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Chemistry: First Semester Exam Prep #4 (Choose the BEST answer.) 1. Which unit is NOT used for volume? a) L b) mL c) mm d) cm3 2. Which piece of equipment would be best to measure 56.7 mL of water? a) graduated cylinder b) balance c) test tube d) electronic balance 3. Which of these are the smallest? a) 1 liter b) 1microliter c) 1milliliter d) 1kiloliter 4. The prefix kilo means: a) 100 times larger b) 1000 times larger c) 1 000 000 times larger d) 10 times smaller 5. The metric prefix milli means: a) 100 times larger b) 1000 times larger c) 1000 times smaller d) 100 times smaller 6. The number 2 X 101 expressed in standard notation is: a) 200 b) 20 c) 2 d) 0.02 7. The number 3 X 10 -3 expressed in standard notation is a) 300 b) 0.300 c) 0.003 d) 3.000 8. How many of the zeros in the measurement 0.000040300 are significant? a) 8 b) 6 c) 5 d) 3 9. The number of significant figures in the measurement 0.070g is: a) 1 b) 2 c) 3 d) 4 10. When adding and subtracting measurements one should limit and round answers to: a) the least number of significant figures in any of the measurements b) the least number of decimal places in any of the measurements c) the tenths place d) three significant figures 11. When multiplying and dividing measurements, one should limit and round answers to: a) the least number of significant figures in any of the factors b) the least number of decimal places in any of the measurements c) the tenths place d) three significant figures 12. When measuring the volume of a liquid in a graduated cylinder, the measurement should be read from the: a) top of the meniscus b) top of the cylinder c) bottom of the meniscus d) table top 13. If the mass of a dry beaker is 19.02 grams and increases to 22.40 grams when a sample is added, what is the mass of the sample? a) 22.40 g b) 41.42 g c) 3.38 g d) 1.10 g 14. The mass of a sample is 4.11grams. The volume is 2.00cm3. What is its density? a) 8.22g/cm3 b) 2.06g/cm3 c) 0.49g/cm3 d) not enough information 15. A student estimated a mass to be 250g, but upon carefully measuring it, found it to be 240g, What is the percent error of the estimated mass? a) 4.0% b) 4.2% c) -4.0% d) -4.2% 16. The number of electrons in a sulfur atom is: a. 32.06 b. 32 c. 16 d. 48 1 Chemistry: First Semester Exam Prep #4 (Choose the BEST answer.) 17. The sum of the protons and neutrons in a lithium atom is: a. 12 b. 13 c. 9 d. 7 18. The electron was discovered by a. Bohr b. Thompson c. Maxwell d. Dalton 19. The number of neutrons in a fluorine atom is: a. 19 b. 10 c. 9 d. 18.998 20. The net charge on any atom is: a. positive b. negative c. neutral d. depends on the element 21. The sum of the protons and neutrons in an atom equals the: a. atomic number b. number of electrons c. atomic mass d. mass number 22. Which of these statements is false? a. Electrons have a negative charge. b. Electrons have a mass of 1 amu. c. The nucleus of an atom is positively charged. d. The neutron is found in the nucleus of an atom. 23. An atom with atomic number 48 and mass number 120 contains: a. 48 protons, 48 electrons, and 72 neutrons b. 72 protons, 48 electrons, and 48 neutrons c. 120 protons, 48 electrons, and 72 neutrons d. 72 protons, 72 electrons, and 48 neutrons 24. An element which has a mass number of 23 and has 13 neutrons is the element: a. Lithium b. Potassium c. Magnesium d. Sodium d. Becquerel 26. Rutherford's alpha scattering experiment showed that the charge on the nucleus of an atom must be: a. positive b. neutral c. negative d. none of the above 27. The nucleus of the atom has a. a high density b. a low density c. a negative charge d. no charge 28. An ion always contains a. unequal number of protons and electrons b. equal number of protons and electrons c. unequal number of protons and neutrons d. equal number of protons and neutrons 29. The whole number that is closest to the atomic mass of an atom is the a. atomic number b. mass number c. Avogadro's number d. number of neutrons 30. Which subatomic particle did Thompson include in his "plumb pudding model"? a. protons b. neutrons c. electrons d. none of the above 31. Which of the following types of reactions results in a single product? a. combination b. decomposition c. single replacement a. double replacement 32. In the reaction 2KClO3 2KCl + 3O2 oxygen is a ___. a. reactant b. product c. coefficient d. subscript 33. Supersaturated solutions are characterized by a. being super hot b. having great stability c. having a larger amount of solute than can be dissolved d. being able to exist at super-low temperatures 25. The experiment that revealed the charge of the electron involved the use of a. gold foil b. Rutherford c. the Cathode ray 2 Chemistry: First Semester Exam Prep #4 (Choose the BEST answer.) 