Final Review Part Two

Multiple ChoiceIdentify the choice that best completes the statement or answers the question.

____ 1. Which hypothesis led to the discovery of the proton?a. When a neutral hydrogen atom loses an electron, a positively-charged particle should

remain.b. A proton should be 1840 times heavier than an electron.c. Cathode rays should be attracted to a positively-charged plate.d. The nucleus of an atom should contain neutrons.

____ 2. Which of the following is correct concerning subatomic particles?a. The electron was discovered by Goldstein in 1886.b. The neutron was discovered by Chadwick in 1932.c. The proton was discovered by Thomson in 1880.d. Cathode rays were found to be made of protons.

____ 3. All atoms are ____.a. positively charged, with the number of protons exceeding the number of electronsb. negatively charged, with the number of electrons exceeding the number of protonsc. neutral, with the number of protons equaling the number of electronsd. neutral, with the number of protons equaling the number of electrons, which is equal to the

number of neutrons

____ 4. Which of the following sets of symbols represents isotopes of the same element?a.

J J Jc.

M M Mb.

L L Ld.

Q Q Q

____ 5. How do the isotopes hydrogen-1 and hydrogen-2 differ?a. Hydrogen-2 has one more electron than hydrogen-1.b. Hydrogen-2 has one neutron; hydrogen-1 has none.c. Hydrogen-2 has two protons; hydrogen-1 has one.d. Hydrogen-2 has one proton; hydrogen-1 has none.

____ 6. Which of the following isotopes has the same number of neutrons as phosphorus-31?a.

Pc.

Sib.

Sd.

Si

____ 7. What is the maximum number of electrons in the second principal energy level?a. 2 c. 18b. 8 d. 32

____ 8. When an electron moves from a lower to a higher energy level, the electron ____.a. always doubles its energyb. absorbs a continuously variable amount of energyc. absorbs a quantum of energyd. moves closer to the nucleus

____ 9. If three electrons are available to fill three empty 2p atomic orbitals, how will the electrons be distributed in the three orbitals?a. one electron in each orbitalb. two electrons in one orbital, one in another, none in the thirdc. three in one orbital, none in the other twod. Three electrons cannot fill three empty 2p atomic orbitals.

____ 10. How many unpaired electrons are in a sulfur atom (atomic number 16)?a. 0 c. 2b. 1 d. 3

____ 11. How many half-filled orbitals are in a bromine atom?a. 1 c. 3b. 2 d. 4

____ 12. Which of the following electron configurations of outer sublevels is the most stable?a. 4d 5s c. 4d 5sb. 4d 5s d. 4d 5s

____ 13. What is the wavelength of an electromagnetic wave that travels at 3 10 m/s and has a frequency of 60 MHz? (1 MHz = 1,000,000 Hz)a.

b. 60 MHz 300,000,000 m/sc.

d. No answer can be determined from the information given.

____ 14. Which variable is directly proportional to frequency?a. wavelength c. positionb. velocity d. energy

____ 15. How do the energy differences between the higher energy levels of an atom compare with the energy differences between the lower energy levels of the atom?a. They are greater in magnitude than those between lower energy levels.b. They are smaller in magnitude than those between lower energy levels.c. There is no significant difference in the magnitudes of these differences.d. No answer can be determined from the information given.

____ 16. In an s orbital, the probability of finding an electron a particular distance from the nucleus does NOT depend on ____.a. a quantum mechanical model c. the Schrodinger equationb. direction with respect to the nucleus d. the electron energy sublevel

____ 17. To what category of elements does an element belong if it is a poor conductor of electricity?a. transition elements c. nonmetalsb. metalloids d. metals

____ 18. Of the elements Fe, Hg, U, and Te, which is a representative element?a. Fe c. Ub. Hg d. Te

____ 19. Which of the following factors contributes to the increase in atomic size within a group in the periodic table as the atomic number increases?a. more shielding of the electrons by the highest occupied energy levelb. an increase in size of the nucleusc. an increase in number of protonsd. fewer electrons in the highest occupied energy level

____ 20. Which of the following elements has the smallest atomic radius?a. sulfur c. seleniumb. chlorine d. bromine

____ 21. In which of the following sets are the charges given correctly for all the ions?a. Na , Mg , Al c. Rb , Ba , Pb. K , Sr , O d. N , O , F

____ 22. In which of the following groups of ions are the charges all shown correctly?a. Li , O , S c. K , F , Mgb. Ca , Al , Br d. Na , I , Rb

____ 23. Which of the following factors contributes to the increase in ionization energy from left to right across a period?a. an increase in the shielding effectb. an increase in the size of the nucleusc. an increase in the number of protonsd. fewer electrons in the highest occupied energy level

____ 24. As you move from left to right across the second period of the periodic table ____.a. ionization energy increases c. electronegativity decreasesb. atomic radii increase d. atomic mass decreases

____ 25. Of the following elements, which one has the smallest first ionization energy?a. boron c. aluminumb. carbon d. silicon

____ 26. Which of the following pairs of elements is most likely to form an ionic compound?a. magnesium and fluorineb. nitrogen and sulfurc. oxygen and chlorined. sodium and aluminum

____ 27. Which elements can form diatomic molecules joined by a single covalent bond?a. hydrogen onlyb. halogens onlyc. halogens and members of the oxygen group onlyd. hydrogen and the halogens only

