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CHEM 1301 POSSIBLE EXTRAS REVIEW Chapter Ten Stuff Differences between solids, liquids and gases. Intermolecular Forces – important in solid state Stronger they are, higher the melting and boiling points: 1. IONIC all ionic compounds have IONIC intermolecular forces 2. METALLIC all metals have METALLIC intermolecular forces Both of these are “infinite” networks – go on and on and on COVALENT molecules have: 3. HYDROGEN BONDS for H attached to F, O, N (or Cl) 4. DIPOLE-DIPOLE if the molecule is polar [has a dipole – opposite ends] 5. LONDON DISPERSON if the molecule is not polar. i. e., has no polar bonds or the polar bonds cancel out. Typical questions: For each of the following solids, assign whether the intermolecular forces are (A) (D) ionic dipole-dipole (B) (E) 1. 5. 9. NaCl H2O KBr 11. Which of the following would be expected to have the highest boiling point? (A) NaCl (B) 12. Which substance has the highest boiling point? (A) CH4 13. Which group of substances is arranged in order from the highest to the lowest melting point? (A) HF>H2>NaF (B) 2. 6. 10. Na (B) metallic London dispersion 3. 7. Cl2 NH3 CH4 (C) He Cl2 (C) Hg FBr (D) all the same (C) HF NaF>H2>HF (C) hydrogen bonding 4. 8. (E) can’t tell (D) Cl2 HF>NaF>H2 (D) 1 Fe CO2 NaF>HF>H2 14. Arrange KCl, NH3, and CH4 in order of increasing boiling point. (A) (C) CH4<KCl<NH3 CH4<NH3<KCl 15. Which has the highest melting point? (A) S8 16. Which of the following molecules will not form hydrogen bonds? (B) NH3<KCl<CH4 (D) NH3<CH4<KCl (B) H2O (C) Ar (D) BaF2 O (A) H3C CH2 C (B) OH HF H O H3C (C) CH2 C (D) CH3 H3C CH2 N 17. In hydrogen iodide __________________ are the most important intermolecular forces. (A) (C) dipole-dipole forces hydrogen bonding (B) (D) CH3 London dispersion forces ionic GASES Pressure Temperature Volume n measured in atmospheres or mmHg [1 atm = 760 mm Hg] measured in Kelvin [K = °C + 273] measured in liters moles Pressure is proportional to temperature and number of moles and inversely proportional to volume Put together the relationships using the correct units: PV = nRT [R = 0.082 l atm / mol K] 18. When the temperature in Kelvin of a fixed quantity of an ideal gas is quadrupled and the pressure is doubled, what is the net effect on the volume of the gas? (A) (C) The volume remains constant. The volume increases four–fold. (B) (D) The volume increases two–fold. The volume increases eight–fold. 2 19. Air is sealed in a vessel at 273 °C and then cooled to 0 °C. If the vessel itself does not contract, the pressure inside the vessel will become (A) (D) twice its original value. one-fourth of its original value. 