Grayslake Central High School
Honors Chemistry
Final Exam Review Packet
Semester 1
Bradley and Piggott
11
Unit 1 – Safety & Equipment Knowledge Objectives: Students will know… 1. How to identify the location of all safety equipment within the classroom. 2. The proper procedures for dealing with accidents in the lab. Skill Objectives: Students will be skilled at… 1. Following all lab safety rules and demonstrating proper lab techniques. Unit 2 – Matter Knowledge Objectives: Students will know… 1. About the composition of matter. 2. Several techniques for separating mixtures and their usefulness. 3. How to design a lab procedure to separate several substances based on physical properties. Skill Objectives: Students will be skilled at… 2. Successfully lighting, adjusting and extinguishing a Bunsen burner. 3. Performing a lab to separate various components within a mixture. 4. Differentiating between chemical and physical properties. 5. Differentiating between chemical and physical changes. 6. Differentiating between mixtures (homogeneous and heterogeneous) and pure substances. 7. Identifying various pieces of scientific equipment and discussing their function in a science laboratory. Important Vocabulary: Atom Filtration Physical Change Centrifuge Filtrate Physical Property Chemical Change Gas Pure Substance Chemical Property Heterogeneous Mixture Qualitative Chromatography Homogeneous Mixture Quantitative Compound Inference Residue Decanting Liquid Separation Density Magnetism Skimming Distillation Mass Solid Element Matter Solution Evaporation Observation 1. Distinguish between qualitative and quantitative observations. GIVE AN EXAMPLE OF EACH! 2. What is matter? 3. How is a compound different from a mixture? 4. What is the difference between an element and a compound? 5. What is constant composition? 6. What is the difference between a homogeneous mixture, heterogeneous mixture, element and compound? Give an example of each. 7. Label the following as either a homogeneous mixture, heterogeneous mixture, element or compound: a. Oxygen gas ___________________________________ b. Carbon dioxide gas ___________________________________ c. Peanut butter and jelly sandwiches ___________________________________ d. Chicken broth ___________________________________ e. Silver nitrate ___________________________________ f.
___________________________________ g. Muddy water ___________________________________ h. Sprite ____________________________________ i.
Platinum ____________________________________ j.
14 K gold ____________________________________ Salt water 8. What is the difference between a chemical property and a physical property? 9. Of the following circle the chemical properties and underline the physical properties: a. Flammable f.
b. Yellow in color g. Attacked by sulfuric acid c. Smells like rotten eggs h. Has density of 2.3 g/ml d. Powdery i.
Not soluble in water e. Melts at 1234 K j.
Inert with nitrogen
Boils at 3456 K 10. What is the difference between a chemical change and a physical change? 3
11. Of the following situations, circle the chemical changes: a. Dying a pair of socks pink in the wash f.
Bread rising b. An egg hard boiling. g. Acid eating through paper c. An fresh egg rotting h. Chewing on paper d. A car rusting i.
Dissolving kool aid in water e. Cutting wood j.
Shaking and opening a can of pop 12. Are all physical changes accompanied by chemical changes? Are all chemical changes accompanied by physical changes? Explain 13. Complete the table below. Separation Property Used to Separate Technique Decanting Explanation Filtration Chromatography Distillation Evaporation Magnetism Skimming 14. What is the key difference between distillation and evaporation? 4
Unit 3 – Measurement Knowledge Objectives: Students will know… 1. Several SI units of measurement. 2. How to calculate the density of an object. 3. How uncertainty in measurement is related to significant figures. 4. How to determine whether a set of measurements are accurate and/or precise. 5. The difference between qualitative and quantitative observations. Skill Objectives: Students will be skilled at… 3. Measuring temperature in Celsius. 4. Measuring an object using the correct number of significant figures. 5. Using dimensional analysis to convert from one unit to another. 6. Converting ordinary numbers into scientific notation. 7. Converting numbers written in scientific notation into ordinary numbers. 8. Converting temperatures from Celsius to Kelvin and from Kelvin to Celsius. 9. Using density in a dimensional analysis conversion. 10. Rounding measurements and answers to dimensional analysis conversions to the correct number of significant figures. Important Vocabulary: International System of Measurement (SI) Absolute Zero Scientific Notation Kelvin Accuracy Significant Figure Mass Celsius Uncertainty Metric System Conversion Factor Units Ordinary Number (Positional Notation) Density Volume Percent Error Dimensional Analysis Precision Equivalence Statement 15.
