* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
Download Practice Multiple Choice Questions for the Chemistry Final Exam
Periodic table wikipedia , lookup
IUPAC nomenclature of inorganic chemistry 2005 wikipedia , lookup
Marcus theory wikipedia , lookup
Debye–Hückel equation wikipedia , lookup
Electrical resistivity and conductivity wikipedia , lookup
Elementary particle wikipedia , lookup
Physical organic chemistry wikipedia , lookup
Inductively coupled plasma mass spectrometry wikipedia , lookup
Metastable inner-shell molecular state wikipedia , lookup
Biochemistry wikipedia , lookup
Electrochemistry wikipedia , lookup
Size-exclusion chromatography wikipedia , lookup
X-ray fluorescence wikipedia , lookup
Artificial photosynthesis wikipedia , lookup
Resonance (chemistry) wikipedia , lookup
History of chemistry wikipedia , lookup
Water splitting wikipedia , lookup
X-ray photoelectron spectroscopy wikipedia , lookup
Light-dependent reactions wikipedia , lookup
Bioorthogonal chemistry wikipedia , lookup
Atomic orbital wikipedia , lookup
Metallic bonding wikipedia , lookup
Gas chromatography–mass spectrometry wikipedia , lookup
Metalloprotein wikipedia , lookup
Hypervalent molecule wikipedia , lookup
Chemistry: A Volatile History wikipedia , lookup
Gaseous detection device wikipedia , lookup
Molecular orbital diagram wikipedia , lookup
Extended periodic table wikipedia , lookup
Chemical bond wikipedia , lookup
Atomic nucleus wikipedia , lookup
Rutherford backscattering spectrometry wikipedia , lookup
Stoichiometry wikipedia , lookup
Hydrogen atom wikipedia , lookup
History of molecular theory wikipedia , lookup
Electrolysis of water wikipedia , lookup
Electron configuration wikipedia , lookup
Practice Multiple Choice Questions for the Chemistry Final Exam 2012 1. One chemical property of matter is a) boiling point. c) reactivity. b) texture. d) density. 2. The state of matter in which a material has definite shape and definite volume is the a) liquid state. b) solid state. c) gaseous state. d) vaporous state. 3. Under ordinary conditions of temperature and pressure, the particles in a gas are a) closely packed. b) very far from each other. c) held in fixed positions. d) able to slide past each other. 4. A horizontal row of elements in the periodic table is called a(n) a) group. b) period. c) family. d) octet. 5. Elements in a group in the periodic table can be expected to have similar a) atomic masses. b) atomic numbers. c) numbers of neutrons. d) properties. 6. The symbol that represents the measured unit for volume is a) mL. b) mg. c) mm. d) cm. 7. A measure of the quantity of matter is a) density. c) volume. 8. 9. 10. b) weight. d) mass. To determine density, the quantities that must be measured are a) mass and weight. b) volume and weight. c) volume and concentration d) volume and mass. The density of aluminum is 2.70 g/m3. The volume of a solid piece of aluminum is 1.50 cm3. Find its mass. a) 1.50 g b) 1.80 g c) 2.70 g d) 4.05 g 1.06 L of water is equivalent to a) 0.00106 mL. c) 106 mL. 11. Convert -25'C to the kelvin scale. a) -323. K c) 248. K b) 10.6 mL. d) 1060 mL. b) -248. K d) 323. K 12. In oxides of nitrogen, such as N2O, NO, NO2, and N2O3, atoms combine in small whole-number ratios. This evidence supports the law of a) conservation of mass. b) multiple proportion. c) definite composition. d) mass action. 1 13. Who was the schoolmaster who studied chemistry and proposed an atomic theory? a) John Dalton b) Jons Berzehus c) Robert Brown d) Dmitri Mendeleev 14. According to Dalton’s atomic theory, atoms a) of different elements b) can be divided into combine in simple protons, neutrons, and whole-number ratios to form electrons. compounds. c) of all elements are identical d) can be destroyed in in size and mass. chemical reactions. 