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1999 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART I Prepared by the American Chemical Society Olympiad Examinations Task Force OLYMPIAD EXAMINATIONS TASK FORCE Arden P. Zipp, State University of New York, Cortland Chair James S. Bock, Gateway High School, PA Edward DeVillafranca (retired), Kent School, CT Peter E. Demmin (retired), Amherst Central High School, NY John Krikau, American Chemical Society, DC Patricia A. Metz, University of Georgia, GA Ronald O. Ragsdale, University of Utah, UT Helen M. Stone (retired), Ben L. Smith High School, NC Diane D. Wolff, Ferrum College, VA DIRECTIONS TO THE EXAMINER–PART I Part I of this test is designed to be taken with a Scantron® answer sheet on which the student records his or her responses. Only this Scantron® sheet is graded for a score on Part I. Testing materials, scratch paper, and the Scantron sheet should be made available to the student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until April 26, 1999, after which tests can be returned to students and their teachers for further study. Allow time for the student to read the directions, ask questions, and fill in the requested information on the Scantron sheet. The answer sheet must be completed using a pencil, not pen. When the student has completed Part I, or after one hour and thirty minutes has elapsed, the student must turn in the Scantron sheet, Part I of the testing materials, and all scratch paper. There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and you are free to schedule rest-breaks between parts. Part I Part II Part III 60 questions 8 questions 2 questions single-answer multiple-choice problem-solving, explanations laboratory practical 1 hour, 30 minutes 1 hour, 45 minutes 1 hour, 15 minutes A periodic table and other useful information are provided on page 2 for student reference. Students should be permitted to use nonprogrammable calculators. DIRECTIONS TO THE EXAMINEE–PART I DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Answers to questions in Part I must be entered on a Scantron answer sheet to be scored. Be sure to write you name on the answer sheet; an ID number is already entered for you. Make a record of this ID number as you will use the same number on both Parts II and III. Each item in Part I consists of a question or an incomplete statement which is followed by four possible choices. Select the single choice that best answers the question or completes the statement. Then use a pencil to blacken the space on your answer sheet having the same letter as your choice. You may write on the examination, but the test booklet will not be used for grading. Scores are based on the number of correct responses. When you complete Part I (or at the end of one hour and 30 minutes), you must turn in all testing materials, scratch paper, and your Scantron answer sheet. Do not forget to turn in your U.S. citizenship statement before leaving the testing site today. Not valid for use as an USNCO National Exam after April 26, 1999. Distributed by the ACS DivCHED Examinations Institute, Clemson University, Clemson, SC. All rights reserved. Printed in U.S.A. amount of substance ampere atmosphere atomic mass unit atomic molar mass Avogadro constant Celsius temperature centi- prefix coulomb electromotive force energy of activation enthalpy entropy ABBREVIATIONS AND SYMBOLS n equilibrium constant K milli- prefix A Faraday constant F molal atm formula molar mass M molar u free energy G mole A frequency ν Planck’s constant N A gas constant R pressure °C gram g rate constant c hour h second C joule J speed of light E kelvin K temperature, K Ea kilo- prefix k time H liter L volt S measure of pressure mmHg volume CONSTANTS m m M mol h P k s c T t V V R = 8.314 J·mol –1·K–1 R = 0.0821 L·atm·mol –1·K–1 1 F = 96,500 C·mol–1 1 F = 96,500 J·V–1·mol–1 N A = 6.022 × 1023 mol–1 h = 6.626 × 10–34 J·s c = 2.998 × 108 m·s–1 PERIODIC TABLE OF THE ELEMENTS 1 H 2 He 1.008 4.003 3 Li 4 Be 5 B 6 C 7 N 8 O 9 F 10 Ne 6.941 9.012 10.81 12.01 14.01 16.00 19.00 20.18 11 Na 12 Mg 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar 22.99 24.31 26.98 28.09 30.97 32.07 35.45 39.95 19 K 20 Ca 21 Sc 22 Ti 23 V 24 Cr 25 Mn 26 Fe 27 Co 28 Ni 29 Cu 30 Zn 31 Ga 32 Ge 33 As 34 Se 35 Br 36 Kr 39.10 40.08 44.96 47.88 50.94 52.00 54.94 55.85 58.93 58.69 63.55 65.39 69.72 72.61 74.92 78.96 79.90 83.80 37 Rb 38 Sr 39 Y 40 Zr 41 Nb 42 Mo 43 Tc 44 Ru 45 Rh 46 Pd 47 Ag 48 Cd 49 In 50 Sn 51 Sb 52 Te 53 I 54 Xe 85.47 87.62 88.91 91.22 92.91 95.94 (98) 101.1 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.3 55 Cs 56 Ba 57 La 72 Hf 73 Ta 74 W 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 132.9 137.3 138.9 178.5 181.0 183.8 186.2 190.2 192.2 195.1 197.0 200.6 204.4 207.2 209.0 (209) (210) (222) 87 Fr 88 Ra 89 Ac 104 Rf 105 Db 106 Sg 107 Bh 108 Hs 109 Mt 110 111 112 114 (223) 226.0 227.0 (261) (262) (263) (262) (265) (266) (269) (272) (277) (289) Page 2 58 Ce 59 Pr 60 Nd 61 Pm 62 Sm 63 Eu 64 Gd 65 Tb 66 Dy 67 Ho 68 Er 69 Tm 70 Yb 71 Lu 140.1 140.9 144.2 (145) 150.4 152.0 157.3 158.9 162.5 164.9 167.3 168.9 173.0 175.0 90 Th 91 Pa 92 U 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr 232.0 231.0 238.0 237.0 (244) (243) (247) (247) (251) (252) (257) (258) (259) (260) Not valid for use as a USNCO National Exam after April 26, 1999 DIRECTIONS § When you have selected your answer to each question, blacken the corresponding space on the answer sheet using a soft, #2 pencil. Make a heavy, full mark, but no stray marks. If you decide to change an answer, erase the unwanted mark very carefully. § Make no marks on the test booklet. Do all calculations on scratch paper provided by your instructor. § There is only one correct answer to each question. Any questions for which more than one response has been blackened will not be counted. § Your score is based solely on the number of questions you answer correctly. It is to your advantage to answer every question. 1. Which substance is most likely to be soluble in a nonpolar solvent? (A) glucose (B) graphite (C) lithium fluoride (D) sulfur 2. A solution of which substance can best be used as both a titrant and its own indicator in an oxidation–reduction titration? (A) I2 (B) NaOCl (C) K2Cr2O7 (D) KMnO4 (A) 0.166 M (B) 0.180 M (C) 0.333 M (D) 0.666 M 7. When ionic hydrides react with water, the products are (A) acidic solutions and hydrogen gas. (B) acidic solutions and oxygen gas. (C) basic solutions and hydrogen gas. (D) basic solutions and oxygen gas. 36 8. 0.250 g of an element, M, reacts with excess fluorine to produce 0.547 g of the hexafluoride, MF6. What is the element? 32 T, °C 3. What value of ∆T should be used for the calorimetry experiment that gives these graphed results? 6. A 20.00 mL sample of a Ba(OH)2 solution is titrated with 0.245 M HCl. If 27.15 mL of HCl is required, what is the molarity of the Ba(OH)2 solution? 28 (A) Cr (B) Mo (C) S (D) Te 24 20 0 2 4 6 8 10 12 14 16 18 20 time, min (A) 10 ˚C (B) 12 ˚C (C) 15 ˚C (D) 19 ˚C Fe3+(aq) + SCN–(aq) ¾ FeSCN2+(aq) 4. The equilibrium constant for this reaction can best be determined by means of (A) chromatography. (B) conductance. (C) ion exchange. (D) spectrophotometry. 9. How many moles of Na + ions are in 20 mL of 0.40 M Na 3PO4? (A) 0.0080 (B) 0.024 (C) 0.050 (D) 0.20 10. What is the mass percent of oxygen in Al2(SO4)3·18H2O? (A) 9.60 Molar Mass, M Al2(SO4)3·18H2O 666.43 g·mol (B) 28.8 (C) 43.2 (D) 72.0 11. What is the coefficient for H +(aq) when the equation is balanced with whole number coefficients? 5. Which solid reacts with dilute hydrochloric acid at 25 ˚C to produce a gas that is more dense than air? (A) Zn (B) Pb(NO3)2 (C) NaBr (D) NaHCO3 __Mn2+(aq) + __BiO3–(aq) + __H+(aq) → __Bi3+(aq) + __MnO4–(aq) + __H2O(l) (A) 3 Not valid for use as an USNCO National Examination after April 26, 1999 (B) 4 (C) 7 (D) 14 Page 3 12. What is the number of O2 molecules in the 2.5 g of O 2 inhaled by the average person in one minute? (A) 1.9 × 1022 (B) 3.8 × 1022 (C) 4.7 × 1022 (D) 9.4 × 1022 (A) –103 J B A C P, atm 13. Which point in the phase diagram best represents supercritical conditions? 19. What is the change in internal energy, ∆E, for a reaction that gives off 65 joules of heat and does 38 joules of work? (B) B (A) D (C) C (B) (C) (A) 1 only (B) 2 only (C) 1 and 3 only (D) 1, 2 and 3 15. What is the maximum number of phases that can be in equilibrium in a one component system? (A) 1 (B) 2 (C) 3 (D) (D) D 14. The vapor pressure of a liquid in a closed container depends on 1. temperature of the liquid 2. quantity of liquid 3. surface area of the liquid (D) 4 (B) 20 g·mol–1 (C) 150 g·mol–1 (D) 190 g·mol–1 (A) CH3OCH3 (B) C 2H5OH (C) CH3CH(OH)CH3 (D) CH2(OH)CH2OH 18. 3N2O(g) + 2NH3(g) → 4N2(g) + 3H2O(g) ∆H = –879.6 kJ What is ∆Hf˚ for N2O in Heats of Formation kJ·mol–1? NH3 –45.9 kJ·mol–1 H2O –241.8 kJ·mol–1 (A) +246 (B) +82 (C) –82 – – + + (D) –246 – + + – (A) 1.3 × 10–7 (B) 2.0 × 10–7 (C) 3.0 × 10–7 (D) 4.5 × 10–7 22. Which statements are true? 1. S˚ values for all elements in their standard states are positive. 2. S˚ values for all aqueous ions are positive. 3. ∆S˚ values for all spontaneous reactions are positive. (A) 1 only (B) 1 and 2 only (C) 2 and 3 only (D) 1, 2 and 3 Ag+(aq) + 2 NH3(aq) ¾ Ag(NH3)2+(aq) 23. 17. Which substance would be expected to exhibit the greatest surface tension at 25 ˚C? (D) +103 J 21. The rate of formation of O3(g) is 2.0 × 10–7 mol·L–1·s–1 for the reaction 3O2(g) → 2O3(g) What is the rate of disappearance of O2(g) in mol·L–1·s–1? 16. The molar mass of a gas with a density of 5.8 g·L–1 at 25 ˚C and 740 mm Hg is closest to (A) 10 g·mol–1 (C) +27 J 20. What are the signs of ∆H and ∆S for this reaction? 2C (s) + O2(g) → 2CO(g) ∆H ∆S T, °C (A) A (B) –27 J For this reaction, K = 1.7 × 107 at 25 ˚C. What is the value of ∆G˚ in kJ? (A) –41.2 (B) –17.9 (C) +17.9 (D) +41.2 24. The value of ∆H for a reaction can be found by appropriate combination of bond enthalpies (the energy required to break a particular bond, represented BE). Which expression will give ∆H for this reaction? C 2H4(g) + H2(g) → C2H6(g) (A) BEC=C + BEH–H – [BEC–C + 2BEC–H] (B) BEC–C + 2BEC–H – [BEC=C + BEH–H] (C) 1 /2BEC=C + BEH–H – 2BEC–H (D) 2BEC–H – 1/2BEC=C + BEH–H Page 4 Not valid for use as an USNCO National Examination after April 26, 1999 25. What is the sign of ∆G˚ and the value of K for an o electrochemical cell for which Ecell = 0.80 V? (A) (B) (C) (D) K Phosphorus reacts with chlorine as shown. What is the equilibrium constant expression, K p, for this reaction? – + + – >1 (A) 4 PPCl (B) 3 6 PPCl ⋅ PCl 3 >1 (C) <1 PPCl3 4 PPCl3 6 PCl2 2 (D) PP4 ⋅ PCl6 2 4 PPCl 3 PCl6 2 <1 (B) 4 (C) 6 (D) 8 27. The decomposition of ethane into two methyl radicals has a first order rate constant of 5.5 × 10–4 sec–1 at 700 ˚C. What is the half-life for this decomposition in minutes? (A) 9.1 P 4(s) + 6Cl2(g) ¾ 4PCl 3(g) ∆ G˚ 26. The reaction between NO(g) and O2(g) to give NO2(g) is second order in NO (g) and first order in O2(g). By what factor will the reaction rate change if the concentrations of both reactants are doubled? (A) 2 31. (B) 15 (C) 21 (D) 30 32. The equilibrium constant for the reaction N2O4(g) ¾ 2NO2(g) is 6.10 × 10–3 at 25˚C. Calculate the value of K for this reaction: NO2(g) ¾ 1 /2N2O4(g) (A) 327 (B) 164 (C) 12.8 (D) 3.05 × 10–3 33. The ion-product constant for water at 45 ˚C is 4.0 × 10–14. What is the pH of pure water at this temperature? (A) 6.7 28. The dependence of the rate constant of a reaction on temperature is given by the equation k = e – E kT . Under what conditions is k the smallest? a (A) high T and large Ea (B) high T and small Ea (C) low T and large Ea (D) low T and small Ea (B) 7.0 (C) 7.3 (D) 13.4 34. The position of equilibrium lies to the right in each of these reactions. N2H5+ + NH3 ¾ NH4+ + N2H4 NH3 + HBr ¾ NH4+ + Br – N2H4 + HBr ¾ N2H5+ + Br– 29. The reaction CHCl3(g) + Cl2(g) → CCl 4(g) + HCl(g) is believed to proceed by this mechanism: Cl2(g) → 2Cl(g) fast Cl(g) + CHCl3(g) → HCl(g) + CCl3(g) slow CCl3(g) + Cl(g) → CCl 4(g) fast What rate equation is consistent with this mechanism? (A) Rate = k[Cl2] Based on this information, what is the order of acid strength? (A) HBr > N2H5+ > NH4+ (B) N2H5+ > N2H4 > NH4+ (C) NH3 > N2H4 > Br– (D) N2H5+ > HBr > NH4+ 35. HCN is a weak acid (K a = 6.2 × 10–10). NH3 is a weak base (K b = 1.8 × 10–5). A 1.0 M solution of NH4CN would be (B) Rate = k[Cl][CHCl3] (A) strongly acidic (B) weakly acidic (C) Rate = k[Cl2][CHCl3] (C) neutral (D) weakly basic (D) Rate = k[Cl2]1/2[CHCl 3] 30. The activation energy of a certain reaction is 87 kJ·mol–1. What is the ratio of the rate constants for this reaction when the temperature is decreased from 37 ˚C to 15 ˚C? (A) 5/1 (B) 8.3/1 (C) 13/1 36. What is the percent ionization of a 0.010 M HCN solution? (Ka = 6.2 × 10–10) (A) 0.0025% (B) 0.025% (C) 0.25% (D) 2.5% (D) 24/1 Not valid for use as an USNCO National Examination after April 25, 1999 Page 5 37. How many moles of HCOONa must be added to 1.0 L of 0.10 M HCOOH to prepare a buffer solution with a pH of 3.4? (HCOOH K a = 2 × 10–4) (A) 0.01 (B) 0.05 (C) 0.1 (D) 0.2 38. The acid–base indicator methyl red has a Ka of 1 × 10–5. Its acidic form is red while its alkaline form is yellow. If methyl red is added to a colorless solution with a pH = 7, the color will be (A) pink (B) red (C) orange (Ksp = 5.0 × 10–13) (B) AgCl (Ksp = 1.8 × 10–10) (C) Ag2CO3 (Ksp = 8.1 × 10–12) (D) Ag3AsO4 (Ksp = 1.0 × 10 ) 43. Which expression gives the value for ∆G˚ in kJ·mol–1 for this reaction at 25 ˚C? (A) –6 × 8.31 × 0.43 × 1000 (B) −6 × 96500 × 0.43 × 1000 8.31 (C) −6 × 96500 × 0.43 1000 (D) −6 × 8.31 × 0.43 1000 (D) yellow 39. Silver ions are added to a solution with [Br–] = [Cl–] = [CO32–] = [AsO43–] = 0.1 M. Which compound will precipitate at the lowest [Ag+]? (A) AgBr Questions 43. and 44. should be answered with reference to the reaction. 2Cr(s) + 3Cu2+(aq) → 2Cr3+(aq) + 3Cu(s) E˚ = 0.43 V 44. What is the voltage for this cell when [Cu2+] = 1.0 M and [Cr3+] = 0.010 M? (A) 1.2 (B) 0.87 (C) 0.47 (D) 0.39 –22 40. Consider a voltaic cell based on these half–cells. Ag+(aq) + e– → Ag(s) E˚ = +0.80 V 2+ – Cd (aq) + 2e → Cd (s) E˚ = –0.40 V Identify the anode and give the voltage of this cell under standard conditions. (A) Ag; Ecell = 0.40 V (B) Ag; Ecell = 2.00 V (C) Cd; Ecell = 1.20 V (D) Cd; Ecell = 2.00 V 45. All of these sets of quantum numbers are permissible except n l ml ms (A) 1 0 0 + 1 /2 (B) 2 2 0 –1 /2 (C) 3 1 1 –1 /2 (D) 3 2 –1 + 1 /2 46. Which element can exhibit more than one oxidation state in compounds? 1. Cr 2. Pb 3. Sr 41. Which two species react spontaneously? (A) Cu (s) + Ag+(aq) (B) Br2(l) + Cl–(aq) (A) 1 only (B) 1 and 2 only (C) H2O(l) + Ca2+(aq) (D) Au(s) + Mg2+(aq) (C) 2 and 3 only (D) 1, 2 and 3 42. When aluminum oxide is electrolyzed in the industrial process for the production of aluminum metal, aluminum is produced at one electrode and oxygen gas is produced at the other. For a given quantity of electricity, what is the ratio of moles of aluminum to moles of oxygen gas? (A) 1:1 (B) 2:1 (C) 2:3 (D) 4:3 47. When the isoelectronic species, K+, Ca2+, and Cl–, are arranged in order of increasing radius, what is the correct order? (A) K+, Ca2+, Cl– (B) K+, Cl–, Ca2+ (C) Cl–, Ca2+,K+ (D) Ca2+, K+, Cl– 48. Which Group 2 element has chemical properties least like the other members of the group? (A) Be Page 6 (B) Ca (C) Sr (D) Ba Not valid for use as an USNCO National Examination after April 26, 1999 49. In the vapor state which atom has the largest ionization energy? (A) Na (B) K (C) Mg (D) Ca 50. All of these species have the same number of valence electrons as NO3– except (A) CO32– (B) HCO3– (C) NF3 (D) SO3 51. Which set contains no ionic species? (A) NH4Cl, OF2, H2S (B) CO2, Cl2, CCl4 (C) BF3, AlF3, TlF3 (D) I2, CaO, CH3Cl 56. How many carbon–carbon bonds are in a molecule of 2-methyl-2-butanol? (A) 2 (B) 3 (C) 4 (D) 5 57. Which molecule can exist as stereoisomers? (A) CHF=CHF (B) F 2C=CCl2 (C) CH2F–CHF2 (D) CF3–CH3 58. What are the most likely products in the reaction between CH3CH2CH2OH and HI? (A) CH3CH2CH2I and H2O (B) CH3CH2CH3 and HOI 52. When these species are arranged in order of increasing bond energy, what is the correct sequence? (A) N2, O2, F2 (B) F 2, O2, N2 (C) O2, F2, N2 (D) O2, N2, F2 (C) CH3OH and CH3CH2I (D) ICH2CH2CH2OH and H2 59. Addition polymers include 1. polyamide 2. polyethylene 53. The geometry of the atoms in the species PCl 4+ is best described as (A) tetrahedral (B) see–saw (C) square (D) trigonal bipyramidal (A) 1 only (B) 2 only (C) 2 and 3 only (D) 1, 2 and 3 60. All of these are aromatic compounds except (A) hexene, C6H12 54. Which are nonpolar molecules? 1. NCl3 2. SO3 3. PCl 5 (B) toluene, C6H5CH3 (A) 1 only (B) 2 only (C) p-dichlorobenzene, C6H4Cl2 (C) 1 and 3 only (D) 2 and 3 only (D) naphthalene, C10H8 55. What are the hybridizations of carbon 1 and carbon 2 in the hydrocarbon? 3. polyester CH3CHCH2 1 (A) sp3, sp (B) sp3, sp2 (C) sp2, sp2 (D) sp, sp2 2 END OF TEST Not valid for use as an USNCO National Examination after April 25, 1999 Page 7 Page 8 Not valid for use as an USNCO National Examination after April 26, 1999 US National Chemistry Olympiad – 1999 National Examination—Part I SCORING KEY Number 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. Answer D D B D D A C B B D D C B A C C D B A B Number 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. 31. 32. 33. 34. 35. 36. 37. 38. 39. 40. Answer C A A A A D C C D C D C A A D B B D A C Property of the ACS Society Committee on Education Number 41. 42. 43. 44. 45. 46. 47. 48. 49. 50. 51. 52. 53. 54. 55. 56. 57. 58. 59. 60. Answer A D C C B B D A C C B B A D B C A A B A 1999 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART III Prepared by the American Chemical Society Olympiad Laboratory Practical Task Force OLYMPIAD LABORATORY PRACTICAL TASK FORCE Lucy Pryde Eubanks, Clemson University, Clemson, SC Chair Robert Becker, Kirkwood High School, Kirkwood, MO Craig W. Bowen, Clemson University, Clemson, SC J. Emory Howell, University of Southern Mississippi, Hattiesburg, MS Sheldon L. Knoespel, Michigan State University, East Lansing, MI Jim Schmitt, Eau Claire North High School, Eau Claire, WI Robert G. Silberman, SUNY-Cortland, NY Christie B. Summerlin, University of Alabama-Birmingham, Birmingham, AL DIRECTIONS TO THE EXAMINER–PART III The laboratory practical part of the National Olympiad Examination is designed to test skills related to the laboratory. Because the format of this part of the test is quite different from the first two parts, there is a separate, detailed set of instructions for the examiner. This gives explicit directions for setting up and administering the laboratory practical. There are two laboratory tasks to be completed during the 75 minutes allotted to this part of the test. Students do not need to stop between tasks, but are responsible for using the time in the best way possible. Each procedure must be approved for safety by the examiner before the student begins that procedure. Part III 2 questions laboratory practical 1 hour, 15 minutes A periodic table is provided on page 8 for reference. Students should be permitted to use non-programmable calculators. DIRECTIONS TO THE EXAMINEE–PART II DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. WHEN DIRECTED, TURN TO PAGE 2 AND READ THE DETAILED DIRECTIONS CAREFULLY BEFORE YOU PROCEED. NOTE THE PERIODIC TABLE ON PAGE 8. There are two laboratory-related tasks for you to complete during the next 75 minutes. There is no need to stop between tasks or to do them in the given order. Simply proceed at your own pace from one to the other, using your time productively. You are required to have a procedure for each problem approved for safety by an examiner before you carry out any experimentation on that problem. You are permitted to use a non-programmable calculator. At the end of the 75 minutes, all answer sheets should be turned in. Be sure that you have filled in all the required information at the top of each answer sheet. Carefully follow all directions from your examiner for the proper disposal of chemicals at your examining site. Not valid for use as an USNCO National Examination after April 26, 1999 Distributed by the ACS DivCHED Examinations Institute, Clemson University, Clemson, SC 1999 UNITED STATES NATIONAL CHEMISTRY OLYMPIAD PART III — LABORATORY PRACTICAL Student Instructions Introduction These problems test your ability to design and carry out laboratory experiments and to draw conclusions from your experimental work. You will be graded on your experimental design, on your skills in data collection, and on the accuracy and precision of your results. Clarity of thinking and communication are also components of successful solutions to these problems, so make your written responses as clear and concise as possible. Safety Considerations You are required to wear approved eye protection at all times during this laboratory practical. You also must follow all directions given by your examiner for dealing with spills and with disposal of wastes. Lab Problem 1 You have been given a sample of 7-Up® that has been allowed to stand open. You also have been provided with some table sugar (sucrose), distilled or deionized water, some measuring devices, and a variety of containers. Graph paper has been provided on page 5 of this test booklet. Devise and carry out an experiment to determine the percent by mass of sugar in a sample of 7-Up. You will be asked to describe the method you developed to solve this problem. Given: The molar mass of sucrose, C12H22O11, is 342.30 g·mol–1 . Lab Problem 2 You have been given a sample of Crystal Drano®. There are two components in the Drano – some small shiny metallic pieces, and some pale green beads. (The green color is a dessicating substance.) The metallic pieces are either zinc, magnesium, or aluminum. The beads are either NaOH, Ca(OH)2, or Al(OH)3. You also have 1.0 M NaOH, 3.0 M HCl, and some phenolphthalein indicator. Devise and carry out an experiment to identify both components of Crystal Drano. You will be asked to describe the method you developed to solve this problem. Special Safety Consideration: Crystal Drano is quite caustic and must only be handled with the scoops or spatulas provided. Also, Crystal Drano will readily absorb moisture from the air so only open the container when you need a sample. Recap the container as quickly as possible. Page 2 Answer Sheet for Laboratory Practical Problem 1 Student's Name: __________________________________________________________________________ Student's School:________________________________________ Date: ___________________________ Proctor's Name: _________________________________________________________________________ ACS Section Name :________________________________Student's USNCO test #: ________________ 1. Give a brief description of your experimental plan. List the equipment and materials you plan to use and the steps you plan to take to solve this problem. Before beginning your experiment, you must get approval (for safety reasons) from the examiner. 2. Record your data and other observations. Page 3 Examiner’s Initials: 3. Calculate the percent by mass of sugar in 7-Up. You may choose to use the graph paper on the next page. Show your methods clearly. Percent by mass sucrose in 7-Up® 4. Explain what assumptions were made in determining the mass percent of sugar in 7-Up. How does each assumption influence the mass percent you calculated? Page 4 Graph Paper for Possible Use with Laboratory Practical Problem 1 Page 5 Answer Sheet for Laboratory Practical Problem 2 Student's Name: __________________________________________________________________________ Student's School:________________________________________ Date: ___________________________ Proctor's Name: _________________________________________________________________________ ACS Section Name : ________________________________Student's USNCO test #: ________________ 1. Give a brief description of your experimental plan. List the equipment and materials you plan to use and the steps you plan to take to solve this problem. Before beginning your experiment, you must get approval (for safety reasons) from the examiner. 2. Record your data and other observations. Page 6 Examiner’s Initials: 3. Identify the two components of Crystal Drano. Support your choices with conclusions drawn from your observations. Identification of metallic pieces: ____________________________________ Identification of beads: ____________________________________ Page 7 PERIODIC TABLE OF THE ELEMENTS 1 H 2 He 1.008 4.003 3 Li 4 Be 5 B 6 C 7 N 8 O 9 F 10 Ne 6.941 9.012 10.81 12.01 14.01 16.00 19.00 20.18 11 Na 12 Mg 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar 22.99 24.31 26.98 28.09 30.97 32.07 35.45 39.95 19 K 20 Ca 21 Sc 22 Ti 23 V 24 Cr 25 Mn 26 Fe 27 Co 28 Ni 29 Cu 30 Zn 31 Ga 32 Ge 33 As 34 Se 35 Br 36 Kr 39.10 40.08 44.96 47.88 50.94 52.00 54.94 55.85 58.93 58.69 63.55 65.39 69.72 72.61 74.92 78.96 79.90 83.80 37 Rb 38 Sr 39 Y 40 Zr 41 Nb 42 Mo 43 Tc 44 Ru 45 Rh 46 Pd 47 Ag 48 Cd 49 In 50 Sn 51 Sb 52 Te 53 I 54 Xe 85.47 87.62 88.91 91.22 92.91 95.94 (98) 101.1 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.3 55 Cs 56 Ba 57 La 72 Hf 73 Ta 74 W 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 132.9 137.3 138.9 178.5 181.0 183.8 186.2 190.2 192.2 195.1 197.0 200.6 204.4 207.2 209.0 (209) (210) (222) 87 Fr 88 Ra 89 Ac 104 Rf 105 Db 106 Sg 107 Bh 108 Hs 109 Mt 110 111 112 114 (223) 226.0 227.0 (261) (262) (263) (262) (265) (266) (269) (272) (277) (289) 58 Ce 59 Pr 60 Nd 61 Pm 62 Sm 63 Eu 64 Gd 65 Tb 66 Dy 67 Ho 68 Er 69 Tm 70 Yb 71 Lu 140.1 140.9 144.2 (145) 150.4 152.0 157.3 158.9 162.5 164.9 167.3 168.9 173.0 175.0 90 Th 91 Pa 92 U 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr 232.0 231.0 238.0 237.0 (244) (243) (247) (247) (251) (252) (257) (258) (259) (260) Page 8 1999 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART III Prepared by the American Chemical Society Olympiad Laboratory Practical Task Force ANSWER KEYS Lab Problem 1 5 pts Plan is expected to include: • method to measure density of 7-Up® • method for determining the composition of sugar in water. • method to compare 7-Up solution to the known sugar/water mixtures. • replications of values of data expected. 8 pts Analysis of sugar/water solutions is expected to include: • data to support the method chosen for determining the composition of the sugar/water solutions. • multiple samples. • volume changes. • well-organized data tables . 6 pts Analysis of 7-Up solution is expected to include: • data to support the method chosen for determining the composition of the 7-Up. • replications of samples. 4 pts Results are expected to include: • clear explanations of calculations used to evaluate data. • reasonable results. Example of a reasonable approach (other approaches were evaluated as possibly acceptable): • Determine mass and volume of samples of 7-Up from which an average density value can be determined. • Determine mass and volume of a wide variety of sugar/water mixtures. Use these data to determine the density and corresponding mass % solute values of these mixtures. • Determine mass a volume of a wide variety of sugar/water mixtures. Use these data to determine the density and corresponding mass % solute for sugar/water mixtures. • Having the calculated average density value for 7_Up, use the graph to determine the mass % of 7-Up. • Data points on graph should bracket the 7-Up value. • Graph paper should be used efficiently. 2 pts Discussion of assumptions made should include: • recognition that other components may be present in the 7-Up other than sucrose. • limitations inherent in the methodology and equipment. Page 9 Lab Problem 2 5 pts Plan is expected to include: • method to separate the metal and green beads in Crystal Draino®. • method for testing the metal with water, HCl, and NaOH. • method to testing the green beads with water. • replications of tests. 8 pts Observations are expected to include that: • the green beads dissolve in water. • the metal turns black in water. • the metal reacts with both HCl and NaoH, forming bubbles. • well-organized observation tables . 12 pts Identification and support section is expected to include: • Reasoning for selecting NaOH. • The green beads dissolved easily in water, so the beads must be NaOH. Neither Ca(OH)2 nor Al(OH)3 are soluble in water. • Using 3 M HCl, a titration can be done to determine the amount of hydroxide in the beads. • Reasoning for selecting Al • All three possible metals would react with 3M HCl, so qualitative observations of reaction with HCl are not definitive. • Only Zn and Al react with 1 M NaOH. Because the metal is observed to react with NaOH, the Mg can be eliminated. • To distinguish between Zn and Al, a quantitative titration can be done with 1 M NaOH or with the 3M HCl. Page 10 2000 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART I Prepared by the American Chemical Society Olympiad Examinations Task Force OLYMPIAD EXAMINATIONS TASK FORCE Arden P. Zipp, State University of New York, Cortland Chair Peter E. Demmin (retired), Amherst Central High School, NY Edward DeVillafranca (retired), Kent School, CT Alice Johnsen, Bellaire High School, TX John A. Krikau (retired), Lyons Township High School, IL Patricia A. Metz, University of Georgia, GA Jerry D. Mullins, Plano Senior High School, TX Ronald O. Ragsdale, University of Utah, UT Diane D. Wolff, Western Virginia Community College, VA DIRECTIONS TO THE EXAMINER–PART I Part I of this test is designed to be taken with a Scantron® answer sheet on which the student records his or her responses. Only this Scantron sheet is graded for a score on Part I. Testing materials, scratch paper, and the Scantron sheet should be made available to the student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until April 16, 2000, after which tests can be returned to students and their teachers for further study. Allow time for the student to read the directions, ask questions, and fill in the requested information on the Scantron sheet. The answer sheet must be completed using a pencil, not pen. When the student has completed Part I, or after one hour and thirty minutes has elapsed, the student must turn in the Scantron sheet, Part I of the testing materials, and all scratch paper. There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and you are free to schedule rest-breaks between parts. Part I Part II Part III 60 questions 8 questions 2 lab problems single-answer multiple-choice problem-solving, explanations laboratory practical 1 hour, 30 minutes 1 hour, 45 minutes 1 hour, 30 minutes A periodic table and other useful information are provided on page 2 for student reference. Students should be permitted to use nonprogrammable calculators. DIRECTIONS TO THE EXAMINEE–PART I DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Answers to questions in Part I must be entered on a Scantron answer sheet to be scored. Be sure to write your name on the answer sheet; an ID number is already entered for you. Make a record of this ID number as you will use the same number on both Parts II and III. Each item in Part I consists of a question or an incomplete statement which is followed by four possible choices. Select the single choice that best answers the question or completes the statement. Then use a pencil to blacken the space on your answer sheet having the same letter as your choice. You may write on the examination, but the test booklet will not be used for grading. Scores are based on the number of correct responses. When you complete Part I (or at the end of one hour and 30 minutes), you must turn in all testing materials, scratch paper, and your Scantron answer sheet. Do not forget to turn in your U.S. citizenship statement before leaving the testing site today. Not valid for use as an USNCO National Exam after April 16, 2000. Distributed by the ACS DivCHED Examinations Institute, Clemson University, Clemson, SC. All rights reserved. Printed in U.S.A. amount of substance ampere atmosphere atomic mass unit atomic molar mass Avogadro constant Celsius temperature centi- prefix coulomb electromotive force energy of activation enthalpy entropy ABBREVIATIONS AND SYMBOLS n equilibrium constant K milli- prefix A Faraday constant F molal atm formula molar mass M molar u free energy G mole A frequency ν Planck’s constant N A gas constant R pressure °C gram g rate constant c hour h retardation factor C joule J second E kelvin K speed of light Ea kilo- prefix k temperature, K H liter L time S measure of pressure mmHg volt CONSTANTS m m M mol h P k Rf s c T t V R = 8.314 J·mol–1·K–1 R = 0.0821 L·atm·mol –1·K–1 1 F = 96,500 C·mol–1 1 F = 96,500 J·V–1·mol–1 N A = 6.022 × 1023 mol–1 h = 6.626 × 10–34 J·s c = 2.998 × 108 m·s–1 PERIODIC TABLE OF THE ELEMENTS 1 H 2 He 1.008 4.003 3 Li 4 Be 5 B 6 C 7 N 8 O 9 F 10 Ne 6.941 9.012 10.81 12.01 14.01 16.00 19.00 20.18 11 Na 12 Mg 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar 22.99 24.31 26.98 28.09 30.97 32.07 35.45 39.95 19 K 20 Ca 21 Sc 22 Ti 23 V 24 Cr 25 Mn 26 Fe 27 Co 28 Ni 29 Cu 30 Zn 31 Ga 32 Ge 33 As 34 Se 35 Br 36 Kr 39.10 40.08 44.96 47.88 50.94 52.00 54.94 55.85 58.93 58.69 63.55 65.39 69.72 72.61 74.92 78.96 79.90 83.80 37 Rb 38 Sr 39 Y 40 Zr 41 Nb 42 Mo 43 Tc 44 Ru 45 Rh 46 Pd 47 Ag 48 Cd 49 In 50 Sn 51 Sb 52 Te 53 I 54 Xe 85.47 87.62 88.91 91.22 92.91 95.94 (98) 101.1 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.3 55 Cs 56 Ba 57 La 72 Hf 73 Ta 74 W 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 132.9 137.3 138.9 178.5 181.0 183.8 186.2 190.2 192.2 195.1 197.0 200.6 204.4 207.2 209.0 (209) (210) (222) 87 Fr 88 Ra 89 Ac 104 Rf 105 Db 106 Sg 107 Bh 108 Hs 109 Mt 110 111 112 (223) 226.0 227.0 (261) (262) (263) (262) (265) (266) (269) (272) (277) Page 2 58 Ce 59 Pr 60 Nd 61 Pm 62 Sm 63 Eu 64 Gd 65 Tb 66 Dy 67 Ho 68 Er 69 Tm 70 Yb 71 Lu 140.1 140.9 144.2 (145) 150.4 152.0 157.3 158.9 162.5 164.9 167.3 168.9 173.0 175.0 90 Th 91 Pa 92 U 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr 232.0 231.0 238.0 237.0 (244) (243) (247) (247) (251) (252) (257) (258) (259) (260) Not valid for use as a USNCO National Exam after April 16, 2000. DIRECTIONS When you have selected your answer to each question, blacken the corresponding space on the answer sheet using a soft, #2 pencil. Make a heavy, full mark, but no stray marks. If you decide to change an answer, erase the unwanted mark very carefully. Make no marks on the test booklet. Do all calculations on scratch paper provided by your instructor. There is only one correct answer to each question. Any questions for which more than one response has been blackened will not be counted. Your score is based solely on the number of questions you answer correctly. It is to your advantage to answer every question. (A) I only (B) III only (C) I and II only (D) I, II, and III 2. Which substance is stored in contact with water to prevent it from reacting with air? (A) bromine (B) lithium (C) mercury (D) phosphorus 6. The molarity of a Cu2+ solution is to be determined from its absorbance, measured under the same conditions as those used to prepare this calibration curve. What will be the percent uncertainty in the concentration of a 0.050 M solution if the uncertainty in the absorbance reading is ±0.01 absorbance units? (A) 5% 3. A solution of concentrated aqueous ammonia is added dropwise to 1 mL of a dilute aqueous solution of copper(II) nitrate until a total of 1 mL of the ammonia solution has been added. What observations can be made during this process? (A) The colorless copper(II) nitrate solution turns blue and yields a dark blue precipitate. (B) The colorless copper(II) nitrate solution yields a white precipitate which turns dark blue upon standing. (C) The light blue copper(II) nitrate solution yields a precipitate which redissolves to form a dark blue solution. (D) The light blue copper(II) nitrate solution turns dark blue and yields a dark blue precipitate. 4. What gas is produced when dilute HNO3 is added to silver metal? (A) NO (B) H2 (C) NH3 (D) N2 5. A substance is analyzed by paper chromatography, giving the chromatogram shown. start solvent front 0.0 2.0 4.0 6.0 8.0 10.0 (B) 10% 0.40 Absorbance 1. Which of these ions is expected to be colored in aqueous solution? I Fe3+ II Ni2+ III Al3+ 0.30 0.20 0.10 0 0 0.05 0.10 0.15 [Cu2+], M (C) 15% (D) 20% 7. A 1.50 g sample of an ore containing silver was dissolved, and all of the Ag+ was converted to 0.124 g of Ag2S. What was the percentage of silver in the ore? (A) 6.41% (B) 7.20% (C) 8.27% (D) 10.8% 8. Methyl-t-butyl ether, C5H12O, is added to gasoline to promote cleaner burning. How many moles of oxygen gas, O 2, are required to burn 1.0 mol of this compound completely to form carbon dioxide and water? (A) 4.5 mol (B) 6.0 mol (C) 7.5 mol (D) 8.0 mol 9. A 0.200 g sample of benzoic acid, C6H5COOH, Substance Molar Mass is titrated with a 0.120 M C 6H5COOH 122.1 g·mol–1 Ba(OH)2 solution. What volume of the Ba(OH)2 solution is required to reach the equivalence point? (A) 6.82 mL (B) 13.6 mL (C) 17.6 mL (D) 35.2 mL 12.0 cm What is the Rf value of the substance represented by the spot at 8.0 cm? (A) 0.80 (B) 0.75 (C) 0.67 (D) 0.60 Not valid for use as an USNCO National Examination after April 16, 2000. Page 3 (A) 5.15 g (B) 14.3 g (C) 19.4 g (D) 26.4 g 11. What is the Na+ ion concentration in the solution formed by mixing 20. mL of 0.10 M Na2SO4 solution with 50. mL of 0.30 M Na3PO4 solution? (A) 0.15 M (B) 0.24 M (C) 0.48 M (D) 0.70 M 12. A solution prepared by Compound Kb dissolving a 2.50 g sample C 6H6 2.53 °C·m–1 of an unknown compound dissolved in 34.0 g of benzene, C6H6, boils 1.38 °C higher than pure benzene. Which expression gives the molar mass of the unknown compound? 2.50 (A) 2.53 × 1.38 (B) 1.38 × 2.53 1 × 34.0 1.38 3 (D) 2.50 × 10 × 1.38 × 2.53 34.0 740 mmHg + 20 mmHg 740 mmHg (C) 300 mL × 740 mmHg 740 mmHg – 20 mmHg (D) 300 mL × 740 mmHg 740 mmHg + 20 mmHg 17. What is the normal melting point of the substance represented by the phase diagram? 1.0 A B C T, °C (A) A 13. What is the total pressure in a 2.00 L container that holds 1.00 g He, 14.0 g CO, and 10.0 g of NO at 27.0 °C? (A) 21.6 atm (B) 13.2 atm (C) 1.24 atm (D) 0.310 atm 14. What type of solid is generally characterized by having low melting point and low electrical conductivity? (A) ionic (B) metallic (C) molecular (D) network covalent 15. How many nearest neighbors surround each particle in a face-centered cubic lattice? Page 4 (B) 300 mL × 34.0 × 2.50 2.53 3 (C) 2.50 × 10 × (A) 4 16. Hydrogen is Compound Vapor Pressure collected over water at 22 °C at 22 °C and a H2O 20. mmHg barometer reading of 740 mmHg. If 300. mL of hydrogen is collected, which expression will give the volume of dry hydrogen at the same temperature and pressure? 740 mmHg – 20 mmHg (A) 300 mL × 740 mmHg P,atm 10. Chlorine can be prepared by reacting HCl with MnO2. The reaction is represented by this equation. MnO2(s) + 4HCl(aq) → Cl2(g) + MnCl2(aq) + 2H2O(l) Assuming the reaction goes to completion what mass of concentrated HCl solution (36.0% HCl by mass) is needed to produce 2.50 g of Cl 2? (B) 6 (C) 8 (D) 12 (B) B (C) C D (D) D 18. A bomb calorimeter has a heat capacity of 783 J·°C–1 and contains 254 g of water, which has a specific heat of 4.184 J·g–1·°C–1. How much heat is evolved or absorbed by a reaction when the temperature goes from 23.73 °C to 26.01 °C? (A) 1.78 kJ absorbed (B) 2.42 kJ absorbed (C) 1.78 kJ evolved (D) 4.21 kJ evolved 19. Consider this equation and the associated value for ∆Ho. 2H2(g) + 2Cl2(g) → 4HCl(g) ∆Ho = –92.3 kJ Which statement about this information is incorrect? (A) If the equation is reversed, the ∆Ho value equals +92.3 kJ. (B) The four HCl bonds are stronger than the four bonds in H2 and Cl2. (C) The ∆Ho value will be –92.3 kJ if the HCl is produced as a liquid. (D) 23.1 kJ of heat will be evolved when 1 mol of HCl (g) is produced. Not valid for use as an USNCO National Examination after April 16, 2000. (A) –1074.0 kJ (B) –22.2 kJ (C) +249.8 kJ (D) +2214.6 kJ 25. A reaction follows this concentration-time diagram. The instantaneous rate for this reaction at 20 seconds will be closest to which value? 0.40 Molarity 20. Determine the heat of reaction for this process. FeO(s) + Fe2O3(s) → Fe3O4(s) Given information: 2Fe(s) + O2(g) → 2FeO(s) ∆Ho = –544.0 kJ 4Fe(s) + 3O2(g) → 2Fe2O3(s) ∆Ho = –1648.4 kJ Fe3O4(s) → 3Fe(s) + 2O2(g) ∆Ho = +1118.4 kJ 0.30 0.20 0.10 0 0 21. For which process will ∆Ho and ∆Go be expected to be most similar? (A) 2Al(s) + Fe2O3(s) → 2Fe(s) + Al2O3(s) (A) 4 × 10–3 M·sec–1 (B) 8 × 10–3 M·sec–1 (C) 2 × 10–2 M·sec–1 (D) 1 × 10–1 M·sec–1 (C) 2NO2(g) → N2O4(g) (A) zero. (B) first. (D) 2H2(g) + O2(g) → 2H2O(g) (C) second. (D) third. Bond Bond Energy H–H O–O O=O H–O 436 142 499 460 (A) –127 kJ (B) –209 kJ (C) –484 kJ (D) –841 kJ 60 26. If the half-life of a reaction increases as the initial concentration of substance increases, the order of the reaction is (B) 2Na (s) + 2H2O(l) → 2NaOH (aq) + H2(g) 22. Use bond energies to estimate ∆H for this reaction. H2(g) + O2(g) → H2O2(g) 20 40 Time, sec kJ·mol–1 kJ·mol–1 kJ·mol–1 kJ·mol–1 23. For a particular reaction, ∆Ho = –38.3 kJ and ∆So = –113 J·K–1. This reaction is (A) spontaneous at all temperatures. (B) nonspontaneous at all temperatures. 27. The radioisotope N-13, which has a half-life of 10 minutes, is used to image organs in the body. If an injected sample has an activity of 40 microcuries (40 µCi), what is its activity after 25 minutes in the body? (A) 0.75 µCi (B) 3.5 µCi (C) 7.1 µCi (D) 12 µCi 28. Propanone reacts with iodine in acid solution as shown in this equation. H+ CH3C(O)CH3 + I2 → CH3C(O)CH2I + HI These data were obtained when the reaction was studied. [CH3C(O)CH3], M [I2], M [H+], M Relative Rate 0.010 0.020 0.020 0.020 (C) spontaneous at temperatures below 66 °C. (D) spontaneous at temperatures above 66 °C. 24. What is ∆Go for this reaction? 1/2N2(g) + 3/2H2(g) = NH3(g) K p = 4.42 × 104 at 25 °C. 0.010 0.010 0.020 0.010 0.010 0.010 0.010 0.020 1 2 2 4 What is the rate equation for the reaction? (A) rate = k[CH3C(O)CH3] [I2] (A) –26.5 kJ·mol–1 (B) –11.5 kJ·mol–1 (B) rate = k[CH3C(O)CH3]2 (C) –2.2 kJ·mol–1 (D) –0.97 kJ·mol–1 (C) rate = k[CH3C(O)CH3] [I2] [H+] (D) rate = k[CH3C(O)CH3] [H+] 29. A particular reaction rate increases by a factor of five when the temperature is increased from 5 °C to 27 °C. What is the activation energy of the reaction? (A) 6.10 kJ·mol–1 (B) 18.9 kJ·mol–1 (C) 50.7 kJ·mol–1 (D) 157 kJ·mol–1 Not valid for use as an USNCO National Examination after April 16, 2000. Page 5 30. Consider this reaction. 2H2(g) + 2NO(g) → N2(g) + 2H2O(g) The rate law for this reaction is rate = k [H2] [NO]2. Under what conditions could these steps represent the mechanism? Step 1. 2NO = N2O2 Step 2. N 2O2 + H 2 → N 2O + H 2O Step 3. N 2O + H 2 → N 2 + H 2O (A) These steps cannot be the mechanism under any circumstances. (B) These steps could be the mechanism if step 1 is the slow step. (C) These steps could be the mechanism if step 2 is the slow step. (D) These steps could be the mechanism if step 3 is the slow step. 31. A reaction has a forward rate constant of 2.3 × 106 s–1 and an equilibrium constant of 4.0 × 108. What is the rate constant for the reverse reaction? (A) 1.l × 10 –15 (C) 1.7 × 10 s 2 s (B) 5.8 × 10 s –1 –3 –1 (D) 9.2 × 1014 s–1 –1 32. For the reaction 2A(g) + 2B (g) = 3C(g) at a certain temperature, K is 2.5 × 10–2. For which conditions will the reaction proceed to the right at the same temperature? [A], M [B], M [C], M 36. What is the conjugate acid of HPO42–? (A) H3PO4(aq) (B) H2PO4–(aq) (C) H3O+(aq) (D) PO43–(aq) 37. The amount of sodium Acid Ka hydrogen carbonate, H2CO3 2.5 × 10–4 NaHCO3, in an antacid – HCO3 2.4 × 10–8 tablet is to be determined by dissolving the tablet in water and titrating the resulting solution with hydrochloric acid. Which indicator is the most appropriate for this titration? (A) methyl orange, pKin = 3.7 (B) bromothymol blue, pKin = 7.0 (C) phenolphthalein, pKin = 9.3 (D) alizarin yellow, pK in = 12.5 38. How many moles of Acid Ka NaOCl must be added to HOCl 2.8 × 10–8 150 mL of 0.025 M HOCl to obtain a buffer solution with a pH = 7.50? (A) 2.6 × 10–5 (B) 1.1 × 10–3 (C) 3.3 × 10–3 (D) 2.2 × 10–2 39. If equal volumes of BaCl2 Substance K sp and NaF solutions are BaF 2 1.7 × 10–7 mixed, which of these combinations will not give a precipitate? (A) 0.10 0.10 0.10 (B) 1.0 1.0 1.0 (C) 1.0 0.10 0.10 (A) 0.0040 M BaCl2 and 0.020 M NaF (D) 1.0 1.0 0.10 (B) 0.010 M BaCl2 and 0.015 M NaF (C) 0.015 M BaCl2 and 0.010 M NaF – 33. What is the Kb of a weak base that produces one OH per molecule if a 0.050 M solution is 2.5% ionized? (A) 7.8 × 10–8 (B) 1.6 × 10–6 (C) 3.2 × 10–5 (D) 1.2 × 10–3 34. What is the [OH–] of a 0.65 M solution of NaOCl? Acid HOCl Ka 2.8 × 10–8 (D) 0.020 M BaCl2 and 0.0020 M NaF 40. What takes place when zinc metal is added to a aqueous solution containing magnesium nitrate and silver nitrate? 1. Zn is oxidized. 2. Mg2+ is reduced. 3. Ag+ is reduced. 4. No reaction takes place. (A) 4.8 × 10–4 M (B) 1.3 × 10–4 M (A) 1 and 2 only (B) 1 and 3 only (C) 3.5 × 10–7 M (D) 2.1 × 10–11 M (C) 1, 2, and 3 only (D) 4 only 35. Which acid is the strongest? (A) H3BO3 (B) H3PO4 (C) H2SO3 (D) HClO3 Page 6 Not valid for use as an USNCO National Examination after April 16, 2000. Questions 41, 42, and 43 should be answered with reference to this information and diagram. Ag+(aq) + e– → Ag(s) Eo = 0.80 V 2+ – Cu (aq) + 2e → Cu(s) Eo = 0.34 V 45. How many unpaired electrons are in a gaseous Fe2+ ion in the ground state? (A) 0 (B) 2 (C) 4 (D) 6 46. Which element has the smallest first–ionization energy? V (A) Mg (B) Al (C) Si (D) P salt bridge Ag Cu Ag+ (aq) Cu2+ (aq) 47. Which set of orbitals is listed in the sequential order of filling in a many-electron atom? (A) 3s, 3p, 3d (B) 3d, 4s, 4p (C) 3d, 4p, 5s (D) 4p, 4d, 5s 48. Which set is expected to show the smallest difference in first–ionization energy? 41. What is the value for ∆G° when [Ag+] = [Cu2+] = 1.0 M? (A) –44.4 kJ (B) –88.8 kJ (C) –243 kJ (D) –374 kJ 42. Which expression gives the voltage for this cell if [Cu2+] = 1.00 M and [Ag+] = 0.010 M? (A) 0.46 V + 0.0591 V (A) He, Ne, Ar + (B) B, N, O 2+ (C) Mg, Mg , Mg (D) Fe, Co, Ni 49. When the atoms Li, Be, B, and Na are arranged in order of increasing atomic radius, what is the correct order? (A) B, Be, Li, Na (B) Li, Be, B, Na (C) Be, Li, B, Na (D) Be, B, Li, Na (B) 0.46 V + 2 × 0.0591 V 50. Which species has the same shape as the NO3– ion? (C) 0.46 V – 0.0591 V (D) 0.46 V – 2 × 0.0591 V 43. Which increases immediately if the surface area of the silver electrode is increased? (A) overall cell voltage (A) SO3 (B) SO32– (C) ClF3 (D) ClO 3– 51. What is the formal charge on the central atom in N2O? N N O + (B) rate of change of [Ag ] (C) mass of Cu electrode mass of Cu (D) change in ratio of electrode masses; ∆ mass of Ag 44. In the galvanizing process, iron is coated with zinc. The resulting chemical protection is most similar to that provided when (A) a magnesium bar is connected to an iron pipe. (B) an iron can is plated with tin. (A) +1 (B) 0 (C) –1 (D) –2 52. How many bonding pairs and lone pairs surround the central atom in the I 3– ion? Bonding Pairs Lone Pairs (A) 2 2 (B) 2 3 (C) 3 2 (D) 4 3 (C) copper pipes are connected using lead solder. (D) a copper pipe is covered with epoxy paint. Not valid for use as an USNCO National Examination after April 16, 2000. Page 7 53. The nitrogen atoms in NH3, NH2–, and NH4+ are all surrounded by eight electrons. When these three species are arranged in order of increasing H–N–H bond angle, what is the correct order? – + (B) NH4 , NH2 , NH3 + – (D) NH2–, NH3, NH4+ (A) NH3, NH2 , NH4 (C) NH3, NH4 , NH2 + 57. Which is the formula for an alkyne? (A) C 2H4 (B) C 3H6 (C) C 3H8 (D) C 4H6 – 58. How many isomers have the formula C3H8O? (A) 2 54. What hybrid orbitals are employed by carbon atoms 1,2, and 3, respectively, as labeled in the compound shown? 3 O H3C C C N 1 2 3 (A) sp , sp, sp 2 2 (B) sp , sp , sp 3 3 2 2 (C) sp , sp , sp (D) sp , sp , sp (C) 4 (D) 5 59. Which type of organic compound is most resistant to oxidation by acidified potassium dichromate? (A) acid (B) alcohol (C) aldehyde (D) alkene 2 55. In which pair, or pairs, is the stronger bond found in the first species? 1. O2–, O2 2. N2, N2+ 3. NO+, NO– (A) 1 only (B) 2 only (C) 1 and 3 only (D) 2 and 3 only 60. What product, in addition to water, is produced by this reaction? CH3OH + C6H5COOH → O C OH (A) H3C (B) 56. What is the molecular formula of this chemical structure? (B) 3 CH3 O C CH3 (C) CH3 (A) C 10H12 (B) C 10H14 (C) C 12H12 (D) C 12H14 CH3 (D) O C O CH3 END OF TEST Page 8 Not valid for use as an USNCO National Examination after April 16, 2000. US National Chemistry Olympiad – 2000 National Examination—Part I SCORING KEY Number 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. Answer C D C A B B B C A B D C B C D A B D C B Number 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. 31. 32. 33. 34. 35. 36. 37. 38. 39. 40. Answer A A C A A A C D C C B D C A D B A C D B Property of the ACS Society Committee on Education Number 41. 42. 43. 44. 45. 46. 47. 48. 49. 50. 51. 52. 53. 54. 55. 56. 57. 58. 59. 60. Answer B D B A C B C D A A A B D C D C D B A D 2000 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART II Prepared by the American Chemical Society Olympiad Examinations Task Force OLYMPIAD EXAMINATIONS TASK FORCE Arden P. Zipp, State University of New York, Cortland Chair Peter E. Demmin (retired), Amherst Central High School, NY Edward DeVillafranca (retired), Kent School, CT Alice Johnsen, Bellaire High School, TX John A. Krikau (retired), Lyons Township High School, IL Patricia A. Metz, University of Georgia, GA Jerry D. Mullins, Plano Senior High School, TX Ronald O. Ragsdale, University of Utah, UT Diane D. Wolff, Western Virginia Community College, VA DIRECTIONS TO THE EXAMINER–PART II Part II of this test requires that student answers be written in a response booklet of blank pages. Only this “Blue Book” is graded for a score on Part II. Testing materials, scratch paper, and the “Blue Book” should be made available to the student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until April 16, 2000, after which tests can be returned to students and their teachers for further study. Allow time for the student to read the directions, ask questions, and fill in the requested information on the “Blue Book”. When the student has completed Part II, or after one hour and forty-five minutes has elapsed, the student must turn in the “Blue Book”, Part II of the testing materials, and all scratch paper. Be sure that the student has supplied all of the information requested on the front of the “Blue Book,” and that the same identification number used for Part I has been used again for Part II. There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and you are free to schedule rest-breaks between parts. Part I Part II Part III 60 questions 8 questions 2 lab problems single-answer multiple-choice problem-solving, explanations laboratory practical 1 hour, 30 minutes 1 hour, 45 minutes 1 hour, 30 minutes A periodic table and other useful information are provided on the back page for student reference. Students should be permitted to use non-programmable calculators. DIRECTIONS TO THE EXAMINEE–PART II DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Part II requires complete responses to questions involving problem-solving and explanations. One hour and forty-five minutes are allowed to complete this part. Be sure to print your name, the name of your school, and your identification number in the spaces provided on the “Blue Book” cover. (Be sure to use the same identification number that was coded onto your Scantron® sheet for Part I. Answer all of the questions in order, and use both sides of the paper. Do not remove the staple. Use separate sheets for scratch paper and do not attach your scratch paper to this examination. When you complete Part II (or at the end of one hour and forty-five minutes), you must turn in all testing materials, scratch paper, and your “Blue Book.” Do not forget to turn in your U.S. citizenship statement before leaving the testing site today. Not valid for use as an USNCO National Exam after April 16, 2000. Distributed by the ACS DivCHED Examinations Institute, Clemson University, Clemson, SC. All rights reserved. Printed in U.S.A. 1. (12%) An unknown metal, M, reacts with excess chlorine to give the metal chloride, MClx. When 0.396 g of the chloride is dissolved in water and passed through an anion exchange column charged with hydroxide ions, the solution requires 23.55 mL of 0.195 M HCl for neutralization. a. Calculate the number of moles of HCl used in the titration. b. Determine the mass of chlorine and the mass of metal in this sample of MClx. c. Assuming that x in MClx is 1, 2 or 3, calculate possible atomic masses for M. d. Use your knowledge of the Periodic Table to write formulas for the possible compounds between chlorine and metals and identify those expected to be stable. 2. (11%) The ionization constant for water is 1.14 × 10–15 at 0 °C and 9.6 × 10–14 at 60 °C. a. Write the equation for the ionization of water and determine the pH of water at 60 °C. b. Calculate each value. i. ∆Hionization over this temperature range ii. ∆G at 60 °C iii. ∆S at 60 °C c. State the significance of the sign of the sign of ∆S obtained in part 2b.iii, and explain how the process indicated in 2a could lead to this sign. 3. (15%) These are the reaction steps in a certain polymerization process, which may occur by either an uncatalyzed or an acid-catalyzed pathway. O R C O O + HA H k1 C+ R k2 O Group 1 O R A– H O H R C O R' O+ H C+ O H A– + R' O k3 H k4 H Group 2 H A– H O R C R' O+ H O O H k5 R C O R' + H2O + HA A– Group 3 H a. Write a balanced equation for the overall reaction. b. Name the functional groups labeled [1], [2], and [3]. c. Given these data for the acid-catalyzed reaction, find the rate law and the value of k, specifying its units. [RCOOH], M [R´OH], M [HA], M Initial Rate, M·min –1 0.35 0.35 0.50 4.60 0.62 0.35 0.50 8.14 0.35 0.81 0.50 10.6 0.35 0.50 0.75 9.84 d. Identify the rate-determining step based on the rate law found in question 3c. Explain your answer. e. The initial reaction rate can be followed spectrophotometrically by quenching the reaction and determining the amount of ROH left by its reaction with dichromate ion, Cr2O72–. i. Write a balanced equation for the reaction of Cr2O72– with R´OH in acid solution. Assume R´ is CH3CH2– and the products of the reaction are Cr 3+ and CH3COOH. ii. Describe the color change expected for the reaction written in question 3e, part i. Page 2 Not valid for use as an USNCO National Examination after April 16, 2000. 4. (12%) 25.00 mL of a solution of a weak monoprotic acid, HX, was titrated with a 0.0640 M solution of NaOH, requiring 18.22 mL. The pH of the solution varied as a function of the percentage of HX titrated. These data were collected. % titrated 0 33.3% 66.7% pH 3.39 5.14 5.74 a. Calculate the initial concentration of the weak acid in the 25.00 mL of solution. b. Determine the value of K a for two of these three conditions. c. Calculate the pH at the equivalence point of this titration and write an equation to account for this pH. d. Calculate the number of moles of a salt, NaX, that must be added to produce a pH of 6.00 in 150.00 mL of the original solution. 5. (14%) Write net equations for each of these reactions. Use appropriate ionic and molecular formulas for the reactants and products and omit formulas for all ions or molecules that do not take part in a reaction. Write structural formulas for all organic substances. You need not balance the reactions. All reactions occur in aqueous solution unless otherwise indicated. a. Phosphorus is burned in excess oxygen. b. Sulfur dioxide is bubbled into water. c. Chlorine gas is bubbled through a sodium bromide solution. d. Solutions of magnesium nitrate and potassium hydroxide are mixed. e. A sodium thiosulfate solution is added to a suspension of silver chloride. f. Bromine is added to a solution of ethylene in hexane. g. Radium-226 emits an alpha particle. 7. 8. (12%) Use the given phase diagram of water to answer these questions. Note that the axis values are not drawn to scale. a. Identify the physical state at points A, B, C, and D. b. Calculate the volume of one mole of water in each of the phases at the triple point, (At the triple point, the density of H2O(l) is 0.9998 g·mL–1 and the density of H2O(s) is 0.917 g·mL–1.) c. Starting with point A, describe the pressure, temperature, and phase changes that correspond to the rectangle around the triple point. P, mmHg 6. (12%) Nitrogen dioxide, NO2, can undergo reactions to form nitrite ion, NO2– , and nitronium ion, NO2+. a. Draw Lewis structures for NO 2– and NO2+ including any resonance forms. b. Predict the shape of each ion and account for each shape using a modern bonding theory. c. Describe and account for the difference in the N–O bond lengths in NO2– and NO2+. d. Determine the oxidation number and the formal charge of nitrogen in the NO2– ion. Outline your reasoning and state the difference between formal charge and oxidation number. 760 4.58 D A C B 0 0.01 T, °C 100 (12%) The behavior of elements can often be predicted based on their positions in the Periodic Table. Use your knowledge about trends in the behavior of elements to answer the following questions about the recently isolated elements 114, 116 and 118. a. Give the names and symbols of the elements in the row above 114, 116, and 118 in the Periodic Table. b. Predict the relative ionization energies of elements 114, 116, and 118 and describe how the ionization energy of one of them is expected to compare with the ionization energy of the element above it, giving reasons for your answers. c. Predict the oxidation states expected for element 114 and indicate which oxidation state is expected to be most stable, giving reasons for your answers. d. Suggest a reason that elements 114, 116, and 118 have been made, but elements 113, 115, and 117 have not. END OF PART II Not valid for use as an USNCO National Examination after April 16, 2000. Page 3 amount of substance ampere atmosphere atomic mass unit atomic molar mass Avogadro constant Celsius temperature centi- prefix coulomb electromotive force energy of activation enthalpy entropy ABBREVIATIONS AND SYMBOLS n equilibrium constant K milli- prefix A Faraday constant F molal atm formula molar mass M molar u free energy G mole A frequency ν Planck’s constant N A gas constant R pressure °C gram g rate constant c hour h second C joule J speed of light E kelvin K temperature, K Ea kilo- prefix k time H liter L volt S measure of pressure mmHg volume CONSTANTS m m M mol h P k s c T t V V R = 8.314 J·mol–1·K–1 R = 0.0821 L·atm·mol –1·K–1 1 F = 96,500 C·mol–1 1 F = 96,500 J·V–1·mol –1 N A = 6.022 × 1023 mol –1 h = 6.626 × 10–34 J·s c = 2.998 × 108 m·s–1 USEFUL EQUATIONS – ∆H 1 ln K = +c R T RT E=E – ln Q nF ο PERIODIC TABLE OF THE ELEMENTS 1 H 2 He 1.008 4.003 3 Li 4 Be 5 B 6 C 7 N 8 O 9 F 10 Ne 6.941 9.012 10.81 12.01 14.01 16.00 19.00 20.18 11 Na 12 Mg 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar 22.99 24.31 26.98 28.09 30.97 32.07 35.45 39.95 19 K 20 Ca 21 Sc 22 Ti 23 V 24 Cr 25 Mn 26 Fe 27 Co 28 Ni 29 Cu 30 Zn 31 Ga 32 Ge 33 As 34 Se 35 Br 36 Kr 39.10 40.08 44.96 47.88 50.94 52.00 54.94 55.85 58.93 58.69 63.55 65.39 69.72 72.61 74.92 78.96 79.90 83.80 37 Rb 38 Sr 39 Y 40 Zr 41 Nb 42 Mo 43 Tc 44 Ru 45 Rh 46 Pd 47 Ag 48 Cd 49 In 50 Sn 51 Sb 52 Te 53 I 54 Xe 85.47 87.62 88.91 91.22 92.91 95.94 (98) 101.1 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.3 55 Cs 56 Ba 57 La 72 Hf 73 Ta 74 W 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 132.9 137.3 138.9 178.5 181.0 183.8 186.2 190.2 192.2 195.1 197.0 200.6 204.4 207.2 209.0 (209) (210) (222) 87 Fr 88 Ra 89 Ac 104 Rf 105 Db 106 Sg 107 Bh 108 Hs 109 Mt 110 111 112 (223) 226.0 227.0 (261) (262) (263) (262) (265) (266) (269) (272) (277) Page 4 58 Ce 59 Pr 60 Nd 61 Pm 62 Sm 63 Eu 64 Gd 65 Tb 66 Dy 67 Ho 68 Er 69 Tm 70 Yb 71 Lu 140.1 140.9 144.2 (145) 150.4 152.0 157.3 158.9 162.5 164.9 167.3 168.9 173.0 175.0 90 Th 91 Pa 92 U 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr 232.0 231.0 238.0 237.0 (244) (243) (247) (247) (251) (252) (257) (258) (259) (260) Not valid for use as an USNCO National Examination after April 16, 2000. 2000 U. S. NATIONAL CHEMISTRY OLYMPIAD KEY for NATIONAL EXAM—PART II 1. a. b. c. 2. a. 0.195 mol = 0.00459 mol H + = 0.00459 mol Cl – L 35.45 g 0.00459 mol Cl – × = 0.163 g Cl – mol 1 50.8 g 0.396 g MCl - 0.163 g Cl – = 0.233 g × = 0.00459 mol mol 1 1 01.3 g 0.396 g MCl 2 - 0.163 g Cl – = 0.233 g × = 0.00230 mol mol 1 152.3.3 g – 0.396 g MCl 3 - 0.163 g Cl = 0.233 g × = 0.00153 mol mol –1 50.8 g·mol most likely V VCl unlikely to be stable 101.3 g·mol–1 most likely Ru RuCl2 stable 152.3 g·mol–1 most likely Eu EuCl3 stable 0.02355 L HCl × H2O → H+ + OH– [H ] = + b. i. 9.6 × 10 –14 = 3.1 × 10–7; pH = 6.51 k ∆H 1 1 ln 2 = – R T1 T2 k1 9.6 × 10 –14 ∆H 1 1 ln – −15 = 1.14 × 10 8.314 273 333 ln 84.2 = 4.43 (8.314) ∆H = ∆H (0.003663 – 0.003003) ; 8.314 0.000660 ∆H = 5.58 × 10 4 J or 55.8 kJ ii. iii. c. 3. ∆G = – RT ln K = (–8.314)(333) ln(9.6 × 10 –14 ) = 8.30 × 10 4 J 8.30 × 10 4 − 5.58 × 10 4 = –81.7 J –333 The negative sign corresponds to an increase in order. An increase in order can result from the H+ and OH– ions structuring the H2O molecules around them. ∆G = ∆H – T∆S ; ∆S = a. RCOOH + R’OH → RCOOR’ + H2O b. Functional group I – carboxyl group (acid) Functional group II – hydroxyl group (alcohol) Functional group III – ester 8.14 0.62 = 1.77 = 1.77 RCOOH is first order. 4.60 0.35 10.6 0.81 = 2.30 = 2.30 R’OH is first order. 4.60 0.35 c. Trials 1 and 4 can be used to determine the order with respect to HA. [RCOOH] is held constant; however, both [R’OH] and [HA] vary. The effect of [R’OH] is: 0.50 M 6.57 M = 1.43 and then (1.43)( 4.60) = 0.35 M min 9.84 0.75 = 1.50 = 1.50 HA is first order 6.57 0.50 Rate = k[RCOOH] [R’OH] [HA] 4.60 M = k (0.35M) (0.35M) (0.50M) min The effect of [HA] is k = 75 M–2 ·min–1 d. Rate determining step could be either: Step 2 that involves R’OH and product from the reaction of RCOOH and HA, or Step 3 that involves the breakup of product in step 2. e. i. 2Cr2O72– + 3CH3CH2OH + 16H+ → 4Cr3+ + 3CH3COOH + 11H2O The color will change from orange (Cr2O72–(aq)) to green (Cr3+(aq)). ii. 4. a. 0.0640 mol ⋅ L–1 × 18.22 mL = 0.0466 M acid 25.00 mL b. With zero % titrated, pH = 3.39; [H+] = 4.07 × 10–4 Assuming negligible dissociation: [H ][X ] = (4.07 × 10 ) = + Ka Assuming significant dissociation: –4 2 – [HX] = 3.56 × 10 (0.0466) [H ][X ] = (4.07 × 10 ) = 3.59 × 10 = [HX] (0.0466 - 4.07 × 10 ) + –6 Ka –4 2 – –6 –4 With 33.3 % titrated, pH = 5.14; [H+] = 7.24 × 10–6 Assuming negligible dissociation: [H ][X ] = (7.24 × 10 )(0.0155) = + Ka Assuming significant dissociation: [HX] (0.0311) [H ][X ] = (7.24 × 10 )(0.0155) = 3.61 × 10 = [HX] (0.0311 - 7.24 × 10 ) + –6 – = 3.62 × 10 –6 Ka – –6 –6 –6 With 66.7 % titrated, pH = 5.74; [H+] = 1.82 × 10–6 Assuming negligible dissociation: Ka = Assuming significant dissociation: [H ][X ] = (1.82 × 10 )(0.0312) = 3.64 × 10 –6 K a = X– + H2O = HX + OH– Kb = + [HX] c. (0.0155) [H ][X ] = (1.82 × 10 )(0.0312) = 3.66 × 10 [HX] (0.0155 -1.82 × 10 ) + –6 – )(25) = 0.0270 M [X ] = (0.(0466 43.22)) [OH ] ; [OH ] = 8.66 × 10 2.78 × 10 = – –6 –6 K w 1.00 × 10 –14 = = 2.78 × 10 –9 Ka 3.60 × 10 –6 – – 2 –9 – 0.0270 d. pH = 6.0 and [H+] = 1.0 3.60 × 10 –6 = –6 ; pOH = 5.06; pH = 8.94 × 10–6 [ ] [ ] (1.0 × 10 –6 ) X – ; X – = 0.168 M 0.0466 0.168 mol Moles X– = × 0.150 L = 0.0252 mol L –6 5. Note: Balanced equations were not required. a. P4 + O2 → P4O10 b. SO2 + H2O → H2SO3 or H+ + HSO3– c. Cl2 + Br– → Cl– + Br2 d. Mg2+ + OH– → Mg(OH)2 e. AgCl + S2O32– → Ag(S2O3)23– + Cl– f. H C C H H 6. H → H H + Br2 → Br C C Br H H g. 226 88 Ra 4 2 He a. For NO2– b. c. NO2– will be bent due to the lone pair of electrons on N. NO2+ will be linear. Nitrogen-to-oxygen bonds in NO2– will be longer than those in NO2+. The average bond + 222 86 Rn O N O – O N O – and for NO2+ O N O + order in NO2– is 11/2. The bond order in NO2+ is 2. d. The oxidation number of N in NO2– is +3. Formal charge is zero. The oxidation number is obtained by assigning bonding electrons to the more electronegative atom. The formal charge is found by dividing the bonding electrons evenly between atoms. The number of electrons left is compared with the original number. 7. a. b. A = liquid; B = gas, C and D = solid 1 mL 1 mL 18.0 g × = 18.0 mL liquid 18.0 g × = 19.6 mL solid 0.9998 g 0.917 g 3 nRT 1 (0.0821) (273.17) (760) ; V= ; V = 3.72 × 10 L P 4.58 A → B pressure decreases, temperature constant, phase change from liquid to gas PV = nRT ; V = c. B → C pressure constant, temperature decreases, phase change from gas to solid C → D pressure increases, temperature constant, no phase change D → A pressure constant, temperature increases, phase change from solid to liquid 8. a. The element above element 114 is lead, Pb; above element 116 is polonium, Po; above element 118 is radon, Rn. b. The ionization energy is expected to increase from element 114 to 116 to 118. This is due to increasing nuclear charge density thus resulting in greater attraction for outer electrons. The ionization energy of elements 114, 116, and 118 are expected to be lower than those of the elements directly above them due to the presence of electrons in higher energy levels, those with higher values of n. c. The oxidation states for element 114 are predicted to be +2 (loss of electrons in p orbitals) and +4 (loss of electrons in both s and p orbitals).The oxidation state +2 is likely more stable; lower oxidation states are usually more stable. d. Even values of Z are usually more stable due to pairing of protons. Odd values of Z are less stable. 2000 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART III Prepared by the American Chemical Society Olympiad Laboratory Practical Task Force OLYMPIAD LABORATORY PRACTICAL TASK FORCE Lucy Pryde Eubanks, Clemson University, Clemson, SC Chair Robert Becker, Kirkwood High School, Kirkwood, MO Craig W. Bowen, Clemson University, Clemson, SC Nancy Devino, ScienceMedia Inc., San Diego, CA Sheldon L. Knoespel, Michigan State University, East Lansing, MI Steve Lantos, Brookline High School, Brookline, MA Jim Schmitt, Eau Claire North High School, Eau Claire, WI Robert G. Silberman, SUNY-Cortland, NY Christie B. Summerlin, University of Alabama-Birmingham, Birmingham, AL DIRECTIONS TO THE EXAMINER–PART III The laboratory practical part of the National Olympiad Examination is designed to test skills related to the laboratory. Because the format of this part of the test is quite different from the first two parts, there is a separate, detailed set of instructions for the examiner. This gives explicit directions for setting up and administering the laboratory practical. There are two laboratory tasks to be completed during the 90 minutes allotted to this part of the test. Students do not need to stop between tasks, but are responsible for using the time in the best way possible. Each procedure must be approved for safety by the examiner before the student begins that procedure. Part III 2 lab problems laboratory practical 1 hour, 30 minutes Students should be permitted to use non-programmable calculators. DIRECTIONS TO THE EXAMINEE–PART III DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. WHEN DIRECTED, TURN TO PAGE 2 AND READ THE INTRODUCTION AND SAFETY CONSIDERATIONS CAREFULLY BEFORE YOU PROCEED. There are two laboratory-related tasks for you to complete during the next 90 minutes. There is no need to stop between tasks or to do them in the given order. Simply proceed at your own pace from one to the other, using your time productively. You are required to have a procedure for each problem approved for safety by an examiner before you carry out any experimentation on that problem. You are permitted to use a non-programmable calculator. At the end of the 90 minutes, all answer sheets should be turned in. Be sure that you have filled in all the required information at the top of each answer sheet. Carefully follow all directions from your examiner for safety procedures and the proper disposal of chemicals at your examining site. Not valid for use as an USNCO National Examination after April 16, 2000. Distributed by the ACS DivCHED Examinations Institute, Clemson University, Clemson, SC 2000 UNITED STATES NATIONAL CHEMISTRY OLYMPIAD PART III — LABORATORY PRACTICAL Student Instructions Introduction These problems test your ability to design and carry out laboratory experiments and to draw conclusions from your experimental work. You will be graded on your experimental design, on your skills in data collection, and on the accuracy and precision of your results. Clarity of thinking and communication are also components of successful solutions to these problems, so make your written responses as clear and concise as possible. Safety Considerations You are required to wear approved eye protection at all times during this laboratory practical. You also must follow all directions given by your examiner for dealing with spills and with disposal of wastes. Lab Problem 1 If anhydrous ammonium nitrate is added to water at room temperature, the temperature of the solution decreases. However, if anhydrous calcium chloride is added to water at room temperature, the temperature of the solution increases. Design and carry out an experiment to determine what mass of ammonium nitrate must be added along with 10.0 g of CaCl2 to 100. mL of water so that the final temperature of the solution is the same as the initial room temperature of the water. You will be asked to describe the method you developed to solve this problem. Lab Problem 2 Although bromocresol green is often used as an acid-base indicator, in this problem it is being used as a reactant. When mixed with a dilute solution of household bleach, bromocresol green is gradually oxidized and changes to a different color, a color that happens to match that of bromocresol green at a pH of 4. Design and carry out an experiment to determine the kinetic order of this redox reaction with respect to bleach. You will be asked to describe the method you developed to solve this problem. Page 2 Not valid for use as an USNCO National Examination after April 16, 2000. Answer Sheet for Laboratory Practical Problem 1 Student's Name: __________________________________________________________________________ Student's School:________________________________________ Date: ___________________________ Proctor's Name: _________________________________________________________________________ ACS Section Name :________________________________Student's USNCO test #: ________________ 1. Give a brief description of your experimental plan. Include a sketch of the equipment you will use and the steps you plan to take to solve this problem. Before beginning your experiment, you must get approval (for safety reasons) from the examiner. Not valid for use as an USNCO National Examination after April 16, 2000. Examiner’s Initials: Page 3 2. Record your data and other observations. 3. What is the mass of ammonium nitrate that must be added along with 10.0 g of calcium chloride to 100 mL of water and so that the temperature of the resulting solution is the same as that of the water? Show your methods clearly. Mass of ammonium nitrate: Page 4 Not valid for use as an USNCO National Examination after April 16, 2000. Answer Sheet for Laboratory Practical Problem 2 Student's Name: __________________________________________________________________________ Student's School:________________________________________ Date: ___________________________ Proctor's Name: _________________________________________________________________________ ACS Section Name : ________________________________Student's USNCO test #: ________________ 1. Give a brief description of your experimental plan. List the equipment and materials you plan to use and the steps you plan to take to solve this problem. Before beginning your experiment, you must get approval (for safety reasons) from the examiner. Not valid for use as an USNCO National Examination after April 16, 2000. Examiner’s Initials: Page 5 2. Record your data and other observations. 3. Determine the order of the reaction with respect to bleach. Show your reasoning clearly. Order of reaction with respect to bleach: Page 6 Not valid for use as an USNCO National Examination after April 16, 2000. 2000 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART III Prepared by the American Chemical Society Olympiad Laboratory Practical Task Force Examiner's Instructions Directions to the Examiner: Thank you for administering the 2000 USNCO laboratory practical on behalf of your Local Section. It is essential that you follow the instructions provided, in order to insure consistency of results nationwide. There may be considerable temptation to assist the students after they begin the lab exercise. It is extremely important that you do not lend any assistance or hints whatsoever to the students once they begin work. As in the international competition, the students are not allowed to speak to anyone until the activity is complete. The equipment needed for each student for both lab exercises should be available at his/her lab station or table when the students enter the room. The equipment should be initially placed so that the materials used for Lab Problem 1 are separate from those used for Lab Problem 2. After the students have settled, read the following instructions (in italics) to the students. Hello, my name is ________. Welcome to the lab practical portion of the U.S. National Chemistry Olympiad Examination. In this part of the exam, we will be assessing your lab skills and your ability to reason through a laboratory problem and communicate its results. Do not touch any of the equipment in front of you until you are instructed to do so. One of this year’s problems requires the use of a Styrofoam® cup calorimeter. These have been assembled for you to preserve maximum time for your experimentation. There are extra cups available for you to use as the inner cup. Show a Styrofoam® cup calorimeter. (See picture of the set-up on page 3 of these instructions.) This problem also requires you to use a balance, which is located _____________________________. Another of this year’s problems uses small-scale chemistry equipment. Small-scale chemistry techniques help to minimize the amount of materials you use, thereby increasing safety and minimizing waste. Specialized equipment for small-scale chemistry that you will use today include Beral-type pipets and reaction plates. Show a 1-mL, 3-mL, or a 5-mL Beral-type pipet, and show a 6-well or a 12-well reaction plate. You will be asked to complete two laboratory problems. All the materials and equipment you may want to use to solve each problem has been set out for you and is grouped by the number of the problem. You must limit yourself to this equipment for each problem. You will have one hour and thirty minutes to complete the two problems. You may choose to start with either problem. You are required to have a procedure for each problem approved for safety by an examiner. (Remember that approval does not mean that your procedure will be successful–it is a safety approval.) When you are ready for an examiner to come to your station for each safety approval, please raise your hand. Page 1 Safety is an important consideration during the lab practical. You must wear goggles at all times. Wash off any chemicals spilled on your skin or clothing with large amounts of tap water. The appropriate procedures for disposing of solutions at the end of this lab practical are: ____________________________________________________________________________________ ____________________________________________________________________________________ We are about to begin the lab practical. Please do not turn the page until directed to do so, but read the directions on the front page. Are there any questions before we begin? Distribute Part III booklets and again remind students not to turn the page until the instruction is given. Part III contains student instructions and answer sheets for both laboratory problems. There is a periodic table on the last page of the booklet. Allow students enough time to read the brief cover directions. Do not turn to page 2 until directed to do so. When you start to work, be sure that you fill out all information at the top of the answer sheets. Are there any additional questions? If there are no further questions, the students should be ready to start Part III. You may begin. After one hour and thirty minutes, give the following directions. This is the end of the lab practical. Please stop and bring me your answer sheets. Thank you for your cooperation during this test. Collect all the lab materials. Make sure that the student has filled in his or her name and other required information on the answer sheets. At this point, you may want to take five or ten minutes to discuss the lab practical with the students. They can learn about possible observations and interpretations and you can acquire feedback as to what they actually did and how they reacted to the problems. After this discussion, please take a few minutes to complete the Post-Exam Questionnaire; this information will be extremely useful to the Olympiad subcommittee as they prepare next year’s exam. Please remember to return the post-exam Questionnaire, the answer sheets from Part III, the Scantron sheets from Part I, and the “Blue Books” from Part II to this address: ACS DivCHED Exams Institute Clemson University 223 Brackett Hall Box 340979 Clemson, SC 29634-0979 Monday, April 24, 2000, is the absolute deadline for receipt of the exam materials at the Examinations Institute. Materials received after this deadline CANNOT be graded. THERE WILL BE NO EXCEPTIONS TO THIS DEADLINE DUE TO THE TIGHT SCHEDULE FOR GRADING THIS EXAMINATION. Page 2 EXAMINER’S NOTES Lab Problem #1: Materials and Equipment Each student will need access to a balance that is capable of weighing an object at least to the nearest 0.01 grams. One balance can serve 2-3 students, but one balance per student is highly desirable. Each student will need: One calorimeter such as shown in this diagram. These should be preassembled for the students to preserve maximum time for their experimentation. 