20. Both the pressure and the absolute temperature of a certain gas sample are doubled. In the absence of dissociation, the volume of the gas is (A) (D) quadrupled. decreased by one–half. 21. If a gas at 1 atm and 273 K occupies 3.36 L, how many moles does it contain? (A) 0.15 22. A gas sample occupies a volume of 16.4 L at 27 °C and 0.300 atm. How many moles of gas are present? (A) 0.200 23. What is the volume of 2.00 mol of helium gas at 27 °C and 3.00 atm? (A) 6.1 x10–2 L 24. How many moles of gas are in a sample with a volume of 500 mL at 25 °C and 0.460 atm? (A) 0.00941 mol (B) (B) (B) (E) 0.81 (B) none of these. (C) one-half of its original value. decreased by one–fourth. unchanged. (C) 0.450 (B) (B) (E) 3.4 (C) 16.4 L 3.50 (C) 10.2 mol 1.48 L (C) 0.160 mol (C) (D) zero. doubled. 6.7 (D) 10.0 (D) 44.8 L (D) 13.4 mol This can be combined into the Roadmap: grams of A g grams of B g Particles NB Particles NA n = g / MM N = 6.02 x 1023 n 23 N = 6.02 x 10 n Moles of A nA Volume of Gas, V at Temp, T, and Pressure, P nB = (b/a)nA pV = nRT M=n/L Volume of Solution, L Moles of B nB Volume of Gas, V at Temp, T, and pV = nRT Pressure, P Volume of Solution, L 3 25. What is the number of molecules in 1.00 mL of an ideal gas at 1 atm and 273 K? (A) 2.69 x 1022 26. How many grams of CO2 would occupy an 8.8 L flask at 300 K and 1.1 atm? (A) 14 g 27. What volume will 5.10 g of sulfuryl fluoride, SO2F2, occupy at 1 atm and 273 K? (A) 0.056 L 28. (B) (B) (B) 1.12 L 2.69 x 1019 17 g (C) 2.24 L 6.02 x 1020 (C) (C) 22 g (D) 11.2 L (D) (D) 2.24 x 1019 32 g (E) 22.4 L What volume does 16.00 g of oxygen gas (O2) occupy at 546 °C and 2.00 atm? (A) 5.6 L (B) 11.2 L (C) 22.4 L (D) 67.2 L (E) 16.8 L 29. When 18.0 g of water is heated to steam at 100 °C, the volume at 1 atm pressure is approximately?: (A) 13.2 mL 30. Benzene, C6H6, can be burned in oxygen according to the equation:, 2C6H6(l) + 15O2(g) → 12CO2(g) + 6H2O(g). If 5.0 L of oxygen measured at 1 atm and 273 K were required to burn a given amount of benzene, the 1 atm and 273 K volume of CO2 formed would be (A) 17.9 L 31. What mass of CaCO3 will produce 8.0 L of CO2 (measured at 1 atm and 273 K) in the reaction, CaCO3(s) → CaO(s) + CO2(g) ? (A) 4.46 g 32. How many liters of CO2 gas at 1 atm and 273 K can be obtained by completely burning one C3H8(g) + O2(g) → CO2(g) + H2O(g) mole of C3H8? (A) 11.2 33. What volume of pure N2O at 1 atm and 273 K could be prepared by the controlled decomposition of 8.00 g of ammonium nitrate - NH4NO3(s) → N2O(g) + 2H2O(l) (A) 1.12 L (B) (B) (B) (B) (B) 18.6 L 4 L 12.5 g 44.8 2.24 L (C) 1800 mL (C) 33.4 (C) 35.7 g (C) 67.2 (C) 3.36 L 4 (D) (D) (D) 11 L (E) 5.5 L 280 g (D) 112 (D) 30.6 L 4.48 L 34. How many liters of hydrogen gas with an excess of nitrogen at 1 atm and 273 K are required to prepare 45.0 g of ammonia? 3 H2(g) + N2(g) → 2 NH3(g) (A) 39.5 35. How many liters of hydrogen at 1 atm and 273 K can be produced by the reaction of 9.00 g of Al with excess dilute H2SO4: 2Al(s) + 3H2SO4(aq) → Al2(SO4)3(aq) + 3H2(g) (A) 5.61 L (B) 30.0 (B) (C) 11.2 L 59.3 (C) (D) 33.6 L (D) 88.9 67.2 L More Reactions We started talking about REDOX equations. You need to be able to: a) b) c) assign oxidation numbers: look at an equation and say whether it is a redox equation by change in charge or oxidation number if redox, which atom is reduced and which is oxidized 36. Calculate the oxidation number of the chlorine in perchloric acid, HClO4, a strong acid. (A) -1 