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UT CH 302 - CH302 Practice Exam 2

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Spring 2009 CH302 Practice Exam 2 1. What would be the pH of a solution prepared by dissolving 120.1 g of CH3COOH and 82 g of NaCH3COO in 1 L of water? Acetic acid has a Ka of 1.8 x 10-5. 1. 5.05 2. 4.78 3. 4.12 4. 4.44 2. Which of the following pairs of solutions would not result in a buffer upon mixing? 1. 100 mL of 10 mM NaOH & 80 mL of 20 mM NH4Cl 2. 20 mL of 0.3 M NaF & 12 mL of 0.4 M HCl 3. 0.4 L of 10 mM HClO3 & 0.5 L of 8 mM C6H5NH2 4. 2 L of 1.35 M Ba(OH)2 & 3 L of 2 M CHOOH Explanation: A buffer prepared by a neutralization reaction requires a weak acid mixed with less strong base or a weak base mixed with less strong acid. The only pair of solutions which fails to satisfy this constraint is 0.4 L of 10 mM HClO3 and 0.5 L of 8 mM C6H5NH2. 3. Consider the following acids and their provided pKas. Rank them in terms of increasing strength of their conjugate bases. CH3COOH pKa = 4.75 CH3CHOHCOOH pKa = 3.85 CHOOH pKa = 3.74 CH3CH2COOH pKa = 4.88 1. CHOOH < CH3CHOHCOOH < CH3COOH < CH3CH2COOH 2. CH3CH2COOH < CHOOH < CH3CHOHCOOH < CH3COOH 3. CH3COOH < CH3CH2COOH < CHOOH < CH3CHOHCOOH 4. CH3CHOHCOOH < CH3COOH < CH3CH2COOH < CHOOH 4. Which of the following buffers could absorb the greatest amount of strong base before being exhausted? 1. 45 mL of 2 mM N2H5Cl, 4 mM N2H4 2. 3.2 L of 0.4 M HClO, 0.5 NaClO 3. 2 L of 9 mM HF, 7 mM NaF 4. 0.3 L of 0.4 M NH4Cl, 0.6 M NH3 5. 20 mL of 5 M CHOOH, 4 M NaCHOO 5. If one added 200 mL of 6 M HCl to 1 L of a buffer composed 4.2 M CH3COOH and 6.6 M NaCH3COO, what would be the resulting pH? The Ka of CH3COOH is 1.8 x 10-5. 1. 5.3 2. 4.9 3. 5.1 4. 4.7 6. How many buffer regions and equivalence points would be visible on the titration curve of a weak tetraprotic acid? 1. 3, 1 2. 3, 4 3. 1, 4 4. 4, 1 5. 4, 4 7. A 100 mL sample of 0.1 M H3PO4 is titrated with 0.2 M NaOH. What is the pH of the solution after 100 mL of NaOH has been added? Phosphoric acid has Ka1 = 7.5 x 10-3, Ka2 = 6.2 x 10-8 and Ka3 =2.1 x 10-13. 1. 4.10 2. 8.51 3. 4.67 4. 7.40 5. 9.94 8. What will be the pH at the first equivalence point of a titration of 0.2 M H2SO4 with 0.2 M NaOH? The Ka for HSO4- is 2 x 10-2. 1. 1.45 2. 1.35 3. 7.00 4. not enough information 9. All of the salts below have the same approximate molar solubility except for one. Which is it? 1. TlBr Ksp = 4.00 x 10-6 2. PbI2 Ksp = 7.47 x 10-9 3. AgSCN Ksp = 1.16 x 10-12 4. CsIO4 Ksp = 5.16 x 10-6 10. The Ksp of MgNH4PO4 at 25 °C is 2.5 x 10-13. What is its molar solubility at this temperature? (Hint: do the RICE diagram for this one.) 1. 3.2 x 10-4 2. 4.0 x 10-5 3. 6.3 x 10-5 4. 1.2 x 10-3 11.2 What would be the molar solubility of Sn(OH)2 (Ksp = 10-26) in pH 13 NaOH solution? 1. 1 x 10-24 2. 4 x 10-24 3. 1 x 10-28 4. 