Differentiated Chemistry Worksheet and Laboratory
Differentiated Chemistry Worksheet and Laboratory Manual (Term 2) Mr. Geist 21 22 23 24 25 26 27 (260) (226) (223) Lr Lawrencium Barium 137.33 88 Cesium 132.91 87 Ra Lutetium 174.97 103 56 Ba 55 Cs Radium 71 Lu 87.62 Fr Yttrium 88.906 Strontium Rubidium 85.468 Francium 40 Y Vanadium (262) Tungsten (263) Seaborgium Sg 183.85 106 Cerium 140.12 Lanthanum 138.91 90 Th Thorium 232.04 89 Ac Actinium (227) 59 Pr 231.04 Proactinium Pa 91 140.91 Praseodymium Actinide series 58 Ce 57 La Lanthanide series (261) Dubnium Db Tantalum W 74 Molybdenum 73 Ta 180.95 105 Rf 42 Mo 95.94 Hafnium Rutherfordium Chromium 51.996 92.906 Niobium Nb 41 50.941 178.49 104 Hf 72 91.22 Zirconium Zr 39 38 Sr 37 Rb Titanium 47.90 Scandium 44.956 Calcium 40.08 Potassium 39.098 Iron 238.03 Uranium U 237.05 Neptunium Np 93 (145) 92 Promethium 144.24 Pm 61 (265) Hassium Hs 190.2 108 Osmium Os 76 101.07 Ruthenium Ru 44 55.847 Neodymium Nd 60 (262) Bohrium Bh 186.21 107 Rhenium Re 75 (97) Technetium Tc 43 54.938 Manganese (244) Plutonium Pu 94 150.4 Samarium Sm 62 (266) Meitnerium Mt 192.22 109 Iridium Ir 77 102.91 Rhodium Rh 45 58.933 Cobalt Co 28 29 79 Platinum (243) Americium Am 95 151.96 Europium Eu 63 (269) Weird Uum 195.09 110 (247) Curium Cm 96 157.25 Gadolinium Gd 64 (272) More weird Uuu 196.97 111 Gold Au 78 Pt 107.87 Silver Ag 47 63.546 Copper Cu 106.4 Palladium Pd 46 58.71 Nickel Ni 30 (247) Berkelium Bk 97 158.93 Terbium Tb 65 (277) Most weird Uub 200.59 112 Mercury Hg 80 112.41 Cadmium Cd 48 65.38 Zinc Zn K Fe 20 Ca 19 Mn 24.305 Cr Magnesium Sodium 22.990 V Aluminum 12 Mg 11 Na Ti 13 Al 9.0122 6.941 Sc 10.81 Beryllium Lithium 5 (251) Californium Cf 98 162.50 Dysprosium Dy 66 204.37 Thallium Tl 81 114.82 Indium In 49 69.72 Gallium Ga 31 26.982 Boron B 4 Be Li 3A 1.0079 3 2A 51 Tin (254) Einsteinium Es 99 164.93 Holmium Ho 67 207.2 Lead Pb 82 118.69 (257) Fermium Fm 100 167.26 Erbium Er 68 208.98 Bismuth Bi 83 121.75 Antimony Sb 50 Sn 74.922 Arsenic As 33 30.974 Phosphorus P 15 14.007 Nitrogen N 7 5A 72.59 Germanium Ge 32 28.086 Silicon Si 14 12.011 Carbon C 6 4A (258) Mendelevium Md 101 168.93 Thullium Tm 69 (209) Polonium Po 84 127.60 Tellurium Te 52 78.96 Selenium Se 34 32.06 Sulfur S 16 15.999 Oxygen O 8 6A (259) Nobelium No 102 173.04 Ytterbium Yb 70 (210) Astatine At 85 126.90 Iodine I 53 79.904 Bromine Br 35 35.453 Chlorine Cl 17 18.998 Fluorine F 9 7A (222) Radon Rn 86 131.30 Xenon Xe 54 83.80 Krypton Kr 36 39.948 Argon Ar 18 20.179 Neon Ne 4.0026 10 Helium 2 He H Hydrogen 8A 1 Do not white-out, add additional paper, or tape. Only write in box to the left, or be unable to use this sheet on the test. Exam: MASTER COPY Period: ________ Name: _____________________________ 1A You may add additional information in your own handwriting in this box. Periodic Table of Elements (Additional Values and Constants on back page) 1A H 2.1 Li 1.0 Na 0.9 K 0.8 Rb 0.8 Cs 0.7 Fr 0.7 3B Ti 1.5 Zr 1.4 Hf 1.3 Th 1.2 4B V 1.6 Nb 1.6 Ta 1.5 Pa 1.5 5B Cr 1.6 Mo 1.8 W 1.7 U 1.7 6B Mn 1.5 Tc 1.9 Re 1.9 7B Fe 1.8 Ru 2.2 Os 2.2 8B Co 1.9 Rh 2.2 Ir 2.2 Ni 1.9 Pd 2.2 Pt 2.2 nitrate sulfate phosphate nitrite sulfite phosphite carbonate acetate hydroxide ammonium silicate cyanide permanganate chromate dichromate 1B Zn 1.6 Cd 1.7 Hg 1.9 2B B 2.0 Al 1.5 Ga 1.6 In 1.7 Tl 1.8 3A C 2.5 Si 1.8 Ge 1.8 Sn 1.8 Pb 1.9 4A N 3.0 P 2.1 As 2.0 Sb 1.9 Bi 1.9 5A O 3.5 S 2.5 Se 2.4 Te 2.1 Po 2.0 6A F 4.0 Cl 3.0 Br 2.8 I 2.5 At 2.2 7A Kb = 0.512C/molal Kf = –1.86C/molal 8A He -Ne -Ar -Kr -Xe -Rn -- Fluorine Chlorine Bromine Iodine Colligative Property Constants for Water Lithium Potassium Calcium Sodium Magnesium Aluminum Zinc Chromium Iron Nickel Tin Lead Hydrogen Copper Mercury Silver Platinum Gold Activity Series of Metals/Halogens (NOTE: Reactivity of the metal/halogen decreases as it gets lower on the list.) Cu 1.9 Ag 1.9 Au 2.4 Average Electronegativities of the Elements 2A Sc 1.3 Y 1.2 Ln 1.0 Ac 1.0 – NO3 : 2– SO4 : PO43–: – NO2 : SO32–: 3– PO3 : CO32–: – C2H3O2 : OH–: + NH : 4 SiO32–: – CN : MnO4–: 2– CrO4 : 2– Cr2O7 : Main Polyatomic Ions Be 1.5 Mg 1.2 Ca 1.0 Sr 1.0 Ba 0.9 Ra 0.9 Other Polyatomic Ions HPO42–: hydrogen phosphate H2PO41–: dihydrogen phosphate HSO31–: hydrogen sulfite HSO41–: hydrogen sulfate HCO31–: hydrogen carbonate ClO41–: perchlorate 1– ClO3 : chlorate ClO21–: chlorite 1– ClO : hypochlorite 2– C2O4 : oxalate Monatomic Ions Cu1+: copper (I) ion Cu2+: copper (II) ion 2+ Fe : iron (II) ion Fe3+: iron (III) ion 2+ Pb : lead (II) ion Pb4+: lead (IV) ion 2+ Sn : tin (II) ion Sn4+: tin (IV) ion 2+ Co : cobalt (II) ion Co3+: cobalt (III) ion Useful Conversion Factors And Conversions Energy: 1 cal = 4.184 J Length: 1 angstrom = 0.100 nm 1 inch = 2.54 cm Mass: 1 lb = 0.4536 kg Pressure: 1 atm = 101.3 kPa 1 atm = 760 mm Hg C = K – 273.15 1 L = 0.001 m3 1 cm3 = 1 mL Temp.: Volume: General Physical Constants 6.022 x 1023 rp/mol 6.626 x 10–34 Js 1.381 x 10–23 J/K 3.0 x 108 m/s 1.6605655 x 10–27 kg 1.097 x 107 m–1 96485.309 C/mol 8.31 (LkPa)/(Kmol) 0.0821 (Latm)/(Kmol) 62.396 (Lmm Hg)/(Kmol) All compounds formed with the negative ion are insoluble except those of the alkali metals and NH4+. All compounds formed with the negative ion are soluble except Ag+, Pb2+, Hg22+, + and Cu . Most compounds formed with the negative ion are soluble; exceptions include SrSO4, BaSO4, CaSO4, RaSO4, Ag2SO4, and PbSO4. All compounds formed with the negative ion are soluble. Solubility Rules Rule Avogadro’s Constant Planck’s Constant Boltzmann’s Constant Speed of Light Atomic Mass Unit Rydberg’s Constant Faraday’s Constant Ideal Gas Constant – Negative Ion NO3 2– – – – I , Br , Cl SO4 CO32–, PO43–, 2– SO3 OH– 2– S All compounds formed with the negative ion are insoluble except those of the alkali metals, NH4+, Sr2+, and Ba2+. (Ca(OH)2 is slightly soluble.) All compounds formed with the negative ion are insoluble except those of the alkali metals, alkaline earth metals, and NH4+. Table of Contents Unit Five Worksheet ..................................................................................... 1 Unit Six Worksheet..................................................................................... 16 Unit Seven Worksheet ............................................................................... 21 Unit Eight Worksheet ................................................................................. 38 Unit Five Experiment – 1: Orbital Structures and Identifications ............... 45 Unit Five Experiment – 2: Metal Ions and Flame Tests ............................ 49 Unit Five Experiment – 3: Viewing Spectra ............................................... 50 Unit Six Experiment – 1: Paper Chromatography of Food Dyes ............... 52 Unit Seven Experiment – 1: The Ideal Gas Law ....................................... 54 Unit Seven Experiment – 2: Specific Heat Experiment ............................. 56 Unit Seven Experiment – 3: Solution Preparation and Rate of Reaction .. 59 Unit Seven Experiment – 4: Le Chatelier’s Principle (Physical and Chemical Changes) ............... 61 Unit Eight Experiment – 1: Using Indicators .............................................. 63 Unit Eight Experiment – 2: Determining Molarity by Neutralization ........... 66 Appendix A – Laboratory Equipment and LPS Safety Contract ............... A-1 Appendix B – SI Units and Conversions .................................................. A-5 Appendix C – Compound Name and Formula Writing ............................. A-8 Appendix D – Chemical Reactions and Quantities ................................ A-10 Appendix E – Study Skills ...................................................................... A-17 Appendix F – Electron Configuration Rules ........................................... A-20 Appendix G – Electron Dot Structure ..................................................... A-22 Appendix H – VSEPR Models ................................................................ A-24 Appendix I – Calorimetry Calculations ................................................... A-26 Appendix J – Water and Solutions ......................................................... A-28 Appendix K – Acid and Base Measurements ......................................... A-32 Appendix L – Acid and Base Notes........................................................ A-33 Appendix M – Neutralization Notes ........................................................ A-36 Appendix N – Practice Tests .................................................................. A-40 Appendix O – Practice Test Keys .......................................................... A-66 Unit Five Worksheet WS – DC – U5 Section 8.1 Short Answer. Answer the following questions. 1. What evidence showed that the particles in the beam of Crookes’s tube were negatively charged? 2. Suppose that two beams pass between a pair of oppositely charged plates. One of the beams is composed of electrons, and the other is composed of protons. Will the two beams bend in the same direction or in opposite directions? Explain. 3. What main feature of Dalton’s atomic model was abandoned after Thomson’s discoveries? 4. Make a diagram of a lithium atom, based on Thomson’s atomic model. 5. If a lithium atom lost one electron, forming a positive ion, how would the diagram drawn in question 4 be changed? page 1 – DC – T2 – BOOK Section 8.2 Identification. Identify which of the three types of radiation – alpha, beta, and/or gamma – each of the following describes. ___________6. Is not deflected by a magnet ___________10. Consists of ions ___________7. Has a negative charge ___________11. Is similar to light rays ___________8. Moves with the greatest speed ___________12. Consists of the same particles as cathode rays ___________9. Has the highest penetrating ability ___________13. Has the lowest penetrating ability Section 8.3 Short Answer. Answer the following questions. 14. Identify two differences between protons and electrons. 15. Which subatomic particle determines the identity of a specific element? What term defines this? 16. Explain what an isotope is. 17. What ultimately determines the instability of an isotope? page 2 – DC – T2 – BOOK Table Completion. Information Atomic Number 236 92 U 142 56 Ba Krypton-91 Iodine-131 27 3+ 13 Al 127 52 I Fill in the table using the information provided in the left-most column to identify the following properties of each isotope. Mass Number Number of Neutrons Number of Electrons Number of Protons 18. 24. 30. 36. 42. 19. 25. 31. 37. 43. 20. 26. 32. 38. 44. 21. 27. 33. 39. 45. 22. 28. 34. 40. 46. 23. 29. 35. 41. 47. Calculation. Show work or receive no credit. Express proper units and correct number of significant figures and decimal places. 48. The element oxygen contains three naturally occurring isotopes: 16 8O 17 8O 18 8O The relative abundances and atomic masses are 99.759% for oxygen-16 (mass = 15.995 amu), 0.037% for oxygen-17 (mass = 16.995 amu), and 0.204% for oxygen-18 (mass = 17.999 amu). Calculate the average atomic mass of oxygen. 49. The element nitrogen contains two naturally occurring isotopes: 14 7N 15 7N The relative abundances and atomic masses are 99.63% for nitrogen-14 (mass = 14.003 amu) and 0.37% for nitrogen-15 (mass = 15.000 amu). Calculate the average atomic mass of nitrogen. page 3 – DC – T2 – BOOK Chapter 8 General Questions Short Answer. Answer the following questions. 50. Elements in the periodic table are ordered according to increasing atomic number rather than increasing atomic mass. There are several places where the atomic number increases but the average atomic mass decreases. Identify two of these places, and explain why these exceptions occur. 51. Assume the nucleus of a fluorine atom is a sphere with a radius of 5 x 10–13 cm. Calculate the density of a fluorine nucleus. Compare this density with the atomic density of iridium, whose density is 22.6 g/cm3. (HINT: Recall how to find the volume of a sphere.) Section 9.1 Short Answer/Calculation. Answer the following questions and problems. Write nuclear equations for the following processes. 52. The alpha decay of polonium-218 53. The beta decay of lead-210 54. The alpha decay of americium-241 55. The beta decay of carbon-14 page 4 – DC – T2 – BOOK Write nuclear equations for the following radioactive processes. 56. alpha decay of francium-208 57. electron capture by beryllium -7 58. beta emission by argon-37 59. positron emission by fluorine-17 60. After 42 days, a 2.0 g sample of phosphorus-32 contains only 0.25 g of isotope. What is the halflife of phosphorus-32? Show work or receive no credit. Include proper units. 61. The mass of cobalt-60 in a sample is found to have decreased from 0.800 g to 0.200 g in a period of 10.5 years. Find the half-life of cobalt-60 and calculate how many years it will take for 0.200 g of cobalt-60 to decrease in the sample to 0.149 g. 62. What happens to the mass number and atomic number of an atom that undergoes beta decay? 63. A radioisotope of an element undergoes alpha particle decay. How do the atomic number and mass number of the particle change? 64. Bismuth-211 is a radioisotope. It decays by alpha emission to yield another radioisotope, which emits beta radiation as it decays to a more stable isotope. Write equations for the nuclear reactions and name the decay products. page 5 – DC – T2 – BOOK 65. A sample initially contains 70.0 g of an isotope of radon. After 6.6 days, the sample only contains 21.0 g radon. What is the half-life of this isotope of radon, and after how many more days will only 9.5 g radon remain? Give the composition of the nucleus of the following isotopes. 64 28 136 53 Ni I Gold-195 66. p+: ________ 68. p+: ________ 70. p+: ________ 67. n0: ________ 69. n0: ________ 71. n0: ________ Section 9.2 Short Answer. Complete the equations for the following transmutation reactions. Li + 01n 42 He + ________ 72. 6 3 73. 235 92 74. 27 13 75. 235 92 76. ________ + 01n U + 01n 141 56 Ba + ________ + 3 01n Al + 42 He 01n + ________ U 90 38 Sr + ________ + 01n + 4-10e 144 58 Ce + 90 38 Sr + 601n + 2-10e Section 9.3 Short Answer. Answer the following questions. 77. Identify two types of nuclear waste produced by nuclear power plants. page 6 – DC – T2 – BOOK 78. Assuming technical problems could be overcome, what are some advantages to producing electricity in a fusion reactor? 79. Describe how a nuclear fission power plant operates. 80. Why are spent fuel rods removed from a reactor core? 81. What do spent fuel rods contain? 82. What happens to spent fuel rods after they are removed? 83. The fission energy of uranium-235 is 2.0 x 107 kcal/g. The heat of combustion of coal is about 8.0 kcal/g. Approximately what mass of coal must be burned to produce the energy released by the fission of 1.0 gram of uranium-235? Section 10.1 Calculations. Solve the following problems. Show work. Include proper units and significant figures. 84. What is the wavelength of the radiation whose frequency is 5.00 x 1015 s-1? 85. An inexpensive laser that is available to the public emits light that has a wavelength of 670 nm. What is the frequency, in hertz, of the radiation? page 7 – DC – T2 – BOOK 86. What is the energy of two moles of photons traveling with a frequency of 2.22 x 1014 s-1? 87. What is the frequency, in hertz, of a photon whose energy is 6.00 x 10-15 J? 88. Suppose that an AM radio station broadcasts at a frequency of 1600 kHz. What is the wavelength in meters of the radiation from the station? 89. What is the energy of a mole of photons whose wavelength is 658 nm? 90. Talking on a cell phone is possible because of the electromagnetic signals sent and received. What is the energy of one photon in a signal that is sent at 850 MHz? 91. What is the energy, in kilojoules, of a one mole of photons from question 87? page 8 – DC – T2 – BOOK Short Answer. Answer the following questions. We spend quite a bit of time staring at the red and green lights in a traffic signals. While that image is in your mind, for questions 92 – 95, decide which would have the following for a red light with a 4.41 x 1014 Hz frequency or a green light with a 6.00 x 1014 Hz frequency. Explain. 92. Longer wavelength 93. Greater speed 94. Greater energy 95. Greater amplitude 96. Explain the statement “Water waves transmit energy, not matter.” Section 10.2 Short Answer. Answer the following questions. 97. What is a bright-line spectrum? 98. Explain why the bright-line spectrum of hydrogen is composed of discrete lines and is not a continuous spectrum. page 9 – DC – T2 – BOOK 99. Explain what happens when the electron of a hydrogen atom changes from a 2s orbital to a 5s orbital. 100. Why is it not possible for two different orbitals to have the same first three quantum numbers? 101. The following four quantum numbers have been supplied for a bound electron. However, one of the values is incorrect. Explain which value is wrong and correct the error. n = 4; l = 3; ml = 4; ms = –1/2 102. The following graphs represent electron probabilities versus average distance from the nucleus. Shown here are the graphs and figures that represent (not necessarily in order) 1s, 2s and 2p orbitals. Match the representations with their correct identities: (a) (b) (i) 1s is shown in choice a b c (ii) 2s is shown in choice a b c (iii) 2p is shown in choice a b c (c) How many orbitals are in each of the following sublevels? _________103. 4p sublevel _________105. 4f sublevel _________104. 3d sublevel _________106. 2s sublevel page 10 – DC – T2 – BOOK Section 10.3 Write the complete (not abbreviated) electron configurations for the following elements. 107. Sulfur: ______________________________________________________________________ 108. Potassium: ___________________________________________________________________ 109. Vanadium: ___________________________________________________________________ 110. Argon: ______________________________________________________________________ 111. Iron: ________________________________________________________________________ 112. Sodium: _____________________________________________________________________ 113. Chromium: ___________________________________________________________________ 114. Iodine: _______________________________________________________________________ 115. Calcium: ______________________________________________________________________ 116. Platinum: _____________________________________________________________________ 117. Tellurium: _____________________________________________________________________ 118. Radium: ______________________________________________________________________ 119. Sr2+: _________________________________________________________________________ 120. Br–: __________________________________________________________________________ 121. P3–: __________________________________________________________________________ 122. O2–: __________________________________________________________________________ 123. Rb+: __________________________________________________________________________ 124. Al3+ Identify the elements described below. 125. Contains a full third energy level 126. Contains the first p electron 127. Element that is a noble gas and has its outermost electrons in the fourth energy level 128. Has an electron configuration of 1s22s22p63s23p4 page 11 – DC – T2 – BOOK 129. Has an electron configuration of 1s22s22p63s23p64s23d104p65s14d5 130. The 3rd period element with a complete outermost energy level 131. The 5th period element containing only 1 s electron and no d electrons in the outermost energy level 132. The 5th period element containing only 3 4d electrons 133. The first element that fills electrons in the f sublevel Consider the following individual orbitals; they are drawn to scale and each orbital has n = 3: s orbital (1) p orbital (2) d orbital (3) d orbital (4) 134. Which orbital has the lowest energy? ________ 135. Which orbital(s) could hold a maximum of 2 electrons? ________ 136. Write a reasonable set of quantum numbers for an electron that would be in the orbital pictured in figure (2) above. 137. What is the maximum number of electrons that could have n = 3? ________ Sections 11.1 and 11.2 Fill in the Blank. Fill in the blank with the appropriate word or phrases. 138. Group 7A elements are known as the _____________________________________________. 139. Group 2A elements are known as the _____________________________________________. 140. Group 8A elements are known as the _____________________________________________. 141. Group 1A elements are known as the _____________________________________________. 142. Group A elements are known as the ______________________________________ elements. 143. Group B elements are known as the ______________________________________ elements. page 12 – DC – T2 – BOOK Identification. Identify the element based on the provided information. 144. The 5th period element containing only 3 4d electrons 145. The first element that fills electrons in the f sublevel 146. The first element that fills electrons in the d sublevel 147. Element that has six outer electrons on the second period of the periodic table of elements 148. Element that has five outer electrons on the fourth period of the periodic table of elements 149. Element that has an outer electron configuration of 3d54s1 150. Element that has an outer electron configuration of 3d104s2 151. Element that has an outer electron configuration of 3s23p6 152. Element that has an outer electron configuration of 3s1 153. Element that has an outer electron configuration of 6s26p1 154. The 6th period Group 5A element 155. The 2th period Group 7A element 156. Element that is an alkali metal and in the fifth period 157. Element that is an alkaline earth metal and is in the fourth period 158. Element that is a halogen and has its outermost electrons in the fifth energy level 159. Element that is a noble gas and has its outermost electrons in the fourth energy level 160. Has an electron configuration of 1s22s22p63s23p4 page 13 – DC – T2 – BOOK 161. Has an electron configuration of 1s22s22p63s23p64s23d104p65s14d5 162. The 3rd period element with a complete outermost energy level 163. The 5th period element containing only 1 s electron and no d electrons in the outermost energy level Short Answer. Write the shorthand electron configurations of the following ions. 164. Cu: __________________________________________________________________________ 165. Cr6+: _________________________________________________________________________ 166. Cr: __________________________________________________________________________ 167. Zn: __________________________________________________________________________ 168. Ag: __________________________________________________________________________ 169. Sr2+: _________________________________________________________________________ 170. P3–: _________________________________________________________________________ Section 11.3 Short Answer. Answer the following questions. For the following pairs of atoms, circle which one of each pair has the largest ionic radius. 