37. Calculate the oxidation number of sulfur in sodium metabisulfite, Na2S2O5. (A) -2 38. Sodium tripolyphosphate is used in detergents to make them effective in hard water. Calculate the oxidation number of phosphorus in Na5P3O10. (A) +3 39. Calculate the oxidation number of iodine in I2. (A) -1 40. The oxidation numbers of P, S and Cl in H2PO2-, H2S and KClO4 are, respectively (A) (E) -1, -1, +3 -1, -2, +3 41. Identify which species is reduced in the following redox reaction. (B) (B) (B) + Cu(s) Cu(s) (D) +10 (C) +1, -2, +7 (D) +4 (C) 0 (B) +5 (C) +5 (B) Hg2+(aq) (C) +2 (B) Hg2+(aq) (A) +4 +1 (C) +1, +2, +7 → Cu2+(aq) + 5 +5 (D) +15 (D) +7 (D) -1, -2, +7 Hg(l) Cu2+(aq) (C) +7 (D) Hg(l) 42. Sodium thiosulfate, Na2S2O3, is used as a “fixer” in black and white photography. Identify which atom is oxidized in the reaction of thiosulfate with iodine. 2S2O32-(aq) + I2(aq) (B) → I- S4O62-(aq) 2I-(aq) + (A) I 43. Which one of the following is not a redox reaction? (A) (B) (C) (D) (E) 2H2(g) + O2(g) → 2H2O(l) Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g) H2O(l) + NH3(g) → NH4+(aq) + OH-(aq) 6FeSO4(aq) + K2Cr2O7(aq) + 7H2SO4(aq) → Cr2(SO4)3(aq) + 3Fe2(SO4)3(aq) + K2SO4(aq) Cl2(g) + 2KBr(aq) → Br2(l) + 2KCl(aq) 44. Which one of the following is not a redox reaction? (A) (B) (C) (D) (E) 2H2O2(aq) → 2H2O(l) + O2(g) N2(g) + 3H2(g) → 2NH3(g) BaCl2(aq) + K2CrO4(aq) → BaCrO4(aq) + 2KCl(aq) 2Al(s) + Fe2O3(s) → Al2O3(s) + 2Fe(s) 2H2O(g) → 2H2(g) + O2(g) 45. Which one of the following is a redox reaction? (A) (B) (C) (D) (E) 2Na(g) + Cl2(g) → 2NaCl(s) Ba2+(aq) + SO42-(aq) → BaSO4(s) K2Cr2O7(aq) + 2KOH(aq) → 2K2CrO4(aq) + H2O(l) Na2CO3(s) + 2HCl(aq) → 2NaCl(aq) + CO2(g) + H2O(l) H2O(l) → H+(aq) + OH-(aq) 46. Which one of the following is not a redox reaction? (A) (B) (C) (D) (E) 2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g) H2(g) + Cl2(g) → 2HCl(g) 2H2O2(aq) → 2H2O(l) + O2(g) Fe2O3(s) + 3H2SO4(aq) → Fe2(SO4)3(aq) + 3H2O(l) 2KMnO4(aq) + 10FeSO4(aq) + 8H2SO4(aq) → K2SO4(aq) + 2MnSO4(aq) + 5Fe2(SO4)3(aq) + (C) S (D) + O 7H2O(l) 8H2O(l) We discussed the relative “strength” of metals and getting electricity from redox reactions. 6 Acids and Bases a) b) c) d) e) Arrhenius definition Bronsted-Lowry definition strong and weak neutralization pH = -log [H+] or –log [H3O+] 47. Which of the following is a weak acid? (A) H2SO4 48. Which of the following is a strong acid? (A) H3PO4 49. Which of the following is a strong base? (A) NH3 50. Which of the following is a weak base? (A) NH3 51. Which one of the following substances is a strong acid? (A) (D) HNO3 CH3COOH 52. What is the pH of a 0.75 M HNO3 solution? (A) 0.12 53. What is the pH of a 0.00200 M HClO4 solution? (A) 0.995 54. What is the pH of a 0.050 M HBr solution? (A) 0.89 (B) (B) (B) (B) HNO3 HNO3 (B) (B) (C) Ba(OH)2 Sr(OH)2 (B) (E) (B) (C) 0.29 1.378 1.12 HF (D) HBr HF (D) CH3COOH (C) Al(OH)3 (D) B(OH)3 (C) Ba(OH)2 (D) NaOH H2CO3 H3PO4 (C) (C) NH3 (D) 0.82 2.699 (D) 6.215 1.30 (D) 3.00 0.63 (C) (C) Answers: 1 – 30: 31-54: AEBBC|CDEAE|ACDCD|CABEE|AABAB|BBEEB| CCBDB|DCBBB|ACCCA|DCBBA|AACC 7