34. In the reaction: N2(g) + 3H2(g) 2NH3(g) + heat a. b. c. d. the reaction is both endothermic and exothermic the reaction is endothermic the reaction is exothermic the reaction is neither endothermic or exothermic 35. How many oxygen atoms are in Al2(SO4)3? a. 3 atoms of O b. 4 atoms of O c. 7 atoms of O d. d. 12 atoms of O 36. The atomic mass of an element: a. depends on the number of isotopes of that element. b. depends on the mass of each isotope of that element. c. depends on the relative abundance of isotopes of that element. d. all of the above. 39. The weighted average mass of all the atoms (isotopes) of an element is the: a. atomic mass b. atomic number c. electron number d. neutron number 40. The percentages of isotopes found in a sample of an element is given below. 22% Carbon-14 with a mass of 14 amu 78% Carbon-12 with a mass of 12 amu The correct method for finding the average atomic mass of the sample would be: a. (22) (14) - (78) (12) b. (0.22) (14) - (0.78) (12) c. (0.22) (14) + (0.78) (12) d. (22) (14) + (78) (12) 41. How do the isotopes hydrogen-2 and hydrogen-3 differ? a. Hydrogen-3 has one more electron than hydrogen-2. b. Hydrogen-3 has two neutrons. c. Hydrogen-3 has three protons. d. Hydrogen-2 has no protons. 42. How many molecules are in 4.50 moles of H2O? a. 4.50 b. 2.71 x 1024 c. 6.02 x 1023 d. 3.00 43. How many moles are in 8.5 x 10 25 molecules of CO? a. 1.4 x 102 b. 7.1 x 10-3 c. 5.1 x 1049 d. 8.5 x 1025 44. What is the molar mass of CO2? a. 36.0 g b. 22.0 g c. 44.0 g d. 6.02 x 1023 g 45. The sum of the atomic masses of all the atoms in a compound is called the ___. a. molar mass b. empirical formula c. molar volume d. percentage composition 46. The number of atoms in one mole of an element is equal to___. a. a measure b. a gram c. a formula unit d. Avogadro's number 47. It is possible to convert moles to particles by: a. multiplying by 6.02 x1023 b. dividing by 6.02 x1023 c. multiplying by the molar mass d. dividing by the molar mass 48. How many molecules of sulfur dioxide are present in 1.60 moles of sulfur dioxide? a. 9.63 x 1023 b. 102.1 x101 c. 7.62 x 101 d. 3.76 x 1023 49. Find the number of moles in 3.30 g of (NH 4)2SO4 a. 132.1 b. 40.0 c. 0.279 d. 0.0250 50. Which contains more atoms? a. 1.00 mole H2O2 b. 1.00 mol C2H6 c. 1.00 mol CO d. 1.00 mol K 51. An element with seven valence electrons would likely be: a. an alkaline earth metal b. an alkali metal c. a noble gas d. a halogen 3 Chemistry: First Semester Exam Prep #4 (Choose the BEST answer.) 52. The most stable atoms are those of the a. metals b. metalloids c. noble gases d. nonmetals 53. The ion with a charge of +1 and the same electron configuration as argon is a. potassium b. sodium c. neon d. magnesium 54. The tendency to lose electrons ______________ as we move across a period on the periodic table a. increases b. remains the same c. decreases d. no trend exists 55. Generally, members of a ____________ have the same number of valence electrons. a. period b. series c. row d. family 56. An element which is considered to be a metalloid is: a. Boron b. Calcium c. Oxygen d. Sodium 57. The element Iodine is a a. period 5 alkali metal b. period 4 halogen c. period 5 halogen d. period 5 transition metal 58. Sodium and Potassium have similar properties because they have the same ___. a. atomic radius b. number of valence electrons c. ionization energy d. electronegativity 59. The maximum number of valence electrons in an atom is ___. a. 2 b. 4 c. 8 d. 12 60. The likeliest charge an atom with 2 valence electrons would develop is a. 2+ b. 6+ c. 2d. 661. The likeliest charge an atom with 6 valence electrons will develop is a. 2+ b. 6+ c. 2d. 662. The likeliest charge of an ion of the element Bromine is a. -1 b. +1 c. +2 d. -2 63. Metals tend to __. a. gain electrons and become positively charged cations b. lose electrons and become negatively charged cations c. lose electrons and become positively charged cations d. gain electrons and become negatively charged anions. 64. The basis of the ionic bond is the ___. a. sharing of an electron pair b. electrical attraction between oppositely charged ions. c. absorption of energy d. absorption of water into their solid structures 65. Elements tend to gain or lose electrons in order to acquire the electron configuration of a a. halogen b. transition metal c. noble gas d. nonmetal 69. When sodium combines with chlorine to form sodium chloride, the sodium attains the electron configuration of a. helium b. neon c. argon d. lithium 70. The type of chemical bonding in which electron pairs are shared is a. ionic b. covalent c. metallic d. none of the above 4 Chemistry: First Semester Exam Prep #4 (Choose the BEST answer.) 