temperature, and pressure of the substance. Heat capacity – energy needed to increase the temperature of 1 mole (1 gram) of the substance by 1°C. Draw labeled (reactants, products, activation energy, Δ H) reaction progress diagrams for endothermic and exothermic reactions. STANDARD CH. 4A Solve the following problems. Show all work. 1. What volume of oxygen gas at STP can be produced when 17.5 grams of potassium chlorate is decomposed? 2KClO3 → 2KCl + 3O2 4.80L O2 2. Chlorine gas will react with 3.45 L of hydrogen gas to yield what mass of hydrogen chloride gas at STP? H2 + Cl2 → HCl 11.2g HCl STANDARD CH. 4C Solve the following problems. Show all work. 1. A mixture of He, Ne, Ar and Xe has a total pressure of 2.5 atmospheres. Calculate the partial pressure of the He if the pressure due to Ne, Ar and Xe, respectively, is 0.3, 0.8 and 1.0 atmospheres. 0.4 atm 2. A sample of carbon dioxide gas was collected over water at 24 °C. If the pressure of the sample was 758 mm Hg and the vapor pressure of water is 20. mm Hg, what is the pressure due to the carbon dioxide. 738 mmHg Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin STANDARD CH. 4D 1. If 4.50 g of methane gas (CH4) is introduced into an evacuated 2.00 L container at 35°C, what is the pressure in the container, in atmospheres? 3.5 atm 2. How many grams of iron would be produced from 52.5g of iron (III) oxide reacting with aluminum? Al + Fe2O3 → Al2O3 + Fe 36.7 g 3. In the reaction between calcium hydride and water, the theoretical yield of hydrogen gas from 75.0 g of calcium hydride is 7.18 g. In running this reaction the actual amount of hydrogen produced was 6.94 g. What is the percent yield of this reaction? CaH2 + H2O → Ca(OH)2 + H2 96.4% 4. Hydrogen gas can be produced in the laboratory by the reaction of magnesium metal with hydrochloric acid. Mg + HCl → MgCl2 + H2 What is the limiting reactant, when 6.00 g of HCl is added to 5.00 g of Mg to product hydrogen gas? LR=HCl STANDARD CH. 5A 1. A gas with a volume of 300 mL at 150°C is heated until its volume is 600 mL. What is the new temperature of the gas if the pressure is constant? 800K 2. Calculate the volume of a gas in liters at 1 atm if its volume at 900 mmHg is 1500 mL. 2L A 3.50 L gas sample at 20°C and a pressure of 650 mmHg is allowed to expand to a volume of 8.00 L. What is the final temperature in degrees Celsius if the final pressure of the gas is 425 mm Hg? 200°C 3. 4. A gas cylinder contains nitrogen gas at 10 atm and temperature of 20°C. The cylinder is left in the sun and the temperature of the gas increases to 50°C. What is the pressure in the cylinder? 10 atm STANDARD CH. 5B 1. 2. 3. How is the normal boiling point of a substance related to its vapor pressure? When you increase the temperature of a system to the point at which the vapor pressure of a substance is equal to the standard atmospheric pressure, you have reached the substance’s normal boiling point. The atmospheric pressure of a mountain in South America is 400 mm Hg. Using the vapor pressure cure of water from your textbook (page 401), determine the boiling point of water at this altitude. 84°C Sketch a graph of the vapor pressure of a liquid. Show the dependency of pressure on temperature. Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin STANDARD CH. 5C Directions: Fill in the blanks with the correct response. The particles in a __solid________ phase are very close together in an orderly, fixed and usually crystalline arrangement. __Melting_______ is an endothermic change of state in which a solid becomes a liquid. The temperature and pressure at which a solid becomes a liquid is its _heat of fusion___. Because particles in the __liquid______ phase have enough kinetic energy to be able to move past each other easily, they take the shape of their container. While many liquids flow readily, many are resistant to flow, or are viscous_____. Because they are held close together, liquid particles are more affected by forces between particles. They have attraction for each other, or cohesion__, as well as attraction for particles of solid surfaces, called adhesion_. Liquids tend to form spherical drops because of surface tension_, or the tendency to decrease their surface area to the smallest possible size. Particles in a liquid can gain enough kinetic energy to leave the surface and become a gas in a process called _vaporization______. Attractive forces between __gas____ particles do not have a great effect, which makes the particles independent of each other. The temperature and pressure at which the number of liquid particles become gas particles is the same as the number of gas particles returning to the liquid phase is called a substance’s __heat of vaporization___. Gas particles lose energy and become liquid during _condensation____. The process during which a liquid loses energy and becomes a sold is called freezing. The temperature at which this change occurs is the _freezing point______ of the substance. When particles of solid become gas particles without first melting, the substance undergoes a process called sublimation___. The reverse of this process, when a gas becomes a solid directly, is called deposition____. STANDARD CH.6A, B 1. List the 3 factors that make the