____ 28. Which of the following electron configurations gives the correct arrangement of the four valence electrons of the carbon atom in the molecule methane (CH )?a. 2s 2p c. 2s 2p 3sb. 2s 2p 3s d. 2s 2p

____ 29. Which of the following covalent bonds is the most polar?a. H—F c. H—Hb. H—C d. H—N

____ 30. When placed between oppositely charged metal plates, the region of a water molecule attracted to the negative plate is the ____.a. hydrogen region of the molecule c. H—O—H plane of the moleculeb. geometric center of the molecule d. oxygen region of the molecule

____ 31. Which polyatomic ion forms a neutral compound when combined with a group 1A monatomic ion in a 1:1 ratio?a. ammonium c. nitrateb. carbonate d. phosphate

____ 32. Consider a mystery compound having the formula M T . If the compound is not an acid, if it contains only two elements, and if M is not a metal, which of the following is true about the compound?a. It contains a polyatomic ion. c. Its name ends in -ic.b. Its name ends in -ite or -ate. d. It is a binary molecular compound.

____ 33. What is the formula for hydrosulfuric acid?a. H S c. HSOb. H SO d. H S

____ 34. What is the correct formula for barium chlorate?a. Ba(ClO) c. Ba(ClO )b. Ba(ClO ) d. BaCl

____ 35. What is the correct formula for calcium dihydrogen phosphate?a. CaH PO c. Ca(H PO )b. Ca H PO d. Ca(H HPO )

____ 36. What is the correct name for Sn (PO ) ?a. tritin diphosphate c. tin(III) phosphateb. tin(II) phosphate d. tin(IV) phosphate

____ 37. Butanol is composed of carbon, hydrogen, and oxygen. If 1.0 mol of butanol contains 6.0 10 atoms of hydrogen, what is the subscript for the hydrogen atom in C H O?a. 1 c. 6b. 10 d. 8

____ 38. Which of the following gas samples would have the largest number of representative particles at STP?a. 12.0 L He c. 0.10 L Xeb. 7.0 L O d. 0.007 L SO

____ 39. Given 1.00 mole of each of the following gases at STP, which gas would have the greatest volume?a. He c. SOb. O d. All would have the same volume.

____ 40. If the density of a noble gas is 1.783 g/L at STP, that gas is ____.

a. Kr c. Arb. Xe d. He

____ 41. Which of the following compounds has the lowest percent gold content by weight?a. AuOH c. AuClb. Au(OH) d. AuI

____ 42. Which of the following compounds has the highest oxygen content, by weight?a. Na O c. BaOb. CO d. H O

____ 43. The ratio of carbon atoms to hydrogen atoms to oxygen atoms in a molecule of dicyclohexyl maleate is 4 to 6 to 1. What is its molecular formula if its molar mass is 280 g?a. C H O c. C H Ob. C H O d. C H O

____ 44. Who was the man who lived from 460B.C.–370B.C. and was among the first to suggest the idea of atoms?a. Atomos c. Democritusb. Dalton d. Thomson

____ 45. Which of the following was NOT among Democritus’s ideas?a. Matter consists of tiny particles called atoms.b. Atoms are indivisible.c. Atoms retain their identity in a chemical reaction.d. Atoms are indestructible.

____ 46. Dalton's atomic theory included which idea?a. All atoms of all elements are the same size.b. Atoms of different elements always combine in one-to-one ratios.c. Atoms of the same element are always identical.d. Individual atoms can be seen with a microscope.

____ 47. Which of the following is NOT a part of Dalton's atomic theory?a. All elements are composed of atoms.b. Atoms are always in motion.c. Atoms of the same element are identical.d. Atoms that combine do so in simple whole-number ratios.

____ 48. Which of the following was originally a tenet of Dalton's atomic theory, but had to be revised about a century ago?a. Atoms are tiny indivisible particles.b. Atoms of the same element are identical.c. Compounds are made by combining atoms.d. Atoms of different elements can combine with one another in simple whole number ratios.

____ 49. The comparison of the number of atoms in a copper coin the size of a penny with the number of people on Earth is made to illustrate which of the following?a. that atoms are indivisibleb. that atoms are very smallc. that atoms are very larged. that in a copper penny, there is one atom for every person on Earth

____ 50. The range in size of most atomic radii is approximately ____.a. 2 to 5 cm c. 5 10 m to 2 10 mb. 2 to 5 nm d. 5 10 m to 2 10 m

____ 51. Why did J. J. Thomson reason that electrons must be a part of the atoms of all elements?a. Cathode rays are negatively-charged particles.b. Cathode rays can be deflected by magnets.c. An electron is 2000 times lighter than a hydrogen atom.d. Charge-to-mass ratio of electrons was the same, regardless of the gas used.

____ 52. Which of the following is true about subatomic particles?a. Electrons are negatively charged and are the heaviest subatomic particle.b. Protons are positively charged and the lightest subatomic particle.c. Neutrons have no charge and are the lightest subatomic particle.d. The mass of a neutron nearly equals the mass of a proton.

____ 53. Who conducted experiments to determine the quantity of charge carried by an electron?a. Rutherford c. Daltonb. Millikan d. Thomson

____ 54. What is the relative mass of an electron?a. 1/1840 the mass of a hydrogen atom c. 1/1840 the mass of a C-12 atomb. 1/1840 the mass of a neutron + proton d. 1/1840 the mass of an alpha particle

____ 55. The particles that are found in the nucleus of an atom are ____.a. neutrons and electrons c. protons and neutronsb. electrons only d. protons and electrons

____ 56. As a consequence of the discovery of the nucleus by Rutherford, which model of the atom is thought to be true?a. Protons, electrons, and neutrons are evenly distributed throughout the volume of the atom.b. The nucleus is made of protons, electrons, and neutrons.c. Electrons are distributed around the nucleus and occupy almost all the volume of the atom.d. The nucleus is made of electrons and protons.

