Chapter 11 Problems

11.1  The Dissolution Process

11.1.  How do solutions differ from compounds? From other mixtures?

 

11.2.  Which of the principal characteristics of solutions are evident in the solutions of K₂Cr₂O₇ shown in Figure 11.2?

 

11.3.  When KNO₃ is dissolved in water, the resulting solution is significantly colder than the water was originally.

(a)  Is the dissolution of KNO₃ an endothermic or an exothermic process?

(b)  What conclusions can you draw about the intermolecular attractions involved in the process?

(c)  Is the resulting solution an ideal solution?

 

11.4.  Give an example of each of the following types of solutions:

(a)  a gas in a liquid

(b)  a gas in a gas

(c)  a solid in a solid

 

11.5.  Indicate the most important types of intermolecular attractions in each of the following solutions:

(a)  The solution in Figure 11.2.

(b)  NO(l) in CO(l)

(c)  Cl₂(g) in Br₂(l)

(d)  HCl(g) in benzene C₆H₆(l)

(e)  Methanol CH₃OH(l) in H₂O(l)

 

11.6.  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):

(a)  vegetable oil (nonpolar)

(b)  isopropyl alcohol (polar)

(c)  potassium bromide (ionic)

 

11.7.  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.

 

11.8.  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 (solution density = 10.8 g cm³).

(a)  Determine the molarity of this solution.

(b)  Determine the molality of this solution.

(c)  Determine the percent by mass of hydrogen atoms in this solution.

 

11.2  Electrolytes

11.9.  Explain why the ions Na⁺ and Cl⁻ are strongly solvated in water but not in hexane, a solvent composed of nonpolar molecules.

 

11.10.  Explain why solutions of HBr in benzene (a nonpolar solvent) are nonconductive, while solutions in water (a polar solvent) are conductive.

 

11.11.  Consider the solutions presented:

(a)  Which of the following sketches best represents the ions in a solution of Fe(NO₃)₃(aq)?

(b)  Write a balanced chemical equation showing the products of the dissolution of Fe(NO₃)₃.

 

11.12.  Compare the processes that occur when methanol (CH₃OH), hydrogen chloride (HCl), and sodium hydroxide (NaOH) dissolve in water. Write equations and prepare sketches showing the form in which each of these compounds is present in its respective solution.

 

11.13.  What is the expected electrical conductivity of the following solutions?

(a)  NaOH(aq)

(b)  HCl(aq)

(c)  C₆H₁₂O₆(aq) (glucose)

(d)  NH₃(aq)

 

11.14.  Why are most solid ionic compounds electrically nonconductive, whereas aqueous solutions of ionic compounds are good conductors? Would you expect a liquid (molten) ionic compound to be electrically conductive or nonconductive? Explain.

 

11.15.  Indicate the most important type of intermolecular attraction responsible for solvation in each of the following solutions:

(a)  the solutions in Figure 11.7

(b)  methanol, CH₃OH, dissolved in ethanol, C₂H₅OH

(c)  methane, CH₄, dissolved in benzene, C₆H₆

(d)  the polar halocarbon CF₂Cl₂ dissolved in the polar halocarbon CF₂ClCFCl₂

(e)  O₂(l) in N₂(l)

 

11.3  Solubility

11.16.  Suppose you are presented with a clear solution of sodium thiosulfate, Na₂S₂O₃. How could you determine whether the solution is unsaturated, saturated, or supersaturated?

 

11.17.  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.

 

11.18.  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.

 

11.19.  Calculate the percent by mass of KBr in a saturated solution of KBr in water at 10°C. See Figure 11.16 for useful data, and report the computed percentage to one significant digit.

 

11.20.  Which of the following gases is expected to be most soluble in water? Explain your reasoning.

(a)  CH₄

(b)  CCl₄

(c)  CHCl₃

 

11.21.  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?

 

11.22.  Refer to Figure 11.10.

(a)  How did the concentration of dissolved CO₂ in the beverage change when the bottle was opened?

(b)  What caused this change?

(c)  Is the beverage unsaturated, saturated, or supersaturated with CO₂?

 

11.23.  The Henry’s law constant for CO₂ is 3.4 × 10⁻² M/atm at 25°C. Assuming ideal solution behavior, what pressure of carbon dioxide is needed to maintain a CO₂ concentration of 0.10 M in a can of lemon-lime soda?

