AP Chemistry



AP Chemistry Name ________________________

Fall Semester Practice Multiple Choice Period _____

Multiple Choice: Briefly show/explain why the multiple choice answer is correct in the space provided (no calculator).

1. Copper has two naturally occurring isotopes, 63Cu and 65Cu. What is the abundance of 65Cu if the average atomic mass of copper is 63.5?

(A) 90% (B) 70% (C) 50% (D) 25%

2. Which of the following particles is emitted by an atom of 39Ca when it decays to produce an atom of 39K?

(A) 10n (B) 11H (C) β- (D) β+

3. After 195 days, a 10.0 g sample of pure 95Zr has decayed to the extent that only 1.25 g of the original 95Zr remains. The half-life of 95Zr is closest to

(A) 195 days (B) 98 days

(C) 65 days (D) 49 days

Questions 4-5 The diagram shows the energy levels (in eV) for hydrogen gas.

[pic]

4. What is the energy, in eV, of a photon emitted by an electron as it moves from the n = 6 to the n = 2 energy level in a hydrogen atom.

(A) 0.38 eV (B) 3.02 eV (C) 3.40 eV (D) 13.60 eV

5. A photon having energy of 9.4 eV strikes a hydrogen atom in the ground state. Why is the photon not absorbed by the hydrogen atom?

(A) The atom's orbital electron is moving too fast

(B) The photon striking the atom is moving too fast.

(C) The photon's energy is too small.

(D) The photon is being repelled by electrostatic force.

Questions 6-8

(A) Heisenberg uncertainty principle

(B) Pauli exclusion principle

(C) Hund's rule (principle of maximum multiplicity)

(D) Shielding effect

6. Can be used to predict that a gaseous carbon atom in its ground state is paramagnetic.

7. Indicates that an atomic orbital can hold no more than two electrons.

8. Predicts that it is impossible to determine simultaneously the exact position and the exact velocity of an electron.

9. Which set of quantum numbers (n, l, ml, ms) best describes the valence electron of highest energy in a ground-state gallium atom (Z = 31)?

(A) 4,0,0,½ (B) 4,0,1,½ (C) 4,1,1,½ (D) 4,1,2,½

Questions 10-12 refer to neutral atoms for which the atomic orbitals are represented below.

(A) 1s(↑↓)

(B) 1s(↑↓) 2s(↑ ) 2p(↑ )( )( )

(C) 1s(↑↓) 2s(↑↓) 2p(↑ )(↑ )(↑ )

(D) [Ar] 4s(↑↓) 3d(↑↓)(↑↓)(↑↓)(↑ )(↑ )

10. Is in an excited state

11. Has exactly five valence electrons

12. Has the highest first ionization energy

13. The first seven ionization energies of element X are shown in the table below.

|Ionization Energy (kJ•mol-1) |

|1st |2nd |3rd |4th |5th |6th |7th |

|787 |1,580 |3,200 |4,400 |16,000 |20,000 |24,000 |

On the basis of these data, element X is most likely a member of which of the following groups of elements?

(A) Alkaline earth metals (B) Boron group

(C) Carbon group (D) Nitrogen group

14. In which of the following are the chemical species correctly ordered from smallest radius to largest radius?

(A) B < C < N (B) At < Xe < Kr

(C) CI < S < S2- (D) Na < Na+ < K

15. Of the following elements, which would be expected to have chemical properties most similar to those of sulfur, S?

(A) Br (B) CI (C) P (D) Se

16. Which pair of ions should have the highest lattice energy?

(A) Na+ and Br- (B) Li+ and F-

(C) Cs+ and F- (D) Li+ and O2-

17. Which molecule has the weakest bond?

(A) CO (B) O2 (C) Cl2 (D) N2

18. Which pair of atoms should form the most polar bond?

(A) F and B (B) C and O

(C) F and O (D) N and F

19. Which species has a valid non-octet Lewis structure?

(A) GeCl4 (B) SiF4 (C) NH4+ (D) SeCl4

20. For which molecule are resonance structures necessary to describe the bonding satisfactorily?

(A) H2S (B) SO2 (C) CO2 (D) OF2

21. The Lewis structure for SeS2 with zero formal charge has a total of

(A) 2 bonding pairs and 7 nonbonding pairs of electrons.