171. Al B 172. S O 173. Br Cl 174. Na Al 175. O F For the following pairs of elements, circle which one of each pair has the greater electronegativity. 176. Ca Ga 177. Li O 178. S Cl 179. As Br 180. Br Cl For the following pairs of elements, circle which one of each pair has the greater ionization energy. 181. Ca Ba 182. Li O 183. S Cl 184. As 186. Would you expect a Cl– ion to be larger or smaller than an Mg2+ ion? Explain. 187. Which effect on atomic size is more significant: the nuclear charge or the energy level that electrons are filling? Explain. page 14 – DC – T2 – BOOK Br 185. Br Cl 188. How does the ionic radius of a typical anion compare with the radius for the corresponding neutral atom? Explain. 189. How does the ionic radius of a typical cation compare with the radius for the corresponding neutral atom? Explain. page 15 – DC – T2 – BOOK Unit Six Worksheet WS – DC – U6 Chapter Thirteen Drawing. Draw electron dot structures for the following molecules or polyatomic ions. 190. HBr (hydrobromic acid) 194. PF3 (phosphorus trifluoride) 191. C2H2 (ethyne, or acetylene) 195. NH4+ (ammonium ion) 192. ClO4– (chlorate ion) 196. SO3– (sulfite ion) 193. HCN (hydrogen cyanide) 197. Br2 (diatomic bromine) page 16 – DC – T2 – BOOK Short Answer. Answer the following questions. 198. How many resonance structures can be drawn for the carbonate ion? Show the structural formulas for each. 199. How many resonance structures can be drawn for the nitrite ion? Show the electron dot structures for each. Drawing. 200. Draw the electron dot structure for each molecule or polyatomic ion. Then identify the molecular geometry of the following molecule or polyatomic ion. NO3- (nitrate ion) 202. Geometry: ___________________ 201. CCl4 (carbon tetrachloride) Geometry: ___________________ MnO4- (permanganate ion) (HINT: Assume Mn has 7 valance e–.) 203. PO43- (phosphate ion) Geometry: ___________________ Geometry: ___________________ page 17 – DC – T2 – BOOK 204. 205. PO33- (phosphite ion) PI5 (phosphorus pentaiodide) 206. Geometry: ___________________ Geometry: ___________________ NH3 (ammonia) 207. Geometry: ___________________ Geometry: ___________________ CO2 (carbon dioxide) Short Answer. Answer the following questions. What type of bond – nonpolar covalent, polar covalent, or ionic – will form between each of the following pairs of atoms? 208. Na and O: _______________________________________________________ 209. Li and Cl: ________________________________________________________ 210. P and O: ________________________________________________________ 211. N and N: _______________________________________________________ 212. Al and Cl: ________________________________________________________ 213. O and F: ________________________________________________________ page 18 – DC – T2 – BOOK 214. Explain why most chemical bonds would be classified as either polar covalent or ionic. (HINT: Consider why most are NOT nonpolar covalent.) 215. Would you expect carbon monoxide to be a polar or nonpolar molecule, and is there any difference in polarity between carbon monoxide and carbon dioxide? Explain. Draw the structural formulas for each molecule and identify polar covalent bonds by assigning the slightly positive (+) and the slightly negative ( –) charges to each atom in each bond. Then identify the overall molecule as polar or nonpolar. 216. NH3 218. CF4 217. CCl4 219. HF 220. Which compound would you expect to have the higher melting point: OCl2 or CaCl2? Explain. page 19 – DC – T2 – BOOK Table Completion. Complete the following table. Symbol 221. Number of valence electrons in atom/ion 222. Electron configuration 223. 224. S2– Ca2+ Electron dot formula 225. 8 226. 1s22s22p63s23p6 227. 228. Na Rb+ 232. 229. 230. 231. 233. 234. Short Answer. Answer the following questions. How many electrons will each of the following elements gain or lose in forming an ion? State whether each is a cation or anion. Number of e– lost or gained: 235. ________ Anion/Cation? 238. _____________ Phosphorus: Number of e– lost or gained: 236. ________ Anion/Cation? 239. _____________ Number of e– lost or gained: 237. ________ Anion/Cation? 240. _____________ Strontium: Bromine: 241. What is the relationship between the group number of the representative elements and the number of valence electrons? page 20 – DC – T2 – BOOK 242. Why do metals tend to form cations while nonmetals tend to form anions? Unit Seven Worksheet WS – DC – U7 Chapter Fourteen Short Answer. Answer the following questions. 243. In your own words, explain what a hydrogen bond is. 244. Depict the hydrogen bonding between three water molecules in a drawing. 245. How is hydrogen bonding responsible for the high boiling point of water? 246. Explain how large bodies of water are able to moderate air temperature. Short Answer. Answer the following questions. 247. Explain why it gets warmer before it rains. page 21 – DC – T2 – BOOK 248. Explain why the density of ice at 0C is less than the density of liquid water at 0C. 249. Explain why water has a relatively high boiling point and heat of vaporization. 250. What is the difference between the structure of liquid water and the structure of ice? How does this explain why ice floats in water? 251. Explain why water has a high surface tension. Chapter Fifteen Identify the solute and solvent in a dilute aqueous solution of sodium hydroxide. 252. Solute: _______________________________________________________________ 253. Solvent: ______________________________________________________________ 254. Give an example of a polar molecular compound that dissolves in water and that is a nonelectrolyte. Which of the following compounds are soluble in water? Which are insoluble? 255. CaCl2: ____________________________________________ page 22 – DC – T2 – BOOK 256. N2: _______________________________________________ 257. HBr: ______________________________________________ 258. NH4C2H3O2: ________________________________________ Write equations to show how the following compounds dissociate in water. 259. NH4NO3(s) 260. K2SO4(s) Write the formulas for the following hydrates. 261. Calcium sulfate decahydrate: ________________________________________________ 262. Cobalt (II) chloride hexahydrate: ______________________________________________ 263. Find the percent by mass of water in NiCl2 6H2O. 264. Why is using water to clean a paintbrush covered with oil-based enamel not an effective cleanup method? 265. How can a supersaturated solution be prepared? 266. You are given a clear aqueous solution containing potassium nitrate (KNO3). How would you determine experimentally if the solution is unsaturated, saturated, or supersaturated? page 23 – DC – T2 – BOOK Short Answer. Write complete balanced net ionic equations based on the following chemicals reacting. 267. Lead (II) nitrate and sulfuric acid 268. Sodium phosphate and iron (III) chloride 269. Ammonium sulfuide and cobalt (II) nitrate 270. Sulfuric acid and barium chloride 271. Aluminum sulfate and ammonium hydroxide 272. Silver nitrate and dihydrogen sulfide 273. Calcium chloride and lead (II) nitrate 274. Calcium nitrate and sodium carbonate 275. Hydrochloric acid and barium hydroxide 276. Iron (III) nitrate and sodium hydroxide page 24 – DC – T2 – BOOK 277. What are colligative properties of solutions? Give examples of three types of colligative properties. 278. How many particles in solution are produced by each formula unit of potassium carbonate, K2CO3? 279. How many moles of particles would 3 mol Na2SO4 give in solution? 280. What kind of property is vapor-pressure lowering? An equal number of moles of NaCl and CaCl2 are dissolved in equal volumes of water. Which solution has the lower 281. freezing point? _______________________________________________________ 282. vapor pressure? ______________________________________________________ 283. boiling point? _________________________________________________________ 284. Why does a solution have an elevated boiling point and a depressed freezing point compared with the pure solvent? Section 6.1 Short Answer. Answer the following questions. 285. When you talk about the volume of a gas, are you referring to the volume of the molecules themselves? Explain. 286. What is the difference between a pressure and a force? 287. What happens to gas particles when a gas is compressed? page 25 – DC – T2 – BOOK 288. Explain how a mercury barometer works. 289. If you collect a gas so that it completely fills a 250 cm3 Erlenmeyer flask, the volume of the gas is actually greater than 250 cm3. Explain why this is true, and explain how you could determine what the volume of the gas is. Calculations. Solve the following problems. Show work and appropriate significant figures and/or decimal places. Hydrogen gas is collected by bubbling it through water. Calculate the partial pressure of the hydrogen gas if: 290. The total pressure is 94000 Pa and the partial pressure of water is 1200 Pa. 291. The total pressure is 100.3 kPa and the partial pressure of water is 2600 Pa. 292. In a flask that has a volume of 273 dm3, you have a sample of two noble gases: neon and xenon. The partial pressure of the neon is 96950 Pa, and the partial pressure of the xenon is 1.025 atm. What is the total pressure (in kPa) exerted by these two gases? Make the following conversions. Show work and appropriate significant figures and/or decimal places. 293. 105 Pa = __________ torr 295. 256.7 mm Hg = __________ atm 294. 12 kPa = __________ psi 296. 285 torr = __________ kPa page 26 – DC – T2 – BOOK Section 6.2 Calculations. Solve the following problems. Show work and appropriate significant figures and/or decimal places. 297. A metal cylinder contains 1 mol of oxygen gas at STP. What will happen to the pressure if another mole of gas is added to the cylinder, but the temperature and volume do not change? 298. If a gas is compressed from 10 L to 1 L, and the temperature remains constant, what happens to the pressure? 299. A truck driver gets paid according to the quantity of a certain kind of gas he can deliver. The more gas he delivers, the more he gets paid. He only has time for one trip, and the dimensions of his tanker are fixed. What properties of gases can he exploit to increase his profit? 300. The gas in a closed container has a pressure of 3.00 x 102 kPa at 30C. What will the pressure be if the temperature is lowered to –172C? 301. Calculate the volume of a gas (in L) at a pressure of 1.00 x 102 kPa if its volume at 1.20 x 102 kPa is 1.50 x 103 mm3. 302. A gas with a volume of 3.00 x 102 mL at 150.0C is heated until its volume is 6.00 x 102 mL. What is the new temperature of the gas if the pressure remains constant during the heating process? 303. A given mass of air has a volume of 6.00 L at 101 kPa. What volume, in cubic centimeters, will it occupy at 25.0 kPa if the temperature does not change? page 27 – DC – T2 – BOOK 304. A 3.50-L gas sample at 20C and a pressure of 86.7 kPa expands to a volume of 8.00 L. The final pressure of the gas is 56.7 kPa. What is the final temperature of the gas, in degrees Celsius? 305. What is the volume of a sample of carbon dioxide at standard temperature and pressure that has a volume of 75.0 mL at 30.0C and 91 kPa? 306. Which of the following samples of gases occupies the largest volume, assuming that each sample is at the same temperature and pressure – 50.0 g neon, 50.0 g argon, or 50.0 g xenon? 307. What volume of carbon dioxide gas contains the same number of oxygen atoms as 250.0 cm3 of carbon monoxide gas, if each gas sample is measured at the same temperature and pressure? Given the following data, Volume of Nitrogen Gas (L) 4.28 L 5.79 L 308. Temperature (K) 303 K 410 K Draw a graph of the relationship between volume and temperature. page 28 – DC – T2 – BOOK 309. Determine the slope of the line. 310. Find the slope of a line in relationship to the temperature-volume law expressed as V/T = k. 311. Calculate the expected volume of the gas when the temperature reaches 1200 K. Section 7.1 Calculations. Solve the following problems. Show work and appropriate significant figures and/or decimal places. 312. If 4.50 g of methane gas (CH4) is introduced into an evacuated 2.00-L container at 35C, what is the pressure in the container? 313. Calculate the number of moles of oxygen gas in a 12.5 L tank if the pressure is 253 kPa and the temperature is 22C. 314. Calculate the mass of nitrogen dioxide present in a 275 mL container if the pressure is 240.0 kPa and the temperature is 28C. 315. What is the density of nitrogen dioxide given the conditions of Problem 30? page 29 – DC – T2 – BOOK 316. During the metabolic process called respiration, your body obtains energy from the breakdown of glucose as shown below. C6H12O6(aq) + 6O2(g) 6H2O(l) + 6CO2(g) What volume of O2, measured at 37C and 790.0 torr pressure, is required to react with 1.00 g of glucose (C6H12O6)? Express the volume in cubic centimeters. 317. Magnesium reacts with oxygen gas to produce magnesium oxide. What volume of oxygen gas, in cubic centimeters, is required to fully react with 20.2 g of magnesium metal? The reaction is taking place at 1.02 atm pressure and 24C. 318. Magnesium reacts with hydrochloric acid to produce hydrogen gas and magnesium chloride. What volume of hydrogen gas is theoretically produced, in cubic millimeters, if 4.9 g magnesium reacts with excess hydrochloric acid? The reaction is taking place at 101.5 kPa pressure and 20C. 319. A 3.50-L gas sample at 20C and a pressure of 86.7 kPa expands to a volume of 8.00 L. The final pressure of the gas is 56.7 kPa. What is the final temperature of the gas, in degrees Celsius? page 30 – DC – T2 – BOOK 320. What is the volume of a sample of carbon dioxide at standard temperature and pressure that has a volume of 75.0 mL at 30.0C and 91 kPa? 321. A truck driver gets paid according to the quantity of a certain kind of gas he can deliver. The more gas he delivers, the more he gets paid. He only has time for one trip, and the dimensions of his tanker are fixed. What properties of gases can he exploit to increase his profit? 322. The gas in a closed container has a pressure of 3.00 x 102 kPa at 30C. What will the pressure be if the temperature is lowered to –172C? 323. Calculate the volume of a gas (in L) at a pressure of 1.00 x 102 kPa if its volume at 1.20 x 102 kPa is 1.50 x 103 mm3. 324. A gas with a volume of 3.00 x 102 mL at 150.0C is heated until its volume is 6.00 x 102 mL. What is the new temperature of the gas if the pressure remains constant during the heating process? 325. A given mass of air has a volume of 6.00 L at 101 kPa. What volume, in cubic centimeters, will it occupy at 25.0 kPa if the temperature does not change? page 31 – DC – T2 – BOOK 326. A 3.50-L gas sample at 20C and a pressure of 86.7 kPa expands to a volume of 8.00 L. The final pressure of the gas is 56.7 kPa. What is the final temperature of the gas, in degrees Celsius? 327. What is the volume of a sample of carbon dioxide at standard temperature and pressure that has a volume of 75.0 mL at 30.0C and 91 kPa? Section 7.2 Calculations. Solve the following problems. Show work and appropriate significant figures and/or decimal places. Calculate the number of liters occupied at STP. 328. 2.5 mol N2(g) 329. 0.600 g H2(g) 330. 2.8 g CO2(g) 331. 2.8 x 1021 molecules CO2(g) page 32 – DC – T2 – BOOK 332. Which of the following samples of gases occupies the largest volume, assuming that each sample is at the same temperature and pressure – 50.0 g neon, 50.0 g argon, or 50.0 g xenon? 333. What volume of carbon dioxide gas contains the same number of oxygen atoms as 250.0 cm3 of carbon monoxide gas, if each gas sample is measured at the same temperature and pressure? Section 7.3 Calculations. Solve the following problems. Show work and appropriate significant figures and/or decimal places. 334. Calculate the ratio of the velocity of helium atoms to the velocity of neon atoms at the same temperature. 335. A certain gas effuses four times as fast as oxygen (O2). What is the molar mass of the gas? 336. During an effusion experiment, it took 75 seconds for a certain number of moles of an unknown gas to pass through a tiny hole. Under the same conditions, the same number of moles of oxygen gas passed through the hole in 30 seconds. What is the molar mass of the unknown gas? page 33 – DC – T2 – BOOK 337. At the same temperature and pressure, which gas moves faster: oxygen or neon? How many times is the speed of the faster gas greater than the slower gas? 338. In an experiment, it takes an unknown gas 1.5 times longer to diffuse than the same amount of oxygen gas. Find the molar mass of the unknown gas. Chapter 16 Calculations. Solve the following problems. Show work and appropriate significant figures and/or decimal places. 339. How many kilojoules of energy are in a donut that contains 205.0 Calories? 340. What is the specific heat of a substance that has a mass of 25.0 g and requires 525.0 calories to raise its temperature by 15.0 K? 341. Suppose 0.20 kg of ice absorbs 125.0 J of heat. What is the corresponding temperature change? The specific heat capacity of H2O(s) is 2.1 J/(gC). 342. How many joules of heat energy are required to raise the temperature of 100.0 g of aluminum by 120.0C? The specific heat capacity of aluminum is 0.90 J/(gC). page 34 – DC – T2 – BOOK 343. A student mixed 75.0 mL of water containing 0.75 mol HCl at 25C with 75.0 mL of water containing 0.75 mol of NaOH at 25C in a foam cup calorimeter. The temperature of the resulting solution increased to 35C. How much heat in kilojoules was released by this reaction? CH2O = 4.18 J/(gC) 344. Calculate the amount of heat evolved when 15.0 g of Ca(OH)2 forms from the reaction of CaO(s) + H2O(l). CaO(s) + H2O(l) Ca(OH)2(s) H = – 65.2 kJ 345. Calculate the amount of heat produced when 52.4 g of methane, CH4, burns in an excess of air, according to the following equation. CH4(g) + 2O2(g) CO2(g) + 2H2O(l) H = – 890.2 kJ 346. When 2 moles of nitric oxide, NO, burn in air to produce 2 moles of nitrogen dioxide, 113.04 kJ of heat is produced. Write a balanced thermochemical equation for this reaction. 347. Calculate the amount of heat needed to melt 35.0 g of ice at 0C. Express your answer in kilojoules. 348. Calculate the amount of heat needed to convert 1.0 kg of liquid water at 15C to steam at 100C. page 35 – DC – T2 – BOOK 349. Calculate the amount of heat needed to convert 96 g of ice at – 24C to water at 100C. The specific heat capacity of H2O(s) is 2.1 J/(gC). 350. How much heat is absorbed when 28.3 g of H2O(s) at 0C is converted to liquid at 0C? The specific heat capacity of H2O(s) is 2.1 J/(gC). 351. When 47.5 g of a metal at 425.0C is dropped into 1000.0 g of water at 18.0C, the final temperature of the metal and water is 21.0C. The specific heat of water is 4.184 J/(gC). What is the specific heat of the metal in J/(gC)? 352. In a calorimetry experiment, you collected the following data: Classification Mass of metal Initial temperature of water in cup in degrees Celsius Initial temperature of metal in degrees Celsius (temperature of boiling water) Maximum temperature of metal and water Mass of water Measurement 100.0 g 20.0C 120.0C 30.0C 150.0 g What is the specific heat of the metal in the experiment in J/(gC)? Short Answer. Answer the following questions. State whether the following physical and chemical changes are endothermic or exothermic. 353. Melting 357. Vaporization 354. Condensation 358. Fusion 355. Freezing 359. Combustion 356. Sublimation 360. Evaporation page 36 – DC – T2 – BOOK 361. If a reaction has H < 0, what kind of reaction occurs? Explain? 362. What is meant by T? Explain two ways of how one can calculate it. 363. The same quantity of heat is added to an iron nail (3.5 g) and to a metric ton (1000 kg) of iron. Which would reach the higher temperature? Explain. 364. Which has more entropy: 1 g of liquid water or 1 g of steam? Explain. 365. Which has more entropy: 1 g of liquid mercury or 1 g of solid mercury? Explain. 366. Analogize entropy and enthalpy in terms of the favorability of a reaction. page 37 – DC – T2 – BOOK Unit Eight Worksheet WS – DC – U8 Chapter Seventeen Calculation. Answer the following problems. Show work or receive no credit. Show proper units. 367. An ice machine can produce 120 kg of ice in 24 hours. Express the rate of ice production in kg/hr. 368. Ethyl acetate (C4H8O2) reacts with a solution of sodium hydroxide (NaOH) in water to form sodium acetate (C2H3O2Na) and ethyl alcohol (C2H6O). Suppose at 25C two moles of ethyl acetate react completely in four hours. How would you express the rate of reaction? Short Answer. Answer the following questions. 369. A friend tells you that you can recognize a fast reaction because it produces more product than a slow reaction. What other factors must be included to make this a correct statement? 370. Ethyl acetate and water are not miscible; thus, the reaction in problem 368 only occurs at the interface of the two liquids. What would be the effect on the reaction rate by adding a solvent to make the reaction homogeneous? 371. What reactant or product would you choose to measure in order to determine the rate of reaction for the following chemical reaction? Zn(s) + 2HCl(aq) H2(g) + ZnCl2(aq) Explain how you would measure the substance you chose. page 38 – DC – T2 – BOOK 372. What reactant or product would you choose to measure in order to determine the rate of reaction for the following chemical reaction? Cu(s) + 2AgNO3(aq) 2Ag(s) + Cu(NO3)2(aq) Explain how you would measure the substance you chose. How would the following actions likely change the rate of reaction in problem 368? 373. the temperature is lowered to 4C 374. the concentration of sodium hydroxide in water is increased List three ways that reaction rates can generally be increased. 375. __________________________________________________________________ 376. __________________________________________________________________ 377. __________________________________________________________________ 378. Explain how a catalyst works as correlated to activation energy and an activated complex. Zinc reacts with hydrochloric acid to produce zinc chloride and hydrogen gas. Predict which reaction will occur at a faster rate based on the following circumstances and explain why. Assume all other factors aside from those specified are the same. 379. Reaction 1: 1.00 M hydrochloric acid is used Reaction 2: 6.00 M hydrochloric acid is used 380. Reaction 1: Chunks of zinc are used Reaction 2: Powdered zinc is used 381. Reaction 1: The reaction occurs at 75C. Reaction 2: The reaction occurs at 25C. page 39 – DC – T2 – BOOK Chapter Eighteen Calculation. Answer the following problems. Show work or receive no credit. Show proper units. 382. Write the expression for the equilibrium constant for the following reaction. 2N2O5(g) 4NO2(g) + O2(g) 383. Calculate the equilibrium constant for the reaction in problem 382 if the equilibrium concentrations are [N2O5] = 0.50 mol/L, [NO2] = 0.80 mol/L, and [O2] = 0.20 mol/L. Short Answer. Answer the following questions. How would the equilibrium position for the equation in problem 382 be affected by the following? 