16.
17.
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20.
What is dimensional analysis? What is a conversion factor? Give an example. 1 cm3 = ______ mL 1 L = ______________ mL 1 km = ________________ m Give the symbol and meaning for each of the following prefixes: a. Kilo c. Milli b. Centi d. Micro 5
21. Complete the following conversions: a. Convert 543,330 mg to lbs b. 12.4 miles to meters c. 40.0 lb to centigrams d. 67.2 km/hr to m/s Use the following information for questions 22 & 23. Given the following: 2 whims = 5 logs 2 logs = 12 whips 9 whips = 4 dags 7 dags = 1 pin 3 pins = 6 horns 22. How many whips are in 39 dags? (Ans: 88 whips) 23. How many pins in 17 whims? (Ans: 16 pins) 6
24. Why are units important in a measurement? 25. In the S.I. system, what are the fundamental units for a. Mass ____________________________________ b. Length ____________________________________ c. amount of matter ____________________________________ d. time ____________________________________ e. volume ____________________________________ f. temperature ____________________________________ 26. What is scientific notation? 27. Why do we use scientific notation? 28. Put the following ordinary numbers in scientific notation: a. 23,340,000 _______________________________ b. 0.00000005430 _______________________________ c. 0.0020002 _______________________________ d. 230.0 _______________________________ 29. Put the following in normal decimal notation: a. 3.450 x 103 _______________________________ b. 9.9900 x 10‐4 _______________________________ c. 6.33 x 10‐7 _______________________________ d. 7.880 x 102 _______________________________ 30. If a sample has a mass of 245.8 grams and takes up 94 mL of space, what is the density of the sample? 7
31. What is the mass in kg of 0.0516 L sample of gasoline which has a density of 0.70 g/cm3? (NOTE: 1 mL = 1 cm3) 32. Summarize the rules for deciding whether or not a particular digit in a number used in a calculation is “significant.” 33. Summarize the rules for rounding off numbers. 34. Summarize the rules for doing arithmetic with the correct number of significant figures. 35. Perform the following calculations. Record the answer to the correct number of significant digits and include the unit. a. 45m + 3.8m + 122m = _______________________________ b. 344mm x 12mm = _______________________________ c. 23444.0mg ‐ 4332.222mg = _______________________________ d. 994g / 21mL = _______________________________ 36. Distinguish between accuracy and precision. 37. Calculate the percent error of each of the following, based on the average of each set of data. Generally, if the % error is less than 5% it is considered accurate for this experiment. Determine whether each data set is accurate, precise, neither, or both if the accepted result is 19.50 g. a. Result 1, 19.77 g, Result 2, 19.00 g, Result 3, 19.82 g b. Result 1, 20.00 g, Result 2, 19.99 g, Result 3, 20.01 g 8
Unit 4 – Chemical Composition Knowledge Objectives: Students will know… 1. How to define and calculate the average atomic mass of an element. 2. The definition of an isotope. 3. What the mole represents. 4. The diatomic elements. 5. Avogadro’s number. 6. How to count particles by weighing them. 7. How to calculate the mass percent of an element in a compound. 8. How to define and calculate the empirical formula of a compound. 9. How to define and calculate the molecular formula of a compound. Skill Objectives: Students will be skilled at… 1. Converting among moles, mass, and number of particles in a given sample. 2. Calculating the molar mass of a substance. 3. Writing the nuclear symbol and hyphen notation for an isotope. Important Vocabulary: Average Atomic Mass Molar Mass Mole Empirical Formula Molecular Formula Percent Composition (Mass Percent) Avogadro’s Number Isotope Atomic Mass Unit (amu) Formula Weight 38. What is an isotope? 39. Magnesium has three naturally occurring isotopes, magnesium ‐ 24, magnesium ‐ 25, and magnesium ‐ 26. The relative abundances of these three isotopes are 78.99%, 10.00% and 11.01% respectively. The atomic mass of magnesium ‐ 24 is 23.985 amu, magnesium ‐ 25 is 24.985 amu, and magnesium ‐ 26 is 25.982 amu. Calculate the average atomic mass of magnesium. 40. Find the molar mass of each of the following substances: a. Potassium chlorate d. Dicarbon hexahydride b. Gold(II) acetate e. ammonium sulfate
