Chapter 14 – Review

14.1 – Solutions: An introduction

  1. How do solutions differ from compounds? From other mixtures?
    Check Answer: [1]
  2. Which of the principal characteristics of solutions can we see in the solutions of K2Cr2O7 shown in Figure 14.1a?
  3. When KNO3 is dissolved in water, the resulting solution is significantly colder than the water was originally.
    1. Is the dissolution of KNO3 an endothermic or an exothermic process?
    2. What conclusions can you draw about the intermolecular attractions involved in the process?
    3. Is the resulting solution an ideal solution?
      Check Answer: [2]
  4. Give an example of each of the following types of solutions:
    1. a gas in a liquid
    2. a gas in a gas
    3. a solid in a solid
  5. Indicate the most important types of intermolecular attractions in each of the following solutions:
    1. The solution in Figure 14.1a.
    2. NO(l) in CO(l)
    3. Cl2(g) in Br2(l)
    4. HCl(aq) in benzene C6H6(l)
    5. Methanol CH3OH(l) in H2O(l)
      Check Answer: [3]
  6. Predict whether each of the following substances would be more soluble in water (polar solvent) or in a hydrocarbon such as heptane (C7H16, nonpolar solvent):
    1. vegetable oil (nonpolar)
    2. isopropyl alcohol (polar)
    3. potassium bromide (ionic)
  7. Heat is released when some solutions form; heat is absorbed when other solutions form. Provide a molecular explanation for the difference between these two types of spontaneous processes. Check Answer: [4]
  8. Solutions of hydrogen in palladium may be formed by exposing Pd metal to H2 gas. The concentration of hydrogen in the palladium depends on the pressure of H2 gas applied, but in a more complex fashion than can be described by Henry’s law. Under certain conditions, 0.94 g of hydrogen gas is dissolved in 215 g of palladium metal.
    1. Determine the molarity of this solution (solution density = 1.8 g/cm3).
    2. Determine the molality of this solution (solution density = 1.8 g/cm3).
    3. Determine the percent by mass of hydrogen atoms in this solution (solution density = 1.8 g/cm3).

14.2 – Solubility

  1. Suppose you are presented with a clear solution of sodium thiosulfate, Na2S2O3. How could you determine whether the solution is unsaturated, saturated, or supersaturated?
  2. Supersaturated solutions of most solids in water are prepared by cooling saturated solutions. Supersaturated solutions of most gases in water are prepared by heating saturated solutions. Explain the reasons for the difference in the two procedures.
    Check Answer: [5]
  3. Suggest an explanation for the observations that ethanol, C2H5OH, is completely miscible with water and that ethanethiol, C2H5SH, is soluble only to the extent of 1.5 g per 100 mL of water.
  4. Calculate the percent by mass of KBr in a saturated solution of KBr in water at 10 °C. See Figure 14.2h for useful data, and report the computed percentage to one significant digit. Check Answer: [6]
  5. Which of the following gases is expected to be most soluble in water? Explain your reasoning.
    1. CH4
    2. CCl4
    3. CHCl3
  6. At 0 °C and 1.00 atm, as much as 0.70 g of O2 can dissolve in 1 L of water. At 0 °C and 4.00 atm, how many grams of O2 dissolve in 1 L of water? Check Answer: [7]
  7. Refer to Figure 14.2c.
    1. How did the concentration of dissolved CO2 in the beverage change when the bottle was opened?
    2. What caused this change?
    3. Is the beverage unsaturated, saturated, or supersaturated with CO2?
  8. The Henry’s law constant for CO2 is 3.4 × 10−2M/atm at 25 °C. What pressure of carbon dioxide is needed to maintain a CO2 concentration of 0.10 M in a can of lemon-lime soda? Check Answer: [8]
  9. The Henry’s law constant for O2 is 1.3 × 10−3M/atm at 25 °C. What mass of oxygen would be dissolved in a 40-L aquarium at 25 °C, assuming an atmospheric pressure of 1.00 atm, and that the partial pressure of O2 is 0.21 atm?
  10. How many litres of HCl gas, measured at 30.0 °C and 745 torr, are required to prepare 1.25 L of a 3.20-M solution of hydrochloric acid? Check Answer: [9]

