Practice Final Exam Version 2

Chemistry 201 (Eikey)

Name___________________________________

Part 1: MULTIPLE CHOICE (48 points). Choose the one alternative that best completes the statement or answers the question. Record your answer on the scantron. Each question is worth 3 points. 1) A student weighed 30.00 µg of sulfur in the lab. This is the same mass as A) 3.000 × 10 4 ng. B) 3.000 × 10 -5 mg. C) 3.000 × 10 -8 g.

1)

D) 3.000 × 10 -5 kg.

2) A student performs an experiment to determine the density of a sugar solution. She obtains the following results: 1.79 g/mL, 1.81 g/mL, 1.80 g/mL, 1.81 g/mL. If the actual value for the density of the sugar solution is 1.80 g/mL, which statement below best describes her results? A) Her results are neither precise nor accurate. B) Her results are precise, but not accurate. C) Her results are accurate, but not precise. D) Her results are both precise and accurate E) It isn't possible to determine with the information given.

2)

3) A person weighs 77.1 kg. What is his weight in pounds? A) 35.0 pounds B) 154 pounds C) 170 pounds

3) D) 162 pounds

4) The density of air under ordinary conditions at 25°C is 1.19 g/L. How many kilograms of air are in a room that measures 11.0 ft × 11.0 ft and has an 10.0 ft ceiling? 1 in. = 2.54 cm (exactly); 1 L = 103 cm3 A) 0.152

B) 0.0962

C) 40.8

D) 4.08 × 104

4)

E) 3.66

5) Which of the following are examples of intensive properties? A) volume B) density C) mass D) None of the above are examples of intensive properties. E) All of the above are examples of intensive properties.

5)

6) Write the formula for strontium nitride. A) Sr(NO2)2

6)

B) Sr(NO3)2 C) Sr3N2 D) SrN E) Sr2N3

1

7) What species is represented by the following information? p+ = 17

n° = 18

A) Cl

7)

e- = 18 B) Kr

C) Ar+

D) Cl-

E) Ar

8) Write the name for Sn(SO4)2.

8)

A) tin (II) sulfite B) tin (IV) sulfate C) tin (I) sulfite D) tin sulfide E) tin (I) sulfate 9) Write the name for Ca3(PO4)2.

9)

A) tricalcium phosphorustetraoxide B) calcium (II) phosphite C) calcium (III) phosphite D) calcium phosphite E) calcium phosphate 10) Give the formula for sulfurous acid. A) H2SO4 B) HSO 4

10) C) H2SO3

D) HSO 3

11) The ion, IO2-, is named

11)

A) iodate ion. C) iodine(II) oxide ion.

B) iodite ion. D) iodine dioxide ion.

12) What mass of phosphorus pentafluoride, PF 5, has the same number of fluorine atoms as 25.0 g of oxygen difluoride, OF 2? A) 0.933 g B) 146 g C) 23.3 g D) 10.0 g

12)

13) How many millimoles of Ca(NO 3)2 contain 4.78 × 10 22 formula units of Ca(NO3)2? The molar

13)

mass of Ca(NO3)2 is 164.10 g/mol. A) 79.4 mmol Ca(NO3)2 B) 12.6 mmol Ca(NO3)2 C) 57.0 mmol Ca(NO3)2 D) 13.0 mmol Ca(NO3)2 E) 20.7 mmol Ca(NO3)2 14) Determine the molecular formula of a compound that is 49.48% carbon, 5.19% hydrogen, 28.85% nitrogen, and 16.48% oxygen. The molecular weight is 194.19 g/mol. A) C8H10N4O2 B) C4H5N2O C) C8H10N2O D) C8H12N4O2

2

14)

15) Combustion analysis of 63.8 mg of a C, H and O containing compound produced 145.0 mg of CO2 and 59.38 mg of H2O. What is the empirical formula for the compound? A) C5H2O

B) C3H6O

C) C3H7O

D) CHO

E) C6HO3

16) In an acid-base neutralization reaction 38.74 mL of 0.500 M potassium hydroxide reacts with 50.00 mL of sulfuric acid solution. What is the concentration of the H2SO4 solution? A) 0.194 M