c. Pentanitrogen dioxide 9
41. One mole = __________________________ particles = _______________________ grams 42. Determine the mass in grams of one atom of gold. 43. Calculate the number of atoms in 10.0 g of sulfur trioxide. 44. Calculate the number of moles in 109 g of iron (II) bromide. 45. Calculate the number of moles in 7.899 x 1023 atoms of zirconium. 46. Calculate the mass in kg of 5.66 x 1025 particles of calcium acetate. 47. Determine the percent composition of each element in ammonium phosphate. 48. Determine the empirical formula of the compound containing 15.8% boron and 84.2% fluorine. 10
49. Determine the molecular formula of a compound having the empirical formula C9H17O and a molar mass of 847.56 g/mol. 50. Calculate the empirical and molecular formulas for a substance that is 49.5% Carbon, 5.15% Hydrogen, 28.9% Nitrogen, and 16.5% Oxygen by mass. The molecular mass for the substance is 195 g/mol. Unit 5 – Atomic Theory & Nuclear Chemistry Knowledge Objectives: Students will know… 1. The internal components of an atom and be able to label them on a diagram. 2. The properties of protons, neutrons and electrons. 3. About the experiments conducted by Rutherford, J.J. Thomson, and Milliken. 4. The contribution(s) to the development of the atomic theory by Democritus, Boyle, Rutherford, Alchemists, J.J. Thomson, Milliken, Chadwick, Priestly, Dalton, Bohr, Lord Kelvin, Praust, Aristotle, and Lavoisier. 5. Several types of radioactive decay. 6. How to write nuclear equations for various types of radioactive decay. 7. The concept of a half‐life and various applications of the concept. 8. How to calculate the age of a sample using the original mass and the half‐life of radioactive isotope. Skill Objectives: Students will be skilled at… 1. Graphing the half‐life of a radioactive isotope. 2. Constructing a Bohr model of an atom. Important Vocabulary: Alchemy Electron Alpha Particle Electron Capture Alpha‐Particle Production Gamma Ray Atomic Number Half‐Life Beta Particle Isotope Beta‐Particle Production Law of Constant Composition Bohr Model Mass Number Chemical Formula Neutron Compound Nuclear Atom Dalton’s Atomic Theory Nuclear Equation Decay Series Nuclear Transformation Nucleus Orbitals Plum Pudding Model Positron Positron Production Proton Radioactive Radiocarbon Dating Radiotracers Transuranium Elements 11
Important People: Alchemists Aristotle Democratus Ernst Rutherford Henri Becquerel J.J. Thompson James Chadwick John Dalton Lavoisier Lord Kelvin Louis de Broglie Marie & Pierre Curie Niels Bohr Robert Millikan
51. State the difference between each of the following terms: atomic number, mass number, and atomic mass. 52. Fill in the following chart. Isotope Nuclear Symbol Atomic # Protons Neutrons Electrons Mass # Charge Tungsten ‐ 184 0 34 45 34 Holmium ‐ 165 +2 20 40 +2 53. Show both the hyphen notation and the nuclear symbol for the isotope of oxygen that has a mass number of 15. 54. Label the parts of the atom in the diagram. 55. Identify the subatomic particles ‐ particles that make up an atom. For each particle give its relative charge and location. Particle Relative Charge Location 12
56. Describe, in detail, the accomplishments and experiments of the following people: a. Democritus b. John Dalton c. Antoine Lavoisier d. Robert Boyle e. Lord Kelvin (William Thomson) f.