14.3 – Molarity

  1. Explain what changes and what stays the same when 1.00 L of a solution of NaCl is diluted to 1.80 L.
  2. What information do we need to calculate the molarity of a sulfuric acid solution?
    Check Answer: [10]
  3. What does it mean when we say that a 200-mL sample and a 400-mL sample of a solution of salt have the same molarity? In what ways are the two samples identical? In what ways are these two samples different?
  4. Determine the molarity for each of the following solutions:
    1. 0.444 mol of CoCl2 in 0.654 L of solution
    2. 98.0 g of phosphoric acid, H3PO4, in 1.00 L of solution
    3. 0.2074 g of calcium hydroxide, Ca(OH)2, in 40.00 mL of solution
    4. 10.5 kg of Na2SO4·10H2O in 18.60 L of solution
    5. 7.0 × 10−3 mol of I2 in 100.0 mL of solution
    6. 1.8 × 104 mg of HCl in 0.075 L of solution
      Check Answer: [11]
  5. Determine the molarity of each of the following solutions:
    1. 1.457 mol KCl in 1.500 L of solution
    2. 0.515 g of H2SO4 in 1.00 L of solution
    3. 20.54 g of Al(NO3)3 in 1575 mL of solution
    4. 2.76 kg of CuSO4·5H2O in 1.45 L of solution
    5. 0.005653 mol of Br2 in 10.00 mL of solution
    6. 0.000889 g of glycine, C2H5NO2, in 1.05 mL of solution
  6. Consider this question: What is the mass of the solute in 0.500 L of 0.30 M glucose, C6H12O6, used for intravenous injection?
    1. Outline the steps necessary to answer the question.
    2. Answer the question.
      Check Answer: [12]
  7. Consider this question: What is the mass of solute in 200.0 L of a 1.556-M solution of KBr?
    1. Outline the steps necessary to answer the question.
    2. Answer the question.
  8. Calculate the number of moles and the mass of the solute in each of the following solutions:
    1. 2.00 L of 18.5 M H2SO4, concentrated sulfuric acid
    2. 100.0 mL of 3.8 × 10−5M NaCN, the minimum lethal concentration of sodium cyanide in blood serum
    3. 5.50 L of 13.3 M H2CO, the formaldehyde used to “fix” tissue samples
    4. 325 mL of 1.8 × 10−6M FeSO4, the minimum concentration of iron sulfate detectable by taste in drinking water
      Check Answer: [13]
  9. Calculate the number of moles and the mass of the solute in each of the following solutions:
    1. 325 mL of 8.23 × 10−5M KI, a source of iodine in the diet
    2. 75.0 mL of 2.2 × 10−5M H2SO4, a sample of acid rain
    3. 0.2500 L of 0.1135 M K2CrO4, an analytical reagent used in iron assays
    4. 10.5 L of 3.716 M (NH4)2SO4, a liquid fertilizer
  10. Consider this question: What is the molarity of KMnO4 in a solution of 0.0908 g of KMnO4 in 0.500 L of solution?
    1. Outline the steps necessary to answer the question.
    2. Answer the question.
      Check Answer: [14]
  11. Consider this question: What is the molarity of HCl if 35.23 mL of a solution of HCl contain 0.3366 g of HCl?
    1. Outline the steps necessary to answer the question.
    2. Answer the question.
  12. Calculate the molarity of each of the following solutions:
    1. 0.195 g of cholesterol, C27H46O, in 0.100 L of serum, the average concentration of cholesterol in human serum