____ 57. The nucleus of an atom is ____.a. the central core and is composed of protons and neutronsb. positively charged and has more protons than neutronsc. negatively charged and has a high densityd. negatively charged and has a low density

____ 58. In which of the following sets is the symbol of the element, the number of protons, and the number of electrons given correctly?a. In, 49 protons, 49 electrons c. Cs, 55 protons, 132.9 electronsb. Zn, 30 protons, 60 electrons d. F, 19 protons, 19 electrons

____ 59. The mass number of an element is equal to ____.a. the total number of electrons in the nucleusb. the total number of protons and neutrons in the nucleusc. less than twice the atomic numberd. a constant number for the lighter elements

____ 60. Using the periodic table, determine the number of neutrons in O.a. 4 c. 16b. 8 d. 24

____ 61. How many protons, electrons, and neutrons does an atom with atomic number 50 and mass number 125 contain?a. 50 protons, 50 electrons, 75 neutrons c. 120 neutrons, 50 protons, 75 electronsb. 75 electrons, 50 protons, 50 neutrons d. 70 neutrons, 75 protons, 50 electrons

____ 62. Which of the following statements is NOT true?a. Atoms of the same element can have different masses.b. Atoms of isotopes of an element have different numbers of protons.c. The nucleus of an atom has a positive charge.d. Atoms are mostly empty space.

____ 63. If E is the symbol for an element, which two of the following symbols represent isotopes of the same element?

1. E 2. E 3. E 4. E

a. 1 and 2 c. 1 and 4b. 3 and 4 d. 2 and 3

____ 64. Select the correct symbol for an atom of tritium.a.

nc.

Hb.

Hd.

H

____ 65. How is the number of neutrons in the nucleus of an atom calculated?a. Add the number of electrons and protons together.b. Subtract the number of electrons from the number of protons.c. Subtract the number of protons from the mass number.d. Add the mass number to the number of electrons.

____ 66. In which of the following is the number of neutrons correctly represented?a.

F has 0 neutrons.c.

Mg has 24 neutrons.b.

As has 108 neutrons.d.

U has 146 neutrons.

____ 67. Which of the following statements is NOT true?a. Protons have a positive charge.b. Electrons are negatively charged and have a mass of 1 amu.c. The nucleus of an atom is positively charged.d. Neutrons are located in the nucleus of an atom.

____ 68. Why do chemists use relative masses of atoms compared to a reference isotope rather than the actual masses of the atoms?a. The actual mass of an electron is very large compared to the actual mass of a proton.b. The actual masses of atoms are very small and difficult to work with.c. The number of subatomic particles in atoms of different elements varies.d. The actual masses of protons, electrons, and neutrons are not known.

____ 69. The atomic mass of an element is the ____.

a. total number of subatomic particles in its nucleusb. weighted average of the masses of the isotopes of the elementc. total mass of the isotopes of the elementd. average of the mass number and the atomic number for the element

____ 70. The atomic mass of an element depends upon the ____.a. mass of each electron in that elementb. mass of each isotope of that elementc. relative abundance of protons in that elementd. mass and relative abundance of each isotope of that element

____ 71. Which of the following is necessary to calculate the atomic mass of an element?a. the atomic mass of carbon-12b. the atomic number of the elementc. the relative masses of the element’s protons and neutronsd. the masses of each isotope of the element

____ 72. The principal quantum number indicates what property of an electron?a. position c. energy levelb. speed d. electron cloud shape

____ 73. Each period in the periodic table corresponds to ____.a. a principal energy level c. an orbitalb. an energy sublevel d. a suborbital

____ 74. The modern periodic table is arranged in order of increasing atomic ____.a. mass c. numberb. charge d. radius

____ 75. Of the elements Pt, V, Li, and Kr, which is a nonmetal?a. Pt c. Lib. V d. Kr

____ 76. The atomic number of an element is the total number of which particles in the nucleus?a. neutrons c. electronsb. protons d. protons and electrons

____ 77. What element has the electron configuration 1s 2s 2p 3s 3p ?a. nitrogen c. siliconb. selenium d. silver

____ 78. Which of the following is true about the electron configurations of the noble gases?a. The highest occupied s and p sublevels are completely filled.b. The highest occupied s and p sublevels are partially filled.c. The electrons with the highest energy are in a d sublevel.d. The electrons with the highest energy are in an f sublevel.

____ 79. Elements that are characterized by the filling of p orbitals are classified as ____.a. groups 3A through 8A c. inner transition metalsb. transition metals d. groups 1A and 2A

____ 80. Which of the following electron configurations is most likely to result in an element that is relatively inactive?

a. a half-filled energy sublevelb. a filled energy sublevelc. one empty and one filled energy subleveld. a filled highest occupied principal energy level

____ 81. Which subatomic particle plays the greatest part in determining the properties of an element? a. proton c. neutronb. electron d. none of the above

____ 82. Which of the following is true about the electron configurations of the representative elements?a. The highest occupied s and p sublevels are completely filled.b. The highest occupied s and p sublevels are partially filled.c. The electrons with the highest energy are in a d sublevel.d. The electrons with the highest energy are in an f sublevel.