1 ring stand 2 clamps 1 thermometer, –10 to 110°C 6 370-mL (12 oz) foam cups 1 370-mL foam cup to serve as cover. Trim off about top one-third of cup, punch two holes of appropriate sizes for the thermometer and stirring rod. 1 5-10 cm piece of wire or string to hold up thermometer 1 30-35 cm piece of 20 gauge copper wire. Bend in a circle on the bottom to serve as a stirrer, hook on top for handle. Notes: The inner cup can be replaced as needed, which is why there are extra foam cups. 250 mL (8 oz) foam cups may also be used but 190 mL (6 oz) cups are not recommended. You may choose to supply a ring clamp to stabilize the base of the cup assembly. Another option is to place the cup assembly in an appropriately sized beaker that serves to support the cups and to provide extra insulation. 2 plastic or metal scoops; can use cut Beral-style pipets 1 500-mL or larger wash bottle, labeled “distilled water” or “deionized water” 1 100-mL graduated cylinder 1 small glass or plastic vial with top, labeled 25.0 g ammonium nitrate 1 small glass or plastic vial with top, labeled 25.0 g calcium chloride 6 small plastic weighing boats, weighing papers, or small dry paper cups 4 Beral-style plastic pipets, 3 mL or 5 mL; eye droppers may be substituted. 1 plastic tub for disposal of liquid wastes (or easy access to sinks) supply of paper towels 1 pair safety goggles 1 lab coat or apron (optional) 1 pair plastic gloves (optional) Page 3 Lab Problem #1: Chemicals Each student will need: 25.0 g anhydrous ammonium nitrate Note: Be sure particles are not clumped. 25.0 g anhydrous calcium chloride Note: Do not use the dihydrate. 500 mL distilled or deionized water Lab Problem #1: Notes 1. Note that the examiner will need to initial each student’s experimental plan. Please do not comment on the plan other than looking for any potentially unsafe practices. 2. Consistent results will depend on starting with anhydrous salts and minimizing exposure to moisture during transfer of the salts to the vials. Be sure the glass or plastic vials are tightly closed after they are filled for each student. 3. The balances can be used as common equipment, although as noted earlier, it will be highly desirable to have one balance per student. Only the first activity requires the use of a balance and this should minimize delays if a number of students must share a balance. Please note any problems on the Examiner’s report sheet if balance use negatively impacts the performance of your students. 4. Safety: It is your responsibility to ensure that all students wear safety goggles during the lab practical. A lab coat or apron for each student is desirable but not mandatory. Please know and follow all safety procedures appropriate for your site. You will need to give students explicit directions for handling spills and for disposing of waste materials, following approved safety practices for your examining site. Lab Problem #2: Materials and Equipment Each student will need: 2 6-well or 1 12-well reaction plate. If reaction plates cannot be obtained or borrowed, 6 50-mL or 100-mL beakers can be substituted. 1 piece of white paper to place under the reaction plates or beakers 3 Beral-style plastic pipets, 1 mL. Eye droppers may be substituted. 1 100-mL or larger wash bottle, labeled “distilled water” or “deionized water” 4 10-mL narrow-mouth plastic dropping bottles. Small glass bottles fitted with eye droppers may be substituted. One labeled “diluted household bleach solution” One labeled “bromocresol green solution” One labeled “pH 4.0 buffer” One labeled “pH 7.0 buffer” 1 timer, stop watch, or access to classroom clock with a second hand 1 plastic tub for disposal of liquid wastes (or easy access to sinks) supply of paper towels 1 pair safety goggles 1 lab coat or apron (optional) 1 pair plastic gloves (optional) Page 4 Lab Problem #2: Chemicals Each student will need: 10 mL of pH 4.0 buffer 10 mL of pH 7.0 buffer 10 mL of diluted bleach solution. Preparation: Obtain fresh commercial bleach containing 5.25% sodium hypochlorite solution. Combine 5.0 mL of this solution with 295 mL distilled water to form the diluted bleach solution. 10 mL of diluted bromocresol green solution. Preparation: Dilute 0.2 g of bromocresol green to form 500 mL of aqueous solution. Do not use existing solutions of bromocresol green, for they may contain alcohol or NaOH commonly used in their preparation. 50 mL of distilled or deionized water Lab Problem #2: Notes 1. Note that the examiner will need to initial each student’s experimental plan. Please do not comment on the plan other than looking for any potentially unsafe practices. 2. Simply combining the two reagents provides enough mixing for the reaction to take place. However, you may wish to provide toothpicks or plastic stirring rods. 3. Premade buffer solutions of pH 4.0 and 7.0 may be used, or purchased from any one of several suppliers. If preparing your own buffer solutions, a useful reference is Silberman, R. G., The Journal of Chemical Education, 1992 (Vol 69, No. 2), p. A42. This article describes a system using boric acid, citric acid monohydrate, and trisodium phoshate dodecahydrate. These directions are also given in ACS Small-Scale Laboratory Assessment Activities, ACS Examinations Institute, 1996, p. C-14. 4. Safety: It is your responsibility to ensure that all students wear safety goggles during the lab practical. A lab coat or apron for each student is desirable but not mandatory. Please know and follow all safety procedures appropriate for your site. You will need to give students explicit directions for handling spills and for disposing of waste materials, following approved safety practices for your examining site. Page 5 Page 6 2000 U. S. NATIONAL CHEMISTRY OLYMPIAD KEY for NATIONAL EXAM—PART III Lab Problem 1 Plan: State: a. Add mass CaCl2 to recorded amount of H2O b. Record initial and final temperatures to obtain ∆T for CaCl2 sample. c. Add mass NH4NO3 to recorded amount of H2O. d. Record initial and final temperatures to obtain ∆T for NH4NO3 sample e. Repeat 2-3 times. Sketch the apparatus as directed in the instructions. Data: Record: a. mass CaCl2 and volume of water. b. Ti and Tf for CaCl2 c. mass NH4NO3 and volume of water d. Ti and Tf for NH4NO3 Amounts of solids added should be large enough to get a reasonable value for ∆T e. results of a “verification” experiment to check prediction. Replication: multiple trials Calculations: ∆T ∆T and for the same amount of H 2 O or Calculate: a. g CaCl 2 g NH 4 NO 3 q q and g CaCl 2 g NH 4 NO 3 for the same amount of H 2 O b. ratio of grams of NH4NO3 to CaCl2 c. amount of NH4NO3 needed for 100 grams of H2O Sample Data: Mass NH4NO3 Volume H2O Initial temperature Final temperature ∆T (NH4NO3) Trial 1 5.005 g 50.0 mL 21.5 °C 14.8 °C –6.7 °C Trial 2 5.073 g 51.9 mL 21.5 °C 15.0 °C –6.5 °C mass CaCl2 volume H2O Initial temperature Final temperature ∆T (CaCl2) 5.064 g 50.0 mL 21.4 °C 33.7 °C 12.3 °C` 5.187 g 52.0 mL 21.0 °C 34.0 °C 13.0 °C Sample Calculations: For NH4NO3, note that the solution became cooler as the NH4NO3 dissolved. This means the dissolution process is absorbing heat from the water. q NH 4 NO3 = – q H 2 O Trial 1 q NH 4 NO3 Trial 2 mc∆T =– g q NH 4 NO3 = – mc∆T g q NH 4 NO3 = – q NH 4 NO3 = (50.0g) (4.184 J ⋅ g –1 ⋅ °C –1 ) (–6.7 o C) 5.005 g 1.40 × 10 3 J 2.79 × 10 2 J = 5.005 g g qaverage for NH 4 NO 3 = q NH 4 NO3 = – q NH 4 NO3 = (51.9 g) (4.184 J ⋅ g –1 ⋅ °C –1 ) (–6.5 o C) 5.073 g 1.41 × 10 3 J 2.78 × 10 2 J = 5.073 g g 2.79 × 10 2 J or 279 J / g g This is the heat absorbed from the water as the NH4NO3 dissolves. For CaCl2, note that the solution became warmer as the CaCl2 dissolved. This means the dissolution process is releasing heat to the water. qCaCl2 = – q H 2 O Trial 1 mc∆T q= g Trial 2 mc∆T q= g qCaCl2 = – (50.0g) (4.184 J ⋅ g –1 ⋅ °C –1 ) (12.3 o C) 5.064 g qCaCl2 = – (52.0g) (4.184 J ⋅ g –1 ⋅ °C –1 ) (13.0 o C) 5.187 g qCaCl2 = – 2.57 × 103 J 5.08 × 10 2 J =– 5.064 g g qCaCl2 = – 2.83 × 10 3 J 5.45 × 10 2 J =– 5.187 g g qaverage for CaCl 2 = – 5.27 × 10 2 J or – 527 J / g g This is the heat released to the water as the CaCl2 dissolves. 1.89 g NH 4 NO 3 527 J / g CaCl 2 = Ratio of heats: qCaCl2 = – q NH 4 NO3 so the ratio is: 279 J / g NH 4 NO 3 1.00 g CaCl 2 Conclusion: For 10.0 g CaCl2: 10.0 g CaCl 2 × 1.89 g NH 4 NO 3 = 18.9 g NH 4 NO 3 1.00 g CaCl 2 Lab Problem 2 Plan: State: a. Add a small amount of bromocresol green to a well of pH 4 buffer and to a well of pH 7 buffer to determine color at each pH. b. Measure known amounts of bromocresol green to well plates (by counting drops). Add various amounts of bleach (by counting drops) and record time to color change. c. Ideally, the reactions will be conducted at constant volume by adding either H2O or pH 7 buffer so that concentration ratios are directly related to volumes of bleach used for each reaction. Replications should be made for each trial. Data: Record: a. color change of bromocresol green from blue at pH = 7 to yellow at pH = 4 b. drops of bromocresol green added to each well c. drops of buffer or water added to each well, if used in plan d. drops of bleach added to each well e. time required for color change Replication: multiple trials Calculations: To determine the order of reaction, use one of these equations. time1 [bleach2 ] = time2 [bleach1 ]x x rate1 [bleach1 ] = rate2 [bleach2 ]x x or Sample Data: The observed color change for bromocresol green is from blue at pH = 7 to yellow at pH = 4. Data gathered if reactions are buffered at pH = 7 Trial Drops of Drops of Bromocresol Green Buffer 1 20 5 2 20 5 3 20 5 4 20 1.5 5 20 1.5 6 20 1.5 Drops of Bleach 20 20 20 10 10 10 Data gathered if reactions are not buffered Trial Drops of Drops of Drops of Bromocresol Green Buffer Bleach 1 20 5 20 2 20 5 20 3 20 5 20 4 20 1.5 10 5 20 1.5 10 6 20 1.5 10 Note: Students do not need to do both buffered and unbuffered reactions. Sample Calculation: Using data from the observations from the pH 7 buffered reactions: Time for color change to occur 14 16 17 28 27 28 Average time for three trials Time for color change to occur 50 52 53 50 51 54 Average time for three trials 15.7 27.7 51.7 51.7 time1 [bleach2 ] = and the concentration of bleach is directly proportional to the number of drops added. time2 [bleach1 ]x x 27.7 seconds (20 drops) = and x ≈ 1 15.7 seconds (10 drops) x x rate = k[bleach]1 if the reaction is carried out in the buffer. Using data from the observations from the unbuffered reactions: time1 [bleach2 ] = and the concentration of bleach is directly proportional to the number of drops added. time2 [bleach1 ]x x 51.7 seconds (20 drops) = and x ≈ 0 51.7 seconds (10 drops) x x rate = k[bleach]0 if the reaction is not carried out in the pH = 7 buffer. Conclusion: If buffered at pH = 7, then the reaction is first order with respect to bleach. If unbuffered, then the reaction is zero order with respect to bleach. 2001 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART I Prepared by the American Chemical Society Olympiad Examinations Task Force OLYMPIAD EXAMINATIONS TASK FORCE Arden P. Zipp, State University of New York, Cortland, NY Chair Jo A. Beran, Texas A&M University-Kingsville, TX Peter E. Demmin (retired), Amherst Central High School, NY Edward DeVillafranca (retired), Kent School, CT Dianne H. Earle, Paul M. Dorman High School, SC Alice Johnsen, Bellaire High School, TX Patricia A. Metz, United States Naval Academy, MD Ronald O. Ragsdale, University of Utah, UT Diane D. Wolff, Western Virginia Community College, VA DIRECTIONS TO THE EXAMINER–PART I Part I of this test is designed to be taken with a Scantron® answer sheet on which the student records his or her responses. Only this Scantron sheet is graded for a score on Part I. Testing materials, scratch paper, and the Scantron sheet should be made available to the student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until April 22, 2001, after which tests can be returned to students and their teachers for further study. Allow time for the student to read the directions, ask questions, and fill in the requested information on the Scantron sheet. The answer sheet must be completed using a pencil, not pen. When the student has completed Part I, or after one hour and thirty minutes has elapsed, the student must turn in the Scantron sheet, Part I of the testing materials, and all scratch paper. There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and you are free to schedule rest-breaks between parts. Part I Part II Part III 60 questions 8 questions 2 lab problems single-answer multiple-choice problem-solving, explanations laboratory practical 1 hour, 30 minutes 1 hour, 45 minutes 1 hour, 30 minutes A periodic table and other useful information are provided on page 2 for student reference. Students should be permitted to use nonprogrammable calculators. DIRECTIONS TO THE EXAMINEE–PART I DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Answers to questions in Part I must be entered on a Scantron answer sheet to be scored. Be sure to write your name on the answer sheet; an ID number is already entered for you. Make a record of this ID number because you will use the same number on both Parts II and III. Each item in Part I consists of a question or an incomplete statement that is followed by four possible choices. Select the single choice that best answers the question or completes the statement. Then use a pencil to blacken the space on your answer sheet next to the same letter as your choice. You may write on the examination, but the test booklet will not be used for grading. Scores are based on the number of correct responses. When you complete Part I (or at the end of one hour and 30 minutes), you must turn in all testing materials, scratch paper, and your Scantron answer sheet. Do not forget to turn in your U.S. citizenship statement before leaving the testing site today. Not valid for use as an USNCO National Exam after April 22, 2001. Distributed by the ACS DivCHED Examinations Institute, Clemson University, Clemson, SC. All rights reserved. Printed in U.S.A. amount of substance ampere atmosphere atomic mass unit atomic molar mass Avogadro constant Celsius temperature centi- prefix coulomb electromotive force energy of activation enthalpy entropy ABBREVIATIONS AND SYMBOLS n equilibrium constant K measure of pressure mmHg A Faraday constant F milli- prefix m atm formula molar mass M molal m u free energy G molar M A frequency ν mole mol N A gas constant R Planck’s constant h °C gram g pressure P c heat capacity C p rate constant k C hour h retention factor Rf E joule J second s Ea kelvin K speed of light c H kilo- prefix k temperature, K T S liter L time t volt V E =E – USEFUL EQUATIONS –∆H 1 ln K = +c R T RT lnQ nF CONSTANTS R = 8.314 J·mol –1·K–1 R = 0.0821 L·atm·mol –1·K–1 1 F = 96,500 C·mol –1 1 F = 96,500 J·V–1·mol–1 N A = 6.022 × 10 23 mol–1 h = 6.626 × 10 –34 J·s c = 2.998 × 10 8 m·s –1 k 2 Ea 1 1 = − k1 R T1 T2 ln PERIODIC TABLE OF THE ELEMENTS 1 H 2 He 1.008 4.003 3 Li 4 Be 5 B 6 C 7 N 8 O 9 F 10 Ne 6.941 9.012 10.81 12.01 14.01 16.00 19.00 20.18 11 Na 12 Mg 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar 22.99 24.31 26.98 28.09 30.97 32.07 35.45 39.95 19 K 20 Ca 21 Sc 22 Ti 23 V 24 Cr 25 Mn 26 Fe 27 Co 28 Ni 29 Cu 30 Zn 31 Ga 32 Ge 33 As 34 Se 35 Br 36 Kr 39.10 40.08 44.96 47.88 50.94 52.00 54.94 55.85 58.93 58.69 63.55 65.39 69.72 72.61 74.92 78.96 79.90 83.80 37 Rb 38 Sr 39 Y 40 Zr 41 Nb 42 Mo 43 Tc 44 Ru 45 Rh 46 Pd 47 Ag 48 Cd 49 In 50 Sn 51 Sb 52 Te 53 I 54 Xe 85.47 87.62 88.91 91.22 92.91 95.94 (98) 101.1 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.3 55 Cs 56 Ba 57 La 72 Hf 73 Ta 74 W 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 132.9 137.3 138.9 178.5 181.0 183.8 186.2 190.2 192.2 195.1 197.0 200.6 204.4 207.2 209.0 (209) (210) (222) 87 Fr 88 Ra 89 Ac 104 Rf 105 Db 106 Sg 107 Bh 108 Hs 109 Mt 110 111 112 114 116 118 (223) 226.0 227.0 (261) (262) (263) (262) (265) (266) (269) (272) (277) (289) (289) (293) Page 2 58 Ce 59 Pr 60 Nd 61 Pm 62 Sm 63 Eu 64 Gd 65 Tb 66 Dy 67 Ho 68 Er 69 Tm 70 Yb 71 Lu 140.1 140.9 144.2 (145) 150.4 152.0 157.3 158.9 162.5 164.9 167.3 168.9 173.0 175.0 90 Th 91 Pa 92 U 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr 232.0 231.0 238.0 237.0 (244) (243) (247) (247) (251) (252) (257) (258) (259) (260) Not valid for use as a USNCO National Exam after April 22, 2001. DIRECTIONS § When you have selected your answer to each question, blacken the corresponding space on the answer sheet using a soft, #2 pencil. Make a heavy, full mark, but no stray marks. If you decide to change an answer, erase the unwanted mark very carefully. § You may write on the test booklet, but it will not be used for grading. § There is only one correct answer to each question. Any questions for which more than one response has been blackened will not be counted. § Your score is based solely on the number of questions you answer correctly. It is to your advantage to answer every question. 1. Which of these compounds is amphoteric? I. Al(OH)3 II. Ba(OH)2 III. Zn(OH)2 7. What is the purpose of this apparatus? water out (A) I only (B) II only (C) I and III only (D) II and III only 2. Calcium hydride reacts with excess water to form (A) CaO and H2 (B) Ca(OH)2 and O 2 (C) Ca(OH)2 only (D) Ca(OH)2 and H 2 3. What is the most likely boiling point of an equimolar mixture of hexane, C6H14, and heptane, C7H16? Boiling Point C 6H14 69 °C C 7H16 98 °C (A) below 69 °C (B) between 69 and 98 °C (C) 69 °C (D) 98 °C 4. Which element melts at the highest temperature? (A) aluminum (B) silicon (C) phosphorus (D) sulfur 5. Which substance participates readily in both acid-base and oxidation-reduction reactions? (A) Na 2CO3 (B) KOH (C) KMnO4 (D) H2C 2O4 6. What mass of magnesium hydroxide is required to neutralize 125 mL of 0.136 M hydrochloric acid solution? Substance Molar Mass Mg(OH)2 (A) 0.248 g (B) 0.496 g (C) 0.992 g (D) 1.98 g 58.33 g·mol–1 water in (A) distilling (B) filtering (C) refluxing (D) titrating 8. Calculate the mass of Substance Molar Mass ammonia that can be produced from the (NH4)2PtCl6 443.9 g·mol–1 decomposition of a sample of (NH4)2PtCl6 containing 0.100 g Pt. (A) 0.0811 g (B) 0.0766 g (C) 0.0175 g (D) 0.00766 g 9. Assume 0.10 L of N2 and 0.18 L of H2, both at 50 atm and 450 °C, are reacted to form NH3. Assuming the reaction goes to completion, identify the reagent that is in excess and determine the volume that remains at the same temperature and pressure. (A) H2, 0.02 L (B) H2, 0.08 L (C) N2, 0.01 L (D) N2, 0.04 L 10. Concentrated hydrochloric acid is 12.0 M and is 36.0% hydrogen chloride by mass. What is its density? (A) 1.22 g·mL–1 (B) 1.10 g·mL–1 (C) 1.01 g·mL–1 (D) 0.820 g·mL–1 Not valid for use as an USNCO National Examination after April 22, 2001. Page 3 11. C4H6O3 → acetic anhydride C7H6O3 + salicylic acid C 9H8O4 + aspirin C2H4O2 acetic acid What is the percent yield if 0.85 g of aspirin is formed in the reaction of 1.00 g of salicylic acid with excess acetic anhydride? Substance Molar Mass C 7H6O3 C 4H6O3 C 9H8O4 C 2H4O2 138.12 g·mol–1 102.09 g·mol–1 180.15 g·mol–1 60.05 g·mol –1 (A) 65 % (C) 85 % (D) 91 % (B) 77 % 12. The triple point of CO2 occurs at 5.1 atm and –56 °C. Its critical temperature is 31 °C. Solid CO2 is more dense than liquid CO2. Under which combination of pressure and temperature is liquid CO2 stable at equilibrium? (A) 10 atm and –25 °C (B) 5.1 atm and –25 °C (C) 10 atm and 33 °C (D) 5.1 atm and –100 °C 13. The vapor pressure of water at 20 °C is 17.54 mmHg. What will be the vapor pressure of the water in the apparatus shown after the piston is lowered, decreasing the volume of the gas above the liquid to one half of its initial volume? (Assume no temperature change.) 16. What is the average velocity of H2 molecules at 100 K relative to their velocity at 50 K? (A) 2.00 times the velocity at 50 K (B) 1.41 times the velocity at 50 K (C) 0.71 times the velocity at 50 K (D) 0.50 times the velocity at 50 K 17. What type of semiconductor results when highly purified silicon is doped with arsenic? (A) n–type (B) p–type (C) q–type (D) s–type 18. The heat of formation of NO from its elements is +90 kJ·mol–1. What is the approximate bond dissociation energy of the bond in NO? Bond N N O O Bond Energy 941 kJ·mol –1 499 kJ·mol –1 (A) 630 kJ·mol –1 (B) 720 kJ·mol –1 (C) 765 kJ·mol –1 (D) 810 kJ·mol –1 water vapor liquid water (A) 8.77 mmHg 19. How much energy must be supplied to change 36 g of ice at 0 °C to water at room temperature, 25 °C? Data for Water, H2O ∆Hofusion 6.01 kJ·mol–1 C P, liquid 4.18 J·K –1·g –1 (A) 12 kJ (B) 16 kJ (C) 19 kJ (D) 22 kJ (B) 17.54 mmHg (C) 35.08 mmHg 20. For a process that is both endothermic and spontaneous, (D) between 8.77 and 17.54 mmHg 14. What is the density of propane, C3H8, at 25 °C and 740. mmHg? (A) 0.509 g·L –1 (B) 0.570 g·L –1 (C) 1.75 g·L –1 (D) 1.96 g·L –1 15. An unknown gas effuses through a small hole one half as fast as methane, CH 4, under the same conditions. What is the molar mass of the unknown gas? (A) ∆H < 0 (B) ∆G > 0 (C) ∆E = 0 (D) ∆S > 0 21. Consider the values for ∆Ho (in kJ·mol–1) and for ∆So (in J·mol –1·K–1) given for four different reactions. For which reaction will ∆Go increase the most (becoming more positive) when the temperature is increased from 0 °C to 25 °C? (A) ∆Ho = 50, ∆So = 50 (B) ∆Ho = 90, ∆So = 20 (A) 4 g·mol –1 (B) 8 g·mol –1 (C) ∆Ho = –20, ∆So = –50 (C) 32 g·mol –1 (D) 64 g·mol –1 (D) ∆Ho = –90, ∆So = –20 Page 4 Not valid for use as an USNCO National Examination after April 22, 2001. 22. A certain reaction is exothermic by 220 kJ and does 10 kJ of work. What is the change in the internal energy of the system at constant temperature? (A) +230 kJ (B) +210 kJ (C) –210 kJ (D) –230 kJ 23. Fe2O3(s) + 3/2C(s) → 3/2CO2(g) + 2Fe(s) ∆Ho = +234.1 kJ C(s) + O2(g) → CO2(g) ∆Ho = –393.5 kJ Use these equations and ∆Ho values to calculate ∆Ho for this reaction. 4Fe(s) + 3O2(g) → 2Fe2O3(s) (A) –1648.7 kJ (B) –1255.3 kJ (C) –1021.2 kJ (D) –129.4 kJ 24. A large positive value of ∆G corresponds to which of these? o 27. Consider this reaction. 2NO2(g) + O3(g) → N 2O5(g) + O2(g) The reaction of nitrogen dioxide and ozone represented is first order in NO2(g) and in O3(g). Which of these possible reaction mechanisms is consistent with the rate law? Mechanism I. NO2 + O3 → NO3 + O2 slow NO3 + NO2 → N2O5 fast Mechanism II. O3 ¾ O2 + O fast NO2 + O → NO3 NO3 + NO2 → N2O5 slow fast (A) I only (B) II only (C) both I and II (D) neither I nor II 28. When the temperature of a reaction is raised from 300 K to 310 K, the reaction rate doubles. Determine the activation energy, Ea , associated with the reaction. (A) small positive K (B) small negative K (A) 6.45 kJ·mol–1 (B) 23.3 kJ·mol–1 (C) large positive K (D) large negative K (C) 53.6 kJ·mol–1 (D) 178 kJ·mol –1 25. What names apply to chemical species corresponding to locations 1 and 2 on this reaction coordinate diagram? 2 1 29. Use the experimental data in the table to determine the rate law for this reaction. A +B → AB These data were obtained when the reaction was studied. [A], M [B], M ∆[ AB] ∆t mol·L–1·s–1 0.100 0.200 0.300 Reaction Progress → Location 1 Location 2 0.100 0.100 0.300 2.0 × 10–4 2.0 × 10–4 1.8 × 10–3 What is the rate equation for the reaction? (A) activated complex activated complex (A) rate = k [A] [B] (B) rate = k [A]2 (B) reaction intermediate activated complex (C) rate = k [B] (D) rate = k [B]2 (C) activated complex reaction intermediate (D) reaction intermediate reaction intermediate 30. Which of the reactions represented in these diagrams will show the greatest increase in rate for a given increase in temperature? 26. Gadolinium-153, which is used to detect osteoporosis, has a half-life of 242 days. Which value is closest to the percentage of the Gd-153 left in a patient's system after 2 years (730 days)? (A) 33.0 % (B) 25.0 % (C) 12.5 % (D) 6.25 % Reaction I Reaction II (A) Reaction I forward (B) Reaction I reversed (C) Reaction II forward (D) Reaction II reversed Not valid for use as an USNCO National Examination after April 22, 2001. Page 5 Questions 31 and 32 should both be answered with reference to this equilibrium system. 2NH3(g) ¾ N2(g) + 3H2(g) K p = 80.0 at 250 °C 38. 31. What is Kp for this reaction? 1/2N2(g) + 3/2H2(g) ¾ NH3(g) Eo = –1.66 V Eo = +0.34 V (A) 0.0125 (B) 0.112 What voltage is produced under standard conditions by combining the half-reactions with these Standard Electrode Potentials? (C) 8.94 (D) 40.0 (A) 1.32 V 32. What is the expression for Kc at 250 °C for this reaction? 2NH3(g) ¾ N2(g) + 3H2(g) (A) Kc = (C) Kc Kp (RT )2 2 = K p (RT) Kp RT (B) Kc = (D) Kc = K p RT HCOOH(aq) ¾ H+(aq) + HCOO–(aq) K a = 1.7 × 10 –4 The ionization of formic acid is represented. Calculate [H+] of a solution initially containing 0.10 M HCOOH and 0.050 M HCOONa. (B) 2.00 V (C) 2.30 V (D) 4.34 V 39. For which of these oxidation/reduction pairs will the reduction potential vary with pH? I. AmO22+/AmO2+ II. AmO22+/Am4+ III. Am4+/Am2+ 40. 33. Al3+(aq) + 3e– → Al(s) Cu 2+(aq) +2e– → Cu(s) (A) I only (B) II only (C) I and II only (D) I, II, and III 2Ag+(aq) + Cu(s) → Cu2+(aq) + 2Ag(s) The standard potential for this reaction is 0.46 V. Which change will increase the potential the most? (A) doubling the [Ag+] (B) halving the [Cu2+] (A) 8.5 × 10 M (B) 3.4 × 10 M (C) doubling the size of the Cu(s) electrode (C) 4.1 × 10 M (D) 1.8 × 10–2 M (D) decreasing the size of the Ag electrode by one half –5 –3 –4 34. Which are strong acids? I. HClO3 II. H2SeO3 III. H3AsO4 (A) I only (B) III only (C) I and III only (D) II and III only 35. Carbonic acid, H2CO3, is a diprotic acid for which K 1 = 4.2 × 10 –7 and K2 = 4.7 × 10 –11. Which solution will produce a pH closest to 9? (A) 0.1 M H 2CO3 (B) 0.1 M Na2CO3 (C) 0.1 M NaHCO3 (D) 0.1 M NaHCO3 and 0.1 M Na2CO3 36. What is the pH of a saturated solution of magnesium hydroxide, Mg(OH) 2 at 25 °C? (K sp = 6.0 × 10 –12 at 25˚C) (A) 10.56 (B) 10.36 (C) 10.26 (D) 10.05 41. 10Cl–(aq) + 2MnO4–(aq) + 16H+(aq) → 5Cl2(g) + 2Mn2+(aq) + 8H2O(l) The value of Eo for this reaction at 25 °C is 0.15 V. What is the value of K for this reaction? (A) 2.4 × 1025 (B) 4.9 × 1012 (C) 1.2 × 105 (D) 3.4 × 102 42. When water is electrolyzed, hydrogen and oxygen gas are produced. If 1.008 g of H 2 is liberated at the cathode, what mass of O2 is formed at the anode? (A) 32.0 g (B) 16.0 g (C) 8.00 g (D) 4.00 g 43. Which property of an element is most dependent on the shielding effect? (A) atomic number (B) atomic mass 37. P 4(s) + 3OH–(aq) + 3H2O(l) → PH 3(g) + 3H2PO2–(aq) For this reaction, the oxidizing and reducing agents are, respectively, (A) P 4 and OH– (B) OH– and P4 (C) P 4 and H 2O (D) P 4 and P4 Page 6 (C) atomic radius (D) number of stable isotopes Not valid for use as an USNCO National Examination after April 22, 2001. 44. How many unpaired electrons are present in a ground state gaseous Ni2+ ion? (A) 0 (B) 2 (C) 4 (D) 6 45. When the elements carbon, nitrogen and oxygen are arranged in order of increasing ionization energies, what is the correct order? (A) C, N, O (B) O, N, C (C) N, C, O (D) C, O, N (A) n = 2, l = 0 (B) n = 2, l = 1 (C) n = 3, l = 0 (D) n = 3, l = 1 47. Which element will exhibit the photoelectric effect with light of the longest wavelength? (B) Rb (C) Mg (D) Ca 48. All these elements have common allotropes except (A) C (B) O (C) Kr (A) Electrons are simultaneously attracted by more than one nucleus. (B) Filled orbitals of two or more atoms overlap one another. (C) Unoccupied orbitals of two or more atoms overlap one another. (D) Oppositely-charged ions attract one another. 46. Given this set of quantum numbers for a multi-electron atom: 2, 0, 0, 1 /2 and 2, 0, 0, –1 /2. What is the next higher allowed set of n and l quantum numbers for this atom in its ground state? (A) K 52. Which is the best description of a covalent bond? (D) S 53. What is the formal charge on the chlorine atom in the oxyacid HOClO2 if it contains only single bonds? (A) –2 (B) –1 (C) +1 (D) +2 54. Which of these compounds is not adequately represented by a valence bond model? I. CO2 II. SO2 III. SiO2 (A) I only (B) II only (C) I and III only (D) II and III only 55. Which compound is not correctly matched with its class name? (A) HCOOH, acid (B) C 6H5CHO, aldehyde (C) C 2H5COCH3, ether 49. How many sigma and pi bonds are shown in this compound? H H O H N C C C C C H H 56. How many toluene derivatives have the formula C7H7Cl? (A) 8 sigma and 7 pi (B) 8 sigma and 3 pi (C) 11 sigma and 3 pi (D) 11 sigma and 4 pi 50. Which reaction involves a change in the electron-pair geometry for the underlined atom? (A) B F 3 + F– → B F 4 (B) NH3 + H+ → NH4+ (C) 2S O2 + O2 → 2 S O3 (D) H2O + H+ → H3O + 51. Which compound has the largest lattice energy? (A) NaF (B) CsI (C) MgO (D) CH3CHOHCH3, secondary alcohol (D) CaS (A) 1 (B) 2 (C) 3 (D) 4 57. When the compounds CH 3COOH, C2H5OH and C 6H5OH are arranged in order of increasing acidity in aqueous solution, which order is correct? (A) C 2H5OH < CH3COOH < C 6H5OH (B) C 6H5OH < CH3COOH < C 2H5OH (C) CH3COOH < C 6H5OH < C2H5OH (D) C 2H5OH < C6H5OH < CH3COOH 58. Which can be used as a catalyst in an esterification reaction? I. NaOH II. H2SO4 (A) I only (B) II only (C) both I and II (D) neither I nor II Not valid for use as an USNCO National Examination after April 22, 2001. Page 7 59. Which term best describes a carbocation? 60. A racemic mixture has equal amounts of (A) electrophilic (B) free radical (A) alkanes and alkenes. (C) hydrophobic (D) nucleophilic (B) cis and trans isomers. (C) functional group isomers. (D) enantiomers. END OF TEST Page 8 Not valid for use as an USNCO National Examination after April 22, 2001. U. S. National Chemistry Olympiad – 2001 National Examination—Part I SCORING KEY Number 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. Answer C D B B D B C C D A A A B C D B A A B D Number 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. 31. 32. 33. 34. 35. 36. 37. 38. 39. 40. Answer C D A A B C C C D B B A B A C B D B B A Property of the ACS Society Committee on Education Number 41. 42. 43. 44. 45. 46. 47. 48. 49. 50. 51. 52. 53. 54. 55. 56. 57. 58. 59. 60. Answer A C C B D B B C D A C A D B C D D B A D 2001 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART II Prepared by the American Chemical Society Olympiad Examinations Task Force OLYMPIAD EXAMINATIONS TASK FORCE Arden P. Zipp, State University of New York, Cortland, NY Chair Jo A. Beran, Texas A&M University-Kingsville, TX Peter E. Demmin (retired), Amherst Central High School, NY Edward DeVillafranca (retired), Kent School, CT Dianne H. Earle, Paul M. Dorman High School, SC Alice Johnsen, Bellaire High School, TX Patricia A. Metz, United States Naval Academy, MD Ronald O. Ragsdale, University of Utah, UT Diane D. Wolff, Western Virginia Community College, VA DIRECTIONS TO THE EXAMINER–PART II Part II of this test requires that student answers be written in a response booklet of blank pages. Only this “Blue Book” is graded for a score on Part II. Testing materials, scratch paper, and the “Blue Book” should be made available to the student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until April 22, 2001, after which tests can be returned to students and their teachers for further study. Allow time for the student to read the directions, ask questions, and fill in the requested information on the “Blue Book”. When the student has completed Part II, or after one hour and forty-five minutes has elapsed, the student must turn in the “Blue Book”, Part II of the testing materials, and all scratch paper. Be sure that the student has supplied all of the information requested on the front of the “Blue Book,” and that the same identification number used for Part I has been used again for Part II. There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and you are free to schedule rest-breaks between parts. Part I Part II Part III 60 questions 8 questions 2 lab problems single-answer multiple-choice problem-solving, explanations laboratory practical 1 hour, 30 minutes 1 hour, 45 minutes 1 hour, 30 minutes A periodic table and other useful information are provided on the back page for student reference. Students should be permitted to use non-programmable calculators. DIRECTIONS TO THE EXAMINEE–PART II DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Part II requires complete responses to questions involving problem-solving and explanations. One hour and forty-five minutes are allowed to complete this part. Be sure to print your name, the name of your school, and your identification number in the spaces provided on the “Blue Book” cover. (Be sure to use the same identification number that was coded onto your Scantron® sheet for Part I.) Answer all of the questions in order, and use both sides of the paper. Do not remove the staple. Use separate sheets for scratch paper and do not attach your scratch paper to this examination. Not valid for use as an USNCO National Exam after April 22, 2001. Distributed by the ACS DivCHED Examinations Institute, Clemson University, Clemson, SC. All rights reserved. Printed in U.S.A. When you complete Part II (or at the end of one hour and forty-five minutes), you must turn in all testing materials, scratch paper, and your “Blue Book.” Do not forget to turn in your U.S. citizenship statement before leaving the testing site today. 1. (10%) The mass percent of MnO2 in a sample of a mineral is determined by reacting it with a measured excess of As2O3 in acid solution, and then titrating the remaining As2O3 with standard KMnO4. A 0.225 g sample of the mineral is ground and boiled with 75.0 mL of 0.0125 M As2O3 solution containing 10 mL of concentrated sulfuric acid. After the reaction is complete, the solution is cooled, diluted with water, and titrated with 2.28 × 10 –3 M KMnO4, requiring 16.34 mL to reach the endpoint. Note: 5 mol of As2O3 react with 4 mol of MnO4–. a. Write a balanced equation for the reaction of As2O3 with MnO2 in acid solution. The products are Mn2+ and AsO43–. b. Calculate the number of moles of i. As2O3 added initially. ii. MnO4– used to titrate the excess As2O3. iii. MnO2 in the sample. c. Determine the mass percent of MnO2 in the sample. d. Describe how the endpoint is detected in the KMnO4 titration. 2. (15%) The presence of CO32–, HCO3– and CO2 in body fluids helps to stabilize the pH of these fluids despite the addition or removal of H+ ions by body processes. Answer the following questions about solutions containing these species in varying combinations at 25 °C. K1 and K2 for H2CO3 are 4.2 × 10 –7 and 4.7 × 10 –11, respectively. a. Write balanced equations to represent the processes responsible for K1 and K2. b. Calculate the [H+] and pH expected for i. 0.033 M solution of H2CO3, which is the saturation point of CO 2 at 25 °C. ii. 1:1 mixture of H2CO3 and HCO 3–. iii. 1:1 mixture of HCO3– and CO32–. iv. 0.125 M solution of CO32–. c. The “normal” pH in blood plasma is 7.40. Identify the components that would provide the best buffer at this pH and calculate their ratio. d. The value of K 1 is based on the assumption that all of the CO 2 dissolved in water exists in the form of H2CO3. However, recent evidence suggests that an additional equilibrium exists as represented by this equation. CO2(aq) + H2O(l) ¾ H2CO3(aq) When the “true” concentration of H2CO3(aq) is taken into account, K 1 =2 × 10 –4. Use this information to determine the percent of dissolved CO2 that is actually present as H2CO3(aq) . 3. (12%) Glucose, C 6H12O6, is readily metabolized in the body. a. Write a balanced equation for the metabolism of C6H12O6 to CO2 and H 2O. b. Calculate ∆Gometabolism for glucose. Given: The free energy of formation, ∆Gf , is –917 kJ·mol –1 for C6H12O6(s); –394.4 kJ·mol–1 for CO 2(g); –237.2 kJ·mol–1 for H2O(l). c. If ∆Ho for this process is –2801.3 kJ, calculate ∆So at 25 °C. d. One step in the utilization of energy in cells is the synthesis of ATP4– from ADP3– and H 2PO4–, according to this equation. ADP3– + H2PO4– → ATP4– ∆Go = 30.5 kJ·mol –1 i. Calculate the number of moles of ATP4– formed by the metabolism of 1.0 g of glucose. ii. Calculate the equilibrium constant, K, for the formation of ATP 4– at 25 ˚C. o 4. (11%) The corrosion of iron is an electrochemical process that involves the standard reduction potentials given here at 25 °C. Fe2+(aq) + 2e– → Fe(s) Eo = –0.44 V + – O2(g) + 4H (aq) + 4e → 2H2O(l) Eo = +1.23 V a. Calculate the voltage for the standard cell based on the corrosion reaction. 2Fe (s) + O2(g) + 4H+(aq) → 2Fe2+(aq) + 2H2O(l) b. Calculate the voltage if the reaction in Part a occurs at pH = 4.00 but all other concentrations are maintained as they were in the standard cell. c. For the reaction Fe(OH)2(s) + 2e– → Fe(s) + 2OH–(aq) , Eo = –0.88 V. Use this information with one of the given standard potentials to calculate the Ksp of Fe(OH)2. Page 2 Not valid for use as an USNCO National Exam after April 22, 2001 d. An iron object may be protected from corrosion by coating it with tin. This method works well as long as the tin coating is intact. However, when the coating is penetrated, the corrosion of the iron is actually accelerated. Use electrochemical principles to account for both of these observations. The standard reduction potential for tin is: Sn 2+(aq) + 2 e – → Sn (s) Eo = –0.14 V 5. 6. (14%) Write net equations for each of these reactions. Use appropriate ionic and molecular formulas for the reactants and products. Omit formulas for all ions or molecules that do not take part in a reaction. Write structural formulas for all organic substances. You need not balance the reactions. All reactions occur in aqueous solution unless otherwise indicated. a. Solid calcium hydrogen carbonate is heated to a very high temperature. b. Solid potassium sulfite is added to a solution of hydrochloric acid. c. Solutions of barium hydroxide and sulfuric acid are mixed. d. A tin (II) chloride solution is added to an acidic solution of potassium dichromate. e. Concentrated hydrochloric acid is added to a solution of sodium hypochlorite. f. Nitrogen-16 undergoes β– decay. (13%) Answer these questions pertaining to chemical kinetics. a. Determine the reaction rate at 10 seconds from the graph. Show your work. b. Using the same units for the reaction rate as in Part a, and assuming concentrations in mol·L–1, give the units for the rate constant of a reaction with an order of: i. zero ii. one iii. two c. Consider this reaction: 4HBr(g) + O2(g) → 2H 2O(l) + 2Br 2(g) i. Express the reaction rates for HBr and Br 2 in this reaction relative to that of O2. ii. Explain why this reaction is unlikely to occur by direct collision of four HBr molecules with one O 2 molecule. d. This mechanism has been suggested for the reaction in Part c: HBr(g) + O2(g) → HOOBr (g) Step 1 HOOBr(g) + HBr(g) → 2HOBr (g) Step 2 HOBr(g) + HBr(g) → H 2O(g) + Br 2(g) Step 3 Give the rate equation in terms of reactants expected for this reaction if the rate-determining step is: i. Step 1 ii. Step 2 iii. Step 3 Assume in each case that the steps before the rate-determining step are in rapid equilibrium. Outline your reasoning in each case. 7. (13%) A certain element, X, forms the fluorides XF 3 and XF 5. Element X also reacts with sodium to form Na3X. a. Give the symbol of an element that behaves in this way. b. For both XF 3 and XF 5; i. write Lewis electron dot structures. ii. describe the electron pair and molecular geometries. iii. give the hybridization of the X atom. c. The bonds in XF5 are not all the same length. Identify the longer bonds and account for this behavior. d. Another element, Y, in the same family as X, forms YF 3 but not YF5. Identify element Y and account for its inability to form YF5. 8. (12%) Account for each observation with appropriate atomic or molecular properties. a. Carbon dioxide has a higher vapor pressure than sulfur dioxide at the same temperature. b. Hydrogen chloride has a lower normal boiling point than either hydrogen fluoride or hydrogen bromide. c. Calcium oxide has a much higher melting point (2580 °C) than potassium fluoride (858 °C). d. Tin (II) chloride is an ionic compound (mp = 240 °C) while tin(IV) chloride is a covalent compound (bp = 114 °C). END OF PART II Not valid for use as an USNCO National Examination after April 22, 2001. Page 3 amount of substance ampere atmosphere atomic mass unit atomic molar mass Avogadro constant Celsius temperature centi- prefix coulomb electromotive force energy of activation enthalpy entropy ABBREVIATIONS AND SYMBOLS n equilibrium constant K measure of pressure mmHg A Faraday constant F milli- prefix m atm formula molar mass M molal m u free energy G molar M A frequency ν mole mol N A gas constant R Planck’s constant h °C gram g pressure P c heat capacity C p rate constant k C hour h retention factor Rf E joule J second s Ea kelvin K speed of light c H kilo- prefix k temperature, K T S liter L time t volt V E =E – USEFUL EQUATIONS –∆H 1 ln K = +c R T RT lnQ nF CONSTANTS R = 8.314 J·mol –1·K–1 R = 0.0821 L·atm·mol –1·K–1 1 F = 96,500 C·mol –1 1 F = 96,500 J·V–1·mol–1 N A = 6.022 × 10 23 mol–1 h = 6.626 × 10 –34 J·s c = 2.998 × 10 8 m·s –1 k 2 Ea = k1 R ln 1 1 − T1 T2 PERIODIC TABLE OF THE ELEMENTS 1 H 2 He 1.008 4.003 3 Li 4 Be 5 B 6 C 7 N 8 O 9 F 10 Ne 6.941 9.012 10.81 12.01 14.01 16.00 19.00 20.18 11 Na 12 Mg 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar 22.99 24.31 26.98 28.09 30.97 32.07 35.45 39.95 19 K 20 Ca 21 Sc 22 Ti 23 V 24 Cr 25 Mn 26 Fe 27 Co 28 Ni 29 Cu 30 Zn 31 Ga 32 Ge 33 As 34 Se 35 Br 36 Kr 39.10 40.08 44.96 47.88 50.94 52.00 54.94 55.85 58.93 58.69 63.55 65.39 69.72 72.61 74.92 78.96 79.90 83.80 37 Rb 38 Sr 39 Y 40 Zr 41 Nb 42 Mo 43 Tc 44 Ru 45 Rh 46 Pd 47 Ag 48 Cd 49 In 50 Sn 51 Sb 52 Te 53 I 54 Xe 85.47 87.62 88.91 91.22 92.91 95.94 (98) 101.1 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.3 55 Cs 56 Ba 57 La 72 Hf 73 Ta 74 W 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 132.9 137.3 138.9 178.5 181.0 183.8 186.2 190.2 192.2 195.1 197.0 200.6 204.4 207.2 209.0 (209) (210) (222) 87 Fr 88 Ra 89 Ac 104 Rf 105 Db 106 Sg 107 Bh 108 Hs 109 Mt 110 111 112 112 116 118 (223) 226.0 227.0 (261) (262) (263) (262) (265) (266) (269) (272) (277) (277) (289) (293) Page 4 58 Ce 59 Pr 60 Nd 61 Pm 62 Sm 63 Eu 64 Gd 65 Tb 66 Dy 67 Ho 68 Er 69 Tm 70 Yb 71 Lu 140.1 140.9 144.2 (145) 150.4 152.0 157.3 158.9 162.5 164.9 167.3 168.9 173.0 175.0 90 Th 91 Pa 92 U 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr 232.0 231.0 238.0 237.0 (244) (243) (247) (247) (251) (252) (257) (258) (259) (260) Not valid for use as an USNCO National Examination after April 16, 2000. 2001 U. S. NATIONAL CHEMISTRY OLYMPIAD KEY for NATIONAL EXAM—PART II 1. a. 2MnO2 + As2O3 + H2O → 2Mn2+ + 2AsO43– + 2H+ b. i. 0.0750 L × ii. 2.28× 10 –3 mol = 3.73 ×10 −5 mol MnO 4– L 5 mol As2O 3 mol MnO4 – × = 4.66× 10−5 mol As 2O 3 left 4 mol MnO 4– 0.01634 L × 3.73× 10 −5 iii. 0.0125 mol = 9.38× 10 −4 mol As 2O3 L 9.38 × 10–4 – 4.66 × 10–5 = 8.91 × 10–4 mol As2O3 react with MnO2 8.91× 10−4 mol As2 O3 × 1.78 ×10 –3 mol MnO 2 × c. 86.94 g MnO2 = 0.155 g MnO2 mol MnO 2 0.155 g MnO2 × 100 = 68.9% MnO 2 in sample 0.225 g sample mass % MnO 2 = 2. 2 mol MnO 2 = 1.78 ×10 −3 mol MnO 2 1 mol As2 O3 d. The endpoint corresponds to a slight purple (pink) color due to excess MnO4–(aq). a. Process responsible for K1 H2CO3(aq) ¾ H+(aq) + HCO3–(aq) Process responsible for K2 HCO3–(aq) ¾ H+(aq) + CO32–(aq) [ H ] [ HCO ] = + b. i. ii. [H 2CO3 ] [H 2CO3 ] = [HCO 3– ] K1 [H ] = + 2 – 3 4.2× 10 –7 (0.033) + iii. [HCO 3– ] = [CO32– ] iv. CO32–(aq) + H2O(l) ¾ HCO3–(aq) + OH–(aq) [OH ] = [ Kw × CO 32 Ka – + [ ] [ H ] = 4.2 ×10 = [H ] [ H ] = 4.7 ×10 K1 = H + K2 [ H ] = 1.2 × 10 + + –4 pH = 3.93 –7 pH = 6.38 –11 pH =10.33 ] – 10 × 0.125 [OH ] = 1.0× 4.7 ×10 pOH = 2.28 [OH ] = 2.7× 10 [ OH ] = 0.0052 [H ] = 1.9× 10 pH = 11.72 Components of the best buffer: H CO pH = 7.40 [ H ] = 4.0 × 10 –14 – –11 – –5 + c. – –12 + –8 2 4.2 × 10 –7 = 4.0 × 10 –8 × HCO 3– H 2CO 3 Key for 2001 USNCO National Exam, Part II 3 and HCO3– HCO 3– = 10.5 H 2CO3 Page 1 d. 3. 4.2 × 10 –7 = 2 ×10 –3 or 0.002 2 × 10 –4 Because the ratio of the two K values is 0.002, 0.2% of dissolved CO2 is actually H2CO3. a. C6H12O6 + 6O2 → 6CO2 + 6H2O b. o ∆Gometabolism = 6∆GCO + 6∆G oH2 O – ∆GCo 6 H12 O6 2 = 6 mol(−394.4 kJ⋅mol –1 )+ 6 mol(–237.2 kJ⋅ mol –1) – ( – 9 1 7kJ⋅ mol –1) = –2366.4 k J – 1 4 2 3 . 2+k 9J 1 7 k J = –2873kJ ∆Go = ∆Ho – T∆So c. ( 2 9 8 K∆S) o = 72 kJ – 2 8 7 3 k J= –2801.3 k J – 2 9 K∆S 8 o d. i. 1.0 g C 6H12 O 6 × ii. 4. a. 2872.6 kJ = 16 kJ mol 1 mol ATP = 0.52 mol ATP formed 30.5 kJ 30.5 ×10 3J –8.314 J = (298 K)ln K mol mol ⋅ K ∆Go = –RT ln K ln K = –12.31 2Fe(s) → 2Fe2+(aq) + 4e– Eo = +0.44 V O2(g) + 4H+(aq) + 4e– → 2H2O(l) Eo = +1.23 V 2Fe(s) + O2(g) + 4H+(aq) → 2Fe2+(aq) b. or 240 J/K 1 mol = 5.6× 10 –3 m o l C6 H 12O6 180 g 5.6 ×10 –3 mol C 6H12 O6 × 16 kJ × ∆So = 0.24 kJ / K [ K = 4.5 ×10 –6 Eo = +1.67 V + 2H2O(l) ] 2+ 2 Fe RT E =E – ln 4 nF H+ PO 2 o [ ]( ) = 1.67V – (8 . 3 1 4 J/ mol⋅ K) ( 2 9 8 K) ( 96,500 J / V) ( 4 mol) ln 1 (1.0 ×10 –4 ) (1) 4 = 1.67 – (0.00642) (+36.84) = 1.43 V c. Fe(OH)2(s) + 2e– → Fe(s) + 2OH–(aq) Fe(s) → Fe2+(aq) + 2e– Eo = –0.88 V Eo = +0.44 V Fe(OH)2(s) ¾ Fe2+(aq) + 2OH–(aq) Eo = –0.44 V ∆Go = –nFEo = – RT ln Ksp ln K sp = Page 2 nFE o RT ln K sp = (2 mol) (96,500J ⋅ V –1 ) (–0.44 V) = –34.28 –1 ( 8 . 3 1 4 J⋅ mol ⋅ K –1) (298 K) Ksp = 1.30 × 10 –15 Key for 2001 USNCO National Exam, Part II d. When iron is coated with Sn, the reaction Sn → Sn2+ + 2e– takes place. If the tin coating is broken, the reaction Sn2+ + Fe → Sn + Fe2+ becomes spontaneous. Iron becomes the anode and is oxidized more readily. 5. 6. Note: Balanced equations were not required. a. Ca(HCO3)2 → CaO + H2O + 2CO2 (partial credit for CaCO3) b. K2SO3 + H+ → K+ + H2O + SO2 (partial credit for H2SO3) c. Ba2+ + OH– + H+ + SO42– → H2O + BaSO4 d. Cr2O72– + Sn2+ + H+ → Cr3+ + Sn4+ + H2O (partial credit for SnCl4) e. H+ + Cl– + OCl–→ Cl2 + H2O (partial credit for HClO) f. 16 7N a. Tangent to curve at 10 seconds: b. i. ii. iii. rate = k rate = k [ ] rate = k [ ]2 c. i. Expressed in symbols: ii. d. i. ii. iii. 7. a. → 0 –1 + 16 8O ∆M –0.39M = = –0.012 M⋅ s–1 ∆T 32 s units are M.s–1 units are s–1 units are M–1.s–1 –d[ O 2 ] –d[ HBr ] d[ Br2 ] = = dt 4dt 2dt This shows that the rate of disappearance of HBr is 4 times that of O2 and the rate of production of Br2 is twice the rate of disappearance of O2. More than mono- or bi-molecular steps improbable. rate = k [HBr] [O2] Rate is proportional to reactants in the rate-limiting step. rate = k [HBr]2 [O2] [HOOBr] in the rate equation must be stated in terms of the previous equilibrium. rate = k [HBr]2 [O2]1/2 [HOBr] in the rate equation must be stated in terms of the previous equilibria. Phosphorus (P) and arsenic (As) might behave in this manner. b. i. F X F F ii. iii. c. d. F F F X F F XF3 Electron pair geometry is tetrahedral; atom geometry is trigonal pyramidal. XF5 Electron pair geometry and atom geometry are both trigonal bipyramidal. XF3 X is sp3 hybridized. XF5 X is dsp3 hybridized. Axial bonds in XF5 are longer than the equatorial bonds. The axial bonds are p/d hybrids and the equatorial bonds are s/p2 hybrids. Another acceptable explanation is that the axial bonds at 90o are repelled more than the equatorial bonds at 120o. Y could be N; N has no d orbitals. Another acceptable explanation is that N is too small to accommodate five F atoms. Y could be Bi; Bi can’t easily be oxidized to +5. Key for 2001 USNCO National Exam, Part II Page 3 8. a. CO2 is linear and therefore nonpolar. O C O S SO2 is bent and therefore polar. O b. c. d. Page 4 O The polar substance will bond more strongly and have the lower vapor pressure. The boiling point of HCl is less than the boiling point of HF because HF forms hydrogen bonds which are harder to break than van der Waals forces. Both HCl and HBr are attracted by van der Waals forces. However, HBr has more electrons and therefore has stronger van der Waals forces. As a result, the boiling point of HCl is less than the boiling point of HBr. Ca2+ and O2– ions are attracted about four times as strongly as K+ and F– ions. Ions with a +2 charge are attracted more strongly than ions with a +1 charge. In addition, the calcium-to-oxygen distance is less than the potassium-to-fluoride distance, leading to an increased force of attraction for the shorter bond. Tin(II) chloride is ionic. Tin(IV) chloride is covalent. The +4 charge on tin causes it to attract electrons more strongly from chloride ion, making the bonds covalent. Key for 2001 USNCO National Exam, Part II 2001 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART III Prepared by the American Chemical Society Olympiad Laboratory Practical Task Force OLYMPIAD LABORATORY PRACTICAL TASK FORCE Lucy Pryde Eubanks, Clemson University, Clemson, SC Chair Robert Becker, Kirkwood High School, Kirkwood, MO Craig W. Bowen, US Naval Academy, Annapolis, MD Nancy Devino, ScienceMedia Inc., San Diego, CA Sheldon L. Knoespel, Michigan State University, East Lansing, MI Steve Lantos, Brookline High School, Brookline, MA Jim Schmitt, Eau Claire North High School, Eau Claire, WI Robert G. Silberman, SUNY-Cortland, NY Christie B. Summerlin, University of Alabama-Birmingham, Birmingham, AL DIRECTIONS TO THE EXAMINER–PART III The laboratory practical part of the National Olympiad Examination is designed to test skills related to the laboratory. Because the format of this part of the test is quite different from the first two parts, there is a separate, detailed set of instructions for the examiner. This gives explicit directions for setting up and administering the laboratory practical. There are two laboratory tasks to be completed during the 90 minutes allotted to this part of the test. Students do not need to stop between tasks, but are responsible for using the time in the best way possible. Each procedure must be approved for safety by the examiner before the student begins that procedure. Part III 2 lab problems laboratory practical 1 hour, 30 minutes Students should be permitted to use non-programmable calculators. DIRECTIONS TO THE EXAMINEE–PART III DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. WHEN DIRECTED, TURN TO PAGE 2 AND READ THE INTRODUCTION AND SAFETY CONSIDERATIONS CAREFULLY BEFORE YOU PROCEED. There are two laboratory-related tasks for you to complete during the next 90 minutes. There is no need to stop between tasks or to do them in the given order. Simply proceed at your own pace from one to the other, using your time productively. You are required to have a procedure for each problem approved for safety by an examiner before you carry out any experimentation on that problem. You are permitted to use a non-programmable calculator. At the end of the 90 minutes, all answer sheets should be turned in. Be sure that you have filled in all the required information at the top of each answer sheet. Carefully follow all directions from your examiner for safety procedures and the proper disposal of chemicals at your examining site. Not valid for use as an USNCO National Examination after April 22, 2001. Page 1 2001 UNITED STATES NATIONAL CHEMISTRY OLYMPIAD PART III — LABORATORY PRACTICAL Student Instructions Introduction These problems test your ability to design and carry out laboratory experiments and to draw conclusions from your experimental work. You will be graded on your experimental design, on your skills in data collection, and on the accuracy and precision of your results. Clarity of thinking and communication are also components of successful solutions to these problems, so make your written responses as clear and concise as possible. Safety Considerations You are required to wear approved eye protection at all times during this laboratory practical. You also must follow all directions given by your examiner for dealing with spills and with disposal of wastes. Lab Problem 1 Design and carry out an experiment to determine the density of the plastic object you have been given. You may use water and the alcohol solution provided at your lab station, as well as the equipment you will find there, but you may not use a balance. You will be asked to describe the method you developed to solve this problem. Given: density of water = 1.00 g·mL–1 density of alcohol solution = 0.85 g·mL–1 Lab Problem 2 Design and carry out an experiment to determine the specific identity of the compound in each of eight numbered vials. Each vial contains one of these ionic compounds. BaCl2, CaCO3, Ca(OH)2, KI, NaCl, NaHCO3, Na2SO4, Pb(NO3)2 In addition to the equipment you will find at your lab station, you may use distilled water. You also have the option of choosing ONE additional reagent from this list. You may do this either before or during your experimentation. 6 M H2SO4, 6 M HCl, 6 M NaOH, phenolphthalein indicator solution You will be asked to describe the method you developed to solve this problem. Page 2 Not valid for use as an USNCO National Examination after April 22, 2001. Answer Sheet for Laboratory Practical Problem 1 Student's Name: __________________________________________________________________________ Student's School: ________________________________________Date: ___________________________ Proctor's Name:__________________________________________________________________________ ACS Section Name : _______________________________ Student's USNCO test #: ________________ 1. Give a brief description of your experimental plan. Before beginning your experiment, you must get approval (for safety reasons) from the examiner. Examiner’s Initials: Not valid for use as an USNCO National Examination after April 22, 2001. Page 3 2. Record your data and other observations. 3. What is the density of the plastic object? Show your methods clearly. Page 4 Not valid for use as an USNCO National Examination after April 22, 2001. Answer Sheet for Laboratory Practical Problem 2 Student's Name: __________________________________________________________________________ Student's School: ________________________________________Date: ___________________________ Proctor's Name:__________________________________________________________________________ ACS Section Name : _______________________________ Student's USNCO test #: ________________ 1. Give a brief description of your experimental plan. Before beginning your experiment, you must get approval (for safety reasons) from the examiner. Examiner’s Initials: When you wish to request the optional reagent, return to the Examiner with this sheet. I request this additional reagent: ______________ Examiner’s Initials: 2. Record your data and other observations. Not valid for use as an USNCO National Examination after April 22, 2001. Page 5 2. Record your data and other observations. (continued) 3. Identify the substance in each numbered vial, giving a brief justification for that choice. Vial # Contains Justification 1. 2. 3. 4. 5. 6. 7. 8. Page 6 Not valid for use as an USNCO National Examination after April 22, 2001. 2001 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART III Prepared by the American Chemical Society Olympiad Laboratory Practical Task Force Examiner's Directions Thank you for administering the 2001 USNCO laboratory practical on behalf of your Local Section. It is essential that you follow the instructions provided, in order to insure consistency of results nationwide. There may be considerable temptation to assist the students after they begin the lab exercise. It is extremely important that you do not lend any assistance or provide any hints whatsoever to the students once they begin work. As is the case with the international competition, students should not be allowed to speak to anyone until the activity is complete. The equipment needed for each student for both lab exercises should be available at his/her lab station or table when the students enter the room. The equipment should be initially placed so that the materials used for Lab Problem 1 are separate from those used for Lab Problem 2. After the students have settled, read the following instructions (in italics) to the students. Hello, my name is ________. Welcome to the lab practical portion of the U.S. Chemistry Olympiad National Examination. In this part of the exam, we will be assessing your lab skills and your ability to reason through a laboratory problem and communicate your results. Do not touch any of the equipment in front of you until you are instructed to do so. One of this year’s problems uses this type of plastic screw anchor as the object being investigated. Show the type of plastic screw anchors being used at your site. (See picture on page 3 of these instructions for approximate size of the anchors.) Another of this year’s problems uses small-scale chemistry equipment. Small-scale chemistry techniques help to minimize the amount of materials you use, thereby increasing safety and minimizing waste. Specialized equipment for small-scale chemistry that you will use today include Beral-type pipets and reaction plates. Show a 1-mL, 3-mL, or a 5-mL Beral-type pipet, and show a 6-well or a 12-well reaction plate. (If your testing site has substituted 50-mL or 100-mL beakers, show those instead.) You will be asked to complete two laboratory problems. The materials and equipment needed to solve each problem has been set out for you and is grouped by the number of the problem. You also may use distilled (or deionized) water. You must limit yourself to this equipment and materials for each problem. A balance may not be used for either problem. You may choose to start with either problem. You are required to have a procedure for each problem approved for safety by an examiner. (Remember that approval does not mean that your procedure will be successful–it is a safety approval.) When you are ready for an examiner to come to your station for each safety approval, please raise your hand. In the second problem, you have the option of choosing ONE additional reagent, either before or during your experimentation. Again, when you are ready to make this choice, please write the formula Examiner’s Directions, 2001 USNCO National Exam, Part III Page 1 of the reagent being requested on your report sheet, and raise your hand. You will have one hour and thirty minutes to complete both problems. Safety is an important consideration during the lab practical. You must wear goggles at all times. Wash off any chemicals spilled on your skin or clothing with large amounts of tap water. The appropriate procedures for disposing of solutions at the end of this lab practical are: ____________________________________________________________________________________ ____________________________________________________________________________________ We are about to begin the lab practical. Please do not turn the page until directed to do so, but read the directions on the front page. Are there any questions before we begin? Distribute Part III booklets and again remind students not to turn the page until the instruction is given. Part III contains student instructions and answer sheets for both laboratory problems. Allow students enough time to read the brief cover directions. Do not turn to page 2 until directed to do so. When you start to work, be sure that you fill out all information at the top of the answer sheets. Are there any additional questions? If there are no further questions, the students should be ready to start Part III. You may begin. After one hour and thirty minutes, give the following directions. This is the end of the lab practical. Please stop and bring me your answer sheets. Thank you for your cooperation during this test. Collect all the lab materials. Make sure that the student has filled in his or her name and other required information on the answer sheets. At this point, you may want to take five or ten minutes to discuss the lab practical with the students. They can learn about possible observations and interpretations and you can acquire feedback as to what they actually did and how they reacted to the problems. After this discussion, please take a few minutes to complete the Post-Exam Questionnaire; this information will be extremely useful to the Olympiad Laboratory Practical subcommittee as they prepare next year’s exam. Please remember to return the post-exam Questionnaire, the answer sheets from Part III, the Scantron sheets from Part I, and the “Blue Books” from Part II to this address: ACS DivCHED Exams Institute Clemson University 223 Brackett Hall Box 340979 Clemson, SC 29634-0979 Wednesday, April 25, 2001 is the absolute deadline for receipt of the exam materials at the Examinations Institute. Materials received after this deadline CANNOT be graded. THERE WILL BE NO EXCEPTIONS TO THIS DEADLINE DUE TO THE TIGHT SCHEDULE FOR GRADING THIS EXAMINATION. Page 2 Examiner’s Directions, 2001 USNCO National Exam, Part III EXAMINER’S NOTES Lab Problem #1: Materials and Equipment Note: Students will NOT be allowed to use a balance for this lab problem. Be sure none are available in the testing area or secure them so they may not be used. Each student will need: 2 10-mL graduated cylinders 2 small beakers (100 mL or 250 mL) one labeled “water”, one labeled “alcohol solution” 2 1-mL Beral-style pipets (Eye droppers may be substituted.) 4 to 6 test tubes, 13 x 100 mm or larger 1 test tube rack 1 250-mL squeeze bottle, labeled “distilled water” or “deionized water” 2 plastic screw anchors (Check your local hardware store for No. 8-10 x 7/8”. These are 2 cm long with a maximum diameter of approximately 0.5 cm. Other sizes varying from No. 4-6 to 10-12 may be used but be sure to check that these sizes fit easily into the test tubes being used.) Lab Problem #1: Chemicals. Each student will need: 250 mL of distilled or deionized water 250 mL of 70% isopropyl alcohol Note: This is sold as “rubbing alcohol” in most stores or pharmacies. You may wish to provide each student with an unopened bottle to emphasize the use of a consumer product. Choose the cheapest brand and check there are no additives such as dyes or perfumes that will change the density. Do not purchase 91% or 99% isopropyl alcohol; these are often available as well. A less desirable alternative, one that does not emphasize the use of a consumer chemical, is to prepare a 70% by volume solution from pure isopropyl alcohol and water, and provide the solution in a 250-mL labeled squeeze bottle. Quick Check to be sure lab problem #1 will work for your examinees: 1) Does the screw anchor fit into the size test tubes being provided? 2) Does the screw anchor float in water, and sink in the alcohol solution? 3) Have you planned to prevent access to all balances in the working area? Lab Problem #1: Notes 1. Note that the examiner will need to initial each student’s experimental plan. Please do not comment on the plan other than looking for any potentially unsafe practices. 2. Safety: It is your responsibility to ensure that all students wear safety goggles during the lab practical. A lab coat or apron for each student is desirable but not mandatory. You will also need to give students explicit directions for handling spills and for disposing of waste materials, following approved safety practices for your examining site. Please check and follow procedures appropriate for your site. Examiner’s Directions, 2001 USNCO National Exam, Part III Page 3 Lab Problem #2: Materials and Equipment. Each student will need: 8 numbered small vials with tops (30-mL plastic vials work well) 1 100-mL or larger wash bottle, labeled “distilled water” or “tap water” 1 24-well reaction plate or 2 12-well plates. If reaction plates cannot be obtained or borrowed, 6 50-mL or 100-mL beakers can be substituted. stirring sticks such as wooden or plastic toothpicks, or coffee stirers 8 1-mL Beral-style pipets, cut to use as scoops (or 8 small spatulas or scoops) 2 1-mL Beral-style pipets (Eye droppers may be substituted.) 1 1-mL Beral-style pipet with label (This will contain 6 M H2SO4, 6 M HCl, 6 M NaOH, or phenolphthalein indicator solution) 1 container designated for disposal of heavy metal waste of Pb2+ and Ba2+ supply of paper towels 1 pair safety goggles 1 lab coat or apron (optional) Lab Problem #2: Chemicals. Each student will need: 1 set of filled, numbered vials. Each numbered vial will contain about 1 g of one of these dry solids. Note: This is the order for filling the numbered vials. Please keep this list secure. 1. NaCl 4. BaCl2 7. Pb(NO3)2 2. CaCO3 5. NaHCO3 8. KI 3. Na2SO4 6. Ca(OH)2 Please have available 100 mL of each of these reagents: 6 M H2SO4, 6 M HCl, 6 M NaOH, and phenolphthalein indicator solution. Students will be asked to choose ONE of these reagents for use with lab problem #2, either before starting experimentation or during their work. You may find it convenient to pre-fill a set of labeled Beral-style pipets for each student but they must not be supplied at the lab station. Supply of distilled water, if available; deionized water may also be used Quick Check to be sure this lab problem will work for your examinees: 1) Are all the solids dry? 2) CaCO3 needs to be provided in powdered form, not as marble chips that are sometimes used. 3) Are you prepared to dispense H2SO4, HCl, NaOH, or phenolphthalein indicator solution quickly and safely when the students have made their choice? 4) Are you prepared to collect all solutions containing Pb2+ and Ba2+ metal ions? Lab Problem #2: Notes 1. Note that the examiner will need to initial each student’s experimental plan. Please do not comment on the plan other than looking for any potentially unsafe practices. The examiner also will need to initial each student’s choice of additional reagent. 2. Safety: It is your responsibility to ensure that all students wear safety goggles during the lab practical. A lab coat or apron for each student is desirable but not mandatory. You will also need to give students explicit directions for handling spills and for disposing of waste materials, following approved safety practices for your examining site. Please check and follow procedures appropriate for your site. Page 4 Examiner’s Directions, 2001 USNCO National Exam, Part III 2001 U. S. NATIONAL CHEMISTRY OLYMPIAD KEY for NATIONAL EXAM—PART III Lab Problem 1 Either of two general plans might be used to solve this problem. Multiple trials are expected for either plan used. Plan A: Make a solution of water and alcohol in which the object is suspended. Using the measured volumes of the water and alcohol, and their given densities, the density of the solution can be found. The density of the solution and the object will be the same. Plan A Sample Data: Trial #1 Trial #2 Volume water to suspend object 2.21 mL 2.23 mL Volume alcohol to suspend object 3.02 mL 3.03 mL Plan A Calculations: ( ) Mass of solution = VH2 O × densityH 2O + ( Valcohol × densityalcohol ) ( ) Volume of solution = VH2 O + (Valcohol ) (assuming volumes are additive) Mass of solution Density of solution = Volume of solution Plan A Sample Results: Trial #1 Trial #2 Mass of solution (g) 4.78 4.81 Volume of solution (mL) 5.23 5.26 Density of solution (g/mL) 0.91 0.91 Plan B: Using Archimedes’ principle, determine the volume of the object by displacing the alcohol. Determine the mass of the object by displacing water, and multiplying the volume of the water displaced by the density of water. Determine the density of the object by dividing the mass of the object by its volume. Plan B Sample Data: Trial #1 Trial #2 Initial volume water 5.02 mL 5.14 mL Final volume water 5.83 mL 5.94 mL Initial volume alcohol 6.01 mL Final volume alcohol 6.88 mL Plan B Calculations: Mass of object = VH 2 O displaced × density H2 O Mass of object Density of object = Volume of object Plan B Sample Results: Trial #1 6.02 mL 6.89 mL Trial #2 Mass of object (g) 0.81 0.80 Volume of object (mL) 0.87 0.87 Density of object (g/mL) 0.93 0.92 Conclusion: The determined density should be between 0.85 g/mL and 1.00 g/mL, because the object sinks in alcohol and floats in water. The density determination must be supported by data gathered and calculations performed. Page 1 2001 USNCO National Exam, Part III (Lab Practical) Lab Problem 2 Credit was awarded for alternate, logical pathways that achieved identification of the compounds. The identifications depend on developing a logical sequence of tests that will lead to the identifications. Some type of tabular form organizes the data for clear presentation, even if a formal flow chart is not included. Sample Plan: Many students started by adding water to each compound. Most often, H2SO4 was chosen as the extra reagent Those tests were followed by adding selected solutions of the unknown to each other to help with identification. Sample Data: FIRST #1 #2 #3 #4 #5 #6 #7 #8 TESTS H2O sol insol sol sol sol insol sol sol H2SO4 no rxn dissolves, no rxn white ppt bubbles dissolves white ppt yellow bubbles The first set of tests allows identification of #2 as CaCO3, #5 as NaHCO3, and #6 as Ca(OH)2. SECOND TESTS #1 #3 – no rxn – #1 #3 #4 #4 no rxn white ppt – #7 #8 #7 #8 white ppt white ppt no rxn no rxn white ppt no rxn – yellow ppt – The second set of tests allows identification of #1 as NaCl, #7 as Pb(NO3)2, and #8 as KI. There is still ambiguity about #3 and #4 at this point. THIRD TESTS H2SO4 #3 no rxn #4 white ppt The third set of tests allows identification of #3 as Na2SO4 and #4 as BaCl2. Conclusion: Identification Page 2 Substance Supporting Evidence #1 NaCl Second tests: forms a white precipitate with #7, Pb(NO3)2 #2 #3 #4 CaCO3 Na2SO4 BaCl2 First tests: insoluble in water, dissolves and bubbles in acid Third tests: no reaction with H2SO4 Third tests: white precipitate forms #5 NaHCO3 First tests: soluble in water, dissolves and bubbles in acid #6 #7 Ca(OH)2 Pb(NO3)2 First tests: insoluble in water, dissolves in acid Second tests: forms three white precipitates and one yellow precipitate #8 KI Second tests: forms one yellow precipitate 2001 USNCO National Exam, Part III (Lab Practical) 2002 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART I Prepared by the American Chemical Society Olympiad Examinations Task Force OLYMPIAD EXAMINATIONS TASK FORCE Arden P. Zipp, State University of New York, Cortland Chair Peter E. Demmin (retired), Amherst Central High School, NY Dianne H. Earle, Paul M. Dorman High School, SC David W. Hostage, Taft School, CT Alice Johnsen, Bellaire High School, TX Elizabeth M. Martin, College of Charleston, SC Jerry D. Mullins, Plano Senior High School, TX Ronald O. Ragsdale, University of Utah, UT DIRECTIONS TO THE EXAMINER–PART I Part I of this test is designed to be taken with a Scantron® answer sheet on which the student records his or her responses. Only this Scantron sheet is graded for a score on Part I. Testing materials, scratch paper, and the Scantron sheet should be made available to the student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until April 21, 2002, after which tests can be returned to students and their teachers for further study. Allow time for the student to read the directions, ask questions, and fill in the requested information on the Scantron sheet. The answer sheet must be completed using a pencil, not pen. When the student has completed Part I, or after 1 hour, 30 minutes has elapsed, the student must turn in the Scantron sheet, Part I of the testing materials, and all scratch paper. There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and you are free to schedule rest-breaks between parts. Part I Part II Part III 60 questions 8 questions 2 lab problems single-answer multiple-choice problem-solving, explanations laboratory practical 1 hour, 30 minutes 1 hour, 45 minutes 1 hour, 30 minutes A periodic table and other useful information are provided on page 2 for student reference. Students should be permitted to use nonprogrammable calculators. DIRECTIONS TO THE EXAMINEE–PART I DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Answers to questions in Part I must be entered on a Scantron answer sheet to be scored. Be sure to write your name on the answer sheet; an ID number is already entered for you. Make a record of this ID number because you will use the same number on both Parts II and III. Each item in Part I consists of a question or an incomplete statement that is followed by four possible choices. Select the single choice that best answers the question or completes the statement. Then use a pencil to blacken the space on your answer sheet next to the same letter as your choice. You may write on the examination, but the test booklet will not be used for grading. Scores are based on the number of correct responses. When you complete Part I (or at the end of 1 hour, 30 minutes), you must turn in all testing materials, scratch paper, and your Scantron answer sheet. Do not forget to turn in your U.S. citizenship statement before leaving the testing site today. Not valid for use as an USNCO National Exam after April 21, 2002. Distributed by the ACS DivCHED Examinations Institute, Clemson University, Clemson, SC. All rights reserved. Printed in U.S.A. amount of substance ampere atmosphere atomic mass unit atomic molar mass atomic number Avogadro constant Celsius temperature centi- prefix coulomb electromotive force energy of activation enthalpy entropy ABBREVIATIONS AND SYMBOLS n equilibrium constant K measure of pressure mmHg A Faraday constant F milli- prefix m atm formula molar mass M molal m u free energy M G molar A frequency mol ν mole Z gas constant h R Planck’s constant N A gram pressure P g °C heat capacity k C p rate constant c hour Rf h retention factor C joule s J second E kelvin speed of light c K Ea kilo- prefix temperature, K T k H liter t L time S volt V E =E – USEFUL EQUATIONS –∆H 1 ln K = +c R T RT lnQ nF CONSTANTS R = 8.314 J·mol –1·K–1 R = 0.0821 L·atm·mol –1·K–1 1 F = 96,500 C·mol –1 1 F = 96,500 J·V–1·mol–1 N A = 6.022 × 10 23 mol–1 h = 6.626 × 10 –34 J·s c = 2.998 × 10 8 m·s –1 1 atm = 760 mmHg k 2 Ea 1 1 = − k1 R T1 T2 ln PERIODIC TABLE OF THE ELEMENTS 1 H 2 He 1.008 4.003 3 Li 4 Be 5 B 6 C 7 N 8 O 9 F 10 Ne 6.941 9.012 10.81 12.01 14.01 16.00 19.00 20.18 11 Na 12 Mg 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar 22.99 24.31 26.98 28.09 30.97 32.07 35.45 39.95 19 K 20 Ca 21 Sc 22 Ti 23 V 24 Cr 25 Mn 26 Fe 27 Co 28 Ni 29 Cu 30 Zn 31 Ga 32 Ge 33 As 34 Se 35 Br 36 Kr 39.10 40.08 44.96 47.88 50.94 52.00 54.94 55.85 58.93 58.69 63.55 65.39 69.72 72.61 74.92 78.96 79.90 83.80 37 Rb 38 Sr 39 Y 40 Zr 41 Nb 42 Mo 43 Tc 44 Ru 45 Rh 46 Pd 47 Ag 48 Cd 49 In 50 Sn 51 Sb 52 Te 53 I 54 Xe 85.47 87.62 88.91 91.22 92.91 95.94 (98) 101.1 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.3 55 Cs 56 Ba 57 La 72 Hf 73 Ta 74 W 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 132.9 137.3 138.9 178.5 181.0 183.8 186.2 190.2 192.2 195.1 197.0 200.6 204.4 207.2 209.0 (209) (210) (222) 87 Fr 88 Ra 89 Ac 104 Rf 105 Db 106 Sg 107 Bh 108 Hs 109 Mt 110 111 112 114 (223) 226.0 227.0 (261) (262) (263) (262) (265) (266) (269) (272) (277) (289) Page 2 58 Ce 59 Pr 60 Nd 61 Pm 62 Sm 63 Eu 64 Gd 65 Tb 66 Dy 67 Ho 68 Er 69 Tm 70 Yb 71 Lu 140.1 140.9 144.2 (145) 150.4 152.0 157.3 158.9 162.5 164.9 167.3 168.9 173.0 175.0 90 Th 91 Pa 92 U 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr 232.0 231.0 238.0 237.0 (244) (243) (247) (247) (251) (252) (257) (258) (259) (260) Not valid for use as a USNCO National Exam after April 21, 2002. DIRECTIONS § When you have selected your answer to each question, blacken the corresponding space on the answer sheet using a soft, #2 pencil. Make a heavy, full mark, but no stray marks. If you decide to change an answer, erase the unwanted mark very carefully. § You may write on the test booklet, but it will not be used for grading. § There is only one correct answer to each question. Any questions for which more than one response has been blackened will not be counted. § Your score is based solely on the number of questions you answer correctly. It is to your advantage to answer every question. 1. Which element commonly exhibits both +1 and +3 oxidation states? (A) Al (Z = 13) (B) Sc (Z = 21) (C) Sn (Z = 50) (D) Tl (Z = 81) 6. A weighed quantity of a gas is collected over water at 25 °C and 742 mmHg. The molar mass of the gas is to be determined at standard temperature and pressure. If the vapor pressure of water is ignored during the calculation, what is the effect on the calculated pressure and calculated molar mass of the gas? pressure 2. Which procedure is best to extinguish burning magnesium? molar mass (A) low low (A) Add water to it. (B) low high (B) Blow nitrogen gas over it. (C) high low (C) Cover it with sand. (D) high high (D) Throw ice on it. 7. A 0.1 M solution of which substance is most acidic? 3. Which two sets of reactants best represent the amphoterism of Zn(OH)2? Set 1. Zn(OH)2(s) and OH–(aq) Set 2. Zn(OH)2(s) and H2O(l) Set 3. Zn(OH)2(s) and H+(aq) Set 4. Zn(OH)2(s) and NH3(aq) (A) Sets 1 and 2 (B) Sets 1 and 3 (C) Sets 2 and 4 (D) Sets 3 and 4 (A) NaHSO 4 (B) Na 2SO4 (C) NaHS (D) NaHCO3 8. The mineral trona has the formula Na 2CO3. NaHCO3. 2H2O and a formula mass of 226 g. mol –1. How many mL of 0.125 M HCl are needed to convert all the carbonate and bicarbonate in a 0.407 g sample of trona into carbon dioxide and water? 4. Which of these statements about sulfur is not correct? (A) 43.2 mL (B) 28.8 mL (C) 21.6 mL (D) 14.4 mL (A) It exists in different allotropic forms. (B) It can behave as either an oxidizing agent or a reducing agent. 9. The percentages by mass of C, H, and Cl in a compound are C 52.2%, H 3.7%, and Cl 44.1%. How many carbon atoms are in the simplest formula of the compound? (C) It can form up to six covalent bonds in compounds. (A) 3 (B) 4 (C) 6 (D) 7 (D) It is a liquid at 25 °C and 1 atm pressure. 10. 5. A solution of sulfuric acid in water that is 25% H2SO4 by mass has a density of 1.178 g·mL–1. Which expression gives the molarity of this solution? (A) 0.25 × 98 × 1178 (B) 0.25 ×1178 98 0.25 98 × 1178 (D) 1178 0.25 × 98 (C) 4KO2(s) + 2CO2(g) → 2K 2CO3(s) + 3O2(g) What is the maximum volume of oxygen that can be produced when 150. mL of CO2 is passed over 0.500 g of KO2? Assume all gases are measured at 0 °C and 1 atm. (A) 118 mL (B) 157 mL (C) 225 mL (D) 475 mL Not valid for use as an USNCO National Examination after April 21, 2002. Page 3 11. The first vertical line in the diagram represents a thermometer with the boiling and freezing points for a pure solvent. The numbered lines represent possible boiling and freezing points for a solution of a nonvolatile solute in the same solvent. Which line best represents the boiling point and freezing point of a solution relative to values for the pure solvent? Note: The differences in temperatures are not to scale. Solvent 1 2 3 4 16. When the temperature of a sample of H2S gas is lowered, the pressure decreases more than predicted by the ideal gas equation. To what is this deviation from expected behavior due? 1. attractive forces between molecules 2. mass of the molecules 3. volume of the molecules (A) 1 only (B) 2 only (C) 1 and 3 only (D) 2 and 3 only bp fp (A) 1 (B) 2 (C) 3 (D) 4 12. Equal masses of gaseous N2, NH3, and N2O are injected into an evacuated container to produce a total pressure of 3 atm. How do the partial pressures of N 2, NH3, and N2O compare? 