(B) 2 bonding pairs and 6 nonbonding pairs of electrons.

(C) 3 bonding pairs and 6 nonbonding pairs of electrons.

(D) 4 bonding pairs and 5 nonbonding pairs of electrons.

Questions 22-24 refer to the following molecules.

(A) CO (B) CH4 (C) HF (D) PH3

22. Contains two π-bonds

23. Has the highest dipole moment (most polar)

24. Has a molecular geometry that is trigonal pyramidal

25. Which of the following molecules contains bonds that have a bond order of 1.5?

(A) N2 (B) O3 (C) NH3 (D) CO2

26. CCl4, CO2, PCl3, PCl5, SF6

Which does NOT describe any of the molecules above?

(A) Linear (B) Octahedral

(C) Square planar (D) Tetrahedral

27. According to the VSEPR model, the progressive decrease in the bond angles in the series of molecules CH4, NH3, and H2O is best accounted for by the

(A) increasing strength of the bonds

(B) decreasing size of the central atom

(C) increasing electronegativity of the central atom

(D) increasing number of unshared pairs of electrons

28. Which of the following is a formula for an acid?

(A) CH3–CO–CH3 (B) CH3–CH2–COOH

(C) CH3–CH2–CH2OH (D) CH3–CH2–O–CH3

29. Which is NOT a structural isomer of 2-methylpentane?

(A) hexane (B) 3-methylpentane

(C) 2,2-dimethylbutane (D) 4-methylpentane

30. Types of hybridization exhibited by carbon atoms in a molecule of propyne, CH3CCH, include which of the following?

I. sp II. sp2 III. sp3

(A) I only (B) II only (C) III only (D) I and III only

31. Which of the following best explains why the normal boiling point of CCI4(I) (350 K) is higher than the normal boiling point of CF4(I) (145 K)?

(A) The C-CI bonds in CCI4 are less polar than the C-F bonds in CF4.

(B) The C-CI bonds in CCI4 are weaker than the C-F bonds in CF4.

(C) The mass of the CCI4 molecule is greater than that of the CF4 molecule.

(D) The electron cloud of the CCI4 molecule is more polarizable than that of the CF4 molecule.

32. Which of the following substances exhibits significant hydrogen bonding in the liquid state?

(A) CH2F2 (B) N2H4

(C) CH3OCH3 (D) C2H4

Questions 33-34 refer to a various points in time during an experiment conducted at 1.0 atm. Heat is added at a constant rate to a sample of a pure substance that is solid at time to. The graph below shows the temperature of the sample as a function of time.

(A) t1 (B) t2 (C) t3 (D) t5

33. Time when the average distance between particles is greatest

34. Time when the temperature of the substance is between its melting point and its boiling point

35. Heat energy is added slowly to a pure solid covalent compound at its melting point. About half of the solid melts to become a liquid. Which of the following must be true about this process?

(A) Covalent bonds are broken as the solid melts.

(B) The temperature of the solid/liquid mixture remains the same while heat is being added.

(C) The volume of the compound increases as the solid melts to become a liquid.

(D) The average kinetic energy of the molecules becomes greater as the molecules leave the solid state and enter the liquid state.

36. Of the following gases, which has the greatest average molecular speed at 298 K?

(A) Cl2 (B) NO (C) H2S (D) HCN

37. At approximately what temperature will 40. g of argon gas at 2.0 atm occupy a volume of 22.4 L?

(A) 600 K (B) 550 K (C) 270 K (D) 140 K

38. Three gases in the amounts shown in the table are added to a previously evacuated rigid tank.

|Gas |Ar |CH4 |N2 |

|Amount |0.35 mol |0.90 mol |0.25 mol |

If the total pressure in the tank is 3.0 atm at 25oC, the partial pressure of N2(g) in the tank is closest to

(A) 0.75 atm (B) 0.50 atm

(C) 0.33 atm (D) 0.25 atm

39. At which of the following temperatures and pressures would a real gas be most likely to deviate from ideal behavior?

Temperature (K) Pressure (atm)

(A) 100 50

(B) 200 5

(C) 300 0.01

(D) 500 0.01

40. In which process are covalent bonds broken?

(A) Solid silver melts.

(B) Solid potassium chloride melts.

(C) Solid carbon (graphite) sublimes.

(D) Solid iodine sublimes.