B) 0.775 M

C) 0.387 M

15)

16)

D) 1.29 M

17) When 31.2 mL of 0.500 M AgNO3 is added to 25.0 mL of 0.300 M NH4Cl, how many grams of AgCl are formed? AgNO 3(aq) + NH4Cl(aq) → AgCl(s) + NH4NO3(aq) A) 2.24 g B) 1.07 g C) 3.31 g D) 6.44 g

17)

18) A 12.39 g sample of phosphorus reacts with 42.54 g of chlorine to form only phosphorus trichloride (PCl3). If it is the only product, what mass of PCl3 is formed?

18)

A) 30.15 g

B) 54.93 g

C) 79.71 g

D) 91.86 g

E) 140.01 g

19) How many molecules of sucrose (C12H22O11, molar mass = 342.30 g/mol) are contained in 14.3

19)

mL of 0.140 M sucrose solution? A) 8.29 × 10 22 molecules C12H22O11 B) 1.63 × 10 23 molecules C12H22O11 C) 5.90 × 10 24 molecules C12H22O11 D) 1.21 × 10 21 molecules C12H22O11 E) 6.15 × 10 22 molecules C12H22O11 20) Identify HCl. A) strong electrolyte, strong acid B) nonelectrolyte C) weak electrolyte, weak acid D) weak electrolyte, strong acid E) strong electrolyte, weak acid

20)

21) How many of the following compounds are soluble in water?

21)

Cu(OH)2 A) 1

B) 3

LiNO3

NH4Br

K2S

C) 4

D) 0

3

E) 2

22) Write a balanced equation to show the reaction of aqueous aluminum acetate with aqueous ammonium phosphate to form solid aluminum phosphate and aqueous ammonium acetate. A) Al(C2H3O2)2 (aq) + (NH4)2PO4 (aq) → AlPO4 (s) + 2 NH4C2H3O2 (aq)

22)

B) Al(C2H3O2)3 (aq) + (NH4)3PO4 (aq) → AlPO4 (s) + 3 NH4C2H3O2 (aq) C) Al(CO3)2 (aq) + (NH3)2PO4 (aq) → AlPO4 (s) + 2 NH3CO3 (aq) D) Al(C2H3O2)2 (aq) + (NH3)2PO4 (aq) → AlPO4 (s) + 2 NH3C2H3O2 (aq) E) Al(CO2)3 (aq) + (NH4)3PO3 (aq) → AlPO3 (s) + 3 NH4CO2 (aq) 23) Give the complete ionic equation for the reaction (if any) that occurs when aqueous solutions of lithium sulfide and copper (II) nitrate are mixed. A) Li+ (aq) + SO 42-(aq) + Cu +(aq) + NO 3-(aq) → CuS(s) + Li +(aq) + NO 3-(aq) B) 2 Li +(aq) + S2-(aq) + Cu2+(aq) + 2 NO3-(aq) → CuS(s) + 2 Li+(aq) + 2 NO3-(aq) C) 2 Li +(aq) + S2-(aq) + Cu 2+(aq) + 2 NO3-(aq) → Cu2+(aq) + S2-(aq) + 2 LiNO3(s) D) Li+ (aq) + S-(aq) + Cu +(aq) + NO 3-(aq) → CuS(s) + LiNO3(aq)

23)

E) No reaction occurs. 24) Give the net ionic equation for the reaction (if any) that occurs when aqueous solutions of Na2CO3 and HCl are mixed.

24)

A) 2 H+(aq) + CO32-(aq) → H2O(l) + CO 2(g) B) 2 Na+(aq) + CO 32-(aq) + 2 H+(aq) + 2 Cl-(aq) → H2CO3(s) + 2 Na+(aq) + 2 Cl-(aq) C) 2 H+(aq) + CO 32-(aq) → H2CO3(s) D) 2 Na+(aq) + CO32-(aq) + 2 H+(aq) + 2 Cl-(aq) → H2CO3(s) + 2 NaCl(aq) E) No reaction occurs. 25) Determine the oxidizing agent in the following reaction.

