AP Chemistry Unit I Test

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					                                    AP Chemistry Unit I Test Preview
   Note: For all questions, assume that the temperature is 298 K, the pressure is 1.00 atmosphere, and solutions are
    aqueous unless otherwise specified.
   Throughout the test the following symbols have the definitions specified unless otherwise noted.

          T       = temperature              m         = molal
          P       = pressure                 L, mL     = liter(s), milliliter(s)
          V       = volume                   g         = gram(s)
          S       = entropy                  nm        = nanometer(s)
          H       = enthalpy                 atm       = atmosphere(s)
          G       = free energy              J, kJ     = joule(s), kilojoule(s)
          R       = molar gas constant       V         = volt(s)
          n       = number of moles          mol       = mole(s)
          M       = molar

Section I: Part A
Directions: Each set of lettered choices below refers to the numbered statements immediately following it. Select the
one lettered choice that best fits each statement and then fill in the corresponding oval on the answer sheet. A choice
may be used once, more than once, or not at all in each set.

Questions 1-3 refer to the following list.

                             (A) Empirical formula
                             (B) Molecular formula
                             (C) Molecular mass
                             (D) Formula mass
                             (E) Molar mass
 ___ 1. The sum of the atomic masses of all the atoms or ions in a ionic compound
 ___ 2. May be determined experimentally by measuring the mass % composition of each element in a compound,
        converting the masses to molar amounts, then determining the ratio of moles of each element in the
 ___ 3. The mass, in grams, of 6.022 x 1023 molecules of a given compound.
Questions 4-6 refer to aqueous solutions containing 1:1 mole ratios of the following pairs of substances. Assume all
concentrations are 1 M.

    (A)   NH3 and NH4Cl
    (B)   H3PO4 and NaH2PO4
    (C)   HCl and NaCl
    (D)   NaOH and NH3
    (E)   NH3 and HC2H3O2 (acetic acid)

 ___ 4. The solution with the lowest pH
 ___ 5. The solution with the highest pH
 ___ 6. The most nearly neutral solution

Questions 7-9 refer to the following elements.

                         (A) Carbon
                         (B) Silver                    (D) Uranium
                         (C) Bromine                   (E) Fluorine

 ___ 7. Is a gas in its standard state at 298 K.
 ___ 8. The element that forms the backbone of all organic molecules.
 ___ 9. Forms insoluble ionic compounds with group 7 elements.
Part B Directions: Each of the questions or incomplete statements below is followed by five suggested answers or
completions. Select the one that is best in each case and then write the letter in the blank on your test form..

___ 10. When the equation below is balanced and all coefficients are reduced to their lowest whole-number terms,
        the coefficient for O2(g) is
    ... C10H12O4S(s) + ... O2(g)  ... CO2(g) + ... SO2(g) + ... H2O(g)
    (A) 6                                   (B) 7
    (B) 14                                  (E) 28
    (C) 12

___ 11. What mass of Au is produced when 0.0500 mol of Au2S3 is reduced completely with excess H2 ?
   (A) 9.85 g
   (B) 39.4 g                             (D) 19.7
   (C) 24.5 g                             (E) 48.9 g
___ 12. The organic compound represented at right is an example of                 |
   (A) an organic acid                                                     CH3—C—CH2—CH3
   (B) an alcohol                         (D) an aldehyde
   (C) an ether                           (E) a ketone

___ 13. A 1.0 L sample of an aqueous solution contains 0.10 mol of NaCl and 0.10 mol of CaCl 2. What is the
        minimum number of moles of AgNO3 that must be added to the solution in order to precipitate all of the Cl –
        as AgCl(s)? (Assume that AgCl is insoluble.)
   (A) 0.10 mol
   (B) 0.30 mol                            (D) 0.40 mol
   (C) 0.60 mol                            (E) 0.20 mol

___ 14. When hafnium metal is heated in an atmosphere of chlorine gas, the product of the reaction is found to
        contain 62.2 percent Hf by mass and 37.4 percent Cl by mass. What is the empirical formula for this
   (A) HfCl
   (B) HfCl4                               (D) HfCl2
   (C) HfCl3                               (E) Hf2Cl3

___ 15. A 40.0 mL sample of 0.25 M KOH is added to 60.0 mL of 0.15 M Ba(OH)2. What is the molar concentration
        of OH–(aq) in the resulting solution? (Assume that the volumes are additive.)
   (A) 0.10 M
   (B) 0.40 M                                (D) 0.55 M
   (C) 0.28 M                                (E) 0.19 M

___ 16. According to the balanced equation below, how many moles of HI would be necessary to produce 2.5 mol of
        I2, starting with 4.0 mol of KMnO4 and 3.0 mol of H2SO4?
      10 HI + 2 KMnO4 + 3 H2SO4  5 I2 + 2 MnSO4 + K2SO4 + 8 H2O
    (A) 20.
    (B) 10.                                 (D) 5.0
    (C) 8.0                                 (E) 2.5

___ 17. Which compound is NOT appreciably soluble in water but is soluble in dilute hydrochloric acid?
   (A) Mg(OH)2(s)
   (B) CuSO4(s)                        (D) (NH4)2SO4(s)
   (C) Sr(NO3)2(s)                     (E) (NH4)2CO3(s)

