CH Practice Exam #2

Transcription

1 CH Practice Exam #2 Part I - Multiple Choice - Choose the best answer and place the letter corresponding to the answer in the space provided. 1. Consider the following reaction and its equilibrium constant: SO2(g) + NO2(g) SO3(g) + NO(g) Kc = 0.33 A reaction mixture contains 0.41 M SO2, 0.14 M NO2, 0.12 M SO3 and 0.14 M NO. Which of the following statements is TRUE concerning this system? A) The reaction will shift in the direction of reactants. B) The equilibrium constant will decrease. C) The reaction will shift in the direction of products. D) The reaction quotient will decrease. E) The system is at equilibrium. 2. What is the conjugate base of H2PO4? A) HPO4 2- B) PO4 2- C) H3PO4 D) H3O+ E) OH 3. Identify the weak diprotic acid. A) CH3COOH B) HCOOH C) H3PO4 D) H2SO4 E) H2CO3 4. Which of the following solutions would have the highest ph? Assume that they are all 0.10 M in acid at 25ºC. The acid is followed by its Ka value. A) HF, B) HCN, C) HNO2, D) HCHO2, E) HClO2, Calculate the hydronium ion concentration in an aqueous solution with a poh of A) M B) M C) M D) M E) M

2 6. A solution with a hydrogen ion concentration of M is and has a hydroxide ion concentration of. A) acidic, M B) acidic, M C) basic, M D) basic, M 7. Which of the following is a STRONG base? A) I- B) NH3 C) CH3OH D) NO3 E) LiOH 8. If an equal number of moles of the weak acid HCN and the strong base KOH are added to water, is the resulting solution acidic, basic, or neutral? A) acidic B) basic C) neutral D) There is insufficient information provided to answer this question 9. Which of the following is TRUE? A) An effective buffer has a [base]/[acid] ratio in the range of B) A buffer is most resistant to ph change when [acid] = [conjugate base] C) An effective buffer has very small absolute concentrations of acid and conjugate base. D) A buffer can not be destroyed by adding too much strong base. It can only be destroyed by adding too much strong acid. E) None of the above are true 10. The molar solubility of ZnS is M in pure water. Calculate the Ksp for ZnS. A) B) C) D) E) Give the equation for a saturated solution in comparing Q with Ksp. A) Q > Ksp B) Q < Ksp C) Q = Ksp D) Q Ksp E) none of the above

3 12. Which of the following solutions is a good buffer system? A) A solution that is 0.10 M NaCl and 0.10 M HCl B) A solution that is 0.10 M HCN and 0.10 M LiCN C) A solution that is 0.10 M NaOH and 0.10 M HNO3 D) A solution that is 0.10 M HNO3 and 0.10 M NaNO3 E) A solution that is 0.10 M HCN and 0.10 M KI 13. You wish to prepare an HC2H3O2 buffer with a ph of If the pka of HC2H3O2 is 4.74, what ratio of C2H3O2 /HC2H3O2 must you use? A) 0.10 B) 0.40 C) 0.40 D) 2.51 E) Which of the following processes have a ΔS > 0? A) CH3OH(l) CH3OH(s) B) N2(g) + 3 H2(g) 2 NH3(g) C) CH4(g) + H2O(g) CO(g) + 3 H2(g) D) Na2CO3(s) + H2O(g) + CO2(g) 2 NaHCO3(s) E) All of the above processes have a ΔS > Consider a reaction that has a positive ΔH and a positive ΔS. Which of the following statements is TRUE? A) This reaction will be spontaneous only at high temperatures. B) This reaction will be spontaneous at all temperatures. C) This reaction will be nonspontaneous at all temperatures. D) This reaction will be nonspontaneous only at high temperatures. 16. Below what temperature does the following reaction become nonspontaneous? 2 HNO3(aq) + NO(g) 3 NO2(g) + H2O(l) ΔH = kj; ΔS = J/K A) 39.2 K B) 151 K C) 475 K D) This reaction is nonspontaneous at all temperatures. E) This reaction is spontaneous at all temperatures.