384. an addition of O2 to the reaction vessel 385. a decrease in the pressure Write the equilibrium constant expression for each of the following reactions. 386. 4NO(g) + 2O2(g) 2N2O4(g) 387. SO2(g) + NO2(g) SO3(g) + NO(g) page 40 – DC – T2 – BOOK 388. Can a pressure change shift the equilibrium position in every reversible reaction? Explain your answer. Short Answer. Answer the following questions. 389. What two factors determine whether a reaction is spontaneous? 390. Where does lost free energy typically end up? Does free energy lost as heat ever serve a useful function? Explain. 391. What can change a reaction from nonspontaneous to spontaneous? 392. Suppose the products in a spontaneous process are more ordered than the reactants. Is the entropy change favorable or unfavorable? Explain. 393. How can an equilibrium constant be used to determine the favorability of a reaction? 394. Why are solids not something to be considered when calculating an equilibrium constant? page 41 – DC – T2 – BOOK Short Answer. Answer the following questions. Classify each of the following acids as monoprotic, diprotic, or triprotic. 395. HCOOH: _____________________________________________________________ 396. HBr: ________________________________________________________________ 397. H2SO3: ______________________________________________________________ 398. HClO4: _____________________________________________________________ 399. What would you expect to happen when lithium metal is added to water? Show the chemical reaction. 400. Identify the hydrogen ion donor(s) and hydrogen ion acceptor(s) for ionization of sulfuric acid in water. Label the conjugate acid-base pairs. 401. Identify all of the ions that may be formed when H3PO4 ionizes in water. Matching. Match each solution with its correct description. _____402. dilute, weak acid _____403. dilute, strong base _____404. concentrated, strong acid A) B) C) D) E) 18M H2SO4(aq) 0.5M NaOH(aq) 15M NH3(aq) 0.1M HC2H3O2(aq) 0.1M HCl(aq) _____405. dilute, strong acid _____406. concentrated, weak base page 42 – DC – T2 – BOOK Short Answer. Answer the following questions. 407. Write the expression for the base dissociation constant for hydrazine, N2H4, a weak base. Hydrazine react with water to form the N2H5+ ion. 408. Write the base dissociation constant expression for the weak base analine, C6H5NH2. C6H5NH2(aq) + H2O(l) C6H5NH3+(aq) + OH–(aq) Short Answer. Answer the following questions. Write, complete, and balance the equations for the following acid-base reactions, including states of matter. 409. Phosphoric acid + aluminum hydroxide 410. Hydroiodic acid + calcium hydroxide 411. Nitric acid + sodium hydroxide page 43 – DC – T2 – BOOK Calculations. Solve the following problems. Show work with correct significant figures and/or decimal places. Include proper units. 412. What is the molarity of a sodium hydroxide solution if 38 mL of the solution is titrated to the end point with 14 mL of 0.75M sulfuric acid? 413. If 24.6 mL of a calcium hydroxide solution are needed to neutralize 14.2 mL of 0.0140M acetic acid, what is the concentration of the calcium hydroxide solution? 414. A 12.4 mL solution of sulfuric acid is completely neutralized by 19.8 mL of 0.0100M calcium hydroxide. What is the concentration of the sulfuric acid? 415. What volume of 0.12M barium hydroxide is needed to neutralize 12.2 mL of 0.25M hydrochloric acid? 416. A 55.0 mg sample of aluminum hydroxide is reacted with 0.200M hydrochloric acid. How many milliliters of the acid are needed to neutralize the aluminum hydroxide? page 44 – DC – T2 – BOOK Unit Five Experiment – 1 Orbital Structures and Identifications EX – DC – U5 – 1 Introduction: The purpose of this simulation is to simulate and identify different kinds of atomic orbitals. Background: An atomic orbital is a region of space around the nucleus of an atom where there is a high probability of finding an electron. The means by which the probability of finding an electron was determined is known as the Schrödinger equation. Although it is a rather complex equation, it can be solved exactly for the hydrogen atom. There are four quantum numbers typically associated with an atomic orbital and used by many chemists and physicists with regard to atomic orbitals, and these are actually used as well in the “Atom in a Box” software. They are as follows: Principal Energy Level (n) This number corresponds to what we have called the principal energy level in our class and also corresponds with the average distance of electron from the nucleus. We often ascribe the principal energy level to be synonymous with the period of the element whose electrons we are trying to convey. Value of n = 1, 2, 3, 4, and so forth. Sublevels (l) The general shape of the electron cloud associated with an atomic orbital is called the sublevel. n and l are related in that l = 0, 1, 2, …, (n – 1). Typically, we can use the following quantum numbers represented by l to indicate the type of sublevel (s, p, d, and f) we are talking about. Quantum number (l) Type of sublevel 0 s 1 p 2 d 3 f Orbitals (ml) This quantum number refers to the number of orbitals possible given which type of atomic orbital is being discussed. We know that an s orbital only has 1 type of orbital whereas a p orbital has 3 different types (x, y, and z). If you know the value of l, indicating the type of orbital, you can choose specifically what kind of orbital since ml, = l, …, +1, 0, -1, …, -l. Using this, one can determine that and s orbital has only one kind of orbital, p has three, d has five, and f has seven. You will need to be familiar with these three quantum numbers as they are in the program you will be using. These values can be seen at the top of the screen and can be modified by pressing the plus (+) or minus (-) sign above the respective values of n, l, and ml. This can be seen in the following screen shot: page 45 – DC – T2 – BOOK Procedure: 1. Follow the directions provided by your instructor for how to load the “Atom in a Box” software. 2. For the first atomic orbital, select n = 1. Record in the data table what kind of an orbital this is as well as its principal energy level. 3. Now change n = 3. Set the value of l = 1. Record in the data table what kind of orbital this is as well as its principal energy level. 4. Change n = 5. Set the value of l = 3 and m = 0. Record in the data table what kind of orbital this is as well as its principal energy level. Then record whether or not it has any nodes. 5. Change the value of m to 2. If you observe and noticeable changes, record them accordingly. 6. Change n = 7. Set the value of l = 3 and m = 0. Record in the data table what kind of orbital this is as well as its principal energy level. Does this kind of orbital have any observable rings? (Rotate if necessary to investigate.) If so, record accordingly. 7. Set the value of l = 6 and m = 0. Record in the data table what kind of orbital this is as well as its principal energy level. Does this kind of orbital have more or less observable rings than the previous orbital? Record accordingly. 8. Experiment with the program to help you answer the questions on the response page. page 46 – DC – T2 – BOOK Unit Five Experiment – 1 – R Orbital Structures and Identifications EX – DC – U5 – 1 – R Procedure question n l m Type of orbital (i.e., s, p, d, f) Principal energy level Answer to (i.e, 1, 2, 3) question (if asked) 2 3 Nodes? 4 Noticeable changes? 5 How many rings? 6 More or less rings? How many? 7 Questions: Identify the quantum numbers for the following atomic orbitals. You may use the Atom in a Box software to help you identify them. 1. 2. page 47 – DC – T2 – BOOK 3. 4. 5. What do these orbitals fully represent? 6. How many different kinds of orbitals are possible in this program given the available information? ***BONUS*** Identify the quantum numbers associated with the following orbital. Answers: 1. n = ____________ l = ____________ 3. n = ____________ l = ____________ 2. n = ____________ l = ____________ 4. n = ____________ l = ____________ 5. _________________________________________________________________________ _________________________________________________________________________ 6. _________________________________________________________________________ BONUS: n = ________________ l = ________________ page 48 – DC – T2 – BOOK Unit Five Experiment – 2 Metal Ions and Flame Tests EX – DC – U5 – 2 Introduction: The purpose of this experiment is to determine the cations of a solution based on flame tests. Background: When a metal or metal salt is added to a flame, a combustion reaction ensues. This reaction excites an electron in the metal from its ground state to a higher orbital. In order to return to its ground state, the electron releases the additional energy in the form of light. Different metal electrons emit different wavelengths of light to return to their respective ground states, so the flame colors are varied. These flames can be used to produce atomic emission spectra of the elements combusted. Using known values of emission spectra, one can perform a flame test on un unknown substance, gather an emission spectrum from it, and determine which elements are in the unknown substance. Procedure: 1. At each lab station, take a splint and tap off excess fluid from the splint. Hold it over the flame and observe the color produced. Record your findings in the table. Discard the used splint in the appropriate place. 2. Repeat step 1 for each additional station. Lab Station Color Produced Metal 1 2 3 4 5 6 Your instructor will provide you data about the metal tests following the experiment. Conclusion/Discussion: 1. Did you observe certain colors consistent with certain groups/families of metals? Explain. 2. How do you think scientists use this knowledge today? 3. How are the results you found related to the photoelectric effect? page 49 – DC – T2 – BOOK Unit Five Experiment – 3 Viewing Spectra EX – DC – U5 – 2 Introduction: The purpose of this experiment is to examine the spectra produced when you view a variety of light sources through a diffraction grating in a spectrometer. Background: A diffraction grating can break up light from a source into its component colors, much the same as a prism can break up sunlight into the colors of the rainbow. In the rectangles below, use crayons to sketch the spectrum you see from each source. -9 The numbers below the rectangles represent the wavelength of the light in nanometers (x 10 m) After you have made your observations, answer the Conclusion/Discussion Questions. Procedure: 1. View an incandescent light bulb through a diffraction grating. Inside the bulb a tungsten filament is heated until it glows. 2. Now view the spectra of three gas discharge tubes. In these tubes, atoms in a low-pressure gas are excited by being bombarded by a stream of high-energy electrons. Record which gas is in the discharge tube you observe. page 50 – DC – T2 – BOOK Conclusion/Discussion: 1. Note how all the colors in the spectrum from the incandescent bulb "run into" one another. We call this a continuous spectrum. Is there evidence to suggest that the incandescent bulb emits light in regions of the spectrum other than the visible? Explain. 2. At which end of the spectrum (blue or red), does light transfer more energy? Explain how you know in terms of the relationships we have discussed this far. 3. How do the spectra produced by the excited gases differ from that produced by the hot metal filament? Do the atoms emit all frequencies of visible light? 4. Determine, as best you can, the wavelength of the red line in the hydrogen spectrum. Calculate the frequency of this light. Determine the energy of the photons of light emitted at this frequency. Show your work clearly. page 51 – DC – T2 – BOOK Unit Six Experiment – 1 Paper Chromatography of Food Dyes EX – DC – U6 – 1 Purpose: The purpose of this experiment is to use paper chromatography to separate and identify food dyes in various samples. Safety: Safety goggles will be worn at all times. During this experiment, no open-toe shoes are to be worn. Procedure: 1. Cut a 5 cm x 10 cm strip of chromatography paper and label it with a pencil as shown in Figure A. Food Color Samples 0.1% NaCl solution Your name X X X X Red Yellow Green Blue 2. Use a different toothpick to place a spot of each of the four food colors on the Xs on the chromatography paper. Allow the spots to dry for a few minutes. 3. Fill the 250-mL beaker so its bottom is just covered with the solvent (0.1% NaCl solution). Wrap the chromatography paper around a pencil. Remove the pencil and place the chromatography paper, color-spot side down, in the solvent. When the solvent reaches the top of the chromatography paper, remove the paper and allow it to dry. page 52 – DC – T2 – BOOK Observation/Data Tables: Record observations from the laboratory experiment below. Observations: ________________________________________________________________________________ ________________________________________________________________________________ ________________________________________________________________________________ ________________________________________________________________________________ ________________________________________________________________________________ ________________________________________________________________________________ Conclusion/Discussion: 1. If a food color yields a single streak or spot, it is usually a pure compound. Which food colors consist of pure compounds? 2. Which food colors are mixtures of compounds? 3. Food colors often consist of a mixture of three colored dyes: Red No. 40, Yellow No. 5, and Blue No. 1. Read the label on the food color package. Which dyes do your food color samples contain? 4. Identify each spot or streak on your chromatogram as Red No. 40, Yellow No. 5, or Blue No. 1. 5. Paper chromatography separates polar covalent compounds on the basis of their relative polarities. The most polar dyes migrate the fastest and appear at the top of the paper. Which dye is the most polar? Which dye is the least polar? page 53 – DC – T2 – BOOK Unit Seven Experiment – 1 The Ideal Gas Law EX – DC – U7 – 1 Purpose: The purpose of this experiment is to use the ideal gas law, understand the variables involved when working with gases, and to experimentally find the molar mass of a gas. Background: Most gas experiments do not occur at standard temperature and pressure. Therefore, we need to use a calculation that allows us to account for changes in pressure and temperature. The ideal gas law allows us to achieve this. The ideal gas law is PV=nRT where P is the pressure of the gas, V is the volume of the gas, n is the number of moles of the gas, R is the ideal gas constant (found on the back of your periodic table for different pressure units), and T is the temperature of gas. You will need to know the atmospheric pressure when you perform this lab. Go to http://www.wunderground.com and place in your zip code to obtain the atmospheric pressure at your location. Pressure will be given in inches of mercury (in Hg). This will need to be converted to different units of pressure. This will be done later. Safety: Safety goggles will be worn at all times. Never shake or tilt the can of compressed air before or during usage. Never use the gas or canister around a possible ignition source. Avoid contact with skin. DO NOT WASTE ANY OF THE COMPRESSED GAS!!! USE ONLY THE GAS REQUIRED!!! Violations of ANY of these safety provisions will result in reduced credit and/or removal from the laboratory. Procedure: 1. Create a data table to record the data for the mass, temperature, pressure and volume of the gas used in this experiment. 2. Fill a tray about 2/3 full of water. 3. Obtain a can of compressed air. Mass the can and record the mass. Also, record the mass when done with the experiment, meaning do NOT use more air than required, or your readings will be offset significantly. Subtract the masses to get the mass of the gas used. 4. Fill a 250 flask with water. Place a watch glass on top of the flask. Turn the flask upside down making sure that there are no bubbles present. Place the flask in the water and remove the watch glass. 5. Holding the flask by the neck, place some glass tubing into the flask. Insert the straw from the compressed air canister into the tubing and seal with some clay. 6. Slowly fill the flask with gas using the can of compressed air by spraying it into the glass tubing. To help insure the clay seal is tight, you may wish to try and hold it in place. 7. Continue to hold down the lever until the gas level almost matches the water level in the tray. Remove the glass tubing. page 54 – DC – T2 – BOOK 8. Adjust the flask so that the water level inside the flask is the same level as the outside water making sure to keep the flask upside down. Using a grease pen or sharpie, mark a line on the flask to record the water level. 9. Remove the flask. Fill the flask with water to the pen mark. Mass a 250 mL beaker. Record the mass. Pour the water into the beaker and mass the water and the beaker. Subtract out the empty beaker mass to get the mass of the water. Since the density of water is 1 g/mL, the mass of the water will equal the volume. 10. Measure and record the temperature of the water bath. Conclusion/Discussion: 1. The atmospheric pressure you recorded earlier was in inches of Mercury. This could be converted to many different units and the units you use for pressure affect what your R value for the ideal gas law will be. Convert your pressure to mm Hg. The conversion factors is 1 in = 25.4 mm. 2. Question 1 reflects the total pressure inside the flask. You know this because when your water levels were equal, the pressure inside and outside the flask were equal. Since your gas was collected through water, some of the pressure in the flask is water vapor. Use your text and read about Dalton’s law of partial pressures. Then use the chart below to calculate the pressure exerted by the gas from the can. Temperature (°C) 0.0 5.0 10.0 12.5 15.0 15.5 16.0 16.5 17.0 17.5 18.0 18.5 19.9 Pressure (mm Hg) 4.6 6.5 9.2 10.9 12.8 13.2 13.6 14.1 14.5 15.0 15.5 16.0 16.5 Water Vapor Pressure Table Temperature Pressure Temperature (°C) (mm Hg) (°C) 27.0 19.5 17.0 28.0 20.0 17.5 29.0 20.5 18.1 30.0 21.0 18.6 35.0 21.5 19.2 40.0 22.0 19.8 50.0 22.5 20.4 60.0 23.0 21.1 70.0 23.5 21.7 80.0 24.0 22.4 90.0 24.5 23.1 95.0 25.0 23.8 100.0 26.0 25.2 Pressure (mm Hg) 26.7 28.3 30.0 31.8 42.2 55.3 92.5 149.4 233.7 355.1 525.8 633.9 760.0 4. Since you now know the pressure, volume, and temperature of the gas, given that the R value is 62.396 (Lmm Hg)/(Kmol), solve for the number of moles of gas using PV = nRT. (Hints: You probably measured your volume of gas in milliliter (mL). Make sure and convert this to liters by dividing by 1000 since 1 L = 1000 mL. Also, your temperature was measured in Celsius, so convert it to Kelvins. The formula for converting Celsius to Kelvin is K = oC + 273.15. 5. Now that you know the number of moles of gas collected and the mass of that gas, calculate the molar mass of the compressed gas. 6. Compare your answer to question 5 with the actual molar mass of difluoroethane. The formula for difluoroethane is C2H4F2. Calculate your percent error. page 55 – DC – T2 – BOOK Unit Seven Experiment – 2 Specific Heat Experiment EX – DC – U7 – 2 Introduction: The purpose of this experiment is to determine the specific heat of a substance. Background: On a sunny day, the water in a swimming pool may warm up a degree or two while the concrete around the pool may become too hot to walk on in your bare feet. This may seem strange because both the concrete and the water are being heated by the same source: the sun. This evidence suggests it takes more heat to raise the temperature of some substances than others. This, in fact, is true: the amount of heat that is required to raise the temperature of 1 g of a substance by 1 degree Celsius is called the specific heat capacity, or simply the specific heat, of that substance. Water, for instance, has a specific heat of 4.184 J/(gC). This value is high in comparison with the specific heats for other materials, such as concrete. In this experiment, you will use a simple calorimeter and your knowledge of the specific heat of water to determine the specific heat of an unknown metal. Safety: Safety goggles will be worn at all times. During this experiment, no open-toe shoes are to be worn. Since Bunsen burners will be used, do not handle hot equipment with only your hands. Use proper protective equipment. Additionally, never use a thermometer as a stirrer. Procedure: As you perform the experiment, record your data in Data Table 1. 1. Measure the mass of the metal cylinder provided by your instructor to the nearest 0.01 g and record the measurement. Transfer the cylinder to a large, dry test tube and use a utility clamp to suspend the test tube in about 400 mL of water in a 600 mL beaker. (If you do not have access to a 600 mL beaker, that is fine. Use a beaker where the metal in the test tube you will be using can be below water level, meaning you may even need more or less than 400 mL of water.) Heat the water until it is boiling gently, and leave the test tube in the boiling water bath for at least 10 minutes. 2. While the cylinder is heating, measure approximately 100 mL of distilled water in a graduated cylinder. Assume the density of water is 1 g/mL, meaning that if you get 98.2 mL of distilled water, you have 98.2 g of water. Record this as the mass of water to one decimal place. Then pour the water into a plastic-foam cup. 3. Measure and record the temperature of the water in the plastic-foam cup and of the water in the boiling bath. This temperature should be recorded to one decimal place. 4. Remove the test tube from the boiling water and quickly pour the cylinder into the water-filled, plastic-foam cup. Place a thermometer and a glass stirring rod into the cup. Use a stirring rod to gently stir the cylinder. Do not stir the cylinder with the thermometer. Note the temperature frequently and record the maximum temperature reached. This temperature should be recorded to one decimal place. page 56 – DC – T2 – BOOK 5. Pour the water off, dry the cylinder, and repeat steps 1 – 4. After a second trial, repeat step 5, but instead of conducting another trial, return the metal cylinder to your teacher. Observations: Data Table I: Measurements of Mass and Temperature Trial 1 Measurement Classification Trial 2 Measurement Mass of cylinder Initial temperature of water in cup in degrees Celsius Initial temperature of metal cylinder in degrees Celsius (temperature of boiling water) Maximum temperature of cylinder and water Mass of water Data Table II: Specific Heat Capacities of Common Substances Substance Water Ice Steam Stainless steel Iron Aluminum Copper Brass Gold Lead Carbon (graphite) Specific heat capacity J/(gC) 4.184 2.06 1.87 0.927 0.449 0.897 0.385 0.376 0.129 0.129 0.709 Conclusion/Discussion: Answer these questions on a separate sheet of paper. 