Go to Main Index

AP Chemistry by Satellite Lectureguide
Student Edition
Properties of Solutions
Chapter 12

Objectives

Following your study of this chapter, you should be able to

1. define the terms; solution, solute, solvent, dissolution and concentration.
2. distinguish between the physical/chemical changes which occur when sodium chloride is added to water and when sodium metal is added to water.
3. list and give examples of the phase combinations from which a solution can be formed.
4. describe the macroscopic observations when pairs of liquids are mixed together.
5. list a set of rules for predicting whether a molecule is polar or nonpolar.
6. describe the intermolecular attractive forces; dipole-dipole, hydrogen-bonding and dispersion forces.
7. describe the interactions responsible for determining whether a particular solute- solvent pair will form a homogeneous mixture or a heterogeneous mixture.
8. describe how an energy analysis of the attractive interactions lead to exothermic or endothermic enthalpies (DHº_solution) of solution.
   
9. define the terms; solubility, unsaturated, saturated and supersaturated.
10. list the important criteria for predicting whether a mixture will be homogeneous or heterogeneous when mixing a solid and a liquid.
11. define the term lattice energy and explain its importance in the enthalpy of solution.
12. list the important criteria for predicting whether a mixture will be homogeneous or heterogeneous when mixing a gas and a liquid.
13. define the concentration terms; weight percent, mole fraction, molarity, and molality and use these concentration expressions to determine the concentration of any homogeneous mixture.
14. define the term colligative property and list those physical properties of a solution which can be classified as colligative properties.
15. explain how the vapor pressure of a liquid solvent in a solution is affected by the presence of a solute.
16. write the mathematical equation which relates the vapor pressure of the solution to the concentration of the solute and the vapor pressure of the solvent.
17. explain how and why the freezing point and the boiling point of a liquid are affected by the addition of a nonvolatile solute.
18. write and use the mathematical equation that describes the dependence of the freezing point or boiling point on the concentration of solution.
19. define the terms semipermeable membrane, osmosis and osmotic pressure.
20. list some examples and several characteristic properties of a colloid.

1. Define the following terms:

solution

solute

solvent

dissolution

concentration

2. Describe your observations of the following experiments performed by the instructor.

a) Sodium chloride added to water

b) Sodium metal added to water

3. Complete the following table by providing physical examples of solute/solvent combinations.

4. Based on the lecture demonstrations describe what happens when the following pairs of liquids are mixed together.

a) water and hexane

b) hexane and carbon tetrachloride

c) water and ethanol

5. List the set of rules which can be used to predict the polarity of a covalent molecule. (Review: See Exercises 9 and 10 in Chapter 9 of the Lectureguide)

6. Briefly describe each of the following types of intermolecular attractive forces. Sketch the orientations of molecules and/or ions involved in the following intermolecular attractive forces. Include at least one specific example where each attractive force is important. For each one, tell what causes the force and describe its strength relative to the others. (Review: See Exercises 10 and 11 in Chapter 11 of the Lectureguide.)

dipole-dipole forces

London dispersion forces

hydrogen-bonding forces

7. The three attractive interactions which are important in solution formation are; solute-solute interactions, solvent-solvent interactions, and solute-solvent interactions. Define each of these interactions and describe their importance in determining whether a particular solute-solvent pair will form a homogeneous mixture or a heterogeneous mixture.

8a. In terms of the attractive interaction discussed in exercise #6 explain how it is the formation of a solution can be exothermic or endothermic.

b. Describe the underlying thermodynamic property which favors the formation of a solution. Explain why some combinations of chemicals do not form homogeneous mixtures.

9a. Define the following terms;

solubility

unsaturated solution

saturated solution

supersaturated solution

b. Given that the beaker to the left contains an aqueous solution of NaCl, describe a simple test to determine whether the solution is unsaturated, saturated or supersaturated.

What would you expect to happen during the test if the solution were unsaturated? saturated? supersaturated?

10a. Given the representations below, sketch the orientations of a chloride ion and a several water molecules and a sodium ion and several water molecules to illustrate the ion-dipole interaction.

b) Briefly describe ion-dipole intermolecular attractive forces which occur when an ionic solid dissolves in water. Indicate what causes the attractive force and describe how the strength depends on the charge and the size of the ion.