15. The discovery of the electron resulted from experiments using a) gold foil. b) cathode rays. c) neutrons. d) alpha particles. 16. Who discovered the nucleus by bombarding gold foil with positively charged particles and noting that some particles were widely deflected? a) Rutherford b) Dalton c) Chadwick d) Bohr 17. Rutherford fired positively charged particles at metal foil and concluded that most of the mass of an atom was a) in the electrons. b) concentrated in the nucleus. c) evenly spread throughout d) in rings around the atom. the atom 18. A nuclear particle that has about the same mass as a proton, but with no electrical charge, is called a(n) a) nuclide. b) neutron. c) electron. d) isotope. 19. An atom is electrically neutral because a) neutrons balance the b) nuclear forces stabilize the protons and electrons. charges. c) the numbers of protons and d) the numbers of protons and electrons are equal neutrons are equal. 20. Atoms of the same element that have different masses are called a) moles. b) isotopes. c) nuclides. d) neutrons. 21. All atoms of the same element have the same a) atomic mass. b) number of neutrons. c) mass number. d) atomic number. 22. An aluminum isotope consists of 13 protons, 13 electrons, and 14 neutrons. Its mass number is a) 13. b) 14. c) 27. d) 40. 2 23. Carbon-14 (atomic number 6), the radioactive nuclide used in dating fossils, has a) 6 neutrons. b) 8 neutrons. c) 10 neutrons. d) 14 neutrons. 24.The number of atoms in 1 mol of carbon is a) 6.02 x 1022. b) 6.02 x 1023. 22 c) 5.02 x 10 . d) 5.02 x 1023. 25. Molar mass a) is the mass in gram of one b) is numerically equal to the mole of a substance. average atomic mass of the element. c) both a and b d) neither a nor b 26. A sample of tin (atomic mass 118.69 amu) contains 3.01 x 1023 atoms. The mass of the sample is a) 3.01 b) 59.3 g. c) 72.6 g. d) 11 g. 27. A bright line spectrum of an atom is caused by the energy released when electrons a) jump to a higher energy b) fall to a lower energy level. level. c) absorb energy and jump to d) absorb energy and fall to a a higher energy level. lower energy level. 28. For an electron in an atom to change from the ground state to an excited state, a) energy must be released. b) energy must be absorbed. c) radiation must be emitted d) the electron must make a transition from a higher to a lower energy level. 29. The set of orbitals that are dumbbell-shaped and directed along the x, y, and z axes are called a) d orbitals. b) p orbitals. c) f orbitals. d) s orbitals. 30. The letter designations for the first four sublevels with the number of electrons that can be accommodated in each sublevel are a) s:1, p:3, d:10, and f:14. b) s:1, p:3, d.5, and f 7. c) s:2, p:6, d.10, and f 14. d) s:1, p:2, d: 3, and f 4. 31. Which of the following rules requires that each of the p orbitals at a particular energy level receive one electron before any of them can have two electrons? a) Hund’s rule b) the Pauli exclusion principle c) the Aufbau principle d) the quantum rule 32. What is the electron configuration for nitrogen, atomic number 7? a) 1s2 2s2 2p3 b) 1s2 2s3 2p2 2 3 l c) 1s 2s 2p d) 1s2 2s2 2p2 3s1 33. Mendeleev predicted that the spaces in his periodic table represented a) isotopes. b) radioactive elements. c) permanent gaps. d) undiscovered elements. 3 34. The discovery of the noble gases changed Mendeleev's periodic table by adding a new a) period. b) series. c) group . d) sublevel block. 35. In the modern periodic table, elements are ordered according to a) decreasing atomic mass. b) Mendeleev's original design. c) increasing atomic number. d) the date of their discovery. 