41. A closed rigid container contains distilled water and N2(g) at equilibrium. Actions that would increase the concentration of N2(g) in water include which of the following?

I. Shaking the container vigorously

II. Raising the temperature of the water

III. Injecting more N2(g) into the container

(A) I only (B) II only (C) III only (D) I and II only

42. What is the mole fraction of ethanol in a 6 molal aqueous solution?

(A) 0.006 (B) 0.1 (C) 0.08 (D) 0.2

43. What additional information is needed to determine the molality of a 1.0-M glucose (C6H12O6) solution?

(A) Volume of the solution

(B) Temperature of the solution

(C) Solubility of glucose in water

(D) Density of the solution

44. A solution of toluene (MM = 90 g) in benzene (MM = 80 g) is prepared. The mole fraction of toluene in the solution is 0.2. What is the molality of the solution?

(A) 0.2 (B) 0.5 (C) 2 (D) 3

45. Which of the following aqueous solutions has the highest boiling point at 1.0 atm?

(A) 0.20 M CaCl2 (B) 0.25 M Na2SO4

(C) 0.30 M NaCl (D) 0.40 M C6H12O6

46. _CH3OCH3(g) + _O2(g) → _CO2(g) + _ H2O(g)

When the equation above is balanced using the lowest whole-number coefficients, the coefficient for O2(g) is

(A) 6 (B) 4 (C) 3 (D) 2

47. What mass of KBr (MM = 119 g•mol-1) is required to make 250. mL of a 0.400 M KBr solution?

(A) 0.595 g (B) 1.19 g (C) 2.50 g (D) 11.9 g

48. Na2CO3(s) + 2 HCl(aq) → 2 NaCl(aq) + CO2(g) + H2O(l)

In a laboratory, a student wants to quantitatively collect the CO2(g) generated by adding Na2CO3(s) to 2.5 M HCI(aq). The student sets up the apparatus to collect the CO2 gas over water. The volume of collected gas is much less than the expected volume because CO2 gas

(A) is soluble in water

(B) is produced at a low pressure

(C) is more dense than water vapor

(D) has a larger molar mass than that of N2 gas, the major component of air

49. Which of the following would produce the LEAST mass of CO2 if completely burned in excess oxygen gas?

(A) 10.0 g CH4 (B) 10.0 g CH3OH

(C) 10.0 g C2H4 (D) 10.0 g C2H6

50. A solution of RbCl (MM = 121 g•mol-1) contains 11.0 % RbCl by mass. From the following list, what is needed to determine the molarity of RbCl in the solution?

I. Mass of the sample

II. Volume of the sample

III. Temperature of the sample

(A) I only (B) II only (C) III only (D) I and II only

51. CS2(l) + 3 O2(g) → CO2(g) + 2 SO2(g)

When 0.60 mol of CS2(l) reacts as completely as possible with 1.5 mol of O2(g) according to the equation above, the total number of moles of reaction products is

(A) 2.4 mol (B) 2.1 mol (C) 1.8 mol (D) 1.5 mol

52. What is the empirical formula of a hydrocarbon that is 10 % hydrogen by mass?

(A) CH3 (B) C2H5 (C) C3H4 (D) C4H9

53. By mixing only 0.15 M HCl and 0.25 M HCl, it is possible to create all of the following solutions EXCEPT

(A) 0.21 M (B) 0.18 M (C) 0.16 M (D) 0.14 M

54. 8 H2(g) + S8(s) → 8 H2S(g)

When 25.6 g of S8(s) (MM = 256 g•mol-1) reacts completely with an excess of H2(g) according to the equation above, the volume of H2S(g), measured at 0oC and 1.00 atm, produced is closest to

(A) 30 L (B) 20 L (C) 10 L (D) 5 L

55. 2 N2H4 + N2O4 → 3 N2 + 4 H2O

What mass of N2 can be produced when 8.0 g of N2H4 (MM = 32 g) and 9.2 g of N2O4 (MM = 92 g) react?

(A) 8.4 g (B) 12.6 g (C) 7.8 g (D) 10.5 g

56. A student weighs out 0.0154 mol of pure, dry NaCI in order to prepare a 0.154 M NaCI solution. Of the following pieces of laboratory equipment, which would be most essential for preparing the solution?