    2. 4.25 g of NH3 in 0.500 L of solution, the concentration of NH3 in household ammonia
    3. 1.49 kg of isopropyl alcohol, C3H7OH, in 2.50 L of solution, the concentration of isopropyl alcohol in rubbing alcohol
    4. 0.029 g of I2 in 0.100 L of solution, the solubility of I2 in water at 20 °C
      Check Answer: [15]
  13. Calculate the molarity of each of the following solutions:
    1. 293 g HCl in 666 mL of solution, a concentrated HCl solution
    2. 2.026 g FeCl3 in 0.1250 L of a solution used as an unknown in general chemistry laboratories
    3. 0.001 mg Cd2+ in 0.100 L, the maximum permissible concentration of cadmium in drinking water
    4. 0.0079 g C7H5SNO3 in one ounce (29.6 mL), the concentration of saccharin in a diet soft drink.
  14. There is about 1.0 g of calcium, as Ca2+, in 1.0 L of milk. What is the molarity of Ca2+ in milk?
    Check Answer: [16]
  15. What volume of a 1.00-M Fe(NO3)3 solution can be diluted to prepare 1.00 L of a solution with a concentration of 0.250 M?
  16. If 0.1718 L of a 0.3556-M C3H7OH solution is diluted to a concentration of 0.1222 M, what is the volume of the resulting solution?
    Check Answer: [17]
  17. If 4.12 L of a 0.850 M-H3PO4 solution is be diluted to a volume of 10.00 L, what is the concentration of the resulting solution?
  18. What volume of a 0.33-M C12H22O11 solution can be diluted to prepare 25 mL of a solution with a concentration of 0.025 M?
    Check Answer: [18]
  19. What is the concentration of the NaCl solution that results when 0.150 L of a 0.556-M solution is allowed to evaporate until the volume is reduced to 0.105 L?
  20. What is the molarity of the diluted solution when each of the following solutions is diluted to the given final volume?
    1. 1.00 L of a 0.250-M solution of Fe(NO3)3 is diluted to a final volume of 2.00 L
    2. 0.5000 L of a 0.1222-M solution of C3H7OH is diluted to a final volume of 1.250 L
    3. 2.35 L of a 0.350-M solution of H3PO4 is diluted to a final volume of 4.00 L
    4. 22.50 mL of a 0.025-M solution of C12H22O11 is diluted to 100.0 mL
      Check Answer: [19]
  21. What is the final concentration of the solution produced when 225.5 mL of a 0.09988-M solution of Na2CO3 is allowed to evaporate until the solution volume is reduced to 45.00 mL?
  22. A 2.00-L bottle of a solution of concentrated HCl was purchased for the general chemistry laboratory. The solution contained 868.8 g of HCl. What is the molarity of the solution? Check Answer: [20]
  23. An experiment in a general chemistry laboratory calls for a 2.00-M solution of HCl. How many mL of 11.9 M HCl would be required to make 250 mL of 2.00 M HCl?
  24. What volume of a 0.20-M K2SO4 solution contains 57 g of K2SO4? Check Answer: [21]
  25. The US Environmental Protection Agency (EPA) places limits on the quantities of toxic substances that may be discharged into the sewer system. Limits have been established for a variety of substances, including hexavalent chromium, which is limited to 0.50 mg/L. If an industry is discharging hexavalent chromium as potassium dichromate (K2Cr2O7), what is the maximum permissible molarity of that substance?