36. Krypton, atomic number 36, is the fourth element in Group 18. What is the atomic number of xenon, the fifth element in Group 18? a) 54 b) 68 c) 72 d) 90 37. The electron configuration of aluminum, atomic number 13, is [Ne] 3s2 3pl. Aluminum is in period a) 2. b) 3. c) 6. d) 13. 38. Calcium, atomic number 20, has the electron configuration [Ar] 4s2. In what period is calcium? a) Period 2 b) Period 4 c) Period 8 d) Period 20 39. The energy required to remove an electron from an atom is the atom’s a) electron affinity. b) electron energy. c) electronegativity. d) ionization energy. 40. A measure of the ability of an atom in a chemical compound to attract electrons is called a) electron affinity. b) electron configuration c) electronegativity. d) ionization potential. 41. Across a period in the periodic table, atomic radii a) gradually decrease. b) gradually decrease, then sharply increase. c) gradually increase. d) gradually increase, then sharply decrease. 42. The number of valence electrons in Group 17 elements is a) 7. b) 8. c) 17. d) equal to the period number. 43. The electrons involved in the formation of a chemical bond are called a) dipoles. b) s electrons. c) Lewis electrons. d) valence electrons. 44. The chemical bond formed when two atoms share electrons is called a(n) a) ionic bond. b) orbital bond. c) Lewis structure. d) covalent bond. 45. If the atoms that share electrons have an unequa1 attraction for the electrons, the bond is called a) nonpolar. b) polar. c) ionic. d) dipolar. 4 46. The electron configuration of nitrogen is 1s2 2s2 2p3. How many more electrons does nitrogen need to satisfy the octet rule? a) 1 b) 3 c) 5 d) 8 47. A formula that shows the types and numbers of atoms combined in a single molecule is a(n) a) molecular formula. b) ionic formula. c) Lewis structure. d) covalent formula. 48. VSEPR theory is a model for predicting a) the strength of metallic b) the shape of molecules. bonds. c) lattice energy values. d) ionization energy. 49. The following molecules contain polar bonds. The only polar molecule is a) CCl4 b) CO2 c) NH3 d) CH4 50. How many atoms of fluorine are present in a molecule of carbon tetrafluoride, CF4? a) 1 b) 2 c) 4 d) 5 51. The formula for carbon dioxide, CO2, can represent a) one molecule of carbon b) 1 mol of carbon dioxide dioxide. molecules. c) one molar mass of carbon dioxide d) all of the above. 52. What is the formula for aluminum sulfate? a) AlSO4 b) Al2SO4 c) Al2(SO4)3 d) Al(SO4)3 53. Name the compound K2SO4. a) potassium sulfate c) potassium sulfide b) potassium sulfur tetroxide d) dipotassium sulfate 54. What is the metallic ion in copper(II) chloride? a) Co2+ b) O22+ c) Cu d) Cl55. Name the compound N2O4. a) sodium tetroxide c) nitrous oxide b) dinitrogen tetroxide d) binitrogen oxide 56. What is the formula for sulfur dichloride? a) NaCl2 b) SCl2 c) S2Cl d) S2Cl2 57. The molar mass of MgI2 is a) the sum of the masses of 1 mol of Mg and 2 mol of I. c) the sum of the masses of 2 mol of Mg and 2 mol of I. b) the sum of the masses of 1 mol of Mg and I mol of I. d) impossible to calculate. 5 58. The molar mass of NO2 is 46.01 g/mol. How many moles of NO2 are present in 114.95 g? a) 0.4003 mol b) 1.000 mol c) 2.498 mol d) 114.95 mol 59. What is the percentage composition of CF4? a) 20% C, 80% F b) 13.6% C, 86.4% F c) 16.8% C, 83.2% F d) 81% C, 19% F 60. What is the empirical formula for a compound that is 43.6% phosphorus and 56.4% oxygen? a) P3O7 b) PO3 c) P2O3 d) P2O5 61. The molecular formula for vitamin C is C6H8O6. What is the empirical formula? a) CHO b) CH2O c) C3H4O3 d) C2H4O2 62. A compound’s empirical formula is HO. If the formula mass is 34 amu, what is the molecular formula? a) H2O b) H2O2 c) HO3 d) H2O3 63. In writing an equation that produces hydrogen gas, the correct representation of hydrogen gas is a) H. b) 2H. c) H2. d) OH. 64. Which equation is NOT balanced? a) 2H2 + O2 2H2O c) 2Mg + O2 2MgO b) Al + 3HCl AlCl3 + H2 d) CaCl2 + 2NaBr CaBr2 + 2NaC 65. The reaction Mg(s) + 2HCI(aq) H2(g) + MgCl2(aq) is a a) composition reaction b) decomposition reaction. c) single-replacement reaction. d) double-replacement reaction. 