17. This curve is produced when a pure substance is heated. Which characteristic of this curve is related to the value for the enthalpy of fusion of the substance? F D B C A Heat added (A) length of AB (B) length of BC (C) slope of AB (D) slope of CD 18. Which statement is correct? (A) PN 2 = PNH3 = PN2 O (B) PN 2 < PNH 3 < PN2 O (A) In a coffee-cup calorimeter, q = ∆H. (C) PNH 3 < PN 2 < PN2 O (D) PN 2O < PN 2 < PNH3 (B) In a coffee-cup calorimeter, w = 0. (C) In a bomb calorimeter, q = ∆S. (D) In a bomb calorimeter, w > 0. 13. According to this phase diagram, which phases can exist at pressures lower than the triple point pressure? 19. Consider this reaction. 4PH3(g) + 8 O2(g) → P4O10(s) + 6H2O(g) ∆H o = –4500 kJ Calculate ∆H of of P4O10(s) in kJ·mol–1. Temperature → (A) gas only (B) solid and gas only (C) liquid only (D) solid and liquid only 14. 1.00 g of water is Vapor Pressure, 50 °C introduced into a 5.00 L H2O 92.5 mmHg evacuated flask at 50 °C. What mass of water is present as liquid when equilibrium is established? (A) 0.083 g (B) 0.41 g (C) 0.59 g (D) 0.91 g 15. Which substance has the greatest lattice energy? (A) NaF Page 4 (B) KCl E (C) MgO Substance PH3(g) H2O(g) ∆H of , kJ·mol–1 +9.2 –241.8 (A) –5914 kJ (B) –4751 kJ (C) –4249 kJ (D) –3012 kJ 20. For which substances and conditions can So = 0? I. elements at 0 K II. compounds at 0 K III. gases at 298 K (A) I only (B) III only (C) I and II only (D) I and III only (D) CaS Not valid for use as an USNCO National Examination after April 21, 2002. 21. 50.0 mL of 0.10 M HCl is Solution Values mixed with 50.0 mL of Cp 4.18 J·g–1·°C–1 0.10 M NaOH. The solution temperature rises density 1.0 g·mL–1 by 3.0 °C. Calculate the enthalpy of neutralization per mole of HCl. (A) –2.5 × 102 kJ (B) –1.3 × 102 kJ (C) –8.4 × 10 kJ (D) –6.3 × 10 kJ 1 1 27. The rate of a reaction at 75 °C is 30.0 times that at 25 °C. What is its activation energy? (A) 58.6 kJ·mol–1 (B) 25.5 kJ·mol–1 (C) 7.05 kJ·mol–1 (D) 1.51 kJ·mol–1 28. 6I–(aq) + BrO3–(aq) + 6H+(aq) → 3I 2(aq) + Br –(aq) + 3H2O(l) These data were obtained when this reaction was studied. [I –], M 22. What can be concluded about the values of ∆H and ∆S from this graph? +100 [BrO 3–], M 0.0010 0.0020 0.0020 0.0010 +50 0 –50 –100 0 100 200 300 400 500 Temperature, K → (A) ∆H > 0, ∆S > 0 (B) ∆H > 0, ∆S < 0 (C) ∆H < 0, ∆S > 0 (D) ∆H < 0, ∆S < 0 23. The boiling point of chloroform, CHCl3, is 61.7 °C and its enthalpy of vaporization is 31.4 kJ·mol–1. Calculate the molar entropy of vaporization for chloroform. (A) 10.7 J·mol–1·K–1 (B) 93.8 J·mol–1·K–1 (C) 301 J·mol–1·K–1 (D) 509 J·mol–1·K–1 (B) 1.1 (C) 0.86 (D) 4.5 × 10–6 0.010 0.010 0.010 0.020 Reaction rate, mol·L–1·s–1 8.0 × 10–5 1.6 × 10–4 1.6 × 10–4 1.6 × 10–4 What are the units of the rate constant for this reaction? (A) s–1 (B) mol·L –1·s–1 (C) L·mol–1·s–1 (D) L2·mol–1·s–1 29. Consider this gas phase reaction. Cl2(g) + CHCl3(g) → HCl(g) + CCl4(g) The reaction is found experimentally to follow this rate law. rate = k [CHCl3] [Cl 2]1/2 Based on this information, what conclusions can be drawn about this proposed mechanism? Step 1. Cl2(g) ¾ 2Cl(g) Step 2. Step 3. 24. ∆Go for a reaction at 25 °C is 30.5 kJ·mol–1. What is the value of K? (A) 2.2 × 105 0.0020 0.0020 0.0040 0.0040 [H +], M Cl(g) + CHCl3(g) → HCl(g) + CCl3(g) Cl(g) + CCl 3(g) → CCl 4(g) (A) Step 1 is the rate-determining step. (B) Step 2 is the rate-determining step. (C) Step 3 is the rate-determining step. 25. This is the rate law for a reaction that consumes X. rate = k [X]2 Which plot gives a straight line? (A) [X] vs. time (B) ln [X] vs. time (C) 1 / [X] vs. time (D) 1 / ln [X]2 vs. time 26. For a first order reaction, the concentration decreases to 30% of its initial value in 5.0 min. What is the rate constant? (A) 0.46 min –1 (B) 0.24 min –1 (C) 0.14 min –1 (D) 0.060 min –1 (D) The rate-determining step cannot be identified. 30. Determine the value of the equilibrium constant for this reaction 2NOCl(g) + O2(g) ¾ 2NO2(g) + Cl2(g) from the K values for these reactions. 2NOCl(g) ¾ 2NO(g) + Cl2(g) K p = 1.7 × 10 –2 2NO2(g) ¾ 2NO(g) + O2 (g) K p = 5.9 × 10–5 (A) 1.0 × 10–6 (B) 1.0 × 10–3 (C) 3.5 × 10–3 (D) 2.9 × 102 Not valid for use as an USNCO National Examination after April 21, 2002. Page 5 31. What is the pH of a 0.15 M solution of hydrazine, N2H4? (A) 3.41 (B) 6.82 Hydrazine Kb N2H4 1.0 × 10–6 (C) 10.59 (D) 11.00 38. How many moles of electrons must be removed from each mole of toluene, C6H5CH3, when it is oxidized to benzoic acid, C6H5COOH? (A) 1 (B) 2 (C) 4 (D) 6 Questions 39 and 40 refer to the reaction represented by this equation. 2Al(s) + 3Cu2+(aq) → 2Al3+(aq) + 3Cu(s) 39. What is the value of Eo for a voltaic cell based on this reaction? 32. The rates of many catalyzed reactions follow the profile shown in the graph. Why does the reaction rate level off? [Reactant], M → (A) The reactant is used up. (B) The reverse reaction becomes dominant. Reaction Cu (aq) + 2e–→ Cu(s) Al3+(aq) + 3e–→ Al(s) 2+ Eo +0.34 V –1.66 V (A) 1.32 V (B) 2.00 V (C) 2.30 V (D) 4.34 V (C) The catalyst decomposes as the reaction proceeds. (D) The active sites on the catalyst are occupied. Questions 33 and 34 refer to aqueous solutions of formic acid, HCOOH, which has a Ka value of 1.9 × 10 –4 at 25 °C. 33. What is the percent ionization of a 0.10 M solution of formic acid at 25 °C? (A) 0.19% (B) 1.4% (C) 4.4% (D) 14% 34. How many moles of sodium formate must be added to 1.0 L of a 0.20 M formic acid solution to produce a pH of 4.00? (A) 0.38 (B) 0.80 (C) 1.9 (D) 3.8 35. During the titration of a weak base with a strong acid, one should use an acid-base indicator that changes color in the (A) acidic range. (B) basic range. (C) buffer range. (D) neutral range. 36. What is the solubility Substance K sp of calcium hydroxide calcium hydroxide 4.0 × 10 –6 in mol·L–1? (A) 1.6× 10 (B) 1.0 × 10 –2 (C) 2.0 × 10 –2 Page 6 (A) 6 (B) 5 (C) 3 (D) 2 41. Use the given standard reduction potentials to determine the reduction potential for this half-reaction. MnO4–(aq) + 3e– + 4H+ → MnO 2(s) + 2H2O(l) Eo Reaction MnO4–(aq) + e–→ MnO 42–(aq) +0.564 V MnO42–(aq) + 2e– + 4H+ → MnO 2(s) + 2H2O(l) +2.261 V (A) 1.695 V (B) 2.825 V (C) 3.389 V (D) 5.086 V 42. How many Faradays are required to reduce all the chromium in 0.150 L of 0.115 M of Cr 2O72– to Cr2+? (A) 0.920 F (B) 0.690 F (C) 0.138 F (D) 0.069 F 43. In which list are the elements arranged in order of increasing first ionization energy? (A) Li, Na, K (B) S, O, F (C) Na, Mg, Al (D) F, Ne, Na (D) 1.0 × 10–3 –3 37. What is the average oxidation number of tungsten in the ion, W6O6Cl122–? (A) 2.7 40. What value should be used for n in the Nernst equation to determine the effect of changes in Al3+(aq) and Cu2+(aq) concentrations in this reaction? (B) 3.3 (C) 3.7 44. Which quantum number is associated with the shape of an atomic orbital? (A) n (B) l (C) ml (D) ms (D) 4.3 Not valid for use as an USNCO National Examination after April 21, 2002. 45. Consider the ions Li+, Na +, Be2+, and Mg2+. Which two are closest to one another in size? (A) Li+ and Na+ (B) Be2+ and Mg2+ (C) Be2+ and Li+ (D) Li+ and Mg2+ 46. What is the electron configuration for a gas phase +3 ion of iron (Z = 26)? (A) [Ar] 3d 5 1 (C) [Ar] 4s 3d 2 (B) [Ar] 4s 3d 4 52. Which species has the strongest oxygen-oxygen bond according to molecular orbital theory? (A) O2 (B) O2– (C) O22– (D) O2+ 53. How many atoms are covalently bonded to the chromium atom in Cr(NH 3)4Cl3? (A) 3 (B) 4 (C) 6 (D) 7 3 (D) [Ar] 4s2 3d 6 47. Magnesium (Z = 12) has isotopes that range from Mg–20 to Mg–31. Only Mg–24, Mg–25, and Mg–26 are not radioactive. What mode of radioactive decay would convert Mg–20, Mg–21, Mg–22, and Mg–23 into stable isotopes most quickly? (A) electron emission (B) alpha particle emission (C) gamma emission (D) positron emission 48. Which oxides exist as individual molecules? 1. Al2O3 2. SiO 2 3. P 4O10 54. When the carbon-oxygen bonds in H3COH, H2CO, and HCO2– are arranged in order of increasing length, what is the correct order? (A) H3COH, H2CO, HCO2– (B) HCO2–, H3COH, H2CO (C) H2CO, HCO2–, H3COH (D) H3COH, HCO2–, H2CO 55. Which reaction is an oxidation? (Only the carboncontaining molecules are shown.) (A) CH2CH2 → CH3CH2OH (A) 2 only (B) 3 only (B) CH3CH2OH → CH 2CHO (C) 1 and 3 only (D) 2 and 3 only (C) CH3CH2OH + HCOOH → CH3CH2OOCH (D) 2 CH3CH2OH → CH3CH2OCH2CH3 49. How many sigma and pi bonds are in this compound? O C OH (A) 9 sigma, 6 pi (B) 10 sigma, 6 pi (C) 10 sigma, 3 pi (D) 15 sigma, 4 pi Use this structure for the indigo molecule to answer questions 56 and 57. O C C N H H N C C O 56. What is the molecular formula of indigo? 50. Which pair of ions has the same shape? (A) CO32– and NO3– (B) CO32– and SO32– (C) NO3– and ClO3– (D) CO32– and ClO3– 51. Which resonance form makes the greatest contribution to the structure of N2O? (A) (C) N N N N O O (B) (D) N N (A) C 8HNO (B) C 16H2N2O2 (C) C 16H10N2O2 (D) C 16H22N2O2 57. What is the hybridization of the carbon atoms bonded to oxygen? (A) sp (B) sp2 (C) sp3 (D) sp3d N O N O 58. Aniline, C 6H5NH2, does not dissolve well in water. Which reagent could be used to increase its aqueous solubility? (A) 1 M HCl (B) 1 M NaOH (C) diethyl ether (D) toluene Not valid for use as an USNCO National Examination after April 21, 2002. Page 7 59. Which molecule reacts most rapidly with water? 60. Which of these elements is found in hemoglobin? (A) CH3CH2CH2Cl (B) CH3CHClCH3 (C) (CH3)2CHCH2Cl (D) (CH3)3CCl (A) Cr (B) Fe (C) Mg (D) Ni END OF TEST Page 8 Not valid for use as an USNCO National Examination after April 21, 2002. U. S. National Chemistry Olympiad – 2002 National Examination—Part I SCORING KEY Number 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. Answer D C B D B C A A D A D D B C C A B A D C Number 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. 31. 32. 33. 34. 35. 36. 37. 38. 39. 40. Answer A A B D C B A C B D C D C A A B C D B A Property of the ACS Society Committee on Education Number 41. 42. 43. 44. 45. 46. 47. 48. 49. 50. 51. 52. 53. 54. 55. 56. 57. 58. 59. 60. Answer A C B B D A D B D A B D C C B C B A D B 2002 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART II Prepared by the American Chemical Society Olympiad Examinations Task Force OLYMPIAD EXAMINATIONS TASK FORCE Arden P. Zipp, State University of New York, Cortland Chair Peter E. Demmin (retired), Amherst Central High School, NY Dianne H. Earle, Paul M. Dorman High School, SC David W. Hostage, Taft School, CT Alice Johnsen, Bellaire High School, TX Elizabeth M. Martin, College of Charleston, SC Jerry D. Mullins, Plano Senior High School, TX Ronald O. Ragsdale, University of Utah, UT DIRECTIONS TO THE EXAMINER–PART II Part II of this test requires that student answers be written in a response booklet of blank pages. Only this “Blue Book” is graded for a score on Part II. Testing materials, scratch paper, and the “Blue Book” should be made available to the student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until April 21, 2002, after which tests can be returned to students and their teachers for further study. Allow time for the student to read the directions, ask questions, and fill in the requested information on the “Blue Book”. When the student has completed Part II, or after 1 hour, 45 minutes has elapsed, the student must turn in the “Blue Book”, Part II of the testing materials, and all scratch paper. Be sure that the student has supplied all of the information requested on the front of the “Blue Book,” and that the same identification number used for Part I has been used again for Part II. There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and you are free to schedule rest-breaks between parts. Part I Part II Part III 60 questions 8 questions 2 lab problems single-answer multiple-choice problem-solving, explanations laboratory practical 1 hour, 30 minutes 1 hour, 45 minutes 1 hour, 30 minutes A periodic table and other useful information are provided on the back page for student reference. Students should be permitted to use non-programmable calculators. DIRECTIONS TO THE EXAMINEE–PART II DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Part II requires complete responses to questions involving problem-solving and explanations. 1 hour, 45 minutes is allowed to complete this part. Be sure to print your name, the name of your school, and your identification number in the spaces provided on the “Blue Book” cover. (Be sure to use the same identification number that was coded onto your Scantron® sheet for Part I.) Answer all of the questions in order, and use both sides of the paper. Do not remove the staple. Use separate sheets for scratch paper and do not attach your scratch paper to this examination. When you complete Part II (or at the end of 1 hour, 45 minutes), you must turn in all testing materials, scratch paper, and your “Blue Book.” Do not forget to turn in your U.S. citizenship statement before leaving the testing site today. Not valid for use as an USNCO National Exam after April 21, 2002. Distributed by the ACS DivCHED Examinations Institute, Clemson University, Clemson, SC. All rights reserved. Printed in U.S.A. 1. 2. (12%) The percentages of NaHCO3 and Na2CO3 are to be determined in a mixture of them with KCl. A 0.500 g sample of the mixture is dissolved in 50.0 mL of deionized water and titrated with 0.115 M HCl, resulting in this pH titration curve. a. Write a balanced equation for the reaction that is responsible for the equivalence point that occurs at about i. pH = 9 ii. pH = 5 b. Calculate the total number of moles of acid used to reach each equivalence point if the volumes are 9.63 mL and 34.27 mL, respectively. c. Determine the number of grams of Na2CO3 and NaHCO3 and their percentages in the original mixture. d. Sketch a titration curve for a solution of Na2CO3 by itself and describe how it differs from the given curve. (15%) Consider the formation of N2O5(g) by this reaction. 2NO2(g) + 1 /2O2(g) → N 2O5(g) For this reaction, ∆H o = –55.1 kJ and ∆S o = –227 J·K–1 Additional data are given in the Table. 14 12 10 8 6 4 2 0 0 10.0 20.0 30.0 Volume HCl(aq), mL → Type of Data Substance ∆Hf o NO2(g) +33.2 kJ·mol –1 So NO2(g) 239.7 J·mol–1·K–1 So O2(g) 205.1 J·mol–1·K–1 a. Calculate these values. 40.0 Value i. ∆Hf o of N2O5(g) ii. S o of N2O5(g) iii. ∆G o of the given reaction at 25 °C iv. K p of the given reaction at 25 °C b. State and explain i. whether this reaction is spontaneous at 25 °C. ii. how the numerical value of K p would be affected by an increase in temperature. iii. how the relative amounts of reactant and product molecules would be affected by an increase in temperature. iv. why the S o values differ for NO2(g) and O2(g) at 25 °C. 3. (11%) Calcium ions form slightly soluble compounds with phosphate ions such as PO43–, HPO42–, and H2PO4–. a. Write the formula and give the K sp expression for the compound formed by Ca2+ and each of these two ions. i. PO43– ii. H2PO4– b. Calculate the equilibrium concentration of Ca 2+ in a saturated solution with each of the phosphate ions given in part a. i. K sp for calcium phosphate equals 1.0 × 10–25. ii. K sp for calcium dihydrogen phosphate equals 1.0 × 10–3. + c. Determine the [H ] needed to just prevent precipitation by H2PO4– in a 0.25 M H3PO4 solution that has [Ca2+] = 0.15 M. The Ka1 of H 3PO4 is 7.1 × 10 –3. 4. (14%) This reaction can be used to analyze for iodide ion. [I –], M [IO3–], M [H +], M Reaction rate, mol·L–1·s–1 IO3–(aq) + 5I–(aq) + 6H+(aq) → 3I2(aq) + 3H2O(l) 0.010 0.10 0.010 0.60 When the rate of this reaction was studied at 25 °C, the 0.040 0.10 0.010 2.40 results in the table were obtained. 0.010 0.30 0.010 5.40 a. Use these data to determine the order of the reaction 0.010 0.10 0.020 2.40 with respect to each of these species. Outline your reasoning in each case. i. I– ii. IO 3– iii. H + b. Calculate the specific rate constant for this reaction and give its units. c. Based on the kinetics, discuss the probability of this reaction occurring in a single step. d. The kinetics of reactions are often studied under pseudo first-order conditions. Describe what is meant by the term pseudo first order and illustrate how the reaction conditions above would be changed so that the [I–] would be pseudo first order. e. The activation energy for this reaction was found to be 84 kJ·mol –1 at 25 °C. How much faster would this reaction proceed if the activation energy were lowered by 10 kJ·mol–1 (for example, by using a suitable catalyst)? Page 2 Not valid for use as an USNCO National Exam after April 21, 2002 5. (12%) Write net equations for each of these reactions. Use appropriate ionic and molecular formulas for the reactants and products. Omit formulas for all ions or molecules that do not take part in a reaction. Write structural formulas for all organic substances. You need not balance the reactions. All reactions occur in aqueous solution unless otherwise indicated. a. Concentrated hydrochloric acid is added to solid manganese(IV) oxide. b. Solutions of magnesium sulfate and barium hydroxide are mixed. c. Solid barium peroxide is added to water. d. A piece of copper metal is added to a solution of dilute nitric acid. e. A solution of sodium thiosulfate is added to a suspension of solid silver bromide. f. 2-butanol is heated with a solution of acidified potassium dichromate. 6. (12%) Chlorine trifluoride, ClF 3, is a vigorous fluorinating agent that has been used to separate uranium from the fission products in spent nuclear fuel rods. a. Write a Lewis dot structure for ClF3. b. Sketch and describe clearly the geometry for the ClF3 molecule. c. Sketch one other possible geometry and explain why it is not observed. d. Identify the hybrid orbitals that are considered to be used by the chlorine atom in ClF3. e. The electrical conductance of liquid ClF3 is only slightly lower than that of pure water. This behavior is attributed to the self-ionization of ClF3 to form ClF2+ and ClF4–. Sketch and describe the expected structures of ClF2+ and ClF4–. 7. (12%) Answer these questions about the voltaic cell based on these half-reactions. MnO2(s) + 4H+(aq) + 2e– → Mn2+(aq) + 2H2O(l) E o = +1.23 V 2+ + – TiO (aq) + 2H (aq) + 4e → Ti(s) + H2O(l) E o = –0.88 V a. Write the equation for the reaction that produces a positive standard potential and then calculate that potential. b. Identify the half-reaction that occurs at the cathode. Explain. c. Specify the conditions that produce the standard potential. d. State whether each of the changes listed in parts i - iii will affect the standard potential calculated in part a for the assembled cell. For each change, state whether the potential will increase, decrease, or remain the same. Outline your reasoning or show your calculations in each case. i. The [Mn2+] is doubled. ii. The size of the Ti(s) electrode is doubled. iii. The pH of both compartments is increased by the same amount. 8. (12%) Three common allotropic forms of carbon are diamond, graphite, and buckminsterfullerene (C60). a. Describe or sketch clearly the structure of each allotrope. b. Compare diamond and graphite with regard to their hardness and electrical conductivity and account for any differences in behavior on the basis of the structures in part a. c. Graphite is more stable than diamond (by 2.9 kJ·mol–1) at room temperature and pressure. Explain why the diamonds in jewelry do not change readily into graphite. d. Use this phase diagram for carbon to determine which has the greater density, 106 diamond diamond or graphite. Explain your reasoning and suggest a means of liquid converting graphite into diamond. 104 graphite 102 vapor 0 2000 4000 6000 Temperature, oC → END OF PART II Not valid for use as an USNCO National Examination after April 21, 2002. Page 3 amount of substance ampere atmosphere atomic mass unit atomic molar mass atomic number Avogadro constant Celsius temperature centi- prefix coulomb electromotive force energy of activation enthalpy entropy ABBREVIATIONS AND SYMBOLS n equilibrium constant K measure of pressure mmHg A Faraday constant F milli- prefix m atm formula molar mass M molal m u free energy M G molar A frequency mol ν mole Z gas constant h R Planck’s constant N A gram pressure P g °C heat capacity k C p rate constant c hour Rf h retention factor C joule s J second E kelvin speed of light c K Ea kilo- prefix temperature, K T k H liter t L time S volt V E =E – USEFUL EQUATIONS –∆H 1 ln K = +c R T RT lnQ nF CONSTANTS R = 8.314 J·mol –1·K–1 R = 0.0821 L·atm·mol –1·K–1 1 F = 96,500 C·mol –1 1 F = 96,500 J·V–1·mol–1 N A = 6.022 × 10 23 mol–1 h = 6.626 × 10 –34 J·s c = 2.998 × 10 8 m·s –1 1 atm = 760 mmHg k 2 Ea 1 1 = − k1 R T1 T2 ln PERIODIC TABLE OF THE ELEMENTS 1 H 2 He 1.008 4.003 3 Li 4 Be 5 B 6 C 7 N 8 O 9 F 10 Ne 6.941 9.012 10.81 12.01 14.01 16.00 19.00 20.18 11 Na 12 Mg 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar 22.99 24.31 26.98 28.09 30.97 32.07 35.45 39.95 19 K 20 Ca 21 Sc 22 Ti 23 V 24 Cr 25 Mn 26 Fe 27 Co 28 Ni 29 Cu 30 Zn 31 Ga 32 Ge 33 As 34 Se 35 Br 36 Kr 39.10 40.08 44.96 47.88 50.94 52.00 54.94 55.85 58.93 58.69 63.55 65.39 69.72 72.61 74.92 78.96 79.90 83.80 37 Rb 38 Sr 39 Y 40 Zr 41 Nb 42 Mo 43 Tc 44 Ru 45 Rh 46 Pd 47 Ag 48 Cd 49 In 50 Sn 51 Sb 52 Te 53 I 54 Xe 85.47 87.62 88.91 91.22 92.91 95.94 (98) 101.1 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.3 55 Cs 56 Ba 57 La 72 Hf 73 Ta 74 W 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 132.9 137.3 138.9 178.5 181.0 183.8 186.2 190.2 192.2 195.1 197.0 200.6 204.4 207.2 209.0 (209) (210) (222) 87 Fr 88 Ra 89 Ac 104 Rf 105 Db 106 Sg 107 Bh 108 Hs 109 Mt 110 111 112 114 (223) 226.0 227.0 (261) (262) (263) (262) (265) (266) (269) (272) (277) (289) Page 4 58 Ce 59 Pr 60 Nd 61 Pm 62 Sm 63 Eu 64 Gd 65 Tb 66 Dy 67 Ho 68 Er 69 Tm 70 Yb 71 Lu 140.1 140.9 144.2 (145) 150.4 152.0 157.3 158.9 162.5 164.9 167.3 168.9 173.0 175.0 90 Th 91 Pa 92 U 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr 232.0 231.0 238.0 237.0 (244) (243) (247) (247) (251) (252) (257) (258) (259) (260) Not valid for use as an USNCO National Examination after April 21, 2002. 2002 U. S. NATIONAL CHEMISTRY OLYMPIAD KEY for NATIONAL EXAM—PART II 1. a. i. ii. b. c. pH = 9 CO32– + H+ → HCO3– pH = 5 HCO3– + H+ → H2CO3 or HCO3– + H+ → H2O + CO2 0.115 mol 9.63× 10 −3 L × = 1.107× 10−3 mol HCl to titrate CO32– L 0.115 mol 3.427× 10 −3 L × = 3.941× 10 −3 mol HCl to titrate HCO3– L 105.99 g Na 2 CO3 1.107 ×10 −3 mol CO3 2– × = 1.17× 10 –1 g Na2CO 3 mol 84.01 g NaHCO3 2.834 × 10−3 mol HCO3 – × = 2.38× 10 –1 g NaHCO3 mol 1.17 ×10 –1 g Na 2 CO3 × 100 = 23.4% Na 2 CO3 5.00 ×10 –1 g mixture 2.38× 10 –1 g NaHCO3 × 100 = 47.6% NaHCO 3 5.00 × 10–1 g mixture d. The total volume required to reach the second equivalence point is twice that required to reach the first equivalence point because the number of moles of HCO3– is equal to the number of moles of CO32–. 14 12 10 8 6 4 2 0 0 10.0 20.0 30.0 40.0 Volume HCl(aq), mL → 2. a.i. ∆Hfo of N2O5(g) o ∆Hrxn = ∆H of (N 2 O5 ) − 2 ∆H of (NO 2 ) –55.1 kJ = ∆H of (N 2 O5 ) – 2 mol (33.2 kJ·mol–1) ∆H of (N 2 O5 ) = +11.3 kJ·mol–1 ii. S o of N2O5(g) ( o ∆Srxn = So(N 2O 5 ) – 2S o (NO2 ) + S o (O2 ) ) –227.0 J·K–1 = S o(N 2 O 5 ) – [2(–239.7 J·mol–1·K–1) + 1/2(205.1 J·mol–1·K–1)] S o(N 2 O 5 ) = 355.4 J·mol–1·K–1 iii. ∆G o at 25 °C ∆Go = ∆Ho – T∆So o ∆G = –55.1 kJ – (298 K)(–0.227 kJ·K–1) ∆Go = 12.5 kJ iv. Kp at 25 °C ∆Go = – RT ln K p 12500 J = –8.314 J ( 2 9 8 K ) lKn p mol⋅ K Key for 2002 USNCO National Exam, Part II l nK p = –5.045 and K p = 6.44 ×10 –3 Page 1 b. i. This reaction is not spontaneous at 25 °C. The value of ∆G o is positive. ii. An increase in temperature will cause ∆G o to become more positive because the value of ∆S o is negative. Therefore, the numerical value of Kp will decrease. iii. An increase in temperature will cause the relative amount of reactants to increase and products to decrease. This can be explained by noting that the value of ∆Hrxn is negative, which means that adding heat will shift the reaction to the left. Another argument is that as the temperature increases, the value of the equilibrium constant Kp will decrease, also shifting the reaction to the left. iv. The S o values for NO2(g) and O2(g) at the same temperature are not the same because NO2, with 3 atoms per molecule, has more possible arrangements than O2, with only 2 atoms per molecule. This leads to a higher value for entropy, although not very much higher. The molar mass of NO2 is also higher than for O2. 3. a.i. Ca3(PO4)2 ii. Ca(H2PO4)2 b.i. 1.0 × 10 –25 [ ] [PO ] K = [Ca ] [ H PO ] = [Ca ] [ PO ] Then, let 3x = [Ca ] and 2 x = [ PO ] 3 Ksp = Ca 2+ 3– 2 4 2+ 2 sp 2+ 3 4 – 2 3– 2 4 2+ 3 – 4 1.0 × 10 –25 = [3 x ]3 [2x ]2 1.0 × 10 –25 = 108x 5 and x 5 = 9.26 ×10 –28 and x = 3.9 ×10 –6 [Ca ] = 3(3.92 ×10 2+ [ –6 ) =1.2 ×10 –5 M ][ 1.0 × 10 –3 = Ca2 + H 2 PO 4– ]2 [ ] [ Then, let x = Ca 2+ and 2x = H2 PO 4 – ] 1.0 × 10 –3 = [ x ] [2x ]2 1.0 × 10 –3 = 4x3 and x = 6.3 ×10 –2 [Ca ] = 6.3 ×10 2+ c. –2 M Ca(H2PO4)2(s) ¾ Ca2+(aq) + 2H2PO4–(aq) [ ][ 1.0 × 10 –3 = Ca2 + H 2 PO 4– [H 2PO 4– ] = ]2 1.0 ×10–3 = 6.67× 10–3 = 8.2× 10 –2 M 0.15 H3PO4(aq) ¾ H+(aq) + H2PO4–(aq) [ H ][ H PO ] + 7.3 × 10 –3 = 7.3 × 10 –3 = 2 [ H3 PO 4] 4 – [ x] [0.082 + x ] and x = 1.7 ×10–2 M [0.25 − x ] [ ] Therefore, to prevent precipitation, H + must be greater than = 1.7 × 10 –2 M 4. a. i. First order in I–. Compare the results of experiments 1 and 2 to see that the rate went up by a factor of 4 when the concentration of I– went up by 4. ii. Second order in IO3–. Compare the results of experiments 1 and 3 to see that the rate went up by 9 when the concentration of IO3– went up by 3. Page 2 Key for 2002 USNCO National Exam, Part II iii. Second order in H+. Compare the results of experiments 1 and 4 to see that the rate went up by 4 when the concentration of H+ went up by 2. [ ] [IO ] [H ] b. rate = k I – 1 + 2 – 2 3 1 2 2 rate = k [0.01] [0.10 ] [ 0.01] 0.60 k= = 6.0× 10 7 mol–4 ⋅ L4 ⋅ s–1 1.0 ×10 –8 c. Reaction is very unlikely to occur in one step. That would require the simultaneous collision of five particles. d. Pseudo-first order refers to carrying out a reaction under conditions such that only one reactant changes concentration. For this reaction, pseudo-first order kinetics can be established by having a large excess of [IO3–] and [H+]. –E a 2 Ea1 Ea2 RT k Ae k2 – RT – RT e. 2 = and after cancelling A and combining exponents, = e – Ea1 k1 k1 RT Ae Ea Ea k Taking the natural log and rearranging yields this expression. ln 2 = 1 – 2 k1 RT RT ln k2 8.4 ×10 4 J ⋅ mol–1 7.4 × 104 J ⋅mol –1 = – –1 –1 k1 (8.314 J⋅ mol ⋅ K ) ( 2 9 8 K ) (8.314 J ⋅ mol –1 ⋅ K –1 ) ( 2 9 8 K ) ln k k2 = 3 3 . 9 0 –29.87 = 4 . 0 3 a n d 2 = e4.03 k1 k1 k2 = 56.3 k1 5. Note: Balanced equations were not required. a. H+ + Cl– + MnO2 → Cl2 + Mn2+ + H2O b. Mg2+ + SO42– + Ba2+ + OH– → BaSO4 + Mg(OH)2 c. BaO2 + H2O → Ba2+ + OH– + O2H– d. Cu + H+ + NO3– → Cu2+ + NO + H2O e. AgBr + S2O32– → Ag(S2O3)23– + Br– f. H H OH H H H O H H C C C C H + Cr2O72– + H+ → H C C C C H + Cr3+ + H2O H H H H H H H F 6. a. F Cl F b. T-shaped F F Cl F F c. F Cl F or F Cl F F These structures are not favored because they provide less volume for the lone pairs and therefore do not minimize all repulsions. Key for 2002 USNCO National Exam, Part II Page 3 + F d. sp3d Cl e. bent – F square planar F 7. F Cl F F a. Ti + 2MnO2 + 6H+ → 2Mn2+ + 3H2O + TiO2+ Eo = 1.23 V + 0.88 V = 2.11 V b. MnO2 + 4H+ + 2e– → Mn2+ + 2H2O Reduction occurs at the cathode. c. The conditions for standard potential are 25°C, 1 atm pressure, and 1 M concentrations. d. i. Doubling [Mn2+] will decrease the potential because Mn2+ is a product. An increase in Mn2+ will shift the equilibrium to the left. ii. Doubling the size of the electrode has no effect. The electrode does not appear in the equilibrium expression nor in the Nernst equation. iii. Increasing pH lowers [H+]. Because H+ appears on the left of the balanced equation, decreasing [H+] will lower the potential. The reaction shifts to the left. 8. a. diamond sp3 hybridization 3-dimensional tetrahedral network solid graphite sp2 hybridization 2-dimensional sheets trigonal planar covalent half-filled p orbital hexagonal rings C60 spherical shape made up of hexagonal and pentagonal rings “soccer ball” design C C C C C C C C C C b. Diamond is the hardest. Diamond has 4 covalent bonds per C atom, making a very strongly bonded 3-D network solid. Graphite’s sheets have only weak forces between the sheets, allowing one to slide by the other. This makes graphite much “softer” than diamond. All valence electrons in diamond are involved in sigma bonds, resulting in a nonconducting material. Graphite has delocalized electrons in the half-filled p orbitals (pi-bonding), allowing for electron movement from one atom to the next when an electromotive force is applied. Graphite is a good conductor. c. Although equilibrium favors graphite at room temperature, the rate of the reaction from diamond to graphite is extremely slow because of a very high activation energy barrier. d. The graph shows that as pressure is increased at a fixed temperature of 1000 °C, graphite is converted into diamond. Since increasing pressure should favor increasing density, one could conclude that diamond is denser than graphite. Since graphite is composed of sheets with considerable empty space between the carbon layers, converting graphite to the tetrahedral form decreases the empty space and increases density. To prepare diamond from graphite, the graph indicates that by carrying out the process at 0 °C, a pressure of only 104 atm would be needed. However, since the rate of change would be very slow, this might not be the most ideal set of conditions. An alternate method might be to heat graphite to 2500 °C and apply a pressure of 105 atm, which should increase the rate of the conversion. Another alternate method would be to apply a pressure of 103 atm, then heat to 5000 °C to allow for liquefication, increase pressure to 106 atm and then cool to the lower temperature. Page 4 Key for 2002 USNCO National Exam, Part II 2002 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART III Prepared by the American Chemical Society Olympiad Laboratory Practical Task Force OLYMPIAD LABORATORY PRACTICAL TASK FORCE Lucy Pryde Eubanks, Clemson University, Clemson, SC Chair Robert Becker, Kirkwood High School, Kirkwood, MO Nancy Devino, Coordinating Board for Higher Education, Jefferson City, MO Sheldon L. Knoespel, Michigan State University, East Lansing, MI Steve Lantos, Brookline High School, Brookline, MA Jim Schmitt, Eau Claire North High School, Eau Claire, WI Robert G. Silberman, SUNY-Cortland, NY Christie B. Summerlin, University of Alabama-Birmingham, Birmingham, AL DIRECTIONS TO THE EXAMINER–PART III The laboratory practical part of the National Olympiad Examination is designed to test skills related to the laboratory. Because the format of this part of the test is quite different from the first two parts, there is a separate, detailed set of instructions for the examiner. This gives explicit directions for setting up and administering the laboratory practical. There are two laboratory tasks to be completed during the 1 hour, 30 minutes allotted to this part of the test. Students do not need to stop between tasks, but are responsible for using the time in the best way possible. Each procedure must be approved for safety by the examiner before the student begins that procedure. Part III 2 lab problems laboratory practical 1 hour, 30 minutes Students should be permitted to use non-programmable calculators. DIRECTIONS TO THE EXAMINEE–PART III DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. WHEN DIRECTED, TURN TO PAGE 2 AND READ THE INTRODUCTION AND SAFETY CONSIDERATIONS CAREFULLY BEFORE YOU PROCEED. There are two laboratory-related tasks for you to complete during the next 1 hour, 30 minutes. There is no need to stop between tasks or to do them in the given order. Simply proceed at your own pace from one to the other, using your time productively. You are required to have a procedure for each problem approved for safety by an examiner before you carry out any experimentation on that problem. You are permitted to use a non-programmable calculator. At the end of 1 hour, 30 minutes, all answer sheets should be turned in. Be sure that you have filled in all the required information at the top of each answer sheet. Carefully follow all directions from your examiner for safety procedures and the proper disposal of chemicals at your examining site. Not valid for use as an USNCO National Examination after April 21, 2002. Page 1 2002 UNITED STATES NATIONAL CHEMISTRY OLYMPIAD PART III — LABORATORY PRACTICAL Student Instructions Introduction These problems test your ability to design and carry out laboratory experiments and to draw conclusions from your experimental work. You will be graded on your experimental design, on your skills in data collection, and on the accuracy and precision of your results. Clarity of thinking and communication are also components of successful solutions to these problems, so make your written responses as clear and concise as possible. Safety Considerations You are required to wear approved eye protection at all times during this laboratory practical. You also must follow all directions given by your examiner for dealing with spills and with disposal of wastes. Lab Problem 1 Design and carry out an experiment to investigate a relationship between the surface area of a piece of raw potato and the rate of decomposition of hydrogen peroxide. You may use only those materials available at your experimental station. You will be asked to describe the method you developed to carry out this investigation. Lab Problem 2 Design and carry out an experiment to determine the equilibrium constant, Keq, for this reaction at room temperature. urea(s) + H2O(l) ¾ urea(aq) You will be asked to describe the method you developed to solve this problem. Given: Page 2 molar mass of urea, CO(NH2)2 molarity of pure H2O = 60.0 g·mol–1 = 55.5 mol·L–1 Not valid for use as an USNCO National Examination after April 21, 2002. Answer Sheet for Laboratory Practical Problem 1 Student's Name: __________________________________________________________________________ Student's School: ________________________________________Date: ___________________________ Proctor's Name:__________________________________________________________________________ ACS Section Name : _______________________________ Student's USNCO test #: ________________ 1. Give a brief description of your experimental plan. Before beginning your experiment, you must get approval (for safety reasons) from the examiner. Examiner’s Initials: Not valid for use as an USNCO National Examination after April 21, 2002. Page 3 2. Record your data and other observations. 3. What relationship did you have find between the surface area of a raw potato and the rate of decomposition of hydrogen peroxide? Support your conclusion with your experimental evidence. Page 4 Not valid for use as an USNCO National Examination after April 21, 2002. Answer Sheet for Laboratory Practical Problem 2 Student's Name: __________________________________________________________________________ Student's School: ________________________________________Date: ___________________________ Proctor's Name:__________________________________________________________________________ ACS Section Name : _______________________________ Student's USNCO test #: ________________ 1. Give a brief description of your experimental plan. Before beginning your experiment, you must get approval (for safety reasons) from the examiner. Examiner’s Initials: Not valid for use as an USNCO National Examination after April 21, 2002. Page 5 2. Record your data and other observations. 3. What value did you calculate for the equilibrium constant? Show your methods clearly. Page 6 Not valid for use as an USNCO National Examination after April 21, 2002. 2002 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART III Prepared by the American Chemical Society Olympiad Laboratory Practical Task Force Examiner's Directions Thank you for administering the 2002 USNCO laboratory practical on behalf of your Local Section. It is essential that you follow the instructions provided, in order to insure consistency of results nationwide. There may be considerable temptation to assist the students after they begin the lab exercise. It is extremely important that you do not lend any assistance or provide any hints whatsoever to the students once they begin work. As is the case with the international competition, students should not be allowed to speak to anyone until the activity is complete. The equipment needed for each student for both lab exercises should be available at his/her lab station or table when the students enter the room. The equipment should be initially placed so that the materials used for Lab Problem 1 are separate from those used for Lab Problem 2. After the students have settled, read the following instructions (in italics) to the students. Hello, my name is ________. Welcome to the lab practical portion of the U.S. Chemistry Olympiad National Examination. In this part of the exam, we will be assessing your lab skills and your ability to reason through a laboratory problem and communicate your results. Do not touch any of the equipment in front of you until you are instructed to do so. Both of this year’s problems use some small-scale chemistry equipment. Small-scale chemistry techniques help to minimize the amount of materials you use, thereby increasing safety and minimizing waste. Specialized equipment for small-scale chemistry that you will use today include Beral-type pipets. Show the Beral-type pipets being used at your site. If you substituted droppers in the first problem, show those as well. You may be unfamiliar with the graduated centrifuge tubes available in both parts of this lab practical exam. This is the type we will be using. Show the type of centrifuge tube being used at your site. The watertight caps will be an advantage in at least one of the problems. You will be asked to complete two laboratory problems. The materials and equipment needed to solve each problem has been set out for you and is grouped by the number of the problem. You also may use distilled (or deionized) water. You must limit yourself to this equipment and materials for each problem. A balance is not needed for either problem. You may choose to start with either problem. You are required to have a procedure for each problem approved for safety by an examiner. (Remember that approval does not mean that your procedure will be successful–it is a safety approval.) When you are ready for an examiner to come to your station for each safety approval, please raise your hand. You will have one hour and thirty minutes to complete both problems. Examiner’s Directions, 2002 USNCO National Exam, Part III Page 1 Safety is an important consideration during the lab practical. You must wear goggles at all times. Wash off any chemicals spilled on your skin or clothing with large amounts of tap water. The appropriate procedures for disposing of solutions at the end of this lab practical are: ____________________________________________________________________________________ ____________________________________________________________________________________ We are about to begin the lab practical. Please do not turn the page until directed to do so, but read the directions on the front page. Are there any questions before we begin? Distribute Part III booklets and again remind students not to turn the page until the instruction is given. Part III contains student instructions and answer sheets for both laboratory problems. Allow students enough time to read the brief cover directions. Do not turn to page 2 until directed to do so. When you start to work, be sure that you fill out all information at the top of the answer sheets. Are there any additional questions? If there are no further questions, the students should be ready to start Part III. You may begin. After one hour and thirty minutes, give the following directions. This is the end of the lab practical. Please stop and bring me your answer sheets. Thank you for your cooperation during this test. Collect all the lab materials. Make sure that the student has filled in his or her name and other required information on the answer sheets. At this point, you may want to take five or ten minutes to discuss the lab practical with the students. They can learn about possible observations and interpretations and you can acquire feedback as to what they actually did and how they reacted to the problems. After this discussion, please take a few minutes to complete the Post-Exam Questionnaire; this information will be extremely useful to the Olympiad Laboratory Practical subcommittee as they prepare next year’s exam. Please remember to return the post-exam Questionnaire, the answer sheets from Part III, the Scantron sheets from Part I, and the “Blue Books” from Part II to this address: ACS DivCHED Exams Institute Clemson University 223 Brackett Hall Clemson, SC 29634-0979 Wednesday, April 24, 2002 is the absolute deadline for receipt of the exam materials at the Examinations Institute. Materials received after this deadline CANNOT be graded. THERE WILL BE NO EXCEPTIONS TO THIS DEADLINE DUE TO THE TIGHT SCHEDULE FOR GRADING THIS EXAMINATION. Page 2 Examiner’s Directions, 2002 USNCO National Exam, Part III EXAMINER’S NOTES Lab Problem #1: Materials and Equipment. Each student will need: 1 stopwatch, timer, or access to clock with second hand 6 25-mL or 15-mL graduated cylinders, with bases Note: If providing this many graduated cylinders per student is not possible, 6 13 x 100 test tubes and a test tube rack may be substituted. 2 small beakers (100 mL or 250 mL); one labeled “water”, one labeled “3% hydrogen peroxide” 2 1-mL Beral-style pipets (eye droppers may be substituted) 2 15-mL graduated centrifuge tubes (see lab problem #2 for specifications; caps not needed for this lab problem) 1 6-in plastic ruler 1 100-mL or larger wash bottle, labeled “distilled water” or “deionized water” 1 25-mL dropping bottle labeled “liquid detergent” 1 kitchen cutting board (or suitable clean hard surface on lab bench) 1 sharp kitchen paring knife, non-serrated edge 1 plastic container (such as a margarine tub or a Deli salad container); capable of holding approximately 200 mL 1 plastic tub for disposal of liquid wastes (or easy access to sinks) supply of paper towels 1 pair safety goggles 1 lab coat or apron (optional) Lab Problem #1: Chemicals . Each student will need: 1 8-fluid oz (237 mL) bottle of 3% hydrogen peroxide Note: Hydrogen peroxide antiseptic is sold in the first aid section in most supermarkets and drug stores. Provide each student with an unopened bottle to emphasize the use of a consumer product. The cheapest brand, so long as it is fresh and unopened, will work. 1 white potato (Russet potatoes work well and are generally available; provide potato whole and unpeeled) 10 mL liquid detergent 100 mL of distilled or deionized water Quick Check to be sure this lab problem will work for your examinees: 1) Are the bottles of hydrogen peroxide fresh and unopened? 2) Is the knife capable of making clean cuts in the potato? 3) Have all detergent or soap residues been removed from the glassware? Lab Problem #1: Notes 1. Note that the examiner will need to initial each student’s experimental plan. Please do not comment on the plan other than looking for any potentially unsafe practices. 2. Safety: It is your responsibility to ensure that all students wear safety goggles during the lab practical. A lab coat or apron for each student is desirable but not mandatory. You will also need to give students explicit directions for handling spills and for disposing of waste materials, following approved safety practices for your examining site. Please check and follow procedures appropriate for your site. Examiner’s Directions, 2002 USNCO National Exam, Part III Page 3 Lab Problem #2: Materials and Equipment. Each student will need: 4 15-mL centrifuge tubes with 0.1 mL or 0.5 mL graduations. The centrifuge tubes should have conical bottoms and screw caps. 2 tubes labeled “~9.0 mL”; 2 tubes labeled “4.0 g urea” Note: Polystyrene, polypropylene, or Pyrex® centrifuge tubes are widely used in biology, chemistry, and biochemistry departments. The dome-seal screw cap prevents loss of any liquid, important for this experiment. The conical-bottom tubes are preferred to the round bottom tubes. It is not necessary to use the far more expensive borosilicate centrifuge tubes. 1 100-mL or larger wash bottle, labeled “distilled water” or “deionized water” 1 10-mL graduated Beral-style pipet 1 small beaker (100 mL or 250 mL), labeled “water” 1 plastic tub for disposal of liquid wastes (or easy access to sinks) supply of paper towels 1 pair safety goggles 1 lab coat or apron (optional) Lab Problem #2: Chemicals. Each student will need: 2 4.0 g samples of urea, (NH2)2CO, provided in closed, labeled conical-bottom graduated centrifuge tubes. The samples should be prepared in advance. 100 mL of distilled or deionized water Quick Check to be sure this lab problem will work for your examinees: 1) Have the correct centrifuge tubes been obtained and labeled? 2) Have two 4.0 g urea samples been prepared and placed in the labeled centrifuge tubes for each student? 3) Have two other centrifuge tubes been labeled “~9.0 mL H2O” for each student? These should not be filled in advance. Lab Problem #2: Notes 1. Note that the examiner will need to initial each student’s experimental plan. Please do not comment on the plan other than looking for any potentially unsafe practices. 2. Be sure that the labels on the centrifuge tubes do not obscure any graduations. 3. Safety: It is your responsibility to ensure that all students wear safety goggles during the lab practical. A lab coat or apron for each student is desirable but not mandatory. You will also need to give students explicit directions for handling spills and for disposing of waste materials, following approved safety practices for your examining site. Please check and follow procedures appropriate for your site. Page 4 Examiner’s Directions, 2002 USNCO National Exam, Part III 2002 U. S. NATIONAL CHEMISTRY OLYMPIAD KEY for NATIONAL EXAM—PART III Lab Problem 1 Part 1. Experimental Plan A good plan included a detailed description of a method to observe the rate of reaction. It would also include a plan for varying the surface area of the potato. Finally, the plan needed to account for the importance of controlling the volume of peroxide. For example, a good plan might consist of these steps. 1) Cut the potato into different size cubes, such as 2 cubes of 0.5 cm on a side and 2 cubes of 1.0 cm on a side. 2) Measure 2.0 mL of H2O2 into a test tube or graduated cylinder and add 3 drops of detergent.. 3) Shake to generate a small amount of foam for a starting point. Measure the height of the foam column or read the volume of the foam directly if using a graduated cylinder. 4) Drop in one cube of potato and start a timer. 5) At appropriate intervals, measure the height of the foam column or read the volume of the foam directly if using a graduated cylinder. 6) Repeat steps 2-5 for other potato pieces. An average plan was either missing one of these three components or had less detail in two or more of these components. A weak plan had minimal detail about how the experiment would be conducted. Part 2. Experiments and Observations A good experimental section included these points. 1) Appropriate measurements of a. reaction rate or progress. b. dimensions of the potato pieces used. c. time. d.volume of H2O2 used. 2) Multiple (at least two) trials for each different surface area of the potato piece. 3) Appropriate quantitative detail such as a. precision in the trials. b.averaging trial data. c. description of any calculation methods used, such as for determining the surface area of the potato piece. An average experimental section was either missing one of these three components or had less detail in two or more of these components. A weak experimental section had minimal detail about how the experiment was conducted and what observations were made. Part 3. Discussion A good discussion included these points. 1) Calculations or graphical determinations of the relationship between surface area and rate of decomposition of hydrogen peroxide. 2) An appropriate description of the scientific reasoning utilized. 3) An accurate conclusion supported by the experimental observations and data reported. Note: Points were not deducted for correct and reasoned discussion of an experiment with seemingly anomolous results. Page 1 2002 USNCO National Exam, Part III (Lab Practical) Lab Problem 2 Part 1. Experimental Plan A good plan recognized that it was necessary to determine how much water was required to completely dissolve 4.0 g of urea. For example, a good plan might consist of these steps. 1) Add water in small increments to 4.0 g of urea in the graduated centrifuge tube. 2) Cap the tube and shake after each addition. 3) If any solid remains, add another small portion of water. Cap and shake the tube. 4) Continue to add water until all the urea is dissolved. 5) Record the total volume of solution and/or the total volume of water added.* 6) Repeat with the second 4.0 g sample of urea.** An average plan was either missing one of these components or had less detail in two or more of these components. A weak plan had minimal detail about how the experiment would be conducted. Part 2. Experiments and Observations Sample Data Trial 1 Total Volume of Solution* 7.3 mL Mass Urea** 4.0 g Trial 2 7.5 mL 4.0 g Many students also observed that as the urea dissolved, the tube felt cool to the touch. Some allowed time for the tube to return to room temperature before making final observations of volume. Part 3. Calculation of Equilibrium Constant Sample calculations for Trial 1 1 mol urea 1) Moles of urea = 4.0 g urea × = 0.067 mol urea 60.0 g urea 0.067 mol urea = 9.2 M 0.0073 L solution 3) Calculation of Keq if assume that the concentration of water is in standard state. 2) Molarity of urea solution = urea(s) + H2O(l) ¾ urea(aq) Keq = [urea (aq) ] Keq = 9.2 * A superior plan recognized that because there is a high concentration of urea in the saturated solution, the solution cannot be treated as a dilute solution in which the concentration of the solvent is the same as that of pure water. Points were awarded to students who realized this and adjusted their experimental approach. This requires knowing the volume of water added, not just the volume of the resulting solution. For example, if 4.3 mL of water was added, the final volume of the solution was reported as 7.3 mL. The molarity of water in the saturated solution can then be calculated, as shown in this example. 4.3 mL H2 O (9.2 M) [urea (aq)] and Keq = × 5 5 M= 32 M and Keq = = 16 7.3 mL solution (1)(32 M /55.5 M) [ urea (s)] [H 2O(l )] ** Some students elected to obtain more samples by dividing the given mass of urea. The mass of smaller samples of urea was estimated by calculating the density of urea, given the known mass and the volume markings on the graduated centrifuge tube. Page 2 2002 USNCO National Exam, Part III (Lab Practical) 2003 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART I Prepared by the American Chemical Society Olympiad Examinations Task Force OLYMPIAD EXAMINATIONS TASK FORCE Arden P. Zipp, State University of New York, Cortland Chair Peter E. Demmin (retired), Amherst Central High School, NY David W. Hostage, Taft School, CT Alice Johnsen, Bellaire High School, TX Jerry D. Mullins, Plano Senior High School, TX Ronald O. Ragsdale, University of Utah, UT Amy Rogers, College of Charleston, SC DIRECTIONS TO THE EXAMINER–PART I Part I of this test is designed to be taken with a Scantron® answer sheet on which the student records his or her responses. Only this Scantron sheet is graded for a score on Part I. Testing materials, scratch paper, and the Scantron sheet should be made available to the student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until April 27, 2003, after which tests can be returned to students and their teachers for further study. Allow time for the student to read the directions, ask questions, and fill in the requested information on the Scantron sheet. The answer sheet must be completed using a pencil, not pen. When the student has completed Part I, or after one hour and thirty minutes has elapsed, the student must turn in the Scantron sheet, Part I of the testing materials, and all scratch paper. There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and you are free to schedule rest-breaks between parts. Part I Part II Part III 60 questions 8 questions 2 lab problems single-answer multiple-choice problem-solving, explanations laboratory practical 1 hour, 30 minutes 1 hour, 45 minutes 1 hour, 30 minutes A periodic table and other useful information are provided on page 2 for student reference. Students should be permitted to use nonprogrammable calculators. DIRECTIONS TO THE EXAMINEE–PART I DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Answers to questions in Part I must be entered on a Scantron answer sheet to be scored. Be sure to write your name on the answer sheet; an ID number is already entered for you. Make a record of this ID number because you will use the same number on both Parts II and III. Each item in Part I consists of a question or an incomplete statement that is followed by four possible choices. Select the single choice that best answers the question or completes the statement. Then use a pencil to blacken the space on your answer sheet next to the same letter as your choice. You may write on the examination, but the test booklet will not be used for grading. Scores are based on the number of correct responses. When you complete Part I (or at the end of one hour and 30 minutes), you must turn in all testing materials, scratch paper, and your Scantron answer sheet. Do not forget to turn in your U.S. citizenship statement before leaving the testing site today. Not valid for use as an USNCO National Exam after April 27, 2003. Distributed by the ACS DivCHED Examinations Institute, University of Wisconsin-Milwaukee, Milwaukee, WI. All rights reserved. Printed in U.S.A. ABBREVIATIONS AND SYMBOLS n Faraday constant F molal A formula molar mass M molar atm free energy G molar mass u frequency ν mole A gas constant R Planck’s constant N A gram g pressure °C heat capacity C p rate constant c hour h retention factor C joule J second E kelvin K speed of light Ea kilo– prefix k temperature, K H liter L time S measure of pressure mmHg volt K milli– prefix m amount of substance ampere atmosphere atomic mass unit atomic molar mass Avogadro constant Celsius temperature centi– prefix coulomb electromotive force energy of activation enthalpy entropy equilibrium constant CONSTANTS m M M mol h P k Rf s c T t V R = 8.314 J·mol–1·K–1 R = 0.0821 L·atm·mol –1·K–1 1 F = 96,500 C·mol–1 1 F = 96,500 J·V–1·mol–1 N A = 6.022 × 1023 mol–1 h = 6.626 × 10–34 J·s c = 2.998 × 108 m·s–1 0 °C = 273.15 K EQUATIONS E = Eo − 1 1A 1 H k E 1 1 ln 2 = a − k1 R T1 T2 −∆H 1 lnK = + c R T RT ln Q nF PERIODIC TABLE OF THE ELEMENTS 18 8A 2 He 3 Li 2 2A 4 Be 13 3A 5 B 14 4A 6 C 15 5A 7 N 16 6A 8 O 17 7A 9 F 6.941 9.012 10.81 12.01 14.01 16.00 19.00 20.18 11 Na 12 Mg 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar 22.99 24.31 26.98 28.09 30.97 32.07 35.45 39.95 19 K 31 Ga 32 Ge 33 As 34 Se 35 Br 36 Kr 1.008 4.003 10 Ne 20 Ca 3 3B 21 Sc 4 4B 22 Ti 5 5B 23 V 6 6B 24 Cr 7 7B 25 Mn 8 8B 26 Fe 9 8B 27 Co 10 8B 28 Ni 11 1B 29 Cu 12 2B 30 Zn 39.10 40.08 44.96 47.88 50.94 52.00 54.94 55.85 58.93 58.69 63.55 65.39 69.72 72.61 74.92 78.96 79.90 83.80 37 Rb 38 Sr 39 Y 40 Zr 41 Nb 42 Mo 43 Tc 44 Ru 45 Rh 46 Pd 47 Ag 48 Cd 49 In 50 Sn 51 Sb 52 Te 53 I 54 Xe 85.47 87.62 88.91 91.22 92.91 95.94 (98) 101.1 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.3 55 Cs 56 Ba 57 La 72 Hf 73 Ta 74 W 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 132.9 137.3 138.9 178.5 180.9 183.8 186.2 190.2 192.2 195.1 197.0 200.6 204.4 207.2 209.0 (209) (210) (222) 87 Fr 88 Ra 89 Ac 104 Rf 105 Db 106 Sg 107 Bh 108 Hs 109 Mt 110 111 112 114 (223) (226) (227) (261) (262) (263) (262) (265) (266) (269) (272) (277) (2??) Page 2 58 Ce 59 Pr 60 Nd 61 Pm 62 Sm 63 Eu 64 Gd 65 Tb 66 Dy 67 Ho 68 Er 69 Tm 70 Yb 71 Lu 140.1 140.9 144.2 (145) 150.4 152.0 157.3 158.9 162.5 164.9 167.3 168.9 173.0 175.0 90 Th 91 Pa 92 U 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr 232.0 231.0 238.0 (237) (244) (243) (247) (247) (251) (252) (257) (258) (259) (262) Not valid for use as an USNCO National Examination after April 27, 2003 DIRECTIONS ! When you have selected your answer to each question, blacken the corresponding space on the answer sheet using a soft, #2 pencil. Make a heavy, full mark, but no stray marks. If you decide to change an answer, erase the unwanted mark very carefully. ! There is only one correct answer to each question. Any questions for which more than one response has been blackened will not be counted. 1. In an experiment to determine the percentage of water in a solid hydrate by heating, what is the best indication that all the water has been removed? (A) The solid melts. (B) The solid changes color. (C) Water vapor no longer appears. 6. According to the solubility curve shown, how many grams of solute can be recrystallized when 20 mL of a saturated solution at 60 ˚C are cooled to 0 ˚C? (D) Successive weighings give the same mass. 60 40 20 10 K o Temperature, C 2. The curve shown results when a liquid is cooled. What temperature is closest to the freezing point of the liquid? Solubility (g solute / 100 mL soln) ! Your score is based solely on the number of questions you answer correctly. It is to your advantage to answer every question. (A) 7.0 M L N (B) 12 30 50 Temperature ( o C) (C) 25 70 (D) 35 7. Which would produce the largest change in the H2O level when added to water in a 25 mL graduated cylinder? (A) 10.0 g of Hg (d = 13.6 g·mL-1) Time, min (A) L (B) M (C) L + M 2 (D) M + N 2 3. What is the proper way to dispose of a two milliliter sample of hexane after completing experiments with it? (A) Return it to the solvent bottle. (B) Place it in a waste bottle with compatible organic materials. (C) Flush it down the drain with large quantities of water. (D) Pour it on a solid absorbent so it can be thrown away with solid waste. 4. Which anion can undergo both oxidation and reduction? (A) Cr2O72(B) NO3(C) OCl - (D) S 2- 5. The mass percentages in a compound are carbon 57.48%, hydrogen 4.22% and oxygen 38.29%. What is its empirical formula? (A) C 2H2O (B) C 4H3O2 (C) C 5H4O2 (D) C 8H7O4 (B) 7.42 g of Al (d = 2.70 g·mL-1) (C) 5.09 g of iron pyrite (d = 4.9 g·mL-1) (D) 2.68 g of oak (d = 0.72 g·mL -1) 8. Diborane, B2H6, can be prepared by the reaction; 3NaBH 4 + 4BF3 r 3NaBF4 + 2B2H6 If this reaction has a 70 percent yield, how many moles of NaBH4 should be used with excess BF3 in order to obtain 0.200 mol of B2H6? (A) 0.200 mol (B) 0.210 mol (C) 0.300 mol (D) 0.429 mol 9. What volume of 6.0 M H2SO4 should be mixed with 10. L of 1.0 M H2SO4 to make 20. L of 3.0 M H2SO4 upon dilution to volume? (A) 1.7 L (B) 5.0 L (C) 8.3 L (D) 10. L 10. An aqueous solution that is 30.0% NaOH by mass has a density of 1.33 g.mL-1. What is the molarity of NaOH in this solution? (A) 8.25 Not valid for use as an USNCO National Examination after April 27, 2003 (B) 9.98 (C) 16.0 (D) 33.2 Page 3 (A) 1 only (B) 2 only (C) both 1 and 2 (D) neither 1 nor 2 12. Benzene melts at 5.50 ˚C and has a freezing point depression constant of 5.10 ˚C. m-1. Calculate the freezing point of a solution that contains 0.0500 mole of acetic acid, CH 3COOH, in 125 g of benzene if acetic acid forms a dimer in this solvent. (A) 3.46 ˚C (B) 4.48 ˚C (C) 5.24 ˚C (D) 6.01 ˚C (B) C 4H10 (C) C 5H12 Temperature (A) decrease both the melting and boiling points (B) increase both the melting and boiling points (C) increase the melting point and decrease the boiling point (D) decrease the melting point and increase the boiling point 13. A 200. mL sample of a gaseous hydrocarbon has a density of 2.53 g.L-1 at 55 ˚C and 720 mmHg. What is its formula? (A) C 2H6 18. According to the phase diagram, what would be the effect of increasing the pressure on this substance? Pressure 11. Which change 1. an increase in water temperature increases the solubility of a gas in 2. a decrease in gas pressure water? (D) C 6H6 19. When the substances below are arranged in order of increasing entropy values, S˚, at 25 ˚C which is the correct order? (A) CO2(s) < CO2(aq) < CO 2(g) 14. A liquid has a vapor pressure of 40 mmHg at 19.0 ˚C and a normal boiling point of 78.3 ˚C. What is its enthalpy of vaporization in kJ . mol -1? (A) 42.4 (B) 18.4 (C) 5.10 15. Sulfur and fluorine form SF 6 and S 2F 10, both of which are gases at 30 ˚C. When an equimolar mixture of them is allowed to effuse through a pinhole, what is (D) 1.45 Molar Mass g.mol-1 SF 6 S 2F 10 146 254 the ratio SF 6/S2F10 in the first sample that escapes? (A) 1.32/1 (B) 1.74/1 (C) 3.03/1 (C) CO2(s) < CO2(g) < CO2(aq) (D) CO2(g) < CO2(s) < CO2(aq) 20. When 50. mL of 0.10 M HCl is mixed with 50. mL of 0.10 M NaOH the temperature of the solution increases by 3.0 ˚C. Calculate the ∆Hneutralization per mole of HCl. (The solution has a density = 1.0 g.mL-1 and C p = 4.2 J . g-1. ˚C-1) (A) 1.3 × 103 kJ (B) -1.3 × 102 kJ (C) -2.5 × 102 kJ (D) -1.3 × 103 kJ (D) 3.48/1 16. The volumes of real gases often exceed those calculated by the ideal gas equation. These deviations are best attributed to the (A) attractive forces between the molecules in real gases. (B) dissociation of the molecules in real gases. (C) kinetic energy of the molecules in real gases. (D) volumes of the molecules in real gases. 17. The electrical conductivity of a solid is slight at 25 ˚C and much greater at 125 ˚C. The solid is most likely a(n) (A) ionic compound (B) insulator (C) metal (D) semiconductor Page 4 (B) CO2(g) < CO2(aq) < CO 2(s) 21. The combustion of 0.200 mol of liquid carbon disulfide, CS2, to give CO 2(g) and SO2(g) releases 215 kJ of heat. What is ∆Hf˚ for CS2(l) in kJ. mol -1? (A) 385 (B) 87.9 ∆ Hf˚ kJ.mol-1 CO2(g) SO2(g) -393.5 -296.8 (C) -385 (D) -475 22. For the reaction: 2NO2(g) r N2O4(g) ∆H < 0. What predictions can be made about the sign of ∆S and the temperature conditions under which the reaction would be spontaneous? ∆ Srxn Temperature Condition (A) negative low temperatures (B) negative high temperatures (C) positive high temperatures (D) positive low temperatures Not valid for use as an USNCO National Examination after April 27, 2003 23. As ∆G˚ for a reaction changes from a large negative value to a large positive value, K for the reaction will change from (A) a large positive value to a large negative value. (B) a large positive value to a small positive value. 29. For the reaction 2A + 2B r Product the rate law is Rate = k[A][B]2. Which mechanism is consistent with this information? (A) B + B s C (C) a large negative value to a large positive value. (D) a large negative value to a small negative value. C + A r Product (slow) (C) A + A s C 24. ∆E˚ is measured at constant volume and ∆H˚ is measured at constant pressure. For the reaction; 2C (s) + O2(g) r 2CO(g) ∆H˚ < 0 kJ How do the ∆E˚ and ∆H˚ compare for this reaction? (A) ∆E˚ < ∆H˚ (B) ∆E˚ > ∆H˚ (C) ∆E˚ = ∆H˚ (D) Impossible to tell from this information. 25. Which statement about second order reactions is correct? (A) Second order reactions require different reactants. (B) Second order reactions are faster than first order reactions. (C) Second order reactions are unaffected by changes in temperature. (D) The half-life of a second order reaction depends on the initial reactant concentration. 26. A first order reaction has a rate constant of 0.0541 s-1 at 25 ˚C. What is the half-life for this reaction? (A) 18.5 s (B) 12.8 s (C) 0.0781 s (D) 0.0375 s 27. The reaction between NO and I2 is second-order in NO and first-order in I 2. What change occurs in the rate of the reaction if the concentration of each reactant is tripled? (A) 3-fold increase (B) 6-fold increase (C) 18-fold increase (D) 27-fold increase 28. For the rate equation, Rate = k[A][B]2, what are the units for the rate constant, k, if the rate is given in mol . L-1. sec-1? (A) L. mol . sec (C) L2. mol -2. sec-1 (B) L. mol -1. sec-1 (D) L3. mol -3. sec-2 (B) A + B r C (slow) C + B r product (D) A + B s C B+BsD B + C r D (slow) C + D r Product (slow) D + A r product 30. Which straight line gives the activation energy for a reaction? (A) rate constant vs T -1 (C) rate constant vs T (B) ln (rate constant) vs T (D) ln (rate constant) vs T-1 31. Based on the equilibrium constant for the reaction below, 2SO3(g) s 2SO2(g) + O2(g) K = 1.8 × 10-5 what is the equilibrium constant for the reaction SO2 (g) + 1/2O2 (g) s SO3(g) K=? (A) 2.1 × 10-3 (B) 4.2 × 10-3 (C) 2.4 × 102 (D) 5.6 × 104 32. CO(g) + Cl2(g) s COCl(g) + Cl(g) Keq = 1.5 × 10-39 If the rate constant, k, for the forward reaction above is 1.4 x 10 -28 L . mol-1. sec-1 what is k (in L. mol-1. sec-1 ) for the backward reaction? (A) 2.1 × 10-67 (B) 1.0 × 10-11 (C) 9.3 × 1010 (D) 7.1 × 1027 33. Calculate the [H +] in a 0.25 M solution of methylamine, CH3NH2 (Kb = 4.4 × 10-4). (A) 1.1 × 10-4 (B) 1.0 × 10-2 (C) 9.1 × 10-11 (D) 9.5 × 10-13 34. A 0.010 M solution of a weak acid, HA, is 0.40% ionized. What is its ionization constant? (A) 1.6 × 10-10 (B) 1.6 × 10-7 (C) 4.0 × 10-5 (D) 4.0 × 10-3 Not valid for use as an USNCO National Examination after April 27, 2003 Page 5 35. 1.0 L of an aqueous solution in which [H 2CO3] = [HCO3-] = 0.10 M, has [H+] = 4.2 × 10-7. What is the [H+] after 0.005 mole of NaOH has been added? (A) 2.1 ×10-9 M (B) 2.2 × 10-8 M (A) 0.629 V (B) 0.689 V (C) 3.8 × 10 M (D) 4.6 × 10 M (C) 0.748 V (D) 0.866 V -7 -7 36. A solution of Pb(NO3)2 is added dropwise to a second solution in which [Cl-] = [F-] = [I -] = [SO42-] = 0.001 M. What is the first precipitate that forms? 37. 41. The voltage for the cell Fe ❘ Fe2+(0.0010 M) ❘ ❘ Cu2+(0.10 M) ❘ Cu is 0.807 V at 25 ˚C. What is the value of E˚? (A) PbCl2 (Ksp = 1.5 × 10-5) (B) PbF2 (K sp = 3.7 × 10-8) (C) PbI2 (K sp = 8.5 × 10-9) (D) PbSO4 (K sp = 1.8 × 10-8) (A) 0.39 M (B) 3 (C) 4 (D) 6 38. Use the standard reduction potentials; Sn 2+(aq) + 2e– r Sn(s) E˚ = -0.141 V Ag+(aq) + e– r Ag(s) E˚ = 0.800 V to calculate E˚ for the reaction; Sn(s) + 2Ag +(aq) r Sn2+(aq) + 2Ag(s) (B) 0.941 V (C) 1.459 V (D) 1.741 V 39. Which of the processes happen during the discharging of a lead storage battery? (B) microwave (C) ultraviolet (D) visible 44. All of the following possess complete d shells EXCEPT (B) 2 only (C) 1 and 3 only (D) 2 and 3 only (B) -585 kJ (C) -390 kJ (D) -195 kJ Page 6 (B) Cu 2+ (C) Ga 3+ (D) Zn2+ (A) 6s (B) 5p (C) 5d (D) 4d 46. Which set of quantum numbers (n, l, ml , ms) is permissible for an electron in an atom? (A) 1 only (A) -1170 kJ (D) 0.89 M 45. Which orbital fills completely immediately before the 4f? 1. H2(g) is produced 2. PbO2 is converted to PbSO4 3. The density of the electrolyte solution decreases 40. What is the value of ∆G˚ for the reaction? 2Al(s) + 3Cu 2+(aq) r 2Al3+(aq) + 3Cu(s) (C) 0.78 M (A) infrared (A) Ag+ (A) 0.659 V (B) 0.46 M 43. Which region of the electromagnetic spectrum is capable of inducing electron transitions with the greatest energy? Cl2 + OH- r Cl- + ClO3What is the coefficient for OH- when this equation is balanced with the smallest integer coefficients? (A) 2 42. A current of 2.0 A is used to plate Ni(s) from 500 mL of a 1.0 M Ni2+(aq) solution. What is the [Ni2+] after 3.0 hours? (A) 1, 0, 0, -1/2 (B) 1, 1, 0, +1/2 (C) 2, 1, 2, +1/2 (D) 3, 2, -2, 0 47. When the elements Li, Be, and B, are arranged in order of increasing ionization energy, which is the correct order? (A) Li, B, Be (B) B, Be, Li (C) Be, Li, B (D) Li, Be, B 48. Which forms the most alkaline solution when added to water? (A) Al2O3 (B) B 2O3 (C) CO2 (D) SiO 2 E˚ = 2.02 V 49. What is the total number of valence electrons in the peroxydisulfate, S2O82-, ion? (A) 58 (B) 60 (C) 62 (D) 64 Not valid for use as an USNCO National Examination after April 27, 2003 50. For which species are both bonds of equal length? 1. ClO22. NO2- (A) 1 only (B) 2 only (C) both 1 and 2 (D) neither 1 nor 2 51. Which compound has the highest melting point? (A) MgO (B) KCl (C) NaCl (D) CaO 52. Which molecular geometry is least likely to result from a trigonal bipyramidal electron geometry? 57. Which substance will react most rapidly with Br 2(aq)? (A) benzene (B) chloropropane (C) propanone (D) propene 58. Which compound includes a carbon atom with an sp hybridized orbital? (A) benzene (B) butyne (C) methyl chloride (D) phenol 59. Which compound has the highest vapor pressure at 25 ˚C? (A) trigonal planar (B) see-saw (A) CH3CH2CH2CH2OH (B) CH3CH2CH2OCH3 (C) linear (D) t-shaped (C) CH3CH2CH2CH2NH2 (D) (CH3) 3COH 53. Which diatomic species has the greatest bond strength? (A) NO (B) NO+ (D) O2- (C) O2 60. Which of the molecules can exist as optical isomers? 54. During the complete combustion of methane, CH4, what change in hybridization does the carbon atom undergo? (A) sp3 to sp (B) sp3 to sp 2 (C) sp2 to sp (D) sp2 to sp 3 O (A) C Br H O OH C H (C) 55. What is the formula for the compound? C Br OH C H O (B) CH3 O (D) CH3 CH3 (A) C 8H10 (B) C 8H12 (C) C 8H14 (D) C 8H16 H OCH3 C OH C Br C Br C Br H H 56. Which is most likely to react by an SN1 mechanism? (A) CH3Cl (B) CH3CHClCH3 (C) (CH3) 3CCl (D) C 6H5Cl Not valid for use as an USNCO National Examination after April 27, 2003 END OF TEST Page 7 CHEMISTRY OLYMPIAD 2003 NATIONAL EXAM PART 1— KEY Number 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. Answer D B B C D A B D C B D B C A A D D B A C B A B A D B D C D D Not valid for use as an USNCO National Examination after April 27, 2003 Number 31. 32. 33. 34. 35. 36. 37. 38. 39. 40. 41. 42. 43. 44. 45. 46. 47. 48. 49. 50. 51. 52. 53. 54. 55. 56. 57. 58. 59. 60. Answer C C D B C D D B D A C C C B A A A A C C A A B A A C D B B B 2003 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART II Prepared by the American Chemical Society Olympiad Examinations Task Force OLYMPIAD EXAMINATIONS TASK FORCE Arden P. Zipp, State University of New York, Cortland Chair Peter E. Demmin (retired), Amherst Central High School, NY David W. Hostage, Taft School, CT Alice Johnsen, Bellaire High School, TX Jerry D. Mullins, Plano Senior High School, TX Ronald O. Ragsdale, University of Utah, UT Amy Rogers, College of Charleston, SC DIRECTIONS TO THE EXAMINER–PART II Part II of this test requires that student answers be written in a response booklet of blank pages. Only this “Blue Book” is graded for a score on Part II. Testing materials, scratch paper, and the “Blue Book” should be made available to the student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until April 27, 2003, after which tests can be returned to students and their teachers for further study. Allow time for the student to read the directions, ask questions, and fill in the requested information on the “Blue Book”. When the student has completed Part II, or after one hour and forty-five minutes has elapsed, the student must turn in the “Blue Book”, Part II of the testing materials, and all scratch paper. Be sure that the student has supplied all of the information requested on the front of the “Blue Book,” and that the same identification number used for Part I has been used again for Part II. There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and you are free to schedule rest-breaks between parts. Part I Part II Part III 60 questions 8 questions 2 lab problems single-answer multiple-choice problem-solving, explanations laboratory practical 1 hour, 30 minutes 1 hour, 45 minutes 1 hour, 30 minutes A periodic table and other useful information are provided on the back page for student reference. Students should be permitted to use non-programmable calculators. DIRECTIONS TO THE EXAMINEE–PART II DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Part II requires complete responses to questions involving problem-solving and explanations. One hour and forty-five minutes are allowed to complete this part. Be sure to print your name, the name of your school, and your identification number in the spaces provided on the “Blue Book” cover. (Be sure to use the same identification number that was coded onto your Scantron® sheet for Part I.) Answer all of the questions in order, and use both sides of the paper. Do not remove the staple. Use separate sheets for scratch paper and do not attach your scratch paper to this examination. When you complete Part II (or at the end of one hour and forty-five minutes), you must turn in all testing materials, scratch paper, and your “Blue Book.” Do not forget to turn in your U.S. citizenship statement before leaving the testing site today. Not valid for use as an USNCO National Exam after April 27, 2003. Distributed by the ACS DivCHED Examinations Institute, University of Wisconsin-Milwaukee, Milwaukee, WI. All rights reserved. Printed in U.S.A. 1. (12%) 0.1152 g of a compound containing carbon, hydrogen, nitrogen and oxygen are burned in excess oxygen. The gases produced are treated further to convert nitrogen-containing products into N2. The resulting mixture of CO2, H2O and N2 and excess O2 is passed through a CaCl2 drying tube, which gains 0.09912 g. The gas stream is bubbled through water where the CO 2 forms H2CO3. Titration of this solution to the second endpoint with 0.3283 M NaOH requires 28.81 mL. The excess O2 is removed by reaction with copper metal and the N2 is collected in a 225.0 mL measuring bulb where it exerts a pressure of 65.12 mmHg at 25 ˚C. In a separate experiment the molar mass of this compound is found to be approximately 150 g. mol -1. a. Calculate the number of moles of i. H2O ii. CO2 iii. N2 b. Determine the mass in the original compound of i. C ii. H iii. N iv. O c. Find the empirical formula of the compound. d. Find the molecular formula. 2. (13%) The enthalpy of combustion of liquid octane, C8H18(l) to gaseous products, is -5090 kJ.mol -1. Use this value to answer the questions below, assuming a temperature of 100 ˚C. a. Write a balanced equation for the complete combustion of liquid octane C 8H18(l). b. Determine the molar enthalpy of formation, ∆Hf˚, for liquid octane, C8H18(l). [∆H f˚ kJ. mol -1; CO2(g) -393.5, H2O(g) -241.8] c. Calculate the value of the internal energy change, ∆E˚, for the combustion reaction. d. If ∆G˚ for the combustion is -5230 kJ. mol -1 of octane, calculate the value of ∆S˚. Comment on the sign of ∆S˚ relative to the equation written above. e. State whether the heat associated with the combustion of liquid octane in a bomb calorimeter represents ∆H˚ or ∆E˚. Explain your reasoning. 3. (13%) Phosphoric acid, H3PO4, ionizes according to the equations, s H+(aq) + H2PO4–(aq) K1 = 7.1 x 10-3 H2PO4–(aq) s H+(aq) + HPO42-(aq) K2 = 6.2 x 10-8 HPO42-(aq) s H+(aq) + PO43-(aq) K3 = 4.5 x 10-13 H3PO4(aq) a. Write the equilibrium expression for the ionization of H3PO4 and find the pH of a 1.5 M solution of H3PO4. b. A student is asked to prepare a phosphate buffer with a pH of 7.00. Identify the species that should be used in this solution and calculate their ratio. c. Assume that 50.0 mL of the buffer solution in b. are available in which the more abundant buffer species has a concentration of 0.10 M. Determine the [H +] in this solution after 2.0 x 10-3 mol of NaOH are added. d. Determine the [H+] in a 0.20 M solution of Na3PO4. 4. (13%) An electrochemical cell is constructed with a piece of copper wire in a 1.00 M solution of Cu(NO3)2 and a piece of chromium wire in a 1.00 M solution of Cr(NO3) 3. The standard reduction potentials for Cr3+(aq) and Cu2+(aq) are: Cr3+(aq) + 3e– ---> Cr(s) Cu 2+(aq) + 2e– ---> Cu(s) –0.744 V 0.340 V a. Write a balanced equation for the spontaneous reaction that occurs in this cell and calculate the potential it produces. b. Sketch a diagram for this cell. i. Label the anode. ii. Show the direction of electron flow in the external circuit. iii. Show the direction of movement of nitrate ions. Explain. c. The cell is allowed to operate until the [Cu 2+] = 0.10 M. i. Find the [Cr3+]. ii. Calculate the cell potential at these concentrations. Not valid for use as an USNCO National Examination after April 27, 2003. Page 2 5. (12%) Write net equations for each of the combinations of reactants below. Use appropriate ionic and molecular formulas and omit formulas for all ions or molecules that do not take part in a reaction. Write structural formulas for all organic substances. You need not balance the equations. All reactions occur in aqueous solution unless otherwise indicated. a. Water is added to magnesium nitride. b. Excess carbon dioxide is bubbled through a solution of calcium hydroxide. c. Acidic solutions of potassium dichromate and iron(II) chloride are mixed. d. Solutions of lead acetate and sulfuric acid are mixed. e. Excess concentrated sodium hydroxide is added to a solution of zinc nitrate. f. Fluorine-18 undergoes positron emission. 6. (14%) Account for the following observations about chemical kinetics. a. Reactions involving molecular chlorine often have nonintegral rate laws. b. The rates of exothermic reactions increase when their temperatures are increased. c. Two reactions, A and B, have rate constants that are equal at 25˚C but the rate constant for reaction A is much greater than that for reaction B at 35˚C. d. The rates of reactions catalyzed by complex molecules, such as enzymes, increase with an increase in temperature up to a certain point above which they decrease again. 