1. Determine the changes in temperature of the water and of the cylinder for each trial (in other words, TH2O (trial 1) , TH2O (trial 2) , TMetal (trial 1) , and TMetal (trial 2) ). 2. Calculate the heat gained by the water in each trial. page 57 – DC – T2 – BOOK 3. Remembering that the heat gained by the water is equal to the heat lost by the cylinder, calculate the specific heat of the cylinder. 4. Calculate the average value for the specific heat of the cylinder in your experiment. 5. Calculate the percent error in the specific heat value that you determined experimentally. Use the accepted value given by your teacher based on the identity of the metal. 6. Identify other possible sources of error in this experiment. 7. Can specific heat be used to identify substances? page 58 – DC – T2 – BOOK Unit Seven Experiment – 3 Solution Preparation and Rate of Reaction EX – DC – U7 – 3 Purpose: The purpose of this experiment is to learn how to prepare solutions from more concentrated solutions and how to investigate factors that speed up or slow down chemical reactions. Safety: Safety goggles will be worn at all times during the course of the experiment. Also, protect yourself from any other external contact with the acid and other chemicals used in this experiment. Procedure: Part I: Effect of Concentration on Reaction Rate at Constant Temperature 1. Pour 5 mL of the 6M acid provided by your instructor into a test tube marked “6M HCl”. 2. Calculate how much 6M hydrochloric acid must be diluted in order to make 3M hydrochloric acid in a 50-mL beaker using a graduated cylinder to measure the volumes of the water with which you will dilute and the amount of acid you use. Record the amount of the 6M HCl to be used in Data Table 1. After making this amount in the 50-mL beaker, pour 5 mL of the 3M acid into a test tube marked “3M HCl”. Pour the remainder of the 3M hydrochloric acid into a separate 250-mL Erlenmeyer flask. 3. Calculate how much 3M hydrochloric acid must be diluted in order to make 1M hydrochloric acid in a 50-mL beaker using a graduated cylinder to measure the volumes of the water with which you will dilute and the amount of acid you use. Record the amount of the 3M HCl to be used in Data Table 1. After making this amount in the 50-mL beaker, pour 5 mL of the 1M acid into a test tube marked “1M HCl”. Pour the remainder of the 1M hydrochloric acid into the original 250-mL Erlenmeyer flask. 4. Take zinc strips or pieces of equal size and put one into each test tube. Record the start time and end time of each reaction. Record your observations in Data Table 2. Part II: Effect of Temperature on Reaction Rate 1. Pour 5.0 mL of 6M hydrochloric acid into a test tube. Pour another 5.0 mL of 6M HCl into a test tube maintained over a hot water bath in a 250-mL beaker maintained at 50ºC. (This may be done over a gas burner.) 2. Drop equal amounts of zinc into each test tube. Record the start time and end time of each reaction. Record your observations in Data Table 3. page 59 – DC – T2 – BOOK Observation/Data Tables: Data Table 1: Preparation of Solutions Solution to be prepared Solution to be diluted 3M 6M 1M 3M How much of solution to be diluted is needed How much water is needed to dilute solution Data Table 2: Effect of Concentration on Reaction Rate Reaction Condition Time Reaction Started Time Reaction Ended Reaction Duration Observations 6M HCl 3M HCl 1M HCl Data Table 3: Effect of Temperature on Reaction Rate Reaction Condition Time Reaction Started Time Reaction Ended Reaction Duration Observations 6M HCl at room temp. 6M HCl at 50ºC Conclusion/Discussion: 1. What equation did you use in order to know by how much you needed to dilute more concentrated solutions to get the concentration you wanted for a solution? 2. Write a balanced chemical equation for the reaction between hydrochloric acid and zinc metal. 3. What happens to the reaction rate as temperature increases? Why does this happen? Explain this in terms of the collision theory of reactions. 4. Describe in your own words the effect of concentration on the rate of a reaction. Explain this effect in terms of the collision theory of reactions. page 60 – DC – T2 – BOOK Unit Seven Experiment – 4 Le Chatelier’s Principle (Physical and Chemical Changes) EX – DC – U7 – 4 Purpose: The purpose of this experiment is to determine how equilibrium systems respond to stress. Safety: Safety goggles will be worn at all times during the course of the experiment. Protect yourself from any other external contact with all chemicals used in this experiment. Return or dispose of all materials according to the instructions of your teacher. Procedure: Part I: Effect of Temperature on a Physical Equilibrium 1. Add 2 – 3 mL of saturated potassium nitrate solution to a clean test tube. Using a spatula, add one crystal of potassium nitrate to the solution to act as a seed crystal. 2. Cool the test tube in a 250-mL of ice water for 10 minutes. Record the results. 3. Remove the tube from the ice water and place it in the test-tube rack. Record what happens as the solution warms to room temperature. Part II: Common Ion Effect on a Chemical Equilibrium 1. Use a graduated cylinder to add 50 mL of distilled water to a 100-mL beaker. Add 1 mL of 0.1M iron (III) chloride and 1 mL of 0.1M potassium thiocyanate (KSCN) to the water. Stir the mixture. The color that appears is due to the presence of ferrothiocyanate ions, FeSCN2+. Your teacher will write on the overhead projector slide the reaction that is observed. Record your observations. 2. Label four identical, clean, dry test tubes with the numerals 1 – 4. Pour 5 mL of the mixture from Step 1 into each. Hold the tubes over a white background and look down into them. The solutions should appear equally dark. 3. Tube 1 is the control for this experiment. To tube 2, add 20 drops of 0.1M iron (III) chloride. To tube 3, add 20 drops of 0.1M potassium thiocyanate. Flick each tube to mix the solutions. To tube 4, add 1 g of potassium chloride crystals. Flick the tube to dissolve the crystals. Compare the colors of the solution in tubes 2, 3, and 4 with the color of the solution in the control tube (tube 1). Record your observations. 4. Discard the solutions as directed by your teacher. page 61 – DC – T2 – BOOK Observation/Data Tables: Data Table: Observations System Observations KNO3 (saturated) (cooled) KNO3 (saturated) (warmed) Fe3+/SCN- reaction Fe3+/SCN- mixture + additional Fe3+ Fe3+/SCN- mixture + additional SCNFe3+/SCN- mixture + KCl(s) Conclusion/Discussion: 1. Write a balanced equation for the equilibrium that exists before the saturated and unsaturated potassium nitrate was cooled. 2. Did lowering the temperature (Step 2, Part I) affect the equilibrium? Explain your answer. 3. Did increasing the temperature (Step 3, Part I) disturb the equilibrium? What evidence do you have for your answer? 4. Explain what happened in the potassium nitrate in terms of Le Chatelier’s Principle. 5. Write a balanced equation for the equilibrium that existed after the ferric (Fe3+) and thiocyanate (SCN-) ions were combined in the beaker. 6. What evidence was there that the equilibrium shifted when iron (III) chloride was added? In which direction did it shift? 7. What evidence was there that the equilibrium shifted when potassium thiocyanate was added? In which direction did it shift? 8. Explain the effect of adding potassium chloride to the system. 9. Explain the changes observed in the ferrothiocyanate ion system in terms of Le Chatelier’s Principle. page 62 – DC – T2 – BOOK Unit Eight Experiment – 1 Using Indicators EX – DC – U8 – 1 Purpose: The purpose of this experiment is to estimate the pH of solutions by using acid-base indicators. Safety: Safety goggles will be worn at all times during the course of the experiment. Protect yourself from any other external contact with all chemicals used in this experiment. Ammonia is an irritant. Do not inhale ammonia. Return or dispose of all materials according to the instructions of your teacher. Procedure: 1. Add 1 – 2 mL of the following to five, separate, clean test tubes: 0.1M hydrochloric acid, 0.1M ethanoic acid, 0.1M ammonia, 0.1M sodium hydroxide, and distilled water. RETURN PIPETS TO THEIR ORIGINAL CONTAINERS! 2. Add 2 drops of phenol red indicator solution to each tube. Flick to mix the contents. Record the final color in Data Table 2 and estimate the pH of the solution by referring to Data Table 1. RETURN PIPETS TO THEIR ORIGINAL CONTAINERS! 3. Using fresh samples of the solutions, repeat the procedure for each of the other indicators named in Data Table 1. If you are using paper indicator strips, use a glass stirring rod to transfer a drop of the solution to the indicator strip. Record the results of all the tests in Data Table 2. RETURN PIPETS TO THEIR ORIGINAL CONTAINERS! 4. Test the common household chemicals that are available to you. Test liquids directly. Solids should be dissolved or suspended in water before testing. Record your results in Data Table 3. RETURN PIPETS TO THEIR ORIGINAL CONTAINERS! 5. Follow your teacher’s instructions for proper disposal of the materials. 6. Clean all glassware with distilled water and dry accordingly. DO NOT LEAVE WATER BOTTLES IN SINKS OR HIDDEN AREAS AND DO NOT DISPOSE OF USED PAPER TOWELS IN ANY OTHER PLACE THAN A TRASH RECEPTACLE. DOING SO WILL RESULT IN A 0% FOR YOUR LABORATORY GRADE AND DISCUSSION QUESTIONS. page 63 – DC – T2 – BOOK Observation/Data Tables: Data Table 1: Common Acid-Base Indicators Color in Acid (Hin form) Indicator Color in Base (Inform) pH range Phenol Red Red Litmus Paper Bromthymol Blue Phenolphthalein Data Table 2: Indicator Reactions with Standard Solutions Solution Methyl Red Red Litmus Paper Bromthymol Blue Phenolphthalein Estimated pH 0.1M HCl 0.1M CH3COOH (ethanoic acid) Distilled water (from bottles) 0.1M NH3 (ammonia) 0.1M NaOH Data Table 3: Indicator Reactions with Household Chemicals Substance Methyl Red Red Litmus Paper Bromthymol Blue Aspirin Tea Baking soda Cola (diet) Vinegar page 64 – DC – T2 – BOOK Phenolphthalein Estimated pH Conclusion/Discussion: 1. Compare the pH of 0.1M ethanoic acid with that of 0.1M hydrochloric acid. Compare the pH of 0.1M ammonia with that of 0.1M sodium hydroxide. Explain any differences. 2. Which of the following indicators used in this experiment could most accurately identify a neutral solution? Explain. 3. Are the household chemicals you tested acidic, basic, or neutral? Explain. page 65 – DC – T2 – BOOK Unit Eight Experiment – 2 Determining Molarity by Neutralization EX – DC – U8 – 2 Purpose: The purpose of this experiment is to measure the molarity of hydrochloric acid using a standardized solution of 0.20 M sodium hydroxide. Safety: Safety goggles will be worn at all times during the course of the experiment. DO NOT CONTAMINATE SOURCE CONTAINERS OR BURETS WITH OTHER CHEMICALS THAN THOSE DESIGNATED TO BE IN THEM. Protect yourself from any other external contact with all chemicals used in this experiment. Return or dispose of all materials according to the instructions of your teacher. Procedure: 1. Clean and mount two 25-mL burets. Place a white sheet of paper beneath each buret. Label the left buret “acid” and the right buret “base.” 2. Rinse the “acid” with three 5-mL portions of the solution of hydrochloric acid. Let each portion drain out of the buret before adding the next rinse. Discard these rinses. Fill the buret with the hydrochloric acid. Before beginning the titration, remove any bubbles trapped in the buret and the stopcock. Also make sure that the solution is below the “0 mL” mark. 3. Using the 0.20 M sodium hydroxide solution, rinse and fill the “base” buret. Use a wash bottle of distilled water to rinse off the tip of each buret; catch the runoff in a sink. Record the initial volume in each buret to the nearest 0.01 mL. 4. Add 10-12 mL of the acid solution to a clean 250-mL Erlenmeyer flask. Use the wash bottle to rinse the last drop of acid from the tip of the buret into the flask. Add at least 50 mL of distilled water and 5 drops of bromthymol blue to the flask. 5. Slowly add sodium hydroxide solution from the “base” buret to the flask. As you add the base, gently swirl the solution in the flask. A blue color will appear and quickly disappear as the solutions are mixed. As more and more base is added, the blue color will persist for a longer time before disappearing. This is a sign that you are nearing the equivalence point, also called the end point. Wash down the sides of the flask and the tip of the buret with distilled water from the wash bottle. Continue to add sodium hydroxide more slowly, until a single drop of base turns the solution a pale green color that persists for 15 – 30 seconds. 6. If you overshoot the end point – that is, if you add too much base so the solution turns bright blue – simply add a few drops of acid from the acid buret to turn the solution yellow again. Approach the end point again, adding base drop by drop, until one drop causes the color chance to pale green. 7. When you are sure that you have achieved the end point, record the final volume reading of each buret. Note: Do not allow the level of the solution in either buret to go below the 25-mL mark. If you do, you will have to discard your sample and begin again. page 66 – DC – T2 – BOOK 8. Discard the solution in the Erlenmeyer flask as directed by your teacher, and rinse the flask well with distilled water. Refill both burets, if necessary. Read the initial volume in each buret and do another titration, as described in steps 4 – 7. Observation/Data Tables: Data Table 1: Molarity of Hydrochloric Acid Trial 1 Acid Trial 2 Base Acid Trial 3 Base Acid Base Final volume (mL) Initial volume (mL) Volume used (mL) Molarity of HCl (M) Average molarity of HCl (M) Conclusion/Discussion: 1. Determine how many moles of sodium hydroxide were used in each trial. (HINT: Remember that you used 0.20 M NaOH, and that molarity = moles solute liters solution, meaning you need to convert the milliliters you used into liters.) 2. Based on how many moles of sodium hydroxide were used in each trial, calculate how many moles of hydrochloric acid were used in each trial. (HINT: Use the balanced chemical reaction HCl + NaOH H2O + NaCl to determine the molar ratio you will need to use.) page 67 – DC – T2 – BOOK 3. Calculate the molarity of hydrochloric acid for each trial. 4. Calculate the average molarity of the trials for hydrochloric acid. 5. Why are the burets rinsed with the acid and base solutions before filling? page 68 – DC – T2 – BOOK Appendix A Laboratory Equipment and LPS Safety Contract Triple beam balance Buret Graduated cylinder Test tube rack Erlenmeyer flask Beaker Crucible and lid Bunsen burner Ring stand Double buret clamp Funnel Wire gauze page A-1 – DC – T2 – BOOK Test tube tongs Distilled water wash bottle Clay triangle Test tube Safety goggles Ring clamp Scoopula Test tube brush page A-2 – DC – T2 – BOOK page A-3 – DC – T2 – BOOK page A-4 – DC – T2 – BOOK Appendix B SI Units and Conversions Density d m v Example: If a substance has a mass of 0.75 g and a volume of 3.0 mL, what is the substance’s density? d m 0.75 g 0.25 g mL v 3.0 mL Example: Gold has a density of 19.3 g/cm3. If one has 10.0 cm3 of gold, what mass of gold is present? m 19.3 g m dv d 10.0 cm3 193 g 3 v cm Example: Mercury has a density of 13.6 g/mL. If there are 7.48 g of mercury present, how many milliliters of mercury are there? m m 7.48 g d v 0.55 mL v d 13.6 g mL Specific Gravity Comparison of densities Formula: Specific gravity density of substance density of water Same units must be used in numerator and denominator Used to diagnoses certain illnesses, such as diabetes; used to check the condition of the antifreeze in a vehicle; used for car batteries Temperature Ways to convert: K = C + 273 C = K – 273 Example: If the temperature is 50C, what is the temperature in Kelvins? K = 50 + 273 = 323 K page A-5 – DC – T2 – BOOK Example: If the temperature is 50K, what is the temperature in degrees Celsius? C = 50 – 273 = – 223 C Units of Measurement SI base unit or SI derived unit Quantity * Length Volume Mass Density Temperature Time Pressure Energy Amount of substance Luminous intensity Electric current Symbol meter cubic meter kilogram* grams per cubic centimeter or grams per milliliter kelvin* second* Pascal m m3 kg g/cm3 Joule mole* J mol candela* cd ampere* A g/mL K s Pa Non-SI unit Symbol liter L degree Celsius C atmosphere millimeter of mercury calorie Atm mm Hg cal * : denotes an SI base unit Commonly Used Prefixes in the Metric System Prefix Symbol Meaning mega M kilo k deci d centi c milli m micro nano n pico p 1 million times larger than the unit it precedes 1000 times larger than the unit it precedes 10 times smaller than the unit it precedes 100 times smaller than the unit it precedes 1000 times smaller than the unit it precedes 1 million times smaller than the unit it precedes 1000 million times smaller than the unit it precedes 1 trillion times smaller than the unit it precedes Important conversions: 1 cm3 = 1 mL 103 mL = 1000 cm3 = 1 L page A-6 – DC – T2 – BOOK Scientific notation Factor 1 000 000 106 1000 103 1/10 10-1 1/100 10-2 1/1 000 10-3 1/1 000 000 10-6 1/1 000 000 000 10-9 1/1 000 000 000 000 10-12 Weight and Mass Mass: amount of matter an object has Weight: force that measures the pull on a given mass by gravity Mass does not change based on location; weight does. Conversions (prelude to Chapter Four) Example: How many centimeters are in a kilometer? Solution: Since there are 100 centimeters in a meter and 1000 meters in a kilometer, find a way that will cancel out units. 1 km 1000 m 100 cm • • 1 1 km 1m 1 kilometer 1000 m 100 cm 100000 cm • • 1 1 km 1m page A-7 – DC – T2 – BOOK Appendix C Compound Name and Formula Writing Metals/Nonmetals: The charge of the metal ions in Group 1A is 1+. The charge of the metal ions in Group 2A is 2+. The charge of the metal ions in Group 3A is 3+. The charge of the transition metals and such elements as Sn, Pb, Hg, and Sb may have more than one charge. The charge of the nonmetal ions in Group 5A is 3-. The charge of the nonmetal ions in Group 6A is 2-. The charge of the nonmetal ions in Group 7A is 1-. Group 8A has no ions. Polyatomic ions: Their charge is always negative except for NH4+. Forming ionic compounds: Compounds have electrical neutrality. Na+ and S2- must be written as Na2S since you need two positive charges to balance the 2- charge on the S. Fe3+ and O2- must be written Fe2O3 since you need two 3+ charges to balance three 2- charges (6 + -6 = 0). The positive ion is always written before the negative ion. If two or more polyatomic ions are used in the formula, enclose the polyatomic ion in parentheses and put the number of polyatomic ions you need on the outside of the parentheses as a subscript. For example, Mg2+ and OH- must be written Mg(OH)2 since you need two negative charges of the OH- ion to balance the 2+ charge on the Mg. Do not write the charge of the ion in the formula. For example, sodium sulfide is Na2S, not Na2+S2-, 2Na+S2-, or Na2+S2-. Naming ionic compounds: When a metal is involved, the name of the metal is used. For example, magnesium becomes “magnesium ion” when it becomes a cation. When the metal ion can have two different charges, the charge of the ion is indicated by writing it in Roman numerals in parentheses after the name of the metal. For example, Cu+ is written as the Copper (I) ion. Cu2+ is written as the Copper (II) ion. When a nonmetal is involved, ide is added as a suffix to the root word of the nonmetal (usually the first syllable). For example, phosphorus become the “phosphide ion” as oxygen becomes the “oxide ion.” Polyatomic ions retain their names. To name a metal and a nonmetal together, combine the ion names. For example, when Copper (II) ion is together with the nitride ion, the compound is Copper (II) nitride. page A-8 – DC – T2 – BOOK Naming binary molecular compounds: The first nonmetal gets its full name. The second nonmetal gets its root word + ide. Both nonmetals get a prefix denoting how many atoms are used to make the compound. However, when only one atom is used in the first nonmetal, the prefix mono is not attached. Examples: o CO is carbon monoxide, not monocarbon monoxide. o N2O5 is dinitrogen pentaoxide. Prefixes: o 1 atom – mono (or mon if it begins with an “o”) o 2 atoms – di o 3 atoms – tri o 4 atoms – tetra o 5 atoms – penta o 6 atoms – hexa o 7 atoms – hepta o 8 atoms – octa o 9 atoms – nona o 10 atoms – deca Naming acids: Use the list of acids to name them. Examples: o HC2H3O2: acetic acid o H2CO3: carbonic acid o HNO3: nitric acid o H2SO4: sulfuric acid o H3PO4: phosphoric acid o HCl: hydrochloric acid o HBr: hydrobromic acid o HI: hydroiodic acid o HF: hydrofluoric acid page A-9 – DC – T2 – BOOK Appendix D Chemical Reactions and Quantities Chemical Reaction Classifications: Synthesis/Combination (Oxidation-Reduction): A + B AB 2Na(s) + Cl 2 2NaCl (s) Decomposition (Oxidation-Reduction): AB A + B 2Hg(l) + O2 2HgO(s) Single-Replacement (Oxidation-Reduction): A + BC AC + B 2K(s) + 2H2O(l) 2KOH(aq) + H2 (g) Double-Replacement (Precipitation): A+B- + C+D- A+D- + C+BK2CO3(aq) + BaCl2(aq) 2KCl(aq) + BaCO3(s) Combustion (Oxidation-Reduction): CxHy + x + y O 2 xCO 2 + y H2O 4 2 CH 4 (g) + 2O2 (g) CO2 (g) + 2H2O(g) Redox reactions: I. The Meaning of Oxidation and Reduction A. Oxidation 1. Classical definition: combination of an element with oxygen to produce oxides 2. Modern definition: complete or partial loss of electrons or gain of oxygen 3. Examples a. Rusting (2Fe + 3O2 2Fe2O3) b. Methane oxidation (CH4 + 2O2 CO2 + 2H2O) c. B. Reduction 1. Classical definition: loss of oxygen from a compound 2. Modern definition: complete or partial gain of electrons or loss of oxygen 3. Examples page A-10 – DC – T2 – BOOK a. b. c. C. D. Reduction of iron ore (2Fe2O3 + 3C 4Fe + 3CO2) 2AgNO3 + Cu 2Ag + Cu(NO3)2 Oxidation and reduction always occur simultaneously. Oxidation-reduction reactions 1. Reactions that involve oxidation and reduction occurring 2. Often called “redox reactions” 3. Electrons of one side must equal electrons of other side a. Example 1 Mg(s) S(s) MgS(s) i. b. II. Oxidizing agent: sulfur (gains electrons) ii. Reducing agent: magnesium (loses electrons) Example 2 i. Oxidizing agent: copper (II) nitrate (gains electrons) ii. Reducing agent: magnesium (loses electrons) Oxidation Numbers A. A positive or negative number assigned to a combined atom according to a set of arbitrary rules B. Generally the charge an atom would have if the electrons in each bond were assigned to the atoms of the more electronegative element C. Rules for assigning oxidation numbers 1. 