Chem lab Final
B
1. Identify a liquid. A) definite volume and definite shape B) definite volume and no definite shape C) no definite shape and definite volume D) no definite shape and no definite volume
B
10. Write the name for Sn(SO4)2. A) tin (I) sulfite B) tin (IV) sulfate C) tin sulfide D) tin (II) sulfite E) tin (I) sulfate
D
100) What mass (in kg) does 5.01 moles of iron have? A) 0.352 kg B) 0.122 kg C) 0.820 kg D) 0.280 kg E) 0.632 kg
A
101) A covalent bond is best described as: A) the sharing of electrons between atoms. B) the transfer of electrons. C) a bond between a metal and a nonmetal. D) a bond between a metal and a polyatomic ion. E) a bond between two polyatomic ions
B
102) What is the empirical formula for Hg2(NO3)2? A) Hg2(NO3)2 B) HgNO3 C) Hg(NO3)2 D) Hg2NO3 E) Hg4(NO3)4
D
103 Determine the name for CoCl2•6H2O. Remember that Co forms several ions. A) cobalt chloride hydrate B) cobalt (I) chloride heptahydrate C) cobalt (II) chloride heptahydrate D) cobalt (II) chloride hexahydrate E) cobalt (I) chloride
D
104) Give the correct formula for aluminum sulfate. A) Al2SO4 B) Al(SO4)3 C) Al3(SO4)2 D) Al2(SO4)3 E) None of these
A
105) Give the formula for sulfurous acid. A) H2SO3 B) HSO3 C) H2SO4 D) HSO4 E) None of these
A
106) Determine the empirical formula for a compound that contains C, H and O. It contains 52.14% C and 34.73% O by mass. A) C2H6O B) CHO C) C4H13O2 D) CH4O3 E) CH3O
B
107) Write a balanced equation to show the reaction of sulfurous acid with lithium hydroxide to form water and lithium sulfite. A) H2SO4(aq) + LiOH(aq) → H2O(l) + Li2SO4(aq) B) H2SO3(aq) + 2 LiOH(aq) → 2 H2O(l) + Li2SO3(aq) C) HSO3(aq) + LiOH(aq) → H2O(l) + LiSO3(aq) D) HSO4(aq) + LiOH(aq) → H2O(l) + LiSO4(aq) E) H2S(aq) + 2 LiOH(aq) → 2 H2O(l) + Li2S(aq)
A
108) Which of the following is one possible form of pentane? A) CH3CH2CH2CH2CH3 B) CH3CH=CHCH2CH3 C)CH3CH2CH2CH2CH2CH3 D) CH3CH2CH2CH2CH2NH2 E) CH3CH2-O-CH2CH2CH3
E
109) How many sodium ions are contained in 99.6 mg of Na2SO3? The molar mass of Na2SO3 is 126.05 g/mol. A) 1.52 × 1027 sodium ions B) 4.76 × 1020 sodium ions C) 2.10 × 1021 sodium ions D) 1.05 × 1021 sodium ions E) 9.52 × 1020 sodium ions
E
11. Write the formula for copper (II) sulfate pentahydrate. A) Cu2SO3•H5 B) Cu2S•H2O C) CuS•5H2O D) (CuSO4)5 E) CuSO4•5H2O
A
110) According to the following reaction, how many grams of sulfur are formed when 37.4 g of water are formed? 2 H2S(g) + SO2(g) → 3 S(s) + 2H2O(l) A) 99.8 g S B) 66.6 g S C) 56.1 g S D) 44.4 g S E) 14.0 g S
C
111) Give the theoretical yield, in moles of CO2 from the reaction of 4.00 moles of C8H18 with 4.00 moles of O2. 2 C8H18 + 25 O2 → 16 CO2 + 18 H2O A) 0.640 moles B) 32.0 moles C) 2.56 moles D) 16.0 moles E) 64.0 moles
B
112) Give the percent yield when 28.16 g of CO2 are formed from the reaction of 4.000 moles of C8H18 with 4.000 moles of O2. 2 C8H18 + 25 O2 → 16 CO2 + 18 H2O A) 20.00% B) 25.00% C) 50.00% D) 12.50% E) None of these
C
113) Determine the limiting reactant (LR) and the mass (in g) of nitrogen that can be formed from 50.0 g N2O4 and 45.0 g N2H4. Some possibly useful molar masses are as follows: N2O4 = 92.02 g/mol, N2H4 = 32.05 g/mol. N2O4(l) + 2 N2H4(l) → 3 N2(g) + 4 H2O(g) A) LR = N2H4, 59.0 g N2 formed B) LR = N2O4, 105 g N2 formed C) LR = N2O4, 45.7 g N2 formed D) LR = N2H4, 13.3 g N2 formed E) No LR, 45.0 g N2 formed
A
114) Determine the molarity of a solution formed by dissolving 97.7 g LiBr in enough water to yield 750.0 mL of solution. A) 1.50 M B) 1.18 M C) 0.130 M D) 0.768 M E) 2.30 M
A
115) Give the net ionic equation for the reaction (if any) that occurs when aqueous solutions of H2SO4 and KOH are mixed. A) H+(aq) + OH-(aq) → H2O(l) B) 2 K+(aq) + SO42-(aq) → K2SO4(s) C) H+(aq) + OH-(aq) + 2 K+(aq) + SO42-(aq) → H2O(l) + K2SO4(s) D) H22+(aq) + OH-(aq) → H2(OH)2(l) E) No reaction occurs.
