Unit 1 Atomic Structure

Unit 1 Atomic Structure Unit 1 Text Questions 1.1 Atoms/Ions/Isotopes Problems Ch 5 Prob: 9,23,24,38 1.2 Average Atomic Mass Problems Ch 5 Prob: 15,17 1.3 Atomic Theory Development Problems Ch 5 Prob: 34,36 1.4 Energy Levels, Sublevels and Orbitals Problems Ch 13 Prob: 4,27 1.5 Orbital Box Notation and Electron Configurations Problems Ch 13 Prob: 5,26,45,47 Section 1.1 Atoms/Ions/Isotopes 1. What is the net charge on a particle that consists of: a. 15 protons, 13 neutrons and 18 electrons? b. 37 protons, 30 neutrons and 36 electrons? 2. Identify the number of protons, neutrons and electrons in: a. potassium 1+ -40 b. sulfur-33 c. lithium-6 Section 1.2 Average Atomic Mass 3. Although the atomic mass of boron is listed as 10.81, there is no boron atom with that relative mass. Explain! 4. What is the mass, in atomic mass units, of an atom having six times the mass of a carbon-12 atom? 5. Calculate the average atomic mass of an element in which 60% of its atoms have a mass of 25 amu and 40% have a mass of 30 amu. Section 1.3 Atomic Theory Development 6. List each model in chronological order. a. Dalton s atomic theory b. charge cloud model c. plum pudding model Section 1.4 Energy Levels, Sublevels and Orbitals 7. How does an electron in an atom absorb and release energy? 8. How many electrons can fit in an s,p,d and f orbital? 9. How many electrons can fit in an s,p,d and f sublevel? Section 1.5 Orbital Box Not. and Electron Config. 10. Write an electron configuration and orbital box notation for an atom of oxygen, magnesium, and vanadium. Identify the energy level of the valence shell and number of valence electrons in each. Unit 1 Supplementary Questions

UNIT 2 The Periodic Table Unit 2 Text Questions 2.1 Origin of the Periodic Table Problems Ch 5 Prob: 28 2.2 Parts of the Periodic Table Problems Ch 5 Prob: 30,48,49/Ch 14 Prob: 10,13 2.3 Electron Configurations Problems Ch 14 Prob: 4 2.4 Periodic Trends Problems Ch 14 Prob: 18,19,20,21,23 2.5 Octet Rule and Ion Charge Problems Ch 14 Prob: 24,32 Unit 2 Supplementary Questions Section 2.1 Origin of the Periodic Table 11. a. How did Mendeleev arrange the elements in his periodic table? b. What did the empty spaces in his table represent? 12. How does the original periodic law differ from the modern periodic law? Section 2.2 Parts of the Periodic Table 13. What is the name is given to: a. a vertical column on the periodic table? b. a horizontal row in the periodic table? 14. Classify each element in as many different ways as possible. a. calcium c. silicon b. bromine d. copper Section 2.3 Electron Configurations 15. What is the relationship between the number of a period on the periodic table and the distribution of electrons of all elements of that period? 16. Write a noble gas notation, and then identify the energy level of the valence shell and the number of valence electrons for each of the following elements. a. fluorine b. iron Section 2.4 Periodic Trends 17. What are the trends in atomic radius and ionization energy from: a. left to right across a period? b. top to bottom down a column? Section 2.5 Octet Rule and Ion Charge 18. What ion is typically formed by an element in: a. group IIA? b. group 17? c. group VIA? d. group 1?

UNIT 3 Chemical Bonding Unit 3 Text Questions 3.1 Electron Dot Diagrams Problems Ch 15 Prob: 1,6,23,25,30 3.2 Ionic Bonding Problems Ch 15 Prob: 7,9,11,12 3.3 Covalent Bonding Problems Ch 16 Prob: 29 3.4 Metallic Bonding Problems Ch 15 Prob: 17,18 3.5 Covalent Structures Problems Ch 16 Prob: 30,47 3.6 Molecular Geometry Problems Ch 16 Prob: 63 3.7 Drawing Structures Problems Ch 16 Prob: 48 3.8 Polarity and Hydrogen Bonding Problems Ch 16 Prob: 21/Ch 17 Prob: 19 Section 3.1 Electron Dot Diagrams 19. Draw an electron dot structure for each of the following elements and the common ion they form: a. nitrogen b. calcium Section 3.2 Ionic Bonding 20. How is an ionic bond formed between an atom of sodium and an atom of chlorine? Illustrate its formation using electron dot diagrams. Section 3.3 Covalent Bonding 21. How is a covalent bond formed between two atoms of fluorine? Illustrate its formation using electron dot diagrams. Section 3.4 Covalent Structures 22. How many electrons are in a: a. single bond? b. triple bonds? b. double bonds? c. lone pairs? 23. Identify the number of single bonds, double bonds and triple bonds, as well as the number of valence electrons (by group heading and by number of electron pairs shown) in the structure-1 to the left. Section 3.5 Molecular Geometry 24. What is the bond angle in a molecule exhibiting molecular geometry. a. tetrahedral b. trigonal planar 25. Identify the molecular geometry around each central atom in each structure-1 to the left Section 3.6 Drawing Structures 26. Draw the following structures and identify the molecular geometry of each. a. PH 3 a. 2- CO 3 b. H 2 S a. 2- SO 4 Section 3.7 Polarity and Hydrogen Bonding 27. Describe the sharing of electrons in a polar bond. Unit 3 Supplementary Questions

