AP Released Stoich Questions 1970 - 2005

AP Questions: Empirical / Molecular Formula and Stoichiometry 1970 A 2.000 gram sample containing graphite (carbon) and an inert substance was burned in oxygen and produced a mixture of carbon dioxide and carbon monoxide in the mole ratio 2.00:1.00. The volume of oxygen used was 747.0 milliliters at 1,092K and 12.00 atmospheres pressure. Calculate the percentage by weight of graphite in the original mixture. 1982 B Water is added to 4.267 grams of UF6. The only products are 3.730 grams of a solid containing only uranium, oxygen and fluorine and 0.970 gram of a gas. The gas is 95.0% fluorine, and the remainder is hydrogen. (a) From these data, determine the empirical formula of the gas. (b) What fraction of the fluorine of the original compound is in the solid and what fraction in the gas after the reaction? (c) What is the formula of the solid product? (d) Write a balanced equation for the reaction between UF6 and H2O. Assume that the empirical formula of the gas is the true formula. 1986 B Three volatile compounds X, Y, and Z each contain element Q. The percent by weight of element Q in each compound was determined. Some of the data obtained are given below. Percent by weight Molecular Compound of Element Q Weight X 64.8% ? Y 73.0% 104. Z 59.3% 64.0 (a) The vapor density of compound X at 27C and 750. mm Hg was determined to be 3.53 grams per liter. Calculate the molecular weight of compound X. (b) Determine the mass of element Q contained in 1.00 mole of each of the three compounds. (c) Calculate the most probable value of the atomic weight of element Q. (d) Compound Z contains carbon, hydrogen, and element Q. When 1.00 gram of compound Z is oxidized and all of the carbon and hydrogen are converted to oxides, 1.37 grams of CO2 and 0.281 gram of water are produced. Determine the most probable molecular formula of compound Z. 1991 B The molecular formula of a hydrocarbon is to be determined by analyzing its combustion products and investigating its colligative properties. (a) The hydrocarbon burns completely, producing 7.2 grams of water and 7.2 liters of CO2 at standard conditions. What is the empirical formula of the hydrocarbon? (b) Calculate the mass in grams of O2 required for the complete combustion of the sample of the hydrocarbon described in (a). 1995 B A sample of dolomitic limestone containing only CaCO3 and MgCO3 was analyzed. (a) When a 0.2800 gram sample of this limestone was decomposed by heating, 75.0 milliliters of CO 2 at 750 mm Hg and 20C were evolved. How many grams of CO2 were produced. (b) Write equations for the decomposition of both carbonates described above. (c) It was also determined that the initial sample contained 0.0448 gram of calcium. What percent of the limestone by mass was CaCO3? (d) How many grams of the magnesium-containing product were present in the sample in (a) after it had been heated? 1993 B I. 2 Mn2+ + 4 OH- + O2(g)  2 MnO2(s) + 2 H2O II. MnO2(s) + 2 I- + 4 H+  Mn2+ + I2(aq) + 2 H2O III. 2 S2O32- + I2(aq)  S4O62- + 2 IThe amount of oxygen, O2, dissolved in water can be determined by titration. First, MnSO4 and NaOH are added to a sample of water to convert all of the dissolved O2 to MnO2, as shown in equation I above. Then H2SO4 and KI are added and the reaction represented by equation II proceeds. Finally, the I2 that is formed is titrated with standard sodium thiosulfate, Na2S2O3, according to equation III. (a) According to the equation above, how many moles of S2O32- are required for analyzing 1.00 mole of O2 dissolved in water? (b) A student found that a 50.0-milliliter sample of water required 4.86 milliliters of 0.0112-molar Na2S2O3 to reach the equivalence point. Calculate the number of moles of O2 dissolved in this sample. (c) How would the results in (b) be affected if some I2 were lost before the S2O32- was added? Explain. (d) What volume of dry O2 measured at 25C and 1.00 atmosphere of pressure would have to be dissolved in 1.00 liter of pure water in order to prepare a solution of the same concentration as that obtained in (b)? (e) Name an appropriate indicator for the reaction shown in equation III and describe the change you would observe at the end point of the titration. 2000 B Answer the following questions about BeC2O4(s) and its hydrate. (a) Calculate the mass percent of carbon in the hydrated form of the solid that has the formula BeC2O4•3H2O. (b) When heated to 220.C, BeC2O4•3H2O(s) dehydrates completely as represented below. BeC2O4•3H2O(s)  BeC2O4(s) + 3 H2O(g) If 3.21 g of BeC2O4•3H2O(s) is heated to 220.C calculate (i) the mass of BeC2O4(s) formed, and, (ii) the volume of the H2O(g) released, measured at 220.C and 735 mm Hg. (c) A 0.345 g sample of anhydrous BeC2O4, which contains an inert impurity, was dissolved in sufficient water to produce 100. mL of solution. A 20.0 mL portion of the solution was titrated with KMnO 4(aq). The balanced equation for the reaction that occurred is as follows. 16 H+(aq) + 2 MnO4-(aq) + 5 C2O42-(aq)  2 Mn2+(aq) + 10 CO2(g) + 8 H2O(l). The volume of 0.0150 M KMnO4(aq) required to reach the equivalence point was 17.80 mL. (i) Identify the reducing agent in the titration reaction. (ii) For the titration at the equivalence point, calculate the number of moles of each of the following that reacted. • MnO4-(aq) • C2O42-(aq) (iii) Calculate the total number of moles of C2O42-(aq) that were present in the 100. mL of prepared solution. (iv) Calculate the mass percent of BeC2O4(s) in the impure 0.345 g sample. 2001 B 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. (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. (c) A student dissolved 1.625 g of pure acetylsalicylic acid in distilled water and titrated the resulting solution to the equivalence point using 88.43 mL of 0.102 M NaOH(aq). Assuming that acetylsalicylic acid has only one ionizable hydrogen, calculate the molar mass of the acid. (d) A 2.00  10-3 mole sample of pure acetylsalicylic acid was dissolved in 15.00 mL of water and then titrated with 0.100 M NaOH(aq). The equivalence point was reached after 20.00 mL of the NaOH solution had been added. Using the data from the titration, shown in the table below, determine (i) the value of the acid dissociation constant, Ka, for acetylsalicylic acid and (ii) the pH of the solution after a total volume of 25.00 mL of the NaOH solution had been added (assume that volumes are additive). Volume of 0.100M NaOH Added (mL) pH 0.00 2.22 5.00 2.97 10.00 3.44 15.00 3.92 20.00 8.13 25.00 ?