bonding of carbon atoms unique. Strong single bond, not extremely reactive, different bonding arrangements Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin 2. Draw Lewis structures, identify shape, and describe polarities of the following molecules: CH 4, C2H6, CH3CH2OH, C6H6 CH4 –tetrahedron, non-polar C2H6,- tetrahedron at each carbon, non-polar CH3CH2OH – tetrahedron at each carbon, polar due to the oxygen C6H6 – hexagonal, planar, non-polar 3. List 6 natural, biological polymers. DNA, RNA, starch, cellulose, protein, poly-lipids, etc. 4. List 6 common pharmaceuticals that are organic compounds. aspirin, vitamins, insulin, Tylenol, ibuprofen, Ritalin, etc. 5. List 6 man-made, synthetic polymers. Nylon, polyethylene, polypropylene, PVC, polystyrene, Kevlar, etc. FOURTH NINE WEEKS STANDARD CH. 4E 1. What is the molality of a solution in which 3.0 moles of NaCl is dissolved in 1.5 kg of water? 2.0 m 2. 3. How many grams of I2 should be added to 750 g of CCl4 to prepare a 0.020 m solution? 3.8 g Calculate the molarity of 5.85 g of NaCl in 2.00 L of solution. 0.0500 M 4. How many moles are needed to make 2.0 L of 0.30M solution of Na 2SO4? 0.60 mol 5. How would you correctly prepare 125 mL of a 0.30M solution of copper(II) sulfate (CuSO 4) from a 2.00M solution of CuSO4? 19 mL 6. What is the boiling point of a solution that contains 1.25 mol of CaCl 2 in 1400 g of water? Kb for water = 0.512 °C/m 101.37°C = 101°C 7. What is the freezing point of Na2SO4 in 1750 g of water? Kf of water = 1.86 °C/m -3.19°C STANDARD CH. 4F DIRECTIONS: Answer the questions below. 3. What is a reversible chemical reaction? A chemical reaction in which the products reform the original reactants. How does the format of the arrow differ from a reaction that goes to completion? Arrows that point in opposite directions. 4. What is Le Châtelier’s principle? States that a system in equilibrium will oppose a change in a way that helps eliminate the change. How would the equilibrium in the following reaction be impacted by: CO (g) + O2 (g) ↔ CO2 (g) d. increasing [CO2]? Shift to the left e. decreasing pressure? Shift to the right f. decreasing [CO]? Shift to the left Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin STANDARD CH. 4G DIRECTIONS: Answer the questions below. 1. Define the following terms: conjugate acid, acid, base, conjugate base, Bronsted-Lowry acid, Bronsted-Lowry base, indicator, neutralization and pH scale. Conjugate acid- compound with a hydrogen ion Acid – donates a proton Base – accepts a proton Conjugate base – same compound of the acid but without the hydrogen ion Bronsted-Lowery acid – substance that is capable of donating a proton Bronsted-Lowery base – substance that is capable of accepting a proton Indicator – dyes that can be added that will change color in the presence of an acid or base Neutralization – adding an acid to a base to produce a salt and water pH scale – a way of expressing the strength of acids and bases 2. Solve the following problems. a. Calculate the pH for each solution: [H+] = 5.0 x 10-6 pH = 5.3and [OH-] = 4.5 x 10-11 pH = 3.7 b. Calculate the [H+] for each solution: pH = 5.0 [H] = 1x10-5 pH = 12.20 [H] = 6.3 x 10-13 1. List the properties of an acid and a base. Acid: produces hydrogen ions, taste sour, corrodes metal, electrolyte, pH<7, reacts with a base, litmus paper turns red Base: produces hydroxide ions, taste bitter, electrolytes, reacts with acids, pH>7, litmus paper turns blue 4. Identify the conjugate acid and conjugate base in the following equation. NH3 + H2O ↔ NH4+ + OHConjugate acid: NH4+ conjugate base: NH3 Conjugate acid: H2O, conjugate base: OH- STANDARD CH. 5F DIRECTIONS: Answer the questions below. 1. What is the freezing point of a solution of 200 grams of ethylene glycol (HOCH 2CH2OH) mixed with 400 grams of water? -14.9°C 2. What is the boiling point elevation of a solution of 26.8 grams of NaCl dissolved in 500 grams of water? 100.9°C 3. Calculate the boiling point and freezing point of a solution of 220 g of sucrose (C 12H22O11) dissolved in 1.0 kg of water. Tb = 100.3°C Tf = -1.2°C How does this compare to the boiling point and freezing point of pure water? Both are close to pure water. CHEMISTRY STANDARD CH. 4G See previous “Try It” in Solutions for CH.4G Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin REFERENCES www.222.nano.gov www.chemmybear.com www.chemistry.merlot.org Holt Chemistry Textbook http://my.hrw.com/ www.nsta.org Prentice Hall Textbook http://www.phschool.com/home.html Virginia Department of Education Blueprint http://www.doe.virginia.gov/testing/sol/blueprints/science_blueprints/blueprint_ch emistry.pdf Virginia Department of Education Framework http://www.doe.virginia.gov/testing/sol/frameworks/science_framewks/framework _science-chem.pdf Virginia Department of Education Scope and Sequence http://www.doe.virginia.gov/testing/sol/scope_sequence/science_scope_sequence/s copeseq_science_chemistry.pdf Virginia Department of Education Standards of Learning http://www.doe.virginia.gov/testing/sol/standards_docs/science/courses/stds_chemi stry.pdf www.scitoys.com www.stevespanglescience.com Designed by Ann-Rene Challenger, Amanda Griffin, Dr. Donna Keene, Marianne Lawrence, Deborah Oliver, and Ashanta Ruffin