11. Define the term lattice energy and explain its importance in the enthalpy of solution.

12. Explain how pressure, temperature and molar mass effect the solubility of a gas in a liquid.

13a. Define the concentration terms;

weight percent

mole fraction

molarity

molality

b. Calculate the molality and mol fraction of HCl for a solution which is 37.1 % HCl by weight (mass).

Ans: 16.2 m

c. If the density of the solution described in b. is 1.18 g/mL, calculate the molarity of the solution.

Ans: 12.0 M

d. Calculate the weight percent NaCl in a solution which is 0.632 molal.

e. A solution of ethylene glycol, C₂H₄(OH)₂, which is 6.77 molar has a density of 1.05 g/mL. Calculate the mole fraction of ethylene glycol in the solution.

Ans: 0.162

Problem Set #19
AP Chemistry by Satellite ^{Name___________________________________}

ALL work must be shown in all problems for full credit.

PS19.1. Predict whether in the following pairs (solute:solvent) a solution is formed. In each case explain why or why not.
a) HCl(g):H₂O(l)

b) CH₃COOH(l):H₂O(l)

c) CH₃OH(l):C₇H₈ (toluene)(l)

d) C₆H₁₄(l):H₂O(l)

e) I₂(s):C₇H₈ (toluene)(l)

f) Br₂(l):H₂O(l)

g) NaNO₃(s):H₂O(l)

h) NaClO₄(s):CCl₄(l)

PS19.2. Predict whether the following compounds are soluble or insoluble in water.

a) Na₂SO₄(s)

b) C₁₂H₂₂O₁₁(s)

c) CoCl₂(s)

d) CH₃OH(l)

e) C₇H₈(l)

f) HF(l)

g) NH₃(g)

h) (CH₃)₂CO(l)

PS19.3. For those compounds in 19.2 that are soluble write the equation which describes the compound's behavior when added to water.

PS19.4. Describe the attractive forces present when KI(s), NH₃(g) and CH₃CH₂OH(l) dissolve in water. Prepare sketches depicting at the atomic level how each of these substances interact with water molecules.

PS19.5. If 52.6 g of AgNO₃ are added to 684 mL of water, determine;

a) how many grams of AgNO₃ (solute).

b) how many grams of H₂O (solvent). (Density of pure water is 1.00 g/mL )

c) how many grams of solution.

d) how many moles of AgNO₃.

e) how many moles of H₂O.

f) how many moles of solution.

g) how many milliliters of solution, if the density is 1.106 g/mL.

h) the weight percent AgNO₃.

i) the mole fraction of AgNO₃.

j) the molality of the solution.

k) the molarity of the solution.

PS19.6. Given that an aqueous solution, which weighs 367 g, is 40.0 % (ethylene glycol) C₂H₆O₂ by mass and has a density of 1.05 g/mL, determine;

a) the mole fraction of ethylene glycol.

b) the molality of the solution.

c) the molarity of the solution.

PS19.7. An aqueous solution of cesium chloride is 5.94 molal and has a density of 1.58 g/mL. Calculate the
a) weight percent cesium chloride.

b) mole fraction of cesium chloride.

c) molarity of the solution.

PS19.8. Describe how you would prepare;
a) 520 mL of a 0.760 M NaCl solution.

b) 37 g of a 15.2 % (by weight) solution of NaCl.

c) 210. g (grams of solution) of a 0.185 molal NaCl solution.

PS19.9. A 10.00 % (by weight) solution of K₂SO₄ in water was found to be 0.574 M by titration at 25 ºC. Calculate

a) the molality of the solution

b) the density of the solution

14. Define the term colligative property and list those physical properties of a solution which can be classified as colligative properties.

15. Illustrate and explain how the presence of a nonvolatile solute affects the vapor pressure of a liquid.

16a. Draw the vapor pressure versus temperature curve for water and label the important features. On the same diagram draw the vapor-pressure versus temperature curve for an aqueous solution containing a nonvolatile solute.

b. Write Raoult's law and define each term.