C
116) What element is undergoing oxidation (if any) in the following reaction? CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g) A) O B) H C) C D) both C and H E) None of the elements
B
117) Which of the following pairs of aqueous solutions will form a precipitate when mixed? A) NH4NO3 + Li2CO3 B) Hg2(NO3)2 + LiCl C) NaCl + Li3PO4 D) AgC2H3O2 + Cu(NO3)2 E) None of the above solution pairs will produce a precipitate.
D
118) What volume (in mL) will a sample of F2 gas occupy in a syringe at 5.5 atm, if the F2 has a volume of 25.0 mL at 1.2 atm? A) 11 mL B) 17 mL C) 3.8 mL D) 5.5 mL E) 7.6 mL
B
119) If a sample of 0.29 moles of Ar occupies 3.8 L under certain conditions, what volume will 0.66 moles occupy under the same conditions? A) 12 L B) 8.6 L C) 17 L D) 5.0 L E) 15 L
E
12. Determine the empirical formula for a compound that is found to contain 10.15 mg P and 34.85 mg Cl. A) P3Cl B) PCl C) PCl2 D) P2Cl3 E) PCl3
E
120) To what temperature must a balloon, initially at 25°C and 2.00 L, be heated in order to have a volume of 6.00 L? A) 993 K B) 403 K C) 75 K D) 655 K E) 894 K
A
121) Determine the density of CO2 gas at STP. A) 1.96 g/L B) 1.80 g/L C) 2.24 g/L D) 4.46 g/L E) 5.10 g/L
A
122) A mixture of 10.0 g of Ne and 10.0 g Ar have a total pressure of 1.6 atm. What is the partial pressure of Ne? A) 1.1 atm B) 0.80 atm C) 0.54 atm D) 0.40 atm E) 1.3 atm
D
123) Which of the gases in the graph below has the largest molar mass? A) A B) B C) C D) D E) There is not enough information to determine.
B
124) Calculate the root mean square velocity of nitrogen molecules at 25°C. A) 729 m/s B) 515 m/s C) 149 m/s D) 297 m/s E) None of these
D
125) The rate of effusion of oxygen to an unknown gas is 0.935. What is the other gas? A) Ne B) Ar C) F2 D) N2 E) Br2
A
126) Which of the following signs on q and w represent a system that is doing work on the surroundings, as well as gaining heat from the surroundings? A) q = +, w = - B) q = -, w = + C) q = +, w = + D) q = -, w = -
C
127) Determine the specific heat capacity of an alloy that requires 59.3 kJ to raise the temperature of 150.0 g alloy from 298 K to 398 K. A) 4.38 J/g°C B) 2.29 J/g°C C) 3.95 J/g°C D) 2.53 J/g°C E) 1.87 J/g°C
D
128) Calculate the change in internal energy (ΔE) for a system that is giving off 25.0 kJ of heat and is changing from 12.00 L to 6.00 L in volume at 1.50 atm pressure. (Remember that 101.3 J= 1 L•atm) A) +25.9 kJ B) -16.0 kJ C) -25.9 kJ D) -24.1 kJ E) 937 kJ
A
129) A 35.6 g sample of ethanol (C2H5OH) is burned in a bomb calorimeter, according to the following reaction. If the temperature rose from 35.0 to 76.0°C and the heat capacity of the calorimeter is 23.3 kJ/°C, what is the value of ΔH°rxn? The molar mass of ethanol is 46.07 g/mol. C2H5OH(l) + 3 O2(g) → 2 CO2(g) + 3 H2O(g) ΔH°rxn = ? A) -1.24 × 103 kJ/mol B) +1.24 × 103 kJ/mol C) -8.09 × 103 kJ/mol D) -9.55 × 103 kJ/mol E) +9.55 × 103 kJ/mol
A
13. Identify the compound with covalent bonds. A) CH4 B) Kr C) KBr D) Li E) NaCl
B
130) A piece of iron (mass = 25.0 g) at 398 K is placed in a styrofoam coffee cup containing 25.0 mL of water at 298 K. Assuming that no heat is lost to the cup or the surroundings, what will the final temperature of the water be? Density of water =1g/ml; specific heat capacity of iron = 0.449 J/g°C and water = 4.18 J/g°C. A) 348 K B) 308 K C) 287 K D) 325 K E) 388 K
B
131) Use the standard reaction enthalpies given below to determine ΔH°rxn for the following reaction: 2 NO(g) + O2(g) → 2 NO2(g) ΔH°rxn = ? Given: N2(g) + O2(g) → 2 NO(g) ΔH°rxn = +183 kJ 1/2 N2(g) + O2(g) → NO2(g) ΔH°rxn = +33 kJ A) -150. kJ B) -117 kJ C) -333 kJ D) +115 kJ E) +238 kJ
A