(A) 50 mL volumetric pipet

(B) 100 mL Erlenmeyer flask

(C) 100 mL graduated beaker

(D) 100 mL volumetric flask

Questions 57-60 The figures show portions of a buret used in a titration 0.0464 moles of monoprotic acid with a solution of Ba(OH)2. Figures I and 2 show the level of the Ba(OH)2 solution at the start and at the endpoint of the titration, respectively. Phenolphthalein was used as the indicator for the titration.

Figure 1 Figure 2

57. What is the evidence that the endpoint of the titration has been reached?

(A) The color of the solution in the buret changes from pink to colorless.

(B) The color of the solution in the buret changes from blue to red.

(C) The color of the contents of the flask changes from colorless to pink.

(D) The color of the contents of the flask changes from blue to red

58. The volume of Ba(OH), used to neutralize the acid was closest to

(A) 22.80 mL (B) 23.02 mL

(C) 23.20 mL (D) 29.80 mL

59. The concentration of the Ba(OH)2 solution is closest to

(A) 1 M (B) 2 M (C) 3 M (D) 4 M

60. What could explain why the student calculated a concentration of Ba(OH)2 that was too large?

(A) An extra drop of phenolphthalein was added.

(B) A small amount of the acid was not transferred to the titration flask.

(C) A drop of Ba(OH)2 remained attached to the buret tip.

(D) Rinsing the buret with distilled water just before filling it with the Ba(OH)2 to be titrated.

61. CaCl2(s) → Ca2+ + 2 Cl-

For the process of solid calcium chloride dissolving in water, represented above, the entropy change might be expected to be positive. However, ΔS for the process is actually negative. Which best helps to account for the net loss of entropy?

(A) Cl- ions are much larger in size than Ca2+ ions.

(B) The particles in solid calcium chloride are more ordered than are particles in amorphous solids.

(C) Water molecules in the hydrated Ca2+ and Cl- ions are more ordered than they are in the pure water.

(D) The Ca2+ and Cl- ions are more free to move around in solution than they are in CaCl2(s)

62. For which of the processes does entropy decrease (ΔS < 0)?

(A) H2O(s) → H2O(l)

(B) Br2(l) → Br2(g)

(C) Crystallization of I2(s) from an ethanol solution

(D) Thermal expansion of a balloon filled with CO2(g)

63. What mass of Cu(s) would be produced if 0.40 mol of Cu2O(s) was reduced completely with excess H2(g)?

(A) 13 g (B) 25 g (C) 38 g (D) 51 g

64. A certain reaction is spontaneous at temperatures below 400 K but is not spontaneous at temperatures above 400 K. If ΔHo for the reaction is -20 kJ•mol-1 and it is assumed that ΔHo and ΔSo do not change appreciably with temperature, then the value of ΔSo for the reaction is

(A) -50 J•mol-1•K-1 (B) -20.0 J•mol-1•K-1

(C) -0.05 J•mol-1•K-1 (D) -20 J•mol-1•K-1

65. Z → X + Y

A pure substance Z decomposes into two products, X and Y, as shown by the equation. Which of the following graphs of the concentration of Z versus time is consistent with the rate of the reaction being first order with respect to Z?

(A) (B)

(C) (D)

Questions 66-67 refer to an experiment to determine the heat of solution of an ionic solid. A student used a calorimeter consisting of a polystyrene cup and a thermometer. The cup was weighed, then filled halfway with water, then weighed again. The temperature of the water was measured, and some of the ionic solid was added to the cup. The mixture was gently stirred until all of the solute dissolved and the lowest temperature reached by the water in the cup was recorded. The cup and its contents were weighed again.

66. The purpose of weighing the cup and its contents again at the end of the experiment was to

(A) determine the mass of solute that was added.

(B) determine the mass of the thermometer.

(C) determine the mass of water that evaporated.

(D) verify the mass of water that was cooled.

67. Suppose that during the experiment, a significant amount of solution spilled from the polystyrene cup before all of the solute dissolved. How does this affect the calculated value for the heat of solution of the ionic compound?

(A) The calculated value is too large because less water was cooled as the remaining solute dissolved.

(B) The calculated value is too large because some solute was lost with the spilled solution.

(C) The calculated value is too small because less solute was dissolved than the student assumed.

(D) The calculated value is too small because the total mass of the calorimeter contents was too small.