7. (12%) The atoms C, N and O can be arranged in three different orders to form negative ions, i.e. CNO–, CON–, and NCO–. Salts of one of these ions are stable. Salts of one are explosive and salts of one are unknown at this time. a. Write Lewis electron dot structures for each atomic arrangement. b. For each arrangement, i. find the formal charge on each atom. ii. use these formal charges to identify the most and least stable arrangements. Explain your reasoning. c. Predict the geometry of the most stable atom arrangement and identify the type of hybridization used by the central atom in this structure. 8. (11%) Three compounds; X, Y, Z, have the formula C3H8O. a. Write structural formulas for three different compounds with the formula C3H8O. b. The boiling points of the compounds are; X 10.8˚C, Y 82.4˚C, Z 97.4˚C. Assign each boiling point to one of the structures in part a. and account for these boiling points on the basis of the molecular structures and the types of intermolecular forces in each. c. A fourth compound with the formula C2H4O2, has the same molar mass as the three compounds above and boils at 117.9˚C. Propose a structure for this compound and account for its higher boiling point. END OF PART II Not valid for use as an USNCO National Examination after April 27, 2003. Page 3 amount of substance ampere atmosphere atomic mass unit atomic molar mass Avogadro constant Celsius temperature centi- prefix coulomb electromotive force energy of activation enthalpy entropy ABBREVIATIONS AND SYMBOLS n equilibrium constant K measure of pressure mmHg A Faraday constant F milli- prefix m atm formula molar mass M molal m u free energy G molar M A frequency ν mole mol N A gas constant R Planck’s constant h °C gram g pressure P c heat capacity C p rate constant k C hour h retention factor Rf E joule J second s Ea kelvin K speed of light c H kilo- prefix k temperature, K T S liter L time t volt V CONSTANTS R = 8.314 J·mol –1·K–1 R = 0.0821 L·atm·mol –1·K–1 1 F = 96,500 C·mol–1 1 F = 96,500 J·V–1·mol–1 N A = 6.022 × 1023 mol–1 h = 6.626 × 10–34 J·s c = 2.998 × 108 m·s–1 USEFUL EQUATIONS E = Eο – k2 Ea 1 1 = − k1 R T1 T2 – ∆H 1 ln K = +c R T RT ln Q nF ln PERIODIC TABLE OF THE ELEMENTS 1 H 2 He 1.008 4.003 3 Li 4 Be 5 B 6 C 7 N 8 O 9 F 10 Ne 6.941 9.012 10.81 12.01 14.01 16.00 19.00 20.18 11 Na 12 Mg 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar 22.99 24.31 26.98 28.09 30.97 32.07 35.45 39.95 19 K 20 Ca 21 Sc 22 Ti 23 V 24 Cr 25 Mn 26 Fe 27 Co 28 Ni 29 Cu 30 Zn 31 Ga 32 Ge 33 As 34 Se 35 Br 36 Kr 39.10 40.08 44.96 47.88 50.94 52.00 54.94 55.85 58.93 58.69 63.55 65.39 69.72 72.61 74.92 78.96 79.90 83.80 37 Rb 38 Sr 39 Y 40 Zr 41 Nb 42 Mo 43 Tc 44 Ru 45 Rh 46 Pd 47 Ag 48 Cd 49 In 50 Sn 51 Sb 52 Te 53 I 54 Xe 85.47 87.62 88.91 91.22 92.91 95.94 (98) 101.1 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.3 55 Cs 56 Ba 57 La 72 Hf 73 Ta 74 W 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 132.9 137.3 138.9 178.5 181.0 183.8 186.2 190.2 192.2 195.1 197.0 200.6 204.4 207.2 209.0 (209) (210) (222) 87 Fr 88 Ra 89 Ac 104 Rf 105 Db 106 Sg 107 Bh 108 Hs 109 Mt 110 111 112 114 (223) 226.0 227.0 (261) (262) (263) (262) (265) (266) (269) (272) (277) (277) 58 Ce 59 Pr 60 Nd 61 Pm 62 Sm 63 Eu 64 Gd 65 Tb 66 Dy 67 Ho 68 Er 69 Tm 70 Yb 71 Lu 140.1 140.9 144.2 (145) 150.4 152.0 157.3 158.9 162.5 164.9 167.3 168.9 173.0 175.0 90 Th 91 Pa 92 U 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr 232.0 231.0 238.0 237.0 (244) (243) (247) (247) (251) (252) (257) (258) (259) (260) Not valid for use as an USNCO National Examination after April 27, 2003. Page 4 2003 U.S. NATIONAL CHEMISTRY OLYMPIAD KEY FOR NATIONAL EXAM – PART II 1. Facts from the problem. • 0.1152 g sample contains C, H, N, O. • 0.0912 g H2 O are recovered. • Conversion of carbon dioxide to carbonic acid has 1 to 1 stoichiometry, • 1 CO2 ! 1 H2 CO3 . • Carbonic acid is titrated with 28.81 mL of 0.3283 M NaOH • H2 CO3 + 2 NaOH ! Na2 CO3 + 2 H2O • Volume of N2 collected is 225.0 mL at 65.12 torr and 25 o C. • Molar mass of the substance is 150 g/mol a. i. H2 O (1 pt) moles H2O = 0.09912 g × ii. 1 mol H2O = 5.501 × 10−3 mol H2O 18.02 g H2O CO2 (2 pts) moles CO2 = 0.3283 mol NaOH 1 mol CO2 × 0.02881 L soln × = 4.729 × 10−3 mol CO2 L soln 2 mol NaOH iii. N2 (1 pt) moles N2 = b. i. pV (65.12 torr)(0.2250 L ) = = 7.879 × 10−4 mol N2 RT (62.4 L ⋅ torr ⋅ mol-1 ⋅ K -1 )(298 K) C (1 pt) g C = 4.729 × 10 -3 mol CO 2 × ii. 1 mol C 12.01 g C × = 0.05680 g C 1 mol CO 2 1 mol C H (1 pt) g H = 5.501 × 10 -3 mol H 2O × 2 mol H 1.008 g H × 1 mol H 2O 1 mol H = 0.01109 g H iii. N (1 pt) g N = 7.879 × 10 -4 mol N 2 × 2 mol N 14.01 g N × 1 mol N 2 1 mol N iv. O (1 pt) g O = g sample – (g C + g H + g N) = 0.1152 g – (0.05680 g + 0.01109 g + 0.02208 g) = 0.1152 g – 0.08997 g = 0.02523 g O c. Find the empirical formula of the compound. (2 pts) Key for 2003 USNCO National Exam, Part II = 0.02208 g N C : 0.05680 g C × 1 mol C 12.01 g C = 4.729 × 10−3 / 1.576 × 10−3 = 3.001 ≈ 3 H : 0.01109 g H × 1 mol H 1.008 g H = 1.100 × 10−2 / 1.576 × 10−3 = 6.980 ≈ 7 N : 0.02208 g N × 1 mol N 14.01 g N = 1.576 × 10−3 / 1.576 × 10−3 = 1.000 = 1 1 mol O = 1.577 × 10−3 / 1.576 × 10−3 = 1.001 ≈ 1 16.00 g O Therefore the empirical formula is C3H7 NO O : 0.02523 g N × d. Find the molecular formula. (2 pts) The molar mass of the empirical formula is 73.10 Compare this molar mass to the measured molar mass. 150 g ⋅ mol = 2.05 ≈ 2 73.1 g ⋅ mol Therefore the molecular formula is C6 H14N2 O2 2. a. Write balanced equations to represent the processes responsible for K1 and K2. 2 C8 H18 (l) + 25 O2 (g) ! 16 CO2 (g) + 18 H2O (g) (1 pt) b. Determine the molar enthalpy of formation, ∆Hf˚, for liquid octane, C8H18(l). (4 pts) ∆H rxn = 8∆H f (CO 2 ) + 9∆H f (H 2O) − ∆H f (C 8H18 ) −5090 kJ = 8(−393.5) kJ + 9(−241.8) kJ − ∆H f (C 8H18 ) ∆H f (C 8H18 ) = −3148 − 2176.2 + 5090 kJ = -234.2 kJ c. Calculate the value of the internal energy change, ∆E˚, for the combustion reaction. (4 pts) ∆H = ∆E + ∆nRT so ∆E = ∆H − ∆nRT = –5090000 J – (4.5 mol)(8.314 J. mol-1 . K-1 )(373 K) = –5090000 J – 13955 J = –5104000 J = -5104 kJ /mol C8 H18(l) -1 d. If ∆G˚ for the combustion is -5230 kJ. mol of octane, calculate the value of ∆S˚. Comment on the sign of ∆S˚ relative to the equation written above . (3 pts) ∆G° = ∆H ° − T∆-1S ° -5230 kJ . mol = -5090 kJ . mol-1 – 373 K (∆So ) so −5090 kJ ⋅ mol-1 + 5230 kJ ⋅ mol-1 ∆S° = 373 K = 0.375 kJ.mol-1.K-1 The increase in ∆S˚ is consistent with the formation of more moles of gas during the reaction. Key for 2003 USNCO National Exam, Part II e. State whether the heat associated with the combustion of liquid octane in a bomb calorimeter represents ∆H˚ or ∆E˚. Explain your reasoning. (1 pt) Heat in a bomb calorimeter is ∆E˚ (q at constant volume) - no credit unless there is a discussion of zero work under constant volume. 3. a. Write the equilibrium expression for the ionization of H3PO4 and find the pH of a 1.5 M solution of H3PO4. (1 pt for equation, 1 pt for solution) [H ][H PO ] = + Ka – 2 4 [H 3PO 4 ] for a 1.5 M solution, assuming no initial concentration of reactants and that the amount of phosphoric acid that reacts is small compared to the original volume, x2 1.5 Solving for x, x = 1.5( 7.1 × 10−3 ) = 0.103. Use successive approximations to check that the amount of reacting phosphoric acid doesn’t change the answer… plug in 1.397 (1.5-0.103=1.397) rather than 1.5, 7.1 × 10−3 = x = 1.397( 7.1 × 10−3 ) = 0.100 the change is small enough to accept this answer. b. A student is asked to prepare a phosphate buffer with a pH of 7.00. Identify the species that should be used in this solution and calculate their ratio. (1 pt for correct identification of salts, 1 pt for solution of ratio) it would be Ka2. Thus the species that would To obtain a pH of 7 the Ka should be close– to 7. In this case 2be present in this buffer should be H2 PO4 (aq) and HPO4 (aq). The ratio can be found by the equation, 2– HPO 2– HPO 4 4 0.62 −8 −7 6.2 × 10 = 1.0 × 10 and the ratio is, = 0.62 or – – H 2PO 4 1 H 2PO 4 [ [ c. ] ] [ [ ] ] Assume that 50.0 mL of the buffer solution in b. are available in which the more abundant buffer species has a concentration + -3 of 0.10 M. Determine the [H ] in this solution after 2.0 x 10 mol of NaOH are added. (1 pt for correct calculation of each concentration) – H2 PO4 (aq) is the more abundant species in the buffer from part b. – – 0.050 L × 0.10 mol ⋅ L-1 H 2PO 4 = 0.0050 mol H 2PO– 4 initially – 0.0020 mol of OH is added, so the amount of H2 PO4 (aq) left is (1 to 1 stoichiometry) – 0.0050 – 0.0020 = 0.0030 mol– H2 PO4 (aq) If the concentration of 2-H2 PO4 (aq) is 0.10 M the ratio calculated earlier indicates that the initial concentration of HPO4 (aq) must be 0.62 × 1 = 0.062 M, so 2– 2– 0.050 L × 0.062 mol ⋅ L-1 HPO 4 = 0.0031 mol HPO 4 initially 2and 0.0020 mol are formed in the reaction so there is 0.0051 mol HPO4 (aq) present. Calculating the amount of hydrogen ion present, 0.0051 6.2 × 10−8 = [H + ] , solving for [H+] gives, 3.65 × 10-8 M 0.0030 Key for 2003 USNCO National Exam, Part II d. + Determine the [H ] in a 0.20 M solution of Na3PO4. (1 pt for determining Kb, 1 pt for noting the need for successive approximation in the calculation of phosphate ion, 1 pt for final solution) PO43– + H2O ! HPO4 2– + OH– and K 1.0 × 10−14 Kb = w = = 0.0222 K a 3 4.5 × 10−13 Obtain concentration of hydroxide, initially assume no reaction of phosphate and let x = [OH–], x2 0.0222 = so x = [OH – ] = 6.67 × 10−2 now for phosphate, 0.20 – 0.0667 = 0.1333 0.20 x2 0.0222 = so x = [OH – ] = 5.44 × 10−2 now for phosphate, 0.20-0.0544 = 0.1456 0.1333 x2 0.0222 = so x = [OH – ] = 5.70 × 10−2 now for phosphate, 0.20-0.0570 = 0.143 0.1456 x2 0.0222 = so x = [OH – ] = 5.63 × 10−2 which is acceptably close to the previous iteration, now use 0.143 Kw to calculate [H+], 1.0 × 10−14 + H = [ ] 5.63 × 10−2 = 1.77 × 10−13 4. a. Write a balanced equation for the spontaneous reaction that occurs in this cell and calculate the potential it produces. (3 pts) 2Cr(s) + 3 Cu2+(aq) r 2 Cr3+(aq) + 3 Cu(s) o E = Eox + Ered = 0.744 + 0.340 = 1.084 V b. Sketch a diagram for this cell. (5 pts for sketch with proper labels. Points taken of for incorrectly labeled components, etc.) i. Label the anode. ii. Show the direction of electron flow in the external circuit. iii. Show the direction of movement of nitrate ions. Explain. volts electrons _ Cr node Cu NO3 r + Key for 2003 USNCO National Exam, Part II u + 2+ c. The cell is allowed to operate until the [Cu ] = 0.10 M. 3+ i. Find the [Cr ]. (2 pts) ii. Calculate the cell potential at these concentrations. (3 pts) [Cu2+] goes from 1.0 M to 0.10 M, so ∆[Cu2+] is –0.90 ∆[Cr3+] = 0.90 × 2/3 = 0.60 so, [Cr3+] = 1.60 plug these values into the equation, 3+ 2 RT [Cr ] o E=E − nF [Cu 2+ ] 3 0.0257 (1.60) 2 E = 1.084 − = 1.05 V 6 (0.10) 3 5. (Note that balanced chemical equations are note required.) (1 pt for each reactant and 2 points for products (usually 1 for each product) then the total point value was multiplied by 2/3 to scale to 12 pts) a. b. c. d. e. f. Mg3 N2 + H2O ! Mg(OH)2 + NH3 CO2 + OH– ! HCO3– Cr2 O7 2– + Fe2+ + H+ ! Cr3+ + Fe3+ + H2O Pb2+ + C2 H3 O2 – + H+ + SO4 2– ! PbSO4 + HC2H3 O2 Zn2+ + OH– ! Zn(OH)4 2– 18 →188 O + +10β 9F 6. a. (3 pts) Cl2 often dissociates to Cl atoms, which react individually. If these atoms participate in the rate determining step the overall rate equation is proportional to [Cl•] = [Cl2]1/2 b. (3 pts) All reactions increase in rate with an increase in temperature due to an increase in the collision frequency and the increase in fraction of species with high velocities. Higher velocity particles impart more energy to collisions in which they participate so those collisions are more likely to exceed Ea. The exothermicity or endothermicity of a reaction has no bearing on its kinetics or the effect of temperature. c. (4 pts) Reactions A and B must have different activation energies. When the log (or ln) of the rate is constant for a reaction is plotted versus 1/T the slope is the Ea. Reactions A and B have different slopes which cross at 25 o C. The Ea for reaction A is greater because an increase in temperature affects it’s rate constant more. d. (4 pts) Reaction rates increase with higher temperature because of an increase in collision rate and an increase in the fraction of molecules that have the necessary energy to react. At still higher temperatures the enzyme is denatured so that it no longer is an effective catalyst. 7. (1 pt for each correct Lewis structure (3 total) 1 pt for correct formal charges on each Lewis structure (3 total) 1 pt for correct identification of most stable structure / 1 pt for correct reasoning 1 pt for correct identification of least stable structure / 1 pt for correct reasoning 1 pt for correctly identifying the structure as linear 1 pt for correct hybridization.) Key for 2003 USNCO National Exam, Part II a. – – C N C O O N b. C O – – i. – N C N O C -1 +1 -1 -1 O +2 – N N C O -2 0 0 -1 ii. N C O is the most stable arrangement because the formal charges are the lowest in this structure. C O N is the least stable structure because the formal charges are the greatest in this structure c. N C O as drawn would be linear because there are two charge centers around the central atom. The central atom will be sp hybridized. 8. a. The structural formulas would be, (1 pt for each) O o X (b.p. 10.8 C) CH3 CH2 CH3 CH CH3 OH CH3 Y (b.p. 82.4 o C) Z (b.p. 97.4 oC) CH3 CH2 CH2 OH Compound Y and Z would have hydrogen bonding where compound X would not. Therefore X would have the weakest intermolecular forces and the lowest boiling point (of 10.8 oC) (2 pts) b. The linear shape of compound Z would allow for stronger dispersion forces than compound Y. Therefore Y should have a lower boiling point than Z. (4 pts) c. The possible Lewis structures include, O O CH3 C C OH H O OH CH2 CH3 C C O OH OH H H C H All of these would have increased hydrogen-bonding relative to the three compounds in part a. Key for 2003 USNCO National Exam, Part II 2003 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART III Prepared by the American Chemical Society Olympiad Laboratory Practical Task Force OLYMPIAD LABORATORY PRACTICAL TASK FORCE Steve Lantos, Brookline High School, Brookline, MA Chair Nancy Devino, ScienceMedia Inc., San Diego, CA Lucy Pryde Eubanks, Clemson University, Clemson, SC Sheldon L. Knoespel, Michigan State University, East Lansing, MI Jim Schmitt, Eau Claire North High School, Eau Claire, WI Christie B. Summerlin, University of Alabama-Birmingham, Birmingham, AL Linda Weber, Natick High School, Natick, MA DIRECTIONS TO THE EXAMINER–PART III The laboratory practical part of the National Olympiad Examination is designed to test skills related to the laboratory. Because the format of this part of the test is quite different from the first two parts, there is a separate, detailed set of instructions for the examiner. This gives explicit directions for setting up and administering the laboratory practical. There are two laboratory tasks to be completed during the 90 minutes allotted to this part of the test. Students do not need to stop between tasks, but are responsible for using the time in the best way possible. Each procedure must be approved for safety by the examiner before the student begins that procedure. Part III 2 lab problems laboratory practical 1 hour, 30 minutes Students should be permitted to use non-programmable calculators. DIRECTIONS TO THE EXAMINEE–PART III DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. WHEN DIRECTED, TURN TO PAGE 2 AND READ THE INTRODUCTION AND SAFETY CONSIDERATIONS CAREFULLY BEFORE YOU PROCEED. There are two laboratory-related tasks for you to complete during the next 90 minutes. There is no need to stop between tasks or to do them in the given order. Simply proceed at your own pace from one to the other, using your time productively. You are required to have a procedure for each problem approved for safety by an examiner before you carry out any experimentation on that problem. You are permitted to use a non-programmable calculator. At the end of the 90 minutes, all answer sheets should be turned in. Be sure that you have filled in all the required information at the top of each answer sheet. Carefully follow all directions from your examiner for safety procedures and the proper disposal of chemicals at your examining site. Not valid for use as an USNCO National Examination after April 27, 2003. Page 1 2003 UNITED STATES NATIONAL CHEMISTRY OLYMPIAD PART III — LABORATORY PRACTICAL Student Instructions Introduction These problems test your ability to design and carry out laboratory experiments and to draw conclusions from your experimental work. You will be graded on your experimental design, on your skills in data collection, and on the accuracy and precision of your results. Clarity of thinking and communication are also components of successful solutions to these problems, so make your written responses as clear and concise as possible. Safety Considerations You are required to wear approved eye protection at all times during this laboratory practical. You also must follow all directions given by your examiner for dealing with spills and with disposal of wastes. Lab Problem 1 Sani-Flush®, a commercial toilet bowl cleaner, contains sodium bisulfate (sodium hydrogen sulfate) and sodium carbonate as its active ingredients. Other ingredients include sodium chloride, sodium lauryl sulfate, talc and some fragrance. Given the information from the manufacturer that the sodium bisulfate is substantially in excess compared to the sodium carbonate, carry out an experiment to determine the percent by weight of the sodium carbonate in a sample of the product. Lab Problem 2 You have been provided with eight vials, each of which is labeled with a number from 1 to 8. Each vial contains one of the following chemicals: Na3PO4, NH4Cl, ZnCl2, KNO3, Mg(OH)2, Pb(NO3)2, CaCO3, Na2SO3. You are allowed to use distilled water, test tubes or well plates, and only TWO additional reagents from the following choices: 6 M H2SO4, 6 M HCl, 6 M AgNO3, phenolphthalein indicator solution You must designate your choice of reagents prior to the start of your testing. Page 2 Not valid for use as an USNCO National Examination after April 27, 2003. Answer Sheet for Laboratory Practical Problem 1 Student's Name: __________________________________________________________________________ Student's School:________________________________________ Date: ___________________________ Proctor's Name: _________________________________________________________________________ ACS Section Name :________________________________Student's USNCO test #: ________________ 1. Give a brief description of your experimental plan. Before beginning your experiment, you must get approval (for safety reasons) from the examiner. Not valid for use as an USNCO National Examination after April 27, 2003. Examiner’s Initials: Page 3 2. Record your data and other observations. 3. Show your calculations. Page 4 Not valid for use as an USNCO National Examination after April 27, 2003. Answer Sheet for Laboratory Practical Problem 2 Student's Name: __________________________________________________________________________ Student's School:________________________________________ Date: ___________________________ Proctor's Name: _________________________________________________________________________ ACS Section Name : ________________________________Student's USNCO test #: ________________ 1. Give a brief description of your experimental plan. Before beginning your experiment, you must get approval (for safety reasons) from the examiner. Examiner’s Initials: When you wish to request the optional reagents, return to the Examiner with this sheet. I request these additional reagents: Not valid for use as an USNCO National Examination after April 27, 2003. Examiner’s Initials: Page 5 2. Record your data and other observations. 3. Identify the substance in each numbered vial. Vial # Contains 1. 2. 3. 4. 5. 6. 7. 8. Page 6 Not valid for use as an USNCO National Examination after April 27, 2003. 4. Explain clearly how you arrived at the identity of each of the vial’s contents. Not valid for use as an USNCO National Examination after April 27, 2003. Page 7 2003 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART III Prepared by the American Chemical Society Olympiad Laboratory Practical Task Force Examiner's Instructions Directions to the Examiner: Thank you for administering the 2003 USNCO laboratory practical on behalf of your Local Section. It is essential that you follow the instructions provided, in order to insure consistency of results nationwide. There may be considerable temptation to assist the students after they begin the lab exercise. It is extremely important that you do not lend any assistance or hints whatsoever to the students once they begin work. As in the international competition, the students are not allowed to speak to anyone until the activity is complete. The equipment needed for each student for both lab exercises should be available at his/her lab station or table when the students enter the room. The equipment should be initially placed so that the materials used for Lab Problem 1 are separate from those used for Lab Problem 2. After the students have settled, read the following instructions (in italics) to the students. Hello, my name is ________. Welcome to the lab practical portion of the U.S. National Chemistry Olympiad Examination. In this part of the exam, we will be assessing your lab skills and your ability to reason through a laboratory problem and communicate its results. Do not touch any of the equipment in front of you until you are instructed to do so. One of this year’s problems requires the use of a plastic syringe with a Luer-lock® tip cap. Show a syringe and Luer-lock® tip cap. This problem also requires you to use a balance, which is located _____________________________. Another of this year’s problems uses small-scale chemistry equipment. Small-scale chemistry techniques help to minimize the amount of materials you use, thereby increasing safety and minimizing waste. Specialized equipment for small-scale chemistry that you will use today include Beral-type pipets and reaction plates. Show a 5-mL Beral-type pipet, and show a 24-well reaction plate or small test tubes. You will be asked to complete two laboratory problems. All the materials and equipment you may want to use to solve each problem has been set out for you and is grouped by the number of the problem. You must limit yourself to this equipment for each problem. You will have one hour and thirty minutes to complete the two problems. You may choose to start with either problem. You are required to have a procedure for each problem approved for safety by an examiner. (Remember that approval does not mean that your procedure will be successful–it is a safety approval.) When you are ready for an examiner to come to your station for each safety approval, please raise your hand. Page 1 Safety is an important consideration during the lab practical. You must wear goggles at all times. Wash off any chemicals spilled on your skin or clothing with large amounts of tap water. The appropriate procedures for disposing of solutions at the end of this lab practical are: ____________________________________________________________________________________ ____________________________________________________________________________________ We are about to begin the lab practical. Please do not turn the page until directed to do so, but read the directions on the front page. Are there any questions before we begin? Distribute Part III booklets and again remind students not to turn the page until the instruction is given. Part III contains student instructions and answer sheets for both laboratory problems. There is a periodic table on the last page of the booklet. Allow students enough time to read the brief cover directions. Do not turn to page 2 until directed to do so. When you start to work, be sure that you fill out all information at the top of the answer sheets. Are there any additional questions? If there are no further questions, the students should be ready to start Part III. You may begin. After one hour and thirty minutes, give the following directions. This is the end of the lab practical. Please stop and bring me your answer sheets. Thank you for your cooperation during this test. Collect all the lab materials. Make sure that the student has filled in his or her name and other required information on the answer sheets. At this point, you may want to take five or ten minutes to discuss the lab practical with the students. They can learn about possible observations and interpretations and you can acquire feedback as to what they actually did and how they reacted to the problems. After this discussion, please take a few minutes to complete the Post-Exam Questionnaire; this information will be extremely useful to the Olympiad subcommittee as they prepare next year’s exam. Please remember to return the post-exam Questionnaire, the answer sheets from Part III, the Scantron sheets from Part I, and the “Blue Books” from Part II to this address: ACS DivCHED Exams Institute Department of Chemistry University of Wisconsin – Milwaukee US Postal Service: P.O. Box 413 Milwaukee, WI 53201 FedEx or UPS: 3210 N Cramer Street Milwaukee, WI 53211 Tuesday, April 29, 2003, is the absolute deadline for receipt of the exam materials at the Examinations Institute. Materials received after this deadline CANNOT be graded. THERE WILL BE NO EXCEPTIONS TO THIS DEADLINE DUE TO THE TIGHT SCHEDULE FOR GRADING THIS EXAMINATION. Page 2 EXAMINER’S NOTES Lab Problem #1: Materials and Equipment Each student should have available the following equipment and materials: • • • • • • • • • • • 60-mL plastic syringe with Luer-lock® tip cap Small plastic cap to hold sample inside syringe Electronic balance, 0.01 grams. One balance can serve 3-4 students. Please do not substitute a milligram balance. Student processing skills, not precision, are being evaluate here. Scoopula or spatula 400-mL or 600-mL beaker Weighing boats or weighing paper Capped vial or small beaker with film covering, at least 50 mL capacity Easy access to sinks Supply of paper towels 1 pair safety goggles 1 lab coat or apron (optional) Lab Problem #1: Chemicals Each student will need: Approximately 5 grams solid sample of Sani-Flush® Note: Solid Sani-Flush® is sold in the cleaning supply section of most supermarkets and convenience stores. Bottle of distilled water, at least 100 mL Examiner access to room temperature and barometric pressure Lab Problem #1: Notes 1. Note that the examiner will need to initial each student’s experimental plan to be sure that safety is considered. 2. You will need to provide students with the room Celsius temperature and barometric pressure (in mm Hg or atm). These data can be recorded on the board and also announced at the start of the experiment. 3. The small plastic cap can be from a soda bottle and should be able to fit with room inside the syringe without interfering with plunger operation. 4. The Sani-Flush should be in solid crystalline form. 5. Be sure that the syringe caps fit tightly over the tip of the syringe. 6. Check that the plastic bottle cap fit easily and completely inside the syringe and does the plunger slide easily when pushed. 7. It is your responsibility to ensure that students wear their safety goggles during the lab practical. A lab coat or apron for each student is desirable but not mandatory. You will also need to give students explicit directions for handling spills and for disposing of waste materials following approved safety practices for your examining site. Please check and follow procedures appropriate for your site. Page 3 Lab Problem #2: Materials and Equipment Each student will need: • Eight vials, with caps, 10-25 mL capacity each • 24-well plate or several spot plates or approximately fifteen 13 x 100mm test tubes • tube rack (if test tubes are used) • A disposal container for chemical waste (designated for heavy metal waste) • 3-4 Beral-style pipets, ungraduated 5-mL capacity • 400-mL beaker to hold vials • paper towels • easy access to sink • one pair safety goggles • one pair lab coat or apron (optional) Note: The 24-well plates and Beral-style pipets can be purchased through Educational Innovations or Micro Mole Scientific among other vendors. Lab Problem #2: Chemicals 3-4 grams each of Na3PO4, NH4Cl, ZnCl2, KNO3, Mg(OH)2, Pb(NO3)2, CaCO3, and Na2SO3 in vials labeled 1 – 8. Since this is an identification experiment, do not identify the contents of each vial! The order should be as follows: Vial 1 NH4Cl Vial 2 CaCO3 Vial 3 Vial 4 Vial 5 Pb(NO3)2 Mg(OH)2 KNO3 Vial 6 Na2SO3 Vial 7 Na3PO4 Vial 8 ZnCl2 Bottle of distilled water, at least 100 mL Additionally, the examiner will prepare 250 mL of 6 M H2SO4, 250 mL of 6 M HCl, 250 mL of 1 M AgNO3, 250 mL of 0.1% phenolphthalein. It is suggested that these reagents be placed in 400 mL beakers. Fill the number of 5 mL Beral-type pipets with each reagent as the number of students present. Student may select only two of these four reagents to use in their experiment. When they request the two solutions, be sure to indicate these on their answer sheets. They may not change their choices once they’ve begun using these two solutions, nor may they use more than one pipet of each of the two solutions selected. Lab Problem #2: Notes 1. Note that the examiner will need to initial each student’s experimental plan to be sure that safety is considered. 2. Make sure that the vials are NOT labeled with their contents. 3. Do not let students know which reagents might be better choices to select. Remind them that they may only select one pipet of each of two of these additional reagents, without refills. 4. The hydrated form of several of these salts (Na3PO4, ZnCl2) should be used, not the anhydrous form. 5. The phenolphthalein solution is a 0.1% solution dissolved in ethanol. Most pre-prepared solutions are the standard 0.1% solution. Page 4 2003 U.S. NATIONAL CHEMISTRY OLYMPIAD KEY FOR NATIONAL EXAM – PART III Problem 1. Sani-Flush®, a commercial toilet bowl cleaner, contains sodium bisulfate and sodium carbonate as its active ingredients and other ingredients such as sodium chloride, sodium lauryl sulfate, talc and some fragrance. Given the information from the manufacturer that the sodium bisulfate is substantially in excess compared to the sodium carbonate, carry out an experiment to determine the percent by weight of the sodium carbonate in a sample of the product. Experimental Plan: 2HSO4 – + CO32– ! CO2 + H2O + 2SO4 2– A good plan consisted of weighing a sample of Sani-Flush® and determining the volume of CO2 produced when the sample is added to water. For example, a good plan might include these steps, 1. Weigh approximately 1 g of Sani-Flush®. 2. Add to syringe. 3. Draw water into the syringe and cover with cap. 4. Determine the volume of gas (CO2 ) produced. 5. Use ideal gas equation to calculate moles of CO 2 6. Convert moles of CO2 to moles of sodium carbonate and then mass of Na2 CO3 7. Divide mass of Na2 CO3 by sample mass to find percentage. 8. Repeat with second sample. An average plan was either missing one of these components or had a procedure based solely on the difference in mass. Such a procedure is subject to greater error due to the smaller change in mass relative to the change in volume. A weak plan had minimal detail about how the experiment would be conducted. Data and other Observations: An example of good work on observations and data recording would be, 1.45 g sample placed in syringe. Set plunger at 5.0 mL. Drew in 14.0 mL of H2O. Gas was evolved and the solution was blue. Plunger reached 47.8 mL after the reaction was complete. mass Sani-Flush® Trial 1 Trial 2 2.05 g 1.60 g initial plunger level 5.0 mL 5.0 mL Temperature is 23 oC and pressure is 755 mmHg Key for 2003 USNCO National Exam, Part III H2 O volume added 14.0 mL 15.0 mL final plunger level 47.8 mL 39.0 mL Sample Calculations: Sample 1: 47.8 – 14 = 33.8 mL removing water volume: 33.8 mL – 5.0 mL = 28.8 mL gas evolved. 755 mmHg -1 (0.0288 L ) 760 mmHg ⋅ atm PV mole of gas: PV = nRT therefore n = = RT (0.0821L ⋅ atm ⋅ mol-1 ⋅ K -1 )(296 K) = 0.00118 mol CO2 106 g convert to mass: 0.00118 mol CO2 = 0.00118 mol Na 2CO 3 × = 0.125 g Na 2CO 3 1 mol calculate percentage: %= 0.125 g × 100 = 6.1% 2.05 g Problem 2: You have been provided with eight vials, each of which is labeled with a number from 1 to 8. Each vial contains one of the following chemicals: Na3 PO4, NH4Cl, ZnCl2 , KNO3 , Mg(OH)2 , Pb(NO3 )2 , CaCO3 , Na2 SO3. You are allowed to use distilled water, test tubes or well plates, and only TWO additional reagents from the following choices: 6 M H2SO4, 6 M HCl, 6 M AgNO3 , phenolphthalein indicator solution You must designate your choice of reagents prior to the start of your testing. Experimental Plan A good plan involved stating that samples of each of the eight unknowns would be placed in wells of the spot plates, distilled water would be added and observations would be made regarding the solubility of the salts. Then, each of th SELECTED reagents would be added and specific identifying tests would be detailed. It should be noted that several combinations of reagents could have been used to identify the unknowns. For example, a good test might include these tests. 1. Place each sample into each of three wells in the spot plate. Add distilled water to each and record solubility r 2. Two of the samples will not dissolve. Add HCl to these. One sample should dissolve and the other should fiz former is magnesium hydroxide, the latter is calcium carbonate. 3. Add HCl to the other solutions. Lead (II) nitrate will form a precipitate. Sodium sulfite will release sulfur dio which can be identified by odor. 4. Add silver nitrate to the other four solutions. The sample which does not react is potassium nitrate. The yello precipitate is produced by sodium phosphate and the two white solids are zinc chloride and ammonium chlorid 5. Prepare additional solutions of the two chlorides and the identified sodium phosphate. Add the latter to the chlorides. The zinc chloride will form a precipitate while the ammonium chloride will not react. An average plan did not detail the expected results of specific tests. A weak plan did not include solution formation or specific results of tests. Key for 2003 USNCO National Exam, Part III Observations and results: An example of good work on observations and results would be: Reagent Sample 1 Sample 2 Sample 3 Sample 4 Sample 5 water soluble soluble fizz No Rx AgNO3 No Rx White ppt soluble White ppt insoluble HCl insoluble dissolve s Na3PO4 No Rx Key for 2003 USNCO National Exam, Part III No Rx Sample 6 Sample 7 Sample 8 soluble Sharp odor soluble soluble No Rx Yellow Ppt No Rx White ppt White ppt 2004 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM Part I Prepared by the American Chemical Society Olympiad Examinations Task Force OLYMPIAD EXAMINATIONS TASK FORCE Arden P. Zipp, State University of New York, Cortland Chair Sherry Berman-Robinson, Consolidated High School, IL David W. Hostage, Taft School, CT William Bond, Snohomish High School, WA Alice Johnsen, Bellaire High School, TX Peter E. Demmin (retired), Amherst Central High School, NY Marian Dewane, Centennial High School, ID Adele Mouakad, St. John’s School, PR Ronald O. Ragsdale, University of Utah, UT Dianne Earle, Boiling Springs High School, SC Jacqueline Simms, Sandalwood Sr. High School, FL Michael Hampton, University of Central Florida, FL DIRECTIONS TO THE EXAMINER–PART I Part I of this test is designed to be taken with a Scantron® answer sheet on which the student records his or her responses. Only this Scantron sheet is graded for a score on Part I. Testing materials, scratch paper, and the Scantron sheet should be made available to the student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until April 19, 2004, after which tests can be returned to students and their teachers for further study. Allow time for the student to read the directions, ask questions, and fill in the requested information on the Scantron sheet. The answer sheet must be completed using a pencil, not pen. When the student has completed Part I, or after one hour and thirty minutes has elapsed, the student must turn in the Scantron sheet, Part I of the testing materials, and all scratch paper. There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and you are free to schedule rest-breaks between parts. Part I Part II Part III 60 questions 8 questions 2 lab problems single-answer multiple-choice problem-solving, explanations laboratory practical 1 hour, 30 minutes 1 hour, 45 minutes 1 hour, 30 minutes A periodic table and other useful information are provided on page 2 for student reference. Students should be permitted to use nonprogrammable calculators. DIRECTIONS TO THE EXAMINEE–PART I DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Answers to questions in Part I must be entered on a Scantron answer sheet to be scored. Be sure to write your name on the answer sheet; an ID number is already entered for you. Make a record of this ID number because you will use the same number on both Parts II and III. Each item in Part I consists of a question or an incomplete statement that is followed by four possible choices. Select the single choice that best answers the question or completes the statement. Then use a pencil to blacken the space on your answer sheet next to the same letter as your choice. You may write on the examination, but the test booklet will not be used for grading. Scores are based on the number of correct responses. When you complete Part I (or at the end of one hour and 30 minutes), you must turn in all testing materials, scratch paper, and your Scantron answer sheet. Do not forget to turn in your U.S. citizenship statement before leaving the testing site today. Not valid for use as an USNCO Olympiad National Exam after April 19, 2004. Distributed by the ACS DivCHED Examinations Institute, University of Wisconsin - Milwaukee, Milwaukee, WI. All rights reserved. Printed in U.S.A. ABBREVIATIONS AND SYMBOLS A Faraday constant F molal atm formula molar mass M molar u free energy G molar mass A frequency ν mole N A gas constant R Planck’s constant °C gram g pressure c heat capacity C p rate constant C hour h retention factor E joule J second Ea kelvin K temperature, K H kilo– prefix k time S liter L volt K milli– prefix m ampere atmosphere atomic mass unit atomic molar mass Avogadro constant Celsius temperature centi– prefix coulomb electromotive force energy of activation enthalpy entropy equilibrium constant CONSTANTS m M M mol h P k Rf s T t V R = 8.314 J·mol–1·K–1 R = 0.0821 L·atm·mol –1·K–1 1 F = 96,500 C·mol–1 1 F = 96,500 J·V–1·mol–1 N A = 6.022 × 1023 mol–1 h = 6.626 × 10–34 J·s c = 2.998 × 108 m·s–1 0 °C = 273.15 K 1 atm = 760 mmHg EQUATIONS E = Eo − 1 1A 1 H k E 1 1 ln 2 = a − k1 R T1 T2 −∆H 1 ln K = + constant R T RT ln Q nF PERIODIC TABLE OF THE ELEMENTS 18 8A 2 He 3 Li 2 2A 4 Be 13 3A 5 B 14 4A 6 C 15 5A 7 N 16 6A 8 O 17 7A 9 F 6.941 9.012 10.81 12.01 14.01 16.00 19.00 20.18 11 Na 12 Mg 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar 22.99 24.31 26.98 28.09 30.97 32.07 35.45 39.95 19 K 31 Ga 32 Ge 33 As 34 Se 35 Br 36 Kr 1.008 4.003 10 Ne 20 Ca 3 3B 21 Sc 4 4B 22 Ti 5 5B 23 V 6 6B 24 Cr 7 7B 25 Mn 8 8B 26 Fe 9 8B 27 Co 10 8B 28 Ni 11 1B 29 Cu 12 2B 30 Zn 39.10 40.08 44.96 47.88 50.94 52.00 54.94 55.85 58.93 58.69 63.55 65.39 69.72 72.61 74.92 78.96 79.90 83.80 37 Rb 38 Sr 39 Y 40 Zr 41 Nb 42 Mo 43 Tc 44 Ru 45 Rh 46 Pd 47 Ag 48 Cd 49 In 50 Sn 51 Sb 52 Te 53 I 54 Xe 85.47 87.62 88.91 91.22 92.91 95.94 (98) 101.1 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.3 55 Cs 56 Ba 57 La 72 Hf 73 Ta 74 W 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 132.9 137.3 138.9 178.5 180.9 183.8 186.2 190.2 192.2 195.1 197.0 200.6 204.4 207.2 209.0 (209) (210) (222) 87 Fr 88 Ra 89 Ac 104 Rf 105 Db 106 Sg 107 Bh 108 Hs 109 Mt 110 111 112 114 (223) (226) (227) (261) (262) (263) (262) (265) (266) (269) (272) (277) (2??) Page 2 58 Ce 59 Pr 60 Nd 61 Pm 62 Sm 63 Eu 64 Gd 65 Tb 66 Dy 67 Ho 68 Er 69 Tm 70 Yb 71 Lu 140.1 140.9 144.2 (145) 150.4 152.0 157.3 158.9 162.5 164.9 167.3 168.9 173.0 175.0 90 Th 91 Pa 92 U 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr 232.0 231.0 238.0 (237) (244) (243) (247) (247) (251) (252) (257) (258) (259) (262) Not valid for use as an USNCO Olympiad National Exam after April 19, 2004. DIRECTIONS ! When you have selected your answer to each question, blacken the corresponding space on the answer sheet using a soft, #2 pencil. Make a heavy, full mark, but no stray marks. If you decide to change an answer, erase the unwanted mark very carefully. ! There is only one correct answer to each question. Any questions for which more than one response has been blackened will not be counted. ! Your score is based solely on the number of questions you answer correctly. It is to your advantage to answer every question. 1. Which element is obtained commercially from seawater? (A) bromine (B) gold (C) iron (D) oxygen 2. Which solution can serve as both reactant and indicator when it is used in redox titrations? (A) FeNH4(SO4)2 (B) KMnO4 (C) H2C2O4 (D) Na 2S 2O3 (A) N2 and H2O (B) N2O and H2O (C) NO and H2 (D) N2, H2 and O2 4. Which method should be used to extinguish burning magnesium metal? (A) Blanket it with CO2 (B) Blow on it. (C) Dump sand on it. (D) Pour water on it. A C B D (A) A (B) B (C) C (A) Ca (B) Mn (C) Ni (D) Zn 8. What is the coefficient for OH- after the equation _ Br2 + _ OH- r _ Br- + _ BrO3- + _ H2O is balanced with the smallest integer coefficients? 3. What is formed when a solution of NH4NO2 is heated gently? 5. Which letter indicates where a thermometer should be placed to determine the boiling point of a distillate? 7. A 1.871 gram sample of an unknown metallic carbonate is decomposed by heating to form the metallic oxide and 0.656 g of carbon dioxide according to the equation MCO3(s) r MO(s) + CO2(g) What is the metal? (A) 3 (B) 6 (C) 12 (D) 18 9. An ionic compound contains 29.08% sodium, 40.56% sulfur and 30.36% oxygen by mass. What is the formula of the sulfur-containing anion in the compound? (A) S 2O32- (B) S 2O42- (C) S 2O52- (D) S 2O62- 10. A solution is prepared Vapor pressure (mmHg) containing a 2:1 mol ratio of C 2H4Br2 173 dibromoethane (C2H4Br2) and dibromopropane (C 3H6Br2). C 3H6Br2 127 What is the total vapor pressure over the solution assuming ideal behavior? (A) 300 mmHg (B) 158 mmHg (C) 150 mmHg (D) 142 mmHg (D) D 6. A 50 mL sample of gas is collected over water. What will be the effect on the calculated molar mass of the gas if the effect of the water vapor is ignored? It will be (A) high because of the mass of water in the collection flask. (B) high because of omitting the vapor pressure of the water in the calculation. (C) low because of the mass of water in the collection flask. 11. A solution of magnesium chloride that is 5.10% magnesium by mass has a density 1.17 g/mL. How many moles of Cl- ions are in 300. mL of the solution? (A) 0.368 (B) 0.627 (C) 0.737 (D) 1.47 12. Which aqueous solution has a freezing point closest to that of 0.30 M C12H22O11? (A) 0.075 M AlCl3 (B) 0.15M CuCl2 (C) 0.30 M NaCl (D) 0.60 M C6H12O6 (D) low because of omitting the vapor pressure of the water in the calculation. Not valid for use as an USNCO Olympiad National Exam after April 19, 2004. Page 3 13. An unknown gas is placed in a sealed container with a fixed volume. Which of the characteristics listed change(s) when the container is heated from 25 ˚C to 250 ˚C? I The density of the gas II The average kinetic energy of the molecules III The mean free path between molecular collisions (B) II only (C) III only (D) I and II only 14. Which gas has the same density at 546 ˚C and 1.50 atm as that of O2 gas at STP? (B) NH3 (C) SO2 (D) SO3 15. Which plot involving vapor pressure (VP) and absolute temperature results in a straight line? (A) VP vs T (B) VP vs T-1 (C) ln VP vs T (D) ln VP vs T-1 16. For a substance with the values of ∆Hvap and ∆Svap given below, what is its normal boiling point in ˚C? (∆Hvap = 59.0 kJ . mol -1; ∆Svap = 93.65 J. mol -1 .K-1) (A) 357 (B) 630 (C) 1314 (D) 1587 17. What is the order of the boiling points (from lowest to highest) for the hydrogen halides? (A) HF < HCl < HBr < HI (B) HI < HBr < HCl < HF (C) HCl < HF < HBr < HI (D) HCl < HBr < HI < HF 18. Of the three types of cubic lattices, which have the highest and lowest densities for the same atoms? Highest (A) S˚200K is smaller because entropy decreases as temperature increases. (B) S˚200K is smaller because the surroundings are more disordered at higher temperatures. (A) I only (A) N2 20. Which is the best description of the relationship between the absolute entropies, S˚, of solid water at 100 K and at 200 K? Lowest (C) S˚100K = S˚200K = because water is in the solid phase at both temperatures. (D) S˚200K is larger because the vibration of the molecules increases as temperature increases. 21. For the reaction, CH4 + Cl2 r CH3Cl + HCl which expression gives ∆H? Bond dissociation energies C-H C-Cl Cl-Cl H-Cl kJ . mol-1 (A) ∆H = (413 + 328) - (242 + 431) (B) ∆H = (413 - 328) - (242 - 431) (C) ∆H = (413 - 242) - (328 - 431) (D) ∆H = (413 + 242) - (328 + 431) 22. Which phase change for water has positive values for both ∆H˚ and ∆G˚? (A) (l) r (s) at 250 K (B) (l) r (s) at 350 K (C) (l) r (g) at 350 K (D) (l) r (g) at 450 K 23. When solid CuSO4 dissolves in water to make a 1M solution, the temperature of the system increases. When solid NH4NO3 dissolves in water to make a 1 M solution, the temperature of the system decreases. Which statement(s) must be correct for these dissolving processes? I ∆H˚ values for both processes have the same sign. II ∆G˚ values for both processes have the same sign. (A) simple cubic body-centered cubic (B) face-centered cubic simple cubic (C) body-centered cubic face-centered cubic (A) I only (B) II only D) face-centered cubic body-centered cubic (C) Both I and II (D) Neither I nor II 19. For which reaction is ∆H (enthalpy change) most nearly equal to ∆E (internal energy change)? 24. Which set of relationships could apply to the same electrochemical cell? (A) H2(g) + 1/2O2(g) r H2O(g) (A) ∆G˚ > 0; E˚ = 0 (B) ∆G˚ < 0; E˚ = 0 (B) Cl2(g) + F2(g) r 2ClF(g) (C) ∆G˚ > 0; E˚ > 0 (D) ∆G˚ < 0; E˚ > 0 (C) H2O(l) r H2O(g) (D) 2SO3(g) r 2SO2(g) + O 2(g) Page 4 413 328 242 431 25. The rate constant for a reaction is affected by which factors? I increase in temperature II concentration of the reactants III presence of a catalyst (A) I and II only (B) I and III only (C) II and III only (D) I, II and III Not valid for use as an USNCO Olympiad National Exam after April 19, 2004. 26. The rate data given were obtained for the reaction, 2NO(g) + 2H2(g) r N2(g) + 2H2O(g) What is the rate law for this reaction? NO pressure (atm) H2 pressure (atm) Rate (atm. sec-1) 0.375 0.500 6.43 × 10-4 0.375 0.250 3.15 × 10-4 0.188 0.500 1.56 × 10-4 (A) Rate = k PNO 2 (B) Rate = k PNO (C) Rate = k PNO PH2 2 2 (D) Rate = k PNO PH 2 time 1 / concentration concentration log (concentration) 27. What is the order of a reaction that produces the graphs shown? 31. H2S(aq) s H+(aq) + HS-(aq) K = 9.5 × 10-8 + 2– HS (aq) s H (aq) + S (aq) K = 1.0 × 10-19 Given the equilibrium constants provided, what is the equilibrium constant for the reaction; S 2–(aq) + 2H+(aq) s H2S(aq) K=? (A) 9.5 × 10-27 (B) 9.7 × 10-14 (C) 9.5 × 1011 (D) 1.0 × 1026 32. Calculate the hydronium ion concentration in 50.0 mL of 0.10 M NaH2AsO4. (K 1 = 6.0 × 10-3, K2 = 1.1 × 10-7 K3 = 3.0 × 10-12) (A) 2.4 × 10-2 (B) 1.6 × 10-3 (C) 1.0 × 10-4 (D) 2.5 × 10-5 33. When the acids; HClO3, H3BO3, H3PO4, are arranged in order of increasing strength, which order is correct? (A) H3BO3 < H3PO4 < HClO 3 time time (B) HClO3 < H3BO3 < H3PO4 (A) zero order (B) first order (C) H3PO4 < HClO 3 < H3BO3 (C) second order (D) some other order (D) H3BO3 < HClO 3 < H3PO4 28. What is the rate law for the hypothetical reaction with the mechanism shown? 2A s intermediate 1 fast equilibrium intermediate 1 + B r intermediate 2 slow fast intermediate 2 + B r A2B2 (A) Rate = k[A]2 (C) Rate = k[A][B] (B) Rate = [B]2 2 (D) Rate = k[A] [B] 29. According to the Arrhenius equation: k = Ae -Ea/RT , a plot of ln k against 1/T yields (A) Ea as the slope and A as the intercept (B) Ea/R as the slope and A as the intercept (C) Ea/R as the slope and ln A as the intercept (D) -Ea/R as the slope and ln A as the intercept (A) II only (B) I and II only (C) I and III only (D) I, II and III 35. A solution is 0.10 M in Ag+, Ca2+, Mg2+, and Al3+ ions. Which compound will precipitate at the lowest [PO43-] when a solution of Na3PO4 is added? (A) Ag3PO4 (Ksp = 1 × 10-16) (B) Ca3(PO4)2 (Ksp = 1 × 10-33) (C) Mg3(PO4)2 (Ksp = 1 × 10-24) (D) AlPO4 (Ksp = 1 × 10-20) 36. Which salt is significantly more soluble in a strong acid than in water? Reaction rate 30. Curves with the shape shown are often observed for reactions involving catalysts. The level portion of the curve is best attributed to the fact that 34. A buffer solution results from mixing equal volumes of which solutions? I 0.10 M HCl and 0.20 M NH3 II 0.10 M HNO 2 and 0.10 M NaNO2 III 0.20 M HCl and 0.10 M NaCl (A) PbF2 Reactant concentration (A) product is no longer being formed. (B) the reaction has reached equilibrium. (C) all the catalytic sites are occupied. (D) all the reactant has been consumed. (B) PbCl2 (C) PbBr 2 (D) PbI2 37. What is the standard cell potential for the reaction, 2Cr(s) + 3Sn2+(aq) r 3Sn(s) + 2Cr3+(aq) given the E˚ values shown? Cr3+(aq) + 3e- r Cr(s) -0.744 V Sn 2+(aq) + 2e- r Sn(s) -0.141 V (A) 0.945 V (B) 0.603 V (C) -0.603 V (D) -0.945 V Not valid for use as an USNCO Olympiad National Exam after April 19, 2004. Page 5 38. How many electrons are needed in the balanced halfreaction for the oxidation of ethanol to acetic acid? C 2H5OH r CH3COOH (A) 1 (B) 2 (C) 3 (B) Cu 2+(aq) (C) H+(aq) (D) Zn2+(aq) (A) N (B) 22 (C) 6.1 × 10 (D) 3.8 × 109 41. When an aqueous solution of potassium fluoride is electrolyzed, which of the following occurs? (A) O2 and H+ are produced at one electrode and H2 and OH- are formed at the other. (B) O2 and OH- are produced at one electrode and H2 and H+ are formed at the other. (C) Metallic K is formed at one electrode and O2 and H+ are formed at the other. (D) Metallic K is produced at one electrode and elemental F2 is produced at the other. 42. A CuSO4 solution is electrolyzed for 20. minutes with a current of 2.0 ampere. What is the maximum mass of copper that could be deposited? (A) 0.20 g (B) 0.40 g (C) 0.79 g (D) 1.6 g 43. Which experimental evidence most clearly supports the suggestion that electrons have wave properties? (A) diffraction (A) 0 (D) deflection of cathode rays by a magnet 44. Which quantum number determines the number of angular nodes in an atomic orbital? (C) ml (D) ms 45. Which element exhibits the greatest number of oxidation states in its compounds? (A) Ca Page 6 (B) V (C) 4 (D) 6 12 7N 18 8O (B) (C) 20 9F (D) 20 10 Ne 49. According to the Lewis dot O C N structure shown, what are the formal charges of the O, C and N atoms, respectively, in the cyanate ion? (A) 0, 0, 0 (B) -1, 0, 0 (C) -1, +1, -1 (D) +1, 0, -2 50. The hybridization of As in AsF5 is best described as (A) sp3 (B) sp4 (C) dsp3 (D) d2sp3 51. In which species do the atoms NOT lie in a single plane? (A) BF3 (B) PF3 (C) ClF3 (D) XeF4 52. For which compound does the reaction, MCO3(s) r MO(s) + CO2(g) occur most readily? (A) BeCO3 (B) MgCO3 (C) CaCO3 (D) BaCO3 53. The color of Co(H2O)62+ is best attributed to electronic transitions (A) between different n levels in the metal. (B) between the metal's d orbitals. (D) during ionization. (C) photoelectric effect (B) l (B) 2 (C) from the Co2+ ion to water molecules. (B) emission spectra (A) n (D) Cl 48. Which species is most likely to lose a positron (β+)? 40. The standard potential for the reaction Cl2(g) + 2Br -(aq) ---> Br2(l) + 2Cl-(aq) is 0.283 volts. What is the equilibrium constant for this reaction at 25 ˚C? 4 (C) S 47. How many unpaired electrons are in a gaseous Fe2+ ion in its ground state? (A) (A) 1.6 × 10-5 (B) P (D) 4 39. Which is the weakest oxidizing agent in a 1 M aqueous solution? (A) Ag+(aq) 46. Of the elements given, which has the lowest ionization energy? (C) Cu 54. When the carbon-oxygen bonds in the species; CH3OH, CH2O and CHO2- are arranged in order of increasing length, which is the correct order? (A) CH3OH < CH2O < CHO2(B) CH2O < CH3OH < CHO2(C) CHO - < CH OH < CH2O 2 3 (D) CH2O < CHO2- < CH3OH (D) As Not valid for use as an USNCO Olympiad National Exam after April 19, 2004. 55. How many different trichlorobenzenes, C6H3Cl3, can be formed? (A) 1 (B) 2 (C) 3 (D) 4 56. What organic product is formed from the mild oxidation of a secondary alcohol? (A) acid (B) aldehyde (C) ether (D) ketone 57. The compound with the formula, H 2NCH2CH2COOH, is best classified as a(n) (A) amide (B) amino acid (C) fatty acid (D) nucleic acid 58. The reaction between which pair of reactants occurs the fastest for [OH-] = 0.010 M? (A) CH3CH2CH2CH2Cl + OH (B) (CH3)3CCl + OH(C) CH3CH2CH2CH2Br + OH(D) (CH3)3CBr + OH59. What is the major organic product formed from the reaction of CH3CH=CH2 and HCl? (A) CH3CHClCH3 (B) CH3CH2CH2Cl (C) CH3CHClCH2Cl (D) CH2ClCH=CH2 60. Fats and oils are formed from the combination of fatty acids with what other compound? (A) cholesterol (B) glucose (C) glycerol (D) phenol END OF TEST Not valid for use as an USNCO Olympiad National Exam after April 19, 2004. Page 7 National Olympiad 2004 Part 1 KEY Number 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. Answer A B A C A D D B A B D A B C D A D B B D D C B D B D C D D C Number 31. 32. 33. 34. 35. 36. 37. 38. 39. 40. 41. 42. 43. 44. 45. 46. 47. 48. 49. 50. 51. 52. 53. 54. 55. 56. 57. 58. 59. 60. Property of the ACS DivCHED Examinations Institute Answer D D A B D A B D D D A C A B B C C A D C B A B D C D B D A C 2004 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART II Prepared by the American Chemical Society Olympiad Examinations Task Force OLYMPIAD EXAMINATIONS TASK FORCE Arden P. Zipp, State University of New York, Cortland Chair Sherry Berman-Robinson, Consolidated High School, IL David W. Hostage, Taft School, CT William Bond, Snohomish High School, WA Alice Johnsen, Bellaire High School, TX Peter E. Demmin (retired), Amherst Central High School, NY Marian Dewane, Centennial High School, ID Adele Mouakad, St. John’s School, PR Ronald O. Ragsdale, University of Utah, UT Dianne Earle, Boiling Springs High School, SC Jacqueline Simms, Sandalwood Sr. High School, FL Michael Hampton, University of Central Florida, FL DIRECTIONS TO THE EXAMINER–PART II Part II of this test requires that student answers be written in a response booklet of blank pages. Only this “Blue Book” is graded for a score on Part II. Testing materials, scratch paper, and the “Blue Book” should be made available to the student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until April 19, 2004, after which tests can be returned to students and their teachers for further study. Allow time for the student to read the directions, ask questions, and fill in the requested information on the “Blue Book”. When the student has completed Part II, or after one hour and forty-five minutes has elapsed, the student must turn in the “Blue Book”, Part II of the testing materials, and all scratch paper. Be sure that the student has supplied all of the information requested on the front of the “Blue Book,” and that the same identification number used for Part I has been used again for Part II. There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and you are free to schedule rest-breaks between parts. Part I Part II Part III 60 questions 8 questions 2 lab problems single-answer multiple-choice problem-solving, explanations laboratory practical 1 hour, 30 minutes 1 hour, 45 minutes 1 hour, 30 minutes A periodic table and other useful information are provided on the back page for student reference. Students should be permitted to use non-programmable calculators. DIRECTIONS TO THE EXAMINEE–PART II DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Part II requires complete responses to questions involving problem-solving and explanations. One hour and forty-five minutes are allowed to complete this part. Be sure to print your name, the name of your school, and your identification number in the spaces provided on the “Blue Book” cover. (Be sure to use the same identification number that was coded onto your Scantron® sheet for Part I.) Answer all of the questions in order, and use both sides of the paper. Do not remove the staple. Use separate sheets for scratch paper and do not attach your scratch paper to this examination. When you complete Part II (or at the end of one hour and forty-five minutes), you must turn in all testing materials, scratch paper, and your “Blue Book.” Do not forget to turn in your U.S. citizenship statement before leaving the testing site today. Not valid for use as an USNCO National Exam after April 19, 2004. Distributed by the ACS DivCHED Examinations Institute, University of Wisconsin-Milwaukee, Milwaukee, WI. All rights reserved. Printed in U.S.A. 1. 2. (12 %) A solution of copper(II) sulfate that contains 15.00% CuSO4 by mass has a density of 1.169 g/mL. A 25.0 mL portion of this solution was reacted with excess concentrated ammonia to form a dark blue solution. When cooled, filtered and dried, 6.127 g of a dark blue solid were obtained. A 0.195 g sample of the solid was analyzed for ammonia by titrating with 0.1036 M hydrochloric acid solution, requiring 30.63 mL to reach the equivalence point. A 0.150 g sample was analyzed for copper (II) by titrating with 0.0250 M EDTA, (which reacts with Cu2+ in a 1:1 ratio). The endpoint was reached after 24.43 mL of the EDTA were added. A 0.200 g sample was heated at 110 ˚C to drive off water, producing 0.185 g of the anhydrous material. a. Determine the molarity of Cu2+ ions in the original solution. b. Find the number of moles of Cu2+ in the 25.0 mL portion. c. Calculate the percentages by mass in the prepared compound of; i. NH3 ii. Cu 2+ iii. H2O iv. SO42d. Use the results in part c. to determine the formula of the compound. e. Assuming that Cu2+ is the limiting reactant in the synthesis, determine the percent yield. (16%) 2NO2(g) + O 3(g) r N2O5(g) + O2(g) ∆H˚ = -198 kJ ∆S˚ = -168 J . K-1 Ozone reacts with nitrogen dioxide according to the equation above. a. Calculate ∆H f˚ for NO2(g) in kJ.mol -1. [∆H f˚ kJ. mol -1; O3(g) 143, N2O5(g) 11] b. Account for the sign of ∆S˚. c. Calculate the value of ∆G˚ at 25 ˚C. d. State and explain how the spontaneity of this reaction will vary with increasing temperature. e. Use the rate data below to determine the rate law for the reaction of NO2(g) and O3(g) NO2(g), M O3(g), M Rate M. s-1 0.0015 0.0025 4.8×10-8 0.0022 0.0025 7.2×10-8 0.0022 0.0050 1.4×10-7 f. Calculate the specific rate constant and give its units. g. The following mechanisms have been proposed for this reaction. Discuss the suitability of each to account for the rate law obtained. Mechanism I Mechanism II NO2 + NO2 r NO3 + NO slow NO2 + O3 s NO3 + O2 fast NO3 + NO2 r N2O5 fast NO3 + NO2 r N2O5 slow NO + O3 r NO2 + O2 fast h. Describe and account for any change expected in the rate of this reaction as the temperature is increased. 3. (10%) A popular lecture demonstration involves the sequential precipitation and dissolution of several slightly soluble silver compounds beginning with a [Ag+ ] = 0.0050 M. Use the information below to answer the following questions about this demonstration. [Ksp values; AgCl 1.8×10-10, AgBr 5×10-13, AgI 8.3×10-17, Ag2SO4 1.4×10-5] a. What must the [SO42-] be in order to start precipitation in a solution in which [Ag+] = 0.005 M? b. State the order in which the halide ions should be added to a concentration of 0.10 M so that each precipitate will form from the [Ag+] in equilibrium with the previous precipitate. Support your answer with appropriate calculations. c. As a way of making this demonstration more striking, one of the silver halides in this series is dissolved by adding aqueous ammonia before precipitating the next silver halide. Which silver halide(s) dissolve in 0.60 M NH3? Support your answer with calculations. [Kf Ag(NH3)2+ 1.7×107] Page 2 Not valid for use as an USNCO National Exam after April 19, 2004 4. (12%) Aluminum metal is obtained commercially by electrolyzing Al2O3 mixed with cryolite (Na3AlF 6). a. Explain why electrolysis is used rather than heating the Al 2O3 either directly or in the presence of C (as is done to extract Fe or Zn from their ores). b. State the purpose of the Na 3AlF 6. c. Write the two half-reactions that occur during electrolysis and indicate which of the two occurs at the cathode d. How many moles of electrons must pass through the cell to produce 5.00 kg of Al? (Assume 100% efficiency.) e. Determine the current required (in amperes) if the aluminum in d. is produced in 10.0 hours. f. Calculate the volume of gas formed in the process in d. at 25 ˚C and 720 mmHg. 5. (12%) Write net equations for each of the combinations of reactants below. Use appropriate ionic and molecular formulas and omit formulas for all ions or molecules that do not take part in a reaction. Write structural formulas for all organic substances. You need not balance the equations. All reactions occur in aqueous solution unless otherwise indicated. a. Concentrated nitric acid is added to iron(II) sulfide. b. Acetic acid is added to solid calcium phosphate. c. Small pieces of aluminum metal are added to concentrated sodium hydroxide solution. d. Solutions of chromium(III) sulfate and barium hydroxide are mixed. e. Concentrated hydrochloric acid is added to an aqueous solution of cobalt(II) nitrate. f. Bromine gas is added to propene. 6. (16%) Account for the following statements or observations in terms of atomic-, ionic- or molecular-level explanations. a. Magnesium exists as +2 ions rather than +1 ions in all of its compounds despite the fact that the second ionization energy of a magnesium atom is more than twice as great as the first ionization energy. b. Titanium forms ions with different charges (+2, +3 and +4). The first two of these ions are colored while the last is colorless. c. Carbon dioxide (CO 2) is a gas at room temperature but silicon dioxide (SiO2) is a high melting solid. d. Nitrogen forms NF 3 but not NF5 whereas phosphorus forms PF3 and PF5. The trifluorides are both trigonal pyramidal and the pentafluoride is trigonal bipyramidal. 7. (12%) Explain these observations about nuclei. a. Elements with even atomic numbers tend to have more stable isotopes than elements with odd atomic numbers. b. Carbon-14 decays with the loss of a β- particle while carbon-10 decays with the loss of a β+. c. Carbon-14 (t1/2 = 5730 years) can be used to determine the age of organic materials that died between approximately 500 and 50,000 years ago. 8. (10%) A student is asked to determine the molar mass of an unknown monoprotic carboxylic acid by freezing point depression. The student dissolves 0.029 g of the unknown acid in 10. mL of water and obtains the cooling curve shown. o Temperature, C 0.00 -0.02 -0.04 a. Account for the shape of the curve including the i. downward slope of the final portion of the curve. Time, min ii. depression before the final portion of the curve. b. Give the value of the freezing point of the solution and calculate the molality of the solute. [Kf = -1.86 ˚C. m-1] c. Determine the molar mass of the acid. d. When the student titrates a solution of the acid in water, a molar mass of 120. g.mol -1is determined. Compare these results with those in part c. and offer an explanation for this behavior. END OF PART II Not valid for use as an USNCO National Examination after April 19, 2004. Page 3 amount of substance ampere atmosphere atomic mass unit atomic molar mass Avogadro constant Celsius temperature centi- prefix coulomb electromotive force energy of activation enthalpy entropy ABBREVIATIONS AND SYMBOLS n equilibrium constant K measure of pressure mmHg A Faraday constant F milli- prefix m atm formula molar mass M molal m u free energy G molar M A frequency ν mole mol N A gas constant R Planck’s constant h °C gram g pressure P c heat capacity C p rate constant k C hour h retention factor Rf E joule J second s Ea kelvin K speed of light c H kilo- prefix k temperature, K T S liter L time t volt V CONSTANTS R = 8.314 J·mol –1·K–1 R = 0.0821 L·atm·mol –1·K–1 1 F = 96,500 C·mol–1 1 F = 96,500 J·V–1·mol–1 N A = 6.022 × 1023 mol–1 h = 6.626 × 10–34 J·s c = 2.998 × 108 m·s–1 USEFUL EQUATIONS k2 Ea 1 1 = − k1 R T1 T2 – ∆H 1 ln K = +c R T RT E = Eο – ln Q nF ln PERIODIC TABLE OF THE ELEMENTS 1 H 2 He 1.008 4.003 3 Li 4 Be 5 B 6 C 7 N 8 O 9 F 10 Ne 6.941 9.012 10.81 12.01 14.01 16.00 19.00 20.18 11 Na 12 Mg 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar 22.99 24.31 26.98 28.09 30.97 32.07 35.45 39.95 19 K 20 Ca 21 Sc 22 Ti 23 V 24 Cr 25 Mn 26 Fe 27 Co 28 Ni 29 Cu 30 Zn 31 Ga 32 Ge 33 As 34 Se 35 Br 36 Kr 39.10 40.08 44.96 47.88 50.94 52.00 54.94 55.85 58.93 58.69 63.55 65.39 69.72 72.61 74.92 78.96 79.90 83.80 37 Rb 38 Sr 39 Y 40 Zr 41 Nb 42 Mo 43 Tc 44 Ru 45 Rh 46 Pd 47 Ag 48 Cd 49 In 50 Sn 51 Sb 52 Te 53 I 54 Xe 85.47 87.62 88.91 91.22 92.91 95.94 (98) 101.1 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.3 55 Cs 56 Ba 57 La 72 Hf 73 Ta 74 W 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 132.9 137.3 138.9 178.5 181.0 183.8 186.2 190.2 192.2 195.1 197.0 200.6 204.4 207.2 209.0 (209) (210) (222) 87 Fr 88 Ra 89 Ac 104 Rf 105 Db 106 Sg 107 Bh 108 Hs 109 Mt 110 111 112 114 (223) 226.0 227.0 (261) (262) (263) (262) (265) (266) (269) (272) (277) (277) Page 4 58 Ce 59 Pr 60 Nd 61 Pm 62 Sm 63 Eu 64 Gd 65 Tb 66 Dy 67 Ho 68 Er 69 Tm 70 Yb 71 Lu 140.1 140.9 144.2 (145) 150.4 152.0 157.3 158.9 162.5 164.9 167.3 168.9 173.0 175.0 90 Th 91 Pa 92 U 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr 232.0 231.0 238.0 237.0 (244) (243) (247) (247) (251) (252) (257) (258) (259) (260) Not valid for use as an USNCO National Examination after April 19, 2004. CHEMISTRY OLYMPIAD KEY FOR NATIONAL EXAM – PART II 1. g 0.1500 g Cu 1 mol × × = 1.099 M L 1.000 g solution 159.62 g mol × 0.025 L = 0.0275 mol CuSO 4 b. 1.099 L a. 1.1169 c. i. mol × 0.03063 L = 0.003173 mol HCl = 0.003173 mol NH 3 L g 0.003173 mol NH 3 × 17.034 = 0.05405 g NH 3 mol 0.1036 % NH 3 = ii. 0.05405 g NH 3 × 100 = 27.7% NH 3 195 g sample mol = 6.11 × 10−4 mol EDTA ≡ 6.11 × 10−4 mol Cu 2+ L g 6.11 × 10−4 mol Cu 2+ × 63.55 = 0.03881 g EDTA mol 0.03881 g EDTA % EDTA = × 100 = 25.9% EDTA 150 g sample 0.02443 L EDTA × 0.0250 0.200 g compound - 0.185 g anhydrous compound = 0.015 g H 2O 0.0150 g H 2O % H 2O = × 100 = 7.5% H 2O 0.200 g sample % SO 4 = 100 - [ 27.7 + 25.88 + 7.5] = 38.92 % iv. d. Assume 100 g; calculate moles using molar mass; divide all results by the smallest number 1 mol 27.7 g NH 3 × = 1.627 / 0.405 = 4.02 17.03 g iii. 1 mol = 0.407 / 0.405 = 1.00 63.55 g 1 mol 7.5 g H 2O × = 0.416 / 0.405 = 1.03 18.02 g 1 mol 38.9 g SO 4 × = 0.405 / 0.405 = 1 96.1 g 25.88 g Cu × Based on these results, we can see the molecular formula is, Cu(NH3)4 SO4.H2 O whose molar mass is 245.28 g/mol e. 245.28 g × 0.275 mol = 6.76 g is the theoretical yield mol % yield = 6.127 g × 100 = 90.7% 6.76 g 2. o o o o o a. ∆H rxn = ∆H f (N 2O 5 ) + ∆H f (O 2 ) - [2∆H f (NO 2 ) + ∆H f (O 3 )] −198 kJ = 11 kJ + 0 - [2∆H of (NO 2 ) + 143 kJ so, ∆H of (NO 2 ) = 33 kJ b. ∆So < 0 because 3 moles of gas are converted into 2 moles of gas. c. ∆Go = ∆Ho – T∆So = -198000 J – 298 K(-168 J/K) = -148 kJ Not valid for use as an USNCO National Examination after April 19, 2004. Page 5 d. Spontaneity will decrease as temperature increases because the entropy change of the process is negative. This makes the contribution of the T∆So term to the free energy positive. As temperature rises, this term contributes more to the free energy and because it is positive, the spontaneity must decrease. e. Looking at the table we see from comparing experiments 1&2 that when [NO 2 ] increases by a factor of 1.5 the rate of reaction also increases by 1.5, so it is first order in [NO2]. From experiments 2&3 we see that an increase in the [O 3 ] by a factor of 2.0 increases the rate of reaction by 2.0 so the reaction is also first order in [O 3 ]. f. The rate law is, rate = k[NO 2 ][O 3 ] 4.8 × 10−8 M −1s−1 = k (0.0015 M )(0.0025 M ) so k = 0.0128 L mol-1 s-1 g. Mechanism #1 would suggest a rate law of rate = k[NO 2 ]2 and mechanism #2 would suggest a rate law of, rate = k[NO 2 ] [O 3 ] . Neither of these mechanisms is consistent with the observed rate law. h. The rate of the reaction would be expected to increase with an increase in temperature. Collision rate increases as temperature increases and a larger fraction of collisions have the energy necessary for reaction (Ea). 2 3. [Ag + ]2 1.4 × 10−5 2= = 0.56 M a. K sp = [Ag + ]2[SO 24 ] so [SO 4 ] = 2 K sp (0.005) b. Cl– then Br– then I– . For Cl, 1.8 × 10−10 = [Ag + ](0.10), so [Ag + ] = 1.8 × 10−9 , for 5 × 10−13 = [Ag + ](0.10), so [Ag + ] = 5 × 10−13 and for 8.3 × 10−17 = [Ag + ](0.10), so [Ag + ] = 8.3 × 10−17 The [Ag+] in equilibrium with AgCl in 0.1 M Cl– is sufficient to cause precipitation of AgBr in 0.1 M Br– . In turn, the [Ag+] in equilibrium with AgBr in 0.1 M Br– will cause precipitation of AgI in 0.1 M I– . However, the reverse order of addition of the anions would not lead to this behavior. c. We need the net equilibrium constant for the combined reaction in each case. AgCl(s) → Ag + (aq) + Cl− (aq) K sp = 1.8 × 10−10 Ag + (aq) + 2NH 3 (aq) → Ag(NH 3 ) +2 K f = 1.7 × 10 7 So K = K sp × K f = 3.06 × 10−2 and AgBr(s) → Ag + (aq) + Br − (aq) K sp = 5 × 10−12 Ag + (aq) + 2NH 3 (aq) → Ag(NH 3 ) +2 K f = 1.7 × 10 7 So K = K sp × K f = 8.5 × 10−7 and AgI(s) → Ag + (aq) + I− (aq) K sp = 8.3 × 10−17 Ag + (aq) + 2NH 3 (aq) → Ag(NH 3 ) +2 K f = 1.7 × 10 7 So K = K sp × K f = 1.41 × 10−9 Now calculate the reaction quotient, Q, for the given case (using chloride as the example), AgCl(s) + 2NH 3 (aq) → Ag(NH 3 ) +2 + Cl− (aq) So, Q = [ [Cl– ] Ag(NH 3 ) 2+ [ NH 3 ] 2 ] = (0.10)(0.005) = 1.39 × 10-3 (0.60) 2 This value is smaller than only one of the calculated values for the net equilibrium constant (the case of chloride) so the silver chloride is the only one that will dissolve in 0.60 M NH3. Page 6 Not valid for use as an USNCO National Examination after April 19, 2004. 4. a. Al-O bonds are too strong to be broken by simple heating of the oxide, even in the presence of carbon. The heat of formation of Al 2 O3 is much more negative than that of CO2, so the reaction: 2Al2O 3 + 3C → 4 Al + 3CO 2 is endothermic. b. Na3 AlF6 is added to lower the melting point of Al2O3 . Lowering the melting point also lowers the amount of energy needed to carry out the process. 2– – 3+ – c. Al + 3e → Al (cathode) and 2O → O 2 + 4e 3 d. 5.00 × 10 g Al × 1 mol 3 mol e – = 185.3 mol Al and 185.3 mol Al × = 556 mol e – 26.98 g 1 mol Al e. We need to determine the charge (in coulomb) and time (in seconds) to calculate current. 96500 C 60 min 60 sec = 5.365 × 10 7 C and 10.0 hours × × = 3.6 × 10 4 sec – 1 hour 1 min 1 mol e (5.365 × 10 7 C) = 1490A so, (3.6 × 10 4 sec) 556 mol e – × f. 556 mol e – × 1 mol O 2 = 139 mol O 2 4 mol e – nRT (139 mol O 2 )(0.0821 L ⋅ atm ⋅ mol-1 ⋅ K -1)(298 K) V = = = 3590 L Now use Ideal Gas Law, 1 atm P 720 mmHg × 760 mmHg 5. a. b. c. d. H + + NO -3 + FeS → NO 2 + Fe 3+ + S + H 2O HC 2 H 3O 2 + Ca 3 (PO 4 ) 2 → Ca 2+ + C 2 H 3O –2 + HPO 2– 4 – – Al + H 2O + OH → Al(OH) 4 + H 2 2+ Cr 3+ + SO 2– + OH – → Cr(OH) 3 + BaSO 4 4 + Ba 2+ – 2– e. Co + Cl → CoCl4 f. Br2 + CH 3CHCH 2 → CH 3CHBrCH 2 Br 6. a. The Mg2+ ion is smaller and has a higher charge than the Mg+ ion, so the lattice energy that arises when Mg2+ ions form compounds is much greater than what would be observed if Mg+ ions formed compounds. The increase in lattice energy more than offsets the larger ionization energy of the Mg2+ ion. b. Ti (atomic number 24) has a valence electron configuration of 4s2 3d2 and can form +2 ions by losing it’s two 4s electrons, +3 ions by losing the two 4s electrons and one 3d electron and +4 by losing all four of the valence electrons. The +2 and +3 ions are colored because of electronic transitions between d orbitals. The +4 ion does not exhibit color because it has no valence d-electrons to undergo electronic transitions. c. Carbon dioxide (O=C=O) molecules are nonpolar and interact with each other only through weak dispersion forces. These weak forces are easily overcome so CO2 is a gas at room temperature. SiO2 doesn’t have the same molecular formula, because Si does not form double bonds as readily as carbon does. Si-O form single bonds that lead to a network solid held together with strong, covalent bonds, so it is a solid that has a high melting point. d. Nitrogen can form three bonds (NF 3 ) but not five (NF5 ) because it lacks d orbitals that are energetically available for the formation of hybrid orbitals (or alternatively, because it is too small to accommodate five atoms.) Both NF3 and PF3 are trigonal pyramidal because the central atom has three bonding pairs and one lone pair of electrons (leading to sp 3 hybridization). PF5 is trigonal bipyramidal because it has five bonding pairs (leading to dsp3 hybridization.) Not valid for use as an USNCO National Examination after April 19, 2004. Page 7 7. a. Even numbers of protons can pair up, which makes nuclei more stable. These more stable nuclei can accommodate a wider range of n/p ratios. b. Light nuclei are more stable when the n/p is close to 1/1. C-14 loses a β– because it’s n/p ratio of 8/6 is too high, so a neutron is converted to a proton, giving a n/p ratio of 7/7. C-10 loses a β+, thereby converting a proton to a neutron and changing its n/p ration from 4/6 to 5/5. c. The difference in radioactivity between a fresh sample of C-14 containing material and a sample between 500 years and 50,000 years old can be used to determine the age of the historical sample. Samples less than 500 years old produce too little difference from new samples to provide reliable dates. Samples older than 50,000 years contain too little C-14 to provide sufficient radioactivity (“counts” measured) to give reliable dates. 8. a. i. As the solution cools, the water freezes, leaving a more concentrated solution. The more concentrated the solution, the lower the freezing temperature, so the line slopes downward once the solution starts to freeze. ii. The depression of the freezing point is due to supercooling of the solution before crystallization begins. b. The freezing point is approximately –0.026 o C. ∆T = k f m so m = ∆T −0.026 = = 0.014 m kf 1.86 c. The molality is 0.014 mol / kg water, so the molar mass can be calculated, mol 0.029 g −4 = 207 g/mol 0.014 × (0.01 kg) = 1.4 × 10 mol and the molar mass is kg 1.4 × 10−4 mol d. The molar mass determination by freezing point depression is roughly twice as great as that determined by titration. Carboxylic acids tend to dimerize in solution so the freezing point depression experiment, which observes the number of particles rather than their chemical properties, overestimates the molar mass by almost a factor of two (dimers take two particles and combine them into one.) Page 8 Not valid for use as an USNCO National Examination after April 19, 2004. 2004 U.S. NATIONAL CHEMISTRY OLYMPIAD KEY FOR NATIONAL EXAM – PART III Problem 1. You have been given a vial containing either maleic acid, C4 H4 O4 , fumaric acid, C4H4 O4 , or tartaric acid, C4 H6 O6 . Your lab instructor will identify the acid you have been given. Design and carry out an experiment with the materials provided to determine the number of ionizable H+ ions possible for each molecule of the acid given. This is a solid acid titration experiment. Students had to determine the number of moles of NaOH base present given a 0.25M solution, then perform a titration against a weighed sample of the solid maleic, fumaric, or tartaric acid. The endpoint is detected using phenolphthalein as the indicator. All of these solid organic acids have two ionizable protons. Experimental Plan: An experiment plan had to describe the steps needed to carry out a successful, microscale titration. Data and Observations: The recording of data and observations needed to include information about the mass of the acid used and the volume of the base used. Where the quantity was obtained by difference, both observations used to determine the amount should be shown. Significant figures had to be used appropriately. Sample Calculations and Conclusions: Students needed to provide calculations that supported their conclusions. An example of excellent student work: Plan: Dissolve a pre-weighed acid sample in distilled water adding two drops phenolphthalein. Fill the 10.0 ml graduated cylinder with .25M NaOH. Add base gradually until a permanent pink color. Record total volume of base used. Repeat. Data: Acid used: Maleic MM=116.0 g/mol Mass of vial + acid Mass of vial less acid Mass of acid Volume base Key for 2004 USNCO National Exam, Part III Trial One 4.50g 4.00 g .50g 34.4 ml Trial Two 4.00 g 3.48g .52g 35.0 ml Calculations: Moles acid = mass acid/MM Moles base= (liters base x Mbase) Mole base/moles acid .0043 .0086 2.0 .0045 .0088 2.0 Conclusion: For maleic acid, there are two moles of ionizable protons/mol acid. It is diprotic. An example of good student work: Plan: Dissolve maleic acid in water. Add phenolphthalein and titrate with .25M NaOH. Calculate. Data: 1.462g maleic in 125 ml solution. Molarity solution = .101 Use 25 ml of solution and titrate. 