132) Use the information provided to determine ΔH°rxn for the following reaction: ΔH°f (kJ/mol) CH4(g) + 4 Cl2(g) → CCl4(g) + 4 HCl(g) ΔH°rxn = ? CH4(g) -75 CCl4(g) -96 HCl(g) -92 A) -389 kJ B) -113 kJ C) +113 kJ D) -71 kJ E) +79 kJ
A
133) Define molar heat capacity. A) the quantity of heat required to raise the temperature of 1 mole of a substance by 1°C B) the quantity of heat required to change a system's temperature by 1°C C) the quantity of heat required to raise the temperature of 1 gram of a substance by 1°C D) the quantity of heat required to lower the temperature of 1 g of a substance by 1°F E) the quantity of heat required to lower the temperature of 1 liter of a substance by 1°C
A
134) Which of the following occur as the wavelength of a photon increases? A) the frequency decreases B) the energy increases C) the speed decreases D) Planck's constant decreases E) None of the above occur as the wavelength of a photon increases.
C
135) Calculate the energy of the orange light emitted, per photon, by a neon sign with a frequency of 4.89 × 1014 Hz. A) 3.09 × 10-19 J B) 6.14 × 10-19 J C) 3.24 × 10-19 J D) 1.63 × 10-19 J E) 5.11 × 10-19 J
C
136) For n = 3, what are the possible sublevels? A) 0 B) 0, 1 C) 0, 1, 2 D) 0, 1, 2, 3 E) none of these
E
137) How many orbitals are contained in the third principal level (n = 3) of a given atom? A) 9 B) 3 C) 18 D) 7 E) 5
A
138) Calculate the wavelength of light associated with the transition from n = 1 to n = 3 in the hydrogen atom. A) 103 nm B) 155 nm C) 646 nm D) 971 nm E) 136 nm
C
139) Give the numbers for ml for an f orbital. A) 0, 1, 2, 3 B) 1, 2, 3, 4 C) -3, -2, -1, 0, 1, 2, 3 D) -2, -1, 0, 1, 2 E) None of these
B
14. The solid compound, K2SO4, contains A) K+, S6+, and O2- ions. B) K+ and SO4 -2 ions. C) K2+ and SO4 -2 ions. D) K2SO4 molecules.
D
140) Identify the correct values for a 2p sublevel. A) n = 3, l = 1, ml = 0 B) n = 2, l = 1, ml = +2 C) n = 1, l = 0, ml = 0 D) n = 2, l = 1, ml = 0 E) n = 4, l = -1, ml = 0
B
141) Describe the shape of a p orbital. A) spherical B) dumbbell shaped C) three balls D) four balls E) eight balls
A
142) No two electrons can have the same four quantum numbers is known as the A) Pauli exclusion principle B) Hund's rule C) Aufbau principle D) Heisenberg uncertainty principle E) None of these
C
143) Only two electrons, with opposing spins, are allowed in each orbital is known as the A) Pauli exclusion principle B) Hund's rule C) Aufbau principle D) Heisenberg uncertainty principle E) None of these
B
144) When filling degenerate orbitals, electrons fill them singly first, with parallel spins is known as A) Pauli exclusion principle B) Hund's rule C) Aufbau principle D) Heisenberg uncertainty principle E) None of these
C
145) The element that corresponds to the electron configuration 1s22s22p63s23p64s13d5 is ________. A) titanium B) vanadium C) chromium D) manganese E) iron
C
146) Place the following elements in order of increasing atomic radius. P Ba Cl A) Ba < P < Cl B) P < Cl < Ba C) Cl < P < Ba D) Cl < Ba < P E) Ba < Cl < P
C
147) Place the following in order of decreasing IE1. Cs Mg Ar A) Cs > Mg > Ar B) Mg > Ar > Cs C) Ar > Mg > Cs D) Cs > Ar > Mg E) Mg > Cs > Ar
B
148) Choose the ground state electron configuration for Zn2⁺. A) [Ar]4s23d8 B) [Ar]3d10 C) [Ar]4s23d6 D) [Ar] E) [Ar]3d8
B
149) Place the following in order of increasing metallic character. Rb Cs K Na A) K < Cs < Na < Rb B) Na < K < Rb < Cs C) Cs < Rb < K < Na D) K < Cs < Rb < Na E) Na < Rb < Cs < K
D
15. What is the chemical formula for iron(III) sulfate? A) Fe3S B) Fe3SO4 C) Fe2S3 D) Fe2(SO4)3 E) FeS
B
150) Choose the paramagnetic species from below. A) Ti4⁺ B) Se C) Ar D) All of the above are paramagnetic. E) None of the above are paramagnetic.