25)

Ni(s) + 2 AgClO4(aq) → Ni(ClO4)2(aq) + 2 Ag(s) A) Ni B) Ag C) Cl D) O E) This is not an oxidation-reduction reaction. 26) How many liters of hydrogen gas can be generated by reacting 9.25 grams of barium hydride with water at 20°C and 755 mm Hg pressure according to the chemical equation shown below? BaH2(s) + 2 H2O(l) → Ba(OH)2(aq) + 2 H2(g) A) 1.60 L B) 3.21 L C) 0.799 L D) 0.219 L

26)

27) A gas mixture contains CO, Ar and H2. What is the total pressure of the mixture, if the mole

27)

fraction of H2 is 0.35 and the pressure of H2 is 0.58 atm? A) 2.1 atm

B) 0.60 atm

C) 0.49 atm

4

D) 1.7 atm

E) 0.20 atm

28) A basketball is inflated to a pressure of 1.90 atm in a 24.0°C garage. What is the pressure of the basketball outside where the temperature is -1.00°C? A) 2.08 atm B) 1.80 atm C) 1.74 atm D) 2.00 atm

28)

29) The density of a gas is 1.43 g/L at STP. What is the gas? A) Cl2 B) O2 C) S

29) D) Ne

30) Zinc reacts with aqueous sulfuric acid to form hydrogen gas:

30)

Zn (s) + H2SO4 (aq) → ZnSO4 (aq) + H2 (g) In an experiment, 201 mL of wet H2 is collected over water at 27°C and a barometric pressure of 733 torr. The vapor pressure of water at 27°C is 26.74 torr. The partial pressure of hydrogen in this experiment is __________ atm. A) 1.00 B) 706 C) 0.929 D) 0.964 E) 760 31) Identify the gas particle that travels the slowest. A) CO B) O2 C) N2

31) D) Ne

E) H2

32) Which of the following statements is TRUE? A) For a given gas, the lower the temperature, the faster it will effuse. B) Particles of different masses have the same average speed at a given temperature. C) The larger a molecule, the faster it will effuse. D) At very high pressures, a gas will occupy a larger volume than predicted by the ideal gas law. E) None of the above statements are true.

32)

33) Which of the following processes is exothermic? A) the sublimation of dry ice (CO2(s))

33)

B) the ionization of a lithium atom C) the breaking of a Cl-Cl bond D) the reaction associated with DH°f for an ionic compound E) All of the above processes are exothermic. 34) Determine the final temperature of a gold nugget (mass = 376 g) that starts at 398 K and loses 4.85 kJ of heat to a snowbank when it is lost. The specific heat capacity of gold is 0.128 J/g°C. A) 297 K B) 133 K C) 377 K D) 398 K E) 187 K

34)

35) When 0.455 g of anthracene, C14H10, is combusted in a bomb calorimeter that has a water jacket containing 500.0 g of water, the temperature of the water increases by 8.63°C. Assuming that the specific heat of water is 4.18 J/(g ∙ °C), and that the heat absorption by the calorimeter is negligible, estimate the enthalpy of combustion per mole of anthracene. A) -7060 kJ/mol B) -8120 kJ/mol C) -39.7 kJ/mol D) +39.7 kJ/mol

35)

5

36) How much heat is absorbed when 45.00 g of C(s) reacts in the presence of excess SO2(g) to produce CS2(l) and CO(g) according to the following chemical equation? 5 C(s) + 2 SO2(g) → CS2(l) + 4 CO(g) ΔH° = 239.9 kJ A) 2158 kJ B) 179.8 kJ C) 898.5 kJ D) 239.9 kJ

36)

37) Choose the reaction that illustrates ΔH°f for NaHCO3. A) Na+(aq) + H2O (l) + CO 2 (g) → NaHCO 3 (s) B) Na(s) + H2(g) + C(s) + O 2(g) → NaHCO 3 (s)

37)

C) Na(s) + 2 H(g) + C(s) + 3 O(g) → NaHCO 3 (s) D) Na(s) + 1/2 H2(g) + C(s) + 3/2 O2(g) → NaHCO 3 (s) E) Na+(aq) + HCO3 -1 (aq) → NaHCO 3 (s) 38) Use the standard reaction enthalpies given below to determine ΔH°rxn for the following

38)

reaction: 4 SO 3(g) → 4 S(s) + 6 O2(g)

ΔH°rxn = ?