___ 18. In which of the following processes are covalent bonds broken?
   (A) I2(s)  I2(g)
   (B) NaCl(s)  NaCl(l)                    (D) C(diamond)  C(g)
   (C) Fe(s)  Fe(l)                        (E) CO2(s)  CO2(g)
___ 19. When 100 mL of 1.0 M Na3PO4 is mixed with 100 mL of 1.0 M AgNO3, a yellow precipitate forms and
        [Ag+] becomes negligibly small. Which of the following is a correct listing of the ions remaining in solution
        in order of increasing concentration?
   (A) [PO43–] < [NO3–] < [Na+]
   (B) [PO43–] < [Na+] < [NO3–]              (D) [Na+] < [NO3–] < [PO43–]
              –         3–
   (C) [NO3 ] < [PO4 ] < [Na ]  +
                                             (E) [Na+] < [PO43–] < [NO3–]

___ 20. The volume of distilled water that should be added to 10.0 mL of 6.00 M HCl(aq) in order to prepare a 0.500
        M HCl(aq) solution is approximately
   (A) 50.0 mL
   (B) 60.0 mL                               (D) 110. mL
   (C) 100. mL                               (E) 120. mL

___ 21. In which of the following compounds is the mass ratio of chromium to oxygen closest to 1.6 to 1.00 ?
   (A) CrO3
   (B) CrO2                               (D) Cr2O
   (C) CrO                                (E) Cr2O3

___ 22. Which of the following actions would be likely to change the boiling point of a sample of a pure liquid in an
        open container?
          I.         Placing it in a smaller container
          II.        Increasing the number of moles of the liquid in the container
          III.       Moving the container and liquid to a higher altitude
      (A) I only
      (B) II only                               (D) I, II, and III
      (C) III only                              (E) II and III only

___ 23. CH3CH2OH boils at 78oC and CH3OCH3 boils at –24oC, although both compounds have the same
        composition. This difference in boiling points may be attributed to a difference in
   (A) molecular mass
   (B) density                               (D) heat of combustion
   (C) specific heat                         (E) hydrogen bonding

___ 24. A hydrocarbon gas with an empirical formula CH2 has a density of 1.88 grams per liter at 0 oC and 1.00
        atmosphere (= standard temperature and pressure, or STP). At STP gases occupy 22.4 L per mole. A
        possible formula for the hydrocarbon is
   (A) CH2
   (B) C2H4                                (D) C4H8
   (C) C3H6                                (E) C5H10

 ___ 25. Which of the following pairs of compounds are isomers?
(A)              CH3-CH2-CH2-CH3            and     CH3-CH-CH3

                         CH3                    CH3
                                               
(B)                  CH3-CH–CH3        and      CH-CH=CH2

(C)                  CH3-O-CH3         and      CH3-C-CH3

(D)                  CH3-OH            and      CH3-CH2-OH

(E)                  CH4      and      CH2=CH2
Take Home FRQs Maximum possible points that can be earned are given in [brackets].

2004 B
               2 Fe(s) +     O (g)  Fe2O3(s)                                 ∆Hf˚ = -824 kJ mol–1
                           2 2

Iron reacts with oxygen to produce iron(III) oxide as represented above. A 75.0 g sample of Fe(s) is
mixed with 11.5 L of O2(g) at a pressure of 2.66 atm and temperature of 298 K.
Note to Mr T’s Students: convert the volume of gases to moles by using the ideal gas law and the formula PV= nRT
(see Equations and Constants sheet).

a) Calculate the number of moles of each of the following before the reaction occurs.
    (i) Fe(s)                                                                                          [1]

     (ii)   O2(g)                                                                                      [1]

b) Identify the limiting reactant when the mixture is heated to produce Fe 2O3. Support your answer with
   calculations.                                                                                [2]

c) Calculate the number of moles of Fe2O3 produced when the reaction proceeds to completion. [1]

Answer the following questions about acetylsalicylic acid, the active ingredient in aspirin.
a) The amount of acetylsalicylic acid in a single aspirin tablet is 325 mg, yet the tablet has a mass of
   2.00 g. Calculate the mass percent of acetylsalicylic acid in the tablet.                        [1]

b) The elements contained in acetylsalicylic acid are hydrogen, carbon, and oxygen. The combustion of
   3.000 g of the pure compound yields 1.200 g of water and 3.72 L of dry carbon dioxide, measured at
   750. mm Hg and 25C. Calculate the mass, in g, of each element in the 3.000 g sample.      [4]
2008 2. Answer the following questions relating to gravimetric analysis.

In the first of two experiments, a student is
assigned the task of determining the
number of moles of water in one mole of
MgCl2 • n H2O. The student collects the
data shown in the following table.

(a) Explain why the student can correctly
conclude that the hydrate was heated a
sufficient number of times in the experiment.                                                       [1]

(b) Use the data above to
        (i) calculate the total number of moles of water lost when the sample was heated, and       [1]

        (ii) determine the formula of the hydrated compound.                                        [2]

(c) A different student heats the hydrate in an uncovered crucible, and some of the solid spatters out of
the crucible. This spattering will have what effect on the calculated mass of the water lost by the hydrate?
Justify your answer.                                                                                 [1]

In the second experiment, a student is given 2.94 g of a mixture containing anhydrous MgCl 2 and KNO3 .
To determine the percentage by mass of MgCl2 in the mixture, the student uses excess AgNO3 (aq) to
precipitate the chloride ion as AgCl(s).

(d) Starting with the 2.94 g sample of the mixture dissolved in water, briefly describe the steps necessary
to quantitatively determine the mass of the AgCl precipitate.                                        [2]

(e) The student determines the mass of the AgCl precipitate to be 5.48 g. On the basis of this information,
calculate each of the following.
        (i) The number of moles of MgCl2 in the original mixture                                    [2]

        (ii) The percent by mass of MgCl2 in the original mixture                                   [1]

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