68. Pb(s) Δ Pb(l)

Which of the following is true for the process represented above at 327oC and 1 atm? (The normal melting point for Pb(s) is 327o(C))

(A) ΔH = 0 (B) TΔS = 0

(C) ΔS < 0 (D) ΔH = TΔS

69. C(diamond) → C(graphite)

For the reaction represented above, the standard Gibbs free energy change, ΔGo298, has a value of -2.90 kJ•mol-1. Which of the following best accounts for the observation that the reaction does NOT occur (i.e. diamond is stable) at 298 K and 1.00 atm?

(A) ΔSo for the reaction is positive.

(B) The activation energy, Ea, for the reaction is very large.

(C) The reaction is slightly exothermic (ΔHo < 0).

(D) Diamond has a density greater than that of graphite.

70. When a solution is formed by adding some methanol, CH3OH, to water, processes that are endothermic include which of the following?

I. Methanol molecules move water molecules apart as the methanol goes into solution.

II. Water molecules move methanol molecules apart as the methanol goes into solution.

III. Intermolecular attractions form between molecules of water and methanol as the methanol goes into solution.

(A) I only (B) II only (C) III only (D) I and II only

Answers

|# |Π |Explanation |

|1 |d |Average = (Mass1 x Abundance1) + (Mass2 x Abundance2) |

| | |63.5 = (63)(1 – x) + (65)(x) = 63 – 63x + 65 x ∴ x = 0.25 (25 %) |

|2 |d |3920Ca → 3919K + 01β |

|3 |c |It takes 3 half-lives to reduce the radioactivity to 1/8 (1.25/10.0). |

| | |195 days/3 = 65 days |

|4 |b |From the diagram: E6 = -0.38 eV and E2 = -3.40 eV |

| | |ΔE = E2 – E6 = -3.40 eV – (-0.38 eV) = -3.02 eV |

|5 |c |The electron can only absorb energy that will move it to a higher |

| | |energy level. 9.4 eV is not enough energy (the minimum needed is |

| | |-3.40 eV – (-13.60 eV) = 10.2 eV). |

|6 |c |The orbital diagram for C,1s(↑↓) 2s(↑↓) 2p(↑ )(↑ )( ), has two |

| | |unpaired electrons (Hund's rule) = paramagnetic. |

|7 |b |Pauli states that no orbital can contain electrons with the same spin.|

| | |Since two spins, this limits the number to two electrons. |

|8 |a |Heisenberg states that the wave nature of matter (DeBroglie) limits |

| | |what we can know about position and velocity. |

|9 |c |Electron # 31 is located in the 4th row (n = 4), 13th column (p |

| | |section, l = 1), which limits ml = 1, 0 or -1 and ms = +½ or -½ ∴ (4,|

| | |1, 1, ½) fits requirement |

|10 |b |1s(↑↓) 2s(↑ ) 2p(↑ )( )( ): The 2p electron is in an excited|

| | |state, otherwise it would go into the 2s sublevel. |

|11 |c |1s(↑↓) 2s(↑↓) 2p(↑ )(↑ )(↑ ): The 2 2s electrons and 3 2p electrons|

| | |are in the valence shell (highest energy level) ∴ five. |

|12 |a |1s(↑↓): The first ionized electron is from the 1s sublevel. It takes |

| | |the most energy to remove electrons that are close to nucleus. |

|13 |c |The biggest jump in ionization energy occurs between 4 and 5, which |

| | |means 4 valence electrons ∴ Carbon group. |

|14 |c |Atomic radius increases going left and down in periodic table. Anion |

| | |is larger than atom and cation is smaller than atom. |

|15 |d |Elements in the same column in the periodic table have similar |

| | |chemical properties. |

|16 |d |Lattice energy is a measure of ionic bond strength, which is |

| | |proportional to charge and inversely proportional to size. |

|17 |c |Single bonds are the weakest (C≡O, O=O, Cl–Cl, N≡N) ∴ Cl2. |

|18 |a |Most polar bond forms between atoms with the greatest |

| | |electronegativity difference (greatest gap on the periodic table). |