 

11.24.  The Henry’s law constant for O₂ is 1.3 × 10⁻³ M/atm at 25°C. Assuming ideal solution behavior, 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?

 

11.25.  Assuming ideal solution behavior, how many liters 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?

 

11.4  Concentration Units

11.26.  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?

 

11.27.  What is the microscopic explanation for the macroscopic behavior illustrated in Figure 11.14?

 

11.28.  Sketch a qualitative graph of the pressure versus time for water vapor above a sample of pure water and a sugar solution, as the liquids evaporate to half their original volume.

 

11.29.  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?

 

11.30.  What are the mole fractions of H₃PO₄ and water in a solution of 14.5 g of H₃PO₄ in 125 g of water?

(a)  Outline the steps necessary to answer the question.

(b)  Answer the question.

 

11.31.  What are the mole fractions of HNO₃ and water in a concentrated solution of nitric acid (68.0% HNO₃ by mass)?

(a)  Outline the steps necessary to answer the question.

(b)  Answer the question.

 

11.32.  Calculate the mole fraction of each solute and solvent:

(a)  583 g of H₂SO₄ in 1.50 kg of water—the acid solution used in an automobile battery

(b)  0.86 g of NaCl in 1.00 × 10² g of water—a solution of sodium chloride for intravenous injection

(c)  46.85 g of codeine, C₁₈H₂₁NO₃, in 125.5 g of ethanol, C₂H₅OH

(d)  25 g of I₂ in 125 g of ethanol, C₂H₅OH

 

11.33.  Calculate the mole fraction of each solute and solvent:

(a)  0.710 kg of sodium carbonate (washing soda), Na₂CO₃, in 10.0 kg of water—a saturated solution at 0°C

(b)  125 g of NH₄NO₃ in 275 g of water—a mixture used to make an instant ice pack

(c)  25 g of Cl₂ in 125 g of dichloromethane, CH₂Cl₂

(d)  0.372 g of tetrahydropyridine, C₅H₉N, in 125 g of chloroform, CHCl₃

 

11.34.  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.)

 

11.35.  What is the difference between a 1 M solution and a 1 m solution?

 

11.36.  What is the molality of phosphoric acid, H₃PO₄, in a solution of 14.5 g of H₃PO₄ in 125 g of water?

(a)  Outline the steps necessary to answer the question.

(b)  Answer the question.

 

11.37.  What is the molality of nitric acid in a concentrated solution of nitric acid (68.0% HNO₃ by mass)?

(a)  Outline the steps necessary to answer the question.

(b)  Answer the question.

 

11.38.  Calculate the molality of each of the following solutions:

(a)  583 g of H₂SO₄ in 1.50 kg of water—the acid solution used in an automobile battery

(b)  0.86 g of NaCl in 1.00 × 10² g of water—a solution of sodium chloride for intravenous injection

(c)  46.85 g of codeine, C₁₈H₂₁NO₃, in 125.5 g of ethanol, C₂H₅OH

(d)  25 g of I₂ in 125 g of ethanol, C₂H₅OH

 

11.39.  Calculate the molality of each of the following solutions:

(a)  0.710 kg of sodium carbonate (washing soda), Na₂CO₃, in 10.0 kg of water—a saturated solution at 0°C

(b)  125 g of NH₄NO₃ in 275 g of water—a mixture used to make an instant ice pack

(c)  25 g of Cl₂ in 125 g of dichloromethane, CH₂Cl₂

(d)  0.372 g of tetrahydropyridine, C₅H₉N, in 125 g of chloroform, CHCl₃

 

11.40.  The concentration of glucose, C₆H₁₂O₆, in normal spinal fluid is 75 mg/100 g. What is the molality of the solution?

 

11.41.  A 13.0% solution of K₂CO₃ by mass has a density of 1.09 g/cm³. Calculate the molality of the solution.

 

11.42.  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?

 

11.43.  Assuming ideal solution behavior, what is the boiling point of a solution of 115.0 g of nonvolatile sucrose, C₁₂H₂₂O₁₁, in 350.0 g of water?

(a)  Outline the steps necessary to answer the question

(b)  Answer the question

 

11.44.  Assuming ideal solution behavior, 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?

(a)  Outline the steps necessary to answer the question.

(b)  Answer the question.

 

11.45.  Assuming ideal solution behavior, what is the freezing temperature of a solution of 115.0 g of sucrose, C₁₂H₂₂O₁₁, in 350.0 g of water?