14.4 – Other Units for Solution Concentrations

  1. Consider this question: What mass of a concentrated solution of nitric acid (68.0% HNO3 by mass) is needed to prepare 400.0 g of a 10.0% solution of HNO3 by mass?
    1. Outline the steps necessary to answer the question.
    2. Answer the question.
      Check Answer: [22]
  2. What mass of a 4.00% NaOH solution by mass contains 15.0 g of NaOH?
  3. What mass of solid NaOH (97.0% NaOH by mass) is required to prepare 1.00 L of a 10.0% solution of NaOH by mass? The density of the 10.0% solution is 1.109 g/mL.
    Check Answer: [23]
  4. What mass of HCl is contained in 45.0 mL of an aqueous HCl solution that has a density of 1.19 g cm–3 and contains 37.21% HCl by mass?
  5. The hardness of water (hardness count) is usually expressed in parts per million (by mass) of CaCO3, which is equivalent to milligrams of CaCO3 per litre of water. What is the molar concentration of Ca2+ ions in a water sample with a hardness count of 175 mg CaCO3/L?
    Check Answer: [24]
  6. The level of mercury in a stream was suspected to be above the minimum considered safe (1 part per billion by weight). An analysis indicated that the concentration was 0.68 parts per billion. Assume a density of 1.0 g/mL and calculate the molarity of mercury in the stream.
  7. In Canada and the United Kingdom, devices that measure blood glucose levels provide a reading in millimoles per litre. If a measurement of 5.3 mM is observed, what is the concentration of glucose (C6H12O6) in mg/dL?
    Check Answer: [25]
  8. A throat spray is 1.40% by mass phenol, C6H5OH, in water. If the solution has a density of 0.9956 g/mL, calculate the molarity of the solution.
  9. Copper(I) iodide (CuI) is often added to table salt as a dietary source of iodine. How many moles of CuI are contained in 1.00 lb (454 g) of table salt containing 0.0100% CuI by mass?
    Check Answer: [26]
  10. A cough syrup contains 5.0% ethyl alcohol, C2H5OH, by mass. If the density of the solution is 0.9928 g/mL, determine the molarity of the alcohol in the cough syrup.
  11. D5W is a solution used as an intravenous fluid. It is a 5.0% by mass solution of dextrose (C6H12O6) in water. If the density of D5W is 1.029 g/mL, calculate the molarity of dextrose in the solution.
    Check Answer: [27]
  12. Find the molarity of a 40.0% by mass aqueous solution of sulfuric acid, H2SO4, for which the density is 1.3057 g/mL.