B
16. What is the mass of 0.500 mol of dichlorodifluoromethane, CCl2F2? A) 4.14 × 10-3 g B) 60.5 g C) 121 g D) 242 g E) 72.5 g
B
17. A sample of pure lithium nitrate contains 10.1% lithium by mass. What is the % lithium by mass in a sample of pure lithium carbonate that has twice the mass of the first sample? A) 5.05% B) 10.1% C) 20.2% D) 40.4% E) 50.5%
A
18. Identify an ether. A) CH3CH2OCH2CH3 B) CH3CH2I C) CH3CH2NH2 D) CH3CH2CH2CH3 E) CH3COOH
C
19. According to the following reaction, how many moles of Fe(OH)2 can form from 175.0 mL of 0.227 M LiOH solution? Assume that there is excess FeCl2. FeCl2(aq) + 2 LiOH(aq) → Fe(OH)2(s) + 2 LiCl(aq) A) 3.97 × 10-2 moles B) 2.52 × 10-2 moles C) 1.99 × 10-2 moles D) 5.03 × 10-2 moles E) 6.49 × 10-2 moles
D
2. Determine the volume of an object that has a mass of 455.6 g and a density of 19.3 g/cm3. A) 87.9 mL B) 42.4 mL C) 18.5 mL D) 23.6 mL E) 31.2 mL
B
20. Identify acetic acid. A) strong electrolyte, weak acid B) weak electrolyte, weak acid C) strong electrolyte, strong acid D) weak electrolyte, strong acid E) nonelectrolyte
A
21. Give the net ionic equation for the reaction (if any) that occurs when aqueous solutions of H2SO4 and KOH are mixed. A) H+(aq) + OH-(aq) → H2O(l) B) 2 K+(aq) + SO42-(aq) → K2SO4(s) C) H+(aq) + OH-(aq) + 2 K+(aq) + SO42-(aq) → H2O(l) + K2SO4(s) D) H22+(aq) + OH-(aq) → H2(OH)2(l) E) No reaction occurs.
A
23. Determine the oxidation state of C in CO3-2. A) +4 B) +2 C) -2 D) -4 E) +6
D
24. What element is undergoing reduction (if any) in the following reaction? Zn(s) + 2 AgNO3(aq) → Zn(NO3)2(aq) + 2 Ag(s) A) Zn B) N C) O D) Ag E) This is not an oxidation-reduction reaction.
E
25. A syringe initially holds a sample of gas with a volume of 285 mL at 355 K and 1.88 atm. To what temperature must the gas in the syringe be heated/cooled in order to have a volume of 435 mL at 2.50 atm? A) 139 K B) 572 K C) 175 K D) 466 K E) 721 K
D
26. What mass of NO2 is contained in a 13.0 L tank at 4.58 atm and 385 K? A) 18.8 g B) 53.1 g C) 24.4 g D) 86.7 g E) 69.2 g
C
27. A 0.334 g sample of an unknown halogen occupies 109 mL at 398 K and 1.41 atm. What is the identity of the halogen? A) Br2 B) F2 C) Cl2 D) I2 E) Ge
D
28. The rate of effusion of oxygen to an unknown gas is 0.935. What is the other gas? A) Ne B) Ar C) F2 D) N2 E) He
D
29. This equation is used to calculate the properties of a gas under nonideal conditions. A) Charles's Law B) Avogadro's Law C) Boyle's Law D) van der Waals equation E) Dalton's Law
A
3. What answer should be reported, with the correct number of significant figures, for the following calculation? (433.621 - 333.9) × 11.900 A) 1.19 × 103 B) 1.187 × 103 C) 1.1868 × 103 D) 1.18680 × 103 E) 1.186799 × 103
E
30. The density of chlorine gas at 1.21 atm and 34.9°C is ________ g/L. A) 0.0479 B) 0.295 C) 0.423 D) 1.70 E) 3.39
A
31. A gas mixture contains CO, Ar and H2. What is the total pressure of the mixture, if the mole fraction of H2 is 0.350 and the pressure of H2 is 0.480 atm? A) 1.37 atm B) 0.168 atm C) 5.95 atm D) 0.729 atm E) 2.1 atm
A
32. Which of the following signs on q and w represent a system that is doing work on the surroundings, as well as gaining heat from the surroundings? A) q = +, w = - B) q = -, w = + C) q = +, w = + D) q = -, w = - E) None of these represent the system referenced above.