SO2(g) → S(s) + O2(g)

ΔH°rxn = +296.8 kJ

2 SO 2(g) + O 2(g) → 2 SO3(g)

ΔH°rxn = -197.8 kJ

Given:

A) -293.0 kJ

B) -494.6 kJ

C) -692.4 kJ

D) 1583 kJ

E) -791.4 kJ

39) Use the ΔH°f information provided to calculate ΔH°rxn for the following: SO2Cl2 (g) + 2 H2O(l) → 2 HCl(g) + H2SO4(l)

ΔH°f (kJ/mol) SO2Cl2(g) H2O(l)

-364

HCl(g) H2SO4(l)

-92 -814

39) ΔH°rxn = ?

-286

A) +161 kJ

B) -422 kJ

C) -256 kJ

D) +800. kJ

E) -62 kJ

40) Each of the following sets of quantum numbers is supposed to specify an orbital. Which of the following sets of quantum numbers contains an error? A) n = 2, l = 1 , ml = -1

40)

B) n = 4, l = 2, ml =0 C) n = 3, l =3 , ml = -2 D) n = 3, l = 0, ml =0 E) n = 1, l = 0, ml =0 41) How many different values of ml are possible in the 4f sublevel? A) 3

B) 1

C) 7

41) D) 2

6

E) 5

42) Calculate the energy of the orange light emitted, per photon, by a neon sign with a frequency of 4.89 × 10 14 Hz.

42)

A) 1.63 × 10 -19 J B) 3.09 × 10 -19 J C) 6.14 × 10 -19 J D) 3.24 × 10 -19 J E) 5.11 × 10 -19 J 43) Determine the shortest frequency of light required to remove an electron from a sample of Ti metal, if the binding energy of titanium is 3.14 × 10 3 kJ/mol.

43)

A) 7.87 x 1015 Hz B) 1.27 x 1015 Hz C) 2.11 x 1015 Hz D) 4.74 x 1015 Hz E) 6.19 x 1015 Hz 44) Identify the correct values for a 4f sublevel. A) n = 3, l = 1, ml = 0

44)

B) n = 4, l = 3, ml = -2 C) n = 2, l = 1, ml = -2 D) n = 1, l = 0, ml = 0 E) n = 2, l = 0, ml = 0 45) Determine the end (final) value of n in a hydrogen atom transition, if the electron starts in n = 1 and the atom absorbs a photon of light with an energy of 2.044 × 10 -18 J. A) 5

B) 4

C) 3

D) 2

45)

E) 6

46) Determine the mass of a ball with a wavelength of 3.45 x 10-34 m and a velocity of 6.55 m/s. A) 12.6 g B) 346 g C) 293 g D) 3.41 g E) 0.293 g

46)

47) Which of the following occur as the energy of a photon increases? A) the frequency decreases. B) the wavelength increases C) the speed increases. D) the wavelength gets shorter. E) None of the above occur as the energy of a photon increases.

47)

48) Which of the following transitions (in a hydrogen atom) represent emission of the longest wavelength photon? A) n = 3 to n = 4 B) n = 3 to n = 1 C) n = 5 to n = 4 D) n = 1 to n = 2 E) n = 4 to n = 2

48)

7

49) For n = 3, what are the possible sublevels? A) 0, 1, 2 B) 0

49) C) 0, 1

D) 0, 1,2, 3

50) How many of the following elements have 1 unpaired electron in the ground state? B

Al

S

A) 2

Cl B) 1

C) 3

D) 4

51) Place the following elements in order of increasing atomic radius. P

Ba

50)

51)

Cl

A) Ba < P < Cl B) P < Cl < Ba C) Ba < Cl < P D) Cl < Ba < P E) Cl < P < Ba 52) Which of the following represents the change in electronic configuration that is associated with the first ionization energy of strontium? A) [Kr]5s15p1 → [Kr]5s1 + e B) [Kr]5s2 → [Kr]5s1 + e C) [Kr]5s2 → [Kr]5s15p1 D) [Kr]5s2 + e - → [Kr]5s25p1

52)

53) How many of the following elements have 2 unpaired electrons in the ground state?