UNIT 4 Nomenclature Unit 4 Text Questions 4.1 Names and Formulas Naming Common Acids Problems Ch 6 Prob: 51,65 4.2 Binary Ionic Compounds Problems Ch 6 Prob: 26,27 4.3 Ternary Ionic Compounds Problems Ch 6 Prob: 43,44,61 4.4 Binary Molecular Compounds Problems Ch 6 Prob: 64,65 Section 4.1 Names and Formulas 28. Write the empirical formula for each substance. a. C 2 H 6 b. CH 4 c. C 3 H 6 O 3 d. AlCl 3 29. Identify the number of atoms of each element in each substance. a. (NH 4 ) 2 CO 3 b. Ca 3 (PO 4 ) 2 30. Write a molecular formula and empirical formula for structure-2 to the left. Section 4.2 Binary Ionic Compounds 31. Write a name when given a formula and vice versa. a. CaCl 2 d. aluminum sulfide b. Fe2O 3 e. cobalt(ii) bromide c. HCl f. lead (II) fluoride Section 4.3 Ternary Ionic Compounds 32. Answer the following questions about magnesium nitrate. a, How many magnesium ions are in each formula unit? b. How many nitrate ions are in each formula unit? c. Identify the number of magnesium, nitrogen and oxygen atoms. 33. Write a name when given a formula and vice versa. a. nitric acid b. (NH 4 ) 2 S c. aluminum nitrate d. FeSO 4 Section 4.4 Binary Molecular Compounds 34. Write a name when given a formula and vice versa. a. CO 2 b. phosphorus trichloride c. P 2 O 7 d. carbon disulfide e. H 2 SO 4 35. Why is this name - formula pairing incorrect? sodium (I) chloride - NaCl Unit 4 Supplementary Questions

UNIT 5 Mathematics Unit 5 Text Questions 5.1 Importance of Measurement Problems Ch 3 Prob: 2,13,16 5.2 Scientific Notation Problems Ch 3 Prob: 3,4,15 5.3 International System of Measurement Problems Ch 3 Prob: 19,22,52,56 5.4 Dimensional Analysis Introduction Problems Ch 4 Prob: 43 5.5 Complex Dimensional Analysis Problems Ch 4 Prob: 44,47 5.6 Density Problems Ch 3 Prob: 23,24 5.7 Temperature Problems Ch 3 Prob: 33,34 Unit 5 Supplementary Questions Section 5.1 Importance of Measurement 36. The mass of a certain object is 42.0 g. Student 1 makes three separate measurements of its mass: 41.9 g, 41.9 g and 42.1 g C. Student 2 makes three measurements as well: 45.0 g, 45.1 g and 44.9 g. Describe their measurements in terms of accuracy and precision. Section 5.2 Scientific Notation 37. Describe each of the following measurements as; less than -1, greater than negative one yet less than zero, greater than zero yet less than positive 1 or greater than positive 1. a. 2.0 x 10 3 b. -2.0 x 10-4 c. 2.0 x 10-5 d. -2.0 x 10 5 Section 5.3 International System of Measurement 38. Which is the largest quantity? 1 gal / 1 ml / 1 L / 1 qt Section 5.4 Dimensional Analysis Introduction 39. Write four conversion factors for the following multiple equality. ( 5280 ft = 1 mile = 1760 yards ) Section 5.5 Complex Dimensional Analysis 40. Convert 25 ft per second to kilometers per hour. Section 5.6 Density 41. What is the mass, in g, of a rock with a density to 3.00 g/ml and a volume of 2 ml? Section 5.7 Temperature 42. List the freezing point and boiling point of water on the Celsius scale, Fahrenheit scale and Kelvin Scale. 43. Make the following temperature conversions. a. 25 F to C b. -25 C to F

