PROPERTIES OF SOLUTIONS - UNIT 08

EXTRA PROBLEMS FOR PRACTICE PROPERTIES OF SOLUTIONS - UNIT 08 ______________________________________________________________

Questions 1 through 27 are multiple choice questions. Questions 28-35 are free response questions and problems.

1. Compared to a 1.0M aqueous solution of glucose, a 1.0M aqueous solution of calcium chloride will have . . .

2. At 298K, 25.00mL of solution containing 27.55mg of protein has an osmotic pressure of 3.22 torr. The molecular mass of the protein, in amu, is . . .

3. What is the relative order of freezing points of these three substances? I. water II. 0.1M aqueous ammonia solution III. 0.1M aqueous ammonium chloride

4. A one thousand liter sample of water contains one gram of iron(III) ions. What is the concentration of Fe3+ (aq) in parts per million?

5. Consider a 0.50M solution of each of the following salts. Which will have the lowest freezing point?

6. What is the vapor pressure of a solution containing 0.500 mol of glucose and 5.00mol of water at 29oC.

7. For a solution containing a nonvolatile solute dissolved in a volatile solvent, what is true of the vapor pressure, boiling point, and freezing point of the solution compared to the pure solvent?

8. The freezing point of a 1.0m aqueous solution of a substance is approximately -5.6oC. Which of these is the most likely substance dissolved in the solution? The molal freezing point constant for water is -1.86 oC/m.

9. What is the mole fraction of water in a solution that contains 32g of methanol in 36g of water?

(a) the same freezing and boiling points. (b) a lower freezing point and a lower boiling point.

(c) a lower freezing point and higher boiling point. (d) a higher freezing point & a higher boiling point.

(e) a higher freezing point and a lower boiling point.

(a) 2,340,000. (b) 159,000. (c) 6360. (d) 254.

(a) III < II < I (b) II < III < I (c) II = III < I

(d) II < III = I (e) I < II < III

(a) 0.001 (b) 0.01 (c) 0.1 (d) 1 (e) 10

(a) NaCl (b) MgCl2 (c) K2SO4 (d) Cr(NO3)3 (e) CaSO4

(a) 30.0 torr (b) 27.3 torr (c) 25.0 torr (d) 2.7 torr

(a) VP = increases BP = increases FP = increases (b) VP = decreases BP = decreases FP = decreases

(c) VP = increases BP = decreases FP = increases (d) VP = decreases BP = decreases FP = increases

(e) VP = decreases BP = increases FP = decreases

(a) CH3CH2OH (b) NaCl (c) Ca(NO3)2 (d) C6H6 (e) Al2(SO4)3

(a) 0.33 (b) 0.50 (c) 0.67 (d) 1.0 (e) 1.1

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10. Which pairs of substances will dissolve in each other? I. CH3OH II. C6H6 III. CH3CH3

11. Adding 1mol of NaCl to 1kg of water lowers the vapor pressure of water more than adding 1mol of C6H12O6. Explain. (a) NaCl is highly polar whereas C6H12O6 is weakly polar and this causes the difference. (b) NaCl has a smaller molar mass than that of C6H12O6 and this causes the difference. (c) NaCl is polar whereas C6H12O6 is nonpolar and this causes the difference. (d) The total solute concentration of NaCl is twice that of C6H12O6 because NaCl forms Na+ and Cl− ions.

12. Which substances will not form a gas upon mixing with an aqueous acid?

13. How many milliliters of 0.40M FeBr3 solution would be necessary to precipitate all of the Ag+ from 30.mL of a 0.40M AgNO3 solution according to the following equation? FeBr3 (aq) + 3 AgNO3 (aq) → Fe(NO3)3 (aq) + 3 AgBr (s)

14. The aqueous solutions of which salt will have the strongest cation-dipole interactions?

15. How does the magnitude of ΔHmix compare with the magnitude of ΔHsolute + ΔHsolvent for exothermic solution processes?

Use the structures of cyclohexane, glucose, pentane, and acetone to answer questions 16 - 18.

