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Quiz 1
____ 1. A 19.0-g sample of lithium is completely burned in air to form lithium oxide. The mass of lithium oxide must be A) less than 19.0 g. B) greater than 19.0 g. C) equal to 19.0 g. D) all of the above. E) none of the above.
____ 2. Express the result of the following calculation in scientific notation: 0.0263 cm^2 ÷ 88.2 cm A) B) C) D) E)
____ 3. Calculate the mass of gold that occupies the same volume as 62.9 g of cobalt. The density of cobalt is 8.90 g/cm^3 and the density of gold is 19.30 g/cm^3. A) 2.73 g B) 136 g C) (^) 1.08 104 g D) 0.0345 g E) 0.366 g
____ 4. An atom that has the same number of neutrons as is A) (^). B) (^). C) (^). D) (^). E) (^).
____ 5. What is the symbol of the nuclide having 15 protons and 16 neutrons? A)
B)
C)
D)
E)
____ 6. Naturally occurring element X exists in three isotopic forms: X-28 (27.977 amu, 92.21% abundance), X-29 (28.976 amu, 4.70% abundance), and X-30 (29.974 amu, 3.09% abundance). Calculate the atomic mass of X. A) 29.09 amu
B) 28.09 amu C) 35.29 amu D) 86.93 amu E) 25.80 amu
____ 7. The formula of magnesium sulfide is A) MgS. B) MgSO 2. C) MgSO 4. D) MgSO 3. E) Mg(SO 4 ) 2.
____ 8. The formula of water, H 2 O, suggests A) there is twice as much mass of hydrogen as oxygen in each molecule. B) there are two oxygen atoms and one hydrogen atom per water molecule. C) there is twice as much mass of oxygen as of hydrogen in each molecule. D) there are two hydrogen atoms and one oxygen atom per water molecule. E) none of these
____ 9. What is the best answer to report for?
A) 2.252 g/mL B) 2.2518 g/mL C) 2.3 g/mL D) 2.25 g/mL E) 2.25183 g/mL
____ 10. The mass spectrum of an element with two naturally occurring isotopes is shown below. What is the best estimate of the element’s atomic mass?
A) 10 amu B) 11 amu C) 10.8 amu D) 10.2 amu E) 10.5 amu
____ 7. A particular compound contains, by mass, 41.4 % carbon, 3.47 % hydrogen, and 55.1 % oxygen. A 0.050-mol sample of this compound weighs 5.80 g. The molecular formula of this compound is A) C 3 H 3 O 3. B) C 3 H 3 O. C) CHO. D) C 4 H 4 O 4. E) C 5 H 5 O 5.
____ 8. A 5.95-g sample of AgNO 3 is reacted with BaCl 2 according to the equation
2AgNO 3 ( aq ) + BaCl 2 ( aq ) 2AgCl( s ) + Ba(NO 3 ) 2 ( aq ) to give 3.36 g of AgCl. What is the percent yield of AgCl? A) 44.6 % B) 33.5 % C) 66.9 % D) 56.5 % E) 100 %
____ 9. Identify the spectator ions in the following reaction.
Ca2+( aq ) + 2NO 3 – ( aq ) + 2Na+( aq ) + CO 32 – ( aq ) CaCO 3 ( s ) + 2Na+( aq ) + NO 3 – ( aq ) A) NO 3 –^ and CO 32 – B) Ca2+^ and Na+ C) Ca2+^ and CO 32 – D) Ca2+^ and NO 3 – E) Na+^ and NO 3 –
____ 10. Which of the following solutions contains the largest number of moles of dissolved particles? A) 25. mL of 5.0 M sodium chloride B) 25. mL of 2.0 M sulfuric acid C) 200. mL of 0.10 M sodium hydroxide D) 50. mL of 1.0 M hydrochloric acid E) 100. mL of 0.5 M nitric acid
Quiz 3
____ 1. A flexible container is charged with 51.00 L of gas at 368. Under conditions of constant pressure and moles of gas, what is the volume of the gas when the temperature is tripled? A) 153 L B) 17.0 L C) 0.0588 L D) 51 L E) 0.850 L
____ 2. A fixed amount of gas in a rigid container is heated from 300 to 900 K. Which of the following responses best describes what will happen to the pressure of the gas? A) The pressure will increase by a factor of 3. B) The pressure will increase by a factor less than 3. C) The pressure will increase by a factor greater than 3. D) The pressure will decrease by a factor of 3. E) The pressure will remain the same.