4 17. Calculate ΔS rxn for the following reaction. The S for each species is shown below the reaction. N2H4(l) + H2(g) 2 NH3(g) S (J/mol K) A) J/K B) J/K C) J/K D) J/K E) J/K 18. Estimate ΔG rxn for the following reaction at 825 K. 2 Hg(g) + O2(g) 2 HgO(s) ΔH = kj; ΔS = J/K A) kj B) -37 kj C) +37 kj D) +645 kj E) -645 kj 19. What element is being oxidized in the following redox reaction? Cr(OH)4 (aq) + ClO (aq) CrO4 2- (aq) + Cl (aq) A) Cr B) O C) H D) Cl 20. Determine the cell notation for the redox reaction given below. 3 Cl2(g) + 2 Fe(s) 6 Cl (aq) + 2 Fe3+(aq) A) Cl2(g) Cl (aq) Pt Fe(s) Fe3+(aq) B) Cl (aq) Cl2(g) Pt Fe3+(aq) Fe(s) C) Fe3+(aq) Fe(s) Cl (aq) Cl2(g) Pt D) Fe(s) Cl2(g) Fe3+(aq) Cl (aq) Pt E) Fe(s) Fe3+(aq) Cl2(g) Cl (aq) Pt

5 21. Use the standard half-cell potentials listed below to calculate the standard cell potential for the following reaction occurring in an electrochemical cell at 25 C. (The equation is balanced.) Pb(s) + Br2(l) Pb2+(aq) + 2 Br (aq) Pb2+(aq) + 2 e Pb(s) Br2(l) + 2 e 2 Br (aq) E = V E = V A) V B) V C) V D) V E) V 22. Which of the following is TRUE? A) If Q>K, Ecel > 0l B) If Q=1, Ecell = Eºcell C) If Q<K, Ecel <0 D) If ΔG > 0, Eºcell>0 E) If ΔG < 0, Eºcell<0 23. A galvanic cell consists of one half-cell that contains Ag(s) and Ag+(aq), and one half-cell that contains Cu(s) and Cu2+(aq). What species are produced at the electrodes under standard conditions? Ag+(aq) + e- Ag(s) E = V Cu2+(aq) + 2 e- Cu(s) E = V A) Ag(aq) is formed at the cathode and, Cu(s) is formed at the anode. B) Ag(s) is formed at the cathode, and Cu2+(aq) is formed at the anode. C) Cu(s) is formed at the cathode, and Ag+(aq) is formed at the anode. D) Cu2+(aq) is formed at the cathode, and Cu(s) is formed at the anode. 24. Describe the reactions during the electrolysis of water. A) Oxygen is reduced and hydrogen is oxidized. B) Oxygen and hydrogen are both oxidized. C) Oxygen and hydrogen are both reduced. D) Oxygen is oxidized and hydrogen is reduced. E) Neither oxygen or hydrogen are oxidized or reduced.

6 Part II (Short Answer) - [For the real exam, you will not be given answers to choose from.] Show all of your calculations for full credit. 1. Determine the [H3O+] in a M HClO solution. The Ka of HClO is A) M B) M C) M D) M E) M 2. Calculate the ph of a solution formed by mixing ml of 0.15 M NH4Cl with ml of 0.20 M NH3. The Kb for NH3 is 1.8 x A) 9.13 B) 9.25 C) 9.53 D) 4.74 E) A ml sample of 0.10 M NH3 is titrated with 0.10 M HNO3. Determine the ph of the solution after the addition of 50.0 ml of HNO3. The Kb of NH3 is A) 4.74 B) 7.78 C) 7.05 D) 9.26 E) Determine the molar solubility of AgBr in a solution containing M NaBr. Ksp (AgBr) = A) M B) M C) M D) M E) M 5. Use the tabulated half-cell potentials to calculate ΔG for the following balanced redox reaction. 3 I2(s) + 2 Fe(s) 2 Fe3+(aq) + 6 I (aq) A) -1.1 x 102 kj B) +4.9 x 101 kj C) -9.7 x 101 kj D) +2.3 x 102 kj E) -3.3 x 102 kj

7 Help equations: aa + bb cc + dd Kw =[H3O+][OH-]=1.0 x Kw/[OH-]= [H3O+] Kw/[H3O+]= [OH-] ph = -log[h3o+] poh = -log[oh-] pka = -logka Ka = Kw/Kb Kb = Kw/Ka ΔE = q + w ΔE = ΔH + PΔV E is a state function eat has S univ = S sys + S surr > 0 ΔS rxn = ΔS sys = S final S initial ΔSº rxn = Σn products Sº products Σn reactants Sº reactants ΔGº reaction = ΔHº reaction - TΔSº reaction ΔGº rxn = 0, and Q = K ΔG rxn = RT ln K (K = e ΔG/RT ) Eºcell = Eºcathode Eºanode Eºcell = Eºoxid Eºred ΔGº = nfeºcell