____ 83. What are the Group 1A and Group 7A elements examples of?a. representative elements c. noble gasesb. transition elements d. nonmetallic elements

____ 84. How does atomic radius change from left to right across a period in the periodic table?a. It tends to decrease. c. It first increases, then decreases.b. It tends to increase. d. It first decreases, then increases.

____ 85. What causes the shielding effect to remain constant across a period?a. Electrons are added to the same principal energy level.b. Electrons are added to different principal energy levels.c. The charge on the nucleus is constant. d. The atomic radius increases.

____ 86. Atomic size generally ____.a. increases as you move from left to right across a periodb. decreases as you move from top to bottom within a groupc. remains constant within a periodd. decreases as you move from left to right across a period

____ 87. What element in the second period has the largest atomic radius? a. carbon c. potassiumb. lithium d. neon

____ 88. Which of the following statements is true about ions?a. Cations form when an atom gains electrons.b. Cations form when an atom loses electrons.c. Anions form when an atom gains protons.d. Anions form when an atom loses protons.

____ 89. The metals in Groups 1A, 2A, and 3A ____.a. gain electrons when they form ions c. all have ions with a 1 chargeb. all form ions with a negative charge d. lose electrons when they form ions

____ 90. Which of the following statements is NOT true about ions?a. Cations are positively charged ions.b. Anions are common among nonmetals.c. Charges for ions are written as numbers followed by a plus or minus sign.

d. When a cation forms, more electrons are transferred to it.

____ 91. Why is the second ionization energy greater than the first ionization energy?a. It is more difficult to remove a second electron from an atom.b. The size of atoms increases down a group.c. The size of anions decreases across a period.d. The nuclear attraction from protons in the nucleus decreases.

____ 92. Which of the following elements has the smallest ionic radius?a. Li c. Ob. K d. S

____ 93. What is the energy required to remove an electron from an atom in the gaseous state called?a. nuclear energy c. shielding energyb. ionization energy d. electronegative energy

____ 94. For Group 2A metals, which electron is the most difficult to remove?a. the firstb. the secondc. the thirdd. All the electrons are equally difficult to remove.

____ 95. Which of the following factors contributes to the decrease in ionization energy within a group in the periodic table as the atomic number increases?a. increase in atomic sizeb. increase in size of the nucleusc. increase in number of protons d. fewer electrons in the highest occupied energy level

____ 96. Which of the following elements has the smallest first ionization energy?a. sodium c. potassiumb. calcium d. magnesium

____ 97. Which of the following elements has the lowest electronegativity?a. lithium c. bromineb. carbon d. fluorine

____ 98. Which statement is true about electronegativity?a. Electronegativity is the ability of an anion to attract another anion.b. Electronegativity generally increases as you move from top to bottom within a group.c. Electronegativity generally is higher for metals than for nonmetals.d. Electronegativity generally increases from left to right across a period.

____ 99. Compared with the electronegativities of the elements on the left side of a period, the electronegativities of the elements on the right side of the same period tend to be ____.a. lower c. the sameb. higher d. unpredictable

____ 100. Which of the following decreases with increasing atomic number in Group 2A?a. shielding effect c. ionization energyb. ionic size d. number of electrons

____ 101. Which of the following statements correctly compares the relative size of an ion to its neutral atom?a. The radius of an anion is greater than the radius of its neutral atom.

b. The radius of an anion is identical to the radius of its neutral atom.c. The radius of a cation is greater than the radius of its neutral atom.d. The radius of a cation is identical to the radius of its neutral atom.

____ 102. What is the electron configuration of the calcium ion?a. 1s 2s 2p 3s 3p c. 1s 2s 2p 3s 3p 4sb. 1s 2s 2p 3s 3p 4s d. 1s 2s 2p 3s

____ 103. What is the electron configuration of the gallium ion?a. 1s 2s 2p 3s 3p c. 1s 2s 2p 3s 3p 4s 4pb. 1s 2s 2p 3s 3p 4s d. 1s 2s 2p 3s 3p 3d

____ 104. The octet rule states that, in chemical compounds, atoms tend to have ____.a. the electron configuration of a noble gasb. more protons than electronsc. eight electrons in their principal energy leveld. more electrons than protons

____ 105. What is the formula of the ion formed when tin achieves a stable electron configuration?a. Sn c. Snb. Sn d. Sn

____ 106. What is the formula of the ion formed when cadmium achieves a pseudo-noble-gas electron configuration?a. Cd c. Cdb. Cd d. Cd

____ 107. Which of the following is a pseudo-noble-gas electron configuration?a. 1s 2s 2p 3s 3d c. 1s 2s 2p 3s 3p 3db. 1s 2s 2p 3s 3p d. 1s 2s 2p 3s 3d 4s

____ 108. What is the electron configuration of the oxide ion (O )?a. 1s 2s 2p c. 1s 2sb. 1s 2s 2p d. 1s 2s 2p

____ 109. What is the electron configuration of the iodide ion?a. 1s 2s 2p 3s 3p 3d 4s 4p 4d 5s 5pb. 1s 2s 2p 3s 3p 3d 4s 4p 4dc. 1s 2s 2p 3s 3p 3d 4s 4p 4d 5sd. 1s 2s 2p 3s 3p 3d 4s 4p