53)

C

O

Ti

A) 4

Si B) 1

C) 3

54) Of the following, which element has the highest first ionization energy? A) beryllium B) boron C) lithium 55) The complete electron configuration of gallium, element 31, is __________. A) 1s22s22p63s23p63d104s24p1 B) 1s42s42p103s43p9

D) 2 54) D) hydrogen 55)

C) 1s42s42p83s43p84s3 D) 1s42s42p63s43p64s43d3 E) 1s22s22p103s23p104s23d3 56) Give the ground state electron configuration for Pb. A) [Xe]6s26p2 B) [Xe]6s25d106p2 C) [Xe]6s24f145d106s26p2 D) [Xe]6s25f146d106p2 E) [Xe]6s24f145d106p2

8

56)

57) How many unpaired electrons are present in the ground state Ge atom? A) 0 B) 1 C) 2 D) 3

57) E) 4

58) Choose the valence orbital diagram that represents the ground state of Zn. A)

58)

B)

C)

D)

E)

59) Which of the following contains an atom that does not obey the octet rule? A) ClF 5 B) CsI C) ClF

59) D) SnO2

60) 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 = 2, ml =1 , ms = + 1 2

60)

B) n = 3, l =2 , ml = 1, ms = - 1 2 C) n = 3, l = 0, ml = 1, ms = - 1 2 D) n = 3, l = 1, ml = 1, ms = + 1 2 E) n = 2, l = 1, ml = 1, ms = - 1 2 61) The iodine atom in I2 would be expected to have a A) partial charge δ-. C) partial charge δ+.

61) B) charge of 0. D) charge of 1-.

9

62) List the following compounds in decreasing electronegativity difference. Cl2

HCl

62)

NaCl

A) NaCl > Cl2 > HCl C) NaCl > HCl > Cl2

B) Cl 2 > HCl > NaCl D) HCl > NaCl > Cl2

63) Place the following in order of decreasing bond length. H-F

H-I

63)

H-Br

A) H-I > H-F > H-Br B) H-Br > H-F > H-I C) H-I > H-Br > H-F D) H-F > H-Br > H-I E) H-F > H-I > H-Br 64) The hybrid orbital set used by the central atom in NCl3 is __________. A) sp3

B) sp

C) sp3d2

D) sp3d

65) List the number of sigma bonds and pi bonds in a triple bond. A) 1 sigma, 1 pi B) 1 sigma, 2 pi C) 2 sigma, 2 pi

64) E) sp2 65) D) 2 sigma, 1 pi

66) Draw the Lewis structure for SO 3. What is the hybridization on the S atom? A) sp

B) sp2

C) sp3d

D) sp3d2

66) E) sp3

67) Draw the Lewis structure for OF2. What is the hybridization on the O atom? A) sp2

B) sp3

C) sp3d

68) Choose the substance with the highest viscosity. A) BeCl2 B) AsCl5 C) OCl2

D) sp

67) E) sp3d2 68)

D) ICl2

E) SbCl3

69) Give the change in condition to go from a liquid to a gas. A) increase heat or reduce pressure B) cool or reduce pressure C) increase heat or increase pressure D) cool or increase pressure E) none of the above

69)

70) What mass of NaCl is contained in 24.88 g of a 15.00% by mass solution of NaCl in water? A) 1.50 g B) 20.00 g C) 3.73 g D) 12.44 g E) 21.15 g

70)

71) Calculate the molality of a solution that is prepared by mixing 25.5 mL of CH3OH (d= 0.792