2. The oxidation number of an element in an elementary substance is 0. a. The oxidation number of chlorine in Cl2 or of phosphorus in P4 is 0. b. The oxidation number of Fe by itself is 0. The oxidation number of an element in a monatomic ion is equal to the charge of that ion. a. In the ionic compound NaCl, sodium has an oxidation number of +1 and chlorine has an oxidation number of –1. page A-11 – DC – T2 – BOOK The oxidation number of the bromide ion (Br-) is –1 while the oxidation number of the iron (III) ion (Fe3+) is +3. The oxidation number of hydrogen in a compound is +1, except in metal hydrides (i.e., NaH) where it is –1. The oxidation number of oxygen in a compound is –2. except in peroxides (i.e., H2O2) where it is –1. For any neutral compound, the sum of the oxidation numbers of the atoms in the compound must equal 0. For a polyatomic ion, the sum of the oxidation numbers must equal the ionic charge of the ion. b. 3. 4. 5. 6. Solubility Rules If a salt is said to be soluble, then it will not be a precipitate of the solution. Salts that are said to be insoluble will precipitate out of the solution. Negative ion NO3– I–, Br–, Cl– SO42– CO32–, PO43–, SO32– OH– S2– Rule All compounds formed with the negative ion are soluble. All compounds formed with the negative ion are soluble except Ag+, Pb2+, Hg22+, and Cu+. Most compounds formed with the negative ion are soluble; exceptions include SrSO4, BaSO4, CaSO4, RaSO4, Ag2SO4, and PbSO4. All compounds formed with the negative ion are insoluble except those of the alkali metals and NH4+. All compounds formed with the negative ion are insoluble except those of the alkali metals, NH4+, Sr2+, and Ba2+. (Ca(OH)2 is slightly soluble.) All compounds formed with the negative ion are insoluble except those of the alkali metals, alkaline earth metals, and NH4+. Rules for Balancing Equations: 1. Be sure to write all the correct formulas for all the reactants and products in the reaction. In some cases, you may also need to write in parentheses the state of matter they are in. (i.e., Fe(s), Br2(l), etc.) 2. Write the formulas for the reactants on the left and the formulas for the products on the right with a yield sign () in between. If two of more reactants are involved, separate their formulas with a plus sign (+). When finished, you will have a skeleton equation. 3. Count the number of atoms of each element in the reactants and products. To be as easy as possible, a polyatomic ion appearing the exact same on both sides of the equation can be counted as a single unit. page A-12 – DC – T2 – BOOK 4. Balance the elements one at a time by using coefficients (the numbers out in front of the formulas). When no coefficient is written, it is assumed to be 1. It is best to begin the balancing operation with elements that appear only once of each side of the equation. You must not attempt to balance an equation by changing the subscripts in the chemical formula of a substance. 5. Check each atom or polyatomic ion to be sure that the equation is balanced. 6. Make sure all the coefficients are in the lowest possible ratio that balances. Stoichiometric/Molar Conversions and Calculations: To go from atoms to moles: # of atoms 1 mol of representative unit 1 6.02 x 1023 atoms 2.3 x 1026 atoms O 1 mol O 380 mol O 1 6.02 x 1023 atoms To go from moles to atoms: # of moles 6.02 x 1023 molecule # of atoms mol molecule 1 23 3.6 mol C6H12O 6 6.02 x 10 molecule 24 atoms 5.2 x 1025 atoms 1 mol molecule What is gram atomic mass (gam)? Gram atomic mass is the average mass of an element per mole. This is shown on the Periodic Table of Elements underneath the symbol of the element. What is the gram molecular mass (gmm) and how is it calculated? The gram molecular mass of any molecular compound is the mass of one mole of that compound. To calculate it, add the gram molecular masses of the atoms that make it up. For example, the mass of water would be calculated by doing the following (since there are two hydrogen atoms and one oxygen atom in each mole of water): 2 mol H 1.0 g H 1 mol O 16.0 g O 18.0 g H2O 1 mol H 1 1 mol O 1 What is the gram formula mass (gfm) and how is it calculated? The gram formula mass of any ionic compound is the mass of one mole of the formula unit of that ionic compound. It is calculated the exact same way as the gram molecular mass of a molecule except that it is done for an ionic compound. To calculate, simply add up the atomic masses of the ions in the formula of the compound. For example, in magnesium hydroxide (Mg(OH2)) where the gmm of Mg is 24.3 g, H is 1.0 g, and O is 16.0 g, the gfm for magnesium hydroxide would be calculated as follows: 1 x 24.3 g Mg + 2 x 1.0 g H + 2 x 16.0 g O = 58.3 g Mg(OH)2 page A-13 – DC – T2 – BOOK To go from moles to grams for a compound: # of moles of substance gam, gfm, or gmm of substance 1 mol 2.85 mol H2O 18.0 g H2O 51.3 g H2O 1 1 mol H2O To go from grams to moles for a compound: mol of substance # of grams of substance gam, gfm, or gmm of substance 1 32.5 g H2O 1 mol H2O 1.81 mol H2O 1 18.0 g H 2O To go from moles to volume of a gas at STP: # of moles of gas 22.4 L of gas 1 1 mol of gas 2.8 moles CO2 22.4 L CO 2 63 L CO2 1 1 mol CO 2 To go from density at STP to molar mass of a gas: density of gas in g 22.4 L of gas L mol of gas 1.43 g O2 22.4 L O2 32.0 L O2 L O2 mol O2 To calculate percent composition of an element in a compound: Experimentally: % mass of Element A = grams of Element A 100% grams of compound For example, if a compound is made up of 7.65 g hydrogen and 5.25 g carbon, the total mass of the compound is 12.90 g. To calculate the percent mass of hydrogen in the compound, you would divide 7.65 g by 12.90 g and multiply by 100% to get a percent composition of 59.3% hydrogen. Theoretically: % mass of Element A = grams of Element A in 1 mol of the compound 100% molar mass of the compound For example, the molar mass of hydrogen peroxide (H2O2) is 2 x 1.01 g + 2 x 16.00 g = 34.02 g. Out of that 34.02 g, the mass of hydrogen that is in that mole of hydrogen peroxide is 2 x 1.01 g = 2.02 g. To calculate the percent composition of hydrogen, you would divide 2.02 g by 34.02 g and multiply by 100% to get a percent composition of 5.94% hydrogen. page A-14 – DC – T2 – BOOK To calculate the mass of an element in a given amount of a compound: mass of compound mass of element in 1 mol of the compound molar mass of the compound 1 For example, if you were asked to calculate the mass of carbon in 48.3 g of methane (CH4), you would know that for every molar mass of methane, which is approximately 16.0 g, 12.0 g of that mole of methane is made up of carbon. Therefore, to calculate the mass present in 48.3 g of methane, g of carbon = 12.0 g C 48.3 g CH4 36.2 g carbon 1 16.0 g CH4 To calculate the empirical formula of a compound: Example: What is the empirical formula of a compound that is 10.0% carbon, 0.80% hydrogen, and 89.1% chlorine. 1. Realize that in a 100.0 g sample of this compound, 10.0 g would be carbon, 0.80 g would be hydrogen, and 89.1 g would be chlorine. 2. Convert the grams of each of the elements to moles. 10.0 g C 1 mol C 0.833 mol C 12.0 g C 1 0.80 g H 1 mol H 0.80 mol H 1 1.0 g H 89.1 g Cl 1 mol Cl 2.51 mol Cl 35.5 g Cl 1 3. The mole ratio is C0.833H0.80Cl2.51. This is not the correct empirical formula though because it is not the lowest whole-number ratio. To do this, we need to divide all the molar quantities by the smallest number of moles. This will give a 1 for the element with the smallest number of moles. 0.833 mol C 1.0 mol C 0.80 0.80 mol H 1.0 mol H 0.80 2.51 mol Cl 3.1 mol Cl 0.80 4. The mole ratio is now CHCl3.1. Given how close the 3.1 is to 3, the empirical formula for this is CHCl3. If the mole ration was something like CHCl0.5, we would need to multiply each molar quantity by a value such as 2 to get all whole numbers, resulting in C2H2Cl. To calculate the molecular formula of a compound given molar mass: Example: What is the molecular formula of the compound whose molar mass is 180.0 g and the empirical formula is CH2O? 1. Calculate the empirical formula mass. In this case, the molar mass of CH2O would be 30.0 g. 2. Divide the compound’s molar mass by the empirical formula mass. In this case, you would divide 180.0 g by 30.0 g to get a value of approximately 6. 3. Multiply the subscripts in the empirical formula by the value you calculated in step 2 to get the molecular formula. Multiplying the example empirical formula subscripts by 6, the answer would be C6H12O6. page A-15 – DC – T2 – BOOK Type of Reaction Synthesis/Combination Hints / What to Look For on the Reactant Side Two elements, element and a diatomic gas/liquid/solid What to Do to Complete the Reaction 1. 2. 1. Decomposition One compound 2. 1. Single Replacement Only one ionic compound; the other reactant is an element or a diatomic gas/liquid/solid 2. 3. 1. Two ionic compounds Double Replacement Product cases: 1. One precipitate formed. 2. One gas formed. 3. One liquid formed. 2. 1. Combustion A hydrocarbon (something with carbon and hydrogen) and oxygen gas (can be complete or incomplete combustion) 2. page A-16 – DC – T2 – BOOK Combine the elements as you would if you were forming any ionic compound. Balance the equation. Break down the compound into its constituent elements and/or compounds. Balance the equation. Switch around the two anions or the two cations that need to be replaced with each other. Remember the Activity Series of Metals and of Halogens when it comes to displacing a metal. Also be sure to balance charges in the new compound formed (i.e. Ca replacing Ag in AgCl has a 2+ charge, resulting in CaCl2 for charges to balance). If a displaced element exists in a diatomic state in nature, be sure to indicate this (i.e. H H2). Balance the equation. Switch around the two cations that need to be replaced with each other. Also be sure to balance charges in the new compound formed (i.e. Ca replacing Ag in AgCl has a 2+ charge, resulting in CaCl2 for charges to balance). Balance the equation. If there is a sufficient amount of oxygen, carbon dioxide and water will be the products (complete combustion). If there is an insufficient amount of oxygen, carbon monoxide and water will be the products (incomplete combustion). Balance the equation. Appendix E Study Skills Reading Skills and Review Skills Contrary to popular belief, studying and reviewing is not simply a matter of reading through something trying to memorize the information and occasionally glancing back at notes and supplemental materials. Reading and reviewing entails a number of different skills, many of which can be close to mastered given continued practice and direction. Reading Skills: Browse through the reading before actually reading and preview each assigned reading assignment. If you know what to look for before actually reading as well as how much time you will have to invest in the reading, you will know what kind of preparation needs to be made for the reading assignment. Reading once is not reading. Reading only once is like riding a boat over an ocean: you know what you have gone over, but you don’t understand what any of it really is. Reread after you have done the initial reading. Write notes while reading. Learning is not simply a matter of osmosis for the brain. It helps to have some extra reference than the book. Writing notes in your own words helps in mastering material. Focus on important aspects while skipping the fluff. Writing down everything the book covers, including the “fluff” (the extra things of no importance) is just like overstudying and can be just as damaging to your learning experience. Pick apart what is important and focus on recording those in your notes. Know what parts of the reading are confusing and record them. Not everything you read will make sense to you. Be sure to record questions you have over the material as soon as you have them so that you can ask your teacher, friends, family, etc. Read to understand, not to memorize. Any field of study requires understanding of the material to perform well, not simply memorizing what to do and when to do it. Read so that you understand the relationships between different aspects of what you read, not simply what they are. Review the material. The saying “If you don’t use it, you lose it” is all too true. Consistently review material so that you don’t forget how to do it, whether it is the first week of the semester or the last week of the school year. Ask questions regarding what you have read. Reading and trying to understand without clarifying the confusing points is as futile as simply listening along and not understanding the relationships between what you have heard. Ask questions so that everything can come together clearly, and, if things don’t make sense after asking those questions, ask more questions, whether of your teacher or of your peers. page A-17 – DC – T2 – BOOK Study Skills: Set priorities. Although work outside of school provides money for a variety of different things, the kind of money you will make during high school will pale in comparison to the money you can make with a quality education, as is shown statistically. Prioritize according to long-term goals, not just for the here-andnow. Also realize that, even though sports and athletics allow for competition and personal development, academics do as well, and you need to prioritize accordingly. Don’t get behind. If you get behind in one area, it will impact everything to be learned after that. Do not allow yourself to slip too far, for catching up may become nearly impossible to achieve later. Don’t get overcommitted. Overstudying one subject (aka cramming) can be just as lethal to achievement as not studying at all. Take adequate breaks and disperse your studying over the range of all subjects being taken. Absence for any reason has an effect. Regardless of whether you are gone for sports, clubs, work, family vacation, or another reason, your absence in class will impact your knowledge and experience with a subject. Make adequate preparations in advance to know what you are missing in the way of knowledge and avail yourself to all options that can help you to recover from your absence. Ask for help. Teachers do not enter the profession of teaching solely to speak in front of a class and do nothing else. If you do not understand material covered in class, ask your teacher as soon as the question arises. If that is not possible, write down the question as soon as you have it and ask your teacher at an appropriate time. Should that still not be possible, talk with other teachers proficient in the field you are studying or talk to classmates who can provide assistance with your questions. Get names and phone numbers of classmates. Sometimes the individuals who can best assist you are those that think like you – your peers, friends, and fellow classmates. Know who to call for help in particular subjects. A study group is not a social group. Although a friend may enjoy talking with you, the friend may not provide any assistance if they are as confused about a subject as you are. Involve yourself in study groups that benefit your learning, not who are simply there for personal discussion. Study in short segments and then take a break or reward yourself. Similar to how overstudying can be as disastrous as not studying at all, take breaks and providing yourself rewards for studying allow your mind to process recently-learned information and help you to retain it. Give yourself breaks that allow you to retain information. Review daily, especially in sequence courses. Continual, methodically study is so much more helpful than trying to cram an entire chapter in your head during an hour before a test. Sometimes only spending a few minutes a day reviewing material discussed previously is all that is required to retain information. Analyze how you best retain information and take appropriate steps to review daily. Also, in sequential courses, material covered at the very beginning of the semester needs to be reviewed occasionally. Take time to review things even if they are from “long ago.” Set definite limits for phone, television, Internet, etc. Distractions can be all too tempting. Limit distractions until you feel you have mastered material, with the exception of distracting yourself for occasional breaks. Select an appropriate place and time for study. Some people need to study in absolute silence; others have learned that music playing in the background allows them to focus better. Know when, where, and how to study for material. Observe proper nutrition. For some people, “junk food,” loaded with carbohydrates and preservatives along with other nutrients, helps during studying. For other people, caffeine is a helpful supplement. However, how page A-18 – DC – T2 – BOOK you eat throughout the day and what you provide your body to remain nourished is extremely important. Eat a balanced diet throughout the day to assure yourself that your brain is in top form. Observe proper sleep habits. How well can your brain perform when it is more focused on its exhaustion or distraction from lack of sleep than it can on material you are learning or reviewing? Provide yourself with enough sleep so that learning can take place more easily for your brain. Keep daily communication with parents. Should you feel too frustrated or need feedback, parents are often the greatest source. Keep your parents apprised of your progress so that they can provide you with feedback, assistance, or whatever they can do to help you. page A-19 – DC – T2 – BOOK Appendix F Electron Configuration Rules Aufbau Principle o “Electrons enter orbitals of lowest energy first.” o Level 1 Level 2 Level 3 … o sp… o Know energy level placements (shown below) Pauli Exclusion Principle o “An atomic orbital may describe at most two electrons.” o Electrons have opposite spin, and can only be up to two electrons per orbital. Hund’s Rule o “When electrons occupy orbitals of equal energy, one electron enters each orbital until all the orbitals contain one electron with parallel spins.” o Even distribution of electrons; have to spread them out in the orbitals Aufbau Diagram page A-20 – DC – T2 – BOOK Electron Configuration “Cheat” Diagram s sublevel: can hold 2 electrons p sublevel: can hold 6 electrons d sublevel: can hold 10 electrons f sublevel: can hold 14 electrons The following chart shows the order in which you should fill the orbitals for an electron configuration. page A-21 – DC – T2 – BOOK Appendix G Electron Dot Structure Instructions: For very simple molecules, Lewis dot structures can often be written by just thinking and realizing what to write. However, many molecules are more complex. The following are the steps used to write Lewis dot structures. 1. Count the number of valence electrons. For a molecule, add up the valence electrons of the atoms that are present. For a polyatomic anion, one electron is added for each unit of negative charge. For a polyatomic cation, a number of electrons equal to the positive charge must be subtracted. 2. Draw a skeleton structure for the molecule or ion, joining atoms by single bonds. In some cases, only one arrangement of atoms is possible. In other cases, there may be two or more ways of drawing the structure. Most of the molecules and polyatomic ions with which you will be concerned consist of a central atom bonded to two or more terminal atoms (atoms located on the outer edges of a molecule or ion). NOTE: The central atoms is usually the atom written first in a formula (i.e., C in CO2; N in NH3); put this in the center of the molecule or ion. Terminal atoms are most often hydrogen, oxygen, or a halogen and should be bonded to the central atom. 3. Determine the number of valence electrons still available to be distributed for bonds. To do this, deduct two valence electrons for each single bond written in Step 2. 4. Determine the number of valence electrons required to fill out an octet for each atom (except H) in the skeleton. NOTE: Remember that shared electrons are counted for both atoms. a. If the number of electrons available in Step 3 is equal to the number required from Step 4, distribute the available electrons as unshared pairs so that every atom (except hydrogen) has an octet. b. If the number of electrons available in Step 3 is less than the number required from Step 4, the skeleton structure must be modified by changing single bonds to double or triple bonds. If you are two electrons short, convert a single bond to a double bond; if there is a deficiency of four electrons, convert a single bond to a triple bond (or converting two single bonds to double bonds). page A-22 – DC – T2 – BOOK Example: Example: Draw the Lewis structure for SO2. Solution: Follow the four steps for creating a Lewis dot structure. 1. The number of valence electrons is 6 (from S) + 12 (6 from each O) = 18 valence electrons. 2. The skeleton structure is 3. The number of electrons available for distribution is 18 (original number) – 2 (from the first S – O bond) – 2 (from the second S – O bond) = 14. 4. Since each bond already provides each oxygen with 2 electrons and sulfur with 4 electrons, each oxygen needs 6 more and sulfur needs 4 more, resulting in a need of 16 electrons. As a result, we are missing two electrons to complete octets for each atom. This means that a single bond in the skeleton must be converted to a double bond. or Therefore, the structure of SO2 is : or page A-23 – DC – T2 – BOOK Appendix H VSEPR Models Geometries of Molecules/Ions With Expanded Octets Species Type AX Molecular Geometry Linear Predicted bond angles 180 AX2 Linear 180 CO2 AX3 Trigonal planar 120 CO32- AX4 Tetrahedral 109.5 CH4 AX5 Trigonal bipyramidal 90 120 180 PCl5 AX6 Octahedral 90 180 SF6 Ball and Stick Representation - page A-24 – DC – T2 – BOOK Example CO Example Representation - Geometries of Molecules/Ions With Expanded Octets Species Type AX2E (E represents a lone pair of electrons on the central atom.) Species Species Specie Type Type Species Type s Type Species Type Bent NO21– 104.5 AX2E2 Bent 104.5 H2O AX2E3 Linear 180 XeF2 AX3E Trigonal pyramidal 109.5 AX3E2 T-shaped 90 180 ClF3 AX4E See-saw 90 120 180 SF4 AX4E2 Square planar 90 180 XeF4 AX5E Square pyramidal 90 180 ClF5 - page A-25 – DC – T2 – BOOK NH3 - Appendix I Calorimetry Calculations Problem: A 15.0-g block of ice is heated from -10C to 115C. In order to accomplish this, how many kilojoules of heat are required to be present? Solution: In order to do this, one must realize that the melting point of ice is 0C and the boiling point of water (melted ice) is 100C. Since the ice is going through two phase changes (ice water and water vapor), one must be able to calculate the heat required as an ice, water, and vapor using q = mCT as well as the heat generated while passing from one phase to another by using heats of fusion and vaporization. This means that one needs to know the following values: Cice = 2.1 J/(gC) for when ice is absorbing heat from -10C to 0C. Hfus = 6.01 kJ/mol for when ice is becoming water at 0C; this is the same as 6.01 x 103 J/mol since we want all heat units to be the same. Cwater = 4.184 J/(gC) for when water is absorbing heat from 0C to 100C. Hvap = 40.7 kJ/mol for when water is becoming steam at 100C; this is the same as 40.7 x 103 J/mol since we want all heat units to be the same. Csteam = 1.7 J/(gC) for when steam is absorbing heat from 100C to 115C. The molar mass of ice and water and steam is 18.015 g/mol; we will use this to convert moles to grams in the heats of fusion and vaporization. We have five different heats to add: q1 mCice T for when ice is absorbing heat from -10C to 0C. 1 q 2 m molar mass H fus for when ice is becoming water at 0C. We multiply the mass by the molar mass to convert moles since heat of fusion is expressed in Joules per mole. q 3 mCwater T for when water is absorbing heat from 0C to 100C. 1 q 4 m molar massHvap for when water is becoming steam at 100C. We multiply the mass by the molar mass to convert moles since heat of vaporization is expressed in Joules per mole. q 5 mCsteamT for when steam is absorbing heat from 100C to 115C. This results in q total q1 q 2 q 3 q4 q 5 . 