19.5 ml NaOH used. Calculations: Mol Maleic = .0025 Mol NaOH = 19.5ml/1000ml/L(.25mol/L) = .004875 Mol NaOH/Mol Maleic = 2 Conclusion: Maleic acid is diprotic. Key for 2004 USNCO National Exam, Part III (Work not shown) (Poor use sig. fig) (Only one trial) Problem 2: You have been given 4 (four) black pens. Design and carry out an experiment to determine whether the dye used in each pen is a compound or a mixture. You will need to provide evidence to justify your conclusions. This is a paper-chromatography identification experiment. Ammonia and water are provided as the carrying liquids. Students had to create an experiment using pieces of filter paper to observe a possible separation of the black ink dyes from each of the four pens provided. Situating the filter paper on or in provided beakers allowed the solvent front to rise, showing a possible separation. Experimental Plan The experimental plan needed to identify a way to use paper chromatography to investigate the inks in the various pens. Some detail about how to carry out the experiment – marking the filter paper above the liquid level, for example, was useful in this component of the exercise. Observations and results: Students needed to summarize observations about the chromatography experiment. Students who did an excellent job carried out more than one trial to verify results and included detail about separations. A table of probable observations includes, Pens #1 Crayola ® #2 Gel-Pen ® #3 Papermate ® #4 Sharpie ® Water Separation No separation or movement Separation No separation Ammonia Separation No separation Conclusions = a mixture = a compound Separation No separation = a mixture = a compound An example of excellent student work: Using Ammonia as solvent: - Pen 1: Different colors appear: orange, blue, green and purple - Pen 2: Solid black color only appears. - Pen 2 (again): Still shows only solid black color. - Pen 3: Different colors appear, but they are light: green, blue and purple - Pen 3 (again): Different colors appear again, darker than previous run. - Pen 4: Mark does not move, only black color appears. - Pen 4 (again): Mark does not move, only black color. Using water as solvent – retry #2 and #4. - Pen 2: Solid black color only appears - Pen 4: No movement observed. - Pen 4 (again): No movement observed. Conclusions: Pen 1 is a mixture because different colored pigments are observed. Pen 2 is a compound because only one color is observed even though the ink moved along the filter paper. Pen 3 is a mixture because different colored pigments are observed. The pigments used are different than those used in Pen 1. Key for 2004 USNCO National Exam, Part III Pen 4 can not be determined in this experiment because no solvent was found that dissolved the ink. Based on the fact that separation of color is not observed, the ink might be a single compound. An example of good student work: Pen 1: Mulitple layers form. Fastest moving layer is light blue green. A darker blue-green is next and a purple color is the slowest. Pen 2: One dark black layer that doesn’t seem to move at all. One light black layer that moves slowly. One gray layer that moves faster. Pen 3: One black layer doesn’t move much. One light black layer that moves slowly. One blue-purple layer that doesn’t move at all. Pen 4: Black mark doesn’t move at all and remains dark black. Conclusion: The pens with multiple layers in the paper chromatography were mixtures – in this case pens 1, 2 and 3. The pen with a single layer in the paper chromatography, pen 4, was a single compound. Key for 2004 USNCO National Exam, Part III 2005 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM PART 1 Prepared by the American Chemical Society Olympiad Examinations Task Force OLYMPIAD EXAMINATIONS TASK FORCE Arden P. Zipp, State University of New York, Cortland Chair Sherry Berman-Robinson, Consolidated High School, IL Alice Johnsen, Bellaire High School, TX William Bond, Snohomish High School, WA Adele Mouakad, St. John’s School, PR Peter E. Demmin (retired), Amherst Central High School, NY Kimberley Gardner, United States Air Force Academy, CO, David W. Hostage, Taft School, CT Jane Nagurney, Scranton Preparatory School, PA Ronald O. Ragsdale, University of Utah, UT Jacqueline Simms, Sandalwood Sr. High School, FL DIRECTIONS TO THE EXAMINER–PART I Part I of this test is designed to be taken with a Scantron® answer sheet on which the student records his or her responses. Only this Scantron sheet is graded for a score on Part I. Testing materials, scratch paper, and the Scantron sheet should be made available to the student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until April 27, 2005, after which tests can be returned to students and their teachers for further study. Allow time for the student to read the directions, ask questions, and fill in the requested information on the Scantron sheet. The answer sheet must be completed using a pencil, not pen. When the student has completed Part I, or after one hour and thirty minutes has elapsed, the student must turn in the Scantron sheet, Part I of the testing materials, and all scratch paper. There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and you are free to schedule rest-breaks between parts. Part I Part II Part III 60 questions 8 questions 2 lab problems single-answer multiple-choice problem-solving, explanations laboratory practical 1 hour, 30 minutes 1 hour, 45 minutes 1 hour, 30 minutes A periodic table and other useful information are provided on page 2 for student reference. Students should be permitted to use nonprogrammable calculators. DIRECTIONS TO THE EXAMINEE–PART I DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Answers to questions in Part I must be entered on a Scantron answer sheet to be scored. Be sure to write your name on the answer sheet; an ID number is already entered for you. Make a record of this ID number because you will use the same number on both Parts II and III. Each item in Part I consists of a question or an incomplete statement that is followed by four possible choices. Select the single choice that best answers the question or completes the statement. Then use a pencil to blacken the space on your answer sheet next to the same letter as your choice. You may write on the examination, but the test booklet will not be used for grading. Scores are based on the number of correct responses. When you complete Part I (or at the end of one hour and 30 minutes), you must turn in all testing materials, scratch paper, and your Scantron answer sheet. Do not forget to turn in your U.S. citizenship statement before leaving the testing site today. Not valid for use as an USNCO Olympiad National Exam after April 26, 2005. Distributed by the ACS DivCHED Examinations Institute, University of Wisconsin - Milwaukee, Milwaukee, WI. All rights reserved. Printed in U.S.A. ABBREVIATIONS AND SYMBOLS A Faraday constant F molal atm formula molar mass M molar u free energy G molar mass A frequency ν mole N A gas constant R Planck’s constant °C gram g pressure c heat capacity C p rate constant C hour h retention factor E joule J second Ea kelvin K temperature, K H kilo– prefix k time S liter L volt K milli– prefix m ampere atmosphere atomic mass unit atomic molar mass Avogadro constant Celsius temperature centi– prefix coulomb electromotive force energy of activation enthalpy entropy equilibrium constant CONSTANTS m M M mol h P k Rf s T t V R = 8.314 J·mol –1·K–1 R = 0.0821 L·atm·mol –1·K–1 1 F = 96,500 C·mol–1 1 F = 96,500 J·V–1·mol–1 N A = 6.022 × 10 23 mol–1 h = 6.626 × 10 –34 J·s c = 2.998 × 10 8 m·s –1 0 °C = 273.15 K 1 atm = 760 mmHg EQUATIONS E = Eo − 1 1A 1 H 1.008 3 Li RT ln Q nF k E 1 1 ln 2 = a − k1 R T1 T2 −ΔH 1 ln K = + constant R T PERIODIC TABLE OF THE ELEMENTS 2 2A 4 Be 13 3A 5 B 14 4A 6 C 15 5A 7 N 16 6A 8 O 17 7A 9 F 18 8A 2 He 4.003 10 Ne 6.941 9.012 10.81 12.01 14.01 16.00 19.00 20.18 11 Na 12 Mg 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar 22.99 24.31 19 K 20 Ca 3 3B 21 Sc 4 4B 22 Ti 5 5B 23 V 6 6B 24 Cr 7 7B 25 Mn 8 8B 26 Fe 9 8B 27 Co 10 8B 28 Ni 11 1B 29 Cu 12 2B 30 Zn 58.69 63.55 65.39 69.72 72.61 74.92 78.96 79.90 83.80 46 Pd 47 Ag 48 Cd 49 In 50 Sn 51 Sb 52 Te 53 I 54 Xe 26.98 28.09 30.97 32.07 35.45 39.95 31 Ga 32 Ge 33 As 34 Se 35 Br 36 Kr 39.10 40.08 44.96 47.88 50.94 52.00 54.94 55.85 58.93 37 Rb 38 Sr 39 Y 40 Zr 41 Nb 42 Mo 43 Tc (98) 44 Ru 101.1 45 Rh 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.3 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 85.47 87.62 88.91 91.22 92.91 95.94 55 Cs 56 Ba 57 La 72 Hf 73 Ta 74 W 132.9 137.3 138.9 178.5 180.9 183.8 186.2 190.2 192.2 195.1 197.0 200.6 87 Fr 88 Ra 89 Ac 104 Rf 105 Db 106 Sg 107 Bh 108 Hs 109 Mt 110 111 112 114 (269) (272) (277) (2??) (223) (226) (227) € 58 Ce 59 Pr (262) 60 Nd (263) 61 Pm (262) 62 Sm (265) 63 Eu (266) 64 Gd 65 Tb 66 Dy 67 Ho 173.0 175.0 101 Md 102 No 103 Lr 144.2 (145) 150.4 152.0 157.3 158.9 162.5 164.9 167.3 91 Pa 92 U 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm (237) (244) (243) (247) (247) (251) (252) (257) (258) 70 Yb (209) 168.9 140.9 238.0 209.0 69 Tm 90 Th 231.0 207.2 68 Er 140.1 232.0 Page 2 (261) 204.4 (259) (210) (222) 71 Lu (262) Not valid as a USNCO National Exam after April 26, 2005 DIRECTIONS When you have selected your answer to each question, blacken the corresponding space on the answer sheet using a soft, #2 pencil. Make a heavy, full mark, but no stray marks. If you decide to change an answer, erase the unwanted mark very carefully. There is only one correct answer to each question. Any questions for which more than one response has been blackened will not be counted. Your score is based solely on the number of questions you answer correctly. It is to your advantage to answer every question. 1. Which solution produces a black precipitate when added to an aqueous copper(II) solution? (A) NH3 (B) (NH4)2S (C) K2SO4 (D) NaOH 6. Which diagram best represents the change in electrical conductivity of a solution of acetic acid as a solution of sodium hydroxide is added? (A) (B) (C) (D) 2. Which oxide is the best reducing agent? (A) CO2 (B) NO2 (C) SiO 2 (D) SO2 3. Solutions of which ion produce a red color when vaporized in a Bunsen burner flame? (A) calcium (B) potassium (C) sodium (D) zinc 4. Which procedure for dispensing a liquid with a volumetric pipet is correct? (A) Draw the liquid up to the line on the pipet using a pipet bulb. Squeeze the bulb to force all the liquid in the pipet into the receiving container. (B) Introduce the liquid into the top end of the pipet until it is filled to the line. Allow the liquid to drain into the desired container. Blow on the pipet to release the last drop. 7. Methylamine, CH3NH2, reacts with O2 to form CO 2, N2, and H2O. What amount of O2 (in moles) is required to react completely with 1.00 mol of CH3NH2? (A) 2.25 (B) 2.50 (C) 3.00 (D) 4.50 (C) Draw the liquid above the line on the pipet using a pipet bulb. With a finger on the top of the pipet allow the curve of the meniscus to drop to the line. Place the tip of the pipet against the side of the receiving container and allow the liquid to drain. 8. Iodine adds to the double bonds in fatty acids (one iodine molecule per double bond). How many double bonds are in a molecule of arachidonic acid (Molar mass = 304.5 g/mol) if 0.125 g of the acid require 0.417 g of iodine? (D) Draw the liquid above the line on the pipet by sucking on the open end of the pipet. Place a thumb on the top of the pipet and allow the curve of the meniscus to drop to the line. Allow the liquid to drain into the receiving container pipet against its side. 9. The solubility of a gas in a I. pressure of the gas liquid increases when II. temperature of the liquid which of the following increases? 5. Which physical characteristic distinguishes copper from brass (an alloy of copper and zinc)? (A) Brass is a liquid at room temperature and copper is not. (A) 2 (B) 3 (C) 4 (D) 8 (A) I only (B) II only (C) both I and II (D) neither I nor II 10. A mineral containing only manganese and oxygen contains 69.6% Mn by mass. What is its empirical formula? (B) Brass is much less dense than copper. (A) MnO (B) Mn2O3 (C) Brass is attracted to a magnet but copper is not. (C) Mn3O4 (D) MnO2 (D) Brass is a much poorer electrical conductor than copper. Not valid as a USNCO National Exam after April 26, 2005 Page 3 11. Toluene, C7H8, is added to gasoline to increase its octane rating. What is the volume ratio of air to toluene vapor to burn completely to form CO2 and H 2O? (Assume air is 20% O2 by volume.) (A) 9/1 (B) 11/1 (C) 28/1 (D) 45/1 12. Acidified solutions of dichromate ion, Cr2O72-, oxidize Fe2+ to Fe3+, forming Cr3+ in the process. What volume of 0.175 M K 2Cr2O7 in mL is required to oxidize 60.0 mL of 0.250 M FeSO4? (A) 14.3 (B) 28.6 (C) 42.9 (D) 85.7 13. Which property is the same for 1.0 g samples of H2 and CH4 in separate 1.0 L containers at 25 ˚C? (A) pressure (B) number of molecules (C) average molecular velocity (D) average molecular kinetic energy 14. When CsI, SiO2, CH3OH and C 3H8 are listed in order of increasing melting point, which is the correct order? (A) CsI, SiO2, CH3OH, C3H8 (B) CH3OH, C3H8, CsI, SiO2 17. When NaF, MgO, KCl and CaS are listed in order of increasing lattice energy, which order is correct? (A) MgO, NaF, KCl, CaS (B) CaS, MgO, KCl, NaF (C) KCl, CaS, NaF, MgO (D) KCl, NaF, CaS, MgO 18. When compared to most I. boiling point other substances of similar II. specific heat capacity molar mass the values of III. surface tension which properties of liquid H2O are unusually large? (A) I only (B) I and II only (C) II and III only (D) I, II and III 19. Calculate ∆H˚ for the reaction; ∆Hf˚ kJ/mol TiCl4(g) + 2H 2O(l) TiCl4(g) –763 r TiO 2(s) + 4HCl(g) H2O(l) –286 TiO2(s) –945 HCl(g) –92 (A) –264 kJ (B) 12 kJ (C) 22 kJ (D) 298 kJ 20. Use bond energies to estimate the value of ∆H˚ for the reaction; N2(g) + 3H 2(g) r 2NH 3(g) (C) CH3OH, C3H8, SiO 2, CsI (D) C 3H8, CH3OH, CsI, SiO2 15. According to the graph (ln vapor pressure vs 1/T) what can be concluded about the enthalpies of vaporization (∆Hvap ) of liquids X and Y? Bond Energies kJ/mol H-H 436 H-N 386 N-N 193 N=N 418 N≡N 941 (A) –995 kJ (B) –590 kJ (C) –67 kJ (D) 815 kJ Questions 21. and 22. should be answered using this thermochemical equation; N2(g) + 2O2(g) r 2NO2(g) ∆Hrxn > 0 21. Which relationship is correct for this reaction at a pressure of 1 atm? (A) ∆Hvap X > ∆Hvap Y (B) ∆Hvap X = ∆Hvap Y (C) ∆Hvap X < ∆Hvap Y (D) No conclusions can be drawn about the relative ∆Hvap values from this diagram. 16. An unknown gas effuses through a pin-hole in a container at a rate of 7.2 mmol/s. Under the same conditions gaseous oxygen effuses at a rate of 5.1 mmol/s. What is the molar mass (in g/mol) of the unknown gas? (A) 16 Page 4 (B) 23 (C) 45 (D) 64 (A) ∆Erxn > ∆Hrxn (B) ∆Erxn < ∆Hrxn (C) ∆Erxn = ∆Hrxn + ∆S rxn (D) ∆Erxn = ∆Hrxn – ∆S rxn 22. Under what temperature conditions is this reaction spontaneous at standard pressure? (A) at low temperatures only (B) at high temperatures only (C) at all temperatures (D) at no temperature Not valid as a USNCO National Exam after April 26, 2005 23. Diethyl ether has a normal boiling point of 35.0 ˚C and has an entropy of vaporization of 84.4 J/mol. K. What is its enthalpy of vaporization? (A) 0.274 J/mol (B) 2.41 J/mol (C) 3.65 J/mol (D) 26.0 kJ/mol 24. A 9.40 g sample of Solution Properties KBr is dissolved in 105 Molar mass KBr 119 g/mol g of H 2O at 23.6 ˚C in a ∆Hsoln KBr 19.9 kJ/mol coffee cup. Find the C p solution 4.184 J/g˚C final temperature of this system. Assume that no heat is transferred to the cup or the surroundings. (A) 20.0 ˚C (B) 20.3 ˚C (C) 26.9 ˚C (D) 27.2 ˚C 25. For the reaction A r B which is first order in A, which of the following change as the concentration of A changes? I. II. III. rate rate constant Half–life (A) I only (B) III only (C) II and III only (D) I, II and III 26. The equation and rate law for the gas phase reaction between NO and H2 are; 2NO(g) + 2H2(g) r N 2(g) + 2H 2O(g) Rate = k[NO]2[H 2] What are the units of k if time is in seconds and the concentration is in moles per liter? (A) L. s. mol -1 (B) L2. mol -2. s-1 (C) mol . L-1. s-1 (D) mol 2. L-2. s-1 27. At a given temperature a first-order reaction has a rate constant of 3.33×10-3 s-1. How much time is required for the reaction to be 75% complete? (A) 100 s (B) 210 s (C) 420 s (D) 630 s 28. Most reactions occur more I. activation energy rapidly at high temperatures II. collision energy than at low temperatures. This III. rate constant is consistent with an increase in which property at higher temperatures? (A) I only (B) II only (C) I and III only (D) II and III only 29. Which graph is diagnostic of an irreversible second order reaction A r B? (A) (B) (C) (D) 30. The reaction; 2NO(g) + 2H2(g) r 2H2O(g) + N2(g) obeys the rate equation Rate = k[NO]2[H 2] This mechanism has been proposed: 1. 2NO(g) r N2O2(g) 2. N2O2(g) + H 2(g) r 2HON(g) 3. HON(g) + H2(g) r H 2O(g) + HN(g) 4. HN(g) + HON(g) r N2(g) + H 2O(g) Which step of the mechanism is the rate-determining step? (A) step 1 (B) step 2 (C) step 3 (D) step 4 31. For the hypothetical equilibrium reactions; AsB K = 2.0 BsC K = 0.010 What is the value of K for the reaction; 2C s 2A? (A) 2500 32. For which reaction is Kp = Kc ? (C) 25 (D) 4.0×10-4 2N2(g) + O 2(g) s 2N2O(g) I. II. C(s) + O2(g) s CO2(g) III. N2O4(g) s 2NO2(g) (A) II only (B) III only (C) I and III only (D) II and III only 33. What is the pH of a 0.010 M solution of a weak acid HA that is 4.0% ionized? (A) 0.60 Not valid as a USNCO National Exam after April 26, 2005 (B) 50 (B) 0.80 (C) 2.80 (D) 3.40 Page 5 34. Given the acid ionization constants, when the conjugate bases are arranged in order of increasing base strength, which order is correct? – Acid Ionization Constant, Ka HClO 3.5×10-8 HClO2 1.2×10-2 HCN 6.2×10-10 – H2PO4 6.2×10-8 – – – – (C) CN , HPO42–, ClO , ClO 2 – – – (D) CN , ClO , HPO42–, ClO 2 Base Ionization Constant, Kb 35. Calculate the NH3 1.8×10-5 concentration of hydrogen ion in mol/L of a 0.010 M solution of NH4Cl. (A) 4.2×10-4 (B) 2.4×10-6 (C) 1.8×10-7 (D) 5.6×10-12 36. For the reaction; – PbI2(s) s Pb2+(aq) + 2I (aq) Ksp = 8.4×10-9 2+ What is the concentration of– Pb in mol/L in a saturated solution of PbI2 in which [I ] = 0.01 M? (A) 8.4×10 -7 (B) 8.4×10 (C) 1.3×10 -3 (D) 2.0×10-3 -5 37. Which statement is correct about the electrochemical cell – represented here? Ag | Ag + || NO3 , NO | Pt (C) 1033 (D) 1078 (B) The major purpose of the Pt is to act as a catalyst. (C) The Ag electrode decreases in mass as the cell operates. (A) I only (B) II only (C) both I and II (D) neither I nor II (A) 0.355 V (B) 0.178 V (C) –0.178 V (D) –0.355 V (C) II and III only (D) II and IV only (B) 0.22 (C) 0.33 (D) 0.66 43. How many orbitals are in an atomic sublevel with l = 3? (A) 3 (B) 5 (C) 7 (D) 9 44. A ground state gaseous atom of which element has the greatest number of unpaired electrons? (A) As (B) Br (C) Ge (D) Se 45. An atom of which element has the highest second ionization energy? (A) Na (B) Mg 46. Which of these properties increase across the period from Na to Cl? (C) Al I. II. III. (D) K atomic radius density electronegativity (A) I only (B) III only (C) I and II only (D) II and III only 47. For the elements in group 14 (C to Pb), which property increases with increasing atomic number? (A) melting points (B) covalent radius (C) magnitude of stable oxidation state (D) ability to form chains of atoms with themselves 48. What mode of radioactive decay is most likely for the isotope 20 11 Na ? + 39. The standard reduction potential for H (aq) is 0.00 V. What is the reduction potential for a 1×10-3 M HCl solution? (B) I and IV only (A) 0.16 (D) The voltage of the cell can be increased by doubling the size of the Ag electrode. 38. The overall reaction for the lead storage battery when it discharges is; + 2Pb(s) + PbO 2(s) + 4H (aq) + 2SO4 (aq) r 2PbSO 4(s) + 2H2O(l) I. PbSO4 is formed only at the cathode. II. The density of the solution decreases. Which statement(s) correctly describe(s) the battery as it discharges? (A) I and III only 42. A current of 0.20 amps is passed through an aqueous solution of nickel(II) nitrate for 45.0 minutes. What mass of Ni metal (in grams) will be deposited? (A) NO undergoes oxidation at the anode. Page 6 (B) 1026 41. Which products are formed by the electrolysis of an aqueous solution of AlCl3? I. Al(s) II. Cl2(g) III. H2(g) IV. O2(g) – (B) ClO 2 , HPO42–, ClO , CN – (A) 1022 – (A) ClO 2 , ClO , HPO42–, CN – Standard Reduction Potential, V 40. What is the – approximate value Ag+(aq) + e r Ag(s) +0.80 – of the equilibrium Cr3+(aq) + 3e r Cr(s) –0.74 constant, Keq , at 25 oC for the reaction; 3Ag+(aq) + Cr(s) r Cr3+(aq) + 3Ag(s) € (A) alpha (B) beta (C) gamma (D) electron capture Not valid as a USNCO National Exam after April 26, 2005 49. Oxygen gas is paramagnetic. This observation is best explained by 56. What is the most characteristic reaction of benzene? (A) resonance. (B) the Lewis structure of O2. (D) the hybridization of atomic orbitals in O2. 50. What is the geometry of the iodine atoms in the I3- ion? (A) bent (B) linear (C) T-shaped (D) triangular O S O O (A) 0, 0 (B) –2, 0 (C) +2, –1 (B) 2 (B) CH3COOH (C) ClCH2COOH (D) ClCH2CH2COOH (B) two (C) three (D) four (A) ClHC = CHCl (B) meta-C6H4Cl2 (C) CH2ClBr (D) CH3CH(Cl)CH2CH3 60. Which type of dietary fat is currently considered the least harmful? (A) monounsaturated fat (B) polyunsaturated fat (C) saturated fat (D) trans fat (D) +6, –2 53. How many different isomers exist for the octahedral + complex [Co(NH3)4Cl2] ? (A) 1 (A) HCOOH 59. Which compound can exist as two optical isomers? (D) SF6 2– O (D) substitution (A) one (C) SbF5 52. In the Lewis structure what are the formal charges on the sulfur and oxygen atoms, respectively? (C) reduction 58. How many structurally isomeric alcohols have the formula C4H9OH? 51. Which species has a dipole moment other than zero? (B) CF4 (B) polymerization 57. Which organic acid is the strongest? (C) the molecular orbital description of O2. (A) BrF 3 (A) addition (C) 3 END OF TEST (D) 4 54. Which order is correct when the species are arranged in order of increasing average N-O bond length? – – + (B) NO , NO3 , NO2 – – + (D) NO , NO2 , NO3 (A) NO3 , NO2 , NO (C) NO2 , NO3 , NO + – – + – – 55. All of the classes of compounds contain at least one oxygen atom EXCEPT (A) esters (B) aldehydes (C) ethers (D) alkynes Not valid as a USNCO National Exam after April 26, 2005 Page 7 NATIONAL OLYMPIAD PART I 2005 KEY Number 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. Answer B D A C D D A C A B D A D D A A D D C C A D D B A B C D C B Number 31. 32. 33. 34. 35. 36. 37. 38. 39. 40. 41. 42. 43. 44. 45. 46. 47. 48. 49. 50. 51. 52. 53. 54. 55. 56. 57. 58. 59. 60. Property of the ACS DivCHED Examinations Institute Answer A A D B B B C B C D C A C A A B B D C B A C B D D D C D D B 2005 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART II Prepared by the American Chemical Society Olympiad Examinations Task Force OLYMPIAD EXAMINATIONS TASK FORCE Arden P. Zipp, State University of New York, Cortland Chair Sherry Berman-Robinson, Consolidated High School, IL Alice Johnsen, Bellaire High School, TX William Bond, Snohomish High School, WA Adele Mouakad, St. John’s School, PR Peter E. Demmin (retired), Amherst Central High School, NY Jane Nagurney, Scranton Preparatory School, PA Kimberley Gardner, United States Air Force Academy, CO, David W. Hostage, Taft School, CT Ronald O. Ragsdale, University of Utah, UT Jacqueline Simms, Sandalwood Sr. High School, FL DIRECTIONS TO THE EXAMINER–PART II Part II of this test requires that student answers be written in a response booklet of blank pages. Only this “Blue Book” is graded for a score on Part II. Testing materials, scratch paper, and the “Blue Book” should be made available to the student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until April 27, 2005, after which tests can be returned to students and their teachers for further study. Allow time for the student to read the directions, ask questions, and fill in the requested information on the “Blue Book”. When the student has completed Part II, or after one hour and forty-five minutes has elapsed, the student must turn in the “Blue Book”, Part II of the testing materials, and all scratch paper. Be sure that the student has supplied all of the information requested on the front of the “Blue Book,” and that the same identification number used for Part I has been used again for Part II. There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and you are free to schedule rest-breaks between parts. Part I Part II Part III 60 questions 8 questions 2 lab problems single-answer multiple-choice problem-solving, explanations laboratory practical 1 hour, 30 minutes 1 hour, 45 minutes 1 hour, 30 minutes A periodic table and other useful information are provided on the back page for student reference. Students should be permitted to use non-programmable calculators. DIRECTIONS TO THE EXAMINEE–PART II DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Part II requires complete responses to questions involving problem-solving and explanations. One hour and forty-five minutes are allowed to complete this part. Be sure to print your name, the name of your school, and your identification number in the spaces provided on the “Blue Book” cover. (Be sure to use the same identification number that was coded onto your Scantron® sheet for Part I.) Answer all of the questions in order, and use both sides of the paper. Do not remove the staple. Use separate sheets for scratch paper and do not attach your scratch paper to this examination. When you complete Part II (or at the end of one hour and forty-five minutes), you must turn in all testing materials, scratch paper, and your “Blue Book.” Do not forget to turn in your U.S. citizenship statement before leaving the testing site today. Not valid for use as an USNCO Olympiad National Exam after April 26, 2005. Page 1 1. (15%) This question involves several reactions of copper and its compounds. a. A sample of copper metal is dissolved in 6 M nitric acid contained in a round bottom flask. This reaction yields a blue solution and emits a colorless gas which is found to be nitric oxide. Write a balanced equation for this reaction. b. The water is evaporated from the blue solution to leave a blue solid. When the blue solid is heated further, a second reaction occurs. This reaction produces a mixture of nitrogen dioxide gas, oxygen gas and a black oxide of copper. i. A sample of the dried gas, collected in a 125 mL flask at 35 ˚C and 725 mm Hg, weighed 0.205 g. Find the average molar mass of the gas and the molar NO2/O2 ratio in it. ii. These data were obtained for the black solid; Mass of empty flask: 39.49 g Mass of flask + copper metal: 40.86 g Mass of flask + oxide of copper: 41.21 g Determine the formula of the oxide of copper. c. If some of the blue solution were lost due to splattering during the evaporation, what would be the effect on the calculated percentage of copper in the black oxide? Explain. d. If all of the blue solid were not decomposed into the black oxide during the final heating, what would be the effect on the calculated percentage of copper in the oxide? Explain. 2. 3. (16%) Liquid hydrazine, N2H4, is sometimes used as a rocket propellant. a. Write an equation for the formation of hydrazine from its elements and use the combustion equations below to derive an equation in which ∆Hf˚ for liquid hydrazine is expressed in terms of ∆H1, ∆H2 and ∆H3. 1 N (g) + O (g) r NO (g) ∆H1 2 2 2 2 H2(g) + 1 2 O2(g) r H 2O(g) ∆H2 N2H4(l) + 3O2(g) r 2NO2(g) + 2H 2O(g) ∆H3 b. In€a rocket, liquid hydrazine is reacted with liquid hydrogen peroxide to produce nitrogen and water vapor. Write a balanced equation for this reaction. € c. Calculate ∆Hrxn˚ for the reaction represented in 2b. ∆Hf˚ kJ/mol N2H4(l) 50.6 H2O2(l) –187.8 H2O(g) –285.8 d. Calculate ∆Hrxn˚ for the reaction in 2b using bond dissociation energies. Bond Dissociation Energy kJ/mol N–N 167 O–O 142 N=N 418 O=O 494 N≡N 942 O–H 459 N-H 386 e. Which value of ∆Hrxn˚ (that calculated in part c or part d) is likely to be more accurate? Justify your answer. f. Calculate the maximum temperature of the combustion gases if all the energy generated in the reaction goes into raising the temperature of those gases. [The heat capacities of N2(g) and H2O(g) are 29.1 J/(mol . ˚C) and 33.6 J/(mol. ˚C), respectively.] (13%) A solution of alanine hydrochloride, [H3NCH(CH3)COOH]+Cl-, is titrated with a solution of sodium hydroxide to produce a curve similar to the one shown. a. Give the formula(s) for the major species present at the points on the titration curve. i. 1 ii. 2 iii. 3 iv. 4 b. If K1 and K 2 of alanine hydrochloride are 4.6×10–3 and 2.0×10–10 respectively, i. write equations to represent the reactions responsible for K1 and K 2. ii. determine the pH at points 1, 2 and 3. c. Describe quantitatively how you could prepare a buffer solution with a pH = 10.0. Solutions of 0.10 M alanine hydrochloride and 0.10 M NaOH are available. Not valid for use as an USNCO Olympiad National Exam after April 26, 2005. 4. (10%) An electrochemical cell based on the reaction; M(s) + Cu2+(aq) r M2+(aq) + Cu(s) E˚ = 1.52 V is constructed using equal volumes of solutions with all substances in their standard states. a. Use the value of the reduction potential of Cu2+(aq) (E˚ = 0.34 V) to determine the standard reduction potential for the – reaction; M2+(aq) + 2e r M(s) b. The cell is allowed to discharge until the [Cu2+] = 0.10 M. Find i. the M 2+ concentration in moles per liter, ii. the cell potential, E. c. 50 mL of distilled water is added to each cell compartment of the original cell. Compare the potential of the cell after the addition of water with the potential of the original cell. Explain your answer. 5. (12%) Write net equations for each of the combinations of reactants below. Use appropriate ionic and molecular formulas and omit formulas for all ions or molecules that do not take part in a reaction. Write structural formulas for all organic substances. You need not balance the equations. All reactions occur in aqueous solution unless otherwise indicated. a. Solid sodium sulfide is added to water. b. An aqueous solution of potassium triiodide is added to a solution of sodium thiosulfate. c. Excess aqueous sodium fluoride is added to aqueous iron(III) nitrate. d. Strontium-90 undergoes beta decay. e. Excess carbon dioxide is bubbled through a solution of calcium hydroxide. f. Ethanol is warmed gently with acidified potassium dichromate. 6. (10%) This mechanism has been proposed for the reaction between chloroform and chlorine. Step 1: Cl2(g) s 2Cl(g) fast Step 2: CHCl3(g) + Cl(g) s CCl 3(g) + HCl(g) slow Step 3: CCl3(g) + Cl(g) s CCl 4(g) fast a. Write the stoichiometric equation for the overall reaction. b. Identify any reaction intermediates in this mechanism. c. Write the rate equation for the rate determining step. d. Show how the rate equation in c. can be used to obtain the rate law for the overall reaction. e. If the concentrations of the reactants are doubled, by what ratio does the reaction rate change? Explain. 7. (14%) Xenon forms several compounds including XeF2, XeF4 and XeO3. a. Draw a Lewis structure for each of these molecules. b. Describe the geometry of each compound including bond angles. c. State and explain whether each molecule is polar or nonpolar. d. Account for the observation that these compounds are highly reactive. 8. (10%) A section of the polymer polypropylene is represented here. a. Sketch a structural formula for the monomer used to make this polymer. b. State and explain whether this polymer is an addition or a condensation polymer. c. Compare the melting points for the following polymers. Give reasons for your answers. i. polypropylene containing 1000 monomer units vs polypropylene containing 10,000 monomer units ii. polypropylene vs a polymer in which the CH3 group is replaced with a CH2CH2CH2CH3 group iii. isotactic polypropylene (all the CH 3 groups on the same side of the carbon backbone) vs atactic polypropylene (CH 3 groups arranged at random) END OF PART II Not valid for use as an USNCO Olympiad National Exam after April 26, 2005. Page 3 amount of substance ampere atmosphere atomic mass unit atomic molar mass Avogadro constant Celsius temperature centi- prefix coulomb electromotive force energy of activation enthalpy entropy ABBREVIATIONS AND SYMBOLS n equilibrium constant K measure of pressure mmHg A Faraday constant F milli- prefix m atm formula molar mass M molal m u free energy G molar M A frequency ν mole mol N A gas constant R Planck’s constant h °C gram g pressure P c heat capacity C p rate constant k C hour h retention factor Rf E joule J second s Ea kelvin K speed of light c H kilo- prefix k temperature, K T S liter L time t volt V E = Eο – USEFUL EQUATIONS –ΔH 1 ln K = +c R T RT ln Q nF CONSTANTS R = 8.314 J·mol –1·K–1 R = 0.0821 L·atm·mol –1·K–1 1 F = 96,500 C·mol–1 1 F = 96,500 J·V–1·mol–1 N A = 6.022 × 10 23 mol–1 h = 6.626 × 10 –34 J·s c = 2.998 × 10 8 m·s –1 k2 Ea 1 1 = − k1 R T1 T2 ln PERIODIC TABLE OF THE ELEMENTS 1 H 2 He 1.008 3 Li 4.003 4 Be 5 B 6 C 7 N 8 O 9 F 10 Ne 6.941 9.012 10.81 12.01 14.01 16.00 19.00 20.18 11 Na 12 Mg 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar 22.99 24.31 19 K 20 Ca 21 Sc 22 Ti 23 V 24 Cr 25 Mn 26 Fe 27 Co 28 Ni 29 Cu 30 Zn 26.98 28.09 30.97 32.07 35.45 39.95 31 Ga 32 Ge 33 As 34 Se 35 Br 36 Kr 39.10 40.08 44.96 47.88 50.94 52.00 54.94 55.85 58.93 58.69 63.55 65.39 69.72 72.61 74.92 78.96 79.90 83.80 37 Rb 38 Sr 39 Y 40 Zr 41 Nb 42 Mo 43 Tc 44 Ru 45 Rh 46 Pd 47 Ag 48 Cd 49 In 50 Sn 51 Sb 52 Te 53 I 54 Xe 85.47 87.62 88.91 91.22 92.91 95.94 (98) 101.1 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.3 55 Cs 56 Ba 57 La 72 Hf 73 Ta 74 W 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 132.9 137.3 138.9 178.5 181.0 183.8 186.2 190.2 192.2 195.1 197.0 200.6 87 Fr 88 Ra 89 Ac 104 Rf 105 Db 106 Sg 107 Bh 108 Hs 109 Mt 110 111 112 112 116 118 (269) (272) (277) (277) (289) (293) (223) 226.0 227.0 58 Ce 59 Pr (262) (262) 150.4 152.0 157.3 158.9 162.5 164.9 167.3 168.9 173.0 175.0 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr 144.2 91 Pa 92 U 238.0 237.0 (244) 63 Eu (266) (243) 64 Gd (247) 65 Tb (247) 66 Dy (251) 67 Ho (252) 68 Er (257) 69 Tm (258) 70 Yb (209) (145) 140.9 62 Sm (265) 209.0 61 Pm 90 Th 231.0 (263) 207.2 60 Nd 140.1 232.0 Page 4 (261) 204.4 (259) (210) (222) 71 Lu (260) Not valid for use as an USNCO Olympiad National Exam after April 26, 2005. CHEMISTRY OLYMPIAD KEY FOR NATIONAL EXAM – PART II 1. a. 3Cu + 8H+ + 2NO3 – → 3Cu2+ + 2NO + 4H 2 O b. i. mRT (0.205 g)(0.0821 L ⋅ atm⋅ mol-1 ⋅ K -1 )(308 K) MM ave = = = 43.5 g ⋅ mol-1 1 atm PV (0.125 L)(725 mmHg × 760 mm Hg) € € For the ratio, let x be the fraction of NO2 , xMM NO 2 + (1- x)MM O 2 = 43.5 46x + 32(1- x) = 43.5 solving for x : x = 0.821 NO 2 0.821 4.59 = = O2 0.189 1 ii. 40.86 g - 39.49 g = 1.37 g Cu 41.21 g - 40.86 g = 0.35 g O 1 mol Cu 1.36 g Cu × = 0.0216 mole Cu 63.55 g Cu 1 mol O 0.35 g O × = 0.0219 mole O 16.0 g O Therefore the ratio is 1 : 1 and the formula must be CuO. c. The lost solution will cause the mass of CuO to be too low relative to the mass of Cu. Therefore the percentage determination for copper will be too high d. The mass of CuO will be too high. Therefore the percentage determination for copper will be too low. 2. a. looking at N2 + 2H2 → N 2 H4 N 2 + O 2 → NO 2 we must use the equations, 2H 2 + O 2 → 2H 2O 2NO 2 + 2H 2O → 3O 2 + N 2 H 4 N 2 + 2H 2 → N 2 H 4 ΔH f = 2ΔH1 + 2ΔH 2 – ΔH 3 2ΔH1 2ΔH 2 – ΔH 3 b. N 2 H4 (l) + 2H2 O2 (l) → N2 (g) + 4H2 O(g) c. The structures of hydrazine and peroxide are, € and Page 5 Thus, ΔH rxn = 4 mol × (–285.8 kJ ⋅ mol-1 ) – [2 mol × (–187.8 kJ ⋅ mol-1 ) + 1 mol × (50.6 kJ ⋅ mol-1 )] = – 1143.2 + 325.0 = – 818.2 kJ d. ΔH rxn = N – N + 4 N – H + 2 O – O + 4 O – H – [N ≡ N + 8 O – H ] € (note : 4O – H cancel from each side) = 167 + 4(386) + 2(142) – [942 + 4(459)] = – 783 kJ e. ΔH from part c should be more accurate. ΔHf values are determined for each compound individually, whereas bond energies are average values. We should expect the actual values for the compounds in this problem to vary from these averages. € f. C = 1 mol × (29.1 J⋅mol -1⋅o C-1) + 4 mol × (33.6 J⋅mol -1⋅o C-1) so C=163.5 J⋅o C-1 q = CΔT 818200 J = 163.5 J⋅o C-1 × ΔT so ΔT = 5004 o C. 3. a. i. ii. iii. iv. [H3 NCH(CH3 )COOH]+ and +H3 NCH(CH3 )COO– + H3 NCH(CH3 )COO– + H3 NCH(CH3 )COO– and H2 NCH(CH3 )COO– H2 NCH(CH3 )COO– b. i. K1 [H3 NCH(CH3 )COOH]+ → H + + +H3 NCH(CH3 )COO– K2 +H3 NCH(CH3 )COO– → H + + H2 NCH(CH3 )COO– ii. [H +] at point 1 = K1 = 4.6×10-3 so pH = 2.34 [H+] at point 2 = K1 × K 2 = 9.2 ×10−13 = 9.6 ×10−7 so pH = 6.02 [H+] at point 3 = K2 = 2×10-10 so pH = 9.07 [1.0 ×10−10 ][B– ] [B– ] 2 so = [HB] [HB] 1 Take a specified volume of 0.10 M alanine hydrochloride(such as 300 mL) in which the predominant species is [H3 NCH(CH3 )COOH]+. Add 0.10 M NaOH so there is 5/3 as much as there is 0.10 M alanine [H NCH(CH 3 )COO – ] 2 (such as 500 mL). This will€give 800 mL of solution with a ratio of + 2 = . [ H 3NCH(CH 3 )COO – ] 1 c. 2.0 ×10−10€= € a. Ε ocell = Ε oox + Ε ored and Ε oox = 1.52 V – 0.34 V = 1.18 V so Ε ored (M 2+ + 2e – → M) = –1.18 V € 1.90 M b. i. If Cu2+ decreases to 0.10 M then M2+ must increase to 0.0257 1.90 − ln ii. Ε = 1.52 € = 1.52 – 0.01285 ln(19) = 1.52 – 0.0378 = 1.48 V € 2 0.10 € c. The E of the cell with dilute solutions will be the same as the original Eo . Because the solutions are diluted by the same amount and the ions have the same coefficients (from the balanced chemical equation), Q in the Nernst equation is 1, and lnQ = 0. € 4. Page 6 Not valid for use as an USNCO Olympiad National Exam after April 26, 2005. 5. a. Na2 S + H2 O → 2Na + + HS– + OH– b. I3 – + S2 O3 2– → 3I– + S4 O6 2– c. F– + Fe3+ → FeF 6 3– 90 0 d. 90 38 Sr → 39Y + -1β e. CO2 + Ca(OH) 2 → Ca2+ + HCO3 – f. CH3 CH2 OH + Cr2 O7 2– + H+ → CH 3 CH=O (or CH3 COOH) + Cr3+ + H2 O 6.€ a. b. c. d. 7. a. Cl2 + CHCl3 → CCl4 + HCl Cl and CCl3 Rate = k[CHCl3 ][Cl] Because this step is at equilibrium, we can express the [Cl] in terms of [Cl2 ] by looking at the [Cl]2 1 1 equilibrium constant expression. K = so [Cl]2 = K[Cl2 ] and [Cl] = K 2 [Cl2 ] 2 . Thus by [Cl2 ] substituting we get the overall expression to be: Rate = k[CHCl3 ][Cl2 ]1/2 e. If [CHCl3 ] and [Cl2 ] are doubled, rate will increase by (2)•(2)1/2 = 2.83 times. € and and b. XeF 2 is linear, 180o . XeF4 is square planar, 90o XeO3 trigonal pyramid, ~107o c. XeF2 is nonpolar. Both Xe–F bond dipoles are the same size, but due to the linear geometry they offset each other. XeF4 is nonpolar. All Xe–F bond dipoles are the same size, but due to the square planar geometry they offset each other. XeO3 is polar. The Xe–O bond dipoles are the same size, and the non planar geometry leads to a net dipole. d. Xe has a formal positive charge in all of these compounds. This makes them good oxidizing agents. 8. a. b. It is an addition polymer. To form, the double bond in a monomer breaks to give a lone electron that forms bonds to other monomers. No other product(s) are formed, so it cannot be condensation. c. i. Polypropylene with 10,000 units melts at a higher temperature than one with 1000 units. The larger molecule, with higher molar mass has stronger dispersion forces. ii. Replacing CH3 with CH2 CH2 CH2 CH3 will lower the melting temperature. The bulk of this larger group will impede the packing of the polymer chains and decrease the strength of the intermolecular forces. iii. Isotactic polypropylene will melt at a higher temperature than atactic polypropylene. The more regular structure of the isotactic form allows better packing and stronger intermolecular forces. Page 7 2005 U.S. NATIONAL CHEMISTRY OLYMPIAD KEY FOR NATIONAL EXAM – PART III Lab Problem 1 You have been given three beakers containing NaHCO3 (sodium hydrogen carbonate), CaCl2 (calcium chloride), and tap water. These two compounds react in the presence of water. Propose a balanced chemical equation to account for this reaction, and support your proposal with all possible qualitative and quantitative observations and measurements. This problem requires the students to carry out a reaction and make qualitative and quantitative observations. 1. Sample answer: CaCl2 + 2NaHCO3 → Ca(OH) 2 (s) + 2 CO2 (g) + 2 Na+(aq) + 2 Cl– (aq) + 2NaHCO3 → CaCO3 (s) + CO2 (g) + H2 O or CaCl2 + 2 Na+(aq) + 2 Cl– (aq) Weigh 1 mmol CaCl2 and add to one corner of the plastic bag, then weigh 2 mmol NaHCO3 and add to the other corner of the plastic bag. Add water to the NaHCO 3 corner without mixing, seal the bag, and weigh. Allow the reagents to mix, observing for bubbling and inflation of the bag (to verify CO 2 production) and formation of a white insoluble solid (to verify either Ca(OH)2 or CaCO3 ). Carefully expel the excess gas and reweigh the bag to measure the amount of CO2 lost, and determine whether that mass corresponds to one or two moles per mole of calcium. The pH of the final mixture will be measured as well; the first reaction should be close to neutral (both a weak base and a weak acid are formed), while the second should be noticeably acidic (from the CO2 ). An excellent answer will clearly note both qualitative and quantitative observations that will be made, and will indicate how they will be interpreted to provide evidence of the nature of the reaction. An average answer will note qualitative or quantitative observations that will be made, but will not have clear indications of how those observations will be used. A poor answer will show little awareness of the significance of mixing the reagents, or will not describe observations that will be attempted. Alternative approaches might include titrating the reagents against each other to determine their mole ratio in the balanced reaction (using cessation of bubbling to mark the endpoint); weight the precipitate formed; measuring the volume of CO 2 produced by inflating the bag and either directly measuring its volume by water displacement or its dimensions, or by measuring the drop in weight of the sealed bag after reaction and calculating the volume change from the mass of air displaced; observing the pH of the precipitate after resuspension in pure water (Ca(OH)2 ) is noticeably basic, CaCO3 is not); observing the exothermicity of the reaction. Key for 2005 USNCO National Exam, Part III 2. Data and observations Weighed 1.15 g CaCl 2 into one corner of the bag. Weighed 1.72 g NaHCO3 and carefully added to the other corner of the bag. Added about 15 mL water and carefully added to the CaCl 2 side; the compound dissolved and the water got hot. Let the solution cool, then sealed the bag and weighed; total mass – 23.03 g. Allowed the CaCl2 solution to mix with the solid sodium bicarbonate; the solution gets colder and bubbles vigorously, partially inflating the bag. The solution turns milky white. After bubbling stops, open the bag and carefully expel the excess gas; the mass is now 22.51 g. The pH of the suspension is about 6. 3. CaCl2 A balanced chemical equation and evidence of reaction. + 2NaHCO3 CaCO3 (s) + CO2 (g) + H2 O + 2 Na+(aq) + 2 Cl– (aq) An excellent equation will be balanced and demonstrate reasonable chemical reactivity; an average equation will be chemically reasonable, but may show inappropriate phases (e.g. NaCl(s)); a poor equation will have chemically unreasonable products (H2 , Cl2 , HOCl; or the production of Ca(OH)2 and H+ at the same time). Evidence for reaction: A colorless gas is evolved, confirming CO 2. A white precipitate forms, consistent with CaCO 3. The pH of the final solution is weakly acidic, consistent with formation of CO2 but not with OHof H 3O+. The solution initially becomes hot, and then cools; CaCl2 dissolves exothermically, but release of CO2 is endothermic. An excellent answer will clearly relate all observations to the features of the reaction; an average answer will relate only some observations to the reaction; a poor answer does not relate the observations to the equation. 4. Calculations 1.15 g CaCl2 /(110.9 g/mol) = 1.04 × 10–2 mol 1.72 g NaHCO3 /(84 g/mol) = 2.05 × 10–2 mol (limiting reagent) Mass CO2 released = 23.03 g – 22.51 g = 0.52 g 0.52 g CO2 /(44 g/mol) = 0.012 mol 0.012 mol CO2 /0.0205 mol NaHCO3 = 0.59 mol CO2 per mol bicarbonate, close to the 0.5 mol expected from the balanced reaction. An excellent answer relates the mass of produced product to the predicted mass of CO2 based on the balanced equation and the number of moles of the limiting reagent; an average answer will note the amount of mass lost but will not relate it to the amounts of reagent; a poor answer will note the number of moles of reagent but not of product. Key for 2005 USNCO National Exam, Part III Lab Problem 2 You have been given five vials, labeled #1-5. These vials contain methanol, 2-propanol, acetone, hexane and water (though not necessarily in this order). You have also been given a container of table sugar (C12H22O11 ). Design and carry out an experiment to determine which liquid is in each labeled vial. You have access to a clock or timer. This problem asks students to distinguish 5 different liquids. Observations and data can be used to identify each liquid by observing: 1. Relative viscosity 2. Miscibility 3. Evaporation rates 4. Solubility with the provided sucrose sample 5. Rates of solubility with sucrose 6. Odor 7. Densities Excellent answers for this problem included at least three of these observations in chart or table form as part of their plan. In the data/observations, we looked for quantitative data (time for evaporation, time for completely dissolved sugar samples using set volumes of each liquid and set masses of sugar samples, etc.). Credit was given for multiple trials with quantitative data and for complete observations with all five samples. Conclusions included correct identification of each of the five liquids using more than one observation for each as evidence for student results. The best student answers were organized, provided observations in chart or tabulated form, and had a high degree of specificity with units were included for quantitative data. Sample data: A single drop of each liquid was placed on a watch glass and the time was recorded it took for that drop to completely evaporate: Vial #1 #2 #3 #4 #5 liquid 2-propanol Water Methanol Hexane acetone evaporation time (one drop) in min. (at room temp.) 2:15 still there 1:05 0:30 0:20 These data correspond nicely to predictions based on the degree of hydrogen bonding and relative intermolecular forces of attractions between molecules of each of the samples present. Solubility with sucrose: Vial #2 dissolved the sugar easiest and quickest and in the greatest amount, evidence that #2 was water; #1 and #3 slightly; #4 and #5 less so. Excellent student answers showed experiments with consistent amounts of sugar and volumes of each vial used with recorded times for completely dissolved sugar observed. Miscibility: Vial #4 was immiscible with #2, forming two distinct layers, conclusive evidence that #4 was hexane on the top layer, water on the bottom layer. Vials #1, 2, 3 were miscible with one another; #4 slightly with #5. Key for 2005 USNCO National Exam, Part III Odors: For excellent data, students might have noted that #2 had no apparent odor, #1, 3 had an ‘alcohol-like’ smell, or even ‘smells like rubbing alcohol (#1), that #4 smelled like ‘gasoline’, or that #5 had a ‘nail polish-like’ smell. Key for 2005 USNCO National Exam, Part III 2006 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM PART 1 Prepared by the American Chemical Society Olympiad Examinations Task Force OLYMPIAD EXAMINATIONS TASK FORCE Arden P. Zipp, State University of New York, Cortland Chair Sherry Berman-Robinson, Consolidated High School, IL David W. Hostage, Taft School, CT William Bond, Snohomish High School, WA Adele Mouakad, St. John’s School, PR Peter E. Demmin (retired), Amherst Central High School, NY Jane Nagurney, Scranton Preparatory School, PA Marian Dewane, Centennial High School, ID Ronald O. Ragsdale, University of Utah, UT Kimberly Gardner, United States Air Force Academy, CO, Todd Trout, Lancaster Country Day School, PA Preston Hayes, Glenbrook South High School, IL DIRECTIONS TO THE EXAMINER–PART I Part I of this test is designed to be taken with a Scantron® answer sheet on which the student records his or her responses. Only this Scantron sheet is graded for a score on Part I. Testing materials, scratch paper, and the Scantron sheet should be made available to the student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until April 25, 2006, after which tests can be returned to students and their teachers for further study. Allow time for the student to read the directions, ask questions, and fill in the requested information on the Scantron sheet. The answer sheet must be completed using a pencil, not pen. When the student has completed Part I, or after one hour and thirty minutes has elapsed, the student must turn in the Scantron sheet, Part I of the testing materials, and all scratch paper. There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and you are free to schedule rest-breaks between parts. Part I Part II Part III 60 questions 8 questions 2 lab problems single-answer multiple-choice problem-solving, explanations laboratory practical 1 hour, 30 minutes 1 hour, 45 minutes 1 hour, 30 minutes A periodic table and other useful information are provided on page 2 for student reference. Students should be permitted to use nonprogrammable calculators. DIRECTIONS TO THE EXAMINEE–PART I DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Answers to questions in Part I must be entered on a Scantron answer sheet to be scored. Be sure to write your name on the answer sheet; an ID number is already entered for you. Make a record of this ID number because you will use the same number on both Parts II and III. Each item in Part I consists o f a question or an incomplete statement that is followed by four possible choices. Select the single choice that best answers the question or completes the statement. Then use a pencil to blacken the space on your answer sheet next to the same letter as your choice. You may write on the examination, but the test booklet will not be used for grading. Scores are based on the number of correct responses. When you complete Part I (or at the end of one hour and 30 minutes), you must turn in all testing materials , scratch paper, and your Scantron answer sheet. Do not forget to turn in your U.S. citizenship statement before leaving the testing site today. Page 1 Not valid for use as an USNCO Olympiad National Exam after April 25, 2006. ampere atmosphere atomic mass unit atomic molar mass Avogadro constant Celsius temperature centi– prefix coulomb electromotive force energy of activation enthalpy entropy equilibrium constant ABBREVIATIONS AND SYMBOLS A Faraday constant F molal atm formula molar mass M molar u free energy G molar mass A frequency ν mole NA gas constant R Planck’s constant °C gram pressure g c heat capacity rate constant Cp C hour retention factor h E joule J second Ea kelvin K temperature, K H kilo– prefix k time S liter L volt K milli– prefix m CONSTANTS m M M mol h P k Rf s T t V R = 8.314 J·mol–1·K–1 R = 0.0821 L·atm·mol–1·K–1 1 F = 96,500 C·mol–1 1 F = 96,500 J·V–1·mol–1 NA = 6.022 × 1023 mol–1 h = 6.626 × 10–34 J·s c = 2.998 × 108 m·s –1 0 °C = 273.15 K 1 atm = 760 mmHg EQUATIONS E= E − o 1 1A 1 H −∆H 1 ln K = + constant R T RT ln Q nF k E 1 1 ln 2 = a − k1 R T1 T2 PERIODIC TABLE OF THE ELEMENTS 3 Li 2 2A 4 Be 13 3A 5 B 14 4A 6 C 15 5A 7 N 6.941 9.012 10.81 12.01 11 Na 12 Mg 13 Al 14 Si 22.99 24.31 26.98 19 K 1.008 18 8A 2 He 16 6A 8 O 17 7A 9 F 14.01 16.00 19.00 20.18 15 P 16 S 17 Cl 18 Ar 28.09 30.97 32.07 35.45 39.95 31 Ga 32 Ge 33 As 34 Se 35 Br 36 Kr 4.003 10 Ne 20 Ca 3 3B 21 Sc 4 4B 22 Ti 5 5B 23 V 6 6B 24 Cr 7 7B 25 Mn 8 8B 26 Fe 9 8B 27 Co 10 8B 28 Ni 11 1B 29 Cu 12 2B 30 Zn 39.10 40.08 44.96 47.88 50.94 52.00 54.94 55.85 58.93 58.69 63.55 65.39 69.72 72.61 74.92 78.96 79.90 83.80 37 Rb 38 Sr 39 Y 40 Zr 41 Nb 42 Mo 43 Tc 44 Ru 45 Rh 46 Pd 47 Ag 48 Cd 49 In 50 Sn 51 Sb 52 Te 53 I 54 Xe 85.47 87.62 88.91 91.22 92.91 95.94 (98) 101.1 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.3 55 Cs 56 Ba 57 La 72 Hf 73 Ta 74 W 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 132.9 137.3 138.9 178.5 180.9 183.8 186.2 190.2 192.2 195.1 197.0 200.6 204.4 207.2 209.0 (209) (210) (222) 87 Fr 88 Ra 89 Ac 104 Rf 105 Db 106 Sg 107 Bh 108 Hs 109 Mt 110 111 112 114 (223) (226) (227) (261) (262) (263) (262) (265) (266) (269) (272) (277) (2??) Page 2 58 Ce 59 Pr 60 Nd 61 Pm 62 Sm 63 Eu 64 Gd 65 Tb 66 Dy 67 Ho 68 Er 69 Tm 70 Yb 71 Lu 140.1 140.9 144.2 (145) 150.4 152.0 157.3 158.9 162.5 164.9 167.3 168.9 173.0 175.0 Property of the ACS DivCHED Examinations Institute 90 Th 91 Pa 92 U 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr 232.0 231.0 238.0 (237) (244) (243) (247) (247) (251) (252) (257) (258) (259) (262) Not valid for use as an USNCO Olympiad National Exam after April 25, 2006. Page 3 DIRECTIONS § When you have selected your answer to each question, blacken the corresponding space on the answer sheet using a soft, #2 pencil. Make a heavy, full mark, but no stray marks. If you decide to change an answer, erase the unwanted mark very carefully. § This is a single use exam, so you may make marks in the test booklet. § There is only one correct answer to each question. Any questions for which more than one response has been blackened will not be counted. § Your score is based solely on the number of questions you answer correctly. It is to your advantage to answer every question. 1. Which substance is NOT paired correctly with its name? (A) baking soda - potassium hydrogen tartrate (B) chalk - calcium carbonate (C) Epsom salt - magnesium sulfate heptahydrate (D) Plaster of Paris - calcium sulfate hemihydrate 2. Which acid should be stored in plastic containers rather than in glass ones? (A) hydrofluoric acid (B) nitric acid (C) phosphoric acid (D) sulfuric acid (B) silicon (C) sulfur (D) tin (B) HCl (C) NH3 (D) O3 5. Which technique is preferred for delivering a solid into a pre-weighed beaker for weighing? (A) Transfer more of the reagent than is needed to the beaker. Return the excess to the bottle with a spatula. (B) Transfer the desired amount of solid from the reagent bottle by holding the neck of the open bottle over the beaker and tapping the bottle. Then weigh the beaker and solid. (C) Weigh a spatula, scoop the desired amount of solid from the bottle, transfer it to the beaker and reweigh the spatula. (D) Weigh a piece of filter paper, tap the neck of the bottle to transfer solid to the filter paper, weigh the filter paper and transfer the solid to the beaker. 6. Bronze is an alloy of (A) copper and tin (B) copper and zinc (C) nickel and tin (D) nickel and zinc Page 4 (B) 0.313 M (C) 0.500 M (D) 0.625 M 8. What is the concentration of the solution that results from mixing 40.0 mL of 0.200 M HCl with 60.0 mL of 0.100 M NaOH? (You may assume the volumes are additive.) (A) 0.150 M NaCl (C) 0.0200 M NaCl and 0.0600 M HCl (D) 0.0600 M NaCl and 0.0200 M HCl 9. Mole fractions are I. freezing point depression typically used to II. osmotic pressure calculate which III. vapor pressure properties for solutions containing nonvolatile solutes? 4. Which gas is odorless? (A) CH4 (A) 0.0301 M (B) 0.0200 M NaCl and 0.0200 M HCl 3. Which element does NOT occur as distinct allotropes at temperatures between 0°C and 150°C? (A) phosphorus 7. What is the molarity of KI in a solution that is 5.00% KI by mass and has a density of 1.038 g·cm-3? (A) I only (B) III only (C) I and II only (D) II and III only 10. An unknown anion in solution is to be identified by adding Ag+ and Ba 2+ ions to separate portions of it. Which anion would produce the results listed for it? (+ indicates the presence of a precipitate) Ag + Ba2+ (A) carbonate + – (B) hydroxide – + (C) iodide + – (D) sulfide – – 11. A 1.0 L portion of a 0.30m solution of which of the following would be most effective at removing ice from a sidewalk? (A) C6H12O6 (B) NaBr (C) KNO3 (D) CaCl2 Property of the ACS DivCHED Examinations Institute 12. C6H6 + Br2 r C6H5Br + HBr In an experiment to prepare bromobenzene according to the equation, a student reacted 20.0 g of C6H6 with 0.310 mol of bromine. If 28.0 g of C6H5Br was obtained, what was the percentage yield? (A) 31.5 (B) 40.3 (C) 57.6 13. When the substances are arranged in order of increasing boiling points, which order is correct? (B) II < III < I (C) III < II < I (D) III < I < II 14. A 225 mL sample of H2 is Vapor Pressure at 25 °C collected over water at H2O 24 mmHg 25 °C and 735 mmHg pressure. Which expression represents the set-up to find the volume of dry H2 at 0 °C and 1 atmosphere? 225 × (735 − 24 )× 273 760 × 298 (A) V= (B) V= (C) V= (D) 225 × (735 + 24 )× 298 V= 760 × 273 Pressure (D) 69.7 I. CH3(CH2)3CH3 II. CH3CH2CH(CH3)2 III. C(CH3)4 (A) I < II < III 17. In the van der Waals equation for real gases, corrections are introduced for both the pressure and the volume terms of the Ideal Gas Equation. Identify the origin of both correction factors and specify whether each is added to or subtracted from the corresponding term. Volume (A) attractive forces / subtracted molecular size / added (B) attractive forces / added molecular size / subtracted (C) molecular size / subtracted attractive forces / added (D) molecular size / added attractive forces / subtracted 18. The structure of a unit cell of an oxide of niobium is depicted here. Niobiums are dark and oxygens are light. What is the empirical formula of this compound? 225 × 760 × 298 (735 − 24) × 273 225 × 273 × 760 (735 + 24 ) × 298 (A) NbO 15. What is ? Hvap for the substance whose vapor pressure is represented by the diagram? (C) NbO3 (D) Nb 2O3 19. For a reaction that is exothermic and non-spontaneous at 25 °C, which quantity must be positive? (A) ?E° (B) ?G° (C) ?H° (D) ?S° 20. Use the thermochemical data given to calculate ? Hf° for N2O5(g) in kJ·mol-1. N2(g) + O2(g) r 2NO(g) ?H° = +180.5 kJ 2NO(g) + O2(g) r 2NO2(g) ?H° = –114.1 kJ 4NO2(g) + O2(g) r 2N2O5(g) ?H° = –110.2 kJ (A) –332.8 (A) 4.8 kJ·mol-1 (B) 33 kJ·mol-1 (C) 44 kJ·mol-1 (D) 50 kJ·mol-1 16. What occurs when liquid CH2F2 evaporates? (B) NbO2 I. Dispersion forces are overcome. II. Dipole-dipole forces are overcome. III. Covalent bonds are broken. (A) II only (B) III only (C) I and II only (D) I, II and III (B) –43.8 (C) 11.3 (D) 22.6 21. Bromine boils at 59°C with ?H°vap = 29.6 kJ·mol-1. What is the value of ?S°vap in J·mol-1·K-1? (A) 11.2 (B) 89.2 (C) 501 (D) 1750 22. The Ksp of calcium fluoride is 3.2×10-11. Calculate the ?G° (in kJ·mol-1) for the dissolving of solid calcium fluoride at 25°C. (A) 2.18 (B) 5.02 (C) 26.0 (D) 59.9 23. For which exothermic reaction is ?E more negative than ?H? Not valid for use as an USNCO Olympiad National Exam after April 25, 2006. Page 5 29. According to the reaction profile given, which reaction step is ratedetermining in the forward direction? (A) Br2(l) r Br2(g) (B) 2C(s) + O2(g) r 2CO(g) (C) H2(g) + F2(g) r 2HF(g) (D) 2SO2(g) + O2(g) r 2SO2(g) 24. For a reaction at 25°C, ?G = 12.7 kJ when the reaction quotient, Q = 10.0. What is the value of ?G° for this reaction? (A) –12.1 kJ (B) 7.0 kJ (A) I r II (B) II r III (C) 18.4 kJ (D) 37.5 kJ (C) III r II (D) III r IV 25. How can the rate of reaction at a specific time be determined from a graph of concentration against time? (A) concentration at that time divided by the time (B) logarithm of the concentration divided by the time (C) absolute value of the slope of the graph at that time (D) logarithm of the slope divided by the time 26. The rate constant for the radioactive decay of C-11 is 0.0341 min -1. How long will it take for a sample of C-11 to decrease to 1/4 of its original activity? (A) 20.3 min (B) 29.3 min (C) 40.6 min (D) 58.6 min (B) ln[A] vs time (C) [A]2 vs time (D) 1/ln[A] vs time (A) This mechanism is consistent with the rate law IF step 1 is the rate determining step. (B) This mechanism is consistent with the rate law IF step 2 is the rate determining step. 27. If a reaction A r B has the rate law k[A]2, which graph produces a straight line? (A) 1/[A] vs time 30. For the reaction; 2H2(g) + 2NO(g) r N2(g) + 2H2O(g), rate = k[H2][NO]2. This mechanism has been proposed: step 1 H2 + NO r H2O + N step 2 N + NO r N2 + O step 3 O + H2 r H2O Which statement about this rate law and mechanism is correct? 28. Two unimolecular reactions, I and II, have the same rate constant at 25 °C but Ea for reaction I is larger than Ea for reaction II. Which statement about these two reactions is correct? (A) k reaction I is the same as k reaction II at all temperatures. (B) k reaction I is larger than k reaction II at lower temperatures but smaller at higher temperatures. (C) k reaction I is smaller than k reaction II at lower temperatures but larger at higher temperatures. (D) k reaction I is larger than k reaction II at temperatures both lower and higher than 25 °C. (C) This mechanism is consistent with the rate law IF step 3 is the rate determining step. (D) This mechanism can not be consistent with the rate law, regardless of which step is rate-determining. 31. C(s) + CO2(g) s 2CO(g) I. raising the temperature II. adding solid C If this system is at III. decreasing the pressure equilibrium, which change(s) will alter the value of Kp? (A) I only (B) II only (C) I and III only (D) II and III only 32. A 0.10 M solution of a weak acid is 5.75% ionized. What is the Ka value for this acid? (A) 3.3×10-3 (B) 3.5×10-4 (C) 4.2×10-5 (D) 3.3×10-5 33. Which base is most suitable to prepare a buffer solution with a pH = 11.00? (A) ammonia (Kb = 1.8×10-5) (B) aniline (Kb = 4.0×10-10) (C) methylamine (Kb = 4.4×10-4) (D) pyridine (Kb = 1.7×10-9) Page 6 Not valid for use as an USNCO Olympiad National Exam after April 25, 2006. 34. Calculate the pH of H2CO3 Acid Ionization Constants a 0.10 M solution K1a 4.4×10-7 of H2CO3. K2a 4.7×10-11 (A) 3.68 (B) 5.76 (C) 7.36 (D) 9.34 35. Which saturated solution has the highest [OH-]? (A) aluminum hydroxide (Ksp = 1.8×10-32) (B) calcium hydroxide (Ksp = 8.0×10-6) (C) iron(II) hydroxide (Ksp = 1.6×10-14) (D) magnesium hydroxide (Ksp = 1.2×10-11) (A) Increasing the [Cu 2+] two-fold has the same effect on the cell voltage as increasing the [Ag+ ] four-fold. (B) Decreasing the [Cu 2+] ten-fold has the same effect on the cell voltage as decreasing the [Ag+ ] by the same ratio. (C) Decreasing the [Cu 2+] ten-fold has less effect on the cell voltage than decreasing the [Ag+ ] by the same amount. (D) Doubling the sizes of the cathode has exactly the same effect on the cell voltage as decreasing the [Cu 2+] by a factor of two. 39. 3Ni2+ + 2Al r 2Al3+ + 3Ni E° = 1.41 V For the reaction given, which expression gives the value of ?G° in kJ·mol-1? 36. Consider these mixtures: Mixture I. 100 mL of 0.006 M Pb(NO3)2 plus 50 mL of 0.003 M NaBr Mixture II. 100 mL of 0.008 M Pb(NO3)2 plus 100 mL of 0.006 M NaBr Which statement is correct? Ksp PbBr2 6.6×10-6 (A) −3× 96.5 1.41 (B) −6× 96.5 1.41 (C) −3× 96.5× 1.41 (D) −6 × 96.5× 1.41 (A) A precipitate will not form in either mixture. (B) A precipitate will form only in mixture I. 40. In which species is the oxidation number for hydrogen different from those in the other three? (C) A precipitate will form only in mixture II. (A) AlH3 (B) H3AsO4 (D) A precipitate will form in both mixtures. (C) H3PO3 (D) NH3 37. The equation for one of the half-reactions in a lead storage battery is: PbO2 + 4H+ + SO42- + 2e - r PbSO4 + 2H2O What happens to the properties of the electrolyte as this cell discharges? 41. A solution Standard Reduction Potential (V) containing Ni2+(aq) + 2e – r Ni(s) –0.236 equimolar Sn 2+(aq) + 2e – r Sn(s) –0.141 amounts of NiCl2 Br2(aq) + 2e – r 2Br –(aq) 1.077 and SnBr2 is Cl2(aq) + 2e – r 2Cl –(aq) 1.360 electrolyzed using a 9V battery and graphite electrodes. What are the first products formed? Density pH (A) increases increases (B) increases decreases (A) Ni(s) at cathode, Cl2(aq) at anode (C) decreases decreases (B) Ni(s) at cathode, Br2(aq) at anode (D) decreases increases (C) Sn(s) at cathode, Br2(aq) at anode 38. For the voltaic cell based on this reaction: 2Ag+ (aq) + Cu r Cu 2+(aq) + 2Ag the concentrations of the aqueous ions and sizes of the electrodes can be changed independently. Which statement is correct? (D) Sn(s) at cathode, Cl2(aq) at anode Not valid for use as an USNCO Olympiad National Exam after April 25, 2006. Page 7 42. True statements about the system shown after the passage of one Faraday of electricity include which of those given? (A) K-38 (B) K-39 I. The number of moles of Al formed is greater than the number of moles of silver formed. II. The final [Al3+] is greater than the final [Ag+ ]. III. The number of electrons reacting with Al3+ ions is the same as the number reacting with Ag+ ions. (A) I only (B) I and III only (C) II and III only (D) I, II and III 43. When l = 3, what are the possible values for the quantum number ml? (A) 2, 1, 0 (B) 3, 2, 1, 0 (C) 2, 1, 0, -1, -2 (D) 3, 2, 1, 0, -1, -2, -3 44. The first ionization energy of cesium is 6.24×10-19 J/atom. What is the minimum frequency of light that is required to ionize a cesium atom? (A) 1.06×10-15 s -1 (B) 4.13×1014 s -1 (C) 9.42×1014 s -1 (D) 1.60×1018 s -1 (A) H2Se, H2S, H2O (B) H2S, H2Se, H2O (C) H2S, H2O, H2Se (D) H2O, H2S, H2Se (B) II only (C) I and III only (D) II and III only 51. What is the shape of the TeF5– anion? (A) see-saw (B) square pyramidal (C) trigonal pyramidal (D) trigonal bipyramidal 52. How many sigma and pi bonds are in maleic acid, HO2CCHCHCO2H? (A) 7 sigma, 2 pi (B) 8 sigma, 3 pi (C) 9 sigma, 2 pi (D) 11 sigma, 3 pi 53. How many isomers exist for the octahedral compound, Pt(NH3)2Cl4? (A) 1 – 2– + 2– – (B) K , S , Cl 2– – + – 2– + (B) 2 (C) 3 (D) 4 54. What is the formal charge on the sulfur atom in SO2? (Assume a Lewis dot structure in which all atoms obey the octet rule.) 45. When the isoelectronic ions, Cl–, S2– and K+ are arranged in order of increasing size, which order is correct? + I. Na(g) r Na + (g) + e– II. F(g) + e– r F–(g) III. Na + (g) + F–(g) r NaF(s) (A) I only (A) +1 (C) S , Cl , K (D) K-43 49. When the species listed are arranged in order of increasing bond angle, which order is correct? 50. Which terms are exothermic for the formation of NaF(s)? (A) K , Cl , S (C) K-42 (C) –1 (D) –2 55. How many structural isomers are possible for C6H14? (A) 2 (D) Cl , S , K (B) +2 (B) 3 (C) 4 (D) 5 56. Which is an ester? 46. Which is most similar for the elements in a group in the periodic table? (A) CH3COOCH2CH3 (B) (CH3)3COOC(CH3)3 (C) CH3OCH3 (D) (CH3)3CCOOH (A) physical state 57. Which type of reaction is typical of aromatic compounds? (B) melting point (A) addition (C) first ionization energy (B) free-radical substitution (D) ground state electron configuration (C) substitution by positively-charged reagents 47. How many unpaired electrons are present in a gaseous Co 3+ ion in its ground state? (A) 1 (B) 3 (C) 4 (D) 5 (D) substitution by negatively-charged reagents 58. What is the IUPAC name of (CH3)2CHCH=CHCH3? 48. Which nucleus is not radioactive? Page 8 Not valid for use as an USNCO Olympiad National Exam after April 25, 2006. (A) 1,2-methyl-isopropylethene (B) 1,1-dimethyl-2-butene (C) 1-isopropylpropene (D) 4-methyl-2-pentene 59. Which compound can exist in optically active forms? (A) CH3CH2CH2CH2OH (B) CH3CH2CH(OH)CH3 (C) (CH3)2CHCH2OH (D) (CH3)3COH 60. How many different tripeptides can be formed from the amino acids glycine, alanine and valine if each is used only once in each tripeptide? (A) 3 (B) 4 (C) 5 (D) 6 END OF PART I Not valid for use as an USNCO Olympiad National Exam after April 25, 2006. Page 9 CHEMISTRY OLYMPIAD 2006 National Test, PART I KEY Number 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. Answer A A B A B A B D B C D D C A C C B A B C B D D B C C A C D B Number 31. 32. 33. 34. 35. 36. 37. 38. 39. 40. 41. 42. 43. 44. 45. 46. 47. 48. 49. 50. 51. 52. 53. 54. 55. 56. 57. 58. 59. 60. Property of the ACS DivCHED Examinations Institute Answer A B C A B A D C D A C C D C A D C B A D B D B A D A C D B D 2006 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART II – ANSWER KEY Prepared by the American Chemical Society Olympiad Examinations Task Force 1. (11%) A(g) + 3B(g) r 2C(g) Use the tabulated data to answer the questions about this reaction, which is carried out in a 1.0 L container at 25°C. Experiment Ao, mol Bo, mol Initial rate of formation of C, mol .L-1.min-1 1 2 3 a. b. c. d. 0.10 0.20 0.10 0.10 0.20 0.20 0.25 2.0 2.0 For experiment 1, give the initial rate of disappearance of i. A ii. B Determine the orders of A and B and write the rate law for the reaction. Calculate the value of the rate constant and give its units. For the initial amounts of A and B in experiment 1, state the initial rate of formation of C under the following conditions. Justify your answer in each case. i. 0.50 mol of neon gas is added to the 1.0 L container. ii. the volume of the container is increased to 2.0 L. a. rate of formation C = 0.25 mol·L-1·min-1 so 1 = 0.13 mol ⋅ L-1 ⋅min −1 2 i. rate of disappearance of A = 3 0.25 × = 0.38 mol ⋅ L-1 ⋅ min −1 2 ii rate of disappearance of B = 0.25 × b. for experiments 2 and 3: [B] is constant while [A] doubles and the rate of the reaction is unchanged. The reaction order with respect to A is zero. for experiments 3 and 1: [A] is constant while [B] doubles and the rate of the reaction increases by a factor of 8. The reaction order with respect to B is 3. Rate = k[B] 3 k= c. 0.25 (0.10) 3 = 250 L2 ⋅ mol -2 ⋅ min d. i. The rate of formation of C is 0.25 mol·L-1·min-1. Adding an inert gas does not change the rate law. ii. The rate of formation of C is 0.031 mol·L-1·min-1. Doubling the volume of the container changes the concentration of B. [B] = 0.050 mol·L-1· Rate = k [B] = 250 L2 ⋅ mol -2 ⋅ min [0.050 mol ⋅ L-1] = 0.031 3 3 2. (13%) A 0.472 g sample of an alloy of tin and bismuth is dissolved in sulfuric acid to produce tin(II) and bismuth(III) ions. This solution is diluted to the mark in a 100 mL volumetric flask and 25.00 mL aliquots are titrated with a 0.0107 M solution of KMnO4, forming tin(IV) and manganese(II) ions. (The bismuth ions are unaffected during this titration.) Not valid for use as an USNCO Olympiad National Exam after April 25, 2006. Page 1 a. b. c. d. e. Write a balanced equation for the reaction of the MnO4- ion with Sn(II) in acid solution. If an average titration requires 15.61 mL of the MnO4- solution, calculate the number of moles of MnO4- used in an average titration. Determine the percentage of tin in the alloy. State how the end point of the titration is detected. Describe and explain the effect on the calculated percentage of tin in the alloy if the same volume of MnO4- solution is used with the following differences: i. During the titration the solution pH increases so that MnO2 is formed rather than Mn(II). ii. The solution volume in the volumetric flask was above the mark on the flask but a volume of 100. mL was assumed. iii. The sample of the original alloy had an oxide coating on it. – 2+ + → 2Mn 2+ + 5Sn 4 + + 8H O 2 a. 2MnO 4 + 5Sn + 16H b. 0.01561 L × 0.0107 mol ⋅ L-1 = 1.669 ×10 -4 mol 1.669 × 10 -4 mol × c. %Sn = 5 mol Sn 2+ 118.7 g Sn 2+ × = 0.04953 g of Sn in 25 mL 2 mol MnO –4 1 mol Sn 2+ 4 × 0.04953 g ×100 = 41.97% 0.472 g sample d. The endpoint is shown by the persistence of a faint purple color. (Indicating that there is an excess of MnO4- .) e. i, The %Sn that is determined is too high. The MnO4-/Sn2+ ratio is 2:3 when MnO2 is formed. Thus, for a given number of moles of 2 2 Sn 3 as many moles of MnO4- will be used rather than 5 . Since the calculations assume MnO4- , the latter ratio is used. So 5 ×2 = 5 ( 2 3 3 ) of the correct moles of Sn would be calculated. ii. The %Sn that is determined would be too low. Each 25 mL aliquot of contain fewer Sn2+ ions, requiring less MnO4- in the titration. iii. The %Sn that is determined would be too low. The oxide coating leads to fewer Sn2+ ions released into solution per g of sample weighed. 3. (15%) In water, HCN is a weak acid with pKa = 9.6. a. Calculate the the Ka and the [H+ ] in a 0.15 M solution of HCN. b. In a closed system, these equilibria are established: HCN(g) s HCN(aq) HCN(aq) s H+ (aq) + CN-(aq) i. Calculate Kp for the equilibrium between HCN(g) and HCN(aq) at 298 K. ? Gf° kJ.mol-1 HCN(g) 124.7 HCN(aq) 119.7 ii. If the total cyanide concentration in solution (i.e. [CN-] + [HCN]) is 0.10M, calculate the partial pressure of HCN(g) in this system at pH = 7. iii. A concentration of 300 ppm of HCN in air is reported to be toxic to humans after a few minutes exposure. Determine the ratio of the pressure calculated in 3.b.ii. to this value. c. Gold can be extracted from its ores by reacting the ore with O2 gas in the presence of aqueous CN- ions according to this equation. 4Au(s) + 8CN-(aq) + O2(g) + 2H2O(l) s 4Au(CN)2-(aq) + 4OH-(aq) i. Write the equilibrium expression for this reaction. ii. At fixed [CN–] and O2 pressure will the amount of Au(CN)2– be greatest at high or low pH? Justify your answer. iii. What purpose does O 2 serve in the extraction process? −9.6 = 2.5× 10 −10 a. pKa = 9.6, so K a = 10 2 Not valid for use as an USNCO Olympiad National Exam after April 25, 2006. Ka = [H + ][CN– ] = 2.5 ×10−10 = x 2 0.15 so x = [H+] = 6.1×10-6 [HCN] b. i. HCN(g) s HCN(aq) so, ? Go = 119.7 – 124.7 = –5.0 kJ·mol-1 ? Go = –RTlnK, so –5000 J·mol-1 = –(8.314 J·mol-1·K-1)(298 K)lnK K = e 2.018 = 7.52 ii. [CN–]+[HCN]=0.10, pH=7.0 so [H+]=1.0×10-7 2.5 ×10−10 = (1.0 × 10−7 )(0.10 − x) x 2.5 ×10−10 x = 1.0 × 10−8 − 1.0 × 10−7 x 1.0025 ×10−7 x = 1.0 ×10−8 so x = [HCN]eq = 0.0998 Kp = [HCN]eq pHCN pHCN = so 0.0998 = 0.0133 atm 7.52 iii. This problem is ambiguous in terms of how it might be solved – by mass or by volume Perhaps, the more obvious method is to solve by volume: 0.0133 atm = 0.0133 = 13300 ppm (vol) 1 atm so the ratio is 13300 ppm / 300 ppm = 44.3 We can also use mass… 0.0133 atm HCN 27 g mol -1 HCN × = 0.0125 = 12500 ppm (mass) -1 1 atm air 28.8 g mol air so the ratio is 12500 ppm / 300 ppm = 41.6 [Au(CN)–2 ] [OH – ] K= [CN– ]8 [O 2] 4 c. i. 4 ii. Low pH should be used because it will lower the [OH–], shifting the equilibrium towards production of products. iii. O2 is an oxidizing agent in this reaction. 4. (13%) The combustion of ethane, C2H6, produces carbon dioxide and liquid water at 25°C. a. b. c. d. Write an equation for this reaction. Given that ? H°comb for ethane under these conditions is -1560.5 kJ/mol ethane, calculate i. ? Hf° for ethane. ? Hf° kJ.mol-1 Bond Energies, kJ.mol-1 CO2(g) –393.5 ii. the bond energy of the C=O bond. C–C 347 H2O(l) –285.8 H–C 413 Given ? G°= –1467.5 kJ/mol, Calculate ? S° for this reaction in J.mol-1.K-1. H–O 464 Compared with combustion to form liquid water at 25 °C, how would combustion O=O 495 to form H2O(g) affect each of the following; i. ? H°combustion ii. ? S°combustion iii. ? G°combustion a. 2C2H6(g) + 7O2(g) r 4CO2(g) + 6H2O(l) (note: dividing all coefficients by 2 was given full credit) b. i. ∆H rxn = 4∆H f (CO 2 ) + 6∆H f (H 2O) – 2∆H f (C 2H 6 ) Not valid for use as an USNCO Olympiad National Exam after April 25, 2006. Page 3 so –3121 = 4(–393.5) + 6(–285.8) – 2∆Hf(C 2H6) and ∆Hf(C 2H6) = –83.9 kJ·mol-1 ii. ∆H rxn = 2BE C– C +12BE C– H + 7BE O= O – 8BE C= O – 12BE H– O so –3121 = 2(347) + 12(413) + 7(495) – 12(464) – 8BEC=O and = BEC=O = 6668 / 8 = 833 kJ.mol-1 c. ∆G o = ∆H o – T∆S o so –1467.5 = –1560.5 – 298 ∆So solving for ∆So yields ∆So = (1560.5 – 1467.5) / –298 = –312 J·K-1 d. i. . The measured ∆Hcombustion is less negative because the heat of vaporization was not released. ii. The ∆Scombustion is more positive (less negative) because H2O(g) has greater entropy than H2O(l) iii. The ∆Gcombustion is less negative due to the combination of these two effects. 5. (12%) Write net equations for each of the combinations of reactants below. Use appropriate ionic and molecular formulas and omit formulas for all ions or molecules that do not take part in a reaction. Write structural formulas for all organic substances. You need not balance the equations. All reactions occur in aqueous solution unless otherwise indicated. a. Excess carbon dioxide is bubbled through a suspension of calcium hydroxide. b. Acidified solutions of cerium(IV) and iron(II) are mixed. c. Solid calcium carbide is added to water. d. Excess concentrated ammonia is added to aqueous nickel(II) nitrate. e. Solutions of silver acetate and hydrobromic acid are mixed. f. Gaseous hydrogen chloride is reacted with gaseous propene. → Ca 2+ + HCO –3 a. Ca(OH) 2 + CO 2 + H 2O → Ce 3+ + Fe 3 + b. Ce 4 + + Fe 2+ c. CaC 2 + H 2 O → Ca 2+ + OH – + C 2 H 2 d. Ni 2+ + NH 3 → Ni(NH 3 ) 2+ 6 e. Ag + C 2 H 3O 2 + H + Br → HC 2 H 3 O 2 + AgBr f. HCl + H 3C – CH = CH 2 → H 3C – CHCl – CH 3 + – + (note: Ca(OH)2 was accepted as a product) – 6. (13%) Answer the following questions. a. For the molecule XeOF4. i. Write a Lewis structure. ii. Predict its geometry and specify the bond angles. iii. State whether it is polar or nonpolar. Explain your answer. b. Nitric acid, HNO3, is a strong acid while phosphoric acid, H3PO4, is a weak acid . i. Draw Lewis structures for each acid . ii. Explain why H3PO4 is stable while H3NO4 is not. iii. Suggest and explain two reasons that nitric acid is stronger than phosphoric acid. c. Ethane and diborane have similar formulas, C2H6 and B2H6, but B2H6 is more reactive. Sketch the structure of C2H6 and explain why B2H6 does not adopt this structure. 4 Not valid for use as an USNCO Olympiad National Exam after April 25, 2006. O O F F Xe F or F F F Xe F F a. i. ii. square pyramidal with 90º bond angles. O F F Xe F F iii. The molecule is polar with charges… O O H O N H O H O P O H O O O N O b. i. and H ii. H3NO4 would have more than 8 e– around N (OR) would put a positive (+) formal charge on N (OR) would be too sterically hindered with 4 oxygen atoms around the nitrogen. iii. (1) N is more electronegative that P, so electron density is shifted from H atoms towards the N, so the H+ can be more readily removed. (2) NO3– is stabilized by resonance more than H2PO4–. (3) HNO3 has two free oxygen atoms that attract electron density from the H atom, whereas H3PO4 has only one free oxygen atom. H H H c. 7. C C H H H B2H6cannot adopt this structure because it has only 12 valence electrons where C2H6 has 14. (11%) A common lecture demonstration involves electrolyzing a 1.0 M aqueous NaI solution containing phenolphthalein with a 9V battery. a. Write a balanced equation for the half-reaction that occurs at the i. anode. ii. cathode. b. Describe what is observed in the solution at the i. anode. ii. cathode. c. If a current of 0.200 amperes is passed through a 25.0 mL solution for 90.0 minutes, calculate the; i. number of moles of electrons passed through the solution. ii. number of moles of each of the products formed. a. i. 2I– r I2+ 2e– ii. 2H2O + 2e– r H2+ 2OH– b. i. At the anode, a yellow-brown color appears due to formation of iodine. If starch is added the color is blue ii. At the cathode, bubbles form due to the formation of hydrogen gas. The solution turns pink if phenolphthalein is added. c. i. 0.200 C ⋅s -1 × 90 min × 60 s ⋅ min -1 = 1080 C 1080 C × 1 mol e – = 1.12 ×10−2 mol e – 96500 C Not valid for use as an USNCO Olympiad National Exam after April 25, 2006. Page 5 1.12 ×10−2 mol e – × 1 mol I 2 = 5.6 ×10−3 mol I 2 2 mol e – 1.12 ×10−2 mol e – × 1 mol H 2 = 5.6 ×10 −3 mol H 2 2 mol e – 1.12 ×10−2 mol e – × 2 mol OH – = 1.12 ×10−2 mol OH – 2 mol e – ii. 8. (12%) There are four isomeric unsaturated compounds (alkenes) with the formula C4H8. a. Draw and name each of these isomers. b. These compounds all react with water in the presence of H2SO4 as a catalyst. i. Name the type of compound formed in this reaction. ii. Three of the four isomers form the same compound during this reaction. Identify these three isomers and outline your reasoning. c. Draw the structure of a saturated compound with the formula C4H8 and describe a chemical test that could be used to distinguish between this compound and one of the alkenes above. (Describe the results obtained for the saturated and unsaturated compound.) a. H H H C C H C H H H C H 1-butene H C C H H C H H H H C H H 2-methylpropene H H C C H H C C H H H H H H C H C C H C H H H cis -2-butene trans-2-butene b. i. An alcohol. ii. 1-butene, cis -2-butene and trans-2-butene all form 2-butanol. The OH of H2O will attack the 2° C of the double bond rather than the 1o C. H H H C H C H iii. H H C H H C C H C C H C H H H H H or Add Br2 to each. Br2 will be decolored with an unsaturated compound because Br2 adds to the double bond. Br2 will not change in the presence of saturated compounds. 