1. Stoichiometry By Danny Chen Andy Wang

2. 1993 B (a) According to the equation, how many moles of S 2 O 3 2- are required for analyzing 1.00 mole of O 2 dissolved in water? (b) A student found that a 50.0 milliliter sample of water required 4.86 milliters of 0.0112-molar Na 2 S 2 O 3 to reach the equivalence point. Calculate the number of moles of O 2 dissolved in this sample. C. How would the results in (b) be affected if some I 2 were lost before the S 2 O 3 2- was added? Explain. D. What volume of dry O 2 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 I.2 Mn 2+ + 4 OH- + O 2 (g) => 2MnO 2 (s) +2 H 2 O II. MnO 2 (s) + 2 I - +4 H + => Mn 2 + + I 2 (aq) + 2 H 2 O III. 2 S 2 O 3 2- + I 2 (aq) => S 4 O 6 2- + 2 I - The amount of oxygen, O 2, dissolved in water can be determined by titration. First, MnSO 4 and NaOH are added to a sample of water to convert all of the dissolved O 2 to MnO 2, as shown in equation I above. Then H 2 SO 4 and KI are added and the reaction represented by equation II proceeds. Finally, the I 2 that is formed is titrated with standard sodium thiosulfate, Na 2 S 2 O 3, according to equation III.

3. 1993B - Answers Answer: (a) 1 mol O 2 x 2 mol MnO 2 /1 Mol O 2 x 1 mol I 2 /1 mol MnO 2 (b) 1 mol S 2 O 3 2- = (0.00486 L)(0.0112 M) = 5.44x10 -5 mol S 2 O 3 2- (c) Less I 2 means less thiosulfate ion required thus indicating a lower amount of dissolved oxygen. (d)molarity of solution in (b) = 1.36x10 -5 mol O 2 / 0.050 L = 2.72x10 -4 M = 6.65x10 -5 L or 6.65 mL O 2 (e)starch indicator color disappears or blue disappears I.2 Mn 2+ + 4 OH- + O 2 (g) => 2MnO 2 (s) +2 H 2 O II. MnO 2 (s) + 2 I - +4 H + => Mn 2 + + I 2 (aq) + 2 H 2 O III. 2 S 2 O 3 2- + I 2 (aq) => S 4 O 6 2- + 2 I - The amount of oxygen, O 2, dissolved in water can be determined by titration. First, MnSO 4 and NaOH are added to a sample of water to convert all of the dissolved O 2 to MnO 2, as shown in equation I above. Then H 2 SO 4 and KI are added and the reaction represented by equation II proceeds. Finally, the I 2 that is formed is titrated with standard sodium thiosulfate, Na 2 S 2 O 3, according to equation III.

4. 1995 B 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 CO 2 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 CaCO 3 ? D. How many grams of the magnesium- containing product were present in the sample in (a) after it had been heated? A Sample of dolomitic limestone containing only CaCO 3 and MgCO 3 was analyzed.

5. 1995B - Answers A. 3.08x10 -3 mol x f(44.0 g CO 2,1 mol) = 0.135 g CO 2 B. CaCO 3 => CaO + CO 2 MgCO 3 => MgO + CO 2 0.0448 gCa x 1 mol Ca/40.08 g Ca x 1 mol CaCO/1 mol Ca C. 0.112 g CaCO3/0.2800 g sample = 40.0% CaCO3 60.0% of 0.2800 g sample = 0.168 g of MgCO 3 0.168g MgCO 3 (1mol MgCO 3 / 84.3g MgCO 3 ) (1 mol MgO/1 mol MgCO 3 ) (40.3g MgO/1 mol MgO) = 0.0803 g MgO

6. 2000 B A. Calculate the mass percent of carbon in the hydrated form of the solid that has the forumla BeC 2 O 4 *3H 2 O. B. When heated to 220*C, BeC 2 O 4 *3H 2 O(s) dehydrates completely as represented below. BeC 2 O 4 *3H 2 O(s) => BeC 2 O 4 (s) + 3 H 2 O(g) C. A 0.345 g sample of anhydruous BeC 2 O 4, which contains an inert impurity, was dissolved in sufficent 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 MnO 4- (aq) + 5 C 2 O 4 2- (aq) => 2 Mn 2 +(aq) + 10 CO 2 (g) + 8 H 2 O(l). Answer the Following questions about BEC2O4(s) and its hydrate.