C
33. Determine the specific heat capacity of an alloy that requires 59.3 kJ to raise the temperature of 150.0 g alloy from 298 K to 398 K. A) 4.38 J/g°C B) 2.29 J/g°C C) 3.95 J/g°C D) 2.53 J/g°C E) 1.87 J/g°C
C
34. Use the ΔH°f and ΔH°rxn information provided to calculate ΔH°f for IF: ΔH°f (kJ/mol) IF7(g) + I2(g) → IF5(g) + 2 IF(g) ΔH°rxn = -89 kJ IF7(g) -941 IF5(g) -840 A) 101 kJ/mol B) -146 kJ/mol C) -190. kJ/mol D) -95 kJ/mol E) 24 kJ/mol
E
35. Which of the following substances (with specific heat capacity provided) would show the greatest temperature change upon absorbing 100.0 J of heat? A) 10.0 g Cu, CCu = 0.385 J/g°C B) 10.0 g H2O, CH2O = 4.18 J/g°C C) 10.0 g ethanol, Cethanol = 2.42 J/g°C D) 10.0 g Al, CAl = 0.903 J/g°C E) 10.0 g Pb, CPb= 0.128 J/g°C
A
36. For a particular process that is carried out at constant pressure, q = 145 kJ and w = -35 kJ. Therefore, A) ΔE = 110 kJ and ΔH = 145 kJ. B) ΔE = 145 kJ and ΔH = 110 kJ. C) ΔE = 145 kJ and ΔH = 180 kJ. D) ΔE = 180 kJ and ΔH = 145 kJ.
D
37. Which of the following processes is exothermic? A) the ionization of a potassium atom B) the breaking of a Br-Br bond C) the sublimation of dry ice (CO2(s)) D) the reaction associated with DH°f for an ionic compound E) All of the above processes are exothermic.
C
38. For n = 3, what are the possible sublevels? A) 0 B) 0, 1 C) 0, 1, 2 D) 0, 1, 2, 3 E) none of the above
A
39. Give the numbers for ml for an s orbital. A) 0 B) -1, 0, 1 C) 0, 1 D) 1 E) -2, -1, 0, +1, +2
D
4. Which of the following are examples of a chemical change? A) copper building materials develop a green patina over time B) a candle burns C) rubbing alcohol evaporates D) Both A and B are examples of chemical change. E) All of the above are examples of chemical change.
E
41. Place the following types of electromagnetic radiation in order of increasing wavelength. visible light gamma rays microwaves A) gamma rays< microwaves <visible light B) microwaves < visible light <gamma rays C) microwaves < gamma rays < visible light D) visible light < gamma rays <microwaves E) gamma rays < visible light < microwaves
A
42. Which of the following transitions (in a hydrogen atom) represent absorption of the smallest frequency photon? A) n = 5 to n = 6 B) n = 5 to n = 3 C) n = 2 to n = 1 D) n = 1 to n = 3 E) n = 1 to n = 2
A
43. Choose the valence orbital diagram that represents the ground state of Zn. A) B) C) D) E)
C
44. The element that corresponds to the electron configuration 1s22s22p63s23p64s13d5 is ________. A) titanium B) vanadium C) chromium D) manganese E) iron
C
45. Place the following in order of decreasing radius. Te2⁻ F⁻ O2⁻ A) F⁻ > O2⁻ > Te2⁻ B) F⁻ > Te2⁻ > O2⁻ C) Te2⁻ > O2⁻ > F⁻ D) Te2⁻ > F⁻ > O2⁻ E) O2⁻ > F⁻ > Te2⁻
D
46. Which reaction below represents the first ionization of O? A) O⁺(g) + e⁻ → O(g) B) O(g) + e⁻ → O⁻(g) C) O⁻(g) → O(g) + e⁻ D) O(g) → O⁺(g) + e⁻ E) O⁻(g) + e⁻ → O2⁻(g)
C
47. Place the following in order of decreasing IE1. Cs Mg Ar A) Cs > Mg > Ar B) Mg > Ar > Cs C) Ar > Mg > Cs D) Cs > Ar > Mg E) Mg > Cs > Ar
A
48. Give the set of four quantum numbers that could represent the last electron added (using the Aufbau principle) to the Cl atom. A) n = 3, l = 1, ml = 1, ms = +1/2 B) n = 3, l = 0, ml = 1, ms = -1/2 C) n = 3, l = 2, ml =1, ms = +1/2 D) n = 2, l = 1, ml = 1, ms = -1/2 E) n = 3, l =1, ml = 1, ms = -1/2
E
49. How many unpaired electrons are there in the ground state of Cl? A) 5 B) 4 C) 3 D) 2 E) 1
A
5. Because of the high heat and low humidity in the summer in Death Valley, California, a visitor requires about one quart of water for every two miles traveled on foot. Calculate the approximate number of liters required for a person to walk 10 kilometers in Death Valley. A) 2.9 L B) 12 L C) 30 L D) 47 L E) 3.9 L
E
50. Give the number of core electrons for Cl-. A) 22 B) 30 C) 17 D) 12 E) 10
A
51. Which of the following statements is TRUE? A) An ionic bond is much stronger than most covalent bonds. B) An ionic bond is formed through the sharing of electrons. C) Ionic compounds at room temperature typically conduct electricity. D) Once dissolved in water, ionic compounds rarely conduct electricity. E) None of the above are true.