6 Not valid for use as an USNCO Olympiad National Exam after April 25, 2006. END OF KEY PART II amount of substance ampere atmosphere atomic mass unit atomic molar mass Avogadro constant Celsius temperature centi- prefix coulomb electromotive force energy of activation enthalpy entropy ABBREVIATIONS AND SYMBOLS n equilibrium constant K measure of pressure mmHg A Faraday constant F milli- prefix m atm formula molar mass M molal m u free energy G molar M A frequency mol ν mole NA gas constant h R Planck’s constant °C gram pressure P g c heat capacity k Cp rate constant C hour retention factor Rf h E joule s J second Ea kelvin c K speed of light H kilo- prefix T k temperature, K S liter t L time volt V CONSTANTS R = 8.314 J·mol–1 ·K–1 R = 0.0821 L·atm·mol–1 ·K–1 1 F = 96,500 C·mol–1 1 F = 96,500 J·V–1 ·mol–1 NA = 6.022 × 1023 mol–1 h = 6.626 × 10–34 J·s c = 2.998 × 108 m·s –1 PERIODIC TABLE OF THE ELEMENTS 1 H 2 He 1.008 4.003 3 Li 4 Be 5 B 6 C 7 N 8 O 9 F 10 Ne 6.941 9.012 10.81 12.01 14.01 16.00 19.00 20.18 11 Na 12 Mg 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar 22.99 24.31 26.98 28.09 30.97 32.07 35.45 39.95 19 K 20 Ca 21 Sc 22 Ti 23 V 24 Cr 25 Mn 26 Fe 27 Co 28 Ni 29 Cu 30 Zn 31 Ga 32 Ge 33 As 34 Se 35 Br 36 Kr 39.10 40.08 44.96 47.88 50.94 52.00 54.94 55.85 58.93 58.69 63.55 65.39 69.72 72.61 74.92 78.96 79.90 83.80 37 Rb 38 Sr 39 Y 40 Zr 41 Nb 42 Mo 43 Tc 44 Ru 45 Rh 46 Pd 47 Ag 48 Cd 49 In 50 Sn 51 Sb 52 Te 53 I 54 Xe 85.47 87.62 88.91 91.22 92.91 95.94 (98) 101.1 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.3 55 Cs 56 Ba 57 La 72 Hf 73 Ta 74 W 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 132.9 137.3 138.9 178.5 181.0 183.8 186.2 190.2 192.2 195.1 197.0 200.6 204.4 207.2 209.0 (209) (210) (222) 87 Fr 88 Ra 89 Ac 104 Rf 105 Db 106 Sg 107 Bh 108 Hs 109 Mt 110 111 112 112 116 118 (223) 226.0 227.0 (261) (262) (263) (262) (265) (266) (269) (272) (277) (277) (289) (293) 58 Ce 59 Pr 60 Nd 61 Pm 62 Sm 63 Eu 64 Gd 65 Tb 66 Dy 67 Ho 68 Er 69 Tm 70 Yb 71 Lu 140.1 140.9 144.2 (145) 150.4 152.0 157.3 158.9 162.5 164.9 167.3 168.9 173.0 175.0 90 Th 91 Pa 92 U 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr 232.0 231.0 238.0 237.0 (244) (243) (247) (247) (251) (252) (257) (258) (259) (260) Not valid for use as an USNCO Olympiad National Exam after April 25, 2006. Page 7 2006 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART III Prepared by the American Chemical Society Olympiad Laboratory Practical Task Force OLYMPIAD LABORATORY PRACTICAL TASK FORCE Steve Lantos, Brookline High School, Brookline, MA Chair Linda Weber, Natick High School, Natick, MA Sheldon L. Knoespel, Mott Community College, Flint, MI Jim Schmitt, Eau Claire North High School, Eau Claire, WI Christie B. Summerlin, University of Alabama-Birmingham, Birmingham, AL DIRECTIONS TO THE EXAMINER–PART III The laboratory practical part of the National Olympiad Examination is designed to test skills related to the laboratory. Because the format of this part of the test is quite different from the first two parts, there is a separate, detailed set of instructions for the examiner. This gives explicit directions for setting up and administering the laboratory practical. There are two laboratory tasks to be completed during the 90 minutes allotted to this part of the test. Students do not need to stop between tasks, but are responsible for using the time in the best way possible. Each procedure must be approved for safety by the examiner before the student begins that procedure. Part III 2 lab problems laboratory practical 1 hour, 30 minutes Students should be permitted to use non-programmable calculators. DIRECTIONS TO THE EXAMINEE–PART III DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. WHEN DIRECTED, TURN TO PAGE 2 AND READ THE INTRODUCTION AND SAFETY CONSIDERATIONS CAREFULLY BEFORE YOU PROCEED. There are two laboratory-related tasks for you to complete during the next 90 minutes. There is no need to stop between tasks or to do them in the given order. Simply proceed at your own pace from one to the other, using your time productively. You are required to have a procedure for each problem approved for safety by an examiner before you carry out any experimentation on that problem. You are permitted to use a non-programmable calculator. At the end of the 90 minutes, all answer sheets should be turned in. Be sure that you have filled in all the required information at the top of each answer sheet. Carefully follow all directions from your examiner for safety procedures and the proper disposal of chemicals at your examining site. Not valid for use as an USNCO National Examination after April 26, 2006. Page 1 2006 UNITED STATES NATIONAL CHEMISTRY OLYMPIAD PART III — LABORATORY PRACTICAL Student Instructions Introduction These problems test your ability to design and carry out laboratory experiments and to draw conclusions from your experimental work. You will be graded on your experimental design, on your skills in data collection, and on the accuracy and precision of your results. Clarity of thinking and communication are also components of successful solutions to these problems, so make your written responses as clear and concise as possible. Safety Considerations You are required to wear approved eye protection at all times during this laboratory practical. You also must follow all directions given by your examiner for dealing with spills and with disposal of wastes. Lab Problem 1 Turmeric, a natural compound, is added to mustard for flavor and color. It changes color from yellow to red at a pH of 7.4. Mustard also contains acetic acid. Given a sample of 0.50 M NaOH and the packets of mustard, create and perform an experiment to determine the mass percentage of acetic acid in mustard. Lab Problem 2 Given a sample of 3.0 M hydrochloric acid, phenolphthalein, and some common laboratory equipment, devise an experiment using both qualitative and quantitative evidence to determine the provided unknown metal given these possible choices: Ag, Al, Ca, or Cr. Page 2 Not valid for use as an USNCO National Examination after April 26, 2006. Answer Sheet for Laboratory Practical Problem 1 Student's Name: __________________________________________________________________________ Student's School:________________________________________ Date: ___________________________ Proctor's Name:__________________________________________________________________________ ACS Section Name :________________________________Student's USNCO test #: ________________ 1. Give a brief description of your experimental plan. This is a titration experiment. Yellow mustard must be between 2.6 – 3.5% acetic acid by law (See: Current CFR 21 for 2005 [ http://www.gpoaccess.gov/cfr/index.html ]http://www.gpoaccess.gov/cfr/index.html). Yellow mustard contains turmeric, here used as an indicator for this experiment. Before beginning your experiment, you must get Examiner’s Initials: Not valid for use as an USNCO National Examination after April 26, 2006. Page 3 approval (for safety reasons) from the examiner. 2. Record your data and other observations. 3. Calculations. The calculations would be: 1. moles base used (V x M) = moles acid present 2. moles acid present x molar mass acetic acid = mass acetic acid 3. Percentage of acetic acid in mustard = mass acetic acid present / mass mustard used Sample Calculation: 0.50 g mustard weighed, titrated with a volume of 0.5 mL NaOH moles OH- = 0.0005 L x 0.5M = 0.00025 mol OH- = 0.00025 mol H+ from HC2H3O2 in mustard mass HC2H3O2 = 0.00025 mol x 60 g/mol = 0.015 g HC2H3O2 finally, % acetic acid in mustard = 0.015 g /.50 g x100 = approx. 3.0% The percentage of acetic acid in your sample of mustard = ___3.0%__________________ Excellent work: Student was able to complete two or more trials and average their results, using a minimum amount of both mustard and NaOH for each titration. Results were clearly shown and observations, i.e. color changes and endpoint were clearly noted. Student thought to make dilute aqueous solutions with each of the samples of mustard in order to completely dissolve the mustard and be able to more clearly note a uniform and lasting color change Average work: Student only completed one trial. Evidence of a titration was performed. Measurements between trials were fairly consistent. Below average work: Student was not able to conclude that this was a titration experiment, or did so, but did not perform the titration correctly to obtain a mass/volume of NaOH added. Only one trial was performed. Measurements were inaccurate or inconsistent between trials. Page 4 Not valid for use as an USNCO National Examination after April 26, 2006. Answer Sheet for Laboratory Practical Problem 2 Student's Name: __________________________________________________________________________ Student's School:________________________________________ Date: ___________________________ Proctor's Name:__________________________________________________________________________ ACS Section Name : ________________________________Student's USNCO test #: ________________ 1. Give a brief description of your experimental plan. Students were to provide both qualitative and quantitative evidence to determine the unknown metal. The metal provided was calcium. The results to this experiment should have included both evidence form data obtained and exclusive information about what was not observed from students’ previous chemical knowledge. Conclusions come from knowledge about each metals’ reactivity to both water and HCl, with phenolphthalein, a possible titration, and gas generation. Students might also have explored reactivity of the metal with NaOH from Problem #1 (this is allowed, though not necessary to successfully complete this problem). Excellent work: Student combined HCl with the unknown metal (Ca) to obtain hydrogen gas in the wellplate, clearly showing evidence of gas production and exothermic reaction. A titration The student then performed this reaction with a measured amount of Ca and excess HCl using the Luer-Lok syringe to quantify the hydrogen gas produced (since room temperature and pressure were not given, student had to make some assumptions about the Kelvin temperature and room pressure, perhaps estimating 298K and 1 atm) to determine the expected volume of hydrogen and compare it to a theoretical volume produced from Ca + HCl CaCl2 + H2 Noting the color change when phenolphthalein is added to the metal reacted to either water or HCl. It is possible that a student might have thought to combine mustard (from Prob.#1) with the metal from this experiment. If so, mustard on the surface of Ca produces over time a crusty white solid, Ca(C2H3O2)2 (there is no evidence of reaction with mustard on the surface of Cr and with pure Ag, no visible reaction). Concluding what DIDN’T occur: If Cr + HCl greenish color indicating CrCl3 (or green color with many chromium salts If Ag + HCl no reaction If Al + HCl no visible reaction due to aluminum oxide layer (though student might have attempted to dissolve the metal with the NaOH from Exp. #1, if Al, would dissolve; Ca + NaOH gives Ca(OH)2, a noticeable milky white precipitate, with phenolphthalein produces a pink color. a) Sample titration experiment conclusions: Reacting a 0.10 g metal turning with water completely, adding phenolphthalein, then titrating with the 3M HCl to obtain a 2 : 1 ratio of OH- : H+ in solution, confirms that OH- must be present in the metal hydroxide form, M(OH)2. b) Sample data for quantifying hydrogen gas generated using the Luer-Lok® syringe: One metal turning, approx. 0.07 g Ca in excess 3M HCl Not valid for use as an USNCO National Examination after April 26, 2006. Page 5 Begin at 12 mL mark on syringe End at 40 mL mark on syringe 40 – 12 = 28 mL hydrogen gas generated, strongly exothermic reaction. Assume room temp. 25oC (298K) and 1 atm: using the ideal gas law, PV = nRT (1 atm)(0.028L) = n (0.0821atm L/mol K) (293K) ; n = 0.00116 mol H2(g) if given 0.07g of calcium, Ca + 2HCl CaCl2(aq) + H2(g) , then 0.0035 g of hydrogen gas is produced, corresponds roughly to number of moles of H2(g) made with these assumed conditions. Average work: Student reacted metal with HCl and concluded hydrogen gas was present but didn’t quantify the gas produced, or did but incorrectly. Student wrote out possible reactions with the other possibilities but did not do so correctly. Below average work: Student was unable to conclude that hydrogen gas was produced, did not use either a titration or quantitative method of data collection, or unable to use the phenolphthalein to qualitatively justify the metal. Before beginning your experiment, you must get approval (for safety reasons) from the examiner. Examiner’s Initials: 2. Record your data and other observations. (See comments above) Page 6 Not valid for use as an USNCO National Examination after April 26, 2006. 3. Conclusions and Evidence. The unknown metal is = __Calcium__________ PERIODIC TABLE OF THE ELEMENTS 1 H 2 He 1.008 4.003 3 Li 4 Be 5 B 6 C 7 N 8 O 9 F 10 Ne 6.941 9.012 10.81 12.01 14.01 16.00 19.00 20.18 11 Na 12 Mg 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar 22.99 24.31 26.98 28.09 30.97 32.07 35.45 39.95 19 K 20 Ca 21 Sc 22 Ti 23 V 24 Cr 25 Mn 26 Fe 27 Co 28 Ni 29 Cu 30 Zn 31 Ga 32 Ge 33 As 34 Se 35 Br 36 Kr 39.10 40.08 44.96 47.88 50.94 52.00 54.94 55.85 58.93 58.69 63.55 65.39 69.72 72.61 74.92 78.96 79.90 83.80 37 Rb 38 Sr 39 Y 40 Zr 41 Nb 42 Mo 43 Tc 44 Ru 45 Rh 46 Pd 47 Ag 48 Cd 49 In 50 Sn 51 Sb 52 Te 53 I 54 Xe 85.47 87.62 88.91 91.22 92.91 95.94 (98) 101.1 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.3 55 Cs 56 Ba 57 La 72 Hf 73 Ta 74 W 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 132.9 137.3 138.9 178.5 181.0 183.8 186.2 190.2 192.2 195.1 197.0 200.6 204.4 207.2 209.0 (209) (210) (222) 87 Fr 88 Ra 89 Ac 104 Rf 105 Db 106 Sg 107 Bh 108 Hs 109 Mt 110 111 112 112 116 118 (223) 226.0 227.0 (261) (262) (263) (262) (265) (266) (272) (277) (277) (289) (293) (269) 58 Ce 59 Pr 60 Nd 61 Pm 62 Sm 63 Eu 64 Gd 65 Tb 66 Dy 67 Ho 68 Er 69 Tm 70 Yb 71 Lu 140.1 140.9 144.2 (145) 150.4 152.0 157.3 158.9 162.5 164.9 167.3 168.9 173.0 175.0 90 Th 91 Pa 92 U 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr 232.0 231.0 238.0 237.0 (244) (243) (247) (247) (251) (252) (257) (258) (259) (260) Not valid for use as an USNCO National Examination after April 26, 2006. Page 7 2007 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM PART 1 Prepared by the American Chemical Society Olympiad Examinations Task Force OLYMPIAD EXAMINATIONS TASK FORCE Arden P. Zipp, State University of New York, Cortland Chair Sherry Berman-Robinson, Consolidated High School, IL David W. Hostage, Taft School, CT Peter E. Demmin (retired), Amherst Central High School, NY Marian Dewane, Centennial High School, ID Jane Nagurney, Scranton Preparatory School, PA Kimberly Gardner, United States Air Force Academy, CO, Preston Hayes, Glenbrook South High School, IL Adele Mouakad, St. John’s School, PR Ronald O. Ragsdale, University of Utah, UT Todd Trout, Lancaster Country Day School, PA DIRECTIONS TO THE EXAMINER–PART I Part I of this test is designed to be taken with a Scantron® answer sheet on which the student records his or her responses. Only this Scantron sheet is graded for a score on Part I. Testing materials, scratch paper, and the Scantron sheet should be made available to the student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until May 1, 2007, after which tests can be returned to students and their teachers for further study. Allow time for the student to read the directions, ask questions, and fill in the requested information on the Scantron sheet. The answer sheet must be completed using a pencil, not pen. When the student has completed Part I, or after one hour and thirty minutes has elapsed, the student must turn in the Scantron sheet, Part I of the testing materials, and all scratch paper. There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and you are free to schedule rest-breaks between parts. Part I Part II Part III 60 questions 8 questions 2 lab problems single-answer multiple-choice problem-solving, explanations laboratory practical 1 hour, 30 minutes 1 hour, 45 minutes 1 hour, 30 minutes A periodic table and other useful information are provided on page 2 for student reference. Students should be permitted to use nonprogrammable calculators. DIRECTIONS TO THE EXAMINEE–PART I DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Answers to questions in Part I must be entered on a Scantron answer sheet to be scored. Be sure to write your name on the answer sheet; an ID number is already entered for you. Make a record of this ID number because you will use the same number on both Parts II and III. Each item in Part I consists of a question or an incomplete statement that is followed by four possible choices. Select the single choice that best answers the question or completes the statement. Then use a pencil to blacken the space on your answer sheet next to the same letter as your choice. You may write on the examination, but the test booklet will not be used for grading. Scores are based on the number of correct responses. When you complete Part I (or at the end of one hour and 30 minutes), you must turn in all testing materials, scratch paper, and your Scantron answer sheet. Do not forget to turn in your U.S. citizenship statement before leaving the testing site today. Not valid for use as an USNCO Olympiad National Exam after May 1, 2007. Distributed by the ACS DivCHED Examinations Institute, University of Wisconsin - Milwaukee, Milwaukee, WI. All rights reserved. Printed in U.S.A. ABBREVIATIONS AND SYMBOLS A Faraday constant F molal atm formula molar mass M molar u free energy G molar mass A frequency ν mole N A gas constant R Planck’s constant °C gram g pressure c heat capacity C p rate constant C hour h retention factor E joule J second Ea kelvin K temperature, K H kilo– prefix k time S liter L volt K milli– prefix m ampere atmosphere atomic mass unit atomic molar mass Avogadro constant Celsius temperature centi– prefix coulomb electromotive force energy of activation enthalpy entropy equilibrium constant CONSTANTS m M M mol h P k Rf s T t V R = 8.314 J·mol –1·K–1 R = 0.0821 L·atm·mol –1·K–1 1 F = 96,500 C·mol–1 1 F = 96,500 J·V–1·mol–1 N A = 6.022 × 10 23 mol–1 h = 6.626 × 10 –34 J·s c = 2.998 × 10 8 m·s –1 0 °C = 273.15 K 1 atm = 760 mmHg EQUATIONS E = Eo − 1 1A 1 H 1.008 3 Li RT ln Q nF k E 1 1 ln 2 = a − k1 R T1 T2 −ΔH 1 ln K = + constant R T PERIODIC TABLE OF THE ELEMENTS 2 2A 4 Be 13 3A 5 B 14 4A 6 C 15 5A 7 N 16 6A 8 O 17 7A 9 F 18 8A 2 He 4.003 10 Ne 6.941 9.012 10.81 12.01 14.01 16.00 19.00 20.18 11 Na 12 Mg 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar 22.99 24.31 19 K 20 Ca 3 3B 21 Sc 4 4B 22 Ti 5 5B 23 V 6 6B 24 Cr 7 7B 25 Mn 8 8B 26 Fe 9 8B 27 Co 10 8B 28 Ni 11 1B 29 Cu 12 2B 30 Zn 58.69 63.55 65.39 69.72 72.61 74.92 78.96 79.90 83.80 46 Pd 47 Ag 48 Cd 49 In 50 Sn 51 Sb 52 Te 53 I 54 Xe 26.98 28.09 30.97 32.07 35.45 39.95 31 Ga 32 Ge 33 As 34 Se 35 Br 36 Kr 39.10 40.08 44.96 47.88 50.94 52.00 54.94 55.85 58.93 37 Rb 38 Sr 39 Y 40 Zr 41 Nb 42 Mo 43 Tc (98) 44 Ru 101.1 45 Rh 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.3 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 85.47 87.62 88.91 91.22 92.91 95.94 55 Cs 56 Ba 57 La 72 Hf 73 Ta 74 W 132.9 137.3 138.9 178.5 180.9 183.8 186.2 190.2 192.2 195.1 197.0 200.6 87 Fr 88 Ra 89 Ac 104 Rf 105 Db 106 Sg 107 Bh 108 Hs 109 Mt 110 Ds 111 Rg 112 Uub (223) (226) (227) € 58 Ce 59 Pr (262) 60 Nd (263) 61 Pm (262) 62 Sm (265) 63 Eu (266) 64 Gd (269) 65 Tb (272) 66 Dy (277) 67 Ho 173.0 175.0 101 Md 102 No 103 Lr (145) 150.4 152.0 157.3 158.9 162.5 164.9 167.3 92 U 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm (243) (247) (247) (251) (252) (2??) 168.9 144.2 (244) 116 Uuh 69 Tm 91 Pa (237) (209) 68 Er 140.9 238.0 209.0 (2??) 90 Th 231.0 207.2 114 Uuq 140.1 232.0 Page 2 (261) 204.4 (257) (258) 70 Yb (259) (210) (222) 118 Uuo (2??) 71 Lu (262) Not valid as a USNCO National Exam after May 1, 2007 DIRECTIONS When you have selected your answer to each question, blacken the corresponding space on the answer sheet using a soft, #2 pencil. Make a heavy, full mark, but no stray marks. If you decide to change an answer, erase the unwanted mark very carefully. There is only one correct answer to each question. Any questions for which more than one response has been blackened will not be counted. Your score is based solely on the number of questions you answer correctly. It is to your advantage to answer every question. 1. Which absorbs gaseous carbon dioxide most effectively? (A) solid KOH (B) solid SiO 2 (C) aqueous HCl (D) aqueous NaF 2. Which laboratory results will tell whether an unknown white solid is NaOH or NH4NO3? 6. When a liquid is delivered from a volumetric pipet a small amount is typically retained in the tip. How should a student proceed in order to deliver the volume of liquid stated on the pipet? (A) Leave the small amount in the tip. (B) Use a pipet bulb to expel the remaining droplet. (A) NaOH is soluble in H 2O but NH4NO3 is not. (C) Shake the pipet to dispense the amount left in the tip. (B) Aqueous NaOH turns litmus blue but NH4NO3 does not. (D) Draw the liquid above the line initially to compensate for the amount that remains in the tip. (C) Aqueous NaOH reacts with copper metal but NH4NO3 does not. (D) NaOH gives a green flame test but NH4NO3 is colorless in a flame. 3. Which sets of chemicals, when mixed, produce the observation(s) listed? Combination Observation I. NH4Cl(s) and H2O(l) endothermic II. 9 M H2SO4(aq) and H2O(l) exothermic III. 1M NaOH(aq) and 1 M HCl(aq) exothermic (A) III only (B) I and II only (C) II and III only (D) I, II and III 4. What happens when 6 M nitric acid is added to an aqueous solution that contains 0.1 M Cl– and 0.1 M Ag(NH3)2+? (A) A deposit of silver metal forms. (B) A precipitate of AgCl forms. (C) Chlorine gas is released. (D) Gaseous ammonia is released. 5. A mixture of which 0.2 M aqueous solutions will form a precipitate that dissolves in 6 M nitric acid? (A) Co(NO3)2 and NH4Cl (B) Pb(NO 3)2 and NaBr (C) Ba(NO3)2 and Na2CO3 (D) Al(NO3)3 and K 2SO4 Not valid as a USNCO National Exam after May 1, 2007 7. What is the molarity of a 0.500 molal aqueous solution of calcium nitrate that has a density of 1.045 g·mL-1? (A) 0.483 M (B) 0.500 M (C) 0.522 M (D) 0.567 M 8. What volume of 0.150 M H2SO4 would be required to completely neutralize a mixture of 20.0 mL of 0.200 M NaOH and 40.0 mL of 0.0500 M Ca(OH)2? (A) 20.0 mL (B) 26.7 mL (C) 40.0 mL (D) 53.3 mL 9. A compound with the formula X2O5 contains 34.8% oxygen by mass. Identify element X. (A) arsenic (B) carbon (C) phosphorous (D) samarium 10. A solution of 0.0400 mol of C2H4Br2 and 0.0600 mol of C 3H6Br2 exerts a vapor pressure of 145.4 mm Hg at a certain temperature. Determine the vapor pressure of pure C 3H6Br2 at this temperature. Assume the vapor pressure of C 2H4Br2 at this temperature is 173 mm Hg and that the solution obeys Raoult's Law. (A) 76.2 mm Hg (B) 118 mm Hg (C) 127 mm Hg (D) 138 mm Hg Page 3 11. When 0.1 M aqueous solutions of aluminum nitrate, magnesium nitrate, sodium nitrate and urea, (NH2)2CO, are arranged in order of increasing boiling point, which order is correct? (A) Al(NO3)3 = Mg(NO3)2 = (NH2)2CO = NaNO 3 (B) Mg(NO 3)2 < (NH2)2CO < NaNO 3 < Al(NO3)3 (C) (NH2)2CO < NaNO 3 < Mg(NO3)2 < Al(NO3)3 (D) NaNO3 < Mg(NO 3)2 < Al(NO3)3 < (NH2)2CO 12. What is the maximum mass Molar Mass / g·mol–1 of Ba3(PO4)2 that can be Ba3(PO4)2 601.84 formed from Na 3PO4 163.94 0.00240 mol of Ba(NO3)2 and 0.131 g of Na3PO4? (A) 0.240 g (B) 0.480 g (C) 1.44 g (D) 7.22 g 13. Which segment of the heating curve obtained at constant pressure corresponds to the transition denoted by the arrow in the phase diagram? 16. The vapor pressure of phosphorus trichloride is 100 mm Hg at 21.0˚C and its normal boiling point is 74.2˚C. What is its enthalpy of vaporization in kJ. mol –1? (A) 0.493 (B) 3.93 (C) 23.0 (D) 32.4 17. If the absolute temperature of a sample of gas is increased by a factor of 1.5, by what ratio does the average molecular speed of its molecules increase? (A) 1.2 (B) 1.5 (C) 2.2 (D) 3.0 18. The curves in the accompanying diagram represent the PV/RT behavior of the gases: He, CH4 and C3H8. Which assignment of behavior to gas is correct? (A) 1 = He (B) 1 = C3H8 2 = CH 4 2 = CH 4 3 = C3H8 3 = He (C) 1 = CH 4 (D) 1 = C3H8 2 = C3H8 2 = He 3 = He 3 = CH 4 19. Calculate the standard enthalpy of formation of acetylene (in kJ. mol –1). 2C 2H2(g) + 5O 2(g) r 4CO2(g) + 2H 2O(l) ∆H˚ = –2243.6 kJ C(s) + O2(g) r CO2(g) ∆H˚ = –393.5 kJ H2(g) + 1/2 O2(g) r H 2O(l) ∆H˚ = –285.8 kJ (A) a (B) b (C) c (D) d 14. What is the molar mass of a gas that has a density of 5.66 g. L–1 at 35˚C and 745 mm Hg? (A) 127 (B) 141 (C) 143 (D) 146 15. Consider the solids: body-centered cubic (bcc), facecentered cubic (fcc), simple cubic (sc) (or primitive), constructed of spheres of the same size. When they are arranged in increasing order of the percentage of free space in a unit cell, which order is correct? (A) fcc, bcc, sc (B) bcc, sc, fcc (C) sc, fcc, bcc (D) bcc, fcc, sc (A) 49.0 (C) 1121.8 (D) 1564.3 20. The boiling point of diethyl ether is 34.6˚C. Which is true for the vaporization of diethyl ether at 25.0˚C? (A) ∆G˚vap > 0 (B) ∆H˚vap < 0 (C) K vap = 1 (D) ∆S˚ vap < 0 21. Estimate the Bond Dissociation Enthalpies / kJ. mol–1 enthalpy of C–C 350 C–O 350 combustion C–H 410 C=O 732 of methane in O–H 460 O–O 180 kJ. mol –1. O=O 498 CH4(g) + 2O2(g) r CO2(g) + 2H 2O(g) (A) 668 Page 4 (B) 98.0 (B) 540 (C) –540 (D) –668 Not valid as a USNCO National Exam after May 1, 2007 22. Which reaction has a positive ∆S˚reaction? 28. For the reaction A r B that is first-order in A, the rate constant is 2.08×10–2 s–1. How long would it take for [A] to change from 0.100 M to 0.0450 M? (A) Ag+(aq) + Br–(aq) r AgBr(s) (B) 2C 2H6(g) + 3O 2(g) r 4CO2(g) + 6H 2O(l) (A) 0.0166 s (C) N2(g) + 2H 2(g) r N 2H4(g) (D) 2H2O2(l) r 2H 2O(l) + O2(g) 23. For reactions I. constant number of moles conducted at constant II. constant temperature pressure, under what III. constant volume conditions are ∆E and ∆H equal? (A) I only (B) II only (C) III only (D) I and II only 24. For the reaction, H2(g) + I 2 (g) s 2HI(g) K p = 50.0 at 721 K. What is the value of ∆G˚ for this reaction (per mole of H2) at 721 K? (A) –32.3 kJ (B) –23.5 kJ (C) –10.2 kJ (D) –0.231 kJ 25. Which of these factors affect the value of the rate constant for a reaction? I. temperature II. reactant concentration III. use of a catalyst (A) I only (B) II only (C) I and III only (D) I, II and III 26. Which is the correct exponential form of the Arrhenius equation? (A) E = Ae a (C) k = Ae –k – RT RT (B) Ea (D) k = Ae Ea = Ae k – Ea RT RT €27. For the reaction A r B, € € what is the order with respect to A that gives this graph? (A) zero (C) 38.4 s (D) 107 s 29. These data were obtained for the reaction: X + Y r Z. X (M) Y (M) Rate: ∆Z/∆t / M·min–1 1.00 1.00 2.36×10-4 2.00 2.00 1.89×10-3 2.00 4.00 3.78×10-3 What is the rate law? (A) Rate = k[X][Y] (B) Rate = k[X]2[Y] (C) Rate = k[X][Y]2 (D) Rate = k[X]2[Y]2 30. A possible mechanism for the conversion of ozone to oxygen in the upper atmosphere is O3(g) s O2(g) + O(g) (fast equilibrium) O(g) + O 3(g) s 2O2(g) (slow) Which rate law is consistent with this mechanism? (A) Rate = k[O3] (B) Rate = k[O3]2 (C) Rate = k[O3][O] (D) Rate = k[O3]2[O 2]–1 31. A 0.050 M solution of an unknown acid is 1.0% ionized. What is the value of its K a ? (A) 2.5×10–7 (B) 5.0×10–6 (C) 5.0×10–4 (D) 5.0×10–2 32. Which mixture(s) form(s) buffer solutions? I. 100 mL of 0.200 M HF and 200 mL of 0.200 M NaF II. 200 mL of 0.200 M HCl and 200 mL of 0.400 M CH3CO2Na III. 300 mL of 0.100 M CH3CO2H and 100 mL of 0.300 M CH3CO2Na (A) I only (B) III only (C) II and III only (D) I, II and III 33. Determine the equilibrium constant for the reaction: HF(aq) + NH3(aq) s NH4+(aq) + F–(aq) given the equilibrium constants for the reactions. K a = 6.9×10–4 HF(aq) + H2O(l) s H3O+(aq) + F–(aq) + – NH3(aq) + H2O(l) s NH4 (aq) + OH (aq) K b = 1.8×10–5 K w = 1.0×10–14 2H2O(l) s H3O+(aq) + OH–(aq) € (B) first (B) 16.7 s (C) second (D) third Not valid as a USNCO National Exam after May 1, 2007 (A) 1.2×10–8 (B) 1.2×106 (C) 8.1×107 (D) 3.8×1015 Page 5 34. Calculate the pH of a 0.15 M solution of HOCl. (A) 3.77 Ka HOCl (B) 4.18 (C) 6.71 2.9×10 –8 (D) 8.36 35. For which reaction does K p = K c ? (A) 2C(s) + O2(g) s 2CO(g) 42. According to the tabulated standard reduction potentials E˚ = –2.38 V Mg2+(aq) + 2e– r Mg(s) – – 2H2O(l) + 2e r H2(g) + 2OH (aq) E˚ = –0.83 V E˚ = 0.53 V Br2(l) + 2e – r 2Br– (aq) + E˚ = 1.23 V O2(g) + 4H (aq) r 2H 2O(l) what products are formed during the electrolysis of an aqueous MgBr2 solution? (B) N2(g) + 3H 2(g) s 2NH3(g) (A) Mg and H 2 (B) H2 and Br2 (C) 2H2(g) + O2(g) s 2H2O(g) (C) H2 and O 2 (D) Mg and O 2 (D) H2(g) + I2(g) s 2HI(g) 36. CaF 2 has a Ksp = 3.9×10–11 at 25˚C. What is the [F–] in a saturated solution of CaF2 at 25˚C? (A) 2.1×10-4 (B) 3.4×10-4 (C) 4.3×10-4 (D) 6.8×10-4 (B) 2 / 1 (C) 3 / 1 (D) 5 / 1 38. Which change could occur at the anode of an electrochemical cell? (A) Cl– r Cl2 (B) H2O r H2 + (C) Na r Na (D) O2 r H2O 39. E˚ = 0.93 V for the reaction: Standard Reduction Potential / E˚ Fe2+(aq) + 2e– r Fe(s) –0.41 V Fe(s) + 2M+(aq) r Fe2+(aq) + 2M(s). What is the standard potential for M+ + e– r M? (A) 0.26 V (B) 0.52 V (C) 0.67 V (D) 1.34 V 3+ (A) V (aq) r V (aq) + e (B) VO3- + 2H+ r VO2+ + H2O (C) Mg 12 (D) Ar 18 n l ml ms (A) 1 0 0 –_ (B) 2 2 1 _ (C) 3 1 1 _ (D) 4 3 –3 _ 45. Which change(s) in electron structure occur when a gas phase Mn atom is converted to a Mn2+ ion in the gas phase? I. The number of occupied energy levels decreases. II. The number of half-filled orbitals decreases. (A) I only (B) II only (C) Both I and II (D) Neither I nor II (A) F, Ne, Na (B) Al, Mg, Na (C) Sr, Ca, Mg (D) Cl, Br, I – 47. How many unpaired electrons are in a gas phase Co2+ ion in its ground state? (C) VO2+ + 2H+ + e – r V3+ + H2O (A) 2 (D) VO2+ + H2O r VO2+ + 2H+ + e– 41. A solution of aqueous CuSO4 is electrolyzed with a 1.50 ampere current for 30.0 minutes. What mass of copper metal is deposited? (A) 0.889 g (B) 6C 46. Which list gives the symbols of the elements in the order of increasing first ionization energy? 40. For which half-reaction will a 1.0 unit increase in pH cause the greatest increase in half-cell potential? 2+ (A) 5B 44. Which set of quantum numbers is NOT allowed? 37. When the reaction: Cl– + ClO3– r Cl2 + H2O is balanced in acid solution what is the ratio of Cl– to ClO3–? (A) 1 / 1 43. Which is the symbol for an element whose ground state atoms have the same total numbers of s electrons and p electrons? (B) 1.19 g (C) 1.78 g (D) 3.56 g (B) 3 (C) 4 (D) 5 48. The energy required to ionize a potassium ion is 419 kJ⋅mol –1. What is the longest wavelength of light that can cause this ionization? (A) 285 nm (B) 216 nm (C) 200 nm (D) 107 nm 49. Which species has the same electron distribution around the central atom as SiF 4? (A) SF4 Page 6 (B) XeF4 (C) ClF4+ (D) BF4– Not valid as a USNCO National Exam after May 1, 2007 50. Which is/are polar species? I. SF2 II. SF4 III. SF6 (A) I only (B) III only (C) I and II only (D) II and III only 51. According to the Lewis dot structure for ozone, what is the formal charge on the central oxygen atom? (A) –2 (B) –1 55. How many unsaturated compounds have the formula C 4H8? (A) 3 (B) 4 (C) 5 (D) 6 56. Which compound is least soluble in water? (A) CH3CH2CH2F (B) CH3CH2CH2NH2 (C) CH3CH2CH2OH (D) CH3CH2CH2COOH 57. Which method for characterizing organic compounds relies on the vibration of atoms in the compound? (C) 0 (D) +1 52. When the species are arranged in order of increasing length of the carbon-oxygen bond, which order is correct? (A) Na 2CO3 < HCO2Na < CH3ONa (A) infrared spectroscopy (B) nuclear magnetic resonance spectroscopy (C) UV-visible spectroscopy (D) X-ray diffraction 58. Which substance reacts most rapidly with water? (B) CH3ONa < HCO 2Na < Na2CO3 (C) HCO2Na < Na2CO3 < CH 3ONa (A) C 6H5Cl (B) (CH3)3CCl (D) Na 2CO3 < CH 3ONa < HCO 2Na (C) (CH3)2CHCH2Cl (D) CH3CH2CH2CH2Cl 53. Which ionic solid would require the most energy to form gaseous ions? (A) NaF (B) Na 2O (C) MgO (D) MgF2 54. Solid calcium occurs as either cubic closest packing or hexagonal closest packing. What is the most significant difference between these two structures? (A) the placement of layers of calcium atoms (B) the distance betweeen calcium atoms in a single layer (C) the distance between calcium atoms in adjacent layers 59. What type of compound is formed by the mild oxidation of 2-pentanol? (A) acid (B) aldehyde (C) ester (D) ketone 60. Which species is lost during the formation of a disaccharide from a monosaccharide? (A) CH2 (B) CH2O (C) CH2OH (D) H2O END OF TEST (D) the coordination number of the calcium atoms in a single layer Not valid as a USNCO National Exam after May 1, 2007 Page 7 NATIONAL OLYMPIAD PART I 2007 KEY Number 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. Answer A B D B C A A B A C C A B D A D A B A A D D C B C D C C B D Number 31. 32. 33. 34. 35. 36. 37. 38. 39. 40. 41. 42. 43. 44. 45. 46. 47. 48. 49. 50. 51. 52. 53. 54. 55. 56. 57. 58. 59. 60. Property of the ACS DivCHED Examinations Institute Answer B D B B D C D A B D A B C B A C B A D C D C C A B A A B D D 2007 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART II Prepared by the American Chemical Society Olympiad Examinations Task Force OLYMPIAD EXAMINATIONS TASK FORCE Arden P. Zipp, State University of New York, Cortland Chair Sherry Berman-Robinson, Consolidated High School, IL David W. Hostage, Taft School, CT Peter E. Demmin (retired), Amherst Central High School, NY Marian Dewane, Centennial High School, ID Adele Mouakad, St. John’s School, PR Jane Nagurney, Scranton Preparatory School, PA Kimberly Gardner, United States Air Force Academy, CO, Preston Hayes, Glenbrook South High School, IL Ronald O. Ragsdale, University of Utah, UT Todd Trout, Lancaster Country Day School, PA DIRECTIONS TO THE EXAMINER–PART II Part II of this test requires that student answers be written in a response booklet of blank pages. Only this “Blue Book” is graded for a score on Part II. Testing materials, scratch paper, and the “Blue Book” should be made available to the student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until May 1, 2007, after which tests can be returned to students and their teachers for further study. Allow time for the student to read the directions, ask questions, and fill in the requested information on the “Blue Book”. When the student has completed Part II, or after one hour and forty-five minutes has elapsed, the student must turn in the “Blue Book”, Part II of the testing materials, and all scratch paper. Be sure that the student has supplied all of the information requested on the front of the “Blue Book,” and that the same identification number used for Part I has been used again for Part II. There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and you are free to schedule rest-breaks between parts. Part I Part II Part III 60 questions 8 questions 2 lab problems single-answer multiple-choice problem-solving, explanations laboratory practical 1 hour, 30 minutes 1 hour, 45 minutes 1 hour, 30 minutes A periodic table and other useful information are provided on the back page for student reference. Students should be permitted to use non-programmable calculators. DIRECTIONS TO THE EXAMINEE–PART II DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Part II requires complete responses to questions involving problem-solving and explanations. One hour and forty-five minutes are allowed to complete this part. Be sure to print your name, the name of your school, and your identification number in the spaces provided on the “Blue Book” cover. (Be sure to use the same identification number that was coded onto your Scantron® sheet for Part I.) Answer all of the questions in order, and use both sides of the paper. Do not remove the staple. Use separate sheets for scratch paper and do not attach your scratch paper to this examination. When you complete Part II (or at the end of one hour and forty-five minutes), you must turn in all testing materials, scratch paper, and your “Blue Book.” Do not forget to turn in your U.S. citizenship statement before leaving the testing site today. Not valid for use as an USNCO Olympiad National Exam after May 1, 2007. Page 1 ! ! 1. (12%) Compound X contains 2.239% hydrogen, 26.681% carbon and 71.080 % oxygen by mass. The titration of 0.154 g of this compound with 0.3351 M KOH produces the curve shown. a. b. c. d. e. a) convert masses to moles: # 1 mol & 2.239 g H " % ( = 2.221 mol (÷2.157) = 1.03 $ 1.008 g ' # 1 mol & 26.681 g C " % ( = 2.157 mol (÷2.157) = 1 $ 12.011 g ' # 1 mol & 71.08 g O " % ( = 4.443 mol (÷2.15) = 2.06 $ 16.00 g ' These numbers are close enough to whole numbers that the empirical formula must be CHO2 b) Obtain molar mass from titration (estimate endpoint at 10.4 mL) ! ! ! ! ! Determine the empirical formula of the compound. Calculate its molar mass and give its molecular formula. When K2Cr2O7 is reacted with X in acidic solution the products are chromium(III) ions and carbon dioxide. Describe the color change that accompanies this reaction. Write a balanced ionic equation for this reaction. Find the volume of dry carbon dioxide that could be collected at 22 ˚C and 738 mm Hg when 0.839 g of compound X is reacted with an excess of K2Cr2O7. Mol NaOH = 0.3351 mol/L " 0.0104 L = 0.00348 mol molar mass = 0.154 g ÷ 0.00348 mol = 44.2 g/mol ( " 2 because titration curve is diprotic) = 88.4 g/mol The molar mass is 88.4 g/mol, the empirical formula molar mass is 45.02. This value is close to half the value of the experimentally determined molar mass, so the molecular formula must be C2H2O4. c) 3+ The color change will be from orange for Cr2O 2– 7 to green for Cr . + 3+ d) The balanced ionic equation is: Cr2O 2– + 6CO 2 + 7H 2O 7 + 3H 2C 2O 4 + 8H " 2Cr e) ! Do the stoichiometry for oxalic acid to carbon dioxide, then calculate volume using ideal gas law. # 1 mol H !C O & # 6 mol H C O & 2 2 4 2 2 4 0.839 g H 2C 2O 4 " % ( "% ( = 0.0186 mol H 2C 2O 4 90.04 g H C O 3 mol H $ $ 2 2 4' 2C 2O 4 ' ( ) –1 –1 nRT ( 0.0186 mol) 0.0821 L " atm " mol " K ( 295 K) V= = = 0.464 L $ $ 1 atm '' P & 738 mmHg # & )) % 760 mmHg (( % 2. (15%) Coffee cup calorimetry experiments can be used to obtain ∆H f˚ for magnesium oxide. a. Write a balanced equation for the formation of magnesium oxide, whose enthalpy change is ∆H f˚. b. To determine the heat capacity of the calorimeter, 49.6 mL of 1.01 M HCl are reacted with 50.1 mL of 0.998 M NaOH. The solution's temperature increases by 6.40˚C. Determine the heat capacity of the calorimeter. You may assume the solution's specific heat capacity is 4.025 J·g–1 ·˚C–1 and the enthalpy of neutralization is –55.9 kJ per mole of H2O. c. When 0.221 g of magnesium turnings are added to 49.9 mL of 1.01 M HCl and 49.7 mL of H2O in the same calorimeter, the temperature increases by 9.67˚C. Write a balanced equation for the reaction that occurs and calculate the ∆H per mole of Page 2 Not valid for use as an USNCO Olympiad National Exam after May 1, 2007. d. e. magnesium. (Assume the solution's specific heat capacity is 3.862 J·g–1 ·˚C–1 and the calorimeter constant is the value obtained in b.) When 0.576 g of MgO react with 51.0 mL of 1.01 M HCl and 50.1 mL of H2O in the same calorimeter the temperature rises 4.72˚C. Write a balanced equation for this reaction and calculate its ∆H per mole of MgO using the same assumptions as in part c. Use the above results and ∆Hf˚ of H2O(l) (–285.8 kJ·mol–1) to calculate ∆H f˚ of magnesium oxide. a) Mg(s) + 1 2 O 2 (g) " MgO(s) b) First, determine the limiting reactant: Mol HCl = 1.01 mol/L " 0.0496 L = 0.00501 mol HCl Mol NaOH = 0.998 mol/L " 0.0501 L = 0.0500 mol NaOH , so because it is a 1:1 stoichiometry, NaOH is limiting. ! Via the enthalpy from the neutralization reaction, HCl(aq) +NaOH(aq) " NaCl(aq) +H 2O(l) #H = –55.9 kJ/mol we can calculate, # –55.9 kJ & 0.0500 mol NaOH " % ( = – 2.795 kJ $ 1 mol NaOH ' ! so the rest is taken up by the calorimeter: Account for heat taken up by the solution, Total volume of solution is 49.6 mL + 50.1 mL = 99.7 mL (no information is provided about density, so the simplest assumption is to use 1.00 g) so we have 99.7 g solution. Using the given specific heat capacity the heat absorbed by the solution is, ! ! ! heat = 99.7 g " 4.025 J # g –1#o C –1 " 6.40 o C = 2568 J (heat absorbed by the solution) Now we can calculate the heat absorbed by the calorimeter: 2795 J – 2568 J = 227 J absorbed by the calorimeter. So the heat capacity of the calorimeter is 227 J / 6.40 oC = 35.5 J· oC–1 ! c) The reaction of magnesium with an acid is: Mg + 2H + " Mg 2+ + H 2 Total mass is: 99.6 g solution + 0.221 g Mg = 99.821 g Total heat is heat absorbed by solution + heat absorbed by calorimeter: heat solution = 99.821 g " 3.862 J #!g –1#o C –1 " 9.67 o C = 3728 J heat calorimeter = 35.5 J"o C –1 # 9.67 o C = 343 J Total heat = 3728 J + 343 J = 4071 J # &( This is heat given off by 0.221 g Mg (using molar mass): 0.221 g Mg " %1 mol 24.31 g' = 0.00909 mol Mg $ –4071 J Thus, = –4.479 " 10 5 J # mol–1 = –447.9 kJ # mol–1 0.00909 mol ! ! ! d) The reaction is: MgO + 2H + " Mg 2+ + H 2O ! # &( First determine moles reacted: 0.576 g MgO " %1 mol 40.31 g' = 0.0143 mol MgO $ Once again, ! total heat is heat absorbed by solution + heat absorbed by calorimeter: (and solution mass includes MgO) heat solution = 101.676 g " 3.862 J # g –1#o C –1 " 4.72 o C = 1853 J ! J"o C –1 # 4.72 o C = 168 J heat calorimeter = 35.5 Total heat = 1853 J + 168 J = 2021 J –2021 J Thus, = –1.413 " 10 5 J # mol–1 = –141.3 kJ # mol–1 0.0143 mol ! ! ! ! ! e) Now construct a series of reactions that when summed are the formation reaction for MgO: ! Mg 2+ + H 2O " MgO + 2H + Mg + 2H + " Mg 2+ + H 2 #H = 141.3 kJ $ mol–1 #H = –447.9 kJ $ mol–1 Summed: These reaction yield: Mg + H 2O " MgO + H 2 #H = –306.6 kJ $ mol–1 Now combine this reaction with the heat of formation for water to yield the desired result: Not valid for use as an USNCO Olympiad National Exam after May 1, 2007. Page 3 ! Mg + H 2O " MgO + H 2 #H = –306.6 kJ $ mol–1 H 2 + 1 2 O 2 " H 2O #H = –285.8 kJ $ mol–1 #H = –592.4 kJ $ mol–1 Mg + 1 2 O 2 " MgO ! ! ! 3. (13%) Hydrogen sulfide, H2S, is a weak acid that can be used to precipitate metal ions from solution selectively by controlling the pH. Acid Ionization Constants, H2S K1 5.7×10–8 K2 1.3×10–13 a. Write equations to represent each of the ionization steps of H 2S. Ksp b. Write an equation to represent the overall ionization of H2S to form S2– and 2H+ Bi S 1.6×10–72 2 3 and calculate the equilibrium constant for this process. MnS 3.0×10–11 c. For a solution with [H2S] = 0.10 M, with [Bi3+] = [Mn2+] = 1.5 mM and [H+] = 10 mM, give the formula for the metal sulfide which precipitates first and calculate the percentage of it that will remain in solution at equilibrium. d. The pH of the solution is raised until the other metal sulfide begins to precipitate. Determine the pH of the solution at which the second metal sulfide begins to precipitate. a) H 2S " H + + HS – K1 = 5.7 # 10 –8 HS – " H + + S 2– K 2 = 1.3# 10 –13 b) H 2S " 2H + + S 2– K = 7.4 # 10 –21 ! c) Calculate sulfide ion concentration: ! ! 2 K= [ H 2S] 2 ( 0.010) [S 2– ] = 7.4 " 10 –21 so [S2–] = 7.4×10–18 ( 0.1) Now calculate Q and compare to K for each cation (with sulfide): Bismuth: Ksp = [Bi3+]2[S2–]3 = 1.6×10–72 Q = (1.5×10–3)2(7.4×10–18)3 = 9.1×10–58 Q > Ksp so there will be a precipitate formed. Manganese: Ksp = [Mn2+][S2–] = 3.0×10–11 Q = (1.5×10–3)(7.4×10–18) = 1.1×10–20 Q < Ksp so there will not be a precipitate formed. Thus – the bismuth is the first metal sulfide to precipitate. Now to calculate what percentage will remain in solution: 3+ 2 [Bi ] = K sp 2– 3 [S ] (1.6 " 10 ) (7.4 " 10 ) –72 = = 3.95" 10 –21 so [Bi3+] = 6.3×10–11 –18 3 The percentage can be calculated using the ratio of the amount remaining in solution divided by the original amount: ! ! [H + ] [S2– ] = (6.9 " 10 ) " 100 = 4.2 " 10 (1.5" 10 ) –11 %= –3 –6 % d) first determine the concentration of sulfide that will result in precipitation: Page 2 Not valid for use as an USNCO Olympiad National Exam after May 1, 2007. [S ] = 2– + 2 [H ] ! 4. = K sp [Mn 2+ (3.0 " 10 ) = 2.0 " 10 = ] (1.5" 10 ) K [ H 2S] [S ] 2_ –11 –3 –8 Now plug this value into the equation for K from Part (c): (7.4 " 10 )(0.10) = 3.7 " 10 (2.0 " 10 ) –21 = –14 –8 (10%) A galvanic cell is based on the half-reactions; Cr3+ + 3e– r Cr ! 2+ E˚ = –0.744 V – Ni + 2e r Ni a. b. c. d. e. f. ! ! ! ! ! and [H+] = 1.92×10–7 so pH = 6.7 E˚ = –0.236 V Write the balanced equation for the overall cell reaction. State which electrode increases in mass as the cell operates. Explain your answer. Calculate E˚cell Determine the value of ∆G˚ for the cell reaction at 25˚C. Calculate the value of K for the cell reaction at 25˚C. Find the voltage of the cell at 25˚C if [Cr3+] and [Ni2+] are both changed to 0.010 M. a) 2Cr + 3Ni2+ " 2Cr 3+ + 3Ni b) The nickel electrode increases in mass as the cell operates because Ni2+ ions in solution are reduced there (it is the cathode) and are deposited as Ni(s). c) E ocell = E ored + E oox = –0.236 V + 0.744 V = 0.508 V d) "G o = –nFE = –(6 mol)(96500 J # mol–1 # V–1)(0.508 V) = – 294000 J = – 294 kJ "G o = –RT ln K e) –294100 J = – (8.314 J # mol–1 # K –1)(298 K) lnK or K = 10 nE o 0.592 = 10 3.048 0.592 = 3.1" 10 51 ln K = 118.7 and K = 3.62 $ 10 51 2% " 0.0257 $ [.01] ' 0.0257 f) E = E o – ln$ = 0.508 V – ln(100 ! ) V = 0.508 – 0.0197 V = 0.488 V 3' 6 6 # [.01] & 5. (12%) Write net equations for each of the combinations of reactants below. Use appropriate ionic and molecular formulas and omit formulas for all ions or molecules that do not take part in a reaction. Write structural formulas for all organic substances. You need not balance the equations. All reactions occur in aqueous solution unless otherwise indicated. a. Solid ammonium chloride and solid calcium hydroxide are mixed. b. Excess carbon dioxide gas is bubbled into a sodium hydroxide solution. c. Sodium sulfite is added to a neutral potassium permanganate solution. d. Aqueous hydrofluoric acid is placed on a piece of silica. e. Chlorobenzene is heated with a mixture of concentrated nitric and sulfuric acids. f. Iodine-131 undergoes radioactive decay. a) NH 4 Cl + Ca(OH) 2 " CaCl2 + NH 3 + H 2O b) CO 2 + OH – " HCO –3 ! – 2– c) SO 2– 3 + MnO 4 " SO 4 + MnO 2 d) HF + SiO 2 " SiF4 + H 2O ! ! ! e) C 6 H 5Cl + H + + NO –3 " or Not valid for use as an USNCO Olympiad National Exam after May 1, 2007. ! Page 5 f) ! 131 53 I " 0 –1# + 131 54 Xe or 131 53 I " 42 He + 12751Sb (for half credit) 6. (12%) The reaction of NO with O2 to give NO2 is an important step in the commercial production of HNO3. a. Describe an experiment to measure the rate of this reaction. b. If the rate equation is found to be Rate = k[NO]2[O2], give the effect on the rate of tripling the concentration of. i. NO ii. O2 c. These two mechanisms have been proposed for this reaction, I 2NO + O2 r 2NO2 II 2NO r N2O2 N2O2 + O2 r 2NO2 i. State and explain which of the two mechanisms is more likely. ii. State and explain which of the two steps in mechanism II must be the slow step if this mechanism is to be consistent with the rate law in b. a) The stoichiometry of the reaction is: 2NO(g) + O2(g) r 2NO2(g) so there is a change in the number of moles of gas as the reaction goes forward. This reaction can be monitored by measuring the total pressure of the system as a function of time. Alternatively, the appearance of the red color of NO2 can be measured (e.g. with a spectropohotometer.) b) Because the rate law is Rate = k[NO]2[O2], tripling the concentration of NO will cause the rate to increase by a factor of 9. Tripling the concentration of O2 will cause the rate to increase by a factor of 3. c) i. Mechanism is II more likely because mechanism one involves a trimolecular collision. Such a collision is uncommon. By contrast, II has a pair of bimolecular reactions which are considerably more likely to occur. ii. Step 2 must be the slow step because it would have a rate law of Rate=k[N2O2][O2], but N2O2 is an intermediate whose concentration arises from the first step. With a slow second step, the first step achieves equilibrium, so [N2O2] = K[NO]2 and the overall rate law would be, Rate = k[NO]2[O2]. 7. (16%) Account for the following observations, a. The bond angle in H2O (104.5˚) is greater than that of H2S (92˚) but less than that in Cl2O (110.8˚). b. The bond dissociation energy of Cl2 (240 kJ·mol–1) is greater than that of F2 (154 kJ·mol–1) or Br2 (190 kJ·mol–1). c. The boiling point of NH3 is higher (–33˚C) than that of NF3 (–129˚C) but lower than that of NCl3 (71˚C). d. SiF4 is tetrahedral while SF4 has a see-saw shape and XeF4 is square planar. a) The angle in H-O-H is greater than H-S-H because the bonding pairs in H-S-H are further from the S atom (the atomic orbitals used in S have electron density that is further from the nucleus) so they can be forced closer together by the lone pair electrons on the S. The Cl-O-Cl bond angle is larger than either of the because the Cl atoms are large which gives rise to steric interference that forces them apart. b) The bond in Cl2 is stronger than that in F2 because the F atoms are sufficiently small that the lone pairs on the F atoms repel one another weakening the bond. The bond in Br2 is weaker than that of Cl2 because the obritals in the larger atom (Br) do not overlap as efficiently. c) The boiling point of NH3 is higher than that of NF3 because NH3 molecules can form hydrogen bonds with each other increasing the attractive forces relative to the dispersion and dipole forces between the NF3 molecules. For NCl3, the dispersion forces are sufficiently large (because of the large, polarizable Cl atoms) that they are stronger than the hydrogen bonding in NH3. d) SiF4 is tetrahedral, with four bonding pairs about the central Si atom. SF4 has 5 pairs of electrons (4 bonding and 1 lone pair) so it has a see-saw shape. XeF4 has 6 pairs of electrons (4 bonding and 2 lone pairs) so it has a square planar shape. 8. (10%) There are six different isomers with the formula C4H8O2 containing a –CO2 group. When added to water two of the six are substantially more soluble than the other four. a. Write structural formulas for the two water-soluble compounds and outline how their structures lead to their greater solubility. b. State the name of the class of compounds represented by the other four isomers. c. Draw structural formulas for any three of the four less soluble isomers. d. Write an equation for the laboratory synthesis of one of these four isomers and name each of the reactants. Page 2 Not valid for use as an USNCO Olympiad National Exam after May 1, 2007. a) The two more soluble forms have –COOH groups that make them more soluble because that group allows for hydrogen bonding with water. b) The less soluble isomers are esters. c) Any 3 of these four structures could be shown. d) An example reaction would be: END OF PART II Not valid for use as an USNCO Olympiad National Exam after May 1, 2007. Page 7 amount of substance ampere atmosphere atomic mass unit atomic molar mass Avogadro constant Celsius temperature centi- prefix coulomb electromotive force energy of activation enthalpy entropy ABBREVIATIONS AND SYMBOLS n equilibrium constant K measure of pressure mmHg A Faraday constant F milli- prefix m atm formula molar mass M molal m u free energy G molar M A frequency ν mole mol NA gas constant h R Planck’s constant °C gram P g pressure c heat capacity k Cp rate constant C hour Rf h retention factor E joule s J second Ea kelvin c K speed of light H kilo- prefix T k temperature, K S liter t L time volt V CONSTANTS R = 8.314 J·mol–1·K–1 R = 0.0821 L·atm·mol–1·K–1 1 F = 96,500 C·mol–1 1 F = 96,500 J·V–1·mol–1 NA = 6.022 × 1023 mol–1 h = 6.626 × 10–34 J·s c = 2.998 × 108 m·s–1 USEFUL EQUATIONS E = E! – ! k2 $ Ea ! 1 1 $ = ' " k1 &% R #" T1 T2 &% " –!H % " 1 % ln K = $ ' $ ' +c # R & # T& RT ln Q nF ln # PERIODIC TABLE OF THE ELEMENTS 1 H 2 He 1.008 4.003 3 Li 4 Be 5 B 6 C 7 N 8 O 9 F 10 Ne 6.941 9.012 10.81 12.01 14.01 16.00 19.00 20.18 11 Na 12 Mg 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar 22.99 24.31 26.98 28.09 30.97 32.07 35.45 39.95 19 K 20 Ca 21 Sc 22 Ti 23 V 24 Cr 25 Mn 26 Fe 27 Co 28 Ni 29 Cu 30 Zn 31 Ga 32 Ge 33 As 34 Se 35 Br 36 Kr 39.10 40.08 44.96 47.88 50.94 52.00 54.94 55.85 58.93 58.69 63.55 65.39 69.72 72.61 74.92 78.96 79.90 83.80 37 Rb 38 Sr 39 Y 40 Zr 41 Nb 42 Mo 43 Tc 44 Ru 45 Rh 46 Pd 47 Ag 48 Cd 49 In 50 Sn 51 Sb 52 Te 53 I 54 Xe 85.47 87.62 88.91 91.22 92.91 95.94 (98) 101.1 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.3 55 Cs 56 Ba 57 La 72 Hf 73 Ta 74 W 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 132.9 137.3 138.9 178.5 181.0 183.8 186.2 190.2 192.2 195.1 197.0 200.6 204.4 207.2 209.0 (209) (210) (222) 87 Fr 88 Ra 89 Ac 104 Rf 105 Db 106 Sg 107 Bh 108 Hs 109 Mt 110 Ds 111 Rg 112 Uub 114 Uuq 116 Uuh 118 Uuo (223) (226) (227) (261) (262) (263) (262) (265) (266) (269) (272) (277) (2??) (2??) (2??) Page 8 58 Ce 59 Pr 60 Nd 61 Pm 62 Sm 63 Eu 64 Gd 65 Tb 66 Dy 67 Ho 68 Er 69 Tm 70 Yb 71 Lu 140.1 140.9 144.2 (145) 150.4 152.0 157.3 158.9 162.5 164.9 167.3 168.9 173.0 175.0 90 Th 91 Pa 92 U 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr 232.0 231.0 238.0 237.0 (244) (243) (247) (247) (251) (252) (257) (258) (259) (260) Not valid for use as an USNCO Olympiad National Exam after May 1, 2007. 2007 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART III Prepared by the American Chemical Society Olympiad Laboratory Practical Task Force OLYMPIAD LABORATORY PRACTICAL TASK FORCE Steve Lantos, Brookline High School, Brookline, MA Chair Linda Weber, Natick High School, Natick, MA Jim Schmitt, Eau Claire North High School, Eau Claire, WI Christie B. Summerlin, University of Alabama-Birmingham, Birmingham, AL DIRECTIONS TO THE EXAMINER–PART III The laboratory practical part of the National Olympiad Examination is designed to test skills related to the laboratory. Because the format of this part of the test is quite different from the first two parts, there is a separate, detailed set of instructions for the examiner. This gives explicit directions for setting up and administering the laboratory practical. There are two laboratory tasks to be completed during the 90 minutes allotted to this part of the test. Students do not need to stop between tasks, but are responsible for using the time in the best way possible. Each procedure must be approved for safety by the examiner before the student begins that procedure. Part III 2 lab problems laboratory practical 1 hour, 30 minutes Students should be permitted to use non-programmable calculators. DIRECTIONS TO THE EXAMINEE–PART III DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. WHEN DIRECTED, TURN TO PAGE 2 AND READ THE INTRODUCTION AND SAFETY CONSIDERATIONS CAREFULLY BEFORE YOU PROCEED. There are two laboratory-related tasks for you to complete during the next 90 minutes. There is no need to stop between tasks or to do them in the given order. Simply proceed at your own pace from one to the other, using your time productively. You are required to have a procedure for each problem approved for safety by an examiner before you carry out any experimentation on that problem. You are permitted to use a non-programmable calculator. At the end of the 90 minutes, all answer sheets should be turned in. Be sure that you have filled in all the required information at the top of each answer sheet. Carefully follow all directions from your examiner for safety procedures and the proper disposal of chemicals at your examining site. Not valid for use as an USNCO National Examination after May 1, 2007 Page 1 2007 UNITED STATES NATIONAL CHEMISTRY OLYMPIAD PART III — LABORATORY PRACTICAL Student Instructions Introduction These problems test your ability to design and carry out laboratory experiments and to draw conclusions from your experimental work. You will be graded on your experimental design, on your skills in data collection, and on the accuracy and precision of your results. Clarity of thinking and communication are also components of successful solutions to these problems, so make your written responses as clear and concise as possible. Safety Considerations You are required to wear approved eye protection at all times during this laboratory practical. You also must follow all directions given by your examiner for dealing with spills and with disposal of wastes. Lab Problem 1 You have been given two ionic solutions, 0.10 M unknown salt, MClx solution and 0.10 M sodium solution, NazY. Devise and carry out an experiment to determine the identity of the unknown metal cation and the unknown anion in these solutions. The possible cations are potassium, zinc, aluminum, or silver. The possible anions are nitrate, carbonate, phosphate, or sulfide. You should provide both quantitative and qualitative evidence to support your answers. Lab Problem 2 LDPE (low density polyethylene, #4) is a petroleum-based polymer used to make flexible bottles, films, and plastic containers. Given water, ethanol (density = 0.789 g·mL–1), and the equipment provided, devise and carry out an experiment to precisely determine the thickness of the LDPE samples provided Page 2 Not valid for use as an USNCO National Examination after May 1, 2007. Answer Sheet for Laboratory Practical Problem 1 Student's Name: __________________________________________________________________________ Student's School:________________________________________ Date: ___________________________ Proctor's Name: _________________________________________________________________________ ACS Section Name :________________________________Student's USNCO test #: ________________ 1. Give a brief description of your experimental plan. Before beginning your experiment, you must get approval (for safety reasons) from the examiner. Not valid for use as an USNCO National Examination after May 1, 2007 Examiner’s Initials: Page 3 2. Record your data and other observations. 3. Calculations and Conclusions. Page 4 Not valid for use as an USNCO National Examination after May 1, 2007. Answer Sheet for Laboratory Practical Problem 2 Student's Name: __________________________________________________________________________ Student's School:________________________________________ Date: ___________________________ Proctor's Name: _________________________________________________________________________ ACS Section Name :________________________________Student's USNCO test #: ________________ 1. Give a brief description of your experimental plan. Before beginning your experiment, you must get approval (for safety reasons) from the examiner. Not valid for use as an USNCO National Examination after May 1, 2007 Examiner’s Initials: Page 5 2. Record your data and other observations. 3. Calculations and Conclusions. Your reported thickness of LDPE = Page 6 Not valid for use as an USNCO National Examination after May 1, 2007. PERIODIC TABLE OF THE ELEMENTS 2 He 1 H 1.008 4.003 3 Li 4 Be 5 B 6 C 7 N 8 O 9 F 10 Ne 6.941 9.012 10.81 12.01 14.01 16.00 19.00 20.18 11 Na 12 Mg 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar 22.99 24.31 26.98 28.09 30.97 32.07 35.45 39.95 19 K 20 Ca 21 Sc 22 Ti 23 V 24 Cr 25 Mn 26 Fe 27 Co 28 Ni 29 Cu 30 Zn 31 Ga 32 Ge 33 As 34 Se 35 Br 36 Kr 39.10 40.08 44.96 47.88 50.94 52.00 54.94 55.85 58.93 58.69 63.55 65.39 69.72 72.61 74.92 78.96 79.90 83.80 37 Rb 38 Sr 39 Y 40 Zr 41 Nb 42 Mo 43 Tc 44 Ru 45 Rh 46 Pd 47 Ag 48 Cd 49 In 50 Sn 51 Sb 52 Te 53 I 54 Xe 85.47 87.62 88.91 91.22 92.91 95.94 (98) 101.1 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.3 55 Cs 56 Ba 57 La 72 Hf 73 Ta 74 W 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 132.9 137.3 138.9 178.5 181.0 183.8 186.2 190.2 192.2 195.1 197.0 200.6 204.4 207.2 209.0 (209) (210) (222) 87 Fr 88 Ra 89 Ac 104 Rf 105 Db 106 Sg 107 Bh 108 Hs 109 Mt 110 Ds 111 Rg 112 Uub 114 Uuq 116 Uuh 118 Uuo (223) 226.0 227.0 (261) (262) (263) (262) (265) (266) (269) (272) (277) (2??) (2??) (2??) 58 Ce 59 Pr 60 Nd 61 Pm 62 Sm 63 Eu 64 Gd 65 Tb 66 Dy 67 Ho 68 Er 69 Tm 70 Yb 71 Lu 140.1 140.9 144.2 (145) 150.4 152.0 157.3 158.9 162.5 164.9 167.3 168.9 173.0 175.0 90 Th 91 Pa 92 U 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr 232.0 231.0 238.0 237.0 (244) (243) (247) (247) (251) (252) (257) (258) (259) (260) Not valid for use as an USNCO National Examination after May 1, 2007 Page 7 2007 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART III - ANSWERS Prepared by the American Chemical Society Olympiad Laboratory Practical Task Force Lab Problem #1 This problem involves knowledge of solubility rules and precipitation reactions. In addition, the identification of the unknown cation and anion requires relating the volumes (drops) of the reacting solutions to the quantity of precipitate produced and hence, to the molar ratios of the reacting ions. Plan The plan should include both an intention to gather qualitative information about the individual solutions and the mixture and quantitative information related to the quantity of precipitate produced upon combining the two solutions. Qualitative observations Both solutions are clear and colorless. There is no odor from the anion solution. When mixed, a white precipitate is formed. No bubbles/gas is produced. Quantitative observations When the two solutions are mixed in test tubes so that the ratio of the cation and anion are varied in a systematic manner the quantity of precipitate should be greatest in the tube with a 3:2 ratio of MClx:NazY. Excellent Student Results Student included a range of qualitative observations and reasoning based on them such as; Clear MClx solution indicates the absence of Ag+ since AgCl is insoluble. Lack of odor in NazY solution indicates the absence of S2-. Appearance of precipitate indicates the absence of K+ and NO3- ions since all their compounds are soluble. Lack of bubbles in NazY solution and upon mixing indicates absence of CO32-. Possible cations are Zn2+ and Al3+ while the anion is most likely PO43-. Student provided a clear explanation of the variation of the number of drops to determine the stoichiometry ratio of MClx:NazY. Student gave a clear data table with several trials to demonstrate the 3:2 ratio of MClx:NazY. Student identified the cation as Zn2+ and the anion as PO43-. Average Student Results Page 8 Not valid for use as an USNCO National Examination after May 1, 2007. Some qualitative information is given to demonstrate a knowledge of solubility and precipitate formation. Student provided evidence of several combinations of the two solutions and may have inferred something about the relationship between the solution ratio and identity of the salts. Below Average Student Results Little or no qualitative information was reported or used to make predictions about the identity of the unknown cation and anion. Student either did not report any quantitative information or was unable to use the quantitative information acquired to infer any information about the reaction stoichiometry from it. ######################### Lab Problem #2 Excellent Students Results: Student proposed a clear, detailed procedure for determining the thickness of the LDPE sheet, recognizing that measuring the volume of such a sheet directly would not be possible because of the small volume. Excellent procedures invariably involved measuring the density of the plastic; good methods included making a series of ethanol-water mixtures and interpolating the mixture of neutral buoyancy, or starting with one liquid and adding the other until neutral buoyancy was achieved. Density of the neutrally buoyant liquid was measured either by using the weighted average of ethanol or water, or by direct measurement of the mass of a known volume of the liquid. Student performed several buoyancy trials, either using a variety of water-ethanol volume ratios in a series of standards, or by redetermining the point of neutral buoyancy. Results were clearly displayed in a data table. Area and mass of LDPE piece(s) were measured in duplicate. Calculations are clearly shown using proper unit measurements and significant figures in final answers. Student demonstrated knowledge of the assumptions used in calculation (for example, the assumption of additive volumes if density of the neutrally buoyant mixture was calculated rather than measured directly). Final value for thickness was within 20% of the accepted value. Average Student Results: Measurement of density was proposed, but not clearly thought out; or, less precise procedures for determining volume directly (e.g., by displacement of liquid in the graduated cylinder) were proposed. Student made only qualitative (floats in water, sinks in ethanol) or grossly erroneous measurements of density. Only one trial was performed. Final value for thickness was within 40% of the accepted value. Below Average Student Results: Procedure was vague or unintelligible. Calculations were unclear or in error. Final value for thickness was over 40% off from the accepted value. Not valid for use as an USNCO National Examination after May 1, 2007 Page 9 2008 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM – PART 1 Prepared by the American Chemical Society Olympiad Examinations Task Force OLYMPIAD EXAMINATIONS TASK FORCE Arden P. Zipp, Chair, State University of New York, Cortland Sherry Berman-Robinson, Consolidated HS, Orland Park, IL (retired) William Bond, Snohomish HS, Snohomish, WA David Hostage, Taft School, Watertown, CT Peter Demmin, Amherst HS, Amherst, NY (retired) Marian Dewane, Centennial HS, Boise, ID Valerie Ferguson, Moore HS, Moore, OK Paul Groves, South Pasadena HS, Pasadena, CA Adele Mouakad, St. John’s School, San Juan, PR Jane Nagurney, Scranton Preparatory School, Scranton, PA Ronald Ragsdale, University of Utah, Salt Lake City, UT Kimberly Gardner, US Air Force Academy, Colorado Springs, CO DIRECTIONS TO THE EXAMINER–PART I Part I of this test is designed to be taken with a Scantron® answer sheet on which the student records his or her responses. Only this Scantron sheet is graded for a score on Part I. Testing materials, scratch paper, and the Scantron sheet should be made available to the student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until April 23, 2008, after which tests can be returned to students and their teachers for further study. Allow time for the student to read the directions, ask questions, and fill in the requested information on the Scantron sheet. The answer sheet must be completed using a pencil, not pen. When the student has completed Part I, or after one hour and thirty minutes has elapsed, the student must turn in the Scantron sheet, Part I of the testing materials, and all scratch paper. There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and you are free to schedule rest-breaks between parts. Part I Part II Part III 60 questions 8 questions 2 lab problems single-answer multiple-choice problem-solving, explanations laboratory practical 1 hour, 30 minutes 1 hour, 45 minutes 1 hour, 30 minutes A periodic table and other useful information are provided on page 2 for student reference. Students should be permitted to use nonprogrammable calculators. DIRECTIONS TO THE EXAMINEE–PART I DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Answers to questions in Part I must be entered on a Scantron answer sheet to be scored. Be sure to write your name on the answer sheet; an ID number is already entered for you. Make a record of this ID number because you will use the same number on both Parts II and III. Each item in Part I consists of a question or an incomplete statement that is followed by four possible choices. Select the single choice that best answers the question or completes the statement. Then use a pencil to blacken the space on your answer sheet next to the same letter as your choice. You may write on the examination, but the test booklet will not be used for grading. Scores are based on the number of correct responses. When you complete Part I (or at the end of one hour and 30 minutes), you must turn in all testing materials, scratch paper, and your Scantron answer sheet. Do not forget to turn in your U.S. citizenship statement before leaving the testing site today. Not valid for use as an USNCO Olympiad National Exam after April 23, 2008. Distributed by the ACS DivCHED Examinations Institute, University of Wisconsin - Milwaukee, Milwaukee, WI. All rights reserved. Printed in U.S.A. ABBREVIATIONS AND SYMBOLS A Faraday constant F molal atm formula molar mass M molar u free energy G molar mass A frequency ν mole NA gas constant R Planck’s constant °C gram g pressure c heat capacity Cp rate constant C hour h retention factor E joule J second Ea kelvin K temperature, K H kilo– prefix k time S liter L volt K milli– prefix m ampere atmosphere atomic mass unit atomic molar mass Avogadro constant Celsius temperature centi– prefix coulomb electromotive force energy of activation enthalpy entropy equilibrium constant CONSTANTS m M M mol h P k Rf s T t V R = 8.314 J·mol–1·K–1 R = 0.0821 L·atm·mol–1·K–1 1 F = 96,500 C·mol–1 1 F = 96,500 J·V–1·mol–1 NA = 6.022 × 1023 mol–1 h = 6.626 × 10–34 J·s c = 2.998 × 108 m·s–1 0 °C = 273.15 K 1 atm = 760 mmHg EQUATIONS E = Eo ! 1 1A 1 H RT ln Q nF "k % E " 1 1 % ln$$ 2 '' = a $$ ( '' # k1 & R # T1 T2 & $ "#H '$ 1 ' ln K = & )& ) + constant % R (% T ( PERIODIC TABLE OF THE ELEMENTS 18 8A 2 He 3 Li 2 2A 4 Be 13 3A 5 B 14 4A 6 C 15 5A 7 N 16 6A 8 O 17 7A 9 F 6.941 9.012 10.81 12.01 14.01 16.00 19.00 20.18 11 Na 12 Mg 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar 22.99 24.31 26.98 28.09 30.97 32.07 35.45 39.95 19 K 31 Ga 32 Ge 33 As 34 Se 35 Br 36 Kr 1.008 4.003 10 Ne 20 Ca 3 3B 21 Sc 4 4B 22 Ti 5 5B 23 V 6 6B 24 Cr 7 7B 25 Mn 8 8B 26 Fe 9 8B 27 Co 10 8B 28 Ni 11 1B 29 Cu 12 2B 30 Zn 39.10 40.08 44.96 47.88 50.94 52.00 54.94 55.85 58.93 58.69 63.55 65.39 69.72 72.61 74.92 78.96 79.90 83.80 37 Rb 38 Sr 39 Y 40 Zr 41 Nb 42 Mo 43 Tc 44 Ru 45 Rh 46 Pd 47 Ag 48 Cd 49 In 50 Sn 51 Sb 52 Te 53 I 54 Xe 85.47 87.62 88.91 91.22 92.91 95.94 (98) 101.1 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.3 55 Cs 56 Ba 57 La 72 Hf 73 Ta 74 W 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 132.9 137.3 138.9 178.5 180.9 183.8 186.2 190.2 192.2 195.1 197.0 200.6 204.4 207.2 209.0 (209) (210) (222) 87 Fr 88 Ra 89 Ac 104 Rf 105 Db 106 Sg 107 Bh 108 Hs 109 Mt 110 Ds 111 Rg 112 Uub 114 Uuq 116 Uuh 118 Uuo (223) (226) (227) (261) (262) (263) (262) (265) (266) (269) (272) (277) (2??) (2??) (2??) Page 2 ! 58 Ce 59 Pr 60 Nd 61 Pm 62 Sm 63 Eu 64 Gd 65 Tb 66 Dy 67 Ho 68 Er 69 Tm 70 Yb 71 Lu 140.1 140.9 144.2 (145) 150.4 152.0 157.3 158.9 162.5 164.9 167.3 168.9 173.0 175.0 90 Th 91 Pa 92 U 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr 232.0 231.0 238.0 (237) (244) (243) (247) (247) (251) (252) (257) (258) (259) (262) Not valid for use as a US National Chemistry Olympiad exam after April 23, 2008. DIRECTIONS When you have selected your answer to each question, blacken the corresponding space on the answer sheet using a soft, #2 pencil. Make a heavy, full mark, but no stray marks. If you decide to change an answer, erase the unwanted mark very carefully. There is only one correct answer to each question. Any questions for which more than one response has been blackened will not be counted. Your score is based solely on the number of questions you answer correctly. It is to your advantage to answer every question. 1. Which substance has the highest melting point? (A) Li2O (B) MgO 2. Which reagents produce a gas when combined? (C) CO2 (D) N2O5 I. HCl and Na2SO3 II. NaOH and Al (A) I only (B) II only (C) Both I and II (D) Neither I nor II 6. A NaOH solution is to be standardized by titrating it against a known mass of potassium hydrogen phthalate. Which procedure will give a molarity of NaOH that is too low? (A) Deliberately weighing one half the recommended amount of potassium hydrogen phthalate. (B) Dissolving the potassium hydrogen phthalate in more water than is recommended. (C) Neglecting to fill the tip of the buret with NaOH solution before titrating. 3. A 1:1 mixture of pentane and hexane is separated by fractional distillation in the apparatus shown. At what temperature does the first drop of condensate appear on the thermometer? (D) Losing some of the potassium hydrogen phthalate solution from the flask before titrating. 7. Which solute is least soluble in water? (A) 1-butanol (B) ethanol (C) methanol (D) 1-propanol 8. The mass of a single molecule of an allotrope of sulfur is 3.20×10–22 g. How many sulfur atoms are present in a molecule of this allotrope? Boiling point / oC pentane 36 hexane 69 (A) 4 (A) less than 36 ˚C (B) 36 ˚C (C) between 36 ˚C and 69 ˚C (D) more than 69 ˚C 4. Which nitrogen halide is least stable thermodynamically? (A) NF3 (B) NCl3 (C) NBr3 (D) NI3 5. Cyclohexane and water can be separated by using a separatory funnel. Which property contributes to this separation? (A) Cyclohexane and water are immiscible. (B) Cyclohexane has a lower viscosity than water. (C) Cyclohexane has a greater molar mass than water. (D) Cyclohexane has a greater vapor pressure than water. (B) 6 (C) 8 (D) 12 9. 100. L of carbon dioxide measured at 740. mmHg and 50 ˚C is produced by the complete combustion of a sample of pentane. 2C5H12 + 16O2 r 10CO2 + 12H2O What mass of pentane reacted? (A) 342 g (B) 265 g (C) 64.4 g (D) 53.0 g 10. Which 0.10 M aqueous solution has the smallest change in freezing point relative to pure water? (A) HC2H3O2 (B) HCl (C) CaCl2 (D) AlCl3 11. Magnetite, Fe3O4, can be Molar Mass / g·mol–1 reduced to iron by heating Fe3O4 232 with carbon monoxide according to the equation: Fe3O4 + 4CO r 3Fe + 4CO2 What mass of Fe3O4 is required in order to obtain 5.0 kg of iron if the process is 88% efficient? (A) 6.1 kg Not valid for use as a US National Chemistry Olympiad exam after April 23, 2008. (B) 6.9 kg (C) 7.9 kg (D) 18 kg Page 3 12. 40.0 g of a solute is dissolved in 500. mL of a solvent to give a solution with a volume of 515 mL. The solvent has a density of 1.00 g/mL. Which statement about this solution is correct? (A) The molarity is greater than the molality. (B) The molarity is lower than the molality. (C) The molarity is the same as the molality. (D) The molarity and molality cannot be compared without knowing the solute. 13. In the graph, the natural log of the vapor pressures of two substances are plotted versus 1/T. What can be concluded about the relative enthalpies of vaporization (∆Hvap) of these substances? 18. The atoms in crystals of silver metal are arranged in a cubic closest packed structure. What is the unit cell in this structure? (A) body-centered cubic (B) face-centered cubic (C) hexagonal-close packed (D) simple cubic 19. Use the information provided to calculate the standard enthalpy of formation of acetylene, C2H2(g), in kJ·mol–1. C2H2(g) + 5/2O2(g) r 2CO2(g) + H2O(l) ∆H˚ = –1299.5 kJ C(s) + O2(g) r CO2(g) ∆H˚ = –393.5 kJ H2(g) + 1/2O2(g) r H2O(l) ∆H˚ = –285.8 kJ (A) –1978.8 (B) –1121.4 (C) 226.7 (D) 453.4 20. Which statement is always true for a spontaneous reaction? (A) The entropy change for the system is negative. (B) The enthalpy change for the system is negative. (C) The entropy change for the universe is positive. (A) ∆Hvap of I is greater than ∆Hvap of II (D) The free energy change for the system is positive. (B) ∆Hvap of I is less than ∆Hvap of II 21. The heat of a reaction is measured in a bomb calorimeter. This heat is equal to which thermodynamic quantity? (C) ∆Hvap of I is is equal to ∆Hvap of II (D) No conclusion can be drawn from this information alone. 14. For which two gases are the rates of effusion 2:1? (A) H2 and He (B) He and O2 (C) Ne and Kr (D) N2 and Ar 15. Which gas has a density of 0.71 g·L–1 at 0 o C and 1 atm? (A) Ar (B) Ne (C) CO (D) CH4 16. Supercritical carbon dioxide exists at which point on the accompanying phase diagram? (A) ∆E (B) ∆G (C) ∆H (D) ∆S 22. 84.12 g of gold at Specific heat capacities / J.g–1.˚C–1 120.1 ˚C is placed Au(s) 0.129 in 106.4 g of H2O H2O(l) 4.184 at 21.4 ˚C. What is the final temperature of this system? (A) 70.8 (B) 65.0 (C) 27.8 (D) 23.7 23. In order to calculate the lattice energy of NaCl using a Born-Haber cycle, which value is not needed? (A) enthalpy of sublimation of Na(s) (B) first ionization energy of Cl(g) (C) bond dissociation energy of Cl2(g) (D) enthalpy of formation of NaCl(s) (A) A (B) B (C) C 17. Which properties increase with an increase in intermolecular forces at 25 ˚C? (D) D I. surface tension II. vapor pressure (A) I only (B) II only (C) Both I and II (D) Neither I nor II Page 4 24. Liquid bromine boils at 332.7 K. Estimate the enthalpy of formation of Br2(g) in kJ·mol–1. (A) 7.40 (B) 12.1 S˚ / J .mol–1.K–1 Br2(g) 58.6 Br2(l) 36.4 (C) 19.5 (D) 22.2 Not valid for use as a US National Chemistry Olympiad exam after April 23, 2008. 