14.5 – Colligative Properties and Osmosis

  1. Which is/are part of the macroscopic domain of solutions and which is/are part of the microscopic domain: boiling point elevation, Henry’s law, hydrogen bond, ion-dipole attraction, molarity, nonelectrolyte, nonstoichiometric compound, osmosis, solvated ion?
  2. What is the microscopic explanation for the macroscopic behaviour illustrated in Figure 14.2f in Chapter 14.2 Solubility?
    Check Answer: [28]
  3. Sketch a qualitative graph of the pressure versus time for water vapour above a sample of pure water and a sugar solution, as the liquids evaporate to half their original volume.
  4. A solution of potassium nitrate, an electrolyte, and a solution of glycerin (C3H5(OH)3), a nonelectrolyte, both boil at 100.3 °C. What other physical properties of the two solutions are identical? Check Answer: [29]
  5. What are the mole fractions of H3PO4 and water in a solution of 14.5 g of H3PO4 in 125 g of water?
    1. Outline the steps necessary to answer the question.
    2. Answer the question.
  6. What are the mole fractions of HNO3 and water in a concentrated solution of nitric acid (68.0% HNO3 by mass)?
    1. Outline the steps necessary to answer the question.
    2. Answer the question.
      Check Answer: [30]
  7. Calculate the mole fraction of each solute and solvent:
    1. 583 g of H2SO4 in 1.50 kg of water—the acid solution used in an automobile battery
    2. 0.86 g of NaCl in 1.00 × 102 g of water—a solution of sodium chloride for intravenous injection
    3. 46.85 g of codeine, C18H21NO3, in 125.5 g of ethanol, C2H5OH
    4. 25 g of I2 in 125 g of ethanol, C2H5OH
  8. Calculate the mole fraction of each solute and solvent:
    1. 0.710 kg of sodium carbonate (washing soda), Na2CO3, in 10.0 kg of water—a saturated solution at 0 °C
    2. 125 g of NH4NO3 in 275 g of water—a mixture used to make an instant ice pack
    3. 25 g of Cl2 in 125 g of dichloromethane, CH2Cl2
    4. 0.372 g of histamine, C5H9N, in 125 g of chloroform, CHCl3
      Check Answer: [31]
  9. Calculate the mole fractions of methanol, CH3OH; ethanol, C2H5OH; and water in a solution that is 40% methanol, 40% ethanol, and 20% water by mass. (Assume the data are good to two significant figures.)
  10. What is the difference between a 1 M solution and a 1 m solution?
    Check Answer: [32]
  11. What is the molality of phosphoric acid, H3PO4, in a solution of 14.5 g of H3PO4 in 125 g of water?
    1. Outline the steps necessary to answer the question.
    2. Answer the question.
  12. What is the molality of nitric acid in a concentrated solution of nitric acid (68.0% HNO3 by mass)?
    1. Outline the steps necessary to answer the question.
    2. Answer the question.
      Check Answer: [33]
  13. Calculate the molality of each of the following solutions:
    1. 583 g of H2SO4 in 1.50 kg of water—the acid solution used in an automobile battery
    2. 0.86 g of NaCl in 1.00 × 102 g of water—a solution of sodium chloride for intravenous injection
    3. 46.85 g of codeine, C18H21NO3, in 125.5 g of ethanol, C2H5OH
    4. 25 g of I2 in 125 g of ethanol, C2H5OH
  14. Calculate the molality of each of the following solutions:
    1. 0.710 kg of sodium carbonate (washing soda), Na2CO3, in 10.0 kg of water—a saturated solution at 0°C
    2. 125 g of NH4NO3 in 275 g of water—a mixture used to make an instant ice pack
    3. 25 g of Cl2 in 125 g of dichloromethane, CH2Cl2
    4. 0.372 g of histamine, C5H9N, in 125 g of chloroform, CHCl3
      Check Answer: [34]
  15. The concentration of glucose, C6H12O6, in normal spinal fluid is [latex]\frac{75\;\text{mg}}{100\;\text{g}}[/latex]. What is the molality of the solution?
  16. A 13.0% solution of K2CO3 by mass has a density of 1.09 g/cm3. Calculate the molality of the solution.
    Check Answer: [35]
  17. Why does 1 mol of sodium chloride depress the freezing point of 1 kg of water almost twice as much as 1 mol of glycerin?
  18. What is the boiling point of a solution of 115.0 g of sucrose, C12H22O11, in 350.0 g of water?
    1. Outline the steps necessary to answer the question
    2. Answer the question
      Check Answer: [36]
  19. What is the boiling point of a solution of 9.04 g of I2 in 75.5 g of benzene, assuming the I2 is nonvolatile?
    1. Outline the steps necessary to answer the question.
    2. Answer the question.
  20. What is the freezing temperature of a solution of 115.0 g of sucrose, C12H22O11, in 350.0 g of water, which freezes at 0.0 °C when pure?
    1. Outline the steps necessary to answer the question.
    2. Answer the question.
      Check Answer: [37]
  21. What is the freezing point of a solution of 9.04 g of I2 in 75.5 g of benzene?
    1. Outline the steps necessary to answer the following question.
    2. Answer the question.
  22. What is the osmotic pressure of an aqueous solution of 1.64 g of Ca(NO3)2 in water at 25 °C? The volume of the solution is 275 mL.
    1. Outline the steps necessary to answer the question.
    2. Answer the question.