2 .1 J 0C 10C 315 J q1 mCice T 15.0 g g C mol 6.01 x 10 3 J q 2 15.0 g 5004 J 5000 J mol 18.015 g page A-26 – DC – T2 – BOOK 4.184 J 100C 0C 6276 J 6280 J q 3 mC water T 15.0 g g C mol 40.7 x 10 3 J 33888 J 33900 J q 4 15.0 J 18.015 g 1. 7 J 115C 100C 382.5 J 383 J q5 mC steam T 15.0 g g C q total q1 q 2 q 3 q 4 q 5 315 J + 5000 J + 6280 J + 33900 J + 383 J 45878 J 45.878 kJ Thus, to heat ice from -10C to 115C, ice must absorb 45.878 kJ of heat. HINTS: 1. Use only the “q”s you need. Don’t think you will have to add up all heats if a certain thermochemical process is not occurring. 2. Anytime you are going up to melting point, going between melting and boiling point, or going beyond boiling point – in other words, are dealing with a temperature difference – remember that q = mCT. 3. Anytime you are simply adding heat to melt something or boil something, take your mass (with appropriate units) divided by your molar mass of your substance times the heat of fusion or heat of vaporization, respectively. 4. Draw a picture. If Mr. Geist draws pictures to remember, you might want to as well. page A-27 – DC – T2 – BOOK Appendix J Water and Solutions I. Liquid Water and Its Properties A. The Water Molecule 1. Simple triatomic molecule 2. Polarity a. Each O-H bond is highly polar b. O: 2- H: + (partial charges) c. Region around oxygen has a slight negative charge; region around the hydrogens has a slight positive charge d. Hydrogen bonding occurs because of polarity e. Results of hydrogen bonding i. High surface tension ii. Low vapor pressure iii. High specific heat capacity iv. High heat of vaporization v. High boiling point B. Surface Properties 1. Surface tension a. Explained by water’s ability to hydrogen bond b. Water molecules experience uneven attractions at the surface, pulling inward and minimizing surface area page A-28 – DC – T2 – BOOK c. Holds a drop of liquid in a spherical shape while being flattened by gravity d. Why water forms a meniscus in a tube e. Surfactant i. Surface active agent ii. Decreases the surface tension of water iii. Used to interfere with bonding between hydrogen molecules, causing beads of water to collapse iv. Example: detergent in water 2. Low vapor pressure a. Strong hydrogen bonds prevent water from escaping easily b. Prevents bodies of water from drying up C. Specific Heat Capacity 1. C = 4.184 J/(gC) 2. Caused by hydrogen bonding 3. Heat released by water during winter; heat absorbed by water in summer II. Water, Vapor, and Ice A. Evaporation and Condensation 1. Evaporation a. Vaporization that occurs at the surface of a liquid that is not boiling b. Occurs as the result of absorption of heat of vaporization c. Involves breaking intermolecular hydrogen bonds between water Molecules d. Explains high heat of vaporization and high boiling point 2. Condensation a. Opposite of evaporation b. Occurs as the result of releasing of heat of vaporization c. Involves forming intermolecular hydrogen bonds between water molecules 3. Relates to why temperatures in tropics are cooler than expected B. Ice 1. Contraction of liquid as cooling occurs 2. Density will increase as temperature decreases because of less volume 3. Below 4C, density begins to decrease because water begins to no longer behave as a liquid. 4. Insulates heat under the ice 5. Less dense than water, allowing it to float in water 6. Prevents freezing of bodies of water page A-29 – DC – T2 – BOOK III. Aqueous Systems A. Solvents and Solutes 1. Chemically pure water does not exist naturally because of all the substances it can dissolve. 2. Aqueous solutions a. Solution in which the solvent is water b. Solvent: dissolving medium, of which water is for aqueous solutions c. Solute: particles dissolved in the solvent d. Homogeneous mixture page A-30 – DC – T2 – BOOK B. The Solution Process 1. Solvation: process that occurs when a solute is dissolved 2. Insoluble: attractions between the ions in water in crystals are stronger than the attractions exerted by the water C. Electrolytes and Nonelectrolytes 1. Electrolytes: compounds that conduct an electric current in aqueous solution or the molten state 2. Nonelectrolytes: compounds that do not conduct an electric current in aqueous solution or the molten state 3. Strong vs. Weak Electrolytes (page 485) D. Water of Hydration 1. Water in a crystal 2. Also called water of crystallization 3. Example: CuSO4 5H2O implies 5 moles of water to every copper and sulfate pair IV. Heterogeneous Aqueous Systems A. Suspensions: mixtures from which particles settle out upon standing B. Colloids: mixtures containing particles that are intermediate in size between those of suspensions and true solutions C. Tyndall effect: the scattering of light in all directions D. Table 17.6 (KNOW) page A-31 – DC – T2 – BOOK Appendix K Acid and Base Measurements Part I: The Difference Between Acids and Bases For any aqueous system, the product of the hydrogen-ion concentration and the hydroxide-ion concentration is known as the ion-product constant for water (Kw). This is always going to have a value of 1.0 x 10-14 M2. OH 1.0 10 Kw H + 14 - M 2 This is because as hydrogen-ion concentration increases, there has to be less hydroxide-ion concentration. Since there is only so much space in water for one to occur, as one goes up, the other concentration goes down. Take a look at the following reaction when hydrogen chloride dissolves in water to become hydrochloric acid: H O + - 2 H(aq) Cl(aq) HCl(g) More hydrogen ions are being dissociated in water than are hydroxide ions. Since more hydrogen ions are dissociating in water, the [H+] is greater, meaning that HCl is an acid. This means that this solution is an acidic solution, a solution in which the concentration of hydrogen ions ([H+]) is greater than the concentration of hydroxide ions ([OH-]). This implies that [H+] > 1.0 x 10-7 M. However, take a look at the following reaction when sodium hydroxide dissolves in water: H O + - 2 Na(aq) OH(aq) NaOH(s) More hydroxide ions are being dissociated in water than are hydrogen ions. Since more hydroxide ions are dissociating in water, the [OH-] is greater, meaning that NaOH is a base. This means that this solution is a basic solution (also known as an alkaline solution), a solution in which the concentration of hydroxide ions ([OH-]) is greater than the concentration of hydrogen ions ([H+]). This implies that [OH-] > 1.0 x 10-7 M. Part II: pH and pOH A person can express hydrogen-ion and hydroxide-ion concentrations in terms of moles per liter, but this involves working with a lot of scientific notation, which can get really boring and sometimes really difficult to use. A Danish scientist named Sren Srensen came up with the idea of a pH scale, a means of rating concentrations of bases and acids on a scale from 0 to 14 with neutral solutions having a pH of 7. By taking the negative logarithm of a hydrogen-ion concentration, one can easily calculate the pH of a solution as follows: pH -log H + The pH for acids generally falls between 0 and 7, with 0 being the most acidic and 7 being neutral. Similarly, bases fall between 7 and 14 with 14 being the most basic and 7 being neutral. By the same token, if one knows the pH and wishes to find the molar concentration (molarity) of the solution, one can take 10 to the negative power of the pH as follows: page A-32 – DC – T2 – BOOK H 10 + pH Sometimes it is helpful, especially when one is more concerned about working with bases, to measure pOH, the negative logarithm of the concentration of hydroxide ions as follows: - pOH -log OH pOH works opposite of pH since pH + pOH = 14. pOH for acids generally falls between 7 and 14, with 14 being the most acidic and 7 being neutral. Similarly, bases fall between 0 and 7 with 0 being the most basic and 7 being neutral. The same way you can work back to find the concentration of hydrogen ions in solution by knowing pH, you can find the concentration of hydroxide ions in solution by knowing pOH: OH 10 - pOH Realize, especially when double-checking figures, that the following is always true in an aqueous solution: pH + pOH = 14 Kw H+ OH- 1.0 1014 M 2 Appendix L Acid and Base Notes I. Acid-Base Theories A. ___________________________ Acids and Bases 1. Acids contain hydronium ions (H3O+) commonly referred to as hydrogen ions (H+) that dissociate in water a. Different acids release different numbers of H+, known as protons since the hydrogen loses its electron, resulting in only one proton (positive charge) Acid HNO3 (nitric acid) HC2H3O2 (acetic acid) HCl (hydrochloric acid) HBr (hydrobromic acid) HF (hydrofluoric acid) HI (hydroiodic acid) H2SO4 (sulfuric acid) H2CO3 (carbonic acid) H3PO4 (phosphoric acid) H3PO3 (phosphorous acid) Common Acids and Types Number of H+ ions per mole page A-33 – DC – T2 – BOOK Type of acid b. c. d. 2. B. Not all compounds containing hydrogen are acids Not all hydrogens in an acid will necessarily dissociate in water Dissociation only occurs when very polar bonds are present because the hydrogen ions are stabilized by dissolving in solution (i.e., forming hydronium ions in solution) Bases contain hydroxide ions (OH-) that dissociate in water a. Differences in solubility in water (page 227, “Solubility Rules for Ionic Compounds”) 1. High solubility: KOH, NaOH, hydroxides with Group 1 elements 2. Low solubility: Ca(OH)2, Mg(OH)2, hydroxides with Group 2 elements b. React with acids to produce salt and water via double-replacement reaction NaOH(aq) + HCl(aq) NaCl(aq) + H2O(l) ______________________________________________ Acids and Bases 1. Some bases do not give off hydroxide ions but are still basic (i.e., NH3, Na2CO3) 2. _________: hydrogen-ion (H+) donor; __________: hydrogen-ion (H+) acceptor 3. __________________________________: what makes the solution acidic; 4. 5. __________________________________: what makes the solution basic Conjugate acid-base pair related by the loss or gain of a single hydrogen ion Examples NH3 (aq ) H2O(aq ) NH+4 (aq ) Ammonia hydrogen ion acceptor; BronstedLowry base Water hydrogen ion donor; BronstedLowry acid Ammonium ion Conjugate acid OH- (aq ) Hydroxide ion Conjugate base since this makes the solution basic HCl(aq ) H2O(l ) H3O+ (aq ) Cl- (aq ) Hydrochloric acid Water hydrogen ion hydrogen ion donor; Bronstedacceptor; BronstedLowry acid Lowry base 6. C. Hydronium ion Conjugate acid since this makes the solution acidic Chloride ion Conjugate base __________________________________ substance a. Substance that can act as either an acid or a base b. Example: _______________________________ Lewis Acids and Bases 1. Lewis _________: can accept a pair of electrons to form a covalent bond 2. Lewis _________: can donate a pair of electrons to form a covalent bond page A-34 – DC – T2 – BOOK D. Summary of acids and bases Type Arrhenius Bronsted-Lowry Lewis II. Acid-Base Definitions Acid H+ producer H+ donor Electron-pair acceptor Base OH- producer H+ acceptor Electron-pair donor Strengths of Acids and Bases A. Strong and Weak Acids and Bases 1. Strong acids completely ionize in water 2. Weak acids only slightly ionize in water B. Acid dissociation constant HCl(aq ) H2O(l ) H3O+ (aq ) + Cl- (aq ) [acid] + [H ] [conjugate base] Acid dissociation constant: _____________________________________________________ 1. Calculation done at equilibrium 2. The smaller the constant, the less likely the acid will ionize in water 3. The smaller the constant, the weaker the acid 4. Each hydrogen ionizing in water has a different ionization constant C. Base dissociation constant NH3 (aq ) H2O(aq ) NH+4 (aq ) + OH- (aq ) [base] [conjugate acid] [OH ] Base dissociation constant: ____________________________________________________ 1. Calculation done at equilibrium 2. The smaller the constant, the less likely the base will ionize in water 3. The smaller the constant, the weaker the base page A-35 – DC – T2 – BOOK Relative Strengths of Common Acids and Bases Substance Formula Relative Strength HCl Hydrochloric acid HNO3 Strong acids Nitric acid Sulfuric acid H2SO4 Phosphoric acid H3PO4 Ethanoic acid CH3COOH | Carbonic acid H2CO3 Increasing strength of acid Hydrosulfuric acid H2S Hypochlorous acid HclO Boric acid H3BO3 Neutral solution Sodium cyanide Ammonia Methylamine Sodium silicate Calcium hydroxide Sodium hydroxide Potassium hydroxide NaCN NH3 CH3NH2 Na2SiO3 Ca(OH)2 NaOH KOH Increasing strength of base | Strong bases Appendix M Neutralization Notes I. Neutralization Reactions A. Acid-Base Reactions 1. An acid reacts with a base to produce water and salt (generally) 2. Examples (strong acids reacting with strong bases) a. HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l) b. H2SO4(aq) + 2KOH(aq) K2SO4(aq) + H2O(l) 3. Solubilities of salts dictated by the Solubility Rules for Ionic Compounds (p. 227) 4. Neutralization reaction a. Reaction in which an acid and a base react in an aqueous solution to produce a salt and water b. A double-replacement reaction c. Neutralization will not necessarily occur between weak acids and/or weak bases d. Type of reaction to prepare pure samples of salt (i.e., NaCl from the reaction shown above) B. Titration 1. Process of adding a known amount of solution of known concentration to determine the concentration of another solution 2. Steps a. A measured volume of an acid solution of unknown concentration is added to a flask. b. Several drops of an indicator are added to the solution. c. Measured volumes of a base of known concentration are mixed into the acid until the indicator just barely changes color and maintains that color. This occurs at the “end point”. 3. Examples for strong acids and strong bases (refer to Chapter 20 Notes) page A-36 – DC – T2 – BOOK 4. When titration is complete at the end point, the contents of the flask are only salt and water. Example 1: How many moles of sulfuric acid are required to neutralize 0.75 mol of potassium hydroxide? Solution 1: It helps to first know the equation of neutralization. A reaction between sulfuric acid (a strong acid) and potassium hydroxide (a strong base) is a double-replacement reaction. Example 2: What is the molarity of sodium hydroxide if 20.0 mL of the solution is neutralized by 17.4 mL of 1.00M phosphoric acid? It helps to first know the equation of neutralization. A reaction between phosphoric acid (a strong acid) and sodium hydroxide (a strong base) is a double replacement reaction. C. Titration Curves 1. Strong-Acid; Strong-Base Titration a. The end point (equivalence point) has a pH at or very close to when the pH = 7. b. Both acid and base completely ionize and therefore do not create hydrolyzing salts (salts produced that remove hydrogen ions from water or donate hydrogen ions to water) page A-37 – DC – T2 – BOOK c. 2. c. Titration curve for a strong-acid, strong-base titration Strong-Acid; Weak-Base Titration a. The end point (equivalence point) has a pH < 7. b. Hydrolyzing salts that are produced from the titration will donate hydrogen ions to the water, increasing the concentration of hydrogen ions ([H+]) and therefore decrease the pH of the solution comparatively. Titration curve for a strong-acid, weak-base titration page A-38 – DC – T2 – BOOK 3. Weak-Acid; Strong-Base Titration a. The end point (equivalence point) has a pH > 7. b. Hydrolyzing salts that are produced from the titration will remove hydrogen ions to the water, decreasing the concentration of hydrogen ions ([H+]) and therefore increase the pH of the solution comparatively. c. Titration curve for a weak-acid, strong-base titration 4. Indicator selection for titration a. The indicator’s point of color change must be taken into consideration. i. Bromythol blue works for strong-acid, strong-base titrations because it will change to green right at or close to pH = 7. Phenolphthalein has an end point at pH = 8 but is often used because it works even if one is colorblind. Also, the pH changes so rapidly near the end point (4 to 10 with one drop of base) that it will work well. ii. Methyl red works for strong-acid, weak-base titrations because it will change color at pH = 5. Other indicators would not show a change or would show it too prematurely. iii. Phenolphthalein works well for weak-acid, strong-base titrations because of its end point in the upper pH range. Alizarin yellow R works also as its color will change in the range of 10 < pH < 12. b. pH meters i. Often used in industry for more precise measurement of endpoint ii. Helpful in creating a titration curve Summary of titrations and solutions a. Strong acid + Strong base = Neutral solution b. Strong acid + Weak base = Acidic solution c. Weak acid + Strong base = Basic solution 5. page A-39 – DC – T2 – BOOK Appendix N Practice Tests Unit Five Practice Test PT – DC – U5 Multiple Choice. On the answer sheet for each question, write the upper-case letter of the answer that best completes or answers the statement or question in the adjacent corresponding blank. LPS Standard(s): 12.2.4a, 12.2.5a, 12.2.5b State Standard(s): 12.8.3 1. Which scientist developed a model with a positive region that encapsulated negatively-charged particles he discovered? (A) Dalton (B) Bohr (C) Rutherford (D) Thomson (E) Schrodinger 2. Which scientist was responsible for discovering the nucleus and developed a model of the atom that incorporated this in his model of the atom? (A) Dalton (B) Bohr (C) Rutherford (D) Thomson (E) Schrodinger 3. Which scientist developed an indivisible model of the atom? (A) Dalton (B) Bohr (C) Rutherford (D) Thomson (E) Schrodinger 4. Which scientist developed the current model of the atom? (A) Dalton (B) Bohr (C) Rutherford (D) Thomson (E) Schrodinger 5. Which scientist described the model of the atom based on electrons in orbits that could be quantized? (A) Dalton (B) Bohr (C) Rutherford (D) Thomson (E) Schrodinger 6. Which scientist discovered the electron? (A) Dalton (B) Bohr (C) Rutherford (D) Thomson (E) Schrodinger 7. Which scientist developed the earliest model of the atom based on experimentation? (A) Dalton (B) Bohr (C) Rutherford (D) Thomson (E) Schrodinger 8. In the Bohr model of the atom, an electron in an orbit has a fixed ___. (A) position (B) color (C) energy 9. The quantum mechanical model exactly predicts which characteristic of electrons in an atom? (A) position (B) energy (C) orbit (D) charge (E) None of the choices listed 10. Which of the following is an accurate description of Thomson’s model of the atom? (A) Electrons occupy fixed positions around the protons, which are at the center of the atom. (B) The electrons orbit in specified energy levels around the protons, which are at the center of the atom. (C) The electrons, like “raisins,” are stuck into a lump of protons, like “dough,” in a “plum pudding” atom. (D) The electrons and protons move throughout the atom. page A-40 – DC – T2 – BOOK LPS Standard(s): 12.2.4a, 12.2.5a, 12.2.5c 11. State Standard(s): 12.3.1a In which of the following is the number of neutrons correctly represented? (A) 167 N has 9 neutrons. (D) 239 94 Pu has 333 neutrons. (B) 146 C has 14 neutrons. (E) 42 He has 6 neutrons. (C) 188 O has 8 neutrons. 12. Which of the following sets of symbols represents isotopes of the same element? 90 90 51 52 (C) 50 (A) 90 42 J 43 J 44 J 19 L 19 L 19 L (B) 168 O 179 O 18 10 O (D) None of those listed 13. How many electrons are in a neutral atom of the isotope plutonium-239? (A) 94 (B) 145 (C) 239 (D) 333 (E) None of those listed 14. How many protons are in a neutral atom of the isotope plutonium-239? (A) 94 (B) 145 (C) 239 (D) 333 (E) None of those listed 15. How many neutrons are in a neutral atom of the isotope plutonium-239? (A) 94 (B) 145 (C) 239 (D) 333 (E) None of those listed 16. What is the mass number of the isotope plutonium-239? (A) 94 (B) 145 (C) 239 (D) 333 (E) None of those listed 17. What is the atomic number of the isotope plutonium-239? (A) 94 (B) 145 (C) 239 (D) 333 (E) None of those listed Refer to the following isotope for questions 18 – 20. 36 17 Cl 18. How many neutrons does the ion of this isotope contain? (A) 17 (B) 18 (C) 19 (D) 36 (E) None of those listed 19. How many electrons does the ion of this isotope contain? (A) 17 (B) 18 (C) 19 (D) 36 (E) None of those listed 20. How many protons does the ion of this isotope contain? (A) 17 (B) 18 (C) 19 (D) 36 (E) None of those listed LPS Standard(s): --- State Standard(s): 12.3.1, 12.3.6 21. In the first principal energy level (n = 1), what orbitals can exist in that energy level? (A) s and p (B) s, p, and d (C) Only s (D) Only p (E) Only d 22. If only two electrons occupy two p orbitals, what is the direction of the spins of these two electrons? (A) Both are always clockwise. (B) Both are always counterclockwise. (C) They are either both clockwise or both counterclockwise. (D) One is clockwise and the other is counterclockwise. page A-41 – DC – T2 – BOOK Refer to the following diagram for questions 23 – 26. 23. What type of orbital does the picture represent? (B) p (C) d (D) f (E) I (A) s 24. What is the maximum number of this kind of orbital in any principal energy level? (A) 1 (B) 2 (C) 3 (D) 5 (E) 7 25. How many electrons can all the orbitals of this sublevel hold collectively in any energy level? (A) 2 (B) 6 (C) 10 (D) 14 (E) infinitely many 26. What does the black spaces near the center in the picture represent? (A) areas of high electron concentration (C) modes of movement for electrons (B) a node (D) areas where protons will move 27. Which electron configuration of the 4d energy sublevel is the most stable? (A) 4d1 (B) 4d2 (C) 4d3 (D) 4d4 (E) 4d5 28. Which electron configuration of the 4d energy sublevel is the most stable? (B) 4d7 (C) 4d8 (D) 4d9 (E) 4d10 (A) 4d6 29. In the second principal energy level (n = 2), what orbitals can exist in that energy level? (A) s and p (B) s, p, and d (C) Only s (D) Only p (E) Only d 30. If only two electrons occupy one p orbital, what is the direction of the spins of these two electrons? (A) Both are always clockwise. (B) Both are always counterclockwise. (C) They are either both clockwise or both counterclockwise. (D) One is clockwise and the other is counterclockwise. 