C
52. Identify the strongest bond. A) single covalent bond B) double covalent bond C) triple covalent bond D) chemical bond between H and O within a water molecule E) all of the above bonds are the same strength
C ( look at answer sheet)
53. Choose the best Lewis structure for XeI2. A) B) C) D) E)
A
54. Give the number of valence electrons for SO42-. A) 32 B) 30 C) 34 D) 28 E) 36
C
55. Identify the compound with atoms that have an incomplete octet. A) ICl5 B) CO2 C) BF3 D) Cl2 E) CO
D
56. Place the following in order of increasing bond length. C-F C-S C-Cl A) C-S < C-Cl < C-F B) C-Cl < C-F < C-S C) C-F < C-S < C-Cl D) C-F < C-Cl < C-S E) C-S < C-F < C-Cl
E
57. Give the approximate bond angle for a molecule with an octahedral shape. A) 109.5° B) 180° C) 120° D) 105° E) 90°
C
58. Determine the electron geometry (eg) and molecular geometry (mg) of ICl2⁻. A) eg=tetrahedral, mg=bent B) eg=tetrahedral, mg=trigonal pyramidal C) eg=trigonal bipyramidal, mg=linear D) eg=trigonal bipyramidal, mg=trigonal planar E) eg=octahedral, mg=linear
D
59. Consider the molecule below. Determine the molecular geometry at each of the 2 labeled carbons. (look at answer Key) A) C1 = tetrahedral, C2 = linear B) C1 = trigonal planar, C2= bent C) C1 = bent, C2 = trigonal planar D) C1 = trigonal planar, C2 = tetrahedral E) C1 = trigonal pyramidal, C2 = see-saw
E
6. Which of the following statements is FALSE according to Dalton's Atomic Theory? A) Atoms combine in simple whole number ratios to form compounds. B) All atoms of chlorine have identical properties that distinguish them from other elements. C) One carbon atom will combine with one oxygen atom to form a molecule of carbon monoxide. D) Atoms of sodium do not change into another element during chemical reaction with chlorine. E) An atom of nitrogen can be broken down into smaller particles that will still have the unique
A
60. How many of the following molecules are polar? BrCl3 CS2 SiF4 SO3 A) 1 B) 2 C) 3 D) 4 E) 0
B
61. Identify the number of electron groups around a molecule with sp hybridization. A) 1 B) 2 C) 3 D) 4 E) 5
C
62. Give the electron geometry (eg), molecular geometry (mg), and hybridization for NH3. A) eg=tetrahedral, mg=trigonal planar, sp2 B) eg=trigonal pyramidal, mg=trigonal pyramidal, sp3 C) eg=tetrahedral, mg=trigonal pyramidal, sp3 D) eg=trigonal pyramidal, mg=tetrahedral, sp3 E) eg=trigonal planar, mg=trigonal planar, sp2
A
63. The orbital hybridization on the carbon atoms in C2H2 is A) sp. B) sp2. C) sp3. D) sp3d2. E) sp3d
B
64. Describe a sigma bond. A) side by side overlap of f orbitals B) end to end overlap of p orbitals C) s orbital overlapping with the side of a p orbital D) overlap of two f orbitals E) p orbital overlapping with a f orbital
C
65. A molecule containing a central atom with sp2 hybridization has a(n) ________ electron geometry. A) linear B) trigonal bipyramidal C) trigonal planar D) octahedral E) bent
A
66. Give the change in condition to go from a liquid to a gas. A) increase heat or reduce pressure B) increase heat or increase pressure C) cool or reduce pressure D) cool or increase pressure E) none of the above
E
67. Describe sweating in humans. A) It is an endothermic reaction. B) The sweat evaporates absorbing heat from the body. C) The skin is cooled. D) None of the above. E) All of the above.