      Check Answer: [38]
  23. What is osmotic pressure of a solution of bovine insulin (molar mass, 5700 g mol−1) at 18 °C if 100.0 mL of the solution contains 0.103 g of the insulin?
    1. Outline the steps necessary to answer the question.
    2. Answer the question.
  24. What is the molar mass of a solution of 5.00 g of a compound in 25.00 g of carbon tetrachloride (bp 76.8 °C; Kb = 5.02 °C/m) that boils at 81.5 °C at 1 atm?
    1. Outline the steps necessary to answer the question.
    2. Solve the problem.
      Check Answer: [39]
  25. A sample of an organic compound (a nonelectrolyte) weighing 1.35 g lowered the freezing point of 10.0 g of benzene by 3.66 °C. Calculate the molar mass of the compound.
  26. A 1.0 m solution of HCl in benzene has a freezing point of 0.4 °C. Is HCl an electrolyte in benzene? Explain.
    Check Answer: [40]
  27. A solution contains 5.00 g of urea, CO(NH2)2, a nonvolatile compound, dissolved in 0.100 kg of water. If the vapour pressure of pure water at 25 °C is 23.7 torr, what is the vapour pressure of the solution?
  28. A 12.0-g sample of a nonelectrolyte is dissolved in 80.0 g of water. The solution freezes at −1.94 °C. Calculate the molar mass of the substance.
    Check Answer: [41]
  29. Arrange the following solutions in order by their decreasing freezing points: 0.1 m Na3PO4, 0.1 m C2H5OH, 0.01 m CO2, 0.15 m NaCl, and 0.2 m CaCl2.
  30. Calculate the boiling point elevation of 0.100 kg of water containing 0.010 mol of NaCl, 0.020 mol of Na2SO4, and 0.030 mol of MgCl2, assuming complete dissociation of these electrolytes.
    Check Answer: [42]
  31. How could you prepare a 3.08 m aqueous solution of glycerin, C3H8O3? What is the freezing point of this solution?
  32. A sample of sulfur weighing 0.210 g was dissolved in 17.8 g of carbon disulfide, CS2 (Kb = 2.43 °C/m). If the boiling point elevation was 0.107 °C, what is the formula of a sulfur molecule in carbon disulfide?
    Check Answer: [43]
  33. In a significant experiment performed many years ago, 5.6977 g of cadmium iodide in 44.69 g of water raised the boiling point 0.181 °C. What does this suggest about the nature of a solution of CdI2?
  34. Lysozyme is an enzyme that cleaves cell walls. A 0.100-L sample of a solution of lysozyme that contains 0.0750 g of the enzyme exhibits an osmotic pressure of 1.32 × 10−3 atm at 25 °C. What is the molar mass of lysozyme? Check Answer: [44]
  35. The osmotic pressure of a solution containing 7.0 g of insulin per litre is 23 torr at 25 °C. What is the molar mass of insulin?
  36. The osmotic pressure of human blood is 7.6 atm at 37 °C. What mass of glucose, C6H12O6, is required to make 1.00 L of aqueous solution for intravenous feeding if the solution must have the same osmotic pressure as blood at body temperature, 37 °C?
    Check Answer: [45]
  37. What is the freezing point of a solution of dibromobenzene, C6H4Br2, in 0.250 kg of benzene, if the solution boils at 83.5 °C?
  38. What is the boiling point of a solution of NaCl in water if the solution freezes at −0.93 °C?
    Check Answer: [46]
  39. The sugar fructose contains 40.0% C, 6.7% H, and 53.3% O by mass. A solution of 11.7 g of fructose in 325 g of ethanol has a boiling point of 78.59 °C. The boiling point of ethanol is 78.35 °C, and Kb for ethanol is 1.20 °C/m. What is the molecular formula of fructose?
  40. The vapour pressure of methanol, CH3OH, is 94 torr at 20 °C. The vapour pressure of ethanol, C2H5OH, is 44 torr at the same temperature.
    1. Calculate the mole fraction of methanol and of ethanol in a solution of 50.0 g of methanol and 50.0 g of ethanol.
    2. Ethanol and methanol form a solution that behaves like an ideal solution. Calculate the vapour pressure of methanol and of ethanol above the solution at 20 °C.
    3. Calculate the mole fraction of methanol and of ethanol in the vapour above the solution.
      Check Answer: [47]
  41. The triple point of air-free water is defined as 273.15 K. Why is it important that the water be free of air?
  42. Meat can be classified as fresh (not frozen) even though it is stored at −1 °C. Why wouldn’t meat freeze at this temperature?
    Check Answer: [48]
  43. An organic compound has a composition of 93.46% C and 6.54% H by mass. A solution of 0.090 g of this compound in 1.10 g of camphor melts at 158.4 °C. The melting point of pure camphor is 178.4 °C. Kf for camphor is 37.7 °C/m. What is the molecular formula of the solute? Show your calculations.
  44. A sample of HgCl2 weighing 9.41 g is dissolved in 32.75 g of ethanol, C2H5OH (Kb = 1.20 °C/m). The boiling point elevation of the solution is 1.27 °C. Is HgCl2 an electrolyte in ethanol? Show your calculations. Check Answer: [49]
  45. A salt is known to be an alkali metal fluoride. A quick approximate determination of freezing point indicates that 4 g of the salt dissolved in 100 g of water produces a solution that freezes at about −1.4 °C. What is the formula of the salt? Show your calculations.