____ 110. How many valence electrons are transferred from the nitrogen atom to potassium in the formation of the compound potassium nitride?a. 0 c. 2b. 1 d. 3

____ 111. How many valence electrons are transferred from the calcium atom to iodine in the formation of the compound calcium iodide?a. 0 c. 2

b. 1 d. 3

____ 112. What is the formula unit of sodium nitride?a. NaN c. Na Nb. Na N d. NaN

____ 113. What is the formula unit of aluminum oxide?a. AlO c. AlOb. Al O d. Al O

____ 114. What is the name of the ionic compound formed from lithium and bromine?a. lithium bromine c. lithium bromiumb. lithium bromide d. lithium bromate

____ 115. What is the formula for sodium sulfate?a. NaSO c. Na(SO )b. Na SO d. Na (SO )

____ 116. Alloys are commonly used in manufacturing. Which of the following is NOT a reason to use an alloy instead of a pure metal?a. Bronze is tougher than pure copper. c. Brass is more malleable than pure copper.b. Sterling silver is stronger than pure silver. d. Cast iron is more brittle than pure iron.

____ 117. Which of the following compounds has the formula KNO ?a. potassium nitrate c. potassium nitriteb. potassium nitride d. potassium nitrogen oxide

____ 118. Which is a typical characteristic of an ionic compound?a. Electron pairs are shared among atoms.b. The ionic compound has a low solubility in water.c. The ionic compound is described as a molecule.d. The ionic compound has a high melting point.

____ 119. 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 pairs of electrons.c. Two atoms share two electrons.d. Two atoms share one electron.

____ 120. Why do atoms share electrons in covalent bonds?a. to become ions and attract each otherb. to attain a noble-gas electron configurationc. to become more polard. to increase their atomic numbers

____ 121. Which of the following elements can form diatomic molecules held together by triple covalent bonds?a. carbon c. fluorineb. oxygen d. nitrogen

____ 122. Which noble gas has the same electron configuration as the oxygen in a water molecule?a. helium c. argon

b. neon d. xenon

____ 123. Which of the following diatomic molecules is joined by a double covalent bond?a. c.b. d.

____ 124. A molecule with a single covalent bond is ____.a. CO c. COb. Cl d. N

____ 125. When one atom contributes both bonding electrons in a single covalent bond, the bond is called a(n) ____.a. one-sided covalent bond c. coordinate covalent bondb. unequal covalent bond d. ionic covalent bond

____ 126. Once formed, how are coordinate covalent bonds different from other covalent bonds?a. They are stronger. c. They are weaker.b. They are more ionic in character. d. There is no difference.

____ 127. When H forms a bond with H O to form the hydronium ion H O , this bond is called a coordinate covalent bond because ____.a. both bonding electrons come from the oxygen atomb. it forms an especially strong bondc. the electrons are equally sharedd. the oxygen no longer has eight valence electrons

____ 128. Which of the following bonds is the least reactive?a. C—C c. O—Hb. H—H d. H—Cl

____ 129. In which of the following compounds is the octet expanded to include 12 electrons?a. H S c. PClb. PCl d. SF

____ 130. How is a pair of molecular orbitals formed?a. by the splitting of a single atomic orbitalb. by the reproduction of a single atomic orbitalc. by the overlap of two atomic orbitals from the same atomd. by the overlap of two atomic orbitals from different atoms

____ 131. The side-by-side overlap of p orbitals produces what kind of bond?a. alpha bond c. pi bondb. beta bond d. sigma bond

____ 132. Where are the electrons most probably located in a molecular bonding orbital?a. anywhere in the orbitalb. between the two atomic nucleic. in stationary positions between the two atomic nucleid. in circular orbits around each nucleus

____ 133. Sigma bonds are formed as a result of the overlapping of which type(s) of atomic orbital(s)?a. s only c. d only

b. p only d. s and p

____ 134. Which of the following bond types is normally the weakest?a. sigma bond formed by the overlap of two s orbitalsb. sigma bond formed by the overlap of two p orbitalsc. sigma bond formed by the overlap of one s and one p orbitald. pi bond formed by the overlap of two p orbitals

____ 135. What causes water molecules to have a bent shape, according to VSEPR theory?a. repulsive forces between unshared pairs of electronsb. interaction between the fixed orbitals of the unshared pairs of oxygenc. ionic attraction and repulsiond. the unusual location of the free electrons

____ 136. What type of hybrid orbital exists in the methane molecule?a. sp c. spb. sp d. sp d

____ 137. What is the shape of a molecule with a triple bond?a. tetrahedral c. bentb. pyramidal d. linear

____ 138. What type of hybridization occurs in the orbitals of a carbon atom participating in a triple bond with another carbon atom?a. c.b. d.

____ 139. How many pi bonds are formed when sp hybridization occurs in ethene, C H ?a. 0 c. 2b. 1 d. 3

____ 140. Which of the following atoms acquires the most negative charge in a covalent bond with hydrogen?a. C c. Ob. Na d. S

____ 141. A bond formed between a silicon atom and an oxygen atom is likely to be ____.a. ionic c. polar covalentb. coordinate covalent d. nonpolar covalent

____ 142. What causes hydrogen bonding?a. attraction between ionsb. motion of electronsc. sharing of electron pairsd. bonding of a covalently bonded hydrogen atom with an unshared electron pair

____ 143. Why is hydrogen bonding only possible with hydrogen?a. Hydrogen’s nucleus is electron deficient when it bonds with an electronegative atom.b. Hydrogen is the only atom that is the same size as an oxygen atom.c. Hydrogen is the most electronegative element.d. Hydrogen tends to form covalent bonds.