31. What type of orbital has the greatest number of orientations? (B) p (C) d (D) f (E) I (A) s 32. What type of orbital is the only one contained in a hydrogen atom? (B) p (C) d (D) f (E) I (A) s 33. Which of the following sublevels exists in boron but not beryllium? (B) p (C) d (D) f (E) I (A) s 34. How many p orbitals can exist in any energy level? (A) 1 (B) 2 (C) 3 (D) 5 (E) 7 35. What is the maximum number of electrons that all the p orbitals of an energy level can hold? (A) 2 (B) 4 (C) 6 (D) 8 (E) 10 page A-42 – DC – T2 – BOOK LPS Standard(s): --- State Standard(s): 12.3.1 36. The electron configuration for rhodium (Rh) is ___. (C) 1s22s22p63s23p64s24d104p65s25d7 (A) 1s22s22p63s23p64s23d104p65s24d7 2 2 6 2 6 2 10 6 1 8 (B) 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d (D) 1s22s22p63s23p64s24d104p65s15d8 37. The electron configuration for copper (Cu) is ___. (C) 1s22s22p63s23p64s24d9 (A) 1s22s22p63s23p64s23d9 (B) 1s22s22p63s23p64s13d10 (D) 1s22s22p63s23p64s14d10 38. The electron configuration for the rubidium ion (Rb+) is ___. (C) 1s22s22p63s23p64s23d104p65s1 (A) 1s22s22p63s23p64s23d104p6 2 2 6 2 6 2 10 6 2 10 6 (B) 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p (D) None of those listed 39. The electron configuration for calcium (Ca) is ___. (C) 1s22s22p103s23p4 (A) 1s22s22p23s23p24s2 2 2 2 6 2 (B) 1s 2s 3s 3p 3d (D) 1s22s22p63s23p64s2 40. The electron configuration for vanadium (V) is ___. (C) 1s22s22p63s23p64s24d3 (A) 1s22s22p63s23p64s23d3 (B) 1s22s22p63s23p64s13d4 (D) 1s22s22p63s23p64s14d4 41. The electron configuration for chromium (Cr) is ___. (C) 1s22s22p63s23p64s24d4 (A) 1s22s22p63s23p64s23d4 2 2 6 2 6 1 5 (B) 1s 2s 2p 3s 3p 4s 3d (D) 1s22s22p63s23p64s14d5 42. The electron configuration for the selenium ion (Se2–) is ___. (C) 1s22s22p63s23p64s23d104p4 (A) 1s22s22p63s23p64s23d104p6 2 2 6 2 6 2 10 5 (B) 1s 2s 2p 3s 3p 4s 3d 4p (D) None of those listed 43. The electron configuration for the sodium ion (Na+) is ___. (B) 1s22s22p6 (C) 1s22s22p63s2 (D) None of those listed (A) 1s2 44. The electron configuration for the fluoride ion (F–) is ___. (B) 1s22s22p4 (C) 1s22s22p6 (D) None of those listed (A) 1s22s22p2 45. The oxide ion (O2–) has the same electron configuration as ___. (A) F (B) C (C) N (D) Ne 46. The calcium ion (Ca2+) is isoelectronic to ___. (A) K (B) Ar (C) Br (D) Kr 47. How many half-filled orbitals are there in a nitrogen atom? (A) 0 (B) 1 (C) 2 (D) 3 (E) 4 48. How many unpaired electrons are there in a sulfur ion? (A) 0 (B) 1 (C) 2 (D) 3 (E) 4 49. How many unpaired electrons are there in a chloride ion? (A) 0 (B) 1 (C) 2 (D) 3 (E) 4 50. What is the number of electrons in the outermost energy level of a boron atom? (A) 1 (B) 2 (C) 3 (D) 4 (E) 5 page A-43 – DC – T2 – BOOK LPS Standard(s): 12.2.5d State Standard(s): 12.3.2b 51. What is a general classification for the elements of neon, calcium, and oxygen? (A) Group A elements (D) noble gases (B) Group B elements (E) halogens (C) Group C elements 52. What is a general classification for the gold, silver, and zinc? (A) Group A elements (D) noble gases (B) Group B elements (E) halogens (C) Group C elements 53. Which of the following elements is a representative element? (A) cesium (B) chromium (C) californium (D) cerium 54. Which of the following groupings contains only transition metals? (A) Magnesium, chromium, silver (C) Copper, cobalt, cadmium (B) Nickel, iron, polonium (D) Aluminum, magnesium, lithium 55. Of the elements Pt, As, V, Li, and Kr, which is a metalloid? (A) Pt (B) As (C) V (D) Li (E) Kr 56. What is another name for a family of elements in the periodic table called? (A) period (B) transition (C) list (D) group 57. Who first arranged the elements according to atomic number and is responsible for our current periodic table of elements? (A) John Dalton (D) Antoine Lavoisier (B) Dmitri Mendeleev (E) Henry Moseley (C) Louis Pasteur 58. What is a horizontal row of elements in the periodic table called? (A) period (B) transition (C) list (D) group 59. The periodic law states that there is a periodic repetition of the physical and chemical properties of elements ___. (A) when they are arranged in the order Dmitri Mendeleev ordered them (B) when they are arranged in the order Henry Moseley ordered them (C) when they are arranged in the order Eugene Kirionov ordered them (D) if only metalloids are considered 60. The symbol for the 4th period Group 4A element is ___. (A) As (B) Ge (C) In (D) Sn (E) none of those listed LPS Standard(s): 12.2.5d 61. State Standard(s): 12.3.1, 12.3.2b What is true of the electron configurations of the inner transition metals? (A) The outermost s and f sublevels are very close in energy and have electrons in them. (B) The outermost s and p sublevels are partially filled. (C) The outermost s and d sublevels are very close in energy and have electrons in them. (D) The outermost s and p sublevels are filled. page A-44 – DC – T2 – BOOK 62. What is true of the electron configurations of cobalt, molybdenum, and titanium? (A) The outermost s and f sublevels are very close in energy and have electrons in them. (B) The outermost s and p sublevels are partially filled. (C) The outermost s and d sublevels are very close in energy and have electrons in them. (D) The outermost s and p sublevels are filled. 63. What is true of the electron configurations of calcium, phosphorus, and fluorine? (A) The outermost s and f sublevels are very close in energy and have electrons in them. (B) The outermost s and p sublevels are partially filled. (C) The outermost s and d sublevels are very close in energy and have electrons in them. (D) The outermost s and p sublevels are filled. 64. The symbol of the first element that fills electrons in the s sublevel is ___. (A) B (B) H (C) He (D) Sc (E) none of those listed 65. The symbol of the second element that fills electrons in the d sublevel is ___. (A) Ca (B) Sc (C) Ti (D) Zn (E) none of those listed 66. The symbol for the 4th period element containing only 6 3d electrons is ___. (A) Fe (B) Mn (C) Ru (D) Tc (E) none of those listed 67. The category of elements that is characterized by the filling of d orbitals is the ___. (A) alkali metals (C) inner transition metals (B) alkaline earth metals (D) transition metals 68. On the periodic table, every period correlates to ___. (A) a principal energy level (C) an energy sublevel (B) an atomic number (D) an atomic mass 69. The category of elements that would end with an s1 electron configuration would be the ___. (A) alkali metals (C) halogens (B) alkaline earth metals (D) noble gases 70. The category of elements that would end with an p6 electron configuration would be the ___. (A) alkali metals (C) halogens (B) alkaline earth metals (D) noble gases LPS Standard(s): 12.2.5d State Standard(s): 12.3.2b 71. Which of the following factors contributes to the greater ionization energy of the lower-atomicnumber elements in a family in the periodic table? (A) Smaller distance from the nucleus (C) Smaller number of protons in nuclei (B) Smaller nuclei (D) Greater number of valence electrons 72. What term is used to describe the energy required to remove an electron from a gaseous atom? (A) excitation energy (D) heat of vaporization (B) ionization energy (E) electrolytic energy (C) polarization energy 73. Which of the following factors contributes to the lower ionization energy of the elements on the left side of a period in the periodic table? (A) Less shielding by inner electrons (C) Smaller number of protons in nuclei (B) Smaller nuclei (D) Greater number of valence electrons page A-45 – DC – T2 – BOOK 74. Which group of the periodic table has the lowest electronegativity? (A) 1A (B) 6A (C) 3A (D) 7A (E) 2A 75. Of the following, which element has the greatest first ionization energy? (A) strontium (B) phosphorus (C) fluorine (D) carbon 76. Of the following, which element’s atoms have the largest ionic radius? (A) lithium (B) potassium (C) rubidium (D) sodium 77. Of the following, which element’s atoms have the largest atomic radius? (A) iodine (B) rubidium (C) strontium (D) tellurium 78. Of the following, what is the most electronegative element? (A) iodine (B) rubidium (C) strontium (D) tellurium 79. As you move from top to bottom down the first group of the periodic table, ___. (A) the ionization energy decreases (C) the electronegativity increases (B) the atomic radii decrease (D) the atomic mass decreases 80. As you move from left to right across the third period of the periodic table, ___. (A) the ionization energy increases (C) the electronegativity decreases (B) the atomic radii increase (D) the atomic mass decreases LPS Standard(s): --- State Standard(s): 12.3.1a 81. What does an unstable nucleus NOT do to become more stable? (A) lose electrons (B) lose neutrons (C) lose protons (D) gain protons 82. Why do radioactive isotopes emit radiation? (A) To achieve a proper protons to neutron ratio (B) To ionize and feel like a noble gas (C) To have enough protons to bond to other atoms (D) To gain energy for more radiation LPS Standard(s): --- State Standard(s): 12.3.1b Choices for this section are as follows: (A) alpha emission (B) beta emission (C) gamma emission (D) None of those listed 83. Emission definitely involved when an element loses 4 in its mass number and 2 in its atomic number 84. Has no mass 85. Least penetrating emission 86. Consists of the same particles J.J. Thomson discovered 87. Emission definitely involved when a mass number does not change but the atomic number does page A-46 – DC – T2 – BOOK LPS Standard(s): --- State Standard(s): 12.3.5a 88. E = mc2 was an equation developed by Einstein to show the relationship between mass and energy with ___ as a constant of proportionality between the two. (A) specific heat (B) Planck’s constant (C) the speed of light (D) gamma radiation 89. In the equation E = mc2, m represents ___. (A) meters (B) mass (C) momentum LPS Standard(s): --- (D) Mr. Geist State Standard(s): 12.1.2a; 12.3.1 90. Which of the following half-lives indicates the most stable element? (A) 2 seconds (B) 2 hours (C) 2 days (D) 2 weeks (E) 2 years 91. If the half-life of sodium-24 is 15 hours, how much remains from a 10.0 g sample after 30 hours? (A) 0.63 g (B) 1.3 g (C) 2.5 g (D) 5.0 g (E) None of the answers listed 92. What is the half-life of iodine-131 if, after 24 days, 0.125 g remains from one 1.00 g starting sample? (A) 3 days (B) 6 days (C) 8 days (D) 12 days (E) None of the answers listed LPS Standard(s): --- Choices for this section are as follows: (A) nuclear fission State Standard(s): 12.3.1 (B) nuclear fusion (C) None of those listed 93. Occurs in most, if not all, stars of the universe 94. Produces radioactive waste that must be contained in a secure facility 95. The combination of atoms LPS Standard(s): --- State Standard(s): 12.3.6d 96. What is the frequency of 6.502 x 10–7 m wavelength light? (A) 4.614 x 105 Hz (B) 4.614 x 1014 Hz (C) 1.951 x 102 Hz 97. What is the approximate energy of a photon having a frequency of 4 x 106 Hz? (A) 3 x 10–27 J (B) 3 x 10–28 J (C) 2 x 10–42 J (D) 3 x 1041 J (E) 1 x 10–19 J (D) 1.951 x 1011 Hz For questions 98 – 100, refer to the following choices: (A) Wave A ( = 5.1 x 105 Hz) (B) Wave B ( = 5.1 x 1010 Hz) (C) Waves A and B are equal for this. (D) Not enough information provided or not a valid question 98. Which wave would possess the greater energy? 99. Which wave would possess the longer wavelength? 100. Which wave would possess the greater speed? page A-47 – DC – T2 – BOOK Unit Six Practice Test (Multiple Choice) PT – DC – U6 Multiple Choice. On the scantron for each question, fill in the rectangle of the corresponding letter of the answer that best completes or answers the statement or question in the adjacent corresponding blank. LPS Standard(s): 12.2.5f State Standard(s): 12.3.3c 1. To form aluminum chloride, ___. (A) one aluminum atom gains five electrons from five chlorine atoms (B) one chlorine atom gains one electron from one aluminum atom (C) three aluminum atoms gain three electrons from one chlorine atom (D) three chlorine atoms gain three electrons from one aluminum atom 2. To form potassium iodide, ___. (A) one potassium atom gains seven electrons from an iodine atom (B) one iodine atom gains seven electrons from a potassium atom (C) one potassium atom gains one electron from an iodine atom (D) one iodine atom gains one electron from a potassium atom 3. To form strontium chloride, ___. (A) one strontium atom gains two electrons from two chlorine atoms (B) one chlorine atom gains two electrons from two strontium atoms (C) two strontium atoms gain two electrons from one chlorine atom (D) two chlorine atoms gain two electrons from one strontium atom 4. To form sodium oxide, ___. (A) one sodium atom gains six electrons from two oxygen atoms (B) one oxygen atom gains six electrons from two sodium atoms (C) one sodium atom gains two electrons from two oxygen atoms (D) one oxygen atom gains two electrons from two sodium atoms 5. How does aluminum obey the octet rule when reacting to form ionic compounds? (A) It gains electrons. (B) It loses electrons. (C) It neither gains nor loses electrons. 6. How does sulfur obey the octet rule when reacting to form ionic compounds? (A) It gains electrons. (B) It loses electrons. (C) It neither gains nor loses electrons. 7. What is the formula of the ion formed when sulfur achieves a noble-gas electron configuration? (B) S+ (C) S– (D) S2– (E) S3– (A) S6+ 8. What is the formula of the ion formed when lithium achieves a noble-gas electron configuration? (B) Li+ (C) Li– (D) Li2– (E) Li3– (A) Li2+ 9. How many electrons does barium have to lose/gain up to achieve a noble-gas electron configuration? (A) lose 1 (B) lose 2 (C) gain 1 (D) gain 2 10. How many electrons does bromine have to lose/gain up to achieve a noble-gas electron configuration? (A) lose 1 (B) lose 2 (C) gain 1 (D) gain 2 page A-48 – DC – T2 – BOOK LPS Standard(s): 12.2.5f 11. State Standard(s): 12.1.2a; 12.3.3c Which of the following covalent bonds is NOT polar? (A) C – C (B) H – Br (C) C – Cl (D) C – Br (E) C – S 12. Which of the following pairs of elements can be joined by an ionic bond if the atoms ionize? (A) sodium and carbon (C) carbon and carbon (B) nitrogen and carbon (D) lithium and oxygen 13. The polarity of the bond between a carbon atom and chlorine atom would best be identified as a(n) ___ bond. (A) ionic (B) nonpolar covalent (C) polar covalent 14. In the bond between sodium and fluorine to make sodium fluoride, the sodium atom would have ___. (D) a complete charge of –1 (A) a partial charge shown with + (B) a partial charge shown with – (E) no partial charge (C) a complete charge of +1 15. In the bond between hydrogen and chlorine in H – Cl, the hydrogen atom would have ___. (D) no partial charge (A) a partial charge shown with + (E) a partial charge shown with +/– (B) a partial charge shown with – (C) a complete charge of –1 LPS Standard(s): 12.2.5f State Standard(s): 12.1.2a; 12.3.3c 16. Which of the following elements can form diatomic molecules held together by single covalent bonds and adhering to the octet rule? (A) hydrogen (B) nitrogen (C) oxygen (D) sodium 17. ___ covalent bonds possess the least energy. (A) Double (B) Single (C) Triple 18. Which of the following compounds is not covalently bonded? (A) nitrogen (B) oxygen (C) strontium (D) sulfur 19. ___ electrons are shared in a triple covalent bond. (A) 0 (B) 2 (C) 4 (D) 6 20. When reacting with atoms of their own element, nitrogen atoms form ___ covalent bonds to create nitrogen molecules. (A) single (B) double (C) triple (D) no 21. Which of the following elements can form diatomic molecules held together by single covalent bonds, not following the octet rule, and not achieving 8 valence electrons per atom? (A) hydrogen (B) nitrogen (C) oxygen (D) sodium 22. What is the total number of covalent bonds normally associated with a single carbon atom as the central atom in a compound? (A) 1 (B) 2 (C) 3 (D) 4 page A-49 – DC – T2 – BOOK 23. How many covalent bonds are there in a covalently bonded molecule containing one phosphorus atom and five chlorine atoms? (A) 3 (B) 4 (C) 5 (D) 6 24. Which of the following elements do not exist as diatomic molecules? (A) bromine (B) iodine (C) oxygen (D) phosphorus 25. A ___ covalent bond is the only bond contained in carbon monoxide. (A) double (B) single (C) triple 26. A molecule or polyatomic ion with only double covalent bonds is ___. (B) SO3 (C) CO2 (D) SO32– (A) CH4 27. A molecule or polyatomic ion with only single covalent bonds is ___. (C) CO2 (D) CO32– (A) HCCH (B) CH4 28. A molecule or polyatomic ion with a single covalent bond and a triple covalent bond is ___. (B) HCN (C) SO3 (D) N2 (A) H2O2 29. How do atoms achieve noble-gas electron configurations in single covalent bonds? (A) One atom completely loses two electrons to the other atom in the bond. (B) Two atoms share two electrons. (C) Two atoms share four electrons. (D) Two atoms share six electrons. 30. ___ covalent bonds are the shortest in length. (A) Double (B) Single (C) Triple LPS Standard(s): --- State Standard(s): 12.3.3c 31. Which of the following molecules has an electron dot structure that does NOT obey the octet rule? (C) PF3 (D) HCN (E) CCl4 (A) NO (B) CS2 32. Which of the following violates the octet rule? (B) IF3 (C) PF3 (D) SbF3 (A) NF3 (E) AsF3 33. In which of the following compounds/ions is the octet expanded to include 14 electrons? (B) BF3 (C) PCl5 (D) IF7 (E) SBr6 (A) H2O 34. In which of the following compounds/ions is the octet expanded to include 12 electrons? (B) SO42– (C) SBr6 (D) SO32– (E) PCl5 (A) SO3 35. Which of the following compounds/ions violates the octet rule? (B) CH4 (C) SCl6 (D) CO (E) SO2 (A) CO2 page A-50 – DC – T2 – BOOK LPS Standard(s): --- State Standard(s): 12.3.3c For the following questions, identify the molecular geometry of the specified compounds using the following choices: (A) bent (B) tetrahedral (C) linear (D) trigonal pyramidal (E) none of the choices listed 36. H2O (water) 39. SO32– (sulfite ion) 37. SO2 (sulfur dioxide) 40. CH4 (methane) 38. HCl (hydrochloric acid) For the following questions, identify the molecular geometry of the specified compounds using the following choices: (A) trigonal planar (D) linear (B) octahedral (C) trigonal bipyramidal (E) none of the choices listed 41. NH3 (ammonia) 44. SiO2 (silicon dioxide) 42. ClO (hypochlorite ion) 45. PCl5 (phosphorus pentachloride) 43. PO43– (phosphate ion) LPS Standard(s): --- State Standard(s): 12.3.3c 46. Which of the following compounds is the most polar? (B) CH4 (C) CO2 (D) HBr (A) Br2 47. Diatomic molecules are ___ compounds. (A) nonpolar covalent (B) polar covalent (C) ionic 48. What would best describe sodium chloride as a compound? (A) nonpolar covalent (B) polar covalent (C) ionic 49. What would best describe oxygen gas as a compound? (A) nonpolar covalent (B) polar covalent (C) ionic 50. What would best describe sulfur dioxide as a compound? (A) nonpolar covalent (B) polar covalent (C) ionic LPS Standard(s): --- State Standard(s): 12.3.3c 51. Which of the following causes the boiling point of HF to be much higher than that of HCl or HBr? (A) hydrogen bonds (C) covalent bonds (B) van der Waals forces (D) coordinate covalent bonds 52. The strongest forms of intermolecular attractions are ___. (A) dipole interactions (B) dispersion forces (C) hydrogen bonds (D) ionic bonds The weakest forms of intermolecular attractions are ___. (A) dipole interactions (B) dispersion forces (C) hydrogen bonds (D) ionic bonds 53. page A-51 – DC – T2 – BOOK 54. Why is hydrogen-bonding only possible with hydrogen? (A) because hydrogen is the only atoms whose nucleus is not shielded by electrons when it is involved in a covalent bond (B) because hydrogen is the only atom that is the same size as an oxygen atom (C) because hydrogen has the highest electronegativity of any element in the periodic table 55. In an electric field, which region of the water molecule is attracted to the positive pole? (A) the oxygen region of the molecule (B) the hydrogen region of the molecule (C) No part of the water molecule is attracted to the positive pole. LPS Standard(s): --- State Standard(s): 12.3.2d For the following questions, identify the following as ionic or molecular compounds using the following (B) molecular compound choices: (A) ionic compound 56. The representative unit of these compounds is the molecule. 57. These compounds tend to have higher conductivity. 58. Li3PO4 59. Have melting points usually below 300C 60. Generally involve entirely nonmetallic elements 61. Involve single, double, and/or triple bonds 62. C2H5OH 63. NH4Cl 64. SCl2 65. Li+ and S2– would be involved in this kind of compound. page A-52 – DC – T2 – BOOK Unit Six Practice Test (Short Answer) PT – DC – U6 Name____________________________________________ Period_______________ LPS Standard(s): --Structures and Geometry. State Standard(s): 12.3.2d For each of the following molecules or anions, draw the correct electron dot structure for the molecule or anion. 1. PCl5 (phosphorus pentachloride) 3. NO21– (nitrite ion) 2. XeF2 (xenon difluoride) 4. ClO4– (perchlorate ion) Short Answer. Answer the following questions. 5. Draw a structural formula of carbon dioxide. Also use + and – in the picture. Make sure the picture is understandable and clear as to which atom is which. Then identify whether the molecule is polar or nonpolar and explain why. 6. Draw a structural formula of carbon monoxide. Also use + and – in the picture. Make sure the picture is understandable and clear as to which atom is which. Then identify whether the molecule is polar or nonpolar and explain why. page A-53 – DC – T2 – BOOK Unit Seven Practice Test (Multiple Choice) PT – DC – U7 NOTE: For all calculations on this test, use the constants and values from the periodic table of elements provided to you. Multiple Choice. On the scantron for each question, fill in the rectangle of the corresponding letter of the answer that best completes or answers the statement or question in the adjacent corresponding blank. LPS Standard(s): 12.2.4c State Standard(s): 12.1.3e; 12.1.2d 1. A gas has a pressure of 555 kPa at 227C. What will its pressure be at 53C if the volume does not change? (A) 59 kPa (B) 130 kPa (C) 362 kPa (D) 578 kPa (E) 2380 kPa 2. A 30-g mass of carbon dioxide occupies 17.0 L at a pressure of 156 kPa. Find the volume of carbon dioxide when the pressure is increased to 215 kPa at the same temperature. (A) 0.000507 L (B) 0.0811 L (C) 12.3 L (D) 23.4 L (E) 1970 L 3. At standard pressure and temperature, a gas occupies 22.4 L. If the volume and pressure of the gas are changed to 3.50 L and 70.1 kPa, respectively, what will the new temperature of the gas be in degrees Celsius? (B) 0C (C) 29.5C (D) 2257C (E) 2530C (A) –243.6C 4. A gas occupies a volume of 0.70 L at 50.C. What volume will the gas occupy at 100.C if the pressure does not change? (A) 0.35 L (B) 0.61 L (C) 0.81 L (D) 1.4 L (E) 2.8 L 5. If a gas has a volume of 24.0 L at 20.C, at what temperature, in degrees Celsius, will the gas have if the volume of the gas is increased to 55.0 L and the pressure does not change? (A) –227°C (B) 8.7°C (C) 46°C (D) 398°C (E) 671°C 6. Why does the pressure inside a container of gas increase if more gas is added to the container? (A) because there is a corresponding increase in the number of particles striking an area of the wall of the container per unit time (B) because there is a corresponding increase in the temperature (C) because there is a corresponding decrease in volume (D) because there is a corresponding increase in the force of the collisions between the particles and the walls of the container 7. If the volume of a container holding a gas is reduced, the pressure within the container ___. (A) decreases (B) increases (C) remains the same 8. The temperature of a gas ___ when the gas is compressed. (A) decreases (B) increases (C) remains the same 9. If a balloon is rubbed vigorously and the volume remains constant, the pressure of the air inside the balloon ___. (A) decreases (B) increases (C) remains the same page A-54 – DC – T2 – BOOK 10. If there is no change in pressure for a sample of gas at 40C, which temperature will cause a decrease in the volume of this gas? (A) 260 K (B) 280 K (C) 300 K (D) All of the temperatures listed. 