A
68. Which substance below has the strongest intermolecular forces? A) A2X, ΔHvap= 39.6 kJ/mol B) BY2, ΔHvap= 26.7 kJ/mol C) C3X2, ΔHvap= 36.4 kJ/mol D) DX2, ΔHvap= 23.3 kJ/mol E) EY3, ΔHvap= 21.5 kJ/mol
C
69. Determine ΔHvap for a compound that has a measured vapor pressure of 24.3 torr at 273 K and 135 torr at 325 K. A) 41 kJ/mol B) 79 kJ/mol C) 24 kJ/mol D) 13 kJ/mol E) 34 kJ/mol
A
7. Calculate the atomic mass of element "X", if it has 2 naturally occurring isotopes with the following masses and natural abundances: X-45 44.8776 amu 32.88% X-47 46.9443 amu 67.12% A) 46.26 amu B) 45.91 amu C) 46.34 amu D) 46.84 amu E) 44.99 amu
A
70. Define sublimation. A) the phase transition from solid to gas B) the phase transition from gas to solid C) the phase transition from gas to liquid D) the phase transition from liquid to gas E) the phase transition from liquid to solid
C
71. What type of intermolecular force causes the dissolution of KF in water? A) hydrogen bonding B) dipole-dipole forces C) ion-dipole force D) dispersion forces E) none of the above
A
72. Choose the substance with the highest surface tension. A) HOCH2CH2OH B) CH2Br2 C) CH3CH2Cl D) CH3CH2OH E) CH3CH2CH3
B
73. Assign the appropriate labels to the phase diagram shown below. A) A = liquid, B = solid, C = gas, D = critical point B) A = gas, B = solid, C = liquid, D = triple point C) A = gas, B = liquid, C = solid, D = critical point D) A = solid, B = gas, C = liquid, D = triple point E) A = liquid, B = gas, C = solid, D = critical point
C
74. Give the major force between ethanol and water. A) dipole-dipole B) dispersion C) hydrogen bonding D) ion-ion E) ion-dipole
E
75. Which of the following compounds is most soluble in hexane (CH3CH2CH2CH2CH2CH2CH3)? A) methanol B) ethanol C) 1-propanol D) 1-butanol E) 1-pentanol
C
76. A solution is formed at room temperature by vigorously dissolving enough of the solid solute so that some solid remains at the bottom of the solution. Which statement below is TRUE? A) The solution is considered unsaturated. B) The solution is considered supersaturated. C) The solution is considered saturated. D) The solution would be considered unsaturated if it were cooled a bit to increase the solubility of the solid. E) None of the above are true.
C
77. How many moles of KF are contained in 347 g of water in a 0.175 m KF solution? A) 1.65 × 10-2 mol KF B) 5.04 × 10-2 mol KF C) 6.07 × 10-2 mol KF D) 3.22 × 10-2 mol KF E) 1.98 × 10-2 mol KF
C
78. Commercial grade HCl solutions are typically 39.0% (by mass) HCl in water. Determine the molarity of the HCl, if the solution has a density of 1.20 g/mL. A) 7.79 M B) 10.7 M C) 12.8 M D) 9.35 M E) 13.9 M
E
79. A solution is prepared by dissolving 49.3 g of KBr in enough water to form 473 mL of solution. Calculate the mass % of KBr in the solution if the density is 1.12 g/mL. A) 10.4% B) 8.57& C) 10.1% D) 11.7% E) 9.31%
C
8. How many xenon atoms are contained in 2.36 moles of xenon? A) 3.92 × 1024 xenon atoms B) 2.55 × 1023 xenon atoms C) 1.42 × 1024 xenon atoms D) 7.91 × 1025 xenon atoms E) 1.87 × 1026 xenon atoms
E
80. Identify the colligative property. A) vapor pressure lowering B) freezing point depression C) boiling point elevation D) osmotic pressure E) all of the above
B
81. Calculate the freezing point of a solution of 500.0 g of ethylene glycol (C2H6O2) dissolved in 500.0 g of water. Kf = 1.86°C/m and Kb = 0.512°C/m. A) 30.0°C B) -30.0°C C) 8.32°C D) -8.32°C E) 70.2°C
A
82. Dynamic equilibrium can be defined as A) rate of dissolution = rate of deposition B) rate of dissolution < rate of deposition C) rate of dissolution > rate of deposition D) rate of bubbling > rate of dissolving E) rate of condensing > rate of bubbling
E
83. A solution is prepared by adding 1.43 mol of KCl to 889 g of water. The concentration of KCl is ________ molal. A) 1.61 × 10-3 B) 622 C) 0.622 D) 1.27 × 103 E) 1.61
A
84. Place the following aqueous solutions of nonvolatile, nonionic compounds in order of decreasing osmotic pressure. I. 0.011 M sucrose II. 0.00095 M galactose III. 0.0060 M glycerin A) I > III > II B) I > II > III C) II > III > I D) III > I > II E) II > I > III
E
85) Which of the following statements is TRUE? A) A scientific law is fact. B) Once a theory is constructed, it is considered fact. C) A hypothesis is speculation that is difficult to test. D) An observation explains why nature does something. E) A scientific law summarizes a series of related observations.