66. The reaction Pb(NO3)2(aq) + 2KI(aq) PbI2(S) + 2KNO3(aq) is a a) double-replacement reaction. b) synthesis reaction. c) decomposition reaction. d) combustion reaction. 67. What is the balanced equation when aluminum reacts with copper(II) sulfate? a) Al + Cu2S A12S + Cu b) 2Al + 3CuSO4 Al2(SO4)3 + 3Cu c) Al + CuSO4 AlSO4 + Cu d) 2Al + Cu2SO4 -4 A12SO4 + 2Cu 68. In the reaction N2 + 3H22NH3, what is the mole ratio of nitrogen to ammonia? a) 1:1 b) 1:2 c) 1:3 d) 2:3 69. In the equation 2KClO3 2KCI + 3O2, how many moles of oxygen are produced when 3.0 mol of KClO3 decompose completely? a) 1.0 mol b) 2.5 mol c) 3.0 mol d) 4.5 mol 6 70. For the reaction 2HNO3 + Mg(OH)2 Mg(NO3)2 + 2H2O, how many grams of magnesium nitrate are produced from 8.00 mol of nitric acid, HNO3? a) 148 g b) 445 g c) 593 g d) 818 g 71. For the reaction 3Fe + 4H2O Fe3O4 + 4H2, how many moles of iron oxide are produced from 500 g of iron? a) 1 mol b) 3 mol c) 9 mol d) 12 mol 72. For the reaction SO3 + H2O H2SO4, calculate the percent yield if 500. g of sulfur trioxide react with excess water to produce 575 g of sulfuric acid. a) 82.7% b) 88.3% c) 91.2% d) 93.9% 73. An ideal gas is an imaginary gas a) not made of particles. b) that conforms to all of the assumptions of the kinetic theory. c) whose particles have zero d) made of motionless mass. particles. 74. Which is an example of gas diffusion? a) inflating a flat tire c) a cylinder of oxygen stored b) the perfume spreading through a room d) gas escaping from a hole in a balloon 75. According to the kinetic-molecular theory, how does a gas expand? a) Its particles become larger. b) Collisions between particles become elastic. c) Its temperature rises. d) Its particles move apart in straight lines 76. Which is an example of effusion? a) air slowly escaping from a pinhole in a tire c) helium dispersing into a room after a balloon pops b) the aroma of a cooling pie spreading across a room d) oxygen and gasoline fumes mixing in an automobile carburetor 77. What does the constant bombardment of gas molecules against the inside walls of a container produce? a) temperature b) density c) pressure d) diffusion 78. Convert the pressure 0.75 atm to mm Hg. a) 101.325 mm Hg b) 430 mm Hg c) 570 mm Hg d) 760 mm Hg 79. Standard temperature is exactly a) 100 C. c) 0 C. b) 273 C. d) 0 K. 80. Standard pressure is exactly a) 1 atm c) 101.325 atm b) 760 atm. d) 101 atm. 7 81. Pressure and volume changes at a constant temperature can be calculated using a) Boyle's law. b) Charles's law. c) Kelvin’s law. d) Dalton's law. 83. The volume of a gas is 5.0 L when the temperature is 5.0 C. If the temperature is increased to 10.0 C without changing the pressure, what is the new volume? a) 2.5 L b) 4.8 L c) 5.1 L d) 10.0 L 84. A 150.0 L sample of gas is collected at 1.20 atm and 25.0 C. What volume does the gas have at 1.50 atm and 20.0 C? a) 94 L b) 118 L c) 143 L d) 183 L 85. In the equation H2(g) + Cl2(g) 2HCI(g), one volume of hydrogen yields how many volumes of hydrogen chloride? a) 1 b) 2 c) 3 d) 4 86. At constant temperature and pressure, gas volume is directly proportional to the a) molar mass of the gas. b) number of moles of gas. c) density of the gas at STP. d) rate of diffusion. 