(a)  Outline the steps necessary to answer the question.

(b)  Answer the question.

 

11.46.  Assuming ideal solution behavior, what is the freezing point of a solution of 9.04 g of I₂ in 75.5 g of benzene?

(a)  Outline the steps necessary to answer the following question.

(b)  Answer the question.

11.47.  Assuming ideal solution behavior, what is osmotic pressure of a solution of bovine insulin (molar mass, 5700 g/mol) at 18°C if 100.0 mL of the solution contains 0.103 g of the insulin?

(a)  Outline the steps necessary to answer the question.

(b)  Answer the question.

 

11.48.  Assuming ideal solution behavior, 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; K_b = 5.02°C/m) that boils at 81.5°C at 1 atm?

(a)  Outline the steps necessary to answer the question.

(b)  Solve the problem.

 

11.49.  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. Assuming ideal solution behavior, calculate the molar mass of the compound.

 

11.50.  A solution contains 5.00 g of urea, CO(NH₂)₂, a nonvolatile compound, dissolved in 0.100 kg of water. If the vapor pressure of pure water at 25°C is 23.7 torr, what is the vapor pressure of the solution (assuming ideal solution behavior)?

 

11.51.  A 12.0-g sample of a nonelectrolyte is dissolved in 80.0 g of water. The solution freezes at −1.94°C. Assuming ideal solution behavior, calculate the molar mass of the substance.

 

11.52.  How could you prepare a 3.08 m aqueous solution of glycerin, C₃H₈O₃? Assuming ideal solution behavior, what is the freezing point of this solution?

 

11.53.  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 (assuming ideal solution behavior)?

 

11.54.  Assuming ideal solution behavior, 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?

 

11.55.  The vapor pressure of methanol, CH₃OH, is 94 torr at 20°C. The vapor pressure of ethanol, C₂H₅OH, is 44 torr at the same temperature.

(a)  Calculate the mole fraction of methanol and of ethanol in a solution of 50.0 g of methanol and 50.0 g of ethanol.

(b)  Ethanol and methanol form a solution that behaves like an ideal solution. Calculate the vapor pressure of methanol and of ethanol above the solution at 20 °C.

(c)  Calculate the mole fraction of methanol and of ethanol in the vapor above the solution.

 

11.56.  The triple point of air-free water is defined as 273.16 K. Why is it important that the water be free of air?

 

11.57.  Meat can be classified as fresh (not frozen) even though it is stored at −1°C. Why wouldn’t meat freeze at this temperature?

 

11.58.  Assuming ideal solution behavior, 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.

(a)  Outline the steps necessary to answer the question.

(b)  Answer the question.

11.7  Determining Molar Masses Using Colligative Properties

11.63.  A sample of sulfur weighing 0.210 g was dissolved in 17.8 g of carbon disulfide, CS₂ (K_b = 2.34°C/m). If the boiling point elevation was 0.107°C, what is the formula of a sulfur molecule in carbon disulfide (assuming ideal solution behavior)?

 

11.64.  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?

 

11.65.  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. Assuming ideal solution behavior, what is the molar mass of lysozyme?

 

11.66.  The osmotic pressure of a solution containing 7.0 g of insulin per liter is 23 torr at 25°C. Assuming ideal solution behavior, what is the molar mass of insulin?

 

11.67.  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 K_b for ethanol is 1.20°C/m. Assuming ideal solution behavior, what is the molecular formula of fructose?

 

11.68.  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. K_f for camphor is 37.7°C/m. Assuming ideal solution behavior, what is the molecular formula of the solute? Show your calculations.

 

11.69.  A sample of HgCl₂ weighing 9.41 g is dissolved in 32.75 g of ethanol, C₂H₅OH (K_b = 1.20°C/m). The boiling point elevation of the solution is 1.27°C. Is HgCl₂ an electrolyte in ethanol? Show your calculations.

 

11.70.  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. Assuming ideal solution behavior, what is the formula of the salt? Show your calculations.

 

11.8  Colloids

11.71.  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.

 

11.72.  Distinguish between dispersion methods and condensation methods for preparing colloidal systems.

 

11.73.  How do colloids differ from solutions with regard to dispersed particle size and homogeneity?

 

11.74.  Explain the cleansing action of soap.

 

11.75.  How can it be demonstrated that colloidal particles are electrically charged?