Joseph Proust g. J. J. Thomson. h. Ernest Rutherford 57. Complete the following table: Particle Name / Decay Type Symbol(s) Alpha Particle / Alpha Decay Beta Particle / Beta Decay Positron / Positron Emission Gamma Ray / Gamma Ray Production Electron / Electron Capture 58. Complete the following nuclear reactions and identify the reaction types: a.  b.  c. d. e. + _______ __________________________________________ __________________________________________ __________________________________________ __________________________________________ + _______ __________________________________________ + _______ + _______   *  + _______ 13
59. Write nuclear reactions for the following: a. The beta decay of cesium ‐ 142. b. The electron capture reaction of europium ‐153 c. The beta decay of manganese‐54 is accompanied by gamma ray production. d. The decay of chlorine‐36 by positron emission. e. The alpha decay of polonium‐210. 60. Define Half‐Life. 61. Silicon‐31 has a half‐life of approximately 2.5 hours. If we begin with a sample containing 2000 kg of Si‐
31, what is the approximate amount remaining after 10 hours? 62. Carbon‐15 (
) has a half‐life of 5.0 seconds. Suppose we have a sample containing 100 grams of carbon‐15. How much will remain after 30 seconds? 14
Unit 6 – Modern Atomic Theory Knowledge Objectives: Students will know… 1. The electromagnetic spectrum and the visible spectrum within it. 2. The relationship between energy, color, and wavelength of light. 3. The wave‐particle nature of light. 4. What happens when electrons move between ground states and excited states. 5. Energy levels of all atoms are quantized and how this relates to the various line‐emmision spectra of each element. 6. The wave mechanical model of the atom. 7. How to predict the location of an element on the periodic table based on ionization energies. 8. The various types of electron orbitals. 9. The quantum numbers. 10. Pauli’s Exclusion Principle, Aufbau’s Principle, and Hund’s Rule and how they apply to locating electrons around an atom. Skill Objectives: Students will be skilled at… 1. Explaining how the periodic table of the elements is organized and labeling several common families of elements. 2. Differentiating between metals, nonmetals, and metaloids. 3. Discussing the following periodic trends: atomic radius, ionic radius, ionization energy, electron affinity, electronegativity, and metalic character. 4. Writing full electron configuration for various elements. 5. Writing orbital notation electron configuration for various elements. 6. Writing noble gas electron configuration for various elements. 7. Using the electron configurations to identify the core and valence electrons for an element. Important Vocabulary: Alkali Metal Orbital Notation Group Alkaline Earth Metal Pauli Exclusion Principle Halogen Atomic Radius Period Hund’s Rule Aufbau Principle Periodic Table of the Elements Ionic Radius Bohr Model Photon Ionization Energy Core Electrons Principle Energy Level Line‐emmision Spectrum Main Group Elements Diatomic Element Quantized Metal Electromagnetic Radiation Quantum Number Metalic Character Electromagnetic Spectrum Representative Elements Metaloid Electron Affinity Sublevel Noble Gas Electron Configuration Transition Element Noble Gas Configuration Electronegativity Valence Electron Noble Metal Excited State Wave Mechanical Model Nonmetal Frequency Wavelength Orbital Ground State 63. List several common types of electromagnetic radiation in order from longest wavelength to shortest wavelength. 64. As the frequency of a wave increases, the wavelength _____________________. 65. Waves with a higher frequency have __________________ energy. 15
66. Explain what is meant by the wave‐particle nature of light. 67. What is the difference between an excited and a ground state? When an atom returns to its ground state, what happens to the excess energy of the atom? Draw a diagram to demonstrate the change. ` 68. Are the colors of flame tests due to taking in energy or releasing energy? Explain. 69. Do atoms in excited states emit radiation randomly, at any wavelength? Explain. 70. What does it mean when we say energy levels are quantized? 71. What are the essential points of Bohr’s theory of the structure of the hydrogen atom? 72. Write the location on of each of the following families or classifications of elements on a periodic table: metals, nonmetals, metalloids, alkali metals, alkaline earth metals, halogens, noble gases, transition metals, & diatomic elements. 16
73. Complete the following table: Property Trend  Period Trend  Group Atomic Radius Electron Affinity Electronegativity Ionization Energy 74. Arrange the following atoms from largest to smallest atomic radius, and from highest to lowest ionization energy. a. Cs, K, Li b. Ba, Sr, Ca c. I, Br, Cl d. Mg, Si, S 75. Explain the difference among the terms energy level, sublevel, and orbital. 76. How many electrons can be placed in a. a given s subshell? _________ b. a given p subshell? _________ c. a specific p orbital? _________ d. a given d subshell? _________ e. a given d orbital? _________ f.