UNIT 6 The Mole Unit 6 Text Questions 6.1 The Concept of the Mole Problems Ch 7 Prob: 11 6.2 Molar Mass Problems Ch 7 Prob: 7 6.3 Mole Calculations Problems Ch 7 Prob: 16,19,59 6.4 Percent Composition Problems Ch 7 Prob: 29,31,62 6.5 Empirical and Molecular Formulas Problems Ch 7 Prob: 37,40,66 6.6 Molarity Problems Ch 18 Prob: 10,11,21 6.1 The Concept of the Mole 44. Compare a mole of copper to a mole of helium in terms of mass, volume and number of particles. 45. What is Avogadro s number? 6.2 Molar Mass 46. What is the molar mass of each compound. a. NaHCO 3? b. H 2 SO 4 6.3 Mole Calculations 47. Express 6.25 mol CO 2 in grams and liters at STP. 48. What is the mass of 125 L of N 2 O 3 at STP? 49. A sample of NO 2 has a mass of 156 g. How many molecules are in this sample? 6.4 Percent Composition 50. What is meant by a substance being 75 % C and 25 % H by mass? 51. What is the percent composition of C 2 H 4 O 2? 52. How many grams of C and H are in a 395 g sample of C 3 H 8? 6.5 Empirical and Molecular Formulas 53. What is the relationship between an empirical and a molecular formula? 54. A compound is 82.8 %C and 17.2 %H, and has a molar mass of 58.0 g/mol. Determine the empirical and molecular formula of this compound. 55. A compound is 40 %C, 6.7 %H and 53.3 %O, and has a molar mass of 90.0 g/mol. Determine the empirical and molecular formula of this compound. 6.6 Molarity 56. Which molarity is greater? a. 100 g of NaCl in 500 ml of water. b. 200 g of C 6 H 12 O 6 in 750 ml of water. Unit 6 Supplementary Questions

UNIT 7 Chemical Equations Unit 7 Text Questions 7.1 Word and Formula Equations Problems Ch 8 Prob: 1,9 7.2 Balancing Equations Problems Ch 8 Prob: 3,10,39 7.3 Types of Reactions Problems Ch 8 Prob: 22 7.4 Moles and Balanced Equations Problems Ch 9 Prob: 6,7 7.5 Limiting and Excess Reactant Concept Problems Ch 9 Prob: 29 7.6 Stoichiometry Problems Ch 9 Prob: 21,22 7.7 Limiting and Excess Reactant Calculations Problems Ch 9 Prob: 23,24 7.8 Percent Yield Problems Ch 9 Prob: 30 Unit 7 Supplementary Questions 7.1 Word and Formula Equations 57. Write six facts about the equation that follows. Zn(s) + HCl 2 (aq) ZnCl 2 (aq) + H 2 (g) 7.2 Balancing Equations 58. Why is it incorrect to balance 2Al(s) + O 2 (g) Al 2 O 3 (s) it in the following ways? a. 2Al(s) + O 3 (g) Al 2 O 3 (s) b. Al(s) + O 2 (g) AlO 2 (s) c. 2Al(s) + 3O(g) Al 2 O 3 (s) 7.3 Types of Reactions 59. What is the relationship between a decomposition and synthesis reaction. 60. Determine the reaction type of each reaction. a. NaCl + AgNO 3 NaNO 3 + Ag Cl b. AgNO 3 + Cu Cu(NO 3 ) 2 + Ag 7.4 Moles and Balanced Equations 61. Balance each equation. a. N 2 + O 2 N 2 O 5 b. Al + F 2 AlF 3 c. O2 + Sb 2 S 3 Sb 2 O 4 + SO 2 7.5 Limiting and Excess Reactant Concept 62. Identify each statement as true of false. a. The limiting reactant is totally consumed. b. The excess reactant is mixed with the product. c. The excess reactant determines the amount of product produced. d. In a perfect mixing ratio, both reactants are considered to be limiting. 7.6 Stoichiometry 63. For the reaction... 4Na + O 2 --> 2Na 2 O How many grams of sodium are required to react with 98.0 g of sodium? 64. For the reaction... 2P + 5Cl 2 2PCl 5 How many liters of chlorine gas are required to form 265 g of phosphorus pentachloride? 7.7 Limiting and Excess Reactant Calculations 65. How many grams of water can be produced by the reaction of 12 g of hydrogen and 64 g of oxygen. 7.8 Percent Yield 66. When a 25 g piece of zinc is placed into a solution of copper sulfate, a student recovers 20.0 g of copper. What is the percent yield for this reaction? Zn + CuSO 4 ZnSO 4 + Cu

End of Semester: Important Information Friday 1/15 Unit 7 Test Tuesday 1/19 Go Over Test Mini Lab Wednesday 1/20 Library for element project Thursday 1/21 All components of Element Project Due Unit 1 and 2 Review Discussion Work on Review Packet Friday 1/22 Unit 1 Due (5 points) Unit 2 Due (5 points) Presentations Unit 1 and 2 Question and Answer Session Unit 3 and 4 Review Discussion Work on Review Packet Monday 1/25 Unit 3 Due (5 points) Unit 4 Due (5 points) Presentations Unit 3 and 4 Question and Answer Session Unit 5 and 6 Review Discussion Work on Review Packet Tuesday 1/26 Unit 5 Due (5 points) Unit 6 Due (5 points) Presentations Unit 5 and 6 Question and Answer Session Unit 7 Review Discussion Work on Review Packet Wednesday 1/27 Unit 7 Question and Answer Session Unit 7 Due (5 points) Exam Review and Discussion Lab Clean-up Day Thursday 1/28 8:00-9:15 Block 1 Exam 9: 30-10:45 Block 3 Exam 11:00-12:15 Block 4 Exam Friday 1/29 8:00-9:15 Block 2 Exam 9: 30-10:45 Block 5 Exam 11:00-12:15 Make-Up Exam