71)

g/mL) and 387 mL of CH3CH2CH2OH (d= 0.811 g/mL). A) 2.01 m

B) 1.57 m

C) 0.812 m

10

D) 0.630 m

E) 4.98 m

72) How many significant figures are in 3.408 x 104 m? A) 8 B) 4 C) 7

D) 3

E) 5

73) What decimal power does the abbreviation f represent? A) 1 × 10 -15 B) 1 × 10 3 C) 1 × 10 -12

D) 1 × 10 6

E) 1 × 10 -1

72)

73)

74) Which of the following statements is FALSE according to Dalton's Atomic Theory? A) All atoms of chlorine have identical properties that distinguish them from other elements. B) An atom of nitrogen can be broken down into smaller particles that will still have the unique properties of nitrogen. 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) Atoms combine in simple whole number ratios to form compounds.

74)

75) What is the concentration of magnesium ions in a 0.125 M Mg(NO 3)2 solution?

75)

A) 0.160 M

B) 0.125 M

C) 0.375 M

D) 0.250 M

E) 0.0625 M

76) In a container containing CO, H2, and O2, what is the mole fraction of CO if the H2 mole fraction

76)

is 0.22 and the O 2 mole fraction is 0.58? A) 0.10

B) 0.30

C) 0.20

D) 0.50

77) Which of the following compounds will behave LEAST like an ideal gas at low temperatures? A) He B) SO2 C) N2 D) H2 E) F 2

77)

78) How many of the following molecules are polar?

78)

PCl5 A) 4

COS

XeO3 B) 2

SeBr2 C) 1

D) 3

E) 0

79) What is the strongest type of intermolecular force present in H2?

79)

A) dispersion B) hydrogen bonding C) dipole-dipole D) ion-dipole E) none of the above 80) If a solution has a temperature of 355 K, what is its temperature in degrees celsius? A) 82°C B) 165°C C) 628°C D) 179°C E) 279°C

80)

81) How many milliliters of 0.132 M HClO4 solution are needed to neutralize 50.00 mL of 0.0789 M NaOH? A) 29.9 B) 0.0120 C) 0.0335 D) 0.521 E) 83.7

81)

11

82) A syringe contains 0.65 moles of He gas that occupy 750.0 mL. What volume (in L) of gas will the syringe hold if 0.35 moles of Ne is added? A) 4.9 L B) 1.2 L C) 2.1 L D) 0.87 L E) 1.9 L

82)

83) Choose the bond below that is the strongest. A) C=O B) C-I C) I-I

83) D) C≡N

E) C-F

84) Place the following in order of increasing F-A-F bond angle, where A represents the central atom in each molecule. PF 3

OF 2

84)

PF 4⁺

A) PF 4⁺ < OF 2 < PF 3 B) PF 4⁺ < PF 3 < OF 2 C) OF 2 < PF 3 < PF 4⁺ D) OF 2 < PF 4⁺ < PF 3 E) PF 3 < OF 2 < PF 4⁺ 85) A metal crystallizes in a face centered cubic structure and has a density of 11.9 g/cm3. If the radius of the metal atom is 138 pm, what is the identity of the metal? A) Cr B) Pd C) Fe D) Mn E) At

85)

86) The Henry's Law constant of methyl bromide, CH3Br, is k = 0.159 mol/(L ∙ atm) at 25°C. What is the solubility of methyl bromide in water at 25°C and at a partial pressure of 270. mm Hg? A) 0.0565 mol/L B) 42.9 mol/L C) 0.355 mol/L D) 0.448 mol/L

86)

87) A student dissolved 4.00 g of Co(NO 3)2 in enough water to make 100. mL of stock solution. He took 4.00 mL of the stock solution and then diluted it with water to give 275. mL of a final solution. How many grams of NO3- ion are there in the final solution? A) 0.108 g B) 0.0197 g C) 0.0542 g D) 0.0394 g

87)

88) How much heat is released when 105 g of steam at 100.0°C is cooled to ice at -15.0°C? The enthalpy of vaporization of water is 40.67 kJ/mol, the enthalpy of fusion for water is 6.01 kJ/mol, the molar heat capacity of liquid water is 75.4 J/(mol ∙ °C), and the molar heat capacity of ice is 36.4 J/(mol ∙ °C). A) 347 kJ B) 54.8 kJ C) 319 kJ D) 273 kJ

88)

12

89) Choose the best Lewis structure for PO 43⁻.