4 x 10-28 5. not enough information 12. Consider the table below. Which anion would be the best for separating Pb2+ from Ca2+? Which would be the worst? Ksp values C2O4- CO32- SO42- IO3- Pb2+ 2.74 x 10-11 3.3 x 10-14 1.6 x 10-8 1.2 x 10-13 Ca2+ 2.57 x 10-9 8.7 x 10-9 4.93 x 10-5 6.44 x 10-7 1. C2O4- & SO42- 2. IO3- & SO42- 3. CO32- & IO3- 4. IO3- & C2O4- 5. CO32- & C2O4- 13. A student used the equation [H+] = (Ka·Ca)1/2 to calculate [H+] and got a value of 0.4 M. The actual value was determined experimentally to be 0.35 M. Which of the following are possibleexplanations for this discrepancy? I. Ka was too small II. Ka was too large III. Ca was too small 1. I only 2. II only 3. III only 4. I & II 5. I & III 6. II & III 7. none 14. For a solution of H3PO4, addition of Na2HPO4 will increase the concentration of which of the following species? I. H3PO4 II. H2PO4- III. PO43- 1. I only 2. II only 3. III only 4. I & II 5. I & III 6. II & III 7. I, II and III 15. Determine the pH of a 5 M solution of Na2HPO4. Assume H3PO4 has a pKa1 of 2.1, a pKa2 of 7.2 and a pKa3 of 12.7. 1. 9.95 2. 4.65 3. 7.4 4. not enough information 16. Write a mass balance for carbon for a solution that initially contains H2CO3. 1. CH2CO3 = [HCO3-] + [CO32-] 2. CH2CO3 = [H2CO3] + [HCO3-] + [CO32-] 3. CH2CO3 = [CO2] + [H2CO3] + [HCO3-] + [CO32-] 4. CH2CO3 = [CO2] + [H2CO3] 17. How many equation are necessary to define a system initially composed of MgNH4PO4? 1. 9 2. 8 3. 7 4. 5 18. Which of the following would be equal to Ka1 times Ka2 for orthocarbonic acid, H4CO4? 1. [H2CO42-]·[H+]/[H4CO4] 2. [H3CO4-]·[H+]/[H4CO4] 3. [H2CO42-]·[H+]/[H3CO4-] 4. [H2CO42-]·[H+]·[H3CO4-]/[H4CO4] 5. [H2CO42-]·[H+]2/[H4CO4]19. What would be the pH of a 2 x 10-8 M solution of Ba(OH)2? 1. 7.009 2. 7.019 3. 7.013 4. 7.004 20. What would be the [H+], [HSO4-] and [SO42-] in a 1 M solution of H2SO4? 1. 1.02, 0.98, 0.02 M, respectively 2. 0.00, 2.00, 1.00 M, respectively 3. 1.14. 0.86, 0.14 M, respectively 4. 0.14, 1.00, 0.14 M, respectively 21. What would be the pH of a 4 mM M Na3C6H5O7 solution (trisodium citrate)? Citric acid has Ka1 = 7.1 x 10-4, Ka2 = 1.7 x 10-5 and Ka3 = 4.0 x 10-7. 1. 6.62 2. 5.00 3. 9.00 4. 7.38 5. 7.00 22. Fully balance the reaction below in acid. How many protons are needed? How many water molecules? (Hint: this one is tricky - the water molecules and protons go on the same side.) N2(g) ֊ 2 NH4OH(aq) 1. 3, 1 2. 6, 2 3. 2, 2 4. 8, 2 5. 8, 1 23. Fully Balance the reaction below in acid. What is the sum of the coefficients? Zn(s) + MnO2(s) + NH4Cl(aq) ֊ ZnCl2(s) + Mn2O3(s) + NH3(aq) 1. 7 2. 12 3. 6 4. 10 5. 9 24. Which of the following statements is untrue concerning ranking the strength/weakness of oxidizing/reducing agents. 1. A reactant with a high reduction potential is a good reducing reagent. 2. A product with a low reduction potential is a good reducing reagent. 3. A reactant with a low reduction potential is a poor oxidzing reagent. 4. A product wit a high reduction potential is a poor …


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