____ 144. In which of the following are the symbol and name for the ion given correctly?

a. Fe : ferrous ion; Fe : ferric ionb. Sn : stannic ion; Sn : stannous ionc. Co : cobalt(II) ion; Co : cobaltous iond. Pb : lead ion; Pb : lead(IV) ion

____ 145. Which of the following correctly provides the name of the element, the symbol for the ion, and the name of the ion?a. fluorine, F , fluoride ion c. copper, Cu , cuprous ionb. zinc, Zn , zincate ion d. sulfur, S , sulfurous ion

____ 146. The nonmetals in Groups 6A and 7A ____.a. lose electrons when they form ionsb. have a numerical charge that is found by subtracting 8 from the group numberc. all have ions with a –1 charged. end in -ate

____ 147. What determines that an element is a metal?a. the magnitude of its charge c. when it is a Group A elementb. the molecules that it forms d. its position in the periodic table

____ 148. What is the Stock name for chromic ion?a. chromium(I) ion c. chromium(III) ionb. chromium(II) ion d. chromium(IV) ion

____ 149. In which of the following are the symbol and name for the ion given correctly?a. NH : ammonia; H : hydride c. OH : hydroxide; O : oxideb. C H O : acetate; C O : oxalite d. PO : phosphate; PO : phosphite

____ 150. Which of the following correctly provides the names and formulas of polyatomic ions?a. carbonate: HCO ; bicarbonate: COb. nitrite: NO ; nitrate: NOc. sulfite: S ; sulfate: SOd. chromate: CrO ; dichromate: Cr O

____ 151. An -ate or -ite at the end of a compound name usually indicates that the compound contains ____.a. fewer electrons than protons c. only two elementsb. neutral molecules d. a polyatomic anion

____ 152. Which of the following is true about the composition of ionic compounds?a. They are composed of anions and cations.b. They are composed of anions only.c. They are composed of cations only.d. They are formed from two or more nonmetallic elements.

____ 153. Which of the following formulas represents an ionic compound?a. CS c. N Ob. BaI d. PCl

____ 154. Which element, when combined with fluorine, would most likely form an ionic compound?

a. lithium c. phosphorusb. carbon d. chlorine

____ 155. Which of the following shows correctly an ion pair and the ionic compound the two ions form?a. Sn , N ; Sn N c. Cr , I ; CrIb. Cu , O ; Cu O d. Fe , O ; Fe O

____ 156. In which of the following are the formula of the ionic compound and the charge on the metal ion shown correctly?a. UCl , U c. IrS , Irb. ThO , Th d. NiO, Ni

____ 157. Which of the following correctly represents an ion pair and the ionic compound the ions form?a. Ca , F ; CaF c. Ba , O ; Ba Ob. Na , Cl ; NaCl d. Pb , O ; Pb O

____ 158. In which of the following is the name and formula given correctly?a. sodium oxide, NaO c. cobaltous chloride, CoClb. barium nitride, BaN d. stannic fluoride, SnF

____ 159. Which of the following compounds contains the lead(II) ion?a. PbO c. Pb2Ob. PbCl4 d. Pb2S

____ 160. Which set of chemical name and chemical formula for the same compound is correct?a. iron(II) oxide, Fe O c. tin(IV) bromide, SnBrb. aluminum fluorate, AlF d. potassium chloride, K Cl

____ 161. What is the correct formula for potassium sulfite?a. KHSO c. K SOb. KHSO d. K SO

____ 162. Which set of chemical name and chemical formula for the same compound is correct?a. ammonium sulfite, (NH ) S c. lithium carbonate, LiCOb. iron(III) phosphate, FePO d. magnesium dichromate, MgCrO

____ 163. What type of compound is CuSO ?a. monotomic ionic c. polyatomic ionicb. polyatomic covalent d. binary molecular

____ 164. Sulfur hexafluoride is an example of a ____.a. monatomic ion c. binary compoundb. polyatomic ion d. polyatomic compound

____ 165. Which of the following correctly shows a prefix used in naming binary molecular compounds with its corresponding number?a. deca-, 7 c. hexa-, 8b. nona-, 9 d. octa-, 4

____ 166. Which of the following is a binary molecular compound?a. BeHCO c. AgIb. PCl d. MgS

____ 167. Which of the following formulas represents a molecular compound?a. ZnO c. SOb. Xe d. BeF

____ 168. When naming acids, the prefix hydro- is used when the name of the acid anion ends in ____.a. -ide c. -ateb. -ite d. -ic

____ 169. Which of the following shows both the correct formula and correct name of an acid?a. HClO , chloric acid c. H PO , phosphoric acidb. HNO , hydronitrous acid d. HI, iodic acid

____ 170. What is the name of H SO ?a. hyposulfuric acid c. sulfuric acidb. hydrosulfuric acid d. sulfurous acid

____ 171. When the name of an anion that is part of an acid ends in -ite, the acid name includes the suffix ____.a. -ous c. -ateb. -ic d. -ite

____ 172. What is the formula for sulfurous acid?a. H SO c. H SOb. H SO d. H S

____ 173. What is the formula for phosphoric acid?a. H PO c. HPOb. H PO d. HPO

____ 174. Which of the following pairs of substances best illustrates the law of multiple proportions?a. H and O c. CaCl and CaBrb. P O and PH d. NO and NO

____ 175. Select the correct formula for sulfur hexafluoride.a. S F c. F Sb. F SO d. SF

____ 176. What is the correct name for the compound CoCl ?a. cobalt(I) chlorate c. cobalt(II) chlorateb. cobalt(I) chloride d. cobalt(II) chloride

____ 177. Suppose you encounter a chemical formula with H as the cation. What do you know about this compound immediately?a. It is a polyatomic ionic compound. c. It is a base.

b. It is an acid. d. It has a 1 charge.