14.6 – Colloids

  1. Identify the dispersed phase and the dispersion medium in each of the following colloidal systems: starch dispersion, smoke, fog, pearl, whipped cream, floating soap, jelly, milk, and ruby.
    Check Answer: [50]
  2. Distinguish between dispersion methods and condensation methods for preparing colloidal systems.
  3. How do colloids differ from solutions with regard to dispersed particle size and homogeneity?
    Check Answer: [51]
  4. Explain the cleansing action of soap.
  5. How can it be demonstrated that colloidal particles are electrically charged?
    Check Answer: [52]

Attribution & References

Except where otherwise noted, this page is adapted by Gregory A. Anderson from “11.1 The Dissolution Process“, “11.3 Solubility“, “6.3 Molarity“, “6.4 Other Units for Solution Concentrations“, “11.4 Colligative Properties” and “11.5 Colloids” In General Chemistry 1 & 2 by Rice University, a derivative of Chemistry (Open Stax) by Paul Flowers, Klaus Theopold, Richard Langley & William R. Robinson and is licensed under CC BY 4.0. ​Access for free at Chemistry (OpenStax)​ . / Extracted and reused review questions and solutions from end of pages.
  1. A solution can vary in composition, while a compound cannot vary in composition. Solutions are homogeneous at the molecular level, while other mixtures are heterogeneous. ↵
  2. (a) The process is endothermic as the solution is consuming heat. (b) Attraction between the K+ and [latex]\text{NO}_3^{\;\;-}[/latex] ions is stronger than between the ions and water molecules (the ion-ion interactions have a lower, more negative energy). Therefore, the dissolution process increases the energy of the molecular interactions, and it consumes the thermal energy of the solution to make up for the difference. (c) No, an ideal solution is formed with no appreciable heat release or consumption. ↵
  3. (a) ion-dipole forces; (b) dipole-dipole forces; (c) dispersion forces; (d) dispersion forces; (e) hydrogen bonding ↵
  4. Heat is released when the total intermolecular forces (IMFs) between the solute and solvent molecules are stronger than the total IMFs in the pure solute and in the pure solvent: Breaking weaker IMFs and forming stronger IMFs releases heat. Heat is absorbed when the total IMFs in the solution are weaker than the total of those in the pure solute and in the pure solvent: Breaking stronger IMFs and forming weaker IMFs absorbs heat. ↵
  5. The solubility of solids usually decreases upon cooling a solution, while the solubility of gases usually decreases upon heating. ↵
  6. 40% ↵
  7. 2.80 g ↵
  8. 2.9 atm ↵
  9. 102 L HCl ↵
  10. We need to know the number of moles of sulfuric acid dissolved in the solution and the volume of the solution. ↵
  11. (a) 0.679 M;  (b) 1.00 M;  (c) 0.06998 M;  (d) 1.75 M;  (e) 0.070 M;  (f) 6.6 M ↵
  12. (a) determine the number of moles of glucose in 0.500 L of solution; determine the molar mass of glucose; determine the mass of glucose from the number of moles and its molar mass; (b) 27 g ↵
  13. (a) 37.0 mol H2SO4; 3.63 × 103 g H2SO4; (b) 3.8 × 10−6 mol NaCN; 1.9 × 10−4 g NaCN; (c) 73.2 mol H2CO; 2.20 kg H2CO; (d) 5.9 × 10−7 mol FeSO4; 8.9 × 10−5 g FeSO4 ↵
  14. (a) Determine the molar mass of KMnO4; determine the number of moles of KMnO4 in the solution; from the number of moles and the volume of solution, determine the molarity; (b) 1.15 × 10−3M ↵
  15. (a) 5.04 × 10−3M; (b) 0.499 M; (c) 9.92 M; (d) 1.1 × 10−3M ↵
  16. 0.025 M ↵
  17. 0.5000 L ↵
  18. 1.9 mL ↵
  19. (a) 0.125 M;  (b) 0.04888 M;  (c) 0.206 M;  (e) 0.0056 M ↵
  20. 11.9 M ↵
  21. 1.6 L ↵
  22. (a) The dilution equation can be used, appropriately modified to accommodate mass-based concentration units: [latex]\%\text{mass}_1 \times \;\text{mass}_1 = \%\text{mass}_2 \times \;\text{mass}_2[/latex]This equation can be rearranged to isolate mass1 and the given quantities substituted into this equation. (b) 58.8 g ↵
  23. 114 g ↵
  24. 1.75 × 10−3M ↵
  25. 95 mg/dL ↵
  26. 2.38 × 10−4 mol ↵
  27. 0.29 mol ↵
  28. The strength of the bonds between like molecules is stronger than the strength between unlike molecules. Therefore, some regions will exist in which the water molecules will exclude oil molecules and other regions will exist in which oil molecules will exclude water molecules, forming a heterogeneous region. ↵
  29. Both form homogeneous solutions; their boiling point elevations are the same, as are their lowering of vapour pressures. Osmotic pressure and the lowering of the freezing point are also the same for both solutions. ↵