____ 3. What volume of methane gas, CH 4 , has the same number of atoms as 6.00 L of krypton gas at the same temperature and pressure? A) 36.0 L B) 1.00 L C) 6.00 L D) 1.20 L E) 30.0 L
____ 4. The volume of 1 mol of nitrogen A) is lower than that of 1mol ammonia at high pressures. B) is decreased by decreasing the pressure of the gas. C) has the value of 22.4 L at 0°C and 1.00 atm. D) is decreased by increasing its kinetic energy. E) is increased by decreasing the temperature.
____ 5. A 1.00-L sample of a gas at STP has a mass of 1.16 g. The molar mass of the gas is A) 5.18 g/mol. B) 26.0 g/mol. C) 22.4 g/mol. D) 44.8 g/mol. E) 193 g/mol.
____ 6. Calcium nitrate will react with ammonium chloride at slightly elevated temperatures, as represented in the equation below. Ca(NO 3 ) 2 ( s ) + 2NH 4 Cl( s ) 2N 2 O( g ) + CaCl 2 ( s ) + 4H 2 O( g ) What is the maximum volume of N 2 O at STP that could be produced using a 3.40-mol sample of each reactant? A) (^) 9.28 102 L B) 152 L
CHM 107-Q4 Fall 17-
____ 1. (1 point) An orbital with the quantum numbers may be found in which
subshell? A) 3f B) 3d C) 3p D) 3g E) 3s
____ 2. (1 point) What is the binding energy (in J/mol or kJ/mol) of an electron in a metal whose threshold frequency for photoelectrons is 2.50 × 10^14 /s? A) 2.75 × 10–^43 J/mol B) 1.66 × 10–^19 J/mol C) 1.20 × 10–^6 J/mol D) 99.7 kJ/mol E) 7.22 × 10^17 kJ/mol
____ 3. (1 point) When photons with a wavelength of 310. nm strike a magnesium plate, the maximum velocity of the ejected electrons is 3.45 × 10^5 m/s. Calculate the binding energy of electrons to the magnesium surface. A) 32.7 kJ/mol B) 321 kJ/mol C) 353 kJ/mol D) 386 kJ/mol E) 419 kJ/mol
____ 4. (1 point) Which ground-state electron configuration is incorrect? A) B) C) D) E)
____ 5. (1 point) Rank the following atoms in order of the largest to smallest atomic radius: Al, P, Cl, K. A) K > Al > P > Cl B) Al > K > P > Cl C) P > Al > K > Cl D) Al > P > Cl > K E) K > P > Al > Cl
____ 6. (1 point) An atom of which of the following elements has the largest second ionization energy? A) Na B) Cl C) S D) Si E) Mg
____ 7. (1 point) An atom of which of the following elements has the most negative electron affinity? A) Rb B) As C) Cl D) Br E) Se
____ 8. (1 point) What is the frequency of light emitted when the electron in a hydrogen atom undergoes a
transition from level ( ) A) B) C) D) E)
____ 9. (1 point) Which of the following statements is true concerning the electron configuration [Kr]5p^2? A) This configuration cannot be the ground-state electron configuration for a Sr atom because it violates the Pauli exclusion principle. B) This configuration cannot be the ground-state electron configuration for a Sr atom because it violates Hund's rule. C) This configuration is the ground-state electron configuration for a Sr atom. D) This configuration cannot be the ground-state electron configuration for a Sr atom because it violates the Heisenberg uncertainty principle. E) This configuration cannot be the ground-state electron configuration for a Sr atom because it violates the Aufbau principle.
____ 10. (1 point) Which of the following forms the most stable anion in the gas phase? A) Br (electron affinity = -325 kJ/mol) B) I (electron affinity = -295 kJ/mol) C) Te (electron affinity = -190 kJ/mol) D) C (electron affinity = -122 kJ/mol) E) As (electron affinity = -77 kJ/mol)
E)
____ 4. (1 point) The Lewis formula of which species does not represent an exception to the octet rule? A) SiF 5 - B) SCl 6 C) SF 4 D) BF 3 E) CF 3 -
____ 5. (1 point) Which of the following molecules has an incorrect Lewis formula? A)
B)
C)
D)
E)
____ 6. (1 point) Which of the following is the best explanation for a covalent bond? A) electrons simultaneously attracted by more than one nucleus B) an interaction between outer electrons C) the overlapping of unoccupied orbitals of two or more atoms D) the overlapping of two electron-filled orbitals having different energies E) a positive ion attracting negative ions
____ 7. (1 point) Which of the following concerning electronegativity is/are correct?
- Differences in element electronegativities may be used to predict the type of bonding, ionic or covalent, in a substance.