_________ a given f subshell? 77. What are the differences between the 2s orbital and the 1s orbital for an atom? How are they similar? 78. How does the energy of a principal energy level depend on the value of n? Does a higher value of n mean a higher or lower energy? 79. The number of sublevels in a principal energy level (increases/decreases) as n increases. 17
80. Define Pauli’s Exclusion Principle and explain how it is used in creating electron configurations. 81. Define Aufbau’s Principle and explain how it is used in creating electron configurations. 82. Define Hund’s Rule and explain how it is used in creating electron configurations. 83. Write complete electron configuration diagrams for the following: a. As b. C 84. Write orbital configuration diagrams for the following: a. Mg b. P 85. Write noble gas configuration diagrams for the following: c. Fe a. Sr d. Lu
b. Si 86. Why are the valence electrons more important to the atom’s chemical properties than the core electrons? How is the number of valence electrons in an atom related to the atom’s position on the periodic table? 18
Unit 7 ‐ Nomenclature Knowledge Objectives: Students will know… 1. How to draw Lewis structures for elements and compounds. 2. The difference between bonding pairs of electrons and lone pairs of electrons. 3. How to define and draw resonance structures. 4. How to write stable electron configurations for ions. 5. What polyatomic ions are. Skill Objectives: Students will be skilled at… 1. Differentiating between ionic and covalent bonds and explaining how they are formed. 2. Predicting the type of ion formed by various elements. 3. Predicting formulas of ionic compounds. 4. Identifying and naming Type I ionic compounds. 5. Writing formulas for Type I ionic compounds from the name. 6. Identifying and naming Type II ionic compounds. 7. Writing formulas for Type II ionic compounds from the name. 8. Identifying and naming Type III binary compounds. 9. Writing formulas for Type III binary compounds from the name. 10. Matching the names and formulas of several common polyatomic ions. 11. Identifying and naming Type IV compounds. (Contain Polyatomic Ions) 12. Writing formulas for Type IV compounds from the name. 13. Identifying and naming acids. 14. Writing formulas for acids from the name. Important Vocabulary: Duet Rule Acid Nonmetal Electronegativity Anion Octet Rule Ion Binary Compound Oxyanion Ionic Bond Binary Ionic Compound Polyatomic Ion Ionic Compound Cation Valence Electron Metal Core Electron Molecule Covalent Bond 87. Compare and contrast ionic and covalent compounds. 88. Determine whether the following bonds are ionic or covalent. SHOW YOUR WORK. d. H – C a. Mg – O e. Cl – Cl b. Ba – Cl c. K – I 89. What is the difference between an atom and an ion? 19
90. What is the difference between a cation and an anion? Give an example of each. 91. Write the symbols for the following ions: a. lithium ion d.
phosphate ion b. palladium (II) ion e.
nitride ion c. silver ion f.
molybdenum (IV) ion 92. Write the electron configuration for the following atoms. Write the electron configuration for the simple ion they become, and write the name of the noble gas corresponding to the ion. THIS HAS 3 PARTS! Electron Configuration Corresponding Element Electron Configuration for Simple Ion Noble Gas N Sr Al Cl Li 93.
In MgI2, how do we know that two iodide ions are needed for each magnesium ion? 94. What is a polyatomic ion? List the names and formulas of 10 polyatomic ions without consulting your polyatomic ion sheet. 95.
If more than one of a polyatomic ion is used in a formula what must be put around it? 96.
What is an acid? Give two examples (one which contains oxygen and one that does not). 20
97.