89)

A)

B)

C)

D)

E)

90) Draw the best Lewis structure for the free radical, NO2. What is the formal charge on the N? A) 0

B) -2

C) -1

D) +2

E) +1

91) Use the bond energies provided to estimate ΔH°rxn for the reaction below.

Bond C-H C-O C=O O=O O-H A) -91 kJ

CH3OH(l) + 2 O2(g) → CO2(g) + 2 H2O(g) Bond Energy (kJ/mol) 414 360 799 498 464 B) +473 kJ

C) +206 kJ

13

91)

ΔH°rxn = ?

D) -486 kJ

90)

E) -392 kJ

92) Choose the compound below that should have the lowest melting point according to the ionic bonding model. A) KBr B) CsI C) RbI D) NaCl E) LiF

92)

93) Place the following in order of decreasing magnitude of lattice energy.

93)

K2O

Rb 2S

Li2O

A) Rb 2S > K2O > Li2O B) Rb 2S > Li2O > K2O C) K2O > Li2O > Rb 2S D) Li2O > Rb 2S > K2O E) Li2O > K2O > Rb 2S 94) Use the data given below to construct a Born-Haber cycle to determine the heat of formation of KCl. DH°(kJ) 89 418 244

K(s) → K(g) K(g) → K⁺(g) + e⁻ Cl2(g) → 2 Cl(g) Cl(g) + e⁻ → Cl⁻(g) KCl(s) → K⁺(g) + Cl⁻(g) A) -437 kJ

94)

B) +158 kJ

-349 717 C) -1119 kJ

D) -997 kJ

E) +631 kJ

95) Consider the molecule below. Determine the molecular geometry at each of the 2 labeled carbons.

95)

A) C1 = tetrahedral, C2 = linear B) C1 = trigonal pyramidal, C2 = see-saw C) C1 = trigonal planar, C2 = tetrahedral D) C1 = bent, C2 = trigonal planar E) C1 = trigonal planar, C2= bent 96) Draw the Lewis structure for SF 6. What is the hybridization on the S atom? A) sp3

B) sp3d2

C) sp2

14

D) sp3d

96) E) sp

97) How many of the following molecules have sp hybridization on the central atom? C2Cl2 A) 3

CO2

O3

97)

H2O

B) 2

C) 0

D) 1

E) 4

98) Use the molecular orbital diagram shown to determine which of the following are paramagnetic.

A) F 22⁺ B) Ne22⁺ C) O22⁻ D) O22⁺ E) None of the above are paramagnetic.

15

98)

99) Place the following substances in order of increasing boiling point. Ne

Cl2

99)

O2

A) Cl2 < Ne < O2 B) Ne < O2 < Cl 2 C) Cl2 < O2 < Ne D) O2 < Cl2 < Ne E) Ne < Cl2 < O 2 100) Place the following substances in order of decreasing vapor pressure at a given temperature. BeF 2

CH3OH

100)

OF 2

A) BeF 2 > OF 2 > CH3OH B) OF 2 > CH3OH > BeF2 C) OF 2 > BeF 2 > CH3OH D) BeF 2 > CH3OH > OF2 E) CH3OH > OF2 > BeF 2

101) Based on the figure above, the boiling point of diethyl ether under an external pressure of 1.32 atm is __________°C. A) 10 B) 20 C) 40 D) 0 E) 30

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101)

102) Determine ΔHvap for a compound that has a measured vapor pressure of 24.3 torr at 273 K and

102)

135 torr at 325 K. A) 24 kJ/mol B) 41 kJ/mol C) 79 kJ/mol D) 13 kJ/mol E) 34 kJ/mol 103) Consider the phase diagram shown. Choose the statement below that is TRUE.