  30. (a) Find number of moles of HNO3 and H2O in 100 g of the solution. Find the mole fractions for the components. (b) The mole fraction of HNO3 is 0.378. The mole fraction of H2O is 0.622. ↵
  31. (a) [latex]X_{\text{Na}_2\text{CO}_3} = 0.0119[/latex]; [latex]X_{\text{H}_2\text{O}} = 0.988[/latex]; (b) [latex]X_{\text{NH}_4\text{NO}_3} = 0.9927[/latex]; [latex]X_{\text{H}_2\text{O}} = 0.907[/latex]; (c) [latex]X_{\text{Cl}_2} = 0.192[/latex]; [latex]X_{\text{CH}_2\text{Cl}_2} = 0.808[/latex]; (d) [latex]X_{\text{C}_5\text{H}_9\text{N}} = 0.00426[/latex]; [latex]X_{\text{CHCl}_3} = 0.997[/latex] ↵
  32. In a 1 M solution, the mole is contained in exactly 1 L of solution. In a 1 m solution, the mole is contained in exactly 1 kg of solvent. ↵
  33. (a) Determine the molar mass of HNO3. Determine the number of moles of acid in the solution. From the number of moles and the mass of solvent, determine the molality. (b) 33.7 m ↵
  34. (a) 6.70 × 10−1m; (b) 5.67 m; (c) 2.8 m; (d) 0.0358 m ↵
  35. 1.08 m ↵
  36. (a) Determine the molar mass of sucrose; determine the number of moles of sucrose in the solution; convert the mass of solvent to units of kilograms; from the number of moles and the mass of solvent, determine the molality; determine the difference between the boiling point of water and the boiling point of the solution; determine the new boiling point. (b) 100.5 °C ↵
  37. (a) Determine the molar mass of sucrose; determine the number of moles of sucrose in the solution; convert the mass of solvent to units of kilograms; from the number of moles and the mass of solvent, determine the molality; determine the difference between the freezing temperature of water and the freezing temperature of the solution; determine the new freezing temperature. (b) −1.8 °C ↵
  38. (a) Determine the molar mass of Ca(NO3)2; determine the number of moles of Ca(NO3)2 in the solution; determine the number of moles of ions in the solution; determine the molarity of ions, then the osmotic pressure. (b) 2.67 atm ↵
  39. (a) Determine the molal concentration from the change in boiling point and Kb; determine the moles of solute in the solution from the molal concentration and mass of solvent; determine the molar mass from the number of moles and the mass of solute. (b) 2.1 × 102 g mol−1 ↵
  40. No. Pure benzene freezes at 5.5 °C, and so the observed freezing point of this solution is depressed by ΔTf = 5.5 − 0.4 = 5.1 °C. The value computed, assuming no ionization of HCl, is ΔTf = (1.0 m)(5.14 °C/m) = 5.1 °C. Agreement of these values supports the assumption that HCl is not ionized. ↵
  41. 144 g mol−1 ↵
  42. 0.870 °C ↵
  43. S8 ↵
  44. 1.39 × 104 g mol−1 ↵
  45. 54 g ↵
  46. 100.26 °C ↵
  47. (a) [latex]X_{\text{CH}_3\text{OH}} = 0.590[/latex]; [latex]X_{\text{C}_2\text{H}_5\text{OH}} = 0.410[/latex]; (b) Vapor pressures are: CH3OH: 55 torr; C2H5OH: 18 torr; (c) CH3OH: 0.75; C2H5OH: 0.25 ↵
  48. The ions and compounds present in the water in the beef lower the freezing point of the beef below −1 °C. ↵
  49. [latex]{\Delta}\text{bp} = K_{\text{b}}m = (1.20\;^{\circ}\text{C}/m)(\frac{9.41\;\text{g}\;\times\;\frac{1\;\text{mol HgCl}_2}{271.496\;\text{g}}}{0.03275\;\text{kg}}) = 1.27\;^{\circ}\text{C}[/latex] The observed change equals the theoretical change; therefore, no dissociation occurs. ↵
  50. Colloidal System Dispersed Phase Dispersion Medium
    starch dispersion starch water
    smoke solid particles air
    fog water air
    pearl water calcium carbonate (CaCO3)
    whipped cream air cream
    floating soap air soap
    jelly fruit juice pectin gel
    milk butterfat water
    ruby chromium(III) oxide (Cr2O3) aluminum oxide (Al2O3)
    ↵
  51. Colloidal dispersions consist of particles that are much bigger than the solutes of typical solutions. Colloidal particles are either very large molecules or aggregates of smaller species that usually are big enough to scatter light. Colloids are homogeneous on a macroscopic (visual) scale, while solutions are homogeneous on a microscopic (molecular) scale. ↵
  52. If they are placed in an electrolytic cell, dispersed particles will move toward the electrode that carries a charge opposite to their own charge. At this electrode, the charged particles will be neutralized and will coagulate as a precipitate. ↵

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Enhanced Introductory College Chemistry Copyright © 2023 by Gregory Anderson; Caryn Fahey; Jackie MacDonald; Adrienne Richards; Samantha Sullivan Sauer; J.R. van Haarlem; and David Wegman is licensed under a Creative Commons Attribution 4.0 International License, except where otherwise noted.

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