(a) I and II only (b) II and III only (c) I & II, I & III, II & III

(d) I and III only (e) I & II, II & III only

(a) NaHCO3 (s) (b) Mg (s) (c) Ca (s) (d) K2CO3 (s) (e) Al2O3 (s)

(a) 10. (b) 20. (c) 30. (d) 40. (e) 50.

(a) NaCl (b) MgCl2 (c) K2SO4 (d) Cr(NO3)3 (e) CaSO4

(a) larger (b) almost the same (c) smaller

cyclohexane

C

C

C

C

C

C

H

H

H

H

H

H

H

H

HH

H

H

acetone

glucose

pentane



16. Pentane is most soluble in . . .

17. When glucose dissolves in water, the strongest intermolecular forces in the solution are . . .

18. Which of the following molecules will not dissolve in water?

Questions 19 and 20 refer to the following statement: When solid ammonium chloride dissolves in water, the solution becomes cool.

19. For such a process, which statement best describes the reason why dissolving occurs spontaneously?

20. Which double-sided arrow in the

following diagram represents the change in enthalpy of solution, ∆Hsolution? The length of the double-sided arrows are proportional to enthalpy changes involved when a solid dissolves to form a solution. (a) A (b) B (c) C (d) D

21. A laboratory procedure requires 0.270 mole of methanol, CH3OH to be added as a solvent. What volume of 1.50M CH3OH is needed to attain this number of moles?

22. Which substance is least soluble in water?

(a) glucose only. (b) glucose and acetone. (c) glucose, acetone, & cyclohexane.

(d) cyclohexane only. (e) cyclohexane and glucose.

(a) dispersion. (b) ion-dipole. (c) hydrogen bonds. (d) ionic. (e) covalent.

(a) acetone (b) cyclohexane (c) pentane

(d) acetone and cyclohexane (e) cyclohexane and pentane

(a) Enthalpy decreases. (b) Enthalpy increases. (c) Vapor pressure decreases. (d) Entropy increases.

(a) 5.55 mL (b) 45.0 mL (c) 405 mL (d) 180 mL

(a) methanol

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(d) methyl bromide

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(c) methane

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(b) methyl amine

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23. The reason for the answer to question 22 is . . . (a) methanol forms the strongest hydrogen bonding with the water molecules. (b) methyl amine forms the strongest hydrogen bonding with the water molecules. (c) methane exhibits London dispersion forces with water molecules. (d) methyl bromide has the highest molar mass.

24. How do the slopes of the lines on the solubility graph to the right vary with molecular mass of the gas? Explain the trend. (a) The slopes increase linearly as the molecular weights increase. Increasing molecular

mass leads to greater polarizability of the gas molecules and greater intermolecular attractions between gas and water molecules.

(b) The slopes exponentially increase as the molecular weights increase. Increasing molecular mass leads to greater polarizability of the gas molecules and greater intermolecular attractions between gas and water molecules.

(c) The slopes increase linearly as the molecular weights decrease. Decreasing molecular mass leads to greater polarizability of the gas molecules and greater intermolecular attractions between gas and water molecules.

(d) The slopes exponentially decrease as the molecular weights decrease. Decreasing molecular mass leads to greater polarizability of the gas molecules and greater intermolecular attractions between gas and water molecules.

25. A pressure of 3.5 atm in CO2 is required to maintain a 0.12M CO2 concentration in soda. Calculate the Henry’s law constant for CO2.

26. On a clear day at sea level, the partial pressure of N2 isn air is 0.78 atm at 25oC. Under these conditions, the concentration of N2

in water is 5.3x10−4M. What is the partial pressure of N2 when the concentration in water is 1.1x10−3M?

27. Consider two solutions: (1) 10 g of glucose (C6H12O6) in 1L of water and (2) 10 g of sucrose (C12H22O11) in 1L of water. Which of the following is true? (a) The glucose solution has the higher vapor pressure. (b) The sucrose solution has the higher vapor pressure. (c) Both solutions have the same vapor pressure. (d) There is not enough information to compare vapor pressures.