25. A student analyzed the data from a zero order reaction and obtained the graph shown. What labels should be attached to the X and Y axes, respectively? (A) time, concentration (B) time, 1 / concentration (C) time, ln (concentration) (D) 1/time, concentration 26. Under certain conditions the reaction of CO with NO2 to give CO2 and NO results in the rate law: rate = k[CO][NO2]. What are the units for the rate constant, k? (A) mol.L–1.min–1 (B) L.mol–1.min–1 (C) mol2.L–2.min–1 (D) L2.mol–2.min–1 27. For the reaction: X + Y r Z, initial rate data are given in the table. [X] / M [Y] / M Rate / mol.L–1.s–1 0.10 0.10 0.020 0.10 0.20 0.080 0.30 0.30 0.540 What is the rate law for this reaction? (A) Rate = k[X]2 (B) Rate = k[Y]2 (C) Rate = k[X][Y] (D) Rate = k[X][Y]2 28. The rate of the reaction of chlorine gas with a liquid hydrocarbon can be increased by all of the changes except one. Which change will be ineffective? (A) Use UV light to dissociate the Cl2. 30. For the reaction; A r B, the rate law is rate = k[A]. If the reaction is 40.0% complete after 50.0 minutes, what is the value of the rate constant, k? (A) 8.00×10–3 min–1 (B) 1.02×10–2 min–1 (C) 1.39×10–2 min–1 (D) 1.83×10–2 min–1 31. When 2.00 mol each of H2(g) and I2(g) are reacted in a 1.00 L container at a certain temperature, 3.50 mol of HI is present at equilibrium. Calculate the value of the equilibrium constant, Kc. (A) 3.7 (B) 14 (C) 56 (D) 2.0×102 32. For which equation is the equilibrium constant equal to Ka for the ammonium ion, NH4+? (A) NH4+(aq) + OH–(aq) s NH3(aq) + H2O(l) (B) NH4+(aq) + H2O(l) s NH3(aq) + H3O+(aq) (C) NH3(aq) + H2O(l) s NH4+(aq) + OH–(aq) (D) NH3(aq) + H3O+(aq) s NH4+(aq) + H2O(l) 33. What is the pH of a solution prepared by mixing 45.0 mL of 0.184 M KOH with 65.0 mL of 0.145 M HCl? (A) 1.07 (B) 1.13 (C) 1.98 (D) 2.92 34. The gas phase reaction shown is endothermic as written. Which change(s) will increase the quantity of CH3CH=CH2 at equilibrium? I. increasing the temperature II. increasing the pressure (B) Increase temperature at constant pressure. (A) I only (B) II only (C) Divide the liquid into small droplets. (C) Both I and II (D) Neither I nor II (D) Double the pressure by adding He gas. 29. One proposed mechanism of the reaction of HBr with O2 is given here. HBr + O2 r HOOBr (slow) HOOBr + HBr r 2HOBr (fast) HOBr + HBr r H2O + Br2 (fast) What is the equation for the overall reaction? (A) HBr + O2 r HOOBr (B) 2HBr + O2 r Br2 + H2O2 (C) 4HBr + O2 r 2H2O + 2Br2 35. The curve represents the titration of a weak monoprotic acid. Over what pH range(s) will the acid being titrated serve as a buffer when mixed with its salt? I. pH 4 – 6 II. pH 7 – 9 III. pH 12 – 13 (A) I only (B) II only (C) I and III only (D) I, II and III (D) 2HOBr r 2H2O + Br2 Not valid for use as a US National Chemistry Olympiad exam after April 23, 2008. Page 5 36. The pH of a saturated solution of Fe(OH)2 is 8.67. What is the Ksp for Fe(OH)2? (A) 5×10–6 (B) 2×10–11 (C) 1×10–16 (D) 5×10–17 37. In an operating voltaic cell electrons move through the external circuit and ions move through the electrolyte solution. Which statement describes these movements? (A) Electrons and negative ions both move toward the anode. (B) Electrons and negative ions both move toward the cathode. (C) Electrons move toward the anode and negative ions move toward the cathode. (D) Electrons move toward the cathode and negative ions move toward the anode. 38. The reduction potentials I. Ao reduces B2+ for the +2 cations, II. B2+ oxidizes Co 2+ – o III. Bo oxidizes Do e.g. A + 2e r A , of four metals decrease in the order A, B, C, D. Which statement(s) is/are true? (A) II only (B) III only (C) I and II only (D) I and III only Questions 39 and 40 should be answered with reference to the reaction: 2Ag+(aq) + M(s) r M2+(aq) + 2Ag E˚ = 0.940 V 42. A 3.00 amp current is used to electrolyze the molten chlorides; CaCl2, MgCl2, AlCl3, and FeCl3. The deposition of which mass of metal will require the longest electrolysis time? (A) 100 g Ca (B) 50 g Mg (C) 75 g Al (D) 125 g Fe 43. Which set of quantum numbers corresponds to an electron in a 4d orbital? (A) n = 4, l = 1, ml = –1, ms = 1/2 (B) n = 4, l = 2, ml = –2, ms = –1/2 (C) n = 4, l = 3, ml = 3, ms = 1/2 (D) n = 4, l = 3, ml = –1, ms = –1/2 44. What is the energy of a photon from a laser that emits light at 632.8 nm? (A) 3.14×10–19 J (B) 1.26×10–31 J (C) 2.52×10–33 J (D) 4.19×10–40 J 45. How many unpaired electrons are in a gaseous Co2+ ion in its ground state? (A) 1 (B) 3 (C) 5 (D) 7 46. Which ion is not isoelectronic with Ar? (A) S2– (B) K+ (C) Sc2+ (D) Ti4+ 47. Which process releases the most energy? (A) Mg2+ (g) + e– r Mg+(g) 39. What is the value of E˚ for the half reaction, 2+ E˚ / V Ag (aq) + e r Ag(s) 0.799 + – – M (aq) + 2e r M(s)? (A) 0.658 V (B) 0.141 V (C) –0.141 V (D) –0.658 V 40. Which change will cause the largest increase in the voltage of a cell based on the reaction above? (A) Doubling the [Ag+] from 1M to 2M + (C) Doubling the volume of the 1M Ag solution (D) Reducing the [M2+] from 1M to 0.5M 41. If a voltaic cell has a positive Eo value, what can be concluded about the values of ∆Go and Keq? o (B) ∆G < 0, Keq > 1 o (D) ∆Go > 0, Keq > 1 (C) ∆G > 0, Keq < 1 Page 6 (C) Na2+(g) + e– r Na+(g) (D) Na+(g) + e– r Na(g) 48. In which list are the ions arranged in order of increasing size? (A) F– < S2– < Al3+ < Mg2+ (B) F– < S2– < Mg2+ < Al3+ (C) Mg2+ < F– < Al3+ < S2– (B) Doubling the amount of M(s) (A) ∆G < 0, Keq < 1 (B) Mg+(g) + e– r Mg(g) o (D) Al3+ < Mg2+ < F– < S2– 49. Molecules with non-zero dipole moments include which of those listed? I. H2C=CHCl II. cis - ClHC=CHCl III. trans - ClHC=CHCl (A) I only (B) III only (C) I and II only (D) I, II and III Not valid for use as a US National Chemistry Olympiad exam after April 23, 2008. 50. Which species is diamagnetic? (B) N2+ (A) NO (D) O22– (C) O2 51. What is the I-I-I bond angle in the I3– ion? (A) 180o (B) 120 o (A) identity of the monomers in the two polymers. (C) 90 o o (D) more than 90 but less than 120 (B) number of monomer units in the two polymers. o (C) orientation of the bonds joining the monomers. 52. Which species has the shortest nitrogen-oxygen bond? (A) NO+ (B) NO2+ 60. Cellulose and starch are biological polymers. Humans are able to digest starch but not cellulose. This difference is due primarily to a difference in the (C) NO2– (D) percentage of carbon in the two polymers. (D) NO3– 53. Which substance will form hydrogen bonds to water molecules but will not form hydrogen bonds with its own molecules? (A) HF (B) C2H5OH (C) CH3NH2 (D) CH3OCH3 END OF TEST 54. In the gas phase PCl5 exists as individual molecules but in the solid it takes on the ionic structure PCl4+PCl6–. What are the geometries of these three species PCl4+ PCl5 PCl6– (A) trigonal bipyramidal see-saw octahedral (B) trigonal bipyramidal tetrahedral octahedral (C) trigonal bipyramidal square planar distorted octahedral (D) square pyramidal see-saw square planar 55. Which molecule contains exactly eight carbon atoms? (A) benzoic acid (B) 2,3-dimethylhexane (C) 3-ethylpentane (D) 3-methyloctane 56. Which formula represents an alkyne? (Assume all are noncyclic.) (A) C2H2 (B) C2H4 (C) C5H10 (D) C8H18 57. How many compounds have the formula C2H3Cl3? (A) 2 (B) 3 (C) 4 (D) 5 58. Which is a condensation polymer? (A) polyethylene (B) polyvinylchloride (C) polystyrene (D) polyethylene terephthalate 59. What is the number of pi (π) bonds in trans-butenedioic acid (C4H4O4)? (A) 1 (B) 2 (C) 3 (D) 4 Not valid for use as a US National Chemistry Olympiad exam after April 23, 2008. Page 7 Olympiad 2008 National Part I KEY Number 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. Answer B C B D A C A B D A C B A C D C A B C C A D B A A B D D C B Number 31. 32. 33. 34. 35. 36. 37. 38. 39. 40. 41. 42. 43. 44. 45. 46. 47. 48. 49. 50. 51. 52. 53. 54. 55. 56. 57. 58. 59. 60. Answer D B C D A D D A C A B C B A B C C D C D A A D B B A A D C C Not valid for use as an USNCO Olympiad National Exam after April 23, 2008. 2008 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART II Prepared by the American Chemical Society Olympiad Examinations Task Force OLYMPIAD EXAMINATIONS TASK FORCE Arden P. Zipp, State University of New York, Cortland Chair Sherry Berman-Robinson, Consolidated HS, Orland Park, IL (retired) Paul Groves, South Pasadena HS, Pasadena, CA William Bond, Snohomish HS, Snohomish, WA David Hostage, Taft School, Watertown, CT Peter Demmin, Amherst HS, Amherst, NY (retired) Marian Dewane, Centennial HS, Boise, ID Valerie Ferguson, Moore HS, Moore, OK Adele Mouakad, St. John’s School, San Juan, PR Jane Nagurney, Scranton Preparatory School, Scranton, PA Ronald Ragsdale, University of Utah, Salt Lake City, UT Kimberly Gardner, US Air Force Academy, Colorado Springs, CO DIRECTIONS TO THE EXAMINER–PART II Part II of this test requires that student answers be written in a response booklet of blank pages. Only this “Blue Book” is graded for a score on Part II. Testing materials, scratch paper, and the “Blue Book” should be made available to the student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until April 23, 2008, after which tests can be returned to students and their teachers for further study. Allow time for the student to read the directions, ask questions, and fill in the requested information on the “Blue Book”. When the student has completed Part II, or after one hour and forty-five minutes has elapsed, the student must turn in the “Blue Book”, Part II of the testing materials, and all scratch paper. Be sure that the student has supplied all of the information requested on the front of the “Blue Book,” and that the same identification number used for Part I has been used again for Part II. There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and you are free to schedule rest-breaks between parts. Part I Part II Part III 60 questions 8 questions 2 lab problems single-answer multiple-choice problem-solving, explanations laboratory practical 1 hour, 30 minutes 1 hour, 45 minutes 1 hour, 30 minutes A periodic table and other useful information are provided on the back page for student reference. Students should be permitted to use non-programmable calculators. DIRECTIONS TO THE EXAMINEE–PART II DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Part II requires complete responses to questions involving problem-solving and explanations. One hour and forty-five minutes are allowed to complete this part. Be sure to print your name, the name of your school, and your identification number in the spaces provided on the “Blue Book” cover. (Be sure to use the same identification number that was coded onto your Scantron® sheet for Part I.) Answer all of the questions in order, and use both sides of the paper. Do not remove the staple. Use separate sheets for scratch paper and do not attach your scratch paper to this examination. When you complete Part II (or at the end of one hour and forty-five minutes), you must turn in all testing materials, scratch paper, and your “Blue Book.” Do not forget to turn in your U.S. citizenship statement before leaving the testing site today. Not valid for use as an USNCO Olympiad National Exam after April 23, 2008. Page 1 ! ! ! ! ! ! Key for 2008 National Olympiad (part 2) 1. (14%) Benzene, C6H6, reacts with Br2 in the presence of FeBr3 as a catalyst to give an organic compound with the percentage composition by mass; C 30.55%, H 1.71%, Br 67.74% and hydrogen bromide. a. Determine the empirical formula of the compound. b. When 0.115 g of this compound are dissolved in 4.36 g of naphthalene the solution freezes at 79.51 ˚C. Pure naphthalene freezes at 80.29 ˚C and has a kf = 6.94 ˚C·m–1. Determine the molar mass and molecular formula of the compound. c. Write a balanced equation for the reaction. d. Calculate the theoretical yield for the organic compound when 4.33 g of of benzene is reacted with an excess of bromine. e. If the actual yield of the reaction is 5.67 g, what is the percentage yield? f. i. Write structures for the possible isomers that could be formed in this reaction. ii. Identify the major isomer(s) formed in this reaction and explain your reasoning. a) convert masses to moles: # 1 mol & 1.71 g H " % ( = 1.70 mol (÷0.848) = 2.00 $ 1.008 g ' # 1 mol & 30.55 g C " % ( = 2.54 mol (÷0.848) = 3.00 $ 12.011 g ' # 1 mol & 67.74 g Br " % ( = 0.848 mol (÷0.848) = 1.0 $ 79.90 g ' These numbers are whole numbers, so the empirical formula must be C3H2Br b) ΔT = 80.29 – 79.51 = 0.78 oC. Plugging this value into the formula for freezing point depression gives, "T = k f • m and m = 0.78 o C 6.94 o C / m = 0.11m mol 0.115 g 0.11 " 0.00436 kg = 0.00048 mol so, MM = = 240 g # mol-1 kg 0.00048 mol 240 / 117.9 = 2.03 which is approximately 2, so the molecular formula must be C6H4Br2 c) C6H6 + 2Br2 r C6H4Br2 + 2 HBr d) The theoretical yield is: # 1 mol C H & # 1 mol C H Br & # 235.89 g C H Br & 6 6 6 4 2 6 4 2 4.33 g C 6 H 6 " % ( "% (" % ( = 13.1 g C 6 H 4 Br2 $ 78.11 g C 6 H 6 ' $ 1 mol C 6 H 6 ' $ 1 mol C 6 H 4 Br2 ' " 5.67 g % Percent yield is: $ ' ( 100% = 43.4% # 13.07 g & The possible isomers for i. are, e) Br ! Br Br Br ii. the major products are, Br Br Br Br Br Page 2 Br because –Br is an ortho-para director. The para isomer should be the most prominent product because of steric hindrance for the ortho product. Not valid for use as an USNCO Olympiad National Exam after April 23, 2008. ! ! ! ! ! ! ! 2. (10%) Photochemical smog is formed through a sequence of reactions, the first three of which are given below. Smog is formed when the O(g) produced in reaction (3) reacts with organic molecules. Bond Dissociation Energy, kJ⋅mol–1 (1) N2(g) + O2(g) r 2NO(g) N–N 193 (2) 2NO(g) + O2(g) r 2NO2(g) N=N 418 (3) NO2(g) + hν r NO(g) + O(g) N≡N 941 a. For reaction (1), ∆Hº = +180.6 kJ·mol–1. Calculate the bond O–O 142 dissociation energy of NO(g). O=O 498 b. Calculate the entropy change for the first reaction. So, J⋅mol–1⋅K–1 c. Determine the minimum temperature at which reaction (1) becomes N2(g) 191.5 spontaneous. O2(g) 205.0 d. For reaction (3), ∆Hº = +306 kJ·mol–1. If the energy for this NO(g) 210.6 reaction were provided by sunlight, estimate the wavelength NO2(g) 240.5 required and specify the region of the spectrum containing this O(g) 161.0 wavelength. a) "H = The overall enthalpy change can be estimated from the bond dissociation energies via the equation, # Energy of bonds broken $ # Energy of bonds formed 180.6 kJ = 941 kJ + 498 kJ " 2 # BDE NO so, BDE NO = 629 kJ " mol–1 b) Similarly, ( ) "S o = 2S o (NO) # S o (N 2 ) + S o (O 2 ) "S o = 2(210.6) # ( (191.5) + (205)) = 24.7 J $ mol-1 $ K -1 c) Utilize the equation, "G o = "H o # T"S o , and set ΔGo to zero to find the minimum temperature. ( ) 0 = 180.6 kJ " mol-1 # T $ 0.0247 kJ " mol-1 " K -1 , so T = 180.6 kJ " mol-1 0.0247 kJ " mol-1 " K -1 = 7311 K d) First convert to energy per molecule, ! # & J 1 mol J -19 3.06 " 10 3 "% . Now calculate wavelength of light with this energy, ( = 5.08 " 10 23 mol $ 6.022 " 10 molecules ' molecule -34 8 -1 % h # c ' 6.626 $ 10 J #s 3.0 $ 10 m #s " = = ' E 5.08 $ 10 -19 J # molecule & ( )( ) (* = 3.91$ 10 -7 * ) m (per molecule) = 391 nm (in the ultraviolet). (12%) Aniline, C6H5NH2, reacts with water according to the equation: C6H5NH2(aq) + H2O(l) s C6H5NH3+(aq) + OH –(aq) 3. In a 0.180 M aqueous aniline solution the [OH–] = 8.80×10–6. ! a. b. c. d. Write the equilibrium constant expression for this reaction. Determine the value of the base ionization constant, Kb, for C6H5NH2(aq). Calculate the percent ionization of C6H5NH2 in this solution. Determine the value of the equilibrium constant for the neutralization reaction; C6H5NH2(aq) + H3O+(aq) s C6H5NH3+(aq) + H2O(l) e. i. Find the [C6H5NH3+(aq)] / [C6H5NH2(aq)] required to produce a pH of 7.75. ii. Calculate the volume of 0.050M HCl that must be added to 250.0 mL of 0.180 M C6H5NH2(aq) to achieve this ratio. a) K b = [C 6H 5NH +3 ][OH -] [C 6H 5NH 2 ] Not valid for use as an USNCO Olympiad National Exam after April 23, 2008. ! Page 3 Key for 2008 National Olympiad (part 2) (8.80 " 10 )(8.80 " 10 ) = 4.3" 10 #6 b) K b = #6 #10 ( 0.180) (8.80 " 10 ) " 100% = 4.9 " 10 #6 c) % ionization = ! #3 ( 0.180) % d) C 6 H 5 NH 2 + H 3O + " C 6 H 5 NH +3 + H 2O so K = ! K b 4.3# 10$10 = = 4.3# 10 4 K w 1.0 # 10$14 e) (i) For a pH = 7.75, the pOH = 6.25 so [OH - ] = 10"pOH = 5.62 # 10"7 M. ! 4.3 " 10#10 = [C 6H 5NH +3 ][OH -]!so, [C 6H 5NH +3 ] = 4.3 " 10#10 = 7.65 " 10#4 [C 6H 5NH 2 ] [C 6H 5NH 2 ] 5.62 " 10#7 (ii) The HCl is a strong acid that will protonate the aniline, so to get the HCl required, we need the amount of C6H5NH2 required multiplied by the value of the ratio from (i): 7.65" 10#4 " 0.250 L " 0.180M C 6 H 5 NH 2 = 3.44 " 10#5 mol HCl Now determine the volume ! of reagent: 3.44 " 10#5 mol HCl " 1 L 0.050 mol HCl = 6.88 " 10#4 L = 0.688 mL ! 4. ! (10%) Gaseous dinitrogen pentoxide, N2O5, decomposes to form nitrogen dioxide and oxygen gas with the initial rate data at 25 ˚C given in the table.! [N2O5], M Rate, mol.L–1.min–1 a. b. c. d. 0.150 3.42×10–4 0.350 7.98×10–4 0.650 1.48×10–3 Write a balanced equation for this reaction. Use the data provided to write the rate law and calculate the value of k for this reaction. Show all calculations. Calculate the time required for the concentration of a 0.150 M sample of N2O5 to decrease to 0.050 M. The initial rate for the reaction of a 0.150 M sample is 2.37×10–3 mol.L–1.min–1 at 40 ˚C. Determine the activation energy for this reaction. a) 2N 2O 5 " 4NO 2 + O 2 ! ! ! ! ! b) The rate will be given by the rate law (Rate=k[N2O5]x) in each case, so by taking the ratio, the rate constant cancels, x Rate1 [ N 2O 5 ]1 = Rate 2 [ N 2O 5 ] x 2 x x 7.98 " 10#4 $ 0.350 ' 1.48 " 10#3 $ 0.650 ' x means that 2.33 = 2.33 so x=1. Checking with a second set of data, = = & ) & ) leads to 3.42 " 10#4 % 0.150 ( 3.42 " 10#4 % 0.150 ( 4.33x = 4.33, confirming that the reaction is first order. Now calculate the rate constant: 1.48 " 10#3 = k(0.650)1 so k = 2.28 " 10 -3 min#1 and we have Rate = 2.28×10-3 min -1[N2O5]. ! " % ! law, ln$ [ N 2O 5 ] init ' = kt . Plugging in ln"$ 0.150 %' = ln(3) = (2.28 ( 10)3 min)1)t so t = 481 minutes. c) Use the integrated rate $ [N O ] ' # 0.050 & 2 5 t & # d) Use the information from the two temperatures given in the Arrhenus equation: " 1.58 ) 10 -2 % "k % E " 1 " 1 1 % Ea 1 % ln$ 2 ' = a $ ( ! ', so, plugging in values gives : !ln$ ) ( '= ' -3 -1 -1 $ # 298 K 313 K & # k1 & R # T1 T2 & # 2.28 ) 10 & 8.314 J * mol * K Ea so ln( 6.93) = # ( 0.00335 $ 0.00319) and solving for E a gives : E a = 1.00 # 10 5 J " mol-1 = 100 kJ " mol-1 . (Note 8.314 J " mol-1 " K -1 that using rates, rather than rate constants in the argument of the natural log is an alternative, correct method.) Page 2 Not valid for use as an USNCO Olympiad National Exam after April 23, 2008. 5. (12%) Write net equations for each of the combinations of reactants below. Use appropriate ionic and molecular formulas and omit formulas for all ions or molecules that do not take part in a reaction. Write structural formulas for all organic substances. You need not balance the equations. All reactions occur in aqueous solution unless otherwise indicated. a. Barium peroxide is added to water. b. Acidic solutions of potassium iodide and potassium iodate are mixed. c. A phosphoric acid solution is added to a solution of calcium hydrogencarbonate. d. Solutions of lead(II) nitrate and potassium chromate are mixed. e. Concentrated hydrochloric acid is added to an aqueous solution of cobalt(II) nitrate. f. 2-butanol is heated with concentrated sulfuric acid. a) BaO 2 + H 2O " Ba 2+ + HO -2 + OH b) I - + IO -3 + H + " I 2 + H 2O c) H 3 PO 4 + Ca 2+ + HCO"3 # Ca 3 (PO 4 ) 2 + H 2O + CO 2 ! ! ! ! ! d) Pb 2+ + CrO 24 " PbCrO 4 e) Co 2+ + Cl- " CoCl24 f) H3C CH2 OH C CH3 + H2SO4 H H3C CH C CH3 H + + H2O H3C CH2 C CH2 H either isomer counts 6. (12%) The apparatus depicted to the right is often used to demonstrate the electrolysis of water. Tubes A and B are initially filled with an aqueous solution of H2SO4 or Na2SO4. a. Describe the purpose of adding the H2SO4 or Na2SO4 rather than using pure water. b. Give the formula of the gas produced in; i. tube A ii. tube B c. Describe a chemical test that could be used to identify the gas collected in tube A. Include the procedure and expected observation. d. Calculate the number of moles of gas expected to be collected in tube B when a 600. milliamp current is applied for 40.0 minutes. (Assume no side reactions occur.) e. Calculate the volume of the gas produced in part d. for a temperature is 20 ˚C and a pressure in the laboratory of 735 mmHg. (The vapor pressure of water is 17.5 mmHg.) f. If H2O2 is formed in a side reaction the quantity of only one of the products is affected. Identify the product affected and state how its quantity compares with that produced with no side reaction. Explain your answer. a) Because pure water is a poor conductor of electricity, the H2 SO4 or Na2SO4 is added to provide electrolyte (so that the solution will conduct). b) i) tube A is the cathode, therefore it is the site of reduction where H2 is produced, while (ii) tube B is the anode, where oxidation occurs, therefore O2 is produced. c) Because H2 is flammable, a burning splint can be inserted into the products from Tube A. If there is a “pop” associated with the reaction, it confirms that the gas is H2. d) Charge = current × time: 0.600 C·s-1 × 2400 s = 1440 C # 1 mol e - & # 1 mol O & 2 = 3.73 " 10)3 mol O 2 and: 1440 C " % (" % - ( $ ' 96500 C 4 mol e $ ' Not valid for use as an USNCO Olympiad National Exam after April 23, 2008. ! Page 5 Key for 2008 National Olympiad (part 2) e) First correct for vapor pressure of water: Ptotal = PO 2 + PH 2 O so PO 2 = 735- 17.5 = 717.5 mmHg nRT (3.73 " 10#3 mol)(0.821 L $ atm $ mol-1 $ K -1)(293 K) = = 0.095 L = 95 mL 1 atm P 717.5 mmHg " ! 760 mm Hg f) The quantity of O2 would be affected, but the quantity of H2 would not. The yield of O2 would be decreased because some of the electricity would oxidize H2O into peroxide (H2O2) instead of O2. V= ! 7. (16%) Explain the following observations in terms of bonding principles. a. Carbon dioxide is a gas at room temperature and pressure but silicon dioxide is a high-melting solid. b. The xenon trioxide molecule has a trigonal pyramidal shape while sulfur trioxide is trigonal planar. c. In many of its ionic compounds oxygen is present as the O2– ion although the addition of two electrons to an oxygen atom in the gas phase is an endothermic process. d. (bmim)+PF6– is a liquid at room temperature while (bmim)+Cl– and Na+PF6– are solids. Note: (bmim)+ is an abbreviation for N-butyl-N-methylimadazolium ion, CH3N2C3H3C4H9+. a) Carbon dioxide is small, non-polar molecule. The intermolecular forces between them are small, so CO2 is a gas. Conversely, silicon dioxide is a network solid. As a network, the connections are covalent bonds, which are quite strong compared to the intermolecular forces between small molecules, and it requires a great deal of energy to break an SiO2 unit away from the rest of the solid. Ultimately the key bonding feature in these molecules that gives rise to this difference is that carbon atoms readily form double bonds, where double bonds to silicon are much less common. b) Looking at the Lewis structure of the two compounds provides the answer: O Xe O O O S O O There are four charge centers (three bonding, and one lone pair) around Xe in XeO3 leading to a trigonal pyramidal shape, while there are three charge centers (all of them bonding pairs) in SO3 which is trigonal planar. c) When the oxide ion, O2–, is in ionic compounds, the 2– charge is interacting with positively charged cations. Thus, even though the ion formation in the gas phase is endothermic, the oxide ion exists in ionic compounds. d) Because the (bmim)+ and PF6–ions are quite large, the lattice energy between the two items will be small. The energy available as heat at room temperature is sufficient to overcome this interaction energy. By contrast (bmim)+ and Cl–and Na+ and PF6– have one large and one small ion, so they can pack more closely and have larger lattice energy. These larger energies mean the compounds are solids at room temperature. 8. (14%) This question deals with the bonding in several organic chemicals. a. Several different compounds have the formula C2H4O2. Two of these contain –CO2 groups. i. Give the structures and names of the two compounds with –CO2 groups. ii. These compounds boil at 31.5 ˚C and 118 ˚C. Assign the two boiling points to the structures in i. and account for the boiling point difference in terms of their structures. iii. Sketch the structure of one of the other compounds. b. Fatty acids are important components of a healthy diet. Three fatty acids are stearic, oleic and linoleic which have the formulas CH3C16H32COOH, CH3C16H30COOH, and CH3C16H28COOH, respectively. i. Describe the differences in bonding suggested by the formulas of these compounds. ii. The compounds melt at –5 ˚C, 13 ˚C and 69 ˚C. Assign these melting points to the respective acids and account for this behavior in terms of their structures and bonding. iii. The salts of fatty acids can be used as soaps or detergents. Describe the chemical basis of this behavior. a) Page 2 (i) The two structures are: Not valid for use as an USNCO Olympiad National Exam after April 23, 2008. H H O C C O H OH C H O C H H H ethanoic acid methylmethanoate ii) 118 o C is ethanoic acid and 31.5 o C is methyl methanoate. The key difference arises from the strength of intermolecular forces present in ethanoic acid, which can participate in hydrogen bonding, while the strongest intermolecular forces present in methylmethanoate are dipole-dipole forces. iii) Possible correct structures include: H OH C H C H C OH HO C HO H C H OH O H C C C H H O H OH H b) (i) CH3C16H32COOH contains a saturated alkyl chain. CH3 C16H30COOH contains a one carbon-carbon double bond. CH3C16H28COOH contains two carbon-carbon double bonds. (ii) CH3C16H32COOH melts at 69 oC. It is the highest melting point because the saturated alkyl chain tails are capable of being closely packed, thereby maximizing the dispersion forces present. Higher intermolecular forces lead to higher melting points. CH3C16H30COOH with one double bond has additional geometrical constraints due to the relative rigidity of that double bond, so the tails cannot pack as efficiently, and the melting point is lower, at 13 oC. Finally, for CH3C16H28COOH with two double bonds, the geometric constraints just noted are even more sizable, so packing is even less efficient. It will, therefore, have the lowest melting point, –5 oC. (iii) The key feature is that the molecules have a charged region (often called the head) where the acid group is and an uncharged and not very polar region (called the tail) where the alkyl chains are. The non-polar tail can interact relatively strongly with non-polar dirts, oils and greases, leaving the polar/charged head group “sticking out”. This polar/charged group interacts strongly with polar water molecules. Thus, while polar water molecules do not wash away non-polar dirts and oils by themselves, taking advantage of the dual behavior of the long-chain fatty acids, the dirt/oil is encapsulated in a micelle-like structure that can be solvated by water. Pictorially, showing far too few fatty acids, it would look something like this. CO2- CO2- CO2CO2CO2CO2- CO2CO2- END OF PART II Not valid for use as an USNCO Olympiad National Exam after April 23, 2008. Page 7 KEY for 2008 National Olympiad (Part 2) amount of substance ampere atmosphere atomic mass unit atomic molar mass Avogadro constant Celsius temperature centi- prefix coulomb electromotive force energy of activation enthalpy entropy ABBREVIATIONS AND SYMBOLS n equilibrium constant K measure of pressure mmHg A Faraday constant F milli- prefix m atm formula molar mass M molal m u free energy G molar M A frequency ν mole mol NA gas constant h R Planck’s constant °C gram P g pressure c heat capacity k Cp rate constant C hour Rf h retention factor E joule s J second Ea kelvin c K speed of light H kilo- prefix T k temperature, K S liter t L time volt V CONSTANTS R = 8.314 J·mol–1·K–1 R = 0.0821 L·atm·mol–1·K–1 1 F = 96,500 C·mol–1 1 F = 96,500 J·V–1·mol–1 NA = 6.022 × 1023 mol–1 h = 6.626 × 10–34 J·s c = 2.998 × 108 m·s–1 USEFUL EQUATIONS E = E! – ! k2 $ Ea ! 1 1 $ = ' " k1 &% R #" T1 T2 &% " –!H % " 1 % ln K = $ ' $ ' +c # R & # T& RT ln Q nF ln # PERIODIC TABLE OF THE ELEMENTS 1 H 2 He 1.008 4.003 3 Li 4 Be 5 B 6 C 7 N 8 O 9 F 10 Ne 6.941 9.012 10.81 12.01 14.01 16.00 19.00 20.18 11 Na 12 Mg 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar 22.99 24.31 26.98 28.09 30.97 32.07 35.45 39.95 19 K 20 Ca 21 Sc 22 Ti 23 V 24 Cr 25 Mn 26 Fe 27 Co 28 Ni 29 Cu 30 Zn 31 Ga 32 Ge 33 As 34 Se 35 Br 36 Kr 39.10 40.08 44.96 47.88 50.94 52.00 54.94 55.85 58.93 58.69 63.55 65.39 69.72 72.61 74.92 78.96 79.90 83.80 37 Rb 38 Sr 39 Y 40 Zr 41 Nb 42 Mo 43 Tc 44 Ru 45 Rh 46 Pd 47 Ag 48 Cd 49 In 50 Sn 51 Sb 52 Te 53 I 54 Xe 85.47 87.62 88.91 91.22 92.91 95.94 (98) 101.1 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.3 55 Cs 56 Ba 57 La 72 Hf 73 Ta 74 W 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 132.9 137.3 138.9 178.5 181.0 183.8 186.2 190.2 192.2 195.1 197.0 200.6 204.4 207.2 209.0 (209) (210) (222) 87 Fr 88 Ra 89 Ac 104 Rf 105 Db 106 Sg 107 Bh 108 Hs 109 Mt 110 Uun 111 Uuu 112 Uub 114 Uuq 116 Uuh 118 Uuo (223) (226) (227) (261) (262) (263) (262) (265) (266) (269) (272) (277) (2??) (2??) (2??) Page 8 58 Ce 59 Pr 60 Nd 61 Pm 62 Sm 63 Eu 64 Gd 65 Tb 66 Dy 67 Ho 68 Er 69 Tm 70 Yb 71 Lu 140.1 140.9 144.2 (145) 150.4 152.0 157.3 158.9 162.5 164.9 167.3 168.9 173.0 175.0 90 Th 91 Pa 92 U 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr 232.0 231.0 238.0 237.0 (244) (243) (247) (247) (251) (252) (257) (258) (259) (260) Not valid for use as an USNCO Olympiad National Exam after April 23, 2008. 2008 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART III Prepared by the American Chemical Society Olympiad Laboratory Practical Task Force OLYMPIAD LABORATORY PRACTICAL TASK FORCE Steve Lantos, Brookline High School, Brookline, MA Chair Linda Weber, Natick High School, Natick, MA John Mauch, Braintree High School, Braintree, MA Nancy Devino, ScienceMedia Inc., San Diego, CA Christie B. Summerlin, University of Alabama-Birmingham, Birmingham, AL DIRECTIONS TO THE EXAMINER–PART III The laboratory practical part of the National Olympiad Examination is designed to test skills related to the laboratory. Because the format of this part of the test is quite different from the first two parts, there is a separate, detailed set of instructions for the examiner. This gives explicit directions for setting up and administering the laboratory practical. There are two laboratory tasks to be completed during the 90 minutes allotted to this part of the test. Students do not need to stop between tasks, but are responsible for using the time in the best way possible. Each procedure must be approved for safety by the examiner before the student begins that procedure. Part III 2 lab problems laboratory practical 1 hour, 30 minutes Students should be permitted to use non-programmable calculators. DIRECTIONS TO THE EXAMINEE–PART III DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. WHEN DIRECTED, TURN TO PAGE 2 AND READ THE INTRODUCTION AND SAFETY CONSIDERATIONS CAREFULLY BEFORE YOU PROCEED. There are two laboratory-related tasks for you to complete during the next 90 minutes. There is no need to stop between tasks or to do them in the given order. Simply proceed at your own pace from one to the other, using your time productively. You are required to have a procedure for each problem approved for safety by an examiner before you carry out any experimentation on that problem. You are permitted to use a non-programmable calculator. At the end of the 90 minutes, all answer sheets should be turned in. Be sure that you have filled in all the required information at the top of each answer sheet. Carefully follow all directions from your examiner for safety procedures and the proper disposal of chemicals at your examining site. Not valid for use as an USNCO National Examination after April 23, 2008 Page 1 2007 UNITED STATES NATIONAL CHEMISTRY OLYMPIAD PART III — LABORATORY PRACTICAL Student Instructions Introduction These problems test your ability to design and carry out laboratory experiments and to draw conclusions from your experimental work. You will be graded on your experimental design, on your skills in data collection, and on the accuracy and precision of your results. Clarity of thinking and communication are also components of successful solutions to these problems, so make your written responses as clear and concise as possible. Safety Considerations You are required to wear approved eye protection at all times during this laboratory practical. You also must follow all directions given by your examiner for dealing with spills and with disposal of wastes. Lab Problem 1 You have been given seven pipets that contain solutions of AgNO3, BaCl2, Cu(NO3)2, CuSO4, Pb(NO3)2, KI, and Na2S2O3, though not necessarily in this order. Using the materials provided, devise and carry out an experiment to correctly determine the contents of each pipet. Lab Problem 2 Given a sample of an unknown metal carbonate, MxCO3, and 3.0M hydrochloric acid, HCl(aq), a balloon, and some laboratory equipment, devise and carry out an experiment by combining these two substances to determine the volume of the gas produced and the unknown metal. The possible metals are Ba, Ca, Li, or Na. Room Temp. = 25oC, Standard Pressure = 1 atm Answer Sheet for Laboratory Practical Problem 1 Page 2 Not valid for use as an USNCO National Examination after April 23, 2008. Student's Name: __________________________________________________________________________ Student's School:________________________________________ Date: ___________________________ Proctor's Name: _________________________________________________________________________ ACS Section Name :________________________________Student's USNCO test #: ________________ 1. Give a brief description of your experimental plan. Before beginning your experiment, you must get approval (for safety reasons) from the examiner. Not valid for use as an USNCO National Examination after April 23, 2008 Examiner’s Initials: Page 3 2. Record your data and other observations. 3. Based on your observations, write the relevant equations that led to your conclusions. 4. Conclusions Pipet Contents Justification #1 #2 #3 #4 #5 #6 #7 Page 4 Not valid for use as an USNCO National Examination after April 23, 2008. Answer Sheet for Laboratory Practical Problem 2 Student's Name: __________________________________________________________________________ Student's School:________________________________________ Date: ___________________________ Proctor's Name: _________________________________________________________________________ ACS Section Name :________________________________Student's USNCO test #: ________________ 1. Give a brief description of your experimental plan. Before beginning your experiment, you must get approval (for safety reasons) from the examiner. Not valid for use as an USNCO National Examination after April 23, 2008 Examiner’s Initials: Page 5 2. Record your data and other observations. 3. Calculations and Conclusions. 4. Conclusions: The volume of gas produced: The unknown metal: 5. Sources of Error in this experiment (please number): Page 6 Not valid for use as an USNCO National Examination after April 23, 2008. PERIODIC TABLE OF THE ELEMENTS 2 He 1 H 1.008 4.003 3 Li 4 Be 5 B 6 C 7 N 8 O 9 F 10 Ne 6.941 9.012 10.81 12.01 14.01 16.00 19.00 20.18 11 Na 12 Mg 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar 22.99 24.31 26.98 28.09 30.97 32.07 35.45 39.95 19 K 20 Ca 21 Sc 22 Ti 23 V 24 Cr 25 Mn 26 Fe 27 Co 28 Ni 29 Cu 30 Zn 31 Ga 32 Ge 33 As 34 Se 35 Br 36 Kr 39.10 40.08 44.96 47.88 50.94 52.00 54.94 55.85 58.93 58.69 63.55 65.39 69.72 72.61 74.92 78.96 79.90 83.80 37 Rb 38 Sr 39 Y 40 Zr 41 Nb 42 Mo 43 Tc 44 Ru 45 Rh 46 Pd 47 Ag 48 Cd 49 In 50 Sn 51 Sb 52 Te 53 I 54 Xe 85.47 87.62 88.91 91.22 92.91 95.94 (98) 101.1 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.3 55 Cs 56 Ba 57 La 72 Hf 73 Ta 74 W 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 132.9 137.3 138.9 178.5 181.0 183.8 186.2 190.2 192.2 195.1 197.0 200.6 204.4 207.2 209.0 (209) (210) (222) 87 Fr 88 Ra 89 Ac 104 Rf 105 Db 106 Sg 107 Bh 108 Hs 109 Mt 110 Ds 111 Rg 112 Uub 114 Uuq 116 Uuh 118 Uuo (223) 226.0 227.0 (261) (262) (263) (262) (265) (266) (269) (272) (277) (2??) (2??) (2??) 58 Ce 59 Pr 60 Nd 61 Pm 62 Sm 63 Eu 64 Gd 65 Tb 66 Dy 67 Ho 68 Er 69 Tm 70 Yb 71 Lu 140.1 140.9 144.2 (145) 150.4 152.0 157.3 158.9 162.5 164.9 167.3 168.9 173.0 175.0 90 Th 91 Pa 92 U 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr 232.0 231.0 238.0 237.0 (244) (243) (247) (247) (251) (252) (257) (258) (259) (260) Not valid for use as an USNCO National Examination after April 23, 2008 Page 7 2008 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART III Prepared by the American Chemical Society Olympiad Laboratory Practical Task Force Examiner's Instructions Directions to the Examiner: Thank you for administering the 2008 USNCO laboratory practical on behalf of your Local Section. It is essential that you follow the instructions provided, in order to insure consistency of results nationwide. There may be considerable temptation to assist the students after they begin the lab exercise. It is extremely important that you do not lend any assistance or hints whatsoever to the students once they begin work. As in the international competition, the students are not allowed to speak to anyone until the activity is complete. The equipment needed for each student for both lab exercises should be available at his/her lab station or table when the students enter the room. The equipment should be initially placed so that the materials used for Lab Problem 1 are separate from those used for Lab Problem 2. After the students have settled, read the following instructions (in italics) to the students. Hello, my name is ________. Welcome to the lab practical portion of the U.S. National Chemistry Olympiad Examination. In this part of the exam, we will be assessing your lab skills and your ability to reason through a laboratory problem and communicate its results. Do not touch any of the equipment in front of you until you are instructed to do so. You will be asked to complete two laboratory problems. Students are to work alone. All the materials and equipment you may want to use to solve each problem has been set out for you and is grouped by the number of the problem. Students can use all materials for both lab problems, but each experiment is designed to work best with equipment and materials provided specifically for each lab problem. You will have one hour and thirty minutes to complete the two problems. You may choose to start with either problem. You are required to have a procedure for each problem approved for safety by an examiner. (Remember that approval does not mean that your procedure will be successful–it is a safety approval.) When you are ready for an examiner to come to your station for each safety approval, raise your hand. Safety is an important consideration during the lab practical. You must wear goggles at all times. Wash off any chemicals spilled on your skin or clothing with large amounts of tap water. The appropriate procedures for disposing of solutions at the end of this lab practical are: __________________________________________________________________________________________ ____________________________________________________________________________ We are about to begin the lab practical. Please do not turn the page until directed to do so, but read the directions on the front page. There is a periodic table and constants on the last page. Are there any questions before we begin? Page 8 Not valid for use as an USNCO National Examination after April 23, 2008. Distribute Part III booklets and again remind students not to turn the page until the instruction is given. Part III contains student instructions and answer sheets for both laboratory problems. There is a periodic table on the last page of the booklet. Allow students enough time to read the brief cover directions. Do not turn to page 2 until directed to do so. When you start to work, be sure that you fill out all information at the top of the answer sheets. Are there any additional questions? If there are no further questions, the students should be ready to start Part III. You may begin. After one hour and thirty minutes, give the following directions. This is the end of the lab practical. Please stop and bring me your answer sheets. Thank you for your cooperation during this test. Collect all the lab materials. Make sure that the student has filled in his or her name and other required information on the answer sheets. At this point, you may want to take five or ten minutes to discuss the lab practical with the students. They can learn about possible observations and interpretations and you can acquire feedback as to what they actually did and how they reacted to the problems. After this discussion, please take a few minutes to complete the Post-Exam Questionnaire; this information will be extremely useful to the Olympiad subcommittee as they prepare next year’s exam. Please remember to return the post-exam Questionnaire, the answer sheets from Part III, the Scantron sheets from Part I, and the “Blue Books” from Part II in the UPS return envelope you were provided to this address: ACS DivCHED Exams Institute Department of Chemistry University of Wisconsin – Milwaukee 3210 N Cramer Street Milwaukee, WI 53211 The label on the envelope should have this address already, you will need only to include your return address and call United Parcel Service - UPS (1-800-742-5877) for it to be picked up (or it can be dropped in a UPS collection box). The cost of shipping will be billed to the Exams Institute. You can write down the tracking number on the label to allow you to track your shipment. Wednesday, April 23, 2008, is the absolute deadline for receipt of the exam materials at the Examinations Institute. Materials received after this deadline CANNOT be graded. Be sure to have your envelope picked up no later than April 21, 2008 for it to arrive on time. THERE WILL BE NO EXCEPTIONS TO THIS DEADLINE DUE TO THE TIGHT SCHEDULE FOR GRADING THIS EXAMINATION. Not valid for use as an USNCO National Examination after April 23, 2008 Page 9 Examiner’s List: 2008 USNCO Lab Practical Equipment and Chemicals Lab Problem #1: Materials and Equipment Each student should have available the following equipment and materials: • Clear acetate sheet • One grease pencil, used to write in the acetate sheet • Seven Microtip or thin-stem Beral-style pipets (approx. 2.5-mL volume) to contain unknown solutions • One 150-mL or 250-mL beaker to hold the filled pipets • 3-4 toothpicks for stirring • Access to distilled water Lab Problem #1: Chemicals Each student will need: • Solution of AgNO3 (mw=170), BaCl2 (mw=208), Cu(NO3)2 (mw=188), CuSO4 (mw=160), Pb (NO3)2 (mw=331), KI (mw=166), and Na2S2O3 (mw=158). Notes to Coordinators: DO NOT IDENTIFY FOR THE STUDENTS WHICH SOLUTION IS IN EACH PIPET! • All of the solutions should be 0.10M. The solutions should be filled to the maximum in each pipet and stored upside-down (bulbs down) in the 150- or 250-mL beakers on the day of the exam. Students DO NOT have access to additional quantities of any of these solutions. • The pipets should be labeled using an indelible marker, i.e., a Sharpie®, writing the letter in a clear, legible capital letter on the bulb portion of each pipet. • IMPORTANT: The key for this experiment is as follows: Na2S2O3 = 1 CuSO4 = 2 KI = 3 AgNO3 = 4 Cu(NO3)2 = 5 Pb(NO3)2 = 6 BaCl2 = 7 Page 10 Not valid for use as an USNCO National Examination after April 23, 2008. Lab Problem #2: Materials and Equipment Each student will need: • • • • • • • • • • One 50-mL beaker (to contain the 3.0 M HCl) labeled ‘3.0M HCl’ One 10-mL graduated cylinder One balloon (Included) One scissors One metric ruler with mm precision One length of string approximately 30 cm (12”) in length 2-3 sheets of waxed weighing paper One metal scoopula Access to a 0.01 or better electronic balance Access to distilled water Lab Problem #2: Chemicals • Approximately 1.5 g sample CaCO3 in an unlabeled, capped 10- or 20-mL capacity vial • 20 – 25-mL of 3.0 M HCl Quick Check to be sure this experiment works for your examinees: Notes to Coordinators: • You can pour the HCl into the 50- mL beaker the day of the exam. • The CaCO3 should be powdered, not granular or in rock form. Be sure that the CaCO3 vial is not labeled and is capped. For your reference, below is CaCO3 “Fisher Science Education Catalog” product number: • S719221 Calcium Carbonate: Reagent Grade - Powder; 100g Calcium Carbonate, White, Application: CO2 generation, CaCO3 ($7.10) • S71922 Calcium Carbonate: Reagent Grade - Powder; 500g Calcium Carbonate, White, Application: CO2 generation, CaCO3 ($8.70) • Please give each balloon a few stretches prior to placing at the student lab bench to ensure that the balloon will adequately inflate. Not valid for use as an USNCO National Examination after April 23, 2008 Page 11 USNCO 2008 Part III Answers Lab Problem #1 This lab problem involves knowledge of precipitation and solubility. The focus of this problem in qualitatively determining each of the seven unknown solutions is to apply knowledge and understanding of predicted reactions from observations made combining the solutions with one another using the provided acetate sheet as a spot plate. Procedure Students should have constructed some kind of data table that examines combinations of solutions with one another. The various color changes and formed precipitates providing evidence which students must use to form conclusions about each unknown solution. Qualitative Evidence: The two copper solutions are blue. To distinguish between both blue solutions of copper sulfate and copper nitrate, the copper (II) sulfate will precipitate with barium but not with copper (II) nitrate; both copper (II) solutions will precipitate with the KI. Copper (II) ions will form a redish-brown precipitate with iodide. Silver nitrate forms white solid silver chloride and darkish silver sulfate. Barium forms insoluble whitish barium sulfate and barium iodide. Nitrates are soluble and will thus not form precipitates in combination with any of the other cations. Lead reacts with iodide to form yellow lead iodide, with chloride to form white lead chloride, and with sulfate to form whitish lead sulfate. Thiosulfate will react with silver to form a darkish precipitate that then dissolves on further addition of thiosulfate. Answers: A = Na2S2O3, B = CuSO4, C = KI, D = AgNO3, E = Cu(NO3)2, F = Pb(NO3)2, G =BaCl2. Lab Problem #2 This lab problem involves identification of an unknown metal carbonate (here, CaCO3) by quantitatively reacting a measured mass of the solid sample to 3.0M HCl, capturing the CO2 gas evolved, and determining the volume of the gas collected in the balloon provided to then calculate the molar mass of the metal carbonate and conclude which metal carbonate was present from the choices given. One experiment might have been to weigh the gas volume produced or measure the gas volume by water displacement (though materials provided were insufficient to approximate the volume produced by water displacement). Another avenue to determine the unknown metal, though not implied with the materials and chemicals provided for this lab problem, could have been to react the solid white calcium carbonate with several of the identified solutions from lab problem #1, but there is no quantitative evidence here, and the lab problem specifies not only to identify the unknown metal, M, but to also determine the volume of the gas produced. Students needed to figure out a method of gas collection without the volume of HCl(aq) occupying the balloon. One possibility was to insert a measured mass of the solid carbonate into the balloon, then carefully pull the lip of the balloon over the top of the provided graduated cylinder (filled with the HCl. By shaking the solid carbonate into the cylinder with the balloon still over the top of the cylinder, the Page 12 Not valid for use as an USNCO National Examination after April 23, 2008. two chemical reactants can combine to produce CO2 gas to fill the balloon. One method for measuring the volume could have been to wait for the completed reaction, then carefully twist shut the balloon at its lip, tie off the balloon to make a more spherical shape, then use the string and ruler to measure the balloon’s circumference to then calculate its approximate spherical volume using the formula for circumference and volume of a sphere. Students might have also measured the circumference, emptied the balloon of the products, then refilled it with water to the approximate volume when filled with CO2, and finally measure the volume of water using the graduated cylinder. The distinguish between the possibility of Na2CO3 (molar mass = 106) vs. CaCO3 (molar mass = 100), students might have taken a small solid sample remaining and test its relative solubility in water. Though not explicitly encouraged, students might also have tested the unknown carbonate with the known Cu (II) solutions from Experiment #1 to observe precipitates form. Points of error abound, including assuming standard pressure (1 atm), room temperature (25 oC), a perfect sphere, CO2 collected as a ‘wet’ gas, completed reaction, and some air and uncaptured CO2 gas remaining in the graduated cylinder. Sample Student Data: General reaction: MxCO3 + HCl Æ MCly + H2O(l) + CO2(g), where the mol ratio of MxCO3 : CO2 is 1 : 1. Mass of MxCO3 used = 1.00 g Circumference of balloon = 23.9 cm Radius of balloon, using C = 2πr = 3.8 cm Calculation of volume from a sphere, V = 4/3πr3 = 240 cm3 Use ideal gas law, PV = nRT, to find moles of CO2 = (1)(0.240) = n(.0821)(298), n = 0.00981 Molar mass of MxCO3: 0.00981 mol = 1.00g/mol. mass M = Ca, therefore unknown compound is CaCO3 . = 102 or close to 100, mol. mass if Not valid for use as an USNCO National Examination after April 23, 2008 Page 13 USNCO 2008 Part III Grading Results Lab Problem #1 Excellent Students Results: Student provided clear explanation of procedure that included an organized and methodical plan to react solutions with each other and obtain qualitative data from which to make conclusions. Student showed an organized data table with all possible trial combinations and complete descriptions of each reaction. Student showed at least four relevant equations that follow a logical conclusion based on student observations. Written equations were all balanced and indicated both precipitates and aqueous ions. Student clearly based their conclusions on chemical knowledge about solubility and color of possible precipitates. Student used deduction and analysis to infer the identity of the unknowns using the qualitative data they obtained in this experiment. Average Student Results: Student conducted most of the trial combinations and made use of chemical knowledge to form conclusions about the unknowns, with some incorrect and missing information. Not all solutions correctly identified. Written equations correctly identified precipitates but were either not completely balanced or did not correctly show aqueous ions. Only several of the relevant equations were included or the minimum were included but not al balanced correctly. Some qualitative information was given to demonstrate an understanding of precipitate formation and solubility Below Average Student Results: Student was unable to apply adequate chemical knowledge from solution combinations. Written equations were incomplete. Little or no qualitative information was used to make predictions about the identity of the unknown solutions. Lab Problem #2 Excellent Students Results: Student proposed a clear, detailed procedure for determining the unknown metal by gas collection and measured gas volume. Data was organized and calculations were clear and followed a logical plan based on the student’s experiment. A detailed and complete listing of points of error was given following the experiment. Multiple trials were performed. Credibly creative alternative methods to determine the unknown carbonate were used, including testing solubility and precipitation with known solutions from Exp. #1 Average Student Results: Student made adequate measurements and had a general understanding of connection between gas collection, volume measurement, and identifying the unknown metal by determining the carbonate’s molar mass. Student incorrectly identified the unknown metal or miscalculated the volume of gas due to one or more sources of error. Page 14 Not valid for use as an USNCO National Examination after April 23, 2008. Only 1-2 points of error were given following the experiment. Below Average Student Results: Student was not able to connect the volume of gas collected to the calculation of the unknown carbonate’s molar mass. An experiment was performed to collect the gas but both reactants were placed in the balloon together. One or no points of error were given. Not valid for use as an USNCO National Examination after April 23, 2008 Page 15 2009 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART II Prepared by the American Chemical Society Olympiad Examinations Task Force OLYMPIAD EXAMINATIONS TASK FORCE Arden P. Zipp, Chair, State University of New York, Cortland, NY James Ayers, Mesa State College, Grand Junction, CO Paul Groves, South Pasadena HS, South Pasadena, CA Sherry Berman-Robinson, Consolidated HS, Orland Park, IL (retired) Preston Hays, Glenbrook South HS, Glenbrook, IL Seth Brown, University of Notre Dame, Notre Dame, IN David Hostage, Taft School, Watertown, CT Peter Demmin, Amherst HS, Amherst, NY (retired) Marian Dewane, Centennial HS, Boise, ID Valerie Ferguson, Moore HS, Moore, OK Adele Mouakad, St. John’s School, San Juan, PR Jane Nagurney, Scranton Preparatory School, Scranton, PA Ronald Ragsdale, University of Utah, Salt Lake City, UT Kimberly Gardner, US Air Force Academy, Colorado Springs, CO DIRECTIONS TO THE EXAMINER–PART II Part II of this test requires that student answers be written in a response booklet of blank pages. Only this “Blue Book” is graded for a score on Part II. Testing materials, scratch paper, and the “Blue Book” should be made available to the student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until April 29, 2009, after which tests can be returned to students and their teachers for further study. Allow time for the student to read the directions, ask questions, and fill in the requested information on the “Blue Book”. When the student has completed Part II, or after one hour and forty-five minutes has elapsed, the student must turn in the “Blue Book”, Part II of the testing materials, and all scratch paper. Be sure that the student has supplied all of the information requested on the front of the “Blue Book,” and that the same identification number used for Part I has been used again for Part II. There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and you are free to schedule rest-breaks between parts. Part I Part II Part III 60 questions 8 questions 2 lab problems single-answer multiple-choice problem-solving, explanations laboratory practical 1 hour, 30 minutes 1 hour, 45 minutes 1 hour, 30 minutes A periodic table and other useful information are provided on the back page for student reference. Students should be permitted to use non-programmable calculators. DIRECTIONS TO THE EXAMINEE–PART II DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Part II requires complete responses to questions involving problem-solving and explanations. One hour and forty-five minutes are allowed to complete this part. Be sure to print your name, the name of your school, and your identification number in the spaces provided on the “Blue Book” cover. (Be sure to use the same identification number that was coded onto your Scantron® sheet for Part I.) Answer all of the questions in order, and use both sides of the paper. Do not remove the staple. Use separate sheets for scratch paper and do not attach your scratch paper to this examination. When you complete Part II (or at the end of one hour and forty-five minutes), you must turn in all testing materials, scratch paper, and your “Blue Book.” Do not forget to turn in your U.S. citizenship statement before leaving the testing site today. Not valid for use as an USNCO Olympiad National Exam after April 29, 2009. Page 1 1. (12%) Butanoic (butyric) acid, C3H7COOH, is a monoprotic acid with Ka = 1.51×10–5. A 35.00 mL sample of 0.500 M butanoic acid is titrated with 0.200 M KOH. a. Calculate the [H+] in the original butanoic acid solution. b. Calculate the pH after 10.00 mL of KOH have been added. c. Determine the pH at the half-equivalence point of the titration. d. Find the volume of KOH solution needed to reach the equivalence point for the titration. e. Calculate the pH at the equivalence point. (a.) Let HA = C3H7COOH, and A–= C3H7COO– Ka H + ][ A - ] [ = [ HA] Let [H+] and [A-] = x. Plugging in we get, 1.51" 10#5 = x2 . Solving for x gives, [H+] = 2.73×10–3 (0.500 # x) (b.) Determine the initial number of moles of acid: 0.03500 L " 0.500 M = 0.0175 mol HA Determine the number of moles of NaOH added: ! 0.0100 L " 0.200 M = 0.0020 mol OH added ! 0.0155 mol HA Determine the molarity of H+: 0.0175!– 0.0020 = 0.0155 mol HA remain, in 0.045 L, so the molarity is = 0.344 M 0.045 L ! 0.0020 mol A The [A-] is changed only through dilution, [A-] = = 0.0444 M 0.045 L ! Plug these values into the equilibrium constant expression and solve for [H+]. [H ](0.0444) ,and + 1.51"10#5 = (0.344) ! "10 [H ] = 1.16 + -4 , so pH = log(1.16 "10 -4 ) = 3.93 (c.) At the half equivalence point: [HA] = [A-] and [H+] = Ka = 1.51×10–5. So, pH = log(1.51×10–5) = 4.82 ! (d.) The initial moles HA (from part b.) = 0.0175 mol HA, so the equivalence point is reached when we have 0.0175 mol OH– added. 0.0175mol OH - " 1L = 0.0875 L . 0.200 mol OH - (e.) Total volume at the equivalence point is 0.0875 L + 0.0350 L = 0.1225 L [ ] At the equivalence point all HA is converted to A-, so: A" = ! 0.0175 mol = 0.143 M 0.1225 L The A- is a base according to the equation, A– + H2O ⇔ HA + OH–, and K b = ! Let [OH–] and [HA] = x. Plugging in we get, 6.63 "10#10 = K w 1.0 " 10#14 = = 6.63" 10#10 K a 1.51" 10#5 x 2 . Solving for x = [OH–] = 9.73×10-6. (0.143) ! -6 So, pOH = -log(9.73×10 ) = 5.01 and pH = 14.00 – 5.01 = 8.99 2. ! (12%) Chromium metal reacts with acid to produce Cr3+ ions and hydrogen gas. a. Write a balanced equation for this reaction. b. When a sample of chromium metal is reacted with excess acid, 94.7 mL of gas is collected over water at 745 mm Hg and 20˚C. Assuming ideal gas behavior determine the mass of metal reacted. (The vapor pressure of water at 20˚C is 24 mmHg.) c. State and explain how the volume of gas would change if Cr2+ (rather than Cr3+) ions were formed in this reaction. d. Determine the number of molecules of water vapor that would be present in the volume of gas produced assuming ideal behavior. e. Calculate the ratio of the average molecular velocity of hydrogen to the average molecular velocity of water vapor at the same temperature. f. Page 2 ( " %2 + *) # V & -, n The van der Waals equation for real gases is *P + a$ ' - . [V / nb] = nRT . Not valid for use as an USNCO Olympiad National Exam after April 29, 2009. ! The coefficients, a and b, for the hydrogen gas are a = 0.242 atm.L2.mol–2 and b = 0.0266 L.mol–1. The corresponding values of a and b for sulfur dioxide are 6.714 and 0.05636, respectively. i. Identify the molecular property that corresponds to the a coefficient and account qualitatively for the difference between its values for these two gases. ii. Identify the molecular property that corresponds to the b coefficient and account qualitatively for the difference between its values for these two gases. a. 2Cr + 6H+ 2Cr3+ + 3H2 b. The pressure of H2(g) produced is: 745 mmHg – 24 mmHg = 721 mmHg " 721 % atm'( 0.0947 L) $ pV # 760 & Determine number of moles of hydrogen: n = = = 0.00373 mol H 2 RT 0.0821 L ( atm mol ( K ( 293 K) 2 mol Cr 52.00 g Cr " = 0.129 g Cr Now get mass of chromium: 0.00373 mol H 2 " 3 mol H 2 1 mol Cr ( ) c. The volume of gas would be only!2/3 as much (63.1 mL) becase each Cr atom would release only two electrons to reduce the H+ ions (rather than 3.) " 24! % atm'( 0.0947 L) $ pV # 760 & d. n = = = 1.24 ) 10 -4 mol H 2 L ( atm RT 0.0821 mol ( K ( 293 K) ( ) and 1.24 " 10 -4 mol H 2 " ! e. ! vH2 vH2O = MM H 2 O MM H 2 = 6.02 " 10 23 molecules H 2 = 7.49 " 1019 molecules H 2 1 mol H 2 18 = 3 so the ratio of velocities is 3:1 2 ! f. (i.) The “a” coefficient is part of the term that is a correction factor for the attractive forces between molecules. SO2 has a larger value for “a” because SO2 molecules have stronger forces (due to it being both larger than H2, and polar.) (ii.) The “b” coefficient is part of the term that is a correction factor for the molecular volume. Because SO2 is a larger molecule, (more volume) the value for “b” is larger. 3. (14%) The reaction of bromate and bromide ions in acid solution is represented by the equation, 5Br–(aq) + BrO3–(aq) + 6H3O+(aq) r 3Br2(aq) + 9H2O(l) In order to measure the rate of the reaction, stock solutions were prepared as shown in the table: Stock Solution Concentrations Stock Br– solution 1.37 M –3 Stock BrO3– solution 7.10×10 M Stock H3O+ solution 0.573 M Reaction mixtures were prepared by mixing the volumes of solutions listed below, and the initial rate of disappearance of bromate ion was measured. Expt. 1 2 3 4 a. b. c. Vol. Br– stock (mL) 0.100 0.200 0.100 0.200 Vol. BrO3- stock (mL) 0.500 0.500 1.000 0.500 Vol H3O+ stock (mL) 1.000 1.000 1.000 0.700 Vol H2O (mL) 1.400 1.300 0.900 1.600 Initial rate of BrO3– disappearance (mol·L–1·s–1) 5.63×10–6 1.09×10–5 1.13×10–5 5.50×10–6 Calculate the rate of appearance of Br2(aq) in experiment one. Write the rate law for this reaction and give the value of the specific rate constant, k. The following mechanism is proposed for the reaction: (I) BrO3–(aq) + H3O+(aq) r HBrO3(aq) + H2O(l) (II) HBrO3(aq) + H3O+(aq) r H2 BrO3+(aq) + H2O(l) Not valid for use as an USNCO Olympiad National Exam after April 29, 2009. Page 3 (III) H2BrO3+(aq) r BrO2+(aq) + H2O(l) (IV) BrO2+(aq) + Br–(aq) r BrOBrO(aq) (V) BrOBrO(aq) + Br–(aq) r Br2 (aq) + BrO2–(aq) Subsequent reactions of BrO2-(aq) are fast. i. Draw a Lewis structure of BrO2+ and predict its geometry. ii. Given the rate law you determined in b, which of the steps (I)-(V) could potentially be rate-limiting? Justify your answer. a. "[ Br2 ] "[ Br2 ] = 3# = 3(5.63 # 10$6 ) = 1.69 # 10$5 mol L %s "t "t b. • Looking at experiments 2 and 1, the volume of Br– is doubled and the rate increases by 1.94 (essentially doubles) so the reaction is first order in Br–. • Looking at experiments 3 and 1, the volume of BrO3– is doubled and the rate doubles so the reaction is first order in BrO3–. • Looking at experiments 2 and 4, the volume ratio of H3O+ is (1.00/0.700 = 1.4) and the rate ratio is (1.09 / 0.55 = 1.98) so the reaction is second order in H3O+ because (1.4)2 = 1.96 Thus, the rate law is : rate = k[Br–][ BrO3–][ H3O+]2 and 5.63 " 10#6 mol L $ s 3 rate k= = = 2.86 L 2 2 mol3 $ s & 0.1 )& 0.5 ) #3 mol )& 1 Br - BrO 3- H 3O + mol mol % 1.37 % 7.1 " 10 % 0.573 ( L +*(' 3 L +*(' 3 L +* ' 3 ! [ ][ ][ ] c. (i.) One resonance structure is shown below, and the shape determined by VSEPR is bent. O ! Br O (ii.) To have the rate law determined in part (b), the rate limiting step of the mechanism must depend on the concentrations of [Br–], [BrO3–], and [H3O+]2. The only way for this to be true is for Step IV to be the rate limiting step. One way to confirm this is to determine what the rate law would be for each step as the rate limiting step. If Step I is limiting, the rate would vary with [BrO3–][ H3O+] If Step II is limiting, the rate would vary with [BrO3–][ H3O+]2 If Step III is limiting, the rate would vary with [BrO3–][ H3O+]2 If Step IV is limiting, the rate would vary with [BrO3–][ Br–][ H3O+]2 If Step V is limiting, the rate would vary with [BrO3–][ Br–]2[ H3O+]2 4. (12%) There is great current interest in developing fuel cells based on the reaction, 2CH3OH(l) + 3O2(g) r 2CO2(g) + 4H2O(l) a. Write a balanced equation for the half-reaction that occurs in acid solution for such a fuel cell at the; i. anode. ii. cathode. b. If the E˚ value for the cell reaction is 1.21 V, calculate the value of ∆G˚. c. The E˚ value for the O2(g) half reaction is 1.23 V in 1 M H+, calculate the E˚ value expected in 1 M OH–. d. State two advantages of carrying out this reaction in a fuel cell rather than burning methanol and converting the heat into electricity. a. i. anode: CH3OH + H2O → CO2 + 6H+ + 6e– ii.) cathode: O2 + 4H+ + 4e– → 2H2O o o b. ΔG = –nFE = –(12 mol)(96500 J/V⋅mol)(1.21 V) = –1.40×103 kJ # & RT % 1 ( o ln c. Use the Nernst equation: E = E " nF %% H + 4 (( $ ' ' (8.314 J/mol # K)(298 K) $ 1 0.0257 ln& "14 4 ) = 1.23" ln 10 56 = 0.40 V so, E = 1.23" (4)(96500) 4 % (10 ) ( [ ] ( ) Page 2 ! ! Not valid for use as an USNCO Olympiad National Exam after April 29, 2009. d. 1. No wasted heat. 2. No energy lost during conversion. 5. (12%) Write net equations for each of the combinations of reactants below. Use appropriate ionic and molecular formulas and omit formulas for all ions or molecules that do not take part in a reaction. Write structural formulas for all organic substances. You need not balance the equations. All reactions occur in aqueous solution unless otherwise indicated. a. Solid calcium is heated in nitrogen gas. b. Solid sodium ethoxide is added to water. c. Solutions of magnesium sulfate and barium hydroxide are mixed. d. An acidic potassium permanganate solution is added to a solution of sodium sulfite. e. Radium-222 undergoes alpha decay. f. 2-propanol is heated with concentrated sulfuric acid. a. Ca(s) + N 2 ( g) " "# Ca 3 N 2 ( s) b. NaOCH 2CH 3 ( s) + H 2O " "# Na + ( aq) +OH - ( aq) + CH 3CH 2OH ! ! ! c. Mg 2+ ( aq) + SO 42- ( aq) + Ba 2+ ( aq) + 2OH - ( aq) " "# BaSO 4 (s) + Mg(OH) 2 (s) d. MnO 4- ( aq) + H + ( aq) + SO 32- ( aq) " "# Mn 2+ ( aq) + H 2O + SO 42- ( aq) e. 222 "# 42 He 88 Ra" + 218 86 Rn H 2SO 4 f. CH 3CH(OH)CH 3 " " "# CH 3CH = CH 2 + H 2O ! 6. ! (12%) a. Explain why many chemical reactions that are nonspontaneous, with ∆G˚ > 0 at room temperature, proceed to a ! significant extent at that temperature. b. Account for the fact that standard enthalpies of formation of compounds at 25˚C may be either positive or negative. c. Explain why all elements and compounds have positive S˚ values at 25˚C. d. Give an example of a chemical species that does not have a positive S˚ value at 25 ˚C and explain why its standard entropy is not positive. a. ΔGo values refer to standard conditions including 1 M concentrations. Reactions that are nonspontaneous under these conditions may be caused to occur by increasing the concentration of the reactants and/or decreasing the concentrations of the products. b. ΔHfo values of compounds are relative to their elements in standards states (for which ΔH fo = 0). Depending on the compound, formation may either release energy (ΔHfo < 0) or absorb energy (ΔHfo > 0). c. The standard for entropy, S˚, is a perfect crystal at 0 K, which by the Third Law of Thermodynamics is zero. As temperature increases, entropy increases, so S˚ is positive at 25˚C. d. S˚ values of many ions (such as F–, Cl–, PO43–) are less than zero, This occurs because the reference for aqueous ions is the standard entropy for H+, which is set to zero. Some ions, like those listed, may organize the solvent molecules more than the hydrogen ions, so their standard entropy will be negative. 7. (12%) Account for the following observations in terms of atomic/ionic/molecular properties. a. Sodium fluoride melts at a higher temperature than potassium chloride. b. Titanium(III) chloride is a solid at room temperature but titanium(IV) chloride is a liquid at room temperature.. c. N2O3 is an acidic anhydride but Bi2O3 is a basic anhydride. d. Lithium chloride is much more soluble in ethanol than is sodium chloride. a. The internuclear distance in NaF is less than that in KCl so the lattice energy of NaF is greater. Overcoming larger lattice energy leads to higher melting points. b. TiCl3 has Ti3+ ions at the center and is an ionic compound whereas for TiCl4, the smaller Ti4+ ion causes the Ti–Cl bonds to have more covalent character. Covalent molecules typically melt at lower temperatures than ionic ones. c. N2O3 reacts with H2O to form HONO, the high electronegativity of N draws electrons from H of H-O bond to give H+. Bi2O3 reacts with H2O to form Bi(OH)3. The less electronegative Bi3+ bonds less strongly to O so OH– is released. Not valid for use as an USNCO Olympiad National Exam after April 29, 2009. Page 5 d. The smaller Li+ ion has a higher charge density than the larger Na+ ion. This makes LiCl more covalent than NaCl (for the same reasons noted in part (a)). Covalent compounds are more soluble in C2H5OH which has lower polarity. 8. (14%) Four compounds with a molar mass of 59 have the formula C3H9N and the structures: CH3CHNH2 CH3CH2CH2NH2 CH3CH2NHCH3 (CH3)3N CH3 a. Name the class to which these compounds belong. b. The boiling points of the four compounds vary from 3˚C to 46˚C. Identify the lowest and highest boiling compounds and account for the difference in terms of the intermolecular forces in each. c. Each of the four compounds is basic. For one of the compounds draw a structural formula for the conjugate acid formed with H+. Account for the observation that all of these compounds are more basic than ammonia. d. There are two amides with the formula C2H5NO and the same molar mass as the above compounds. i. Draw structural formulas for these two compounds. ii. State whether these compounds have boiling points above or below 46˚C. Rationalize your prediction. iii. State whether these compounds are more or less basic than those with the structures given above. Rationalize your prediction. a. These molecules are amines. b. (CH3)3N is the lowest boiling, CH3CH2CH2NH2 is the highest boiling. (CH3)3N is lowest because it has no hydrogen bonding interactions. The remaining three all have hydrogen bonding, so the highest boiling will have the largest dispersion forces, which the longer chain alkane provides. c. Any of these four drawings would count… CH3 H H H CH3 CH2 CH2 N CH3 H CH N H CH3 CH2 N H CH3 N H H CH3 CH3 CH3 H These compounds are more basic than NH3 because the carbon containing groups are better electron donors than hydrogen. This inductive effect causes the lone pairs on the nitrogen to be donated to H+ more readily. d. i. O O CH3 C N H H C N CH3 H H or ii. These compounds have higher boiling points because they will hydrogen bond more strongly. iii. They will be less basic. The C=O functional group will draw electrons from the nitrogen making the lone pairs less available. END OF PART II Page 2 Not valid for use as an USNCO Olympiad National Exam after April 29, 2009. amount of substance ampere atmosphere atomic mass unit atomic molar mass Avogadro constant Celsius temperature centi- prefix coulomb electromotive force energy of activation enthalpy entropy ABBREVIATIONS AND SYMBOLS n equilibrium constant K measure of pressure mmHg A Faraday constant F milli- prefix m atm formula molar mass M molal m u free energy G molar M A frequency ν mole mol NA gas constant h R Planck’s constant °C gram P g pressure c heat capacity k Cp rate constant C hour Rf h retention factor E joule s J second Ea kelvin c K speed of light H kilo- prefix T k temperature, K S liter t L time volt V CONSTANTS R = 8.314 J·mol–1·K–1 R = 0.0821 L·atm·mol–1·K–1 1 F = 96,500 C·mol–1 1 F = 96,500 J·V–1·mol–1 NA = 6.022 × 1023 mol–1 h = 6.626 × 10–34 J·s c = 2.998 × 108 m·s–1 USEFUL EQUATIONS E = E! – ! k2 $ Ea ! 1 1 $ = ' " k1 &% R #" T1 T2 &% " –!H % " 1 % ln K = $ ' $ ' +c # R & # T& RT ln Q nF ln # PERIODIC TABLE OF THE ELEMENTS 1 H 2 He 1.008 4.003 3 Li 4 Be 5 B 6 C 7 N 8 O 9 F 10 Ne 6.941 9.012 10.81 12.01 14.01 16.00 19.00 20.18 11 Na 12 Mg 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar 22.99 24.31 26.98 28.09 30.97 32.07 35.45 39.95 19 K 20 Ca 21 Sc 22 Ti 23 V 24 Cr 25 Mn 26 Fe 27 Co 28 Ni 29 Cu 30 Zn 31 Ga 32 Ge 33 As 34 Se 35 Br 36 Kr 39.10 40.08 44.96 47.88 50.94 52.00 54.94 55.85 58.93 58.69 63.55 65.39 69.72 72.61 74.92 78.96 79.90 83.80 37 Rb 38 Sr 39 Y 40 Zr 41 Nb 42 Mo 43 Tc 44 Ru 45 Rh 46 Pd 47 Ag 48 Cd 49 In 50 Sn 51 Sb 52 Te 53 I 54 Xe 85.47 87.62 88.91 91.22 92.91 95.94 (98) 101.1 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.3 55 Cs 56 Ba 57 La 72 Hf 73 Ta 74 W 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 132.9 137.3 138.9 178.5 181.0 183.8 186.2 190.2 192.2 195.1 197.0 200.6 204.4 207.2 209.0 (209) (210) (222) 87 Fr 88 Ra 89 Ac 104 Rf 105 Db 106 Sg 107 Bh 108 Hs 109 Mt 110 Uun 111 Uuu 112 Uub 114 Uuq (223) (226) (227) (261) (262) (263) (262) (265) (266) (269) (272) (277) (2??) 58 Ce 59 Pr 60 Nd 61 Pm 62 Sm 63 Eu 64 Gd 65 Tb 66 Dy 67 Ho 68 Er 69 Tm 70 Yb 71 Lu 140.1 140.9 144.2 (145) 150.4 152.0 157.3 158.9 162.5 164.9 167.3 168.9 173.0 175.0 90 Th 91 Pa 92 U 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr 232.0 231.0 238.0 237.0 (244) (243) (247) (247) (251) (252) (257) (258) (259) (260) Page 7 2009 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART III Prepared by the American Chemical Society Olympiad Laboratory Practical Task Force OLYMPIAD LABORATORY PRACTICAL TASK FORCE Steve Lantos, Brookline High School, Brookline, MA Chair Linda Weber, Natick High School, Natick, MA John Mauch, Braintree High School, Braintree, MA Mathieu Freeman, Greens Farms Academy, Greens Farms, CT Erling Antony, Arrowhead Union High School, Hartland, WI Christie B. Summerlin, University of Alabama-Birmingham, Birmingham, AL DIRECTIONS TO THE EXAMINER–PART III The laboratory practical part of the National Olympiad Examination is designed to test skills related to the laboratory. Because the format of this part of the test is quite different from the first two parts, there is a separate, detailed set of instructions for the examiner. This gives explicit directions for setting up and administering the laboratory practical. There are two laboratory tasks to be completed during the 90 minutes allotted to this part of the test. Students do not need to stop between tasks, but are responsible for using the time in the best way possible. Each procedure must be approved for safety by the examiner before the student begins that procedure. Part III 2 lab problems laboratory practical 1 hour, 30 minutes Students should be permitted to use non-programmable calculators. DIRECTIONS TO THE EXAMINEE–PART III DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. WHEN DIRECTED, TURN TO PAGE 2 AND READ THE INTRODUCTION AND SAFETY CONSIDERATIONS CAREFULLY BEFORE YOU PROCEED. There are two laboratory-related tasks for you to complete during the next 90 minutes. There is no need to stop between tasks or to do them in the given order. Simply proceed at your own pace from one to the other, using your time productively. You are required to have a procedure for each problem approved for safety by an examiner before you carry out any experimentation on that problem. You are permitted to use a non-programmable calculator. At the end of the 90 minutes, all answer sheets should be turned in. Be sure that you have filled in all the required information at the top of each answer sheet. Carefully follow all directions from your examiner for safety procedures and the proper disposal of chemicals at your examining site. Not valid for use as an USNCO National Examination after April 29, 2009 Page 1 2009 UNITED STATES NATIONAL CHEMISTRY OLYMPIAD PART III — LABORATORY PRACTICAL Student Instructions Introduction These problems test your ability to design and carry out laboratory experiments and to draw conclusions from your experimental work. You will be graded on your experimental design, on your skills in data collection, and on the accuracy and precision of your results. Clarity of thinking and communication are also components of successful solutions to these problems, so make your written responses as clear and concise as possible. Safety Considerations You are required to wear approved eye protection at all times during this laboratory practical. You also must follow all directions given by your examiner for dealing with spills and with disposal of wastes. Lab Problem 1 You have been given six numbered pipets containing 0.50M solutions of the sodium salts Na2CO3, NaHCO3, NaHSO3, NaH2PO4, Na2HPO4, Na3PO4, not necessarily in this order, a 50-mL beaker containing 0.40M HCl, and a pipet containing methyl orange indicator. Devise and carry out an experiment to determine to contents of each pipet, providing both qualitative and quantitative data to justify your conclusions. Lab Problem 2 You have been given a thermometer, styrofoam cup with lid, a beaker, a graduated cylinder, and access to room temperature water, heated water and ice cubes. Using these materials, design and carry out an experiment to determine the heat of fusion, Hf, for water. Page 2 Not valid for use as an USNCO National Examination after April 29, 2009 Answer Sheet for Laboratory Practical Problem 1 Student's Name: __________________________________________________________________________ Student's School: _______________________________________ Date: ___________________________ Proctor's Name: _________________________________________________________________________ ACS Section Name: ________________________________ Student's USNCO test #: ________________ 1. Give a brief description of your experimental plan. Before beginning your experiment, you must get approval (for safety reasons) from the examiner. Examiner’s Initials: 2. Record your data and other observations. Not valid for use as an USNCO National Examination after April 29, 2009 Page 3 (data and observations – continued) 3. Based on your observations, write the relevant equations that led to your conclusions: 4. Conclusions Pipet # Page 4 Contents Justification Not valid for use as an USNCO National Examination after April 29, 2009 Answer Sheet for Laboratory Practical Problem 2 Student's Name: __________________________________________________________________________ Student's School: _______________________________________ Date: ___________________________ Proctor's Name: _________________________________________________________________________ ACS Section Name: ________________________________ Student's USNCO test #: ________________ 1. Give a brief description of your experimental plan. Before beginning your experiment, you must get approval (for safety reasons) from the examiner. Not valid for use as an USNCO National Examination after April 29, 2009 Examiner’s Initials: Page 5 2. Record your data and other observations. 