11. If a gas’s temperature is decreased and volume is held constant, its pressure will ___. (A) decrease (B) increase (C) remain the same 12. If the volume of a gas is increased and its temperature remains constant, what will happen to its pressure? (A) decrease (B) increase (C) remain the same 13. If a capped syringe is heated, in which direction will the syringe plunger move? (A) In (B) Out (C) No sliding will occur. 14. How do gas particles respond to an increase in pressure? (A) An increase in kinetic energy and an increase in temperature (B) A decrease in kinetic energy and a decrease in volume (C) An increase in temperature and an increase in volume (D) A decrease in kinetic energy and an increase in temperature 15. One way to increase the pressure of a gas is to ___. (A) lower the kinetic energy of the gas particles (B) decrease the number of gas particles (C) increase the temperature of the gas (D) increase the volume of the gas LPS Standard(s): --- State Standard(s): 12.1.3 16. A sealed vessel contains 0.200 mol of oxygen gas, 0.100 mol of nitrogen gas, and 0.200 mol of argon gas. The total pressure of the gas mixture is 5.00 atm. The partial pressure of the argon is ___ atm. (A) 0.200 (B) 0.500 (C) 1.00 (D) 2.00 (E) 5.00 17. Atmospheric pressure on the surface of Mars is 6.0 torr. The partial pressure of carbon dioxide is 5.7 torr. What percent of the Martian atmosphere is carbon dioxide? (A) 5.0% (B) 6.0% (C) 95% (D) 96% (E) 98% 18. A breathing mixture used by deep-sea divers contains helium, oxygen, and carbon dioxide. What is the total pressure of the air if PHe = 84.5 kPa, PO2 = 2.8 kPa,and PCO2 = 0.1 kPa? (A) 87.4 kPa (B) 23.7 kPa (C) 81.6 kPa (D) None of the choices are correct. 19. What happens to the partial pressure of nitrogen in the air if the air temperature is decreased? (A) It decreases. (B) It increases. (C) It remains the same. 20. If oxygen is added to a scuba tank, what happens to the total pressure of the air in the tank? (A) It decreases. (B) It increases. (C) It remains the same. LPS Standard(s): 12.2.4c 21. State Standard(s): 12.1.3e; 12.1.2d At standard temperature and pressure, 22.4 L of a gas is discovered to have a mass of 64.06 grams. What is the gas? (B) CO2 (C) NO2 (D) P2O5 (E) Cl2 (A) SO2 page A-55 – DC – T2 – BOOK 22. Which of the following gases would have the largest volume at 10C and 760 mm Hg? (B) 10 g He (C) 64 g O2 (D) 56 g N2 (A) 88 g CO2 23. If 0.214 mol of argon occupies 652 mL at a given temperature and pressure, what is the volume of 0.214 mol of butane at the same temperature and pressure? (A) 652 mL (B) 1304 mL (C) 3047 mL (D) 6094 mL 24. Calculate the approximate temperature of a 0.500 mol sample of gas at 750. mm Hg and a volume of 12.0 L. (B) 11C (C) 15C (D) 288C (A) –7C 25. What is the approximate volume of gas in a 1.50 mol sample that exerts a pressure of 0.922 atm and has a temperature of 10.0ºC? (A) 13.0 L (B) 14.2 L (C) 37.8 L (D) 378 L 26. In collisions between ideal gas molecules, the total energy of the gas ___. (A) decreases (B) increases significantly (C) increases slightly (D) remains the same 27. The ideal gas law will be least likely to predict the behavior of ___ gas. (A) helium (B) hydrogen (C) neon (D) water vapor 28. What is the pressure exerted by 1.2 mol of a gas with a temperature of 20.ºC and a volume of 9.5 L? (A) 0.030 atm (B) 1.0 atm (C) 3.0 atm (D) 30. atm 29. What does the ideal gas law lead a scientist to calculate that the other gas laws cannot lead to calculating? (A) mass (B) temperature (C) volume (D) energy (E) pressure 30. A sample of gas at 25ºC has a volume of 11 L and exerts a pressure of 660 mm Hg. How many moles of gas are in the sample? (A) 0.39 mol (B) 3.9 mol (C) 9.3 mol (D) 87 mol LPS Standard(s): 12.2.4b State Standard(s): 12.1.2a; 12.1.2b 31. At low temperatures and pressures, how does the volume of a real gas compare with the volume that would be predicted for an ideal gas under the same conditions? (A) It is much greater. (B) It is much less. (C) There is no difference. 32. Another way of describing temperature is as ___ kinetic energy. (A) adequate (B) average (C) static (D) total 33. An ideal gas cannot be ___. (A) expanded (B) compressed (C) heated (D) frozen 34. Which term is used to describe changing a gas into a liquid? (A) condensation (B) freezing (C) melting (D) vaporization 35. Which of the following is not one of the assumptions of kinetic theory? (A) Particles in a gas are assumed to have a significant volume. (B) All gas particles move in constant random motion. (C) Gases consist of hard spherical particles. (D) No attractive and repulsive forces exist between gas particles. page A-56 – DC – T2 – BOOK LPS Standard(s): --- State Standard(s): 12.3.5; 12.3.2e 36. A 5.00-g sample of liquid water at 25.0C is heated by the addition of 84.0 J of energy. What is the final temperature of the water if the specific heat capacity of liquid water is 4.184 J/(gC)? (A) 4.02C (B) 21.0C (C) 29.0C (D) 95.2C 37. If the heat involved in a chemical reaction has a negative sign, ___. (A) heat is lost to the surroundings (B) heat is gained from the surroundings (C) no heat is exchanged in the process 38. If you want to cool a hot drink, it is best to use a spoon with a relatively ___ specific heat. (A) low (B) high (C) The specific heat of the spoon does not matter. 39. When heat is added to melting ice, its temperature ___. (A) decreases (B) increases (C) remains the same 40. If the heat of a substance decreases, its temperature ___. (A) decreases (B) increases (C) stays the same 41. Which of the following would you measure using a calorimeter? (A) specific heat (B) weight (C) specific gravity (D) density 42. Compared to 100 grams of iron, a 10-gram sample of iron has ___. (A) a higher specific heat (B) a lower specific heat (C) the same specific heat 43. How many calories are there in 148 Joules? (A) 6.61 J (B) 35.4 J (C) 148 J (D) 619 J (E) 3320 J 44. In an exothermic reaction, the energy stored in the chemical bonds of the reactants is ___. (A) equal to the energy stored in the bonds of the products (B) greater than the energy stored in the bonds of the products (C) less than the energy stored in the bonds of the products (D) less than the heat released (E) less than the heat absorbed 45. If a substance gets colder, what happens to the average kinetic energy of the particles of the substance? (A) It decreases. (B) It increases. (C) It remains the same. 46. How much energy would be released as the temperature of 150. grams of copper (specific heat = 0.0924 cal/(gC) drops from 79.0°C to 50.0°C? (A) 3.00 cal (B) 402 cal (C) 3650 cal (D) 3950 cal 47. What is the energy required to melt one mole of a solid at its melting point? (A) molar heat of melting (C) molar heat of vaporization (B) molar heat of fusion (D) molar heat of solution 48. A process that releases heat is a(n) ___ process. (A) ectothermic (B) endothermic (C) exothermic page A-57 – DC – T2 – BOOK (D) polythermic 49. Materials with a very high specific heat capacity can absorb little energy and show ___ change in temperature. (A) great (B) little (C) no 50. The freezing of a liquid is a(n) ___ process. (A) ectothermic (B) endothermic (C) exothermic (D) polythermic 51. The amount of heat absorbed by a vaporizing liquid ___ the amount of heat lost by the same vapor if it is condensing. (A) is the same as (B) is less than (C) is greater than 52. Which of the following equations correctly represents an exothermic reaction? (B) A + B + heat C + D (A) A + B C + D + heat 53. If Mr. Geist jumps into a cold pool, the pool will be experiencing a(n) ___ process. (A) ectothermic (B) endothermic (C) exothermic (D) polythermic 54. When heat is released from steam, its temperature ___. (A) increases (C) depends on the amount of water (B) decreases (D) remains constant 55. If you were to touch the flask in which an exothermic reaction were occurring, ___. (A) the flask would feel the same as before the reaction started (B) the flask would probably feel cooler than before the reaction started (C) the flask would probably feel warmer than before the reaction started LPS Standard(s): --- State Standard(s): 12.1.2a 56. Which of the following is NOT a result of water’s high specific heat capacity? (A) The temperature of water goes up rapidly as it absorbs solar energy. (B) The temperature of cities near large bodies of water are moderated. (C) For the same increase in temperature, iron needs to absorb only about one-tenth as much energy as water. (D) Water is an excellent medium for the storage of solar energy. 57. How does the vapor pressure of water compare with the vapor pressures of other molecules of similar size? (A) It is higher. (B) It is lower. (C) It is about the same. 58. Which of the following is primarily responsible for the high heat of vaporization of water? (A) polar covalent bonds (C) ionic bonds (B) dispersion forces (D) hydrogen bonds 59. Which of the following is primarily responsible for the low vapor pressure of water? (A) polar covalent bonds (C) ionic bonds (B) dispersion forces (D) hydrogen bonds 60. At what temperature does liquid water freeze? (B) 4C (C) 37C (D) 100C (A) 0C page A-58 – DC – T2 – BOOK LPS Standard(s): --- State Standard(s): 12.3.2; 12.1.2b 61. Which of the following substances is the most insoluble in water? (A) HCl (B) CH4 (C) CaCO3 (D) NaOH 62. Which of the following compounds would be the most soluble in water? (B) PCl5 (C) NaCl (D) SF6 (A) CO2 63. Which of the following substances is the most soluble in water? (A) sodium fluoride (B) oxygen gas (C) carbon (D) liquid bromine 64. What is the term for the substance being dissolved into a solution? (A) solvator (B) solute (C) emulsifier (D) solvent 65. Why are two polar substances able to dissolve in each other? (A) Polar substances cannot dissolve in each other. (B) They combine to produce a polar substance. (C) There is no attractive or repulsive force between them. (D) There are attractive and repulsive forces between them. LPS Standard(s): --- State Standard(s): 12.1.3e 66. If the temperature of a liquid decreases, the solubility of a solid in the liquid generally ___. (A) decreases (B) increases (C) stays the same 67. If the pressure above a liquid decreases, the solubility of a gas in the liquid ___. (A) decreases (B) increases (C) stays the same 68. If the temperature of a liquid decreases, the solubility of a gas in the liquid generally ___. (A) decreases (B) increases (C) stays the same 69. The solubility of a gas at constant temperature and 2.0 atm pressure is 3.28 g/mL. What would the solubility of the gas be at 4.0 atm pressure and without a temperature change? (A) 0.410 g/mL (B) 0.610 g/mL (C) 2.44 g/mL (D) 6.56 g/mL 70. Which of the following substances is more soluble in hot water than in cold water? (A) CO (B) O2 (C) NH3 (D) LiCl page A-59 – DC – T2 – BOOK Unit Seven Practice Test (Short Answer/Calculation) PT – DC – U7 Name____________________________________________ Period_______________ Short answer/Calculation. Answer the following questions. For questions involving mathematics, show work or receive no credit, and include correct significant figures, decimal places, and proper units. LPS Standard(s): --- State Standard(s): 12.1.3e 1. How could you experimentally determine if a solution is unsaturated, saturated, or supersaturated? 2. Among the following gases, which of the following, if they possess the same temperature, would have the slowest velocity (rate of effusion)? Circle your answer AND explain your answer. Cl2 3. C2H6 CO2 CO How much heat (in kJ) does it take to convert 1550 g of water at 10.0C to steam at 105.0C? Cice = 2.1 J/(gC); Cwater = 4.184 J/(gC); Csteam = 1.7 J/(gC); Hfus = 6.01 kJ/mol; Hvap = 40.7 kJ/mol page A-60 – DC – T2 – BOOK 4. During the metabolic process called respiration, your body obtains energy from the breakdown of glucose as shown below. C6H12O6(aq) + 6O2(g) 6H2O(l) + 6CO2(g) What volume of O2, measured at 37C and 790.0 torr pressure, is required to react with 1.00 g of glucose (C6H12O6)? Express the volume in milliliters. 5. Why is ice less dense than water? 6. Explain why water has a relatively high heat of vaporization. page A-61 – DC – T2 – BOOK Unit Eight Practice Test (Multiple Choice) PT – DC – U8 LPS Standard(s): --Multiple Choice. State Standard(s): 12.1.2 Identify the letter of the choice that best completes the statement or answers the question. 1. When an acid reacts with a base, what is one compound that is always formed? (A) salt (B) sugar (C) carbohydrate (D) protein 2. Which of the following is a property of an acid? (A) reactive with bases (B) nonreactive with bases (C) nonelectrolyte (D) nonconductive 3. What is the charge on the hydronium ion? (A) 2– (B) 1– (C) 0 (D) 1+ (E) 2+ 4. Which of the following is a property of an acid? (A) sour taste (B) strong color (C) nonelectrolyte (D) unreactive 5. What is the name of the compound H3PO3? (A) hydrophosphoric acid (C) phosphoric acid (B) hydrophosphorus acid (D) phosphorous acid 6. What is the name of the compound KOH? (A) potassium oxygen hydride (B) potassium (I) hydroxide (C) potassium (II) hydroxide (D) potassium hydroxide 7. What is the formula of lithium hydroxide? (D) Li(OH)2 (A) LiH (B) LiOH (C) LiH2 8. What is the formula of chloric acid? (A) HCl (B) HClO (C) HClO2 9. 10. What is the formula of carbonic acid? (C) H2CO3 (A) HC (B) H4C (D) HClO3 (D) H3CO3 What is the formula of hydrobromic acid? (C) HBrO3 (D) HBrO4 (A) HBr (B) HBrO2 LPS Standard(s): --- State Standard(s): 12.1.2 11. An Arrhenius acid ___. (B) produces OH– (A) produces H+ (C) accepts H+ (D) accepts OH– 12. An Arrhenius base ___. (B) produces OH– (A) produces H+ (C) accepts H+ (D) accepts OH– 13. A Bronsted-Lowry acid ___. (B) donates OH– (A) donates H+ (C) accepts H+ (D) accepts OH– 14. A Bronsted-Lowry base ___. (B) donates OH– (A) donates H+ (C) accepts H+ (D) accepts OH– page A-62 – DC – T2 – BOOK 15. A substance that can behave as an acid or a base is ___. (A) andropic (B) amphoteric (C) analgesic (D) analytic 16. Which of the following is a Bronsted-Lowry base, but not an Arrhenius base? (B) NaOH (C) LiOH (D) HCl (A) NH3 17. In the chemical equation LiOH + H2O OH– + LiOH2+, what is the conjugate acid? (A) LiOH (B) H2O (C) OH– (D) LiOH2+ 18. In the chemical equation NH3 + H2O NH4+ + OH–, what is the conjugate base? (B) H2O (C) NH4+ (D) OH– (A) NH3 19. In the chemical equation HCl + H2O Cl– + H3O+, what is the conjugate acid? (C) Cl– (D) H3O+ (A) HCl (B) H2O 20. In the chemical equation HCl + H2O Cl– + H3O+, what is the conjugate base? (A) HCl (B) H2O (C) Cl– (D) H3O+ LPS Standard(s): --Identification. State Standard(s): 12.1.2 Identify the following properties, definitions, or compounds as being those of (A) acidic, (B) basic, or (C) neutral. 21. Solution where pOH = 2 26. Solution where pH = 3.8 22. HNO3 27. Solution where pOH = 7 23. H2O 28. Produce hydronium ions in solution 24. H2SO4 29. Solution where [H+] = 3.1 x 10–10 M 25. KOH 30. Solution where [OH–] = 4.8 x 10–4 M 31. Distilled water containing only HC2H3O2 dissolved in it 32. Distilled water containing only NaOH dissolved in it 33. In HNO3 + H2O NO3– + H3O+, what H2O would be 34. A solution containing substantially more hydronium ions than hydroxide ions 35. A solution containing substantially more hydrogen ion acceptors than hydrogen ion donors 36. Hydrogen ion concentration of 1 10–7 M 37. Hydroxide ion concentration of 3 10–4 M 38. Hydronium ion concentration of 3 10–4 M 39. pH = 6.0 40. pOH = 6.0 page A-63 – DC – T2 – BOOK Unit Eight Practice Test (Short Answer) PT – DC – U8 Name_________________________________________________ Period__________________ LPS Standard(s): --Calculations. State Standard(s): 12.3.3a Solve the following problems. Show work or receive no credit. Include proper units and significant figures and/or decimal places. 41. What volume (in milliliters) of 0.275M sulfuric acid is needed to neutralize 17.2 mL of 0.550M sodium hydroxide? 42. A 100.0 mL solution of sodium hydroxide is completely neutralized by 100.0 mL of 0.5000M phosphoric acid. What is the concentration of the sodium hydroxide? LPS Standard(s): --- State Standard(s): 12.1.2 A lab assistant tests poolwater near a park and determines it to have a pOH of 5.9. 43. What is the hydrogen-ion concentration [H+] of the solution? 44. What is the hydroxide-ion concentration [OH-] of the solution? 45. What is the pH of the solution? page A-64 – DC – T2 – BOOK 46. What is the ion-product constant of the solution? 47. Is this solution acidic, basic, or neutral? LPS Standard(s): --- State Standard(s): 12.1.2 Identification. Identify the following piece of equipment involved in the titration process. 48. 49. 50. Equipment: 48. _________________________________________________________________________ 49. _________________________________________________________________________ 50. _________________________________________________________________________ page A-65 – DC – T2 – BOOK Appendix O Practice Test Keys Unit Five Practice Test Key PTK – DC – U5 1. D 2. C 3. A 4. E 5. B 6. D 7. A 8. C 9. E 10. C 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. A C A A B C A C B A 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. C C C D C B E E A D 31. 32. 33. 34. 35. 36. 37. 38. 39. 40. D A B C C A B A D A 41. 42. 43. 44. 45. 46. 47. 48. 49. 50. 51. 52. 53. 54. 55. 56. 57. 58. 59. 60. B A B C D B D C A C A B A C B D E A A B 61. 62. 63. 64. 65. 66. 67. 68. 69. 70. A C B B C A D A A D 71. 72. 73. 74. 75. 76. 77. 78. 79. 80. A B C A C C B A A A 81. 82. 83. 84. 85. 86. 87. 88. 89. 90. A A A C A B B C B E 91. C 92. C 93. B 94. A 95. B 96. B 97. A 98. B 99. A 100. C Unit Six Practice Test Key (Multiple Choice) PTK – DC – U6 1. 2. 3. 4. 5. 6. 7. D D D D B A D 8. B 9. B 10. C 11. A 12. D 13. C 14. C 15. 16. 17. 18. 19. 20. 21. A A B C D C C 22. 23. 24. 25. 26. 27. 28. D C D C C B B 29. 30. 31. 32. 33. 34. 35. 36. 37. 38. 39. 40. 41. 42. B C A B D C C A A C D B E D 43. 44. 45. 46. 47. 48. 49. E D C D A C A 50. 51. 52. 53. 54. 55. 56. B A C B A A B 57. 58. 59. 60. 61. 62. 63. A A B B B B A 64. B 65. A Unit Six Practice Test Key (Short Answer) PTK – DC – U6 1. 2. 3. 4. 5. Nonpolar; partial charges cancel out because of symmetry 6. Polar; partial charges do not cancel out because of creation of poles on each side page A-66 – DC – T2 – BOOK Unit Seven Practice Test Key (Multiple Choice) PTK – DC – U7 1. 2. 3. 4. 5. 6. 7. C C A C D A B 8. B 9. B 10. D 11. A 12. A 13. B 14. A 15. 16. 17. 18. 19. 20. 21. C D C A A B A 22. 23. 24. 25. 26. 27. 28. B A C C D D C 29. 30. 31. 32. 33. 34. 35. A A B B D A A 36. 37. 38. 39. 40. 41. 42. C A A C A A C 43. 44. 45. 46. 47. 48. 49. B B A B B C B 50. 51. 52. 53. 54. 55. 56. C A A B B C A 57. 58. 59. 60. 61. 62. 63. B D D A B C A 64. 65. 66. 67. 68. 69. 70. B D A A B D D Unit Seven Practice Test Key (Short Answer/Calculation) PTK – DC – U7 1. Add more solute. If the solute dissolves in the solution, the solution is unsaturated. If the solute will not dissolve in and goes to the bottom, the solution is saturated. If the the solute causes the solution to crystallize, the solution is supersaturated. 2. Cl2: 2(35.453) = 70.906 g/mol C2H6: 2(12.011) + 6(1.0079) = 24.022 + 6.0474 = 30.069 g/mol CO2: 1(12.011) + 2(15.999) = 12.011 + 31.998 = 44.009 g/mol CO: 1(12.011) + 1(15.999) = 12.011 + 15.999 = 28.010 g/mol Cl2 would have the slowest velocity (rate of effusion) because it has the greatest molar mass. 3. q3: Heat needed for water to elevate from 10.0C to 100.0C q4: Heat needed to vaporize water into steam q5: Heat needed for water to elevate from 100.0C to 105.0C 1550 g 4.184 J 100.0C 10.0C 1550 g 4.184 J 90.0C 1 g C 1 1 g C 1 = 584000 J = 584 kJ 1550 g 1 mol 40.7 kJ q4 3500 kJ 1 18.015 g 1 mol 1550 g 1.7 J 105.0C 100.0C 1550 g 1.7 J 5.0C q 5 mCT 1 g C 1 1 g C 1 = 13000 J = 13 kJ q 3 mCT Total heat = 584 kJ + 3500 kJ + 13 kJ = 4097 kJ 4. 1.00 g C 6H12 O 6 1 mol C 6H12 O 6 6 mol O 2 0.0333 mol O 2 1 180.155 g C 6H12 O 6 1 mol C 6H12 O 6 P: 790.0 torr (mm Hg) n: 0.0333 mol O2 R: 62.396 (Ltorr)/(Kmol) T: 37C + 273.15 = 310. K PV = nRT (790.0 torr)V = (0.0333 mol O2)(62.396 (Ltorr)/(Kmol))(310. K) page A-67 – DC – T2 – BOOK V 0.0333 mol62.396 (L torr)/(K mol)310. K 0.815 L 815 mL O 2 790.0 torr 5. The structure of ice is a very regular, open framework in which the water molecules are farther apart from each other than they are in liquid water, in great part due to hydrogen bonds. When ice melts, this open framework collapses and the water molecules move closer together. As a result, the water is denser than the ice. 6. Water has a high heat of vaporization as a result of its hydrogen bonding. Because of its extensive network of hydrogen bonds, the molecules of water are held together more tightly than are the molecules of many other liquids. The attractive force of these hydrogen bonds must be overcome in order for water to vaporize. Unit Eight Practice Test Key (Multiple Choice) PTK – DC – U8 1. 2. 3. 4. A A D A 5. 6. 7. 8. D D B D 9. C 10. A 11. A 12. B 13. 14. 15. 16. A C B A 17. 18. 19. 20. D D D C 21. 22. 23. 24. B A C A 25. 26. 27. 28. B A C A 29. 30. 31. 32. B B A B 33. 34. 35. 36. B A B C Unit Eight Practice Test Key (Short Answer) PTK – DC – U8 41. H2SO4(aq) + 2NaOH(aq) Na2SO4(aq) + 2H2O(l) 1 L H2 SO 4 0.550 mol NaOH 0.0172 L NaOH 1 mol H2 SO 4 1L NaOH 1 2 mol NaOH 0.275 mol H2 SO 4 = 0.0172 L H2SO4 = 17.2 mL H2SO4 42. H3PO4(aq) + 3NaOH(aq) Na3PO4(aq) + 3H2O(l) 0.5000 mol H3PO 4 0.1000 L H3PO 4 3 mol NaOH 1 1 L H3PO 4 1 1 mol H3PO 4 0.1000 L NaOH = 1.500 M NaOH 43. [H+] = 10–8.1 = 7.94 10–9 M (8.1 from answer to question 45) 44. [OH-] = 10–5.9 = 1.26 10–6 M 45. pH + pOH = 14 pH + 5.9 = 14 pH = 8.1 48. Funnel 49. Ring stand 46. 1 10–14 M2 50. Beaker 47. Basic page A-68 – DC – T2 – BOOK 37. 38. 39. 40. B A A B
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