|19 |d |SeCl4 has 34 valence electrons, which require an expanded octet system|

| | |(sp3d) to accommodate all the electrons. |

|20 |b |Only SO2 has a single and double bond, which can exchange places, thus|

| | |forming resonance forms. (H2S and OF2: 2 single bonds, CO2: 2 double |

| | |bonds) |

|21 |d |::S=Se:=S:: has zero formal charge. There are 4 bonds and 5 pairs of |

| | |nonbonding electrons. |

|22 |a |C≡O: The triple bond between C and O is composed of one sigma bond and|

| | |two pi bonds. |

|23 |c |H-F: The electronegativity difference is greatest between H and F ∴ |

| | |the most polar, which produces the highest dipole moment. |

|24 |d |PH3: The three H are pushed away from the pair of non-bonding |

| | |electrons around phosphorus resulting in a pyramidal structure. |

|25 |b |A bond order of 1.5 means 1 sigma bond and 50% share of a pi bond, |

| | |which is the case for O3 (O=O–O). |

|26 |c |CCl4 (tetrahedron), CO2 (linear), PCl3 (trigonal pyramid), PCl5 |

| | |(trigonal bipyramid), SF6 (octahedron) |

|27 |d |The non-bonding electron pairs take up more space than bonding pairs ∴|

| | |H2O (2 pairs) < NH3 (1 pair) < CH4 (0 pair). |

|28 |b |Acids contain the COOH functional group ∴ (B) (a is a ketone, c is an|

| | |alcohol, and d is an ether) |

|29 |d |4-methylpentane is the same as 2-methylpentane because you number from|

| | |the closest end ∴ 4 becomes 2. |

|30 |d |C1H3–C2≡C3H: C1 is sp3, C2 is sp, C3 is sp |

|31 |d |Both molecules are non-polar, but the larger molar mass of CCl4 means |

| | |that there are more electrons, which are more polarizable ∴ generating|

| | |a stronger dispersion force. |

|32 |b |Hydrogen bonding occurs when H is bonded to N, O or F. Only N2H4 has |

| | |that arrangement. (CH2F2: H is not bonding to the F) |

|33 |d |Farthest from each other in the gaseous phase, which is at t5. |

|34 |c |Melting occurs along 1st plateau (t2) and boiling along 2nd plateau |

| | |(t4) ∴ the time between these two temperatures is t3. |

|35 |b |(a) Covalent bonds aren't broken if its molecular. |

| | |(c) Volume only increases if liquid state is less dense. |

| | |(d) Temperature doesn't increase ∴ KE is not greater. |

|36 |d |At the same temperature, lighter molecules have greater speed. |

|37 |b |40 g of Ar = 1 mol. One mole at STP = 22.4 L. Since P is 2 x |

| | |standard, then T has to be 2 x standard (V = nRT/P) ∴ 2 x 273 K |

|38 |b |PN2 = XN2Ptot = (0.25/(0.35 + 0.90 + 0.25))(3.0 atm) = 0.50 atm |

|39 |a |Real gases deviated from ideal behavior at low temperatures (near |

| | |their boiling point) and high pressure. |

|40 |c |(A) metallic bond, (B) ionic bond, (C) covalent bond, and (D) |

| | |molecular bond. |

|41 |c |Gas solubility increases with greater partial pressure of the gas in |

| | |the container or lower temperature. |

|42 |b |6 molal = 6 mol ethanol in 1000 g H2O |

| | |1000 g H2O x 1 mol/18 g = 55 mol H2O |

| | |∴ mole fraction = 6/(6 + 55) = 0.1 |

|43 |d |Molality is moles solute/kg solvent. Molarity is moles solute/L |

| | |solution. To determine kg solvent from L solution, you need to know |

| | |density; then subtract g solute. |

|44 |d |0.2 mole fraction = 0.2 mol toluene in 0.8 benzene |

| | |0.8 mole x 80 g/1 mol = 60 g benzene |

| | |molality = 0.2 mol/0.060 kg = 3 m |

|45 |b |The highest boiling point = highest concentration of ions. |

| | |(A) .2 x 3 = .6, (B) .25 x 3 = .75, (C) .3 x 2 = .6, (D) .4 x 1 = .4 |

|46 |c |1 CH3OCH3(g) + 3 O2(g) → 2 CO2(g) + 3 H2O(g) |