Ans: 23.5 mmHg

d. A solution was prepared by adding 20.0 g of urea to 125 g of water at 25 ºC, a temperature at which pure water has a vapor pressure of 23.76 mm of Hg. The observed vapor pressure of the solution was found to be 22.67 mm of Hg. Calculate the molecular weight of urea.

Ans: 59

e. Show the derivation of a mathematical relationship for the vapor pressure lowering (Pº_solvent - P_solution) of a liquid following the addition of a nonvolatile solute.

17. Draw the vapor pressure versus temperature curve for water and label the important features. On the same diagram draw the vapor-pressure versus temperature curve for an aqueous solution containing a nonvolatile solute. Explain how the addition of a nonvolatile solute affects the freezing point and boiling point of water.

18a. Write the general mathematical relation which describes the dependence of the freezing point or boiling point on the concentration of solution.

b. Calculate the freezing point and boiling point of a solution prepared by mixing 6.00 g of C₆H₁₂O₆ with 35.0 g of H₂O.

Ans: T_fp = -1.77 ºC : T_bp = 100.486 ºC

c. A solution containing a nonelectrolyte dissolved in water has a boiling point of 100.305ºC. Calculate the freezing point of the same solution.

Ans: T_fp = -1.11 ºC

d. What is the molecular mass of nicotine if 5.04 grams of this compound changes the freezing point of 90.0 g of water by 0.647 ºC?

Ans: 161 g/mol

e. Calculate the freezing point and the boiling point of a saturated solution of Li₂CO₃. The solubility of lithium carbonate is 0.72 g per 100 g of water at 100 ºC.

Ans: T_fp = -0.544 ºC : T_bp = 100.149 ºC

f. 2.57 g of an ionic compound with the formula KX are dissolved in 120 g of water. The freezing point of the solution was lowered by 1.37 ºC. Determine the formula weight of X.

Ans: 19 g/mol

19a. Define the terms semipermeable membrane, osmosis and osmotic pressure.

b. Using a kinetic molecular model illustrate the movement of solvent molecules across a semipermeable membrane which separates pure water from a solution of sugar and water. Using this illustration explain what happens in reverse osmosis.

20. Give three examples of colloids.

Problem Set #20
AP Chemistry by Satellite ^{Name___________________________________}

ALL work must be shown in all problems for full credit.

PS20.1. Calculate the vapor pressure for each of the following solutions at 25 ºC; a) 16.8 g of urea (NH₂)₂CO dissolved in 108 g of water.

b) 8.54 g of MgCl₂ dissolved in 108 g of water.

PS20.2. To what temperature (ºC) would a solution containing 150 g of glycerol, C₃H₅(OH)₃ (assume to be a nonvolatile solute), in 100 g of water have to be heated to have a vapor pressure of 91.1 mmHg?

PS20.3. Determine the freezing point of a solution which is 0.50 molal urea (a nonelectrolyte). Determine the boiling point of this solution.

PS20.4. A solution containing a nonelectrolyte dissolved in water has a freezing point of -1.62 ºC. Calculate the boiling point of the same solution.

PS20.5. Given the following data;

a) If each of the solutions is prepared by adding 1 mole of compound to 1 kg of water why does each have a different DT_f ?

b) Predict the ideal DT_f for the above compounds.

PS20.5. (Continued)

b) Predict the ideal DT_f for the above compounds.

c) Why does the ideal DT_f differ from the experimental DT_f?

d) Classify each compound as a strong, weak or nonelectrolyte.

PS20.6. Determine the ideal freezing point of a solution prepared by mixing 3.53 g of MgCl₂ in 430 g of water.

PS20.7. Calculate the "molecular formula" of solid sulfur if the freezing point of carbon tetrachloride is lowered by 0.28 ºC when 0.24 g of sulfur is added to 100 g of carbon tetrachloride. (Note: K_f(CCl₄) = 29.8 ºC/m .)

PS20.8. When 7.20 g of HBr are added to 100 g of water, the freezing point of the solution is lowered by 3.30 ºC. How can you explain this temperature change? Is HBr a strong, weak or nonelectrolyte?

PS20.9. 5.76 g of an ionic compound with the formula K₃X is dissolved in 350 g of water. The freezing point of the solution was lowered by 0.370 ºC. Determine the formula weight of X.