87. According to Avogadro's law, 1 L of H2(g) and 1 L of O2(g) at the same temperature and pressure a) have the same mass. b) have unequal volumes. c) contain 1 mol of gas each. d) contain equal numbers of molecules. 88. The standard molar volume of a gas at STP is a) 22.4 L. b) g/22.4 L. c) g-mol wt/22.4 L. d) 1 L. 89. A 1.00 L sample of a gas has a mass of 1.25 g at STP. What is the mass of 1 mol of this gas? a) a little less then 1.0 g b) 1.25 g c) 22.4 g d) 28.0 g 90. According to Graham’s law, two gases at the same temperature and pressure will have different rates of diffusion because they have different a) volumes. b) molar masses. c) kinetic energies. d) condensation points. 91. Why are water molecules polar? a) They contain two kinds of b) The electrons in the covalent atoms. bonds spend more time closer to the oxygen nucleus. c) The hydrogen bonds are d) They have covalent bonds. weak. 92. What is the freezing point of water at standard pressure? a) -10 C b) 0 C c) 4'C d) 32 C 8 93. What is the boiling point of water at standard pressure? a) 100 C b) 112 C c) 212 C d) 200 C 94. Which of the following is a pure substance? a) water b) milk c) soil d) concrete 95. Sugar in water is an example of which solute-solvent combination? a) gas-liquid b) liquid-liquid c) solid-liquid d) liquid-solid 96. Solutions that conduct electricity are called a) ions b) super solutions c) electrolytes d) nonelectrolytes 97. If the amount of solute present in a solution at a given temperature is less than the maximum amount that can dissolve at that temperature, the solution is said to be a) saturated. b) unsaturated. c) supersaturated. d) concentrated. 98. Which of the following is likely to produce crystals if disturbed? a) an unsaturated solution b) a supersaturated solution c) a saturated solution d) a concentrated solution 99. What is the molarity of a solution that contains 0.202 mol KC1 in 7.98 L solution? a) 0.0132 M b) 0.0253 M c) 0.459 M d) 1.363 M 100. How many moles of HCI are present in 0.70 L of a 0.33 M HCI solution? a) 0.23 mol b) 0.28 mol c) 0.38 mol d) 0.47 mol 101. An NaOH solution contains 1.90 mol of NaOH, and its concentration is 0.555 M. What is its volume? a) 0.623 L b) 0.911 L c) 1.05 L d) 3.42 L 102. How many milliliters water are needed to make a 0.171 M solution that contains 1.00 g of NaCl? a) 100 mL b) 1000 mL c) 171 mL d) 17.1 mL 103. How many moles of ions are produced by the dissociation of 1 mol of MgCl2? a) 0 mol b) 1 mol c) 2 mol d) 3 mol 104. Colligative properties depend on a) the crystal size of the solute c) the physical properties of the solute particles. b) the concentration of solute d) the boiling point and freezing point of the solution. 9 105. Compared with a 0.01 m sugar solution, a 0.01 m KCl solution has a) the same freezing-point b) about twice the depression. freezing-point depression. c) the same freezing-point d) about six times the elevation. freezing-point elevation. 106. Electrolytes affect colligative properties differently than do nonelectrolytes because electrolytes a) are volatile. b) have lower boiling points. c) produce fewer moles of d) produce more moles of solute particles per mole of solute particles per mole of solute. solute. 107. Acids taste a) sweet. c) bitter. b) sour. d) salty. 108. Acids react with a) bases to produce salts and water. c) water to produce bases and salts. b) salts to produce bases and water. d) neither bases, salts, nor water. 109. Bases taste a) soapy. c) sweet. b) sour. d) bitter. 110. A binary acid contains a) two hydrogen atoms. c) hydrogen and two other elements. b) hydrogen and one other element. d) hydrogen and three other elements. 