71. Which of the following correctly matches the names and formulas of the compounds? a. AlCl3 . aluminum trichloride and N2O4 , nitrogen oxide b. AlCl3 , aluminum trichloride and N2O4 dinitrogen tetroxide c. AlCl3, aluminum chloride and N2O4 nitrogen oxide d. AlCl3 , aluminum chloride and N2O4 dinitrogen tetroxide d. balanced 74. Choose the correct set of coefficients needed to balance the following equation: PCl5 PCl3 + Cl2 a. 2,2,1 b. 1,2,1 c. 1,1,2 d. balanced 75. Choose the correct set of coefficients needed to balance the following equation: PbCl2 + Li2SO4 LiCl + PbSO4 a. 1,1,2,1 b. 1,2,2,1 c. 2,2,1,1 d. balanced 73. Choose the correct set of coefficients needed to balance the following equation: H2 + O2 ---> H2O a. 2,1,1 b. 1,1,2 c. 2,1,2 Short Answer: 76. A copper penny has a mass of 3.1 g and a volume of 0.35cm3. What is the density of copper? 8.86 g/cm3 77. A liquid has a density of 4.8 g/l. What is the mass of a 2 liter sample? 9.6 g 78. What is the volume of a substance that has a mass of 80 g and a density of 10 g/cm3? 10 cm3 79. Calculate the following quantities: a. 1,100 cm = 1.1 m b. 1 m = 1000 mm c. 10 m = 1000 cm d. 2.5 km = 2500 m. e. 4.05 kg = 4050 g 80. Indicate the number of significant figures in each of the following: a. 12600 3 b. 0.09 1 c. 2001 4 d. 0.00500100 6 83. Which of the following are empirical formulas and which are molecular formulas? 81. The accepted value or true value for the density of lead (Pb) is 11.35 g/ml. Your experimental value or observed value found during a class lab is 9.65 g/mL 84. Find the empirical formula of each compound from its % composition. a. 72.4 % Fe and 27.6% O ___________Fe What is the error of your measurement? b. 94.1% O and 5.9% ___________Fe -1.7 What is the percent error of your measurement? (-) 15% 82. Name the two temperature scales used in science? Give the freezing pt., and boiling pt. of water for each of them. Celcius 0oC / 100oC Kelvin a. CH4N empirical b. NaO empirical c. C6H3O3 molecular d. H2O2 molecular e. Na2SO3 empirical f. C6H10O4 molecular 85. If given the empirical formula and molar mass for a compound, calculate the compound's molecular formula? a. CH2O , mass = 90 g/mol C3H6O3 b. C3H5O2 mass = 146 g/mol 273 / 373 5 C6H10O4 Chemistry: First Semester Exam Prep #4 (Choose the BEST answer.) 86. Find the missing density, mass, or volume of the following: a. The mass of a substance is 45.6 g and the volume is 15 cm3: Density = ________ b. The volume of a substance is 2.9 ml its density is 6 g/ml: Mass = ________ 95. Calculate the % composition of each element by mass of: Propane C3H8 _______________ / _______________ Water ______________ / _______________ 96. Which of the following are isotopes? a. 2412X b. 2713X c. 2813X d. 2411X 97. Describe the major parts of Dalton's Atomic Theory: c. The density of a substance is 7.8 g/cm3 and the mass is 125 g: Volume = _______ (Hint: D/1 = M/V (Given any two of the numbers; D, M or V, you can cross multiply and divide to find what’s missing) 87. If you have 6.7 L of O2 at STP, how many moles do you have __________________ 88. What is the molar mass of Sn3(PO4)2? _________ 98. Balance the following chemical equations: 3 CO + _____ Fe2O3 2 Fe + 3 CO2 3 Zn(OH)2 + 2 H3PO4 ____ Zn3(PO4)2 + 6 H2O 89. How many moles are in 137.5 g of Mn? ________ 90. What is the mass of 2 moles of C2H6? _________ _____ H3PO4 + 3 KOH _____ K3PO4 + 3 H2O 99. Name and describe the 6 types of chemical reactions. 91. What are the correct formulas or names for the following compounds? a. potassium sulfate K2SO4 b. calcium phosphate Ca3(PO4)2 c. disulfur heptoxide S2O7 d. trinitrogen pentahydride N3H5 92. List the diatomic molecules: a. _______ b. _______ c. _______ d. _______ 100. Describe difference between homogenous, heterogeneous mixtures, and pure substance: e. _______ f. _______ g. _______ 93. List the names & formulas of the 5 common acids: a. ____________________- ______________ b. ____________________- ______________ c. ____________________- ______________ d. ____________________- ______________ e. ____________________- ______________ 94. What is the oxidation (nuclear) charge of each substance (ion) given? Al __________ Ca __________ O __________ S __________ K __________ N __________ 6 Chemistry: First Semester Exam Prep #4 (Choose the BEST answer.) 7