7. 2000 B a) Total mass of carbon / molar mass X 100 = 24.022/151.03 x 100 =15.9% b) i) 3.21g x 1mol hyd./151.03 g x 1mol anhyd./ 1 mol hyd. X 97.03 g anhyd./1 mol anhyd. = 2.06 g ii) Mol H 2 O = 3.21g – 2.06 g/18.0g mol =0.06375 mol V= nRT/P = (0.0635 mol) (0.0821) (220 + 273K) 735/760 atm =2.67 L c) (i) C 2 O 4 2- (aq) (ii) 17.80 mL x 0.0150 mol MnO 4 - 1000 mL =2.67 x10 –4 mol MnO 4 - 2.67´10–4 mol MnO4- x 5 mol C 2 O 4 2- 2 mol MnO 4 - =6.68x10 –4 mol C 2 O 4 2- (iii) 100mL x 6.68x10-4 mol C 2 O 4 2- /20 mL = 3.34x10-3 mol C 2 O 4 (iv) 3.34x10–3 mol C 2 O 4 2 x 1 mol BeC 2 O 4 /1 mol C 2 O 4 2- x 97.03g/1mol = 0.324g BeC 2 O 4 (s) 0.324g BeC 2 O 4 (s)/0.345 g sample x 100 = 93.9%

8. Volume of 0.100M NaOH Added (mL) pH 0.002.22 5.002.97 10.003.44 15.003.92 20.008.13 25.00? 2001 B 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 sissolved 1.625 g of pure acetylsalicylic acid in distilled water and titrated the resulting solution to the equivalaence 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 x 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 pointwas 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 acetylsalicyclic acid (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). Answer the following questions about acetylsalicylic acid, the active ingredient in aspirin.

9. 2001B Answers A. 0.325 g/2.00g x 100% = 16.3% B. 1.200 g H 2 O x (1.0079)(2) + 16 g H 2 O = 0.134 g H n= p*V/R*T = 750/760atm (3.72L)/(0.0821L*atm*mol-1*K-1)(298K)= 0.150 mol CO 2 0.150 mol CO 2 X 12.0 g C/1mol CO2 = 1.801 g C 3.000 g ASA – (1.801 g C + 0.134 g H) = 1.065 g O C. 0.08843 L x 0.102 mol/1L = 0.00902 mol base 1 mol base = 1 mol acid 1.625 g ASA/0.00902 mol = 180 g/mol D. (i) Hasa Asa- + H+ 2.00 x 10-3 moll/0.015L = 0.133M pH = -log[H+]; 2.22 = -log[H+] [H+] = M = [Asa-] [Hasa] = 0.133 M – 6.03 x 10-3 M = 0.127 M K = [H+][Asa-]/[Hasa] = (6.03x10-3)2/0.127 = 2.85x 10-4 OR when the solution is half-neutralized, pH = pK a at 10.00 mL, pH = 3.44; K = 10 –pH = 10 –3.44 = 3.63  10 -4 (ii) 0.025 L  0.100 mol/L = 2.50  10 -3 mol OH - 2.50  10 -3 mol OH - - 2.00  10 -3 mol neutralized = 5.0  10 -4 mol OH - remaining in (25 + 15 mL) of solution; [OH - ] = 5.0  10 -4 mol/0.040 L = 0.0125 M pH = 14 – pOH = 14 + log[OH - ] = 14 – 1.9 = 12.1 Answer the following questions about acetylsalicylic acid, the active ingredient in aspirin.