103)

A) The line separating the solid and liquid phases represents the ΔHvap . B) At 10 atm of pressure, there is no temperature where the liquid phase of this substance would exist. C) The solid phase of this substance is higher in density than the liquid phase. D) The triple point of this substance occurs at a temperature of 31°C. E) None of the above are true. 104) Choose the statement below that is TRUE. A) A solution will form between two substances if the solute-solute interactions are strong enough to overcome the solvent-solvent interactions. B) A solution will form between two substances if the solute-solvent interactions are small enough to be overcome by the solute-solute and solvent-solvent interactions. C) A solution will form between two substances if the solute-solvent interactions are of comparable strength to the solute-solute and solvent-solvent interactions. D) A solution will form between two substances only if the solvent-solvent interactions are weak enough to overcome the solute-solvent interactions. E) None of the above are true.

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104)

105) Which of the following compounds will be most soluble in pentane (C5H12)?

105)

A) acetic acid (CH3CO2H) B) benzene (C6H6) C) ethyl methyl ketone (CH3CH2COCH3) D) pentanol (CH3CH2CH2CH2CH2OH) E) None of these compounds should be soluble in pentane. 106) Which of the following ions should have the most exothermic ΔHhydration? A) Ca2⁺

B) Sr2⁺

C) Mg2⁺

D) Na⁺

106) E) Al 3⁺

107) Place the following aqueous solutions of nonvolatile, nonionic compounds in order of decreasing osmotic pressure. I. 0.011 M sucrose

II. 0.00095 M glucose

107)

III. 0.0060 M glycerin

A) III > I > II B) I > II > III C) II > I > III D) I > III > II E) II > III > I 108) A solution is 0.0433 m LiF. What is the molarity of the solution if the density is 1.10 g/mL? A) 0.0441 M B) 0.0476 M C) 0.0390 M D) 0.0519 M E) 0.0417 M

108)

109) Determine the vapor pressure of a solution at 55°C that contains 34.2 g NaCl in 375 mL of water. The vapor pressure of pure water at 55°C is 118.1 torr. A) 87.1 torr B) 115 torr C) 92.8 torr D) 112 torr E) 108 torr

109)

110) Calculate the freezing point of a solution of 500.0 g of ethylene glycol (C2H6O2) dissolved in

110)

500.0 g of water. Kf = 1.86°C/m and Kb = 0.512°C/m. A) 30.0°C B) 8.32°C C) -30.0°C

D) -8.32°C

E) 70.2°C

111) A 150.0 mL sample of an aqueous solution at 25°C contains 15.2 mg of an unknown nonelectrolyte compound. If the solution has an osmotic pressure of 8.44 torr, what is the molar mass of the unknown compound? A) 294 g/mol B) 223 g/mol C) 195 g/mol D) 341 g/mol E) 448 g/mol

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111)

Answer Key Testname: CHEM 201 PRACTICE FIN

1) A 2) D 3) C 4) C 5) B 6) C 7) D 8) B 9) E 10) C 11) B 12) C 13) A 14) A 15) B 16) A 17) B 18) B 19) D 20) A 21) B 22) B 23) B 24) A 25) B 26) B 27) D 28) C 29) B 30) C 31) A 32) D 33) D 34) A 35) A 36) B 37) D 38) D 39) E 40) C 41) C 42) D 43) A 44) B 45) B 46) C 47) D 48) C 49) A 19

Answer Key Testname: CHEM 201 PRACTICE FIN

50) C 51) E 52) B 53) A 54) B 55) A 56) E 57) C 58) A 59) A 60) D 61) B 62) C 63) C 64) A 65) B 66) B 67) B 68) A 69) A 70) C 71) A 72) B 73) A 74) B 75) B 76) C 77) B 78) D 79) A 80) A 81) A 82) B 83) D 84) C 85) B 86) A 87) A 88) C 89) B 90) E 91) E 92) B 93) E 94) A 95) C 96) B 97) B 98) A 20

Answer Key Testname: CHEM 201 PRACTICE FIN

99) B 100) B 101) C 102) A 103) C 104) C 105) B 106) E 107) D 108) B 109) B 110) C 111) B

21