Complete the table below: Formula Mg(ClO4)2 HIO3 SF6 NaCl P4O10 CO CoBr2 ICl CuO NO2 K2O MnO2 N2O3 MgBr2 FeCl3 Zr(C2H3O2)4 HI BaCl2 Ions (if needed) Name Type 21
98. Complete the table below: Name Carbon monoxide Manganese (II) chloride Dinitrogen trioxide Potassium sulfide Barium phosphate Sodium hydroxide Nitric acid Ammonium acetate Hydrosulfuric acid Silver chloride Copper (II) bromide Hydrochloric acid Calcium sulfate Aluminum sulfide Silver sulfite Magnesium oxalate Ions (if needed) Formula Type Unit 8 – Chemical Reactions and Driving Forces Knowledge Objectives: Students will know… 1. Types of evidence that a chemical reaction has occurred. 2. The parts and symbols within a chemical equation. 3. The 4 possible driving forces for chemical reactions. 4. How to classify chemical reactions according to driving force. 5. The 5 types of chemical reactions. 22
Skill Objectives: Students will be skilled at… 1. Writing a chemical equation from the word equation. 2. Writing a word equation from a chemical equation. 3. Balancing chemical equations. 4. Classifying chemical reactions according to type. 5. Using the activity series to determine the likelihood of a single displacement reaction occurring. 6. Using the solubility chart to determine the likelihood of a double displacement reaction occurring. 7. Predicting the products of a chemical reaction. Important Vocabulary: Activity Series Salt Insoluble Aqueous Single Displacement Reaction Neutralization Reaction Catalyst Slightly Soluble Oxidized Chemical Equation Solubility Precipitate Coefficient Subscript Precipitation Reaction Combustion Reaction Synthesis Reaction Product Reactant Decomposition Reaction Yields Redox Reaction Double Displacement Reaction Reduced Gas Production Reaction 99. What are 4 indicators that a chemical reaction has occurred? 100. In a chemical equation: a. where do we find the reactant(s)? b. where do we find the product(s)? c. what does the arrow represent? 101. What does each of the following represent? (s) = ______________________ (g) = _______________________ (l) = _______________________ (aq) = ______________________ 102. Write balanced equations for each of the following: a. Barium solid reacts with octasulfur solid to produce barium sulfide solid. b. Ammonia gas reacts with chlorine gas to produce solid ammonium chloride and gaseous nitrogen trichloride. 23
c. When a strip of magnesium metal is heated in oxygen, it bursts into an intensely white flame and produces a finely powdered dust of magnesium oxide. d. When lead(II) sulfide is heated to high temperatures in a stream of pure oxygen gas, solid lead(II) oxide forms with the release of gaseous sulfur dioxide. 103. Write the word equation from the following chemical equations and balance. a. Na2O2(s) + H2O(l)  NaOH(aq) Br2(l)  FeBr3(s) + O2(g) b. Fe(s) + 104. What do the coefficients in a balanced chemical equation represent? What do the subscripts in a balanced chemical equation represent? Which can be changed when balancing a chemical equation? 105. Balance the following equations: a. FeCl3(aq) + KOH(aq)  Fe(OH)3(s) + KCl(aq) b. AgC2H3O2(aq) + HCl(aq)  AgCl(s) + c. SnO(s) + C(s)  + d. K2O(s) + H2O(l)  KOH(aq) g. Na2S(s) + HCl(aq)  NaCl(aq) h. H2SO4(aq) + NaCl(s)  i. N2(g) I2(s)  HC2H3O2(aq) Sn(s) CO2(g) + H2S(g) Na2SO4(aq) + HCl(aq) + NI3(s) 24
106. What are the 4 possible driving forces for chemical reactions? 107. Given the reaction 2 Ca + O2  2 CaO a. Identify the substance that is oxidized? b. Identify the substance that is reduced? 108. What are the 5 patterns of chemical reactions? 109. Complete the table below. EXAMPLE: Sodium metal reacts with aqueous aluminum chloride  aluminum (s)
3 Na(s)
+
AlCl3 (aq)

Al (s)
+
3 NaCl (aq)
+ Sodium chloride (aq)
Driving
Force
Redox
Pattern
CSD
Aluminum metal plus copper(II) iodide solution  Driving
Force
Pattern
Lithium hydroxide solution plus phosphoric acid  Driving
Force
Pattern
Sodium chloride solution plus silver nitrate solution  Driving
Force
Pattern
Propane gas burns in oxygen gas  Driving
Force
Pattern
25
Magnesium metal reacts with oxygen gas  Driving
Force
Pattern
Acetic acid solution (HC2H3O2) is added to sodium hydroxide solution  Driving
Force
Pattern
Cadmium oxide decomposes when heated  Driving
Force
Pattern
Nickel(II) sulfide solution plus sodium metal  Driving
Force
Pattern
Hydroiodic acid solution plus lead(II) nitrate solution  Driving
Force
Pattern
Chromium(III) chloride solution plus fluorine gas  Driving
Force
Pattern
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