- The larger the differences in electronegativity between two bonded atoms the more polar the bond.
- The electrons in a polar bond tend to spend more time around the least
electronegative element.
A) 1 only B) 2 only C) 3 only D) 1 and 2 E) 1, 2, and 3
____ 8. (1 point) Rank the following covalent bonds in order of decreasing polarity: C-H, N-H, O-H, F-H. A) F-H, O-H, N-H, C-H B) O-H, F-H, N-H, C-H C) N-H, F-H, O-H, C-H D) C-H, N-H, O-H, F-H E) C-H, F-H, O-H, N-H
____ 9. (1 point) The Lewis formula for phosphine, PH 3 , has A) four lone pairs. B) four bonding pairs. C) two bonding pairs and two lone pairs. D) three bonding pairs and one lone pair. E) one bonding pair and three lone pairs.
____ 10. (1 point) Which one of the following has a Lewis formula most similar to that of NO–? A) O 2 B) O 22 – C) O 2 – D) NO+ E) NO
A) sulfur tetrafluoride, SF 4 B) iodine trichloride, ICl 3 C) nitrogen trifluoride, NF 3 D) phosphorus pentafluoride, PF 5 E) sulfur dioxide, SO 2
____ 7. (1 point) When an atom in a molecule or ion is described as sp^3 d^2 hybridized, its molecular geometry is A) octahedral. B) trigonal bipyramidal. C) linear. D) tetrahedral. E) trigonal planar.
____ 8. (1 point) Consider the molecule:
Specify the hybridization of each carbon atom.
C-1 C-2 C-3 C-4 C-
A) sp^2 sp^2 sp^2 sp^3 sp B) sp^2 sp^2 sp^2 sp^3 sp^3 C) sp^2 sp^2 sp^3 sp^3 sp^2 D) sp^2 sp^2 sp^3 sp^3 sp^3 E) sp^2 sp sp sp^2 sp
____ 9. (1 point) The bond angles in ICl 2 –^ are expected to be A) a little less than 109.5°. B) 109.5°. C) a little more than 109.5°. D) 120°. E) 180°.
____ 10. (1 point) An sp^3 hybridized central carbon atom with no lone pairs of electrons has what type of bonding? A) (^1) and 2 bonds B) (^1) and 3 bonds C) (^2) and 2 bonds D) (^3) and 2 bonds E) (^0) and 4 bonds
Chm107 Quiz 7
____ 1. (1 point) Which one of the following liquids would you expect to have the highest vapor pressure at room temperature? (all boiling points are normal boiling points) A) n-pentane, b.p. = 36.1°C B) methanol, b.p. = 65.0°C C) carbon tetrachloride, b.p. = 76.7°C D) acetic acid, b.p. = 118°C E) mercury, b.p. = 357°C
____ 2. (1 point) Methane (CH 4 ) is able to be liquefied at low temperatures due to which intermolecular force? A) ionic bonding B) covalent bonding C) hydrogen bonding D) dipole–dipole E) London dispersion
____ 3. (1 point) The space-filling representation of a crystalline polonium provided below is an example of a _____ unit cell, which contains the equivalent of _____ atom(s) within a single unit cell.
A) simple cubic, 1 atom B) body centered cubic, 2 atoms C) face centered cubic, 4 atoms D) simple cubic, 8 atoms E) body centered cubic, 3 atoms
____ 4. (1 point) The metal iron crystallizes in a body-centered cubic lattice. If the density of iron is 7.87 g/cm^3 , what is the unit cell edge length? A) 287 pm B) 77.6 pm C) 75.0 pm D) 61.6 pm E) 228 pm
____ 5. (1 point) What is the simplest formula of the compound represented by the unit cell provided below?
____ 9. (1 point) What is the standard enthalpy of formation of liquid methylamine (CH 3 NH 2 )?
C( s ) + O 2 ( g ) CO 2 ( g ); H ° = –393.5 kJ 2H 2 O( l ) 2H 2 ( g ) + O 2 ( g ); H ° = 571.6 kJ N 2 ( g ) + O 2 ( g ) NO 2 ( g ); H ° = 33.10 kJ 4CH 3 NH 2 ( l ) + 13O 2 ( g ) 4CO 2 ( g ) + 4NO 2 ( g ) + 10H 2 O(l); H ° = –4110.4 kJ A) +3899.2 kJ/mol B) – 3899.2 kJ/mol C) – 47.3 kJ/mol D) +47.3 kJ/mol E) +3178.4 kJ
____ 10. (1 point) Which of the following has a standard enthalpy of formation value of zero at 25°C? A) Cl( g ) B) Cl 2 ( l ) C) Cl 2 ( g ) D) Cl( s ) E) Cl 2 ( s )
Chm107 Quiz 8
____ 1. (1 point) What is H ° for the following phase change?
KCl( s ) KCl( l ) Substance (^) H ° f (kJ/mol) KCl( s ) – 436. KCl( l ) – 421.