____ 178. Which of the following is the correct name for N O ?a. nitrous oxide c. nitrogen dioxideb. dinitrogen pentoxide d. nitrate oxide

____ 179. How many moles of tungsten atoms are in 4.8 10 atoms of tungsten?a. 8.0 10 moles c. 1.3 10 molesb. 8.0 10 moles d. 1.3 10 moles

____ 180. How many moles of silver atoms are in 1.8 10 atoms of silver?a. 3.0 10 c. 3.0 10b. 3.3 10 d. 1.1 10

____ 181. How many atoms are in 0.075 mol of titanium?a. 1.2 10-25 c. 6.4 10b. 2.2 10 d. 4.5 10

____ 182. How many molecules are in 2.10 mol CO ?a. 2.53 10 molecules c. 3.49 10 moleculesb. 3.79 10 molecules d. 1.26 10 molecules

____ 183. How many atoms are in 3.5 moles of arsenic atoms?a. 5.8 10 atoms c. 2.1 10 atomsb. 7.5 10 atoms d. 1.7 10 atoms

____ 184. Which of the following is NOT a true about atomic mass?a. The atomic mass is 12 g for magnesium.b. The atomic mass is the mass of one mole of atoms.c. The atomic mass is found by checking the periodic table.d. The atomic mass is the number of grams of an element that is numerically equal to the

mass in amu.

____ 185. What is true about the molar mass of chlorine gas?a. The molar mass is 35.5 g.b. The molar mass is 71.0 g.c. The molar mass is equal to the mass of one mole of chlorine atoms.d. none of the above

____ 186. What is the molar mass of AuCl3?a. 96 g c. 232.5 gb. 130 g d. 303.6 g

____ 187. What is the molar mass of (NH ) CO ?a. 144 g c. 96 gb. 138 g d. 78 g

____ 188. The molar mass of C H and the molar mass of CaCO contain approximately the same number of ____.

a. carbon atoms c. cationsb. anions d. grams

____ 189. What is the mass in grams of 5.90 mol C H ?a. 0.0512 g c. 389 gb. 19.4 g d. 673 g

____ 190. What is the number of moles in 432 g Ba(NO ) ?a. 0.237 mol c. 1.65 molb. 0.605 mol d. 3.66 mol

____ 191. What is the number of moles of beryllium atoms in 36 g of Be?a. 0.25 mol c. 45.0 molb. 4.0 mol d. 320 mol

____ 192. How many moles of CaBr are in 5.0 grams of CaBr ?a. 2.5 10 mol c. 4.0 10 molb. 4.2 10 mol d. 1.0 10 mol

____ 193. For which of the following conversions does the value of the conversion factor depend upon the formula of the substance?a. volume of gas (STP) to molesb. density of gas (STP) to molar massc. mass of any substance to molesd. moles of any substance to number of particles

____ 194. What is the mass of silver in 3.4 g AgNO ?a. 0.025 g c. 2.2 gb. 0.64 g d. 3.0 g

____ 195. What is the mass of oxygen in 250 g of sulfuric acid, H SO ?a. 0.65 g c. 16 gb. 3.9 g d. 160 g

____ 196. What is the volume, in liters, of 0.500 mol of C H gas at STP?a. 0.0335 L c. 16.8 Lb. 11.2 L d. 22.4 L

____ 197. What is the number of moles in 500 L of He gas at STP?a. 0.05 mol c. 22 molb. 0.2 mol d. 90 mol

____ 198. What is the number of moles in 9.63 L of H S gas at STP?a. 0.104 mol c. 3.54 molb. 0.430 mol d. 14.7 mol

____ 199. What is the density at STP of the gas sulfur hexafluoride, SF ?a. 0.153 g/L c. 3270 g/Lb. 6.52 g/L d. 3.93 10 g/L

____ 200. The molar mass of a certain gas is 49 g. What is the density of the gas in g/L at STP?a. 3.6 10 g/L c. 2.2 g/Lb. 0.46 g/L d. 71 g/L

____ 201. A 22.4-L sample of which of the following substances, at STP, would contain 6.02 10 representative particles?a. oxygen c. cesium iodideb. gold d. sulfur

____ 202. If the density of an unknown gas Z is 4.50 g/L at STP, what is the molar mass of gas Z?a. 0.201 g/mol c. 26.9 g/molb. 5.00 g/mol d. 101 g/mol

____ 203. If 60.2 grams of Hg combines completely with 24.0 grams of Br to form a compound, what is the percent composition of Hg in the compound?a. 28.5% c. 71.5%b. 39.9% d. 60.1%

____ 204. What is the percent composition of chromium in BaCrO ?a. 4.87% c. 20.5%b. 9.47% d. 25.2%

____ 205. If 20.0 grams of Ca combines completely with 16.0 grams of S to form a compound, what is the percent composition of Ca in the compound?a. 1.25% c. 44.4%b. 20.0% d. 55.6%