|47 |d |0.250 L x 0.400 mol/L = 0.1 mol KBr x 119 g/mol = 11.9 g |

|48 |a |If CO2 is soluble in water, then some gas would remain in the water |

| | |and not bubble into the gas collecting bottle. |

|49 |b |CH3OH has the lowest proportion of C (12/32) ∴ given equal masses; |

| | |CH3OH would generate the least amount of CO2. |

|50 |d |molarity = mol solute/volume solution (L). The mass is needed to |

| | |determine the number of moles of solute. The volume is needed to |

| | |determine the volume of solution. |

|51 |d |0.6 mol CS2 x 3 mol Products/1 mol CS2 = 1.8 mol Products |

| | |1.5 mol O2 x 3 mol Products/3 mol O2 = 1.5 mol Products |

| | |1.5 mol Products is the lesser number. |

|52 |c |Assume 100 g of compound |

| | |10 g H x 1 mol/1 g = 10 mol/7.5 = 1.33 x 3 = 4 |

| | |90 g C x 1 mol/12 g = 7.5 mol/7.5 = 1 x 3 = 3 |

|53 |d |The resulting solution must have a concentration between the two |

| | |solutions added together. 0.14 M is less then both. |

|54 |b |25.6 g S8 x 256 g/mol x 8 mol H2S/1 mol S8 = 0.8 mol |

| | |0.8 mol x 22.4 L/mol = 20 L |

|55 |a |0.25 mol N2H4 x 3 mol N2 x 28 g N2 = 10.5 g |

| | |2 mol N2H4 1 mol N2 |

| | |0.1 mol N2O4 x 3 mol N2/1 mol N2O4 x 28 g N2 = 8.4 g |

|56 |d |0.0154 mol NaCl x 1 L/0.154 mol = 0.1 L (100 mL) |

| | |The most accurate way to measure 0.1 L of solution is to use a |

| | |volumetric flask. |

|57 |c |Phenolphthalein changes from clear (acid) to pink (base) |

|58 |c |Final volume – Initial volume = Change in volume |

| | |35.75 mL – 12.55 mL = 23.20 mL |

|59 |a |0.0464 mol H+ x 1 mol OH- x 1 mol B(A).. = 0.0232 mol B(A).. |

| | |1 mol H+ 2 mol OH- |

| | |0.0232 mol B(A)../0.0232 L = 1 M |

|60 |b |MBa(OH)2 = (½ mol H+)/VBa(OH)2 |

| | |∴ too large MBa(OH)2 = too small VBa(OH)2 |

| | |Titrating less acid would result in smaller VNaOH |

|61 |c |Dissolving involves two process; (1) separation into ions, which |

| | |increases disorder (+ΔS) and (2) ions combining with water |

| | |(solvation), which decreases disorder (-ΔS). |

|62 |c |Disorder decreases when I2(aq) → I2(s). Disorder increases when s → l|

| | |→ g (A) and (B). Disorder also increases when gas molecules spread |

| | |out (D). |

|63 |d |Cu2O + H2 → 2 Cu + H2O (balancing wasn't necessary because mole Cu in |

| | |reactants and products are equal) |

| | |0.80 mol Cu x 63.5 g/mol = 51 g |

|64 |a |Tthreshold = ΔH/ΔS |

| | |ΔS = ΔH/Tthreshold = -20 kJ•mol-1/400 = -0.05 kJ•mol-1•K-1 |

| | |-0.05 kJ•mol-1•K-1 x 1000 J/1 kJ = -50 J•mol-1•K-1 |

|65 |d |For a first order reaction, the straight line graph is ln[Z] vs. t. |

| | |(zero order is [Z] vs. t and second order is 1/[Z] vs. t) |

|66 |a |At the end of the experiment, the cup contained solute and water. If |

| | |this value is subtracted from the mass of the cup and water, then the |

| | |difference is the mass of the solute. |

|67 |a |The heat needed to dissolve the remaining solute had to come from less|

| | |water, which would make ΔT greater than it should have been (ΔH = |

| | |-mcΔT) ∴ ΔH would be too large. |

|68 |d |At the normal melting point: ΔG = 0 = ΔH – TΔS ∴ ΔH = TΔS |

|69 |b |The reaction is spontaneous, but it must not occur at a fast rate. |

| | |This could be because the activation energy is so high, that it takes |

| | |too much energy to start the process. |

|70 |d |Breaking solute-solute bonds in methanol and water is endothermic, but|

| | |forming solute-solute bonds between methanol and water is exothermic. |

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