A) 858.47 kJ B) 14.89 kJ C) – 858.47 kJ D) – 14.89 kJ E) 0 kJ
____ 2. (1 point) Which of the following is true for the sublimation of a solid substance? A) (^) S = 0 and H = 0. B) (^) S < 0 and H < 0. C) (^) S < 0 and H > 0. D) (^) S > 0 and H > 0. E) (^) S > 0 and H < 0.
____ 3. (1 point) At the normal boiling point of o-xylene, H °vap = 36.2 kJ/mol and S °vap = 86.7 J/(molK). What is the normal boiling point of o-xylene? A) 314 K B) 373 K C) 115 K D) 867 K E) 418 K
____ 4. (1 point) Assuming H and S are constant with respect to temperature, under what conditions will a chemical reaction be spontaneous only at low temperatures? A) (^) H is negative, and S is negative. B) (^) H is positive, and S is positive. C) S = 0, and H is positive. D) (^) H = 0, and S is negative. E) none of these
____ 5. (1 point) In which reaction is S ° expected to be negative? A) (^) 2C 2 H 6 ( g ) + 7O 2 ( g ) 4CO 2 ( g ) + 6H 2 O( l ) B) (^) Ga( l ) Ga( s ) C) (^) H 2 O( l ) + 2SO 2 ( g ) H 2 SO 4 ( l ) D) (^) CO 2 ( g ) CO 2 ( s ) E) all of above
____ 6. (1 point) What is the change in entropy when 0.646 g of water decomposes to form hydrogen gas and oxygen gas at 298 K?
C) – 41.3 kJ D) – 47.4 kJ E) (^) 9.16 103 kJ
Quiz 1
Answer Section
- ANS: B PTS: 1 DIF: easy REF: 1. OBJ: Apply the law of the conservation of mass. (Example 1.1) TOP: general concepts | matter
- ANS: D PTS: 1 DIF: easy REF: 1. OBJ: Know how to represent numbers using scientific notation. TOP: general concepts | measurement KEY: significant figures | scientific notation MSC: general chemistry
- ANS: B PTS: 1 DIF: difficult REF: 1. OBJ: Use density to relate mass and volume. (Example 1.5) TOP: general concepts | measurement KEY: SI unit | density MSC: general chemistry
- ANS: B PTS: 1 DIF: easy REF: 2. OBJ: Write the nuclide symbol for a given nuclide. TOP: early atomic theory | atomic theory of matter KEY: atomic symbol MSC: general chemistry
- ANS: C PTS: 1 DIF: moderate REF: 2. OBJ: Write the nuclide symbol of an element. (Example 2.1) TOP: early atomic theory | atomic theory of matter KEY: atomic symbol MSC: general chemistry
- ANS: B PTS: 1 DIF: moderate REF: 2. OBJ: Determine the atomic mass of an element from the isotopic masses and fractional abundances. (Example 2.2) TOP: early atomic theory | atomic theory of matter KEY: atomic weight MSC: general chemistry
- ANS: A PTS: 1 DIF: easy REF: 2. OBJ: Write the formula of an ionic compound from its name. (Example 2.5) TOP: early atomic theory | chemical substance KEY: nomenclature of simple compound | ionic compound MSC: general chemistry
- ANS: D PTS: 1 DIF: easy REF: 2. OBJ: Define and provide examples for the term formula unit. TOP: early atomic theory | chemical substance KEY: chemical formula MSC: general chemistry
- ANS: D PTS: 1 DIF: moderate REF: 1. OBJ: Use significant figures in calculations. (Example 1.2) TOP: general concepts | measurement KEY: significant figures MSC: general chemistry
- ANS: C PTS: 1 DIF: moderate REF: 2. OBJ: Describe how a mass spectrometer can be used to determine the fractional abundance of the isotopes of an element. TOP: early atomic theory | atomic theory of matter