____ 206. What is the percent composition of carbon, in heptane, C H ?a. 12% c. 68%b. 19% d. 84%

____ 207. What is the percent by mass of carbon in acetone, C H O?a. 20.7% c. 1.61%b. 62.1% d. 30.0%

____ 208. Which expression represents the percent by mass of nitrogen in NH4NO3?a. 14 g N/80 g NH NO 100% c. 80 g NH NO /14 g N 100%b. 28 g N/80 g NH NO 100% d. 80 g NH NO /28 g N 100%

____ 209. What is the empirical formula of a compound that is 40% sulfur and 60% oxygen by weight?a. SO c. SOb. SO d. S O

____ 210. What is the empirical formula of a substance that is 53.5% C, 15.5% H, and 31.1% N by weight?a. C HN c. C H Nb. C H N d. CH N

____ 211. Which of the following sets of empirical formula, molar mass, and molecular formula is correct?a. CH, 78 g, C H c. CaO, 56 g, Ca O

b. CH N, 90 g, C H N d. C H O, 120 g, C H O

____ 212. Which of the following is the correct skeleton equation for the reaction that takes place when solid phosphorus combines with oxygen gas to form diphosphorus pentoxide?a. P(s) O (g) PO (g) c. P(s) O2(g) P O (s)b. P(s) O(g) P O (g) d. P O (s) P (s) O (g)

____ 213. What are the missing coefficients for the skeleton equation below?Cr(s) Fe(NO ) (aq) Fe(s) Cr(NO ) (aq)a. 4, 6, 6, 2 c. 2, 3, 3, 2b. 2, 3, 2, 3 d. 1, 3, 3, 1

____ 214. What are the missing coefficients for the skeleton equation below? Al (SO ) (aq) KOH(aq) Al(OH) (aq) K SO (aq)a. 1, 3, 2, 3 c. 4, 6, 2, 3b. 2, 12, 4, 6 d. 1, 6, 2, 3

____ 215. When potassium hydroxide and barium chloride react, potassium chloride and barium hydroxide are formed. The balanced equation for this reaction is ____.a. KH BaCl KCl BaH c. 2KOH BaCl 2KCl Ba(OH)b. KOH BaCl KCl BaOH d. KOH BaCl KCl BaOH

____ 216. Which of the following is a balanced equation representing the decomposition of lead(IV) oxide?a. PbO Pb 2O c. Pb O 2Pb Ob. PbO Pb O d. PbO Pb O

____ 217. What are the correct formulas and coefficients for the products of the following double-replacement reaction?RbOH H PO a. Rb(PO ) H O c. Rb PO 3H Ob. RbPO 2H O d. H Rb PO OH

____ 218. When the equation for the complete combustion of one mole of C H OH is balanced, the coefficient for oxygen is ____.a. 13

2c. 7

2b. 11

2d. 9

2

____ 219. Which of the following statements is NOT true about the decomposition of a simple binary compound?a. The products are unpredictable.b. The products are the constituent elements.c. The reactant is a single substance.d. The reactant could be an ionic or a molecular compound.

____ 220. Which of the following statements is true about single-replacement reactions?a. They are restricted to metals. c. Two reactants produce two products.b. They involve a single product. d. Any metal replaces any other metal.

____ 221. In the activity series of metals, which metal(s) will displace hydrogen from an acid?a. only metals above hydrogen c. any metalb. only metals below hydrogen d. only metals from Li to Na

____ 222. Use the activity series of metals to complete a balanced chemical equation for the following single replacement reaction. Ag(s) KNO (aq) a. AgNO Kb. AgK NOc. AgKNOd. No reaction takes place because silver is less reactive than potassium.

____ 223. Which of the following statements is NOT true about double-replacement reactions?a. The product may precipitate from solution.b. The product may be a gas.c. The product may be a molecular compound.d. The reactant may be a solid metal.

____ 224. In a double-replacement reaction, ____.a. the reactants are usually a metal and a nonmetalb. one of the reactants is often waterc. the reactants are generally two ionic compounds in aqueous solutiond. energy in the form of heat or light is often produced

____ 225. A double-replacement reaction takes place when aqueous Na CO reacts with aqueous Sn(NO ) . You would expect one of the products of this reaction to be ____.a. NaNO c. Sn(CO )b. NaSn d. CNO

____ 226. The complete combustion of which of the following substances produces carbon dioxide and water?a. C H c. CaHCOb. K CO d. NO

____ 227. Which of the following is the correctly balanced equation for the incomplete combustion of heptene, C H ?a. C H 14O 7CO 7H O c. 2C H 21O 14CO 14H Ob. C H 7O 7CO 7H O d. C H O C O 7H

____ 228. If a combination reaction takes place between rubidium and bromine, the chemical formula for the product is ____.a. RuBr c. RbBrb. Rb Br d. RbBr

____ 229. What is the balanced chemical equation for the reaction that takes place between bromine and sodium iodide?a. Br NaI NaBr I c. Br NaI NaBrIb. Br 2NaI 2NaBr I d. Br NaI NaBr I

____ 230. What is the driving force in the following reaction?

Ni(NO ) (aq) K S(aq) NiSs 2KNO (aq)a. A gas is formed. c. Ionic compounds are reactants.b. A precipitate is formed. d. Ionic compounds are products.