A
86) Which of the following statements about the phases of matter is TRUE? A) In both solids and liquids, the atoms or molecules pack closely to one another. B) Solids are highly compressible. C) Gaseous substances have long-range repeating order. D) There is only one type of geometric arrangement that the atoms or molecules in any solid can adopt. E) Liquids have a large portion of empty volume between molecules.
C
87) Which of the following represents a chemical property of hydrogen gas? A) It is gaseous at room temperature. B) It is less dense than air. C) It reacts explosively with oxygen. D) It is colorless. E) It is tasteless.
D
88) The outside air temperature is 40.0°F, what is the temperature in Kelvin? (0oC=273.15 K) A) 313 K B) 377 K C) 283 K D) 278 K E)243K
B
89) How many cm3 are contained in 3.77 × 104 mm3? A) 3.77 × 104 cm3 B) 3.77 × 101 cm3 C) 3.77 × 10-10 cm3 D) 3.77 × 1020 cm3 E) 3.77 × 106 cm3
D
9. What is the identity of element Q if the ion Q2+ contains 10 electrons? A) C B) O C) Ne D) Mg E) Cl
B (look at answer sheet)
90) How many significant figures are there in the answer for the following problem? A) one B) two C) three D) four E) all of these because the zeros are ambiguous
A
91) Which of the following are examples of intensive properties? A) density B) volume C) mass D) None of these E) All of these
A
92) A student performs an experiment to determine the density of a sugar solution. She obtains the following results: 1.71 g/mL, 1.73 g/mL, 1.67 g/mL, 1.69 g/mL. If the actual value for the density of the sugar solution is 1.40 g/mL, which statement below best describes her results? A) Her results are precise, but not accurate. B) Her results are accurate, but not precise. C) Her results are both precise and accurate D) Her results are neither precise nor accurate. E) It isn't possible to determine with the information given.
D
93) When two elements form two different compounds, the masses of element B that combine with 1 g of element A can be expressed as a ratio of small whole numbers. Which law does this refer to? A) Law of Definite Proportions B) Law of the Conservation of Mass C) Law of Modern Atomic Theory D) Law of Multiple Proportions E) First Law of Thermodynamics
D (look at answer sheet)
94) What does "X" represent in the following symbol? X A) mercury B) chlorine C) scandium D) bromine E) selenium
B (look at answer sheet)
95) Determine the number of protons, neutrons and electrons in the following: X A) p+ = 18 n° = 18 e- = 22 B) p+ = 18 n° = 22 e- = 18 C) p+ = 22 n° = 18 e- = 18 D) p+ = 18 n° = 22 e- = 40 E) p+ = 40 n° = 22 e- = 18
A
96) Calculate the atomic mass of element "X", if it has 2 naturally occurring isotopes with the following masses and natural abundances: X-45 44.8776 amu 32.88% X-47 46.9443 amu 67.12% A) 46.26 amu B) 45.91 amu C) 46.34 amu D) 46.84 amu E) 44.99 amu
B
97) Which of the following statements is FALSE? A) Halogens are very reactive elements. B) The alkali metals are fairly unreactive. C) Sulfur is a main group element. D) Noble gases do not usually form ions. E) Zn is a transition metal.
A
98) How many electrons are in the ion, Cu2+? A) 27 B) 29 C) 31 D) 64 E) None of these
C
99) How many Fe atoms are contained in 787 g of Fe? A) 5.90 × 1025 Fe atoms B) 7.09 × 1021 Fe atoms C) 8.49 × 1024 Fe atoms D) 4.27 × 1022 Fe atoms E) 4.18 × 1024 Fe atoms
A
The titration of 25.0 mL of an unknown concentration H2SO4 solution requires 83.6 mL of 0.12 M LiOH solution. What is the concentration of the H2SO4 solution (in M)? A) 0.20 M B) 0.40 M C) 0.10 M D) 0.36 M E) 0.25 M
C
What is the maximum number of f orbitals that are possible? A) 1 B) 3 C) 7 D) 5 E) 9