3. Calculations and Conclusions. 4. Conclusions: The Hf for water is: _____________ 5. Sources of Error in this experiment (please number) Page 6 Not valid for use as an USNCO National Examination after April 29, 2009 2009 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART III Prepared by the American Chemical Society Olympiad Laboratory Practical Task Force Examiner's Instructions Directions to the Examiner: Thank you for administering the 2009 USNCO laboratory practical on behalf of your Local Section. It is essential that you follow the instructions provided, in order to insure consistency of results nationwide. There may be considerable temptation to assist the students after they begin the lab exercise. It is extremely important that you do not lend any assistance or hints whatsoever to the students once they begin work. As in the international competition, the students are not allowed to speak to anyone until the activity is complete. The equipment needed for each student for both lab exercises should be available at his/her lab station or table when the students enter the room. The equipment should be initially placed so that the materials used for Lab Problem 1 are separate from those used for Lab Problem 2. After the students have settled, read the following instructions (in italics) to the students. Hello, my name is ________. Welcome to the lab practical portion of the U.S. National Chemistry Olympiad Examination. In this part of the exam, we will be assessing your lab skills and your ability to reason through a laboratory problem and communicate its results. Do not touch any of the equipment in front of you until you are instructed to do so. You will be asked to complete two laboratory problems. Students are to work alone. All the materials and equipment you may want to use to solve each problem has been set out for you and is grouped by the number of the problem. Students can use all materials for both lab problems, but each experiment is designed to work best with equipment and materials provided specifically for each lab problem. You will have one hour and thirty minutes to complete the two problems. You may choose to start with either problem. You are required to have a procedure for each problem approved for safety by an examiner. (Remember that approval does not mean that your procedure will be successful–it is a safety approval.) When you are ready for an examiner to come to your station for each safety approval, raise your hand. Safety is an important consideration during the lab practical. You must wear goggles at all times. Wash off any chemicals spilled on your skin or clothing with large amounts of tap water. The appropriate procedures for disposing of solutions at the end of this lab practical are: We are about to begin the lab practical. Please do not turn the page until directed to do so, but read the directions on the front page. There is a periodic table and constants on the last page. Are there any questions before we begin? Not valid for use as an USNCO National Examination after April 23, 2009 Page 1 Distribute Part III booklets and again remind students not to turn the page until the instruction is given. Part III contains student instructions and answer sheets for both laboratory problems. There is a periodic table on the last page of the booklet. Allow students enough time to read the brief cover directions. Do not turn to page 2 until directed to do so. When you start to work, be sure that you fill out all information at the top of the answer sheets. Are there any additional questions? If there are no further questions, the students should be ready to start Part III. You may begin. After one hour and thirty minutes, give the following directions. This is the end of the lab practical. Please stop and bring me your answer sheets. Thank you for your cooperation during this test. Collect all the lab materials. Make sure that the student has filled in his or her name and other required information on the answer sheets. At this point, you may want to take five or ten minutes to discuss the lab practical with the students. They can learn about possible observations and interpretations and you can acquire feedback as to what they actually did and how they reacted to the problems. After this discussion, please take a few minutes to complete the Post-Exam Questionnaire; this information will be extremely useful to the Olympiad subcommittee as they prepare next year’s exam. Please remember to return the post-exam Questionnaire, the answer sheets from Part III, the Scantron sheets from Part I, and the “Blue Books” from Part II in the UPS return envelope you were provided to this address: ACS DivCHED Exams Institute Department of Chemistry Iowa State University0213 Gilman Hall Ames, IA 50011 The label on the envelope should have this address already, you will need only to include your return address and call United Parcel Service -UPS (1-800-742-5877) for it to be picked up (or it can be dropped in a UPS collection box). The cost of shipping will be billed to the Exams Institute. You can write down the tracking number on the label to allow you to track your shipment. Wednesday, April 29, 2009, is the absolute deadline for receipt of the exam materials at the Examinations Institute. Materials received after this deadline CANNOT be graded. Be sure to have your envelope picked up no later than April 28, 2009 for it to arrive on time. THERE WILL BE NO EXCEPTIONS TO THIS DEADLINE DUE TO THE TIGHT SCHEDULE FOR GRADING THIS EXAMINATION. Not valid for use as an USNCO National Examination after April 23, 2009 Page 2 Examiner’s List: 2009 USNCO Lab Practical Equipment and Chemicals USNCO 2009 PART III: EXAMINER’S NOTES Lab Problem #1: Materials and Equipment Lab Problem #1 Each student will have: Materials • Six Beral-style pipets to be filled with the unknown sodium salts. • One 150-mL or 250-mL beaker for holding the six pipets containing the unknown solutions. • A twelve-hole clear well plate • Several toothpicks for stirring • Two empty Beral-style pipets • One Beral-style pipet to be filled with methyl orange indicator. • Access to sink and running tap water. Chemicals • Six filled Beral-style transfer pipets each containing 0.50M solution of the following sodium salts: Na2CO3, NaHCO3, NaHSO3, NaH2PO4, Na2HPO4, Na3PO4 IMPORTANT: Use this key to number the pipets: Na3PO4 NaHSO3 Na2CO3 NaH2PO #1 #2 #3 #4 NaHCO3 Na2HPO4 #5 #6 4 Vial • A single Beral-style transfer pipet of methyl orange indicator, filled and labeled ‘methyl orange’. Methyl Orange Indicator: 0.1% aqueous solution. • Approximately 20-30 mL of 0.40M HCl, in a 50-mL beaker labeled ‘0.40M HCl’. Notes to Coordinators • Place the numbered pipets upside down in the small beaker at the student’s desk. • Make sure the solutions are freshly made (especially the NaHSO3 solution) prior to the lab. • Obviously, do NOT label the pipets with the chemical formulas. Not valid for use as an USNCO National Examination after April 23, 2009 Page 3 Lab Problem #2 Each student will have: Materials • One Celsius thermometer (-20oC to 100 oC range is sufficient) • One standard size (8 oz.) styrofoam cup with tight-fitting lid. • One 25-mL graduated cylinder • One 250-mL beaker Chemicals • Access to tap water, ice cubes, and hot water. Notes to Coordinators • Make sure that students can easily access the ice cubes and hot water. Ideally, they can obtain each and return to their work areas using their 250-mL beakers. The hot water should be heated to approximately 60-70 oC. • Students might inquire about the need for an electronic balance to determine the mass of the water, but assuming that the density of water is 1 g/mL eliminates the need for determining masses here. Do NOT tell students this – it is for them to figure out. Note that the examiner will need to initial each student’s experimental plan. Please do not comment on the plan other than looking for any potentially unsafe practices. Safety: It is your responsibility to ensure that all students wear safety goggles during the lab practical. A lab coat or apron for each student is desirable but not mandatory. You will also need to give students explicit directions for handling spills and for disposing of waste materials, following approved safety practices for your examination site. Please check and follow procedures appropriate for your site. Not valid for use as an USNCO National Examination after April 23, 2009 Page 4 2009 USNCO Part III Lab Practicals Answers Lab Problem #1 Key: Vial Na3PO4 #1 NaHSO3 #2 Na2CO3 #3 NaH2PO4 #4 NaHCO3 #5 Na2HPO4 #6 This problem tests students’ understanding of acids and bases, titration, differentiation of mono-protic, di-protic, and tri-protic acids in titration, and some qualitative observations in acid-base neutralization. A sample data table quantifying the HCl added might look like this: sample # drops HCl Justifications / Observations Na3PO4 33 No bubbles, color ∆ from red to orange-yellow Na2HPO4 18 No bubbles, color ∆ from red to orange-yellow NaH2PO4 1 No bubbles, color ∆ from red to orange -yellow NaHSO3 8 Sharp, biting odor, very slight bubbling NaHCO3 15 Bubbles, red to orange-yellow color change Na2CO3 29 Bubbles, red to orange-yellow color change Excellent Student Results: • Student presented an organized plan to add a fixed number of drops of each of the unknowns, a fixed number of drops of the methyl orange indicator to each unknown, and then add HCl drop wise until a color change occurs. Student planned to note the color changes, drops added, and a detectable odor for the sulfite ______ solution. • Student showed a carefully constructed data table OR written account of data collected including all color changes, number of drops added to reach color changes, extent of bubbling including size of bubbles, and odor produced. Multiple trials were performed. • Student included each of the six ionic equations that correlated to the reactions performed from data table, correctly indicating ion charges, states symbols, and stoichiometric relationship. • Student correctly identified each of the six unknowns and gave justifications that were consistent with the data taken in this experiment. Average Student Results: • Student presented a plan but might not have indicated the odor to identify the sulfite or was not clear about how to distinguish the bisulfite from the sulfite. • Student might have written most of the equations correctly but neglected to include ion charges or states symbols. One trial was performed. • Most of the identifications were correct; most of the justifications were valid. Below Average Student Results: • Student did not connect the procedure with the conclusions. Plan was vague or unclear about how the added HCl would be used to conclude unknowns. Plan was difficult to follow. • Data was unorganized or difficult to follow. Not valid for use as an USNCO National Examination after April 23, 2009 Page 5 • • Student neglected to include or incorrectly wrote chemical equations. A majority of the unknowns were incorrectly identified. Notes: The sulfite solution should have been made fresh. Bisulfite is easily oxidized in water and older solutions that were used take little or no acid added to cause the color change and distinguishing odor. Lab Problem # 2 This is a calorimetry problem. The novelty in this question was the intended lack of access to a balance. Students were to make assumptions about the mass of ice based on measured volumes of water prior to adding ice, and the final volume after the ice had been added. Students also had to account for the heat required to increase the temperature of the cold water from the melted ice. Conservation of energy applies such that q(ice) + q(cold water) = q(hot water). Once the ice begins to melt (don’t spill, and put the cap on quickly to minimize heat loss) in the calorimetry cup, the hot water becomes colder. Since you don’t end at 0 oC, the equation breaks into: [mice x Hfus] + [mice-water x Cp x Tfinal – initial ] = [mhot water x Cp x Tfinal – initial ]. Then solve for Hfus Below are sample student data: 10.0 gice Hfus + 10.0 gice-water (4.184) 27.0 – 0.0) = -[ 100ghotwater (4.184) (27.0 – 40.0) 10.0 gice Hfus + 1130 J = 5440 J 10.0 gice Hfus = 4310 J Hfus = 431 J/g, about a 30% error based on 334 J/g as a correct value. The question did not specify the units for the reported answer. Other acceptable values include: 80.0 cal/g, 6.01 J/mol, 1440 cal/mol, 6.01 kJ/mol, or 1.44 kcal/mol. Excellent Student Results: • Student included initial and final temperatures, determined quantity of ice used, water used, carried out replicate determinations. The coffee cup/lid were constructed as a calorimeter in this experiment. • Data and calculations clearly used the overall heat lost-heat gained equations. • Students provided a cogent discussion of the major sources of error such as heat lost or gained by the system, inaccuracies of the measuring equipment used. Average Student Results • Student did not perform multiple determinations. • Student did perform more than one trial, but averaged observations rather than results. • Student might have neglected the heating of the water from the melted ice to the Tfinal in their calculations. Below Average Student Results • Student neglected to provide any measure (volume or converted mass) for ice used. • Student did not clearly use the heat lost-heat gained equation to solve for heat of fusion. • Student did not provide ‘heat loss’ or uncertainty in temperature measurement as sources of error, or provided frivolous sources such as ‘spilled water’. Not valid for use as an USNCO National Examination after April 23, 2009 Page 6 2010 U.S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM PART I Prepared by the American Chemical Society Chemistry Olympiad Examinations Task Force OLYMPIAD EXAMINATIONS TASK FORCE Arden P. Zipp, Chair, State University of New York, Cortland, NY James Ayers, Mesa State College, Grand Junction, CO Sherry Berman-Robinson, Consolidated HS, Orlando Park, IL (retired) Seth Brown, University of Notre Dame, Notre Dame, IN Peter Demmin, Amherst HS, Amherst, NY (retired) Marian Dewane, Centennial HS, Boise, ID Xu Duan, Queen Anne School, Upper Marlboro, MD Valerie Ferguson, Moore HS, Moore, OK Julie Furstenau, Thomas B. Doherty HS, Colorado Springs, CO Kimberly Gardner, United States Air Force Academy, CO Regis Goode, Ridge View HS, Columbia, SC Paul Groves, South Pasadena HS, South Pasadena, CA Preston Hayes, Glenbrook South HS, Glenbrook, IL (retired) David Hostage, Taft School, Watertown, CT Dennis Kliza, Kincaid School, Houston, TX Adele Mouakad, St. John's School, San Juan, PR Jane Nagurney, Scranton Preparatory School, Scranton, PA Ronald Ragsdale, University of Utah, Salt Lake City, UT DIRECTIONS TO THE EXAMINER-PART I Part I of this test is designed to be taken with a Scantron answer sheet on which the student records his or her responses. Only this Scantron sheet is graded for a score on Part I. Testing materials, scratch paper, and the Scantron sheet should be made available to the student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until April 26, 2010, after which tests can be returned to students and their teachers for further study. Allow time for students to read the directions, ask questions, and fill in the requested information on the Scantron sheet. The answer sheet must be completed using a pencil, not pen. When the student has completed Part I, or after one hour and thirty minutes has elapsed, the student must turn in the Scantron sheet, Part I of the testing materials, and all scratch paper. There are three parts to the National Chemistry Olympiad Examination. You have the option of administering the three parts in any order, and you are free to schedule rest breaks between parts. Part I 60 questions single answer, multiple-choice 1 hour, 30 minutes Part II 8 questions problem-solving, explanations 1 hour, 45 minutes Part III 2 lab problems laboratory practical 1 hour, 30 minutes A periodic table and other useful information are provided on page 2 for student reference. Students should be permitted to use non-programmable calculators. DIRECTIONS TO THE EXAMINEE DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Answers to questions in Part I must be entered on a Scantron answer sheet to be scored. Be sure to write your name on the answer sheet, an ID number is already entered for you. Make a record of this ID number because you will use the same number on Parts II and III. Each item in Part I consists of a question or an incomplete statement that is followed by four possible choices. Select the single choice that best answers the question or completes the statement. Then use a pencil to blacken the space on your answer sheet next to the same letter as your choice. You may write on the examination, but the test booklet will not be used for grading. Scores are based on the number of correct responses. When you complete Part I (or at the end of one hour and 30 minutes), you must turn in all testing materials, scratch paper, and your Scantron answer sheet. Do not forget to turn in your U.S. citizenship statement before leaving the testing site today. Distributed by American Chemical Society, 1155 16th Street, N.W., Washington, DC 20036 All rights reserved. Printed in U.S.A. Property of ACS USNCO -Not for use as an USNCO National Exam after April 26, 2010 amount of substance ampere atmosphere atomic mass unit Avogadro constant Celsius temperature centi– prefix coulomb density electromotive force energy of activation enthalpy entropy equilibrium constant 1 1A 1 H 1.008 3 Li n A atm u NA °C c C d E Ea H S K ABBREVIATIONS AND SYMBOLS Faraday constant F molar mass free energy G mole frequency ν Planck’s constant gas constant R pressure gram g rate constant hour h reaction quotient joule J second kelvin K speed of light kilo– prefix k temperature, K liter L time measure of pressure mm Hg vapor pressure milli– prefix m volt molal m volume molar M CONSTANTS M mol h P k Q s c T t VP V V R = 8.314 J·mol–1·K–1 R = 0.0821 L·atm·mol–1·K–1 1 F = 96,500 C·mol–1 1 F = 96,500 J·V–1·mol–1 NA = 6.022 × 1023 mol–1 h = 6.626 × 10–34 J·s c = 2.998 × 108 m·s–1 0 °C = 273.15 K PERIODIC TABLE OF THE ELEMENTS 2 2A 4 Be 13 3A 5 B 14 4A 6 C 15 5A 7 N 16 6A 8 O 17 7A 9 F 18 8A 2 He 4.003 10 Ne 6.941 9.012 10.81 12.01 14.01 16.00 19.00 20.18 11 Na 12 Mg 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar 22.99 24.31 19 K 20 Ca 3 3B 21 Sc 4 4B 22 Ti 5 5B 23 V 6 6B 24 Cr 7 7B 25 Mn 8 8B 26 Fe 9 8B 27 Co 10 8B 28 Ni 58.69 63.55 65.39 69.72 72.61 74.92 78.96 79.90 83.80 46 Pd 47 Ag 48 Cd 49 In 50 Sn 51 Sb 52 Te 53 I 54 Xe 39.10 40.08 44.96 47.88 50.94 52.00 54.94 55.85 58.93 37 Rb 38 Sr 39 Y 40 Zr 41 Nb 42 Mo 43 Tc 44 Ru 45 Rh 11 1B 29 Cu 12 2B 30 Zn 26.98 28.09 30.97 32.07 35.45 39.95 31 Ga 32 Ge 33 As 34 Se 35 Br 36 Kr 85.47 87.62 88.91 91.22 92.91 95.94 (98) 101.1 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.3 55 Cs 56 Ba 57 La 72 Hf 73 Ta 74 W 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 132.9 137.3 138.9 178.5 180.9 183.8 186.2 190.2 192.2 195.1 197.0 200.6 204.4 207.2 209.0 (209) (210) (222) 87 Fr 88 Ra 89 Ac 104 Rf 105 Db 106 Sg 107 Bh 108 Hs 109 Mt 110 Ds 111 Rg 112 113 114 115 116 117 118 (223) (226) (227) (261) (262) (266) (264) (277) (268) (281) (272) (Uut) (Uuq) (Uup) (Uuh) (Uus) (Uuo) 58 Ce 60 Nd 61 Pm 62 Sm 63 Eu 64 Gd 65 Tb 66 Dy (277) 67 Ho 68 Er 69 Tm 70 Yb 71 Lu 140.1 140.9 144.2 (145) 150.4 152.0 157.3 158.9 162.5 164.9 167.3 168.9 173.0 175.0 90 Th 91 Pa 92 U 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr (237) (244) (243) (247) (247) (251) (252) (257) (258) (259) (262) 232.0 Page 2 59 Pr Uub 231.0 238.0 Property of ACS USNCO – Not for use as an USNCO National Exam after April 26, 2010 DIRECTIONS When you have selected your answer to each question, blacken the corresponding space on the answer sheet using a soft, #2 pencil. Make a heavy, full mark, but no stray marks. If you decide to change an answer, erase the unwanted mark very carefully. There is only one correct answer to each question. Any questions for which more than one response has been blackened will not be counted. Your score is based solely on the number of questions you answer correctly. It is to your advantage to answer every question. 1. A student prepares a 100 mL aqueous solution containing a small amount of (NH4)2SO4 and a second 100 mL solution containing a small amount of NaI, then mixes the two solutions. Which statement describes what happens? 6. Four elements were tested in the laboratory and gave the results in the table below. Which element is a metalloid? Element Appearance Conductivity A High B Slight luster Shiny (B) Both compounds dissolve initially but NH4I precipitates when the solutions are mixed. C Dull None (C) Both compounds dissolve initially but Na2SO4 precipitates when the solutions are mixed. D Shiny High (A) Both compounds dissolve and remain in solution when the two solutions are mixed. (D) The NaI dissolves but the (NH4)2SO4 does not. There is no change upon mixing. 2. A colored gas is observed with which combination? (A) calcium hydride and water (B) lead metal and nitric acid Low (A) Element A (B) Element B (C) Element C (D) Element D 7. What is the molarity of Na+ ions in a solution made by dissolving 4.20 g of NaHCO3 (M = 84.0) and 12.6 g of Na2CO3 (M = 126) in water and diluting to 1.00 L? (C) sodium carbonate and sulfuric acid (A) 0.050 M (B) 0.100 M (D) zinc sulfide and hydrochloric acid (C) 0.150 M (D) 0.250 M 3. Mixing which pair of 0.10 M solutions produces two precipitates that cannot be separated from one another by filtration? (A) aluminum chloride and copper(II) nitrate (B) strontium bromide and lead(II) acetate (C) magnesium perchlorate and lithium carbonate (D) barium hydroxide and copper(II) sulfate (B) O2 (C) CO2 (D) CH4 5. For aqueous solutions of which of the following substances could the concentration be determined by visible spectrophotometry? I Cr(NO3)3 II KMnO4 III Zn(NO3)2 (A) I only (B) III only (C) I and II only (D) I, II, III Page 3 8. Which solute has the greatest solubility (in mol/L) in water at 25 ˚C and 1 atm? (A) CH4 (B) NH3 (C) AgCl (D) CaSO4 9. Which 2.00 M solution can be used to separate Al3+ from Fe3+ in an aqueous solution? (A) HCl (B) H2SO4 (C) NaCl (D) NaOH 10. The percent composition of the high explosive HNS is 4. Which gas turns limewater, a saturated solution of Ca(OH)2, cloudy? (A) H2 Behavior with HCl Bubbles slowly No reaction No reaction Bubbles rapidly C H N O 37.35% 1.34% 18.67% 42.65% The molar mass of HNS is 450.22. What is the molecular formula of HNS? (A) C13H4N7O12 (B) C14H6N6O12 (C) C15H10N6O11 (D) C16H12N5O11 11. A student prepares four 0.10 M solutions, each containing one of the solutes below. Which solution has the lowest freezing point? (A) CaCl2 (B) KOH (C) NaC2H3O2 (D) NH4NO3 Property of ACS USNCO – Not for use as an USNCO National Exam after April 26, 2010 12. What is the molarity of a hydrochloric acid solution if 20.00 mL of it neutralizes 18.46 mL of a 0.0420 M Ba(OH)2 solution? 19. Calculate ∆E when one mole of liquid is vaporized at its boiling point (80 ˚C) and 1 atm pressure. (A) 0.0194 M (B) 0.0388 M [∆Hvap = 30.7 kJ/mol] (C) 0.0455 M (D) 0.0775 M (A) 33.6 kJ 13. Ar and He are both gases at room temperature. How do the average molecular velocities (V) of their atoms compare at this temperature? (A) VHe = 10VAr (B) VAr = 10VHe (C) VHe = 3VAr (D) VAr = 3VHe 14. Lithium reacts with water to produce hydrogen gas and lithium hydroxide. What volume of hydrogen collected over water at 22˚C and 750 mm Hg pressure is produced by the reaction of 0.208 g of Li? [VPH2O = 19.8 mm Hg] (A) 367 mL (B) 378 mL (C) 735 mL (D) 755 mL 15. Correct statements about samples of ice and liquid water at 0 ˚C include which of the following? I Molecules in ice and liquid water have the same kinetic energy. II Liquid water has a greater entropy than ice. III Liquid water has a greater potential energy than ice. (A) I and II only (B) I and III only (C) II and III only (D) I, II, and III 16. A sample of a volatile liquid is introduced to an evacuated container with a movable piston. Which change occurs as the piston is raised? (Assume some liquid remains.) I The fraction of the molecules in the gas phase increases II The pressure in the container decreases (A) I only (B) II only (C) Both I and II (D) Neither I nor II 17. The kinetic energy of the molecules in a sample of H2O in its stable state at –10 ˚C and 1 atm is doubled. What are the initial and final phases? (A) solid → liquid (B) liquid → gas (C) solid → gas (D) solid → solid 18. Barium metal crystallizes in a body-centered cubic lattice with barium atoms only at the lattice points. If the density of barium metal is 3.50 g/cm3, what is the length of the unit cell? (A) 3.19 × 10–8 cm (B) 4.02 × 10–8 cm (C) 5.07 × 10–8 cm (D) 6.39 × 10–8 cm Page 4 (B) 31.4 kJ (C) 30.0 kJ (D) 27.8 kJ 20. Use the following data to calculate the molar enthalpy of combustion of ethane, C2H6. 2C2H2(g) + 5O2(g) → 4CO2(g) + 2H2O(l) ∆H = –2511 kJ/mol C2H2(g) + 2H2(g) → C2H6(g) ∆H = –311 kJ/mol 2H2(g) + O2(g) → 2H2O(g) ∆H = –484 kJ/mol (A) –1428 kJ/mol (B) –2684 kJ/mol (C) –2856 kJ/mol (D) –3306 kJ/mol 21. A 10.00 g piece of metal is heated to 80.00 ˚C and placed in 100.0 g of water at 23.00 ˚C. When the system has reached equilibrium the temperature of the water and metal are 23.50 ˚C. What is the identity of the metal? [Specific heat capacity of H2O = 4.184 J/g ˚C] (A) Ag (Cp 0.236 J/g ˚C) (B) Cu (Cp 0.385 J/g ˚C) (C) Fe (Cp 0.449 J/g ˚C) (D) Al (Cp 0.901 J/g ˚C) 22. For a reaction at constant pressure to be spontaneous, which relationship must be correct? (A) ∆Hrxn < 0 (B) ∆Grxn < 0 (C) ∆Srxn < 0 (D) ∆Suniv < 0 23. Tungsten is obtained commercially by the reduction of WO3 with H2 according to the equation: WO3(s) + 3 H2(g) → W(s) + 3 H2O(g) The following data related to this reaction at 25 ˚C are available. H2O(g) WO3(s) ∆H˚ kJ/mol –840.3 –241.8 ∆G˚ kJ/mol –763.5 –228.5 The temperature at which this reaction is at equilibrium at 1 atm is closest to which of the following? (A) 124 K (B) 213 K (C) 928 K (D) 2810 K 24. The gaseous compound NOBr decomposes according to the equation NOBr(g) NO(g) + 1/2 Br2(g) At 350 K the equilibrium constant, Kp, is 0.15. What is the value of ∆G˚? (A) –5.5 × 103 J/mol (B) –2.4 × 103 J/mol (C) 2.4 × 103 J/mol (D) 5.5 × 103 J/mol Property of ACS USNCO – Not for use as an USNCO National Exam after April 26, 2010 25. The rate of decomposition of a certain compound in solution is first order. If the concentration of the compound is doubled, what happens to the reaction's half-life? (A) It doubles 30. (B) It decreases to ½ of the original value (C) It decreases to ¼ of the original value (D) It remains the same 26. Consider the reaction: 2 ICl(g) + H2(g) → 2 HCl(g) + I2(g) At a certain temperature the rate constant is found to be 1.63 x 10–6 L/mol.s. What is the overall order of the reaction? (A) zero (B) first (C) second (D) third 27. For the reaction: N2O4(g) → 2 NO2(g) the number of moles of N2O4(g) is Time, min Moles N2O4(g) 0 5 10 0.200 0.170 0.140 What is the number of moles of NO2(g) at t = 10 min? (Assume moles of NO2(g) = 0 at t = 0.) (A) 0.280 (B) 0.120 (C) 0.110 (D) 0.060 28. A compound decomposes with a first-order rate constant of 0.00854 s–1. Calculate the concentration after 5.0 minutes for an initial concentration of 1.2 M. (A) 0.010 M (B) 0.093 M (C) 0.92 M (D) 1.1 M 29. Ozone in the earth's atmosphere decomposes according to the equation: 2 O3(g) → 3 O2(g) This reaction is thought to occur via the two-step mechanism: O2(g) + O(g) Fast, reversible Step 1 O3(g) Step 2 O3(g) + O(g) → 2 O2(g) Slow What rate law is consistent with this mechanism? (A) –∆[O3]/∆t = k[O3] (B) –∆[O3]/∆t = k[O3]2 (C) –∆[O3]/∆t = k[O3]2/[O2] (D) –∆[O3]/∆t = k[O3]2/[O2]3 The rates of many substrate reactions catalyzed by enzymes vary with time as shown. Which factor(s) best account(s) for the constant reaction rate after a certain time? I The enzyme's active sites are filled. II The amount of substrate is constant. (A) I only (B) II only (C) Both I and II (D) Neither I nor II 31. Consider the system at equilibrium: NH4HS(s) NH3(g) + H2S(g) ∆H > 0 Factors which favor the formation of more H2S(g) include which of the following? I adding a small amount of NH4HS(s) at constant volume II increasing the pressure at constant temperature III increasing the temperature at constant pressure (A) I only (B) III only (C) I and II only (D) I and III only 32. A 2.0 L container is charged with a mixture of 6.0 moles of CO(g) and 6.0 moles of H2O(g) and the following reaction takes place: CO2(g) + H2(g) CO(g) + H2O(g) When equilibrium is reached the [CO2] = 2.4 M. What is the value of Kc for the reaction? (A) 16 (B) 4.0 (C) 0.25 (D) 0.063 33. Determine K for the reaction: H2C2O4(aq) + 2OH–(aq) Æ C2O42–(aq) + 2 H2O(l) H2C2O4(aq) Ka1 = 6.5 × 10–2 H2O Kw = 1.0 × 10–14 Ka2 = 6.1 × 10–5 (A) 4.0 × 10–34 (B) 4.0 × 10–6 (C) 4.0 × 106 (D) 4.0 × 1022 34. Which range includes the value of the equilibrium constant, Keq, for a system with ∆G˚ << 0? (A) –1 < Keq < 0 (B) 0 < Keq < 1 (C) Keq < –1 (D) 1 < Keq Property of ACS USNCO – Not for use as an USNCO National Exam after April 26, 2010 Page 5 35. What volumes of 0.200 M HNO2 and 0.200 M NaNO2 are required to make 500. mL of a buffer solution with pH = 3.00? [Ka for HNO2 = 4.00 x 10–4] (A) 250. mL of each (B) 143 mL of HNO2 and 357 mL of NaNO2 (C) 200. mL of HNO2 and 300. mL of NaNO2 (D) 357 mL of HNO2 and 143 mL of NaNO2 36. A sample of sparingly soluble PbI2(s) containing radioactive I-133 is added to 0.10 M KI(aq) and stirred overnight. Observations about this system include which of the following? I The radioactivity of the liquid phase increases significantly. II The concentration of the I– ion in solution increases significantly. (A) I only (B) II only (C) Both I and II (D) Neither I nor II 37. An unknown metal, M, and its salt, M(NO3)2, are combined with a half-cell in which the following reaction occurs: Ag+(aq) + e– → Ag(s) [E˚red = 0.80 V] If E˚cell = 1.36 V, what is E˚ red for M2+(aq) + 2e– → M(s)? (A) 0.56 V (B) 0.24V (C) –0.24V (D) –0.56V 38. Given the standard reduction potentials: O2 + 4H+ + 4e– → 2H2O Br2 + 2e– → 2Br – 2H+ + 2e– → H2 Na+ + e– → Na E˚ = 1.23 V E˚ = 1.08 V E˚ = 0.00 V E˚ = –2.71 V What products are formed in the electrolysis of 1 M NaBr in a solution with [H3O+] = 1 M? (A) Na(s) and O2(g) (B) Na(s) and Br2(g) (C) H2(g) and Br2(g) (D) H2(g) and O2(g) 39. According to the standard reduction potentials: Pb2+(aq) + 2e– Æ Pb(s) Fe2+(aq) + 2e– Æ Fe(s) Zn2+(aq) + 2e– Æ Zn(s) E˚ = –0.13 V E˚ = –0.44 V E˚ = –0.76 V Which species will reduce Mn3+ to Mn2+ [E˚ = 1.51 V] but will NOT reduce Cr3+ to Cr2+ [E˚ = –0.40 V]? (A) Pb only (B) Zn only (C) Pb and Fe only (D) Pb, Fe, and Zn 40. Zn(s) / Zn2+(aq) // H+(aq) / H2(g) E˚ = 0.76 V What must be the pH in the hydrogen compartment of the cell designated above if the cell voltage is 0.70 V? (Assume that both the [Zn2+] and the H2(g) pressure are at standard values and T = 25 ˚C.) (A) 0.51 (B) 1.01 (C) 2.50 (D) 3.21 41. The equilibrium constant, K, is 2.0 × 1019 for the cell Ni(s) / Ni2+(aq) // Hg22+(aq) / Hg (l) The value of E˚ at 25 ˚C for this cell is closest to (A) –1.14V (B) –0.57V (C) 0.57V (D) 1.14V 42. In a battery with a zinc anode, what is the minimum mass of zinc required if a current of 250 mA is drawn for 12.0 minutes? (A) 0.0610 g (B) 0.122 g (C) 0.244 g (D) 1.02 g 43. Which set of quantum numbers (n, l, ml, ms) is possible for the outermost electron in a strontium atom in its ground state? (A) 5, 0, 0, –1/2 (B) 5, 0, 1, 1/2 (C) 5, 1, 0, 1/2 (D) 5, 1, 1, –1/2 44. How many orbitals are in an f sublevel (l = 3)? (A) 3 (B) 5 (C) 7 (D) 14 45. What is the energy of photons with a wavelength of 434 nm? (A) 2.76 × 105 kJ/mol (B) 2.76 × 102 kJ/mol (C) 2.76 × 10–1 kJ/mol (D) 2.76 × 10–4 kJ/mol 46. In which choice are the species listed in order of increasing radius? (A) Na+, Mg2+, Al3+ (B) Cl–, S2–, P3– (C) Ar, K+, Cl– (D) Cl–, Ar, K+ 47. Which element has the highest melting point? (A) Na (B) K (C) Mg (D) Ca 48. Which element forms a compound with the formula H3XO4? (A) As (B) Cl (C) N (D) S 49. Which molecule contains the smallest F-S-F angle? (A) SF2 (B) SOF2 (C) SO2F2 (D) SF6 50. Which species has the longest N–O bond? (A) NO Page 8 (B) NO+ (C) NO2 (D) NO2+ Property of ACS USNCO – Not for use as an USNCO National Exam after April 26, 2010 51. How many pi bonds and how many lone pairs are in the Lewis structure of hydrazine, N2H4? (A) 2 pi bonds, 0 lone pairs 59. All of the following statements concerning benzene, C6H6, are correct EXCEPT (A) Each carbon atom forms three sigma bonds. (B) 1 pi bond, 0 lone pairs (B) Each carbon is sp2 hybridized. (C) 1 pi bond, 1 lone pair (C) Pi electrons are delocalized over all 6 carbon atoms. (D) 0 pi bonds, 2 lone pairs (D) Benzene forms cis and trans isomers when it reacts. 52. In the Lewis structure of nitrous acid: 60. Which functional group is not commonly found in nucleic acids? What is the formal charge on nitrogen? (A) –1 (B) 0 (C) +1 (D) +3 53. How many isomers of octahedral Co(NH3)3Cl3 are there? (A) 2 (B) 3 (C) 4 (A) alcohol (B) amine (C) carboxylic acid (D) dialkyl phosphate END OF TEST (D) 5 54. The bond angle in H2O is approximately 105˚ while the bond angle in H2S is approximately 90˚. Which explanation best accounts for this difference? (A) H–S bonds are longer than H–O bonds. (B) H–S bonds are less polar than H–O bonds. (C) S has d orbitals available for bonding, O does not. (D) O uses sp3 hybrid orbitals for bonding, S uses its 3p orbitals. 55. Which compound exists in optically active forms? (A) CH3CFCClCH3 (B) CH2FCH2CH2Cl (C) CH2FCHClCH3 (D) CHF2CH2CH2Cl 56. What product results when 2-butene reacts with chlorine? (A) 2-chlorobutane (B) 1,2-dichlorobutane (C) 2,2-dichlorobutane (D) 2,3-dichlorobutane 57. Which chloroalkane undergoes substitution with OH– exclusively by an SN1 mechanism? (A) (CH3)2CHCH2Cl (B) (CH3)3CCl (C) CH3CH2CHClCH3 (D) CH3CH2CH2CH2Cl 58. Which is a monosaccharide? (A) fructose (B) lactose (C) maltose (D) sucrose Property of ACS USNCO – Not for use as an USNCO National Exam after April 26, 2010 Page 9 2010 U.S. National Chemistry Olympiad National Exam Part I KEY Number 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. Answer A B D C C B D B D B A D C B D A C C D A B B C D D C B B C A Number 31. 32. 33. 34. 35. 36. 37. 38. 39. 40. 41. 42. 43. 44. 45. 46. 47. 48. 49. 50. 51. 52. 53. 54. 55. 56. 57. 58. 59. 60. Not for use as an USNCO National Exam after April 26, 2010 Answer B A D D D A D C A B C A A C B B D A D C D B A D C D B A D C 2010 U.S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM - PART II Prepared by the American Chemical Society Olympiad Examinations Task Force OLYMPIAD EXAMINATIONS TASK FORCE Arden P. Zipp, Chair, State University of New York, Cortland, NY James Ayers, Mesa State College, Grand Junction, CO Sherry Berman-Robinson, Consolidated HS, Orlando Park, IL (retired) Seth Brown, University of Notre Dame, Notre Dame, IN Peter Demmin, Amherst HS, Amherst, NY (retired) Marian Dewane, Centennial HS, Boise, ID Xu Duan, Queen Anne School, Upper Marlboro, MD Valerie Ferguson, Moore HS, Moore, OK Julie Furstenau, Thomas B. Doherty HS, Colorado Springs, CO Kimberly Gardner, United States Air Force Academy, CO Regis Goode, Ridge View HS, Columbia, SC Paul Groves, South Pasadena HS, South Pasadena, CA Preston Hayes, Glenbrook South HS, Glenbrook, IL David Hostage, Taft School, Watertown, CT Dennis Kliza, Kincaid School, Houston, TX Adele Mouakad, St. John's School, San Juan, PR Jane Nagurney, Scranton Preparatory School, Scranton, PA Ronald Ragsdale, University of Utah, Salt Lake City, UT DIRECTIONS TO THE EXAMINER - PART II Part II of this test requires that student answers be written in a response booklet with blank pages. Only this "Blue Book" is graded for a score on Part II. Testing materials, scratch paper, and the "Blue Book" should be made available to the student only during the examination period. All testing materials including scratch paper should be turned in and kept secure until April 26, 2010, after which tests can be returned to students and their teachers for further study. Allow time for the student to read the directions, ask questions, and fill in the required information on the "Blue Book". When the student has completed Part II, or after one hour and forty-five minutes has elapsed, the student must turn in the "Blue Book", Part II of the testing materials, and all scratch paper. Be sure that the student has supplied all of the information requested on the front of the "Blue Book," and that the same identification number used for Part I has been used again for Part II. There are three parts to the National Olympiad Examination. You have the option of administering the three parts in any order, and you are free to schedule rest breaks between parts. Part I 60 questions single-answer multiple-choice 1 hour, 30 minutes Part II 8 questions problem-solving, explanations 1 hour, 45 minutes Part III 2 lab questions laboratory practical 1 hour, 30 minutes A periodic table and other useful information are provided on the back page for student reference. Students should be permitted to use non-programmable calculators. DIRECTIONS TO THE EXAMINEE - PART II DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. Part II requires complete responses to questions involving problem-solving and explanations. One hour and forty-five minutes are allowed to complete this part. Be sure to print your name, the name of your school, and your identification number in the spaces provided on the "Blue Book" cover. (Be sure to use the same identification number that was coded onto your Scantron sheet for Part I.) Answer all of the questions in order, and use both sides of the paper. Do not remove the staple. Use separate sheets for scratch paper and do not attach your scratch paper to this examination. When you complete Part II (or at the end of one hour and forty-five minutes) you must turn in all testing materials, scratch paper, and your "Blue Book ". Do not forget to turn in your U.S. citizenship statement before leaving the testing site today. Distributed by American Chemical Society, 1155 16th Street, N.W., Washington, DC 20036 All rights reserved. Printed in U.S.A. Property of ACS USNCO – Not for use as an USNCO National Exam after April 26, 2010 amount of substance ampere atmosphere atomic mass unit Avogadro constant Celsius temperature centi– prefix coulomb density electromotive force energy of activation enthalpy entropy equilibrium constant n A atm u NA °C c C d E Ea H S K ABBREVIATIONS AND SYMBOLS Faraday constant F molar mass free energy G mole frequency ν Planck’s constant gas constant R pressure gram g rate constant hour h reaction quotient joule J second kelvin K speed of light kilo– prefix k temperature, K liter L time measure of pressure mm Hg vapor pressure milli– prefix m volt molal m volume molar M CONSTANTS M mol h P k Q s c T t VP V V R = 8.314 J·mol–1·K–1 R = 0.0821 L·atm·mol–1·K–1 1 F = 96,500 C·mol–1 1 F = 96,500 J·V–1·mol–1 NA = 6.022 × 1023 mol–1 h = 6.626 × 10–34 J·s c = 2.998 × 108 m·s–1 0 °C = 273.15 K PERIODIC TABLE OF THE ELEMENTS 1 1A 1 H 18 8A 2 He 3 Li 2 2A 4 Be 13 3A 5 B 14 4A 6 C 15 5A 7 N 16 6A 8 O 17 7A 9 F 6.941 9.012 10.81 12.01 14.01 16.00 19.00 20.18 11 Na 12 Mg 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar 22.99 24.31 26.98 28.09 30.97 32.07 35.45 39.95 19 K 31 Ga 32 Ge 33 As 34 Se 35 Br 36 Kr 1.008 4.003 10 Ne 20 Ca 3 3B 21 Sc 4 4B 22 Ti 5 5B 23 V 6 6B 24 Cr 7 7B 25 Mn 8 8B 26 Fe 9 8B 27 Co 10 8B 28 Ni 11 1B 29 Cu 12 2B 30 Zn 39.10 40.08 44.96 47.88 50.94 52.00 54.94 55.85 58.93 58.69 63.55 65.39 69.72 72.61 74.92 78.96 79.90 83.80 37 Rb 38 Sr 39 Y 40 Zr 41 Nb 42 Mo 43 Tc 44 Ru 45 Rh 46 Pd 47 Ag 48 Cd 49 In 50 Sn 51 Sb 52 Te 53 I 54 Xe 85.47 87.62 88.91 91.22 92.91 95.94 (98) 101.1 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.3 55 Cs 56 Ba 57 La 72 Hf 73 Ta 74 W 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 132.9 137.3 138.9 178.5 180.9 183.8 186.2 190.2 192.2 195.1 197.0 200.6 204.4 207.2 209.0 (209) (210) (222) 87 Fr 88 Ra 89 Ac 104 Rf 105 Db 106 Sg 107 Bh 108 Hs 109 Mt 110 Ds 111 Rg 112 113 114 115 116 117 118 (223) (226) (227) (261) (262) (266) (264) (277) (268) (281) (272) (Uut) (Uuq) (Uup) (Uuh) (Uus) (Uuo) Page 2 Uub (277) 58 Ce 59 Pr 60 Nd 61 Pm 62 Sm 63 Eu 64 Gd 65 Tb 66 Dy 67 Ho 68 Er 69 Tm 70 Yb 71 Lu 140.1 140.9 144.2 (145) 150.4 152.0 157.3 158.9 162.5 164.9 167.3 168.9 173.0 175.0 90 Th 91 Pa 92 U 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr 232.0 231.0 238.0 (237) (244) (243) (247) (247) (251) (252) (257) (258) (259) (262) Property of ACS USNCO – Not for use as an USNCO National Exam after April 26, 2010 1. (12%) 5.60 g of solid carbon is placed in a rigid evacuated 2.5 L container. Carbon dioxide is added to the container to a final pressure of 1.50 atm at 298 K. a. Calculate the number of moles of each reactant in the container originally. 2 CO(g) ∆H˚ = 173 kJ b. The container is heated to 1100 K and the following reaction occurs: C(s) + CO2(g) i. Calculate the pressure in the container at this temperature before the reaction takes place. ii. When equilibrium is reached the pressure inside the container is 1.75 times that calculated in b.i. Determine the equilibrium partial pressures of CO2(g) and CO(g). iii. Write the equilibrium expression for this reaction, Kp. iv. Calculate the value of Kp for this reaction at 1100 K. c. Predict the effect on the number of moles of carbon monoxide of each of the following changes made to this system at equilibrium. Give reasons for your predictions. i. The volume of the container is increased to 5.0 L. ii. The pressure inside the container is increased by adding helium. iii. The temperature of the system is increased to 1200 K. iv. The amount of solid carbon is increased to 6.00 g. 2. (14%) Green plants utilize sunlight to convert CO2 and H2O to glucose (C6H12O6) and O2. a. Write a balanced equation for this process. b. Use the information in the accompanying table to calculate i. ∆H˚ ii. ∆S˚ iii. ∆G˚ at 298 K for this reaction. Substance CO2(g) H2O(l) C6H12O6(s) O2(g) ∆Hf˚ kJ/mol –393.5 –285.8 –1273.3 S˚ J/mol⋅K 213.2 69.9 212.1 205.0 c. Comment on the spontaneity of this reaction at 25˚C and other temperatures. d. Green plants use light with wavelengths near 600 nm for this process. Calculate i. the energy of a 600 nm photon, ii. ∆G˚ for the formation of one molecule of glucose by the reaction in 2a, iii. the minimum number of 600 nm photons required to make one molecule of glucose by the reaction in 2a. e. All of the photosynthesis on earth in a year stores 3.4 × 1018 kJ of solar energy. i. Use the ∆G˚ for the photosynthetic reaction to calculate the number of moles of CO2 removed from the atmosphere by photosynthesis each year. ii. Determine the mass of carbon that is fixed annually by photosynthesis. 3. (14%) A 0.125 g piece of vanadium reacts with nitric acid to produce 50.0 mL of a yellow solution of vanadium ions in their highest oxidation state. a. Calculate the number of moles of vanadium dissolved and the molarity of vanadium ions in this solution. b. Write the electron configuration of a neutral gaseous vanadium atom. c. Give the oxidation state of vanadium in the yellow solution and outline your reasoning. d. A 25.0 mL portion of this yellow solution is reduced with excess zinc amalgam under an inert atmosphere to give a violet – solution. A 10.0 mL aliquot of this violet solution is titrated with a solution of 2.23 × 10 2 M KMnO4 in acid forming Mn2+. – A volume of 13.20 mL of the MnO4 solution is required to convert the vanadium back to yellow. Determine the: – i. number of moles of MnO4 used in this titration, – ii. mole ratio of vanadium ions to MnO4 ions in this titration, iii. oxidation number change for vanadium in this titration and the oxidation state of vanadium ions in the violet solution. e. When 2.00 mL of the violet solution are mixed with 1.00 mL of the original yellow solution, a green solution results. When this ratio is reversed a bright blue solution is formed. Determine the oxidation states of the green and blue vanadium ions. Support your answers with calculations. Property of ACS USNCO Not for use as an USNCO National Exam after April 26, 2010 Page 3 4. (12%) The reaction NO(g) + O3(g) → NO2(g) + O2(g) is first order in each reactant with an activation energy, Ea, of 11.7 kJ/mol and a – – rate constant of k = 1.2 × 1010 .L mol 1 s 1 at 25 ˚C. – a. Calculate the value of the pre-exponential factor, A, in the equation k = Ae Ea/RT. b. Would the A factor for the chemical reaction NO(g) + N2O(g) → NO2(g) + N2(g) be expected to be larger or smaller than the A factor in the above reaction if each reaction occurs in a single step? Outline your reasoning. c. Calculate the rate constant for this reaction at 75 ˚C. d. The following two-step mechanism has been proposed for this reaction: Step 1 O3(g) → O2(g) + O(g) NO(g) + O(g) → NO2(g) Step 2 State and explain whether this mechanism is consistent with the observed rate law. 5. (12%) Write net equations for each of the reactions below. Use appropriate ionic and molecular formulas and omit formulas for all ions or molecules that do not take part in a reaction. Write structural formulas for all organic substances. You need not balance the equations. a. Solutions of hydrochloric acid and silver acetate are mixed. b. A small piece of potassium is added to water. c. Concentrated hydrochloric acid is added to a solution of cobalt(II) sulfate. d. An acidified potassium dichromate solution is added to a tin(II) chloride solution e. Methyl ethanoate (methyl acetate) is reacted with a sodium hydroxide solution. f. Carbon-14 undergoes beta decay. 6. (12%) Account for the following observations on the basis of electrochemical principles. The Standard Reduction Potentials are provided. – E˚ = 1.61 V 2 HOCl(aq) + 2 H+(aq) + 2 e → Cl2(g) + 2 H2O(l) – – E˚ = 1.36 V Cl2(g) + 2 e → 2 Cl (aq) – O2(g) + 4 H+(aq) + 2 e → 2 H2O(l) E˚ = 1.23 V – Cu (aq) + 2 e → Cu(s) E˚ = 0.34 V – Sn (aq) + 2 e → Sn(s) E˚ = –0.14 V – Fe (aq) + 2 e → Fe(s) E˚ = –0.44 V – Zn (aq) + 2 e → Zn(s) E˚ = –0.76 V 2+ 2+ 2+ 2+ In a voltaic cell made with Cu metal in a 1.0 M CuSO4 and Zn metal in 1.0 M ZnSO4 the Zn is the anode and the cell potential is more than 1.0 V. When aqueous sodium sulfide is added to the CuSO4 solution the cell potential decreases substantially. b. Iron metal corrodes readily in moist air but this corrosion can be prevented when iron is coated with tin or zinc. Corrosion is prevented when the zinc coating is intact or broken. In contrast, corrosion is prevented by coating iron with tin only as long as the tin coating remains intact but actually occurs faster when there is a break in the tin coating. c. In acid solution chloride and hypochlorite ions react to form chlorine gas whereas in basic solution chlorine gas reacts to form chloride and hypochlorite ions. a. 7. (12%) Two stable allotropes of oxygen are dioxygen (O2) and ozone (O3). a. Describe the geometry of ozone and state the hybridization of each of the oxygen atoms. b. Ozone has a nonzero dipole moment. Account for this fact and predict the direction of the dipole moment. c. Dioxygen is weakly attracted to strong magnetic fields (i.e. is paramagnetic), while ozone is weakly repelled by magnetic fields (i.e. is diamagnetic). Account for these observations in terms of the bonding in the two molecules. d. The most stable allotrope of sulfur is the cyclic S8 molecule while S2 is a highly unstable gas. In contrast, O2 is the most stable allotrope of oxygen and O8 is unknown. Account for these differences in the relative stability of the allotropes of these two elements. 8. (12%) There are four structural isomers with the formula C4H9Cl, one of which exists in optically active forms. a. Write structural formulas for these four isomers. b. Identify the isomer that exists in optically active forms and describe the difference in behavior of these two forms. – c. Each of these isomers reacts with OH ions to eliminate a molecule of HCl. i. Give the name and molecular formula for the family of compounds formed by this elimination reaction. ii. Write a structural formula for each of the elimination products. iii. Identify the elimination product that can exist in different isomeric forms and draw structures for these forms. Page 4 Property of ACS USNCO – Not for use as an USNCO National Exam after April 26, 2010 2010 U.S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM - PART II - KEY 1 1mol = .466mol 12.01g CO2 n = PV RT n = (1.50atm )(2.5L ) / 298K (.0821) n = 0.153mol (1100K ) P2 = 5.54atm b. (i). P2 = P1 T2 / T1 P2 = 1.50atm 298K (9.70atm)(2.5L) n = PV RT (ii). PT = 1.75(5.54) = 9.70atm n= = 0.268mol 1100K (.0821) 0.268 = 0.153 − x + 2x x = 0.268 − 0.153 x = .115 nRT (.038)(.0821)1100 = 1.3atm n CO = 0.230 PCO 2 = PCO 2 = n CO 2 = 0.038 V 2.5L (.230)(.0821)1100 PCO = 8.31atm PCO = = 8.31atm 2.5L ⎛ (.0821)(1100) ⎞ 9.70atm = (0.153 − x ) + 2x⎜ 0.269 = 0.153 + x x = .116 ⎟ 2 .5 ⎠ ⎝ a. 5.60gC × ( (iii). K p = PCO 2 ) PCO 2 (iv). K p = (8.31) 2 K p = 50.4 1.37 c. (i). nCO will increase. As V is ↑ , P ↓ s so system shifts → . (ii). nCO does not change. He is not in K p so has no effect. (iii). nCO will increase.. ΔΗ is positive so ↑ T will favor → . (iv). nCO will not change. Solids do not affect equilibrium. 2 a. 6CO2+6H2O → C6H12O6+6O2 b. (i). ∆H˚ = –1273.3+0–[6(–393.5)+6(–285.8)] = –1273.3–[–2361–1714.8] = –1273.3+4075.8 = 2802.5 kJ (ii). ΔS˚ = 212.1+6(205.0)–[6(213.2+6(69.9) = 212.1+1230.0–[1279.2+419.4] = 1442.1–1698.6 = –256.5 J/mol K (iii). ΔG˚ = ΔH˚–TΔS˚ ΔG˚ = 2802.5 kJ–298(–.2565 kJ/mol) ΔG˚ = 2802.5 kJ+76.44 = 2878.9 kJ/mol c. Reaction is not spontaneous at 25˚C because ΔG˚>0 Reaction is not spontaneous at other Ts because ∆H˚>0 and ΔS˚<0 d. (i). Ε = hν c = νλ Ε = hc λ (6.626 × 10 −34 )(3 × 108 ) 600 × 10 − 9 E = 3.31 × 10 −19 J kJ 1mol × = 4.78 × 10 − 21 kJ mol ec = 4.78 × 10 −18 J molec (ii). ΔG o / molecule = 2878.9 mol 6.022 × 10 23 molec J 1phot × = 14.4photons (iii). # of photons = 4.78 × 10 −18 molec 3.31 × 10 −19 J 1molC6 H12 O 6 6molCO 2 e. (i). 3.4 × 1018 kJ yr × × = 7.08 × 1015 molCO 2 2.88 × 103 kJ 1molC6 H12 O 6 E= (ii). 7.08 × 1015 molC × 12.01 g mol = 8.50 × 1016 gC 3 1mol = .00245mol 50.94g .00245mol M= = 0.049M .050L b. V Z = 23 1s2 2s2 2p6 3s2 3p6 4s2 3d 3 c. V is in +5 oxid st. due to loss of 4s and 3d electrons. d. (i). mol MnO 4− = 2.23 × 10 −2 mol L × .01320L = 2.94 × 10 −4 mol a. 0.125gV × (0.010L)(.0491 mol L) 4.91 × 10 −4 1.67 = = 1 2.94 × 10 − 4 mol 2.94 × 10 − 4 5 1.67 = 3.0 Δ for V (iii). Mn goes from +7 → +2 Δ = 5 (ii). e. 2.00( x − 2) = 1.00(5 − x ) 1.00( x − 2) = 2.00(5 − x ) 4 a. k = Αe − Ea 2x − 4 = 5 − x 3x = 9 x = 3 green x − 2 = 10 − 2x 3x = 12 x = 4 blue 11700 J 1.2 × 1010 = Αe 8.−314 ( 298) RT V 2 + violet 1.2 × 1010 = Αe −4.722 Α = 1.2 × 1010 / .008894 Α = 1.35 × 1012 b. The A factor for NO and N2O would be smaller than that for NO and O3 because there are fewer geometric arrangements involving NO and N2O molecules that could lead to a successful reaction. The probability of successful reactions are lower for NO and N2O. E ⎛1 k 1 ⎞ c. ln 2 = a ⎜⎜ − ⎟⎟ k1 R ⎝ T1 T2 ⎠ ln k2 11700J ⎛ 1 1 ⎞ = − ⎜ ⎟ 10 J 1.2 ×10 ⎝ 298 348 ⎠ 8.314 MolK = 1407.3(.0033557 − .0028736 ) ( ln k2 1.2 × 1010 ) = 1407.3 4.821 × 10 −4 k2 = 0.6785 = 1.971 1.2 × 1010 k 2 = 2.37 × 1010 d. This mechanism would give either R = k[O 3 ] if first step is the slow one or R = k [ NO][O 3 ] if second step [O 2 ] is slow since neither of these rate laws R = k[ NO][O 3 ] this can’t be the mechanism. 5 a. H + + Cl − + Ag + + C 2 H 3O 2− → AgCl + HC 2 H 3O 2 b. K + H 2 O → K + + OH − + H 2 c. H + + Cl − + Co 2 + + SO 24 − → CoCl 24 − + HSO −4 d. Cr2 O 72− + H + + Sn 2+ → Cr 3+ + Sn 4+ + H 2 O e. O O H3CCOCH3 + OH- → H 3CCO - f. 6 + H3COH 14 14 0 6 C→ 7 N + −1 β E cell = 0.76 + 0.34 = 1.08V a. Zn + Cu2+ → Zn2+ + Cu 22+ When S is added to Cu /Cu half cell CuS forms reducing [Cu2+], shifting the reaction to the left and decreasing Ecell b. Fe → Fe2+ +2e– is 0.44V so oxidation (corrosion) is spontaneous. Covering surface with Sn or Zn prevents reaction with O2. If Zn coating is broken, Zn will still oxidize preferential. If Sn coating is broken, Fe will oxidize more readily c. 2HOCl + 2H+ + 2Cl– →2Cl2 + 2H2O Ε o = 0.25 spontaneous in acid. + In basic solution [H ] is very low so reaction shifts to the left and Cl2 forms Cl– and OCl–. 7 a. O O ↔ O O O + O 2 O O O 2 O3 is bent. Central O is Sp hybridized. All 3 are Sp hybridized in delocalized structure. b. Formal charge for central O is +1, –½ for each of terminal Os in delocalized structure. (-1 and O in each resonance form). DM has + end on central O. c. O3 is diamagnetic because all e– are paired. O2 is paramagnetic because it has 2 unpaired e– (M.O theory). KΚ 4 σ 2 σ*22s π 22 p π 22 p σ*2 π*1π*1 d. P orbitals overlap better in smaller O atoms. So double bond in O2 is stronger that 2 single bonds / e–-e– repulsion between O atoms weakens single bonds. S is larger than O so p orbitals don’t overlap as well S-S . Double bond is weaker that 2 single bonds. S-S bonds are longer so e–-e– repulsion is lower. 8 Cl Cl a. C C C C Cl C C C C C C C C CH 3 CH 3 Cl b. C C C C Occurs in optically active forms. Cl Cl C 2 H5 C Vs CH 3 CH3 C C2 H 5 H H They differ in the direction they rotate plane polarized light. c. (i). Alkenes C4H8. (ii). C C C C C C C C C (iii). C C C C Exists in isomeric forms CH 3 CH 3 C H CH 3 and C H C C H C CH3 C C C C C Cl 2010 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM – PART III Prepared by the American Chemical Society Olympiad Laboratory Practical Task Force OLYMPIAD LABORATORY PRACTICAL TASK FORCE Steve Lantos, Chair, Brookline High School, Brookline, MA Linda Weber, Natick High School, Natick, MA John Mauch, Braintree High School, Braintree, MA Mathieu Freeman, Greens Farms Academy, Greens Farms, CT Erling Antony, Arrowhead Union High School, Hartland, WI Christie B. Summerlin, University of Alabama-Birmingham, Birmingham, AL DIRECTIONS TO THE EXAMINER–PART III The laboratory practical part of the National Olympiad Examination is designed to test skills related to the laboratory. Because the format of this part of the test is quite different from the first two parts, there is a separate, detailed set of instructions for the examiner. This gives explicit directions for setting up and administering the laboratory practical. There are two laboratory tasks to be completed during the 90 minutes allotted to this part of the test. Students do not need to stop between tasks, but are responsible for using the time in the best way possible. Each procedure must be approved for safety by the examiner before the student begins that procedure. Part III 2 lab problems laboratory practical 1 hour, 30 minutes Students should be permitted to use non-programmable calculators. DIRECTIONS TO THE EXAMINEE–PART III DO NOT TURN THE PAGE UNTIL DIRECTED TO DO SO. WHEN DIRECTED, TURN TO PAGE 2 AND READ THE INTRODUCTION AND SAFETY CONSIDERATIONS CAREFULLY BEFORE YOU PROCEED. There are two laboratory-related tasks for you to complete during the next 90 minutes. There is no need to stop between tasks or to do them in the given order. Simply proceed at your own pace from one to the other, using your time productively. You are required to have a procedure for each problem approved for safety by an examiner before you carry out any experimentation on that problem. You are permitted to use a non-programmable calculator. At the end of the 90 minutes, all answer sheets should be turned in. Be sure that you have filled in all the required information at the top of each answer sheet. Carefully follow all directions from your examiner for safety procedures and the proper disposal of chemicals at your examining site. Distributed by American Chemical Society, 1155 16th Street, N.W., Washington, DC 20036 All rights reserved. Printed in U.S.A. Property of ACS USNCO – Not for use as an USNCO National Exam after April 26, 2010 amount of substance ampere atmosphere atomic mass unit Avogadro constant Celsius temperature centi– prefix coulomb density electromotive force energy of activation enthalpy entropy equilibrium constant 1 1A 1 H n A atm u NA °C c C d E Ea H S K ABBREVIATIONS AND SYMBOLS Faraday constant F molar free energy G molar mass frequency ν mole gas constant R Planck’s constant gram g pressure hour h rate constant joule J reaction quotient kelvin K second kilo– prefix k speed of light liter L temperature, K measure of pressure mm Hg time milli– prefix m vapor pressure molal m VP volt volume CONSTANTS M M mol h P k Q s c T t R = 8.314 J·mol–1·K–1 R = 0.0821 L·atm·mol–1·K–1 1 F = 96,500 C·mol–1 1 F = 96,500 J·V–1·mol–1 NA = 6.022 × 1023 mol–1 h = 6.626 × 10–34 J·s c = 2.998 × 108 m·s–1 0 °C = 273.15 K V V PERIODIC TABLE OF THE ELEMENTS 18 8A 2 He 3 Li 2 2A 4 Be 13 3A 5 B 14 4A 6 C 15 5A 7 N 16 6A 8 O 17 7A 9 F 6.941 9.012 10.81 12.01 14.01 16.00 19.00 20.18 11 Na 12 Mg 13 Al 14 Si 15 P 16 S 17 Cl 18 Ar 22.99 24.31 26.98 28.09 30.97 32.07 35.45 39.95 19 K 31 Ga 32 Ge 33 As 34 Se 35 Br 36 Kr 1.008 4.003 10 Ne 20 Ca 3 3B 21 Sc 4 4B 22 Ti 5 5B 23 V 6 6B 24 Cr 7 7B 25 Mn 8 8B 26 Fe 9 8B 27 Co 10 8B 28 Ni 11 1B 29 Cu 12 2B 30 Zn 39.10 40.08 44.96 47.88 50.94 52.00 54.94 55.85 58.93 58.69 63.55 65.39 69.72 72.61 74.92 78.96 79.90 83.80 37 Rb 38 Sr 39 Y 40 Zr 41 Nb 42 Mo 43 Tc 44 Ru 45 Rh 46 Pd 47 Ag 48 Cd 49 In 50 Sn 51 Sb 52 Te 53 I 54 Xe 85.47 87.62 88.91 91.22 92.91 95.94 (98) 101.1 102.9 106.4 107.9 112.4 114.8 118.7 121.8 127.6 126.9 131.3 55 Cs 56 Ba 57 La 72 Hf 73 Ta 74 W 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 81 Tl 82 Pb 83 Bi 84 Po 85 At 86 Rn 132.9 137.3 138.9 178.5 180.9 183.8 186.2 190.2 192.2 195.1 197.0 200.6 204.4 207.2 209.0 (209) (210) (222) 87 Fr 88 Ra 89 Ac 104 Rf 105 Db 106 Sg 107 Bh 108 Hs 109 Mt 110 Ds 111 Rg 112 113 114 115 116 117 118 (223) (226) (227) (261) (262) (266) (264) (277) (268) (281) (272) (277) (Uut) (Uuq) (Uup) (Uuh) (Uus) (Uuo) Page 2 Uub 58 Ce 59 Pr 60 Nd 61 Pm 62 Sm 63 Eu 64 Gd 65 Tb 66 Dy 67 Ho 68 Er 69 Tm 70 Yb 71 Lu 140.1 140.9 144.2 (145) 150.4 152.0 157.3 158.9 162.5 164.9 167.3 168.9 173.0 175.0 90 Th 91 Pa 92 U 93 Np 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 100 Fm 101 Md 102 No 103 Lr 232.0 231.0 238.0 (237) (244) (243) (247) (247) (251) (252) (257) (258) (259) (262) Property of ACS USNCO -Not for use as an USNCO National Exam after April 26, 2010 2010 U. S. NATIONAL CHEMISTRY OLYMPIAD PART III – LABORATORY PRACTICAL Student Instructions Introduction These problems test your ability to design and carry out laboratory experiments and to draw conclusions from your experimental work. You will be graded on your experimental design, on your skills in data collection, and on the accuracy and precision of your results. Clarity of thinking and communication are also components of successful solutions to these problems, so make your written responses as clear and concise as possible. Safety Considerations You are required to wear approved eye protection and gloves at all times during this laboratory practical. You also must follow all directions given by your examiner for dealing with spills and with disposal of wastes. In particular, special precautions are required when using the silver nitrate solution in Problem #2. Do not get it on your skin or clothing as it will cause stains. Please wash off any chemicals spilled on your skin or clothing with large amounts of tap water. On Problem #2, neither solution ‘AgNO3 solution’ nor ‘K2CrO4 solution’ can be disposed of down the drain. Both solutions need waste beakers (provided) in which the used pipettes should be placed. Lab Problem 1 You have been given a well plate, several test tubes and pipets, a concentrated ammonia solution, access to distilled water, and four numbered vials containing iron (III) chloride hexahydrate, cobalt (II) sulfate heptahydrate, copper (II) chloride dihydrate, and potassium oxalate monohydrate, though not necessarily in this order. Devise and carry out an experiment to produce at least FIVE new different complex compounds, using your understanding of coordination compound geometry and qualitative evidence in your results. Lab Problem 2 You have been given a sample of seawater, a pipet that contains 5.0 × 10–4 mole/gram AgNO3, a pipet that contains K2CrO4(aq), several small test tubes, and access to an electronic balance. Devise and carry out an experiment to determine the percentage of chloride ion, Cl–(aq), in seawater. Ksp at 25 °C AgCl = 1.77 × 10–10 Ag2CrO4 = 1.12 × 10–12 Property of ACS USNCO -Not for use as an USNCO National Exam after April 26, 2010 Page 3 Answer Sheet for Laboratory Practical Problem 1 Student's Name: __________________________________________________________________________ Student's School:________________________________________ Date: ___________________________ Proctor's Name: _________________________________________________________________________ ACS Section Name:_________________________________Student's USNCO test #: ________________ 1. Give a brief description of your experimental plan. 2. Data and Observations. Vials #1 Contents as correctly written formulas Evidence #2 #3 #4 Before beginning your experiment, you must get Approval (for safety reasons) from the examiner Page 4 Examiner’s Initials: Property of ACS USNCO -Not for use as an USNCO National Exam after April 26, 2010 3. Show any relevant reactions and sketch the possible geometry of each formed compounds. 4. Conclusions Reactant(s) Used Proposed coordinate complex formula Evidence Property of ACS USNCO -Not for use as an USNCO National Exam after April 26, 2010 Page 5 Answer Sheet for Laboratory Practical Problem 2 Student's Name: __________________________________________________________________________ Student's School:________________________________________ Date: ___________________________ Proctor's Name: _________________________________________________________________________ ACS Section Name:_________________________________Student's USNCO test #: ________________ 1. Give a brief description of your experimental plan. 2. Record your data and all relevant balanced chemical equations. Before beginning your experiment, you must get Approval (for safety reasons) from the examiner Page 6 Examiner’s Initials: Property of ACS USNCO -Not for use as an USNCO National Exam after April 26, 2010 3. Calculations: 4. Conclusion. The percentage of chloride ion present in seawater = 5. List the assumptions that you made in this experiment: Property of ACS USNCO -Not for use as an USNCO National Exam after April 26, 2010 Page 7 2010 U. S. NATIONAL CHEMISTRY OLYMPIAD NATIONAL EXAM—PART III Prepared by the American Chemical Society Olympiad Laboratory Practical Task Force Examiner’s Instructions Directions to the Examiner: Thank you for administering the 2010 USNCO laboratory practical on behalf of your Local Section. It is essential that you follow the instructions provided, in order to insure consistency of results nationwide. There may be considerable temptation to assist the students after they begin the lab exercise. It is extremely important that you do not lend any assistance or hints whatsoever to the students once they begin work. As in international competition, the students are not allowed to speak to anyone until the activity is complete. The equipment needed for each student for both lab exercises should be available at his/her lab station or table when the students enter the room. The equipment should be initially placed so that the materials used for Lab Problem 1 are separate from those used for Lab Problem 2. After the students have settled, read the following instructions (in italics) to the students. Hello, my name is ________ . Welcome to the lab practical portion of the U.S. National Chemistry Olympiad Examination. In this part of the exam, we will be assessing your lab skills and your ability to reason through a laboratory problem and communicate its results. Do not touch any of the equipment in front of you until you are instructed to do so. You will be asked to complete two laboratory problems. All the materials and equipment you may want to use to solve each problem has been set out for you and is grouped by the number of the problem. You may use equipment from one problem to work on the other problem, but the suggested ideal equipment and chemicals to be used for each problem has been grouped for you. You will have one hour and thirty minutes to complete the two problems. You may choose to start with either problem. You are required to have a procedure for each problem approved for safety by an examiner. (Remember that approval does not mean that your procedure will be successful – it is a safety approval.) When you are ready for an examiner to come to your station for each safety approval, please raise your hand. Safety is an important consideration during the lab practical. You must wear goggles and gloves at all times. In particular, special precautions are required when using the silver nitrate solution in Problem #2. Do not get it on your skin or clothing as it will cause stains. Please wash off any chemicals spilled on your skin or clothing with large amounts of tap water. On Problem #2, neither solution ‘AgNO3 solution’ nor ‘K2CrO4 solution’ can be disposed of down the drain. Both solutions need waste beakers (provided) in which the used pipettes should be placed. The appropriate procedures for disposing of solutions at the end of this lab practical are: ____________________________________________________________________________________ ____________________________________________________________________________________ ____________________________________________________________________________________ We are about to begin the lab practical. Please do not turn the page until directed to do so, but read the directions on the front page. Are they any questions before we begin? Property of ACS USNCO - 2010 National Exam Part III Examiner’s Notes Page 1 Distribute Part III booklets and again remind students not to turn the page until the instruction is given. Part III contains student instructions and answer sheets for both laboratory problems. There is a periodic table on the second page of the booklet. Allow students enough time to read the brief cover directions. Do not turn to page 2 until directed to do so. When you start to work, be sure to fill out all of the information at the top of the answer sheets. Are they any additional questions? If there are no further questions, the students should be ready to start Part III. You may begin. After one hour and thirty minutes, give the following directions. This is the end of the lab practical. Please stop and bring me your answer sheets. Thank you for your cooperation during this portion of the exam. Collect all the lab materials. Make sure that the student has filled in his or her name and other required information on the answer sheets. At this point, you might wish to take a few minutes to discuss the lab practical with the students. They can learn about possible observations and interpretations and you can acquire feedback as to what they actually did and how they reacted to the problems. After this discussion, please take a few minutes to complete the Post-Exam Questionnaire; this information will be extremely useful to the USNCO subcommittee as they prepare for next year’s exam. Please remember to return the post-exam Questionnaire, the answer sheets form Part III, the Scantron sheets from Part I, and the ‘Blue Books” from Part II in the overnight return envelope you were provided to this address: American Chemical Society U.S. National Chemistry Olympiad Office 1155 16th Street, NW Washington, DC 20036 The label on the UPS Express Pak envelope should have this address and your return address already. The cost of the shipping is billed to ACS - USNCO. You can keep copy of the tracking number to allow you to track your shipment. Wednesday, April 28, 2010, is the absolute deadline for receipt of the exam material. Materials received after this deadline CANNOT be graded. Be sure to have your envelope sent no later than Tuesday, April 27, 2010 for it to arrive on time. THERE WIL BE NO EXCEPTIONS TO THIS DEADLINE DUE TO THE TIGHT SCHEDULE FOR GRADING THIS EXAMINATION. Property of ACS USNCO - 2010 National Exam Part III Examiner’s Notes Page 2 Lab Problem #1: Materials and Equipment Each student should have available the following equipment and materials: Materials • Three standard size (18 or 20 × 150mm) test tubes • One 150 or 250 mL beaker to hold the test tubes • Four 5 mL capacity Beral-style pipets • Four pieces of waxed weighing paper • One 100 or 150 mL Erlenmeyer flask with stopper • One 12-hole white spot plate (porcelain or polyurethane) • Several wooden stirring sticks, the kind used for stirring coffee • Access to distilled water, preferably individual bottles at each lab station Chemicals • Four capped numbered vials (20–30 mL capacity) containing approximately 3–5 g of each of the following solids: FeCl3⋅6H2O, CoSO4⋅7H2O, CuCl2⋅2H2O, and K2C2O4⋅H2O Vial #1 #2 #3 #4 Substance FeCl3⋅6H2O CoSO4⋅7H2O CuCl2⋅2H2O K2C2O4⋅H2O Obviously, do not identify or label the contents of each vial! • 25–30 mL of concentrated ammonia solution per student. To make this solution for a group of 25–30 students, using stock (14.8M) ammonium hydroxide, NH4OH(aq) mix 250 mL of the ammonium hydroxide with 500 mL water to make a total volume of 750 mL ammonia solution. The 100 or 150 mL flask should be labeled ‘Ammonia Solution’. Lab Problem #1 Notes • Be sure that the solids are powdered and dry. In the case of the iron (III) chloride you may want to grind it fine as it has a tendency to clump. • Students will be warned that the ammonia solution is pungent and to take caution as it is used. Safety: It is your responsibility to ensure that all students wear safety goggles and gloves during the lab practical. A lab coat or apron for each student is desirable but not mandatory. You will also need to give students explicit directions for handling spills and for disposing of waste materials, following approved safety practices for your examination site. Please check and follow procedures appropriate for your site. Property of ACS USNCO - 2010 National Exam Part III Examiner’s Notes Page 3 Lab Problem #2: Materials and Equipment Each student should have available the following equipment and materials: Materials • Three small (10 × 75 mm) test tubes • One 50 mL beaker to hold the test tubes • One 100 or 150 mL beaker to hold the labeled silver nitrate, seawater, and potassium chromate pipets • Three 5 mL Beral–style pipet (graduated or ungraduated, thin or regular stem) Chemicals • Approximately 30 mL of seawater in a covered 100 or 150 mL flask labeled ‘seawater’ • Two 5 mL capacity Beral-style pipets filled with 5.0 ×10–4 mol AgNO3 per gram solution. To make this solution, dissolve 2.13 g of solid AgNO3 in 25.0 mL DI water. Store in a dark, covered container until ready to apportion to students. Pipets should be labeled ‘AgNO3 solution’. It should be clearly labeled "Caution - Do not get on Skin or Clothes - Will cause Stains" • One Beral-style pipet containing potassium chromate solution • The concentration of this solution should be approximately 1M (Mr K2CrO4 = 194). This pipet should be labeled ‘K2CrO4 solution’. To make this solution, dissolve 0.4 g K2CrO4 / 2 mL per student or 10 g in 50 mL for 25 students. • Neither solution ‘AgNO3 solution’ nor ‘K2CrO4 solution’ can be disposed of down the drain. Both solutions need waste beakers in which the used pipettes are placed. Lab Problem #2 Notes: • For the ocean water, if you live near the coast, obtain a sample directly from the ocean. Let any solids settle before apportioning to students. If you do not have an available source of ocean water, you may simulate a sample by dissolving 2.80 g of NaCl in 100 mL of distilled water (don’t use tap water as it likely already contains measurable amounts of chloride). You may also use common commercial laboratory seawater preparations that include approximate percentages of minerals found in ocean water. • Make sure to use DI water in making the silver nitrate solution. Ideally, this solution is made just prior to use the day of the exam. • The concentration of the potassium chromate solution is not critical but should be approximately 1M. Safety: It is your responsibility to ensure that all students wear safety goggles and gloves during the lab practical. A lab coat or apron for each student is desirable but not mandatory. You will also need to give students explicit directions for handling spills and for disposing of waste materials, following approved safety practices for your examination site. Please check and follow procedures appropriate for your site. Property of ACS USNCO - 2010 National Exam Part III Examiner’s Notes Page 4 2010 USNCO Part III Answers Lab Problem 1 Students were expected to create different complex compounds by combinations of the hydrate compounds provided with the aqueous ammonia solution. Students might have thought also to react the hydrates with an aqueous oxalate solution or to slowly add water to samples of each hydrate. For example, adding H2O slowly to a small amount of copper(II) chloride in one of the test tubes produces a color change as the dihydrate of this salt forms a tetrahydrate. Vial 1 2 3 4 Compound Yellow-Orange, FeCl3 . 6H2O Pinkish, CoSO4 . 7H2O Green, CuCl2 . 2H2O White, K2C2O4 . H2O Compound . CuCl2 2H2O + H2O CuCl2 . 2H2O + ammonia solution . CoSO4 7H2O + ammonia solution . FeCl3 6H2O + ammonia solution . K2C2O4(aq) + FeCl3 6H2O . K2C2O4(aq) + CuCl2 2H2O Possible Formula . CuCl2 4H2O Cu(NH3)4 Observation green Æ blue deep blue color change Co(NH3)6 bright blue-green color change Fe(NH3)6 Brown-purple gelatinous ppt. K3[Fe(C2O4)3] K2[Cu(C2O4)2] greenish color change greenish color change? Ammonia replaces water in these hydrated compounds to produce complexes that include NH3 surrounding the central metal cation. Possible formulas might also include ligand combinations of ammonia, water, and chloride or sulfate as part of the complex. Possible sketches of these structures show the central Cu, Fe, and Co cations surrounded by the NH3 ligands in tetrahedral, square planar geometry, or octahedral geometries. Excellent Results: Students attempted as many matches as possible, creating a legible and systematically organized table to communicate their results. The nomenclature of the possible complexes was clear and the formulas used the proper coordinate complex compound was accurate. Drawings for the possible structures included three-dimensional representations of the complexes and indicated how and where isomers might also be formed. It was drawn/noted that the oxalate ion was a bidentate and three ions were able to coordinate with the iron (III) cation. 1 Average Results: Students attempted to create and draw only the requested five complexes. A listing of these combinations and the observed results was provided. At least one possible complex geometry was shown. Below Average Results: Students created fewer than the requested five complexes. No systematic table or chart was provided to indicate the possible compounds formed. One or no possible geometries were provided. Errors in the nomenclature and formulas were evident. Lab Problem 2 Sample Data: Initial mass of the seawater and pipet Final mass of the seawater and pipet Initial mass of silver nitrate and pipet Final mass of silver nitrate and pipet 3.121 g 2.243 g 2.820 g 1.918 g Sample Student Calculations: 3.121 g – 2.243 g = 0.878 g seawater used 2.820 g – 1.918 g = 0.902 g silver nitration solution used Molar masses: NaCl = 58.44 g/mol AgNO3 = 169.87 g/mol 0.902 g AgNO3 x 5.0 x 10-4 mol/g AgNO3 = 4.51 x 104 mole AgNO3 4.51 x 104 mole AgNO3 = mole Ag+ ; mole Cl4.51 x 104 mole Cl- x 35.5 g/mol Cl- = 1.60 x 10-2 g Cl1.60 x 10-2 g Cl-/0.878 g seawater x 100 = 1.82% Cl- in solution by mass (A value of 1.80-1.90% is generally accepted) Excellent Results: At least two trials were attempted using both of the filled AgNO3 pipets and averaging results. These titrations were completed in the small test tubes provided. A clear and organized table showed the initial and final masses of both the seawater and silver nitrate pipets. Observations of the silver chloride and the silver chromate formation were noted. Calculations were legible, organized, and followed a logical sequence in order to determine the mass of chloride ion 2 present and the final mass percentage present. Use of the Ksp values was made to clearly show the AgCl then Ag2CrO4 selective order of precipitation. Assumptions included the complete precipitation of AgCl from solution, the equivalent molar ratio of Ag+: Cl-, the presence of a drop or two of the chromate solution to the seawater to be sufficient to create a white/red color change endpoint. Error included the subjective judgment of the endpoint to distinguish the complete precipitation of the chloride from the beginning of the precipitation of the chromate, and the additional drop or two overshooting of the endpoint to turn the solution entirely reddish. Average Results: Only one trial was attempted. Leftover materials were not utilized to complete additional experiments. Observations were made to determine the endpoint. Calculations show a logical sequence to solve this problem, but may not have been clearly organized. Only one point was included for the assumptions and error. Below Average Results: Not all of the materials were used. Calculations did not show evidence of logical sequence in problem solving. Observations were incomplete. One or no points were made regarding error or assumptions. 3