28. 175g of calcium chloride is dissolved in 975g water. The density of the solution is 1.10 g/mL. For water at 25oC: VP = 23.76 torr FP constant = 1.86oC/m BP constant = 0.51oC/m a. What is the vapor pressure of the solution at 25oC? b. At what temperature will the solution freeze? c. At what temperature will the solution boil? d. What will be the osmotic pressure?

29. An unknown compound contains only C, H, and O. Combustion analysis of the compound gives mass percents of 31.57%C and 5.30%H. A solution made by dissolving 10.56g of the compound in 25.0mL of water freezes at -5.20oC. Determine the empirical formula, molar mass, and molecular formula of the unknown compound. Assume the compound is a nonelectrolyte. The freezing point constant for water is 1.86oC/m.

(a) 29.2 (b) 0.42 (c) 2.9x101 (d) 3.4x10−2 (e) 4.2

(a) 0.63atm (b) 0.78 atm (c) 1.0 atm (d) 2.1 atm (e) 1.6 atm



30. Answer the following questions about these laboratory observations. Solid ammonium chloride dissolves in water with a marked decrease in temperature. Calcium chloride solid dissolves in water with a marked increase in temperature. Little or no temperature change is observed when solid sodium chloride dissolves in water. a. Write an equation that describes the dissolving process of ammonium chloride. b. Is the dissolving of calcium chloride endothermic or exothermic? c. Describe the opposing forces of attraction that are at work in the dissolution of calcium chloride. Which are greater? Why? d. What can be said about the opposing forces of attraction when sodium chloride dissolves in water? Why? e. Use the observation for ammonium chloride to discuss these seemingly contradictory statements:

(i) Thermodynamically, exothermic processes tend to be spontaneous. (ii) Most processes occur spontaneously when there is an increase in entropy.

31. The molecular formula of an unknown compound is determined by combustion analysis and freezing point depression. A solution containing 0.496g benzoic acid (C6H5COOH), and 25.0g camphor (C10H16O), freezes at 173.3oC. The freezing point of pure camphor is 179.8oC. An unknown molecular compound is found to contain 80.77%C, 3.846%H, and 15.38%O. A solution of 0.243g of unknown and 15.1g of camphor melts at 176.7oC. a. What is the empirical formula of the unknown compound? b. What is the freezing point constant for camphor? Specify units. c. What is the concentration of the unknown compound in camphor in units of molality? d. What is the molar mass of the unknown compound? e. What is the empirical formula of the unknown compound?

32. Three unknown metals are labeled A, B, and C. One is silver, one is aluminum, and the other is nickel. Using only aqueous solutions of 0.50M iron(II) nitrate and 1.0M hydrochloric acid, write a short, concise, experimental procedure, the results of which will be sufficient to identify each of the unknown metals. Tell whether you would expect to see and what the observation means. Write net ionic equations to illustrate your answers.

33. Two important vitamins are shown in the figure below. a. Tell whether each vitamin is more likely to be water soluble or fat soluble. Justify your reasoning. b. What is the principle intermolecular force acting when a nonpolar substance dissolves in a nonpolar solvent? c. Describe the principle intermolecular force(s) acting when a water-soluble vitamin dissolves in water? d. Identify one aspect that the space-filling models in the figure represent accurately and one aspect that limits their usefulness. e. Identify one aspect that the structural formulas represents accurately and one aspect that limits their usefulness.

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34. The figure shows the dynamic equilibrium in a saturated solution with excess ionic solute. a. Describe the forces that must be overcome for the salt

to dissolve. b. Are the larger ions depicted most likely cations or

anions? Give two reasons for your answers. c. Describe the principle force that holds the ions in

solution. d. Under what conditions will this system never reach

equilibrium? Explain. e. Comparing a salt solution to pure water with both

systems containing the same amount of water, which system will require more heat to completely evaporate the system to dryness? Explain.

35. When HNO2 is dissolved in water, it partially ionizes according to the equation: HNO2 ⇌ H+ + NO2−. A solution is prepared

that contains 7.050g HNO2 in 1.000kg of water. Its freezing point is −0.2929oC. Calculate the fraction of HNO2 that has ionized.


PROPERTIES OF SOLUTIONS - UNIT 08

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