111. According to the traditional Arrhenius definition, an acid contains a) hydrogen and does not b) hydrogen and ionizes to ionize. form hydrogen ions. c) oxygen and ionizes to form d) oxygen and ionizes to form hydroxide ions. oxygen ions. 112. A substance that ionizes nearly completely in aqueous solutions and produces H3O+ is a a) weak base. b) strong base. c) weak acid. d) strong acid. 113. A Bronsted-Lowry acid is a) an electron-pair acceptor. c) a proton acceptor. b) an electron-pair donor. d) a proton donor. 10 114. What is neutralization? a) an acid-base reaction that does not include dissocation c) a reaction of oxygen and hydrogen ions to form water molecules 115. Pure water contains a) water molecules only. c) hydroxide ions only. b) a reaction of hydronium ions and hydroxide ions to form a salt d) a reaction of an acid and a base to form water molecules and a salt b) hydronium ions only. d) water molecules, hydronium ions, and hydroxide ions. 116. What is the concentration of H3O+ in pure water? a) 10-7 M b) 0.7 M c) 55.4 M d) 107 M 117. Which expression represents the pH of a solution? a) log[H3O+1] b) -log[H3O+1] c) log[OH-] d) -log[OH-] 118. If [H3O+1] of a solution is less than [OH-1], the solution a) is always acidic. b) is always basic. c) is always neutral. d) might be acidic, basic, or neutral. 119. What is the pH of a neutral solution at 25 C? a) 0 b) 1 c) 7 d) 14 120. The pH scale in general use ranges from a) 0 to 1. b) - 1 to 1. c) 0 to 7. d) 0 to 14. 121. The pH of an acidic solution is a) less than 0. c) greater than 7. b) less than 7. d) greater than 14. 122. The pH of a basic solution is a) less than 0. c) greater than 7. b) less than 7. d) greater than 14. 123. If [H3O+] = 1.7 x 10-3 M, what is the pH of the solution? a) 1.81 b) 2.13 c) 2.42 d) 2.77 124. What is the pH of a 0.027 M KOH solution? a) 6.47 b) 12.43 c) 12.92 d) 14.11 125. What is the hydronium ion concentration of a solution whose pH is 4.12? a) 4.4 x 10-8 M b) 5.1 X 10-6 M c) 6.4 x 10-5 M d) 7.6 x 10-5 M 11 126. Dyes with pH-sensitive colors are used as a) primary standards. b) indicators. c) titrants. d) None of the above 127.In an acid-base titration, equivalent quantities of hydronium ions and hydroxide ions are present a) at the beginning point. b) at the midpoint. c) at the endpoint. d) throughout the titration. 128. What is the molarity of an HCL solution if 125 mL is neutralized in a titration by 76.0 mL of 1.22 M KOH? a) 0.371 M b) 0.455 M c) 0.617 M d) 0.742 M 129. What is the molarity of a Ba(OH)2 solution if 93.9 mL is completely titrated by 15.3 mL of 0.247 M H2SO4? a) 0.0101 M b) 0.0201 M c) 0.0402 M d) 0.0805 M 130. How is a Celsius temperature reading converted to a Kelvin temperature reading? a) by adding 273 b) by subtracting 273 c) by dividing by 273 d) by multiplying by 273 131. The pH of a solution is 9. What is its H3O+ concentration? a) 10-9 M b) 10-7 M -5 c) 10 M d) 9 M 132. After balancing the equation FeCl3 + Zn ZnCl2 + Fe, the coefficients, in order from left to right, are a) 2, 2, 1, 2. b) 1, 1, 1, 1. c) 4, 3, 3, 4. d) 2, 3, 3, 2. 133. Which of the following is the electron configuration of carbon in the ground state? a) 1s2 2s2 2p2 b) 2s2 2sl 2p3 2 2 c) 1s 2s 2p3 d) 1s2 2s2 2p6 134. How many valence electrons does a carbon atom have? a) 3 b) 4 c) 5 d) 6 135. To draw the Lewis dot structure for Nitrogen, you would need to draw ___dots. a) 7 b) 5 c) 3 d) 2 136. Which of the following would not speed up the rate that a reaction occurs? a) catalyst b) heat it up c) use larger pieces d) increase concentration 137. The shape of an ammonia (NH3) molecule is a) trigonal pyramidal b) bent c) trigonal planar d) tetrahedral 12